Thermodynamic Properties for Transfer of Electrolytes fro m Water to Acetonitrile and to Acetonitrile+ Water Mixtures BY BRIAN G. Cox* AND RAJAGOPALAN NATARAJAN Department of Chemistry, University of Stirling, Stirling FK9 4LA, Scotland W. EARLE WAGHORNE* AND Department of Chemistry, University College, Belfield, Dublin 4, Ireland Received 22nd March, 1978 The free energies, enthalpies and entropies of transfer of a variety of electrolytes, including alkali metal halides, silver halides, tetraphenylarsonium and tetraphenylboride salts, from water to acetonitrile and water + acetonitrile mixtures have been measured. The enthalpies and entropies of transfer show a complex dependence upon solvent composition, which is discussed in terms of the effects of solvent sorting in the immediate neighbourhood of the ions, and of the effect of the solvated ions on the bulk solvent properties.The large structural effects occurring when ions are dissolved in water do not, however, appear to make a net significant contribution to the free energies of solution (and transfer) of electrolytes. Thus the overall free energies of transfer and their dependence upon solvent composition may be very simply interpreted in terms of the assumption that alkali metal and halide ions interact more strongly with water, and silver ions with acetonitrile, without reference tolany special structural properties of water. The entropies of solution of simple electrolytes in water are considerably more positive than those in a variety of dipolar aprotic s01vents.~'~ The differences, which are of the order of 40 cal K-l mol-1 (corresponding to a TAS term of some 12 kcal mol-1 at 25"C), may be ascribed primarily to effects resulting from the hydrogen bonded structure of water.Entropy losses resulting from the " freezing " of solvent molecules in the vicinity of the ions should be lower in water, and it has been suggested that the disruption of the water structure by the solvated ions may cause a significant increase in entropy in a region further from the ion?* It is not clear, however, to what extent these " structural " effects contribute to the free energies of solution of the ions, because of the tendency of the resulting enthalpy and entropy changes to compensate for one another. Indeed, much available evidence suggests that changes in structure of aqueous solutions frequently occur with little change in free energy.'-l We have recently studied the changes in thermodynamic properties accompanying the transfer of electrolytes (Ag+, Na+ and Li+ salts) between two dipolar aprotic solvents, dimethylsulphoxide and propylene carbonate, and their mixtures.The observed free energies (AG;), enthalpies (AHtr) and entropies (Astor) of transfer of the cations can be very simply interpreted in terms of the interactions between the ions and their immediate neighbour solvent molecules. In particular, a sharp minimum in ASFr occurs in the mixtures, resulting from large differences in the solvent composition of the coordination sphere of the ions, and of the bulk solvent. In the present paper, we report a comprehensive study of AG,", AH,", and ASFr for transfer of a variety of electrolytes from water to acetonitrile and their mixtures.86B. G. COX, R. NATARAJAN AND W. E. WAGHORNE 87 The results are discussed with reference to those obtained in mixtures of dipolar aprotic solvents. Acetonitrile +water mixtures are particularly interesting in relation to studies of ion-solvent interactions, because of the contrasting behaviour of alkali metal cations, which interact more strongly with water, and silver (and