CATALYSIS NOTES.pdf

UNIT IV: CATALYSIS 2014

CATALYSIS 

INTRODUCTION: The systematic study of the effect of various foreign substances on the rates of chemical

reactions was first made by Berzelius in 1835. He suggested the term catalyst for such substances. In Greek, kata = wholly, lein = to loosen. 

CATALYST: Substances, which accelerate the rate of a chemical reaction and themselves remain

chemically and quantitatively unchanged after the reaction, are known as catalysts. For example, MnO2 acts as a catalyst for the following reaction

The phenomenon of increase in the rate of a reaction that results from the addition of a catalyst is known as catalysis. The action of the catalyst can be explained on the basis of intermediate complex theory. According to this theory, a catalyst participates in a chemical reaction by forming temporary bonds with the reactant resulting in an intermediate complex which decomposes to yield product and the catalyst. It is believed that the catalyst provides an alternative pathway or reaction mechanism by reducing the activation energy between reactants and products and hence lowering the potential energy barrier, and the reaction rate is increased as shown below

∆H

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UNIT IV: CATALYSIS 2014 Although a catalyst lowers the activation energy E a for a reaction, it does not affect the energy difference ∆H between the products and reactants. It is clear from Arrhenius equation, lower the value of activation energy (E a) faster will be the rate of a reaction. (Note: Arrhenius equation is K

= A e-Ea/RT, where A is the Arrhenius factor or the

frequency factor, R is gas constant, Ea is activation energy.) 

CHARACTERISTICS OF CATALYSIS: 1. Catalyst remains chemically and quantitatively unchanged after the reaction. However, it may undergo a physical change. 2. A small amount of the catalyst can catalyse a large amount of reactants. For example, 1 g of metallic platinum is sufficient to decompose 108 liters of H2O2. while, some catalysts are required in relatively large amount to be effective. For example, In Friedel-Crafts reaction, anhydrous aluminium chloride catalyst is required to the extent of 30% of the mass of benzene. 3. The activity of a solid catalyst is enhanced with increase in its surface area.Thus, finely divided nickel is a better catalyst than lumps of nickel. 4. Catalyst is specific in its action.while a particular catalyst can be used for one reaction, it will not necessarily work for another reaction. For example, decomposition of KClO3 is catalyzed by MnO2 but not by platinum. Sometimes, for the same substrate different catalyst yield different products.

5. Catalyst cannot initiate a reaction. In most cases, it accelerates the reaction already in progress. 6. Catalyst does not change the equilibrium constant of a reaction. It helps in attaining the equilibrium faster, that is, it catalyses the forward as well the backward reactions to the same extent so that the equilibrium state remains same but is reached earlier.

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UNIT IV: CATALYSIS 2014 7. Maximum activity of a catalyst is obtained at a particular temperature called as optimum temperature. 8. Catalyst does not change the enthalpy of reaction (∆H), i.e. it does not affect the energy difference ∆H between the products and reactants. 9. A catalyst does not alter Gibbs energy, ∆G of a reaction.It catalyses the spontaneous reactions but does not catalyse non-spontaneous reactions. 

HOMOGENEOUS AND HETEROGENEOUS CATALYSIS: Catalysis can be broadly divided into groups a) Homogeneous catalysis: When the reactants and the catalyst are in the same phase (i. e. liquid or gas) and are evenly distributed throughout, the process is said to be homogeneous catalysis. Examples of homogeneous catalysis in gas phase: i) Oxidation of SO2(g) to SO3(g) in the presence of NO(g) as the catalyst.

ii) Oxidation of CO(g) to CO2(g) in the presence of NO(g) as the catalyst.

Examples of homogeneous catalysis in liquid phase: i) Hydrolysis of an ester is catalysed by H+ ions furnished by HCl.

ii) Hydrolysis of sugar is catalysed by H+ ions furnished by sulphuric acid.

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UNIT IV: CATALYSIS 2014 b) Heterogeneous catalysis: The catalytic process in which the reactants and the catalyst are in different phases is known as heterogeneous catalysis. The catalyst is usually a solid, and the reactants are either gases or liquids. Heterogeneous catalysis is the most important type of catalysis in industrial chemistry. It is also used in catalytic converters in automobiles. Examples of heterogeneous catalysis with gaseous reactants: i) Oxidation of SO2(g) to SO3(g) in the presence of Pt.

ii) Combination between nitrogen and hydrogen to form ammonia in the presence of finely divided iron in Haber’s process.

iii) Oxidation of ammonia into nitric oxide in the presence of platinum gauze in Ostwald’s process.

iv) Hydrogenation of ethylene catalysed by using finely divided Ni or Pt.

v) Hydrogenation of vegetable oils in the presence of finely divided nickel as a catalyst. vegetable oils (l)

+

H 2(g)

Ni(s)

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vegetable ghee(s)

UNIT IV: CATALYSIS 2014 Examples of heterogeneous catalysis with liquid reactants: i) The decomposition of aqueous solution of hydrogen peroxide is catalysed by platinum.

ii) Reaction of benzene and acetyl chloride is catalysed by anhydrous aluminium chloride.



