Comprehensive Notes on Metals and Non-Metals

5. Metals and Non-Metals

5.1 Introduction

• Natural resources are classified as renewable and non-renewable.
• Renewable resources can be replenished by natural processes at the same rate they are used; examples include plants, animals, soil, and water.
• Non-renewable resources exist in a fixed amount and cannot be easily replenished, such as natural gas, petroleum, and coal.
• Metallic minerals contain one or more metallic elements.
• Non-metallic minerals do not contain metals.
• Metallic minerals occur in rare, naturally formed concentrations known as mineral deposits.
• The primary source of metallic and nonmetallic minerals is the Earth.
• Most metals and non-metals are extracted from their minerals.
• Examples of metallic minerals include iron haematite (Fe2O3), aluminum bauxite (Al2O3.2H2O), zinc blende (ZnS), and malachite (Cu(OH)2.CuCO_3).
• Examples of nonmetallic minerals include diamond, gravel, and mica.
• Metals and non-metals can also be found in seawater and the atmosphere (e.g., oxygen and nitrogen).

5.2 General Properties of Metals and Production of Some Metals

5.2.1 Properties and Extraction of Metals

I. Some physical properties of metals:

• Lustrous appearance and can be polished.
• Malleable (can be beaten into thin sheets), used in aluminum foils and gold/silver ornaments.
• Ductile (can be drawn into wires), examples include copper, gold, iron, and silver.
• Hard and have tensile strength, except for lithium, potassium, and sodium.
• Generally have a high density, except for lithium, potassium, and sodium.
• Usually sonorous, producing a metallic sound when struck (e.g., school bell).

II. Chemical properties of metals:

• Metals possess positive valency and lose electrons; e.g., Mg(g) \rightarrow Mn^{2+}(g) + 2e^-.
• They have 1, 2, or 3 valence electrons.
• They are oxidized by losing electrons and act as reducing agents.
• They mostly form basic oxides and some amphoteric oxides.
• They form chlorides that are true salts and are electrovalent.
• They form hydrides that are ionic, unstable, and reactive.
• They usually replace hydrogen from dilute non-oxidizing acids like HCl and H2SO4, except for copper, silver, and gold.

III. Natural occurrence and extraction of metals:

• Noble metals such as Ag, Au, Bi, Cu, Pd, Pt exist in nature in an uncombined or free state.
• More active metals like alkali and alkaline earth metals never exist in an uncombined state and always exist in compounds.
• Active metals may exist as carbonates, halides, oxides, phosphates, silicates, sulphides, and sulphates.
• Constituents of the Earth’s crust containing these metals or their compounds are known as minerals.
• Examples: sodium as halite (NaCl), potassium as sylvite (KCl), magnesium as magnesite (MgCO3), calcium as limestone (CaCO3).
• Ores are minerals from which a metal can be profitably extracted.
• Ores contain impurities like sand and other undesirable materials called gangue.
• Metallurgy is the science and technology of extracting metals from their ores and compounding alloys.
• Three principal steps in the extraction of a metal from its ore:
  • Preparation (concentration) of the ore (e.g., oil floatation, magnetic separation)
  • Pre-treatment and production of the metal (e.g., roasting, calcination)
  • Purification of the metal (e.g., chemical reduction, electrolytic reduction)
• The most active metals (K, Na, Ca, Mg) are extracted from their compounds by electrolysis because there is no economic reducing agent.