cuprous) cations which are more strongly solvated by acetonitrile.' EXPERIMENTAL CHEMICALS Acetonitrile was purified by successive distillation from phosphorus pentoxide and calcium hydride, after initial drying with calcium hydride.13 Alkali metal salts were Specpure grade (Johnson Mathey).They were dried at 100°C under vacuum prior to use. AnalaR silver nitrate was used without further purification. Bu4NBr (Cambrian Chemicals) was dried under vacuum. Ph4AsI was prepared from Ph4AsC1 and KI. It was purified by recrystallization from water. KBPh4 was prepared from KCI and NaBPh4, and purified by recrystallization from acetone+ H20. NaBPh4 (B.D.H.) was purified by recrystallization from acetone+ toluene as previously described.14 Ph4AsC1 (Aldrich) was purified by recrystallization from dichloromethane, and dried at 100°C under vacuum. This was necessary to prevent contamination of the sample with Ph4AsCl 2H20. FREE ENERGIES OF TRANSFER SOLUBILITY MEASUREMENTS The standard free energies of transfer AGPr(MX),from water to solvent S of AgBr, AgI, P u s 1 and KBPh4 were calculated from eqn (l), where Ksp is the solubility product referred to infinite dilution in the appropriate solvent (S or H20) Ksp values for AgCl are available in the 1iterat~re.l~ Ksp for silver halides were determined from potentiometric titrations of AgN03 into standard tetra-alkylammonium halide solutions, AGZ(:(MX) = RT In [&pCH20)IKsp(~)I.(1) TABLE 1 .-SOLUBILITIES OF ELECTROLYTES IN ACETONITRILE+ WATER MIXTURES AT 25°C PKS a JAN^ AgBr &I Ph&I KBPh4 0 12.21 16.03 5.26 7.34= 0.10 11.i6 14.8, 4.4, 0.20 10.84 14.2, 3.60 6.12 0.40 10.27 13.69 2.22 4.61 0.60 10.2~ 13.41 1.67 3.10 0.80 10.78 13.5, 1.53 2.74 0.90 11.51 14.00 1.00 13.90f 14.49 g 2.69' 3-11 a Solubility products corrected to zero ionic strength ; estimated uncertainty kO.1 log unit ; b volume fraction of acetonitrile; Cref.(4); d c f : pKs = 5.1, ref. (12); e c f . pK, = 7.5, ref. (12); fref. (23) ; g ref. (12) ; h cf. pK, = 2.72 ref. (23) ; 1 cf. pKs = 3.24 ref. (23). The difference arises from the activity coefficients used to correct to zero iomc strength. the titrations being monitored with an Ag/Ag+ electrode. KBPh4 and PLAsI solubilities were determined by analysing saturated solutions of the salts for K+ by atomic absorption spectroscopy, and by titration against standard AgN03 solutions to determine I-, respectively. The solubility products were corrected to infinite dilution using activity coefficients (yk) calculated from the Davies eqn (2),16 where A is the Debye-Hiickel parameter88 ELECTROLYTES I N WATER AND ACETONITRILE which depends upon the solvent dielectric constant and temperature, and I is the ionic strength.Values of A may be calculated from the known dielectric constants of acetonitrile+ water mixture~.'~ The results are listed in table 1. ELECTROCHEMICAL MEASUREMENTS Free energies of transfer of alkali metal halides, MX, were determined from e.m.f. measurements on cell A.18s21 In cell A, the glass electrode was either a sodium ion-selective electrode (Activation) or a general cationic Ag-AgX 1 M x ~ ~ ~ ~ ~ = m, I glass electrode electrode (Beckman 39137). Earlier studies have shown that such electrodes give results in good agreement with those of conventional amalgam electrodes,' 8-20 the results for Cs+, however, being generally less reliable than for the other cations.Ag-AgX electrodes were prepared. by electrolysing silver electrodes in MX solutions.21 The free energy of transfer of MX from water to solvent S is given by eqn (3),22 where I& refers to the e.m.f. of cell A, (3) with electrolyte MX of concentration ms, and similarly for EH20. In eqn (3, y; represen s the mean ionic activity coefficient of electrolyte MX referred to infinite dilution in solvent S. The activity coefficients were calculated from eqn (2). Electrolyte concentrations in cell A varied between and mol dm-3. AGS(MX) = F(Es-E~20)+2RTln (mHZ0/mS)+2RTln (ypo/y$) TABLE 2.