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~ QUARTERLY REVIEWS THE CHEMISTRY OF CARBON-OXYGEN SURFACE COMPOUNDS By R. NELSON SMITH Ph.D. (POMONA COLLEGE CLAREMONT CALIFORNIA) IN the many reviews of the preparation properties and structure of activated carbons and carbon blacks little attention has been paid to the properties of the “carbon-oxygen complexes”. Unless very special measures are taken to remove them these complexes are in some degree present on the surface of every carbon used in the laboratory or in commerce and are often the source of the property by which the carbon becomes useful or effective. As might be inferred from a glance at the major sub-divisions of this Review these complexes may often determine the adsorption catalytic and electrical properties of carbon; they act as intermediates in a variety of important reactions; they are vital in the compounding of rubber and the production of inks and paints; and they are important to the lubricating property of carbon.The nature and quantity of the complexes which any given sample possesses are affected by its surface area particle size porosity ash con- tent and degree of carbonisation and graphitisation but in this Review an effort has been made to separate and minimise these effects in order to emphasise the intrinsic properties of the complexes themselves. This separation cannot always be done satisfactorily especially with respect to earlier work because substantial amounts of ash were usually present in the carbons. Also in many cases no distinction is made between “high- temperature” carbons whose oxygen lies almost entirely on the surface and partially pyrolysed carbonaceous materials which have residual hydrogen and/or oxygen distributed throughout their bulk.Mellitic anhydride CI2O9 (50 % 0) carbon suboxide polymer (C,O,) (47 % 0) and graphite oxide (- 30 % 0) are solids with reasonably well- defined properties and structure. This Review however is concerned primarily with carbon combined with 20% or less of oxygen in forms which are not well defined. Their existence was first reported1 in 18 12 and a few additional papers appeared later in that century but the first systematic and comprehensive work2 did not appear until 1913. Formation and decomposition Temperature is the most important factor which determines the nature and formation of the complexes by oxidising gases. If oxygen is added to de Saussure Annals OfPhilosophy 1815 6 241.Rhead and Wheeler J. 1913 103,461 1210. 287 288 QUARTERLY REVIEWS carbon at -78" or below only physical adsorption occurs. Reaction with the surface sets in at higher temperatures and the first amounts of oxygen are bound with no measurable pressure. Indeed adsorption is so effective that traces of oxygen nitric acid and nitrogen dioxide can be removed from nitrogen by beds of activated charcoal at room temperature or higher making the nitrogen suitable for synthesis of ammonia or for analytical purposes. Oxygen-free hydrogen and inert gases can similarly be prepared. Relative to the total oxygen-combining capacity of a carbon ~ample,~ only small amounts of oxygen are fixed at room temperature but the amount increases with temperature until at 400-500 O carbons containing up to about 20% of oxygen may be prepared.These high-oxygen carbons are obtained with simultaneous formation of considerable amounts of carbon dioxide and monoxide. At still higher temperatures the amount converted into gaseous oxides increases and that converted into solid oxides steadily decreases; at 1000" there are virtually no surface oxides. This variation of composition with the temperature of oxidation of pure carbon is not the same as the variation with temperature of pyrolysis of cellulosic or hydrocarbon materials. In the pyrolysis process oxygen and hydrogen are eliminated from the bulk as well as the surface and appreci- able amounts of hydrogen are retained to much higher temperatures than is oxygen. Whatever the nature of the original carbons the optimum temperature to fix large quantities of oxygen is 400-500" with oxygen or nitrogen dioxide and room temperature with ozone.Nitrous oxide carbon dioxide and water are unsuitable for making high-oxygen carbons. Increasing particle size degree of graphitisation and purity and decreasing surface area all tend to decrease the rate of complex forma- tion. If large graphite crystallites are given a prolonged treatment with potassium chlorate in a solution of concentrated sulphuric and nitric acids instead of an exposure to oxidising gases graphite oxide is the principal product with carbon dioxide and mellitic acid as common by- products. Graphite oxide may also be prepared by electrolysis of solutions of hydrofluoric nitric perchloric sulphuric or phosphoric acid graphite being used as the anode.Gaseous oxygen is not produced in these electro- lyses but is instead held interlamellarly by the graphite in what is often a highly swollen solid whose properties depend on the acid concentration. These solution treatments permit a separation of the graphite lamellz and a substantial penetration by the oxidising agent which are not possible with the gases. The peculiar structure and properties of graphite oxide which set it apart from other oxygen complexes arise from this basically different preparation procedure. Penetration of chemisorbed oxygen into the carbon (anode) lattice also occurs during the high temperature (670-770") electrolysis of alumina in molten cryolite. Weller and Young J . Amer. ChPm. Soc. 1948 70 4155. SMITH CARBON-OXYGEN SURFACE COMPOUNDS 289 The main feature about the decomposition* of the complexes is that regardless of the source of carbon the products are always oxides of carbon with any hydrogen coming off as steam below about 700" and above that as hydrogen.Most of the hydrogen comes off above 900". In general surface oxides made at any temperature up to about 500" are thermally stable for long periods and under high vacuum at all tempera- tures below that at which they were made; regardless of the preparative method they are fairly stable up to about 300". At the lowest decomposition temperatures the gas is largely carbon dioxide and above 600" it is largely the monoxide-the ratio of the two gases changing continuously with temperature. Samples once heated to 1000-1200" will seldom show any but trace quantities of hydrogen in the decomposition products.When carbon suboxide polymer decomposes the gaseous product is exclusively carbon dioxide below loo" but at higher temperatures the monoxide is produced in increasing amounts. Graphite oxide decomposes to give carbon dioxide steam and small amounts of carbon monoxide. In both cases the solid product is an oxygen-complexed carbon which becomes graphite at high temperature. Structure The formulation of a structural theory of surface oxygen complexes which gives an unequivocal integrated explanation of all the properties described in this Review is far from achievement. The principal functional groups which have been found5y6 by direct chemical analysis are >C=O -OH and -CO-OH with some evidence for a small number of ester groups; the difference (often as much as 75%) between total oxygen and that determined by functional-group analysis is usually assigned to "ether oxygen".The presence of these groups as well as aromatic rings and C-H bonds is also shown by infrared spectra. There is a great need for more information about the analytical behaviour of functional groups on large polynuclear aromatic molecules such as those which correspond to a graphite layer in a crystallite. The evidence already at hand with such model compounds shows that quinone ketone and phenol groups often do not react at all or only partly react with the standard reagents. There is also the question of specificity of reagent. Infrared spectra give only limited help because the wide absorption bands and high background absorption do not permit unequivocal assignment of frequencies and there is usually more than one way to interpret changes in relative band intensities resulting from chemical treatment.Furthermore carbons having more than about 94% of carbon are opaque to infrared which in turn means that all infrared information has been obtained from coal coal extracts partially pyrolysed sugar or highly oxidised blacks rather than from the surface complexes of relatively pure carbon. Anderson and Emmett J. Phys. Chem. (a) 1947 51 1308; (b) 1952 56 753. Studebaker Huffman Wolfe and Nabors Ind. Eng. Chem. 1956,48 162. Blom Edelhausen and van Krevelen Fuel 1957,36 135. 290 QUARTERLY REVIEWS There seems to be no evidence for alcoholic hydroxyl groups The separate determination of phenolic and hydroquinonic (quinolic) -OH is difficult though a real advance was made in the recent development of a constant-potential coulometric method for hydroquinone (quinol) and quinone groups on blacks.Potentiometric titrations indicate that the acidic groups have a pK similar to that of phenol and occasionally pK values similar to that of benzoic acid. The high-potential emission spectra of some carbon blacks showed the existence of twelve fragments in the discharge and of interest here is the fact that in addition to considerable -OH and CO there was a high concentration of C02+ which must have come from -C02H and of -CHO which must have been originally associated in that form on the surface. A number of carbon-oxygen structures have been proposed for which no direct experimental proof has been given.A recent example of this is a comprehensive theory’ which suggests that the acidic groups on the surface are primarily lactones and only to a small extent phenolic. Two types of lactone are postulated (1) a fluorescein type which can form a sodium salt and be methylated and (2) a “normal” (probably 8) type which can be hydrolysed by sodium hydroxide but not methylated. The theory explains the properties of carbons with basic groups by postulating the presence of a benzopyran-benzopyrylium redox system. Some specific chemical compounds have been removed from carbon surfaces but their removal demonstrates an exceptionally limited type of chemical reaction with oxygen in the presence of the extractant (water) . rather than proof of the existence of certain groups on the surface.For example hydrogen peroxide is produced in very small amounts by oxygen in acid solution or by the combined action of oxygen and water vapour; and minute traces of oxalic acid can be repeatedly extracted with water from a given sample of ashless char. Diphenylene oxide 2,6-xylenol and p-phenylphenol have been extracted in trace quantities from coals by pyridine and ethylenediamine but they probably represent intermediate fragments in the coalification process rather than the surface complexes formed by oxygen or other oxidising agents. Another approach to structure determination is to identify the degrada- tion products resulting from oxidative methods and then to deduce from these what the parent compounds must have been. In the main however such studies substantiate the graphitic nature of coals but give no specific information about the surface complexes.High-temperature phenomena such as combustion probably do not depend so much on the properties of specific surface structures which would tend to be unstable at these tempcratures as they do on mobility of oxygen atoms interlamellar penetration and lattice properties. Graphite oxide (GO) has a controversial formula [(C,O 2H)z (C704H2)% (C,O,H&J and many properties in common with the carbon-oxygen Garten and Weiss Rev. Pure Appl. Chem. 1957 7 69. SMITH CARBON-OXYGEN SURFACE COMPOUNDS 29 1 surface complexes of chars and blacks with high percentages of oxygen. In spite of the compounds' having a "formula" these properties are retained on thermal decomposition over a composition range of 40-1 8 % without significant change in X-ray structure and even down to low percentages of oxygen as the GO structure collapses.The chief difference between GO and the other oxygenated carbons is that GO has a high concentration of oxygen-containing groups attached interlamellarly in the relatively easily accessible space (6-2 A) between the carbon planes. The interplanar distance in graphite is 3-44 A. There is general agreementS that hydroxyl is one of the major groups attached to carbon in GO and it appears that some hydroxyl groups are involved in a keto-enol equilibrium. The remaining oxygen atoms are believed to be attached to "meta-oxygen" bridges (i.e. oxygen attached to the first and third of a chain of three carbon atoms) which are randomly. distributed both above and below each carbon layer.There are a few carboxyl groups probably on the periphery of the GO layers. The carbon layers themselves are puckered rings unlike the planar aromatic layers of graphite crystallites. Controlled thermal decomposition first eliminates the meta-oxygen bridges as carbon oxides. Finally at 130-180" the hydroxyl groups are eliminated with concornitan t collapse of the c-spacing and disappearance of swelling properties. Rapid heating to 200" causes GO to explode and produce exfoliated carbon exceedingly thin wafers of graphite carbon containing 20% of oxygen. Selected oxidation reactions Combustion of Carbon.-In spite of a long history of intensive research on the combustion of carbon there is still no comprehensive theory which satisfactorily accounts for all of the kinetic structural catalytic and temperature-effect data which have accumulated.The major facts,g which have been established under conditions which exclude secondary pro- duction of carbon dioxide by carbon and oxidation of the monoxide by oxygen indicate that the primary step in the combustion of pure carbon is dissociative adsorption of oxygen to give oxygen complexes. Below about 500" the rate of formation of these complexes is greater than their rate of decomposition but as the temperature of combustion increases these complexes become increasingly unstable and reactive and the composition of the primary decomposition products (CO and CO,) alters to become predominantly CO especially above 1000". The two oxygen atoms in the dioxide come from different oxygen molecules; thus appreciable formation of carbon dioxide requires a low enough tempera- ture to permit a significant surface concentration of oxygen complexes for reaction with oxygen.At high temperatures the mobility of oxygen atoms no doubt plays a part and at still higher temperatures (of the order of Clam Plass Boehm and Hofmann 2. anorg. Chem. 1957,291 205. Rossberg 2. Elektrochem. 1956 60 952. 292 QUARTERLY REVIEWS 2100”) self-diffusion of carbon atoms and increase in degree of graphitisa- tion set in. Since experimental conditions or the nature or quantity of the sample may also superimpose diffusion processes secondary reactions and varying surface areas it is no wonder that the kinetics of this “simple” reaction have not been completely elucidated even for pure carbon. A rate equation derived receiitlylO to describe the oxidation of carbon Jilaments takes account of most of these variables and reasonably assumes the presence of reactive hydrogen atoms throughout the lattice as major impurity sites but it makes no guess as to the structure of the oxygen or hydrogen complexes.Most of the detailed oxidation mechanisms are oversimplified and postulate attack on prism faces of the crystallites at edge atoms of the layer planes and assume functional group reactivities characteristic of relatively simple organic molecules. There is evidence however that the base-plane atoms of the crystallites also have strong binding and reactivity. High-temperature experiments indicate that it is the disorganised carbon atoms and the smallest crystallites which first disappear partly through reaction with oxygen and partly through formation of large crystallites ; when a base plane disappears it does so all at once.Crystallite diameters and gross particle sizes stay remarkably constant through the major part of combustion. Combustion is exceedingly sensitive to catalytic impurities the most effective being potassium and sodium oxides and carbonates. The principal effects are increase in rate an increase in the relative amount of carbon dioxide produced and change in the order of reaction. Mercury vapour (from manometer etc.) has a similar effect. Water vapour catalyses the combustion but has little effect on the CO/C02 ratio. Traces of halogen- containing vapours (phosphoryl chloride carbon tetrachloride and chlorine) retard the decomposition of the complexes but do simplify such studies by preventing the secondary oxidation of carbon monoxide.Oxidation by Carbon Dioxide.-Carbon dioxide oxidises carbon to the monoxide by a relatively simple mechanism with oxygen complexes (CO) as intermediates.ll The steps are fast co,+ c-+co + (CO) . ’ * slow . . . . . . . (CO)-+CO (2) c o + ( c o ) - + c o ~ + c . (3) . . . . . Diffusion processes being excluded the reason that the mechanism is not more complicated is probably that the dioxide does not react appreciably below 600° and hence it is not possible to form those low-temperature lo Blyholder Binford and Eyring J. Phys. Chem. 1958 62 263. l1 Reif J . Phys. Chem. 1952 56 778, SMITH CARBON-OXYGEN SURFACE COMPOUNDS 293 complexes which yield carbon dioxide on decomposition ; further free oxygen is not involved.A side reaction which plays no significant rale in the oxidation is the exchange of carbon atoms between the dioxide and the carbon lattice but not with carbon monoxide. The temperature must be greater than 700" and at 1600" a really significant exchange takes place. Potassium sodium and iron oxides are excellent catalysts. Steam-Carbon Reaction.-It cannot be said that the mechanism of this reaction is completely established beyond doubt but the evidence12J3 available strongly suggests the existence of two surface complex inter- mediates which might be designated as the TO-complex" (CO) and the TO,-complex" (CO,). The postulated mechanism steps are H2O + C-+ H2 + (CO) H2O + (CO)--+M2 + (CO,) . . . . (4) (CO) -+ co (5) (6) (CO,) -+ CO (7) (8) . . . . .. . . . . . . . . . . . . . . . . . . . H + (CO) -+ H,O At low temperatures only steps 4 6 and 7 are important because the "CO-complex" is thermally stable and unreactive towards hydrogen. At temperatures af about 600-1400" all steps are important in a closed system or in a system with appreciable added or accumulated partial pressure of hydrogen; in a fast-flow system or under conditions where hydrogen cannot accumulate step 8 is unimportant; at very low partial pressures of steam in a flow system only steps 4 and 5 are important. At high temperatures (above 1400") only steps 4 5 and 8 are important because the very rapid decomposition of the 'TO-complex" leaves little opportunity for 6 and 7 to occur; 8 is again unimportant under conditions of low partial pressure of hydrogen. The steam-carbon reaction is exceptionally sensitive to catalytic impurities such as potassium sodium iron and copper oxides the order of the reaction may be changed and with very large amounts of potassium carbonate the whole nature of the reaction is changed so that hydrogen and carbon dioxide are the main products along with an appreciable amount of methane.The large increase in carbon dioxide may be due largely to the CO-H,O reaction especially when diffusion is important. As steam pressures rise from 1 to 50 atm. in the range of 750-850" there is a rapid increase in methane production. The mechanism of this competing reaction is still in doubt. Complexes are probably involved only if methane is made by the direct action of water on partially hydro- genated carbons and not if by direct action with hydrogen.Oxidation by Nitrous Oxide.-The oxidation of carbon by nitrous oxide proceeds by a very simple mechanism l2 Singh Parkash and Puri Clzern. and Znd. 1959 18. l3 Wicke and Rossberg 2. Elektrochein. 1953,57 641. 294 QUARTERLY REVIEWS fast slow N20 + C-+ N + (CO) . . . . (9) NZO + (CO) -+N2 + COa . (10) (CO) -3 co . . . . . . . (1 1) Only step 9 occurs at low temperatures and the reaction stops as soon as all reactive sites have been converted into oxygen complexes. Reaction at 0" has been reported but normally 100-300" is required. Step 10 sets in at some intermediate temperature and the production of carbon dioxide is a first-order reaction in the range 300-700". With a filament and low nitrous oxide pressure at high temperatures (900-2000") there is small probability that nitrous oxide wi!l react with a complex whose concentra- tion is very low because of rapid thermal decomposition and consequently carbon monoxide is the principal product with only traces of the dioxide.As in the other oxidations ash acts catalytically. Oxidation by Nitric Oxide.-The mechanism for this oxidation is 2N0 + 2C' +N2 + 2(C'O) . . . . (12) NO + (C'O) ?+ (OC'. *ON) . . . (13) NO + (OC' - ON) + C" -+ C'O + N + (C"0). NO + (OC' * ON) + C"' -+ C'O + N2 + C"'0 . (14) . (15) Below 200" only step 12 occurs and this ceases when the reaction sites have been covered with complexes. Step 13 involves the reversible physical adsorption of nitric oxide at the oxygen complex sites and the (C"0) complexes formed at C" adjacent to C' can serve as adsorption sites in 13.After providing oxygen complexes at the outset of a reaction when there are no oxygen complexes the rble of step 12 is to replenish those complex sites not regenerated in step 15. Step 15 is minor compared to step 14 contributing about 8 % at 450" or 20 % at 600". Tke steady-state concen- tration of the complex decreases with increasing temperature. Ash acts catalytically . Oxidation by Bacteria.-It has been neatly demonstrated that certain soil bacteria will at room temperature in the presence of sterile carbon dioxide- free air oxidise chars or blacks that have been pyrolysed at 1200". The ash was not removed and the water vapour needed to preserve the bacteria may have participated in the oxidation process. The rate more than doubled from 20" to 40"; the bacteria died at 100".No test was made for the pre- sence of surface complexes or for whether the bacteria actually attacked complexes rather than bare carbon atoms but both seem very likely. Adsorption of electrolytes Adsorption of electrolytes varies tremendously with the nature and amount of the surface-oxygen complexes. Ash-free carbons outgassed at SMITH CARBON-OXYGEN SURFACE COMPOUNDS 295 1000" and subsequently exposed to oxygen at temperatures below about 200" or above about 750" will adsorb strong acid from solution but no strong base. The same carbons heated in oxygen in the range 200-750" will adsorb strong base but no strong acid. The optimum temperature of oxygen-treatment for base adsorption is about 400-500" the same temperatures as for the maximum production of surface oxides.There is however no clean-cut temperature boundary between acid and base adsorption but instead a certain overlap where one drops off and the other begins. Carbons made by low-temperature carbonisation adsorb base only the amount decreasing as the temperature is raised from 500" to 750" whereat acid adsorption sets in. If the acid and base adsorbing capacities of a given carbon are compared under conditions for optimum adsorption of each base adsorption is found to be far greater than acid adsorption. The many and diverging views of electrolyte adsorption were clearly s~mmarised'~ in detail in 1944. With the additional research reported since then it appears that the picture presented below satisfactorily explains the published data. Acid-adsorbing Carbons.-Adsorbates on these carbons may be classified in the order of decreasing adsorbability (a) non-electrolytes (6) large ions of organic strong electrolytes (c) hydrogen ions on oxygen complexes (d) hydrogen ions on bare carbon and (e) metal ions inorganic anions and hydroxyl ions (these last can be held only in the diffuse double layer).The following equilibrium expressions using these relative adsorbabilities will explain the adsorption phenomena. (C,O) + H+ + A- + (C,OH)+ + A- . . . (16) (C,) +H+ +A-+(C,)H+ + A - . . . . (17) (C,O) + H,O + (C,OH)+ + OH- . . (18) The electrokinetic potential of the carbons (parenthetic entities on the right) is positive as shown. (C,O) appears to be quinonic in nature. Equilibrium 16 observed only in dilute solution (< O-O~N) has been confirmed by experiments which show up to a point a 1 1 correspondence between surface oxygen atoms and adsorbed acid; if there are no oxygen atoms there is no adsorption.At higher acid concentration adsorption by step 16 still takes place but is masked by the larger additional adsorp- tion of step 17 which varies with concentration according to a Langmuir- type expression. The acid adsorbed by step 17 is easily washed out with water but that held by step 16 cannot be washed out-it must be removed with sodium hydroxide the excess of which is easily washed out with water. The anions held in the diffuse double layer are easily exchangeable for other anions from neutral salts. The primarily-held hydrogen ion may be displaced by cations of strong organic bases such as methylene-blue l4 Steenberg " Acsorption and Exchange of Ions on Activated Charcoal " Almquist and Wiksells Uppsala 1944.296 QUARTERLY REVIEWS cation leaving A- exchangeable. If the primarily-held hydrogen ion is displaced by anions of strong organic acids (picrate) it is the positive ions which become exchangeable. Equilibrium 1 8 explains why aqueous suspen- sions of the carbons are alkaline though only slightly so because most of the hydroxyl ions are held in the diffuse double layer. When the carbons are suspended in solutions of neutral salts the solution becomes appreciably basic because the hydroxyl ion goes into solution by anionic exchange. The surface oxygen complexes involved in steps 16 and 18 may be formed before acid adsorption measurements or simultaneously with them. If the adsorption is performed in air so that gaseous oxygen as well as acid is available another type of complex (C,O,) can be formed which leads to the production of small amounts of hydrogen peroxide as follows (C,O.J + H,O+ + A- -+ (C,OH)+ + A- + H202 .(19) If a carbon is used which has no surface oxygen the initial amounts of gaseous oxygen are used to form (C,O) which does not lead to hydrogen peroxide formation. The amount of peroxide produced is equivalent to less than 10% of the acid adsorbed in step 16. Base-adsorbing Carbons.