摘要:

R.I.C. Reviews R.I.C. Reviews, published twice yearly, will review areas of chemistry of interest to the chemist who has no specialist knowledge of the field under review, but who wishes to keep abreast of the growth of chemistry as a dis- cipline. These reviews will prove useful to students in familiarizing them- selves with a particular field. R.I.C. Reviews will interpret the significance of chemistry in a wide context and will publish articles on the economic, social and historical aspects of chemistry, as well as on the research and applied sectors. Suggestions for future titles are welcomed. Prospective contributors should write to the Editor, enclosing a synopsis (of about 250 words) indicating the scope of their subject. The preferred length for reviews is 8,000 words. Annual Subscription: €2 (R.I.C. members, Ll)

ISSN:0035-8940    

DOI:10.1039/RR96801FX001

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

ROLE OF TRANSITION METAL IONS IN BIOLOGICAL PROCESSES R. J. P. WILLIAMS, M.A., D.Phil, A.R.I.C. catalysed by CU(II) -+ CU(I). Wadham College and lnorganic Chemistry Laboratory, Oxford The transition nietal ions which are available to a biological system are those from scandium to copper and from yttrium to molybdenum in the periodic table. Of these, manganese, iron, cobalt, copper, and molybdenum are of major importance, while vanadium, chromium, and niobium are of much less interest. The virtual absence from living systems of some relatively abundant metals, such as chromium and nickel, has been explained on the basis of the preferential binding of these metals to the regular sites of miner- a1s.l Zinc and cadmium have been excluded from this review, although these elements are sometimes included with transition metals, for they are quite distinct in both chemistry and biochemistry.Zinc and, to a lesser degree, cadmium are metals which act as strong Lewis acids, in both chemical and biochemical reactions, but do not take part in reactions involving valence changes. In strong contrast the major role of transition metal ions is in oxidation-reduction reactions involving such change. Examples are : (9 0 2 + HzO, Nz +NH3, catalysed by the valence state changes Fe(I1) + Fe(m), Cutr) -+ CU(II). (ii) Dehydrogenation, e.g. ascorbate -+ dehydroascorbate, catalysed by the changes Cu(r~) -+ CU(I), MO(VI) -+ Mo(v), Fe(rr1) + Fe(Ir). (iii) HzO -+ 0 2 , probably catalysed by valence state changes of Mn. ( i l l ) (4 Ribose -+ deoxyribose catalysed by CO(I) + CO(III).In each case the metal ion is bound in an enzyme or coenzyme. Those enzymes into which transition metal ions have been substituted in vitro have also been excluded from this review, although this substitution has such advantages as to suggest that it should be made whenever activity can be retained. Transition metal ions have such physical properties as to make them excellent probes of their envir~nments.~J Thus they give rise to EPR signals, CU(II), Fe(IrI), Mo(v); Mossbauer signals, Fe; d --f d spectra, CO(II) ; proton NMR broadenings and/or shifts, Mn(Ir). BINDING OF TRANSITION METAL IONS As yet few binding centres of these cations have been fully delineated. Table 1 shows the present state of our knowledge.The binding groups will be seen to be strongly polarizable donors such as >N-, RS-, or phenolate, as had Williams 13 Element Mn Table I : Bindinn of transition metal ions Fe co c u Mo Zn(Cu) Zn(Co) Zn(Co) Zn(Cd) Tyrosine(2-3), N-base. R2S, - i m i d azo I e, por p h y r i n . I m idazole? Po r p h y r i n , -i m idazol e . Sulphur-ligands. Por p h y r i n , - i m id azol e, -CO,? Corrin ring, carbanion been anticipated from the general chemistry of the metals. An exception is provided by the Mn(I1) ion which binds preferably to weaker donor ligands such as carboxylate and phosphate. It is a characteristic of Mn(I1) that, of all the transition metal cations, its chemistry most strongly follows that of Group IIA cations such as Mg(I1).Biological systems are no exception and Mn(I1) can often replace Mg(1r) especially in reactions involving phosphate. Because of its considerable difference from the other transition metal cations, the biochemistry of Mn(r1) will not be discussed further in this review. Al- though not much experimental work has yet been done it would appear that the little Cr(II1) and Ni(r1) in biological systems is associated with ribonucleic acid (RNA)4 and is there bound to pho~phate.l1~ The iron in RNA is certainly very different in character. Although the binding groups for the metals may be conventional their stereochemistry is not. Table 2 gives some of the suggestions which have been made on the basis of physico-chemical properties.Undoubtedly these geometries are a consequence of the inability of the protein to provide highly symmetric binding sites, or even usual ligand-bond distances, through the need for it to conform to internal structural requirements. In the opinion of the writer, this misfitting has great functional significance. Howevei, it would be an obvious mistake to suppose that the description of the metal and its immediate ligands is sufficient to explain the function of the metal- containing systems. The next near-neighbour groups of the metal complex may well be reactive protein side chains. In either acid-base or redox reactions these groups could react in a cooperative manner with the metals. Before discussing the highly catalytic systems, those cases where the metal is concerned with storage or transport only will be examined.Here the transition metals can be compared with other metals when they have only a structural role as is true of the zinc in amylase and some of the zinc in dehydrogenases. l6 14 >.-- groups (RS-), -NH2 or -imidazole - NH2 or -imidazole and RS- -NH2 or -imidazole (- RS-):! Protein Pyruvate decarboxylase ., Conalbumin Cytochrome c Hemerythrin Heme oxygen carriers Fe r r ed oxi ns Oxidases, catalases and peroxidases B12 enzymes ' BI ue' enzymes, hemocyan in Xanthine oxidase Insulin Carboxypept idase Carbonic anhydrase Alcohol dehydrogenase Probable ligands Not known; imidazole? Method Electron paramagnetic resonance (EPR) Chemical, EPR Spectra, EPR Chemical X-ray Chemical, EPR Chemical and spectra X-ray Spectra, EPR EPR EPR Spectra Spectra Spectra R.I.C.Reviews Table 2: Irregular physical prc lerties of metalloenzvmes Protein Metal ion Fe2f Fe3f Hemoglobin, m yog lo bi n H emog I o b i n , m yog I o b i n Ru bredoxin Conalbumin Aromatic substrate Fe3+ Fe3f Fe3+ oxidases Cytoc hrome chain Ferredoxins Fe3f Fe3+ c u2+ Ten different ‘blue’ proteins Car boxy pe pt id ase Carbonic Anhydrase Insulin Zn2+ (CO2+) Zn2f (CU2+) C03f MO(V) Fe(lll) 612 enzymes Xan t h ine oxidase Cytoc h romes Method Geometry known or inferred X-ray 5-Coordinate Very low symmetry, maximum rhombic field X-ray and EPR I EPR (g=4*3) Very iron not strong in-plane axial field, 7 7 Not understood, low symmetry Low symmetry-not tetragonal 5-Coord i nate? 5-Coordi nate Trigonal field 5-Coordi nate pyramid ? Lower than axial symmetry Rhombic - Table 3: Metal and substrate, storage and transport proteins Fe He mog I o b i n Hemerythrin Chlorocruorin J M yog lo bi n Conal bu min Transferrin cf.Ferroxamins Ferritin - c u Hemocyan i n 0 2 carrier Ceruloplasmin Hematocuprein and mitochondria1 copper proteins 1 Cu carrier? Hemocuprein - V ‘Hemovanadin’ - TRANSPORT AND STORAGE 0 2 carriers 0 2 storage Fe carriers Fe storage I Cu storage? Carrier of oxidizing equivalents? EPR (g=4*3) EPR (g= I *94) Spectra, EPR, CD Spectra EPR Spectra, NMR EPR EPR There are two types of transport and storage proteins involving metals: those which control metal ion concentrations and could be called sequestering proteins ; and those which control concentrations of certain substrates such as oxygen and carbon dioxide (see Table 3).In the first group, the iron- and copper-carrier proteins which are mobile in the blood of mammals, have been most thoroughly studied as they are highly coloured. (It appears that there are other proteins which control the concentrations of calcium, mag- 15 Williams L - Ref. EPR (g=4* 3) EPR (g=4*3) 8 9 10 II I I 12 13 14 15 - nesium, zinc and perhaps most heavier metals in biological systems, although these are not well-recognized as yet.) It is not always possible to say that these metallo-proteins are just con- cerned with the control of the supply of a particular cation.For example, the iron-transporting protein, transferrin, could act catalytically in ‘placing’ iron in porphyrin to make heme. Also the copper-carrier protein, caerulo- plasmin, may act in monitoring the concentration of transferrin. Thus it has been postulated that, as the copper in caeruloplasmin undergoes ready oxidation and reduction, reacting with molecular oxygen and iron, it controls the Fe(rI1) concentration in serum and thence transferrin formation. l7 Even in the role of sequestering agents proteins can have major control over biological activities.For example, they can be involved in membrane trans- port or they can so lower the concentration of a cation as to prevent infection with another biological system. Conalbumin (egg-white) is an iron-binding protein and could prevent bacterial invasion of egg-yolk by removing iron. l8 The metal-proteins that carry oxygen often show weak catalytic properties also. Here, however, transport and storage are of such dominant importance that they are obviously the major functions. Copper and iron proteins are involved in a total of three series of such compounds. They are the por- phyrin-iron oxygen carriers, e.g. hemoglobin and myoglobin; the non-heme iron proteins, e.g. hemerythrin; and the copper proteins, e.g.hemocyanin. Just three similar series of proteins are involved in oxidase and electron- transfer processes. Already common evolutionary stages are being sought, assuming oxygen transport to be a relative newcomer.19 The metal geometry in the oxygen carriers would appear to be ideally designed for function. Thus, in essence, the metal has to undergo addition, which conventionally involves substitution by molecular oxygen. In fact the five-coordination of Fe(r1) in myoglobin7 (and hemoglobin) makes it the S N ~ intermediate for a conventional octahedral substitution reaction, so that the curious Fe(1r) geometry appears to be directly designed to overcome the normal activation energy for addition. Although the geometries of the active sites in hemerythrin and hemocyanin are unknown, the fast reaction with oxygen and the very high redox potentials of these proteins (see later) suggest that the metal ions in the proteins also have an open coordination position.In both the latter series of compounds two moles of metal absorb one molecule of oxygen so that in this respect they are unlike the heme-containing oxygen carriers. The reaction can be written: 02 -1- [+cu cue] - [5cu 02 cu c3 The physical properties of the oxygenated species are extremely confusing. 2 0 Oxygenated hemocyanin has the visible spectrum normally associated with a cupric ‘blue’ protein yet it is formed by the reaction of oxygen with a CU(I) dimer. Oxygenated hemerythrin is spectroscopically very like the Fe(Ir1) compound.Oxygenated myoglobin has the spectrum in part of Fe(11) and in part of Fe(rrr) heme, and its Mossbauer spectrum has a curiously high quadrupole splitting. Neither the copper nor the iron proteins give electron 16 R.I. C. Reviews paramagnetic resonance (EPR) signals. The situation is closely reminiscent of that in the oxygen, hydrogen21 and nitrogen22 uptake reactions of model transition metal complexes, e.g. ___+ IrX4 { - H, IrX4H2 In each case it is permissible to look upon the reaction product just as an addition complex or as involving change of oxidation state. The distinction here may be merely the apparent one of electron distribution (resonance) or could be one of change of nuclear positions or even of spin states, i.e.a definite equilibrium in the product. The discussion of these different possi- bilities has not been clearly thought out either in inorganic or biological systems but it is of immense importance to the understanding of the catalysis of reactions. We shall refer to the metal in these situations as being in a condition of ‘equilibrated’ oxidation state. (One is tempted to wonder if the debate over classical and non-classical carbonium ions is bound to occur again over labile inorganic systems where both electron and nuclear distri- butions could be changing very rapidly.) Oxygen absorption by hemoglobin has certain other features which help the study of the general complexities of multi-enzyme systems involved in oxidation processes and which have more than one metal atom.The hemo- globin unit is composed of four protein chains each carrying a heme group. Uptake of oxygen at one iron atom increases the affinity of the other iron atoms for oxygen although they are some 25 A away. Now, on absorption of oxygen, the iron atom changes its electronic structure, pairing four elec- trons. In model complexes such a change is known to alter metal-ligand bond lengths by at least 0.1 A. Such a structural change could be amplified grossly in the geometry of the protein, resulting in important cooperative changes at long distance. However, it is the very fact that such small altera- tions bring about a considerable change in reactivity 25 A away which is important in a discussion of some oxidative enzymes which involve more than one metal atom.The general heading under which such effects are discussed is allosterism. Apart from the general storage and transport systems mentioned above there are a variety of metal compounds of similar function which are appar- ently common to a more limited series of biological systems. Many such metal compounds are poorly studied as yet. Iron, for example, is held in bulk reserve in mammals in ferritin, which is not too dissimilar from a Fe(r1r) hydroxide polymer wrapped in a protein.23 There are also a number of copper proteins in liver and brain cells which are of no known function and could be storage proteins. The liver copper protein24 strikingly resembles in amino-acid composition the cadmium protein discovered by Margosches and Vallee.25 The discovery of vanadium in tunicates, of niobium in MoZguZa manhattensis and of traces of other metals in a great variety of biological materials has only just begun to be followed up by detailed studies.Williams 17 Carriers for molecules other than oxygen are poorly studied as yet. It could be that the transferrin and hemoglobin reactions with bicarbonate are important to carbon dioxide movement. l8 The most likely transport problem where metallo systems will be found to be important is in the transport of anions, e.g. phosphate, across membranes. This is still a little explored subject. REDOX REACTIONS The redox function of transition metal ions is not simple, for the metals act in a diversity of catalytic roles: electron transfer; oxygen atom and hydroxyl group incorporation; hydrogen ion and atom removal.In complex systems, to which we shall turn later, the metal containing units are organized in many-headed and multi-enzymes. The redox activity of these units is con- nected to energy conservation, oxidative- and photo-phosphorylation, as in mitochondria and chloroplasts respectively, in a manner which is not yet understood. Because of this complexity the discussion is divided as follows : (i) hydride transfer, (ii) oxidation by oxygen incorporation (single-headed enzymes), (iii) electron transfer, (iv) many-headed enzymes, i.e. those enzymes which carry out a single reaction between two or more substrates but do so by activating the substrates at well separated sites on the enzyme.Hydrogen removal from a substrate by molecular oxygen is a very common example. (v) multi-enzymes, i.e. a system of enzymes which are bound to one another and therefore behave as a single unit-for example electron-transfer frequently occurs along a chain of enzymes in a tightly bound particle. Both the many- headed and the multi-enzyme systems can incorporate all the other three activities in a sequence: Oxidation by 02 ' - or by H,O, 1 (Phosphorylation, rnulti - enzymes only) It is the ability of a biological system to place these three types of metal catalysts in a particular spatial relationship which has generated a major difference between the biology and the test tube chemistry of transition metal catalysts. Physical properties and redox potentials The physical properties of the metals in the redox enzymes have already been described as unconventional and we have associated this peculiarity with unusual geometry.In model systems it is known that an enforced departure from a conventional geometry usually destabilizes the higher valence state of an element more than the lower, whence the redox potential becomes more p~sitive.~' The redox potentials of a large number of copper and iron proteins are known. In the case of the copper proteins of unusual physical properties-the copper blue proteins-the redox potentials are R. I. C. Re views 18 invariably high, around + 0.4 V. This is consistent with the picture of a non-ideal environment for the cupric ion, i.e.some considerable distortion of the tetragonal field in which this cation is generally found in model sys- tems.1° The redox potential of the oxygen carrier, hemocyanin, is higher still, suggesting low irregular coordination. The iron proteins must be considered in two major groups-the heme and non-heme systems. Figure 1 shows that amongst the heme iron proteins, the low-spin cytochromes often have much higher redox potentials than the corresponding model heme complexes.28 Once again we may presume that the Fe(1rr) state has been destabilized by steric misfitting. Other physical properties, Mossbauer spectra and spin state equilibria, for example, indicate that the misfitting could lie in the inability of the protein groups perpendicular to the heme to come to within normal bond distances of the iron.29 In the myoglobins and hemoglobins, discussed above, the misfitting-as shown by crystallographic data-is such as to allow only one basic group to reach the heme iron. As a consequence, these heme-proteins are high-spin.The study of series of model complexes has suggested that as a second protein donor group is allowed to approach closer to the iron, the redox potential will at first become more positive and then more negative-a maximum in redox potential being associated with the change in spin states29 (Fig. 2). The size of the maximum in redox potential is increased if one of the ligands is ‘unsaturated’ as such ligands preferentially stabilize the lower valence state.The very high potentials of some heme proteins can now arise from three factors: unsaturation in the heme, e.g. heme a; unsaturation in the protein groups, e.g. thioethers; and misfitting of the donors from the protein to the iron. (On top of such considerations we should also add perhaps the unknown effect of the total environment of the heme moiety.) However, the fact that with the same porphyrin a slowly graded series of redox potentials is observed in a variety of proteins (in which it is known from other physical measure- ments that the iron binding is curiously balanced between spin states-see below), suggests that the misfitting is of great importance. Apart from the cytochromes, hemoglobins, and myoglobins, there is a wide variety of heme enzymes of the high spin type-oxidases, peroxidases and catalases that are spectroscopically like myoglobin.26 It has been con- cluded that the heme iron has as immediate ligands a number of very different groups on one side (e.g.imidazole or carboxylate) and may be very low indeed, - 0.4 V, to more usual values, e.g. + 0.1 V. The low ‘open’ on the other. The redox potentials of these proteins range from potentials could well be brought about by carboxylate groups while higher potentials would result from neutral ligands near the metal. The other major group of iron proteins, non-heme iron, are not as yet a well-defined chemical group.s Although many of them have physical proper- ties in common and are said to be ferredoxin-like there are others, e.g.phenol-oxidases, of rather different properties. In the ferredoxin proteins redox potential, from - 0-5 to + 0.2 V, suggests that a great variety of themselves it is certain that the iron is bound to sulphur. The wide range of other ligands are also involved. On the other hand it is known that in the non-heme iron oxygen-carrier protein, hemerythrin, the iron is not bound Williams 19 Number of models 10 5 I I 10 30 20 Low T spin .f High spin LL d- Ligand field --P High- spin Low -spin Fig. 2. Schematic representation of the variation of redox potentials of model (lower curve) and biological (upper curve) heme complexes with changing ligand character.On the curves some probable bases that are coordinated to the heme iron are given and on the upper curve the relative positions of different cytochromes and oxygen carriers are shown. to such sulphide ligands30 and is more probably bound by imidazole (and carboxylate). Many of the oxidases seem to resemble hemerythrin, especially in the Fe(rI1) state, more than ferredoxin. It could well be that the oxidases are related to hemerythrin much as heme-containing oxidases are related to myoglobin, the differences between the two types of protein resting in rela- tively small changes close to the iron. There are no electron transfer proteins, apparently, which belong to the hemerythrin series, in contrast to the cyto- chromes, presumably because the ligand field around the iron does not produce low-spin complexes and will not transfer electrons.Of the other transition metal protein complexes, vanadium31 and man- g a n e ~ e ~ ~ compounds have not been studied in sufficient detail as yet to merit lengthy summary. On the other hand, through the work of Bray in parti- ~ u l a r , ~ ~ , ~ ~ our knowledge of molybdenum proteins is increasing rapidly. The molybdenum, presently thought to be bound to s u l p h ~ r , ~ ~ ~ ~ ~ cycles between oxidation states v and VI. In the former it gives rise to an EPR signal of low symmetry. The study of these proteins has been made complex for they contain much iron as well as some flavin and thus there are several sources of EPR signals. Cobalt-containing enzymes, except those associated with vitamin B12, are not well-known.In the BIZ coenzyme/enzyme complexes it has been postu- lated that the CO(III) can become five-coordinate, an unrecognized geometry for Co(11r) except in models based upon the coenzyme (Fig. 3). In the co- enzyme there is also the suggestion that when bound to its enzyme the cobalt-carbon bond is weakened so that the CO(III) may readily change to 21 Williams Fig. 3. An outline diagram of the vitamin B,, series of complexes. O n the left the series of cobalamins is illustrated where Y is usually benzimidazole while Z can be any ligand, including carbanions. On the right is the series of cobinamides showing two typical ligands. In each case the corrin ring is represented by the circle.a square coordination. If this is so then the geometries of the coenzyme, nominally always written as a CO(III) compound, correspond to the three stable geometries of the three valence states, CO(III), CO(II), and CO(I). The spectrum of the corrin ring, which is bound to the cobalt, most closely resembles that of the Co(1r) model compounds, so that one is left with the impression that the enzyme induces a condition of the cobalt and its ligands which would facilitate rapid redox changes. A special feature of the total environment of all the four major redox metals of biology-iron, copper, cobalt and molybdenum-is seemingly generated by their highly polarizable ligands and their distorted geometries. Electron-transfer to and from the metal and/or substitution reactions at the metal are made to be of relatively low activation energy.Thus, even when studied in an environment quite unlike that found in biology, cytochromes react rapidly in electron transfer.34 Similarly the exchange of ligands from around cobalt in vitamin B1, is extremely fast, unlike that in most CO(III) complexes of inorganic origin. Such observations have led us to postulate that the metal ions in the enzymes are poised in a condition most resembling the transition states of their normal complexes and very different from their conventional ground state complexes. 