摘要:

Quarterly Reviews No 4 Vol22 1968 TILDEN LECTURE The Photochemistry of Some Allylic Compounds By R. C. Cookson Chemistry of Tetracyclines By D. L. J. Clive Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments By E. Kent Barefield D. H. Busch and S. M. Nelson Macromolecular Structure and Properties of Ribonucleic Acids By R. A. Cox Isopoly-vanadates -niobates and -tantalates By M. T. Pope and B. W. Dale Decomposition Reactions of Radicals By J. A. Kerr and A. C. Lloyd Page 423 435 457 499 527 549 The Reactions of Hydrated Electrons with Inorganic Compounds By M. Anbar 579 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 1V OBN. Such pre- liminary communications should be accompanied by an outline of the ground to be covered (about two quarto pages) rather than by the completed manuscript. Price to non-fellows &4 10s. Od. per annum The Chemical Society-Endowed Lectures. The Council of The Chemical Society has decided that the Endowed Lectures of the Society shall in future be published in full in Quarterly Reviews to ensure that there will be a permanent record. These will be in addition to the normal Review articles. 0 Copyright reserved by The Chemical Society 1968 Published by The Chemical Society Burlington House London. Printed in England by The Thanet Press Margate

ISSN:0009-2681    

DOI:10.1039/QR96822FP007

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

TILDEN LECTURE The Photochemistry of Some Allylic compounds By R. C. Cookson CHEMISTRY DEPARTMENT THE UNIVERSITY SOUTHAMPTON SO9 5 N H Although the importance of organic photochemistry has always been obvious as the ultimate basis of all life and of the oldest and still the largest industry it is only in the last ten or fifteen years that its interest and potentialities have caught the imagination of organic chemists. Not only are the journals and lecture rooms increasingly full of accounts of old and new photochemical reactions and speculations (or even experiments) on their mechanisms but photochemical processes are beginning to be accepted as another normal laboratory technique rather than the esoteric secret of the specialist that they seemed a few years ag0.l After all it is just as easy to plug in a U.V.lamp as a heating mantle. The present intense activity can probably be traced to several causes. (a) Im- proved instruments and experimental methods have greatly simplified the work (b) the physical chemists and spectroscopists have now established the important fundamental processes (c) the molecular orbital theory especially the brilliant popularisation of orbital symmetry rules by Woodward and Hoffmann has provided a useful qualitative way of thinking about primary photochemical processes (d) the extraordinary structures formed by many photochemical re- actions have emphasised the potential of photosynthesis and (e) even the sceptics have been impressed by the Toyo Rayon Co.’s demonstration of the possibility of large-scale industrial photochemistry2 with its photochemical nitrosation of cyclohexane on a plant to produce 1.1 x 108 lb.a year of nylon-6. To these temptations one can add the further inducement that some traditional fields of organic chemistry have become over-grazed or like the determination of struc- tures of natural products have become obsolete. Some interested in molecular 1 In providing rather belatedly this written version of an unwritten lecture apart from some abbreviations I have kept to the original particularly in not giving references. Anyone interested in learning more about photochemistry can be referred first of all to the com- prehensive and authoritative gospel by J. Calvert and J. N. Pitts ‘Photochemistry’ Wiley New York 1966. The direct use of sunlight for industrial photochemistry on a large scale (apart from agriculture and the smog over Los Angeles) seems as far off as when Giacomo Ciamician one of the pioneers of organic photochemistry made his prediction in 1912 ‘Solar energy is not evenly distributed over the surface of the earth; there are privileged regions and others that are less favoured by the climate.. . . On the arid lands there will spring up industrial colonies without smoke and without smokestacks; forests of glass tubes will extend over the plains and glass buildings will rise everywhere.’ 423 The Photochemistry of Some Allylic Compounds structure and stereochemistry refugees from the X-ray crystallographers have turned to photochemistry. Those interested in reaction mechanisms too have found many thermal reactions increasingly well explored and now enjoy the intricacies of electronically as well as vibrationally excited states there is a limit to the number of alkyl halides one can solvolyse.Indeed one of the attractions of photochemistry is the way it brings together all we know about molecular structure and behaviour. Before considering a few reactions that have interested my colleagues and me recently it may be helpful to remind ourselves of some relevant features of mole- cular energy states as summarised in the familiar Figure 1. The energy level G in this diagram represents the ground state of molecule X and S and S2 the first and second electronically excited states each with its vibrational sub-levels. The arrow (a) represents the excitation of X by absorption of a photon of appropriate energy to the second excited state S2.Usually the equilibrium geometry of the excited state will differ sometimes drastically from that of the ground state and the Frank-Condon principle ensures that S2 is reached at a high vibrational level. In the liquid state at room temperature the excess vibra- tional energy is usually lost so quickly (k > 10l2 sec.-l) that the lowest vibra- tional level is reached before anything else can happen to the molecule. Since fluorescence (6) can be observed from Sl but not S2 conversion of S2 to S must also be extremely fast. Photochemical reactions in solution that proceed through the singlet will amost always therefore involve Sl rather than higher excited states even when light of short wavelength is used which initially produces a higher excited state. A ‘forbidden’ transition (f) with inversion of the spin of an electron can OCCUT from S to the corresponding triplet state T, forbidden emission from which causes phosphorescence (c).The other and usually more important way for T to get back to the ground state G is of course by switching spin without emitting radiation and going over to a high vibrational level of G. In solution exchange of energy between molecules is often very important shown by absorption of light by molecules of one kind and emission or photo- chemical reaction from another. Exchange of energy between singlets may take place over relatively large distances by coupled dipole interaction. Exchange of triplet energy requiring close contact of the two molecules with overlap of their outer electron orbitals is particularly prevalent in organic photochemistry.In Figure 1 the process is shown by the dotted lines (e) in which the triplet of molecule Y undergoes a transition to the ground state coupled with the simul- taneous transition of X with which it is in contact from the ground state to Tl (with preservation of multiplicity of the complex). Y sensitises the phosphores- cence or photochemical reaction of X through T, while X conversely quenches the phosphorescence or reaction of Y. For efficient exchange the triplet excitation energy of Y naturally has to be greater than that of X (e > c). By irradiation of a solution containing a photosensitiser Y with light of wavelength absorbed by Y (d) but not by the reactant X [which would need light quanta of energy (b)] X can be put in the triplet state T without going through any excited singlet state.Because of the small singlet-triplet splitting of the nn* excited states and 424 Cookson 7; - sl G X Y Figure 1 Partial energy diagram for two molecules X and Y (a) Excitation of X to second excited singlet state (b) Fluorescence from first excited state (c) Phosphorescence from lowest triplet state ( d ) Excitation of Y to first excited singlet state (e) Coupled transition of X from G to TI and Y from TI to G. the high efficiency of crossing of the excited singlet to the triplet ketones are very popular triplet photosensitisers. G. S. Hammond has provided some nice examples involving dienes. Whereas thermal dimerisation of cyclopentadiene yields the endo-Diels-Alder dimer (I) irradiation of the diene in the presence of benzophenone (ET 69 Kcal.mole-l) with light absorbed only by the latter produces a mixture of dimers consisting mainly of the exo-isomer (11) and the cyclobutane-dimer (111). The nn* triplet of benzophenone formed very efficiently from the singlet produces the triplet diene which adds to a molecule in the ground state to form the dl and mesu allylic diradicals (IV). Spin-inversion and bond-formation lead to the structural and stereochemical isomers. The behaviour of butadiene is more complicated because it can exist in s-cis and s-trans conformations (V and VI) the latter with a higher triplet energy than the former. The two planar conformations are in rapid equilibrium in the Dimerisation of cyclopentadiene by heat and light 425 The Photochemistry of Some Allylic Compounds v) W c W .- -a 80 c 0 0 0 ZI U - 40 50 60 70 Excitation energy of sensitizer Figure 2 Variation of composition of mixture of photodimers of isoprene with tripiet energy of photosensitiser 96 ?o ? 70 54 kcul.mole" 1 is I A (VI I) Is 4 ground state the s-trans predominating but the triplets with a higher 2,3 bond order crudely depicted in (V T) and (VI T) are not. The composition of the photo-products therefore depends on the triplet energy of the sensitiser. Semi- tisers with triplet energies greater than that of s-trans-butadiene (V) produce both cis (VI T) and trans (V T) triplets but mainly the latter since it comes from the predominant ground-state conformation. Sensitisers with triplet energies below that of planar s-trans-butadiene (about 60 Kcal. mole-l) and above that of the s-cis (about 54 Kcal.mole-l) produce only cis-triplets (VI T). The assump- tion that addition of cis-triplets to ground state s-trans-butadiene gives mainly 426 Cookson vinylcyclohexane (VIII) whereas tramtriplets give mainly cis- and trans- divinylcyclobutane (VII) and cyclo-octa-l,5-diene then explains the variation of product composition with sensitiser energy shown in Figure 2 (the points in Figure 2 actually refer to the very similar case of 2-methylbutadiene). Sensitisers with triplet energies above 60 Kcal. mole-l all give the same product composition consisting very largely of cyclobutanes and cyclo-octadiene the proportion of which falls sharply below 60 until at about 52 Kcal. mole-l it levels out again to become independent of the precise energy. But what is happening at very low sensitiser energies where the curve rises steeply again? Hammond's answer is to point out that this type of energy transfer is not a spectroscopic transition and that the Franck-Condon principle need not apply the exchange of electronic energy in the transient complex has long enough to allow simultaneous vibra- tional and rotational changes.The low-energy triplet of the diene may for example have one methylene group orthogonal to the remainder of the molecule consisting essentially of an alkyl radical joined to an ally1 radical. We now turn to the main theme of this lecture the thermal and photochemical rearrangements of diallyl and its derivatives. The behaviour of these compounds is particularly instructive because the three different accessible electronic states isomerise in three separate and unique ways emphasising that the various electronic states of a molecule have quite distinct chemical properties.Figure 3 shows an energy-level scheme for the six electrons associated with the two double bonds and the central C-C single bond of diallyl (IX). On the left are the energies of the uncoupled a and n orbitals in the planar molecule where the a orbital is in the nodal plane of the two n orbitals and in the middle those in the conformation where the 3,4-bond is in a plane at right-angles to the plane of both double bonds and then mdn* levels are therefore split by coupling through the a bond. The resulting molecular orbitals shown in the right of the Figure are simply arranged in increasing numbers of nodes from none to five (apart from the n nodes in the plane at right-angles).Whether the relative energies of the pairs of n and T* levels labelled as Symmetric (S) or Antisym- metric (A) about the dyad axis or mirror plane through the centre of the 3,4-bond are correctly given by this zero-order approximation or should be reversed does not affect the expectation that there should be a spectroscopic transition (of uncertain intensity) at longer wavelength in the coupled system than in the uncoupled one or in a normal olefin. We have now collected a large number of examples of molecules with three C-C single bonds separating two double bonds or two atoms with unshared electron pairs (0 S N) in various com- binations. Only those with the necessary geometry show a new absorption band in the region 220-260 nm. The thermal Cope rearrangement (Xa -+ Xb) of derivatives of diallyl is consistent with the overlap of the orbitals of C-1 and C-6 as the transition state is approached which in the highest occupied level ns are of like phase (XI).As Woodward and Hoffmann and others have even more convincingly shown correlation by their symmetry of all the affected levels of the reactant and pro- duct through the transition state requires no use of high-energy anti-bonding 427 The Photochemistry of Some Allylic Compounds a- Figure 3 Correlation of niolecular orbitals of plattar and non-planar diallyl 428 Cookson orbitals consistent with the concerted intramolecular thermal reaction of relatively low activation energy involving reversal of linkage in both allyl groups. The photochemical rearrangement of several unsymmetrical substituted derivatives of diallyl with one double bond conjugated (to bring its U.V.absorp- tion into a more easily accessible part of the spectrum) was investigated with great skill and efficiency by Dr. N. A. Mirza who returned to the Pakistan Council of Scientific and Industrial Research in 1966. cis- and trans- Geranonitrile (XII) for example rearranges with migration of the dimethylallyl group to give (XIII) or of the cyanomethylallyl group to give and (XV) but not with migration of both as in the Cope rearrangement which would have given (XVI). To check whether this was really a necessary feature of the photo- chemical path a simple allyl system was needed free from the possibility of steric influences. The rearrangement of the allyl group in (XVII) to (XVIII) seemed a suitable case and in a remarkably short time in 1966 Dr.R. F. C. Brown on sabbatical leave from Canberra devised and carried out the synthesis of the specifically deuteriated compounds (XVIIa) and (XVIIIa) as summarised in the flow sheet. Irradiation of (xVIIa) gave the product (XVIII) with the deuterium exclusively still at the 3-position of the allyl group although it had lost its geometrical purity during rearrangement. PhJ CN NC CN &% d/) (xvii) (xv I I I) In striking contrast to this is the photochemical behaviour of the corresponding carbonyl compounds (XIX) such as citral (XIX R = H) which give not pro- ducts of allylic shift but instead form the isomers gcxrr) and (XXIII). Here it is tempting to suppose that crossing of the nn* singlet (or of the ~TT* via the latter) to the nn* triplet crudely represented by (XX) is faster than the allylic shift that occurs in the nitriles.( X X ) then adds to the isolated double bond producing the diradical (XXI) with the two substituents trans. After spin 429 The Photochemistry of Some Allylic Compounds CN D cis and trans (XVI la) (XIXI (xx) R= H M e ,OH ,OMe ,O‘ nn* r / 1 (xxI) Jlj* s’ (xx I I) (XXIII) (XXlV) (xxv) (X xv 0 inversion the reaction can be completed either by formation of a bond between the two tertiary radical centres to give the bicyclohexane or by hydrogen transfer to give the isopropenyl compound (XXIII). The two predominant products were proved by Dr. Mirza to have the configurations shown. 430 Cookson [The process cannot in fact be as simple as implied above because of the variation with sensitiser of the proportions of products from geranic acid found by Dr.Mirza which requires the intervention of at least two triplet states. These products include the 19y-unsaturated isomers (XXIV) shown to be formed only from the starting isomer with the carbonyl cis to the methyl group and (XXV) only from the one with the carbonyl trans to the methyl group presumably arising from the enols produced by the familiar hydrogen shift.] If the carbonyl compounds (XIX) which do not undergo the allylic shift really do react through the triplet state then generation of the triplet of the nitrile (XII) in a way that avoids the excited singlet state should divert the reaction from the allylic shift. Sure enough irradiation of (XII) in acetone as a triplet sensitiser with light not absorbed directly by (XII) gave none of the products of direct irradiation only the bicyclohexanes (XXVII) and (XXVIII) mainly the former.05 8 I ‘/o (XXVIII) In summary we can now show the separate forms of isomerisation exhibited by the three different electronic states illustrated for the deuteriated dinitrile (XVIIa) each characteristic product as far as one can tell being quite free from contamination by either of the others CN NC CN Ground MN mD CN NC CN CN ’ NC CN Excited triplet @ - :Is. ,‘J$ Reactions of different electronic states 43 1 The Photochemistry of Some Allylic Compounds There remain some important stereochemical aspects of the 1,3-allyl shift which can be clarified by following Woodward and Hoffmann's predictions based on the symmetry of thew orbitals of the allyl radical shown in Figure 4.As illustrated in Figure 5 since the orbital of highest energy in the excited state (#,) is symmetric about the mirror plane normal to the allyl group a saturated carbon atom can undergo a 1,3-shift while retaining bonding by simply present- ing the same face to the allyl group throughout the movement. The same process in the ground state however would have to proceed through a transition-state with the symmetric alkyl carbon orbital overlapping the antisymmetric allyl orbital ($&. Bonding can be maintained only by a clockwise rotation of the alkyl carbon atom through 90" in going to the transition-state continuing through another 90" to reach the product. In short a concerted 1,3-allyl shift should result in inversion of configuration of the migrating alkyl group in the thermal rearrangement but retention in the photochemical rearrangement.Figure 4 Molecular orbitals of the allyl radical Berson and Willcott have just provided a well chosen example of the thermal rearrangement the CHD group in (XXIX) moves from one end to the other of the cyclopentenyl group at about 500" the isomer (XXIX) with deuterium C I ,& 7 c /F - c' ,c - = c \. / \=/ C C" Figure 5 1,3-AIlylic migration of an alkyl group 432 Cookson trans to the acetoxy-group giving the product (XXXI) with deuterium and acetoxy-group cis. The adoption of the transition-state (XXX) that forces the C-H bond against the allyl system rather than of the more obvious and less compressed alternative involving the same side of the CHD group throughout tells eloquently of the value of orbital symmetry in understanding concerted reactions.(XXIX) (XXX) (xxxr) Such thermal 1,3-allylic shifts may well occur quite commonly in diallyl derivatives but have escaped recognition because the normal Cope rearrangement is usually much faster. Some possible examples have been investigated by Dr. J. E. Kemp in the geranyl system. For example at about 530" for 5 seconds trans- or cis-geranonitrile (XI9 and the isomer (XIII) give the same mixture of products consisting largely of ( X I I ) with a small amount of (XIII). The corresponding thermal conversion of trans- and cis-geranic ester (XXXII) into a mixture containing a few percent of (XXXIII) [which again reverts to (XXXII) when heated] misled earlier workers into claiming the conversion of the isopropylidene -11) into the isopropenyl compound (XXXIV).We hope to test whether the thermal reaction does involve inversion and the photo- chemical one retention of configuration by use of a system such as (XXXV).* An example of other systems under investigation is the dinitrile (XXXVII) which Mrs. M. Sharma finds rearranges on irradiation to the unconjugated isomer (XXXVIII). Unfortunately it has not been possible to reverse this rearrangement by heating the photo-isomer (XXXVIII) to see whether the trans-isomer of (XXXVII) would result. Such structures as (XXXVII) may nevertheless prove useful in testing the stereochemistry of the rearrangement at one end of the allyl group after stereospecific deuteriation of the appropriate methylene group in (XXXVII) the two olefinic methylene protons in the photo- isomer (XXXVIII) are easily distinguished in the n.m.r.spectrum by the strong shielding of the one cis to the nitrile groups. * In some structures where the two allyl groups are highly conjugated such as in derivatives of (XXXIX) and (XL) rearrangement is quite unstereospecific the same mixture of isomers epimeric at the benzylic centre being formed from either epimer whether by heat or light. These unconcerted rearrangements presumably occur through the diradical (XLI) which has time to undergo rotation about the bonds in the side-chain before recyclising 433 The Photochemistry of Some Allylic Compounds (XXXIII) (XXXIV) A Ph (XXXV) (XXXVI) NC CN Perhaps I shall be able to tell you the results of these and other experiments now under way on some future occasion. Meanwhile I would like to thank my past and present collaborators some of whom have already been introduced by name and to add that most of the work at Southampton is part of a joint project with my old friend and colleague John Hudec. 434

ISSN:0009-2681    

DOI:10.1039/QR9682200423

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

Chemistry of Tetracyclines By D. L. J. Clive 29 BROOK R I S E CHIOWELL ESSEX ENGLAND 1 Introduction The tetracyclines are a group of compounds which has afforded the medical profession a number of powerful antibiotics active against a wide of human and animal pathogens. The compounds are closely related to the parent substance tetracycline(I) and they are obtained as metabolic products of various natural forms of Streptomyces or produced by chemical modification of the metabolites. Nomenclature is based on the numbering shown in (A). Because of their complex structure and great therapeutic value [e.g. the number of prescriptions for antibiotics dispensed in England during 1966 is estimated3 to be (millions) penicillins 14.6; tetracyclines 12.6; others 3.3.1 the tetracyclines have been the object of numerous studies.* Their mode of action biogenesis,6-1° and chemical properties have been examined extensively and the prospect of synthesising a natural member of the group has received a great deal of attenti~n.~ This Review sumarises the results of structural work and examines the published reactions.Such a study will show the circumstances (a) L. M. Pruess and C. H. Demos in ‘Encyclopedia of Chemical Technology’ ed. R. E. Kirk and D. F. Othmer Interscience New York 1954 vol. 13 p. 785; (b) P. P. Regna ibid. pp. 800 808. a H. K. Spitzy Antibiotica et Chemotherapia 1962 10 193. *Annual Report of The Ministry of Health for the Year 1966 Cmnd. 3326 H.M.S.O. London 1967 p. 95. H. Muxfeldt and R. Bangert Fortschr. Chem.-org. Naturstoffe 1963 21 80. G. C. Barrett J. Pharm.Sci. 1963 52 309. J. R. D. McCormick in ‘Biogenesis of Antibiotic Substances’ ed. Z. Vanek and Z. Hostalek Academic Press New York 1965 p. 73. 7 J. H. Martin L. A. Mitscher P. A. Miller P. Shu and N. Bohonos Antimicrobial Agents and Chemotherapy Amer. SOC. Microbiol. Ann Arbor Michigan 1966 p. 563. 13 P A. Miller A Saturnelli J. H. Martin L. A. Mitscher and N. Bohonos Biochem. Biophys. Res. Comm. 1964 16 285. 9 J. R. D. McCormick U. H. Joachim E. R. Jensen S. Johnson and N. 0 Sjolander J Amer. Chem. SOC. 1965 87 1793. lop. A. Miller J. H. Hash M. Lincks and N. Bohonos Biochem. Biophys. Res. Comm. 1965 18 325. 435 Chemistry of Tetracyclines in which the tetracyclines are unstable and will illustrate the extent to which their sensitivity can be reduced by temporary modification.There will also be an opportunity to screen the reactions for possible use in total or partial syn- thetic work. 2 Structural Results The fundamental discoveries in the tetracycline field were made over a period of six years.ll The first member of the group 7-chlorotetracycline (2) was NMe OH OH CONH CONH (2) Aureomycin (3) Terramycin isolated12J3 in 1947 and being effective against a wider range of pathogenic micro-organisms than any agent in use up to that time it soon became important in medical practice in which it was known by the trade name aureomycin. A few years later 5-oxytetracycline or terramycin (3) was discovered14J5 and it too acquired great importance. Finally in 1953 hydrogenolysis of aureomycin was reported to give16J7 tetracycline (1). This compound obtained subsequently also by fermentation,18 represents the molecular structure common to the other antibiotics and formally it is the parent for a number of compounds isolated later.(Not all the tetracyclines nor indeed all their laboratory derivatives are useful as antibiotics; tetracycline therapy involves mainly the first three members and the more recent discovery 7-chloro-6-demethyltetracycline.) These include the 7-brom0-,~~,~~ 6-dimethyl-,21*22 and 7-chlor0-6-demethyl~~,~~ derivatives a l1 J. H. Boothe Antimicrobial Agents and Chemotherapy Amer. SOC. Microbiol Ann Arbor Michigan 1962 p. 213. l a B. M. Duggar Ann. New York Acad. Sci. 1948,51 177. la R. W. Broschard A. C. Dornbush S. Gordon B. L. Hutchings A. R. Kohler G. Krupka S. Kushner D. V. Lefemine and C. Pidacks Science 1949,109 199.l*A. C. Finlay G. L. Hobby S. Y. P’an P. P. Regna J. B. Routien D. B. Seeley G. M. Shull B. A. Sobin I. A. Solomons J. W. Vinson and J. H. Kane Science 1950,111,85. lri P. P. Regna I. A. Solomons K. Murai A. E. Timreck K. J. Brunings and W. A. Lazier J. Amer. Chem. SOC. 1951 73 4211. l6 L. H. Conover W. T. Moreland A. R. English C. R. Stephens and F. J. Pilgrim J. Amer. Chem. SOC. 1953,75,4622. l7 J. H. Boothe J. Morton jun. J. P. Petisi R. G. Wilkinson and J. H. Williams J. Amer. Chem. SOC. 1953,75,4621. l* P. P. Minieri M. C. Firman A. G. Mistretta A. Abbey C. E. Brickler N. E. Rigler and H. Sokol ‘Antibiotics Annual 1953-1954‘ Medical Encyclopedia Inc. New York 1953 p. 81. IsA. P. Doerschuk J. R. D. McCormick J. J. Goodman S. A. Szumski J. A. Growich P. A. Miller B. A. Bitler E.R. Jensen M. Matrishin M. A. Petty and A. S. Phelps J. Amer. Chem. Soc. 1959,81 3069. 80 P. Sensi G. A. De Ferrari G. G. Gallo and G. Rolland I1 Farmaco Edizione Scientijka 1955 10 337. 21 J. R. D. McCormick N. 0. Sjolander U. Hirsch E. R. Jensen and A. P. Doerschuk J. Amer. Chem. Soc. 1957 79 4561. g2 J. S. Webb R. W. Broschard D. B. Cosulich W. J. Stein and C. F. Wolf J. Amer. Chem. SOC. 1957 79 4563. 436 Clive number of 2-acetyl-2-decarboxyamidotetracyclines~3~24 and several dehydro- compounds26 of which 7-chlorodehydrotetracycline (4) has been described26-28 in most detail. The position of the extra double bond in the last substance has not been settled; probably the different tauto~ners~~ have the double bond at positions 5,5a or 5aJla. The structures of terramy~in~~ and a u r e ~ m y c i n ~ ~ ~ ~ were established by chemical experiments.Then with the reactions and spectral attributes of the tetracyclines extensively defined assignment of structures to other members of the group was greatly simplified and followed closely on their isolation. The structure of tetracycline was apparent from its simple relationship to aureomycin (from which it can be obtained by hydrogenoly~is).~~~~~ An X-ray analysis of aureomycin hydrochloride confirmed32 structure (2) and established the relative configuration at each of the asymmetric centres. The results are expressed by the formula which also represents the absolute ~onfiguration.~~ Tetracycline is comparably defined33 by formula (1) because the only molecular change involved in preparing the compound from aureomycin is replacement of chlorine by hydrogen.Similarly the results for aureomycin applP3 to 7-bromotetracycline. In the case of terramycin the gross structure was again confirmed by X-ray studies but these did not define the relative stereochemistry at C(5). The formu- lation given (3) in which the hydroxyl group at C(5) is trans to that at C(6) is supported by n.m.r. and by a reexamination of the early X-ray M. W. Miller and F. A. Hochstein J. Org. Chem. 1962,27,2525. F. A. Hochstein M. Schach von Wittenau F. W. Tanner jun. and K. Murai J. Amer. Chem. SOC. 1960,82,5934. B.P. 863,419; U.S.P. 3,226,435. 96 J. R. D. McCormick P. A. Miller J. A. Growich N. 0. Sjolander and A. P. Doerschuk J. Amer. Chem. SOC. 1958 80 5572. l7 M. Schach von Wittenau F. A. Hochstein and C. R.Stephens J. Org. Chem. 1963 28 2454. 88 A. I. Scott and C. T. Bedford J. Amer. Chem. SOC. 1962,84,2271. ae F. A. Hochstein C. R. Stephens L. H. Conover P. P. Regna R. Pasternack P. N. Gordon F. J. Pilgrim K. J. Brunings and R. B. Woodward J. Amer. Chem. Sac. 1953,75,5455. 8o C. R. Stephens L. H. Conover R. Pasternack F. A. Hochstein W. T. Moreland P. P. Regna F. J. Pilgrim K. J. Brunings and R. B. Woodward J. Amer. Chem. SOC. 1954 76 3568. a1 C. W. Waller B. L. Hutchings R. W. Broschard A. A. Goldman W. J. Stein C. F. Wolf and J. H. Williams J. Amer. Chem. SOC. 1952 74,4981. 3a J. Donohue J. D. Dunitz K. N. Trueblood and M. S . Webster J. Amer. Chem. SOC. 1963 85 851. asV. N. Dobrynin A. I. Gurevitch M. G. Karapetyan M. N. Kolosov and M . M. Shemyakin Tetrahedron Letters 1962 901. M.Schach von Wittenau R. K. Blackwood L. H. Conover R. H. Glauert and R. B. Woodward J. Arner. Chem. SOC. 1965 87 134. 36 Cf. M. Schach von Wittenau and R. K. Blackwood J. Org. Chem. 1966,31 613. 437 Chemistry of Tetracyclines data.36 Probably all the fermentation tetracyclines have the same relative stereochemistry at asymmetric centres that are common and non-epimerizable. The general structure exemplified by formulae (l) (2) and (3) shows a number of features in a simplified way. A n a l ~ s i s ~ ~ ~ of the pK values of the three antibiotics indicates that the amphoteric forms are zwitterionic essentially as shown (S) with both charged sites on ring A. Secondly the diketo-amide ~ i ~ e loo- CONH OH ( 5 ) Ho O- 0 moiety is involved in unusual keto-enol t a u t o m e r i ~ m ~ ~ ~ ~ and it is difficult to define its status.The dimension^^^*^ of ring A (see the Figure) in the hydrochlorides of aureomycin and terramycin indicate that the double bond is delocalised and suggest that the acidic hydrogen atom is attached to the oxygen atom of the amide group at least in the crystal. Other evidence for extensive electron de- L- 2 . 6 7 2 Figure Bond distances (A) in crystalline aureomycin hydrochloride localisation is provided by the high acidity of the enolic hydrogen atom and work with model c ~ m p o u n d s ~ ~ ~ ~ ~ has illustrated further the strong hydrogen bonding and highly enolic character of the system. The p-diketone at C(ll)-C(12) would be expected to be more stable in an enolised form41,42 and X-ray rneasurernent~~~,~~ show a localised double bond at C(l1a)-C(12).The importance of this structure in solution with respect to the alternative [double bond at C(ll)-C(lla)] has not been reported. 36 H. Cid-Dresdner Z. Krist. 1965 121 170. 37 N. E. Rigler S. P. Bag D. E. Leyden J. L. Sudmeier and C. N. Reilly Analyt. Chem. 1965 37 872. ’*C. R. Stephens K. Murai K. J. Brunings and R. B. Woodward J. Amer. Chem. SOC. 1956,78,4155. a0 G. 0. Dudek and G. P. Volpp J. Org. Chem. 1965,30,50. 40 H. Muxfeldt G. Grethe and W. Rogalski J. Org. Chem. 1966 31 2429. 41 C’ M. Gorodetsky Z . Luz and Y. Mazur J. Amer. Chem. Soc. 1967 89 11 83. C’ H. Stetter and U. Milbers Chem. Ber. 1958 91,977. 438 Clive 3 Reactions The C(6) hydroxyl group of the fermentation tetracyclines together with other details of the oxygenation pattern render these materials sensitive to acids and bases [cf.the 6-deoxytetracyclines (p. 455)]. Consequently reaction conditions that may be used during synthesis are limited and the restrictions will become increasingly severe as the natural product is approached. However the synthetic problem is simplified a little by the following observations (which are developed later). The hydroxyl group at the 1%-position can be replaced by a hydrogen atom and it is possible to re-hydroxylate the site in the correct stereochemical sense. It is also known that the asymmetric centre at C(4) is readily epimerised. [Tetracyclines with the unnatural configuration at C(4) have negligible bio- a~tivity.4~1 The change is reversible both epimers cafl be isolated and no other modification takes place. In principle therefore synthesis of tetracycline could involve the preparationu of compounds of type (6).Me OH YMe2 A. Reactions in Acidic Media.-(Cf. the 6-deoxytetracyclines p. 455.) The hydroxyl group at C(6) in the natural tetracyclines is secondary or tertiary. The position is benzylic and the group bears a trans relation to the adjacent hydrogen atom. Moreover there is available45 a transition state in which ring c assumes aromatic character [see (7)]. These circumstances account for the acid lability of the hydroxyl group and make dehydration to anhydro-tetracyclines [e.g. anhydroaureomycin (8)] an easy process. Warm mineral acids in aqueous or 43 J. R. D. McCormick S. M. Fox L. L. Smith B. A. Bitler J. Reichenthal V. E. Origoni W. H. Muller R. Winterbottom and A. P. Doerschuk J. Amer.Chem. Soc. 1957,79 2849. U Cf. R. B. Woodward Pure Appl. Chem. 1963,6,561. 46 R. K. Blackwood J. J. Beereboom H. H. Rennhard M. Schach von Wittenau and C. R. Stephens J. Amer. Chem. Soc. 1963 85 3943. 439 Chemistry of Tetracyclines anhydrous solution are usually employed for preparation of anhydro-tetra- cyclines,16~17~23~31~48-48 but strong organic acids are also s ~ i t a b l e . ~ ~ ~ ~ The reaction Table 1 Type of tetracycline Conditions Tetracycline O*2N-H2SOp 7-ChlOrO O*2N-H$O4 7-Bromo O.2N-H $0 4 Tetracycline 1 *ON-H,SO 5a-Epi 1 *ON-H,SO 6-Demethyl 1 *ON-H2SO4 6-Demethyl 3.ON-HCI 7-Chloro 1 *ON-H,SOd 7-C hloro- 6-deme t h yl I sON-H~SO~ 4-Epi 6-Demet hyl-4-epi 7-Chloro-4-epi 7-Chloro-6-demet hyl-4-epi 4-Epi-5-oxy Tetracycline 4-Epi 7-Chloro 7-Chloro-4-epi 4-Epi-5-oxyt 5-OXY 5-Oxyt 7-Chloro 5-OXY 5-OXY 5-OXY 5-OXY 5-oxy ** ButTerlpH 1 BufferlpH 2.5 BufferlpH 4-6 BufferlpH 5-5 BufTerIpH 7.0 Temp.Half-lve* 100" <2 min. 100 8.2 min. 100 18.8 min. 100 <1 min. 100 67 min. 100 24.8 min. 100 1-4 min. 100 2.1 min. 100 445 min. 100 0.9 min. 100 25-8 min. 100 5.8 min. 100 322 min. 100 4.5 min. 100 2.6 min. 24 15.5 hr. 24 24 hr. 50 7.3 hr. 50 12-8 hr. 50 6.3 hr. 50 6.0 hr. R 14 day 37 114 hr. 37 134 hr. 37 45 hr. 37 45 hr. 37 26 hr. Ref. a a a b? c b d b b b b b b b b e e e e e e C f g g g g g * The results are approximate; they were obtained by measuring changes in bio-activity. Aqueous solutions were used throughout and products were not isolated. t The anhydro-derivatives are degraded even under these conditions. ** Solution of hydrochloride (pH 2.5-2.8); R = room temperature.4 Ref. 19; * Ref. 21; C Ref. 26; d J. R. D. McCormick E. R. Jensen P. A. Miller and A. P. Doerschuk J. Amer. Chem. SOC. 1960 82 3381; 6 Ref. 43;f Ref. la p. 782; P. P. Regna and A. I. Solomons Ann. New York Acad. Sci. 1950,53,229. 46 R. K. Blackwood and C. R. Stephens Canad. J. Chem. 1965,43,1382. 47 J. J. Hlavka and H. M. Krazinski J. Org. Chem. 1963,28,1422 footnote 4a. 48 A. Green R. G. Wilkinson and J. H. Boothe J. Amer. Chem. SOC. 1960,82,3946. '@ B.P. 785,047. 440 Ciive is successful with a wide variety of tetracyclines including those modified in ring A (see p. 449) but is sometimes c~mplicated~~ by epimerisation at C(4). Tetracycline itself is stable for long periods in O.O3~-hydrochloric acid at room temperature but degradation occurs if the pH is lower.& Further data are collected in Table 1.The comparative stability of 6-demethyltetracycline and of the partially synthetic compound Sa-epitetracycline is not unexpected for in one case the hydroxyl group is secondary and in the other it is cis to the adjacent hydrogen atom. What is noticeable however is the influence of a halogen at C(7). The effect is probably a steric one32 and in the case of the 6-demethyltetracyclines it is very pronounced. The entries show that 7-chloro-6-demethyltetracycline (which is a fermentation product) can withstand acid conditions of comparative severity and the same may be true of immediate precursors in synthesis. This exception apart only mild conditions can be tolerated for any length of time. Table 1 also shows that the orientation of the dimethylamino-group does have some effect but the published results are inconsistent.Those for the lower temperature range are likely to be the more reliable. B. Reactions in Basic Media.-(Cf. the 6-deoxytetracyclines p. 455.) The tetracyclines are also very sensitive to bases and data are summarised in Table 2. Again a number of trends are apparent. The absence of a methyl group from C(6) greatly improves base-stability and the presence of a halogen at C(7) also has a strong effect. Tetracycline is more base-resistant than aureomycin; but in the 6-demethyl series the chlorinated compound is the more stable. The results also suggest that it might be preferable to work with certain compounds in their Cepi-modification and to epimerise them at a later stage. f OH Me NMe -0" Me NMe2 d \ CONH HO 0 O O H 0 (9) iso-aureomycin Scheme I The initial products of base degradation are the isotetracyclines [e.g.iso- aureomycin (9)] and they are formed as shown in Scheme 1 (ionisation of acidic groups not shown in the Scheme. An alternative process takes place in triethyl- 6o U.S.P. 2,990,426. 441 Chemistry of Tetracyclines Table 2 Type of tetracycline Tetracycline 7-Chloro 7-Bromo Tetracycline 6-Demethyl 7-Chloro 7-Chloro-6-demethyl 5-OXY Tetracycline 4-Epi 7-Chloro 7-Chloro4epi 5-oxy 4-Epi-5-oxy Tetracycline Tetracycline 7-chlOrO Tetracycline 7-ChlOrO 5-OXY 5-OXY 5-OXY Conditions Buffer/pH 10 Buffer/pH 10 Buffer/pH 10 0.1 N-NaOH 0.1 N-NaOH O*lN-NaOH 0.1 N-NaOH 0.1 N-NaOH 0.1 N-NaOH 0.1 N-NaOH BufTerlpH 8-8 Buffer/pH 8.8 0.1N-NaOH O.1N-NaOH 0.1 N-NaOH 5% NaHCO 5 % NaHCO Buffer/pH 8.85 pH 8.5 Buf€er/pH 7 BufTerlpH 8.5 BuRer/pH 10 Temp.* Hau- Iifet 22" >600 min. 22 18.6 min. 22 10.9 min. 100 6.8 min. 100 2.2 min. 100 32 min. 100 <0.3 min. 100 40 min. 60 101 min. 60 225 min. 29 53 min. 29 154 min. 23 10.9 hr. 23 21-8 hr. 25 30-34hr. R 24-3Ohr. R 1- 3hr. R ca. 12 hr. R 4 hr. 37 26 hr. 37 33 hr. 37 14 hr. Ref. a a a b b. c b. c b b d d d d d d e f f g h i i i * R = Room temperature. f The results are approximate; they were obtained by measuring changes in bio-activity. Aqueous solutions were used throughout and products were not isolated. aRef. 19; bRef. 21; CTable 1 ref. d; "Ref. 43; 6U.S.P. 3,122,578;fRef. 30; #Ref. 18; h Ref. la p. 782; i Table 1 ref. g. aminesl under reflux). In the case of terramycin the hydroxyl group on ring B permits further degradation and isoterramycin has not been reported.C. Formation of Metal Complexes.-It might be expected that degradation by alkali would be slower if the /%dicarbonyl system at C(ll)-C(12) were stabilised by chelation. The tetracyclines indeed form complexes with many metal cations52 and the complexes do show enhanced alkaline stability. However the site of complex-formation remains to be settled. Model compounds representing 51 J. S. P. Schwarz and H. E. Applegate J. Org. Chem. 1967,32 1241. 62 L. Z . Benet and J. E. Goyan J. Pharm. Sci. 1965,54983 and refs. therein. 442 Clive ring A or the BcD-system are themselves chelating agents,53,54 but these two sections of the tetracycline structure are not completely insulated by the tetra- hedral 12acarbon atom for there is a large degree of interaction owing to extensive intramolecular hydrogen bonding.37 Early work drew attention to the effect of metal ions on the ultraviolet spectrum specifically the band at ca.370 mp of terrmycin with the corresponding effect of the ions on model compounds and led to the suggestion that the C(ll)-C(12) system is the major site of complex-formation.” Spectroscopic measurements on solid complexes55 indicate that the molecules co-ordinate through oxygen and there is some additional evidence from solu- tion studies that the dimethylamino-group is not involved.58 Another approach:’ requiring potentiometric titration of tetracyclines in the presence and absence of metal cations has also been published; unfortunately the results are difficult to interpret.Recent work has emphasised the fact that the macroscopic pK,‘s observed for tetracyclines e.g. the values for tetracycline hydrochloride in aqueous solution ares ca. 3.3 ; 7.7; 9.7 each represents ionisa- tion at more than one position on the m o l e c ~ l e . ~ ~ ~ ? ~ ~ Probably more than one site is co-ordinated both in the solid and solution phases (the nature of tetra- cycline complexes in solution depends on the pH80) and the nature of the cation itself could be a deciding structural factor. The matter is important because it is linked to the mode of action of the antibiotics; however what is to be noted here is the established fact that complex-formation suppresses the reactions that normally take place in basic media. A number of high-molecular-weight complexes have been preparede1Ssa of general formula (tetracycline group antibiotic) (aluminium),(calcium)m)(gluconic acid) (no comment on the ionic status of the components is implied by this formulation) where a 6 and c the molar ratios of the respective constituents can vary over a wide range.Greatly enhanced alkaline stability occurs when a and b have certain values (Table 3). The molar ratio of gluconic acid appears to play an important part only in solubilising the derivative. Of greater interest is the uses3 of concentrated ammonia solution (28 % w/w) in the presence of magnesium chloride to hydrolyse the diester (10; R = Ac) to the rnono-ester (10; R = H). In the absence of the metal cation degradation is ‘significant’ after about 1 hr. and the procedure described actually calls for a 53 J.L. Colaizzi A. M. Knevel and A. N. Martin J. Pharm. Sci. 1965,54 1425. 64L. H. Conover in ‘Symposium on Antibiotics and Mould Metabolites’ Chem. SOC. Special Publ. No. 5 1956 p. 48. 66 W. A. Baker jun. and P. M. Brown J. Arner. Chem. SOC. 1966,88,1314. 66 F. Z. Benet and J. E. Goyan J. Pharm. Sci. 1966 55,1184. 6l J. T. Doluisio and A. N. Martin J. Medicin. Chem. 1963 6 16. 68 L. J. Leeson J. E. Krueger and A. Nash Tetrahedron Letters 1963 1155. 60 K. K. Kalnin’sh and B. G. Belen’kii Proc. Acad. Sci. (U.S.S.R.) 1964 157 721. 6o A. Albert Nature 1953 172 201. Pharm. Sci. 1962 51 86. 1964,53 1453. E. G. Remmers G. M. Sieger A. P. Doerschuk L. Ritter and J. F. Weidenheimer J. C’ E. G. Remmers W. C. Barringer G. M. Sieger and A. P. Doerschuk J. Pharm. Sci, U.S.P. 3,047,617. 443 Chemistry of Tetracyclines Table 3 Compound Conditions Temp.Half-life Ref. 7-Chloro-6-demethylTC* 0.1N-NaOH 90-100" 0.67 hr. a 7- Chloro-(6-demethylTC)(Al) (fi)(Glu)t 1 :4 Oh-NaOH 90-100 50 hr. a :2 :12 * 7-Chloro-6-demethyltetracycline. 7 No comment on the ionic status of the components is implied by this formulation. Glu = gluconic acid. a Ref. 61. reaction time of 5 min. With magnesium chloride however a reaction time of 5 hr. is specified. Me OH OH NMe M.e OH NMe a O H \ \ ! HO ! 0 CONH (10) m O H \ HO 0 \ 0 ~ 0 ~ ~ CONH 2H,NNHZ (11) Ro 0 NMe / OAc CONH Ho N-N It is also relevant that tetracycline forms addition products with hydrazinee4 and can be recovered by the action of water. Aureomycin on the other hand cannot be regenerated from its (single) addition compound because the latter is too readily converted into a pyrazoline[ (11) -+ (12)].frhe configuration at C(4) is probably a in compound (11). The tautomer shown (12) is likely to be more important than that given in the literat~re.~~] Although the C(6) hydroxyl group plays a central role in reactions brought about by acids and bases it also participates in a number of non-degradative transformations which temporarily modify its own activity as well as that of other positions on the tetracycline skeleton. These reactions are examined next. D. 6,12-Hemiketals.-When the amphoteric compounds specified by formula (13) as well as certain others of less interest are treated with perchloryl fluoride (FCIO,) in the presence of a base (at least one equiv. of base per mole of sub- strate is required; the experiment is done with ice-bath cooling) there are obtained crystalline products for which the 1 la-€luoro-6,12-hemiketal structures (14) have been e ~ t a b l i s h e d ? ~ * ~ ~ ~ ~ Formation of these derivatives is sterically O4 U.Valcavi G. Campanella and N. Pacini Guzzerru 1963 93,916. 66 U.S.P. 3,109,007. *( U.S.P. 3,165,551. 444 Clive R,' OH R2 NMe R' 0 R2 NMe W H H'. H €-I" CONH \ \ ! CONH \ Ho 0 HO b &OHo (14) HO OHO (13) R' R2 (a) Me H (b) Me OH possible only if the C(5a) hydrogen atom and the C(6) hydroxyl group are trans so that isolation of compound (14c) shows that the 6-demethyl-tetracyclines have the same relative stereochemistry as does the parent series. The 1 la-fluorine atom in 6,lZhemiketals is not especiaIIy labile but it can be removed. In the case of (14a) for example hydrogenolysis regenerates the anti- biotic though a significant amount of anhydrotetracycline is also formed.An improved yield of tetracycline (ca. 60 % of the total product as against ca. 26 %) is obtained by using zinc-hydrochloric acid ( 0 . 2 ~ ) for the r e d ~ c t i o n ? ~ ~ ~ When a C(6) hydroxyl group is absent (see p. 455) a hemiketal cannot form and compounds such as (15; X = F) are produced i n ~ t e a d . ~ ~ ~ ' It is noteworthy that the 1 la-fluoro-hemiketals (14) resist the usual dehydrating action of acids and can even survive treatment with boiling methanolic hydrogen ~ h l o r i d e . ~ ~ ~ ~ Their stability is attributed to the fluorine atom which prevents enolisation of the C(11) carbonyl group. Treatment of amphoteric tetracyclines with one equivalent of N-chloro- succinimide in glyme (ethylene glycol methyl ether) a f f o r d ~ ~ ~ ~ ~ analogous chloro-hemiketals (16) (use of two equiv.leads to 7,l la-dichloro-hemiketals). Again in the &deoxy-series hemiketal formation is impossible so that the -OH CONH \ Ho 0 &OHo (16) R' R2 (a) Me H (6) Me OH (4 H H 67 H. H. Rennhard R. K. Blackwood and C. R. Stephens J. Amer. Chem. Soc. 1961 83 2775. 445 Chemistry of Tetracyclines producP is of type (15; X = CI). The lla-bromo- and lla-iodo-analogues of both series prepared similarly are sensitive compounds. Unlike its fluoro-analogue compound (16a) exists in solution partly as the C(12) ketonic form (17). [Probably the same is true of compounds (166) and (16c) though spectroscopic evidence has not been reported.] In this tautomer as with other compounds [e.g.(15)] having a carbonyl group at C(12) the halogen atom is activated and removable by reduction under mild ~ o n d i t i o n s . ~ ~ ~ ~ Catalytic reduction of (16a 17) yields a mixture of about equal parts of tetra- cycline and anhydr~tetracycline.~~ Usees of sodium dithionite might give a better result. Me 0 NMe m o H H= H.' -OH H- \ CONH \ CONH HO ci -- Ho 0 Br OH 0 OH 0 (18) CH R NMe R X (a) H Cl @JJJH CONH ( b ) (C) OH H C1 F \ HO OH (20) ( 4 O H F 0 0 0 Interestingly the 11a-brorno-4-dedimethylamino-6,12-hemiketal (18) has been obtained crystalline not only as shown but also in its C(12) ketonic f 0 1 - m . ~ ~ ~ ~ ~ As another consequence of its tatuomeric nature hemiketal (16a) does not have the acid stability of its fluoro-analogue and is converted into lla-chloro- isotetracycline (19) by hot methanolic hydrogen In contrast exocyclic dehydration takes place45 in liquid hydrogen fluoride (hydrogen fluoride is the preferred dehydrating acid) and the reaction which is almost instantaneous yields the 1 1 a-chloro-6-methylene derivative (20a).Compounds (20b)-(20d) are f ~ r m e d ~ ~ ~ in the same way from the correspond- ing 6,12-hemiketals but the rate of dehydration is lower. Consequently it is possible to manipulate some 6,lZhemiketals in an acidic medium in order to effect substitution of ring D. In the absence of a C(6)-methyl group 1 la-halogeno- 6J2-hemiketals are even more stable to acids so that substitution of ring D poses fewer problems. Halogenation (N-halogenosuccinimide-liquid hydrogen fluoride) takes place at C(7) while nitration (potassium nitrate-liquid hydrogen fluoride) occurs predominantly at C(9).The following sequencee5 is typical of those reported (no yield is given) 7-Chloro-6-demethyltetracycline was con- verted into 7,l la-dichloro-6-demethyltetracycline-6,12-hemiketal (N-chloro- succinimide-glyme) and nitrated at C(9). Reduction (sodium dithionite) afforded 9-amino-7-chloro-6-demethyltetracycline. '* R. K. Blackwood J. J. Beereboom H. H. Rennhard M. Schach von Wittenau and C. R. Stephens J. Amer. Chem. SOC. 1961 83 2773. 446 Clive The 6-methylene derivatives have a much greater acid stability than the natural tetracyclines and can therefore undergo acid-catalysed reaction^^^,^^ of the type referred to for the hemiketals. 1 la-Halogeno-6-methylene compounds are easily reduced45 to the corre- sponding 6-methylene-tetracyclines [e.g.(21)] zinc-mineral acid being the preferred reagent for removal of an 1 la-fluorine atom.65p66 The exocyclic course of the dehydration has been attributed to the presence of the 110-substituent. By preventing aromatisation of ring c it witholds a powerful driving force for 5~,6-dehydration.~~ The acid stability of the 1 1 a-halogeno-6-methylene com- pounds (20) is understandable on this basis but the explanation is incomplete CH R NMe H** H- / OH \ \ ' C O N H CONH R' R' X ( 4 H Me H ( b ) H Me C1 (c)CI H H (d) Br H H because the halogen-free compounds (21) also show enhanced acid-stability. They are intermediate in this respect between tetracyclines and 6-deoxy-tetra- cyclines (see p. 455).E. 4,6-Hemiketals.-The hemiketals of the previous section were prepared from amphoteric substrates. A slight change however affords an entirely different series of ~ o m p o u n d s . ~ ~ ~ ~ ~ ~ ~ For example treatment of an aqueous solution of tetracycline hydrochloride with N-chlorosuccinimide precipitates 4-dedimet hylamino-4-oxot etracycline-4,6-hemike tal (22a) a substance also known as 4-hydroxy-6-methyltetracycloxide. n h e presence of a little hydro- chloric acid is beneficial. Other oxidising agents have been but an acid must be present to bind the dimethylamine as it is released and the presence of water is e~sential.~~] Similar treatment of the 6,12-hemiketal (16~) in the presence of hydrochloric acid affords the new derivative (22b). The reaction is believed to proceed by the stages of Scheme 2 and this me~hanisrn~~s~~ is supported by the successful trap- ping of a 4-NN-dimethyliminium i~~terrnediate.'~ 69 B.P.995,032. 'O R. K. Blackwood and C. R. Stephens J. Amer. Chern. SOC. 1964 86,2736. 71 U.S.P. 3,247,226. 72 R. C. Esse J. A. Lowery C. R. Tamorria and G. M. Sieger J. Amer. Chem. SOC. 1964 86 3874. 73 J. S. P. Schwarz H. E. Applegate J L. Bouchard and F. I . Weisenborn J. Org. Cliem 1967,32 1238 447 Chemistry of Tetracyclines + NMe 0 + + H HNMe C1 HNMe ) "z0) A ,4,6- herniketal A-A Scheme 2 The structures (22) which were deduced from chemical experiments demand the stereochemistry shown and X-ray data74 for the 7-halogeno-deri~atives~~,~~,~~ confirm the assignment and indicate an unusual situation in ring A. It appears that the chloro-compound (22c) is enolized with a double bond at C(2)-C(3) but that the bromo-analogue (226) has the double bond at C(l)-C(2).The tetracycloxides have been examined in detail and some important trans- formations are established. Hydrazine and hydroxylamine react at room temperature with tetracycloxide (22a) to produce the hydrazone (23a) and the oxime (23b) respectively and both products46 can be reduced to the 4-epi-compound (24). The same material is obtained46 directly from tetracycloxide (22a) by hydrogenation in dimethyl- formamide containing ammonium hydroxide and magnesium chloride. In another study compound (25) was hydrogenated in the presence of various primary a m i n e ~ . ~ ~ ~ ~ A slight excess of amine provided the requisite alkaline /R M,e OH Me OH H. NH -0I-i H ..W O H CONH \ \ . CONH HO \ \ * HO HO h0 (24) HO OHo (23) 0 R (b) OH (4 NH conditions but promoted destruction of the starting material. As a result the products [e.g. (26; R1 = H; R2 = Me or Et)] were obtained only in 2 0 4 0 % yield. Certain of these secondary amines were alkylated reductively and then epimeri~ed.~~ In this manner (25) was reconverted into 6-demethyltetracycline by way of intermediates (26; R1 = H; R2 = Me) and (26; R1 = R2 = Me). F. Photo-oxygenation.-Although the C(6)-hydroxyl group survives many reactions which involve changes at various positions of the tetracycline skeleton it might be desirable to introduce the group in the last stage of synthesis. The feasibility of such an approach was established by the d i s c o ~ e r y ~ ~ ~ ~ ~ that 74 J. H. van den Hende J.Amer. Chem. SOC. 1965 87,929. 76 U.S.P. 3,159,675. 76 R. C. Esse J. A. Lowery C. R. Tarnorria and G. M. Sieger J. Amer. Chem. SOC. 1964,86 3875. 77 M. Schach von Wittenau J. Org. Chem. 1964,29,2746. 448 Clive anhydroaureomycin reacts with oxygen in the presence of light to give the hydroperoxide (27). (Tautomers with the internal double bond in the 5,5a- position are believed to exist.77) OH CONH (25) 0 (26) Hydrogenation cleaves the peroxide (G-0-OH -+ C-OH) and then effects removal of the chlorine atom and saturation of the internal double bond.The last process takes place from both faces of the molecule so that a mixture of tetracycline and its Sa-epimer is formed. The photoreaction is accelerated by traces of 3,4-benzopyrene and works well in the cases that have been examined except when a hydrogen atom is at C(7).In this instance ring D is probably attacked for anhydrotetracycline gave no crystalline product although the presence of some tetracycline was detected chr~matographically.~~,~~ This result is disappointing because anhydrotetracycline has been ~ynthesised~~ so that the reaction would complete the first total synthesis of the antibiotic. G. Modiiications to Ring A.-Reversible epimerisation at C(4) mentioned earlier has been carried out on numerous corn pound^.^^ The process occurs in a variety of solvent-systems normally within the pH range 2-6 and is acceler- ated43,80 by certain anions such as citrate phosphate or acetate. Below pH 2 and above pH 9 the rate is extremely The technique is successful with most fermentation tetracyclines and their 6-deoxy- (see p.455)81 and 5a,6- anhydro-derivatives,43 although tedious separation methods may be required. A more convenient procedure r e p ~ r t e d ~ ~ ~ ~ for a few tetracyclines involves preparation of a metal complex in situ (usually with calcium ions) and adjust- ment of the pH to 8-5-10.0 Under these conditions 4-epi-compounds are isomerised to products of normal configuration; apparently the reverse process does not occur to any significant extent. 78 Cf.Belg. P. 631,118. 79A. I. Gurevitch M. G. Karapetyan M. N. Kolosov V. G. Korobko V. V. Onoprienko S. A. Propravko and M. M. Shemyakin Tetrahedron Letters 1967 13 1. E. G. Remmers G. M. Sieger and A. P. Doerschuk J. Pharm. Sci. 1963 52 753. Table 1 ref. d. U.S.P. 3,009,956. 449 Chemistry of Tetracyclines Epimerisation involves an enolic tautomer such as (28) with a double bond at C(3)-C(4).It is understandable because of electron delocalisation in ring A that changes at the site of the carboxyamide group should affect the ease of the reaction. Attempts to epimerise 2-acetyl-2-decarboxyamidotetracyclines~~~ and some tetracyclino-nitriles (formed by dehydration of the carboxyamide group; see p. 454)43 failed so that it appears necessary for the carboxyamide group to be intact. Some modification is permissible however for the t-butyl derivative (29) undergoes the changes in ultraviolet absorption that characterise equilibration of the C(4) epimers of anhydrotetracycline.84 HO Me y e The extent of alkylation of the C(4)-nitrogen atom is also an important factor. Primary and secondary amino-groups are not epimeri~able'~ (at least under normal conditions; for a possible exception see ref.46) but tertiary amino- systems involving a range of alkyl groups are readily epimeri~ed.~~ In addition to the conformational changes the dimethylamino-group can also be removed. One procedure discovered with tetracycline and aureomycin calls initially for prolonged treatment of the antibiotics with methyl iodide. Crystalline products (30a) and (30b) are obtained though in the case of terra- mycin general decomposition occurs.8s Tetracycline methiodide has the a-con- + I - W O H NMe \ \ CONH Ho 0 HO OHo (30) OH 8s J. Keiner R. Hiittenrauch and W. Poethke Arch. Pharm. 1967 300 840. 84 Ref. 43 footnote 16. 86 J. H. Boothe G. E. Bonvincino C. W. Waller J. P. Petisi R. W. Wilkinson and R.B. Broschard J. Amer. Chem. SOC. 1958 80 1654. 450 Clive figuration at C(4) and is readily epimerised.& Presumably the same is true of aureomycin methiodide. When compounds (30a) and (306) are each treated briefly with zinc dust in aqueous acetic acid the dedimethylamino-tetracyclines (31a) and (316) are obtained in good yield.85 This method is also applicable to 6-deo~ytetracyclines.~~ All the dedimethylamino-compounds (31) have been produced directly from the natural products,8' though in poor yield,29,30,85 by prolonged action of zinc in aqueous acetic acid. In this method the by-products area5 the dedimethyl- amino-1%-deoxy-tetracyclines (32) [other tautomers are possible; in methanolic hydrogen chloride (0.01~) it has been established88 that C(l1a) is tetrahedral] and by increasing the reaction time still further these are obtained as major Two methods are established for selective removal of the 12a-hydroxyl group.One of these require^^^,^^ the use of zinc dust in dilute ammonium hydrox- ide and in the case of tetracycline for example affords the 12a-deoxy-compound (33) in 42% yield. This compound has a time-variable ultraviolet absorption spectrum [methanolic hydrogen chloride (0.01~)] owing to equilibration of tautomer~.~~ The orientation of the dimethylamino-group has not been settled. In one case,g4 rehydroxylation at C(12a) gave the biologically inactive compound prod~cfs.29,30,89,90 H&OH -OH / C ONH \ \ ; corn2 OH 0 (33) Ho 0 Ho 0 (34) CHO 4epitetracycline but other experimentsg5 appear to have given the active (normal) epimer.It has been observedss that 12a-deoxytetracyclines are far more resistant to epimerisation than their parent antibiotics. The other p r o ~ e d u r e ~ ~ ~ ~ ~ ~ ~ ~ involves preparation of a 12a-0-acyl derivative usually the O-formate [e.g. (34)]. This derivative is obtained by the action of acetic-formic anhydrideg8 in pyridine [with terramycin the reaction is more C. R. Stephens J. J. Beereboom H. H. Rennhard P. N. Gordon K. Murai R. K. Blackwood and M. Schach von Wittenau J. Amer. Chem. SOC. 1963 85 2643. ~3' U.S.P. 2,786,077. 88 H. Muxfeldt W. Rogalski and K. Striegler Chem. Ber. 1962 95 2581. 89 Ref. 88 footnote 23. Cf. T. L. Fields A. S. Kende and J. H. Boothe J. Amer. Chem. SOC. 1960 82 1250. 91 Belg. P. 572,382. 9a A. Green and J. H. Boothe J. Amer. Chem. SOC.1960 82 3950. 98 R. K. Blackwood H. H. Rennhard and C. R. Stephens J. Amer. Chem. SOC. 1960 82 5194. 94 H. Muxfeldt G. Buhr and R. Bangert Angew. Chem. Internat. Edn. 1962 1 157. 96 U.S.P. 3,043 877. 96 U.S.P. 3,043,876. O7 U.S.P. 3,002,021. O8 L. F. Fieser and M. Fieser 'Reagents for Organic Synthesis' J. Wiley and Sons New York 1967 p. 4. 451 Chemistry of Tetracyclines complicateds3 but other 12a-O-acyl derivatives can be used;Q7*QQ during some acylations the C(4) position is epimerised] and is then hydrogenolysed [a C(7)- halogen atom is lost in the processs6] under mild conditions [e.g. (34) -+ (3311. Again the stereochemistry at C(4) has not been established. The 12a-position in the deoxy-compounds is an activated site. Treatment of dedimethylamino-l2a-deoxytetracycline with one equivalent of N-bromo- succinimide affords the 1 2a-&bromo-derivative.48 More significantly a hydroxyl group can be re-introduced at C(12a) in the correct stereochemical sense.[Strong evidence for the desired stereochemical result at C(12a) has not always been published; the C(4) position can be epimerised if necessary.] Microbio- logical hydroxylation has been reportedloo for 12a-deoxytetracycline but this compound as well as many of its derivatives including 5a,6-anhydro-tetra- cyclines can be hydroxylated by chemical ~ ~ i d a n t ~ . ~ ~ J ~ ~ J ~ ~ Use of metals in the elemental form or as salts together with gaseous oxygen appears to be a very convenient procedureg4Jo3Jo4 and in one which was examined with special care none of the unnatural lh-epimer could be detected.The molecule in question was judged therefore to have a /%oriented dimethylamino-group at C(4) [whose bulk shields the face of C(12a)l. Peracids have been used for 12a- epi-hydroxylation of certain compounds.34~54~10g H. Aromatisation.-Besides Sa,6-dehydration which usually leads to naph- thacenic products other types of aromatic compound have been made from tetracyclines. Heating 12a-O-acyl derivatives causes cis-elimination of a carboxylic acid and yieldsQ3J07 products such as (35a) (this compound has a time-variable ultra- violet absorption spectrums3). Analogous compounds lacking the dimethyl- amino-group are also ~btainable.~~ Treatment of 12a-deoxytetracycline with methyl iodide in tetrahydrofuran containing an acid scavenger givesg3~Q6 the derivative (356) which also has a time-variable ultraviolet spectrum.This compound is available from the same starting material by the actiong2 of per- benzoic acid or in higher yield by treating 12af-bromodedirnethylamino-12a- deoxytetracycline with pyridine.48J08 The first of these three methods involves quaternisation (and epimerisation) of the C(4) substituent while the second proceeds by Cope rearrangement of an N-oxide. Both compounds (3%) and (356) are fdly aromatised by the action of acids gQ U.S.P. 2,976,318. loo C. E. Holmlund W. W. Andress and A. J. Shay J. Amer. Chem. Soc. 1959,81,4750. lol C. E. Holmlund W. W. Andress and A. J. Shay J. Amer. Chem. SOC. 1959,81 4748. lo8 Cf. P. Shu J. Amer. Chem. SOC. 1966 88 4529. lo9 B.P. 947,601. lOd U.S.P. 3,188,348. lob R. B. Woodward Pure Appl. Chem. 1963,6,561.lo6 H. Muxfeldt and A. Kreutzer Chem. Ber. 1961,94 881. le7 J. R. D. McCormick S. Johnson and N. 0. Sjolander J. Amer. Chem. Soc. 1963 85 1692 footnote 3. lo* CJ J. R. D. McCormick J. Reichenthal S. Johnson and N. 0. Sjolander J. Amer. Chem. SOC. 1963 85 1694. 452 Clive to the naphthacenes (36a) and (366) re~pectively.~,~~ Compound (366) is known as 6-methylpretetramid and is a biogenetic precursor6 of tetracyclines having a methyl group at C(6). The next stage of biogenesis isSs9 oxidation to 4-hydroxy-6-methylpretetramid (36c) and this too can be made from tetracycline. Dissolution of tetracycline methiodide in water and adjustment of the pH to 4-5 affords 85e109 the crystal- line betaine (37) and pyrolysis1l0 in anhydrous acetonitrile converts this sub- stance into the phenolic diketone (38).The latter gives 4-hydroxy-6-methyl- pretetramid (36c) on treatment with acid.l1° The thermal reaction probably + M,e OH NMe Me OH H" ; y a O H \ \ CONHz . . \ \ * ' -cow2 \ Ho 0 *O O H 0 (37) Ho 0 Ho Ho (38) proceeds by Hofmann elimination [(37) --+ (39)] followed by ring cleavage [(39) - (@)I. Similar degradations have made a number of other pretetramids avai1able.lleu2 Me OH M e OH m O H W O H CONH ,(38) \ / CONH \ Ho H d b oHO) (39) HO 0 c,o OH (40) Me OH 0 log Cf. Ref. 46 and footnote 3 therein. ll1 J. J. Hlavka P. Bitha and J. H. Boothe Tetrahedron Letters 1967 1139. J. J. Hlavka P. Bitha and J. H. Boothe J. Amer. Chem. SOC. 1965 87 1795. Cf. U.S.P. 3,226,305. 453 2 Chemistry of Tetracyclines Interestingly the A-D aromatic compound (38) has been oxidised113 to the quinone (41).In principle hydration of the 4a 12a-double bond could give a tetracycloxide bringing the sequence full circle to tetracycline. I. Miscellaneous Reactions.-Although the carboxyamide group survives all the reactions described so far there axe of course circumstances in which it is modified. Dehydration to a nitrile group takes place on treatment with dicyclo- hexylcarbodi-imide114 or with an acid chloride in pyridine and in the latter case esterification of the phenolic hydroxyl group at C(10) may compete with de- hydrat ion.l16 Of pharmaceutical interest is the improved water-solubility generally con- ferred on a tetracycline by aminomethylation.116-118 For example treatment of tetracycline with morpholine and formalin yieldslle the zwitterionic and highly water-soluble compound N-(morpholinomethy1)tetracycline (42a) apparently without epimerisation at C(4).The antibiotic can be regenerated by hydro- NMe2 R n (a) -CH,-N 0 W 'ONHR (b) -CH,- OMe genolysis over Raney nickel (43 % yield) or by treatment with aqueous sodium hydrogen sulphite (96 % yield).llg What information there is available shows aminomethyl derivatives to be subject to easy hydrolysis120-122 to the parent antibiotic and this reaction probably accounts for their biological activity. In an analogous type of compound exemplified by the derivative (42b) the substituent on the carboxyamide nitrogen atom is probably less labile (under acidic conditions hydrolysis takes place) and has been shown to be suitable for protecting the group from deh~drati0n.l~~ Another subject of experimental interest is the response of the tetracyclines to hydrogenation conditions.Very substantial changes can take place but methods are available for selective reactions. A C(ir)-halogen atom (which can be a useful feature during certain stages of synthesis) can be removed cleanly by high-pressure reduction over a palladium catalyst in basic ~ o l u t i o n . ~ ~ J ~ ~ ~ ~ ~ ~ ~ ~ ~ lls J. E. Baldwin D. H. R. Barton L. Bould and P. D. Magnus Chem. Comm. 1967 319. 114 G.P. 1,091,564. B.P. 766,512. 118 G.P. 1,088,481. 11' B.P. 809 585. 11* U.S.P. 3,104,240. 119 W. J. Gottstein W. F. Minor and L. C. Cheney J. Arner. Chem. Soc. 1959 81 1198. 120 R. Huttenrauch and J. Keiner Naturwiss. 1966 53 552. 121 A. Brunzell Acta Chem. Scand. 1962 16 245. 122 See M. J. Martell jun.A. S. Ross and J. H. Boothe J. Medicin. Chem. 1967 10 485. l*s C. R. Tamorria and R. C. Esse J. Medicin. Chem. 1965 8 870. le4 Cfi Ref. 16. G.P. 1,007,775. 454 Clive Hydrogenolysis of the benzylic hydroxyl group at C(6) is also possible and for this purpose the use of a noble-metal catalyst (usually palladium or platinum) under acidic conditions is required.8g,12s Further addition of a chelating agent such as boric acid is recommended81,126 in order to protect the C(ll)-C(12)- dicarbonyl system from reduction. [A C(7) halogen atom is usually127 lost except when a rhodium catalyst is employed in which case partial retention is observed.128 With a rhodium catalyst a wider pH range is permissible (e.g. as high as pH 9-5) and a chelating agent is not required.] The acidic conditions used do not by themselves promote 5a,6-dehydration but the noble metal in the presence of hydrogen appears to have a catalytic effect on this processs1 and the anhydrotetracyclines produced initiate a series of competing side reactions; consequently yields are very low.In the case of 6-demethyltetracyclines which are more resistant to acid yields of 3 0 4 % have been ~btained.*~J~~ Hydrogenolysis of the hydroxyl group at C(6) is accompanied by stereo- chemical inversion at that site;86 therefore tetracycline affords 6-deoxy-6P- tetracycline (43). 6-Deoxy-compounds of natural configuration are available SPh H Me NMe CH2 H w e 2 -OH CONH =OH \ \ CONHz \ \ * HO HO OHo (43) Ho 0 Ho O H 0 (44'1 from the corresponding 6-methylenetetracyclines. The latter react with thiols in the presence of a free-radical initiator to give45 compounds such as (44).Desulphurisations6 does not disturb the stereochemistry at C(6). Alternatively direct reduction of 6-methy1enetetracyclines yields a mixture of the epimeric 6-deoxy-compounds.69~86 The 6-deoxytetracyclines undergo many of the reactions already discussed such as formation of complexes and the various transformations characteristic of ring A but in contrast to the natural tetracyclines they are dramatically resistant to the action of acids and bases (see Table 4). Table 4 Compound Conditions * 6-deox y- 6P-te t r ac ycl ine 3N-HCI tetracycline 3N-HC1 6-deoxy-6@-tetracycIine 0.1 N-NaOH tetracycline a Table 1 ref. d. 0.1 N-NaOH * All at looo Half-life( m in.) 1 600 a 570 6.8 Ref. a a a a lB6 B.P. 855,170. 12' Cf. B.P.871,423. "* U.S.P. 3,019,260. ''~3 U.S.P. 2,999,111. 455 Chemistry of Tetracyclines Their enhanced stability has made possible substitution in the aromatic ring D a process which usually requires the presence of a strong acid. m e strongly acid conditions used should not allow epimerisation at C(4); however definite evidence for the stereochemistry has not often been sought.] Nitration which is often carried out in concentrated sulphuric acid generally affords mixtures of 9- and 7-nitro-products (if both sites are a ~ a i l a b l e ) l ~ ~ - l ~ and the new compounds provide access to a wide range of derivatives. Reduc- tion gives the corresponding a m i n e ~ ~ ~ J ~ ~ J ~ ~ which can be a ~ y l a t e d l ~ ~ J ~ l or d i a z o t i ~ e d ~ ~ ~ ~ ~ ~ ~ and in the latter case nucleophilic substitution of ring D becomes p0ssib1e.l~~ C(9)-Diazonium compounds are useful as intermediates in the preparation of the corresponding halogeno- (fluoro- chloro- bromo- iodo-) deriva- t i v e ~ .~ ~ ~ ~ ~ The methods are also applicable138 in the C(7) series but bromina- tion and iodination at C(7) is achievable directly through the use of the appro- priate N-halogeno-amide in concentrated sulphuric a ~ i d . ~ ~ ~ J ~ ~ [Bromination studies revealed134 that in acetic acid halogenation occurs at C(l la).] Chlorina- tion at C(7) (N-chlorosuccinimide-trifluoroacetic acid) is more complicatedes for the product contains an 1 la-chlorine atom which must be removed and some substitution occurs at C(9). Besides making possible this partial synthetic work the convenient stability of 6-deoxytetracyclines has focused attention5 on them as objectives for total synthesis and in this particular area the problems have been solved.The pro- totype of the series 6-demethyl-6-deoxytetracycline has been obtained in two laboratories (as its racemate) by tptal s y n t h e ~ i s l ~ ~ J ~ ~ and the achievement marks a highlight in the synthetic field for the compound though not occurring naturally possesses the full biological activity characteristic of the tetracycline antibiotics. 130 J. L. Spencer J. J. Hlavka J. Petisi H. M. Krazinski and J. H. Boothe J. Medicin. Chem. 1963 6 405. 131 J. Petisi J. L. Spencer J. J. Hlavka and J. H. Boothe J. Medicin. Pharmaceur. Chem. 1962 5 538. 13* J. H. Boothe J. J. Hlavka J. P. Petisi and J. L. Spencer J. Amer. Chem. SOC. 1960 82 1253. 133 J. J. Beereboom J. J. Ursprung H. H. Rennhard and C. R. Stephens J. Amer. Chem. SOC. 1960 82 1003. la4 J. J. Hlavka A. Schneller H. Krazinski and J. H. Boothe J. Amer. Chem. SOC. 1962 84 1426. 135 J. J. Hlavka H. Krazinski and J. H. Boothe J. Org. Chem. 1962 27 3674. la6 B.P. 935,384. la7 H. Muxfeldt and W. Rogatski J. Amer. Chern. SOC. 1965 87 933.

