Which statement correctly defines ΔG and identifies the sign of ΔG° that indicates spontaneity at constant temperature and pressure?

• A. ΔG = ΔH − TΔS; ΔG° < 0 indicates non-spontaneity.
• B. ΔG = ΔS − TΔH; ΔG° > 0 indicates spontaneity.
• C. ΔG = ΔH + TΔS; ΔG° < 0 indicates non-spontaneity.
• D. ΔG = ΔH − TΔS; ΔG° < 0 indicates spontaneity.

Answer: (D)

Gibbs free energy change tells you whether a process can occur spontaneously at constant temperature and pressure. It is defined as ΔG = ΔH − TΔS. When ΔG is negative, the process can proceed without adding energy, so it’s spontaneous; when ΔG is positive, it requires input of energy and is non-spontaneous; ΔG equals zero means the system is at equilibrium.

For standard conditions, the standard free energy change ΔG° follows the same idea: ΔG° < 0 indicates spontaneity under those standard conditions. That negative value also connects to the equilibrium constant through ΔG° = −RT ln K, so a negative ΔG° corresponds to K > 1 (products favored at equilibrium).

Think about how the two terms compete: a negative ΔH (exothermic) or a positive ΔS (increasing disorder) both push ΔG negative, especially at higher temperatures where the TΔS term becomes more influential. Conversely, a positive ΔH or negative ΔS tends to make ΔG positive, hindering spontaneity.