Final answer:
The higher the activation energy of a chemical reaction, the fewer molecules have sufficient energy to overcome this barrier, making it more difficult for the reaction to proceed.
Step-by-step explanation:
When considering the activation energy of a chemical reaction, the higher the energy barrier, or the larger the activation energy, the fewer molecules that will have sufficient energy to overcome it. Activation energy is the minimum amount of energy needed for a reaction to occur. It represents an energy threshold that reacting molecules must surpass to be able to effectively collide and form a new product, thus requiring molecules to have enough kinetic energy to overcome electrostatic repulsion and break existing chemical bonds.
For example, if the dissociation energy were larger, it would be more difficult to break the solid apart because more energy would be needed to overcome the interactions holding the particles together.
An increase in temperature usually results in a higher fraction of molecules acquiring enough energy to surpass this barrier. Therefore, reactions generally proceed faster at higher temperatures due to the increased number of molecules with energy equal to or greater than the activation energy.