p-Block Elements (Groups 13 to 18)

1. General p-Block Trends and First-Row Anomalies

• Location and configuration: Groups 13–18, valence configuration ns²np^1–6. Only block containing metals, metalloids and non-metals together. Maximum oxidation state = (Group number − 10): Group 13 → +3, Group 14 → +4, Group 15 → +5, Group 16 → +6, Group 17 → +7.
• Inert pair effect: down Groups 13–15, the ns² pair increasingly resists bonding (poor shielding by intervening d/f electrons), so the (Group − 2) oxidation state becomes progressively MORE stable in heavier members: Tl⁺ > In⁺ > Ga⁺; Pb²⁺ > Sn²⁺ > Ge²⁺; Bi³⁺ > Sb³⁺ > As³⁺.
• Why the first element of each group (B, C, N, O, F) is anomalous: (1) exceptionally small atomic size, (2) high electronegativity, (3) high ionization enthalpy, (4) NO vacant d-orbitals in the valence shell, so maximum covalency is capped at 4 (N never shows covalency 6; nitrogen belongs to Period 2 and simply cannot expand its octet). Consequences you will see repeatedly below: N forms no pentahalide (NCl₅ doesn't exist, NF₃ is the only stable trihalide), CCl₄ resists hydrolysis while SiCl₄ hydrolyses readily, and O/N/C form strong pπ–pπ multiple bonds WITH THEMSELVES (O=O, N≡N, C=C) that heavier congeners cannot — S₈, not S=S, is sulfur's stable form because 3pπ–3pπ self-bonding is too weak. Sulfur CAN still form pπ–pπ bonds with a small, electronegative partner like C or N (C=S, or the resonance-stabilised thiocyanate S=C=N⁻ ↔ S^⊖–C≡N) — it is only S–S self-bonding that fails, not S bonding to small atoms in general.
• Diagonal relationship carries over from s-block logic: B–Si is the Group-13/14-relevant diagonal pair (both metalloid-like, both form covalent halides and complex anions with F⁻: [BF₄]⁻, [SiF₆]²⁻).

2. Group 13: The Boron Family

• Elements: B (non-metal), Al, Ga, In, Tl (metals) — ns²np¹. B has an exceptionally high melting point (2453 K, rigid B₁₂ lattice); Ga has an unusually LOW melting point (303 K, liquid in summer) but a very high boiling point, so it is used in high-temperature thermometers.
• The four Group 13 anomalies (all driven by poor 3d shielding in Ga and poor 4f shielding in Tl): atomic radius B < Ga < Al < In < Tl; IE₁ order B > Tl > Ga > Al > In; electronegativity B > Tl > In > Ga > Al; stability of +1 state rises Al⁺ < Ga⁺ < In⁺ < Tl⁺ (Tl⁺ is the everyday Tl species; Tl³⁺ is a strong oxidizer).
• Oxidation states and oxides: B, Al show only +3. Ga, In, Tl show both +1 and +3. Oxide basicity rises down the group: B₂O₃ acidic, Al₂O₃/Ga₂O₃ amphoteric, In₂O₃/Tl₂O₃ basic.
• Lewis acidity of BX₃: electron-deficient (only 6 valence electrons around B), so BX₃ compounds act as Lewis acids. Strength order BI₃ > BBr₃ > BCl₃ > BF₃, opposite to what halogen electronegativity alone predicts, because back-bonding (halogen lone pair → boron's empty 2p orbital) is strongest for the small, close-fitting F, which most effectively quenches B's electron deficiency. BF₃ forms a stable adduct with NH₃ (BF₃·NH₃) and is trigonal planar (sp²).
• BCl₃ hydrolysis: BCl₃ is covalent and hydrolyses completely in neutral water to B(OH)₃ + 3HCl (the tetrahedral [B(OH)₄]⁻ ion forms only in BASIC medium, and boron never forms an aquo-cation like [B(H₂O)₆]³⁺ since it is not a transition metal).
• Diborane B₂H₆: electron-deficient, held together by two 3-centre-2-electron ("banana") B–H–B bridge bonds plus four normal 2-centre-2-electron terminal B–H bonds.

