Molecular Complexes of Water in Organic Solvents and in the Vapour Phase By Sherril D. Christian Ahmed A. Taha and Bruce W. Gash DEPARTMENT OF CHEMISTRY THE UNIVERSITY OF OKLAHOMA NORMAN OKLAHOMA 73069 1 Introduction Much of the literature of aqueous solution chemistry concerns solvation or hydration phenomena and the interpretation of these phenomena in terms of water-dissolved solute interactions. Previous reviews have treated the inter- action of water molecules with ionic salts polar organic molecules and com- pounds of biological imp0rtance.l The interaction of water molecules with each other in liquid water and in aqueous solutions has been investigated by virtually all the applicable methods of physical chemistry.1b,2 That formidable problems arise in interpreting physical and chemical information about concentrated solutions is evident from the variety of mutually contradictory models which have been proposed for liquid water and aqueous ~olutions.~ It has long been recognised that solute-solute molecular interactions can more readily be elucidated than solvent-solvent or solvent-solute interactions.Fortunately non-ionic solute species ordinarily follow Henry’s law at concen- trations up to at least several mole per~ent.~ When spectral methods are employed Beer’s law generally applies to the individual solute species in a comparable concentration range. Similarly other specific properties of a dissolved solute are constant in the dilute region e.g. partial molar volume energy and heat capacity; dipole moment; chemical shift in proton magnetic resonance o R.A. Robinson and R. H. Stokes ‘Electrolyte Solutions’ 2nd edn. Butterworths Scientific Publications London 1959; b J. L. Kavanau ‘Water and Solute-Water Interactions’ Holden-Day Inc. San Francisco 1964; c F. Franks ed. ‘Physico-Chemical Processes in Mixed Aqueous Solvents’ Heinemann Educational Books Ltd. London 1967; d F. Franks and D. J. G. Ives Quart. Rev. 1966 20 1 ; e R. W. Gurney ‘Ionic Processes in Solution’ McGraw-Hill Book Co. Inc. New York 1953;fH. E. Whipple ‘Forms of Water in Biologic Systems’ Ann. New York Acad. Sci. 1965 125 (2)’ 249. a C. Pimentel and A. L. McClellan ‘The Hydrogen Bond’ W. H. Freeman and Co. San Francisco 1960. a H. S . Frank and Wen-Yang Wen Discuss. Faraday Soc. 1957 24 133; b H. S. Frank Proc. Roy. SOC. 1958 A . 247 481; c L. Pauling and R. E. Marsh Proc.Nat. Acad. Sci. U.S.A. 1952 38 112; d W. F. Claussen J Chem. Phys. 1951,19,259 1425; e G. Nkmethy and H. A. Scheraga ibid. 1962,36 3382 3401 ; f G. Nemethy and H. A. Scheraga J. Phys. Chem. 1962,66 1773; g H. S. Frank and A. S . Quist J. Chem. Phys. 1961,34,604; h M. D. Danford and H. A. Levy J. Amer. Chem. SOC. 1962,84,3965; i J. A. Pople Proc. Roy. Soc. 1951 A 205 163; j B. E. Conway Ann. Rev. Phys. Chem. 1966,17,481; k W. A. P. Luck in ‘Physico-Chemical Processes in Mixed Aqueous Solvents’ ed. F. Franks Heinemann Educational Books Ltd.. London 1967. J. H. Hildebrand and R. I,. Scott ‘The Solubility of Non-electrolytes’ 3rd edn. Dover Publications Inc. New York 1964 29. 20 Christian Taha and Gash (lH n.m.r.) studies and many others. Advantage has been taken of the constancy of parameters such as these in numerous studies of molecular interactions in dilute non-electrolyte solutions and in the vapour phase.lbBa As early as 1890 Beckmann inferred from cryoscopic data that benzoic acid dissolves in benzene primarily as the dimer;6 in contrast there is still considerable discussion regard- ing the size and nature of the kinetic units in pure associated liquids.It should be emphasised that investigations of molecular complex formation in dilute solution yield information about the interactions of solute molecules which are solvated not free as they would be in the vapour phase. For example the equilibrium 2PhC0,H = (PhC02H)2 is strongly affected by solvents;s as the medium is changed from vapour to cyclohexane to chloroform solvation more and more effectively opposes the formation of the dimer.In comparing hydrogen-bonding interactions in dilute organic solutions with those in either aqueous solutions or the vapour state methods are needed for predicting the influence of solvents on hydrogen-bonding equilibria. The present Review will summarise what is known about the molecular complexity of water in the vapour state and in dilute solution in non-polar and polar organic solvents; both the behaviour of water as an individual solute and its interactions with other polar solutes will be considered. In order to relate information about the solute properties of water to results for aqueous solutions a discussion will be given of methods for correlating effects of solvation on molecular complex formation reactions. 2 The Molecular Complexity of Water in Dilute Solution Water vapour is acknowledged to be virtually ideal at least at pressures less than ca.90% of saturation in the vicinity of room temperature.' High-tem- perature PVT8 and i.r. spectrala data show however that water vapour deviates considerably from ideality and that aggregates exist in significant concentrations at temperatures above loo" near saturation. Statistical mechanical calculations used to correlate the second virial coefficient of water vapour and the lattice energy of ice indicate that the distance of contact between water molecules in the vapour is comparable to the intermoIecular distance of water molecules in ice.l0 Two other methods have provided information related to the association of water in vapour or in an inert matrix.Mass spectrometric measurements of irradiated water vapour have led to the determination of relative concentrations of the hydrates H+(H20)n and thermodynamic constants for the reactions E. Beckmann 2. phys. Chem. (Leipzig) 1890 6 437. a Y . I'Haya and T. Shibuya Bull. Chem. SOC. Japan 1965 38 1144; b G. Allen J. G. Watkinson and K. H. Webb Spectrochim. Acta 1966 22 807. ' E. N. Dorsey ed. 