7. Metal

7. METAL OCCURRENCE OF METALS Metals occur in nature, in the native (in free state) as well as in the combined state. Those occurring in the free state are very few e.g. gold, silver and platinum. Whether a given metal occurs in nature in a native or combined state, it is always contaminated with impurities–mostly clays and siliceous matter. Naturally occurring materials containing metals are called minerals. A mineral from which a given metal is obtained economically is called on ore. The natural occurrence of some selected metals is given in Table . Natural Occurrence of Selected Metals Element Nature of Ore Mineral Composition Iron (Fe) Iron pyrites FeS2 Copper (Cu) Copper pyrites CuFeS2 Silver (Ag) Argentite Ag2S Zinc (Zn) Zinc blende ZnS Mercury (Hg) Cinnabar Hgs Lead (Pb) Galena PbS Sulphides Oxides Iron (Fe) Haematite Fe2O3 Magmetite Fe2O4 Aluminium (Al) Bauxite Al2O3.2H2O Titanium (Ti) Rutile TiO2 Manganese (Mn) Pyrolusite MnO2 Copper (Cu) Cuprite Cu2O Tin (Sn) Cassiterite SnO2 Dolomite CaCO3.MgCo3 Magnesite MgCO3 Carbonates Magnesium (Mg) Calcium (Ca) Lime stone CaCO3 Zinc (Zn) Calamine ZnCO3 Sodium (Na) Rock salt NaCl Magnesium (Mg) Camallite KCl.MgCl2.6H2O Calcium (Ca) Gypsum CaSO4.2H2O Magnesium (Mg) Epsom salt MgSO4.7H2O Halides Sulphates A variety of minerals are found in India, some in huge quantities and some in small. Properties of Metals Physical Properties Mercury is the only liquid metal, all other are solids. They have high melting and boiling points. They are malleable and ductile. They are quite lustrous and are good conductors of heat and electricity. Chemical Properties Metals are electropositive and form positive ions by losing electrons. They combine with oxygen to form basic oxides. They react with dilute mineral acids to liberate hydrogen gas. With halogens they form halides. Their reactivity varies. The first group of elements, that is alkali metals are most reactive. Mineral Wealth of India State Mineral Andhra Pradesh Diamonds, Limestone, asbestos, copper, mica, gold Bihar Asbestos, bauxite, chromite, coal, graphite, haematite, limestone, manganese ore, mica Gujarat Gypsum, manganese ore, fluorspar Karnataka Asbestos, gold, chromite, iron ore, magnesite, corundum Kashmir and Gypsum Himachal Pradesh Kerala Manazite, illmenite, rutile, garnet Madhya Pradesh Iron and manganess ores, bauxite, limestone, coal, dolomite Maharashtra Chromite and manganese ores Orissa Dolomite, limestone, graphite, haematite, chromite Rajasthan Copper, lead, manganese and zine ores, rock phosphate, barytes, gypsum, mica Tamil Nadu Iron and manganese ores, magnesite, mica, limestone, lignite West Bengal Coal, dolomite tungsten Uttar Pradesh Extraction of Metals Gypsum, magnesite, dolomite, rock phosphate Extraction of metals comprises the processing of natural raw materials to obtain elemental metals. The initial stages of this process involve three steps: 1. Concentration of the ore (or beneficiation) 2. Chemical or electrochemical reduction to the metal (i) In many cases the ore is roasted (heating in air) to drive off volatile impurities, burn off organic matter, and to from compounds that are more easily smelted. Roasting in air usually converts sulphides to oxides. (ii) Concentration of bauxite (A12 O3. 2H2O) by hot aqueous sodium hydroxide. This method is based on the amphoteric nature of aluminium hydroxide. The crushed ore is treated with a hot aqueous solution of NaOH to dissolved alumina, A12O3 or bauxite and impurities such as iron oxide and silicates remain unaffected. The solution is filtered, cooled and CO2 gas is blown when aluminium hydroxide gets precipitated. On heating aluminium hydroxide, pure alumina is obtained. 3. Removal of impurities from the metal (i.e. purification or refining). In certain cases these steps may overlap, for example, removal of impurities may be a part of the ore concentration process such that subsequent refining is not necessary. Concentration of the Ore The ore usually contains rocky and siliceous matter called gangue. If the gangue is objectionable at a later stage, the first step in ore concentration is to remove it. To do so the ore is usually crushed and ground until the particles of the mineral are broken apart from the gangue. If possible, these particles are separated by physical means such as washing, magnetic attraction or froth floatation. These methods are based on differences between the physical properties of the mineral and the gangue. If the ore cannot be sufficiently concentrated by these physical methods, chemical processes are used. Some examples are given below: (iii) Leaching of gold and silver ores with aqueous sodium cyanide in the presence of air to