Group 13

THE p-BLOCK ELEMENTS → Focus: Group 13 (B, Al, Ga, In, Tl)

A. Quick context: what makes the p-block special

• The last electron enters a p orbital. There are 3 p orbitals → up to 6 e⁻ → 6 p-block groups (13–18).

• Valence configuration (except He): ns²np¹-⁶.

• Inner cores differ (some have d and f electrons) → affects radii, ionization enthalpy (IE), electronegativity (χ), bonding, etc.

• Max oxidation state = total valence electrons (s + p). Alternative states often differ by ±2 (inert-pair patterns).

• Non-metallic character decreases down a p-group; the heaviest member is usually most metallic.

• Oxides: non-metal oxides acidic/neutral; metal oxides basic.

• 2nd-period elements (B, C, N, O, F) lack valence d orbitals → max covalence 4; heavier congeners may expand beyond 4 (e.g., [AlF₆] ³⁻).

• π-bonding: strong pπ–pπ for 2nd row (C=C, C≡C, C=O, etc.); heavier rows rely on dπ–pπ / dπ–dπ (weaker).

B. Group 13 at a glance (Boron family)

Members: B, Al, Ga, In, Tl (Nh synthetic; chemistry not established—omitted).
Valence configuration: ns²np¹.
Cores:

• B, Al: noble-gas core

• Ga, In: noble-gas + d¹⁰

• Tl: noble-gas + 4f¹⁴5d¹⁰

Key consequences

• d (and f) poor shielding → d-block (and f-block) contraction, causing Ga radius < Al radius, IE and χ quirks.

• Oxidation states: +3 common for all; +1 grows in stability down the group (inert-pair effect): Al < Ga < In < Tl; Tl(I) dominates, Tl (III) strongly oxidising.

• Bonding: B tends strictly to covalent (B³⁺ not viable in water); Al, Ga, In form covalent anhydrous halides but ionic aquo species in water (large hydration enthalpy).

C. Atomic / physical data

Metallic (atomic) radius (pm) ↑

B (88) < Ga (135) < Al (143) < In (167) < Tl (170)
• Anomaly: Ga < Al (poor 3d shielding → higher Zeff → smaller Ga).

Ionic radius M³⁺ (pm) ↑

B³⁺ (27) < Al³⁺ (53.5) < Ga³⁺ (62.0) < In³⁺ (80.0) < Tl³⁺ (88.5)

Ionic radius M⁺ (pm) ↑

Ga⁺ (120) < In⁺ (140) < Tl⁺ (150)
(B and Al do not commonly form stable M⁺.)

First ionisation enthalpy, IE₁ (kJ·mol⁻¹) ↑

In (558) < Al (577) ≲ Ga (579) < Tl (589) ≪ B (801)
• Irregularities: Ga slightly > Al (3d contraction); Tl > In (4f contraction).

Second ionisation enthalpy, IE₂ (kJ·mol⁻¹) ↑

Al (1816) < In (1820) < Tl (1971) < Ga (1979) ≪ B (2427)

Third ionisation enthalpy, IE₃ (kJ·mol⁻¹) ↑

In (2704) < Al (2744) < Tl (2877) < Ga (2962) ≪ B (3659)

Electronegativity (Pauling) ↑

Al (1.5) < Ga (1.6) < In (1.7) < Tl (1.8) < B (2.0)
• Trend is not perfectly smooth because of d/f-electron effects.

Density at 298 K (g·cm⁻³) ↑

B (2.35) < Al (2.70) < Ga (5.90) < In (7.31) < Tl (11.85)

Melting point (K) ↑

Ga (303) < In (430) < Tl (576) < Al (933) ≪ B (2453)
• Ga anomaly: exceptionally low mp (liquid near room temp).

Boiling point (K) ↑

Tl (1730) < In (2353) < Ga (2676) < Al (2740) ≪ B (3923)
• Ga: despite low mp, bp is high → huge liquid range.

Standard reduction potential E° (M³⁺/M) (V) ↑ (more reducing → less reducing)

Al (−1.66) < Ga (−0.56) < In (−0.34) < Tl (+1.26)
• Indicates: Al is strongly electropositive (forms Al³⁺ readily); Tl(III) is unstable/oxidising (Tl(I) preferred).

Trends to remember

• Ga radius anomaly: Ga < Al (poor 3d shielding → higher Z_eff).

• IE pattern: doesn’t decrease smoothly (Ga, Tl show irregularities via d/f contraction).

