Hydrogen Class 11 Notes

• Protium is the most abundant (99.98%), then deuterium (~0.016%), tritium is radioactive.
• Deuterium replaces hydrogen in heavy water D₂O.
• Physical properties differ: D₂O freezes at 3.8 °C and boils at 101.4 °C.

• Like alkali metals: one valence electron, forms H⁺, strong reducing agent in some contexts, diatomic H₂.
• Like halogens: needs one electron to complete its shell, forms hydride ion H⁻ and diatomic covalent molecules.
• Unlike both: very low ionisation enthalpy is misleading — its IE₁ is actually higher than alkali metals; it is a unique element, so it is placed separately above Group 1.

• Very light, colourless, odourless gas; burns with a pale blue flame to give water.
• Acts as a reducing agent: reduces metal oxides (CuO → Cu) and metal ions.
• Combines with halogens, nitrogen (Haber process) and metals to form hydrides.
• Used in the Haber process, hydrogenation of oils, rocket fuel, and as a clean fuel with only H₂O as product.

• Ionic hydrides are high-melting solids that react with water to release H₂.
• Covalent hydrides are molecular: water, ammonia, methane.
• Electron-precise hydrides (CH₄), electron-rich (NH₃, with lone pair) and electron-deficient (B₂H₆) are sub-types of covalent hydrides.

• Higher boiling point, density and viscosity than H₂O but similar ionic chemistry.
• Physical constants differ measurably for the isotopes.
• Used in NMR and tracer studies to follow reaction paths.

• Stored in dark bottles with a stabiliser; catalase speeds decomposition in biological tissue.
• Oxidising: H₂O₂ + 2e⁻ → 2OH⁻; reducing: H₂O₂ → O₂ + 2H⁺ + 2e⁻.
• Antiseptic (dilute), bleaching agent for whitening, and rocket propellant oxidiser.
• Volumes strength: '10 volume' H₂O₂ releases 10 L of O₂ per litre.

Example: Why is hydrogen not placed with the halogens permanently?

Solution: Hydrogen shares one electron with halogens (forms H⁻ like halides) but one valence electron with alkali metals (forms H⁺). Neither analogy is complete, so periodic tables place it above Group 1 in its own position.

Example: What is the oxidation state of hydrogen in NaH versus HCl?

Solution: In NaH (an ionic hydride), hydrogen is the more electronegative partner, so it is −1. In HCl, chlorine is more electronegative, so hydrogen is +1.

Exercises (NCERT Q9.1–Q9.36)

