SUCCESS CHEMISTRY

(c) chlorine acts as the oxidising agent because Sulphur dioxide usually acts as a reducing agent. However,
it oxidises iron and is itself reduced. in the following reactions, sulphur dioxide acts as an
oxidising agent because hydrogen sulphide and carbon
(d) iron acts as the reducing agent because it are stronger reducing agents than sulphur dioxide.
reduces chlorine and is itself oxidised.

5 When chlorine gas is passed into potassium
bromide solution, the following reaction occurs:

(oxidation number increases) (a) oxidation number of S increases
oxidation (oxidation)

0 –1 –1 0 0
–2

Cl2(g) + 2KBr(aq) → 2KCl(aq) + Br2(aq) SO2(g) + 2H2S(g) ⎯⎯⎯→ 2H2O(l) + 3S(s)
+4 0

reduction
oxidation number of S decreases
(oxidation number decreases) (reduction)

(b) oxidation


0 +4

The oxidation number of potassium in the above +4S O2(g) + C(s) ⎯⎯⎯→ CO2(g) + S(s)
reaction does not change because potassium does not 0
take part in the reaction.

3 reduction

In this reaction, 7 A reaction is not a redox reaction if the substances
(a) bromide ion (Br–) is oxidised to bromine involved in the reaction do not undergo any
because the oxidation number of bromine changes in oxidation numbers. For example,
increases from –1 to 0.
(b) chlorine (Cl2) is reduced to chloride ion +1 –2 +1 +1 +6 –2 +1 +6 –2 +1 –2
(Cl–) because the oxidation number of
chlorine decreases from 0 to –1. 2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)
(c) chlorine acts as an oxidising agent and
potassium bromide acts as a reducing agent. The reaction between sodium hydroxide (NaOH)
and sulphuric acid (H2SO4) is a neutralisation
6 Ammonia reacts with copper(II) oxide as reaction and not a redox reaction. As a result,
represented by the equation: the oxidation numbers of all the elements
(sodium, oxygen, hydrogen and sulphur) are
(oxidation number increases) the same before and after the reaction.
oxidation

–3 +2 0 0

2NH3(g) + 3CuO(s) → N2(g) + 3H2O(l) + 3Cu(s) Oxidation and reduction always take place together. A
redox reaction must have
reduction • a substance that undergoes oxidation and acts as
(oxidation number decreases)
the reducing agent, and
In this reaction, • another substance that undergoes reduction and
(a) ammonia is oxidised to nitrogen because
the oxidation number of nitrogen increases acts as the oxidising agent.

from –3 to 0. Oxidation and Reduction in Terms of
(b) copper(II) oxide is reduced to copper Electron Transfer

because the oxidation number of copper 1 In terms of electron transfer,
(a) oxidation is defined as the loss of electrons
decreases from +2 to 0. SPM from a substance. If a substance loses
(c) copper(II) oxide acts as an oxidising agent ’05/P1 electrons during a reaction, it has been
oxidised.
and ammonia acts as a reducing agent.
(d) the oxidation numbers of hydrogen and

oxygen remain unchanged.

Oxidation and Reduction 394

(b) reduction is defined as the gain of 5 3
electrons by a substance. If a substance
gains electrons, it has been reduced. Write the half-equation for the reduction of acidified
manganate(VII) ion (MnO4–) to manganese(II) ion
2 During a redox reaction, transfer of electrons (Mn2+) in the presence of acid.
occurs between the reactants.
Solution
An oil rig is used for getting oil and gas out of the ground Step 1: Write the reactants and products involved in
in the petroleum industry. Use ‘OIL RIG’ to help you
remember oxidation and reduction in terms of electron the reaction.
transfer.
OIL : OXIDATION IS LOSS OF ELECTRONS MnO4– + H+ → Mn2+ + H2O
RIG : REDUCTION IS GAIN OF ELECTRONS Step 2: Balance the number of atoms on both sides

3 The reactant that loses electrons undergoes of the equation:
oxidation and acts as a reducing agent.
For example, MnO4– + 8H+ → Mn2+ + 4H2O
Step 3: Balance the number of charges on both
Na(s) o⎯xi⎯da⎯tio→n Na+(aq) + e– … (1)
sides of the equation:
In this reaction,
(a) sodium atoms undergo oxidation by losing MnO4– + 8H+ → Mn2+ + 4H2O
electrons to form sodium ions (Na+). total charge total charge
(b) sodium acts as the reducing agent. = (–1) + (+8) = +7 = +2

4 A substance that accepts electrons undergoes to balance the charges,
reduction and acts as an oxidising agent. add 5e– to the left of
For example, the equation

Cl2(g) + 2e– ⎯red⎯uc⎯tio→n 2Cl–(aq) … (2) MnO4–(aq) + 8H+(aq) + 5e– →
Mn2+(aq) + 4H2O(l)
In this reaction,
(a) eealecchtrochnlsotroinfeormmoltewcoulceh(loCrli2d)eaicocenpst(sCtlw–)o. 6 (a) Zinc reacts with hydrochloric acid as
(b) chlorine acts as the oxidising agent and is represented by the equation
itself reduced.
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
5 Balancing half-equations for oxidation and
SPM reduction The ionic equation for the reaction is
’06/P1 Equations (1) and (2) as shown above are
oxidation
known as half-equations. Half-equations must (loss of electrons)
be balanced in terms of
(a) the number of atoms, and Zn(s) + 2H+(aq) → Zn2+(aq) + H2(g)
(b) the number of charges.
reduction
Photographic films are coated with silver bromide, (gain of electrons)
AgBr. When the film is exposed to light, the following
redox reaction occurs: (b) The transfer of electrons can be represented
by the following half-equations:
2Ag+ + 2Br–­ → 2Ag + Br2
The amount of silver produced depends on how much Zn(s) → Zn2+(aq) + 2e– ... oxidation
light gets through the camera lens. In this reaction,
silver ions are reduced to silver by the gain of electrons reducing agent
and bromide ions are oxidised to bromine by the loss
of electrons. 2H+(aq) + 2e– → H2(g) ... reduction
oxidising agent

(c) In the reaction between hydrochloric acid
and zinc, zinc is oxidised to zinc chloride
whereas hydrochloric acid is reduced to
hydrogen.

395 Oxidation and Reduction

(d) Hydrochloric acid acts as an oxidising 9 Combustion of metals in chlorine
agent by accepting electrons and is itself Figure 3.3 shows the combustion of copper in
reduced. Conversely, zinc acts as the reducing chlorine. When the hot copper foil is placed
agent by donating electrons and is itself in a gas jar of chlorine, a vigorous reaction
oxidised. occurs and a green precipitate of copper(II)
chloride, CuCl2 is formed.
7 In terms of transfer of electrons, oxidising
agents are electron acceptors while reducing
agents are electron donors.

8 If a coil of copper is placed in a solution of
silver nitrate, the copper slowly dissolves and
the solution turns blue. At the same time, the
copper coil becomes coated with a layer of
silver metal (Figure 3.2).

Figure 3.3 Combustion of copper in chlorine

3 (a) In the reaction between copper and
chlorine to form copper(II) chloride, a
transfer of electron occurs between copper
metal and chlorine gas.

Figure 3.2 Oxidation of copper oxidation
(loss of electrons)
(a) The overall equation for the reaction is
Cu(s) + Cl2(g) → CuCl2(s)
Cu(s) + 2AgNO3(aq) →
Cu(NO3)2(aq) + 2Ag(s) reduction
(gain of electrons)
The reaction can be represented by the
ionic equation:
(b) The transfer of electrons can be represented
oxidation by the half-equations as shown below:
(loss of electrons)
Cu(s) → Cu2+(s) + 2e– … oxidation

Cu(s) + 2Ag+(aq) → Cu2+(aq) + 2Ag(s) Cl2(g) + 2e– → 2Cl– (s) ... reduction

(c) In the reaction between copper and
reduction chlorine, the copper atom (Cu)
(gain of electrons)
(i) loses electrons
(b) In this reaction, each silver ion (Ag+) accepts (ii) undergoes oxidation
one electron to form a silver atom (Ag). (iii) is oxidised to copper(II) ion, Cu2+
(iv) acts as a reducing agent
Ag+(aq) + e– → Ag(s) ... reduction
An oxidising agent is an electron acceptor. (d) Conversely, the chlorine molecule (Cl2)
(i) gains electrons
Hence, silver ion acts as an oxidising (ii) undergoes reduction
agent in this reaction. (iii) is reduced to chloride ions, Cl–
(c) Conversely, each copper atom donates two (iv) acts as an oxidising agent
electrons and are converted to copper(II) 10 Combustion of metals in oxygen
ion (Cu2+) in the aqueous solution.
When metals burn in oxygen,
Cu(s) → Cu2+(aq) + 2e– ....oxidation (a) the metals undergo oxidation by losing
Reducing agents are electron donors. Hence,
electrons to form metal ions,
copper acts as a reducing agent. (b) the oxygen undergoes reduction by gaining

electrons to form oxide ions (O2–).
The combustion of lead in oxygen is studied
in Activity 3.1.

Oxidation and Reduction 396

To investigate the combustion of metals in oxygen and
chlorine

Apparatus Gas jar, tongs, combustion spoon, gas 2 The magnesium ribbon is held with a pair of
Materials jar and Bunsen burner. tongs and lit in the Bunsen burner.
Procedure
Magnesium ribbon, sodium and 3 It is quickly placed into a gas jar filled with oxygen.
chlorine. 4 Any changes that occur are recorded.

(B) Reaction of sodium with chlorine

Figure 3.5 The combustion of sodium in chlorine

Figure 3.4 The combustion of magnesium in oxygen 1 A small piece of sodium metal is placed in a Activity 3.1 3

(A) Combustion of magnesium in oxygen combustion spoon and heated.
1 A piece of 5 cm magnesium ribbon is cleaned 2 When the sodium metal starts to burn, it is quickly

with sandpaper. placed in a gas jar filled with chlorine gas.
3 Any changes that occur are recorded.
Observation

Experiment Observation
Combustion of magnesium in oxygen
• The magnesium ribbon burns with a bright white flame.
Reaction of sodium with chlorine • White fumes are produced.
• A white powder is formed.

• The sodium metal burns with a yellow flame.
• White fumes are produced.
• A white powder is formed.

Discussion 5 In this reaction,
• magnesium acts as the reducing agent because
(A) Combustion of magnesium in oxygen it reduces oxygen to oxide ion.
• oxygen acts as the oxidising agent because it
1 The combustion of magnesium in oxygen produces oxidises magnesium to magnesium ion.

magnesium oxide (white powder). (B) Combustion of sodium in chlorine
1 The combustion of sodium in chlorine produces
2 Magnesium is oxidised by losing electrons to
sodium chloride (white powder).
form magnesium ions, Mg2+. 2 Sodium is oxidised by losing electrons to form

Half-equation: Mg(s) → Mg2+(s) + 2e– sodium ions, Na+.
3 Oxygen is reduced by gaining electrons to form
Half-equation: Na(s) → Na+(s) + e–
oxide ion, O2–.
3 Chlorine is reduced by gaining electrons to form
4 Half-equation: O2(g) + 4e– → 2O2–(s) chloride ion, Cl–.
The overall equation for the reaction is
Half-equation: Cl2(g) + 2e– → 2Cl–(s)
oxidation
(loss of electrons)


2Mg(s) + O2(g) → 2MgO(s)


reduction
(gain of electrons)

397 Oxidation and Reduction

4 The overall equation for the reaction is Conclusion
oxidation
(loss of electrons) 1 In the combustion of magnesium in oxygen,
(a) magnesium undergoes oxidation to form
Mg2+ ions,
2Na(s) + Cl2(g) → 2NaCl(s) (b) oxygen undergoes reduction to form O2–
ions.

reduction 2 In the combustion of sodium in chlorine,
(gain of electrons) • sodium act as reducing agent by losing
5 In this reaction, electrons,
• chlorine act as oxidising agent by gaining
• sodium acts as the reducing agent because it electrons.
reduces chlorine to chlorine ion.

• chlorine acts as the oxidising agent because it
oxidises sodium to sodium ion.

Conversion of Fe2+ Ions to Fe3+ Ions SPM
and Vice Versa ’08/P2,
’09/P1

3 There are some chemicals such as hydrogen peroxide 1 Iron metal (Fe) exhibits two oxidation states,
(H2O2) and nitric(III) acid (nitrous acid, HNO2) which +2 and +3.
can act as an oxidising agent or a reducing agent
depending on the conditions of the reaction. For example, 2 Fe2+ ions can be converted to Fe3+ ions. Similarly,
in reaction (1), hydrogen peroxide acts as an oxidising Fe3+ ions can be converted to Fe2+ ions.
agent, but in reaction (2), it acts as a reducing agent.
Oxidation of Fe2+ to Fe3+
H2O2 + 2I– + 2H+ → I2 + 2H2O … (1)

oxidising reducing 1 Iron(II) ion, Fe2+, can be converted to iron(III)
agent agent ions, Fe3+, by oxidation reaction.

5H2O2 + 2MnO4– + 6H+ → 2Mn2+ + 8H2O + 5O2 … (2) oxidation (loss of electrons)

Fe2+(aq) → Fe3+(aq) + e–

reducing oxidising 2 Potassium manganate(VII) is an oxidising
agent agent agent that can oxidise Fe2+ ions to Fe3+ ions.

3 (a) When acidified potassium manganate(VII)
solution is added to a solution of iron(II)

salt, decolourisation occurs. MnO4– ions are
reduced to Mn2+ ions while Fe2+ ions (pale

green) are oxidised to Fe3+ ions (brown).

oxidation

(loss of electrons)

MnO4–(aq) + 8H+(aq) + 5Fe2+(aq) → Mn2+(aq) + 4H2O(l) + 5Fe3+(aq)
green brown
reduction

(gain of electrons)

A catalytic converter

Catalytic converters are fitted to the exhaust pipes of Dilute sulphuric acid is always used to acidify KMnO4
cars to reduce air pollution. In the catalytic converter, solution.
the following redox reactions take place to convert
poisonous gases (NO, CO and petrol vapour) to non- (b) The half-equations for the reactions are:
poisonous gases. For example, Fe2+(aq) → Fe3+(aq) + e–
(oxidation – loss of electrons)
2NO(g) + 2CO(g) → N2(g) + 2CO2(g)
oxidising reducing


agent agent

Oxidation and Reduction 398

MnO4–(aq) + 8H+(aq) + 5e– → reduction (gain electron)
Mn2+(aq) + 4H2O(l)
(reduction – gain of electrons) SO32– + H2O + 2Fe3+ → 2Fe2+ + H2SO4
brown green
4 The formation of Fe3+ ions can be confirmed
by using sodium hydroxide solution. When oxidation (lose electron)
sodium hydroxide solution is added to the
reaction product, a brown precipitate of
iron(III) hydroxide, Fe(OH)3, insoluble in Sodium sulphite acts as the reducing
excess NaOH(aq) is obtained. agent and reduces iron(III) ions to iron(II)
ions and is itself oxidised to sulphate ions
Fe3+(aq) + 3NaOH(aq) → Fe(OH)3(s) + 3Na+(aq) (SO42–).

5 Other oxidising agents that can be used to (b) The half-equations for the reaction are:
oxidise Fe2+ to Fe3+ are as follows.
(a) Chlorine gas or chlorine water Fe3+(aq) + e– → Fe2+(aq)
(reduction – gain of electrons)
Cl2(aq) + 2Fe2+(aq) → 2Fe3+(aq) + 2Cl– (aq)
SO32–(aq) + H2O(l) →
(b) Liquid bromine SO42–(aq) + 2H+(aq) + 2e–
(oxidation – loss of electrons)
Br2(l) + 2Fe2+(aq) → 2Fe3+(aq) + 2Br–(aq)
3 The formation of Fe2+ ions can be confirmed 3
(c) Acidified potassium dichromate(VI) by using sodium hydroxide solution. When
solution (acidified with dilute sulphuric
acid) sodium hydroxide solution is added to the
reaction product, a dirty green precipitate
Cr2O72–(aq) + 14H+(aq) + 6Fe2+(aq) → of iron(II) hydroxide, Fe(OH)2, insoluble in
6Fe3+(aq) + 2Cr3+(aq) + 7H2O(l) excess NaOH(aq), is obtained.

