Occurrence and Location of Metals in Tanzania
Locations of Important Metal Ores in Tanzania
Identify locations of important metal ores in Tanzania
Most metals are found naturally as compounds called minerals.
Rocks are made up of crystals of metals. An ore is a rock that contains enough of a metal compounds for it to be worth extracting the metal.
The most common ores contain oxides. An example is the ore haematite, which contains iron (III) oxide. Some contain other metal compounds.
Malachite contains copper (II) carbonate.
Chemical Properties of Metals
The Difference between Physical and Chemical Strengths of Metals
Differentiate the physical and chemical strengths of metals
Metals are known to have a greater tendency to lose electrons in the course of which they go into solution as ions. A metal, therefore, may be defined as an element which can ionize by electron loss. The number of electrons lost per atom is the valency of the metal and the ion carries an equal number of positive charges, as:
- Na →Na+ + e– (univalent)
- Mg →Mg2+ + 2e– (divalent)
- Al →Al3+ + 3e– (trivalent)
The following are important chemical properties of metals:
- They
react with oxygen to form oxides. For example, magnesium burns in air
to form magnesium oxide. Metal oxides are bases, which mean they react
with water to form an alkaline solution and with acids to form salts. - Metals
form positive ions when they ionize. Consider the ionization of sodium,
magnesium and aluminium in the above equations in which case ionization
resulted into Na+, Mg2+ , and Al3+ ions respectively. However, there are some exceptions. For example, hydrogen is a non-metal which forms positive ions in solution, H+. This is the only exception in this case.
Physical
strength is different from chemical strength. Physical strength is
tensile strength of the metal. There are metals with high tensile
strengths like iron, copper and aluminium. Other metals like sodium and
potassium have low tensile strengths..Chemical strength is the
reactivity of the metal. Sodium and potassium have very low tensile
strengths but they are the strongest metals chemically.
strength is different from chemical strength. Physical strength is
tensile strength of the metal. There are metals with high tensile
strengths like iron, copper and aluminium. Other metals like sodium and
potassium have low tensile strengths..Chemical strength is the
reactivity of the metal. Sodium and potassium have very low tensile
strengths but they are the strongest metals chemically.
Sodium and potassium are very weak physically. Sodium is very soft metal and can be cut easily with a knife. It melts at 98oC.
It is so light that if floats on water, but it reacts immediately with
the water forming an alkaline solution. When freshly cut, it has a
silvery lustre but rapidly furnishes due to its reaction with
atmospheric carbon dioxide and moisture.
It is so light that if floats on water, but it reacts immediately with
the water forming an alkaline solution. When freshly cut, it has a
silvery lustre but rapidly furnishes due to its reaction with
atmospheric carbon dioxide and moisture.
Potassium is a white metal with lustre. It melts at 63oC.
Like sodium, potassium is a very light metal and it can also float on
water, with which it reacts to form an alkaline solution.
Like sodium, potassium is a very light metal and it can also float on
water, with which it reacts to form an alkaline solution.
However,
sodium and potassium are among the strongest metals chemically. These
metals are very reactive and they are always found combined with other
elements. The metals are so reactive that they will combine with any
non-metal nearby. They never occur free in nature. Both sodium and
potassium are so reactive that they have to be stored under oil to
prevent them coming into contact with water or air.The compounds of
sodium and potassium are quite abundant in nature.
sodium and potassium are among the strongest metals chemically. These
metals are very reactive and they are always found combined with other
elements. The metals are so reactive that they will combine with any
non-metal nearby. They never occur free in nature. Both sodium and
potassium are so reactive that they have to be stored under oil to
prevent them coming into contact with water or air.The compounds of
sodium and potassium are quite abundant in nature.
The reactivity and tensile strengths of some metals
- Calcium:Calcium
is amongst a group of metals that are too reactive to occur in the free
state. It occurs mainly as carbonate, sulphate, fluoride and silicate.
It is a soft, greyish metal. In comparison with potassium and sodium, it
has a lower tensile strength and high density. - Iron:Iron is a typical metal. Its density is 7.87. It melts at 1530oC. Iron is a moderately reactive metal. The metal reacts with excess steam at red heat to produce triiron tetraoxide.
