Endergonic and Exergonic Reactions in the Cell

Exergonic reactions release free energy, producing a negative free energy change (ΔG < 0), and occur spontaneously without added energy. Endergonic reactions require energy input, producing a positive free energy change (ΔG > 0), and do not occur spontaneously. Cells harness exergonic reactions to perform work and couple them with endergonic reactions to drive essential processes.

Cells couple endergonic reactions with exergonic reactions to make them thermodynamically favorable. For example, the endergonic conversion of glucose to glucose-6-phosphate (ΔG° = 13.8 kJ/mol) is coupled with the exergonic hydrolysis of ATP (ΔG° = −30.5 kJ/mol), resulting in an overall exergonic process with ΔG° = −16.7 kJ/mol that proceeds spontaneously.

A negative free energy change (ΔG < 0) indicates that a reaction is exergonic and spontaneous. The products have less free energy than the reactants because energy was released during the reaction. This negative value means the reaction can occur without external energy input and can be harnessed to perform cellular work.

Anabolic processes build complex molecules from simpler ones, requiring energy input to form new chemical bonds. These biosynthetic reactions produce products with higher free energy than reactants, resulting in a positive ΔG. Examples include synthesizing sugars and proteins, which are energy-storing molecules essential for cellular structure and function.

A spontaneous reaction has a negative free energy change and can occur without adding energy to the system. Contrary to everyday usage, spontaneous does not mean sudden or fast; rusting iron is a spontaneous reaction that occurs slowly over time. Spontaneity depends on thermodynamics, not reaction rate.

Catabolic reactions break down complex molecules into simpler ones, releasing energy in a series of exergonic reactions with negative ΔG values. Anabolic reactions build complex molecules from simpler ones, requiring energy input and producing endergonic reactions with positive ΔG values. Together, these opposing processes regulate cellular energy balance.

Standard free energy change (ΔG°) reports the free energy change of a reaction under standardized conditions, allowing comparison across different reactions. For example, the hydration of fumarate to malate has ΔG° = −3.8 kJ/mol, indicating it is exergonic. This standardization helps predict reaction spontaneity and cellular energy availability.