We know that if,
- ΔG is negative, then the reaction is spontaneous and proceeds in the forward direction.
- ΔG is positive, then reaction is considered non-spontaneous. Instead, as reverse reaction would take place
- ΔG is 0, reaction has achieved equilibrium; at this point, there is no longer any free energy left to drive the reaction.
A mathematical expression of this thermodynamic view of equilibrium can be described by the following equation:
ΔG = ΔG⁰ + RT ln Q
Where, G⁰ is standard Gibbs energy.
At equilibrium, when ΔG = 0 and Q = Kc
ΔG = ΔG⁰ + RT ln K = 0
ΔG⁰ = -RT ln K
ln K = ΔG⁰ / RT
Taking antilog of both sides, we get,
K = e-ΔG⁰ /RT
- If ΔG0 < 0, then –ΔG0/RT is positive, and e-ΔG⁰ /RT >1, making K >1, which implies a spontaneous reaction or the reaction which proceeds in the forward direction to such an extent that the products are present predominantly.
- If ΔG0 > 0, then –ΔG0/RT is negative, and e-ΔG⁰ /RT < 1, that is, K < 1, which implies a non-spontaneous reaction or a reaction which proceeds in the forward direction to such a small degree that only a very minute quantity of product is formed.