We know that in order for a reaction to occur, reactant
molecules must collide and that both an increase in the concentration
of reactant molecules and an increase in the temperature of the
system can cause an increase in reaction rate. But it takes more
than just a regular collision to cause a chemical reaction to occur—in
fact, only a very small fraction of collisions that occur in the
solution lead to a reaction. This is true for two reasons. First
of all, for a reaction to occur, the colliding molecules must be
oriented in exactly the correct way: they must be oriented in suitable
way for the product molecule bonds to be formed. Second, the two
molecules must collide with sufficient energy to overcome the activation energy
of the reaction. The activation energy is defined as
the minimum energy needed to initiate a chemical reaction, and it
is symbolized by Ea.
Now let’s talk about the energy diagram below.
This energy diagram is a graph of the progress of a chemical
reaction, versus the total energy of the system. The reactant in
this case is BrNO, and the products are NO and Br2. As
you can see, after the reaction occurs, the energy of the system
is lower than it was before the reaction. This energy diagram shows
an exothermic reaction, one in which energy is given off. In the
energy diagram for an endothermic reaction, the energy of the products
would be higher than that of the reactants.
In this diagram, the activation energy is signified by
the hump in the reaction pathway and is labeled. At the peak of
the activation energy hump, the reactants are in the transition state,
halfway between being reactants and forming products. This state
is also known as an activated complex.
The figure below shows the energy diagram for a reaction
in the presence of a catalyst and in the absence of a catalyst.
As you can see, the catalyst has decreased the activation energy
of the reaction, which means that more molecules are able to surmount
it and react.