What Delta G = 0 Means: Equilibrium, Q vs K and Delta G Standard

ΔG = 0 means the system is at equilibrium. There is no net driving force in either direction, the forward and reverse rates are equal, and the reaction quotient Q equals the equilibrium constant K.

The master equation

ΔG = ΔG° + RT ln Q

At equilibrium, ΔG = 0 and Q = K, so ΔG° = -RT ln K. That is the link between standard free energy and K.

Q vs K tells you the direction

You can rewrite the master equation as ΔG = RT ln(Q/K). It makes the table easy to rebuild: if there are fewer products than equilibrium needs, Q is smaller than K and the reaction runs forward.

ΔG vs ΔG°

• ΔG° is fixed for a reaction at given conditions and tells you K. It is the value at 1 M and standard state, not at equilibrium.
• ΔG depends on the current mixture through Q and is the quantity that predicts the real direction.
• A reaction with positive ΔG° (K less than 1) can still run forward if Q is small enough that ΔG is negative. This is how cells drive unfavorable steps by keeping product concentrations low.

Want this topic as spaced-repetition cards? The matching StudyStore MCAT pack has original Anki cards, a study guide and practice questions for it.

Worked example

A reaction has K = 10 and Q = 0.1 at 298 K. Q is 100 times smaller than K, so ln(Q/K) = ln(0.01), about -4.6. With RT near 2.5 kJ/mol, ΔG is about -11.5 kJ/mol. It is negative, so the reaction proceeds forward until Q rises to 10.

Common traps

• Saying a reaction with ΔG° = 0 is at equilibrium. That means K = 1, not that Q = K.
• Using ΔG° to predict direction in a cell. Use ΔG with the real concentrations.
• Linking ΔG = 0 to no reaction. The reaction continues in both directions at equal rates.

One-minute recap

ΔG = 0 means equilibrium and Q = K. If Q is below K, ΔG is negative and the reaction goes forward. If Q is above K, ΔG is positive and it goes in reverse. ΔG° sets K; ΔG sets the direction now. The broader picture is in the Gibbs free energy note.