J. Phys. Chem. 1985, 89, 5093-5097 Finally, the results on solubility measurements, together with the spectroscopic data, are given for Sudan Red B in various organic solvents in Table IV. The solubility of the dye in the micelle “phase”, which can be given by Se/V,is calculated to be M for M for spherical micelles and 1.7 X (0.7-1.1) X rodlike micelles, respectively, where P i s the partial molar volume of surfactant in the micelles, 0.273 L These ranges of solubilization correspond to the solubility of Sudan Red B in acetone and chloroform, respectively. When we refer to the spectroscopic data, the state of solubilized Sudan Red B is more similar to that in a water-acetone ( 2 5 v/v) mixture. Sudan Red B in the spherical micelle would be in a microenvironment similar to acetone mixed with a less amount of water, and Sudan Red B in the rodlike micelle is in a state such as in chloroform. Thus, solubilized Sudan Red B would by only partly buried in a spherical micelle, while it is located more deeply inside the rodlike micelle.
5093
This is reasonable when the size of the Sudan Red B molecule is compared with that of the two kinds of micelles. However, Zachariasse et al.44measured the absorption spectra of a probe, a derivative of pyridinium N-phenolbetaine called E;(30), and concluded that the inside of rodlike micelles of sodium dodecyl sulfate is less polar than that of its spherical micelles. It is more likely that the size of the solubilizate molecule is too large to be incorporated in a spherical micelle but can be well accommodated in a rodlike micelle. Registry No. Sudan Red B, 3176-79-2;NaCI, 7647-14-5;dodecyldimethylammonium chloride, 2016-48-0; methanol, 67-56-1; ethanol, 64-17-5; acetone, 67-64-1; 1-butanol, 71-36-3; chloroform, 67-66-3; benzene, 71-43-2; n-hexane, 110-54-3;dimethyldodecylamine, 112-18-5. (44) K. A. Zachariasse, N. van Phuc, and B. Kozankiewicz, J . Phys. Chem., 85, 2676 (1981).
Preferential Solvation and Hydrophobic Hydration of Polyols in Mixtures of Water and N ,N-Dimet hylformamide R. W. Balk and G. Somsen* Department of Physical Chemistry, Free University, De Boelelaan 1083, 1081 H V Amsterdam, The Netherlands (Received: February 11, 1985; In Final Form: July 12, 1985)
Enthalpies of solution of nine polyols in binary solvent mixtures of N,N-dimethylformamide and water and of one polyol in mixtures of formamide and water at 298.15 K over the whole mole fraction region are reported. The deviation of linear behavior at high water content is considered to result from hydrophobic hydration of the apolar sites at the poly01 molecule and the curvature at low water content from preferential hydrogen bonding of the hydroxy groups of the solute. The effects are described by a simple two-parameter hydrophobic hydration model and a thermodynamic theory of preferential solvation, respectively. With regard to the enthalpies, the relative importance of the two solvation mechanisms appeared to depend on cosolvent and stereochemical detail of the solute.
Introduction In recent years much attention in our laboratory has been paid to the solvation behavior of hydrophobic solutes in binary aqueous solvent mixtures. In this respect, enthalpies of solution (AHso,) of a number of electrolytes and nonelectrolytes have been measured in D M F / H 2 0 (DMF = N,N-dimethylformamide) mixtures over the whole mole fraction region.’-) D M F appeared to be a suitable cosolvent, because it is aprotic and fully miscible with water. The hydrophobicity of solutes manifests itself in hydrophobic hydration (HH). This hydration is characterized by an enhanced hydrogen bonding between water molecules in the neighborhood of the nonpolar groups in water and gives rise to an exothermic contribution to the enthalpy of solution in pure water. Addition of an aprotic cosolvent breaks down this kind of hydration rapidly so that plots of AH,,, vs. the mole fraction of cosolvent exhibit pronounced endothermic shifts. In previous papers our group proposed a model to describe the enthalpic effect of HH.4-6 In this model cooperative formation of hydrogen-bonded clathrate-like envelopes of water molecules around the apolar groups is assumed. The model described the experimental results in terms of two parameters; especially one of these, Hb(W), appeared to be a good quantitative measure of the hydrophobic character of the s01ute.~ Description of the experimental results by the model was found to be good in cases of pronounced hydrophobic character of the (1) Heuvelsland, W. J. M.; Bloemendal, M.; De Visser, C.; Somsen, G.J . Phys. Chem. 1980.84, 2391. (2) Rouw, A. C.; Somsen, G.J . Chem. Soc., Faraday Trans. 1 1982, 78, 3397. (3) Booij, M.; Somsen, G. J . Chem. Soc., Faraday Trans. 1 1982,78,2851. (4) de Visser, C.; Somsen, G. J . Phys. Chem. 1974, 78, 1719. (5) de Visser, C.; Heuvelsland, W. J. M.; Somsen, G. J . Solution Chem. 1975, 4, 311. ( 6 ) Heuvelsland, W. J. M.; Somsen, G.J . Chem. Thermodyn. 1976,8, 873.
