Nitrogen and Sulfur: Reactivity and Atmospheric Chemistry

This guide covers subtopic 12.1, Nitrogen and sulfur, from Topic 12 of Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. This is AS Level content.

Before studying this

This resource assumes the Brønsted-Lowry theory from Acids and Bases: The Brønsted-Lowry Theory — ammonia’s basicity is a direct application of it — and coordinate (dative covalent) bonding from Chemical Bonding: Shapes and Intermolecular Forces, which already covers the ammonium ion’s formation in detail; this resource reuses that rather than re-deriving it.

Syllabus coverage

CAMBRIDGE INTERNATIONAL AS & A LEVEL CHEMISTRY 9701 — AS Level, Topic 12

12.1 Nitrogen and sulfur — explaining the lack of reactivity of nitrogen, with reference to triple bond strength and lack of polarity; describing and explaining the basicity of ammonia using the Brønsted-Lowry theory, the structure of the ammonium ion and its formation by an acid-base reaction, and the displacement of ammonia from ammonium salts by an acid-base reaction; stating and explaining the natural and man-made occurrence of oxides of nitrogen and their catalytic removal from the exhaust gases of internal combustion engines; understanding that atmospheric NO and NO₂ can react with unburned hydrocarbons to form peroxyacetyl nitrate (PAN), a component of photochemical smog; describing the role of NO and NO₂ in the formation of acid rain, both directly and by catalysing the oxidation of atmospheric sulfur dioxide.

Why nitrogen is so unreactive

Nitrogen gas, N₂, is strikingly unreactive despite being roughly 78% of the atmosphere. Two factors explain this together:

• The N≡N triple bond is exceptionally strong (around 945 kJ mol⁻¹, among the strongest bonds known) — breaking it to start a reaction requires a large amount of energy.
• The molecule is non-polar — two identical nitrogen atoms share the bonding electrons perfectly evenly, so N₂ has no permanent dipole to attract an approaching reactive species.

Both factors point the same way: a strong, non-polar bond gives nitrogen a very high activation energy for almost any reaction, which is why it takes extreme conditions (lightning, an internal combustion engine, or an industrial catalyst in the Haber process) to make it react at all.

The basicity of ammonia

Ammonia’s nitrogen atom carries a lone pair, which makes NH₃ a Brønsted-Lowry base — a proton acceptor:

NH₃ + H⁺ → NH₄⁺

The lone pair forms a new N–H bond by donating both electrons to the incoming H⁺ — a coordinate (dative covalent) bond — giving the ammonium ion, NH₄⁺, in which all four N–H bonds are identical once formed, even though one was made differently from the other three. This acid-base reaction happens directly when ammonia meets an acid, for example:

NH₃ + HCl → NH₄Cl

Displacing ammonia from an ammonium salt runs the same acid-base logic in reverse: heating an ammonium salt with a stronger base than ammonia releases ammonia gas.

NH₄Cl + NaOH → NaCl + NH₃ + H₂O

This is a standard test for an ammonium salt: warm gently with aqueous sodium hydroxide, and ammonia’s characteristic smell (and its effect of turning damp red litmus paper blue) confirms the presence of NH₄⁺.

Oxides of nitrogen

Natural formation: the enormous energy of a lightning strike is enough to overcome nitrogen’s high activation energy, driving N₂ + O₂ → 2NO directly in the atmosphere.

Man-made formation: the same reaction happens inside an internal combustion engine, where the high temperature and pressure of combustion provide the activation energy that ordinary atmospheric conditions can’t. Some of the NO formed is further oxidised to NO₂ by atmospheric oxygen.

Catalytic removal: a catalytic converter reduces NOx back to harmless N₂ using carbon monoxide (itself a combustion product) as the reducing agent, over a platinum/rhodium/palladium catalyst:

2CO + 2NO → 2CO₂ + N₂

NOx and photochemical smog

Atmospheric NO and NO₂ react with unburned hydrocarbons from vehicle exhaust — driven by sunlight — to form peroxyacetyl nitrate (PAN), a key component of photochemical smog and an eye and lung irritant characteristic of sunny, traffic-heavy cities.

NOx and acid rain

NO and NO₂ contribute to acid rain in two distinct ways, and it’s worth keeping them separate:

• Directly: nitrogen oxides dissolve in atmospheric water and are oxidised further, forming nitric acid, which falls as acid rain.
• Indirectly, as a catalyst: NO₂ oxidises atmospheric sulfur dioxide to sulfur trioxide, NO₂ + SO₂ → NO + SO₃, regenerating NO₂ from the NO produced by further reaction with O₂ — a catalytic cycle. The SO₃ produced then dissolves in water to form sulfuric acid, SO₃ + H₂O → H₂SO₄, which falls as acid rain. NO₂ isn’t consumed overall in this pathway; it speeds up sulfur dioxide’s contribution to acid rain without being used up itself.

Common mistakes

• Explaining nitrogen’s unreactivity with only “the triple bond is strong.” Both factors — bond strength and lack of polarity — are expected in a full answer; each addresses a different reason a reactive species would struggle to attack N₂.
• Describing ammonia’s basicity without naming the Brønsted-Lowry theory. “Ammonia is basic because it’s an alkali” doesn’t answer what’s being asked — the expected explanation is specifically proton acceptance via the nitrogen lone pair.
• Confusing the direct and indirect (catalytic) routes to acid rain. Nitrogen oxides forming nitric acid directly, and nitrogen dioxide catalysing sulfur dioxide’s oxidation to sulfur trioxide, are two separate mechanisms — a full answer distinguishes them.
• Forgetting that the catalytic converter equation is itself a redox reaction. Carbon is oxidised (CO → CO₂) while nitrogen is reduced (NO → N₂) — it’s worth being able to identify the oxidation-number changes if asked.

Quick revision checklist

• Nitrogen’s unreactivity: strong triple bond + non-polar molecule
• Ammonia as a Brønsted-Lowry base; formation and structure of NH₄⁺ via a coordinate bond
• Displacing ammonia from an ammonium salt with a stronger base
• Natural (lightning) vs man-made (combustion engines) formation of NOx; catalytic converter equation
• NOx and photochemical smog (PAN)
• NOx and acid rain: direct route vs catalysing SO₂ oxidation

Related resources

• Acids and Bases: The Brønsted-Lowry Theory — the theory ammonia’s basicity is built on
• Chemical Bonding: Shapes and Intermolecular Forces — the ammonium ion’s coordinate bond in full
• Redox Processes: Oxidation Numbers and Electron Transfer — oxidation-number changes in the catalytic converter reaction
• Cambridge AS & A Level Chemistry hub

Written against Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. Always check the current syllabus for your examination year.