In simple terms, activation energy is the amount of energy required to start a chemical reaction, measured in joules (J) or kilojoules (kJ) per mole (the molecular weight in grams) of reactants. All chemical reactions involve breaking chemical bonds in reactants and making new ones to form products. Activation energy is involved in the bond-breaking process.
Activation energy can be defined as the minimum energy required to form an “activated complex”--a high-energy intermediate between reactants and products--when reactant molecules collide.
Activation energy varies from reaction to reaction. Some elements and compounds react spontaneously, requiring no activation energy, while others require heat energy to be supplied before they react.
Types of Reaction
Chemical reactions can be fast or slow. In a fast reaction, many of the reactant molecules are moving fast enough to form an activated complex when a collision occurs, while in a slow reaction few are moving fast enough and most collisions do not produce a reaction.