Ammonia (NH3)

Chapter in a nutshell: Ammonia (NH₃) is a colourless, pungent-smelling, alkaline gas — the hydride of nitrogen — that is lighter than air and extremely soluble in water. It is prepared in the laboratory by warming an ammonium salt with caustic alkali (or by the action of warm water on a metallic nitride), and manufactured industrially by Haber's process; its alkaline aqueous solution (ammonium hydroxide), its reducing power and its many salts make it the cornerstone of fertilizers, nitric acid and refrigeration.

1. Introduction, Occurrence and Molecular Structure

Ammonia is a compound of nitrogen and hydrogen, formula NH₃, relative molecular mass 17. First prepared by Joseph Priestley (1774) by heating ammonium chloride with slaked lime; its elementary nature was established by Berthollet (1785).

Occurrence: Free state — small amounts in air and traces in natural water. Combined state — as ammonium salts (ammonium chloride/Nausadar, ammonium sulphate). Being highly soluble, ammonia and its compounds do not occur as minerals. The pungent smell near decaying matter is from ammonia produced by bacterial decomposition of urea:

$$\mathrm{CO(NH_2)_2 + 2H_2O \rightarrow 2NH_3 + H_2O + CO_2}$$

Structure and bonding: A polar covalent compound. Nitrogen forms three N–H covalent bonds and retains one lone pair, giving a trigonal pyramidal shape. The lone pair causes ammonia's basic (proton-accepting) nature, its polarity, and its ability to form co-ordinate bonds (NH₄⁺, complex ions).

Forms of ammonia:

2. Laboratory Preparation of Ammonia

2.1 General methods

$$\mathrm{(NH_4)_2SO_4 \xrightarrow{\Delta} 2NH_3\uparrow + H_2SO_4}$$ $$\mathrm{(NH_4)_3PO_4 \xrightarrow{\Delta} 3NH_3\uparrow + H_3PO_4}$$

2.2 Standard lab preparation — from ammonium chloride

$$\mathrm{2NH_4Cl + Ca(OH)_2 \xrightarrow{\Delta} CaCl_2 + 2H_2O + 2NH_3\uparrow}$$

Procedure: Reactants are finely ground and taken in a round-bottom flask tilted with its mouth downwards, then heated. The downward tilt stops water condensing in the cooler part from trickling back and cracking the hot flask.

Why: Ca(OH)₂ is cheap and non-deliquescent (unlike NaOH/KOH); the 2 : 3 ratio offsets loss of NH₄Cl by sublimation.

Drying: Through a tower of quicklime (CaO). Conc. H₂SO₄, P₂O₅ and anhydrous CaCl₂ are NOT used because basic ammonia reacts with them:

$$\mathrm{2NH_3 + H_2SO_4 \rightarrow (NH_4)_2SO_4}$$ $$\mathrm{6NH_3 + P_2O_5 + 3H_2O \rightarrow 2(NH_4)_3PO_4}$$ $$\mathrm{CaCl_2 + 8NH_3 \rightarrow CaCl_2 \cdot 8NH_3}$$ (addition compound)

Collection: In an inverted gas jar by downward displacement of air because NH₃ is lighter than air (V.D. 8.5; air 14.4) and highly soluble (cannot be collected over water).

Why NH₄NO₃ / NH₄NO₂ are NOT used: On heating they give nitrous oxide or nitrogen (not ammonia), and NH₄NO₃ is explosive:

$$\mathrm{NH_4NO_3 \xrightarrow{\Delta} N_2O + 2H_2O}$$ $$\mathrm{NH_4NO_2 \xrightarrow{\Delta} N_2 + 2H_2O}$$

2.3 Preparation from metallic nitrides

$$\mathrm{Mg_3N_2 + 6H_2O \xrightarrow{warm} 3Mg(OH)_2 + 2NH_3\uparrow}$$ $$\mathrm{AlN + 3H_2O \xrightarrow{warm} Al(OH)_3 + NH_3\uparrow}$$ $$\mathrm{Ca_3N_2 + 6H_2O \xrightarrow{warm} 3Ca(OH)_2 + 2NH_3\uparrow}$$

The metal hydroxide stays in the flask while NH₃ escapes, collected mouth-downwards and dried over quicklime.