• Borax and the borax bead test: borax is Na₂B₄O₇·10H₂O. On heating it loses water and melts into a glassy bead of sodium metaborate + boric anhydride: Na₂B₄O₇ → 2NaBO₂ + B₂O₃. B₂O₃ then combines with transition-metal-oxide impurities to give characteristic coloured beads: Cu → blue Cu(BO₂)₂ in the non-luminous (oxidizing) flame, which is reduced by carbon in the luminous flame to colourless/red Cu(I) metaborate CuBO₂ (Cu goes +2 → +1); Co gives a dark blue bead, Cr gives a green bead.
• Alums: general formula M₂SO₄·M′₂(SO₄)₃·24H₂O. M⁺ = Na⁺, K⁺, NH₄⁺ (Li⁺ is too small and cannot form an alum); M′³⁺ = Al³⁺, Fe³⁺, Cr³⁺.

3. Group 14: The Carbon Family

• Elements: C, Si, Ge, Sn, Pb — ns²np². Covalent radius rises sharply C→Si, then only slightly further down (poor d/f shielding).
• +4 vs +2 stability (inert pair effect, the single most-tested Group 14 fact): for C and Si only +4 is stable (no stable +2 state). Down the group +4 stability falls and +2 rises: Ge⁴⁺ > Ge²⁺, Sn⁴⁺ > Sn²⁺, but the trend flips at Pb: Pb²⁺ > Pb⁴⁺. Consequently PbO₂ (Pb⁴⁺) is a strong oxidizing agent (readily falls back to Pb²⁺) while SnCl₂ (Sn²⁺) is a strong reducing agent (readily rises to Sn⁴⁺): 2HgCl₂ + SnCl₂ → Hg₂Cl₂↓ (white) + SnCl₄, and excess SnCl₂ reduces this further to black Hg.
• Δ_rG° sign rule for these reactions: PbO₂ + Pb → 2PbO is spontaneous (Δ_rG° < 0) because it converts the unstable Pb⁴⁺/Pb⁰ pair into the stable Pb²⁺ product. SnO₂ + Sn → 2SnO is non-spontaneous (Δ_rG° > 0) because it converts stable Sn⁴⁺ into the less-stable Sn²⁺.
• Amphoteric oxide pairs: among [SiO₂, CO₂] (both acidic), [GeO, GeO₂] (both amphoteric), [SnO, SnO₂] (both amphoteric) and [PbO, PbO₂] (PbO amphoteric, but PbO₂ is oxidizing, not simply amphoteric) — exactly 2 pairs have BOTH oxides amphoteric.
• Hydrolysis of tetrahalides: CCl₄ is completely inert to water (carbon has no d-orbitals to expand its octet and accept the incoming OH⁻/H₂O lone pair); SiCl₄, GeCl₄, SnCl₄ all hydrolyse readily using empty d-orbitals.
• Non-existence of PbI₄: Pb⁴⁺ would oxidize the strongly reducing I⁻ before a stable tetraiodide could form, and the Pb–I bond energy released is insufficient to promote a 6s² electron into 6p; so PbI₄ does not exist, even though PbCl₄ and PbBr₄ do.
• Catenation: C >> Si > Ge ≈ Sn; Pb shows essentially none. Governed by C–C bond energy (348 kJ/mol), the strongest single bond among Group 14 self-links.

4. Group 15: The Nitrogen Family

• Elements: N, P, As, Sb, Bi — ns²np³ (extra-stable half-filled configuration → unusually high IE₁ for the group, especially N at 1402 kJ/mol).
• N is anomalous: absence of d-orbitals caps N's covalency at 4 ([NH₄]⁺), so N forms no pentahalide (unlike P, As, Sb which form both EX₃ and EX₅). Among nitrogen trihalides only NF₃ is stable; NCl₃, NBr₃, NI₃ are unstable/explosive and hydrolyse readily.
• +5 vs +3 stability down the group: common oxidation states of Group 15 are −3, +3 and +5, but (inert pair effect again) the stability of the +5 state FALLS down the group while +3 rises — Bi⁵⁺ compounds (BiF₅ only) are strong oxidizers and Bi³⁺ is Bismuth's stable everyday state, mirroring Tl/Pb in Groups 13/14.
• Oxide acidity: acidity of E₂O₃ falls down the group as the element becomes more metallic — N₂O₃, P₂O₃ acidic; As₂O₃, Sb₂O₃ amphoteric; Bi₂O₃ basic. Electronegativity likewise falls N > P > As > Sb, so among N/P/As/Sb, N is most electronegative (its oxide most acidic) and Sb is least electronegative (its oxide amphoteric).

