Boron Family (Group 13): Trends, Boron Compounds, and Aluminium Industry Meet Group 13 (Boron Family): Why this matters Group 13 has the valence configuration ns 2 ,np 1 and includes boron (B), aluminium (Al), gallium (Ga), indium (In), and thallium (Tl). From B to Tl, atomic radius generally increases and metallic character increases (B is a non-metal; others are metals). Ionization enthalpy generally decreases down the group, with an important anomaly: Ga has slightly higher first ionization enthalpy and a smaller atomic radius than Al due to d-block contraction. Oxidation state +3 dominates near the top; lower states (+1) become more stable down the group due to the inert pair effect (Tl is mostly +1). Big-picture trends for Group 13 (and 14 for comparison): metallic character, electronegativity, density, and melting point directions. Use this as a visual compass while studying anomalies. Smallest in group Highest in group Relatively higher in group Non-metal High +3 (dominant), covalent chemistry Al Larger than B Drops from B Lower than B Metal Moderate ( ~933 K) +3 (dominant) Ga Slightly smaller than Al (d-block contraction) Slightly higher than Al (anomalous) Similar to Al Metal Very low (melts near room temp) +3, +1 appears In Increases from Ga Decreases from Ga Lower Metal Moderate +3, +1 more common Tl Largest Not strictly lowest; +1 stabilized (inert pair effect) Lower Metal Moderate +1 (most stable), +3 less stable Element Member Atomic radius IE (1st) Electronegativity Metallic character Melting point (relative) Common oxidation states Qualitative trend map across Group 13 (NEET view) Ionization enthalpy and atomic radii across Groups 13 and 14 with anomalies highlighted (Ga vs Al; Tl vs In). Anomalies are NEET-tricky — memorize the reasons. Group 13 shows anomalies: Ga has smaller radius and slightly higher first ionization enthalpy than Al due to d-block contraction; Tl trends are influenced by f-block contraction and inert pair effect. Periodic trends down a group change smoothly for every property. Inert pair effect (preserved) Boron vs Aluminium: diagonal similarity but not identical Boron is small, has no low-lying vacant d-orbitals, and forms electron-deficient covalent compounds; its maximum covalency is 4. Aluminium is metallic, often forms partially ionic or coordinate compounds, and (unlike B) can reach coordination number 6 in many complexes because it can expand its coordination sphere via available orbitals in bonding schemes taught at this level. Both show some diagonal relationship features with elements in Group 14/2 (e.g., B with Si; Be with Al) — similar charge/radius patterns can mimic chemistry. But do not assume B and Al are identical in behavior. Typical bonding Predominantly covalent; many electron-deficient species (e.g., B 2 H 6 ) More ionic/covalent borderline; strong Lewis acid centers (e.g., AlCl 3 ) Max covalency (common) 4 (no d-orbitals; small size) Up to 6 in many complexes (higher coordination possible) Reaction with water BF 3 hydrolyses; B 2 H 6 hydrolyses vigorously Al forms protective oxide; powdered Al reacts with steam; Al 2 O 3 amphoteric Acid/base behavior of oxide B 2 O 3 : acidic oxide Al 2 O 3 : amphoteric (reacts with both acids and bases) Aspect Feature Boron (B) Aluminium (Al) Boron vs Aluminium — key compound behavior Diagonal relationship (preserved) Boron and aluminium chemistry mirror each other closely. Boron is unique: very small, electron-deficient, lacks accessible d-orbitals and has max covalency 4; aluminium is metallic and forms different types of compounds (e.g., AlCl 3 as a strong Lewis acid, amphoteric Al 2 O 3 ). Electron-deficient bonding: Diborane (B2H6) Diborane ( B 2 H 6 ) is the flagship electron-deficient molecule. Each B has 3 valence electrons, so total 12 valence electrons must connect 8 atoms — not enough for all conventional 2c–2e bonds. Solution: two bridging H atoms form three-centre two-electron (3c–2e) bonds between B–H–B. These curved, off-axis bonds are called “banana bonds.” Terminal B–H bonds are normal 