Comprehensive Study Notes – p-Block Elements (Groups 13 & 14)

Position of the p-Block and General Characteristics

• In p-block elements the last electron enters an outermost $p$ orbital. With three $p$ orbitals, a maximum of six electrons can be accommodated, hence six groups (13 → 18).
• Valence-shell configuration (except He): $ns^2np^{1-6}$.
• Heading elements: B, C, N, O, F, He.
• Inner-core differences (presence/absence of filled $d$ and $f$ shells) strongly influence:
  • Atomic/ionic radii
  • Ionisation enthalpies $\Delta_iH$ and electronegativities
  • Oxidation-state patterns and the ability to involve $d$ orbitals (possible only from the 3rd period onwards).

Oxidation-State Trends

• Maximum (group) oxidation state = total number of valence electrons $s + p$.
• Elements often also show the state two units lower ((\text{group} - 2)), increasingly stable down a group – the inert-pair effect (poor screening by interposed $d$/$f$ electrons keeps the $ns^2$ pair non-bonding).
• General pattern (Table 11.1):
  • Group 13: $+3$ (stable for B, Al) and $+1$ (stable for Ga, In, TI).
  • Group 14: $+4$ (C, Si) and $+2$ (stable for Sn, predominant for Pb).
• Light p-block elements excel at $p\pi{-}p\pi$ multiple bonding (C=C, C≡C, N≡N, etc.). Heavier congeners form weaker $d\pi{-}p\pi$/$d\pi{-}d\pi$ bonds instead.

Group 13 – The Boron Family

1. Occurrence and Isotopes

• Boron is rare (≤0.0001 % of crust). Minerals: orthoboric acid $(\mathrm{H<em>3BO</em>3})$, borax $\mathrm{Na<em>2B</em>4O<em>7\cdot10H</em>2O}$, kernite $\mathrm{Na<em>2B</em>4O<em>7\cdot4H</em>2O}$.
• Isotopes: $^{10}\mathrm{B}$ (19 %), $^{11}\mathrm{B}$ (81 %).
• Aluminium: third most abundant element (8.3 %), minerals bauxite $\mathrm{Al<em>2O</em>3\,2H<em>2O}$, cryolite $\mathrm{Na</em>3AlF_6}$.
• Ga, In, Tl scarce. Nh (Z = 113) synthetically made; chemistry not established (t½ ≈ 20 s).

2. Electronic Configuration

$\text{Outer: } ns^2np^1$

• B, Al: noble-gas cores.
• Ga, In: +10$d$ electrons; Tl: +14$f$ +10$d$ → complex shielding trends.

3. Periodic Trends (Table 11.2 excerpts)

Key anomalies and causes:

• Ga radius < Al owing to poor screening by 3$d^{10}$.
• $\Delta_iH$ discontinuities Al→Ga and In→Tl arise from weak screening by $d$/$f$.
• Ga melts near room temperature ($303\,\text{K}$), yet boils at $2676\,\text{K}$ → useful high-T thermometer liquid.

4. Oxidation-State Chemistry

• B: essentially only covalent; cannot form $\mathrm{B^{3+}}$ (huge $\Sigma\Delta<em>iH</em>{1-3}$).
• Al: readily forms $\mathrm{Al^{3+}}$ (highly electropositive, E^\ominus_{\text{Al^{3+}/Al}} = -1.66\,\text{V}).
• Ga, In, Tl: +1 and +3; stability order of +1 increases Al < Ga < In < Tl (Tl(I) predominant, Tl(III) strong oxidant, E^\ominus_{\text{Tl^{3+}/Tl}} = +1.26\,\text{V}).

5. Reactivity Patterns

(i) Air/Oxygen

$\mathrm{2E + 3O<em>2 \;\xrightarrow{\Delta}\; 2E</em>2O_3}$ (E = B, Al, …)

• Crystalline B inert; amorphous B and Al form protective $\mathrm{B<em>2O</em>3}$ / $\mathrm{Al<em>2O</em>3}$ layers.
• Oxide nature trend: $\mathrm{B<em>2O</em>3}$ acidic > $\mathrm{Al<em>2O</em>3}$ amphoteric > $\mathrm{Ga<em>2O</em>3}$ amphoteric > $\mathrm{In<em>2O</em>3, Tl<em>2O</em>3}$ basic.

