Equilibrium and Changes to Concentration / Partial Pressure

Le Châtelier’s Principle tells us that if you add more of a reactant to a chemical reaction at equilibrium (or increase the partial pressure of a gaseous reactant), then the system will try to counteract this change by shifting the equilibrium position to the right and producing more products. However, LCP does not provide an explanation for this prediction – for that, we must look to collision theory.

According to collision theory, an increase in the concentration of a reactant (or partial pressure of a gaseous reactant) will increase the frequency of collisions between reactants and lead to an increase in the rate of forward reaction.

(It's worth pointing out that partial pressure in a gaseous system practically works the same as concentration does in an aqueous system.)

Since the rates of forward and reverse reaction are equal in a system at chemical equilibrium, when the concentration of a reactant (or partial pressure of a gaseous reactant) is suddenly increased, the rate of forward reaction will suddenly exceed the rate of reverse reaction. This will cause the equilibrium position to shift to the right, increasing the concentration of products and decreasing the concentration of reactants.

As the concentration of products (or partial pressure of gaseous products) increases and the concentration of reactants (or partial pressure of gaseous reactants) decreases, the rate of forward reaction will begin to decrease, and the rate of reverse reaction will begin to increase. This will continue until the rate of reverse reaction equals the rate of forward reaction – the restoration of equilibrium.