A few elements like carbon, sulphur, gold and noble gases,
occur in free state while others in combined forms in the earth’s crust. The
extraction and isolation of an element from its combined form involves various
principles of chemistry. A particular element may occur in a variety of
compounds. The process of metallurgy and isolation should be such that it is
chemically feasible and commercially viable. Still, some general principles are
common to all the extraction processes of metals. For obtaining a particular
metal, first we look for minerals which are
naturally occurring chemical substances in the earth’s crust obtainable by
mining. Out of many minerals in which a metal may be found, only a few are
viable to be used as sources of that metal. Such minerals are known as ores.
Rarely, an ore
contains only a desired substance. It is usually contaminated with earthly or
undesired materials known as gangue. The
extraction and isolation of metals from ores involve the following major steps:
•
Concentration of the ore,
•
Isolation of the metal from its concentrated ore, and
•
Purification of the metal.
The entire
scientific and technological process used for isolation of the metal from its
ores is known as metallurgy.
In the
present Unit, first we shall describe various steps for effective concentration
of ores. After that we shall discuss the principles of some of the common
metallurgical processes. Those principles shall include the thermodynamic and
electrochemical aspects involved in the effective reduction of the concentrated
ore to the metal.
6.1
Occurrence of Metals
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Elements vary in
abundance. Among metals, aluminium is the most abundant. It is the third most
abundant element in earth’s crust (8.3% approx. by weight). It is a major
component of many igneous minerals including mica and clays. Many gemstones
are impure forms of Al2O3 and the impurities range from Cr (in ‘ruby’) to Co (in
‘sapphire’). Iron is the second most abundant metal in the earth’s crust. It
forms a variety of compounds and their various uses make it a very important
element. It is one of the essential elements in biological systems as well.
The principal ores of aluminium, iron, copper and zinc have
been given in Table 6.1.
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6.2
Concentration
of Ores
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iron, usually the oxide
ores which are abundant and do not produce polluting gases (like SO2
that is produced in case iron pyrites) are taken. For copper and zinc, any of
the listed ores (Table 6.1) may be used depending upon availability and other
relevant factors. Before proceeding for concentration, ores are graded and
crushed to reasonable size.
Removal
of the unwanted materials (e.g., sand, clays, etc.) from the ore is known as concentration, dressing or benefaction.
It involves several steps and selection of these steps depends upon the
differences in physical properties of the compound of the metal present and
that of the gangue. The type of the
metal, the available facilities and the environmental factors are also taken
into consideration. Some of the important procedures are described below.
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6.2.1 Hydraulic
Washing
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This
is based on the differences in gravities of the ore and the gangue particles. It is therefore a
type of gravity separation. In one
such process,
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an
upward stream of running water is used to wash the powdered ore. The lighter
gangue particles are washed away and the heavier ores are left behind.
belt which passes over a magnetic roller
(Fig.6.1). Fig. 6.1: Magnetic
separation (schematic)
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6.2.3
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Froth
Floatation
Method
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This method has been in use
for removing gangue from sulphide ores. In this process, a suspension of the
powdered ore is made with water. To it, collectors
and froth stabilisers are added.
Collectors (e. g., pine oils, fatty
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6.2.2 Magnetic
This is based on differences in Separation magnetic properties of the ore components.
If either the ore or the gangue (one of these two) is capable of being
attracted by a magnetic field, then such separations are carried out (e.g., in
case of iron ores). The ground ore is carried on a conveyer
acids, xanthates, etc.) enhance non-wettability of the mineral particles and froth stabilisers ( e. g., cresols, aniline) stabilise the froth.
The mineral particles become wet
by oils while the gangue particles by water. A rotating paddle agitates the
mixture and draws air in it. As a result, froth is formed which carries the
mineral particles. The froth is light and is skimmed off. It is then dried for
recovery of the ore particles.
Sometimes, it is possible to
separate two sulphide ores by adjusting proportion of oil to water or by using
‘depressants’. For example, in case
of an ore containing ZnS and PbS, the depressant used is NaCN. It selectively
prevents ZnS from coming to the Fig.
6.2: Froth floatation process (schematic) froth
but allows PbS to come with the froth.
The Innovative Washerwoman
One can do
wonders if he or she has a scientific temperament and is attentive to
observations. A washerwoman had an innovative mind too. While washing a miner’s
overalls, she noticed that sand and similar dirt fell to the bottom of the
washtub. What was peculiar, the copper bearing compounds that had come to the
clothes from the mines, were caught in the soapsuds and so they came to the
top. One of her clients was a chemist, Mrs. Carrie Everson. The washerwoman
told her experience to Mrs. Everson. The latter thought that the idea could be
used for separating copper compounds from rocky and earth materials on large
scale. This way an invention was born. At that time only those ores were used
for extraction of copper, which contained large amounts of the metal. Invention
of the Froth Floatation Method made
copper mining profitable even from the low-grade ores. World production of
copper soared and the metal became cheaper.
