1. Kinetics What do you understand about rate of reaction?
L.O: To know how the rate of a reaction can be affected.
To be able to explain changes in rate of reaction in terms of particle collisions
What do you understand about rate of reaction?
What does rate mean?

2. What affects the rate of a reaction?
Concentration of reactants.
Pressure in a reaction (gaseous).
Temperature of reaction.
Particle size (surface area) of a solid.
Intensity of radiation (specific reactions).
Using a catalyst.

3. but, overall, is NOT consumed
RATE CAN BE INCREASED BY :
1. Increasing
reactant concentrations.
2. Increasing
reactant pressures.
3. Increasing
temperature.
4. Adding
a catalyst
– species which causes rate to increase
but, overall, is NOT consumed
5. Using powdered
solids
- increases surface area
Explained using : A.
the COLLISION THEORY
and B.
the BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGIES

4. Collision Theory
A chemical reaction will ONLY happen if the reacting particles:
1. COLLIDE
and 2. Collide with a kinetic energy EQUAL TO or GREATER THAN the reaction’s ACTIVATION ENERGY, (Ea).
Activation Energy (Ea) = the MINIMUM energy of collision needed for reaction to occur.
and 3. Collide with PROPER ORIENTATION (bonds are correctly aligned to react)

5. Energy barrier separates R from P
An Enthalpy Profile for a Chemical Reaction
Energy barrier separates R from P
Enthalpy
Time
Larger Ea
 SLOWER reaction Ea
Smaller Ea
 FASTER reaction
Reactants
Products
Remember: Ea = activation energy

6. An Enthalpy Profile for the reaction between chlorine atoms and ozone
Point X
Cl O O O
The arrangement at point X only exists for a very short time.
This is a simple one step reaction.
Others (combustion of methane) go through a series of steps and there will be a point X for each step.
Enthalpy
Time Ea
Cl + O3
ClO + O2

7.  can react when they collide
Boltzmann distribution of molecular energies
Explains why some molecules have energy  Ea
 can react when they collide
while others have energy < Ea
 cannot react even if they collide.
ENERGY, E 
NUMBER OF MOLECULES
WITH ENERGY E
TOTAL AREA under curve = TOTAL number of molecules in system A
Area under curve and to right of Ea value = number of molecules with sufficient energy to react on collision Ea
Most molecules (this area) have insufficient energy to react

8. How TEMPERATURE influences rateEa
ENERGY, E 
NUMBER OF MOLECULES
WITH ENERGY E A
N.B: Area under curve and to right of Ea value = number of molecules with sufficient energy to react on collision
How TEMPERATURE influences rate
Graph B represents the distribution of molecular energies at B
HIGHER TEMPERATURE but SAME CONCENTRATION than for graph A
 more mols with ≥ Ea
 more collisions per second with ≥ Ea
 faster reaction for B

9. Effect of temperature on distribution of energies
As the temperature increases, more molecules have a higher kinetic energy.
The increase in energy helps the collision to happen more effectively:
More molecules have more energy, so it is MORE LIKELY that a collision will have enough energy to overcome the Ea.
Higher speed of molecules means it takes less time for molecules to get close enough to react. There are MORE collisions.
Raising the temperature will always increase the rate of a reaction because it makes molecular collisions MORE EFFECTIVE

10. Variation of boltzmann distribution with temperature
ENERGY, E 
NUMBER OF MOLECULES
WITH ENERGY E
TOTAL AREA under curve =
TOTAL number of molecules in system T1
As temperature increases T1  T2  T3 T2
 total number molecules unchanged
but more molecules with higher energy T3
 distribution shrinks vertically and
moves to right

11. How CONCENTRATION influences rate
ENERGY, E 
NUMBER OF MOLECULES
WITH ENERGY E Ea A
N.B: Area under curve and to right of Ea value = number of molecules with sufficient energy to react on collision
How CONCENTRATION influences rate
Graph B represents the distribution of molecular energies at
LOWER CONCENTRATION but SAME TEMPERATURE than for graph A B
 fewer mols with ≥ Ea
 fewer collisions per second with ≥ Ea
 slower reaction for B

12. How PRESSURE influences rate
ENERGY, E 
NUMBER OF MOLECULES
WITH ENERGY E Ea A
N.B: Area under curve and to right of Ea value = number of molecules with sufficient energy to react on collision
How PRESSURE influences rate
Graph B represents the distribution of molecular energies at
LOWER PRESSURE but SAME TEMPERATURE than for graph A B
 fewer mols with ≥ Ea
 fewer collisions per second with ≥ Ea
 slower reaction for B

13. Variation of boltzmann distribution with concentration / pressure
TOTAL AREA under curve =
TOTAL number of molecules in system
ENERGY, E 
NUMBER OF MOLECULES
WITH ENERGY E
As conc / pressure decreases
 total number molecules decreases
 distribution shrinks vertically
C / P decreasing

14. Catalysis
Catalyst = a substance which INCREASES rate of a reaction without being consumed
Causes reaction to follow a different reaction route with lower activation energy, Ea
Energy
Time
Ea without catalyst
Ea with catalyst
Reactants
Products

15. Rate Examination Questions
1(a) The collision theory states that reactions can only occur when particles collide. Give a reason why collisions between particles do not always lead to a reaction.
Energy of collision may be LESS than activation energy of the reaction
1(b) Explain why, at a given temperature, hydrochloric acid reacts faster with powdered calcium carbonate than it does with lumps of calcium carbonate.
Carbonate has greater surface area
 more collisions per second with HCl
 more collisions per second having E ≥ Ea

16. 1(d). State and explain the effect on rate of
1(d) State and explain the effect on rate of increasing the temperature.
1(c) State and explain the effect on rate of increasing the concentration of a reactant at a fixed temperature.
Increases
because TOTAL number of collisions per second increases.
 greater number of collisions per second having E ≥ Ea
Increases
% of molecules with energy greater than activation energy increases
 greater number of collisions per second having E ≥ Ea

17. Q2. The curves labelled 1- 5 could be obtained by
Q2 The curves labelled 1- 5 could be obtained by reacting excess zinc with different solutions of HCl at different temperatures and measuring the volume of hydrogen gas produced at regular time intervals
100cm3 0.1M HCl at 25C + excess Zn  curve 3
(a) 50cm3 0.2M HCl at 25C + excess Zn  curve ____
(b) 100cm3 0.1M HCl at 15C + excess Zn  curve ____
(c) 100cm3 0.2M HCl at 25C + excess Zn  curve ____
100cm3 0.1M HCl at 25C + excess Zn
+ catalyst  curve ____ 2 4 1 2

18. Vol /cm3
t /s 1 2 3 4 5
Graphs for Q2

19. How A CATALYST influences rate
ENERGY, E 
NUMBER OF MOLECULES
WITH ENERGY E Ea A
N.B: Area under curve and to right of Ea value = number of molecules with sufficient energy to react on collision
How A CATALYST influences rate
At SAME TEMPERATURE and SAME CONCENTRATION but with a CATALYST
 lower Ea
Ea with catalyst
 more mols with ≥ Ea
 more collisions per second with ≥ Ea
 faster reaction with catalyst

20. Using powdered solids
(reactant or catalyst)
 total surface area of solid increased
 more collisions per second with surface
 more collisions per second with E  Ea
 faster reaction.

21. Boltzmann distribution of molecular energies
Ea with catalyst
ENERGY, E 
NUMBER OF MOLECULES
WITH ENERGY E Ea
Rate  number of collisions per second with E  Ea B A
without catalyst, rate  area A but with catalyst, rate  area (A+B)
Catalysts influence Ea and hence rate, not overall yield!!!