An understanding of the solvating effect helps rationalize the relation between solvation and acidity of the compound. In addition, this also explains the relative stability of conjugate bases for compounds with different pKa values. This lesson details, in-depth, the principle of solvating effects. The strength of an acid and the stability of its corresponding conjugate base are determined using pKa values. This observed relationship is a consequence of solvation, which is the interaction between a dissolved ion and solvent molecules. During this process, the solvent molecules surround the ions and stabilize them.
Solvation of dissolved ions can be classified into three types: (i) donor interaction, (ii) charge–dipole interaction, and (iii) hydrogen-bonding interaction. In the donor interaction, a solvent donates its unshared electron pairs to the dissolved ion. The solvent acts as a Lewis base, and the ion acts as a Lewis acid. In the second type, charge–dipole interactions are observed in polar solvents, where their dipole moments can interact with the charged ions. This involves rearranging the positive partial charge on the solvent molecules to align with the negative charge of the ions, thus stabilizing the ions. For instance, as noted in the solvation of ethanol, the ethoxide anion, which is the conjugate base, is solvated by the positive center of the solvent’s dipole that stabilizes it effectively. Lastly, when the ions are stabilized by hydrogen bonding between the solvent molecules and the dissolved ions, the interaction is called a hydrogen-bonding interaction.
The interactions between the dissolved ions and solvent molecules influence their stability, which is directly proportional to the strength of the acidity. Accordingly, the stability of such ions increases with a larger number of interactions when they are surrounded by more solvent molecules. Therefore, during solvation, the steric hindrance from bulky substituents on the molecule plays an important role. Compounds with less bulky groups are sterically unhindered, allowing for more interaction with solvent molecules.
In contrast, the compounds that possess bulky groups have steric hindrance and are consequently poorly solvated. As a result, the sterically unhindered ion demonstrates more stability, making its corresponding acid stronger. This is demonstrated with the comparison of acidity of ethanol, isopropanol, and tert-butanol. With the increasing size of substituents, the corresponding conjugate base of each of these compounds has more steric hindrance. Hence, it is less solvated. As a result, isopropanol is a weaker acid (pKa=17.10) than ethanol (pKa=16.00), and tert-butanol (pKa=19.20) is a weaker acid than isopropanol (pKa=17.10). In summation, the steric hindrance of the conjugate base anions defines the degree of solvation. Low solvation leads to instability of the dissolved ion that makes the corresponding acid weak.
