JEE Chemistry Hydrogen Notes Set 04

Position and Nature of Hydrogen

Hydrogen has the configuration 1s¹, is the lightest element and is the most abundant element in the universe, although on earth it is found mostly in combined form. It resembles the alkali metals (one valence electron, tendency to form H⁺, oxidation state +1, reducing nature, liberation at the cathode) and also the halogens (one electron short of the helium configuration, forms H⁻, exists as a diatomic molecule, forms covalent compounds). Because of this dual behaviour it has no fixed position in the periodic table.

The H–H bond is very strong (436 kJ mol⁻¹, 74 pm), so molecular hydrogen is unreactive at ordinary temperature, while atomic hydrogen is extremely reactive.

Isotopes and the Isotopic Effect

Isotopes of hydrogen differ so much in mass that their physical properties (melting point, boiling point, bond energy, heat of vaporisation) change noticeably, with the order H < D < T. This is the isotopic effect; chemical properties differ only in rate (a C–H bond breaks faster than a C–D bond). Isotopic effects are small for other elements because the relative change in mass is small.

Forms of Hydrogen

By oxidation state: proton H⁺ (+1; in water it exists as hydrated H⁺(H₂O)n, for example H₃O⁺), hydride H⁻ (−1) and atomic hydrogen (0).

By reactivity: atomic hydrogen (obtained at high temperature and low pressure) and nascent hydrogen (hydrogen at the moment of formation, for example Zn + dilute acid), which is a more powerful reductant than ordinary hydrogen.

Spin isomers: molecular hydrogen exists as ortho (parallel nuclear spins) and para (antiparallel nuclear spins) forms. At room temperature the ratio is 3 : 1, and at 0 K pure para hydrogen exists. They differ in physical properties but are chemically identical.

Preparation

Laboratory: granulated zinc with dilute H₂SO₄ or HCl: Zn + H₂SO₄ → ZnSO₄ + H₂. Impure zinc reacts faster because local galvanic cells form. Amphoteric metals (Be, Zn, Al, Sn, Pb, Si) liberate hydrogen with hot alkali: Zn + 2NaOH → Na₂ZnO₂ + H₂; 2Al + 2NaOH + 2H₂O → 2NaAlO₂ + 3H₂. Ionic hydrides give hydrogen with water: CaH₂ + 2H₂O → Ca(OH)₂ + 2H₂.

Metals and water: highly reactive metals (Na, K, Ca) react with cold water, moderately reactive metals (Mg, Al, Zn) with hot water, and less reactive metals (Fe, Ni, Sn) with steam.

Pure hydrogen: electrolysis of water made slightly conducting by acid or alkali: cathode 2H⁺ + 2e⁻ → H₂.

Industrial: Lane process (steam over red-hot iron: 3Fe + 4H₂O → Fe₃O₄ + 4H₂); water gas (C + H₂O → CO + H₂) and the Bosch water-gas shift (CO + H₂O → CO₂ + H₂ over Fe₂O₃/Cr₂O₃); steam reforming of natural gas (hydrocarbon + steam over Ni catalyst to CO + H₂).

Properties and Uses

Hydrogen is a colourless, odourless, tasteless, neutral gas of very low melting and boiling points (liquid hydrogen is a cryogenic fluid). It burns with a pale blue flame forming water. It combines with electropositive metals to give ionic hydrides (2Na + H₂ → 2NaH), with non-metals (halogens in the order F₂ > Cl₂ > Br₂ > I₂; N₂ in the Haber process; sulphur to H₂S), reduces metal oxides (CuO + H₂ → Cu + H₂O), forms methanol with CO over ZnO/Cr₂O₃, and hydrogenates unsaturated compounds over Ni, Pd or Pt (hardening of vegetable oils to vanaspati ghee).

Uses: ammonia manufacture, hydrogenation of oils, rocket fuel (liquid H₂ with liquid O₂), methanol synthesis, fuel cells, and filling balloons (with helium for safety). On a per-gram basis hydrogen releases far more energy (about 143 kJ g⁻¹) than LPG or methane.

