What does the term "activation energy" refer to?

Activation energy is defined as the minimum energy that reactants must possess for a chemical reaction to occur. This concept is crucial because it explains why some reactions may proceed readily while others occur very slowly, or not at all, under certain conditions. Activation energy represents the energy barrier that must be overcome for the reactants to transform into products.

In a reaction, even if the reactants have sufficient energy to collide and interact, they may not have enough energy to overcome this barrier. By providing this necessary energy, often from heat or light, the reactants can reach the transition state and proceed to form products.

The other options do not accurately reflect the concept of activation energy. The energy released during a reaction, for example, refers to the energy change that occurs when reactants are converted into products and does not describe the energy required to initiate the reaction. Similarly, the maximum energy of the products pertains to the energy state of the products after the reaction has taken place and is unrelated to the energy needed to start the reaction itself. Lastly, the energy consumed in a phase change relates to transformations of matter, like melting or boiling, rather than the initiation of chemical reactions.