What is known as the energy required to start a chemical reaction?

The energy required to start a chemical reaction is called activation energy. It is the minimum energy reactant particles need to reach a high-energy transition state, where existing bonds can break and new bonds can begin forming.

Activation energy explains why many reactions do not happen instantly, even when the products are energetically favored. Heating a substance increases the number of particles with enough kinetic energy to cross the reaction barrier, usually making the reaction faster. A spark, such as the one used to ignite a gas burner, supplies an initial burst of activation energy.

Catalysts and enzymes speed reactions by providing an alternative pathway with a lower activation energy. They are not used up overall and do not change the reaction’s final energy difference or chemical equilibrium. This is why a catalyst can accelerate both forward and reverse reactions without changing which side is ultimately favored.

The Swedish chemist Svante Arrhenius introduced the term in 1889. Activation energy is commonly expressed in kilojoules per mole, and its effect on reaction rate is described quantitatively by the Arrhenius equation.