1. Higher Chemistry Controlling the Rate
NEW LEARNING
Collision Theory and Activation Energy
Kinetic energy and temperature
Reaction profiles
Catalysts
REVISION
Reactions monitored and graphs interpreted
Average Rate of Reaction Calculated from a Graph

2. Starter Questions
S3 Revision

3. Starter Questions
S3 Revision

4. Lesson 1: Collision Theory
Today we will learn to
Explain why reaction rate is increased by increasing temperature, concentration or surface area.
We will do this by
Using collision theory to explain some demonstrations and familiar reactions
We will have succeeded if
We can explain any rate change using collisions.

5. Activation Energy
Activity 1.1: Activation Energy
Your teacher may demonstrate some reactions that use heat or light to give the reactant particles the required activation energy.
Watch the demonstration of bromine reacting with a cycloalkane.
At National 5, we learned that alkenes rapidly decolourise bromine water. The reaction with saturated hydrocarbons is slower because light is needed to provide the activation energy.
Heat is also sometimes the source of the activation energy.

6. Factors Affecting Reaction Rate
Particle Size
The smaller the particle size, the greater the surface area, and the faster the reaction
Concentration
The higher the concentration of the reactants, the faster the reaction.
Temperature
The higher the temperature, the faster the reaction
These observations are explained by the collision theory.

7. Collision Theory In order to react particles must collide.
A chemical reaction will only occur if the reacting particles collide with enough kinetic energy or speed. The energy is required to overcome the repulsive forces between the atoms and molecules and to start the breaking of bonds.
The minimum kinetic energy required for a reaction to occur is called the activation energy (EA).
When the reactant particles collide with the required activation energy they form an activated complex. This unstable intermediate breaks down to form the products of the reaction.

8. Collision Geometry and the Activated Complex
E,g, The reaction of hydrogen and bromine H Br 
2HBr
Sometimes the collisions do not result in a reaction, despite having the minimum kinetic energy.
This is thought to be because the particles have not collided with the correct geometry (angle) to allow the activated complex to be formed.
In the above reaction of hydrogen and bromine the particles collided side on but if they collided end on…
H-H + Br-Br  H----H-----Br----Br
no reaction occurs as the activated complex cannot be formed if only 2 of the atoms come into contact with one another.

9. Surface Area/Particle Size
Increasing Surface Area
A reaction’s speed can be altered by increasing the surface area of one of the reactants. This can be achieved by breaking up lumps into chips or even better grinding it into a powder. When this occurs more of the reactant’s particles are exposed and therefore more particles can come in contact with the other reactant. This means that there will be a greater number of collisions, and therefore more successful collisions, resulting in a faster rate of reaction.
Activity 1.2: After looking at the available chemicals and equipment and discussion with others in your class and your teacher, devise an experiment that would show clearly and fairly that increasing the surface area of a solid reactant speeds up the reaction.

10. Collision Theory and Particle Size
The smaller the particle size, the faster the reaction as the total surface area is larger so more collisions will occur.

11. Starter Task
Complete questions 1 and 2 in the tutorial booklet and we will review as a class. You have 10 minutes!
S3 Revision

12. Starter Answers
S3 Revision

13. Starter Answers
S3 Revision

14. Lesson 2: Concentration
Today we will learn to
Investigate the effect of concentration on the rate of a reaction
We will do this by
Designing an experiment to monitor different concentrations but keep other variables the same.
We will have succeeded if
Alter concentration effectively to influence rate.

15. Concentration Increasing Concentration
A reaction’s speed can be altered by increasing the concentration of a solution or pressure in a gas. When this occurs there are more reactant particles in a given volume of space and therefore more particles can come in contact with the other reactant. This means that there will be a greater number of collisions, and therefore more successful collisions, resulting in a faster rate of reaction.

16. Concentration
Activity 1.3: Your teacher will show you what happens during the iodine clock reaction. After looking at the available chemicals and equipment and discussion with others in your class and your teacher, can you alter the concentration of the chemicals but not the total volume to get the clock to show exactly 1 minute.

17. Collision Theory and Concentration
The straight line graph means rate is directly proportional to the concentrations of the reactants, i.e. double the concentration and you double the rate. This is true of many reactions.
The faster rate is due to the increased number of collisions which must occur with higher concentrations of reactants.

18. Starter Task
Complete questions 1 and 2 in Tutorial 3 booklet and we will review as a class. You have 10 minutes!
S3 Revision

19. Starter Answers
S3 Revision

20. Starter Answers
S3 Revision

21. Lesson 3: Temperature Today we will learn to
Investigate the effect of temperature on the rate of a reaction.
We will do this by
Designing an experiment to monitor different temperature but keep other variables the same.
We will have succeeded if
We can plot a graph of rate vs temperature to show the effect.

22. Temperature Increasing Temperature
A reaction’s speed can be altered by increasing the temperature of the reaction mixture. When this occurs all the reactant particles have more energy and so a much greater number will have the required activation energy. This means that there will be a greater number of collisions with the activation energy, and therefore more successful collisions, resulting in a faster rate of reaction.
Shaded areas represents number of particles which can collide successfully at each temperature.
Activity 1.4: After looking at the available chemicals and equipment and discussion with others in your class and your teacher, devise an experiment that would show clearly and fairly that increasing the temperature speeds up the reaction.

