Collision Theory, Transition States, and Energy

Profiles

In order for a reaction to occur, molecules must collide with each other. The collision theory of
chemical kinetics states that the rate of a reaction is proportional to the number of collisions per
second between the reacting molecules.

Not all collisions, however, result in a chemical reaction. An effective collision (one that leads to the
formation of products) occurs only if the molecules collide with correct orientation and sufficient
force to break the existing bonds and form new ones. The minimum energy of collision necessary for
a reaction to take place is called the activation energy, or the energy barrier, designated Ea. Only a
fraction of colliding particles have enough kinetic energy to exceed the activation energy. This
means that only a fraction of all collisions are effective.

When molecules collide with sufficient energy, they go through what is known as a transition state
(also called the activated complex), in which the old bonds are weakened and the new bonds are
beginning to form. The transition state then dissociates into products and the new bonds are fully
formed. For example, in a reaction A 2 + B 2 → 2AB, the transition state or activated complex may
look like the middle species in the following diagram, where the dashed lines represent partial
bonds (bonds that are not quite as strong as a single bond):

The activated complex exists only in one fleeting instant in time and is thus only a snapshot of how
the molecules are arranged somewhere along the reaction; it is not a stable isolatable species by
itself.

A potential energy diagram is very helpful in visualizing the progress of a reaction. The x-axis in
these diagrams corresponds to the “reaction coordinate,” which essentially measures how far along
one is in a reaction by charting the progress from reactants to products: As one moves from left to