ISSN:0009-2681    

DOI:10.1039/QR9682200435

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments By E. Kent Barefield and D. H. Busch CHEMISTRY DEPARTMENTy OHIO STATE UNIVERSITY COLUMBUS OHIO S. M. Nelson CHEMISTRY DEPARTMENT QUEEN’S UNIVERSITY BELFAST The measurement of magnetic susceptibility constitutes one of the most important experimental routes by which the chemist can easily and quickly acquire information on the structure and bonding in transition-metal com- plexes. There are available many books and reviews which deal in detail with theory measurement and interpretation of normal magnetic beha~iour.l-~ We here review representative examples of those classes of compound which exhibit so-called anomalous magnetic moments and examine the mechanisms producing these non-standard values. Reference will be made only to systems which do not require consideration of interactions between two or more metal ions; i.e.they are at least approximately magnetically dilute. Attention will be directed mainly to complexes of nickel(@ which provide the largest number of known examples of anomalous magnetic behaviour but reference will also be made to the increasing number of cobalt@) iron(@ and iron(m) complexes showing unusual magnetic behaviour. First it is necessary to define what we mean by ‘anomalous’ in referring to magnetic moments. For the present we take it to mean any magnetic moment for a metal ion in a discrete molecular species having a value which falls outside the range of values predicted from considerations of the spin and orbital angular momenta in ligand fields of given strength and symmetry.For the metal com- plexes to be considered the most common stereochemistries are octahedral tetrahedral square planar and distorted forms derived from these. Let us briefly illustrate what we mean by ‘anomalous’ by reference to what is ‘normal’ for regular complexes of the metals to be considered (Table 1). For nickel(@ in a regular octahedral complex consideration of spin-orbit coupling and contributions from the 3A2g and next higher 3T2g state (Figure 1) gives the maximum value somewhat above the spin-only moment of 2.83; the value6 is about 3-3 B.M. for a d value of ca. 8000 crn.-l and = 75 % of the free P. W. Selwood ‘Magnetochemistry’ 2nd edn. Interscience New York 1956. R. S. Nyholm J. Inorg. Nuclear Chem. 1958 8 401. 3 B. N. Figgis and J. Lewis ‘Modern Co-ordination Chemistry’ ed.J. Lewis and R. G. Wilkins Interscience New York 1960. B. N. Figgis and J. Lewis Progr. Znorg. Chem. 1964 6 37. B. N. Figgis ‘Introduction to Ligand Fields’ Interscience New York 1966. 6A.B.P. Lever Inorg. Chem. 1965 4 763. 457 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments Table 1 Expected ranges of 'normal' and 'anomalous' room-temperature moments (B.M.) for complexes of the nickel@) cobalt(n) iron(@ and iron(m) lono 6-Co- ordinate pseudo- octahedral NiII h.s. 3.0-3.3 NiII1.s. - COII h.s. 4-7-5.2 COII I.s. 1 *8-2.0 FeII h.s. 5.1-5.7 FeII 1.s. -0 FenCr h.s. -5.9 FG1.s. -2.3 ~ - C O - ordinate 3-0-3-45 -0 4-2-46 1.7-2.1 5.1-5.5 2.9-3.1 - - 4- CO- 4- co- Range for ordinate ordinate anomalous pseudo- planar moments tetrahedral -0-2.8 3 . 4 5 4 .0 - 0 - } -2.9-4.2 4 . 2 4 . 8 - 2.1-2.9 - a h.s. is high-spin and 1.s. is low-spin. b These are the ranges for both trigonal bipyramidal and square pyramidal geometries. C The moment of no species should fall within these ranges without being defined as anomalous. However a low-spin CoII complex known to have oh or pseudo-octahedral symmetry with a moment as high as 2.9 B.M. would also be considered anomalous etc. d This range excludes the low-spin 5-co-ordinate values as this geometry does not occur in any examples discussed. I- ___- A + Figure 1 Qualitative energy level diagram for d a systems showing triplet states and low-lying singlet states ion value. This also includes a small temperature-independent contribution. Ordinarily peff is in the range 3-0-3.3 B.M. although apparent values in excess of this have been noted.' Somewhat similar considerations apply to iron(@ iron(@ and cobalt(I1) ; however in some instances orbital angular momenta for the ground states of these ions in strict Oh symmetry are not quenched and the expression for /.i,eff is more complicated.For these cases there also exists the ' B. N. Figgis Nature 1958 182 1568. 458 Barefield Busch and Nelson possibility of ‘high spin’ (h.s.) and ‘low spin’ (I.s.) complexes with variations in the ligand field strength (Figures 2 3 and 4). a- Figure 2 Qualitative energy level diagram for ds systems such as iron(1II) a- Figure 3 Qualitative energy level diagram for d? systems like cobalt(1r) A --+ Figure 4 Qualitative energy level diagram for d@ systems like iron([$ 459 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments For the high-spin cases of C S F e y and F f l there is a rather complicated dependence on h and on kT.5,7 Calculations give room-temperature values for the high-spin moments of 4 7 and 5.1 B.M.for Con. For this ion as for all those where the ground-state configuration is Tl resulting from a free ion F ground state two cases must be considered in calculations; that where the ligand field is small compared with the interelectronic interactions and that where the field is large compared with interelectronic interactions. Values found experimentally fall in the range 4.7-5.2 B.M. depending on the ligand field strength and the amount of distortion from Oh symmetry. For iron(@ the free ion term is a 5D state and the absence of adjacent higher terms of the same multiplicity simplifies calculations.The ground state in Oh symmetry is V2 and pelf is expected to have a value of 5.6 B.M. at room temperature. Experimentally moments of 5-1-5.7 B.M. are observed. Iron(m) in weak ligand fields has a 6Al ground state coming from a free ion 6S state and moments of the order of 5-9 B.M. are generally found. The added possibility of spin pairing must be considered to determine the ranges in which moments are expected to fall. Iron@) is simple since with strong ligand fields a singlet state falls lowest and the complexes are diamagnetic. This occurs only for very strong fields such as those due to CN- CNR and o-phen- anthroline. For C# when the ligand field is strong enough (and this is relatively rare) the ground state is an E state with one unpaired electron.Since the orbital angular momentum is essentially quenched moments approaching the spin-only value of 1.73 B.M. are expected and values of 1.8-2-0 B.M. are observed. This may be only fortuitous however since the actual electronic structure can oiily be represented in terms of covalency theories (M.O. theory) and because a large amount of mixing between ligand and metal orbitals impairs the predictions of simple ligand field theory. Iron(m) has a low-spin 2T2g ground state with a normal moment of about 2.3 B.M. which corresponds to one unpaired electron. Since the ground state retains some orbital angular momentum moments may range in value from [4S(S + 1)13 to [4S(S + 1) + L(L + 1)]* the exact value depending on the temperature of measurement and the magnitude of the spin-orbit coupling.Mention should be made of the only known example of what appears to be a diamagnetic pseudo-octahedral Nin complex. Tris-(o-phenylenebisdimethyl- arsine)nickel(n) ion first reporteds by Nyholm in 1950 has been prepared in greater yields and its possible structure discussed in light of its diarnagneti~m.~ The nickel atom has peff 0.6 B.M. presumably owing to TIP (temperature- independent paramagnetism). Nuclear magnetic resonance (nmr.) results indicate that the structure must involve D symmetry with a pseudo-octahedral array of donors about the nickel atom. A qualitative molecular orbital scheme has been offeredg to account for the observed diamagnetism. The principal feature of the model involves treatment of the orbitals of the nickel atoms as a set of functions quantised about the C axis in D symmetry.In such a treatment * R. S. Nyholm J. Chem. SOC. 1950,2061. B. Bosnich R. Bramley R. S. Nyholm and M. L. Tobe J. Amer. Chem. SOC. 1966,88,3926. 460 Barefeld Busch and Nelson the n and 0 bonding distinction between specific orbitals is lost and mixing of all orbitals of like symmetry can occur. A qualitative M.O. scheme is presented in Figure 5. This scheme accommodates the last two electrons in a single orbital Figure 5 A qualitative molecular orbital scheme for Ni(diars) as+ showing the correlation between the nickel and arsenic atom orbitals in D I symmetry. sufficiently separated from the next level so as to maintain spin pairing. As Nyholm and his co-workers point out it is not sufficient in this case to work with donor-group microsymmetry in explaining spectral and magnetic properties.In most instances involving bidentate chelates the chelate rings act as a minor perturbation on strictly octahedral functions whereas in Ni [diars],2+ this is apparently not true. For Ninin tetrahedral fields the ground state is an orbitally degenerate Tstate and relatively large contributions to observed moments from orbital angular momentum are expected and observed. For regular tetrahedral complexes perf should occur within the interval 3-5-4-0 R. M. However in pseudo-tetrahedral complexes containing different ligands well separated in the spectrochemical series there may be a loss of ground-state degeneracy giving a lower orbital contribution and hence lower moments. Tetrahedral Con complexes usually exhibit moments closer to the spin-only values and these are of the order 4-2-44 B.M.Tetrahedral iron(@ and iron@) complexes are somewhat rare and are not important to this Review. 461 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments Five-co-ordinate complexes of nickel(@ and cobalt(rr) are fairly numerous and are now structurally well characterised. In addition a smaller number of iron(@ complexes are found to be five-co-ordinate. Although extensive spectral work has been done no detailed magnetic study has appeared. For this reason the exact behaviour of such complexes cannot be predicted but approximate ranges for the experimental moments appear to be defined. Five-co-ordinate nickel(I1) may be either high-spin with two unpaired electrons and moments in the range 3-0-3.4 or low-spin with no unpaired electrons and moments close to zero.These values apply to geometries approximating both trigonal bi- pyramidal and square pyramidal forms. Cobalt(n) complexes which are five- co-ordinate may also exhibit high-spin moments of 4 - 2 - 4 6 (S = #) or low-spin moments of 1.7-2.1 (S = 4). Again the differences in geometry are not always obvious. Five-co-ordinate complexes of iron(@ are not included among the examples discussed in this Review but the high-spin range of moments appears to be 5.1-5-5 B.M. and the low-spin 2.9-3.1 B.M. These correspond to 4 and 2 unpaired electrons respectively. Square planar NiIIcomplexes have no unpaired electron and exhibit essentially zero moments or small contributions due to temperature-independent para- magnetism.Square planar Con species are however paramagnetic and are reported to exhibit moments in the range 2.1-2-9 B.M. having slightly larger orbital contributions than low-spin octahedral complexes.1° Theoretically high- spin square planar CoIIspecies are possible but none have yet been authenticated although there have been reports11,12 of such compounds which were later shown in actuality to be low-spin specieslaa (presumably contamination yielded erroneously high values) or to involve higher co-ordination numbers.13b The ranges of occurrence of magnetic moments for the normal complexes that provide the standard against which anomalous behaviour is measured are surnrnarised in Table 1. The ranges within which anomalous moments will fall for each metal are also shown.It is useful to recall the elementary fact that a condition for our defined normal paramagnetic behaviour (except for the specified cases involving Tground states) is found in the temperature-dependence of the magnetic susceptibility of the metal ion xu. The susceptibility should be linearly related to the reciprocal of the absolute temperature (T) conforming to the Curie-Weiss law i.e. where the Weiss constant 8 can be determined as a finite intercept in a plot of l/xa against T. As will be seen below a departure of the measured magnetic susceptibilities from this simple dependence on temperature can frequently provide valuable clues concerning the origin of anomalous magnetic behaviour. lo B. N. Figgis and R. S. Nyholm J. Chem. SOC. 1959 338. l1 H. B. Gray R. Williams I. Bernal and E.Billig J. Amer. Chem. SOC. 1962 84 3596 l2 F. A. Cotton and R. H. Holm J. Amer. Chem. SOC. 1960 82 2979. 18 (a) A. Davison N. Edelstein R. H. Holm and A. H. Maki J. Amer. Chem. SOC. 1963 85 3049; (b) F. A. Cotton and R. C. Elder Inorg. Chem. 1965,4 1145. 462 Barefield Busch and Nelson It must be noted however that anomalous magnetic moments may arise in the solid state where non-equivalent lattice sites are occupied by the metal atoms and that the Curie-Weiss law will still be obeyed in such cases. This too can provide valuable information about the source of the anomaly. Anomalous magnetic moments for NiU compounds have been recognised since 1942 when French Magee and Sheffield14 observed that bis(formy1- camphorethylenediamine)nickel(rI) diamagnetic in the solid state shows a moment of 1-9 B.M.in methanol solution. For iron(@ such phenomena have been known since 1931 when Cambi and his co-workers first published work on the iron(m)-NN-dialkyldithio~arbamates,~~-~~ which show room-temperature moments between 2.3 and 5-9 B.M. The early investigators attributed the values to a thermal equilibrium between two magnetically isomeric forms. That the phenomenon is by no means rare is demonstrated by the large number of publications on the subject in recent years. A variety of models is required to account for the many examples of anomalous behaviour. These may be con- veniently considered under the following categories (1) Equilibrium between two spin states; (2) magnetically non-equivalent sites in the unit cell; (3) solute- solvent interaction ; (4) solute-solute interaction; (5) configurational equi- librium.1 Equilibrium between Two Spin States A. Nickel(II).-In D4h symmetry nickel may undergo a singlet-triplet spin state isomerism. This model has been invoked several times to explain unusual magnetic behaviour but it has been convincingly demonstrated only recently. The theoretical basis for the understanding of this type of behaviour was laid by Makilg for a weak ligand field and by Liehr and Ballhausen20 for a strong ligand field. The application of a ligand field causes splitting of the five degenerate d levels of a transition-metal ion into two or more separate sets of levels depending on the symmetry of the field. For fields of Oh symmetry two such levels e and t2, are produced. These exist as an upper doublet (e.g.the dx2-fl and d, orbitals) and a lower triplet (t2, the d,, dZz and dYz orbitals) as shown in Figure 6. For weak ligand fields the spin multiplicity will be the maximum obtainable for any given number of d electrons in accordance with Hund's rule. For strong fields on the other hand for which the energy separation do of the e and tzg sets of orbitals exceeds the electron-pairing energy (P.E.) the more stable configuration for d4 d5 d6 and d7 systems will be one in which the t2 level is completely filled before any occupation of the higher-energy e level occurs. Thus Fe(NH,);+ exhibits a magnetic moment of 5.45 B.M. (high-spin) 14 H. S. French M. Z. Magee and E. Sheffield J. Amer. Chem. SOC. 1942 64 1924. 16 L. Cambi L. Szego and A. Cagnasso Atti Accad. naz. Lincei 1932 15 266.l7 L. Cambi and L. Szego Ber. 1933 66 656. la L. Cambi and L. Malatesta Ber. 1937 70 2067. 1s (a) G. Maki J . Chem. Phys. 1958 28 651 ; (b) G. Maki ibid 1958 29 162 and 1129. Zo A. D. Liehr and C. J. Ballhausen J . Amer. Chem. SOC. 1959 81 538. L. Cambi and L. Szego Ber. 1931 64,2591. 463 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments d-or Free ion octahedral weak tetrugotial strong tetragonut (square planar) Figure 6 Splitting of orbitals in octahedral tetragonal and square planar ligand fields indicating the presence of four unpaired electrons whereas Fe(CN),"- is dia- magnetic (low-spin). For NiII complexes of Oh symmetry there is clearly only one way in which the eight d electrons can be arranged for states of lowest energy. All regular octahedral complexes of NiIT must therefore show a paramagnetism corresponding to the presence of two unpaired spins.This situation is alterna- tively represented by the energy level diagram in Figure 1 where it can be seen that the triplet 3A state (paramagnetic) must always lie below the singlet lE state (the latter state would impose diamagnetism) for all values of do. The situation changes however whenever an axial distortion is applied to the originally regular octahedral field. Such distortion lowers the symmetry from Oh to Ddi and is accompanied by a further loss of degeneracy of the d orbitals When this tetragonal distortion is large the energy separation between the dX*-,,% and d, orbitals (A,) or between dx2-yZ and dzy orbitals (d,) may exceed the electron-pairing energy. In this case there will be a change in magnetic moment from ca.3.0 B.M. to zero. Conceptually tetragonal distortion can be envisaged as arising by moving the two ligands on the z axis to greater distances from the metal ion. In fact this occurs only for certain electronic configurations and is an example of Jahn-Teller distortion. Alternatively and more subject to control the presence at these axial positions of different ligands which exercise a different perturbation on the d orbitals from those in the xy plane will generate the effect. In high-spin NiIIcomplexes only the latter type of tetragonal distortion is found as in mixed complexes of the type Nia,b,. Figure 6 represents the changes in orbital energy levels as the symmetry is progressively lowered from that of a regular octahedron to a strong tetragonal structure.It is clear that tetragonal distortion leads in the limit to square planarity and a co-ordination number of four. Proceeding in the reverse direction for a square planar complex Nia, which interacts along the z axis with two ligands b there will be a reduction in the energy separation (dl or d,) of the two highest-energy orbitals. For weak interactions such as would be caused by ligands low in the spectrochemical series e.g. I- d may still exceed the pairing energy. On the other hand if the separation of the ligands a and b in the spectrochemical series is relatively small 464 Barefield Busch and Nelson then d (or dl) may be less than the pairing energy and a change in spin multi- plicity will occur. Even in this simple model it is necessary to remain aware of the dependence of the pairing energy on the nature of the donor atoms for at a fixed d value reduction in magnitude of the pairing energy may produce the low-spin state.There are many examples of series of tetragonal Nin complexes wherein a relatively small change in the nature of the axial ligands brings about the pre- dicted magnetic cross-over. For example Goodgame and Venanzi21 found that for a series of NN’diethylethylenediamhe complexes of the type Ni(diamine),X (X is a univalent anion) the complexes fall in two categories diamagnetic and paramagnetic ( p = 3.1-3.3 B.M.) and the cross-over in ground states occurs between X = Br (diamagnetic) and X = CI (p = 3.29 B.M.). The sharpness of the change-over is demonstrated by the fact that chloride and bromide occupy adjacent positions in the spectrochemical series.Similar results have been observed for complexes of the planar quadri-dentate ligand (I).22 Here also the magnetic moment varies with the nature of the axially co- ordinated univalent groups. Thus the perchlorate and iodide are diamagnetic while the (anhydrous) bromide (in CHCl solution) chloride aide and thio- cyanate are paramagnetic ( p = 3.1-3.2 B.M.). Some unusual features of the dibromo-complex will be discussed in Section (2). An equally sharp change is seen for complexes of 1,4,8,11 -tetra-azacyclotetradecane (cyclam) (11)23 where the chloride and bromide complexes are paramagnetic and the iodide is dia- magnetic. Quinoxaline (In) complexes of Ninhalides provide a third example of a rather different kind.24 The complex with nickel iodide has the stoicheiometry NiQ,I (Q is quinoxaline); it is diamagnetic.The chloride and bromide of stoicheiometry NiQCl and NiQBr, are paramagnetic tetragonal polymers involving bridging halide and quinoxaline groups. In all of these systems the room-temperature magnetic behaviour is consistent with expectation for either 21 D. M. L. Goodgame and L. M. Venanzi J. Chem. SOC. 1963 616 5909. 22 G. R. Brubaker and D. H. Busch Inorg. Chem. 1966 5 2114. a3 B. Bosnich M. L. Tobe and G. A. Webb Inorg. Chem. 1965,4 1109. a4 A. B. P. Lever J. Inorg. Nuclear Chem. 1965,27 149. 465 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments the triplet or singlet ground state; i.e. no anomalous behaviour was observed. Certain systems are explicable on the basis of equilibria between coexisting singlet and triplet states whose energy states differ by amounts comparable with thermal energies.For such an equilibrium between spin isomers the distribution between the two states can be described by Maxwell-Boltzmann statistics if it is assumed that there is no change in structure other than the alteration in spin state. It has been shown25 that the magnetic susceptibility is given by 2g2N' [ 1 + 3 exp - kT X u = + Na where g is the Lande' splitting factor N is Avogadro's number ,8 is the Bohr magneton k is the Boltwnann constant T is the absolute temperature AE is the energy separation between the singlet and triplet states and Na is the tem- perature-independent paramagnetism (TIP). The factor Q arises because a small increase in entropy (Rln3) is associated with the increase in spin multiplicity.Other sources of entropy change are ignored and this amounts to the assumption of a vertical transition between states as represented in a single Orgel diagram. It follows from this equation that the magnetic susceptibility will increase with temperature except where d < < kT i.e. when the singlet and triplet states are essentially degenerate. This is of course the opposite direction of change for XA from that predicted by the Curie-Weiss law. The first convincing example of spin-state isomerism in a crystalline com- pound was reported in 1964 by Holt Bouchard and Carlin.26 Dichlorotetrakis- (NN-diethylthiourea)nickel(n) is spin-paired at temperatures lower than ca. 194"~ but attains partial paramagnetism reversibly as the temperature is raised.The corresponding bromide and iodide are diamagnetic at room tem- perature but the bromide becomes slightly paramagnetic at 373 OK. The relative abundances of the triplet states among the three halides is that predicted by the Maki19 and Ballhausen and Liehr20 treatment of Nin in D4h symmetry since the separation between the singlet and triplet states is least for the chloride which lies closest in ligand field strength to the substituted thiourea. The authors calculated AH = 4 kcal. mole-l and AS = 10 e.u. for the chloride. It may be noted that the gain in entropy is substantially larger than the value of 2-15 e.u. (R In 3) predicted for a vertical transition (change in spin multiplicity alone). This suggests not unreasonably that the magnetic cross-over is accompanied by other perhaps minor structure changes Le.the transition is not strictly vertical. This is common in fact to all reported examples of this phenomenon. Among the obvious changes that almost certainly must accompany the spin- state alteration is the decrease in in-plane metal-donor distance when the triplet changes into the singlet. Other Nin compounds showing spin-state isomerism in the solid state are found among complexes with planar quadridentate macrocyclic ligands. A 25 B. N. Figgis and R. L. Martin J. Chem. SOC. 1956 3837. 26 S . L. Holt R. J. Bouchard and R. L. Carlin J. Amer. Chem. SOC. 1964 86 519. 466 Barefield Busch and Nelson series of six-co-ordinate tetragonal complexes Ni(TAAB)X { TAAB is tetra- benzo[b f j n]l 5 9 13-tetra-azacyclohexadecane (IV)) have been examined where X is a singly charged anion.,' Table 2 summarises the data obtained for these systems.The chloro- and bromo-complexes exhibit anomalous room- temperature moments and susceptibility studies as a function of temperature give the non-Curie-Weiss plots shown in Figure 7. Rather than assuming the Table 2 Magnetic data for Ni(TAAB),+ complexes Complex Ni(TAAB)(C104) Ni(TAAB)(BF,) Ni(TAAB)(BPh,) Ni(TAAB)(NCS) N i(T A AB)(NO J Ni(TAAB)12 Ni(TAAB)C12,H20 Ni (TAAB)Br ,,H ,O peiia (B.M.) Diamagnet icb Diamagnetic * Diamagnetic 3.2 3-2 3.2 1.7 1.5 a Room temperature vaIues. b There is a small residual paramagnetism. susceptibility equation for a vertical transition Busch and his co-workers obtained equilibrium constants from and the variation of log Kes with the reciprocal of temperature; Figure 8 was used to derive enthalpy and entropy changes for the equilibrium.AH Values of 800 and 700 cal. mole-l and AS values of +037 and -0.45 e.u. for the chloride and bromide were obtained. If the transition were vertical (and the simple equation properly applied) the AS value obtained would have agreed 27 (a) G. A. Melson and D. H. Busch J . Amer. Chem. Suc. 1964,86,4830; (b) D. H. Busch Adv. Chem. Ser. 1967 62 616; (c) E. B. Fleischer personal communication. 467 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments I IOC *- ex 60C 0 soc I 1 5 8 1 0 3 1 ~ Figure 7 Temperature-dependence of the magnetic susceptibility of 0 Ni(TAAB)CI ,,H 0 Ni(TAAB)Br,,H,O and with the theoretical value of 2-15 e.u. It is noteworthy that for both TAAB complexes dS is closer to the predicted value of 2.15 e.u.than that observed for the first example cited.26 This might well be associated with for example smaller structural changes on passing from the singlet to the triplet state in the case of complexes of the cyclic ligands than in those of the unidentate substituted thiourea. %Ray crystallographic data do indeed indicate an increase in metal- donor bond distance for Ni(TAAB)2+ complexes on going from singlet to triplet ~OI-I-II.~~~ The nickel-nitrogen distance in Ni(TAAB)(BF,) is 1.90 A while that for [Ni(TAAB)(H,O)]I is 2.09 A. The role of lattice water in these systems has been discussed.27b This constitutes the first case where specific chemical interactions were blamed for this kind of spin-state change. Noting that the chloride and bromide form surprisingly stable hydrates they assumed that this hydrate water takes part in the equilibrium.The residual charge on the anion should be greatest for the singlet state since the metal-halide bond is weaker in that case. Consequently it was suggested that the water may hydrogen bond strongly to the singlet chloride but not to the triplet chloride thereby contributing specifically to the energy balance of the poised system. Karn2* has shown that an equilibrium between singlet and triplet spin states exists for the NiI complex of a macrocyclic ligand (V). This complex has a room-temperature moment of 2.79 B.M. A temperature-dependent 28 J. Karn Ph.D. Thesis Ohio State University 1966. 468 Barefield Busch and Nelson 1.1 sc Q 0 . E 0. F I O ~ / T ( O ~ ) .Figure 8 Variation of log Keel with 1/T fur 0 Ni(TAAB)Cl a,H ,O and Ni(TAAB)Br ,,H ,O susceptibility study from 100" to 340°K indicated Curie-Weiss dependence at low temperatures but increasing deviation from linearity at higher temperatures. This was attributed to a spin-state equilibrium between a triplet ground state and a Zow-lying singlet excited state. This is the first example for which the triplet lies lowest. Analytical treatment of the data yielded values of 1140 cal./mole and -0.68 e.u. for AH and AS respectively. This complex represents one of the most extreme cases of tetragonal distortion documented with the apparent ligand field splitting parameter of the axial iodides (Dqz = 216 cm.-l) being only about one-fifth that of the in-plane ligands (DqZy = 1124 cm.-l). Additional examples of such equilibrium behaviour have been reported recently for other Nin complexes.Ni(VPP),(NO,) [VPP is P(Ph),CH=CHP- (PH),] has been shown29 to have non-Curie-Wiess behaviour in the temperature range 300-421"~ with magnetic moments ranging from 1.67-3.14 B.M. This so H. N. Ramaswamy H. B. Jonassen and A. M. Aguiar Inorg. Chim. Ada 1967,1 141. 469 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments complex acts as a 2:l electrolyte in nitromethane solution and gives spectra typical of square planar nickel. The magnetic behaviour was explained on the assumption that the nitrate produces an axial perturbation in the solid state. A perchlorate salt of the same complex cation is reported to have a room- temperature moment of 1-50 B.M. but to show Curie-Wiess behaviour over an unspecified temperature range.Room-temperature magnetic measurements on a sample of the perchlorate complex prepared in these laboratories give a molar susceptibility of -491 x e.s.u. which yields a value of 052 B.M. after correction for the diamagnetism (+611.9 x e.s.u.). An interesting series of NiII complexes with benzimidazole has recently been reported.30 Several of these compounds show unusual magnetic behaviour and one of them having the stoicheiometry NiL,C12,S2 (L is benzimidazole and S is acetone of solvation) is reported to show spin-state equilibrium behaviour. Although a large number of 5-co-ordinate nickel@) complexes of both high- and low-spin type have been characterised only one compound has been found to exist with high- and low-sph forms in equilibrium.Nelson and Kelly have recently found31 that the 5-co-ordinate complex Ni(PNP)CI, where PNP is the terdendate ligand (VI) has a moment of 3.14 B.M. at 2 9 4 " ~ which falls reversibly to 1-30 B.M. at 9 9 " ~ . No change in the 2 9 9 " ~ X-ray diffraction pattern except that caused by anisotropic thermal expansion coefficients was observed on cooling to 77"~. Application of the theory for a Maxwell-Boltzmann distribu- tion between singlet and triplet states yielded a temperature-dependent energy separation from which it appears that the ground state is a singlet below about 2 5 0 " ~ and a triplet above this temperature. Additional examples of such pheno- mena may appear when extensive magnetic studies are performed on 5-co- ordinate systems. While these appear to be the only examples of an equilibrium between NiU complexes of the same symmetry in singlet and triplet states in the crystalline phase the phenomenon may well be more common in solution where solvent molecules provide the axial perturbation as was suggested by MALIs In this case it is difficult to distinguish between a spin-state equilibrium and one involv- ing some combination of 4-co-ordinate 540-ordinate (monosolvated) and/or 6-co-ordinate (disolvated) nickel@) species.B. Cobalt(n).-As previously described the normal magnetic moments of 6-co-ordinate Con complexes should fall near the extremes of the range from ca. 1-8 B.M. to ca. 5.2 B.M. One survey of Co* chemistry32 notes that there are moments falling throughout this entire range. The energy-level diagram in Figure 3 shows that at some value of Dy the 2E state crosses the 4TI state and in this region a thermal distribution between spin states could exist.This statement is oversimplified in view of spin-orbit interaction and the non-crossing rule. 30 D. M. L. Goodgame M. Goodgame and M. J. Meeks J. Chem. SOC. (A) 1967 (a) 1125 (b) 1676. s1 S. M. Nelson and W. S. J. Kelly Chem. Comm. 1968,436. 32 D. H. Busch 'Cobalt' ed. R. S. Young ch. 6 ACS Monograph series Reinhold New York 1960. 470 Barefeld Busch and Nelson This matter is discussed further in following examples. In 1961 Stoufer Busch and Hadley showed 33a that a previously postulatedsb equilibrium between a doublet and quartet state does indeed exist for the complex bis-(2,6-pyridin- dialdihydrazone)cobalt(n) iodide ([Co(PdAdH),]I,) which has a moment of 3-04 B.M.34 at 300"~.A study of the temperature-dependence of the magnetic susceptibility by these workers showed that the moment for the complex varies from 1.9 B.M. at 80"~ to 3.7 B.M. at 337"~. A simple thermodynamic treatment ignoring spin- orbit coupling gave AH and AS values of -2-14 kcal./mole and -0-44 e.u. respectively for the equilibrium spin-free + spin-paired However the magnitude of h should require its use in interpretation of the data.8 Other examples were not authenticated until 1966 when Stoufer et al?6 reported temperature-dependence studies on the magnetic properties of six Con complexes (including additional work on the compound previously studied33a) which exhibit anomalous magnetic behaviour. Table 3 gives observed values Table 3 Magnetic data for some magnetically anomalous cobalt (n) complexes Complex* Magnetic moment (B.M.) 100°K 300 O K [Co(BMI)jj I,,H,O 1-91 2.91 [Co(PMI)d (BFd 2.1 6 4.31 [ Co(terpy)d Br2,H,036 - 1.9 - 2.7 [ Co(PdAdH),] I 2.22 2-85 [CO(PBI)dI 2-38 3.72 [ Co(DTPH)] (c104)2 2.36 1 -90 * BMI Biacetylbis (methylimine) ; PMI pyridine-2-aldehyde methylimine; terpy 2,2',2"- terpyridyl; PdAdH pyridine-2,6-dialdehyde dihydrazone; PBI pyridine-2,6-dialdehyde bis(benzy1imine) ; DTPH 1 12-bis-(2-pyridyl)-lY2 1 1 12-tetra-aza-5,8-dithia-Az~10-dodecadiene.for peff at 100" and 300°K for the six complexes. These compounds a11 contain- ing ligands with unsaturated nitrogen donors give non-linear Curie-Weiss plots (see Figure 9) and can best be explained qualitatively at present by considering the anomalous moments to arise from a Boltzmann distribution between the *TI and 2E states in the vicinity of the cross-over point.For the terpyridyl complexes a marked dependence of the observed susceptibility on the counter- anion has been Additional studies on this phenomenon have been reported by Judge and Baker.% Results obtained by these workers are shown in 39(u) R. C. Stoufer D. H. Busch and W. B. Hadley J. Amer. Chem. SOC. 1961 83 3732; (b) R. C. Stoufer Thesis Ohio State University Colombus Ohio 1958; P. E. Figgins and D. H. Busch J. Amer. Chem. SOC. 1960 82 820. 34 R. C. Stoufer and D. H. Busch J. Amer. Chem. SOC. 1956 78 6016. 85 R. C. Stoufer D. W. Smith E. A. Clevenger and T. E. Norris Znorg. Chem. 1966,5 1167. 36 Data from R. Hogg and R. G. Wilkins J. Chem. SOC. 1962 341. 37 J. G. Schmidt W.S. Brey jun. and R. C. Stoufer Inorg. Chem. 1967 6 268. 38 J. S. Judge and W. A. Baker jun. Znorg. Chim. Actu 1962 1 68. 471 Iron Cobalt and Nickel Complexes having AnOmaIOUS Magnetic Moments loop 100 I 200 I L - I T (OK) Figure 9 Temperature-dependence of the magnetic susceptibility of some cobalt(x1) complexes Table 4. There is a pronounced dependence of magnetic moment on the asso- ciated anion and the degree of hydration. This is attributed to low-symmetry fields resulting from crystal-packing effects. Enthalpy calculation^^^ on linear portions of the susceptibility curves (Figure 9) indicate that the 2E(2G') state is the one of lower energy suggesting that the ligand field is not strong enough to cause complete spin pairing. Further studies3' by the electron paramagnetic resonance technique support the enthalpy data for a lower-lying ,E state at low temperatures and a decreasing population of this state at higher temperatures.It was found however that at still higher temperatures the *T state becomes lowest for the susceptibility continues to increase beyond the value expected for full degeneracy between the doublet and quartet states. More subtle effects such as mixing of the state with the 4T state under the influence of spin-orbit coupling in the vicinity of the cross-over point have been treated by Stoufer and his co-worker~.~~ Using the equation for the susceptibility x = NC1[W12/kT - 2 W,] exp - ( Wo/kT) Cexp - (WO/IET) and deriving experimental curves for the dependence of x on T they obtained quite good agreement with the experiment.Here Wis a function of two variables; [ the one electron spin-orbit coupling constant and 6 the separation between D. L. Williams D. A. Smith and R. C . Stoufer Inorg. Chem. 1967,6,590. 472 Table 4 Magnetic moments of Co(terpy),2+ with various anions as a function of temperaturea Compound Temp( OK) Eflective magnetic moments 15 3.75 4.17 1.89 2.14 - - - - 2.07 2.57 1 -87 - 20 3-78 4.20 1-91 2.14 1 -94 2.07 2-58 1.87 2.1 2 - - - CFrom Ref. 38 Corrected for TIP and ligand diamagnetism. 30 60 100 150 200 250 300 350 3.81 3.83 3-84 3.87 3.94 4.04 4-17 4.31 4.30 4.37 4.46 4.57 4-61 4.62 4.65 4.70 - - - 2-26 2.41 2.66 2-94 3.24 2-15 2.16 2.31 2-84 3.38 3-74 3.97 4.11 3.3 - - - - - - - 1.94 1.96 2.03 2.23 2.63 3.05 3.40 - - 2.15 - 1.94 - - I - 3.56 3.71 3.86 4.00 - - - - 2.07 2-08 2.25 2-89 3.62 4.17 4.49 4.66 2.61 2.72 2.82 3.13 3.43 3.75 4.01 4.25 1-87 1.87 1.88 1-89 2.04 2-42 2.96 3.52 2.27 2-43 2.58 2.73 2-90 3.06 3.21 3.53 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments the original 2E and 4T states.In addition 8 depends explicitly on T being of the form 8 = so + aT + bT2 + cT3. Although agreement appears on the surface quite good this could be coincidental because as Judge and Baker point any treatment probably should include low-symmetry ligand field effects. Calculations similar to those of Williams et al.39 have been reported by Barra- ~lough.~O These are the only examples among Corl compounds which have been relatively well characterised with respect to this type of anomalous magnetic behaviour. C. Iron.-An extensive series of compounds having the general formula Fephen,X has been reported where X covers a large part of the spectro- chemical series.When X is C1 Br I N, OCN HCO, and CH,C02 high-spin compounds are f ~ r m e d ~ l - ~ ~ while CN NCO and NO give low-spin diamagnetic corn pound^.^^-^^ The high-spin derivatives have reported room-temperature moments in the range of 5.1-5.3 B.M. which is lower than that predicted for spin-free d6 octahedral systems considering spin-orbit c o ~ p l i n g . ~ ~ ~ This lower- ing of the moment is attributed in part to the non-cubic field arising from the difference in ligand field strength of o-phen (o-phenanthroline) and the anion X. Temperature-dependence studies have been completed on a series of such compounds where X is I N, SCN and SeCN.48 The magnetic behaviour for the CI Br I and N at least approximate to what is expected for octahedral complexes although the moments are never as high as 5.6 B.M.and there is a slight decrease at lower temperatures rather than the increase expected due to spin-orbit coupling.8 The SCN and SeCN complexes on the other hand show moments of ca. 5.0 B.M. at room temperature and drop sharply to ca. 1.5 B.M. at 110"~ and 1 5 0 " ~ respectively. Similar behaviour was also noted for Fe(bipy),(NCS),. From a consideration of infrared data these workers formu- lated a dimeric structure with bridging thiocynates and considered the anomalous magnetic behaviour to result from co-operative interactions between the two metal atoms. It was recognised that the behaviour is somewhat different from that previously reported for dimeric systems.49 It might be noted that there have been more recent reports on such co-operative interaction^,^^ though again these are explicitly omitted from this Review.Other investigators concerned with the Fe(~hen),(NcS)~ and Fe(phen),- (NCSe) systems51 have suggested that the anomalous behaviour is better 40 C. G. Barraclough Trans. Faraday SOC. 1966 62 1033. 41 F. Basolo and F. P. Dwyer J. Amer. Chem. SOC. 1954,76 1454. 4a K. Madeja and E. Konig J . Inorg. Nuclear Chem. 1963 25 377. 48 W. A. Baker jun. and H. M. Bobonich Inorg. Chem. 1963,2 1071. 44 A. A. Shilt J. Amer. Chem. SOC. 1960 82 3000. 4s W. Beck and E. Schuierer Chem. Ber. 1962,95 3048. 46 K. Madeja Chem. Zvesti 1965 19 186. 4' J. S. Griffith Trans. Faraday SOC. 1958,54. 1109. 4* W. A. Baker jun.and H. M. Bobonich Inorg. Chem. 1964,3 1184. 4O A. Earnshaw and J. Lewis J. Chem. SOC. 1961 396. so J. Lewis F. E. Mabbs and A. Richards J. Chem. SOC. (A) 1967 1014; A. V. Khedekar J. Lewis F. E. Mabbs and H. Weigold ibid p. 1561. 61 E. Konig and K. Madeja Inorg. Chem. 1967 6 48. 474 Barefield Busch and Nelson represented as an equilibrium between 5T2 and lA ground states. High-spin forms would be obtained when the ligand field strength (A) is less than the pair- ing energy (P.E.) for the iron@) d6 ions. When A > P.E. compounds of the low-spin form result. If d and P.E. are of approximately the same magnitude then such an equilibrium could result between tzQ6 (spin-paired) and t2s4eg2 (spin-free). It is well that d for spin-paired forms is greater than that for spin-free forms and as Ewald et al.53 have shown for such equi-energetic configurations d (high spin) < P.E.< d (low spin). The temperature-dependence of the susceptibility for Fe(phen),(NCS) and Fe(phen)2(NCSe)2 determined by Konig and Madeja is shown in Figure 10. Figure 10 Temperature-dependence of the magnetic susceptibility for 0 Fe(phen) ,(NCS) and o Fe(phen) ANCSe) The general behaviour is similar to that reported previouslf8 and it was noted that individual preparations gave somewhat different values in certain tempera- ture ranges but the shapes of the curves are the same. Such transitions do take place reversibly with temperature. When data were incomplete it was thought that acceptable agreement between measured and calculated values of peff would be found:* but the data from the complete study could not be fitted well to the traditional Van Vleck equation.55 52 C.K. Jnrrgenson ‘Absorption Spectra and Chemical Bonding in Complexes’ Pergamon Press Oxford 1962 p. 128. 5s A. H. Ewald R. L. Martin 1. G. Ross and A. H. White Proc. Roy. SOC. 1964 A 280,235. 54 E. Konig and K. Madeja Chem. Comm. 1966 61. 55 J. H. Van Vleck ‘The Theory of Electric and Magnetic Susceptibilities’ Oxford Univ. Press 1932. 475 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments Further work including attempts to apply the treatment of Ewald et ~ 1 . ~ account being taken of the variation of the metal-ligand bond distance with population of a particular ground state did not improve the fit. Careful inves- tigations of the electronic Mossbauer and i.r. spectra indicated a change in molecular structure between 300" and 7 7 " ~ and one may assume that this must occur at the point where there is a sharp reduction in magnetic moment.These systems may qualitatively exhibit a lA1-V2 equilibrium which is associated with a modification in molecular structure such that the usual theoretical equations do not specifically apply. One possibility could be a crystalline phase change between two solid solutions (low-spin dissolved in high-spin + high-spin dis- solved in low-spin) at the point where there is an abrupt change in susceptibility. However X-ray powder data appear to contradict this. Recent magnetic studies on Fe(dipy),(NCS),S6 have revealed a behaviour quite similar to that for Fe(phen),(NCS) and Fe(phen),(NCSe) as noted previously.48 Fe(dipy),(NCS) has been found56 to exist in three polymorphic forms exhibiting related but distinctly different magnetic behaviour.One of these is found to have a slight field-dependent susceptibility. In all cases at some critical tempera- ture there is a more or less sudden decrease in susceptibility and concomitant change in ground state for the complex. This change is found to be essentially complete for one of the polymorphs but some molecules remain in the quintet state in each of the other two forms. Again such a sharp change in ground state is not characteristic of a bonafide equilibrium between two spin states but rather must indicate a significant variation in structural parameters. Nonetheless again no crystalline phase change is indicated. Depending upon the substituents on the rings some iron(@ poly-( 1 -pyrazolyl) borates (VII) have also been found to undergo a 1A1-6T2 spin-state equilibrium.670 Where R = H and X and Y are CH the complex is fully paramagnetic while when R = Ph and X = Y = H the compound is diamagnetic.However when R = X = Y = H or when R = X = H and Y is isopropyl intermediate room- temperature moments (ca. 2.7 B.M.) are found. An extensive study of the tem- perature-dependence of the susceptibility in chloroform by the n.m.r. technique57b indicated that the intermediate moments might arise from an equilibrium between lA1 and 5A1 states (the 5A1 in the DSd point-group is derived by descent in symmetry from the V1 state in Oh symmetry). The structure of the complex precludes a conformational equilibrium such as has been observed for some systems (see p.491). The positions of the resonances in the n.m.r. spectrum for the complex undergoing the spin equilibrium were intermediate between those of the diamagnetic and paramagnetic compounds and exhibit a non-Curie- Weiss dependence on temperature. A range in moments from 1.5 B.M. at 2 5 0 " ~ to ca. 3.5 B.M. at 340"~ was observed. AH and A S values of 3.85 kcal./ mole and 11.4 e.u. were derived by analysis of the results. Since the entropy in going from a l A to 5A1 spin state is again in excess of that expected for the change E. KGnig K. Madeja and K. J. Watson J . Amer. Chem. SOC. 1968 90 1146. 57 (a) J. P. Jesson S. Trofimenko and D. R. Eaton J. Amer. Chem. SOC. 1967 89 3158; (b) D. R. Eaton W. D. Phillips and D. J. Caldwell ibid. 1963 85 397 476 Barefield Biisch and Nelson P in spin degeneracy alone there must be a large contribution due to ligand field alterations bond-distance changes and/or solvation effects.The same general phenomenon is evident in the solid state but the equilibrium is shifted to favour the diamagnetic form. It is noteworthy that this system has allowed the first direct observation by Mossbauer spectroscopy of the cross-over between spin- states5* giving distinct spectra for the quintet state at higher temperatures and for the singlet state at lower temperatures. Recently the unusual magnetic properties of some iron(@ complexes of 2-(2-pyridylamino)-4-(2-pyridyl) thiazole (yapth) VIII) have been in~estjgated.~~ Bis-complexes of this ligand have been found to be spin-free for Con and NiI1; however for FG there is a strong dependence on the anion and in some cases the degree of hydration.Table 5 gives some room-temperature magnetic data including 6 values for the normal compounds. A number of interesting features were noted in this system. In addition to a large variation in magnetic behaviour with associated anion and degree of hydration the anhydrous nitrate requires about 3 hr. to attain an equilibrium or constant moment at low temperatures. Also there appear to be spin isomers of the [Fe(papth),]Cl, these being yellow (normal high-spin behaviour) and red- brown (predominantly spin-paired). In all cases a sample which has a high moment at room temperature appears yellow and as the temperature is decreased it becomes red-brown. No difference in the reflectance spectrum is evident however between yellow and red-brown samples so these workers attribute the vastly different appearance to differences in band intensities.58 J. P. Jesson and J. F. Weihner J. Chem. Phys. 1967 46 1995. 5s R. N. Sylva and H. A. Goodwin Austral. J. Chem. 1967 20 479. 477 Iron Cobalt and Nickel Complexes hnving Anomalous Magnetic Moments Table 5 Magnetic data for compounds of the type [Fe(papth)AX,,nH,O Anion C2042- 21- 2ClO 4- 2C1- (yellow) so,2- SO?- PtC1,2- 2c1- 2C1- (red-brown) 2SCN2- 2N0,- 2Br- 2N0,- No. of water molecules 1 1 1 0 3.5 5 2 0 2 1 1 2 0 5.28 - 5.0 - 5.1 - 5.4 - 5-0 N 3.4 N 1.9 N 1.8 N 3.1 4-48 N 2-4 1.3 5.18 - 16 - 20 - 38 - 5 -11 Quantitative treatment of these systems via a Boltzmann distribution over spin-states is not adequate but an empirical treatment in which relative concen- trations of the two spin forms were considered was used to obtain AH values for the various compounds studied.With the exception of [Fe(papth),]Br2,2H20 the plots of log K against Tare not straight lines over the whole temperature range but may show continuous curvature or have segments which approximate to straight lines of different slopes. Such deviations indicate something besides a vertical spin-state equilibrium and may be attributed in part to a modification of the crystal lattice. Such an interpretation indicates that the time-dependence found for the anhydrous nitrate is the consequence of a slow phase change at low temperatures. This suggests that the two anhydrous forms of the chloride are in fact lattice isomers first (allotropes) and spin isomers as a consequence of the structure modification.There are additional possibilities of solid solutions and perhaps some less obvious intermolecular interactions. Determinations of the susceptibilities at different field strengths might be helpful for further elucidation of the anomalies. Tris-(2-aminornethylpyridine)iron(11) is reported to undergo a spin-equili- briums0 with the associated anion (Cl- Br- I-) determining the specific behaviour observed. Moments show behaviour similar to those for Fe(phen),- (NCS),48s51 in dropping to a constant value at some intermediate temperature and remaining there to 2 0 " ~ . No success was achieved in fitting the data to an equilibrium between lAl and 5T2 states. The existence of two spin states is sup- ported by Mossbauer data for the Cl and I salts.Dilution experiments of the C1- salt with the corresponding zinc complex supports the conclusion that inter- molecular factors do not determine the magnetic behaviour. 60 G. A. Renovitch and W. A. Baker jun. J. Amer. Chem. SOC. 1967 89,6377. 478 BareFeld Busch and Nelson At present these constitute the documented cases of spin-state equilibria among iron@) complexes. There is an interesting compound phthalocyan- ine iron(@ which has been the topic of continued discussion since Klemm's first work showed a moment about midway between high-spin (S = 2) and low-spin (S = 0). The possibility of a spin-state equilibrium has been considered.61 Because of this possibility a more complete temperature study was performed on this compound along with other phthalocyanines.62a This work agreed well with that of Klemm (room-temperature moment = 3.85 B.M.) and covered a greater temperature range (1OO-340"~).Approximate Curie-Weiss depend- ence was found with some slight curvature at low temperatures. This observa- tion fails to support a spin-state equilibrium and the small Weiss constant of 9 OK casts doubt on the possibility of antiferromagnetism. Lever62a considered the moment to arise from the S = 1 spin configuration which is probably b22 - e * alg giving a 3Eg term (for S = 2 the configuration would be b2g2 * eg2 al big). No electron spin resonance signal has been observed for the complex .63 A triplet ground state (S = 1) has also been claimed recently for iron(@ in a six-co-ordinate environment. Konig and Madeja62b report moments of 3.98 B.M. and 3-80 B.M.respectively at 292"~ for Fe(phen),ox,SH,O and Fe(phen) mal,7H20 (ox = oxalate and ma1 = malonate). Curie-Weiss behaviour with small 8 is observed for both compounds and the moments are unchanged on dissolution in methanol. 57Fe Mossbauer and i.r. spectra are not consistent with the occurrence either of 5T2 or of lA states but show a greater similarity to those expected for the latter. No doubt there will be other cases of spin-equilibrium for iron(@ complexes when more extensive studies have been carried out. Workers should beware of the pitfalls of ferromagnetism and antiferromagnetism which can exist par- ticularly with iron(@. Also 'magnetic purity' is a matter of considerable concern for analytically acceptable samples can sometimes yield reproducible values whose magnitudes depend on the purification procedure.D. Iron(IrI).-The prime examples of spin-state equilibria involving iron(II1) occur with the NN-dialkyldithiocarbamates. These compounds first investi- gated by Cambi and his co-workers in the late 1930'~,~~-l* have been extensively studied more r e ~ e n t l y . ~ ~ ~ * ~ ~ The later workers have presented detailed studies of the magnetic behaviour of 18 iron@)-dialkyldithiocarbamates over the tem- perature range 8 0 4 0 0 " ~ . There exist three types of behaviour depending on L. Orgel Proc. 10th Solvay Conference Brussels 1956 ed. R. Stoops Brussels 1956. (a) A. B. P. Lever J. Chem. Suc. 1965 1821 ; (b) E. Konig and K. Madeja J. Amer. Chem. SOC. 1966 88 4528. D. J. E. Ingram and J. E. Bennett J . Chem. Phys. 1954 22 1136 Discuss. Faraday Suc.1955 19 140. 