Definition of activation energy

Summary

Activation energy is the minimum energy that must be available to reactant molecules for a chemical reaction to occur. In the Arrhenius model, it represents the energy barrier that reactants must overcome to form products, and reactions typically require sufficiently high temperature so that enough molecules have kinetic energy equal to or greater than this barrier. Activation energy is commonly denoted as E_a and is measured in units such as kJ/mol or kcal/mol. The Arrhenius equation links activation energy to how the reaction rate changes with temperature, allowing E_a to be inferred from temperature-dependent rate data. Catalysts lower the activation energy by stabilizing the transition state, increasing reaction rate without changing the energies of reactants and products (so equilibrium is not altered). In transition state theory, the related concept of Gibbs energy of activation (ΔG‡) is used; although Arrhenius and Eyring forms look similar, ΔG‡ includes entropic contributions in addition to enthalpic ones. In some cases, observed activation energies can be negative, often indicating barrierless or multi-step mechanisms where the overall rate decreases with increasing temperature.

Key Points

• 1

  Activation energy (E_a) is the minimum energy barrier reactants must overcome for a reaction to proceed.

• 2

  Arrhenius equation relates E_a to temperature dependence of reaction rates, while catalysts increase rates by lowering E_a via transition-state stabilization.

• 3

  Transition state theory uses Gibbs energy of activation (ΔG‡), which includes both enthalpic and entropic effects; negative E_a can occur for barrierless or certain multi-step mechanisms.