3. Manufacture of Ammonia — Haber's Process

Manufactured by the direct combination of nitrogen and hydrogen:

$$\mathrm{N_2 + 3H_2 \rightleftharpoons 2NH_3 + heat}$$

The reaction is reversible, exothermic and proceeds with a decrease in volume (4 vol → 2 vol).

Sources: Nitrogen — from fractional distillation of liquid air; Hydrogen — from water gas (Bosch process) or natural gas; mixed in the ratio N₂ : H₂ = 1 : 3 by volume.

Favourable conditions:

Recovery: Only ~15 % converts per pass. Ammonia is separated by liquefaction (NH₃ has a much higher boiling point than N₂/H₂) or by dissolving in water (NH₃ very soluble; N₂/H₂ nearly insoluble). Unreacted N₂ and H₂ are recirculated (yield ~98 %). No external heat is needed after start-up. Catalytic poisons (CO₂, CO, H₂S) must be removed first.

4. Physical Properties of Ammonia

5. Chemical Properties of Ammonia

5.1 Thermal dissociation

5.2 Basic nature

$$\mathrm{NH_3 + H_2O \rightleftharpoons NH_4OH \rightleftharpoons NH_4^+ + OH^-}$$

Action on indicators:

5.3 Reaction with acids (salt formation)

$$\mathrm{NH_3 + HCl \rightarrow NH_4Cl}$$ (dense white fumes) $$\mathrm{NH_3 + HNO_3 \rightarrow NH_4NO_3}$$ $$\mathrm{2NH_3 + H_2SO_4 \rightarrow (NH_4)_2SO_4}$$

The same salts form via NH₄OH: $\mathrm{NH_4OH + HCl \rightarrow NH_4Cl + H_2O}$; $\mathrm{2NH_4OH + H_2SO_4 \rightarrow (NH_4)_2SO_4 + 2H_2O}$.

5.4 Reaction with soluble metal salts (precipitation of hydroxides)

In excess NH₄OH the zinc and copper precipitates dissolve forming soluble complex salts: $$\mathrm{Zn(OH)_2 + 4NH_4OH \rightarrow [Zn(NH_3)_4](OH)_2 + 4H_2O}$$ $$\mathrm{Cu(OH)_2 + (NH_4)_2SO_4 + 2NH_4OH \rightarrow [Cu(NH_3)_4]SO_4 + 4H_2O}$$ (azure-blue tetraamminecopper(II) sulphate)

5.5 Combustibility / Reaction with oxygen

Catalytic oxidation: With platinum at 800 °C, ammonia is oxidised to nitric oxide (basis of nitric-acid manufacture): $$\mathrm{4NH_3 + 5O_2 \xrightarrow[800^\circ C]{Pt} 4NO + 6H_2O + heat}$$ The NO is then oxidised in air to brown nitrogen dioxide: $\mathrm{2NO + O_2 \rightarrow 2NO_2}$. The platinum glows on after heating stops (exothermic).

5.6 Reducing nature of ammonia

5.7 Reaction with chlorine

(a) Ammonia in excess: chlorine is reduced and the HCl formed reacts with extra ammonia: $$\mathrm{8NH_3 + 3Cl_2 \rightarrow N_2 + 6NH_4Cl}$$ Observation: greenish-yellow chlorine disappears and white fumes of ammonium chloride form.

(b) Chlorine in excess: the products are hydrogen chloride and the yellow, highly explosive liquid nitrogen trichloride: $$\mathrm{NH_3 + 3Cl_2 \rightarrow NCl_3 + 3HCl}$$ These reactions show that chlorine has a strong affinity for hydrogen and that ammonia is a reducing agent.