• 9.1) Justify the position of hydrogen in the periodic table on the basis of its electronic configuration.
• 9.2) Write the names of isotopes of hydrogen. What is the mass ratio of these isotopes?
• 9.3) Why does hydrogen occur in a diatomic form rather than in a monoatomic form under normal conditions?
• 9.4) How can the production of dihydrogen, obtained from 'coal gasification', be increased?
• 9.5) Describe the bulk preparation of dihydrogen by electrolytic method. What is the role of an electrolyte in this process?
• 9.6) Complete the following reactions: (i) H₂(g) + MₓO_y(s) → (Δ) (ii) CO(g) + H₂(g) → (catalyst) (iii) C₃H₈(g) + 3H₂O(g) → (catalyst) (iv) Zn(s) + NaOH(aq) → (heat)
• 9.7) Discuss the consequences of high enthalpy of H–H bond in terms of chemical reactivity of dihydrogen.
• 9.8) What do you understand by (i) electron-deficient, (ii) electron-precise, and (iii) electron-rich compounds of hydrogen? Provide justification with suitable examples.
• 9.9) What characteristics do you expect from an electron-deficient hydride with respect to its structure and chemical reactions?
• 9.10) Do you expect the carbon hydrides of the type (C_nH_2n+2) to act as 'Lewis' acid or base? Justify your answer.
• 9.11) What do you understand by the term 'non-stoichiometric hydrides'? Do you expect this type of the hydrides to be formed by alkali metals? Justify your answer.
• 9.12) How do you expect the metallic hydrides to be useful for hydrogen storage? Explain.
• 9.13) How does the atomic hydrogen or oxy-hydrogen torch function for cutting and welding purposes? Explain.
• 9.14) Among NH₃, H₂O and HF, which would you expect to have highest magnitude of hydrogen bonding and why?
• 9.15) Saline hydrides are known to react with water violently producing fire. Can CO₂, a well known fire extinguisher, be used in this case? Explain.
• 9.16) Arrange the following: (i) CaH₂, BeH₂ and TiH₂ in order of increasing electrical conductance. (ii) LiH, NaH and CsH in order of increasing ionic character. (iii) H–H, D–D and F–F in order of increasing bond dissociation enthalpy. (iv) NaH, MgH₂ and H₂O in order of increasing reducing property.
• 9.17) Compare the structures of H₂O and H₂O₂.
• 9.18) What do you understand by the term 'auto-protolysis' of water? What is its significance?
• 9.19) Consider the reaction of water with F₂ and suggest, in terms of oxidation and reduction, which species are oxidised/reduced.
• 9.20) Complete the following chemical reactions. (i) PbS(s) + H₂O₂(aq) → (ii) MnO₄⁻(aq) + H₂O₂(aq) → (iii) CaO(s) + H₂O(g) → (v) AlCl₃(g) + H₂O(l) → (vi) Ca₃N₂(s) + H₂O(l) → Classify the above into (a) hydrolysis, (b) redox and (c) hydration reactions.
• 9.21) Describe the structure of the common form of ice.
• 9.22) What causes the temporary and permanent hardness of water?
• 9.23) Discuss the principle and method of softening of hard water by synthetic ion-exchange resins.
• 9.24) Write chemical reactions to show the amphoteric nature of water.
• 9.25) Write chemical reactions to justify that hydrogen peroxide can function as an oxidising as well as reducing agent.
• 9.26) What is meant by 'demineralised' water and how can it be obtained?
• 9.27) Is demineralised or distilled water useful for drinking purposes? If not, how can it be made useful?
• 9.28) Describe the usefulness of water in biosphere and biological systems.
• 9.29) What properties of water make it useful as a solvent? What types of compound can it (i) dissolve, and (ii) hydrolyse?
• 9.30) Knowing the properties of H₂O and D₂O, do you think that D₂O can be used for drinking purposes?
• 9.31) What is the difference between the terms 'hydrolysis' and 'hydration'?
• 9.32) How can saline hydrides remove traces of water from organic compounds?
• 9.33) What do you expect the nature of hydrides is, if formed by elements of atomic numbers 15, 19, 23 and 44 with dihydrogen? Compare their behaviour towards water.
• 9.34) Do you expect different products in solution when aluminium(III) chloride and potassium chloride treated separately with (i) normal water (ii) acidified water, and (iii) alkaline water? Write equations wherever necessary.
• 9.35) How does H₂O₂ behave as a bleaching agent?
• 9.36) What do you understand by the terms: (i) hydrogen economy (ii) hydrogenation (iii) 'syngas' (iv) water-gas shift reaction (v) fuel-cell?