(ii) Acids/Alkali

• B unaffected; Al amphoteric:
  $\mathrm{2Al + 6HCl \rightarrow 2Al^{3+} + 6Cl^- + 3H<em>2}$$\mathrm{2Al + 2NaOH + 6H</em>2O \rightarrow 2Na^+[Al(OH)<em>4]^- + 3H</em>2}$

(iii) Halogens

$\mathrm{2E + 3X<em>2 \rightarrow 2EX</em>3}$ (except $\mathrm{TlI_3}$ unstable).

• $\mathrm{AlCl<em>3}$ dimerises to $\mathrm{Al</em>2Cl_6}$; fumes in moist air via hydrolysis $\rightarrow$ $\mathrm{HCl}$.

6. Lewis-Acid Behaviour

• Trivalent halides $\mathrm{BX<em>3, AlX</em>3, …}$ are electron-deficient (6 e⁻ around M) → strong Lewis acids.
  $\mathrm{BCl<em>3 + :NH</em>3 \rightleftharpoons BCl<em>3\cdot NH</em>3}$
• Acid strength decreases down the group (size ↑).
• $\mathrm{BF<em>4^-}$ forms easily; $\mathrm{BF</em>6^{3-}}$ impossible (B lacks $d$ orbitals for octet expansion).

7. Important Boron Compounds

(a) Borax

$\mathrm{Na<em>2[B</em>4O<em>5(OH)</em>4]\cdot8H<em>2O}$ (trad. $\mathrm{Na</em>2B<em>4O</em>7\cdot10H_2O}$).

• Alkaline in water: $\mathrm{Na<em>2B</em>4O<em>7 + 7H</em>2O \rightarrow 2NaOH + 4H<em>3BO</em>3}$.
• Heating → puffy ➔ glassy borax bead $\rightarrow$ $\mathrm{NaBO<em>2 + B</em>2O_3}$ (bead test for transition metals).

(b) Orthoboric Acid $\mathrm{H<em>3BO</em>3}$

• Layer lattice joined by H-bonds; weak monobasic Lewis acid: $\mathrm{B(OH)<em>3 + 2H</em>2O \rightleftharpoons [B(OH)<em>4]^- + H</em>3O^+}$.
• $\mathrm{\Delta >370\,K}$: $\mathrm{H<em>3BO</em>3 \rightarrow HBO<em>2 \rightarrow B</em>2O_3}$.

(c) Diborane $\mathrm{B<em>2H</em>6}$

• Prep: $\mathrm{4BF<em>3 + 3LiAlH</em>4 \rightarrow 2B<em>2H</em>6 + 3LiF + 3AlF<em>3}$ or laboratory $\mathrm{2NaBH</em>4 + I<em>2 \rightarrow B</em>2H<em>6 + 2NaI + H</em>2}$.
• Structure: four terminal $\mathrm{B{-}H}$ (2c–2e) + two bridging $\mathrm{B{-}H{-}B}$ (3c–2e “banana”) bonds; B uses $sp^3$ hybrids.
• Highly pyrophoric; hydrolysis: $\mathrm{B<em>2H</em>6 + 6H<em>2O \rightarrow 2B(OH)</em>3 + 6H_2}$.
• Reacts with bases to give borane adducts $\mathrm{BH<em>3\cdot L}$; with $\mathrm{NH</em>3}$ ⇒ initially $[BH<em>2(NH</em>3)<em>2]^+[BH</em>4]^-$, on heating ➔ borazine $\mathrm{B<em>3N</em>3H_6}$ (“inorganic benzene”).

8. Uses of Boron & Aluminium

• Boron fibres: bullet-proof vests, aerospace composites.
• $^{10}\mathrm{B}$ neutron absorber → control rods.
• Borax/boric acid: Pyrex®, glass-wool, soldering flux, glazed enamels, mild antiseptic.
• Al: lightweight, high conductivity, machines, aircraft, foil; forms alloys (Cu, Mg, Mn, Si, Zn). Toxicological concerns limit culinary use.

Group 14 – The Carbon Family

1. Natural Abundance & Isotopes

• Carbon 17th most abundant; forms coal, diamond, graphite; atmospheric $\mathrm{CO_2}$ ≈ 0.03 %.
• Isotopes: $^{12}\mathrm{C}$ (standard), $^{13}\mathrm{C}$ (1.1 %), $^{14}\mathrm{C}$ (t½ = 5770 y; radiodating).
• Silicon: 2nd most abundant element (27.7 %) as silica/silicates; vital for glass, ceramics, cement.
• Ge trace, Sn as cassiterite $\mathrm{SnO_2}$, Pb as galena $\mathrm{PbS}$.

2. Electronic Configuration and Radii

• Valence: $ns^2np^2$.
• Covalent radius: large jump C (77 pm) → Si (118 pm), only slight increase thereafter due to poor $d$/$f$ shielding (cf. Ga anomaly).
• $\Delta<em>iH</em>1$ declines C→Si→Ge→Sn, slight rise Sn→Pb (inert-pair effect).

3. Physical Gradation

• All solids; C, Si non-metal, Ge metalloid, Sn/Pb soft metals.
• Melting/boiling points much higher than group 13; metallic character intensifies downward.

4. Oxidation-State Chemistry

• Common: $+4$ and $+2$; C additionally negative states (e.g., $\mathrm{CH_4}$ –4).
• Stability pattern: $+4$ (C, Si), $+4 > +2$ (Ge), $+4\approx +2$ (Sn), $+2$ dominant (Pb; $+4$ a strong oxidant).
• Heavier elements capable of expanding octet (utilise $d$ orbitals) → complex ions $\mathrm{[SiF<em>6]^{2-}}, [Sn(OH)</em>6]^{2-}$ etc.

Important Reactivities

(i) With O₂ → monoxides/dioxides $\mathrm{MO, MO_2}$.

• Acid–base nature: $\mathrm{CO<em>2, SiO</em>2, GeO<em>2}$ acidic > $\mathrm{SnO</em>2, PbO<em>2}$ amphoteric; among monoxides $\mathrm{CO}$ neutral, $\mathrm{GeO}$ acidic, $\mathrm{SnO,PbO}$ amphoteric. (ii) With Water: C, Si, Ge inert; Sn + steam $\rightarrow$ $\mathrm{SnO</em>2 + H<em>2}$; Pb inert (protective film). (iii) Halides: $\mathrm{MX</em>4}$ (tetrahedral, covalent except $\mathrm{SnF<em>4, PbF</em>4}$ ionic). $\mathrm{MX<em>2}$ stability ↑ downwards; $\mathrm{PbI</em>4}$ non-existent (insufficient energy to free 6s² pair).
• Hydrolysis: $\mathrm{SiCl<em>4 + 4H</em>2O \rightarrow Si(OH)<em>4 + 4HCl}$; $\mathrm{CCl</em>4}$ resists (no $d$ orbitals).

5. Special Behaviour of Carbon

• Small size, high $\chi$, high $\Delta_iH$, lack of $d$ orbitals.
• Catenation (self-linking) order: $\mathrm{C \gg Si > Ge \approx Sn > Pb}$; governed by bond enthalpy
  $\mathrm{C{-}C}$ 348 kJ·mol⁻¹ > $\mathrm{Si{-}Si}$ 297 > $\mathrm{Ge{-}Ge}$ 260 > $\mathrm{Sn{-}Sn}$ 240.
• Extensive $p\pi{-}p\pi$ bonding possible; unavailable to heavier congeners.

6. Allotropes of Carbon

(a) Diamond

• 3-D network of $sp^3$ tetrahedral carbons; $d<em>{CC}=154\,\text{pm}$; hardest material; electrical insulator; $\Delta</em>fH^\ominus=+1.90\,\text{kJ mol}^{-1}$.

