Exercise Solutions Available
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Chapter Introduction
Hydrogen is the lightest element, the most abundant element in the universe, and the only element that can legitimately be placed in two different groups of the periodic table. CBSE Class 11 Chemistry Chapter 9 covers hydrogen's unique position, its isotopes, the preparation and properties of dihydrogen, the classification of hydrides, the chemistry of water and hydrogen peroxide, and the emerging role of hydrogen as a fuel.
This chapter connects periodic table concepts (Chapter 3) with inorganic chemistry and has practical relevance through water treatment and hydrogen energy.
Position of Hydrogen in the Periodic Table
Hydrogen has a dual nature — it resembles both alkali metals (Group 1) and halogens (Group 17), but fits neither group completely.
Resemblance with Alkali Metals (Group 1)
| Property | Hydrogen | Alkali Metals |
|---|---|---|
| Electronic configuration | 1s¹ | ns¹ |
| Formation of cation | H → H⁺ + e⁻ | Na → Na⁺ + e⁻ |
| Reaction with halogens | H₂ + Cl₂ → 2HCl | 2Na + Cl₂ → 2NaCl |
| Reaction with oxygen | 2H₂ + O₂ → 2H₂O | 4Na + O₂ → 2Na₂O |
| Formation of oxides | Forms oxide (H₂O) | Forms oxide (Na₂O) |
Resemblance with Halogens (Group 17)
| Property | Hydrogen | Halogens |
|---|---|---|
| Electronic configuration | 1s¹ (one electron short of noble gas) | ns²np⁵ (one electron short of noble gas) |
| Formation of anion | H + e⁻ → H⁻ | Cl + e⁻ → Cl⁻ |
| Diatomic molecule | H₂ | Cl₂, Br₂, F₂ |
| Reaction with metals | 2Na + H₂ → 2NaH | 2Na + Cl₂ → 2NaCl |
| High ionization enthalpy | 1312 kJ/mol | Similar range |
Why Hydrogen is Placed Separately
Hydrogen differs from alkali metals because it is a non-metal, a diatomic gas, and has much higher ionization enthalpy. It differs from halogens because it has low electron affinity and lacks the p-orbital characteristics of halogens. Due to this dual character, hydrogen is placed separately at the top of Group 1 with a distinct arrangement.
Isotopes of Hydrogen
| Isotope | Symbol | Protons | Neutrons | Mass (u) | Abundance |
|---|---|---|---|---|---|
| Protium | ¹H or H | 1 | 0 | 1.0078 | 99.985% |
| Deuterium | ²H or D | 1 | 1 | 2.0140 | 0.015% |
| Tritium | ³H or T | 1 | 2 | 3.0160 | Trace (radioactive) |
Key differences:
- D₂O (heavy water) has higher boiling point (101.4 °C) and density (1.107 g/mL) than H₂O.
- Reaction rates differ: reactions of H₂ are faster than D₂ (kinetic isotope effect).
- Tritium is radioactive (β-emitter, half-life 12.3 years), used as a tracer.
Preparation of Dihydrogen (H₂)
Laboratory Methods
| Method | Reaction | Notes |
|---|---|---|
| Dilute H₂SO₄ on zinc | Zn + H₂SO₄ → ZnSO₄ + H₂ | Most common lab method |
| Zinc + NaOH | Zn + 2NaOH → Na₂ZnO₂ + H₂ | Gives hydrogen;ZnO₂²⁻ is sodium zincate |
| Metal hydride + water | CaH₂ + 2H₂O → Ca(OH)₂ + 2H₂ | Used in portable H₂ generators |
Industrial Methods
| Method | Reaction | Notes |
|---|---|---|
| Electrolysis of water | 2H₂O → 2H₂ + O₂ | Pure H₂; expensive |
| Lane's process | 3Fe + 4H₂O → Fe₃O₄ + 4H₂ | Steam over iron at high temperature |
| Water-gas reaction | C + H₂O → CO + H₂ (water gas) | Mixture of CO and H₂ |
| Bosch process | CO + H₂ + H₂O → CO₂ + 2H₂ | Water-gas shift; CO₂ removed |
| Coal gasification | C(s) + H₂O(g) → CO(g) + H₂(g) | At high temperature |
| From natural gas | CH₄ + H₂O → CO + 3H₂ | Steam reforming of methane |
Water gas (CO + H₂) is also called synthesis gas or syngas.
