In chemistry, the equilibrium constant is a theoretically-based number that helps chemists determine the concentration of various reactants or products in a reaction where chemical equilibrium exists.

Contents

1 Reaction Conditions 2 Applications 3 See Also 4 Sources

Reaction Conditions

A typical equilibrium situation is as below:

<math>kA_{aq} + mB_{aq} \leftrightarrow nC_{aq} + pD_{aq}<math>

when the forward reaction is occurring at the same rate as the reverse reaction. For simplicity in this example all reactants are aqueus, in solution, because solids and liquids to not enter the equation for the equilibrium constant as discussed in the article on solubility equilibrium (however gases do).

The equation of the equilibrium constant is equal to the product of the product concentrations to the power their respective stoichiometric coeffecients divided by the product of the reactant concentrations to the power of their stoicheometric coefficients.

<math>K = \frac{\left[A\right]^k \left[B\right]^m} {\left[C\right]^n \left[D\right]^p}<math>

The constant is constant only at a given temperature.

Applications

There are certain implications of the equilibrium constant. If the value is very large, over 1, the reaction is said to lie to the right (of the arrow) indicating a greater concentration of products; values less than 1 lie to the left higher formation rates of reactants, and values of one indicate equal concentrations. Knowledge of the equilibrium constant can help us determine, in an industrial setting for example, how to best produce a desirable material.

For example, in the Haber process for the formation of ammonia, the value of K is around 30 at pressures and temperatures standard for the process. This knowledge can help us determine how to best make more ammonia. We cannot simply place all the reactants in a tank and expect an eventual yield of 100% ammonia, it is an equilibrium and therefore by definition this would not occur. Therefore, according to Le Chatelier's Principle the optimum way is to allow the ammonia to form, and then remove it from the tank. We know this is efficient because of the very high value of the constant, whereas if the value were significantly lower, perhaps in the area of 0.001 or really anything below 1, it wouldn't be worth the effort due to the extended periods of time it would take for economically feasible product yields.

Note that all the above applies only to static equilibria, when concentrations change such as when an equilibrium is established and then additional substances are added, calculations and concepts become less straightforward. This latter situation is known as a dynamic equilibrium.

See Also

• Acidity Constant
• Solubility equilibrium

Sources

Zumdahl, Steven; Zumdahl, Susan. Zumdahl Chemistry, 4th Edition.