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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.
This is called the
activation energy.



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

 

 

 

 

 

 

 

 

A highly effective catalyst can reduce the activation energy so much that the majority of the molecules can react.

 

 

 

 

 

 

 

 

If the activation energy is reduced to Ea2 by the use of a catalyst, a higher proportion of molecules will have sufficient energy to react.
From a reaction which proceeds very slowly we have gone to a reaction, which at the same temperature, is almost instantaneous.

Some enzymes can achieve this, increasing rates by up to 1020 times.