p-Block Elements

Group 13: B, Al, Ga, In, Tl. Electronic config: ns² np¹.

Oxidation states: +3 mainly. Tl shows +1 (inert pair effect — 6s² electrons reluctant to participate in bonding). Down the group, +1 becomes more stable.

BORON — non-metal; unique in its group.

Anomalous properties (small size, high IE, no d-orbitals):

• Doesn't form B³⁺ ion easily (covalent compounds).
• Forms electron-deficient compounds (BF₃, BCl₃, B₂H₆).
• Maximum covalency 4 (no d-orbital expansion).

Important compounds:

Borax (Na₂B₄O₇ · 10H₂O):

• Important industrial source of boron compounds.
• Heated: → Na₂B₄O₇ (anhydrous) → 2NaBO₂ + B₂O₃.
• Borax bead test: transparent bead with characteristic colors based on metal ions. Cu → blue; Cr → green; Mn → violet.

Orthoboric acid (H₃BO₃ or B(OH)₃):

• White, slippery to touch.
• Weak monobasic acid — but not because it releases H⁺. Instead, it accepts OH⁻ from water:
  B(OH)₃ + H₂O ⇌ [B(OH)₄]⁻ + H⁺.
• Used as antiseptic (eye drops).

Diborane (B₂H₆):

• Electron-deficient. Has 3-center 2-electron bonds (banana bonds).
• Reacts with NH₃ giving "inorganic benzene" (B₃N₃H₆ - borazine).

Boron carbide (B₄C):

• One of the hardest known materials. Used in bulletproof vests, tank armor.

ALUMINUM — most abundant metal in Earth's crust.

Occurrence: mainly as bauxite (Al₂O₃·xH₂O). Pure form: extracted via Bayer process (dissolve in NaOH) then Hall-Héroult process (electrolysis of molten Al₂O₃ in cryolite Na₃AlF₆).

Diagonal relationship of Al with Be (both small, high charge density):

• Both form amphoteric oxides.
• Both react with water to form basic hydroxides + H₂.

Important compounds:

Aluminum chloride (AlCl₃):

• In gas phase / non-polar solvents: dimeric (Al₂Cl₆) with bridging Cl atoms.
• In water: ionizes completely (Al³⁺ + 3Cl⁻).
• Lewis acid — accepts electron pairs. Catalyst in Friedel-Crafts alkylation/acylation.

Aluminum oxide (Al₂O₃, alumina):

• Amphoteric — dissolves in both acids and bases:
  Al₂O₃ + 6 HCl → 2 AlCl₃ + 3 H₂O.
  Al₂O₃ + 2 NaOH + 3 H₂O → 2 Na[Al(OH)₄].

Aluminum sulfate / alum (KAl(SO₄)₂·12H₂O):

• Double salt. Used as water purifier (coagulates colloidal impurities).

Important reaction — Thermite reaction:
2 Al + Fe₂O₃ → Al₂O₃ + 2 Fe + heat.
Highly exothermic; used for welding railway tracks and military ammunition.

Why Al is more reactive than Fe but doesn't corrode?

Al forms a thin, adherent layer of Al₂O₃ on its surface that prevents further oxidation. Iron's oxide (rust) is flaky and falls off, exposing fresh metal.

Down the group trends:

• Atomic radius increases: B < Al < Ga ≈ In < Tl. Note Ga ≈ Al because of poor shielding by 3d electrons.
• Density: B (2.3) < Al (2.7) << Ga (5.9) — anomalous jump because Ga starts the post-transition series.
• Inert pair effect: stability of lower oxidation state increases down. Tl(I) is more stable than Tl(III).
• Acidic-basic character of oxides: B₂O₃ (acidic) > Al₂O₃ (amphoteric) > Ga₂O₃, In₂O₃ (amphoteric) > Tl₂O (basic).

Group 14: C, Si, Ge, Sn, Pb. Electronic config: ns² np².

Oxidation states: +4 and +2 both common. Inert pair effect down the group:

• C, Si: +4 dominates.
• Ge: both +2 and +4 known; +4 more stable.
• Sn: +2 and +4 comparable.
• Pb: +2 more stable than +4 (inert pair effect).

ANOMALOUS PROPERTIES OF CARBON:

1. Small size + high IE → forms strong covalent bonds (not ionic).
2. Catenation — exceptional ability to bond with itself in chains/rings of any length. Source of >10 million organic compounds.
3. Can form double and triple bonds with itself and with N, O.
4. Tetravalency + sp³/sp²/sp hybridization → diverse 3D structures.
5. Si shows limited catenation (Si-Si bond weak — 220 kJ/mol vs C-C 348 kJ/mol). Si-O bonds are stronger than C-O, so silicon prefers Si-O-Si chains (silicates).

ALLOTROPES OF CARBON

Diamond:

• sp³ hybridization. Each C bonded tetrahedrally to 4 others.
• Hardest natural substance. Very high melting point (~3550°C).
• Excellent thermal conductor (better than copper at room T).
• Electrical insulator (no free electrons).
• Refractive index 2.42 → brilliance in jewelry.

Graphite:

• sp² hybridization. Hexagonal layers held by weak van der Waals.
• Slippery — layers slide easily → used as lubricant.
• Electrical conductor (delocalized electrons in layers).
• Used in pencils ("lead" is graphite + clay), electrodes.

Fullerene (C₆₀, buckyball, 1985 Nobel 1996):

• Hollow cage of 60 carbons, like a soccer ball — 20 hexagons + 12 pentagons.
• All atoms sp² hybridized.
• Crystalline solid, soluble in some organics.
• Father of nanotechnology.

Other allotropes:

• Graphene (single layer of graphite, 2004 Nobel 2010): strongest 2D material known. Conducts electricity better than copper.
• Carbon nanotubes: rolled-up graphene sheets. Used in advanced composites.

