This presentation is on Chemical Kinetics for the Chemistry Undergraduate students of Mumbai University.

1. Chemical Kinetics
Dr.Pravin U. Singare
Associate Professor
Department of Chemistry
N.M. Institute of Science,Bhavan’s College,
Andheri (west), Mumbai 400 058

2. Chemical Kinetics
• Study of chemical reactions.
• Study of reaction rate.
• Study of effect of various factors on reaction rate.
• Study of effect of presence of catalyst on reaction rate.

3. COLLISION THEORY
Postulates of collision theory
1. Reacting molecules are considered as a rigid sphere having only translational motion.
2. Rate of reaction is directly proportional to the number of collisions taking place between
the reacting molecules.
3. The collisions between the reactant molecules are considered to be elastic collisions in
which there is only exchange of energy between the reactant molecules, so that the total
energy remains same.
4. Not all colliding molecules react, but only those colliding molecules having certain
minimum energy will undergo effective collisions with each other and react to give
products.
5. In order to undergo effective collisions, the reacting molecules must be properly oriented
so that they will have head on collisions.
6. When the properly oriented molecules undergo effective collisions, there will be
rearrangement of bonds in the reacting molecules giving rise to the formation of
products.

4. The number of collisions (Z) taking place per cm3 per second between the reactant molecules
is given by the equation
Z = 2n2σ2 π𝑅𝑇/𝑀
here n is the number of molecules /cm3
σ is the collision diameter i.e. distance between the center of two colliding molecules
R is the gas constant
M is the molecular weight
T is the absolute temperature.
For bimolecular reactions in a mixture of gases 1 & 2, number of collisions (Z) taking place
per cm3 per second will be
2 Z(1,2) = n1.n2 σ2 8π𝑅𝑇/(𝑀1 +M2/M1.M2)
here M1 and M2 are the molecular weight of the reacting gas molecules 1 and 2
σ = σ1+ σ2/2 = average molecular diameter

5. Collision theory with respect to Bimolecular reaction
• According to the collision theory, out of Z number of molecules colliding with each other per cm3 per second,
only fraction of a molecules (q) having energy equal to the energy of activation will undergo effective
collisions resulting in the formation of product.
• Then the rate constant (k) of the bimolecular reaction in terms of colliding molecules will be given by the
equation
k = Z.q ------------ (1)
• According to Law of distribution of energy, if ‘n’ number of molecules are present per cm3 at temperature T,
then the number of molecules (n’) possessing energy of activation (Ea) is given by the relation
n’ = n. e-Ea/RT
𝑛′
𝑛
= e-Ea/RT ---------- (2)
But the fraction of molecules in activated state (q) is given by
q =
𝑛′
𝑛
---------- (3)
Substituting eq (3) in eq (2)
q = e-Ea/RT ---------- (4)
Substituting value of q from eq (4) in eq (1) we get
k = Z. e-Ea/RT ------------ (5)
• Comparing eq (5) with the Arrhenius equation k = A. e-Ea/RT it appears that ‘Z’, the number of molecules
colliding per cm3 per second which is twice in bimolecular reaction will be same as the frequency factor ‘A’,
i.e. A = 2Z

6. • The eq (5) can be applied to simple bimolecular reactions, but it is unable to explain the
complex molecular reactions in which the experimental determined k value is lower than
the theoretical k value calculated by eq (5).
• This is because the complex polyatomic molecules are considered to have vibrational,
rotational energies in addition to translational energy. Since this energies are not considered
while calculating the reaction rate k values, we get the deviation in experimental and
theoretical calculated k values.
• To account for this deviation, collision theory eq (5) is modified as
k = P.Z. e-Ea/RT ------------ (6)
Here P is the probability factor having value varying in the range of 1 to 10-9.
• The probability factor (P) take in to consideration the fact that not all activated molecules
will undergo reaction, but only the molecules which are properly oriented will react to give
products.
• The reactions in which the P value = 1 will obeys collision theory.