-These carbons have a negative electrokinetic potential and their aqueous suspensions are acidic. It is not possible to wash out most of the adsorbed base with water and large amounts of non-electrolytes can be added without affecting the adsorption of base. They have cation-exchange properties but with less capacity than for base adsorption.These properties may be accounted for by equilibria involving two types of site (C,OH) being hydroquinonic (quinolic) or phenolic in nature and (C,O) of the nature of an acid anhydride or fluorescein lactone. (C,OH) may be made by reduction of (C,OH)+. (C,O) + H,O + (C,02H)- + H+ . . . (20) (C,OH) + Na+ + OH- + (C,O)- + Na+ + H20 . (21) Equilibrium 20 accounts for the slight acidity of aqueous suspensions only slight because most of the hydrogen ion is held in the diffuse double layer. . The hydrogen ion may be exchanged for sodium ion from sodium acetate as is commonly done for analytical purposes. If it is exchanged for sodium ion in sodium hydroxide then the total base adsorption is determined by both types of complex; this is also done for analytical purposes.Graphite oxide also has pronounced cation-exchange properties. The behaviour of large organic strong-electrolyte ions is not clear but it is certain that they are involved in both primary and exchange adsorption and that the complexes can cause an enormous increase in adsorption for no change in area. This sensitivity of ionic dyes to oxygen complexes as well as to pore size makes them unsatisfactory for measurement of surface area. SMITH CARBON-OXYGEN SURFACE COMPOUNDS 297 Adsorption of non-electrolytes In general surface oxygen complexes have no effect on the physical adsorption of non-polar gases. Nevertheless the complexes have been known to cause errors in interpretation of physical adsorption data. For example even though the charcoal samples had been outgassed at 300" in vacuo low-pressure adsorption of carbon tetrachloride (less than 0.01 mm.) was shown to displace part of the complexes as oxides of carbon and to provide a gas pressure of these oxides comparable to the carbon tetrachloride pressure.Chlorobenzene hydrogen chloride and silicon tetrafluoride will do the same as will doubtless many other gases not proved to do so. Such chemical displacement is not an effective method for removal of complexes. Certain oxidising gases which form complexes on adsorption may also yield misleading data. Thus the "adsorption" of nitric oxide on carbon liberates nitrogen leaving the oxygen on the surface and leads to quite erroneous conclusions if the assumption is made that the equilibrium gas pressure is nitric oxide.The adsorption of polar gases on carbon is usually seriously affected by oxygen complexes. The adsorption of H20 CH,-OH NH3 CH3.NH2 SO2 CO and HCl is greatly increased and though these gases are thus unreliable for surface area measurements they may be useful for character- ising the polar sites on a carbon surface. Sulphur dioxide is especially recommended for this. Certain of these gases have special peculiarities. Ammonia and methylamine are adsorbed so abnormally that it has been suggested that perhaps they dissolve in the surface oxides or form a salt; they are desorbed at pressures lower than that at which they were adsorbed. At 200" or above ammonia will react with surface oxides to give surface nitrogen compounds (possibly amides) which are stable to above 400". Sulphur dioxide on chars in the presence of oxygen forms such a stable complex that it cannot be desorbed up to 220" and then only as the trioxide-in fact it serves as a basis for a method of making sulphuric acid.Water,15 being carbophobic is negligibly adsorbed on pure carbon at low pressures P. If the carbon is non-porous adsorption rises sharply when P > 0-9P0 and ceases at Po (vapour pressure of pure liquid) with something like two statistical layers instead of the thirty or more which would have been expected with a non-polar vapour. If the carbon is porous adsorption rises sharply when P reaches about 0-5-0.6P0 and then remains fairly constant up to Po with an adsorbed liquid volume of water very nearly the same as that for a non-polar vapour. The effect of oxygen complexes (or ash!) on these isotherms1* is to increase adsorption at low pressures to shift the rapid rise of adsorption to lower pressures and to increase the total adsorption at Po.Up to the beginning of the rapid rise there is a 1 1 correspondence between chemisorbed oxygen atoms and l5 Pierce Smith Wiley and Cordes J. Amer. Chem. Soc. 1951,73,4551. 16 Dubinin and Zaverina Zzvest. Akad. Nauk S.S.S.R. Otdel. khim. Nauk 1955,594. 298 QUARTERLY REVIEWS adsorbed molecules. Because of strong water-water interaction additional adsorption occurs on these oxygen-atom sites and this leads to “island” formation rather than to “monolayer” formation. This island formation which is peculiar to water on carbon provides a situation in which the adsorption process is different from the desorption p r o c e s ~ .~ ~ * ~ ~ Thus in the case of porous carbons the capillaries are filled by coalescence of islands but they empty from menisci; and desorption from non-porous carbons is from large islands which were formed on adsorption by coales- cence of small ones. Water held in capillaries or in large islands over several polar sites is held more strongly than in small islands over one polar site and lower pressure is needed for desorption. As a consequence water desorption curves always lie to the left of the adsorption curves at pressures above those needed to adsorb one water molecule on every chemisorbed oxygen atom. Adsorption of water by graphite oxide is a very slow interlamellar process and several weeks are needed to attain equilibrium at each pressure. Interlamellar spacing is 6-2 A for 0-4 % of water 6.8 A at 11 % of water and approximately 11 A with 60% or more.Desorption is also slow and removal of the last traces of water is almost impossible. The adsorption of many common non-electrolytes from solution is affected by the presence of the oxygen complexes. Three of these are worthy of special comment. Iodine has been widely used for surface-area determin- ations because of simplicity in analytical determination but data from many sources indicate that oxygen complexes severely decrease iodine adsorbed either from water or from organic solvents and render this method unsatisfactory for area measurements. A standard measurement involving the adsorption of diphenylguanidine (an accelerator in rubber manufacture) from benzene is commonly used as a sensitive measure of the amount of surface oxygen.Recent experiments in which different surface groups were chemically blocked indicate that the surface carbonyl group is primarily responsible for diphenylguanidine adsorption. Toluene and benzene are selectively adsorbed from heptane solution if complexes are on the surface and essentially the reverse is true without the complexes. It is judged that this behaviour is caused by the strong interaction of the n electrons of aromatic nuclei with the “acidic complex” in a manner similar to the interaction with silica gel. Graphite oxide can reversibly adsorb polar molecules either from pure liquid or from hydrocarbon solution. Dioxan can form either one or two flatly oriented layers between carbon planes giving rise to lamellar separations of 9.8 or 14.5 A.In the homologous series of alcohols the swelling is about constant up to propanol then becomes increasingly large with butanol pentanol etc. indicating a change from “fiat” to “upright” orientation. l7 Pierce and Smith J. Phys. Colloid Chem. 1950 54 784. lS Dubinin Zaverha and Serpinsky J. 1955 1760. SMITH CARBON-OXYGEN SURFACE COMPOUNDS 299 Catalytic activity Reactions involving Oxygen.-In most cases of carbon-catalysed air oxidation it has not been explicitly demonstrated that surface oxygen complexes participate in the oxidation. In the case of Fe(CN)64- NO2- As0,- Sn2+ and quinol it has been shown19 that the maximum activity is observed with carbons with maximum amounts of oxygen and it is likely that the complexes are involved with all examples reported here.(a) Organic compounds catalytically oxidised in aqueous solution are amines amino-acids anilines N-substituted anilines and aminophenol ; formic oxalic malonic and thioglycollic acids ; salts of uric and substituted uric acids in alkaline solution; thiourea and phenylurea ; quinol; allantoin; hexosephosphoric acid. Catalysis may be promoted by ash especially by the Fe-C-N linkage. Cyanide and thiocyanate ions eliminate the promoted catalysis of iron and pentyl alcohol and urethanes inhibit by desorbing the species to be oxidised. It appears that hydrogen peroxide is not an intermediate because the presence of potassium iodide has no effect on the velocity. Compounds not oxidised are ethanol amine hydrochlorides fatty acids carbohydrates xanthine hypoxanthine caffeine and t h e e bromine.(b) Inorganic substances or ions catalytically oxidised in alkaline solution are I- s= so3“ NO2- Sn02= A s O ~ ~ - Fe(CN)64- N2H,. In acid solution Sn2+ Fe2+ Hgz+ I- Fe(CN),4- SOz HzS. (c) The oxidation of nitric oxide at room temperature and of sulphur dioxide from room temperature up to 200” is strongly catalysed and there is qualitative evidence for the catalytic oxidation. of ethylene propene and ethanol vapours. Reactions which do not involve Oxygen.-(a) The oxidation of CO by H20 N20 and NO proceeds in each case through a surface oxide. For H20 and N20 (represented as X20) the mechanism is x20 + C-t(C0) + x . . . . . (22) c o + ( c o ) - + c o + c . . . . . . (23) The mechanism for NO has not yet been worked out. (b) The reaction of hydrogen with bromine at 150° used as a simple measure of “catalytic activity” of carbons is inhibited by the complexes.(c) The following ion-exchange mechanism20 accounts for most of the facts known about the hydrogen peroxide decomposition [(CzOH)+ + OH-] + HOz- + [(C.OH)+ + HO2-] + OH- (24) King J. 1936 1688. von Brinkmann KolIoid Z. 1951,123 116. 300 QUARTERLY REVIEWS Decomposition is rapid in basic solution where the chain steps 24 and 25 are favoured by having H02- in reasonable concentration and by having a carbon with a large number of basic groups (C,O) on the surface. Decom- position is slow in acid solution where KO2- concentration is very low and equilibrium 16 not 24 is favoured. Step 26 accounts for the observa- tion that when carbons catalyse the decomposition of hydrogen peroxide their acid-absorbing capacity decreases and their base-adsorbing capacity increases by an amount which is at least double the loss in acid-adsorbing capacity.Step 26 being a chain-breaker stops catalysis and accounts for the fact that eventually a catalysing carbon is virtually unable to effect further catalysis. If at the outset hydrogen peroxide creates a few new (C,O) sites by direct action on the carbon the apparent increase in rate constant with the peroxide concentration is explained. At elevated tem- peratures carbon dioxide is also produced being in excess of oxygen in dilute solution. Carbons made from nitrogen compounds give immeasur- ably fast action and physically adsorbed gelatin inhibits catalysis. (d) Para-hydrogen is usually produced by conversion from ortho- hydrogen in the presence of suitable charcoals at liquid-air temperature.Recent work21 shows that the best charcoals for this purpose are prepared by carbonising sugar (or other cellulosic compounds) at about 600" to give a char containing 90-94x of carbon. The oxygen in this product cannot be considered to be primarily in the surface or involved in the para-ortho conversion. Instead paramagnetic resonance studies have shown that such a carbon possesses the maximum concentration of free radicals caused most probably by the presence of large condensed resonat- ing ring structures with unsaturated edge bonds and stabilised unpaired electrons. The para-ortho conversion involving a paramagnetic mechanism is improved by the presence of paramagnetic oxygen physically adsorbed at liquid-air temperatures but oxygen adsorbed above -80" eliminates the paramagnetic centres (and conversion ability) because each oxygen molecule shares a pair of electrons with a charcoal free radical.Two pairs of electrons are probably not shared by one oxygen molecule because of the low radical concentration; instead it is thought that a stabilised peroxidic complex -R-0-O. is formed which is incapable of effecting the para-ortho conversion. (e) The carbon-catalysed hydrogen-deuterium exchange reaction at 50" is inhibited by the oxygen in and on the carbon for catalytic activity increases with carbonising temperature from zero at 650" to high activity at 950". The effect of surface oxygen alone has not been published. It is postulated that a pool of interacting electrons is needed for catalysis here 21 Turkevich and Laroche 2.phys. Chem. (Frankfurt) 1958,15 399. SMITH CARBON4XYGEN SURFACE COMPOUNDS 30 1 rather than individual non-interacting electrons as in the case of ortho- para conversion. Thermodynamic properties Chemisorption.-The first quantities of oxygen admitted to a carbon surface at room temperature are chemisorbed with exceedingly high heats of formation (70-100 kcal./mole) of the same order of magnitude as the heat of formation of carbon dioxide. These heats decrease rapidly as additional oxygen is added finally levelling out at about 5.5-6-0 kca1.l mole when physical adsorption sets in. Nitric acid appears to have a very high heat of adsorption (-70 kcal./mole) for the initial amounts but this has been shown to be the heat of the reaction which forms nitrogen gas and surface complexes.However whether made by nitric oxide or by oxygen the initial heat of formation of the complex is about 40 kcal. per g.-atom of oxygen. The heat of formation of complexes made at 400" (calculated from heats of combustion) is -32.5 kcal. per g.-atom of oxygen independent of the amount of oxygen combined. The stability of these 400" complexes is thus almost as great as that of the small initial quantities adsorbed at room temperature and vastly greater than that of the bulk of the oxygen adsorbed at room temperature. In contrast to this the heat of combustion of graphite oxide varies only from 92 to 94.1 kcal. per g.-atom of carbon even though the oxygen content may vary from 0 to 30%. This heat of combustion almost identical with that for graphite indicates the very weak oxygen binding in graphite oxide and is in keeping with its unstable and explosive nature.Physical A dsorption.-Direct calorimetric measurement22 of the integral heats of adsorption of NH, CH3.NH, CO, and SO on carbon blacks indicates that these gases have a high initial heat of adsorption if the surface has oxygen complexes. With increasing coverage the heats for the last two fall steadily and approach the heat of liquefaction. Ammonia however after an initial sharp drop continues to have a high heat about 3 kcal./ mole greater than the heat of liquefaction independent of coverage in the region studied. In general the heat of adsorption of non-polar gases is unaffected by the presence of complexes. Because the adsorption of water on carbon is small and accompanied by very small heat effects direct calorimetry is not very reliable and heats of adsorption must be determined indirectly.The common method of determining isosteric (differential) heats of adsorption by calculation from adsorption isotherms obtained at two temperatures has been applied to water and has led to conflicting conclusions as to whether more or less heat is liberated on adsorption than in condensation on bulk water. The very small amount of adsorption the continual reaction of water with the surface causing an increase in surface polarity during the course of iso- therm measurements and the inapplicability of the Clausius-Clapeyron 22 Spencer Amberg and Beebe J. Phys. Chem. 1958,62 719. 302 QUARTERLY REVIEWS equation under these conditions contribute to this difficulty.A more reliable method combines heat of immersion measurements with one adsorption isotherm and it shows that the first amounts of water vapour adsorbed at low pressure are adsorbed with less than the heat of condensa- tion to pure liquid. Spontaneous adsorption of this nature is possible here because a compensating increase in entropy of the adsorbed film over that of the corresponding pure water is provided by the practically complete isolation of the oxygen-complex adsorption sites. The heat of immersion in water and methanol is a linear function of the number of oxygen in agreement with the 1 :1 correspondence of oxygen atoms with adsorbed water molecules in the low-pressure regions. There is also a direct relation between the heat of immersion and the adsorbability of baryta and also the quantity of oxygen that can be removed as carbon dioxide.Oxygen which may be evolved as carbon monoxide or water appears to be un- related to the heat of immersion. Adsorption of water by graphite oxide involves the additional factor of swelling the heat of which has been judged to be very small. The heat of adsorption of inorganic bases from solution on to ash-free charcoals previously wet with water corresponds in magnitude to the heat of neutralisation of carbonic acid by the respective bases. Electrical properties Oxygen EZectrode.-The only known reversible oxygen gas electrode at room temperature involves carbon as the conducting solid. This was first discovered in 1943 when it was shown that the reversible reaction requires the presence of OH- and H02- in solution.Measurements of magnetic susceptibility and of electron paramagnetic resonance indicate that the oxygen so involved may be adsorbed as a peroxidic complex (C,02). This is further substantiated by recent work24 with l80 which shows that none of the hydrogen peroxide involved in the reversible electrode reaction involves the breaking of the 0-0 bond. Thus the electrode reaction is probably (C,) 4-02 +(C,O,) . . . . . . (27) (C,02) + H20 + 2e s (C,) + H02- + OH- It is likely that step 28 is composed of two one-electron steps involving an adsorbed HO,. radical as intermediate. A certain similarity of this mechanism to equation 19 for hydrogen peroxide formation or to the mechanism for catalysed hydrogen peroxide decomposition is at once evident but carbons used for reversible electrode work have negligible area and are probably unable to show the other two effects.Electrical Resistances.-Carbon-carbon contact resistance under low a4 Davies Clark Yeager and Hovorka J. Electrockem. SOC. 1959 106 56. . (28) Kraus J . Phys. Chem. 1955 59 343. SMITH CARBON-OXYGEN SURFACE COMPOUNDS 303 weight loads is not affected by simple evacuation. High-temperature evacuation removes the oxygen complexes and lowers the resistance for a 1 mg. load by about 75% but for loads greater than 100 mg. the contact resistance is virtually unaffected; mm. of O2 will raise the contact resistance again. In general oxygen complexes increase the electrical resistance of carbon but they do not seem to interfere with the slight decrease in electrical resistance usually caused by the physical adsorption of other gases.A careful studyz5 of the variation of the bulk electrical resistance with surface oxygen showed that a fairly sharp maximum in resistance occurs with a thoroughly outgassed (at 1200”) powdered carbon when it is re-oxidised in the temperature range of 500-70Q0 near the optimum range for forming large amounts of complexes. This maximum amounting to about a 1000- fold increase in resistance corresponded to about 4 % of oxygen. Addition of oxygen at 20” to a well-outgassed carbon filament causes an increase of resistance of a few units % over a period of several days but if the oxygen is added at a high temperature (1000-1200”) an immediate increase in resistance of about 25% is observed. Both the filament and the powder work indicate that part of the effect of oxygen is produced by interlamellas penetration of oxygen.It is noteworthy that in the case of graphite oxide where almost all of the oxygen is interlamellar a tremendous drop in resistance occurs on slow thermal decomposition as the oxygen content goes from 24 % to 22 % though there is no significant change in interlamellar spacing. ~isceZZaneous.-Oxygen complexes formed by adding oxygen to out- gassed carbon also cause an immediate large rise in the surface potential which subsequently disappears during several hours. In addition 1 0-2 mm. of oxygen will cause a 50% reduction in the ultraviolet photoelectric sensitivity of carbon. Rubber of the total carbon black production (approximately 2-2 x lo9 lb. in 1957) is used as a filler for rubber products.Such products to be satisfactory require a strong bond between the black and the rubber polymer. Surface oxygen complexes play an important r6le in this binding and to a certain extent determine the preparation and properties of the final product. Thus blacks with a high concentration of surface oxygen will yield a rubber product which ages more rapidly and has a higher electrical resistance and a mix which has a longer “scorch time” (induc- tion period for premature vulcanisation) than those with low concentration of surface oxygen. The complexes have a profound retarding effect26 on the rate of “cure” 26 Hirabayashi and Toyoda Tanso 1954 4 2. t 6 Cines Rubber Age 1951,69 183. About 95 304 QUARTERLY REVIEWS (vulcanisation). If for a given formulation all variables are held constant except for the proportion of surface oxygen the resulting rubber product is observed to be increasingly poor with respect to mechanical tests such as tensile strength stretching modulus and abrasion resistance with increas- ing content of surface oxygen.Primarily this decrease in quality of product indicates only an increasingly poor state of vulcanisation and if in each case the accelerator concentration is adjusted so as to give the same state of cure for a given cure-time the mechanical test results from the product seem to be essentially independent of the state of oxidation of the black. The nature of the binding between carbon particles and polymer is not at all well known. Some relevant experimental facts which must be con- sidered are (a) dihydromyrcene (a diolefin) is adsorbed at 100” on a channel black (3-5% of oxygen) with polymerisation and a high heat of adsorption whereas only reversible adsorption without polymerisation occurs on the same black without oxygen complexes; (6) but-1-ene is isomerised to cis-but-2-ene at temperatures greater than about 100” on a channel black (3 % of oxygen) but not on the same black without oxygen; (c) butadiene eliminates the “ -OH wave” and 2,2,4-trimethylpentane eliminates the “quinone wave’’ observed in polarographic examination of oxidised blacks; (4 the rate of binding increases with the concentration of olefin bonds in the polymer; (e) blacks with oxygen complexes react much more rapidly in polymer formation than if the oxygen is not present ; (f) spontaneous gelation accompanied by carbon-polymer bonding occurs when a styrene partially polymerised by oxygen is suddenly mixed with a black ; ( g ) diphenylguanidine and a number of other accelerators are preferentially adsorbed on oxygen complexes.One recent hypothesis27 of carbon-rubber binding postulates the presence of adjacent -OH and > C =O groups on the surface and suggests the transfer of a hydrogen ion from a hydroxyl group to a double bond of the elastomer leaving a negative charge on the carbon particle and creating a carbonium ion of the elastomer. This carbonium ion may then attack the aromatic ring and bind the elastomer to the carbon particle by a C-C bond or it may attack a quinone group and bind the elastomer by a C-0-C bond. If the binding does not come from the double bond it is suggested that the quinone group abstracts a hydrogen atom from the most substituted carbon of a saturated elastomer chain forming a hydroxyl group on the carbon particle and creating a free radical of the elastomer.The resulting elastomer radical would then attack the aromatic ring and establish a C-C bond to the particle by alkylation. Pigments About 2.5% of the total production of carbon black is used by manu- facturers of printing ink. The “long-flow” pigments28 are channel blacks 27 Hallum and Drushel J. Phys. Chern. 1958 62 110. rta Brogin Amer. Ink Maker 1952 30 30. SMITH CARBON-OXYGEN SURFACE COMPOUNDS 305 partly characterised as small soft particles with moderate areas and a “percent volatile” (weight loss on heating primarily as carbon oxides) of about 5-15 %. These pigments make an ink which is finely dispersed spreads rapidly distributes evenly and is non-wearing on plates.The large quantity of oxygen complexes retards drying and makes it necessary to add driers to these inks. Furnace pigment blacks have larger harder more abrasive particles with smaller areas and less “percent volatile” ; they are more easily handled. The surface concentration of oxygen com- plexes is comparable on the two types and they are the key to the stability of the pigment suspensions. These factors involving flow and suspension stability in inks must also be considered in the formulation of paints and varnishes. Friction Experimental evidence is against the popular view that the most important factor in the lubricating properties of graphite is the slippage permitted by the relatively weak van der Waals forces between internal cleavage planes.If any form of carbon (including diamond) is heated to 1200” in vacuo there is an enormous increase in friction at the carbon- carbon contact under slow-moving loads carried at room temperature in a vacuum. Subsequent addition of oxygen or water vapour at pressures of 0.1 mm. or less will immediately cause a great decrease in friction. It is now believed29 &hat the oxygen complexes reduce the friction between the individual crystallites rather than between cleavage planes and that on high-temperature evacuation the loss of oxygen from the crystallite surfaces leaves free valencies available to increase adhesion between neighbouring crystallites. Physical adsorption of gases produces an additional small decrease in friction.In instances of high-velocity friction (e.g. carbon brushes in motors) physical adsorption is tremendously important.30 Here simple evacuation at room temperature produces friction so high that seizure occurs and mechanical disintegration of the carbon proceeds rapidly; yet a partial pressure of 3 mm. of water vapour (or 300 mm. of oxygen) will reduce this wear to zero. Thus effective graphite lubrication at high speeds requires a layer of carbon-oxygen complexes (formed from either oxygen or water) to which is attached a physically adsorbed water layer. This physically adsorbed water determines the difference between high- and low-speed lubrication. Grateful acknowledgment is made to the donors of the Petroleum Research Fund administered by the American Chemical Society for support of this study. This Review was written at Cambridge University; discussions with the staff and research students of the Department of Colloid Science and use of the University libraries are also gratefully acknowledged. A comprehensive bibliography keyed to this article may be obtained from the Reviewer on request. gs Goodman and Rowe “ Chemisorption Proceedings of the Symposium held at Keele ” 1956 Butterworths 1957 p. 272. 30 Savage J. Appl. Physics 1948 19 1.

