How does ∆G° behave when K is greater than 1?

When the equilibrium constant (K) is greater than 1, it indicates that the products of a reaction are favored over the reactants at equilibrium. This preference for products suggests that the reaction is spontaneous under standard conditions, contributing to a negative change in Gibbs free energy (∆G°).

The relationship between ∆G° and K is defined by the equation:

[ \Delta G^\circ = -RT \ln K ]

In this equation, R is the gas constant and T is the temperature in Kelvin. When K is greater than 1, the natural logarithm of K (ln K) is positive, leading to a negative value for ∆G°. This negative ∆G° signifies that the reaction proceeds spontaneously in the forward direction toward products under standard conditions.

Therefore, the correct conclusion is that when K is greater than 1, ∆G° is indeed less than 0, aligning with the understanding of spontaneous reactions and their favorability towards product formation at equilibrium.