Metallurgy Class 10 Note

What is Metallurgy?

The process of extracting metals from their ores and refining them for use is known as metallurgy. In other words, the process of obtaining a metal from its ores is called metallurgy of the metal.

Common terms used in discussing metallurgical operations :

  1. Charge : The mixture of materials fed to a furnace for obtaining the metal is called charge.
  2. Dressing of the ore : The removal of impurities associated with ore is called dressing or enrichment of the ore.
  3. Calcination : The process of heating a metal-rich ore to a high temperature to convert the metal into its oxide, either in absence or insufficient supply of air is called calcination.
  4. Roasting : The process of heating a finely ground ore to a high temperature in excess of air is called roasting. Roasting converts the metal present in the ore to its oxide.
  5. Flux : A flux is a substance that is mixed with the fumance charge (calcined or roasted ore and coke) to remove the infusible impurities present in the ore.
  6. Slag : Flux combines with the infusible impurities to convert them into a fusible substance called slag. Being light, slag floats over the molten metal and is removed from there.

Impurities present in metal oxides may be acidic or basic. For acidic impurities, such as SiO2 or P2O5, a basic flux (e.g., CaO) is added to the charge. If basic impurities such as MnO are present, silica is added to the charge.

Impurity Flux Slag
SiO2       + CaO  → CaSiO3
P2O5       + 3CaO  → Ca3(PO4)2
MnO       + SiO2 MnSiO3
  1. Smelting The process of obtaining the metal by reducing its oxide ore with coke is known as smelting.

Principles of metallurgy

The extraction of a metal from its ore depends upon the reactivity of the metals.

  1. Metals at the top of the activity series (K, Na, Ca, Mg, etc.) are highly reactive. They do not occur in the free state. They are extracted by the electrolysis of the molten ore.
  2. Metals in the middle of the activity series (Zn, Fe, Pb, etc.) are moderately reactive. These are obtained by roasting and calcination of their sulphide or carbonate ore.
  3. Metals at the bottom of the activity series (Au, Ag, Pt, Cu) being the least reactive are found in the free state. Copper and silver also occur as their sulphide or oxide ores. These are obtained by the process of roasting.

Metallurgical Operations

The various steps used in metallurgy are listed below.

  1. Enrichment or dressing of the ore
  2. Conversion of the enriched ore into the oxide of metal
  3. Extraction of metal from the metal oxide
  4. Refining or purification of the metal

Enrichment or dressing of an ore 

An ore mined from the earth’s crust contains a number of impurities (gangue), which must be removed. The ore, free from gangue, then becomes suitable for subsequent treatment. Enrichment or dressing of an ore is carried out by the following methods.

Levigation : The powdered ore is washed in a jet of water. The lighter, rocky and earthy impurities are washed away by water, while heavier ore particles are left behind to settle down at the bottom. This process is also called hydraulic washing.

Froth floatation : Sulphide ores of copper, lead and zinc are generally concentrated by this method.

The finely powdered ore is mixed with water and a small amount of oil in a tank. Air is blown into the mixture. A froth or scum is produced at the surface. The ore particles are carried by the froth to the surface. The earthy impurities sink to the bottom. The froth along with the ore is removed. An acid is added to break up the froth. The concentrated ore is filtered and dried.

Liquation : This process is used to concentrate the ore whose melting point is lower than that of the impurities. Stibnite, an ore of antimony, is concentrated by this method. The impure ore is heated. The ore melts and flows along the surface. The impurities are left behind.

Magnetic separation : This method is used when the magnetic properties of the ore and the impurities are different. For example, tinstone, an ore of tin, contains wolfram as an impurity that is magnetic. To remove this impurity, the ore is finely powdered to make the magnetic and the nonmagnetic particles distinctly separate. The powdered tinstone is spread on a belt moving over electromagnetic rollers in figure. The wolfram, being magnetic, is attracted and gets collected in the pot near the magnet. Tinstone falls away from the magnet.