ISSN:0009-2681    

DOI:10.1039/QR9591300287

出版商:RSC    

年代:1959

数据来源: RSC

摘要:

CURRENT THEORIES OF SOLUTIONS OF By N. G. PARSONAGE and L. A. K. STAVELEY (THE INORGANIC CHEMISTRY LABORATORY OXFORD UNIVERSITY) THE purpose of this Review is to consider some of the recent theoretical and experimental studies of the connection between the thermodynamic properties of a mixture of two liquids and the intermolecular forces existing in it. For many years the usual point of departure in discussing the thermodynamic properties of a liquid mixture has been Raoult’s law. If the vapours are assumed for simplicity to behave as perfect gases then the mixture obeys Raoult’s law if the partial vapour pressure of each component is proportional to its mole fraction x. Thus for component 1 NON-ELECTROLYTES p = p,ox . . . . . . . . . (1) where p 1 is its partial vapour pressure and pl0 the vapour pressure of the pure component 1.Such a solution is described as ideal or perfect. Alter- natively conformity with Raoult’s law requires that the chemical potential p of any component in the liquid should have the dependence on its mole fraction x which is given for component 1 by the equation . . . . . . . p = plo + RTln x1 (2) where pl0 is the chemical potential of the pure liquid component 1 at the temperature Tand at a prescribed pressure. If it is remembered that the chemical potential p,’ of the same component in the vapour phase (in which it is assumed to behave as a perfect gas) is given by . . . . . . . p,’ = pl’O + RTlnp (3) and that for equilibrium . . . . . . . . . p = p1‘ (4) then the connection between equations (1) and (2) is apparent. It follows from equation (2) that if x1 moles of a liquid 1 and x ( = 1 - x1 ) moles of a liquid 2 are mixed at a constant temperature Tand at a constant pressure to form one mole of mixture the increase AGi in the Gibbs free energy G accompanying the formation of this ideal mixture is dGi = x,(plo + RT In x,) + x2(p20 + RTln x2) - xlplo - X ~ U ~ By applying (5) to the general thermodynamic relation .. . . . . . . . . . . = RT(x In x1 + x In x,) ( 5 ) 306 PARSONAGE AND STAVELEY SOLUTIONS OF NON-ELECTROLYTES 307 (which is one form of the Gibbs-Helmholtz equation) it follows that for an ideal solution the heat content. change AH accompanying the mixing process must be zero. Likewise from the relation . . . . . . . . A Y = [ a ( A G ) / a p ] ~ (7) the volume change d Y for the mixing process must also be zero.Finally since AH = 0 from the equation dG = d H - TAS and from equation (9 the entropy of mixing ASi for an ideal solution is given by ASi = -R(x,lnx + x,lnx,) . . . . . (8) In reality it is almost unknown for two liquids to mix without showing detectable departures from the criteria given above for an ideal solution. (Even mixtures of isotopes may not necessarily form an ideal so1ution.l) Systems which were once cited as examples of ideal solutions such as benzene-ethylene dichloride have been found on a more careful experi- mental examination to show measurable departures from perfect be- haviour. In other words if two liquids are mixed at constant pressure and temperature there is an evolution or absorption of heat and a volume change and the free energy and the entropy of mixing are not correctly expressed by equations ( 5 ) and (8).It has become customary to represent quantitatively the departure of a system from ideal behaviour by means of the so-called excess functions first introduced by Scatchard,2 and now usually denoted by a superscript E attached to the appropriate thermo- dynamic symbol. Thus if dG is the actual increase in Gibbs free energy on formation of one mole of mixture the excess Gibbs free energy C E is given by . . . . . . . . G E = A G - AGi (9) The heat absorbed on formation of one mole of mixture is HE and YE is the corresponding increase in volume where . . . . YE= v - x l v l o - x x Y ~ o (10) Y10 and V20 being respectively the molar volumes of the pFre components 1 and 2 at the prescribed constant temperature and pressure and V the actual volume of one mole of mixture.Alternatively a non-ideal solution may be defined as one for which the chemical potentials of the components are no longer given exactly by equations such as (2) but by such equations modified by the introduction of the activity coefficient y. Thus for component 1 in a non-ideal mixture p1 = plo + R T l n x y . . . . . . (11) It then follows from equations (9 (9) and (1 1) that Prigoghe “The Molecular Theory of Solutions” North HoIland Publ. Co. Amsterdam 1957. * Scatchard Chem. Rev. 1931,28,321. 308 QUARTERLY REVIEWS G E = RT(x,ln y1 + x21ny2) . . . (12) The experimental determination of G E is therefore in effect a matter of determining the activity coefficients of the two components in the mixture and different methods can be employed to do this.One commonly used is to find the composition of vapour in equilibrium with mixtures of known composition i.e. to establish both the dew-point and boiling-point curves of the liquid-vapour equilibrium diagram. However for the binary mixtures of condensed gases with which we shall be largely concerned it is more convenient to measure only the total vapour pressure as a function of the composition of the liquid mixture. Provided that information is available to allow for the imperfection of the vapour phase it is then possible to evaluate G E if some suitable type of relation between the activity coefficients and the liquid composition is assumed. Here use is made of expressions originally due to Margules. In pure liquid 1 y1 is unity and In y = 0.As the component 2 is added and x 2 increases from zero In y1 departs from 0. It is reasonable therefore to express the dependence of In y1 on x2 in the form In y1 = Alx2 -t B1xz2 + C,x,3 + . . . and similarly to write for ln y 2 In y z = A,x + B2x12 + Czxl3 + . . . (14) In y1 and In y 2 are however related by the Gibbs-Duhem equation x,dln yl + x,dln y2 = 0 . . . . . . (1 5 ) and if (13) and (14) are substituted in (15) and the differentiation per- formed (remembering that x1 + x 2 = 1 and dx = - dx2) it is then found that Al = A 2 = 0 and that certain relationships must exist between the coefficients of the remaining terms. In particular if the terms in xZ2 and x12 [which are now the first terms in (13) and (14)] are considered to be adequate to represent the dependence of y1 and y 2 on composition then B and B2 have to be equal (= B say) and in this event the excess free energy G E will b'e derivable from (1 2) by the simple expression G E T B R T x ~ x ~ = BRTxl(1 - XJ .. . . . - (16) A plot of G E against the mole fraction of either component will then be a symmetrical parabola with a maximum at xl = x 2 = +. This is often nearly true of many systems in which the departures from ideality are not too pronounced not only for GE but also for H E and Y E . If necessary allowance for asymmetry may be made by introducing into equation (16) terms in ( x - x,) ( x - x~)~ etc. e.g. G E == RTx~x,[B + P ( x ~ - ~3 + Q(x1 - ~3'1 . . . (17) where the coefficients P and Q can be related to the coefficients in equations PARSONAGE AND STAVELEY SOLUTIONS OF NON-ELECTROLYTES 309 (13) and (14).Although in explaining the criteria of an ideal solution we assumed for simplicity that the vapour phase was a perfect gas in practice it is usually important to allow for the imperfection of this phase. This requires a knowledge of the second virial coefficients B,, B22 B, at the temperature in question where B, and B, refer to the pure gases while B, concerns the interaction between a molecule of component 1 and one of component 2. Information about the virial coefficients of binary mixtures of vapours is still meagre. The excess volume V E can be determined from the density of mixtures of known composition or by direct measurement. Since G E and the heat of mixing HE are related by an equation similar to (6) HE can in principle be found from the temperature variation of GE(or of the activity coefficients).But since the temperature range over which the measurements can be made is usually quite limited it is better to determine HE by direct calori- metry. In the recent work on condensed gas mixtures this has required the development of calorimeters for measuring heats of mixing at low tem- perature~.~.~ It should be noted that strictly the values of the excess functions should relate to mixing carried out not only at constant temperature but also at constant pressure. If the vapour pressures of the pure components 1 and 2 are say -$ atm. and -1 atm. respectively and if the mixtures are formed and studied at the vapour pressure of the mixture itself then of course the pressure on the system is not constant throughout but varies from -t to -1 atm.Variations as small as this however have a negligible effect on the excess functions. We shall see that the theoretical expressions for these functions are usually formulated so as to be valid at p = 0 and values calculated from them can be compared without serious risk of error with the experimental values obtained under a pressure of an atmosphere or so. But non-constancy of the pressure cannot always be overlooked. Thus the extreme molecular simplicity of a mixture of monatomic sub- stances such as argon and krypton makes such a system a favourable one with which to test the theories we shall describe. But at the lowest tem- perature at which this system could be studied in the liquid state over the whole range of composition namely the triple-point of krypton the vapour pressure of krypton is -0-7 atm.while that of argon is -9 atm. Pressures such as these would not only make the experiments more difficult but in addition considerable pressure changes would accompany the mixing the effect of which could not be ignored. In the last decade several valuable books have been written about solutions. Two of outstanding importance are “The Solubility of Non- electrolytes” by Hildebrand and S ~ o t t ~ in which a wide variety of solubility J Pool and Staveley Trans. Faraday SOC. 1957 53 1186. 4 Jeener Rev. Sci. Instr. 1957 28 263. 6 Hildebrand and Scott “Solubility of Non-electrolytes” Reinhold New York 1930. 3 10 QUARTERLY REVIEWS data is correlated with notable success and “Mixtures” by GuggenheimYB which includes a particularly useful survey of the lattice-model approach to the problem of liquid mixtures.Our concern here however is to review that recent work which has attempted to account quantitatively for excess thermodynamic functions in the light of our knowledge of intermolecular energies. In this work Prigogine and his school and Kirkwood’ have played leading parts. Prigogine has co-ordinated his contributions in his book “The Molecular Theory of Solutions”,l and we acknowledge our indebtedness to this in the preparation of this Review. We must now briefly consider the intermolecular energy E of a pair of uncharged molecules as a function of their separation r both when the molecules are identical and when they are different. The general cause of attraction between molecules is the so-called dispersion effect which may be pictured as arising from the oscillation in phase of the electronic systems of the two molecules.This as London showed gives an energy term which for not too small values of r is proportional to r6. The repulsion between the molecules which becomes dominant at sufficiently close approach and which is due to the overlapping of the electronic clouds gives rise to an energy term which according to quantum-mechanical calculations may be approximately represented by an expression of the form (polynomial in r ) x e-br. This cumbersome expression is never in fact used but as the exponential term is the more important one the simplified expression Ae-*r is sometimes employed. But even this proves to be rather unwieldy and more frequently the energy of repulsion is represented by the purely empirical expression d/rn; n must then be found empirically and as there often proves to be some latitude in the acceptable values of n which usually fall between 9 and 12 the most common choice is to take n = 12.for algebraic convenience. This gives the following expression for E (the Lennard-Jones potential) A plot of E against r gives a curve like that in Fig. 1. Equation (18) can be re-written in terms of E* and r* where E* is the value of the minimum potential energy and r* the distance at which this minimum occurs giving and also in terms of E* and o where o is the intermolecular separation at which E = 0 when Guggenheim “Mixtures” Clarendon Press Oxford 1952. ‘I Salsburg and Kirkwood J. Chem. Phys. 1952,20 1538; 1953,21,2169. PARSONAGE AND STAVELEY SOLUTIONS OF NON-ELECTROLYTES 3 1 I FIG.1. Intermolecular energy ( E ) as a function of intermolecular separation (r) showing significance oj’quantities c* r* and u. It will be seen from the last two equations that for all substances for which the Lennard-Jones potential is an adequate representation of the inter- molecular potential energy this energy is a universal function of the inter- molecular separation the only quantities characteristic of each substance being the scale factors E* and r* (or E* and 0). This will later be seen to have very important consequences owing to its close connection with the theorem of corresponding states. In a mixture of components 1 and 2 we are concerned not only with the intermolecular potentials 1-1 and 2-2 but also with those for the unlike molecules 1 and 2.This is a crucial matter in the theory of solutions and here unfortunately we are on less certain ground. It is usually assumed that if the Lennard-Jones potential is valid for the interaction 1-1 and 2-2 it is also valid for the interaction 1-2 with values of c12* and r12* given by the following so-called combination rules . . . . . . . . . . E12* = (Ell* x E2X*)* (21) r12* = + r,,*) (22) . . . . . . . . . . The justification for these rules is empirical rather than theoretical. London’s theory of dispersion forces related the coefficient c of equation (18) to the polarisability of the molecules and to the characteristic fre- quencies appearing in the dispersion equation and showed that the relation between the coefficients c12 cll and c22 is in general c12<(cll x c2& the equality only holding as a limiting case.Nothing whatever can be predicted about the relationship between the corresponding coefficient d for the repulsive energy term since this is quite empirical. Prigogine,l and Rowlin- son and Townley,* have reviewed the evidence in support of the rules (21) * Rowlinson and Townley Trans. Faraday SOC. 1953 49 20. 312 QUARTERLY REVIEWS and (22). This evidence is not yet as comprehensive as one could wish and a further drawback is that most of it is derived from experimental studies of gas mixtures for which a property such as the second virial coefficient depends upon the three coefficients Bll B22 and BI2. B12 is therefore only part of the quantity which is experimentally determined and this clearly affects the accuracy with which E ~ ~ * and r12* can be evaluated.The chief tests to which the validity of relations (21) and (22) has been submitted are the following (1) If equations (19) and (20) are valid for the pure com- ponents 1 and 2 then these substances conform to the law of corresponding states and (r11*)3 and (r22*)3 will be proportional to the critical volumes Ylc and V2c. Similarly ell* and E ~ ~ * will be proportional to the critical temperatures T, and T2c. For such substances the reduced virial coefficient B/Vc will be a universal function of T/Tc. If equations (19) and (20) are valid for the 1-2 interaction B12/(r12*)3 should be a universal function of T/e12*. Introducing the combination rules (21) and (22) this means that if &2/[&(V1,* + V2,Q)I3 is plotted against T’,(Tlc x T2,)g the points should fall on the curve obtained by plotting B/V against T/T for the pure substance^.^ (2) Values of B12 have been obtained by measuring the volume change on mixing of the two gases and compared with the values calculated by using (21) and (22).1° (3) Similar use has been made of experimental data on transport properties such as viscosity thermal diffusion and diffusion.(A study of interdiffusion coefficients has the advantage that they depend only upon the 1-2 interaction in the gas mixture.) As a result of his survey of the available evidence Prigogine concludes that the combination rules (21) and (22) “seem to be a fair approximation for non-polar molecules”. It should be noted however that if one of the molecules is polar we cannot strictly expect (21) to hold good.To test theories of solution it is possible to select for experimental study a number of binary systems where the molecules are sufficiently small and also non- polar. But even a diatomic molecule will have a quadrupole moment. Thus in the system argon-nitrogen in addition to the dispersion forces there will be quadrupole-quadrupole forces between the nitrogen molecules giving rise to a potential energy term which is proportional to r-l0 (and also quadrupole-induced dipole forces giving an energy term dependent on r8 though this is probably less important). Here too equation (21) cannot be strictly applicable. It is now realised that quadrupole forces can have a more important influence on the physical properties of substances with simple non-polar molecules than was at one time supposed.ll The Lennard-Jones potential (I 8) (or indeed any other tractable func- tion in which the intermolecular energy is expressed in terms of the distance r between the molecular centres) can only be expected to apply to small Guggenheim and McGlashan Proc.Roy. SOC. 1951 A 206,448. lo Michels and Boerboom BuZZ. Suc. chim. belges 1953 62 119. l1 Buckingham Quart. Rev. 1959 13 183. PARSONAGE AND STAVELEY SOLUTIONS OF NON-ELECTROLYTES 3 13 molecules. Indeed even for such simple molecules as carbon tetrachloride and sulphur hexafluoride it has been suggested that a better agreement with the thermodynamic properties of the vapours would be obtained by considering the centres of force to be situated on the periphery of the molecules. A fair test of theories based on the Lencard-Jones potential should therefore be made on binary mixtures of small preferably non- polar molecules that is molecules of substances which will probably be gaseous at ordinary temperatures.This is the reason for the recent experi- mental studies of the thermodynamic properties of binary mixtures of condensed gases. The theories which we shall consider employ the methods of statistical mechanics. If on the basis of a suitable model it is possible to evaluate the partition function Q the thermodynamic properties of the system may be calculated from general relations such as the following F = Helmholtz free energy = - k Tln Q s = - (g)v = LIn Q + kT ( a - ;)e)v . . . . . . . (23) (24) . . . . . . . . . . . . . . (25) The partition function Q is in general given by the expression .. . . . . . . . . . Q = Cexp - E,/kT (27) 9. where E is the energy of the rth level available to the molecules. In the theories we shall consider it is supposed that Q can be written as Qint x Qtr where Qtr is the translational partition function and Qint that associated with the other degrees of freedom such as the internal vibrations of the molecules and their rotation. It is further assumed that when a mixture is formed the internal partition functions Qint remain unchanged. (This should be noted; it implies that any intramolecular vibrations are not affected by mixing which is probably a very good approximation and that the rotational movement of the molecules is the same in the mixture as in the pure liquids. The validity of this second assumption may sometimes have to be reconsidered.) Interest therefore centres on the quantity Qt, which for a classical system may be written in the form (2rrmkT)f * Qt = [-jp--] Q,y + .I . . I (28) 3 14 QUARTERLY REVIEWS 1 - U/kT where eU= N ! - J . . . . J e dx dy dz1. . . . dx dyfi dZN (29) U is the potential energy of all the molecules in the system and the integra- tion extends over its whole volume. QU is called the configurational partition function and it is with the determination of this that theories of solutions are essentially concerned since once this has been done the configurational contributions to the thermodynamic properties of the system can be derived by using equations such as (23) to (26). Throughout this Review the related quantities are the configurational contributions to the thermodynamic functions the non- configurational parts making no contribution to the excess properties in which we are ultimately interested.The correlation of the thermodynamic properties of a fluid with the intermolecular energy parameters was first attempted for a single com- ponent by Lennard-Jones and Devonshire.12 The model for the fluid which they employed the so-called cell model has played an important part in modern statistical theories of solutions and the significance of these should be more readily perceived if we first briefly outline the Lennard-Jones and Devonshire approach to the problem of the single component fluid. It is first supposed that any one molecule moves in a cell formed by its nearest neighbours the average volume of which is V/N where Y is the total volume of the system of N molecules.The potential field in which this molecule moves is determined by its interaction with its neighbours. Strictly speaking this field varies with time so it is replaced by the average field in which the molecule would move if each of its neighbours was at rest at the centre of its own cell. (Only nearest-neighbour interactions were considered at first but the treatment can be extended to include inter- actions with more distant molecules.) On this basis and by using equation (19) for the potential energy of a single pair of molecules and assuming that the total potential energy of one molecule is the sum of its energy with respect to each of its neighbours in turn the following expression is obtained for w(r) the potential energy of the molecule when its centre is at a distance r from the centre of its cell w(0) is the energy of the molecule when at the centre of its cell (r = 0).A = z J E * ~ where z is the number of nearest neighbours and I€*] the depth of the trough in the intermolecular potential energy curve (Le. [€*I is the numerical value of the E* of Fig. 1 which is a negative quantity). V* = N ( Y * ) ~ / ~ where Y* is the intermolecular separation at the minimum le tennard-Jones and Devonshire Pmc. Roy. SOC. 1937 A 163,53; 1938 A 165,l. PARSONAGE AND STAVELEY SOLUTIONS OF NON-ELECTROLYTES 3 15 in the potential energy curve and y is a numerical factor depending on the type of molecular packing assumed. (For a face-centred cubic lattice y = 4 2 . ) y = r2/a2 where a is the average distance between nearest neighbour molecules; Z(y) and m(y) are functions given by the equations I(y) = (1 + 12y + 2 5 .2 ~ ~ ~ + 12y3 + f)(1 - y)-lo - 1 (31) m(y) = (1 +y)(l - 9 ) - 4 - 1 . . . . . (32) The configurational partition function for a single molecule referred to an energy zero with the molecule at the centre of its cell which we will denote as Qu' to distinguish it from the corresponding function QU of equations (28) and (29) which relates to all the molecules is given by the equation where U(r) = W(Y) - w(0). Changing to polar co-ordinates we have The integration in (34) is then carried out between the limits r - 0 and Y = af2 (= half the cell radius). (The justification for this arbitrary choice of the. upper limit is that small values of Y i.e. < the chief contributions to the integral come from a/2).This gives Qu' = 277a3g . . . . . . . (35) If it is supposed that the molecules are not localised but that the whole volume V is available to them then the relation between Quf and Qv is (36) Qu= ~ . . . . . . . ( QulN)N - Ntl~(0)/2kT N ! since Nw(0)/2kTis the potential energy of the whole system when all the molecules are at the centres of their cells referred to the free molecules as energy zero. Thus equations (31) (32) ( 3 9 and (36) combine to give the complete function for the fluid and its thermodynamic properties follow from relations (23) to (26). In particular it will be noticed that use of equation (26) establishes a relation between p V and T i.e. an equation of state which in principle enables the volume of the system at a given temperature and pressure to be estimated.It is in this way that the problem of VE for the liquid mixture is approached. 