Group 16 Elements (O, S, Se): JEE Main Guide

Group 16 — the chalcogens — generates two to four JEE Main questions per session, with sulfur and its compounds being the dominant focus. Oxygen's anomalous behaviour, the structures of sulfur allotropes, and the oxyacids of sulfur (particularly H₂SO₄ and H₂SO₃) are the highest-priority topics. This guide organises Group 16 chemistry systematically so you can answer questions on any subtopic in under a minute.

Oxygen: Anomalous Behaviour and Allotropes

Oxygen is anomalous compared to the heavier chalcogens: (1) small atomic size allows multiple bond formation (O=O double bond in O₂; heavier chalcogens form S-S, Se-Se single bonds in their allotropic forms); (2) high electronegativity (second highest after F); (3) no available d-orbitals (maximum valency 2, forms only oxides and peroxides), while S, Se, Te can use d-orbitals and form +4 and +6 compounds (SO₄²⁻, SeO₄²⁻). O₂ is paramagnetic (two unpaired electrons in π* orbitals — MOT explanation that JEE tests). Allotropes of oxygen: O₂ (dioxygen, stable) and O₃ (ozone, angular V-shape, bond angle 117°, diamagnetic due to resonance delocalisation).

Ozone: prepared by passing O₂ through silent electric discharge (ozoniser). O₃ is a strong oxidising agent (decomposes to give nascent O). Test for O₃: turns moist starch-iodide paper blue (oxidises I⁻ to I₂). Ozone layer in stratosphere absorbs UV-B and UV-C radiation. CFCs deplete ozone via radical chain mechanism. Take a free inorganic chemistry mock. See our p-block guide.

Allotropes and Properties of Sulfur

Rhombic sulfur (α-S₈): stable at room temperature, crown-shaped S₈ rings, yellow. Monoclinic sulfur (β-S₈): stable above 96°C (transition temperature), needle-shaped crystals, also S₈ rings but different packing. Plastic sulfur: formed by pouring molten S into cold water → long zigzag chains (amorphous). Milk of sulfur: colloidal sulfur from acid precipitation. All forms convert to rhombic sulfur below 96°C (thermodynamic stability). JEE asks about the transition temperature, the structure (S₈ rings), and relative stability of allotropes. Sulfur shows wide range of oxidation states: −2 (H₂S, sulfides), 0 (elemental), +4 (SO₂, H₂SO₃), +6 (SO₃, H₂SO₄).

Sulfur Dioxide and Trioxide

SO₂: angular molecule (sp³ hybridisation of S with one lone pair), bond angle ~119°, has resonance (S=O and S–O bonds are equivalent at 143 pm). SO₂ is acidic oxide — dissolves in water to give H₂SO₃ (sulfurous acid). SO₂ is a reducing agent (oxidised to SO₃ by strong oxidants) and also a mild oxidising agent (e.g., oxidises H₂S to S). SO₂ is a bleaching agent — bleaches by reduction (decolourises coloured organic dyes by reducing them; colour returns on exposure to air). This is in contrast to Cl₂, which bleaches by oxidation (permanent). SO₃: trigonal planar monomer but polymerises to chains and rings in solid state. Acidic oxide — reacts with water vigorously to form H₂SO₄. Industrial preparation: Contact Process (2SO₂ + O₂ ⇌ 2SO₃ using V₂O₅ catalyst).

Sulfuric Acid: Structure and Properties

H₂SO₄ is a diprotic acid, strong acid, viscous oily liquid, density 1.84 g/mL, miscible with water in all proportions with large exotherm (always add acid to water, not water to acid). Properties: (1) Dehydrating agent — removes water from organic compounds (sugars are charred to carbon; HCOOH → CO + H₂O; CH₃COOH → CH₂=C=O, ketene). (2) Oxidising agent (concentrated) — oxidises metals (Cu → CuSO₄ + SO₂), non-metals (S, C, P), and many organic compounds. Dilute H₂SO₄ does NOT show oxidising properties beyond acid properties. (3) Sulphonation of benzene: C₆H₆ + H₂SO₄ → C₆H₅SO₃H + H₂O (electrophilic substitution using SO₃ as electrophile). For the complete inorganic strategy see our oxidation states guide and our acid-base guide.