• Ga physical anomaly: mp 303 K (liquid near room temp), bp 2676 K → huge liquid range.

• Densities increase down the group.

D. Occurrence & extraction (high-yield)

Occurrence

• Boron: mainly as H₃BO₃, borax Na₂B₄O₇·10H₂O, kernite Na₂B₄O₇·4H₂O, colemanite Ca₂B₆O₁₁·5H₂O; crustal abundance < 0.0001% (¹⁰B 19%, ¹¹B 81%).

• Aluminium: most abundant metal (~8.3% by mass), ores: bauxite Al₂O₃·xH₂O, cryolite Na₃AlF₆, corundum Al₂O₃.

• Ga, In: trace in Al/Zn ores; Tl in sulphide ores (Zn, Pb, Cu, Fe).

Extraction snapshots

• Boron: reduce B₂O₃ with Mg/Na; pure crystalline B by reducing BBr₃ with H₂ on hot Ta (1275–1475 K).

• Aluminium (Hall–Héroult):

  1. Purify Al₂O₃ from bauxite (Bayer):

     • Al₂O₃ + 2NaOH + 3H₂O → 2 Na[Al(OH)₄]

     • SiO₂ + 2NaOH → Na₂SiO₃ + H₂O

     • 2NaAl(OH)₄ + CO₂ → Al(OH)₃↓ + Na₂CO₃ + H₂O

     • Al(OH)₃ → Al₂O₃ + 3H₂O (Δ)

  2. Dissolve Al₂O₃ in molten cryolite (with CaF₂) at ~1100–1300 K; C-lined cell cathode, C anodes.

     • Cathode: Al³⁺ + 3e⁻ → Al(l)

     • Anode: O²⁻ + C → CO/CO₂ (anode consumption)

E. Chemical properties & reactions

1) Oxidation states & Lewis acidity

• +3: common; +1 stabilises down group (inert-pair). Tl(I) dominant, Tl(III) oxidising.

• Electron-deficient trihalides (e.g., BF₃, BCl₃, AlCl₃) are Lewis acids; BCl₃·NH₃ forms readily.

• In water, AlCl₃ etc. hydrolyse to [M(OH)₄]⁻/[Al(H₂O)₆]³⁺, showing sp³/sp³d² (via acceptor levels; boron lacks d—max covalence 4).

2) Action of air/oxygen & nitrogen

• Surface oxide: Al forms protective Al₂O₃.

• On heating in air: B₂O₃, Al₂O₃ form.

• With N₂ at high T → nitrides:

  • 2E + 3/2 O₂ → E₂O₃ (E = B, Al, Ga, In)

  • 2E + N₂ → EN (B, Al)

  • BN also via: 2B + 2NH₃ → 2BN + 3H₂; or B₂O₃ + 3C + N₂ → 2BN + 3CO

• Acid–base of oxides: B₂O₃ acidic; Al₂O₃, Ga₂O₃ amphoteric; In₂O₃, Tl₂O₃ basic. Tl₂O dissolves → TlOH (strong base).

3) Reactivity with water, acids, bases

• B inert to acids/alkali at moderate T; hot oxidising acids slowly → H₃BO₃.

  • 2B + 3H₂SO₄ → 2H₃BO₃ + 3SO₂

  • B + 3HNO₃ → H₃BO₃ + 3NO₂

• Al/Ga/In/Tl + warm dilute HCl/H₂SO₄ → M³⁺ (Tl → Tl⁺) + H₂.

  • 2M + 6HCl → 2M³⁺ + 6Cl⁻ + 3H₂

  • 2Tl + 2HCl → 2Tl⁺ + 2Cl⁻ + H₂

• Passivation: conc. HNO₃ renders Al, Ga passive (oxide film).

• Al/Ga + aqueous alkali → aluminate/gallate + H₂:

  • Al + 2NaOH + 2H₂O → 2NaAlO₂ + 3H₂

  • More fully: 2Al + 2NaOH + 6H₂O → 2Na⁺[Al(OH)₄]⁻ + 3H₂

4) Carbon, borides, thermite

• Carbides:

  • 4B + C → B₄C (very hard, covalent, abrasive)

  • 4Al + 3C → Al₄C₃ (ionic; Al₄C₃ + 12H₂O → 4Al(OH)₃ + 3CH₄)

• Borides: hard, refractory (e.g., TiB₂, ZrB₂).

• Thermite (aluminothermy):

  • Fe₂O₃ + 2Al → Al₂O₃ + 2Fe(l) (≈3300 K)

  • Also used for MnO₂, Cr₂O₃, etc.

