Metallurgy, Chemistry tutorial


As we are familiar that metals as a group of elements have acquired a unique significance in the modern world. Though, nature doesn't usually offer us metals in the Free State. Metals generally take place in nature in combined state as ores mixed by other earthy materials. The branch of science dealing by the methods of extraction of metals from their ores is known as metallurgy.

Occurrence of Metals:

The earth's crust and sea are the two major sources of metals. In the earth's crust metals take place both in the combined state in the form of minerals and also in the native or free state. Earth's crust is the outermost portion of the earth, which has an average thickness of around 17 km. The crust is thinner under the oceans and thicker beneath the continents. The minerals from which the extraction of any metal is chemically feasible and economically competitive are termed as ores of that metal. Metals occur in broadly varying quantities in the earth's crust. The relative profusion of the most common elements in the earth's crust is provided in the table shown below. We might note that around 75% of the earth's crust is comprised of non-metals, oxygen and silicon. The relative abundance of only three industrially significant metals, that is, aluminium, iron and magnesium is more than 2%. The profusion of most other helpful metals in the earth's crust is extremely low. Thus, if the metals had been uniformly distributed in earth's crust, it would have not been possible to extract them. However luckily, the metals usually in the form of their minerals, are unevenly distributed and are accumulated and some positions, making their extraction easier. Such accumulations of minerals are known as mineral deposits. Generally, the mineral is covered by a layer of soil, termed as overburden. The thickness of overburden might differ from a few meters as in case of iron ore to thousands of meters as in the case of deposits of gold. The mineral deposit is brought to the surface through mining.

Table: Relative abundance of various elements in earth's crust

2300_Relative abundance of various elements in earth crust.jpg

Minerals are solid substances differing in the chemical composition, colour, density, lustre, hardness and other features. Based on the chemical composition, minerals can be categorized into given groups:

Native Minerals:

These minerals have the metal in free or elemental state, example: copper, gold, silver, platinum and iron. The metals are generally found mixed with clay, sand and so on. At times lumps of nearly pure metals are as well found. These lumps are termed as nuggets. Native iron is of meteorite origin and its occurrence is rare. Deposits of native iron are found in the Greenland.

Sulphide Minerals:

In such minerals metals are present as their sulphides. For illustration, iron pyrites (FeS2), calcocite (Cu2S), chalcopyrite (CuFeS2), zinc blende (ZnS), argentite (Ag2S), cinnabar (HgS), galena (PbS), millerite (NiS) and so on.

Oxide Minerals:

These minerals comprise of oxides of metals, which are made up of either by oxidation of sulphide minerals or through direct oxidation of metals. Highly electropositive metals, like Al and Mg, take place only as oxides instead of as sulphides. Some significant oxide minerals are haematite (Fe2O3), magnetite (Fe3O4), bauxite (Al2O3•2H2O), cassiterite (SnO2), cuprite (Cu2O), zincite (ZnO), rutile (TiO2), pyrolusite (MnO2), chromite (FeO Cr2O3), uraninite or pitchblende (2UO3 .UO2) and so on.


In these minerals, the metals are present as their oxosalts, like carbonates, nitrates, sulphates, phosphates, borates and silicates. Some significant minerals of this group are siderite (FeCO3), magnesite (MgCO3), dolomite (MgCO3.CaCO3), cerussite (PbCO3), malachite [CuCO3. Cu(OH)2], calamine (ZnCO3), barytes (BaSO4), gypsum (CaSO4.2H2O), epsomite (MgSO4. 7H2O), anglesite (PbSO4), soda nitre (NaNO3), monazite (LaPO4.CePO4.NdPO4.PrPO4.Th3(PO4)4), spodumene (LiAlSi2O6), zircon (ZrSiO4), beryl (Be3Si6O18) and so on. Phosphate minerals are, in general, rare and take place in low concentrations. Silicate minerals are rich in nature. Though, the extraction of metals from silicates is difficult and the cost of extraction is incredibly high. Thus, only the less common metals like lithium are extracted from the silicate minerals.

