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🧪 Chemistry · Class 11 · NEET & JEE

p-Block Elements (Groups 13 & 14)

Covers the Boron family (Group 13: B, Al) and the Carbon family (Group 14: C, Si, Ge, Sn, Pb). Key topics include Lewis acid behaviour of BF₃, borax bead test, silicates, silicones, allotropes of carbon, and trends in properties down the group.

Practice p-Block Elements (Groups 13 & 14) Quiz — 100% Free →

Reading time~11 min

Revision time~4 min

Last updated2026-07-19

1 Read the chapter ~11 min

🎯 Key Points

p-block spans Groups 13-18: boron family, carbon family, nitrogen family, oxygen family, halogens, noble gases
First element of each group is anomalous (B≠Al, C≠Si, N≠P, O≠S, F≠Cl) — usually due to small size and no available d-orbitals
Reactivity/oxidising power of halogens: F₂ > Cl₂ > Br₂ > I₂
Maximum covalency: N=4 (no d-orbitals), P/As/Sb/Bi can expand to 5/6 using d-orbitals
Carbon shows strong catenation + stable π bonds; Silicon catenates poorly and rarely forms π bonds
Noble gases are inert due to fully-filled valence shells; only Xe/Kr form well-characterised compounds (with F, O)

Diamond's rigid 3D tetrahedral network makes it the hardest natural material, while graphite's flat hexagonal sheets are held together only by weak forces, letting them slide past each other (used as a lubricant and in pencils).

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

Nitrogen: Makes up 78% of air; forms NH₃ by Haber process
HNO₃ (nitric acid): Strong oxidising acid; used in fertilisers and explosives
Phosphorus allotropes: white P (toxic), red P (stable), black P
PCl₃ and PCl₅ are important chlorides of phosphorus

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

Oxygen: Most electronegative element after F; essential for life and combustion
Ozone (O₃): Allotrope of oxygen; absorbs harmful UV radiation in the stratosphere
Sulphur allotropes: rhombic (stable at room temp), monoclinic
H₂SO₄ (sulphuric acid): Most important industrial chemical; strong diprotic acid

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

Most electronegative non-metals; one electron short of noble gas configuration
Reactivity and oxidising power: F₂ > Cl₂ > Br₂ > I₂
Cl₂ used in water purification; bleaching powder = Ca(OCl)Cl
Interhalogen compounds: ClF, BrF₃, IF₅, IF₇

Group 18: Noble Gases (He, Ne, Ar, Kr, Xe)

Completely filled outer shells; chemically inert under normal conditions
He used in balloons and diving tanks; Ar in welding; Ne in sign boards
Xenon compounds (XeF₂, XeF₄, XeOF₄) are the most stable noble gas compounds

Quick Tips

F₂ is the strongest oxidising agent among all elements
Cl₂ can displace Br₂ and I₂ from their salt solutions (higher reactivity)
Noble gases: He has the lowest boiling point of any substance (-269°C)

Anomalous Behaviour of the First Element in Each Group

Boron (Group 13) differs from Al: B is a metalloid/non-metal while Al is a metal; B forms only covalent compounds and shows a maximum covalency of 4 (e.g., BF₄⁻)
Carbon (Group 14) differs from Si: C shows strong catenation and forms stable pi bonds (multiple bonds), while Si rarely forms stable pi bonds due to larger atomic size
Nitrogen (Group 15) differs from P: N₂ is a stable diatomic gas with a triple bond, while phosphorus exists as reactive P₄ tetrahedral molecules; N has a maximum covalency of 4 (no d-orbitals), P can expand to 5 or 6
Oxygen (Group 16) differs from S: O forms strong pi bonds, giving diatomic O₂, while S prefers catenation forming S₈ rings
Fluorine (Group 17) differs from Cl: F shows only -1 oxidation state (no d-orbitals), while Cl can show +1 to +7 states; F has an anomalously low bond dissociation energy due to lone pair-lone pair repulsion in the small F₂ molecule

Important Compounds: Borax and Silicones

Borax (Na₂B₄O₇.10H₂O): On heating, swells and forms a transparent glassy bead of sodium metaborate and B₂O₃ (used in the borax bead test to identify coloured metal ions)
Boric acid (H₃BO₃): A weak monobasic Lewis acid; acts as an acid by accepting OH⁻ from water rather than donating H⁺
Silicones: Synthetic polymers with a repeating -Si-O-Si-O- backbone and organic side groups; water repellent, heat resistant, used as sealants, lubricants, and in waterproofing fabrics
Silicates and zeolites: Built from SiO₄ tetrahedra; zeolites (porous aluminosilicates) are used as molecular sieves and catalysts in petrochemical cracking