6.2.4
Leaching Leaching is often used
if the ore is soluble in some suitable solvent.
The following examples
illustrate the procedure:
(a) Leaching of alumina from bauxite
The principal ore of aluminium,
bauxite, usually contains SiO2, iron oxides and titanium oxide (TiO2) as impurities. Concentration is carried out
by digesting the powdered ore with a concentrated solution of NaOH at 473 – 523
K and 35 – 36 bar pressure. This way, Al2O3 is
leached out as sodium aluminate (and SiO2 too as sodium
silicate) leaving the impurities behind:
Al2O3(s) + 2NaOH(aq) + 3H2O(l)
→ 2Na[Al(OH) 4](aq) (6.1)
The aluminate in solution is
neutralised by passing CO2 gas and hydrated Al2O3
is precipitated. At this stage, the solution is seeded with freshly prepared
samples of hydrated Al2O3 which induces the
precipitation:
2Na[Al(OH)4](aq) + CO2(g) → Al2O3.xH2O(s)
+ 2NaHCO3 (aq) (6.2)
The sodium silicate remains in
the solution and hydrated alumina is filtered, dried and heated to give back
pure Al2O3:
1470 K
Al2O3.xH2O(s)
Al2O3(s)
+ xH2O(g) (6.3)
Al2O3(s)
+ xH2O(g) (6.3)
(b) Other examples
In the metallurgy of silver and
that of gold, the respective metal is leached with a dilute solution of NaCN or
KCN in the presence of air (for O2)
from which the metal is obtained later by replacement :
4M(s) + 8CN–(aq)+
2H2O(aq)
+ O2(g) → 4[M(CN)2]– (
aq) +
4OH–(aq) (M= Ag or Au) (6.4) 2[M(CN)2 ]−(aq) +Zn(s) →[Zn(CN)4]2− (aq)+2M(s) (6.5)
Intext Questions
6.1 Which
of the ores mentioned in Table 6.1 can be concentrated by magnetic separation
method?
6.2 What
is the significance of leaching in the extraction of aluminium?
6.3 Extraction of Crude Metal
from
Concentrated
Ore
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The concentrated ore must
be converted into a form which is suitable for reduction. Usually the
sulphide ore is converted to oxide before reduction. Oxides are easier to
reduce (for the reason see box). Thus isolation of metals from concentrated
ore involves two major steps viz.,
(a)
conversion to oxide, and
(b)
reduction of the oxide to metal.
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(a) Conversion to oxide
(i)
Calcination: Calcinaton involves heating when the volatile
matter escapes leaving behind the metal oxide:
Fe2O3.xH2O(s)
Fe2O3
(s) + xH2O(g) (6.6)
ZnCO3 (s)
ZnO(s) + CO2(g) (6.7)
CaCO3.MgCO3(s) CaO(s) + MgO(s ) + 2CO2(g) (6.8)
(ii) Roasting: In roasting, the ore is heated in a regular
supply of air in a furnace at a temperature below the melting point of the
metal. Some of the reactions involving sulphide ores are:
2ZnS + 3O2
→ 2ZnO + 2SO2 (6.9)
2PbS + 3O2 →
2PbO + 2SO2 (6.10)
2Cu2S
+ 3O2 →
2Cu2O + 2SO2 (6.11)
The sulphide
ores of copper are heated in reverberatory furnace. If the ore contains iron,
it is mixed with silica before heating. Iron oxide ‘slags of’* as iron silicate and copper is produced in the
form of copper
Fig. 6.3: A section of a modern matte which contains
Cu2
S and FeS.
reverberatory furnace
FeO + SiO2
→ FeSiO3 (6.12)
(slag)
The SO2 produced is utilised for manufacturing
H2SO4 .
(b) Reduction of oxide to the metal
Reduction of the metal oxide usually involves
heating it with some other substance acting as a reducing agent (C or CO or
even another metal). The reducing agent (e.g., carbon) combines with the oxygen
of the metal oxide.
MxOy
+ yC → xM + y CO (6.13)
Some metal oxides get reduced
easily while others are very difficult to be reduced (reduction means electron
gain or electronation). In any case, heating is required. To understand the
variation in the temperature requirement for thermal reductions (pyrometallurgy) and to predict which
element will suit as the reducing agent for a given metal oxide (MxOy),
Gibbs energy interpretations are made. 6.4 Some basic concepts of thermodynamics
help us in understanding the Thermodynamic theory of metallurgical
transformations. Gibbs energy is the most Principles
of significant term here.The
change in Gibbs energy, ΔG
for any process at any specified temperature, is described by the equation:
Metallurgy ΔG = ΔH – TΔS (6.14)
where, ΔH is the enthalpy change and ΔS is the entropy change for the
process. For any reaction, this change could also be explained through the
equation:
ΔGV = – RTlnK (6.15)
where, K is the equilibrium constant of the ‘reactant – product’ system at the
temperature,T. A negative ΔG
implies a +ve K in equation 6.15. And this can happen only when reaction
proceeds towards products. From these facts we can make the following
conclusions:
* During metallurgy, ‘flux’ is added which combines with ‘gangue’ to form
‘slag’. Slag separates more easily from the ore than the gangue. This way,
removal of gangue becomes easier.
