Catalysis is the process by which a catalyst increases the rate of a chemical reaction without undergoing any permanent chemical change. A catalyst works by providing an alternative reaction pathway or mechanism that has a lower activation energy. There are two main types of catalysis: homogeneous catalysis where the catalyst is in the same phase as the reactants, and heterogeneous catalysis where the catalyst is in a different phase. Catalysts can be classified as positive or negative depending on whether they increase or decrease the reaction rate.

2. • A catalyst is a substance which alters the speed
of a chemical reaction without itself
undergoing any chemical change and the
phenomenon is known as catalysis.
• Example,
2KlClO3  2KCl + 3O2
In the above reaction, MnO2 acts as a catalyst.

3. General characteristics of catalytic
reactions
• 1. The catalyst remains unchanged in mass and in
chemical composition at the end of the reaction.
• 2. Only a small quantity of catalyst is generally
needed.
• 3. A catalyst cannot initiate a reaction. The
function of a catalyst is only to alter the speed of
the reaction which is already occurring at a
particular rate.
• 4. A catalyst does not alter the position of
equilibrium in a reversible reaction.
• 5. The catalyst is generally specific in its action.

4. Types of catalytic reactions
• Catalytic reactions are classified into two
broad types;
1. Homogeneous catalysis
2. Heterogeneous catalysis

5. 1. Homogeneous Catalysis
• In these reactions, the reactants and catalyst
remain in the same phase. The following are
some of the examples of homogeneous
catalysis.
• Oxidation of SO2 to SO3 with oxygen in the
presence of nitric oxide as the catalyst in the
lead chamber process.

6. 1. Homogeneous Catalysis
Hydrolysis of methyl acetate is catalysed
by H+ ions furnished by hydrochloric
acid.

7. 2. Heterogeneous Catalysis
• The catalytic process in which the reactants and the
catalyst are in different phases is known as
heterogeneous catalysis. Some of the examples of
heterogeneous catalysis are given below.
• Oxidation of SO2 to SO3 in the presence of Pt metal
or V2O5 as catalyst in the contact process for the
manufacture of sulphuric acid.

8. 2. Heterogeneous Catalysis
• Combination between nitrogen and hydrogen
to form ammonia in the presence of finely
divided iron in Haber’s process.

9. Types of catalysts
• Positive catalyst
• Negative Catalyst
• Auto catalyst
• Induced Catalyst

10. Positive catalyst
• A catalyst which enhances the speed of the
reaction is called positive catalyst and the
phenomenon is known as positive catalysis.
Various examples are given below :
• Decomposition of H2O2 in presence of
colloidal platinum

11. Positive catalyst
• Decomposition of KClO3 in presence of
manganese dioxide.

12. Negative Catalyst
• There are certain substances which, when
added to the reaction mixture, retard the
reaction rate instead of increasing it. These are
called negative catalysts or inhibitors and the
phenomenon is known as negative catalysis.
The examples are given below.
• The oxidation of sodium sulphite by air is
retarded by alcohol.

13. Negative Catalyst

14. Auto catalyst
• In certain reactions, it is observed that one of the products
formed during the reaction acts as a catalyst for that reaction.
Such type of catalyst is called auto catalyst and the
phenomenon is known as auto catalysis.
• In the oxidation of oxalic acid by potassium permanganate,
one of the products MnSO4 acts as a auto-catalyst because it
increases the speed of the reaction.

15. Induced Catalyst
• When one reactant influences the rate of other
reaction, which does not occur under ordinary
conditions, the phenomenon is known as induced
catalysis.
• 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.
Thus sulphite has induced the arsenite and hence
is called induced catalyst.

16. Promoters
• The activity of a catalyst can be increased by addition
of a small quantity of a second material. A substance
which, though itself not a catalyst, promotes the
activity of a catalyst is called a promoter. Some
examples of the promoters are given below.
• In the Haber’s process for the synthesis of ammonia,
traces of molybdenum increase the activity of finely
divided iron which acts as a catalyst.

17. Catalytic Poisons
A substance which destroys the activity of the
catalyst is called a poison and the process is
called catalytic poisoning. Some of the
examples are
• The platinum catalyst used in the oxidation of
SO2 in contact process is poisoned by
arsenious oxide.

18. Catalytic Poisons
The iron catalyst used in the synthesis of
ammonia in Haber process is poisoned by H2S

19. Active Centres
• The catalytic surface has unbalanced chemical
bonds on it. The reactant gaseous molecules are
adsorbed on the surface by these free bonds. This
accelerates the rate of the reaction. The
distribution of free bonds on the catalytic surface
is not uniform. These are crowded at the peaks,
cracks and corners of the catalyst. The catalytic
activity due to adsorption of reacting molecules is
maximum at these spots. These are, therefore,
referred to as the active centres. If a catalyst has
more active centres, then its catalytic activity is
increased.

20. THEORIES OF CATALYSIS
• There are two main theories to explain
catalysis.
• 1. Intermediate compound formation theory
• 2. Adsorption theory
• In general, the intermediate compound
formation theory applies to homogeneous
catalytic reactions and the adsorption theory
applies to heterogeneous catalytic reactions

21. The Intermediate Compound
Formation Theory
• According to this theory, the catalyst first
forms an intermediate compound with one of
the reactants. The 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.

24. Adsorption Theory
• This theory explains the mechanism of
heterogeneous catalysis. Here, the catalyst
functions by adsorption of the reacting
molecules on its surface.

26. • Step - 1. Adsorption of reactant molecules
• The reactant molecules A and B strike the surface of the catalyst.
They are held up at the surface by weak vanderwaal’s forces or by
partial chemical bonds.
• Step - 2. Formation of Activated complex
• The particles of the reactants adjacent to one another join to form an
intermediate complex (A-B). The activated complex is unstable.
• Step - 3. Decomposition of Activated complex
• The activated complex breaks to form the products C and D. The
separated particles of the products hold to the catalyst surface by
partial chemical bonds.
• Step - 4. Desorption of Products
• The particles of the products are desorbed or released from the
surface.

27. Applications of catalysis