28 This proposal is discussed in more detail later. It is a well-known observation that pyridinium compounds react as if the p-carbon were strongly positively charged ; for example, they add hydride ions : QJR’ R I R I R I 7 Hydride transfer and rearrangement reactions - OR’ a OR’ Y NAD + R H- ~ Flavin NADH R I H 22 I COZH R.I.C.Reviews In biological systems the most common pyridinium compound which carries out this reaction is the pyridine nucleotide coenzyme, NAD+. A similar reaction is observed with flavin coenzymes. A metal cation is obviously able to stabilize the reduced forms of the pyridinium or the flavinium molecule by providing positive charge perturba- tion of the heteroaromatic rings. Thus it could do so by coordinating to a basic nitrogen as shown in Fig. 4. (Note that it is possible to look upon a coenzyme in an enzyme as being near in geometry to its normal transition state.) H W R EEL!? 2 Cyt.Fe (m) Fig. 4. The attack of hydride on a pyridine nucleotide aided by a cation (here zinc). An alternative representation would place a water molecule between the zinc and the nitrogen. It is therefore of great interest that many NAD and flavin enzymes which act by hydride transfer are found associated with zinc (particularly NAD systems), manganese, molybdenum and iron. By replacing zinc with cadmium and examining the absorption spectrum of the resulting inactive dehydro- genase, it has been shown by Vallee that the zinc is probably bound to RS- in one of these enzymes-liver alcohol dehydrogenase, which is a free enzyme in cytoplasm. In complex particulate systems (see p.30) the hydride enzymes, e.g. lactic and succinic dehydrogenases, are directly coupled first to iron proteins of the ferredoxin type and then to cytochromes: via via ' Ferredoxin ' 2 Cyt. Fe (II) Ferredoxin ' The vitamin B,, series of enzymes is also involved in several reactions which presently may be classified as hydride transfer although the product of the overall reaction may be either reduction or rearrangement (Rearrange men t) e.g. the methyl-malonate/succinate rearrangement. (Other two-electron trans- fer reactions such as methyl transfer can be similarly represented). The cobalt of the BI2 complex is exceptionally well adapted for the catalysis of these reactions in that it is readily reduced in either one two-electron or two one- electron steps, The cobalt is bound to the corrin ring and to a carbanion in Williams 23 I I I I I I I 2090 21 10 2130 cyanide, C,H-, C,H;, Stretching frequency R - Co- CN (cm-’) Fig.5. The variation of the C-N stretching frequency in the vitamin B,, complexes of formulae in R-Co-X. The different equilibria plotted are: 0, H,O + CN-; +, H,O + Nucleotide; R-Co-CN with log,,K, where K is the equilibrium constant for the replacement of X by Y 0, H,O + CN-; and A, nucleotide + CN-. R, reading from left t o right i s nucleotide, CH; and C,H,. Free cyanide has a stretching frequency of 2078 cm-I. many of the enzymes and this binding further conditions the cobalt for the reaction.35 Thus it has been observed that (a) trans to the carbanion, 2 in Fig.3, the cobalt accepts almost all ligands, Y , equally and all are bound very weakly (Fig, 5). It is this enormous trans effect which causes CO(III) to be effectively five-coordinate in extreme cases. (b) The corrin ring itself takes on an electron density and possibly a conformation more like that found in the lower CO(I) and Co(11) states than in the normal CO(III) state complexes. CO(III) is the formal valence in this series of compounds but the actual way in which the electron distribution is best visualized is open to debate (see above). Co ( I l l ) R- - Co (rr) R’ +--+ complexes : +--+ [Fe(u)] (SR) (RS) [ F e ( m ) ] (JR), Again we are up against the problem of ‘equilibrated’ oxidation states. In writing reaction mechanisms care must clearly be taken to avoid argument Co (I) R-i- Thus Co(r11) with its strongly-reducing ligand can be pictured much as the iron in ferredoxin bound to sulphur, or indeed of the iron in oxygenated Fe(1) (SR), R.I.C. Reviews 24 Table 4: Single-headed redox enzymes Enzyme Function Metal - Fe(heme) Catalase (one o r four independent heme units) Peroxidase Tryptophan pyrrolase Fe(heme) Fe( h e me) Tryptophan + 0; + kynu renine* Oxidation by H20z o r 0; +I1 /‘H/OCO*OH I /\/OH II Fe M etapyrocatecholase \/\OH \/\OH lnositol + glucuronate* Uses pyridoxal Fe Cu(ll) lnositol oxygenase Diamine oxidase C 4 ) Tyrosinase (sometimes contains 2-4 Cu atoms) //\OH //\OH _ _ _ - - - - - - - \\/ \/OA I - - I Electron transfer, see Table 9 Electron transport? Electron transfer, see Table 9 Not known - - - - - - - - - - - - - Cytochromes, e.g.c Rubredoxin Pseudomas blue protein Azurin _ - - - Fe ( h e m e) Fe[RS- (?)I c u c u * 1 8 0 incorporation from 1 8 0 2 o r H21802 has been shown thus demonstrating direct oxygen attack on the substrate.26 between say carbanion and carbonium ion mechanisms. The cobalt catalyst, metal plus ligand, has a concerted reaction path open to it but which group is the acid and which the base is far from obvious. It is also clear that a homolytic path is not ruled out in a four-centre reaction. There are several other metal enzymes which do not require B,, but carry out rearrangement reactions, e.g.aconitase which requires both iron and a sulphydryl system, and it could be that these rearrangements also require the metal as part of a two-electron reagent. Single-headed oxidases? Whereas hydride transfer in biological systems is often associated with zinc ions, which are incapable of undergoing redox reactions, oxygen and electron transfer are mediated by iron and copper (Table 4) and, to a lesser extent, by manganese and molybdenum. As pointed out earlier the first two metals also absorb oxygen reversibly. It could well be that the carrier proteins are closely related to those enzymes which act in oxygen molecule or atom trans- fer from oxygen and hydrogen peroxide directly to a substrate (incorpora- t i ~ n ) . ~ ~ Many of these enzymes are thought to be free in cytoplasm.Clearly, oxygen atom transfer direct to a substrate must depend upon the -f Under this heading we shall describe the simplest oxidases which contain either one active site or several non-interacting active sites. In Table 4 some single-headed electron transfer proteins have been included for comparison. The latter belong to multi-enzyme systems. Williams 25 concerted action of two substrates, oxygen (or hydrogen peroxide) and the substrate to be oxidized, at one region of the protein which itself must be very near the site of oxygen (or hydrogen peroxide) absorption. (Reversible oxygen absorption needs only one site.) As this form of oxidation does not lead to energy conservation, i.e. the free energy of oxidation is lost (as heat, for example), interest centres on the nature of the environment of the metal ion which, on uptake of oxygen (or hydrogen peroxide) allows it to act in oxidation by oxygen molecule transfer, or by hydroxyl (oxygen atom) transfer, or simply as an oxygen carrier, as in the earlier section.The immediate coordination sphere of the metal is partly responsible for all these reactivities. It would appear that oxygen is initially absorbed in the same way by many of the proteins, especially heme-containing oxidases, for the intermediate complexes of oxygenases are often very like oxygenated compounds. 3 7 9 38 Thus the immediate coordination sphere may be common in this type of protein, i.e. the binding ligands of the metal leave a single coordination posi- tion virtually open for oxygen or hydrogen peroxide reaction.The close similarity between the structures of the metal complexes of different proteins with different activities, e.g. myoglobin and tryptophan pyrrolase, have led to the further proposal that it is the next near-neighbour protein side chains that are important for their differences. It is noteworthy that the protein side chains around the heme of the oxygen carriers of known structure7- myoglobin and hemoglobin-are both hydrophobic and poor electron donors. Thus, critical factors in the reactivity of the bound oxygen or hydrogen peroxide may well be the physical nature,39 or the electron-donor ability of the next near-neighbours of the iron complex, or the ease with which these groups allow substrate approach to the bound oxygen.Undoubtedly the consideration of these same protein groups must be important in the long-range electron-transfer observed in many-headed and multi-enzymes. Very similar considerations may apply to some of the non- heme iron oxidases, e.g. phenol oxidases which in many respects resemble hemerythrin, or to the copper oxidases which can be closely likened to hemo- cyanin. In the reactions of the iron proteins, peroxidase and catalase, with oxygen or hydrogen peroxide, intermediates have often been observed which carry a large number of oxidizing equivalents per metal atom. Although recent Mossbauer data supports George’s4 O contention that one of the intermediates is a ferry1 ion, i.e.Fe02+ or iron in oxidation state IV, there is as yet little detailed understanding of the states which contain bound hydrogen peroxide or oxygen. In some cases considerable steric or chemical modification of organic groups such as the porphyrin or protein side chains seem to be necessary postulates in order to understand the concomitant changes in both absorption spectrum and circular dichroism. 41 Addition of the second sub- strate, which is to be oxidized, often causes further change in spectrum so that either the substrate is bound close to the heme (this seems to be essential for oxygen incorporation) or there is a long range allosteric effect. The nature of bound oxygen has been discussed already. Bound hydrogen per- oxide could be present in the form of inorganic or organic peroxide or as a hydroxylated intermediate.l5 R.I.C. Reviews 26 The reverse reaction to oxygen absorption-oxygen evolution-is brought about in chloroplasts. It is still uncertain if any metal is involved in the last step of the reaction but it is usually thought that Mn(1rr) or Mn(rv) are in- volved. Not all oxidases act by forming a complex between a metal and oxygen. The metal, now in the higher oxidation state, can form a complex with the second s ~ b s t r a t e , ~ ~ ~ ~ ~ e.g. The excited states of such complexes, in which the electron is transferred to the metal, are known to be low-lying from studies of charge-transfer spectra. The next step in the reaction can be that of oxygen addition to the organic radical and thence by dissociation to the observed products.Other metal cations that could function in this way are CU(II) and Mo(vI). In all cases the metal acts as an initial one-electron acceptor so activating the substrate to oxygen attack. Some flavoproteins contain both flavin and a metal ion and there is often strong interaction between these two reactive centres, the flavin interacting with the metal ion much as is pictured above for the metal-phenolate interaction. Such systems can be either single-headed or many-headed depending on the distance apart of the metal and the flavin. These enzymes, containing particularly iron and molybdenum, will be de- scribed in greater detail later (p. 31) for most frequently they form a part of complex systems.Many amine oxidases also have two cofactors, CU(II) and p y r i d ~ x a l . ~ ~ Here the metal seems to activate pyridoxal for condensation reactions and no valence change is involved44-compare zinc activation of NAD. Just as we have observed similarities between the hydride accepting, "3-dependent, enzymes of cytoplasm and of complex structures such as mitochondria-the major difference resting in the coupling of the latter to electron transfer systems of the ferredoxin type-so there is a relationship between the simple oxidases (and oxygen carriers) and the terminal oxidases which act at the other extreme of the mitochondria1 chain. These oxidases are known as cytochromes a3, a, and 0. While the hydride transferring centres are connected to ferredoxins, the terminal oxidases are bound to either copper or additional cytochrome electron-transfer proteins.(The scheme is represented in Table 7, p. 33). Electron-transfer proteinss* 41 In the case of the iron and copper proteins active in direct reaction with oxygen or hydrogen peroxide, evidence is accumulating that the metal is either open-sided or is readily substituted. In keeping with this reactivity the iron for example is always in a high-spin state [g-values 'v 2.0 and e 6-0 in heme complexes and 4.3 in others-Fe(nr)] and gives a readily recognized (non-hemochromogen) spectrum (Table 5). The copper is probably in the CU(I) state and as such is not easily followed by physical methods.The physical properties of the electron-transfer proteins are at first sight extremely different. Thus the heme iron compounds in this class (Table 5) Williams 27 Protein (a) Electron-transfer proteins Table 5: Classification of heme-~roteinsl~ Cytochrome a C ytoc h rome b Cytochrome c (b) Oxidases, etc Cytochrome a3 Cytochrome a2 Cytochrome o Cytochrome cc’ Cytochrome P450 Catalase Cytochrome c Peroxidase Peroxidases Tryptophan pyrrolase (c) Oxygen-carriers H emog lo b i ns M yog lo b i ns C h lorocr u ori ns HC HC HC HC(50%) HC? Not HC Not HC Not HC Not HC Not HC Not HC Not HC Not HC Not HC Not HC HC = hemochromogen spectrum; * P. Day, D. W. Smith and R. J. P. Williams, t o be published.have always been classified as low spin, the copper of the blue proteins is CU(II) and may be inaccessible to water, and the iron of the ferredoxins gives rise to a set of g-values about which there is much discussion but no firm conclusion.s Tt was generally thought that the coordination sphere of the metal in these proteins remains intact during the electron transfer process. Recently, however, the certainty with which this position could be held has been lessened by the knowledge that the cytochromes may not be more than about 90-95 per cent low spin (Table 5). In the ferredoxin series it is also possible that the iron can relax between different spin states. Again Moss- bauer spectroscopy, unlike EPR studies, has failed to recognize the spin state equilibrium in any of the heme systems suggesting perhaps that at s as opposed to s relaxation has averaged out the signals.41 In many cases very low temperatures have been required to observe the systems even by EPR.Similar difficulties arise with the study of CU(II) electron- transfer pro- teins and ‘non-wR detectable CU(II)’ is often described. Considering the repeated reports of anomalous physical states of these proteins Vallee and 10 Very small Very small Very small 28 Spectra Fe(ll) 1 Fe(lll) Not HC Not HC Not HC Low spin content 10 (%) Fe(lll) at 25°C N N Near 100 90 90* 0 Near 50 Near 50 Near 20 Near 20 Near 50 Near Near Variable 10-90 Near ~~ ~ ~ R . I.C. Re views Williams have been led to ask whether these anomalies have a special signi- ficance for function. 28 Fast electron-transfer, i.e. of low activation energy, to and from a metal coordination sphere demands that the two valence states involved should be made structurally equivalent with but low energy expendit~re.~~ Thus the optimal condition is that a compromise geometry between that most desired by the two states should be forced upon them in their ground state. The optimal geometries are different for different metals as the valence states of the metals have a variety of geometric demands. Thus the two valence states of copper demand very different site symmetries-Cu(I1) tetragonal, CU(I) linear or tetrahedral. Here some irregular geometry must be the best com- promise between the two states.In the copper blue proteins such an irregular geometry has been achieved as noted in Table 2, p. 15. By way of contrast, the two valence states of iron demand optimal geo- metry of the same symmetry, e.g. octahedral, but they demand different bond distances. Thus the optimal condition for very fast electron transfer to and from an iron in, for example, a porphyrin complex (an axial field in which the four in plane ligands cannot move but the out-of-plane ligands are somewhat free to move) is that the z-axis ligands should be held by ex- ternal bonding, e.g. to a protein, at a distance which would not be optimal for either valence state in the absence of the restrictions built into the external bonding by the protein structure. In the case of the higher valence state, z-axis bonds should be longer than normal, by say 0.1-0-2 A.This situation will be expected to show itself in the Fe(Irr) complex by, (1) lowered thermo- dynamic stability of the complex; (2) a stabilization of high-spin relative to low-spin electronic structures; (3) rapid relaxation between spin states, for this relaxation could be achieved by low-energy vibration of the z-axis ligand. The optimal condition for electron transfer is less critical with respect to a low valence state as, (4) the bond energy of the lower valence state is smaller and, in any event, (5) the bond distance of the low-spin, low-valence state matches almost exactly that of the high-spin, high-valence state.From (I), (4) and (5) we expect that, (6) the redox potential of the complex will be abnormally high if optimal conditions for electron transfer are achieved. The high redox potential of the heme enzymes has been commented upon earlier (Fig. 1). The abnormality of the Fe(I1I) sites in cytochromes is shown further by the EPR and Mossbauer signals themselves. In the EPR spectrum there are three widely separated g-values, e.g. for cytochrome b, = 1.8, 2.2, and 3.05 indicating a strongly asymmetric field despite the fact that the 4 7 porphyrin give would Mossbauer spectrum have of been cytochrome expected c to has a quadrupole a simple tetragonal splitting field. of - The 2.1 mm/s which can be explained readily on the basis of weak coordination perpendicular to the heme plane.In all particulars then-redox potential, spin-state equilibria, relaxation rate, EPR and Mossbauer spectra-the physical measurements suggest that the iron is relatively weakly bound perpendicular to the porphyrin plane. Of course this does not exclude weak binding to a further protein group (a seventh ligand) which could also aid electron transfer but it does appear that the coordination sphere itself may well be peculiarly suited for rapid electron William 29 The ferredoxinssp The ferredoxins are a much more difficult series of proteins to describe. Although the iron is bound to sulphur, the nature of this sulphur is uncertain, other groups in the coordination sphere are unknown, and the interactions between the different iron atoms (some ferredoxins contain eight such atoms) are not understood. It is possible that much of the difficulty in understanding is generated from comparison with naive models of high symmetry.As was pointed out in the previous section advantages accrue to these metal centres if the coordination sphere is a compromise geometry between those demanded by different valence states. This has led to the suggestion that they are generally involved in the H- + H+ The ferredoxins are found associated with numerous low-potential systems. reaction in biology. Examples of their occurrence are in the systems which carry out nitrogen fixation, hydrogen fixation, methane production, and hydride transfer from NADH, flavins and hydroquinones.They do not appear to be involved in direct reaction with oxygen. By way of contrast the heme-iron, the non-sulphur non-heme iron, the copper proteins, and perhaps even the molybdenum proteins are more often involved in oxygen reactions. In model complexes it is the low formal oxidation states, comparable to the ferredoxins, which are particularly active in hydrogen and nitrogen reactions. Thus ferredoxins could have a great significance in the most primitive biologi- cal reactions. It must always be remembered that oxygen is a relatively recent metabolic product of life and catalysts for its reactions may be relative