(b) Graphite

• Layers of planar $sp^2$ hexagonal sheets (C–C 141.5 pm) separated by 340 pm; delocalised $\pi$-cloud → good conductor; layers slip → lubricant.

(c) Fullerenes $\mathrm{C<em>{60}, C</em>{70}, …}$

• Made by laser/electric-arc vaporisation of graphite in He/Ar.
• $\mathrm{C<em>{60}}$ Buckminsterfullerene: soccer-ball cage (20 hexagons + 12 pentagons); all $sp^2$ carbons; mixed single (143.5 pm) / double (138.3 pm) bonds; $\Delta</em>fH^\ominus=38.1\,\text{kJ mol}^{-1}$.

Other forms: carbon black, coke, charcoal (amorphous graphite variants).

Uses

• Graphite electrodes, high-T crucibles, fibre composites, dry lubricants.
• Diamond abrasives, drill bits, jewellery (measured in carats, $1\,\text{carat}=200\,\text{mg}$).
• Activated charcoal: adsorption of poisons/odours.

Key Compounds of C & Si

1. Carbon Monoxide $\mathrm{CO}$

• Formed by limited-oxygen combustion: $\mathrm{2C + O_2 \rightarrow 2CO}$.
• Lab prep: dehydration of formic acid with conc. $\mathrm{H<em>2SO</em>4}$ at $373\,\text{K}$.
• Industrial: steam on red-hot coke $\rightarrow$ water gas (CO + H₂); air on coke $\rightarrow$ producer gas (CO + N₂).
• Linear molecule: one $\sigma$ + two $\pi$ bonds, lone pair on C → ligand; forms metal carbonyls.
• Highly toxic: binds haemoglobin ~300× stronger than O₂.
• Strong reductant: $\mathrm{Fe<em>2O</em>3 + 3CO \rightarrow 2Fe + 3CO_2}$.

2. Carbon Dioxide $\mathrm{CO_2}$

• Complete combustion of C/fuels; lab: $\mathrm{CaCO_3 + 2HCl}$.
• Linear $O=C=O$ (equal C–O 115 pm), $sp$ hybridised; resonance $O=C=O \leftrightarrow ^-O{-}C≡O^+$.
• Weak dibasic acid $\mathrm{H<em>2CO</em>3}$ (buffers blood 7.26–7.42).
• Photosynthesis: $6CO<em>2 + 12H</em>2O \xrightarrow{h\nu,\,\text{chlorophyll}} C<em>6H</em>{12}O<em>6 + 6O</em>2 + 6H_2O$.
• Applications: dry-ice refrigerant, fire extinguishers, carbonation of beverages.
• Excess atmospheric CO₂ → greenhouse effect ⟹ global warming.

3. Silicon Dioxide $\mathrm{SiO_2}$ (silica)

• Giant $\mathrm{SiO_4}$ tetrahedral network; polymorphs quartz, tridymite, cristobalite.
• Chemically inert (high $E<em>{Si{-}O}$) but reacts with$\;\mathrm{HF}$: $\mathrm{SiO</em>2 + 4HF \rightarrow SiF<em>4 + 2H</em>2O}$; with alkali: $\mathrm{SiO<em>2 + 2NaOH \rightarrow Na</em>2SiO<em>3 + H</em>2O}$.
• Uses: piezoelectric oscillators, chromatographic supports, silica gel desiccant, kieselghur filtration.

4. Silicones (Organopolysiloxanes)

• Repeating unit $(R<em>2SiO)</em>n$; made from hydrolysis/condensation of alkyl- chlorosilanes. Example: \mathrm{(CH3)2SiCl2 \xrightarrow{H2O} (CH3)2Si(OH)2 \xrightarrow[-H2O]{\text{polymerise}} [-(CH3)2SiO-]_n.
• End-capping with $\mathrm{(CH<em>3)</em>3SiCl}$ controls chain length.
• Hydrophobic, thermally stable, high dielectric strength; used as sealants, greases, waterproofing agents, biocompatible implants.

