Acid dissociation constant

The equilibrium constant for an acid’s proton transfer to a solvent. Its logarithmic form, pKa, makes acid strengths easier to compare.

Acid–base reaction: A chemical reaction involving transfer of a proton or formation of a shared electron pair. Acid dissociation is an equilibrium case of proton transfer.
Brønsted–Lowry acid–base theory: A theory defining acids as proton donors and bases as proton acceptors. Ka applies the theory’s proton-donor definition to an equilibrium.
Buffer solution: A solution that resists changes in pH when small amounts of acid or base are added. A buffer is most effective near the pKa of its weak acid.
Strong acid: An acid that dissociates nearly completely in a specified solvent at ordinary concentrations. Strong acids contrast with weak acids, whose partial dissociation is summarized by Ka.
Equilibrium constant: A ratio of product and reactant activities at chemical equilibrium, raised to their stoichiometric coefficients. Ka is the equilibrium constant for an acid’s dissociation reaction.
Conjugate acid–base pair: Two species that differ by one proton, such as an acid and the base formed when it donates that proton. Each dissociation equilibrium connects an acid to its conjugate base.
Acid–base titration: A method for determining concentration by reacting an acid with a base of known concentration, or vice versa. Titration curves reveal pKa values and acid-equivalence behavior.
Weak acid: An acid that dissociates only partially in a specified solvent at equilibrium. Its measurable partial dissociation is the usual setting for Ka and pKa comparisons.
Chemical equilibrium: A dynamic state in which forward and reverse reaction rates are equal. The balance between proton donation and return establishes the measured Ka.
Hydronium: The hydrated proton commonly represented in water as the ion H₃O⁺. In aqueous solutions, proton donation is represented through formation of hydronium.

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