Activation Energy

Activation energy, $E_a$, is the minimum extra energy, over and above the average

energy the reactant molecules already possess, that colliding molecules must have in order to react --

it is the energy barrier that must be crossed for reactants to be converted into products.

The potential-energy profile. Activation energy is most easily pictured on a graph plotting the

potential energy of the reacting system (on the vertical axis) against the progress of the reaction,

often called the reaction coordinate (on the horizontal axis). Reactants start on the left of this

diagram at their own characteristic energy level. As the reaction proceeds, the system's potential

energy rises to a maximum at a point called the activated complex (or transition state) -- an

unstable, high-energy, fleeting arrangement of atoms in which old bonds are partway to breaking and new

bonds are partway to forming, existing for only the briefest instant before either falling forward to

products or reverting back to reactants. The energy difference between the reactants and this peak is

the (forward) activation energy, $E_a$; only collisions energetic enough to reach this peak can

successfully cross over to become products. Beyond the peak, the potential energy falls again to the

level of the products, on the right of the diagram.

Reading $\Delta H$ off the same diagram. The vertical difference between the reactants' energy

level and the products' energy level is the reaction's enthalpy change, $\Delta H$. If the products lie

below the reactants on the diagram, the reaction is exothermic ($\Delta H < 0$): energy is

released overall, even though energy must still be temporarily invested to climb over the activation

barrier on the way. If the products lie above the reactants, the reaction is endothermic

($\Delta H > 0$). The diagram also defines a reverse activation energy, $E_{a,\text{reverse}}$ -- the

energy barrier the products would need to climb to revert back to reactants -- and the two activation

energies are related to the reaction enthalpy by $E_{a,\text{forward}} - E_{a,\text{reverse}} = \Delta H$. For an exothermic reaction, the reverse activation energy is always larger than the forward one

(the barrier looks steeper when climbed from the product side).

Why lower $E_a$ means faster reaction. Because only molecules with energy exceeding $E_a$ can react

(the collision-theory requirement of the previous section), a lower activation-energy barrier means a …