5.2.2 Production of Aluminum, Iron, and Copper

A. Aluminum (Al)

• Aluminum is the most abundant metal and the third most plentiful element in the Earth’s crust (about 7%).
• It is the second-most important metal after iron in terms of consumption.
• Occurrence: Aluminum does not occur as a free metal in nature.
• Its principal ore is bauxite (Al2O3.2H2O). Other minerals containing aluminum are orthoclase (KAlSi3O8), cryolite (Na3AlF6), corundum (Al2O3), beryl (Be3Al2Si6O8) and china clay (Al2Si2O7.2H_2O).
• Extraction: Aluminum is extracted industrially from bauxite by the Hall–Héroult process.
• Bauxite is first purified by heating with sodium hydroxide solution to convert silica into a soluble silicate: SiO2(s) + 2NaOH(aq) \rightarrow Na2SiO3(aq) + H2O(l).
• Aluminum oxide is converted to soluble sodium aluminate: Al2O3(s) + 2NaOH(aq) \rightarrow NaAlO2(aq) + H2O(l).
• Impurities like iron oxides and titanium (IV) oxide are filtered off.
• The solution is treated with acid to precipitate aluminum hydroxide: AlO2^-(aq) + H3O^+(aq) \rightarrow Al(OH)_3(s).
• Aluminum hydroxide is collected, washed, dried, and heated strongly to obtain Al2O3: 2Al(OH)3 \rightarrow Al2O3 + 3H2O.
• Pure aluminum oxide is mixed with cryolite (Na3AlF6) to reduce its melting point from 2045°C to 1000°C.
• The molten mixture is then electrolyzed to obtain aluminum.
• The electrolytic cell contains graphite electrodes as both anode and cathode.
• Anode reaction: 3C(s) + 6O^{2-} \rightarrow 3CO_2(g) + 12e^-.
• Cathode reaction: 4Al^{3+}(l) + 12e^- \rightarrow 4Al(l).
• Overall reaction: 4Al^{3+}(l) + 6O^{2-}(l) + 3C(s) \rightarrow 4Al(l) + 3CO2(g) or 2Al2O3(l) + 3C(s) \rightarrow 4Al(l) + 3CO2(g).
• Physical Properties:
  • Soft silvery-white metal with a density of 2.7 g/cm3.
  • Melts at 660°C.
  • Can be shaped into wires, rolled, pressed, or cast into different shapes.
  • Good conductor of heat and electricity.
• Chemical Properties:
  • Aluminum is a reactive metal.
  • Reaction with Oxygen: Aluminum reacts with atmospheric oxygen to form a thin film of aluminum oxide on its surface: 4Al(s) + 3O2(g) \rightarrow 2Al2O_3(s).
  • This oxide film inhibits further reaction with oxygen.
  • Reaction with dilute acids: Aluminum reacts with dilute acids like HCl and H2SO4, forming salts and liberating hydrogen gas:
  • 2Al(s) + 3H2SO4(aq) \rightarrow Al2(SO4)3(aq) + 3H2(g)
  • 2Al(s) + 6HCl(aq) \rightarrow 2AlCl3(s) + 3H2(g)
  • Aluminum does not react with dilute or concentrated HNO_3 due to the formation of a protective oxide layer.
  • Aluminum burns in chlorine gas to form aluminum chloride: 2Al(s) + 3Cl2(g) \rightarrow 2AlCl3(s).
  • Aluminum reacts with sodium hydroxide solution: 2Al(s) + 2NaOH(aq) + 6H2O(l) \rightarrow 2NaAl(OH)4(aq) + 3H_2(g).
• Uses:
  • Lightweight and corrosion-resistant, used to make light alloys like duralumin (mixture of Al, Cu, and Mg).
  • Extensively used in the transport industry to make aircraft, ships, and cars.
  • Used in the manufacture of household cooking utensils due to its high thermal conductivity and resistance to corrosion.
  • Also used to make door and window frames, roofing for buildings, packaging material in food industries, and for electrical transmission lines.
  • In the thermite welding process, powdered aluminum mixed with iron (III) oxide produces a temperature of about 3000°C: 2Al(s) + Fe2O3(s) \rightarrow 2Fe(l) + Al2O3(s).
  • The mixture of powdered aluminum and iron oxide is called thermite.

B. Iron (Fe)

• Iron is the second-most abundant metal next to aluminum in the Earth’s crust and the fourth most abundant element (about 4.7%).
• Occurrence: Iron is never found as a free metal in nature.
• It occurs in nature only in the form of compounds such as oxides, carbonates, and sulphides.
• The chief ores of iron are hematite (Fe2O3), limonite (Fe2O3.H2O), magnetite (Fe3O4), and siderite (FeCO3). It is also found as iron pyrites (FeS_2), commonly called fool’s gold.
• Extraction: Iron is generally extracted from hematite, magnetite, and siderite in a blast furnace.
• Raw materials for the extraction of iron are iron ore, coke, limestone, and hot air.
• Reactions in the blast furnace:
  1. Oxidation of coke to carbon dioxide: C(s) + O2(g) \rightarrow CO2(g) + heat
  2. Reduction of carbon dioxide to carbon monoxide: CO_2(g) + C(s) \rightarrow 2CO(g)
  3. Reduction of iron oxides to metallic iron by carbon monoxide:
     • 3Fe2O3(s) + CO(g) \rightarrow 2Fe3O4(s) + CO_2(g)
     • Fe3O4(s) + CO(g) \rightarrow 3FeO(s) + CO_2(g)
     • FeO(s) + CO(g) \rightarrow Fe(l) + CO_2(g)
  4. Decomposition of limestone to remove impurities: CaCO3(s) \rightarrow CaO(s) + CO2(g)
  5. Reaction of calcium oxide with silica: CaO + SiO2 \rightarrow CaSiO3
• The iron obtained directly from the blast furnace is called pig iron, which is impure and contains about 2% silicon, up to 1% phosphorus and manganese, and traces of sulphur.
• When pig iron is re-melted with scrap iron and cast into moulds, it forms cast iron.
• Wrought iron is the purest form of commercial iron, obtained by removing most of the impurities from pig iron by heating impure iron with hematite and limestone, increasing the purity to 99.5%.
• Conversion of Pig Iron to Steel: Pig iron is converted to steel by removing impurities through oxidation at high temperatures.

