Acid dissociation constant

In chemistry, an acid dissociation constant (also known as acidity constant, or acid-ionization constant; denoted ${\displaystyle K_{a}}$) is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction

${\displaystyle {\ce {HA <=> A^- + H^+}}}$

known as dissociation in the context of acid–base reactions. The chemical species HA is an acid that dissociates into A⁻, the conjugate base of the acid and a hydrogen ion, H⁺. The system is said to be in equilibrium when the concentrations of its components will not change over time, because both forward and backward reactions are occurring at the same rate.

The dissociation constant is defined by

${\displaystyle K_{\text{a}}=\mathrm {\frac {[A^{-}][H^{+}]}{[HA]}} ,}$ or

${\displaystyle \mathrm {p} K_{{\ce {a}}}=-\log _{10}K_{\text{a}}=\log _{10}{\frac {{\ce {[HA]}}}{[{\ce {A^-}}][{\ce {H+}}]}}}$

where quantities in square brackets represent the concentrations of the species at equilibrium. As a simple example for a weak acid with K_a = 10⁻⁵, log K_a is the exponent which is -5, so that pK_a = 5. And for acetic acid with K_a = 1.8 x 10⁻⁵, pK_a is close to 5. A higher K_a corresponds to a stronger acid which is more dissociated at equilibrium. For the more convenient logarithmic scale, a lower pK_a means a stronger acid.