0022-3654/85/2089-5093$01.50/0
TABLE I: Names, Abbreviations, and Molecular Conformations of Polyols
solute: CH,OH(CHOHj,CH,OH ethylene glycol . -_ glycerol meso-erythritol xylitol
D-arabitol L-arabitol ribitol
D-sorbitol D-mannitol
i 0 1 2 3 3 3 3 4 4
abbrev
conformation of nonterminal C
EN GLYC
m-ER XYL
D-AR L-AR RIB D-SORB D-MAN
dd did ddl dll ddd ddld ddll
solute. However, upon introduction of functional groups (-OH, -NH, -NH2) in the solute molecule, the description became wor~e.~~~*~ In connection with this it seemed interesting to extend our studies to hydrophilic nonelectrolytes. We have chosen polyols, because these can be studied as a function of (1) the number of functional groups and of (2) their stereochemical conformation. The general formula of these compounds is given by CH,OH(CHOH),CH,OH ( i = 0-4). Polyhydroxy compounds are also of biological importance: in life processes they are implicated in recognition phenomena between cells and the various components of the immune system. They also find application in the in vivo protection of organisms against desiccation whether through freezing or drought.8-10 Thus, recent times have witnessed a ( 7 ) Rouw, A. C.; Somsen, G. J . Chem. Thermodyn. 1981, 13, 67. (8) Crowe, J. H.; Crowe, L. M. Cryobiology 1982, 19, 317. (9) Duman, J. G.Cryobiology 1982, 19, 613.
0 1985 American Chemical Society
5094
Balk and Somsen
The Journal of Physical Chemistry, Vol, 89, No. 23, 1985 W
DMF
A Htr (W-M) / k J mor’
ce
12
EN
DMF
W
t
-
l4
GLYC m-ER
L - AR ____ _. D-AR
12/
D-AR
XY L
D - SORB
I
4 -
xw
0 0
02
04
06
00
RIB
0
10
0
Figure 1. Dependence on the size of the polyol: enthalpies of transfer of five polyols vs. the mole fraction of water in D M F H20at 298.15
+
K.
renewed interest in the details of solute-solute interactions and hydration of these compounds in aqueous solution.l*17 The behavior of the bifunctional molecules appeared to be very complex; the purpose of our investigation is to contribute to a further elucidation of their hydration properties. The polyols of this study are presented in Table I, together with the configuration of their nonterminal carbon atoms and the abbreviations used in the text.
