What is activation energy in chemical reactions?

Activation energy is a crucial concept in understanding how chemical reactions proceed. It refers specifically to the minimum amount of energy that reactant particles must possess in order to successfully collide in a way that causes a chemical reaction. When reactants collide, they often need to overcome an energy barrier before they can rearrange themselves into products. This energy barrier is what activation energy represents.

In a typical reaction, molecules must collide with enough energy to break existing bonds and form new ones. If the colliding molecules do not have sufficient energy, they will simply bounce off each other without reacting. Therefore, the activation energy sets the threshold that needs to be reached for a reaction to occur.

The other choices do not accurately describe activation energy. The maximum energy required for a reaction to occur refers to the peak of the energy profile but does not define the energy needed specifically for a collision that initiates a reaction. The energy associated with the formation of products relates to the energy changes occurring after the reaction has taken place, rather than the energy required to start it. Lastly, the energy released when bonds are formed refers to exothermic reactions, which is the opposite of activation energy, highlighting the energy change when new bonds are created after the reaction. Thus, option C captures the