J. Phys. Chem. 1996, 100, 433-438
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Intermolecular Interactions in 2-Butoxyethanol-DMSO-H2O† Peter Westh‡ and Yoshikata Koga* Department of Chemistry, UniVersity of British Columbia, 2036 Main Mall, VancouVer, B.C., Canada V6T 1Z1 ReceiVed: August 31, 1995X
Excess partial molar enthalpy, HBE, and chemical potential, µBE, of 2-butoxyethanol (B) were determined in ternary mixtures of B, dimethyl sulfoxide (D), and H2O. The data were obtained in small enough mole fraction increments to evaluate the so-called interaction functions, ∂HBE/∂xB, ∂HBE/∂xD, ∂µBE/∂xB, and ∂µBE/ ∂xD. These interaction functions previously proved useful in elucidating the “mixing schemes” in binary aqueous solutions of B and D. For the binary mixtures, it was found that both B and D influenced H2O in the following manner: in the water-rich composition range (region I) within a certain threshold (xB < 0.0175 and xD < 0.28 at 25 °C), both solutes enhance the hydrogen-bonded network of water in their vicinity, and the mixtures retain the percolated nature of the network. At higher B or D concentrations (region II) a qualitatively different mixing scheme becomes operative. The results from this work suggest that, in the ternary mixtures, solute B and D influences the percolated hydrogen bond network of water competitively or cooperatively. The observed effects are in accordance with those characteristic of mixing scheme I in the binary mixtures, as long as the concentration of both solutes is within the threshold values. When either one of the solutes is concentrated beyond its threshold, mixing scheme II seems to set in. It was found that D diminished the positive (unfavorable) enthalpy of B-B interactions and that this effect was almost completely compensated by changes in interaction entropy. Hence, D had little net effect on the Gibbs energy of mutual B interactions, or in other words, D did not weaken the hydrophobic attraction between B molecules. Evaluation of “heterogeneous” B-D interactions in region I suggested that they were weaker than B-B interactions and governed by reorganization of water-water hydrogen bonding rather than interactions between specific groups in B and D. Some implications of these findings on cosolvent effects in aqueous solutions of biopolymers are discussed.
Introduction In previous studies it has been shown that molecular interactions in solution can be evaluated from precise measurements of partial molar quantities, such as chemical potential and partial molar enthalpy.1-3 So far this approach has been used to elucidate the complicated nature of binary aqueous solutions of some nonelectrolytes. In particular, composition derivatives of excess partial molar functions have been utilized to evaluate solute-solute interactions. In aqueous solutions of 2-butoxyethanol (B), for example, the composition dependence of the excess partial molar enthalpy1 and entropy2 of B, HBE, and SBE supported the view that the hydrogen bond network of solvent water is enhanced by B in the low B composition region (xB < 0.0175 at 25 °C, called region I). B-B interactions expressed by the composition derivatives of HBE and SBE were found to be unfavorable (i.e., repulsive) in terms of enthalpy while attractive entropy-wise in region I. Moreover, the entropic attraction surpassed the enthalpic repulsion, giving rise to a net attraction in terms of Gibbs energy.1-3 These findings are consistent with a classical picture of “iceberg formation” by a nonpolar solute4 and hydrophobic attraction among the solute molecules.5 It was also suggested that in this region the percolated nature of the hydrogen bond network is preserved.3 In the intermediate composition region (denoted region II), 0.0175 < xB j 0.5, the mixture seemed to consist of two kinds of clusters rich in each component. When the temperature is increased to about 50 °C, the lower critical solution temperature, †
This paper is dedicated to the late Dr. Aase Hvidt. Permanent address: Department of Chemistry, University of Copenhagen, 5 Universitetsparken, DK-2100, Copenhagen, Denmark. * Corresponding author. X Abstract published in AdVance ACS Abstracts, December 1, 1995. ‡
0022-3654/96/20100-0433$12.00/0
