1. Rates of Reaction

2. Objectives
To understand that a chemical reaction involves collisions between particles
To be able to describe the factors which will affect the rate of a chemical reaction.

3. Some Everyday Chemical Reactions
Burning wood
Fruit ripening
Getting a tan
Cooking food

4. How Reactions Happen
A chemical reaction involves a collision between particles.
The particles collide and make new substances
The particles which react are called the reactants
The substances which are made are called the products

5. Reactions Happen at Different Speeds
There are chemical reactions that occur very slowly and others that occur very quickly
RUST FORMATION
FIREWORKS

6. Rate = Change in concentration (molL-1)
Rates of Reaction
The rate of a reaction is how quickly a reaction proceeds
It may be defined as the change in concentration in unit time of any one reactant or product
Rate = Change in concentration (molL-1)
Time taken (s)

7. Calculate the rate of the following
H2O2 → H2O + ½ O2 The initial concentration of the reaction is 5 molL-1, ten seconds later this has decreased 3.5 molL-1 What is the rate?

8. Calculate the rate of the following
2SO2 + O2 → 2SO3 The initial concentration of the reaction is 15 molL-1, 20 seconds later this has decreased 12.5 molL-1 What is the rate?

9. Average and Instantaneous Rate
The average rate of reaction is average rate over the course of the reaction
The instantaneous rate of reaction is the rate at a particular point in time during the reaction

10. Controlling the Rates of Reaction
Being able to control the speed (rate) of chemical reactions is important both in everyday life (cooking) and when making new materials on an industrial scale.

11. Factors Affecting Rate of Reaction
Concentration of reactants
Temperature of reaction
Particle size of solid reaction
Nature of reactants
Presence of catalyst

12. 1. Concentration of Reactants
The higher the concentration of reactants the higher the probability of collisions between reactant molecules
There are less red particles in the same volume so there is less chance of a collision
There are more red particles in the same volume so there is more chance of a collision so the reaction goes faster

13. 2. Temperature of Reaction
An increase in temperature brings about an increase in reaction rate.
You give more energy to the system in the form of heat

14. 3. Particle Size Particle size (finely divided particles react faster)
Molecules can only collide at the surface.
Smaller particles bigger surface area

15. Dust Explosions
A dust explosion is the explosive combustion of a dust suspended in air in an enclosed location
Any solid material that can burn in air will do so at a rate that increases with increased surface area

16. Grain Dust Peril
In 1998, at a series of explosions occurred at a grain elevator facility in Haysville, Kansas
There were seven fatalities as a result of the explosions.
It is not the actual grain that is ignited, but the fine, thick dust which is released during the loading process when grain particles rub against each other

17. 4. Nature of Reactants Ionic compounds react faster than covalent
In reactions bonds are broken and form
When an ionic compound is placed in water it dissociates
It takes more energy to break covalent bonds

18. Activation Energy (Eact)
For a reaction to happen, energy is required.
Activation energy is the minimum energy with which particles need to collide to cause a reaction
This is different
according to the
type of bonds
of the reactants.

19. The activation energy may be shown on a reaction profile diagram (right)
These diagrams show energy as a barrier that needs to be overcome by the reactants before they become products

20. The difference between the energy of the reactants and the energy of the products is the heat of reaction (ΔH)

21. Endothermic reaction A reaction in which heat is taken in.
In an endothermic reaction heat is taken in from the surroundings and the products formed have more energy than the reactants.
It is written as + ΔH

22. Exothermic Reaction A reaction in which heat is liberated.
In an exothermic reaction heat is lost to the surroundings and the products formed have less energy than the reactants.
It is written as - ΔH

23. 5. Catalysts
A catalyst is a substance that alters the rate of a chemical reaction but is not consumed (used up) in the reaction.
In most cases it makes the reaction go faster. Some catalysts make a reaction go slower and are called negative catalysts or inhibitors. Eg Calcium propionate added to bread to make it stay fresh longer (ie. It slows down staleness)

24. The catalyst does not get used up in the reaction.
It gives the reaction the energy to get started

25. General Properties of catalysts
Catalysts are recovered chemically unchanged at the end of a reaction
(Eg. Manganese dioxide used to speed up decomposition of Hydrogen peroxide has exactly the same chemical properties before and after the reaction)

26. 2. Catalysts tend to be specific, even though a catalyst may catalyse one reaction it may not have any effect on a similar reaction
Enzymes in the body are examples of catalysts that are very specific
Know two examples
Protease breaks down proteins such as blood stains on clothes and are used in washing powders
Catalase breaks down hydrogen peroxide in the body

27. 3. Catalysts need only be present in very small amounts
Increasing the amount of catalyst does not greatly affect the rate of a reaction and in cases where it does it is usually something to do with the reaction itself

28. 4. Catalysts help reactions reach equilibrium quicker but do not change what the equilibrium of a reaction is

29. 5. Action of catalysts may be destroyed by catalytic poisons for example lead in petrol can destroy the catalytic converters in cars
Arsenic is a poison that inhibits the action of certain enzymes in the body

30. Types of Catalysis Chemists have discovered 3 types of catalysis
Homogenous Catalysis
Heterogeneous Catalysis
Autocatalysis

31. Homogenous Catalysis
This describes when reactants and catalysts are in the same phase and there is no boundary between them
Eg Iodine Snake where Potassium Iodide catalyses the decomposition of H2O2 to release oxygen both catalyst + Reactant are liquids

32. Heterogeneous Catalysis
Catalysis where the catalyst and reactants are in different phases Eg. a liquid and a solid
There is a boundary between the catalyst and the reactants
Eg. MnO2 catalyses the decomposition of H2O2 to release oxygen both catalyst + Reactant are liquids

33. Autocatalysis
When one of the products of the reaction catalyses the reaction
In this type of reaction it occurs slowly at first but as the reaction proceeds it gets quicker this is because the products drive the reaction forward.

