EXTRACTION OF METALS

METALS ARE EXTRACTED BY ELECTROLYTIC REDUCTION AND CHEMICAL REDUCTION!

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 EXTRACTION OF METALS
INTRODUCTION TO METALS
 A metal is an element that readily loses electron(s) to form a positive ion.
 For example
Na → Na+ + e- (univalent)
Mg → Mg2+ + 2e- (divalent)
Al → Al3+ + 3e- (trivalent)
MINERALS
 A mineral is a naturally occurring inorganic solid, with a definite chemical composition and
an ordered atomic arrangement. OR
 A mineral is the natural compound of a metal and non-metal.
 Metal compounds occur in the Earth’s crust along with some impurities such as rock particles.
ORES
 These are minerals that contain large proportions of metals or metallic compounds worth to
be extracted.
 Metals exist in nature either as free elements or as compounds.
 Some of the metals that exist as free elements include silver, gold and platinum. These
elements are the least reactive, hence do not form compounds as others.
 The compounds of metals can be oxides, carbonates, sulphates, or complex compounds
made from two or more simple compounds.
IMPORTANCE OF METALS
i. Many materials such as mobile phones, computers, electric wires, nails, planes, cooking
pots, rails, door locks, and earrings are made of metals.
ii. Metals are used in different areas such as; in agriculture, industries, automobiles, security
systems, drilling and construction, buildings and bridges.
iii. Metals are also used in electronics and communication, furniture, jewellery, clothing,
electric lightening, medicine, and medical instruments.
WHY ARE METALS EXTRACTED?
 Most metals are found in the Earth’s crust combined with other elements in rocks known
as ores. For example, iron is found combined with oxygen in ores called haematite and
magnetite.
 Metals need to be extracted from ores before they can be turned into useful products,
such as cars or cutlery.

OCCURRENCE AND ABUNDANCES OF METALS IN THE EARTH’S CRUST

There are different elements in the Earth’s crust.
i. Aluminium makes about 8% of the Earth’s crust and is the most abundant metal, followed
by iron which constitutes about 6%.
ii. The abundances of calcium, sodium, potassium, and magnesium in the Earth’s
crust are 3.6%, 2.8%, 2.0%, and 2.3%, respectively.
iii. Other metals found in low abundances in the Earth’s crust include the following: titanium
(0.44%), manganese (0.095%), barium (0.043%), chromium (0.61%), nickel
(0.01%), copper (0.006%), zinc (0.007%), cobalt (0.003%), lead (0.0013%), uranium
(0.0003%), gold (<5 × 10−6%), mercury (<8 × 10−6%), and silver (<7 × 10−6%).
iv. The remaining metals make about 3% of the Earth’s crust.

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OCCURRENCE AND LOCATIONS OF METALS IN TANZANIA
 Tanzania is among the countries that have abundant mineral resources.
 The minerals found in Tanzania are categorised into metallic, gemstones, industrial, energy
source, and construction minerals.
 The metallic minerals include gold, iron, silver, cobalt, aluminium, uranium, copper, and
nickel.
 Gemstones include tanzanite, diamond, garnet, and ruby.
 Examples of industrial minerals found in Tanzania are gypsum, limestone, phosphates,
soda ash, and salt.
 Energy source minerals include coal and uranium.
 Gravel, dimension stones, and sand are in the category of construction minerals.
METALS LOCATION IN TANZANIA
Gold Geita, Kahama, Buhemba (Butiama), Biharamulo, Nyamongo (Tarime),
Majimoto (Ngara), Mananila, Magambazi and Lupa (Chunya), Nzega,
Handeni, and Ikungi.
Iron Liganga (Ludewa and Mbinga), Uluguru mountains, Mbabala near Lake
Tanganyika, Karema, Manyara, and Itewe.
Nickel Kabanga (Ngara)
Aluminium Mabughai-Mlomboza (Lushoto)
Copper Mpanda
Uranium Mkuju (Namtumbo), Bahi, Galapo, Minjingu, Mbulu, Simanjiro, Manyoni,
Songea, Tunduru, Madaba, and Nachingwea.
Salt Uvinza
Tanzanite Merelani
Diamond Mwadui
Ruby Mahenge (Ulanga).

WHY ARE SOME AREAS MINED AND NOT OTHERS?

 Mining for ores is expensive and so is only carried out where minerals are abundant
enough for this to be profitable.
PROPERTIES OF METALS
 Pure metals have specific physical and chemical properties.
 Their physical properties include physical strength, colour, melting point, boiling point,
and solubility.
 The chemical properties include their ability to react with other substances.
PHYSICAL PROPERTIES OF METALS
i. They have a shiny lustre, especially when freshly cut.
 The lustre of a sharpened end of an axe and knife exemplify this property.
ii. They are ductile, which means they can be stretched into thin wires or rollers.
DUCTILITY
 This is the ability of metals to be made into thin wires.
 Copper wires, tungsten wires and iron rollers are examples of such thin materials
made due to the ductility of metals.
iii. They are malleable as they can be hammered into thin sheets without breaking.
MALLEABILITY
 This is the ability of metals to be hammered into thin sheets without breaking.
 The corrugated iron sheets are examples of such sheets from a malleable metal.

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iv. They are good conductors of heat and electricity.
CONDUCTIVITY
 This is the ability of metals to conduct heat and electricity.
 They can be used in making cables for conducting electricity.
v. They are solids with high melting points with the exception of mercury which is a liquid
metal.
 The high melting points of metals are due to strong forces of attraction among the
metal atoms.
vi. Some metals are sonorous.
 A sonorous substance is the one that is capable of producing a deep ringing sound
when hit.
 For a metal to suit as a construction material for a bell, it must be sonorous.
 Examples of sonorous metals are aluminium, iron, and silver.
SONORITY
 This is the ability of a metal to produce ringing sound when hit.
vii. They are strong and tough with high tensile strength.
TENSILE STRENGTH
 This is a resistance of a material to breaking under tension. It includes the resistance
against bending.
 For a metal to qualify as a building material for bridges, it must have a high tensile
strength.
CHEMICAL PROPERTIES OF METALS
i. Metals are electropositive elements.
ii. In chemical reactions, metals donate electrons to other elements.
 The tendency of an element to lose electron and form a positive ion is called
electropositivity.
iii. In solutions, metals form cations allowing them to become stable.
CHEMICAL STRENGTHS VERSUS PHYSICAL STRENGTHS OF METALS
i. PHYSICAL STRENGTH
 The physical strength of a metal is the ability to withstand the force applied on it without
breaking.
 The physical strength of a metal is referred to as its tensile strength.
 Iron has a very high tensile strength, while sodium has a low tensile strength.
 Metals which are physically strong are not necessarily chemically strong, and metals
which are chemically strong are not necessarily physically strong.
ii. CHEMICAL STRENGTH
 The chemical strength of a metal is a measure of how readily the metal takes part in
chemical reactions.
 Some metals are very reactive hence react rapidly to produce new substances.
 Potassium and sodium are some of the most reactive metals.
 Metals such as platinum, gold and silver are very unreactive. Such metals do not release
their electrons easily.
 The reactivity of metals has potential roles towards the methods that are used in their
extraction.