these clusters grow to macroscopic droplets and phase separation occurs. Hence, the hydrogen bond network is no longer percolated.3 In region II, B-B interactions are diminished and characterized by entropic attraction, which causes phase separation at high temperatures, and enthalpic repulsion, which is the driving force for the miscibility at lower temperatures.3,6 In dilute aqueous solutions of dimethyl sulfoxide (D)7 the excess partial molar enthalpy and entropy of D, HDE and SDE, show some resemblance with the excess functions of B mentioned above. Namely, in the region xD < 0.28, D-D interactions are enthalpically repulsive and attractive in terms of entropy. Likewise, the composition dependence of the excess partial molar volume of D,7 VDE, was found to resemble that observed for B,8 VBE. Unlike the B-H2O case, however, the enthalpic repulsion between solute molecules is larger than the entropic attraction in D-H2O mixtures, resulting in net repulsion among D molecules in terms of Gibbs energy. In the present work, we shall apply the same methodology, i.e., assessment of molecular interactions from the composition dependence of partial molar excess functions, to a threecomponent system. The partial molar quantities of a given component will provide information on the actual situation of this component regardless of the number of components in the solution. Integral molar quantities, on the other hand, would be less informative in a multicomponent system than in a binary system. One of the motivations for initiating thermodynamic investigations of molecular interactions in three-component systems is their potential as model systems for processes of biological interest. Dimethyl sulfoxide9,10 as well as other small alkyl compounds11 has been demonstrated to destabilize the conformational stability of globular proteins. Hydrophobic “binding” © 1996 American Chemical Society
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of the cosolvent to the protein surface12 has been suggested as a likely mechanism for this effect.9,13 Interactions between D and the mainly hydrophobic B elucidated in this study might throw some light on the propensity of D to “hydrophobically bind” to nonpolar surfaces in an aqueous environment. Data Analysis A partial molar quantity of a component A in solution, FA, signifies the contribution per mole of A to the integral quantity, F.
FA ) (∂F/∂nA)T,p,nB...
(1)
where nA is the number of moles of A. For component A, the sign of the derivative of FA with respect to nA, (∂FA/∂nA)T,P,nB..., would thus indicate whether incoming A molecules enhance or decrease the contribution of A to the F function. If we consider excess enthalpy (F ) HE), a situation where ∂HAE/∂nA > 0 implies that additional A makes the contribution (per mole) of A to the total enthalpy of the solution more positive; in other words, it makes the enthalpic situation of A in the solution less favorable. In such a case it may be said that A-A interactions are unfavorable or repulsive in terms of enthalpy. Similarly, if ∂HAE/∂nA < 0, A-A interactions make the enthalpic situation of A more favorable, and A-A interactions are denoted favorable or attractive in terms of enthalpy. In the present context, it is important to note that the above arguments are valid regardless of the number of components in the solution and that similar conclusions can be obtained for “heterogeneous” A-B interactions using the derivative ∂HAE/ ∂nB. Numeric values for derivatives of this type can be obtained if the concentration dependence of a partial molar quantity is known with sufficient accuracy. For convenience, the partial molar quantities are normally obtained as a function of mole fractions, in this case xB and xD. Hence, conversion to a mole fraction based variable system is performed in the following manner:
F(nW,nD,nB) ) F(xW,xD,N)
(2)
where x is the mole fraction of a component; N ) nW + nD + nB and F ) G, H, S, GB, HB, or SB. Differentiation of eq 2, keeping nW and nD constant gives
(∂F/∂nB)nW,nD ) (∂F/∂N)xB,xD + (∂F/∂N)xD,N (1 - xB)/N - (∂F/∂xD)xB,N (xD/N) (3) For an intensive property, such as FB, this reduces to
FBB ≡ N(∂FB/∂nB)nW,nD ) (∂FB/∂xB)xD,N (1 - xB) - (∂FB/∂xD)xB,NxD (4) Similarly, if nW and nB are kept constant, differentiation of eq 2 gives