1.
When the value of ΔG
is negative in equation 6.14, only then the reaction will proceed. If ΔS is positive, on increasing
the temperature (T), the value of TΔS
would increase (ΔH
< TΔS) and then ΔG will become –ve.
2.
If reactants and products of two reactions are
put together in a system and the net ΔG
of the two possible reactions is –ve, the overall reaction will occur. So the
process of interpretation involves coupling of the two reactions, getting the
sum of their ΔG
and looking for its magnitude and sign.
Such coupling is easily understood through Gibbs energy (ΔGV) vs T plots for formation of the oxides
(Fig. 6.4).
Ellingham Diagram
The graphical representation of
Gibbs energy was first used by H.J.T.Ellingham. This provides a sound basis for
considering the choice of reducing agent in the reduction of oxides. This is
known as Ellingham Diagram. Such diagrams help us in predicting the feasibility
of thermal reduction of an ore. The criterion of feasibility is that at a given
temperature, Gibbs energy of the reaction must be negative.
(a)
Ellingham diagram normally consists of plots of ΔfGV vs T for formation of oxides of elements i.e., for the reaction,
2xM(s) + O2(g) → 2MxO(s)
In this reaction, the gaseous
amount (hence molecular randomness) is decreasing from left to right due to the
consumption of gases leading to a –ve value of ΔS which
changes the sign of the second term in equation (6.14). Subsequently ΔG shifts towards higher side
despite rising T (normally, ΔG decreases
i.e., goes to lower side with increasing temperature). The result is +ve slope
in the curve for most of the reactions shown above for formation of MxO(s).
(b)
Each plot is a straight line except when some change
in phase (s→liq or liq→g) takes place. The temperature at
which such change occurs, is indicated by an increase in the slope on +ve side
(e.g., in the Zn, ZnO plot, the melting is indicated by an abrupt change in the
curve).
(c)
There is a point in a curve below which ΔG is negative (So MxO is
stable). Above this point, MxO will
decompose on its own.
(d)
In an Ellingham diagram, the plots of ΔGV for oxidation (and therefore reduction of the corresponding
species) of common metals and some reducing agents are given. The values of Δf GV, etc.(for formation of oxides) at different temperatures are
depicted which make the interpretation easy.
(e)
Similar diagrams are also constructed for sulfides and
halides and it becomes clear why reductions of MxS is
difficult. There, the Δf GV of MxS is not
compensated.
Limitations of Ellingham Diagram
1.
The graph simply indicates whether a reaction is
possible or not i.e., the tendency of reduction with a reducing agent is
indicated. This is so because it is based only on the thermodynamic concepts.
It does not say about the kinetics of the reduction process (Cannot answer
questions like how fast it could be ?).
2.
The interpretation of ΔGV is based on K (ΔGV = – RT lnK). Thus it is presumed that the reactants and
products are in equilibrium:
MxO + Ared l xM + AOox
This is not always true because the reactant/product
may be solid. [However it explains how the reactions are sluggish when every
species is in solid state and smooth when
the ore melts down.It is interestng to note here that ΔH (enthalpy change) and the ΔS ( entropy change) values for any chemical reaction remain nearly constant even on varying temperature. So the only dominant variable in equation(6.14) becomes T. However, ΔS depends much on the physical state of the compound. Since entropy depends on disorder or randomness in the system, it will increase if a compound melts (s→l) or vapourises (l→g) since molecular randomness increases on changing the phase from solid to liquid or from liquid to gas].
the ore melts down.It is interestng to note here that ΔH (enthalpy change) and the ΔS ( entropy change) values for any chemical reaction remain nearly constant even on varying temperature. So the only dominant variable in equation(6.14) becomes T. However, ΔS depends much on the physical state of the compound. Since entropy depends on disorder or randomness in the system, it will increase if a compound melts (s→l) or vapourises (l→g) since molecular randomness increases on changing the phase from solid to liquid or from liquid to gas].
The reducing agent forms its oxide
when the metal oxide is reduced. The role of reducing agent is to provide ΔGV negative
and large enough to make the sum of ΔGV of the two reactions (oxidation of the
reducing agent and reduction of the metal oxide) negative.