5.8 Reaction with carbon dioxide (manufacture of urea)

6. Tests for Ammonia Gas and Ammonium Ion

1. Sharp pungent odour.
2. Turns moist red litmus blue, turmeric paper brown, phenolphthalein pink.
3. Dense white fumes with conc. HCl: $\mathrm{NH_3 + HCl \rightarrow NH_4Cl}$.
4. With CuSO₄ solution → blue ppt. dissolving to deep azure blue in excess.
5. All ammonium salts liberate ammonia when warmed with alkali (NaOH).
6. Ammonium salts give a brown colour/precipitate with Nessler's reagent (K₂HgI₄).

7. Uses of Ammonia and the Ostwald Process

Uses: Refrigerant in ice plants (high latent heat of vaporisation, 5700 cal/mole at −33 °C; ozone-friendly but not compatible with copper); laboratory reagent in qualitative analysis; cleansing agent (emulsifies grease); manufacture of nitrogenous fertilizers (ammonium sulphate, urea), explosives (ammonium nitrate), nitric acid (Ostwald), sodium carbonate (Solvay), nylon, dyes, plastics; source of hydrogen.

Manufacture of nitric acid — Ostwald process (1914): Dry, dust-free air : ammonia = 10 : 1 by volume passed over a platinum gauze catalyst at ~800 °C: $$\mathrm{4NH_3 + 5O_2 \xrightarrow[800^\circ C]{Pt} 4NO + 6H_2O + heat}$$ The nitric oxide is cooled and oxidised at ~50 °C in the oxidising chamber: $$\mathrm{2NO + O_2 \xrightarrow{50^\circ C} 2NO_2}$$ The NO₂ with more air is absorbed in warm water in the absorption tower: $$\mathrm{4NO_2 + 2H_2O + O_2 \rightarrow 4HNO_3}$$ This gives ~50 % acid, concentrated by distillation to 68 % (lab grade); fuming nitric acid (~98 %) is obtained by distilling 68 % HNO₃ over conc. H₂SO₄.

Worked Examples

1 — Mass of NH₃ from NH₄Cl. From 53.5 g NH₄Cl: 2NH₄Cl + Ca(OH)₂ → CaCl₂ + 2H₂O + 2NH₃; 107 g NH₄Cl → 34 g NH₃, so 53.5 g (1 mol) → 1 mol = 17 g NH₃.

2 — Volume at STP. 1 mol NH₃ from Example 1 = 22.4 L (22 400 cm³) at STP.

3 — Haber stoichiometry. N₂ + 3H₂ → 2NH₃. By Gay-Lussac, 50 L N₂ needs 150 L H₂ and gives 100 L NH₃.

4 — % N in urea. NH₂CONH₂, molar mass = 60; N = 28; % N = 28/60 × 100 = 46.66 %.

5 — Cu from CuO reduction. 3CuO + 2NH₃ → 3Cu + N₂ + 3H₂O; 6.8 g NH₃ = 0.2 mol → Cu = (3/2)×0.2 = 0.3 mol = 19.05 g Cu.

6 — Solubility. 1 vol water dissolves 702 vol NH₃, so 250 cm³ water dissolves 250×702 = 175 500 cm³ ≈ 175.5 L NH₃.

7 — Identify a gas that turns moist red litmus blue, fumes white with conc. HCl, and turns CuSO₄ deep blue: ammonia (NH₃).

Key Terms — Quick Glossary

Common Mistakes to Avoid

• Never dry NH₃ with conc. H₂SO₄, P₂O₅ or CaCl₂ — basic NH₃ reacts with all. Use quicklime (CaO).
• Collect by downward displacement of air, NOT over water — lighter than air but too soluble.
• NH₄NO₃ / NH₄NO₂ are NOT used to make ammonia — they give N₂O / N₂; NH₄NO₃ is explosive.
• Don't confuse the chlorine reactions: NH₃ excess → N₂ + NH₄Cl (white fumes); Cl₂ excess → NCl₃ + HCl (yellow explosive liquid).
• Dry ammonia is neutral (covalent) — only aqueous NH₄OH changes indicators.
• Balance metal-hydroxide equations to valency: FeCl₃ needs 3 NH₄OH, FeSO₄ needs 2.
• Haber ratio N₂ : H₂ = 1 : 3; Ostwald ratio NH₃ : air = 1 : 10.