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REDOX REACTIONS INVOLVING METALS
 REDOX REACTIONS these are reactions in which oxidation and reduction take place
simultaneously.
 OXIDATION is a process which involves the loss of electron(s) from a substance. For
example, sodium metal loses an electron to form sodium ion.
Na(s) Na+(g) + e-
 In this reaction, sodium has been oxidised to sodium ion. An equation which represents the
loss of electron(s) is referred to as an oxidation half reaction.
 REDUCTION involves the gaining of electron(s). For example, chlorine gas can be reduced
to chloride ions by gaining electrons.
Cl2(g) + 2e- 2Cl-(g)
 An equation that represents a reduction reaction is referred to as a reduction half
reaction.
 A redox reaction is therefore made up of the two half reactions, one representing oxidation
and the other representing reduction.
Consider the reaction
2Na(s) + Cl2(g) 2NaCl(s)
This is a redox reaction and its half reactions are:

REDUCING AGENT
 The substance that loses electrons.
 For example from the above equation sodium is a reducing agent.
OXIDIZING AGENT
 This is a substance that gains electrons.
 For example from the above reaction, sodium is an oxidising agent.
THE REDUCING POWERS OF METALS
 In chemical reactions, metals tend to lose their outermost shell electron(s).
 The ease with which metals lose electrons increases down the group within the Periodic
Table.
 Potassium with an electronic configuration of 2:8:8:1 is a stronger reducing agent than
lithium which has an electronic configuration of 2:1.
 This is because the number of electron shells increases down the group causing the
outermost electrons to be weakly attracted to the nucleus.
 The outermost electrons are therefore easily lost during chemical reactions.
 The reducing powers of metals thus increase down the group.
 Across the period, that means from left to right, the number of electrons in the outermost
shell increases.
 The number of protons also increases because the number of protons is equal to the number
of electrons in a neutral atom. However, the number of shells does not change. This results
in stronger attractions of the electrons in the outermost shell to the protons in the nucleus.

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REACTIVITY SERIES OF METALS
 The reducing power of a metal determines its reactivity.
 The higher the reducing power, the more reactive the metal is.
 The arrangement of metals according to their reducing powers or reactivity is known as the
reactivity series of metals.
 At the top of the reactivity series are metals with the highest reducing powers.

DISPLACEMENT REACTIONS INVOLVING METALS

 A displacement reaction takes place when an atom or group of atoms takes the position of
another in a compound.

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 Examples
i. Consider the reaction between iron metal and hydrochloric acid to form iron(III)
chloride.

ii. When zinc is added to copper(II) sulphate solution and stirred, a brown solid is formed.
The beaker becomes warm indicating that the reaction is exothermic. The brown solid
is copper metal.
𝐙𝐧(𝐬) + 𝐂𝐮𝐒𝐎𝟒(𝐚𝐪) → 𝐙𝐧𝐒𝐎𝟒(𝐚𝐪) + 𝐂𝐮(𝐬)
iii. Iron will displace copper from its salt.
Fe(s) + CuSO4(aq) → Cu(s) + FeSO4(aq)

CARBON AND HYDROGEN IN THE REACTIVITY SERIES

 Carbon and hydrogen are included in the reactivity series although they are non-metals
because they are reducing agents just like metals.
 Carbon is not a metal, but its characteristics lie between metals and non-metals. Based
on its reducing power, carbon is placed between aluminium and zinc in the reactivity series.
 Metals above carbon in the reactivity series are usually extracted by electrolysis.
 Metals below carbon can be extracted by reducing their oxides using carbon or carbon
monoxide.
 For example, charcoal (carbon) is heated with copper(II) oxide to produce brown specks of
copper metal.
 White specks of lead metal are also formed when lead(II) oxide is heated under carbon.

METHODS USED IN EXTRACTION OF DIFFERENT METALS

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STUDY QUESTIONS
1. A metal has both physical and chemical strengths. What does this statement imply?
2. With examples, explain any four chemical properties of metals.
3. When a piece of magnesium ribbon is added into a solution of dilute hydrochloric acid,
some gas bubbles are observed, but when a piece of copper wire is used instead of
magnesium, no bubbles are observed. Explain.

EXTRACTION OF METALS
 This is the process in which a metal is separated from its ore.
 Metals are extracted from their ores.
AN ORE
 This is a naturally occurring solid material from which a metal or valuable mineral is
extracted profitably.
METALLURGY
 This is a branch of science that deals with the properties, production, extraction and
purification of metals.
MINING
 This is the process of obtaining minerals or ores from the earth’s crust.

PROCESSES/STAGES FOR EXTRACTION OF METALS

The main processes involved in the extraction of metals from their ores are:
a. Crushing and grinding.
b. Concentration or dressing of the ore.
c. Calcination or roasting of the ore.
d. Reduction of a metal oxide to a free metal.
e. Purification and refining of the metal.

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1. CRUSHING AND GRINDING
 Metal ores are generally obtained as large pieces of rocks.
 The metal ore has to be crushed and ground (pulverised) to convert the components into
powder.
 This is done by using crushers and grinders.
2. CONCENTRATION OR DRESSING OF THE ORE
 This is the process of removing impurities from the metal ore.
 Ores are usually found mixed with earthy impurities like sand, clay, and limestones.
 The impurities in the ore are called gangue or matrix.
 This helps to obtain the materials richer in the metal or those worth working up with.
METHODS USED IN THE CONCENTRATION OF THE ORE
i. Gravity separation
 This method is based on the differences in weight between the impurities and the
metal.
 Impurities are lighter than the metal particles, the principle of gravity is used to
separate them.
 In this method, the light earthy impurities are removed from the heavier metallic
ore particles by washing with water.
ii. Magnetic separation
 In this method, the magnetic properties of certain metals are used to separate them
from the non-magnetic impurities.
iii. Froth flotation
 This is the process for selectively separating hydrophobic materials from hydrophilic
materials.
 It uses the difference in the wetting characteristics between minerals and impurities.
 Froth flotation is primarily used for concentrating sulphide ores.
 Sulphide ores are hydrophobic in nature and are easily wetted by oil.
 Thus, sulphide ores can be selectively separated from hydrophilic impurities by froth
flotation.

iv. Leaching
 Leaching is the chemical method of concentrating the ore.
 It involves dissolving the ore in a chemical solution in which the mineral dissolves
leaving behind the impurities which are filtered out.
 The metal ore is then recovered from the chemical solution by a suitable chemical.
 For example, the concentration of the bauxite ore for extraction of aluminium.
Bauxite is contaminated with iron, titanium, and silicon oxide.