'Properties of Ordinary Water-Substance' Rheinhold Publishing Corp. New York 1950 p. 54. * a Cr. S. Kell G. E. McLaurin and E. Whalley J. Chem. Phys. 1968,48 3805; b G. S. Kell G. E. McLaurin and E. Whalley ibid. 1968 49 2839. lo J. S. Rowlinson Trans. Faraday SOC. 1949,45 974; J. S. Rowlinson ibid. 1951,47 120. W. A. P. Luck and W. Ditter Ber. Bunsengesellschaft Phys. Chem. 1966 70 11 13. 21 Molecular Complexes of Water in Organic Solvents and in the Vapour Phase H,O + H+ (H20)n-1 = H+ ( H 2 0 ) n in which 1 < n < 8.11 The matrix isolation method (in which an associating vapour is diluted with an inert gas quickly frozen and examined spectrally) indicates that water dimers (probably cyclic) and higher polymers exist in a nitrogen matrix at 20K.l2 it is generally believed that water dissolves in the aliphatic and aromatic hydrocarbons and CCl primarily as the monomer.Cryo~copic,~~ spectral,16 vapour pressure,le dielectric constant,17 tracer,ls and partial molar volume1B studies indicate that only small concentra- tions of associated species are present even at concentrations approaching saturation. Activity data for water dissolved in several hydrocarbons and chlorinated hydrocarbons are shown in Figure 1.The activity aw is the ratio of water partial pressure to the saturation pressure of pure water at 25"; fw represents the total or analytical concentration of water. Data are displayed in a log-log plot; the fact that the points for the non-polar solvent systems fall nearly on straight lines drawn with unit slopes implies that water is essentially monomeric in these systems. There is evidence for positive curvature in the plots for aromatic solvents near saturation but only a few percent of the water molecules appear to be in associated forms. In slightly polar organic solvents such as the partially chlorinated hydro- carbons water is somewhat polymerised. Vapour pressure,labP2O i.r.,21 and lH n.m.r.22 data permit calculation of self-association constants for water in 1 ,Zdichloroethane (DCE) 1,1,2,2-tetrachIoroethane (TCE) and 1,2,3-trichIoro- propane.Spectral evidence indicates that trimeric species are favoured over dimers although one of the two vapour pressure investigations of water in DCE yielded data which can equally well be interpreted as indicating the presence of dimers or trimers,20a and the other vapour pressure results indicate that the most IikeIy associated species are trimers or tetramers.lsb Ca. 8-15% of the dissolved water is associated at saturation at 25" ; trimer formation constants vary in the range 3-5 1 mole-%. Water activity data for DCE and TCE are included in Figure 1. In spite of some reports to the l1 a P. Kebarle S. K. Searles A. Zolla J. Scarborough and M. Arshadi J. Amer. Chem.Soc. 1967,89,6393; b M. G. Inghram and R. Gomer Z. Naturforsch. 1955,10a 863. l2 M. Van Thiel E. D. Becker and G. C. Pimentel J. Chem. Phys. 1957,27,486. 13a M. Gordon C. S. Hope L. D. Loan and Kyong-Joan Roe Proc. Roy. SOC. 1960 A 258 215; b T. Ackerman Z . phys. Chem. (Frankfurt) 1964 42 119; c A. Risbourg and R. Liebaert Compt. rend. 1967 264 C 237. l4 J . M. Peterson and W. H. Rodebush J. Phys. Chem. 1928,32,709. l5 a E. Greinacher W. Luttke and R. Mecke 2. Elektrochem. 1955,§9,23; b S. D. Christian H. E. Affsprung J. R. Johnson and J. D. Worley J. Chem. Educ. 1963,40,419. leS. D. Christian H. E. Affsprung and J. R. Johnson J . Chem. SOC. 1963 1896; b J. R. Johnson S. D. Christian and H. E. Affsprung ibid. (A) 1966 77. l7 M. D. Gregory H. E. Affsprung and S. D. Christian J. Phys.Chem. 1968,72 1748. l8 J. W. Roddy and C. F. Coleman Talanta 1968 15 1281. ID W. L. Masterton and H. K. Seiler J . Phys. Chem. 1968,72,4257. 2o a W. L. Masterton and M. C. Gendrano J . Phys. Chem. 1966 70 2895; b R. L. Lynch Ph.D. Dissertation University of Oklahoma in preparation. 21 C. Jolicoeur and A. Cabana Canad. J . Chem. 1968,46 567. zeT. F. Lin S. D. Christian and H. E. Affsprung J. Phys. Chem. 1965. 69 2980. 22 Christian Taha and Gash 0.10 0.08 0.06 0.04 0.03 0.02 2 0.0 1 0.008 0.006 0.004 0.003 0.002 0.00 1 0.1 0.2 0.3 0.4 0.6 0.8 1.0 QW Figure 1 Formal solubility of water at 25" in 1 ,ZdichIoroethane 1,1,2,2-tetrachIoroethane benzene toluene cyclohexane-J. R. Johnson Ph.D. Dissertation University of Oklahoma 1966; diphenylmethane 1,2,3-trichloropropane-R. L .Lynch Ph.D. Dissertation in prepara- tion carbon tetrachloride-R. D. Grigsby Ph.D. Dissertation University of Oklahoma 1966 23 Molecular Complexes of Water in Organic Solvents and in the Vapour Phase Activity-concentration data have been obtained for water in numerous slightly volatile polar organic solvents for which the saturation concentration of water exceeds 1 mole The solvents fall into two classes those for which an assumed monomer-dimer equilibrium is adequate to correlate activity data and those for which monomer-trimer equilibrium provides a better fit. Introducing equilibrium constants for larger aggregates does not improve the correlation. Table 1 summarises formation constants at 25" for the dimer and trimer in Table 1 Solubility and association of H20 in organic solvents Solvent Tributylamine NN-Dimethylaniline 1,1,2,2-Tetra- doroethane 1,2-Dichloroethane Nitrobenzene Di butylphthalate NN-Dimethy lcyclo- hexylamine Methylphenylketone Cyclohexanone Aniline Benzyl alcohol KH (torr 1 mole- 504.0 435.1 256.2 218.8 182.3 128.3 51.9 43.6 21.8 18.1 11.9 fw" r Kw,* -l) (mole 1-l) (1 mole- 0-0522 10 0.0668 19 0.1010 8-4 0.1262 14.2 0.1666 21 1.054 0.2440 20 0.68 13841 75 0-9840 41 2.4590 50 2-9240 50 3.812 355 Kw,* Ref.16-11 23 26-39 23 -l) (1 mole-a) 3.60 16b 4.60 16b 23 23 4.92 23 0.776 23 0-289 23 0.194 23 0.046 23 *Formation constant for associated species from monomeric water. several solvents. Included are Henry's Law constants for the water monomer KH; the approximate formal concentration of water at saturation in each solvent fw0; and the approximate percent of the dissolved water in associated forms at saturation r.As KH decreases and fwo increases the association con- stants tend to decrease but in spite of this r is generally greater in the more reactive solvents. Thus mass action is more effective in promoting association than is solvation effective in preventing it. Little is known about water association in polar solvents which are completely or nearly completely miscible with water. Approximate values of dimerisation constants for water in several proton-accepting solvents have been inferred from limiting slopes of plots of proton chemical shift vs. water c0ncentration.