Experimental Section The enthalpies of solution were measured with an LKB 8700-1 precision calorimetry system equipped with a 1OO-cm3 reaction vessel. The experimental procedure and test of the calorimeter have been reported earlier.l* N,N-Dimethylformamide (Baker, analyzed reagent) was stored over molecular sieves for at least 48 h and used without further purification; formamide (FA; Baker, analyzed reagent) was freshly distilled under reduced pressure before use. The mole fraction of water for these solvents as determined by a modified Karl Fischer19 titration was less than 5 X The solvent mixtures were prepared by mass, and distilled deionized water was used. meso-Erythritol (Aldrich), the pentitols (Sigma), and the hexitols (Sigma) were dried in vacuo at room temperature for 48 h and used without further purification; glycerol (Aldrich) was kept over molecular sieves 3A for 1 week and distilled twice in vacuo. The mass fraction of water in the middle fraction of the distillate was less than 2 X Enthalpies of solution of ethylene glycol were obtained during an earlier investigation in our laboratory.20
02
04
06
00
10
Figure 2. Dependence on the conformation of the polyol 1: enthalpies of transfer of four pentitols vs. the mole fraction of water in DMF + H 2 0 at 298.15 K. DMF
t
D - SORB
D -MAN
l4
XW
Figure 3. Dependence on the conformation of the polyol 2: enthalpies of transfer of two hexitols vs. the mole fraction of water in DMF H,O at 298.15 K.
+
AHtr(W-.M)/kJ 10
mor’
-
FA/H20 DMF/ H2 0
Results The values of AH,,, for the various polyols are presented in (10) Franks, F.; Pedley, M. D. J . Chem. Soc., Faraday Trans. 1 1983.79, 2249. (1 1) Franks, F.; Ravenhill, J. R.; Reid, D. S . J . Solution Chem. 1972, 1 ,
3. (12) DiPaola, G.; Belleau, B. Can. J . Chem. 1977, 55, 3825. (13) Barone, G.; Bove, 9.; Castronuovo, G.; Elia, V. J . Solution Chem. 1981, 10, 803. (14) Jasra, R. V.; Ahluwalia, J. C. J . Solution Chem. 1982, 11, 325. (15) Tasker, I. R.; Wood, R. H. J . Solution Chem. 1983, 12, 315. (16) Abate, V.: Barone. G.;Castronuovo. G.:Elia. V.; Rizzo, E. J. Solution Chem.’1983, 12, 645. (17) Barone, G.; Elia, V.; Rizzo, E. J. Solurion Chem. 1982, 11, 687. (18) Heuvelsland, W. J. M.; de Visser, C.: Somsen, G. J . Phys. Chem. 1978, 82, 29.
(19) Verhoef J. C.; Barendrecht, E. J . E/ectroana/. Chem. Interfacial Electrochem. 1977, 75, 705. (20) Rouw, A. C.; Somsen, G., unpublished results.
Figure 4. Dependence on the cosolvent: enthalpies of transfer of mesoerythritol in two different binaries vs. the mole fraction of water at 298.1 5
K.
Table 11. They are generally taken to be the average of 2-6 measurements, which are carried o u t in such a way that the final concentrations ranged from to IO-* mol dm-). No concen-
The Journal of Physical Chemistry, Vol. 89, No. 23, 1985 5095
Poiyols in D M F / H 2 0 Mixtures TABLE II: Enthalpies of Solution of Polyols in DMF
+ H20at 298.15 K AHWl/kJmol-' D-AR
L-AR
RIB
D-SORB
D-MAN
x,
mol-', EN
[28.96]' [28.32]' [27.91]" [27.52]" [27.08]" [26.54]" [25.89]"
27.69 27.73 28.15 28.75 29.81 31.13 32.38
25.43 25.53 25.90 26.63 27.58 28.56 29.72
24.21 24.31 24.40 24.86 25.87 27.10 28.30
25.40 25.16 25.34 25.95 27.01 28.30 29.33
21.93 22.03 22.47 23.33 24.60 26.35 27.70
28.91 28.97 29.22 29.88 31.23 32.55
1.563 1.183 -1.097 -5.870
31.55 [25.15]" 30.75 [24.55]" 28.13 [23.63]" 22.66
32.77 31.86 28.96 23.33
30.65 29.48 26.48 20.63
28.70 28.12 24.80 19.10
29.74 28.98 25.48 19.20
28.28 26.91 23.24 16.76