1 3 ~ A. H. White R. Roper E. Kokot H. Waterman and R. L. Martin Austral. J. Chem. 1964 17 294. 6s R. M. Golding W. C. Tennant C. R. Kanekar R. L. Martin and A. H. White J. Chem. Phys. 1966 45 2688. 479 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments the alkyl substituent. Room-temperature moments range from high-spin ( p e n = 5.83 R = pyrrolidyl) to a minimum value of peff = 3.02 (R = isobutyl). Almost any intermediate value can be achieved by proper choice of alkyl group or combination thereof. It is well established that the compounds show the same general behaviour in solution and that they are monomeric non-conducting molecular species. Pressure-dependent magnetic studies show a tendency toward the spin-paired state with increased pressure.This is as expected for it is estab- lished that d(spin-paired) should be greater than d(~pin-free)?~ and this can best arise by a decrease in the metal-donor bond distance giving in turn a smaller molal volume. It should be noted however that the change in molal volume with spin state is very small at constant pressure. These workers have produced one of the most complete studies so far on spin equilibria treating the data by application of Van Vleck's equation p2 = $g2 + 105 exp[-(1 + E/6)] + 8r1[1 - exp(-$x)] 1 + 2 exp(-$x) + 3 exp[-(l + E/@] x where x = 8/kT E is the separation between zero-point levels of the two states undergoing the equilibrium and g is a parameter to account for interactions of the 2T2 state with higher terms. The variation of magnetic susceptibility with temperature is explained qualitatively quite well by this equation which predicts with proper choice of constants a maximum and minimum similar to that observed in the experimental temperature studies.However no exact fit could be obtained. Three reasons were suggested to help explain this inadequacy (I) Failure of parameter g to account for interactions with higher states; (2) reduction of effective symmetry from Oh; and (3) differences in vibrational frequencies of the OAl and 2T2 states. Of these three ideas only the third was thought to be profitable for further treatment. These workers were able to assign metal-donor stretching modes for the 2T2 and 6Al states and by choosing probable values for the S-Fe-S bending modes they were able to modify the above equation to account for free-energy changes and thus alter the population levels at the measurement temperature.After such a modification a satisfying agreement could be obtained. The treatment still does not give exact molecular parameters. This would require precise values for the populations of the energy levels. An n.m.r. study of this equilibrium in solution was recently reported and supports the previous conclusions.65 Some interesting comments on the mean pairing energy for iron(In) long a topic of discussion were presented by these workers.s3 Ho and Livingstone66 have recently reported magnetic behaviour best ex- plained by a thermal equilibrium between t2$ and f2g3eg2 spin states for iron(n1) complexes of monothio-/3-diketones. For complexes of the type Fe[RIC(S)- CHC(0)R2] the magnetic behaviour is dependent on the nature of the R 66 R.K. Y. Ho and S. E. Livingstone Chem. Comm. 1968,217. 480 Barefield Busch and Nelson groups and electron-withdrawing substituents appear to increase the population of the spin-paired configuration. One of the most interesting cases of spin equilibrium occurs among the ferri- haemoprotein hydroxides. Pauling and his co-workers measured the magnetic moment of ferrihaemoglobin hydroxide in 1937 and obtained a value of 4.7 B.M.67 This differs considerably from that of the F- (5.92 B.M.) and from those of the CN- and SH- 2.5 and 2.26 B.M. derivatives respectively. Since that time there has been a continuing discussion of the electronic structures of such complexes. The first workers postulated a configuration with 3 unpaired electrons corresponding to an intermediate spin state between high-spin (S = z) and low-spin (S = i) configurations.Theoretical treatment of this electronic state shows it to be unstable in a strict octahedral donor environment.68a This does not mean that it is not likely in species of lower symmetry but Griffith has concluded S = $ states are highly improbable for these compounds68b whereas high-spin and low-spin configurations are definitely established. Spectra of a series of ferrihaemoproteins are for the most part identical for a particular axial ligand ; however the hydroxide shows band shifts throughout the same series.gg These workers have used such spectral data in connection with magnetic data to support the postulated existence of an equilibrium between high- and low-spin forms.In fact through sensitive experimental technique- they were able to measure and reduce spectral data into component contribus t ions from high-spin and low-spin ferrimyglobin hydroxide. This treatment is supported by the results of indirect susceptibility determinations over a rather narrow temperature range (1 "-30"c) which reveals a non-Curie-Weiss depend- ence as expected for a thermal equilibrium between spin states. From peff values of 5.92 and 2-24 for high-spin and low-spin forms respectively the percentages of the configurations for several haemoproteins were calculated to be as shown in Table 6. Table 6 High-spin (%) Low-spin (%) Myglobin 70 30 Haemoglobin 50 50 Peroxidase 7 93 Such equilibria have been postulated for other derivatives where the magnetic moment differs appreciably from ideal high- or low-spin values but no other detailed study has been reported.70 13' C.D. Coryell F. Stitt and L. Pauling J. Amer. Chem. SOC. 1937 59 633. 68 (a) J. S. Griffith J. Inorg. Nuclear Chem. 1956 2 1 and 229. (b) J. S. Griffith Proc. Roy. SOC. 1956 A 235 23. P. George J. Beetlestone and J. S. Griffith 'Haematin Enzymes' Pergamon Press New York 1961 p. 111. 'O (a) W. Scheler G. Schoffa and F. Jung Biochem. Z. 1957,329,232; (b) R. Havemann and W. Haberditzl Z. phys. Chem. 1958 209 135. 48 1 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments This completes the discussion of spin-state equilibria wherein a single sym- metry characterises both low- and high-spin forms. Many of these systems leave something to be desired in the quantitative interpretation but at present the equations generally employed are not adequate to describe the many perturba- tions a system undergoes in a transition from one spin state to another.The phenomenon remains a molecular dimorphism of the most interesting kind. 2 Magnetically Non-equivaIent Sites in Unit Cell The coexistence of metal ions with two different stereochemistries within the unit cell is well documented in a number of cases. This is well known only for NiII and so far involves two general situations. The first circumstance requires the metal ion to exhibit the same co-ordination number and the same set of attached ligands but different stereochemistries. The second produces two different co-ordination numbers and hence different stereochemistries within the unit cell.Kilbourn Powell and Darbyshire'l have carried out an X-ray investigation of the green complex dibromobis(benzyldiphenylphosphine)nickel(rI) which has a moment of 2.7 B.M. (a red diamagnetic form is also known). They showed that there are three 4-co-ordinate nickel atoms in the unit cell one Ni [P(PhCH,.)Ph,],Br is square planar and two Ni [P(Ph-CH,.)Ph,],Br units are tetrahedral. When allowance was made for only two of the three nickel atoms being high-spin the moment of the tetrahedral nickel atoms was calcu- lated to be 3.3 B.M. in good agreement with the values generally observed for nickel in pseudo-tetrahedral environments. Powell suggested the term 'inter- allogon' for such compounds. An analogous explanation has been suggested to account for the low moment (2.78 at 20"c) of bis-(Z-methyl-l,2-diamino- propane)nickel(n) nitrate perchlorate on the basis of an incomplete X-ray investigation and temperature-dependent magnetic studies.72 An example of the second type of isomerism has been observed in the yellow form of the Lifschiltz compound bis-(meso-stilbenediamine)nickel(n) dichloro- acetate $C,H,OH,QH,O which has been extensively studied by magnetic spectral and X-ray measurement~.~ Here the room-temperature moment is 2-58 B.M.and the unit cell contains both 6-co-ordinate and planar 4-co-ordinate nickel atoms in the ratio 2:l. In the distorted octahedra the axial positions are oc- cupied by water molecules. Recalculation of the magnetic data on this basis leads to a value of 3-16 B.M. for each paramagnetic ion a value that is com- pletely consistent with high-spin six-co-ordinate nickel(@.Quinoxaline complexes of nickel@) have been mentioned previously. Two forms of Ni Q(Br) appear to exist.,* It has been suggested from spectral evidence that the brown isomer (peff = 3-51 B.M.) also contains two stereochemical components within the unit cell. In this case alternating tetrahedra and tetra- 71 B. T. Kilbourn H. M. Powell and J. A. C. Darbyshire Proc. Chem. SOC. 1963,207. 72 D. Hall and J. M. Waters Inorg. Chem. 1964 3 615. 73 W. C. E. Higginson S. C. Nyburg and J. S. Wood Inorg. Chem. 1964,3,463; S . C . Nyburg and J. S. Wood ibid. 1964 3 468. 482 Barejield Busch and Nelson gonally distorted octahedra are suggested to coexist in a polymeric structure involving bridging bromides. Since the nickel atoms are high-spin in both environments a resultant moment in excess of 2-83 B.M.is to be expected. It has been suggested on the basis of magnetic susceptibility measurements that Ni(PEX)Br (1) may belong to this class of ~ ~ m p ~ ~ n d ~ ~ specifically to the second type. Anhydrous Ni(PEX)Br has a moment of 1.57 B.M. at room temperature and the magnetic susceptibility obeys the Curie law with only a small Weiss constant. If one assumes three diamagnetic ions for each para- magnetic species a moment of 3.14 B.M. for the paramagnetic species is calcu- lated. This is supported by the fact that addition of mole of lattice water causes the complex to become fully spin-paired. It was suggested that the trans- formation involves hydrogen bonding of the water to the bromide and effective removal of that ion from the co-ordination sphere of the nickel@) ion.The magnetic behaviour with temperature for such compounds stands in strong contrast to that of materials of class 1 and to others to be considered. This emphasises the fact that the measurement of susceptibility over a wide range of temperature can be a very useful aid to an understanding of the origin of an anomaly as indicated previously. Examples of such effects among derivatives of cobalt and iron appear not to have been observed. 3 Solute-Solvent Interactions A large number of NiU complexes which are diamagnetic in the solid state become paramagnetic when dissolved without gross decomposition in a variety of solvents. In many cases the paramagnetism corresponds to a moment per nickel atom of less than 2.9 B.M. Similarly certain complexes of cobalt(@ appear to undergo changes from S = Q to S = $ upon dissolution.Such changes may come as a result of several possible types of interaction. The character of the solvent is a key factor in determining the type of interaction that occurs. In most common polar solvents a strong ligating atom such as nitrogen or oxygen is present and this may exercise a moderate d-orbital perturbation with a resultant change in the co-ordination number and spin state of a particular species. Such cases are well documented experimentally and are explained theoretically for NiII by the treatments of Makilg and Liehr and Ballhausen.20 Such arguments can logically be extended to cobalt(r1) or even iron@). Willis and Mel10n'~~ noted that several diamagnetic bis(salicylaldimine)nickel(rr) complexes are fully paramagnetic ( p = 3-0-3.2 B.M.) in pyridine solution a phenomenon which they attributed to the formation of 6-co-ordinate pyridine adducts.Attempts to isolate the adducts in the solid state failed; however Basolo and M a t o ~ s h ~ ~ ~ were subsequently able to obtain the dipyridine adducts in pure form. Other solid diamagnetic NiU complexes of similar type partial paramagnetism in pyridine. In these cases it was assumed that the origin 74 G. R. Brubaker Ph.D. Thesis Ohio State University 1965. 76 (a) J. B. Willis and D. P. Mellon J. Amer. Chem. Soc. 1947 69 1237; (b) F. Basolo and W. R. Matoush ibid. 1953 75 5663; (c) H. C. Clark and A. L. Odell J. Chern. SOC. 1955 3431 3435. 483 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments of the paramagnetism is the same and that it corresponds to the equilibrium Solvent + Planar (diamagnetic) + pseudo-octahedral (paramagnetic) Similar behaviour has been observed in water pyridine and acetonitrile for the nickel complex of the macrocyclic ligand CRH shown in Structure (5).276,28 Detailed studies on the equilibrium Ni(CRH)(H20),2+ + Ni(CRH)2+ + xH,O in water by spectroscopic and magnetic methods reveal AH and AS values of 4.5 kcal./mole and 16 e.u.respectively for a 3.02 x 10-2~ solution of complex. These values are concentration-dependent because of changes in solvent activity. Solvent activity may be particularly important in determining the position of the equilibrium and therefore the apparent magnetic moments of such systems. Notable examples are those of Ni(trien)2f 76 and bis-(meso-2,3-diaminobutane)- while others have been discussed.27b Jlzrrgensen found that addition of quantities of extraneous salts to aqueous solutions of Ni(trien)(H20),2+ caused a new band to appear in the visible spectrum with a decrease in intensity of those already present.This new band could be assigned to a transition of square planar nickel produced by loss of 2 water molecules from the complex nickel ion. Similar results are obtained with bis(mesa-2,3-diaminobutane)- diaquonickel(n) where the relative amounts of spin-free and spin-paired forms as a function of added salt concentrations have been estimated spectroscopic- ally.77 This loss of water is apparently a result of the reduced activity of the solvent water. Nelson and his co-workers have studied similar equilibria spectrophoto- metrically for some CoII complexes (CoL2X, where L is pyridine substituted pyridine or isoquinoline) only in this case the equilibrium proceeds from a tetrahedra1 form to the octahedral structure.78 The equilibrium involves two paramagnetic species with only slightly different magnetic moments; however careful measurements should indicate a change in moment with an increase in the amount of octahedral complex.Of greater interest and complexity are the anomalous moments frequently observed for solutions of nickel and cobalt complexes in so-called 'inert' solvents such as CHCl, CCI, CS, C,H, etc. and even in co-ordinating solvents when they do not act as donor species. The first work on such systems used nickel@) compounds. The model initially suggested for this behaviour was that of a diamagnetic planar-paramagnetic tetrahedral rearrangement.76,79~ao This model is consistent with the temperature-dependence of the moment in 76 C.K. J~rgenson Actu Chem. Scund. 1957,11 399. 77 D. L. Leussing J. Harris and P. Wood J. Phys. Chem. 1962 66 1544. '* H. C. A. King E. Kor6s and S. M. Nelson J. Chem. Soc. 1963,5449; 1964,4832; S. M. Nelson and T. M. Shepherd ibid. 1965 3284; J. de 0. Cabral H. C. A. King S. M. Nelson T. M. Shepherd and E. Koros ibid. (A) 1966 1348. 79 S. Fujii and M. Surnitani Sci. Reports Tokoku Univ. Ser. I 1953 37 49. 8o H. C. Clark and A. L. Odell J . Chem. SOC. 1956 520. 484 Barefield Busch and Nelson solutions of bis-(N-methylsalicylaldimine)nickel(n).77~7g Also in the case of solutions of bis(formylcamphor)ethylenediaminenickel(n) in methylbenzenes Basolo and M a t o ~ s h ' ~ ~ found no direct correlation between the paramagnetism and the base strength of the solvent.From this they also inferred that the forma- tion of tetrahedral species rather than solvation is responsible for the para- magnetism. In none of these systems however was the presence of tetrahedral species actually demonstrated. In fact it was not until some years later that tetrahedral nickel@) complexes were first conclusively characteri~ed.~l-~~ Indeed it was shown by Sacconi Paoletti and Del ReE5 that bis-(N-alkyl- salicylaldimine)nickel(n) complexes possess zero dipole moment in benzene solution and that tetrahedral molecules could therefore not be present in signi- ficant concentrations. The same conclusion was subsequently reached by Holmg6 on electronic spectral grounds because solutions of this class of compound show none of the spectral characteristics of NiII in tetrahedral or pseudo-tetrahedral symmetry.87 Sacconi and his c o - w o r k e r ~ ~ ~ ~ ~ ~ showed further that paramagnetism occurs in the molten state as well as in solution and they concluded that inter- action with solvent is not a necessary condition for the development of para- magnetism.It soon became clear that no single model could account satisfactorily for all of the observed behaviours even within the N-substituted salicylaldimine class of compound. However as a result of concentrated efforts particularly by Sacconi and Holm and their co-workers it appears that the two models con- sistent with most of the experimental evidence are (i) molecular association (solute-solute interaction) and (ii) planar-tetrahedral equilibria.Similar behaviour in both categories has been found for cobalt@) and this is discussed at an appropriate point. It should be noted that most of the work on Con has been guided by analogy with the Ni*I examples many times using the same ligand systems. 4 Solute-Solute Interaction A. Nickel(II).-An early indication that the anomalous moment of bis-(N- methylsalicyladimine)nickel(n) in solution might be due to molecular association was provided by the isolation of a buff-coloured paramagnetic ( p = 3-2 B.M.) solid form of the complex by heating the diamagnetic green form to 150- ZO0°.90,91 In contrast to the diamagnetic isomers the paramagnetic form was observed to have a very low solubility in organic solvents and a polymeric L.M. Venanzi 4th International Conference on Coordination Chemistry Rome 1957; J. Chem. SOC. 1958 719. 82 N. S. Gill and R. S. Nyholm J. Chem. SOC. 1959 3997. D. M. Gruen and R. L. McBeth J. Phys. Chem. 1959,63 393. A. D. Liehr and C. J. Ballhausen Ann. Phys. 1959 6 134. 86 L. Sacconi P. Paoletti and G. Del Re J. Amer. Chem. SOC. 1957 79 4062. a(1 R. H. Holm J. Amer. Chem. SOC. 1961 83,4683. I~'F. A. Cotton 0. D. Faut and D. M. L. Goodgame J. Amer. Chem. SOC. 1961 83 344. 88 L. Sacconi R. Cini and F. Maggio J. Amer. Chem. SOC. 1957 79 3933. 89 L. Sacconi R. Cini M. Ciampolini and F. Maggio J . Amer. Chem. SOC. 1960 82 3487. *l L. Sacconi P. Paoletti and R. Cini J. Inorg. Nuclear Chem. 1958 8 492. C. H. Harris S. L. Lenzer and R. L. Martin Austral.J. Chem. 1958 11 331. 485 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments structure was suggested.90 This suggested that six-co-ordination is attained by the stacking of the planar units in a manner such that the oxygen atoms of adjacent molecules interact axially with the metal ions as shown in structure (IX). Further examples of salicylaldimine complexes that are paramagnetic in the solid state were later found by Holm and by Clark and O'Brien. It is con- venient to consider the N-alkyl-substituted salicylaldimine complexes as a func- tion of the nature of the substituent on nitrogen. An excellent review on Schiff base and /I-ketoamine chemistry is now available.92 (i) N-n-AlkylsaZicylaldimine complexes. The N-methyl compound has been studied to the greatest extent.It exhibits moments of ca. 1.9-2.3 B.M. at 20" depending on the s0lvent.8~ The moments of higher members of the n-alkyl series (ethyl to n-pentyl) are lower. In all cases the dipole moments in benzene or dioxan are practically zero thus ruling out the earlier s ~ g g e s t i o n ~ ~ s ~ ~ s ~ ~ of a planar-tetrahedral conformational equilibrium. A re-examination of the elec- tronic spectra made possible after publication of the s p e ~ t r a ~ ~ - ~ * of proven tetrahedral nickel complexes conclusively confirms the absence of tetrahedral species in these s o l ~ t i o n s . ~ ~ ~ ~ ~ - ~ ~ Thirdly the alternative possibility of a thermally regulated population of singlet and triplet states seemed unsatisfactory in view of the observation by Holms6 that in order to make the magnetic data fit the equation for the susceptibility in such systems the energy separation d E must itself be dependent on temperature to an unrealistic degree.Positive evidence se R. H. Holm G. W. Everett jun. and A. Chakravorty Progr. Inorg. Chem. 1967,7 83. 93 H. C. Clark and R. J. O'Brien Cunud. J. Chem. 1961 39 1030. B5 J. Ferguson Spectrochim. Acta 1961 17 316. R. H. Holm and T. M. McKinney J. Arner. Chem. SOC. 1960,82 5506. 486 Barefield Busch and Nelson for solute association was provided by the observation of simultaneous increases with increasing concentration in magnetic moment n.m.r. contact shifts and molecular weight for the N-methyl complex in chloroform and benzene as well as by spectral data.86~95~96 Solute association in solutions of the higher N-n- alkyl-substituted analogues is apparently less,97 these being mainly monomeric and diamagnetic at ordinary temperatures.A test of the solute-solute interaction model for N-alkylsalicyl-aldimine complexes was provided by introducing steric barriers to solute association by means of suitable substituents (X) on the benzene ring. Holms6 showed that the presence of a substituent in the 3-position eliminates the paramagnetism in the N-methyl series whereas substituents in the 5-position are less effective. The former position is of course expected to be sterically the more critical (see structure 10). In recent solution studies on these n-alkyl-substituted complexes Ewald and Sinng8 found that the pressure-dependence was as expected. The associated paramagnetic form prevails at higher pressures and the magnetic moment varies accordingly.Solid complexes subjected to high pressures (as a suspension in a non-solvent) also show a small increase in susceptibility but this does not result in isolation of a solid exhibiting permanent polymerisation. (ii) N-AryZ-substituted saZicyZaZdimine complexes. Diverse behaviour has been found for N-aryl-substituted salicylaldimine complexes the controlling factor being the nature and position of the substituents X and Y in the aryl ringsQ9 (structure 11). A certain regularity in behaviour exists. With few exceptions complexes containing Y in the 2-position are diamagnetic both in the solid and in solution. They are also monomeric in solution. The diamagnetism is retained even at high temperatures. Models show that the presence of Y substituents in the 2-N-aryl position or of X substituents in the 3-position of the salicylaldimine ring demand that N-aryl ring be rotated out of the plane of the chelate rings thereby preventing the lateral association in much the same way (though more effectively) as do the N-(straight-chain alkyl) substituents other than methyl.One exception is the complex in which Y = tolyl at position 2 and X = iso- propyl at position 3. This complex is weakly paramagnetic. Compounds in which there are no substituents in the 2- and 3-positions of the N-aryl and salicyl- aldimine rings respectively all show solution paramagnetism to a greater or less degree and a satisfactory correlation between solution moment and degree of association as judged by molecular-weight or spectral data was found for most of these systems.Exceptions are the N-phenyl and N-p-fluorophenyl complexes which show a solution paramagnetism greater than can be accounted for solely on the assumption of solute association. An extension of Holm and Swaminathan’s studies on the N-arylsalicylaldimine complexes was made by Sacconi and Ciampoliiii.lOO Magnetic and spectral measurements in xylene or D6 W. Ludwig Hefv. Ciiim. A d a 1962 45 665. D7 H. C. Clark K. Macvicar and R. J. O’Brien Cunud. J. Chem. 1962,40,822. Ds A. H. Ewald and E. Sinn Znorg. Chem. 1967 6 40. 9g R. H. Holm and K. Swaminathan Inorg. Chem. 1962 1 599. loo L. Sacconi and M. Ciampolini J . Amer. Chem. Soc. 1963 85 1750. 487 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments bibenzyl solution were made at temperatures up to 200"~.By comparison with the spectra of NiU chelates of known tetrahedral geometry bands at 6700 and 11,100 cm.-l were attributed to the presence of tetrahedral species. These bands increased in intensity with rising temperature while other spectral bands assigned to the six-co-ordinate associated molecules decreased. Further in the lower temperature range peff decreases with an increase in temperature in accord with expectation for the association model whereas at higher temperatures small increases in peff become apparent. It was concluded therefore that three forms of the N-aryl complexes (Y = H 3-C1 4-C1 3-methyl 4-methyl 3,4-benzo) can coexist in solution Associated + Planar + Tetrahedral (paramagnetic) (diamagnetic) (paramagnetic) At room temperature the proportion of molecules in the associated form is greater than 80%.With an increase in temperature the percentage of the tetra- hedral form increases. Significantly the same shift in equilibrium should be achieved isothermally by dilution. An interesting application of n.m.r. spectroscopy to the problem of the N-3- tolylsalicylaldimine complexes in non-co-ordinating solvents has been noted by Phillips and his co-workers.lo1 Following studies on nickel(@ chelates of the aminotroponeimineates (to be described) they attributed the large proton shifts to delocalisation of unpaired spin density on to the salicylaldimine ring a result consistent with the presence of paramagnetic associated chelate molecules in solution. A discussion of other nickel@) complexes of salicylaldimine in which the N- substituent is an u-branched alkyl group is deferred to section (5).(iii) 18- Diketone complexes. Until relatively recently bis(acetylacetone)nickel(n) which has a normal magnetic moment of 3.07 B.M. at 20"c was considered to be an example of a planar nickel compound with a triplet ground state (see ref. 19). The compound can be sublimed and it has been shown to be monomericlo2 and planarlo3 in the vapour phase. However in 1961 Bullen Mason and Pauling showed conclusively from X-ray investigations that the structure is trimeric in which each nickel atom is surrounded by a distorted octahedron of oxygen atoms; the central octahedron shares a triangular face with each of the two terminal octahedra (XII).1049105 Fackler and Cottonlos demonstrated from electronic spectral studies that the trimeric unit persists in solution but becomes reversibly dissociated into the red monomer at high temperatures.Molecular weight measurements in dichloromethane also indicated association. lor E. A. LaLancette D. R. Eaton R. E. Benson and W. D. Phillips J. Amer. Chem. Soc. 1962 84 3968. loa F. Gach Monatsh. 1900 21 98. lo9 S. Shibata Bull. Chern. SOC. Japan 1957 30 753. lo* G. J. Bullen Nature 1956 177 537. loti G. J. Bullen R. Mason and P. Pauling Nature 1961 189 291. losF. A. Cotton and J. P. Fackler jun. J . Amer. Chem. SOC. 1961 83 2818. 488 Barefield Busch and Nelson (EL) (rn) Replacement of the methyl groups R1 of the acetylacetonate ion (XIII) by t-butyl hinders the close approach of the individual units to such a degree that the complex is diamagnetic and remains so in solution in non-co-ordinating solvents.107 While the bulky t-butyl groups completely prevent an increase in co-ordination number by means of molecular association six-co-ordination can still be achieved by the introduction of small ligands such as water alcohols or amines.Thus such complexes are blue-green and paramagnetic. When R1 = R2 = isopropyl the steric restriction to molecular association is reduced to a degree intermediate between R = methyl and R = t-butyl. The solid is green and paramagnetic but a significant proportion of monomeric species is observed at much lower temperatures than in the case of R1 = Me. Steric hindrance of a comparable degree is also present in the paramagnetic form of the chelate R1 = H R2 = Ph since the phenyl group cannot assume coplanarity with the chelate ring.lo8 B.CobaIt(u).-Sacconi demonstrated that bis-(N-methylsalicylaldiminato) cobalt(@ achieves 5-co-ordination by association in the solid state.log As a result this complex has a moment of 4.62 B.M. Dissolution in benzene gives a brown solution and spectral studies indicate the species is largely monomeric and tetrahedral at low concentrations. No moment is reported for the complex in solution. Gachlo2 also prepared the bisacetylacetonate complex of Con but its solid- lo’ J. P. Fackler and F. A. Cotton J. Amer. Chem. SOC. 1960 82 5005. lo* J. P. Fackler and F. A. Cotton J. Amer. Chem. SOC. 1961 83 3775. loS P. L. Orioli M. Di Vaira and L. Sacconi Chern. Cumm. 1965 103. 489 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments state structure remained in question until recently.Cotton and Holm first postulated it to be a high-spin square planar C& complex,12b having a moment of 4-8-49 B.M. in the solid state and in solution. However comparison of the spectrum with that of Co(dipivaloylmethane), which is known to be pseudo- tetrahedral,ll* indicated that this compound is also tetrahedral in the monomeric state. In addition a strong deviation from Beer’s law was observed in non-co- ordinating solvents and molecular-weight measurements showed a concentration- dependent association. Co(acac) can be isolated in several forms; Co(acac),2H20 [Co(acac) ,- H,O], [Co(acac),],,,,H,Q and [Co(acac),],. X-Ray studies on a twinned crystal show [Co(acac)] to have a structure somewhat like that of the nickel ~omp1ex.l~~ Cotton and Elder suggest considering the tetramer as two diastereo- meric fragments joined along a common octahedral edge (XIV).X-Ray studieslll of [Co(acac),,H,O] show it to be the centre portion of the tetrameric species with bridge bonds on the two faces occupied by water molecules (XV) so regardless of the stoicheiometry Co(acac) achieves 6-co-ordination and high- spin character. It is possible for association to lead to a low-spin form for Con. Recently studies on Co(PEt,),(NCS) have shown that the magnetic moment and electronic and i.r. spectra depend strongly on temperature and concentration in dichloro- methane and dichloroethane s o l ~ t i o n s . ~ ~ ~ J ~ ~ This has been interpreted in terms of a monomeric (high-spin) tetrahedral f dimeric (low-spin) 5-co-ordinate equilibrium where the 5-co-ordinate dimer uses bridging thiocyanates.Further substantiation of such an equilibrium was obtained through the study of reac- tions of tertiary phosphines with Co(PR,),X complexes to obtain 5-co-ordinate species. The reaction of Co(PEt&,(NCS) with excess of PEt causes a decrease in the observed moment from 3.4 to 2.05 B.M. Although no solid was isolated cryoscopic measurement indicated Co(PEt,),(NCS) to be present. 5 Configurational Equilibrium An equilibrium between diamagnetic planar molecules and paramagnetic tetrahedral molecules was the mechanism first postulated and later disproved to explain the intermediate magnetic moments of bis-(N-methylsalicyla1dimine)- nickel@) in non-co-ordinating solvents. The model was subsequently invoked to account for some of the properties of solutions of N-aryl-substituted salicyl- aldimineslOO (see section 4).It is important to recognise that planar and tetra- hedral molecules have the same co-ordination number and that whereas a planar configuration may be stabilised by crystal lattice forces a lower-energy state in solution may sometimes be achieved simply by an intramolecular rearrangement. A difficulty to the conclusive recognition of the presence of planar-tetrahedral equilibria for solutions of solid diamagnetic complexes for 110 F. A. Cotton and R. H. Soderberg J. Amer. Chem. SOC. 1962 84 872. ll1 F. A. Cotton and R. C. Elder Znorg. Chem. 1966 5,423. lla M. Nicolini C. Pecile and A. Turco J. Amer. Chem. SOC. 1965 87 2379; Co-ord. Chem. Rev. 1966 1 133. 113 T. Boschi M. Nicolini and A.Turco Co-ord. Chem. Rev. 1966,1,269. 490 Barefeld Busch and Nelson many years was the fact that the stereochemical species had not been character- ised s p e ~ t r a l l y . 8 ~ ~ ~ ~ ~ ~ After about 1959 the isolation114 in some cases of planar and tetrahedral isomeric forms for example bis(benzyldipheny1phosphine)- nickel(@ halides provided a useful guide in recognizing planar-tetrahedral equilibria in benzene solutions of these and complexes of related type which exhibit moments between 0 and cu. 2.9 B.M?l”ll6 A particularly thorough study of solutions of nickel(@ chelates of the amino- troponeimineates has been carried out by Phillips and his c o - ~ o r k e r s . 1 ~ ~ - ~ ~ ~ For these the magnetic susceptibility and the electronic and n.m.r. spectra were observed to be dependent on temperature solvent and the nature of the N- substituent (XVI).Very large proton contact shifts in deuteriochloroform were interpreted as being due to the presence of electron spin density on the seven- membered hydrocarbon rings; the magnitude of the shift depends on the pro- portion of tetrahedral molecules in solution. The contact shifts which were measured relative to the free ligand or to the corresponding diamagnetic zinc chelate permitted the evaluation of equilibrium constants for the planar- tetrahedral equilibrium. In all cases d G decreased with temperature (200-350”~) and the AH values which varied from cu. 1 to 6 kcal. rnole-l were calculated from the temperature-dependence. Derived entropy changes were positive falling in the range 8-17 cal. deg.-l mole-l.(It has been pointed out by H o r r ~ c k s l ~ ~ that some of the thermodynamic values calculated for this system and for some others are based on an incorrect form of the equation for the isotropic shifts.) The magnetic susceptibility data (for CDCl solutions) were similarly treated in terms of a Maxwell-Boltzmann distribution of nickel ions between the two spin states. Thermodynamic functions were also evaluated from the intensities of electronic spectral bands at cu. 800 and 1200 mp attributed respectively to the diamagnetic and paramagnetic forms. The thermodynamic functions evalu- ated by the three different methods were in good agreement. That the para- magnetic species in solution are tetrahedral was established by evidence from several sources (i) the complexes are monomeric in solution; (ii) satisfactory assignment of the near-infrared spectra of almost completely diamagnetic and almost completely paramagnetic systems to transitions predicted for planar and tetrahedral symmetries respectively could be made; (iii) the magnetic sus- ceptibilities increase with an increase in temperature; and (iv) no correlation exists between the position of equilibrium and the expected co-ordinating power of the solvent.114 M. C. Browning J. P. Mellor D. J. Morgan S. A. J. Pratt L. E. Sutton and L. M. Venanzi J. Chem. SOC. 1962 693. C. R. C. Coussmaker M. H. Hutchinson J. P. Mellor L. E. Sutton and L. M. Venanzi J. Chem. SOC. 1961 2705. 116 R. G. Hayter and F. S. Humiec J. Amer. Chem. SOC. 1962 84 2004. 117 W. D. Phillips and R. E. Benson J. Chem. Phys. 1960 33 607.118 R. E. Benson D. R. Eaton A. D. Josey and W. D. Phillips J. Amer. Chem. SOC. 1961 83 3714. 119 D. R. Eaton A. D. Josey W. D. Phillips and R. E. Benson J. Chem. Phys. 1962,37,347. I 2 O D. R. Eaton W. D. Phillips and D. J. Caldwell. J. Amer. Chem. Soc. 1963 85 397. lZ1 W. D. Horrocks,jun. J. Amer. Chem. SOC. 1965 87 3779. 491 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments Recent pressure-dependent spectral and magnetic studiesQ8 indicate the planar form to be favoured at high pressures suggesting a smaller partial molar volume than for the tetrahedral form. This is consistent with what has been found for the F F dithiocarbamate~,~~ where bond length increases have been found upon the unpairing of spins. Further there was no indication of a transformation to a co-ordination number greater than 4 with increasing pressure.The variation in behaviour with change in the nature of the group R can be fairly satisfactorily accounted for on steric grounds. Clearly as the R groups become bulky steric repulsions between them are expected to become important in a planar structure. Rotation of the planes of the chelate rings toward a tetra- hedral configuration for the central nickel atom should relieve the steric strain. It is predicted therefore that the equilibrium should shift in favour of the tetrahedral form as the size of R is increased. This proved to be the case as shown in Table 7. The extensive diamagnetism of the N-aryl-substituted complexes was attributed to a rotation of the plane of the aryl ring out of the plane of the 7-membered ring.Table 7 Magnetic character of aminotroponeiminate complexes as a function of N-substituent N-substituent Magnetic character of chelate at 23 'c in CDCl R1 = R2 = H Diamagnetic R1 = H R2 = alkyl Diamagnetic R1 = R2 = CH Slightly paramagnetic (< 1 a R' = R2 = ~1 25-75 % paramagnetic R1 = R2 = ethyl Almost completely paramagnetic (> 99 %) The discussion of the magnetic behaviour of N-alkylsalicylaldimine complexes having secondary and tertiary alkyl groups has been deferred to this point because of their different behaviour. In contrast to solutions of the N-methyl- and most N-aryl-salicylaldimine complexes of nickel@) where the paramag- netism can be accounted for mainly in terms of molecular association the paramagnetism of a-branched N-alkyl-substituted complexes appears to be due predominantly to the presence of tetrahedral species.122 The effect of branching of the N-alkyl substituent is reflected in the solid state as well.Thus the n-propyl compounds for example are planar even when substituents are present in the 3-position of the salicylaldimine ring. The t-butyl chelates on the other hand are all pseudo-tetrahedral. Both stereochemistries are found for the solid s-alkyl complexes the choice of configuration in these cases depending on the nature and position of the ring substituent. In inert solvents at room temperature the n-propyl derivatives are predominantly planar the t-butyl derivatives largely pseudo-tetrahedral and the s-alkyl chelates coexist in comparable proportions of both forms. Assignment of stereochemistry in the solid state was made on the basis of magnetic moments and reflectance spectra and in solution on the bases of dipole-moment and molecular-weight evidence.Although there is evidence lte L. Sacconi M. Ciampolini and N. Nardi J. Amer. Chem. Soc. 1964 86 819. 492 Barefield Busch and Nelson for some association in solutions of the s-alkyl derivatives it has been shown that at or above 37"c they are essentially monomeric under the same conditions of solvent concentration and temperature where they were found to be appreciably paramagneti~.'~~ An indication of configuration is shown by the dipole moments in Table 8. With the exception of the t-butyl chelate an Table 8 Dipole moments for some (N-alkylsalicylaldimine)nickel(n) chelates in benzene at 25 'c R Prn Bun Pri Bus But P (0) 0.00 0.00 2-51 2-34 4-74 increase in temperature leads to a shift in the position of the plan=-tetrahedral equilibrium toward the tetrahedral form as evidenced by (a) an increase in solution magnetic moments and (b) an increase in intensity of the band near 6200 cm.-' attributable to the pseudo-tetrahedral species (a decrease was observed in the case of the t-butyl compound).122 Sacconi et calculated the percentages of the tetrahedral form for the n-propyl and isopropyl series from the electronic polarization susceptibility and electronic spectral data and good agreement of the derived thermodynamic functions by the three methods was obtained.The data in Table 9 show that for the straight-chain series the planar form is more stable than the tetrahedral form by about 2.8 kcal. mole-l. The relatively large enthalpy difference 46-52 kcal.mole-l means that at room temperature the concentration of tetrahedral molecules is negligible and not more than 10% at 170"c. The data also show that for the isopropyl series the tetrahedral form is comparatively strongly favoured over the pJanar. Steric Table 9 Thermodynamic quantities associated with the planar-tetrahedral equi- librium for N-alkylsalicylaldiminenickel(n) complexes as a function of the ring substituent R X H 5-Me 5-CI H 3-C1 5-Me 5-c1 AG (kcal. mole-l) 2.9 2.8 2.7 - 0.52 0-35 0.20 0.16 A H A S (kcal. mole-') (cal. mole-l deg.) 4.6 4 4.6 4 5.2 6 3.2 10 3-0 7 2.5 6 2.0 4 118 R. H. Holm and K. Swaminathan Inorg. Chem. 1963 2 181. 493 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments considerations seem to control the behaviour of the s- and t-alkyl derivatives.Intramolecular repulsion between the large t-butyl groups attached to the N-position and the chelate rings prevents a planar arrangement. Such repulsion is smaller in the s-alkyl compounds and still smaller in the n-allcyl chelates so that in the latter case only the planar configuration is present except at rather high temperatures. Thus the percentage of tetrahedral form increases in the order R = n-alkyl < s-awl < t-alkyl. In general substitution in the phenyl ring is less critical than is the nature of the N-R substituent (Table 9) except perhaps in the case of benzo-substitution. Benzo-substitution in 5,6-positions enhances the stability of the diamagnetic planar isomer whereas the benzo-group in the 3,4-ring position has the opposite effect.122 In the latter case severe steric interaction between the benzo-group and the N-substituent is predicted.While steric effects are generally dominant in deciding the position of the planar-tetrahedral equilibrium electronic factors also must play a significant role in some cases. For example for the chelates in which R = isopropyl and X = 3-C1 and 3-Br the planar form is more stable than expected on the basis of steric considerations alone.122 An extension of the investigation of N-s-alkylsalicylaldiminenickel(n) systems by n.m.r. spectroscopy has been carried out by Holm Chakravorty and Dudek.12* Their results provide general support for the interpretation of the properties of this class of salicylaldimine chelate in terms of planar-tetrahedral equilibria.As with the aminotroponeimineates discussed previously large iso- tropic proton hyperfine contact shifts (to both high- and low-field) relative to the resonance position of the free ligands were observed. For the protons at the 3- and 5-ring-carbon atoms the shifts are positive while those for the protons at the 4- and 6-positions are negative. The derived spin-density distribution indi- cates resonance among the valence-bond structures (17-21). These are the same R3 124 R. H. Holm A. Chakravorty and G. 0. Dudek J. Amer. Chem. SOC. 1964 86 379. 494 Barefield Busch and Nelson valence-bond structures predicted for a transfer of spin density on to the ring by d,,-p, bonding in the tetrahedral complex. Holm and Chakravorty have also studied a series of bis-(0-hydroxynaphthaldi- mine)nickel(n) complexes.125 These have the general form shown in XXIV).Table 10 lists their magnetic behaviour as a function of ring substituent Table 10 Magnetic properties of bis-(0-hydvoxynaphthaldimine)nickel(n) complexes Nap h thaldimine 22 22 23 23 24 24 24 R Pri But Pri Et,CH Prl But Et 2CH Solid Solutiona Diamag. 0.86 (30") Paramag. 3.23 (25 ") Paramag. 2-78 (20")-2.83 (50") Paramag. 2.93 (20")-2-96 (50") Diamag. Diamag. 1.80 (50") Paramag. 3.23 (25") a Solvent is chloroform. * Paramagnetic species have moments of 3.30 f 0.03 B.M. at 25". and R group both in the solid and in solution. The anomalous moments in solution have been found to arise from a planar-tetrahedral equilibrium. Spin- density calculations were performed and thermodynamic data were obtained by use of n.m.r.contact shift methods. The n.m.r. technique which appears to be a powerful one for the evaluation of the thermodynamic functions for diamagnetic-paramagnetic equilibria has recently been extended by Everett and Holm126 to bis-(/%ketoamino)nickel(n) complexes. These complexes depending on the R substituents (XXV) display greatly varied behaviour. Two groups of complexes having R2 = R3 = CH and R2 = CH,; R3 = C,H were studied and in each series a planar-tetrahedral equilibrium was found. For both series when R1 is s-alkyl the complexes are essentially tetrahedral at room temperature but when R1 is ary1 for R2 = R3 = CH or R1 is n-alkyl and R2 = CH, R3 = C,H5 then the complexes are largely planar. A small amount of association was found when R1 is aryl. More recently a planar-tetrahedral equilibrium has also been demonstrated for analogous bis-(p-ketoamino)cobalt(n) c0mp1exes.l~~ The types of compound studied and their solid and solution configurations are listed in Table 11.The formation of the tetrahedral cobalt species is endothermic as for the nickel com- pounds. A comparison of the data for the nickel and cobalt systems127 indicates that any substituents which promote tetrahedral behaviour for Nin cause Con to be fully tetrahedral and substituents on CoII inducing planar behaviour cause Nin to be fully planar. Following this idea Holm and his co-workers prepared 126 A. Chakravorty and R. H. Holm Inorg. Chem. 1964,3 1010. 1z8 G. W. Everett and R. H. Holm Proc. Chem. SOC. 1964 238; J. Amer. Chem. Soc. 1965 87 2117. 137 G. W. Everett jun.and R. H. Holm J. Amer. Chem. SOC. 1966,88,2442. 495 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments Table 11 Configurations of bis-(fi-ketoamino)cobalt(n) complexes in solid state and in Solution R' R2 Me,Prn,Ph Me H Me H Me Ha H The methine proton is R3 Configuration Solid Solution Me Ph Tetrahedral Tetrahedral Me Planar Tetrahedral and planar Ph Planar Tetrahedral and planar Me Planar Tetrahedral and planar replaced by a methyl group in this complex. (==I (=) bis(pyrrole-2-aldimino)-complexes of nickel(@ and cobalt(xr).128 These have the structure (XXVI) which is quite similar to the /3-keto-amine complexes. The only Con complex which could be prepared that with R = t-butyl was tetrahedral. The solution properties of the Nin series show a similar dependence on R to that of the salicylaldimine complexes in that no tetrahedral form is present if R = n-alkyl and no planar for R = t-butyl.Likewise when R is s-alkyl both configurations are found in solution at room temperature. It might be noted that when R = CH there appears to be some association but no para- magnetic species is found. The s-alkyl complexes have a moment per nickel atom less than 2 B.M. and all spectra show tetrahedral components. Molecular-weight measurements do not indicate association and spectral studies indicate an endo- thermic equilibrium of planar (S = 0) and tetrahedral (S = 1) forms. This equilibrium was demonstrated for R = isopropyl s-butyl 1 -ethylpropyl cyclo- hexyl and lY2-dimethylpropyl as R groups. At lower temperatures n.m.r. signals are broadened and dF plots are non-linear for these derivatives in a direction which indicates a less paramagnetic species than the planar-tetrahedral equilibrium should give at the same temperature.In addition for R = -CH(CHJCH20CH3 markedly different spectral and proton resonance patterns are observed and it has been suggested that there may be planar-tetrahedral-octahedral equilibria involved. This has been observed for Nin systems containing the ligand where R is CH,- or -H.129 Here ringclosure can occur with the two oxygen atoms completing an octahedral co-ordination sphere. In these cases the equilibria involved are between mono- meric octahedral planar and tetrahedral species and the magnetic moment varies accordingly. Monomeric octahedral-square planar equilibria also occur in solutions with la8 R.H. Holm A. Chakravorty and L. J. Theriot Znorg. Chem. 1966,5,625. laoA. Chakravorty J. P. Fennessey and R. H. Holm Znorg. Chem. 1965 4 26. 496 Bare$eld Busch and Nelson non-co-ordinating solvents of complexes of the related ligand series (XXVIII).130 This occurs only when R1 = H and R2 = C,H,- a-naphthyl or substituted phenyl. The nature of the X substituent determines the solid-state behaviour. In solution an increase in temperature shifts the equilibrium towards the 4-co- ordinate form giving a smaller moment. A measure of steric hindrance to ring- closure seems to be responsible for the coexistence of the two conformations in solution of this series of complexes. When the R1 and R2 substituents are made less bulky as when R1 = H R2 = n-alkyl or benzyl the complexes are mono- meric octahedral both in solid state and in solution.On the other hand when R1 = H and R2 = o-tolyl or a-naphthyl the complexes are all square planar and diamagnetic in the solid state. Further increase in steric hindrance as in R1 = phenyl and R2 = methyl or phenyl prevents the terminal (p) nitrogen atom from co-ordinating at all; i.e. the Schiff base acts as a bidentate ligand and the complexes are completely square-planar and diamagnetic in both the solid state and in solution. Sacconi and his co-worker~l~~ subsequently found that when R1 = R2 = ethyl and the ring substituents are 3-C1 5-C1 or 3,4-benzo that the solid complexes are paramagnetic ( p = ca. 3.3 B.M.) and 5-co-ordinate. While diamagnetic 5-co-ordinate nickel(@ complexes are now well known these are the first examples of this co-ordination number in the high-spin state.A distorted square pyramidal structure for the 5-chloro-derivative is indicated by X-ray studies.132 The nickel atom lies a little above the mean basal plane. Moments in benzene intermediate between 0 and ca. 3.3 B.M. were observed and electronic absorption spectra suggested the following equilibria:130 where one or both of the /%nitrogen atoms is detached to give a 5- or 4-co- ordinate species and the stereochemical form is determined by the free energy of the system at a particular temperature and in a particular solvent. Interestingly the diethyl derivative can be studied in the molten state where a similar de- pendence of the magnetic moment on temperature is found as in solution. The presence of greater amounts of planar form at higher temperatures indicates the endothermic nature of the change from pseudo-octahedral to the planar form.A d H value of 3.4 kcal. mole-l was calculated. More recently Sacconi and his co-workers have studied both types of equilibrium for the Schiff bases formed from salicylaldehydes and NN-substituted ethylenediamines by the n.m.r. technique.lS They were able to assign spin I3O L. Sacconi P. Nannelli and U. Campigli Inorg. Chem. 1965 4 818. 131 L. Sacconi P. Nannelli N. Nardi and M. Campigli Inorg. Chem. 1965 4,943. 132 L. Sacconi P. L. Orioli M. Di Vaira J. Amer. Chern. Soc. 1965 87 2059. lS3 J. D. Thwaites and L. Sacconi Inorg. Chem. 1966 5 1029; J. D. Thwaites I. Bertini and L. Sacconi Inorg. Chern. 1966,5 1036. 497 Iron Cobalt and Nickel Complexes having Anomalous Magnetic Moments densities to positions on the phenylene ring and to determine dG values for the equilibrium react ions.Most of the preceding discussion has dealt with systems containing quite similar ligand systems with nitrogen and/or oxygen donors. Recently it was reported that a planar-tetrahedral equilibrium was found for a Cco-ordinate system having phosphorus and halogen d0n0rs.l~~ When Ni[(C,H5)2P(CH~,P(C,H5)2~x2 (X is Br C1 or I and the ligand denoted dpp) is dissolved in methylene dichloride nitrobenzene or chloroform spectral studies indicated a band in the region 12,50~10,000 cm.-l attributable to tetrahedral species none of which is present in the solid state. Gouy measure- ments for Ni(dpp)Br2 in nitrobenzene (22") and Ni(dpp)I in nitrobenzene (23") and chloroform (24") gave moments of 2-12 2.62 and 2.73 B.M. respectively. The equilibrium was studied by the temperature-dependence of isotropic shifts in the n.m.r. in nitrobenzene and chloroform. Entropy and enthalpy contribu- tions favour the tetrahedral form of the iodide more than the bromide. Since Ni [(C,H,)2P(CHc32P(C,H~)2]X2 complexes are completely planar in solution it was suggested that steric factors control the equilibrium to a large extent. We thank the U.S. Public Health Service for support. G. R. Van Hecke and W. D. Horrocks jun. Inorg. Chem. 1966,5 1968. 498