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 These impurities are removed by heating the powdered ore with an aqueous

solution of sodium carbonate or sodium hydroxide at a certain temperature and
pressure.
 Aluminium oxide dissolves in the base, whereas the impurities remain insoluble and
are removed by filtration.
3. CALCINATION AND ROASTING OF THE ORE
 The concentrated metal ore is converted into metal oxide by either calcination or
roasting.
CALCINATION
 Calcination is a process in which the ore is heated below its melting point in the
absence of air.
 This method is commonly used for converting metal carbonates into their respective
oxides.
 Calcination is generally carried out in a reverberatory furnace or in multiple heater
furnace.
 During calcination, moisture, volatile gaseous substances, and water of hydration are
driven out. Examples of calcination include the following:

ROASTING
 Roasting involves converting the concentrated ore into a metal oxide by heating it
below its melting point in the presence of excess air.
 While calcination is mostly used for converting metal carbonates to oxides, roasting is
a method used for converting sulphide ores to their respective oxides as shown in the
following examples:
𝒉𝒆𝒂𝒕
𝟐𝑷𝒃𝑺(𝒔) + 𝟑𝑶𝟐(𝒈) → 𝟐𝑷𝒃𝑶(𝒔) + 𝟐𝑺𝑶𝟐(𝒈)
𝒉𝒆𝒂𝒕
𝟐𝒁𝒏𝑺(𝒔) + 𝟑𝑶𝟐(𝒈) → 𝟐𝒁𝒏𝑶(𝒔) + 𝟐𝑺𝑶𝟐(𝒈)

4. REDUCTION OF THE METAL OXIDES TO FREE METALS

 There are various ways by which the reduction of a metal oxide can be done.
 The methods used include the reduction by heating in the air (auto-reduction),
reduction by carbon (smelting), reduction by electrolysis (electrolytic
reduction), reduction by precipitation, and amalgamation.
 The choice of the reduction method to be used for a given metal compound depends
on the position of the metal in the reactivity series.

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 METHOD METAL COMPOUND
Reduction by using Oxides of less electropositive metals such as lead, zinc, iron, tin, and
carbon copper.
(smelting)
Reduction by Oxides and sulphides of less electropositive metals
heating in the air such as mercury, copper, and lead
(auto-reduction)
Reduction by Oxides, hydroxides or chlorides of metals high in the reactivity series
electrolysis such as sodium, calcium, magnesium, and aluminium.
Reduction by Metal ores of less reactive metals. A solution containing metal ions is
precipitation reacted with a more reactive metal. The more reactive metal displaces
the less reactive metal from the solution.
Amalgamation Less reactive metals like silver and gold. The finely crushed ore is
brought into contact with mercury which forms an alloy (amalgam) with
the metal. The metal is then recovered by distilling the amalgam.

5. REFINING AND PURIFICATION

 Metals obtained through reduction process are generally impure. They are known as crude
metals.
 Crude metals contain impurities such as other metals, non-metals like silicon and phosphorus
and unreduced oxides and sulphides of the metals.
 Purification of the crude metals is done through three main methods which are distillation,
oxidation, and electrorefining.
DISTILLATION
 Distillation involves heating the crude metal in a furnace until a pure metal evaporates, leaving
behind the impurities.
 The vapour is then collected and condensed in a separate chamber.
OXIDATION
 The molten crude metal is exposed to air in the furnace so that the impurities are oxidised
and escape as vapour or form a scum over the molten metal.
 The scum is then removed by skimming. This method is used only when the impurities have
a greater affinity for oxygen than the metal.
ELECTRO-REFINING
 Most metals are purified by electrolysis (electro-refining).
 The crude metal is moulded into blocks and made the anode of an electrolytic cell.
 The cathode is usually made up of a thin plate of the pure metal.
 The soluble impurities go into the solution, while the insoluble impurities settle down below
the anode as mud or sludge.
EXTRACTION OF METALS BY ELECTROLYTIC REDUCTION
 This is a common extraction process for the more reactive metals.
 The extraction of metals by electrolytic reduction involves the application of electric
current as a source of electrons for the reduction process.
 The electrolytic reduction is a type of electrolysis.
 The oxides, hydroxides and chlorides of metals in a fused state are electrically reduced by
using this method.

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 Some metals like sodium, potassium, magnesium, calcium, and aluminium are obtained
by this process.

1. EXTRACTION OF SODIUM
POSSIBLE ORES
 Sodium is abundant in different compounds such as sodium chloride (rock salt) and
sodium sulphate.
EXTRACTION
 Sodium is extracted by the electrolytic reduction of purified molten rock salt.
 The electrolysis is mainly carried out in the Down’s cell.
 The cell has a steel cathode and a graphite anode.
 Calcium chloride is added to sodium chloride to lower the melting point of sodium
chloride from 774°C to about 600°C.
 The composition of the electrolyte is 40% sodium chloride and 60% calcium chloride.
 During the electrolysis, chloride ions move to the anode, while sodium ions move to the
cathode.
 The reactions taking place at the electrodes are:
(a) At the anode, chloride ions are discharged as chlorine gas as shown by the following
equation:

(b) At the cathode, sodium ions are discharged as a metal as shown by the following
equation:

A diagram of Down’s cell for extraction of sodium

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  If the sodium metal and chlorine gas produced are allowed to come into contact at 600
°C, a violent reaction occurs which oxidises the metal.
 To prevent this, the cell has a steel gauze as a diaphragm around the anode to keep the
two products apart.
 The large graphite anode is used to facilitate maximum oxidation of chloride ions to
chlorine gas.
 The sodium metal is collected upwards in the Down’s cell because of its low density which
makes it float over the mixture.
 The sodium metal from the Down’s cell contains some calcium, which is also formed
through electrolysis.
 The calcium crystallises when the mixture cools and a relatively pure sodium metal is
obtained.

USES OF SODIUM
i. Is alloyed with lead in the preparation tetraethyl (IV) lead, which is added to petrol as an
anti-knock.
ii. Provides the glow in sodium vapours lamps, for street lighting (orange-yellow street lights).
iii. Is an excellent conductor of heat and electricity with low melting point hence used;
 In nuclear reactors to conduct away heat.
 Modern aeroplane engines.
 Manufacture of sodium peroxide, and sodium cyanide used in the extraction of silver
and gold.

