For a system undergoing a reversible reaction described by the general chemical equation
a thermodynamic equilibrium constant, denoted by , is defined to be the value of the reaction quotient Q 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  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  in bar is used. If it can be assumed that the quotient of activity coefficients, , 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  by 
where R is the universal gas constant, T is the absolute temperature (in kelvins must be a pure number and cannot have a dimension, since logarithms can only be taken of pure numbers.  must also be a pure number. On the other hand, the reaction quotient at equilibrium
does have the dimension of concentration raised to some power (see , below). Such reaction quotients are often referred to, in the biochemical literature, as equilibrium constants.
For an equilibrium mixture of gases, an equilibrium constant can be defined in terms of partial pressure or fugacity.
An equilibrium constant is related to the forward and backward rate constants, k and k of the reactions involved in reaching equilibrium: