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

Published in: Science


  1. 1. CATALYSIS Joel Cornelio I M Sc Analytical Chem
  2. 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. 3. Diagram
  4. 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. 5. • Take the iron rod out and file it. • We get powdered iron.
  6. 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. 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. 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. 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. 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. 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. 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. 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. 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.
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  17. 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. 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.
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  20. 20. Transition metals used as catalysts, provide vacant d orbitals for bonding with the Π electrons of the substrate 2
  22. 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. 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. 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.
  26. 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. 27. 6. March, Jerry: Advanced Organic Chemistry, 4th ed, Wiley 2005. 7. Mukherji and Singh: Reaction Mechanism in Organic Chemistry, Macmillan 1984 8. INTERNET