2. EXTRACTION OF ALUMINIUM
POSSIBLE ORES
 The main ores of aluminium are bauxite (Al2O3.2H2O), cryolite (Na3AlF6),
feldspar (KAlSi3O8), and kaolin (Al2Si2O7.2H2O).
EXTRACTION PROCESS
 The chief ores from which aluminium is extracted are bauxite and cryolite.
 Aluminium is extracted from bauxite by electrolysis.
 This is done by using the Hall-Héroult method that involves the following stages:
i. Purification of bauxite
 Bauxite contains oxides of iron and silicon as impurities.
 These impurities are removed by using the Hall’s process where the powdered ore is
heated to bright red with sodium carbonate.
 Since aluminium oxide is amphoteric, it dissolves to form sodium aluminate (NaAlO2).

 The molten mass is then removed together with water, while the insoluble iron and
silicon oxides are left behind as residue.
 When the filtrate is heated at a temperature between 50°C and 60°C and a stream of
carbon dioxide is passed through the mixture, precipitation of aluminium hydroxide
occurs.

 Aluminium hydroxide is filtered off, washed, dried, then calcinated at 1500°C to give
alumina.

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ii. Electrolysis of alumina
 Alumina (Al2O3) cannot be reduced to aluminium metal by carbon because aluminium
has a greater affinity to oxygen than carbon.
 The electrolysis of alumina is carried out in the presence of cryolite (Na3AlF6) which
improves the electrical conductivity of alumina, which is otherwise a bad conductor of
electricity.
 Cryolite also lowers the melting point of aluminium oxide. This is very important
because if electrolysis is carried out at very high temperatures, the aluminium formed
will vapourise.

A diagram of an electrolytic cell for electrolysis of Alumina

 The electrolytic cell for the electrolysis of alumina consists of an iron bath lined with
graphite (carbon) which acts as the cathode.
 The anode consists of a number of carbon rods attached to copper clamps and
suspended in the fused mass.
 The carbon rods are arranged in such a way that they can be raised, lowered or
replaced when needed.
 The reactions taking place at the electrodes are as shown in the following equations:

 The production of aluminium by electrolysis is an expensive process. The temperature

of the fused electrolyte has to be maintained between 800 °C and 900 °C.
 This is achieved by passing a steady current of about 100,000 ampere through the
electrolyte.

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  This makes aluminium expensive despite the fact that it is the most abundant metal on
Earth.
 The process is only economical when electricity is cheap and readily available.
iii. Refining of crude aluminium
 The aluminium obtained from the Hall-Héroult process is 99% pure. It contains small
quantities of iron, silicon, alumina and carbon.
 It is purified further by electrolysis using Hoope’s electrolytic cell.
 The Hoope’s electrolytic cell is made up of an iron tank lined on the inside with carbon
which serves as the anode. The tank has three different layers of molten substances:
(a) The bottom layer contains impure molten aluminium containing copper and
silicon to increase the density.
(b) The middle layer contains a mixture of fluorides of sodium, barium, and
aluminium in a fused form. This layer acts as an electrolyte.
(c) The top layer contains pure molten aluminium which together with the carbon
rod serves as the cathode. The carbon cathodes are suspended from above
inside the cell mixture.

A diagram Hoope’s electrolytic cell for purification of Aluminium

USES OF ALUMINIUM
i. It is used to manufacture cables for electric transmission as it is a good conductor of
electricity.
ii. It is used to make household appliances like cooking utensils.
Reasons:
 It is a light metal
 It is a good conductor of heat
 It is resistant to corrosion
iii. It is used in construction and transport because it is a light metal and resistant to corrosion.
iv. It is used to make alloys with other metals like stainless steel.
v. It is used to manufacture storage cans. This is because it is a soft metal and can be molded
into any shapes. Example wrapping chocolates, medicines, caps of milk bottles.
vi. It is used in extraction of manganese and chromium from its oxide.
vii. The dusts of aluminium are used in paints.
viii. Aluminium folis are used in food industries for wrapping food.

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STUDY QUESTIONS
1. Briefly explain how sodium metal is extracted from its ore.
2. Explain what would happen in the Down’s cell if the molten sodium is mixed with chlorine at
very high temperature during the electrolytic extraction of sodium.
3. Describe the processes involved in the extraction of aluminium.
4. Comment on the extraction of aluminium from bauxite in the absence of cryolite.

EXTRACTION OF METALS BY CHEMICAL REDUCTION

 Less electropositive metals such as iron and copper can best be extracted through chemical
reduction which involves heating the ore with reducing agents.
 The common material used in reducing the ores is carbon.

1. EXTRACTION OF IRON
THE POSSIBLE ORES
 The ores of iron include haematite (Fe2O3), magnetite (Fe3O4), limonite (Fe2O3.H2O),
siderite or spathic iron ore (FeCO3), and iron pyrites (FeS2).
 The most prominent ores of iron are haematite which contains about 70% iron, and
magnetite which contains 72.4% iron.
 Although iron pyrites is abundant on the Earth’s crust, it is not used as a source of iron, but
it is mainly used in the production of sulphuric acid.
Question
Explain why iron pyrites are abundant on the earth’s crust but it is not used as
an ore for extraction of iron?
Answer
 Due to formation of excess polluting gas of sulphur dioxide because the pyrites have
high content of sulphur.
 Due to high contents of sulphur, it requires additional purification methods to
remove the impurities.
THE STAGES/PROCESS OF EXTRACTION OF IRON
 Iron is mainly extracted from the haematite ore by reduction using carbon (coke).
 Three main stages are involved in the extraction process, which are concentration,
roasting, and smelting (reduction).
i. Concentration
 The haematite ore (Fe2O3) is first washed to remove earthy impurities such as clay, sand,
and other non-magnetic impurities.
ii. Roasting
 The concentrated ore is strongly heated in excess air in shallow kilns and the following
chemical changes occur during the process.
 In addition, calcination of hydrated iron(III) oxide and iron(II) carbonate occurs.
(a) Most of the moisture is removed.

(b) Sulphur and arsenic impurities are oxidised to their oxides, which escape as gases.

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 𝐡𝐞𝐚𝐭
𝟒𝐀𝐬(𝐬) + 𝟑𝐎𝟐(𝐠) → 𝟐𝐀𝐬𝟐 𝐎𝟑(𝐠)
(c) Carbonates are decomposed into oxides giving off carbon dioxide gas. For example,

(d) Iron(II) oxide is oxidised to iron(III) oxide

𝒉𝒆𝒂𝒕
𝟒𝑭𝒆𝑶(𝒔) + 𝑶𝟐(𝒈) → 𝟐𝑭𝒆𝟐 𝑶𝟑(𝒔)
 This reduces the loss of iron, since iron(II) oxide has a tendency of forming a slag with
silicon dioxide (sand).
𝒉𝒆𝒂𝒕
𝑭𝒆𝑶(𝒔) + 𝑺𝒊𝑶𝟐(𝒔) → 𝑭𝒆𝑺𝒊𝑶𝟑(𝒍)
iii. Smelting
 The roasted ore is reduced (smelted) by carbon in the blast furnace.
 The BLAST FURNACE is a metallurgical furnace used for smelting to produce metals such
as iron, lead, and copper.
 The structure of the blast furnace for iron smelting is about 30m high and 10m in diameter
at its widest part.
 The inside lining of the furnace is made up of firebrick which is resistant to high
temperature.
 The ore is mixed with carbon (coke) and limestone (CaCO3) which are introduced into
the furnace from the top.
 Blasts of hot air are blown into the furnace through small openings near the base known
as tuyéres.

