Power Point Kinetika

8/8/2019 Power Point Kinetika

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In a unimolecular reaction a single reactantmolecule isomerizes or decomposes to hive aproduct or products. In terms of reaction-rate

theory, the transition state or activated complexhas a configuration similar to the reactant so thatthe process can be represented by :

y A A ++ product



In 1922 Lindemann showed thatunimolecular reactions do acquire theirenergy of activated by bimolecularcollisions, but that this can give rise tofirst- order kinetics expent at lowpressures. His theory was in important

development and still forms the basisof all modern theories of unimolecularreactions.



His theory assumed that reactant molaculesare activated by collisions. He postulatedthere is a time lag between the activationand the reaction of these energisedmolecules to give product. As aconsequence, most of the energized

molecules collide with a normal reactantmolecule before they can be react, lose theirexcess energy and are deactivated



y At low pressures the rate of deactivation decreases at the molecularcollision rate decreases, and rate of conversion of energized moleculesto product becomes comparable to their rate of deactivation. Underthese conditions the rate of reaction depends on the rate of activationof the energised molecules (a bimolecular process) and the overall

kinetics becomes second order.The mechanism for the reaction can be represented by the following proceses :

Activation

A + A A* + A (1)

Deactivation A* + A A + A (-1)

k 1

k 1

k -1

k -1



y A* products (2)

W here A and A* represent a normal and energised molecule,

respectively.Since A* molecules are formed by reaction (1) and lost

by reactions (-1) and (2) , their rate of formation isgiven by the rate oh reaction (1) minus the sum of the

rates of reactions (-1) and (2) ; that isk1 [A]2 k-1[A*] [A] k2 [A] (6.1)

(6.2)

k 2

k 2



Combining equations 6.1 and 6.2 gives

[A*] = (6.3)

y The rate of reaction v (that is, the rate of formationof product) is given by the rate of reaction (2)

v = k 2 [A*]= (6.4)



y At high pressures where the rate of deactivation is muchgreater than the rate of conversion to products, that is k -

1[A] [A*] > k

2[A*], equation 6.4 becomes

v = = = k[A] (6.5)

Therefore the reaction is first order and the limiting or

high-pressure first order rate constant k is equal tok 1k 2/k -1



At low pressures, the rate of deactivation becomes less thanthe rate of conversion to product, that is k -1[A] [A*] < k 2[A*], so that equation 6.4 becomes.

v = k 1 [A]2 (6.6)

Therefore at low pressures the reaction is second order.

It has been shown that Lindemann theory predict at changein order when the pressure is increased or descreased.

Let the rate of reaction at any pressure be given by v = k [A] (6.7)



W here k is a rate coefficient that varies with pressure.From equation 6.4 it is seen that k is given by

k =

k = (6.8)

equation 6.8 predicts that a plot of k against [A] will be asshown in figurasi 6.1 and that k will have a limiting valueof k at high preesure, but falls of to zero at low pressure.



Alternatively Lindemann theory can be tested by inverting equation 6.8 giving

(6.9)

Further evidence for the basic soundness of Lindemann theory can beontained when the reaction is carried out with a constant reactantpressure and the total pressure varied by the addition of an inert gas Msuch as nitrogen, argon or xenon. The mechanism is now representedby

A + M A*+ M A* + M A + M

A* products

k 1k 1

-k 1-k 1

k 2k 2



Hinselwood theory

In 1927 Hinselwood postulated that the rate of energisation of a molecule depends on the number of vibrational degrees of freedom in the molecule.amolecule with a large number of vibrational degrees of

freedom has a much greater probability of acquiring theenergy needed for activation , since this energy can bedistributed among all these degrees of freedom.

For a molecule with one degree of freedom, therate constant for the energisation ion process (1) given by:

= exp ( / RT)



In the light of later theories the basic mechanismfor a unimolecular reaction is best presented by thefollowing modification of Lindemanns mechanism. A + A A* + A (1)

A + A* A + A (-1)

A* A + (2a)

A + products (2b)



RRK theory seeks to explain why plots of the typeshown in figure 6.2 are not linear. It proposes that amolecule is activated when the critical amount of energy is

concentrated in one particular bond. It is assumed that theenergy redestributes itself freely between the normal

vibrational modes during each vibration of the molecule.The rate constant K2b is therefore of the same magnitude asthe averagevibration frequencyof the molecule.

On the other hand slaters theory proposes thatenergy is not free to f low within the molecule. Slatersuggests that reaction occurs when a critical coordinate inthe molecule, usually a bond length, becomes extender by acritical amount.

RRK and slater theory

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