F. Halides of Group 13

1) Boron trihalides: BX₃ (X = F, Cl, Br, I)

• Preparation (selected):

  • B₂O₃ + 3CaF₂ + 3H₂SO₄ → 2BF₃ + 3CaSO₄ + 3H₂O

  • Na₂B₄O₇ + 6CaF₂ + 8H₂SO₄ → 4BF₃ + 6CaSO₄ + 6NaHSO₄ + 7H₂O

  • B₂O₃ + 3C + 3Cl₂ (775 K) → 2BCl₃ + 6CO

  • BI₃: LiBH₄/NaBH₄ + I₂ (400–443 K): BI₃ + LiI/NaI + HI

• Properties: volatile; hydrolyse:

  • BF₃ + 3H₂O → H₃BO₃ + 3HBF₄

  • BCl₃ + 3H₂O → H₃BO₃ + 3HCl

• Structure: trigonal planar (sp²), Lewis acidic.

• Back-bonding & acidity: π back-donation BF₃ > BCl₃ > BBr₃ > BI₃ (strongest at F), so Lewis acidity increases: BF₃ < BCl₃ < BBr₃ < BI₃.

• Lower dihalides: B₂F₄ (gas), B₂Cl₄ (liquid), less stable than BX₃.

2) Aluminium trihalides: AlX₃

• AlF₃: ionic, non-volatile (from Al₂O₃ + HF at 1000 K).

• AlCl₃/AlBr₃/AlI₃: covalent dimers Al₂X₆ (halogen bridges); in donor solvents give adducts (AlCl₃·OEt₂, AlCl₃·NMe₃).

• Hydrolysis / aquation: [Al(H₂O)₆]³⁺ (acidic), stepwise deprotonation to [Al(H₂O)₂(OH)₄]⁻ in base; gelatinous Al(OH)₃ precipitates near neutral pH (amphoteric).

• Friedel–Crafts: strong Lewis acid (AlCl₃).

G. Oxides & hydroxides

• B₂O₃ (boron trioxide): white, hygroscopic, acidic, polymeric (-B–O–B–).

  • Dissolves: B₂O₃ + 3H₂O → 2 B(OH)₃

  • Fused with metal oxides → metaborates (borax bead colours: Co, Cu, Cr, etc.).

• Al₂O₃ (alumina): γ-Al₂O₃ (from ≤750 K dehydration) → reactive, dissolves in acid/base; at ~1500 K converts to α-Al₂O₃ (corundum): very hard, mp ≈ 2328 K, inert.

• Trend: acidity → basicity increases down the group:
  B₂O₃ acidic → Al₂O₃/Ga₂O₃ amphoteric → In₂O₃/Tl₂O₃ basic; Tl₂O strongly basic (TlOH ≈ strong base).

H. Oxoacids of Boron & Borates

1) Orthoboric acid: H₃BO₃ (= B(OH)₃)

• Prep: Na₂B₄O₇ + 2HCl + 5H₂O → 2NaCl + 4B(OH)₃
  (also by hydrolysis of many boron compounds).

• Structure: layered sheets of trigonal BO₃ units (H-bonded).

• Acidity: Lewis acid (accepts OH⁻), not Brønsted:

  • B(OH)₃ + 2H₂O ⇌ [B(OH)₄]⁻ + H₃O⁺ (very weak, pK_a ~9.25).

  • Strength increases with polyols (mannitol, glycerol) → analytical titration.

• Heating: H₃BO₃ (Δ ≥ 370 K) → HBO₂ (metaboric); further Δ > 500 K → B₂O₃.

2) Borax: Na₂B₄O₇·10H₂O (actually Na₂[B₄O₅(OH)₄]·8H₂O)

• Hydrolysis: Na₂B₄O₇ + 7H₂O → 2NaOH + 4H₃BO₃ (alkaline solution).

• Heating:

  • Na₂B₄O₇·10H₂O → Na₂B₄O₇ (Δ)

  • Na₂B₄O₇ → 2NaBO₂ + B₂O₃ (Δ)

• Borax bead: metaborates of transition metal ions give diagnostic colours.

• From colemanite: Ca₂B₆O₁₁ + 2Na₂CO₃ → Na₂B₄O₇ + 2NaBO₂ + 2CaCO₃; then 4NaBO₂ + CO₂ → Na₂B₄O₇ + Na₂CO₃.

