SUCCESS CHEMISTRY

Essay Questions

1 (a) Using polyethene as an example, explain the What are composite materials? Use two suitable

terms polymer and monomer. [4 marks] examples to explain the above statement.

(b) Name three examples of [10 marks]
(i) natural polymers and
(ii) synthetic polymers. [6 marks] (b) Mr Vellu found that an art sculpture made of
pure metal is easily dented in his workshop but if
State a use for each type of polymer. he were to use an alloy, the sculpture will not be
dented. Using one suitable example, describe an
(c) Explain briefly how sulphuric acid is manufactured experiment to show how you can compare the
hardness of an alloy with that of a pure metal.
in the industry. [10 marks]
[10 marks]
2 (a)
9 Composite materials are produced for the
purpose of improving the original materials
and to fulfill specific needs.

Experiments (a) Measure the diameters of the two depressions

1 Brass is a copper alloy that is used to make accurately and record in the spaces provided in
souvenirs and decorative items. Diagram 1 shows the
experimental set-up used to compare the hardness Diagram 2. [3 marks]
of pure copper and brass.
(b) Construct a table to show all the data in the
Diagram 1
The 1 kg weight is dropped onto the steel ball experiment. [3 marks]
bearing and the diameter of the depression formed
on the block is measured. (c) State the operational definition for alloy. [3 marks]
The experiment is repeated using a copper block to
replace the brass block. (d) What is the relationship between the diameter
The cross section of the diameter of depression of
the two materials is shown in Diagram 2. of the depression and the hardness of the

Diagram 2 materials? [3 marks]

(e) Referring to results obtained from the experiment,

state the conclusion that can be drawn from the

experiment. [3 marks]

(f) State three variables that must be kept constant

in this experiment. [3 marks]

2

Iron nails that are used in the construction of
buildings rust more than stainless steel nails
when exposed to rain.

Referring to the situation above, plan an experiment

to compare the rate of rusting of an iron nail and a

stainless steel nail. Your explanation should have the

following items:

(a) Statement of problem

(b) All the variables

(c) Statements of hypothesis

(d) List of materials and apparatus

(e) Procedure

(f) Tabulation of data [17 marks]

Manufactured Substances in Industry 294

1CHAPTER FORM 5

THEME: Interaction between Chemicals

Rate of Reaction

SPM Topical Analysis

Year 2008 2009 2010 2011

Paper 1 2 31 2 31 2 31 2 3

Section ABC ABC ABC ABC

Number of questions 2 1 – – – 6 1 – – 1 5 – 1 – 1 4 1 – – 1

ONCEPT MAP

Rate of reaction Measuring the speed of reaction Concentration-time graph
• Average speed is the amount of • from changes in the mass of • The gradient of the graph

reactant used up or the product reactant or product against time. indicates the rate of reaction.
formed per unit time. • from changes in the volume of • The rate of reaction decreases as
• Rate is proportional to —t—im——e—1t—a—k—e—n .
gas produced against time. the reaction proceeds.

Concentration Applications in daily activities
An increase in concentration • Combustion of charcoal
increases the speed of reaction. • Keeping food in a refrigerator
• Cooking food in a pressure cooker
Particle size RATE OF
A decrease in the particle size (larger REACTION Pressure
total surface area) increases the An increase in pressure increases
speed of reaction. the speed of reaction (applies only
to gases).
Temperature
An increase in temperature increases Catalyst
the speed of reaction. Catalyst increases the rate of reaction.

Uses of catalysts in industry

Collision theory • Iron in the Haber process
• Explains rate of reaction in terms of effective collisions between reactant
N2 + 3H2 2NH3
particles.
• For effective collisions, the particles must have energy equal to or greater than • V2O5 in the Contact process
2SO2 + O2 2SO3
the activation energy.
• Any factor that increases the rate of effective (successful) collisions will increase • Pt in the Ostwald process

the speed of reaction. 4NH3 + 5O2 4NO + 6H2O

1.1 Rate of Reaction Reactants: CaCO3(s) + 2HCl(aq) →
Products: CaCl2(aq) + CO2(g) + H2O(l)
The Meaning of Rate of Reaction
(b) During the reaction, the following
1 During a chemical reaction, the reactants are observable changes take place.
used up as the products are formed.
(i) The mass of calcium carbonate (the
Example: CaCO3 + 2HCl → CaCl2 + H2O + CO2 reactant) decreases.
Thus, the amounts of reactants decrease
(Figure 1.1(a)) while the amounts of (ii) The concentration of hydrochloric
products increase as the reaction proceeds acid (the reactant) decreases.
(Figure 1.1(b)).
(iii) The volume of carbon dioxide (the
product) produced increases.

(c) Thus, the rate of reaction between calcium
carbonate and hydrochloric acid can be
determined by measuring

(i) the decrease in mass of calcium
carbonate per unit time, or

(ii) the increase in volume of carbon
dioxide per unit time.

That is,

1 Reaction rate = —M——a—s—s—o—f—C—a–—C—O——3—r—e—a—c—te——d , or
Time taken
Figure 1.1 The graph of amount of substance (mol) Reaction rate = —V—o—l—u—m——e—o—f–—C—O——2—p—r—o—d—u—c—e——d
against time (minutes) Time taken

2 Definition The gas produced during a reaction can be collected
The rate of reaction is defined as the amount by using a burette or a gas syringe.

of a reactant used up or the amount of a 5 The rate of reaction is inversely proportional to
product obtained per unit time. the time taken for the reaction to be completed.

Rate of reaction = —A—m——o—u—n—t—o—f——re—a—c—t—a—n—t—u—s—e—d—u—p— Reaction rate ∝ —T—i—m—e—1—ta—k—e—n
Time taken

or

Rate of reaction = —A—m——o—u——n—t—o—f—p—r—o—d—u—c—t—o—b—t—a—i—n—e—d
Time taken

3 Methods of measuring reaction rates The reaction is fast if it takes a short time to
complete. Conversely, the reaction is slow if it
S PM The amount of a reactant used up or a product takes a long time for the reaction to complete.
’07/P1, 6 Example of a reaction involving a change in
’09/P1, obtained can be measured in terms of colour
’11/P1 (a) changes in the mass or concentration of
(a) The reaction between potassium
the reactant or product
manganate(VII), KMnO4, and ethanedioic
(b) volume of gas produced acid, H2C2O4, can be represented by the
ionic equation below.
(c) changes in colour

(d) formation of precipitate

(e) changes in mass of the reaction mixture 5C2O42–(aq) + 16H+(aq) + 2MnO4–(aq)

4 Reaction between calcium carbonate and ethanedioate ion manganate(VII)

dilute hydrochloric acid ion (purple)

(a) The reaction between calcium carbonate → 10CO2(g) + 8H2O(l) + 2Mn2+(aq)
colourless
(marble chips) and dilute hydrochloric

acid can be represented by the equation:

Rate of Reaction 296

(b) Observable changes: When excess ethanedioic Table 1.1 Example of some fast reactions SPM
acid solution is added to an aqueous solution ’09/P1
of potassium manganate(VII), KMnO4, the
purple colour of KMnO4 decolourises slowly Type of reaction Example
at room temperature.
Neutralisation Reaction between an acid and an
Reaction rate ∝ —T—i—m——e—t—a—k—e—n—f—1o—r—t—h—e—p—u——rp—l—e
colour to disappear alkali.

A concentrated solution of manganese(II) ions, Mn2+, HCl + NaOH → NaCl + H2O
is pink in colour. However, a very dilute solution of
Mn2+ ions appears colourless. Double Reaction between silver nitrate

decomposition solution and sodium chloride solution

to form silver chloride precipitate.

AgNO3(aq) + NaCl(aq)
→ AgCl(s) + NaNO3(aq)

Combustion Burning fuel to form carbon

dioxide and water.

CH4 + 2O2 → CO2 + 2H2O

7 Example of a reaction involving the formation Other fast reactions include 1
of a precipitate • burning of magnesium
(a) The reaction between sodium thiosulphate
and dilute hydrochloric acid is a slow 2Mg + O2 → 2MgO
reaction. • reaction of sodium or potassium with water

Na2S2O3(aq) + 2HCl(aq) → 2Na + 2H2O → 2NaOH + H2
2NaCl(aq) + H2O(l) + SO2(g) + S(s)
yellow Table 1.2 Example of slow reactions SPM
precipitate
’09/P1
(b) Observable changes: When dilute hydro­-
chloric acid is added to sodium thiosul­ Type of reaction Example

phate solution, the solution becomes Iron rusting Rusting takes place slowly in the
cloudy because sulphur is precipitated. presence of oxygen and water.
Sulphur is a yellow solid, but in small
quantities, it appears yellowish-white. Fermentation of 4Fe + 3O2 + 2H2O → 2Fe2O3•H2O
glucose solution
Rate of reaction ∝ —T—i—m——e—t—a—k—en—1—f—o—r—a—g—i—v—e—n rust
amount of sulphur
precipitate to form In the presence of yeast,
fermentation of glucose solution
8 The units used for the rate of reaction will produces alcohol and carbon
depend on the changes measured. For example, dioxide.
(a) cm3 per unit time (second or minute) for
a gas evolved Photosynthesis C6H12O6 → 2C2H5OH + 2CO2
(b) g per unit time or mol per unit time for a
solid reactant glucose alcohol (ethanol)
(c) mol dm–3 per unit time for a reactant in
aqueous solution During photosynthesis, carbon
dioxide reacts with water to form
9 Different chemical reactions take place at glucose and oxygen gas.
different rates. Some reactions occur rapidly and
some slowly. Table 1.1 shows some examples 6CO2 + 6H2O → C6H12O6 + 6O2
of fast reactions. Table 1.2 shows some examples
of slow reactions. glucose

• The reactions of Groups 1 and 2 metals with oxygen is
a fast reaction. However, the reactions of other metals
(such as copper) with oxygen are slow reactions.

• The rate of decay of the radioactive carbon-14 is very
low. For example, 1.0 g of carbon-14 takes 5730 years
to disintegrate (decay) to 0.50 g. The rate of decay of
carbon-14 is used in archaeology to estimate the age of
ancient artifacts. This method is called carbon dating.

297 Rate of Reaction

2

The rate of a reaction depends on various factors A piece of magnesium ribbon weighing 0.1 g is added
(Section 1.2). For example, the rate of rusting of iron to dilute hydrochloric acid. After 5 seconds, all the
is increased if the iron is exposed to acid (such as magnesium had dissolved. What is the average rate
polluted air in industrial areas) or to the salt, sodium of reaction?
chloride in sea air.
Solution

Measuring Reaction Rates Average rate of = Mass of magnesium reacted
reaction Time taken

1 The rate of reaction can be expressed in two = 0.1 g = 0.02 g s–1
5s
SPM ways:
’04/P1,
’10/P1, (a) the average rate of reaction over a period The value obtained is the average rate of
’11/P1 of time, or reaction over a period of 5 seconds.

(b) the rate of reaction at any given time.

2 The average rate of reaction is the average of 4 The average rate
of reaction can
the reaction rates over a given period of time. also be determined
from the graph.
We can measure the average rate of reaction Based on Figure
1.2, the average
by measuring the change in amount (or rate of the first t1
second
1 concentration) of a reactant or a product over
= —V—1 cm3 s–1
a period of time. t1

3 For example, the average rate of reaction between

magnesium and hydrochloric acid can be

determined by measuring the time taken for a

piece of magnesium to dissolve completely in Figure 1.2

the acid.

1 SPM The average rate of reaction from t1 to t2
’04/P1 = —(—V—2—–——V—1) cm3 s–1
(t2 – t1)
Calcium carbonate reacts with dilute hydrochloric
acid to form carbon dioxide. 5 The rate of reaction at any given time is the
After 1.2 minutes, the volume of gas produced is actual rate of reaction at a given time. The
100 cm3. Calculate the average rate of reaction in
units of (a) cm3 min–1, (b) cm3 s–1. reaction rate at any given time is also known as
the instantaneous rate of reaction.
Solution 6 The rate of reaction at a given time can be
determined by the following methods.
Volume of CO2
(a) Average reaction rate = ————p—ro—d—u—c—e—d———— (a) By measuring the gradient of the graph of
mass of reactant against time (Figure 1.3).
Time taken
= —1—0—0——c—m—3 Figure 1.3 Measuring the rate of
1.2 min reaction involving a change
= 83.3 cm3 min–1 in mass at a given time


Volume of CO2
(b) Average reaction rate = ————p—ro—d—u—c—e—d———

Time taken
= —8—3—.3——c—m—3——m—i—n—–1
60 s min–1
= 1.39 cm3 s–1

Rate of Reaction 298

Determining the gradient of the tangent at time, t: Analysing a reaction rate curve
The following steps are used to determine the (i) The steeper the gradient, the higher the rate
gradient of the tangent at time, t (Figure 1.3).
of reaction.

Step 1
Draw the tangent XY at point P.

Step 2 Figure 1.5 Comparing the rates of reaction for
Complete the right-angled triangle XYZ. a given reaction at different times

Step 3 (ii) Figure 1.6 shows the graph of volume of hydrogen
Measure the lengths of XZ and ZY. against time for the reaction between excess
zinc powder and dilute hydrochloric acid.

Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)

1

Step 4 Figure 1.6
Find the gradient of the line XY.
(a) Initially (t1),
Gradient of the line XY = —ab • the graph is steep,
• the rate of reaction is high.

= rate of reaction at (b) As the reaction proceeds (t2),
• the graph is less steep,
time, t (g s–1) • the rate of reaction
decreases because
(b) By measuring the gradient of the graph of the concentration
of hydrochloric acid
volume of gas produced (product) versus time decreases.

(Figure 1.4). (c) Finally (t3),
• the graph becomes
Figure 1.4 Measuring the rate of reaction at a horizontal,
given time involving a change in the • the gradient of the graph
volume of a gas becomes zero,
• the reaction stops because
Grad ient = —ab all the hydrochloric acid
= rate of reaction at time, t has reacted.
(cm3 min–1)
Rate of Reaction
299

3

Consider the reaction between excess magnesium There are some parameters which cannot be
and dilute sulphuric acid: measured accurately to determine the instantaneous
rate of reaction, for example the change in colour or
Mg(s) + H2SO4(aq) → MgSO4(aq) + H2(g) the formation of precipitate.
The reaction stops at 20 seconds.
Plot the graph of (c) (d)
(a) mass of magnesium against time,
(b) concentration of sulphuric acid against time,
(c) concentration of magnesium sulphate against time,
(d) volume of hydrogen gas against time.

Solution (b)
(a)

Activity 1.1 To find the reaction rates at (a) 90 s, (b) 180 s and (c) the
1 average rate of the reaction between zinc and dilute sulphuric
acid

Apparatus Figure 1.7 Procedure

Materials Conical flask, measuring cylinder, 1 The burette is filled with water and inverted over
Results delivery tube, burette, basin, retort a basin of water.
stand, retort clamp and stopwatch.
2 Using a measuring cylinder, 20.0 cm3 of 0.3 mol
Granulated zinc and 0.3 mol dm–3 dm–3 sulphuric acid is measured out and poured
sulphuric acid. into a conical flask.

3 5.0 g of granulated zinc is then added to the
sulphuric acid in the conical flask.

4 The conical flask is then closed and the hydrogen
gas produced is collected in the burette by the
displacement of water as shown in Figure 1.7.

5 The stopwatch is started immediately.
6 The volume of hydrogen gas collected in the

burette is recorded at 30-second intervals.

