GIBBS ENERGY CHANGE AND EQUILIBRIUM
In thermodynamics, Gibbs free energy (G) is a measure of the maximum amount of work that can be extracted from a thermodynamic system at a constant temperature and pressure, while it undergoes a process from its initial state to a final state. The Gibbs free energy change (∆G) is the difference in Gibbs free energy between the reactants and products in a chemical reaction.
For a chemical reaction, the change in Gibbs free energy can be used to predict the spontaneity of the reaction. A negative value of ∆G indicates that the reaction is spontaneous, while a positive value of ∆G indicates that the reaction is non-spontaneous. If ∆G is zero, the reaction is at equilibrium.
At equilibrium, the forward and backward reactions occur at the same rate, and there is no net change in the concentration of reactants and products. The Gibbs free energy change at equilibrium (∆G°) can be calculated using the equilibrium constant (K) for the reaction and the standard Gibbs free energy change (∆G°f) for the reactants and products.
∆G° = -RTlnK
where R is the gas constant, T is the temperature in Kelvin, and ln is the natural logarithm.
If the value of Q (reaction quotient) is equal to K, then the system is at equilibrium, and ∆G = ∆G°. If Q is less than K, the reaction will proceed forward to reach equilibrium, and ∆G will decrease until it reaches ∆G°. If Q is greater than K, the reaction will proceed backward to reach equilibrium, and ∆G will increase until it reaches ∆G°.
In summary, the Gibbs free energy change (∆G) is a measure of the spontaneity of a reaction, and the value of ∆G at equilibrium (∆G°) can be used to calculate the equilibrium constant for the reaction.