Group 13 (Boron Family)
Group 13 of the periodic table — the Boron Family — comprises boron (B), aluminium (Al), gallium (Ga), indium (In), and thallium (Tl). All members have the general valence electron configuration ns²np¹, giving a group oxidation state of +3. However, a critical trend runs through this group: as atomic number increases, the +1 oxidation state becomes increasingly stable relative to +3. This phenomenon, called the inert pair effect, arises because the two ns² electrons become progressively harder to ionise as nuclear charge and relativistic contraction increase down the group. Thallium, for instance, is far more commonly found as Tl⁺ than as Tl³⁺ in its stable compounds. Boron stands apart from every other member of Group 13 in a collection of properties collectively called the boron anomaly. Unlike aluminium, which is a metallic conductor, boron is a metalloid with extremely high melting point (~2076 °C) and very low electrical conductivity. Its small atomic radius and high charge density give it an enormous polarising power, meaning boron forms predominantly covalent bonds rather than ionic ones. Boron also has only three valence electrons — it forms compounds such as BF₃ and BCl₃ that are electron-deficient (only six electrons around boron), making them strong Lewis acids capable of accepting an electron pair from donors such as ammonia or fluoride ions. Additionally, boron is the only non-metal in the group and the only element in Group 13 that does not form simple M³⁺ aqueous ions. A particularly fascinating consequence of boron's electron deficiency is the existence of multi-centre covalent bonds. In diborane (B₂H₆), the molecule cannot be described by ordinary two-centre, two-electron (2c-2e) bonds because there are not enough electrons. Instead, it contains three-centre, two-electron (3c-2e) bonds — sometimes called banana bonds or bridging bonds — where a single pair of electrons holds three atoms together simultaneously. This unusual bonding makes boron hydrides (boranes) a whole sub-discipline of inorganic chemistry. The borax bead test, a classical qualitative analytical technique, exploits the Lewis acid chemistry of boron to identify transition metals by the characteristic colours of their borate glasses.
- The Inert Pair Effect
- The Boron Anomaly (Diagonal Relationship)
- Structure of Diborane and 3c-2e Bonds
- Borax Bead Test
The Inert Pair Effect
In Period 2 (boron) and Period 3 (aluminium), the ns² electrons are readily removed and the +3 state dominates. Moving down to Ga, In, and Tl, the ns² electrons experience increased effective nuclear charge and relativistic stabilisation. The energy cost of removing these two electrons becomes prohibitively large compared to the lattice or solvation energy gained. Consequently, the +1 oxidation state — where only the np¹ electron is removed or shared — grows more stable. Tl⁺ compounds (e.g., TlCl, Tl₂O) are thermodynamically more stable than Tl³⁺ compounds. The standard reduction potential E°(Tl³⁺/Tl⁺) = +1.25 V, meaning Tl³⁺ is a fairly strong oxidising agent.
The Boron Anomaly (Diagonal Relationship)
Boron resembles silicon diagonally across the periodic table more than it resembles aluminium below it. Key similarities: (a) Both B and Si form covalent, network-solid oxides (B₂O₃ and SiO₂) with high melting points. (b) Both form volatile, flammable hydrides (B₂H₆ and Si₂H₆). (c) Both oxides are acidic and dissolve in alkalis to form borate and silicate ions respectively. (d) Both form halides that are completely hydrolysed by water: BCl₃ + 3H₂O → H₃BO₃ + 3HCl, and SiCl₄ + 2H₂O → SiO₂ + 4HCl. Aluminium, by contrast, shows amphoteric behaviour and forms stable aqueous Al³⁺ ions. Boron's high charge-to-radius ratio (Z/r) of approximately 150 × 10⁹ C m⁻¹ causes extreme polarisation of surrounding electron clouds, forcing covalent character into bonds that would be ionic for heavier analogues.
