Which statement best describes activation energy?

Activation energy is defined as the minimum energy required for reactants to collide effectively and form products in a chemical reaction. This concept is crucial because, in order for a reaction to proceed, the molecules must overcome an energy threshold that allows them to break existing bonds and form new ones.

When reactants possess energy equal to or greater than the activation energy, they can transition into a higher energy state, known as the transition state, which facilitates the formation of products. This is why the statement that describes activation energy as the minimum energy needed for a reaction to occur is accurate.

The other statements do not align with the definition of activation energy. For instance, sustaining a reaction relates more to the overall energy dynamics rather than the initial requirement to start a reaction. Breaking chemical bonds is part of the reaction process but does not specifically define the energy threshold needed to initiate the reaction. Finally, the energy released during an exothermic reaction refers to the energy change that occurs after the reaction has taken place, not the initial energy required to start the reaction.