~4 To explain proton-exchange rate data for water in proton-accepting solvents it has been proposed that solvated cyclic trimers are present at higher water concentrations and that protons transfer within the ring by a concerted mech- ani~m.*~b lH N.m.r.and molar volume data are available for the system H,O- ethylene oxide (EO) at 0" over the entire concentration range.as Water protons tend to bond to lonepair electrons of other water molecules in preference to the 23 J. R. Johnson S. D. Christian and H. E. Affsprung J. Chem. SOC. (A) 1967,1924. 24a G. Mavel Compt. rend. 1959 248 1505; b J. R. Holmes D. Kivelson and W. C. Drinkard J . Amer. Chem. SOC. 1962,84,4677. D. N. Glew H. D. Mak and N. S. Rath Canad. J. Chem. 1967,45,3059. 24 Christian Taha and Gash EO oxygen lone-pair electrons. The linear water trimer appears to be an important associated species at mole fractions of water less than ca. 0.11. At higher water concentrations chain-branching cross-linking and formation of cyclic 'aggregates predominate until ultimately the three-dimensional water networks characteristic of pure liquid water are formed.Density and Raman spectral data for H,O-acetic acid have been interpreted by assuming the presence of water dimers as well as hydrated acid species.26 3 Water-Polar Solute Interactions in Dilute Solution Early evidence for polar solute-water interactions was provided by Bodtker who noted that the solubility of water in ether increased as oxalic acid was added;27 he surmised that a compound forms between water and the acid. Before 1940 there were several reports of the effect of dissolved water on results of cryoscopic,2s partition,2a and s o l ~ b i l i t y ~ ~ ~ ~ ~ ~ studies of polar organic mole- cules.sl In particular S z y s z k o ~ s k i ~ ~ ~ calculated dimer formation constants for several aromatic carboxylic acids in benzene from partition and acid solubility data; he showed that the apparent association constant calculated by ignoring the effect of dissolved water was ordinarily considerably smaller than the corrected constant in one case by a factor of nearly 1:50.Later investiga- indicated that hydrates of both the acid dimer and the monomer are probably present in wet benzene solutions of chlorinated aliphatic carboxylic acids. Lassettre31 reviewed the early partition results and cautioned that the technique should not be employed for determining association constants unless corrections are made for water-polar solute interactions in the organic phase.In spite of this warning there have been numerous recent attempts to use the uncorrected partition method to study self-association; for certain systems however reasonable formation constants are obtained with the method,33 presumably owing to a fortunate cancellation of errors. The techniques employed to investigate molecular complexes of water in dilute solution resemble those usually employed in hydrogen bond studies2 However complications arise in using solvents such as benzene carbon tetra- chloride and the aliphatic hydrocarbons which dissolve less than 0.5 mole 28 J. J. Kipling J. Chem. SOC. 1952 2858. 27 E. Bodtker 2. phys. Chem. (Leipzig) 1897 22 505. a M. Rozsa 2. Elektrochem. 1911 17,934; b R. P. Bell and M. H. M. Arnold J. Chem. SOC. 1935 1432. as a S. Horiba Mem. Coll.Sci. Univ. Kyoto 1914,1,49; b B. U. Szyszkowski Z.phys. Chem. (Leipzig) 1928 131 175. 30 a 0. N. Lewis and G. H. Burrows J. Amer. Chem. SOC. 1912,34,1515; b L. A. K. Staveley J. H. E. Jeffes and J. A. E. Moy Trans. Faraday SOC. 1943,39,5; c E. Cohen and W. D. J. van Dobbenburgh 2. phys. Chem. (Leipzig) 1925 118 37; d E. Cohen and S . Miyake Verlag Akad. Wetenschapen Amsterdam 1925 34 933 ; e J. H. Hildebrand Science 1936 83 21. 81 E. N. Lassettre Chem. Rev. 1937 20 259. 32 R. P. Bell 2. phys. Chem. (Leipzig). 1930,150 A 20. 33 a M. Davies and H. E. Hallam J. Chem. Educ. 1956,33 322; b M. Davies and D. M. L. Griffiths 2. phys. Chem. (Frankfurt) 1954 2 352; c M. Davies P. Jones D. Patnaik and E. A. Moelwyn-Hughes J . Chem. Soc. 1951 1249. 25 Molecular Complexes of Water in Organic Solvents and iii the Vapour Phase percent water.It is difficult to prepare homogeneous solutions of water in a non-polar solvent by direct mixing of the liquid components; even when con- siderably less water is added than is required to saturate the solvent vigorous stirring for 24 h will not promote dissolution of all the drops.l7 This problem may account for some of the discrepancies among literature reports on the properties of water in dilute s ~ l u t i o n . ~ ~ ~ ~ @ The solute isopiestic method has been employed to prepare homogeneous solutions of water in organic solvents at known water activities; vapour phase equilibration of benzene or CCI with an aqueous solution of known water activity occurs within only a few hours even without stirring and without removing air from the equilibrator v e s ~ e l .~ ~ ~ J Manometric techniques are convenient for measuring activities of water in the very dilute solution range (formal water concentrations 0.00005 to 0-00300); one method incorporates a mercury-covered sintered-glass disc inlet valve formerly used in mass spectrometer ~ y s t e m s . ~ ~ ? ~ ~ A. Treatment of Data.