Read Diborane's Bridge Bonds Before Boron Chemistry
Group 13 chemistry is dominated by two features: boron's electron deficiency (only 6 valence electrons, driving Lewis acidity and unusual bridge bonding) and the inert pair effect further down the group (making thallium's +1 state as stable as its +3).
Group 13: The Boron Family is the central idea in this question pattern. Identify the condition in the stem before applying the method below.
Diborane's bridge bonds: 4 normal + 2 banana bonds
B₂H₆ has four terminal 2-centre-2-electron B-H bonds and two bridging 3-centre-2-electron B-H-B bonds — this electron-deficient bridging is exactly why diborane exists as a dimer while BCl₃ doesn't: chlorine is too large to fit in the bridging position, and even if it fit, BCl3's boron already satisfies its octet through pπ-pπ back-bonding from chlorine's lone pairs, removing any driving force to dimerise.
Boron's electron deficiency makes it a strong Lewis acid
BF₃, BCl₃ and SiF₄ are all Lewis acids (empty orbital available to accept a lone pair), but ethylene (C₂H₄) is not — a molecule needs a genuinely empty, accessible orbital to act as a Lewis acid, not just an electron-poor-looking formula. Adduct formation with NH₃ (BX₃ + NH₃ → BX3·NH₃) releases more heat for the LESS electronegative halogens (BI₃ > BBr₃ > BCl₃ > BF₃ in adduct-formation enthalpy) because B-X bond weakening/rehybridisation costs less energy when X is a poorer π-donor to begin with.
Boric acid is monobasic — through water, not through releasing its own H⁺
H₃BO₃ is a weak MONOBASIC Lewis acid, not a tribasic Brønsted acid as its formula might suggest — it acts as an acid by accepting OH⁻ from water (B(OH)₃ + H₂O → [B(OH)₄]- + H⁺), not by donating one of its own three O-H protons. It's soluble in hot water, has a planar layered structure held together by hydrogen bonds, and is used medicinally as a mild antiseptic for eye infections.
The inert pair effect: why thallium prefers +1 over +3
Going down group 13, the ns₂ electron pair becomes increasingly reluctant to participate in bonding (poor shielding by the intervening d/f electrons, plus relativistic effects for the heaviest elements) — this is why thallium readily forms Tl⁺ (using only the single p-electron) alongside the 'expected' Tl³⁺, while boron and aluminium show essentially only the +3 state.
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- ✓Diborane's structure: 4 terminal 2c-2e B-H bonds + 2 bridging 3c-2e B-H-B bonds — a direct consequence of boron's electron deficiency.
- ✓A true Lewis acid needs a genuinely empty, accessible orbital — an electron-poor-looking formula isn't enough on its own (C₂H₄ is not a Lewis acid).
- ✓Boric acid is monobasic via OH⁻ acceptance from water, not via donating its own O-H protons — despite having three O-H groups.
- ✓The inert pair effect makes the ns₂ pair increasingly reluctant to bond down the group, so Tl⁺ is as stable as (or more stable than) Tl³⁺.
Don't assume boric acid is tribasic just because it has three O-H groups in its formula — it's monobasic, and acts as an acid via a completely different mechanism (accepting OH⁻ from water) than a normal Brønsted acid donating its own protons.
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