Activation energy is defined as:

• A. The energy barrier that must be overcome for a reaction to proceed.
• B. The energy released when products form.
• C. The energy required to reach equilibrium.
• D. The minimum energy required to start a chemical reaction.

Answer: (A)

Activation energy is the energy barrier that must be overcome for a reaction to proceed. It represents the difference between the energy of the reactants and the high-energy transition state along the reaction path. Only when molecules gain enough energy to reach that transition state can bonds break and new bonds form to create products. This barrier is what makes many reactions slow at room temperature and why increasing temperature speeds them up, as described by the Arrhenius relationship. The energy released when products form is the overall energy change of the reaction (enthalpy change), not the barrier to reaction. Equilibrium involves equal forward and reverse rates, not crossing a fixed energy threshold, so it’s not about an energy barrier. The correct concept, then, is the energy barrier that must be overcome to reach the transition state.