FBD ≡ N(∂FB/∂nD)nW,nB ) (∂FB/∂xD)xB,N (1 - xD) - (∂FB/∂xB)xD,N xB (5) According to the arguments above, eqs 4 and 5 reflect effects, in terms of the thermodynamic property F, of B-B (eq 4) and B-D (eq 5) interactions in the ternary solution. In the
following, molecular interactions, FBB or FBD, shall be evaluated by the right-hand side of eqs 4 and 5. Experimental Section The excess partial molar enthalpy of B, HBE, was measured directly at 25 °C by titration calorimetry. The cell was loaded with a given water-D mixture and titrated with pure B. The procedures are described in detail elsewhere.14 Partial vapor pressures of B were measured by head space gas chromatography. (Total pressures of B-D-water mixtures were also measured utilizing a method similar to that employed for the binary (B-water and D-water) systems.2,7 However, due to azeotropic behavior for xB ≈ 0.01, attempts to estiamte partial pressures from these data remained unsuccessful.) For the gas chromatographic measurements, about 150 µL of a solution with known composition was transferred to a 2 mL vial. The vial was subsequently hermetically sealed with a silicon septum and allowed to thermostat at 25.00 ( 0.01 °C for several hours. The silicon septum, which was a few millimeters above the thermostat bath surface, was heated to ca. 40 °C to avoid condensation. A 45.0 ( 0.1 µL aliquot of vapor was removed from the head space by a warm (45 °C), gas-tight syringe and immediately injected in a Hewlett-Packard 5890 II gas chromatograph equipped with a capillary column and a flame ionization detector. The peak area due to B in binary B-water mixtures was calibrated against published data for the partial pressure of B, pB.2 (Water cannot be detected with a flame ionization detector.) For B-D-water mixtures, a peak due to D was negligibly small, and hence the vapor phase practically consisted of B and W. Thus, pB could be determined from the above calibration curve. The phase behavior of B-water and B-D-water mixtures was studied by light absorption measurements. Two cuvettes, containing a given solution, were placed in a Beckman DK-2A ratio recording photometer. One was thermostated at 25 °C, while the other was heated at 0.5 °C/min. A strong sudden increase in the differential light absorption at 460 nm was ascribed to phase separation (cloud point formation) in the heated cuvette. The temperature was measured to the nearest 0.1 °C inside the closed cuvettes. Results Excess partial molar enthalpies of B, HBE, in water and mixed W-D solvents are plotted against the mole fraction of B in Figure 1. Each curve represents HBE for solutions with a constant D:H2O mole ratio (r ) nD/nW). It appears that, for a given xB, HBE becomes less negative with increasing D content in the solvent (increasing r). Also, the sigmoidal shape of the curve, characteristic for the binary B-W system,1 is gradually lost with increasing r. Partial pressures of B (pB) for four r values are plotted against the mole fraction of B in the liquid phase in Figure 2. The excess chemical potential of B is calculated from these data as
µEB ) RT ln(pB/xBpB°)
(6)
where pB° is the vapor pressure of pure B at 25 °C. (A value of 0.859 Torr is used.2) Values of µBE calculated from eq 6 are plotted as a function of xB in Figure 3. The phase behavior of B-D-water mixtures is illustrated in Figure 4. The data points indicate observed cloud points for solutions with three different r values. The lines are fitted by the eye and show the phase boundaries between full miscibility (below the lines) and the region with two liquid phases.
2-Butoxyethanol-DMSO-H2O
Figure 1. Excess partial molar enthalpy of 2-butoxyethanol (B) in water and mixed water-dimethyl sulfoxide solvents plotted as a function of xB. The r values (r ) nD/nW) are as follows: circles, 0 (from ref 7); squares, 0.0128; up triangles, 0.0256; down triangles, 0.0576; diamonds, 0.0988; hexagons, 0.251; dotted circles, 0.489; dotted squares, 1.947. Dotted triangles represents data from titration of pure B into pure D.
J. Phys. Chem., Vol. 100, No. 1, 1996 435
Figure 4. Phase behavior of the water-rich composition range of the B-D-water system. The points designate observed cloud points for solutions with r values of 0, 0.0048, and 0.0121 for circles, squares, and triangles, respectively. The mixtures consist of two liquid phases above the fitted lines and are fully miscible below the lines.