Hydrides

1. Ionic (saline) hydrides are formed by s-block metals other than Be and Mg. They have rock-salt type structures, liberate hydrogen at the anode on electrolysis in the molten state, and react with water to give hydrogen. BeH₂ and MgH₂ are covalent and polymeric.

2. Metallic (interstitial) hydrides are formed by d- and f-block metals, are non-stoichiometric (for example TiH_1.8, ZrH_1.3), and conduct like the parent metal. Groups 7, 8 and 9 form no hydrides (hydride gap).

3. Covalent (molecular) hydrides are formed by p-block elements. They are electron deficient (group 13, such as B₂H₆), electron precise (group 14, such as CH₄, SiH₄) or electron rich (groups 15 to 17, such as NH₃, H₂O, HF, which have lone pairs and can act as Lewis bases).

Water and Hardness

Water is a V-shaped polar molecule (bond angle 104.5°) with extensive hydrogen bonding, which explains its high boiling point, specific heat and latent heats; its maximum density is at 4 °C. It is amphoteric, reacts with metals, non-metals (Cl₂ gives HCl and HOCl), acidic and basic oxides, and hydrolyses hydrides, carbides (CaC₂ to C₂H₂, Al₄C₃ to CH₄), nitrides (Mg₃N₂ to NH₃) and phosphides (Ca₃P₂ to PH₃).

Hard water forms scum with soap because Ca²⁺ and Mg²⁺ precipitate the soap anion. Temporary hardness (bicarbonates of Ca and Mg) is removed by boiling or by Clarke's process (Ca(HCO₃)₂ + Ca(OH)₂ → 2CaCO₃ + 2H₂O). Permanent hardness (chlorides and sulphates of Ca and Mg) is removed by washing soda (Na₂CO₃), Calgon (sodium hexametaphosphate), permutit (zeolite) or synthetic ion exchange resins. In the resin method the cation exchanger (R–SO₃H or R–COOH) replaces metal ions by H⁺ and the anion exchanger replaces anions by OH⁻, giving demineralised water; they are regenerated by dilute acid and dilute NaOH respectively.

Heavy Water

Heavy water D₂O is obtained by repeated electrolysis of water, since H⁺ and OH⁻ are discharged faster and D₂O accumulates in the residue. Its physical constants are slightly higher than those of water (b.p. 101.4 °C), and it reacts more slowly in chemical reactions. It reacts like water to give deuterium compounds: 2D₂O + 2Na → 2NaOD + D₂. It is used as a moderator in nuclear reactors and as a tracer in studying reaction mechanisms.

Hydrogen Peroxide

Preparation: Na₂O₂ + H₂SO₄ (ice cold) → Na₂SO₄ + H₂O₂; Merck process using hydrated BaO₂·8H₂O with H₂SO₄ or H₃PO₄; industrially by auto-oxidation of 2-alkylanthraquinol, by electrolytic oxidation of 50% H₂SO₄ through peroxodisulphuric acid, or by oxidation of isopropyl alcohol.

Structure and properties: a non-planar, open-book molecule with an O–O single bond (dipole moment 2.1 D). It is a pale blue, syrupy liquid, a weak dibasic acid, with a boiling point (150 °C) higher than that of water. It is stored in dark plastic or wax-lined bottles with stabilisers such as acetanilide or phosphoric acid because it decomposes (disproportionates) in light and on contact with rough surfaces or metal ions: 2H₂O₂ → 2H₂O + O₂. The 30% solution is called perhydrol.

Oxidising action: PbS to PbSO₄ (restoring old paintings), sulphite to sulphate, nitrite to nitrate, H₂S to sulphur, iodide to iodine. Reducing action towards stronger oxidants: Ag₂O to Ag, O₃ to O₂, acidified KMnO₄ to Mn²⁺.

Volume strength: a 'V volume' solution gives V mL of oxygen at STP per mL of solution. Molarity = V/11.2 and percentage (w/v) = 1.7 × V/5.6 ≈ 0.304 V. For example 11.2 volume is 1 M and 20 volume is about 6.07%. Uses: bleaching, antiseptic, rocket propellant, and as a green oxidant.