Where Oxygen Comes From
Oxygen sits in Group VI, period 2. It makes up 21% of atmospheric air by volume in its free (uncombined) form, and roughly half the mass of the Earth’s crust in combined forms — sulphur dioxide, silicon dioxide, calcium carbonate, water, and more.
Preparing Oxygen
Industrially: by fractional distillation of liquid air. In the laboratory: by gently heating powdered potassium chlorate (KClO₃) with manganese(IV) oxide, MnO₂, as a catalyst (a substance that speeds up a reaction without being consumed by it): 2KClO₃(s) —MnO₂→ 2KCl(s) + 3O₂(g).
Properties of Oxygen
Physical: colourless, odourless, tasteless; no effect on litmus paper; slightly soluble in water (important for aquatic life); slightly heavier than air.
Chemical:
- With metals (especially Group I/II): forms basic oxides, e.g. 4K(s) + O₂(g) → 2K₂O(s); 2Ca(s) + O₂(g) → 2CaO(s). These oxides dissolve in water to give alkaline solutions, e.g. K₂O(s) + H₂O(l) → 2KOH(aq).
- With non-metals: forms acidic oxides, e.g. C(s) + O₂(g) → CO₂(g); S(s) + O₂(g) → SO₂(g); P₄(s) + 5O₂(g) → P₄O₁₀(g). These dissolve in water to form acids — CO₂ gives carbonic acid (H₂CO₃), SO₂ gives sulphurous acid (H₂SO₃), and P₄O₁₀ gives phosphoric acid (H₃PO₄).
- Combustion examples: propane burning, C₃H₈(g) + 5O₂(g) → 3CO₂(g) + 4H₂O(l); iron(II) sulphide roasting, 4FeS(s) + 7O₂(g) → 2Fe₂O₃(s) + 4SO₂(g).
Test for oxygen: it rekindles a glowing splint. Uses: respiration; steelmaking; oxy-acetylene cutting torches; manufacturing nitric and sulphuric acid.
Types of Oxides
| Type | Description | Examples |
|---|---|---|
| Acidic | Oxides of non-metals; dissolve in water to form acids | CO₂, SO₂, P₄O₁₀ |
| Basic | Oxides of metals; neutralised by acids to form salt and water (some also dissolve in water to form alkalis, e.g. Na₂O(s) + H₂O(l) → 2NaOH(aq)) | Na₂O, K₂O, CaO |
| Amphoteric | Behave as an acid in the presence of a base, and as a base in the presence of an acid | Al₂O₃, ZnO, PbO |
| Neutral | Neither acidic nor basic; no effect on litmus | CO, N₂O, H₂O |
| Peroxide | Contain a higher proportion of oxygen than the normal oxide | H₂O₂, Na₂O₂ |
| Compound (mixed) | A combination of two oxides of the same metal | Pb₃O₄ (2PbO + PbO₂); Fe₃O₄ (FeO + Fe₂O₃) |
Aluminium oxide’s amphoteric behaviour: acting as an acid, Al₂O₃(s) + 2NaOH(aq) + 3H₂O(l) → 2NaAl(OH)₄(aq) (sodium aluminate); acting as a base, Al₂O₃(s) + 3H₂SO₄(aq) → Al₂(SO₄)₃(aq) + 3H₂O(l).
Reduction of Oxides
Reduction means removing oxygen from an oxide. Oxides of moderately reactive metals are reduced using carbon monoxide; less reactive metals’ oxides can be reduced using hydrogen gas:
- ZnO(s) + CO(g) → Zn(s) + CO₂(g)
- CuO(s) + CO(g) → Cu(s) + CO₂(g)
- Fe₂O₃(s) + 3H₂(g) → 2Fe(s) + 3H₂O(l)
- PbO(s) + H₂(g) → Pb(s) + H₂O(l)
Non-metal oxides, mostly gases, are reduced by passing them through hot coke (carbon): CO₂(g) + C(s) → 2CO(g).
