Chlorine and the Halogens

Chlorine is the most important element in the halogen family. It does not occur as free element in nature because it is too reactive. It is usually found in combined state as chlorides.

Laboratory Preparation of Chlorine

  1. By the oxidation of concentrated HCl with strong oxidizing agent such as MnO2 or KMnO4

    2. MnO2(s) + 4HCl(aq) → MnCl2(aq) + 2H2O(l) + Cl2(g)

  2. By heating concentrated H2SO4 with a mixture of NaCl and MnO2

    3. 2NaCl(s) + MnO2(s) + 2H2SO4(aq) → Na2SO4(aq) + MnSO4(aq) + H2O(l) + Cl2(g)

Industrial Preparation

Chlorine is manufactured industrially by the electrolysis of brine and molten NaCl, MgCl2 or CaCl2.

Physical Properties

  1. Chlorine is a greenish-yellow gas with unpleasant chocking smell.
  2. It is a poisonous gas.
  3. It can be liquefied under a pressure of about 6atm.
  4. It is moderately soluble in water.
  5. It is about 2.5 times denser than air.

Chemical Properties

  1. It is very reactive and tends to attain stability by forming electrovalent compound with metals and a single covalent bond compounds with non-metals.

    2. 2Na(s) + Cl2(g) → 2NaCl(s)

    Cl2(g) + H2(g) → 2HCl(g)

  2. It displaces other halogens from solution of their acids and salts. It is the most reactive halogen.

    3. Cl2(g) + NaI(aq) → 2NaCl(aq) + I2(g)

  3. It combines directly with other elements except oxygen, nitrogen carbon and the noble gases; to form chlorides.

    4. Ca(s) + Cl2(g) → CaCl2(s)

  4. It has a very strong affinity for hydrogen; it removes hydrogen from its compounds.

    5. C10H12(l) + 8Cl2(g) → 10C(s) + 16HCl(g)

  5. It is a powerful oxidizing agent: it oxidizes green Fe2+ to yellow Fe3+

    6. 2FeCl2(aq) + Cl2 →2FeCl3(aq)

  6. It reacts with hot concentrated NaOH solution to give a mixture of trioxochlorate (V) and chloride of the metal.

    7. 6NaOH + 3Cl2(g) → NaClO3(aq) + NaCl(aq) + H2O(l)

    • With cold dilute solution of NaOH, a pale yellowish mixture of oxochlorate (I) and chloride of the metal is formed.

      • 2NaOH(aq) + Cl2(g) → NaOCl(aq) + NaCl(aq) + H2O(l)

  7. It has a bleaching action: in the presence of water, chlorine bleaches most dyes and inks except printer’s ink. The bleaching action of chlorine is due to its ability to react with water to form oxochlorate (I) acid which decomposes to release oxygen which oxidizes the dye to form a colourless compound.
    • H2O(l) + Cl2(g) → HCl(aq) + HOCl(aq)
    • HOCl(aq) → HCl(aq) + [O]
    • Dye + [O]→ [Dye + O]
  8. It reacts with slaked lime solutions to produce bleaching powder

    9. Ca(OH)2(aq) + Cl2(g)→ CaOCl2.H2O(s)

Tests for Chlorine

  1. It turns damped starch-iodide dark blue. Chlorine turns starch-iodide paper blue because it displaces iodine from the iodide. The iodine liberated then turns the starch blue
  2. It turns damped blue litmus paper pink and then bleaches it. It is acidic gas.

Uses of Chlorine

  1. It is used as a bleaching agent for cotton, wool, pulp etc.
  2. It is a powerful germicide [due to its oxidizing nature].
  3. It is used in the manufacture of polyvinyl chloride (PVC) and synthetic rubber.
  4. It is used for making domestic antiseptics e.g acidified NaClO solution.
  5. It is used in producing KClO3, for making matches and fireworks.
  6. It is used in the manufacture of organic compound e.g CHCl3, CCl4.
  7. It is used for making NaClO3, a weed killer.

Compounds of Chlorine

Hydrogen Chloride

Hydrogen chloride (marine-acid gas) exists as a gas at room temperature. It dissolves in water to form hydrochloric acid. It occurs in traces in the air as industrial by-product and is considered as air pollutant; but it can be easily washed down as acid rain since it is very soluble in water.

Laboratory Preparation

The gas is prepared by the action of hot concentrated H2SO4 on any soluble chloride.

Example: 2NaCl(s) + H2SO4(aq) → Na2SO4(aq) + 2HCl(g)

Note: NaHSO4 is first formed at a lower temperature and later at higher temperature HCl gas is formed. The gas is dried by passing it through concentrated H2SO4 in another flask and collected.

Industrial Preparation

Pure HCl gas can be produced in large scale by direct combination of hydrogen and chloride gas obtained from the electrolysis of brine.
H2(g) + Cl2(g) → 2HCl(g)

Physical Properties

  1. Pure HCl gas is colourless and has sharp irritating smell
  2. It turns damp blue litmus paper red
  3. It is very soluble in water, forming aqueous HCl acid
  4. It is readily dissolved in non-polar solvent like chloroform and toluene; but the solution does not conduct electricity and has no acidic properties because hydrogen chloride which is a covalent molecule does not ionize when it dissolve in non-polar solvents. But it dissolves in water and ionizes. The ions formed in aqueous solution are responsible for the acidic property and conductivity of its aqueous solution.
  5. It forms misty fumes in moist air because it dissolves in the moisture to form tiny droplets of HCl acid.
  6. It is about 1.25times denser than air
  7. It does not support combustion.

Chemical Properties

  1. It combines directly with NH3 and produces a white fumes of ammonium chloride.

    2. HCl(g) + NH3(g) → NH4Cl(s)

  2. It reacts with various heated metals to form their respective chloride and hydrogen.

    3. Zn(s) + 2HCl(g) → ZnCl2(s) + H2(g)

Tests for Hydrogen Chloride

  1. A gas rod that has been dipped in ammonia solution is brought near the gas jar containing the unknown gas, if there are dense white fumes on the glass rod, then the gas is hydrogen chloride gas.
  2. Few drops of silver trioxonitrate (V) is added to the gas jar containing the unknown gas and shaken. If white precipitate of silver chloride is observed, then the gas is hydrogen chloride gas.

Chlorides

Chlorides are normal salts formed when metallic ion replace the hydrogen ion in hydrochloric acid. Chlorides are prepared by neutralization reaction. Chlorides are soluble in water with exception of few.

  1. Soluble Chlorides: NaCl, NH4Cl, KCl CaCl2 etc
  2. Insoluble Chlorides: CuCl2, AgCl, PbCl2

Properties of Chlorides

  1. Chlorides are not decomposed by heat. They can only be recovered from solution by evaporation to dryness or sometimes by crystallization.
  2. They react with hot concentrated tetraoxosulphate (VI) acid to produce hydrogen chloride gas.

    3. 2NaCl(s) + H2SO4(aq) → Na2SO4(aq) + 2HCl(g)

    On heating a chloride with concentrated tetraoxosulphate (VI) in the presence of a strong oxidizing agent, chlorine is produced.
    ZnCl2(s) + KMnO4(s) + 2H2SO4(aq) → ZnSO4(aq) + K2SO4(aq) + 2MnO2(aq) + 2H2O(l) + Cl2(g)

Test for Chlorides

The test solution is acidified with dilute trioxonitrate (V) acid to prevent precipitation of other salts. Few drops of AgNO3(aq) is then added to the acidified solution in a test tube, a white precipitate of AgCl which readily dissolves in excess NH3(aq) solution indicates the presence of a chloride.