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Hydrophobic Effect Secondary article Judith Herzfeld, Brandeis University, Waltham, Massachusetts, USA Article Contents Donald J Olbris, Brandeis University, Waltham, Massachusetts, USA . Hydrophobicity . Thermodynamics of Transfer The hydrophobic effect refers to the relatively poor solubility of nonpolar substances in . Thermodynamics of Solvation water. The effect is seen in the organization of biomolecules such that nonpolar portions . Thermodynamics of Hydration are largely sequestered from the aqueous environment. The origin of the effect lies in the . Physical Features of Water response of the three-dimensional hydrogen bonding network of water to different types . Mechanism of Hydration of solutes. Hydrophobic Interactions . Protein Folding Hydrophobicity Hydrophobic (literally, water-fearing) substances are ones solvent will be enthalpically unfavourable owing to the that are poorly soluble in water compared with their disruption of solvent–solvent dipole interactions. solubility in nonpolar solvents. For example, the solubility As it turns out, water is not an ordinary polar solvent, of ethane at 1 atm and 258C is 0.21 mol L 2 1 in carbon and the foregoing expectations are not generally met for tetrachloride (Wilhelm and Battino, 1973), but only transfer of nonpolar solutes from a nonpolar solvent to 1.9 Â 10 2 3 mol L 2 1 in water (Ben-Naim and Marcus, water. Figure 1 shows the thermodynamic results obtained 1984b). In general, ethane equilibrates between nonpolar for transfer of ethane from carbon tetrachloride to water solvents and water at a molar concentration ratio of about when the molar concentrations are the same in the two 100:1, and the partitioning of larger nonpolar solutes phases (so that dilution effects make no contribution to the disfavours water still more strongly. As a result, nonpolar entropy). The term ‘local’ is used when this molar standard solvents are very effective at extracting nonpolar solutes is applied because the thermodynamic quantities then from water. correspond to the hypothetical process of transferring the The biological interest in hydrophobicity stems from its solute from a fixed position in one solvent to a fixed role in the intramolecular and intermolecular associations position in the other (Ben-Naim and Marcus, 1984a). At of biomolecules. Lipids, with polar head groups and equilibrium the Gibbs free energy change (DG) is zero, and nonpolar tails, assemble into bilayers, with the head groups the equilibrium concentrations in the water and the organic in contact with water, while the tails are sequestered from phase obey the condition described by eqn [1]. water and solvate one another. Proteins, with both polar 0 5 DG 5 DG8 1 RT ln (cwater/corganic) [1] and nonpolar side-chains, fold with the nonpolar groups largely turned inward so they avoid water. Hydrophobic The standard free energy change is thus related to the water organic groups remaining on protein surfaces are often buried in partition coefficient Kd 5 (c /c ) by eqn [2]. supramolecular assembly, including insertion into mem- branes, self-assembly into filaments, and docking with DG8 52RT ln Kd [2] other molecules. This free energy change includes an enthalpic contribution (DH8) and an entropic contribution (DS8) in the combina- tion shown in eqn [3]. Thermodynamics of Transfer DG8 5 DH8 2 TDS8 [3] Often solubilities are rationalized in terms of the favour- These contributions can be teased apart in two ways. One is ability of interactions between solute and solvent, as to obtain DG8 for a range of temperatures and obtain DS8 compared with solvent–solvent interactions and solute– from the relation DS8 52@DG8/@T. More recently, the solute interactions. The maxim that ‘like dissolves like’ is development of calorimeters sensitive enough for studies of based on the idea that dipole interactions occur only very dilute solutions has made it possible to measure DH8 between polar molecules and these are lost when polar directly. molecules are mixed with nonpolar molecules. The Figure 1 shows that the transfer of ethane from carbon expectation then is that immiscibility occurs because the tetrachloride to water is generally unfavourable (DG840). enthalpic disadvantages of mixing prevail over the entropic However, it is enthalpically unfavourable (DH840) only advantages of mixing so that, overall, the free energy of at relatively high temperatures, while it is entropically mixing is unfavourable. Likewise, we expect that transfer unfavourable (DS850) in varying degrees throughout the of a nonpolar solute from a nonpolar solvent to a polar temperature range of liquid water. In fact, both DH8 and ENCYCLOPEDIA OF LIFE SCIENCES © 2002, John Wiley & Sons, Ltd. www.els.net 1 Hydrophobic Effect 30 6 ∆G°/RT 20 4 ∆G° ) –1 10 2 ∆ G ° / RT (kJ mol ° (unitless) G 0 0 ∆ , ° S ∆ T , ∆H° ° –10 –2 H ∆ –20 –4 T∆S° Th Ts –30 –6 –10 20 50 80 110 140 T (°C) Figure 1 The thermodynamics of transfer of ethane from carbon tetrachloride to water. Th and Ts represent the temperatures at which DH8 and DS8 of transfer are zero, respectively. At Ts, DG8 is at its maximum, and at Th, DG8/T is at its maximum and the partition coefficient Kd 5 exp ( 2 DG8/RT) is at its minimum. DS8 are strongly temperature dependent, indicating that Th, the hydrophobicity is driven entirely by the unfavour- transfer is accompanied by a large change in the constant able entropy change (DS850). pressure heat capacity DCp. Nevertheless, DG8 only varies weakly with temperature, indicating that the changes in DH8 and DS8 are largely compensating. This occurs because DCp appears in both @DH8/@T 5 DCp and Thermodynamics of Solvation @(TDS8)/@T 5 DS8 1 DCp, so this large contribution can- cels in the full temperature dependence of DG8, as shown in Interpretation of the thermodynamics of transfer between eqn [4]. two solvents is complicated because it involves changes in both of the solvents. Therefore it is desirable to consider @ÁG @ÁH @ TÁS dissolution in each phase separately from a solvent-free À ÀÁS 4 @T @T @T reference state. Figure 2 shows the relationships between these processes. Here o and w represent the solute dissolved The interesting feature in the temperature dependence of in organic solvent and water, respectively, while g, l and s DG8 is the curvature. This also reflects the large change in represent the gaseous, liquid and solid states of pure solute. the heat capacity through the relationship shown in eqn [5]. Thus the horizontal line represents the transfer between solvents discussed above, the vertical lines represent the @ÁH @ÁS @2ÁG ÁCp T ÀT 5 condensation–vaporization and freezing–melting transi- @T @T @T 2 tions of pure solute, and the oblique lines represent the The positive value of DCp upon transfer from organic to dissolution of solute. Clearly the thermodynamic functions aqueous solution produces a maximum in DG8 when @DG8/ for transfer can be parsed as differences between the @T 52DS8 5 0. This temperature is designated Ts. The thermodynamic functions for dissolution in the two positive value of DCp also generates a maximum in DG8/ solvents. Using the notation dX 5 Xb 2 Xa, these relation- RT 52ln Kd because when @(DG8/RT)/@T 52(DH8/ ships can be summarized as in eqn [7]. T2) 5 0, the second derivative must be less than zero, as 8 9 < Áw Ág Áw Áo = shown in eqn [6]. g o g À g Áw Áw Ál Áw Áo 7 o : l o l À l ; 2 Áw Ás Áw Áo @ ÁG=RT ÁH ÁC s o s À s À p < 0 6 @T 2 T 3 T 2 As a matter of convenience, the reference state for Thus, the partitioning is most hydrophobic (i.e., Kd is the experiments is chosen as the state in which the pure solute smallest) when DH8 5 0. At this temperature, designated is stable under ambient conditions. Translation of 2 Hydrophobic Effect g ethane, again using the molar concentration standard (for the gas phase as well as the solutions). It is clear that the various nonpolar solvents show similar behaviour and water is the anomalous solvent. A great deal of attention has therefore focused on understanding hydration (i.e. solvation in water). However, before we focus on water we l should note that the nonpolar solvents all show a favourable enthalpy of solvation that presumably reflects attractive dispersion forces, and an unfavourable entropy of solvation that presumably reflects the formation of cavities to accommodate the solute molecules. Overall, the solvation of ethane in nonpolar solvents is favourable s (DG8 5 DH8 2 TDS850) at 258C. The mystery in the case of hydrophobic hydration is: (1) the more favourable enthalpy and the much less favourable entropy, such that hydration overall is unfavourable (DG8 5 DH8 2 TDS840); and (2) the relatively large heat capacity, such that there are large temperature dependen- cies in the enthalpy and entropy which are largely owcompensating. Figure 2 Relationships between the transfer of solute between solvents and the dissolution of pure solute in these solvents. o and w represent the solute dissolved in organic solvent and water, respectively, while g, l and s Thermodynamics of Hydration represent the gaseous, liquid and solid states of pure solute. While the hydration of nonpolar molecules like ethane and thermodynamic functions to other reference states can propane is unfavourable, the hydration of polar molecules then be made by using the thermodynamic functions for is favourable, and a comparison is instructive. Table 2 the condensation–vaporization and freezing–melting shows thermodynamic data for two groups of molecules. transitions. It is also common to assume that nonpolar The members of each group are isoelectronic. The liquids are all pretty much alike. In this case, 0, which differences are in the replacement of a methyl group by a gives the expedient result that .
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