Amines can behave as Brønsted–Lowry bases by accepting a proton from the acid to form corresponding conjugate acids. Due to a lone pair of nonbonding electrons, aliphatic amines can also act as Lewis bases by forming a covalent bond with an electrophile.

To measure the basicity of amines, two conventions are generally used. The first defines K_b as the basicity constant for the deprotonation reaction of water by the amine, as presented in Figure 1. Conventionally, lower K_b indicates higher basicity of the amine. For example, ammonia and methylamine have basicity constants of 4.7 and 3.3, respectively.

Figure 1. The deprotonation of water by an amine.

The other convention correlates the basicity of the amines with the acidity of the corresponding conjugate acid. The higher the acidity constant, pK_a, of the conjugate acid, the higher the basicity of the amine from which the conjugate acid has been formed. For example, as shown in Figure 2, the conjugate acids of ammonia, methylamine, and ethylamine have pK_a values of 9.26, 10.64, and 10.75, respectively. This indicates the higher basicity of ethylamine, followed by methylamine and ammonia.

Figure 2. The pK_a values of various amines.

Aliphatic amines are more basic than ammonia due to the electron-releasing ability of the alkyl groups attached to the N atom. The alkyl group stabilizes the conjugate acid by dispersing its positive charge; as a result, enhancing the basicity of amines. It might be anticipated that the tertiary amines, having three alkyl groups, should be more basic than the primary amines having only one alkyl group. However, this is only true in the gas phase, where we see the acidity of the conjugate acids of ammonia, methylamine, dimethylamine, and trimethylamine in the order listed below.

[NH₄]⁺ > [MeNH₃]⁺ > [Me₂NH₂]⁺ > [Me₃NH]⁺

Therefore, ammonia is the weakest base, and trimethylamine is the strongest base in the gas phase. However, the basicity of amines in the aqueous phase is also influenced by the solvation of the corresponding conjugate acids in water. Therefore, despite having three alkyl groups attached to the N atom, conjugate acids of tertiary amines have only one H atom to be donated for intermolecular H bonding with water. In contrast, conjugate acids of primary amines have three H atoms for intermolecular hydrogen bonding with water. Therefore, the pK_a values of the conjugate acids of amines in an aqueous solution run as shown in Figure 3.

Figure 3. The pK_a values of conjugate acids of amines.

The conjugate acids of secondary amines enjoy an optimal balance of electron-releasing ability of alkyl groups and solvation capability with water.