5. Silicates & Zeolites

• Fundamental unit $\mathrm{SiO_4^{4-}}$ tetrahedron. Corner-sharing generates chains, rings, sheets (mica, asbestos) or 3-D frameworks (feldspar).
• Partial $\mathrm{Si^{4+}\rightarrow Al^{3+}}$ substitution yields negatively charged aluminosilicates; cations $\mathrm{Na^+, K^+, Ca^{2+}}$ balance.
• Zeolites (e.g., ZSM-5): microporous aluminosilicates – shape-selective catalysts (cracking, isomerisation) and ion-exchangers for water softening.

Conceptual & Practical Connections

• Comparison with s-block: first-row anomalies (B vs heavier; Li, Be analogues) arise from small size + absence of $d$ orbitals.
• Inert-pair effect parallels trend in post-transition metals (Tl, Pb, Bi etc.).
• $d$-Orbital availability from 3rd period affects maximum valence (contrast B limited to 4 vs Al 6-coordination).
• Acid–base character of oxides illustrates metallicity gradient within a group.
• Environmental relevance: CO toxicity, CO₂ greenhouse, boron neutron capture therapy.

Numerical / Statistical Highlights

• $r<em>{\text{Ga}} = 135\,\text{pm} < r</em>{\text{Al}} = 143\,\text{pm}$ due to poor 3$d$ screening.
• $\Delta<em>iH</em>1$ (B) $= 801\,\text{kJ mol}^{-1}$; (Al) $= 577$; (Tl) $= 589$.
• $E^{\ominus}<em>{\text{Al}^{3+}/\text{Al}} = -1.66\,\text{V};\; E^{\ominus}</em>{\text{Tl}^{3+}/\text{Tl}} = +1.26\,\text{V}$.
• $\Delta<em>fH^{\ominus}(\text{diamond}) = +1.90\,\text{kJ mol}^{-1};\; \Delta</em>fH^{\ominus}(\text{C}_{60}) = +38.1$.
• $d<em>{\text{C–C}}$: diamond 154 pm; graphite 141.5 pm (in-layer); $d</em>{\text{C–O}}$ in CO₂ 115 pm.
• $\text{Ga m.p.} = 303\,\text{K},\; b.p. = 2676\,\text{K}$.

Ethical, Practical & Industrial Implications

• Toxic effects of aluminium cookware, CO inhalation, and greenhouse CO₂ levels.
• Boron neutron capture for cancer therapy; borax in detergents raises environmental phosphate issues.
• Use of zeolites as eco-friendly catalysts vs acid-catalysed cracking.

Sample Reactions (LaTeX Notation)

$\mathrm{2Al(s) + 2NaOH(aq) + 6H<em>2O(l) \;\rightarrow\; 2Na^+[Al(OH)</em>4]^- + 3H<em>2(g)}$$\mathrm{SiCl</em>4 + 4H<em>2O \;\rightarrow\; Si(OH)</em>4 + 4HCl}$
$\mathrm{B<em>2H</em>6 + 6H<em>2O \;\rightarrow\; 2B(OH)</em>3 + 6H<em>2}$$\mathrm{CO + Fe</em>2O<em>3 \;\rightarrow\; 2Fe + 3CO</em>2}$
$\mathrm{SiO<em>2 + 4HF \;\rightarrow\; SiF</em>4 + 2H_2O}$

Quick Revision Points

• $\text{Group 13}:$ Covalent B; metallicity ↑ down group; +1 state stabilises (inert pair).
• $\text{Group 14}:$ Catenation strongest for C; +2 state dominates Pb; allotropes of carbon with distinct properties.
• Electron-deficient $\mathrm{BX<em>3}$ Lewis acids; $\mathrm{AlCl</em>3}$ dimerises $\rightarrow$ $\mathrm{Al<em>2Cl</em>6}$.
• Diborane contains 3c–2e bonds; boric acid a weak Lewis acid, not protonic.
• Silicones: $(R<em>2SiO)</em>n$ hydrophobic polymers; zeolites: 3-D aluminosilicate catalysts.