The equilibrium constant expression is written as the molar concentrations of the products, C and D, over the reactants, A and B, at equilibrium, each raised to their respective stoichiometric coefficients. When solved, the expression is equal to the equilibrium constant, K_c.

An expression in the same form can also be written for the reactants and products at any concentration, and the calculated quantity is known as the reaction quotient, Q_c.

Like Q_c, the Q_p expression can be written for gaseous reactions using partial pressures.

While K remains constant at a specific temperature irrespective of concentration, the value of Q changes as the reaction proceeds towards the products or the reactants.

The reaction quotient can be used to determine the direction a reaction will proceed in order to reach equilibrium.

At the start of a given reaction, if the concentration of the products is zero, the reaction quotient is zero.

Whenever the concentration of the reactants in the denominator is high, such that Q is smaller than K, the reaction will move to the right to synthesize more products until the system reaches equilibrium.

If the concentration of the reactants is zero, the reaction quotient is infinite.

Whenever the concentration of the products in the numerator is high, such that Q is larger than K, the reaction will move to the left to produce more reactants.

If Q is equal to K, the system is at equilibrium, and the rate of the forward and reverse reactions are equal.

Consider the given reaction with an equilibrium constant 50. If the reaction mixture contains 0.20 molar hydrogen, 0.20 molar iodine, and 1.7 molar hydrogen iodide, the direction of the reaction shown can be determined by calculating Q.

Substituting the given concentrations into the expression, Q equals 72, which is greater than K. Therefore, the reaction will shift towards the left.