(d) Concentrated nitric acid Fe2+(aq) + 2NaOH(aq) →
Fe(OH)2(s) + 2Na+(aq)
HNO3(aq) + 3Fe2+(aq) + 3H+(aq) →
3Fe3+(aq) + 2H2O(l) + NO(g) 4 Other reducing agents that can be used to reduce
Fe3+ ions to Fe2+ ions include the following.
(e) Acidified hydrogen peroxide (a) Metals more reactive (electropositive)
than iron. For example, zinc.
H2O2(aq) + 2H+(aq) + 2Fe2+(aq) →
2Fe3+(aq) + 2H2O(l) Zn(s) + 2Fe3+(aq) → 2Fe2+(aq) + Zn2+(aq)

(b) Sulphur dioxide

Reduction of Fe3+ to Fe2+ SO2(g) + 2H2O(l) + 2Fe3+(aq) →
2Fe2+(aq) + 2H+(aq) + H2SO4(aq)
1 Iron(III) ions, Fe3+, can be converted to iron(II) (c) Potassium iodide
ions, Fe2+, by reduction.
2KI(aq) + 2Fe3+(aq) →
reduction (gain electron) 2Fe2+(aq) + 2K+(aq) + I2(s)
Fe3+(aq) + e– → Fe2+(aq)
(d) Hydrogen sulphide

2 (a) When sodium sulphite (Na2SO3) solution H2S(aq) + 2Fe3+(aq) →
is added to iron(III) chloride, and the 2Fe2+(aq) + 2H+(aq) + S(s)
mixture is acidified with dilute sulphuric
acid, the colour of the solution changes (e) Tin(II) chloride solution
from brown to light green.
Sn2+(aq) + 2Fe3+(aq) →
2Fe2+(aq) + Sn4+(aq)

399 Oxidation and Reduction

To study the oxidation of Fe2+ ions to Fe3+ ions and the SPM
reduction of Fe3+ ions to Fe2+ ions
’09/P2

Apparatus Test tubes and droppers. 4 Sodium hydroxide solution is then added to the
reaction mixture slowly until in excess.
Materials FeSO4, KMnO4, FeCl3, Na2SO3, dilute
H2SO4 and dilute NaOH solution. 5 The observations are recorded in the table
below.
Procedure
(B) Conversion of Fe3+ ions to Fe2+ ions
(A) Conversion of Fe2+ ions to Fe3+ ions 1 About 2 cm3 of iron(III) chloride solution is added
1 About 2 cm3 of iron(II) sulphate solution is poured
to a test tube.
into a test tube. 2 Sodium sulphite (Na2SO3) solution is added to
2 About 2 cm3 of potassium manganate(VII) solution
iron(III) chloride solution, followed by dilute
is poured into another test tube, followed by about sulphuric acid. The mixture is shaken gently.
2 cm3 of dilute sulphuric acid. 3 Sodium hydroxide solution is then added slowly
3 Using a dropper, about 2 cm3 of the acidified to the reaction mixture until in excess.
potassium manganate(VII) solution is added slowly 4 The observations are recorded in the table
to the iron(II) sulphate solution. The mixture is below.

shaken gently.

Activity 3.2 3 Observations

Solution Test Observation

FeSO4(aq) (a) Fe2+ ion + acidified KMnO4 • The light green iron(II) sulphate solution changes to
yellow.

• The purple colour of acidified KMnO4 solution turns
colourless (decolourised).

(b) Add excess NaOH(aq) to (a) • A brown precipitate, insoluble in excess NaOH(aq) is
formed.

FeCl3(aq) (c) Fe3+ ion + Na2SO3(aq) • The colour of the solution changes from yellow to light
green.

(d) Add excess NaOH(aq) to (c) • A dirty green precipitate, insoluble in excess NaOH(aq)
is obtained.

Conclusion

1 Fe2+ ions are oxidised to Fe3+ ions by the acidified KMnO4 solution.
2 Fe3+ ions are reduced to Fe2+ ions by the sodium sulphite (Na2SO3) solution.

Displacement of Metals from Their Salt 3 Electropositive metals are strong reducing
Solutions agents. In contrast, the metallic ions of
electropositive metals are weak oxidising
1 The arrangement of metals according to their agents. Figure 3.6 shows that in the
electrochemical series,
SPM tendency to lose electrons to form positive (a) the strength of a metal as a reducing agent
’04/P1 increases on going up the electrochemical
’07/P1 ions is called the electrochemical series. series,
(b) the strength of the metallic ion as an
2 The higher the position of the metal in the oxidising agent increases on going down
the series.
electrochemical series,

(a) the more electropositive the metal,

(b) the more readily the metal donates electrons

to form positive ions,

(c) the more easily the metal will undergo

oxidation.

Oxidation and Reduction 400

• Tendency of a metal • Tendency of an ion
to ionise (by donating to accept electrons
electrons) increases. increases.

• Strength of a metal • Strength of an ion as
as a reducing agent an oxidising agent
increases. increases.

Figure 3.6 Electrochemical series

4 Consider the formation of sodium ions (Na+) position in the electrochemical series) will Experiment 3.2 3
from sodium metal (Na). displace a less electropositive metal (lower
position in the electrochemical series) from the
Na metal has a strong tendency salt solutions of the less electropositive metal.
6 Transfer of electrons occurs during displacement
to lose an electron to form sodium ion reactions.
N⎯a⎯(s⎯) →⎯⎯Na⎯+(⎯aq⎯) +→e– (a) The more electropositive metal donates
←Na⎯+ io⎯n⎯ha⎯s a⎯we⎯ak⎯te⎯ndency
electrons and acts as a reducing agent.
to accept an electron to form Na metal The metal undergoes oxidation and is
oxidised to its metal ions.
(a) Sodium metal is placed at a high position (b) The metal ion (from the less electropositive
in the electrochemical series. metal) in aqueous solution acts as an
oxidising agent. The metal ions undergo
(b) This means that sodium metal donates reduction and is reduced to its metal.
electrons very easily. As a result, sodium
is a strong reducing agent. A more electropositive metal is also a more reactive
metal. We can therefore state that a more reactive metal
(c) Conversely, sodium ions (Na+) have a will displace a less reactive metal from the solution of its
weak tendency to accept electrons. Since salts.
oxidising agents are electron acceptors,
sodium ions are weak oxidising agents.

5 A displacement reaction is a reaction in which
one element (metal or non-metal) displaces
another element (metal or non-metal) from its
salt solution. In the displacement reactions of
metals, the more electropositive metal (higher

3.2

To study the redox reaction in terms of displacement reaction of a metal from its salt
solution

Problem statement
How does redox reaction occur in a displacement reaction in which a metal is displaced from its salt
solution?

Hypothesis
(a) The metal that acts as a reducing agent will form metal ion.
(b) The metal ion that acts as an oxidising agent will be precipitated as metal.

Variables
(a) Manipulated variable : A pair of metals and salt solutions
(b) Responding variable : Precipitation of metal and colour changes in the solutions
(c) Constant variables : Volumes and concentrations of solutions containing the metal ions

Apparatus
Beakers

401 Oxidation and Reduction

Materials Procedure
Copper(II) sulphate solution and silver nitrate
solution, zinc plate and copper plate. 1 A strip of zinc plate and a strip of copper plate
are cleaned with sandpaper.
Figure 3.7 The displacement of a
metal from its salt solution 2 The zinc plate is then immersed in copper(II)
sulphate solution (beaker P) and the copper plate
is immersed in silver nitrate solution (beaker Q).

3 The mixture is left aside for half an hour.
4 The changes that take place on the zinc plate,

the copper plate, and in the copper(II) sulphate
solution and the silver nitrate solution are
recorded.

Observation

Reactants Observation Explanation

(a) Zinc plate • A section of the zinc plate dissolves. • Zinc displaces copper metal (brown
in copper(II) • Brown precipitate is deposited on the precipitate) from copper(II) sulphate
sulphate solution solution.
zinc plate.
• The blue solution fades until it • Copper metal is deposited on the zinc
plate.
becomes colourless.
• The blue colour fades as the
3 concentration of Cu2+ ions decreases.

(b) Copper plate • The copper plate dissolves. • Copper displaces silver metal (greyish-
in silver nitrate • A greyish-black precipitate is black) from the silver nitrate solution.
solution
deposited on the copper plate. • Silver metal is precipitated on the
• The colourless solution turns blue. copper plate.

• The colourless solution turns blue
because of the formation of Cu2+ ions.

Conclusion

During the displacement reaction, the more electropositive metal will act as a reducing agent.
It reduces the metal ion (oxidising agent) which is less electropositive to form metal. The hypothesis is accepted.

Displacement of Copper by Zinc from SPM The reaction can be represented by the
Copper(II) Sulphate Solution following half-equations:
’08/P2 (a) Zn(s) → Zn2+(aq) + 2e–

1 The following equation shows the reaction (oxidation – loss of electrons)
between copper(II) sulphate solution and zinc. Zinc acts as a reducing agent (electron donor)
(b) Cu2+(aq) + 2e– → Cu(s)
Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)
(reduction – gain of electrons)
Zinc is more electropositive than copper. Copper ion acts as an oxidising agent (electron
It displaces copper from its salt (that is, acceptor)
copper(II) sulphate). 3 When copper(II) ion is displaced, the
concentration of Cu2+ ions in the solution
2 A displacement reaction is a redox reaction. decreases. This causes the blue colour to fade.

oxidation (loss of electrons) Displacement of Silver by Copper from Silver
Nitrate Solution

1 The displacement reaction between copper and
Zn(s) + CuSO4(aq) ⎯⎯→ ZnSO4(aq) + Cu(s) silver nitrate solution is shown as follows.

reduction (gain of electrons)

Oxidation and Reduction 402

oxidation (loss of electrons) (a) Cu(s) → Cu2+(aq) + 2e–
(oxidation – loss of electrons)
Copper acts as a reducing agent (electron
Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s) donor).

(b) Ag+(aq) + e– → Ag(s)
reduction (gain of electrons) (reduction – gain of electrons)
Silver ion acts as an oxidising agent (electron
Copper is more electropositive than silver. acceptor).
It displaces silver from its salt (that is, silver
nitrate). 3 When copper dissolves in silver nitrate solution,
2 The reaction can be represented by the the formation of copper(II) ion causes the
following half-equations: solution to turn blue. The intensity of blue
colour increases as more copper is dissolved.

2 ’07

An experiment is carried out to determine the positions of the metals, Q, X, Y, Z in the electrochemical series.
The results of the experiment on displacement reactions are shown in the table below.

Metal Solution AgNO3 Pb(NO3)2 FeSO4 MgSO4 3
No change No change
Q Silver is displaced No change No change
X Silver is displaced Lead is displaced No change
Y Silver is displaced Lead is displaced Iron is displaced No change
Z

What is the correct position of the metals, in ascending Q is the least reactive. It has no reactions with Pb2+,
Fe2+ and Mg2+.
order, of the tendency of the metals to be oxidised? Y is the most reactive. It can displace three metals
from their solutions.
A Q, X, Y, Z C X, Y, Z, Q Z is more reactive than X. It can displace two metals
from their salt solutions.
B Q, X, Z, Y D Y, Z, X, Q
Answer B
Comments
The most reactive metal is the strongest reducing
agent, that is, it has the highest tendency to form
metal ions by losing electrons, that is to be oxidised.

Displacement of Halogens from Halide halogen will displace a less reactive halogen
Solutions from the solution of its halide ions.
3 Hence, chlorine displaces bromine from an
1 In general, the stronger the oxidising strength aqueous solution of bromide ions. It also
displaces iodine from an aqueous solution of
SPM of the halogen, the weaker the reducing iodide ions. Similarly, bromine displaces iodine
’04/P1 from an aqueous solution of iodide ions.
’05/P1 strength of the corresponding halide ion is.
Cl2(aq) + 2KBr(aq) → 2KCl(aq) + Br2(aq)
Thus, chlorine is a stronger oxidising agent Cl2(aq) + 2KI(aq) → 2KCl(aq) + I2(aq)
Br2(aq) + 2KI(aq) → 2KBr(aq) + I2(aq)
than iodine, but the iodide ion is a stronger
4 Conversely, bromine cannot displace chlorine
reducing agent than the chloride ion. Figure 3.8 from an aqueous solution of chloride ions and
iodine cannot displace bromine or chlorine
shows the trend in the reactivity and oxidising

strength of the halogens and the reducing

strength of the halide ions.

2 The reactivity of the halogens can be used to

predict whether the displacement reactions of

halogens can occur or not. A more reactive

403 Oxidation and Reduction

from an aqueous solution of bromide ions or 1,1,1-trichloroethane. The colours of halogens
chloride ions respectively. in 1,1,1-trichloroethane are shown in Table 3.10.

Br2(aq) + KCl(aq) → No reaction Table 3.10 The colours of halogens in 1,1,1-
I2(aq) + KCl(aq) → No reaction trichloroethane
I2(aq) + KBr(aq) → No reaction
Halogen Colour

Chlorine Colourless

5 The colours of halogens in water are shown in Bromine Brown
Table 3.9.
Iodine Purple

Table 3.9 The colour of halogens in water

Halogen Concentrated Dilute aqueous

aqueous solution solution

Chlorine Greenish-yellow Colourless The structural formula of 1,1,1-trichloroethane is:
H Cl
Bromine Brown Yellow ⎮ ⎮
H ⎯ C ⎯ C ⎯ Cl
Iodine Brown Yellow ⎮ ⎮
H Cl
Experiment 3.3 3 6 Halogens can be identified by adding 1,1,1- It is very volatile and is used as a solvent in paper correction
trichloroethane (CH3CCl3) to its aqueous fluid. It is produced as one of the organic products when
solution. Water and 1,1,1-trichloroethane are ethane reacts in chlorine in the presence of sunlight.
immiscible and two layers are formed. The
upper layer is water and the lower layer is A The tendency of electrons
being removed from
A Reactivity of halide ions to form
halogens increases halogens increases

A Oxidising strength of A The strength of halide
halogens increases ion as a reducing
agent increases

Figure 3.8

3.3

To study the displacement reactions between halogens and halide ions
Problem statement
How do redox reactions occur in displacement reactions between halogens and aqueous solutions of halide
ions?

Hypothesis
A more reactive halogen will displace a less reactive halogen from an aqueous solution of its halide ions.

Variables
(a) Manipulated variable : A pair of halogens and their halide ions
(b) Responding variable : Changes in colour in 1,1,1-trichloroethane, CH3CCl3
(c) Constant variable : Volume of reaction mixture

Apparatus
Test tubes

Oxidation and Reduction 404

Materials 2 2 cm3 of chlorine water is added to the potassium
1,1,1-trichloroethane, potassium bromide, KBr(aq), bromide solution. The mixture is shaken gently.
potassium chloride, KCl(aq), potassium iodide, KI(aq),
chlorine water, liquid bromine and iodine solution. 3 2 cm3 of 1,1,1-trichloroethane (CH3CCl3) is then
added to the mixture obtained in step 2. The
Figure 3.9 Displacement of bromine from mixture is then shaken vigorously.
potassium bromide solution
4 The test tube is allowed to stand for a few minutes
Procedure and the colour of the 1,1,1-trichloroethane layer
1 A test tube is filled with 2 cm3 of potassium is recorded.

bromide, KBr solution. 5 Steps 1 to 4 are repeated using the following
mixtures.
(a) Chlorine water and potassium iodide, KI
solution
(b) Liquid bromine and potassium chloride, KCl
solution
(c) Liquid bromine and potassium iodide solution
(d) Iodine solution and potassium bromide
solution
(e) Iodine solution and potassium chloride
solution

Observation 3

Mixture Colour of CH3CCl3 Halogen in CH3CCl3 Has displacement
layer layer reaction occurred?

Cl2(aq) + KBr(aq) Brown Bromine Yes

Cl2(aq) + KI(aq) Purple Iodine Yes
Br2(l) + KCl(aq) Brown ψ Bromine No

Br2(l) + KI(aq) Purple Iodine Yes

I2(aq) + KBr(aq) Purple *Iodine No
I2(aq) + KCl(aq) Purple *Iodine No

ψThe bromine present in the CH3CCl3 layer is due to the bromine added.
*The iodine present in the CH3CCl3 layer is due to the iodine added.

Discussion oxidation

1 When chlorine water is added to potassium Cl2(aq) + 2KBr(aq) → 2KCl(aq) + Br2(l)
bromide solution, the colour of the solution
changes from colourless to brown because reduction
chlorine displaces bromine from an aqueous
solution of bromide ions. 4 The half-equations for the reactions are as
follows:
Cl2(aq) + 2KBr(aq) → 2KCl(aq) + Br2(l) (a) 2Br–(aq) → Br2(l) + 2e–

2 If 1,1,1-trichloroethane is added to the reaction (oxidation – loss of electrons)
mixture and shaken, two liquid layers are formed.
The lower organic layer (1,1,1-trichloroethane) (b) Cl2 (aq) + 2e– → 2Cl–(aq)
has a brown colour and shows the presence of (reduction – gain of electrons)
bromine. This means that the bromide ions have
been oxidised to bromine. Chlorine accepts electrons and acts as an
oxidising agent. Bromide ion (Br–) loses
3 Displacement reactions can also be considered electrons and acts as a reducing agent.
in terms of oxidation and reduction. Chlorine 5 Other displacement reactions that occur in this
oxidises bromide ion (Br–) to bromine and experiment are
chlorine is itself reduced to chloride ion (Cl–).

405 Oxidation and Reduction

oxidation (b) I2(aq) + KBr(aq) → No reaction
I2(aq) + KCl(aq) → No reaction
Cl2(aq) + 2KI(aq) → 2KCl(aq) + I2(aq) This is because iodine is less reactive than
bromine and chlorine.
reduction
oxidation Conclusion

Br2(l) + 2KI(aq) → 2KBr(aq) + I2(aq) 1 Chlorine displaces bromine from potassium
bromide solution and iodine from potassium
reduction iodide solution. Bromine displaces iodine from
iodide solution but does not displace chlorine
from chloride solution. Iodine does not displace
6 The following displacement reactions do not occur. chlorine from chloride solution or bromine from
(a) Br2(l) + KCl(aq) → No reaction bromide solution.
This is because bromine is less reactive than
chlorine. 2 The experimental results prove that a more
reactive halogen can displace a less reactive
halogen from its halide solution. The hypothesis
is accepted.