316 QUARTERLY REVIEWS The cell model has been extended to mixtures by Prigogine and co- w o r k e r ~ . ~ ~ In their work random mixing of the components is assumed this supposition being justified by the calculations of Rushbrooke14 and of Prigogine and Garikian,13 who have shown that at least for molecules of the same size there is only a very small effect on the thermodynamic excess functions due to the departures from random mixing which occur in real solutions. This assumption is embodied in the expressions for the mean potential fields in which a molecule of each type moves. Attention was first directed to the simplest case that where rll* = r22* for which it is reasonable to assume that all the cells are of the same size.The potential field for a molecule of component 1 in such a mixture is then obtained by replacing the parameter ell* in the Lennard-Jones and Devonshire expression for pure liquid 1 by xlell* + x2e12* the potential function for a molecule of component 2 being obtained in a similar way. Unfortunately the complicated nature of the expressions which result obscures their dependence on the mole fractions. Because of this Prigogine and Mathot have used the so-called smoothed potential model in which a square well is substituted for the full Lennard-Jones and Devonshire type of potential. In this approximation to the proper cell model the depth of the potential well is that of the minimum in the Lennard-Jones and Devonshire potential I t I I I I I I I t i I 1 I I I I I I I I I I I I I I I I I I I I FIG.2. Comparison of the Lennard-Jones and Devonshire potential (full curve) witlz so-called smoothed potential (broken line). The potential energy w(r) of the molecule as a function of its distance ( r ) from the centre of its cell. l3 Prigogine and Garikian Physica 1950 16 239; PrigDgine and Mathot J. Chem. Phys, 1952,20,49; Prigogine and Bellemans J. Cliern. Phys. 1953,21 561. l4 Rushbrooke Proc. Roy. Soc. 1938 A 166 296, PARSONAGE AND STAVELEY SOLUTIONS OF NON-ELECTROLYTES 3 17 "(O)] and the diameter of the well is the radius of the cell (a) minus the molecular diameter (a) for interactions between the molecule and the wall of the cell (Fig. 2). On the smoothed potential model the configurational partition function for a single molecule of component I (Q'u,) is given by Ha-d Q'U1 = 4 ~ 1 r2 dr = &r(a - a)3 .. . . . . . . (37) 0 referred to an energy zero with the particle at the centre of its cell. Since the potential energy of the system when each molecule is at the centre of its cell is +N[x,w,(O) + x2w2(0)] referred to the completely free molecules as energy zero (where w,(O) = the potential energy of a molecule of com- ponent 1 at the centre of its cell referred to the completely free molecule as energy zero and w2(0) = the potential energy of a molecule of component 2 at the centre of its cell referred to the completely free molecule as energy zero) the configurational partition function for the whole system is w,(O) is easily shown to be given by or if interactions with second and third shell neighbours are taken into account ~1(0) = ~ ( ~ 1 ~ 1 1 + ~ 2 ~ 1 2 ) Thus w1(O) and likewise w2(0) can be expressed in terms of a which is proportional to Vg.It can be seen therefore that (38) gives Q as a func- tion of V T and the parameters characteristic of the inter-molecular interactions. By using (23) to (26) the thermodynamic functions are ob- tained in terms of the same variables. When in addition the usual combination rules for the inter-molecular parameters (21) and (22) are assumed this theory leads to the rather surprising prediction of a positive excess free energy coupled with a negative excess volume. On previous theories a positive excess volume would be expected. The possibility of contractions on mixing in simple systems was confirmed by experiments on the mixtures CClp-C(CH3)415 and CO-CHp.16 l5 Mathot and Desmyter J.Chenz. Phys. 1953 21 782. l* Mathot Staveley Young and Parsonage Tram. Faraday Soc. 1956,52,1488. 318 QUAKTEKLY REVIEWS A qualitative explanation can be given of this contraction which is predicted and observed for mixtures of molecules of equal size whose interaction energy contains only terms due to dispersion and short-range repulsion forces. Ignoring zero-point energy the inter-molecular separation of a 1-1 pair at O'K would be rll*. However as the temperature is raised the mean inter-molecular distance increases owing to the asymmetry of the potential well about r = Y*. (This is in fact partly responsible for the thermal expansion of the liquid.) If the inter-molecular potential energy is represented by the Lennard-Jones 6-12 equation it can readily be shown that if the energy of vibration of the pair of molecules relative to each other is A then the mean distance apart of the molecules (rll) is For small values of A/cll we may use the approximate formula Remembering that for molecules of the same size rlZ* = rll* = rZ2* and putting c12* = ( E ~ ~ * .E ~ ~ * ) * it then follows that 2 ~ ~ < r ~ ~ + r22 i.e. a contraction on mixing would be expected. Indeed if the assumption 2el2* = ell* + E ~ ~ * is made r12 is even smaller. The extension of the theory to the case where rll* # r22* was made by Prigogine and Bellemans. They assumed that cells of two sizes existed in the liquid mixture cells containing molecules of species 1 being different in size from those containing molecules of species 2.The ratio of the diameters of the cells was then chosen so as to minimise the free energy of the mixture. Since the existence of cells of different sizes is really in- compatible with these cells' having the simple close-packed structures of the pure liquids the Prigogine and Bellemans treatment can only be a valid approximation if the sizes of the two species of molecule do not differ too much. Otherwise the structure of the mixture might be quite different from that of the pure liquids. One important point was however established by this work namely that a relatively large positive excess volume would result from quite a small difference in the molecular size leading us to expect that the negative excess volumes which should occur for mixtures of molecules of the same size would be only rarely observed.In the cell theories of mixtures dealt with so far discrepancies between the experimental and the theoretical values of the excess thermodynamic functions could arise either from the failure of the model to lead to the correct values of the functions for any liquid even for the pure com- ponents or from errors introduced by the assumptions made about the molecular movement and relationships within the mixture. As a result of this a number of theories have been developed based on the theorem - - - PARSONAGE AND STAVELEY SOLUI IONS OF NON-ELECTROLYTES 319 of corresponding states which are liable only to the second type of error. In these theories the properties of the mixture and of the pure componznts are expressed in terms of those of a suitable reference liquid and as a consequence the values given for the pure components are necessarily correct provided only that they obey the theorem of corresponding states.For the derivation of these theories nf mixtures the form of the theorem of corresponding states due to Pitzerl' is used. He showed that if the inter- molecular potential energy of a pair of molecules ( E ) at a distance apart (r) could be expressed in the form where cj5 is a universal function and E* and r* are two parameters character- istic of the interacting molecules (and it will be noted that the Lennard- Jones 6-12 potential conforms to this equation) then the configurational contribution to the bulk thermodynamic functions could also be expressed in terms of universal functions of W/E* and T / T * .Thus for the Helmholtz free energy we would have E = E*&r/r*) . . . . . . . . (43) (44) . . . . . where $ is a universal function. the equations We may define the reduced temperature (?) and reduced volume (;> by n. ?= kT/E* . . . . (45) v* = ~ / r * ~ . . . . 446) where e = the volume per molecule = r3/y y being a constant character- istic of the liquid strccture. It follows from equation (44) thst F/eTis a universal function of T' and G which function we denote as (F/RT)(T,v"). The important step is now taken upon which depends the form of the expressions for the excess functions the Helmholtz frce enErgy is related to the thermodynamic proper_ties of a reference liquid (T = To w" = u",,) by a Taylor series expansion in l/Tand 5 Pitzer J. Chin. Phys. 1939 7 583.2 320 QUARTERLY REVIEWS The coefficients of the series expansion are readily expressible in terms of the thermodynamic functions of the reference substance. Thus for a given liquid the reference liquid the coefficient of the third term is {*j-)? = (,,,> E*{ a(l/T) 1 - E* a R where E is the internal energy of the reference liquid. If the two liquids are at the same temperature and occupy the same volume this is tantamount to an expansion in terms of the differences of their force parameters E* - E,* and r* - ro*. The functions S and p can be obtained by differentiation of I; with respect to T and V respectively. Of considerable interest are the functions G H and V which are the quantities handled experimentally. These quantities can be obtained provided the equation of state-is first solved.Alternatively GIRT can be expanded in terms of p” and 1/T but here the problem is more complicated than the one dealt with above because of the nature of p” [ = P ( ~ * ) ~ / E * ] . Longuet-HigginslS was the first to employ a treatment of this type. In his theory of conformal solutions he considered only terms in the expres- sions for the excess functions which were proportional to ( E ~ ~ * - ell*) and (rZ2* - rll*) where these represent the differences between the force parameters of the two pure substances. For a binary solution the results which he obtained may be written as . . . * (48) a(FtRT) a(F/RT) k W R T ) -5 E VE - S E HE (dQo/dT - R) = T(dQo/dT) - Qo - G E RT - Qo JQPo - Tao) = x1x& . . . I . . . . . . (49) where Q is the latent heat of conversion of a mole of the reference liquid at temperature T and pressure p to its vapour at T and p cco and Po are respectively the coefficient of thermal expansion and the isothermal compressibility of the reference substance and S12 is an adjustable para- meter characteristic of the mixture.The insertion of typical numerical values into the terms in the deno- minators of equation (49) shows that all of the denominators are negative except (dQ,/dT) - R which is positive. It can be seen therefore that in this form the theory could not explain the results on the CO-CH4 and CC1,-C(CH,) systems for each of which GE and YE have opposite signs. Prigogine and his co-workers19 and Scott2 have therefore considered also higher terms in the series for the excess functions which are proportional to ( E ~ ~ * - (r22* - r,l*)2 and (eZ2* - ell*) (rZ2* - rll*) if e12* and r12* are eliminated by using (21) and (22).In Prigogine’s so-called average potential model it is assumed that all cells are of the same size whether Longuet-Higgins Proc. Roy. Soc. 1951 A 205 247. lo Prigogine Bellemans and Englert-Chwoles J. Chern. Phys. 1956 24 518. 2o Scott J. Chem. Phys. 1956 25 193. PARSONAGE AND STAVELEY SOLUTIONS OF NON-ELECTROLYTES 32 1 occupied by a molecule of component 1 or by one of component 2. Hence the volume ZI to be used in evaluating the reduced volume of the liquid is just the molecular volume of the mixture. The usual assumptions of random mixing and additivity of the inter-molecular potential energies are also made which lead to the expression for the inter-molecular potential energy of a pair of molecules in the pure liquid e*((r*/r)12 - 2(r*/r)6} being replaced by + X2€22*{(-) r22* l2 - 2 (+*y] * .This may be written as ~m*{(r,*/r)12 - 2 (r,*/r)6) . . . . where { X12Ell*(rll*)6 + 2X1X2E12*(r12*)6 + -qJ2c22*(r22*)6}2 Ern* = x12Ell*(r11*)12 + ~ x ~ x ~ E ~ ~ * ( ~ ~ 2*)12 + x ~ ~ E ~ ~ * ( ~ ~ ~ * and Thus these expressions give the values of E* and r* to be used mixture. Prigogine has also carried out a more refined treatment in which the cells occupied by molecules of types 1 and 2 are no longer assumed to be of the same size. The sizes of the two types of cell cannot now be expressed in terms of the molar volume of the mixture alone an additional equation now being necessary since there are two quantities to be chosen. This extra equation is provided by requiring that the cell sizes must be chosen so as to minimise the free energy as was done in the Prigogine and Belle- mans theory.Since two types of cell are present it is necessary to choose average values of E* and r* for each type of cell. By the same process as before that is assuming additivity of the potentials and random mixing the following expressions are derived where cl* and rl* are the average parameters to be used for the cells con- taining molecules of component 1. Similar expressions are obtained for the 322 QUARTERLY REVIEWS corresponding parameters for cells containing molecules of component 2. The partial molar thermodynamic functions for component 1 can then be expressed in terms of el* rl* and the volume of a type 1 cell and the corresponding functions for component 2 can be calculated in a similar way.If the partial molar Helmholtz free energies are found to be Fl and Fn respectively then the total Helmholtz free energy of the mixture is - F=xlF1 +x2F2 . . Scott has also considered a third type of approach in which the mean distance apart of unlike molecules is allowed to adjust itself independently of the mean separations of 1-1 and 2-2 pairs. But because of the strain on the lattice brought about by the presence of in effect cells of three different sizes Scott himself considers this to be not a good representation of real liquids. He has also compared the theories previously discussed in which only one (one-liquid model) or two (two-liquid model) sizes of cell are permitted and he has come to the conclusion that the latter should be preferred since the one-liquid model overestimates the effect of difference in size of the molecules by not allowing any adjustment of the lattice which might lower the free energy.In the remainder of this Review the two-liquid approximations will be meant when reference is made to the Prigogine corresponding-states treatment or the Scott treatment. As with the cell theories the corresponding-states treatments are limited in their validity to mixtrxes the components of which have molecules of not too widely different size. Byers-Brown21 has examined very closely the errors which might be introduced into the results of a corresponding-states theory by departures from random mixing and has concluded that these errors might well be large for mixtures of molecules of appreciably different size.He has also pointed out that previous workers who have discussed this problem and reached the conclusion that non-random mixing is un- important have always considered systems of molecules of equal size whilst it is only in the case of mixtures of molecules of different sizes that non-random mixing would be expected to be important. A comparison of the thermodynamic properties of several binary mixtures of simple substances with the values predicted for these quantities on the refined version of Prigogine’s corresponding-states theory is made in the Table where the figures given are for the equimolar mixture. For some systems the values obtained from Scott’s two-liquid model are also included. Bearing in mind that the effects concerned are relatively small one may note that for the majority of systems the theoretical values are of the correct sign and order of magnitude.Particularly remarkable are the successful predictions of the volume contractions found for the systems CO-CH, 02-N2 and CCl,C(CH,), all of which approximate quite closely to the simple case of a mixture of molecules of equal size. Indeed a1 Byers-Brown Phil. Trans. 1957 250 221. PARSONAGE AND STAVELEY SOLUTIONS OF NON-ELECTROLYTES 323 CO-CHp (~cH** = 4.29 UCO* = 4-21) one of the few systems for which a complete set of experimental data is available also shows quite good agreement between theory and experiment for the other excess functions. The same is true of the system CC14-C(CH3)4. Krypton and methane molecules are also similar in size ( r ~ ~ * = 4-04 r c ~ ~ * = 4.29) and it is unfortunate therefore that only G E has been determined.It will be noticed that the theoretically predicted values are frequently very sensitive to the choice of reference substance. This would not be so if the reference substances conformed exactly to the theorem of correspmd- ing states and a sufficiently large number of terms was retained in the expansion (47). The series however is not very rapidly convergent and the number of terms which can be evaluated is severely limited by the lack of accurate data for the higher derivatives with respect to temperature and volume of the ordinary thermodynamic functions. The most striking case of the effect of the choice of reference substance on the value predicted is that of VE for the system Ar-N, where the change of reference substance leads to a change in the sign of the effect to be expected.The system Ar-CH4 appears to provide a good test of the adequacy of the theory to predict effects due to size differences of the molecules since complications due to the polar character of the components or their lack of spherical symmetry should be almost entirely absent. The comparison between theory and experiment for this system is however very disappoint- ing the experimental values being all very much smaller than those pre- dicted. Here it may well be that the difference in size of the component molecules is too large to justify the supposition that the mixtures and the pure liquids all have the same structure ( r ~ ~ * = 3.82 ~ c H * * = 4.29 r ~ r - t ~ * / r ~ ~ * = 1.12).If modifications to the structure of the mixture occur because of the difference in size of the components the free energy of the mixture and hence GE would be reduced and it is also likely that V E would become smaller. Both of these expectations are in agreement with the data. At first sight it might seem surprising that the real value of S E is less than is predicted in spite of the increase in disorder resulting from the distortions of the liquid structure. However it must be borne in mind that any con- traction in volume brought about by the distortions would lead to a large drop in the thermal entropy of the molecules and this might well be more than sufficient to counterbalance the increase in entropy previously mentioned. Appreciable effects may also arise from non-random mixing in the system.As already mentioned such effects have been carefully con- sidered by Byers-Brown. The contrast between theory and experiment for the system CO-N seems particularly significant. The close similarity b2tween these iso- electronic components (which is reflected in the small differences in many of the physical properties of the pure substances) would lead us to expect that their solutions would be almost ideal as is indeed quantitatively predicted by the theories under discussion. The actual values of G E and Y E 324 QUARTERLY REVIEWS are however much larger and are comparable in magnitude with those of many of the other systems in the Table. Admittedly carbon monoxide has a dipole moment but it is far too small to account for the discrepancy and the only feasible reason for this seems to be that carbon monoxide and nitrogen have considerably different quadrupole moments (& = 1-7 1 x e.s.u.).Calculations by one of us (N.G.P.) and Dr. A. D. Buckingham show that effects of the size observed could arise from quadrupole-quadrupole interactions. It is not surprising that the results for the Ar-0 system fit the theory well since the constituent molecules are similar in size (rg ,*/r~* = 1-01 6) and the quadrupole moment of oxygen is very small (<+@N~). But all the other systems containing diatomic molecules are more complicated in that the errors arising both from differences in size and also from neglected quadrupole effects are involved. The two causes of error seem to lead to opposing shifts in the values of the excess functions.(Yet another possible complication is that the rotational partition functions of the diatomic molecules are not in fact independent of their environment.) When considering solutions containing carbon tetrafluoride it must be remembered that although it has zero dipole and quadrupole moments (on account of its symmetry) it must have large higher electric moments because of the polarity of the individual C-F bonds. It seems probable that such higher moment interactions would cause modifications to the excess properties which are of the same sign as those produced by the quadrupolar interactions already mentioned. Comparing therefore the systems CI-1,-CF and Kr-CF, the fact that in the latter G E (expt.) falls short of the theoretical figure whilst in CH,-CF the experimental value exceeds that predicted is consistent with the greater disparity in molecular size in the Kr-CF system.It would seem to be highly desirable that theories capable of dealing with orientational forces should be developed. Some work has been done Comparison of the predicted and observed values of the excess thermodynamic functions of eqzcimolar mixtures of simple substances. e.s.u. 0~~ = 1.27 x GE = Excess Gibbs free energy (cals./mole of mixture) HE = Excess heat content (cals./mole of mixture) SE = Excess entropy (cals./deg.mole of mixture) V E = Excess volume (c.c./mole of mixture) Expt. = Experimental value Theory Po() = Value predicted from Prigogine's refined corresponding states treatment using X as reference liquid Theory S = Value predicted from Scott's two-liquid model System T('K) G E HE TSE V E Theory P(C0) 32 27 -5 -1.2 Theory P(CH,) 19 16 -3 -0.1 CO-CH4 90.7" Expt.2816 263 -2 -0.3016 PARSONAGE AND STAVELEY SOLUTIONS OF NON-ELECTROLYTES 325 System T('K) Ar-CH 91 Ar-CO 83.8 Ar-0 83.8 OZ-N 83.8 A-N 83.8 CO-N 83.8 Kr-CH 115.5 CH,-CF 110 Kr-CF 117.1 CC1,- 273 Expt. Theory P(Ar) Theory P(CH4) Expt. Theory P(Ar) Theory P(C0) Expt. Theory P(Ar) Theory P(0,) Expt. Theory P(0,) Theory P(N,) Expt. Theory P(Ar) Theory P(N,) Theory S Expt. Theory P(C0) Theory P(N,) Expt. Theory P(CH,) Expt. Theory P(CH,) Theory P(CF4) Theory S Expt. Theory P(Kr) Theory P(CF4) Expt. GE 1 822 49 53 gz4 7 7 8.2,' 30 28 5*428 0-9 0.6 1429 7 8629 55 74 113 7529 115 126 7615 C(CH3)4 Theory P[C(CH3),] 84 Theory P(CC1,) 75 Ar-Kr* 83 Expt. 4532 Theory P 41 Theory S 27 Theory P 139 Kr-Xe* 105 Expt.8132 * Solid solutions. 22 Mathot personal communication. 23 Shields and Staveley unpublished work. 24 Pool Shields and Staveley unpublished work. * 5 Pool and Stavelev. Nature. 1957. 180. 11 18. HE TSE V E 2522 7 +0*18 71 22 +1.02 +090 +0.30 +0*17 1 42s 5 +0*1425 11 4 +0.13 11 4 +0.10 -0.3 126 - 0.09 -0.4 1 -0.1826 +0x2 -0.32 +O* 17 +0.13as +0.1023 +0*02 +0-88ao +3.7 7531 - 1 -0.515 71 -13 -2.2 -0.5 803 35 31 -10 30 3 17032 89 149 10 26 Pool and Stavelei; unpublished work. 27 Herrington Saville and Staveley unpublished work. 28 Herrington and Staveley unpublished work. 29 Thorp and Scott J. Phys. Chem. 1956 60 670. 30 Croll and Scott J. Phys. Chem. 1958 62 954. 31 Englert-Chwoles J . Chem. Phys. 1952 20 925; 1955 23 1168. 38 Walling and Halsey J . Phys. Chem. 1958 62 752.326 QUARTERLY REVIEWS in this field notably by P ~ p l e ~ ~ by Rowlinson and his co-worker~,~~ and by Prig0gine.l However all of these treatments consider the orientationally dependent part of the energy of interaction to be a small perturbation and from what we have seen of the results for the system CO-N2 in particular this would not appsar to be so. The development of more precise theories valid for strong orientationally dependent forces is hindered by the co- operative nature of these forces. Thm the probability of any partichlar relative orientation of two molecules is influenced by the orientations of their neighbours which in turn are influenced by the orientations of more remote molecules. Perhaps the best possibility for advance in this field lies in the work of Wood and Parker35 using high-speed electronic com- puters.So far this work has reached the stage of deriving by Monte Carlo methods values for the thmnodynamic functions of three-dimen- sional arrays of molecules interacting according to the Lennard-Jones 6-12 law of force. Systems of 108 molecules are normally studied but it has been shown that the results obtained for a 32-molecule system differ only insignificantly from those for the larger system. Thus it would appear that the sizes of the samples are quite adequate. The theoretical difficulties of treating any system containing molecules for which the fields of force are non-spherical makes particularly desirable the determination of the properties of mixtures of the inert gases. Un- fortunately measurements on the Ar-Kr system (and similarly on the Kr-Xe system) would be beset by considerable difficulties for the reasons given earlier.It is interesting therefore to have the data for the solid solutions Ar-Kr and Kr-Xe. The results of this work by Halsey and his co-workers are included in the Table. It was of course not possible to measure the heats of formation of the solid solution directly as is normally done for liquid mixtures. Instead 4 P was derived from values of G E obtained over a range of temperature [equation (6)]. Since this range was only 26" for Ar-Kr and 30" for the Kr-Xe system this method is less satisfactory than the direct method. Halsey et al. give 12 cals. mole-1 as the uncertainty in the heat of mixing. In all these experiments great care was exercised to ensure that true equilibrium was reached a waiting period of -24 hours being allowed before measurements were made. Presumably therefore the quantity most accurately known is G E . For Ar-Kr there is good agreement between the observed and the calculated values of GE but for Kr-Xe the agreement is only moderate. Finally the large differences between the observed and the calculated entropy values are a reminder that the swcessful prediction of excess entropies is an exacting test for any theory of solutions. Is Pople Proc. Roy. SOC. 1954 A 223 498. 84 Cook and Roulinson Proc. Roy. SOC. 1953 A 218 405; Rowliilson and Sutton 36 Wood and Parker J. Chew. Phys. 1957,27 720. Proc. Roy. Soc. 1955 A 229 271.