ACTION OF CATALYTIC PROMOTERS AND INHIBITORS: a. Promoters or activators: Promoters are substances that enhance the activity of a catalyst. Such substances act as catalyst for the catalyst. For example, In Haber’s process for manufacture of ammonia, molybdenum acts as a promoter for iron which is used as a catalyst.

There is no clear explanation for the mechanism of promoter action. It is assumed that the following changes take place on addition of promoter. i) Change in catalyst spacing: The lattice spacing is increased thus, enhancing the space between the catalyst particles. The bonds between the molecules of the reactants (For e.g. H2) are weakened and cleaved. This makes the reaction go faster. Thus, the role of Mo as a promoter in the manufacture of NH3 by Haber’s process for which iron is used as a catalyst can be explained.

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UNIT IV: CATALYSIS 2014 H ---------- H H ----- H Fe

Fe

Fe

Fe

Fe

Fe

Mo Fe

Fe Mo

Fe

Fe

Distance between catalyst particles

ii) Increase of active sites: The number of peaks and cracks on the catalyst surface is increased on the addition of promoter to the catalyst. This increases the concentration of reactants and rate of the reaction. b. Inhibitors or catalytic poisons: Inhibitors are substances that decrease the activity of a catalyst or make the catalyst inactive. Arsenius oxide and HCN are two of the most powerful catalytic poisons. Some examples of catalytic poisons 1. Arsenius oxide poisons the catalytic activity of platinized asbestos in the manufacture of H2SO4 by contact process. 2. Platinum catalyst used in the oxidation of H2 is poisoned by CO.

O C

Pt

O C

Pt

Pt

O C

Pt

Pt

Pt

Fig. Poisoning of platinum catalyst by carbon monoxide

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UNIT IV: CATALYSIS 2014 

AUTO CATALYSIS: When one of the products formed during the course of reaction itself act as a catalyst for that reaction the phenomenon is called as autocatalysis. In normal reaction, the rate of reaction decreases with the passage of time. However, in autocatalysis, as the reaction proceeds, concentration of catalytic product increases and so the rate of reaction increases. Examples: 1) Hydrolysis of methyl acetate is catalysed by H+ ions furnished by acid. As the reaction proceeds, concentration of catalyst (H+) increases and hence, the rate of reaction increases.

2) In the titration of oxalic acid (H2C2O4) with acidified potassium permanganate (KMnO4), the reaction is slow in the beginning but becomes fast as the reaction progresses. Manganese sulphate or Mn2+ ions produced during the reaction acts as autocatalyst for the reaction. As the concentration of Mn2+ ions increases with time, the rate of reaction increases with time. Hence, the time required for decolourisation of first drop of KMnO4 is much higher and it goes on decreasing with time.



ENZYME CATALYSIS: a. Enzymes are complex nitrogenous organic compounds produced by living plants and animals which facilitate the reactions, but unlike other catalysts are not consumed in reactions. b. They are typically proteins or RNA.

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UNIT IV: CATALYSIS 2014 c. Numerous reactions that occur in the bodies of animals and plants to maintain the life process are catalysed by enzymes. The enzymes are, thus, termed as biochemical catalysts and the phenomenon is known as biochemical catalysis. d. The first enzyme was synthesised in the laboratory in 1969. The following are some of the examples of enzyme-catalysed reactions: 1. Inversion of cane sugar into glucose and fructose by invertase enzyme.

2. Conversion of glucose into ethyl alcohol and carbon dioxide by zymase enzyme.

3. Conversion of starch into maltose by diastase enzyme. diastase 2(C6H10O5)n(aq) + nH2O(aq) starch

nC12H22O11(aq) maltose

4. Conversion of maltose into glucose by maltase enzyme.

5. Decomposition of urea into ammonia and carbon dioxide by urease enzyme.

6. In stomach, the pepsin enzyme converts proteins into peptides while in intestine, the pancreatic trypsin converts proteins into amino acids by hydrolysis. 7. Conversion of milk into curd: It is an enzymatic reaction brought about by lacto bacilli enzyme present in curd.