Halide Minerals:

Highly electropositive alkali and alkaline earth metals tend to make halide salts, which being soluble in the water are washed away to the oceans due to leaching of the top soil. Though, most of the deposits of halide minerals are as well found beneath the soil. Some significant halide minerals are rock salt (NaCl), sylvine (KC1), horn silver (AgCl), carnallite (KCl.MgCl2-6H2O), fluorspar (CaF2) and cryolite (AlF3.3NaF).

Ores as mined, usually have variable amounts of unwanted minerals like silica, clay, granite and so on. These unwanted materials are known as gangue. The proportion of the desired metal in the ore should be adequately high in such a way that the extraction of metal is chemically feasible and economically competitive. The ores of very low concentration are employed only if they can be processed simply and inexpensively or if the metal produced is scarce. The lower limit of the percentage of metal in mineral below that extraction becomes unprofitable based on the value of the metal. Therefore, ores having 1% tin are often worked on to get tin and ores having 5% tin are considered rich deposits of tin. If gold is present to the extent of even 0.0015%, this is considered worth extraction. On the other hand, iron and aluminium will not be worth extracting except they have 30% or more of the metal.

As stated earlier, in addition to the earth's crust, oceans as well give a huge storehouse of minerals in which the metals take place primarily as soluble sulphates and halides. This is estimated that one cubic kilometer of sea water consists of 1 million tons of magnesium, 1,500 tons of strontium and 5 tons each of gold, copper, manganese, zinc and lead. Magnesium is already being extracted from the sea water. In future, greater attention will be paid to sea as a source of raw materials whenever supplies of ore deposits on land are depleted. In addition to sea water, nodules or lumps regarding the size of an orange have been found on sea bed at depths of 4,000 to 5,000 meters. The nodules are comparatively rich in manganese (25%) and iron (15%). In recent times technology for deep sea mining of such nodules has been developed.

Form the discussion above; it must be clear to you that there is a relationship between the reactivity of metals and the form in which they take place in nature. Reactive metals take place in nature in the form of their compounds like oxides, sulphides, halides and oxosalts. On the other hand coinage and noble metals having instead low reactivity are found in nature in both combined and also native states.

Beneficiation of Ores:

Most of the ores available in nature have huge amounts of impurities, that is, gangue. Direct extraction of metals from the ores via metallurgical processes is uneconomical and technically difficult. Thus, the ores must be processed first by some cheaper methods that eliminate the gangue partly or completely. The pretreatment of ores via cheaper methods, based mostly on physical properties and without bringing out any major chemical change in the ore is termed as beneficiation or concentration of the ore or ore dressing. Beneficiation of ores yields in saving the cost of the transportation, fuel, fluxing agents and the increased production.

The processes employed for beneficiation of ores are based on differences in such properties of ores and gangue as lustre, colour, size, density, and wettability via water or oil. The simplest process of ore beneficiation comprises of hand picking of ore particles that is based on difference in colour, lustre or shape and size of ore particles and gangue. Hand picking can be adopted in the areas where labor is cheap. Though, this process is outdated and is practiced only in very particular cases whenever other methods are not possible, example: hand picking of diamonds from gravel and clay. Significant methods of beneficiation of ores are gravity separation, magnetic separation and froth flotation.

Gravity Separation:

This is one of the simplest processes of concentration of ores. This is based on the difference in the specific gravities of the ore and gangue. In this process, the crushed ore is kept on the top of a sloping table that is made to vibrate. A stream of water is passed in the direction perpendicular to the slope. The lighter particles are thrown up through vibration and are eliminated by the water stream. The heavier mineral particles settle to the bottom and are collected. This process of gravity separation is termed as tabling. Casseterite or tin-stone, chromite and pitchblende are concentrated through this process.

Modifications of the above process are sink and float method. In this, the powdered ore is suspended in a liquid whose specific gravity is intermediate between the densities of gangue and the ore. The lighter material floats and the heavier material sinks. In this process, the difficulty is in determining a liquid of the proper specific gravity. A solution of calcium chloride in water is frequently used. Suspensions of the sand in water giving liquids of specific gravities up to 3.2 are as well employed. Though, due to technical problems this process is rarely employed in concentration of low grade ores, however it is broadly employed in the cleaning coal.

