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Acids differ in their ability to donate H⁺. Stronger acids, such as HCl, react almost completely with water, whereas weaker acids, such as acetic acid (CH₃COOH), react only slightly. The exact strength of a given acid HA in water solution is described using the acidity constant ($K_a$) for the acid-dissociation equilibrium.

$\begin{array}{l}{\text{HA+H}}_2\text{O}⇄{\text{A}}^{−}{\text{+H}}_{\text{3}}{\text{O}}^{−}\\ K_a\underset{}{}\text{=}\underset{}{}\frac{{\text{[H}}_{\text{3}}{\text{O}}^{+}{\text{][A}}^{−}\text{]}}{\left[HA\right]}\end{array}$

Acid strengths are normally using $\text{p}{K}_a$values rather than$K_a$ values, where the $\text{p}{K}_a$is the negative common logarithm of$K_a$:

$\text{p}K\text{a=-log}{K}_a$

The stronger base has a larger $\text{p}{K}_a$, and the weaker base has a smaller $\text{p}{K}_a$value.

The stronger the acid (smaller ), the weaker its conjugate base. Since is a stronger acid than , methylamine is a stronger base than ammonia.

Note: Since methylamine has an electron-donating group (methyl). Therefore electron density on nitrogen atom increase. Hence it is a stronger base, whereas, in ammonia, an electron-donating group is absent. So it is a weak base when compared to methyl amine.