Leaching or chemical separation : In this method, the powdered ore is treated with a suitable solvent. The ore dissolves in it while the impurities remain undissolved. For example, the bauxite ore contains Fe2O3, SiO2, etc., as impurities. The ore is powdered and treated with a solution of sodium hydroxide. A12O3 and SiO2 present in the ore dissolve, forming sodium aluminate and sodium silicate respectively. The impurities are left behind undissolved. The impurities are filtered off. The filtrate containing sodium aluminate and sodium silicate is stirred with some freshly prepared aluminium hydroxide for several hours. Sodium aluminate undergoes hydrolysis producing aluminium hydroxide as precipitate. The addition of aluminium hydroxide accelerates the precipitation of hydroxide. Soluble sodium silicate remains in solution. The precipitate, when filtered, washed, dried and ignited, gives pure alumina (Al2O3).

Al2O3 + 2NaOH → 2NaAlO2 + H2O

NaAlO2 + 2H2O → Al(OH) + NaOH

2Al(OH) → Al2O3 + 3H2O

Conversion of the enriched ore into the oxide of metal

It is easier to obtain metals from their oxides than from their carbonates or sulphides. Hence, the concentrated ore is converted into the oxide of metal which is then reduced to metal. This conversion to oxide is done by the process of calcination or roasting. In this process the ore is heated very strongly in the absence of air, keeping the temperature below its melting point so that volatile impurities are driven off.

EXAMPLES :

(i)   Oxide ores are calcined to remove moisture and other volatile impurities.

Al2O3 . 2H2O → Al2O3 + 2H2O

(ii)  Carbonate ores are calcined to expel carbon dioxide.

CaCO3 → CaO + CO2

CaCO3 × MgCO3 → CaO + MgO + 2CO2

\underset{{calamine\,\,ore}}{\mathop{{ZnC{{O}_{3}}}}}\,    → ZnO + CO2

CuCO3 × Cu(OH) → 2CuO + H2O + CO2

Sulphide ores are usually converted to oxides by roasting. The process involves heating the ore at a temperature below its fusion point, but always in the presence of air so that it may be oxidized.

Arsenic and similar other elements present in free state or combined state are also oxidized to volatile oxides.

(i)   Zinc blende (ZnS) is roasted in air to convert it into zinc oxide.

2ZnS + 3O2 → 2ZnO + 2SO2

(ii)  Galena (PbS) is converted into litharge (PbO) by roasting.

2PbS + 3O2 → 2PbO + 2SO2

(iii) Cinnabar (HgS) is roasted to convert it directly into mercury (Hg).

HgS + O2 → Hg + SO2

(iv) Iron pyrite (FeS2) is converted into ferric oxide (Fe203) by roasting.

4FeS2 + 11O2 → 2Fe2O3 + 8SO2

Thus, both calcination and roasting produce oxide of the metal. However, there are a few points of difference between the two processes.

Calcination

Roasting

1.    The ore is heated in       the absence of air. The ore is heated in the presence of air.
2.    It is used for oxide or   carbonate ores. It is used for sulphide ores.

Chloride ores remains unchanged by calcination or roasting.

Extraction of metal from metal oxide

A metal oxide thus produced is then reduced into metal. For this, the method used depends upon the reactivity of the metal being extracted. The following methods are used.

(i) Reduction by heat alone Metals occupying lower positions in the activity series can be obtained by heating their oxides.

2HgS + 3O2 → 2HgO + 2SO2

2HgO → 2Hg + O2

(ii) Chemical reduction Metals in the middle of the activity series (Fe, Zn, Ni, Sn, etc.) cannot be obtained by heating their compounds alone. They require to be heated with a reducing agent, usually carbon (coke). When a metal oxide is heated with carbon, it is reduced to free metal.

The reduction of metal oxides with carbon is known as smelting. The impurities are removed as slag

EXAMPLES :

(i)   When zinc oxide is heated with carbon, zinc metal is obtained.

ZnO + C → Zn + CO

(ii)  When stannic oxide is heated with carbon, tin metal is produced.

SnO2 + 2C → Sn + 2CO

(iii) Ferric oxide (Fe2O3) is reduced to iron by heating with coke in a blast furnace.

Fe2O3 + 3C →  2Fe  + 3CO

Reduction with aluminium (thermit process or alumino-thermic process)

Some metal oxides cannot be reduced satisfactorily by carbon. For them, aluminium, a more reactive metal, is used. The process is called thermic process or alumino-thermic process.

