This document discusses reaction rates and factors that affect reaction rates. It begins by explaining what reaction rate means and how it can be measured by determining the amount of product formed, reactant used, or time taken for a reaction to complete. It then discusses several factors that affect reaction rates, including temperature, concentration, particle size, and catalysts. Higher temperatures, concentrations, smaller particle sizes, and the presence of catalysts generally increase reaction rates by increasing the frequency and success of particle collisions. The document provides examples and experiments to illustrate these concepts.

1. LEARNING OUTCOMES
 Explain what is meant by “rate of reaction”;
 Interpret graphical diagrammatic presentation of
data obtained in studying rates of reactions
 Identify the factors which affect the rate of
reaction
 Predict the effect of factors on rates of reaction
from given data
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Chapter 14

2. Measuring the Speed of Reaction
 Different kinds of reactions take place at different speeds.
 Some reactions are very fast e.g. explosion of gases and
chemicals.
 On the other hand, some reactions are slow e.g. rusting of
iron and fermentation of sugar to form ethanol.
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3. Rate of reaction
 The rate of a reaction tells us how fast or slow a reaction
is taking place.
 We can measure the rate of a reaction in 3 ways:
2. Measuring the amount of product formed per unit time
3. Measuring the amount of reactant used up or
remaining per unit time
1. Measuring the time taken for a reaction to complete
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4. Measuring the Rate of Reaction
An experiment was set up to measure the rate of reaction between
magnesium and two different solutions of dilute hydrochloric acid.
hydrochloric acid
(2 mol/dm3
)
hydrochloric acid
(1 mol/dm3
)
magnesium ribbon magnesium ribbon
Experiment I Experiment II
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5. Measuring the time taken for a reaction to complete
If the time taken in Experiment 1 for the magnesium to completely dissolve
in the acid was 60 s, and the time taken in Experiment II was 30 s, then
the speed of the reaction in Experiment II was two times as fast as in
Experiment I.
hydrochloric acid
(2 mol/dm3
)
hydrochloric acid
(1 mol/dm3
)
magnesium ribbon magnesium ribbon
Experiment I Experiment II
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6.  It can be seen that the shorter the time
taken for a reaction to complete, the faster
the speed of the reaction.
 Thus the speed of a reaction is inversely
proportional to the time taken:
Reaction Rate (speed ) = ___1____
Time taken
Measuring the time taken for a reaction to be
complete
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7. Measuring the amount of product formed in a reaction
An experiment was set up to measure
the rate of reaction between calcium
carbonate and dilute hydrochloric acid.
The CO2 produced is collected in a
gas syringe.
 The speed of the reaction can be determined by measuring the
volume of carbon dioxide produced at regular time intervals
during the reaction.
 A graph of the volume of gas formed is plotted against time taken.
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8.  The gradient becomes zero at 2.5 minutes, showing that no more gas is
produced and the reaction has stopped.
 The gradient of the graph is greatest
at the start of the experiment, showing
that the rate of the reaction is fastest at
the start of the experiment.
Measuring the amount of product formed in a reaction
 The gradient decreases with time, showing that the rate of the
reaction is decreasing over time.
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9. Rate (Speed) of reaction = Quantity of product formed
Time taken
 The rate of reaction at a particular
point P on the graph is given by the
gradient of the graph at P.
 Rate of reaction at P = Gradient
= y
x
= 26 cm3
/min
The average rate of reaction over a time interval is given by the formula:
For e.g. average rate for the first 2.5 minutes of the reaction = (70 – 0) cm3
2.5 min
= 28 cm3
/min
Measuring the amount of product formed in a reaction
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10. Measuring the amount of reactant left
 The change in mass of the reaction mixture can be read off
from the electronic top pan balance and a graph of mass of
the flask with its contents is plotted against time.
The speed of reaction
between calcium carbonate
and hydrochloric acid can also
be determined by measuring
the loss of mass of the flask as
carbon dioxide escapes from
the reaction mixture.

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11.  The gradient of the graph is greatest at the start
of the experiment, hence the speed of the
reaction is greatest at the start of the experiment.
 The gradient decreases with time, showing that
the speed of the reaction decreases as time
proceeds.
 The gradient is zero after about 4.2 min, showing
that the reaction has stopped.
 The reaction has stopped because one of the
reactants (either HCl or CaCO3 ) has been used
up in the reaction.
Measuring the amount of reactant left
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12. Quick check 1
1. Explain what is meant by the “rate of reaction”. How is the reaction
rate related to the time taken for a reaction to complete?
2. How may the speed of chemical reactions be measured
experimentally? Give two examples to illustrate your answer.
3. The graph shows the total volume of
hydrogen produced plotted against
time in a reaction.
Calculate the average rate of the
production of hydrogen.
Solution
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Chapter 14