3 Redox Reactions by the Transfer of
Electrons at a Distance

In Figure 3.9 (Experiment 3.3), the aqueous layer 1 If a solution containing an oxidising agent is
contains KCl but the organic layer (CH3CCl3) does not
contain KCl. This is because KCl is an ionic compound SPM separated from a solution containing a reducing
which is soluble in water but not in organic solvent. ’06/P2,
’11/P2 agent by an electrolyte, the redox reaction can

still occur by transfer of electrons at a distance.

3 ’04 The apparatus set-up is shown in Figure 3.10.

Which of the following equations represent redox

reactions?

I Cu(s) + Cl2(g) → CuCl2(s)
II Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s)
III Cu(OH)2(aq) + 2HCl(aq) →
CuCl2(aq) + 2H2O(l)
IV Cu(NO3)2(aq) + K2CO3(aq) →
CuCO3(s) + 2KNO3(aq)
A I and II only

B III and IV only

C I, II, III only­ Figure 3.10 The apparatus set-up in which the
transfer of electrons occurs at a distance
D I, II, III and IV
2 In Figure 3.10, the electrons are transferred
Comments by the connecting wire in the external circuit,
Combustion of metals in chlorine or oxygen are from a reducing agent to an oxidising agent.
redox reactions (I is correct).
Displacement reactions are redox reactions (II is 3 The electrode placed in a solution containing a
correct). reducing agent acts as the negative terminal.
Neutralisation reactions are not redox reactions (III The reducing agent undergoes oxidation with
is incorrect). the loss of electrons.
Double decomposition reactions are not redox
reactions (IV is incorrect). 4 The electrons produced will flow through the
connecting wire to the electrode placed in a
Answer A solution of oxidising agent. This electrode
acts as the positive terminal.

Oxidation and Reduction 406

5 Sulphuric acid acts as the salt bridge. The 2 The deflection of the galvanometer needle
functions of a salt bridge are shows that electrons flow in the external
(a) to separate the oxidising agent from the circuit from the carbon electrode immersed
reducing agent, in the potassium iodide solution (negative
(b) to complete the electric circuit so that ions electrode) to the carbon electrode immersed
can move through it. in the acidified potassium manganate(VII)
solution (positive electrode).
6 Besides dilute sulphuric acid, the following
strong electrolytes can also be used as salt 3 Changes at the negative electrode
bridges: (a) The colourless layer of potassium iodide
(a) Sodium or potassium chloride solution solution slowly changes to yellow.
(b) Sodium or potassium nitrate solution (b) The oxidation reaction occurs at the
negative electrode.

Redox reactions can occur under two conditions as 2I–(aq) → I2(aq) + 2e–
shown below.
(c) The electrons released during oxidation
Conditions Experiment then flow through the connecting wire
(external circuit) to the positive electrode
(a) When the oxidising Mixing the oxidising and and are accepted by the MnO4– ions. 3
agent and the reducing agents in a Hence, MnO4– ion acts as the oxidising
reducing agent are test tube. agent.
in contact.
4 Changes at the positive electrode
(b) When the oxidising Transfer of electrons at (a) The purple layer of potassium man­
agent and the a distance in a redox ganate(VII) slowly becomes colourl­ ess
reducing agent are cell. (decolourises).
not in contact. (b) The reduction reaction occurs at the
positive electrode.

Reaction between Potassium Iodide and MnO4–(aq) + 8H+(aq) + 5e– →
Acidified Potassium Manganate(VII) by Transfer Mn2+(aq) + 4H2O(l)
of Electrons at a Distance
(c) Overall reaction
1 Figure 3.11 shows the apparatus set-up using The ionic equation for the redox reaction
acidified potassium manganate(VII) as an
oxidising agent and potassium iodide as a is shown below.
reducing agent.
oxidation


2MnO4– + 10I– + 16H+ → 2Mn2+ + 5I2 + 8H2O



reduction

reducing
agent

oxidising agent

Figure 3.11 Apparatus set-up consisting of KI and KMnO4

407 Oxidation and Reduction

To study the redox reactions by the transfer of electrons at a
distance

Apparatus U-tube, dropper, carbon electrodes, A few drops of starch solution are added to the
galvanometer, retort stand with clamp potassium iodide solution.
and connecting wire with crocodile 9 Steps 1 to 6 are again repeated using chlorine
clips. and potassium bromide solution.

Materials Liquid bromine, potassium Results
dichromate(VI) solution, chlorine,
iron(II) sulphate solution, potassium 1 Reaction of iron(II) sulphate solution with liquid
iodide solution, potassium bromide bromine
solution, sulphuric acid, sodium
hydroxide solution and starch Experiment Observation
solution.
(a) Colour change in The colour of the
Activity 3.3 3 Procedure iron(II) sulphate solution changes from
solution light green to yellow.
1 The U-tube is half-filled with dilute sulphuric acid.
2 Using a dropper, iron(II) sulphate solution is added (b) Colour change in The colour of solution
liquid bromine changes from brown
to the left arm of the U-tube until a 2 cm high to colourless.
layer is formed on top of the dilute sulphuric acid. (c) Reaction with
3 Using the same method, liquid bromine is added sodium hydroxide A brown precipitate
to the right arm of the U-tube. solution insoluble in excess
4 Two carbon electrodes are then immersed into NaOH(aq) is
the upper layer of the U-tube. The electrodes are produced.
connected to the galvanometer by electric wire
(Figure 3.12). 2 Reaction of potassium iodide solution with
acidified potassium dichromate(VI) solution

Experiment Observation

(a) Colour change in Colour changes
potassium iodide from colourless to
solution brown

(b) Colour change Colour changes

reducing agent oxidising agent in potassium from orange to

Figure 3.12 Oxidation of iron(II) sulphate by liquid dichromate(VI) solution green
bromine
(c) Reaction with starch Blue-black colour
5 The direction of the deflection of the solution is produced.
galvanometer needle is observed to determine
the direction of the flow of electrons. 3 Reaction of potassium bromide solution with
chlorine
6 The apparatus is left aside for a few minutes and
the changes that occur at the carbon electrodes Experiment Observation
are recorded. Colour changes from
(a) Colour change in colourless to brown
7 Using a clean dropper, a small quantity of the potassium bromide
iron(II) sulphate solution is removed and placed solution Colour changes from
in a test tube. Sodium hydroxide solution is then greenish-yellow to
added to the iron(II) sulphate solution. (b) Colour change in colourless
chlorine
8 Steps 1 to 6 are repeated using acidified potassium
dichromate(VI) solution and potassium iodide
solution. A small quantity of the potassium iodide
solution is removed and placed in a test tube.

Oxidation and Reduction 408

Conclusion oxidation

(a) Bromine oxidises iron(II) sulphate The oxidation number of
(b) Acidified potassium dichromate(VI) oxidises –1 0 iodine increases from –1
to 0. Iodide ion (I–) is a
potassium iodide 2I–(aq) → I2(aq) + 2e– reducing agent.
(c) Chlorine oxidises potassium bromide by transfer
colourless brown
of electrons at a distance
(b) At the positive terminal/pole/electrode
Discussion (i) Reduction of dichromate(VI) ions to

1 Reaction of iron(II) sulphate solution and chromium(III) ions occurs.
bromine
(a) At the negative terminal/pole/electrode reduction

(i) Oxidation of Fe2+ ions to Fe3+ ions occurs
and electrons are released.
+6 +3

Cr2O72–(aq) + 14H+(aq) + 6e– → 2Cr3+(aq) + 7H2O(l)
orange green
+2 +3 Oxidation number of
iron increases from (ii) The oxidation number of chromium
Fe2+ → Fe3+ + e– +2 to +3. Fe2+ ion is a
reducing agent. decreases from +6 to +3.
green brown (c) Overall reaction

(ii) The carbon electrode dipped in iron(II) Cr2O72–(aq) + 14H+(aq) + 6I–(aq) → 3
sulphate solution is negatively-charged 2Cr3+(aq) + 7H2O(l) + 3I2(aq)
because electrons released during the
reaction gather at the electrode. 3 Reaction of potassium bromide solution and

(iii) The electrons flow through the electric chlorine
wire (external circuit) to the positive
terminal and are accepted by the (a) At the negative terminal/pole/electrode
bromine molecules.
At the negative electrode, bromide ions (Br–)
(iv) The colour of the solution changes from
pale green to yellow or brown when Fe2+ donate electrons and are themselves oxidised
ions are oxidised to Fe3+ ions.
to liquid bbrroomminineeinc(Brera2)s.esTfhroem oxidation
(b) At the positive terminal/pole/electrode number of –1 to 0.
(i) Reduction of bromine molecules to bromide
oxidation
ions occurs. Electrons are accepted by the
bromine molecules.
–1 0
–1 The oxidation number of
0 bromine decreases from 2Br–(aq) → Br2(l) + 2e–
0 to –1. Bromine is an colourless brown
oxidising agent.
Br2(aq) + 2e– → 2Br–(aq) (b) At the positive terminal/pole/electrode
brown colourless At the positive electrode, chlorine molecules,
rCeld2(uacqe)dactocecphtloerliedcetrioonnss,aCndl–(aarqe).themselves
(ii) The colour of the solution changes
from brown to colourless when bromine Cl2(g) + 2e– → 2Cl–(aq)
molecules are reduced to bromide ions. greenish-yellow colourless

(c) Overall reaction (c) Overall reaction
The ionic equation for the reaction between The overall reaction can be represented by
bromine and iron(II) sulphate is the following ionic equation:

Br2(aq) + 2Fe2+(aq) → 2Fe3+(aq) + 2Br–(aq) oxidation

2 Reaction of potassium iodide and acidified
potassium dichromate(VI) Cl2(aq) + 2Br–(aq) → 2Cl–(aq) + Br2(l)
(a) At the negative terminal/pole/electrode
Oxidation of iodide ions (I–) to iodine reduction
occurs.

409 Oxidation and Reduction

Oxidation reaction Reduction reaction
• Addition of oxygen • Removal of oxygen
• Removal of hydrogen • Addition of hydrogen
• Increase in oxidation number • Decrease in oxidation number
• Loss of electrons • Gain of electrons

occur simultaneously

Redox reactions Redox cell
(a) Combustion of metals in oxygen or chlorine
(b) Heating of metallic oxide with carbon (transfer of electrons at a distance)
(c) Changes of Fe2+ ions to Fe3+ ions and Fe3+ ions • The electrode immersed in a reducing agent is the

to Fe2+ ions negative electrode
(d) Displacement (i) of metal from its salt, • The electrode immersed in an oxidising agent is

(ii) halogen from its halide solution the positive electrode
(e) Transfer of electrons at a distance • Electrons flow from negative electrode to positive

electrode

3 Oxidising agent Reducing agent Rules for determining oxidation numbers
• Brings about oxidation • Brings about reduction • The oxidation number of an uncombined

in another substance in another substance element is zero.
• It is an electron • It is an electron donor • The sum of the oxidation numbers of all the
• Examples: FeSO4, KBr,
acceptor atoms in a polyatomic ion is the charge on the
• Examples: Cl2, Br2, KI, C, CO, H2, metals ion.
• The sum of the oxidation numbers of all the
KMnO4, K2Cr2O7 atoms in a compound is zero.

4 ’03

Which of the following reagents can be used to reduces Fe3+ to Fe2+ by displacement reaction. (I is
convert Fe3+ ions in solution to Fe2+ ions? correct)
I Magnesium powder
II Potassium iodide solution reduction
III Potassium hexacyanoferrate(II) solution
IV Acidified potassium manganate(VII) solution 0 +3
+2 +2
A I and II only
B II and III only Mg(s) + 2Fe3+(aq) → 2Fe2+(aq) + Mg2+(aq)
C I, II and III only
D I, II, III and IV oxidation

Comments Potassium iodide is a reducing agent. (II is correct)
The conversion of Fe3+ ions to Fe2+ ions is a reduction
reaction. reduction
Hence, reducing agents are required for the conversion.
Potassium hexacyanoferrate(II) is not a reducing
agent. (III is incorrect)
Acidified potassium manganate(VII) is an oxidising +3 +2
agent. (IV is incorrect)
Mg is a more electropositive metal than iron. It 2KI(aq) + 2Fe3+(aq) → 2Fe2+(aq) + I2(s) + 2K+(aq)


oxidation
Answer A

Oxidation and Reduction 410

3.1 (b) Explain your answer in terms of (i) transfer of
electrons, (ii) change in oxidation number.
1 Consider the following reactions and state whether
the chemicals underlined have been oxidised or 6 Zinc reacts with iron(III) chloride as represented by
reduced? the following ionic equation
(a) 2Fe + 3Cl2 → 2FeCl3
(b) Zn + CuSO4 → ZnSO4 + Cu Zn(s) + 2Fe3+(aq) → 2Fe2+(aq) + Zn2+(aq)
(c) PbO + CO → Pb + CO2
(d) 4NH3 + 5O2 → 4NO + 6H2O (a) Describe what you would see when zinc powder
is added to iron(III) chloride solution.
2 Calcium reacts with chlorine to form calcium
chloride. (b) Write the half-equations for the chemical changes
that take place.
Ca(s) + Cl2(g) → CaCl2(s)
(c) Is zinc being oxidised or reduced?
(a) Identify the element that has been (i) oxidised Explain your answer in terms of transfer of
and (ii) reduced.
electrons.
(b) Identify the oxidising and reducing agents in this
reaction. 7 (a) Identify the redox reactions from the equations

3 (a) State the oxidation numbers of the underlined given below.
elements in each of the following compounds:
((((ii(vviiiii))))) 22M4CFNFl2geea(+COC+l3lH23++K)O2BPS+2rbn→→(CHNl22O2S2→OK3F)eC422l2O→→+F3eCMBPlrb2g2CS+lO2 S4+n+C2l2N4HaN2OO3 3
(i) CaCrO4
(ii) HNO2
(iii) NaClO3
(iv) Cr2(SO4)3
(v) FeCl3.6H2O
(vi) K2MnO4 (b) Write the half-equations for these redox

(b) Give the oxidation numbers of the underlined reactions.
elements in the following ions:
(c) Identify the oxidising agents and reducing agents
((iii)) SCOuC42l–4 3– ((iivii)) CNlOO3––
4 (a) Write the molecular formulae of the following in the following reactions. Explain your answer in

compounds. terms of oxidation number.
(i) Copper(I) sulphate
(ii) Manganese(II) chloride (i) SnCl2 + 2FeCl3 → 2FeCl2 + SnCl4
(iii) Nickel(II) sulphate (ii) Zn + Pb(NO3)2 → Zn(NO3)2 + Pb
(iii) Cl2 + 2KBr → 2KCl + Br2
(b) Write the names of the compounds with the (iv) 2KI + H2O2 + H2SO4 → I2 + K2SO4 + 2H2O
following molecular formulae:
8 The figure shows an experiment on the transfer of
(i) CrCl3
(ii) CoO electrons at a distance.
(iii) Fe2(SO4)3
’06
(c) The formulae of two compounds are shown
below. (a) Name the oxidising agent and reducing agent in
the experiment.
’06
(b) Write the half-equations for the reactions that
Fe2O3 and Al2O3 occur at the negative and positive terminals.

(i) State the oxidation numbers for iron and (c) Based on your answer in (b), describe the
aluminium in the compounds above. oxidation and reduction processes in terms of
electron transfer at the negative and positive
(ii) Name both the compounds using IUPAC terminals.
nomenclature system.
(d) State the changes that occur at the negative and
(iii) Explain the difference between the names positive terminals after 10 minutes.
of the two compounds.

5 The chemical equation for the combustion of
magnesium in air is

2Mg(s) + O2(g) → 2MgO(s)

(a) Is magnesium oxidised or reduced?

411 Oxidation and Reduction

3.2 Rusting as a Redox The golden mask of King
Reaction Tutankhamun was buried in his
tomb for more than 3000 years. It
Conditions for the Rusting of Iron is still uncorroded. This is because
gold is an unreactive metal and
1 Rusting is a redox reaction between iron, never corrodes.
oxygen and water to form a brown substance
called rust. Rust is hydrated iron(III) oxide, 2 Most metals corrode readily in air. When
Fe2O3.xH2O. The composition of water in rust corrosion occurs, the metal surface loses its
is not constant. luster (shine) and becomes dull. This is because
metals react slowly with oxygen in the air to
4Fe(s) + 3O2(g) + 2xH2O(l) → 2Fe2O3.xH2O(s) from metal oxides on the metal surfaces.

2 Two conditions required for rusting are 3 Metals that are more electropositive (reactive)
(a) the presence of air (oxygen) and will corrode more readily.
(b) the presence of water.
3 K very easy to
Corrosion of Metals Na corrode
electropositivity of metal decreases Ca very difficult
1 The corrosion of metals is a redox reaction tendency for corrosion decreases Mg to corrode
in which a metal is oxidised spontaneously at Al
room temperature with the release of electrons Zn
to form the metal ions. Fe
Sn
M(s) → Mn+(aq) + ne– Pb
metal metal ion Cu
Hg
Thus, metal M is said to have corroded and Ag
the process is known as corrosion of metal. Au

Group 1 metals in the Periodic Table (for reactivity series
example, sodium and potassium) are very
reactive and must be kept in paraffin oil to 4 The higher the position of the metal in the
protect them from oxidation by air and water. reactivity series, the easier it is for the metal to
When sodium or potassium is exposed to the donate its electrons and be corroded.
atmosphere, the metals corrode rapidly.
Aluminium corrodes rapidly in air to form a
thin layer of aluminium oxide on its surface.
This oxide layer is hard, impermeable (does
not allow liquid or gas to pass through)
and difficult to crack. Thus, the thin layer
of aluminium oxide protects the aluminium
below it from further corrosion.

