ΔG° and Keq - MCAT General Chemistry

ΔG° and Keq

8 min read Updated Mar 26, 2026

In the previous section, you learned that ΔG determines spontaneity. But which ΔG? There is a critical distinction between ΔG° (standard free energy change) and ΔG (actual free energy change under real conditions). Confusing the two is one of the most common mistakes on the MCAT.

ΔG° and the Equilibrium Constant

Interpreting ΔG° and K

The sign of ΔG° tells you which side of the reaction is favored at equilibrium:

ΔG°	K	Meaning
Negative	K > 1	Products favored at equilibrium
Zero	K = 1	Neither side favored
Positive	K < 1	Reactants favored at equilibrium

ΔG Under Non-Standard Conditions

Real reactions rarely occur under standard conditions. The actual free energy change depends on the current concentrations via the reaction quotient Q:

How ΔG Changes as a Reaction Proceeds

When a reaction starts:

If Q < K: ΔG < 0 (reaction proceeds forward to make more products)
If Q > K: ΔG > 0 (reaction proceeds backward to make more reactants)
If Q = K: ΔG = 0 (equilibrium - no net change)

As the reaction approaches equilibrium, ΔG approaches zero. At equilibrium, Q = K and:

ΔG = ΔG° + RTlnK = 0

This is exactly where the equation ΔG° = -RTlnK comes from - it is the ΔG = 0 condition rearranged.

Connecting Everything

Here is how the three big thermodynamic equations fit together:

ΔG = ΔH - TΔS (relates free energy to enthalpy and entropy)
ΔG° = -RTlnK (relates standard free energy to equilibrium)
ΔG = ΔG° + RTlnQ (relates actual free energy to current conditions)

These three equations are the thermodynamic backbone of the MCAT. Know them cold.

Temperature and K

Because ΔG° depends on temperature (through the TΔS° term in ΔG° = ΔH° - TΔS°), and K depends on ΔG° (through ΔG° = -RTlnK), the equilibrium constant changes with temperature.

For an exothermic reaction (ΔH° < 0): increasing temperature makes ΔG° less negative (or more positive), which decreases K. Heat shifts equilibrium toward reactants.

For an endothermic reaction (ΔH° > 0): increasing temperature makes ΔG° more negative, which increases K. Heat shifts equilibrium toward products.

This connects directly to Le Chatelier’s principle from Chapter 6: heat acts as a reactant (endothermic) or product (exothermic).

At equilibrium, what is the value of ΔG?

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ΔG = 0. At equilibrium, the forward and reverse reactions are balanced, and there is no thermodynamic driving force in either direction. Note: ΔG° is NOT zero at equilibrium (unless K = 1). ΔG° is a fixed property of the reaction. ΔG is the quantity that equals zero at equilibrium.

A reaction has ΔG° = -30 kJ/mol. Is this reaction spontaneous under all conditions?

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Not necessarily. ΔG° = -30 kJ/mol means the reaction is spontaneous under STANDARD conditions (1 M, 1 atm) and that K > 1. But under non-standard conditions, ΔG = ΔG° + RTlnQ. If Q is very large (lots of products already present), RTlnQ could be positive enough to make ΔG > 0, meaning the reaction would actually go in reverse. ΔG° tells you about the equilibrium position; ΔG tells you about the current direction.