13. 1. The “rate of reaction” tells us how fast or slow a reaction is taking place.
The reaction rate is inversely proportional to the time taken.
2. The speed of a chemical reaction can be measured by:
(i) determining the quantity of product formed per unit time;
E.g. to find the speed of reaction between magnesium and hydrochloric acid, we
can measure the volume of hydrogen produced over a period of time and
determine from the gradient of the volume-time graph, the speed of the reaction
at any particular time interval.
(ii) determining the quantity of reactant used up per unit time.
E.g. to find the speed of reaction between dilute hydrochloric acid and calcium
carbonate, we can measure the loss of mass form the reacting mixture over a
period of time. From the gradient of the mass-time graph, the speed of reaction
can be obtained at any particular time interval.
3. Average rate of the production of hydrogen = 32 cm3
80 s
= 0.4 cm3
/s
Return
Solution to Quick check 1
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Chapter 14

14. Effect of Temperature on the Speed of Reactions
 We know that food cooks faster when the temperature is
higher. For this reason, a pressure cooker is able to
cook red beans in 30 minutes compared to an ordinary
cooker which may take more than 2 hours. The
temperature in a pressure cooker is about 120 ºC
compared to 100 ºC in an ordinary cooker.
 Temperature is a very important factor in the speed of
reaction. In general, the rate of reaction increases two
times for about every 10 ºC rise in temperature.
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Factors Affecting the Speed of Reactions

15. How temperature affects the Speed of Reactions
 At higher temperature, the reacting particles move at higher
speeds as they have more kinetic energy.
 At higher speeds, the particles collide more often and with
greater force. This leads to more successful collisions and
hence increases the rate of reaction.
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16. Effect of particle size on the speed of reactions
 We know that meat and vegetables can be cooked more quickly
by cutting them into smaller pieces.
 This is because the smaller the size of the particles, the faster
the rate of a chemical reaction.
 When a solid is broken into smaller sizes, the surface area of
the solid is increased, thus exposing more particles of the solid
to the reactant, and more reactions can occur.
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17. Effect of Concentration on Speed of Reactions
 We all know that concentrated acids react more vigorously and faster
than dilute acids with metals and other reactants.
Experiment I Experiment II
 The speed of reaction in Experiment II was about two times as fast as in
Experiment I. This is because the concentration of the hydrochloric acid in
Experiment II was higher than that of Experiment I.
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18. Effect of concentration on the speed of reaction
 In general, the rate of reaction increases when the
concentration of one or more of the reactants is
increased.
 This is because a more concentrated solution contains
more particles per unit volume, so there will be more
particles to react with one another.
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19. Effect of pressure on the speed of reactions
 Pressure has very little effect on the rate of reactions in solids and liquids, because
they cannot be compressed.
 Pressure is important in gases because it has a great effect on the volume of gases.
 At higher pressure, gas particles are compressed closer together so there are more
particles per unit volume. This is equivalent to increasing its concentration thus
increasing the rate of reaction.
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20. The Collision Theory
 A chemical reaction only occurs when two particles (atoms
or molecules) collide into each other and bond together by
chemical forces.
 In order for the particles to be bonded together, the force
of collision must be great enough to overcome the initial
repulsive forces (the activation energy of the reaction).
 We can use the collision theory to explain the effect of
temperature and concentration on the rate of reaction.
Pow!Pow! CompoundCompound
formedformed
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21. How concentration affects the speed of reactions
 At higher concentration, the number of reacting particles
increases.
 The reacting particles are more crowded and there will be
a greater chance for them to meet, therefore resulting in
more collisions.
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22. Some everyday applications of the speed of reactions
 When cooking food, we cut them into smaller pieces and use
a higher temperature to make the food cook faster.
 To slow down the process of decay, food is kept at a low
temperature in a refrigerator.
 To make certain medicines work faster, they are often taken
in powder form and with warm water.
 Precaution must be taken in coal mines and flour mills to
prevent explosions due to the fine coal or flour dust particles.
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Chapter 14