ISSN:0009-2681    

DOI:10.1039/QR9591300306

出版商:RSC    

年代:1959

数据来源: RSC

摘要:

LIQUID-LIQUID EXTRACTION IN INORGANIC CHEMISTRY By F. S. MARTIN and R. 3. W. HOLT [UNITED KINGDOM ATOMIC ENERGY AUTHORITY DEVELOPMENT AND 1. Introduction ENGINEERING GROUP (R. & D. BRANCH) SPRINGFIELDS] LIQUID-LIQUID extraction most frequently involves an aqueous phase an immiscible organic liquid phase and one or more solutes in a system which is commonly operated under prevailing atmospheric conhtions. There are a few recent developments in which the immiscible phases are liquid metals; these represent a threshold of future work and expansion and as such they merit brief mention in this Review. The analytical implications of liquid-liquid extraction are extensive but since these have previously been comprehensively reviewed by Irving1 in these Reviews no more than cursory attention will be paid to them here.2. Historical In 1872 Berthelot2 stated that when a third substance is present in a system of two immiscible liquids it distributes itself between them in a definite manner if it is soluble in both of them. If c1 and c are the con- centrations of the solute in phases 1 and 2 at equilibrium at constant temperature then CJC = constant (the distribution or partition coefficient) . (1) This is the simplest form of the distribution or partition law. Equation (1) may be derived from the equality of the chemical potential (p) of the solute in two phases at equilibrium i.e. p1 = p2. Since p = po + RTln a (where p0 = const. and a = activity) then If the solutions obey Henry’s or Raoult’s law activities can be replaced in practice by mole fractions and then al/a2 = const.(at constant T ) . . . . . . (2) x1/x2 = const. . . . . . . . . (3) For dilute solutions the ratio of mole fractions approximates to that of concentrations (in molarities or molalities) ; consequently c,/c = constant (=equation 1). Hence equation (1) will only be strictly true for idea€ dilute solutions; however many substances obey the law in this form for example iodine Trving Quart. Rev. 1951 5 200. * Berthelot and Jungfieisch Ann. Chim. Phys. 1872 26 396. 327 328 QUARTERLY REVIEWS in the carbon tetrachloride-water pair sulphur dioxide in water-chloro- form and mercuric chloride in water-benzene. According to N e r n ~ t ~ the distribution law applies only to those species common to both phases. If a represents the fraction of solute which under- goes association or dissociation in a phase equation (1) becomes c (1 - al) /cz(l-a2) = constant .. . . . . (4) The inexactitude of equation (4) is greater than that of (1) since dissociable or associable solutes depart from ideal behaviour. The use of activity coefficients in modern treatments4 results in improved forms of the distribu- tion expression. In spite of the inexactitude of equation (4) it has formed the basis of several studies of association or dissociation e.g. the distribution of benzoic acid between water and benzene a study of which has shown that in benzene benzoic acid is largely dimeri~ed.~ However similar studies in other systems may lead to false conclusions. For instance Anderson and Yost6 showed that the distribution of osmium tetroxide between carbon tetrachloride and water was consistent with the hypothesis that the solute exists in the organic phase as the tetramer (OsO,),.Later studies’ showed that a similar hypothesis was not tenable in the case of the very similar ruthenium tetroxide ; Hildebrand and Scott8 suggested that differences in internal pressure of solute and solvent were more likely to account for the experimental results than the existence of tetramers in this case. Experimental details of studies of this type will be found in most text books of physical chemistry. 3. Classification of solvent extraction (a) Aqueous Phase.-Of recent years solvent extraction (this term is fre- quently used instead of liquid-liquid extraction for aqueous-organic pairs) has been used for a variety of studies and practical applications; these are based on various reactions between solute species or between solute and solvent (Table 1).This classification is not rigid or exact. For example uranyl nitrate may be extracted under different conditions as the dinitrato-complex U02(N03) 2 or the trinitrato-anionic one H[UO,(NO,),]; in neither case can the formation of these species be said to result solely from ion association [e.g. (U0,++)(2N03-)] or solely from simple co-ordination. The extraction of a species usually depends on its conversion into a a Nernst Z. phys. Chem. 1891 8 110. McKay Chem. and Ind. 1954 1549. Hendrixson Z. anorg. Chem. 1896 13 73. Anderson and Yost. J. Amer. Chem. SOC.. 1938. 60. 1822. I Martin J. 1954 2564. Hildebrand and Scott “Solubility of Non-Electrolytes,” Reinhold Publ. Corp. New York 1950 p. 220. MARTIN AND HOLT LIQUID-LIQUID EXTRACTION 329 Reactions involved compound capable of solution in the organic phase e.g.anionic ion- association complexes may form oxonium salts with ethereal solvents or the metal may be associated with hydrophobic complexes as in many chelates. In a few cases the inorganic compound seems to dissolve as such in the organic phase e.g. ruthenium tetroxide in carbon tetrachloride; in Examp le TABLE 1. Types of extractable species involving inorganic solutes in aqueous phases. 1. Ion association with the metal in either the cation or the anion 2. Simple co-ordination 3. Chelating co-ordination ~~ Cationic Copper hexanoate Anionic H [U0,(N03),] Germanium tetrachloride Zirconium thenoyltrifluoro- acetonate these cases the compound usually possesses a high degree of covalency.Some compounds may be extracted into an organic phase with the forma- tion of “solvates” or mixed solvate hydrates as exemplified by the extrac- tion of several metal nitrates into ethers and ketones. (b) Types of Extraction.-Table 2 summarises the main types of extrac- tion; following this each will be dealt with in more detail. TABLE 2. Categories of solvent extraction 1. Physical solution in organic phase without solvation 2. Physical solution accom- panied by partial or complete solvation 3. Oxonium or other compound formation with the solvent l including ion exchange Example RuO in CCl,; S-hydroxyqui- U02(N03)2,2H,0,2S where S = molecule of solvent (ether ketone ester etc.) Nitric acid in ether (Et,OH+) (NO,-) FeC1,-HCl in ether (Et ,OH+)( FeCl 4-) nolinates in CHCl, etc.Irving Rossotti and Williams9 have developed a quantitative generalised treatment of partition equilibria in inorganic systems by combining the concept of step equilibrium with the partition law. In this the partition equilibria are described quantitatively in terms of step equilibria governing Irving Rossotti and Williams J. 1955 1906. 330 QUARTERLY REVIEWS the relative concentrations of different species in the aqueous phase and of a series of partition coefficients referring to the species which are common to the two phases. Development of this theme enables the relative degrees of association in each phase of metal ion with hydrogen and ligand ions to be determined ; partition coefficients and stability constants of species present in simple “ideal” systems may be derived and the probable nature of species present in more complicated systems may be deduced.Although the comprehensive treatment of lrving et al. should be read by all serious students of inorganic extraction we consider it more appropriate here to give individual attention to the various classes of extraction. (c) Extraction of Covalent Compounds (into Inert Solvents).-The extraction of such compounds is often described as being effected by “simple” physical solution in the organic phase. However although direct solvation or co-ordination of solvent to solute molecules may not seem to be involved there must nevertheless be some intermolecular interaction leading in many cases to departures from ideality. Examples of species which fall into this category include elements (e.g.chlorine bromine iodine) and compounds (e.g. sulphur dioxide mercuric chloride and the tetroxides of Group VIII); the organic solvents which will extract these from aqueous solutions are usually inert (hydrocarbon halogenated hydrocarbon etc.). Accounts of studies of systems such as iodine in carbon tetrachloride- water will be familiar to many readers; the general principles involved are applicable to less familiar but equally interesting cases e.g. the tetroxides of Group VIII. For ruthenium tetroxide the distribution coefficient for the; carbon tetrachloride-water system is constant over the twenty-fold con- centration range 0.006-0*012~ in the organic phase and has the value D = 58.4;’ i.e. equation (1) is applicable and it may be deduced that the reactions represented by (i) RuO + H,O + H,RuO 1 (ii) H,RuO + HRu0,- + H+ J’ (Aqueous phase) (iii) H,RuO + HRu04+ + OH- (iv) xRuO + (RuO,) (CCI phase) do not occur appreciably.However reactions (ii) and (iii) are obviously dependent upon pH and by varying the latter their equilibrium constants may be obtained from partition experiments. If K = [H+][OH-] D = [RUO,]~/[R~O,]~,. (0 = organic phase) D’ = [RuO,]o/([RuOa]as. + [HRuO,-]) (at high pH) [HRuO,-] [H+] [RuO 4 I aq. K = = acid dissociation constant of Ru0,aq. then K, = K,(D - D’)/[OH-ID‘ . . . . . . . . . . (5) Experimentally K (from 5) was found to be constant when the hydroxyl MARTIN AND HOLT LIQUID-LIQUID EXTRACTION 33 1 ion concentration was varied over the tenfold range 0-001-0~01~ as would be expected if the simple partition theory is applicable.Salting-out of extractable non-electrolytes. The solubility of non- electrolytes in water is affected by the presence of neutral basic and acidic salts (“Setchenow” salting-out) ;lo hence the distribution of a non- electrolyte between organic and aqueous phases will also be affected by their presence. There is a marked difference between this type of salting-out and that caused by the common-ion effect which applies to the extraction of dissociable species. The latter plays a prominent part in many other systems and is discussed in the appropriate sections. “Setchenow” salting-out is described qualitatively by . . . . . . . log S/So = kc (6) where So = solubility of non-electrolyte in water; S = solubility of non- electrolyte in salt solution of concentration c; and k = constant (salting- out coefficient).For gases in water the expression is found to hold up to surprisingly high values of c (e.g. up to 3N). For distribution experiments expression (6) becomes . . . . . . . log D - log D = kc (7) where D applies to water and D to salt solutions. D is proportional to the activity coefficient (y) of the non-electrolyte in the aqueous phase; for dilute solutions in pure water y may be taken as unity. Hence we also have where y (= D,/D) is the activity Coefficient of the non-electrolyte in aqueous solutions of salt concentration c. As in the case of gases the logy = kc . . . . . . . (8) 0.3 rl k 0 2 4 Soft in ogueous pboSe,M 6 FIG. 1 Eflect of salting-out on the extraction of RuOo by carbon tetrachloride from linearity of this expression is verified for ruthenium tetroxide up to quite high salt concentrations in water (Fig.1). water. A NaCl. B NaNO,. C KC1. D KNO,. E LICl. F LiN08. lo McDevit and Long Chem. Rev. 1952,51 119. 332 QUARTERLY REVIEWS Setchenow salting-out results usually in a decrease in solubility of the non-electrolyte in the aqueous phase (and hence enhanced extraction into the second phase). However k is sometimes negative (“salting-in”). Whilst salting-out can be tentatively described in terms of a lowering of water activity by the binding of water molecules into the hydration spheres of the added electrolyte species such an explanation does not unambigu- ously apply to salting-in; a satisfactory explanation of Setchenow salting- out or -in is still lacking. (d) Extraction of Ion Association Complexes into Neutral Solvents.- Since no classification of types of solvent extraction can be rigid the term “neutral” here is relative and does not mean for instance that the solvent cannot form solvates with the solute; it does mean that it does not (in the present context) participate in any kind of ionic reaction.Some solvents e.g. diethyl ether act by solvating an electrically neutral ion-association complex and also by participating in ionic reactions as in oxonium-salt formation examples of the two modes are (i) U02++ + 2N03- + xH20 + (4 - x) Et20 + UO2(NO3),,(Et2O),_ (H2O)cc (ii) Hf + NO,- + Et20 + [Et,OH]+ + (NO,)- where the left-hand and the right-hand side represent aqueous and organic phases respectively. Many solvents e.g. tributyl phosphate are always neutral and act by solvation of electrically neutral species.A large amount of work has been carried out on the uranyl nitrate-tributyl phosphate (TBP) extraction system UOz++aq. + 2NO3-aq. 3- nTBPorg. + U02(N03), nTBPorg. . . . (9) [Square brackets denote appropriate activities.] Thus Du depends upon the nth power of the concentration of tributyl phosphate. By using an inert diluent such as kerosene the variation of Du with ester concentration can be observed and thus the value of n deter- mined. For uranyl nitrate DU varies as the square of the phosphate concentration ; consequently it is concluded that the disolvate U02(N03) 2 2TBP exists in the solvent phase.I1 Similarly nitric acid is extracted from aqueous solutions by tributyl phosphate and D N H O ~ varies directly as the solvent concentration (for dilute solutions) ; nitric acid therefore seems to extract as HN0,,TBP.12 Tributyl phosphate also extracts the nitrates of tervalent cations [e.g.La(N03),,3TBP] and of quadrivalent cations [e.g. Th (N03),,2TBP] as well as of other bivalent oxyions such as l1 Moore 1951 USAEC Unclassified Document AECD 3196. la Alcock Grimley Healey Kennedy and McKay Trans. Furaday Soc. 1956,52,39. MARTIN AND HOLT LIQUID-LIQUID EXTRACTION 333 plutonyl and neptunyl. In this type of extraction the number of solvate molecules attached to the central atom of the extractable solute makes up the usual co-ordination number of that atom. When solvents like ethers alcohols and ketones act in their so-called neutral capacity mixed solvate-hydrate conditions sometimes arise e.g. as in U 0 2 ( N 0 3 ) 2 3H ,O,Et 20 and UQ 2(N03) 2,2H,0,2Et 20.13J4 (e) Extraction of Ion-association Complexes into Inert Solvents.- Some ion-association complexes notably those containing bulky anions or cations e.g.tetraphenylarsonium per-rhenate pertechnetate or per- manganate are soluble in non-reactive solvents such as chloroform. The extractable complexes are usually formed in situ by reaction of the ap- propriate salt with e.g. tetraphenylarsonium chloride. Relatively simple equilibria are involved in these reactions; their treatment bears a formal resemblance to that of chelate extraction systems which are dealt with on page 337. (f) Thermodynamics of the Extraction of Ion-association Complexes.- Uranyl nitrate again provides us with an example of the treatment of the extraction of ion-association complexe~.~J~ From equation (9) we have .. . . . . . y = ~ ~ x ( ~ 0 ~ - ) 2 ~ 3 ~ 2 ~ ~ 2 / y ~ (10) where x and y are the molar concentrations of uranyl nitrate in aqueous and organic phases respectively; ye is the activity coefficient of the solvate in the organic phase; y is the mean molar activity coefficient of uranyl nitrate in the aqueous phase (y is related in the usual manner to y+ and y-) ; YT is the activity coefficient of solvent ; and T is the concentration of free solvent in the organic phase. If the sole source of nitrate is from dissociation of uranyl nitrate (NO3-) = 2x and Also Tis equal to To - Tbound where To is the original concentration and Tbound is the quantity found in the Sohate. In the present case Tbound = 2y; hence where E = 2y/T0 = degree of binding of solvent = amount of saturation of solvent with solvate.E is an important quantity in the consideration of the similtaneous extraction of two solutes. For pure undiluted solvent and low concentrations of solute equation (1 1) reduces to Thus the plot of YYe against x3y3 should be linear with slope = 4 0 ~ . Values of y are available from vapour-pressure studies of uranyl nitrate . . . . . . . . y = 40ux3y3T2yT2/yO (1 1) . . . . . . T = To -2y = T,(l - € ) (12) . . . . . . . . y = 4Du x3y3/y(j (13) l3 Katzin and Sullivan J. Phys. Colloid. Chem. 1951 55 346. l4 Bachelet and Cheylan J. Chim. phys. 1947 44 248. l5 Rozen and Khorkhorina Zhur. neorg. Khim. 1957,2 1956. 3 34 QUARTERLY REVIEWS solutions; Ye may be taken to equal unity at low concentrations of solute. Having found DU for low concentrations we can use it to find y at higher concentrations.The validity of this linear law has been demonstrated for uranyl nitrate over a concentration range of lo2 for ethereal solvents.16 Equations (11) and (13) imply the complete immiscibility of the two phases. In practice the sohbility of the solvent in the aqueous phase can usually be neglected; however in some cases dissolved water in the organic phase may appreciably affect the activity of the solvent. Some of this water may be associated as hydration of the solute ( e g . the mixed hydrate-solvates already mentioned) a phenomenon more usually associated with ethereal ketonic and alcoholic solvents than with esters of the phosphate type. For uranyl nitrate the solute activity in the organic phase appears to be a function of the fourth power of the water a~tivity,~ as a result of which equation (1 3) becomes (14) y = 4Du x3y3aw4/yeo .. . . . . where yeo is the limiting value of ye with pure water and a is the water activity (or water vapour pressure). Salting-out. The addition of a second salt with a common ion (in- soIuble in the organic phase) to these systems results qualitatively in the formation of a higher proportion of neutral or ion-association solute mole- cules through the common-ion effect and hence in enhanced extraction. As an example the addition of a bivalent metal nitrate to uranyl nitrate will be considered. In the aqueous phase the activity of uranyl nitrate is given by [UQ,++] [NO3-I2 = ~ X ( X + where z is the concentration of the second nitrate. This leads to the parti- tion equation1’ where y is now the activity coefficient of uranyl nitrate in the mixed nitrate system.If values of y are available partitions can be estimated; alternatively if the partitions are determined values of y can be derived. If the second nitrate is itself extracted into the solvent account must be taken of altered activity coefficients in that phase and also of the factor E (equation 12). If the second extractable nitrate is nitric acid the main effect in the case of neutral phosphate solvents is on the value of E and there is “competition” for solvent molecules between uranyl nitrate and nitric acid. This leads to the type of extraction curve shown in Fig. 2 where for small additions of nitric acid the salting-out effect is apparent but for larger additions the competition eflect becomes predominant and the extraction of solute falls.In the case of ethereal solvents the “neutral” y = 4 D d x + ~ ) ~ y ~ a ~ ~ / y ~ ” . . . . . ’ (15) l6 GlueFkauf McKay and Mathieson Tram. F’rurlay Soc. 1951,47,437. l‘ J-s and McKay Tram. Farahy Soc. 1954,50,107. MARTIN AND IIOLT LIQUID-LIQUID EXTRACTION 335 r6le is transformed into a reactive one in the presence of concentrations of strong acid and the solvent may extract conxplex anions as oxonium salts (see below); the competition effect is therefore less marked and salting-out is effective over a wide range of concentrations. FIG. 2 Tj-pica1 curve for the extraction of an actiriide nitratefrom various concentrations of nitric acid into a neutral phosphate ester such as tributyl phosphate (TBP) being a plot of partition coefficient of extractable nitrate ( D in arbitrary units) against the molarity of nitric acid in the aqueous phase (M,HNO,).To the left of the peak salting-out predominates; to the right competition. (g) Extraction into Reactive Solvents.-Some solvents particularly ethers become reactive under conditions of high acidity by virtue of formation of oxonium-type salts; in this a solvent molecule becomes co- ordinatively attached to a hydrogen ion and electrical neutrality is maintained by ion association with an anion. Such solutions may have determinable (but low) electrical conductivities whereas solutions of inorganic salts in neutral solvents such as tributyl phosphate will have negligible conductivities . Nitric acid can be extracted from water by ether in the following way With another extractable nitrate such as uranyl nitrate the presence of nitric acid leads to the formation of acids having complex anions such as H[UO,(NQ,),] and this tri-nitrate acid species may be found in the solvent phase.Another example is the extraction of ferric chloride into di-isopropyl ether from hydrochloric acid where one of the extractable species is H[FeCl,] and the oxonium salt (Pr',OH)+ (FeC1,)- exists in the solvent phase. (h) Acidic Solvents.-Whereas solvents like ethers ketones and alcohols are most useful in the extraction of nitrates (as a result of the ready 336 QUARTERLY REVIEWS formation of neutral or anionic nitrate complexes) some newer organophos- phorus acidic solvents are of great interest in the extraction of metal compounds from aqueous sulphate chloride and phosphate solutions.Typical examples are monoheptadecylphosphoric acid (HDPA) (hepta- decyl dihydrogen phosphate) and di-2-ethylhexylphosphoric acid (D2EHPA) [di-(2-ethylhexyl) hydrogen phosphate]. For viscosity reasons these solvents are normally used in kerosene xylene or some other inert diluent. It has been shown that acid alkyl phosphates exist in these phases as dimers and that uranium extraction varies as the square of the acid alkyl phosphate concentration and inversely as the square of the hydrogen- ion concentration in the aqueous phase.18J9 It may be concluded that the extraction mechanism is UO2++aq. + 2(HX) + UO2X2,2HXorg. + 2H+aq. (16) where (HX) represents the dimerised acid. This is a type of cation exchange and explains the ability to extract uranium from solutions containing sulphate ions etc.as well as from the nitrate solutions. Generally the extracting power of dialkylphosphoric acids is three or four times greater than that of neutral alkyl phosphates. Synergism. A remarkable feature exhibited by Qialkyl hydrogen phosphates is their synergistic behaviour in the presence of small amounts of neutral alkyl phosphates or other organophosphorus compounds. Synergism here is the co-operative and cumulative action of separate extractants such that the total effect is greater than the sum of the individual effects. Di-(2-ethylhexyl) hydrogen phosphate (D2EHPA) forms syner- gistic mixtures with e.g. tributyl phosphate (TBP) dihexyl hexylphos- phonate (DHHP) and trioctylphosphine oxide (TOP0).20 Synergistic enhancement occurs in the extraction of uranyl plutonyl and quadrivalent plutonium solutions21 but not of U(IV) V(IV) Al or Mo.Considering that diethylhexyl hydrogen phosphate (D2EHPA) extracts 3-4 times as strongly as tributyl phosphate the effect of small additions of the latter is quite notable. The addition of as little as O-O5~-tributyl- phosphine oxide to the diethylhexyl ester in the extraction of uranium from 1.5hl-sulphuric acid may increase the partition coefficient by a factor of fifty. As a rule as the concentration of neutral additive increases the enhancement increases to a maximum and then decreases; an extraction curve of the general shape similar to that of Fig. 2 is obtained. A tentative explanation22 of the synergism implies that two moles of Blake Baes Brown Coleman and White UNICPUAE 1958 P/1550.* l9 Baes Zingaro and Coleman J.Phys. Chem. 1958 62 129. 2o Anon. Reactor FueIProcessing 1959 2 (l) 12. 21 Blake Horner and Schmitt 1959 USAEC Unclassified Document ORNL 2259. 2z Kennedy 1958 U.K.A.E.A. Document Number AERE CjM 369. * Here and later UNICPUAE refers to the Geneva Conferenws on the Peaceful Uses of Atomic Energy.. MARTIN AND HOLT LIQUID-LIQUID EXTRACTlON 337 dimer normally (i.e. in the absence of synergistic additives) depolymerise before forming the octaco-ordinated complex (RO) ,(OH)P=O ,--tUO 2 1 (i.e. the extracted species according to equation 16). In the presence of the neutral additive two of the links in this complex could be replaced by links from the additive it being assumed that the solvating (co-ordinating) properties of e.g.tributyl phosphate would not differ much from those of the di-(2-ethylhexyl) ester. Thus the necessity of monomerising one of the two moles of dimer would be obviated with an estimated saving in free-energy change of about 8000 cal./mole. The diminished energy requirements could account for the enhanced extraction. The equation representing the synergistic extraction would then be of the form The decrease in enhancement at high neutral additive concentrations is accounted for by a type of competition effect due to interaction (hydrogen bonding) between the dialkyl hydrogen phosphate and the additive,,l e.g. TBP + (HX) + TBP(HX) and TBP + 9(HX) + TBP(HX) thus effect- ively increasing the value of E (equation 12). (i) Basic Extractants.-Like the acidic organophosphorus reagents the basic organonitrogen compounds are relative newcomers into the extrac- tion field.These are amines (e.g. tri-iso-octylamine) and they act esssentially as liquid anion-exchanger materials just as the acid phosphates are liquid cation-exchanger materials. They will extract simple and complex anions (e.g. trinitratouranyl) from highly acid solutions and the solute may be recovered by washing the solvent with dilute acid or complexing solutions. For example quadrivalent neptunium and plutonium both form anionic nitrate complexes in 4-6M-nitric acid and may be extracted as such by a solution of tri-iso-octylamine in ~ylene,,~ whilst the same solvent will extract anionic uranium chloro-complexes from hydrochloric acid.,* (j) Extraction of Chelate Compounds.-Although extractable metal chelates (usually soluble in inert solvents such as benzene chloroform etc.) are essentially covalent compounds they cannot be treated in quite the same simple way as the extractable covalent compounds described on pp.330-332. The latter exist essentially in a single form in both phases whereas the presence of the extractable chelates depends upon a number of equilibria (hydrolysis and association) as well as upon the presence of excess of chelating compound in the inert solvent. OS Sheppard 1957 USAEC Unclassified Document HW 51958. 24 Moore 1957 USAEC Unclassified Document CR-57-1-61; Analyt. Chem. 1951 29 1661. 338 QUARTERLY REVIEWS General treatment. The equilibria involved are Ionisation of chelating compound (assumed to be a weak acid HA) HA + H+ + A-; KS= [H+] [A-]/[HA] .. . . . . (18) Stepwise formation of chelate with metal ions of valency n Mff+ + A-' + MA("1)+ ; Kl = [MA('+l)+]/[Mn+][A-] MA("-l>+ = A- $ MA,("i-2)+; Kz = [MA,("2)+]/[MA("l)+][A-]; MA+,,1 + A- + MA,; K = [MA,]/[MA+n-l][A-] . . (19) The overall formation constant K = K1.K2.K3 . . . . . K,. Hydrolysis and so on. Anion co-ordination (anion = X-) and so on. Distribution of chelating compound Mn+ +- OH- + M(OH)(*l)+ . . . . . . . . (20) Mn+ + X- + MX(*l)+ . . . . . . . (21) (22) DHA = [HA]org./[HA]a, . . . . . . distribution of the metal chelate DMA, = [MAnlorg./[MAn]Itq. . . . . . . (23) From these equilibria an expression for the apparent partition co- efficient D (i.e. the ratio of the total metal concentrations in the two phases) can be derived.If we assume that M is present in the organic phase only as MA, it follows that [MAnlorg. [Ma+] + [MA("')+] + . . . . . [MAfn-l] + C [M(OH)&k-i)+] + C[MXj("Q+] D = @ I . . . . (24) Substituting from expressions (19) and (23) after division of numerator and denominator by [MA,],,, we have MARTIN AND HOLT LIQUID-LIQUID EXTRACTION 339 Substituting from (18) and (19) we have This expression shows clearly the importance of the hydrogen-ion concentration the stability of the chelate (reflected by Kf) and the relative solubility of the chelate (reflected by DMAJ in the organic phase in governing the extraction of the metal. Equation (26) may be simplified by assuming that hydrolysis and anion co-ordination of the metal ion in the aqueous phase are negligible (zk x = 0) that [MA,laq.is negligible [H+l K,&" [HA lorg. &-€An (i.e.,{ }' 9 - ) and that negligible amounts of the inter- mediate chelate species are formed;26 Thus if [HAIorg. is constant the extraction is a function of pH alone and curves of distribution versus pH are of great analytical significance.aa In the case of the extraction of e.g. zirconium by thenoyltrifluoracetone (TTA) where hydrolysis in aqueous solution cannot be neglected Connick and McVey2' derived by a rather similar mathematical route the expression Treatment of hydrolysable species. a In D a in [HA] = 4f' + 3 7 + 2 f ' Z +f't - (4f0 + 3ft + 2fe +f& * (28) wheref' andfrefer to aqueous and solvent (TTA in benzene) phases respectively and are equal to the fraction of total activity contributed by each zirconium species containing the indicated number of chelate groups.The dependence of Connick and McVey's experimental values of In D on TTA activity fitted a line of slope 4 and it was concluded that the extraction reaction is Zr(OH),(4-fl)+ + 4HA = ZrA + (4 - n)H+ + nHzO By studying the dependence of the extraction on the hydrogen-ion activity in the aqueous phase Connick and McVey were able to determine 86 Kolthoff and Sandell J. Arner. Chem. Soc. 1941 63 1906. 2* Irving and Williams J. 1949 1841. 97 Connick and McVey J. Amer. Chern. SOC. 1949,71,3182. 