In terms of chemical reactions, what does "activation energy" refer to?

Activation energy refers to the minimum amount of energy that reactants need to possess in order to successfully collide and form products in a chemical reaction. This concept is key in understanding how chemical reactions proceed, as reactions do not occur spontaneously without overcoming this energy barrier. When reactants collide with enough energy that meets or exceeds the activation energy, they can break bonds and rearrange to form new substances.

The other options do not accurately describe activation energy. Total energy released during a reaction relates to the overall energy change from reactants to products, rather than the energy needed to start the reaction. The energy remaining after a reaction does not apply to activation energy, as it is specifically concerned with the energy needed to initiate the reaction, not what remains after it is completed. Lastly, while catalysts can affect the activation energy required for a reaction, the energy consumed by catalysts is not the same as the activation energy itself; catalysts lower the activation energy but are not directly related to the energy that reactants initially need to start a reaction.