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Acetic acid, a weak acid, donates a proton (white) to water in an equilibrium reaction to give the acetate ion and the hydronium ion. Red: oxygen, grey: carbon.

An acid dissociation constant, K_a, (also known as acidity constant, or acid-ionization constant) is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction known as dissociation in the context of acid-base reactions. The equilibrium can be written symbolically as

HA ⇌ A⁻ + H⁺,

where HA is a generic acid which dissociates into A⁻, known as the conjugate base of the acid, and the hydrogen ion or proton, H⁺, which, in the case of aqueous solutions, exists as a solvated hydronium ion. The chemical species HA, A⁻ and H⁺ are said to be in equilibrium when their concentrations do not change with the passing of time. The dissociation constant is usually written as a quotient of the equilibrium concentrations, denoted by [HA], [A⁻] and [H⁺]:

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

Due to the many orders of magnitude spanned by K_a values, a logarithmic measure of the acid dissociation constant is often preferred in practice. pK_a, which is equal to −log₁₀ K_a, may also be referred to as an acid dissociation constant. The larger the value of pK_a, the smaller the extent of dissociation. A weak acid has a pK_a value in the approximate range −2 to 12 in water. Acids with a pK_a value of less than about −2 are said to be strong acids; a strong acid is almost completely dissociated in aqueous solution, to the extent that the concentration of the undissociated acid becomes undetectable. pK_a values for strong acids can, however, be estimated by theoretical means or by extrapolating from measurements in non-aqueous solvents such as acetonitrile and dimethyl sulphoxide.

The acid dissociation constant for an acid is a direct consequence of the underlying thermodynamics of the dissociation reaction; the pK_a value is directly proportional to the standard Gibbs free energy change for the reaction. The value of the pK_a changes with temperature and can be understood qualitatively based on Le Chatelier's principle: when the reaction is endothermic, the pK_a increases with increasing temperature; the opposite is true for exothermic reactions. The underlying structural factors that determine the magnitude of pK_a values include Pauling's rules for acidity constants, inductive effects, mesomeric effects, and hydrogen bonding.

The quantitative behaviour of acids and bases in solution can only be understood if their pK_a values are known. For example, many compounds used for medication are weak acids or bases, so a knowledge of the pK_a and water–octanol partition coefficient is essential for an understanding of the extent to which the compound enters the blood stream. There are many other applications, including aquatic chemistry, chemical oceanography, buffer solutions, acid-base homeostasis and enzyme kinetics. A knowledge of pK_a values is also a prerequisite for a quantitative understanding of the interaction between acids or bases and metal ions to form complexes in solution. Experimentally, pK_a values can be determined by potentiometric (pH) titration, but for values of pK_a less than about 2 or more than about 11 spectrophotometric or NMR measurements may be required due to practical difficulties with pH measurements.