This figure is sometimes referred to as a potential energy diagram. The potential
energy of the activated complex is greater than that of the reactants, and the
energy difference between the two is called the activation energy (symbolized as
Ea). The activation energy is the minimum energy that must be supplied for the
activated complex to be formed. It is an energy barrier that must be overcome by
reactant molecules; reactant molecules acquire the necessary activation energy
by absorbing heat from the surroundings.
The size of the activation energy barrier indicates how difficult it is for the
reaction to proceed. A relatively small barrier, indicating a low activation energy,
means that collisions between reactant molecules need less energy to produce an
activated complex. Thus, a greater percentage of reactant collisions is likely to
lead to product formation, resulting in a relatively high rate of reaction.
Let’s return to the question How do catalysts work? Catalysts increase the
reaction rate by enabling the reaction to proceed through a series of different
steps with a lower activation energy than they ordinarily would require. So a
catalyst reduces the minimum energy requirement of the reaction. This leads to a
greater percentage of product-forming collisions and thus an increased rate of
reaction. Since a catalyst is not consumed by the process, a small amount of
catalyst can be used to speed up a reaction with a fairly large reactant
concentration.
Catalyst Facts
Catalysts increase the rate of a reaction by lowering the activation energy.Catalysts are not consumed in a reaction.Catalysts do not change the equilibrium of a reaction.The rate of reaction depends on both the frequency of molecular collision and
the energy of molecular collision.