newcomers. Before discussing the organized metal-containing systems of biology, the great variety of the proteins which is hidden by the strong classification set up in this review should be noted.Most of the proteins of each of the different kinds which have been described are species-specific and any one species may contain several proteins of one kind, e.g. hemoglobins. The small variations are due to changes in individual amino-acids in the protein chains. Occasionally these changes can be detrimental to function so that all muta- tions are not viable. Diseased conditions can also result from detrimental changes in composition of a single protein. In the biological system the species-specific protein is then wrapped in an individual matrix under the influence of which it expresses its properties. Thus classes of compounds such as ‘hemoglobins’ and ‘cytochromes c’ cover a wide range of oxygen affinity and electron transfer activity respectively.While exercising one’s mind with the classifications of the inorganic material in a living system the diversity of the material must be constantly remembered. Even against this background the rather limited range of classes of biological compounds is curious in view of the ever-evolving range of new inorganic compounds. Organized units of metal lop rote in^^^ We have now described the many individual proteins which have one or occasionally several metal-ion centres. Often these proteins have been separated from larger units and the separation has removed a part of the functional significance of the protein even if there is but little change in the R .I. C. Re views 30 individual metal-ion centres. These centres have been seen to retain one of three functions : hydrogen transfer ; oxidation, with molecular oxygen- though nitrite or even nitrogen could be used; electron transfer. It is the ordering of these three activities that will now be considered. In the first instance the so-called many-headed enzymes will be described for they can be studied in simple solution. They are of high molecular weight (several hundred thousand) so that they could be considered as small droplets of a separate phase in water. It is likely that they consist of several separable sub-units, each one of these sub-units being comparable with the simpler enzymes already described, but the total unit is thought to have but one functional action.Secondly we shall describe the multi-enzyme systems which are more clearly particulate and are readily seen to represent separate phase systems in water. These multi-enzyme systems are obtained from membranes where they are usually in a more or less lipid matrix. They carry out a complex and connected set of reactions. Many- headed enzymes26 9 Some many-headed metal enzymes are listed in Table 6. They may contain different metals or many atoms of the same metal either in the same or different oxidation states. The metal ions are often at some distance from one another, as is shown by the similarity in physical properties (spectra, EPR signal, redox potential) of the metal ions in the many-headed and the single- headed enzymes.However, in general, the many-headed enzymes carry out a different form of oxidation. Thus they usually oxidize by dehydrogenation rather than by oxygen incorporation. Ascorbic acid oxidase is a good example, the first reaction product being the ascorbate free radical due to hydrogen atom abstraction. The oxidizing agent is still molecular oxygen. Thus a logical assumption is that the many-headed enzymes absorb oxygen at one Table 6: Manv-headed metallo-enzvmes Metals Enzyme Fe(2-7) Ferredoxins Fe(2-6) Fe(4) flavin Fe(6) Mo(2) flavin Fe(2) flavin NAD H-oxid ases Flavin-oxidases (SUC- ci nic, dehyd roq u i none) Xanthine-oxidase Di hydro-orotic dehydrogenase Ascorbic acid oxidase Phenol oxidases Laccases Cu(6) Cu(2-8) Cu(2-4) - - - - - - - - - - - - - Monoamine oxidase Cu(4) Williams Plastocyanin Ceruloplasmin RH P-protei n 3 Function Mol.wt. x 10 5-10 Removes H, o r electron trans- fer, see Table 9 Removes H, see Table 9 Removes H, see Table 9 ? ? 250 62 Oxidation of xanthine Removes H I50 30-300 I20 Removal of H atoms Removal of H atoms-melanins Removal of H atoms 20 I60 - - - - - - - - - - - - - 225 Am i ne+aldehyde (pyridoxal coenzyme) Electron transfer, see Table 9 Cu transport? Fe oxidation? Bacterial respiration 30 ~ 31 site and the substrate at another. Electrons or hydrogen atoms must then pass between the sites. The mechanism of such reactions, i.e.dehydrogenation by oxygen without oxygen incorporation, demands that the reaction, 4H' 4- 4e $- O2 - H,O a four-electron change, does not go through a radical of any appreciable life, i.e. Og, '0,H or OH', for this radical could clearly attack the radical of the substrate known to be produced by dehydrogenation. Thus it appears to be a necessary feature of that active site of the many-headed enzyme which is connected to oxygen that it should be capable of rapid release of two or four electrons to oxygen. The high concentration of metals in the many- headed enzymes is then partly required as a many-electron reserve or sink. Furthermore, all the groups in this sink must react very rapidly and together. Assuming this description to be correct it is naturally difficult to understand or perhaps even observe Mossbauer and EPR signals for relaxation and saturation of all the connected groups involved, Cu2+, flavin', Mo5+, Fe3+, is likely to be abnormal.Effectively many one-electron reactions are made equivalent to a many-electron process. It is in just such a process that one would expect the unusual geometric and chemical environments ('equili- brated' valence states) of the copper blue, cytochrome and ferredoxin pro- teins to be so advantageous. The exact way in which oxidizing equivalents migrate in a protein medium is not likely to be an easy problem to solve for we have already seen that the simpler interaction between iron atoms in hemoglobin is not understood, although a relatively complete structure of the protein is available.However, the differences between the two types of cooperative interaction, one requir- ing just an adjustment of geometry, an allosteric effect as in hemoglobin, and the other demanding a conduction path for hydrogen atoms or electrons with or without an allosteric effect, suggests that it is the chemical reactivities of the next near-neighbour protein side chains which are important. This is illustrated in Fig. 6. The complexities of multi-enzymes which have four functions-reduction, oxidation, electron transfer and energy conservation- at different sites must imply an even greater function significance of protein side-chain s. There are some many-headed and multi-enzyme systems which do not oxidize by hydrogen abstraction alone. In particular some of the mixed function oxidases act in a simultaneous dual fashion, i.e.both as true oxidases incorporating oxygen and as dehydrogenases. s 1- H-donor -t 0; 6 SO:'H + H20 -+ donor(oxidized1 The donor here is effectively a third substrate so that a complicated inter- dependence of oxidase and dehydrogenation activity is brought about. Some examples are given in Table 7. The most interesting of these catalysts are those which occur in microsomes. Here the third substrate is NADPH or NADH so that the total multi-enzyme is NADPH FD ? cyt b P450 R. I.C. Reviews 32 I / I A 8 Fig. 6. An idealized diagram of the relationship between the transition metal containing biological catalysts.A is an oxygen carrier; B is an oxygenase which inserts 0, into a sub- strate; C is a dehydrogenase catalyst using 0, as acceptor-a mixed function oxidase would which have the an right energy hand trap unit i s inserted of C replaced i n t o C. by P - 6. P D represents illustrates the the formation energy conservation of ATP and chain in each i n diagram M is a metal. Unknown organic groups are shown as shaded areas. D The P450 cytochrome is typical of a heme oxidase in spin state, in the sensi- tivity of its spectrum to substrates and in its reaction with carbon monoxide. It probably reacts with oxygen and substrate simultaneously. Perhaps the chain of catalysts is only able to handle two electrons at a time so that two- electron removal from substrate is obligatory for the 0, + H,O reaction.The further discussion of such chains properly belongs to the next section. Table 7: Some mixed function oxidases Enzyme A r y 1-4- h y d roxy las e Steroid-2-hyd roxylase Steroid- I 7-O( hydroxylase Fatty acid hydroxylase Kyn urenite hydroxylase Metals Fe(b5) Fe( P450) Fe(?) Fe( P450) Fe(?) Fe(?) Substrate Aniline Estriol Progesterone Fatty acids Kynurenate NADPH NADPH NADPH NADPH NADH Organized or m ul t i-enzyme syst ems4 For more than 30 years evidence has been accumulating to the effect that many metal proteins aggregate in highly organized systems. Some of these are given in Table 8 and their inter-relationships are shown in Table 9.Willianis I I I C P - P 33 Table 8: Multi-enzyme systems System Cytochrome chain (Mitochondria) Bacterial oxidases (no m i toc h on d ria) Plant chloroplasts Bacterial photosyn- thesis (no c h I orop I asts) Microsomal chain N2 and H2 fixation (bacteria) Letters refer to: (a) (c) (f) (0) corresponding cytochromes; (fd) 'erredoxins; (hv) light; RH P, r hodos pi re1 I um )rum cytochrome system; P450 a special heme-iron protein of microsomes. Some of the proteins contain many metal ions, e.g. fd has from 2-7 iron atoms. The chains are not complete as non-metallic systems, e.g. quinones, are omitted. Metal ion chain in probable order 0 2 t Fe(a3) Cu(a3) Fe(a) Cu(a) Fe(c) Fe(c1) Fe(fd) Fe(b) Fe(fd) t H- 0 2 t Fe(o) Fe(b) Fe(fd) t H- 0 2 t Fe(a2) Fe(c) Fe(fd) t H- 0 2 t Mn(?) Cu(?) [(hv) Mg] Fe(b6) Fe(f) [(hv) Mg] Fe(fd) t H- Fe(RHP) Fe(c2) [(hv) (Mg)] Fe(fd) t H- Cu(?) 0 2 c Fe(P450) Feb5 Fe(fd) t H- substrate (which i s hydroxylated) (N2) t [Fe(fd) Mo(?) Fe(?heme)] t (H2) Table 9: Some proposed energy-producing systems System Cytochrome chain Chloroplast Nitrate reduction Nitrogen fixation Bacterial photosynthesis S,? Fe? Iron-dependent bacteria Energy Oxidizing system (ADP- P) f 0 2 4 Fe/Cu 0 2 f Mn? (ADP- P) f hv (Mg) J (ADP- P) f NOa-J Mo? ADP- P J.N2 J. (Mo o r Fe) Unstable mineral sources in oxidized state (e.g. iron oxide deposits) .f In ( 4 ) implies that the system uses the substrate while out ( f implies that it generates the substrate.Energy in (4) implies that the system absorbs energy in this particular form while energy out ( f ) implies that energy is generated. Many of the systems can be reversed, ADP-P means energy trapped as ATP. I Unstable mineral sources in reduced states, e.g. iron sulphide deposits J- 34 .1 Function 0 2 + H20 H-+ Hf co2 -+ 'CHOH / ADP + ATP H20 -+ 0 2 I or ADP H2S + -+ ATP S, Hyd roxylation Reduction of N2 t o NH3 (little under- stood) Reducing system Fe CHOHJ I I I Fe CHOHf H204 I I Fe CHOHJ I I Fe CHOH 4 or H2 J I I Fe CH0H.f H2SJ. R. I.C. Reviews Catalyst Abnormality Cytochrome a3 Cytochrome Q Copper (11) Spin state mixture, unusual spectrum Split Soret absorption band Unusual spectrum and EPR signal-in some cases EPR signal not Cytochrome c Cytochrome 6 Ferredoxins Flavins NADH Cytochrome P450 Chlorophyl I detected Spin state mixture, EPR signal very broad Spin state mixture, EPR signal not detected EPR signal has very fast relaxation and is unique Unusual spectrum Unusual spectrum Highly anomalous spectral properties Unusual spectrum probably of the molecule which acts as a trap Peculiar features of these systems have been observed in some and may be common to all.They are: (1) One end of the chain is at a very different redox potential from the other and reacts with entirely different substrates.Thus one end is connected to an oxidizing agent, O,, NO;, N,(?) while the other is connected to a reducing agent, succinate NADH, reduced flavin, H,. Chloroplast and bacterial chromatoids may contain more than one such chain (Table 7) for the absorp- tion of light generates one reducing and one oxidizing equivalent, i.e. charge separation. (2) The multi-enzymes transfer electrons not by conduction along proteins but by a so-called hopping mechanism, in which an electron goes from one reactive site to a n ~ t h e r . ~ ~ , ~ ’ This hopping is under feedback metabolic control. Thus electron transfer in the cytochrome chain of mitochondria is regulated by the concentration of phosphate, adenosine diphosphate (ADP) and regulatory hormones. (3) The series of redox metal-enzymes is connected, via the hopping mechanism, to careful energy conservation in the form of phosphorylation, ADP + P -+ ATP (Table 8).This process may be reversed, energy utilization, i.e. pyrophosphate hydrolysis, yielding reducing equivalents, light, or energy for reactions. Nitrogen fixation may be brought about using the breakdown Of ATP. (4) Many of the catalysts in the chains have unusual physical properties indicating that they are in a reactive condition as discussed earlier (see Table 10). ( 5 ) The whole multi-enzyme system is in a separate phase from the cyto- plasm of the cell. This phase is often called a membrane by biologists for it separates two regions of ‘cytoplasm’ aqueous media. The function of the membrane is not like that of the familiar membrane of an osmometer, however, for it contains the enzymes for certain reactions as well as acting as a diffusion barrier (compare the many-headed enzymes).In fact some ‘membranes’ may be almost entirely composed of a multi-enzyme system in a lipid medium. (6) The metal proteins are arranged in a sequence of redox potential suggesting that energy can best be handled in such a sequence. Williams 35 At the present time the understanding of these ‘inorganic’ highly-organized, energy-generating catalysts, the main energy sources of life, is so slight that no attempt can be made to copy them. Nevertheless it is already clear that all living things have such ‘inorganic’ units, and despite the diversity of proteins in them, they have common inorganic active sites.Thus all have cytochromes, copper proteins, and ferredoxins incorporated into their multi-enzymes. The very persistence of these chains of catalysts suggests that in their properties lies the secret of organized activity-the ability to retain energy and use it constructively, which must have come before the highly sophisticated DNA code. Have these systems developed slowly from non-living single-headed catalysts, through many-headed catalysts to multi- enzymes trapped in a non-aqueous phase?49 Could it be that these energy traps have developed into a membrane form and thus allowed life to appear as a property of cells?49 The structure of these non-aqueous phases suggests that it is the very incorporation of inorganic catalysts into a specific spatial relationship which generates their special properties, e.g.in chloroplasts, mitochondria, microsomes, and other particles, and which gives them the ability to generate the necessary energy for life. Consideration of such organized systems is leading to new suggestions about sequential chemical reactions, which have rather little in common with present-day laboratory inorganic reactions. It is proposed that through rapid electron-transfer, protons of high chemical potential are produced in special regions of an reaction by removing water from phosphate + ADP. The total system is so aqueous or a non-aqueous phase. The proton brings about a condensation constructed that the condensation reaction (which removes H+) is a necessary step for electron transfer between different, many-headed, regions of the particulate Insight into systems as difficult as these requires the combined skill and thought-patterns of specialists from most scientific disciplines, but it is a direct challenge to inorganic chemists to attempt to make models of such organizations.REFERENCES 1 R. J. P. Williams, Nature, 1959, 184, 44; Fifth International Congress of Biochemistry (Moscow) Vol. 4, p. 133. Pergamon Press, Oxford, 1963. 2 B. L. Vallee and J. E. Coleman, in Comprehensive Biochemistry (ed. M. Florkin and E. H. Stotz), p. 165. Elsevier, Amsterdam, 1964. B. G. Malmstrom and J. Neilands, A . Rev. Biochem., 1964, 33, 331 ; Several Chapters of The Enzymes, Vols.1-8 (ed. P. D. Boyer, H. Lardy, and K. Myrback), Academic Press, New York, 1959-64, deal with metal-enzymes. 3 R. J. P. Williams, Endeavour, 1967, 26, 96. 4 W. E. C. Wacker and B. L. Vallee, J. bid. Chem., 1959, 234, 3257. 5 R. G. Shulman, H. Stemlicht, and B. J. Wyluda, J . chem. Phys., 1965, 43, 3116. 6 W. E. C. Wacker and R. J. P. Williams, presented at the International Biochemistry Congress, Tokyo, 1967. 7 J. C . Kendrew, Science, 1963, 139, 1259. 8 A, San Pietro (ed.), Non-heme Iron Proteins, Antioch Press, Yellow Springs, Ohio, 1965. 9 B. B. Buchanan, in Structure and Bonding, Vol. I, Springer-Verlag, New York, 1966. 10 A. S. Brill, B. R. Martin, and R. J. P. Williams, in Electronic Aspects of Biochemistry (Ed.B. Pullman), p. 519. Academic Press, New York, 1964; W. E. Blumberg, in The Biochemistry of Copper (Ed. J. Peisach, P. Aisen, and W. E. Blumberg), p. 49. Academic Press, New York, 1966. R.I.C. Reviews 36 11 A. E. Dennard and R. J. P. Williams, in Transition-Metal Chemistry (Ed. R. L. Carlin), Vol. 2, p. 115. Dekker, New York, 1966. 12 A. S. Brill and J. H. Venable, in The Biochemistry of Copper, (Ed. J. Peisach, P. Aisen and W. E. Blumberg), p. 67. Academic Press, New York, 1966. 13 R. A. Firth, H. A. 0. Hill, J. M. Pratt, R. G. Thorp and R. J. P. Williams, Chem. 14 R. C. Bray, G. Palmer and H. Beinert, J. biol. Chem., 1964, 237, 2657, 2667. Comm., 1967, 127; F. Wagner, A. Rev. Biochem., 1966, 35, 405. 15 A. Ehrenberg, in Heme and Hemoproteins (Ed.B. Chance, R. Estabrook and T. Yone- tani), pp. 331,462. Academic Press, New York, 1966; A. S. Brill and R. J. P. Williams, Biochem. J., 1961, 78, 246, and references therein. 16 D. E. Drum, J. H. Harrison, T. K. Li, J. L. Bethune, and B. L. Vallee, Proc. natn. Acad. 17 See discussion of Ceruloplasmin in, The Biochemistry of Copper (ed. J. Peisach, P. Aisen and W. E. Blumberg), p. 513. Academic Press, New York, 1966. 18 (a) A. L. Schade in Protides of the Biological Fluids (ed. H. Peeters), Vol. 14, p. 13. Sci. U.S.A., 1967, 57, 1434. Elsevier, Amsterdam, 1967; (b) R. E. .Feeney and St. K. Komatzu in Structure and Bonding, Vol. I, p. 149. Springer-Verlag, New York, 1966. 19 J. Ozols and P. Strittmatter in Symposium on Cytochromes, (ed.K. Okunuki and M. Kamen), p. 304. Tokyo University Press, Tokyo, 1967. 20 See, for example, Haemocyanin (ed. F. Ghiretti). Academic Press, New York, 1967, and, in particular, articles by K. E. Van Holde, and G. Morpurgo and R. J. P. Williams. 21 L. Vaska, Science, 1963, 140, 800. J. A. Osborne, F. H. Jardine, J. F. Young and G . Wilkinson, Chem. Comm., 1965, 131. 22 A. Misomo, Y. Uchida and T. Saito, J. Jap. Chem. Soc., 1966, 46, 700. 23 P. Saltman, J . chem. Educ., 1965, 42, 68L. 24 H. Porter, in The Biochemistry of Copper (ed. J. Peisach, P. Aisen and W. E. Blumberg), p. 159. Academic Press, New York, 1966. 25 M. Margosches and B. L. Vallee, J. Am. chem. Soc., 1957, 79, 4813. 26 0. Hayaishi, in Oxygenuses (ed. 0. Hayaishi), p. 1 . Academic Press, New York, 1962; T. E. King, H. S. Mason and M. Morrison (eds.) Oxidases and Related Redox Systems, WiIey, New York, 1965. 27 J. C. Tomkinson and R. J. P. Williams, J. chem. Soc., 1958, 2010; B. R. James and R. J. P. Williams, ibid., 1961, 2007. 28 B. L. Vallee and R. J. P. Williams, Proc. natn. Acad. Sci. U.S.A., 1968, February. 29 R. J. P. Williams, in Haemitin Enzymes (ed. J. E. Falk, R. Lemberg and R. K. Morton), p. 41. Pergamon Press, Oxford, 1961; B. R. James, J. R. Lyons and R. J. P. Williams, Biochemistry, 1962,1, 379; R. J. P. Williams, in The Enzymes (ed. P. D. Boyer, H. Lardy and K. Myrback), Vol. 1, p. 391. Academic Press, New York, 1959, 30 W. R. Groskopf, J. W. Holleman, E. Margoliash and I. M. Klotz, Biochemistry, 1966, 5, 3779. 31 H.-J. Bieleg, E. Bayer, H.-D. Dell, G. Rohns, H. Mollinger and W. Riidiger, in Protides of the Biological Fluids (ed. H. Peeters), Vol. 14, p. 197. Elsevier, Amsterdam, 1966. 32 A. S. Mildvan, M. C. Scrutton and M. F. Utter, J. biol. Chem., 1966, 241, 3480, 3481. 33 R. C. Bray and L. S. Merriweather, Nature, 1966, 212, 465. 34 N. Sutin and D. R. Christman, J . Am. chem. SOC., 1961, 83, 1773. 35 H. A. 0. Hill, J. M. Pratt and R. J. P. Williams, Proc. R. SOC., 1965, 288A, 352. 36 0. Hayaishi (ed.), Oxygenases, Academic Press, New York, 1962. 37 Y. Ishimura, M. Nozaki, 0. Hayaishi, M. Tamura and I. Yomazaki, in Symposium on Cytochromes (ed. K. Okunuki and M. Kamen), p. 373. University of Tokyo Press, Tokyo, 1967. 38 R. Lemberg, in Symposium on Cytochromes (ed. K. Okunuki and M. Kamen), p. 1 . University of Tokyo Press, Tokyo, 1967; K. Okuniki, in Oxygenases (ed. 0. Hayaishi), p. 409. Academic Press, New York, 1962. 39 J. H. Wang, in Oxygenases (ed. 0. Hayaishi), p. 470. Academic Press, New York, 1962. 40 P. George, Biochem. J., 1953, 54, 267. 41 For a general review see, Hemes and Hemoproteins (ed. B. Chance, R. Estabrook and T. Yonetani). Academic Press, New York, 1966. 42 Y. Kojima, H. Fujisawa, A. Nakazawa, T. Nakazawa, F. Kanetsuna, H. Taninchi, M. Nozaki, and 0. Hayaishi, J. biol. Chem., 1967, 242, 3270. 43 H. S. Mason, A. Rev. Biochem., 1965, 34, 1965. 44 D. E. Metzler, M. Ikawa and E. E. Snell, J. Am. chem. Soc., 1954, 76, 648. 45 F. Basolo and R. G. Pearson, Mechanisms of Inorganic Reactions, p. 306. Wiley, New York, 1958. Williams 37 46 R. J. P. Williams, in Protides of the Biological Fluids (ed. H. Peeters), Vol. 14, p. 13. Elsevier, Amsterdam, 1967; and in Biochemistry of Copper (ed. J. Peisach, P. Aisen and W. E. Blumberg), p. 131. Academic Press, New York, 1966. 47 R. J. P. Williams, in Non-heme Iron Proteins (ed. A. San Pietro), p. 7. Antioch Press, Yellow Springs, Ohio, 1965. 48 Series of papers by B. Chance, W. Slater, D. Green and €1. Beinert should be consulted in order to appreciate the methods and skills which have gone into the attack on this problem. See, for example, references 23, 43. 49 A. I. Oparin, Adv. Enzymol., 1965, 27, 347. 50 P. Mitchell, Nature, 1961, 191, 144. R. J. P. Williams, J . Theoret. Biol., 1961, 1, 1. 38 R. I. C. Reviews