yield the metal cyano complexes. (iv) Concentration of magnesium ions from sea water by calcium hydroxide. Mg (OH)2 is converted into anhydrous MgCl2 which is used for the electrolytic reduction for obtaining magnesium. Calcination and Roasting When the ore has been sufficiently concentrated for the isolation of a metal, it is subjected either to calcination or roasting process depending on the nature of the ore. In these operations, volatile impurities are also removed. Calcination is the heating of the ore in the absence of air. This method is employed for obtaining the metal oxides from carbonates and hydroxides. Roasting is the heating of the ore in the presence of air. On roasting, part of the ore is oxidized to form an oxide. This oxide is then reduced to the metal. Conversion of the ore into the oxide form is an important step in the final reduction to the metal. The reducing agent used should be easily available and cheap, the obvious choice is carbon. Smelting The industrial reduction process for obtaining metal from the treated ore is called smelting. There are several chemical and electrochemical methods available which can be used to produce metals. Ores generally contain considerable gangue even after concentration. To remove the last amounts of the gangue, a flux is added during the smelting process. Flux is a substance that combines with gangue and forms a molten material called slag as the mixture is heated in a furnace. At high temperatures, the slag is a liquid that is insoluble in the molten metal and it forms a separate layer. If the gangue is an acidic oxide, such as silica (SiO2), a cheap basic oxide like quicklime (CaO) may be used as flux. Impurities and flux react in a furnace to form slag. If on the other hand, the gangue is basic, the flux to be added would be acidic. Refining of Metals The purification or adjustment of the composition of impurities in crude metals is called refining. Metals with low boiling points, such as mercury, zinc and magnesium can be separated from most impurities by simple distillation. Like salts, metals can also be refined by fractional crystallization. Probably the most widely used refining method is the electrolytic method. Among modern purification methods, Zone refining, chromatographic, ion-exchanges and solvent extraction methods are also used for specific purpose. COPPER Symbol Cu Atomic No.29 Electronic configuration of copper The atomic number of copper is 29. Thus, an atom of copper contains 29 electrons. The electronic configuration of copper is : K L M Cu (29) 2 8 18 N 1 Copper atom has only one electron in its outermost shell. Position in the periodic table. Copper is a transition metal. It is placed in group IB along with silver and gold. VIII in the form of sulphide, oxide, carbonate, etc. The important ores of copper are: (i) Sulphide ores Copper glance (chalcocite), Cu 2S Copper pyrites (chalcopyrite), CuFeS2 IB IIB Cu Zn (26) (27) (28) (29) (30) Ru Rh Pd Ag Cd (44) (45) (46) (47) (48) Os Ir Au Hg The most important ore of copper is copper pyrites (CuFeS2). (79) (80) Copper in India Fe Co Ni Pt (76) (77) (78) (ii) Oxide ores Cuprite, Cu2O (iii) Carbonate ores Malachite, CuCO3.Cu(OH)2 Azurite, 2CuCO3.Cu(OH)2 Being a member of the same subgroup (IB), copper shows some marked similarities in properties with those of silver and gold. Further, the properties of subgroup (IB) metals are entirely different from those of subgroup (IA) metals. Hence, the subdivision into (IA) and (IB) is justified. In India, copper is found in Sikkim, Singhbhum in Bihar, Nellore in Andhra Pradesh and Khetri in Rajasthan. Copper is nowadays extracted through a modern process by The Indian Copper Corporation Ltd., Musabani, Ghatsila. Occurrence of copper Copper is obtained chiefly from the pyrite ore. This ore contains different materials. Hence, it needs to be concentrated or dressed. Copper is an unreactive metal. Hence, it occurs in\ nature in native or free state. It is also found in combined state Extraction of copper from copper pyrites 1. Dressing of the ore The dressing of the ore is done by froth floatation process. Properties of Copper 2. Roasting The dressed or the concentrated ore is roasted on the bed of a reverbveratory furnace in a current of air. As a result of this operation, a portion of sulphur burns to form SO2 and the metals are partially oxidized. The residue contains oxides and sulphides of copper and iron. Copper is a reddish, shining metal. It melts at 1083°C. Its density is 8.95 g/cm3. It is ductile and malleable. It is a good conductor of heat and electricity. 