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A reversible chemical reaction is a process in which reactants convert to products and products simultaneously convert back to reactants. The double arrow notation signifies this reversible nature. In the thermal decomposition of phosphorus pentachloride, for example, gaseous PCl₅ decomposes into PCl₃ and Cl₂ while these products recombine to reform PCl₅, demonstrating how both forward and reverse reactions occur simultaneously.

Chemical equilibrium is dynamic because both forward and reverse reactions continue occurring at equal rates. Although the system appears static from outside, molecular-level activity remains high. When the rate of product formation equals the rate of reactant formation, concentrations stabilize. Reactants and products form and are consumed simultaneously, maintaining constant but not necessarily equal concentrations.

As a reversible reaction proceeds, reactant concentration decreases, slowing the forward reaction rate. Simultaneously, product concentration increases, accelerating the reverse reaction rate. This continues until both rates match, achieving equilibrium. At this point, the rate of NO₂ formation equals its consumption rate, and N₂O₄ formation equals its consumption rate, stabilizing all concentrations.

A homogeneous equilibrium occurs when all reactants and products exist in the same phase—either gaseous or aqueous. A heterogeneous equilibrium involves reactants and products in two or more different phases. For example, the decomposition of N₂O₄ gas to NO₂ gas is homogeneous, while reactions involving solids and gases represent heterogeneous equilibria.

At equilibrium, concentrations remain constant because reactants and products form at the same rates they are consumed. The forward reaction rate equals the reverse reaction rate, creating a balance. Although individual molecules continue reacting, the overall composition of the system does not change, making equilibrium a dynamic steady state rather than a static condition.

Initially, reactant concentration is finite while product concentration is zero, so the forward reaction proceeds at a finite rate. As the reaction progresses, reactant concentration decreases while product concentration increases. This changing concentration affects reaction rates: the decreasing reactant concentration slows the forward reaction, while increasing product concentration speeds the reverse reaction until equilibrium is reached.

When a system reaches equilibrium, the forward and reverse reaction rates become equal. Although both reactions continue occurring, they proceed at identical rates, so no net change in concentrations occurs. This equal-rate condition defines the equilibrium state and explains why equilibrium concentrations from initial concentrations can be predicted using equilibrium expressions and constants.