Nitrogen and Its Compounds

1. Introduction to Nitrogen (N₂)

Nitrogen is a diatomic gas with formula N₂. It makes up about 78% of Earth's atmosphere by volume. Nitrogen is largely unreactive at room temperature because of the strong triple bond between the two nitrogen atoms (N≡N).

Important roles: component of proteins and nucleic acids (biological), raw material for fertilizers and explosives (industrial), and present in many inorganic compounds (ammonia, nitric acid, nitrates).

2. Laboratory Preparation of Ammonia (NH₃)

Equation & Principle

Ammonia is commonly prepared in the laboratory by heating a mixture of an ammonium salt and a strong alkali, for example:

NH₄Cl(s) + Ca(OH)₂(s) → CaCl₂(aq) + H₂O(l) + NH₃(g)

Or using sodium hydroxide:

NH₄Cl(s) + NaOH(aq) → NaCl(aq) + H₂O(l) + NH₃(g)

Apparatus & Procedure (brief)

• Heat the mixture gently in a flask with a delivery tube angled into a gas jar over water or into upside-down bottle (to collect NH₃ by upward delivery since NH₃ is lighter than air). Use a drying tube if needed.
• Collect gas carefully — ammonia dissolves in water to form ammonium hydroxide (alkaline solution).

Simple diagram (lab set-up)

NH₃(g) + HCl(g) → NH₄Cl(s)

3. Industrial Manufacture of Ammonia — Haber Process

Overall reaction

N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

Conditions used industrially

• High pressure: typically around 150–250 atm (modern plants may use ~150–200 atm).
• Moderately high temperature: around 400–500 °C (compromise between rate and equilibrium yield).
• Catalyst: iron catalyst with promoters (e.g., K₂O, Al₂O₃) is used to increase rate.
• Continuous removal of ammonia by liquefaction shifts equilibrium to the right (Le Chatelier's principle).

Simple schematic (Haber plant)

Factors affecting yield (explain)

1. Pressure: Increasing pressure favours the side with fewer gas molecules (2 NH₃ vs 4 reactant molecules). → High pressure increases equilibrium yield.
2. Temperature: The forward reaction is exothermic (releases heat). Lower temperature favours higher yield but slows rate. Industrial choice (~450 °C) balances acceptable rate and reasonable yield.
3. Catalyst: Does not change equilibrium position, but increases reaction rate so equilibrium is reached faster—essential for economical production.
4. Removal of product: Liquefying and removing NH₃ shifts equilibrium to the right (Le Chatelier), increasing overall yield.
5. Purity of reactants: Presence of impurities (e.g., water, oxygen) can poison the catalyst and reduce yield.

4. Properties and Uses of Ammonia (NH₃)

Properties

• Colourless gas with pungent smell; very soluble in water forming alkaline solution (ammonium hydroxide, NH₄OH).
• Less dense than air (so collect by upward delivery in lab).
• Turns damp red litmus blue (basic).
• Reaction with acid to form ammonium salts (e.g., NH₃ + HCl → NH₄Cl).

Uses

• Production of fertilizers (ammonium nitrate, urea).
• Used in manufacture of nitric acid (via oxidation) and other chemicals.
• Refrigerant in industrial systems (anhydrous ammonia).
• Laboratory reagent and cleaning agents (aqueous ammonia).

5. Nitric Acid (HNO₃) — Preparation and Uses

Industrial preparation (Ostwald process — overview)

1. Oxidation of ammonia to nitric oxide:

   4NH₃(g) + 5O₂(g) → 4NO(g) + 6H₂O(g)  (platinum-rhodium catalyst, ~800–900 °C)

2. Oxidation of NO to NO₂:

   2NO(g) + O₂(g) → 2NO₂(g)

3. Absorption in water to form nitric acid:

   3NO₂(g) + H₂O(l) → 2HNO₃(aq) + NO(g)

   (NO formed is recycled back into the process)

Laboratory preparation (brief)

Conc. nitric acid is usually prepared industrially. In school labs, small amounts may be obtained by careful distillation of nitrate salts with concentrated sulfuric acid (requires expert supervision).

Properties

• Strong oxidising acid; corrosive and fuming when concentrated.
• Reacts with metals and organic compounds; causes nitration of aromatic rings (advanced/topic beyond O-Level).

Uses

• Manufacture of fertilizers (ammonium nitrate), explosives (nitroglycerin, TNT — industrial relevance only), and dyes.
• Laboratory reagent for testing (e.g., to form nitrates) and in analytical chemistry.

6. Oxides of Nitrogen (Detailed)

Nitrogen forms several oxides with different properties. Important oxides are:

Nitrous oxide — N₂O (laughing gas)

• Colourless gas with slight sweet smell; used as an anaesthetic and analgesic in dentistry and surgery (medical use).
• Supports combustion weakly; not highly toxic at low concentrations but misuse is dangerous.
• Formation (example): thermal decomposition of ammonium nitrate produces N₂O and water.

Nitric oxide — NO

• Colourless gas produced by combustion and lightning; formed in Ostwald process as intermediate.
• Very reactive with oxygen to form nitrogen dioxide:

  2NO + O₂ → 2NO₂

• Soluble in water as it oxidises further; biologically, low concentrations act as signalling molecule (advanced topic).

Nitrogen dioxide — NO₂

• Brown, pungent, choking gas; toxic and a major air pollutant (component of "NOₓ").
• Dimerises to form N₂O₄ at low temperatures. Reacts with water to give nitric and nitrous acids:

  2NO₂ + H₂O → HNO₃ + HNO₂

• NO₂ contributes to acid rain and photochemical smog; produced by vehicle engines and industry.

Dinitrogen pentoxide — N₂O₅

• White solid; reacts with water to form nitric acid:

  N₂O₅ + H₂O → 2HNO₃

• Exists in equilibrium with NO₂ and O₂ at certain conditions; not commonly met at O-Level but useful to know as an anhydride of nitric acid.

Environmental & health notes

• NOₓ gases (NO and NO₂) are harmful to human health (respiratory irritants) and contribute to acid rain and ground-level ozone formation.
• Control of NOₓ emissions is a major environmental concern (vehicle catalysts, industrial scrubbers).