C. Copper (Cu)

• Occurrence: Copper is occasionally found as native copper but mainly in compounds such as sulphides, oxides, and carbonates.

• The most important sulphide ores are chalcopyrite (CuFeS2), chalcocite (Cu2S), covellite (CuS), and bornite (Cu5FeS4).

• The principal oxide ores are cuprite (Cu2O) and tenorite (CuO). In carbonate form, it exists as malachite (CuCO3.Cu(OH)_2).

• Extraction: Copper is principally extracted from chalcopyrite.

• The crushed and ground sulphide ore is first concentrated by froth flotation, increasing the copper concentration from 2% to as high as 30%.

• The concentrated ore is then roasted with a limited supply of air:

  • 2CuFeS2(s) + 4O2(g) \rightarrow Cu2S(s) + 2FeO(s) + 3SO2(g)

• The roasted mixture is smelted by adding limestone and sand to form a molten slag to remove many of the impurities.

• FeO can be removed as slag in the form of iron silicate (FeSiO3), and silica as calcium silicate (CaSiO3):

  • CaCO3(s) + SiO2(s) \rightarrow CaSiO3(l) + CO2(g)
  • FeO(s) + SiO2(s) \rightarrow FeSiO3(l)

• Cu2S obtained by roasting chalcopyrite is then reduced by heating it in a limited supply of oxygen: Cu2S(s) + O2(g) \rightarrow 2Cu(l) + SO2(g).

• The copper produced is called blister copper (98.5 – 99.5 % purity).

• Blister copper contains iron, silver, gold, and sometimes zinc as impurities and is refined further by electrolysis.

• Electrolytic refining of copper:

  • Anode: A thick block of impure copper.
  • Cathode: A thin strip of pure copper.
  • Electrolyte: An aqueous solution of copper sulphate, with a small quantity of dilute sulphuric acid.
  • Oxidation at anode: Cu(impure metal) \rightarrow Cu^{2+}(aq) + 2e^-.
  • Reduction at cathode: Cu^{2+}(aq) + 2e^- \rightarrow Cu(metal).

• Physical Properties of Copper:

  • Soft, ductile, malleable, reddish-brown metal with a density of 8.96 g/cm3.
  • Second to silver in electrical conductivity.
  • Melts at 1086°C and boils at 2310°C.

• Chemical properties of copper: Copper is less reactive; it is found in the native state.

  • Powdered copper when heated in air forms a black powder of copper (II) oxide, CuO: 2Cu(s) + O_2(g) \rightarrow 2CuO(s).
  • Copper does not react with dilute acids like HCl and H2SO4, but it can be oxidized by oxidizing acids such as dilute and concentrated nitric acid and hot concentrated sulphuric acid, H2SO4.
  • 3Cu(s) + 8HNO3(aq) (dilute) \rightarrow 3Cu(NO3)2(aq) + 2NO(g) + 4H2O(l)
  • Cu(s) + 4HNO3(aq) (concentrated) \rightarrow Cu(NO3)2(aq) + 2NO2(g) + 2H_2O(l)
  • Cu(s) + 2H2SO4(aq) (Hot and concentrated) \rightarrow CuSO4(aq) + 2SO2(g) + 2H_2O(l)

• Copper corrodes in moist air over a long period of time, forming verdigris: a basic copper carbonate (CuCO3.Cu(OH)2) or Cu2(OH)2CO_3:

  • 2Cu(s) + H2O(l) + O2(g) + CO2(g) \rightarrow CuCO3.Cu(OH)_2

• Copper exhibits different oxidation states (cuprous Cu^+ and cupric Cu^{2+}). 2Cu^+(aq) \rightarrow Cu^{2+}(aq) + Cu(s).