ISSN:0009-2681    

DOI:10.1039/QR9682200457

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

Macromolecular Structure and Properties of Ribonucleic Acids By R. A. Cox NATIONAL INSTITUTE FOR MEDICAL RESEARCH MILL HILL LONDON N.W.7 1 Introduction The variety of roles that ribonucleic acids play in the living cell became apparent during the past decade. Living organisms develop and reproduce themselves with great fidelity and geneticists have long been able to deduce laws governing the inherited characteristics of the organism. However growth development and reproduction depend on directed chemical reactions which are brought about by specific catalysts called enzymes. Each enzyme is a unique protein molecule which is synthesised by the cell presumably with the aid of other enzymes. The individual polypeptide chains for a protein are polymers of 20 common amino-acids so that the number of possible amino-acid sequences is unlimited yet the living cell is able precisely to reproduce a particular sequence.l Biologists have known for many years that ribonucleic acid (RNA) is im- plicated in the biosynthesis of proteins.2 During the past decade methods of isolating it were so improved that studies of the biological chemical and macromolecular properties of RNA have been possible.At least 4 classes of RNA are now recognised; these are ‘messenger’ ribosomal transfer and viral RNA. Each class of RNA has unique properties which reflects its function in the cell. Studies of the biological activity of RNA have led to the elucidation of the mechanism of protein biosynthesis. Abbreviations,-The abbreviation RNA for ‘ribonucleic acids’ is used generically for polymers of nucleotides linked by diesterified phosphate bonds joining the C(3’) hydroxyl of one ribose residue with the C(5’) hydroxyl of the next ribose residue.Particular species of ribonucleic acid are referred to thus t-RNA transfer RNA; phenylalanyl t-RNA etc. is the ester formed between phenyl- alanine etc. and the t-RNA specific for phenylalanine etc. ; m-RNA messenger M. F. Perutz ‘Proteins and Nucleic Acids Structure and Function’ Elsevier Amsterdam 1962; J. N. Davidson ‘The Biochemistry of Nucleic Acids’ Methuen and Co. London 5th ed. 1965; J. D. Watson ‘Molecular Biology of the Gene’ Benjamin New York 1965; ‘The Living Cell’ Readings from Scientific American W. H. Freeman and Co. San Francisco 1965. 2 J. Brachet ‘The Biological Role of Ribonucleic Acids’ Elsevier Amsterdam 1960; H.Chantrenne ‘The Biosynthesis of Proteins’ Pergamon Press London 1961; H. R. V. Amstein Brit. Med. Bull. 19 21 217; P. N. Campbell and J. R. Sargent ‘Techniques in Protein Biosynthesis’ Academic Press London 1967; R. Schweet and R. Heintz Ann. Rev. Biochemistry 1966 35 723. 499 Macromolecular Structure and Properties of Ribonucleic Acids RNA; r-RNA ribosomal RNA; A C G and U refer respectively to adenine cystosine guanine and uracil residues; AMP CMP GMP UMP respectively are adenylic acid cytidylic acid guanylic acid and uridylic acid. Poly-A (poly- adenylic acid) etc. refers to a polymer of AMP etc. having the same inter- nucleotide linkages as RNA poly-AU etc. refers to a copolymer of AMP and UMP etc. poly(A 4- U) etc. refers to a complex formed between one strand of poly-A and one strand of poly-U etc.poly(A + 2U) etc. refers to a complex formed between one strand of poly-A and two strands of poly-U etc. 2 Biological Properties The genetic information which specifies the sequence in which the amino-acids must occur in order to produce a particular protein is stored within the chromo- some of the cell and is used as required. The chromosome consists of very long chains of DNA3 (deoxyribonucleic acid) in combination with protein. Each chromosome is made up of many genes and each gene has the information essential for the synthesis of a particular protein.* The DNA component of the chromosome is the repository of genetic information. DNA is synthesised from four principal deoxyribonu~leotides~ and the sequence in which these occur specifies the amino-acid sequence of the protein which is under the control of a gene.A sequence of three deoxyribonucleotides (a codon) is required to specify an amino-acid. The code words have now been identified and a genetic dic- tionary exists.6 Thus a linear sequence of the nucleotides of the gene is translated into the linear amino-acid sequence of the protein. Whereas genetic information is stored in the chromosome proteins are synthesised principally in the cytoplasm. The genetic message is carried from one site to another in the form of messenger RNA.’ The linear deoxyribo- nucleotide sequence of the gene is copied by means of an enzyme (DNA-depend- ent RNA polymerase) that synthesises an RNA molecule which has a nucleotide sequence that is complementary to one of the strands of the gene.a Messenger RNA is then transported to the ribosomes of the cytoplasm where polypeptide synthesis takes place.Ribosomes are ribonucleoprotein particless having a molecular weight of ( 3 4 ) X lo6 daltons and consisting of 30-50% protein. In general the ‘messenger’ interacts with more than one ribosome to form 3 H. J. F. Cairns J. Mol. Biol. 1963 6 208; C. I. Davern Proc. Nat. Acad. Sci. U.S.A. 1966 55 792. 4 F. Jacob and J. Monod J. Mol. Bid. 1961,3 318; N. R. Cohen Biological Reviews 1966 41 503. 6 p. A. Edwards and K. V. Shooter Quart. Rev. 1965,19,369. 6 ‘The Genetic Code’ Cold Spring Harbor Symposia 1966 voI. 31. 7 S. Brenner F. Jacob and M. Meselson Nature 1961 190 576; F. Gros H. Hiatt W. Gilbert C. G. Kurland R. W. Riseborough and J. D. Watson ibid. 1961 190 581; E. Volkin and L.Astrachan Virology 1956,2 149. 8 v. V. Koningsberger and L. Bosch ‘Regulation of Nucleic Acid and Protein Biosynthesis’ Elsevier Publishing Co. London 1967; J. Josse A. D. Kaiser and H. L. Kornberg J. Biol. Chem. 1961 236 864; S. B. Weiss and T. Nakamoto Proc. Nut. Acad. Sci. U.S.A. 1961 47 1400; M. Chamberlin and P. Berg ibid. 1962,48 81. 9 M. L. Petermann ‘The Physical and Chemical Properties of Ribosomes’ Elsevier Amsterdam 1964; H. R. V. Arnstein Ann. Reports 1963 60 512. cox polyribosomes.1° An adaptor molecule (transfer RNA) is required to translate the polynucleotide sequence into an amino-acid sequence.ll The adaptor is able to combine chemically with a specific amino-acid to form a amino-acid ester of transfer (t-)RNA and it is able to recognise the triplet of nucleotides (the codon) which correspond to its specific amino-acid.The principal steps in normal protein biosynthesis (see Figure 1) have been confirmed by Khorana and his colleagues12 who synthesised oligo-deoxyribonucleotides of known sequence and then showed that the amino-acid sequence of the polypeptide which could be synthesised in vitro was determined by the DNA. The genetic code is given in Table 1. The normal cycle of protein synthesis may be interrupted when the cell becomes infected by an RNA virus.13 The viral RNA behaves as messenger and takes over the protein synthetic apparatus of the cell for the production of viral protein (Figure 2). The viral RNA has also to replicate itself (by means of its own RNA-dependent RNA polymerase).l* The single strand of viral RNA is itself copied and a double-stranded intermediate is pr0d~ced.l~ Viral RNA may also exist in the virus as a double-helical structure.16 A nucleic acid must be isolated before its properties can be studied extensively.The isolation of a pure nucleic acid requires methods that do not damage primary or secondary structure and give preparations that are homogeneous with respect to molecular weight nucleotide sequence and biological activity. Before 1956 RNA was generally believed to have a molecular weight of about 15,000 daltons although preparations of about 300,000 daltons were occasionally obtained. The sensitivity of RNA to hydrolysis either by potent degradative enzymes (nucleases) or by alkali was not appreciated. The development of improved methods for the fractionation of subcellular c0mponents,1~ and the use of reagents such as phenolla for the dissociation of protein from RNA and 10 A.Gierer J. Mol. Biol. 1963 6 148; J. R. Warner A. Rich and C. E. Hall Science 1964 138 1399; J. R. Warner P. M. Knopf and A. Rich Proc. Nut. Acad. Sci. U.S.A. 1963 49 122; F. 0. Wettstein T. Staehelin and H. Noll Nature 1963 197 430. 11 (a) M. Hoagland ‘The Nucleic Acids’ E. Chargaff and J. N. Davidson Academic Press New York 1960 vol. 3 p. 349; (b) G. L. Brown and S. Lee Brit. med. Bull. 1965 21 236; G. L. Brown ‘Progress in Nucleic Acid Research’ ed. J. N. Davidson and W. E. Cohn Academic Press vol. 1 1 p. 259; G. L. Brown S. Lee and D. H. Metz ‘Genetic Elements Properties and Function’ ed. D. Shugar Academic Press New York 1967 p. 57. l a H Kossel A. R. Morgan and H. G.Khorana J. Mol. Biol. 1967,26,449; H. G. Khorana ‘Genetic Elements; Properties and Function’ ed. D. Shugar Academic Press 1967 p. 209. 1s ‘Basic Mechanisms in Animal Virus Biology’ Cold Spring Harbor Symposium L. Frisch Long Island Biol. ASSOC. New York vol. 27; E. M. Martin Brit. Med. Bull. 1967 23 192. 14 S. Spiegelman 4’””. Nat. Acad. Sci. U.S.A. 1965 54 919. 16 E. M. Martin Genetic Elements Properties and Function’ ed. D. Shugar Academic Press London 1967 p 117; L. Montagnier and F. K. Sanders Nature 1963 199 664. 16 P. J. Gomatos and I. Tamm Proc. Nat. Acad. Sci. U.S.A. 1963,49 707. 17 ‘Methods in Enzymology’ Biochem. SOC. Symposium ‘Methods of Separation of Subcellular Structural Compounds’ Cambridge Univ. Press 1963 p. 23. 18 K. S. Kirby ‘Progress in Nucleic Acid Research,’ ed.J. N. Davidson and W. E. Cohn Academic Press New York 1964 vol. 3 p. 17; H. R. V. Arnstein and R. A. Cox Brit. Med. Bull. 1966 22 158; G. L. Cantoni and D. R. Davies ‘Procedures in Nucleic Acid Research’ Harper and Row New York 1966; L. Grossman and K. Moldave ‘Methods in Enzymology’ Academic Press New York 1967 vol. 12. Macromolecular Structure and Properties of Ribonucleic Acids Phe-specific transfer -RNA Phenylalanine Growing Polypeptide chain (- NH2 terminus) A 30s - ri bosomol subparticle Figure 1 Mechanisms of protein biosynthesis The synthesis of a hypothetical polypeptide with the amino-acid sequence formylmethionine- alanine-valine-phenylalanine-phenylalanine-phenylalanine-phenylalanine - - - - by E. coli is illustrated. The next amino-acid to be incorporated is phenylalanine which reacts specifically with phenylalanine t-RNA forming phenylalanyl-t-RNA the reaction is brought about by activating enzymes.Phenylalanyl-t-RNA is bound to the larger subparticle (W. Gilbert J. Mol. Biol. 1963 6 360) but interacts with the codon UUU of m-RNA through the anti- codon AAA (H. Kaji and A. Kaji Proc. Nat. Acad. Sci. U.S.A. 1964,52,1541; M. Nirenberg and P. Leder Science 1964 145 1399). The NH2 group of the phenylalanyl residue of phenylalanyl-t-RNA forms a peptide bond with the COaH group of the last amino-acid of the growing polypeptide chain which is now attached to phenylalanyl-t-RNA. The genetic message is read from the 5' end to the 3' end (M. Salas M. A. Smith W. M. Stanley A. J. Wahba and S . Ochoa J. Biol. Chem. 1965,240 3988; R. Thach M. A.Cecere T. A. Sun- dararajan and P. Doty Proc. Nat. Acad. Sci. U.S.A. 1965 54 1167; H. Lamfrom C. S. McLaughlin and A. Sarabhai J. Mol. Biol. 1967 22 355). The codon AUG marks the start of the message (B. F. C. Clark and K. A. Marcker J. Mol. Biol. 1966,17,394) and the codons UAA and UAG may signal the release of the polypeptide chain (S. Brenner T. Stretton and S . Kaplan Nature 1965 206 994). When the ribosome reaches the 3' end of m-RNA it is likely that the subparticles fall apart. The smaller subparticle may then form a complex m-RNA (5' end) and formylmet-t-RNA. The larger subparticle then reacts with the complex to complete the ribosome (G. Mangiarotti and D. Schlessinger J. Mol. Biol. 1966 20 123; H. P. Ghosh and H. G. Khorana Proc. Nat. Acad. Sci. U.S.A. 1967,58 2455; R. 0.Kaempfer M. Meselson and H. J. Raskas J. Mol. Biol. 1968,31,277). The site for peptide bond synthesis may be in the interior of the-ribosome (L. i. Malkin and A. Rich J. Mol. biol. 1967 26 329) 502 cox Table 1 The genetic code Second letter U C A G U UUA UAA OCHRE UGA ? A UUG Leu UCG UAG AMBER UGG Tryp G UGC EE lphe ucu LJJ; 7 jSer uAu}Tyr UAC ccu CAU F.c EEE }Leu CCC ] Pro CAC lHis 2 CUA CCA cAA GluN CUG CCG CAG +a .a AUU ACU AAU A AUA ACA AUG Met ACG AAG CGU 7 U CGC CGA 2 2 CGG G ? 2 AGU )ser AGC u s AGA)Arg AGG 2 GUU GCU GAU GCU U GCC} GCA GAC GAA 1" GGC} GGA A C GUG GCG GAG GGG G V i p RNA \. \ \ Figure 2 Diagram showing the way in which infecting virus donates either DNA or RNA to the synthetic machinery of the host cell and thereby diverts the normal course of cell replication (T.S . Work Adv. Sci. 1965 22 98) 503 Macromolecular Structure and Properties of Ribonucleic Acids improvementslg in the fractionation of RNA have led to the isolation of bio- logically active RNA. 3 Primary Structure of RNA The oligonucleotides that were isolated before 1956 proved sufficient20 for the general features of the chemical structure of RNA to be elucidated. This is illus- trated in Figure 3 which shows an oligomer of the 4 common nucleotides viz. adenylic (AMP) cytidylic (CMP) guanylic (GMP) and uridylic acids (UMP). The sugar was shown to be D-ribose21 in the furanose form22 and the nucleosides SLOH END I 0=P-0 I 0 - 4 24 Adenine ( I CH,. I /0\ I 6 OH A-7 - I Figure 3 The chemical structure of ribonucleic acid (The older numbering system is used for pyrimidine derivatives in conformity with purines [cf.A. M. Michelson ‘The Chemistry of Nucleosides and Nucleotides’ Academic Press New York 19631) l9 R. E. Click and D. P. Hackett Biochim. Biophys. Acta 1966 129 74; J. R. B. Hasting J. H. Parish K. S. Kirby and E. S. Klucis Nature 1965,208,645; M. D. Dabeva and R. G. Tsanev Andy?. Biochem. 1966,17,390; D. H. L. Bishop J. R. Claybrook and S. Spiegelman J. Molec. Biol. 1967 26 373; U. E. Loening Biochem. J. 1967 102 251; R. L. Erikson J. Molec. Biol. 1966 18 372; J. F. Weiss and A. D. KeImers 1967 Biochemistry 1967 6 2507. 2o (a) P. A. Levene and L. W. Bass ‘Nucleic Acids’ American Chemical Society Monograph Series Chemical Catalog. Co. New York 1931; E. Chargaff and J. N. Davidson ‘The Nucleic Acids’ Academic Press New York 1955; D.0. Jordan ‘Chemistry of Nucleic Acids’ Butterworths London 1960; (b) A. M. Michelson ‘The Chemistry of Nucleosides and Nucleotides’ Academic Press New York 1963. 21 G. R. Barker K. R. Farrar and J. M. Gulland J. Chem. SOC. 1947 21 ; G. R. Barker and J. M. Gulland ibid. 1943 625; P. A. Levene and W. A. Jacobs Bet. 1909 42 2102; 2469;2412;2103. 2a A. R. Todd J. Chem. SOC. 1946 2 647; P A. Levene and R. S . Tipson J. Biol. Chem. 1932 94,809; 1932,97,491; 1933 101 529; 1934 105,419; 1934,106 113; 1937,121 131; B. Lythgoe and A. R. Todd J. Chem. SOC. 1944,2,592; D. M. Brown and B. Lythgoe ibid. 1950 3 1990. 504 cox formed by the reaction of a purine or pyrimidine base with D-ribose were established as ring-N glycoside~~~ involving N(3) of pyrimidines or N(9) of the purines. The glycoside linkage has the /3 ~onfiguration.~~ It has been suggested on the basis of studies of the optical rotatory dispersion of the nucleosides that both purine and pyrimidine base-residues adopt the anti-conf~rmation~~ as shown in Figure 3.The nucleotides are linked together through diesterified phosphate residues which are formed by reaction with the C(3’) hydroxyl group of one D-ribo- furanose moiety and the C(5’) hydroxyl group of another.26 The structure has been confirmed by ~ynthesis.2~ 4 Chemical Properties of RNA The internucleotide bond is readily hydrolysed by alkali because the presence of the C(2’) hydroxyl facilitates the formation of an intermediate 2’,3’-cyclic phos- phate during hydrolysis.28 This intermediate has been isolated.28 Whereas hydrolysis of the glycoside linkage between purines and deoxyribose is rapid in mildly acidic solutions (PH 3-5 or less) the corresponding bond between purines and ribose is much more stable.More drastic conditions e.g. treatment with more concentrated hydrochloric acid (PH 1-6) for 24 hr. at 37” liberate the pyrimidine as well as the purine bases from DNA but the resistance of the pyrimidine ribosides is such that it is very difficult to get quantitative yields of the free bases from RNA. The glycoside linkage is relatively stable between pH 3 and 7 but can be hydrolysed in more acidic solutions.29 The internucleotide bond is hydrolysed at an appreciable rate in neutral aqueous salt solutions at 60” or more.3o The hydrolysis is catalysed by ions such as magnesium,31a M P. A. Levene and R. S . Tipson J.Biol. Chem. 1934 104 385; S. Furberg Acta Chem. Scand. 1950 4 751; 1952 6 634; S. Furberg Acta Cryst. 1950 3 329; J. M. Gulland J. Chem. Soc. 1938 1722; G. A. Howard G. S. Kenner B. Lythgoe and A. R. Todd ibid. 1945 556; J. Davoll B. Lythgoe and A. R. Todd ibid. 1948 967; B. Lythgoe H. Smith and A. R. Todd ibid. 1946 355. 24 H. Davoll B. Lythgoe and A. R. Todd J. Chem. Soc. 1944,833; V. M. Clark A. R. Todd and J. Zussman ibid. 1951 2952. 25 T. R. Emerson R. J. Swan and T. L. V. Ulbricht Biochem. Biophys. Res. Comm. 1965. 19 643; T. L. V. Ulbricht T. R. Emerson and R. J. Swan Tetrahedron Letters 1966 1561. 26 P. A. Levene and H. S . Simms J. Biol. Chem. 1962 70 327. 27 R. Lohrmann D. ,5011 H. Hayatsu E. Ohtsuka and H. G. Khorana J. Amer. Chem. SOC. 1966 88 819. 28 A. Fon6 Arkiv. Kemi.Min. Geol. 1947,24 A No. 34 p. 1 ; D. M. Brown and A. R. Todd J. Chem. SOC. 1952 52; D. M. Brown D. I. Magrath and A. R. Todd ibid. p. 2708; R. Markham and J. D. Smith Biochem. J. 1952 52 552. 29 E. Visher and E. Chargaff J. Biol. Chem. 1948 176 715. 30 J. Eigner H. Boedtker and G. Michaels Biochim. Biophys. Acta 1961,51 165. (a) T. Lindahl J. Biol. Chem. 1967 242 1970; (6) J. J. Butzow and G. L. Eichorn Bio- polymers 1965 3 395; G. L. Eichhorn P. Clark and E. D. Becker Biochemistry 1966 5 245; T. Yamani and N. Davidson J. Amer. Chem. SOC. 1961 83 2599; (c) G. L. Eichhorn and J. J. Butzow Biopolymers 1965 3 79; ( d ) A. M. Michelson ‘The Chemistry of Nucleo- sides and Nucleotides’ Academic Press New York p. 340; (e) P. D. Boyer H. Lardy and K. Myrbtick ‘The Enzymes’ Academic Press New York 1961 vol.5; T. Uchida and F. Egami ‘Methods in Enzymology’ ed. L. Grossmann and K. Moldave Academic Press New York 1967 vol. 12A pp. 228 239; S. Linn ibid. p. 247; W. E. Razzell ibid. p. 255; K. K. Reddi ibid. p. 257; M. F. Singer and G. Tolbert Biochemistry 1965 4 1319; H. 0. Robertson R. E. Webster and N. Zinder J. Biol. Chem. 1968 243 82. 505 Macromolecular Structure and Properties of Ribonucleic Acids copper(@ manganese@) nickel@) and zinc;31b by lanthanum and other tervalent inner transition metals;31c by heavy metal hydroxides;31d and by a range of hydrolytic enzymes31 (nucleases). The chemical modification of the base-residues has been studied because of their importance in chemical mutagenesis e.g. mutants of tobacco mosaic virus have been produced through the action of nitrous acid on the RNA component,3a and because a modified base-residue provides a useful marker in piecing frag- ments of RNA together in order to ascertain the nucleotide ~equence.5~ The chemical reactivity of the base-residues is modified by conformation so that a knowledge of the rate with which a reagent reacts with a base-residue and of the sites which become modified have proved useful in studies of secondary struc- ture% as well as in correlating biological activity with primary and secondary structure (a study of ‘active sites’).35 Representative reactions of the nucleotides are given in Table 2.The sodium and potassium salts of polynucleotides are soluble in formamide and dimethyl sulphoxide as well as in 80% methanol and the cetyltrimethyl- ammonium salt is soluble in a wide range of organic solvents.36 Single-stranded polynucleotides but not double-helical complexes tend to precipitate at 0” from M sodium or potassium chloride solutions.The base-residues strongly labsorb ultraviolet light below 300 m ~ . ~ ’ The ordered oligo nucleotides and polynucleo- tides are optically active38 owing to their conformation. 5 The Principal Conformations of Ribonucleic Acids The macromolecular properties of polynucleotides are governed by the negatively charged electrostatic field due to the diesterified phosphate residues (PK - ca. 1.6) and by the ability of the base-residues to react with one another to form helical structures. The two forces tend to oppose one another so that conforma- tion depends on ionic strength temperature and also on pH (owing to the acid- base properties of the purine and pyrimidine residues).The conformation and physicochemical properties of model polynucleotides have been extensively reviewed.5O Four conformations of RNA can be defined as described below. 84 H. Schuster Biochem. Biophys. Res. Comm. 1960 2 320; W. Vielmetter and H. Schuster ibid. 1960,2 324. 9a S. W. Brostoff and V. M. Ingram Science 1967,58 666. * J. T. Penniston and P. Doty Biopolymers 1963 1 145,209. iu K. Miura Progr. Nucleic Acid Res. 1967 6 39; P. B. Moore J. Mol. Biol. 1966 22 145; G. L. Brown S. Lee and D. Metz ‘Genetic Elements Properties and Function’ ed. D. Shugar Academic Press New York 1967 p. 57. 86 P. 0. P. Ts’O G. K. Helmkamp C. Sander and F. W. Studier Biochim. Biophys. Actu 1963,76 54. 87 G. H. Beavan E. R. Holiday and E.A. Johnson ‘The Nucleic Acids’ ed. E. ChargafT and J. N. Davidson Academic Press New York 1955. a* I. Tinoco ‘Molecular Biophysics’ ed. B. Pullman and M. Weissbluth Academic Press New York 1965. 89 G. Felsenfeld and H. T. Miles Ann. Rev. Biochemistry 1967 36 407; A. M. Michelson ‘The Chemistry of Nucleosides and Nucleotides’ Academic Press London 1963 p. 339; M. L. Petermann ‘The Physical and Chemical Properties of Ribosomes’ Elsevier Amsterdam 1964. 506 cox Table 2 Examples of the reactivity of the common purine and pyrimidine base- residues Site of reaction -~ ~ Reagent Adenine Cytosine Guanine Uracil -~ ~ ~~ ~~~ ~~ Formaldehydea C(6)-NH2 C(6)-NH2 C(2)-NH2 N(l) Nitrous acidb C(6)-NH2 C(6)-NH2 C(2)-NH2 - - Dimethyl sulphatecsd N(l) N(3) N(7) DiazomethanecSd N(1) N(3) “7) N(3) Hydrazinee - C(4)=C(5) - C(4)=C(5) Hydroxylaminef - C(4)-C(5) - C(4) =C(5) H24g C(4)-C(5) - C(4) =C( 5) Osmium tetroxideZ - C(4)-C(5) - C(4) =C(5) Semicarbazid& - C(6)-NH2 - - Monoperphthalic acidk N( 1) - - I HalogensdSh C(6) or C(8) C(5) C(6) or C(8) C(5) Kethoxalz# - - N( l)-C(2)-NH CMECnsP - - Girard-p*lr - C(6)-NH2 - H+8 N(l)=C(6)- N(l) C(6)- C(2)-NH2 - N(1) - N(1) NH2 OH-8 - - -NH*C(6)0 -NH*C(6)0 NH2 a H.Fraenkel-Conrat Biochim. Biophys. Acta 1954,15,307; L. G-rossman S. S. Levine and W. S. Allison J. Mol. Biol. 1961,3,47; E. J. Eyring and J. Ofengand Biochemistry 1967 6 2500; H. Boedtker Biochemistry 1967,6,2718; b J. A. Carbon Biochim. Biophys. Acta 1965 95,550; H. Schiister and 0. Schramm Z. Naturforsch. 1958,13b 697; R. Shapiro and R. S. Klein Biochemistry 1966 5 2358; R.Shapiro and S. H. Pohl Biochemistry 1968 I 448; C B. E. Oriffin Methods in Enzymology 1967,12 A 141 ; I. Wempen and J. J. Fox Methods in Enzymology 1967 12 A 76; 6 D. M. Brown Methods in Enzymology 1967 12 A 31 ; f J. H. Phillips Methods in Enzymology 1967 12 A 34; 8 A. H. Schein and F. T. Schein Methods in Enzymology 1967,12 A 38; * C. T. Yu and P. C. Zamecnik Biochim. Biophys. Acta 1963,76,209; Science 1964,144,856; J. H. Weil Bull. SOC. Chim. biol. 1965,47 1303; * K. Burton N. F. Varney and P. C. Zamecnik Biochem. J. 1966,99,29~-31c; K. Burton Biochem. J. 1967,104,686; H. Hayatsu K. Takeishi and T. Ukita Biochim. Biophys. Acta 1966,123,445; k F. Cramer and H. Seidel Biochim. Biophys. Acra 1964,91,143; Z p-Ethoxy- a-ketobutyraldehyde (kethoxal); m R. Shapiro and J.Hachmann Biochemistry 1966 5 2799; M. Litt and V. Hancock Biochemistry 1967 6 1848; N-CyclohexyI-N-~-(4-methyl- morpho1inium)ethylcarbodi-imide (CMEC); p P. T. Gilham J. Amer. Chem. SOC. 1962 84 687; 0. Augusti-Tocco and G. L. Brown Nature 1965 206 683; D. G. Knorre E. G. Matygin G. S. Mushinskaya and V. V. Favorov; Biokhimiya 1966,31,334; Qacetohydrazide pyridinium chloride; r K. Kikugawa H. Hayatsu and T. Ukita Biochim. Biophys. Acta 1967,134,221 ; 8 A. R. Peacocke Chem. SOC. Special Publ. 1957 No. 8,139. A. The Amorphous Form.-The properties of polyuridylic acid are consistent with an amorphous structure i.e. the base-residues have no preferred orientation with respect to one another. When uridylic acid is polymerised to polyuridylic acid a hypochromic effect is observed i.e.the extinction below 300 m p is decreased by about lo% and the optical rotatory dispersion per residue is en- 507 Macromolecular Structure and Properties of Ribonucleic Acids han~ed.~O These optical properties are scarcely affected41 by increasing the tem- perature from 25 to 95". The base-composition of an oligonucleotide in the amorphous form can be deduced from its optical properties.P2 At moderate ionic strength the pK of the titratable groups of the base-residues will be slightly higher than the pK of the corresponding nucleotide because of the additional work required to move a proton away from the negative field of the sugar phosphate backbone. The form of the pH-titration curve is g i ~ e n ~ l ~ ~ by the equation (1) where a is the degree of dissociation of the base residue pKm pK = pKm + dpK = pH + log(1-a)/a (1) is the negative logarithm of the dissociation constant of the monomer and dpK is dependent on the electrostatic potential $ of the polynucleotide.It has been suggested that dpK = 0.434 e$/kTwhere e is the charge of the electron.44 It appears that the negatively charged diesterified phosphate residues (PK = ca. 1.5) make the major contribution to $ so that there is little variation in dpK over the range pH 3-14. Thus for polyuridylic acid the plot of a against pH has the same form as for UMP although the pK is found to be 9.7 in O.~M-KC~ at 25" compared with 9.4 for UMP. However $ (and hence dpK) is dependent on the ionic strength of the solution so that pK may increase by as much as 1-5 units when the electrolyte concentration is decreased.The hydrodynamic properties of poly-U are also influenced by the poly- electrolyte properties of the diesterified sugar phosphate residues. Repulsion between the phosphate residues causes the hydrodynamic volume to expand and in the absence of added electrolyte the chain is highly extended as shown by the high viscosity and low sedimentation coefficient. When the concentration of electrolyte (e.g. KCI) is increased above about 0 3 m ~ the viscosity of a flexible polyelectrolyte is found45 to decrease according to the empirical relation (2) whence [TI and [qO] respectively are the limiting viscosity numbers when the electrolyte concentration is Csalt and Cosalt (a reference state) and m lies within the range 04-0.6. The viscosity of poly-U obeys41 equation (2) where m = 0.38.In O.l5~-NaCl-@O1 S~-sodiurn citrate solutions the dependence of S,, and [TI on molecular weight are given43 by the equations (3) and (4). 40 H. Simpkins and E. G. Richards J. Mol. Biol. 1967 29 349. O1 E. G. Richards C. P. Flessel and J. R. Fresco Biopolymers 1963 1 431. 42 S. Mandeles and C. R. Cantor Biopolymers 1966 4 759. 43 R. C. Warner and E. Breslow Proc. 4th Int. Congr Biochem. Vienna 1958 vol. 9 p. 157. 44 C. Tanford 'Physical Chemistry of Macromolecules' J. Wiley New York 1961; A. Katchalsky J. Mazur and P. Spitnik J. Polymer Sci. 1957 23 513; A. Katchalsky Z. Alexandrowicz and 0. Kedem 'Transactions of the Symposium on Electrolyte Solutions' The Electromechical Society New York 1965. 45 R. A. Cox J. Polymer Sci. 1960 47 441. 508 cox B. Single-stranded Helical Structures.-Many polynucleotides approach the amorphous state at high temperatures.However the base-residues have an increasing tendency to interact with one another as the temperature is lowered.46 The term ‘stacking’ is often used to describe this interaction because the flat purine or pyrimidine rings tend to ‘stack‘ like a pile of pennies with the plane of the base-residues perpendicular to the helix axis.47 It appears that the distance between adjacent base-residues is about 3.4 A and that there is a translation of 30-45” between adjacent residues there being 8-12 residues per turn of the helixa4* The structure does not seem to be rigid except possibly at low tempera- tures if the degree of stacking is very high. Interaction takes place principally between nearest neighb0u1-s~~ so that the dinucleotide may be regarded as a model for the single-stranded polymer.The dinucleotide adenyl(3’ 5’) adenylate (ApA) has been extensively studied and the enthalpy and entropy for the reaction ‘stacked’ -+ amorphous form were calculated50 to be f 8 kcal./mole and +25 to 30 cal. mole-l deg.-l. Similar values were calculated for the enthalpy and entropy of polyadenylic acid confirming that pairs of base-residues contribute independently and non-co-operatively to the equilibrium between ‘stacked’ and amorphous conformation^^^ Measurements of mass per unit length confirm that poly A is ~ingle-stranded.~~ The optical properties of ApA are typical of poly-A which has a single helical ~onformation>~7~~7~~ Thus the maximum extinction (cp) per residue (observed at about 260 mp) is 13,800 at 4” and increases gradually to 14,700 when the temperature is increased to 65” whereas cp for the monomer is 15,400.The mononucleotide has a small positive rotation owing to the ribose moiety. The dimer has a much larger rotation which arises from the dissymmetry of the con- formation. A Cotton effect is observed below 300 mp within the region of the spectrum where the base-residues absorb and circular dichroism is observed. ‘Stacking’ tends to suppress the ionisation of the base-residues and may modify the shape of the pH-titration curve52 which may have a slope (da/dpH) that is no more than twice that found for mononucleotides. lonisation of the base-residues tends to reduce ‘stacking’. Because polynucleotides are also poly- electrolytes it might be expected that the stability of the single-helical con- formation will depend on the electrolyte concentration.The available data53 show that the effect is small. An idea of the hydrodynamic properties of the single-helical conformation is derived from the properties of poly-A. The dependence of S,o,w and [q] upon 46 A. M. Michelson J . Chem. SOC. 1959 1371. 47 J. Brahms A. Michelson and K. E. Van Holde J. Mol. Biol. 1966 15 467. 4a C. A. Bush and I. Tinoco jun. J. Mol. Biol. 1967 23 601. 49 M. Leng and G. Felsenfeld J . Mol. Biol. 1966 15 455. 6o K. E. Van Holde J. Brahms and A. M. Michelson J. Mol. Biol. 1965 12 726. 61 V. Luzzati A. Mathis F. Masson and J. Witz J. Mol. Biol. 1964 10 28. 62 R. A. Cox Biochem. J. 1966 100 146; €3. Simpkins and E. G. Richards Biochemistry 1967 6 25k3. 63 (a) R.A. Cox and K. Kanagalingam Biochem. J. 1967 A 103,431 ; (b) R. A. Cox. and K. Kanagalingam ibid. 1967 A 103 749; (c) D. Barszcs and D. Shugar Acta Biochim. Polon. 1964 11 481. 509 Macromolecular Structure and Properties of Ribonucleic Acids molecular weight was foundu to be as given by equations (5) and (6) indicating s20,w = 2.1 x 1 0 - 2 ~ 0 . 4 5 that at 20" in 0.1 5hl-NaCI-O.01 5hl-sodium citrate pH 7.1 owing to base-stacking poly-A forms a more compact coil than poly-U. The viscosity was found to increase when the electrolyte concentration decreased as expected of a flexible polyelectrolyte. At lower temperatures when the tendency to 'stack' is strong the hydrodynamic properties are no longer consistent with a random The differences in the spectra of 'stacked' and amorphous forms reported53b for oligo-A oligo-C and oligo-G are distinct so that the hypochromism of oligo- and poly-nucleotides due to 'stacking' depends on base composition.Studies of the hypochromism of di- and tri-nucleotides suggest that base- sequence also influences hypo~hrornisrn.~~ The optical properties of the single-stranded form of a polynucleotide can be obtained to a first approximation by preventing the formation of base-pairs e.g. by reaction with f~rmaldehyde.~' C. The Double-helical Confoamation.-X-Ray diffraction showsK8 that high- molecular-weight RNA in the double-helical form and fragments of RNA probably of ribosomal origin have a structure that strongly resembles the A form of DNA it consists of two antiparallel polynucleotide chains stabilised by Watson-Crick base-pairs the distance between base-pairs being 3.05 k 0.05 A there being 10 or 11 base-pairs per turn of the helix.The 2'-OH groups of ribose may be free to form intermolecular hydrogen bonds. When native (double-helical) DNA is heated in solution denaturation (Le. a transition to a single-stranded form) takes place over a very narrow temperature range; the mid-point of this transition range is termed the melting temperature (Tm). In neutral solutions the T' of an RNA double helix is about 10" greater than that of a DNA double helix of similar nucleotide comp~sition.~~ The properties of the RNA double-helix in general resemble those of DNA. Thus the thermal stability of the double-helix form appears to increase as the properties of G:C base-pairs increase; the transition from the double-helical to single-stranded 54 J.R. Fresco and P. Doty J. Amer. Chem. SOC. 1957 79 3928. 65 H. Eisenberg and G. Felsenfeld J. Mol. Biol. 1967 30 17. 66 G. B. Zavil'gel'skii T. V. Venkstern and A. A. Baer Doklady Biochemistry 1966 166 38; G. B. Zavil'gel'skii T. V. Venkstern and A. A. Baer Doklady Akad. Nauk S.S.S.R. 1966 166 978; W. M. Stanley jun. and R. M. Bock Analyt. Biochem. 1965 13 43. 67 G. D. Fasman C. Lindblow and E. Seaman J. Mol. Biol. 1965 12 630; H. Boedtker Biochemistry 1967 6 2718; R. A. Cox and K. Kanagalingam Biochem. J. 1968 108 599. 58 P. J. Gomatos and I. Tamm Proc. Nat. Acad. Sci. U.S.A. 1963 49 707; P. J. Gomatos and R. Langridge Science 1963 141,694; K. Tomita and A. Rich Nature 1964,201 1160; S . Amott F. Hutchinson M. Spencer M. H. F. Wilkins W. Fuller and R.Langridge Nature 1966 211 227. 59 H. C. Kaerner and H. Hoffman-Berling 2. Naturforsch. 1964 19b 593; A. R. Bellamy and W. K. Joklik J. Mol Biol. 1967,29 19. 510 cox forms is abrupt owing to the co-operative nature of the transition and thermal stability depends on the electrolyte concentrationsm (the temperature at which denaturation takes place increases by about 12" when the concentration of electrolyte is increased ten-fold within the range 1 mM-M). Multivalent cationssob and polyamines60c may either stabilise or labilise the double-helical form. The hydrodynamic properties of double-helical RNA would be expected to be the same as those of native DNA for which the relation between Sozo,w and M was found61 to be as in equations (7)-(10). The double-helical form can be So20,w = 0-116M0'325 (where M < 4 x los daltons) So20,w = 0.034M0*405 (where M > 4 x lo6 daltons) [q] = 1.05 x 10-7M1'32 (where M < 2 x lo6 daltons) [q] = 6.9 x 104M0'70 (where M > 2 x lo6 daltons) (7) (8) (9) (10) seen in the electron microscope62 and its mass per unit length can be deduced by low-angle X-ray ~cattering.