If ΔG° > 0 at standard conditions, the reaction is...

• A. Spontaneous
• B. Non-spontaneous
• C. At equilibrium
• D. Requires a catalyst

Answer: (B)

A positive ΔG° under standard conditions means the reaction is not spontaneous in that setting. Spontaneity hinges on the sign of Gibbs free energy change: negative ΔG° indicates the process can proceed without added energy, while positive ΔG° means energy input is required for the forward reaction. The reversed reaction would have the opposite sign for ΔG°, so it would be spontaneous if the forward is not. At equilibrium, ΔG° equals zero, so a positive value tells you you’re not at equilibrium in the forward direction. A catalyst changes the rate by providing a lower-energy pathway but does not alter ΔG°, so it doesn’t change whether the reaction is spontaneous. Therefore, the reaction is non-spontaneous under standard conditions.