Topic 2: Nitrogen, Sulphur & Phosphorus

Complete Study Guide - Group V & VI Elements

🌬️ NITROGEN (N)

Position: Group V, Period 2

Atomic Number: 7

Electron Configuration: 2.5

Valency: 3 or 5

Abundance: 78% of air

Symbol: N

Sources of Nitrogen

Air (atmosphere): Makes up 78% of the air we breathe

Earth's crust: Found in mineral deposits

Lightning: Converts atmospheric nitrogen into compounds

Biological sources: Proteins, amino acids, leguminous plants (groundnuts, pigeon peas)

Fertilizers: NH₄NO₃ (ammonium nitrate), (NH₄)₃PO₄ (ammonium phosphate), NaNO₃ (sodium nitrate)

Properties

Physical Properties:

Colorless gas
Odorless
Insoluble in water
Less dense than air

Chemical Properties:

Diatomic gas (N₂) with strong triple covalent bonds (N ≡ N)

Very unreactive (inert) under normal conditions due to strong bonds

Reacts at high temperatures with alkali metals, alkaline earth metals, and hydrogen

Important Chemical Reactions

With Alkali Metals

Potassium + Nitrogen → Potassium nitride

6K(s) + N₂(g) → 2K₃N(s)

With Alkaline Earth Metals

Magnesium + Nitrogen → Magnesium nitride

3Mg(s) + N₂(g) → Mg₃N₂(s)

With Hydrogen (Haber Process)

Nitrogen + Hydrogen ⇌ Ammonia

N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

⇌ means reversible reaction

Uses of Nitrogen

Ammonia production: Used in Haber process to make fertilizers

Food preservation: Prevents oxidation and spoilage

Fire prevention: Used in oil tankers due to inertness

Liquid nitrogen: Refrigerant and for shrink-fitting machine parts

Nitrogen Compounds

Ammonia (NH₃)

Preparation

Method: Heat ammonium salt with alkali

Calcium hydroxide + Ammonium chloride → Calcium chloride + Water + Ammonia

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

Note: Water vapour removed by passing through calcium oxide

Properties

Physical: Colorless gas with pungent (choking) smell

Solubility: Very soluble in water

Nature: Basic substance (turns red litmus blue)

Chemical Reactions

With hydrogen chloride: Forms white smoke

NH₃(g) + HCl(g) → NH₄Cl(s)

White smoke = tiny ammonium chloride particles

With acids: Forms salts

NH₃(g) + HNO₃(aq) → NH₄NO₃(aq)

Forms ammonium nitrate (fertilizer)

Uses of Ammonia

Nitric Acid

Manufacturing

Fertilizers

Agricultural use

Plastics

Industrial production

Water Softening

Hard water treatment

Explosives

Manufacturing

Dry Cells

Ammonium chloride

Nitric Acid (HNO₃)

Preparation

Method: Heat potassium nitrate with concentrated sulphuric acid

Potassium nitrate + Sulphuric acid → Potassium hydrogen sulphate + Nitric acid

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

Uses of Nitric Acid

Nitrate fertilizers manufacturing

Explosives (TNT, dynamite)

Plastics manufacturing

Metal purification (silver, gold, platinum)

Dyes and drugs manufacturing

Textile industry (oxidizing agent)

Gemstone refining

Preparation of Nitric Acid Using the Ostwald Process

The Ostwald Process is used to produce nitric acid (HNO₃) from ammonia (NH₃) through a series of steps.

Steps of the Ostwald Process

Starting Material:

• Ammonia (NH₃), usually produced by the Haber Process.

Step 1: Oxidation of Ammonia:

• Ammonia is oxidized to form nitrogen monoxide (NO).

Chemical Reaction:

4NH₃(g) + 5O₂(g) → 4NO(g) + 6H₂O(g)

• This reaction occurs at high temperatures (about 900°C) in the presence of a platinum or rhodium catalyst.

Step 2: Oxidation of Nitrogen Monoxide:

• Nitrogen monoxide (NO) is further oxidized to form nitrogen dioxide (NO₂).

Chemical Reaction:

2NO(g) + O₂(g) → 2NO₂(g)

Step 3: Formation of Nitric Acid:

• Nitrogen dioxide (NO₂) is absorbed in water to produce nitric acid.

Chemical Reaction:

3NO₂(g) + H₂O(l) → 2HNO₃(aq) + NO(g)

Result:

• The final product is nitric acid, which can be concentrated and used in various applications.

⚡ SULPHUR (S)

Position: Group VI, Period 3

Atomic Number: 16

Electron Configuration: 2.8.6

Molecule: S₈ (crown-shaped)

Allotropes: Rhombic & Monoclinic

Symbol: S

Sources of Sulphur

Volcanic regions: Natural deposits from volcanic activity

Crude oil: Found in petroleum deposits

Metal ores: Combined with various metals

Natural gas: As hydrogen sulphide (H₂S)

Properties

Physical Properties:

Brittle yellow solid
Crown-shaped S₈ molecules
Insoluble in water
Soluble in organic solvents
Low melting point
Non-conductor (heat & electricity)

Chemical Properties:

With metals: Forms sulphides

Mg(s) + S(s) → MgS(s)

With oxygen: Burns to form SO₂

S(s) + O₂(g) → SO₂(g)

Uses of Sulphur

Sulphuric Acid Production

Most important industrial use

Rubber Vulcanization

Strengthens rubber by adding sulphur

Manufacturing

Matches, pesticides, drugs, paper

Construction

Sulphur concrete

Explosives

Gunpowder manufacturing

Decorative

Plastic flowers

Allotropes of Sulphur

Rhombic Sulphur (α-sulphur): Most stable form at room temperature. Forms orthorhombic crystals with bright yellow color. Composed of S₈ ring molecules arranged in a specific crystal structure.

