Nitrogen and sulfur: properties, preparation, reactions, oxides, acids, uses, environmental impact

Cambridge International AS & A Level Chemistry – Nitrogen & Sulfur

1. Physical‑Chemistry Foundations (AS)

Topic	Key Points for Nitrogen & Sulfur
Atomic structure & isotopes	Nitrogen (N): Z = 7; ¹⁴N (99.6 %), ¹⁵N (0.4 %). Sulfur (S): Z = 16; major isotopes ³²S (95 %), ³³S, ³⁴S, ³⁶S. Electronic configurations: N [He] 2s²2p³, S [Ne] 3s²3p⁴.
Mole concept & stoichiometry	Molar masses: N = 14.01 g mol⁻¹, S = 32.07 g mol⁻¹. Use balanced equations for preparation, combustion and redox (e.g. N₂ + 3 H₂ → 2 NH₃; S + O₂ → SO₂).
Bonding & structure	N₂: triple covalent bond, DB ≈ 941 kJ mol⁻¹ → kinetic inertness (high bond energy + non‑polar molecule). S₈: puckered crown‑shaped ring; solid is a covalent network. Hybridisation: N in NH₃ – sp³ (tetrahedral NH₄⁺), S in SO₃ – sp² (trigonal planar). Ionic examples: NH₄⁺Cl⁻, Na₂S, (NH₄)₂SO₄.
Energetics (ΔH°, Hess’s law, bond‑energy cycles)	Standard enthalpies of formation (kJ mol⁻¹): N₂ = 0, NO = +90.3, NO₂ = +33.2, N₂O₅ = +11.3; S = 0, SO₂ = –296.8, SO₃ = –395.7. These values allow calculation of combustion, oxidation and acid‑formation enthalpies.
Redox & electrochemistry	Standard electrode potentials (V): N₂/NH₃ = –0.06, NO₃⁻/NO₂⁻ = +0.80, SO₄²⁻/S = +0.14. Electrolysis examples: molten NaNO₃ → N₂, molten Na₂SO₄ → Na₂S + O₂.
Equilibria (Kc, Kp, pH) & Le Chatelier	NO₂ ⇌ N₂O₄ Kp ≈ 0.15 atm⁻¹ (298 K). SO₂ + ½ O₂ ⇌ SO₃ Kc ≈ 4 × 10⁻³ (500 K). Acid‑base: HNO₃ (strong), HNO₂ (pKa ≈ 3.15), H₂SO₄ (pKa1 ≈ –3, pKa2 ≈ 1.99), NH₃ (Kb = 1.8 × 10⁻⁵, pKb ≈ 4.74, pKa(NH₄⁺) ≈ 9.25).
Kinetics & catalysis	Haber‑Bosch: N₂ + 3 H₂ → 2 NH₃ (Fe‑based catalyst, 400–500 °C, 150–300 atm). Contact process: 2 SO₂ + O₂ → 2 SO₃ (V₂O₅ catalyst, 450 °C). Three‑way catalytic converter (automotive): 2 NO + 2 CO + 4 CₓHᵧ → 2 N₂ + 2 CO₂ + 2 H₂O. Selective catalytic reduction (SCR): 4 NH₃ + 4 NO + O₂ → 4 N₂ + 6 H₂O (Cu‑zeolite catalyst).

2. Inorganic Chemistry (AS & A Level)

2.1 Periodic Trends (relevant to N & S)

• Atomic radius: decreases across a period, increases down a group – N < O < F; S > Cl > Ar.
• Ionisation energy: highest for N (half‑filled 2p) and for O (high electronegativity).
• Electronegativity (Pauling): N = 3.04, O = 3.44, F = 3.98; S = 2.58, Cl = 3.16.
• Oxidation‑state range: Group 15 (‑3 to +5) – N prefers –3, +3, +5; Group 16 (‑2 to +6) – S prefers –2, +4, +6.

2.2 Nitrogen (N)

2.2.1 Properties

Property	Value
Atomic number	7
Atomic mass (u)	14.01
Standard state	Colourless diatomic gas (N₂)
Melting point (K)	63.15
Boiling point (K)	77.36
Density (g cm⁻³, 0 °C, 1 atm)	0.0012506
Electronegativity (Pauling)	3.04

2.2.2 Why N₂ is inert

The N≡N triple bond has a very high bond‑dissociation energy (≈ 941 kJ mol⁻¹) and the molecule is non‑polar (no permanent dipole moment). Consequently, a large amount of energy is required to break the bond, giving N₂ kinetic inertness under normal conditions.

