Catalysts increase the rate of a reaction without being used up in the process.
In order for a reaction to occur bonds in the reactants must be broken
- this requires energy, and new bonds must be made, releasing energy.
Because of the necessary bond breaking a certain minimum energy is needed
before reaction is possible, even for an exothermic reaction.
Catalysts reduce the activation energy required for reaction.
For example introducing a small amount of a catalyst into the reaction mixture for the reaction
H2(g) + I2(g) 2HI(g)
increases the rate by a factor of 1000. The catalyst causes this increase in rate by reducing the activation energy of the reaction, in this case by about 40 kJ mol-1.
Why does a moderate reduction in activation energy cause such a big increase in rate? The energy for reaction comes from the kinetic energy with which molecules collide. The figure shows the kinetic energy distribution for a gas mixture. A typical activation energy for an uncatalysed reaction is shown as Ea1 on the diagram. Only those molecules which collide with a kinetic energy greater than Ea1 can react. These are shaded green on the diagram, and it is a very small fraction of the molecules.
Effect on Rate