- Copper Copper
is a less reactive metal. It is a red-brown metal with a lustre. It can
be polished. Its tensile strength is fairly high. When heated in air,
copper forms a layer of black copper (II) oxide on the surface: 2Cu(s) + O2(g)→2CuO(s)It reacts with hot concentrated sulphuric acid to form copper (II) sulphate and liberate sulphur dioxide. Cu(s) + 2H2SO4(aq)→CuSO4(aq) + 2H2O(l) + SO2(g)
The table below compares the tensile strengths of calcium, iron and copper metals.
Tensile strengths, densities and melting points of calcium, iron and copper
| Calcium | Copper | Iron | |
| Tensile strength | Low | Fairly high | High |
| Melting point (oC) | 850 | 1080 | 1535 |
| Density (g cm-3) | 1.55 | 8.95 | 7.9 |
The Reducing Power of Different Metals
Compare the reducing Power of different metals
As we learned early, metals tend to lose electrons during chemical reactions. This process of losing electrons is called oxidation.
Metals normally lose electrons to non-metals, which accept those
electrons. Therefore, metals are said to be electron donors while
non-metals are electron acceptors. In this case, metals can be termed as
reducing agents, because they donate electrons which, when
accepted by non-metals, tend to lower their oxidation numbers.
Non-metals are called oxidizing agents, because they oxidize or increase the oxidation number of metallic atoms through accepting the electrons donated by metals.
Metals normally lose electrons to non-metals, which accept those
electrons. Therefore, metals are said to be electron donors while
non-metals are electron acceptors. In this case, metals can be termed as
reducing agents, because they donate electrons which, when
accepted by non-metals, tend to lower their oxidation numbers.
Non-metals are called oxidizing agents, because they oxidize or increase the oxidation number of metallic atoms through accepting the electrons donated by metals.
Metals
differ in the ease of losing the electrons, depending on their
electronic configurations. This is because; it is only the outer
electron(s), which take part in a chemical reaction. The nucleus of an
atom, being positively charged, normally attracts the electrons towards
itself, hence making the electrons difficult to remove from their
shells. The further the electrons are from the central nucleus, the
easier it is removing them from their shells and vice versa. Therefore,
atoms with larger atomic radii donate their electrons more easily than
those with small atomic radii.
differ in the ease of losing the electrons, depending on their
electronic configurations. This is because; it is only the outer
electron(s), which take part in a chemical reaction. The nucleus of an
atom, being positively charged, normally attracts the electrons towards
itself, hence making the electrons difficult to remove from their
shells. The further the electrons are from the central nucleus, the
easier it is removing them from their shells and vice versa. Therefore,
atoms with larger atomic radii donate their electrons more easily than
those with small atomic radii.
For
example, compare the electronic configuration of sodium (2:8:1) with
that of potassium (2:8:8:1). Sodium ionizes by losing one electron from
its outer shell to attain the structure of the nearest noble gas (2:8).
Likewise, potassium ionizes by losing one electron from its outer shell
to attain the noble gas structure, 2:8:8. But, in which of the two cases
is it easy to remove electron and why? It is easy to remove the
electron from the outer shell of potassium than sodium because the
atomic radius of calcium is too large as compared to that of sodium.
This is because in calcium the outer electron is more loosely held by
the central nucleus and hence easy to remove from its shell. This is
true because the ability of the nucleus to hold the electrons firmly
depends on the distance of the electrons from the nucleus.
example, compare the electronic configuration of sodium (2:8:1) with
that of potassium (2:8:8:1). Sodium ionizes by losing one electron from
its outer shell to attain the structure of the nearest noble gas (2:8).
Likewise, potassium ionizes by losing one electron from its outer shell
to attain the noble gas structure, 2:8:8. But, in which of the two cases
is it easy to remove electron and why? It is easy to remove the
electron from the outer shell of potassium than sodium because the
atomic radius of calcium is too large as compared to that of sodium.
This is because in calcium the outer electron is more loosely held by
the central nucleus and hence easy to remove from its shell. This is
true because the ability of the nucleus to hold the electrons firmly
depends on the distance of the electrons from the nucleus.
Metals
that release the electrons more readily are said to be strong reducing
agents compared to those that release their electrons least readily. For
example, potassium with an electronic configuration of 2:8:8:1 is a
stronger reducing agent than sodium, which has an electronic
configuration of 2:8:1.
that release the electrons more readily are said to be strong reducing
agents compared to those that release their electrons least readily. For
example, potassium with an electronic configuration of 2:8:8:1 is a
stronger reducing agent than sodium, which has an electronic
configuration of 2:8:1.