34.28 34.45 32.69 29.07 22.79
0.0000 0.0599 0.1891 0.3251 0.4500 0.5500 0.6500 0.7706 0.8500 0.9000 0.9500 1.0000
-0.85 -1.06 -1.03 -0.74 -0.13 0.42 0.92 0.68 -0.66 -2.00 -4.00 -6.85
f0.017
f0.05 [f0.05]"
f0.06
f0.06
f0.05
fO.05
f0.04
f0.06
GLYC
0.000 0.100 0.200 0.300 0.400 0.500 0.600 0.650 0.700 0.800 0.900 1.000
-1.422 -1.522 -1.459 -1.120 -0.544 0.335 1.138
27.88 27.86 28.06 28.37 29.03 30.02 31.01
66
In FA
m-ER
*0.03
+ H20. Mean deviation averaged over all mole fractions (in kJ mol-')
tration dependence of the enthalpies of solution was observed, so the measured enthalpies of solution were regarded as those at infinite dilution, AHml".They all refer to a temperature of 298.15 K. Comparison of our values of AHw,in pure water for the various polyols with those of Jasra and A h l ~ w a l i a shows '~ that there is good agreement for D-SORB and reasonable agreement for m-ER but that our values of AHsoIfor D-MAN and the pentitols are 0.7 (RIB) to 1.9 kJ mol-' (D-AR) more endothermic. Contrary, our value of AH,,, of GLYC in pure water is about 10% more exothermic than the two values quoted by Franks and Reid.2' Enthalpies of transfer from water (W) to mixtures of D M F W (M) or FA W (M') vs. the mole fraction of W are presented in Figures 1-4.
+
AHsollkJ
XYL
x,
+
Discussion Figures 1-3 show that, in the DMF-rich region of the D M F / H 2 0 mixtures (x, C 0.6), the enthalpy of transfer curves of the polyols are not linear and increase with increasing water content. At high x,, AH,,(W M) (or AH,,l") reaches an endothermic maximum at about x, = 0.7, after which a decrease to the value in pure water occurs. Thus, in this respect the general behavior of these hydrophilic solutes is much alike that of alcohols,' ureas, and amides:2 the curves are characterized by a curvature at low water content and exothermic shifts in the water-rich region. Figure 2 also shows that roughly speaking, both the exothermic shifts at x, > 0.7 and the curvature at x, C 0.5 increase upon increasing size of the polyol. On the other hand, Figures 2 and 3 show that the exothermic shifts and/or the curvature depend on the conformation of the polyol. Considering the enthalpies of transfer from pure W to pure DMF, AH,,(W-DMF), no obvious trends are observed. They vary between 4.36 (XYL) and 6.20 kJ mol-' (RIB) in a rather irregular way. In comparison, the enthalpies of transfer from W to D M F of ethanol to n-pentanol vary between 11.46 (ethanol) and 12.31 kJ mol-' (n-propanol) and are not very dependent on the size of the alkyl group either.' The only difference is that, where AH,,(W-DMF) of the polyols is somewhat smaller than that of methanol (Le., 6.75 kJ mol-'), those of the higher (primary) alcohols are about 5 kJ mol-' larger. Behavior at x, > 0.6. Endothermic shifts of the enthalpies of solution upon small additions of an aprotic cosolvent to water are known to appear for 'hydrophobic" solutes.22 They are attributed to the enthalpic effect of hydrophobic hydration (HH) of the solutes in the highly aqueous mixture^.^ We ascribe the shifts observed for the polyols also to this effect. Upon increasing size of the polyol molecule, the number of apolar C-H sites at the molecule also increases, resulting in an increase of H H and thus also of the endothermic shifts. This behavior is
-
(21) Franks, F.; Reid, D. S. In "Water, A Comprehensive Treatise"; Franks, F., Ed.; Plenum Press: New York, 1975; Vol. 2, Chapter 5. (22) Cifra, P.; Romanov, A. J . Solurion Chem. 1984, 13, 431.