ISSN:0035-8940    

DOI:10.1039/RR9680100013

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

INFRARED AND RAMAN SPECTRA INSTRUMENTATION OF INORGANIC COMPOUNDS H. E. Hallam, M.Sc., Ph.D., F.R.I.C. Department of Chemistry, University College of Swansea, Singleton Park, Swansea. Infrared and Raman spectroscopy are the two techniques commonly used for the study of molecular vibrations. During the last two decades i.r. spectro- scopy has achieved paramount importance as a structural diagnostic tool, particularly in the field of organic chemistry. For two reasons, its application to inorganic compounds has developed more slowly. First, all organic com- pounds give i.r. spectra but some inorganic materials (monatomic anions and cations) do not ; secondly, many metal-ligand (M-L) frequencies occur below 400 cm-l, a region which has only recently become examinable by commercial spectrometers.As a result of instrumental developments there has been a surge of interest in the applications of i.r. spectroscopy to prob- lems in inorganic chemistry and the technique now plays an important r61e in undergraduate inorganic practical courses alongside that already estab- lished in organic chemistry. Ideally an i.r. study should be carried out in conjunction with a study of the Raman spectrum since this gives complementary information. The Raman effect in fact provided much of the early experimental data and is of particular value to inorganic chemistry because of its facility for studying aqueous solutions. The very recent introduction of commercial laser Raman spectrometers has also caused a resurgence of interest in the Raman spectra of inorganic compounds,l although the expense of this instrumentation will preclude it from becoming routine in teaching laboratories for some time.This review is limited to a discussion of the elementary principles and features of the i.r. and Raman spectra of inorganic molecules and ions as a guide to qualitative and structural analysis. No attempt is made to discuss the detailed theory nor to describe fully experimental procedures and instru- mentation. The general text edited by Davies2 adequately covers these aspects of both i.r. and Raman spectroscopy and also includes a chapter by Ebsworth on applications to inorganic chemistry. Sampling techniques and i.r. instrumentation are fully described in an excellent new manual3 edited by Miller.An outline of instrumentation and practical techniques is given here to provide a background to the discussion of vibrational spectra. A spectrometer is an. instrument which measures the amount of radiation transmitted by a substance as a function of the wavelength (A), frequency (Y = c/A), or wavenumber (ij = v / c = 1/A) of the radiation. Hallam 39 Infrared A schematic diagram of an i.r. spectrometer is shown jn Fig. 1. Radiation, from an electrically-heated filament which emits frequencies over a wide spectral range, is reflected through the sample and reference cells by mirrors. The reference beam is attenuated by a comb which can be moved into and out of the optical path until it is equal in intensity to the sample beam.The two beams are alternately sent to the thermocouple detector by a rotating chopper mirror. This allows the transmitted intensity at each frequency to be compared with the intensity in the absence of sample. Before reaching the detector the alternating beam is dispersed into its component frequencies by means of an alkali halide prism or a diffraction grating. The spectrum is scanned by rotating the prism or grating so as to alter the frequency being received by the detector. When the sample is not absorbing, the detector receives a constant signal; if the sample absorbs at a particular frequency the sample transmission is lowered. The thermocouple thus receives an alternating signal which is amplified and fed to a servo motor which drives the attenuating comb into the reference beam until the two beams are balanced.The comb is linked to a pen which records the attenuation, i.e. sample transmission (T = transmitted intensity/incident intensity I/IoO) on a chart moving in synchronism with the rotation of the prism/grating. The resultant plot of Ill0 versus v is the i.r. spectrum of the sample. There are several routine i.r. spectrometers available commercially which cover the range 4000 to 400 cm-1 and one which extends to 250 cni-l. More sophisticated (and more expensive) instrumentation is required to reach the very far i.r., i.e. down to tens of cm-1. Raman When an intense beam of light of frequency v is incident upon a transparent medium a small amount of the radiation energy is scattered, even if the medium is rigorously freed from dust particles. The scattered energy consists almost entirely of radiation of the incident frequency (Rayleigh scattering) but also contains frequencies ( V - Av) and ( V + Av) where the Raman shift A v corresponds to a molecular vibrational frequency.If the frequency of the incident beam is within the visible range the scattered radiation can be examined by means of a visible spectrometer. To excite Raman spectra it is necessary to use an extremely intense source of mono- chromatic radiation; the mercury arcs developed for this purpose are now being replaced by 1asers.l Figure 2 shows schematically the essential com- ponents of a recording Raman spectrometer.SAMPLING TECHNIQUES The vibrational spectrum of a molecule arises from its internal vibrations and we can best interpret it when the molecule is in its simplest and most repro- ducible surroundings. Ideally this is in the gas phase where, at low pressures, we can consider the molecule to be isolated and thus free from the external influence of neighbouring molecules. Unfortunately few molecules have 40 R. I. C. Review2 L 0 -.- f I V u' Hallam q- I Flat contact Sample tube I A$-- Hemispherical lens Fig. 2. Schematic diagram of a Raman spectrometer. sufficient vapour pressure at room temperature to provide enough absorbing molecules for a gas-phase spectrum to be obtained, and thermal decomposi- tion often occurs if higher temperatures are employed to increase the vapour- phase population.The closest possible approach to the ideal state usually is to study the substance in dilute solution in an inert solvent, As far as possible highly polar solvents should be avoided because of the resulting strong solute-solvent interactions. However, since the majority of inorganic compounds are highly polar ionic molecules they are generally only sufficiently soluble in polar solvents-in particular, water. Aqueous solution spectra, however, are still preferable to spectra of the solid state, in which environment the molecule/ion is greatly perturbed by strong inter- actions with neighbouring molecules/ions in the crystal lattice. The difficulty with the solution technique is that the solvent has its own vibrational spectrum and, in spectral regions where the solvent is strongly absorbing, the solute cannot be examined.Water, being a simple triatomic molecule, has a limited number of vibrational frequencies and might be expected to be a suitable i.r. solvent. Unfortunately the very large changes in dipole moment accompanying the vibrations of the water molecule lead to extremely intense absorption in the i.r. The situation is worsened by features arising from the presence of higher molecular aggregates, (H20)n, due to the strong hydrogen-bonding properties of water. It i s thus impossible to obtain good i.r. spectra of aqueous solutions over a wide frequency range although the use of D2O does extend the available range.The polarizability of the water molecule, however, changes little during its vibrations. This leads to a very weak Raman scattering and makes water an excellent solvent for Raman spectroscopy. 1.r. cells A simple gas cell consists of a 10 cm length of wide bore glass tubing with i.r.- transparent windows fitted to the ends. Window materials commonly used are rock-salt (transparent to 600 cm-l), KBr (to 350 cm-l), CsI (to Dispersing prism * o r grating 42 jd Recorder F Amplifier Laser source 7 1 A ~ / / ’ Y P h oto m u I t ~ p I I er detector +- R.I.C. Reviews 200 cm-1)-these are relatively expensive and water-soluble. CaFz (to 1000 cm-I), AgCl (to 450 cm-1) and Irtran-4 (500 cm-l) are water-insoluble but also expensive.Polythene (Rigidex) is a cheap and useful window material below 400 cm-l. Solution cells consist of two windows separated by a spacing washer of suitable thickness-usually 0.1 to 10 mm. Pure liquids can usually be ex- amined simply as a capillary film between two windows. Raman cells These are much simpler than i.r. cells and usually consist of glass tubes with a flat glass end-window. Greater care however must be taken in the Raman technique to filter solutions to remove all dust particles, since these act as scattering centres. With both techniques the examination of powdered solids presents some difficulty due to reflection from the numerous crystal faces. However, using ‘front illumination’ it appears1 that excellent Raman spectra can now be obtained directly from a powder in a sample tube.Two techniques have been developed which permit satisfactory i.r. spectra of solids to be obtained : the oil-mull technique and the pressed-disc technique. Both involve (a) finely grinding the sample so that the average particle size is less than the wave- length of radiation being used (i.e. less than ca 3 p) so that reflections and refractions at the particle faces are minimized, and (b) dispersing the particles in a medium (liquid or solid) of approximately the same refractive index so that scattering by reflection and refraction is further reduced and Rayleigh scattering is also reduced. Oil mull technique An oil mull is made by grinding the solid sample with a mulling agent such as liquid paraffin (Nujol) and placing a few drops of the resulting paste squeezed between two plates as for a liquid film.The method has the dis- advantage that the ‘solute’ spectrum is overlaid by the absorptions due to the mulling oil. In the case of Nujol (a mixture of c20-c30 alkanes) these are at 2900, 1460, 1375, and 720 (w) cm-1 and result from stretching and bending modes of the C-H groups. If the solid being mulled has absorptions at or near these frequencies a second mull has to be prepared using a per- fluorocarbon oil (Fluorolube ; Kelex). Pressed-disc technique This consists of grinding a few milligrams of the solid with about 0.5 g of a dry powdered alkali halide, placing the mixture in a special die and com- pressing it with a hydraulic press to a small disc ca 1 mm thick under a pressure of several tons.The method is commonly known as the KBr disc technique since it was developed for organic compounds studied by spectro- meters going down to 650 or 400 cm-l. For inorganic compounds where it is usually necessary to go to lower frequencies it is necessary to use CsI (Special grade CsI for this purpose is now available from BDH). The technique has the advantage that the halide (if pure) does not have specific absorption bands. Its main disadvantage is moisture pick-up during preparation. It is Hnllam 43 probably the method to be recommended for inorganic solids although thought must always be given to the possibilities of halogen exchange and the effect of the alkali-halide matrix environment on the ‘solute’ frequencies.VIBRATIONAL SPECTROSCOPY AND MOLECULAR STRUCTURE The combination of i.r. and Ranian spectroscopy can be used to determine unequivocally the structure of small molecules of high symmetry; the term molecule in this sense includes a radical or an ion. The spatial arrangement of atoms determines the molecular symmetry which in turn determines whether or not a particular vibrational frequency will appear in the i.r. and/or Raman spectrum. It is not the purpose here to discuss symmetry theory (refs 4 and 5 provide good accounts); suffice it to say that the combina- tion of all symmetry elements (rotations and reflections) determines the point group to which the molecule belongs.This in turn determines the number of vibrational modes to be expected, and their i.r. and Raman activity, for a particular molecular structure. Thus, observation of the number of bands appearing in the i.r. and/or Raman spectrum of a molecule may decide the structure or may eliminate certain possible structures. For complex polyatomic molecules it is generally impossible to apply group theory and symmetry principles. Nevertheless, vibrational frequencies can yield valuable structural information on a semi-empirical basis using either i.r. or Raman spectroscopy since in general all the frequencies will be i.r.- and R(aman)-active. A molecule consists of n atoms of varying masses linked together by (n - 1) chemical bonds of varying strengths.The atoms are not held rigid but can move together in straight-line translation and rotate and vibrate periodically about a mean position. There are two types of fundamental vibrations : stretching, in which the distance between two atoms increases or decreases but the atoms remain on the same bond axis, and bending or deformation, in which the position of the atom changes with respect to the original bond axis. To describe these complex motions three coordinates must be specified for each atom, i.e. 3n coordinates or degrees of freedom. Of these, three describe the translation as a rigid molecular unit whilst another three describe the rotation of a non-linear molecule (two for a linear molecule). Thus each molecule has 3n - 6 (3n - 5 for a linear molecule) internal degrees of freedom which, when executed, cause a distortion of the molecule. These distortions are the normal modes of vibration; such a mode is one in which the centre of gravity of the molecule does not move and in which all of the atoms move with the same frequency and in phase.Each mode is independent of the others and a molecule may execute each of these normal modes of vibration simultaneously. They may, however, not always be different for, as a consequence of geometrical symmetry, two or more vibrational frequencies may coincide-these are said to be degenerate. Of all vibrations, only those which involve a change of dipole moment are i.r.- active and are capable of absorbing all or part of any i.r.radiation of the corresponding frequency incident upon the molecule. Vibrations which involve a change of polarizability are R-active. R. I. C. Reviews 44 In simple molecules the activity of a vibration may be determined by inspection of the normal mode. In a polyatomic molecule having a centre of symmetry (a), the vibrations symmetric with respect to the cs ( g vibra- tions) are R-active/i.r.-inactive, but the vibrations antisymmetric with respect to the cs (u vibrations) are i.r.-active/R-inactive. This is called the mutual exclusion rule. It should be noted however that this rule may not hold in polyatomic molecules having several symmetry elements besides a cs. Inorganic compounds will generally have their i.r. spectra measured in the solid state (as an oil mull or in an alkali halide pressed disc).It is therefore important to bear in mind site symmetry-the local symmetry of the crystal- line environment around the cs of the molecule in the unit cell. The site symmetry is usually lower than the molecular symmetry in the free state, so that the selection rule for the gaseous state is relaxed in the crystalline state. Hence bands forbidden in the gaseous state may appear weakly, and the degenerate vibrations may split in the crystalline state. For example, calcite and aragonite, two different crystalline forms of calcium carbonate, exhibit different spectra although their chemical compositions are the same. In addition to the internal molecular vibrations, solid state spectra will be further complicated by the presence of lattice vibrations; these are due to the translation and torsional motions of the molecule as a whole.Vibra- tional frequencies may also be split due to coupling between neighbouring molecules. Miller and Wilkins6 made the first systematic survey of the i.r. spectra of inorganic substances. Their reference collection covers 159 compounds and lists characteristic frequencies of 33 polyatomic ions in the range 5000- 650 cm-I. It has recently been extended7 down to 300 cm-1. Most polyatomic ions exhibit' characteristic frequencies which are thus useful for qualitative analysis; a selection of these is shown in Figs. 3-5. A bibliography of inorganic i.r. spectra has been compiled by Lawsong but unfortunately it is badly classified and incomplete.An excellent text by Nakamoto5 covers simple ions in detail and also several classes of coordina- tion compounds. A comprehensive text by Adamsg has just appeared which gives an authoritative account of metal-ligand frequencies. A good introduc- tion to Raman spectroscopy in inorganic chemistry is given by Woodwardlo and a recent bibliographic review has been compiled by R. N. and M. K. Jones.ll From this brief introduction it is apparent that knowledge of all the vibrational frequencies, together with their i.r. and Raman activity, makes it possible to decide between alternative proposed structures having different symmetries. The most important of these structures are outlined in the following sections.DIATOMIC MOLECULES Diatomic molecules have only one vibrational mode-along the chemical bond; its frequency, v, in cm-1 is given by Haflarn 45 where f is the force constant of the bond, p the reduced mass, and c the velocity of light. Mode - * p = MzMl//(Mz + M d X2 Molecules (Point group Dooh) In a homonuclear molecule the dipole moment of the molecule is unchanged during the vibration, therefore the vibration is not i.r.-active. The polariza- bility change during the vibration, however, causes it to be R-active. X Y Molecules (C, v> In a heteronuclear diatomic molecule both the dipole moment and the polarizability change during the vibration and the vibration is thus both i.r.- and R-active. Diatomic vibrational frequencies are well-known and can be found in reference 5; Table 1 lists some of these.From the standpoint of inorganic chemistry the most important are the diatomic ions, e.g. the [O-HI- ion is characterized by a strong sharp band at 3700-3500cm-l whereas the [C=N]- gives a medium sharp absorption at 2250-2050cm-l. The nitro- sonium ion (NO)+ in nitric acid absorbs at 2220cm-1. LINEAR TRIATOMIC MOLECULES The (3n - 5) = 4 normal modes of vibration are shown below: .-c-b v Description VI Symmetric stretching VQ - (The + and - signs indicate displacement out of the plane of the molecule.) X3 Molecules (D mh) The symmetrical stretching mode VI clearly involves no dipole moment change and is thus i.r.-inactive but is R-active. vz and v3 are i.r.-active/R- inactive (centro-symmetric molecule : mutual exclusion rule). Azide ion [N=N=N]- is a good example: 1360 cm-l ( v I ) , 650 cm-1 ( VZ), 2040 crn-l(v3).XY2 Molecules (Dooh) The three modes again have the same activity as the corresponding modes of X3; common examples are carbon dioxide and carbon disulphide (Table 2). 46 Symbol S I Asymmetric stretching 1 vas R.I.C. Reviews V Table I : Vibrational frequencies (cm-1) of diatomic n Ilecules and ions Molecule Ion Li+[ 0 HI- Na+ [ 0 HI- N a+ [0 D]- K+[C N] - Na+[CN]- [NO]+ 4161 2994 3962 2886 - 0 2883 * 9 2230 233 I 892 317 665 V 3678 3637 268 I 2080 2080 2220 Table 2: Vibrational frequencies (cm-1) of linear XS and XY2 molecules and ions Molecu I e or ion (1343)" 658 497 I344 667 397 213 645 (667) 2349 I533 555 204 I 2360 coz cs2 XeF2 K+[N3]- [NOzl+ 1400 * Average of a Ferrni resonance doublet.The presence and structure of the nitronium ion [NOz]+ in mixtures of nitric acid and sulphuric acid has been demonstrated from their Raman spectra. X YZ Molecules (Cmv) The three modes are both i.r.- and R-active; hydrogen cyanide is a typical such molecule; this and several others are given in Table 3. The symbols in brackets give an approximate description of the vibrational frequency, v2 7 I 2 (6C-H) 569 (6C-D) 524 589 637 628) Table 3: Vibrational frequencies (cm-1) of linear XYZ molecules and ions Molecule or ion Vl HCN DCN cos 33 I I (vC-H) 2630 ( vC-D) 859 I286 N2O K+[NCO]- 2165 2052 K+ [ N CS]- 487 470) I v3 2097 ( VC 3 N) I925 ( VC = N) 2064 2224 I207 739 HaIlam 4 47 e.g.v (C-H) represents a stretching vibration which is largely localized in the C-H bond and v(C==N) one which is essentially concerned with the stretching of the triple bond. Table 4: Vibrational frequencies (cm-1) of bent triatomic molecules and ions Molecule or ion V l v2 I I10 I043 3756 2788 (2627) 969 1618 1261 3657 267 I 2615 688 1318 I328 3213 I151 3263 I362 739 705 I595 I178 I183 320 750 828 I540 518 I242 332 3626 I800 605 Note: (a) that stretching (v) frequencies are higher than bending (6) frequencies; (b) that the anti- symmetric stretching (vss) frequency is usually higher than the symmetric (vs) one (exceptions above are 0 3 and [NOzI-). BENT TRIATOMIC MOLECULES The (3n - 6) = 3 normal modes of vibration are shown below: The vibrations are both i.r.- and R-active whether the molecule is symmetri- cal, X3 or XY2 (C2,) or asymmetrical, X X Y or XYZ (Cs).Table 4 lists some typical examples. 48 R.I.C. Reviews PLANAR TETRA-ATOMIC MOLECULES (3n - 6) = 6 normal modes of vibration are to be expected: E) XY3 Molecules (D3h) Of the four vibrations above ~1 (the ‘breathing’ frequency) clearly involves no change in dipole moment and is thus i.r.-inactive but is R-active; vz on the other hand will be i.r.-active but R-inactive; v3 and v4 are both i.r.- and R-active.Table 5 gives some examples. Table 5: Vibrational frequencies (cm-1) of planar XY3 molecules and ions Molecule or ion ~~ BC13 480 652 47 I I069 so3 244 532 995 I330 La3+ [B03]3- ;A:] N a+ [ N 0 3 1 - Ca2+[C03]2- 83 I 879 939 I068 I087 692 706 1275 I405 I460 (cal cite) X Y2Z (C2 v ) and X YZ W ( Cs) molecules Successive replacement of the Y atoms by different atoms lowers the sym- metry to CzV and then to Cs with a consequent change in selection rules. For both cases degeneracy is removed and all six vibrations are i.r.- and R-active (see Table 6). Hallam . 49 v2 v1 Molecule I- COBrCl I828 PYRAMIDAL TETRA-ATOMIC MOLECULES The six normal modes of vibration are shown below: 3336 3338) * 2327 PH3 PC13 XY3 Molecules (GV) All four vibrations are both i.r.- and R-active; common examples a reshown in Table 7.NH3 Table 7: Vibrational frequencies (cm-I) of pyramidal X Y 3 molecules and ions u1 Molecule or ion v3 v2 v4 3414 242 I 494 975 826 96 I 737 K+[C103]- [IO3]- [S0312- [Se03]2- * Bands split due t o inversion doubling. 507 930 779 1010 807 Chlorine trifluoride is an interesting example of an XY3 molecule in that it is found to have six i.r.- and R-active fundamentals. This means 50 v5 v4 v3 --- 932 * 968) 1628 1121 189 486 330 496 374 260 620 390 633 43 2 - V6 R. I. C. Reviews that the molecule is neither symmetrical planar nor pyramidal.Diffraction studies have shown that the molecule is T-shaped with one Cl-F bond longer than the other two; from symmetry considerations it should therefore be formulated as an XYZZ molecule belonging to the CzV point group (see Table 8). Molecule v3 v5 v4 v2 v1 V6 SOFz SOCl2 FClF2 530 344 752 806 492 364 I333 1251 528 XYzZ (CZ,) and XYZW (Cs) molecules Substitution of one Y atom lowers the symmetry to CzV and substitution of two Y atoms by different atoms lowers it further to Cs. In both cases the degeneracy is removed yielding six vibrational frequencies, all of which are i.r.- and R-active (Table 8). PLANAR PENTA-ATOMIC MOLECULES The (3n - 6) = 9 normal modes of vibration of a square-planar molecule are shown below: --- 390 284 434 748 455 326 v5 410 I94 703 v 4 si, s - / +Y- Hallam 51 XYq MoIecuIes (D4h) Of the above seven vibrations the three symmetrical modes, VI, vz and v4, involve no change of dipole moment and are i.r.-inactive but R-active.The v3, V6, and v7 modes are R-active but i.r.-inactive. The v5 fundamental vibrational frequency is inactive in both i.r. and Raman but can be observed in the hyper-Raman effect or as an overtone, It will be seen that no funda- mental vibration in the i.r. appears in the Raman and vice versa. This 'rule of mutual exclusion' indicates that the molecule has a cs and confirms that it is not tetrahedral. Some of the few molecules, and ions, which possess this structure are given in Table 9.v1 Molecule or ion - 543 v3 29 I v2 502 XeF4 R b+ [A u C 141- 347 3 24 144 TETRAHEDRAL PENTA-ATOMIC MOLECULES The nine normal modes are depicted below: 52 v7 v4 v5 --- V6 (221) 235 586 171 358 I23 (I:: R . I. C. Reviews XY4 Molecules ( T d ) All four vibrations are R-active but only ~3 and v4 are i.r.-active. Examples are shown in Table 10. Molecule or ion I v1 1 v2 I v3 Table 10: Vibrational frequencies (cm-1) of tetrahedral XY4 molecules and ions SiH4 C12F4 2180 908 459 cc14 SnC14 [NH4]+CI- [ PO41 3- Nai[S04]2- 3 68 3040 970 983 [Zn C1412- 28 I ~ ~~ ~~ * Fermi resonance doublet. 910 628 2183 1281 970 43 5 218 3 I 4 I06 I680 358 454 I30 82 X Y3Z (C3 ,,), X Y2Z2 (C2 v) X Y2Z W ( Cs) and X YZ W V ( Cs) molecules Successive replacement of the Y atoms lowers the symmetry which splits the degenerate vibrations and activates the i.r.-inactive vibrations to six in xY3z and nine in X Y2Z2, X Y2Z W and X YZ WV.A large number of these molecules exists and the vibrational spectra of many have been recorded; Tables 11 and 12 give a small selection. Molecule --- v1 486 408 I082 995 v2 1290 1035 786 669 or ion POC13 voc13 [FSO3]- [ S z 0 3 ] 2 - Table 12: Vibrational frequencies (cm-1) of tetrahedral XYZZZ and XY2Z'Jv molecules v2 v3 Molecule ------- S 0 2 F 2 553 497 848 608 SOzBrF v1 1269 1228 v5 v4 360 450 274 461 1502 1460 v4 -- 58 I 504 1287 I123 V6 I93 I29 409 335 v3 267 I65 566 446 v5 337 249 592 54 I v7 539 270 (:3* 131 I400 500 622 403 3 I45 I080 I106 277 Hallam V6 53 x Y5 MOLECULES Five (v3 to v7) are i.r.-active and six (v1, VZ, v5 to V8) R-active.Examples are listed in Table 13. P V7 Trigonal bipyramidal X Y5 (D3h) molecules The (3n - 6 ) = 12 normal vibrations are as follows: P b Mole- cule v2 v1 ---- PF5 PC15 SbC15 640 370 307 817 395 356 v3 945 441 V 5 V8 V6 4 P v4 576 30 I --- 534 28 I I00 I82 514 26 I I66 1026 58 I 399 74 Tetragonal pyramidal X Y5 (C4v) molecules Molecules possessing this structure are rare, examples being BrF5 and IF5.They should exhibit nine R-active vibrations of which only six should also be i.r.-active. OCTAHEDRAL xY6 (Oh) MOLECULES The (3n - 6) = 15 normal modes are illustrated below: 4 An octahedral molecule is centro-symmetric, thus the mutal exclusion rule applies, i.e. R-active vibrations (vl, v2 and ~ 5 ) are i.r.-inactive and the i.r.-active modes (v3 and v4) are R-inactive; V6 is hyper-Raman active only but can often be estimated from i.r.-active combination frequencies. Octahedral coordination is common in inorganic chemistry and many i.r. and Raman studies have been made of compounds which exhibit it; examples are listed in Table 14. POLYATOMIC MOLECULES There also exists a rapidly-expanding wealth of information on more complex compounds.These include ammine and amido, nitro and nitrito, aquo and hydroxo, cyano, carbonyl, acetylacetonates, and many others. Few studies on complex compounds of low symmetry yield complete structural informa- tion but approximate approaches based on the concept of localized group Hallam 55 v2 v1 v3 v4 Molecule or ion s Fs u Fs 775 668 31 I 344 644 532 3 29 320 [ S n C I 612- [ Pt GIG] 2- vibrations often yield useful information. The concept has proved extremely useful when dealing with metal-ligand vibrations particularly with complexes of heavy metal atoms. For example, they frequently provide clues as to the bonding site of a ligand. For illustration, in the case of the nitro group NO,, there are two possible sites of ligand attachment, at the oxygen or the nitrogen atom: 615 I89 I 86 825 825 940 626 294 330 M\o/N\o (11) 1330 1180 /O M-N \o Or (1) N-attachment (I) is unlikely to have such a pronounced effect on the NO, frequencies (vsN02 1300; vasN02 1260; 8NO2 825; see p.48) as would oxygen attachment (11). The following sets of frequencies have been observed.12 (1) 1400 1480 (11) and have been assigned to structures I and 11 respectively and serve to dis- tinguish readily between, for example, the isomeric nitropentammine and nitritopentammine cobalt (111) complex anions, [CO(NH3)5N02I2+ and [CO/NH3)50N0I2+. A more detailed application13 is in the structure of the carbonates. The structure of the free ion (D3h) and its spectrum has been discussed (p.49). When the group is covalently bonded to other groups as in organic carbonates and hydrogen carbonates, or is coordinated to a metal ion by a bond which has partial covalent character, the symmetry of the group will be changed from that of the free ion. As a result of the decreased symmetry additional bands will arise due to the removal of degeneracy. The ranges of these for the various possible structures are shown in the correlation diagram in Fig. 3. These readily allow one to distinguish between unidentate and bidentate carbonate ligands. Studies of polynuclear metal carbonyls have shown that terminal (C=O) groups consistently absorb at 2100-2000 cm-1 whereas bridging (C= 0) groups which are of lower bond order, absorb at lower frequencies, 1900- 1800 cm-1.Similar but less comprehensive studies have been made on cyano complexes where it appears that a bridging cyano group absorbs at a higher frequency than does a terminal cyano group. 56 V 6 v5 I- 524 202 (363) (144) I 58 I62 R. I. C. Reviews \ +--=- \ . / / / HaIlam i a, 0 -0 - 0 -0 9 \ \ \ \ \ h iI----- - 0” v 57 I I I I I T I I I 1 I I I so;- HSO, so; - s,o,z - s,o;- Se0;- SCN- c to, cio; Po:- cr0:- Mn0:- w0:- BrO, 10, vo; Cr202,- 400 1200 1400 1600 200 600 1000 800 cm-’ d Fig. 4. Characteristic i.r. frequency ranges of some polyatomic inorganic ions. CORRELATION TABLES As with organic compounds, measurements of i.r.(and Raman) spectra of large numbers of compounds of similar structural types allow correlation tables of characteristic group frequencies to be drawn up. It should be stressed however that these are of far less value than their counterparts in organic chemistry. This is to be expected since organic molecules have a very limited range of atomic masses and force constants compared with inorganic molecules. Furthermore, most metal-ligand vibrations are at lower frequencies than C-X frequencies and therefore the possibilities of vibrational interaction are greater. The i.r. correlation table illustrated (Fig. 4) will serve as a preliminary guide towards spectral interpretation and qualitative identification; references 58 R.I . C. Reviews 5-7 and 9 should be used for any serious qualitative analysis. Great care must, however, be exercised in this empirical approach because frequency ranges are affected by many factors such as physical state, coupling of vibrations, and the nature of the cation or anion. Most of the data in Fig. 4 refer to the solid state which is the physical state in which inorganic com- pounds are usually examined. AQUEOUS SOLUTIONS Water is the commonest solvent used in inorganic chemistry and it is there- fore essential that vibrational spectroscopy should contribute to our know- ledge of the structure and reactions of inorganic species in aqueous solution. Because of the very weak scattering by liquid water, aqueous solutions are ideally studied by the Raman technique.With the introduction of the new laser sources these studies now embrace1 highly coloured ions. Recent developments in sampling techniques however have revived interest in the i.r. of aqueous solutions. Multi-reflection ATR (attenuated total reflection) units allow the i.r. beam to be reflected many times from the interface of a solution in contact with a crystal of relatively higher refractive index such as silver chloride or germanium. This circumvents the practical difficulties of filling and cleaning the very short path length transmission cells necessary for aqueous solutions and ATR cells provide quite satisfactory i.r. spectra of aqueous solutions over a reasonable spectral range.It appears, however, that the ATR technique possesses no inherent advantage over the transmission technique for aqueous solutions and very recently it has been demonstrated14 that careful use of conventional transmission methods using water-insoluble window materials (IRTRAN-2) can yield acceptable spectra. Goulden and Manning14 have studied a large number of inorganic compounds in aqueous solutions of varying pH over the range 950-1550 cm-l. In contrast to the corresponding solid state spectra they find the absorption bands of the dissolved ions appear over very limited frequency ranges and are in- dependent of the other ions present. Correlation charts (Fig. 5) constructed from such data permit unequivocal identification of most common inorganic ions.Care must be exercised in the choice of pH and high concentrations should be avoided. At higher concentrations ionic interactions may lead to the appearance of additional bands as the symmetry of the free ion is reduced; the technique, of course, allows such interactions to be studied. The impetus now given to i.r. and Raman spectroscopy of aqueous solu- tions will provide precise knowledge of ionic species in aqueous media. For example, alkaline solutions of zincates and aluminates are generally now described as existing as [M(OH)4In- ions although earlier texts write these as Mot- species. Raman spectral5 of zincates are consistent with the presence of tetrahedral [Zn(OH)4I2- species but recent i.r. and Raman studies16 of aluminates fail to identify the presence of the corresponding [Al(OH)4]- ion.In very alkaline solutions (PH > 12.5) the species is linear AlO,. Near pH 12.5 A10, is in equilibrium with a second form which completely takes over at pH values down to pH 8. The second form is probably an octa- hedrally coordinated polymeric anion although a square planar [Al(OH)4]- Hallam 59 Ion so;- HSO, so: - HSO s,o; - so, aq CI 0, ClO, P o i - HP0:- H,PO, COf - HCOS NH; Cr,O 5- Fig. 5. i.r. Frequency correlation chart for aqueous solution spectra of some polyatomic inorganic ions (adapted from Goulden and Manning14). is not entirely ruled out. Clearly however, in the dissolution of Al(OH)3 in the alkaline conditions described in qualitative analysis schemes, it is the A10; species which is present. BOND PROPERTIES Numerous correlations exist between vibrational frequencies and other bond properties which can provide simple and rapid estimates of such properties.For example, the symmetrical stretching frequencies for a large number of halides are known and their values change in the same order as the stability constants where these are known. Metal-oxygen stretching frequencies for a limited number of hydrates are known and these indicate that the order of stability of the hydrates is, e.g. Mn2+ < Cu2+ > Zn2+, i.e. in agreement with the normal Irving-Williams order for the stability of the complexes. SUMMARY Both i.r. and Raman spectroscopy give essentially the same information about molecular structure.For centro-symmetric molecules they are exactly complementary and for complete structural information it is desirable to 60 I I I I I I I I I I I I I I I I t I I I I I I I I I I I1 crn-’ --t R.I.C. Reviews have both i.r. and Raman spectra for the compound. Qualitative inorganic analysis can, with profit, incorporate either technique and should be encour- aged at the undergraduate level for the identification of ions, since the spectra also lead to the structure of the ion. Had such techniques been used more frequently in the past many over- simplified descriptions in inorganic chemistry might have been avoided. One obvious example is the description of the metaborate ion as BO,. A simple examination of the i.r. and/or Raman spectra would have eliminated the existence of the linear [0-B=O]- ion in most metaborate systems although a more detailed spectroscopic study would have been necessary to suggest the chain or cyclic trimeric units such as that below which actually exist in these salts. 0 I I 3- REFERENCES Elsevier, 1963. 1 I. R. Beattie, Chem. in Britain, 1967, 3, 347. 2 Manse1 Davies (ed.), Infrared Spectroscopy and Molecular Structure, Amsterdam : 1965. 3 R. G. J. Miller (ed.), Laboratory Methods in Infrared Spectroscopy, London: Heyden, 4 I. J. Worrall, Molecular Symmetry, R.I.C. Lecture Series, 1967, No. 2. 5 K. Nakamoto, Infrared Spectra of Inorganic and Coordination Compounds, New York : Wiley, 1963. 6 F. A. Miller and C. H. Wilkins, Analyt. Chem., 1952, 24, 1253. 7 F. A. Miller, G. L. Carson, F. F. Bentley and W. H. Jones, Spectrochim. Acta, 1960, 10 L. A. Woodward, Quart. Rev., 1956, 10, 185. 16, 135. 8 K. E. Lawson, Infrared Absorption of Inorganic Substances, New York: Reinhold, 1961. 9 D. M. Adams, Metal-Ligand and Related Vibrations, London: Edward Arnold, 1967. 1 1 R. N. Jones and Magda K. Jones, Analyt. Chem., 1966, 38, 393. 12 R. B. Penland, T. J. Lane and J. V. Quagliano, J. Am. chern. Soc., 1956, 78, 887. 13 B. M. Gatehouse, S. E. Livingstone and R. S . Nyholm, J. chem. SOC., 1958, 3137. 14 J. D. S. Goulden and D. J. Manning, Spectrochim. Acta, 1967, 23A, 2249. 15 E. R. Lippincott, J. A. Psellos and M. C. Tobin, J. chem. Phys., 1952, 20, 536. 16 L. A. Carreira, V. A. Maroni, J. W. Swaine and R. C. Plumb, J. chem. Phys., 1966, 45, 2216. 61 Hallam