2c–2e bonds; the two bridges hold the B atoms together despite electron deficiency. gpt-image-2 2026-05-26T17:04:51.750Z Diborane structure showing two terminal and two bridging hydrogens; highlight the two 3c–2e banana bonds. Mechanism-style bonding diagram of diborane B2H6: show two BH2 units joined by two bridging H atoms. Label terminal B–H (2c–2e) and bridging B–H–B (3c–2e) banana bonds with charge delocalization. White background, vector style, red curved arrows for bonding description, labels with IUPAC/common terms. No embedded text captions. Hydrolysis of diborane gives boric acid and hydrogen — vigorous reaction. With ammonia, diborane first forms an adduct, then gives borazine (inorganic benzene) on heating: Low temperature: adduct (borane–ammonia). Heating the adduct gives borazine ( "inorganic benzene ") and hydrogen. Preparation of diborane: in the lab, reducing boron trifluoride with lithium aluminium hydride — a typical equation is: 4 ,BF 3 + 3 ,LiAlH 4 2 ,B 2H 6 + 3 ,LiF + 3 ,AlF 3 . remember Knowing B–H reactivity helps in organic hydroboration (separate syllabus link). Application of B–H bonds (anti-Markovnikov addition across C=C via hydroboration–oxidation). BF3 — classic Lewis acid Boron as Lewis acid (preserved) Boron trifluoride ( BF 3 ; SMILES: FB(F)F) is electron-deficient at boron (only 6 electrons around B), so it accepts a lone pair to complete its octet. With fluoride, it forms the tetrafluoroborate ion: BF 3 + F - BF 4 - . BF 3 and BF 4 - salt chemistry illustrates strong Lewis acidity and adduct formation — a theme repeated for AlCl 3 (Friedel–Crafts catalyst). trifluoroborane Lewis acid boron trifluoride ammonia azane Lewis base amminotrifluoroborane product adduct Ammonia donates a lone pair to electron-deficient boron in BF3, giving a coordinate covalent bond and octet completion at B. Lone pair on N of NH3 approaches the empty p-orbital on B in BF3. Curved arrow from N lone pair to B. Coordinate bond forms: BF3·NH3; boron achieves an octet; geometry around B changes from trigonal planar to approximately tetrahedral in the adduct. Lewis adduct formation: BF3 + NH3 Boric acid, H3BO3 — weak and monobasic (Lewis-acidic) Boric acid (orthoboric acid, H 3 BO 3 ) is a very weak acid and is monobasic — not because it donates H + , but because it accepts OH - from water to form [B(OH) 4] - , releasing H + into solution. Structurally, layers of trigonal planar BO 3 units are hydrogen-bonded, giving a sheet-like arrangement. Used in eyewash and as a mild antiseptic. 2026-05-26T17:04:51.643Z Diagram of boric acid layers: each B center trigonal planar with three OH groups; dashed lines for interlayer H-bonds. Labels: B, O, H; show one full layer and partial second layer. Clean vector, neutral palette, no embedded text. Layered structure of boric acid: trigonal planar BO3 units linked by hydrogen bonds between –OH groups. gpt-image-2 Boric acid acts as a Lewis acid by accepting OH - ; effective acidity is via this hydrolysis. Boric acid is monobasic. It does not donate protons directly; instead, it accepts OH - from water to give [B(OH) 4] - + H +. Boric acid is tribasic because it has three OH groups. Borax (Na2B4O7·10H2O) — preparation and bead test Borax is sodium tetraborate decahydrate, Na 2B 4O 7 10H 2O . Sources include natural tincal; it is also obtained industrially by treating calcium borates (e.g., colemanite) with sodium carbonate to form soluble sodium borate, which crystallizes as borax. Uses: glass and ceramics (especially borosilicate glass with B 2O 3 + SiO 2 ), laundry additive, flux, and mild antiseptic. Borax bead test: With transition metal salts, heating a borax bead in a flame produces characteristic colors (e.g., Co 2+ deep blue, Cr 3+ green). Four-panel vector: preparation of a borax bead on a loop; then beads colored deep blue (Co2+) and green (Cr3+) under different flames. Labels: metal ion, flame type. Clean lab-illustration style, white background. Borax bead test outcomes showing colored beads for Co2+ (blue) and Cr3+ (green) in oxidizing/reducing flames. 