2CuFeS2 + O2 Cu2S+ 2FeS+SO2 3. Smelting The roasted ore is mixed with a little coke and sand. The mixture is heated strongly in a blast furnace. Most of the ferrous sulphide is converted into ferrous oxide. The ferrous oxide forms a slag with the sand and is removed. The residue is almost pure copper sulphide and is called the white metal. 2FeS + 3O2  2FeO + 2SO2 FeO + SiO2 ® FeSiO3 (ferrous silicate) slag 4. Preparation of blister copper The white metal is carefully heated in a reverberatory furnace where a portion of Cu 2S is oxidized to Cu 2O. As the temperature rises, Cu2O reacts with the remaining portion of Cu2S to form copper and SO2. Physical properties Chemical properties 1. Valency Copper exhibits variable valencies. The valencies of copper are 1 and 2. Thus, copper forms monovalent ions (Cu2+) and divalent ions (Cu) . The mono- valent ion (Cu) is Known as cuprous ion and the divalent ion (Cu2+ ) is known as cupric ion. Compouds corresponding to cuprous and cupric ions are called cuprous compounds and cupric compounds respectively .The cupric ion (Cu2+ ) is more stable than cuprous (Cu+). An atom of copper does not easily lose its electrons to form ions. Thus, the reactivity of copper is very low and it is treated as an uncreative metal. Copper lies very low in the activity series of metals. 2. Action of air Dry air has no action on copper. When the metal is exposed to moist air, its surface gradually gets covered with a green layer of basic copper carbonate CuCO3 Cu(OH)2. When copper is heated in air at about 300°C it combines with oxygen to form black cupric oxide. 300°C 2Cu2O + Cu2S  6Cu + SO2 Due to the evolution of S02, copper obtained has a blister-like appearance. This copper is called blister copper. 5. Refining of blister copper Blister copper contains iron, sulphur, arsenic, lead, etc., as impurities. It is melted in a reverberatory furnace in a current of air. The impurities present are oxidized away as gases or they form a slag with silica of the furnace. The slag is removed. Anthracite coal is sprinkled on the surface. The molten mass is stirred with poles of green wood. Pure copper is obtained. Very pure copper is, however, obtained by electrolysis. Sheets of impure copper are made the anode while thin plates of pure copper are made the cathode. A solution of copper sulphate is used as the electrolytic solution. On passing electric current, sheets of impure copper gradually dissolve. Pure copper is deposited on the cathode. The impurities settle down at the bottom or remain in the solution. Cu2++2e Cu (at cathode) Cu Cu2+ +2e (at anode) At the cathode, Cu2+ ions are taken from the copper sulphate solution while Cu2+ ions go into the solution at the anode. In this way, pure copper is transferred from the anode to the cathode. 2Cu + 02  2CuO At a temperature of about 1000°C, cuprous oxide is formed. 1000°C 4Cu + O2 2Cu2O 3. Action of water Copper does not react with water or steam. 4. Action of acids Copper lies below hydrogen in the electrochemical series. Hence, it cannot displace hydrogen from acids. Copper, however reacts with boiling concentrated HCl in presence of air or some oxidizing agent, with the evolution of hydrogen. 2Cu+ 2HCl  2CuCl+H2 Copper does not react with dilute hydrochloric dilute sulphuric acid. CuCl + CI-  [CuCl2]This is the only reaction in which copper displaces H2 from HCl. Copper reacts with dilute H2SO4 only in the presence of air: 2Cu + 2H2SO4 + O2  2CuSO4 + 2H2O When copper is heated with concentrated sulphruic acid, copper sulphate, sulphur dioxide and water are formed. Cu + 2H2SO4  CuSO4 + SO2 + 2H2O CU + H2O + 1/2O2  CU2+ + 2OH- With dilute nitric acid copper reacts to form nitrate and nitric oxide gas. CU2++4NH3  [CU(NH3)4]2+ 3Cu + 8HNO3 3Cu(NO3)2 + 2NO + 4H2O 5. Action of chlorine When copper is heated in a current of chlorine, cupric chloride is produced. Cu + Cl2  CuCl2 6. Action of sulphur When copper is heated with sulphur, cupric sulphide is formed. [Cu(NH3)4]2++ 2OH-  [Cu(NH3)4]-(OH)2 Test for copper (II) salts or cupric ions (Cu2+) When a copper (II) salt solution or a solution containing Cu2+ions is treated with ammonium hydroxide, a pale blue precipitate of copper (II) hydroxide is obtained , On adding excess of ammonium hydroxide, the precipitate dissolves to produce a deep blue solution. Uses of copper Cu + S  CuS 7. Displacement of metals Copper displaces some such as silver, gold, platinum, etc., from the solutions of their salts. Cu + 2AgNO3  Cu(NO3)2 + 2Ag It should be noted that only those metals are displaced from their salt solutions which lie below copper in the elecrochemical series. 1. Copper is largely used in electrical industry for making electric wires, electric motors, dynamos and several other electrical goods. 2. Copper is used in electroplating and electrotyping. 