What defines a spontaneous reaction?

• A. The reactants have a higher free energy than the products
• B. The products have a higher free energy than the reactants
• C. ΔG = 0
• D. The reaction requires activation energy

Answer: (A)

A spontaneous reaction is one that can occur without ongoing external energy input because the system’s free energy decreases as the reaction proceeds. When reactants start with higher free energy than the products, the Gibbs free energy change is negative, meaning the process is thermodynamically favorable and can proceed on its own under constant temperature and pressure.

If the products had higher free energy than the reactants, the reaction would not be spontaneous. If the free energy change were zero, the system would be at equilibrium with no net change. Activation energy is about how fast a reaction occurs (kinetics) and does not define spontaneity; a spontaneous reaction can still require an initial energy input to get over an energy barrier, but once that barrier is overcome, the overall free-energy change drives the reaction forward.