As we know, during
reduction, the oxide of a metal decomposes:
MxO(s) → xM (solid or liq) + 
O2 (g) (6.16)
O2 (g) (6.16)
The reducing agent takes away the
oxygen. Equation 6.16 can be visualised as reverse of the oxidation of the
metal. And then, the Δf GV value is written in the usual way:
xM(s or l) + 
O2(g) →
MxO(s) [ΔGV (M,MxO)] (6.17)
O2(g) →
MxO(s) [ΔGV (M,MxO)] (6.17)
If reduction
is being carried out through equation 6.16, the oxidation of the reducing agent
(e.g., C or CO) will be there:
C(s) + 
O2(g) →
CO(g) [ΔG(C, CO)] (6.18)
O2(g) →
CO(g) [ΔG(C, CO)] (6.18)
CO(g) + 
O2(g) →
CO2(g) [ΔG(CO, CO2)] (6.19)
O2(g) →
CO2(g) [ΔG(CO, CO2)] (6.19)
If carbon is
taken, there may also be complete oxidation of the element to CO2:
C(s) + 
O 2(g) → CO2(g) [ ΔG(C, CO2)] (6.20)
On subtracting
equation 6.17 [it means adding its negative or the reverse form as in equation
6.16] from one of the three equations, we get:
MxO(s)
+ C(s) → xM(s or l) +
CO(g) (6.21) MxO(s)
+ CO(g) → xM(s or l) +
CO2(g) (6.22)
MxO(s) + 
C(s) →
xM(s or l) + 
CO2(g) (6.23)
C(s) →
xM(s or l) + CO2(g) (6.23)
These reactions describe the
actual reduction of the metal oxide, MxO that we want to
accomplish. The ΔrG0 values for these reactions in general,
can be obtained by similar subtraction of the corresponding Δf G0 values.
As we have
seen, heating (i.e., increasing T) favours a negative value of ΔrG0. Therefore, the temperature is chosen such that the sum
of ΔrG0 in the two combined redox process is
negative. In ΔrG0 vs T plots, this is indicated by the
point of intersection of the two curves (curve for MxO
and that for the oxidation of the reducing substance). After that point, the ΔrG0 value becomes more negative for the combined process
including the reduction of MxO. The difference in the two ΔrG0 values after that point determines whether reductions
of the oxide of the upper line is feasible by the element represented by the
lower line. If the difference is large, the reduction is easier.
Example 6.1 Suggest a condition under which magnesium could reduce alumina.
Solution
The
two equations are:
(a) Al + O2 → Al2O3 (b)
2Mg +O2 →
2MgO
At the point of intersection of the Al2O3 and
MgO curves (marked “A” in diagram 6.4), the ΔG0 becomes ZERO for the reaction:
Al2 O3 +2Mg → 2MgO + Al
Above that point
magnesium can reduce alumina.
Example 6.2 Although
thermodynamically feasible, in practice, magnesium metal is not used for the reduction
of alumina in the metallurgy of aluminium. Why ?
Solution Temperatures above the point of
intersection of Al2O3 and MgO curves, magnesium can reduce
alumina. But the temperature required would be so high that the process will be
uneconomic and technologically difficult.
Example 6.3 Why
is the reduction of a metal oxide easier if the metal formed is in liquid state
at the temperature of reduction?
Solution The entropy is higher if the
metal is in liquid state than when it is in solid state. The value of entropy
change (ΔS) of the
reduction process is more on +ve side when the metal formed is in liquid state
and the metal oxide being reduced is in solid state. Thus the value of ΔG0 becomes
more on negative side and the reduction becomes easier.
6.4.1
Applications
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(a)
Extraction of iron from its oxides
Oxide ores of iron, after
concentration through calcination/roasting ( to remove water, to decompose
carbonates and to oxidise sulphides ) are mixed with limestone and coke and
fed into a Blast furnace from its
top. Here, the oxide is reduced to the metal. Thermodynamics helps us to
understand how coke reduces the oxide and why this furnace is chosen. One of
the main reduction steps in this process is:
FeO(s) + C(s) →
Fe(s/l) + CO (g) (6.24)
It
can be seen as a couple of two simpler reactions. In one, the reduction of
FeO is taking place and in the other, C is being oxidised to CO:
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FeO(s) → Fe(s) + 
O2(g) [ΔG(FeO, Fe)] (6.25)
O2(g) [ΔG(FeO, Fe)] (6.25)
C(s) + O2 (g) →
CO (g) [ΔG (C, CO)] (6.26)
O2 (g) →
CO (g) [ΔG (C, CO)] (6.26)
When both the
reactions take place to yield the equation (6.23), the net Gibbs energy change
becomes:
ΔG (C, CO) + ΔG (FeO, Fe) = ΔrG (6.27)
Naturally,
the resultant reaction will take place when the right hand side in equation
6.27 is negative. In ΔG0 vs T plot representing reaction
6.25,
the plot goes upward and that representing the change C→CO
( C,CO) goes downward. At temperatures above 1073K
(approx.), the C,CO line comes below the Fe,FeO line [ΔG (C, CO) < ΔG(Fe, FeO)] . So
in this range, coke will be reducing the FeO and will itself be oxidised to CO.
