Sulphur
Sulphur sits in Group VI, period 3, making up about 0.1% of the Earth’s crust, and occurs as free deposits in places like the USA, Poland, Japan, New Zealand and Sicily. Underground deposits are mined using the Frasch process (developed by Hermann Frasch): hot water and compressed air are pumped down a specialised pipe, dissolving and lifting the sulphur to the surface as a mixture with air and water, after which the water is evaporated off.
Allotropes of Sulphur
Allotropy is the existence of an element in two or more different forms in the same physical state. Sulphur’s two main allotropes are rhombic and monoclinic sulphur (a further form, “plastic sulphur”, also exists).
| Rhombic sulphur | Monoclinic sulphur | |
|---|---|---|
| Colour | Bright yellow | Amber |
| Shape | Octahedral crystals | Needle-shaped crystals |
| Density | 2.08 g/cm³ | 1.98 g/cm³ |
| Melting point | 113°C | 119°C |
| Stable range | Below 96°C | 96°C–119°C |
Properties of Sulphur
Physical: a yellow solid, insoluble in water, melting at 119°C and boiling at 444°C.
Chemical:
- With metals: forms sulphides directly, e.g. Fe(s) + S(s) → FeS(s); 2Cu(s) + S(s) → Cu₂S(s).
- With oxygen: burns in air to form sulphur dioxide: S(s) + O₂(g) → SO₂(g).
- With hydrogen: reacts slowly at high temperature to form hydrogen sulphide: H₂(g) + S(s) → H₂S(g).
- With carbon: mixed with coke in an electric furnace, forms carbon disulphide: C(s) + 2S(g) → CS₂(l).
Uses: manufacturing sulphuric acid; vulcanising rubber; making matches, gunpowder and fireworks; manufacturing dyes, fungicides and insecticides.
Hydrogen Sulphide (H₂S)
Found in natural gas, coal gas, sulphur springs, and gases given off as sulphur-containing organic matter decays.
Preparation: reacting a dilute acid with a metal sulphide such as iron(II) sulphide: 2HCl(aq) + FeS(s) → FeCl₂(aq) + H₂S(g). The gas is dried using fused calcium chloride — not concentrated sulphuric acid, which would reduce it to sulphur instead of just drying it.
Properties: a colourless, poisonous gas smelling of rotten eggs, slightly denser than air, moderately soluble in water, turning moist blue litmus red, and burning with a pale blue flame in plentiful oxygen: 2H₂S(g) + 3O₂(g) → 2H₂O(g) + 2SO₂(g).
Chemical behaviour:
- As a weak acid: bubbled through an alkali, forms a normal salt and water, e.g. 2NaOH(aq) + H₂S(g) → Na₂S(aq) + 2H₂O(l).
- As a precipitating agent: forms distinctively coloured, insoluble sulphides with many metal salt solutions — useful for identifying metal ions. E.g. Pb(CH₃COO)₂(aq) + H₂S(g) → PbS(s) (black) + 2CH₃COOH(aq); ZnSO₄(aq) + H₂S(g) → ZnS(s) (white) + H₂SO₄(aq).
- As a reducing agent: a strong one — hydrogen sulphide is itself oxidised to sulphur in the process, e.g. reacting with acidified potassium permanganate: 2KMnO₄(aq) + 3H₂SO₄(aq) + 5H₂S(g) → K₂SO₄(aq) + 2MnSO₄(aq) + 8H₂O(l) + 5S(s); or with chlorine: Cl₂(g) + H₂S(g) → 2HCl(g) + S(s).
Test: its rotten-egg smell, and it turns filter paper moistened with lead(II) nitrate solution black: Pb(NO₃)₂(aq) + H₂S(g) → PbS(s) + 2HNO₃(aq).
Sulphur Dioxide (SO₂)
Preparation: in the lab, by heating copper with concentrated sulphuric acid: Cu(s) + 2H₂SO₄(aq) → CuSO₄(aq) + 2H₂O(l) + SO₂(g). Industrially, by burning sulphur or roasting metal sulphides: S(s) + O₂(g) → SO₂(g); 2PbS(s) + 3O₂(g) → 2PbO(s) + 2SO₂(g).
Properties: a colourless gas with an irritating, burnt-match smell; very soluble in water; poisonous; about 2.5 times denser than air; turns moist blue litmus red.
Chemical behaviour:
- As an acid: reacts with alkalis to form a sulphite salt and water, e.g. SO₂(g) + 2NaOH(aq) → Na₂SO₃(aq) + H₂O(l).
- As a reducing agent: in the presence of water, forms sulphite ions that readily donate electrons to oxidising agents, e.g. reducing acidified potassium permanganate.
- As a bleaching agent: dissolves in water to form sulphurous acid (H₂SO₃), which bleaches dyes. This bleaching is less permanent than chlorine’s, since atmospheric oxygen can re-oxidise the bleached dye back to its original colour.
- As an oxidising agent: in the presence of a strong reducing agent, e.g. 2H₂S(g) + SO₂(g) → 2H₂O(l) + 3S(s).
Test: recognisable by its irritating smell; confirmed because it bleaches coloured flowers and decolourises purple potassium permanganate solution.
Uses: manufacturing sulphuric acid; bleaching; as a refrigerant (liquefies easily); as a food preservative (canned goods, jams, fruit juice); as a germicide and fumigant.