ISSN:0035-8940    

DOI:10.1039/RR9680100039

出版商:RSC    

年代:1968

数据来源: RSC

摘要:

STRUCTURE AND PROPERTIES OF WATER D. J. G. Ives, D.Sc., A.R.C.S., F.R.I.C. Birkbeck College, Malet St., London W C I and T. H. Lemon, B.Sc., M.Phil. Post Office Research Station, Brook Rd., Dollis Hill, London N W 2 Introduction Bases of comparison A general view of related hydrides, 65 General comment on solid-liquid-gas transitions, 66 A structural theory of liquids, 68 Liquids over a wide temperature range, 71 . . .. . . . . . . .. . . . . . . . . .. The abnormal hydrides . . .. . . . . . . The hydrogen bond Infrared and Raman spectra Theories of the structure of water . . . . . . .. . . . . .. .. .. .. .. .. The water molecule, 87 Ice, 87 x-Ray scattering, 9 1 The Bernal and Fowler theory, 91 The distorted bond model, 92 Interstitial models, 92 Cluster models, 94 Thermal anomalies Hydrophobic bonding and co-solvent behaviour * ... .. .. * . . . . . 100 Conclusion . . .. .. .. .. .. .. .. . . 102 Water is familiar to us from our earliest years. We get it by turning a tap, and what could be less remarkable? We see that ice floats on water, and we are shown at school that water has a maximum density at 4°C. We may be told that these are happy circumstances because, if ice did not act as an insulating upper layer and, if slightly warm water did not remain at the bottom, lakes and rivers might freeze solid, with fatal effects to fish. Mildly interesting this, but natural, i.e. ‘according to, or provided by, nature’.l We are not by experience inclined to think of water as anything but normal, and there is certainly a lot more of it than of any other liquid.Little may happen during a chemist’s education to disturb this phlegmatic attitude to water. It was the solvent predominantly used in the early develop- ment of the physical chemistry of solutions (with, oddly, sucrose as the most normal of solutes). It is, no doubt, gradually borne in upon us as students that water has outstandingly convenient properties. It is fortunate that it can remain in the liquid state over a curiously wide and high range of tempera- 62 62 65 . . .. .. .. .. .. .. . . .. 75 .. 81 . . 78 .. 86 .. 99 .. .. RIC Reviews tures, and that it is the nearest approach we have to the universal solvent- especially that it is an ‘ionizing solvent’, ostensibly because of its high dielec- tric constant.It is also fortunate that water has a broad domain of thermodynamic stability2 and can participate in acid-base equilibria over a range of 16 pH units, and sustain redox equilibria over a potential range of more than two volts. A great deal of chemistry depends on these providential properties of water, but it is the chemistry rather than the water which receives attention. There is little change in emphasis when we come to more recent fundamental studies of equilibria and reaction rates in aqueous solutions, or even to work leading to present knowledge of electrolytic solution^.^ Interionic attraction theory has reached a sophisticated level by treating the solvent-water or another-as if it were merely a ‘dielectric continuum’.The physico-organic chemists have founded an unassailable structure of theoretical organic chemistry very largely on studies of reactions in solution.4 They have used a variety of solvents, including water and aqueous-organic mixtures, and have recognized the profound importance of solvent effects.5 Until quite recently,6 however, they have not been forced to give detailed consideration to the problems of solvent structure and of solvation pheno- mena in their own right. It is as if there were some general, simplifying prin- ciple which has shielded chemists from the effects of the more complex phenomena of nature, allowing them often to get away with treating solvents just as media, of status little more than that of reaction vessel.It will be suggested later that there is indeed such a principle. In these circumstances, students are not to blame if they graduate knowing no more about water than that it is an abnormal, hydrogen-bonded liquid. Further, the growth of science imposes a progressive squeeze on syllabuses, leaving little time for overly-debatable topics. Not even normal liquids are well understood, so water, as a mystery within a mystery, gets little attention. The same applies to the hydration of ions-a messy and contentious subject, fortunately not obtrusive-and still more to the hydration of molecules. The study of binary liquid systems is an associated field, formally intimidating and difficult of interpretation without exhaustive mustering of diverse lines of evidence,7 so usually this also gets short shrift.If it is agreed that solutions are important-aqueous solutions above all others-must it not be admitted that solute-solute, solute-solvent and solvent- solvent interactions should all be studied in a reasonably balanced way? Would it be an overstatement to say that, in general and for many years past, only the first and least important of these interactions has been studied seriously by the average student ? Activity coefficients, Debye-Huckel, and all that-thermodynamic formalism and a theory of simple electrolytes admittedly based on a grossly oversimplified model. Perhaps the progress of physical chemistry in a vitally important direction has been retarded by the constraints and limitations of undergraduate courses-students can hardly become interested in a subject if they barely know of its existence.Water is the keystone of this argument, because of its ubiquitous and supremely important role in natural processes. All living structures are composed mostly of water. Many lines of enquiry are now converging on Zves and Lemon 63 5 0- Y n 0 60, ,3 49 - 1 44 5.58 I O! TM 600 500 - 400 + 2 0 300 F E 2oc I oc 0.43 C 3 Ne L 0 079 the problem of the structure of water, how this structure varies with tempera- ature and is modified in the vicinity of charged and uncharged solute particles and of charged and uncharged interfaces, and how water may play a specific part in life-processes.Over the last several years there has been rapidly growing awareness of this situation, and at the present time there is intense activity directed to these problems; there is also as little agreement as there is much debate. The special difficulty of reviewing a subject in this state has been met by a plan discernible from the table of contents-first to establish a broad foundation to support consideration of the more special problems. It might be described as a conscious effort to avoid not being able to see the wood for the trees. If this could be well done it might turn out that much had been learned about water ‘in advance’, but whether the reviewers succeed or fail in this, the main objective may still be achieved-to demonstrate the importance, difficulty and fascination of the challenge that this subject presents.BASES OF COMPARISON Abnormality must be considered by reference to standards of normality. Water is a volatile hydride and it is a liquid. It is reasonable first to give a general view of how water is abnormal by reference to other hydrides and to other liquids, first deciding what may be considered normal for each. A general view of related hydrides Water, as a hydride, is to be compared with its congeners in the periodic table. The basis for such a comparison is provided in Fig. 1, which displays melting and boiling points, critical temperatures (TM, TB, Tc) and appropriate latent heats (AHm, AHE, kcal mole-l) for hydrides of elements in the first, second and third periods.A noble gas as ‘zeroth hydride’ is included in each of the three isoelectronic groups. There is a close resemblance between the second (Ar-SiH4) and third (Kr-GeH4) groups; they hardly differ except for an expected displacement on the temperature scale. This must be due to similarities in the cohesional forces responsible for the existence of the liquid and solid phases of the hydrides in these groups. These forces are accounted for by dispersion, dipole-dipole and dipole-induced dipole attractions*- collectively, van der Waals interactions. In general, except for highly polar molecules, the dispersion or London forces predominate; they have the important characteristic of additivity, i.e.the total interaction between a number of molecules is the sum of the interactions of all the pairs of mole- cules.* * These forces, acting between a pair of molecules, give rise to the following interaction energy in which the ternis can be easily identified : cy is polarizability, v is the zero-point frequency of oscillation responsible for the dispersion effect, p is dipole moment, r is intermolecular distance and other symbols have their usual meanings. It may be useful to note the trends of dipole moments and polarizabilities shown in Fig. 1 . Ives a d Lemon 65 It is very significant that the resemblance between these two groups of hydrides extends, broadly, to the differences between boiling points and melting points (TB - TM): to explain why, and to provide a background for discussion of the behaviour of the abnormal hydrides, calls for general comment on transitions between solid, liquid and gaseous states.9 General comment on solid-liquid-gas transitions At sufficiently low temperatures, the ordered structure of a crystalline solid is such as to minimize potential energy. With rising temperature, entropy must increase and potential energy must follow suit.The entropy requirement increasingly militates against the tendency to minimization of energy-in other words, rising temperature compels a shift of the competitive balance between the influence of order-producing forces and the randomizing effect of thermal agitation in favour of the latter. Although melting marks a defeat for the ordering forces, it may not be a major defeat, since order is relaxed in stages, of which melting is not the last,g or, in some cases, the first.Melting may be viewed thermodynamically in terms of the relation AGM = AHM - TASM representing the Gibbs free energy, enthalpy and entropy changes accompany- ing the process at temperature T. For a hypothetical solid-to-liquid transition below TM, AGM is positive-the process is unnatural and does not occur. But AHm and ASM are different functions of temperature,* and, as tempera- ture rises, TASM overtakes AHM, and brings AGM to zero at TM, where solid and liquid phases are in equilibrium. Figure 2a presents approximate data for the melting of ice. A broader view of phase transitions is given by Fig.2b, consisting of formal plots of free energy G (on no absolute scale), against temperature. The lines (for simplicity, straight) have relative slopes, (aG/aT), = - S, broadly consistent with relative entropies of solid, liquid and gaseous phases at 1 atm pressure. The intersections of the lines determine TM and TB. At some prevailing pressure much lower than 1 atm-below the triple-point pressures of normal substances-a very steep ‘gas line’ would cut the ‘solid line’ directly. A native observer on the moon might be unaware of the existence of liquids. Under more normal conditions, a ‘liquid line’ cuts off the corner between the other two, representing the fact that a liquid, with assistance from external pressure, can retain effective cohesion despite the absence of long-range order in the liquid state; the cohesion is always less, however, than that of the ordered solid state from which the liquid was derived by melting.Inspection of Fig. 2b indicates that if a liquid is of high entropy, so that the ‘liquid line’ is steeper, then, other things being equal, the intersection with the ‘gas line’ would be at a higher temperature than for a liquid of low entropy, i.e. TB would be raised. It is generally true that the necessity for jumping into the next state of greater randomness is deferred to a higher RZC Reviews Solid - Gas 1 atrn L G\ u ' - 5 - I I I I I I I I I I I I Temperature ("Cl Fig. 2. (a) Thermodynamic functions for ice-+water; (b) formal G(T) plot embracing solid- liquid-gas transitions.temperature the higher the entropy of a phase, or the greater its capacity for accommodating entropy. Similar thought can be applied to other features of the basic diagram. At the boiling point, TB, cohesion of the liquid fails, but it is to be noted that TB is not normally far ahead on the temperature scale of the intersection of the gas line and the extrapolated solid line, suggesting that cohesion in the liquid is related to, if less than, cohesion in the solid. An approximate assessment of either can be made because ASB, the large entropy increment accompanying evaporation at 1 atm pressure does not vary greatly (except in special cases) from one substance to another ; its near-constancy is expressed in Trouton's rule.Since TB = AHB/ASB for reversible evaporation, TB serves as a rough measure of the energy (- AHB) required to overcome the cohesional forces.* No similat statement can be made about TM because of the very wide variation of ASM, from one substance to another. All that can be said, in terms of TM = AHM/ASM, is that the melting point will depend on the magnitude of the cohesional forces to be overcome in melting and on the disparity of entropy between liquid and solid phases. No ad hoc conclusion can be drawn from whether TM is high or low, because four factors are involved-energies and entropies in two phases. This uncertainty is reflected in the great variation of (TB - TM) shown by substances, even those of not dissimilar TB.For instance, boron and tin have boiling points of 2823 and 2533°K respectively, but (TB - TM) values of 250 and 2028". Boron is a * It could be argued that TC would be better than TB as such a criterion. It can be replied that Tc has less fundamental significance than is usually assigned to it because there is no similar temperature known to limit the existence of solid phases, i.e. above which a tion is low and compressibility approaches infinity. Since, however, TB - +Tc (Guldberg's normally given encountered because Moreover, VC - solid 12 Nw, cannot where exist. N is Avogadro's often TC number TB S- may TB and and be w more conditions is the suited volume near to of discussion TC a are molecule, far from of liquids coordina- normal as : rule), approximate argument can be conducted on either basis.Ives and Lemon 67 hard solid built of interlinked BIZ icosahedra-a structure of great geometrical specialization formed by strong, directional bonds; tin is a metal, not close- packed in the solid state, owing its cohesion when molten to strong, non- directional metallic bonding. No doubt this is an extreme example, but it is true to say in general that TIM and (TB - TM) depend critically on the nature of structures and forces in solid and liquid states. It is therefore evident that the previously noted resemblance between the second and third groups of hydrides (Ar-SiH4, Kr-GeH4) in Fig.1 is more remarkable than was at first apparent. This may strengthen the justi- fication for using their common pattern of behaviour to assess abnormalities in the first group, but, before making such comparisons, it is well to reflect that there is nothing absolute about such a criterion. The considerable variation of (TB - TM) between the members of the ‘normal’ groups suggests that they have their own structural problems. Since interest centres on the liquid phase, there is a need to define a normal liquid and to know something about its behaviour, particularly from the structural angle. It would hardly be sensible to approach the problem of the peculiarities of water as a liquid without attempting to satisfy this need.A structural theory of liquids Comparatively recently it was said that ‘we are completely in the dark as to the degree of order in a liquid’.lO This is disconcerting, since the problem is of such wide interest and also because the generally high specific heats of liquids suggest that they do have considerable order to lose and entropy to gain. In these circumstances, it is reasonable to turn in the first instance to the noble gases. Their monatomic molecules attract each other only by short- range, non-directional dispersion forces. Within close limits they have the same reduced melting point TM/Tc (0.556 & 0.002) and the same entropy of fusion, ASM (3.29 & 0.05 cal OK-lmole-1). This uniformity in properties (which could be illustrated further) indicates that they are likely to form the simplest kind of liquid phase, suitable as a standard of normal behaviour.ll Considering the noble gases together, we note that low TB and AH, (Fig. 1) denote feeble cohesion.But (TB - TM) is small so that, relatively, TM is high. Noting that triple-point pressures are high (Ar, 0.68 atm), that RIC Reviews TM is proportional to AHA[ because ASM is constant and that AH, and AHB vary with atomic number in a manner to be expected from dispersion forces alone (Eqn. (1); Fig. l), we see that ‘high TM’ is not due to order- promoting, stabilizing forces of attraction characteristic of the solid state. On the contrary, the solid state is such as to be able to gain entropy with the least restriction from such forces.The noble gas elements form close-packed solid phases, each atom having 12 equivalent nearest neighbours: this is a structure determined by forces of repulsion rather than attraction, i.e. the simplest regular structure to be expected from the most space-economizing packing together of hard spheres. When the solid noble gas elements melt, they expand by 15 per cent-an additional uniformity in behaviour which suggests that they provide us also with the simplest melting process, adequately defined as a transition from regular to irregular closest packing. This occurs 68 when the constancy in number, and the symmetry in arrangement of nearest neighbours essential to preservation of long-range order can no longer be sustained against the increasing attack of thermal agitation.This is essentially the picture used by Bernal in his ‘polyhedral hole theory’ of these simple liquids,l2 which is considered to present a new structural viewpoint basic to the understanding of liquids in general. In this theory, Bernal views a liquid neither as a ‘blurred solid’ nor as a condensed gas, but directly, in terms of how a homogeneous, coherent, densely but irregularly packed assembly of spherical molecules can be described structurally. The essential feature of the liquid state is variation of coordination-a continuous change in identity and number of nearest neighbours on which the self-diffusion and fluidity of liquids must depend. It may seem odd to look for structure under circumstances which doubly forbid the existence of long-range order, but it is to be remembered that the yield and shear involved in fluid flow are slower by orders of magnitude than the mainly vibrational thermal motions of the molecules.On a short enough time-scale, liquids show solid-like properties (as in ultrasonic and neutron diffraction studies); at any instant they are geometric- ally, rather than physically, dissimilar to crystalline solids, and the dis- similarity is that of irregularity as compared to regularity. A liquid may be regarded as passing through a random sequence of irregular molecular arrangements evolving continuously from each other. There is a vast number of such energetically equivalent arrangements, compared with the very few regular ones, and a constant flux between them.It is the ready transition from one to another under the slightest stress which is the basis of fluidity. The question arises whether these irregular arrangements have any common structural features; whether liquids present a static as well as a dynamic problem. Bernal studied the static, structural aspect by a method briefly explicable in the following way. Consider four spheres of identical radius in mutual contact. Imaginary straight lines between their centres trace the equal edges of a regular tetra- hedron enclosing the gap or ‘hole’ between the four spheres. The hole may be called a ‘tetrahedral hole’. In any regular assembly of many such spheres each making the maximum number of contacts with others, only two kinds of ‘polyhedral hole’ are possible-tetrahedral and octahedral.Such a ‘spheres- in-contact’ method can give only a model of a regular crystal, but it may be noted that the structure can be described in terms of an assembly of poly- hedra (containing holes) just as well as in the more conventional way. In building a model of a liquid, there must be no geometrical constraint on each added sphere to occupy a position precisely defined by its previously built-in neighbours. The model must be looser to represent the higher energy, less cohesion, great volume and, above all, the irregularity of the liquid. This representation becomes possible if nearest-neighbour distances are allowed latitude to increase by up to 15 per cent above the minimum corresponding to spheres-in-contact.With no longer any repeating unit of pattern, the model must be large. Each sphere must be placed randomly within the general requirement that, when all have been placed, the best economy of Ives and Lemon 69 space in irregular packing is achieved, with the least possible stretching of nearest-neighbour distances (maximization of entropy, minimization of energy). To build such a model is an exercise in statistical geometry-ideally the model should be allowed to ‘build itself’, without human interference. The most successful attempt to solve this problem objectively was experi- mental. Conducted by Bernal and his team, it was based on a random heap of a large number of ball-bearings, consolidated by dousing with paint- subsequently allowed to drain and harden.Clearly-marked ball-to-ball contacts could ultimately be counted, and confirmed the essential variability of coordination. By an ingenious method, the spatial coordinates of each ball in the heap were determined, distances and directions were computed, and a ‘transparent’, much enlarged, ball-and-wire replica of the heap was constructed. Examination of this open model showed that it contained (apart from a few significant ‘accidents’) holes of only five polyhedral types : tetrahedron, octahedron, and three somewhat larger ‘deltahedra’ (poly- hedra with triangular faces) lacking the symmetry for regular packing. These are the only polyhedra, generated by sphere packing, small enough not to admit an extra sphere.It was shown independently that these five poly- hedra can be packed together, with but little distortion from equality of edge-length, in an indefinitely great number of space-filling ways, but the minimum volume for such packing is always 15 per cent greater than that for regular cl ose-packing. That this model successfully simulated the simplest liquid was supported by calculation of a radial distribution curve* for the model-it was in agree- ment (with appropriate change of distance scale) with that derived for liquid argon by neutron scattering. The following less-expected features of the model therefore merit close attention. The irregularity of the model was not absolute.Tetrahedral holes, which predominated, tended to occur in company, orderly arranged. Tetrahedra can pack together economically of space, forming a number of structures, some closed (e.g. 20 tetrahedra each sharing three faces, with a common vertex at the centre), others not (e.g. a ‘triple helix’ formed by face-sharing tetra- hedra). The density of packing in such local regions is higher than can be attained in regular close-packing, but all the ordered structures formed in this way must remain local, because they incorporate an element of five-fold symmetry which forbids their indefinite extension in space. Bernal calls them ‘pseudonuclei’ because they are like crystal nuclei which can never grow, and can never form the basis of long-range order.In considering whether these features of the model indeed reflect what happens in the simplest liquid, we recall the standing conflict between the interests of energy and entropy, and ask why, in the liquid, energy-minimizing forces should not promote a kind of order forbidden to the solid simply because it is of the wrong symmetry? When the solid has melted, the prohi- bition of five-fold symmetry is lifted, and local order based upon it could conceivably be more probable than any other. At least the thought should be retained that order in a liquid need not be related to that of the solid from * RadiaI distribution curves show the probability, as a function of distance from any molecule, of encountering another molecule. RIC Reviews 70 which the liquid was derived by melting.The occurrence in liquids of local regions of less than average energy and entropy is acceptable thermodynami- cally ; constant free energy accommodates such fluctuations,* which are, indeed, to be expected.13 If it is accepted that simple liquids contain evanescent, super-dense pseudo- nuclei, it must be remembered that the liquid as a whole is less dense and of higher energy than the corresponding regular solid. There must be other local regions with density lower and energy higher than the average. Reference to the model indicates a higher than average proportion of the larger, more irregular, polyhedral holes in the immediate vicinity of the pseudonuclei, i.e. the denser, more ordered regions are surrounded by emptier, more disordered zones.Several of Bernal’s experiments suggest that in assemblies in which there is an order-disorder balance, any locally ordered region may be adjoined by a misfit zone of greater than average disorder. If this is generally so, it may have implications in relation to solvation and to special effects at interfaces. The model could, of course, represent only one possible configuration of an assembly of molecules loosened just enough to be able to acquire the irregularity of the liquid state, or of a just-melted solid. Valuable as it has proved to be, it could not be expected to give guidance on the changes occurring in liquids with rise of temperature towards Tc-with one possible exception. The ‘accidents’ previously mentioned were holes large enough to admit an extra sphere; they correspond with vacancies-the kind of accident to be expected particularly in liquids.Liquids over a wide temperature range The minor expansion normally associated with melting is just adequate to establish the irregularity of the liquid state. With rising temperature, it is followed by a gradual but major expansion-commonly, volume is more than trebled between triple and critical points. A general theory must say how liquids use this excess volume. It is not unreasonable to consider that multiplication of volume between TM and TC may involve profound change in the state of a liquid. At temperatures little above TM, liquids can be viewed as melts, best to be understood in terms of a lattice theory, and to be compared with the solid phases from which they came.Towards the upper end of the temperature range, comparison with the gas phase might be more appropriate -the classical ‘continuity of state theorem’ certainly suggests this. Both approaches have been used in discussing liquids. Liquids, however, are very diverse-even in the restricted field of molecular liquids-because of wide variation in kind and symmetry of intermolecular force and of molecular size and geometry. Perhaps, therefore, no more than a general theory of broad principles is to be looked for. Particular liquids may require their own theories within the general framework-perhaps one for temperatures near TIM, and another for higher temperatures where entropy requirements have prevailed more or less decisively. For the first kind of particular theory it may be relevant to consider that structure in crystalline solids has always * G = H - T S .Ives and Lemon 71 been revealed by external form; structure in liquids has been concealed by lack of it. Eyring’s ‘significant structures theory of liquids’l* encourages this attitude. It is based on a generalized model of liquids which ascribes both solid-like and gas-like behaviour to the molecules they contain, and is de- signed to accord with two main observations. First, x-ray scattering studies of simple liquids such as argon over the temperature range between TM and Tc show that the average coordination- 10 or 11 just above TM-steadily decreases with rising temperature to about four at about 5” below Tc, but the distance between nearest neighbours remains almost constant.Expansion is therefore largely confined to the holes between the molecules. Nearer to Tc attractions begin to fail, nearest-neigh- bour distances increase, and coordination rises to six. The theory describes the holes as ‘fluidized vacancies’ and suggests that they move about in the liquid as freely as molecules in a gas. This involves the second main observation, which is the classical ‘law of rectilinear dia- meters’-to the effect that the average of the densities of a liquid and its equilibrium vapour is nearly independent of temperature, but decreases slowly and linearly from TM to Tc.The liquid seems to bear some kind of reciprocal relationship with the vapour. It is adopted as a cardinal principle of the theory that the liquid indeed ‘mirrors’ the gas or vapour with which it is in equilibrium. To quote:14 ‘individual molecules translating in the vapour are mirrored as vacancies translating in the liquid. Rotating molecules are mirrored as rotating vacancies. Association of vapour molecules is inatched in the liquid by association of the corresponding vacancies which translate and rotate like their molecular counterparts.’ The translation of vacancies is of course due to the contrary movement of molecules (just as the migration of positive holes is due to movement of electrons), but there is only one phenomenon and it is immaterial which description is applied to it.The theory requires some of the molecules some of the time to have translational freed om. The solid-like behaviour of molecules has nothing to do with any crystal- linity in the liquid; it depends only on the vibrational motion of a molecule temporarily trapped in the potential well created by its nearest neighbours. The motion is adequately represented by an Einstein oscillator* of the same characteristic temperature, 8, as that of the parent solid. Occasionally, such a molecule will acquire enough energy to push its neighbours aside, and a vacancy is generated. The energy required to create a vacancy is directly related to the sublimation energy of the solid. The theory gives a very simple answer to the problem of what proportions of the molecules are gas-like or solid-like (i.e.having translational or vibra- tional motions) at any temperature. It is that (VL - VS)/VL is the fraction * The Einstein theory of the heat capacity of a crystal, considered as N independent oscillators gives ,o co = 3Nk ( ;)2 J 1‘ - 1)2 where 8 = hv/k; v is a characteristic frequency, and other symbols have their usual mean- ings. When T 9 8, Cu -+ ca 6 cal°K-lmole-l. RIC Reviews 72 of gas-like, and VS/VL is the fraction of solid-like molecules, where VL and VS are molar volumes of liquid and solid respectively. It is remarkable how well the simple expression fits the heat capacity data for liquid argon from TM to Tc. The theory is tested by setting up a partition function for the liquid as a product of partition functions for solid- and gas-like states (the ‘significant structures’), weighted according to their contributions, i.e.where N is the Avogadro number. For the boundary conditions VL = VS and VL > VS, f L assumes the limiting valuesfs and fG respectively. Appropriately expanded with feeding-in of molecular weight, Einstein 8, sublimation energy and molar volumes, this partition function can be used to calculate the thermodynamic properties of a liquid from TM to Tc. The results are then compared with the experi- mental data. In general, tests of this kind have been remarkably successful for a wide variety of liquids; the model on which the theory is based must therefore be given weight.On the grounds that there is no majority group of liquids that can safely be described as normal, it is expedient to abandon the attempt to define a normal liquid in favour of tracing the events that may occur and the sequence of states that may be traversed on the way from solid to gas. To do this adequately would require much closer attention to the phenomenon of melting15 than is at present permissible. The brief resumd which is alone practicable must start with the statement that crystalline solids in their lowest free energy states contain defects-the cost in energy to produce defects is met by entropy profit-at concentrations which increase exponenti- ally with rising temperature. There is, of course, no reason to suppose that most solids are not in a state of frozen-in disequilibrium but this is likely to be relieved as TM is approached from below, and, particularly if co- operative positional defects develop, ‘pre-melting’ phenomena may be observed-they are in any case to be expected from fluctuation theory.It is of interest that some melts show pre-freezing effects as TIM is ap- proached from above, but they do not normally ‘mirror’ pre-melting as might be expected. The general inference is that melting may not be as sharp as unsophisticated thermodynamic theory suggests. In the present context, however, the main point is that the expansion which normally accompanies melting so greatly lowers the energy to produce defects that positional disordering becomes almost, if not quite, complete and all long-range order is lost.In this order-disorder transformation the distinction between right and wrong positions vanishes, so that ‘positional defect’ loses meaning. Nevertheless, in the fluid state attained by melting, it is fundamental to expect the fluctuating balance of the ‘energy-entropy conflict’ to be even more significant, if less externally apparent. Whereas in the ordered solid lves and Lemon 73 state fluctuations had to be towards disorder, in the liquid we should expect them to be of both kinds-toward local order in the interests of energy minimization, and towards even greater disorder as entropy prevails. We look for an even greater prevalence of ‘defects’, but they now need redefining by reference to a system not considered defective.It is reasonable to adopt Bernal’s model of the ideal ‘monatomic liquid’ as the new basis of comparison-or an averaged-out version of the assemblies represented by alternative packings of the five ‘canonical holes’. The less regular of these polyhedral holes cannot be considered as defects, since they are formed with insignificant energy effect. What is to be looked for is something of enhanced energy and entropy, increasing in concentration with rising temperature until cohesion is undermined and the system is ready to take the major entropy jump of transition to the gas phase. Bernal’s ‘acci- dents’-holes large enough to admit an extra molecule-or Eyring’s ‘fluidized vacancies’ answer this description.The ‘hole theory of liquids’ in this sense is generally acceptable and is basic to mass-transfer phenomena in liquids- self-diffusion and fluid flow. Bernal’s ‘pseudonuclei’ represent the opposite kind of fluctuation, but a greater width of possibilities needs consideration. The relaxation of symmetry restrictions imposed by the long-range order of crystals, and the greater volume and freer movement of molecules make possible the formation of various types of more or less evanescent clusters with a variety of structures. If crystal-type symmetries are not excluded, but must merely take their chance, the following types of cluster may be envisaged in liquids, probably confined to temperatures not too far above TIM. ( 1) Quasi-crystalline, or ‘crystallizable’ clusters, with structure related to that of the solid phase in equilibrium with the liquid at TIM.The adjective ‘quasi-crystalline’ is often used loosely; it should be confined to the result of incomplete ‘mucking-up’ of a crystal, by expansion and introduction of defects of various kinds in number insufficient to complete the destruction of order. This creates a picture of a tattered remnant of a crystal, but there is the possibility of a tighter, less imperfect cluster too small to serve as a nucleus for growth. (2) Quasi-crystalline clusters with structures related to some other real or conceivable crystalline phase the substance might form, other than that normally existing at TM-still, in principle ‘crystallizable’ and perhaps either ‘tattered’ or ‘tight’.(3) Anti-crystalline or non-crystallizable clusters-pseudo-nuclei, of growth-forbidding symmetry. If these are small ‘closed’ units, they might tend to associate together in a quasi-crystalline way. They may, as Bernal showed, be capable of extension in one dimension. Disordered clusters, distinguished only by lowered energy and changed density, cannot be ex- cluded; they might, indeed, be the most probable of all, because of the lesser penalty in entropy. There are no grounds to deny that more than one-perhaps several-of these possibilities might be open to a single liquid. From one liquid to another, preferred types of clustering would clearly depend on the directional nature of forces and the shapes of molecules.There are but two further RIC Reviews 74 comments to be made on this largely unexplored because nearly inaccessible problem. It is believed that attention should be paid to Bernal’s finding that, adjacent to local order, special misfit disorder is to be looked for. The second comment is that all clusters must sooner or later vanish as the temperature is raised and it might be thought that this would happen but little above Tm,, but undoubtedly there are exceptions. For these exceptions, it is of particular interest how quasi- or anti-crystalline clusters would vanish. The order of these structures does not extend throughout a phase, so their ‘melting’ into disorder is not constrained by the phase rule to take place at a single characteristic temperature for each.On the other hand, their order, if not long-range, is co-operative in nature, and its failure and dissolution might be expected to show the general feature of order-disorder transformation-that once the rot sets in, complete collapse is not long delayed. The ‘melting’ of a cluster would then occur over quite a narrow temperature range and give rise to an effect on the properties of the liquid concerned-not expected to be very marked or easily discernible. If, in a given liquid, more than one kind of cluster is formed each will have its own energy-entropy balance, to swing over at its own ‘submerged’melting point. A sequence of second or higher order transitions in the properties of the liquid might then be detectable along the rising temperature scale.This is discussed later (p. 99). THE ABNORMAL HYDRIDES Comparison can now be made between normal and abnormal hydrides. The previously noted similarity between the second and third groups of hydrides in Fig. 1 does not extend to the first group; in particular HF, H2O and NH3 show enormously greater cohesion in condensed phases than would be expected. This is, of course, due to hydrogen bonding. The extent of the abnormalities can be assessed, first by means of the thermal data given in the first three columns of Table 1 and secondly by inventing a ‘normal’ first group of hydrides by extrapolation from the other two. Neon and methane should help in this, but are abnormal-the former because of low Table I : Some thermal data for HF.HzO and N k I 7 1 8 1 9 HF 1 1 9 0 * 1 = g & I 293*1=- I790 -- 226 0.24 0-57 0*140* 503 54 129 184 108 647 311 0-35 0.59 0.148 HzO 5.258 6 - 10 9717 373. I = ~ 26 * 04 79 128 0-62 0.62 0.242 NH3 I436 273. I =-- Qy;4 239*7= __ 558 I 23 * 28 l 9 5 * 4 = 7 * AHB calcd. for evap. to monomolecular vapour = 7.80 kcal mole-1. Units are O K , cal mole-1, cal O K - 1 mole-’. Ives and Lemon 75 mass, the latter because of rotational freedom in the solid. Nevertheless, the rough constancy of TM, TB and TC differences between corresponding members of the third and second groups (23 & 6, 24 5 4,47 3 6", respec- tively) can be used for a linear extrapolation of each which may have a little sigrdicance.It leads to 'abnormality increments', ATM, ATB and ATc, probably underestimated and relatively uncertain to about 10 ", entered in columns 4 to 6 of Table 1. Somewhat unusual 'reduced increments' appear in columns 7 and 8 and column 9 contains ratios of latent heats. It is obvious that water shows the greatest anomaly; very high AHB and high ASB indicate strong cohesion and residual organization in the liquid at TB. High TC and normal TB/Tc (0.577) show that cohesion does not fall away abnormally with further substantial rise of temperature- weak hydrogen bonding could be expected to be more vulnerable. This is all the more striking because water vapour at 1 atm pressure does not depart greatly from ideal gas behaviour.It seems that pair-wise hydrogen bonding between water molecules is weak, but is somehow much strengthened in condensed phases. The behaviour of HF stands in strong contrast to that of H2O. Very low AHB and ASB are due to persistence of association in the saturated vapour which consists of polymeric, zig-zag chains of average length (HF)3.5. Evi- dently, polymeric molecules in the liquid at TB need little excess energy (1.8 kcal mole-1) to slip into the vapour state, but a further 6.0 kcal mole-1 are required to break them up into monomers. This is in line with superior stability of F-H ..... F bonds, but not with considerable stabilization of them in the liquid as compared with the gaseous state. Cohesion in liquid HF could be said to be strong in one direction-along the chains-but weak in others; this is a situation conducive to ordering in one dimension but not in three.The rather high ASM perhaps reflects an appropriate disparity of order between liquid state and solid (an ordered assembly of infinite zig- zag chains). Hydrogen bonds between like atoms decrease in strength in the electro- negativity sequence F > 0 > N. Accordingly, NH3 shows least non-ideality as a gas and abnormality as a liquid, although still relatively high AHB and ASB indicate appreciable hydrogen-bonding stabilization of the liquid state. Liquid ammonia is probably the most normal of the three liquids; this view is supported by their dielectric constants (NH3, 22 at - 34°C; HF, 83.6 at 0°C; HzO, 87-7 at OOC), which broadly reflect the extent of co-operative association.On the other hand, AT, and, especially, ATM/ATc place NH3 differently, suggesting that it is particularly well stabilized in the solid state. Disparity between solid and liquid in this respect is confirmed by high AS,* and AHM/AHB and is supported by relatively low (TB - TIM). In solid NH3, each nitrogen atom has six equivalent neighbours at 3.38 A and six more at the greater distance of 3.95 A. This suggests that the nitrogen atoms form six weak hydrogen bonds-they would have to be weak, but would contribute to cohesion in three dimensions. If this picture is correct, it is a remarkable example of the fortification of weak interactions by * The situation is, however, complicated by the ready inversion of the pyramidal NH3 molecule in liquid and gaseous, but not solid, states.RIC Reviews 76 co-operation depending on precise geometrical requirements-positional, orientational-to be attained in the regularity of a crystal but not conceivably in the irregularity of a liquid. Water comes into comparison again because ASM is relatively low, despite the strong hydrogen bonding in ice and its comparatively uncomplicated structure. Although allowance must be made for positional disorder of hydro- gen atoms in ice (equivalent to orientational disorder of molecules), this is consistent with retention of considerable order after melting, for which there appears to be a reserve of ‘hydrogen-bonding power’. Some inferences may be drawn from this intercomparison of the three abnormal hydrides. Their properties are dominated by hydrogen bonding.The hydrogen bond, normally formed between a hydrogen atom of one molecule and a lone pair of electrons of another, can be regarded as a Brarn- sted acid-base interaction, and it is convenient to refer to one participating molecule as the proton donor and to the other as the proton acceptor. All the three hydride molecules can act in both capacities, but the numbers of protons and lone pairs of electrons available per molecule are, respectively : HF, 1 and 3; HzO, 2 and 2; NH3, 3 and 1. In very many crystalline substances hydrogen bonding is a contributory factor determining the solid state structures that are adopted. In the regularity of the crystal, with participating molecules held in optimum positional and orientational relations, even very weak bonding is co-operatively protected and cumulative in effect.This appears to be so in solid NH3. In liquids, weak bonding has no such privileged position and must sustain the knocks of its hostile, uncoordinated surroundings. This may be the basis of the general rule that, except in the solid state, hydrogen bonding in one-com- ponent systems is confined to equal numbers per molecule of protons donated and accepted. Whereas HF and NH3 molecules form but two hydrogen bonds each under such conditions (acting once as donor and once as acceptor), H20 can give two protons and accept two. It is thus unique in its capability of promoting three-dimensional order, and is the only molecule which, from a single atomic centre, can give rise to four hydrogen bonds directed in space.Whatever the ways in which water may exploit this facility in the liquid state, there can be no doubt that co-operation is involved. Two kinds of co-operative effect must be distinguished. The first is a function of the nature of hydrogen bonding between molecules which can both donate and accept protons. It is co-operative in the sense that a molecule that has acted as donor is more ready than before to act as acceptor, and vice versa. Donation and acceptance are mutually supporting and hydrogen bonding between like molecules is autocatalytic. This is no doubt why, in the association of monohydric alcohols, the dimer stage is very nearly skipped.7 The second effect is the crystal-like kind in which co-operation depends on the satisfaction of precise, three-dimensional geometrical requirements.The problem is, whether hydrogen bonding in liquid water brings this second kind of co-operation into play, so that it is significantly fortified by the formation of evanescent quasi-crystalline or anti-crystalline clusters, or both. If so, does this occur over the whole temperature span between TIM and Tc, Ives and Lemon 77 or only over restricted lower ranges of temperature? On the other hand, is the first kind of co-operation, intrinsic to hydrogen bonding, sufficient, leading to amorphous clusters of no distinguishable symmetry, representing no more than energy and density fluctuations? These are some of the main questions in debate.THEHYDROGENBOND It is mandatory to consider a bond so basic to the subject under review. Strictly, the name ‘bond’ is unsuitable in connexion with the highly variable ‘lone pair interactions’ concerned but ‘hydrogen bond’ is conventionally, if ambiguously, applied to the three atom system X-H *.... Y, where X and Y are covalently bound atoms, rarely other than F, 0 or N, in separate molecules or in the same molecule. In this formulation, the dotted line represents the interaction between a lone pair of electrons, provided by Y , and the hydrogen atom covalently bound to X. In so far as the electrons of this bond are with- drawn towards X (of higher electronegativity), Y may ‘see’ the hydrogen atom as an imperfectly screened proton of radius lO-l3cm-unique as a centre of attraction for its lone pair electrons. It is common to speak of ‘hydrogen bonding between molecules’, i.e.one molecule carrying the group X-H and the other, the atom Y, implying that it is the H - . . . . Y interaction which is being defined as the hydrogen bond. Seemingly logical, this is not acceptable, because the bond X-H is profoundly affected by the interaction. Even to consider X--H--.*Y as a hydrogen bond between X and Y may be satisfactory only as a first approximation because so much depends on what else is attached to X and Y. This is implicit in the reference already made to the co-operative nature of hydrogen bonding.Energetically, the hydrogen bond can be understood only by assessment and algebraic summation of a number of interaction energies, no one term of the summation being of decisively predominant weight. Theoretically based calculation of small, net bond energies (1-6 kcal mole-1) appearing as differences between larger quantities, turns out to be of almost impossible difficulty. This is one reason why the present state of knowledge on hydrogen bonding is regrettably inadequate. Difficulties are enhanced by wide variation in hydrogen-bond geometry. Thus, 0-H ..... 0 bonds vary in length (0 *..-- 0 internuclear distance) from 2.44 to 3.36 A-to be compared with 3.5 A, the normal distance of closest approach between non-bonded atoms.Although colinearity of the three atoms favours maximum stability, some intramolecular hydrogen bonds (e.g. that in salicylic acid) must be ‘bent’, with the hydrogen atom displaced from the line of centres of the terminal (oxygen) atoms. Presumably, then, intermolecular hydrogen bonds can also bend, but how this would affect the bond energy depends on the nature of the bond. The simplest assumption, not inconsistent with known bond energies, is that only electro- static interaction is involved-dipole-dipole attraction between the X-H bond and the ‘atomic dipole’ (lone pair-nuclear charge) of Y. Energy would then bear a simple cosine relation to angle of bend, and considerable de- parture from colinearity would have little effect.If, on the other hand, the bonding were covalent in nature, there would be a much more critical RIC Reviews 78 dependence of energy on angle. It is clearly fundamental to the description of the hydrogen bond to decide which of these views is correct. There is strong evidence that the electrostatic model is inadequate. It does not satisfactorily explain the increase in intensity of infrared absorption due to 0-H vibration which accompanies hydrogen bonding, nor the shortening by nearly 1 A of the H.--*O distance below the sum of the normal van der Waals radii of hydrogen and oxygen atoms. Purely electro- static attraction implies that the attracted species suffer no mutual deforma- tion. At the close approach involved this is impossible-they are bound to polarize each other, and the distortions of the charge clouds must give rise to delocalization of electrons and perhaps change in hybridization.Coordi- nated motions of electrons in atoms brought close together will give rise to dispersion forces, and overlapping of charge clouds will bring repulsion into play. Accordingly, electrostatic, delocalization, dispersion and repulsion energies have been taken into account in attempts to calculate hydrogen-bond energy. Coulsonl6 considered delocalization in terms of contributions from ionic - 4 - - + (e.g. 0 H *.... 0) and covalent (e.g. 0 H-0) valence-bond structures to the complete molecular wave function of the three-atom system and found, in agreement with other calculations, that these contributions are significant.Additional difficulties in estimating dispersion and repulsion terms augment the uncertainty of his assessment of the four main contributions to hydrogen bond energy; these, for the hydrogen bond in ice, are shown in Table 2.l6 I kcal mole-1 I Table 2: Contributions to hydrogen bond energy in ice16 Electrostatic Delocalization Dispersion Repulsion Experimental (from sublimation energy) +6 $8 -8.4 + 3 +8.6 +6*1 It might be remarked that a better result is obtained if the bonding is assumed to be entirely electrostatic, and perhaps this is the reason why this assumption has not been universally discarded. It is unfortunate if trivial coincidence between two energy terms obscures other indications of the theoretical treatment.The most important of these is that decrease of 0 **.** 0 distance is accompanied by increase in delocalization and increase in 0-H bond length. As the 0 ***.. 0 separation shortens towards 2.45 A, the proton moves towards a central position, and probably attains it in a number of cases (e.g. the acid maleate ion). There is no doubt that the [FHFJ- ion is symmetrical, with equal F-H distances of 1.13 (the normal HF bond length is 0.92 A). Electrostatic attraction will not do for such cases. Ives and Lemon 79 6 It is appropriate to interpolate a general comment on hydrogen-bond energy. The heat of formation of the HF, ion from HF and F- is 58 kcal mole-l ; this provides a reason for modifying any impression that the hydrogen bond is always characteristically weak.Another reason comes from an alternative theoretical modell7 which views the hydrogen atom in 0-H ..... 0 as common to two 0-H bonds, one strong and slightly stretched, the other weak and highly stretched. Energies for each are calculated from semi- empirical interatomic potentials and with repulsion and electrostatic attrac- tion terms, provide hydrogen-bond energies as a function of 0 **... 0 distance with considerable success. Particular interest lies in the equivalent treatment of the two halves of the hydrogen bond, and the indication provided that the ‘lone pair attraction-long covalent bond’ half is a good deal stronger than the over-all bond energy might suggest.Examples, in terms of inter- oxygen distance, ‘weak bond’ and total bond energies, quoted in sequence are: 2-70 A, 11.0 and 5.6 kcal mole-1; 2.50 A, 35.5 and 14-4 kcal mole-1. Even if these figures are challenged as derived on too simple a basis, they illustrate a valid point relevant to the vulnerability of hydrogen bonds to thermal disruption. In this connexion, there is evidence that the volume of a system, rather than its temperature, is the decisive independent variable. The theoretical conclusions about delocalization have strong experimental support from the infrared absorption spectra of hydrogen-bonded systems. The increase in intensity of absorption in the 3 p region which accompanies the formation of O-H....*O bonds requires the 0-H vibration to be accompanied by a fluctuating charge separation.This implies considerable and mobile charge migration within the three-atom system in response to the motions of the proton. There is also a frequency displacement (approxi- mately in proportion to the increase in intensity) which reflects the stretching of the 0-H bond-longer bonds have lower frequencies than shorter ones. The third general effect is an increase in the width of absorption bands, less for intra- than for intermolecular hydrogen bonding because in any macro- scopic molecular system there must be considerable statistical variation of 0 ....