1. Chapter 6
Chemical kinetics

2. 6.1
Reaction Rates

3. Key Terms Chemical kinetics Reaction rate Average reaction rate
Instantaneous reaction rate

4. Chemical Kinetics
Thermodynamics (Chapter 5) – does a reaction take place?
Kinetics (Chapter 6) – how does a reaction proceed (reaction mechanism) and how fast (reaction rate).
Reaction rate is the change in the concentration of a reactant or a product with time (Units: mol/Ls).
R P
rate = -
D[R] Dt
D[R] = change in [reactants] over a time period Dt
rate =
D[P] Dt
D[P] = change in [products] over a time period Dt
Because [R] decreases with time, D[R] is negative.

5. Describing Reaction Rates
Reaction rate (r) is determined by measuring the rate at which a product is formed or the rate at which a reactant is consumed over a series of time intervals
properties like mass, colour, conductivity, volume, pressure or concentration may be measured to determine the reaction rate
reaction rate is expressed mathematically in terms of a change in property of reactant or product per unit time

6. A B
time
rate = -
D[A] Dt
rate =
D[B] Dt

7. TASK 1 Copy down the example from the board. (Sample Problem 1 and 2)
Do Practice Problem 1 on page 350

8. Reaction Rates Using a Graph
Reaction Rate for Reactant

9. Br2 (aq) + HCOOH (aq) 2Br- (aq) + 2H+ (aq) + CO2 (g)
slope of
tangent
slope of
tangent
slope of
tangent
average rate = -
D[Br2] Dt
= -
[Br2]final – [Br2]initial
tfinal - tinitial
instantaneous rate = rate for specific instance in time

10. Average vs Instantaneous Reaction Rate
AVERAGE REACTION RATE
The average rate of reaction between two time points is equal to the slope of the secant line drawn between those two points. (Figure 4 on page 351).
INSTANTANEOUS REACTION RATE
The instantaneous rate at any time is equal to the slope of the tangent to the curve at that particular instant in time. (Figure 7 on page 353)

11. TASK 2 Copy down the sample problems frompage 352
Do practice problem 1 on page 352
Do sample problem 1 and 2 on page 354
Do Practice Problem 1 and 2 on page 356

12. Factors Affecting Reaction Rates
6.2 and 6.3
Factors Affecting Reaction Rates

13. Collision Theory
Collision Theory tells us that in order to react, molecules and atoms must touch each other. They must hit each other hard enough to react.
Anything that increase these things will make the reaction faster.
Activation Energy is the minimum energy that reactant molecules must possess for a reaction to be successful.

14. Activation Energy 1 Defined as:
The minimum energy required to bring about a chemical reaction.
If there were no such thing as ‘activation energy’ life would be very difficult:
Gasoline for your car would ignite as soon as it came into contact with air.
You would burst into flames.
Trees would spontaneously combust.
Activation energy is why these things do not happen, there is an energy barrier so most reactions need to be ‘started off’ by putting in some energy.

15. Energy
Reactants
Products
Reaction coordinate

16. Activation Energy - Minimum energy to make the reaction happen
Reactants
Products
Reaction coordinate

17. Activated Complex or Transition State
Energy
Reactants
Products
Reaction coordinate

18. Energy
Reactants
Overall energy change
Products
Reaction coordinate

19. Things that Effect Rate
Chemical Nature of the Reactants
Concentration of Reactants
Surface Area
Temperature of the Reaction System
Presence of a Catalyst

20. 1. Chemical Nature of the Reactants

21. 2. Concentration of Reactants
More concentrated closer together the molecules.
Collide more often.
Faster reaction.

22. Effect of Concentration
Consider the reaction between zinc and hydrochloric acid:
Zn + 2HCl  H2 + ZnCl2
(How could you follow the progress of this reaction? Click to find out)
Acid Particles
Zinc
1M hydrochloric acid
2M hydrochloric acid
There are more particles of acid per unit volume in the 2M acid than there are in the 1M acid. So, there will more collisions between the acid and zinc particles in the stronger acid, giving a faster reaction.
Higher concentration = faster reaction

23. 3. Surface Area Particle size
Molecules can only collide at the surface.
Smaller particles bigger surface area.
Smaller particles faster reaction.
Smallest possible is molecules or ions.
Dissolving speeds up reactions.
Getting two solids to react with each other is slow.

24. Effect of Surface Area 1
When solids take part in chemical reactions only the surface particles are exposed so they are the only ones that can collide with particles of other reactants.
‘Inner’ particles are protected and cannot collide with other particles until they become ‘exposed’.
The surface particles are ‘exposed’ and can react.

25. Effect of Surface Area 2 Larger surface area = faster reaction.
If we break up this ‘lump’ into smaller pieces the number of particles has not changed but the there are now more ‘surface’ particles.
There is now a greater surface area with more exposed particles so more collisions can occur, hence faster reaction.
Larger surface area = faster reaction.

26. 4. Temperature of the Reaction System
Higher temperature faster particles.
More and harder collisions.
Faster Reactions.

27. Effect of Temperature Higher temperature = faster reaction
According to kinetic theory (do you remember this?) as the temperature increases the particles in a substance move about more quickly.
Reaction at 300C
Reaction at 500C
As the temperature increases the number of collisions increases as well as the energy of the collisions. So temperature has a big effect on the rate of reaction. For every 100C increase the rate approximately doubles.
Higher temperature = faster reaction

28. 5. Presence of a Catalyst
Catalysts- substances that speed up a reaction without being used up.
Speeds up reaction by giving the reaction a new path.
The new path has a lower activation energy.
More molecules have this energy.
The reaction goes faster.
Inhibitor- a substance that blocks a catalyst.

29. Effect of a Catalyst 2 Catalyst = faster reaction.
Activation energy without catalyst.
The lower activation energy in the presence of a catalyst means the reaction will be faster. More of the collisions have enough energy to react. There is a lower ‘energy barrier’.
energy
Activation energy with catalyst.
Catalyst = faster reaction.

30. Energy
Reactants
Products
Reaction coordinate

31. Catalysts Pt surface H H H H H H Hydrogen bonds to surface of metal.
Break H-H bonds H H H H H
Pt surface

32. Catalysts H H C C H H H H H H
Pt surface

33. Catalysts C C Pt surface H H H H H H H H
The double bond breaks and bonds to the catalyst. H H H C C H H H H H
Pt surface

34. Catalysts C C Pt surface H H H H H H H H
The hydrogen atoms bond with the carbon H H H C C H H H H H
Pt surface

35. Catalysts H H C C H H H H H H
Pt surface

36. Summary
Increasing the surface area gives a faster reaction because more particles are ‘exposed’ to the other reactant.
Increasing the concentration increases the rate of reaction because there are more collisions between the reactant particles.
Increasing the temperature increases the rate of reaction because the particles move move quickly and so collide more often and with greater energy.
A catalyst increases the rate of a reaction because it reduces the activation energy so more of the collisions have enough energy to react.