Nitrogen and Its Compounds

Group V elements include Nitrogen, Phosphorus, Arsenic, Antimony, and Bismuth. These elements show more differences than similarities in their physical and chemical properties.

1. Nitrogen exists as a diatomic gas and does not exhibit allotropy. Phosphorus is a typical non-metal and exists in several allotropic forms. Arsenic and Antimony are metalloids, while Bismuth is a metal.
2. Nitrogen is a colorless gas. Phosphorus appears as white or red solids depending on its allotrope. Arsenic is a dull gray metallic solid, and Antimony and Bismuth are silvery-white solids.
3. Nitrogen is the most electronegative element in the group. It forms various compounds with both metals and non-metals.

Nitrogen is primarily found in the atmosphere, making up about 78% by volume. It also occurs in combined forms such as ammonia, urea, and proteins.

Electronic Configuration: 1s² 2s² 2p³

Laboratory Preparation

Nitrogen can be extracted from air by removing CO₂ using caustic soda and oxygen using heated copper turnings. However, this method produces impure nitrogen (containing ~1% rare gases).

Pure nitrogen is prepared chemically using the following reactions:

1. Thermal decomposition of ammonium dioxonitrate(III):
   • NaNO₂(aq) + NH₄Cl(aq) → NH₄NO₂(aq) + NaCl(aq)
   • NH₄NO₂(aq) → N₂(g) + 2H₂O(l)
2. (NH₄)₂Cr₂O₇(s) → Cr₂O₃(s) + 4H₂O(l) + N₂(g)
3. 2NH₃(g) + 3CuO(s) → 3Cu(s) + 3H₂O(g) + N₂(g)
4. N₂O(g) + Cu(s) → CuO(s) + N₂(g)

Industrial Production: Nitrogen is produced by the fractional distillation of liquid air.

Physical Properties

1. Colorless, odorless, and tasteless gas
2. Slightly soluble in water
3. Melting point: -210°C; Boiling point: -196°C
4. Pure nitrogen is lighter than air

Chemical Properties

1. Reacts with reactive metals to form nitrides:
   3Mg(s) + N₂(g) → Mg₃N₂(s)
2. Combines with non-metals to form ammonia and nitrogen oxides:
   N₂(g) + 3H₂(g) → 2NH₃(g)

Uses of Nitrogen

1. Liquid nitrogen serves as a coolant
2. Used industrially to produce ammonia
3. Preserves packaged food by preventing oxidation

The Nitrogen Cycle

The nitrogen cycle describes how nitrogen moves between the atmosphere, soil, and living organisms. It involves several steps:

1. Nitrogen Fixation: Atmospheric nitrogen is converted to compounds like nitrates and ammonium by bacteria or lightning.
2. Nitrification: Nitrifying bacteria convert ammonia to nitrates, which plants absorb to build proteins.
3. Denitrification: Bacteria convert nitrogen compounds back to nitrogen gas, returning it to the atmosphere.
4. Decay: Decomposers break down dead matter into ammonia and ammonium compounds.
5. Lightning Fixation: Lightning forms nitrogen oxides, which dissolve in rain to produce nitrates in the soil.

Oxides of Nitrogen

Nitrogen(I) Oxide (N₂O) - Laughing Gas

Preparation

1. KNO₃(s) + NH₄Cl(s) → KCl(s) + NH₄NO₃(s)
2. NH₄NO₃(s) → N₂O(g) + 2H₂O(g)

Physical Properties

1. Colorless gas with a faint, sweet odor
2. Fairly soluble in cold water
3. 1.5 times denser than air
4. Neutral to moist litmus paper

Chemical Properties

1. Decomposes when heated:
   2N₂O(g) → 2N₂(g) + O₂(g)
2. Supports combustion of hot substances:
   Mg(s) + N₂O(g) → MgO(s) + N₂(g)
3. Reduced by heated copper:
   Cu(s) + N₂O(g) → CuO(s) + N₂(g)

Test for N₂O

If a glowing splint rekindles in the gas and the gas has a sweet smell without producing brown fumes, it is likely nitrogen(I) oxide.

Use: As a mild anesthetic in minor surgeries.

Nitrogen(II) Oxide (NO)

Preparation

Reacting copper with dilute trioxonitrate(V) acid:

3Cu(s) + 8HNO₃(aq) → 3Cu(NO₃)₂(aq) + 2NO(g) + 4H₂O(l)

Physical Properties

1. Colorless and toxic
2. Slightly denser than air
3. Insoluble in water
4. Neutral to litmus

Chemical Properties

1. Combines with oxygen:
   2NO(g) + O₂(g) → 2NO₂(g)
2. Decomposes at high temperatures:
   2NO(g) → N₂(g) + O₂(g)
3. Reduced by hot metals:
   2Cu(s) + 2NO(g) → 2CuO(s) + N₂(g)
4. Reduces acidified KMnO₄:
   3MnO₄^- + 4H⁺ + 5NO → 3Mn²⁺ + 5NO₃^- + 2H₂O

Tests for NO

1. Reacts with air to form reddish-brown NO₂ fumes
2. Turns acidified FeSO₄ solution dark brown

Nitrogen(IV) Oxide (NO₂)

Preparation

By heating lead(II) trioxonitrate(V):

Pb(NO₃)₂(s) → 2PbO(s) + 4NO₂(g) + O₂(g)

The gas is cooled in a U-tube to liquefy NO₂ while oxygen escapes.