• Uses of Copper:

  • Used to manufacture alloys like bronze (copper and tin), used to make coins, medals, bells, machinery parts, etc.
  • Brass (copper and zinc) is used for hardware tops, terminals, and pipes.
  • Widely used in the electrical industry as electric wires and cables.
  • Copper compounds, such as copper chloride, copper carbonate, and copper hydroxide, are used as pesticides.

5.3 Production of Some Important Non-metals

5.3.1 General Properties of Nonmetals and Common Uses of Some Nonmetallic Compounds

• Nonmetals generally have opposite characteristics to metals.
• Physical properties:
  • Exist as solids, liquids, and gases
  • Non-lustrous
  • Nonmalleable and non-ductile
  • Varying hardness and low density
  • Low melting and boiling points
  • Non-sonorous
  • Poor conductors of heat and electricity
• Chemical properties:
  • React with oxygen on heating or burning to form their oxides
  • Do not displace hydrogen on reaction with dilute acids
  • React with oxygen to form acidic or neutral oxides
  • Combine with hydrogen to form stable hydrides
  • Do not react with water
  • Electronegative (form negative ions by gaining electrons)
  • Oxidizing agents

5.3.2 Production of Nitrogen, Phosphorous, Oxygen, Sulphur, and Chlorine

I. Nitrogen (N)

• Nitrogen is the most abundant element in the atmosphere.
• Occurrence:
  • Exists as a diatomic molecule, N_2, in atmospheric air (about 80% by volume).
  • In compounds, it exists as sodium nitrate (Chile salt peter, NaNO3) and potassium nitrate (KNO3) also called saltpetre.
  • Also found in DNA molecules and proteins of all living things.
• Production:
  • Industrially produced from atmospheric air via fractional distillation.
  • Step 1: Removing dust and other particles from the air.
  • Step 2: Compressing air under high pressure and low temperature removes CO2 and H2O.
  • Step 3: Fractional distillation of liquid air to separate nitrogen (collected and stored in steel cylinders under pressure).
  • Argon distils off at –186°C, leaving behind oxygen that boils at –183°C.
  • In the laboratory, nitrogen is prepared by warming an aqueous solution containing ammonium chloride and sodium nitrite: NH4Cl(aq) + NaNO2(aq) \rightarrow NaCl(aq) + N2(g) + 2H2O(l).
• Physical Properties of Nitrogen:
  • Colorless, odorless, and tasteless gas.
  • Inert under ordinary conditions.
  • The inertness of nitrogen at low temperatures is due to the strength of the triple bond.
• Chemical Properties of Nitrogen:
  • Reacts with metals of group IA and IIA as well as oxygen at higher temperatures.
  • Forms nitrides when heated with reactive metals like lithium, calcium, and magnesium:
  • 6Li(s) + N2(g) \rightarrow 2Li3N(s)
  • 3Ca(s) + N2(g) \rightarrow Ca3N_2(s)
  • 3Mg(s) + N2(g) \rightarrow Mg3N_2(s)
  • Combines with oxygen at elevated temperatures or in an electric arc to form oxides:
    • N2(g) + O2(g) \rightarrow 2NO(g)
    • N2(g) + 2O2(g) \rightarrow 2NO_2(g)
  • Nitric oxide (NO) forms nitrogen dioxide (NO2), a reddish-brown gas: 2NO(g) + O2(g) \rightarrow 2NO_2(g).
  • Nitrogen also forms oxides like dinitrogen monoxide (N2O), dinitrogen trioxide (N2O3), and dinitrogen pentoxide (N2O_5).
  • Reacts directly with hydrogen in the Haber process to form ammonia: N2(g) + 3H2(g) \rightleftharpoons 2NH_3(g).
• Uses of Nitrogen:
  • Used in food packaging to prevent oxidation.
  • Creates an inert atmosphere in the production of semiconductors.
  • Liquid nitrogen is used as a refrigerant to preserve samples.
  • A major use is in the production of ammonia.

