Ever wondered why your kitchen cleaner smells so sharp or why car batteries are full of liquid that can melt metal? The answer lies in three famous chemicals you’ll meet in ICSE Chemistry – ammonia, nitric acid and sulphuric acid.
💡 In Simple Words: Ammonia is a light, stinky gas that can neutralise acids. Nitric acid is a strong, corrosive liquid that makes nitrates, which are used in fertilizers. Sulphuric acid is a thick, oily liquid that loves to pull water out of everything – it’s the work‑horse of the lab and industry.
What is Ammonia (NH₃)?
Physical properties
Ammonia is a colour‑less gas with a pungent, irritating smell – think of the scent of a cleaning spray. It is lighter than air, so it rises quickly. In water it forms a weak base called ammonium hydroxide (NH₄OH), which means it can accept hydrogen ions (protons).
How is ammonia prepared?
The classic laboratory method is the Haber process (though on a smaller scale in school labs). In simple terms, nitrogen gas (N₂) from the air reacts with hydrogen gas (H₂) under high pressure and temperature, using an iron catalyst, to give ammonia:
\(N₂ + 3H₂ \xrightarrow{\text{Fe, 450°C}} 2NH₃\)
In the lab you’ll more often see the reaction of a metal nitride with water:
\(Ca₃N₂ + 6H₂O \rightarrow 3Ca(OH)₂ + 2NH₃\)
Common uses
- Household cleaners – the smell you notice is ammonia dissolving grime.
- Fertiliser production – it is converted into ammonium nitrate, a key nitrogen source for plants.
- Cooling packs – when ammonia evaporates it absorbs heat, making it useful in refrigeration.
What is Nitric Acid (HNO₃)?
Physical properties
Nitric acid is a colourless to pale yellow liquid with a sharp, acrid smell. It is a strong acid, meaning it completely donates its hydrogen ion to water, forming nitrate ions (NO₃⁻). It also acts as an oxidising agent – it can pull electrons away from other substances, similar to how rust forms when iron meets oxygen.
How is nitric acid prepared?
In the laboratory you make it by the oxidation of ammonia using a copper catalyst. The steps are:
1. \(4NH₃ + 5O₂ \xrightarrow{\text{Cu}} 4NO + 6H₂O\) (ammonia burns to give nitric oxide, NO)
2. \(2NO + O₂ \rightarrow 2NO₂\) (NO reacts with oxygen to give nitrogen dioxide, NO₂)
3. \(3NO₂ + H₂O \rightarrow 2HNO₃ + NO\) (NO₂ dissolves in water to give nitric acid and a little NO back).
Important reactions
- Reaction with metals: Most metals dissolve, producing metal nitrates and hydrogen gas – e.g., \(Cu + 4HNO₃ \rightarrow Cu(NO₃)₂ + 2NO₂ + 2H₂O\).
- Formation of explosives: Mixing nitric acid with sulphuric acid concentrates nitrate ions, which then react with organic compounds to make TNT, gun‑cotton, etc.
Common uses
- Manufacturing of fertilizers like ammonium nitrate.
- Etching and cleaning of metals in the electronics industry.
- Production of explosives – a classic chemistry‑class example.
What is Sulphuric Acid (H₂SO₄)?
Physical properties
Sulphuric acid is a thick, oily liquid that feels sticky. It is a strong acid and a powerful dehydrating agent – it loves to take water out of other substances, much like a sponge that sucks up moisture. Because of this, it can char sugar, turning it black and releasing carbon dioxide.
How is sulphuric acid prepared?
The industrial Contact process starts with sulphur (S) burning in air to give sulphur dioxide (SO₂). This is then oxidised to sulphur trioxide (SO₃) using a vanadium(V) oxide catalyst, and finally SO₃ is absorbed in water to make H₂SO₄:
\(S + O₂ \rightarrow SO₂\)
\(2SO₂ + O₂ \xrightarrow{\text{V₂O₅}} 2SO₃\)
\(SO₃ + H₂O \rightarrow H₂SO₄\)
Key reactions
- Dehydration of sugars: \(C₁₂H₂₂O₁₁ \xrightarrow{\text{conc. H₂SO₄}} 12C + 11H₂O\) – the sugar turns black, showing the acid’s water‑stealing power.
- Reaction with metals: Dilute H₂SO₄ reacts with active metals to give hydrogen gas, e.g., \(Zn + H₂SO₄ \rightarrow ZnSO₄ + H₂\).
- Formation of acid‑base salts: Mixing with bases gives sulphate salts, such as \(NaOH + H₂SO₄ \rightarrow NaHSO₄ + H₂O\).
Common uses
- Battery electrolyte – the acid in a car battery conducts electricity.
- Manufacturing of detergents, dyes, and synthetic fibres.
- Petroleum refining – it helps crack heavy oils into gasoline.
Quick Comparison of the Three Compounds
| Property | Ammonia (NH₃) | Nitric Acid (HNO₃) | Sulphuric Acid (H₂SO₄) |
|---|---|---|---|
| State at RT | Gas | Liquid | Liquid (viscous) |
| Acid / Base | Weak base (accepts H⁺) | Strong acid (donates H⁺) | Strong acid (donates H⁺) & strong dehydrating agent |
| Key industrial use | Fertiliser (ammonium nitrate) | Fertiliser & explosives | Battery electrolyte & petrochemical processing |
| Common lab preparation | Metal nitride + water or Haber process | Oxidation of ammonia (NO → NO₂ → HNO₃) | Contact process (S → SO₂ → SO₃ → H₂SO₄) |
| Safety note | Pungent, irritant – use fume hood | Corrosive, oxidising – wear gloves | Highly corrosive, can cause severe burns – handle with extreme care |
📝 Likely Exam Questions
- Write the balanced chemical equation for the laboratory preparation of nitric acid from ammonia.
Answer: 4NH₃ + 5O₂ → 4NO + 6H₂O; 2NO + O₂ → 2NO₂; 3NO₂ + H₂O → 2HNO₃ + NO. - Why is sulphuric acid called a dehydrating agent? Give one example.
Answer: It removes water from substances because it reacts vigorously with H₂O to form hydrated ions. Example: Concentrated H₂SO₄ turns sugar black by removing water, leaving carbon. - Compare the acidic or basic nature of ammonia, nitric acid and sulphuric acid.
Answer: Ammonia is a weak base (accepts H⁺). Nitric acid and sulphuric acid are strong acids (donate H⁺ completely in water). - List two important industrial uses of each of the three compounds.
Answer: Ammonia – manufacture of ammonium nitrate fertilizer, refrigerant. Nitric acid – production of ammonium nitrate fertilizer, etching of metals. Sulphuric acid – electrolyte in lead‑acid batteries, petroleum refining (alkylation). - Explain how the contact process produces sulphuric acid.
Answer: Sulphur is burnt to give SO₂, which is oxidised to SO₃ using V₂O₅ catalyst, then SO₃ is absorbed in water to form H₂SO₄.