# Endergonic reaction

An endergonic reaction (such as photosynthesis) is a reaction that requires energy to be driven. Endergonic (from the prefix endo-, derived from the Greek word ἔνδον endon, "within", and the Greek word ἔργον ergon, "work") means "absorbing energy in the form of work." The activation energy for the reaction is typically larger than the overall energy of the exergonic reaction(1). Endergonic reactions are nonspontaneous. The progress of the reaction is shown by the line. The change of Gibbs free energy (ΔG) during an endergonic reaction is a positive value because energy is gained (2).

In chemical thermodynamics, an endergonic reaction (also called a heat absorb nonspontaneous reaction or an unfavorable reaction) is a chemical reaction in which the standard change in free energy is positive, and energy is absorbed. In layman's terms, the total amount of energy is a loss (it takes more energy to start the reaction than what you get out of it) so the total energy is a negative net result. For an overall gain in the net result see exergonic reaction. Another way to phrase this is that energy is absorbed from the surroundings into the workable system.

Under constant temperature and constant pressure conditions, this means that the change in the standard Gibbs free energy would be positive may be negative in some cases

${\displaystyle \Delta G^{\circ }>0}$

for the reaction at standard state (i.e. at standard pressure (1 bar), and standard concentrations (1 molar) of all the reagents).

In metabolism, an endergonic process is anabolic, meaning that energy is stored; in many such anabolic processes energy is supplied by coupling the reaction to adenosine triphosphate (ATP) and consequently resulting in a high energy, negatively charged organic phosphate and positive adenosine diphosphate.

## Equilibrium constant

The equilibrium constant for the reaction is related to ΔG° by the relation:

${\displaystyle K=e^{-{\frac {\Delta G^{\circ }}{RT}}}}$

where T is the absolute temperature and R is the gas constant. A positive value of ΔG° therefore implies

${\displaystyle K<1\,}$

so that starting from molar stoichiometric quantities such a reaction would move backwards toward equilibrium, not forwards.

Nevertheless, endergonic reactions are quite common in nature, especially in biochemistry and physiology. Examples of endergonic reactions in cells include protein synthesis, and the Na+/K+ pump which drives nerve conduction and muscle contraction.