Therefore, it is clear that when atoms of elements lose electrons from their outermost shells, they obtain octet or doublet structures
and hence become chemically stable. It is not easy to remove
electron(s), by ordinary chemical means from such stable noble gas
structures.
and hence become chemically stable. It is not easy to remove
electron(s), by ordinary chemical means from such stable noble gas
structures.
From
the context of the above explanations, the reducing power of a metal
can be defined as its readiness to release electrons from its outer
shell. Metals whose atoms release electrons more readily have high
reducing powers than those metals whose atoms do not readily release
their electrons.
the context of the above explanations, the reducing power of a metal
can be defined as its readiness to release electrons from its outer
shell. Metals whose atoms release electrons more readily have high
reducing powers than those metals whose atoms do not readily release
their electrons.
Reactivity series of metals
Reactivity
series refers to arranging or listing the metals in order of
reactivity. The reactivity series are obtained by consideration of the
action of air, water and acids on the metals,
and how easily the oxides of these metals can be reduced. Consider the
table of reactivity series below (Table 8.2). Oxides of the first group
of five metals cannot be reduced by carbon. Those of the second group of
three metals can react with acids, displacing hydrogen. The third and
last group comprises of least reactive metals. In table 8.2, the metals
are arranged in order of reactivity series. It indicates the inverse
order in which the elements were isolated. Thus, metals low in the
series such as gold, silver and lead have been known since early times.
Metals high in the series proved very difficult to isolate. It was
Davy’s work on electrolysis that led to isolation of potassium, sodium,
calcium, magnesium and aluminium over a period of years from 1807, when
Davy isolated potassium and sodium, to about 1850, when aluminium was
isolated.
series refers to arranging or listing the metals in order of
reactivity. The reactivity series are obtained by consideration of the
action of air, water and acids on the metals,
and how easily the oxides of these metals can be reduced. Consider the
table of reactivity series below (Table 8.2). Oxides of the first group
of five metals cannot be reduced by carbon. Those of the second group of
three metals can react with acids, displacing hydrogen. The third and
last group comprises of least reactive metals. In table 8.2, the metals
are arranged in order of reactivity series. It indicates the inverse
order in which the elements were isolated. Thus, metals low in the
series such as gold, silver and lead have been known since early times.
Metals high in the series proved very difficult to isolate. It was
Davy’s work on electrolysis that led to isolation of potassium, sodium,
calcium, magnesium and aluminium over a period of years from 1807, when
Davy isolated potassium and sodium, to about 1850, when aluminium was
isolated.
The reactivity series of metals
Metals
low down in the series are frequently found as the free elements,
although they may also be obtained from ores because the amounts found
as the free metal are not sufficient for industrial purposes. Gold,
however, the last element of the series is found and mined almost
entirely as the free element.
low down in the series are frequently found as the free elements,
although they may also be obtained from ores because the amounts found
as the free metal are not sufficient for industrial purposes. Gold,
however, the last element of the series is found and mined almost
entirely as the free element.
Generally,
it is these relatively uncreative metals that we find the most uses
for. Iron and copper, for example, can be found in many household and
everyday objects.Metals higher up in the series are more reactive than
those lower down. A metal higher up in the series will displace a metal
lower down from a solution of one of its salts. For example, iron will
displace copper from its salt.
it is these relatively uncreative metals that we find the most uses
for. Iron and copper, for example, can be found in many household and
everyday objects.Metals higher up in the series are more reactive than
those lower down. A metal higher up in the series will displace a metal
lower down from a solution of one of its salts. For example, iron will
displace copper from its salt.
Fe(s) + CuSO4(aq)→ Cu(s) + Fe SO4(aq)
Zinc and iron will displace hydrogen from dilute acids such as dilute hydrochloric acid or dilute sulphuric acid.
Zn(s) + 2HCl(aq) →ZnCl2(aq) + H2(g)
Fe(s) + H2SO4(aq) →FeSO4(aq) + H2(g)
Copper
has no reaction on either dilute sulphuric acid or dilute hydrochloric
acid. With hot concentrated sulphuric acid, sulphur dioxide is liberated
and copper (II) sulphate is formed.
has no reaction on either dilute sulphuric acid or dilute hydrochloric
acid. With hot concentrated sulphuric acid, sulphur dioxide is liberated
and copper (II) sulphate is formed.