Corrosion of metals and
positions of metals in the

reactivity series

Less reactive metals such as chromium, zinc Unreactive metals (such as gold and
and nickel also form hard metal oxides platinum) do not corrode because they are
that are impermeable to water and air and resistant to oxidation.
resistant to cracks. These metal oxide layers
can prevent and hence protect the metals
from further corrosion.

Oxidation and Reduction 412

5 The layer of aluminium oxide can be made (b) The side of the drop of water
thicker by electrolysis. This process is called This is the area in the metal surface where
anodising. In industry, anodising is used to
protect aluminium from rusting. Anodising is it is rich in oxygen. This area will act as
an electrolytic process using aluminium as the the positive terminal (cathode).
anode (positive electrode). During electrolysis 3 The following stages are involved during the
(anodising), a layer of aluminium oxide is rusting of iron.
deposited on the surface of aluminium.
(a) In the centre of the water droplet (anode)
• Window and door frames made from anodised Iron rusts via the oxidation process to form 3
aluminium do not require painting because the iron(II) ions.
aluminium oxide layer prevents them from attack by Fe(s) → Fe2+(aq) + 2e– … oxidation
air and water. The electrons flow to the edge of the water
droplet through the iron surface.
• Iron is less reactive than aluminium and window
frames made from iron must be painted frequently. (b) At the edge of water droplet (cathode)
This is because the hydrated iron(III) oxide (rust) Oxygen accepts electrons from the oxidation
formed during corrosion is permeable to air and of iron and is reduced to hydroxide ions.
water, and can crack easily. Thus, rusting will continue O2(g) + 2H2O(l) + 4e– → 4OH–(aq)
below the rusted surface, until all the metal is ‘eaten … reduction
up’ by rust.
(c) Formation of Fe(OH)2
Rusting in Terms of Oxidation and The Fe2+ and OH– ions in the water droplet
Reduction combine to form iron(II) hydroxide.
Fe2+(aq) + 2OH–(aq) → Fe(OH)2(s)
1 Rusting of iron is an electrochemical process

SPM that occurs spontaneously. When iron is in
’05/P1
’06/P1 contact with water, a simple chemical cell
’07/P2 is formed. Figure 3.13 shows the reactions

involved in the formation of rust.

(d) Formation of rust
The iron(II) hydroxide produced is oxidised
by oxygen to form iron(III) hydroxide, which
then decomposes to hydrated iron(III) oxide.


4Fe(OH)2 + 2H2O + O2 ⎯ox⎯id⎯ati⎯on→ 4Fe(OH)3

Fe(OH)3 ⎯de⎯co⎯mp⎯os⎯iti⎯on→ Fe2O3.xH2O
rust

Figure 3.13 Rusting of iron 4 (a) The equations for the redox reactions are
shown below.
2 Consider a drop of water on the metal (iron)
surface. Anode : 2Fe(s) → 2Fe2+(aq) + 4e– … oxidation
(a) The centre of a drop of water Cathode : O2(g) + 2H2O(l) + 4e– → 4OH–(aq)
This is the area in the metal surface where
there is a lack of oxygen. This area will … reduction
act as the negative terminal (anode).
2Fe(s) + O2(g) + 2H2O(l) → 2Fe(OH)2(s)

413 Oxidation and Reduction

(b) The overall equation for the rusting of iron Prevention of Rusting of Iron
is
1 The rusting of iron can be prevented if iron
4Fe(s) + 3O2(g) + 2xH2O(l) → 2Fe2O3.xH2O(s) is in contact with a more electropositive metal.
Conversely, the rate of rusting of iron is
5 The rate of rusting of iron is increased if a increased if the iron is in contact with a less
strong electrolyte (such as salt and acid) is electropositive metal.
present. Thus, the rusting of iron occurs more
rapidly in areas near the sea or in industrial Iron does not corrode if it is Iron corrodes rapidly
areas. This is because sea air contains salts such in contact with Zn, Al or Mg if it is in contact with
as sodium chloride and magnesium chloride. (Refer to Experiment 3.4) Sn, Pb or Cu
In the industrial area, the air is polluted by
acidic gases such as sulphur dioxide and K Na Mg Al Zn Fe Sn Pb Cu
nitrogen dioxide. These substances increase the
electrical conductivity of water, thus, making Electropositivity decreases
water a better electrolyte.
Experiment 3.4 3 2 When two metals are in contact, the greater the
6 The rate of rusting is also increased if iron is in difference in electropositivity between these
contact with a metal less electropositive than two metals, the faster the more electropositive
iron. For example, if iron is in contact with metal will rust. For example, a piece of iron
copper (a metal less electropositive than iron), joined to copper will corrode more rapidly
the rate of rusting is increased. This process is than a piece of iron joined to tin.
known as electrochemical corrosion.
3 Experiment 3.4 shows the experimental set-up
7 Besides the corrosion of iron and steel, corrosion for the study of the effect of other metals on
of other metals can also occur. The main causes the rate of rusting of iron.
of corrosion of metals are attack by chemicals,
such as acids, damp air or electrochemical
corrosion.

3.4

To investigate the effect of other metals with different electropositivity on the SPM
rusting of iron ’04/P2

Problem statement (a) Manipulated variable : Different metals used to
What is the effect of other metals with different wrap around iron nails
electropositivity on the rusting of iron?
(b) Responding variable : Colour change in the
Hypothesis gelatin solution

(c) Constant variable : Iron nails

(a) A metal more electropositive than iron will Apparatus Test tubes
protect iron from rusting. Materials
Iron nails, magnesium, zinc, tin
(b) A metal less electropositive than iron will and copper foils, gelatin, potassium
increase the rate of rusting. hexacyanoferrate(III), phenolphthalein
indicator and sandpaper.
Variables

Figure 3.14 Effect of contact with other Procedure
metals on the rusting of iron
1 Five pieces of iron nails are cleaned using sandpaper.
2 The first clean iron nail is placed in test tube A.
3 Strips of magnesium (Mg), zinc (Zn), tin (Sn)

and copper (Cu) foils are cleaned with sandpaper.
4 Each iron nail is wrapped with a different metal

foil and placed in test tubes B, C, D and E
respectively.

Oxidation and Reduction 414

5 A solution of gelatin in hot water is prepared. A few drops of potassium hexacyanoferrate(III) solution,

6 K3Fe(CN)6 and phenolphthalein indicator are added to the hot gelatin solution.
The mixture is stirred and then poured into each of the test tubes (Figure 3.14).

7 The test tubes are set aside for three days and then examined. The observations are recorded in the table

below.

Results

Test tube ABCDE

Observation

Metal Fe only Fe + Mg Fe + Zn Fe + Sn Fe + Cu

Intensity of blue colour Low None None High High

Intensity of pink colour None High High Low Low

Gas bubbles None Plenty Plenty Few Few

Discussion (c) At the anode: Magnesium foil and zinc foil 3
The oxidation of magnesium to Mg2+ ions
1 (a) Potassium hexacyanoferrate(III) is used to and the oxidation of zinc to Zn2+ ions occur.
detect the Fe2+ ions. Potassium hexacyano- Electrons are released.
ferrate(III) produces a dark blue colour in the
presence of Fe2+ ions. Mg(s) ⎯ox⎯id⎯ati⎯on→ Mg2+(aq) + 2e–

(b) Phenolphthalein is used to detect the OH– Zn(s) ⎯ox⎯id⎯ati⎯on→ Zn2+(aq) + 2e–
ions. Phenolphthalein produces a pink colour
in the presence of OH– ions. Thus, magnesium and zinc are corroded instead
of iron. The electrons released then flow to the
(c) The bubbles of gas produced are hydrogen iron nail and prevent it from forming Fe2+ ions,
gas. that is, prevent the rusting of iron.
(d) At the cathode: Iron nails
2 Reactions in test tube A (Fe only) Water molecules dissociate to form hydrogen
(a) Test tube A is used as a control to study the ions (H+) and hydroxide ions (OH–).
effect of other metals on the rusting of iron.
(b) In the presence of water and oxygen, rusting H2O(l) H+(aq) + OH–(aq)
of iron occurs to produce iron(II) ions (Fe2+)
and hydroxide ions (OH–).

Fe(s) ⎯ox⎯id⎯ati⎯on→ Fe2+(aq) + 2e– Hydrogen ions accept electrons and are
reduced to hydrogen gas.
O2(g) + 2H2O(l) + 4e– ⎯re⎯du⎯cti⎯on→ 4OH–(aq)
2H+(aq) + 2e– ⎯re⎯du⎯ct⎯ion→ H2(g)
(c) Fe2+ ions react with potassium hexacyano­
(e) When hydrogen ions are discharged to form
ferrate(III) to produce a deep blue precipitate. hydrogen gas, the concentration of hydroxide
ions in water increases. Consequently, the
(d) Phenolphthalein does not produce a pink area around the iron nail becomes alkaline
and causes the colour of phenolphthalein to
colour because the OH– ions produced react change from colourless to pink.

3 with Fe2+ ions to form FBe((OFHe )+2. Mg) and C 4 Reactions in test tubes D (Fe + Sn) and E
Reactions in test tubes (Fe + Cu)
(a) Deep blue colour appears. This implies that
(Fe + Zn) Fe2+ ions are produced, that is, rusting of iron
nails has occurred. The high intensity of the
(a) Deep blue colour does not appear in test blue colour shows that the rusting of iron nail
is speeded up.
tubes B and C. This implies that Fe2+ ions are

not produced, that is, the iron does not rust.

(b) Magnesium and zinc are more electropositive

than iron. Thus, magnesium and zinc act as

the negative terminal (electrode) and iron

acts as the positive terminal (electrode).

415 Oxidation and Reduction

(b) Iron is more electropositive than tin and Consequently, tin foil and copper foil do not
copper. Hence, iron has a greater tendency corrode.
to lose electrons and acts as the negative (e) When hydrogen ions are discharged, the
terminal (anode). Tin and copper, on the other concentration of hydroxide ions (OH–) in
hand, act as the positive terminal (cathode). water increases. However, most of the OH–
ions produced will combine with Fe2+ ions to
(c) At the anode: Iron nail form rust. Consequently, the area around tin
Oxidation of iron to Fe2+ ions occurs and foil or copper foil is slightly alkaline and a
electrons are released. slight pink colour is observed.

Fe(s) ⎯o⎯xid⎯ati⎯on→ Fe2+(aq) + 2e– Conclusion

(d) At the cathode: Tin foil or copper foil 1 The rusting of iron can be prevented if iron is in
Electrons released by the iron nail will flow contact with more electropositive metals such as
to the tin or copper foil which acts as the magnesium or zinc.
positive terminal. Hydrogen ions from water
will then accept these electrons and are 2 The rusting of iron is speeded up if iron is in
themselves reduced to hydrogen gas. contact with less electropositive metals such as
tin or copper. The hypothesis is accepted.
2H+(aq) + 2e– ⎯re⎯du⎯cti⎯on→ H2(g)

3 (ii) Another reason is that tin plating
makes the articles shiny and more
Zinc is not used to coat food cans although zinc can attractive in appearance.
prevent the rusting of iron. This is because zinc is a
poisonous substance. Zinc is more reactive than tin (iii) However, there is one disadvantage
and is susceptible to attack by acids such as fruit juices. in tin plating because tin is less
Food cans are usually electroplated with tin because electro­positive than iron. If the tin
tin is non-poisonous. Furthermore, tin is resistant to coating is broken, the iron beneath it
oxidation by oxygen and water. Tin can prevent iron will rust even more rapidly because
from rusting as long as the tin surface is not cracked. iron is more electropositive than tin.

Methods Used for the Prevention of Rusting SPM (b) Plating iron with chromium
’11/P1 (i) Chromium is a metal that is resistant

1 Using a protective layer to rusting. When chromium is exposed
(a) Rusting of iron and steel can be prevented to water and air, an impermeable, non-
by keeping them away from air (oxygen) brittle oxide layer is formed.
and water. (ii) The oxide layer acts as a protective
(b) A layer of paint, oil, grease or plastic coating layer to prevent iron beneath it from
protects the iron surface from coming into coming into contact with water and
contact with air and water. Without the air in the atmosphere.
presence of both air and water, rusting of (iii) Car bumpers, bicycle handles and
iron cannot occur. pipes are chromium-plated to prevent
(c) Oil and grease are usually used for movable corrosion.
machine parts. Plastics are usually used 3 Using more electropositive (reactive) metals
for household articles such as flower pot SPM (a) Galvanising is the coating of iron or steel
stands, coat hangers and so on. Paints are ’10/P1 with zinc for protection from corrosion.
used for bigger objects such as car bodies, Galvanising is carried out by dipping the
fences, gates and bridges. iron object into molten zinc or by electro­
plating. Zinc-coated iron is known as
2 Using less electropositive metals galvanised iron.
(a) Plating iron with tin (i) Even if the layer of zinc is scratched
or broken, the iron beneath it does not
(i) Tin is a not an electropositive metal rust. This is because zinc is more electro­
and is resistant to oxidation by water positive than iron and will corrode first.
and air. Hence, iron plated with tin is
used in making food cans. Zn(s) → Zn2+(aq) + 2e– ... oxidation

Oxidation and Reduction 416

(ii) This method of rust prevention is called Magnesium is more electropositive than 3
cathodic protection. The metal zinc iron and will be corroded in preference
is known as sacrificial metal because to iron.
zinc is sacrificed in the protection of
iron from rusting. Figure 3.16 Using the sacrificial metal
(magnesium blocks) to
(iii) Galvanised iron are used for making protect the underground pipes
zinc roof and gutter.
4 Using alloys
(b) Rusting in ships is prevented by fixing bars (a) The best known rust-resistant alloy of iron
of zinc to the part of the ship submerged is stainless steel. Stainless steel contains
in water (Figure 3.15). Zinc is oxidised 10–20% nickel and 10–25% chromium.
(corroded) in preference to iron. When exposed to the air, a hard layer of
chromium(III) oxide is formed on the
Figure 3.15 Using a sacrificial metal, zinc, to prevent surface of iron and prevents the iron from
the rusting of ships rusting.
(b) Stainless steel is used to make surgical
(c) Rusting in underground iron pipes is instruments and kitchen wares such as
prevented by having blocks of magnesium knives, forks and spoons.
attached to the iron pipes (Figure 3.16).

Conditions for rusting of iron Corrosion of metals and rusting of iron
• Presence of air (oxygen) • Corrosion of metals is a redox reaction in which a metal is oxidised
• Presence of water
spontaneously by losing electrons to form metal ions.
• When iron corrodes, the process is called rusting.

Explaining the rusting of iron Methods for rust prevention
• Corrosion of metals and rusting of iron are redox • Using a protective layer such as

reactions. paint, oil, grease, or plastic covering
• At the anode (negative electrode), iron is corroded. • Coating/plating iron with tin
• Coating/plating iron with chromium
Fe(s) → Fe2+(aq) + 2e– … oxidation • Using sacrificial metals
• Using alloys
• At the cathode (positive electrode), OH– ions are
produced. Oxidation and Reduction

O2 + 2H2O + 4e– → 4OH– … reduction

• Fe2+ combines with OtoHF–et2oO3fo.xrHm2OFe((rOusHt)).2.
• Fe(OH)2 is oxidised

417

3.2 4 Oxygen used for burning the metals is supplied

1 State the (a) physical changes, (b) chemical by heating potassium manganate(VII), potassium
changes that occur when iron corrodes.
nitrate or a mixture of potassium chlorate(V)
2 (a) Write the chemical equation for the rusting of
iron. Assume the formula of rust as Fe2O3.H2O. and manganese(IV) oxide (catalyst).

(b) Explain why a layer of grease applied to an (a) 2KMnO4(s) ⎯h⎯eat→ K2MnO4(s) + MnO2(s)
iron object will prevent iron from rusting. + O2(g)

3 (a) What is meant by galvanised iron?
(b) Explain why galvanised iron does not rust (b) 2KNO3(s) ⎯he⎯at→ 2KNO2(s) + O2(g)
even though its surface is scratched.
(c) 2KClO3(s) ⎯Mhne⎯aOt→2 2KCl(s) + 3O2(g)
4 Figure 3.17 shows the arrangement of apparatus
to study the effect of magnesium on the corrosion Reactivity Series of Metals
of zinc.
3 1 The reactivity series is a list of metals arranged
Figure 3.17 according to their chemical reactivity with
oxygen.
(a) Write the half-equations for the reactions that
occur at (i) the zinc rod, (ii) the magnesium foil. 2 Metals high in the reactivity series are very
reactive. These metals react vigorously with
(b) Explain the effect of magnesium on the corrosion oxygen. Potassium is the most reactive of these
of zinc. metals. It is placed at the top of the series.

3 Metals that are less reactive are placed at the
lower part of the series. These metals react slowly
with oxygen. Gold is the most unreactive of these
metals. It is placed at the bottom of the series.

