In a closed system, reversible reactions attain a state of dynamic equilibrium when the rates of forward and reverse reactions are equal.

An equilibrium can only form in a closed system, where no reactants or products can escape, because this would decrease the concentration of the reactants of either the forward or reverse reactions which would decrease the rate of that reaction and cause the opposite reaction to overtake, as in the following example.

$$ \begin{align} \ce{CaCO_3} &\rightleftharpoons \ce{CaO + CO_2} \end{align} $$

At equilibrium, the concentrations of reactants and products remain constant, but are rarely equal.

Although the concentration of the reactants and products are unchanging, this does not mean that the reactions have stopped. Reactant molecules still react to form products and vice versa at the same rate — the process is said to be dynamic, because reactions are still taking place the equilibrium is described as a dynamic equilibrium.

For a given reversible reaction, the effect of altering temperature or pressure or of adding/removing reactants/products (the concentration) can be predicted.

The general principle ‘the position of the equilibrium of a system changes to minimise the effect of any imposed change in condition’ was discovered by French chemist Henri Le Chatelier.

The addition of a catalyst increases the rates of the forward and reverse reactions equally.

For equilibrium reactions, a catalyst increases the rate at which equilibrium is achieved but does not affect the position of equilibrium.

To maximise profits, chemists employ strategies to move the position of equilibrium in favour of the products.