Boron is a chemical element; it has symbol B and atomic number 5. In its crystalline form it is a brittle, dark, lustrous metalloid; in its amorphous form it is a brown powder. As the lightest element of the boron group it has three valence electrons for forming covalent bonds, resulting in many compounds such as boric acid, the mineral sodium borate, and the ultra-hard crystals of boron carbide and boron nitride.
Boron is synthesized entirely by cosmic ray spallation and supernovas and not by stellar nucleosynthesis, so it is a low-abundance element in the Solar System and in the Earth's crust. It constitutes about 0.001 percent by weight of Earth's crust. It is concentrated on Earth by the water-solubility of its more common naturally occurring compounds, the borate minerals. These are mined industrially as evaporites, such as borax and kernite. The largest known deposits are in Turkey, the largest producer of boron minerals.
Elemental boron is found in small amounts in meteoroids, but chemically uncombined boron is not otherwise found naturally on Earth.
Several allotropes exist: amorphous boron is a brown powder; crystalline boron is silvery to black, extremely hard (9.3 on the Mohs scale), and a poor electrical conductor at room temperature (1.5 × 10−6 Ω−1 cm−1 room temperature electrical conductivity). The primary use of the element itself is as boron filaments with applications similar to carbon fibers in some high-strength materials.
Boron is primarily used in chemical compounds. About half of all production consumed globally is an additive in fiberglass for insulation and structural materials. The next leading use is in polymers and ceramics in high-strength, lightweight structural and heat-resistant materials. Borosilicate glass is desired for its greater strength and thermal shock resistance than ordinary soda lime glass. As sodium perborate, it is used as a bleach. A small amount is used as a dopant in semiconductors and reagent intermediates in the synthesis of organic fine chemicals. A few boron-containing organic pharmaceuticals are used or are in study. Natural boron is composed of two stable isotopes, one of which (boron-10) has a number of uses as a neutron-capturing agent.
Borates have low toxicity in mammals (similar to table salt) but are more toxic to arthropods and are occasionally used as insecticides. Boron-containing organic antibiotics are known. Although only traces are required, boron is an essential plant nutrient.
History
The word boron was coined from borax, the mineral from which it was isolated, by analogy with carbon, which boron resembles chemically.
Borax in its mineral form (then known as tincal) first saw use as a glaze, beginning in China circa 300 AD. Some crude borax traveled westward, and was apparently mentioned by the alchemist Jabir ibn Hayyan around 700 AD. Marco Polo brought some glazes back to Italy in the 13th century. Georgius Agricola, in around 1600, reported the use of borax as a flux in metallurgy. In 1777, boric acid was recognized in the hot springs (soffioni) near Florence, Italy, at which point it became known as sal sedativum, with ostensible medical benefits. The mineral was named sassolite after Sasso Pisano in Italy. Sasso was the main source of European borax from 1827 to 1872, when American sources replaced it. Boron compounds were rarely used until the late 1800s when Francis Marion Smith's Pacific Coast Borax Company first popularized and produced them in volume at low cost.Hildebrand, G. H. (1982) "Borax Pioneer: Francis Marion Smith." San Diego: Howell-North Books. p. 267
Boron was not recognized as an element until it was isolated by Sir Humphry Davy and by Joseph Louis Gay-Lussac and Louis Jacques Thénard. In 1808, Davy observed that electric current sent through a solution of borates produced a brown precipitate on one of the electrodes. In his subsequent experiments, he used potassium to reduce boric acid instead of electrolysis. He produced enough boron to confirm a new element and named it boracium. Gay-Lussac and Thénard used iron to reduce boric acid at high temperatures. By oxidizing boron with air, they showed that boric acid is its oxidation product. Jöns Jacob Berzelius identified it as an element in 1824.Berzelius produced boron by reducing a borofluoride salt; specifically, by heating potassium borofluoride with potassium metal. See: Berzelius, J. (1824) "Undersökning af flusspatssyran och dess märkvärdigaste föreningar" (Part 2) (Investigation of hydrofluoric acid and of its most noteworthy compounds), Kongliga Vetenskaps-Academiens Handlingar (Proceedings of the Royal Science Academy), vol. 12, pp. 46–98; see especially pp. 88ff. Reprinted in German as: Berzelius, J. J. (1824) "Untersuchungen über die Flußspathsäure und deren merkwürdigste Verbindungen", Poggendorff's Annalen der Physik und Chemie, vol. 78, pages 113–150. Pure boron was arguably first produced by the American chemist Ezekiel Weintraub in 1909.
Characteristics of the element
Isotopes
Boron has two naturally occurring and stable isotopes, 11B (80.1%) and 10B (19.9%). The mass difference results in a wide range of δ11B values, which are defined as a fractional difference between the 11B and 10B and traditionally expressed in parts per thousand, in natural waters ranging from −16 to +59. There are 13 known isotopes of boron; the shortest-lived isotope is 7B which decays through proton emission and alpha decay with a half-life of 3.5×10−22 s. Isotopic fractionation of boron is controlled by the exchange reactions of the boron species B(OH)3 and [B(OH)4]−. Boron isotopes are also fractionated during mineral crystallization, during H2O phase changes in hydrothermal systems, and during hydrothermal alteration of rock. The latter effect results in preferential removal of the [10B(OH)4]− ion onto clays. It results in solutions enriched in 11B(OH)3 and therefore may be responsible for the large 11B enrichment in seawater relative to both oceanic crust and continental crust; this difference may act as an isotopic signature.