4 Figure 3.18 shows the reactivity series of metals
that do not include hydrogen and carbon.

3.3 The Reactivity Series
of Metals and Its
Applications

Reactivity of Metals with Oxygen

1 Most metals form metal oxides when heated Figure 3.18 The reactivity series of metals
’0S5P /MP 2 or burnt in air. For example, The reactivity series of metals that includes
/SA oxidation hydrogen and carbon is shown in Figure 3.24.
5 The position of a metal in the reactivity series
+1 –2 SPM can also be determined by the reaction between
’05/P1 the metal and the oxide of another metal
0 0 (Figure 3.19).

4Na(s) + O2(g) ⎯⎯⎯→ 2Na2O(s) Figure 3.19 Heating a mixture of a metal and the
reducing oxidising oxide of another metal

agent agent reduction

2 Different metals have different reactivity with

oxygen. Reactive metals have strong affinity

for oxygen. Hence, reactive metals will burn

rapidly and vigorously in oxygen. Conversely,

less reactive metals will react slowly with oxygen.

3 The reactivity of metals with oxygen can be

compared by observing the flame or the glow

produced when a metal is heated in oxygen.

The more reactive the metal is with oxygen, the

more brightly and rapidly the metal burns.

Oxidation and Reduction 418

6 If metal X is more reactive than metal Y with→→ 8 Magnesium is more reactive than copper.
oxygen, metal X will displace metal Y from its→→ Thus, magnesium can displace copper from
oxide when a mixture of powdered X and the copper(II) oxide. Conversely, copper is less
oxide of metal Y is heated, that is, reactive than magnesium. Thus, copper cannot
displace magnesium from magnesium oxide.
more reactive less reactive
9 The displacement reaction can be considered in
terms of oxidation and reduction. For example,
X + oxide of Y ⎯⎯→ oxide of X + Y
oxidation
7 For example, when a mixture of copper(II)
oxide and magnesium is heated, the following Zn(s) + PbO(s) ⎯h⎯eat→ ZnO(s) + Pb(s)
reaction occurs because magnesium is more reduction
reactive than copper.

more reactive less reactive
10 In the reaction between zinc and lead(II) oxide, Experiment 3.5 3
zinc acts as a reducing agent and reduces lead(II)
Mg(s) + CuO(s) ⎯he⎯at→ MgO(s) + Cu(s)
oxide to lead. Conversely, lead(II) oxide acts as an
Conversely, when a mixture of magnesium oxide
and copper is heated, reaction does not occur oxidising agent and oxidises zinc to zinc oxide.
because copper is less reactive than magnesium. 11 The reactivity series can be used to predict

MgO(s) + Cu(s) → No reaction the reaction between a metal and the oxide
of another metal. A more reactive metal will
6 displace a less reactive metal from its oxide.

When a mixture of the oxide of a metal P and the Test 2
powdered metal Q is heated, there is a glow in the Metal R has no reaction with the oxide of P. Thus, R is
mixture. When the experiment is repeated by using less reactive than P.
metal R as a substitute for metal Q, no change occur. Conclusion
From these observations, arrange the reactivity of The reactivity of P, Q and R in ascending order is
P, Q and R in ascending order.
R<P<Q
Solution
Test 1 reactivity increases
The mixture glows because of the reaction between
Q and oxide of P. Thus, metal Q is more reactive
than metal P.

3.5

To deduce the reactivity series of metals (c) Constant variables : The amount of metal and
Problem statement potassium manganate(VII)
How is the reactivity series of metals deduced from used
the reactions of metals with oxygen?
Figure 3.20 The combustion of a metal in oxygen
Hypothesis
The more reactive a metal, the more brightly and
more rapidly the metal will burn in oxygen.

Variables
(a) Manipulated variable : Type of metal
(b) Responding variable : The intensity of the flame

419 Oxidation and Reduction

Apparatus Combustion tube, retort stand with 2 A spatula of zinc powder is placed on a sheet
clamp, spatula and Bunsen burner.
of asbestos paper and put inside the combustion
Materials Potassium manganate(VII), powdered
zinc, iron, lead, copper, magnesium, tube. The combustion tube is then clamped to a
glass wool and asbestos paper.
retort stand.
Procedure 3 The zinc powder is heated strongly (Figure 3.20).
4 When the zinc powder has become very hot,
1 Two spatulas of potassium manganate(VII)
crystals are placed in a combustion tube. A potassium manganate(VII) is heated strongly to
small quantity of glass wool is then placed
inside the combustion tube to prevent potassium produce oxygen gas. The intensity of the flame
manganate(VII) from spilling over. (Caution! If
potassium manganate(VII) is mixed with metal or glow is recorded in the following table.
powder, an explosion may occur during heating). 5 When the reaction has been completed, the

Results combustion tube is set aside to cool and the

contents of the combustion tube taken out.
6 Steps 1 to 5 are repeated using (a) iron powder,

(b) lead powder, (c) copper powder and (d)

magnesium powder.

Metal Intensity of flame/glow Observation
Zinc
Iron • Burns rapidly Colour of hot oxide Colour of cold oxide
• Bright glow
3 Lead Yellow White
Copper • Burns less rapidly
Magnesium • The glow is less bright than the Reddish-brown Reddish-brown

burning of zinc Brown Yellow

• Burns slowly Black Black
• Faint glow White White

• Faint glow

• Burns very rapidly
• Very bright white flame produced

Conclusion (B) Colour of metal oxide
Metals burn in oxygen to form metal oxides. The
Based on the results obtained in this experiment, the chemical equations for the burning of these five
reactivity of the five metals with oxygen is as follows: metals in oxygen are shown below:

(Very Mg > Zn > Fe > Pb > Cu (Very 1 Burning of magnesium
unreactive) 2Mg(s) + O2(g) → 2MgO(s)
reactive) Reactivity decreases Magnesium oxide is white in colour.
2 Burning of zinc
The hypothesis is accepted. 2Zn(s) + O2(g) → 2ZnO(s)
Zinc oxide is yellow when hot and white when cold.
Discussion 3 Burning of iron
4Fe(s) + 3O2(g) → 2Fe2O3(s)
(A) The intensity of flame/glow or the vigour of Iron(III) oxide is reddish-brown.
reaction 4 Burning of lead
2Pb(s) + O2 → 2PbO(s)
1 Magnesium burns very rapidly in oxygen and Lead(II) oxide is brown when hot and yellow
produces a very bright white flame. This shows
that magnesium is very reactive. In contrast, copper when cold.
only glows weakly when it reacts with oxygen. 5 Burning of copper
This shows that copper is very unreactive. 2Cu(s) + O2(g) → 2CuO(s)
Copper(II) oxide is black.
2 Comparing zinc with iron and lead, zinc burns
rapidly and produces a bright glow. This shows that
zinc is more reactive, compared with iron and lead.

3 Iron burns less rapidly and lead burns slowly.
Thus, iron is more reactive than lead.

Oxidation and Reduction 420

The Position of Carbon in the Reactivity agent and the oxide of metal X is reduced to
Series of Metals metal X.

Heating Carbon with Metal Oxides reduction

1 The position of carbon in the reactivity series Carbon + oxide of metal X ⎯⎯→ metal X
can be determined by heating carbon with
metal oxides. reducing +

2 When a mixture of carbon and the oxide of agent carbon dioxide
metal X is heated strongly, a reaction will
occur if carbon is more reactive than the metal oxidation
X. In this reaction, carbon acts as the reducing
3 Conversely, if carbon does not remove oxygen
from a metal oxide, this means that carbon is
less reactive than the metal in the oxide.

3.6

To determine the position of carbon in the reactivity series of metals

Problem statement Variables

How is the position of carbon in the reactivity series (a) Manipulated variable : Type of metal oxide Experiment 3.6 3
of metals determined? (b) Responding variable : Intensity of flame
(c) Constant variable : Carbon powder
Hypothesis

(a) A reaction will occur if carbon is more reactive
than the metal.

(b) A reaction will not occur if carbon is less reactive
than the metal.

(c) Carbon is placed between aluminium and zinc in
the reactivity series of metals.

Apparatus

Spatula, asbestos paper, wire gauze, tripod stand and Figure 3.21 The reaction between carbon and
Bunsen burner. the metal oxide

Materials zinc oxide and carbon is heated strongly for a
few seconds (Figure 3.21).
Powdered carbon, powdered zinc oxide, copper(II) 4 After this, the Bunsen flame is removed and the
oxide and aluminium oxide. mixture examined to determine whether it will
continue to glow.
Procedure 5 Steps 1 to 4 are repeated using a mixture of carbon
and
1 Two spatulas of carbon powder are placed on a (a) copper(II) oxide, (b) aluminium oxide.
piece of asbestos paper.

2 One spatula of zinc oxide is added to the carbon
powder. The zinc oxide and carbon powder are
mixed uniformly.

3 The asbestos paper with its contents is placed on
a wire gauze over a tripod stand. The mixture of

Results

Mixture Observation Reactivity of carbon
(a) C + ZnO Carbon is more reactive than zinc.
• The reaction mixture glows brightly.
(b) C + CuO • A grey solid is formed. Carbon is more reactive than copper.

(c) C + Al2O3 • The reaction mixture burns with a bright flame. Carbon is less reactive than aluminium.
• A brown solid is obtained.

• No visible change

421 Oxidation and Reduction

Discussion 2 Carbon does not react with aluminium oxide.
1 Carbon has reduced zinc oxide and copper(II) Thus, carbon is less reactive than aluminium.

oxide to their respective metals. C(s) + Al2O3(s) ⎯⎯Δ → No reaction

C(s) + 2wZnhOite( s) ⎯⎯Δ → 2Zgnre(ys) + CO2(g) Conclusion
The position of carbon is between aluminium and
C(s) + 2CbluaOck(s ) ⎯⎯Δ → 2bCruow(s)n+ CO2(g) zinc. Thus, the position of carbon in the reactivity
series is
(Note: the symbol ‘Δ’ denotes heating.) K > Ca > Mg > Al > C > Zn > Cu > Hg > Ag > Au
This means that carbon is more reactive than
copper and zinc. Reactivity decreases

The hypothesis is accepted.

3 Heating Carbon Dioxide with Metals 4 When a piece of burning magnesium ribbon
1 The ability of a metal to remove oxygen from is put into carbon dioxide in a gas jar, the

carbon dioxide can be used to determine the magnesium will continue to burn for a short
position of carbon in the reactivity series. time. Black specks of carbon can be seen on
2 Sodium, potassium, calcium, magnesium and the sides of gas jar and magnesium burns to
aluminium are more reactive than carbon. These form a white powder (magnesium oxide).
metals will therefore react with carbon dioxide
and remove oxygen from carbon dioxide. 2Mg(s) + CO2(g) → 2MgO(s) + C(s)
Metal + carbon dioxide → metal oxide + carbon
3 Figure 3.22 shows the arrangement of apparatus 5 This reaction shows that magnesium
used to investigate the reaction between (a) is more reactive than carbon with oxygen,
magnesium and carbon dioxide. (b) acts as the reducing agent,
(c) reduces carbon dioxide to carbon,
Figure 3.22 The burning of magnesium in (d) is itself oxidised to magnesium oxide.
carbon dioxide gas
oxidation


2Mg(s) + CO2(g) → 2MgO(s) + C(s)
reducing
agent

reduction

6 The higher the element is in the reactivity
series, the stronger it acts as a reducing agent
in the redox reaction.

7 Conversely, if a metal does not remove oxygen
from carbon dioxide, it implies that the metal
is less reactive than carbon.

5 ’05

When powdered metal X is heated with the black A metal Y can displace metal X from its salt solution.
oxide of metal Y, B metal X can react with magnesium oxide when
• a glow is seen,
• the residue produced is yellow when it is hot and heated.
C the oxide of metal X can react with iron powder
white when it is cold.
Based on the above observations, it can be deduced when heated.
that D the oxide of metal Y can react with carbon powder

when heated.

Oxidation and Reduction 422

Comments cannot react with heated magnesium oxide. Iron is
The residue is zinc oxide, which is yellow when hot less reactive than zinc. It cannot react with oxide of
and white when cold. zinc when heated.
Metal Y is less reactive than metal X. It cannot displace The oxide of metal Y is copper(II) oxide (black).
X from its salt solution. Carbon can reduce copper(II) oxide to copper.
Metal X (zinc) is less reactive than magnesium. It
Answer D

The Position of Hydrogen in the spread throughout the metal oxide and the
Reactivity Series of Metals metal is produced.
4 If hydrogen gas does not remove oxygen from
1 The position of hydrogen in the reactivity series the metal oxide, hydrogen is less reactive with
oxygen than the metal.
SPM can be determined by passing dry hydrogen gas 5 Hydrogen used for reducing metal oxides to
’05/P1 over hot metal oxides. metals can be produced from the reaction
between dilute sulphuric acid or dilute
2 If hydrogen is more reactive than metal X, hydrochloric acid and zinc.

hydrogen will reduce the oxide of metal X to Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
metal X.
6 Hydrogen gas is dried by passing it through a
reduction drying agent such as concentrated sulphuric
acid or anhydrous calcium chloride.
Experiment 3.7 3
Hydrogen + oxide of metal X → water + metal X
reducing

agent


oxidation

3 If the reaction between hydrogen gas and
a metal oxide occurs, a flame or a glow will

3.7

To determine the position of hydrogen in the reactivity series of metals

Problem statement Safety precautions

How is the position of hydrogen in the reactivity A mixture of hydrogen and air will explode when
series of metals determined? ignited. The following steps must be taken before the
hydrogen is ignited.
Hypothesis 1 Make sure that all the stoppers are fitted tightly
Hydrogen is placed between zinc and iron in the
reactivity series of metals. so that there is no leakage.
2 Make sure that hydrogen gas flows through the
Variables
(a) Manipulated variable : Different types of metal apparatus continuously.
3 Make sure that all the air in the apparatus is removed
oxides
( b) Responding variable : Intensity of flame before igniting hydrogen at the small hole in the
(c) Constant variable : Hydrogen gas combustion tube. This can be carried out as follows:
(a) A sample of gas is collected from the small
Apparatus
A combustion tube with a small hole, round- hole in the combustion tube.
bottomed flask, U-tube, delivery tube, retort stand (b) The gas is then tested with a lighted wooden
with a clamp and Bunsen burner.
splint.
Materials If a ‘pop’ sound is heard, the test is repeated until
Zinc, dilute sulphuric acid, copper(II) oxide, lead(II)
oxide, iron(III) oxide and zinc oxide. no ‘pop’ sound is heard on ignition of the gas.
4 The combustion tube must be tilted downwards

to prevent the flow of water formed back to the
hot part of the tube, thus causing the combustion
tube to crack.

423 Oxidation and Reduction

Procedure round- bottomed flask. It is dried by passing
through anhydrous calcium chloride in a U-tube.
Figure 3.23 Reaction between hydrogen and metal 3 The dry hydrogen gas is passed through the
oxides combustion tube to displace all the air in the
apparatus.
3 1 A small quantity of copper(II) oxide is put on 4 When all the air in the combustion tube has been
an asbestos paper which is then placed in a removed, the hydrogen gas coming out from the
combustion tube with a small hole at one end of small hole of the test tube is ignited.
the tube. 5 Copper(II) oxide is heated strongly while
hydrogen gas is passed over it until the reaction
2 The hydrogen gas is produced by the reaction is complete, that is, until no further changes in
between zinc and dilute sulphuric acid in a colour is observed.
6 Hydrogen gas is allowed to continue to flow
Results through, in order to cool the apparatus and to
prevent air from entering the combustion tube.
Mixture Observation 7 Steps 1 to 6 are repeated using
(a) lead(II) oxide,
(b) iron(III) oxide,
(c) zinc oxide.

Reactivity of hydrogen

(a) H2 + CuO A bright flame is produced. Hydrogen is more reactive than copper.
The black powder changes to brown.

(b) H2 + PbO A bright flame is produced. Hydrogen is more reactive than lead.
The yellow powder changes to grey.

(c) H2 + Fe2O3 A bright glow spreads over iron(III) oxide. Hydrogen is more reactive than iron.
The brown powder changes to grey.

(d) H2 + ZnO No glow is observed. Hydrogen is less reactive than zinc.
The white powder becomes yellow when
heated and white when cold.

Discussion 2 Hydrogen does not react with zinc oxide. Hence,
1 Hydrogen has reduced copper(II) oxide, lead(II) hydrogen is less reactive than zinc.

oxide and iron(III) oxide to their respective ZnO(s) + H2(g) ⎯⎯→ No reaction
metals and is itself oxidised to water. yellow
when hot
CuO(s) + H2(g) → Cu(s) + H2O(l)
black brown Conclusion
Hydrogen is placed above iron but below zinc in the
PbO(s) + H2(g) → Pb(s) + H2O(l) reactivity series of metals. That is,
yellow grey
Zn > H > Fe > Pb > Cu
Fe2O3(s) + 3H2(g) → 2Fe(s) + 3H2O(l)
brown grey Reactivity decreases

This means that hydrogen is more reactive than The hypothesis is accepted.
copper, lead and iron.

Oxidation and Reduction 424

Positions of Carbon and Hydrogen in the Reactivity Series
1 The reactivity series that includes both carbon and hydrogen is shown in Figure 3.24.