-Simultaneous partition water solubility and isopiestic measurements for a given organic solvent-polar solute system have been used to investigate hydration equilibria. To illustrate the rationale of the method we shall outline the computations involved in studying a hypothetical compound which is present in the organic phase as both monomers and dimers (hydrated and unhydrated). Assume that the solute A is distributed between water and the organic solvent and that the concentration of A in the aqueous phase C A ~ is directly proportional to the activity of A.The analytical or formal concentration of A in the organic phase may be written f A = CA (1 + K11CW + K12W2 + - . ) 2CA2 (K20 -I- & ~ C W + K2zw2 . . ) (1) or f A = KDcAw (1 + K1lm f K12Wa + ) f 2KD2CAw2 (Kzo + K ~ ~ c w + K22m2 + . - ) (2) where CA and cw are concentrations of monomeric A and water (W) respectively in the organic phase; K2 is the equilibrium constant for formation of the unhydrated dimer from the unhydrated monomers; KI1 K12 . . . KZl K22 . . . are formation constants for the hydrates AW AW2 . . . A2W A2W2 etc. from monomers of A and W; and KD = CA / C A ~ is the distribution constant for monomeric A between water and the organic phase.The concentration of water monomer CW may be calculated from the expression cw = awcwo where -0 is the limiting value of cw at unit water activity. Henry's Law is assumed to apply to each hydrated and unhydrated species individually. The measured concentration of water in the organic phase is greater in the presence of A than it would be if only water and the solvent were present; the excess water solubility dfw is due to the formation of hydrated species and may be expressed as s4 a A. A. Taha R. D. G-rigsby J. R. Johnson S. D. Christian and H. E. Affsprung J . Chem. Educ. 1966,43,432; b E. E. Tucker Ph.D. Dissertation University of Oklahoma 1969. 26 Christian Taha and Gash Afw = CA ( K ~ ~ C W + 2K1acwa + . . . ) + C A ~ ( K ~ ~ c w + 2K22cw2 + . . . ) (3) Equations (1) and (3) apply equally to the partition systems or to solutions of A in solvent equilibrated at known reduced water activities.If sets of distribution data VA CA* CW) and water solubility data (fA CW dfw) are available Km and the hydration constants may be inferred by non-linear least-squares analysis.36 Initial values of KD K2 and the hydration constants are chosen and equation (1) is solved for each point to yield trial values of CA. The mean square deviation E given by E = CWS(AfWexpt1. - Af+alc.)a + CWpdfAexptl. - fAcalc.)2 (4) all points partition data points is then calculated using expressions (3) and (2) to obtain dfwcalc. and f~calc. respectively. (The weight factors WS and WP are determined from the precision of the water solubility and partition data respectively.) After each calculation of E the procedwe is repeated with a set of modified values of all the constants generated with a numerical optimum-seeking technique.The process is continued until the absolute minimum in E and the corresponding least-squares values of the constants are located. Standard errors in the constants may be calculated by the method of Sillen.36 Ordinarily no more than 2 or 3 hydration constants are needed to fit a given collection of data. Equations similar to (1-4) may be developed to fit i.r. and lH n.m.r. data for systems in which hydrates of polar solutes are p r e ~ e n t . ~ ~ . ~ ~ It is commonly assumed that both Henry’s Law and Beer’s Law apply individually to solute species and that the proton chemical shift of each species is concentration inde~endent.~~ Important qualitative information about hydrated species can be obtained from spectral data even when it is not possible to calculate hydrate formation constants.Bands in the 3 pm region have been attributed to 1 :1 and 2:l hydrates (B H-0-H and B * * - H-0-H * * - B) of organic bases (B) in CCl,.39 The relative strengths of water-polar solute interactions have been inferred from proton chemical shifts24 and OH-stretching f r e q ~ e n c i e s ~ ~ ~ of water dissolved in hydrogen-bonding solvents. B. Thermodynamic Results for Water-Polar Solute Complexes.-Before the past decade not much was known about specific molecular complexes of water. as R. Van Duyne S. A. Taylor S . D. Christian and H. E. Affsprung,J. Phys. Chem. 1967,71 3427. 38 L. 0. Sill& Acra Chem. Scand. 1964 18 1085.37 a F. Takahashi and N. C. Li J . Amer. Chem. SOC. 1966,88 11 17; b J. D. Worley Ph.D. Dissertation University of Oklahoma 1964; c A. Fratiello and D. C. Douglas J. MoZ. Spectroscopy 1963 11 465; d J. M. Sorensen Ph.D. Dissertation Case-Westem Reserve University 1967. 88 N. Muller and P. Simon J. Phys. Chem. 1967 71 568. as a S. C. Mohr W. D. Wilk and G. M. Barrow J. Amer. Chem. SOC. 1965,87 3048; b P. Saumagne and M. L. Josien Bull. Soc. chim. France 1958 813 40 E. D. Becker J. Chem. Phys. 1959 31 269. 27 Molecular Complexes of Water in Organic Solvents and in the Vapour Phase This section summarises recent research which has yielded thermodynamic parameters for specific complexes of water with carboxylic acids phenols and alcohols ketones and ethers amines amides and various poly-functional polar solutes.Since 1960 several specific complexes of water (W) with monocarboxylic acids (A) have been reported based on spectral,4l vapour vapour pressure,43 and partition-water solubility44 measurements. 