Properties of Dihydrogen
Physical Properties
| Property | Value |
|---|---|
| Colour, odour | Colourless, odourless |
| State at room temperature | Gas |
| Density | Lightest known gas (0.09 g/L) |
| Boiling point | 20.3 K (−252.8 °C) |
| Melting point | 13.8 K (−259.4 °C) |
| Solubility | Very slightly soluble in water |
| Nature | Non-poisonous, non-supportive of combustion |
Chemical Properties
1. Reaction with halogens: H₂ + X₂ → 2HX (reactivity: F₂ > Cl₂ > Br₂ > I₂) 2. Reaction with oxygen: 2H₂ + O₂ → 2H₂O; ΔH = −286 kJ/mol (exothermic) 3. Reaction with nitrogen: N₂ + 3H₂ ⇌ 2NH₃ (Haber process, 400–500 °C, 200 atm, Fe catalyst) 4. Reaction with metals: 2Na + H₂ → 2NaH (forms ionic hydrides) 5. Reaction with metal oxides: CuO + H₂ → Cu + H₂O (reducing property) 6. Hydrogenation: Vegetable oil (unsaturated) + H₂ → Vanaspati ghee (saturated), Ni catalyst 7. Reaction with CO: CO + 2H₂ → CH₃OH (methanol, ZnO/Cr₂O₃ catalyst)
Hydrides — Classification and Comparison
Type 1: Ionic (Salt-Like) Hydrides
- Formed by alkali metals (Group 1) and alkaline earth metals (Group 2, except Be and Mg under normal conditions).
- Crystalline solids with high melting points.
- Conduct electricity in molten state (H⁻ ions migrate to anode).
- React vigorously with water to give H₂: NaH + H₂O → NaOH + H₂.
Examples: NaH, CaH₂, LiH
Type 2: Covalent (Molecular) Hydrides
- Formed by p-block elements.
- Low boiling points, volatile, molecular.
- Three sub-categories based on electron count:
| Category | Description | Examples |
|---|---|---|
| Electron-deficient | Fewer electrons than needed for Lewis structure | BH₃ (dimerizes to B₂H₆), AlH₃ |
| Electron-precise | Exactly enough electrons for Lewis structure | CH₄, SiH₄, C₂H₆ |
| Electron-rich | More electrons than needed (lone pairs) | NH₃, H₂O, HF |
Trend in boiling points: NH₃ < H₂S < HF < H₂O — explained by hydrogen bonding in H₂O and HF.
Type 3: Metallic (Interstitial) Hydrides
- Formed by d-block and f-block elements (Groups 3–5).
- Hydrogen atoms occupy interstitial sites in the metal lattice.
- Retain metallic properties (conductivity, malleability).
- Non-stoichiometric (composition varies): e.g., LaH₂.⁷⁶, TiH₁.₇₃.
- Palladium absorbs ~900 times its volume of hydrogen (used for H₂ storage).
| Property | Ionic Hydrides | Covalent Hydrides | Metallic Hydrides |
|---|---|---|---|
| Bonding | Ionic (H⁻) | Covalent (shared electrons) | Metallic/interstitial |
| State | Crystalline solid | Gas/liquid | Solid |
| Melting point | High | Low | Variable |
| Conductivity | Conducts when molten | Non-conductor | Metallic conductor |
| Reaction with water | Produces H₂ | Varies; some produce H⁺ | Generally inert |
| Examples | NaH, CaH₂ | CH₄, NH₃, H₂O | TiH₂, ZrH₂, PdHx |
Water (H₂O)
Structure of Water
- V-shaped (bent) molecule with bond angle 104.5° (slightly less than tetrahedral 109.5° due to lone pair repulsion).
- O—H bond length: 0.96 Å.
- Polar molecule (dipole moment = 1.84 D) due to electronegativity difference (O = 3.5, H = 2.1).
- Extensive hydrogen bonding gives water its unique properties.
Physical Properties of Water
| Property | Value | Reason |
|---|---|---|
| Boiling point | 100 °C (anomalous) | Hydrogen bonding |
| Freezing point | 0 °C | Hydrogen bonding |
| Density | Maximum at 4 °C (1 g/cm³) | Open structure of ice collapses on melting |
| Specific heat | 4.18 J/(g·K) (very high) | Hydrogen bonding |
| Surface tension | 72.8 mN/m at 20 °C (high) | Hydrogen bonding |
Anomalous properties: Water has unusually high boiling point, high specific heat, and high surface tension compared to H₂S, H₂Se, H₂Te — all due to hydrogen bonding.
Chemical Properties of Water
1. Amphoteric: Acts as both acid and base. H₂O + HCl → H₃O⁺ + Cl⁻ (base); H₂O + NH₃ → NH₄⁺ + OH⁻ (acid). 2. Hydrolysis: Reacts with many salts. Na₂CO₃ + H₂O → NaHCO₃ + NaOH (basic solution). 3. Reaction with metals: Active metals (Na, K, Ca) react with cold water; less active metals (Fe, Zn) react with steam. 4. Reaction with non-metals: 2F₂ + 2H₂O → 4HF + O₂; Cl₂ + H₂O → HCl + HOCl. 5. Hydrate formation: CuSO₄ + 5H₂O → CuSO₄·5H₂O (blue vitriol).