SPM

’07/P1

Figure 3.24 The reactivity series that includes both carbon and hydrogen

2 Table 3.11 summarises the reactions between 3
(a) metal oxides and carbon,
(b) metal oxides and hydrogen.

Table 3.11 Reduction of metal oxides with carbon and hydrogen

Metal oxide Reaction with carbon Reaction with hydrogen

Potassium oxide (K2O) No reaction No reaction
Sodium oxide (Na2O)
Calcium oxide (CaO) 2ZnO + C → 2Zn + CO2 No reaction
Magnesium oxide (MgO) 2Fe2O3 + 3C → 4Fe + 3CO2 Fe2O3 + 3H2 → 2Fe + 3H2O
Aluminium oxide (Al2O3) SnO2 + C → Sn + CO2 SnO2 + 2H2 → Sn + 2H2O
2PbO + C → 2Pb + CO2 PbO + H2 → Pb + H2O
Zinc oxide, ZnO 2CuO + C → 2Cu + CO2 CuO + H2 → Cu + H2O
2Ag2O + C → 4Ag + CO2 Ag2O + H2 → 2Ag + H2O
Iron(III) oxide (Fe2O3)
Tin(IV) oxide (SnO2)
Lead(II) oxide (PbO)
Copper(II) oxide (CuO)
Silver(I) oxide (Ag2O)

1 The position of a metal in the reactivity series can In general, the electrochemical series is the same as
be used to predict the ability of a metal to react the reactivity series. However, the position of hydrogen
with water. in these two series are different. In the electrochemical
series, hydrogen is placed between lead and copper.
2 Metals that are more reactive than hydrogen In the reactivity series, the position of hydrogen is
will reduce water to hydrogen. The higher the between zinc and iron.

position of the metal in the reactivity series, the Extraction of Metals from Their Ores

faster and more vigorous the metal will react with 1 Most metals in metal ores exist in the forms
of oxides, carbonates and sulphides in the
water. Conversely, metals that are less reactive Earth’s crust.

than hydrogen do not react with water.
3 Hence, potassium and sodium react vigorously

with cold water and magnesium reacts with steam

but not cold water. Lead and copper do not react

with water.

425 Oxidation and Reduction

2 The extraction of metals involves the reduction • Modern blast furnaces are very tall (50 – 70 m high).
of metal ores to metals. They are made of steel and lined with fireproof bricks.

3 Two main methods are used to extract metals • Coke is a form of carbon made by heating coal in
from their ores. the absence of air. The process is called destructive
(a) Electrolysis of metal compounds in the distillation.
molten state.
(b) Reduction of metal oxides by carbon. 1 The important iron ores are haematite and
magnetite. Haematite contains iron(III) oxide,
4 The important factor for determining the most Fe2O3 whereas magnetite contains triiron
suitable method in the extraction of metals tetroxide, Fe3O4.
is the position of the metal in the reactivity
series. 2 The extraction of iron from haematite or
magnetite is carried out in the blast furnace
5 For metals lower than carbon in the reactivity (Figure 3.25) by reduction using carbon.
series of metals, the method used is the
reduction of metal oxides by carbon. Figure 3.25 The blast furnace
(a) Raw materials required
6 For metals higher than carbon in the reactivity
series, the extraction of metals must be carried A mixture of iron are, coke (carbon) and
out by the electrolysis of molten metal limestone (calcium carbonate) is put in
compounds. a blast furnace. Hot air is blown into the
furnace from the bottom.
7 Table 3.12 shows the methods used to extract (b) Production of carbon dioxide
some metals. In the lower section of the blast furnace,
the oxygen in hot air reacts with coke to
3 form carbon dioxide.

Table 3.12 Methods of metal extraction according to
the position of the metal in the reactivity
series

Metal Method of extraction

Potassium, K Electrolysis of metal chlorides
Sodium, Na in the molten state
Calcium, Ca
Magnesium, Mg

Aluminium, Al Electrolysis of Al2O3 in the
molten state

Zinc, Zn Heating metal oxides with
Iron, Fe carbon
Tin, Sn
Lead, Pb

Copper, Cu Heating metal sulphides in air
Mercury, Hg

Silver, Ag Exist as free elements in
Gold, Au Earth’s crust

Extraction of Iron from Its Ore

C(s) + O2(g) → CO2(g)

At high temperatures, limestone decomposes
into quicklime (calcium oxide, CaO) and
carbon dioxide.

Molten iron flowing out from the blast furnace CaCO3(s) → CaO(s) + CO2(g)

Oxidation and Reduction 426

(c) Production of carbon monoxide • Limestone Waste gases Slag
In the upper section of the blast furnace, • Iron ore Blast furnace Molten iron
carbon dioxide reacts with coke to produce • Coke
carbon monoxide. Hot air

C(s) + CO2(g) → 2CO(g) (b) The iron produced in this process is not

(d) Reduction of iron ore to iron pure iron and contains about 5% carbon.
• In the upper section of the blast furnace, The iron is called cast iron.
where the temperature is about 400 –
800 °C, the iron ore is reduced by carbon
monoxide to iron.

Fe2O3(s) + 3CO(g) → 2Fe(l) + 3CO2(g) • Cast iron is hard and brittle. Cast iron is used for making
haematite
Fe3O4(s) + 4CO(g) → 3Fe(l) + 4CO2(g) the base of Bunsen burner, lamp posts or pipes.
magnetite
• Cast iron is usually converted into other more useful
forms such as steel in an oxygen furnace.

• In the lower section of the blast furnace, Thermite Process 3
the iron ore is reduced by coke (carbon)
to iron. 1 The thermite process is a displacement reaction
between aluminium and iron(III) oxide to
Fe2O3(s) + 3C(s) → 2Fe(l) + 3CO(g) produce iron.

Fe3O4(s) + 2C(s) → 3Fe(l) + 2CO2(g) 2Al(s) + Fe2O3(s) → Al2O3(s) + 2Fe(l)

• In these reactions, carbon and carbon Thus, the thermite process can be used for the
monoxide act as the reducing agents. small-scale extraction of iron.

• The molten iron produced flows to 2 The thermite process can be carried out in the
the bottom of the blast furnace and is school laboratory by using the apparatus as
collected. The molten iron is poured shown in Figure 3.26. Magnesium acts as the
into moulds and set aside to solidify. fuse to ignite the mixture for this experiment.

3 Removal of impurities
(a) In the blast furnace, calcium oxide is produced
from the decomposition of limestone. It
then reacts with silica (sand) to form slag
(calcium silicate).

CaO(s) + SiO2(s) → CaSiO3(s) Figure 3.26 Thermite process

(b) Molten slag floats on top of iron. The slag 3 When the mixture of magnesium powder and
and iron are separated through a tap at
the bottom of the furnace. barium peroxide (BaO2) burns, a large amount
of heat is produced to initiate the thermite
(c) In this reaction, calcium oxide acts as a
basic oxide, whereas silica acts as an acidic process to produce molten iron.
oxide. 4 We can also consider the thermite process as a

(d) The slag produced during the extraction of redox reaction:
iron is used mainly for road surfacing.
oxidation
4 (a) The extraction of iron can be summarised
in the flowchart as follows. 2Al(s) + Fe2O3(s) → Al2O3(s) + 2Fe(l)


reduction

427 Oxidation and Reduction

5 The thermite process is highly exothermic, (c) The carbon monoxide produced can also
that is, it gives out a lot of heat during the reduce tin(IV) oxide to tin.
reaction. It is used for welding steel objects,
for example, railway lines. SnO2(s) + 2CO(g) → Sn(l) + 2CO2(g)

Coke is a form of carbon. It is the solid substance left (d) The molten tin is then tapped off and
after destructive distillation of coal. The term ‘destructive poured into a mould and solidified into
distillation’ means heating a substance in the absence ingots.
of air. Coke contains more than 80% of carbon. Coke
is used as a reducing agent in the extraction of metals 5 The extraction of tin can be summarised in
and as a non-smoky fuel. the flowchart below.

Extraction of Tin from Its Ore tin ore

1 The most important tin ore is cassiterite. water + oil
Cassiterite contains tin(IV) oxide, SnO2 and
unwanted materials such as sand, soil, oil, soil froth floatation sand
sulphur and carbon.
3 concentrated tin ore
2 The two main steps involved in the extraction
of tin are carbon roasting
(a) concentration process, sulphur
(b) reduction process. carbon
furnace dioxide
3 Concentration process +
(a) At the first stage of tin extraction, the tin carbon
ore is concentrated by froth floatation coke monoxide
method. In this process, the tin ore is
crushed to a fine powder and mixed with tin ingots
water and special oils (known as frothing
agents) in a large tank. The Use of Carbon as the Main Reducing Agent
(b) The mixture is agitated by blowing air to
form a froth. The unwanted materials sink in Metal Extraction
to the bottom of the tank.
(c) The froth contains particles of concentrated The main reasons for using carbon as the main
tin ore and floats to the top of the tank reducing agent in metal extraction are:
where it is removed. The concentrated tin 1 Chemical reason
ore is then dried and roasted to remove
impurities such as carbon, sulphur and Carbon is more reactive than zinc, iron, tin
special oils. and lead. Therefore, carbon can easily reduce
the oxides of these metals.
4 Reduction process 2 Economic reason
(a) The concentrated tin ore is mixed with Carbon is cheap and can be obtained easily.
coke (a form of carbon). The mixture is Reduction of metal ores using coke (carbon) is
heated to a high temperature (about 1360 cheaper than using electricity for the electrolysis
°C) in a furnace. of molten ores.
(b) During heating, tin(IV) oxide is reduced 3 Environmental reason
by carbon to molten tin and carbon is The carbon dioxide gas produced during metal
oxidised to carbon dioxide and carbon extraction is non-poisonous and does not
monoxide. pollute the atmosphere.

SnO2(s) + C(s) → Sn(l) + CO2(g)

SnO2(s) + 2C(s) → Sn(l) + 2CO(g)

Oxidation and Reduction 428

Reactions of metals with oxygen
The more rapidly the metal burns in oxygen and the brighter the
flame produced, the more reactive the metal is with oxygen.

is used to build up reactivity series of metals

Position of carbon in the The reactivity series of metals Position of hydrogen in the
reactivity series The reactivity series of metals is reactivity series
a series of metals arranged in
• Carbon will reduce the oxide the order of how vigorously the • The position of hydrogen can
of metal X if carbon is more metals react with oxygen.
reactive than X. be determined by passing the
K > Na > Ca > Mg > Al > C >
• The position of carbon is Zn > H > Fe > Sn > Pb > Cu hydrogen gas over hot metal
between aluminium and zinc.
reactivity decreases → oxides.

• The position of hydrogen is
between zinc and iron.

uses of reactivity series in the extraction of metals

3

Extraction of iron Extraction of tin
• Raw materials: Iron ore (haematite or magnetite), • Raw materials: Tin ore (cassiterite) and coke.
• The froth floatation method is used to produce
limestone and air.
• Iron is produced by the reduction of iron ore by concentrated tin ore.
• Tin is produced by the reduction of cassiterite by
coke (carbon) or carbon monoxide:
coke or carbon monoxide:
Fe2O3(s) + 3C(s) → 2Fe(l) + 3CO(g)
SnO2(s) + C(s) → Sn(l) + CO2(g)
Fe2O3(s) + 3CO(g) → 2Fe(l) + 3CO2(g)
SnO2(s) + 2CO(g) → Sn(l) + 2CO2(g)
• Limestone is used to form slag.

CaCO3(s) → CaO(s) + CO2(g)

CaO(s) + SiO2 → CaSiO3(s)

3.3 (a) Based on the experimental results, arrange the
metals P, Q, R and T in descending order of the
1 The table below shows the experimental results when reactivity of the metals with oxygen.
a mixture of a metal and the oxide of another metal
is heated strongly. (b) Predict whether metal Q will react with the oxide
of metal P. Explain your answer.
Mixture Observation
The mixture glows 2 The oxide of metal X can be reduced by carbon but
Metal P and oxide of not hydrogen.
metal Q The mixture glows
(a) Identify the metal X.
Metal Q and oxide of The mixture does
metal R not glow (b) (i) Write the equation for the reaction between
The mixture does not the oxide of metal X and aluminium powder.
Metal Q and oxide of glow
metal T (ii) Identify the oxidising and reducing agents
in this reaction.
Metal P and oxide of
metal T 3 State four different methods that can be used to
obtain lead from lead(II) oxide. Write the equations
for all the reactions that take place.

429 Oxidation and Reduction

3.4 Redox Reactions in (a) the external circuit,
Electrolytic Cell and (b) the reactions occurring in the electrolyte,
Chemical Cell
and
1 In Chapter 6 of Form 4, you have already (c) the reactions at the electrodes.
studied the concepts of electrolytic cell and 3 When electrolysis occurs, electrical energy is
chemical cells. converted into chemical energy. The electrical
energy is used to decompose the electrolyte in
2 Chemical cells are also known as voltaic the electrolytic cell.
cells. 4 Inert electrodes such as carbon or platinum are
usually used as electrodes in electrolytic cells.
3 The reactions that occur in electrolytic cells and However, in some electrolytic cells, one of the
chemical cells are redox reactions involving electrodes is an inert electrode and the other
the transfer of electrons. electrode is a metal electrode. There are also
electrolytic cells in which both the electrodes
4 Whether it is an electrolytic cell or a chemical are of the same metal.
cell, oxidation occurs at the anode, and reduction
occurs at the cathode. Chemical Cells SPM

Differences between an Electrolytic Cell ’04/P1, ’05/P1
and a Chemical Cell ’05/P2, ’07/P1

3 Electrolytic Cells 1 The chemical cell is set up by dipping two
different metals in an electrolyte. The metals
1 The basic structure of an electrolytic cell is as
follows: act as the electrodes in the chemical cell.
(a) A battery to supply electrical energy, 2 Figure 3.28 shows the apparatus in a chemical
(b) An electrolyte to supply free (mobile)
ions for conducting electric current, cell, using zinc and copper electrodes and
(c) Two electrodes for the transfer of electrons
to occur. Electrons are transferred from the dilute sulphuric acid as the electrolyte.
anions (negative ions) to the anode and from 3 The basic structure of a chemical cell is as
the cathode to the cations (positive ions).
follows:
2 Figure 3.27 shows the apparatus set-up in (a) A connecting wire for electrons to flow
an electrolytic cell. The electrolysis process
involves three main aspects, that is, through in the external circuit,
(b) An electrolyte for electric current to flow

through,
(c) Two electrodes for the transfer of electrons

from a reducing agent to an oxidising agent.
4 In a chemical cell, chemical energy is converted

into electrical energy.

External circuit
During electrolysis, electrons flow from the anode (positive terminal) to the cathode (negative terminal)

Anode (positive terminal) Cathode (negative terminal)
During electrolysis, During electrolysis,
• anions (negative ions) move towards the • cations (positive ions) move towards the

anode cathode
• anions release electrons at the anode • cations accept electrons from the
• oxidation occurs at the anode
cathode
• reduction occurs at the cathode

Electrolyte
During electrolysis
• cations move towards the cathode
• anions move towards the anode
• the flow of ions to the electrodes constitute the flow of electric current in the electrolyte

Figure 3.27 The movement of ions and electrons during electrolysis

Oxidation and Reduction 430

External circuit

Electric current is produced because electrons flow from the more electropositive
electrode (negative terminal) to the less electropositive electrode (positive terminal).

Anode (negative terminal) Cathode (positive terminal)
• H+ ions accept electrons from
• Zinc dissolves to form zinc
ions with the release of zinc to form hydrogen gas.
electrons. 2H+(aq) + 2e– → H2(g)

Zn(s) → Zn2+(aq) + 2e– • Effervescence occurs around
the copper electrode.
• The electrons flow through the
external circuit to the copper • Reduction occurs at the
electrode. cathode.

• Oxidation occurs at the Electrolyte
anode. • The concentration of Zn2+ ions increases, while the concentration of H+ ions decreases.
• The overall reaction that occurs in the chemical cell is
3
Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)

• Zinc acts as the reducing agent and hydrogen ions act as the oxidising agent.

Figure 3.28 A simple chemical cell

Redox Reactions in Electrolytic Cells SPM 2 The redox reactions that occur at the electrodes
’09/P1 are shown in Table 3.13.

1 Ionic compounds, in the molten state, or Table 3.13 Electrolysis of molten lead(II) bromide
dissolved in water, are electrolytes.
At the cathode At the anode
2 During electrolysis, oxidation occurs at the
anode and reduction occurs at the cathode. (a) Pb2+ ions gain (a) Br– ions lose
(a) At the anode: An– → A + ne– … oxidation electrons to form electrons to form
(b) At the cathode: Bn+ + ne– → B ... reduction lead metal. bromine molecules.

3 In an electrolytic cell, electrons flow from Pb2+(l) + 2e– → Pb(l) 2Br–(l) → Br2(l) + 2e–
the anode (positive electrode) to the cathode
(negative electrode) through the connecting (b) This is a reduction (b) This is an oxidation
wire (Figure 3.29). process. The process. The
oxidation number of oxidation number of
lead decreases from bromine increases
+2 to 0. from –1 to 0.

