Industrial Manufacture of Ammonia (Haber Process)

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Nitrogen (N₂) is the major component of the atmosphere (~78%) and enters the nitrogen cycle through fixation, nitrification, assimilation, and denitrification. The two key industrial processes are:

Haber process: N₂ + 3H₂ ⇌ 2NH₃ (Fe catalyst, 200 atm, 500 °C)
Ostwald process: 4NH₃ + 5O₂ → 4NO + 6H₂O, then 2NO + O₂ → 2NO₂, then 3NO₂ + H₂O → 2HNO₃ + NO (Pt/Rh catalyst)

Ammonia is a Lewis base (lone pair on N) that forms the deep-blue [Cu(NH₃)₄]²⁺ complex with Cu²⁺. In the lab, NH₃ is made by heating NH₄Cl + Ca(OH)₂ and is dried over CaO — never over CaCl₂ or concentrated H₂SO₄ (both react with NH₃). Nitrates are detected by the brown ring test (FeSO₄ + conc. H₂SO₄). High-yield pointers: distinguish N₂O, NO, NO₂; remember the products of heating Group I vs Group II nitrates.

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Industrial Manufacture of Ammonia (Haber Process)

Ammonia is synthesised by the reversible reaction N₂ + 3H₂ ⇌ 2NH₃, ΔH = –92 kJ mol⁻¹. The optimum industrial conditions are 200 atm, 500 °C, iron catalyst with Mo/K₂O promoter. The forward reaction is favoured by high pressure and low temperature, but a compromise temperature of 500 °C is used to give an acceptable rate. A typical conversion is about 15% per pass; unreacted gases are recycled. The Haber process is a classic UACE question testing Le Chatelier’s principle — students are often asked to justify the chosen compromise conditions.

Manufacture of Nitric Acid (Ostwald Process)

Ammonia is oxidised in a Pt/Rh gauze catalyst at 900 °C to give NO, which is cooled and further oxidised to NO₂. The NO₂ is absorbed in warm water: 3NO₂ + H₂O → 2HNO₃ + NO. The NO is recycled, giving overall conversion of NH₃ to ~60% HNO₃.

Laboratory Preparation of Ammonia

NH₃ is generated by warming a mixture of solid NH₄Cl and Ca(OH)₂: 2NH₄Cl + Ca(OH)₂ → CaCl₂ + 2NH₃ + 2H₂O. The gas is dried by passing through quicklime (CaO). CaCl₂ and conc. H₂SO₄ both react with NH₃ (forming CaCl₂·8NH₃ and (NH₄)₂SO₄ respectively) and cannot be used as drying agents.

Thermal Decomposition of Nitrogen Compounds

Compound	On Heating	Products
NH₄NO₃	N₂O + 2H₂O	(moderately explosive)
NH₄NO₂	N₂ + 2H₂O
NH₄Cl	Sublimes	NH₃ + HCl (recombine on cooling)
Group I nitrates (e.g. NaNO₃)	→ MNO₂ + O₂	nitrite + oxygen
Group II nitrates (e.g. Cu(NO₃)₂)	→ MO + 2NO₂ + ½O₂	oxide + NO₂ + O₂

Brown Ring Test for Nitrates

A freshly prepared FeSO₄ solution is mixed with the sample and concentrated H₂SO₄ is carefully added down the side of the tube. A brown ring of [Fe(H₂O)₅NO]²⁺ forms at the junction. This is a positive test for NO₃⁻, not NO₂⁻.

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The Nitrogen Cycle

The cycle maintains the supply of usable nitrogen in the biosphere and is a recurring structured question in Paper 2:

Fixation – conversion of atmospheric N₂ to NH₃/NH₄⁺. Natural routes include lightning (≈5–8% of natural fixation) and symbiotic Rhizobium in legume root nodules; industrial route is the Haber process.
Nitrification – two-step bacterial oxidation: Nitrosomonas converts NH₄⁺ → NO₂⁻, then Nitrobacter oxidises NO₂⁻ → NO₃⁻. Both are aerobic and inhibited in waterlogged soils.
Assimilation – plants absorb NO₃⁻ and reduce it back to NH₃ for amino acid synthesis.
Denitrification – anaerobic Pseudomonas and Thiobacillus reduce NO₃⁻ → N₂/N₂O, returning nitrogen to the atmosphere.

Complex Formation and Acid–Base Character

NH₃ donates its lone pair, behaving as a Lewis base:

[Cu(H₂O)₄]²⁺ + 4NH₃ → [Cu(NH₃)₄]²⁺ + 4H₂O (deep blue, used as a confirmatory test for Cu²⁺)
AgCl + 2NH₃ → [Ag(NH₃)₂]Cl (dissolves AgCl in Tollens’ reagent)

NO₂ is an acidic oxide: 2NO₂ + H₂O → HNO₃ + HNO₂. With NaOH, NO₂ disproportionates: 2NO₂ + 2NaOH → NaNO₃ + NaNO₂ + H₂O — examiners often exploit this “disproportionation” in structured questions.

Common Mistakes

Drying NH₃ over CaCl₂ or H₂SO₄ — both react with NH₃.
Confusing nitrogen(I) oxide N₂O (colourless, sweet, “laughing gas”), NO (paramagnetic, colourless) and NO₂ (brown, toxic).
Writing the Haber equation with a single arrow instead of ⇌ and forgetting the iron catalyst.
Confusing the Ostwald process (NH₃ → HNO₃) with the Haber process (N₂ + H₂ → NH₃).
Mixing up Hofmann degradation (NH₃ → N₂ with Br₂/NaOH) with simple oxidation of ammonia.

Practice Prompts

Explain why an iron catalyst and 500 °C are used in the Haber process despite higher yields being predicted at lower temperatures.
Outline the nitrogen cycle, naming the bacteria responsible for nitrification and denitrification, and state two ways in which human activity disturbs the natural balance.

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Industrial Manufacture of Ammonia (Haber Process)