~~ When solutions are made acid or alkaline Tm is decreased as the pH is made more extreme until denaturation takes place at the ambient temperat~re.~~ Ionisation is suppressed in the double-helical form so that denaturation (which remains co-operative at all pH values) is also accompanied by an abrupt increase in the degree of ionisation of base-residues.The difference in the spectra of the double-helical and single-stranded forms over the range 220-300 mp accurately reflects the nucleotide composition of the do~ble-helix.~~ D.The Hairpin-loop Conformation.-Many species of RNA are single-stranded at high temperatures and low ionic strengths but as the temperature is decreased or the ionic strength increased the polynucleotide folds upon itself forming short hairpin loops that are stabiIised by interaction principally between comple- mentary base-residues located on different segments of the same molecule.g5 The transitions are illustrated in the Figure 4. The partly double-helical form shown in (c) may be over-simplified but serves to illustrate the principal features. Three parameters the number N of base-pairs per loop the number b of un- paired residues per loop and the number c of unpaired residues which link one hairpin loop to another characterise the structure.Approximate values of N b and c can be deduced from degradation studies.66 6o (a) A. R. Bellamy and W. K. Joklik J. Mol. Biol. 1967 29 19; C. Schildkraut and S. Lifson Biopolymers 1965,3 195; W. F. Dove and N. Davidson J. Mol. Biol. 1962,5,467; (b) G. L. Eichhorn Nature 1962 194,474; (c) E. Gabbay and R. Kleinman J. Amer. Chem. SOC. 1967 89 7123. 61 J. Eigner and P. Doty J. Mol. Biol. 1965 12 549. 62 N. Granboulan and R. M. Franklin J. Mol. Biol. 1966 22 173. 63 Ch. Zimmer and H. Venner J. Mol. Biol. 1963,7 603. 64 G. Felsenfeld and G. Sandeen J. Mol. Biol. 1962 5 587; S. Z . Hirschman and G. Felsenfeld ibid. 1966 16 347; H. R. Mahler B. Kline and B. D. Mehrotra ibid. 1964 9 801. 65 (a) A. S . Spirin Progr. Nucleic Acid Res. 1963 1 301 ; P.0. P. Ts'o Ann. Rev. Plant Physiol. 1962 13 45; (b) P. Doty H. Boedtker J. R. Fresco R. Haselkorn and M. Litt Proc. Nut. Acad. Sci. U.S.A. 1959 45 482. R. A. Cox Biochem. J. 1968,106 725; R. A. Cox H. J. Gould and K. Kanagalingam ibid. p. 733. 51 1 Macromolecular Structure and Properties of Ribonucleic Acids Figure 4 The principal conformations of a single-stranded polynucleotide (cf. R. A. Cox and K. Kanagalingam Biochem. J. 1967 103 431). (a) the amorphous form in which the base- residues (indicated by the short horizontal lines) have no preferred orientation with respect to one another (6) the single-stranded ‘stacked’ form in which the base-residues tend to pile one upon another to form a single helix (c) the hairpin-loop conformation where N is the number of base-pairs per loop b is the number of unpaired residues within the loop and c is the number of unpaired residues linking one hairpin loop to another The optical propertiese7 and acid-base propertiese8 of the partly double- helical form appear to be the sum of the properties of the double-helical and single-stranded regions.However the temperature at which the double-helical structure reverts to the single-stranded form depends on N the number of base- pairs per segment as well as upon pH and electrolyte concentrati~n!~ The relation between Tm for a particular value of N and Tm(co) (the value of Tm when N - a) has the form (1 1) where A is a c~nstant.’~ Although equation (11) 1/Tm = 1/Tm(a) + A/N (1 1) may not be precise it serves to show that Tm depends on N when N is small (e.g. <20).(a) J. R. Fresco L. C. Klotz and E. G. Richards Cold Spring Harbor Symp. Quantitative Biology 1963 28 83; G. Felsenfeld and G. L. Cantoni Proc. Nat. Acad. Sci. U.S.A. 1964 51 818; ( b ) W. Guschlbauer Biophysik. 1966 3 156; C. R. Cantor S. R. Jaskunas and I. Tinoco J. Mol. Biol. 1966 20 39. 68 R. A. Cox and U. Z. Littauer Biochim. Biophys. Acta 1963,72 188. 6Q P. Doty J. Polymer Sci. 1961 55 1 ; M. N. Lipsett L. Heppel and D. F. Bradley J. Biof. Chem. 1961,236 857; M . N. Lipsett ibid. 1964,239 1256; E. K. F. Bautz and F. A. Bautz Proc. Nat. Acad. Sci. U.S.A. 1964 52 1476. 70 J. Appliquist and V. Damle J. Amer. Chem. SOC. 1965 87 1450. 512 cox The hairpin-loop form appears to behave as a compact random coil in solu- tion and the mass per unit length approaches that of the double-helical form.Various authors'l have related the sedimentation coefficients S and molecular weight M by means of equation (12) and values of o! within the range 0.40-0.55 S = KM4 (12) have been reported. Thus far the relations described have limited application. The formation of the hairpin loops leads to an anomalous contraction of the molecule (Figure 5). The ont traction^^^^^ is evident from the dependence of [q] I I 1-01 I I 1 .O -10 -2.0 -3.0 Log [NaCI] Figure 5a The dependence of the intrinsic viscosity of RNA and of DNA on the concentration of added sodium chloride at 25'. (-) sodium ribonucleate in neutral solution (R. A. Cox J . Polymer Sci. 1960 47 441); (0) sodium deoxyribonucleate data of Conway and Butler J. Polymer Sci. 1953 11 277; 1954 12 199; (- -) values calculated (R.A. Cox loc. cit.) for RNA if simple polyelectrolyte behaviour were followed as in equation (1) assuming m = 0.6 Figure 5b Temperature-dependence of reduced viscosity of high-polymer RNA preparations of E. Coli RNA (2.3 mg.lml.) in phosphate bufer pH 7.3 of ionic strength 0.1 in the presence of0.01M-versene (after Spirin Colloq. int. Cent. nat. Rech. Sci. No. 106 Acides Ribonucleique et Polyphosphates ; Structure SynthBse et Fonctions Strasbourg p. 8 1) on ionic strength since equation (2) is no longer obeyed and the observed viscosity is about one-quarter of the predicted value. When the temperature is increased the viscosity of the polynucleotide (e.g. in O-lM-NaCl) increases about four- fold until the polynucleotide becomes entirely single-stranded. Judged by many criteria the transition between the entirely single-stranded and partly double-helical forms takes place reversibly.The extent to which double-helical regions are formed depends on the precise nucleotide sequence. Random 'I1 A. Maeda J. Biochem. (Tokyo) 1961 50 377; A. S . Spirin Biochemistry (U.S.S.R.) 1961 26 454; C. G. Kurland J. Mol. Biol. 1960 2 8 3 ; R. F. Gesteland and H. Boedtker ibid. 1964 8 496. 72 (a) R. A. Cox and U. Z . Littauer Biochim. Bi0phy.s. Acta 1962 61 197; U. Z. Littauer and H. Eisenberg Biochim. Biophys. Acta 1959 32 320; (b) H. Boedtker J. Mol. Biol. 1960 2 171. 51 3 Macromolecular Structure and Properties of Ribonucleic Acids copolymers of AMP and UMP or GMP and CMP readily form such structures. The transition between one form and another can be followed by measuring the changes in e x t i n c t i ~ n ~ ~ .~ ~ optical rotatory dispersion,67b circular dichr~ism,~~ lH n.m.r. spectra,74 mass per unit length,75 or hydro-dynamic ~ o 1 ~ m e ~ ~ ~ or by the techniques of p~larography,~~ infrared spectro~copy,~~ and electron micro- scopy.62,’S 6 The Macromolecular Structure and Properties of the Principal Species of RNA A. Messenger. RNA.-The isolation of pure messenger RNA has proved difficult so that the macromolecular properties of this species are not known in detail. The size of messenger is determined by the size of the protein or proteins for which it codes. Messenger RNA may be polycistronic i e . code for more than one protein so that a wide range of molecular weights may be enc~untered.~~ There is increasing evidence that the messenger for haemoglobin has been iso- latedsO and its sedimentation coefficient is about 8-10s in O-IM-KCI in accord with a chain length of ca.450 residues the appropriate size of messenger coding for one of the polypeptide chains (about 150 amino-acid residues) of haemo- globin. Messenger from rat liver appears to be heterogeneouss1 and some species may be as large as lo6 daltons or more judged by their sedimentation properties. There is no information about the secondary structure of messenger RNA.s2 However it has been shown that the presence of stable secondary structure inhibits the translation of the genetic message. The synthetic messenger poly- uridylic acid which has little or no organised secondary structure due either to stacking or hydrogen bonding very efficiently directs the synthesis of poly- phenylalani~~e.~~ If poly-U is first allowed to react with poly-A to form the double-helical complex poly(A 4- U) there is no synthesis of polyphenylalanine.The nucleotide sequence of ‘messenger’ RNA has not been studied. 73 J. Brahms Proc. Roy. SOC. 1967 A 297 152. ’* C. C. Macdonald W. D. Phillips and M. Penswick Biopolymers 1965 3 609. 76 V. Luzzati J. Witz and A. Mathis ‘Genetic Elements Properties and Function’ ed. D. Shugar Academic Press New York 1967 p. 41 ; S. N. Timasheff J. Witz and V. Luzzatti Biophys. J. 1961 1 525. 76 E. Palac%k J. Mol. Biol. 1966 20 263. 77 H. T. Miles Biochim. Biophys. Actu 1958,30,324; F. B. Howard J. Frazier M. F. Singer and H. T. Miles J. Mol. Biol. 1966 16,415. 78 D. Danon Y. Marikovsky and U. Z. Littauer J. Biophys.Biochem. Cytol. 1961 9 253; N. A. Kisselev L. P. Gavrilova and A. S. Spirin J. Mol. Biol. 1961 3 778. 79 B. H. Hoyer B. J. McCarthy and E. T. Bolton Science 1963 140 1408; K. Asano J. Mol. Biol. 1965 14 71; G. Attardi H. Parnas M. I. Hivany and B. Attardi ibid. 1966 20 145; K. Scherrer K. Marcaud F. Zajdela 1. M. London and F. Gros Proc. Nut. Acud. Sci. U.S.A. 1966 56 1571; D. D. Brown and E. Littner J. Mol. Biol. 1966 20 81 95. 8o G. Marbaix and A. Burny Biochem. Biophys. Res. Comm. 1964 16 522; A. Burny and G. Marbaix Biochim. Biophys. Actu 1965 103 409; G. Huez A. Burny G. Marbaix and E. Schramm European J. Biochem. 1967 1 179. 82E. P. Guiduschek J. W. Moohr and S. B. Weiss Proc. Nut. Acud. Sci. U.S.A. 1962 48 1078. J. H. Matthaei and M. W. Nirenberg Proc. Nut. Acad. Sci. U.S.A.1961 47 1580 1588. T. Staehelin F. 0. Wettstein H. Oura and H. Noll Nature 1964 201 264. 514 cox B. Hybridisation of RNA with DNA.-The formation of a double-helix between a polyribonucleotide strand and a complementary polydeoxyribonucleotide strand is predicted on the basis of the theory of protein bio~ynthesis.~~ The first demonstration that a hybrid helix could be formed was achieved by use of model polynucleotides. The interaction in vivd5 and in vitro between RNA and de- natured DNA was developed as a means of isolating messenger RNA? t-RNA and ribosomal RNA interact with only a small fraction of the total DNA. C. Ribosomal RNA.-Ribosomes are nucleoprotein particles containing 50- ’60 % of RNA which can be dissociated reversibly into two unequal subparticles having the same RNA:protein ratio.The structural RNA from each sub- particle appears to be a continuous polynucleotide The molecular weight of RNA from the smaller subparticle is about 0.5 x los daltons irre- spective of its origin. Two species of RNA have been isolated from the larger subparticle. The larger species has a molecular eight^,^^^ of about (1-1.6) x los daltons (depending on the source of the ribosomes). The smaller speciess7 is known as SS-RNA and has a molecular weight of about 35,000 daltons. The RNA species differ in their nucleotide composition (see Table 3) and appear to have few sequences in common since each species hybridises with its own DNA cistron.88 The nucleotide composition of bacterial ribosomal RNA does not vary to the same extent as the base compo~ition~~ of bacterial DNA.Presum- ably the ribosomal RNA cistron which accounts for no more than 3% of the total DNA is copied very many times. Higher organisms (unlike bacteria which seem to have only one species of ribosome) contain ribosomes of different sizes. The ribosomes of the cytoplasm have a molecular weight of about 4 x lo6 daltons whereas ribosomes of the mitochondrion (or of chloroplasts) are about 3 x lo6 d a l t o n ~ . ~ ~ The larger species of RNA from cytoplasmic ribosomes has a molecular weight of about 1.6 x lo6 daltons compared with about 1.0 x los daltons for RNA isolated from mitochondria1 ribosomes. e4A. Rich Proc. Nut. Acad. Sci. U.S.A. 1960 46 1044; C. L. Schildkraut J. Marmur J. Fresco and P. Doty J. Biol. Chem. 1961 236 P C ~ . 8s G. Richter and H. §anger Biochim.Biophys. Acta 1965 95 362; M. N. Hayashi and M. Hayashi Proc. Nut. Acad. Sci. U.S.A. 1966 55 635; Y. Miura H. Itoh K. Sunaga T. Nishizawa and I. Ohki Biochim. Biophys. Acta 1967 134 258. 86 H. F. Lodish and W. D. Zinder Biochem. Biophys Res. Comm. 1965,19,269; D. Gillespie and S. Spiegelman J. Mol. Biol. 1965; 12 830; B. D. Hall and S. Spiegelman Proc. Nat. Acad. Sci. U.S.A. 1961 47 137. 87 R. J. Bachvaroff and V. Tongur Nature 1966,211,248; R. Rosset and R. Monier Biochim. Biophys. Acta 1963 68 653. 88 S. A. Yankofsky and S. Spiegelman Proc. Nut. Acad. Sci. U.S.A. 1963 48 1069; 1963 49,538; G. Attardi P. Huang and S. Kabat ibid. 1965,54 185; H. Wallace and M. Birnstiel Biochim. Biophys. Acta 1966 114 296; I. Merts W. Schulze and L. R. Overby Arch. Biochem. Biophys. 1966 115 197; D.Apirion J. Mol. Biol. 1967 30 255. A. N. Belozersky and A. S. Spirin ‘The Nucleic Acids’ ed. E. Chargaff and J. N. Davidson Academic Press New York 1960 vol. 3 p. 147. 90 (a) G. Brawerman Biochim. Biophys. Acta 1963 72 317; J. L. Chen and S. G. Wildman Science 1967 155 1271; T. W. O’Brien and G. F. Kalf J. Biol. Chem. 1967 242 2180; (6) P. J. Rogers B. N. Preston E. B. Titchener and A. W. Linnane Biochem. Biophys. Res. Comm. 1967 27 405. 515 Table 3 Base composition of ribosomal RNA isolated from animals plants and bacteria* MolesJ 100 moles of nucleotides 2 Q\ Molar ratios Organism S - value CMP AMP GMP UMP (A 4 U)/(C + G) Ref. Krebs ascites cells 30 30.6 16.0 .36-1 17-1 0.50 a Rabbit reticulocytes 28 31.6 16.4 35.3 16.6 0-49 b 18 27.9 21.2 29.4 21.4 0.74 18 28.8 20-5 30.7 20.0 0.68 Xenopus laevis Potato tuber Pea seedlings Cauliflower Escherichia coli Bacillus subtilis 28 30 17 37 16 0.49 C 18 29 22 31 18 0.67 25 22.0 25.1 31.7 21.2 0.86 16 22.2 25.4 27.2 25.2 1 -02 25 22-6 23.6 32-1 21.6 0.82 d 16 20.1 23-7 31.1 25-1 0.96 28 22.6 25-7 33.1 18-4 0.79 e 18 20.6 25.6 32.9 21.0 0.87 16 22.7 24-8 31.0 21.5 0.86 16 22.3 26.5 29.6 21-6 0-93 23 21.5 25.4 33.5 19.6 0.82 f 23 22.5 26-5 32.0 19.3 0.84 f Pseudomanas aeruginosa 23 21.2 26.3 31.2 21.3 0.91 f Drosophila melanogaster 28 19.6 30.8 22.5 27.1 1.38 g 16 21.6 25.1 32.8 20.5 0.84 16 20- 3 28- 8 23.5 27.4 1.28 Hyalophora cecropia (total RNA) 23-4 23.2 31.2 22.2 0.83 h * Data taken from (a) L.Montagnier and A. D. Bellamy Biochim. Biophys. Acta 1964 80 157; (b) R. H. DeBellis N. Gluck and P. A. Marks J. Clin.Invest. 1964,43 1329; ( c ) D. D. Brown and J. B. Gurdon Proc. Nat. Acad. Sci. U.S.A. 1964,51 139; (e) C. J. Pollard Biochem. Biophys. Res. Comm. 1964 17 171 ; (f) J. E. M. Midgeley Biochim. Biophys. Acta 1962,61,513; ( g ) G. Hastings and K. Kirby Biochem. J. 1966,100 532; (h) G. R. Wyatt and B. Lenzin Biochim. Biophys. Acta 1965,103 588. Ribosomal RNA from different organs of the same species have the same nucleotide composition (Hirsch Biochim. Biqhys 123,246) and &esamese$imentation e ~ c ~ ~ - % R 4 f f ; t ; T w r t 5. a. Biol.. 1967 25. 111). cox Comparatively little is known about the secondary structure of ribosomal RNA. Apparently structural RNA of both the smaller and larger subparticles have a hairpin-loop structure within the ribosome which is retained in solution after the protein moiety is About 50-70 % of the base-residues form b a s e - p a i r ~ .~ ~ ~ ~ The average size of the hairpin loops was estimated on the basis of the statistics of degradation of linear polymers to be 25 f 5 residues and 35 rt 7 residues respectively for RNA from the smaller and larger ribosomal subparticles of rabbit reticulocytes.66 These values are towards the lower limit obtained from measurements of mass per unit length.51,75 The secondary structure of ribosomal RNA ‘melts’ over a broad temperature range suggesting that the double-helical segments differ appreciably either in the number or composition of the ba~e-pairs.~~ Analysis of the nearest neighbours of the nucleotides of rat liver ribosomal RNA shows that long runs of a particular sequence can be excluded.g3 Partial enzymic hydrolysis leads to the production of discrete fragments which differ in nucleotide composition.53a~94 The stepwise hydrolysis of rat liver ribosomal RNA suggests that uracil residues are unevenly distributed along the chain.95 The heterogeneity of the hairpin loops is matched by the heterogeneity of ribosomal proteins.g6 Ribosomal RNA exhibits h y s t e r e s i ~ ~ ~ ~ ~ on titration over the range pH 3-7 since curve I (Figure 6) is followed on titration from pH 7 to pH 3 whereas curve II (Figure 6 ) is followed on titration with alkali from pH 3 to pH 7.Curves I and I1 are accurately reproduced on successive titration cycles and the hysteresis loop can be scanned by titration from intermediate pH values. The observations suggest that more than one conformation of ribosomal RNA is stable over the range pH 3-7.Fragments of ribosomal RNA having a molecular weight of about 13,000 daltons obtained by hydrolysis in O~~N-KOH form crystallites which have an X-ray diffraction pattern that is characteristic of double-helical RNA.98 The relation of these fragments to the intact molecule has yet to be established but they may be single intact or damaged hairpin loops. O1 C. E. Hall and H. S. Slater J. Mol. Biol. 1959 1 329; H. E. Huxley and G. Zubay ibid. 1960,2 10; G. Zubay and M. H. F. Wilkins p. 105; F. Bonhoeffer and H. K. Schachman Biocliem. Biophys. Res. Comm. 1960 2 366; D. Schlessinger J. Mol. Biol. 1970 2 92; A. Klug K. C. Holmes and J. T. Finch J. Mol. Biol. 1961 3 87; A. Blake and A. R. Peacocke Nature 1965 208 1319; P. McPhie and W.Gratzer Biochemistry 1966 5 1310; I. R. Cotter P. McPhie and W. Gratzer Nature 1967 216 864. p3 A. A. Hadjiolov P. V. Venkov and L. B. Dolopchiev Biochim. Biophys. Acta 1965 108 220. 04 P. McPhie J. Hounsell and W. B. Gratzer Biochemis&y 1965 5 988; H. Gould ibid. p. 1103; H. Gould Biochim. Biophys. Acta 1966 123 441; N. Delihas and J. Bertman J. Mol. Biol. 1966 21 391; N . Delihas Biochemistry 1967 6 3356. p6 A. A. Hadjiolov P. V. Venkov L. B. Dolapochiev and D. D. Genchev Biochim. Biophys. Acta 1967 142 1 11. O6 M. G. Hamilton and M. E. Ruth Biochemistry 1967 6 2585; W. Moller and A. Chrambach J. Mol. Biol. 1967 23 377; P. B. Moore R. R. Traut H. Noller P. Pearson and H. Delius ibid. 1968 31,441. R. A. Cox Biochim. Biophys. Acta 1963 A 68 401; R. A. Cox ibid. 1963 72 203.M. Spencer and F. Poole J. Mol. Biol. 1965 11 314; W. J. Pigram J. Littlechild and M. R. A. Cox Biochem. 1966,98,841. Spencer J. Mol. Biol. 1968 in the press. 517 4 Macromolecular Structure and Properties of Ribonucleic Acids Studies of nucleotide sequence have been confineds9 to the identification of the base-residues at the C(3')-OH and C(S)-OH termini. 2.0 ' 1 I I I 1 8.0 30 4.0 5.0 6.0 7.0 PH Figure 6 Scanning curves of ribosomal RNA from E. coli Solvent O.lM-NaC1 at 0.1". Curve I Titration with acid from pH 8. Curves I and IIa Titration from pH 8.00 to pH 3.80 to pH 5.2 to pH 3.80. Curves I IIb and Ia respectively Titration from pH 8.0 to pH 3-58 pH 3.58 to pH 4.88 to pH 3.58. Curves I and IIc Titration from pH 8.0 to pH 3.37 to pH 8.0. Curves I and JId Titration from pH 8.0 to pH 2.80 to pH 8.0 61 I20 (0) .G+- C G-G-G-C- C - I ) .G - G - V U - A - G U - A - C - V - U G G Q - U C - D G - l t t - A C - GG-UG-C-U-GU-A-G-G-C-U-U (- U 10" 67 99 110 I18 (b) Figure 7a Homologies between the two halves of the sequence of 5s-RNA The residues are numbered as in Figure 1. Homologies are shown by the boxed areas. Dashes are where gaps have to be left in the sequence in order to maximise these homologies. The underlining shows similarities between the two ends of the molecule (G. G. Brownlee F. Sanger and B. G. Barrell Nature 1967 215 735). Figure 7b Sequence of KB cell 5s RNA (B. G. Forget and S. M. Weissman Science 1967 158 1695) gQ B. G. Lane Biochemistry 1965 4 212; J. A. Hunt Biochem. J. 1965 95 541; J. E. M. Midgley Biochim. biophys. Acta 1965 95 232; J.E. M. Midgley ibid. 1966 123 210; J. E. M. Midgley and D. J. McIlreavy ibid. 1967 145 512; D. J. McIlreavy and J. E. M. Midgeley ibid. 1967 142,47; M. Takanami J. Mol. Biol. 1967,23 135. 51 8 cox D. 5S-RNA.-The function of SS-RNA associated with the larger ribosomal subparticle is unknown but the nucleotide sequence of SS-RNA from Escherichia coli (V) and from KB carcinoma cells (VI) has been reported (Figure 7). In both cases the molecules may have a clover-leaf conformationloo like that suggested for t-RNA (Figure 8) although this notion is far from established. Residues 31-37 (VI) and 41-47 are identical and so are the sequences 66-71 and 93-98 in (V) the sequences 10-19 and 61-70 are identical and so are the sequences 35-42 and 90-97 (see Figure 7). SS-RNA may be able to combine specifically with t-RNA.E. Transfer RNA.-The role of transfer RNA in protein biosynthesis was de- scribed above. This adaptor molecule is capable of combining chemically with the particular amino-acid specified by the anticodon. The nucleotide sequences of 5 t-RNA molecules from yeast have been determined four of which are given in Figure 8. Each species appears to be capable of forming a clover-leaf structure. The possible conformation of the anticodon loop deduced from model building is given in Figure 9. The general features of the clover-leaf structure include a high concentration of dihydrouracil in one loop and the anticodon is flanked by uncommon base-residues the anticodon itself may includelOl a 'wobble' i.e. a base-residue that is capable of forming base-pairs other than A:U and G:C.The regulation of secondary structure by means of the large number of modified base-residues is of particular interest. t-RNA possesses a specific tertiary structure.lo2 When certain species of t-RNA are heated to 70" in the presence of ethylenediaminetetra-acetic acid and cooled their biological activity and chromatographic behaviour are altered. The original active form may be regained by heating to 70" in the presence of Mg2+. The secondary structure of t-RNA has been intensively studied by opticaPo3 and physical lo4 methods as well as by chemical methodsllb based on measure- ments of the rate of reaction with reagents that are specific for particular base- residues (see Table 2). In general the results are in agreement with the clover- leaf structure.F. Viral RNA.-Each species of viral RNA has unique biological propertieslog which arise from a unique primary structure (the nucleotide composition and loo H. Boedtker and D. G. Kelling Biochem. Biophys. Res. Comm. 1967 29 758; C. R. Cantor Nature 1967 216 513. lor F. H. C. Crick J. Mol. Biol. 1966 19 548. loa W. Gartland and N. Sueoka Proc. Nat. Acad. Sci. U.S.A. 1966 55 948; T. Lindahl A. Adams and J. R. Fresco ibid. p. 941; A. Adams T. Lindahl and J. R. Fresco ibid. 1967,57 1684. lo3 C. R. Cantor S. R. Jaskunas and I. Tinoco J. Mol. Biol. 1966 20 39; J. N. Vournakis and H. A. Scheraga Biochemistry 1966 5 2997. lo4 J. A. Lake and W. W. Beeman J. Mol. Biol. 1968 31 115; V. G. Tumangan N. G. Esipova and L. L. Kiselev Dokludy Biochemisrry 1966 168 180; C. C. MacDonald W. D.Phillips and J. Penswick Biopolymers 1965 3 609; D. Bell and G. J. Russell Biochemistry 1967 6 3363. lo6 A. J. D. Bellet J. Virology 1967 1 245. 519 Macromolecular Structure and Properties of Ribonucleic Acids i P' 9 a' m 4' \ I 4' ' 2' ,3 h F' 0-5 F ? 0 rT. Y n 4 IX X L+ indicates the anticodon triplet Figure 8 Sequences and possible secondary structure of t-RNAs (VII) Two serine t-RNAs [The three nucleotides in serine t-RNA (11) which differ from those in serine t-RNA (I) are indicated in parentheses] (H. G. Zachau D. Dutting and H. Feldmann Angew. Chem. Internat. Edn. 1966 5 422; 2. physiol. Chem. 1966 347 212; H. Feldmann D. Dutting and H. G. Zachau ibid. p. 236; D. Dutting H. Feldman and H. G. Zachau ibid. p. 249). (VIII) Tyrosine t-RNA (J. T. Madison G. A. Everett and H.Kung Science 1966 153 531). (IX) Alanine t-RNA (R. W. Holley J. Apgar G. A. Everett J. T. Madison M. Marquisee S. H. Merrill J. R. Penswick and A. Zamir Science 1964 147 1462) (X) Phenylalanine t-RNA (U. L. Raj Bhandary S. M. Chang A. Stuart R. D. Faulkner. R. H. Hoskinsau and H. G. Khorana Proc. Nat. Acad. Sci. U.S.A. 1967,§7,571) 520 cox Figure 9 Schematic diagram of the t-RNA anticodon loop illustrating its relationship to the codon and the helical character of the structure The letters A B C and D identify the same points on the structure as in Figure 3. The bases of nucleotides 1 to 10 are stacked on one another and follow the regular helix which is shown black. The chain of the anticodon double-helix between D and B is shaded like the codon to indicate that they follow the same helix.This helix is complementary to the black one. The two nucleotides not in the standard conformation are represented by dark line shading. The representation of their conformation is very schematic because they lie behind nucleotides 8 9 and 10 in the black chain. The dotted lines indicate the generic helix from which the structure can be imagined to be derived (W. Fuller and A. Hodgson Narure 1967 215 817) 52 1 4* Macromolecular Structure and Properties of Ribonrrcleic Acids molecular weights of RNA from viruses are compared in Tables 4 and 5). A few of the viruses so far examined have a double-stranded RNA component,lm e.g. reovirus and wound tumour virus. The molecular architecture of double- stranded RNA has been studied.58 The smallest known viruses are called satellites because they can replicate only when another virus e.g.tobacco mosaic virus or alfalfa mosaic virus is present. The RNA component is about lo5 daltons in size lo’ and is sufficiently long to code for one or two small proteins.lo8 Next in size come the spherical bacterial viruseslog such as f2 R17 and M.S.2 whose RNA component is about 1 x 10s daltons. There are a few plant viruses in this class such as bromegrass mosaic virus.l1° There are two types of virus containing RNA of about 2 x lo6 daltons the rod-shaped plant viruses,111,112 such as tobacco mosaic virus and spherical viruses such as picornavirusesf13 (e.g. poliovirus) and some spherical plant viruses such as turnip yellow mosaic virus. The picornaviruses are the smallest RNA-containing animal viruses and a great deal more is known about their biochemistry than about their structure.The arboviruses (e.g. Sindbis and Semliki Forest virus) are more complex than picornaviruses since they have a thick lipoprotein membrane which sur- rounds a central core that resembles a picornavirus. Little is known about the RNA moiety of these viruses. Similarly our knowledge of the RNA of the rnyxo~iruses~~~ (e.g. influenza virus or Newcastle disease virus) is limited because of the difficulties in obtaining quantities of the virus and in isolating undegraded RNA.114 In the majority of animal plant and bacterial RNA viruses so far studied the RNA is single-stranded. It is probable that the RNA moiety of these viruses is a single molecule The molecular weight of the RNA moiety varies from about 1 to 3 x lo6 daltons.The RNA moiety is a genetic element and the amount of genetic information is proportional to the length of the poly- nucleotide chain. The introduction of viral RNA into its host cells may lead to the synthesis of complete virus. It was first shown in 1956 that when RNA from Io6 P. J. Gomatos and I. Tamm Proc. Nat. Acad. Sci. U.S.A. 1963 50 878. Io8 J. M. Clark A. Y. Chang S. Speigelman and M. E. Reichmann Proc. Nut. Acad. Sci. U.S.A. 1965 54 1193. lo* N. D. Zinder Ann. Rev. Microbiol. 1965 19,455; H. Hoffmann-Berling H. C. Kaerner and R. Knippers Adv. Virus Res. 1966 12 329. 110 L. E. Boekstahler and P. Kaesberg Nature 1961 190 192; L. E. Boekstahler and P. Kaesberg Biophys. J. 1962 2 1. R. Markham ‘Progress in Nucleic Acid Research’ ed. J. N. Davidson and W. E.Cohn Academic Press New York 1963 vol. 2 p. 61. Il*H. Schuster ‘The Nucleic Acids’ ed. E. Chargaff and J. N. Davidson Academic Press New York 1960 vol. 3 p. 245. 113 C. H. Andrewes and H. G. Pereira ‘Viruses of Vertebrates’ Ballitire Tindall and Cassell London 1964 2nd ed.; F. M. Burnet and W. M. Stanley ‘The Viruses’ Academic Press New York 1959 vol. 3; F. L. Schaffer and C. E. Schwerdt ‘Viral and Rickettsia1 Infec- tions of Man’ F. L. Horsfall and I. Tamm Pitman Medical Pub. Corp. London 1965 4th ed. p. 94. 11* P. H. Duesberg and W. S. Robinson Proc. Nat. Acad. Sci. U.S.A. 1965 54 794; M. W. Pons Virology 1967,31 523. M. E. Reichmann Proc. Nut. Acad. Sci. U.S.A. 1964 52 1009. 522 Table 4 Some RNA viruses of plants and bacteria Group Virus Particle weight M . w. of RNA Isolated RNA shown to be infective Nucleotide composition (moles/100 mmoles nucleotide) CMP AMP GMP UMP Satellite (plant) Rod-shaped (plant) Double-helical plant RNA Spherical (bacterial) Turnip yellow mosaic Tomato bushy stunt Southern bean mosaic Tobacco necrosis Tobacco ingot bromegrass Satellite tobacco necrosis Tobacco mosaic Potato X wound tumour QB f2 R17 M.S.2 fr 5.0 x lo6 10.7 x 10' 6-6 x 10' 8.0 x lo6 3-4 x lo6 4.6 x 10' 2-0 x 108 40.0 x 10' 36-39 x 106 70.0 x 10' 1.7 x 10' 1.7 x 10' 1.4 x 10' 1.5 x 10' 1.5 X 10' 1.0 x 10' 2-2 x lo6 2.0 x lo6 10.0 x 10' 3.94 x 105 - 3.6-4.19 X 10' 3.6 x 10' 4.1 x lo6 1.1 x lo6 1-8 X 10' 1.1 x 1@ 1-1 x 10' 1.2 x log Y E S YES YES YES 38-1 22.6 17.2 22.1 20.8 25.7 27.9 25.7 23.0 25.8 26.0 25.3 22-0 27.9 24.4 25.7 23.2 23.9 24.7 28.2 22.1 28.0 24.0 24.9 18.5 29.8 25.3 26.3 22.8 34.2 21.8 21.3 19.1 31.1 18.6 31.3 24.7 22.3 23.7 29.4 25.9 22.1 26-8 25.1 24-9 23.1 26.3 25.7 24.9 22.8 27.1 25.2 24.9 24.3 27.1 23.7 Lh Table 5 RNA viruses of vertebrates Nucleotide composition* 8 2 CJ Group Example Particle weight M.W. of RNA (moles/100 moles nucleotides) 5 B to be infective CMP AMP GMP U M . 3 Poliomyelitis YES 22-0 29.0 24.0 25.0 $ Picornaviruses Foot-and-mouth disease YES 28.0 26.0 24.0 22.0 3 Encephalomyo-carditis 22-27 mp - 2 x lo8 YES 23.0 27.0 24.0 26.0 Coxsakie YES 23.0 29.0 24.0 24.0 Rhinoviruses YES RNA has been shown or diameter 4 Eastern equine 25-50 mp 2 x log YES g Reovirus Reovirust 70 x lo6 60-90mp -10 x lo6 21.0 28-02 22.3 29.0 2 encephalomyo-cardit is % 2 Sindbis 4 0 - 4 8 mp 24.9 29.6 25.8 25.5 Arboviruses Semliki forest 50 mp YES West Nile 20-30 mp YES Myxoviruses Influenza 80-120 mp 2 x 106 ? 24-0 23.1 20.1 32.8 Paramyxo- Newcastle disease virus 100-200 mp 3-8 X lo6 27.0 26.1 24.9 22.0 viruses * Data from F.L. Schaeffer and C. E. Schwerdt 'Viral and Rickettsia1 Infections in Man' ed. F. H. Horsfall and I. Tamm Pitman Publishing Corporation London 4th edn. 1965 p. 94. t Reovirus RNA may exist on more than one molecular species within the virus (A. R. Bellamy L. Shapiro J. T. August and W. J. Joklik J. Mol. Biol. 1967 29 1). $ The high ratio of AMP is due to the presence of low molecular weight RNA rich in adenine residues (A. R. Bellamy L. Shapiro J. T. August and W. J. Joklik J. Mol. Biol. 1967 29 1). R (3 2 cox tobacco mosaic virus was rubbed into tobacco leaves the cells became infected with tobacco mosaic virus.l15 More recently RNA from a bacterial virus R17 was added to a cell-free system of Escherichia coli ribosomes and supernatant enzymes and the synthesis of viral coat protein was demonstrated.ll8 The viral RNA not only codes for its coat protein but also for its own RNA p01ymerase.l~~ Comparatively little is known about the secondary structure of viral ribo- nucleic acids118 although they adopt a hairpin-loop form in neutral salt solu- The structure in solution is not necessarily the same as in the virus.For example the RNA component of tobacco mosaic virus is entirely single- stranded within the virusllg but has a well defined partly double-helical form in soIution.65,i2b,120 In the case of spherical viruses such as turnip yellow mosaic virus the secondary structure of the RNA moiety may be similar within the virus121 and also after isolation.lZ2 The determination of the nucleotide sequence in the vicinity of the C(5’)-OH123 and C(3’)-OH124 termini has been attempted for RNA from tobacco mosaic virus and for RNA from the bacterial viruses f2 R17 and M.S.2.The replicative form of viral RNA has been isolated in many cases and was shown to be double-helical by its resistance to degradation by ribonuclease its sharp melting properties and its buoyant density.125 7 Conclusions The ribonucleic acids have fascinating physical and chemical properties in addi- tion to their important biological role. The correlation between structure and function remains a challenge even in the case of t-RNA. There is a need for nucleotides of known sequence that also form hairpin loops so that physical chemistry of short double-helical segments can be established.The secondary l15A. Gierer and G. Schramm Nature 1956 177 702; A. Gierer and G. Schramm 2. Naturforsch. 1956 l l b 138; H. Fraenkel-Conrat B. Singer and R. C. Williams Biochim. Biophys. Acta 1957 25 87. 116 D. Nathans G. Notani J. H. Schwartz and N. Zinder Proc. Nut. Acad. Sci. U.S.A. 1962,48 1424; Y. Ohtaka and S. Spiegelman Science 1963 142 493; D. Nathans J. Mol. Biol. 1965,13,521; M. R. Capecchi and D. Gussin Science 1965; 149,417; M. R. Capecchi J. Mol. Biol. 1966 21 173. 117 G. N. Gussin J. Mol. Biol. 1966 21 435. 118 S. Sprecher-Goldberger Archiv. f i r gesamte Virusforschung 1967,20 225. 119 W. Ginoza Nature 1958 181 958. lZo A. Gierer Z. Naturforsch.1958 13b 477. 121J. T. Finch and A. Klug J. Mol. Biol. 1966 15 344; A. mug W. Longley and R. Leberman ibid. p. 315; G. Zubay and M. H. F. Wilkins ibid. 1960 2 105. 122 R. E. F. Matthews and R. K. Ralph Adv. Virus Res. 1966 12,273. 12s R. Roblin J. Mol. Biol. 1968 31 51. lZ4 B. Singer M. Sherwood and H. Fraenkel-Conrat Biochim. Biophys. Acta 1965,108,306; P. R. Whitfield ibid. p. 202; H. L. Weith and P. T. Gilham J. Amer. Chem. SOC. 1967 89 21 5473; R. de Wachter and W. Fiers J. Mol. Biol. 1967 30 507. 125 L. Montagnier and F. K. Sanders Nature 1963 199 664; R. Langridge M. A. Billeter P. Borst R. H. Burdon and C. Weissmann Proc. Nut. Acad. Sci. U.S.A.. 1964 52 114; Y. Watanabe Biochim. Biophys. Acta 1965,95 515; D. Baltimore J. Mol. Biol. 1966 18,421 ; R. K. Ralph R. E. F. Matthews A.I. Matus and H. 0. Mandel ibid. 1965 11 202; B. Francke and P. H. Hofschneider ibid. 1966 16 544; G. Feix R. Pollet and C. Weissman Prw. Nat. Acud. Sci. U.S.A. 1%8 59 145; N. R. Pace D. H. L. Bishop and S. Spiegelman ibid. p. 139. 525 Macromolecular Structure and Properties of Ribonucleic Acids structure of ribosomal and viral RNA and other RNA species having a hairpin- loop conformation may then be better understood. The structure of ribonucleo- proteins remains ill-defined except possibly for small viruses. Future develop- ments may depend on the extension of sequencing technique to larger molecules on improved methods for the synthesis of polynucleotides of known sequence and on the production of crystals of nucleic acids and nucleoproteins for X-ray crystallographic studies.526