II. Phosphorus

• Relatively abundant element, ranking 12th in the Earth’s crust.
• Occurrence: Exists naturally only in the combined state, such as in rock phosphate, Ca3(PO4)2, fluoroapatite, Ca{10}(PO4)6F2 or 3Ca3(PO4)2.CaF_2. Also found in teeth, bones, and DNA.
• Allotropy: Phosphorus exhibits allotropy.
  • The two common allotropic forms of phosphorus are white phosphorus and red phosphorus.
  • White phosphorus: a very poisonous, white waxy substance that melts at 44.1°C and boils at 287°C (density is 1.8 g/cm3). Consists of individual tetra-atomic (P_4) molecules and is an unstable form.
  • Red phosphorus: denser (2.16 g/cm3) and much less reactive than white phosphorus at normal temperatures. Consists of P_4 molecules linked together to form a polymer.
• Extraction: Industrially, white phosphorus is manufactured by heating a mixture of crushed rock phosphate, Ca3(PO4)2, silica, SiO2, and coke in an electric furnace: 2Ca3(PO4)2(s) + 6SiO2(s) + 10C(s) \rightarrow 6CaSiO3(l) + P4(g) + 10CO(g).
  • The vaporized phosphorus (P_4) is condensed, collected, and stored under water.
  • Red phosphorus is prepared by heating white phosphorus in sunlight for several days.
  • White phosphorus is stored under water because it spontaneously ignites in the presence of oxygen.
• Physical properties of phosphorus:
  • Phosphorus is a white, red, or black solid.
• Chemical Properties of Phosphorus:
  • Reacts with limited and excess supplies of oxygen to form tetraphosphorus hexoxide (P4O6) and tetraphosphorus decoxide (P4O{10}) respectively:
    • P4(s) + 3O2(g) \rightarrow P4O6(s)
    • P4(s) + 5O2(g) \rightarrow P4O{10}(s)
  • P4O6 and P4O{10} dissolve in water to form phosphorous acid, H3PO3, an orthophsophoric acid, H3PO4 respectively:
    • P4O6(s) + 6H2O(l) \rightarrow 4H3PO_3(aq)
    • P4O{10}(s) + 6H2O(l) \rightarrow 4H3PO_4(aq)
  • Phosphorus also reacts with limited and excess supplies of chlorine to form phosphorus (III) chloride (PCl3) and phosphorus (V) chloride (PCl5) respectively:
    • P4(s) + 6Cl2(g) \rightarrow 4PCl_3(s)
    • P4(s) + 10Cl2(g) \rightarrow 4PCl_5(s)
• Uses of Phosphorus:
  • Red phosphorous is used to make matches.
  • Most of the white phosphorous produced is used to make phosphoric acid or other phosphorous compounds.
  • Also used in making fireworks, smoke bombs, rat-poisons, and tracer bullets.
  • Essential for plant growth.

III. Oxygen

• Oxygen is the most abundant element in the Earth’s crust (about 46.6% by weight).
• Occurrence:
  • Exists in nature in the elemental state in atmospheric air (about 20% by volume).
• Production:
  • Manufactured industrially by the fractional distillation of liquid air.
  • Oxygen has two allotropic forms, diatomic (O2) and triatomic (O3), ozone.
• Physical Properties:
  • Colorless, odorless, tasteless gas.
  • Changes from a gas to a liquid at a temperature of -182.96°C (slightly bluish color).
  • Density of oxygen is 1.429 grams per liter.
• Chemical Properties of Oxygen:
  • Relatively reactive and combines directly with most elements to form oxides.
  • Combines with metal to form metal oxides (basic oxides). For example:
    • 2Mg(s) + O_2(g) \rightarrow 2MgO(s)
    • 2Ca(s) + O_2(s) \rightarrow 2CaO(s)
  • Combines with non-metals to form acidic oxides; for example:
    • S8(s) + 8O2(g) \rightarrow 8SO_2(g)
    • P4(s) + 5O2(g) \rightarrow P4O{10}(g)
  • Supports combustion and is necessary for the burning of substances like charcoal:
    • C(s) + O2(g) \rightarrow CO2(g)
    • CH4(g) + 2O2(g) \rightarrow CO2(g) + 2H2O(g)
• Uses:
  • Medical: used to treat respiratory illnesses and is essential for surgery and trauma.
  • Manufacturing: used in the production of steel, plastics, and textiles, and in the manufacturing of chemicals.
  • Rocket propulsion: used as a rocket propellant.
  • Mining: used in mining.

IV. Sulphur

• Occurrence:

  • Found in nature and in the form of compounds such as galena (PbS), pyrites (FeS2) cinnabar (HgS), gypsum (CaSO4.2H2O), barite (BaSO4), and hydrogen sulphide (H_2S) in natural gas and crude oil.

