p-Block Elements

1. Chapter at a Glance

The p-block contains groups 13 to 18, with outer configuration ns² np^1–6. It has metals, metalloids, non-metals and the noble gases, so its chemistry is very varied. The current JEE syllabus covers groups 13 and 14 (Class 11) and groups 15 to 18 (Class 12). JEE Main typically asks about structures of oxoacids and interhalogen compounds, inert pair effect, anomalous behaviour of the first member, and properties of important compounds.

2. General Trends

• Inert pair effect: down a group, the ns² electrons are less willing to take part in bonding (poor shielding by d and f electrons). So the lower oxidation state (group number − 2) becomes more stable: Tl⁺ > Tl³⁺, Pb²⁺ > Pb⁴⁺, Bi³⁺ > Bi⁵⁺. Hence PbO₂ and NaBiO₃ are strong oxidising agents.
• First members (B, C, N, O, F) differ from the rest: small size, high electronegativity and no d orbitals — they can form pπ–pπ multiple bonds (N≡N, O=O, C=C) and their maximum covalency is 4.
• Non-metallic character decreases down a group; metallic character increases.

3. Group 13: Boron Family (B, Al, Ga, In, Tl)

• Boron is a non-metal; the others are metals. Common oxidation state +3; +1 becomes more stable for Tl.
• Trihalides (BF₃, AlCl₃) are electron deficient Lewis acids. Lewis acidity: BF₃ < BCl₃ < BBr₃ (back-bonding from F's 2p to B's 2p is strongest in BF₃).
• Borax, Na₂[B₄O₅(OH)₄]·8H₂O: gives the borax bead test with transition-metal oxides.
• Boric acid, B(OH)₃: a weak monobasic Lewis acid — it accepts OH⁻ rather than donating H⁺: B(OH)₃ + H₂O → [B(OH)₄]⁻ + H⁺.
• Diborane, B₂H₆: four terminal B–H bonds (2c–2e) and two bridging B–H–B bonds (3c–2e, "banana bonds"). Each B is sp³ hybridised.
• Aluminium forms a protective oxide layer; Al₂O₃ is amphoteric.

4. Group 14: Carbon Family (C, Si, Ge, Sn, Pb)

• Oxidation states +4 and +2; +2 becomes more stable down the group (Pb²⁺). SnCl₂ is a reducing agent; PbO₂ is an oxidising agent.
• Catenation (self-linking): C ≫ Si > Ge ≈ Sn; Pb does not catenate. C–C bond is strong.
• Allotropes of carbon: diamond (sp³, 3D network, hardest, non-conductor), graphite (sp², layers, conducts electricity, lubricant), fullerenes (C₆₀, sp², soccer-ball shape with 20 hexagons and 12 pentagons).
• CO is a neutral, poisonous gas that binds to haemoglobin; CO₂ is acidic and linear.
• SiO₂ is a 3D network solid (each Si bonded tetrahedrally to four O). Silicones are polymers with –R₂SiO– units; they are water-repellent and heat-stable.
• CCl₄ is not hydrolysed (C has no d orbitals), but SiCl₄ is readily hydrolysed to Si(OH)₄.

5. Group 15: Nitrogen Family (N, P, As, Sb, Bi)

• Oxidation states −3 to +5. N₂ is very inert (N≡N bond energy 941 kJ mol⁻¹).
• Hydrides EH₃: basic strength and thermal stability decrease NH₃ > PH₃ > AsH₃ > SbH₃ > BiH₃; reducing power increases. Bond angle decreases: NH₃ (107°) > PH₃ (93.6°). Boiling point: NH₃ is high due to H-bonding, but SbH₃ has the highest.
• Ammonia is made by the Haber process (Fe catalyst, Mo promoter, about 200 atm, 700 K). Nitric acid by the Ostwald process (Pt/Rh gauze catalyst oxidises NH₃ to NO).
• Allotropes of phosphorus: white (P₄ tetrahedron, very reactive, glows in the dark, stored under water), red (polymeric, less reactive), black (most stable).
• PCl₅ is trigonal bipyramidal in gas; in the solid it exists as [PCl₄]⁺[PCl₆]⁻.

6. Group 16: Oxygen Family (O, S, Se, Te, Po)

• Hydrides H₂E: acidic strength increases H₂O < H₂S < H₂Se < H₂Te (weaker H–E bond); thermal stability decreases. H₂O has the highest boiling point (H-bonding).
• Ozone, O₃: bent molecule, a powerful oxidising agent; the ozone layer absorbs UV radiation.
• Sulphur: rhombic (α, stable at room temperature) and monoclinic (β, above 369 K); both contain S₈ puckered crown rings.
• Sulphuric acid by the contact process: 2SO₂ + O₂ ⇌ 2SO₃ using V₂O₅ catalyst; SO₃ is absorbed in H₂SO₄ to give oleum (H₂S₂O₇), which is then diluted.
• Oxoacids of S: H₂SO₃, H₂SO₄, H₂S₂O₇ (pyrosulphuric, S–O–S), H₂SO₅ (Caro's acid, peroxide), H₂S₂O₈ (Marshall's acid, peroxide link).

7. Group 17: Halogens (F, Cl, Br, I, At)

• Most reactive non-metals. Oxidising power: F₂ > Cl₂ > Br₂ > I₂. F shows only −1.
• Bond dissociation enthalpy: Cl₂ > Br₂ > F₂ > I₂ (F₂ is low because of lone-pair repulsion in a short bond).
• Acid strength of HX: HF < HCl < HBr < HI (bond strength decreases). HF is weak due to the strong H–F bond and H-bonding.
• Acid strength of oxoacids of Cl: HOCl < HClO₂ < HClO₃ < HClO₄ (more O atoms stabilise the anion).
• Interhalogen compounds: XX′ (ClF, linear), XX′₃ (ClF₃, T-shaped), XX′₅ (BrF₅, square pyramidal), XX′₇ (IF₇, pentagonal bipyramidal). They are more reactive than halogens (except F₂) because the X–X′ bond is weaker.
• Chlorine bleaches by oxidation (in presence of moisture): permanent. SO₂ bleaches by reduction: temporary.

8. Group 18: Noble Gases

• Stable ns² np⁶ configuration (He: 1s²); very low reactivity, high IE, positive electron gain enthalpy.
• Neil Bartlett made O₂⁺[PtF₆]⁻ and, noting that Xe has nearly the same IE as O₂, made the first noble gas compound Xe⁺[PtF₆]⁻.
• Xenon fluorides and shapes: XeF₂ (linear, sp³d), XeF₄ (square planar, sp³d²), XeF₆ (distorted octahedral, sp³d³). XeO₃ is pyramidal; XeOF₄ is square pyramidal.
• Uses: He in balloons and diving gas (low solubility in blood); Ar in bulbs and welding; Ne in advertising signs.

9. Master Table: Shapes of Key p-Block Molecules

10. Mind Map

11. Tips and Tricks

12. Important One-liners

• Diborane has two three-centre two-electron bonds.
• Graphite conducts electricity because of delocalised π electrons.
• White phosphorus is stored under water because it catches fire in air.
• Oleum is H₂S₂O₇ (fuming sulphuric acid).
• Fluorine has a lower bond dissociation enthalpy than chlorine.
• XeF₄ is square planar with two lone pairs on Xe.

13. PYQ-pattern Questions with Solutions

These questions are original and modelled on concepts frequently tested in JEE Main.

14. Quick Revision Summary