Cu(s) + 2H2SO4(l) →CuSO4(aq) + 2H2O(l) + SO2(g)
Useful things to remember about the reactivity series
- The
more reactive the metal, the more compounds it forms. So only copper,
silver and gold are ever found as free elements in the earth’s crust.
The other metals are always found as compounds. - When a metal reacts, it gives up electrons to form ions. The more reactive the metal, the more easily it gives up electrons.
- The more reactive the metal, the more stable its compounds are. Stable means difficult to breakdown. For example, when you heat sodium nitrate you get sodium nitrite:2NaNO3(s)→ 2NaNO2(s) + O2(g)But copper (II) nitrate breaks down further, to oxide, giving off nitrogen dioxide:2Cu(NO3)2(s)→ 2CuO(s) + 4NO2(g) + O2(g)
- The more reactive the metal, the more difficult it is to extract
from its compounds (since the compounds are stable). For the most
reactive metals, you will need the toughest method of extraction: electrolysis. - The
less reactive metals have been known and used since ancient times,
because they are easiest to extract.6. If you stand two metals in an
electrolyte and join them up with a copper wire, electrons will flow
from the more reactive metal to the less reactive one.
Extraction of Metals by Electrolytic Reduction
The Criteria for the Choice of the Best Methods of Extracting a Metal from its Ore
Outline the criteria for the choice of the best methods of extracting a metal from its ore
Most metals are naturally found in compounds called ores. The metals are obtained from the ores by a process calledextraction.The extraction consists of separating the metal from minerals in the ore
A few metals are found “native”. This means that they are uncombined with other elements, for example, silver and gold.To
extract ores containing minerals, a chemical reaction must be used to
separate the metal from other elements. The choice of the best method
for extraction of a metal from its ore depends on the chemical
reactivity of the metal to be extracted. Most ores contain metal oxides.
To extract the metal, oxygen must be removed from it. This reaction if
called reduction.
extract ores containing minerals, a chemical reaction must be used to
separate the metal from other elements. The choice of the best method
for extraction of a metal from its ore depends on the chemical
reactivity of the metal to be extracted. Most ores contain metal oxides.
To extract the metal, oxygen must be removed from it. This reaction if
called reduction.
Because
the metal is always the most electropositive part of an ore, and so has
a positive oxidation state, the formation of the free metal from its
ore is always a reduction.
the metal is always the most electropositive part of an ore, and so has
a positive oxidation state, the formation of the free metal from its
ore is always a reduction.
Mn+ + ne– →Mo
Very
reactive metals like sodium and potassium are strongly bonded in their
ores. The more reactive the metal is the more stable its compounds are,
and the more energy is needed to break down the bond between the metal
and oxygen. Therefore, the extraction of these metals requires a strong
method of reduction: electrolysis (or electrolytic reduction).
reactive metals like sodium and potassium are strongly bonded in their
ores. The more reactive the metal is the more stable its compounds are,
and the more energy is needed to break down the bond between the metal
and oxygen. Therefore, the extraction of these metals requires a strong
method of reduction: electrolysis (or electrolytic reduction).
Hence,
metals at the top of the reactivity series are obtained from their ores
by electrolysis (electrolytic reduction). Electrolysis is a more
expensive process than reduction with carbon or carbon monoxide.
However, it is the only economic way to obtain metals such as aluminium.
For less reactive metals at the middle of the reactivity series, the
oxygen can be removed by chemical reduction with carbon or carbon
monoxide. This method is used for extraction of metals such as zinc,
iron and copper as shown and discussed in table 8.3. Least reactive
metals such as copper, silver and gold may be found in uncombined state.
metals at the top of the reactivity series are obtained from their ores
by electrolysis (electrolytic reduction). Electrolysis is a more
expensive process than reduction with carbon or carbon monoxide.
However, it is the only economic way to obtain metals such as aluminium.
For less reactive metals at the middle of the reactivity series, the
oxygen can be removed by chemical reduction with carbon or carbon
monoxide. This method is used for extraction of metals such as zinc,
iron and copper as shown and discussed in table 8.3. Least reactive
metals such as copper, silver and gold may be found in uncombined state.