EXAMPLES :

(i)   Manganese dioxide is reduced to manganese by heating with aluminium.

3MnO2 + 4Al → 3Mn + 2Al2O3

(ii) Ferric oxide (FeP3) is reduced by aluminium to free iron.

Fe2O3 + 2Al → 2Fe + Al2O3

(iii) Chromium sesquioxide is reduced by aluminium to chromium metal.

Cr2O3 + 2Al → 2Cr + Al2O3

In the thermit process, aluminium powder is mixed with metal oxide. A piece of magnesium is set alight to start the reaction. The aluminium reduces the oxide to free metal.

In case of iron oxide, iron is obtained in the molten state. (The mixture of iron oxide and aluminium powder is called thermite). The molten iron may be allowed to trickle down to weld two iron objects together. Cracked machine parts, railway tracks, etc., are joined by this method.

(iii) Electrolytic reduction : The reactive metals (high up in the activity series) cannot be produced by any of the above methods. They are obtained by electrolytic reduction of their molten oxides or chlorides. During electrolysis, the cathode supplies electrons to metal ions for their reduction to the metal.

EXAMPLES :

(i)  Sodium metal is obtained by the electrolysis of molten sodium chloride.

(ii)  Magnesium metal is obtained by the electrolysis of molten magnesium.

(iii)Aluminium oxide (Al2O3) is reduced to aluminium by the electrolysis of molten aluminium oxide.

Al2O3 → 2Al3+ + 3O2–

The aluminium ions present in aluminium oxide go to the cathode and are reduced there to aluminium atoms.

Note : During electrolytic reduction of the molten salts, the metals are always liberated at the cathode.

(iv) Some specific methods Silver and gold are obtained by treating the ore with a solution of sodium cyanide. Sodium argentocyanide (in case of silver) or sodium aurocyanide (in case of gold) is obtained in the solution. On adding zinc dust to the solution, silver or gold is precipitated.

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

2Na[Au(CN)2] + Zn → Na[Zn(CN)4] + 2Au

Refining of Metals :

The metal obtained from the ore is not pure. It contains various substances as impurities. The process of removing these impurities is called refining of the metal. Some of the methods generally applied for refining metals are discussed below.

  1. Liquation : This process is used to separate metals of low melting points (e.g., tin and lead) from the metals of high melting points.

In this process, a sloping hearth is used. The hearth is kept at a temperature a little above the melting point of the metal. The impure metal is placed at the top of the hearth. The metal melts and flows down the hearth. The infusible impurities are left behind. This method is used in the purification of tin.

  1. Cupellation : This method is used to purify silver, containing lead as an impurity. The impure silver is heated in the presence of air in a vessel made of bone-ash. This vessel is called cupel. Lead is oxidized to lead monoxide. Most of the lead monoxide is carried away in the blast of air. The remaining portion of the lead monoxide melts and is absorbed by the bone-ash. Pure silver is left behind.
  2. Poling :Copper is purified by this method. The molten impure copper (called blister copper) is stirred thoroughly with poles of green wood. The gases escaping from the poles reduce the oxide of metal to the metal. The surface of the molten copper is kept covered with powdered charcoal so that copper may not be reoxidized in contact with air.
  3. Electrolytic refining : This method is widely used for purification of metals. Several metals such as aluminium, copper, tin, lead, gold, zinc and chromium are purified by this method. The impure metal is made the anode while a strip of pure metal acts as the cathode. A solution of the salt of the metal acts as the electrolyte.

On passing electric current through the solution, pure metal gets deposited on the cathode. The more reactive impurities present in the metal to be purified go into solution and remain there. The less reactive impurities fall to the bottom of the electrolytic cell.

Ultra-pure Metals

In the present age of technological advancement, metals of high purity are required for special purposes. For example, pure germanium is needed for semiconductor devices. Uranium of high-grade purity is used as fuel in nuclear reactors.

Two special techniques have been devised to prepare metals of very very high purity.

  1. Van Arkel method  : This method was developed by van Arkel to obtain ultra-pure metals. It is based on the thermal decomposition of metal compounds. It is used for obtaining pure titanium which is used in space technology. The impure titanium metal is converted into titanium tetra-iodide.