Answers at a glance

1. 1
   Q9.1 → H is 1s¹: like alkali metals it has one electron in its valence shell (ns¹), but it is also one electron short of the stable ns² configuration like halogens — so it is placed in group 1 (and above group 17).
2. 2
   Q9.2 → Protium (¹H), Deuterium (²H or D) and Tritium (³H or T); the mass ratio of the three isotopes is 1 : 2 : 3.
3. 3
   Q9.3 → Free monoatomic hydrogen is extremely reactive; the H–H bond enthalpy (435.9 kJ/mol) is high, so at ordinary conditions H atoms combine to the more stable diatomic molecule H₂.
4. 4
   Q9.4 → Coal gasification (C + H₂O → CO + H₂) is endothermic, so the yield is increased by raising the temperature; the CO + H₂O ⇌ CO₂ + H₂ water-gas shift step is exothermic, so the additional H₂ from it is favoured by lowering the temperature (Le Chatelier).
5. 5
   Q9.5 → Acidified (or alkalised) water is electrolysed: H₂ at cathode, O₂ at anode (2H₂O → 2H₂ + O₂). Pure water is a poor conductor, so the electrolyte supplies ions that carry the current through the solution.
6. 6
   Q9.6 → (i) MₓO_y + yH₂ → xM + yH₂O (oxide reduced) (ii) CO + 2H₂ → CH₃OH (methanol) (iii) C₃H₈ + 3H₂O → 3CO + 7H₂ (steam reforming) (iv) Zn + 2NaOH → Na₂ZnO₂ + H₂.
7. 7
   Q9.7 → The high bond enthalpy (435.9 kJ/mol) makes H₂ kinetically stable and rather unreactive at room temperature; it reacts only at high temperature, with a catalyst, or in presence of light.
8. 8
   Q9.8 → (i) Electron-deficient: fewer than 8 electrons around the central atom (e.g. B₂H₆). (ii) Electron-precise: all electrons are in bonding pairs and the octet is complete (e.g. CH₄, C₂H₆). (iii) Electron-rich: central atom has lone pairs after bonding (e.g. NH₃, H₂O, HF).
9. 9
   Q9.9 → Structurally they need multi-centre bonds (e.g. 3-centre–2-electron B–H–B bridges in B₂H₆) to make up the electron shortage; chemically they are Lewis acids — they readily accept electron pairs from Lewis bases in addition reactions.
10. 10
    Q9.10 → Neither. In alkanes every electron is in bonding pairs, so there are no lone pairs to donate (not a base) and no vacant orbital to accept (not an acid) — they are electron-precise and inert.
11. 11
    Q9.11 → Hydrides of indefinite composition in which H atoms occupy interstitial sites of a metal lattice (e.g. TiH₁.₇, LaH₂.₈₇). Alkali metals form only stoichiometric ionic hydrides (e.g. NaH), so not such hydrides.
12. 12
    Q9.12 → They absorb and hold hydrogen in the crystal lattice; the stored hydrogen is released on heating (or lowering pressure), so alloys like LaNi₅ act as a reversible hydrogen reservoir for mobile fuel-cell applications.
13. 13
    Q9.13 → H₂ is dissociated to atomic H in an electric arc; on striking the metal, atomic H recombines to H₂ releasing enormous heat (up to ~4000 °C), which melts the metal for cutting and welding.
14. 14
    Q9.14 → H₂O — each O can form two hydrogen bonds (two H atoms plus two lone pairs), so water develops the most extensive, strongest hydrogen bonding of the three.
15. 15
    Q9.15 → No. The hydride (e.g. NaH, CaH₂) contains the strongly reducing hydride ion H⁻, and CO₂ would itself be reduced while the runaway H₂ production and heat continue, so CO₂ cannot extinguish this fire.
16. 16
    Q9.16 → (i) BeH₂ < TiH₂ < CaH₂ (ii) LiH < NaH < CsH (iii) F–F < H–H < D–D (iv) H₂O < MgH₂ < NaH.
17. 17
    Q9.17 → H₂O is bent (V-shaped), bond angle 104.5°, oxygen sp³ hybridised with two lone pairs; H₂O₂ is a non-planar open-book molecule about the O–O bond with a dihedral angle of 94.8° (gas phase).
18. 18
    Q9.18 → The self-ionisation of water into H₃O⁺ and OH⁻ (2H₂O ⇌ H₃O⁺ + OH⁻). Its significance is that it makes water amphoteric, underlies the acid–base (pH) scale and all aqueous equilibria.
19. 19
    Q9.19 → 2F₂ + 2H₂O → 4HF + O₂. Fluorine is reduced (0 to −1); the oxygen of water is oxidised (−2 to 0).
20. 20
    Q9.20 → (i) PbS + 4H₂O₂ → PbSO₄ + 4H₂O (redox) (ii) 2MnO₄⁻ + 5H₂O₂ + 6H⁺ → 2Mn²⁺ + 5O₂ + 8H₂O (redox) (iii) CaO + H₂O → Ca(OH)₂ (hydration) (v) AlCl₃ + 3H₂O → Al(OH)₃ + 3HCl (hydrolysis) (vi) Ca₃N₂ + 6H₂O → 3Ca(OH)₂ + 2NH₃ (hydrolysis).