Hard Water vs Soft Water
| Property | Hard Water | Soft Water |
|---|---|---|
| Contains | Ca²⁺, Mg²⁺ ions (bicarbonates, chlorides, sulphates) | Na⁺, K⁺ ions (no Ca²⁺, Mg²⁺) |
| Lather with soap | No (forms scum) | Yes |
| Boiling point | Slightly higher | Normal |
| Types | Temporary (bicarbonates) and Permanent (chlorides, sulphates) | — |
Water Softening Methods
| Method | Removes | Reaction |
|---|---|---|
| Boiling | Temporary hardness (bicarbonates) | Ca(HCO₃)₂ → CaCO₃↓ + H₂O + CO₂ |
| Clark's method | Temporary hardness | Ca(HCO₃)₂ + Ca(OH)₂ → 2CaCO₃↓ + 2H₂O |
| Washing soda | Both types | CaCl₂ + Na₂CO₃ → CaCO₃↓ + 2NaCl |
| Ion exchange (permutit) | Both types | 2NaZ + Ca²⁺ → CaZ₂ + 2Na⁺ (Z = zeolite) |
| Calgon method | Both types | Na₆P₆O₁₈ + Ca²⁺ → Na₄[CaP₆O₁₈] + 2Na⁺ |
| Synthetic resins | Both types | Cation exchange (H⁺) and anion exchange (OH⁻) → demineralized water |
Hydrogen Peroxide (H₂O₂)
Structure
- Open book (skewed) structure.
- Dihedral angle ≈ 111.5° in gas phase, 90.2° in solid.
- O—O bond length: 1.475 Å; O—H bond length: 0.95 Å.
- Hydrogen bonding gives it higher boiling point (150 °C) than expected.
Preparation
| Method | Reaction | Notes |
|---|---|---|
| Laboratory | BaO₂·8H₂O + H₂SO₄ → BaSO₄↓ + H₂O₂ + 8H₂O | Phosphoric acid preferred (avoids catalytic decomposition) |
| Industrial (auto-oxidation) | 2-Ethylanthraquinol + O₂ → 2-Ethylanthraquinone + H₂O₂ | Catalyst: Pd; H₂O₂ extracted with water |
| Electrolysis | 50% H₂SO₄ electrolyzed → Perdisulphuric acid → H₂O₂ + H₂SO₄ | Perdisulphuric acid (H₂S₂O₈) hydrolyzed |
Physical Properties
| Property | Value |
|---|---|
| Appearance | Colourless syrupy liquid (pure) |
| Odour | Slightly pungent |
| Density | 1.44 g/mL (pure), ~1.11 g/mL (30%) |
| Boiling point | 150 °C (pure) — decomposes before boiling at 1 atm |
| Freezing point | −0.89 °C (pure) |
| Stability | Decomposes to H₂O + ½O₂; catalyzed by light, heat, metal ions |
Chemical Properties
H₂O₂ acts as both an oxidizing agent and a reducing agent.
As an oxidizing agent: 1. PbS + 4H₂O₂ → PbSO₄ + 4H₂O (black → white; used in restoring old paintings) 2. 2Fe²⁺ + 2H⁺ + H₂O₂ → 2Fe³⁺ + 2H₂O 3. 2I⁻ + 2H⁺ + H₂O₂ → I₂ + 2H₂O
As a reducing agent: 1. H₂O₂ + Cl₂ → 2HCl + O₂ (bleaching action) 2. 2MnO₄⁻ + 6H⁺ + 5H₂O₂ → 2Mn²⁺ + 8H₂O + 5O₂ 3. H₂O₂ + 2HOCl → 2HCl + O₂ + H₂O (dechlorination)
Other reactions:
- Decomposition: 2H₂O₂ → 2H₂O + O₂ (catalyzed by MnO₂, Pt, Fe²⁺, light)
- Bleaching: H₂O₂ → H₂O + [O] (nascent oxygen bleaches)
Uses of H₂O₂
1. Bleaching agent for textiles, paper, hair, and teeth. 2. Disinfectant and antiseptic (3% solution for wounds). 3. Oxidizing agent in chemical synthesis. 4. Rocket fuel (as oxidizer with hydrazine). 5. Pollution control — treating industrial effluents. 6. Etching agent in electronics. 7. Restoring old paintings (PbS → PbSO₄ conversion).
Storage of H₂O₂
- Stored in wax-lined glass or plastic bottles (surface of glass catalyzes decomposition).
- A stabilizer (acetanilide or phosphoric acid) is added to prevent decomposition.
- Stored in dark, cool place (light and heat accelerate decomposition).
