The hydrophobic interaction
المؤلف:
Peter Atkins، Julio de Paula
المصدر:
ATKINS PHYSICAL CHEMISTRY
الجزء والصفحة:
ص635-636
2025-12-16
45
The hydrophobic interaction
Nonpolar molecules do dissolve slightly in polar solvents, but strong interactions between solute and solvent are not possible and as a result it is found that each individual solute molecule is surrounded by a solvent cage (Fig. 18.9). To understand the consequences of this effect, consider the thermodynamics of transfer of a nonpolar hydrocarbon solute from a nonpolar solvent to water, a polar solvent. Experiments indicate that the process is endergonic (∆transfer G > 0), as expected on the basis of the increase in polarity of the solvent, but exothermic (∆transfer H < 0). Therefore, it is a large decrease in the entropy of the system (∆transfer S < 0) that accounts for the positive Gibbs energy of transfer. For example, the process
CH4(in CCl4) → CH4(aq)
has ∆transfer G =+12 kJ mol−1, ∆transfer H =−10 kJ mol−1, and ∆transfer S =−75 J K−1 mol−1 at 298 K. Substances characterized by a positive Gibbs energy of transfer from a nonpolar to a polar solvent are called hydrophobic. It is possible to quantify the hydrophobicity of a small molecular group R by defining the hydrophobicity constant, π, as
π=log
where S is the ratio of the molar solubility of the compound R-A in octanol, a non-polar solvent, to that in water, and S0 is the ratio of the molar solubility of the com pound H-A in octanol to that in water. Therefore, positive values of π indicate hydrophobicity and negative values of π indicate hydrophilicity, the thermodynamic preference for water as a solvent. It is observed experimentally that the π values of most groups do not depend on the nature of A. However, measurements do suggest group additivity of π values. For example, π for R = CH3, CH2CH3, (CH2)2CH3, (CH2)3CH3, and (CH2)4CH3 is, respectively, 0.5, 1.0, 1.5, 2.0, and 2.5 and we conclude that acyclic saturated hydrocarbons become more hydrophobic as the carbon chain length increases. This trend can be rationalized by ∆transfer H becoming more positive and ∆transfer S more negative as the number of carbon atoms in the chain increases.
At the molecular level, formation of a solvent cage around a hydrophobic molecule involves the formation of new hydrogen bonds among solvent molecules. This is an exothermic process and accounts for the negative values of ∆transfer H. On the other hand, the increase in order associated with formation of a very large number of small solvent cages decreases the entropy of the system and accounts for the negative values of ∆transfer S. However, when many solute molecules cluster together, fewer (albeit larger) cages are required and more solvent molecules are free to move. The net effect of formation of large clusters of hydrophobic molecules is then a decrease in the organization of the solvent and therefore a net increase in entropy of the system. This increase in entropy of the solvent is large enough to render spontaneous the association of hydrophobic molecules in a polar solvent. The increase in entropy that results from fewer structural demands on the solvent placed by the clustering of nonpolar molecules is the origin of the hydrophobic interaction, which tends to stabilize groupings of hydrophobic groups in micelles and biopolymers (Chapter 19). The hydrophobic interaction is an example of an ordering process that is stabilized by a tendency toward greater disorder of the solvent.
Fig. 18.9 When a hydrocarbon molecule is surrounded by water, the H2O molecules form a clathrate cage. As a result of this acquisition of structure, the entropy of the water decreases, so the dispersal of the hydrocarbon into the water is entropy opposed; its coalescence is entropy favoured.
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