Acid–base equilibrium of methylamine and water:

pK_b

It is a measure of basic strength. If methylamine is dissolved in water, equilibrium is set up (Fig. 5).

Figure:   Acid-base equilibrium of methylamine and water.

Methylamine on the left side of the equation is known as the free base, whereas the methyl ammonium ion formed on the right side is known as the conjugate acid. The extent of ionization or dissociation in the equilibrium reaction is described by the equilibrium constant (K_eq);

K_eq = [Products]/ [reactants] = [CH₃NH₃⁺][HO^-]/ [CH₃NH₂][H₂O]

K_b = K_eq [H₂O] = [CH₃NH₃⁺] [HO^-]/ [CH₃NH₂]

K_eq is generally measured in a dilute aqueous solution of the base and thus the concentration of water is high and supposed to be constant. Hence, we can rewrite the equilibrium equation in a much simple form in which K_b is the basicity constant and involves the concentration of pure water (55.5 M). pK_b is the negative logarithm of K_b and is employed like a measure of basic strength (pK_b = Log₁₀K_b).

A large pK_b points out a weak base. For instance, the pK_b values of ammonia and  methylamine are 4.74 and 3.36, correspondingly that points out that ammonia is a weaker base as compared to the methylamine.

pK_b and pK_a are related via the equation pK_a+ pK_b = 14. Therefore, if one knows the pKa of an acid, the pK_b of the conjugate base can be calculated and conversely.