HydrogenNEET Chemistry · Class 11 · NCERT Chapter 10

Hydrogen resembles Group 1 (alkali metals) in that it has one electron in its outermost shell and tends to form H⁺ by losing that electron. It resembles Group 17 (halogens) in that it needs one more electron to complete its valence shell and can form H⁻ (hydride ion). However, hydrogen differs fundamentally from both groups: it is a non-metal, forms a diatomic molecule (H₂), and has a very low ionisation energy compared to halogens but does not form salts the way alkali metals do. So it occupies a unique, anomalous position.

In ice, water molecules form a hydrogen-bonded lattice with an open, cage-like hexagonal structure. This structure has larger intermolecular spaces than liquid water, making ice less dense than liquid water (density of ice ≈ 0.917 g/cm³ vs water = 1.0 g/cm³). When ice melts, some hydrogen bonds break and the open structure collapses, allowing molecules to pack more closely. This is why water has maximum density at 4 °C, not at 0 °C.

H₂O₂ can act as both an oxidising agent and a reducing agent. As an oxidising agent (more common): H₂O₂ + 2H⁺ + 2e⁻ → 2H₂O (gains electrons, oxidises the substrate). Example: oxidises Fe²⁺ to Fe³⁺, and bleaches coloured matter. As a reducing agent: H₂O₂ → O₂ + 2H⁺ + 2e⁻ (loses electrons when reacting with a stronger oxidising agent). Example: reduces Cl₂ to HCl, and reduces acidic KMnO₄.

(1) Ionic (saline) hydrides: formed by highly electropositive metals (Group 1 and some Group 2). The H is H⁻ (hydride ion). Examples: NaH, CaH₂. They are crystalline solids that react vigorously with water to give H₂. (2) Covalent (molecular) hydrides: formed by non-metals. H is covalently bonded. Examples: CH₄, NH₃, H₂O, HF. (3) Metallic (interstitial) hydrides: formed by transition metals (d-block) and lanthanides/actinides. H atoms occupy interstitial positions in the metal lattice. Non-stoichiometric. Examples: PdHₓ, TiH₂.

Water (M = 18 g/mol) has a boiling point of 100 °C, which is much higher than expected by comparison with other Group 16 hydrides (H₂S: −60 °C, H₂Se: −41 °C). The reason is extensive hydrogen bonding between water molecules. Each water molecule can form up to 4 hydrogen bonds (2 as donor, 2 as acceptor). A lot of energy is needed to break these hydrogen bonds before vaporisation can occur, resulting in a high boiling point.

H₂O₂ has an open-book (butterfly) structure. The two O-H bonds are not in the same plane; the dihedral (torsional) angle between the two O-H bonds is about 111.5° in the gas phase (different in solid due to hydrogen bonding). The O-O bond is a single bond (longer than the O=O in O₂). The non-planar structure means H₂O₂ is chiral in principle, but the two mirror-image forms interconvert rapidly.