Methods of extraction of different metals from their ores
| Metal | Method of extraction from ore |
| Potassium | Electrolysis |
| Sodium | |
| Calcium | |
| Magnesium | |
| Aluminium | |
| Zinc | Chemical reduction with carbon or carbon monoxide |
| Iron | |
| Lead | |
| Copper | Roasting in air |
| Silver | Occur naturally as elements |
| Gold |
Stages of the extraction of moderately reactive metals
There are four main stages for the extraction of moderately reactive metals:
- Mining and concentration of the ore
- Roasting in air
- Reduction of oxides to metals
- Purifying the metal
Mining and concentration of the ore
Mining is just the first step. The ore must then be decomposed to give the metal. This is called extraction.The unwanted rock material in the ore is called gangue. The ore is treated to remove as much gangue as possible. For example:
- It
is crushed and washed. In this case, the ore is broken down into small
pieces, which are then grinded down to fine powder. Then it is either
dropped into water, where the fragments containing the metal sink faster
or jets of air are blown at it, where the lighter waste material is
carried to one side. - A method called froth flotation
is used with sulphide ores (e.g. CuS or ZnS). The ore is powdered, fed
into water tanks and made into slurry with water. Then “frothing”
chemicals (a suitable oil) are added. Sulphides are attracted to these
chemicals. When air is blown through the slurry, froth rises to the top
of the tank carrying the metal sulphides with it. They are skimmed off
and dried. The gangue sinks. - Magnetic separation can be used. The iron ore can be separated from other material in the crushed ore by using electromagnet.
Roasting in air
For
the ores that occur as sulphides or carbonates of the metal, the
concentrated ore is heated (roasted) in air to convert the ore into an
oxide, for example:
the ores that occur as sulphides or carbonates of the metal, the
concentrated ore is heated (roasted) in air to convert the ore into an
oxide, for example:
- 2PbS(s) + 3O2(g)→ 2PbO(s) + 2SO2(g)
- 2ZnS(s) + 3O2(g)→ 2ZnO(s) + 2SO2(g)
- ZnCO3(s)→ ZnO(s) + CO2(g)
It
is usual to convert sulphides and carbonates into oxides before
reduction because oxides are more easily and efficiently reduced than
sulphides. The oxides resulting from roasting (heating) the sulphides or
carbonates in air are then reduced chemically with carbon or carbon
monoxide.
is usual to convert sulphides and carbonates into oxides before
reduction because oxides are more easily and efficiently reduced than
sulphides. The oxides resulting from roasting (heating) the sulphides or
carbonates in air are then reduced chemically with carbon or carbon
monoxide.
Most copper is extracted from copper pyrites, CuFeS2. The ore is roasted in a limited supply of air to produce to produce copper (I) sulphide:
2CuFeS2(s) + 4O2(g) → Cu2S(s) + 3SO2(g) + 2FeO(s)
The copper (I) sulphide is then reduced to copper by heating in air: Cu2S(s) + O2(g) → 2Cu(s) + SO2(g)
The
copper produced by this method is never pure. It must be refined
(purified) by electrolysis if it has to be used for electrical wiring.
copper produced by this method is never pure. It must be refined
(purified) by electrolysis if it has to be used for electrical wiring.
Reduction of oxides to metals
This
is another important stage in the extraction of metal. Here, the
roasted ore must now be reduced to respective metals. Reduction occurs
at a very high temperature. The materials employed for reduction are
mainly carbon or carbon monoxide. Thus;
is another important stage in the extraction of metal. Here, the
roasted ore must now be reduced to respective metals. Reduction occurs
at a very high temperature. The materials employed for reduction are
mainly carbon or carbon monoxide. Thus;
Fe2O3(s) + 3CO (g) →2Fe (l) + 3CO2(g)
ZnO(s) + C(s) →Zn(s) + CO(g)
Purification (Refining)
Very
often, the product of the reduction process is never pure. The product
has to be purified first before being put into use. Purification is
normally done through a number of ways, which include:
often, the product of the reduction process is never pure. The product
has to be purified first before being put into use. Purification is
normally done through a number of ways, which include:
- Electrolysis:
Electrolysis is used to produce a pure metal directly from its molten
compounds. Examples of metals which are purified by electrolysis are
copper and zinc. Copper produced in large scale is purified by
electrolysis, a process often called copper refining. - Oxidation:
The molten crude metal is exposed to hot air in a furnace. The
impurities in the crude metal are oxidized with oxygen from the air.
