What is activation energy in a chemical reaction?

Activation energy is defined as the minimum energy required for a chemical reaction to occur. It is the energy barrier that reactants must overcome for the transition to products. This energy is essential because it helps to initiate the reaction process, allowing the reactant particles to collide with sufficient energy to break bonds and form new ones.

While the other options mention related concepts, they do not encapsulate the full definition of activation energy. It is specifically about the energy threshold that reactant particles need to surpass for a reaction to be initiated. This understanding is crucial in the context of reaction kinetics, where different reactions have varying activation energies, influencing their rates. High activation energy typically leads to slower reactions, as fewer particles will have enough energy to meet the required threshold when they collide. In contrast, a lower activation energy can result in a faster reaction because more molecules will have the necessary energy to react.