For UG & PG students

1. 1
Dr. Y. S. THAKARE
M.Sc. (CHE) Ph D, NET, SET
Assistant Professor in Chemistry,
Shri Shivaji Science College, Amravati
Email: yogitathakare_2007@rediffmail.com
CHEMICAL DYNAMICS
COLLISION THEORY

3. Chemical Kinetics
Chapter 16

5. THEORIES OF REACTION RATES
1885-1889 1916-1920 1931-1937 1921-1952
ARRHENIUS
THEORY
COLLISION
THEORY
MODIFIED
COLLISION
THEORY
ABSOLUTE REACTION RATE
THEORY/ ACTIVATED
COMPLEX THEORY/
TRANSITION STATE THEORY
UNIMOLECULAR
THEORIES
Wynne-Jones
& Eyring
H. Eyring,
M. G. Evans &
M. Polanyi

6. Model for Kinetics
Arrhenius Theory
Effect of Temperature
Collision Theory
Rate determined by particle collisions;
Collision frequency and energy
Modified Collision Theory
Steric requirements
Transition State Theory
How reactants convert to products
Unimolecular Theory

7. Svante August Arrhenius
1859-1927

8. .
In 1884, based on this work, he submitted
a 150-page dissertation on electrolytic
conductivity to Uppsala for the
doctorate. It did not impress the
professors, among whom was Per
Teodor Cleve, and he received a fourth
class degree, but upon his defence it
was reclassified as third class. Later,
extensions of this very work would
earn him the Nobel Prize in Chemistry.

9. Many of these conceptual and experimental difficulties
would disappear with the brilliant work of van’t Hoff,
who introduced the concept of order of reaction and,
through it, the possibility of knowing the mechanism of a
chemical reaction just on the basis of chemical kinetics In
fact, van’t Hoff used the term molecularity for what we
would call today reaction order (the power to which a
concentration of a component enters into the rate
equation). When referring to the actual concept of
molecularity , this author used the explicit expression “the
number of molecules that participate in the reaction” The
Term order is due to Ostwald. Van’t Hoff received the
first Nobel Prize in 1901 for his discovery of the laws
of chemical dynamics

10. Chemical dynamic
Introduction: -
Collision theory is proposed by Arrhenius and Van't Hoff.
According to this theory the rate of reaction directly proportional to number
of collision per minute time.
1) Colliding molecule must possess sufficient K.E. Greater than or equal to
activation molecule energy to cause the reaction.
2) The reacting molecule must collide with proper orientation.
Explanation: -
• The molecule must collide with sufficient kinetic energy (K.E).

12. The reacting molecule must collide with proper orientation.

14. FACTORS AFFECTING RATE OF CHEMICAL REACTIONS
The rate of a chemical reaction is affected by several factors like:
1) Concentration of reactants
2) Pressure
3) Temperature
4) Catalyst
5) Nature of reactants
6) Orientation of reacting species
7) Surface area
8) Intensity of light
9) Nature of solvent

15. Arrhenius Equation
k: rate constant
Ea: activation energy (minimum required)
T: absolute temperature
R: universal gas constant=8.314JK-1mol-1
A: orientation factor or Arrhenius constant
Frequency factor or pre-exponential factor
Energy & orientation requirements for reaction
RT
Ea
Aek



17. Van-Hoff Rule:
10
1
2
1
2
t
t
t
t
t
k
k 
 


If t2t1
t
t
k
k 10

Temperature coefficient:
The ratio of rate constants of a reaction at two different temperatures
separated by 10°C'

18. CATALYST
Catalyst is a substance which alters the rate of a reaction
without being consumed or without undergoing any chemical
change during the reaction.
A catalyst increases the rate of reaction by providing a new
path with lower activation energy (Ea’) for the reaction.

19. Maxwell–Boltzmann Distributions
This fraction of molecules can be found through the expression
where R is the gas constant and T is the Kelvin temperature.
f = e
-Ea
RT

20. Ea
ACTIVATION ENERGY - Ea
The Activation Energy is the minimum energy required for a reaction to take place
The area under the curve beyond Ea corresponds to the number of molecules with
sufficient energy to overcome the energy barrier and react.
MAXWELL-BOLTZMANN
DISTRIBUTION OF
MOLECULAR ENERGY
NUMBER OF
MOLECULES WITH
SUFFICIENT
ENERGY TO
OVERCOME THE
ENERGY BARRIER
INCREASING TEMPERATURE
MOLECULAR ENERGY
NUMBEROFMOLECUESWITH
APARTICULARENERGY

21. Explanation
increasing the temperature gives more particles an energy greater than Ea
more reactants are able to overcome the energy barrier and form products
a small rise in temperature can lead to a large increase in rate
T1
T2
TEMPERATURE
T2 > T1
Ea
MAXWELL-BOLTZMANN
DISTRIBUTION OF
MOLECULAR ENERGY
INCREASING TEMPERATURE
MOLECULAR ENERGY
NUMBEROFMOLECUESWITH
APARTICULARENERGY
EXTRA
MOLECULES WITH
SUFFICIENT
ENERGY TO
OVERCOME THE
ENERGY BARRIER

23. ASSUMPTIONS : COLLISION THEORY
 The molecules are HARD SPHERE
 For a Reaction to occur between molecules, the two molecules must collide
 Not all collisions produce reaction; instead, reaction occur if and only if the RELATIVE
TRANSLATIONAL KINETIC ENERGY exceeds THRESHOLD ENERGY
 The MAXWELL-BOLTZMANN Distribution of Molecular Velocities is maintained
during the reaction
 Certain Steric requirements should be fulfilled
 The rate to be proportional to the rate of collisions, and therefore to the mean speed of
molecules
 The collision theory of gases gives the rate constant for bimolecular gas-phase reactions;
 A reaction occurs on the collision of two molecules only if they possess a certain minimum
amount of energy in excess to the normal energy of the molecules.
 The collision between the molecules other than activated molecules do not lead to chemical
reaction at all.
 The minimum energy in excess to their normal energy which the molecules must possess
before chemical reaction on collision is known as and equal to the activation energy.
 For the formation of product reactant molecule must collide with proper orientation.

29. Collision Theory (Bimolecular Collsions)
Z: no. of bimolecular collisions per second
fa: fraction with Ea
P: fraction with correct orientation
Ea: activation energy
pfZrate a 

31. Limitations of collision theory
1) It is applicable for the simple gases reaction only.
2) It is also applicable to the reaction in solution having reaction species exist in
simple molecule only.
3) Calculated values for the rate constant are usually too high compared with
measured values
4) The value of rate constant calculated from collision theory is equal to the observed
for the simple gaseous reaction only. For the reaction having complex molecule the
calculated rate is very large than that of observed rate.
5) There is no method of determine the (rho) (probability) for a reaction whose rate
constant has not been determine experimentally.
6) Collision theory is considering only the K.E. of colliding molecule there is no
reason why rotational and vibrational energy of molecule should be ignoring.
7) The collision theory is silent on clevage and formation of bond involved in the
reaction.
8) For the reaction involving complex molecule the experimental rate constant is quite
different from the calculate values.
9) Calculate value of (rho) from the structure and properties of the reacting molecule
is not in complete arrangement.
10) Measured activation energies are lower than the energies of the bonds that have to
be broken in reactions
11) Collision theory does not provide any prediction of p, the “steric factor”