A diagram of blast furnance for extraction of iron

17
ZONES OF THE BLAST FURNANCE
 The blast furnace is divided into three zones, according to the temperature profile in
the furnace.
i. The upper zone of reduction
 This zone is near the top of the furnace with a temperature range of 300°C−800 °C.
 The reactions which take place in this zone depend directly on the reactions that occur
in the lower and middle zones. In the lower zone, the carbon burns in oxygen of the
hot air to form carbon dioxide.
𝒉𝒆𝒂𝒕
𝑪(𝒔) + 𝑶𝟐(𝒈) → 𝑪𝑶𝟐(𝒈)
 The carbon dioxide formed reacts with more coke to form carbon monoxide at the
middle zone.

 The carbon monoxide rises up to the upper zone of the furnace for reducing iron(III)
oxide to spongy iron.

 It should be noted that, the main reducing agent in the blast furnace is carbon
monoxide and not carbon.
 The metal obtained is called spongy iron because the temperature in this zone is too
low to melt the iron.
 As the iron ore is being reduced, part of the limestone decomposes at 600 °C to
quicklime (CaO).

18
  The limestone also reacts with coke to form quicklime.

ii. In the middle zone

 This is the lower zone of reduction with the temperature range of 900°C ̶ 1200°C.
 The following are the reactions that take place in this zone:
(i) Carbon dioxide is reduced by coke to carbon monoxide.

(ii) At about 1000 °C, calcium carbonate is almost completely decomposed

to quicklime (CaO)
𝟏𝟎𝟎𝟎℃
𝑪𝒂𝑪𝑶𝟑(𝒔) → 𝑪𝒂𝑶(𝒔) + 𝑪𝑶𝟐(𝒈)
(iii) The calcium oxide acts as a flux and combines with the silica (silicon
dioxide) to form calcium silicate (slag).
A flux is a mixture of chemicals that react with impurities to form slag.
𝟏𝟎𝟎𝟎℃
𝑪𝒂𝑶(𝒔) + 𝑺𝒊𝑶𝟐(𝒔) → 𝑪𝒂𝑺𝒊𝑶𝟑(𝒍)
iii. The lower zone of combustion
 This is the lowest and the hottest part of the furnace. The temperature range at this zone
is 1200 °C ̶ 1500 °C. The following reactions take place in this zone:
i. Coke burns in the blast of hot air to produce carbon dioxide.
 This reaction is exothermic, therefore, it raises the temperature of the
furnace.

ii. The spongy iron melts in this zone at about 1300 °C and collects at the bottom of the
furnace.
iii. Iron(II) oxide which might have escaped the reduction process in the middle zone is
reduced by coke in this zone.

 The molten iron sinks at the bottom of the furnace, while the less dense fusible slag floats
over the molten iron forming a separate layer.
 This layer of slag prevents iron from being oxidised by hot air.
 The slag and iron are periodically removed through different outlets.
 The mixture of waste gases containing nitrogen, carbon dioxide, and carbon monoxide are
known as blast furnace gases.
 This mixture of gases is burnt in air to produce heat which is used for pre-heating the air
blast coming in through the tuyéres.
 The iron obtained through the blast furnace is an impure type known as pig iron.
 Pig iron is further purified by re-smelting it using coke and lime in another furnace called
cupola. The molten iron from the cupola is poured into moulds of desired shapes. The
pure iron obtained is called cast iron.
TYPES OF IRON
i. Cast iron or pig iron
 This is the impure iron which contains varying amount of impurities like carbon.

19
  It melts at lower temperature than pure iron and it is brittle.
 It cannot be welded and posses little tensile strength.
 It is used to make railings, hot water pipes, bunsen burner bases and ironing box.
ii. Wrought iron
 This is the purest form of iron.
 It is produced from cast iron by heating it with iron (III) oxide in a furnance by a process
of puddling.
 It has higher melting point than cast iron.
 It is malleable and can be forged, hammered and welded when hot.
 It is used to make iron nails, shhetings, horse shoes, agricultural tools.
iii. Steel
 This is a material containing iron and small proportion of carbon.
 Over 90% of pig iron is converted into steel.
 It is hard, tough and strong.
 It is used to make girders, wire mesh, cutting and boring tools, crushing machinery,
stainless cutlery (cutting tools)
USES OF IRON
i. In form of steel it used to make bridges, machinery, T.V towers and alloys.
ii. Pig iron is used in making pipes, stoves, radiators, railings and drain pipes.
iii. Wrought iron is used to make springs, anchors and electromagnets.
iv. Steel is used to make different tools like hoes, racks also nuts and bolts.
CHEMICAL PROPERTIES OF IRON
i. It reacts with oxygen to form iron (III) oxide.
2Fe(s) + 2O2(g) → 2Fe2O3(s)
ii. It reacts with sulphur to form black iron (II) sulphide
Fe(s) + S(s) → FeS(s)
iii. It reacts with chlorine to form iron (III) chloride
2Fe(s) + 3Cl2 (g) → FeCl3(s)
iv. It reacts with hydrogen chloride gas to form iron (II) chloride
Fe(s) + 2HCl(g) → FeCl2(s) + H2(g)
v. Reacts reversibly with steam to form hydrogen and triirontetroxide
3Fe (s) + H2O(g) → Fe3O4(s) + 4H2(g)
vi. It reacts with dilute acid to liberate hydrogen
Fe(s) + 2H+(aq) → Fe2+(aq) + H2(g)

2. EXTRACTION OF COPPER
POSSIBLE ORES
 Copper occurs in both free and combined states.
 Most copper occurs in combined forms, mainly as sulphides and oxides.
 The main ores of copper are copper pyrites (CuFeS2), copper glance (Cu2S),
cuprite (Cu2O), malachite [Cu2CO3(OH)2] or [CuCO3.Cu(OH)2] and azurite
[Cu3(CO3) 2(OH)2] or [2CuCO3.Cu(OH)2].