Figure 2. Results from the gas chromatographic measurements. The figure illustrates partial pressures of 2-butoxyethanol, pB, as a function of xB determined from smooth curves through the data points. The r values are as follows: circles, 0 (data from ref 2); squares, 0.0162; up triangles, 0.0545; down triangles, 0.0988.
Figure 5. Examples of plots used for obtaining connected (HBE,xB,xD) and (µBE,xB,xD) data sets. The thermodynamic function (circles, left ordinate) and xD (dotted line, right ordinate) are plotted against xB. r values are 0.0545 for the µBE plot (A) and 0.0576 for HBE (B).
Figure 3. Excess chemical potential of 2-butoxyethanol, µBE, as a function of xB. r values are as follows: circles, 0 (data from ref 2); squares, 0.0162; up triangles, 0.0545; down triangles, 0.0988. The error decreases from about twice the height of the symbols at the lower B concentrations to approximately the height of the symbols at the higher B concentrations.
Discussion The main purpose of this study is to discuss molecular interactions based on estimated values of the interaction functions, eqs 4 and 5. To do so, for a series with a fixed value of r, we calculated the mole fraction of D, xD ) (1 + xB)/(1 + 1/r), and plotted this together with either the experimental values
of HBE as exemplified in Figure 5A or those of µBE as in Figure 5B. From graphs like these, values of HBE (or µBE) were read off smooth curves at fixed values of xB and plotted as a function of xD. This latter type of graphs (not shown), which illustrate connected data sets of the type (HBE,xB,xD) or (µBE,xB,xD), was used for calculating the partial excess entropy of B, SBE (TSBE ) HBE - µBE), as well as for the enumeration of the derivatives of eqs 4 and 5 (i.e., estimation of (∂FB/∂xD)xB and (∂FB/∂xB)xD). B-B Interactions. Figure 6 illustrates the B-B interaction functions (µBB and HBB) as a function of xB for selected values of xD. It appears that µBB is invariably negative, i.e., that B-B interactions are net attractive, more so the lower the B concentration. Since the µBB curves are practically superimposed, D appears to have little effect on the free energy of B-B interactions. This observation is in contrast to previous
436 J. Phys. Chem., Vol. 100, No. 1, 1996
Figure 6. B-B interaction functions (right-hand side of eq 4) plotted as a function of xB for selected values of xD. Open symbols represent µBB for constant xD values of 0 (circles), 0.04 (squares), and 0.08 (triangles). These three curves are practically superimposed. Filled and filled cross-hair symbols represent HBB at different D concentrations. The values of xD changes from 0 to 0.08 at 0.02 increments in the sequence circles, squares, up triangles, down triangles, and diamonds. For higher D concentrations, HBB is indicated by hexagons (xD ) 0.16), cross-hair circles (xD ) 0.21), and cross-hair triangles (xD ) 0.29).
suggestions,15 stating that D tends to “break hydrophobic bonds” between molecules with large nonpolar moieties such as B. The limited effect of D on the free energy of B-B interactions turned out to be a result of compensation between enthalpy and entropy effects as discussed below. A more detailed picture of B-B interactions emerges from the enthalpy data (Figures 1 and 6). Large negative values of HBE at low xB (Figure 1) have been ascribed to enhanced hydrogen bonding between water molecules solvating the nonpolar parts of the solute1,2 (so-called hydrophobic solvation). Judging from Figure 1, addition of D strongly reduces the propensity of the solvent to form enthalpy low solvation structures around B. Similar observations have previously been reported.15 This behavior might be expected simply from the reduced water concentration of the solvent, but the rapid increase in HBE with increasing r (at low xB) may also depend on enhancement of water-water hydrogen bonding already introduced by D (see refs 7 and 16 and references therein) in the solvent. The enthalpy interaction function (Figure 6) shows that B-B interactions are unfavorable (repulsive) in terms of enthalpy at low xB (see refs 1 and 3 for a detailed discussion). Interestingly, the effect of D for the lowest B concentrations is negligible up to xD ≈ 0.1. (The curves in Figure 6 are practically superimposed for xB < 0.005.) This suggests that while the enhancement of water-water hydrogen