Time (s) 0 30 60 90 120 150 180 210 240 270 300 330 360

Burette reading (cm3) 50.00 33.00 24.50 18.00 13.00 10.00 6.50 5.00 4.00 3.50 3.00 3.00 3.00

Volume of H2 released 0.00 17.00 25.50 32.00 37.00 40.00 43.50 45.00 46.00 46.50 47.00 47.00 47.00
(cm3)

Based on the experimental results, a graph of the volume of hydrogen released against time is plotted.

Rate of Reaction 300

Calculation

(a) The rate of reaction at 90 s

= gradient of the curve at 90 s

= —XY—ZY— = —((5—12—8—–0—–—2–—03—)0—c)—ms—3


= —31—25—c0—m—s3 = 0.213 cm3 s–1


(b) The rate of reaction at 180 s

= gradient of the curve at 180 s

= —QP—Q—R = —((—42—84—–0——3–—01—)8—c)—m—s3


= —12—82—c2—m—s3 = 0.081 cm3 s–1


(c) The average rate of reaction

= —T—o—t—a—l —vT—oo—ltu—am—l —tei—mo—fe——Ht—a2k—pe—nr—o—d—u—c—e—d Conclusion

The rate of reaction decreases as the reaction
= —43—70—c0—m—s3 = 0.157 cm3 s–1 proceeds.


To measure the rate of reaction between calcium carbonate 1
(CaCO3) and excess hydrochloric acid Activity 1.2

Apparatus Conical flask, electronic balance, Procedure
Materials measuring cylinder and stopwatch.
Calcium carbonate (CaCO3) pieces, 1 Using a measuring cylinder, 50 cm3 of 2 mol
Results 2.0 mol dm–3 hydrochloric acid and dm–3 hydrochloric acid is measured out and
cotton wool. poured into a dry conical flask. The mouth of the
conical flask is covered with some cotton wool.
Figure 1.8 The cotton wool is inserted into the mouth of the
conical flask to prevent liquid from splashing out
during the reaction.

2 The conical flask is placed on the electronic
balance as shown in Figure 1.8.

3 The mass of conical flask, calcium carbonate,
hydrochloric acid and cotton wool is recorded.

4 The calcium carbonate is then transferred to the
hydrochloric acid in the conical flask and the
stopwatch is started immediately.

5 The mass of the conical flask (and its contents)
is recorded at one-minute intervals.

Time (min) 012345678

Mass of conical 60.0 59.1 58.3 57.9 57.4 57.0 56.8 56.5 56.3
flask + contents (g)

Based on the experimental results, a graph of the mass of conical flask and its contents against time is plotted
(Figure 1.9).

301 Rate of Reaction

= —1—.00—.—9m—–gi—n


= 0.9 g min–1

(b) The average rate of reaction between 1.4 minutes

and 2.2 minutes.

Rate of decrease in mass

= —((—25—.82—.8—–—–1—.54—8)—.—3m—)i—gn From the graph (Figure 1.9)


= 0.625 g min–1

(c) The reaction rate at the 5th minute

= gradient of the graph at the 5th minute = —ba


a = 57.5 – 56.4

Figure 1.9 = 1.1 g

Calculation b = 7.0 – 3.4

= 3.6 minutes

(a) The average rate of reaction for the first minute. Gradient = —3—.16—.1m—gi—n

Decrease in mass = mass of carbon dioxide

1 produced = 0.306 g min–1

= (60.0 – 59.1) g See table of results. Conclusion
= 0.9 g
The rate of reaction decreases as the reaction
Average rate of reaction for the first minute proceeds. Finally, the reaction will stop when all the
calcium carbonate added have reacted.
= —M——a—ss—T—o—fim—C—eO—t2—apk—re—on—d—u—c—e—–d


SPM Solving Numerical Problems Involving
Rate of Reaction
’08/P2
1 The rate of reaction can be stated in terms of
The graph below shows the total volume of oxygen (a) the average rate of reaction for a given
gas produced against time for the decomposition of period of time, or
hydrogen peroxide. (b) the rate of reaction at any given time
(instantaneous rate).

2 The average rate of reaction can be calculated
(a) directly from the data given (Example 2) or
(b) from the graph drawn (Example 4).

3 The reaction rate at a given time can only be
obtained from the gradient of the graph at the
given time (Example 5).

At time, t, the maximum volume of oxygen is collected. The rate of reaction is useful to a chemist because he
The gradient of the curve at time, t is zero. Hence, is not satisfied with merely converting one substance
the rate of reaction is zero, that is, the reaction has to another. In most cases, he wants to obtain the
stopped at time, t. products in the fastest and most economical way.

Rate of Reaction 302

4 5

3.0 g of excess marble (CaCO3) are added to 100 Hydrogen peroxide decomposes as represented by
cm3 of dilute hydrochloric acid. Figure 1.10 shows the equation:
the graph of volume of carbon dioxide produced
against time. 2H2O2(aq) → 2H2O(l) + O2(g)

The results of an experiment on the decomposition
of hydrogen peroxide are given below.

Time (s) 0 15 30 45 60 90

Volume of O2 (cm3) 0 16 30 40 48 56

Calculate the rate of reaction at 40 seconds in units
of (a) cm3 s–1, (b) cm3 min–1.

Solution
(a)

1

Figure 1.10

Calculate
(a) the average rate of reaction,
(b) the concentration of hydrochloric acid in mol

dm–3.
[1 mol of any gas occupies 24 dm3 at room
conditions]

Solution

(a) Total volume of carbon dioxide evolved
= 360 cm3
Time taken = 8.0 minutes
A verage rate of reaction = —3—86—0—m—ci—mn—3
= 45 cm3 min–1

(b) Number of moles of CO2 evolved
= —(—2—4——3—61—00—c0—m0—)—3 —c—m—3 = 0.015

CaCO3 + 2HCl → CaCl2 + H2O + CO2 The rate of reaction at 40 s
Mole ratio of HCl : CO2
= gradient at 40 s

=2:1 From equation = —ab = —(—4(—95—8–—–—2—11—)8—)c—ms—3 obtained from
? : 0.015 the graph

According to the equation, number of moles of = 0.70 cm3 s–1
hydrochloric acid used = 2  0.015 = 0.03 mol.
Concentration of hydrochloric acid (b) 1 minute = 60 seconds
= —NV—u—om—lu—bm—e—er—o(—ifn—m—d—om—le—3)s = —00—..0—13—d—mm——o3l
= 0.3 mol dm–3 ∴Rate of reaction in cm3 min–1

From the question: 100 cm3 = 0.1 dm3 = 0.70 cm3 s–1  60 s min–1
= 42 cm3 min –1

303 Rate of Reaction

1 ’09

Which of the following is correctly matched with its rate of reaction?

High reaction rate Low reaction rate
A Combustion of fuels
B Combustion of fuels Respiration

C Rusting of iron Double decomposition between silver nitrate and sodium
D Respiration chloride solution

Fermentation of glucose solution

Neutralisation reaction between an acid and an alkali

Comments
Combustion, double decomposition and neutralisation are fast reactions. Rusting and respiration are slow
reactions.

Answer A

1

1.1 Figure 1.11

1 Which of the following reactions occur at (a) a high (a) What is the total time required for the magnesium
rate, (b) a low rate? ribbon to react completely with hydrochloric
acid?
I Fe3+(aq) + 3OH–(aq) → Fe(OH)3(s)
II 2Cu(s) + O2(g) → 2CuO(s) (b) Based on the graph, is the reaction rate at the
III S2O32–(aq) + 2H+(aq) → S(s) + H2O(l) + SO2(g) first minute higher or lower than the reaction at
IV 4K(s) + O2(g) → 2K2O(s) the second minute? Explain your answer.
2 You are given the chemicals and apparatus as listed
(c) Is this a normal behaviour? Suggest one reason
below. for this behaviour.
• A piece of zinc of mass 2.0 g
• A beaker containing sulphuric acid (d) The reaction between hydrochloric acid and
• A stopwatch another metal produces 12 cm3 of hydrogen
(a) Using the chemicals and apparatus given, after 1.0 minute. Is this reaction rate higher or
lower than the reaction between magnesium
describe an experiment to measure the rate of and hydrochloric acid? Explain your answer.
reaction between zinc and sulphuric acid.
(b) State the units for the rate of reaction.
(c) State two assumptions for this experiment.

3 A student intends to study the rate of reaction
between iron and dilute sulphuric acid. The equation
for the reaction is as follows.

Fe(s) + H2SO4(aq) → FeSO4(aq) + H2(g)

Suggest two methods that he can use to measure
the rate of reaction.

4 The graph in Figure 1.11 shows the results of
an experiment to measure the rate of reaction
of magnesium ribbon with an excess of dilute
hydrochloric acid.

Rate of Reaction 304

5 The table below shows the results for two experiments What is the average rate of reaction?
carried out under room conditions. (b) 4.0 g of magnesium is added to excess dilute

Experiment Reaction Result sulphuric acid. If the average rate of reaction is
0.0030 mol s–1, what is the mass of magnesium
I 1 g of nickel powder Time taken to unreacted after 0.5 minute?
+ 50 cm3 of 1 mol collect 60 cm3 [Relative atomic mass of Mg = 24]
dm–3 hydrochloric of hydrogen
acid gas = 120 s 7 In the presence of manganese(IV) oxide, hydrogen
peroxide decomposes according to the equation:
II 1 g of zinc powder Time taken to
+ 50 cm3 of 1 mol collect 45 cm3 MnO2
dm–3 hydrochloric of hydrogen 2H2O2(aq) ⎯⎯→ 2H2O(l) + O2(g)
acid gas = 56 s
A sample of hydrogen peroxide decomposed in the
Based on the information given in the table above, presence of a catalyst and the volume of oxygen
predict which metal is more reactive, nickel or zinc? gas produced was collected at regular time intervals.
The results of the experiment were recorded in the
6 (a) The volume of hydrogen gas collected at regular following table.
intervals for the reaction between granulated
zinc and dilute hydrochloric acid is shown below. Time (min) 0 1 1—12 2 2—12 3 4 5

Time (s) Volume of H2 (cm3) Volume of 1
0 0 O2 (cm3)
20 16 0 32 46 56 64 69 74 74
40 26
60 32 Calculate
70 36 (a) the average rate of reaction for the first 144
80 36
seconds.
(b) the average rate of reaction for the overall

reaction in cm3 s–1.
(c) the average rate of reaction between the first

minute and the 3rd minute.
(d) the rate of reaction at the 150th second.

1.2 Factors that Affect the
Rate of Reaction SPM
Uranium is the radioactive isotope used for making
’08/P1 nuclear bombs. Uranium decays slowly to form lead.
The decay of uranium and other radioactive elements
1 The rate of reaction is affected by the following is unique. These nuclear reactions are not influenced by
factors: factors such as surface area, temperature and catalyst.
(a) Total surface area (or particle size) of the
solid reactant 4 Reaction involving gases
(b) Concentration of reactant (a) Changes in pressure will not affect
(c) Temperature of reaction reactions in aqueous solutions.
(d) Use of catalyst (b) Changes in pressure will only affect
(e) Pressure (for reactions involving gases) reactions involving gases.
(c) Increasing the pressure will compress
2 When the condition of reaction changes, the more gas molecules into a given space.
rate of reaction also changes. Hence the gaseous particles will collide
more frequently and the rate of reaction
3 Table 1.3 explains briefly how these conditions increases.
of reaction affect the rate of reaction between
zinc metal and dilute sulphuric acid.

305 Rate of Reaction

Table 1.3

Surface area (particle size) • When zinc foil is broken into
Concentration of reactant smaller pieces, the total surface area
Temperature of reaction increases.

• The smaller the size of zinc foil, the
greater the total surface area exposed
to the hydrogen ions. Hence the rate
of reaction increases.

• In dilute acid, there are not so many
hydrogen ions present.

• In more concentrated acid, there are
more hydrogen ions in the solution.
Hence the rate of reaction increases.

1

• When the temperature of a reaction
increases, the particles move faster
because they have higher kinetic
energy. Hence the rate of reaction
increases.

Catalyst A catalyst will increase the rate of
Some catalysts used in industry. The catalysts are in pellet form reaction. This will be explained in
for larger surface area. Section 1.3.

SPM Factors that Influence the Rate of Reaction • Manipulated variable: Size (total surface area)
of magnesium
’10/P2
• Responding variable: Time taken to collect
Effect of Surface Area on the Rate of Reaction 60 cm3 of hydrogen gas

1 Two experiments are carried out to study the • Constant variables: Temperature, concentration
rate of reaction between magnesium and dilute and volume of sulphuric acid as well as
sulphuric acid under different conditions. mass of magnesium

Mg(s) + H2SO4(aq) → MgSO4(aq) + H2(g) 2 The results of the experiments are shown
below.

Experiment Conditions of experiment Time taken to collect Average rate of
60 cm3 of H2 (s) reaction (cm3 s–1)
I 1.0 g of magnesium ribbon and 50 cm3 30
of 1.0 mol dm–3 sulphuric acid 2
10
II 1.0 g of magnesium powder and 50 cm3 6
of 1.0 mol dm–3 sulphuric acid

The size (surface area) of magnesium is manipulated

Rate of Reaction 306

3 The results show that the time taken to collect 1 mol dm–3 hydrochloric acid. This means that
60 cm3 of gas using magnesium powder is the higher the concentration of hydrochloric
shorter than using magnesium ribbon. This acid, the higher the rate of reaction.
is due to the smaller size of particles (total 4 (a) It is important to know
surface area is greater) in magnesium powder
than in magnesium ribbon. • how to plot graphs (on the same axis),
or
Effect of Concentration of Reactant on the Rate
of Reaction • how to interpret graphs on rate of
reaction if information on the conditions
1 Two experiments are carried out to study the of reaction are given.

SPM rate of reaction between magnesium and (b) The two features on the graph to be
’04/P1 considered are
’05/P3 hydrochloric acid. • the gradient of the graph which shows
’06/P1 the rate of reaction,
• the height of the graph which shows
Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g) the total amount of product formed.

• Manipulated variable: Concentration of Effect of Temperature on the Rate of Reaction 1
hydrochloric acid
1 Two experiments are carried out to study the
• Responding variable: Time taken for rate of reaction between zinc and sulphuric
magnesium to dissolve completely acid at different temperatures.

• Constant variables: Size of magnesium Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)
ribbon, volume of hydrochloric acid and
temperature of experiment • Manipulated variable: Temperature of sulphuric
acid
2 The results of the experiments are shown below.
• Responding variable: Volume of hydrogen
Experiment Conditions of Time taken for gas evolved
I experiment magnesium to
• Constantvariables: Massofzinc,concentration
II 5 cm magnesium dissolve and volume of sulphuric acid
ribbon and 50 cm3 of completely (s)
1 mol dm–3 2 The results are shown in Figure 1.12.
hydrochloric acid 78

5 cm magnesium 39
ribbon and 50 cm3 of
2 mol dm–3
hydrochloric acid

Concentration of acid is manipulated Figure 1.12 Comparing the rates of reaction at
different temperatures
3 The shorter the time taken for a reaction to
complete, the higher the rate of reaction. 3 The results show that the higher the temperature
The results show that the time taken for of sulphuric acid, the steeper the graph and
magnesium to react completely in 2 mol hence, the higher the rate of reaction.
dm–3 hydrochloric acid is shorter than that in

Experiment Mass of zinc Volume of Concentration of Temperature of
sulphuric acid sulphuric acid sulphuric acid
I 1.0 g of zinc powder
II 1.0 g of zinc powder 20 cm3 0.1 mol dm–3 28 °C

20 cm3 0.1 mol dm–3 35 °C

Conditions remain unchanged Temperature of acid is manipulated

307 Rate of Reaction

4 In general, the rate of reaction increases if the occurs. In this reaction, manganese(IV) oxide acts
temperature of the reactants is increased. as a catalyst and catalyses the decomposition
of hydrogen peroxide to give oxygen gas and
Effect of Catalysts on the Rate of Reaction water.

