Endergonic and exergonic reactions are defined according the change in Gibbs free energy. In an endergonic reaction, the free energy of the products is higher than the free energy of the reactants ((∆G > 0; energy is stored in the products), so the reaction is not spontaneous and additional energy must be supplied to make the reaction proceed. In an exergonic reaction, the free energy of the reactants is higher than the free energy of the products (∆G < 0). Energy is released to the environment, which overcomes the activation energy of the reaction and makes it spontaneous.
Here is a closer look at endergonic and exergonic reactions, examples of each types, and how the reactions are coupled to force unfavorable reactions to occur.
An endergonic reaction is a chemical reaction with a positive standard Gibbs free energy, at constant temperature and pressure:
∆G° > 0
In other words, there is a net absorption of free energy. Chemical bonds in the products store energy. Endergonic reactions are also called unfavorable or nonspontaneous reactions because the activation energy for an endergonic reaction usually is larger than the energy of the overall reaction. Because Gibbs free energy relates to the equilibrium constant, K < 1.
There are several ways to make unfavorable reactions proceed. You can supply energy by heating the reaction, couple it to an exergonic reaction, or making it share an intermediate with a favorable reaction. You can pull the reaction to proceed by removing the product from the system.
Examples of endergonic reactions include photosynthesis, the Na+/K+ pump for muscle contraction and nerve conduction, protein synthesis, and dissolving potassium chloride in water.
An exergonic reaction is a chemical reaction with a negative standard Gibbs free energy, at constant temperature and pressure:
∆G° < 0
In other words, there is a net release of free energy. Breaking chemical bonds in the reactants releases more energy than that used to form new chemical bonds in the products. Exergonic reactions are also known as exoergic, favorable, or spontaneous reactions. As with all reactions, there is an activation energy which must be supplied for an exergonic reaction to proceed. But, the energy released by the reaction is enough to meet the activation energy and keep the reaction going. Note that while an exergonic reaction is spontaneous, it may not proceed quickly without the aid of a catalyst. For example, the rusting of iron is exergonic, but very slow.
Examples of exergonic reactions include cellular respiration, the decomposition of hydrogen peroxide, and combustion.
Endergonic/Exergonic vs Endothermic/Exothermic
Endothermic and exothermic reactions are types of endergonic and exergonic reactions, respectively. The difference is the the energy absorbed by an endothermic reaction or released by an exothermic reaction is heat. Endergonic and exergonic reactions may release other kinds of energy besides heat, such as light or even sound. For example, a glow stick is an exergonic reaction that releases light. It is not an exothermic reaction because it does not release heat.
Forward and Reverse Reactions
If a reaction is endergonic in one direction, it is exergonic in the other direction (and vice versa). For this reaction, endergonic and exergonic reactions may be called reversible reactions. The amount of free energy is the same for both the forward and reverse reaction, but the energy is absorbed (positive) by the endergonic reaction and released (negative) by the exergonic reaction. For example, consider the synthesis and degradation of adenosine triphosphate (ATP).
ATP is made by joining a phosphate (Pi) to adenosine disphosphate (ADP):
ADP + Pi → ATP + H2O
This reaction is endergonic, with ∆G = +7.3 kcal/mol under standard conditions. The reverse process, the hydrolysis of ATP, is an exergonic process with a Gibbs free energy value equal in magnitude, but opposite in sign of -7.3 kcal/mol:
ATP + H2O → ADP + Pi
Coupling Endergonic and Exergonic Reactions
Chemical reactions proceed in both the forward and reverse direction until chemical equilibrium is reached and the forward and reverse reactions proceed at the same rate. At chemical equilibrium, the system is in its most stable energy state.
Equilibrium is bad news for biochemistry, because cells need metabolic reactions to occur or else they die. Cells control the concentration of products and reactants to favor the direction of the reaction needed at the time. So, for a cell to make ATP, it needs to supply energy and add ADP or remove ATP and water. To continue converting ATP to energy, the cell supplies reactants or removes products.
Often, one chemical reaction feeds the next and endergonic reactions are coupled to exergonic reactions to give them enough energy to proceed. For example, firefly bioluminescence results from endergonic luminescence by luciferin, coupled with exergonic ATP release.
- Hamori, Eugene (2002). “Building a foundation for bioenergetics.” Biochemistry and Molecular Biology Education. 30 (5):296-302. doi:10.1002/bmb.2002.494030050124
- Hamori, Eugene; James E. Muldrey (1984). “Use of the word “eager” instead of “spontaneous” for the description of exergonic reactions”. Journal of Chemical Education. 61 (8): 710. doi:10.1021/ed061p710
- IUPAC (1997). Compendium of Chemical Terminology (2nd ed.) (the “Gold Book”). ISBN 0-9678550-9-8. doi:10.1351/goldbook