Why is there an Activation Energy Barrier?

During the course of a reaction considerable redistribution of electrons may occur. Consider, for example, the reaction of CH3Br with Cl- in water at 298K.

As the bimolecular reaction occurs (i) there is angle bending: the initially pyramidal CH3 grouping become planar; (ii) there is bond-making and breaking: a partial CBr bond is weakened. The energy released by the formation of the partial CCl bond will not fully compensate for the other two (endothermic) changes and yet there is no lower energy pathway from reactants to products. The reactants can get to the point of highest potential energy (the "activated complex" or "transition state" - in curly braces above) only if they initially have sufficient kinetic energy to turn into the potential energy of the activated complex. The activated complex can not be isolated; it is that arrangement of reactants which can proceed to products without further input of energy.

It is often useful to make a schematic plot of the total energy (enthalpy) of the combined reactant molecules during the various stages of the chemical reaction. The points on the plot which we can pinpoint are:

(i) the difference between the average energy of the products and the average energy of the reactants, Hreaction (~ + 25 kJ mol- for CH3Br + Cl- ) and

(ii) the activation energy (obtained experimentally; 103 kJ mol- for CH3Br + Cl-).

If we assume the total energy varies smoothly with the course of the reaction we obtain the following "energy profile":

It is important to note that there will be more than one way for the reactants to interact and so pass to the products. However, there can only be one minimum energy pathway and essentially all of the reaction will occur via this pathway.