-FREE ENERGIES OF TRANSFER (AGt",)' OF ELECTROLYTES FROM WATER TO MIXTURES OF WATER AND ACETONITRILE AT 25°C A AN 0 0.10 0.20 0.30 0.40 0.50 0.60 0.70 0.80 0.90 0.95 1 .oo LiCl 0.00 0.30 0.69 3.12 4.28 7.4 10.4 17.2f AGtJkcal mol-1 b (molar scale) NaCl NaBr KCl KBr KI CsCl AgBreld 0.00 0.00 0.00 0.00 0.00 0.00 0.00 0.27 0.l9 0.l9 0.09 0.03 0.l9 -1.40 0.51 0.39 0.45 0.29 0.05 0.37 -1.92 0.82 0.68 0.73 0.52 0.19 0.61 -2.40 1.24 1.00 1.13 0.82 0.30 0.89 -2.70 1.68 1.38 1.59 1.19 0.52 1.18 -2.81 2.84 2.39 2.68 2-20 1.25 2.21 -2.40 3.94 3-22 3.59 3.06 1.94 3.48 -2.00 7.02 5.49 6.28 5.57 3.19 6.08 -1.00 9.74 7.79 8.68 7.28 4.20 8.41 0.30 2.23 1.81 2.09 l.60 0.81 l.60 -2.65 13.4g 17.0h 12.0i 9.2g 6.4h 11.6g 2.3g KBPh4 c*e 0.00 -0.8 - 1.8 - 2.9 - 4.0 - 5.2 - 6.0 - 6.4 - 6.5 - 6.4 - 6.2 - 6.0 g Ph&I '1' 0.00 - 0.9 - 2.3 - 3.4 - 4.0 - 4.4 - 4.7 - 4.9 - 4.9 - 4.4 - 3.9 -3.3 i a Estimated error i- 0.05 kcal mol-l unless otherwise stated ; 1 cal = 4.184 J ; C from table 1 ; d estimated error & 0.1 kcal mol-' ; e estimated error rt 0.2 kcal mol-1 ; f M.Solomon, J. Phys. Chem., 1970,74, 2519 ; g ref. (23) ; h J. Pavlopoulous and H. Strehlow, 2. phys. Chem., 1954, 202, 474 ; i ref. (12). Cell A was not used for solvent mixtures containing > -85 vol % acetonitrile, be- cause of the tendency of AgX to dissolve (giving AgX,) when the anion activity becomes too high. However, AG,", values for transfer to pure acetonitrile have been previously determined from solubility meas~rements,~~ and AGP, values for the region between 85 vol % and pure acetonitrile were obtained by interpolation. AGP, values for Na+ salts have been shown to vary smoothly and monotonically over this region, using AG,", ma+- Ag+) values obtained from measurements on cell B, together with values of AG: (AgX) obtained from solubility measurement~.~~~ 24 The combined results, listed at selected volume fractions of acetonitrile are given in table 2.B .G . COX, R . NATARAJAN A N D W . E. WAGHORNE 89 In determining AGg values from solubility measurements [eqn (l)] and electrochemical measurements [eqn (3)], it is assumed that all of the salts are fully dissociated at the con- centrations used. The only measurements carried out in pure acetonitrile, where ion-pair formation should be strongest, involved solutions of tetra-alkylammonium, tetraphenylboride or tetraphenylarsonium salts, which are known to be strong electrolytes in acetonitrile.l2? 23 All other measurements were carried out in solvent mixtures containing at least 20vol % water, and at electrolyte concentrations < mol drr3. Under these conditions complete dissociation of the salts may be safely assumed.l2. 23 TABLE 3.-ENTHALPIES OF TRANSFER OF ELECTROLYTES FROM WATER TO MIXTURES OF WATER AND ACETONITRILE AT 25°C (kcal mol-l)b $AN 0 LiCl d 0 0.0 0.10 0.20 -0.5 0.30 0.40 -0.9 0.50 0.60 -1.8 0.70 0.80 -2.5 0.90 -4.6 0.95 -7.0 1.00 +4.9 0 0.0 0.10 -0.4 0.20 -1.0 0.30 -1.6 0.40 -1.9 0.50 -2.4 0.60 -3.0 0.80 -4.6 0.90 -6.6 1.00 -1.3e & A X C CScI 0.70 0.95 NaCl 0.0 - 0.5 -0.8 - 1.0 - 1.5 - 2.2 - 3.0 -4.5 + 1.6f AgCl 0.0 - 5.9 - 7.4 - 8.5 -9.1 - 11.0 - 11.4 - 7.9 g NaBr 0.0 - 0.3 -0.6 - 1.0 - 1.4 - 1.9 - 2.4 - 2.9 - 3.6 - 5.2 - 5.6 - 0.9 f 0.0 - 6.0 - 8.0 -9.1 - 10.1 - 10.6 -11.2 -11.9 - 10.9 g AgBr d KCI 0.0 - 0.9 - 1.4 - 1.6 -2.1 - 2.7 - 3.2 - 4.0 - 4.9 - 0.7 f AgI 0.0 - 7.0 - 8.6 - 9.3 - 9.5 - 10.7 - 11.7 - 12.5 - 13.7 - 14.6 KBr 0.0 - 1.3 - 1.9 - 2.2 -2.8 - 3.3 -4.1 - 4.6 - 5.3 - 3.4f 0.0 3.6 6.5 6.4 6.0 5.9 BudNBr 6.4 g KI RbCl 0.0 0.0 - 1.4 - 1.1 -2.8 - 1.8 - 4.2 - 4.7 - 5.7 -7.1 -7.1 f Ph4AsCl 0.0 4.7 5.7 4.9 2.5 - 2.8 - 4.0 - 5.8 -4.8 -1.5f NaBPhdd 0.0 4.2 2.1 0.5 - 2.0 1.3 - 5.0 0.8 - 6.0 1.9 - 5.6 g a Estimated uncertainty +_ 0.1 kcal mol-' unless otherwise stated ; enthalpies of solution in water :25 LiCI, - 8.85 ; NaCl -0.14 ; NaBr -0.14 ; KC1 4.16 ; KBr 4.75 ; KI 4.86 ; RbCl4.13 ; CsC14.25 ; AgCI4 15.7 ; AgBr4 20.2 ; Ag14 26.9 ; Bu4NBr4 -2.2 ; Ph4AsC14 - 2.6 ; NaBPh, -4.8 ; b 1 cal = 4.184 J ; C volume fraction of acetonitrile ; destimated uncertainty k 0.5 kcal mol-I ; e Y .4 . Choi and C. M. Criss, J. Chem. Eng. Data, 1977, 22, 297 and ref. (4) ; fref. (3) and (4) ; g ref. (4). ENTHALPIES OF TRANSFER The enthalpies of solution, AHs, (or precipitation in the case of silver halides) of electro- lytes in acetonitrile+ water mixtures were determined by standard calorimetric techniques using a Tronac 450-4 calorimeter. Electrolyte concentrations were in the range w mol dm-3. The values have not been extrapolated to infinite dilution, as the variations with concentration observed within this range were negligible within experimental error. Thus values are taken to be standard state values. Except for the silver halides, HCl and NaBPh4, AH: values were in the range - 5 kcal mol-1 < AH: < 5 kcaI mol-1 and were reproducible to +O.l kcal mol-l.In water, AH? values agreed to within kO.1 kcal mol-1 of literature values.25 Values of AH: for the Ag halides, HCI and NaBPh4 varied between - 5 and 27 kcal rnol-l, and consequently the errors in AH: values for these electrolytes may be as high as k0.5 kcal rno1-l. The collected values are given in table 3.90 ELECTROLYTES IN WATER AND ACETONITRILE ENTROPIES OF TRANSFER Standard entropies of transfer, AS:, have been obtained by the application of eqn (4) to the data in tables 2 and 3. The values are listed in table 4. To facilitate comparisons AG;' = AH,", - TAS,O, (4) of the relative effects of AH" and AS" in determining the total free energy changes, they are quoted as values of 298ASP, (corresponding to 25°C).TABLE 4.-ENTROPIES OF TRANSFER (AS;)' OF ELECTROLYTES FROM WATER TO MIXTURES OF WATER AND ACETONITRILE AT 25°C values of 298 AS,,/kcal mol-I (molar scale) AN C LiCl 0 0.0 0.10 -0.5 0.20 -0.9 0.30 -1.6 0.40 -2.3 0.50 -3.0 0.60 -4.1 0.70 -5.5 0.80 -7.6 0.90 -12.1 0.95 -16.3 1.00 -12.3 +AN c RbCl 0 0.0 0.10 -0.6 0.20 -1.4 0.30 -2.2 0.40 -3.0 0.50 -3.9 0.60 -4.9 0.70 -6.1 0.80 -8.0 0.90 -12.1 0.95 -13.4 1.00 -13.2 NaCl 0.0 - 0.5 - 1.0 - 1.6 - 2.2 - 2.9 - 3.7 - 5.0 - 7.0 -11.5 - 14.7 -11.8 CSCl 0.0 - 0.5 - 1.3 -2.1 - 2.8 - 3.6 - 4.6 - 5.8 -8.1 - 12.7 - 14.9 - 12.9 NaBr 0.0 - 0.5 - 1.0 - 1.7 - 2.4 - 3.2 - 4.2 - 5.3 - 6.8 - 10.6 - 13.4 - 11.9 AgCl'* f 0.0 - 4.6 - 5.6 - 5.9 - 6.1 - 6.5 - 7.2 - 8.2 - 9.6 - 12.0 - 13.5 - 12.8 KCl 0.0 -0.6 - 1.3 - 2.0 - 2.8 - 3.7 - 4.8 - 5.9 - 7.5 -11.2 - 13.9 - 12.7 AgBr f 0.0 - 4.5 - 5.9 - 6.3 - 6.5 - 6.9 - 7.5 - 8.2 - 9.2 - 10.9 - 12.0 - 13.2 KBr 0.0 - 0.6 - 1.5 - 2.4 - 3.1 - 4.0 - 5.0 - 6.2 - 7.5 - 10.9 - 13.0 - 12.6 &I f 0.0 - 5.2 - 6.0 - 6.4 - 6.7 - 6.7 - 7.7 - 8.5 - 9.4 -11.3 - 11.9 - 12.6 KI 0.0 - 0.8 - 1.5 - 2.4 - 3.1 - 4.0 - 4.9 - 6.0 - 7.7 - 10.3 - 11.8 - 13.6 Bu4NBr 0.0 4.1 6.8 7.9 8.6 9.2 9.7 9.8 9.4 8.8 aValues obtained by application of eqn (4) to the data in tables 2 and 3 ; estimated error 1 cal = 4.184 J ; C volume fraction of acetonitrile ; dAGtr values from ref.