Boron is a hard, brittle metalloid used in borosilicate glass, detergents, semiconductors, and as a neutron absorber in nuclear applications.
The ground-state configuration is \([He]2s^2 2p^1\). With only three valence electrons, boron often forms electron-deficient compounds that favor covalent bonding and multicenter bonding, leading to properties intermediate between metals and nonmetals—hence its classification as a metalloid.
Boron in BF3 has an incomplete octet (six valence electrons). It readily accepts a lone pair from donors like NH3 to complete its octet:
\(\mathrm{BF_3 + :NH_3 \rightarrow F_3B\!\leftarrow NH_3}\)
The strong B–F bond and the ability of fluorine to stabilize negative charge also enhance the Lewis acidity of boron trifluoride.
Boranes are hydrides of boron (e.g., B2H6, B4H10). Due to electron deficiency, they feature three-center two-electron (3c–2e) bonds where two electrons are shared by three atoms.
In diborane, two bridging hydrogens form banana bonds between the two boron atoms, explaining its structure and reactivity.
Boron oxide (B2O3) is introduced into silica networks to make borosilicate glass. B2O3 acts as a network former, reducing the coefficient of thermal expansion and increasing resistance to thermal shock and chemical attack. This is why laboratory glassware and cookware (e.g., Pyrex®) often use borosilicate.
Borax (Na2B4O7·10H2O) forms a molten glassy bead that can dissolve metal oxides to give characteristic colors—useful in qualitative analysis.
In detergents, borates act as builders and buffering agents, softening water and enhancing surfactant action.
The isotope \(^{10}\!\mathrm{B}\) has a high neutron-capture cross section, undergoing reactions like:
\(\mathrm{^{10}B + n \rightarrow ^7Li + \alpha + \gamma}\)
This efficiently removes thermal neutrons, so boron (or boron carbide, B4C) is used in control rods and shielding.
Boric acid, \(\mathrm{H_3BO_3}\), is a weak Lewis acid that accepts hydroxide from water rather than donating a proton:
\(\mathrm{B(OH)_3 + H_2O \rightleftharpoons B(OH)_4^- + H^+}\)
Thus, it behaves as a Lewis acid by accepting an electron pair from OH−, generating acidity in solution.
Boron has three valence electrons vs. silicon’s four. When substituted into the Si lattice, boron creates an acceptor level just above the valence band. Thermal excitation promotes electrons into this level, leaving behind holes (positive charge carriers), producing p-type conductivity.
Elemental boron can be produced by reducing boron oxide with magnesium, followed by acid leaching to remove MgO:
\(\mathrm{B_2O_3 + 3\,Mg \rightarrow 2\,B + 3\,MgO}\)
High-purity boron requires further purification; crystalline forms show high hardness and complex icosahedral structures.