By Andrea Becker | Updated Aug 30, 2022
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In most introductory chemistry courses, reactions are shown with a single arrow, implying a one‑way process. In reality, chemical reactions are reversible, and the direction they favor depends on the system’s Gibbs free energy (ΔG). By evaluating ΔG, scientists can predict whether a reaction will proceed forward, backward, or reach equilibrium.
Enthalpy represents the total energy stored in a system, largely arising from the random motion of molecules. It is not the same as potential energy of bonds or kinetic energy of bulk motion. Enthalpy changes when heat or work is added or removed, and it is influenced by pressure and volume—especially in gases.
Entropy measures the degree of disorder or randomness in a system. When a system loses heat, such as water freezing into ice, its entropy decreases because the molecules adopt a more ordered arrangement. At the universal scale, entropy never decreases; it always tends to increase.
Both enthalpy and entropy are temperature dependent. Adding heat raises both ΔH and ΔS. The Gibbs free energy change is calculated as ΔG = ΔH – TΔS, where T is the absolute temperature in Kelvin. Because temperature multiplies the entropy term, it can tip the balance between enthalpy and entropy, altering the reaction’s spontaneity.
By examining ΔG, chemists can determine reaction feasibility:
When enthalpy and entropy favor opposite directions, temperature becomes the decisive factor. For example, an endothermic reaction (ΔH > 0) with positive ΔS can become spontaneous at high temperatures because the TΔS term outweighs ΔH.
Understanding Gibbs free energy equips researchers to design efficient processes, predict reaction pathways, and control industrial synthesis.