(19) ; e AGpr values from ref. (15), (24); festimated uncertainty t0.5 kcal mol-' ; B ref. (26). 0.2 kcal mol-1 unless otherwise stated ; DISCUSSION FREE ENERGIES OF TRANSFER The results in table 2 show that, with the exception of AgI, all of the simple electrolytes are more soluble in water than in acetonitrile.KBPh4, Ph4AsI and B u , N B ~ , ~ ~ on the other hand, are all more soluble in acetonitrile. More informative, however, is the way in which the free energies vary with solvent composition in the mixtures.15* 27-29 Thus if an ion is strongly solvated by one of the solvent com-B . G . COX, R . NATARAJAN A N D W . E . WAGHORNE 91 ponents, addition of the second, more weakly interacting, component will have little effect on its free energy until the last few percent of the " better '' solvent is removed, when the free energy will increase rapidly. Conversely, addition of small amounts of a "good'' solvent to an ion in a poorly solvating medium will result in a sharp decrease in its free energy.Strong solvent-solvent interactions may complicate the situation, but water and acetonitrile do not interact strongly with one another (their mixtures show positive deviations from Raoult's law).30 The results in table 2 can be readily understood on this basis, if it is assumed that the alkali metal cations and halide anions are more strongly solvated by water, and Ag+ and the " organic " ions, Ph,As+, BPh; and Bu,N+, more strongly solvated by acetonitrile. I . I GI- AgCl Na+ t ao 0.2 0.4 0.6 0.8 1.0 ~ C H ~ C N FIG. 1.-Free energies of transfer of NaCl and AgCl from water to acetonitrile+water mixtures at 25°C. This behaviour may be clearly seen in fig. 1, in which the free energies of transfer of NaCl and AgCl are plotted against the volume fraction of acetonitrile.Also included are the single ion values for Ag+, Na+ and C1-, based on the assumption AG,",(Ph,As+) = AG,"(BPh;) applied to the results in table 2. The free energy of NaCl increases slowly with increasing acetonitrile content of the solvent, until = 80 vol % acetonitrile, beyond which it increases rapidly to its value in pure acetonitrile. The free energy of AgCl, on the other hand, passes through a minimum. This is typical of the behaviour expected for a binary electrolyte, of which one ion is bound more strongly by one solvent compound and the other ion prefers the second component.2 9 s 31 We have shown earlier that in mixtures of dipolar aprotic solvents, where specific solvent-solvent interactions are small, free energies of transfer ions can be very satisfactorily accounted for in terms of a simple model in which the solvents are regarded as competing ligands for coordination sites of the ions.11* 15* 2 8 Excellent agreement is obtained between experimentally estimated Act", values, and values calculated from measured stability constants for coordination of the ions with the solvents involved.The present results are clearly in qualitative agreement with such a concept, although quantitative calculations are difficult because of the lack of stability constant data, and the non ideality of the solvent mixtures. However, it is known from potentiometric titrations 32* 33 that the complex Ag(CH,CN)Z is quite stable in aqueous solutions, and this, together with the absence of similar complexes with alkali metal cations, supports the above interpretation.92 ELECTROLYTES IN WATER A N D ACETONITRILE ENTHALPIES A N D ENTROPIES The relatively simple behaviour of the free energies of transfer is clearly not reflected in the enthalpies and entropies of transfer (tables 3 and 4).The transfer of both silver and alkali metal halides to acetonitrile results in very large decreases in entropy, the total entropy loss (TASP, = - 12.5 & 1.0 kcal mol-l) being relatively independent of the electrolyte involved. The variation of the entropy with solvent composition is, however, very different for silver and alkali metal salts. This may be seen in fig. 2, where AH; and -TAS,", for NaCl and AgCl are plotted against the volume fraction of acetonitrile.1 0 : : $1 a 0.0 0.2 0.4 0.6 0.8 1.0 k H X N FIG. 2.