Why is the Gibbs free energy change zero at equilibrium?

The Gibbs free energy change is zero at equilibrium because there is no net change in the system.

In more detail, the Gibbs free energy (ΔG) is a thermodynamic potential that measures the maximum reversible work that a system can perform at constant temperature and pressure. It is used to determine whether a reaction will occur spontaneously or not. If ΔG is negative, the reaction is spontaneous; if it is positive, the reaction is non-spontaneous; and if it is zero, the system is at equilibrium.

At equilibrium, the forward and reverse reactions occur at the same rate, meaning there is no net change in the concentrations of reactants and products. This is why ΔG is zero at equilibrium. The system is in a state of maximum stability, and no further work can be done. This is a fundamental principle of thermodynamics.

The relationship between Gibbs free energy, enthalpy, and entropy is given by the equation ΔG = ΔH - TΔS, where ΔH is the change in enthalpy, T is the absolute temperature, and ΔS is the change in entropy. At equilibrium, ΔG equals zero, which implies that ΔH = TΔS. This means that the energy available for work (enthalpy) is being used to increase the disorder (entropy) of the system.

In the context of chemical reactions, the Gibbs free energy change provides valuable information about the feasibility and direction of a reaction. A reaction will proceed in the direction that lowers the Gibbs free energy of the system. At equilibrium, the Gibbs free energy of the system is at its minimum, and the reaction will not proceed in either direction.

In conclusion, the Gibbs free energy change is zero at equilibrium because the system is in a state of maximum stability and no net change is occurring. This is a key concept in thermodynamics and is crucial for understanding the behaviour of chemical reactions.