The exotic 17B exhibits a nuclear halo, i.e. its radius is appreciably larger than that predicted by the liquid drop model.
NMR spectroscopy Both 10B and 11B possess nuclear spin. The nuclear spin of 10B is 3 and that of 11B is . These isotopes are therefore of use in nuclear magnetic resonance spectroscopy, and spectrometers specially adapted to detecting the boron-11 nuclei are available commercially. The 10B and 11B nuclei also cause splitting in the resonances of attached nuclei.
Allotropes
Boron forms four major allotropes: α-rhombohedral and β-rhombohedral (α-R and β-R), γ-orthorhombic (γ), and β-tetragonal (β-T). All four phases are stable at ambient conditions. β-rhombohedral is the most stable and the most common. An α-tetragonal phase also exists (α-T) but is difficult to produce without significant contamination. Most of the phases are based on B12 icosahedra, but the γ phase can be described as a rocksalt-type arrangement of the icosahedra and B2 atomic pairs. It can be produced by compressing other boron phases to 12–20 GPa and heating to 1500–1800 °C; it remains stable after releasing the temperature and pressure. The β-T phase is produced at similar pressures, but higher temperatures of 1800–2200 °C. The α-T and β-T phases might coexist at ambient conditions, with the β-T phase being the more stable. Compressing boron above 160 GPa produces a boron phase with an as yet unknown structure, and this phase is a superconductor at temperatures below 6–12 K.
Boron phaseα-Rβ-Rγβ-TSymmetryRhombohedralRhombohedralOrthorhombicTetragonalAtoms/unit cell12~10528Density (g/cm3)2.462.352.522.36Vickers hardness (GPa)424550–58Bulk modulus (GPa)185224227Bandgap (eV)21.62.1
Atomic structure Atomic boron is the lightest element having an electron in a p-orbital in its ground state. Its first three ionization energies are higher than those for heavier group III elements, reflecting its electropositive character.
Chemistry of the element
Preparation Elemental boron is rare and poorly studied because the pure material is extremely difficult to prepare. Most studies of "boron" involve samples that contain small amounts of carbon. Highly pure boron is produced with difficulty because of contamination by carbon or other elements that resist removal.
Some early routes to elemental boron involved the reduction of boric oxide with metals such as magnesium or aluminium. However, the product was often contaminated with borides of those metals. Pure boron can be prepared by reducing volatile boron halides with hydrogen at high temperatures. Ultrapure boron for use in the semiconductor industry is produced by the decomposition of diborane at high temperatures and then further purified by the zone melting or Czochralski processes.
Reactions of the element Crystalline boron is a hard, black material with a melting point of above 2000 °C. Crystalline boron is chemically inert and resistant to attack by boiling hydrofluoric or hydrochloric acid. When finely divided, it is attacked slowly by hot concentrated hydrogen peroxide, hot concentrated nitric acid, hot sulfuric acid, or hot mixture of sulfuric and chromic acids.
Since elemental boron is rare, its chemical reactions are of little significance. The elemental form is not typically used as a precursor to compounds. Instead, boron compounds are produced from borates.
When exposed to air, under normal conditions, a protective oxide or hydroxide layer forms on the surface of boron, which prevents further corrosion. The rate of oxidation of boron depends on the crystallinity, particle size, purity and temperature. At higher temperatures boron burns to form boron trioxide: 4 B + 3 O2 → 2 B2O3
Chemical compounds
General trends In some ways, boron is comparable to carbon in its capability to form stable covalently bonded molecular networks (even nominally disordered (amorphous) boron contains boron icosahedra, which are bonded randomly to each other without long-range order.). In terms of chemical behavior, boron compounds resemble silicon due to their diagonal relationship. Aluminium, the heavier congener of boron, does not behave analogously to boron: it is far more electropositive, it is larger, and it tends not to form homoatomic Al-Al bonds. In the most familiar compounds, boron has the formal oxidation state III. These include the common oxides, sulfides, nitrides, and halides, as well as organic derivatives.
Boron compounds often violate the octet rule.
Halides Boron forms the complete series of trihalides, i.e. BX3 (X = F, Cl, Br, I). The trifluoride is produced by treating borate salts with hydrogen fluoride, while the trichloride is produced by carbothermic reduction of boron oxides in the presence of chlorine gas:
The trihalides adopt a planar trigonal structures, in contrast to the behavior of aluminium trihalides. All charge-neutral boron halides violate the octet rule; hence, they are typically Lewis acidic. For example, boron trifluoride (BF3) combines eagerly with fluoride sources to give the tetrafluoroborate anion, BF4−. Boron trifluoride is used in the petrochemical industry as a catalyst. The halides react with water to form boric acid. Other boron halides include those with B-B bonding, such as B2F4 and B4Cl4. Larger polyhedral boron(I) halides also exist, e.g. B7Br7, B8Cl8 and B9I9.
Oxide derivatives