3. Copper is also used in making household utensils. 4. An important use of copper is in making alloys with metals. Some of the alloys and their compositions; 8. Action of aqueous ammonia Copper dissolves in aqueous ammonia in the presence of air, forming complex compound, tetra ammine cooper (II) hydroxide. given in Table. Table Alloys of copper Alloy Composition Uses Brass Cu= 80%, Zn=20% Utensils, cartridges, condenser tubes, etc (ii) Bronze Cu= 80%, Zn=10%, Sn= 10% Utensils, statues, coins, etc. (iii) German silver Cu= 60%, Zn=20%, Ni= 20% Utensils, resistance , coils, etc (iv) Bell metal Cu= 80%, Sn=20% Bells , etc. (v) Gun metal Cu= 87%, Sn=10% , Zn= 3% Bearings, machine parts Formerly use for cannons. (i) IRON Symbol: Fe Atomic no. 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: K L M N Fe (26) 2 8 14 2 Thus, an atom of iron contains two electrons in its outermost shell. 2C + 02  2CO Position of iron in the periodic table The electronic configuration of iron shows that there are two electrons in its outermost shell or the valence shell. So, it can be reasonably expected that iron should have a place in group II of the periodic table. But it is not so. Since iron is a transition metal, its properties are different from the normal elements of group II. In fact, iron has been assigned a place in group VIII of the periodic table. It is a member of the first transition series in the fourth period known as a triad. Group VIII Fourth period Fe Co Ni Fifth period Ru Rh Pd Sixth period Os Ir Pt Occurrence of iron Fe2O3 + 3CO  2Fe + 3CO2 The iron thus obtained is called spongy iron. (ii) At the 1000°C region, the silica is converted to slag. CaC03  CaO + CO2 CaO + SiO2 CaSiO3 slag (iii) At the 1300°C region, spongy 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. Iron is the second most abundant metal in the earth’s crust, the first one being aluminium. Free iron has been found in most meteorites. It is also believed that a core of iron is present in the centre of the earth. 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 on account of its high sulphur content. Iron in India In India, large deposits of iron ores are found in Bihar, Orissa, Madhya Pradesh, Tamilnadu and Sikkim. The important iron and steel plants are located at Bhillai, Bokaro, Jamshedpur, Rourkela, Durgapur, Asansole and Bhadravati. Extraction of iron from haematite 1. 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. 2. (i) Smelting in the blast furnace The concentrated ore is mixed with coke and limestone. The mixture is charged at the top of the blast furnace. The following reactions occur in the blast furnace: As the charge comes down to the 600°C region, the iron oxide is reduced by the ascending carbon monoxide gas produced by the burning of coke. Function of limestone Limestone is decomposed to give quicklime. CaCO3 ® CaO+ CO2 quicklime Quicklime combines with impurities like sand to form a molten slag (calcium silicate) CaO + SiO2 ® CaSiO3 slag The slag floats on the surface of molten iron. It is taken out through a hole from time to time. The of calcium silicate as slag, not only removes unwanted silica but also keeps iron away from being oxidized. Varieties of iron 1. Pig iron and Cast iron The iron produced in the blast furnace is pig iron. It contains a comparative 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 and is hard and brittle. 2. Wrought iron It is almost pure form of iron. It contains only 0.12 to 0.25% carbon. It melts at a higher temperature (1500°C) than that at which cast Iron melts. It is obtained by melting cast iron on a hearth lined with ferric oxide (Fe2--O3). The impurities, such as carbon, phosphorus, silicon and manganese, are oxidized by Fe2O3. 3. Steel It is an alloy of iron and carbon. It contains about 0.1 to 0.15% of carbon. There are different types of steel. (a) Mild steel It contains about 0.15% carbon. It is used for making sheets and wires. (b) Hard steel It contains higher percentage (1.5%) carbon. It is used in making tools and instruments. (c) Alloy steels Alloy steels are prepared by adding small amounts of nickel, cobalt, chromium, tungesten, 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 0.1% carbon. It is hard Fe2O3 + 3C 2Fe + 3COWrought 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, wires, anchors and cores of electromagnets. and is used in making rails, bridges, etc. Table: Typical percentage compositions of the three types of iron C Si S P Cast iron 2-4.5% 0.7-3.5% 0.7-3.5% 0.05-1.5% 0.5-1.0% Wrought 0.12-0.25% 0.03-2.0% 0.03-2.0% 0.04-0.04% 0.1-4% Mild steel 0.15% 0.03% 0.05% 0-0.05% 0-50% whereas the compounds in which iron shows trivalent 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. So, it is called ferric chloride. Tempering The hardness 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 state of hardness and elasticity. 