The air in the barrel used in this process is removed to create a high vacuum. An iodine bulb is broken. Titanium metal is heated which reacts with iodine to form gaseous titanium tetra-iodide.

Ti + 2I2 → TiI4

The impurities do not react with iodine. The vapour of titanium tetra-iodide is passed over a heated tungsten filament (1674 K). Titanium tetra-iodide gets decomposed into titanium and iodine. Pure titanium is deposited upon the filament and can be removed. The regenerated iodine can be reused to react with more titanium. The process is repeated.

  1. Zone refining method : This method is capable of producing metals of high purity. Germanium, which is used in semiconductor devices, is purified by this method. In this method, advantage is taken of the fact that impure molten metal, when allowed to cool, deposits crystals of pure metal.

An impure germanium rod is provided with a circular heater. The heater is slowly moved along the metal rod. A band of the rod melts. As the heater moves away, the metal crystallizes out of the melt. The impurities are swept along the molten zone. Finally, the impurities reach the other end of the rod, and are removed.

SOME COMMON METALS

IRON

Symbol Fe Atomic number 26

 

Electronic Configuration of Iron 

The atomic number of iron is 26. This means that an atom of iron contains 26 electrons in its shells. The electronic configuration of iron is shown below.

K L M N
Fe(26) 2 8 14 2

Thus, an atom of iron contains two electrons in its outermost shell.

Occurrence of Iron

Iron is second to aluminium in terms of abundance in the earth’s crust. It makes up 4.7% of the earth’s crust. Free iron has been found in most meteorites.

Iron is a reactive metal. So it does not occur free in nature. In combined state, it occurs as oxide, sulphide, carbonate, etc. The important ores of iron are:

(i) Haematite, Fe2O3

(ii) Magnetite, Fe3O4

(iii) Limonite, 2Fe2O3 . 3H2O

(iv) Siderite, FeCO3

(v) Iron pyrites, FeS2

The most important ore of iron is haematite, which is used most commonly in the extraction of iron. The pyrite ore (FeS2) is not used for the extraction of iron because of its high sulphur content.

Iron in India

Iron metal has great economic importance. The world output of iron exceeds two hundred million tonnes per annum. In 2002-03, India’s total production of iron reached almost 97 million tonnes. Besides, India has a vast deposit of iron ore: about 12,318 million tonnes of haematite and 5,396 million tonnes of magnetite. Most of these deposits are located in Jharkhand, Orissa, Chhattisgarh, Tamil Nadu, Karnataka and Maharashtra. The important iron and steel plants are located at Bhillai, Bokaro, Jamshedpur, Rourkela, Durgapur, Asansol and Bhadravati.

Extraction of iron from Haematite

Dressing of the ore :

The big lumps of the ore are broken into small pieces and then washed with water to remove clay, sand and other adhering impurities. The ore thus becomes ready for treatment in the blast furnace.

Smelting in the blast furnace

The concentrated ore is mixed with coke and limestone. The mixture is charged at the top of a blast furnace. The following reactions occur in the furnace.

(i)   As the charge comes down to the 873 K region, the iron oxide is reduced by the ascending carbon monoxide gas produced by the burning of coke.

2C + O2 → 2CO

Fe2O3 + 3CO → 2Fe + 3CO2

The iron thus obtained is called sponge iron.

(ii)  At the 1273 K region, the silica is converted to slag.

CaCO3 → CaO + CO2

CaO + SiO2\underset{{slag}}{\mathop{{\text{CaSi}{{\text{O}}_{\text{3}}}}}}\,

(iii) At the 1573 K region, sponge iron melts and dissolves carbon, phosphorus, silica, etc. The slag also fuses. The molten mass collects at the base of the furnace. The slag floats over it. The molten iron is taken out as required. This iron is called pig iron.

Function of limestone : Limestone is decomposed to give quicklime.

CaCO3 → CaO + CO2

Quicklime combines with impurities like sand to form a molten slag (calcium silicate).

CaO + SiO2 → CaSiO3

The slag floats on the surface of molten iron. It is taken out through a hole from time to time.