1.r. spectra of moist CClp solutions of three highly acidic halogenated acetic acids and three weaker aliphatic acids have been obtained in the acid C=O stretching and the water hydroxyl-stretching regions.4f At very low water activities new bands attributed to A W and AaW appear in both spectral regions. Formation constants for A W and A2W as well as AW2 have been r e p ~ r t e d . ~ ~ ~ ~ ~ The 1 :2 and 2:l complexes of benzoic acid and salicyclic acid in several non-polar and slightly polar solvents have formation constants (from the monomers) of the order of several hundred in la mole- units.A vapour phase constant at 25" of 0.010 torrs (equivalent to 3.5 x lo6 1 mole-2) has been obtained for the formation of the dihydrate of trifluoroacetic Values of the formation constant for A W for aromatic acids in CC14 benzene and 1,2-dichloroethane are in the range 1-12 1 mole-2. The large magnitudes of the A2W and AW formation constants as well as the acid self-association constants indicate the difficulty of determining accurate parameters for the A W formation constant; i.e. it is difficult to reach a con- centration region in which a major portion of the hydrated acid is in the form of the 1:l complex. Similar but less severe complications arise in studies of hydrates of other polar solutes having both donor and acceptor sites. Relatively little is known about specific hydrates of phenols and alcohols in spite of the fact that numerous studies have been made of aqueous and moist organic solutions of these mmpounds.ld~* Hydration studies are complicated by a lack of reliable values of self-association constants.Early p a r t i t i ~ n ~ ~ ~ ~ ~ and cryoscopic2sa investigations showed that strong interactions occur between dissolved water and phenols or alcohols in organic solvents. Partition-water solubility data at unit water activityYc6 and water solubility data at reduced water activities,46b-d have indicated the presence of PW Paw and PW in solutions of phenol (P) and water (W) in several organic solvents. In benzene 1,1,2,2- tetrachloroethane and 1 ,2-dichloroethaneY the trirneric hydrates are more 41 J.de Villepin A. Lautie and M. L. Josien Ann. Chim. 1966,1 365. 42 S. D. Christian H. E. Affsprung and C. Ling J. Chem. SOC. 1965 2378. 4s G. 0. Wood D. D. Mueller S. D. Christian and H. E. Affsprung J . Phys. Chem. 1966 70 2691. *4 a R. Van Duyne S. A. Taylor S. D. Christian and H. E. Affsprung J . Phys. Chem. 1967 71 3427; b R. Van Duyne Ph.D. Dissertation University of Oklahoma 1969. 45 a R. D. Vold and E. R. Washburn J. Amer. Chem. SOC. 1932,54,4217; b E. R. Washburn V. Hnizda and R. D. Vold ibid. 1931,53,3237; c E. R. Washburn and H. C. Spencer ibid. 1934 56 361. 46 R. M. Badger and R. C. Greenough J. Phys. Chem. 1961 65 2088; b J. R. Johnson Ph.D. Dissertation University of Oklahoma 1966; c J. R. Johnson S. D. Christian and H. E. Affsprung J. Chem. SOC. (A) 1967 764; d J. R. Johnson S.D. Christian and H. E. Affsprung ibid. 1965 1. 28 Christian Taha and Gash important than PW and account for most of the bound water. In CC14 solutions of phenol Afw rises rapidly with increase in phenol concentration and aw indicating the presence of hydrates involving at least 2 water molecules and several phenol molec~les.~~d Partition and water solubility data for alcohols and 1-naphthol in toluene indicate that 2:l alcohol-water and 1 :1 naphthol-water species are probably present,,' although the stoicheiometry of water in these complexes has not been confirmed by solubility measurements at reduced water activities. Similar data for 1-decanol (D) provide evidence for the existence of DW in 1 ,2-dichloroethane and D2W in i~o-octane.~~ Conductance measurements for solutions of several nitro-substituted phenols and water (fw > 0-3 mole 1-l) dissolved in acetonitrile indicate that the phenols form hydrates PWn in which n is one greater than the number of nitro-groups in the substituted phenol.*@ Ethers and ketones are not highly associated in organic solvents at concentra- tions less than several tenths molar.Therefore the analysis of hydration data for these compounds does not involve large corrections for solute self-associa- tion. The acetone monohydrate in 1 ,Zdichloroethane has been investigated by dielectricto lH n.m.r.,22 and partition-water solubilityZ2 methods; a formation constant of 0.85 1 mole-' at 25" and a dipole moment of 3.4 D have been reported for the complex. The dipole moment of the hydrate nearly equals the vector sum of the moments of water and acetone.Table 2 lists formation constants Table 2 Equilibrium constants for formation of hydrates of ketones ethers and dimethylsulphoxide at 25 O Ketone or Ether Solvent Cyclopen t anone Acetone 2,3-Butanedione Acety lacet one Acetone Tetrahydrofuran Dioxan Dimethylsulphoxide CCl CCI CC14 CC14 Acet one-cyclohexane THF-cyclohexane Cyclohexane Dimethylsulphoxide-CCI KBW KB,w* Ref. 2.2 1.4 51b 2.4 1.6 51a 1.2 2.0 51a 1.5 0.86 51a 0-08t 37a 0-13t 37a (I mole-l) (1 mole-') 1*3t 0.85t 38 0-26t 52 *For the reaction BW + B = B2W. ?Converted from mole fraction units and interpolated to 25" for the ketone (B)-water complexes BW and B,W in CC14 and cyclohexane at 25°.37aJj1 Hydration equilibria of ethers have been investigated by lH n.m.r. measurements;37as38 constants for dioxan and tetrahydrofuran complexes are 47 D.J. Turner A. Beck and R. M. Diamond J. Phys. Chem. 1968,72,2831. 48 D. J. Turner and R. M. Diamond J . Phys. Chern. 1968,72,3504. '@ A. D'Aprano and R. M. Fuoss Proc. Nat. Acad. Sci. US. 1968,61 1183. 6o T. F. Lin S. D. Christian and H. E. Afhprung J. Phys. Chem. 1967,71 1133. 61 a T. F. Lin S. D. Christian and H. E. Affspmg J. Phys. Chem. 1967,71 1968; b R. L. Lynch S. D. Christian and H. E. Affsprung ibid. 1969,73 3273. 29 Molecular Complexes of' Water in Organic Solvents and in the Vapour Phase included in Table 2 as well as the B2W formation constant for dimethyl sul- phoxide.6a AH for the addition of the first water to acetone is ca. -3.5 k a l / whereas values in the range - 1.6 to - 2-8 kcal/mole have been obtained for addition of a second base unit to the monohydrates of ketones and ethersg7V* and dimethyl s~lphoxide.