Heavy Water (D₂O)
| Property | H₂O | D₂O |
|---|---|---|
| Molecular mass | 18 | 20 |
| Density (g/mL) | 0.997 | 1.107 |
| Boiling point | 100 °C | 101.4 °C |
| Freezing point | 0 °C | 3.8 °C |
| Viscosity | Lower | Higher (≈ 1.25 times) |
Uses of heavy water: 1. Moderator in nuclear reactors (slows down fast neutrons). 2. Tracer in metabolic studies. 3. Preparation of deuterium compounds.
Dihydrogen as a Fuel
| Advantage | Limitation |
|---|---|
| High calorific value (150 kJ/g vs 50 kJ/g for petrol) | Low ignition temperature — explosive with air |
| Clean fuel — produces only water | Difficult to store and transport |
| Abundant (as component of water) | Current production is expensive (electrolysis) |
| Can be used in fuel cells | Hydrogen embrittlement of metals |
Hydrogen economy: Concept of using hydrogen as the primary energy carrier. Water is electrolyzed using solar/nuclear energy to produce H₂, which is stored and transported, then used in fuel cells to generate electricity with only water as the by-product.
Common Mistakes
1. Calling H₂O₂ an oxidizing agent only: H₂O₂ acts as both oxidizing and reducing agent depending on the reaction. With KMnO₄ it is a reducing agent; with PbS it is an oxidizing agent.
2. Confusing temporary and permanent hardness: Temporary hardness is due to bicarbonates (removed by boiling). Permanent hardness is due to chlorides and sulphates (not removed by boiling; requires ion exchange or chemical treatment).
3. Wrong hydride classification: BeH₂ and MgH₂ are covalent/molecular hydrides, not ionic. Only Group 1 (except LiH has partial covalent character) and heavier Group 2 metals form ionic hydrides.
4. Electron-deficient vs electron-rich hydrides: BH₃ is electron-deficient (only 6 electrons around B); H₂O is electron-rich (2 lone pairs on O). These are different from electron-precise hydrides like CH₄ (8 electrons around C).
5. Density anomaly: Ice floats because it is less dense than water (open structure due to H-bonds). At 4 °C, water has maximum density. Students sometimes write "ice is denser than water."
6. Lane's process vs Bosch process: Lane's process uses iron and steam; Bosch process uses water gas (CO + H₂) with steam. Don't confuse them.
Exam Tips
1. Position of hydrogen (3–5 marks): Be ready to explain why hydrogen is placed separately, with at least 2 similarities to alkali metals and 2 similarities to halogens.
2. Hydride classification (2–3 marks): Know all three types with examples. A comparison table is a very common exam question.
3. H₂O₂ reactions (2–3 marks): Practice at least 2 reactions where H₂O₂ acts as oxidizing agent and 2 where it acts as reducing agent. Mention the oxidation number changes.
4. Water hardness (2–3 marks): Know the difference between temporary and permanent hardness, and at least 2 methods of softening with equations.
5. Preparation of H₂ (1–2 marks): Lab method (Zn + dil. H₂SO₄) and industrial methods (Lane's process, electrolysis) are frequently asked.
6. Structure of H₂O₂ (1 mark): Draw the open book/skewed structure with bond angle.
7. Hydrogen as fuel (2 marks): List advantages (clean, high calorific value) and limitations (storage, explosiveness, cost). This is a common short-answer question.
Quick Revision Notes
- Position of hydrogen: Resembles alkali metals (ns¹, H⁺) and halogens (needs 1 e⁻, forms H⁻, diatomic).
- Isotopes: Protium (¹H), Deuterium (²H or D), Tritium (³H or T, radioactive).
- Preparation of H₂: Lab — Zn + dil. H₂SO₄. Industrial — electrolysis, Lane's process, Bosch process.
- Properties of H₂: Lightest gas, colourless, odourless, non-poisonous, reducing agent, forms hydrides.
- Hydrides: Ionic (NaH, CaH₂) — crystalline, conduct when molten; Covalent (CH₄, NH₃, H₂O) — volatile, low boiling; Metallic (TiH₂, PdHx) — interstitial, non-stoichiometric.
- Water: Bent structure, 104.5° bond angle, hydrogen bonding → high bp, mp, surface tension, specific heat.
- Hard water: Temporary (bicarbonates, removed by boiling), Permanent (chlorides, sulphates, needs ion exchange).
- H₂O₂: Open book structure, acts as both oxidizing and reducing agent, stored with stabilizer in wax-lined bottles.
- Heavy water (D₂O): bp 101.4 °C, used as nuclear moderator.
- Hydrogen as fuel: High calorific value, clean (only H₂O produced), but storage and transport challenges.
- Density of water: Maximum at 4 °C. Ice is less dense than water (hydrogen bonding creates open structure).