thus quite comparable with that of "hydrophobic" solutes. For example, the value of Hb(W) of D-SORB with 8 C-H bonds (-18.3 kJ mol-') lies just between those of n-PrOH with 7 (-17.4 kJ mol-') and n-BuOH with 9 C-H bonds (-19.6 kJ mol-'). The Hb(W) value is a measure of the enthalpic effect of HH; in this case the values are approximated by a linear extrapolation of the linear part of the enthalpy curve (0.4 C x, C 0.6) to x, = 1.0 [Hb(W) = AHsoI"(W) - AHsoFtraP ( x , = 1.0)]. Furthermore, if one compares the enthalpy curves of m-ER in D M F / H 2 0 and F A / H 2 0 mixtures, the endothermic shifts almost disappear upon going to the F A / H 2 0 mixtures (see Figure 4). A similar behavior has been found for e.g. (Bu,N)Br, which is assumed to be a typical hydrophobic solute.23 The enthalpy curve of this solute exhibits a pronounced endothermic maximum at about x, = 0.7 in the D M F / H 2 0 mixtures, which totally disappears going to the F A / H 2 0 system. At the same time, the strong endothermic shifts are about halved. Thus, although the polyols cannot be viewed as "hydrophobic" solutes,21 HH occurs upon their solvation in water and water-rich mixtures. In our opinion, HH is not limited to hydrophobic solutes but occurs as soon as an apolar part of the molecular surface (e.g. a C H group) is exposed to liquid water. This is in fact observed for formamide,2 a compound which possesses only one C H bond. A very recent Monte Carlo study on the hydration of formamide strongly supports this view.24 Behavior at x, C 0.6. In this region, the course of the enthalpies of solution of the polyols exhibits an exothermic curvature, which in some cases passes through a slight minimum (EN, GLYC, RIB). Such a curvature shows up upon introduction of functional groups in the solute molecule (Le., -OH, -NH, -NH2) and also increases upon increasing number of functional groups. This effect is nicely shown by the work of Rouw2 on alkyl-substituted ureas in D M F / H 2 0 mixtures: in the case of TMU (tetramethylurea), no curvature is observed but exothermic shifts at x, > 0.6 are, while for U (urea), a compound which does not possess any -CH group and can be viewed as consisting of functional groups only, the enthalpy curve exhibits a curvature and no shifts. Thus, without HH, Le., in the absence of exothermic shifts at x, > 0.6, the enthalpy curves of the polyols would be similar to that of U.I7 Very pronounced curvature in the water-poor region has been found for e.g. LiBr and NaI in acetonitrile/water mixture^.^^*^^ As pointed out by Covington and co-workers, both the alkali and the halide ions are preferentially hydrated to a large extent in these mixtures.27 Thus, assuming that ideal behavior (Le., no preferential hydration) would give a linear or nearly linear path over the whole mole fraction region of the binary,4 the excess enthalpy curves of these electrolytes will be due to the enthalpic effect of (23) de Visser, C.; Somsen, G.J . Solution Chem. 1974, 3, 847. (24) Marchez, F. T.; Mehrotra, P. K.; Beveridge, D. L. J . Phys. Chem. 1984, 88, 5692. (25) Heuvelsland, W . J. M. Thesis, Rodopi, Amsterdam, 1980. (26) Stennikova, M. F.; Poltorazskii, G. M.; Mishchenko, K. P. Zh. Obshch. Khim. 1971, 4 1 , 2588; Zh. Strukt. Khim. 1972, 13, 143. (27) Covington, A. K.; Thain, J. M. J. Chem. SOC.,Faraday Tram. 1 1974, 70, 1879.