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UNIT IV: CATALYSIS 2014 

CHARACTERISTICS OF ENZYME CATALYSIS: 1. One molecule of an enzyme may transform one million molecules of the reactant per minute. 2. Each enzyme is specific for a given reaction, i.e., one catalyst cannot catalyse more than one reaction. For example, the enzyme urease catalyses the hydrolysis of urea only. 3. The rate of an enzyme reaction become maximum at a definite temperature, called the optimum temperature. On either side of the optimum temperature, the enzyme activity decreases. The optimum temperature range for enzymatic activity is 298-310 K. Human body temperature being 310 K is suited for enzyme-catalysed reactions. 4. The rate of an enzyme-catalysed reaction is maximum at a particular pH called optimum pH, which is between pH values 5-7. 5. The enzymatic activity is increased in the presence of certain substances, known as co-enzymes. 6. The inhibitors interact with the active functional groups on the enzyme surface and often reduce or completely destroy the catalytic activity of the enzymes.

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MECHANISM OF ENZYME CATALYSIS: There are a number of cavities present on the surface of enzymes. These cavities are of characteristic shape and possess active groups like -NH2, -COOH, -SH, -OH, etc. These are actually the active centres on the surface of enzyme particles. Molecules of the substrate, which have complementary shape, fit into these cavities just like a key fits into a lock. Thus, the enzyme-catalysed reaction proceeds in two steps. Step 1: Binding of enzyme (E) to substrate (S) forms enzyme substrate complex (ES)

Step 2: Decomposition of the enzyme substrate (ES) complex to form product (P)

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UNIT IV: CATALYSIS 2014

Fig. Mechanism of enzyme catalysed reaction 

KINETICS OF ENZYME CATALYSED REACTIONS: MICHAELIS-MENTEN EQUATION In 1913, Michaelis and Menten proposed a mechanism for the kinetics of enzyme catalyzed reaction. It involves two steps Step I: The enzyme (E) reacts with the substrate (S) to form enzyme-substrate complex (ES). This is a reversible process and equilibrium is established rapidly.

Step II: The ES complex undergoes dissociates to give product (P) and enzyme (E). K2 ES

P

enzyme-substrate complex

product

+

Second step is the slow and rate determining step. Thus, Rate of reaction = - d[S] = d[P] = K2[ES] .....(3) dt dt But, [ES] is not an experimentally measurable quantity. Solve for intermediate [ES], d[ES] = K1[E][S] – K-1[ES] – K2[ES] dt Use of steady state approximation for [ES] We get, d[ES] = 0 dt d[ES] = K1[E][S] – K-1[ES] – K2[ES] = 0 dt

…..(4)

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E enzyme

.....(2)

UNIT IV: CATALYSIS 2014 Instead of solving for [ES] in terms of Free enzyme [E], solve for [ES] in terms of Total enzyme [E]0. Total enzyme [E]0 =

Free enzyme [E] + Bound enzyme

[ES]

[E] = [E]0 – [ES] …..(5) Replace [E] = [E]0 – [ES] from equation (5) in equation (4) d[ES] = K1{[E]0 –[ES]} [S] – K-1[ES] – K2[ES] = 0 dt d[ES] = K1[E]0 [S] – K1[ES] [S] – K-1[ES] – K2[ES] = 0 dt Rearrange [ES] terms to one side of the equation K1 [ES] [S] + K-1 [ES] + K2 [ES] = K1 [E]0 [S] [ES] {K1 [S] + K-1 + K2} = K1 [E]0 [S] [ES] =

K1 [E]0 [S]

…..(6)

K-1 + K2 + K1[S] Dividing the numerator and denominator by K1 [ES] =

K1 [E]0 [S] ÷ K1

=

{ K -1 + K2 + K1 [S]} ÷ K1 [ES] =

K1 [E]0 [S]

K1 [E]0 [S] K1 { K-1 + K 2 + K1 [S]} ÷ K1 =

K 1 { K-1 + K2 + K1 [S]}

[E]0 [S]

…..(7)

{ K-1 + K2} + {K 1 [S]}

K1

K1

K1

Introduce new term Km (Michaelis-Menten constant) in equation (7) Where, Km = K -1 + K2 K1 [ES] =

[E]0 [S]

…..(8)

(Since, K-1 + K 2 = Km )

[S] + Km K1 Substitute [ES] into equation (3) Rate of product formation = d[P] = K2[ES] = K2 [E]0 [S] …..(8) dt

[S] + Km

Equation (8) is known as Michaelis-Menten equation. Page 11

UNIT IV: CATALYSIS 2014 (a) When [S] >> Km, then neglecting Km as compared to [S] Rate of product formation = K2 [E]0 [S] = K2 [E]0 This is called Vmax [S] + Km (small and is neglected) Maximal rate = Vmax = K2 [E]0 (b) When [S] = Km, then Rate of product formation = K2 [E]0 [S] = K2 [E]0 [S] = 1 K2 [E]0 half maximal rate [S] + [S]

2[S]

2

Definition of Michaelis constant (Km) = concentration of [S] for which the rate is half-maximal. ..………………………………..Chapter completed……………………………………….

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