Limestone, industrial chemistry and Earth's resources

Limestone and its uses

Limestone is mainly calcium carbonate (CaCO₃). It is quarried and used directly (building stone, aggregate) or processed:

Thermal decomposition (calcination): CaCO₃(s) → CaO(s) + CO₂(g) (heated in a kiln)

CaO = quicklime (calcium oxide). Add water to produce slaked lime (calcium hydroxide): CaO(s) + H₂O(l) → Ca(OH)₂(s) ΔH = −65 kJ/mol (exothermic)

Limewater = dilute aqueous Ca(OH)₂ — used as the test for CO₂ (turns milky/cloudy due to CaCO₃ precipitate forming).

Uses of calcium compounds:

• CaCO₃: glass-making, cement, road aggregate, blast furnace (removes acidic SiO₂)
• CaO: neutralising acidic soil, steel-making (removes Si impurities as slag), cement manufacture
• Ca(OH)₂: agriculture (to raise soil pH), water treatment, mortar (with sand)

Cement and concrete: Cement = CaO + SiO₂ + Al₂O₃ (from limestone + clay). Add water → hydration reactions form calcium silicate hydrate → concrete hardens.

The Haber process

The Haber process manufactures ammonia (NH₃) from nitrogen and hydrogen: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ/mol

Key features:

• Reversible reaction (double arrow ⇌) — an equilibrium is established.
• Conditions: temperature ~450 °C; pressure ~200 atm; iron catalyst (with Al₂O₃/K₂O promoters).
• Gases are recycled; ammonia liquefied and separated as it forms.

Why these conditions? (Le Chatelier's Principle)

• Lower temperature favours the forward reaction (exothermic → equilibrium shifts right → more NH₃). BUT too low a temperature → very slow rate. 450 °C is a compromise.
• Higher pressure favours the side with fewer moles of gas (left: 4 mol; right: 2 mol → forward shift → more NH₃). BUT very high pressure is costly and hazardous. 200 atm is a compromise.
• Iron catalyst does not shift equilibrium but allows it to be reached faster at lower temperature.

Equilibrium yield vs rate: WJEC examiners frequently ask you to evaluate why conditions are a compromise between yield and rate.

The Contact process (sulfuric acid manufacture)

Stage 1: S + O₂ → SO₂ (burning sulfur or roasting sulfide ores) Stage 2: 2SO₂ + O₂ ⇌ 2SO₃ (vanadium(V) oxide catalyst, V₂O₅; 450 °C; 1–2 atm) Stage 3: SO₃ + H₂SO₄ → H₂S₂O₇ (oleum), then H₂S₂O₇ + H₂O → 2H₂SO₄

Sulfuric acid is used in fertiliser manufacture, paint, detergents, car batteries.

Fertilisers

Plants need nitrogen (N), phosphorus (P), and potassium (K) for growth. Fertilisers supply these.

Nitrogen-containing fertilisers: made from ammonia (from the Haber process):

• Ammonium nitrate (NH₄NO₃): NH₃ + HNO₃ → NH₄NO₃
• Ammonium sulfate: 2NH₃ + H₂SO₄ → (NH₄)₂SO₄
• Urea: CO(NH₂)₂

Issues with fertilisers:

• Eutrophication: excess fertiliser washes into waterways (leaching/run-off) → algae bloom → algae die → bacteria decompose → bacteria use up oxygen → aquatic animals die (deoxygenation).
• Nitrate in drinking water: high concentrations can cause health problems (blue baby syndrome in infants).

Sustainable chemistry

WJEC links resource use to sustainability: finite vs renewable resources, recycling metals, green chemistry principles (reduce waste, atom economy, catalysis, renewable feedstocks).

Common examiner traps

1. Catalyst in Haber process does not shift equilibrium: it only speeds up attainment of equilibrium — yield (%) is unchanged; just reached faster.
2. Eutrophication sequence: many students stop at "algae grow." The key step that causes deaths is bacterial decomposition depletes O₂.
3. Low temperature gives higher yield but slower rate: the two are in conflict — this is why 450 °C is used, not a temperature that maximises either alone.
4. Contact process uses V₂O₅ not iron: iron is the Haber catalyst. V₂O₅ is used in the Contact process.