Activation energy is the minimum amount of energy that reactants must possess in order for a chemical reaction to occur. It can be thought of as the energy barrier that must be overcome for the reactants to be transformed into products. This energy is required to break bonds in the reactants so that new bonds can form in the products.
Activation energy is often denoted by the symbol \(E_a\) and is a key concept in chemical kinetics. It influences the rate of a reaction; reactions with higher activation energies tend to proceed more slowly than those with lower activation energies, assuming other conditions are constant. The concept of activation energy is also closely related to the Arrhenius equation, which describes how the rate constant of a reaction varies with temperature and activation energy.
In practical terms, factors that can affect activation energy include temperature, catalysts, and the nature of the reactants. Catalysts, for example, lower the activation energy, providing an alternative pathway for the reaction that requires less energy to occur.