bonds introduced by the first B molecules decreases significantly upon addition of D to the solvent, the unfavorable enthalpy of B-B interaction remains unchanged. The sharp maximum in the enthalpy interaction function (around xB ) 0.018) for the binary B-W system has been ascribed to a change in “mixing scheme” from region I to II.2,3 In region I (xB < 0.018) thermodynamic properties of the solution are dominated by strong, long-range (hydrophobic) interactions between B molecules. In region II, B-B interactions become much weaker, probably because incoming B molecules now settle into B-rich clusters in which they experience interactions comparable to those in the pure B liquid. It appears from Figure 6 that the large unfavorable enthalpy of B-B interaction observed around the mixing scheme boundary is strongly reduced by D but that the B concentration where this anomaly occurs is only marginally influenced. The fact that anomalies at xB ) 0.018 at 25 °C change shape gradually from a peak at xD ) 0 to a step at xD ) 0.08 has been
Westh and Koga
Figure 7. B-D interaction functions (right-hand side of eq 5) in terms of free energy (µBD, open symbols) and enthalpy (HBD, filled symbols) as a function of xB for selected values of xD. The xD values for µBD are (top to bottom) 0, 0.04, and 0.08. For HBD the curves (top to bottom) represent xD values from 0 to 0.08 at 0.02 increments.
related to the detailed way in which the solute enhances the hydrogen bond probability and its strength in region I.3 Briefly, if the first solute molecule enhances the hydrogen bond strength of solvent H2O in its vicinity and if such an influence tails off to an infinite range, the second solute will influence the hydrogen bond network to a lesser extent than the first. Hence, the enthalpic interaction HBB ) δHBE/δnB > 0. However, even within the group of solutes that exhibit positive enthalpy interaction functions, the detailed manner by which the hydrogen bond network is perturbed may vary. The rate of change in the part of the hydrogen bond network influenced, |δHBE|, may (A) increase, (B) stay constant, or (C) decrease with the solute concentration. Case A corresponds to the binary B-water system, in which HBB ) δHBE/δnB, increases sharply up to the mixing scheme boundary, as shown in Figure 6 for xD ) 0. Previous studies on binary aqueous solutions of isobutyric acid and 2-butanone showed that HBB remained constant (i.e., that δHBE stayed constant) up to the boundary and then decreased sharply at the change to the second mixing scheme.6,17 This results in a step-type anomaly and corresponds to case B. For binary D-water mixtures, the D-D interaction, HDD, decreased gradually to the mixing scheme boundary at xD ) 0.28. Thereupon, a much sharper decrease in HDD associated with the transition to the second mixing scheme was observed.7 This corresponds to case C. Such differences between cases A, B, and C may be sought in competing effects of hydrophobic and hydrophilic moieties of each particular solute species. The behavior of HBB shown in Figure 6 indicates a subtle interplay between effects of B and D on the hydrogen bond network of water. Thus, the effect of D is such that the increase in δHBE as xB increases is diminished eventually to the case B, and yet the location (in terms of xB) of the mixing scheme boundary does not seem to be affected by D. B-D Interactions. B-D interactions according to eq 5 (µBD and HBD) are illustrated in Figure 7. Values of the free energy interaction function are close to zero (2 orders of magnitude smaller that similar values for B-B interactions). This suggests that B-D interactions contribute only marginally to the excess free energy of the solution or, in other words, that interactions between B and D in an aqueous environment are characterized by almost completely compensating enthalpy and entropy effects. The enthalpy of interaction between B and D, HBD, is relatively large and positive (about half the magnitude of the B-B interaction function) for the lowest B concentrations (Figure 7). Hence, B-D interactions appear to be characterized
2-Butoxyethanol-DMSO-H2O
J. Phys. Chem., Vol. 100, No. 1, 1996 437
Figure 8. Excess partial molar enthalpy of B at infinite dilution, HBE(xBf0), as a function of the D mole fraction. The inset shows the slope of the curve estimated graphically.