1 Definition MnO2
A catalyst is a substance that increases the rate 2H2O2(aq) ⎯⎯⎯→ 2H2O(l) + O2(g)
of a reaction but is itself chemically unchanged
at the end of the reaction. 4 The reaction between zinc and dilute acid is a
slow reaction.
2 In contrast, a substance that decreases the rate
of a chemical reaction is called an inhibitor. Zn(s) + 2H+(aq) → Zn2+(aq) + H2(g)

3 At room temperature, hydrogen peroxide When it is catalysed by copper(II) sulphate
decomposes very slowly. But when a very small solution, the reaction speeds up.
amount of manganese(IV) oxide is added to
hydrogen peroxide, a vigorous effervescence

The Characteristics of Catalysts SPM • The catalyst remains chemically unchanged after
the reaction.
Only a small amount ’06/P1,
of catalyst is needed ’08/P1, • Thus the chemical properties, mass and chemical
to increase the rate of ’10/P1, composition of the catalyst remain unchanged at
reaction. ’11/P1 the end of the reaction.

1

The physical appearance Characteristics A catalyst lowers the
of a catalyst may change at of a catalyst activation energy of a
the end of the reaction. For reaction (see Section 1.3
example, small pieces of – The Collision Theory)
catalyst may become fine
powder after the reaction.

A catalyst increases the rate of a chemical reaction (a) In general, catalysts are highly specific.
but it does not change (increase or decrease) the For example, iron catalyses the reaction:
yield of a chemical reaction.
N2 + 3H2 2NH3
Zn(s) + H2SO4(aq) ⎯cCa⎯utaSl⎯Oys4t→ ZnSO4(aq) + H2(g)
but not the reaction:

2SO2 + O2 2SO3
(b) However, some catalysts can catalyse

several different reactions. For example,

MnO2 can catalyse the following
reactions:

2H2O2(aq) ⎯M⎯nO⎯2→ 2H2O(l) + O2(g)
2KClO3(s) ⎯M⎯n⎯O2→ 2KCl(s) + 3O2(g)

Catalysts can be poisoned by impurities. When a catalyst is poisoned, its effectiveness as a catalyst is decreased.

Rate of Reaction 308

Examples of Catalysts

1 Transition metals and compounds of transition metals are often used as catalysts for
industrial processes.

2 Table 1.4 shows some examples of catalysts and the reactions catalysed by them.

Table 1.4 Some common catalysts Catalyst used
Type of reaction Iron, Fe
(a) Haber process for the manufacture of ammonia.

N2(g) + 3H2(g) 2NH3(g)

(b) Contact process for producing sulphur trioxide. Vanadium(V) oxide, V2O5
Platinum, Pt
2SO2(g) + O2(g) 2SO3(g)

Sulphur trioxide is used for the manufacture of sulphuric acid.

(c) Ostwald process for producing nitrogen monoxide.

4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g)

Nitrogen monoxide is used for the manufacture of nitric acid.

(d) Manufacture of margarine Nickel, Ni 1
In the presence of a catalyst at 200 °C, vegetable oils react with
hydrogen to produce margarine. This process is called hydrogenation. Aluminium oxide, Al2O3
or
(e) Cracking process
When big alkane molecules are passed over a catalyst at 600 °C, a Silicon(IV) oxide, SiO2
mixture of small alkane and alkene molecules is produced.
This process is called catalytic cracking.
(Refer Sections 2.2 and 2.3 on alkanes and alkenes)

2 ’06 Effect of Pressure on the Rate of Reaction

Which of the following statements about catalysts 1 The changes in pressure will only affect
reactions involving gases. An increase in
are true? pressure increases the rate of reaction. In
contrast, a decrease in pressure decreases the
I A catalyst is specific in its reaction. rate of reaction. In the following reactions
involving gases, the rate of reaction increases
II A catalyst changes the quantity of product if the pressure is increased.

formed. N2(g) + 3H2(g) 2NH3(g) (Haber process)

III Only a small amount of a catalyst is needed to 2SO2(g) + O2(g) 2SO3(g) (Contact process)

change the rate of reaction. 2 Pressure has no effect on reactions involving
only solids or liquids. For example, the
IV The chemical properties of the catalyst remain following reactions are not affected by changes
in pressure.
unchanged at the end of a reaction.
CaCO3(s) + 2HCl(aq) →
A I and II only C I, II and III only CaCl2(aq) + H2O(l) + CO2(g)

B II and IV only D I, III and IV only 2H2O2(aq) → 2H2O(l) + O2(g)

Answer D

A catalyst takes part in a chemical reaction. In actual fact, a
catalyst combines with the reactants to form an unstable
intermediate species. This species then decomposes to
re-form the catalyst and to produce the products.

309 Rate of Reaction

1.1

To investigate the effect of the surface area of a reactant on the rate of reaction SPM

’06/P2

Problem statement Procedure

How does the surface area of a solid reactant affect
the rate of reaction?

Hypothesis

The smaller the size of the marble chips, that is, the
larger the total surface area of the marble chips, the
higher the rate of reaction.

Experiment 1.1 Variables Figure 1.13
1
(a) Manipulated variable : Size of the marble chips 1 A burette is filled with water and inverted over
used a basin containing water. The burette is clamped
to the retort stand. The water level in the burette
(b) Responding variable : Volume of gas given off at is adjusted and the initial burette reading is
30-second intervals recorded.

(c) Constant variables : Temperature of the 2 5.0 g of marble chips are placed in a small
experiment, mass of marble chips, concentration conical flask.
and volume of hydrochloric acid
3 50 cm3 of 0.08 mol dm–3 hydrochloric acid is
Apparatus Conical flask, delivery tube, retort added to the marble chips.
stand and clamp, burette, measuring
cylinder and stopwatch. 4 The conical flask is then stoppered and the
stopwatch is started immediately (Figure 1.13).
Materials Marble chips, powdered marble and
0.08 mol dm–3 hydrochloric acid. 5 The burette readings are recorded at 30-second
intervals.
Experiment I
The rate of reaction using large marble chips

Results

Time (s) 0 30 60 90 120 150 180 210 240

Burette reading (cm3) 50.00 45.50 41.50 38.00 35.00 33.00 31.00 29.00 28.00

Volume of gas (cm3) 0.00 4.50 8.50 12.00 15.00 17.00 19.00 21.00 22.00

Experiment II The rate of reaction using powdered marble Constant variable is also known as
Procedure fixed variable or controlled variable.

1 Steps 1 to 4 in Experiment I are repeated using 5.0 g of powdered marble. All other conditions such as

temperature, volume and concentration of hydrochloric acid are kept constant.
2 The results of the experiment are recorded in the following table.

Results

Time (s) 0 30 60 90 120 150 180 210 240

Burette reading (cm3) 50.00 42.00 35.00 29.50 25.50 22.00 19.50 17.50 16.00

Volume of gas (cm3) 0.00 8.00 15.00 20.50 24.50 28.00 30.50 32.50 34.00

Rate of Reaction 310

Based on the results obtained, a graph of the total volume of the hydrochloric acid used in both the 1
volume of carbon dioxide produced against time experiments are the same.
for each experiment is plotted on the same axes 3 The gradients of the graphs for Experiments
(Figure 1.14). I and II become less steep as the reactions
proceed. This shows that the rates of reaction
Figure 1.14 (a) are very high at the beginning of the
Discussion
1 Figure 1.15 shows the graphs that will be reactions,
(b) decrease as the reactions proceed,
obtained if the reactions in Experiments I and II (c) become zero when the reactions are
are completed.
completed. At this time, the graphs become
horizontal.
4 The rate of reaction between the marble and
hydrochloric acid decreases because
(a) the mass of the remaining unreacted marble
decreases,
(b) the concentration of hydrochloric acid
decreases.
5 The reaction in Experiment I stops after t2 minutes
while the reaction in Experiment II stops after t1
minutes, where t1 < t2. This shows that the rate of
reaction for Experiment II (powdered marble) is
higher than the rate of reaction for Experiment I
(marble chips).
6 The total volume of carbon dioxide collected
in the burette is usually slightly less than the
theoretical value (48 cm3 for the experiment
above). This is because carbon dioxide is slightly
soluble in water. To overcome this problem, a gas
syringe is used to collect carbon dioxide released
during the experiment (Figure 1.16).

Same maximum volume of CO2 Figure 1.16 Measuring the volume of gas using a
collected because mass of CaCO3, gas syringe
concentration and volume of HCl are
kept constant. Conclusion

Figure 1.15 Graph II is steeper than graph I. This shows that the
rate of reaction in Experiment II is higher than the
2 Figure 1.15 shows that both graphs level off at rate of reaction in Experiment I as powdered marble
the same value. This indicates that the maximum is used in Experiment II. Thus, the rate is higher with
volume of carbon dioxide collected at the end powdered marble than with marble chips. Hence, we
of reaction for both Experiments I and II are can conclude that the smaller the particle size, the
the same (that is, 44 cm3). This happens because larger the total surface area exposed for reaction
the mass of the marble, concentration and and the higher the rate of reaction. The hypothesis
is accepted.

311 Rate of Reaction

1.2

To study the effect of concentration on the rate of reaction between sodium SPM
thiosulphate solution and dilute sulphuric acid
’07/P1,
’11/P2

Problem statement Apparatus

How does the concentration of a reactant affect the 10 cm3 and 100 cm3 measuring cylinders, 100 cm3
rate of reaction between sodium thiosulphate and conical flask, white paper marked with a cross ‘X’
dilute sulphuric acid? and stopwatch.

Materials

0.2 mol dm–3 sodium thiosulphate solution, 1.0 mol
dm–3 sulphuric acid and distilled water.

Experiment 1.2 Figure 1.17 Procedure
1
Hypothesis 1 50 cm3 of 0.2 mol dm–3 sodium thiosulphate
solution is measured out using a 100 cm3 measuring
The more concentrated the sodium thiosulphate cylinder. The solution is then poured into a clean,
solution, the higher the rate of reaction. dry conical flask.

Variables 2 The conical flask is placed on a piece of paper
(a) Manipulated variable: Concentration of sodium with a cross ‘X’ marked on it (Figure 1.17).

thiosulphate solution 3 5 cm3 of dilute sulphuric acid is measured out
(b) Responding variable: Time taken for the cross by using a 10 cm3 measuring cylinder. The acid
is then quickly poured into sodium thiosulphate
‘X’ to disappear solution. The stopwatch is started immediately.
(c) Constant variables: Concentration and volume of
4 The reaction mixture is swirled once and the cross
dilute sulphuric acid as well as the temperatures ‘X’ is viewed from above. A yellow precipitate
of the solutions will appear slowly in the conical flask.

Results 5 The stopwatch is stopped as soon as the cross
disappears from view and the time taken is recorded.

6 Steps 1 to 5 are repeated with different mixtures
of sodium thiosulphate solution and distilled
water as shown in the following table.

Experiment 1 2 3 4 5 Different volumes (V1) of
50 40 30 20 10 Na2S2O3 solution are
Volume of Na2S2O3 (cm3) 0 10 20 30 40 diluted with water to make
Volume of water (cm3) 5 5 5 5 5 up to 50 cm3 solution (V2).
0.20 0.16 0.12 0.08 0.04
Volume of H2SO4 (cm3) 24 30 42 62 111
Concentration of Na2S2O3 (mol dm–3)
Time taken (s) 0.042 0.033 0.024 0.016 0.009

—T—i1—m—e (s–1)

Discussion The ionic equation is as follows:

1 The following equation shows the reaction S2O32–(aq) + 2H+(aq) → S(s) + SO2(g) + H2O(l)
between sodium thiosulphate, Na2S2O3 and dilute
sulphuric acid: The sulphur is precipitated as fine yellow
Na2S2O3(aq) + H2SO4(aq) → particles that cause the solution to turn cloudy.
Na2SO4(aq) + H2O(l) + SO2(g) + S(s)

Rate of Reaction 312

2 As the amount of sulphur increases, the cross ‘X’
becomes more and more difficult to see. Finally,
the cross ‘X’ disappears from view when a certain
mass of sulphur is precipitated. Hence, the time
recorded for the disappearance of the cross ‘X’ is
the time taken for the formation of a fixed mass
of sulphur.

3 R ate of reaction = —M——a—s—s—o—fT—si—um—lpe—h—tau—kr—ep—nr—o—d—u—c—e—d
H ence, rate of reaction ∝ —T—i—m—e——ta—k—e—n—1f— o—r—t—h—e—c—r—o—ss

‘X’ to disappear

4 The concentration of sodium thiosulphate solution Figure 1.19 1
after mixing with water can be obtained by using
the following formula: 7 The conical flask used for each experiment must
C oncentration of Na2S2O3 = —M—V—1V—2 —1 have the same size (for example, 100 cm3 volume).
If the conical flask of a larger size is used, the
= —0—.—2———V——o—lu—m——e5—o0—f—N——a2—S—2—O—3—u—s—e—d mol dm–3 time, t, taken for the cross ‘X’ to disappear will
increase. Conversely, if a smaller conical flask
5 Based on the results obtained, two graphs can be is used, the time taken for the cross to disappear
plotted. will be shorter.
(a) The graph of concentration of sodium
thiosulphate against time (Graph I, Figure 1.18). 8 If the experiment is repeated with dilute
sulphuric acid of different concentrations, but
the concentration of sodium thiosulphate is kept
constant, the rate of reaction will also be directly
proportional to the concentration of the acid used.

Figure 1.18 Conclusion

(b) The graph of concentration of sodium 1 (a) From graph I, we can conclude that the
t hiosulphate against —ti—m1—e (Graph II, Figure 1.19). higher the concentration of sodium
6 Different volumes of distilled water are added thiosulphate, the shorter the time taken for a
to sodium thiosulphate solution so that the certain mass of sulphur to be precipitated, that
final volume of the diluted sodium thiosulphate is, for the cross ‘X’ to disappear from view.
solution is 50 cm3 in each experiment. Hence,
the concentration of sodium thiosulphate solution is (b) This means that the higher the
directly proportional to its volume before dilution. concentration of sodium thiosulphate, the
higher the rate of reaction.

2 From graph II, it can be concluded that the
concentration of sodium thiosulphate is directly
p roportional to —ti—m1—e .
Concentration of Na2S2O3 ∝ —ti—m1—e …(1)

3 B ut the rate of reaction is ∝ —ti—m1—e … (2)
Hence, combining equations (1) and (2), we have,
c oncentration of Na2S2O3 ∝ —ti—m1—e ∝ reaction rate
That is, rate of reaction ∝ concentration of
Na2S2O3 solution. The hypothesis is accepted.

313 Rate of Reaction

3 ’06

Which of the following reactants produces the highest Mg(s) + 2H+(aq) → Mg2+(aq) + H2(g)
rate of reaction with magnesium powder?
Ethanoic acid is a weak acid and produces very few
A 50 cm3 of 0.5 mol dm–­3 nitric acid H+ ions. HNO3, H2SO4 and HCl are strong acids. But
B 50 cm3 of 0.5 mol dm–­3 ethanoic acid each mole of H2SO4 produces two moles of H+ ions.
C 50 cm3 of 0.5 mol dm–­3 sulphuric acid Hence, 50 cm3 of 0.5 mol dm–3 sulphuric acid contains
D 50 cm3 of 0.5 mol dm–3 hydrochloric acid the highest concentration of H+ ions.
Answer C
Comments
Magnesium reacts with the hydrogen ions of the acids.

