Reversible reaction

A reversible reaction is a reaction in which the conversion of reactants to products and the conversion of products to reactants occur simultaneously. \mathit aA{} + \mathit bB \mathit cC{} + \mathit dD A and B can react to form C and D or, in the reverse reaction, C and D can react to form A and B. This is distinct from a reversible process in thermodynamics. Weak acids and bases undergo reversible reactions. For example, carbonic acid: H2CO3 (l) + H2O(l) ⇌ HCO3−(aq) + H3O+(aq). The concentrations of reactants and products in an equilibrium mixture are determined by the analytical concentrations of the reagents (A and B or C and D) and the equilibrium constant, K. The magnitude of the equilibrium constant depends on the Gibbs free energy change for the reaction. So, when the free energy change is large (more than about 30 kJ mol−1), the equilibrium constant is large (log K > 3) and the concentrations of the reactants at equilibrium are very small. Such a reaction is sometimes considered to be an irreversible reaction, although small amounts of the reactants are still expected to be present in the reacting system. A truly irreversible chemical reaction is usually achieved when one of the products exits the reacting system, for example, as does carbon dioxide (volatile) in the reaction CaCO3 + 2HCl → CaCl2 + H2O + CO2↑ The concept of a reversible reaction was introduced by Berthollet in 1803, after he had observed the formation of sodium carbonate crystals at the edge of a salt lake (one of the natron lakes in Egypt, in limestone): 2NaCl + CaCO3 → Na2CO3 + CaCl2 He recognized this as the reverse of the familiar reaction Na2CO3 + CaCl2→ 2NaCl + CaCO3 Until then, chemical reactions were thought to always proceed in one direction. Berthollet reasoned that the excess of salt in the lake helped push the "reverse" reaction towards the formation of sodium carbonate. In 1864, Waage and Guldberg formulated their law of mass action which quantified Berthollet's observation.