Presentation on basic ideas in catalysis. suitable for UG and +1 and +2 classes.

1. CATALYSIS
Joel Cornelio
I M Sc Analytical Chem

2. Definitions and Properties
• To learn both these ideas, let’s imagine an
experiment.
• The experiment is a laboratory demo of the
Haber’s process of NH3 production.
• Requirements include three cylinders (C1, C2 and
C3) and large jar (J) made of metal, a file
• Chemicals required are N2 and H2 and H2S gases
separately in the three cylinders, an Iron rod,
Molybdenum.
• Other things required are a burner and a lot of
your attention.

3. Diagram

4. Steps in the Experiment.
• Open the taps and let the N2 and H2 into the
jar J.
• No reaction occurs.
• Now we insert the iron rod into the jar.
• Reaction does not begin, no ammonia formed.
• Now we heat the jar to about 400 to 450oC.
• Reaction begins, rate increases.

5. • Take the iron rod out and file it.
• We get powdered iron.

6. • Weigh this iron powder, and put it inside the
jar, keep the temperature between 400-450oC
• There’s a tremendous increase in the reaction
rate.
• Now add a weighed amount of molydenum
into the jar.
• Reaction rate is highest.

7. Remove the iron powder out using a magnet and
weigh it.
There is no change in it’s weight.
Put the iron back and Open the third tube, and
let the H2S gas into the chamber. Rate of the
reaction rapidly decreases.
If the iron is weighed now, a small change in it’s
weight is seen.

8. Terms in Catalysis 1
• Catalyst: A substance that alters the rate of a
reaction without any undergoing any change
in itself.
• Substrate: The molecule subject to catalysis.
• Positive Catalyst: increases rate of reaction.
• Negative Catalyst: decreases rate of reaction.

9. • Promoter: Substances that increase the
catalytic activity, by increasing the number of
active sites in the catalyst.
(In other words, the catalyst of a catalyst.)
• Catalytic Poison: decreases or stops catalytic
activity by blocking the active sites of the
catalyst. 4

10. • Homogenous Catalysis: Reaction where the
catalysts are in the same phase as the
reactants.
e.g: Hydrolysis of ester into alcohol and acid
catalyzed by H+ ions
• Heterogeneous Catalysis: Reaction where the
catalyst and reactant are in different phases.
e.g: Contact process for H2SO4 production
catalyzed by Pt gauze or V2O5.

11. Properties
• A small quantity of the catalyst is enough to
catalyze a large quantity of reactants into
products.
• Most catalysts are highly selective.
• Though reactants are same, the type of catalyst
used may decide the products formed. 8
Eg.
Pt
C2H2 + 2H2 C2H6
Lindlar’s
Catalyst
C2H2 + H2 C2H4

12. • A catalyst cannot start the reaction by itself.
• Catalytic activity increases as surface area of
catalyst increases.
• Catalysts are thermolabile, this effect is very
well pronounced in enzymes.
• Catalytic activity is maximum at a catalyst’s
optimum temperature.
• A catalyst does not alter the position of the
equilibrium, instead it helps in achieving the
equilibrium faster.

13. • If two catalysts can catalyze the same reaction,
the better catalyst is the one which produces a
greater reduction in activation energy.
e.g. Activation energy for the decomposition of
H2O2 for different catalysts. 8
Catalyst EA, kJ/mol
None 75
MnO2 58
I- 56
Colloidal Pt 49

14. • The reaction occurs on the surface of the
catalyst, at special sites called active sites.
Increasing surface area, increases the number
of active sites available for catalysis.
• In homogenous catalysis, esp. in liquids, the
rate of the reaction is proportional to the
concentration of the catalyst.

15. 2

16. 2

17. So how do catalysts work..?
• All molecules posses energy, and during
reaction they undergo collision.
• Only those collisions result in chemical
reaction, in which the colliding molecules are
associated with a certain minimum amount of
energy, called the Threshold Energy.
• The excess of energy that molecules have to
gain to reach the threshold energy is called
Activation energy. 1
• Ea = Et - Ei

18. • Catalysts provide an alternate pathway for
reaction. 3
• This alternate pathway requires lower
activation energy and is easier to achieve, and
a larger number of molecules can react and
yield products.
• All these factors aid in increasing the rate of
the reaction, and that is what catalysis is all
about.

19. 2

20. Transition metals used as catalysts, provide vacant
d orbitals for bonding with the Π electrons of the
substrate
2

21. WILKINSON’S CATALYST

22. Speeding Up Organic Reactions 7
• Chloromethane reacts with hydroxyl to yield
methanol.
CH3Cl + OH- CH3OH + Cl-
• The reaction is much faster when iodide ions
are used, since they act as excellent
nucleophiles and leaving groups.
• CH3Cl + I- CH3I + Cl-
• CH3I + OH- CH3OH + I-

23. Applications.
• Every chemical industry uses catalysts to
produce chemicals.
• In biological systems, enzymes catalyze a
variety of biochemical reactions. 8
• e.g.
• Animals use enzymes for their defense, esp
snakes. (snake venom is also an enzyme.)

24. • Photosynthesis is the best and most
significant aspect of catalysis, by which we get
food.
• Pharmaceutical industries use catalyst for
production of various drugs and medicines.
• In catalytic converters of vehicles which aid in
reducing emissions of harmful exhaust gases.

25. CATALYTIC CONVERTER

26. References
1. Puri, Sharma and Pathania: Principles of
Physical Chemistry, VPC 2003
2. Matthews, Philip: Advanced Chemistry,
Cambridge University Press, 2006
3. Upadhyay, Santhosh: Chemical Kinetics and
Reaction Dynamics
4. Richard I. Masel: Principles of Adsorption and
Reaction on Solid Surfaces
5. D’Mello, Caroline: Organic reaction
mechanisms, Pearson, 2006

27. 6. March, Jerry: Advanced Organic Chemistry,
4th ed, Wiley 2005.
7. Mukherji and Singh: Reaction Mechanism in
Organic Chemistry, Macmillan 1984
8. INTERNET