1.3

To study the effect of temperature on the rate of reaction between sodium SPM
thiosulphate solution and dilute sulphuric acid
’05/P1

Experiment 1.3 Problem statement 5 The stopwatch is started immediately and the
1 conical flask is swirled gently.
How does temperature affect the rate of reaction between
sodium thiosulphate solution and sulphuric acid? 6 The cross ‘X’ is viewed from above. The stopwatch
is stopped as soon as the cross disappears from view
Hypothesis The higher the temperature of the and the time taken is recorded.
reactant, the higher the rate of
reaction. 7 The solution in the conical flask is poured out.
The conical flask is washed thoroughly and dried.
Variables 50 cm3 of 0.1 mol dm–3 sodium thiosulphate
solution is poured into the conical flask.
(a) Manipulated variable: The temperature of
sodium thiosulphate solution 8 The solution is heated over a wire gauze until the
temperature reaches about 45 °C (Figure 1.21).
(b) Responding variable: The time taken for the
cross ‘X’ to disappear

(c) Constant variables: The concentrations and
volumes of both sodium thiosulphate solution
and dilute sulphuric acid

Apparatus Conical flask, 10 cm3 measuring
cylinder, thermometer, stopwatch,
white paper marked with a cross
‘X’, wire gauze, tripod stand and
Bunsen burner.

Materials 0.1 mol dm–3 sodium thiosulphate Figure 1.20
solution and 1.0 mol dm–3 sulphuric
acid.

Procedure Figure 1.21

1 50 cm3 of 0.1 mol dm–3 sodium thiosulphate
solution is poured into a clean, dry conical flask.

2 The temperature of the sodium thiosulphate
solution is measured with a thermometer.

3 The conical flask is placed on a white paper
marked with a cross ‘X’ (Figure 1.20).

4 5 cm3 of 1 mol dm–3 sulphuric acid is quickly
poured into the sodium thiosulphate solution.

Rate of Reaction 314

9 The hot conical flask is placed over a white paper

marked with a cross ‘X’.
10 5 cm3 of 1 mol dm–3 sulphuric acid is measured

out using a 10 cm3 measuring cylinder.
1 1 When the temperature of sodium thiosulphate

solution falls to 40°C, the sulphuric acid is
quickly poured into the thiosulphate solution.
12 The stopwatch is started immediately and the

conical flask is swirled gently.
1 3 The cross ‘X’ is viewed from the top and the

time taken for the cross to disappear from view is

recorded.
14 Steps 7 to13 are repeated at higher temperatures

as shown in the following table.

Results Figure 1.22 Graph of temperature against —ti—m1—e

Experiment 1 2 3 4 5 Discussion

Temperature 30 40 50 55 60 1 The graph shows that the temperature of sodium 1
(°C) thiosulphate solution is proportional (but not

Time (s) 52 27 16 13 10 linearly) to —ti—m1—e.
2 Temp erature ∝ —ti—m1—e … (1)
—T—i—m1—e (s–1) 0.019 0.037 0.063 0.077 0.100
But rate of reaction ∝ —ti—m1—e … (2)
Based on the results of the experiment, a graph of Combining equations (1) and (2), we have,
temperature of sodium thiosulphate solution against Rate of reaction ∝ temperature
—ti—m1—e is plotted (Figure 1.22).
Conclusion

The higher the temperature of the experiment, the
higher the rate of reaction.

4 ’05

The rate of reaction between sodium thiosulphate
solution and dilute sulphuric acid can be determined
by using the arrangement of apparatus as shown
below.

Which of the following conditions will cause the mark
‘X’ to take the shortest time to disappear from sight?

Sulphuric acid Sodium thiosulphate solution

Volume (cm3) Concentration Volume (cm3) Concentration Temperature (°C)
(mol dm–3) (mol dm–3)
30
A5 1.0 45 0.5 35
B5 1.0 45 0.5 30
C5 0.5 45 0.5 35
D 10 0.5 40 0.5

315 Rate of Reaction

Comments
The shorter the time taken for the mark ‘X’ to disappear, the faster the reaction.
The rate of reaction is affected by temperature and concentration.
The higher the temperature, the faster the reaction. (Answer B or D is correct).
The higher the concentration of sulphuric acid or sodium thiosulphate in the reaction mixture, the faster the
reaction (Answer D is incorrect).

Answer B

1.4

To study the effect of a catalyst on the rate of decomposition of hydrogen peroxide

Problem statement 4 0.5 g of manganese(IV) oxide, MnO2, is added to
hydrogen peroxide and shaken. The changes that
How do catalysts affect the rate of decomposition of take place in the test tube and on the glowing
hydrogen peroxide? splint are recorded.

Experiment 1.4 Hypothesis Results
1 Manganese(IV) oxide increases the rate of
decomposition of hydrogen peroxide. Experiment Observation

Variables Inside the test On the glowing
(a) Manipulated variable: The presence of tube splint

manganese(IV) oxide H2O2 without No effervescence The glowing
(b) Responding variable: The release of oxygen gas MnO2 splint does not
(c) Constant variables: Volume and concent­ration of
light up.
hydrogen peroxide
H2O2 with Bubbles of The glowing
Apparatus MnO2 oxygen gas are splint is
Test tube and wooden splint produced rekindled and
burns brightly.
Materials
Hydrogen peroxide and manganese(IV) oxide Discussion

Procedure 1 The following equation shows the decomposition
1 A test tube is half-filled with hydrogen peroxide. of hydrogen peroxide:
2 A glowing splint is placed at the mouth of the

test tube to test for the gas evolved (Figure 1.23).

2H2O2(aq) → 2H2O(l) + O2(g)

2 The glowing splint is rekindled in the presence
of oxygen gas.

Figure 1.23 The effect of a catalyst on the Conclusion
decomposition of hydrogen peroxide
The rate of evolution of oxygen gas increases when
3 The changes that take place inside the test tube manganese(IV) oxide is added to hydrogen peroxide.
and on the glowing splint are recorded. This proves that manganese(IV) oxide acts as a
catalyst and speeds up the decomposition of hydrogen
peroxide to water and oxygen. The hypothesis is
accepted.

Rate of Reaction 316

The reaction mixture remaining after Experiment 1.4 reaction in Experiment II. We can therefore
can be filtered to obtain the manganese(IV) oxide. It conclude that the higher the concentration
is found that (a) the mass of manganese(IV) oxide of hydrogen peroxide, the higher the rate
before and after the experiment is the same (0.5 g), of reaction.
(b) the chemical properties of manganese(IV) oxide (c) The maximum volume of oxygen gas
remain unchanged. produced in Experiment I is twice that
produced in Experiment II. This is because
The Effect of Concentration of Hydrogen SPM the number of moles of hydrogen peroxide
Peroxide on the Rate of Reaction used in Experiment I is twice that used in
’05/P1 Experiment II.

1 The graphs in Figure 1.24 show the effect of Explaining the Effectiveness of Different Catalysts
concentration of hydrogen peroxide on the on the Rate of Decomposition of Hydrogen Peroxide
rate of decomposition of hydrogen peroxide.
1 Figure 1.25 shows the results of an experiment
Graph I: more O2 produced and higher rate carried out to study the effect of different
of reaction because larger volume and catalysts (of the same mass) on the rate of
higher concentration of H2O2 is used. decomposition of hydrogen peroxide.

Maximum volumes of O2 collected are the 1
same for Experiments I and II because the
concentration and volume of H2O2 used
are the same.

Figure 1.24 The effect of concentration of hydrogen Figure 1.25 The effect of different catalysts
peroxide on the rate of decomposition
of hydrogen peroxide on the rate of reaction

In Experiment I, 50 cm3 of 0.14 mol dm–3 of In Experiment I, 50 cm3 of hydrogen peroxide
and 0.5 g of manganese(IV) oxide are used.
hydrogen peroxide and 0.2 g of manganese(IV) In Experiment II, 50 cm3 of hydrogen peroxide
and 0.5 g of iron(III) oxide are used.
oxide are used. In Experiment II, a solution For both the experiments, the concentration
and volume of hydrogen peroxide as well as
containing 25 cm3 of the same hydrogen the temperature are kept constant.
2 Analysis of the reaction rate curve in Figure 1.25.
peroxide mixed with 25 cm3 of water and 0.2 g (a) At any particular instant, the gradient of

of manganese(IV) oxide are used. graph I is greater than the gradient of graph
II. This means that the rate of reaction in
For both the experiments, the temperature is Experiment I is higher than the rate of reaction
in Experiment II. Thus, the experiment
kept constant. proves that manganese(IV) oxide is a more
effective catalyst than iron(III) oxide in the
2 (a) For Experiment I decomposition of hydrogen peroxide.
(b) The maximum volumes of oxygen gas
Concentration of H2O2 collected in both the experiments are the
= 0.14 mol dm–3 same because the volume and concentration
higher concentration of hydrogen peroxide used are the same. This
experiment shows that a catalyst does not
For experiment 2, hydrogen peroxide is change the yield of the products.
diluted.
(M1V1)before dilution = (M2V2)after dilution

Concentration of H2O2 after dilution
= —0—.1—4——m——o—l5—d0—m—c—m–3—3———2—5—c—m—3


= 0.07 mol dm–3 lower concentration

(b) At any particular instant, the gradient of

graph I is greater than the gradient of graph

II. This means that the rate of reaction

in Experiment I is higher than the rate of

317 Rate of Reaction

Effect of surface area Effect of temperature
Reaction of HCl(aq) with marble chips Reaction of HCl(aq) with zinc powder

Graph I : Large marble chips mass of Graph I : at 30°C mass of zinc (in excess), volume
Graph II : Small marble chips marble is
Graph III : Powdered marble the same Graph II : at 40°C and concentration of HCl are

Comment: The smaller the size of marble chips, the kept constant
higher the reaction rate.
Comment: When the temperature is raised, the rate
of reaction also increases.

Factors affecting the rate of reaction

1 Effect of concentration of reactant Effect of catalyst
Reaction of HCl(aq) with magnesium Decomposition of hydrogen peroxide

Graph I : 0.5 mol dm–3 HCl mass of Mg and Graph I : Fe2O3 used as catalyst
Graph II : 1.0 mol dm–3 HCl volume of HCl (in excess) Graph II : MnO2 used as catalyst
Graph III : 2.0 mol dm–3 HCl are kept constant Graph III : No catalyst
Comment: A catalyst increases the reaction rate. MnO2
Comment: When the concentration of hydrochloric acid
increases, the rate of reaction also increases. is a more effective catalyst than Fe2O3.

Effect of concentration and volume of acid used

Reaction of HCl(aq) with magnesium Reaction of HCI(aq) with magnesium

Graph I : Mg in excess Graph I : Mg in excess
20 cm3 of 0.2 mol dm–3 HCI 10 cm3 of 0.2 mol dm–3 HCI
Graph II : Mg in excess Graph II : Mg in excess
20 cm3 of 0.1 mol dm–3 HCl 30 cm3 of 0.1 mol dm–3 HCl

Comment: I II Comment: I II
Graph Graph

Reaction rate Higher Lower Reaction rate Higher Lower

Amount of 0.002 mol 0.001 mol Amount of 0.001 mol 0.0015 mol
H2 released (48 cm3) (24 cm3) H2 released (24 cm3) (36 cm3)

Rate of Reaction 318

Applications of Factors that Affect Rates is kept in a refrigerator will last longer because 1
of Reaction in Daily Life and in Industrial the decaying reaction that destroys the food
Processes can be slowed down.
4 In the supermarkets, fish, meat and other
Combustion of Charcoal types of fresh foods are kept in deep-freeze
compartments where the temperature is about
1 Combustion of charcoal in excess oxygen –20 °C. This keeps the food fresh for a few
produces carbon dioxide and water. Heat months because the very low temperature
energy is released during combustion. slows down the chemical reactions that cause
the food to decay.
2 Large pieces of charcoal will not burn easily
because the total surface area exposed to Cooking Food in Pressure Cookers
oxygen is small.
1 Pressure cookers are used to speed-up cooking.
3 If small pieces of charcoal are used, they can 2 In the pressure cooker, the higher pressure enables
burn easily. This is because the total surface area SPM water or oil to boil at a temperature higher than
exposed to the air increases. Thus, the rate of ’06/P1 their normal boiling points. Furthermore, an
reaction with oxygen (combustion) increases.
increase in pressure causes an increase in the
Coal is mainly carbon. Coal mining is dangerous number of water molecules or cooking oil
because coal dust present in the coal mine catches molecules coming into contact and colliding
fire very easily. Because of this, serious accidents in with the food particles.
coal mines can happen due to the explosion of coal 3 At a higher temperature and pressure, the rate
dust. Human lives are often lost in such explosions. of reaction becomes higher. Thus, food cooks
faster in pressure cookers.

Storing Food in Refrigerators Uses of Catalysts in Industry

1 The decomposition and decay of food is a 1 Catalysts do not increase the yields of reactions.
However, catalysts are used widely in industrial
SPM chemical reaction caused by the action of processes to increase the rates of reactions
’05/P2 so that the same amount of products can be
/SB microorganisms such as bacteria and fungi. obtained in a shorter time. As a result, the use of
catalysts brings down the cost of production.
These microorganisms multiply very rapidly
2 In the chemical industry, small pellets of solid
at the temperature range of 10–60 °C. catalysts are used instead of big lumps. This is
to give a larger surface for catalytic reaction to
2 Room temperature is the optimum temperature occur and hence a faster reaction will result.

for the breeding of microorganisms in food. 3 The table below summarises the raw materials
and the conditions needed for the Haber,
As a result, food turns bad quickly at room Contact and Ostwald processes.

temperature.

3 At low temperatures, for example, 5 °C (the

normal temperature of a refrigerator), the activities

of bacteria are slowed down. Hence, food that

Industrial process Substances Optimum conditions/equation reaction
Nitrogen and hydrogen
Manufacture of ammonia Temperature: 450–500 °C
(Haber process) Pressure: 250 atmospheres
Catalyst: Finely divided iron (Fe)

N2(g) + 3H2(g) 450 °C, 250 atm 2NH3(g)
Fe (catalyst)

Manufacture of sulphuric Sulphur (to make SO2), Temperature: 400–450 °C
acid (Contact process) air and water Pressure: 1–2 atmospheres
Catalyst: Vanadium(V) oxide, V2O5

319 Rate of Reaction

Industrial process Substances Optimum conditions/equation reaction

Manufacture of nitric • The following reaction scheme shows the steps
acid (Ostwald process) involved in the manufacture of sulphuric acid:

S o⎯xsitde⎯apti1→o n SO2 oxidation SO3 ⎯step→3 H2S2O7 s⎯tep→4 H2SO4

⎯ste⎯p 2→

• In step 2, sulphur dioxide is oxidised to sulphur trioxide.

450 °C, 1 atm 2SO3(g)
2SO2(g) + O2(g) V2O5 (catalyst)

Ammonia, air and water Temperature: 900 °C
Pressure: 5 atmospheres
1 Catalyst: platinum
• The following reaction scheme shows the steps

involved in the manufacture of nitric acid.

NH3 ⎯oxs⎯itdeap⎯ti1o→ n NO ⎯oxs⎯itdeap⎯ti2o→ n NO2 o⎯xsit⎯deapt⎯i3on→ HNO3

In step 1, ammonia is oxidised to nitric oxide.

