Describe the relationship between Gibbs energy change and chemical equilibrium.

(N/A) When reactions proceed in both directions,a dynamic equilibrium is established. This occurs only when the Gibbs free energy of the system is at a minimum.
The criterion for equilibrium for a reaction $A + B \rightleftharpoons C + D$ is $\Delta_{r} G = 0$.
Gibbs energy for a reaction where all reactants and products are in their standard states,$\Delta_{r} G^{\circ}$,is related to the equilibrium constant $K$ as follows:
$\Delta_{r} G = \Delta_{r} G^{\circ} + RT \ln Q$
At equilibrium,$\Delta_{r} G = 0$ and $Q = K$,therefore:
$0 = \Delta_{r} G^{\circ} + RT \ln K$
$\Delta_{r} G^{\circ} = -RT \ln K = -2.303 RT \log K$
Also,$\Delta_{r} G^{\circ} = \Delta_{r} H^{\circ} - T\Delta_{r} S^{\circ}$.
For strongly endothermic reactions,$\Delta_{r} H^{\circ} > 0$,which leads to $K < 1$,meaning the reaction does not favor product formation.
For exothermic reactions,$\Delta_{r} H^{\circ} < 0$,which leads to $K > 1$,meaning the reaction favors product formation.