21. 21
    Q9.21 → Ice has an open hexagonal (wurtzite-like) structure: each O is tetrahedrally bonded by hydrogen bonds to four neighbouring O atoms, leaving large empty spaces — which is why ice is less dense than water and floats.
22. 22
    Q9.22 → Temporary hardness — bicarbonates of calcium and magnesium, Ca(HCO₃)₂ and Mg(HCO₃)₂; permanent hardness — chlorides and sulphates, CaCl₂, MgCl₂, CaSO₄ and MgSO₄.
23. 23
    Q9.23 → Hard water is passed through a cation-exchange resin that exchanges Na⁺ (or H⁺) for Ca²⁺/Mg²⁺ and then an anion-exchange resin that exchanges OH⁻ for anions; the soft, deionised water results, and the resins are regenerated with brine and acid/base.
24. 24
    Q9.24 → As a base (accepts H⁺): H₂O + HCl → H₃O⁺ + Cl⁻; as an acid (donates H⁺): H₂O + NH₂⁻ → NH₃ + OH⁻ (or NaH + H₂O → NaOH + H₂).
25. 25
    Q9.25 → Oxidising: 2I⁻ + H₂O₂ + 2H⁺ → I₂ + 2H₂O (H₂O₂ gains electrons, is reduced — itself the oxidant). Reducing: Ag₂O + H₂O₂ → 2Ag + H₂O + O₂ (H₂O₂ loses electrons, is oxidised — itself the reductant).
26. 26
    Q9.26 → Water free from all mineral ions. It is obtained by passing water first through a cation-exchange (H⁺) resin and then an anion-exchange (OH⁻) resin, which replace the cations and anions by H⁺ and OH⁻ (that combine to water).
27. 27
    Q9.27 → No — it lacks the essential minerals. It can be made useful by aerating it and adding the required mineral salts (calcium, magnesium, etc.) back.
28. 28
    Q9.28 → In living systems water is the medium for biochemical reactions, hydrolyses biomolecules, regulates body temperature and transports nutrients; in the biosphere it shapes climate, weathers rocks and carries plant nutrients in soil.
29. 29
    Q9.29 → High polarity and a very high dielectric constant (78) weaken ionic bonds, while strong hydrogen bonding dissolves polar molecules. It (i) dissolves ionic and polar covalent compounds and (ii) hydrolyses salts, oxides, chlorides, nitrides and carbides of less electropositive elements.
30. 30
    Q9.30 → No. D₂O modifies the rate of enzyme-catalysed (metabolic) reactions and inhibits cell division, so it is not fit for drinking.
31. 31
    Q9.31 → Hydrolysis is a chemical reaction in which water cleaves the compound (its H and OH add across the bond, e.g. Ca₃N₂ → Ca(OH)₂ + NH₃); hydration simply adds whole water molecules as water of crystallisation without breaking the compound (e.g. CuSO₄·5H₂O).
32. 32
    Q9.32 → Saline hydrides react with water completely and irreversibly, e.g. CaH₂ + 2H₂O → Ca(OH)₂ + 2H₂, so they act as efficient drying agents for organic solvents.
33. 33
    Q9.33 → Z = 15 (P) gives covalent PH₃ — behaves normally, not hydrolysed; Z = 19 (K) gives ionic/saline KH — reacts violently with water (KOH + H₂); Z = 23 (V) and 44 (Ru) give metallic (interstitial/non-stoichiometric) hydrides — inert towards water.
34. 34
    Q9.34 → Yes for AlCl₃. (i) Normal water: AlCl₃ + 3H₂O → Al(OH)₃ + 3HCl. (ii) Acidified water: excess H⁺ suppresses hydrolysis, Al³⁺ stays as hydrated ions. (iii) Alkaline water dissolves the hydroxide: Al(OH)₃ + OH⁻ → [Al(OH)₄]⁻. KCl does not hydrolyse in any of the three — K⁺ and Cl⁻ are spectator ions.
35. 35
    Q9.35 → In the presence of moisture H₂O₂ decomposes to give nascent oxygen (H₂O₂ → H₂O + [O]), which oxidises the coloured dye or pigment to colourless compounds — so it bleaches hair, wool, silk and fabrics in acidic medium.
36. 36
    Q9.36 → (i) Using H₂ as a clean fuel (with O₂ in fuel cells) to replace fossil fuels. (ii) Addition of H₂ to unsaturated compounds (e.g. oils → vanaspati ghee) over Ni/Pd/Pt. (iii) Syngas = the CO + H₂ mixture (water gas) used to synthesise methanol and other chemicals. (iv) CO + H₂O ⇌ CO₂ + H₂, which raises the H₂ yield in dihydrogen production. (v) A galvanic cell that continuously produces electricity from the H₂–O₂ reaction to give water (2H₂ + O₂ → 2H₂O).

