What is activation energy in a chemical reaction?

Activation energy is defined as the minimum amount of energy required for reactants to undergo a chemical reaction and transition into products. This energy barrier must be overcome for the reaction to proceed, allowing the reactants to reach the transition state, where molecular bonds are breaking and forming. When this minimum energy level is met, the reaction can proceed, leading to the formation of products.

Understanding this concept is crucial because it illustrates why some reactions occur rapidly under certain conditions while others may require catalysts or increased temperatures to provide the necessary energy for the reaction to take place. This is different from the energy released during the formation of products, which refers to the exothermic nature of a reaction, and does not represent the energy required to initiate the reaction itself. Similarly, the overall energy change and the energy stored in products pertain to the thermodynamics of the reaction but do not define the activation energy, which focuses specifically on the energy needed to get the reaction started.