20
STAGES/PROCESSES FOR EXTRACTION OF COPPER
 Copper is mainly extracted from copper pyrites by using the dry process.
 There are five main stages that are involved in the extraction process, namely;
concentration, roasting, smelting, conversion, and refining.
i. Concentration
 The powdered ore is concentrated by froth flotation or other processes to remove nonore
impurities such as galena (PbS).
ii. Roasting
 The concentrated ore is heated strongly in the presence of excess air in a special furnace
where it decomposes to form copper(I) and iron(II) sulphides.
𝐡𝐞𝐚𝐭
𝟐𝐂𝐮𝐅𝐞𝐒𝟐(𝐬) + 𝐎𝟐(𝐠) → 𝐂𝐮𝟐 𝐒(𝐬) + 𝟐𝐅𝐞𝐒(𝐬) + 𝐒𝐎𝟐(𝐠)
 Part of the sulphides may be oxidised to their respective oxides. The reactions are shown
in the following equations:

 As the roasting proceeds, Impurities such as sulphur, arsenic, and antimony are removed
in the form of their volatile oxides. For example, sulphur is oxidised as shown in the
following equation:

 The roasting process results in a mixture of sulphides and oxides of copper and iron.
iii. Smelting
 The roasted ore is mixed with coke and sand in the presence of excess air in a water
jacketed blast furnace.
 In the blast furnace, the oxidation of iron(II) sulphide which started during the roasting
process continues as shown in the following equation.
𝐡𝐞𝐚𝐭
𝟐𝐅𝐞𝐒(𝐬) + 𝟑𝐎𝟐(𝐠) → 𝟐𝐅𝐞𝐎(𝐬) + 𝟐𝐒𝐎𝟐(𝐠)
The iron(II) oxide formed combines with sand to form a fusible slag.
𝐡𝐞𝐚𝐭
𝐅𝐞𝐎(𝐬) + 𝐒𝐢𝐎𝟐(𝐬) → 𝐅𝐞𝐒𝐢𝐎𝟑(𝐥) (slag)
Copper(I) oxide also reacts with some of the iron(II) sulphide to form iron(II) oxide.
𝐡𝐞𝐚𝐭
𝑪𝒖𝟐 𝑶(𝒔) + 𝑭𝒆𝑺(𝐬) → 𝐅𝐞𝐎(𝐬) + 𝑪𝒖𝟐 𝑺(𝒔)
 The reaction between copper(I) oxide and iron(II) sulphide results in the conversion of
most of the iron sulphide into the oxide, whereby iron(II) oxide is removed as slag.
MATTE
 This is a molten mass containing mainly copper(I) sulphide with little amount of iron(II)
sulphide.
 It is taken out through the outlet at the bottom.

21
 A diagram of a blast furnance for the extraction of copper
iv. Conversion
 The molten matte is then transferred to a Bessemer converter.
 This is a pear-shaped furnace made of steel.
 The furnace is fitted with pipes known as tuyéres on the sides through which sand and
hot air are blown.

A diagram of a bessemer converter

 The following reactions take place in the converter
(a) Iron(II) sulphide reacts with oxygen to form iron(II) oxide.
𝐡𝐞𝐚𝐭
𝟐𝐅𝐞𝐒(𝐥) + 𝟑𝐎𝟐(𝐠) → 𝟐𝐅𝐞𝐎(𝐬) + 𝟐𝐒𝐎𝟐(𝐠)
(b) The iron(II) oxide formed reacts with sand (silicon dioxide) to form slag.
𝐡𝐞𝐚𝐭
𝐅𝐞𝐎(𝐥) + 𝐒𝐢𝐎𝟐(𝐬) → 𝐅𝐞𝐒𝐢𝐎𝟑(𝐥) (slag)
(c) Some of the copper(I) sulphide is oxidised to copper(I) oxide
𝐡𝐞𝐚𝐭
𝟐𝑪𝒖𝟐 𝑺(𝒔) + 𝟑𝑶(𝐠) → 𝐒𝐎𝟐(𝐠) + 𝑪𝒖𝟐 𝑶(𝒍)

22
 (d) Copper(I) oxide is reduced by copper(I) sulphide to copper.
𝐡𝐞𝐚𝐭
𝟐𝑪𝒖𝟐 𝑶(𝒍) + 𝟐𝑪𝒖𝟐 𝑺(𝒍) → 𝐒𝐎𝟐(𝐠) + 𝟔𝑪𝒖(𝒍)
 The reaction taking place in (d) is known as self-reduction of copper because the
substances involved are both copper compounds.
 The molten metal is poured into sand moulds and allowed to stand. On cooling, any
dissolved sulphur dioxide escapes leaving blisters on the surface of copper.
 The copper obtained from the Bessemer converter is therefore known as blister copper.
v. Refining/purification of copper
 Blister copper is 97% to 99% pure. This means that it contains a small amount of impurities
which affect its quality.
 Trace amounts of sulphur in copper significantly reduce its electrical conductivity.
 Blister copper is refined further by electrolysis in which copper(II) sulphate solution is an
electrolyte.
 The blister copper is cast into blocks and made the anode.
 The cathode is made up of a sheet of pure copper.

A diagram of electrolyitic cell for refining blister copper

At the anode, blister copper is oxidised to copper(II) ions.

The ions move to the cathode where they are reduced to copper metal.
𝑪𝒖𝟐+ (aq) + 𝟐𝒆− → 𝑪𝒖(𝒔)
 The copper is deposited on the cathode sheet. As the electrolysis continues, the whole
blister copper block dissolves.
 This eventually increases the size of the pure copper plate.
USES OF COPPER
i. It is extensively used to manufacture electric cables and other electric appliances.
ii. It is used to make utensils, containers, calorimeters and coins.
iii. It is used in electroplating.
iv. It is alloyed with gold and silver for making jewels.

23
STUDY QUESTIONS
1. Describe any four chief ores of each of the following metals:
(a) Iron (b) Copper
2. Differentiate the extraction of copper from that of iron.
3. What is the role of limestone in the extraction of iron?

ENVIRONMENTAL DEGRADATION CAUSED BY EXTRACTION OF METALS

 The extraction of metals causes environmental degradation due to production of solid
waste, chemical waste, energy, gaseous pollutants, such as greenhouse gases, and dust.
i. Solid wastes
 A small amount of ore obtained from extraction process produces a huge mountain of solid
wastes.
 Mining wastes are not inert as they often contain toxic chemicals such as cyanide that are
added to the ore during extraction of metals.
ii. Chemical wastes
 Many of the chemicals that are used in the separation of metals are toxic.
 Froth flotation uses a range of chemicals, and leaching utilises acids as well as toxic metals
such as mercury.
 If such chemicals contaminate the environment (water, air, soil, and organisms), they may
cause health problems such as headache, nausea, genetic defects, and respiratory
disorders in human beings when they get exposed to them.
iii. Energy production
 Some processes such as smelting produce a large amount of mechanical, heat, and sound
energy.
 These interfere with the environment in some ways.
 The mechanical energy interferes with the earth’s structure, while the noise pollution
causes health problems such as hearing loss and cardiovascular diseases.
iv. Gaseous pollutants
 Smelting releases gases such as carbon dioxide, nitrogen dioxide and sulphur dioxide.
 These gases contribute to the formation of acid rain.
 The increase of carbon dioxide in the atmosphere contributes to global warming.
v. Dusts
 The crushing and separation processes of ores produce dust.
 The dust pollutes the environment and may cause health problems such as respiratory
problems, and irritation of eyes, nose, throat, and skin to human beings.

