The equilibrium constant of a chemical reaction is the value of its reaction quotient at chemical equilibrium, a state approached by a dynamic chemical system after sufficient time has elapsed at which its composition has no measurable tendency towards further change. For a given set of reaction conditions, the equilibrium constant is independent of the initial analytical concentrations of the reactant and product species in the mixture. Thus, given the initial composition of a system, known equilibrium constant values can be used to determine the composition of the system at equilibrium. However, reaction parameters like temperature, solvent, and ionic strength may all influence the value of the equilibrium constant.

A knowledge of equilibrium constants is essential for the understanding of many chemical systems, as well as the biochemical processes such as oxygen transport by hemoglobin in blood and acid–base homeostasis in the human body.

Stability constants, formation constants, binding constants, association constants and dissociation constants are all types of equilibrium constants.

For a system undergoing a reversible reaction described by the general chemical equation

a thermodynamic equilibrium constant, denoted by ${\displaystyle K^{\ominus }}$, is defined to be the value of the reaction quotient Q_t when forward and reverse reactions occur at the same rate. At chemical equilibrium, the chemical composition of the mixture does not change with time, and the Gibbs free energy change ${\displaystyle \Delta G}$ for the reaction is zero. If the composition of a mixture at equilibrium is changed by addition of some reagent, a new equilibrium position will be reached, given enough time. An equilibrium constant is related to the composition of the mixture at equilibrium by

where {X} denotes the thermodynamic activity of reagent X at equilibrium, [X] the numerical value of the corresponding concentration in moles per liter, and γ the corresponding activity coefficient. If X is a gas, instead of [X] the numerical value of the partial pressure ${\displaystyle P_{X}}$ in bar is used. If it can be assumed that the quotient of activity coefficients, ${\displaystyle \Gamma }$, is constant over a range of experimental conditions, such as pH, then an equilibrium constant can be derived as a quotient of concentrations.

An equilibrium constant is related to the standard Gibbs free energy change of reaction ${\displaystyle \Delta G^{\ominus }}$ by

where R is the universal gas constant, T is the absolute temperature (in kelvins), and ln is the natural logarithm. This expression implies that ${\displaystyle K^{\ominus }}$ must be a pure number and cannot have a dimension, since logarithms can only be taken of pure numbers. ${\displaystyle K_{c}}$ must also be a pure number. On the other hand, the reaction quotient at equilibrium