Magnetic Separation:

This method is based on the difference in magnetic properties of minerals. If the ore however not the gangue is attracted through a magnetic field, it can be concentrated to yield a sample that is rich in the metal. The pulverized mineral is passed over a rubber belt that moves on a pulley in a magnetic field (figure shown below). The non-magnetic gangue particles fall off in a vertirle position whenever the belt passes over the pulley; however the magnetic ore clings to the belt. Whenever the belt passes out of the influence of the magnetic field the ore drops off. Magnetite (Fe3O4), haematite (Fe2O3), wulframite (FeWO4), chromite (FeO.Cr2O3) and ilmenite (FeO.TiQ2) are some of the minerals that are separated from non-magnetic impurities via this method.

1293_Magnetic separation of ores.jpg

Fig: Magnetic separation of ores

Froth Flotation Process:

Froth flotation method is the most significant method for beneficiation of ores. This method has made possible the beneficiation of the low grade ores that could not be processed earlier. The method of froth flotation is broadly used to concentrate sulphide ores. Though, most of the oxide ores can as well be concentrated by this method. This is based on the difference in wettability of different minerals. In this method, the ore is finely ground to provide a thick pulp having 30 to 40% solids. A small amount of pine oil, oleic acid or cresylic acid, that cause frothing, is added to the pulp. A substance that is capable of repelling water from the surface of mineral and therefore promotes attachment of mineral particles to air bubbles is as well added to pulp. This substance is termed as collector. Sodium ethyl xanthate, C2H5OCS2Na, is generally employed as a collector in the floating copper, lead and nickel sulphide ores. The other substance known as activator that assists in the action of collector can as well be added. The whole material, that is, the mixture of pulp, frother and collector, is taken in the container and then air is blown. Air bubbles stick to the mineral particles and make them float in the form of a froth that is collected. The gangue is wetted via water and sinks (figure shown below).

1029_Froth flotation process for concentration of sulphide ores.jpg

Fig: Froth flotation process for concentration of sulphide ores

Some of the ores contain more than one mineral; therefore separation of one mineral from the other in addition to separation from the gangue is essential. To accomplish this, a depressing agent or depressor that suppresses the notation of one of the minerals is added. A significant illustration is the concentration of lead-zinc ore. Whenever the ore is concentrated devoid of a depressor, both the lead and zinc sulphides collect in the froth. If a small amount of sodium cyanide or zinc sulphate is added, zinc sulphide is depressed, permitting flotation of lead sulphide. After eliminating the lead sulphide, copper sulphate is added to activate the depressed zinc sulphide and air is blown if zinc sulphide floats. This process is termed as differential flotation.

Reduction to metals:

After the removal of gangue, that is, impurities physically mixed by the metal compounds, the concentrated ore becomes ready for the isolation of metal. In the concentrated ore, metals are present in the form of their compounds. Extraction of metals comprises the reduction of metal compounds to free metals. In common, based on the reactivity of metals, their compounds can be reduced via one or more than one of the three kinds of metallurgical operations. Such operations are pyrometallurgy, hydrometallurgy and electrometallurgy.


In the operation of pyrometallurgy, the concentrated ore is heated to a high temperature and reduction is done by an appropriate reducing agent. The various steps involved in pyrometallurgy are calcinations, roasting and smelting. The concentrated ore is transformed to the metal oxide via calcination or roasting, if it doesn't already exist as an oxide. This is due to the reason that other metal compounds such as sulphides, sulphates, carbonates and so on are difficult to reduce. Finally, the metal oxide is reduced to metal via smelting.


This is the method of heating the concentrated ore in a limited supply of air to a high temperature however below the fusion temperature. In the operation of calcination, volatile constituents of an ore are expelled. The hydroxide and hydrated ores lose their water making metal oxides. In case of carbonate ores, carbon-dioxide is lost and metal oxides are made. 

