Nitrogen gas and nitrogen fixation
Nitrogen is in group 15 (VA) with five outer electrons. Two N atoms share three electron pairs: N≡N. This triple bond is one of the strongest bonds known, so N₂ is colourless, odourless and very unreactive.
Living things need nitrogen for proteins and DNA, but most cannot use N₂ gas. Nitrogen fixation turns N₂ into compounds such as NH₃ or nitrates.
- Lightning: N₂ + O₂ → 2NO, which ends up as nitrates in soil.
- Biological: bacteria such as Rhizobium in the root nodules of peas, beans and pulses.
- Industrial: the Haber process.
Ammonia, the Haber process and ammonium salts
N₂ + 3H₂ ⇌ 2NH₃ (exothermic, the forward reaction gives out heat).
- Pressure about 200 atm: 4 gas molecules become 2, so high pressure pushes the equilibrium to the right.
- Temperature about 450 °C: a lower temperature would give more NH₃ but too slowly, so this is a compromise.
- Catalyst: iron, to speed it up. Unreacted gases are recycled.
Properties of ammonia
Colourless gas with a sharp smell, lighter than air, very soluble in water (fountain experiment). NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, so the solution is a weak base and turns red litmus blue. NH₃ + HCl → NH₄Cl (white smoke). Ammonia can also act as a reducing agent: 4NH₃ + 5O₂ → 4NO + 6H₂O (with a platinum catalyst).
Ammonium salts
Examples: NH₄Cl, (NH₄)₂SO₄, NH₄NO₃. All are soluble. Test: warm with NaOH. Ammonia gas comes off and turns damp red litmus blue: NH₄Cl + NaOH → NaCl + NH₃↑ + H₂O. On heating, NH₄Cl breaks into NH₃ and HCl, which join again on the cool part of the tube.
Nitrogen oxides, nitric acid and acid rain
| NO (nitrogen monoxide) | NO₂ (nitrogen dioxide) |
|---|---|
| Colourless, insoluble in water | Red-brown, choking, dissolves in water |
| N₂ + O₂ → 2NO (lightning, engines) | 2NO + O₂ → 2NO₂ (in air, fast) |
3NO₂ + H₂O → 2HNO₃ + NO. Industrially (Ostwald process): NH₃ → NO → NO₂ → HNO₃.
Nitric acid, HNO₃
- A strong acid: reacts with bases, carbonates and metal oxides to form nitrates.
- A strong oxidising agent: with copper, dilute HNO₃ gives NO and concentrated HNO₃ gives brown NO₂, not hydrogen.
- Concentrated HNO₃ turns yellow on standing because light breaks it into NO₂; store it in brown bottles.
- Iron and aluminium become passive in cold concentrated HNO₃ (a thin oxide layer protects them).
Nitrates (NO₃⁻) are used as fertilisers and in explosives; most decompose on heating. Nitrites (NO₂⁻) preserve meat but are toxic in excess.
Acid rain
Normal rain has a pH of about 5.6 because of dissolved CO₂. NO₂ (and SO₂ from coal) make nitric and sulfuric acid, so pH can fall to 4 or lower. Effects: damaged forests and lakes, corroded metal, worn-away marble and limestone (CaCO₃ + 2HNO₃ → Ca(NO₃)₂ + H₂O + CO₂). Remedies: catalytic converters in cars, cleaner fuels, scrubbers in power plants and adding lime to lakes.
Phosphorus and N and P fertilisers
Phosphorus is just below nitrogen in group 15. White phosphorus (P₄) is very reactive, glows in air and is stored under water; red phosphorus is safer and is used on matchboxes. Phosphorus burns to P₄O₁₀, which makes phosphoric acid, H₃PO₄, with water. Phosphates (PO₄³⁻) are in bones, teeth, DNA and ATP.
Plants need N for leaves (proteins) and P for roots and energy. Common fertilisers: urea CO(NH₂)₂ (46% N), ammonium nitrate, ammonium sulfate, and superphosphate or diammonium phosphate (DAP) for P. Too much fertiliser washes into lakes and causes algal blooms (eutrophication).
Try it at home
Grow two cups of moong seeds: one with plain water and one with a pinch of urea or plant food dissolved in water. After a week, compare leaf colour and height. Then hold a strip of red cabbage paper above a bowl of rainwater and of tap water and compare colours.
Key formulas and definitions
- N₂ + 3H₂ ⇌ 2NH₃ (Fe, ~450 °C, ~200 atm, exothermic)
- NH₃ + HCl → NH₄Cl; NH₄⁺ + OH⁻ → NH₃↑ + H₂O (test)
- N₂ + O₂ → 2NO; 2NO + O₂ → 2NO₂
- 3NO₂ + H₂O → 2HNO₃ + NO
- % N in urea CO(NH₂)₂ = 28 / 60 × 100 ≈ 46.7%
Worked examples
1. How many litres of NH₃ (at the same conditions) can be made from 30 L of H₂ with excess N₂?
N₂ + 3H₂ → 2NH₃. Gas volumes follow mole ratios: 3 L H₂ → 2 L NH₃. So 30 L H₂ → 20 L NH₃.
2. Why is high pressure used in the Haber process?
The left side has 4 moles of gas and the right side 2. High pressure favours the side with fewer gas molecules, so more NH₃ forms.
3. Calculate the percentage of nitrogen in ammonium nitrate, NH₄NO₃. (N = 14, H = 1, O = 16)
Molar mass = 14 + 4 + 14 + 48 = 80 g/mol. N = 28 g. % N = 28/80 × 100 = 35%.
4. A gas is given off when a white salt is warmed with NaOH and it turns damp red litmus blue. What is the salt type and the gas?
The salt is an ammonium salt and the gas is ammonia: NH₄⁺ + OH⁻ → NH₃ + H₂O.
Common mistakes
- Thinking NO is brown. NO is colourless; NO₂ is the brown gas.
- Saying metals + nitric acid give hydrogen. HNO₃ is an oxidising acid and gives NO or NO₂ instead.
- Saying 450 °C gives the highest yield. Lower temperature gives more yield but too slowly; 450 °C is a compromise.
- Calling all acidic rain 'acid rain'. Normal rain is already about pH 5.6; acid rain is below that.