Overall reaction
The overall reaction is the breakdown of lead(II)
bromide to give lead and bromine:

PbBr2(l) → Pb(l) + Br2(l)

Figure 3.29 Electrolytic cell Electrolysis of Copper(II) Sulphate Solution
using Inert Electrodes
Electrolysis of Molten Lead(II) Bromide SPM
1 Electrolysis of copper(II) sulphate solution can
’09/P1 be carried out using the apparatus as shown in
Figure 3.30. Platinum (Pt) electrodes are used
1 When molten lead(II) bromide is electrolysed, as inert electrodes in this experiment.
the cations (lead(II) ions, Pb2+) are attracted
to the cathode and the anions (bromide ions,
Br–) are attracted to the anode.

431 Oxidation and Reduction

SPM Electrolysis of Copper(II) Sulphate Solution

’10/P1

using Copper Electrodes

Figure 3.30 Electrolysis of CuSO4 using inert electrodes 1 If the electrolysis of copper(II) sulphate solution
is carried out using reactive electrodes such as
copper electrodes, a different reaction occurs
at the anode, that is, both the OH– ions and
the SO42– ions are not discharged. Instead, the
copper anode dissolves (corrodes) to form
copper(II) ions.

2 An aqueous solution of copper(II) sulphate Cu(s) → Cu2+(aq) + 2e–
contains four types of ions.
2 At the copper anode, three possible reactions

From CuSO4 : Cu2+(aq) and SO42–(aq) that can occur are:

From H2O : H+(aq) and OH–(aq) (a) SO42– ions are discharged by donating
(b) OH– ions are discharged
3 The cations, Cu2+ and H+ ions are attracted to (c) Copper metal is oxidised electrons
the cathode and the anions, SO42– and OH–
ions, are attracted to the anode. to Cu2+ ions

4 Table 3.14 shows the redox reactions that 3 Because copper is a reactive electrode,
occur at the electrodes.
3 the reaction that can occur most easily

is the conversion of copper to Cu2+ ions.

Table 3.14 Redox reactions at platinum electrodes Consequently, SO42– ions and OH– ions
remain in the solution and are not discharged.

during the electrolysis of CuSO4(aq) Instead, the copper anode dissolves to form

At the cathode At the anode Cu2+(aq) ions and the electrode becomes

(negative electrode) (positive electrode) thinner.

(a) Copper is below (a) OH– ion is below Cu(s) → Cu2+(aq) + 2e–

hydrogen in the SO42– ion in the 4 The redox reactions that occur at the electrodes
electrochemical electrochemical during the electrolysis of copper(II) sulphate
using copper electrodes are shown in Table 3.15.
series. Hence, Cu2+ series. Hence, OH–

ions are preferentially ions are preferentially

discharged at the discharged at the Table 3.15 Redox reactions at copper electrodes

cathode. anode. during electrolysis of CuSO4(aq)

reduction oxidation At the cathode At the anode
Cu2+(aq) + 2e– ⎯⎯⎯→ 4OH–(aq) ⎯⎯⎯→

Cu(s) O2(g) + 2H2O + 4e– Cu2+(aq) + 2e– → Cu(s) Cu(s) → Cu2+(aq) + 2e–

(b) At the cathode, Cu2+ (b) At the anode, OH– … reduction … oxidation

ions are reduced to ions are oxidised to Overall reaction
• The overall reaction is the transfer of copper from
copper metal. oxygen gas.
the anode to the cathode.
(c) H+ ions remain in the (c) SO42– ions remain in • The concentration of copper(II) sulphate does not
solution. the solution.
change and the blue colour of the electrolyte does
Overall reaction not fade.
Copper metal is deposited at the cathode, oxygen
gas is given off at the anode and the solution Electrolysis of Concentrated Sodium Chloride
becomes more and more acidic. Solution

reduction 1 Figure 3.31 shows the apparatus set-up for the
electrolysis of concentrated sodium chloride
electrolysis solution.
2CuSO4 + 2H2O ⎯⎯⎯⎯→ 2Cu + O2 + 2H2SO4


oxidation

Oxidation and Reduction 432

A simple way to remember redox reactions at the

electrodes is to remember ‘RED CAT’.
REDuction occurs at CAThode

Conversely, oxidation occurs at the anode.

Electrolysis of Dilute Sodium Chloride Solution

Figure 3.31 Electrolysis of concentrated sodium 1 A dilute sodium chloride solution contains:
chloride solution
(a) Na+(aq) and Cl–(aq) from NaCl(aq)

2 An aqueous solution of sodium chloride (b) H+(aq) and OH–(aq) from H2O(l) that
contains four types of ions: 2 Table 3.17 shows the redox reactions
(a) From NaCl: Na+(aq) and Cl–(aq)
(b) From H2O: H+(aq) and OH–(aq) occur and the products obtained at the carbon

3 Na+ ions and H+ ions are attracted to the electrodes.
cathode. Cl– ions and OH– ions are attracted
to the anode. Table 3.17 Redox reactions during electrolysis of
dilute sodium chloride solution
4 Table 3.16 shows the redox reactions that occur
at the anode and cathode when concentrated At the cathode At the anode 3
sodium chloride solution is electrolysed using (Na+ and H+ ions are (Cl– and OH– ions are
inert electrodes.
present) present)

H+ ions gain electrons OH– ions donate electrons

from the cathode to form to the anode to form

Table 3.16 Redox reactions during electrolysis of hydrogen gas. oxygen gas and water.
concentrated sodium chloride solution
2H+(g) + 2e– → H2(g) 4OH–(aq) →
O2(g) + 2H2O(l) + 4e–
At the cathode (negative At the anode (positive
electrode), H+ and electrode), Cl– and Overall reaction
Na+ ions are present OH– ions are present
oxidation
(a) Hydrogen ions (H+) (a) Chloride ions (Cl–)
are preferentially are preferentially electrolysis
discharged at the discharged at the
cathode. anode. 2H2O(l) ⎯⎯⎯⎯⎯⎯⎯→ 2H2(g) + O2(g)
reduction
2H+(aq) + 2e– → H2(g) 2Cl–(aq) → Cl2(g) + 2e–
• Electrolysis of dilute sodium chloride solution
(b) At the cathode, H+ (b) At the anode, Cl– produces two volumes of hydrogen at the cathode
ions are reduced to ions are oxidised to and one volume of oxygen at the anode.
hydrogen gas. chlorine gas.
• Since water is being removed (by decomposition
(c) Na+ ions remain in (c) OH– ions remain in to form H2 and O2), the concentration of sodium
the solution. the solution. chloride increases gradually.

Overall reaction
oxidation

electrolysis Electrolysis of
2NaCl(aq) + 2H2O(l) ⎯⎯⎯⎯→
2NaOH(aq) + H2(g) + Cl2(g) • tmdmhileoouldtteceeanrNtahtNaeoCaldylCepl,crpoaorndomcudecuinxecttsuerarsHeteN2d(oagf()NlO)aa2naC(ndlgd)OpCra2o(nl2dgd()ugc)Cels2(gH) 2(agt )thaet
reduction •
•
Electrolysis of concentrated sodium chloride
solution produces one volume of hydrogen at the anode
cathode, one volume of chlorine at the anode and
sodium hydroxide solution. • concentrated NaCl produces H2(g) and Cl2(g).

433 Oxidation and Reduction

Redox Reactions in Chemical Cells 5 At the negative electrode (zinc plate)
(a) Zinc is more electropositive than copper.
1 In chemical cells, the electric current is produced Hence, zinc has a greater tendency to donate
from chemical reactions that occur in the cell. its electrons compared to copper and is
Examples of chemical cells are: oxidised to zinc ions.
(a) Daniell cell
(b) Dry cell/alkaline cell Zn(s) → Zn2+(aq) + 2e– … oxidation
(c) Lead-acid accumulator
(b) At the negative electrode, oxidation occurs
2 In chemical cells, oxidation occurs at the negative and the zinc metal acts as the reducing
terminal (anode) while reduction occurs at agent.
the positive terminal (cathode).
6 At the positive electrode (copper plate)
Daniell Cell (a) At the positive electrode, Cu2+ ions gain
electrons from zinc and is reduced to
1 The Daniell cell is made up of a zinc plate copper metal.
dipped into zinc sulphate solution and a
copper plate dipped into copper(II) sulphate Cu2+(aq) + 2e– → Cu(s) … reduction
solution (Figure 3.32).
(b) At the positive electrode, reduction occurs
3 and Cu2+ ions act as the oxidising agent.

7 Overall reaction
(a) The overall reaction that occurs in the
Daniell cell is a redox reaction.

oxidised

Figure 3.32 Daniell cell Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
2 The function of the porous pot is to
SPM (a) separate the zinc sulphate solution from reduced
’11/P1 the copper(II) sulphate solution so that
(b) The redox reactions that occur in the
the solutions do not mix. Daniell cell and many other chemical cells
(b) complete the electric circuit by allowing are displacement reactions. A metal higher
up in the electrochemical series displaces
the ions to pass through it. another metal lower in the electrochemical
3 The Daniell cell can also be set up by using a series from an aqueous solution of its salt.

salt bridge to replace the porous pot as shown (c) When the Daniell cell is in use,
in Figure 3.33. (i) the concentration of Zn2+ ions in the

Figure 3.33 Daniell cell solution increases.
4 Zinc is more electropositive than copper (ii) the blue colour of copper(II) sulphate

and is placed higher than copper in the solution fades gradually as more copper
electrochemical series. Hence, zinc plate acts is deposited and the concentration of
as the negative electrode and copper plate acts Cu2+ decreases.
as the positive electrode. (iii) the mass of zinc electrode decreases
gradually.
(iv) the mass of copper electrode increases
gradually.
8 Cell symbols
(a) The cell symbols are used to represent
the chemical cells. For example, the cell
symbol for Daniell cell is

Zn(s)/Zn2+(aq) // Cu2+(aq)/ Cu(s)

(b) According to the IUPAC rules, the negative
electrode is written on the left of the cell
symbol and the positive electrode is written
on the right. The symbol ‘//’ represents the
porous pot or the salt bridge.

Oxidation and Reduction 434

anode cathode 9 Voltage of Daniell cell
(a) If the concentrations of both ZnSO4 and
(negative electrode) (positive electrode) CuSO4 solutions are 1.0 mol dm–3, the
maximum voltage of the Daniell cell is
1.10 V.
Zn(s)/Zn2+(aq) // Cu2+(aq)/ Cu(s) (b) The voltage of the cell will decrease with
time when the cell is being used because
the concentration of Cu2+ ions decreases.

electrolyte at anode electrolyte at cathode

salt bridge (or porous pot)

To study the reactions that occur in the Daniell cell

Apparatus 2 The electrodes are then washed with distilled Activity 3.4 3
Beaker, porous pot, sandpaper, voltmeter, connecting
wires with crocodile clips and electronic balance. water, dried and weighed.
Materials 3 A beaker is filled with 1.0 mol dm–3 copper(II)
1.0 mol dm–3 copper(II) sulphate solution, 1.0 mol dm–3
zinc sulphate solution, zinc plate and copper plate. sulphate solution.
Procedure 4 A porous pot is filled with 1.0 mol dm–3 zinc

Figure 3.34 The Daniell cell sulphate solution.
1 A piece of zinc plate and a piece of copper plate 5 The zinc plate is dipped into the solution of zinc

are cleaned using sandpaper. sulphate while the copper plate is dipped into the

solution of copper(II) sulphate.
6 The zinc and copper electrodes are then connected

to the voltmeter as shown in Figure 3.34.
7 The Daniell cell is allowed to operate for 30

minutes.
8 After 30 minutes, the zinc and copper electrodes

are removed from the electrolytes.
9 The electrodes are rinsed with distilled water,

dried and weighed again.
10 The changes that occur in the electrodes, the

electrolytes and the voltmeter are recorded.

Results

1 The blue colour of copper(II) sulphate fades gradually until it becomes colourless.
2 The zinc plate becomes thinner and the copper plate becomes thicker.
3 The voltmeter needle is deflected. The deflection of the voltmeter needle shows that electrons flow from the

zinc electrode to the copper electrode.
4 The changes in mass of the zinc and copper electrodes are shown below.

Electrode Mass of electrode Mass of electrode Inference
Zinc plate before experiment (g) after experiment (g)
• The mass of zinc plate
20.50 18.29 decreases.

Copper plate 25.74 27.89 • Oxidation occurs at the
anode.

• The mass of copper plate
increases.

• Reduction occurs at the
cathode.

435 Oxidation and Reduction

Discussion 3 The overall reactions is
1 Zinc metal is more electropositive than copper
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
metal. Hence, zinc metal has a greater tendency … redox reaction
to give up electrons to form zinc ions.
4 In this reaction, Cu2+ ions act as the oxidising
Zn(s) → Zn2+(aq) + 2e– … oxidation agent and zinc acts as the reducing agent.

2 The electrons released are accepted by copper(II) Conclusion
ions to form copper metal. 1 Oxidation occurs at the anode (zinc plate).
2 Reduction occurs at the cathode (copper plate).
Cu2+(aq) + 2e– → Cu(s) … reduction 3 The reaction that occurs in the Daniell cell is a

redox reaction.

3 Different types of chemical cells
1 There are two main classes of chemical cells:

Primary cells and secondary cells.
2 Primary cells are not rechargeable. Examples of primary

cells are dry cell, alkaline cell and mercury cell.
3 Secondary cells are rechargeable. Examples of

secondary cells are lead-acid accumulator and

nickel-cadmium (Ni-Cd battery).

Dry Cell Figure 3.35 The structure of a dry cell

1 A dry cell is made up of a zinc container as 2 When the dry cell is used to generate electrical
the anode (negative terminal) and a carbon energy, oxidation occurs at the negative
rod as the cathode (positive terminal). The terminal (zinc container) and reduction occurs
electrolyte in the dry cell is a paste consisting at the positive terminal (carbon rod).
of ammonium chloride, zinc chloride and a
little water (Figure 3.35). 3 The reactions at the electrodes are shown
below.

Reactions at anode (negative terminal) Reactions at cathode (positive terminal)

• Zinc is oxidised to Zn2+. • NH4+ is reduced to NH3 and H2.

oxidation reduction


0 +2 +1 0
Zn(s) → Zn2+(aq) + 2e–
2NH4+(aq) + 2e– → 2NH3(g) + H2(g)

• Electrons flow from the zinc container to the • The H2 gas produced is removed by the reaction
carbon rod. with MnO2.
2MnO2 + H2 → Mn2O3 + H2O

Summary
• The overall reaction is a redox reaction.

Zn + 2NH4+ + 2MnO2 → Zn2+ + 2NH3 + Mn2O3 + H2O

• Oxidising agent: Ammonium ion, NH4+
Reducing agent: Zinc

Oxidation and Reduction 436

Alkaline Cell

1 Figure 3.36 shows the structure of an alkaline Figure 3.36 The structure of an alkaline cell
cell. An alkaline cell is also known as an alkaline
battery.
Negative terminal: Zinc container
Positive terminal: Manganese(IV) oxide powder
Electrolyte: Lithium hydroxide (LiOH) or
potassium hydroxide (KOH)

2 The reactions at the electrodes during discharge
are shown below.

Reactions at anode (negative terminal) Reactions at cathode (positive terminal)
• Zinc is oxidised to Zn2+.
• MnO2 is reduced to manganese(III) oxide, Mn2O3.

oxidation reduction

0 +2 +3
Zn(s) → Zn2+(aq) + 2e– +4

2MnO2(s) + H2O(l) + 2e– → Mn2O3(s) + 2OH–(aq)

• Electrons flow from the zinc container to MnO2. 3

Summary • Oxidising agent: Manganese(IV) oxide, MnO2
• The overall reaction is a redox reaction. Reducing agent: Zinc

Zn + 2MnO2 + H2O → Zn2+ + Mn2O3 + 2OH–

Mercury Cell

1 Figure 3.37 shows the structure of mercury cell. Figure 3.37 The structure of mercury cell
Positive terminal: Mercury(II) oxide, HgO
Negative terminal: Zinc metal Lead-acid Accumulator
Electrolyte: Potassium hydroxide

2 The reactions at the anode and cathode of
mercury cell are summarised below.

Reactions at anode Reactions at cathode SPM
(negative terminal) (positive terminal)
’07/P1
• Zinc is oxidised to • HgO is reduced to
Zn2+. mercury. 1 A lead-acid accumulator is often known as a car
battery. The accumulator is a battery (chemical
Zn(s) + 2OH–(aq) → HgO(s) + H2O(l) + 2e– cell) that can be recharged by passing a current
Zn(OH)2(aq) + 2e– → Hg(l) + 2OH–(aq) through it from an external direct current
(d.c.) supply.
• Electrons flow from
the zinc electrode to 2 Figure 3.38 shows the structure of a lead-acid
HgO. accumulator.

Summary
• The overall reaction is a redox reaction.

Zn(s) + HgO(s) + H2O(l) → Zn(OH)2(s) + Hg(l)

• Oxidising agent: Mercury(II) oxide
Reducing agent: Zinc

Figure 3.38 The structure of a lead-acid accumulator

437 Oxidation and Reduction

Positive terminal: Lead plate coated with PbO2 3 Table 3.18 shows the reactions at the electrodes
Negative terminal: Lead plate during discharge, that is, when the battery
Electrolyte: Sulphuric acid supplies electricity.

