Which statement about activation energy is true?

Activation energy refers to the minimum amount of energy that reacting particles must possess in order to collide successfully and trigger a chemical reaction. This energy is crucial for overcoming the energy barrier associated with breaking existing chemical bonds in reactants so that new bonds can form during the reaction process.

When a reaction occurs, the molecules must go through a transition state that requires energy input to break the bonds in the reactants. Therefore, the statement highlighting that activation energy is necessary for breaking chemical bonds is accurate. It illustrates the concept that without sufficient energy, reactions would not proceed, as the necessary bonds cannot be broken.

The other statements do not accurately describe activation energy. For example, it is not a constant for every reaction, as different reactions can have varying activation energies depending on the nature of the reactants and the overall reaction pathway. Additionally, while activation energy can appear to decrease with an increase in temperature, this is more about the fraction of molecules that have energy exceeding the activation barrier, rather than a change in the activation energy itself. Lastly, activation energy certainly affects the reaction rate; a higher activation energy typically correlates with a slower reaction rate, since fewer molecules will have sufficient energy to overcome this barrier, thus highlighting its significance in kinetics.