1. RATES OF REACTION

2. Students should be able to:
Describe the effect of concentration changes on the rate of a reaction.
Explain why an increase in the pressure of a gas increases the rate of the reaction.
State that a catalyst speeds up a reaction without being used up by the overall reaction.
Using enthalpy profile diagrams, explain how a catalyst increases the rate of a reaction.

3. Understand that catalysts affect conditions, often needing lower temperatures, reducing the energy costs and emissions resulting from the combustion of fuels.
Explain that catalysts allow different reactions to be used, with better atom economy and less waste.
State that catalysts have great economic importance.
Understand that catalysts are often enzymes operating close to room temperatures and pressures.

4. Explain the Boltzmann distribution and its relationship with activation energy.
Use the Boltzmann distribution to explain qualitatively the effect of temperature on rates of reactions.
Interpret catalytic behaviour in terms of the Boltzmann distribution.

5. RATES OF REACTION CONTENTS Prior knowledge Collision Theory
Methods for increasing rate
Surface area
Temperature
Catalysts
Light
Pressure
Concentration
Check list

6. Before you start it would be helpful to…
RATES OF REACTION
Before you start it would be helpful to…
know how the energy changes during a chemical reaction
know the basic ideas of Kinetic Theory
know the importance of catalysts in industrial chemistry

7. QUALITATIVE or QUANTITATIVE
CHEMICAL KINETICS
Introduction
Chemical kinetics is concerned with the dynamics of chemical reactions such as the way reactions take place and the rate (speed) of the process.
QUALITATIVE or QUANTITATIVE
Chemical kinetics plays an important part in industrial chemistry because the time taken for a reaction to take place and the energy required are of great economic importance. The kinetic aspect of chemistry is often at odds with the thermodynamic side when considering the best conditions for industrial production.

8. COLLISION THEORY also Collision theory states that...
particles must COLLIDE before a reaction can take place
not all collisions lead to a reaction
reactants must possess at least a minimum amount of energy –
ACTIVATION ENERGY
also
particles must approach each other in a certain relative way -
STERIC EFFECT

9. COLLISION THEORY
According to collision theory, to increase the rate of reaction you need...
more frequent collisions
increase particle speed have more particles present
more successful collisions
give particles more energy
or lower the activation energy

10. When magnesium is dropped into hydrochloric acid . .
successful collision occur when . .
unsuccessful collision is when . .

12. INCREASING THE RATE
The following methods may be used to increase the rate of a chemical reaction
INCREASE THE SURFACE AREA OF SOLIDS
INCREASE TEMPERATURE
ADD A CATALYST
INCREASE THE PRESSURE OF ANY GASES
INCREASE THE CONCENTRATION OF REACTANTS
SHINE LIGHT

13. MEASURING THE RATE
RATE How much concentration changes with time. It is the equivalent of velocity.
THE SLOPE OF THE GRADIENT OF THE
CURVE GETS LESS AS THE
REACTION SLOWS DOWN
WITH TIME
CONCENTRATION y x
gradient = y x
TIME
the rate of change of concentration is found from the slope (gradient) of the curve
the slope at the start of the reaction will give the INITIAL RATE
the slope gets less (showing the rate is slowing down) as the reaction proceeds

14. RATE CHANGE DURING A REACTION
Reactions are fastest at the start and get slower as the reactants concentration drops.
In a reaction such as A + 2B ——> C the concentrations might change as shown
Reactants (A and B)
Concentration decreases with time
Product (C)
Concentration increases with time
the steeper the curve the faster the
rate of the reaction
reactions start off quickly because
of the greater likelihood of collisions
reactions slow down with time as
there are fewer reactants to collide
TIME
CONCENTRATION B A C
units of rate are - mol dm-3 /s or mol dm-3 s-1

15. Experimental Investigation
MEASURING THE RATE
Experimental Investigation
the variation in concentration of a reactant or product is followed with time
the method depends on the reaction type and the properties of reactants/products
Extracting a sample from the reaction mixture and analysing it by titration.
- this is often used if an acid is one of the reactants or products
Using a colorimeter or UV / visible spectrophotometer.
Measuring the volume of gas evolved.
Measuring the change in conductivity.