Allotropes Recap with Structural Detail

Carbon: Diamond (sp³, tetrahedral network, hardest natural substance, poor conductor); graphite (sp², hexagonal layers, good conductor, lubricant); fullerene (C₆₀, soccer-ball-shaped, sp²)
Phosphorus: White P (P₄ tetrahedra, reactive, glows in dark, kept under water); red P (polymeric, less reactive, non-poisonous); black P (most stable, layered structure)
Sulphur: Rhombic (alpha-sulphur, stable form, S₈ crown-shaped rings); monoclinic (beta-sulphur, stable above 369 K)

Oxoacids and Hydrides Trends

Acidic strength of hydrides decreases down a group for the same family in some cases, but increases across a period: NH₃ < PH₃ < AsH₃ (basicity decreases down group 15 instead)
Oxoacids of chlorine in increasing acidic and oxidising strength order: HOCl < HOClO < HOClO₂ < HOClO₃ (more oxygen atoms stabilise the conjugate base better through resonance)

General Trends in Groups 13 and 14

Electronic configuration: Group 13 is ns²np¹ (max oxidation state +3); Group 14 is ns²np² (max oxidation state +4)
Atomic radius: Increases down each group, but Ga is anomalously smaller than Al (poor shielding by intervening 3d electrons); similarly the size increase from Si to Pb is not regular
Ionisation enthalpy: Decreases down the group overall, though small increases appear (e.g. Ga, Tl) due to poor d/f-electron shielding
Electronegativity: Decreases from B to Al, then stays roughly constant; Group 14 elements are slightly more electronegative than Group 13 in the same period
Inert pair effect: The stability of the lower oxidation state (+1 for Group 13, +2 for Group 14) increases down the group. So Tl⁺ is more stable than Tl³⁺, and Pb²⁺ is more stable than Pb⁴⁺; consequently PbO₂ and Tl₂O₃ are strong oxidising agents
Metallic character: Increases down each group — B is a metalloid while Al, Ga, In, Tl are metals; C is a non-metal, Si and Ge are metalloids, Sn and Pb are metals

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

Anomalous behaviour of boron: Small size, high ionisation enthalpy and absence of d-orbitals restrict B to a maximum covalency of 4 (e.g. BF₄^-); B is a non-metallic metalloid forming only covalent compounds, unlike the metallic rest of the group
Electron deficiency and Lewis acidity: BX₃ molecules have only 6 electrons around B (incomplete octet), so they are strong Lewis acids that accept a lone pair to complete the octet
Amphoteric nature of aluminium: Al reacts with both acids (2Al + 6HCl gives 2AlCl₃ + 3H₂) and alkalis (2Al + 2NaOH + 2H₂O gives 2NaAlO₂ + 3H₂), so Al₂O₃ and Al(OH)₃ are amphoteric
Reactions of aluminium: Forms a thin protective oxide film (passivation) that resists further corrosion; reacts with dilute mineral acids to give H₂, but concentrated HNO₃ renders it passive; used in the thermite process (2Al + Fe₂O₃ gives Al₂O₃ + 2Fe) for welding rails
Aluminium chloride: Exists as the dimer Al₂Cl₆ (chlorine-bridged) in the vapour and non-polar solvents, relieving Al's electron deficiency

Diborane and Boron Hydrides

Boranes: Boron forms a series of covalent hydrides such as diborane (B₂H₆), tetraborane (B₄H₁₀) and higher members; they are electron-deficient molecules
Structure of diborane (B₂H₆): Contains four terminal B-H bonds (normal 2-centre-2-electron bonds) in a plane and two bridging B-H-B bonds; each bridge is a 3-centre-2-electron "banana" bond in which two electrons are shared among three atoms
Diborane has 12 valence electrons but would need more for eight conventional bonds, so the two bridging bonds are the hallmark of its electron-deficient bonding
Preparation: 3LiAlH₄ + 4BF₃ gives 2B₂H₆ + 3LiF + 3AlF₃; diborane catches fire spontaneously in air, releasing large amounts of energy

Structure of diborane B2H6: two boron atoms each bonded to two terminal hydrogens and bridged by two hydrogen atoms through three-centre two-electron banana bonds.

Diborane (B₂H₆) is electron-deficient: each boron carries two normal terminal B–H bonds, and the two borons are held together by two bridging hydrogens through three-centre two-electron “banana” bonds. Image: Grajkowf, CC BY-SA 4.0, via Wikimedia Commons.

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

Catenation: The tendency to form long chains/rings through self-linking is strongest for carbon (very strong C-C bonds, ~348 kJ/mol) and falls sharply down the group (Si-Si is weaker); this underlies the vast number of organic compounds
Tendency to form pi bonds: Carbon forms stable p pi-p pi multiple bonds (C=C, C=O), whereas the heavier elements rarely do because of their larger size and poor orbital overlap
Allotropes of carbon: Diamond (sp³ 3D network, hardest natural solid), graphite (sp² layers, conductor and lubricant), and fullerene (C₆₀ "buckyball", sp² closed cage) — see the allotropes section above for structural detail
Oxidation states: All show +4, but the +2 state grows more stable down the group owing to the inert pair effect; Sn(II) is a reducing agent while Pb(IV) is an oxidising agent
Silicates and silicones: Silicates are built from SiO₄^4- tetrahedra linked in chains, sheets or 3D networks; zeolites are porous aluminosilicates used as molecular sieves and cracking catalysts; silicones are water-repellent -Si-O-Si- polymers (covered above)

