Laboratory preparation of ammonia

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your exam.

Ammonia (NH₃) is a colourless, pungent gas with a lone pair on N that makes it a Lewis base, forming the tetrahedral ammonium ion NH₄⁺ by dative bonding. It is prepared in the lab by heating an ammonium salt with a strong alkali and dried over calcium oxide (CaO) — never CaCl₂, which forms adducts.

The Haber-Bosch equation is N₂ + 3H₂ ⇌ 2NH₃, ΔH = −92 kJ mol⁻¹. Nitric acid (HNO₃) is made by the Ostwald process: 4NH₃ + 5O₂ → 4NO + 6H₂O (Pt/Rh gauze, ~900 °C), followed by 3NO₂ + H₂O → 2HNO₃ + NO.

Test / Reaction	Observation	Equation (key species)
Cu + dilute HNO₃	Colourless gas (NO)	3Cu + 8HNO₃ → 3Cu(NO₃)₂ + 2NO + 4H₂O
Cu + conc. HNO₃	Brown gas (NO₂)	Cu + 4HNO₃ → Cu(NO₃)₂ + 2NO₂ + 2H₂O
Brown ring test	Brown ring at junction	NO₃⁻ + Fe²⁺ + H⁺ → [Fe(H₂O)₅(NO)]²⁺

Oxidation states run from −3 in NH₃ to +5 in HNO₃.

🟡 Standard — Regular Study (2d–2mo)

Standard content for students with a few days to months.

Laboratory preparation of ammonia

Ammonia is generated by heating a mixture of an ammonium salt (NH₄Cl) and a strong base (NaOH or Ca(OH)₂). The gas is passed upward (it is lighter than air) and collected by downward displacement of air. It must be dried over calcium oxide because CaO is basic and will not react with the basic NH₃ — unlike CaCl₂, which forms the adduct CaCl₂·8NH₃.

Trap: Examiner favourite — students write “CaCl₂” as the drying agent and lose a mark.

The Haber-Bosch process

The reversible reaction N₂(g) + 3H₂(g) ⇌ 2NH₃(g) is exothermic (ΔH = −92 kJ mol⁻¹) and proceeds with a decrease in moles of gas (4 → 2). Applying Le Chatelier’s principle: high pressure (~200 atm) and low temperature favour yield, but a low temperature slows the rate. The compromise is ~450 °C, 200 atm, iron catalyst with K₂O/Al₂O₃ promoters to maximise rate × yield.

Unreacted gases are recycled; the condensed NH₃ is stored as liquid under pressure.

Ammonia as a Lewis base

The lone pair on nitrogen accepts a proton from HCl or H₂O, forming NH₄⁺. Four equivalent N–H bonds form — three covalent and one dative (coordinate) bond from N → H⁺. NH₃ also acts as a ligand in complex ions: with Cu²⁺ it forms the deep-blue [Cu(NH₃)₄(H₂O)₂]²⁺ (used to identify Cu²⁺ after excess NH₃ is added).

Nitric acid — an oxidising acid

HNO₃ does not liberate H₂ with metals. Dilute acid gives NO (colourless, browns in air) and concentrated acid gives NO₂ (brown, toxic):

Dilute: 3Cu + 8HNO₃ → 3Cu(NO₃)₂ + 2NO + 4H₂O
Concentrated: Cu + 4HNO₃ → Cu(NO₃)₂ + 2NO₂ + 2H₂O

Nitrogen’s oxidation state changes from +5 in HNO₃ to +2 (NO) or +4 (NO₂), confirming HNO₃ acts as the oxidising agent.

Brown ring test

A nitrate is reduced by FeSO₄ in cold concentrated H₂SO₄. The brown ring at the acid-solution interface is the complex ion [Fe(H₂O)₅(NO)]²⁺, in which NO acts as a ligand and iron is formally Fe(I).

Exam patterns

UNEB UACE commonly asks: (i) explain Haber conditions using Le Chatelier with a sketch of the yield–temperature graph; (ii) state with reason why CaO is preferred to CaCl₂ as NH₃ desiccant; (iii) write equations for Cu + dilute and Cu + concentrated HNO₃; (iv) describe the brown ring test. Numeric items test % yield = (actual/theoretical) × 100 and gas-volume calculations using PV = nRT at given T and P (e.g. 5 mol H₂ at 25 °C, 1 atm).

🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for students on a longer study timeline.

Why the Haber compromise exists

A pure thermodynamic optimum (~200 °C, high pressure) is impractical because the rate is too slow to be economic, and plant walls would fail above ~350 atm. The Fe catalyst is poisoned by trace H₂S, so syngas (N₂ + H₂) is rigorously desulphurised. Molar concentrations at equilibrium can be computed from Kc data: Kc ≈ 1.45 × 10⁻⁵ mol² dm⁻⁶ at 500 °C, and the % conversion increases with pressure because n(gas) decreases.

Oxidation-state map of nitrogen

Nitrogen displays every oxidation state from −3 to +5 in this syllabus. Recognising these is essential for balancing redox equations using the electron-book method:

Species	NH₃	N₂	N₂O	NO	NO₂⁻	NO₂	NO₃⁻
Oxidation state	−3	0	+1	+2	+3	+4	+5

Exam trap: In peroxides (O₂²⁻) oxygen is −1, not −2 — misassigning it breaks redox balancing.

Extended mechanism notes

N₂O (laughing gas) has structure :N≡N–O: with formal charges N(−1) and O(0); it is linear, prepared by cautious heating of NH₄NO₃.
NO is a neutral odd-electron molecule (paramagnetic) that dimerises weakly in the solid.
NO₂ exists in equilibrium with its pale-yellow dimer N₂O₄; both are paramagnetic/odd-electron species and must be drawn with correct dot-and-cross bonding.

Brown ring test — details and limitations

The test requires fresh FeSO₄ (Fe²⁺ oxidises in air), cold concentrated H₂SO₄ to avoid thermal decomposition of the complex, and works for soluble nitrates only. Bromides and iodides interfere and should be removed first with Ag⁺ or by oxidation. The brown ion [Fe(H₂O)₅(NO)]²⁺ has Fe in the +1 oxidation state.

Connections and applications

The nitrogen cycle: fixation (Haber + biological N₂ → NH₃), nitrification (NH₃ → NO₂⁻ → NO₃⁻ by Nitrosomonas and Nitrobacter), denitrification (NO₃⁻ → N₂ by Pseudomonas).
NOₓ from combustion: high-temperature N₂ + O₂ → 2NO contributes to acid rain (→ HNO₃) and photochemical smog; catalytic converters reduce NOₓ back to N₂.
Cu(NH₃)₄²⁺ links this topic to transition-metal complex-ion chemistry under “ligands and coordination numbers”.

Worked micro-example

A 2.50 dm³ flask contains 0.600 mol H₂ at 450 °C and 200 atm. Using PV = nRT, calculate the maximum moles of NH₃ if stoichiometric N₂ is supplied.

V = 2.50 × 10⁻³ m³, T = 723 K, P = 200 × 101 325 = 2.027 × 10⁷ Pa.
Total n = PV/RT = (2.027 × 10⁷ × 2.5 × 10⁻³)/(8.314 × 723) ≈ 8.43 mol of gas.
Stoichiometric ratios: 3 H₂ + 1 N₂ → 2 NH₃. So 0.600 mol H₂ needs 0.200 mol N₂ and gives at most 0.400 mol NH₃, since N₂ is limiting when only 0.600 mol H₂ is present.

Common mistakes to eliminate

Forgetting the reversible arrow and state symbols in Haber’s equation.
Calling NH₃ “planar” instead of trigonal pyramidal (bond angle ~107°).
Treating NH₄⁺ as having four covalent bonds — one is dative.
Confusing oxidation numbers in NO₂ (N = +4, not +3) and NO (N = +2, not +3).
Writing HNO₃ as a typical Brønsted acid — it is an oxidising acid and rarely yields H₂.

Practice prompts

Sketch and explain how % yield of NH₃ varies with (a) pressure at fixed temperature and (b) temperature at fixed pressure, citing two Le Chatelier arguments each.
A mixture of Cu turnings reacts with 100 cm³ of 2.0 mol dm⁻³ HNO₃ (dilute). Calculate the moles and volume at RTP (24 dm³) of NO produced, and show that HNO₃ is the limiting reagent.

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Laboratory preparation of ammonia