Likely Exam Questions (with crisp answers)

1. State the molecular formula and relative molecular mass of ammonia. NH₃; RMM = 17.
2. Write the balanced equation for the lab preparation of ammonia. 2NH₄Cl + Ca(OH)₂ → CaCl₂ + 2H₂O + 2NH₃↑.
3. Why is ammonia collected by downward displacement of air? Because it is lighter than air and highly soluble in water (cannot be collected over water).
4. Name the drying agent for ammonia and explain why others fail. Quicklime (CaO); conc. H₂SO₄, P₂O₅, CaCl₂ react with basic ammonia.
5. Why is Ca(OH)₂ preferred to NaOH in the lab preparation? It is cheap and non-deliquescent.
6. State the conditions of Haber's process. 450–500 °C, 200–1000 atm, finely divided iron catalyst, molybdenum promoter; N₂ : H₂ = 1 : 3.
7. Why is high pressure used in Haber's process? Forward reaction reduces volume (4→2 vol), so high pressure increases NH₃ yield.
8. What is the fountain experiment? What does it prove? Water injected into a flask of NH₃ creates a vacuum, drawing red litmus up as a blue fountain — proving NH₃'s extreme solubility and basic nature.
9. What happens when ammonia is passed over heated CuO? Black CuO is reduced to reddish-brown copper: 3CuO + 2NH₃ → 3Cu + N₂ + 3H₂O — shows NH₃ is a reducing agent.
10. Products of NH₃ with (a) excess NH₃ and (b) excess Cl₂? (a) 8NH₃ + 3Cl₂ → N₂ + 6NH₄Cl; (b) NH₃ + 3Cl₂ → NCl₃ + 3HCl.
11. Write the catalytic oxidation of ammonia. 4NH₃ + 5O₂ →(Pt, 800°C) 4NO + 6H₂O — the first step of the Ostwald process.
12. Which fertilizer forms from NH₃ + CO₂? Conditions? Urea, NH₂CONH₂; at 150 °C and 150 atm.
13. Distinguish liquid ammonia from liquor ammonia fortis. Liquid ammonia = liquefied NH₃ (molecules only, strong refrigerant); liquor ammonia fortis = saturated aqueous solution (molecules + ions, weak refrigerant).
14. Obtain ammonia from a metal nitride. Mg₃N₂ + 6H₂O → 3Mg(OH)₂ + 2NH₃↑.
15. Three tests for ammonia. Pungent smell; white fumes with conc. HCl; moist red litmus → blue.
16. Why is liquid ammonia a good refrigerant? High latent heat of vaporisation; liquefies easily.
17. Colour of ppt. with FeCl₃, FeSO₄, CuSO₄. Reddish brown Fe(OH)₃; dirty green Fe(OH)₂; pale blue Cu(OH)₂.

Chemical Equations & Formulas (quick reference)

Formulas: NH₃ (RMM 17); NH₄OH; NH₄Cl; (NH₄)₂SO₄; NH₄NO₃; NH₂CONH₂ (urea); NCl₃; Mg₃N₂; AlN; Ca(OH)₂; CaO; [Cu(NH₃)₄]SO₄; [Zn(NH₃)₄](OH)₂; Nessler's reagent K₂HgI₄.

Decomposition of urea (bacterial): $\mathrm{CO(NH_2)_2 + 2H_2O \rightarrow 2NH_3 + H_2O + CO_2}$

Lab preparation: $\mathrm{2NH_4Cl + Ca(OH)_2 \xrightarrow{\Delta} CaCl_2 + 2H_2O + 2NH_3\uparrow}$ $\mathrm{(NH_4)_2SO_4 + 2NaOH \rightarrow Na_2SO_4 + 2H_2O + 2NH_3\uparrow}$ $\mathrm{(NH_4)_2SO_4 + 2KOH \rightarrow K_2SO_4 + 2H_2O + 2NH_3\uparrow}$

From nitrides: $\mathrm{Mg_3N_2 + 6H_2O \rightarrow 3Mg(OH)_2 + 2NH_3\uparrow}$ $\mathrm{AlN + 3H_2O \rightarrow Al(OH)_3 + NH_3\uparrow}$ $\mathrm{Ca_3N_2 + 6H_2O \rightarrow 3Ca(OH)_2 + 2NH_3\uparrow}$