900 °C, 5 atm
4NH3(g) + 5O2(g) Pt (catalyst)

4NO(g) + 6H2O(g)

Solving Problems on Rate of Reaction Mass and Concentration Temperature
nature of Zn of H2SO4
6
C 5.0 g of
Curve I in Figure 1.26 is obtained by treating 5.0 g granulated Zn
of granulated zinc with 2.0 mol dm–3 sulphuric acid
(in excess) at 30 °C.

2 mol dm–3 40 °C

D 5.0 g of 1 mol dm–3 30 °C
granulated Zn
Figure 1.26

Which of the following conditions will produce Comments
graph II? Curves I and II show that:
(a) The total volume of hydrogen produced in Experiment
Mass and Concentration Temperature
nature of Zn of H2SO4 II is the same as that produced in Experiment I.
This means that the amount of zinc used is 5 g
A 2.5 g of 2 mol dm–3 30 °C and not 2.5 g. Answer A is incorrect.
granulated Zn (b) Reaction II is slower than reaction I.
This means that zinc powder or a higher
B 5.0 g of Zn 2 mol dm–3 30 °C temperature of 40 °C is not used in Experiment II.
powder Answers B and C are incorrect.
The low rate is achieved by using sulphuric acid
more dilute than 2 mol dm–3 (1 mol dm–3).

Answer D

Rate of Reaction 320

7 (b) Differences in terms of rate of reaction
Graph II is steeper than graph I because the rate of
Two experiments were carried out to determine
the rate of oxygen gas production during the reaction in Experiment II is expected to be higher
decomposition of hydrogen peroxide. In Experiment than Experiment I. When the concentration of
I, 20 cm3 of 2 mol dm–3 hydrogen peroxide were hydrogen peroxide is increased from 2 mol dm–3
used and the results of the experiment are shown on to 4 mol dm–3, the rate of reaction also increases.
graph I in Figure 1.27.
Difference in terms of volume of oxygen released
Figure 1.27
(a) Sketch a graph on the same axes to show the Step 1 To calculate the volume of oxygen
produced in Experiment I
results of the experiment that will be obtained
if 5 cm3 of 4 mol dm–3 hydrogen peroxide were 2H2O2(aq) → 2H2O(l) + O2(g)
used for the reaction.
(b) Explain your answer in (a). Number of moles of H2O2 used in Experiment I
(c) State the constant variables for both the = —2—1—0—0—20—0 = 0.04
experiments.
Solution ∴ Volume of oxygen collected at room
(a) temperature in Experiment I

= —21  0.04  24 000 = 480 cm3 (V cm3)

Step 2 To calculate the volume of oxygen 1
produced in Experiment II

Number of moles of H2O2 used in Experiment II
= —4—1—0—0—05 = 0.02

∴ Volume of oxygen collected at room
temperature in Experiment II

= —21  0.02  24 000 = 240 cm3 (—21 V cm3)

(c) Constant variables:
In both the experiments, the same mass of the
catalyst and the same temperature of reaction

are used.

1.2 Figure 1.28

1 State three ways that can be used to increase the rate of Rate of Reaction
reaction of zinc powder with dilute sulphuric acid.

2 Excess calcium carbonate is added to hydrochloric
acid at room temperature. The volume of carbon
dioxide collected is recorded at regular time intervals.
The results of the experiment are shown in Figure
1.28.
(a) At what time does the reaction stop?
(b) Why does the reaction stop at this particular
time?
(c) The experiment is repeated by using the same
hydrochloric acid but at a lower temperature
than room temperature. On the same axes,
plot a graph to show the results of the second
experiment.
(d) State the constant variables for both the
experiments.

321

3 Hydrogen peroxide decomposes as represented by (c) Give one inference that can be made from the
the equation: results in Experiment I.

2H2O2(aq) → 2H2O(l) + O2(g) (d) Explain why the initial readings on the electronic
balance are different for the three experiments.
(a) On the same axes, sketch two graphs of total
amount (in mol) of oxygen gas given off against 5 Four experiments are carried out to study the rate of
time to show the results of Experiments I and II reaction between zinc (in excess) and sulphuric acid
under the conditions stated below. at different conditions.
In each experiment 80 cm3 of 0.1 mol dm–3 sulphuric
Experiment I: acid is used. The time taken to collect 192 cm3 of
100 cm3 of 1.0 mol dm–3 H2O2 hydrogen gas produced are shown below.
Experiment II:
300 cm3 of 0.2 mol dm–3 H2O2 Experiment Substances Temperature Time
(°C) (s)
(b) Explain your answer based on the graphs that I Zinc + H2SO4 35 40
you have sketched. II Zinc + H2SO4 38 18
III Zinc + H2SO4 + 35 12
4 In an experiment carried out at room temperature 1 cm3 of CuSO4
(28 °C), 8.0 g of marble chips are added to 100 cm3 IV Zinc + H2SO4 + 35 40
of dilute hydrochloric acid in a conical flask. The mass 1 cm3 of Na2SO4
of the conical flask and its contents is determined
1 using an electronic balance at the beginning of the
experiment (that is, as soon as the marble chips are
added) and then after 1 minute. (a) Sketch the graphs of total volume of hydrogen
released against time for Experiments I, II and III
The experiment is repeated at 35 °C and 40 °C. The on the same axes.
experimental results are shown below.
(b) Explain why the reaction rate for (i) Experiment
Electronic Electronic I is different from that of Experiment II, (ii)
Experiment II is different from that of Experiment
Experiment Temperature balance balance III.
(°C) reading at reading
after 1 (c) What conclusion can you make by comparing
the Experiments III and IV?

beginning (g) minute (g) (d) The reaction mixture in Experiment III is filtered.
Excess sodium hydroxide is added to the filtrate.
I 28 270.35 270.04 (i) Predict what you would observe.
(ii) Write an ionic equation for the reaction.
II 35 271.42 270.01
(e) Based on your answer in (d), what inference
III 40 268.20 266.00 can you make with regards to the property of a
catalyst?
(a) State a hypothesis for the experiment.
(b) State the constant variables for the experiment.

1.3 The Collision Theory SPM particles (atoms, molecules or ions) collide
’05/P1 with each other. However, not all collisions
1 According to the kinetic theory of matter, all
matter is made up of tiny, discrete particles. will result in a chemical reaction to form the
These particles are continually moving and so products of the reaction. It is likely that the
have kinetic energy. particles collide and bounce back without
producing any changes. The collisions that are
2 Based on the assumption that the particles in successful in producing a chemical reaction
matter are moving all the time and collide with are called effective collisions.
each other, the collision theory was introduced 4 Collisions of particles that are unsuccessful
to explain in producing a chemical reaction are called
(a) how chemical reactions occur, and ineffective collisions.
(b) the factors (such as particle size, 5 The collision theory states that for a chemical
concentration, temperature, catalyst and reaction to occur, the reacting particles must
pressure) affecting the rates of reactions. (a) collide with each other so that the breaking

3 When the reactants are mixed, the reactant and formation of chemical bonds can occur.

Rate of Reaction 322

(b) possess energy that is equal to, or more than 5 (a) If two molecules with sufficient energy
the minimum energy called the activation (that is, energy equal to or more than
energy. the activation energy) collide in the
correct orientation, the chemical bonds
(c) collide in the correct orientation. in the reactant molecules will break and
reaction will occur to form new bonds in
Activation Energy the product molecules (Figure 1.31). For
example,
1 Activation energy is the minimum energy that
the reactant particles must possess at the time H2(g) + I2(g) 2HI(g)
of collision in order for a chemical reaction to
take place. Figure 1.31 Effective collision (sufficient energy 1
and correct orientation)
2 The activation energy can also be considered as
an energy barrier that must be overcome by the (b) However, if two molecules, with energy
colliding particles in order that collision will equal to or more than the activation energy,
result in the formation of product molecules. but collide with each other in an incorrect
orientation, then reaction will not occur.
3 Figure 1.29 shows the energy profile diagram
for the exothermic reaction: A + B → C + D. (c) If two reactant molecules, with energy
An exothermic reaction is the reaction that less than the activation energy, collide in
releases heat energy (Refer Section 4.1). In the the correct orientation, then reaction will
energy profile diagram, the y-axis represents the also not occur. The colliding molecules
energy content of the reactants and products, will simply rebound and move away from
while the x-axis represents the progress of the each other.
reaction (reaction coordinate).
Relating the Frequency of Effective
SPM Collisions with Factors Influencing the
Rate of Reaction
’07/P2
1 Based on the collision theory, two important
The energy of the products factors that determine the rate of a chemical
is lower than the energy reaction are
of the reactants. Therefore (a) the frequency of effective collisions and
heat is released during (b) the magnitude of the activation energy.
the reaction.
2 Frequency of effective collisions
Figure 1.29 Energy profile diagram for an exothermic For a given reaction, if the frequency of collisions
reaction between the reactant molecules is high, it follows
that the frequency of effective collisions that
4 Figure 1.30 shows the energy profile diagram causes a reaction to occur will also be high. As a
for the endothermic reaction: E + F → G + H. result, the rate of reaction increases.
An endothermic reaction is a reaction that
absorbs heat energy. 3 Magnitude of activation energy
Reactions that have high activation energy will
The energy of the products occur at a slow rate. This is because only a small
is higher than the energy fraction of the molecules possess sufficient
of the reactants. Therefore energy to overcome the activation energy for
heat is absorbed during the reaction to occur. In contrast, reactions that
the reaction.

Figure 1.30 The energy profile diagram for an
endothermic reaction

323 Rate of Reaction

1 possess low activation energy will occur at a Effect of Concentration on the Rate of Reaction
high rate. This is because most of the molecules
have sufficient energy to overcome the activation 1 Magnesium reacts with dilute hydrochloric
energy and this enables the reaction to occur. acid as represented by the equation
4 In general, any factor that increases the rate
of effective collisions will also increase the Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)
rate of reaction.
When the concentration of hydrochloric acid
Effect of Size of Reactant (Surface Area) on the increases, the rate of reaction also increases.
Reaction Rate 2 Figure 1.33 shows the arrangement of particles
in 1 mol dm–3 hydrochloric acid and 2 mol dm–3
1 The sodium chloride crystal as shown in Figure (more concentrated) hydrochloric acid.
1.32(a) has a total surface area of 16 cm2.
When this crystal is divided into smaller Figure 1.33 Arrangement of particles in dilute and
crystals as shown in Figure 1.32(b), the total concentrated solutions
surface area is increased to 24 cm2.
When the concentration of hydrochloric acid
Total surface area of the NaCl crystal in increases, the number of particles per unit
Figure 1.32(a) volume also increases and the particles are
= (1  2)  4 + (2  2)  2 closer together.
= 16 cm2 3 When the number of particles increases, the
frequency of collisions also increases. As a result,
Total surface area in Figure 1.32(b) the frequency of effective collisions increases.
= (1  1)  6  4 This causes the rate of reaction to increase.
= 24 cm2

Figure 1.32 5 ’05

The smaller the particle size, the greater the The decomposition of hydrogen peroxide produces
total surface area exposed for reaction to occur. oxygen gas. Curve P is obtained when 25 cm3 of 0.1
2 Dilute sulphuric acid reacts with magnesium mol dm–3 hydrogen peroxide undergoes decomposition.
as represented by the equation
If the experiment is repeated using another hydrogen
Mg(s) + H2SO4(aq) → MgSO4(aq) + H2(g) peroxide solution, which solution will produce
curve Q?
If small pieces of magnesium or magnesium
powder are used, the rate of reaction between A 10 cm3 of 0.25 mol dm–3 hydrogen peroxide
magnesium and sulphuric acid will increase. B 15 cm3 of 0.15 mol dm–3 hydrogen peroxide
3 The smaller the size of the solid, the larger the C 20 cm3 of 0.20 mol dm–3 hydrogen peroxide
total surface area exposed for collisions. This D 25 cm3 of 0.15 mol dm–3 hydrogen peroxide
means that the frequency of effective collisions Comments
(that is, collisions with the correct orientation • Steepness of curves P and Q
and with energy equal to or greater than the Curve Q is more steep than curve P. This means
activation energy) between reacting particles
will increase. As a result, the rate of reaction
also increases.

Rate of Reaction 324

that the rate of reaction is higher. That is, the 3 As a result, increasing the pressure causes the
hydrogen peroxide used for curve Q is more gaseous molecules to collide more frequently.
concentrated. Consequently, the frequency of effective
• Maximum volume of oxygen gas produced in collisions increases and the rate of reaction
curve Q is less than that in curve P. This means also increases.
that the number of moles of hydrogen peroxide
used in curve Q is less than that in curve P. An increase in pressure will not increase the speed
of the reacting particles. In actual fact, the increase in
Number of moles of hydrogen peroxide rate at high pressures is caused by the particles being
squeezed closer together. This increases the frequency
= —c—o—n—c—e—n—tr—a—ti—o—n—(—m——o—l1—d0—m0—0–—3)————v—o—lu—m——e——(c—m——3) of effective collisions and hence the rate increases.

Solution Relative Number of Effect of Temperature on the Rate of Reaction
concentration moles of H2O2
For curve P 1 Calcium carbonate reacts with hydrochloric
A 0.1 mol dm–3 2.5  10–3 acid to form carbon dioxide as represented by
B More concentrated 2.5  10–3 the following equation
C More concentrated 2.25  10–3
D More concentrated 4.0  10–3 CaCO3(s) + 2HCl(aq) →
More concentrated 3.75  10–3 CaCl2(aq) + H2O(l) + CO2(g)

Answer B Figure 1.35 shows the graphs of total volume 1
of carbon dioxide given off against the time taken
Effect of Pressure on the Rate of Reaction for the reaction between calcium carbonate and
dilute hydrochloric acid at 25 °C and 30 °C.
1 In chemical reactions involving gases, increasing
the pressure increases the rate of reaction. SPM
Conversely, decreasing the pressure decreases
the rate of reaction. For example, the rate of ’05/P2
reaction between nitrogen and oxygen to ’06/P1
produce nitrogen monoxide can be increased
by increasing the pressure. Figure 1.35 Effect of temperature on
the rate of reaction
N2(g) + O2(g) → 2NO(g)
2 At low temperatures, particles of reactants
2 At low pressures, the gaseous molecules are move at a slower speed. However, when the
spread far apart (Figure 1.34(a)). When the temperature is increased, the particles absorb
pressure is increased, the volume of the gas is the heat energy. As a result, the kinetic energy
reduced (Figure 1.34(b)). of the particles increases. Hence,
(a) the reacting particles move faster, and
Figure 1.34 Effect of pressure on gaseous molecules (b) the number of reacting particles with the
activation energy required for the reaction
This means that at high pressures, increases.
(a) the number of gaseous molecules per unit
3 Consequently, the frequency of effective
volume is increased, and collisions increases and hence, the rate of
(b) the gaseous molecules are packed closer reaction also increases.

together. 4 Temperature has a great effect on the rate of
reaction. For most reactions, the rate of reaction
approximately doubles when the temperature
of reaction increases by 10 °C.