I. Hydrides & boranes; key reducing agents

1) Diborane (B₂H₆)

• Preparations

  • 4BF₃ + 3LiAlH₄ → 2B₂H₆ + 3LiF + 3AlF₃ (Et₂O)

  • 2NaBH₄ + H₂SO₄ → B₂H₆ + 2H₂ + Na₂SO₄ (also with H₃PO₄)

  • 2NaBH₄ + I₂ → B₂H₆ + 2NaI + H₂

  • Industry: 2BF₃ + 6NaH (450 K) → B₂H₆ + 6NaF

• Properties: colourless, toxic gas (bp 180 K), spontaneously flammable; hydrolyses completely:

  • B₂H₆ + 6H₂O → 2B(OH)₃ + 6H₂

  • Combustion: B₂H₆ + 3O₂ → B₂O₃ + 3H₂O (ΔH_c ≈ −1976 kJ·mol⁻¹)

• Structure (must-know!):

  • 4 terminal B–H 2c–2e bonds (in the B–B plane).

  • 2 bridging B–H–B 3c–2e “banana” bonds (above/below plane).

• Adduct formation (cleavage by bases):

  • B₂H₆ + 2NMe₃ → 2BH₃·NMe₃

  • B₂H₆ + 2CO → 2BH₃·CO

• With NH₃ (on heating): B₃N₃H₆ (borazine) + H₂ (via [BH₂(NH₃) ₂]⁺[BH₄] ⁻).

2) Borohydrides / hydridoborates

• [BH₄] ⁻ tetrahedral.

• 2MH + B₂H₆ → 2M⁺[BH₄] ⁻ (M = Li, Na).

• NaBH₄ (selective; aldehydes/ketones → alcohols) vs LiAlH₄ (powerful, less selective; reacts with H₂O):

  • LiAlH₄ + 4H₂O → LiOH + Al(OH)₃ + 4H₂

  • 4BCl₃ + 3LiAlH₄ → 2B₂H₆ + 3LiCl + 3AlCl₃

3) Aluminium hydride (AlH₃)

• Polymeric (AlH₃)ₙ; from LiAlH₄ + H₂SO₄ / AlCl₃; reacts with protic reagents to give H₂.

J. Borazine (B₃N₃H₆) — “inorganic benzene”

• Prepared via BCl₃ + NH₄Cl → B-trichloroborazine, then reduction with NaBH₄; or from B₂H₆ + NH₃ (see above).

• Planar hexagon, alternating B–N; partial π-delocalisation (more reactive than benzene).

• Polarisation: B δ⁺ / N δ⁻ overall (σ-bond polarity dominates) → nucleophilic attack at B; adds H₂O/MeOH/HX to give 1:3 adducts (lose H₂ on heating).

K. Complex formation & alums

• Boron: max CN = 4 (no d); e.g., [BF₄]⁻, [BH₄]⁻, BH₃·L adducts, chelates [B(O-C₆H₄O)₂]⁻.

• Al/Ga/In/Tl: CN up to 6 (acceptor orbitals): adducts AlCl₃·L (CN 4/5), [AlF₆]³⁻ (CN 6).

• Alums: MAl(SO₄)₂·12H₂O (M = K, Rb, Cs, NH₄…) are double salts (not complexes), isomorphous; e.g., potash alum KAl(SO₄)₂·12H₂O, ammonium alum NH₄Al(SO₄)₂·12H₂O; analogues with Cr³⁺, Fe³⁺ (chrome/ferric alum).

L. Applications

• Boron: very hard, low density, low conductivity → fibres (composites, armor), neutron absorbers (¹⁰B; borides as control/shield).

• Boron compounds: borax/boric acid in Pyrex/borosilicate glass, glazes, bead test; peroxyborates (cleaning).

• Aluminium: light, strong, conductive → structural alloys (duralumin, magnalium), transport, packaging/foil, cookware (use declining due to toxicity concerns), anodising (hard Al₂O₃ film).

• Ga: electronics (GaAs, GaN); long liquid range (thermometry).

• In: bearings, low-m.p. alloys, electronics.

• Tl: fungicides (toxic), high-index optical glass.

M. Why Ga is “smaller than Al” (and other anomalies)

• Inserting a filled 3d¹⁰ subshell between Al and Ga yields poor shielding, so Ga’s outer electrons feel larger Zeff → smaller radius, higher IE/χ than a naive trend would suggest.

• Similar reasoning with 4f affects Tl.