7. Lindemann Concept of Unimolecular reaction
• According to Lindemann concept
1. molecules get activated by collision with each other.
2. the activated molecules will decompose to give product only after some
time lag.
3. If the time lag is large, then there is a possibility that the activated molecules may get
deactivated by collision with low energy molecules.
• Because of the above reason, the rate of reaction will not depend on the number of activated molecules but will
depend only on those molecules which remain in the activated state and decompose finally to give product.
• Consider a Uni molecular reaction A  Product.
• According to Lindemann, the above reaction take place in following two steps:
1. In the 1st Step two molecules of A will collide with each other and there will be exchange of energy between the
reacting molecules. As result, one molecule get activated as A* and the rate of reaction is kf .
2. In 2nd step after a small time lag, the activated molecule will decompose to give product, and the rate of reaction is
k.
kf (activation) k (decomposition)
Step 1: A + A A* + A Step 2: A* Product
kr (deactivation)
• However, if in case of large time lag, there is a possibility that the activate A* will get deactivated by collision with
another molecule A of low energy and the reverse process will take place having rate constant kr.
• Therefore Rate of formation of A * = kf [A].[A] = kf [A]2 & Rate of removal of A* = kr[A*].[A] + k [A*]

8. • In the above reaction, A* is short lived and high energy intermediate which is formed during the
reaction.
• According to the steady state principle, for such short lived intermediate A*
rate of formation of A* = rate of removal of A *
kf [A]2 = kr[A*].[A] + k [A*]
kf [A]2 = [A*].{kr[A] + k}
[A*] = kf [A]2 / kr .[A] + k ------- (1)
• Since the rate of reaction [-dA/dt] is proportional to the concentration of intermediate activated
complex A*
Rate of reaction =
−𝑑𝐴
𝑑𝑡
α [A*]
Rate of reaction =
−𝑑𝐴
𝑑𝑡
= k1[A*] ---------- (2)
Substituting eq (1) in eq (2)
−𝑑𝐴
𝑑𝑡
= k1.kf [A]2 / kr . [A] + k ---------- (3)

9. −𝑑𝐴
𝑑𝑡
= k1.kf [A]2 / kr . [A] + k ---------- (3)
CASE I
• At high pressure & high concentration of A,
deactivation of A * take place to larger extent as
compared to decomposition of A* i.e. kr . [A]
>>> k and eq (3) will be
−𝑑𝐴
𝑑𝑡
= k1.kf [A]2 / kr . [A]
−𝑑𝐴
𝑑𝑡
= k1.kf [A] / kr -----(4)
• From eq (4) it is clear that at high pressure &
high concentration of A , the rate of reaction
−𝑑𝐴
𝑑𝑡
𝛼 [𝐴]
and the reaction will be unimolecular.
CASE II
• At low pressure & low concentration of A,
decomposition of A* take place to larger extent
as compared to deactivation i.e. k >>> kr . [A]
and eq (3) will be
•
−𝑑𝐴
𝑑𝑡
= k1.kf [A]2 / k ---------- (5)
• From eq (5) it is clear that at low pressure &
low concentration of A, the rate of reaction
−𝑑𝐴
𝑑𝑡
𝛼 [𝐴]2
and the reaction will be bimolecular.

12. Flash Photolysis
• It is a photochemical method which is used
to study the rate of chemical reactions which
are induced by light.
• The absorption of photon of light by the
reactant molecules will provides the
necessary activation energy for
photochemical reaction.
• In this method, the reaction mixture is
placed in a reaction vessel which is equipped
with suitable thermostatic arrangement for
maintaining the constant temperature during
the reaction.
• The reaction mixture is irradiated with
intense flash of UV or Visible light radiations
having energy in the order of 2000J/pulse for
short duration of 1 to 100 µ s.
• The rate of reaction occurring after flash
photolysis can be measured
spectrophotometrically by monitoring the
change in absorbance of reaction mixture as
a function of time.
Trigger
Photo flash tube
UV/Visible light radiations
Reaction vessel
Spectrophotometer Photomultiplier tube
Amplifier
Recorder