In a similar way the reduction
of Fe3O4 and Fe2O3 at
relatively lower temperatures by CO can be explained on the basis of lower
lying points
oxides (Ellingham diagram)
In the Blast furnace, reduction of
iron oxides takes place in different temperature ranges. Hot air is blown from
the bottom of the furnace and coke is burnt to give temperature upto about
2200K in the lower portion itself. The burning of coke therefore supplies most
of the heat required in the process. The CO and heat moves to upper part of the
furnace. In upper part, the temperature is lower and the iron oxides (Fe2O3
and Fe3O4) coming
from the top are reduced in steps to FeO. Thus, the reduction reactions taking
place in the lower temperature range and in the higher temperature range,
depend on the points of corresponding intersections in the ΔrG0 vs T plots. These reactions can be
summarised as follows:
At 500 – 800 K (lower
temperature range in the blast furnace)–
3 Fe2O3
+ CO → 2 Fe3O4
+ CO2 (6.28)
Fe3O4
+ 4 CO → 3Fe + 4 CO2 (6.29)
Fe2O3
+ CO → 2FeO + CO2 (6.30)
Fig. 6.5: Blast furnace Fe2O3
+ 3 C → 2 Fe + 3 CO
(6.33)
Limestone is added as a
flux and sulphur, silicon
and phosphorus are oxidised and passed into the slag. The
metal is removed and freed from the slag by passing through rollers.
(b) Extraction of copper from cuprous oxide [copper(I)
oxide]
In the graph of ΔrG0 vs T for formation of oxides (Fig. 6.4),
the Cu2O line is almost at the top. So it is quite easy to reduce
oxide ores of copper directly to the metal by heating with coke (both the lines
of C, CO and C, CO2 are at much lower positions in the graph
particularly after 500 – 600K). However most of the ores are sulphide and some
may also contain iron. The sulphide ores are roasted/smelted to give oxides:
2Cu2S
+ 3O2 →
2Cu2O + 2SO2 (6.34)
The oxide can then be
easily reduced to metallic copper using coke:
Cu2O
+ C → 2 Cu + CO (6.35)
In actual
process, the ore is heated in a reverberatory furnace after mixing with silica.
In the furnace, iron oxide ‘slags of’ as iron silicate and copper is produced
in the form of copper matte. This
contains Cu2S and FeS.
FeO + SiO2
→ FeSiO3 (6.36)
(Slag)
Copper matte
is then charged into silica lined convertor. Some silica is also added and hot
air blast is blown to convert the remaining FeS2, FeO and Cu2S/Cu2O
to the metallic copper. Following reactions take place:
2FeS + 3O2
→ 2FeO + 2SO2
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(6.37)
|
FeO + SiO2
→ FeSiO3
|
(6.38)
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2Cu2S
+ 3O2 →
2Cu2O + 2SO2
|
(6.39)
|
2Cu2O
+ Cu2S →
6Cu + SO2
|
(6.40)
|
The solidified copper obtained has
blistered appearance due to the evolution of SO2 and so it is
called blister copper.
(c) Extraction of zinc from zinc oxide
The reduction of zinc oxide is done using coke. The
temperature in this case is higher than that in case of copper. For the purpose
of heating, the oxide is made into brickettes with coke and clay.
nO + C coke, 673 K Zn + CO (6.41)
The metal is distilled off and collected by rapid
chilling.
Intext Questions
6.3 The
reaction,
Cr2 O3 +
2 Al → Al2 O3 +
2 Cr (ΔG0 = – 421 kJ)
is thermodynamically feasible as is apparent from the
Gibbs energy value. Why does it not take place at room temperature?
6.4 Is
it true that under certain conditions, Mg can reduce SiO2
and Si can reduce MgO? What are those conditions?
6.5 We have seen how principles of
thermodyamics are applied to
Electrochemical pyrometallurgy. Similar principles
are effective in the reductions of metal Principles
of ions in solution or molten
state. Here they are reduced by electrolysis or by adding some reducing
element.
Metallurgy
In the reduction of a molten
metal salt, electrolysis is done. Such methods are based on electrochemical
principles which could be understood through the equation,
ΔG0 = – nE0F (6.42) here n is the number of electrons and E0 is the electrode potential of the redox couple formed
in the system. More reactive metals have large negative values of the electrode
potential. So their reduction is difficult. If the difference of two E0 values corresponds to a positive E0 and consequently negative ΔG0 in
equation 6.42, then the less reactive metal will come out of the solution and
the more reactive metal will go to the solution, e.g.,
Cu2+
(aq) + Fe(s) →
Cu(s) + Fe2+ (aq) (6.43)
In simple
electrolysis, the Mn+ ions are discharged at negative electrodes
(cathodes) and deposited there. Precautions are taken considering the
reactivity of the metal produced and suitable materials are used as electrodes.
Sometimes a flux is added for making the molten mass more conducting.
