Only about 1 in 1013 molecular collisions results in a reaction. For a reaction to occur, colliding molecules must have sufficient kinetic energy to overcome the activation energy barrier. Insufficient energy results in the molecules bouncing off each other without reacting. The rate of reaction increases with higher temperature as more molecules then have enough energy to react. A catalyst also increases reaction rate by providing an alternative reaction pathway with lower activation energy. Increasing the concentration of reactants or the surface area of solids increases the number of collisions and reaction rate.

1. QUESTION?QUESTION?
ANSWER!ANSWER!
Only a small fraction of
collisions results in a
reaction. (~1 in every 1013
collisions!)
Does every collision result in a
reaction?
Collision Theory

2. REACTION RATE & TEMPERATURE
QUESTION?QUESTION?
Why doesn't every molecular
collision result in a reaction?
2AB A2 + B2
No Reaction
Reaction occursA - A B - B
Illustrati

3. Collision Theory
No Reaction
Reaction occursA - A
B - B
 If molecules are moving too slowly, they collide with insufficient
energy (unable to overcome the repulsion between electron
clouds), and just bounce off each other instead of reacting.
 If the reacting species are colliding with sufficient kinetic energy--
activation energy--(to stretch, bend and break the present bonds),
this leads to the formation of products.
Illustration2:

4. Conditions for a Reaction to Occur
Reacting species must
 Collide
 Collide in a favourable orientation (appropriate geometry
or juxtaposition) – steric factor
 Collide with a certain minimum kinetic energy – (more
than the activation energy barrier).
Collision Theory

5. Reaction Profile
Energy
Reaction Pathway
CH3• + HCl → [H3C---H---Cl]#
→ CH4 + Cl•
CH3• + HCl
[ H3C---H---Cl ]#
CH4 + Cl•
∆H
Ea1
Ea2
A) One Essential Step
or Energy Profile Diagram

6. Characteristics of Activated Complex
 Very unstable i.e. It has a pretty short half-life.
 Its potential energy is greater than reactants or
products.
 The activated complex and the reactants are in chemical
equilibrium.
 It decomposes to form products or reactants.

7. ENERGY PROFILE DIAGRAM
B) Two or more essential/elementary steps
Step 1: A → B
Step 2: B → C
Overall: A → C
Energy
Reaction Coordinate
∆H
Ea
A
B C
B: reactive
intermediate

8. Example: Reaction Mechanism
Step 1: CH3CHO + I2 → CH3I + HI + CO
Step 2: CH3I + HI → CH4 + I2
Draw an energy profile diagram for the above reaction.Draw an energy profile diagram for the above reaction.
Overall:Overall: CH3CHO → CH4 + CO ∆∆H = +veH = +ve
Reactive Intermediates: CH3I and HI
Catalyst: I2

9. ENERGY PROFILE DIAGRAMenergy
Reaction progress
∆H
Ea
CH3CHO
CH4 + CO
CH3I + HI

10. Objectives:
 Explain the effects of concentration
change/pressure, temperature change ,
surface area and the presence of catalyst on
the reaction rate
Factors that Affect the Reaction Rate

11. Factors that Affect the Reaction Rate
 Reactant concentration @ pressure
 System temperature
 Catalyst
 Reactants’ physical appearance (particle size,
medium)
 Other agents: light, pH, presence of impurities
 Chemical nature of reactants

12. a) Reactant concentration / Pressure of gases
PV = nRT
P = (n/V) RT ⇒ P = c.R.T
⇒ pressure, P ∝ concentration, c
P or c ↑,
Distance between particles is shorter,
Effective collisions increase
More molecules are able to overcome Ea
The no. of collisions per unit time will increase
Reaction rate increases
Explanation:

13. b) TemperatureMolefraction
Kinetic energyEa
T1
T2
T2 > T1
Maxwell-Boltzmann
distribution
At higher T
More molecules posses kinetic
energy equal to or higher than
Ea
i.e.: More effective collisions
⇒ reaction rate increases

14. c) Catalyst
 A catalyst speeds up a chemical reaction by
providing an alternative pathway
(mechanism) that has a lower activation
energy.
 Catalyst are not consumed in a chemical
reaction.
(A substance, which slows down the rate of a
reaction, is called an inhibitor )

15. Function of Catalyst
 Catalysts provide alternative reaction pathways of
lower activation energies than those of the
uncatalysed reactions
energy
Reaction progress
X
Y
catalysed
∆H
Ea
uncatalysed

16. d) Surface Area
Collisions can only occur on the surface of a
solid. If you cut a potato in half, then the boiling
water can reach the “inside” surface that it
couldn’t previously. So again the number of
collisions per unit time will increase