23. Kinetic Energy and Temperature
Temperature is a measure of the average kinetic energy or speed of the particles of a substance.
At any given temperature, the particles of a substance will have a range of kinetic energies and this can be shown on an energy distribution graph.
NB The maximum height of T2 is always lower than T1
E.g. 20oC
E.g. 30oC

24. The graph above shows the kinetic energy distribution of the particles of a reactant at two different temperatures.
It shows that at the higher temperature (T2), many more molecules have energies greater than the activation energy and will be able to react when they collide.

25. Explain why this student statement is wrong…
Exit Task
Mr Wrong
Explain why this student statement is wrong…
‘Reaction rate is directly proportional to temperature, concentration and particle size’
Pupils could answer this, and if time, do some examples using the show me boards.

26. Starter Questions
From your S2-S4 Coursework, describe what a catalyst is.
A substance which speeds up a chemical reaction without being used up in the reaction.
Give an example of a catalyst.
Pt in a catalytic converter, Fe in the Haber process.
Can the same catalyst be used in different chemical reactions?
Sometimes.
S3 Revision

27. Lesson 4/5: Catalysts Today we will learn to
Explain how a catalyst speeds up a reaction.
We will do this by
Investigating different catalysts in the decomposition of H2O2
We will have succeeded if
We can explain the effect of a catalyst on the energy changes in a reaction.

28. Catalysts
A reaction’s speed can be altered by using a catalyst. A catalyst allows the reaction to take place via an alternative route that has a lower activation energy. This means that there will be a greater number of collisions with the activation energy at this temperature, and therefore more successful collisions, resulting in a faster rate of reaction. The enthalpy of the reaction stays the same.

29. Catalysts and Activation Energy
A catalyst provides an alternative pathway for the reaction with a lower activation energy.
(We will look at more of this type of energy diagram in the next week or so) ∆H

30. Catalysts and Reaction Rates
A catalyst is a substance which changes the speed of a chemical reaction without being permanently changed itself.
Catalysts speed up chemical reactions by providing an alternative reaction pathway which has a lower activation energy.
There are 2 main types of catalyst:
Homogeneous Catalysts
Heterogeneous Catalysts

31. Homogeneous Catalysts
Homogeneous catalysts are in the same state as the reactants.
e.g.
Heat
Very little reaction occurs
Fast reaction, solution turns green, gases evolve rapidly
Solutions of potassium sodium tartrate and hydrogen peroxide (colourless)
CoCl2(aq) +
Reaction complete, solution turns pink again

32. Heterogeneous Catalysts
Heterogeneous catalysts are in a different state to the reactants.
e.g. Decomposition of hydrogen peroxide (solution) using manganese (IV) oxide (solid) as a catalyst.
Manganese (IV) oxide (s)
2H2O2 (aq) H2O (l) + O2(g)

33. Common Catalysed Reactions
Zymase
Alcohol & Carbon dioxide
Maltose & glucose
Brewing
Platinum
Carbon Dioxide
Carbon Monoxide
Catalytic Converter
Vanadium (V) oxide
Sulphuric acid
Sulphur Dioxide & oxygen
Contact
Aluminium oxide or silicate
Short –chain hydrocarbons
Long-chain Hydrocarbons
Cracking
Nitric acid
Ammonia & Oxygen
Ostwald
Iron
Ammonia
Nitrogen & Hydrogen
Haber
Catalyst
Products
Reactants
Process
Homogeneous catalysis – hydroxide ions used in manufacture of soap from fats & oils.

34. Catalysts
Activity 1.5:
After looking at the available chemicals and equipment and discussion with others in your class and your teacher, devise an experiment that could identify the best catalyst for the decomposition of hydrogen peroxide.

35. How Heterogeneous Catalysts Work
Active sites
This type of catalyst is called a surface catalyst.
It works by adsorbing the reacting molecules on to active sites and holding them with weak bonds on its surface.
This not only causes the bonds within the molecule to weaken but also helps the collision geometry.
The reaction occurs on the surface with less energy needed to form the activated complex (lower activation energy).
The products are formed and leave the catalyst surface free for further reactions

36. Catalyst Poisoning
A surface catalyst can be poisoned when another substance attaches itself to the ‘active sites’. This is very often irreversible so prevents reactant molecules from being adsorbed onto the surface.
For this reason, catalysts have to be regenerated or renewed.
E.g.
Lead and its compounds are poisons of transition metal catalysts. This is why unleaded petrol must be used in cars with catalytic converters.
Arsenic and its compounds are also common poisons.
Catalysts can also be made ineffective by side-reactions.
E.g. the iron catalyst used in the Haber Process rusts due to the presence of air and water, so needs to be replaced every so often.

37. Enzymes
Enzymes catalyse the chemical reactions which take place in living cells.
Enzymes are complex protein molecules which are very specific- they usually only speed up one particular reaction and work best at specific temperatures and pH (optimum).

38. Examples in nature are:
Amylase – breaks down starch during digestion.
Catalase – breaks down hydrogen peroxide
Many enzymes are used in industry:
Invertase – used in chocolate industry for the hydrolysis of sucrose to form fructose and maltose.
Zymase - converts glucose into alcohol in the brewing industry.
Protease (and others) – used in biological washing powders to dissolve natural stains like protein .

39. Objective Traffic Lights
How do you feel about the lesson objectives?
Red = don’t think I have grasped this
Amber = feeling OK about this, have just about got there
Green = Confident I have achieved this
If there are no coloured cards available (e.g. in planners), Mark corners in the room for each colour and ask pupils to move to indicate confidence.