They escape as vapour or form a scum over the molten metal, which is
then removed by skimming. However, this method is used only when the
impurities have a greater affinity for oxygen than the metal. The method
is applied in the manufacture of steel from pig iron and in the
purification of tin and lead. - Distillation:In
distillation, the crude metal is heated in a furnace until the pure
metal evaporates, leaving behind the impurities. The vapour is then
collected and condensed in a separate chamber. This method forms an
integral part in the extraction of zinc, cadmium and mercury. A further
distillation, usually in vacuum, gives a very pure product. - Formation of carbonyls: very pure nickel and iron are made by forming their volatile carbonyls, which are then decomposed by heating.
- Zone refining:
This recently developed method is used to produce silicon and germanium
of extreme purity. In this method, a small high-frequency induction
furnace is placed round one end of a long rod of the metal and a thin
cross-section of the metal is melted. The furnace is then moved slowly
along the rod. Pure crystals of the metal separate from the melting
metal but impurities remain in the liquid and are carried along to the
other end.
The Extraction of Sudium from its Ore
Explain the extraction of sodium from its ore
Alkali
metals are strong reducing agents and cannot be extracted by chemical
reduction of their oxides or other compounds. The only possible method
of their extraction is by electrolysis of their fused chlorides.Sodium
is extracted industrially by electrolysis of either fused sodium
hydroxide (Castner’s process) or fused sodium chloride (Down’s process),
in which sodium chloride is electrolysed in the molten condition.
metals are strong reducing agents and cannot be extracted by chemical
reduction of their oxides or other compounds. The only possible method
of their extraction is by electrolysis of their fused chlorides.Sodium
is extracted industrially by electrolysis of either fused sodium
hydroxide (Castner’s process) or fused sodium chloride (Down’s process),
in which sodium chloride is electrolysed in the molten condition.
Electrolysis of sodium chloride by the down’s process
In
this case, fused sodium chloride is used. And because the melting point
of sodium chloride is high (about 800°C), calcium chloride is added to
lower the melting point to about 600°C and thus economize on electrical
power. The composition of the electrolyte is 40% sodium chloride and 60%
calcium chloride.The Down’s cell (figure 8.1) used for the extraction
of sodium consists of an iron box through the bottom of which rises a
circular carbon anode. A ring-shaped iron cathode surrounds this carbon
anode. The cathode is enclosed in iron gauze diaphragm, which also
separates the two electrodes. At 600°C, the sodium and chlorine produced
would react violently if allowed to come in contact. A diaphragm around
the anode, which keeps the two products apart, prevents this.
this case, fused sodium chloride is used. And because the melting point
of sodium chloride is high (about 800°C), calcium chloride is added to
lower the melting point to about 600°C and thus economize on electrical
power. The composition of the electrolyte is 40% sodium chloride and 60%
calcium chloride.The Down’s cell (figure 8.1) used for the extraction
of sodium consists of an iron box through the bottom of which rises a
circular carbon anode. A ring-shaped iron cathode surrounds this carbon
anode. The cathode is enclosed in iron gauze diaphragm, which also
separates the two electrodes. At 600°C, the sodium and chlorine produced
would react violently if allowed to come in contact. A diaphragm around
the anode, which keeps the two products apart, prevents this.
During
electrolysis, chlorine is librated at the anode and escapes via the
hood. Sodium is liberated at the cathode, collects in the inverted
trough placed over the cathode, rises up the pipe, and overflows into
the storage tank, from which it is tapped off through the iron vessel.
electrolysis, chlorine is librated at the anode and escapes via the
hood. Sodium is liberated at the cathode, collects in the inverted
trough placed over the cathode, rises up the pipe, and overflows into
the storage tank, from which it is tapped off through the iron vessel.
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
Down’s cell contains some calcium, which is also formed through
electrolysis. The calcium crystallizes when the mixture cools and a
relatively pure sodium metal is obtained.
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
Down’s cell contains some calcium, which is also formed through
electrolysis. The calcium crystallizes when the mixture cools and a
relatively pure sodium metal is obtained.
The Down’s cell
Fused sodium chloride produces Na+ and Cl– ions, where Na+ moves to the cathode and Cl– moves to the anode. The reactions taking place at the electrodes are as follows:
Chlorine gas is therefore formed at the anode while sodium metal is produced at the cathode.
