D22.3 ΔG° and K°

We know that there is a qualitative relationship between ΔrG° and equilibrium constant for a given reaction. The standard Gibbs free energy change for a reaction indicates whether a reaction is product-favored at equilibrium (ΔrG° < 0) or reactant-favored at equilibrium (ΔrG° > 0). A strongly product-favored reaction (large negative ΔrG°) has a large equilibrium constant (K>> 1) and a strongly reactant-favored reaction (large positive ΔrG°) has a very small equilibrium constant (K<<1, a very small fraction because K cannot be negative).

Quantitatively, this relationship between the equilibrium constant and ΔrG° is expressed by the equation:

ΔrG° = −RT(lnK°)        or        K° = erG°/RT

Note that in these equations the equilibrium constant is represented by K°. The standard equilibrium constant, Kº, is either the concentration equilibrium constant (Kc) with each concentration divided by the standard-state concentration of 1 M or the pressure equilibrium constant (Kp) with each pressure divided by the standard-state pressure of 1 bar. Hence, Kº is truly unitless. Dividing by the standard-state concentration or pressure means that if concentrations in Kc are expressed in M (mol/L) the numerical values of Kº and Kc are the same. Similarly, if partial pressures in Kp are expressed in bar, the numerical values of Kº and Kp are the same.

> 1 < 0 Product-favored at equilibrium.
< 1 > 0 Reactant-favored at equilibrium.
= 1 = 0 Reactants and products are equally abundant at equilibrium.

Exercise: Gibbs Free Energy and Equilibrium


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Chemistry 109 Fall 2021 by John Moore, Jia Zhou, and Etienne Garand is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License, except where otherwise noted.