~~ Presumably B2W has the bridged structure B 9 HOH - - - B proposed earlier on the basis of i.r.It is obvious that weaker hydrogen bonding is involved in ketone- or ether-water complexes than in hydrates of phenols and carboxylic acids. The 1 :1 complex formation constants reported in Table 2 are comparable in magnitude to those for com- plexes of similar bases with Considerable effort has been expended in studying the hydration of amines aniides and biologically important compounds containing the amide group and peptide linkages. The fact that amines are known to form strong hydrogen bonds with proton donors of many types and the tendency of amines to obey Henry's Law throughout considerable ranges of concentration have made these compounds logical choices as solutes in hydration studies.There is spectral evidence for the aromatic amine (D) hydrates DW and D2W in organic solvent~,39,~~ although some aliphatic amines apparently form DW but not D2W owing to the strong inductive effect of the polar D W hydrogen bond.65 Table 3 summarizes formation constants for DW DW2 and D2W (from the monomers) as well as selected values of these parameters for methanol-amine complexes. Dipole enhancements for 1 1 complexes of hydroxylic proton donors and amines have been reported; the excess of the complex dipole moment over the vector sum of the donor and acceptor moments ranges from less than 0.3 D for aliphatic alcohol complexes with pyridine and triethylaminess to 0.4-1.3 D for aliphatic amine complexes with ~ a t e r l ' e ~ ~ and phenols.68 1.r.spectral evidence for the separation of charge in various n-propylamine-proton donor com- plexes50 complements the dielectric results; the amine N - H and the OH a e N stretching bands are significantly altered in complexes with proton donors having pKa values < ca. 10. Nevertheless the charge separations indicated by the dipole moment enhancement for water-amine complexes are unexpectedly large compared to those reported for the aliphatic alcohol-amine complexes. Table 4 gives formation constants for the amide (M) hydrates MW and M2W. Difficulties again arise in treating hydration data for lactamsso and acetamides61 in that accurate self-association constants are required in calculating hydrate S.F. Ting S. Wang and N. C. Li Canad. J. Chem. 1967,45,425. 6a E. S. Hanrahan and F. Deskins Proc. West Virginia Acad. Sci. 1967,39 371. 64 A. N. Sidorov Optics and Spectroscopy 1960 8 24. 6s H. E. Affsprung J. Derkosch and F. Kohler Discuss. Furaday Soc. 1965 40,224. sB J. W. Smith J. Chim. phys. 1964 61 125. M. D. Gregory Ph.D. Dissertation University of Oklahoma 1968. J. R. Hulett J. A. Pegg and L. E. Sutton J. Chem. SOC. 1955 3901. T. Zeegers-Huyskens Spectrochim. Acta 1965,21 221. 8o a J. D. Worley Ph.D. Dissertation University of Oklahoma 1964; b D. Mueller Ph.D. Dissertation University of Oklahoma 1966. a R. D. Grigsby Ph.D. Dissertation University of Oklahoma 1966; b R. D. Grigsby S. D. Christian and H. E. Affsprung J . Phys. Chem. 1968,72,2465. 30 Christian Taha and Gash Table 3 Equilibrium constants for formation of amine-water a d amine-methanol complexes at 25" Amine Diethy lamine Diethy lamine Diethylamine Pyridine Pyridine Pyridine Pyridine Pyridine C yclohexy lamine N-Methylcyclo- hexy lamine NN-Dimethylcyclo- hexy lamine Triethy lamine Triethylamine Triethy lamine Diethy larnine Diethylamine Diethy lamine Pyridine 2-Picoline 2-Ethylpyridine 2-Isoprop ylpyridine 2- t-Butylpyridine 2.6-Lu tidine Water Complexes Solvent KDW Hexadecane 11.0 Diphenylmethane 8.5 Benzyl Ether 2.9 Cyclohexane 5.3 CC14 2.6 Toluene 1.5 Benzene 1.2 1-2 Dichloroethane 1.0 Benzene 6.7 (1 mole-l) KDW~* KD,w* Ref.(12 mole-2) (P mole-2) a a 9-7 a 6-8 b 18.3 3-0 b 7.3 1.6 b 6.3 1-5 b 1 -0 1.0 b C Benzene 5.3 C Benzene 3.9 Benzene 3.5 Cyclohexane 7.0 Toluene 3.7 Methanol Complexest Hexadecane 8.7 Diphenylmethane 4.8 Benzyl Ether 2.7 CC14 2.3 CCL 3.1 CC14 2.6 CC14 1-7 CC14 1.4 CCl 4-4 300 22-5 8-5 C C C c sEthyl-6-methylpyridine CCl; 4.1 d 2,6-Diethylpyridine CC14 3.0 d 2,6-Di-isopropylpyridine CCl 0.8 d a ref.34b; b ref. 74b; C ref. 17; T. Kitao and C. H. Jarboe J. Org. Chem. 1967,32,407. *For the reactions from the monomers fIn the case of the methanol systems KDW and KDW indicate formation constants for the 1 1 and 1 2 amine -methanol complexes respectively. Table 4 Equilibrium constants for formation of amide hydrates at 25" Amide Solvent KMW KM~w* Ref. (1 mole-') (1 mole-') 2-Pyrrolidone CC14 9.0 9 a N-Methyl-2-pyrrolidone CCI 9.2 39 b N-Methyl-2-pyrrolidone Benzene 6.0 4.3 b N-Methyl-2-p yrrolidone 1,2-Dichloroethane 2.4 b NN-Dimethylacetamide 1,2-Dichloroethane 2.2 b N-Methylacetamide cc14 11 C NN-Di me t h ylacet amide NN-Dime t hylacetamide- cyclohexane 0.3 d a ref.37b; b ref. 606; ref. 616; ref. 37a *For the reaction MW + M = M2W 31 2 Molecular Complexes of Water in Organic Solvents and in the Vapour Phase formation constants. N-Methylacetamide (M) in CC& for example associates extensively in dilute solution to form chain polymers. The formation constant for the reactions M + W = MW Mz + W = MzW . . . Mn + W = MnW appears to be nearly constant (ca. 11 1 mole-’) and independent of chain length which suggests that each unit in the chain contributes an equal number of nearly equivalent hydration sites.B1b 2-Pyrrolidone is extensively dimerized in CCl, and the equilibrium constant for the reaction M + W = MW nearly equals that for MP + W = MPW.The species MIW probably consists of water attached to the cyclic 2-pyrrolidone dimer ; whereas N-methyl-2-pyrrolidone which associates to a lesser degree probably forms the bridged hydrate Several reviews have treated the extensive literature on the energetics dynamics and structure of water molecules bound to proteins and other macro- molecules;1bJf~s2 no attempt will be made to summarise research in this area. In the context of the present Review a series of studies of the interactions of water with various polar groups of deoxyribonucleic acid (DNA)s3 and its molecular constituentsB4 is of interest. 