Sulphur

General Properties of Group VI Elements (Sulphur Group)

The Group VI elements include: Oxygen, Sulphur, Selenium, Tellurium, and Polonium.

1. Metallic character increases as you move down the group. Oxygen and sulphur are non-metals, selenium and tellurium are metalloids, while polonium is a metal.
2. All members are solids at room temperature except oxygen, which is a gas.
3. Oxygen and sulphur exhibit allotropy.
4. They all have six electrons in their outermost shell, giving them a common oxidation number of -2. However, sulphur can also show oxidation states of -4 and +6 in some compounds.
5. Electronegativity decreases down the group, making oxygen the most electronegative and a strong oxidizing agent.

Electronic Structure of Group VI Elements

• Oxygen (8): 1s² 2s² 2p⁴
• Sulphur (16): 1s² 2s² 2p⁶ 3s² 3p⁴
• Selenium (34): 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁴
• Tellurium (52): 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ 5s² 5p⁴
• Polonium (84): 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ 5s² 5p⁶ 5d¹⁰ 5f¹⁴ 6s² 6p⁴

Sulphur

Sulphur is a non-metal that occurs naturally as free elements and in compounds like sulphides and sulphates.

Extraction of Sulphur (Frasch Process)

In this method, a special triple-pipe device is inserted into the sulphur bed:

• Superheated water at 170°C melts sulphur (which melts at 115°C).
• Hot compressed air (at 15 atm) forces the molten sulphur to the surface.
• The extracted sulphur (99.5% pure) solidifies in a tank.

Allotropes of Sulphur

• Rhombic Sulphur: Yellow, octahedral crystals. Stable below 96°C. Made of S₈ rings.
• Monoclinic Sulphur: Amber, needle-like crystals. Stable above 96°C. Converts back to rhombic form below 96°C.

Comparison of Physical Properties

• Melting point: Rhombic = 113°C, Monoclinic = 119°C
• Color: Rhombic = Bright yellow, Monoclinic = Amber
• Shape: Rhombic = Octahedral, Monoclinic = Needle-like
• Transparency: Rhombic = Translucent, Monoclinic = Transparent

Non-crystalline forms of sulphur include:

• Amorphous sulphur
• Plastic sulphur

Physical Properties of Sulphur

• Yellow solid
• Insoluble in water, soluble in toluene and carbon disulphide
• Poor conductor of heat and electricity
• Melting point: 119°C; Boiling point: 444°C

Chemical Properties of Sulphur

• With metals:
  Fe(s) + S(s) → FeS(s)
• With excess oxygen:
  S(s) + O2(g) → SO2(g)
• With hydrogen:
  H2(g) + S(s) → H2S(g)
• With carbon:
  C(s) + 2S(s) → CS2(l)

Uses of Sulphur

• Manufacture of H₂SO₄
• Vulcanization of rubber
• Production of germicides
• Bleaching agents

Hydrogen Sulphide (H₂S)

Found in volcanic gases, sulphur springs, coal and natural gas.

Laboratory Preparation

Reacting dilute acids with metallic sulphides like iron(II) sulphide:

2HCl(aq) + FeS(s) → FeCl2(aq) + H2S(g)

Kipp's apparatus is used for continuous supply.

Physical Properties

• Colourless gas with rotten egg smell
• Highly toxic
• 1.18 times denser than air
• Moderately soluble in water (weakly acidic)
• Burns with pale blue flame

Chemical Properties

• With alkali:
  2NaOH(aq) + H2S(g) → Na2S(aq) + 2H2O(l)
• With oxygen:
  2H2S(g) + 3O2(g) → 2H2O(l) + 2SO2(g)
2H2S(g) + O2(g) → 2H2O(l) + 2S(s)
• Acts as a reducing agent with oxidizing agents like acidified KMnO₄, K₂Cr₂O₇, chlorine, FeCl₃, SO₂, HNO₃

Test for H₂S

Moisten filter paper with Pb(NO₃)₂ and expose to gas. Blackening indicates H₂S.

Sulphur(IV) Oxide (SO₂)

Laboratory Preparation

By heating sodium trioxosulphate(IV) with HCl:

Na2SO3(aq) + 2HCl(aq) → 2NaCl(aq) + H2O(l) + SO2(g)

Physical Properties

• Colourless gas with choking smell (like burning matches)
• Poisonous and very soluble in water
• 2.5 times denser than air

Chemical Properties

• With alkali:
  2NaOH(aq) + SO2(g) → Na2SO3(aq) + H2O(l)
• As a reducing agent: reacts with acidified KMnO₄, K₂Cr₂O₇, FeCl₃, etc.
• Acts as a bleaching agent in presence of water
• Acts as oxidizing agent with H₂S:
  2H2S(g) + SO2(g) → 2H2O(l) + 3S(s)
• With carbon:
  C(s) + SO2(g) → CO2(g) + S(s)

Test for SO₂

• If it bleaches a flower, suspect SO₂.
• Bubble gas through acidified K₂Cr₂O₇ or KMnO₄. If K₂Cr₂O₇ turns green or KMnO₄ becomes colourless, SO₂ is confirmed.