340 QUARTERLY REVIEWS the average degree of hydrolysis of the zirconium species in the system by using the expression a In D/i3 In [H+] = - 4 +fl + 2f2 + 3f3 + . . . . . . . . . (29) where eachf = fraction of total zirconium in the aqueous phase having the number of OH groups per zirconium indicated by the subscript.For very dilute solutions of zirconium in 2~-perchloric acid the average species has between 0 and 1 hydroxyl groups attached to it i.e. it lies between 21.4+ and Zr(OH)3+. For the study of general anion (chloride nitrate sulphate fluoride) complex formation of zirconium the expression a In D/a In [HX] = -fl - 2f2 - 3f3 (where X = anion) was used and the degree of complex-ion formation was determined. 4. Solvent extraction in practice (a) Practical Calculations.-The simple distribution equation (equation 1) applies only occasionally the value of D often varying with solute con- centration. If this variation is known it is possible to calculate the separa- tion and extraction of different solutes. It is often convenient to plot extraction systems graphically by using x and y co-ordinates for aqueous and solvent concentration.The mass balance for a single equilibration of a volume Val of aqueous phase containing concentration xo of a solute with volume Vsl of pure solvent is . . (30) By substitutingyJx = D and Vsl/Val = R, we have x = x,(l +R,D,). If the two phases are separated after equilibration and a volume Vsz of pure solvent is equilibrated with the aqueous phase from the first operation then x1 = x2(1 + R2D,) and xo = x2(1 + R2D2)(1 + R,D,) . . (31) In the ideal case (D = constant)? if R is constant for all stages then x0 = xn(l + RD)” . . . . . ‘ (32) The distribution diagram (Fig. 3A) refers to a three-stage batch extraction process. The point A (x = xo y = 0) represents the composition of the starting phases before mixing; the point By where the line through A of slope -1/R cuts the line y / x = D gives the equilibrium concentrations y and x,.If the equilibrium had not been fully attained the concentrations would still be on the line AB but the solvent concentration would be less than yl. The second stage of the extraction starts at C the point ( x l 0) and after equilibrium is reattained gives x, y z as the appropriate concentra- tions. As “batch” extraction this is familiar in the form of extraction of e.g. fatty acids with ether in organic chemistry and of ferric chloride with MARTIN AND HOLT LIQUID-LIQUID EXTRACTION 34 1 ether in inorganic analysis. Batch extraction is a poor method for separating materials with similar D values.The ease of separation of two solutes is indicated by the magnitude of the “separation factor” i.e. the ratio of the two D values. As the separation factor approaches unity so the difficulty Concrnirotion in oyueous pbore Concontrotion in apueous pboso FIG. 3 Distribution diagrams for batch extractions. of separation increases. Separation even in more or less difficult cases can however be accomplished by countercurrent liquid extraction where the effective number of stages is large. It is convenient to consider first the mass balance and the distribution diagram (Fig. 3B) for a batch extraction in which both phases initially contain solute i.e. (xo - xl) = R(y - y o ) andy,/x = D . . . . . . (33) Fig. 4A shows the flow diagram for countercurrent extraction for a hypothetical three-stage process.Each stage can be regarded as a batch extraction whose products are the feeds to adjacent stages; this leads to the following mass balances Stage 2 (x - x,) = R(y2 - y3) and y2/x2 = D2 } . . . (34) Stage 1 (x, - x,) = R(yl - y 2 ) and yl/xl = D Stage 3 (x - x,) = R(y3) and y$x = D These are shown in the form of a distribution diagram in Fig. 4B whilst a simplification is shown in Fig. 4C. This type of diagram (due to Varteressian and Fenske28) is based on the fact that in the rectangle representing stages in Fig. 4B the diagonals opposite the mixing lines all fall on a straight line of slope IIR; this line @ Varteressian and Fenske Ind. Eng. Chem. 1936,28,928 1353. 342 QUARTERLY REVIEWS A stuge 3 Stoqe 2 Stogu I Aqueous raf finata "3 1 1 s p - x2 M feed x3 "2 "I "0 "3 x2 %I x Concenfrofion i n a p e o u s phase FIG.4 Flow diagram (A) distribution diagram (B) and Varteressian and Fenske diagram (C) for three-stage countercurrent system. S = Settler; M = Mixer. is the "operating line" and each stage is represented by a step between the operating and the equilibrium lines. Ideally (i.e. D = constant) the mass balances reduce to . . . . . . . (35) I Stage 1 (xo - x,) = RD(xl - x,) Stage 2 (x - x,) = RD(x - x,) Stage 3 ( x - x3) = RD(x,) which combine to give the geometric progression xo - X I X I - x e xe - x 3 X I -x x2 - x 3 x3 --I - - RD. (36) - . . . . . . . whose summation in the general case for n stages is (RD)n+l - 1 xo = x . . RD - 1 . . . . . . . . . . . (37) More efficient separations and purifications are obtained if the solvent product from such countercurrent extraction is "stripped" in further countercurrent stages.In this operation a solute-free aqueous phase washes part of the solute from the solvent together with the bulk of any impurities which may have been extracted with a low partition coefficient, MARTIN AND HOLT LIQUID-LIQUID EXTRACTIOX 343 The aqueous product from this stripping operation is re-cycled with the aquems feed solution to the extraction stages. Fig. 5A shows the flow arrangement for a countercurrent extraction system with three extraction A Concentrotion in cpueous phase f FIG. 5 Flow diagram (A) and Varteressiun and Fenske diagram (B) for a system having three strip and three extraction stages. Two soIiites are considered. The relative flaw rates are strip 1 ; solvent 2; aqueous feed 1 .and three stripping stages. The Varteressian and Fenske diagram (Fig. 5B) shows the stage-to-stage concentration changes for two solutes whose D values difl’er only slightly. The values of D chosen in the illustration approximate to those for praseodymium and neodymium nitrates in the nitric acid (13*8~)-tributyl phosphate system.29 Fig. 5B shows that the use of six stages brings about a considerable separation between praseodymium and neodymium; since countercurrent extraction can readily be made a continuous process with numerous stages the technique is obviously a powerful tool for difficult separations and for preparing high-purity materials either in small quantity in the laboratory or in bulk industrially. In most cases complications in treatment are caused by the dependency of D on concentration and on the presence of other solutes in both phases (page 334).(b) Equipment.-Continuous countercurrent solvent extraction is usually 29 Hesford Jackson and McKay J. Inorg. Nucl. Chem. 1959 9 279. 344 QUARTERLY REVIEWS carried out either in a vertical column or in a mixer-settler cascade. The scale of operations can vary between a small analytical apparatus handling less than one ml. of liquids per minute up to large industrial plants processing 2000 1. or more per minute. A comprehensive review of equipment cannot be given here but a brief description of some simple columns and mixer-settlers is desirable. (i) Columns. In columns one phase is dispersed in and is made to flow under gravity through a continuous column of the other.The light phase is introduced at the bottom of the column and taken off at the top; the denser phase is admitted at the top and flows downwards. In order to improve solute transfer rates across the solvent-aqueous interface and so to reduce the height of column equivalent to a theoretical stage (HETS) various packing systems and/or mechanical agitation are used. (ii) Mixer-settlers. The many varieties of this type mostly fall into two classes first those in which the mixer lifts the mixed phases into a settler from which the separated phases can flow over weirs to the next mixers and secondly those in which the mixers perform virtually no lifting. In the first class the position of the interface in the settler is controlled by the relative position of the weirs and is therefore independent of the neighbouring stages.This makes possible a very flexible unit; additionally the contents of a settler can be sampled for analysis. The other class of mixer-settler is used mostly for large-scale work but may also be used in the laboratory. They suffer on the small scale from instability in operation as a result of surface-tension effects interfering with inter-stage flows. For a fuller discussion of laboratory-scale mixer- settlers the review by Jamrack et aL30 should be consulted; large-scale equipment has been reviewed by Pratt.31 (c) Applications.-Solvent extraction is of increasing importance in general extraction metallurgy and in other fields where difficult separations are to be made or high-purity products are required. It is used extensively in analysis1 and as a laboratory technique in such fields as the study of lanthanide and actinide chemistry29 and of complex formation as described in earlier sections.Analytical aspects have been reviewed by 1rving;l the review by Fletcher32 of the potential uses of solvent extraction for purification in extraction metallurgy should be consulted for the broader aspects in that field. (i) Nuclear energyJield. Much work on solvent extraction has been initiated by the special needs of the atomic energy industries. It is fre- quently necessary to attain high levels of purity in materials in order to remove trace amounts of contaminants having undesirable nuclear 30 Jamrack Logsdail and Short “Progress in Nuclear Energy,” Series 111 1958 Vol. 2 (Process Chemistry) Pergamon Press London p. 332. 31 Pratt UNICPUAE 1958 8 520.32 Fletcher “Extraction and Refining of the Rarer Metals,” Institute of Mining and Metallurgy 1957 London p. 15. MARTIN AND HOLT LIQUID-LIQUID EXTRACTION 345 properties e.g. nuclides with high capture cross sections for thermal neutrons. The important nuclear fuels are the fissile isotopes uranium-233 uranium-235 and plutonium-239; 233U and 239Pu are obtained by neutron irradiation of the natural fertile materials 232Th and 238U. Solvent-extraction procedures have been devised to separate these from irradiated fertile fuels in high states of purity (usually as their nitrates). Similar procedures may also be used to obtain other very pure materials for nuclear reactor use e.g. zirconium for the protection of bare uranium fuel or structural use. (ii) Uranium.Naturally occurring uranium is the most important source of fissile and fertile materials; its separation from ores and sub- sequent purification have been the subject of much research in which solvent extraction studies have played a large part. Uranium is usually extracted from solutions as an ion association nitrate complex by ethers (as solvate or oxonium salts) or by tributyl phosphate (as a solvate). An early British p r ~ c e s s ~ ~ used batch extraction of molten uranyl nitrate hexahydrate by diethyl ether. Present processes34 use continuous countercurrent extraction of uranyl nitrate from filtered acidic ore leach liquor into tributyl phosphate diluted with kerosene. The purified uranyl nitrate is washed back from the loaded solvent by water. In this system the uranyl nitrate partition coefficient decreases with increasing temperature ; the wash-back section is therefore operated at a high temperature to obtain more efficient solute recovery.Processes based on tributyl phosphate are also used extensively in the U.S.A. and in France for uranium purification. Complexing anions in the crude aqueous solution e.g. Po,’- Sod2- F- etc. reduce the partition coefficient of uranyl nitrate into most solvents including tributyl phosphate.35 However recently developed processes36 for the extraction of uranium from sulphate phosphate and chloride liquors use liquid-ion exchange solvents of the type described on p. 335 and the quaternary ammonium anion-exchange solvents for example are suffi- ciently strong bases to be able to extract anionic uranyl carbonate from alkaline sodium uranyl carbonate solution.The values of the partition coefficients for uranium obtainable with these ion-exchange solvents can be as high as several thousand; it is thus possible to extract uranium efficiently from dilute solutions and to recover it in a more concentrated form. (iii) Thorium. Thorium being a fertile nuclear material must be freed very thoroughly from rare-earth metals with high neutron cross sections and from uranium in order to avoid dilution of “bred” 233U by the ss Grainger UNICPUAE 1958 8 149. 34 Hamilton Nature 1959 lS3 789. 35 Quoted in ref. 36. s6 Brown Coleman Crouse Blake and Ryan UNICPUAE 1958 3 472; see also refs. 18 and 32. 346 QUARTERLY REVIEWS L Feed solution Final product solution (overall Th recovery = 99.7%) 838U originally associated with the thorium.The values of the partition coefficients under one set of conditions for uranium and thorium nitrates into 40 % tributyl phosphate in xylene are respectively 20 and 0.5 whereas they are 6 and 0-04 into 5 % phosphate in xylene. Thus an initial extraction with the dilute solvent will remove the uranium and a second operation with the more concentrated solvent will recover the thorium. This principle forms the basis of a process37 described by Audsley and Table 3 shows the power of the solvent extraction method in effecting the necessary separa- tions. CeO 830 g. RE,&* 780g. U 48 mg. 1.1 mg. 3.5 mg. 0.08 mg. I I * RE == Rare earths Residual phosphate (derived from the monazite) in the feed solution to the solvent extraction may be complexed by adding an equivalent amount of ferric nitrate3* in order to prevent its interfering with extraction efficiency through the formation of inextractable thorium phosphate complexes.Methods for the separation of uranium-233 from irradiated thorium39 depend upon first extracting both uranium and thorium leaving the bulk of the fission products in the aqueous phase; the thorium is then preferen- tially washed back. Some of these processes use tributyl phosphate with aluminium nitrate as salting-out agent in an acid-deficient aqueous phase. These conditions together with saturation by thorium nitrate of the solvent in part of the system i.e. the value of E (p. 333) is high give good decontamination from rare-earth fission products and protactinium which otherwise tend to follow the thorium.(iv) Plutonium. PIutonium is found in irradiated fuel elements in association with a large bulk of “unburned” uranium and with fission products of low bulk but high specific radioactivity. The approximate composition of natural uranium after an irradiation corresponding to an energy release of 10o0 megawatt-days per ton is U 99.8 % Pu 0.08 % Audsley Lind and England ref. 32 p. 351. Menzies J . A p ~ l . Chcm. 1959 9 249. Sadowski ibid. 1958,17,49; Buller ibid. 1955,9,464. s9 Greskey. UNICPUAE 1955 9 505 Bruce Shank Brooksbank Parrott and MARTIN AND HOLT LIQUID-LIQUID EXTRACTION 347 fission products 0*08%.4* From this mixture it is required to separate the uranium and plutonium and to purify each of them from highly radioactive fission products. Many chemical processes for accomplishing this have been described the most successful of which are based on solvent e~traction.~~ Older separation processes use ketonic or ethereal solvents like “butex” (diethylene glycol dibutyl ether) diethyl ether isobutyl methyl ketone (hexone) and triglycol dichloride.These are oxonium-type extractants and they accordingly extract solutes more efficiently under highly acid or salting-out conditions. In a process as used at Windscale (U.K. Atomic Energy Authority) irradiated uranium is converted into a solution of uranyl Pu(rv) and plutonyl and fission product nitrates in excess of free nitric acid (3M). Both uranyl and plutonium nitrates are extracted into “butex” leaving the bulk of the fission products in the aqueous raffinate. The solvent from this extractor is partially neutralised with aqueous ammonia and is treated with ferrous sulphamate to reduce the extractable PuO,++ and Pu(iv) to the inextractable PU(ITI).This behaviour reflects the greater tendency of the higher-valency states to form nitrate complexes. The plutonium is re-extracted into an aqueous phase containing 8M-ammonium nitrate; the uranium remains in the solvent under the influence of the salting-out agent and is subsequently recovered in a third column by washing the organic phase with dilute nitric acid (0.05~). The plutonium and uranium solutions require further purification by extraction and wash-back cycles. Tributyl phosphate is used in most of the more recent plutonium separation processes41 to carry out the primary separation of plutonium uranium and fission products. This solvent acts by solvation and highly acid conditions are less necessary than in the case of ethereal or ketonic solvents.Neptunium and americium may also be separated pure from irradiated fuel-element solutions by solvent-e~traction.~~ (v) Niobium and tantalum. In Nature these two metals are always associated and separation from each other by classical methods is very Solvent extraction however provides a satisfactory process,44 either a hydrochloric acid aqueous phase and methyldioctylamine in xylene45 being used as the extractant or fluoride solutions (containing free hydrofluoric acid) with either or tributyl phosphate*’ as extract- ants. In fluoride extractions tantalum has the higher partition coefficient ; either it is preferentially extracted leaving niobium in the aqueous phase 4 O Howells Hughes Mackey and Saddington UNICPUAE 1958 17 3.41 See UNICPUAE 1955 8 and 1958 17 for a selection of papers. 42 McKay UNICPUAE 1955,9 314. 43 Miller Ind. Chemist 1959 175. 44 Foos and Wilhelm 1954 USAEC Unclassified Document ISC 694. 46 Ledlicotte and Moore J. Amer. Chem. SOC. 1952,74 1618. 48 Werning Highie Grace Speece and Gilbert Ind. Eng. Chem. 1954,46.644. l7 Morris. Wain and Fletcher Bull. Insf. Min. Met. 1956 No. 497,487. 348 QUARTERLY REVTEWS or both metals can be extracted leaving other impurities in the aqueous raffinate; the niobium is then selectively washed back from the solvent by suitably adjusting the composition of the wash-back solution. (vi) Zirconium and hafnium. Zirconium occurs in Nature associated with hafnium a powerful neutron absorber.For use in atomic energy processes efficient removal of hafnium is essential. Various methods are available; solvent extraction is used on the large scale in the U.S.A.,48 France,49 and Britain.50 In the American process isobutyl methyl ketone preferentially extracts hafnium thiocyanate from a hydrochloric acid solution. In the French and British processes tributyl phosphate pre- ferentially extracts zirconium from a nitrate solution leaving the hafnium and other impurities in the aqueous raffinate. The French use an aliphatic diluent for the phosphate and the ready formation of an organic third phase owing to the limited solubility of the zirconium solvate in the diluent limits the zirconium concentration in the feed solution to about 30 g./l. ; if xylene is used as a diluent the concentration can be increased to about 80 g./l.50 The separation of the rare-earth metals is a difficult problem which has recently been eased by the use of ion-exchange techniques but the amounts of material handled are small compared with the size of equipment used.Solvent extraction offers a possible method; a partial separation has been reported51 by extraction with tributyl phos- phates of rare-earth nitrates in aqueous nitric acid. Several papers53 have reported the behaviour of the rare earths in tributyl phosphate-nitrate systems. Useful separation factors have been obtained and some interesting differences between rare earths of odd and even atomic numbers have been discovered as a result of this work. (viii) Miscellaneous separations. Many other metals have been the subject of solvent-extraction studies for general chemical purposes as well as for their separation or purification and even for isotope separation.Table 4 lists a few metals for which solvent-extraction studies are reported. Some interesting processes have been reported for the recovery of acids from aqueous solutions. Moore reports a general study of the use of the anion-exchange solvents (long-chain amines) for the extraction of various acids including hydrofluoric nitric sulphuric and phosphoric as well as various complex acids containing the metals Po Pu Zr and Pa.53 The extraction of phosphoric acid into n-butyl or isopentyl alcohols has been operated on a pilot-plant scale for the production of phosphoric acid from phosphate rock treated with hydrochloric acid.54 48 Shelton Dilling and McClain UNICPUAE 1955 8 505.Hure and Saint James ibid. p. 551. 50 Hudswell and Hutcheon ibid. p. 563. 51 Foos and Wilhelm USAEC Unclassified Report ISC 695. 52 Hesford Jackson and McKay J . Inorg. Nucl. Chem. 1959,9 279. 63 Moore Analyt. Chem. 1957 29 1660. 54 Daniel Blumberg and Aim Brit. Chen.1. Engineer April 1959 223 (vii) Rare-earth metals. MAKTIN AND HOLT LIQUID-LIQUL1) EXTKACTION 349 Miscellaneous extraction studies (those discussed in the text are TABLE 4. Bi excluded). Solvent phase Separation from Ge process wastes I 57 I Ether 1 Separations from Ethers 1 G ~ ~ ~ ~ a ~ r t ~ ~ e m i C a study Metal 1 Aqueous phase I i I ~ - - Polonium1 HCl or H2S04 57 58 Actinium HC1; controlled I PH ~- Lithium Gallium ~M-HCI Scandium1 I Thiocyanate -i- chloride Thiocyanate 40% TBP Ion-exchange solvents Attempted isotope 59 separation Separation 60 I - Ketones and I General chemical alcohols 1 study Iron 61 NH,CNS I Beso and Cobalt - in vest iga t ion RaE Radio-Pb Thiocyanate and perchlorate I General chemical 1 study Diethyl ether 62 Various I 10% TBP 1 Removal of Fe 1 63 1 from Be I I ‘ _ ! 55 Bagnall “Chemistry of the Rare Radioelements,” Butterworths Scientific Publica- 56 Farr Giorgi Bowman and Money 1953 USAEC Unclassified Report LA 1545.57 Powell reported in Ref. 32. 58 Foos and Wilhelm USAEC Unclassified Report ISC 694. 59 Dury USAEC Unclassified Report Y 785. 6o Anon. “Reactor Fuel Processing,” 1959 2 (l) 13. 61 Brubaker and Johnson J. Inorg. Nucl. Chem. 1959,9 184. 6* Diamond J. Aner. Chem. SOC. 1959 63 659. 63 Byersmith Pomelee and Birmbaum USAEC Unclassified Report NYO 1 1 16.tions London 1957. 350 QUARTERLY REVIEWS 5. Molten metal systems Liquid-liquid extraction using immiscible or partially miscible metal is still very much in its infancy although the origins of one application (the Parkes process for the desilverisation of lead) go back more than 100 years. Even more so than for immiscible aqueous-organic systems the recent progress in molten-metal extraction systems is derived from the impact of nuclear energy. Briefly this arises from the consideration that metallic nuclear fuel may conceivably be reprocessed in the metallic state by processes which are inherently cheaper and simpler than those which in- volve chemical destruction of the metal followed by purification and reconversion into metal albeit the processes would be less efficient.(a) Principles.-Little fundamental work has been done on liquid- metal extraction systems but it is clear that the basic extraction law is similar to that of aqueous-organic systems and equation (1) applies ideally. For the system silver distributed between immiscible molten lead and zinc phases (i.e. the basis of the Parkes process) MellorG4 quotes results showing that the partition coefficient for silver is about 300 (in favour of the zinc phase) and is reasonably constant for all silver concen- trations up to 6% in lead. To describe systems exactly appropriate activities should be used but data are scanty and there is scope for much investigation. Just as aqueous- organic systems have been used to study complex formation ionic effects solvation etc.so liquid-liquid metal systems may be used to study inter- metallic compound formation “salting-out” (the effect of adding a second metallic solute) and heats of solutions etc. (b) Theoretical Predictions.-Within limits one can forecast whether certain pairs of molten metals are miscible and hence the qualitative behaviour of systems containing three liquid metals. MacKen~ie~~ used an expression due to Mildebrand and Scotta to predict miscibilities and to deduce that in a system of uranium plutonium and silver the plutonium should be extractable into the silver phase (which is immiscible with the uranium phase). Direct prediction may be made from phase diagrams when available. Hildebrand and Scott8 deduced as a criterion for the complete inis- cibility of two molten metals the expression where Vl and Vz are the atomic volumes of metals 1 and 2 6 and 8 are their solubility parameters R is the gas constant and T i s the absolute temperature.The solubility parameter 6 = ( dE,,/V)* where AE is the energy of vaporisation of the metal. cal Chemistry,” Vol. 111 p. 313 Longmans Green & Co. 1952. 64 Mellor “Comprehensive Treatise on Inorganic and Theol O6 MacKenzie Canad. J. Chem. 34 749. 1956. MARTIN AND HOLT LIQUID-LIQUID EXTRACTION 35 1 The same criterion has been used to predict whether plutonium and some metallic fission products might be extracted from uranium by molten calcium the last two metals forming an immiscible pair. The determined partition coefficients of a number of fission-product elements showed that qualitative predictions based on miscibility considerations were largely substantiated.In addition in accordance with equation (l) the distribu- tion coefficient for plutonium (Dpu = 0.0073) was reasonably constant over the concentration range 0-01 % to 20.4% of plutonium in uranium and the one for cerium (Dee = 9.4) did not vary significantly over the range 0.04 % to 0.46 % of cerium in uranium.66 (c) Applications of Liquid-Liquid Metal Extraction.-So far the main developments in this field have been made with the aim of removing plutonium and/or fission products from irradiated uranium or in some cases of removing uranium-233 from irradiated thorium. Silver,s5 rare- earth metals:’ and magnesiums8 have been used to extract plutonium from molten uranium in addition to the use of calcium for removing some fission products.For magnesium the applicability of equation (1) has been demonstrated. Experiments involving both magnesium and uranium would normally be carried out in pressurised systems since magnesium has a vapour pressure of > 760 mm. at the m.p. of uranium ( 1 1 3 5 ” ~ ) . Feder68 overcame this difficulty by alloying the uranium phase with 5 % of chromium and thereby lowering its m.p. to 950”c. In a similar manner a molten uranium-chromium eutectic phase may be used to extract pro- tactinium from molten magnesium-thorium e u t e c t i ~ ~ ~ and calcium- magnesium eutectic may be used to extract certain metallic fission pro- ducts from molten phtonium-cobalt and plutonium-iron alloys.70 * (d) Fundamental Studies.+) Analogy of salting-out or complex formation. The distribution of plutonium in the uranium-silver system is dependent on the plutonium c~ncentration;~~ it is known from vapour- pressure measurement~~~ that plutonium forms ideal (dilute) solutions in uranium.Consequently in the uranium-silver system the variation of Dpu with plutonium concentration must reflect on some non-ideal behaviour of plutonium in the silver phase. Insufficient data exist as yet to enable this effect to be assessed quantitatively. The addition of small amounts of gold to the uranium-silver system has an interesting effect on plutonium distribution (Fig. 6). Whilst the distribution of gold itself is virtually independent of concentration the extraction of plutonium into the silver pbase is increased by a factor of over 2 as the gold concentration increases from zero’to about 3.5 wt.% in the Martin Jenkins and Keen UNICPUAE 1958 17 352. 67 Voigt ibid. 1955 9 591. 68 Feder Greenberg Nathans and Nutthall Paper presented at Amer. Chem. SOC. Meeting April 7-12th 8957. ’* Chiotti and Voigj UMCPUAE 1958 17 368. ’* Anen. “Reactor Fuel Processing,” 1958 1 (2) 71 McKenzie Canad. J. Chern. 1956,34 515. 3 352 QUARTERLY REVIEWS uranium; beyond this point the plutonium extraction suddenly decreases. Qualitatively and terms being used analogous to those found in the sections on aqueous-organic systems this behaviour may be described as initial “salting-out” followed by “competition” of gold atoms with plutonium atoms for interaction with silver atoms as the gold concentration 14 r 0 2 4 6 W t . % o f Au uddd fo U FIG. 6 Variation of Kpu in silver-uranium with addition of gold (reprinted with permission from McKenzie Canad.J. Chem. 1956,34,749). increases It appears that the gold atoms “compete” successfully thereby indicating the greater stability of gold-silver “complexes” than of those of plutonium-silver. (ii) Determination of thermodynamic properties. By measuring the variation with temperature of Dpu in the system uranium-silver the limiting heat of solution of plutonium in silver has been determined. The distribution data65 were found to fit the equation log Dpu = 2.21 - 2-74 x 103/T. From the slope of this a value of AH (= 12.5 kcal.) is obtained where A H represents the heat of transfer of one g.-atom of plutonium from liquid uranium to liquid silver. On the basis of the ideal behaviour of plutonium in uranium and recognising that the experiments were carried out at low plutonium concentrations_(about 0.05 atom % in silver) this value of AH-may be equated with LO,pu the limiting partial molal heat content relative to liquid plutonium of plutonium in liquid silver. Similarly the limiting heat of solution of plutonium in some rare-earth metals has been e~timated‘~ by determining the variation of its distribution with temperature in the systems uranium-cerium -neodymium and -mischmetall. 72 McKenzie Canad. J. Chem. 1956 34 1176.