34. Theories of Catalysts
Speed up a reaction by giving the reaction a new path.
The new path has a lower activation energy and more molecules have this energy.
The reaction goes faster.

35. Think of a catalyst as a tunnel through a mountain
By lowering the activation energy a catalyst makes it possible to carry out a reaction at lower temperatures (lower energy)

36. Mechanisms of Catalysis
The mechanism of catalysis tells us how the catalyst works
There are two main mechanism of catalysis for you to study
Intermediate Formation theory
Surface Adsorption theory

37. Intermediate Formation Theory of Catalysts
Homogeneous catalysts sometimes work by reacting with reactants to form unstable intermediate products
The intermediate exists for a very short time and reacts with the other reactant to give the final product and regenerate the catalyst
X+A + B → C AX
+ B

38. See it with you own eyes
Oxidation of Potassium Sodium Tartrate using hydrogen peroxide
Catalyst in this reaction is Cobalt ions which give a pink colour
The intermediate is a green colour which appears as carbon dioxide + steam are given off
The pink colour is restored at the end indicating the Cobalt ions have not been used up

39. Surface Absorption Theory of Catalysts
Heterogenous catalysis of gas reactions by metals
The reaction happens on the surface of the metal
The reaction occurs at the active site of the catalyst
A catalyst can have multiple active sites

40. Stages of reactions of ethene
CATALYST
Reactants get absorbed onto catalyst surface. Bonds are weakened
CATALYST

41. Bonds Break
CATALYST
New bonds formed
CATALYST

42. Second bond forms and product diffuses away from catalyst surface, leaving it absorb fresh reactants
Catalytic Poisons
Catalysts can be poisoned they can become less efficient and sometimes they no longer work at all
In heterogeneous catalysis particles that poison the catalyst (lead / arsenic) are absorbed more strongly onto the catalyst surface than the reactant particles
Catalytic poisons block the active sites of enzymes

43. How it works (Catalytic Converter)
The catalyst in the converter speeds up reactions that reduce atmospheric pollution
The catalyst remains unchanged at the end of the reaction
The catalyst is a mix of transition metals (platinum, rhodium, palladium)

44. Reactions Catalysed
Carbon monoxide is converted to carbon dioxide by reaction with oxygen
CO + ½ O2 → CO2
Carbon monoxide can react with nitrogen monoxide to give carbon dioxide
2CO + 2NO → 2CO2 + N2
Unburnt hydrocarbons are oxidised to carbon dioxide and water
C8H ½ O2 → 8CO2 + 9H2O

45. Environmental Benefits
Reduction in emissions of toxic gases including unburnt hydrocarbons
Reductions in photochemical smog

46. Smog in Beijing

47. Mandatory Experiment 14.1
Monitoring the rate of production of oxygen from hydrogen peroxide, using manganese dioxide as a catalyst

48. Hydrogen peroxide→Oxygen + Water

49. How to layout your results

50. T 15 30 45 60 90
120
150
180
210
Vol
110
134
170
184
190
197
O2 Cm3 60
Time secs 15

51. When the reaction is complete the amount of oxygen stops increasing
As the reaction proceeds more oxygen is produced

52. Average and Instantaneous Rates
The instantaneous rate of reaction is the rate at a particular point in time during the reaction
You use your graph to find the instantaneous rate

53. Draw a tangent to the curve

54. V2 V1
Connect the tangent to the X and Y axis using straight lines t1 t2

55. Calculations
Instantaneous rate= tan θ= Δv = v2 – v1 (Slope) Δt t2 – t1 = (cm3) (min) = 7 (cm3) 1 (min) instantaneous rate = 7cm3 min-1

56. Mandatory Experiment 14.2
Studying the effects on the reaction rate of (i) concentration and (ii) temperature, using sodium thiosulfate solution and hydrochloric acid

58. What’s happening?
2HCl(aq) + Na2S2O3(aq) → 2NaCl(aq) + SO2(aq) + S(s)↓ + H2O(l)
Place 100 cm3 of the sodium thiosulfate solution into a conical flask.
Add to the flask, while starting the stop clock at the same time.

59. Rate of Reaction (1/time)
Repeat the experiment using 80, 60, 40 and 20 cm3 of. sodium thiosulfate solution respectively. In each case, add water to make the volume up to 100 cm3 and mix before adding HCl.
Concentration of Thiosulphate
Reaction time
(s)
Rate of Reaction (1/time)
(s-1 )
0.1 M
0.08 M
0.06 M
0.04 M
0.02 M

60. Draw Graph 1
time
Concentration of Thiosulphate