Harper's Illustrated Biochemistry, Volume 1 - Murray R. 1993

Bioenergetics and Metabolism of Carbohydrates and Lipids
Bioenergetics
Free Energy and the Laws of Thermodynamics

Free energy change (∆G) represents that portion of a system's internal energy change that can be converted into work; in other words, it is useful energy, referred to in chemical systems as chemical potential.

The First Law of Thermodynamics states that the internal energy of a system combined with its surroundings remains constant. This is one of the formulations of the law of conservation of energy, asserting that during any changes within a system, internal energy is neither lost nor gained. At the same time, energy within the system under consideration may be transferred from one part to another or transformed from one form into another. For instance, chemical energy can be converted into heat, electrical energy, radiant energy, or mechanical energy.

The Second Law of thermodynamics states that the Entropy of a system increases during spontaneous processes. Entropy serves as a measure of disorder and randomness in a system, reaching a maximum when the system attains true equilibrium. At constant Temperature and pressure, the relationship between The change in the system's free energy (∆G) and the entropy change (∆S) is expressed by the following equation, which combines both Laws of Thermodynamics:

Class="center">∆G = ∆Н - T∆S,

where ∆Н is the enthalpy change (heat) and T is the absolute temperature.

Under the conditions where biochemical reactions take place, ∆Н is approximately equal to ∆Е, the change in the system's internal energy resulting from the reaction. Under these conditions, the expression given above can be written as

∆G = ∆Е — T∆S.

If ∆G is negative, the reaction proceeds spontaneously and is accompanied by a decrease in free energy. Such reactions are termed exergonic. Moreover, if ∆G is large in absolute value, the reaction goes virtually to completion and can be considered irreversible. Conversely, if ∆G is positive, the reaction will only proceed upon the input of external free energy; such a reaction is called endergonic. If ∆G is large in this case as well, the system is stable and the reaction practically does not take place. When ∆G is equal to zero, the system is at equilibrium.

At reactant concentrations of 1.0 mol/L, the free energy change is designated as ∆G°, the Standard Free Energy change. For biochemical reactions, the standard state is defined at pH 7.0. The standard free energy change under such standard conditions is denoted as ∆G°'.

The standard free energy change can be determined by knowing the Equilibrium Constant K'_eq:

∆G°' = - 2,303 RT log К'_eq

where R is the gas constant and T is the absolute temperature (see p. 84). It is important to note that ∆G may be greater or less than ∆G°' depending on the concentrations of the reactants.

It should be clearly understood that in the presence of Enzymes, biochemical reactions merely reach equilibrium faster, while the final equilibrium concentrations of the reactants remain unchanged.