Introduction:
The standard free energy change (∆G°) is a thermodynamic property that indicates the direction and extent of a chemical reaction. It is related to the equilibrium constant (K) by the equation: ∆G° = -RTlnK, where R is the gas constant and T is the temperature. Reactions with a ∆G° less than zero are considered spontaneous and can proceed in the forward direction, while reactions with a positive ∆G° are non-spontaneous and require an input of energy to proceed.

Reactions with ∆G° less than zero:
Reactions that have a standard free energy change (∆G°) less than zero are considered exergonic or spontaneous. This means that the products of the reaction have a lower free energy than the reactants, and the reaction can proceed in the forward direction without the need for an input of energy. In such cases, the equilibrium constant (K) is always greater than unity.

Explanation:
1. Equilibrium constant equal to unity: If a reaction has a ∆G° equal to zero, it means that the reactants and products have the same free energy and are in a state of equilibrium. In this case, the equilibrium constant (K) is equal to unity, as the forward and reverse reaction rates are equal. However, for reactions with ∆G° less than zero, the reaction proceeds spontaneously in the forward direction, and the equilibrium constant is not equal to unity.

2. Equilibrium constant greater than unity: Reactions with a ∆G° less than zero have a negative value for ∆G°, indicating that the reaction is spontaneous and proceeds in the forward direction. According to the equation ∆G° = -RTlnK, a negative ∆G° leads to a positive value for the natural logarithm of the equilibrium constant (lnK). Since the natural logarithm of a number greater than one is positive, it follows that the equilibrium constant (K) must be greater than unity.

3. Equilibrium constant less than unity: Reactions with a positive ∆G° (non-spontaneous) have an equilibrium constant (K) less than unity. This is because the natural logarithm of a number less than one is negative, leading to a positive value for ∆G° according to the equation ∆G° = -RTlnK. Non-spontaneous reactions require an input of energy to proceed in the forward direction, and their equilibrium constant is less than unity.

Conclusion:
Reactions with standard free energy changes (∆G°) less than zero have a spontaneous nature and proceed in the forward direction without an input of energy. The equilibrium constant (K) for these reactions is always greater than unity as a result of the relationship between ∆G° and K. This understanding of the relationship between thermodynamic properties and reaction spontaneity is crucial in predicting and analyzing chemical reactions.

2 ) greater than unity