INTERVENTION/CONTROL MEASURES TO ENVIRONMENTAL DEGRADATION

 Extraction of metals is inevitable for economic reasons.
 However, it is important to consider environmental health for better life.
 The following are some intervention measures to environmental degradation due to metal
extraction:
i. Laws and regulations
 To minimise the environmental effects due to metal extraction, the laws and regulations
that require the extraction companies to restore the land to its previous state should be
implemented.
 This involves flattening the heaps of wastes, rehabilitation of landfill sites and landscapes,
eliminating ponds, and afforestation after extraction.

24
 ii. Recycling
 Minerals do not exist forever, thus once exploited they are gone.
 Recycling of metals such as iron reduces the need for further extraction.
 Recycling saves the valuable finite natural mineral resources and reduces the problems of
waste disposal.
iii. Controlling chemicals before release
 Toxic gases such as carbon monoxide which can be used as fuel must be controlled before
being released into the atmosphere.
 Harmful acidic gases such as sulphur dioxide and nitrogen dioxide can also be treated with
alkali solutions.
 For example, limewater can be used to neutralise the sulphur dioxide gas to harmless
calcium sulphate.
iv. Safety regulations and practices must be maintained to avoid the risk of
accidental release of harmful materials.
v. Care must be taken to relocate streams, wildlife and other valuable resources.
 Quarries and opencast workings can be reclaimed by the process of filling the holes with
solid wastes.
 The eroded bare soil can be conserved by planting trees and grasses to serve as a soil
cover, which would counteract the impacts of wind, running water, rain and animals to the
soil.
 Reclaimed land can have many uses such as agriculture, forestry, wildlife, habitation and
recreation.
STUDY QUESTIONS
1. Briefly explain the impacts of extraction of metals to the environment.
2. Describe the major ways which can be used to control the environmental impacts that are
caused by extraction of metals through mining.
3. Suggest the best solution for solid waste management in the mining industry.
4. Describe three methods that are used in extracting metals from metal ores. Give one
example of a metal that is extracted by using each of the methods.
5. Metal ores contain impurities that must be removed during the extraction process.
(a) Give three common impurities found in metal ores.
(b) Why should the impurities be removed before the reduction process?
6. The following figure is a set-up used in the extraction of sodium by the Down’s cell.

(a) What is the major ore from which sodium is extracted in this cell?
(b) What form of the ore is placed in the cell for the extraction?

25
 (c) Why is sodium not extracted by reduction using carbon?
(d) What role does calcium chloride play in the extraction of sodium from sodium chloride?
(e) Explain the roles of the parts marked A, B, C, and D.
7. Explain the following processes which are involved in the extraction of aluminium:
(a) Bauxite ore should be heated to bright red with sodium carbonate.
(b) An electrolyte must be added in the electrolysis of bauxite.
8. Iron is extracted from various ores by reduction in the blast furnace.
(a) Why is the furnace called a blast furnace?
(b) Identify the chief ore from which the iron metal is extracted.
(c) Describe the role of the blast furnace in the extraction process.
(d) What roles are played by tuyéres in the furnace?
(e) Write the equation for the reduction of the ore into iron.
(f) Explain the purpose of adding limestone and coke in the blast furnace.
9. The following figure is a sketch of the furnace for the extraction of copper:

(a) Explain the roles of parts E, F and G.

(b) Give four ores from which copper can be extracted.
(c) What is the chief ore from which copper is extracted?
(d) Write an equation to show the reaction between copper(I) oxide and iron(II) sulphide.
(e) Why is the reaction in (d) possible?
10. Some metals were mixed with salt solutions as outlined below:
(a) Copper and sodium chloride
(b) Iron and copper(II) sulphate
(c) Potassium and sodium carbonate
(d) Calcium and magnesium chloride
(e) Lead and silver nitrate
In which mixture(s) did a reaction take place? Write the ionic equation for each reaction.
11. The following equation represents a reaction between potassium and iodine:

26
 (a) Identify the oxidising agent, and the reducing agent.
(b) Write the balanced half reaction equations for the oxidation and reduction reactions.
12. Suggest the ways which can be used to minimise environmental pollution
due to liquid and solid wastes from mining.
13. Aluminium is mostly abundant on the Earth’s crust but not cheaply
available. Explain.
14. All ores are minerals, but not all minerals are ores. Explain.
15. Although iron pyrites is abundant in the Earth’s crust, it is not used as a source of iron.
Explain.
16. (a) Explain the function of coke and hot air in the extraction of iron from its ore.
(b) Account for the fact that aluminium is a vital element in our daily life. Give four points.
17. (a) Give three ways in which environmental destruction is likely to occur during extraction
of metals
(b) the following equations represent the steps involved in the conversion stages of iron
extraction in Bussener converter. Arrange the equations in the chronological order from
the first step to the last by writing the respective letter so as to get a complete explanation
of the conversion stage.

18. Describe the extraction of iron from the haematite ore and write all the chemical equations
for the reactions involved in each stage of extraction.
19. Describe four common stages for the extraction of metals. Does the extraction of gold
follow all four stages? Give reasons.
20. (a) List down four (4) common stages in the extraction of less reactive metals like zinc and
copper.
(b) Name the ore commonly used in the extraction of iron metal.
(c) The following are series of chemical reactions which occur in the blast furnace during
the process of extraction of iron metal.
C(s) + O2 → CO2 + heat (1)
CO2 + C → 2CO (2)
Fe2O3 + 3C → 2Fe + 3CO (3)
Fe2O3 + CO → Fe + 2CO2 (4)
(i) Indicate the two reducing agents in the blast furnace
(ii) Explain the importance of steps (1) to (3).
(iii) In this process a compound “L” which produces a chemical substance that
removes impurities as slag is added. Give the name of the substance.
(iv) Write the complete chemical reactions that compound “L” undergoes to form slag.
21. (a) Name the ore commonly used in the extraction of copper metal.
(b) Steps (i) to (iv) below are used during the extraction of copper metal from its ore.
Write a balanced chemical equation for each step.
(i) Roasting of the concentrated ore (CuFeS 2 ) in air.
(ii) Heating the roasted ore with silica in the absence of air.
(iii) Burning copper sulphide ore (CuS) in regulated supply of air.
(iv) Purification of copper by electrolysis using copper sulphate solution electrolyte, pure
copper cathode and impure copper obtained from the extraction anode.