50%
The Journal of Physical Chemistry, Vol. 89, No. 23, 1985
their preferential solvation (PS).28 We ascribe the curvature observed for the nonelectrolytes also to the enthalpic effect of PS, meaning PS of their functional groups. These groups will be hydrogen bonded to solvent molecules in solution but will be preferentially hydrogen bonded to one of the components of the solvent mixture. Such a behavior is not unexpected, since the hydrogen-bond acceptor and donor properties of e.g. DMF and water are different, which appears from the work of Kamlet et al.29 Model Approach. As shown by De Visser et al.: the enthalpies of solution of hydrophobic solutes in mixtures of water and D M F can be expressed as a function of the mole fraction of water, x,, by ~ s 0 l 0 ( M= ) ( x , ~ s 0 l 0 ( W * )+ (1 - x,)A~sol0(DMF)1+ xwrnHb(W) (1) In this equation, which is based on a simple clathrate model, AHmIo(DMF)denotes the enthalpy of solution in pure D M F and AHsOlo(W*)a hypothetical one in pure water, Le., without the enthalpic effect of H H , Hb(W): AHsolo(W*) = AHsolo(W)- Hb(W) (2) In the earlier papers of our group about this s ~ b j e c t , ~m- was ~ considered to be equal to the number of water molecules forming a clathrate-like cage around the hydrophobic sites or parts of the solute molecule, but later on it was shown that the value of m is not unambiguously defined by the model.'-1s This is not surprising, because m includes effects from (1) HH by cooperative action of a certain number of water molecules, (2) interaction with cosolvent molecules, and (3) competition because of hydration of cosolvent molecules. The first part of eq 1 describes the linear variation of AHsolo(M)with x, without HH and, regarding the polyols, also without the enthalpic effect of PS. The second part in fact gives a phenomenological description of the exothermic shifts upon increasing water content, a description which we shall also adopt for the shifts of the polyols. For a quantitative approach of the enthalpic effect of PS,which shows up as an excess effect in the enthalpy curve of e.g. U (vide supra), we used the simplest version of the thermodynamic theory on PS of Covington et aL30 Although developed for electrolytes in binary mixtures, it was recently shown by Remerie et aL3' to work also quite well for nonelectrolytes. The theory analyzes the change in composition of the solvation shell with solvent composition for a solute S in a mixture of solvent components W and P. When the mixture becomes richer in one of its components, the change in the composition of the solvation shell of S is described by a successive series of n equilibria, where n is the solvation number:
+
K,
+
S(W,-lPn+l-i) W .G S(WiP,+) P (1 I i 5 n ) (3) Assuming that the Gibbs energy change accompanying the substitution of one molecule P by one molecule W is equal for all equilibria, it can be shown that Ki = K I I n [ ( n+ 1 - i ) / i ]
(4)
in which K is the equilibrium constant of the overall process K =
n
~ I K ~
i= 1
After considerable simplification the following equation for the Gibbs energy of transfer of S from water (W) to a mixed solvent with mole fraction 1 - x, of P can be obtained
-
AG,,(W M) = -nRT In [K-'/"(1 - x,) + x,] (6) in which the contributions resulting from long-range electrostatic interactions are left out and solvent mole fractions are used instead (28) Cox, B. G.; Waghorne, W. E. J. Chem. Soc., Faraday Trans. I 1984, 80, 1267. ( 2 9 ) Kamlet, M. J.; Taft, R. W. J . Am. Chem. Soc. 1976, 98, 377. (30) Covington, A. K.; Newman, K. E. Adu. Chem. Ser. 1976, No. 155, 153. (31) Remerie, K.; Engberts, J. B. F . N . J . Phys. Chem. 1983, 87, 5449.
Balk and Somsen of activities. The reason for the last simplification is that the use of mole fractions leads to the best reproduction of experimental data, which is apparently the result of the large difference in activity coefficients of solvent components in bulk solvent and in the solvation she11.29s30Furthermore, this expression does not account for effects from nonstatistical distributions of solvated species or changing solvation numbers. An expression for the entropy of transfer from W to the mixture M is derived by Cox and Waghorne2* and is given by A&(W