MCQ Questions

1 Mark Questions

1. 1
   Write the names and symbols of the isotopes of hydrogen. (Ans: protium ¹H, deuterium ²H (D), tritium ³H (T))
2. 2
   What is water gas? (Ans: an equimolar mixture of CO and H₂, also called synthesis gas)
3. 3
   What causes permanent hardness of water? (Ans: chlorides and sulphates of calcium and magnesium)
4. 4
   Define auto-protolysis of water. (Ans: the self-ionisation of water into H₃O⁺ and OH⁻, Kw = 10⁻¹⁴ at 25 °C)

2 Mark Questions

1. 1
   Write reactions showing that H₂O₂ can act as an oxidising agent. (Ans: H₂O₂ + 2I⁻ + 2H⁺ → I₂ + 2H₂O; H₂S + H₂O₂ → S + 2H₂O)
2. 2
   Write reactions showing that H₂O₂ can act as a reducing agent. (Ans: Ag₂O + H₂O₂ → 2Ag + H₂O + O₂; H₂O₂ reduces KMnO₄ in acid)
3. 3
   How do saline hydrides act as drying agents? (Ans: CaH₂ + 2H₂O → Ca(OH)₂ + 2H₂ removes water completely)
4. 4
   Distinguish between temporary and permanent hardness. (Ans: temporary from bicarbonates and removed by boiling; permanent from chlorides/sulphates and not removed by boiling)

3 Mark Questions

1. 1
   Describe the principle and method of softening hard water by synthetic ion-exchange resins. (Ans: cation exchanger exchanges Na⁺/H⁺ for Ca²⁺/Mg²⁺, anion exchanger exchanges OH⁻ for anions; regenerated by brine/acids)
2. 2
   Explain the structure of ice and why it is less dense than liquid water. (Ans: open hexagonal H-bonded structure, wurtzite-like, with empty spaces)
3. 3
   Write equations to show hydrogen peroxide as both an oxidising and a reducing agent, and justify with oxidation states. (Ans: oxidising — oxygen reduces from −1 to −2; reducing — oxygen oxidises from −1 to 0)
4. 4
   How does H₂O₂ behave as a bleaching agent? (Ans: it gives nascent oxygen (H₂O₂ → H₂O + [O]) that oxidises the coloured dye to a colourless compound)

5 Mark Questions

1. 1
   Classify hydrogen compounds into electron-deficient, electron-precise and electron-rich types with suitable examples, and compare their structure and chemical reactivity. (Ans: B₂H₆ deficient (Lewis acid); CH₄/C₂H₆ precise (inert); NH₃/H₂O/HF rich (donate lone pairs))
2. 2
   Explain the term 'hydrogen economy'. Describe the routes of dihydrogen production, the water-gas shift reaction, and the role of a hydrogen–oxygen fuel cell, writing the electrode reactions.
3. 3
   Discuss temporary and permanent hardness of water, their causes, and the ion-exchange and other methods used to soften hard water, with equations.

Assignments walk the chapter's topics in order; the tests follow the current CBSE blueprint (MCQ/VSA, 2M, 3M and 5M questions) with answer keys. They will appear here as direct download links.