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. Annealing of steel Hard steel can be softened by heating it to a high temperature and then allowing it to cool down slowly. This called annealing. Quenching of steel Hard steel is heated to a high temperature. It is then suddenly cooled by plunging it into oil or water. Steel becomes as hard and brittle as glass. This is known as quenched steel, and the process of obtaining such a steel is known as quenching or hardening of steel 3Fe + Iron 1. Valency Iron shows variable valency. It shows valancies of 2 and 3, i.e., it forms divalent ion (Fe2+) as well trivalent ion (Fe3+). The compounds in which iron shows are known as ferrous compounds, steam Fe3O4 ferrosoferric + 4H2 hydrogen 4. Action of acids Iron lies above hydrogen in the chemical series. So, it can displace hydrogen from the 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. Physical properties Chemical Properties 4H2O  oxide Properties of Iron Pure iron has a grey color. It is malleable and ductile, conductor of heat and electricity. It melts at 1525°C and has a density of 7.86 g/cm3. Mn Fe + 2HC1  FeCl2 + H2 (b) Concentrated hydrochloric acid also evolves 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 (b)  FeSO4 + H2 Iron reacts with concentrated sulphuric acid to form ferrous sulphate with the evolution of sulphur oxide. There is no formation of hydrogen gas. Fe + 2H2SO4  FeSO4 +SO-2 +2H2O Some ferric sulphate is also formed due to the oxidation of FeSO4 by concentrated H2SO4. 2FeSO4 + 2H2SO4  Fe2(SO4)3 + 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. 5. 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 + 3C12 6.  2FeCl3 Action with sulphur When iron filings are heated with sulphur, iron sulphide is produced. Rusting When iron is exposed to moist air, a reddish-brown coating of a mixture of ferric oxide (Fe2O3) and ferric hydroxide (Fe(OH)3) 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)3, by water and atmospheric oxygen, is known as rusting. Rusting of iron is an oxidation reaction which 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 due to 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. Fe + S  FeS 7. Displacement of less electropositive metal When an iron piece is dipped in 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 + FeS04 or Fe + Cu2+  Cu+ Fe2+ Uses of iron Tests to distinguish between ferrous and ferric salts 1. Iron is used in making household utensils and equipment. 2. Wrought Iron and cast iron are largely used in the manufacture of locomotives, railway lines, springs, tubes, etc. 1. Iron finds wide application in house construction, e.g., in the reinforcement of roofs and other parts of buildings. 1. 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)2 + Na2SO4 When a ferric salt solution is treated with a solution of sodium hydroxide, a brown precipitate of ferric hydroxide is produced. 2. ALUMINIUM Fe2(SO4)3 + 6NaOH  2Fe(OH)3 + 3Na2SO4 Symbol Al The ferrous salts are generally green-coloured, whereas the ferric salts are brown in colour. History Aluminium was first prepared in 1827. Impressed by its properties, scientists all over the world attempted to develop a process for the commercial production of Atomic no. 13 the metal. But it remained a costly metal till 1886, when Heroult in France and Hall in USA succeeded in developing independently a process for the extraction of aluminium. Since then, large-scale production of aluminium began in several parts of the world. It came into large-scale use towards the end of the nineteenth century. Electronic configuration of aluminium The atomic number of aluminium is 13. This means that an atom of aluminium contains 13 electrons in its shells. The electronic configuration is then, K L A1 (13) M 2 8 3 Thus, the outermost shell or valence shell of an aluminium atom has 3 electrons. Position in the periodic table First period IA IIA IIIA H (1) Second period Li (3) Be (4) B(5) Third period Na (11) Mg (12) Al (13) Fourth period K(19) Ca (20) Ga (31) Occurrence of aluminium Aluminium comes third amongst the most abundant elements in nature. But aluminium is the most abundant metal in the earth’s crust. It is a reactive metal, hence does not occur native or free in nature. In the state, it is widely distributed in nature in the form of compounds. Some of its important minerals are: i. Cryolite, Na3AlF-6 hydrolyzed to produce a white precipitate of aluminium hydroxide. NaAlO-2 + 2H2O  Al(OH)3 + NaOH Sodium Aluminate water aluminium sodium hydroxide