ISSN:0009-2681    

DOI:10.1039/QR9591300327

出版商:RSC    

年代:1959

数据来源: RSC

摘要:

THE INFRARED SPEC‘I’KA OF HETEKOAROMATIC COMPOUNDS By A. R. KATRITZKY M.A. D.PHIL. PH.D. B.Sc. (THE UNIVERSITY CHEMICAL. LABORATORY CAMBRIDGE) THE object of this Review is to increase the usefulness of infrared spectro- scopy to organic chemists working with heterocyclic compounds (for previous reviews see refs. 1 and 2). It is felt that the best way of doing this is to attempt a simplified explanation of the spectra of compounds con- taining planar five- and six-membered rings. Throughout the emphasis is on simplicity and not on rigour; thus approximate statements are made without qualification. a This Review commences with a short introduction on the origin of infrared spectra and the classification of the relevant molecular vibrations into types. Available data on the infrared absorption of heteroaromaric nuclei in various regions of the spectrum are then discussed and finally an account is given of the effect of heteroaromatic nuclei on the infrared absorption of substituents.It should be emphasised that the discussion is largely based on the fundamental work carried out on the infrared spectra of benzene and deuterobenzenes.3 The motion of an atom is fully described by the variation of three co-ordinates e.g. the x y and z components of its distance from a fixed point. The motion of a molecule of n atoms is defined by the variations of 3n co-ordinates; i.e. molecules possess 3n degrees of freedom. Three of these involve translational motion and three* involve rotational motion of the molecule as a whole; therefore 3n -6 degrees of freedom are associated with changes in the relative positions of the atoms i.e.with vibration. It can be shown that for simple harmonic motion a molecule of n atoms can vibrate in just 3n-6 ways in which each of the atoms is moving in phase and at the same frequency. These 3n-4 vibrations are the normal vibrations of the molecule and the displacements of the atoms from their equilibrium positions are the normal co-ordinates. Any other molecular vibration can be represented as the superposition of two or more normal * Two appear in linear molecules; hence these molecules show 3n - 5 vibrational modes. Bellamy “The Infrared Spectra of Complex Molecules” 2nd edn. Methuen London 1958 pp. 277-285. R. N. Jones and Sandorfy in Weissburger “Techniques of Organic Chemistry. Vol. IX. Chemical Applications of Spectroscopy” Interscience Publ.Inc. London See ref. 4 and references therein especially papers by Ingold and his co-workers. Herzberg “Infrared and Raman Spectra of Polyatomic Molecules” Van Nostrand Sheppard and Simwon Quart. Rev. 195.7 7 19. 1956 pp. 533-537. London 1945. 353 3* 354 QUARTERLY REVIEWS vibrations. Each normal vibration is quantised and the vibrational energy levels of a molecule are given by i = 3 n - 6 2 viwi i- 1 where w represents the normal vibrational frequencies and Y the quantum numbers. In theory infrared absorption bands could result from the promotion of a molecule from one energy level to any higher one. However for simple harmonic motion selection rules forbid transitions in which more than one vibrational quantum number is changed (i.e. combination bands) or in which the quantum number is changed by more than one unit (i.e.overtone bands); bands corresponding to these abnormal transi- tions are therefore absent or weak except in a few cases where Fermi resonance is irnp~rtant.~ Thus most of the strong bands in a spectrum correspond to one of the 3n -6 transitions in which each quantum number is separately altered by one unit i.e. to fundamentals. However all fundamentals do not cause strong absorption in the infrared region. The intensity of an infrared band depend upon the change of dipole moment during the vibration ; for certain modes of symmetrical molecules there is no change and the band is “forbidden”. If the change in dipole moment is small only a weak band results.* The complete interpretation of a spectrum involves a knowledge of the normal vibrations of the molecule and of the assignment of the indi- vidual bands in the spectrum to them.For small molecules it is possible to calculate the normal vibrations from force constants bond lengths and atomic weights. This has not been done for heteroarornatic molecules but by using results obtained for simpler molecules (especially benzene and de~terobenzenes~) and by empirical correlation of series of compounds many spectra can be reasonably well explained. For compounds composed of a substituent and an aromatic ring the vibrational motion is concen- trated either in the ring or in the substituent for many of the normal modes ; thus series of related molecules show characteristic absorption bands. Because of the large mass difference between H and C N and 0 the normal vibrations of a heteroaromatic ring can be roughly divided into (a) those in which hydrogen atoms move relatively to the heavy skeleton and (b) those in which each CH or NH group moves as a unit.Substitution of N for CH or of 0 for NH has relatively little effect upon the normal vibrations of class (b) because of the similar masses involved. Vibrations of Predominantly CH Character.-The presence of one hydrogen atom in a planar ring means that there are three normal vibra- tions in addition to those the ring would have without this hydrogen atom. * The fundamentals which are forbidden in the infrared spectrum usually cause lines in the Raman spectrum; the simultaneous use of Raman and infrared data often simplifies the interpretation of both types of ~pectra.~g~J KATRITZKY SPECTRA OF HETEROAROMATrC COMPOUNDS 355 It is found that these vibrations are concentrated in the bond joining the hydrogen atom to the ring and are stretching (I) in-plane bending (11) and out-of-plane bending motions (111).Similarly six CH-type vibrations are associated with two nuclear hydrogen atoms but if both of these are attached to the same planar ring the two vibrations of each class interact to give in-phase (symmetrical) and out-of-phase (asymmetrical) modes shown (IV-IX) for two para-hydrogen atoms. In general if n hydrogen atoms are attached to a planar ring there will be 3n CH-type modes and in each group the n modes will be coupled. The fundamental absorption bands corresponding to the three classes occur in different parts of the spectrum CH stretching near 3000 cm.-l in-plane CH bending at ca.1300-1000 cm.-l and out-of-plane CH bending in the 1000-700 cm.-l region. Vibrations ofpredominantly CC CN etc. Character.-The discussion of six-membered rings is most conveniently based upon the benzene molecule which has 30 (i.e. 3 x 12 - 6) normal vibrational modes 18 of which are of mainly CH character as explained above. The remaining 12 modes are depicted diagrammatically in Fig. 1; because of the high symmetry of benzene some modes are “degenerate” i.e. cause absorption at the same frequency. Similar vibrational modes occur for substituted benzenes pyridines and azines but the degeneracy now disappears because of the lower symmetry.* It is found that the spectral positions of these modes are relatively constant and that they can be conveniently sub-divided into C-C etc.stretching modes (X-XIV) which absorb at ca. 1600-1300 * When the only element of symmetry in an aromatic ring is the plane of the molecule (as in e.g. 2- and 3-substituted pyridines) the symmetry is denoted C, and the vibra- tional modes can be classified as those (X-XVIII) in which this plane of symmetry is preserved (denoted A’) and those (XIX-XXI) in which it is not preserved (denoted A”). When the aromatic ring has an additional plane of symmetry perpendicular to the ring (as in 4-substituted pyridines or 5-substituted pyrimidines) the symmetry is denoted by Czo and the vibrational modes may be divided into (a) those (X XIT XIV-XVII) which preserve both planes of symmetry (denoted A J (b) those (XI XIII XVIII) which preserve the symmetry plane of the molecule (denoted Bl) (c) those (XIX XX) which preserve the symmetry plane perpendicular to the molecule (denoted B2) and ( d ) those (XXI) which do not preserve any symmetry (denoted A2).Pyrazine 1,3,5-triazine and benzene are examples of molecules possessing still higher symmetry; their symmetry is denoted by Drh (or Vh) D,, and Dab re~pectively.~ 356 QUARTERLY REVIEWS FIG. 1. SkeIetul frequencies of benzene (XVI) (XVII) (XVIII) (XIX) ow (XXI) Modes which are degenerate for benzene are bracketed together. The symmetry type position of absorption band and notation4 for benzene are given beneath each mode [an alternative notation of Wilson (Phys. Rev. 1934 45 706) is often used]. qs. is infrared-active; v16 v2 and vI8 are Raman-active; vg v6 v8 and v20 are totally inactive (obtained by indirect means e.g.deuteration). FIG. 2. Skeletal modes for thiophen furan and pyrrole # (XXII) (XXI I I) (xxrv) (XXV) v15& VeA 1 %A 1 v14B1 Thiophen 1590 (8) 1404 (35) 1358 (Ra)a 1252 (50) Pyrrole 1530 (25) 1467 (10) 1384 (5) 1418 (20) Furan 1579 (10) 1486 (60) 1381 (25) ? g+ @ G +Q -Q (XXVI) (XXVII) (XXVIII) (XXIX) (XXX) t VaA 1 V I O & VIA 1 VrsA v23 B2 Thiophen 832 (95) 872 (3) 604 565 453 Furan 994 (170) 872 (75) 724 ? 605 Pyrrole 1144 (lo)* 647 71 1 51OC 565a Lord and Miller's notationlo for the vibrational modes the symmetry type (these molecules are of CZv symmetry; see footnote on p. 3 5 9 and the positions (and where available intensities determined in CHC13 solution) are given below each diagram. The assignments are essentially those given by Lord and Miller,lo Thompson and Temple (Trans.Faraduy Soc. 1945 41 27 and Waddington et aL8 and Lecomte Bull. SOC. chim. France 1946 415). This assignment is not beyond question. Cf. ref. 8. a Observed in Raman spectrum only. Calculated see ref. 10. KATRITZKY SPECTRA OF HETEROAROMATIC COMPOUNDS 357 cm.-l ring-breathing modes (XV XVI) near 1000 cm.-l C-C in-plane bending modes (XVII XVIII) in the 700-600 cm.-l region and C-C out-of-plane bending modes (XIX-XXI) below ca. 550 c1n-l. The discussion of the skeletal modes for five-membered rings is based upon thiophen furan and pyrrole although the assignments for these molecules are not unquestionable. Furan and thiophen have 21 normal modes each and pyrrole has 24; in each case 9 are skeletal and these are shown in Fig.2. Again these modes can be divided into C-C etc. stretching (XXII-XXV) at ca. 1600-1 250 cm.-l the breathing frequency (XXVI) at ca. 1150-800 crn.-l C-C in-plane bending (XXVII XXVIII) at ca. 900-600 cm.-l and C-C out-of-plane bending (XXIX XXX) at ca. 600-450 cm.-l. Characteristic absorption of heterocyclic nuclei* The 3000 cm.-l Region.-The CH stretching vibrations occur in much the same range i.e. ca. 3100-3000 cm.-l as they do for benzenoid compounds; this has been shown inter alia for pyridine,' thi~phen,*>~ pyrrole (XXXI) ,9J0 thiazoles,ll and diazines.12 Few data have been reported on the influence of substituents on the intensity or position of these modes. 0 -1 (XXXI) (xxxr I) (xxx I I I) (XXXIV) The NH stretching frequencies of pyrrolesla and ind0les~~3~5 cause absorption at ca.3450-3400 cm.-l (EA ca. 120) for dilute solutions and at *Throughout this Review positions of bands are quoted in cm.-l units. Data for a class of compounds are often expressed as arithmetical means and standard deviations (the original papers should be consulted for full details as it was not possible to indicate many exceptions and trends here). Precise intensity data are given where available in parentheses after the frequencies; these data are in the form of extinction coefficients of band maxima €A where EA = (l/c/).log~o(~/Zo) c being the molar concentration and 1 the path length in cm. The area under the absorption peak is of greater physical sig- nificance than the peak height,2*B but in complex molecules band overlap often makes area measurements impracticable.Many of the values quoted are for 0.2~-chloro- form solutions measured in a 0.1 mm. cell and are reproducible to an accuracy of about 10%. Brown Chem. Rev. 1958,58,581. Wilmshurst and Bernstein Canad. J. Chem. 1957 35 1183; Corrsin Fax and Waddington et al. J. Amer. Chem. SOC. 1949 71 797. Lebas and Josien Bull. SOC. chim. France 1957 251. Lord J . Chem. Phys. 1953 21 1170. lo Lord and Miller J. Chem. Phys. 1942 10 328. l1 Taurins Fenyes and R. N. Jones Canad. J. Chem. 1957,35,423. l2 Lord Marston and Miller Spectrochim. Acta 1957,9,113 ; Ito Shimada Kuraishi l3 Eisner and Erskine J. 1958 971. l4 Brown Henbest and E. R. H. Jones J. 1952 3172. l6 Millich and Becker J. Org. Chern. 1958 23 1096. and Mizushima J. Chem. Phvs. 1956 25 597. 358 QUARTERLY REVIEWS ca. 3400-3100 cm.-l for solid compounds; the frequency lowering in the solid state has been attributed to intermolecular hydrogen bonding.Within both the ranges quoted electron-withdrawing substituents tend to lower the frequency because of the increased importance of canonical forms of type (XXXII). Substituents in the /%position of indoles (cf. XXXII) have a larger effect on the frequency than those in the a-position but for pyrroles (XXXI) there is little difference; these facts are in accord with the known conjugative abilities of the heterocyclic nuclei at the various nuclear positions.16 Imidazole17 (XXXIII) absorbs at 3 125 cm.-l. The >N+-H group in the pyridinium ion (XXXIV) causes,18 in the solid state one or several “immonium” bands at ca. 2200-1900 cm.-? The 2000-1 650 cm.-l Regisn.-Benzenoid compounds show overtone and combination frequencies (EA - 5) of the out-of-plane CH bending modes in this region;l9 pyridines have been shown to possess analogous bands which depend upon the type of substitution.1ga*20 The 1600-1 350 cm.-l Region for Six-rnembered Rings.-The bands corresponding to modes (X)-(XIII) of benzene occur in this region and are relatively invariant near 1605 1575 1480 and 1430 cm.-l (Table 1).The third band is often found at lower frequencies for electron-acceptor than for electron-donating substituents ; otherwise there is usually little dependence of position on substituent type. The intensities of these bands vary widely but in each series the in- tensities can be correlated with the electronic nature of the substitu- (i) For the monosubstituted compounds EA of the band near 1605 cm.-l can be high for electron-donating groups low for weakly interacting substituents and high again for electron-accepting groups as in mono- substituted benzenes 4-substituted pyridine 1-oxides and 3-substituted pyridines.It can be very high for donors medium for weak substituents and low for acceptors as in 2- and 4-substituted pyridines and 3-substituted pyridine 1-oxides (2-substituted pyridine 1-oxides show intermediate behaviour). For disubstituted compounds EA can vary directly as the algebraic sum of the electronic effects of the substituents (e.g. rneta- ent( s> 21XCW4 l6 Katritzky and Lagowski “Heterocyclic Chemistry” Methuen London 1960. l7 Garfinkel and Edsall J. Amer. Chem. Soc. 1958 80 3807. l8 Witkop J. Amer. Chem. SOC. 1954 76 5597; Experientia 1954 10 420.l* Ref. 1 p. 67; ref. 2 p. 397. l9aC00k and Church J. Phys. Chem. 1957,61 458. ** Podall Analyt. Chem. 1957 29 1423. a1 (a) Katritzky and Lagowski J. 1958 4155; (b) Katritzky and Jones J. 1959 3670; (c) Katritzky and Simmons J. 1959 2058; (d) Katritzky and Simmons J. 1959 2051. 22 (a) Katritzky and Hands J. 1958,2202; (b) Katritzky Hands and Jones J. 1958 3165; (c) Katritzky and Gardner J. 1958 2198. 23 (a) Katritzky and Hands J. 1958 2195; (b) Katritzky Beard and Coats J. 1959 3650; (c) Katritzky and Gardner J. 1958 2192. 24 Katritzky J. 1959 2049. KATRITZKY SPECTRA OF HETEROAROMATIC COMPOUNDS 359 disubstituted benzenes) as the algebraic difference (e.g. para-disubstituted benzenes) or intermediately (ortho-disubstituted benzenes). (ii) In each series of compounds the band near 1575 cm.-l is generally weaker than the preceding band but shows similar intensity variations with the nature of the substituents.(iii) The intensity of the band near 1480 cm.-l is always high for electron- donor substituents (including the N-oxide group) and weak or absent where these are not present. (iv) The intensity of the band near 1430 cm.-l is relatively independent of the nature of the substituent. It is possible to account for these variation^,^^ particularly for the compounds of C, symmetry (cf. footnote on p. 355). The intensity of an infrared band is determined by the product of the rate of change of the dipole moment of the molecule as it undergoes the vibration in question and by the amplitude of that vibration. Vibrations (X) and (XI) are for- bidden in the infrared spectrum of benzene (no associated change of dipole moment) but vibrations (XII) and (XIV) together give a bandof Q-80.When the combined effect of the substituent and/or heteroatom is to produce little change in the charge distribution in the ring modes (X) and (XI) should give weak bands whereas vibrations (XII) and (XIII) should absorb with an intensity approximately half that of benzene itself. If the substituent and/or heferoatom do produce a marked charge disturbance in the ring this disturbance will be symmetrical with respect to the vertical plane of the molecule for CZv compounds and should therefore affect the intensities of (X) and (XII) more than those of (XI) and (XIII). However the relative change in the intensities of (X) and (XI) should be greater than those of (XII) and (XHI) because the former are allowed owing only to the presence of the heteroatom and substituent groups.Thus in CZtl compounds intensities should vary with the amount of disturbance of charge in the ring; strongly for (X) less strongly for (XI) and (XII) and comparatively little for (XIII). Moreover in (X) the intensity should increase more or less continuously with increasing charge disturbance but in (XII) the nature of the variation cannot be predicted. These predictions hold for monosubstituted and para-disubstituted benzenes and for 4- monosubstituted pyridines pyridine 1-oxides and pyridine-boron trichlorides (a measure of the charge disturbance of the ring was obtained from dipole-moment data).25 The 1600-1 350 cm-l Region for Five-membered Rings.-Five-membered heteroaromatic compounds generally show three bands in this region near 1590 1490 and 1400 cm.-l corresponding to modes (XXII-XXIV); available data on the positions of these bands are summarised in Table 2.The effect of the electronic nature of the substituent on the intensity of these bands has been studied for 2-substituted furansZ6 and thiophens ;27 the Katritzky J. 1958 4162. 2o Katritzky and Lagowski J. 1959 657. *' Katritzky and Boulton J. 1959 3500. w o\ 0 TABLE 1. Frequencies (in cm.-l) of ring-stretching bands in the 1600-1350 cm.-l region. Monosubstituted benzenes . . .. .. o-Disubstituted benzenes . . .. .. m-Disubstituted benzenes . . . . .. p-Disubstituted benzenes . . * . .. 1,2,4-Trisubstituted benzenes .. .. Pyridine .. .. ... . .. 2-Substituted pyridines . . . . .. 3-Substituted pyridines .. .. . . 4-Substituted pyridines . . .. . . Polysubstituted pyridines . * . .. Pyridine 1-oxide . .. .. . . 2-Substituted pyridine 1 -oxides . . . 3-Substituted pyridine 1-oxides . . . . 4-Substituted pyridine 1-oxides . . .. Pyridine-boron trihalides . . . . .. 4-Substituted pyridine-boron trichlorides . . Pyridazine . . . . .. .. .. Pyrimidine . . . . * . . . . . Substituted pyrimidines . . * . . . Pyrimidine oxide . . * . .. .. Pyrazine . . .. . . .. .. 1,3,5-Triazine . . . . . . . . 1 1604&3 1407 6 9 1 600+ 1620 1606&6 161658 1599 1615+1585 1595&5 1603&5 1610-1 597 1612 1 6 4 0 4 600 1605 A4 1 6 4 5 4 6 10 ca. 1640 1572 1610 1590-1 555 1577 1584s 1556 1636-1 624 2 1585&3 1577 &4 1586&5 1579 &6 1577&8 1583 1572 5 4 1577 &5 1561 &8 1588-1 564 - 1567hlO 156353 ca.1590 1564&5 1565 1569 1534 1523s - 1565-1 520 - 3 15 1 O+ 1480 15 10+ 1460 1495+1470 1520+1480 15lOj=8 1482 1471 5 6 1485-1465 1 52O+ 1480 1468 15404480 1480k6 1483 -+6 ca. 1500 1444 1461 1480-1400 1468 1490 1555-1490 1493-1 488 - 4 1452&4 1447 -+ 10 1465+1430 1409 433 1456 &4 1441 1433 A5 1421 &4 1415h4 ? 1435AIO 1434&5 1443 5 7 - 1465-1 459 1 4 60-+ 1 4 3 0 1414 1400 1410-1375 1410 1418 1410 r 5 R 4 cc 1 Probably combination bands; see refs. quoted. 9 Found in Raman spectrum only. Arithmetical means and standard deviations are given for each band unless there is a marked correlation between the position of the band and % the electronic nature of the substituent.In such cases a notation of the type 1510i.1475 is used to indicate that electron-donor substituents cause absorption up to ca. 1510 cm.-l and electron-acceptor substituents down to ca. 1475 cm.-l. Compounds in which a heavy atom (e.g. S C1 Br) is directly attached to the ring often absorb at approximately 10 cm.-l shorter wavelength than other related compounds; this is not 2 taken into account in the Tables. In general data are taken from the first reference in the last column. References a Randle and Whiffen Report Conf. Mol. Spectroscopy 1954 Inst. Petroleum paper no. 12. b Katritzky and Lagowski J. 1958 4155. c Katritzky and Jones J . 1959 3670. d Katritzky and Simmonds J. 1959 2058. e Idem J. 1959 2051.J Wilmshurst and Bernstein Canad. J . Chem. 1957 35 1183; Corrsin Fax and Lord J .Chem. Phys. 1953 21 1170. g Katritzky and Hands J. 1958 2202. m h Shinda and Ikekawa Phavm. Bid!. (Japan) 1956 4 192; Shindo ihid. 1957 5 472. i Cook and Church J . Phys. Chem. 1957 61 458. 