Table 3.18 Redox reactions that occur during discharge in lead-acid accumulator

At the anode (negative terminal) At the cathode (positive terminal)

(a) Lead is oxidised to lead(II) ions with the (a) Lead(IV) oxide is reduced to Pb2+ ions by
release of electrons. accepting electrons.

oxidation reduction

0 +2
Pb(s) → Pb2+(aq) + 2e– +4 +2
PbO2(s) + 4H+(aq) + 2e– → Pb2+(aq) + 2H2O(l)
(b) The electrons given out at the cathode flow
through the external circuit to the positive (b) A white solid is produced when Pb2+ ions
terminal. react with SO42– ions in sulphuric acid to form
lead(II) sulphate.
(c) A white precipitate is produced when Pb2+ ions
react with SO42– ions in the sulphuric acid to Pb2+(aq) + SO42–(aq) → PbSO4(s)
form lead(II) sulphate.
(c) The white solid, lead(II) sulphate, then deposits
3 Pb2+(aq) + SO42–(aq) → PbSO4(s) on the surface of the positive electrode to form
a white coating.
(d) The negative electrode becomes white because
white solid lead(II) sulphate is deposited on its (d) The overall reaction at the positive electrode
surface. during discharge is

(e) The overall reaction at the negative electrode PbO2(s) + 4H+(aq) + 2e– → Pb2+(aq) + 2H2O(l)
during discharge is Pb2+(aq) + SO42–(aq) → PbSO4(s)

Pb(s) → Pb2+(aq) + 2e– PbO2(s) + 4H+(aq) + SO42–(aq) + 2e– →
brown PbSO4(s) + 2H2O(l)
Pb2+(aq) + SO42–(aq) → PbSO4(s)
Pb(s) + SO42–(aq) → PbSO4(s) + 2e– white
grey white

Summary
• The overall cell reaction is a redox reaction.

Pb(s) + PbO2(s) + 4H+(aq) + 2SO42–(aq) → 2PbSO4(s) + 2H2O(l)

2H2SO4(aq)
• Oxidising agent: Lead(IV) oxide, PbO2
Reducing agent: Lead
• During discharge, sulphuric acid is used up.

4 Recharging of lead acid accumulator (a) At the negative terminal
When the battery is fully charged, the Lead(II) sulphate is reduced to lead.
concentration of sulphuric acid is 4.1 mol dm–3
and the density is 1.3 g cm–3. When the density PbSO4(s) + 2e– → Pb(s) + SO42–(aq) ... reduction
of sulphuric acid drops from 1.3 to 1.1 g cm–3, white grey
the accumulator must be recharged.
(b) At the positive terminal
5 When an accumulator is recharged, the direct Lead(II) sulphate is oxidised to lead(IV)
current is passed through it, in the direction
which is opposite to the discharge. oxide.

6 The reactions that occur when the accumulator PbSO4(s) + 2H2O(l) → PbO2(s) + 4H+ + SO42–(aq)
is recharged are as follows. white brown + 2e– ...oxidation

Oxidation and Reduction 438

(c) The overall reaction during recharging is Figure 3.39 Electrolysis of molten sodium
2PbSO4(s) + 2H2O(l) chloride

recharging


Pb(s) + PbO2(s) + 4H+(aq) + 2SO42–(aq)
2H2SO4(aq)

6 ’05 3 In the electrolytic cell, the electrode connected

A chemical cell is shown below. to the positive terminal of a chemical cell is
called the anode. Conversely, the electrode
3
connected to the negative terminal of the
Which of the following occur in the chemical cell? chemical cell is called the cathode.
4 (a) For both the chemical and electrolytic
I The iron rod becomes thinner.
cells,
II The copper rod becomes thicker. (i) oxidation occurs at the anode,
(ii) reduction occurs at the cathode.
III The intensity of the blue colour in beaker 1
(b) In the electrolytic cell, the anions from
decreases. the electrolyte donate the electrons and is
oxidised at the anode.
IV The concentration of iron(II) ions (Fe2+) in
(c) In the electrochemical cell, the more
beaker 2 decreases. electropositive metal is oxidised at the
anode and donate the electrons to the
A I and III only C I, II and III only
cathode.
B II and IV only D I, II, III and IV
(d) Table 3.19 compares the chemical cell and

the electrolytic cell in terms of oxidation

and reduction.

Comments Table 3.19 Comparison between chemical cell and
At the iron rod: electrolytic cell in terms of redox reactions

Fe(s) → Fe2+(aq) + 2e– Chemical cell (for Electrolytic cell
example, Daniell cell) (for example, electrolysis
Therefore, the iron rod becomes thinner and the
concentration of Fe2+ in FeSO4 solution increases. of molten sodium
At the copper rod: chloride)

Cu2+(aq) + 2e– → Cu(s) At the anode: oxidation At the anode: oxidation

Therefore, the copper rod becomes thicker and occurs occurs
the intensity of blue colour decreases as the
concentration of Cu2+ decreases. Zn(s) → Zn2+(aq) + 2e– 2Cl–(l) → Cl2(g) + 2e–

Answer C At the cathode: reduction At the cathode: reduction

occurs occurs

Compare and Contrast Electrolytic Cells and Cu2+(aq) + 2e– → Cu(s) Na+(l) + e– → Na(s)
Chemical Cells in Terms of Redox Reactions
5 However, electrodes in chemical and electrolytic
1 Redox reactions occur in both the chemical cells have different signs (positive or negative)
and electrolytic cells. as shown in Table 3.20. For example, the
anode in a chemical cell is the negative electrode
2 Figure 3.39 shows the arrangement of apparatus whereas the anode in an electrolytic cell is a
for the electrolysis of sodium chloride in molten positive electrode.
condition.

439 Oxidation and Reduction

Table 3.20 Positive and negative terminals for 6 For both chemical and electrolytic cells, anions
chemical and electrolytic cells move to the anode while cations move to the
cathode. Figure 3.40 shows the movement of
Chemical cell Electrolytic cell ions in a chemical cell. Cations (Zn2+ and Cu2+)
move to the cathode and the anions (SO42–)
Anode: negative terminal Anode: positive terminal move to the anode.

In the chemical cell, In the electrolytic cell, Figure 3.40 The movement of cations and anions
in a chemical cell.
electrons are released at the electrons flow out from

anode. Thus, the anode is the anode to the battery.

the negative terminal. Thus, the anode is the

positive terminal.

Cathode: positive terminal Cathode: negative terminal

In the chemical cell, In the electrolytic cell,

electrons are removed from electrons flow from the

the cathode by positive ions battery and enter the

present in the electrolyte. cathode. Hence, the cathode

Hence, the cathode is the is the negative terminal.

positive terminal.

3 7 ’07

A chemical cell is shown below. B The electrolyte is hydrochloric acid.
C The positive terminal is lead plate.
Which of the following is true about the chemical cell? D The cell is non-rechargeable.
A At the positive terminal, lead(IV) oxide is reduced to
Pb2+ ions. Comments
The electrolyte is sulphuric acid. The positive terminal
(cathode) is the plate coated with PbO2. At the anode,
reduction occurs:

PbO2 + 4H+ + 2e– → Pb2+ + 2H2O
Answer A

3.4 (b) What is the energy conversion in this
experiment?
1 Explain the redox reactions at the anode and cathode
when electric current is passed into the following 3 Nickel-cadmium cell (Ni-Cd battery) is a rechargeable
solutions. battery. It consists of
(a) Concentrated potassium iodide solution with Anode(–): Cadmium
carbon electrodes. Cathode (+): Nickel(IV) oxide
(b) Copper(II) sulphate solution with copper electrodes. Electrolyte: Potassium hydroxide
The overall equation for the reaction that occurs
2 The following figure shows the arrangement of when Ni-Cd battery is supplying current is
apparatus for the electrolysis of iron(II) sulphate.
Cd + NiO2 + 2H2O → Cd(OH)2 + Ni(OH)2
(a) Describe (i) the oxidation and reduction processes,
(ii) the transfer of electrons that occur at the (a) Identify the oxidising and reducing agents in this
carbon electrodes. reaction.

(b) Write the half-equation for the reaction that
occurs at (i) the anode, (ii) the cathode.

(c) State the direction of flow of electrons in the
external circuit.

Oxidation and Reduction 440

Comparison between electrolytic cell and chemical cell

Electrolytic cell Chemical cell

• Cathode (–) • Anode (+) • Cathode (+) • Anode (–)

– Reduction occurs – Oxidation occurs – Reduction occurs – Oxidation occurs

– Cations accept electrons – Anions release electrons – Accepts electrons – Releases electrons

from the cathode at the anode

Redox reactions in electrolytic and chemical cells

Electrolysis of Chemical cell Negative terminal Positive terminal (reduction)
concentrated NaCl(aq) • Daniell cell (oxidation)
• Dry cell
• Anode: Carbon Zn(s) → Zn2+(aq) + 2e– Cu2+(aq) + 2e– → Cu(s)
Cathode: Carbon • Lead-acid
accumulator Zn(s) → Zn2+(aq) + 2e– 2NH4+(aq) + 2e– → 2NH3(g) + H2(g) 3
• At the anode: 2MnO2(s) + H2(g) → Mn2O3(s) + H2O(l)
Oxidation occurs Pb(s) → Pb2+(aq) + 2e– PbO2(s) + 4H+(aq) + 2e– →
2Cl– → Cl2 + 2e– Pb2+(aq) + SO42–(aq) → Pb2+ + 2H2O(l)
PbSO4(s) Pb2+(aq) + SO42–(aq) → PbSO4(s)
• At the cathode:
Reduction occurs
2H++ 2e– → H2

3.5 Appreciating the Ability oxidation
of the Elements to
Change their Oxidation
Numbers 0 +2

Various Applications of the Changes of +F3e 2O3(s) + 3C(s) → 2F0e(s) + 3CO(g)
Oxidation Numbers in Substances
reduction
1 Groups 1 and 2 elements in the Periodic Table
have fixed oxidation numbers of +1 and +2 4 In the corrosion of iron, the following changes
respectively. However, most elements (metals in oxidation numbers occur.
and non-metals) have variable oxidation
numbers. oxidation

2 The changes in the oxidation number of a +3
substance can be applied in the following 0
processes:
(a) Extracting metal from its ore 4Fe + 3O0 2 + 2xH2O → 2Fe2–O23.xH2O
(b) Corrosion of metal
(c) Preventing corrosion of metal
(d) Generation of electricity by cells reduction
(e) Recycling of metals
5 The following chemical changes occur when
3 In the extraction of iron from its ores, the zinc is used in the prevention of rusting.
changes in the oxidation numbers of both iron
and carbon are shown as follows. Zn → Zn2+ + e– ... (1)

O2 + 2H2O + 4e– → 4OH– ... (2)

The oxidation number of zinc changes from 0
to +2 while the oxidation number of oxygen
changes from 0 to –2.

441 Oxidation and Reduction

6 When a Daniell cell is used to generate electricity, Methods for Rust Prevention
the overall cell reaction is
Table 3.21 shows a summary of the common
Zn + Cu2+ → Zn2+ + Cu methods used for the prevention of metal corrosion,
especially the rusting of iron.
The oxidation number of zinc changes from 0
to +2 while the oxidation number of copper Table 3.21 Methods for rust prevention
changes from +2 to 0.
Method Objects
The Occurrence of Various Ores in Our Country
1 Paint Big objects like motor
1 Gold vehicles, ships and
Gold mines are found in Pahang (Lipis, Raub, steel bridges
Jerantut), Terengganu (Mandi), Sabah (Merungin
and Paginatan) and Sarawak (Bau, Serian and 2 Oil and grease Tools and machine
Lund). parts

2 Iron, bauxite (Al2O3) and ilmenite (FeTiO3) 3 Phosphoric acid The bottom (chassis)
Iron mines are found in Johor (Kota Tinggi), (H3PO4) of cars

Kedah (Semeling), Pahang and Perak. Bauxite 4 Galvanising (zinc- Buckets, ‘zinc’ roof
is aluminium ore and is found in Johor (Teluk plating)
Ramunia). Ilmenite is a titanium ore and is
3 found in Terengganu and Pulau Pinang. 5 Tin-plating Food cans
3 Tin
Tin ores are found in Perak (Lembah Kinta) 6 Chrome-plating Taps, bicycle handle
and Selangor (Lembah Langat). bars, car bumpers
4 Coal
Coal mines are found in Sarawak (in the areas 7 Block of magnesium Underground pipes,
of Kapit, Lucky Selantek and Sri Aman). or zinc (sacrificial ships
5 Kaolin and barite (barytes) metals)
Kaolin is a type of clay used for making
ceramics. Kaolin is found in Johor (Mersing) 8 Stainless steel Cutlery, surgical
and Perak (Bidor-Tapah). Barite (BaSO4) is the instruments
chief ore of barium. It is found in Kelantan
(Gua Musang) and Terengganu. Chemical Cell as a Source of Electrical
Energy
The Contribution of the Metal Extraction
Industry in Enhancing the National 1 Chemical cells are alternative sources of
Economy renewable energy. Chemical cells are also known
as batteries. Nowadays, there are many types
1 The metal extraction industry provides job of chemical cells. Examples are: zinc-carbon
opportunities and lowers the unemployment battery, mercury battery, alkaline battery, lithium
rate. The export revenue from tin and other battery, nickel-cadmium (Ni-Cd) battery, lead-
metals enhances the national economy. acid accumulator, fuel cell and photo/solar
battery.
2 Our country exports tin ingots to countries
like Japan, America and Britain. This will earn 2 A fuel cell is a device in which the fuel is
foreign exchange which can be used for oxidised in a chemical cell so as to produce
national development. electricity directly. In the hydrogen-oxygen fuel
cell, chemical energy from the redox reaction
3 The metal extraction industry produces metals between hydrogen and oxygen to form water
as raw materials for many other industries, such is used to generate electric current.
as the motor and construction industries. Fuel cells are used widely in spacecraft. The
water produced can be used for drinking. Fuel
cells are also used to power electric cars. Fuel
cells differ from the usual chemical (voltaic)
cells in two ways:

Oxidation and Reduction 442

(a) The fuel and oxygen are fed into the cell strikes the cell, the electrons flow from the
continuously, donor to the acceptor crystal. Solar cells have
been used for years to power spaceships.
(b) The electrodes are made from inert material 4 Research is still being carried out on chemical
such as platinum that does not react during cells to develop their potential as an alternative
the process. energy source, which is now produced chiefly
by the burning of fossil fuels.
3 Solar cells (also known as photovoltaic cells) 5 One of the industries that is actively engaged
convert sunlight directly into electricity. Solar cells in developing a light, powerful, efficient and
are made from extremely pure silicon crystals. long-lasting battery is the automobile industry.
One type of crystal has about 1 ppm (parts per In the near future, electric cars that use battery
million) of arsenic added to it. This crystal is to power, instead of using petrol as the fuel
called the donor crystal. Another crystal is made will be on the road. If that happens, air pollution
by adding about 1 ppm of boron. This crystal is would be greatly reduced.
called the acceptor crystal. The two crystals are
connected by an external circuit. When sunlight

3

1 Oxidation is 11 A more reactive metal will displace a less reactive
• gain of oxygen, or metal from its oxide. For example, if P has a
• loss of hydrogen, or reaction with the oxide of Q, P is more reactive than
• loss of electrons or Q.
• increase in the oxidation number of the element.
P + oxide of Q → Q + oxide of P
2 Reduction is
• loss of oxygen, or 12 The reactivity series that includes both carbon and
• gain of hydrogen, or hydrogen is
• gain of electrons or
• decrease in the oxidation number of the element. K, Na, Ca, Mg, Al, C, Zn, H, Fe, Sn, Pb, Cu, Ag

3 An oxidising agent is a substance that causes ⎯⎯⎯⎯⎯ reactivity decreases ⎯⎯⎯⎯→
oxidation in another substance.
13 Reactive metals such as K, Na, Ca, Mg and Al are
4 A reducing agent is a substance that causes extracted from their ores by electrolysis.
reduction in another substance.
14 Metals such as Zn, Fe, Sn and Pb are extracted from
5 Oxidation and reduction take place simultaneously their ores by heating the metal oxides with carbon.
in a redox reaction.
15 The reactions that occur in electrolytic cells or
6 A displacement reaction is a redox reaction. chemical cells (voltaic cells) are redox reactions
A more electropositive metal will displace a less involving the transfer of electrons.
electropositive metal from its salt solution.
16 In an electrolytic cell:
7 In a redox reaction, electrons are transferred from • At the anode (positive electrode), oxidation occurs
the reducing agent to the oxidising agent. and anions are discharged by losing electrons to
form molecules.
8 Rusting is a redox reaction. For iron to rust, oxygen • At the cathode (negative electrode), reduction
(air) and water must be present. occurs and cations are discharged by gaining
electrons to form metal or hydrogen gas.
9 During rusting, iron reacts with oxygen and water to
form a brown substance, called rust (Fe2O3. xH2O). 17 In a chemical cell:
• The more electropositive metal is the negative
4Fe + 3O2 + 2xH2O → 2Fe2O3.xH2O electrode and the less electronegative metal is the
positive electrode.
10 Rusting can be prevented by • Oxidation occurs at the anode (negative electrode).
• a protective layer (oil, grease or plastic layer), • Reduction occurs at the cathode (positive
• coating/plating iron with tin/chromium electrode).
• using sacrificial metals
• using alloys

443 Oxidation and Reduction