5.7:
Solvating Effects
The strengths of acids and the relative stabilities of their corresponding conjugate bases can be estimated from the pKa values of the acids. Lower pKa 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 pKa 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.
5.7:
Solvating Effects
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.
Suggested Reading
- Brown, W.H., & Iverson, B.L., & Anslyn, V.E., & Foote S.C. (2014). Organic Chemistry. Mason, Ohio: Cengage Learning, 118-123.
- Solomons, G., & Fryhle, C. & Snyder, S. (2015). Organic Chemistry. New Jersey, NJ: Wiley, 192-194.
- Loudon, M., & Parise, J. (2016). Organic Chemistry. New York, NY: Macmillan Publishers, 230-234.
- Klein, D. (2017). Organic Chemistry. New Jersey, NJ: Wiley, 183-188.
- Clayden, J., & Greeves, N., & Warren, S. (2012). Organic Chemistry. Oxford: Oxford University Press, 300-305.