Theories of Catalysis

Theories of Catalysis

For a chemical reaction to occur, the reactants are to be activated to form the activated complex. The energy required for the reactants to reach the activated complex is called the activation energy. The activation energy can be decreased by increasing the reaction temperature. In the presence of a catalyst, the reactants are activated at reduced temperatures in otherwords, the activation energy is lowered. The catalyst adsorbs the reactants activates them by weakening the bonds and allows them to react to form the products.

As activation energy is lowered in presence of a catalyst, more molecules take part in the reaction and hence the rate of the reaction increases.

The action of catalysis in chemical reactions is explained mainly by two important theories. They are

(i) the intermediate compound formation theory

(ii) the adsorption theory.

1. The intermediate compound formation theory

A catalyst acts by providing a new path with low energy of activation. In homogeneous catalysed reactions a catalyst may combine with one or more reactant to form an intermediate which reacts with other reactant or decompose to give products and the catalyst is regenerated.

Consider the reactions:

A+B → AB            (1)

A+C → AC (intermediate)            (2)

C is the catalyst

AC+B → AB+C                (3)

Activation energies for the reactions (2) and (3) are lowered compared to that of (1). Hence the formation and decomposition of the intermediate accelerate the rate of the reaction.

Example 1

The mechanism of Fridel crafts reaction is given below

C₆ H₆ + CH₃Cl  ^{Anhydrous AlCl3}→ C₆ H₅ CH₃ +HCl

The action of catalyst is explained as follows

CH₃ Cl+AlCl₃ → [CH₃ ]⁺[AlCl₄ ]^-

It is an intermediate.

C₆ H₆ +[CH₃⁺][AlCl₄ ]^- → C₆ H₅CH₃ +AlCl₃ +HCl

Example 2

Thermal decomposition of KClO₃ in presence of MnO₂ proceeds as follows.

Steps in the reaction 2KClO ₃ → 2KCl+3O₂ can be given as

2KClO₃ +6MnO₂ → 6MnO₃ +2KCl

It is an intermediate.

6MnO₃ → 6MnO₂ +3O₂

Example 3:

Formation of water due to the reaction of H₂ and O₂ in the presence of Cu proceeds as follows. Steps in the reaction H₂+½O₂ → H₂O can be given as

2Cu+ 1/2 O ₂ → Cu ₂O

It is an intermediate.

Cu₂ O+H ₂ → H₂O + 2Cu

Example 4:

Oxidation of HCl by air in presence of CuCl₂ proceeds as follows. Steps in the reaction 4HCl + O₂ →2H₂O + 2Cl₂ can be given as

2CuCl ₂ → Cl₂ +Cu₂ Cl₂

2Cu₂ Cl₂ +O₂ → 2Cu₂OCl₂

It is an intermediate.

2Cu₂ OCl₂ +4HCl → 2H₂ O + 4CuCl₂

This theory describes

(i) the specificity of a catalyst and

(ii) the increase in the rate of the reaction with increase in the concentration of a catalyst.

Limitations

(i) The intermediate compound theory fails to explain the action of catalytic poison and activators (promoters).

(ii) This theory is unable to explain the mechanism of heterogeneous catalysed reactions.

2. Adsorption theory

Langmuir explained the action of catalyst in heterogeneous catalysed reactions based on adsorption. The reactant molecules are adsorbed on the catalyst surfaces, so this can also be called as contact catalysis.

According to this theory, the reactants are adsorbed on the catalyst surface to form an activated complex which subsequently decomposes and gives the product.

The various steps involved in a heterogeneous catalysed reaction are given as follows:

1. Reactant molecules diffuse from bulk to the catalyst surface.

2. The reactant molecules are adsorbed on the surface of the catalyst.

3. The adsorbed reactant molecules are activated and form activated complex which is decomposed to form the products.

4. The product molecules are desorbed.

5. The product diffuse away from the surface of the catalyst.

Active centres

The surface of a catalyst is not smooth. It bears steps, cracks and corners. Hence the atoms on such locations of the surface are co-ordinatively unsaturated. So, they have much residual force of attraction. Such sites are called active centres. So, the surface carries high surface free energy.

The presence of such active centres increases the rate of reaction by adsorbing and activating the reactants.

The adsorption theory explains the following

i. Increase in the surface area of metals and metal oxides by reducing the particle size increases acting of the catalyst and hence the rate of the reaction.

Figure 10.4 Finely divided catalyst is more effective due to increase in the number of active centres.

ii. The action of catalytic poison occurs when the poison blocks the active centres of the catalyst.

iii. A promoter or activator increases the number of active centres on the surfaces.