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Laboratory vs industrial preparation of ammonia

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Laboratory vs industrial preparation of ammonia

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your exam.

Nitrogen (N₂) is a diatomic non-metal held together by a triple bond (+945 kJ mol⁻¹), which is why it is unreactive at room temperature. Its most important compounds are ammonia (NH₃) and nitric acid (HNO₃).

  • Haber–Bosch: N₂(g) + 3H₂(g) ⇌ 2NH₃(g); ΔH = −92 kJ mol⁻¹; Fe/K₂O/Al₂O₃ catalyst, ~450 °C, 200 atm.
  • Ostwald: 4NH₃ + 5O₂ → 4NO + 6H₂O (Pt/Rh gauze, 900 °C), then 2NO + O₂ → 2NO₂, then 3NO₂ + H₂O → 2HNO₃ + NO.
  • Dilute HNO₃ + Cu → NO (colourless); concentrated HNO₃ + Cu → NO₂ (brown).
  • Brown ring test: NO₃⁻ + 3Fe²⁺ + 4H⁺ → NO + 3Fe³⁺ + 2H₂O; brown complex = [Fe(H₂O)₅(NO)]²⁺.

Mnemonic: “Dilute gives NO, dense gives NO₂” — only oxidising acids release nitrogen oxides, never H₂.

🟡 Standard — Regular Study (2d–2mo)

Standard content for students with a few days to months.

Laboratory vs industrial preparation of ammonia

In the lab, ammonia is generated by warming an ammonium salt with a strong base, e.g. NH₄Cl(s) + NaOH(s) → NaCl(s) + NH₃(g) + H₂O(l). The gas must be dried over calcium oxide (CaO), not concentrated H₂SO₄, because ammonia is basic and would be absorbed by the acid. Industrially, the Haber–Bosch process combines atmospheric N₂ with H₂ (from methane steam reforming) at ~200 atm and 450–500 °C over an iron catalyst promoted by K₂O and Al₂O₃. These are a Le Chatelier compromise: lower temperature favours yield but slows rate; higher pressure favours yield (4 moles of gas → 2 moles).

Ammonia as a base and ligand

The lone pair on nitrogen makes NH₃ both a Brønsted–Lowry base (accepts H⁺ to form NH₄⁺) and a Lewis base / ligand. With Cu²⁺(aq) it forms the deep-blue tetraamminecopper(II) ion [Cu(NH₃)₄]²⁺ — the standard test for Cu²⁺. With HCl(g) it produces dense white fumes of NH₄Cl.

The Ostwald process and nitric acid

StepEquationConditions
14NH₃ + 5O₂ → 4NO + 6H₂OPt/Rh gauze, ~900 °C
22NO + O₂ → 2NO₂cool, excess air
33NO₂ + H₂O → 2HNO₃ + NOabsorption tower

The product is ~68% HNO₃; distillation with concentrated H₂SO₄ removes water to give concentrated acid. HNO₃ is a strong monobasic acid and an oxidising agent; its oxidising power scales with concentration.

Common mistake

Writing the Brown ring as [Fe(NO)]²⁺ ignores the five water ligands and loses the +2 charge calculation. Always write [Fe(H₂O)₅(NO)]²⁺.

🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for students on a longer study timeline.

Thermal decomposition of nitrates — a group trend

The pattern of nitrate decomposition reveals the stability of the corresponding metal cation:

  • Group 1 nitrates (Li → Cs) give the nitrite + O₂: 2NaNO₃ → 2NaNO₂ + O₂.
  • Group 2 nitrates (and Al) give the metal oxide + NO₂ + O₂: 2Ca(NO₃)₂ → 2CaO + 4NO₂ + O₂.
  • Heavy metals (Ag, Hg) often give the metal + NO₂ + O₂.

Mechanistically, larger cations polarise the NO₃⁻ less, so the N–O bond stays intact enough to lose only one oxygen. Smaller, more charge-dense cations pull the N–O bond apart, giving the metal oxide. Exam questions routinely ask you to deduce the cation identity from a decomposition equation.

Worked example — percentage yield

In an industrial Haber–Bosch run, 28 kg of N₂ was mixed with excess H₂. At equilibrium, 17 kg of NH₃ was recovered. Theoretical moles from N₂ = 28000 g ÷ 28 g mol⁻¹ = 1000 mol N₂ → 2000 mol NH₃ = 34 kg. Actual = 17 kg = 1000 mol. % yield = (1000/2000) × 100 = 50%. Use PV = nRT (R = 8.314 J K⁻¹ mol⁻¹) when volume is given instead of mass.

Environmental chemistry

NOₓ from vehicle engines and lightning dissolves in rainwater to give HNO₃ (acid rain). Run-off of NH₄NO₃ and (NH₄)₂SO₄ fertilisers causes eutrophication — algal blooms deplete O₂ and kill aquatic life. Anhydrous NH₃ is injected directly into soil as a high-N fertiliser.

Two practice prompts

  1. Explain, with equations, why ammonia is dried over calcium oxide and not concentrated sulphuric(VI) acid.
  2. A colourless gas X turns brown in air and dissolves in water to give an acidic solution. Identify X and write equations for the two changes observed.

Exam tip: 1–2 structured questions per UACE paper test this topic — usually one Haber/Ostwald equation question and one nitrate/oxide identification.

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