2026-05-26T17:04:51.957Z gpt-image-2 From bauxite to aluminium: Bayer’s purification + Hall–Héroult extraction Feedstock: bauxite (impure Al 2 O 3 with Fe 2 O 3 as red mud and silica). Bayer’s process (purification): Hot concentrated NaOH digests bauxite; Al 2 O 3 dissolves as sodium aluminate, leaving Fe 2 O 3 as red mud. The solution is cooled/neutralized or seeded to precipitate Al(OH) 3, which upon heating (calcination) gives pure alumina ( Al 2 O 3 ). Hall–Héroult process (extraction): Al 2 O 3 is dissolved in molten cryolite ( Na 3 AlF 6 ) around 950 °C to lower melting point and increase conductivity. Electrolysis in a carbon-lined cell deposits molten Al at the cathode; the carbon anodes are consumed, evolving CO 2 (and some CO). Backbone of global Al industry (India: NALCO, Hindalco). Bayer’s process flow: bauxite digestion in NaOH → sodium aluminate solution → precipitation of Al(OH)3 → calcination to Al2O3. Block-flow diagram with arrows: Bauxite + NaOH (digestion) → sodium aluminate liquor; separation of red mud; seeding/neutralization → Al(OH)3 precipitate; calcination furnace → Al2O3. Vector, labeled streams, white background. 2026-05-26T17:04:52.103Z gpt-image-2 Remember: Fe2O3 (red mud) is insoluble and removed early. Bauxite to pure alumina via digestion in NaOH, precipitation, and calcination. Hall–Héroult cell: carbon-lined cathode at the bottom, consumable carbon anodes dipping into molten cryolite–alumina bath; molten Al layer collects at bottom. Cutaway cross-section of Hall–Héroult cell. Label: carbon cathode liner, cryolite–Al2O3 molten bath (~ 950 C ), carbon anodes (CO2 evolution), molten Al pool, tap hole. Vector engineering schematic, clear labels, neutral palette. 2026-05-26T17:04:52.244Z gpt-image-2 Cathode (preserved key reaction) Anode (preserved key reaction) Electrolytic reduction of alumina dissolved in cryolite; C anodes consumed to CO2. Carbon anodes need periodic replacement due to oxidation. Hall–Héroult must-haves: molten cryolite (Na3AlF6), ~950 °C, carbon anode and cathode, molten Al at cathode, CO2 at anode. neet-alert Aluminium chemistry: amphoterism, AlCl3 Lewis acidity, and uses Amphoteric nature of Al (preserved) Al and Al 2 O 3 are amphoteric. With acids: 2Al + 6HCl 2AlCl 3 + 3H 2 . With bases (aq. NaOH): sodium aluminate forms (or [Al(OH) 4] - in solution). AlCl 3 is a strong Lewis acid and catalyzes Friedel–Crafts reactions (organic chemistry link). Uses: Al foil, aircraft and bicycle frames (duralumin: Al–Cu–Mg–Mn alloys), beverage cans, electrical cables. AlCl3 acts as a Lewis acid to generate strong electrophiles in aromatic substitution. Cross-link: Organic chemistry — electrophilic aromatic substitution. Al2O3 is a basic oxide. Al2O3 is amphoteric. It reacts with both acids and bases (e.g., dissolves in aq. NaOH to form aluminate). Thermite reaction — welding rails with molten iron A highly exothermic redox: aluminium reduces iron(III) oxide to iron, generating enough heat to produce molten iron on-site (railway track welding). Thermite is a classic high-temperature application of redox chemistry. Applied redox heat source; thermite mixtures burn intensely to generate molten metal. Illustration: crucible above rail joint, thermite mix ignited, molten Fe flowing into mold around rail ends. Labels: thermite mixture, ignition, slag, molten Fe. Diagrammatic, safety colors, no photorealism. Thermite welding setup on railway tracks: ignition of Al + Fe2O3 mix, molten iron filling the rail gap. 2026-05-26T17:04:53.217Z gpt-image-2 Down the group: oxidation states and the inert pair For B and Al, +3 is most stable. As you move to Ga and In, +1 appears and becomes more stable; for Tl, +1 is the most stable state. Reason: the inert pair effect — the ns 2 electrons are held more tightly (poorly shielded by inner d/f electrons), so they are less available for bonding in heavier elements. Increasing stability of lower oxidation states down Group 13 (and 14) due to the inert pair effect — Tl+ and Pb2+ are especially stable. Tl is mostly +1 