What is the relationship between ΔG and the equilibrium constant K?

The relationship between the change in Gibbs free energy (ΔG) and the equilibrium constant (K) stems from fundamental thermodynamic principles. At equilibrium, the change in Gibbs free energy for a reaction is related to the equilibrium constant through the equation ΔG = -RT ln K.

In this equation, R represents the universal gas constant, T is the absolute temperature in Kelvin, and K is the equilibrium constant for the reaction. This relationship indicates that a negative value of ΔG corresponds to a spontaneous reaction in the forward direction and a positive value indicates non-spontaneity. Specifically:

• If K is greater than 1, it implies that the products of the reaction are favored at equilibrium, resulting in a negative ΔG, which suggests the reaction can occur spontaneously.

• Conversely, if K is less than 1, the reactants are favored, and ΔG will be positive, indicating that the reaction is non-spontaneous under standard conditions.

Understanding this relationship helps chemists predict whether a reaction will proceed spontaneously based on its equilibrium constant. In essence, this equation integrates both thermodynamics and kinetics to provide insights into chemical reactions and their favorability.