-
M) = -n,R In (n,/nx,)
-
n,R In (np/n(l - x,)) (7)
in which n, and n,, represent the number of molecules of W and P in the solvation shell of S ( n = n, + n,). Expressing n, and np in terms of fractions of solvated species and using AH,,(W
-
-
M) = AG,,(W
-
M)
+ TAS,,(W
-
M) (8)
we obtain the following formula for the enthalpy of transfer: AH,,(W
M) = n R T [ ( 1 - x,,)/{(l
-
x,)
+ K'/"x,)]
in K1in (9)
For the excess enthalpic effect of PS it follows that
AH,,~(M) = n R T [ { ( l- x,)/((l - x,)
+ K1/"x,)l- ( 1 - x,)]
In K1/" (10)
This function is negative over the whole mole fraction region (K1/" > 0), thus resulting in an exothermic contribution to the enthalpy of solution. The enthalpy of solution of a solute which exhibits H H due to apolar sites and PS due to functional groups is now given by
AHsoIo(M)= {X,AH,,~~(W*)+ (1 - X,)AH,,~~(DMF)J+ AHpsE(M) x,"Hb(W) ( 1 1 )
+
Analysis of the Enthalpy Curves. We have analyzed our experimental data in the D M F / H 2 0 system with this model approach. The fit procedure was that we firstly optimized the parameters Klin and Hb(W), assuming that n equals the number of functional groups times three (an -OH group possesses three sites for hydrogen bonding), and finally m, because the curvature in the region below x, = 0.5, where the effect of H H was expected to be negligibly small,18could be much better reproduced by the PS model than the exothermic shifts by the H H model. (The enthalpy curves of e.g. U in several binaries could be reproduced by the PS model within the experimental error. This is mostly not the case for the enthalpy curves of hydrophobic solutes by the HH The results for the parameters, together with the deviations of calculated values from the experimental ones (A) are given in Table 111. For reason of comparison, the values for methanol (MeOH) are also included in this table. Although one has to be very careful in drawing conclusions from these results because of interfering enthalpy effects of PS and HH, there are some noteworthy things in them. Firstly, the deviation, A, of which the larger part is determined by deviations in the water-rich region (x, > O S ) , is reasonable and much better than that found by application of the H H model only (compare e.g. ref 7). On the other hand, the results are such that in our opinion they do not justify a refinement of the model of PS. Secondly, the parameter Kiln is larger than one, which means that the functional groups are preferentially hydrated. On the contrary, the analysis of the enthalpy curve of m-ER in the F A / H 2 0 suggests PS by FA. In Figure 5 this analysis is shown. The -Hb(W) value used for this fit is the same as the one deduced from the DMF/H,O system. For KIIn = 0.44 the model reproduces the enthalpy of transfer curve within the experimental error (compare Table 11). Figure 5 also shows the distinct enthalpy contributions from PS and H H for the given set of parameters. This result for m-ER is supported by that for U in FA/H,0:32 (32) de Visser, C.; Grunbauer, H. J. M.; Sornsen, G . Z . Phys. Chem. (Munich) 1975, 97, 69.
The Journal of Physical Chemistry, Vol. 89, No. 23, 1985 5097
Polyols in D M F / H 2 0 Mixtures TABLE 111: Fit Parameters of Polyols in DMF + HzO config of no. of O H groups
n
p i n
no. of C H bonds
MeOH
1
3
1.88
3
EN dd ddl ddld
2 3 4 5 6
6 9 12 15 18
1.99 1.78 1.68 1.67 1.73
4 5 6 7 8
dld ddl dll ddd
5 5 5 5
15 15 15 15
1.69 1.67 1.74 1.84
ddld
6 6
18 18
1.73 1.77
solute
nonterminal C
GLYC m-ER D-AR D-SORB XYL D-AR L-AR RIB D-SORB D-MAN a
ddll
A = (IA&(calcd)
-Hb(W)/kJ mol-I 9.12
m
A"/kJ mol-'
8.02
0.09
11.5 11.3 13.6 16.2 20.3
6.83 8.08 7.22 6.55 5.99
0.06
0.12 0.13 0.12 0.21
7 7 7 7
16.1 16.2 16.5 18.8
6.45 6.55 6.78 6.21
0.16 0.12 0.16 0.09
8 8
20.3 21.2
5.99 5.41
0.21 0.20
- AHml(exptl)l).
