Acid–base equilibrium

Acid–base equilibrium is the balance between proton-transfer forms of acids and bases in solution. Its position is described by equilibrium constants such as Ka and Kb.

Brønsted–Lowry acid–base theory: A theory defining acids as proton donors and bases as proton acceptors. Its proton-transfer definition supplies the framework for describing these equilibria.
Chemical equilibrium: A dynamic state in which forward and reverse reaction rates are equal. Proton-transfer forms coexist because their interconversion reaches this balance.
Buffer solution: A solution that resists pH change through a weak acid–base pair. Buffers work because both protonation forms remain available near equilibrium.
Strong acid: An acid that nearly completely transfers its protons to water in dilute solution. Its near-complete ionization contrasts with the partial dissociation typical of weak acids.
Activity coefficient: A factor relating a species' chemical activity to its concentration. It corrects equilibrium calculations when ionic interactions make solutions nonideal.
Acid dissociation constant: An equilibrium constant that quantifies an acid's tendency to donate a proton in solution. Ka measures the position of an acid's proton-transfer equilibrium.
pH: A logarithmic measure of hydrogen-ion activity in a solution. pH locates a solution relative to an acid's dissociation tendency.
Acid–base titration: A quantitative method that determines an acid or base concentration by controlled neutralization. Its titration curve tracks changing equilibrium composition as titrant is added.
Weak acid: An acid that only partially transfers its protons to water in solution. Its appreciable equilibrium between forms is the clearest case for equilibrium analysis.
Ionic strength: A measure of the total concentration of ions weighted by their charges squared. Changing ionic strength shifts activity coefficients and apparent equilibrium behavior.

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