• Extraction:

  • Extracted from underground deposits of elemental sulphur by the Frasch process.
  1. Three concentric pipes are sent down to the sulphur deposit.
  2. Superheated water at about 170°C is pumped through the outermost pipes.
  3. Hot air is compressed in the innermost tube. A froth of sulphur, air, and water come out to the surface of the earth forced by hot compressed air in the middle tube.
  4. The molten sulphur is then cooled and solidified.

• Allotropic Forms of Sulphur:

  • Sulphur exhibits allotropy and the important allotropes of sulphur are rhombic and monoclinic sulphur.
  • Rhombic sulphur is the most stable form consisting of S_8 molecules.

• Physical Properties of Sulphur:

  • Pure sulphur is a tasteless, odourless, brittle solid that is pale yellow, a poor conductor of electricity, and insoluble in water.

• Chemical Properties of Sulphur:

  • Relatively stable and un-reactive at room temperature but reacts with metals and non-metals when heated.
    • Combines with metals when heated to form sulphides: 8Fe(s) + S_8(s) \rightarrow 8FeS(s)
    • Burns in oxygen to form oxides: S8(s) + 8O2(g) \rightarrow 8SO_2(g)
  • Sulphur is the raw material for the production of sulphuric acid (H2SO4) by the Contact Process:
    1. Oxidized to produce sulphur dioxide: S8(s) + 8O2(g) \rightarrow 8SO_2(g)
    2. Sulphur dioxide is converted to sulphur trioxide at high temperatures in the presence of a catalyst: 2SO2(g) + O2 \rightarrow 2SO_3(g)
    3. Sulphur trioxide is absorbed into concentrated sulphuric acid to produce oleum (H2S2O7): SO3(g) + H2SO4(l) \rightarrow H2S2O_7(l)
    4. The oleum is then diluted with water to produce the desired concentration of sulphuric acid: H2S2O7(l) + H2O(l) \rightarrow 2H2SO4(aq)

• Uses:

  • Fertilizer: used to make phosphate and ammonium sulphate fertilizers.
  • Pesticides and fungicides
  • Sulphuric acid: primarily used to make sulphuric acid, a key raw material in many industries.
  • Rubber: used in rubber processing and vulcanization.
  • Cosmetics and pharmaceuticals
  • Paper: used in bleaching paper

V. Chlorine

• Chlorine belongs to group VIIA, known as the halogens and is the most abundant element among them.

• Occurrence:

  • Found in nature in the form of compounds only.
  • Found chiefly in the form of chlorides of sodium, potassium, calcium, and magnesium.
  • Sodium chloride is the chief source, obtained from seawater or as deposits of rock salt.

• Extraction:

  • Commercially, chlorine is manufactured by the electrolysis of a concentrated aqueous solution of sodium chloride.
  • Oxidation at anode: 2Cl^–(aq) \rightarrow Cl_2(g) + 2e^–
  • Reduction at cathode: 2H2O(l) + 2e^– \rightarrow H2(g) + 2OH^–(aq)
  • Cell reaction: 2NaCl(aq) + 2H2O(l) \rightarrow 2NaOH(aq) + Cl2(g) + H_2(g)

• Physical Properties of Chlorine:

  • Chlorine is a greenish-yellow gas at room temperature.
  • Melts at –102°C and boils at –34°C.
  • Fairly soluble in water.
  • Extremely poisonous.

• Chemical Properties of Chlorine:

  • Highly reactive non-metal; reacts directly with almost all elements except the noble gases, carbon, and nitrogen.
  • A powerful oxidizing agent.
  • Reactions with heated metals: 2Fe(s) + 3Cl2(g) \rightarrow 2FeCl3(s)
  • 2Al(s) + 3Cl2(g) \rightarrow 2AlCl3(s)
  • Reacts smoothly with hydrogen: H2(g) + Cl2(g) \rightarrow 2HCl(g)
  • Displaces less reactive halogens (Br2 and I2) from aqueous solutions of their compounds:
    • Cl2(g) + 2KBr(aq) \rightarrow 2KCl(aq) + Br2(l)
    • Cl2(g) + 2Kl(aq) \rightarrow 2KCl(aq) + I2(aq)
  • Dissolves in water and reacts with it: Cl2(g) + H2O(l) \rightarrow HCl(aq) + HOCl(aq).

• Uses:

  • Commercially used as a bleaching agent and disinfectant.
  • Strong enough to oxidize dyes and to bleach coloured materials from wood pulp, paper and cotton.