Al(OH)3 → (1500K) → Al2O3 + 3H2O ↑

CaCo3 → (1300-1550K) → CaO + CO2


Roasting is the method of heating ores in the presence of surplus air and comprises oxidation. It is mainly applied to sulphide ores, which are transformed to oxides or sulphates. Some of the impurities such as sulphides of arsenic and antimony as well get oxidized and volatilized. For illustration:

4FeS2 + 11O2 → 2Fe2O3 + 8SO2

ZnS + 2O2 → ZnSO4

2ZnO + 3O2 → 2ZnO + 2SO2

2As2S3 + 9O2 → 2As2O3 ↑ + 6SO2

Whenever cuprous sulphide is roasted in a limited supply of air, it is partially oxidized to Cu2O, is then reduced to copper via the remaining cuprous sulphide: 

2Cu2S + 3O2 → 2Cu2O + 2SO2

Cu2S + 2Cu2O → 6Cu + SO2

At times, the oxides formed throughout the roasting are unstable and decompose into elements at a fairly high temperature. For illustration: in the roasting of cinnabar, the red sulphide ore of mercury, the oxide formed decomposes to provide the metal:

2HgS + 3O2 → 2HgO + SO2

2HgO → (800K) → 2Hg + SO2


The roasted ore that is generally an oxide is strongly heated with an appropriate reducing agent as an outcome of which the metal is obtained in the molten state. This method is known as smelting. In smelting, an appropriate chemical substance known as flux is as well added. The flux reacts by the gangue which remains after concentration to form a low melting compound known as slag. The liquid metal and the liquid slag are immiscible and are simply separated. Generally the slag is lighter than the liquid metal and can be simply skimmed off from the surface of the molten metal. The gangue usually contains either basic oxides such as CaO, FeO and so on or an acidic oxide such as silica. If the gangue consists of a basic oxide, the flux employed is an acidic oxide such as silica. For gangues having an acidic oxide, a basic flux such as FeO, CaO or lime stone is added.

SiO2 + CaO → CaSiO3

SiO2 + FeO → FeSiO3

We have studied under roasting that HgO can be reduced to mercury via simply heating it to 800K - a temperature that can be easily managed. Most of the oxides can be reduced to free metals via thermal decomposition at extremely high temperatures, however then the procedure becomes extremely expensive. Though, by using an appropriate reducing agent, reduction of metal oxides can be accomplished at much lower temperatures. The choice of a reducing agent is guided via two considerations. First, the reducing agent must be capable to produce the desired metal at a low temperature. The second consideration is the cost of the reducing agent. It must be less costly than the metal to be produced. Carbon in the form of coke is the least expensive reducing agent. Iron, zinc, lead, tin, cadmium, antimony, nickel, cobalt, molybdenum and numerous other metals are formed by carbon reduction of their oxides at temperatures up to 1800K. For illustration: zinc oxide is reduced to zinc:

ZnO(s) + C(s) → Zn(g) + CO(g)

Though, the reactions that take place in a high temperature carbon reduction method are not as simple as illustrated above. In most of the cases, the efficient reducing agent is carbon monoxide, not carbon. This is as both the metal oxide and coke are solids, thus, contact between them is poor and direct reaction is slow:

MO(s) + C(s) → M(l) + CO(g)

Though, carbon monoxide, which is a gas, forms a better contact with the solid metal oxide and the reaction carry on more readily: 

2C (s) + O2 (g) → 2CO (g)

MO (s) + CO (g) → M (l) + CO2 (g)

A few metals like Cr, Mo, W, Ti, Mn, Mg, Al and so on, can be produced theoretically via reduction of their oxides by carbon; however they react with carbon to produce metallic carbides. Thus, reduction with carbon is not a satisfactory process for producing such metals in a pure form. Hydrogen, although more costly than carbon, is employed as a reductant for the extraction of some of these metals, example: Ge, Mo and W:

GeO2 + 2H2 → Ge + 2H2O

MoO3 + 3H2 → Mo + 3H2O

WO3 + 3H2 → W + 3H2O

Though, most of the metals join with hydrogen too to form metal hydrides.

Thus, hydrogen as well can't be employed for the reduction of compounds of these metals. Highly reactive metals such as Na, Mg, Ca and Al are employed to displace these metals from their oxides or halides. Such reactive metals are relatively more expensive reducing agents as they themselves are difficult or costly to manufacture. The reduction of an oxide via aluminium is known as Goldschmidt's aluminothermic method.

