Section 14.3

Thermodynamics Review

Gibbs Free Energy

In the previous section, we discussed that spontaneous reactions always proceed in a direction that will give the products less potential energy, or energy available to do work, than they started with. This means that in a spontaneous chemical reaction, energy is released, and the products of the reaction have less energy than the original reactants. Think of an old-fashioned mill house at a waterfall. The falling water, which has high potential energy, is used to do work (it turns the water wheel, which in turn rotates a grinding stone). This process results in the water falling to the pond below, where its potential energy is much lower.

Let’s look at a hypothetical spontaneous chemical reaction:

The quantity of usable energy (chemical potential) in a reaction is called the Gibbs Free Energy (DG). DG is the difference between the energy contained in the products of a reaction and the reactants:

DG = (energy of products) - (energy of reactants)

Chemical reactions are classified as being either exergonic or endergonic. That just means that a reaction can either release energy useful for work (an exergonic reaction) or requires energy to proceed (an endergonic reaction). The spontaneous reaction above is an exergonic reaction. Note how for an exergonic reaction, DG will be negative. Thus, a negative DG value tells you that that reaction is possible.

Exergonic Reaction
Endergonic Reaction


Not Spontaneous

Release Free Energy ( -DG)

Consume Gibbs Free Energy (+DG)

Products have less energy than reactants

Products have more energy than reactants

For example, here is the reaction when the high energy compound ATP is hydrolyzed to release an inorganic phosphate molecule and energy:

ATP + H2O ADP + Pi         DG = -30.5 kJ/mol

Note that when the reaction releases energy, DG is negative!

When the reaction is written in reverse, the sign of DG changes:

ADP + Pi ATP + H2O         DG = +30.5 kJ/mol

Copyright 2002, John Wiley & Sons Publishers, Inc.