Gibbs Free Energy and Thermodynamic Favorability

The change in Gibbs free energy, or delta G, is the energy absorbed or released when energy-storing bonds are formed or broken, respectively.

The sign of delta G depends on the signs and the relative values of enthalpy, entropy, and temperature.

If delta H is negative and delta S is positive, delta G is negative at all temperatures. Thus, exothermic reactions where the entropy of the system increases are always spontaneous.

If both delta H and delta S are negative, delta G depends on the temperature. Reactions with negative enthalpy and entropy changes are spontaneous only at low temperatures.

Delta G is also dependent on temperature if both delta H and delta S are positive. Reactions with positive enthalpy and entropy changes are spontaneous only at higher temperatures.

When delta H is positive and delta S is negative, delta G is always positive, and the reaction is nonspontaneous at all temperatures.

For any reaction mixture composition, the delta G for the reaction is the sum of the standard free energy and RT times the natural log of the reaction quotient.

When the reactants and products are at equilibrium, the free energy change is zero, and the reaction quotient equals the equilibrium constant. So, the standard free energy change equals negative RT ln(K).

If delta G naught is less than zero, ln(K) is positive, meaning K is greater than 1. In this case, product formation is favored at equilibrium. The larger the equilibrium constant, the greater the decrease in Gibbs free energy.

Reactions with negative delta G values involve the release of free energy to the surroundings and are called ‘exergonic’ reactions.

Conversely, if delta G naught is greater than zero, ln(K) is negative, meaning K is less than 1, and the reverse direction of the reaction is favored.

Reactions with a positive delta G absorb free energy from the surroundings and are called ‘endergonic’ reactions.