16. INCREASING THE PRESSURE
increasing the pressure forces gas particles closer together
this increases the frequency of collisions so the reaction rate increases
many industrial processes occur at high pressure to increase the rate... but
it can adversely affect the position of equilibrium and yield
The more particles there are in a given volume, the greater the pressure
The greater the pressure, the more frequent the collisions
The more frequent the collisions, the greater the chance of a reaction

17. INCREASING CONCENTRATION
Increasing concentration = more frequent collisions = increased rate of reaction
Low concentration = fewer collisions
Higher concentration = more collisions
However, increasing the concentration of some reactants
can have a greater effect than increasing others

18. Low concentration . . .
High concentration . . .

20. INCREASING SURFACE AREA CUT THE SHAPE INTO SMALLER PIECES
Increases chances of a collision - more particles are exposed
Powdered solids react quicker than larger lumps
Catalysts (e.g. in catalytic converters) are finely divided for this reason +
In many organic reactions there are two liquid layers, one aqueous, the other non-aqueous. Shaking the mixture increases the reaction rate as an emulsion is often formed and the area of the boundary layers is increased giving more collisions.
CUT THE SHAPE INTO SMALLER PIECES 1 1 1 1 3 3
SURFACE AREA
= 30 sq units
SURFACE AREA
9 x ( ) = 54 sq units

22. INCREASING TEMPERATURE
Effect increasing the temperature increases the rate of a reaction
particles get more energy so can overcome the energy barrier
particle speeds also increase so collisions are more frequent
ENERGY CHANGES
DURING A REACTION
As a reaction takes place the enthalpy of the system rises to a maximum, then falls
A minimum amount of energy is required to overcome the ACTIVATION ENERGY (Ea).
Only those reactants with energy equal to, or greater than, this value will react.
If more energy is given to the reactants then they are more likely to react.
Typical energy profile diagram for an exothermic reaction

23. INCREASING TEMPERATURE ZARTMANN’S EXPERIMENT
According to KINETIC THEORY, all particles must have energy; the greater their temperature, the more energy they possess.
The greater their KINETIC ENERGY the faster they travel.
ZARTMANN heated tin in an oven and directed the gaseous atoms at a rotating disc with a slit in it. Any atoms which went through the slit hit the second disc and solidified on it. Zartmann found that the deposit was spread out and was not the same thickness throughout.
This proved that there was a spread of velocities and the distribution was uneven.
ZARTMANN’S EXPERIMENT

24. INCREASING TEMPERATURE
MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY
NUMBER OF MOLECUES WITH
A PARTICULAR ENERGY
MOLECULAR ENERGY
Experiments showed that, due to the many collisions taking place between molecules, there is a spread of molecular energies and velocities.
no particles have zero energy/velocity
some have very low and some have very high energies/velocities
most have intermediate velocities.

25. NUMBER OF MOLECUES WITH
MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY
MOLECULAR ENERGY
NUMBER OF MOLECUES WITH
A PARTICULAR ENERGY
no particles have
some have . . .
most have . . .

26. INCREASING TEMPERATURE
MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY T1
NUMBER OF MOLECUES WITH
A PARTICULAR ENERGY T2
TEMPERATURE
T2 > T1
MOLECULAR ENERGY
Increasing the temperature alters the distribution
get a shift to higher energies/velocities
curve gets broader and flatter due to the greater spread of values
area under curve stays constant - corresponds to the total number of particles

27. DECREASING TEMPERATURE
MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY T1
NUMBER OF MOLECUES WITH
A PARTICULAR ENERGY
TEMPERATURE
T1 > T3
MOLECULAR ENERGY
Decreasing the temperature alters the distribution
get a shift to lower energies/velocities
curve gets narrower and more pointed due to the smaller spread of values
area under curve stays constant

28. INCREASING / DECREASING TEMPERATURE
MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY T1
NUMBER OF MOLECUES WITH
A PARTICULAR ENERGY T2
TEMPERATURE
T2 > T1 > T3
MOLECULAR ENERGY
Lets REVIEW !
no particles have zero energy/velocity
some particles have very low and some have very high energies/velocities
most have intermediate velocities
as the temperature increases the curves flatten, broaden and shift to higher energies

29. NUMBER OF MOLECUES WITH
Increasing the temperature alters the distribution EXPLAIN !
Decreasing the temperature alters the distribution EXPLAIN ! T3 T1
NUMBER OF MOLECUES WITH
A PARTICULAR ENERGY T2
TEMPERATURE
T > T > T
MOLECULAR ENERGY

30. INCREASING TEMPERATURE
MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY
NUMBER OF MOLECUES WITH
A PARTICULAR ENERGY
NUMBER OF MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER Ea
MOLECULAR ENERGY
ACTIVATION ENERGY - Ea
The Activation Energy is the minimum energy required for a reaction to take place
The area under the curve beyond Ea corresponds to the number of molecules with sufficient energy to overcome the energy barrier and react.