Cr2O3 (s) + 2Al (s) → 2Cr (l) + Al2O3 (l)

3MnO2 (s) + 4Al (s) → 3Mn (l) + 2Al2O3 (l)

3BaO (s) + 2Al (s) → 3Ba (l) + Al2O3 (l)

The reactions are highly exothermic generating metals in the molten state. We have already studied that the reaction of Fe2O3 with Al is employed in spot welding of iron pieces. Other oxides commercially reduced via metals comprise UO3 (by Al or Ca), V2O5, MoO3 and WO3 (by Al), Sc2O3, La2O3, ThO2 (by Ca) and Ta2O5 (by Na).

A few metals can be more conveniently produced by reduction of their halides like TiCl4, ZrCl4, HfCl4, LaCl3, LaCl3, UF4  and so on by Mg, Ca or Na. This method is termed as Kroll's process.

TiCl4 (g) + Mg (l) → (He or Ar at 1000K) → Ti (s) + 2MgCl2 (l)

The most reactive metals that can't be reduced by any other reducing agent are made by electrolytic reduction of their compounds in the molten state. Lithium, sodium, magnesium and aluminium are produced by this process. These metals are too reactive to be discharged by electrolysis of an aqueous solution.

Thermodynamics of Reduction Process:

As we know that metallurgy of most metals comprises reduction of their oxides. The nature of the reduction process based upon the ease by which the oxide can be reduced. Some of the oxides are so simply reduced that they decompose just via heating at relatively low temperatures. For illustration: Priestley, in his experiments on oxygen produced metallic mercury and oxygen from mercuric oxide via simply heating it by sunlight. Whenever sun light was focused on HgO by means of the magnifying glass, it decomposed spontaneously according to the equation:

2HgO (s) → 2Hg (g) + O2 (g)

The practicality of producing a free metal via thermal decomposition based on the extent to which the reaction carries on to completion at a particular temperature. As we know, the feasibility of the reaction is controlled by the free energy change taking place throughout the reaction. When ΔG° for a reaction is negative, the reaction is feasible from a practical stand point as significant amounts of products will be formed. We know that the standard free energy change, ΔG°, is associated to the standard enthalpy change, ΔH°, and the standard entropy change, ΔS°, according to the given equation:

ΔG' = ΔHo - TΔSo

In another words, the sign and magnitudes of ΔH° and ΔS° manage the sign and magnitude of ΔG°.

As in the decomposition of an oxide, oxygen is produced in the gaseous form and at times the metal might as well be produced in vapor form, the process takes place by a sizeable increase in entropy, therefore ΔS° will be positive. Enthalpy of decomposition, ΔH°d is simply the negative of enthalpy of formation of the oxide, ΔH°f as ΔH°f  is usually negative for the metal oxides, enthalpy of decomposition will be positive. As an outcome, the sign of ΔG° is found out by the difference between the two positive quantities ΔH° and TΔS°, 'T' the absolute temperature being for all time positive.

From the above, we can assume that if the enthalpy of formation of the metal oxide is small as in the case of HgO, Ag2O, CuO and Au2O3, then the enthalpy of decomposition will be small positive quantity and ΔG°, which is represented by the difference of ΔH° and TΔS°, will become negative at comparatively low temperatures. These oxides are stated to have relatively low thermal stabilities. On the other hand, if the oxide consists of a huge negative enthalpy of formation, then the enthalpy of decomposition of the oxide will be a large positive quantity. As an outcome, the value of ΔG° will become negative at an extremely high temperature here TΔS° becomes larger than ΔH°. Therefore, the metal oxide would be stable with respect to the thermal decomposition. In order to decompose such a metal oxide, it would have to be heated to a very high temperature at which cost becomes too expensive. Therefore, knowledge of how the standard free energy change, ΔG°, for the reduction reaction differs with temperature is extremely significant.

Ellingham Diagrams:

Ellingham studied the variation of standard free energy change for the formation of a number of compounds, example: oxides, sulphides and chlorides, with temperature and plotted ΔG° against temperature. These diagrams exhibiting the variation of ΔG° with T are termed as Ellingham diagrams after his name. As stated above, ΔG° is associated to ΔH°, ΔS° and T according to the given equation:


We as well are familiar that for most of the chemical reactions, ΔH° and ΔS° don't change significantly by temperature and can be regarded as constant. Therefore, ΔG° plotted against T provides a graph of constant slope that is equivalent to -ΔS°. However, due to abrupt changes in ΔS°, breaks in the graph take place at temperatures at which reactants or products melt or boil, that is, undergo phase change.