• EH₃ hydride trends (N→Bi): thermal stability falls down the group (E–H bond weakens); reducing character rises (BiH₃ strongest reducer); basicity falls (lone-pair density is highest on the small N atom and gets diffused in larger orbitals lower down); bond angle falls sharply from NH₃ (107.8°) to PH₃ (93.6°) onward (heavier hydrides use nearly pure p-orbitals for bonding). Boiling point is the one exception that does NOT rise smoothly down the group: PH₃ < AsH₃ < NH₃ < SbH₃ < BiH₃, because NH₃ is lifted anomalously high by hydrogen bonding, but not high enough to beat SbH₃/BiH₃'s much larger van der Waals mass.
• Ammonia (NH₃): made industrially by the Haber process (N₂ + 3H₂ ⇌ 2NH₃, ~200 atm, 700 K, Fe catalyst). Cannot be dried over conc. H₂SO₄, P₄O₁₀, or anhydrous CaCl₂ (all react chemically with it) — only CaO (quicklime) is used. Forms deep-blue [Cu(NH₃)₄]²⁺ with Cu²⁺, and a soluble [Ag(NH₃)₂]⁺ complex that dissolves AgCl.
• Nessler's reagent test (a favourite PYQ identification tool): NH₄⁺/NH₃ + alkaline K₂[HgI₄] gives a characteristic brown precipitate of Millon's base (H₂N·HgO·HgI). A "p-block element E forming a binary cation EH₄⁺ that gives this brown precipitate, with IE₁ = 1402 kJ/mol" identifies E = Nitrogen.
• Nitric acid and metals: conc. HNO₃ passivates Fe, Cr and Al (protective oxide layer forms, no visible reaction) — the same passivation behaviour you saw for Al in Group 13.

5. Group 16: The Oxygen Family

• Elements: O, S, Se, Te, Po — ns²np⁴. IE₁ of Group 16 is LOWER than the corresponding Group 15 element in the same period (O < N, S < P) despite higher nuclear charge, because Group 15's half-filled np³ is extra stable — the same logic you used for Period 2/3 ionization-energy anomalies in Unit 9.
• Electron gain enthalpy anomaly: S has a MORE negative EGA than O (S > Se > Te > O in magnitude), because O's very compact 2p subshell suffers strong inter-electronic repulsion when accepting an extra electron — the exact same reasoning as Cl > F in Group 17 below.
• H₂E hydride trends (O→Te): acidic character rises down the group (H₂O < H₂S < H₂Se < H₂Te); thermal stability falls; reducing power rises (H₂Te strongest). Boiling point exception: H₂S < H₂Se < H₂Te < H₂O — water's extensive hydrogen bonding lifts its boiling point (373 K) far above even the much heavier H₂Te.
• Hexafluorides: only the hexaFLUORIDES (EF₆) are stable across the group; SF₆ is exceptionally inert because six tightly packed F atoms sterically shield the central S from attack.