31. INCREASING TEMPERATURE
T2 > T1
MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY T1
NUMBER OF MOLECUES WITH
A PARTICULAR ENERGY T2
EXTRA MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER Ea
MOLECULAR ENERGY
Explanation
increasing the temperature gives more particles an energy greater than Ea
more reactants are able to overcome the energy barrier and form products
a small rise in temperature can lead to a large increase in rate

33. ADDING A CATALYST
Catalysts provide an alternative reaction pathway with a lower Activation Energy (Ea)
Decreasing the Activation Energy means that more particles will have sufficient
energy to overcome the energy barrier and react
Catalysts remain chemically unchanged at the end of the reaction.
WITHOUT A CATALYST
WITH A CATALYST

34. ADDING A CATALYST
MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY
NUMBER OF MOLECUES WITH
A PARTICULAR ENERGY
NUMBER OF MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER
MOLECULAR ENERGY Ea
The area under the curve beyond Ea corresponds to the number of molecules with sufficient energy to overcome the energy barrier and react.
If a catalyst is added, the Activation Energy is lowered - Ea will move to the left.

35. ADDING A CATALYST
MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY
NUMBER OF MOLECUES WITH
A PARTICULAR ENERGY
EXTRA MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER
MOLECULAR ENERGY Ea
The area under the curve beyond Ea corresponds to the number of molecules with sufficient energy to overcome the energy barrier and react.
Lowering the Activation Energy, Ea, results in a greater area under the curve after Ea showing that more molecules have energies in excess of the Activation Energy

36. CATALYSTS - A REVIEW
work by providing an alternative reaction pathway
with a lower Activation Energy
using catalysts avoids the need to supply extra heat
- safer and cheaper
catalysts remain chemically unchanged at the end of the reaction.
Types Homogeneous Catalysts
same phase as reactants
e.g. CFC’s and ozone
Heterogeneous Catalysts
different phase to reactants
e.g. Fe in Haber process

37. ADDING A CATALYST
Catalysts provide an alternative reaction pathway with a lower Activation Energy (Ea)
Decreasing the Activation Energy means that more particles will have sufficient
energy to overcome the energy barrier and react
Catalysts remain chemically unchanged at the end of the reaction.
WITHOUT A CATALYST
WITH A CATALYST

38. CATALYSTS DO NOT AFFECT THE POSITION OF ANY EQUILIBRIUM
CATALYSTS - A REVIEW
lets go to – 3.12
CATALYSTS DO NOT AFFECT THE POSITION OF ANY EQUILIBRIUM
but they do affect the rate at which equilibrium is attained
a lot is spent on research into more effective catalysts - the savings can be dramatic
catalysts need to be changed regularly as they get ‘poisoned’ by other chemicals
catalysts are used in a finely divided state to increase the surface area

39. CATALYSTS - WHY USE THEM?
Catalysts are widely used in industry because they…

40. CATALYSTS - WHY USE THEM?
Catalysts are widely used in industry because they…
allow reactions to take place at lower temperatures SAVE ENERGY (lower Ea) REDUCE CO2 OUTPUT

41. CATALYSTS - WHY USE THEM?
Catalysts are widely used in industry because they…
allow reactions to take place at lower temperatures SAVE ENERGY (lower Ea) REDUCE CO2 OUTPUT
enable different reactions to be used BETTER ATOM ECONOMY
REDUCE WASTE

42. CATALYSTS - WHY USE THEM?
Catalysts are widely used in industry because they…
allow reactions to take place at lower temperatures SAVE ENERGY (lower Ea) REDUCE CO2 OUTPUT
enable different reactions to be used BETTER ATOM ECONOMY
REDUCE WASTE
are often enzymes GENERATE SPECIFIC PRODUCTS
OPERATE EFFECTIVELY AT ROOM TEMPERATURES