23. Quick check 2
1. State 3 factors which affect the rate of reactions.
2. The graph below shows the results of an experiment done to compare the
rate of reaction between marble chips and marble powder with dilute
hydrochloric acid.
Graph
A
Graph B
VolofCO2
(a) Which graph shows the reaction between
the acid and (i) marble chips, (ii) marble
powder?
(b) Which graph shows that the rate of
reaction is faster? Explain why.
(c) At what time does the reaction between the marble chips and the acid stop?
(d) State one variable that must be kept constant when carrying out the experiment.
0 1 2 3 4 5 6 7 8 Time/ min
Solution
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24. 3. The following table shows the results of an experiment done to compare
the effect of concentration of sulphuric acid on magnesium.
Test tube No. 1 2 3 4 5
Volume of HCl/ cm3
50 40 30 20 10
Volume of H2O/ cm3
0 10 20 30 40
Total volume/ cm3
50 50 50 50 50
Time taken/ s 10 12 18 25 50
(a) Why are different volumes of water added to each test tube of acid?
(b) In which test tube is the concentration of the acid most concentrated?
(c) In which test tube is the concentration of the acid least concentrated?
(d) Plot a graph of the time taken for the magnesium to dissolve with the volume
of the acid used.
(e) What conclusion can you get from your graph? Solution
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Chapter 14

25. 1. Concentration of reactants, temperature and particle size.
2. (a) (i) Marble chips: graph B, (ii) Marble
powder: graph A
(b) Graph A is faster because it has a steeper gradient.
(c) 8 minutes after the start of the reaction.
(d) Concentration of the acid/ Temperature/ Mass of the
calcium carbonate.
Return
Solution to Quick check 2
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Chapter 14

26. (a) To make the total volume of each acid solution equal to 50 cm3
.
(b) Most concentrated - Test tube 1
(c) Least concentrated - Test tube 5
(d) A curve with decreasing gradient is obtained.
(e) The speed of the reaction decreases as the concentration of the acid
is decreased.
Graph of Vol. of acid vs Time taken
0
10
20
30
40
50
60
0 10 20 30 40 50 60
Time/s
Vol.ofAcid/cm3
3.
Return
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Chapter 14

27. What is a catalyst?
 A catalyst is a substance which changes the speed of
a chemical reaction, but is itself chemically unchanged
at the end of the reaction.
 Catalysts are very important for making slow chemical
reactions go faster.
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28.  A catalyst works by one or both
ways:
1. It provides an alternative
reaction pathway with lower
activation energy. More
particles are able to react
because of the lower activation
energy required.
2. A catalyst (often in finely
divided form) provides a large
surface area for the reactants
to adsorb and brings them into
close contact with one another.
How does a catalyst work?
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Chapter 14

29.  The chemical industry depends on catalysts for many of
the industrial processes.
 Examples are:
 Manufacture of ammonia: iron catalyst;
 Manufacture of sulphuric acid: vanadium(V) oxide;
 Manufacture of margarine: nickel catalyst;
 Catalytic converter in motorcars: platinum.
Importance of catalysts
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30. Enzymes
 Enzymes are biological catalysts found in plants and
animals. They are mainly made up of proteins.
 The enzymes in our bodies enable us to carry out our
bodily functions such as digestion of food and
absorption of nutrients.
 E.g. the enzyme amylase catalyses the conversion of
starch that we eat into sugars;
 Enzymes present in yeast are used in the making of
bread and wine.
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31. Catalytic decomposition of hydrogen peroxide
 We can show the effect of a catalyst
on the speed of a chemical reaction
by carrying out the experiment as
shown in the diagram.
 Hydrogen peroxide decomposes
rapidly when a little manganese(IV)
oxide is added as a catalyst.
2H2O2  2H2O + O2
 This process is usually used in the
preparation of oxygen in the
laboratory.
H2O2 + MnO2
oxygen
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32. Quick check 3
1. What is a catalyst? Give an example
of the use of a catalyst in a particular
chemical reaction.
2. The graph shows the catalytic
decomposition of hydrogen
peroxide.
2H2O2  2H2O + O2
(a) Which reaction is faster? State two
ways how
you can make the reaction faster.
(b) What is the total volume of oxygen
produced? Calculate the mass of
hydrogen peroxide decomposed.
Solution
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Chapter 14

33. 1. A catalyst is a substance which changes the speed of a chemical reaction, but is itself
unchanged after the reaction. E.g. iron in the manufacture of ammonia in the Haber
process.
2.
(a) Reaction A is faster. (i) Use manganese(IV) oxide as catalyst, (ii) Heat the
reacting mixture.
(b) Total volume of oxygen produced = 48 cm3
(0.002 mol)
Mass of hydrogen peroxide = 0.002 mol x 2 x 34 g mol-1
= 0.136 g
Solution to Quick check 3
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Chapter 14

34. 1. http://www.nelsonthornes.com/secondary/science/scinet/scinet/rea
ction/rates/content.htm
2. http://www.jghs.edin.sch.uk/mathscience/chemistrynotes/topic2.ht
ml
To Learn more about Speed of Chemical
Reactions, click on the links below!
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Chapter 14

35. References
 Chemistry for CSEC Examinations by
Mike Taylor and Tania Chung
 Longman Chemistry for CSEC by Jim
Clark and Ray Oliver
35