8 j Katritzky Hands and Jones J . 1958,3165. k Katritzky and Gardner J. 1958 2198. I Sartori Costa and Blasina 2. phys. Chem. (Frank- furt) 1955 4 24. m Katritzky J. 1958 4162; Sartori Costa and Blasina Gazzetra 1955 85 1085. n Shindo Pharm. Bull. (Japan) 1956 4 @ 460. o /dew ibid. 1958,6,117. p Katritzky and Hands J. 1958,2195. q Katritzky Beard and Coats J. 1959 3680. r Katritzky and Gardner J. 1958 2192. s Katritzky J. 1959 2049. t Short and Thompson J. 1952 168. u Wiley and Slaymaker J . Amer. Chem. Soc.! 1957 79 2233. v Brownlie J. 1950 3062; Blout and Fields J . Amer. Chem. SOC. 1950 72 479; Lacher Bitner Emery Seffl and Park J.Phys. Chem. 1955,59 615. w Lancaster and Colthup J. Chem. Phys. 1954,22 1149; Padgett and Hamner J. Amer. Chem. SOC. 1958 80 803. $ .. 2-Substituted furans Polysubstituted furans 2-Substituted thiophens 3-Substituted thiophens Substituted pyrroles Thiazole . . . . Substituted thiazoles Furoxanss . . .. Substituted isoxazoles Substituted furazans; Imidazole . . .. 1,2,3-Triazole .. TABLE 2. Frequencies (cm.-l> for five-membered heteroaromatics. .. .. . . . . . . . . . . . . . . . . .. .. . . . . .. . . . . . . . . . . . . .. .. .. . . . . .. . . ,. .. .. . . .. . . .. .. . . .. .. . . .. . . . . .. .. .. .. .. 1605+1570 ca. 1560 1523 &9 ca. 1530 ca. 1565 1615 ca. 1610 ca. 1610 ca. 1600 ca. 1570 1550 1520 1 5 10+ 1475 ca.1510 1442k12 ca. 1410 ca. 1500 1485 ca. 1500 ca. 1460 ca. 1460 ca. 1425 1492 1450 1 380+ 1400 ? 1354 &-77 ca. 1370 1385 ca. 1380 ca. 1420 ca. 1420 ca. 1380 1451 1410 ? Ref. a b c d d 13 11 11 e 8 h i f §l-Oxa-2,5-diazole 2-oxide. # l-Oxa-2,5-diazole. Tf Another mode presumably (XXV) is shown at 1231 f 10 cm.-l. References a Katritzky and Lagowski J. 1959,657. b Kubota Tetrahedron 1958,4,68; Cross Stevens and Watts J. Appl. Chem. 1957 7 562; Daasch Chem. and Ind. 1958 1113; Cross and Watts ibid. p. 1161. c Katritzky and Boulton J. 1959 3500. d Hidalgo,J. Phys. Radium 1955 16 366; Hochgesang in Hartough’s “Thiophene and its Derivatives” Interscience Publ. Inc. New York 1952. e Boyer Czerniak Gutowsky and Snyder J. Arner. Chem. Sac. 1955 77,4238. f Katritzky et ul. unpublished work.g Milone and Borello Gazzeffu 1951,81,368. h Garfinkel and Edsall J. Amer. Chem. Soc. 1958,80,3807. i Hartzel and Benson ibid. 1954,76 667. TABLE 3 . Absorption bands near lo00 cm.-l assigned to ring-breathing modes. Unsubstituted 2-Substituted 3-Substituted 4-Substituted Monosubstituted benzenesQ ca. 995 (< 1 5)b ? 995&3 (20k15) ? 1015 (145) 1 1015&2 (150&25) 7 994k4 (45k25) 1025k2 (65*35) 993Ifi2 (35-&10y 1030 (45) 993 (45) ca. 990 - ca. 990 1055 { Pyridines Pyridine oxides Pyrimidines 991 a 4 0 7 1 Pyrazine Furan Thiophen . 1022 m 1 ? 994 (170) 101554 (85&15)F ? 832 (95) 823h20 (50h20) ? ? - For references see Tables 1 and 2. In this row the column headings refer to further substitution. * Higher EA for electron-donor substituents. Highe reA for electron- acceptor substituents.363 Benzenes Pyridines Pyridine 1 -oxides Azines TABLE 4. Out-of-plane modes for six-membered rings. Number and orientation of hydrogen atoms. Five Four Four Four Ref. adjacent adjacent 1,2333 f ,27475 r LXIII LIX XIX LXVIII LXXIII XIX - 751&15 697511 751&7 782&9 690k15 817&13 a 749 700 780-740 820-770 730-690 850-790 7 b 759 669 790-750 820-760 680-660 855-820 c A (Not applicable) 760 (vs) 72 1 680 804 (vs) 12 d $ References a Randle and Whiffen Report Conf. Mol. Spectroscopy 1954 Inst. Petroleum paper no. 12. by Podall Analyt. Chem. 1957 29 1423; Shindo and Ikekawa Pharm. Bull. (Japan) 1956 4 192; Shindo ibid. 1957 5 472. c Idem ibid. 1956 4 460; 1958 6 117. d Short and Thompson J . 1952 168. 5 z 2 r; v1 TABLE 5. Out-of-plane modes for Jive-membered rings.Furan Thi ophen Pyrrole Unsubstituted 2-Substituted A r \ LXVI LXVIII LXIV LXIX LXXIV - 837 ( 5 ) 744 925&-9 (55&17) 884&2 (30513)" ca. 800 832 (95)b 710 925&8 (1055) 853 5 7 (75 525) 800 838 (5) 768 - - - For references see Table 2. See footnotes to Table 1. LA higher for electron-acceptor substituents. * Most of this intensity is probably due to the ring-breathing mode see Fig. 2. KATRITZKY SPECTRA OF HETEROAROMATIC COMPOUNDS 365 intensity of all these bands is increased by strongly electron-withdrawing substituents. This is according to expectation since these rings can con- jugate tightly with electron-accepting substituents. FIG. 3. In-plane hydrogen deformation modes. Five aaacent hydrogen atoms. 6 A 4 4 A (xxxv) (xxxvr) (xxxvrr) (xxxvrii) (xxx r x) Monosubstituted Pyridine 1218 1218 1148 (20) 1085 ( < 5 ) 1068 (30) benzenes 1240f8 w 1177f6a 1156f5a 107334 (20f10) 1027f3 (20f10) Four adjacent hydrogen atoms.0 hi) A I ow ortho-Disubstituted benzenes 1269 f 17 (1 5 f 5 ) 1160f4 (20f10)b 2-Substituted pyridines 1279i14 (15h5) 1147f3 (25f10)c 2-Substituted pyridine I-oxides ? 1150f4 (25f10) Pyridazine 1239a 1160 m Thiophen 1283 (3) 1077 (30) Furan 1270 ( ( 5 ) 1137d PyrroIef 1046 (130) 1237d B (XLII) A (XUII) ortho-Disubstituted benzenes 1125 f14 (25 f15) 1033fll (50f40) 2-Substituted pyridines 1093f4 (15fl5) 1048f5 (20klO) 2-Substituted pyridine I-oxides 1106f10 (25f20)C 1044f5 (25f15) Pyridazine 1063 s 1063 Thiophen 909 ? 1032 (15) Furan ?e 1067 (35) Pyrrolef 1015 (150) 1076 (20) 366 QUARTERLY REVIEWS Four hydrogen atoms in the 1,2,3,5-positions.rnetu-Disubstituted benzenes 3-Substituted pyridines 3-Substituted pyridine oxides Pyrimidine meru-Disubstituted benzenes 3-Substituted pyridines 3-Substituted pyridine oxides Pyrimidine 8 (XLIV) 1278f12 (30f15) ca. 1190 ? ? 1 220 1096f7 (30f15) 1103f5 (20f5) ? 1140 6 (W 1157f5 (cu. 1124f5 (20f10) 1 156f2C 1165 Four hydrogen atoms in the 1,2,4,5-~ositions. B (XLVIt) 1076f7 (ca. 25)g 1038f7 (20f5) ? 1021 para-Disubstituted benzenes CSubstituted pyridines CSubstituted pyridine 1-oxides Pyrazine para-Disubstituted benzenes 4-Substituted pyridines Mubstituted pyridine I-oxides Pyra zine B,g @ I ) (XLVI I I) 1258 f 11 (35 f25) ca. 1220 m (obscured) 1 232d 62 @,) (L) 1111rf7 (25f15)h ? 1101f7 (20f10) 1148 (s) A (A,) (XLI x) 1175f6 (var.) ? 1169f5 (200f50) 11186 81 (4) (LO 1013f5 (var.) 1067f3 (25f10) 1033rt5 (55kl5) 1067 (vs) KATRITZKY SPECTRA OF HETEROAROMATIC COMPOUNDS 367 Three adjacent hydrogen atoms.Vicinal trisubstituted benzenes ca. 1200 ? 1160f5 (w) 1073f10 (s) 2-Substituted furans ca. 1220 1158f7 (95f25) 1076rt3 (45f20) 2-Substituted thiophens 3 1081$13 (10f5) 1043&7 ( 5 5 3 4 0 ) For references see Tables 1 and 2. a These bands have EA-5 for weakly interacting substituents but become stronger for electron-donor substituents. Higher shown by compounds with two electron-donor substituents. Higher shown by compounds with electron-donor substituents. Found in Raman spectrum only. Band originally assigned (Thompson and Temple Trans. Faraday SOC. 1945,41,27) to this mode was later* suggested to be an out-of-plane mode.Agreement is poor for pyrrole which may indicate coupling of the CH and NH in-plane modes or an incorrect assignment. High €A shown for compounds with electron-acceptor substituents. High E A for nitro-substituents. The 1250-1000 crn.-l Region.-Heterocyclic compounds show series of characteristic bands in this region which are assigned to in-plane CH deformations and ring-breathing modes (XV XVI XXVI). Fig. 3 gives the form of the in-plane CH modes for several types of compound and a tentative assignment of the bands; data for substituted benzenes are included for comparison. There is an overall similarity in the positions of these bands for compounds with the same number of hydrogen atoms in the same orientation. Table 3 records bands tentatively assigned to ring-breathing modes.As an empirical correlation it can be noted that these bands are of appreciable intensity for six-membered rings if CH groups or nitrogen atoms with lone electron pairs are present in the 1,3,5-positions. Two bands are sometimes shown by unsubstituted compounds. The 1 000-600 crn.-l Region.-Out-of-plane CH deformations (which are depicted in Fig. 4 for several orientations of the hydrogen atoms) and in-plane ring deformations could cause absorption in this region. The out-of-plane CH deformations of symmetry type A in CZv compounds (see footnote on p. 355) are forbidden in the infrared spectrum; all six-membered rings are of at least approximate C, symmetry and these bands are there- fore weak or absent. Of the B type vibrations the in-phase “umbrella” mode always results in a very strong band but the other vibrations are usually weak and hard to detect for six-membered ring compounds.Tables 4 and 5 contain data on six- and five-membered ring compounds respectively. The relation between the nature of the substituents and the position of a band within the range given in Table 4 is not clear; e.g. changing methyl to ethyl displaces the bands for pyridines to higher Mono- and meta-di-substituted benzenes show28 an out-of- Randle and Whiffen Conf. Mol. Spectroscopy 1954 Imst. Petroleum paper No. 12. 27aShindo and Ikekawa Pharrn. Bull. (Japan) 1956,4 192. 368 QUARTERLY REVIEWS plane ring deformation mode (cf. XIX) near 700 cm.-l and corresponding heterocyclic compounds show similar bands (Table 4). Little is known of the remaining ring deformation modes which cause absorption below 600 cm.-l.FIG. 4. Out-ofplane CH modes. I 8 2 (LVI) 6 2 (LVII) A,(LVIII) B2 (DVI) A2(UVI I) 6 (WXI) B (LXXII) Infrared Spectra of Heterocycles with Condensed Rings.-A treatment of the above type is not yet possible for condensed-ring compounds because of the paucity of data. However it has been shown that each type of sub- stituted naphthalene causes absorption in characteristic narrow regions of the spectrum,29 and that the out-of-plane CH vibration of polycyclic benzenoid compounds and of heteroaromatic compounds can be classified according to the number of adjacent hydrogen atoms The de- tailed assignment of naphthalene has been Qwin01ines~~ show three bands near 1600 crn.-l five bands in the 1500- 1350 cm.-l region and in- and out-of-plane CH deformation frequencies in the 1300-700 cm? region; the last are determined by the number of nuclear hydrogen atoms present and can be correlated with the corres- ponding naphthalenes.Indoles14 show bands at ca. 1460 1420 and 29 Hawkins Ward and Whiffen Specfrochim. Acta 1957,10,105; Werner Kennard and Rayson Aust. J. Chem. 1955 8 346. 30 Groenewege Spectrochim. Acta 1958,11 579; Cannon and Sutherland ibid. 1951 4 373. 31 Mitra and Bernstein Canad. J. Chem. 1959 37 553 and references therein. 32 Katritzky and Jones J. in the press; Shindo Pharm. Bull. (Japan) 1956 4 292; Karr Estep and Papa J. Amer. Chem. Soc. 1959,81 152; Braunholtz and Mann J* 1958 3368. KATRITZKY SPECTRA OF HETEROARQMATIC COMPOUNDS 369 1350 cm.-l; quinazoline~~~ absorb at 1628-1612 1581-1566 and 1517-1478 cm.-l.Purines34 show strong bands in this region but some of them are due to v C=O and NH deformation modes. Effects of heterocyclic rings on the characteristic substituent vibrations* In general substituents show the same characteristic bands regardless of whether they are present in heterocyclic benzenoid or even aliphatic compounds. The following discussion is an attempt to point out the chief changes in these vibrations caused by heterocyclic nuclei. The effect is largest when strong electronic interaction occurs between the ring and the substituent and the magnitude of the effect can often be correlated with the degree of the interaction (e.g. by using the Hammett equation). It is convenient to use the corresponding substituted benzene as a standard.Two common situations are illustrated. If an atom with a lone electron pair is adjacent to the ring as in (LXXV) the double-bond character of the ring-X bond increases with the electron-acceptor character of the ring (cf. LXXVI) and therefore the compound absorbs at higher frequencies. { y x - -{Ti [ y x = 2 (LXXV) (Lxxv I) ( W X V I I) If a multiple bond is adjacent to the ring (cf. LXXVII) (with Y more electronegative than X) increasing the electron-donor properties of the ring will increase the single-bond character of the XY bond and cause absorption to occur at lower frequencies. Tautomerism-Heterocyclic compounds which contain hydroxyl mercapto- amino- or acylamino-groups a or y to ring nitrogen atoms are potentially tautomeric (cf. LXXVIII-LXXXV). Infrared spectra have F \NJ% ( U X V I I I) (LXXIX) (LX x x) (LXX x I) G O H G A S H H G O H *Throughout this Review Py PyO Ph T Pyrr and F denote monosubstituted pyridine pyridine 1 -oxide benzene thiophen pyrrole and f uran rings respectively and the preceding numeral denotes the position of nuclear substitution.33 Culbertson Decius and Christensen J. Amer. Chm. Soc. 1952 74 4834. 34 Blout and Fields J. Arner. Chem. SOC. 1950 72 479; Lacher Bitner Emery Seffl and Park J. Phys. Chem. 1955,59 615. 370 QUARTERLY REVIEWS been valuable in investigating this tautomerism e.g. of amino-,j5 h y d r ~ x y - ~ ~ me~capto-,~’ and acylamino-pyridine~,~~ -pyridine l - ~ x i d e s ~ ~ and -pyrimidine~.~O In work of this nature the results are most conclusive when the spectrum of the potentially tautomeric compound is compared with the spectra of derivatives of both possible forms the structures of which are “fixed” by alkylation.Thus the aminopyridines show the bands expected for NH groups and not those for -NH- and =NH groups (cf. LXXXIII) and the acylamino- pyridines show the bands of the particular acylamino-group ; moreover all these compounds show the absorption characteristic of the appropriately substituted pyridine ring and not that of a pyridone-like molecule. The pyridones show no bands attributable to a hydroxyl group and their spectra resemble those of their N-methyl derivatives more than those of the 0-methyl derivatives e.g. 4-methoxypyridine but not pyrid-4-one shows the characteristic bands of the 4-substituted pyridine nucleus indicating that the latter does not exist as 4-hydroxypyridine ; similar considerations apply to the pyridthiones.Pyridones and pyridthiones also show bands attributable to the C=O and C=S bonds and similar to those in the corresponding N-methyl derivatives.* Recent extensive work using diverse physical methods indicates that in most cases (at least in six-membered rings) amino- and acylamino- compounds exist predominantly as such but hydroxy- and mercapto- compounds exist in the tautomeric carbonyl and thiocarbonyl forms.16 These facts should assist in the future interpretation of infrared spectra. Carbonyl Compounds.-The position of the C= 0 stretching frequency in compounds (LXXXVI) gives a measure of the electron-releasing ability of the groups X and Y; the lower the band position the greater is the (LXXXWI) O=C-CH ( u x x I x) (LXXXVI) (UXXVIII) single-bond character and the importance of form (LXXXVII).the groups is kept constant and the other varied through a heterocyclic rings the relative electron-donating ability of these rings at If one of series of * Bellamy’s method41 of observing the shifts in band positions from a solvent capable of hydrogen bonding with the C=O or C=S (e.g. CHC1,) to one incapable of so doing (e.g. C2C14) is useful in the identification of these bands.42 35 Angyal and Werner J. 1952 2911 ; Goulden J. 1952 2939. 3u Mason J. 1957 4874. Jones and Katritzky J. 1958 3610. Jones and Katritzky J. 1959 1317. 39 Gardner and Katritzky J. 1957 4375. *“Short and Thompson J. 1952 168; Brownlie J. 1950 3062; Brown and Short 41 Bellamy et al. Trans. Faradav Soc. 1958 53 1120; 1959,54 14.42 L J Bellamy personal communication; Katritzky and Jones J. in the press. J. 1953 331; Brown Hoerger and Mason J. 1955 4035. KATRITZKY SPECTRA OF HETEROAROMATIC COMPOUNDS 371 various nuclear positions is determined. Mea~urementsl~~~~ of methyl and ethyl esters aldehydes and methyl ketones (LXXXVI Y = OMe OEt H Me) show that the order is 2-Pyrr - 3-Pyrr > 2-T > 3-T > Ph > 4-PyO - 3-Py > 4-Py > 3-PyO > 3-Py-BCl - 4-Py-BCl,; this is in agreement with other evidence.16 When the carbonyl group is attached directly to a nitrogen atom of a heterocyclic ring as in (LXXXVIII) the C=O stretching frequency increases and the C-N stretching frequency decreases as the ring becomes increasingly electron-attracting i. e. pyrrole < iinidazole < triazole < tetra~ole.~~ In compounds of type (LXXXIX) the heterocyclic ring and the C=O group “compete” for the lone electron pair on the nitrogen atom giving a measure of the electron-accepting power of the heterocyclic ring Ph < C s N Stretching Frequency.-As the ring to which a cyano-group is attached increases in electron-donating character the stretching frequency decrease~,~~ and the intensity increase^.^^^**^ NO a Frequencies.-The asymmetricaP and symmetricalMb NOz stretch- ing modes absorb at higher frequencies as the heterocyclic rings to which the group is attached become increasingly electron-demanding.Ring-0 and Ring-N Stretching Modes.-The positions of the C-0 bands near 1300 cm.-l for ethoxyl and methoxyl compounds indicate the following order of electron-accepting ability Ph < 3-Py < 4-Py - 2-Py < 4-PyO - 2-PyO - ~-PY-BC~,,*~ which is in agreement with chemical evidence.ls For amino- and dimethylamino-compounds the positions of the C-N bands near 1280 and 1350 cm.-l respectively are similarly related to the electron-accepting ability of the ring.60 NH Stretching Frequencies.-The asymmetrical and symmetrical N-H stretching modes for amino- and the N-H stretching mode in alkylamino- heterocyclic compounds increase in both wavelength and intensity as the ring becomes more strongly electron-a~~epting.~~~~~ For compounds of type (XC)t5 just as in the corresponding a~etanilides,~~ the position and 4s (a) Katritzky Monro Beard Dearnaley and Earl J.1958 2182; (6) Katritzky Lagowski and Beard unpublished work Gronowitz and Rosenberg Arkiv Kemi 1955,8,23; Shindo Pharm.Bull. (Japan) 1957,5,472; 1958,6 117. Otting Gem. Ber. 1956 89 1940; Staab Otting and Ueberle 2. Electrochem. 1957,61,1000; Bellamy Spectrochim. Acta 1958,13,60. 3-py < 4-py - py0.45 40 Katritzky and Jones J. 1959 2067. 46 Shindo Bull. Phurm. (Japan) (a) 1957,5,472; (b) 1958,6,117. 47 Sensi and Gallo Gazzetfa 1955 85 235. 49 Katritzky and Coats J. 1959 2062. so Katritzky and Jones J. 1959 3674. 81 Mason J. 1958 3619; Short J. 1952 4584; Bellamy and Williams Spectrochim. Russell and Thompson J. 1955,483; Spectrochim. Acta 1956,8,138; Moccia and (a) Katritzky Rec. Trav. chim. in the press; (b) Katritzky and Simmons unpub- lished work. Acta 1957,9 341; Osborn Schofield and Short J. 1956,4191. Thompson ibid. 1957 LO 240. 372 QUARTERLY REVIEWS intensity of the NH stretching band are little affected by the nature of the ring unless intramolecular hydrogen bonding can occur as in (XCI).0 (XCI 11) “HOMe (XCIV) (xcv) N-Oxide Freq~encies.~~v~~?~~-The Nf-0- stretching frequency in N- oxides (XCII) occurs at ca. 1300-1200 cm.-l (EA ca. 280) but within this range the frequency is increased by the presence of electron-accepting substituents in the ring; these substituents increase the importance of canonical forms of type (XCIII) and therefore increase the double-bond character of the N-0 bond. However alkyl substituents in the 3-position also cause a shift towards higher frequencies; the reason for this is not known but similar effects are observed for the C-N and C-0 stretching frequencies in meta-substituted anilines and phenyl methyl and ethyl ethers.45,49$50 The Nf-0- stretching band shows large shifts to lower frequencies when measured in a solvent with which hydrogen-bonding can occur (e.g.XCIV). N-Oxides show another characteristic band near 850 cm.-l which was originally assigned54 to the N+-O- stretching mode but may be the N+-0- in-plane bending mode.46b fur ox an^^^ (XCV) show bands at 1360-1300 1190-1150 and 890- 840 cm.-l which are analogous to the above N-oxide frequencies. Summary of the principal correlations for use in diagnostic work The presence of an aromatic ring is indicated by the appearance of four bands for six-membered rings or three for five-membered rings in the 1600-1350 cm.-l region of the spectrum. The positions of these bands give information concerning the nature of the ring and the substitution pattern.The number and positions of the in- and out-of-plane CH de- formation bands in the 1200-700 cm.-l region indicate the number of ring hydrogen atoms and thus the substitution pattern. High intensity of the band near 1490 cm.-l for six-membered rings 63 Sartori Costa and Blasina Gazzetia 1955 85 1085; 2. phys. Chem. (Frankfurt) 66 Boyer Czerniak Gutowsky and Snyder J. Amer. Chem. SOC. 1955 77,4238. 1955,4,24; 1956 7 123. Ito and Hata Bull. Chem. SOC. Japan 1955 28 353. KATRITZKY SPECTRA OF HETEROAROMATIC COMPOUNDS 373 indicates the presence of eIectron-donor substituents ; the intensities of the bands at ca. 1600 and 1580 cm.-l are directly related to the amount of disturbance to the charge symmetry of the ring. The electron-donor and/or electron-acceptor powers of the ring are indicated by the positions and intensities of the characteristic substituent vibration modes. I thank Sir Alexander Todd F.R.S. for his encouragement and Dr. N. Sheppard for reading the manuscript and for very helpful discussion.

ISSN:0009-2681    

DOI:10.1039/QR9591300353

出版商:RSC    

年代:1959

数据来源: RSC