The formation of calcium silicate as slag not only removes unwanted silica but also keeps iron away from being oxidized.

Varieties of Iron

Pig iron and cast iron : The iron produced in the blast furnace is pig iron. It contains a comparatively high percentage of carbon due to which it is hard and brittle. It also contains phosphorus, silicon and manganese as impurities. Pig iron is melted, mixed with steel scrap and allowed to cool in moulds to give cast iron. Cast iron is impure iron, and is hard and brittle.

Wrought iron : It is almost a pure form of iron. It contains only 0.12% to 0.25% carbon. It melts at a higher temperature (1773 K) than that at which cast iron melts. Wrought iron is obtained by melting cast iron on a hearth lined with ferric oxide (Fe2O3). The impurities such as carbon, phosphorus, silicon and manganese are oxidized by Fe2O3.

Fe2O3 + 3C → 2Fe + 3CO

Wrought iron is soft, grey and tough. It is malleable and ductile. Hence, it can be drawn into sheets and stretched into wires. It is used in making chains, wire, anchors and cores of electromagnets.

Steel : It is an alloy of iron and carbon. It contains about 0.15 to 1.7% of carbon. There are different types of steel.

(a)  Mild steel : It contains less than 0.3% carbon. It is also called soft iron. Mild steel is used for making sheets and wires.

(b) Hard steel : It contains higher percentage (0.7-1.7%) of carbon. It is used in making tools and instruments.

(c) Alloy steels : Alloy steels are prepared by adding small amounts of nickel, cobalt, chromium, tungsten, molybdenum, manganese and silicon to steel. Alloy steels are used extensively in making rock-crushing machinery, helmets, armour plate, cutlery, springs, etc.

(d) Medium steel : It contains 0.3-D.7% carbon. It is hard and is used in making rails, bridges, etc.

Tempering

The hardness and elasticity of steel can be controlled by heat treatment. The steel is heated to a temperature below redness. It is then cooled slowly. The process is called tempering of steel. It is used to bring the steel to a suitable state of hardness and elasticity.

Annealing of steel 

Hard steel can be softened by heating it to a high temperature and then allowing it to cool down slowly. This process is called annealing.

Quenching of steel

Hard steel is heated to a high temperature. It is then suddenly cooled by plunging into oil or water. Steel becomes as hard and brittle as glass. Steel produced in this way is known as quenched steel and the process of making such steel is known as quenching or hardening of steel.

Properties of Iron

Physical properties 

Pure iron has a grey colour. It is malleable and ductile. It is a good conductor of heat and electricity. It melts at 1808 K and boils at 3023 K. It has a density of 7.9 × 103 kg cm–3.

Chemical properties

  1. Valency : Iron shows variable valency: 2 and 3. It forms divalent ion (Fe2+) as well as trivalent ion (Fe3+). The compounds in which iron shows divalency are known as ferrous compounds, whereas the compounds in which iron shows trivalency are known as ferric compounds. For example, in FeCl2 the valency of iron is 2. So, it is called ferrous chloride. In FeCl3, the valency of iron is 3. Hence, it is called ferric chloride.
  2. Action of air : In the presence of moist air and carbon dioxide, iron gets covered with a thin deposit of rust. The rust consists of hydrated ferric oxide (2Fe2O3×3H2O).
  3. Action of water : Red hot iron decomposes steam, forming ferrosoferric oxide and evolving hydrogen gas.

\displaystyle \underset{{iron}}{\mathop{{2Fe}}}\,\,+\underset{{steam}}{\mathop{{4{{H}_{2}}O}}}\,\,\xrightarrow[{}]{{}}\underset{{ferrosoferric\,oxide}}{\mathop{{F{{e}_{3}}{{O}_{4}}}}}\,\,+\underset{{hydrogen}}{\mathop{{4{{H}_{2}}}}}\,\,

  1. Action of acids : Iron lies above hydrogen in the activity series of metals. So, it can displace hydrogen from dilute hydrochloric acid or dilute sulphuric acid. The corresponding ferrous salts are also produced.

(i)   With hydrochloric acid

(a)  Iron dissolves in dilute hydrochloric acid, forming ferrous chloride with the evolution of hydrogen gas.