6. Group 17: The Halogen Family

• Elements: F, Cl, Br, I — ns²np⁵, smallest radii in their respective periods (highest Z_eff).
• Four classic Group 17 anomalies, all traced to F's small size: electron gain enthalpy magnitude Cl > F > Br > I (compact 2p subshell of F suffers strong electron-electron repulsion on accepting one more electron); bond dissociation energy Cl₂ > Br₂ > F₂ > I₂ (high lone-pair/lone-pair repulsion weakens the very short F–F bond); hydration energy F⁻ > Cl⁻ > Br⁻ > I⁻ (smallest ion, most charge density); standard reduction potential F₂ > Cl₂ > Br₂ > I₂ (F₂ remains the strongest oxidizer overall — its weak F–F bond is more than compensated by F⁻'s huge hydration energy).
• Acid strength trends (HX): thermal stability HF > HCl > HBr > HI (falls with weakening bond); acidic/reducing strength HF < HCl < HBr < HI (rises as bond weakens, HI most easily ionized/oxidized). Boiling point: HCl < HBr < HI < HF (H-bonding lifts HF above even the heavier HI). Melting point: HCl < HBr < HF < HI — here HI (largest, most polarizable, strongest lattice packing) edges out HF in the solid state, even though HF wins on boiling point. Both orders were directly tested together in a 2026 PYQ as a two-statement question.
• Cannot prepare HBr/HI using conc. H₂SO₄: it oxidizes the strongly reducing HBr/HI to Br₂/I₂; use non-oxidizing conc. H₃PO₄ instead.
• Oxoacid acidity (same halogen, rising oxidation state): HClO < HClO₂ < HClO₃ < HClO₄ — more O atoms on the same central halogen stabilize the conjugate base further by resonance, so acidity rises with oxidation state.

7. Group 18: Noble Gases

• Configuration: ns²np⁶ (He: 1s²), closed shell → chemical inertness, very high IE₁, near-zero/positive electron gain enthalpy.
• Bartlett's discovery (1962): IE₁ of Xe (1170 kJ/mol) is close to that of O₂ (1175 kJ/mol); since O₂⁺[PtF₆]⁻ was already known, Bartlett successfully made Xe⁺[PtF₆]⁻, the first noble-gas compound. Only Xe (and to a lesser extent Kr) form a significant range of compounds, exclusively with the most electronegative elements F and O.
• Xenon fluorides — structures (JEE Main tests this table directly, not the reaction mechanisms):

• How they're made (formula-level only): XeF₂ from Xe + F₂ (2:1); XeF₄ from Xe + F₂ (1:5); XeF₆ from Xe + F₂ (1:20, higher pressure). Complete hydrolysis of XeF₆ gives pyramidal XeO₃; partial hydrolysis (1 H₂O) gives square-pyramidal XeOF₄.
• Clathrates: noble-gas atoms trapped inside ice/quinol crystal cages (~6H₂O : 1 gas atom). He and Ne are too small and escape the cavities, so they do NOT form clathrates.

8. How JEE Frames p-Block Questions

• Two-statement / assertion-reason MCQs: the dominant p-block format — one statement states a trend, the other an exception or a compound fact (borax bead test, BCl₃ hydrolysis, hydride boiling points). Verify each statement independently before combining.
• Identify-the-element questions: give a reaction/test (Nessler's reagent, a characteristic precipitate) plus a numeric IE₁/radius value, and ask you to name the element — always cross-check the qualitative test AND the number.
• Amphoteric-oxide-pair counting: list several [MO, MO₂]-type oxide pairs across Group 14 and ask how many pairs are both amphoteric — work each pair individually (SiO₂/CO₂ are NOT amphoteric, they're acidic).
• Reaction-sequence identification (Pb chemistry): a short chain of reactions (PbCl₂ + K₂CrO₄ → A, A + NaOH → B, etc.) testing whether you know Pb²⁺'s common salts (PbCrO₄, [Pb(OH)₄]²⁻) rather than any exotic chemistry.
• Δ_rG° sign questions: given two comproportionation-type reactions (e.g. PbO₂ + Pb → 2PbO vs SnO₂ + Sn → 2SnO), decide spontaneity purely from which oxidation state is more stable for that element — no thermodynamic data needed.

9. Quick Sheet and Checklist

Before the exam, check you can:

• State the Group-number-minus-10 rule and separate "maximum" from "most stable" oxidation state using the inert pair effect.
• Explain why BF3 is the weakest Lewis acid of the BX3 series, and describe the borax bead test's colour and oxidation-state changes for Cu.
• Predict which Group 14 oxide pairs are amphoteric and identify the Sn/Pb +4-vs-+2 stability flip.
• Recall the Group 15 and Group 16 hydride property trends, including which one (boiling point) is NOT monotonic and why.
• Rank halogen EGA, BDE, hydration energy and SRP, and explain why F2 stays the strongest oxidizer despite a weak F-F bond.
• Read off hybridisation, lone pairs and shape for XeF2, XeF4, XeF6, XeO3 and XeOF4 from the structures table.