325 Rate of Reaction

Effect of Catalysts on Reaction Rates 4 A catalyst provides an alternative reaction
route (or pathway) for the reaction to occur.
1 The decomposition of hydrogen peroxide to In the presence of a positive catalyst, this
water and oxygen occurs very slowly at room alternative route has a lower activation energy.
temperature. In other words, a positive catalyst lowers the
activation energy required for the reaction
2H2O2(aq) → 2H2O(l) + O2(g) (Figure 1.37). As a result, more reacting particles
possess sufficient energy to overcome the lower
In the presence of a catalyst, the decomposition activation energy required for effective collisions.
of hydrogen peroxide occurs rapidly. Hence, the frequency of effective collisions
2 Figure 1.36 shows the rate of evolution of oxygen increases and the rate of reaction increases.
for the decomposition of hydrogen peroxide
without a catalyst and in the presence of a catalyst Ea = Activation energy
such as manganese(IV) oxide or iron(III) oxide. without catalyst

Ea’ = Activation energy
with catalyst

Ea’ < Ea

1 Figure 1.37 Effect of catalyst on the activation energy
of a reaction
Figure 1.36 The effect of a catalyst on the
decomposition of hydrogen peroxide Enzymes are biological catalysts. Enzymes are protein
molecules produced in living cells. The enzyme,
3 A chemical reaction occurs when reactant catalase, is found in the liver. This enzyme can catalyse
particles collide with one another. In the presence the decomposition of hydrogen peroxide (a toxic
of a catalyst, the reactant particles can collide substance produced by metabolism in human bodies)
with the catalyst and also with each other. This to harmless substances, that is, water and oxygen.
causes the reactants to react in a different way. Enzymes are also used in detergents to remove protein
Thus, the activation energy of the reaction can stains (for example, food or bloodstains) on clothings.
be increased or decreased depending on the
type of catalyst used.

Effective collisions: Collision theory: a reaction only Activation energy: the
Collisions that produce a occurs if the particles (a) have minimum energy the
reaction. sufficient energy to overcome the reac­ting substances must
Reaction rate increases when activation energy and (b) collide possess before reaction can
the effective collisions increase. in the correct orientation. occur.

is used to explain the following factors

particle size concentration/pressure temperature catalyst

When the particle size is When the concentration/ When the temperature is A catalyst will lower the
decreased, the total pressure is increased, the increased, the number of activation energy required
surface area exposed for particles with the activation for the reaction by providing
reaction increases. number of particles per energy required increases. an alternative route with a
lower activation energy.
unit volume increases.

Frequency of effective collisions increases Rate of reaction increases

Rate of Reaction 326

6 ’03

Two experiments were carried out to study the rate Comments
of reaction between zinc and excess sulphuric acid at In these two experiments, the constant variables are
room temperature. The table below shows the results • volume, concentration and temperature of sulphuric
of the experiments.
acid used,
Experiment I II • mass and surface area of zinc used.

Experimental Solution
set-up
The time taken for Experiment II is shorter. This
Time taken for 30 12 implies that the reaction for Experiment II is faster.
all the zinc to 35 °C 35 °C Thus, copper(II) sulphate acts as the catalyst.
dissolve (s) • If a catalyst is added, the rate of reaction increases

Temperature because the catalyst provides an alternative route
with a lower activation energy for the reaction to
occur. 1
• Hence, the minimum energy required for the
By using collision theory, explain why the time taken reaction is less. As a result, more reacting particles
for Experiment II is different from that of Experiment I. possess sufficient energy to overcome the lower
activation energy required for effective collisions.
Hence, the frequency of effective collisions
increases and the rate of reaction increases.

1.3 3 Four experiments were carried out to study the rate of
reaction between nitric acid and calcium carbonate of
1 Which of the following changes will increase the rate different sizes. In Experiment I, V cm3 of 1.0 mol dm–3
of reaction between sodium thiosulphate solution and nitric acid is added to big lumps of excess calcium
sulphuric acid? carbonate in a conical flask. The total volume of
carbon dioxide produced is plotted against time
I By using sulphuric acid of a higher concentration taken (Figure 1.38).
II By increasing the temperature of sodium thiosulphate
Figure 1.38
solution used (a) (i) What is the difference between the rate of
III By increasing the volume of sodium thiosulphate
reaction at the first minute and the rate of
used reaction at the second minute?
IV By adding a small amount of sodium hydroxide (ii) Explain this difference in terms of collision
theory.
solution to the reaction mixture (b) The experiment was repeated three times by
Explain your answer in terms of the collision theory. changing the reaction conditions for each experiment
as shown below.
2 An experiment is carried out to study the decomposition
of hydrogen peroxide. In this experiment, 2.0 g of
manganese(IV) oxide is added to 30 cm3 of 0.2
mol dm–3 hydrogen peroxide. The volume of oxygen
produced is recorded at 30-second intervals.
(a) Calculate the maximum volume of oxygen gas
produced in the experiment at room conditions.
[1 mol of gas occupies 24 dm3 at room
conditions]

(b) (i) Sketch a graph of volume of oxygen produced
against time.

(ii) Explain the shape of the graph.
(c) (i) What is the function of manganese(IV) oxide

in this experiment?
(ii) Based on collision theory, explain the effect of

manganese(IV) oxide on the decomposition
of hydrogen peroxide.

327 Rate of Reaction

Experiment Change in conditions of reaction (i) Sketch the graphs for Experiments II, III and IV
II on Figure 1.38 and label each of these graphs.
III Nitric acid at a lower temperature
( ii) Explain the difference between the reaction rates
IV Small lumps of calcium carbonate but for Experiments I and II in terms of collision
of the same mass theory.

2.0 mol dm–3 nitric acid but of the 4 With the aid of an energy profile diagram, explain
same volume (V cm3) how a negative catalyst (inhibitor) affects the rate of
reaction.

1 1.4 Practising Scientific 4 Nowadays, enzymes are used extensively in
Knowledge to Enhance industry to enable reactions to proceed rapidly
Quality of Life at room temperature and pressure.

1 In our homes, we require machines to increase 5 In our daily living, we face many social and
the rates of reactions and machines to reduce the environmental issues that threaten the quality
rates of reactions. Examples of such machines are of living. For example, air and water pollution,
microwave ovens and refrigerators respectively. food shortages, diseases and so on.

2 In the hospitals, oxygen tents are used to save 6 We must use science and technology to
lives. The high concentration of oxygen helps overcome these problems in a rational and
patients with difficulty in breathing to breathe systematic way so that the quality of life can
normally. be improved.

3 In human bodies, enzymes (biological catalysts) 7 We should be thankful for the contribution of
are needed to catalyse complex biochemical scientists in enhancing the quality of life in
reactions. modern living.

1 The rate of reaction is defined as the amount of • Particle size (surface area) of solid reactant

a reactant used up or the amount of a product • Concentration of reactants

obtained per unit time. • Temperature of reactants
2 The rate of a reaction is inversely proportional to the
• Presence of catalyst
time taken for the reaction.
3 Different chemical reactions take place at different • Pressure (for reactions involving gases)
7 Transition metals (Fe, Ni and Pt) and compounds
rates. A fast reaction takes a shorter time to complete
of transition metals (MnO2 and V2O5) are often used
than a slow reaction. as catalysts.
4 The rate of reaction can be determined in the school 8 According to the collision theory, a reaction will occur

laboratory by measuring the if the reacting particles
• collide with each other
• changes in the mass of the reactant, • possess activation energy
• collide in the correct orientation
• changes in volume of gas produced, 9 The collisions that are successful in producing a
chemical reaction are called effective collisions.
• time taken for formation of precipitate. 10 Any factor that increases the rate of effective
5 The rate of reaction can be expressed in terms of
collisions will also increase the rate of reaction.
(a) the average rate of reaction over a period of
time, or (b) the instantaneous rate of reaction at a

given time.
6 The rate of a reaction is affected by the following

factors:

Rate of Reaction 328

1

Multiple-choice Questions

1.1 Rate of Reaction 4 The diagram below shows
the apparatus set-up used to
1 Calcium carbonate reacts with dilute hydrochloric acid to produce carbon determine the rate of reaction.
dioxide. The total volume of carbon dioxide collected during the reaction is

’11 shown below.

Time (s) 0 5 10 15 20 25 30 35 40

Volume of 0.00 12.00 20.00 26.50 31.00 32.50 33.00 33.00 33.00
CO2 (cm3)

What is the overall average rate of reaction? The apparatus set-up is not
suitable to be used for determining
A 0.825 cm3 s–1 C 1.100 cm3 s–1 the rate of reaction for
A Na2SO3(s) + H2SO4(aq) →
B 0.943 cm3 s–1 D 1.300 cm3 s–1 1
Na2SO4(aq) + H2O(l) + SO2(g)
2 The average rate of reaction of calcium carbonate with hydrochloric acid is B Mg(s) + 2HCl(aq) →

0.0080 MgCl2(aq) + H2(g)
C NaHCO3(s) + HNO3(aq) →
mol s–1. What is the time taken for 9.60 g of calcium carbonate to
NaNO3(aq) + H2O(l) + CO2(g)
completely react with excess hydrochloric acid?
D 2H2O2(aq) ⎯M⎯nO⎯2(s)→
[Relative atomic mass: C, 12; 2H2O(l) + O2(g)

O, 16; Ca, 40] 5 The graph shows the total
volume of carbon dioxide
A 12 s B 24 s C 120 s D 240 s evolved when 10.0 g of calcium
carbonate (in excess) reacts
3 Calcium carbonate reacts with dilute hydrochloric acid to produce carbon with 20.0 cm3 of 1.0 mol dm–3
dioxide gas. dilute hydrochloric acid.
The plot of volume of CO2 produced against time is shown as follows.
volume of gas(cm3)

Y Z
X TC 55

0 t1 t2 t3 time(s)
Which of the following

statements is correct?

A The reaction is faster at point

Which of the following can be deduced from the graph? Y than at point X.

I The average rate of reaction is 1.5 cm3 s–1. B The reaction is fastest at

II The rate of reaction decreases with time. point Z.

III The rate of reaction at 35 seconds is zero. C The reaction reaches

IV The gradient of the curve decreases because the concentration of acid D completion at time t3c.arbon
The total volume of
decreases.
dioxide evolved is the
A I and II only C I, II and III only
same if 12.0 g of calcium
B III and IV only D II, III and IV only
carbonate is used.

329 Rate of Reaction

6 Two experiments on the The rate of reaction is best determined by
decomposition of hydrogen
peroxide were carried out. The A measuring the cvoonlucmenetraotfioSnOo2f produced at regular time intervals.
graphs in the following diagram B measuring the hydrochloric acid at regular time intervals.
show the total volume of
oxygen collected against time for C recording the time as soon as the ‘cross’ mark disappears.
each of the experiments.
D recording the time as soon as precipitate appears.
Which of the following graphs
shows how the rates of reaction 1.2 Factors that Affect the Rate of Reaction
vary with time for the experiments?
A 8 Ammonia is produced using the Haber process. The table shows the mass

B of ammonia produced at four conditions of temperature and pressure.

C ’04 I II III IV

D Condition

Mass of

ammonia 300 250 150 200

1 produced (kg)

Time taken 4 hours 2 ½ hours 8 minutes 12 minutes

At which condition is the rate of production of ammonia the highest?

A Condition I C Condition III

B Condition II D Condition IV

9 Zinc powder reacts with hydrochloric acid according to the equation:

Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
In order to have the highest initial rate, which of the following solutions
should be used for the reaction with zinc powder?
A 30 g of HCl in 1000 cm3 of water C 15 g of HCl in 100 cm3 of water
B 20 g of HCl in 1000 cm3 of water D 4.0 g of HCl in 50 cm3 of water

10 Two experiments were carried out at 25 °C to study the rate of reaction
between magnesium carbonate powder (in excess) and an acid. The
volume of carbon dioxide liberated was measured at regular intervals.

Experiment Acid used

I 100 cm3 of 0.5 mol dm–3 hydrochloric acid

II 100 cm3 of 0.5 mol dm–3 sulphuric acid

Which of the following graphs represents the results obtained in

Experiments I and II?

A C

7 The diagram shows the apparatus B
set-up for an experiment to 330 D

’07 determine the rate of reaction
between sodium thiosulphate
and hydrochloric acid

S2O32–(aq) + 2H+(aq) →
H2O(l) + SO2(g) + S(s)

Rate of Reaction

11 2.5 g of zinc were allowed to react with 100 cm3 of hydrochloric acid III temperature of hydrogen
under different conditions as shown below. peroxide decreases with time.
Under which conditions will hydrogen gas be given off at the highest rate?
IV mass of manganese(IV)
Temperature (°C) Concentration (mol dm–3) Form of zinc oxide decreases.
Small pieces
A 30 0.5 A I only
B 25 1.0 Powder B I and II only
C 35 1.0 Small pieces C III and IV only
D 35 1.0 D I, III and IV only
Powder
15 The following equation shows
12 For the following reaction: the reaction between powdered

Zn +SuHrf2aScOe4a→reaZonSf Ozi4nc+ H2 , which factor does not affect the rate of reaction? ’03 calcium carbonate and dilute
A C Volume of sulphuric acid hydrochloric acid
’08
CaCO3(s) + 2HCl(aq) →
B Concentration of sulphuric acid D Temperature of sulphuric acid CaCl2(aq) + H2O(l) + CO2(g)

13 Graph X in the diagram below shows the result of the decomposition of The production of carbon dioxide
10 cm3 of 0.4 mol dm–3 hydrogen peroxide. The experiment was carried can be slowed down by
out at 30 °C. I reducing the temperature of

hydrochloric acid used. 1
II adding distilled water to
Which of the following conditions produces graph Y if 0.1 g of
manganese(IV) oxide was used as the catalyst for both experiments? hydrochloric acid before the
reaction.
Volume of H2O2 Concentration of H2O2 Temperature (°C) III using larger pieces of calcium
(cm3) (mol dm–3) carbonate.
30 IV reducing the pressure on the
A 10 0.25 30 reaction mixture.
B 12.5 0.40 30 A I, II and III only
C 20 0.25 40 B I, III and IV only
D 20 0.40 C I, II and IV only
D I, II, III and IV
14 The graph in the diagram below shows the changes in the concentration
of hydrogen peroxide, H2O2, when powdered manganese(IV) oxide is 16 A piece of magnesium ribbon is
allowed to react with 100 cm3 of
’04 added to it. 1.0 mol dm–3 hydrochloric acid.
Which of the following changes
The gradients of the hgyradprohgaetntimpeersoxt1idaenddet2carereasdeifsf.erent because the will increase the rate of reaction?
I concentration of
I Increasing the temperature of
hydrochloric acid.
II volume of hydrogen peroxide decreases.
II Replacing the magnesium
ribbon with magnesium
powder.

III Replacing the acid with
50 cm3 of 2.0 mol dm–3
hydrochloric acid.

IV Adding 50 cm3 of 1.0 mol
dm–3 hydrochloric acid.