2NaCl(l) →2Na(s) + Cl2(g)
Chlorine is a valuable by-product of the decomposition process
Extraction of Metals by Chemical Reduction
The Extraction of Iron from its Ore
Describe the extraction of iron from its ore
Iron
is only second to aluminium as the most abundant metal in the earth’s
crust. Its chief ores are haematite, an impure iron (III) oxide, Fe2O3, which contains about 70% of iron; magnetite (or magnetic iron ore), triiron tetraoxide, Fe3O4, which contains 72.4% of iron; and spathic iron ore, iron (II) carbonate, FeCO3. It also occurs as limonite, Fe2O3.xH2O and as the sulphide in iron pyrites, FeS2.
However, though abundant in the earth’s crust, iron pyrite is not used a
source of iron. It is mainly used in the production of sulphuric acid.
is only second to aluminium as the most abundant metal in the earth’s
crust. Its chief ores are haematite, an impure iron (III) oxide, Fe2O3, which contains about 70% of iron; magnetite (or magnetic iron ore), triiron tetraoxide, Fe3O4, which contains 72.4% of iron; and spathic iron ore, iron (II) carbonate, FeCO3. It also occurs as limonite, Fe2O3.xH2O and as the sulphide in iron pyrites, FeS2.
However, though abundant in the earth’s crust, iron pyrite is not used a
source of iron. It is mainly used in the production of sulphuric acid.
To extract iron, three substances are needed:
- Iron ore: The chief ore is haematite. It is mainly iron (III) oxide, Fe2O3 mixed with sand.
- Limestone: This is mainly calcium carbonate, CaCO3.
- Coke: This is made from coal and is almost pure carbon.
The extraction of iron in a blast furnace
Before being introduced into the blast furnace, all forms of iron ore must be converted to the oxide, Fe2O3, by roasting in air.
4FeCO3(s) + O2(g) →2Fe2O3(s) + 4CO2(g)
4FeS2(s) +11O2(g) →2Fe2O3(s) + 8SO2(g)
Iron ore, coke and limestone are mixed together to give a mixture called charge. The charge is introduced into a tall tower called a blast furnace
(figure 8.2). At the bottom of the furnace, hot air is blasted in
through several pipes known as tuyeres. A well at the bottom of the
furnace serves to hold the molten iron and slag until these can be run
off. The charge is fed in continuously from the top.
(figure 8.2). At the bottom of the furnace, hot air is blasted in
through several pipes known as tuyeres. A well at the bottom of the
furnace serves to hold the molten iron and slag until these can be run
off. The charge is fed in continuously from the top.
Reactions that occur in the furnace
- At the bottom of the furnace where temperature is the highest, air attacks the coke to produce carbon dioxide.C(s) + O2(g)→ CO2(g)
- In the middle of the furnace, the rising up carbon dioxide gas is reduced by more coke, producing carbon monoxide.C(s) + CO2(g)→ 2CO(g)
- At the top of the furnace, carbon monoxide reduces iron (III) oxide to metal. Fe2O3(s) + 3CO(g)→2Fe(s) + 3CO2(g)
The
molten iron trickles down the furnace and gathers at the bottom.
Periodically, this molten iron is tapped off and run into moulds (or
containers), where it is allowed to cool in long bars of about 1 metre
long and 10 cm in diameter. At this stage, it is called ‘cast iron’ or
‘pig iron’.
molten iron trickles down the furnace and gathers at the bottom.
Periodically, this molten iron is tapped off and run into moulds (or
containers), where it is allowed to cool in long bars of about 1 metre
long and 10 cm in diameter. At this stage, it is called ‘cast iron’ or
‘pig iron’.
The hot waste gases, mainly nitrogen and oxides of carbon, escape from the top of the furnace and are used to heat incoming air. This helps to reduce the energy costs of the process.
The Blast furnace
Action of limestone
The
limestone, which is introduced together with the ore, is first
decomposed at this high temperature to form calcium oxide.CaCO3(s)→ CaO(s) + CO2(g)
limestone, which is introduced together with the ore, is first
decomposed at this high temperature to form calcium oxide.CaCO3(s)→ CaO(s) + CO2(g)
The main impurity in iron ore is sand, made of silicon dioxide (SiO2). This reacts with calcium oxide to form calcium silicate.CaO(s) + SiO2(s) →CaSiO3(l)
The calcium silicate, which is in liquid state, falls to the bottom of the furnace. This liquid, called slag
(because it is less dense than the molten iron) is tapped off
separately. Slag is a useful by-product. It is used for making roads,
production of cement, and as a fertilizer.
(because it is less dense than the molten iron) is tapped off
separately. Slag is a useful by-product. It is used for making roads,
production of cement, and as a fertilizer.