ISSN:0009-2681    

DOI:10.1039/QR9682200499

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

Isopoly-vanadates -niobates and otantalates By M. T. Pope and B. W. Dale* DEPARTMENT OF CHEMISTRY GEORGETOWN UNIVERSITY WASHINGTON D.C. 20007 U.S.A. 1 Introduction In aqueous solution vanadates molybdates tungstates and to a smaller extent niobates tantalates and chromates undergo complex hydrolysis-polymerisation reactions upon acidification. The major solute species generated by such re- actions are the so-called isopoly-anions of which dichromate [Cr20,I2- and heptamolybdate [Mo,O,]6- are well-characterised examples. So far as they are known the structures of isopolyanions as well as those of heteropoly- anions such as [siW1204,,]4- and [C~MO,O~~H,]~- can be regarded as specific fragments of metal oxide 1attices.l The metal atoms in all these compounds can be considered to be occupying interstices in close-packed arrays of oxide ions.In the oxides mixed oxides and some insoluble isopolyanions the close-packed oxygen atoms form infinite chains sheets or three-dimensional lattices whereas the structures of the soluble isopoly- and of the heteropoly-anions are discrete and usually of high symmetry. As well as behaving as conventional electrolytes albeit often of high charge and large size isopoly- and heteropoly-anions exhibit many of the properties of infinite lattices e.g. intraionic antiferromagnetism and the ability to undergo partial reduction forming soluble analogues of the tungsten and vanadium bronzes (heteropoly and isopoly ‘blues’). For such reasons polyanions can be used to examine the redox acid-base and catalytic properties of metal oxides by methods which are inapplicable to solids such as absorption spectroscopy and polarography.Further these anions can be used as convenient models for testing theories of bonding and metal-metal interactions in oxide lattices. A rationale of the formation structures and reactions of such a large funda- mental category of compound is far from clear at present. In spite of consider- able study during the past 50 years or so it is probably fair to say that of the six hydrolysis systems mentioned above only the isopolychromates are reason- ably well understood. Presumably the relative simplicity of chromate hydrolysis is related to the fact that while chromium(w) appears to be limited to tetrahedral co-ordination in oxide structures the larger sizes of V5+ Nbw Ta- Mow and Ws+ permit variable co-ordination numbers (4-6).Consequently the structural possibilities for isopolyanion formation are much greater with these elements . In this Review we consider critically information on the structures reactions * Present address Nuclear Physics Division A.E.R.E. Harwell England. 1 L. C. W. Baker in ‘Advances in the Chemistry of Co-ordination Compounds’ ed. S. Kirschner Macmillan New York 1961 p. 604. 527 Isopoly-vanadates -niobates and -tantalates and equilibria of those isopoly-vanadates -niobates and -tantalates which may be present in aqueous solution. Polyanions obtained from melts and other non- aqueous media are specifically excluded. A brief reading of the literature including some recent surveys soon reveals serious differences between the results and interpretations of different research groups particularly with regard to the vanadates.2a-2e [The subject has not been comprehensively and critically reviewed since Rosenheim’s ambitious attempt2f to correlate the formulae and reactions of all isopoly- and heteropoly-anions in terms of the now abandoned ‘Miolati-Rosenheim Theory’.] Some of these differences arise because the experimental conditions used are not strictly comparable but many have been (and some still are) due to invalid or overambitious reasoning from experimental data or to measurements inadvertently made under non-equilibrium conditions.In order to resolve some of the current debates we fmd it necessary to preface our Review with a brief discussion of experimental difficulties and the limitations of some commonly-used methods of investigation.2 Experimental Problems Table 1 lists those experimental methods which have most often been used for investigating isopolyanions in solution. A few less-common but potentially valuable methods such as n.m.r. are also included. Table 1 Methods for investigating isopolyanions in solution 1. EMF (pH) measurements ; equilibrium or rapid-flow 2. Salt cryoscopy 3. Light scattering 4. Equilibrium ultracentrifugation 5. Ion exchange 6. Ultraviolet and visible spectroscopy 7. Sedimentation ultracentrifugat ion 8. Infrared and Raman spectroscopy 9. Nuclear magnetic resonance (n.m.r.) spectroscopy 10. Diffusion and dialysis The methods listed can roughly be divided into those which provide informa- tion concerning the stoicheiometry of the hydrolysis-polymerisation reactions (numbers 1-6) and those which may reveal ionic structures or sizes (7-10).At the outset it should be made clear that in the opinion of the Reviewers no single method is capable of providing an unambiguous answer especially when the solution being investigated contains several solute species in comparable concentrations (i.e. is polydisperse). a(a) A. Morette and R. Rohmer in ‘Nouveau Trait6 de Chimie MinBrale’ ed. P. Pascal Masson et Cie Paris 1958 vol. 12; (b) L. G. Sillb Quart. Rev. 1959,13,146; (c) P. Souchay ‘Polyanions et Polycations’ Gauthier-Villars Paris 1963 ; (d) K. F. Jahr Angew. Chem. Internat. Edn. 1966,5,689; (e) Gmelins Handbuch der Anorganischen Chemie System Nr. 48 ‘Vanadium’ part B section 1 and 2 Verlag Chemie Weinheim 1967; cf) A. Rosenheim in Abeggs Handbuch der Anorganischen Chemie vol.4 part 1 Leipzig 1921 p. 977. 528 Pope and Dale Before discussing the limitations of the methods listed in Table 1 the possible complexities of the systems investigated must be considered. In a formal sense hydrolysis of a metal oxyanion such as V04% involves reactions in which protons are consumed and the net anionic charge per metal atom (Z) is reduced (e.g. from 3 in V043- to 0.5 in [HVl,02s]5-). Since the neutral end-products of hydro- lysis of metal oxyanions are the polymeric hydrous oxides (2 = 0) it is not surprising that polymeric anions are produced during the course of hydrolysis. What is perhaps surprising is that only a relatively small number of discrete isopolyanions seems to be formed rather than the complete range of inter- mediates found in other polymerisations.The various equilibria involved in anion hydrolysis may be formally divided into two categories (1) and (2). Protonation e.g. vo43- + H+ + HV02-; Condensation e.g. 2HV042- -+ V,O?- + H20 [HV,O0,,l5- + H+ + [H2V1@2J4- Polymerisations that do not appear to invalve condensations such as (3) and (4) have been postulated in some systems. However the degree of solvation of such species is not known and the mech- anisms of these reactions could well involve condensation steps. It has been recognised for many years that these reactions can proceed at very different rates and since the rapid-flow pH-measurements of Schwarzenbach? more quantitative estimates of such differences can be made. Thus it appears that while protonations are very rapid (< sec.) some polymerisations such as the formation of decavanadate from metavanadate or [H,W,204,]6- from [HW,021]5- can be extremely slow and may take several days weeks or even months to reach equilibrium at room temperatures.There is no completely satisfactory explanation why some reactions should be so slow although it will be noted that in the examples cited the reactions involve large increases in the degree of polymerisation and must involve gross structural changes. Solu- tions of isopolyanions may therefore contain species which are simultaneously reacting at several different rates. This situation leads to problems of irreversi- bility and the appearance of pseudo-equilibria. Because of the wide variation of anionic charge and size encountered in solutions of isopolyanions moderately concentrated solutions of supporting electrolytes (0.5-3.0~) have been used in those methods ( 1 4 6-7) requiring essentially ideal-solution behaviour for valid interpretation of data.Light- scattering and ultra~entrifugation~,~ techniques have clearly demonstrated that G. Schwarzenbach J. Inorg. Nuclear Chem. 1958 8 302. W. H. Nelson and R. S. Tobias Canad. J. Chem. 1964,42 731. W. H. Nelson and R. S. Tobias Inorg. Chem. 1963 2 985; 1964 3 643; J. Aveston and J. S. Johnson Inorg. Chem. 1964 3 1051. 529 Isopoly-vanadates -niobates and -tantalates highly charged polyanions bind appreciable numbers of counterions from the supporting electrolytes. This is known to affect the position of equilibrium quite significantly in some cases? Finally many methods (24 7) require extrapola- tion of data to infinite dilution.In those solutions where concentration-dependent equilibria are involved such an extrapolation is clearly not possible. Unfortunately many investigators have not taken adequate account of the above complications. As a general rule we have disregarded the conclusions of such studies and we do not quote numerical results if no indication was provided concerning (a) the experimental attainment of equilibrium (b) the temperature at which the measurements were made or (c) the nature or concentration of the supporting electrolyte. The most widely used method for investigating anion hydrolysis is through the interpretation of inflexions or breaks in pH or conductimetric titration curves. Although some of the earliest investigations of isopoly-vanadates’ and -niobates* were made by such methods in many cases the results obtained have not proved to be of much value for reasons such as those discussed above.In recent years Sillin and his colleagues at Stockholm aware of such difficulties have signi- ficantly improved the technique of pH measurement and interpretati~n.~ The method basically measures the number of protons or hydroxyl ions consumed following addition of a known amount of acid or base to a polyanion solution. The results obtained at several different total concentrations of metal ion are conventionally presented in the form of titration curves (2 against log [H+]). In the general case such curves are computer-fitted in terms of the formation constants of a limited number of solute species each defined by its 2 and N (number of metal atoms).Although in principle the interpretation of e.m.f. data should lead to an unambiguous result this is true only for relatively simple systems. For polydisperse solutes such as are often encountered with the vanadates the choice of solute species to be included in the data analysis has usually been made on the basis of other experimental information. Recently the Sillin group has developed a computer programme (LETAGROPVRID) which can systematically test a large number of solute species and discard those that are unsatisfactory. However the ambiguities discussed here still remain in complex systems. Consequently the main limitation of the method appears when several species are required to fit the whole range of data as for example in the ‘metavanadate’ region which is discussed more fully below (p.536). In order to distinguish between several plausible multicomponent schemes severe demands are made upon the accuracy and precision of the experimental data. Aveston et al.lo have recently demonstrated the ambiguity which can arise in such cases. However with the possible exception of the rapid-flow and n.m.r. methods mentioned below the method is the only one which has proved capable of detecting individual components of a polydisperse solute. G. Schwarzenbach and G. Geier Helv. Chim. Acra 1963 46 906. ’I P. Dullberg 2. phys. Chem. 1903 45 129. P. Sub Ann. Chim. 1937 7 493. L. G. Sillen Quart. Rev. 1959 13 146. lo J. Aveston E. W. Anacker and J. S. Johnson Inorg. Chem. 1964,3,735. 530 Pope and Dale Recently Schwarzenbach has developed a technique of rapid-flow pH measure- ment for polyanion system^.^ Besides determining stoicheiometries and forma- tion constants of stable species the method can detect short-lived intermediates and thereby provide mechanistic and possibly structural information.Specific applications are discussed later. Salt cryoscopy is probably the second most often used method of investiga- tion. It has recently received a thermodynamic treatment1' and can certainly be regarded as a valid method. In a commonly used system the transition point of reaction (5) is lowered by the presence of foreign ions. If sodium salts are used only free H+ HS04- and the isopolyanions are cryoscopically active and since the medium furnishes its own concentrated supporting electrolyte the behaviour is close to ideal.Difficulties arise (a) in making adequate allow- ance for free H+ and HS0,- ions particularly when their contributions to d t are larger than those due to the polyanions and/or when N is large and (b) when concentration-dependent equilibria prevent extrapolation to infinite dilution. Na2S04 10H20 + Na2S04 + 10H20 In spite of criticisms12 the method is valuable particularly for identifying or conhning the degree of polymerisation of essentially monodisperse solutes. Light-scattering13 and equilibrium ~ltracentrifugationl~ are complementary methods providing the same information by independent means. Unambiguous results are obtained only if the solutions are monodisperse and if the degree of polymerisation of the solute is unaffected by dilution. The data yield self- consistent values for N and z* the effective ionic charge which allows for the effects of counterion binding referred to above.In an analogous fashion sedi- mentation ~ltracentrifugationl~ can give the effective ionic weight M* of a monodisperse solute species. In each of these methods extrapolation to infinite dilution is necessary. Measurement of the number of metal ions bound by an anion-exchange resin together with the number of equivalents of counterions liberated can give a direct measure of 2. However if moderately rapid depolymerisations can occur or if the solution is polydisperse the information obtained by this method is presumably of little value. The remaining methods listed in Table 1 are of more limited scope or have been less frequently used. Since most isopolyanions absorb in the near ultra- violet region spectroscopy is not in general capable of differentiating between individual components of a polydisperse system.The method has however been used to investigate equilibria involving only two major solute species. The use of Raman spectroscopy to identify well-characterised ionic structures in solution by comparison with the solid-state spectra has been particularly valuable for 11 R. Fernandez-Prini and J. E. Prue J. Chem. SOC. (A) 1967 1974. la R. S. Tobias J . Znorg. Nuclear Chem. 1961 19 348; G. Schwarzenbach and G. Parissakis Helv. Chim. Actu 1958 41 2425. l8 R. S. Tobias and S. Y. Tyree jun. J. Amer. Chem. SOC. 1959 81 6385. 14 J. S. Johnson K. A. Kraus and G. Scatchard J. Phys. Chem. 1954,58,1034. 16 M. C. Baker P. A. Lyons and S. J. Singer J.Amer. Chem. Sac. 1955 77 201 1. 53 1 Isopoly-vanadates -niobates and dantalates niobates and tantalates (see below). Direct vibrational analysis and interpreta- tion of Raman and i.r. spectra is limited to simple symmetrical species. High- resolution n.m.r. spectroscopy has many advantages for studying polydisperse systems and can also give some structural information. So far the technique has only been used on vanadate solutions by use of 51V. The results were on the whole encouraging and the method when further developed and tested with well-characterised solute species could prove to be very useful. Measurements of diffusion and dialysis co-efficients although they may be capable of reveal- ing gross changes in the size charge or structure of the solute species as the pH of the solution is varied can give no quantitative information.3 Vanadates A Solution Properties.-Before a detailed discussion we give a brief overall picture of present knowledge regarding the chemistry of the vanadates in aqueous solution. Figure 1 shows the ranges of existence of the various vanadate species 14 \ \ \ \ \ \ \ \ \ \ \ ! H VO:- I I l l 12 10 \ \ '. - I 6 4 H 2 0 Figure 1 The approximate conditions of pH and total vanadium concentration under which a given species would be the major solute component of a vanadate solution at 25". Demarcations about which there is considerable doubt are represented by broken lines. The data used in the construction of this diagram were taken from L. Newman W. J. LaFleur F. J. Brousaides and A. M. Ross J. Amer. Chem. SOC. 1958 80 4491; N .Ingri and F. Brito Acta Chem. Scand. 1959,13 1971; D. Dyrssen and T. Sekine Acta Chem. Scand. 1961 15 1399; F. J. C. Rossotti and H. Rossotti Acta Chem. Scand. 1956 10 957. The data regarding the solubility of V,O are approximate 532 Pope and Dale as a function of pH and the total vanadium concentration. Such a diagram was first suggested for the vanadates by Schiller and Thilo,’6 but their version con- tained several errors. (For example if the law of mass action is assumed to hold all demarcation lines in the diagram should be straight. Also all equilibria not involving the proton or the hydroxyl ion must be represented by horizontal lines.) We have we hope corrected the errors and brought the diagram up to date. Nevertheless this diagram is intended as a visual aid not as a definitive and accurate statement of fact.For example the results quoted in the caption were not all obtained at the same ionic strength. We now discuss each area of Figure 1 starting from the alkaline side. There seems to be little doubt that [VOJ3- is the predominant species present in vanadate solutions at high pH. There is no evidence for the polymerisation of this ion and no credible measurement of its degree of hydration. It is a moderately strong base hydrolysing in solution to form [HV0412-. The hydrolysis constant or alternatively the pK of [HV0412- has been determined by severd workers; the most accurate measurements are listed in Table 2. Evidence for the dimerisation of the [HVO4l2- ion to yield the species [V207]4- (and its protonated forms) now seems to be conclusive.Values for the dimerisa- tion constant are given in Table 2. Ingri and Brito17 in the original report of their work on alkaline vanadate solutions could find no evidence for the un- protonated dimer although they postulated the existence of the protonated form [HV,O7I3-. But in a later recalculation of their data18 they concluded that there was evidence for both the protonated and the unprotonated form. Schwar- zenbach and Geier6 also found evidence for both forms and report that [V20,l4- is a strongzr base than [HV041Z-. Protonation of [HVO,]*- yields the ion [H2V0,]- (Schwarzenbach and Geier’s work6 indicates that the ion formed is a simple protonation product of [HVO4I2- and we therefore write it as [H2V04]- rather than [VO,]-) but this polymerises extensively if the vanadium concentration is greater than about g.-atom/l.The simple protonation reaction can therefore be studied only in very dilute solution,16 or by using a rapid-flow technique,6 although Ingri and Britol’ obtained a value for the hydrolysis constant by the analysis of equilibrium data obtained at vanadium concentrations in the range 6 x to 8 x g.- atom/]. The most reliable values for the hydrolysis constant of [HVO$- and the pK of [H2V04]- are listed in Table 2. The values depend quite strongly upon the nature of the medium employed. Schwarzenbach and Geier6 have interpreted their results in terms of complex formation between [HV0,I2- ions and the cations of the supporting electrolyte. There seems to be a case for using tetra-alkylammonium salts as supporting electrolytes for work on alkaline vanadate solutions.The nature of the species formed by the polymerisation of [H2V0,]- is perhaps l6 K. Schiller and E. Thilo Z. anorg. Chern. 1961 310 261. l i N. Ingrj and F. Brito Acra Chem. Scand. 1959 13 1971 ; F. Brito and N. Ingri Andes. Real. SOC. ESP. Fis. Quim. 1960 R 56 165. la F. Brito N. Jngri and L. G. Sillen Acra Chem. Scand. 1964 18 1557. 533 ul w P Table 2 Selected equilibrium data for aqueous vanadate solutions Equilibrium -Log K Method Temp. 13.2 13-0 & 0-3 HVO + v03- + H+ H,VO + H,V04- + H+ 8-31 8.36 7-88 8.04 8.23 3.70 3-40 VO,+ 3- H2O + HVO + H+ 3.30 Na2S04 cryos. u. v.Spect. U.v.Spect. U. v. Spect. Emf. E.m.f E.m.f. E.m.f. E.m.f. (rapid flow) 1E.m.f. (rapid flow) E.m.f. (rapid flow) E.m.f. (rapid flow) E.m.f. (rapid flow) Phase distr.E.m.f. (rapid flow) Phase distr. 32.4" 25 25 25 25 25 25 25 20 20 20 20 20 25 20 25 & 6' E sat. Na2S04 a 0.5 (NaCl) b 0. 3-0 (NaClO,) C $ & 3.0 (NaClO,) C 0.5 (NaCl) e Q E 0.5 (NaC1) b d 2 0 5 (NaCl) e 1.0 (T.M.A.Cl)t f 0.1 (T.M.A.Cl)t f Q 0.5 (NaCl) b d $ c 1.0 (NaCl) f 1.0 (KCl) f 0.1 (NaClOJ f 0-5 (NaClOJ L? 0.1 (NaClOJ f 0.5 (NaCIO,) g Method E.m.f. E.m.f. E . in. f. E.m.f. E.m.f. (rapid flow) E.m.f. (rapid flow) E.m.f. (rapid flow) E . m . f. E.m.f. (rapid flow) E.m.f. (rapid flow) E.m.f. (rapid flow) a P. Souchay and R. Schaal Bull. SOC. chim. France 1950 542. C L. Newman W. J. LaFleur F. J. Brousaides and A. M. Ross J. Amer. Chem. Sor.. 1958. 80 4491 F. Brito and N. Ingri Anales. Real. SOC. ESP. Fis. Quim. 1960 B 56 165. N. Ingri and F. Brito Acta Chem. Scand.1959 13 1971. F. Brito N. Ingri and L. G. Sillen Acta Chem. Scand. 1964 18 1557. f A. Schwarzenbach and A. Geier Helv. Chim. Acta 1963 46,906. g D. Dyrssen and T. Sekine Acta Chem. Scand. 1961 15 1399. * The magnitude of the ionic strength employed and the supporting electrolyte used. 1. T.M.A.CI = tetra-methyl ammonium chloride. 4 However see text. F. J. C. Kossotti and H. Rossotti Acta Chem. Scand. 1956 10 957. Values are also given for other media. Temp. 25 25 25 25 20 20 20 25 20 20 20 Medium* 1.0 (NaClO,) 3-0 (NaCIO,) 1.0 (NaClO,) 3.0 (NaClO,) 0.1 (NaClO,) 1 SO (NaClO,) 1 SO (NaClO,) 0.1 (NaClO,) 1 -0 (NaClO,) 0.1 (T.M.A.Cl)f 0.1 (T.M.A.Cl) Ref. h h h h f § f § f § f $ f f h Isopoly- vanada tes -n io ba tes and -tan tala tes the most important unsolved problem in this field.Dullberg’s original postulate,’ based upon pH titrations and conductivity measurements was that the major polymeric species in the pH range 6.5-8.2 was the trimer [V3O9I3- and he has been supported in this by several other workers.16J9 On the other hand a con- siderable body of reliable evidence has accumulated since 1953 to support the suggestion originally made in 1933 by Jander and Jahr,,O that the main product of polymerisation in this pH range is the tetramer [V,O,,]*-. The fact that most of the evidence for the tetramer has come from cryoscopic measurements may however be significant ; such measurements necessarily involve very high ionic strengths and usually high vanadium concentrations. The most reliable of these measurements are those of Jahr and his schoo1,21 and those of Naumann and Hallada., The latter workers studied as large a concentration range as they found was compatible with accuracy and also examined the equilibria on the acid side of the metavanadate region.This was to ensure that their results were not being distorted by the highly condensed polyvanadates known to be present in acidic vanadate solutions. Credible evidence for the [V30,]3- ion is provided by Schiller and Thilo.lS Using U.V. spectroscopy and very dilute solutions they obtained a set of consistent values for the trimerisation constant of [H,VO,]-. Unfortunately however they did not work at constant ionic strength and they also failed to record the temperature of measurement. The evidence provided by extensive e.m.f. measurements cannot yet be regarded as conslusive.Ingri and Brito in their original report,17 obtained a satisfactory theoretical fit for their experimental data on the assumption that the only important species in solution between pH 7 and pH 12 were [HVO4I2- [HV,O,I3- [H,VO,]- and [V3Ogl3-. However in a later recalculation18 using the computer programme LETAGROP they found that the overall standard deviation could be reduced from 0.012 to 0.009 if the set ofspecieswasenlarged to include [V20,]4- [V4013]6- and [V4012]4-. This slight improvement in the standard deviation cannot however be considered to be positive evidence for a tetrameric ‘metavanadate’ in view of the almost two-fold increase in the number of species being considered. [The prefixes ortho- pyro- and meta- with ref- erence to vanadates and niobates were originally introduced in the belief that these ions were analogous to the corresponding phosphates.The terms have no structural significance but have been retained for convenience. They refer to anions with Z-values of 3(ortho) 2(pyro) and l(meta).] More recent work by B r i t ~ ~ ~ on the acid side of the ‘metavanadate’ region (carried out at 40” to hasten the attainment of equilibrium) has also been inter- preted in terms of a trimeric and a tetrameric ‘metavanadate’ in equilibrium with one a n ~ t h e r . ~ ~ ? ~ In this case it was found that a satisfactory fit could be achieved l9 R. A. Robinson and D. A. Sinclair J. Clzem. Soc. 1934 642; P. Souchay and G. Carphi BiiII. SOC. chim. France 1946 13 160; P. Souchay ibid. 1947 14,914. 2o G. Jander and K. F. Jahr 2. anorg. Chem. 1933,212 1 .21 K. F. Jahr and L. Schoepp 2. Nafurfbrsch. 1959 14b 467; K. F. Jahr H. Schroth and J. Fuchs ibid. 1963 18b 1133. 22 A. W. Naumann and C. J. Hallada Inorg. Cliem. 1964 3 70. 23 F. Brito Anales. Real. SOC. ESP. Fis. Quim. 1966 B 62 123. 24 F. Brito Acfa Chem. Scand. 1967 21 1968. 536 Pope and Dale by assuming a set of five species each of which had been postulated at various times by other workers. However it was then found that a slightly better fit was obtained by the addition of several other rather less likely species. It is clearly necessary therefore to have some independent way of deciding what constitutes an acceptable species. The method assumes that all species having a given value of 2 and N have the same properties irrespective of the degree of counterion- binding.Such an assumption may have no validity. The measurements under discussion were all carried out in O - ~ M - N ~ C ~ (although some preliminary measurements were made in 3:Ohl-NaCI). Schwarzenbach and Geiefl have recently shown that there is appreciable complex formation between sodium ions and several vanadate species. Ingri Brito and Silldn do not seem to have adequately considered this point. Measurements made with a different support- ing electrolyte seem to be very desirable in this complex system. These criticisms notwithstanding the suggestion that trimeric and tetrameric ‘metavanadates’ are in equilibrium with one another in aqueous solution is one which explains all known credible data on this problem but the matter cannot yet be regarded as closed. (There is of course some inconsistency in criticising Ingri Brito and Sill6n’s evidence for the [V,O,#- ion while at the same time quoting their values for the hydrolysis constants of other species.However the figures that we quote are ones that remain consistent within the given limits of error irre- spective of what other species are considered to be present in solution. We would regard such figures as valid.) Simple protonation of [H,V04]- yields the uncharged vanadic acid H3V04. The second of these species is even more susceptible to polymerisation than the first but Dyrssen and Sekinez6 were able to study this reaction under equilibrium conditions by working with exceedingly dilute solutions and extracting the un- charged vanadic acid into an organic solvent. The distribution of vanadium between the two phases was determined by a sensitive radioactive tracer method.Dyrssen and Sekine give a value of 3.7 for the pK of H3VO4 (Table 2). Schwar- zenbach and Geier6 have also been able to study this equilibrium using their rapid-flow technique. They give a value of 3.3 for the first pK of H3V0, and attach considerable importance to the difference between this value and that of Dyrssen and Sekine. We shall return to this point. The nature of the species existing on the acid side of the ‘metavanadate’ region at ordinary concentrations is another subject that has provoked much discussion. Early workers based their conclusions almost entirely upon breaks and inflexions in conductimetric and pH-titration curves. They soon divided themselves into two camps; those who supported Dullberg’ in his belief that the main species present in weakly acidic vanadate solutions was the hexavanadate [v601,]k and those who preferred the decavanadates suggested by Britton26 and his co-workers.Very little real progress was made until 1956 when Rossotti 26 D. Dyrssen and T. Sekine Acta Chem. Scand. 1961 15 1399. )* H. T. S. Britton and R. A. Robinson J. Chern. Soc. 1932 1955; H. T. S. Britton and G. Welford ibid. 1940 764. 537 Isopoly-vanadates -niobates and -tantalates and R o ~ s o t t i ~ ~ published the results of their work. They chose to start in the low pH region where the predominant vanadium species is almost certainly the V02+ ion,28 and then studied the polymerisation that took place as the solution was made more basic. Using graphical methods they were able to account for their e.m.f.and spectral data on the basis of four species V02+ [H2V10o28l4- [HVlo02s]5- and [v10028]6-. (They could not be completely unequivocal about this but their results certainly could not be interpreted in terms of hexavanadate or other even less polymerised species.) Crysocopic measurements22 and e.m.f. measurements in a rapid-flow apparatus29 later provided independent verifica- tion of these results. However recent papers6s23124 have cast doubt upon some of this work. The Rossottis made their measurements between pH 1 and pH 7. It seems probable in the light of recent developments that they did not in fact achieve true equilibrium in the upper part of this range although they reported that no electrode readings were recorded unless they had remained stable for at least 3 hr.Schwarzenbach and Geiere have shown that although the mono- hydrogen decavanadate [HV,002s]5- is quite stable in fairly concentrated solu- tion ([Vltotal = 0.5 g.-atom/l.) it hydrolyses to the extent of about 4% in a more dilute solution ([V]total=O*l5 g.-atomll.) with the evolution of protons. Equilibrium is only achieved after six months or so. A dilute solution of the unprotonated decavanadate [Vl,0,,]6- disproportionates extensively during 6 months into the monoprotonated form and presumably ‘metavanadate’. On the assumption that in both cases the hydrolysis product was indeed ‘meta- vanadate’ and further assuming that this was in the trimeric form Schwarzenbach and Geier were able to obtain a value of 10-lo8 for K the equilibrium constant 6H20 4- 3[HVloo28]5- + 15H+ -/- 10[v30g]3- [ p = 1.0 (NaC10,); 2o°C] (6) of reaction (6).Now by combining the Rossottis’ results27 with those of Ingri and Britol’ and those of Dyrssen and Sekine,25 we can obtain a value of for K. This is surprisingly good agreement in view of the large number of steps required to derive the second figure and the differences in ionic strength and temperature that are involved. Hence it should have been obvious for some time that the unprotonated decavanadate could not be a thermodynamically stable species at pH 6-5 as is implied by the Rossottis. This point has been underlined by recent work of Brit0,2~,~~ some of which we discussed earlier. By computer analysis of e.m.f. measurements made at 40” in the pH range 2-8 he concluded that the major vanadate species present under such conditions were [V30g]3- [V4Ol2I4- [HV60,,]3- [HVlo02s]5- and [H2V10028]4-.He found no evidence for the unprotonated decavanadate [Vlo02s]e-. We have commented previously upon the validity of results obtained by the type of analysis used by Brito and in our opinion his conclusions require independent verification before they can be regarded as well-established. Such verification is not unfortunately provided a7 F. J. C. Rossotti and H. S. Rossotti J. Inorg. Nuclear Chem. 1956 2 201; Acta Chem. Scand. 1956 10 957. 28 M. J. LaSalle and J. W. Cobble J . Phys. Chem. 1955 59 519. as G. Schwarzenbach Pure Appl. Chem. 1962 5 377. 538 Pope and Dale by Naumann and Hallada’s cryoscopic work,22 even though they did undertake a study of the metavanadate-decavanadate equilibrium.It is extremely unlikely that any cryoscopic measurement would be able to distinguish a direct meta- vanadate-decavanadate equilibrium from one going through an intermediate hexavanadate. However Brito’s results are obviously in good general agreement with those of Schwarzenbach and Geier. Both investigations indicate that [v10028]‘- has a very limited range of existence if it is stable at all. In Table 2 we have recorded the formation constant of [HZV10028]4- from V02+ and also its two p&’s as determined by Rossotti and Rossotti. Despite what has been said above about the instability of the [Vlo02,]6- ion these figures are still valid. They are seen to agree quite well with those of Schwarzen- bach and Geier which are also shown. B. The Structures of Vanadate Ions in the Solid State.-This aspect of the subject was last reviewed in 1961,3O and we shall be concerned mainly with developments that have taken place since then.On a simple radius-ratio argument the V5f ion would be expected to be somewhat large for tetrahedral co-ordination to oxygen somewhat small for octahedral co-ordination. The possibility of very irregular structures therefore arises. This point has been discussed by Orge131 and we shall confine ourselves to empirical observations here. A detailed structure of a crystalline orthovanadate has yet to appear in the literature. Most of the present information comes from partial structural deter- minations performed on the very insoluble orthovanadates of tervalent metal cations or from the observed isomorphism of hydrated sodium ortho-vanadates -phosphates and -arsenate~.~~ These studies indicate that such compounds contain discrete [V0,’J3- anions and that these are essentially tetrahedral although slight distortions are sometimes found.The V - 0 bond lengths are reported to vary between 1.64 and 1.74 A although a value of ca. 1.66 I$ could be regarded as typical. The [V20,l4- ion was shown to be isostructural with the [P20,l4- ion by Peyronela in 1942 and the complete structure of Cd2V20 has recently been published by Au and C a l ~ o . ~ ~ The crystal contains discrete [V20,l4- ions which consist of two V04 tetrahedra sharing a corner. The vanadium ions are dis- placed slightly away from the centre of symmetry. However the bridging oxygen shows a very high thermal motion and it is not at all clear whether the V-0-V link should be regarded as linear or not.No structure is yet known in which a discrete ‘metavanadate’ ion can be dis- tinguished. All structures so far studied have been found to contain infinite 30 H. G. Bachmann and W. H. Barnes Z. Krist. 1961 115 215. s1 L. E. Orgel Discuss. Faraday SOC. 1958 26 138. 3a E. Broch Z. physik. Chem. (Leipzig) 1933 20 B 345; R. W. G. Wyckoff ‘Crystal Structures’ 2nd edn. Interscience New York vol. 3 1965 pp. 17,67; A. F. Wells ‘StructuraI Inorganic Chemistry 3rd edn. Oxford Univ. Press 1962 p. 686. 83 G. Peyronel Gazzetta 1942 72 83. s4 P. K. L. Au and C. Calvo Canad. J . Chem. 1967,45,2297. 539 Isopoly-vanadates -niobates and -tantalates chains of vanadium-oxygen polyhedra linked either through corners or edges. Two distinctly different types of chain have so far been described.In the an- hydrous metavanadates of which we shall take KVO and NH4V0 as examples since they have received the closest attenti0n,3~,~~ the vanadium atoms are quite definitely four-co-ordinate (Figure 2a). Of the four oxygen atoms surrounding h Figure 2 The structures of some isopolyanions in the solid state. (a) The structure of the infinite chains of VO1 tetrahedra found in anhydrous ‘metavanadates’. (b) The structure of the infinite chains of VOs polyhedra found in hydrated ‘metavanadates’. (c) The structure of the [Vlo0,8]s- ion. ( d ) The structure of the [Nb,01B18- and [Ta,01B]8- ions In (a) and (b) 8 = V; 0 = 0; in (c) and ( d ) each octahedron represents an MO unit. each vanadium atom two are responsible for linking the tetrahedra together and are therefore shared between two vanadium atoms.The V-0 bond lengths in this case are quite long 1-81 A. The other two oxygen atoms are unshared and the bond lengths here are found to be 1.66 A close to the typical distance found in the pure tetrahedral co-ordination of the orthovanadates. After these 36 V. SyneEek and F. Hanic Czech. J. Phys. 1954 4 120. 36 H. T. Evans Z. Krist. 1960 114 257. 540 Pope and Dale four the next nearest oxygen atom is more than 3 8 away and is clearly not involved in the co-ordination around the vanadium. In the hydrated meta- vanadates of which we take KVO, H20 as an e~ample,3~,~~ the vanadium atoms are quite clearly five-co-ordinate. Again the oxygen atoms can be divided into two groups; those that are shared between two vanadium atoms and those that are unshared (Figure 2b).There are three of the former type per vanadium atom and again the V-0 bond lengths here are long (1.93 1.93 and 1.97 8,) while the two unshared oxygen atom remain at 1.65 8 and 1-67 8,. The next nearest oxygen atom is 3.73 8 away. The co-ordination around the vanadium atom is approximately trigonal bipyramidal. The five-co-ordinate chain structure thus formed is very similar to that found in crystalline v,05.38 As described in the previous section slightly acidic vanadate solutions probably contain a great many different species at equilibrium. Moreover there is an additional complication in that equilibrium is reached very slowly. It is not surprising therefore to find that the nature of the solids that crystallise out of such solutions is very strongly dependent upon the exact conditions employed.When a vanadate solution is acidified beyond pH 6-5 it turns bright orange and at suitable concentrations will yield crystalline salts of the deca- vanadate ions [H2VloOz> [HV1002J5- and W,,O2,l6- in various degrees of hydration. However if sufficient acid is added to make the Z-value exactly 0.6 (corresponding to [VlOO2d6-) and the solution is then kept at 60° an almost insoluble light-orange product is obtained having the empirical formula M1V308.39 But if a solution containing slightly less acid is rapidly evaporated at 40" dark red crystals of a pentavanadate M13V501* are produced.39 Of these products the most important from a general point of view are the decavana- dates because of the support they provide for the conclusions of Rossotti and Rossotti and because of the strong possibility that these isopolyanions have the same structure in solution as they do in the solid state.The structure of the decavanadate ion w10028]6- was determined simul- taneously and independently by Evans>O working with K2Zn2V10028 16H20 by Swallow and Barnes$O using crystals of Ca,Vlo028 16H20 and by Pullman$O working with Na6V1&8 18H20. The results of these three investigations are entirely consistent and the structure of the decavanadate ion is as shown in Figure 2c. Six VO octahedra are linked by edge-sharing into a 3 x 2 rect- angular sheet. The other four octahedra are attached to this sheet two to the upper side two to the lower by sharing of sloping edges. Evans41 and Swallow Ahmed and Barnes42 subsequently published refinements of their data which showed that all VO octahedra were strongly distorted.The distortions can be C. L. Christ J. R. Clark and H. T. Evans Acta Cryst. 1954 7 801. as A. Bystrom K. A. Wilhelmi and 0. Brotzen Acta Chem. Scand. 1950 4 119; H. G. Bachmann F. R. Ahmed and W. H. Barnes 2. Krist. 1961 115 110. 39 H. T. Evans and S. Block Znorg. Chem. 1966 5 1808. 40 H. T. Evans A. G. Swallow and W. H. Barnes J. Amer. Chem. SOC. 1964 86 4209; N. Pullman Ph.D. Thesis Rutgers University 1966 (Diss. Abs. 1967,28 B 140). 41 H. T. Evans Znorg. Chem. 1966 5 967. 4a A. G. Swallow F. R. Ahmed and W. H. Barnes Acta Cryst. 1966 21 397. 541 Isopoly-vanadates -niobates and -tantalates described briefly by saying that all vanadium atoms are displaced away from the centre of the ion.Hence in each of the four octahedra that form the corners of the rectangular sheet and in each of the four that are on its top and bottom sides the vanadium atom is displaced towards the outermost oxygen atom. The short V-0 distances are found to be between 1.59 and 1.62 A. In each of the two octahedra that are in the centre of the rectangular sheet the vanadium atom is displaced towards the outside pair of oxygen atoms and the short V-0 bond lengths here are found to lie between 1.68 and 1.70 A. There are further less symmetric distortions but these are probably due to hydrogen bonding in the crystal. The structures of KV,O and CsV,O have been described by Evans and Blo~k.3~ In contrast to the decavanadates no discrete isopolyanion can be distinguished here.It is best first to consider the vanadium to be octahedrally co-ordinated. The octahedra then link together by sharing edges to form a buckled sheet. Adjacent sheets are held together entirely by the cations. How- ever within each octahedron the vanadium atom is strongly displaced towards one oxygen atom and two types of vanadium can be distinguished according to the extent of this distortion. In one case the short V-0 distance is 1.624 A and in the other it is 1.575 A. This leaves correspondingly long V-0 distances 2.28 A and 2.973 A respectively on the opposite sides of the octahedra. In the second case it is doubtful whether this oxygen atom can really be considered to be part of the vanadium co-ordination which now might be regarded as five- fold (square pyramidal). The structure of K3V6OI4 has not been reported in detail but Bystrom and Evans43 have shown that the vanadium-oxygen polyhedra are arranged in groups of five so that the vanadium atoms lie roughly in a plane forming an almost regular pentagon.Three of these vanadium atoms are in approximately square pyramidal co-ordination while the other two are in approximately tri- gonal pyramidal co-ordination. In each case there is a short bond to one apex of the pyramid (1-5-1.6 A) and a very long bond (cu. 3.4 A) through the opposite side to an oxygen atom in another layer of V-0 polyhedra. C. The Structures of Vanadate Species in Solution.-Experimental techniques which are capable of providing unequivocal evidence for the structures of solute species are at the moment non-existent except for the special case of highly symmetric molecules and ions.We therefore have very few hard facts regarding the structures of vanadate species in solution. In this situation it is very tempting to draw heavily on the few known crystal structures containing discrete vanadate ions. However such an approach could prove to be misleading unless very careful note is taken of the precise conditions under which the crystals were obtained. studies indicate that VOZ- is tetrahedral in solution Infrareda and 43 A. M. Bystrom and H. T. Evans Acta Chem. Scand. 1959,13 377. 44 H. Siebert Z. anorg. Chem. 1954 275 225. 46 W. P. Griffiths and I. D. Wickins J. Chem. SOC. (A) 1966 1087. 542 Pope and Dale as it is in the solid state. This is supported by slV n.m.r. work;46 the resonance corresponding to the vOp3- is very narrow indicating that the symmetry around the vanadium atom is cubic.According to Pauling’s well-tested theory the successive pK,’s of a polybasic acid such as H3V04 should be separated by equal increments provided that the basic structure does not change. Schwarzenbach and Geier’s values6 for the first and second pK of H3V04 obtained by protonation of [HVO4I2- in a rapid- flow apparatus are 3.4 and 8.23 [ p = 0 . 1 ~ (NaC104) 25’1. The work of Newman4’ and his co-workers yields a value of 13.5 for the third pK, calculated on the assumption that the ionic product of water is in 3h4-Nac1O4 at 25” and with correction of +0.3 in order to convert from ~M-N~CIO to 0 . 1 ~ - NaC104.10 The increments between these values are seen to be sensibly constant. Moreover the parallel between these values and those for phosphoric acid H3P04 (2-1 6-71 and 11.8) is seen to be quite close.This strongly suggests that the species p04]3- [HVO4I2- [H2V04]- and H3V04 are all based upon the same structure i.e. a tetrahedrally co-ordinated vanadium atom. Schwarzenbach and Geier’s results were however obtained under non-equilibrium conditions and it is necessary to ask whether these species are the same as those existing at equilibrium. In the case of [HV04J2- and [H2V04]- they probably are; Newman’s value for the third pK was the result of measurements made on equilibrated solutions and values for the second pK determined from equilibrium measure- ments agree with those obtained by the rapid-flow technique when differences in ionic strength are taken into account. However in the case of the first pK, there appears to be a genuine difference between the value obtained by the rapid- flow method and the equilibrium value obtained by Dyrssen and Sekine.25 The latter employed an ionic strength of 0-5 and obtained a value of 3-7 (Schiller and Thilo,le presumably using solutions of very low ionic strength found the first pK to be between 4-3 and 4.8 at equilibrium; which is consistent with the present argument).A positive correction must be made to Dyrssen and Sekine’s value of 3.7 to make it comparable with Schwarzenbach and Geier’s value of 3.4 obtained at an ionic strength of 0.1. When this is done it is apparent that the two results do not agree and this has been taken by Schwarzenbach and Geier to mean that ‘vanadic acid’ is not all in the form of H3VO4 at equilibrium.They suggest that some of it may exist as V(OH), which would be expected to be a considerably less acidic species than H3Vop. Howarth and Richards46 have shown that when solutions containing the [VO4I3- ion are progressively acidified a plot of the 51V chemical shift against average charge per vanadium atom (2) is linear for all species up to and including ‘metavanadate’. There is a definite discontinuity at decavanadate. Since the WO4l3- ion is tetrahedral these results suggest that the structures of the ions [HVO4I2- [HV20,I3- [V,0,l4- and [V30,J3- (and/or [V4O1$-?) (these are the *13 0. W. Howarth and R. E. Richards f. Chem. SOC. 1965 864. *’ L. Newman W. J. LaFleur F. J. Brousaides and A. M. Ross f. Amer. Chern. SOC. 1958 80,449 1. 543 Isopoly-vanadates -niobates and -tantalates only ions that would be present in appreciable amounts at the concentrations used by Howarth and Richards) are also based upon VO tetrahedra while the structure of the decavanadate ion is based upon some other co-ordination.If this is correct there could be quite a close parallel between the structures of these ions in solution and the corresponding structures in the solid state although the vanadium atoms in solid ‘metavanadates’ can of course be either four- or five- co-ordinate. Before the work of Howarth and Richardspg Ingri and Britol’ had put forward a different theory. Noting that in solid hydrated ‘metavanadates’ the vanadium atoms can be considered to be five-co-ordinate they pointed out that the species thought to be present in vanadate solutions under various conditions could all be based upon V05 units rather than VO tetrahedra.However given that WO4I3- is tetrahedral it is significantly more difficult to account for the above evidence on the basis of a pentaco-ordinate vanadium species although neither the thermodynamic nor the n.m.r. evidence can be regarded as conclusive An experiment performed by Schwarzenbach and Geiefl is very interesting in this context. To a solution in which VO$ ions were the predominant vanadium- containing species they added sufficient hydroxide ion to give a mixture of wO4l2- and [H2V04]- in equal proportions (k 2.5 moles of hydroxide per mole of V02+ more than required to neutralise the free acid present). After sec. the pH of the solution was 4.5; only after 10-1 sec. did it rise to 8.2.A possible interpretation of these results would be that the first-formed species were WO2(0H)J2- and [VO,(OH),]- which then had to undergo rearrangement in order to yield wV04]2- and [H2V04]-. At present however this can only be regarded as an attractive hypothesis. It is noteworthy that in solution there seems to be evidence for both trimeric and tetrameric forms of metavanadate while there is no evidence whatever for a dimer. This suggests that polymeric metavanadates are cyclic. This possibility was first put forward by Ingri and Britol’ in the case of w3O,]3- and by Simon and J a P 8 in the case of W4OI2l4-. chemical shift against the average charge per vanadium atom shows a discontinuity at the metavanadate-dmvana- date transition. This probably indicates a pronounced difference in the co- ordination around the vanadium atoms in these two species.As was also pointed out earlier the reaction of metavanadate with acid to form decavanadate is extremely slow. This could be taken as evidence for a change in the co-ordination number of the vanadium atoms although the corresponding reaction in the case of the molybdates where a change in 60-ordination is known to OCCUT has been reported3 to be fast. (In the case of the molybdates a monomer reacts with acid to form a polymer. In the case of the vanadates both reactant and product are polymers.) In the solid state of course there is a pronounced differ- ence between co-ordination around the vanadium in the metavanadates as opposed to the decavanadates and it is very tempting to suppose that the As pointed out earlier a plot of the J.Simon and K. F. Jahr Z. Nuturforsch. 1964 19b 165. 544 Pope and Dale [V,o0,,]6- ion and its protonated forms have the same structure in solution as they do in the solid state. However what little evidence there is on this point all of it from n.m.r. measurements seems to point to a different conclusion. Howarth and Richards46 report that the n.m.r. spectra of weakly acid vanadate solutions consist of two broad resonances and one narrow one having respective intensity ratios of 1 :3 :l. These features are practically independent of the total vanadium concentration. Such results seem to be quite incompatible with the supposition that the decavanadate ion has the same structure in solution as it does in the solid state since this would require intensity ratios of 2:2:1 (not necessarily in that order).However Hatton et ~ 1 . ~ who have also studied 51V n.m.r. spectra of vanadate solutions come to some quite different conclusions. The chemical shifts they report are unfortunately quoted relative to the reson- ance given by an aqueous solution of ammonium metavanadate but they seem to agree roughly with those of Howarth and Richards. However the line-widths and intensities are quite different in the two papers. While the more careful and extensive work of Howarth and Richards seems to be preferable there is con- siderable scope for a reinvestigation here. It would also be advantageous to test the n.m.r. technique on a well-established system. The rather elaborate cage structure suggested for the decavanadate anion by Jahr Fuchs and PreussSo is almost certainly wrong.The structure was suggested to account for the complexes that are formed between the decavanadate anion and alkaline earth cations and it consists of three interlocking rings of V04 tetrahedra. Such complexity is not required in order to account for the observed complex formation since it is now known that the hexaniobate anion which as we shall see later is structurally related to the decavanadate ion forms complexes with metal cations by accommodating them in the octahedral pockets that exist on its outer faces.51 Now that suitable laser sources are available there seems to be no reason why Raman spectroscopy could not be used to investigate the structure of the de- cavanadate ion in solution. This technique has proved to be very successful in determining the structure of heptamolybdate and also as we shall see hexanio- bate and hexatantalate ions in solution.4 Niobates and Tantalates We shall discuss the properties and structures of niobates and tantalates together since as far as they are known these are very similar. In contrast to the situation with the vanadates the hydrolyses of these anions appear to be quite simple and the literature is both less extensive and less controversial. A. Solution Properties.-Niobium and tantalum pentoxides are considerably more inert than vanadium pentoxide. Aqueous solutions of niobates and J. V. Hatton Y. Saito and W. G. Schneider Canad. J. Chem. 1965 43 47. 6o K. F. Jahr J. Fuchs and F. Preuss Chem. Ber. 1963,96 556. 51 B. W. Dale and M. T. Pope Chem. Comm. 1967 792; G. D. Stucky and C.M. Flynn jun. American Crystallographic Association Meeting Minneapolis Minn. 1967 Paper ~ 1 0 . 545 Isopoly-vanadates -niobates and -tantalates tantalates can only be prepared by fusion of the oxides with excess of alkali- metal hydroxide or carbonate and subsequent dissolution of the melt in water. [Although freshly precipitated hydrous Nb205 (but not Ta205) is soluble in concentrated alkali the precipitate ages very rapidly and becomes inert.] The pH of the resulting very alkaline solutions cannot be lowered by the direct addition of acid as this always results in the irreversible precipitation of the hydrous oxides. Indirect acidification such as addition of ammonium chloride and removal of ammonia with a current of or slow addition of carbon must be used. By such methods the pH of niobate solutions can be reduced to ca.7 and that of tantalates to ca. 10 before the onset of precipitation. There now seems little doubt that over these limited pH-ranges with the possible exception of the very alkaline region (to be discussed below) niobate and tantalate solutions contain a single polymeric species. Although both pentamers and hexamers have been suggested recent experimental evidence is overwhelmingly in favour of the hexamers [Nb601g]8- and [Ta601g]8-. Such evidence is provided by e.m.f. rnea~urements,4~~~~~ equilibrium ultracentrifuga- tion,4,6 light-scattering,4p5 and Raman spectro~copy54~~~ of both niobate and tantalate solutions. for the pentameric tantalate anion [Ta,0,6]7- were based upon ionic weights derived from diffusion coefficients a procedure now known to have no empirical or theoretical validity.57 Shortly after this work Lehn6 and his co-w~rkers,~~ on the basis of e.m.f.measurements and salt cryoscopy in potassium chloride at -10.7" suggested that an analogous pent- aniobate existed in the pH range 11-13.5. Later publications from the same laboratory favour the hexamer however.59 Several worker^^,^^,^^ have noted the protonation of [Nb,019]*- although the only quantitative measurements of equilibrium constants are (7) and (8) derived from e.m.f. measurements in ~M-KCI at 25°.52 [H2Nb6019]6- + [HNb,01g]7- 4- H+ [HNb6Ol9l7- + [Nb60,,Is- -I- H+ pK = 10.88 & 0.05 pK2 = 13-8 & 0.2 (7) (8) There was no evidence for any further protonation at pH ca. 10.7 the most acid solution studied. Somewhat similar mea~urements,~ in O-~M-KCI at 25" showed no protonation of ~a601,]s- in the pH range 10-13.In the course of his e.m.f. measurements of niobate solutions Neumann5* observed irreversible behaviour above pH 12-5 accompanied by some precipita- tion. Previously from salt cryoscopy Lehn6 and his co-workers,5s had suggested that the species [H2Nb207]2- [HNb2O,l3- and wbO4l3- were predominant in solutions above pH 13-5. However in such solutions 80% or more of the cryo- 68 G. Neumann Acta Chem. Scand. 1964,18,278. 69 G. Jander and D. Ertel J. Inorg. Nuclear Chem. 1960 14 71 77 85. 64 J. Aveston and J. S. Johnson Inorg. Chem. 1964 3 1051. 66 R. S. Tobias Canad. J. Chem. 1965 43 1222. 66 G. Jander and D. Ertel J. Inorg. Nuclear Chem. 1956 3 139. 87 L. C. W. Baker and M. T. Pope J. Amer. Chem.SOC. 1960 82,4176. 68 M. Lehnt and H. Goetz Bull. SOC. chim. France 1961 334; J. Leicht M. Lehnt and R. Rohmer ibid. 1963 213. 69 Y. Dartiguenave M. Lehnb and R. Rohmer Bull. SOC. chim. France 1965 62. 546 Pope and Dale scopic depression is due to free OH- ions. Consequently interpretation of the results depends upon precise [OH-] measurement in I-~M-KOH and is open to serious doubt. Moreover it seems unlikely that equilibrium was reached in these measurements. Very recently the Raman spectrum of a saturated solution of potassium niobate in ~ ~ M - K O H has been found45 to be similar to that in IM-KOH suggesting that the hexamer still predominates in the more basic solution. Unfortunately since the report gave no indication of attempts to reach equilibrium the question of the nature of niobate species in very alkaline solution must remain open for there is no reason to believe that the depolymerisation of hexaniobate will be other than very slow at room tem- perature.B. Structures.-With the exception of a few insoluble lanthanon niobates and tantalates MIIIXO (X = Nb Ta) which have the scheelite (CaWO,) structure and discrete tetrahedral oxyanions the co-ordination number of Nb and Ta in all known oxide structures is restricted60 to 6. (The high-temperature form of Nb20 has one Nb atom in 28 occupying a tetrahedral site the remainder beings1 octahedrally co-ordinated.) This is to be expected from simple radius- ratio considerations and argues against the likelihood of species such as NbOd- in aqueous solutions. Consequently although many niobates and tantalates are often formulated as salts containing xo43- X03- and X20,4- anions they actually have infinite mixed oxide structures with octahedrally co-ordinated Nb and Ta.Common structures are perovskite for MIXO, cassiterite for MI1(XO,) and MIIIXO, and atopite for M211X207.60 Apart from such mixed oxides which are generally prepared from melts several soluble alkali-metal niobates can be crystallised or precipitated from aqueous solutions. Only two incomplete X-ray structural investigations have been carried out on such compounds; Nal,Nb120,7,32H,062 and K8Ta,01, 16H20s3 were both found to contain octahedral groupings of six Nb or Ta atoms. From the metal-metal distances (3.3 A) the positions of the oxygen atoms were inferred and the symmetrical [M601g]8- structure illustrated in Figure 2d was suggested.This structure is closely related to that of the decavana- date ion. The niobate is therefore more realistically formulated as an acid salt Na7HNb601, aq. A normal salt K8Nb6ol,,16H@ can be isolated from more alkaline solutions and is isomorphous with the tantalate.* Crystals of both &Nb,& and K,Ta,O, may contain excess of alkali in the range K/Nb,Ta = 4.04-2 without changes being detected in the X-ray powder patterns,s4 and this has been responsible for some confusion in the early literature. The Raman spectra of solid K,Ta601,,16H,0 and K,HNb60,,,f3H20 are 6o R. W. G. Wyckoff ‘Crystal Structures’ 2nd edn. Interscience New York vol. 1 1963 p. 252; vol. 3 1965 pp. 23 361,439. *l B. M. Gatehouse and A. D. Wadsley Acta Cryst. 1964 17 1545. m I.Lindqvist Arkiv Kemi 1952 5 247. 68 I. Lindqvist and B. Aronsson Arkiv Kemi 1954 7 49. 64 F. Windmaisser 2. anorg. Chem. 1941 248 283; F. Halla A. Neth and F. Windmaisser 2. Krist. 1952 104 A 1961. 547 Isopoly- vanada tes -niobates and -tan tala tes relatively simple and are virtually identical to the spectra of saturated solutions of these salts in ~ M - K O H . ~ ~ ~ ~ There is little doubt therefore that the structures of these ions in solution are identical with those in the solid state. Soluble ‘metaniobates’ of potassium and sodium can be crystallised from weakly basic solution^^^^^ (e.g. by evaporation under an atmosphere of carbon dioxide). Although usually formulated as MNbO, aq. they are probably more correctly regarded as salts of [H2Nb60,g]6- in view of what is now known about the protonation of mb6019]8-.In this connection we may note that (a) the corre- sponding ‘soluble metatantalates’ have not been reported and (b) complete dehydration of a metaniobate renders it insoluble suggesting the presence of constitutional water in the anion. Anhydrous alkali-metal niobates have the pzrovskite structure.60 A soluble sodium orthoniobate N+NbO, has been prepared by fusing Nb20 with a large excess of NaOH and extracting unused alkali with absolute The resulting crystals are said to be yellow and to hydrolyse immedi- ately and irreversibly in solution yielding the hexaniobate anion. No structural investigation of this material has been reported. We thank Professor L. C. W. Baker for discussions concerning isopoly- and heteropoly-anions and Dr. H. T. Evans jun. for comments on the manuscript. This work was supported by the U.S. Air Force Office of Scientific Research. wj R. Rohmer and J. E. Guerchais Bull. SOC. chim. France 1961 324. 548