Figure 9. Entropic contribution, TSBE, to the excess chemical potential of B calculated from interpolated values of the enthalpy and free energy data. The mole fraction of D is (bottom to top) 0, 0.02, 0.04, 0.06, 0.08, and 0.10.
by a sizable enthalpic repulsion which is compensated by entropic attraction. Compensation of numerically large enthalpy/entropy effects is characteristic for nonpolar groups in dilute aqueous solutions, and the thermodynamics of B-D interactions, at low concentration of B, therefore appears to be governed by interactions between the alkyl moieties. As mentioned in the Introduction, there has been some discussion of hydrophobic binding of D to proteins. If B is considered a model for a nonpolar region on a protein surface, the results suggest that interactions with D are strong in the sense that sizable enthalpy and entropy effects are involved but that “binding” is weak since these two effects compensate almost completely. The B-D enthalpy interaction functions in Figure 7 drop steeply at about the mixing scheme boundary for the B-W system. Hence, the unfavorable enthalpy of B-D interactions, like in the case of B-B interactions, seem to be a characteristic of mixing scheme I. The similarity of the curves in Figure 7 also suggests that the nature of the mixing scheme occurring in region I is retained in the presence of D at least up to xD ≈ 0.08. As discussed above, studies of binary (D-H2O and B-H2O) mixtures suggest that effects of D (up to xD ) 0.28) on the hydrogen bond network of water are qualitatively the same as those of B (up to xB ) 0.018). This view is supported by the present data, since, beyond xD ≈ 0.28, HBB is close to zero (Figure 6), as it is characteristic of mixing scheme II. When enough D molecules (xD > 0.28) influence the hydrogen bond network of H2O, then, even for B concentrations of xB < 0.018, where in the binary B-H2O system mixing scheme I prevails, the mixing now shows the characteristics of type II, e.g., almost zero HBB. These competitive or cooperative effects of B and D on water may be even more clearly illustrated in Figure 8 in which values of HBE extrapolated to zero B concentration (i.e., HBE at infinite dilution of B, HBE(xBf0)) are plotted against xD. It appears that HBE(xBf0) increases with xD up until xD ≈0.28, where it runs through a maximum. The slope of HBE(xBf0) vs xD is estimated from the smooth curve and plotted in the inset of Figure 8. It is noteworthy that the curve in the inset, which is HBD, has exactly the same course (case C) as the HDD function of the binary D-water system.7 Hence, upon addition of a very small amount of B to mixed D-water solvents, interactions between B and D as illustrated by Figure 8 merely reflect the mixing scheme status of the binary solvent. It seems clear that as long as either one kind of solute reaches concentrations beyond the mixing scheme boundary, the B-D-
water system behaves as mixing scheme II, characterized by very small values of HBB and HBD. Within mixing scheme I, i.e., for xD < 0.28 and xB < 0.018, solute-solute interactions may change from case A, through case B, and to case C depending on D concentration as shown in Figures 6-8. This must be caused by a competition between D and B on modifying the (percolated) hydrogen bond network of water. The data in Figure 7 for xD ) 0 and Figure 8 show that anomalies in heterogeneous (B-D) interactions at infinite dilution of D (Figure 7) and B (Figure 8) occur concurrently with solutesolute anomalies in the binary aqueous mixtures. This supports the view that the pronounced interactions in mixing scheme I are governed by the common component, solvent water, and it is in accordance with the suggestion that the change from mixing scheme I to II involves loss of percolation in the hydrogenbonded network of water.3 Namely, when the connectivity of the hydrogen-bonded network falls below the critical level of percolation, long-range (hydrophobic) interactions, between any solute species, are lost. Mixing Scheme II. As mentioned in the Introduction, binary B-water mixtures in mixing scheme