In the metallurgy
of aluminium, purified Al2O3 is mixed with Na3AlF6 or
CaF which lowers the melting point of the mix and brings conductivity. The
fused matrix is electrolysed. Steel cathode and graphite anode are used. The
graphite anode is useful here for reduction to the metal. The overall reaction
may be taken as:
2Al2O3
+ 3C → 4Al + 3CO2 (6.44)
This process of
electrolysis is widely known as Hall-Heroult
process.
The
electrolysis of the molten mass is carried out in an electrolytic cell using
carbon electrodes. The oxygen liberated at anode reacts with the carbon of
anode producing CO and CO2. This way for each kg of aluminium
produced, about 0.5 kg of carbon anode is burnt away. The electrolytic
reactions are: Cathode: Al3+ (melt) + 3e–
→ Al(l) (6.45)
–
Anode: C(s) + O (melt) →
CO(g) + 2e (6.46)
C(s) + 2O2–
(melt) → CO2
(g) + 4e– (6.47)
Copper from Low Grade Ores and Scraps
Copper is extracted by hydrometallurgy from low grade ores. It
is leached out using acid or bacteria. The solution containing Cu2+ is
treated with scrap iron or H2 (equations 6.42; 6.48).
Cu2+(aq)
+ H2(g) →
Cu(s) + 2H+ (aq) (6.48)
Example 6.4 At a site, low grade copper ores are available and zinc and iron scraps are also available. Which of the two scraps would be more suitable for reducing the leached copper ore and why?
Solution Zinc being
above iron in the electrochemical series (more reactive metal is zinc), the
reduction will be faster in case zinc scraps are
used. But zinc is costlier metal than
iron so using iron scraps will be advisable and advantageous.
6.6
Oxidation Reduction
|
Besides reductions, some extractions are based on oxidation
particularly for non-metals. A very common example of extraction based on
oxidation is the extraction of chlorine from brine (chlorine is abundant
|
in sea
water as common salt) .
2Cl–(aq)
+ 2H2O(l) →
2OH–(aq) + H2(g) + Cl2(g) (6.49)
The ΔG0
for this reaction is + 422 kJ. When it is converted to E0 (using ΔG0 = – nE0F), we get E0 = – 2.2 V. Naturally, it will require an external
e.m.f. that is greater than 2.2 V. But the electrolysis requires an excess
potential to overcome some other hindering reactions. Thus, Cl2
is obtained by electrolysis giving out H2 and aqueous NaOH as
byproducts. Electrolysis of molten NaCl is also carried out. But in that case,
Na metal is produced and not NaOH.
As studied
earlier, extraction of gold and silver involves leaching the metal with CN–.
This is also an oxidation reaction (Ag →
Ag+
or Au →
Au+).
The metal is later recovered by displacement method.
4Au(s) + 8CN–(aq)
+ 2H2O(aq) + O2(g) →
4[Au(CN)2]–(aq)
+ 4OH–(aq) (6.50)
2[Au(CN)2]–(aq) + Zn(s) → 2Au(s) + [Zn(CN)4]2– (aq) (6.51) In this reaction zinc acts as a
reducing agent. 6.7 Refining A
metal extracted by any method is usually contaminated with some impurity. For
obtaining metals of high purity, several techniques are
used depending upon the
differences in properties of the metal and the impurity. Some of them are
listed below.
(a) Distillation (b) Liquation
(c) Electrolysis (d) Zone refining
(e) Vapour
phase refining (f)
Chromatographic methods
These are described in
detail here.
(a)
Distillation
Distillation
This is very useful for low boiling metals like zinc and
mercury. The impure metal is evaporated to obtain the pure metal as distillate.
(b)
Liquation
In this method a low melting metal like tin can be made to
flow on a sloping surface. In this way it is separated from higher melting
impurities.
(c)
Electrolytic
refining
In this method, the impure metal
is made to act as anode. A strip of the same metal in pure form is used as
cathode. They are put in a suitable electrolytic bath containing soluble salt
of the same metal. The more basic metal remains in the solution and the less
basic ones go to the anode mud. This process is also explained using the
concept of electrode potential, over potential, and Gibbs energy which you have
seen in previous sections. The reactions are: Anode: M → Mn+ + ne–
Cathode: Mn+
+ ne– →
M (6.52)
Copper is refined using an electrolytic method.
Anodes are of impure copper and pure copper strips are taken as cathode. The
electrolyte is acidified solution of copper sulphate and the net result of
electrolysis is the transfer of copper in pure form from the anode to the
cathode:
Anode: Cu →
Cu2+
+ 2 e–
Cathode: Cu2+
+ 2e– → Cu (6.53)
Impurities from
the blister copper deposit as anode mud which contains antimony, selenium,
tellurium, silver, gold and platinum; recovery of these elements may meet the
cost of refining. Zinc may also be refined this way.
