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Endergonic reactions are processes in physical chemistry that require an input of energy to form products whose free energy exceeds that of the reactants. When heat is the primary energy carrier, the reaction is specifically termed endothermic.
These non‑spontaneous transformations demand external energy. A classic biological illustration is photosynthesis, where plants absorb sunlight to convert water and carbon dioxide into glucose—a compound with higher free energy than its starting materials. In chemical terms, the bonds formed in an endergonic reaction are weaker than those broken, making the process energetically unfavorable without an energy source.
Another everyday example is the melting of ice: when solid water absorbs heat, it transitions to liquid water, an endothermic process driven by rising temperatures.
Exergonic reactions are spontaneous; they release energy into their surroundings and generate bonds that are stronger than those broken. The system’s free energy decreases. Common examples include the synthesis of table salt from sodium and chlorine, and chemiluminescent reactions that emit visible light. When heat is released, the reaction is also exothermic.
While “endergonic” and “exergonic” refer to the net change in free energy (ΔG), “endothermic” and “exothermic” refer to the enthalpy change (ΔH). Thus, a reaction can be endergonic yet endothermic, or exergonic yet exothermic. For instance, mixing sodium carbonate with citric acid in water absorbs heat (endothermic) and is endergonic. Conversely, a glow‑stick reaction releases light without significant heat and is exergonic but not exothermic.
In everyday life, you can observe exergonic exothermic reactions during laundry: adding detergent and water produces a mild heat sensation, indicating energy release.
