(i) The catalyst remains unchanged in amount and chemical composition at the end of the reaction; it
may, however, undergo considerable change in physical form.

(ii) A small quantity of the catalyst is capable of producing the desired effect.

(iii) The action of a catalyst is specific to a large extent. Thus, the decomposition of `KClO_3` is catalysed by `MnO_2` but not by platinum.

(iv) The catalyst does not initiate a reaction; it merely accelerates the reaction that is already occurring.

(v) A catalyst does not alter the final state of equilibrium in a reversible reaction.

A certain minimum energy must be possessed by the reactants so that they may react and produce the
products. This is called the activation energy `(E_a)` for the reaction. A catalyst is said to lower the activation
energy and thus increase the rate of the reaction.

`text(Thus, a catalyst increases the rate of a reaction by providing a pathway whose activation energy is lower than the activation energy of the uncatalyzed reaction)`.

(i) Positive catalysis : The catalyst increases the rate of a reaction.

Examples: `2SO_2+O_2 overset(V_2O_5)-> 2SO_3` (Contact Process)

`C_2H_4 + H_2 underset(Delta) overset(N i) rightarrow C_2H_6` (Ethane)

(ii) Negative catalysis: (Inhibitor or retarder) : Chemical reactions are sometimes retarded by the presence of a foreign substance. The substance is known as a negative catalyst.

Examples: Alcohol, Acetanilide : Prevents oxidation of `Na_2SO_3` by air

`H_3PO_4` : Prevents decomposition of `H_2O_2`

(iii) Auto catalysis: In this type of catalysis, one of the product of the reaction catalyses the reaction. In the oxidation of oxalic acid by `KMnO_4`, `Mn^(2+)` ion formed is known to accelerate the reaction. So, when `KMnO_4` solution is run into warm solution of oxalic
acid (+ dil. ` H_2SO_4`), initially there is a time lag before decolourisation occurs; as more `KMnO_4` is added, the decolourisation becomes almost instantaneous.

(iv) Homogeneous catalysis: A catalytic process in which the catalyst is in the same phase as the reactant is called homogeneous catalysis.

`2SO_2(g)+O_2(g) overset(NO) -> 2SO_3(g)` (Lead chamber process)

`C_12H_22O_11(aq)+H_2O overset(H^+ (aq))-> C_6H_12O_6(aq)+C_6H_12O_6(aq)` (Inversion of cane sugar)

(v) Heterogeneous catalysis: A catalytic process in which the catalyst and the reactants are in different phases is called heterogeneous catalysis. This process is also called contact or surface catalysis.

Examples: (a) `2H_2O_2( l) overset(Pt)-> 2H_2O(l) + O_2(g) `

(b) `2SO_2(g) + O_2(g) overset( Pt quad text(asbestos)) rightarrow 2SO_3(g)`

(c) `N_2(g) + 3H_2(g) overset( Fe-Mo) rightarrow 2NH_3(g)`

(d) `4NH_3(g) + 5O_2(g) overset(Pt quad text(gauze)) rightarrow 4NO(g)+ 6H_2O(l)`

(e) `CO(g) + 2H_2(g) overset(ZnO + Cr_2O_3) rightarrow CH_3OH(l)`


(vi) Induced catalyst: When one reaction influences the rate of other reaction, which does not occur under ordinary conditions, the phenomenon is known as induced catalysis.

Examples of induced catalysis:

(a) Sodium arsenite solution is not oxidised by air. If, however, air is passed through a mixture of the solution of sodium arsenite and sodium sulphite, both of them undergo simultaneous oxidation. The oxidation of sodium sulphite, thus, induces the oxidation of sodium arsenite.

(b) The reduction of mercuric chloride (`HgCl_2`) with oxalic acid is very slow, but potassium permanganate is reduced readily with oxalic acid. If, however, oxalic acid is added to a mixture of `KMnO_4` and `HgCl_2`, both are reduced simultaneously. The reduction of potassium permanganate, thus, induces the reduction of `HgCl_2`.

It is not possible to give a uniform explanation of the mechanism of the phenomenon of catalysis as catalytic reactions are of varied nature. However, two broad theories of catalytic action have been proposed. First theory, known as intermediate compound formation theory explains successfully the homogeneous catalysis. The second theory termed as adsorption theory explains the heterogeneous catalysis.

According to this theory, the catalyst first forms an intermediate compound with one of the reactants. The intermediate compound is formed with less energy consumption than needed for the actual reaction. The intermediate compound being unstable combines with other reactant to form the desired product and the catalyst is regenerated.

Consider, a reaction of the type `A + B ⇋ AB` which occurs in presence of a catalyst `C`, may take place as

`A + undersettext(Catalyst)C ⇋ undersettext(Intermediate Compound)(AC)` (slow reaction)

`AC + B ⇋ undersettext(Product)(AB) + undersettext(Catalyst)C` (fast reaction)

Rate = `k[A] [C]`

Many catalytic reactions can be explained on the basis of this theory. Consider the catalytic oxidation of `SO_2` to `SO_3` in the lead
chamber process. This occurs as follows

`undersettext(Catalyst)(2NO) + O_2 -> undersettext( Intermediate product)(2NO_2)`

`NO_2 + SO_2 -> undersettext(Product)(SO_3) + undersettext(Catalyst)(NO)`

This theory explain why the catalyst remains unchanged in mass and chemical composition at the end of the reaction and is effective even in small quantities. The scope of this theory is, however, limited, as the formation of intermediate compound is possible in the case of homogenous catalysis only. It also fails to explain the action of catalytic promoters, catalytic poisons and action of finely divided catalysts.

This theory explains the mechanism of heterogeneous catalysis. The old point of view was that when the catalyst is in solid state and the reactants are in gaseous state or in solutions, the molecules of the reactants are adsorbed on the surface of the catalyst. The increased concentration of the reactants on the surface influences the rate of reaction. Adsorption being an exothermic process, the heat of adsorption is taken up by the surface of the catalyst, which is utilised in enhancing the chemical activity of the reacting molecules. The view does not explain the specificity of a catalyst.

The modern adsorption theory is the combination of intermediate compound formation theory and the old adsorption theory. The catalytic activity is localised on the surface of the catalyst. The mechanism involves five steps:

(i) Diffusion of reactants to the surface of the catalyst.

(ii) Some form of association between the catalyst surface and the reactants occurs. This is assumed to be adsorption.

(iii) Occurrence of chemical reaction on the catalyst surface.

(iv) De-sorption of reaction products away from the catalyst surface.

(v) Diffusion of reaction products away from the catalyst surface.