Structure of Diborane and 3c-2e Bonds
Diborane (B₂H₆) has only 12 valence electrons (3 from each B + 1 from each of 6 H). A conventional ethane-like structure would require 14 electrons — two more than available. The actual structure has: (a) four terminal B–H bonds (2c-2e, normal), and (b) two bridging B–H–B units. Each bridging B–H–B unit is a three-centre, two-electron bond: one molecular orbital formed by linear combination of one H(1s) orbital and one sp³-hybrid orbital from each boron accommodates a single electron pair that simultaneously bonds all three atoms. The molecule has D₂h symmetry. Bond angles: Hb–B–Hb = 97°, Ht–B–Ht = 122°, where Hb = bridging hydrogen, Ht = terminal hydrogen. Bond lengths: B–Ht = 119 pm (normal B–H), B–Hb = 133 pm (longer, weaker, 3c-2e). Diborane is prepared by: 2BCl₃ + 6NaH → B₂H₆ + 6NaCl, or commercially by reaction of BF₃ with LiAlH₄.
Borax Bead Test
Borax (sodium tetraborate decahydrate, Na₂B₄O₇·10H₂O) when heated on a platinum wire loop first loses water of crystallisation, then melts and fuses into a glassy bead of boron(III) oxide mixed with sodium metaborate (NaBO₂). When a tiny amount of a transition metal salt is introduced into this hot bead, the metal oxide dissolves in the molten borate matrix forming a coloured glass of characteristic metal borate. Colours (in oxidising flame / reducing flame): Cobalt → deep blue / deep blue; Chromium → emerald green / green; Manganese → violet/amethyst / colourless; Iron → yellow-brown / green; Copper → green / red (Cu metal appears). The reaction for borax fusion is: Na₂B₄O₇ → 2NaBO₂ + B₂O₃ (on strong heating). Metal oxide incorporation: CoO + B₂O₃ → Co(BO₂)₂ (cobalt metaborate glass, blue). The test exploits boron's role as a network-forming Lewis acid oxide in molten glass chemistry.
The Key Equation
- \ce{BCl_3} — Boron trichloride (mol): Planar, sp² hybridised Lewis acid; electron-deficient boron has only 6 valence electrons.
- \ce{B(OH)_3} — Boric acid (orthoboric acid) (mol): Weak Lewis acid (accepts OH⁻ from water rather than donating H⁺); planar trigonal structure, pKₐ ≈ 9.24.
- \ce{HCl} — Hydrogen chloride (mol): Strong acid produced during hydrolysis; evolves as gas or dissolves in water as hydrochloric acid.
Worked Example
Problem. Diborane (B₂H₆) reacts with excess ammonia. Predict the product and explain whether the bond formed is a 3c-2e bond or a conventional 2c-2e bond. Also, count all the valence electrons in B₂H₆ and explain why a conventional single-bonded structure is impossible.
Answer. B₂H₆ reacts with 2NH₃ to give 2(H₃B·NH₃). The 3c-2e bridging bonds of diborane are broken; the B–N bond in the product is a conventional 2c-2e dative bond.
Origins · 1 minHerbert C. Brown and how this was discovered
Herbert C. Brown (born Herbert Brovarnik) won the Nobel Prize in Chemistry in 1979 for his discovery and development of organoboranes as reagents in organic synthesis. His work on hydroboration — the addition of boron hydrides (derived from diborane) across carbon–carbon double bonds — opened entirely new synthetic routes to alcohols, alkenes, and other functional groups with precise stereochemical control. He also made fundamental contributions to understanding the Lewis acid character of boron compounds and the concept of steric strain. Brown was so dedicated to chemistry that he reportedly proposed to his wife with the words 'If you marry me, you will not live in poverty; I will receive the Nobel Prize.' He kept his promise — nearly 40 years later.
Related Concepts
- Group 14 — The Carbon Family — Adjacent p-block group; illustrates how inert pair effect and diagonal relationships shift from non-metal (C) to metalloid (Si, Ge) to metal (Sn, Pb) character, directly comparable to the B→Tl trend in Group 13.
- Lewis Acids and Bases — Boron compounds such as BF₃ and B₂H₆ are the archetypal Lewis acids used in virtually every introductory treatment of electron-pair donor–acceptor chemistry.
- Covalent Bonding and Molecular Orbital Theory — The 3c-2e bonds in diborane can only be properly understood using molecular orbital (MO) theory — three atomic orbitals combine to give bonding, non-bonding, and antibonding MOs.
- Qualitative Analysis — Flame and Bead Tests — The borax bead test is a classic qualitative technique for identifying transition metals; understanding it requires both Group 13 chemistry and knowledge of transition metal colour and oxidation states.