What does the ionization constant of an acid or a base indicate about either the acid or the base?

Let's look at the ionization of an acid, HA (all species here are aqueous):

#HA rightleftharpoons H^+ + A^-#

In aqueous solution, the acid, #HA# dissociates into #H^+# and #A^-#. The reaction is reversible and the chemical species, #HA#, #H^+# and #A^-# are in equilibrium when the forward and reverse reactions are equal rates, giving the appearance of constant concentrations.

Therefore, the acid ionization constant can be determined from this equation:

#K_a=([H^+][A^-])/([HA])#

where the square brackets indicate the equilibrium concentration for each chemical species.

As the strength of acid increases, #HA# will dissociate more into #H^+# and #A^-#, therefore #[H^+]# and #[A^-]# will be higher than in a weaker acid, making the value of #K_a# higher.

Therefore,

• the higher the value of #K_a# , the stronger it is as an acid in a solution.
• the lower the value of #K_a# , the weaker it is as an acid in a solution.

Similar rationale can be applied to the ionization of a general base, B, except that the equation will look like this:

#B + H_2O(l) rightleftharpoons HB^+ + OH^-#

Therefore, the base ionization constant can be determined from this equation:

#K_b=([HB^+][OH^-])/([B])#

where as before, we have equilibrium concentrations for each species.