(d)
Zone refining
This method is based on the
principle that the impurities are more soluble in the melt than in the solid
state of the metal. A circular mobile heater is fixed at one end of a rod of
the impure metal
( Fig. 6.7). The
molten zone moves along with the heater which is moved forward. As the heater
moves forward, the pure metal crystallises out of the melt and the impurities
pass on into the adjacent molten zone. The process is repeated several times
and the heater is moved in the same direction. At one end, impurities get
concentrated. This end is cut off. This method is very useful for producing
semiconductor and other metals of very high purity, e.g., germanium, silicon,
boron, Fig. 6.7: Zone refining
process gallium and indium.
(e)
Vapour phase
refining
In this method, the metal is
converted into its volatile compound and collected elsewhere. It is then
decomposed to give pure metal. So, the two requirements are:
(i)
the metal should form a volatile compound with an
available reagent,
(ii)
the volatile compound should be easily decomposable, so
that the recovery is easy.
Following examples will illustrate this technique.
Mond Process
for Refining Nickel: In this process, nickel is heated in a stream of
carbon monoxide forming a volatile complex, nickel tetracarbonyl:
330 – 350 K
Ni + 4CO 
Ni(CO)4 (6.54)
Ni(CO)4 (6.54)
The carbonyl is subjected to higher
temperature so that it is decomposed giving the pure metal:
Ni(CO)4
450–470K Ni + 4CO (6.55)
van Arkel
Method for Refining Zirconium or Titanium: This method is very useful for removing all
the oxygen and nitrogen present in the form of impurity in certain metals like
Zr and Ti. The crude metal is heated in an evacuated vessel with iodine. The metal
iodide being more covalent, volatilises:
Zr + 2I2 → ZrI4 (6.56)
The metal iodide is decomposed on
a tungsten filament, electrically heated to about 1800K. The pure metal is thus
deposited on the filament.
ZrI4
→ Zr + 2I2 (6.57)
(f) Chromatographic
methods
This method is based on the principle
that different components of a mixture are differently adsorbed on an
adsorbent. The mixture is put in a liquid or gaseous medium which is moved
through the adsorbent.
Different components are adsorbed at different levels on the
column. Later the adsorbed components are removed (eluted) by using suitable
solvents (eluant). Depending upon the physical state of the moving medium and
the adsorbent material and also on the process of passage of the moving medium,
the chromatographic method* is given the
name. In one such method the column of Al2O3
is prepared in a glass tube and the moving medium containing a solution of the
components is in liquid form. This is an example of column chromatography. This is very useful for purification of the
elements which are available in minute quantities and the impurities are not
very different in chemical properties from the element to be purified. There
are several chromatographic techniques such as paper chromatography, column
chromatography, gas chromatography, etc. Procedures followed in column
chromatography have been depicted in Fig. 6.8.
Fig. 6.8: Schematic diagrams showing column
chromatography
* Looking it the other way, chromatography in general, involves a mobile
phase and a stationary phase. The sample or sample extract is dissolved in a
mobile phase. The mobile phase may be a gas, a liquid or a supercritical fluid.
The stationary phase is immobile and immiscible (like the Al2O3 column in the
example of column chromatography above). The mobile phase is then forced
through the stationary phase. The mobile phase and the stationary phase are
chosen such that components of the sample have different solubilities in the
two phases. A component which is quite soluble in the stationary phase takes
longer time to travel through it than a component which is not very soluble in
the stationary phase but very soluble in the mobile phase. Thus sample components
are separated from each other as they travel through the stationary phase.
Depending upon the two phases and the way sample is inserted/injected, the
chromatographic technique is named. These methods have been described in detail
in Unit 12 of Class XI text book (12.8.5).
6.8 Uses of Aluminium, Copper,
Zinc and Iron
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Aluminium foils are used as
wrappers for chocolates. The fine dust of the metal is used in paints and
lacquers. Aluminium, being highly reactive, is also used in the extraction of
chromium and manganese from their oxides. Wires of aluminium are used as electricity
conductors.
Alloys containing
aluminium, being light, are very useful.
|
Copper is used
for making wires used in electrical industry and for water and steam pipes. It
is also used in several alloys that are rather tougher than the metal itself,
e.g., brass (with zinc), bronze (with tin) and coinage alloy (with nickel).
Zinc is used
for galvanising iron. It is also used in large quantities in batteries, as a
constituent of many alloys, e.g., brass, (Cu 60%, Zn 40 %) and german silver
(Cu 25-30%, Zn 25-30%, Ni 40–50%). Zinc dust is used as a reducing agent in the
manufacture of dye-stuffs, paints, etc.
Cast iron, which is the most important
form of iron, is used for casting stoves, railway sleepers, gutter pipes ,
toys, etc. It is used in the manufacture of wrought iron and steel. Wrought
iron is used in making anchors, wires, bolts, chains and agricultural
implements. Steel finds a number of uses. Alloy steel is obtained when other
metals are added to it. Nickel steel is used for making cables, automobiles and
aeroplane parts, pendulum, measuring tapes, chrome steel for cutting tools and
crushing machines, and stainless steel for cycles, automobiles, utensils, pens,
etc.
