Chemical Kinetics 2

The number of reacting species (atoms, ions or molecules) taking part in an elementary reaction, which must collide simultaneously in order to bring about a chemical reaction is called molecularity of a reaction

i) The reaction can be unimolecular when one reacting species is involved, for example, decomposition of ammonium nitrite. NH4NO2 → N2+ 2H2O ii)Bimolecular reactions involve simultaneous collision between two species, for example, dissociation of hydrogen iodide. 2HI → H2+ I2 iii)Trimolecular or termolecular reactions involve simultaneous collision between three reacting species, for example, 2NO + O2 —-> 2NO2 The probability that more than three molecules can collide and react simultaneously is very small. Hence, reactions with the molecularity three are very rare and slow to proceed.

Similarly, the overall rate of the reaction is controlled by the slowest step in a reaction called the rate determining step. Consider the decomposition of hydrogen peroxide: 2H2O2 ——I-/Alkaline medium——> 2H2O + O2’ This reaction is first order with respect to both H2O2 and I– . Evidences suggest that this reaction takes place in two steps (1) H2O2+ I– → H2O + IO– (2) H2O2 + IO– → H2O + I– + O2 Both the steps are bimolecular elementary reactions. Species IO- is called as an intermediate since it is formed during the course of the reaction but not in the overall balanced equation. The first step, being slow, is the rate determining step. Thus, the rate of formation of intermediate will determine the rate of this reaction. There for Rate = k [H2O2] [I-]

(i) Order of a reaction is an experimental quantity. It can be zero and even a fraction but molecularity cannot be zero or a non integer. (ii) Order is applicable to elementary as well as complex reactions whereas molecularity is applicable only for elementary reactions. For complex reaction molecularity has no meaning. (iii) For complex reaction, order is given by the slowest step and molecularity of the slowest step is same as the order of the overall reaction.

Zero order reactions are relatively uncommon but they occur under special conditions. Some enzyme catalysed reactions and reactions which occur on metal surfaces are a few examples of zero order reactions.

The decomposition of gaseous ammonia on a hot platinum surface is a zero order reaction at high pressure. 2NH3 (g) ——1130k/Platinum catslyst→N2(g) +3H2 (g) In this reaction, platinum metal acts as a catalyst. At high pressure, the metal surface gets saturated with gas molecules. So, a further change in reaction conditions is unable to alter the amount of ammonia on the surface of the catalyst making rate of the reaction independent of its concentration. The thermal decomposition of HI on gold surface is another example of zero order reaction.

Hydrogenation of ethene is an example of first order reaction. C2H4(g) + H2(g) → C2H6(g) Rate = k [C2H4] All natural and artificial radioactive decay of unstable nuclei take place by first order kinetics.

The half-life of a reaction is the time in which the concentration of a reactant is reduced to one half of its initial concentration. It is represented as t1/2. For zero order reaction t 1/2 prop. [R]0 . For first order reaction t1/2 is independent of [R]0 .

A chemical reaction in which value of order of reaction is one but molecularity is more than one are known as pseudo unimolecular/pseudo first order reaction.

consider the acidic hydrolysis of ethyl acetate In reality, it is a second order reaction and concentration of both ethyl acetate and water affect the rate of the reaction. But water is taken in large excess for hydrolysis, therefore, concentration of water is not altered much during the reaction. Thus, the rate of reaction is affected by concentration of ethyl acetate only. CH3COOC2H5+ H2O —-H+——>CH3COOH + C2H5OH Rate=k’[CH3COOC2H5][H2O] but [h2o] is a const so rate=k[CH3COOC2H5] where k=k’[H2O]

C12H22O11 + H2O + →H C6H12O6+ C6H12O6 cane sugar glucose fructose rate=k[C12H22O11]