because the inert pair ( ns 2 ) remains unbonded; +3 is less stable. Thallium is mostly +3 like Al and Ga. Environmental/medical tie-in: Stable lower oxidation states ( Tl + in Group 13, Pb 2+ in Group 14) change redox behavior and toxicity profiles — important in analytical chemistry and toxicology. remember Carbon catenation (preserved for cross-trend context) Group 13 order — “B Al, Ga In Tl”: Say it as “BAl — Ga-In — Tl” while pointing down the group. Diborane has two 3c–2e bridges (“banana bonds”). Boric acid is monobasic via [B(OH) 4] - formation. Al/Al2O3 amphoteric; AlCl3 is a Lewis acid (Friedel–Crafts). Hall–Héroult: cryolite bath, C anode consumed to CO2. Thermite: Fe 2O 3 + 2Al Al 2O 3 + 2Fe (welding). Tl + is most stable at the bottom (inert pair effect). High-yield checkpoints Bayer’s process Bauxite (Al2O3 + Fe2O3 + SiO2) Digest in hot conc. NaOH → sodium aluminate; remove Fe2O3 (red mud); precipitate Al(OH)3; calcine Pure alumina (Al2O3) Hall–Héroult Al2O3 (dissolved in Na3AlF6, ~ 950 C ) Electrolysis in carbon-lined cell; Al3+ → Al at cathode; O2− oxidizes C anodes to CO2 Molten aluminium metal Step Industrial extraction overview — Bayer vs Hall–Héroult Process Feedstock Principle (chemistry) Main product/output Ga shows higher IE than Al and a smaller radius due to d-block contraction; Tl anomalies reflect f-block effects and the inert pair. All periodic trends always apply without exceptions (e.g., IE strictly decreases down Group 13). Supplement: Acid reaction of Al metal Side-by-side line graphs: atomic radius and first ionization enthalpy for Groups 13 and 14. Emphasize Ga < Al radius and IE(Ga) > IE(Al); show Tl anomaly. Bright markers, vector style, no clutter. Group 13 vs 14 radii and IE anomalies (focused callout on Ga–Al and Tl–In) to reinforce the trend traps. 2026-05-26T17:04:54.430Z gpt-image-2 Preserved cross-link Preserved diagonal link Maximum Covalency of B and C Absence of vacant d -orbitals in the valence shell ( n=2 ) Limited to 4 ; heavier members ( Al to Tl , Si to Pb ) can expand covalency to 6 using d -orbitals. Ability to form p -p Multiple Bonds Small atomic size and high electronegativity allow effective lateral overlap of p -orbitals C forms C=C , C C , C=O ; heavier elements ( Si , Ge ) form weak p -p bonds or prefer single bonds. Exceptional Catenation Power of Carbon Very high C-C bond enthalpy ( 348 kJ mol -1 ) Decreases sharply down the group: C >> Si > Ge Sn ; Pb does not show catenation. Nature of Oxides ( CO 2 vs SiO 2 ) Carbon forms p -p double bonds with oxygen CO 2 is a discrete monomeric gas; SiO 2 is a high-melting giant covalent 3D network solid. Acid-Base Character of Oxides Higher electronegativity and small size of the first member B 2O 3 and CO 2 are acidic; Al 2O 3 is amphoteric; Tl 2O and PbO are basic. Formation of Boranes (Hydrides) High B-H bond strength and electron deficiency leading to multicenter bonds B forms a variety of stable covalent hydrides ( B 2H 6 ); Al forms polymeric (AlH 3) n . Metallic vs Non-metallic Character Highest ionization enthalpy and smallest size in the respective groups B and C are strictly non-metals; Al is a metal; Si and Ge are metalloids. Stability of Tetrahalides towards Hydrolysis CCl 4 cannot accept lone pairs from water CCl 4 is resistant to hydrolysis; SiCl 4 hydrolyzes easily due to availability of d -orbitals. Allotropy and Physical State Strong covalent bonding in 3D or 2D lattices B and C have exceptionally high melting points; Ga has a very low melting point ( 303 K ). Ion Formation Very high sum of first three/four ionization enthalpies B and C do not form B 3+ or C 4+ ions under normal conditions; Tl + and Sn 2+ /Pb 2+ are common. Understanding why the first member of a group behaves differently. Small size and no d-orbitals make Boron and Carbon the 'First-Row Rebels' who can't expand their homes (covalency). TREND Group 13 & 14 Anomalies Anomalous Property Reason Comparison with Group p-Block Elements Group 13 Group 14 Anomalous Properties Inorganic Chemistry NEET Preparation Element