Calculating the Equilibrium Constant

Understanding the Equilibrium Constant

The equilibrium constant (K) is a measure of the ratio of concentrations of products to reactants at equilibrium at a certain temperature.
Each concentration in the equilibrium constant expression is raised to the power of its stoichiometric coefficient found in the balanced chemical equation.
The equilibrium constant is dimensionless. It does not have a unit.
Different reactions have different equilibrium constants. A large K denotes a reaction which leans towards products, whereas a small K implies a reaction favoring reactants.

Form of the Equilibrium Constant

There are two forms of the equilibrium constant commonly used: Kc and Kp.
Kc refers to molar concentration (expressed in moles per litre) while Kp refers to partial pressure (expressed in atmospheres or bar).
Kc is used when concentrations are given, whereas Kp is used when gases are involved and their partial pressures are provided.

Calculating the Equilibrium Constant

To calculate K, you should know the concentrations of products and reactants at equilibrium.
Keep in mind that concentrations are always positive values.
Any pure solids or liquids do not appear in the equilibrium constant expression.
Put the concentrations of the products in the numerator and those of the reactants in the denominator to form the expression.
Solving for K involves plugging these equilibrium concentrations into the equilibrium constant expression and performing the calculations.
Check to ensure your answer is a positive value.

Changing Conditions and the Equilibrium Constant

The equilibrium constant, K, remains unchanged by the addition of a catalyst. Catalysts affect the speed at which equilibrium is reached, but not the position of equilibrium.
Changing the concentration of reactants or products initially does not change the value of K, but shifts the equilibrium position to re-establish the ratio of concentrations that equals K.
The equilibrium constant, K, does change with changes in temperature, as it is dependent upon the reaction’s endothermic or exothermic nature. Higher temperature favours an endothermic reaction, reducing K, and favours an exothermic reaction, increasing K.