Solvating Effects

The strengths of acids and the relative stabilities of their corresponding conjugate bases can be estimated from the pK_a values of the acids. Lower pK_a values indicate stronger acids and stable corresponding conjugate bases.

This can be explained based on the ability of polar acid molecules to stabilize their conjugate base anions in their solutions through solvation.

For example, consider the solution of ethanol and its conjugate base ethoxide ion. During solvation, the positive end of the solvent’s dipole interacts with the ethoxide anion.

This charge-dipole attraction heavily solvates the sterically unhindered ethoxide ion and effectively stabilizes it. Consequently, the deprotonation reaction of ethanol to ethoxide ion is favored and ethanol has a pK_a value of 15.5.

Next, consider the solution of isopropanol and its conjugate base isopropoxide ion.

Compared to the ethoxide ion, the isopropoxide ion has an additional methyl group at the alpha carbon, making it sterically more hindered.

Since fewer solvent molecules interact with the moderately hindered ion, the isopropoxide anion is less stable than the ethoxide ion, and thus isopropanol is a weaker acid than ethanol.

Lastly, examine the solution of tert-butanol and its conjugate base tert-butoxide ion.

Compared to the isopropoxide ion, the tert-butoxide ion has three bulky methyl groups, making it poorly accessible for the solvent to interact with.

The poor conjugate base stabilization makes the tert-butoxide ion a less stable base than the isopropoxide ion. Hence, tert-butanol is a weaker acid than isopropanol.

To summarize, an increase in the steric hindrance of the conjugate base anions decreases their degree of solvation, which invariably makes the ions less stable, and their corresponding acids weaker.