Physical Properties

1. Reddish-brown gas with a pungent smell
2. Acidic and turns damp blue litmus red
3. Liquefies at 21°C into a yellow liquid
4. Heavier than air

Chemical Properties

1. Exists in equilibrium with N₂O₄ at low temperatures:
   N₂O₄(g) ⇌ 2NO₂(g)
2. Decomposes on heating:
   2NO₂(g) → N₂(g) + 2O₂(g)
3. Reduced by copper:
   4Cu(s) + 2NO₂(g) → 4CuO(s) + N₂(g)
4. Reacts with water to form a mixture of acids:
   H₂O(l) + 2NO₂(g) → HNO₂(aq) + HNO₃(aq)
5. Reacts with alkalis to form nitrate and nitrite salts:
   2KOH(aq) + 2NO₂(g) → KNO₂(aq) + KNO₃(aq) + H₂O(l)

Ammonia

Ammonia is a hydride of nitrogen. It forms naturally during the decay of nitrogen-containing materials in the absence of air. Although it exists in trace amounts in the atmosphere, it is highly soluble in water and is washed into the soil by rain.

Laboratory Preparation of Ammonia

Ammonia is produced in the lab by heating calcium hydroxide (slaked lime) with ammonium chloride:

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

Ammonia is dried using calcium oxide (CaO). Since ammonia is alkaline, it cannot be dried with concentrated H₂SO₄ or fused CaCl₂, as these react with it.

Haber Process (Industrial Production of Ammonia)

In the Haber process, nitrogen and hydrogen gases are combined in a 1:3 ratio by volume. The reaction is reversible, so specific conditions are needed to maximize the yield of ammonia:

N₂(g) + 3H₂(g) ⇌ 2NH₃(g) + heat

Conditions:

1. Finely divided iron catalyst
2. Temperature of around 450°C
3. Pressure of about 200 atm

Physical Properties of Ammonia

1. Colorless gas with a pungent, choking smell
2. Only known alkaline gas
3. Toxic in large quantities; affects respiratory muscles
4. Approximately 1.7 times less dense than air
5. Melting point: -77.7°C, Boiling point: -33.4°C

Chemical Properties of Ammonia

1. Burns in oxygen to produce water vapor and nitrogen gas:
   4NH₃(g) + 3O₂(g) → 6H₂O(g) + 2N₂(g)
2. Acts as a reducing agent, reacting with:
   • Copper(II) oxide: 3CuO(s) + 2NH₃(g) → 3Cu(s) + 3H₂O(l) + N₂(g)
   • Chlorine gas: 3Cl₂(g) + 8NH₃(g) → 6NH₄Cl(s) + N₂(g)
3. Reacts with carbon(IV) oxide (CO₂) to form urea and water vapor

Tests for Ammonia

1. Litmus test: Damped red litmus paper turns blue in the presence of ammonia gas.
2. Hydrochloric acid test: A glass rod dipped in concentrated HCl produces white fumes when introduced to ammonia gas, forming ammonium chloride.

Uses of Ammonia

1. Used in manufacturing nitric acid and sodium carbonate (via the Solvay process)
2. Aqueous ammonia softens temporary hard water
3. Liquid ammonia serves as a refrigerant
4. Used in laundries to remove grease and oil stains

Trioxonitrate(V) Acid (HNO₃)

Laboratory Preparation

HNO₃ is a volatile acid and is prepared in the lab by displacing it from a nitrate salt using concentrated sulfuric acid (which is less volatile):

KNO₃(s) + H₂SO₄(aq) → KHSO₄(aq) + HNO₃(aq)

Industrial Preparation

HNO₃ is industrially produced by catalytic oxidation of ammonia in three main steps:

1. Ammonia reacts with excess air in the presence of a platinum-rhodium catalyst at 700°C:
   4NH₃(g) + 5O₂(g) → 4NO(g) + 6H₂O(g)
2. Nitrogen(II) oxide is cooled and reacts with excess air:
   2NO(g) + O₂(g) → 2NO₂(g)
3. Nitrogen(IV) oxide is dissolved in water with oxygen to yield nitric acid:
   4NO₂(g) + 2H₂O(l) + O₂(g) → 4HNO₃(aq)

Physical Properties

1. Colorless, fuming liquid with a sharp, choking smell
2. Turns yellow due to decomposition to NO₂, which redissolves in the acid
3. Boiling point: 86°C, Melting point: -47°C
4. Density: 1.52 g/cm³
5. Miscible with water, forming a constant boiling mixture at 121°C
6. Turns blue litmus red

Chemical Properties

1. Neutralizes bases to form nitrates and water:
   NaOH(aq) + HNO₃(aq) → NaNO₃(aq) + H₂O(l)
2. Reacts with carbonates to release carbon dioxide:
   CaCO₃(s) + 2HNO₃(aq) → Ca(NO₃)₂(aq) + H₂O(l) + CO₂(g)
3. Only reacts with Ca, Mg, or Mn to release hydrogen gas when very dilute
4. As an oxidizing agent, it reacts with:
   • Sulfur: S(s) + 6HNO₃(aq) → H₂SO₄(aq) + 2H₂O(l) + 6NO₂(g)
   • Copper, lead, mercury, and silver: Forms metal nitrates and nitrogen oxides (NO or NO₂ depending on concentration)
   • Hydrogen sulfide: H₂S(g) + 2HNO₃(aq) → S(s) + 2H₂O(l) + 2NO₂(g)
   • Iron(II) salts: 6Fe²⁺(aq) + 8H⁺(aq) + 2NO₃^-(aq) → 6Fe³⁺(aq) + 4H₂O(l) + 2NO(g)
5. Aluminum and iron form an oxide layer when exposed to concentrated HNO₃, making them passive and unreactive. Thus, containers made of these metals can store the acid safely.

Uses of Nitric Acid

1. Used as an acid, oxidizing agent, and nitrating agent in the lab
2. Used as a component in rocket fuels
3. Used in the production of nylon and Terylene
4. Used in making fertilizers, dyes, drugs, and explosives