Crystal structure: Orthorhombic system with well-defined geometric shapes. Density: 2.07 g/cm³. Melting point: 112.8°C

Formation: Crystallizes slowly from solution at temperatures below 96°C. Most common natural form found in volcanic deposits.

Monoclinic Sulphur (β-sulphur): Stable at temperatures above 96°C. Forms needle-like crystals in the monoclinic crystal system. Also composed of S₈ molecules but arranged differently.

Crystal structure: Monoclinic system with prismatic, needle-like crystals. Density: 1.96 g/cm³. Melting point: 119.6°C

Formation: Forms when rhombic sulphur is heated above 96°C (transition temperature). Slowly converts back to rhombic form when cooled below 96°C.

Production of Sulphuric Acid (H₂SO₄)

Contact Process - Four Steps

Burning Sulphur in Oxygen

Sulphur is burnt in oxygen to produce sulphur dioxide

S(s) + O₂(g) → SO₂(g)

Oxidation of Sulphur Dioxide

Sulphur dioxide reacts with oxygen to produce sulphur trioxide

2SO₂(g) + O₂(g) → 2SO₃(g)

Formation of Oleum

Sulphur trioxide is mixed with concentrated sulphuric acid to produce oleum

SO₃(g) + H₂SO₄(l) → H₂S₂O₇(l)

Formation of Sulphuric Acid

Oleum is added to water to produce sulphuric acid

H₂S₂O₇(l) + H₂O(l) → 2H₂SO₄(l)

Uses of Sulphuric Acid

Inorganic fertilizers: Manufacturing ammonium sulphate

Car batteries: Used as battery acid

Synthetic fibres: Production of nylon

Petroleum refining: Industrial process

Paints and dyes: Manufacturing process

Soaps and detergents: Production

Dehydrating agent: Removes water from substances (e.g., sucrose dehydration)

🔥 PHOSPHORUS (P)

Position: Group V, Period 3

Atomic Number: 15

Electron Configuration: 2.8.5

Valency: 3 or 5

Allotropes: White & Red phosphorus

Symbol: P

Sources of Phosphorus

Agricultural sources: Composite farm manure

Earth's crust: Found in form of phosphates

Mineral rocks: Mainly phosphate minerals

Properties

Physical Properties:

Yellow solid at room temperature
Does not conduct heat or electricity
Two allotropes: white and red phosphorus
Insoluble in water
Melting point: 44°C
Boiling point: 280°C

Chemical Properties:

With oxygen: Forms oxides (P₂O₅, P₂O₃)

With halogens: Combines easily

With metals: Forms phosphides

Allotropes of Phosphorus

White Phosphorus (P₄): Soft, waxy solid that glows in the dark (phosphorescence). Highly reactive and poisonous. Stored under water to prevent reaction with air. Ignites spontaneously at 35°C.

Appearance: Colorless to pale yellow, translucent solid with a garlic-like odor

Uses: Military smoke screens, incendiary weapons, and rat poison (though banned in many countries due to toxicity)

Red Phosphorus: More stable allotrope formed by heating white phosphorus in absence of air. Non-poisonous and does not glow in the dark. Less reactive than white phosphorus.

Appearance: Dark red powder or crystalline solid, odorless

Uses: Safety matches (striking surface), fireworks, pesticides, and fertilizer production. Safer to handle than white phosphorus.

Uses of Phosphorus

Inorganic fertilizers: Manufacturing ammonium phosphate and other phosphate fertilizers

Phosphoric acid production: Important industrial chemical

Match manufacturing: Used in match heads for ignition

Consumer products: Toothpaste, detergents, and baking powder

Quick Comparison Table

Property	Nitrogen (N)	Sulphur (S)	Phosphorus (P)
Group & Period	Group V, Period 2	Group VI, Period 3	Group V, Period 3
Atomic Number	7	16	15
Electron Configuration	2.5	2.8.6	2.8.5
Physical State	Colorless gas	Yellow solid	Yellow solid
Reactivity	Very unreactive (inert)	Moderately reactive	Reactive
Main Industrial Use	Ammonia production	Sulphuric acid production	Fertilizer production
Valency	3 or 5	2, 4, or 6	3 or 5

Important Chemical Reactions Summary

Nitrogen Reactions

Haber Process:

N₂ + 3H₂ ⇌ 2NH₃

Ammonia Preparation:

Ca(OH)₂ + NH₄Cl → CaCl₂ + H₂O + NH₃

Nitric Acid Prep:

KNO₃ + H₂SO₄ → KHSO₄ + HNO₃

Sulphur Reactions

Contact Process Step 1:

S + O₂ → SO₂

Contact Process Step 2:

2SO₂ + O₂ → 2SO₃

With Metals:

Mg + S → MgS

Key Points to Remember

Nitrogen:

Triple bond (N≡N) makes it unreactive

Sulphur:

Forms S₈ crown-shaped molecules

Phosphorus:

White and red allotropes exist

📚 Study Tips & Exam Points

Remember These Key Points:

Nitrogen makes 78% of air and is very unreactive due to triple bonds
Ammonia is basic (turns red litmus blue) and very soluble in water
Contact process has 4 steps for making sulphuric acid
Sulphur exists as S₈ molecules and has two allotropes
Phosphorus is used mainly for fertilizers and matches

Common Exam Questions:

Preparation methods for ammonia and nitric acid
Properties and uses of these three elements
Chemical equations for important reactions
Contact process steps and equations
Sources of these elements in nature

📖 Complete Study Guide: Nitrogen, Sulphur & Phosphorus

Review this comprehensive guide for thorough understanding!