Activation Energy: Arrhenius, Catalysts, and Beyond

Definition and physical meaning of activation energy

Activation energy (Ea) is the minimum energy that must be available to reactant molecules for a chemical reaction to occur, as described by the Arrhenius model of reaction rates. Physically, it represents the energy barrier that reactants must overcome to reach the transition state; only molecules with sufficient kinetic energy (typically requiring higher temperature) can cross this barrier and form products. Ea is commonly reported in units such as kJ/mol or kcal/mol. In the energy-profile picture, the transition state corresponds to the highest-energy point along the reaction coordinate, and Ea is the energy difference between the reactants and that peak. Catalysts lower Ea by stabilizing the transition state (without changing the energies of the reactants or products), which increases reaction rate and does not affect equilibrium. In transition state theory, the closely related quantity is the Gibbs energy of activation (ΔG‡), which includes both enthalpic and entropic contributions; while Arrhenius Ea and ΔG‡ can be similar in magnitude, they are not identical and the overall reaction free-energy change is independent of activation energy. In some cases, observed (apparent) activation energies can be negative, typically when reactions are barrierless or when multistep kinetics produce a temperature-dependent rate decrease.

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Units and typical measurement conventions

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