Uses of SO₂

• Manufacture of H₂SO₄
• Used as germicide and fumigant
• Used to bleach straw, silk, and wood
• Preservative in drinks like orange juice
• Liquid SO₂ is used as a refrigerant

Sulphur (VI) Oxide (SO₃)

Preparation

Sulphur (VI) oxide is produced by reacting sulphur (IV) oxide (SO₂) with oxygen (O₂) under specific conditions:

1. Presence of a catalyst such as platinized asbestos or vanadium (V) oxide (V₂O₅).
2. Pressure of 1 atm.
3. Temperature range of 400°C – 450°C.

Equation: 2SO₂(g) + O₂(g) → 2SO₃(g)

Physical Properties

1. It exists as white, needle-like crystals at room temperature.
2. It has a low boiling point and vaporizes on gentle heating.
3. It dissolves readily in water to form tetraoxosulphate (VI) acid (H₂SO₄).

Trioxosulphate (IV) Acid (H₂SO₃)

Laboratory Preparation

H₂SO₃ is a dibasic acid prepared by the following process:

1. Sodium trioxosulphate (IV) is heated with dilute hydrochloric acid to release sulphur (IV) oxide (SO₂).
2. The SO₂ gas is then dissolved in water.

Equations:

Na₂SO₃(s) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + SO₂(g)

H₂O(l) + SO₂(g) → H₂SO₃(aq)

Physical Properties

1. It is a colourless liquid.
2. It turns blue litmus paper red.
3. It mixes readily with water.
4. It has an irritating, choking smell.

Chemical Properties

1. Reacts with alkalis to form salt and water:
   2NaOH(aq) + H₂SO₃(aq) → Na₂SO₃(aq) + 2H₂O(l)
2. Oxidized by air to form tetraoxosulphate (VI) acid:
   2H₂SO₃(aq) + O₂(g) → 2H₂SO₄(aq)
3. Reduces oxidizing agents like KMnO₄ and K₂Cr₂O₇.
4. Bleaches dyes in the presence of water.

Test for SO₃²⁻ Ion

Add barium chloride solution. A white precipitate that dissolves in dilute HCl confirms the presence of trioxosulphate (IV).

Uses of H₂SO₃

1. Used for bleaching straw and fabrics.
2. Used as a germicide.

Tetraoxosulphate (VI) Acid (H₂SO₄)

Industrial Preparation (Contact Process)

1. Burn sulphur in dry air to form SO₂:
   S(s) + O₂(g) → SO₂(g)
2. Purify the gas and dry using concentrated H₂SO₄.
3. Pass through contact tower with V₂O₅ catalyst at 450–500°C:
   2SO₂(g) + O₂(g) → 2SO₃(g)
4. Dissolve SO₃ in concentrated H₂SO₄ to form oleum:
   H₂SO₄(aq) + SO₃(g) → H₂S₂O₇(aq)
5. Dilute oleum to form 98% H₂SO₄:
   H₂S₂O₇(aq) + H₂O(l) → 2H₂SO₄(aq)

Note: SO₃ is not directly dissolved in water due to the intense heat evolved, which forms an acid mist.

Physical Properties

1. Colourless, viscous liquid with density 1.84 g/cm³.
2. Highly corrosive and causes burns.
3. Highly soluble in water with much heat evolution.

Chemical Properties

1. Reacts with metals above hydrogen to produce hydrogen gas:
   Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)
2. Reacts with basic oxides:
   MgO(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂O(l)
3. Reacts with alkalis to form salts:
   H₂SO₄(aq) + NaOH(aq) → NaHSO₄(aq) + H₂O(l)
   H₂SO₄(aq) + 2NaOH(aq) → Na₂SO₄(aq) + 2H₂O(l)
4. Reacts with carbonates to release CO₂:
   H₂SO₄(aq) + CuCO₃(s) → CuSO₄(aq) + H₂O(l) + CO₂(g)
5. Acts as an oxidizing agent:
   Zn(s) + 2H₂SO₄(conc) → ZnSO₄(aq) + 2H₂O(l) + SO₂(g)
   C(s) + 2H₂SO₄(conc) → CO₂(g) + 2SO₂(g) + 2H₂O(l)
   H₂S(g) + H₂SO₄(conc) → S(s) + SO₂(g) + H₂O(l)
6. Acts as a dehydrating agent:
   C₁₂H₂₂O₁₁(s) → 12C(s) + 11H₂O(l)
7. Displaces volatile acids from their salts:
   KCl(s) + H₂SO₄(aq) → KHSO₄(aq) + HCl(g)

Test for SO₄²⁻ Ion

Add barium chloride solution. A white precipitate that is insoluble in excess dilute HCl confirms the presence of sulphate ion.

Uses of H₂SO₄

1. Used in making fertilizers (e.g. ammonium tetraoxosulphate (VI)).
2. Used in crude oil purification.
3. Used as an electrolyte in lead-acid batteries.
4. Used as a drying agent (except for NH₃ and H₂S gases).
5. Used to clean metals before electroplating.

Uses of Tetraoxosulphate (VI) Salts

1. Ammonium salt is used as fertilizer.
2. Sodium salt is used in paper manufacturing and as a purgative.
3. Calcium salt (gypsum) is used to make plaster of Paris.
4. Aluminium salt is used in water treatment.
5. Iron (II) salt is used to treat anaemia.