27
 (c) With the help of chemical equations explain what will happen to
(i) an iron earring dropped into a container of copper sulphate solution?
(ii) copper knife dipped into zinc nitrate solution?
(iii) copper turnings dropped into a container of dilute hydrochloric acid?

SOLVED EXAMPLES
1. Although electrolysis is an expensive way of obtaining metals, it must be used for some
metals. Explain.
Answer
Group 1 and 2 metals together with Al are themselves such powerful reducing agents that
their oxides cannot be reduced by chemical reducing agents.

2. Below is a simplified diagram of the Downs cell in which sodium metal is manufactured.

i. Identify; -
Electrolyte Y: - Molten sodium chloride
Gas X: -Chlorine gas
ii. Write an equation for the reaction at the cathode.
Na+(l) + e- → Na(l).
iii. In what state is sodium collected?
Molten state/liquid state.
iv. Give two properties of Na that makes it possible to be collected as in (b) (iii) above.
 Its less dense than molten sodium chloride.
 Has a low melting point.
v. The cathode is made of steel but the anode is made of graphite. Why is this yet steel is a
better conductor?
 At high temperature steel would react with chlorine formed but graphite is inert even
at high temperatures.
vi. In this process, the naturally occurring, raw material is usually mixed with another
compound. Identify the compound and state its use.
ANS: Compound; - Calcium chloride
Use; -To lower melting point of NaCl2 from 800oC to 600oC
vii. What is the function of the steel gauze cylinder?
ANS: Prevents sodium reacting with chlorine forming NaCl
viii. Give one industrial use of sodium
 A coolant in nuclear reactors;
 Alloy with lead in tetraethyl (IV) lead;
ix. Explain why sodium metal is stored under paraffin;
ANS: If kept it out on air; reacts very fast with air forming a dull surface.
Also it can react with water.

28
x. A current of 100 Amperes flows through an electrolyte of molten sodium chloride for 15
hours. Calculate the mass of sodium produced in kg (Na = 23; 1F = 96500 C)
Solution:
Q = It
=100 x 15 x 60 x 60
=5400000C.
Cathode equation
Na+(l) + e- → Na(l)
96500C = 23g of Na
5400000 C = 23 x 5400000/96500 =1287.04g
The mass of sodium =1.287kg
3. Give two requirements for vapour phase refining.
Solution:
(i) The metal should react easily with the reagent to form the complex.
(ii) The volatile complex should easily decomposable so that the recovery is easy.
4. Write down the reactions taking place in Blast furnace-related to the metallurgy of iron in
the temperature range 500-800 K.
Solution:
These are the reactions that happen in the blast furnace at the 500K-800K range.
a. 3Fe2O3 + CO → 2Fe3O4 + CO2
b. Fe3O4 + CO →3FeO + 2CO2
c. Fe2O3 + CO → 2FeO + CO2
5. What is the role of flux in metallurgical processes?
Solution:
 To remove gangue, certain substances are mixed with it. These are called fluxes. Flux
can be basic or acidic. Acidic flux removes basic impurity and basic flux removes acidic
impurity.
 Flux is also used for increased conductivity.
6. What should be the considerations during the extraction of metals by electrochemical
method?
Solution:
(i) Reactivity of the metal produced
(ii) Electrodes to be made of suitable material
(iii) Addition of flux for making molten mass conducting
7. Why are sulphide ores converted to oxide before reduction?
Solution:
 Sulphide ores cannot be easily reduced so they are usually converted to oxides
because oxides can be reduced easily.
8. Why is sulphide ore of copper heated in a furnace after mixing with silica?
Solution:
 Sulphide ore of copper is heated in a furnace after mixing with silica because the iron
impurities present in the ore can form slag with silica and will be easily removed. The
copper is produced as copper matte.
FeO + SiO2→ FeSiO3(slag)
9. Although carbon and hydrogen are better-reducing agents they are not used to reduce
metallic oxides at high temperatures. Why?
Solution

29
  Carbon and hydrogen will not reduce oxides to metals but they will instead form
hydrides and carbides. Therefore, they are not used as reducing agents.
10. Wrought iron is the purest form of iron. Write a reaction used for the preparation of wrought
iron from cast iron. How can the impurities of sulphur, silicon and phosphorus be removed
from cast iron?
Solution:
Fe2O3 + 3C → 2Fe + 3CO
Limestone is added as a flux to remove the impurities of sulphur, silicon and phosphorous.
They form a slag which can be easily removed. The metal is removed from the slag by
passing through rollers.
11. (a) carbon monoxide was passed over heated iron (II) oxide
i. Write equation for the reaction that took Place.
ii. Write equation for the reaction between the solid product in (a)(i) and dilute
sulphuric acid.
(b) Chlorine was bubbled through the product in (a)(i)
(i) State what was observed
(ii) Write equation for the reaction that took place
Solution
(a) (i) FeO(s) + CO (g) → Fe(s) + CO2(g)
(ii) Fe(s) + H2SO4(aq) → FeSO4(aq) + H2(g)
(b) (i) black crystals formed
(ii) 2Fe(s) + 3Cl2(g) → 2FeCl3(g)
12. (a) Name the raw materials which are used in extraction of iron using a blast furnace.
(b) Briefly describe the reactions that lead to the formation of iron during extraction using
a blast furnace.
(c) state what would be observed and write equations for the reactions that would take
place when the following gases are passed over heated iron
(i) dry chloride
(ii) steam
(d) Dilute hydrochloric acid was added to iron filling and a mixture warmed. Write the
equation for the reaction that took place.
Solution
(a) haematite, Fe2O3 and magnetite, Fe3O4, and from the siderite, FeCO3
(b) (i) Reduction of the ore
 The roasted is mixed with coke and lime stone in a blast furnace and heated
with hot air.
 Coke or carbon burns in oxygen to form carbon dioxide
C(s) + O2(g) → CO2(g)
 Carbon dioxide reacts with carbon to form carbon monoxide
CO2(s) + C(s) → 2CO(g)
 Carbon monoxide reduces iron (III) oxide to iron
Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g)
(ii) Removal of impurities of silicon dioxide, SiO2.
 Lime stone or calcium carbonate decompose on heating
CaCO3(s) → CaO(s) + CO2(g)
 Calcium oxide reacts with impurities to form slag of calcium silicate that is pour away
CaO(s) + SiO2(s) → CaSiO3(l)
(c) (i) iron glow red forming black crystals and purple vapor

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 2Fe(s) + 3Cl2(g) → 2FeCl3(s)
(ii) a black solid forms
3Fe(s) +4H2O(l) → Fe3O4(l) + 4H2
(d) Fe(s) + 2HCl (aq) → FeCl2(g) + H2(g)

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