hydroxide The precipitate of Al(OH)3 is filtered off, washed, dried and ignited to get pure aluminium oxide. ii. Corundum, A12O3 iii. Bauxite, A12O3.2H2O iv. Feldspar or Felspar, KAlSi3O8 2Al(OH)3  Al2O3 + 3H2O Aluminium aluminium water However, bauxite is the only mineral from which aluminium is obtained profitably. Hence, bauxite is an ore of aluminum. In India, Hindustan Aluminium Corporation, Indian Aluminium Company and the Bharat Aluminium Corporation produce aluminium from bauxite. Hydroxide 1. oxide Electrolytic reduction of alumina Alumina is mixed with cryolite. The mixture is melted in an iron cell. Carbon lining Extraction of aluminium from bauxite The extraction of aluminium from bauxite consists of the following steps: 1. Preparation of alumina Crude bauxite is finely powdered and treated with a concentrated caustic soda at 150°C. The bauxite goes into forming sodium aluminate while the behind undissolved. Al2O3.2H2O + 2NaOH  2NaAlO2+ 3H2O bauxite sodium hydroxide sodium water aluminate The solution is filtered. The filtrate is diluted with water and a little freshly precipitated aluminium hydroxide is added to cause precipitation. Sodium aluminate gets Powerderd coke Molten aluminum Alumina melts at 200°C. It is a bad conductor of electricity. But when mixed with cryolite and some calcium fluoride, the mixture becomes a good conductor of electricity, and melts at 900-950°C. The iron cell is lined inside with gas carbon which serves as cathode. Carbon rods act as anode. The electrolyte, thus, contains Na+, Al3+, F- and O2- ions. On passing electric current, Al3+ ions are discharged at the cathode and the O2- ions at the anode. Al3+ + 3e  Al (at cathode) Chemical properties 1. Valency The outermost shell or the valence shell of aluminium contains 3 electrons. These three electrons can be given up to form aluminium ion, Al3+. Thus, the valency of aluminium is 3, i.e., aluminium is a trivalent metal. 2. Action of air Aluminium is a very reactive metal. Yet it does not get corroded when exposed to air. Dry air has no action on aluminium. In moist air, a thin film of its oxide is formed over its surface. This prevents the metal from further attack by air or water or even by dilute acids. The film of the oxide can be removed by rubbing the surface of the metal with mercury. Then rapid corrosion by moist air occurs. The oxide film may also be removed by sodium chloride. That is why aluminium rusts in coastal areas. 2O2-  O2 + 4e (at anode) Some quantity of oxygen formed escapes and some reacts with the anode to form CO2. c + o2  co2 Hence, the anode has to be replaced from time to time. This increases the cost of production of aluminum. Molten aluminium sinks to the bottom, and is taken out from there. 3. Electrolytic refining of aluminium Aluminium thus obtained is 99% pure. It contains iron and silicon as impurities. It is purified by further electrolysis using Hope’s method. The electrolysis is carried out in an iron box lined with carbon. Three layers, differing in densities, are placed in the box. When heated to 800°C in air, aluminium burns with a brilliant light. In this reaction, aluminium oxide and a little aluminium nitride are formed. Bottom layer It consists of impure aluminium. It acts anode. Middle layer It consists of a mixture of the fluorides of Al, Na and Ba. It serves as electrolyte. Upper layer It consists of pure molten aluminium and serves as cathode. 4Al + 3O2 aluminium  2 Al2O3 + Heat oxygen aluminium oxide 2Al + N2 ® Aluminium nitrogen 3. Pure aluminium 2A1 Carbon lining aluminium nitride Action of water Aluminium is a reactive metal. But it does not react with water. This is because the surface of aluminium remains covered with a thin film of unreactive oxide. The thin oxide film does not allow water to come in contact with aluminium. If the thin oxide film is removed by rubbing with a sand paper then aluminium reacts with boiling water to produce aluminium hydroxide and hydrogen. Fluorides of Al, Na, Ba Impure Aluminium 2AlN +  6H2O aluminium water 2Al(OH)3 + 3H2 aluminium hydroxide hydrogen 4. Action of acids Aluminium reacts with acids to produce aluminium salts and hydrogen. (i) Reaction with hydrochloric acid Aluminium dissolves in hydrochloric acid to form aluminium chloride and hydrogen. On passing electric current, aluminium ions from the middle layer are discharged at the cathode. An equivalent amount of aluminium goes into the middle layer from the bottom layer. The impurities are left behind. Oxygen is liberated at the anode. Oxygen reacts with the carbon anode. Hence, the anode is gradually consumed. aluminium hydrochloric aluminium chloride hydrogen Properties of Aluminium (ii) Physical properties Aluminium possesses a bluish white lustre. It is a light metal of density 2.7. It is malleable and ductile. It melts at 660°C. It is a good conductor of heat and electricity. 