II are characterized by clusters rich in each component premonitory of the phase separation at the lower critical solution temperature (LCST).1,2 The phase behavior illustrated in Figure 4 shows that the presence of D increases LCST (the minima of the curves) but that it has little effect on the B concentration where LCST is observed. For a phase separation with LCST to occur, the solute-solute interaction must be attractive in terms of entropy and repulsive enthalpy-wise, which is indeed the case for B-water. In a most simplistic view,3 if the entropic attraction weakens, the LCST may be expected to rise. As shown in Figure 9, this is in accordance with the observed behavior since the slope of the SBE vs xB plots decreases as xD increases. The actual mechanism by which the entropic B-B interaction decreases within mixing scheme II is yet to be elucidated. Closing Remarks. Information on molecular interactions obtained from thermodynamic measurements on binary aqueous mixtures has contributed significantly to the identification of driving forces in important processes such as folding, binding, and assembly of biopolymers. Application of ternary mixtures as model systems for processes of biological importance appears attractive since it offers a possibility to evaluate interactions, in aqueous solutions, between two different solute species, as well as the effect of one species on mutual interactions between the molecules of another. Hence, studies on more realistic model system can be performed. We suggest that the methodology presented in this study might be expedient for the
438 J. Phys. Chem., Vol. 100, No. 1, 1996 treatment of experimental data obtained from such model systems. Also, it might be applied directly to solutions of macromolecules. The present findings that B and D competitively or cooperative influence the percolated hydrogen bond network of H2O may have important implications for the understanding of the role of water in aqueous solutions of biopolymers. The effect of cosolvents, for example, on the conformational stability of proteins may be a manifestation of the competitive or cooperative influence of the two solute species on solvent H2O. Acknowledgment. This study was supported by the National Research Council of Canada, a NATO collaborative grant (Y.K.), and a postdoctoral fellowship (P.W.) from the Danish National Research Council. References and Notes (1) Siu, W.; Koga, Y. Can. J. Chem. 1989, 67, 671. (2) Koga, Y. J. Phys. Chem. 1991, 95, 4119. (3) (a) Koga, Y. J. Crystallogr. Soc. Jpn. 1995, 37, 157. (b) Koga, Y. J. Phys. Chem., submitted for publication. (4) Shinoda, K. J. Phys. Chem. 1977, 81, 1300.
Westh and Koga (5) Tanford, C. The Hydrophobic Effect: Formation of Micelles and Biological Membranes; John Wiley and Sons: New York, 1979. (6) Wong, T. Y. H.; Wong-Moon, K. C.; Beach, L. J.; Chuang, Y. F.; Koga, Y. J. Phys. Chem. 1992, 96, 10025. (7) Lai, J. T. W.; Lau, F. W.; Robb, D.; Westh, P.; Nielsen, G.; Trandum, C.; Hvidt, Aa.; Koga, Y. J. Solution Chem. 1995, 24, 89. (8) (a) Koga, Y.; Kristiansen, J.; Hvidt, Aa. J. Chem. Thermodyn. 1993, 25, 51. (b) Koga, Y. J. Phys. Chem. 1992, 96, 10466. (9) Fujita, Y.; Izumiguchi, S.; Noda, Y. Int. J. Peptide Protein Res. 1982, 19, 25. (10) Hamaguchi, K. J. Biochemistry 1964, 56, 441. (11) (a) Velicelebi, G.; Sturtevant, J. M. Biochemistry 1979, 18, 1180. (b) Gerlsma, S.; Stuur, E. R. Int. J. Peptide Protein Res. 1972, 4, 377. (12) (a) Lehmann, M. S.; Mason, S. A.; McIntyre, G. J. Biochemistry 1985, 24, 5862. (b) Lehmann, M. S.; Stansfield, R. F. D. Biochemistry 1989, 28, 7028. (13) Arakawa, T.; Carpenter, J. F.; Kita, Y. A.; Crowe, J. H. Cryobiology 1990, 27, 401. (14) (a) Koga, Y. Can. J. Chem. 1986, 64, 204. (b) Koga, Y. J. Chem. Thermodyn. 1987, 19, 571. (c) Koga, Y. Can. J. Chem. 1988, 66, 1187. (15) Mastroianni, M. J.; Pikal, M. J.; Lindenbaum, S. J. Phys. Chem. 1972, 76, 3050. (16) Westh, P. J. Phys. Chem. 1994, 98, 3222. (17) (a) Siu, W. W. Y.; Wong, T. Y. H.; Chao, L. C. F.; Koga, K. Can. J. Chem. 1991, 69, 1065.
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