1.r. spectra of thin films of Na and Li salts of DNA have been measured at various water activities (relative humidities).Bab As aw increases from 0 to 0.60 first the POz- group and then the POC and COC oxygens become hydrated.At a water activity of ca. 0.65 a hydration shell of 5 or 6 water molecules probably surrounds the phosphate group; above aw = 0.65 the C-0 and ring nitrogen atoms become hydrated. Polarised i.r. and U.V. spectra indicate that between aw = 0.55 and 0-75 the structure of DNA changes from the disordered lower humidity form to the ordered B configuration in which the base pairs become stacked one above another and oriented perpendicular to the axis of the helix.s3c Gravimetric results for numerous constituents of DNA show that most of the solid purines and pyrimidines and their nucleosides and nucleotides are little hydrated at water activities <0.93 whereas several salts of the nucleotides and other com- pounds containing the ionic phosphate group absorb >20 molecules of water per molecule of compound at aw = 0.93.These studies underline the importance of the ionic phosphate groups as sites for hydration in nucleotides and poly- nucle~tides.~~ Tributylphosphate (TBP) and other organophosphorus compounds have been widely used as extracting agents for metal ions.6s There have been several reports of the hydration of these compoundS,sa although in most studies the M . . W . . M.606 6a J. Steinhardt and S. Beychok in ‘The Proteins’ ed. H. Neurath Academic Press New York and London 1964 vol. 2 p. 276. 63 a M. Falk K. Hartman and R. C. Lord J. Amer. Chem. SOC. 1962,84,2843; b M. FaIk K. Hartman and R. C. Lord ibid. 1963,85 387; c M. Falk K. Hartman and R. C. Lord ibid. 1963 85 391.64 M. Falk Canad. J. Chem. 1965 43 314. 66 H. Freiser Analyt. Chem. (Ann. Rev.) 1968 40 522R. a L. Kuca Coll. Czech. Chem. Comm. 1967,32,720; b D. C. Whitney and R. M. Diamond J. Phys. Chem. 1963 67,209; c T. J. Conocchioli M. I. Tocher and R. M. Diamond ibid. 1965 69 1106; d J. J. Bucher and R. M. Diamond ibid. 1969,73,675; e J. J. Bucher and R. M. Diamond ibid. 1969 73 1494. 32 Christian Taha and Gash water activity was not varied systematically so that the stoicheiometry of water in the hydrate formation reactions could be inferred. A IH n.m.r. investigations7 yielded equilibrium constants equal to 6.1 and 0-73 1 mole-1 (converted from mole fractions to molarity units; interpolated to 25") for the reactions TBP + W = TBP-W and TBP-W + TBP = (TBP),.W respectively.Enthalpy values of -4.1 and -2.0 kcal/mole respectively were given for the two hydration reactions. Water solubility-partition data for TBP yield 1 1 hydrate formation constants of 34 in iso-octane,66d 18 in C14,86b 9-1 in benzene,66e and 2.6 1 mole-1 in CHC13.66e IH N.m.r.68 and partition-water solubilityss studies indicate that HN03 H20 and (HN03)2 H20 exist in benzene and toluene. Recent reviews provide numerous references to investigations of the hydration of ions in organic medias5 and the extraction of metal salts and chelate~.~~ There is ample evidence for the formation of strong complexes between water and extracting agents (including ions ion pairs and polar molecules) but accurate thermodynamic data for well-characterised hydrate species are generally lacking. 4 Solvent Effects Variation in polarity and reactivity of solvents strongly affects values of the thermodynamic functions characteristic of hydrogen bond formation reactions.Even the so-called 'inert' solvents influence molecular complex formation equilibria in general reducing the tendency for aggregates to form as compared to the vapour More polarisable and more polar solvents ordinarily lead to a further decrease in the percentage of associated species present at fixed total concentration of donor and acceptor. Henry's Law constants,'l X-H stretching frequencies,7oa solubility parameter~,~~C and IH n.m.r. chemical shifts of dissolved solutesa4 have been related to the reactivity of solvents; in a general way changes in these parameters parallel variations in thermodynamic constants for association reactions of the solutes in the various media.Quantitative treatments of the effect of solvent on complex formation reactions have falIen into two main classes:70C972 (i) methods in which corrections are made for presumed competitive equilibria involving the formation of complexes between solvent molecules and one or more of the dissolved solute species; and (ii) techniques in which solvation of the donor acceptor and complex molecules is considered to occur in a non-specific way. Neither approach by itself has 67 S. Nishimura C. Ke and N. C. Li J. Amer. Chem. SOC. 1968 90 234. g9 a E. Hogfeldt and B. Bolander Arkiv Kemi 1963,21 161 ; b C. J. Hardy B. F. Greenfield and D. Scargill J. Chem. SOC. 1961 90. 70 u A. R. H. Cole and A. J. Michell Austral. J. Chem. 1965,18,102; b T.Gramstad Specfro- chim. Acru 1963 19 1363; c H. Buchowski J. Devaure P. V. Huong and J. Lascombe Bull. SOC. chim. France 1966 2532. 71 M. L. Josien and N. Fuson J. Chem. Phys. 1954,22,1169. 78 a R. S. Drago T. F. Bolles and R. J. Niedzielski J. Amer. Chem. SOC. 1966 88 2717; b P. J. Trotter and M. W. Hanna ibid. 1966,88 3724; c S. Carter J. N. Murrell and E. S. Rosch J. Chem. SOC. 1965 2048; d G. Kortiim and W. M. Vogel Z. Elektrochem. 1955 59,16; e S. D. Christian J. R. Johnson H. E. Affsprung and P. J. Kilpatrick J. Phys. Chem. 1966,70,3376; f J. Barrio1 and A. Weisbecker Compr. rend. 1967,265 C 1372. J. C. Eriksson L. Odberg and E. Hogfeldt Acra Chem. Scand. 1967,21,1925. 