FA
W
A Htr(W+M)/ k J mol-'
I
. ideal behaviour ._- + P S
,'
...I
j j
+
-
HH
final fit
I +
experimental points
,I
.,*----
/
--, ; '\
fit parameters
1'
m decreases upon increasing size of the polyol: the molecules possess an increasing number of more or less isolated apolar sites.I3 These will be responsible for a relatively large hydrophobic surface and will also favor the formation of small hydration cages. Finally, the variation of the parameters within a group of conformers also reflects their different enthalpy curves. However, there is no apparent correlation between the parameter values and the molecular configuration of the solutes in the crystal, although it seems as if -Hb(W) decreases when d and 1 configurations of the nonterminal carbon atoms are alternating. Similar observations for pair interaction coefficients have been made by other a~thors.'~~'~
Conclusions Hb(W)
- 13.6 kJ mol-' 0 0 5 k J mol-'
XW
Figure 5. Possible fit for the enthalpy of transfer curve of meso-erythritol in FA HzO.
+
the enthalpy curve of the solute can only be reproduced assuming that PS occurs by FA (thus, KIIn C 1). This means that the sharp decrease of the enthalpy of solution at x, N 0.7 upon going from DMF/H,O to F A / H 2 0 (see Figure 4) is not primarily caused by a decrease of HH in the F A / H 2 0 mixtures as suggested earlier1s~23 but by a large contribution of the enthalpic effect of PS in this mole fraction region. Thirdly, the -Hb(W) values reflect the variation of the endothermic maxima in Figures 1-3. Going from MeOH with 3 CH bonds to the hexitols with 8 C H bonds, -Hb(W) increases, although EN seems to be an exception:I2 its -Hb(W) value is even somewhat larger than that of GLYC. The increase of -Hb(W) is not regular going from GLYC to D-SORB, nor is a kind of leveling off for the higher polyols observed as was found for the primary alcohols or tetraalkylammonium halides. The -Hb(W) value also depends on the conformation of the polyol. On the other hand, m decreases upon increasing -Hb(W), which is also contrary to the behavior of the primary alcohols or tetraalkylammonium halides. However, going from primary to secondary or tertiary alcohols, the leveling off of -Hb(W) disappears and m decreases. These effects are ascribed to an increase of the effective hydrophobic surface of the alcohol molecules due to branching and the formation of smaller subcages around the smaller hydrocarbon side chains of the alcohol molecule, respectively.' If it is assumed that m at least reflects the size of the hydration cage formed around an apolar site (vide supra), then it is not surprising that
It has been shown that the enthalpy of solution curves of a number of polyols in D M F / H 2 0 mixtures could be described by a model which combines a simple version of the theory on preferential solvation of Covington and co-workers and the cage model on hydrophobic hydration. In our opinion this combined model is of general validity and in fact describes the enthalpy of solution curves of a wide variety of solutes in various binaries. For the polyols, the results point to the existence of two different solvation mechanisms in water: (1) hydrophobic hydration of the apolar C H sites and (2) specific hydration of the functional groups. This last mechanism is in fact already proposed by Stokes and Robinson to describe their experimental activity data for aqueous glucose and sucrose solutions.33 The two hydration mechanisms may act cooperatively: such an effect was confirmed for MeOH and TBA (tert-butyl alcohol) by Monte Carlo studies,34and also for FA.24 A quantitative analysis of the experimental enthalpy curves in principle provides the parameters for the various solutes giving information about the balance of hydrophilic and hydrophobic properties in a certain mixture. Although the parameters values are somewhat ambiguous because of interfering exothermic contributions of preferential solvation and hydrophobic hydration, especially in the mole fraction region x, > 0.5, they point to the fact that the polyols behave hydrophilic in D M F / H 2 0 mixtures (Le., preferential hydration of the hydroxy groups), contrary to their behavior in F A / H 2 0 . They also show that both hydrophobicity and hydrophilicity of the polyols are sensitive to stereochemical detail. Registry No. HzO, 7732-18-5; DMF, 68-12-2; FA, 75-12-7; EN, 107-21-1; GLYC, 56-81-5; m-ER,149-32-6; XYL, 87-99-0; D-AR, 488-82-4; L-AR, 7643-75-6; RIB, 488-8 1-3; D-SORB, 50-70-4; D-MAN, 69-65-8. (33) Stokes, R. H.; Robinson, R. A. J. Phys. Chem. 1966, 70, 2126. (34) Okazaki, S.; Touhara, H.; Nakanishi, K. J. Chem. Phys. 1984, 81, 890 and references therein.
