A substance which increases the speed of a reaction without being consumed in the reaction is called a catalyst. The phenomenon of increasing the rate of reaction by the use of catalyst is called catalysis. Living cells contain thousands of different kinds of proteins called enzymes which act as catalysts.
A catalyst provides an entirely new path for the reaction in which the reactants are converted to the products quickly. A chemical reaction takes place by a reaction path, first converted to activated state and then finally to the products.
A catalyst forms a new activated complex of lower potential energy. The catalyst provides a new pathway of lower activation energy. The fraction of the total number of collisions possessing lower activation energy is increased and hence, the rate of reaction also
The following points should be kept in mind regarding the function of a catalyst :
1) A catalyst may undergo intermediate physical changes and it may even form temporary chemical bonds with the reactants but it is recovered unchanged in original form at the end of the reaction.
2) A catalyst speeds up the reaction, but it does not shift the position of equilibrium.The presence of a catalyst reduces the height of barrier by providing an alternative path for the reaction and lowers the activation energy. However, the lowering in activation energy
is to the same extent for the forward as well as for the backward reaction.
As a result, the increase in the rate of the forward and backward reactions is same and hence, the position of equilibrium remains unaltered. However, by increasing the two rates, the equilibrium is attained earlier.
3) A catalyst which can catalyse one reaction may have no effect on another reaction even, if that reaction is very similar.
4) The catalyst does not change ΔE of the reaction. The addition of a catalyst does not change the energies of reactants (Er) and products (Ep) so that ΔE (Ep– Er) remains same.
Promoters and Poisons
Promoters are substances that enhance the activity of the catalyst.
For example: In Haber’s process for the manufacture of ammonia, molybdenum is used as promoter for iron, which acts as a catalyst.
N2 (g) + 3H2 (g) ⇔ 2 NH3 (g)
The substances which decrease the activity of the catalyst are called catalytic poisons.
Classification of Catalysts
The catalysts may be of two main types :
1. Homogeneous catalysts
2. Heterogeneous catalysts
1. Homogeneous catalysts
When the catalyst is present in the same phase as the reactants and products, it is called homogeneous catalyst.
In these reactions, the catalysts, reactants and products are in the same phase and are called homogeneous catalytic reactions.
(1) In lead chamber process, SO2 , is oxidised to SO3 , in the presence of nitric oxide as catalyst :
2 SO2 (g) + 2 O2 (g) —–> 2 SO3 (g)
The reactants sulphur dioxide and oxygen and the catalyst nitric oxide all are in the same phase.
(2) Catalytic decomposition of ozone occurs by chlorine atoms in the gas phase :
O3 + O ——-> 2O2
(3) Carbon monoxide is oxidised by O2 , in the presence of nitric oxide (NO) as catalyst :
2CO (g) + O2 (g) ——> 2 CO2 (g)
(4) Hydrolysis of sucrose is catalysed by the presence of dil. HCl or H2SO4
C12H22O11 + H2O ——> C6H12O6 (aq) + C6H12O6 (aq)
(5) Preparation of diethyl ether from ethyl alcohol using conc. H2SO4, at 413 K.
2 CH3CH2OH ——> CH3CH2OCH2CH3 + H2O
Mechanism of Homogeneous Catalytic Reactions
The catalyst combines with one of the reactants to form an intermediate. Intermediate compound being unstable either decomposes or combines with the other reactant to form the product and the catalyst is regenerated.
For example: the combination of SO2 and O2 to form SO3 is a slow process. However, in the presence of NO (catalyst) the reaction becomes fast.
2 SO2 (g) + O2 (g) ———-> 2 SO3 (g)
In this reaction nitric oxide combines with one of the reactants to form intermediate compound (NO2). This intermediate (NO2) combines readily with SO2 to form SO3 and the catalyst NO is regenerated in the last step:
1) 2NO+ O2 ——> 2 NO2 (fast)
2) NO2 + SO2 ——> SO3 + No (fast)
2SO2 + O2 ——> 2 SO3 (fast)
When the catalyst is in different phase than the reactants, it is called heterogeneous catalyst. Such reactions are called heterogeneous catalytic reactions.
In heterogeneous catalysis, catalyst is generally a solid and the reactants are generally gases. This is also known as surface catalysis because the reaction starts at the surface of the solid catalyst. These catalysts have very large surface area of the order of 1 to 500 m2 per gram for contact.
Thus, despite an enormous surface area, once the reactant gas molecules cover the surface, the rate does not increase on increasing the reactant concentrations.
Some other examples are:
(1) Manufacture of NH3 from H2 and N2 by Haber’s process using finely divided iron catalyst.
3H2 (g) + N2 (g) —–> 2 NH3 (g)
Here reactants are in the gaseous state while the catalyst is in the solid state.
(2) Methanol is prepared from CO and H2 by using a mixture of copper, zine oxide as catalyst and Cr2O3 as promoter:
CO (g) + 2 H2 (g) ——-> CH3OH (l)
(3) Manufacture of SO3 from SO2 in the Contact process using platinised asbestos or V2O5 as catalyst :
2 SO2 + O2 ——-> 2 SO3
(4) Dehydrogenation of ethanol by using nickel catalyst:
CH3CH2OH ——> CH3CHO + H2
(4) Oxidation of ammonia into nitric oxide in the presence of platinum gauze in Ostwald’s process:
4NH3 (g) + 5 O2 ——> 4 NO (g) + 6 H2 O (g)
(5) Hydrogenation of vegetable oils in the presence of finely divided nickel as catalyst is also an example of heterogeneous catalytic reaction because one of the reactants is in liquid state and the other in gaseous state, while the catalyst is in the solid state.
Vegetable oils (l) + H2 (g) ——-> Vegetable ghee (s)