33 Molecular Complexes of Water in Organic Solvents and in the Vapour Phase proven to be entirely satisfactory.Methods of type (i) require the assumption that equilibrium constants for specific molecular complexes involving the solvent and other species are truly constant over wide ranges of concentration and in different media. On the other hand type (ii) methods tend to underestimate the effect that formation of specific molecular complexes between solute and solvent has on the energetics of the dissolved donor acceptor or complex. We shall limit our discussion of solvent effects to some recent applications of non-specific solvation theories. In treating data for reactions such as A(so1vated acceptor) + D(so1vated donor) = AD(so1vated complex) it is convenient to define the dimensionless parameter a = LIEADO / (~EAO + LIED*) in which ~EADO ~EAO and represent internal energies of transfer of the individual species from vapour to the sohent s at infinite d i l ~ t i o n .~ * e ~ ~ ~ The parameter a may be interpreted as the fraction of the energy of solvation of the free donor plus acceptor that is retained by the complex AD; in the absence of strong dipole enhancement in the complex it is expected that a < 1 since the process of bringing the reactive groups of donor and acceptor together to form the complex should 'squeeze out' solvent molecules. Apparently a is not strongly dependent upon choice of medium; a similar parameter a' = ~GADO 1 (AGAO + AGDO) (where the AGO terms represent solvation Gibbs free energies for IM ideal dilute solution standard states) is also ntarly constant from solvent to solvent. Procedures have been developed for calculating a from a lattice theory utilising group interaction energies,74 and for predicting E' from v 4 v-+s v-+s v+s v-bs V+S v-4 v4s v+s Several useful relations involving a and a' are - ~ E I O = (a - 1) (AEA' + LIEDO) I-tII I+II AGIIO - AGIO = (a' - 1) ( ~ G A ' +AGDO) I411 I+II KADII I KADI = (KI),AKD,D)"-~ 8 in which ~EIIO and ~EIO are standard internal energy changes for the reaction A + D = AD in solvent I1 and solvent I respectively; ~ G I I O anddGIO are the corresponding standard free energy changes; KADII and KADI are equilibrium constants for the formation reaction in the two solvents in 1 mole-* units; and &,A and KD,D are distribution constants between solvent I and solvent I1 for the acceptor and donor respectively.(KD,A for example may be equated to the ratio of the Henry's Law constant of the acceptor in solvent I to the Henry's Law constant of acceptor in solvent 11; a large value of KD for a given solute indicates that solvent I1 is much more effective than solvent I in solvating the species.) 79 a S.D. Christian and J. Grundnes Acta Chem. Scand. 1968 22 1702; b S. D. Christian and H. E. Affsprung ?3olute Properties of Water' U.S. Government Printing Office Wash- ington D.C. 1968 p. 57. 74 a T. L. Stevens Ph.D. Dissertation University of Oklahoma 1968; b J. R. Johnson P. J. Kilpatrick S. D. Christian and H. E. Affsprung J. Phys. Chem. 1968,72 3223. 34 Christian Taha and Gash If a’ is constant equation (7) requires that for a given complex formation reaction in a series of solvents a plot of log KAD vs. log (KD,AKD,D) will be linear with slope a’ - 1.Such a plot is shown in Figure 2 for the reaction water (A) + pyridine (D) = pyridine monohydrate (AD).72e The points lie approxi- mately on a straight line having a slope of - 0-29 corresponding to a’ - 0-71. For the solvents employed lattice calculations lead to predicted values of a and a‘ of ca. 0-75. Strong complexes for which the dipole moment is considerably greater than the vector sum of the moments of the unreacted donor and acceptor exhibit 35 Molecular Complexes of Water in Organic Solvents and iit the Vapour Phase anomalous solvation effects.75 To account for the large values of a and a’ obtained for amine-water complexes it has been proposed that the solvation energy of the complex includes an extra contribution due to excess dipole- induced dipole interactions not properly accounted for in lattice energy calcula- tion~.~’ Table 5 lists values of a and a’ for the 1 :1 hydrate of diethylamine (DEA) Table 5 Eflect of solvent on donor acceptor and complex for the reaction DEA + W = DEA-W or(experimenta1) a’(experimenta1) ~EODEA*H,O (kcal/mole) ~ E O D E A (kcal/mole) AEOH,O (kcal/mole) v+s v4s v+s d@DEA’H,O (kcal/mole) v4s v+s v-bs d G’DEA (kcal/mole) dGO13,o (kcal/mole) n-Hexadecane 1.01 1-05 - 7.80 - 5.60 -2.13 - 3.80 - 3.16 - 0.45 Dipheny lmethnne 0-85 1 -01 - 8.92 - 6.29 -4.17 - 5.20 - 3.36 - 1.79 Benzyl ether 0.80 0.90 - 10.18 - 6.43 - 6.33 - 5.26 - 3.33 - 2.54 in several solvents as well as solvation energies and free energies for the in- dividual species involved in the Predicted values of a and a’ for n-hexadecane and diphenylmethane obtained from lattice calculations again lie in the range 0-7-0.8.The difference between the energy of solvation of the complex in each solvent and that predicted from the solvation energies of the unreacted components using a = 0.75 is ca. - 1 to - 2 kcal/mole. A major part of this extra solvation energy is probably due to the interaction of the excess dipole of the complex with the polarisable solvent. It is clear that non- polar and slightly polar solvents are less effective in preventing the formation of highly polar complexes than they are in opposing complex formation between weakly-interacting donor and acceptor molecules. The properties of aqueous solutions and of dilute solutions of water in non- polar media are apparently little related. However an understanding of the progressive changes which occur in the solvation of water and other polar molecules in various media including dipole-solvent intera~tions~~f and inductive effects,”a should form the basis for building a bridge between these two interesting areas of chemistry. l6 J. Grundnes and S. D. Christian J. Amer. Chem. SOC. 1968,90,2239. 36