43. CATALYSTS - WHY USE THEM?
Catalysts are widely used in industry because they…
allow reactions to take place at lower temperatures SAVE ENERGY (lower Ea) REDUCE CO2 OUTPUT
enable different reactions to be used BETTER ATOM ECONOMY
REDUCE WASTE
are often enzymes GENERATE SPECIFIC PRODUCTS
OPERATE EFFECTIVELY AT ROOM TEMPERATURES
have great economic importance in the industrial production of POLY(ETHENE)
SULPHURIC ACID
AMMONIA
ETHANOL

44. CATALYSTS - WHY USE THEM?
Catalysts are widely used in industry because they…
allow reactions to take place at lower temperatures SAVE ENERGY (lower Ea) REDUCE CO2 OUTPUT
enable different reactions to be used BETTER ATOM ECONOMY
REDUCE WASTE
are often enzymes GENERATE SPECIFIC PRODUCTS
OPERATE EFFECTIVELY AT ROOM TEMPERATURES
have great economic importance in the industrial production of POLY(ETHENE)
SULPHURIC ACID
AMMONIA
ETHANOL
can reduce pollution CATALYTIC CONVERTERS

45. CATALYSTS - WHY USE THEM?
Catalysts are widely used in industry because they…
allow reactions to take place at lower temperatures SAVE ENERGY (lower Ea) REDUCE CO2 OUTPUT
enable different reactions to be used BETTER ATOM ECONOMY
REDUCE WASTE
are often enzymes GENERATE SPECIFIC PRODUCTS
OPERATE EFFECTIVELY AT ROOM TEMPERATURES
have great economic importance in the industrial production of POLY(ETHENE)
SULPHURIC ACID
AMMONIA
ETHANOL
can reduce pollution CATALYTIC CONVERTERS

49. certain reactions only
SHINING LIGHT
certain reactions only
shining a suitable light source onto some reactants increases the rate of reaction
the light - often U.V. - provides energy to break bonds and initiate a reaction
the greater the intensity of the light, the greater the effect
Examples a) the reaction between methane and chlorine - see alkanes
b) the darkening of silver salts - as used in photography
c) the reaction between hydrogen and chlorine
Equation H2(g) Cl2(g) ———> 2HCl(g)
Bond enthalpies H-H kJ mol-1 Cl-Cl kJ mol-1
Mechanism Cl2 ——> 2Cl• INITIATION
H Cl• ——> HCl H• PROPAGATION H• Cl2 ——> HCl Cl•
2Cl• ——> Cl TERMINATION
2H• ——> H2
H• + Cl• ——> HCl

50. units of rate are - mol dm-3 /s or mol dm-3 s-1
REVISION CHECK
What should you be able to do?
Recall and understand the statements in Collision Theory
Know six ways to increase the rate of reaction
Explain qualitatively how each way increases the rate of reaction
Understand how the Distribution of Molecular Energies is used to explain rate increase
Understand how the importance of Activation Energy
Recall and understand how a catalyst works by altering the Activation Energy
Explain how the rate changes during a chemical reaction
units of rate are - mol dm-3 /s or mol dm-3 s-1

51. When magnesium is dropped into hydrochloric acid . .
successful collision occur when . .
unsuccessful collision is when . .

52. Low concentration . . .
High concentration . . .

56. NUMBER OF MOLECUES WITH
MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY
MOLECULAR ENERGY
NUMBER OF MOLECUES WITH
A PARTICULAR ENERGY
no particles have
some have . . .
most have . . .

57. NUMBER OF MOLECUES WITH
Increasing the temperature alters the distribution
Decreasing the temperature alters the distribution T3 T1
NUMBER OF MOLECUES WITH
A PARTICULAR ENERGY T2
TEMPERATURE
T > T > T
MOLECULAR ENERGY

58. On collision frequency On collision energy On activation energy
Effect:
On collision frequency
On collision energy
On activation energy
On fraction of successful collisions
On rate
Decrease particle size (solids only)
Increases
No effect
Increase concentration
(liquids and gases)
Increase pressure (gases)
Increase temperature
Add a catalyst
Decreases

62. FINDING A GOOD CATALYST
IDEA
When iron (III) nitrate reacts with sodium thiosulphate the initial purple colour becomes colourless.
The speed of reaction is determined by timing how long it takes to go colourless.
Different metal ions can be tested to see if they catalyse the reaction.

