In any solution, the value of pK_a indicates whether an acid is completely dissociated or not. A negative pK_a corresponds to a stronger acid, whereas a positive pK_a corresponds to a weaker acid. Consider the reaction between ammonia and an ethoxide ion. In this reaction, ethanol with a pK_a of 15.9 is a stronger acid than ammonia with a pK_a of 38. Recall that the strong acid forms a weak conjugate base, and a weak acid forms a strong conjugate base. Hence, the ethoxide ion is a weak base.

Figure 1. Acid–base reaction between ammonia and an ethoxide ion

In an equilibrium-controlled acid–base reaction, the equilibrium position always favours the formation of the weaker acid and the weaker base. This is because the weaker acid and the weaker base are the most stable species due to their lower potential energies. Therefore, in this reaction, the equilibrium favours the formation of the ethoxide ion and ammonia.

Additionally, the position of the equilibrium can also be identified using the equilibrium constant that is calculated from pK_a values. For an acid–base reaction, the equilibrium constant, K_eq, can be calculated by subtracting the pK_a value of the acid on the left side from the pK_a value of the acid on the right side and then taking the antilog of the result. Depending on the value of the equilibrium constant, the position of the equilibrium is determined. When the value of K_eq is greater than 1, the position of the equilibrium lies far to the product side. On the other hand, with the equilibrium constant smaller than 1, the equilibrium lies far to the reactant side.