OXIDES OF CARBON

CO (carbon monoxide):

• Neutral, colorless, odorless.
• Highly toxic — binds Hb 200× more than O₂ → blocks O₂ transport.
• Reducing agent in industry: 3CO + Fe₂O₃ → 2Fe + 3CO₂ (steelmaking).

CO₂ (carbon dioxide):

• Acidic oxide. Greenhouse gas.
• Dissolves in water → carbonic acid: CO₂ + H₂O ⇌ H₂CO₃.
• Solid CO₂ = "dry ice" — sublimes at −78°C without melting.

SILICON: the "rock" of Group 14

Second most abundant element in Earth's crust (~28%). Forms basis of:

• Quartz (SiO₂) — crystalline.
• Silicates — building blocks of rocks: feldspar, mica, asbestos.
• Clay — hydrated aluminum silicates.
• Cement, glass, ceramics.

Silicates are tetrahedral SiO₄ units linked in various ways:

• Single tetrahedra → orthosilicate.
• Linear chains (pyroxenes, asbestos).
• Sheets (mica).
• 3D networks (quartz, feldspar).

Cement: mixture of CaO, SiO₂, Al₂O₃, Fe₂O₃. Hardens with water.

Glass: non-crystalline (amorphous) solid. Soda-lime glass = SiO₂ + Na₂O + CaO. Pyrex = borosilicate (more thermal-shock resistant).

SILICONES — synthetic polymers with -Si-O-Si- backbone.

General formula: (R₂SiO)_n where R is methyl or other group.

Preparation: RnSiCl_(4-n) hydrolysis → Si-OH → condensation.

Properties:

• Thermally stable up to 300°C.
• Water-repellent (hydrophobic).
• Electrically insulating.
• Chemically inert.
• Range from liquids (oils) to elastomers (rubber) to resins.

Uses:

• Lubricants and greases.
• Sealants (bathroom caulk).
• Waterproofing fabrics.
• Medical implants (silicone breast implants, contact lenses).
• High-voltage electrical insulation.

TIN AND LEAD

Tin (Sn):

• Three allotropes: white (β, common), grey (α, brittle, below 13°C — "tin pest"), rhombic (γ).
• Resistant to corrosion → used to coat steel cans (tinplate).

Lead (Pb):

• High density (11.3 g/cm³), low melting (327°C).
• Toxic — replaces Ca in bones, damages nervous system.
• Lead poisoning historically from water pipes (Roman aqueducts), gasoline (now phased out — leaded petrol).
• Modern uses: batteries (Pb-acid), shielding (X-rays, gamma).

Tetraethyl lead (TEL): Pb(C₂H₅)₄ — formerly added to gasoline as anti-knock. Banned worldwide due to lead pollution. India phased out in 2000.

Carbon vs Silicon vs Tin/Lead — periodic trends visible:

Halogens = Group 17: F, Cl, Br, I, At (radioactive). Outermost configuration: ns²np⁵ — one electron short of noble gas.

Physical states at room T:

• F₂: pale yellow gas
• Cl₂: greenish-yellow gas
• Br₂: red-brown liquid (only halogen liquid at RT)
• I₂: dark purple solid (sublimes)

Trends down the group:

Why F is anomalous:

1. Small size + high electron density → repulsion lowers EA.
2. Weak F-F bond (159 kJ/mol vs Cl-Cl 243 kJ/mol).
3. Only −1 oxidation state (no d-orbitals in 2nd shell for higher OS).
4. Stronger oxidizer than expected from EA (because hydration is very exothermic).
5. HF is weak acid (others are strong) due to strong H-F bond.
6. HF forms hydrogen bonds → high BP, polymeric (HF)_n in liquid phase.

Oxidation states:

• F: only −1.
• Cl, Br, I: −1, 0, +1, +3, +5, +7 (use of d-orbitals from 3rd shell).

Important compounds:

HCl:

• Made by NaCl + conc H₂SO₄ → NaHSO₄ + HCl (lab).
• Strong acid; muriatic acid commercially.

Bleaching powder (CaOCl₂): Cl₂ + Ca(OH)₂. Slowly releases Cl₂ on hydrolysis → bleach, disinfect.

Interhalogen compounds (XX', XX'_3, XX'_5, XX'_7):

• ClF, ICl, IF₅, IF₇, BrF₃.
• More reactive than parent halogens (weaker bonds).

Oxoacids of chlorine:

Acidity increases with O count (more electron-withdrawal → more H acidic).

Reactions to know:

1. Halogen + metal → halide.
2 Na + Cl₂ → 2 NaCl.

2. Halogen displaces lower halogens.
Cl₂ + 2 KBr → 2 KCl + Br₂.
F₂ > Cl₂ > Br₂ > I₂ as oxidizing agents.

3. Halogen + H₂ → HX.
F₂ reacts explosively in dark; Cl₂ needs light; Br₂ slow at high T; I₂ reversible/incomplete.

4. Cold dilute alkali: disproportionation.
Cl₂ + 2 NaOH → NaCl + NaOCl + H₂O.

5. Hot conc alkali.
3 Cl₂ + 6 NaOH → 5 NaCl + NaClO₃ + 3 H₂O.

6. With ammonia.
3 Cl₂ + 8 NH₃ (excess) → 6 NH₄Cl + N₂. (limited NH₃ → NCl₃ — explosive)

Test for halide ions (Cl⁻, Br⁻, I⁻):

• Add AgNO₃ solution → coloured precipitate:
  • Cl⁻ → white AgCl, soluble in NH₃
  • Br⁻ → pale yellow AgBr, partially soluble in NH₃
  • I⁻ → yellow AgI, insoluble in NH₃