2Al + 6HCl - 2AlCl3 + 3H2 The reaction is slow with dilute hydrochloric acid but rapid with concentrated hydrochloric acid. Reaction with sulphuric acid Aluminium dissolves in dilute sulphuric acid to form aluminium sulphate and hydrogen. The reaction is, however, slow. 2Al + 3H2SO4 Al2(SO4)3 + 3H2 Aluminium dissolves in hot concentrated sulphuric acid to form aluminium sulphate with the evolution of sulphur dioxide gas. 2Al + 6H2SO4  Al2(SO4)3 + 6H2O + aluminium sulphuric acid water sulphur dioxide 3SO2 Aluminium is a good conductor of electricity; it is extensively used in making electrical wires. 3. Aluminium foils are used for wrapping foodstuffs. 4. Aluminium is a light metal. Hence, it is used in making bodies and parts of aircraft, buses furniture, etc. Since pure aluminium is not very strong, alloys of aluminium are used for these purposes. 5. Aluminium powder is used in anti-corrosion paints and in explosives. 6. Aluminium powder is also used in alumino-thermic process for welding two iron objects. aluminium sulphate (iii) Reaction with nitric acid In dilute or concentrated nitric acid, aluminium is rendered passive due to the formation of an insoluble layer of oxide. The oxide layer does not permit further reaction. (iv) Action of alkalis Aluminium dissolves in caustic soda or caustic potash solution with the evolution of hydrogen, forming soluble sodium or potassium aluminate respectively. 2Al + 2NaOH + 2H2O  2NaAlO2 + 3H2 2Al + 2KOH + 2H2O  2KaAlO2 + 3H2 On account of these reactions, aluminum containers cannot be used to store alkali solutions. (v). 2. Action of chlorine When chlorine gas is passed over heated finely powdered aluminum, aluminium chloride is produced. 2Al + 3Cl2 2AlCl3 USES OF METALS From time immemorial, metals have been in the service of man for various purposes. We know has been using metals for the manufacture of tools and weapons. With the passage of time, the use became widespread. Metals are used in three different forms: 1. As pure metals 2. As alloys 3. As compounds of metals Uses of pure metals 1. Metals are good conductors of electricity. copper is used to make good quality electric wires. Aluminium is also used to make electric wires which are cheaper than copper wires. 2. Iron, aluminium and copper are used for making cooking utensils and factory equipment. It reduces many metal oxides to metals. For example, it reduces ferric oxide (Fe2O3) to iron metal. 3. Mercury is used in thermometers and in the action of some medicines. Fe2O3 + 2Al  Al2O3 + 2Fe 4. Silver and gold are used in jewellery. A large amount of heat is liberated in this reaction. As a result, temperature rises to about 35000C. 5. Thin foils of aluminium are used for packaging 6. Iron, in the form of steel, wrought iron and cast iron is used for the manufacture of locomotives, railway lines, implements of war, springs, tubes, cranes, etc. 7. Gold, silver, copper, nickel, etc., are being used for coinage all over the world. 8. Zinc and tin are used for coating iron objects to prevent rusting. 9. Cadmium, titanium and zirconium are used in atomic energy and space research projects. 10. Nowadays, titanium is used in aerospace and aircraft industries for making aircraft frames and engines where its high strength to weight ratio is useful, in the manufacture of military hardware, and in chemical reactors and industries. It finds application in making marine instruments. It is also used for hardening of steel. (vi) Reducing action Aluminum is a powerful reducing agent. It reduces CO and CO2 to Carbon . 3CO+ 2Al  Al2 + 3C 3CO2+4Al 2Al2O3 +3C Tests for aluminium salts 1. Charcoal test A solid aluminium salt is heated the cavity of a charcoal block. A white residue is obtained. When this residue is moistened with cobalt nitrate solution, a blue mass is obtained. 