12.1 Nitrogen and sulfur
- Syllabus
- 9701–2028–2029
- Topic
- 12.1
- Level
- AS
Learning objectives
12.1.1Lack of reactivity of nitrogen, using triple• Explain the lack of reactivity of nitrogen, using triple bond strength and lack of polarity12.1.2And explain• Describe and explain:: (a) the basicity of ammonia, using the Brønsted–Lowry theory; (b) the structure of the ammonium ion and its formation by an acid–base reaction; (c) the displacement of ammonia from ammonium salts by an acid–base reaction12.1.3NOx sources and catalytic removal• State/explain: the natural and man-made occurrences of oxides of nitrogen and their catalytic removal from the exhaust gases of internal combustion engines12.1.4Atmospheric NOx and photochemical smog• Understand: atmospheric oxides of nitrogen (NO and NO 2) can react with unburned hydrocarbons to form peroxyacetyl nitrate, PAN, which is a component of photochemical smog12.1.5Role of NO and NO2 in the formation of acid• Describe the role of NO and NO2 in the formation of acid rain both directly and in their catalytic role in the oxidation of atmospheric sulfur dioxide
Nitrogen is relatively unreactive because N₂ has a strong non-polar triple bond
Nitrogen gas exists as non-polar N₂ molecules containing a very strong N≡N triple bond. Breaking this bond requires substantial energy, so many reactions have a high activation energy.
The absence of a permanent dipole also limits attraction to polar reagents. Reactivity can still be increased by high temperature, pressure, catalysts or an appropriate reaction partner; “unreactive” is not “inert in every condition”.
The Haber process makes ammonia from nitrogen and hydrogen only with elevated pressure, temperature and an iron catalyst. The catalyst changes rate, not the equilibrium position.
Do not explain nitrogen’s behaviour using only “it is a gas”, and do not confuse a strong N≡N bond with zero possible chemistry.
One proton-transfer model links ammonia, ammonium and ammonia displacement
A Brønsted–Lowry base accepts a proton. The lone pair on nitrogen lets ammonia accept H⁺: NH₃ + H⁺ → NH₄⁺. In water, NH₃ + H₂O ⇌ NH₄⁺ + OH⁻; the equilibrium lies mainly to the left, so ammonia is a weak base.
NH₄⁺ is tetrahedral with bond angles of about 109.5°. Its fourth N–H bond forms when nitrogen donates its lone pair to H⁺, so it begins as a coordinate bond; after formation all four N–H bonds are equivalent.
| Representation | Ammonium ion acts as | Hydroxide ion acts as |
|---|---|---|
| NH₄⁺ + OH⁻ → NH₃ + H₂O | proton donor / acid | proton acceptor / base |
| NH₄Cl + NaOH → NH₃ + H₂O + NaCl | ammonium salt source | stronger base that displaces NH₃ |
Warm an ammonium salt with aqueous alkali: released ammonia turns damp red litmus paper blue. Damp paper is needed because the gas must dissolve before its basic behaviour is shown.
Ammonia does not need an OH group to be basic; its lone pair accepts H⁺. A coordinate bond describes the origin of the shared pair, not a permanently different N–H bond in NH₄⁺.
High temperatures form NOₓ; catalytic converters return nitrogen
| Occurrence | Why NOₓ forms | Representative chemistry |
|---|---|---|
| lightning | very high temperature supplies energy for N₂ and O₂ to react | N₂ + O₂ → 2NO; then 2NO + O₂ → 2NO₂ |
| vehicle engines and high-temperature combustion | compressed burning fuel heats atmospheric N₂ and O₂ | NO and NO₂ leave with exhaust gases |
| power stations / industrial combustion | sustained high-temperature burning also forms thermal NOₓ | the same atmospheric N₂/O₂ source is involved |
On a hot precious-metal catalyst, nitrogen oxides are reduced to N₂ while carbon monoxide and unburned hydrocarbons are oxidised. A representative coupled equation is 2CO + 2NO → 2CO₂ + N₂.
The catalyst surface adsorbs reactants and provides a lower-activation-energy route. It speeds conversion but is regenerated overall; it does not prevent NOₓ forming inside the engine.
Do not describe lightning as man-made or the converter as filtering NOₓ physically. The removal is a surface-catalysed redox reaction that changes NOₓ into nitrogen.
Sunlight turns primary NOₓ and hydrocarbons into PAN smog
| Component | Role in the pathway |
|---|---|
| NO and NO₂ | primary nitrogen-oxide pollutants from combustion |
| unburned hydrocarbons and their oxidised products | volatile organic reactants from incomplete fuel combustion |
| sunlight | supplies energy for the photochemical reaction sequence |
| peroxyacetyl nitrate, PAN (CH₃CO₃NO₂) | secondary pollutant formed in air; a component of photochemical smog |
NO₂ absorbs light and initiates radical chemistry that oxidises hydrocarbons. Products of that oxidation combine with NO₂ to form PAN, so PAN is not emitted directly from the engine: it forms later in the atmosphere.
Photochemical smog requires NOₓ, hydrocarbon-derived species and sunlight. Do not equate PAN with NO₂ itself, ordinary smoke, sulfurous smog or acid rain.
NOₓ forms nitric acid directly and catalyses sulfuric-acid formation
NO is first oxidised in air: 2NO + O₂ → 2NO₂. In cloud water, further oxidation of NO₂ produces nitric acid: 4NO₂ + 2H₂O + O₂ → 4HNO₃. Dilute HNO₃ lowers rainwater pH directly.
| Step | Equation | Role of nitrogen oxide |
|---|---|---|
| oxidise sulfur dioxide | NO₂ + SO₂ → SO₃ + NO | NO₂ is reduced to NO while SO₂ becomes SO₃ |
| regenerate catalyst | NO + ½O₂ → NO₂ | atmospheric O₂ restores NO₂ |
| form acid | SO₃ + H₂O → H₂SO₄ | sulfuric acid enters cloud/rain water |
SO2+21O2→SO3
Adding the first two equations cancels NO and NO₂, giving the net oxidation of SO₂. Their regeneration is why the pair acts catalytically and can accelerate formation of more H₂SO₄ without net consumption.
Keep the two roles distinct: direct NOₓ chemistry forms HNO₃, whereas the catalytic cycle helps convert atmospheric SO₂ into H₂SO₄. Nitrogen oxides are not the only cause of acid rain.