Hydrophobic Interactions

Fats, hydrocarbons, and other materials whose molecules consist largely of nonpolar groups have a low solubility in water and a high solubility in nonpolar solvents. Similarly, the long alkyl groups of fatty acid esters aggregate within membranes and nonpolar side chains of proteins are often packed together in the centers of protein molecules. Because it is as if the nonpolar groups “fear” water, this is known as the hydrophobic effect. The terms hydrophobic forces, hydrophobic interactions, and hydrophobic bonding have also been used. However, the latter term can be misleading because the hydrophobic effect arises not out of any special attraction between nonpolar groups but primarily from the strong internal cohesion of the hydrogen-bonded water structure.

–CH2–(water) --→–CH2–(hexane)
ΔG° = –3.8 kJ/mol

This equation is a quantitative statement of the fact that the CH2 group prefers to be in a nonpolar environment than to be surrounded by water. A similar Gibbs energy change would be expected to accompany the bringing together of a methylene unit from a small molecule and a hydrophobic surface on a protein molecule. However, in the latter case the accompanying What causes the decrease in Gibbs energy when nonpolar groups associate in water? Jencks60 suggested that we think of the transfer of a nonpolar molecule from a nonpolar solvent into water in two steps: (1) Create a cavity in the water of about the right size to accommodate the molecule. Since many hydrogen bonds will be broken, the Gibbs energy of cavity formation will be high. It will be principally an enthalpy (ΔH) effect. (2) Allow the water molecules in the solvent to make changes in their orientations to accommodate the nonpolar molecule that has been placed in the cavity. The water molecules can move to give good van der Waals contacts and also reorient themselves to give the maximum number of hydrogen bonds. Since hydrogen bonds can be formed in many different ways in water, there may be as many or even more hydrogen bonds after the reorientation than before. This will be true especially at low temperature where most water exists as large icelike clusters. For dissolved hydrocarbons, the enthalpy of formation of the new hydrogen bonds often almost exactly balances the enthalpy of creation of the cavity initially so that ΔH for the overall process (transfer from inert solvent into water) is small. For the opposite transfer ΔH° is usually a small positive number for aliphatic hydrocarbons and is nearly zero for aromatic hydrocarbons.
Since ΔG° = ΔH°– TΔS°, it follows that the negative value of ΔG° for hydrophobic interactions must result from a positive entropy change, which may arise from the restricted mobility of water molecules that surround dissolved hydrophobic groups. When two hydrophobic groups come together to form a “hydrophobic bond,” water molecules are freed from the structured region around the hydrophobic surfaces and the entropy increases. The ΔS°. Attempts have been made to relate this value directly to the increased number of orientations possible for a water molecule when it is freed from the structured region. However, interpretation of the hydrophobic effect is complex and controversial.
The formation constant Kf for hydrophobic associations often increases with increasing temperature. This is in contrast to the behavior of Kf for many association reactions that involve polar molecules and for which ΔH° is often strongly negative (heat is released). An example of the latter is the protonation of ammonia in an aqueous solution

NH3 + H+ → NH4+
ΔH° = –52.5 kJ/mol

Since R1nKf = –ΔG°/T = –ΔH°/T + ΔS°, Kf decreases with increasing temperature if ΔH° is negative. Because for a hydrophobic interaction with a positive value of ΔH° Kf increases with increasing temperature, an increase in stability at higher temperatures is sometimes used as a criterion for hydrophobic bonding. However, this criterion does not always hold. For example, base stacking interactions in polynucleotides, whose strength does not increase with increasing temperature, are still thought to be hydrophobic.
The water molecules that are in immediate contact with dissolved nonpolar groups are partially oriented. They form a cagelike structure around each hydrophobic group. When particles surrounded by such hydration layers are 1–2 nm apart, they sometimes experience either a fairly strong repulsion or an enhanced attraction caused by these hydration layers Direct experimental measurements have shown that these effects extend to distances of 10 nm and can account for the previously mentioned long-range van der Waals forces.
Various efforts have been made to develop scales of hydrophobicity that can be used to predict the probability of finding a given amino acid side chain buried within a protein or in a surface facing water. A new approach has been provided by the study of mutant proteins. For example, deletion of a single –CH2– group from an interior hydrophobic region of a protein was observed to decrease the stability of the protein by 4.6 kJ/mol.