Salts NOT giving ammonia: $\mathrm{NH_4NO_3 \xrightarrow{\Delta} N_2O + 2H_2O}$ $\mathrm{NH_4NO_2 \xrightarrow{\Delta} N_2 + 2H_2O}$

Why drying agents fail: $\mathrm{2NH_3 + H_2SO_4 \rightarrow (NH_4)_2SO_4}$ $\mathrm{6NH_3 + P_2O_5 + 3H_2O \rightarrow 2(NH_4)_3PO_4}$ $\mathrm{CaCl_2 + 8NH_3 \rightarrow CaCl_2 \cdot 8NH_3}$

Haber's process: $\mathrm{N_2 + 3H_2 \rightleftharpoons 2NH_3 + heat}$

Basic nature / dissolution: $\mathrm{NH_3 + H_2O \rightleftharpoons NH_4OH \rightleftharpoons NH_4^+ + OH^-}$

Thermal dissociation: $\mathrm{2NH_3 \rightleftharpoons N_2 + 3H_2}$

With acids: $\mathrm{NH_3 + HCl \rightarrow NH_4Cl}$ $\mathrm{NH_3 + HNO_3 \rightarrow NH_4NO_3}$ $\mathrm{2NH_3 + H_2SO_4 \rightarrow (NH_4)_2SO_4}$

With soluble metal salts: $\mathrm{FeCl_3 + 3NH_4OH \rightarrow Fe(OH)_3\downarrow + 3NH_4Cl}$ $\mathrm{FeSO_4 + 2NH_4OH \rightarrow Fe(OH)_2\downarrow + (NH_4)_2SO_4}$ $\mathrm{Pb(NO_3)_2 + 2NH_4OH \rightarrow Pb(OH)_2\downarrow + 2NH_4NO_3}$ $\mathrm{Zn(NO_3)_2 + 2NH_4OH \rightarrow Zn(OH)_2\downarrow + 2NH_4NO_3}$ $\mathrm{CuSO_4 + 2NH_4OH \rightarrow Cu(OH)_2\downarrow + (NH_4)_2SO_4}$ $\mathrm{Cu(OH)_2 + (NH_4)_2SO_4 + 2NH_4OH \rightarrow [Cu(NH_3)_4]SO_4 + 4H_2O}$ $\mathrm{Zn(OH)_2 + 4NH_4OH \rightarrow [Zn(NH_3)_4](OH)_2 + 4H_2O}$

Combustion / catalytic oxidation: $\mathrm{4NH_3 + 3O_2 \rightarrow 2N_2 + 6H_2O}$ $\mathrm{4NH_3 + 5O_2 \xrightarrow[800^\circ C]{Pt} 4NO + 6H_2O}$ $\mathrm{2NO + O_2 \rightarrow 2NO_2}$

Reducing nature: $\mathrm{3CuO + 2NH_3 \rightarrow 3Cu + N_2 + 3H_2O}$ $\mathrm{3PbO + 2NH_3 \rightarrow 3Pb + N_2 + 3H_2O}$

With chlorine: $\mathrm{8NH_3 + 3Cl_2 \rightarrow N_2 + 6NH_4Cl}$ (NH₃ excess) $\mathrm{NH_3 + 3Cl_2 \rightarrow NCl_3 + 3HCl}$ (Cl₂ excess)

With CO₂ (urea): $\mathrm{2NH_3 + CO_2 \xrightarrow[150\,atm]{150^\circ C} NH_2CONH_2 + H_2O}$

Ostwald process (nitric acid): $\mathrm{4NH_3 + 5O_2 \xrightarrow[800^\circ C]{Pt} 4NO + 6H_2O}$ $\mathrm{2NO + O_2 \xrightarrow{50^\circ C} 2NO_2}$ $\mathrm{4NO_2 + 2H_2O + O_2 \rightarrow 4HNO_3}$

Identification: $\mathrm{NH_3 + HCl \rightarrow NH_4Cl}$ (dense white fumes)