Many chemical reactions can proceed in both directions. A reversible reaction is written with opposing arrows,

$$\mathrm{A+B\rightleftharpoons C+D},$$

because products can react to regenerate reactants as well as reactants forming products.

If a closed reacting system is allowed to evolve under fixed conditions, it can reach chemical equilibrium. At equilibrium, the forward and reverse reaction rates are equal:

$$r_{\mathrm{forward}}=r_{\mathrm{reverse}}.$$

A reaction rate describes how quickly a reaction changes the amounts or concentrations of its species. Equal forward and reverse rates do not mean that the reactions have stopped. Molecules continue reacting in both directions, but the two opposing processes occur at equal average rates. The equilibrium is therefore dynamic.

Equal rates also do not imply equal concentrations of reactants and products. Consider a schematic reaction

$$\mathrm{A\rightleftharpoons B}.$$

One equilibrium mixture might contain much more A than B; another reaction might favor B strongly. What remains constant at equilibrium is the macroscopic composition, not necessarily a 1:1 amount ratio.

Suppose a system initially contains only A. At first, the forward reaction can occur but the reverse reaction cannot because almost no B is present. As B accumulates, the reverse rate increases. Eventually the two rates become equal, after which the amounts of A and B fluctuate microscopically but remain constant on average.

Starting from mostly B can lead to the same equilibrium composition when temperature and other relevant conditions are the same. Equilibrium is therefore a property of the reaction system and conditions, not of which side was chosen as the starting mixture.

The quantitative relation among equilibrium amounts is described by an equilibrium constant.