ISSN:0009-2681    

DOI:10.1039/QR9682200527

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

Decomposition Reactions of Radicals By J. A. Kerr and A. C. Lloyd DEPARTMENT OF CHEMISTRY THE UNIVERSITY BIRMINGHAM 15 DEPARTMENT OF CHEMISTRY THE UNIVERSITY LEEDS 2 1 Introduction Radical decomposition reactions have long been recognised as playing an important part in the complex mechanisms of pyrolysis combustion photo- oxidation polymer degradation and other processes. Thus in the radical-chain mechanisms for the decompositions of organic compounds originally suggested by Rice and Herzfeld; the chain-propagating steps involve radical decomposition reactions. In the pyrolysis of acetaldehyde for example the chain-propagating steps are CH,* + CH,. HO -+ CH + CH,-kO CH,*CO __+ CH,. + CO The free valence in the radical permits the decomposition to proceed with the formation of a multiple bond in the molecular product thereby reducing the endothermicity of the reaction.Accordingly bond strengths in free radicals are very much lower than in the corresponding molecules. It is apparent that for a complete understanding of complex processes such as pyrolyses it is necessary to have kinetic and thermodynamic data on the elementary reactions including radical decompositions. Quantitative information on these reactions has been slow in appearing in comparison with other radical reactions. A recent compilation of gas-phase unimolecular reactions together with subsequent data shows that about 50 rate constants for radical decompositions have been reported. The relative paucity of data is readily appreciated from a consideration of the difficulties involved. The activation energies for radical decompositions usually lie within the range 10-40 kcal.mole-l whereas those for radical-radical and radical-molecule reactions fall within the ranges 0-2 and 5-15 kcal. mole-l respectively. At temperatures convenient for measuring the rates of the decompositions (400-700"~) it is apparent that these other types of radical reaction will also be occurring and adding to the complexities of the system. It is also clear from the temperature range why most of the work on radical decompositions has been performed in the gas phase and the bulk of this Review deals with gaseous systems. 1 F. 0. Rice and K. F. Herzfeld J. qrner. Chern. Soc. 1934 56 284. 2 H. E. O'Neal and S. W. Benson Kinetic Data on Gas Phase Unimolecular Reactions' NSRDS-NBS U.S. Department of Commerce in the press.549 Decomposition Reactions of Radicals It is only comparatively recently that attention has been focussed on devising specific reaction systems suitable for studying radical decompositions. While useful information has sometimes been derived indirectly from studies initiated for different ends on the whole this is much less desirable than the direct approach. The purpose of this Review is to illustrate the kinetic approach to radical decomposition reactions by considering (i) the methods for studying these reactions (ii) the interpretation of the results in terms of thermodynamics and theories of unimolecular reactions and (iii) existing data on these processes. Discussion will be restricted to the decompositions of thermally equilibrated radicals as the decompositions of chemically activated alkyl radicals have already been reviewed? 2 Experimental Methods The problems involved in a kinetic study of a radical decomposition are best illustrated by reference to a specific example The initial requirement is to measure the rate constant kl which is defined by equation where square brackets denote concentration (mole ml.-l or mole l.-l) and R denotes rate of formation of product (mole ml.-l sec.-l or mole 1.-l sec.-l).Such a decomposition reaction is of course unimolecular and the rate constant kl is in first-order units usually sec.-l. The rate constants are then measured over as wide a temperature range as possible and the temperature dependence is assumed to obey the Arrhenius equation k = A exp (-E/RT) from which the Arrhenius parameters the A-factor and the activation energy are derived.As for any kinetic study of a radical reaction there are three problems to be solved (i) a suitable source of radicals is required (ii) the rate of the reaction must be measured and (iii) the concentration of the radicals must be determined although this is usually done indirectly. These problems are discussed below. A. Radical Sources.-(i) Pyrolysis. In general pyrolytic sources of radicals are not suitable for studying radical decomposition reactions. A few results on alkyl radicals have been derived from the pyrolyses of hydrocarbon^.^ Since however the reactions have to be carried out at very high temperatures the overall mechanisms are extremely complex and it is usually difficult to decipher un- ambiguous data on isolated elementary reactions.8 B. S. Rabinovitch and M. C. Flowers Quart. Rev. 1964 18 122; B. S. Rabinovitch and D. W. Setser Adv. Phofochem. 1964 3 1. J. H. Pumell and C. P. Quinn Proc. Roy. SOC. 1962 A 270 267; C. P. Quinn Trans. Faraday SOC. 1963,59,2543. 550 Kerr and Lloyd An exception to this general rule is the formation of alkoxy-radicals from the decomposition of peroxides5 or nitrites,g e.g. [(CH3)2CHO]2 + 2(CHa)&HO* (CH.J,CHONO __+ (CH3),CHO- + NO These reactions occur at relatively low temperatures since the 0-0 and N-0 bonds are weak. Occasionally information on radical decompositions has been obtained from toluene-carrier studies of the pyrolyses of molecules. Thus the rate of the decomposition CH2 0 CH2Br - has been reported from a study of the pyrolysis of dibromo-p-xylene.' (ii) Photolysis.The photodecompositions of aldehydes ketones and acetates have been extensively used as sources of alkyl,8 acylYg and alkoxy-radicals:1° RCHO + hv+ R. + CHO RCOR1 + hv + RCO. + R.' RCO,R1 + hv + RCO. + R1O* Unfortunately these are not the only primary photochemical modes of decom- position,ll as other radical and molecular elimination reactions can occur particularly with compounds containing H atoms in the y-position relative to the carbonyl group e.g. CH3CO*CH2.CH2*CH3 + hv __+ CH3.CO.CH3 + C2H4 Another complication arises from 'down-chain' methyl-radical splits :g CH3.CH2.CH2.CH2.CHO + hv + CH3 * + CH2*CH,*CH2CHO The generation of these additional radicals and molecules leads to complications in studying the reactions of the radicals of primary interest.In the case of aldehyde photolyses this difficulty is somewhat alleviated since at higher temperatures the majority of the alkyl radicals in the system are produced in a thermal chain sequence R. + R*CHO + RH + RCO RCO -+R* f C O 5 M. J. Yee Quee and J. C. J. Thynne Trans. Faraday SOC. 1967 63 2970. 6 D. L. Cox R. A. Livermore and L. Phillips J. Chem. SOC. (A) 1966 245. 7 M. Levy M. Szwarc and J. Throssell J. Chem. Phys. 1954 22 1904. 8 J. A. Kerr and A. F. Trotman-Dickenson Progr. Reaction Kinetics 1961 1 107. 10 M. H. J. Wijnen J. Amer. Chem. SOC. 1960 82 3034. 11 R. B. Cundall and A. S. Davies Progr. Reaction Kinetics 1967 4 149. J. G. Calvert Chem. Rev. 1959 59 569. 551 Decomposition Reactions of Radicals Although a lot of work has been done on radical decomposition reactions from direct photolyses of carbonyl compounds the results on the whole have not been entirely unambiguous and these compounds have lost favour to the azo- compounds which photodissociate cleanly to give alkyl radicals :12 RN=NR + hv + 2R* + N2 With 3660 A radiation there appears to be only one major primary split and there is no evidence for the formation of activated radicals.One minor dis- advantage is that it is not always possible to work at as high a temperature as one would like owing to the onset of the pyrolysis of the azo-compound and it is not clear what effect this has on the photolytic system. (iii) Photosensitised decompositions. The mercury-photosensitised decom- positions of alkanes was the first method to yield quantitative results on the decompositions of alkyl radi~a1s.l~ Mercury vapour is excited to the triplet state by 2537 A radiation from a lowpressure mercury arc and sensitises the decompositions of alkanes and other hydrogen-containing compounds C3H8 + Hg(3P1) + Prn + Ha + HgPS,) -+ Pri 3.H. + Hg(lS,) As this example shows a mixture of radicals is formed when there is more than one type of C-H bond in the parent molecule. The proportions of radicals are mainly determined by the dissociation energies of the C-H bonds. It is possible to study the decomposition of one of the radicals provided its decomposition product is not formed by other reactions. With the development of gas chromato- graphy there has been a revival in the application of mercury-photosensitisation as a method of generating radicals for studying decomposition rea~ti0ns.l~ A considerable merit of the method is that it lends itself to extensive variation in the overall pressure of the system which as will be shown later is important in studying the decompositions of smaller radicals.To overcome the difficulties of multiple radical formation which plague direct photolyses Calvert9 introduced the radical-sensitised decomposition of aldehydes for producing acyl and alkyl radicals. Thus azomethane has been selectively photolysed (3660 A) in the presence of propionaldehyde to yield the propionyl radical :15 The method is based on two properties of aldehydes (i) they are transparent to light of 3660 A and (ii) the acyl-H atom is much more rapidly abstracted than la D. H. Slater S. S. Collier and J. G. Calvert J. Amer. Chem. SOC. 1968,90,268.S . Bywater and E. W. R. Steacie J . Chem. Phys. 1951 19 319; 326. l4 L. F. Loucks and K. J. Laidler Canad. J . Chem. 1967 45,2767. J. A. Kerr and A. C. Lloyd Trans. Faraday SOC. 1967 63 2480. 552 Kerr and Lloyd the alkyl-H atoms. The decomposition of aldehydes can also be sensitised by the selective photolysis of chlorine or bromine.16 Radical- or atom-sensitised decompositions of aldehydes are not particularly suitable for extensive pressure studies but the method has the considerable merit that the radicals are formed in a thermal reaction and are consequently thermally equilibrated. The decompositions of alkanes sensitised by methyl radicals have also been used in studying alkyl-radical decomposition^.^^^^* As with mercury photosensitisation there is always the problem of generating mixtures of radicals when the alkane contains different types of C-H bond.(iv) Additions to olefins. Hydrogen-atom addition reactions to olefins the reverse of radical decomposition have been widely applied for the production of alkyl radicals These reactions are highly exothermic (ca. 40 kcal. mole-l) and give rise to chemically activated radicals. Rabinovitch and his co-workers3 have made extensive studies of the decompositions of chemically activated alkyl radicals and have used the results to test the various theories of unimolecular reactions. Recent studies of the additions of NF radicals to olefins have yielded information on the decompositions of the adduct radicals e.g. CH,-CH(NF2).CH.CH3 radicals from the addition of NF to but-2-ene.19 Likewise the rates of decomposition of chloro- and bromo-alkyl radicals have been determined from studies of the photochemical chlorination and bromination of olefins.20 In these systems the information on the radical decompositions has been derived from a detailed kinetic treatment of the overall mechanisms and these studies come within the category of indirect determinations.B. Measurement of the Decomposition Rate.-Of the three problems to be solved in measuring the rate constant of a radical decomposition the rate of the reaction is the least difficult to solve. This is achieved by following the rate of formation of the product molecule; for the ethyl decomposition [reaction (l)] the rate of formation of ethylene is measured. Ideally the rate of formation of H atoms should also be determined and there should be agreement between the two rates.Since the atoms and radicals can undergo several reactions to yield products that may also be formed in other ways it is seldom possible to carry out this check. It is of course extremely important to ascertain that the molecular product 16 J. C. Amphlett and E. Whittle Trans. Faraday SOC. 1967 63 80. 17 M. C. Lin and K. J. Laidler Canad. J. Chem. 1966 44,2927. 18 A. S. Gordon S. R. Smith and C. M. Drew J. Chem. Phys. 1962 36 824. 1s A. J. Dijkstra J. A. Ken and A. F. Trotman-Dickenson J. Chem. SOC. (A) 1967 105,864. 2O G. Huybrechts L. Meyers and G. Verbeke Trans. Faraday Soc. 1962,58,1128; H. Schmitz H. J. Schumacher and A. Jager Chem. Ber. 1942 B 51,281. 553 Decomposition Reactions of Radicals from the radical decomposition is not formed in other reactions or to make corrections where such reactions are known to occur.For ethyl radicals the rate of formation of ethylene from the decomposition has to be corrected to allow for the disproportionation although this correction is quite small. Ambiguity in product formation is much less of a problem with oxygen-containing radicals than with alkyl radicals. It has frequently been pointed out that secondary radicals generated by hydrogen-atom abstraction from the parent molecule could yield the same decomposition product as the principal radicals. This point is illustrated by the photolysis of azo-n-butane for studying n-butyl decomposition 2C,H5* + C2He + C2H4 C,H,.N=N.C,H + hv + 2C4H,* + N2 C4H9* +C2H,* +C2H4 It is conceivable that the radical *CH2CH2CH2CH2-N = NC,H, from radical attack on the parent azo-compound could also decompose to give ethylene.It has been argued that since the formation of the latter radical involves abstraction of a primary C-H atom it is much less likely to be formed than those from the abstraction of the secondary C-H and consequently it should not be an important source of ethylene?' C. Measurement of the Radical Concentration.-Since the steady-state concen- trations of radicals in the types of system described above rarely exceed 10-lo mole.ml. -l direct determination of the concentration is not possible. Con- siderable success has recently been achieved in directly determining atom concentrations in gaseous systems by applying electron spin resonance If the difficulties in interpreting the extremely complex e.s.r.spectra of gaseous radicals can be resolved,2s direct measurement of radical concentrations would be possible and absolute rate-constant measurements would assume a different dimension. At present the difficulty of measuring the radical concentration is circumvented by measuring the rate constant for the decomposition relative to the rate constant of some other reaction of the radical. Frequently the reference reaction is the formation of the dimer from radical combination 2C,H,. + n-C4Hl For reactions (1) and (2) we have and thus Rc*& = kl CC2H,* 1 and Rn-C& = k [C2H,* l2 klIk2') = RCIHI/Rn-c~,' 3 a1 W. E. Morganroth and J. G. Calvert J. Amer. Chem. SOC. 1966,88,5387. t a J. A. Ken Ann. Reports 1967,64,72; S . W. Benson and W. B. DeMore Ann. Rev. Phys. Chem.1965,16 397. aa A. Carrington Proc. Roy. SOC. 1968 A 302 291. 554 Kerr and Lloyd To obtain kl it is then necessary to know k2 the rate constant for the combination of the radicals. Comparatively few rate constants for radical com- bination reactions have been measured. The most commonly used method is the rotating-sector technique which determines the lifetime of the radicals. Although the method does not provide precise values of these rate constants it seems clear that for small alkyl radicals the rate constants for combination are very high approaching the collision rates. Rotating-sector measurements of radical combinations cannot be carried out over very wide ranges of temperature but the few results available indicate small or zero temperature coefficients. There is also some indirect evidence from cross- combination ratios for pairs of different radicals which supports the contention that E = ca.0 kcal. mole-l for radical combinations. For the combination of methyl radicals the rate constant has been determined to be k = (mole-1 ml. sec.-l) independent of temperature:* and this value is often assumed for other radical combinations. In the photo-initiated chain-decompositions of aldehydes* at temperatures above about 600"~ the rates of formation of the radical dimers become very slow in comparison with those of other products. Consequently at these tem- peratures it is not possible to achieve a measurable rate of formation of dimer at low percentage conversions of the aldehyde. The concentrations of the radicals have been measured from their reaction with the parent aldehyde R.+ RCHO -+ RH + RCO (3) for which we have RRH = k,[R*][R*CHO] The values of k3 are extrapolated from results at lower temperatures where radical dimer formation was observed and hence from the measured rate of formation of alkane RRH and the known concentration of aldehyde [RCHO] it is possible to calculate [R-1. The calculated values of [R-] are however sensitive to the predicted values of k, and quite large errors can be introduced in determining the rate constants for the radical decomposition reaction. 3 Thermodynamics of Radical Decompositions The kinetics and thermodynamics of radical decomposition reactions and the reverse radical- or atom-addition reactions are related via the equilibrium constant (4) (-4) R $ r + M (4) From kinetics the equilibrium constant is given by the ratio of rate constants 24 A.Shepp J. Chem. Phys. 1956,24,939. Decomposition Reactions of Radicals and the temperature-dependence of the rate constants has been assumed to obey the Arrhenius equation. From thermodynamics the equilibrium constant is related to the standard free-energy change - RTIn Kp = AGO and since AGO = AHO - TAS" where AH" and AS" are the standard enthalpy and entropy changes it follows that K4,- = (A4/A-,) exp [- (E4 - E-JRT] = exp (ASOIR). exp (-AHo/RT) AH" = E - E- and ASo = R In ( A 4 / L 4 ) By equating coefficients we have so that activation energies are considered in terms of the enthalpy change and the A-factors in terms of the entropy change. A. Activation Energies.-When there is an increase in the number of moles during the reaction (An) the relation between AH" and the activation energy difference (AE) is modified owing to the change in standard states AH" = AE i- AnRT and hence for radical decompositions where An = + 1 AH" = AE4 + RT= E4 - E-4 + RT The RTcorrection which is usually ca.1 kcal. mole-l allows for the fact that enthalpies are measured at constant pressure whereas activation energies are measured at constant volume. It should also be noted that AH" defines the bond dissociation energy D(r-M) in the radical R (equation 4). The application of these relations will be illustrated by considering the de- compositions of the n-propyl radical n-C3H7- + CH3* + C2H n-C3H7- + He 4- C3H6 AH" = D(CH3-CH,.CH,.) = A Hf"(CH3.) + AHf"(C2H4) (5) (6) The enthalpy changes can be calculated from the thermochemical equation - A Hi "(n-C3H,-) - AHf"(n-C3H,-) AH" = D[CH3*CH(CH2)-H] = AHf"(H.) + AHf"(C3H6) Bond-dissociation energies in a variety of radicals have been calculated in this way since the enthalpies of formation of many atoms and radicals are When this information is not available reasonable estimates of the S.W. Benson J. Chem. Educ. 1965,42,502; J. A. Kerr Chem. Rev. 1966,66,465. 556 Kerr and Lloyd enthalpy data can be made by the method of additivity of bond or group pro- perties whereby thermodynamic properties are taken to consist of individual additive contributions of the component parts of the molecules.2s This method of estimating enthalpies entropies and heat capacities has met with considerable success and is now widely accepted as yielding generally reliable estimates of these properties.For the n-propyl decompositions the thermochemical cal- culations yield D(CH,-CH,CH,.) = 26 and D [CH,CH(CH,)-HI = 36 kcal. mole-l with estimated errors of about f 2 to f 3 kcal. mole-l. The activation energies from kinetic studies of the radical decompositions can now be compared with these enthalpy changes calculated from thermochemistry if the activation energies for the reverse radical- and atom-addition reactions [(-5) and (-6)] are known. A growing body of information on these types of addition reaction is available,22 and when direct measurements have not been made it is often possible to obtain a reasonable estimate of the activation energies from analogous reactions. The relations between the enthalpy changes and activation energies for the n-propyl decompositions are shown in Figure 1.E- 6 I I I 1 I I I I I I I -. - CU3-CH2CH2 Y CH k-CH 7 + 3 N' I H-C3H6 \ I 1 I I I I I I 1 I \ I I 1 I I I I c 36 \ zd=39 I I I I I I 2 CH3.Ct- '. P 2 Figure 1 Energy diagram for the decomposition of the n-propyl radical n-C,H,. __+ CH,. + C,H ( 5 ) and n-C,I-X,- + H. + C,H (6) 26 S. W. Benson and J. H. Buss J . Chem. Phys. 1958 29 546. 557 Decomposition Reactions of Radicals A correction should be applied to the activation-energy difference to allow for the fact that thermodynamic enthalpy data refer to 298"~ whereas the activation energies will have been measured at higher temperatures. This correction corresponds to the net enthalpy change in decreasing the temperature of the system from the mid-point of the experimental temperature range in the kinetic experiments to 298"~ and is given by - jf ACpdT where A C is the heat-capacity change for the reaction.The complete expression relating the enthalpy change to the activation energy diRerence becomes AH" = E3 - E-3 + RT - ST ACpdT For most radical decompositions the temperature correction is small and is largely offset by the RTcorrection for the change in the number of molecules. Consequently it is usually adequate to equate the two correction terms and identify AH" with the experimentally observed difference in activation energies. Tn view of the errors involved in the calculated values of dH" and in determin- ing the activation energies the agreement between AH" and (Ef - E,) would not be expected to be much better than f2 kcal. mole-l.B. A-Factors.-Precise statistical mechanical calculations of entropies of free radicals are usually not possible since the required information on the structures and vibrational and rotational properties is not available. Reasonable estimates can be made however since the major contributions to the total entropy are from translational and electronic terms that can be evaluated precisely and the largest uncertainties reside with the minor terms for vibrational and rotational contributions. The entropies of methyl ethyl n-propyl and isopropyl radicals have been cafculated over the temperature range 200-1000"~ by use of con- sistent vibrational and rotational models.27 These calculated values were considered to be good to 30.5 cal. mole-l deg.-l for CH3. and =tl cal. mole-l deg.-l for C2H5.and C3H,.. As previously mentioned an alternative approach for calculating entropies involves the method of additivity of bond or group properties and entropies of radicals calculated by this method should not be in error by more than &1 cal. mole-l deg.-l. Less precise estimates (f2 cal. mole-l deg.-l) can be made by taking the values of entropies of closely related molecules and making corrections for symmetry changes and the electronic entropy of the radical. The entropies of radicals derived by the above methods are usually tabulated €or a standard state of 1 atm.,2P2* whereas comparisons of ratios of A-factors with AS" require entropies corresponding to a standard state of 1 mole ml.-l or 1 mole 1-l. The entropies in pressure units (S,) are converted to concentration units (Sc) by the equation 21 J.H. Purnell and C. P. Quinn J. Chem. SOC. 1964 4049. J. G. Calvert and J. N Pitts 'Photochemistry' Wiley New York 1966 p. 819. 558 Kerr and Lloyd where the last term is the entropy increase in expanding the gas. It follows that Sop = (Soc + 20.1) or (Soc + 8.2) cal. mole-I deg.-l when the concentration units are mole ml.-l and mole 1.-l respectively. To illustrate a typical comparison between AS" calculated from thermo- dynamic entropy data and from experimental A-factors we shall consider the decomposition and formation reactions of the acetyl radical (7) (-7) CH,CO + CH,. + CO The required entropy data at 2 9 8 " ~ is S"(CH,-) = 46.4 S"(C0) = 47.3 and S"(CH,*CO) = 64.5 cal. mole-l deg.-l hence AS" = 29.2 cal. mole-I deg.-l (standard state 1 atm.) or 9-1 cal.mole-I deg.-l (standard state 1 mole ml.-I). For the decomposition reaction A = 1010'3 (sec.-l) while for the reverse radical formation reaction A_ = 1PDM (mole-l ml. sec.-I) leading to a kineticestimate of AS" = 7.9 cal. mole-I deg.-l (standard state 1 mole rnl.-l) which is in reason- able agreement with the thermodynamic data. The results are typical of the agreement to be expected in these comparisons. C. A-Factors and Transition-state Theory.-The A-factor for a unimolecular reaction is given from transition-state theory by the expression A = (ekT/h) exp (AStlR) where k and h are Boltzmann's and Planek's constants respectively and is the entropy of activation. For radical decompositions the assumption is some- times made that the structures of the initial and transition states are similar and hence dSt = 0 giving A = ekT/h = (sec.-l) at 6 0 0 " ~ .Since the overall entropy change for radical decompositions AS" is positive it seems likely that dSt 3 0 and consequently estimates of A-factors based on the assumption AS1 = O are lower limits. The application of transition-state theory to these systems comes down to estimating AS the entropy of activation. O'Neal and BensonZ9 have recently developed a method of estimating AS1 for four- and six-centre unimolecular reactions based on empirical assignments of bending stretching and torsional frequencies. The same principles have been applied to radical decompositions.e For the decompositions of the n-propyl radical [reactions (5) and (6)] values of AS$ = 0.8 and ASS = 0.2 cal. mole-l deg.-l have been estimated leading to A = 1013'8 and A = sec.-l at 600"~.While such calculations do not give precise values of A-factors they nevertheless form a useful basis for examining the general validity of experimental A-factors. 4 Theories of Unimolecular Reactions Since radical decompositions are unimolecular reactions they can be treated by the various theories of unimolecular reactions. This section will deal with 29 H. E. O'Neal and S. W. Benson J. Phys. Chem. 1967,71,2903. 559 Decomposition Reactions of Radicals the applications of the theories to these reactions rather than with the detailed developments of the theories. One of the early successes of unimolecular rate theory was the prediction of the decrease of the apparent unimolecular rate constant with a corresponding change from first- to second-order kinetics as the pressure in the system is decreased.As the theories of unimolecular reactions developed it became clear that the pressure region at which this fall-off behaviour could be observed depends on two things (i) the number of vibrational degrees of freedom of the reacting molecules and (ii) the term E/RT where E is the critical energy needed to bring about the reaction which is usually identified with the experimental activation energy. Experiment has confirmed the theoretical deduction that the pressure-sensitive region shifts to lower pressures as the size of the reacting molecules and the activation energy increase. Benson30 has compiled a useful table based on the Rice-Ramsperger-Kassel (RRK) theory of unimolecular reactions whereby the pressures at which the unimolecular rate constants fall off markedly from their high-pressure limits can be readily estimated from a knowledge of the activation energy of the reaction and the number of vibrational modes of the reactant species.It has already been pointed out that bond strengths in radicals are much lower than in the corresponding molecules and consequently the activation energies for radical decompositions are considerably lower than for molecular decompositions. This lowering of the activation energy results in a considerable increase in the pressure at which the rate constants for radical decompositions show pronounced falling-off from their high-pressure values. In a recent study of the thermal decomposition of ethane,3l information was obtained simultaneously on the reactions (8) and (9) where the activation energies are E = 86 and C,H +2CH3.E9 = 38 kcal. mole-l. At 9 5 8 ” ~ the pressures at which the first-order rate constants k and k fell to half their high-pressure values were approximately 20 and 400 mm. respectively. It is clear that pressure-dependent kinetics are much more of a problem with radical decompositions than with molecular decompositions. Although several instances of pressure-dependence of unimolecular rate constants for radical decompositions had previously been reported the first attempt to make quantitative measurements on this type of system wasby O’Neal and Benson who studied the effect of pressure on the decomposition of the acetyl radical.32 The results were treated by the Lindemann-Hinshelwood (LH) theory 80 S.W. Benson ‘Foundations of Chemical Kinetics’ McGraw-Hill New York 1960 p. 234. 81 M. C. Lin and M. H. Back Canad. J. Chem. 1966,44,2357. sa H. E. O’Neal and S. W. Benson J. Chem. Phys. 1962,36,2196. 560 Kerr and Lloyd The asterisk denotes an energised molecule i.e. a molecule containing sufficient vibrational energy to react. Following the LH treatment first-order kinetic behaviour of the decomposition reaction can be explained by assuming that there is a small stationary concentration of energised molecules that is the rate of de-energisation [reaction ( l l ) ] is rapid compared with the rate of de- composition [reaction (12)]. From a steady-state treatment of this simple mechanism it can be shown that the rate of the radical decomposition Rco is given by (13). In expression (13) it is assumed that M corresponds to the total pressure in the system and there is usually no attempt to allow for the varying efficiencies of different molecules in the deactivation reaction.At high concentrations i.e. high pressures kll[M] >> k, and the rate becomes where k is the high-pressure limiting rate constant and the reaction obeys first-order kinetics. At low pressures the opposite condition applies kll[M] < k12 and the rate is given by R, = klo[M][CH3.dO] = ko[M][CH3*i'Ol where k is the low-pressure limiting rate constant and the reaction now follows second-order kinetics. The observed first-order rate constant (k') can be defined as R, = k'[CH,.CO] (14) and from equations (13) and (14) it can be shown that Ilk' = kll/(kl&l2) + l/(kl,[MI) It follows that when Ilk' is plotted against 1/[M] for a fked temperature there should be a linear relation with the intercept equal to l/k and the slope equal to l/k,.A series of such lines is then obtained by making measurements over a range of pressures for different temperatures and the values of k and k at each temperature are found from the intercepts and slopes. A typical set of data are shown in Figure 2. An Arrhenius treatment of the rate constants 561 Decomposition Reactions o f Radicals 5 10 106/[M] (ml. mole-') Figure 2 Lindemann-Hinshelwood plots showing the efect of pressure on the decomposition of the propionyl radical C.&CO. + M -+ C,H,* 3- CO + M; F - function equivalent to Ilk' [From J. A. Ken and A. C . Lloyd Trans. Faraday SOC. 1967,63,2480] k' yields Em and E, the limiting high- and low-pressure activation energies and A and A, the limiting high- and low-pressure A-factors.Most of the results given in section 5 for the limiting high- and low-pressure Arrhenius parameters of radical decompositions have been derived by the above method The shortcomings of the LH theory are well known. Although the LH plots described above are remarkably linear for the radical decomposition reactions so far studied this is only because the pressure ranges in most of the systems are very limited. For other unimolecular reactions where it has been possible to carry out large pressure variations the LH plots deviate considerably from linearity at high pressures. This deviation arises from the deficiency of the theory to take account of the effect on the rate constants of differing energy contents of the energised molecules.The effect of this deficiency is worse on the values derived from the slopes of the LH plots (k,) than those from the intercepts (km). Consequently the derived values of Eo are not so reliable as those for Em. From the thermochemical point of view this is not so serious since it is the high-pressure limiting activation energy which is related to the enthalpy change of the reaction. 562 Kerr and Lloyd An alternative procedure for extrapolating to k and k involves plotting l/k’ against l/[hl]* and gives better linear relations over wide pressure ranges for molecular reactions.% Such extrapolations have been made for the de- composition of the C2H,- radical,34 where the pressure rangc was considerably larger than is normally possible for radical decompositions.The C2H,. radicals were generated by mercury-photosensitisation which appears to be the most useful source of radicals for studying the effect of pressure variations on the rate constants. In studying the fall-off behaviour of the rate constant for a thermal uni- molecular reaction the normal procedure is to obtain the high-pressure limiting rate constants experimentally and to observe the decrease in the rate constants over a pressure range varying by a factor of several hundreds from the high- pressure limit. So far it has not proved possible actually to measure the high-pressure limiting rate constants for any of the radical decompositions that exhibit fall-off behaviour. Accordingly all the measurements in these pressure-sensitive systems have been made in the fall-off region.The difficulty of reaching the high-pressure limit is a practical one; the pressure range may be limited by the volatility of the radical source or the addition of a large pressure of third-body molecules interferes with the analysis. Clearly thermal decom- positions of radicals are not the most suitable reactions for testing theories of unimolecular reactions although it should be pointed out that much valuable information has been derived from studies of the decompositions of activated radicals from the additions of H atoms to 01efins.~ The Rice-Ramsperger-Kassl (RRK) theory and the extension by Marcus (RRKM) have been applied to the results on the decompositions of the C,H,- and CH,.O-CH radicals where the ranges of pressure were sufficient to warrant the application of the more elaborate theories.4,14,31,34,35 These theories have been extensively treated e l ~ e w h e r e ~ ~ ~ ~ ~ ~ ~ and only a brief outline of their relevance to the present reactions will be given here.The simple LH mechanism is modified as follows (17) A* -+ AS kx A2 + product where a distinction is made between energised molecules A* and activated complex molecules At that have attained the transition-state configuration. ss E. W. Schlag and B. S . Rabinovitch J. Amer. Chem. SOC. 1960 82 5996. 34 L. F. Loucks and K. J. Laidler Canad. J. Chem. 1967 45 2795. 35 M. C. Lin and K. J. Laidler Trans. Faraday SOC. 1968 64 79. 38 K. J. Laidler ‘Chemical Kinetics’ McGraw-Hill New York 1965 p. 150. 563 Decornposition Reactions of Radicals From a molecular model consisting of a system of loosely-coupled harmonic oscillators the RRK theory gives for the apparent unimolecular rate constant e-x dx xu-1 1.k,[Al b+x 1 + kS (,L) s-l kt e4*/kT k‘ = (s - l)! where x = ( E - c*)//iT 6 = E*/kT and the terms are defined as follows the rate constants refer to the above general mechanism k = Boltzmann’s constant E = total energy content of the reacting molecule E* = critical energy for the reaction s the number of effective oscillators is related to the total number of vibrational modes (3n - 6 where n is the number of atoms) and the remaining terms have their usual meaning. For a given value of s the integral corresponds to a particular variation with the concentration [A]. The normal procedure is to determine empirically which value of s predicts the observed variation of k’ with pressure.The integrals can now be evaluated readily with the aid of com- puter programmes. A typical RRK plot is shown in Figure 3. The least satisfactory feature of the RRK theory is the arbitrary nature of the parameter s -6,’O I I 0 1.0 2.0 3 0 109 10 p (mm.) Figure 3 Rice-Ramsperger-Kassel plot of log, ($unction equivalent to k‘) against log, P for k‘ decomposition of ethyl radical C,H,* + M + Ha + C,H + M; circles are experimental results curves are calculated with S = 7 8 and 9 [From L. F. Loucks and K. J. Laidler Canad. J. Chem. 1967,45,2795] which although usually corresponding to half to two-thirds of the total vibrational modes cannot be found apriori. The results of the decompositions of the C2H,. and .CH,.O.CH radicals generated by mercury-photosensitisation have recently been treated by the 564 Kerr and Lloyd RRKM theory.% This theory is undoubtedly the most satisfactory approach to unimolecular reactions.The RRK model consisting of a system of classical oscillators is replaced by a model which takes into account the individual molecular vibrations and internal rotations and also the zero-point energy levels. In the high-pressure limit the RRKM theory incorporates the transition- state description of the unimolecular decomposition but discards the assumption inherent in the RRK theory that the partition function of the activated complex is unaffected by the distortions of the bonds. For thermal reactions the procedure in applying RRKM theory is analogous to that for RRK theory; the fall-off behaviour of the rate constants is compared with the theoretical curves based on different molecular models but this time without the adjustable parameter s.As seen from Figure 4 the fall-off curves calculated for the C,H,- decomposition by the RRKM theory give excellent fits with the experimental data on the basis of a model involving active rotations in the energised radical and the activated complex. n .r* 8 2 W 0.2 - 0.4 - 0.6 - 0.8 - 1.0 - 1.2 - 1*4 - Figure 4 Rice-Ramsperger-Kassel-Marcus plots of log, (k'lkm) against log, P for the k' decomposition of ethyl radical C,H,. + M + H- + C,H4 + M; circles are experimental results curves are theoretical based on two models [From M. C. Lin and K. J. Laidler Trans. Faraday SOC. 1968 64 791 It is clear that the RRKM theory will be more widely applied as data on radical decompositions becomes more extensive.5 Results on Radical Decompositions Quantitative data are listed in the Table which is intended to summarise existing information rather than give all available results. Where more than one study 565 2 Table Kinetic and thermodynamic data on radical decompositions a\ Reaction Radical source Temp. E log A Ref. (sec.-l and (kcal. mole-l (OK) mole-l) ml. sec.-l) Alkyl Radicals C2H5. -+ H a + C2H4 Hg(3P1) + C2H6 673-773 40.9 (Em) 14.4 (Am) c 31.8 (E,) 17.8 (A,) H- + C2H4 -+ C2H5- 3.0 13.6 b 3 7'-tH. + C3H6 CD3. + CH,*CD,.CH 569-693 37.0 14-1 d H 4- C3H6 -+ n-C,H 3.0 13.6 e 3 7' -+ CH3. + C2H4 CH3. + C3HS 533-573 31.4 13.5 .f {"-C {n-c { { {n-c CH H + C,H -+ n-C4H,.i n-C,H,CHO - - - CH,.+ C,H -+ n-C,H,- 7.9 11.9 g Photolysis 673-773 36.9 13.8 h iso-C,H,- -+ H- 4 C3H H* + C3H6 -+ iso-C,H,. C,H,*CHO 2-6 13.5 b iso-C3H,- -+ CH,. + C2H4' Pyrolysis 713-814 34.5 12.0 j CH,. + C,H -+ iso-C,H,.r (CH31,CH - - - IC2H5. + C2Hl -+ n-C,H,. (n-C4H9N)2 8.6 12.1 1 Photolysis 432-520 28.7 13.6 k Jn-C,H,. + C2H5. + C,H4 CH3* + C3H6{ Photolysis 571-689 27.1 12.1 m - n-C,H,. -+ H2 + CH3CH*CH=CH2 CD,. 1- alkanes 523-773 -30 n sec-C,H -f CH,. + C,H CH,. + n-C,H, 533-613 32.6 14.8 0 CH3. + C3H6 -+ sec-C,H,- 7-4 11.5 g iso-C,H,- -+ Ha + iso-C4H Photolysis 602-691 31.0 13.0 p Ha + iso-C,H -+ iso-C4H,- iso-C,H,.CHO 1.4 13.6 q { P AH"2gS ASo AH0298 AS" 8 3 'cl (cal. (cal. 2 (kcal. mole-l (kcal. mole-l 2 mole-') deg.-l)e mole-') deg.-l)a Kinetic data datab f@ Thermodynamic 3-7 38.9 0.2 9 3.37.9 2 34-0 2.3 36.3 3.9 2 % s 23.5 7.3 25.8 11.6 2; 34.5 1.4 39.3 3-7 - - 28-9 13-4 7.1 22.5 13.6 20.1 - 14.8 5.3 25.2 15.1 26.2 13.8 - 29.6 -2.8 34.4 3-8 Table-continued Reaction Jiso-C,H,. -f CH,. + C3H XCH,. + C3H -+ iso-C,Hg* t-C,H,- -+ H. + iso-C,H He + iso-C,H -+ t-C4H,* t-C4H9.+ CH3. + C,H,* CH3. + C3H6 -+ t-C,Hg*i { { neo-C5H11- -+ CH3. + iso-C,H { CH3* f iso-C4H -f neo-C5Hll. JcYc~o-C~H,. C3H5 + CZH4 IC3H5. + CgH4 -+ CYC~O-C~H,. sec-C,H, -+ n-C3H,- + C,H n-C3H,- 4- C3H -+ n-C,H,3. L IC,H,Cl+ CIS + C2H4 XCl. + C2H4 -+ C,H4Cl v CH,-ClCHCl -+ CI*+CzH,CI 3 { C1. + C2H3Cl + CH2ClCHCI Radical source Photolysis (~SO-C,H~N)~ Photolysis t-C,Hg*CHO Photolysis t-C,H,.CHO HCl catalysed decompn. neo-C,H, CD3. + CYC~O-C~H~ Temp. E log A Ref. dH029B AS" AH029e As" (kcal. mole-' (kcal.mole-l Thermodynamic (set.-' (kcal. mole-' and (cal. (cal. (OK) mole-l) ml. sec.-l) mole-l) deg.-l)a mole-') deg.-lP Kinetic data datab Alkyl radicals Decompn. C,H, addition C2H5. to CzH4 and isomeriza- tion n-C,H,,* -+ sec-C6HI3 543-598 31.0 8.8 666-797 43.6 1.4 742-797 46.3 - 762 34 -8.3 573-718 37.7 -10.3 822 22.4 -8.6 Halogeno-radicals Photochlorination 298-327 23.9 of C2H4 1.5 Photochlorination 298-327 23.8 of CzH3C1 0.9 12.4 11-4 15.5 13.6 15.2 - 13 -1 1.4 14.5 -10.5 13.5 -12.1 r 22.2 4.6 25.1 16.1 b s 42-2 9.1 41.3 7.4 4 S 32.0 19.7 - - - t 25.7 7-4 24.0 17.3 b u 27.4 18.3 28.7 21.8 b v 13.8 6.3 22.9 15.9 b .Y w 22.4 -1.4 21.3 3.7 X % 13 13.3 13 w 22.9 12.4 X A e 5.5 2 2.7 22.7 Table-continued % Reaction CHCI,*CHCI + C1. + CHCI=CHCI 7 C1. + CHCl=CHCl 1 + CHC1,CHCI I CHCI=CC12 7 CI.+ CHCI=CC12 1 + CHCI,.CCI CHCl2.CC12 -+ CI. + C,C15* + C1- + C,C14 C1- + C,C14 -+ C,Cl -C,H4Br + Br. + CzH4 Br. + C,H4 -+ .C2H4Br *CH,-C,H,.CH,Br -+ Br.CH,Br. { + CHz :C6H4 :CH P s c Radical source Temp. E log A Ref. AH0298 ASo dH"29 AS" b (sec.-l and (cal. (cal. 2 (kcal. mole-' (kcal. mole-l (kcal. mole-l 2 (OK) mole-') ml. sec.-l) mole-l) deg.-l)= mole-l) deg.-')a 5' Thermodynamic Kinetic data datab F 8 Photochlorination 313-335 20.3 13 w 19.3 2.3 20-1 4.9 s* 1-0 12.5 X % 8 Halogeno-radicals of CHCl = CHCl 2Y Q 7.6 Pho tochlorination 433497 20.4 13.7 y 19.7 5.0 19.8 of C,HCI 2;. 0-7 12.6 X Photochlorination 300-388 16.8 12.8 b 16.8 1.8 17.5 6-3 of C2C14 4 12.4 z Addition of HBr 298-327 1 1 . 1 12.9 aa b 8.2 1-8 8.8 4.0 to CZHa 2.9 12.5 bb b - - - Toluene-carrier pyrolysis 748-8 14 40 13 cc - BrCH,C,H4CH2Br Addition NF (N2F4) 373-428 13.7 12.9 dd -1.8 10.5 4 2 * - to olefin 15.5 10.6 dd 99 ,9 334-391 13.5 13.4 dd -0-2 14.6 0.9* - 13-7 10.2 dd Table-continued Reaction Radical source Temp.E log A (sec.-l and (kcal. mole-1 (OK) mole-') ml. sec.-l) Halogeno-radicals Addition NF2(N2F4) 334-391 10.3 12.5 13.6 10.1 to olefin 1 s s7 336391 13.6 14.0 11.9 9.5 7 7 7 9 334-391 13.6 14.0 11.9 9.5 99 7 9 314-373 9.7 13.3 10.1 9.0 9 9 31 314-373 8.3 12.1 8.3 8.3 9s 7 9 334-391 9.8 11.4 11.0 8.9 Ref. ASo AHbZos AS" (cal. (Cal. (kcal. mole" (kcal. mole-' mole-') deg.-l)" mole-l) deg.-l)a Thermodynamic Kinetic data datab dd -3.3 11.0 0*9* - dd dd 1.7 20.6 1*6* - dd dd 1.7 20.6 2*8* - dd dd -0.4 19.7 4*0* - dd dd 0 17.4 6*3* - dd 3 9 8 dd -1.2 11.4 7.1* - dd a Tablecontinued Reaction 2 1 H ~ O -+ H.+ co 1H. + CO -+ HCO rCH3.CO + CH3. + CO < CCH,. + co -+ C H ~ O { C3?::60 -+ CF,. + CO T CO -+ CF3*C0 rC2H5*C0 -+ C2H5' + CO tc,H5* + CO + CzH,*CO *CH20H -+ He + H-CHO *H + HCHO + CH,OH CH,*O*CH,* -+ CH3. + HCHO CH3. + H*CHO -+ CH,*OCH,* CH3.0-O*CH2* P CH30* I + HCHO 1 CH30* + HCHO 1 -+ CH,-OO-CH,* CH3.CO.CH2 + CH,. + CHZCO CH3. + CHzCO + CH3'COCHZ Radical source Temp. E log A (sec.-l and (kcal. mole-' (OK) mole-') ml. set.-') Oxygen-containing radicals 13.0 1.0 13.8 /"\ Hg(3P1) + CH2-CH2 298 14.4 Photolysis CH3.C0.CH3 473-568 15.0 (Ern) 10.3 (Am) in presence of HI 12.0 (E,) 14.5 (A,) 3.9 8-6 Photolysis C1 or Br 298-521 10 10.4 in presence of CF,.CHO 4 9 Photolysis (C2H5N) 303-353 14.7 (Em) 13.3 (Am) in presence of C2H5.CH0 10.5 (E,) 15.6 (A,) 5 11 Hg(3P1) + CH30H 673-773 29 13-1 3 13.6 Hg(3P,) + CH3.O.CH3 473-573 25.5 (Em) 13.2 (Am) 18.1 (Eo) 16.4(A0) Photolysis 394-453 5.8 - (CH301 - - Photolysis 365-435 41 - CH3*CO*CH3 - - Ref.ee b 8 ii ii j j b kk V R dHozga AS" AH0298 ASo g (cal. (kcal. mole-l mole-') deg.-l)a Kinetic data 134 -3.7 c;; (cal. 2 (kcal. mole-' 3 mole-') deg.-l)a Thermodynamic !Ri n" datab rs 6' % 18-5 1.1 6 6-4 14* 15" 9.7 10.5 10.3 11.0 1.0 26.0 -2.5 28.5 II - - -37.3 9.5 mm - - - 35.2 10.9 Table-continued Reaction Radical source Temp. E log A (sec.-l and (kcal. mole-l ( O K ) mole-') ml. sec.-') Oxygen-containing radicals C,H5*O- -> CH,. + H*CHO Photolysis 288-468 13.0 10.5 CH,. + HCHO + C,H5.0- C2H5-C02*C2H5 - - r(CH,),CHO. + CH,. + CH,.CHO Pyrolysis PI-i O N 0 in presence of NO 8.3 (E,) 13.1 (A,) 433-473 17.3 ( E X ) 11.8 ( A m ) - - 4 (CH3.+ CH,*CHO + (CH3)ZCHO. [ sec-C,H,O. -f C2H5- + CH,.CHO Pyrolysis sec-C,H,.ONO 10.6 (Eo) 14.5 (A,) 423-463 17.5 ( E K ) 13.4(Aa) - - 3 IC2H,* + CH,CHO + sec-C,H,O. rt-C,H,O. -+ CH,. + CH3.C0.CH Pyrolysis [Bu'O] 398-436 22.8 (Em) 14.7 (Am) in presence of NO 13.4 (Eo) 15-7 (Ao) I C H . -k CH,*CO*CH -+ t-C,H,O 5.0 9-2 CH3.CHO.C,H5 + CH,*CHO CH + (C2H5),0 416-453 23.5 10.9 20.6 8.6 J + C2H5* { 1 CzH6* + CH,*CHO I + CH3*CH0.C,H5 C6H,*OCH2. -+ C6k16*CH0 + -Hi CH,. 1- C,H5*OCH 453-539 21 12.5 H- 4- C,H5.CH0 -+ C6H5*OCH2- - - CH,-SO,* + CH3* + SO CH,. -t so 298-437 22.4 13-0 CHB. + SO CH,*SO,. 1.5 10.8 Ref. ASo (cal. (kcal. mole-l mole-l) deg.-l)a Kinetic data - nn - - - 00 - - PP - qq 17.8 25-3 b rr 2-9 10.5 b tt 20.9 10.1 tt dN",,s AS" (cal.(kcal. mole-l mole-l) deg.-;)" Thermodynamic datab 11.2 12.4 6.6 17.3 4.2 19.1 3.8 20.3 3.7 14.8 22 - VI JC2H5.S0,. + C2H5* + SO C&5* + SO2 301-348 19.9 14.4 2 1C2H6* + SO + C2H6S0,- 3.1 11.0 uu 16.8 15-6 UU 2 References for Table b t3 8 0 3 'tr $. 5 aStandard state 1 mole ml.-l; bref. 2; Cref. 34; dW. M. Jackson and J. R. McNesby J. Amer. Chem. SOC. 1961,83,4891; eK. Yang J. Amer. Chem. SOC. 1962,84,3795; fref. 17; gR. J. CvetanoviC and R. S. Irwin J. Chem. Phys. 1967,46,1964; href. 45; ireaction involves isomerisation see section 6; rref. 47 ref. 48; kref. 21; IJ. A. Kerr and A. F. Trotman-Dickenson J. Chem. Soc. 1960 1611; mJ. A. Kerr and A. F. Trotman- Dickenson J. Chem. SOC. 1960 1602; *A. S. Gordon J. Chem. Phys. 1961,34 331; OM.C. Lin and K. J. Laidler Canad. J. Chem. 1967,45 1315; PE. L. Metcalfe and A. F. Trotman-Dickenson J. Chem. SOC. 1960,5072 J. R. McNesby and W. M. Jackson J. Chem. Phys. 1963,38,692; QD. G. Dalgleish and J. H. Knox Chem. Comm. 1966 917; rref. 12; SR. N. Birrell and A. F. Trotman-Dickenson J . Chem. SOC. 1960,4218; J. R. McNesby and W. M. Jackson J. Chem. Phys. 1963 38 692; tK. H. Anderson and S. W. Benson J. Chem. Phys. 1964,40 3747; "A. S. Gordon Canad. J. Chem. 1965,43,570; W. P. Quinn Trans. Faraday SOC. 1963,59,2543; *R. Eckling P. Goldfinger G. Huybrechts G. Martens L. Meyers and S. Smoes Chem. Ber. 1960 93 3014; 2P. B. Ayscough A. J. Cocker F. S. Dainton S. Hurst and M. Weston Proc. Chem. SOC. 1961 244; YG. Huybrechts L. Meyers and G. Verbeke Trans. Faraday SOC. 1962 58 1128; zP.Goldfinger G. Huybrechts and G. Martens Trans. Faraday SOC. 1961 57,2210; aaD A. Armstrong and J. W. T. Spinkes Canad. J. Chem. 1959 37 1210; bbR. J. CvetanoviC Adv. Photo- chem. 1?63 1 173; Ccref. 7; ddref. 19; eeR. J. CvetanoviC Canad. J. Chem. 1955 33 1684; ffref. 32; ggref. 16; hhestimated from CH,. + CO -+ CH,CO; tiref. 15; ijM. K. Phibbs and B. de B. Darwent J. Chem. Phys. 1950,18,495; kkL. F. Loucks and K. J. Laidler Canad. J. Chem. 1967 45,2767; IlY. Takezaki T. Mizazaki and N. Nakohara J. Chem. Phys. 1956 25 536; mmR. K. Brinton J. Amer. Chem. SOC. 1961 83 1541; nnM. H. J. Wijnen J. Amer. Chern. SOC. 1960 82 3034; Ooref. 6; ppR L. East and L. Phillips J. Chem. SOC. (A) 1967 1939; qqref. 5; rrJ. Long and G. Skirrow Trans. Faraday SOC. 1962,58 1403; SSM. F. R. Mulcahy B.G. Tucker D. J. Williams and J. R. Wilmshurst Austral. J. Chem. 1967 20 1155; ttA. Good and J. C. J. Thynne Trans. Faraday SOC. 1967 63 2708; uuA. Good and J. C. J. Thynne Trans. Farahy SDC. 1967 63 2721. * Estimates based on group-additivity rules. 0 2 2 Kerr and Lloyd has been carried out the latest result is usually quoted. While many of the data have appeared in the comprehensive compilation of O'Neal and Bensoq2 to which the reader is referred for greater detail a significant number of subsequent results are also included. Comparisons between the kinetic and thermodynamic estimates of AH" and AS" as detailed in the previous sections are given wherever possible. Discussion of the results will be limited to reactions where discrepancies exist in these comparisons. There is reasonably good agreement between the kinetic and thermodynamic data for the decompositions of the alkyl radicals notable exceptions being some of the decompositions yielding hydrogen atoms and the decomposition of the 1-methylpentyl radical.It seems likely in the latter cases that the kinetic data for the decomposition reactions are in error. The kinetic data for chloro- and bromo-alkyl radicals are remarkably consistent with the thermochemical estimates and this lends considerable support to the kinetic results. Unfortunately there is still little thermodynamic information on difluoroamino-radicals to compare with the reasonably extensive kinetic results for the radical decompositions; approximate enthalpy changes have been obtained from group-additivity calculations and these show satis- factory agreement with kinetic estimates.The largest anomalies between the kinetic and thermodynamic determinations of AH" and AS" for radical decompositions arise with oxygenated radicals. Reliable kinetic information on the decomposition of the forinyl radical is lacking. There are considerable practical difficulties in studying the reaction particularly in determining the concentrations of the radicals and the best estimate of the rate constant for the decomposition is derived from the thermo- dynamic data.2 Amphlett and WhittlelG were. unable to detect any pressure- dependence in the decomposition of the CF,-CO radical in the pressure range 10-60 mm. generating the radicals from the selective photolysis of chlorine or bromine in the presence of the aldehyde. This result is difficult to reconcile with the observed pressure-dependence of the CH3k0 and C,H,.CO radical decompositions.Subsequent work on the azomethane-trifluoroacetaldehyde system 37 has revealed pressure-dependence in the CF,.CO decomposition and accordingly the Arrhenius parameters reported by Amphlett and Whittlels appear to have been measured in the fall-off region. The reactions of alkoxy-radicals including decompositions have been reviewed recently.38 Although the rates of decomposition of several alkoxy- radicals have been measured and thermodynamic estimates of AH" and AS" are available there have been few kinetic studies of the reverse radical addition reactions to the carbonyl compounds. These reactions are difficult to study since they compete unfavourably with hydrogen-abstraction reactions.Only for the t-butoxy-radical is all the information available and as seen from the Table a major discrepancy exists between the kinetic and thermodynamic s7 J. A. Kerr and A. C. Lloyd unpublished results. 8a P. Gray R. Shaw and J. C. J. Thyme Progr. Reaction Kinetics 1967 4 63. 573 6 Decomposition Reactions of Radicals data. The thermochemical estimate of AH" is least likely to be in error so this implies that the kinetic data for the decomposition and/or addition reactions are seriously at fault. Several studies of the decomposition have been made but only the most recent has attempted a quantitative treatment of the pressure- dependence. 6 Radical Isomerisation Reactions Rearrangements of radicals in gas-phase oxidation and related processes have been discu~sed,3~ and there is ample evidence for the structural isomerisations of aryl radicals in solution,4° although kinetic information is lacking.Of considerable kinetic interest are the structural isomerisations of small radicals in the gas phase on which several studies have recently been reported. It was first suggested by Kossiakoff and Rice41 that n-alkyl radicals can isomerise by intramolecular hydrogen abstraction the so-called 'tail biting' reaction CH3CH2CH2CH2.CH2 + CH3*dHCH,CH2CH3 (1 8) Qualitative support for such reactions has been obtained mainly from studies of alkyl radical additions to simple ole fin^.*^ Quantitative data are even more difficult to obtain than for radical decompositions since the product of the reaction is another radical. Endrenyi and LeRov3 have reported the first Arrhenius parameters for this type of reaction the isomerisation of n-pentyl to 1-methylbutyl [reaction (18)l.The radicals were generated by the addition of methyl radicals to ethylene and the subsequent addition of the n-propyl radicals to another molecule of ethylene. The occurrence of reaction (18) was inferred from the presence of 2-methylpentane in the products CH,. + CH3*CH*CH2CH2CH3 + (CH3)&HCH2*CH2.CH3 (19) In the presence of a high concentration of methyl radicals it was assumed that reaction (19) was the major fate of the 1-methylbutyl radicals. The rate constant for reaction (18) is thus given by and the concentration of n-pentyl radicals can be monitored by the reactions n-C5Hll. + CH,. = n-C,Hl CH3. + CH3* = C2He A. Fish Quart. Rev. 1964 18 243.'O A. F. Trotman-Dickenson 'Free Radicals' Methuen London 1959 p. 113; W. A. Pyror 'Free Radicals' McGraw-Hill New York 1966 p. 266. 41 H. Kossiakoff and F. 0. Rice J. Amer. Chem. SOC. 1943 65 590. B. Sefton and D. J. LeRoy Canad. J. Chem. 1956 34 41; A. S. Gordon and J. R. McNesby J . Chem. Phys. 1959 31 853; M. H. J. Wijnen J. Amer. Chem. SOC. 1961 83 3752. L. Endrenyi and D. J. LeRoy J. Phys. Chem. 1966,70,4081. 574 from which it can be shown that Kerr and Lloyd (23) From equations (20) and (23) it follows that and the temperature coefficient of the rate-constant ratio kl,k223/k21 can be determined by analysis of the ethane and n- and iso-hexane products. The above scheme is somewhat simpler than that of Endrenyi and LeRoP3 since they had to take account of the fact that hexane is also formed by the com- bination of n-propyl radicals.From the known value of kz2 and an assumed value of kZl it was deduced that log k, = 7.15 - 10,800/2*3RT(~~.-') An activation energy of about 10 kcal. mole-l seems reasonable for this type of process on the basis of its analogy with a radical hydrogen-abstraction reaction. On the other hand an A-factor of 107s2 sec.-l is incompatible with transition-state theory calculations on related systems. An estimate based on the method of O'Neal and B e n ~ o n ~ ~ (see section 3) indicates an A-factor in the range In the same study Endrenyi and LeRof3 were able to estimate asrate constant for the analogous isomerisation of the radical CH3COCH2CH,CH, although they could not determine the Arrhenius parameters. More work is needed on this type of process before the rate constants can be assigned with confidence.Although it is reasonable to postulate isomerisation involving intramolecular hydrogen transfer with C and larger radicals the idea seems less attractive for C4 and C3 radicals. Considerable discussion has taken place on the feasibility of reactions such as sec.-l for this type of process. iso-C,H,- -+ n-C,H,- -+ CH,. + C2H t-C4H,. -+ iso-C,H,. -+ CH,. + C,H6 Two types of experiment bearing on this topic have been carried out. In the first the occurrence of isomerisation-decomposition reactions has been inferred from product analysis of rather complex systems known to involve the initial radical. For instance in the photolyses of di-isopropyl ketoneu and isobutyl- aldehyde4 and in the pyrolyses of di-isopropylmercury46 and i~obutane:~ where it has been established that isopropyl radicals are produced reaction (24) has been suggested as the source of ethylene that is observed in the products at high temperatures.In all of the experiments of this type it has been necessary to postulate simultaneous isomerisation and decomposition of the initial radical as other products of the isomerised radical such as radical dimers 44 C. A. Heller and A. S. Gordon J. Phys. Chem. 1958 62 709. 45 J. A. Kerr and A. F. Trotman-Dickenson Trans. Faruday SOC. 1959 55,921. 46 B. H. M. Billinge and B. G. Gowenlock J. Chem. SOC. 1962 3252. 47 R. S. Konar R. M. Marshall and J. H. Purnell Trans. Faraduy SOC. 1968,63,405. 575 Decomposition Reactions of Radicals have not been observed in significant quantities.It should also be pointed out that in some instances it is not inconceivable that the products assumed to arise from radical isomerisation could be produced in secondary reactions.2 Thus ethylene in the isopropyl systems might arise from the reactions :48 H. + C,H -+ n-C,H,. In the second type of experiment concerning C and C alkyl isomerisations isotopically labelled radicals have been generated either by pyrolysis of a suitably labelled alkane or by decomposition of the alkane sensitised by methyl radicals. Thus reaction (24) has been investigated with CD,CH,CH radicals produced in the pyrolysis of CD,-CH,CH :48 n-C,H,. -+ CH,. + C2H -+ CD,.CH.CH + RH R. + CD3CH2CH3 __+ CD2*CH2CH3 + RD -+ CD3CH2*CH2 + RH Mass-spectrometric analysis was carried out for the isotopically labelled ethylenes and it was justifiably argued that the product CH,-CDH could only arise from the isomerisation-decomposition sequence CD3-CH*CH3 + [CD,-CHD*CH,] + CD,=CHD + CH,.(240) In this way the rate of reaction (24a) was measured relative to the rate of the alternative decomposition CH3-dH*CH3 + CH,=CH-CH + H- (25) Jackson and McNesbP* determined k24a /k25 < 0.06 below 826"~ whereas other determinations without isotopic labelling44s4' have indicated much higher values of this ratio. In similar experiments with isotopically labelled alkyl radicals McNesby and his co-workers have shown that up to 7 7 3 " ~ the isomerisations of i s o b ~ t y l ~ ~ ~ * and sec-buty15" radicals are negligible. In conclusion it may be said that while the simultaneous isomerisation and decomposition of isopropyl and n- and t-butyl radicals (see Table) may occur at high temperatures these reactions are by no means established and there is no general agreement regarding their rate constants.It may be further stated that present evidence is wholly against simple isomerisation of C and C4 alkyl radicals at lower temperatures. Substantial evidence based on product analyses has been obtained for the isomerisations of cyclopropy151 and cyclobutyl radicals :18 46 W M. Jackson and J. R. McNesby J. Chem. Phys. 1962 36 2272. Jackson J. R. McNesby and B. de B. Darwent ibid. 1962 37,2256. 6o A. S. Gordon and J. R. McNesby J. Chem. Phys. 1960 33 1882. J. C. J. Thynne Trans. Faraday SOC. 1966 62 3338; 1967 63 1369. J. R. McNesby C. M. Drew and A. S. Gordon J. Chem. Phys. 1956 24 1260; W.M. J. R. McNesby and A. S. Gordon J. Amer. Chem. SOC. 1957 79 825; G. Greig and 576 Kerr and Lloyd CH2 /\ + CH2CH=CH2 H2C - CH H2C-CH2 I I __+ CHa*CH,*CH=CH2 H2C-CH Activation energies of ca. 20 kcal. mole-l have been estimated for these reactions but there was insufficient information to obtain the A-factors as well. An interesting type of isomerisation has been reported for alkenyl radicals produced from the additions of propyl and butyl radicals to alkynes in the gas phase:52 (CH3),CH 3. CH=CH __+ (CH,),CH-CH=CH (CH3)2CHCH=kH + CH,*CH*CH,.CH=CH Benson and D e M ~ r e ~ ~ have suggested the following general mechanism to explain the isomerisation I I I I I I I l l I I I I I l l H-C142-C3=C4 + .C1-C2-C3=c H .1 -c2-c3-c4 I I I I .CLCLCS=C4 f- I H I 1 \/ I C1 H and all the systems so far studied can be explained on this basis.Arrhenius parameters corresponding to E ca. 10 kcal. mole-l and A ca. sec.-l have been estimated for the isomerization process (26) by an approximate RRK treatment of the data on the additions of isopropyl and t-butyl radicals to acetylene and pr~pyne.~ Similar isomerizations have been shown to occur in the liquid-phase reactions between iodine and ole fin^,^^ and between carbon tetrachloride and hept-l-yne." We thank Drs. S. W. Benson H. E. O'Neal and R. Walsh for access to results before publication and Professor K. J. Laidler for permission to use Figures 3 and 4. 52 J. A. Garcia Dominguez and A. F. Trotman-Dickenson J. Chem. Soc. 1962 940; R. R. Getty J. A. Ken and A. F. Trotman-Dickenson J. Chem. Soc. (A) 1967 1360. 63 L. H. Slaugh R. D. Mullineaux and J. H. Raley J. Amer. Clrem. Soc. 1963 85 3180. 54 E. A. I. Heiba and R. M. Dessau J . Amer. Chem. Soc. 1966 88 1589. 577