A I and II only
B III and IV only
C II and IV only
D I, II and III only

17 What is the most suitable
method for cooking 100 g of

’06 potatoes within a short time?
A Steam the potatoes in a
steamer
B Fry the potatoes in a copper
pot

331 Rate of Reaction

C Boil the potatoes in a saucepan 21 Which of the following pairs of catalyst and processes are correctly
D Boil the potatoes in a matched?

pressure cooker Catalyst Process

18 Iron(III) oxide is a brown solid and I Iron Manufacture of ammonia in the Haber process
iron(III) salts are brown in colour. II
When iron(III) oxide is added to III Nickel Manufacture of nitric acid in the Ostwald process
hydrogen peroxide solution in a
test tube, a fast reaction occurs IV Vanadium(V) Manufacture of sulphuric acid in the Contact
and oxygen gas is liberated. oxide process
What is left in the test tube at
the end of the reaction? Lead(IV) oxide Production of oxygen by the decomposition of
A A brown solution only. hydrogen peroxide
B A brown solid and a brown
solution. A I and II only C I, II and IV only
C A brown solid and a B II and III only D I, III and IV only
colourless solution.
D A white solid and a 22 Which of the following reactions C
colourless solution . D
require a catalyst to speed up
19 The diagram below shows
the total volume of carbon the reaction?
dioxide given off when dilute
1 hydrochloric acid reacts with IIIIII 22SOSHO32O2+2+H→O2S22O→H4 2→O2S+HO2O3S22O7
calcium carbonate powder. IV Na2S2O3 + H2SO4 →
Na2SO4 + H2O + SO2 + S
I and II only
A

B I and III only

C I, III, and IV only

D II, III and IV only 25 Magnesium ribbons of the same
length are added separately to
23 Which of the following reactions each of the following solutions
of hydrochloric acid.
require a catalyst to speed up In which solution will the
magnesium ribbon disappear first?
the reaction? [The hydrochloric acid used is in
excess]
As the reaction proceeds, the IIVIIIII 2SNNOHa2 232+SO+22O3HH→3 22+S→2O2H4H22→OCNl H+H→32OS22O7
2NaCl + H2O + SO2 + S
gradient of the graph becomes I and II only Volume Concentration Temperature
A
less steep because of HCI of HCI of HCI
B I and III only
I the mass of calcium (cm3) (mol dm–3) (°C)
C I, III and IV only
carbonate decreases.
D II, III and IV only A 300 1.0 30
II the total surface area of B 200 1.0 25
24 The reaction between sulphuric C 100 2.0 30
calcium carbonate decreases. acid and magnesium carbonate D 200 2.0 25
is carried out at different
III the volume of hydrochloric conditions. 26 The energy profile diagram
Which reaction is fastest? for an uncatalysed reaction is
acid decreases. A shown below.

IV the temperature of the

mixture increases.
A I and II only
B I, II and III only
C II, III and IV only
D I, III and IV only

20 Which of the following reactions

takes place in the Ostwald

process?

BAC 2N22SNS2(OOHg)323(((+ggg)))3++H2OO(g22)((gg)) 2NH3(g) B
D 42NNOH3((gg))++35HO2(2(gg))

4NO(g) + 6H2O(l)

Rate of Reaction 332

The reaction was repeated using C 1.3 The Collision Theory
a catalyst. D
What is the effect of the catalyst 30 Based on the collision theory,
on the heat of reaction (∆H) what are the effects of a rise
and activation energy (Ea) for
the reaction? ’06 in temperature on the reactant
particles?
∆H Ea
A No change Decrease I The kinetic energy of the
B Decrease No change reactant particles increases.
C Decrease Decrease
D Increase Increase II The number of reactant
particles per unit volume
29 Three experiments are carried increases.

27 Carbon dioxide is produced out to study the rate of III The frequency of collisions
when magnesium carbonate between reactant particles
reacts with dilute hydrochloric decomposition of hydrogen increases.
acid.
peroxide by the catalyst, IV The activation energy of the
MgCO3(s) + 2HCl(aq) → reactant particles increases.
MgCl2(aq) + H2O(l) + CO2(g) manganese(IV) oxide. In all
A I and II only
Which of the following changes these three experiments, the B I and III only
will increase the initial rate of C II and IV only
carbon dioxide production? mass of the catalyst used is the D III and IV only
A Heat the reaction mixture
B Increase the size of solid same. The experimental results 31 When zinc powder is added to 1
dilute sulphuric acid, gas bubbles
magnesium carbonate are shown in the following
C Increase the volume of ’11 are produced slowly. When a few
diagram. drops of copper(II) sulphate are
hydrochloric acid added to the reaction mixture,
D Increase the pressure on the Solutions of hydrogen peroxide gas bubbles are produced
vigorously. Which statement best
reaction mixture used explains the effect of copper(II)
sulphate on the reaction?
28 Calcium carbonate is added Solution P: 50 cm3 of 2.0 mol A It lowers the activation energy.
to excess hydrochloric acid B It increases the collision
at 30 °C. The experiment is dm–3 H2O2 1.0 mol frequency between the
repeated at 40 °C. The volume Solution Q: 100 cm3 of reacting particles.
of carbon dioxide released for C It increases the concentration
each experiment is measured at Solution R: 1d0m0–3cmH23Oo2f 3.0 mol of sulphate ions and hence
room temperature and pressure. increases the rate of reaction.
Which of the following graphs dm–3 H2O2 D It causes the reacting
represents the results of these particles to collide in the
two experiments? Which of the following correct orientation.
A
statements are correct? 32 The diagram below shows the
B energy profile for the following
I The curve X is obtained by reaction

using solution R. X(g) + Y(g) → Z(g)

II The curve Y is obtained by The curves, P and Q, represent
two different paths for this
using solution P.
Rate of Reaction
III The curve Z is obtained by

using solution Q.

IV The curve Z is obtained by

using solution R.
A I and II only
B II and IV only
C I, II and III only
D I, II and IV only

333

1 reaction. What conclusion A I and II only Why is iron used in this process?
can be drawn based on the B III and IV only A To increase the rate of
diagram? C I, II and III only
A The reaction by path D II, III and IV only reaction between nitrogen
and hydrogen
P occurs at a higher 34 The reaction between Fe3+ and B To absorb the smell of
temperature. ’06 SeOqu32a–tioisnr:epresented by the ammonia
B The reaction by path P C To oxidise nitrogen to form
occurs at a higher rate than b2rFoew3+n+ SO32–+ H2O → g2rFeee2+n+ H2SO4 ammonia
by path Q. D To increase the yield of
C The activation energy for It is found that the change of colour ammonia
path P is (x + y) kJ. from brown to green occurs at
D The activation energy for a higher rate when the reaction 36 The rate of reaction between
path Q is (x – y) kJ. mixture is heated. This is due to the 1 mol dm–3 hydrochloric
I decrease in activation energy. acid and 3 g of magnesium
33 Consider the reaction between II increase in the frequency of powder is higher than the rate
magnesium and dilute of reaction between 1 mol
hydrochloric acid. collisions between Fe3+ and dm–3 ethanoic acid and 3 g of
SO32– ions. magnesium powder.
Mg(s) + 2HCl(aq) → III increase in the kinetic energy What is the explanation for this
MgCl2(aq) + H2(g) of Fe3+ and SO32– ions. observation?
IV increase in the frequency of A Hydrochloric acid is more
Which of the following will effective collisions. soluble in water than
increase the frequency of A I and II only ethanoic acid.
collisions between the reactants? B II and III only B The kinetic energy of
I Increase the concentration of C I, III and IV only hydrochloric acid is higher
D II, III and IV only than ethanoic acid.
hydrochloric acid. C Hydrochloric acid forms a
II Increase the temperature of 35 Iron is used in the Haber process soluble salt whereas
to manufacture ammonia from ethanoic acid forms an
reaction. nitrogen and hydrogen. insoluble salt.
III Use magnesium ribbon D The concentration of H+ ions
in hydrochloric acid is higher
instead of magnesium than ethanoic acid.
powder.
IV Remove the hydrogen gas
produced.

Structured Questions Experiment II

1 Diagram 1 shows two experiments to investigate one
factor that affects the rate of reaction between zinc

’06 and dilute sulphuric acid.

Experiment I

Diagram 1

(a) What is the factor that affects the rate of reaction

in Experiments I and II? [1 mark]

(b) State two constant variables in both experiments.

[2 marks]

(c) The following equation represents the reaction
that occurs in both the experiments.

Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)

Rate of Reaction 334

(i) Choose one of the products shown in recorded. The experiment is repeated using
different acids as shown below. All experiments
the equation that is most suitably used to are carried out at the same temperature.

measure the rate of reaction. [1 mark]

(ii) Give one reason for your answer in (i). Experiment Reactants Time taken (s)

[1 mark]

(d) Graph 1 shows the results for both experiments. I 50 cm3 of 1.0 mol dm–3 HCl t1
+

5.0 cm of magnesium
ribbon

II 50 cm3 of 1.0 mol dm–3 t2
CH3COOH
+
5.0 cm of magnesium
ribbon

Graph 1 (i) Write the ionic equation for the reaction

Based on Graph 1: between magnesium and an acid. [1 mark]

(i) Which experiment has a higher rate of (ii) Which is shorter t1 or t2? Why? [3 marks]

reaction? [1 mark] (c) Three experiments are carried out to investigate the
factors that affect the rate of the following reaction:
(ii) How do you come to this conclusion? [1 mark]
’08 Mg + 2HCl → MgCl2 + H2
(iii) Explain what happens after time t. [1 mark] The conditions of this experiment are shown below. 1

(iv) Why are both curves at the same level after

time t? [1 mark]

(e) State the conclusion for the experiments. [1 mark] Experiment Reactants Temperature
(°C)
(f) Another experiment was carried out using excess
zinc powder and dilute sulphuric acid with I Excess magnesium 35
different concentrations.
Sketch the curve of concentration of dilute powder
sulphuric acid against the time taken to collect a
fixed quantiy of the product. +

25 cm3 1.0 mol

dm–3 HCl

II Excess magnesium 25
ribbon
+
25 cm3 0.5 mol
dm–3 HCl

[2 marks] III Excess magnesium 35
ribbon
2 (a) Pcarorbpoanned, ioCx3iHde8, burns in excess oxygen to form +
and water as represented by the 25 cm3 1.0 mol
dm–3 HCl

equation The results of this experiment are shown in
Diagram 2.

C3H8 + 5O2 → 3CO2 + 4H2O

At time t, the rate of reaction of propane is 0.20
mol s–1.

Calculate

(i) the rate of consumption (using up) of

oxygen, and

(ii) the rate of production of carbon dioxide at

time t. [2 marks]

(b) 50 cm3 of 1.0 mol dm–3 of hydrochloric acid is Diagram 2
poured into a conical flask. A piece of 5.0 cm
magnesium ribbon is added to the acid. The time (i) Which curves (X, Y or Z) represent the results
taken to dissolve the magnesium completely is
of Experiments I, II and III? [3 marks]

335 Rate of Reaction

(ii) Give one reason why the final volume of (a) What is the average rate of reaction for Experiment
II? [2 marks]
gas obtained in curve X is half the final

volume of gas in curve Z. [1 mark] (b) Use the collision theory to explain why the

3 Three experiments were carried out to investigate the time taken for Experiment II is shorter than for
factors influencing the rate of reaction.
Three pieces of 0.12 g of magnesium ribbon are Experiment I. [3 marks]
added separately to excess hydrochloric acid.
The time taken for all the magnesium to dissolve is (c) Explain why the time taken for Experiment III is
taken. Table 1 shows the results of the experiments.
longer than for Experiment II. [3 marks]

(d) Suggest another method that can be used to

increase the rate of decomposition of hydrogen

Experiment I II III in Experiment II. [1 mark]
Reactants 0.12 g Mg
0.12 g Mg + excess 0.12 g Mg (e) Write a chemical equation for the catalytic
+ excess HCl(aq) + excess decomposition of hydrogen peroxide. [1 mark]
HCl(aq) HCl(aq) +
35 CuSO4(aq) (f) Experiments II and III were allowed to continue

32 35 until the decomposition of hydrogen peroxide

10 was completed. Sketch a graph of total volume

Temperature 30 of gas against time for Experiments II and III on
(oC) 60
the same axes. [2 marks]
Time (s)
5 Dilute sulphuric acid reacts with sodium thiosulphate
1 Table 1 solution to produce sulphur. The presence of sulphur
causes the solution to become cloudy.
(a) Write the chemical equation for the reaction between Five experiments were carried out to study the rate
of reaction between dilute sulphuric acid and sodium
magnesium and hydrochloric acid. [1 mark] thiosulphate solution. The reaction takes place in a
conical flask placed over a white paper marked with
(b) Calculate the maximum volume of hydrogen gas a cross, ‘X’. The time taken for the cross to disappear
from view was recorded.
produced in Experiment I. [2 marks] In each experiment, the volume and concentration of
sodium thiosulphate solution were kept constant. The
(c) Predict the maximum volume of hydrogen gas experimental results are shown in Table 3.

produced in Experiment III. Explain your answer.

[2 marks]

(d) Calculate the average rate of reaction in

Experiment II. [1 mark] Concentration Temperature Time taken
of acid (°C) (s)
(e) (i) Which experiment has the highest rate? Experiment
(mol dm–3) 65
Justify your answer. 45
85
(ii) Sketch the graphs for the volume of A 0.15 20 55
B 0.10 30 105
hydrogen gas against time for Experiments C 0.10 20
D 0.05 30
I, II and III on the same axes. [3 marks] E 0.05 20

[Relative atomic mass: Mg = 24; Molar gas volume,

24 dm3 at r.t.p.]

4 Three experiments were carried out to study the effect Table 3
of iron(III) oxide, Fe2O3, on the rate of decomposition
of 0.5 mol dm–3 hydrogen peroxide. Table 2 shows (a) The reaction between dilute sulphuric acid and
the mixtures of substances used and the time taken sodium thiosulphate (Na2S2O3) produces sulphur,
to collect 30 cm3 of the colourless gas given off in sodium sulphate, sulphur dioxide and water.
each experiment.

Experiment Mixture of substances Time taken Complete the following equations:

I 20.0 cm3 of 0.5 mol dm–3 H2O2 A few ((iii)) NS2aO2S322–O+3 +2HH+2S→O4_→___ ____ + ____ + ____
weeks + ____ + ____

[2 marks]

II 20.0 cm3 of 0.5 mol dm–3 H2O2 35 s (b) (i) Which of the experiments shown above

+ 0.2 g of Fe2O3 should be chosen to compare the effect

of concentration of the acid on the rate of

III 20.0 cm3 of 0.5 mol dm–3 H2O2 45 s reaction? [1 mark]

+ 25.0 cm3 of water + 0.2 g of (ii) Give one reason why you chose these

Fe2O3 experiments. [1 mark]

(iii) What conclusion can be made from the

Table 2 experiments in (i)? [1 mark]

Rate of Reaction 336

(c) (i) Which of the experiments should be chosen (a) Table 5 shows the conditions used for carrying
out Experiment 1. Complete Table 5 to predict
to compare the effect of temperature on the conditions used for obtaining the results of
Experiments 2, 3 and 4. In each case, state the
the rate of reaction? [1 mark] constant variables and manipulated variables
used and briefly explain your answer.
(ii) Give one reason for your choice. [1 mark]

(iii) What conclusion can be made from the

experiments you have chosen in (i)?

[1 mark]

(d) Explain your answer in (c)(iii) based on collision Experi­ Experi­ Experi­ Experi­
ment ment ment ment
theory. [2 marks]
1234
6 Four experiments were carried out to investigate the
decomposition of hydrogen peroxide to form water Volume of
and oxygen gas. hydrogen peroxide 40 40 … …
(cm3)

2H2O2(aq) ⎯M⎯n⎯O2→ 2H2O(l) + O2(g) Volume of water 40 40 40 0
(cm3)

Temperature (°C) 30 30 32 30

The total volume of oxygen evolved at one-minute Mass (ogf)MnO2 1.0 … 1.0 1.0
intervals were recorded in Table 4. used

Volume of oxygen gas released (cm3) Table 5 [5 marks] 1
Time
(min) Experiment Experiment Experiment Experiment (b) Copper(II) oxide is a less effective catalyst than

1234 manganese(IV) oxide for the decomposition of
00 0 0 0
1 18 0 25 34 hydrogen peroxide. If copper(II) oxide is used to
2 33 0 35 60
3 36 0 37 69 replace manganese(IV) oxide in Experiment 1,
4 37 0 38 75
5 38 0 38 76 what is the effect of this change on
6 38 0 38 76
(i) the volume of oxygen collected at 1.0 minute?
Table 4
(ii) the volume of oxygen collected after the

reaction has completed?