Fe + 2HCl → FeCl2 + H2

(b)  Concentrated hydrochloric acid also produces hydrogen with iron.

(ii) With sulphuric acid

(a) Iron dissolves in dilute sulphuric acid, forming ferrous sulphate. Hydrogen gas is evolved in the reaction.

Fe + H2SO4 → FeSO4 + H2

(b) Iron reacts with concentrated sulphuric acid to form ferrous sulphate with the evolution of sulphur dioxide. Hydrogen gas does not evolve in this reaction.

Fe + 2H2SO4 → FeSO4 + SO2 + 2H2O

Some ferric sulphate is also formed due to the oxidation of FeSO4 by concentrated H2SO4,

2FeSO4 + 2H2SO4 → Fe2(SO4) + 2H2O + SO2

(iii) With nitric acid :

(a)  Iron reacts with dilute nitric acid to form ferrous nitrate and ammonium nitrate.

4Fe + 10HNO3 → 4Fe(NO3)2 + NH4NO3 + 3H2O

(b)  With concentrated nitric acid, iron is rendered passive due to the formation of insoluble ferrosoferric oxide (Fe3O4) on the surface of iron.

  1. Action of halogens : Halogens combine with heated iron, forming the halides of iron. For example, chlorine combines with heated iron to form ferric chloride.

2Fe + 3Cl2 → 2FeCl3

  1. Action with sulphur : When iron filings are heated with sulphur, iron sulphide is produced.

Fe + S → FeS

  1. Displacement of less electropositive metal : When an iron piece is dipped into a solution of copper sulphate, copper is displaced from the salt and gets deposited on the surface of iron. This is because copper is less electropositive than iron.

Fe + CuSO4 → Cu + FeSO4

Tests to distinguish between ferrous and ferric salts

(i)   When a ferrous salt solution is treated with a solution of sodium hydroxide, a greenish precipitate of ferrous hydroxide is obtained.

FeSO4 + 2NaOH → Fe(OH) + Na2SO4

When a ferric salt solution is treated with a solution of sodium hydroxide, a brown precipitate of ferric hydroxide is produced.

Fe2(SO4)3 + 6NaOH→ 2Fe(OH) + 3Na2SO4

(ii)  The ferrous salts are generally green-coloured, whereas the ferric salts are brown in colour.

Rusting

When iron is exposed to moist air, a reddish-brown coating of a mixture of ferric oxide (Fe2O3) and ferric hydroxide ((Fe(OH3)) is deposited on the surface of the metal. This reddish-brown coating is known as rust, and this process is known as rusting. Thus, the slow conversion of iron into a mixture of Fe2O3 and Fe(OH) by water and atmospheric oxygen is known as rusting.

Rusting of iron is an oxidation reaction that occurs due to the attack of water and oxygen. It has been found that rusting does not take place in air-free water. It also does not occur in presence of oxygen alone. Both water and oxygen are essential for rusting. Thus, the following conditions are necessary for rusting:

(i)   Presence of oxygen or air

(ii)  Presence of water or moisture

The process of rusting is continuous. The strength of iron decreases gradually and finally the metal is destroyed completely.

Prevention of Rusting

Iron can be prevented from rusting by keeping it out of contact with air and water, and also by converting it into an alloy. This can be achieved in the following ways.

  1. By covering the surface of iron with grease, paint, varnish, enamel, etc.
  2. By galvanizing iron: a thin coating of zinc is deposited on the surface of the iron object. This is done by electroplating. Since zinc does not corrode on exposure to air, zinc metal prevents iron from rusting.
  3. By coating the surface of the iron object with chromium, tin, nickel, or aluminium. These metals resist corrosion. Hence, they protect iron from rusting.
  4. By converting it into an alloy with chromium and nickel. This alloy is called stainless steel.

Uses of Iron

  1. Iron is used in making household utensils and equipments.
  2. Wrought iron and cast iron are largely used in the manufacture of locomotives, railway lines, springs, tubes, etc.
  3. Iron finds wide application in house construction, e.g., in the reinforcement of roofs and other parts of buildings.

ALUMINIUM

Symbol Al Atomic number 13