Oxides of Carbon: CO and CO₂

Carbon monoxide (CO): A neutral oxide formed by incomplete combustion of carbon; highly poisonous because it binds haemoglobin ~300 times more strongly than O₂, forming carboxyhaemoglobin and blocking oxygen transport; a good reducing agent used in metallurgy (Fe₂O₃ + 3CO gives 2Fe + 3CO₂)
Carbon dioxide (CO₂): A weakly acidic oxide (gives carbonic acid with water); a linear non-polar molecule; produced by respiration, combustion and fermentation; a major greenhouse gas whose rising level drives global warming
Solid CO₂ ("dry ice") sublimes directly and is used as a refrigerant

🚀 JEE Advanced Edge

Back-bonding strength order: In boron trihalides (BX₃), the Lewis acidic strength order is BF₃ < BCl₃ < BBr₃ < BI₃ — the OPPOSITE of what electronegativity alone would predict. This is because back-bonding (halogen lone pair donating into boron's empty p-orbital) is strongest for F (best orbital size match with B) and weakest for I, so BF₃'s electron deficiency is most "patched up," making it the weakest Lewis acid of the four.

Inert pair effect in stability of oxidation states: Down Group 15, the +5 state becomes progressively less stable relative to +3 (e.g., BiCl₅ is far less stable than BiCl₃) because the heavier elements' ns² electron pair resists getting involved in bonding — this is why Bi(V) compounds like NaBiO₃ are actually strong oxidising agents (they "want" to drop down to the more stable +3 state).

Worked problem: Explain why XeF₂ is linear despite Xe having 3 lone pairs. Approach: Xe has 8 valence electrons; 2 are used in Xe-F bonds, leaving 3 lone pairs. Total electron pairs = 5 → sp³d hybridisation, trigonal bipyramidal electron geometry. The 3 lone pairs occupy the 3 equatorial positions (where repulsion is minimised), forcing the 2 bonding pairs into the axial positions — giving a linear molecular shape.

2 Revise ~4 min before the exam

📐 Formula Sheet

Group 13 (B, Al…): ns²np¹, mainly +3; the inert-pair effect makes +1 more stable down the group (Tl⁺)
Group 14 (C, Si…): ns²np², +4 and +2; +2 becomes more stable down (Pb²⁺ > Pb⁴⁺)
Inert-pair effect: the ns² electrons resist bonding down a group, favouring the lower oxidation state
Boron anomaly: only non-metal in group 13; forms electron-deficient compounds (BF₃, diborane B₂H₆)
Catenation: carbon > silicon ≫ others (strong C–C bonds); the basis of organic chemistry
Allotropes of carbon: diamond (sp³, hard, insulating), graphite (sp², conducting), fullerene (C₆₀)
Key compounds: borax Na₂B₄O₇·10H₂O, alum, CO, CO₂, silicates, silicones
CO vs CO₂: CO is a strong reducing agent and toxic (binds haemoglobin); CO₂ is acidic and a greenhouse gas

3 Practice apply it

✍️ Worked Examples

Example 1 — Inert-pair effect
Q: Why is Pb²⁺ more stable than Pb⁴⁺, but Sn⁴⁺ more stable than Sn²⁺?
Step 1 — The inert-pair effect strengthens down a group, as the ns² electrons become increasingly reluctant to bond.
Step 2 — Lead is below tin, so its 6s² pair is especially inert, favouring the +2 state.
Step 3 — Tin, higher up, still readily uses all four valence electrons, so +4 dominates.
Answer: the inert-pair effect makes the lower oxidation state more stable further down — hence Pb²⁺ but Sn⁴⁺. Note: this is why PbO₂ is a strong oxidiser (it reverts to the stable +2).

Example 2 — Electron deficiency of boron
Q: Why does BF₃ act as a Lewis acid?
Step 1 — Boron has only 3 valence electrons, so in BF₃ it has just 6 around it — an incomplete octet.
Step 2 — The empty 2p orbital can accept an electron pair.
Step 3 — Accepting a lone pair (e.g. from NH₃) is exactly what a Lewis acid does.
Answer: its incomplete octet lets BF₃ accept an electron pair, making it a Lewis acid. Note: this drives its use as a catalyst in many organic reactions.

Example 3 — Catenation
Q: Why does carbon show far more catenation than silicon?
Step 1 — Catenation depends on the strength of element–element bonds.
Step 2 — The C–C bond (~348 kJ/mol) is much stronger than Si–Si (~226 kJ/mol), because carbon is small and its orbitals overlap well.
Step 3 — Strong, stable C–C chains form readily; weaker Si–Si chains do not.
Answer: carbon's stronger self-bond allows extensive catenation — the foundation of organic chemistry.