The figure below exhibits the Ellingham diagrams for the formation of metal oxides from free elements. By observing the Ellingham diagram for the formation of an oxide, we can determine the temperature at which the standard free energy change for the reaction will become positive. For illustration, consider the ΔG°/T graph (figure shown below) for the reaction of zinc with oxygen:

2Zn (s) + O2 (g) → 2ZnO (s)

2373_Ellingham diagram.jpg

Fig: Ellingham diagram showing the variation of free energy

At 273 K, the value of standard free energy change for this reaction is around - 600 kJ, that becomes less negative as temperature increases and finally at 2173 K it becomes zero. Above this temperature, ΔG° will become more positive, thus ZnO will spontaneously decompose to zinc and oxygen. This behavior is typical for all the elements apart from carbon; at adequately high temperatures the oxides become unstable relative to the constituent elements.

By the assistance of Ellingham diagrams, we can determine the standard free energy changes for a large number of reactions. For illustration, we can read off from the diagram the standard free energy changes for the given two reactions at 298 K:

2C (s) + O2 (g) → 2CO (g): ΔG° = - 275KJ

2ZnO (s) → 2Zn (s) + O2 (g):  ΔG° = + 640 KJ

We may note that the standard free energy change in the above equation is positive as it stands for the decomposition of zinc oxide. On adding the above two equations and respective ΔG° values we obtain.

2ZnO (s) + 2C (g) → 2Zn (s) + 2CO (g): ΔG° = + 365 KJ

As the standard free energy change for the above reaction is positive, the reaction has little tendency to take place at 298 K. As, two moles of gaseous product, that is, CO, are produced throughout the reaction, ΔS° is positive. Thus, ΔG° reduces with the increase in temperature and will become zero at certain temperature. In this specific case, AG' becomes zero at 1173 K as can be seen from the figure above. This temperature corresponds to the point of intersection of the two graphs for C/CO and Zn/ZnO systems. Over this temperature, ΔG° will become negative. Thus, carbon will reduce zinc oxide over 1173 K - a temperature 1000 K lower than the temperature of thermal decomposition of zinc oxide. Likewise, by the help of Ellingham diagrams for Zn/ZnO and H2/H2O systems, we can determine that H2 will reduce ZnO at a temperature, of 1400 K. As reduction of ZnO with carbon, which is as well much cheaper than hydrogen, can be taken at a lower temperature, it is clear that reduction by employing carbon is much cheaper than reduction employing hydrogen. 

Ellingham diagrams are very helpful for determining the temperature at which appreciable reaction takes place. The lower the ΔG°/T graph of an element is on the diagram, the more stable its oxide is relative to the dissociation into element and oxygen. These elements will reduce the oxides of other elements whose ΔG°/T graph appears above them on the diagram at a particular temperature. As we have notice above, carbon reduces ZnO above 1173 K; however below 1173 K zinc will reduce CO. As the ΔG°/T graph of C/CO system slopes downwards, it will finally be below all oilier graphs at adequately high temperatures. Thus, theoretically carbon will decrease all oxides. However difficulties in obtaining extremely high temperatures cheaply and the formation of carbides prevent the preparation of the more electropositive metals via this method. It is as well clear from the Ellingham diagrams that hydrogen can be employed as a reducing agent for the oxides of those elements whose ΔG°/T graphs are above that of hydrogen in the diagram. Therefore hydrogen can reduce the oxides of tungsten, lead, antimony, copper, nickel, zinc and cadmium.

Figure below exhibits the Ellingham diagram for sulphides of different elements. We can see from the diagram that carbon and hydrogen are not efficient reducing agents for metal sulphides. Thus, sulphides are first roasted in air to transform them to oxides, which are then reduced.

548_Ellingham diagram for sulphides.jpg

Fig: Ellingham diagram for sulphides

The Ellingham diagram for chlorides is illustrated in the figure shown below. This can be considered from the diagram that carbon is useless as a reductant for chlorides; however hydrogen can be employed for this purpose, particularly at higher temperatures.