Summary
Metals are
required for a variety of purposes. For this, we need their extraction from the
minerals in which they are present and from which their extraction is
commercially feasible.These minerals are known as ores. Ores of
the metal are associated with many impurities. Removal of these impurities to
certain extent is achieved in concentration steps. The
concentrated ore is then treated chemically for obtaining the metal. Usually
the metal compounds (e.g., oxides, sulphides) are reduced to the metal. The reducing
agents used are carbon, CO or even some metals. In these reduction processes,
the thermodynamic and electrochemical concepts
are given due consideration. The metal oxide reacts with a reducing agent; the
oxide is reduced to the metal and the reducing agent is oxidised. In the two
reactions, the net Gibbs energy change is negative, which becomes more negative
on raising the temperature. Conversion of the physical states from solid to
liquid or to gas, and formation of gaseous states favours decrease in the Gibbs
energy for the entire system. This concept is graphically displayed in plots of
ΔG0 vs T (Ellingham diagram) for such oxidation/reduction
reactions at different temperatures. The concept of electrode potential is
useful in the isolation of metals (e.g., Al, Ag, Au) where the sum of the two
redox couples is +ve so that the Gibbs energy change is negative. The metals
obtained by usual methods still contain minor impurities. Getting pure metals
require refining. Refining process depends upon the differences in properties
of the metal and the impurities. Extraction of aluminium is usually carried out
from its bauxite ore by leaching it with NaOH. Sodium aluminate, thus formed,
is separated and then neutralised to give back the hydrated oxide, which is
then electrolysed using cryolite as a flux. Extraction of iron is done by
reduction of its oxide ore in blast furnace. Copper is extracted by smelting
and heating in a reverberatory furnace. Extraction of zinc from zinc oxides is
done using coke. Several methods are employed
Exercises
6.1 Copper can be extracted by
hydrometallurgy but not zinc. Explain.
6.2 What is the role of depressant in
froth floatation process?
6.3
Why
is the extraction of copper from pyrites more difficult than that from its
oxide ore through reduction?
6.4 Explain: (i) Zone refining ( ii) Column chromatography.
6.5 Out of C and CO, which is a better
reducing agent at 673 K ?
6.6
Name
the common elements present in the anode mud in electrolytic refining of
copper. Why are they so present ?
6.7
Write
down the reactions taking place in different zones in the blast furnace during
the extraction of iron.
6.8 Write chemical reactions taking
place in the extraction of zinc from zinc blende.
6.9 State the role of silica in the
metallurgy of copper.
6.10 What is meant by the term
“chromatography”?
6.11 What
criterion is followed for the selection of the stationary phase in
chromatography?
6.12 Describe a method for refining
nickel.
6.13
How
can you separate alumina from silica in a bauxite ore associated with silica?
Give equations, if any.
6.14 Giving examples, differentiate
between ‘roasting’ and ‘calcination’.
6.15 How is ‘cast iron’ different from
‘pig iron”?
6.16 Differentiate between “minerals”
and “ores”.
6.17 Why copper matte is put in silica lined converter?
6.18 What is the role of cryolite in the
metallurgy of aluminium?
6.19 How is leaching carried out in case
of low grade copper ores?
6.20 Why is zinc not extracted from zinc
oxide through reduction using CO?
6.21 The value of ΔfG0 for formation of Cr2 O3 is – 540
kJmol−1and that of
Al2 O3 is – 827 kJmol−1. Is the reduction of Cr2 O3 possible
with Al ?
6.22 Out of C and CO, which is a better reducing
agent for ZnO ?
6.23
The
choice of a reducing agent in a particular case depends on thermodynamic
factor. How far do you agree with this statement? Support your opinion with two
examples.
6.24
Name
the processes from which chlorine is obtained as a by-product. What will happen
if an aqueous solution of NaCl is subjected to electrolysis?
6.25 What is the role of graphite rod in
the electrometallurgy of aluminium?
6.27
Outline
the principles of refining of metals by the following methods: (i) Zone refining
(ii)
Electrolytic
refining
(iii)
Vapour phase refining
6.28 Predict conditions under which Al
might be expected to reduce MgO.
( Hint: See Intext question 6.4)
Answers
to Some Intext Questions
6.1
Ores in which one of the components (either the
impurity or the actual ore) is magnetic can be concentrated, e.g., ores
containing iron (haematite, magnetite, siderite and iron pyrites).
6.2
Leaching is significant as it helps in removing the
impurities like SiO2, Fe2O3, etc. from
the bauxite ore.
6.3
Certain amount of activation energy is essential even
for such reactions which are thermodynamically feasible, therefore heating is
required.
6.4
Yes, below 1350°C Mg can reduce Al2O3 and above
1350°C, Al can reduce MgO.
This can be
inferred from ΔGV Vs T plots (Fig. 6.4).
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