N. Important points

1. Passivation of Al by conc. HNO₃: a thin, coherent Al₂O₃ film forms, stopping further attack.

2. Why BCl₃ is gaseous but AlCl₃ dimerises (Al₂Cl₆): B (no d) uses pπ–pπ back-bonding to “complete” octet; Al lacks effective pπ overlap, so dimerises via Cl bridges.

3. Why BF₆³⁻ doesn’t exist: B (n=2) cannot expand octet; max covalence 4.

4. Standard potentials: Al³⁺/Al (−1.66 V) → Al³⁺ favoured (electropositive metal); Tl³⁺/Tl (+1.26 V) → Tl(III) unstable/oxidising; Tl(I) more stable in solution.

O. All reactions (collected for revision)

Oxides/nitrides (air/N₂):

• 2E + 3/2 O₂ → E₂O₃ (E = B, Al, Ga, In)

• 2E + N₂ → EN (B, Al); BN also from B₂O₃ + 3C + N₂ → 2BN + 3CO; or 2B + 2NH₃ → 2BN + 3H₂

Acids/alkali:

• 2B + 3H₂SO₄ → 2H₃BO₃ + 3SO₂

• B + 3HNO₃ → H₃BO₃ + 3NO₂

• 2M + 6HCl → 2M³⁺ + 6Cl⁻ + 3H₂ (Tl → Tl⁺ + H₂)

• 2Al + 2NaOH + 6H₂O → 2Na⁺[Al(OH)₄]⁻ + 3H₂

Carbides/borides:

• 4Al + 3C → Al₄C₃; Al₄C₃ + 12H₂O → 4Al(OH)₃ + 3CH₄

• 4B + C → B₄C

• Borides: MgB₂, TiB₂, ZrB₂, etc.

Thermite (examples):

• Fe₂O₃ + 2Al → Al₂O₃ + 2Fe(l)

• 3MnO₂ + 4Al → 2Al₂O₃ + 3Mn

Trihalides (prep):

• B₂O₃ + 3CaF₂ + 3H₂SO₄ → 2BF₃ + 3CaSO₄ + 3H₂O

• Na₂B₄O₇ + 6CaF₂ + 8H₂SO₄ → 4BF₃ + 6CaSO₄ + 6NaHSO₄ + 7H₂O

• B₂O₃ + 3C + 3Cl₂ (775 K) → 2BCl₃ + 6CO

• Al₂O₃ + 6HF (1000 K) → 2AlF₃ + 3H₂O

• 2Al + 3X₂ → 2AlX₃; Al₂O₃ + 3C + 3Cl₂ → 2AlCl₃ + 3CO

Hydrolysis/adducts:

• BCl₃ + 3H₂O → H₃BO₃ + 3HCl

• BF₃ + 3H₂O → H₃BO₃ + 3HBF₄

• BCl₃ + :NH₃ → BCl₃·NH₃

• Al₂Cl₆ + (Et₂O) → 2(Et₂O·AlCl₃)

Boric acid & borax:

• Na₂B₄O₇ + 2HCl + 5H₂O → 2NaCl + 4B(OH)₃

• B(OH)₃ + 2H₂O ⇌ [B(OH)₄]⁻ + H₃O⁺

• H₃BO₃ (Δ) → HBO₂ (Δ) → B₂O₃

• Na₂B₄O₇·10H₂O →(Δ) Na₂B₄O₇ →(Δ) 2NaBO₂ + B₂O₃

• Ca₂B₆O₁₁ + 2Na₂CO₃ → Na₂B₄O₇ + 2NaBO₂ + 2CaCO₃

• 4NaBO₂ + CO₂ → Na₂B₄O₇ + Na₂CO₃

Diborane & hydrides:

• 4BF₃ + 3LiAlH₄ → 2B₂H₆ + 3LiF + 3AlF₃

• 2NaBH₄ + I₂ → B₂H₆ + 2NaI + H₂

• 2NaBH₄ + H₂SO₄ → B₂H₆ + 2H₂ + Na₂SO₄

• B₂H₆ + 6H₂O → 2B(OH)₃ + 6H₂

• 2MH + B₂H₆ → 2M⁺[BH₄]⁻ (M = Li, Na)

• LiAlH₄ + 4H₂O → LiOH + Al(OH)₃ + 4H₂

• 4BCl₃ + 3LiAlH₄ → 2B₂H₆ + 3LiCl + 3AlCl₃

• B₂H₆ + 2NMe₃ → 2BH₃·NMe₃

• B₂H₆ + 2CO → 2BH₃·CO