2.2.3 Basicity of Ammonia and the NH₄⁺ ion

• Brønsted–Lowry reaction: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
• K_b(NH₃) = 1.8 × 10⁻⁵ → pK_b ≈ 4.74
• pK_a(NH₄⁺) ≈ 9.25 (conjugate acid).
• Structure of NH₄⁺: tetrahedral (sp³), N–H bond length ≈ 1.04 Å, strong hydrogen‑bonding in salts and aqueous solution.

2.2.4 Industrial Preparation of Nitrogen

1. Fractional distillation of liquid air – exploits N₂’s lower boiling point (77 K) compared with O₂ (90 K).
2. Thermal decomposition of ammonium nitrate (lab):

   NH₄NO₃  →[250 °C]  N₂O + 2 H₂O

   Followed by catalytic reduction:

   N₂O + H₂  →  N₂ + H₂O

3. Electrolysis of molten sodium nitrate (demonstration):

   2 NaNO₃  →[electrolysis]  2 NaNO₂ + N₂↑

2.2.5 Key Reactions of Nitrogen

1. Combustion (formation of nitrogen oxides)

   N₂ + O₂ →[Δ] 2 NO

   2 NO + O₂ → 2 NO₂

2. Haber‑Bosch (ammonia synthesis)

   N₂ + 3 H₂ →[Fe catalyst, 450 °C, 200 atm] 2 NH₃

3. Metal nitrides (high‑temperature)

   3 Mg + N₂ →[Δ] Mg₃N₂

   2 Al + N₂ →[Δ] AlN

4. Reaction with halogens – nitrogen trihalides

   N₂ + 3 Cl₂ →[Δ] 2 NCl₃

   (Analogous: NBr₃, NF₃ – NF₃ is stable and used as a fluorinating agent.)

2.2.6 Oxides of Nitrogen

Oxide	Formula	Oxidation state of N	Key properties / reactions
Nitric oxide	NO	+2	Paramagnetic gas; oxidises to NO₂; weak acid‑oxide.
Nitrogen dioxide	NO₂	+4	Brown toxic gas; dimerises to colourless N₂O₄ (< –11 °C); acidic oxide – reacts with bases to give nitrite.
Dinitrogen tetroxide	N₂O₄	+4	Colourless liquid; equilibrium with NO₂; used as rocket oxidiser.
Dinitrogen pentoxide	N₂O₅	+5	Anhydride of HNO₃; decomposes on heating → 2 NO₂ + ½ O₂.
Dinitrogen trioxide	N₂O₃	+3	Deep‑blue liquid (‑30 °C); hydrolyses to nitrous acid (HNO₂).
Nitrous oxide	N₂O	+1	Colourless, mildly anaesthetic; produced from thermal decomposition of NH₄NO₃.

2.2.7 Acids Derived from Nitrogen Oxides

• Nitric acid (HNO₃) – from N₂O₅:

  N₂O₅ + H₂O → 2 HNO₃

  (Strong acid, pK_a ≈ –1.4)
• Nitrous acid (HNO₂) – from NO₂:

  2 NO₂ + H₂O → HNO₃ + HNO₂

  (Weak acid, pK_a ≈ 3.15)
• Ammonium nitrate (NH₄NO₃) – neutralisation of HNO₃ with NH₃; key fertiliser.

2.2.8 Amphoteric Behaviour of Nitrogen Oxides

NO₂ (and N₂O₃) act as acidic oxides, forming acids with water, but they also react with strong bases:

2 NO₂ + 2 NaOH → NaNO₂ + NaNO₃ + H₂O

This dual character is required by the syllabus under “acid‑base properties of nitrogen oxides”.

2.2.9 Catalytic Removal of NOₓ (Environmental Chemistry)

• Three‑way catalyst (automotive) – simultaneously:
  • 2 NO → N₂ + O₂ (reduction)
  • 2 CO + O₂ → 2 CO₂ (oxidation)
  • 2 CₓHᵧ + (2x + y/2) O₂ → 2x CO₂ + y H₂O (oxidation)
  Catalyst composition: Pt/Pd/Rh on Al₂O₃.
• Selective catalytic reduction (SCR) using NH₃:

  4 NH₃ + 4 NO + O₂ → 4 N₂ + 6 H₂O

  Industrially employed in power‑plant flue‑gas treatment.

2.2.10 Major Uses of Nitrogen

• Fertiliser production – Haber‑Bosch → NH₃ → NH₄NO₃, urea, ammonium sulphate.
• Ostwald process – NH₃ → NO → NO₂ → HNO₃ (industrial nitric acid).
• Inert atmosphere – high‑purity N₂ for electronics, food packaging, fire‑suppression.
• Explosives – nitration of aromatics (TNT, nitroglycerine) and ammonium‑nitrate based mixtures.
• Medical & laboratory gases – N₂ as carrier gas; N₂O as anaesthetic.