Explain your answers. [2 marks]

(c) ‘A catalyst remains chemically unchanged at the

end of the reaction’.

(i) What is meant by chemically unchanged?

(ii) How would you prove that your answer in

(i) is correct? [3 marks]

Essay Questions

1 (a) Two experiments are carried out to study the rate
of reaction between iron and dilute acids.

Experiment Reactants

I 1.12 g of iron and 50 cm3 of
2.0 mol dm–3 sulphuric acid

II 1.12 g of iron and 50 cm3 of
2.0 mol dm–3 hydrochloric acid

The following graphs show the results of the
experiments.

337 Rate of Reaction

Based on the graph: (c) mEthaanngeandaioteic(VaIcIi)d,slHo2wCly2Oa4t, decolourises potassium
room temperature. The
(i) Calculate the average rate of reaction for

’07 Experiment I. [2 marks] reaction is catalysed by manganese(II) sulphate,

(ii) Explain the difference in the rate of reaction MnSO4. Describe how you would prove that the

between Experiment I and Experiment II sulphate ions, SO42–, do not act as a catalyst in

before 100 s. [6 marks] this reaction. [3 marks]

(b) Describe an experiment to show that lead(IV) 4 (a) Two experiments are carried out to study the rate
of reaction between zinc and two acids, R and T.
oxide is a more effective catalyst than copper(II)
’07 The data for the experiments are shown below.
oxide for the decomposition of hydrogen peroxide.

Your answer should include a labelled diagram

on the apparatus set-up for the experiment. Experiment Reactants Observation Products

[12 marks] I Excess zinc The Zinc

2 (a) Iron powder will dissolve in cold dilute hydrochloric and 25 cm3 temperature sulphate and

acid while coarse iron filings do not dissolve until of 1.0 mol of the hydrogen

the acid is heated. Explain these observations. dm–3 acid R reaction

[7 marks] mixture

(b) There is a high risk of explosions occurring in coal increases

mines. Explain why this is so. [6 marks] II Excess zinc The Zinc

(c) ‘Temperature is important in preserving food’. Give and 25 cm3 temperature chloride and

one example from your daily life to justify this of 1.0 mol of the hydrogen

1 statement. [4 marks] dm–3 acid T reaction

(d) When a drop of blood is added to hydrogen mixture

peroxide solution, a vigorous effervescence occurs. increases

Explain this observation. [3 marks] (i) State the names of the acids used in

3 (a) (i) Define the term rate of reaction? experiments I and II. [2 marks]

[2 marks] (ii) Write the chemical equation for the reaction

(ii) At high temperatures and pressures, nitrogen that occurs in Experiment I. [2 marks]

reacts with hydrogen to form ammonia. (iii) Draw the energy profile diagram for the

N2(g) + 3H2(g) 2NH3(g) reaction in Experiment II. On the energy

Use the collosion theory to explain how profile diagram, show the

high pressure increases the rate of reaction • activation energy wwiitthhoautcactaatlaylsyts,t,EaE’,a,
• activation energy
between nitrogen and hydrogen. [5 marks]
• heat of reaction, ΔH.
Explain the energy profile diagram.
(b) Describe an experiment to demonstrate the effect
[10 marks]
of the concentration of sodium thiosulphate on

the rate of reaction between hydrochloric acid and (b) Explain the factors that affect the rate of reaction
in the following daily activities:
sodium thiosulphate. Draw the apparatus used for
(i) Combustion of charcoal
this experiment and state the hypothesis and (ii) Cooking food in a pressure cooker [6 marks]

variables in this experiment. [10 marks]

Experiments The experiment was repeated using sodium
thiosulphate solutions at 30 °C, 35 °C, 40 °C and
1 Sodium thiosulphate solution reacts with dilute 45 °C.
sulphuric acid to produce a yellow precipitate Diagram 1 shows the stopwatch readings for each of
of sulphur. 50 cm3 of 0.10 mol dm–3 sodium the experiments.
thiosulphate solution at 25 °C was measured into a
250 cm3 conical flask. The conical flask was placed
on a white paper marked with the ‘X’ sign.
5 cm3 of 0.50 mol dm–3 sulphuric acid was added
to the sodium thiosulphate solution and the mixture
shaken. At the same time, the stopwatch was started.
The time was taken as soon as the ‘X’ sign was no
longer visible.

Rate of Reaction 338

at 25 °C at 30 °C at 35 °C at 40 °C at 45 °C

T ime, t1 ______ s Time, t2 ______ s Time, t3 ______ s Time, t4 ______ s Time, t5 ______ s
Diagram 1

(a) Record the readings of the stopwatch in the spaces provided in Diagram 1. [3 marks]

(b) State the variables in this experiment. [3 marks]
Manipulated variable:
Responding variable:
Constant variable:

(c) Construct a table containing the information on temperature, time and —t—im—1——e for the experiments. [2 marks]

(d) (i ) Draw a graph of temperature against ——ti—m1——e on a graph paper. [4 marks] 1

(ii) Based on the graph in (i), what conclusion can be drawn from this experiment? [3 marks]

(e) Predict the time taken for the ‘X’ sign to disappear if the experiment is repeated at 50 °C. [3 marks]

(f) State the hypothesis for this experiment. [3 marks]

(g) Based on your hypothesis, explain why meat and fish are always kept in refrigerators. [3 marks]

2 An experiment was carried out to investigate the rate of reaction between granulated zinc and dilute
hydrochloric acid. The results of the experiment are shown below.

’09

Time (s) 0 10 20 30 40 50 60 70 80

Burette reading 50.00 40.00 33.50 …… …… 21.50 20.00 17.50 17.50
(cm3)

Volume of gas 0.00 10.00 16.50 …… …… 28.50 30.00 32.50 32.50
evolved (cm3)

Diagram 2 shows the burette readings at 30 seconds and 40 seconds respectively.

29 25

TC 56

28 24

At 30 s At 40 s

Diagram 2 [3 marks]
[3 marks]
(a) Based on this experiment, what is meant by the rate of reaction? [3 marks]
(b) Based on Diagram 2, what are the volumes of gas evolved at 30 seconds and 40 seconds?
(c) State one conclusion, based on the experimental results.

339 Rate of Reaction

2CHAPTER FORM 5

THEME: Interaction between Chemicals

Carbon Compounds

SPM Topical Analysis

Year 2008 2009 2010 2011

Paper 1 2 31 2 31 2 31 2 3

Section ABC ABC AAC ABC

Number of questions 4 1 – – 1 2 1 1 – – 7 – —21 – – 4 – – 1 –

ONCEPT MAP

CARBON COMPOUNDS

Organic carbon compounds: Isomerism (same molecular
Produce CO2 and H2O on complete combustion formula, different structural formulae)

Hydrocarbons (elements C and H only) Non-hydrocarbons Natural rubber
Physical properties: (Elements: C, H, O) (natural polymers)
Insoluble in water, low melting and boiling points, • Coagulation
non-conductors of electricity • Vulcanisation

Alkanes hydrogenation Alkenes hydration Alcohols oxidation Carboxylic acids
(saturated) (non-saturated) dehydration Reactions: Reactions:
Reactions: Reactions: • Combustion • With metals/metal
• Combustion • Combustion • Oxidation
• Substitution • Addition • Dehydration carbonates/alkalis
• Polymerisation • Esterification to form salts
• Esterification

esterification

Esters
Physical properties:
• Sweet fruity smell
• Insoluble in water

Fats and Oils
• Fats: saturated, higher melting point
• Oils: unsaturated, lower melting point

2.1 Carbon Compounds 8 Almost all organic compounds contain the
elements carbon and hydrogen. Hence the
1 Carbon compounds are compounds that complete combustion of carbon compounds
contain the element carbon. produces carbon dioxide and water.

2 Carbon compounds can be classified into two Carbon compounds
groups: inorganic compounds and organic
compounds. Organic Inorganic
compounds compounds
3 Organic compounds are carbon compounds
in which carbon is bonded to other elements Example: Example:
by covalent bonds. Examples of organic • Hydrocarbons • Hydrogen
compounds are hydrocarbons, alcohols, • Alcohols
carboxylic acids and esters. • Carboxylic acids carbonates
• Esters • Carbonates
4 Most carbon compounds are derived from • Carbohydrates • Carbides
living organisms. Nowadays many organic • Oxides of carbon
compounds can be synthesised in laboratories. • Cyanides Activity 2.1 2

5 Most inorganic compounds do not contain Organic compounds are the largest group of chemicals
carbon. Examples of inorganic compounds we know today, numbering in thousands. All the food
containing the element carbon are carbonates, and medicines we consume are organic compounds
hydrogen carbonates, oxides of carbon and as well as most of the synthetic products such as
cyanides. clothing and household materials.

6 Carbon atom (proton number 6) has the
electron arrangement: 2.4. Hence, each carbon
atom can form four covalent bonds in organic
compounds.

7 The covalent bonds in carbon compounds may
be
(a) single bond,
(b) double bond or
(c) triple bond.

To investigate the products formed by complete combustion
of organic compounds

Materials Ethanol and palm oil, limewater, Apparatus Filter funnel, test tubes, delivery tubes,
Procedure ice and water, anhydrous cobalt(II) spirit lamp, suction pump and beaker.
chloride paper.

Figure 2.1 Combustion of organic compounds 1 A filter funnel is connected to the suction pump

via test tubes A and B where test tube A is dipped

in ice water and test tube B contains limewater.
2 A spirit lamp filled with ethanol is lit and placed

under the filter funnel and the suction pump

turned on.
3 The changes in test tubes A and B are noted.
4 The liquid collected in test tube A is tested with

anhydrous cobalt(II) chloride paper.
5 Steps 1 to 4 are repeated using a spirit lamp with

palm oil.

341 Carbon Compounds

Results Inference Conclusion

Test Observation The colourless 1 The combustion of organic compounds such as
tube liquid formed in ethanol and palm oil produces water and carbon
test tube A is water dioxide.
A A colourless liquid
is formed and it Carbon dioxide gas 2 During the combustion of an organic compound,
changes anhydrous is produced (a) the carbon combines with oxygen to form
cobalt(II) chloride carbon dioxide,
paper from blue to (a) the hydrogen combines with oxygen to form
pink water.

B Limewater turns
milky

Hydrocarbons excess oxygen (complete combustion), carbon
dioxide and water are produced.
1 Hydrocarbons are organic compounds that 8 Incomplete combustion of hydrocarbons will
contain the elements carbon and hydrogen
only. produce water, carbon dioxide, carbon monoxide

2 Hydrocarbons that have only single covalent and carbon (as soot).
bonds between all the carbon atoms in the
2 molecule are called saturated hydrocarbons.

Example Propane

H H H carbon-carbon Hydrocarbon

| | | single bond

H—C—C—C—H Saturated hydrocarbons Unsaturated
| | | have only single bonds hydrocarbons have double
between the carbon or triple bonds between
H H H atoms. the carbon atoms.
Examples: Ethane, propane Examples: Ethene, propene
3 Hydrocarbons that have at least one carbon-
carbon double bond (C = C) or triple bond 1 SPM
(C ≡ C) in the molecule are called unsaturated ’10/P1
hydrocarbons.

Example Propene

H H H carbon-carbon The complete combustion of 0.1 mol of a
| | | double bond hydrocarbon Z in excess oxygen produces 0.3 mol
of carbon dioxide and 0.4 mol of water. Determine
H—C=C—C—H the molecular formula of hydrocarbon Z.
|
H

4 The main sources of hydrocarbons are: Solution

(a) Petroleum (crude oil) Since 0.1 mol of Z produces 0.3 mol of carbon
dioxide and 0.4 mol of water, 1 mol of Z will
(b) Natural gas produce 3 mol of carbon dioxide and 4 mol of water.

(c) Coal C + O2 → CO2
5 Petroleum is a complex mixture of
The number of moles of carbon in 1 mol of Z
hydrocarbons.
6 Fractions of hydrocarbons are separated by = the number of moles of carbon dioxide = 3.

a process called fractional distillation. The 2H + —12 O2 → H2O

fractions are separated based on the difference

in boiling points. The fractions with lower The number of moles of hydrogen in 1 mol of Z

boiling points will be distilled off earlier. = 2 3 the number of moles of water = 2 3 4 = 8.
7 Hydrocarbons contain carbon and hydrogen
Hence the molecular formula of Z is C3H8.
only. Thus when hydrocarbons are burnt in

Carbon Compounds 342

2.1 7 Table 2.1 shows the prefixes used to indicate
the number of carbon atoms per molecule
1 Classify the following substances into organic of an organic compound, the names and
compounds and inorganic compounds. molecular formulae of the first ten alkanes.

Rubber, sugar, limestone, carbon dioxide, Table 2.1 Prefixes used to indicate the number
calcium carbonate, sand, sodium chloride, of carbon atoms, names and molecular
vinegar, polyvinylchloride, urea, ammonium formulae of alkanes
sulphate.
Number of Prefix Name of Molecular
2 (a) What is meant by hydrocarbon? carbon atoms alkane formula
(b) State three sources of hydrocarbon. per molecule

3 State two main products that are formed when 1 Meth Methane CH4
rubber is burnt in excess air. Explain your answer.
2 Eth Ethane C2H6

3 Prop Propane C3H8

2.2 Alkanes 4 But Butane C4H10

5 Pent Pentane C5H12

1 Alkanes are saturated hydrocarbons with the 6 Hex Hexane C6H14

general formula Cn Hs2ant+u2,rwatheedre n = 1, 2, 3… 7 Hept Heptane C7H16
2 Alkanes are called hydrocarbons
8 Oct Octane C8H18 2

SPM because the molecules contain only single 9 Non Nonane C9H20

’11/P1 covalent bonds between carbon atoms.

3 The molecular formula is a chemical formula 10 Dec Decane C10H22

that shows the actual number of atoms of

each element present in one molecule of the 8 The structural formula of an organic compound
is the chemical formula that shows the
substance. arrangement of atoms and covalent bonds
between atoms in a molecule of the compound.
4 The molecular formula of an alkane can be
9 Note that when writing the structural formula
obtained by substituting n in the general of alkanes,
(a) each carbon atom should have four single
formula Cn H2n+2 with the number of carbon covalent bonds.
atoms. (b) each hydrogen atom should have one single
covalent bond.
5 In the naming of alkanes according to the (c) the carbon atoms are connected by single
bonds.
IUPAC system, all members of the alkane
10 The structural formulae of the first ten straight
series have their names ending with -ane chain alkanes are shown as follows:

(IUPAC is the abbreviation for International

Union of Pure and Applied Chemistry).

6 The first part (prefix) of the name of an

alkane depends on the number of carbon

atoms in the molecule.

Table 2.2 The structural formulae of the first ten straight chain alkanes

H H H H H H
| | | | | |

H–C–H H–C – C–H H–C–C–C–H
| | | | | |
H H H H H H

methane (CH4) ethane (C2H6) propane (C3H8)

H H H H H H H H H H H H H H H
| | | | | | | | | | | | | | |

H–C–C–C–C–H H–C–C–C–C–C–H H–C–C–C–C–C–C–H
| | | | | | | | | | | | | | |
H H H H H H H H H H H H H H H

butane (C4H10) pentane (C5H12) hexane (C6H14)

343 Carbon Compounds