2384_Ellingham diagram  for chlorides.jpg

Fig: Ellingham diagram for chlorides


The main application of hydrometallurgy is in the case of low grade ores that can't be concentrated economically. In this method, the powdered ore is first treated by an aqueous solution of an appropriate chemical whereby the metal is obtained in the form of its soluble salt leaving behind the gangue panicles. This method is known as leaching. Some of the illustrations of leaching are represented below:

Low grade oxide, carbonate and sulphide ores of copper are treated by dilute sulphuric acid in the presence of oxygen:

CuO + H2SO4 → CuSO4 + H2

CuCO3 + H2SO4 → CuSO4 + CO2 + H2

Cu2OS + 2H2SO4 + 2O2 → 2CuSO4 + SO2 + 2H2O

If the silver ore, AgCl, is treated by an aqueous solution of sodium cyanide, AgCl dissolves in it due to the formation of Na[Ag(CN)2]:

AgCl + 2NaCN → Na[Ag(CN)2] + NaCl

Sulphide ore, Ag2S, dissolves slowly only as the reaction is reversible:

Ag2S + 4NaCN ↔ 2Na[Ag(CN)2] + Na2

If air is passed via this solution, sodium sulphide is oxidized to sodium sulphate and the forward reaction goes to completion dissolving all the sulphide ore. In the presence of air, native silver is as well leached out in the form of Na[Ag(CN)2]:

4Ag + 8NaCN + 2H2O + O2 → 4Na[Ag(CN)2] + 4NaOH 

The leached out metals are recovered from the solution either through precipitation on treatment by a more electropositive metal or through electrolysis. For illustration, copper can be recovered from its solution via adding metals such as Fe, Al and so on. Silver is obtained from its solution via treatment with Zn or Al:

CuSCO4 + Fe → Cu + FeSO4

2Na[Ag(CN)2] + Zn → 2Ag + Na2[Zn(CN)4]

On the other hand, the dilute solution can be concentrated and then electrolyzed to get pure metals. From leached solution of copper ores, copper is frequently recovered via electrolysis of the solution. In electrolysis the anode utilized is of lead alloy and the cathode is of a pure copper sheet. Whenever direct current is passed via the solution, copper gets deposited on the cathode. Sulphuric acid is produced throughout electrolysis that is recycled in leaching of ore. Following reactions occur throughout electrolysis: 

Anode: 2H2O → O2 (g) + 4H+ (aq) + 4e

Cathode: Cu2+ (aq) + 2e → Cu(s)

2H+ (aq) + SO42- (aq) → H2SO4 (aq)


The two above metallurgical methods, namely pyrometallurgy and hydrometallurgy can be employed in the extraction of a fairly large number of metals. These methods, though, can't be utilized in cases: 

a) Where the metal is highly reactive, example: Na, Li and so on. There are not any chemical-reducing agents strong adequate to prepare these metals.

b) Where the oxide gets reduced only at extremely high temperatures at which the formation of carbides as well takes place, example: Al, Mg and so on.

In these cases, metals can be extracted via electrolysis of their salts in the molten state. Therefore, sodium and magnesium are made by electrolysis of fused chlorides, where the metals are discharged at the cathode and chlorine gas is evolved at the anode (figure shown below). Following reactions occur throughout electrolysis:

Anode: 2Cl- → Cl2 (g) + 2e

Cathode: 2Na+ + 2e → 2Na (l)  

Mg2+ + 2e → Mg (l)

554_Electrolysis of molten sodium chloride.jpg

Fig: Electrolysis of molten sodium chloride

In theory, the aluminium metal could be made the similar way. However, aluminium trichloride is covalent and it doesn't conduct electricity. As we will recall aluminium is obtained via electrolytic reduction of alumina in fused cryolite at 1100-1300 K by employing carbon anode and iron cathode. Electrolysis yields aluminium at cathode and 62 at anode that reacts by carbon to produce CO2.The reactions at electrodes are as shown:

Anode: 2O2- → O2 (g) + 4e

C(s) + O2 (g) → CO2 (g)

Cathode: A13+ + 3e → Al (s) 

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