2.2.11 Environmental Impact of Nitrogen Compounds

• NOₓ (NO + NO₂) are precursors to photochemical smog and acid rain. In the atmosphere:

2 NO₂ + H₂O → HNO₃ + HNO₂

• NOₓ also participate in tropospheric ozone formation:

NO + O₃ → NO₂ + O₂

• Mitigation strategies: catalytic converters, low‑NOₓ burners, SCR, and flue‑gas desulphurisation (combined with SO₂ control).

2.3 Sulfur (S)

2.3.1 Properties

Property	Value
Atomic number	16
Atomic mass (u)	32.07
Standard state	Yellow orthorhombic solid (S₈)
Melting point (K)	388.36
Boiling point (K)	717.8
Density (g cm⁻³, 25 °C)	2.07
Electronegativity (Pauling)	2.58

2.3.2 Industrial Preparation of Elemental Sulfur

1. Claus process (recovery from H₂S in refinery gas)

   2 H₂S + SO₂ →[catalyst] 3 S + 2 H₂O

   (First step: partial combustion of H₂S → SO₂; second step: catalytic conversion.)
2. Thermal decomposition of metal sulfates (laboratory)

   CuSO₄ →[Δ] CuO + SO₃

   SO₃ + C → SO₂ + CO

   2 SO₂ →[Cu catalyst, Δ] 2 S + 2 O₂

3. Direct recovery from volcanic gases or native deposits – mined as “native sulphur”.

2.3.3 Key Reactions of Sulfur

1. Combustion

   S + O₂ →[Δ] SO₂

2. Oxidation to sulphur trioxide (contact process)

   2 SO₂ + O₂ →[V₂O₅, 450 °C] 2 SO₃

3. Formation of sulphuric acid (industrial)

   SO₃ + H₂O → H₂SO₄

   (H₂SO₄ is a strong diprotic acid, pK�a1 ≈ –3, pK_a2 ≈ 1.99.)
4. Reaction with metals – metal sulfides

   Fe + S → FeS

   Zn + S → ZnS

5. Reaction with halogens – sulphur halides

   S₈ + 8 Cl₂ → 8 SCl₂

   (SCl₂ is a useful chlorinating agent.)

2.3.4 Oxides of Sulfur

Oxide	Formula	Oxidation state of S	Key properties / reactions
Sulphur dioxide	SO₂	+4	Colourless gas; acidic oxide; dissolves in water to give H₂SO₃ (weak).
Sulphur trioxide	SO₃	+6	Colourless liquid; reacts violently with water to form H₂SO₄ (strong diprotic acid).
Disulphur dichloride	S₂Cl₂	+1	Liquid; used to produce mustard gas and as a chlorinating agent.
Sulphur hexafluoride	SF₆	+6	Colourless, inert gas; powerful greenhouse gas.

2.3.5 Acids Derived from Sulphur Oxides

• Sulphurous acid (H₂SO₃) – formed by dissolution of SO₂:

  SO₂ + H₂O ⇌ H₂SO₃

  (Weak acid, Ka₁ ≈ 1.5 × 10⁻².)
• Sulphuric acid (H₂SO₄) – from SO₃:

  SO₃ + H₂O → H₂SO₄

  (Strong acid, fully dissociates for the first proton; second pKₐ ≈ 1.99.)

2.3.6 Amphoteric Behaviour of Sulphur Oxides

SO₂ is an acidic oxide (forms H₂SO₃) but can act as a base toward very strong acids, e.g.:

SO₃ + HF → HSO₃F

(Formation of fluorosulphonic acid demonstrates basic character of the oxide.)

2.3.7 Catalytic Removal of SO₂ (Environmental Chemistry)

• Flue‑gas desulphurisation (FGD) – wet scrubbers using limestone:

  SO₂ + CaCO₃ + H₂O → CaSO₃·½H₂O + CO₂

  Followed by oxidation to gypsum (CaSO₄·2H₂O).
• Contact process optimisation – V₂O₅ catalyst maximises SO₃ yield, reducing SO₂ emissions.

2.3.8 Major Uses of Sulphur

• Production of sulphuric acid – “the world’s most important industrial chemical”.
• Vulcanisation of rubber (addition of cross‑links via sulphur).
• Manufacture of fertilizers (ammonium sulphate, super‑phosphate).
• Manufacture of detergents, batteries, and pigments.
• SF₆ as an electrical insulator in high‑voltage equipment.

2.3.9 Environmental Impact of Sulphur Compounds

• SO₂ emitted from combustion of fossil fuels contributes to acid rain:

SO₂ + H₂O → H₂SO₃

2 H₂SO₃ + O₂ → 2 H₂SO₄

• Acid rain damages ecosystems, corrodes buildings, and acidifies lakes.

• Mitigation: flue‑gas desulphurisation, low‑sulphur fuels, catalytic converters for SO₂, and strict emission limits.