63. Measuring the speed of the reaction without a catalyst
Put 10 cm3 of iron (III) nitrate in the conical flask.
Stand the flask on the white tile.
Add 10 cm3of sodium thiosulphate into the conical flask.
Start the stop clock.
stop the clock when the solution goes colourless.
Record, in the table below, the time taken for the solution to go colourless.

64. Measuring the speed of the reaction with a catalyst
Pour 10 mls of iron (III) nitrate into the conical flask.
Stand the flask on the white tile.
Add 5 drops of nickel (II) sulphate solution.
Add 10 mls of sodium thiosulphate, into the conical flask.
Start the stop clock.
Record, in the table below, the time taken for the solution to go colourless

65. Substance Time (in secs) for solution to go colourless
No catalyst .
Nickel (II) sulphate .
Copper (II) sulphate .
Iron (II) sulphate .
Cobalt (II) chloride .

66. FOR THE TECHNICIAN
Conical flasks
test tube
Stop clock
white tile
teat pipette
sodium thiosulphate soln (40 g/litre)
iron (III) nitrate soln [consisting of 20g of Fe (NO3) 3 and 20g of potassium thiocyanate in 1 litre]
iron (II) sulphate sol [0.01M]
copper (II) sulphate sol [0.01M]
cobalt (II) chloride. [0.01M]
nickel (II) sulphate [0.01M]
HAZARDS
Iron (II) salts,
Copper (II) salts,
Nickel (II) salts
Cobalt (II) salts
Harmful if swallowed
  Iron (III) salts
Act as an irritant to eyes and skin

67. Page 34 & 35
How fast?
Objectives
What do we mean by the rate of a chemical reaction ?
How can we find out the rate of a chemical reaction ?
page
C2 4.1 How fast?

68. Calcium carbonate and hydrochloric acid.
page
C2 4.1 How fast?

69. .
page
C2 4.1 How fast?

70. Calcium carbonate and hydrochloric acid
page
C2 4.1 How fast?

71. The rate is not a constant throughout the reaction - it changes!
2. The reaction is fastest at the start, gradually becoming slower as the reaction proceeds.
3. From the graph, the fastest part of the reaction is shown by the steepest curve.
4. The curve on the graph goes flat when the reaction is complete. This is because, as time goes on the volume of the gas evolved does not change.
page
C2 4.1 How fast?

72. page
C2 4.1 How fast?

73. The results are shown in the graph below.
Calcium carbonate powder was mixed with hydrochloric acid. A reaction occurs that produces carbon dioxide, Calcium chloride and water.
The mass of the dish, and its contents were recorded every half minute.
The results are shown in the graph below.
Write down the word equation for the reaction.
Explain why the mass of the dish and its contents lose mass.
What was the mass of the evaporating dish and its contents at the
a) start of the experiment?
b) end of the experiment?
What mass of carbon dioxide was produced during the experiment ?
When does the reaction seem to stop?
During which of these times in the experiment was the rate of reaction most rapid:
minutes?
minutes?
minutes?
On the graph sketch the curve you would expect if
the acid used was much hotter. (Label it X)
marble chips had been used instead of the powder.(Label it Y)
more concentrated hydrochloric acid was used.
If the amount of marble powder used was doubled how much carbon dioxide would be produced ?
page
C2 4.1 How fast?

74. Amount of reactant used or amount of product formed.
Summary
• The rate of a chemical reaction can be found by measuring the amount of a reactant used or the amount of product formed over time:
Rate of reaction =
Amount of reactant used or amount of product formed
Time
You should be able to
• interpret graphs showing the amount of product formed (or reactant used up) with time, in terms of the rate of the reaction.
page
C2 4.1 How fast?

75. page
C2 4.1 How fast?

76. Collision theory
Collision theory explains how the behaviour of particles effects the rate of a reaction
For a reaction to occur particles must collide
We need a successful collision so
the collision must have sufficient energy . . .
this is to break the existing bonds
The 4 main factors that effect the rate of a reaction are –
Temperature
Concentration
Surface area
Catalytic action
page
C2 4.2 Collision theory

77. page
C2 4.2 Collision theory

78. page
C2 4.2 Collision theory

79. Concentration after time
page
C2 4.2 Collision theory

80. page
C2 4.2 Collision theory