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Boron



Symbol: B; m.p. 2300°C; b.p. 2658°C; r.d. 2.34 (20°C); p.n. 5; r.a.m. 10.811.
A hard rather brittle metalloid element of group 3 (formerly IIIA) of the periodic table. It has the electronic structure 1s²2s²2p¹. 
Boron is of low abundance (0.0003%) but the natural minerals occur in very concentrated forms as either borax (Na₂B₄O₇.10H₂O) or colemanite (Ca₂B₆O₁₁). The element is obtained by conversion to boric acid followed by dehydration to B₂O₃ then reduction with magnesium. High-purity boron for semiconductor applications is obtained by conversion to boron trichloride, which can be purified by distillation, then reduction using hydrogen. Only small quantities of elemental boron are needed commercially; the vast majority of boron supplied by the industry is in the form of borax or boric acid. As boron has a small atom and has a relatively high ionization potential its compounds are predominantly covalent; the ion B³⁺ does not exist. Boron does not react directly with hydrogen to form boron hydrides (boranes) but the hydrolysis of magnesium boride does produce a range of boranes such as B₄H₁₀, B₃H₉, and B₆H₁₀. Thermal decomposition of these higher boranes produces, among other things, the simplest borane, B₂H₆ (diborane). The species BH₃ is only a short-lived reaction intermediate. 
Finely divided boron burns in oxygen above 600°C to give the oxide, B₂O₃, an acidic oxide which will dissolve slowly in water to give boric acid (B(OH)₃) and rapidly in alkalis to give borates such as Na₂B₄O₇. A number of polymeric species with B–B and B–O links are known, e.g. a lower oxide (BO)x, and a polymeric acid (HBO₂)n. Although the parent acid is weak, many salts containing borate anions are known but their stoichiometry gives little indication of their structure, many of which are cyclic or linear polymers. These contain both BO₃ planar groups and BO₄ tetrahedra. Boric acid and the borates give a range of glassy substances on heating; these contain cross-linked B–O–B chains and nets. In the molten state these materials react with metal ions to form borates, which on cooling give characteristic colors to the glass. See borax-bead test. Boron reacts with nitrogen on strong heating (1000°C) to give boron nitride, a slippery white solid with a layer structure similar to that of graphite, i.e., hexagonal rings of alternating B and N atoms. The material has an extremely high melting point and is thermally very stable but there is sufficient bond polarity in the B–N links to permit slow hydrolysis by water to give ammonia. There is also a ‘diamond-like’ form of B–N which is claimed to be even harder than diamond. Elemental boron reacts directly with fluorine and chlorine but for practical purposes the halides are obtained via the BF₃ route from boric acid:


B₂O₃ + 3CaF₂ + 3H₂SO₄ → 3CaSO₄ + 3H₂O + 2BF₃


BF₃ + AlCl₃ → AlF₃ + BCl₃


These halides are all covalent molecules, which are all planar and trigonal in shape. Boron halides are industrially important as catalysts or promoters in a variety of organic reactions including polymerization and Friedel–Crafts type alkylations. The decomposition of boron halides in atmospheres of hydrogen at elevated temperatures is also used to deposit traces of pure boron in semiconductor devices. Boron forms a range of compounds with elements that are less electronegative than itself, called borides. Borides such as ZrB₂ and TiB₂ are hard refractory substances, which are chemically inert and have remarkably high electrical conductivities. Borides have a wide range of stoichiometries, from M₄B through to MB₆, and can exist in close-packed arrays, chains, and two-dimensional nets. Natural boron consists of two isotopes, ¹⁰B (18.83%) and ¹¹B (81.17%). These percentages are sufficiently high for their detection by splitting of infrared absorption or by n.m.r. spectroscopy. Both borax and boric acid are used as mild antiseptics and are not regarded as toxic; boron hydrides are however highly toxic.