ISSN:0009-2681    

DOI:10.1039/QR9682200549

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

The Reactions of Hydrated Electrons with Inorganic Compounds By Michael Anbar* STANFORD RESEARCH INSTITUTE MENLO P A R K CALIFORNIA 94025 The hydrated electron is undoubtedly one of the most fundamental species ever discovered in chemistry and its discov^ery and investigation should be considered one of the most important achievements in pure chemistry in the present decade. The discovery of e-aq and the methods of production and investigation of the hydrated electron as well as its physicochemical properties have been recently comprehensively reviewed.l The reactions of eaq were reviewed over three years ago2 and the rapid developments in this field made many of the earlier con- clusions outdated. As these reactions have far-reaching implications in general chemistry a revised discussion of the reactions of e-aq and their mechanisms must be of general interest.This Review covers a limited yet important phase of the chemistry of hydrated electrons namely their reactions with inorganic compounds. We shall discuss the available information on the rates of e-aq reactions and the identity of the products of these reactions. Many of these products are hitherto unknown species the chemical behaviour of which may engage a great number of inorganic chemists in the coming years. All e-aq reactions are electron-transfer reactions by definition. The experimental informa- tion on e;tq reactions will be used in the interpretation of these mechanisms to which we shall devote the second part of this Review. 1 The Reactions of Hydrated Electrons and Their Products Hundreds of inorganic compounds have been investigated for their reactivities with hydrated electrons in aqueous solutions.The reactivities of these compounds range from diffusion controlled rates down to 16 1. mole-l sec.-l which is the calculated specific rate of reaction (1). In all cases investigated the primary eaq + H,O - H + OH- (1) product of the electron-transfer reaction (2) acquires an additional electron before undergoing any subsequent chemical change. In numerous cases the primary product is thermodynamically unstable and undergoes further reactions such as protonation dissociation disproportionation or charge-transfer. *This review has been prepared in part at the NASA Ames Research Center Moffett Field California 94035. D. C. Walker Quart. Revs. 1967 21 79. a M. Anbar Adv. Chem.Ser. 1965 50 55. 578 Anbar The reactions of inorganic compounds with hydrated electrons will be de- scribed following the elements and their compounds throughout the Periodic Table. A. Hydrogen and Its Oxides.-Molecular hydrogen is completely inert toward e-aq. The reaction of hydrogen atom with e-aq on the other hand is a diffusion- controlled reaction which has been shown to produce hydride ions (3) as the e-aq + H --t H- (3) primary pr~duct.~ The reaction of e-aq with H+aq proceeds at a high rate (k = 2-3 1O1O 1. mole-l se~.-l)~ which is however somewhat lower than the calculated diffusion-controlled limit. It has been suggested that owing to the fact that the electron vacancy in H,O$ is rather diffuse the electron transfer involves an appreciable entropy of activation.2 This rapid reaction limits the lifetime of e-aq in the acid range.Even at pH 5 the lifetime of e-aq is less than 3 pec. which makes it rather difficult to investigate any other e-aq reaction much below this pH. The final products of this reaction (4) are hydrogen atoms which are pro- eaq + H+aq --t H (4) duced3s5 with an H/D isotope effect of 3.7 which can hardly be explained unless an intermediate e.g. H,O is formed as a primary product. The lifetime of H30 in water is probably shorter than lo-* psec. as no significant difference has been observed between the rates of reaction of H atoms produced directly by radiolysis and those formed from the reaction of eaq with H30+ up to relatively high concentrations of solute^.^ Reaction (5) can be considered as a special case of reactions of type (6) where HX may be any protonic acid e.g.HF NH+, eaq + HX 3 HX- -+ R + X- (6) H,PO, HCO,H etc. A Br~nsted relation has been shown to exist between the specific rates of certain protonic acids and their dissociation constants.6 This relation has been suggested to imply that the rate-determining step involves a proton transfer from the acid to the hydrated electron resulting in the formation of an H atom? It was found however that a number of protonic acids which follow the same Br~rnsted relation produce products other than H atoms from their reactions with e-as.2 Moreover even when hydrogen atoms are formed a considerable H/D isotope effect on the H atom formation is observed which is not reflected in their respective ~ K ’ S . ~ ~ It must be concluded therefore that the M.Anbar and D. Meyerstein Trans. Faraday SOC. 1966 62,2121. S. Gordon E. J. Hart M. S. Matheson J. Rabani and J. K. Thomas J. Amer. Chem. SOC. M. Anbar and D. Meyerstein J. Phys. Chem. 1965 69 698. J. Rabani Adv. Chem. Ser. 1965 50 242. M. Anbar and P. Neta Trans. Faraday Soc. 1967 63 141. 1963,85 1375. 579 The Reactions of Hydrated Electrons with Inorganic Compounds Br~rnsted relation reflects the tendency of a protonic acid to accommodate an additional electron and to form HX- as primary product. This tendency is correlated with the electron vacancy in the same hydrogen atom which in turn is a prerequisite for its protonic dissociation. Once HX- has been formed it may lose a hydrogen atom or dissociate in another way depending on its structure and the strength of its chemical bonds.The reaction of the hydrated electron with water is probably the most im- portant reaction of e-aq as it determines its natural lifetime in the pure solvent. Because this reaction is relatively slow it is possible to observe the hydrated electron and to study its physical and chemical properties. Reaction (7) is over- shadowed under most conditions by reactions (8)-(10) as well as by the scaveng- ing of e-aq by free radicals like H atoms or OH radicals simultaneously produced. Trace impurities in the solvent such as oxygen hydrogen peroxide metal ions or organic compounds may also react with e-aq and shorten its life. It was necessary therefore to take utmost precautions of reagent purity and to use most refined experimental techniques in order to obtain a reliable value for the unimolecular conversion of e-aq into H atoms in pure water.In a most careful study a unimolecular specific rate constant of 890 sec.-l was measured from which an apparent bimolecular rate constant k = 16 f 1 1. mole-l sec.-l can be calculated.s The mechanism of the reaction of eaq with H,O is far from understood. Unlike the majority of eaq reactions there is no room for the accommodation of an additional electron in the water molecule at its ground state. The formation of an electronically excited H20- is energetically unfeasible; thus a concerted multimolecular process is much more likely. It seems plausible that this reaction involves a number of water molecules which solvate the OH- in the transition state making the reaction thermodynamically feasible. The reaction of e-aq with hydrogen peroxide is diffusion-controlled~~g an expected result since fast single-electron reduction reactions of H202 are well- known processes.B. The Alkali-metal Ions Alkaline Earths and Rare Gases.-The hydrated electron has been shown to have a redox potential of about 2-7 v;l0 therefore it can reduce any species with a lower redox potential. All the alkali-metal cations Li+ Na+ K f Rb+ and Cs+ have a higher redox potential than 2.7 v; thus they * E. J. Hart S. Gordon and E. M. Fielden J. Phys. Chem. 1966,70 150. E. J. Hart and E. M. Fielden ‘Pulse Radiolysis’ ed. M. Ebert Academic Press New York 1965 p. 253. lo J. H. Baxendale ‘Current Topics in Radiation Research’ ed. M. Ebert and A. Howard North Holland Amsterdam 1967 vol. 3 no. 1. 580 Anbar cannot be reduced by hydrated electrons to the metal atoms.In fact there are indications1 that under certain conditions the reaction of sodium goes spon- taneously in the opposite direction (11). Na + H20 3 e-aq + Na+ The reactivity of hydrated electrons toward beryllium compounds in aqueous solutions has not been investigated yet. Mg2+ Ca2+ Sr2+ and Ba2+ ions were found to be non-reactive toward hydrated electrons.11J2 On the other hand Mgn was shown to be reduced to MgI by electrons in frozen solutions,13 most probably because of the diminished stability of Mg2+ in the ice lattice compared with that of this ion in aqueous solution. It can be predicted therefore that Mg2+ and perhaps also the ions of other alkaline earths may be reduced by solvated electrons in alcoholic solutions. The rare gases have a highly negative electron affinity and species like He- or A- have no chemical stability whatsoever.Therefore the elemental inert gases are not expected to show any reactivity toward hydrated electrons. On the other hand rare-gas compounds such as H3Xe0,- or XO which are extremely efficient oxidants are expected to react with e-aq at diffusion-controlled rates. C. Compounds of the Elements of Group III and the Lanthanides.-Unlike the alkali metals and alkaline earths which do not react with hydrated electrons the tervalent ions of the elements of Group 111 with the exception of boron react with eaq at relatively high rates. The very low reactivities of B2072- H2B03- and BF4- ions14 are most probably not a consequence of thermodynamic but of kinetic parameters. The reaction (12) for instance is a highly exoergic process.BF4- + eaq- --t BF3- + F- (1 2) Aluminium salts react with e-aq in neutral and alkaline solutions the re- activity ranging from 2 x lo9 1. mole-l sec.-l at pH 6-8 to 5.5 x lo6 1. mole-l sec.-l at pH 14.15 The high reactivity in neutral solutions is not surprising since a rough estimate of the free energy of the reaction (13) yields a negative value of over 100 kcal./mole. The decrease in reactivity with pH can be explained by two effects. In the intermediate region polynuclear species prevail which evidently have an apparent lower specific rate when calculated on the basis of the concen- tration of the monomer. AI(OH),- which prevails in alkaline solution exhibits a very low reactivity toward e-aq for reasons which will be discussed in Part 2.The profound effect of ligands on reactivity is demonstrated again in the case of the aluminum-EDTA complex (k = 3 x lo7 1. mole-l sec.-l).16 11 M. Anbar and D. Meyerstein J. Phys. Chem. 1964 68 1712. 13 M. Anbar and E. J. Hart J . Phys. Chem. 1965 69 1244. Is P. N. Moorthy and J. J. Weiss Adv. Chem. Ser. 1965 50 180. l4 M. Anbar and E. J. Hart A h . Chem. Ser. 1968 81 79. 16 M. Anbar and E. J. Hart J. Phys. Chem. 1965 69,973. lo M. Anbar and D. Meyerstein to be published. 58 1 The Reactions of Hydrated Electrons with Inorganic Compounds The aquo- or hydroxy-complexes of the tervalent ions of the other members of Group I11 are insoluble in neutral or alkaline solutions; thus their reactivity with eaq cannot be evaluated. However the EDTA complexes of gallium and indium have been studied for their rates of reaction with e-aq and an increase in reactivity in the sequence Al Ga In has been observed (k = 3.0 7-8 and 41.0 x lo7 1.mole-l sec.-l respectively).16 This increase in rate may be due to the increase in the stability of the bivalent state of these elements. TPaq is unstable in neutral solution; however TI+ is quite stable and has been investigated. Tlaq+ reacts with e-aq at a diffusion-controlled rate17 produc- ing T1° as the primary product. The latter species forms a complex1* with T1+ giving T1$. Whereas it is impossible to study the behaviour of gallium and indium aquo- complexes in neutral solutions this has been possible for other members of Group 111 namely yttrium lanthanum and the lanthanides. The reactions of these aquo-complexes can be divided into two groups; those that react rapidly with e-aq (k > lo9 1.mole-l sec.-l) including samarium europium thulium and ytterbium on one hand and those that react more slowly (in the range lo7-lo9 I. mole-l sec.-l). For the first a correlation has been pointed out19 between the redox potential of the M3+/M2+ couple and the rate of the reaction with e-aq. If the specific Table 1 Eu Yb Sm Tm 23" (M*/M2+) (v) 0.43 1-15 1-55 > 1.5 lolokobs (1. mole-l sec.-l)13 6.1 4.3 2.5 0.3 lV0k (1. mole-l sec.-l) 24-6 9.0 3.6 0.3 rate constants for these reactions are corrected for the diffusion-controlled rate according to equation (14) a fairly good linear correlation is obtained indi- cating that the change in free energy makes a major contribution to the reac- tivity of these complexes with hydrated electrons.The reduction of Eum to Eurc by e-aq has been demonstrated spectrophoto- metrically by pulse radidlysis.20 The identification and study of the chemical behaviour of other bivalent cations of the lanthanides produced by e-aq in aqueous solutions still awaits investigation. The EDTA complexes of the transition elements of Group IIIA scandium yttrium lanthanum and the lanthanides have been investigated.16 The rates for these EDTA complexes excluding the highly reactive lanthanides discussed J. H. Baxendale E. M. Fielden and J. P. Keene Proc. Roy. SOC. 1965 A 286 320. S. Gordon ref. 9 p. 285. 18 B. Cercek M. Ebert and A. J. Swallow J. Chem. SOC. (A) 1966 612. la J. K. Thomas S. Gordon and E. J. Hart J. Phys. Chem. 1964 68 1524. 582 Anbar above were correlated with their thermodynamic stability constants.It was found that the reactivity of the members of this family of complexes increases linearily with increasing stability of their EDTA complex.16 As an inverse correla- tion would be expected if the rate of reaction of eaq depended solely on the redox potentials of these complexes this finding implies that other parameters besides the gross effect of free-energy change determine the rate of the electron transfer. D. Compounds of the Elements of Group N.-Relatively little is known about the reactivity of hydrated electrons with compounds of the elements of Group IV with the exception of the organic compounds. The main reason for this is that relatively few inorganic compounds of this group are stable in a monomeric form in neutral aqueous solutions.Of the 'inorganic' compounds of carbon investigated for their reactivity with e a q one should consider CO C02 CCl, C(NOh4 CS2 HC03- C032- CN- SCN- and OCN-. These compounds range in their reactivity from quite inert species like CN- SCN- OCN- C02- and HC0,- (k > lo6 1. mole-l sec.-l)19 to the extremely reactive CCI, CS, and C(NO&, which react at diffusion- controlled rates (k = 3.0 3.1 and 4.6 x 1O1O 1. mole-l sec.-l respectively).21 Carbon dioxide reacts with eaq at a rate which is close to the diffusion-controlled limit whereas carbon monoxide reacts2 at an intermediate rate (k = 1 x lo9 1. mole-l sec.-l). The product of the reaction of CO with e-aq is probably CO- which reacts with water to give HCO; HCO then dimerizes to glyoxa1.22 HCO is a strong reducing species and is oxidized by H202 to formic acid.The chain reaction of CO in alkaline solution to yield formate ions23 may involve CO- as chain carrier. CO is reduced by eaq to CO, which has been shown to combine with various organic radicals to form carb~xylates,~~,~~ e.g. reaction (15). C0,- C02- + CH2CO2 -+ CH,(C02-)2 (1 5 ) is also an efficient reducing agent which was found to reduce H202,25 I, Br2 RBr NO3- and NO2- as well as c6H5No.26*27 The only silicon-containing compound investigated for its reaction with e-aq is SiFt- which is quite inert (k > 1.5 x lo6 1. moleb1 sec.-l).14 This is not surprising in view of the thermodynamic instability of Sim. Whereas none of the germanium compounds have been investigated both tin(1r) and tin(1v) derivatives have been studied for their reaction with e-aq.The Snn hydroxy-complex (Sn0d2-aq has been found to react with e a q at a diffusion- al E. J. Hart s. Gordon and J. K. Thomas J. Phys. Chem. 1964 68 1271. a2 E. J. Hart J. K. Thomas and S . Gordon Rqdiation Res. Suppl. 1964 4 74. 14 G. Scholes and M. Simic J. Phys. Chem. 1964 68 1731. 26 A. Husain and E. J. Hart J. Amer. Chem. SOC. 196sf87 1180. a* M. Anbar and P. Neta J. Chem. SOC. (A) 1967 841. 27 K. D. Asmus G. Beck A. Henglein and A. Wigger Ber. Bunsengeseilschaft Phys. Chem. 1966 70 869. J. J. Weiss Radiation Res. Suppl. 1964 4 141. 583 The Reactions of Hydrated Electrons with Inorganic Coinpounds controlled rate? The fluoro-complexes of Snw and Sn* SnFZ- and SnF,- react with e-aq at rates approaching the diffusion-controlled limit (k = 4.1 and 9.3 x lo9 1.mole-l sec.-l respectively).14 Lead the last member of Group IV has been studied for its reactivity only in the bivalent state. It was found that the aquo- hydroxy- and EDTA com- plexes of P@I react at diffusion-controlled rates.l4-I6 There is an interesting difference between the behaviour of the hydroxy- and EDTA complexes of lead and those of aluminum and zinc.15 Whereas the hydroxy- and EDTA complexes of A1111 and ZnII are much less reactive than the corresponding aquo- or aquo- hydroxy-complexes the hydroxy-complex of lead is just as reactive. Since little change takes place in all these cases in the electronic configuration of the central atom on going from the aquo- to the hydroxy- or EDTA complex the change in reactivity must be in this case due to the ligand.E. Compounds of the Elements of Group V.-The chemistry of nitrogen is characterized by its large number of states of oxidation most of which include water-soluble compounds. There is no wonder therefore that a large number of different nitrogen compounds have been investigated for their reactivity toward hydrated electrons. Starting with nitrogen at its lowest state of oxidation N-m one finds ammonia quite inert toward hydrated electrons. This is to be expected from a molecule with all its low-lying orbitals completely occupied. Liquid ammonia is a medium in which solvated electrons are even more persistent than in water and the life- time of hydrated electrons in water is not affected even in the presence of molar concentration of ammonia. Ammonium ions do react with hydrated electrons although slowly (k = 1.5 X lo8 1.mole-l sec.-1).28 The product of this reaction is probably the NH radical29 which dissociates to NH + H with an HID isotope effect of 4.2 & 0 ~ 1 . ~ Hydrazine and hydroxylamine are relatively inert toward ~ a q . ~ ’ Elementary nitrogen has a negative electron affinity; thus it will not react with the hydrated electron. On the other hand N20 which represents NI is a most useful reagent in radiolytic studies as it ‘converts’ e-aq into OH radicals. It reacts (16) with e-aq rather quickly (k = 5.6 x lo9 1. mole-l form- ing nitrogen and OH radicals as final products. (1 6) NzO + e-aq 4 N20- +- N2OH -+ N2 + OH A significantly slower rate of reaction of e-aq with N20 has been observed in acid solution; this has been interpreted to indicate the existence of a less reactive hydrated form of N20 in acid solution.32 Different estimates exist concerning the lifetime of the intermediate 28 E.M. Fielden and E. J. Hart personal communication. 29 C. E. Melton and H. W. Joy J. Chem. Phys. 1967,46,4275. 30 J. H. Baxendale el al. Nature 1964 201 468. 31 J. P. Kenne Radiation Res. 1964 22 1. 8a D. A. Head and D. G. Walker Canad. J. Chern. 1967,45,2051. 584 Anbar N20- 2932-36 The lifetime of N20- in the gas phase was estimated to be greater than set.,' In view of this evidence it is likely that the lifetime of N20- is more than sec. especially in alkaline solutions and that it may react with solutes when these are present at sufficiently high concentration^.^^,^^ The high reactivity of NO with eAaq (k = 3.1 x 1O1O 1.mole-l sec.-l)* is not surprising in view of its electron affinity in the gas phase. The characteristic compounds of Nm in neutral aqueous solutions are nitrite ions. These react with e-aq at a rate approaching the diffusion-controlled limit (k = 4-6 x 109).19 The product of this reaction is probably N02- and it may be most interesting to study the chemistry of the latter species by pulse radiolysis and to compare it with the chemical behaviour of NO dissolved in water. There is no stable Nm compound in neutral aqueous solution; therefore we shall now consider the behaviour of two NV compounds NO,- and N3-. Whereas the former oxyanion reacts at a diffusion-controlled rate (k = 8.5 x lo9 1. mole-l sec.-l) the latter pseudo-halide reacts extremely slowly (> 1.3 x lo6 1. mole-l sec.-l).This difference may be considered a rather unexpected result as azide may be formally considered a derivative of amino-nitrate. The difference in reactivity is thus hard to explain in terms of the redox potentials of these anions. It is suggested that the intermediate N,2- has a much more positive free energy than NO,2-. N3- is a typical electron donor; NO3- on the other hand is a rather poor electron donor and may accept much more readily an additional electron. There is little information on the chemical behaviour of NO,*- the product of the reaction of NO,- with e-aq reaction. Although it has been postulated as intermediate in certain radiolytic there is no evidence for its being long-lived enough to react with solutes.26 Phosphorus oxyanions do not exhibit any spectacular behaviour on reacting with e-aq.H2P0,- and H2P03- react at comparable rates (k = 7-7 x lo6 7.2 x lo6 1. mole-l sec.-l respectively).14 This result indicates that the valency of phosphorus has little effect on the reactivity. As these two ions are protonic acids of comparable strength (pK = 7.2 and 6.6 respectively) it seems that the electron is incorporated into the orbitals of the hydrogen atom. The H atoms e-aq + H2P04- -+ H + HP042- formed in the overall reaction (17) exhibit41 an H/D isotope effect of over 6 strongly suggesting the existence of H,P04’- as intermediate. Hypophosphate 33 B. Cercek ref. 9 p. 289. 34 Y. Raef ref 9 p. 145. 36 E. J. Land ref. 9 p. 285. 3e C. S. Munday J. T. Richards G. Scholes and M. Simic ‘The Chemistry of Ionization and Excitation’ ed. G. R. A. Johnson and G.Scholes Taylor and Francis London 1967 p. 151. 37 W. J. Holtstander and G. R. Freeman J . Phys. Chem. 1967 71 2562. 98 M. Anbar R. A. Munoz and P. Rona J . Phys. Chem. 1963 67 2708. 3s G. E. Adams ‘Radiation Research’ ed. G. Silini North Holland Amsterdam 1967 p. 195. 40 M. Daniels and F. F. Wigg J. Phys. Chem. 1967 71 1024. *l C. Lifshitz and G. Stein Israel J. Chem. 1964 2 337. 585 The Reactions of Hydrated Electrons with Inorgniric Compounds ion H,PO,- which is not a protonic acid at all reacts at a much slower rate (k > 1 x lo5 1. mole-l ~ec.-l),1~ if it reacts at all. Arsenic is a more metal-like element than phosphorus and the oxidizing potential of the couple Mv/Mm rises rapidly going from phosphorus through arsenic and antimony to bismuth. This trend is reflected in the reactivity of these elements toward e-aq.Although no Asw compound has been isolated AsV reacts with e-aq at an appreciable rate. The oxy-derivative HA SO,^- is much less reactive than the hexafluoroarsenate AsF,- which reacts at a diffusion- controlled rate (k = 2.0 x lo8 and 9 x lo9 1. molev1 sec.-l respectively).15 This difference in rates probably originates from the fact that the ionic character of the As-F bond is about twice as large making the central atom more electro- positive. The reactivity of oxy-derivative of Sbv already reaches the diffusion-controlled limit ;15 Biv will evidently behave analogously. The measured specific rate for Sbvaq (1-2 x 1O1O 1. mole-l sec.-l) suggests the participation of a singlypositively charged species. Alternatively one has to assume in this case electron-tunnelling which allows larger reaction cross-sections than the geometrical dimensions of the reactants.Arsenite ions were found to react significantly more quickly than arsenate ions at the same pH (k = 2-0 and 5.9 x lo8 1. molee1 sec.-l respectively).14 This difference may originate from the tendency of Mm to behave as M3+ a tendency which starts with As111 and is most pronounced with BinI. Of the transition metals of Group VA only vanadate ions were examined and found to react with e-aq at a diffusion-controlled rate (k = 4.9 x lo9 1. mole-l sec.-l).14 The rate suggests that the reacting species is singly negatively charged. F. Compounds of the Elements of Group VI.-Elementary oxygen reacts with e a q at a diffusion-controlled rate (k = 1-9 x 1O1O 1.mole-l ~ e c . - ~ ) . ~ ~ It being a biradical with a relatively high electron amity this high reactivity is expected. 02- which is the primary product may subsequently undergo disproportiona- tion to give H,02 and 02.42 The absorption spectrum of 02- formed as a transient from the reaction of 0 with e-aq reaction has been inve~tigated.~~,~~ Water and H202 have been discussed in Section (A). Elementary sulphur is insoluble in monomoric form in water but is expected to be highly reactive toward e-aq owing to its electron affinity (> 45 k~al./mole).~~ H,S is highly reactive toward e-aq (k = 1.35 x 1O1O 1. mole-l ~ e c . - l ) . ~ ~ This reactivity is evidently considerably higher than expected for H,S as a protonic acid. The products of the reaction however are predominantly H + HS- analogous to the behaviour of other protonic acids.The other products formed by the decomposition of H2S- are H + S- which are produced with a yield pp G. E. Adams ref. 9 p. 57. 43 S. Gordon E. J. Hart and J. K. Thomas J. Phys. Chem. 1964 68 1262. 44 G. Czapski and L. M. Dorfman J. Phys. Chem. 1964 68 1169. 45 L. M. Branscomb and S. J. Smith J. Chem. Phys. 1956 25 598. 40 G. Meissner and A. Henglein Ber. Bunsengesellschaft. Phys. Chem. 1965 69 3. 586 Anbar of 35 %. There is no information on the reactivity of S2- and HS- with e-aq but they are expected to be inert. Sulphate and sulphite ions are non-reactive (k < lo6 1. mole-1 sec.-l). There are conflicting results on the reactivity of thi~sulphate,~~ and it is pos- sible that the higher rate observed is due to colloidal sulphur usually present in traces in thiosulphate solutions.The only other well established highly re- active sulphur compound is the persulphate ion S,OS2- which reacts at a diffusion-controlled rate. Selenite and selenate ions are much more reactive than their sulphur analogues (k = 1.2 x lo7 and 1.1 x lo9 1. mole-l sec.-l)14 resembling the case of phosphates and arsenates. As expected from this analogy tellurite reacts at an even faster rate (1.1 x lo9 1. mole-l sec.-l)14 and tellurates react at diffusion- controlled rates (k = 1.6 X 1O1O 1. mole-l s ~ . - ~ ) with an unexpectedly large reaction cro~s-section.~~ Of Group VIA chromium compounds have been investigated. In this section we shall refer only to Cr* ions which have been shown to react at diffusion- controlled as expected from such strong oxidizing agents.The formation of a CrV transient formed from Cr042- has been dem~nstrated.~~ G. The Halogens and Their Compounds.-Fluoride ions are completely non- reactive toward e-aq (k < 2 x lo4 1. mole-l sec.-l).12 Their inertness is shared by the other halide ions as well.47 The halide ions have all their low-lying orbitals completely filled in a rare-gas fashion; thus their reactivity with any reducing agent including e-aq is thermodynamically prohibited. Hydrofluoric acid reacts with e-aq at a moderate rate 6 x lo7 1. mole-l se~.-l)~ which is expected from it as a weak protonic acid.6 HF2- is slightly less reactive (k = ca. 3 x lo7 1. mole-l ~ec.-l),~ again in line with its being a weaker acid than HF. The other halogen hydrides are completely dissociated in aqueous solutions; thus their reactivities which are probably very high cannot be measured.The only elementary halogen which is stable enough in neutral solution and could have been measured is iodine and the tri-iodide ions both of which react at diffusion-controlled rates ;12 so also does the Br2- i~n-radical.~~ Of the oxyanions of chlorine C10- CIO, and CI04- have been investigated. Whereas the reaction of C10- is diffusion-controlled (k = 7.2 x lo9 1. mole-1 sec.-l)14 both (210,- and Clop- are practically non-reactive. The high reactivity of CIO- may be due to an oxygen-induced electron vacancy in the chlorine atom. Although the redox potential Br03-/Br0- is comparable with that of ClO,-/ClO- bromate reacts with e-aq at a much faster rate (k = 2.1 x lo9 1. mole-l sec.-l),15 interestingly enough at a rate similar to that of selenate (1.1 x lo9 1.mole-l sec.-l).14 Iodate and periodate ions like tellurate ions react at diffusion-controlled rates.l49l5 The increase in reactivity along the line C10 < 47 M. Anbar and P. Neta Inter. J . Appl. Rad. Isotopes 1967 18 493. 48 J. H. Baxendale ref. 9 p. 15. 40 M. S. Matheson W. A. Mulac J. L. Weeks and J. Rabani J . Phys. Chem. 1966 70,2092. 587 The Reactions of Hydrated Electrons with Inorganic Compounds BrO < 10,- is probably the result of the stabilization of the quadrivalent states of the halogens. The chemical behaviour of Brw produced by the reaction of BrO,- with e-aq has been in~estigated.~~ This species Br0,2- or (BrO.Jaq was found to react both as an oxidizing and as a reducing agent.In acid solution Br032- was found to oxidize I- and Br- as well as Mn2+ ions. In neutral solution Br0,2- reduces I and Br to I,- and Br,- respectively. In the absence of reactive solutes Br03,- disproportionates to BrO and Br0,-. The permanganate ion representing Group VIIA reacts with e-aq at a diffusion-controlled rate.17s20 H. The Transition-metal Ions and Their Complexes.-In this section we shall discuss the behaviour of the complex ions of Groups IB IIB and VIII as well as those of Cru Crm and Mnn. All these ions with the exception of Group IIB have vacant d orbitals which facilitates their reduction by e-aq. Group IIB element may still be considered as transition metals.51 These ions may exchange the ligand water molecules for other ligands and these substitutions have been shown to affect their reactivities toward e-aq.Some of the reduced products of the eaq reactions are well-known chemical species and the existence of others has been demonstrated for the first time through the e-aq reactions. The tervalent ions of the first transition series Crm F S and CdIIreact with e-aq at diffusion-controlled rates. This includes their amino- ethylenediamino- EDTA cyano- and other c ~ r n p l e x e s . ~ ~ - ~ ~ ~ ~ ~ @ Most complexes of bivalent ions of the first transition series Crn Mnn Fen Con Con Nin and CuII react with e-aq at fast rates but only certain complexes reach the diffusion-controlled limit. Of the aquo-complexes only and perhaps CUIIaq approach the diffusion-controlled limit whereas MGaq is rather inert.30 The EDTA complexes show a similar trend,16 whereas in the case of cyano-complexes one finds Mn(CN),4- extremely reactive and Fe(CWe4- non-reactive alt~gether.~~ The behaviour of these complexes as well as those of d1°Zn2+ are summarized in Table 2.Table 2 Rates of reaction of complexes of the first transition-metals series in the bivalent state of oxidation ( x lo9 1. mole-I sec.-1)14716,47 H2O 42 0.08 0.12 12 22 33 1.5 CN* 4 5 <0~0001- 4.1 - 0.08 EDTA - (0.002 - 0.51 0.1 10.4 <0*002 Cr Mn Fe Co Ni Cu Zn * Results corrected for salt effect. 6o M. Anbar and P. Neta J. Znorg. Nuclear Chem. 1966 28 1645. 6 1 L. E. Orgel ‘An Introduction to Transition Metal Chemistry’ 2nd edn. Methuen London 1966. s2 A. Szutka J. K. Thomas S. Gordon and E. J. Hart J . Phys. Chem. 1965 69 289. 53 R. G . Pearson W. L. Waltz and E.J. Hart to be published. 54 M. Anbar Chem; Comm. 1966 416. Anbar The kinetic behaviour of the complexes has been correlated with the electron distribution in their d orbitals which are a function of their ligand-metal inter- actions.2@ The effect of the ligand on the electron distribution in the d orbitals does not explain all the experimental fmdings. First it is found that Zn2+aq is also reduced fairly rapidly although it has no vacancy in the d shell. Secondly it is evident that EDTA a ligand which has an effect on the electron distribution similar to that of H20 has a profound effect on the reactivity of some of its complexes. Moreover EDTA affects ions when no ligand field effects are ex- pected. Hydroxy-complexes behave in a similar manner as has been pointed out in connection with the hydroxy-complexes of aluminium and zinc.It has been shown further that the reactivity of both Nin and Cdn ethylenediamine complexes show the trend M(H2O):+% M(en),2+ M(en)32+.55 It is unlikely that the substitution of the last two H20 molecules by ethylenediamine has a significant effect on the electron affinity of the complex. The effect of certain ligands on the rate of electron transfer thus implies the participation of additional parameters. The complexes of the second and third series of trahsition metals have not been investigated as extensively as those of the first. Of the second series RuO\JH3):+ Rh(NH&:+ and R.h(bipy)?+ were studied% as well as PdC1,2- and Pd(CN)42-.14 All these complexes were found to react at diffusion-controlled rates. On the other hand Ru(CN)t- is extremely non-reactive in analogy to Fe(CN):-.l4 Several silver salts were also examined including Ag+aq Ag(NHd2+ Ag(CN), and Ag(EDTA)3- all of which also react at diffusion-controlled rates.l5?l6 The complexes of the third series studied include Os(NH,):+ which reacts at a diffusion-controlled rate;53 Os(CN)" which is very non-reactive like RU(CN)G~- and Fe(CN),&; IrClZ- and Irc1:- which react at diffusion-con- trolled rates;14 and Ir(NHd:+ which was found to be slightly less reactive (k = 1.3 x 1O1O 1.mole-' s~c.-').~~ All platinum complexes studied namely PtC142- PtCl,2- and Pt(CN)42- react at diffusion-controlled rates.14 Au(CN)~- and Hg(CN)z- react at fast but not a diffusion-controlled rates (k = 3-5 x 109 and 1.9 x lo9 1. mole-l set.-' respectively) whereas the reaction of Hg(EDTA)2- is diffusion-c~ntrolled.~~~~~ The primary products of the e-aq reactions with the bivalent transition-metal ions are evidently the univalent ion formed k s t at excited states as will be discussed later.The formation of Mn+ Co+ Ni+ Zn+ and Cd+ has been demonstrated through their ultraviolet absorption ~pectra.~~,~~~4*~56 Some of these were also observed in frozen solution by e.s.r.13,57,5s AqOaq formed in the reaction of Aq+ with eaq reaction has been detected spectrophotometrically by pulse r a d i o l y s i ~ . ~ ~ ~ ~ ~ 55 D. Meyerstein and W. A. Mulac to be published. 66 G. E. Adams J. H. Baxendale and J. W. Boag Proc. Chem. Soc. 1963 241. 67 W. Low and J. T. Suss Physics Letters 1963 7 310. 68 T. Feldman A. Treinin and V. Volterra J. Chem.Phys. 1965 42 3366. 5@ J. H. Baxendale E. M. Fielden and J. P. Keene ref. 9 p. 207. ' 0 J. H. Baxendale and P. L. T. Bevan ref. 36 p. 253. 589 The Reactions of Hydrated Electrons with Inorganic Compounds The chemical behaviour of these unique species has also been studied and as expected from their calculated redox potentia1s:l they react with various oxidants including permanganate,62 0xygen,5~,~ N,O and H20284,s5a and alkyl halide derivatives2s as well as with nitrite nitrate bromate iodate and cupric ions.65b In the latter study it was shown that the Zn+ Cd+ and Ni+ ions exhibit a salt effect of singly positively charged species. It has been further shown that Zn+ is more reactive than Cd+ and that both these univalent reducing agents are much more reactive than Ni+. As Ni+ has a d9 configuration and the addi- tional electron is accommodated in the incomplete d shell the latter result is not surprising.There is little information available on the fate of these univalent transition- metal ions in the absence of oxidizing additives.ssb They probably dimerize to form M22+ intermediates which disproportionate to give M2f + MO; the mechanism of these processes is still open to investigation. 2 The Mechanism of Reactions of Hydrated Electrons with Inorganic Compounds. Over 250 inorganic compounds have been studied for their reactivity toward e-aq.47 The information gathered from this extensive investigation which included representative compounds of 65 elements and has been summarised in the preceding pages allows us to predict the reactivity of many additional compounds with reasonable confidence.Further the e a q reactions have given us a simple reagent and a single type of reaction (electron transfer) to be tried on the vast diversity of inorganic compounds. This is an unprecedented chemical achievement. In the following pages we shall try to discuss the mechanisms of these reactions. A. The Energy of Activation of the Reactions of Hydrated Electrons.-The measurements of the activation energy of e-aq reactions provided invaluable information which must be discussed before one tries to elucidate the mechanisms of these electron-transfer processes. The energy of activation of a large number of eaq reactions has been measured both by competition kineticss6yg7 and by pulse r a d i o l y ~ i s . ~ ~ ~ ~ ~ ~ The great majority of eaq reactions investigated about 25 different compounds were found to have an activation energy of 3.5 & 0-5 kcal./mole.The eaq reactions which have been shown to have such a uniform energy of activation range in their rates of reaction from the diffusion-controlled limit to 16 1. mole-l sec.-l. Only three compounds were found to deviate significantly from this average 61 J. H. Baxendale and R. S. Dixon Z. phys. Chem. 1964,43 161. 62 J. H. Baxendale J. P. Kenne and D. A. Scott ref. 9 p. 107. 63 J. H. Baxendale J. P. Kenne and D. A. Scott Chem. Comm. 1966 715. 64 G. V. Buxton F. S. Dainton and G. Thielens Chem. Comm. 1967 201. 72 784. 66 M. Anbar and P. Neta Chem. Comm. 1965 365. 67 M. Anbar Z. B. Alfassi and H. Bregman-Reisler J. Amer. Chem. SOC. 1967 89 1263. 68 M. Anbar and E.J. Hart J. Phys. Chem. 1967,71 3700. (a) F. S. Dainton ref. 36 p. 3; (b) D. Meyerstein and W. A. Mulac J. Phys. Chem. 1968 590 Anbar but their different behaviour can readily be explained. Phenyl acetate was found to have an apparent lower A E,66 which is most probably the result of a secondary reaction with the electron adduct of trichloroacetate which was formed as the competing compound.69 The higher activation energies observed for Co2+ and Mn2+ ions17 are probably due to pre-equilibria with dimeric or polymeric ions formed in the pH range under study. It is expected that metal complexes of ligands which have a lower reactivity than their aquo-analogues will have apparent energies of activation higher than the actual A E owing to the tempera- ture effect of their dissociation constants.It was reported that the reaction of e-aq with H20 (k = 16 1. mole-l sec.-l) has an activation energy of 6.7 k~al./rnole.~~ This value has not been corrected for the temperature effect on the dissociation constant of water. The corrected value is less than 4 kcal./mole in agreement with other independent results.71 It has been pointed O U ~ ~ ~ ~ ~ that reactions whose rates are near the diffusion- controlled limit are expected to have an apparent activation energy of 3-4 kcal./ mole which is the activation energy of diffusion (dEdiff) of most solutes in water irrespective of their actual enthalpy of activation ( AEaCt). For reaction proceeding at rates slower than lo8 1. mole-l sec.-l the measured activation energy represents the actual enthalpy barrier of the reaction.There are two possibilities regarding the actual energies of activation ( dEact) of e-aq reactions. One is that the fast reactions proceed with AEact = ca. 0 like many ion-radical reactions and that the slower reactions have activation energies of 3-4 kca1.l mole. The invariance of the apparent activation energy over the whole range of reactivities would thus be fortuitous. The other possibility is that the majority of e-aq reactions fast and slow have an invariant energy of activation of about 3.5 kcal./mole. The value of dEdiff of eaq in water is expected to be considerably lower than dEdiff of conventional solutes in water as no cavity has to be formed in the former process. In most e-aq reactions the diffusion coefficient of e-aq con- tributes over 80% to the sum of diffusion coefficient.Thus it is likely that the measured dE of the diffusion-controlled e-aq reactions is a measure of dEact. One can find supporting evidence for this hypothesis in the fact that reaction H,O+ + H20 -+ H20 + H,O+ (18) (18) which determines the rate of diffusion of H+ in water has an activation energy of only 2.4 k~al./mole.~~ As the diffusion of H,O+ is the predominant factor in the diffusion in the reaction of H,O+ with eaq the observed d E = 3-2 for the latter reaction2* cannot be due to dEdiff. An energy of activation of 2.6 & 0.2 kcal./mole which has been reported recently for this reaction,74 has 1 3 ~ S. R. Logan and P. B. Wilmot Chem. Comm. 1966 558. 'O E. M. Fielden and E. J. Hart Trans. Faraday SOC. 1967 63 2975. 71 E. J. Hart personal communication. 72 S.R. Logan Trans. Faraday SOC. 1967 63 1712. 73 Z. Luz and S. Meiboom J. Amer. Chem. SOC. 1964 86,4768. 74 B. Cercek and M. Ebert J . Phys. Chem. 1968,72,766. 59 1 The Reactions of Hydrated Electrons with Inorganic Compounds not been corrected for the temperature effect on pKw a correction which would evidently increase this value. It has been found that the relative rates of reactions of electrons in ice with different solutes at 77°K are identical within experimental error with the relative rates of reaction in the same solutes with eaq in water at 300"~.'~ Several of the compounds examined in ice react in H20 at diffusion-controlled rates (including Cd2+ Cr0,2- NO,- and NO2-) whereas the others react more slowly. The results strongly suggest that all the reactions in ice which take place with 'mobile' electrons have practically no activation energy.AEaCt in ice may be different from that in water but it is still invariant for a large number of reactants of different chemical nature. It is suggested that most e-aq reactions in water have the same energy of activation independent of their rate and the nature of the substrate. It is unlikely that such an invariant energy of activation results from the energy requirements of a transition state involving different substrates of different chemical nature. One has to attribute therefore the observed invariant activation energy as an energy requirement of the hydrated electron. If e a q required a minimum amount of energy before being ready for transfer into a substrate molecule one could explain the experimental findings.B. The Primary Products of Reactions of the Hydrated Electron.-Reactions of e-aq invariably proceed by the transfer of an electron (19) from its site in the solvent into the acceptor molecule irrespective of the charge of the latter (n is a positive or negative integer or zero and A is an atom or a polyatomic molecule). e-aq + An+ 3 A n-l + (1 9) The electron has to transfer without violating the Franck-Condon restriction namely that none of the atoms involved changes position during the instant (> 10-14 sec.) of the actual electron transfer. In the case of diffusion-controlled reactions the acceptor molecule is believed to be in its vibrational and electronic ground states at the instant of electron transfer. It is most probable that after the molecule has accommodated an additional electron its interatomic distances at the ground state and in many cases its overall configuration will differ from those of the acceptor molecule.Thus at the very first instant following electron transfer a vibrationally and occasionally an electronically excited product has to be formed. A non-excited primary product can be envisaged only when an electron is incorporated into a vacant orbit of an atom e.g. reaction (20) and eaq + I 3 I- (20) even then a rearrangement of the solvation shell around the solute has to take place immediately (ti < The de-excitation of the vibrationally excited state may take place within the time of a few vibrations (< 10-l2 sec.). These excited species have to be con- sec.) following the electron transfer. L. Kevan J. Amer.Chem. SOC. 1967 89 4238. 592 Anbar sidered as intermediates and not as transition states because they evidently persist longer than a single vibration. It should be noted however that on the ‘chemical’ time scale one seldom considers intermediates with lifetimes less than 10-l2 sec.; thus these excited primary products may be overlooked from the stoicheiometric ‘chemical’ standpoint One cannot however neglect these excited products when considering the detailed mechanism of the electron transfer. These species can be demonstrated only in two cases when the de- excitation is accompanied by an irreversible cleavage of chemical bonds and when photons are emitted as a result of intramolecular electron transitions. There is very little experimental evidence for the formation of the excited primary products.The only indicative cases are the formation of Fe(CN),,H203- from Fe(CN),3- 76 and the long-lived excited C ~ ( d i p y ) ~ + . ~ ~ ~ ~ No chemo- luminescence which would be emitted on radiative de-excitation has been demonstrated in inorganic systems. The radiationless de-excitation may be facilitated by the intimate interaction of the product with the highly polar solvent. C. The Mechanism of Electron Transfer.-The reactivity of hydrated electrons toward inorganic compounds has been shown to be correlated in many cases with the oxidation potentials of the substrates. These oxidation potentials are a measure of the change in free energy on going from the ground state of the reactant to that of the reduced product. There are however additional factors which determine the rates of these reactions; these include inter alia the effect of ligands on the reactivity of complex ions or the requirement of a finite positive electron affinity of the substrate at its (initial) ground state.In trying to under- stand the mechanism of e-aq reactions one has also to consider the low and practically invariant energy of activation which strongly suggests a similar mechanism for most of these reactions. Such a mechanism should be consistent with each of the experimental findings cited above. The specific rates of reaction of inorganic compounds with e-aq range over many orders of magnitude. This is true also of homologous elements like the rare earths as well as of different complexes of the same element at the same state of oxidation. What is the reason for these differences in the rate of reaction? Three alternatives have been advanced to explain these differences in rate.According to the first it is the Franck-Condon restriction that requires consider- able free energy of activation to organise the electron acceptor to a configuration appropriate for electron t r a n ~ f e r . ~ ~ ~ ~ The second attributes the free energy of activation to a restricted transport of the electron through the ligands.15 Accord- ing to a third mechanism one assumes that electron tunnelling takes place and that the observed slow rates are due to transmission coefficients smaller than unity . l6 76 M. Haissinsky A. M. Kaulkes and E. Masri J. Chim. phys. 1966 63 1129. 77 W. L. Waltz Ph.D. Thesis Northwestern Univ. 1967. 78 R. A. Marcus Adv. Chem.Ser. 1965 50 138. 7s R. A. Marcus J . Chem. Phys. 1965,43 3477. 593 The Reactions of Hydrated Electrons with Inorganic Compounds An attempt has been made79,80 to draw an analogy between outer-sphere redox reactions which presumably follow the Marcus theory of electron transfer,81 and the reactions of the hydrated electron. It was assumed that the rate of electron transfer is limited by the energy necessary to rearrange the acceptor molecule as well as the water molecules of e-aq to a configuration compatible with the Franck-Condon restriction. This energy requirement is compensated by the gain in free energy of the electron-transfer reaction. At first approximation this theory79 predicts equation (21) where W the work neces- sary to bring the two charged reactants together makes a negligible contribu- tion to dF3.The reorganisation energy of both reactants A is of the order of 40-60 kcaI./m~le~~ and dF” is the gain in free energy in the reaction. This amounts to the differences between the standard redox potentials of e-aq and k = loll exp (- dF$/RT); dF$ = W + the oxidant minus the free-energy change of about 0 . 2 ~ due to the loss in trans- latory degrees of freedom going from the reactants to the p r o d ~ c t . ~ ~ ~ ~ According to this treatment the rate of reduction of a given inorganic complex by e-aq is primarily determined by the gain in free energy of the reaction. Although only very rough estimates could be made this theory requires any eaq reaction with a free-energy gain of less than 20 kcal./mole to be immeasur- ably slow and those with a free-energy gain of over 50 kcal./mole to be diffusion- controlled.If Marcus’s assumption is followed it would be implied that the Cd2+/Cd+ couple has a redox potential of about O-~V Ni2+/Ni+ about O - ~ V and Zn2+/Zn+ about l-Ov whereas all the slower reactants and their EDTA com- plexes have redox potentials in the range of 10--2*0v. These predictions are however rather unlikely following any physicochemical estimation of these redox potentials.61,82 A semiquantitative evaluation of the rates of reaction of EDTA complexes showed that the experimental data are in discord with the predictions of the Marcus theory.16 If a Franck-Condon restriction would reduce the rates of reaction this effect should be manifested in the activation energy.83 It has been shown however that even the reaction of Mnn-EDTA with e-aq which is one of the slowest studied has an activation energy less than 4 kcal./mole;68 moreover it seems that the energy of activation of e-aq reactions does not depend on the nature of the substrate.It is plausible therefore that in the case of MnIT-EDTA as well as of other EDTA complexes the favourable free energy of the reduction by e-aq makes the electron transfer feasible despite the Franck-Condon restriction at the ground state of the substrate. A second mechanism for electron transfer which we may name the ‘conduct- ance’ mechanism attributes the limitations in the rate of electron transfer to the central atom to a limited conducting capacity of the ligands.15 This mechanism 8o N. Sutin ‘Exchange Reactions’ IAEC Vienna 1966 7. 82 C. S. G. Phillips and R.J. P. Williams ‘Inorganic Chemistry’ Oxford Univ. Press 1966. R. A. Marcus J. Chem. Phys. 1956 24 966; Trans. Faraday SOC. 1960 29 129. M. E. Sacher and K. J. Laidler Trans. Faraday Soc. 1963 59 396. 594 Anbar which can be looked at as an analogue of an inner-sphere or bridged electron transfer requires an overlap of the orbitals of the hydrated electron and the ligand as well as those of the ligand and the central metal ion. According to this mechanism an electron is added to the orbitals of the ligand simultaneously with the release of an electron to the metal ion. The ligand which does not necessarily carry an excessive electron at any time is still required to have a favourable electronic configuration to allow such an electron transfer to take place. This mechanism has however two major drawbacks.First it does not imply a correlation between dF$ and dF” which seems to be a major factor in many cases. Moreover many EDTA complexes react at diffusion-controlled rates despite their suggested limited electron-conducting capacity and invariably these are cases of reactions with a high dF”. Secondly it does not explain why a certain ligand namely H20 enhances the rates of electrons compared with cyanide,15 or hydroxide which is not readily expected from the electronic structure of these ligands. From the available information on the behaviour of outer-sphere redox reactions there is no evidence for a superior electron-transfer capacity of H20. It is still possible that water is outstanding in being capable of bonding an eaq and a ligand of a metal ion at the same time and therefore best adapted to act as a bridge in this mechanism.Whether this bridged complex is a transition state or a short-lived intermediate is still an open question. Before going further with the discussion of the mechanism one has to answer the question ‘Do all hydrated electron reactions have a transition state? In any assessment of eaq reactions we have to make clear whether we assume that these like conventional reactions proceed via a transition state or whether they can be considered as intermolecular ‘adiabatic’ electron-transfer processes analogous to intramolecular ‘vertical’ electron transitions. A conventional transition state involves the existence of a state with a higher free energy than both reactants and products which has a lifetime of the order of a single vibration (10-13 sec.).The transition state of an eaq reaction if it existed is expected to have a lifetime of the order of an electronic transition namely sec. and therefore the equilibrium between it and the reactants might involve the electronic states only as the atoms do not move during the transition. As the suggested transition state should resemble both reactants and product we cannot envisage the electron in an excited state having the very low and invariant activation energy found for so many e-aq reactions. A conventional transition state is thus rather unlikely to be involved in e-aq reactions. Another approach from the standpoint of the potential-energy profile of the system shows that there is little effect of the potential energy barrier between reactants and products on the rate of electron transfer.The hydrated electron is held in the solvent by an energy of hydration of about 40 kcal./mole and there are additional potential-energy restrictions owing to the repulsive force of the native electrons of the acceptor molecule. The outer sphere of solvent molecules and gegenions is also expected to contribute to the potential barrier. All these should result in a potential-energy barrier much higher than the invariant 3.5 kcal./mole observed. It may be concluded therefore that owing to its small 595 The Reactions of Hydrated Electrons with Inorganic Compounds mass the electron manages to tunnel through this energy barrier without forma- tion of a transition state. Electron-tunnelling is evidently dependent on the height and width of the potential barrier and also depends on the overall change in potential energy.83,84 The width of the barrier is a function of the distance of closest approach of the e-ag to a vacant orbital.A large dF" is also accompanied by a decrease in the height of the potential barrier thus the electron transfer may take place over longer distances. It has been pointed out that in certain raq reactions the cross-sections for reactions with oxidants having very high oxidation potentials (dF" > 100 kcal./mole) exceed the geometrical dimensions of the reactants strongly suggesting electron-t~nne1ling.l~ The restriction of this path- way by the Franck-Condon principle for e-aq reactions with a large dF" is negligible owing to the availability of sufficient free energy to allow electron transfer into excited states.In the quickly reacting EDTA complexes cited in sections 1C and 1H AF" is most probably large enough to allow tunnelling with a probability of unity. The electron-tunnelling mechanism may however also explain the behaviour of many of the slower reactions. EDTA or (ethylenediamine) complexes differ from aquo-complexes by their dimensions much more than by their standard redox potentials. When dF" is not very large it being remembered that the transmission coefficient is an exponential function of the width of the potential barrier,84 the increase in distance between the hydrated electron and the central atom may result in a substantial diminution of the transmission coefficient. In its simplest form the tunnelling mechanism should be feasible only when the energy level of the accommodated electron in the acceptor molecule is lower than the energy level of the hydrated electron.In other words e-aq would be expected to react only with substrates which have an electron affinity of over 40 kcal./mole before any change in their atomic configuration took place. This would be true of most diffusion-controlled reactions but might not be applicable to many substrates that have sufficient electron affinity only after reaching the new ground state including a rearrangement of their solvation shell. We have therefore to modify the description of the electron-transfer mechanism to include these cases. Let us assume that the substrate in its ground state has a relatively low electron affinity. When such a substrate approaches a hydrated electron to within a critical distance there will be a certain probability of finding it in the substrate a probability much smaller than unity and which will drop to zero as soon as the hydrated electron recedes.This probability is synonymous with the tunnell- ing coefficient discussed above. Now if the substrate molecule rearranges to a configuration which accommodates the additional electron in a lower level the probability of its escape is diminished. As soon as these rearrangements includ- ing those of the solvation shell result in stabilisation of the additional electron at an energy level below 40 kcal./mole the reaction goes to completion. The a4 W. L. Reynolds and B. W. Lumry 'Mechanism of Electron Transfer' The Ronald Press Co. New York 1966. 596 Anbar overall rate of reaction according to this mechanism will depend on the level of the initial electron trap as well as on the rate of rearrangement to a stable product.If the former parameter is the more important one the reactions will have no activation energy of their own. If the activation energy of intramolecular rearrangement of most molecules is within the range of 3-4 kcal./mole this mechanism would again be consistent with the experimental results. This mech- anism is the only plausible explanation for many of the reactions of e-aq with aromatics5 or carbonylic compounds,86 the rates of which range over many orders of magnitude but which proceed with a small and practically invariant AES.67 Many of these substrates have insufficient electron affinity to compete with the hydration of the electron unless the product rearranges and becomes solvated.From the standpoint of the absolute rate theory one may look at the course of a reaction of the latter type as tunnelling followed by a series of transition states each with a lower potential energy separated from each other by free- energy barriers which consist predominantly of entropy of activation. The inclusion of the intramolecular rearrangement following the electron- tunnelling in the requirements for completion of certain e-aq reactions makes electron-tunnelling plausible as the first step in any of the eaq reactions hitherto investigated. There is at present no case which requires ‘conductance’ mechanism and which could not be explained by electron-tunnelling. The efficiency of water as a bridging ligand may be explained by a substantial decrease in the width of the potential barrier in the reactions of aquo-complexes which is accompanied by an increase in the transmission coefficient.The ‘classical’ transition state involving the rearrangement of a molecular configuration before the electron transfer seems to be even less likely although it cannot be excluded. The prevalence of electron-tunnelling in e-aq reactions does not imply that other ‘outer sphere’ electron-transfer processes in inorganic chemistry proceed by the same mechanism. It is conceivable that while the electron donor re- organises to facilitate electron-transfer at the lowest energy expense a re- organisation of the electron acceptor also takes place to facilitate the incorpora- tion of an electron into the lowest level possible.The difference between these and eaq reactions is that in the latter the electron is much less restricted and localised. None of the conventional reducing agents in homogeneous aqueous solution has an electron bound by as little as 40 kcal./mole. It should be stated however that there is no experimental evidence or theore- tical reason that all eaq reactions should proceed by the same mechanism. In fact each of the main three mechanisms discussed namely the two ‘classical’ mechanisms and the electron-tunnelling may contribute to the different electron- transfer reactions. 3 Conclusion The reactions of hydrated electrons have contributed to three different aspects 85 M. Anbar and E. J. Hart J . Amer. Chem. SOC. 1964 86 5633. 86 M. Anbar and E. J. Hart J. Phys.Chem. 1967,71 3993. 597 The Reactions of Hydrated Electrons with Inorganic Compounds of inorganic chemistry. First the study of these reactions has given us a new insight into electron-transfer processes and has raised the possibility of the existence of electron-tunnelling in chemical processes. Secondly in homologous series the rates of e-aq reactions may be used as a semiquantitative measure of the relative electron affinity of various species in aqueous solution. And lastly but not least hydrated-electron reactions provided inorganic chemistry with an extensive series of novel reduced inorganic compounds in aqueous solution. These compounds are presently in their very early stages of study but they will undoubtedly become an important part of modern inorganic chemistry. 598

ISSN:0009-2681    

DOI:10.1039/QR9682200578

出版商:RSC    

年代:1968

数据来源: RSC