2. When the aqueous solution of an aluminium salt is treated with ammonium hydroxide, a white gelatinous precipitate of aluminium hydroxide is produced. AlCl3 + 3NH4OH  Al(OH)3 + 3NH4Cl Aluminium chloride aluminium hydroxide ammonium hydroxide ammonium chloride Uses of aluminium 1. Aluminium is a good conductor of heat and it does not get corroded. Hence, it is widely used to make utensils like pans, kettles, etc. The large-scale use of titanium is due to (i) its high strength to weight ratio, (ii) its resistance to corrosion and (iii) its high melting point. Due to its specific (especially military and aerospace) application called a strategic element. Uses of alloys An alloy is a mixture of a metal with other metals or nonmetal. It has properties different from those of its constituents. The properties of a metal are much more improved and pronounced when the metal is transformed into an alloy. For example, aluminium is a light metal and it is not very strong. But an alloy of aluminium (duralumin) is light and very strong. Similarly, silver is soft, but a silver alloy is hard and more useful than pure silver. In order to prepare an alloy we mix up various metals molten state. The molten mass is then cooled to room temperature. The solid substance formed is an alloy. Some of the commonly used alloys, their composition various uses, are listed in Table. Alloy Composition Us es 1. Brass Copper = 80% zinc= 20% Brass is harder than pure copper. It is used for making utensils, cartridges, etc. 2. Bronze Copper= 80%Zinc = 10%tin= 10% It is used for making statues, medals, ships, coins, and machines, 3. Solder tin= 50%lead = 50% It is used for joining metals, soldiering wires, electronic components. Etc. 4. Duralumin Aluminium= 94.4%silicon = 4%and It is used for making bodies of Copper, Magnesium, manganese . aircraft and kitchenware Copper = 60%Zinc=20%Nickel= 20% It is used for making utensils, ornaments, etc. 5. German Silver 6. Monel metal Nickel=70%Copper=30%Iron=2% 7. Nichrome Nickel =67.5%Chromium=15%Iron =16%It is used for making coils of electric Manganese=1.5% furnaces. 8. Gun metal CopperTinZinc It is used for making bearing, valves, etc. 9. Rose metal SiliconLeadTin It is used for making fuse wires. 21.5 SOME USEFUL COMPOUNDS OF METALS Common Salt. Common salt is sodium chloride (NaCl). Common salt when kept open absorbs moisture from air. This is because of the presence of magnesium chloride in common salt. Magnesium chloride is deliquescent, that is, it absorbs water. Sodium chloride is very important chemical. It is an essential ingredient in our daily meal. It can be converted into useful chemicals like caustic soda, washing soda and baking soda. It also yields hydrochloric acid, an important chemical used in chemical industry, and chlorine which is used as a bleaching agent. Caustic Soda. Chemically caustic soda is sodium hydroxide (NaOH). It is obtained by electrolysis of brine (NaCI solution). Sodium hydroxide is a deliquescent solid. On dissolving in water it gives a strong alkaline solution. Washing Soda. Washing soda is sodium carbonate decahydrate Na2CO3 .10H2O. Sodium carbonate (soda ash) is alkaline in nature and has detergent (cleaning) properties. In a crystalline state it possesses water of crystallization. On exposure to air, crystals lose water and turn into white opaque powder. Such an action, that is the giving up of water of crystallization to the atmosphere is termed as efflorescence. Sodium carbonate is used as a cleansing agent for domestic purposes. It is also used for softening of hard water, in manufacturing of chemicals like borax, caustic soda, glass and water glass. It is a constituent of many dry soap powders. Baking Soda. The chemical name of baking soda is sodium hydrogen carbonate (NaHCO3). Sodium bicarbonate is used in the manufacture of baking powders and health salts. In addition to sodium bicarbonate, baking powders contain tartaric acid (or a similar Compound.) When heated sodium bicarbonate reacts with acid producing carbon dioxide which causes cake to rise and become light. Lime. Calcium oxide is commonly known as quicklime or just lime. It is prepared by heating lime stone (CaCO3) in a kiln. CaCO3  CaO + CO2 Calcium oxide reacts vigorously with water producing slaked lime [Ca(OH)2] which is used for making mortar and for white washing. A solution of slaked lime is known as lime water which is used for testing carbon dioxide. Calcium oxide is used for making cement and glass. Bleaching Powder. The chemical name of bleaching powder is calcium oxychloride (CaOCl2). Bleaching powder is prepared by passing chlorine over solid slaked lime. Ca(OH)2 + Cl2  CaOCl2 + H 2O Bleaching powder When exposed to air, bleaching powder deteriorates giving off chlorine. Bleaching powder is commonly used for bleaching in laundries. It is used in paper and textile industries, for disinfecting water and for preparation of chloroform. Plaster of Paris. It is a calcium sulphate hemihydrate (CaSO-4)2 H2O. It is prepared by heating gypsum (CaSO4 2H2O). Plaster of paris has a remarkable property of hardening when mixed with proper quantity of water. The hardening of plaster of paris is due to its hydration to form gypsum. It is used in the laboratories for sealing the gaps, where sunlight arrangement is required. It is also used for making casts for statues, and in surgery to maintain joints in a fixed position