- 0 separation. Particularly because of the sensitivity of delocalization to this separation, there will be a correspondingly wide variation in the frequency of the 0-H vibration.At this point it might seem justifiable to make the inference that hydrogen bonding is not purely electrostatic; on the contrary, there must be an import- ant covalent contribution to it. It has been strongly argued that the co- operative strengthening of hydrogen bonding in systems of like molecules depends fundamentally on delocalization. If this inference is correct, it follows that hydrogen bonding is the more directional in nature, and that hydrogen bonds will have the greater tendency to be straight because of a considerable dependence of energy on angle of bend. Unfortunately, true as this may be for the ‘ideal’ hydrogen bond, the situation remains confused, and the fact must be faced that there is a wide range of interactions legitimately included under the name of ‘hydrogen bonding’. Proposals have been made for a closer classification-short, straight, long, bent, shading off into attractions which are purely electrostatic.The difficulty is, where to draw an agreed line between interactions which are, or are not, to be considered ‘true’ RIC Reviews 80 hydrogen bonds. Although electrostatic interactions are, in general, somewhat weaker they are not substantially weaker than ‘authentic’ hydrogen bonds, nor are they completely non-directional, because of the localization of elec- tronic charge in hybridized lone-pair orbitals. INFRARED AND RAMAN SPECTRA The expectation that infrared and Raman spectroscopy would be unrivalled in providing information about hydrogen bonds is realized only in the study of dilute solutions of hydrogen-bonding substances in inert solvent-evidence for the co-operative nature of the bonding has been gained in this way.18 For pure, hydrogen-bonded liquids such as water the situation is less favour- able because of greatly enhanced difficulties.Coupling of molecular motions, overlap of overtone, combination and resonance bands, added to the broaden- ing caused by the structural disorder of the liquid state, make it hard to identify peak frequencies and to assign them to specific modes of motion. The literature is correspondingly large and contentious ; the present writers have no alternative to arbitrary selection of topics with the hope of presenting an unbiased view of the present situation.Table 3 lists one selection of principal vibration frequencies, from Raman and infrared spectra, for HzO in the three states of aggregation. v2 Bend v1 Symmetrical stretching v3 Asymmetrical stretching I Ref 19 3755 3397 3434 I595 I642 164’ Table 3: Spectroscopic frequencies for H2O (cm-1) Water vapour Liquid water (3°C) Liquid water (70°C) Ice Ref 20 3650 3448 3448 3360 - The vapour assignments are not in doubt. Raman intensities, v1 9 v3; i.r. intensities, vl < v3; frequencies, v1 < v3. But in ice, the main Raman band is at 3360 cm-1, and is expected t o be vl; the main i.r. band is at 3210 cm-1 and is expected t o be v3. This puts the frequencies in the reverse order, v1 > v3.Hornig et a/. adopted this, but other workers seem either not t o have noticed the transposition, o r t o have chosen to disregard the intensity evidence, preferring t o assume that the frequencies for ice must remain in the order V I < v3. This problem i s not properly resolved-there are arguments either way, and, of course, the uncertainty extends to liquid water. The uncertainty is worsened by the asymmetry of local fields of force in the liquid; the effective symmetry of a water molecule is lowered from tetrahedral (CZ,), so that v 1 and v3 are no longer purely symmetrical and antisymmetrical stretching vibrations-they lie close to each other and to the overtone 2 ~ 2 . Assignments remain correspondingly questionable.Ives and Lemon 81 Bending frequency is little affected by phase change, but the stretching frequencies decrease in the sequence vapour > liquid > solid, with liquid water about two-thirds the way to ice. Increase of temperature has but little effect on the vibrational frequencies of liquid water and hence, it would be thought, little hydrogen-bond breaking effect Wall and Hornig21 have made precise photoelectric Raman spectroscopic studies of HDO using five mole per cent solutions of D2O in H2O and of H2O in D2O. This method of applying isotopic substitution has the advantage that the five nearest-neighbour positions of each OH oscillator are pre- dominantly deuterated, or of each OD oscillator, protonated, so that un- coupled fundamental stretching frequencies can be observed. The simplified, uncoupled bands are still broad (although very narrow in the case of deuter- ated ice), and it is asserted that their breadth directly reflects the spread of 0 *.*..0 distances in liquid water. Using a well-substantiated correlation between this distance and frequency, Wall and Hornig derived a distribution function for nearest-neighbour distances in good agreement with that from the best x-ray scattering measurements. This supported their assertion, but perhaps more important was the fact that the Raman bands and the distri- bution curves were smooth, continuous functions, indicating that allowable values for intermolecular distances in water are densely distributed within a finite range, with one most probable distance represented by a single maximum not far from the centre of the range.Frequency shift with temperature was small, confirming the earlier indication that hydrogen bond strength varies little with temperature. It is clear that these results must be taken into account in discussion of the nature of any clusters that may be postulated in liquid water. The conclusions of Wall and Hornig have been supported by Falk and Ford’s studies21a of the infrared absorption of dilute solutions of HDO in H2O and D2O between 0 and 130°C. This method has an advantage over Raman spectroscopy in greater resolution, better ‘signal to noise’ ratio, and consequent improved detectability of small shoulders on band profiles.Well separated v1 (2505 cm-l) and v3 (3400 cm-1) bands were identified, of broad, smooth, singly-peaked, nearly Gaussian contour, decreasing in intensity with rising temperature in a manner consistent with a gradual weakening of hydrogen bonds. High resolution examination of the region (3550-3750 cm-l), where absorption due to non-hydrogen-bonded OH would be expected, showed no trace even of a shoulder-and similarly for OD. This, with the parallel Raman result, was taken as conclusive evidence that there is in water a continuous distribution of hydrogen-bond strengths from weak to as strong as in ice, and that the existence of discrete species differing in the extent of hydrogen bonding is ruled out. No less weight, however, must be given to Walrafen’s equally careful broad 152-175 cm-l was assigned to hydrogen-bond stretching, A in studies22 Raman band centred of intermolecular at - 60 cm-1 vibrations to hydrogen-bond water.bending. band and These of a frequency weaker, bands decreased in intensity quite rapidly with rising temperature, and also with increasing concentration of added electrolyte. Both these effects strongly suggest extensive breaking of hydrogen bonds, and Walrafen was able to support this view with plausible thermodynamic argument. RIC Reviews a2 More recently,22a Walrafen has substantially supplemented the evidence for his viewpoint in two main ways. First, he has shown that the broad Raman band between 300 and 1000 cm-1, associated with intermolecular libration in liquid water, can be resolved into three Gaussian components which vary little in frequency, half-width or relative proportion between 0 and 90°C.All three have essentially the same negative temperature coefficient of intensity as the previously studied hydrogen-bond stretching and bending vibrations. ‘411 these frequencies (717, 538, 439, 152-175 and 60 cm-I), which are associated with water molecules in a tetrahedral, hydrogen-bonded environment of CzV symmetry, fade out together. It is relevant to later discussion of water structure theories (e.g. that of NCmethy and Scheraga) that if this fading out is attributed to the breaking of hydrogen bonds, then the constancy in relative proportions of the librational components indicates that each molecule breaks its bonds in one single step, and not in consecutive steps.This implies that it is not acceptable to assume different energy levels for water molecules forming one, two, three or four hydrogen bonds. The second way in which Walrafen has supported his case is by re-examina- tion of the intramolecular Raman bands of liquid water in the 2800-3900 cm-1 region, from 10 to 90°C. These, by analogue computer technique, have been resolved into four Gaussian components. Two, at 3247 and 3435 cm-1, are strong, decrease in intensity with rising temperature and are not found for water near Tc; two, at 3535 and 3622 cm-l, are weak, but increase in intensity with rising temperature and are found for water near TC (with but slight change in frequency).There is an intermediate frequency, 3460 cm-l, of intensity independent of temperature ; this is equivalent to an isosbestic point. This behaviour is judged by Walrafen to be irrefutable evidence of an equilibrium, shifting with temperature, between hydrogen-bonded (‘lattice’) water and non-hydrogen-bonded water. It is clear that there is disagreement in the interpretation of observations of fundamental intramolecular and intermolecular frequencies, perhaps to be resolved by the suggestion that alternative methods of study are not ‘seeing’ the same phenomenon. Walrafen himself points out that the inconsistency can be clarified only by better knowledge of so-called hydrogen bonding, and that his ‘non-hydrogen-bonded’ water molecules (with spectra unlike those of high pressure steam or of dilute solutions of water in inert solvents) are still subject to strong intermolecular forces.The situation emphasizes the quandary mentioned in the last section. It is perhaps that intermolecular Raman effects are associated with straight ‘covalent’ hydrogen bonds, and that it is the disappearance of tetrahedral symmetry of vibrating groups with rising temperature that is being observed.22b If this is so, we need some hydrogen bonds other than the straight ‘covalent’ ones, hardly less effective in promoting liquid phase cohesion and in their influence on intramolecular vibrations. Alternatively, we require some other kind of bonding altogether.A further spectral region of interest to the study of water extends from 5000 to 11 000 cni-1 in the near infrared, where there are several systems of well-marked overtone and combination bands of not universally agreed assignment. Quite different interpretations, each with important implications in relation to the structure of water, have been given to essentially identical Ives and Lemon 83 O 8t 0 31- 0.2 f I 01 - L t L U I I 2 0 I 2 5 I 3 0 115 I 05 I10 I 30 Fig. 3. Near infrared absorption of water, after Luck. E = optical density, p = density, d = optical path length. observations-notably those of Buijs and Choppin23 and of Some of Luck's spectra, which extend over the wider temperature range (> TC - TM) are reproduced in Fig.3. Buijs and Choppin assigned frequencies of 8620 cm-1 (1.16 p), 8330 cm-1 (1.20 p) and 8000 cm-1 (1.25 p) to water molecules forming zero, one and two hydrogen bonds respectively. From appropriate intensity measurements, they calculated the mole fractions of these species to be (in the same order) 0.27, 0.42 and 0.31 at 6"C, and 0.40, 0.42 and 0-18 at 72"C, these being the extremes of their experimental temperature range. Extrapolation gave 46 per cent hydrogen-bond breakage accompanying the melting of ice; further extensive breaking with rise of temperature from 0°C was clearly inferred. Luck contended that Buijs and Choppin's solution of their simultaneous equations for the mole fractions of three water species was not unique, and that, in any case, non-hydrogen-bonded molecules in liquid water cannot be detected readily by infrared absorption. Although there are nine 'hydro- gen-bonding states' of a water molecule,* it is only proton donation, not acceptance, that has an appreciable effect on vibration frequency, largely confined to that of the OH group acting as donor, and almost independent of the state of the other OH group in the same molecule. The appropriate assignment of the peak near 1.15 p (Fig.3) to free OH groups was supported by the cfose comparison afforded by the similar band system of methanol in the same region. The large proportional increase in intensity between TB and TC is to be noted-it suggests that this is the temperature range for most of the hydrogen-bond breaking.One independently determined per- *The molecule can act zero, one or two times as proton donor and, independently, zero, one or two times as proton acceptor. 84 RIC Reviews centage of free OH groups at one temperature was needed to calibrate the intensity scale. Studies of the dielectric constant of water up to Tc25 had led to an estimate of nine per cent at 0°C; adoption of this figure brought the dielectric constant and infrared data into agreement over the whole tempera- ture range. Twenty per cent of non-hydrogen-bonded OH groups at 100°C provides a very different picture of water from that presented by Buijs and Choppin, but more consistent with the high TC of water. For hydrogen-bonded OH groups, Luck assigned absorption at 1.25 p to those making linear (ice-like) bonds of maximum strength, and at 1.20 p to those involved in ‘unfavourable’ bent or long hydrogen bonds, considered collectively.This rough classification was based first on observations by the ‘matrix isolation method’ of the infrared absorption of low, varied concen- trations of H2O in solid nitrogen at 20”K26, which gave the fundamental frequencies of monomers, di-, tri- . . . n-mers. These showed the expected trend (e.g. in increasing order of complexity, higher v only: 3725, 3545, 3510, 3390 . . . 3355 cm-1-cf Table 3 ) and further indicated that the dimer, under these conditions, was cyclic (Fig. 4)-two weak, bent bonds doing the work of one straight bond. H \ 0-H ‘ \ , , ‘H -0 H \ Fig.4. Water dimer in solid NZ at 75°K. By assuming that all the lower polymers were cyclic, Luck obtained a relation between frequency and angle of bend. Secondly, he studied the hydrogen- bonding equilibria of numerous organic compounds in inert solvents, relating by stereochemical argument bond energy with angle of bend ; use of the Badger- Bauer rule on the dependence of frequency on energy then gave results agreeing with his other frequency-angle relationship. Luck completed his detailed assignments with 1.143 ,u to OH vibration in free HzO molecules, and 1.150 p to free OH groups sharing a molecule with a hydrogen-bonded OH group. These assignments, and intensities, led by a Buijs and Choppin type of calculation to the results illustrated in Fig.5 -stressed to be a possible solution, but not the only one. Substantial evidence for the persistence of molecular association well above TC will be mentioned later. If these calculations are considered to be overly conjectural, they are hardly more so than any others. The extent of disagreement on matters fundamental to a theory of water is well illustrated by running down a listZ1a of estimates by 19 authorities of the percentage of hydrogen bonds broken whenice melts: 71.5,66, 60, 57.5, 56,47,46,46, 38, 32, 30, 18, 18, 15, 10-5,9,9, 2.5 and 0. On the whole, opinion seems to be moving away from wholesale breaking of hydrogen bonds and considerable concentrations of ‘monomers’ Ives and Lemon 85 Temperature ("C) Fig. 5.Possible concentrations of named entities in water, after in liquid water. Vacuum ultraviolet spectroscopy of water, vapour and liquid, has led to an estimate of less than one per cent monomers up to 100°C.27 Recent studies by Glew and his collaborators27a of the fundamental valence- stretching frequencies in infrared absorption shown by water in dilute solution in a wide range of solvents demonstrate that the monomer bands (unequivocally identifiable with the vapour or nitrogen matrix bands) are to be found for solutions in one group of solvents including halogenated hydrocarbons and some alcohols. In a second group-of proton-accepting solvents (e.g. ethylene oxide, acetone, tetrahydrofuran)-the spectra show the presence of a single type of dissolved water species hydrogen bonded to the solvent, with no sign of the presence of non-hydrogen-bonded monomers.Extrapolation strongly suggests that for water in water (the best proton acceptor) there can be no significant concentration of non-bonded monomers. THEORIES OF THE STRUCTURE OF WATER It is because there is no generally agreed theory of the structure of water that the plan stated in the Introduction was adopted, and the main concern has been with factual terms of reference (e.g. that water is a liquid) which ultimately must be satisfied, and with some difficulties of inadequate know- ledge, and of interpretation of experiments, that continue to vex the theory makers. In this section, further evidence and theories are non-exhaustively presented in whatever order best suits a continuing narrative.It is suitable first to ask, what is to be expected of a theory? Two of the oldest (Rowland, 1880; Rontgen, 1892) can be combined and summarized in a statement that 'water is a mixed solution of ice and steam that varies in proportion with temperature'-very close to some modern views, and perhaps qualitatively (or better) satisfactory. Eucken's theory2* that there is in water an equilibrium between monomers, dimers, tetramers and octomers (the latter necessary for Hohlraum) gave an excellent account of many proper- ties of water. It had three equilibrium constants available for adjustment. Some recent theories (mostly as well equipped with adjustable parameters) RIC Reviews 86 satisfy more exacting thermodynamic tests, but such tests are not discriminat- ing and can be inadequate to prove a theory wrong-particularly a structural theory. More is now required than success in calculating smoothed-out macroscopic properties, because of the virtual certainty that the behaviour of water to solutes and surfaces and in biological systems presents a whole range of microscopic, structurally-based problems.It is in this direction that progress is needed but is most difficult-inordinately so on the theoretical front because a realistic model is unlikely to be simple or mathematically tractable. More structure-sensitive experiments are needed, as well as recon- sideration of things we are reasonably sure about.The narrative begins with the latter need in mind. ...... . . . . . . Fig. 6. (a) The water molecule; (b) primary unit of water association. The water molecule The structure of the water molecule (Fig. 6a) lies at the root of the properties and behaviour of water. Two protons and two lone pairs of electrons occupy approximately tetrahedrally-disposed, hybridized orbitals. The vacuum dipole moment, 1-87D, is not large, and can hardly be relevant to close-range interactions ; it results from opposed components-that due to the unscreened lone pairs and the oxygen nuclear charge has been estimated as 3~38D.~9 Then, if dipole-dipole association between water molecules should be in- voked, the question arises, what dipoles? Also, would such association lead to something distinguishable from what is generally included under the heading of hydrogen bonding? This bonding leads to the formation of a five-molecule, approximately tetrahedral, primary structural unit, somewhat crudely represented by Fig.6b. It can be seen that the net moments of the molecules, lying symmetrically between the OH bonds, neither fully support nor fully oppose each other. The tetrahedral unit is not consistent with a simple dipolar model for the water molecule, which will, accordingly not be suitable for consideration of the behaviour of water in condensed systems. Lone pair hybrids Ice Tetrahedral association of water molecules and sp3 hybridization of oxygen orbitals are perfected in ordinary ice-ice Ih-the structure of which is represented in Fig. 7, where hydrogen-bond distances have been stretched for clarity.Each oxygen atom is surrounded by four others at a distance of 2.75& three in the same puckered layer of six-membered rings, and the Ives and Lemon 87 n n n Fig. 7. Structure of ice I, Ih, plan and elevation of model. (Reprinted with permission. Copyright 0 1956 by Scientific American, Inc. All rights reserved.) fourth in an adjacent layer. Adjacent layers form between them rather large, polyhedral cavities, each with 12 vertices, which link up to form channels threading their way through the hexagonal rings. The structure is that of a form of silica-tridymite-formed by the sharing of oxygen atoms between SiO4 tetrahedra.Apart from normal crystal dislocations and defects, this representation is ideaIized. The protons are disordered. This can be understood by consider- ing that, radiating from the central oxygen atom of a five-molecule, tetra- hedral assembly (see Fig. 6b) there are four 0 . . - - * 0 directions, each accom- modating one proton; two of the four protons must be near the central oxygen, being covalently bound to it. Since there are six ways of taking two out of four things, there are six possible arrangements of the protons, giving six alternative orientations to the central water molecule. But each of these 88 RIC Reviews O C Fig. 8. Phase diagrams of water-substance. alternatives requires two of the peripheral molecules to have favourable, as opposed to equally likely unfavourable orientations.This divides the absolute probability of any of the six arrangements of protons by four. The total number of configurations in an ice-like assembly of N molecules is, accord- ingly, (6/4)N. This statistical disorder of protons, frozen-in at low tempera- tures, confers on ice a residual entropy at 0°K of S,O = R In 3/2 = 0.81 cal "K-1 mole-1, which is precisely the value required to bring 'third law' and spectroscopic entropies of water vapour into coincidence, and has been otherwise confirmed. The electrical properties of ice show that the protons are mobile at higher temperatures. The dielectric constant of ice, although small (-4) at very low temperatures, is greater than that of water at 0°C (91.2, 88*2), and proton mobility (hydrogen ion equivalent conductance) is much greater in ice than in liquid water.These facts require the presence in ice of two kinds of intrinsic defect-ionic and orientational. The former arise from proton transfer. HzO + HzO = H3Of + OH- The latter arise by rotation of a water molecule through 120" about an ..... 0 axis, whereby empty (leer) and doubly-occupied (doppelt) 0-H oxygen-oxygen 'bonds' are generated (i.e. 0 *.... 0 and 0-H ....- H-0). These are called L- and D-Bjerrum defects30 and, by rotation of adjacent water molecules, they can become separated and migrate through the ice crystal. Ives and Lemon On the grounds that one cannot afford to ignore any aspect of the be- haviour of water-substance in seeking guidance on the problem of liquid water, attention should be given to the other forms of ice.Figure 8 presents recent phase diagrams.31 Ices Ic and IV are metastable phases occupying areas in the stability regions of ices Ih and V respectively. There is a recently 89 Fig. 9. Radial distribution curves for water, after Morgan and Warren.35 reported ice V I I P of the same structure as ice VII, but orientationally immobile, and probably ‘proton-ordered’ (as is ice 11). It is of interest that ordinary water-substance is crystalline at 400°C under a pressure of 200 kbar, and also that ice VII, which can exist metastably at - 175°C and 1 atm pressure, has a density of 1.50 g cm-3, still substantially less than would be expected for close-packed water molecules (- 1.84 g cm-3).Progress has been made in determining the structures of these ices, notably by Kamb.33 In ice VII, each water molecule has eight nearest neighbours, but is bonded to only four of them-it probably consists of two interpenetrating lattices of cubic ice Ic (which has a very similar open lattice of that of Ih), each lattice being fully hydrogen bonded within itself. Ice VI also consists of two interpenetrating but not interconnected lattices. These two then can be described as ‘self-clathrates’, and this appears to be the way in which relatively high density is achieved by tetrahedrally-linked structures. It seems that all RIC Reviews 90 the ices retain substantially tetrahedral disposition of water molecules-if in some cases with considerable distortion, i.e.bending of hydrogen bonds. Conclusions might be that water molecules persist in preferring four- fold bonded coordination, and that the hydrogen bond is highly adaptable to extreme conditions. A report34 that water-substance at 1000°C and 98 kbar has a conductance of 0.7 ohm-1 cm-1 suggests an ionic bond. x-Ray scattering When ice melts, there is a contraction, but this abnormality is less remarkable than the supplementary contraction which accompanies the warming of the melted ice towards 4°C. This is very like the continuation of a melting process, and a negative coefficient of thermal expansion needs must be based on structural change. There is a compulsion to admit that, at least in this temperature range, liquid water must have a structure.The retention of structure in water is supported by radial distribution curves derived from x-ray scattering. The curves obtained by Morgan and Warren35 are reproduced in Fig. 9. They established that the preferred first and second nearest neighbour distances are 2.9 A and 4.5-4.9 A-somewhat greater than those in ice. Average coordination increased from 4.4 at 195°C to 4.9 at 83°C-to be compared with four in ice, and a number decreasing from about 11 with rising temperature for a normal liquid. There cannot be much doubt about the ice-likeness of water. If water were a normal liquid, its density would be in the region of 1.84 g CITL-~, so that x-rays would hardly be needed to prove its open structure, but, in a sense, clarify the questions to be answered.Since x-rays indicate a structural expansion accompanying the melting of ice, why is water denser than ice? Why does rising temperature seem to close up the molecular packing, rather than loosen it ? How is it that forces of intermolecular attraction succeed in holding molecules apart? I The Bernal and Fowler theory On the basis of earlier x-ray scattering studies, and of wider consideration of the properties of water, Bernal and Fowler in 193336 proposed the first, and now very well known structural theory of water. It involved preferred molecular arrangements in liquid water simulating three structures. These were described as ‘water I, tridymite-ice-like, rather rare, present to a certain degree at low temperatures below 4°C ; water 11, quartz-like, predominating at ordinary temperatures ; water 111, close-packed ideal liquid, ammonia- like, predominating at high temperatures for some distance below the critical point at 374°C.These forms pass continuously into each other with change of temperature. Throughout . . . . there is no question of a mixture of volumes with different structures: at all temperatures the liquid is homo- geneous but the average mutual arrangements of the molecules resemble water I, I1 and 111 in more or less degree.’ This passage has been quoted verbatim because the theory has been misinterpreted-it has always been a ‘uniformist theory’, never postulating the existence of water molecules in alternative, sharply distinguishable situations, nor any long range order.Ives and Lemon 91 that of quartz, - 1.08 g cm-3, although adequately explaining the maximum scope. arrangements There Thus, was if a which the difficulty density water in of molecules that tridymite the use is tend adjusted of to forms assume to of that silica did of ice, not as models - give 0.91 adequate g for cm-3, the densityphenomenon in terms of the shift water I + water TI with rising tempera- ture, was too high. It would also lead to a second nearest-neighbour distance discovery obviated The less than that of a new observed form by of x-rays. silica, keatite difficulty (density has on the been same scale - by 1.01 the g cm-3) incorporating five-membered, non-planar rings.Models of dis- ordered keatite give a calculated radial distribution function to a first approxi- mation the same as that of water. More recently,12 Bernal has generalized the theory, seeing the preferred molecular arrangements in water as simulating a network of linked four-coordinated molecules, forming rings of four, five, six, seven, or even more, molecules (perhaps with five preferred) arranged in sets of random order. The distorted bond model This model is the basis of the uniformist theory par excellence. It was proposed by Pople37 that the key to the structure of water is that energy can be absorbed, entropy increased, and long-range order destroyed by the bending of hydrogen bonds.The melting of ice removes the restriction of synchronization of these bending motions, so that, in water, the increased amplitude and randomness of bond-bending accounts for the greater heat capacity of water than ice (about double) and by bringing more molecules into the first and second nearest-neighbour zones, for the greater density and average coordination. No hydrogen-bond breaking is envisaged. It is proposed that there is in water ‘a network of bonds extending throughout the whole liquid, which is, in a sense, one large molecule.’ Radial distribution curves in agreement with those for x-ray scattering are calculable on the basis of this theory which has also given a good, quantitative account of the dielectric constant of water over a range of temperatures.38 It is supported by Wall and Hornig21 and by Falk and Fordzla on grounds already explained.Interstitial models Morgan and Warren35 extended their analysis of the radial distribution curves derived from x-ray scattering by comparing the curve of water at 1-5 “C with that to be expected from a statistically randomized, ‘softened-up’ ice lattice, with the result shown in Fig. 10. In this diagram A is the experi- mental distribution curve, B is the curve for softened-up ice, and C is the difference, A - B. It shows an excess density of water molecules at a distance of 3.5 A from the arbitrary reference molecule at the origin. Such a density could be provided by water molecules, dislodged from ice-like lattice sites, occupying the centres of the large cavities in the ice Ih structure. The idea that ice ‘melts internally’, self-stabilizing residual structure, well explains the anomaly of structural expansion accompanied by over-all contraction, and also provides a means of understanding the maximum density pheno- menon.RIC Reviews 92 I 0 20 3 0 4 0 5 0 6 0 7 0 r [HI- Fig. 10. Calculated and experimental radial distribution curves for water at I - 5 " C after Morgan and Warren.35 A = experimental, B = 'softened-up' ice, C = A - B. This was the basis of the first interstitial model of liquid water proposed by Samoilov.39 F ~ r s l i n d , ~ ~ considering the progressive generation of defects in the structure of ice by rise of temperature, arrived at a similar model.Later, the refined x-ray studies of Danford and Levy41 were shown to be consistent with the existence in liquid water of an expanded tridymite-like lattice, with partial occupation of the cavities by non-lattice molecules. Similar work by the same has recently indicated that water retains average coordination of 4.4 to 4-5 and changes but little in mean nearest-neighbour distance (2.82-2.94 A) between 4" and 200°C. The tridymite lattice arrange- ment that fits the radial. distribution curves is, however, anisotropically expanded, and the non-lattice molecules are not at the centres of the cavities- instead of six nearest lattice neighbours, they have three. This is a puzzling result, but it is not advanced as an unique solution to the distribution curves, which cannot decisively discriminate between alternative models.Not dissimilar proposals came from two aspects of the behaviour of water with non-polar solutes, including the noble gases and hydrocarbons. First, the formation of 'gas hydrates'-crystalline hydrates, such as C3Hs,17HzO, by substances noted for their lack of affinity for water. The structures of many of these hydrates are now well e~tablished~~; they are clathrates, with guest molecules held in the cavities of hydrogen bonded cages. The simplest such cage is the pentagonal dodecahedron of 20 water molecules illustrated in Fig. 11, where it is represented as enclosing a methyl halide m0lecule.~3 This unit cannot grow further in three dimensions because of five-fold sym- metry, but it has hydrogen-bonding sites left over for spatial linking of permissible symmetry with other such units.Secondly, non-polar solutes, although very sparingly soluble, dissolve in water with surprisingly great loss of enthalpy and entropy (e.g. CH4; AH" = 3.19 kcal mole-1; ASo = - 31-8 cal OK-1mole-1 at 25"C), only explicable in terms of their effect on the solvent, and water is unique as a solvent in this respect. The suggestion by Frank and Evans44 that each solute particle becomes embedded in an 'iceberg' has become celebrated, but is now 93 Ives and Lemon Fig. I I Pentagonal dodecahedra1 water cage. less graphically described in terms of structure-promotion in the ambience of the non-polar solute particle, with no implication that it is a normal ice- structure being promoted.There is no doubt about: the generality of this effect, now referred to as ‘hydrophobic hydration’. Since it is particularly typical of the gas-hydrate formers, the view is held45 that the structure promoted is that of the appropriate hydrate-but not in the sense that crystalline solute entities exist in solution. It was Pauling46 who proposed the ‘water hydrate’ theory-that water is, in effect, a clathrate hydrate of itself. The structure envisaged is shown in Fig. 12. Despite a certain advantage in maximizing hydrogen-bond-forming possibilities, the general view seems to be that this model involves more long-range order than is appropriate for a liquid.It is said not to be consistent with x-ray scattering mea~urements.41~ Interstitial models, generalized as ‘any quasi-crystalline framework with single molecules occupying interstitial sites and making no contribution to the total volume’ were shown by Frank and Q u i ~ t ~ ~ to be capable of giving a good account of the thermodynamic properties of water-except perhaps for the heat capacity, but to call for free rotation of the interstitial monomers, as if they were situated in a field-free region. Cluster models Frank and Wen’s ‘suggested picture of water stru~ture’~83 49 laid emphasis on the dynamic nature of any possible structure in water, and on the co- operative nature of hydrogen bonding.To quote49: ‘Liquid water is thus pictured as consisting of flickering clusters of bonded molecules mixed with and alternating roles with non-bonded fluid which encloses them, and con- RIC Reviews 94 Fig. 12. Pauling’s water structure model. stitutes the rest of the sample. A (larger or smaller) cluster is pictured as leaping to attention, so to speak, when the stage is set by an energy fluctuation which creates a suitably “cold” region, and relaxing “at ease” when the necessary energy of “melting” becomes available.’ Emphasis is placed on a closer definition of the hydrogen bonding con- cerned in cluster formation : that there is considerable mutual polarization of bonded molecules, resulting in charge separation and rehybridization ; a considerable non-linear dependence of such charge separation on 0-H and 0 a * .- . 0 distances ; a considerable mutual neutralization of net effective charge between each proton and lone pair involved, so that in a fully-bonded system-ice or a cluster-these charges will be largely suppressed, their interaction with other charges or dipoles virtually disappearing. This provides a basis for the integrity, or distinguishability of a cluster, and leaves open possibilities for interactions responsible for cohesion of the denser fluid (hardly to be regarded as a normal ‘monomer liquid’) surrounding it. Voluminous clusters and denser fluid of higher energy (perhaps not sub- stantially higher) provide the means to interpret the maximum density and the high heat capacity.Hydrophobic hydration appears in a new guise. Non-polar solute molecules, lacking strong external fields of force, are relatively incapable of transmitting disruptive influences. Replacing the higher energy more-disordered water molecules from which clusters normally receive their ‘heat of melting’, the inert molecules increase the half-life of clusters bordering on them, and add to the statistical ice-likeness of the system. This theory of water is based on a ‘mixture model’, which, like the inter- stitial or clathrate models, is also ‘bistructural’ in that it envisages water Ivzs and Lemon 95 7 Fig. 13. Nemethy and Scheraga’s water structure model. molecules in two main roles or situations, no doubt with rapid exchange between them.There is a sharp contrast with ‘uniformist’ theories. Evidence is available however, to support two states for water molecules in liquid water. A mixture, or ‘two-fluid7, model for water was found to be necessary by Hall50 to explain its excessive absorption of sound (which does, however, fall away quite rapidly with rising temperature) in terms of a shift of equili- brium between bulky and dense states imposed by ultrasonic compression waves. Inelastic scattering of cold neutrons51 has also indicated that water can support lattice-like vibrations, and also, with less certainty, that it contains non-translating (presumably trapped) monomers. Clusters have to be viewed on an appropriate time-scale. They are not, in the ordinary sense, polymeric entities in equilibrium with each other and with monomers, since relaxation methods do not find them.Water (unlike alcohols) has a single dielectric relaxation-time of about 10-11 sec52-not a whole family of relaxation times as might be expected-and a common activation energy, 4.6 kcal mole-1, for dielectric relaxation, self-diffusion, shear viscosity and bulk viscosity. The implication is that the same funda- mental process is concerned, and this, in the theory, is regarded as the forma- tion and dissolution of clusters. This assigns to the clusters a half-life of the order of 10-11sec, about 100 to 1000 times the period of a molecular vibration and considered adequate to confer meaningful existence to the inside and clusters.NCmethy a cluster Scheraga53 (containing have - 25 adopted molecules this model, within requiring the temperature the molecules range RIC Reviews 96 0-70°C) to be fully hydrogen bonded, and excluding the cavities within a cluster from occupation by water or non-polar solute molecules. The surfaces of the clusters contain water molecules forming three or two bonds, or one bond, and the clusters are separated by water molecules having no hydrogen bonds. This model is illustrated in Fig. 13. Water molecules were ascribed equally spaced energy states* according to the number of hydrogen bonds in which they participate: four, three, two, one or zero. Given suitable assumptions, a partition function set up for this five-state model allowed thermodynamic properties of water to be calculated by the ‘significant structures’ method with considerable success. Vand and Senior54 have modified this model, replacing the inter-cluster monomeric fluid by an equilibrium mixture containing dimers and straight or branched polymers not forming ring- or cage-systems.They also replaced the concept of discrete energy levels by energy bands, centred on the three levels assumed by Buijs and Choppin,23 and were able to calculate thermodynamics functions from 0-100°C, assuming those at 50°C, with remarkable accuracy. The dense, non-hydrogen-bonded fluid, required by cluster theories to act as a flotation medium for the ice-like clusters, presents a difficulty which has clearly caused discomfort.The molecules in this fluid cannot seriously be considered as ‘free monomers’- they must be in some strongly interacting state to hold the whole system together in a manner consistent with the high internal pressure (20 000 atm) and strong cohesion of water. Are dipole-dipole and dispersion forces, as suggested by Nkmethy and Scheraga, adequate ? Comments previously made on the basis of Fig. 1 show that this assumption is debatable. It has however been maintained55 that we misunderstand the strength of such forces, and evidence has been advanced that something other than hydrogen bonding can provide a high cohesive pressure between water molecules. For instance, DzO, normally less volatile and with greater ‘hydrogen bondedness’ than water, becomes more volatile than water above 220°C and has a lower critical temperature.This is thought to mean that the molecules have more than one way of pulling each other together, and that at low temperatures one of them preponderates, and at high temperature another. This may be so, but it is not comfortable to suggest that there is another area of ignorance which lies between us and the problem to be solved. A two-state model not involving non-hydrogen-bonded water has been proposed by Davis and Litovitz.56 Briefly, it invokes the maximal strength of linear ‘covalent’ hydrogen bonds, which, because of the bonding geometry of the water molecule, will strongly favour the formation of closed, six- membered, hydrogen-bonded rings. Ring closure would strengthen all the bonds by co-operative effect.These rings would be identical with those in the puckered layers of normal ice, so that the first step in the formation of an ice-like cluster would be accomplished. But, in liquid water, rings could conceivably come together in an alternative way providing another, less deep, potential minimum. In normal ice (Ih), adjacent layers are in mirror- image relation. Consider a hexagonal, non-planar ring in one layer, situated vertically above the corresponding ring in the layer beneath. Three oxygen * This equal spacing, is hardly consistent with the co-operative nature of hydrogen bonding, on which the formation of ‘flickering clusters’ would seem to depend. Ives and Lemon 97 7§ Fig.14. Luck’s cluster model for water just above 0°C. atoms in one ring are hydrogen bonded to their ‘image’ oxygen atoms in the other, completing delineation of a cavity. If these three inter-ring bonds break, the upper ring is free to rotate in its own plane, and a twist of 60” will bring it into ‘stacking chair relation’ with the lower ring. Collectively, such rearrangements lead to a nearly close-packed body-centred cubic structure in which each oxygen has two hydrogen-bonded nearest neigh bours and five non-bonded nearest neighbours with which it can interact otherwise. This model can be reconciled with the radial distribution curves and gives a good account of thermal expansion and heat capacity. Luck’s cluster model5’ also excludes ‘free’ water molecules to an extent consistent with Fig.5. An ice-like cluster form is adopted, and on the basis of the co-operative nature of hydrogen bonding, it is assumed that broken bonds and free OH groups will not be randomly distributed, but will be situated on fissure surfaces, or ‘defect planes’ such as Frenke158 has proposed to be characteristic of liquids. This provides a snapshot model of the kind shown in Fig. 1G‘flickering’ with a period of sec, with closing and opening of the fissure surfaces, where the excess density is located. Account is taken of ‘unfavourable’ hydrogen bonds, and the liquid is thought of as possessing a range of different states in accordance with the Vand and Senior picture.54 RIC Reviews This theory, based on an interpretation of spectroscopic observations, gives, in terms of calculations which the author describes as ‘needing re- finement’, an impressive account of the behaviour of water.Very interesting observations in the supercritical region lend support to the postulation of such extensive clustering. In effect, liquid persists, and, without a meniscus, occupies the lower part of the infrared cell at temperatures 30-50” above 98 Tc, confirming many older observations, and suggesting that the critical phenomenon is a matter for the surface chemists. The latent heat of evapora- tion of the invisible, supercritical liquid appears as a set of maxima in heat capacity which allows the normal vapour pressure curve to be smoothly extended on P(T) diagram up to T/Tc ri 1.7.F r a n ~ k , ~ ~ studying supercritical solutions of water in argon (allowing density to be widely adjusted as an independent variable), has used excess proton mobility (proton jumps along a hydrogen-bonded path) in diagnosis of cluster formation. Considerable molecular aggregation has been found on this basis at 4OO0C, at densities greater than 0.5 g cm-3. It is impracticable to follow the changes which continue to be rung on the not very clear classification of water-structure models, each supported by its own evidence, and each encountering its own difficulties. Pending new and decisive diagnostic experiments, the present need seems to be a widening of the contexts of thought and investigation-first (as has been attempted) to consolidate bases of comparison and terms of reference that are reasonably sure; secondly by studying ‘tangential’ problems-flank attacks.There is, indeed, intense activity in the latter kind of study, particularly in relation to systems of which water is one component. This is tantamount to using solutes (ionic or molecular) and surfaces as probes to find out more about the inwardness of the behaviour of water. The remaining sections are briefly concerned with some tangential studies. THERMAL ANOMALIES These are transitions of higher order in the properties of water, occurring within a restricted temperature range, and the problem is, whether or not there is acceptable evidence for their existence.This has been in debate for a long time, notably since Magat’s 1935 paper.60 Certainly not all the reported ‘kinks’ in these properties plotted against temperature can be easily swallowed. There is a company of witch-hunters maintaining that authors reporting anything but perfectly smooth and continuous functions have under-estimated their experimental errors.6l The importance of this problem to the theory of water is apparent from the previous discussion (p. 75) of the possibilities of quasi- or anti-crystalline clusters in liquids, but there is also a possible biological significance in that one of the best supported ‘submerged transi- tions’, at 30-35”C, may have something to do with the slightly higher tempera- ture that evolution has selected as the optimum body temperature of mam- mals.This view is held by Drost-Hansen, who has carried out the most exhaustive assessments of data on water, aqueous solutions, and many kinds of systems with behaviour determined by their properties.62 His firm belief in the reality of these transitions is shared by the present writers63 and new evidence of them continues to come to hand. Two items may be mentioned. Luck64 finds that the wavelength of the maximum infrared absorption in the 1-14-1.25 p region is displaced in the direction of shorter wavelength with rising temperature, but between 36” and 38°C the displacement is greater by an order of magnitude than anywhere else on the temperature scale. Ives and Lemon 99 I I 1 1 I I I I I I I 1 I I 3 0 10 I 20 30 40 50 60 70 Temperature ("C) Fig.15. Differences in i.r. absorptions of water at 2100 and 1900cm-1. Figure 15 shows, as a function of temperature, differences in infrared absorptions of water at 2100 and 1900 cm--l found and plotted realistically by Salama and Goring.65 The method of moving quadratics confirmed the already obvious inflexion between 30" and 40°C. The band at 2100cm-1 is a combination of v2 (1646 cm-l) and librational modes at 500-700 cm-l. The result is associated with a fading-out of the latter with rise of temperature, observed earlier by Magat and attributed to the onset of rotation of water molecules about their axes of symmetry. HYDROPHOBIC BONDING AND CO-SOLVENT BEHAVIOUR The hydrophobic bond, brought into prominence by NCmethy and Scheraga,66 depends for its existence on the peculiarities of water.It can be understood by considering that the dissolution of hydrocarbons in water is disfavoured by the large entropy loss involved in hydrophobic hydration, but is favoured by the heat evolution and the entropy of mixing. Solute particles are randomly distributed and move freely about the solution. The solubility at any given temperature is determined by the balance of these factors; solute particles do not aggregate together because their loss of kinetic freedom would disturb this balance. If, however, the hydrocarbon is not in the form of separate molecules, but is attached as a side-chain to some very large molecule, such as that of a protein, this factor unfavourable to aggregation is obviated.Side-chains of proteins (e.g. of leucine or valine units), normally hydrophobi- cally hydrated in an aqueous medium, do seek their own kind. Contact established between two such side-chains, possibly situated on consecutive RIC Reviews 100 6 i Fig. 16. Sound absorption of t-butyl alcohol-water mixtures as functions of mole fraction of alcohol, xe, after Blandamer et 01.67 turns of a helix, allows some hydrophobic hydration to be sloughed off. The increase of entropy arising from the release of water molecules makes a contribution to the stability of the contact, and dispersion forces do the rest. This is the hydrophobic bond which, with hydrogen bonding between functional groups, is instrumental in determining conformation.Not dissimilar effects occur in ostensibly simpler systems-such as mixtures of water with an organic ‘co-solvent’. The example of t-butyl alcohol in water may be q ~ o t e d . ~ At the limit of low concentration, the difference between the partial molar volume of the alcohol and its molar volume in the pure state (V2 - V i ) is negative-in line with many non-polar solutes. Use is being made of existing cavities in water. But with increasing mole fraction of the alcohol, 7 2 decreases to a minimum which, at 15”C, lies at x2 = 0.04, a mole fraction at which other properties show extrema or inflexions. Quite the most dramatic of these is ultrasonic absorption,67 illustrated in Fig. 16. This can be interpreted in terms of an initial acceptance by the water, as host, of the t-butyl alcohol molecules, the hydrocarbon part of the latter being hydrophobically hydrated, and the hydroxyl group participating in the structure of the hydrogen-bonded clathrate cage.With rising concentra- tion of solute, minor adjustment occurs, in the sense of the water becoming more accommodating for the guest molecules. Then there is a ‘switch’ which is suggested to be hydrophobic dimerization-in effect, hydrophobic bonding between the t-butyl groups of partnered molecules. There can be no doubt that the dimerized solute forms its own, new, larger cage of hydrophobic hydration. Other examples of such dimerization have been reported, and it no doubt plays a part in the pre-micellar behaviour of long-chain soaps9 The general comment on this is that there seems to be practically no limit to the versatility of water in forming hydrogen-bonded cages.Ives and Lemon 101 CONCLUSION An attempt at summing up can take only the form of questions, such as the following : Do current theories of the structure of water take enough guidance from theories of simpler liquids? In the Bernal theory, pseudonuclei are ‘cold’ fluctuations ; in them, the strongest intermolecular attraction available succeeds in producing evanescent, ordered regions of diminished energy and enhanced density. In water, the strongest intermolecular attraction available is almost certainly that of hydrogen bonding, producing evanescent, ordered regions of diminished energy and diminished density-provided ‘clusters’ are accepted.Are these clusters to be considered as pseudonuclei? If the strength of hydrogen bonding depends much on linearity, should the clusters not be expected to be quasi- or anti-crystalline? If so, should not thermal anomalies be a natural consequence? How strong are the grounds for excluding the trapping of ‘molecules-in-transit’ in the cavities of clusters- formed and dissipated with immense rapidity-thus rejecting interstitial models ? In view of the solid-like behaviour of liquids in general, how convincing is the evidence of bistructural or mixture models, to the exclusion of uni- formist theories? Clusters, with their cavities, do not serve to explain the increase in volume between TM and Tc, nor, because of their enhanced rigidity, the facts of fluidity.Are not vacancies of a different order from ice-like cavities required, oppositely dependent in number on temperature ? If so, should they not be built-in to a proper structural theory? One view of hydro- phobic hydration is that an inert solute particle does no more than create a cavity in the water, and that the internal, closed surface of such a cavity is as naturally ‘icy’ as the free surface of water is held to be. Other reviews would be needed to survey the theories of the behaviour of water with respect to a second component. Does a solid surface have a long- range effect on the structure of water adjacent to it? If so, is this imposed order subtended by a misfit region of enhanced disorder, so that we need a theory of a ‘structural double layer’-perhaps before we can fully under- stand the electrical double layer so fundamental to colloid systems ? What is the state of water held in clays and finely porous materials? Is there really, close to some surfaces, a radically different kind of water with inter alia, 12-15 times the viscosity of normal water?69 Similar sets of questions are posed by the reactions of water to ionic solutes-the structure-makers and structure-breakers-but at least in this area we have a theory hard to resist44: that order imposed by ionic fields is subtended by a misfit region of disorder.It is clear that the time has come for an apologia.The writers are conscious of all the evidence ignored or partially and selectively presented, but make the plea that the subject is such that they had no option. There is, however, one final question that may arise in the reader’s mind. How is it that water, being such an infernally complicated thing, so little attention seems to have been paid to it until quite recent times? This returns us to the protective, simplifying principle mentioned in the Introduction (p. 62). It is the Compen- sation Law on which ‘linear free energy relation^'^ depend. In terms of the RIC Reviews 102 standard thermodynamics functions controlling an equilibrium or a reaction rate, AGt = AH$ - TAS$ there is a natural tendency for the A H and TAS terms to bs of the same sign (e.g.if forces of attraction prevail, producing order from disorder, both AH and A S are negative). Their effects on the AG term therefore partly cancel, so that AG is a more stolid, well behaved function than AH or AS. For the reactions in solution (particularly aqueous solution) on the study of which so much physico-chemical progress has depended, there can be no doubt at all that changes in hydration accompany the transformation of reactants into products and contribute largely to enthalpies and entropies of reaction qr activation. It is not however the AH and A S terms that have been predominantly studied, but rather the AG terms-or, what amounts to much the same thing, the equilibrium constants and the velocity constants. It is the well-behaved nature of AG that has afforded the protection from complication. The argument is this70: for an equilibrium in aqueous solution, we can, in principle, split the AGO of reaction into two parts-a ‘reaction- proper’ part (with its intramolecular energy and entropy contributions) and a hydration part-to include all the changing effects of the solute reaction system on the water used as solvent, i.e.AGO = AGr + AGh But if it is assumed that all the hydrational systems remain in equilibrum with one and the same bulk solvent all the time, all the hydrationalprocesses occur at equilibrium, so that AGh = 0. On the other hand AHh and ASh are certainly not equal to zero; it is just that, as in the more familiar freezing and melting processes, AHh = TASh.This seems to be the basis of the special application of the Compensation Law-which can no longer keep us in blink- ers. AGO = AH0 - TASO; ACKNOWLEDGEMENTS The authors record thanks to Dr L. J. Bellamy and Dr D. N. Glew for advice and discussion. The authors also wish to thank the following for permission to reproduce copy- right material: Heinemann Educational Books Ltd for Figs. 3, 5 and 14; Scientific American Inc. for Fig. 7; Verlag Chemie for Fig. 8; Plenum Publishing Corporation for Figs. 9 and 10; The Faraday Society for Fig. 11 ; Pergamon Press for Fig. 12: the American Institue of Physics for Fig. 13; the American Chemical Society for Fig. 15; and The Chemical Society for Fig. 16. REFERENCES Pergamon, 1966.4 E.g. J. E. Leffler and E. Grunwald, Rates and Equilibria of Organic Reactigns. New York : 1 Pocket Oxford Dictionary, Oxford, 1966. 2 M. Pourbaix, Atlas of Electrochemical Equilibria in Aqueous Solutions. London : 3 Cf., however, R. A. Robinson and R. H. Stokes, Electrolyte Solutions, 2nd Edn. London: Butterworths, 1959, p. 103. Wiley, 1963. 5 E.g. A. J. Parker, Q. Rev. chenr. SOC., 1962, 16, 163. Ives and Lemon 103 6 E.g. E. M. Arnett, in Physico-chemical Processes in Mixed Aqueous Solvents, ed. F. Franks. London: Heinemann, 1967, p. 105. 7 F. Franks and D. J. 0. Ives, Q. Rev. chem. SOC., 1966, 20, 1. 8 J. A. V. Butler, Rep. Prog. Chem., 1937, 34, 75. 9 Cf. C. N. Hinshelwood, The Structure ofPhysica1 Chemistry, Oxford: OUP, 1951. 10 E.A. Moelwyn-Hughes, Physical Chemistry. London : Pergamon, 1957. 11 J. S. Rowlinson, Liquids and Liquid Mixtures. London: Butterwarths, 1959. 12 J. D. Bernal, Proc. R. Soc., 1964, A280, 299; Scierzt. Am., No. 267, August 1960. 13 Ref. 10, p. 62. 14 H. Eyring and R. P. Marchi, J. chem. Educ., 1963, 40, 562. 15 A. R. Ubbelohde, Melting and Crystal Structures. Oxford : Clarendon, 1965. 16 C. A. Coulson, Research, 1957, 10, 149; also in Hydrogen Bonding, ed. D. Hadzi and H. W. Thompson. London: Pergamon, 1959, p. 339. 17 E. R. Lippincott and R. Schroeder, J. chem. Phys., 1955,23, 1099; E. R. Lippincott and A. S. Rao, J. chem. Phys., 1964, 41, 3006. 18 L. J. Bellamy and R. J. Pace, Spectrochim. Acta, 1966, 22, 525; L. J. Bellamy, K. J. Morgan and R. J. Pace, Spectrochim. Actu, 1966, 22, 535. 19 P. G. Owston, Q. Rev. chem. Soc. 1951,5, 344. 20 C. Haas and D. F. Hornig, J. clzem. Phys. 1960, 32, 1763. 21 T. T. Wall and D. F. Hornig, J. chem. Phys., 1965,43,2079. 21a M. Falk and T. A. Ford, Can. J. Chem., 1966, 44, 1699. 22 G. E. Walrafen, J. chem. Phys., 1962, 36, 1035; 1964, 11, 3249: 1966, 44, 1546. 22a G. E. Walrafen, J. chem. Phys., 1967, 47, 114. 22b H. S. Frank, Discuss. Faraduy Soc., 1967, 43, 147. 23 K. Buijs and G. R. Choppin, J. chent. Phys., 1963, 39, 2035. 24 W. Luck, Ber. Bunsenges. Phys. 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ISSN:0035-8940    

DOI:10.1039/RR9680100062

出版商:RSC    

年代:1968

数据来源: RSC