Equilibrium constants
The equilibrium constant of a chemical reaction is the value of the reaction quotient when the reaction has reached equilibrium. An equilibrium constant value is independent of the analytical concentrations of the reactant and product species in a mixture, but depends on temperature and on ionic strength. Known equilibrium constant values can be used to determine the composition of a system at equilibrium.
For a general chemical equilibrium
the thermodynamic equilibrium constant can be defined such that, at equilibrium,
where curly brackets denote the thermodynamic activities of the chemical species. The right-hand side of this equation corresponds to the reaction quotient Q for arbitrary values of the activities, and becomes the equilibrium constant as shown when the reaction is at equilibrium.
An equilibrium constant is related to the standard Gibbs free energy change for the reaction.
If deviations from ideal behaviour are neglected, the activities of solutes may be replaced by concentrations, [A], and the activity quotient becomes a concentration quotient, Kc.
Kc is defined as equal to the thermodynamic equilibrium constant but with concentrations of reactants and products instead of activities. (Kc appears here to have units of concentration raised to some power while Ko is dimensionless; however as discussed below under Definitions, the concentration factors in Kc are properly divided by a standard concentration so that Kc is dimensionless also.)
Again assuming ideal behavior, the activity of a solvent may be replaced by its mole fraction, or approximately by 1 in dilute solution. The activity of a pure liquid or solid phase is exactly 1. The activity of a species in an ideal gas phase may be replaced by its partial pressure.
A knowledge of equilibrium constants is essential for the understanding of many chemical systems, as well as biochemical processes such as oxygen transport by hemoglobin in blood and acid-base homeostasis in the human body.
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