-Enthapies and entropies of transfer of NaCl(0) and AgCl (+) from water to acetonitrile+ water mixtures at 25°C. In discussing the entropy results, it is useful to consider firstly the effects expected in the immediate neighbourhood, or coordination sphere, of the ions. In all cases, there will be a loss in entropy on dissolving the electrolyte, resulting from the loss of translational entropy of the solvent molecules. More important in the present context, is that in a mixed solvent, where the ions are preferentially solvated by one of the components, there will be an unfavourable entropy term (relative to the ions in either of the pure solvents) because the composition of the coordination sphere will be different from that of the bulk solvent mixture, i.e.there will be an entropy decrease resulting from solvent sorting. If the composition of the solvation sphere of the ion is known, e.g. from n.m.r. measurements 29 or calculated from known stability constants for complex formation,ll' 15* 28 then the entropy losses may be readily calculated from simple statistical thermod ynamics.l We have shown, for example, that the entropy of transfer of Ag+ salts from propylene carbonate to dimethyl- sulphoxide (AG& large, -ve, AStf z 0) passes through a sharp minimum at 3 vol % dimethylsulphoxide (with ASFr falling to z - 20 cal K-l mol-1 or 29SAS,O, = - 6 kcal mol-l), directly attributable to solvent sorting around Ag+.l The effect is discussed in detail elsewhere, but model calculations show that as the strength of interactionB .G . COX, R . NATARAJAN AND W . E . WAGHORNE 93 with the more strongly solvating component increases, the minimum in AS& increases in magnitude, and moves to lower fractions of this component.ll In the present system, the free energy data (table 2 and fig. 1) suggest that Agf will be preferentially coordinated to acetonitrile and the alkali metal and halide ions will be preferentially coordinated to water. This should result in a large loss in entropy (increase in -TAS,") at low volume fractions of acetonitrile for Ag+ salts and at high fractions of acetonitrile (low water content) for the other ions. This is shown schematically in fig.3, and such an effect can clearly be seen in the results in fig. 1, 2 and table 4. Single ion values, obtained as before from the Ph4Ag+, BPh; assumption, suggest that these entropy losses are comparable for the alkali metal and chloride ions, but much smaller for bromide and iodide ions. 0.0 0.2 0.4 0.6 0.8 1.0 k H C N FIG. 3.-Entropy changes resulting from preferential solvation of ions in acetonitrile + water mixtures. Superimposed on the minima in the entropies resulting from solvent sorting, is the large net decrease in entropy accompanying the transfer of electrolytes from water to pure acetonitrile, which may be associated with bulk solvent properties. The more extensive structure of water compared with acetonitrile means that the loss in translational entropy on coordination to an ion will be smaller for water molecules than acetonitrile (cf.the considerably lower entropy losses on freezing of water compared with dipolar aprotic In addition, more positive entropies may result from the disruption of the water structure by the hydrated ions, whose bound water molecules are not correctly orientated to fit into the hydrogen bonded structure of ti What is perhaps surprising, however, is that the effect is important even at high fractions of acetonitrile. In fact, at least half of the observed total increase in entropy on transfer from acetonitrile to water has occurred upon addition of only 20 vol % of water, where the average distance apart of the water molecules is five times that in pure water. This may be seen for example from the results for KI, where solvent sorting effects should be smaller, or from the results for the alkali metal chlorides and bromides, when reasonable allowance is made for entropy losses arising from solvent sorting.This suggests that long range order in water is not of major importance in determining entropy effects. Rather, similar effects may result from considerably smaller, but presumably still strongly hydrogen bonded, groups of water molecules. The entropies of transfer of Bu,NBr and Ph4As+ and BPh; salts (values for which may be obtained by appropriate combinations from the results in tables 2 and 3) differ from the alkali metal salts in two ways : first transfer to acetonitrile leads to a considerable increase in entropy, and secondly the major change occurs at low fractions of acetonitrile.This behaviour is very similar to that of neutral organic94 ELECTROLYTES I N WATER AND ACETONITRILE molecule^,^ and is consistent with the notion that the addition of such molecules increases the hydrogen bonded structure of water.' The ions are presumably preferentially solvated by acetonitrile in the mixtures, so that the effect will disappear after the addition of relatively small amounts of acetonitrile. The enthalpies of transfer may be discussed in an analogous way to the entropies. In mixtures of dipolar aprotic solvents, these are determined almost entirely by the relative strengths of the interactions with the two solvent cornp0nents.l This effect would also account for the rapid changes in AH,, of the alkali metal halides at low water concentrations, and for the silver halides at low acetonitrile concentrations. It is also noticeable that diflerences between the heats of transfer of pairs of alkali metal and silver halides from water to acetonitrile agree closely with differences between the corresponding free energies of transfer, e.g. AG,,(NaCl) - AG,,(AgCl) = 8.9 kcal mol-1 ; AHt,(NaC1) -AHt,(AgC1) = 9.5 kcal mol-l.The absolute enthalpies of transfer, however, bear no simple relationship to the free energies of transfer. This is not surprising, as the entropy results show that large effects associated, with solvent reorganisation are involved, and these must have corresponding enthalpy effects, e.g., disruption of the water structure leads to an increase in entropy, but also to an unfavourable increase in enthalpy, because of the breaking of hydrogen bonds.The apparently simple behaviour of the free energies of transfer of electrolytes, strongly suggests that the contributions of solvent structural effects to ASFr and AH; terms largely cancel, thus making little overall contribution to the free energies of the ions. In conclusion, the enthalpies and entropies of transfer of electrolytes from water to acetonitrile show a complex dependence upon the solvent composition, reflecting the interactions between the ions and the neighbouring solvent molecules, as well as the effect of the solvated ions on the bulk solvent properties. 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