Prediction of solubility using the regular solution theory usually fails when the solute and solvent are polar in character. The dipole-dipole, dipole-induced-dipole, charge-transfer, and hydrogen bonding interactions that exist between solute and solvent molecules may reduce the free energy of the solution, and increase the solubility. In these solutions, the activity coefficient may be less than one, a fact that cannot be explained using regular solution theory. The range of dipole-dipole, dipole-induced-dipole, and hydrogen bonding interactions in polar solutions may also lead to molecular orientation, which would tend to decrease the entropy of mixing. Clearly the nature of the forces involved in solution, and the influence of the forces on solubility, are important in order to arrive at a better understanding of solubility behavior.
Coulombic interaction is a valence force between counterions, and in extreme situations a cation-anion pair might form a strong ion-dipole interaction in solution. Such interactions would tend to be major for ionic substances dissolved in non-polar solvent systems, but less so in polar solvents where the forces of solvation serve to disrupt ion pairs into individual solvated ions. These trends provide an insight into why salts tend to be soluble in polar solvents, but not in non-polar solvents.
Van der Waals forces represent important intermolecular interactions between nonelectrolyte substances, and can be categorized into dipoledipole, dipole-induced-dipole, and induced-dipole-induced-dipole forces. Polar molecules, by definition, will have a permanent dipole moment, and will interact with the oppositely charged portions or other molecules having permanent dipole moments. The dipole-dipole interaction is known as the orientation effect, or as the Keesom force.
Molecules having delocalized electron systems or large molar volumes often are characterized by high degrees of polarizability. Their interaction with polar molecules can induce shifts in electron density that result in the transient presence of induced dipole moments, and the charged portions of the induced dipoles can form an attractive interaction with the neighboring polar molecule. This type of interaction is termed the dipole-induced-dipole force, the induction effect, or the Debye force. It is found that Keesom and Debye forces provide efficient molecular packing in crystals, accounting for the high stability, low thermodynamic activity, and the high melting point of many organic crystals. These attractive effects may yield substantial lattice energies for such crystals, and therefore tend to reduce their solubility in potential solvents.
All molecules, whether polar or nonpolar, are also attracted to each other by induced-dipole-induced-dipole interactions, which are known as dispersion forces, or London forces.
Nonpolar molecules can only interact by dispersion forces, while the interactions of polar molecules are often dominated by the Keesom forces. However, under certain circumstances it is still possible that dispersion forces might predominate over the other forces, even for polar molecules such as HCl. The Debye forces are often stronger than the London forces for highly polar molecules, and would predominate over Keesom forces for weakly polar molecules. Debye forces are selective, and important in explaining why certain nonpolar but polarizible molecules can still be soluble in polar solvents (Krishnan and Fredman, 1971).
Hydrogen atoms are small in size, and would be positively polarized in molecules where it is bound adjacent to an electronegative atom, A. Should another strongly electronegative atom, B, approach the hydrogen atom at a short distance, a strong interaction may develop that is termed a hydrogen bond. The strongest hydrogen bonds are formed when the electronegative atoms involved are fluorine, oxygen, or nitrogen, although chlorine and sulfur are known to form weak hydrogen bonds in some molecules.
The strengths ofhydrogen bonds are similar in magnitude to those ofvan der Waals forces, but is also directional in the manner of a covalent bond. Hydrogen bonding tends to stabilize molecular pairs and reduces the enthalpy, but also tends to orient the molecules involved and decrease the entropy. The effect of hydrogen bonding on solubility is complicated, and the analysis must proceed on a case-by-case basis. Extensive intermolecular hydrogen bonding in a crystal would tend to decrease the free energy, with this stabilization effect reducing the activity of the solute, and tending to reduce the solubility. However, the hydrogen bonds formed between solute and solvent molecules would tend to reduce the activity coefficient, and this effect would lead to increased solubility.
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