Boron Dinitrogen Activation

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, it 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 Pisanoan 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.

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. 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 (OH)₄sup>−. 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 sup>10B(OH)₄sup>− 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, and β-rhombohedral is the most common and stable. An α-tetragonal phase also exists (α-T), but is very 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.

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. Very 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 very rare, its chemical reactions are of little significance practically speaking. The elemental form is not typically used as a precursor to compounds. Instead, the extensive inventory of 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 resembles silicon. 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 typically are 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.

Oxide derivatives

Boron-containing minerals exclusively exist as oxides of B(III), often associated with other elements. More than one hundred borate minerals are known. These minerals resemble silicates in some respect, although it is often found not only in a tetrahedral coordination with oxygen, but also in a trigonal planar configuration. The borates can be subdivided into two classes, anhydrous and the far more common hydrates. The hydrates contain B-OH groups and sometimes water of crystallization. A typical motif is exemplified by the tetraborate anions of the common mineral borax. The formal negative charge of the tetrahedral borate center is balanced by sodium (Na⁺). Some idea of the complexity of these materials is provided by the inventory of zinc borates, which are common wood preservatives and fire retardants: 4ZnO·B₂O₃·H₂O, ZnO·B₂O₃·1.12H₂O, ZnO·B₂O₃·2H₂O, 6ZnO·5B₂O₃·3H₂O, 2ZnO·3B₂O₃·7H₂O, 2ZnO·3B₂O₃·3H₂O, 3ZnO·5B₂O₃·14H₂O, and ZnO·5B₂O₃·4.5H₂O.

As illustrated by the preceding examples, borate anions tend to condense by formation of B-O-B bonds. Borosilicates, with B-O-Si, and borophosphates, with B-O-P linkages, are also well represented in both minerals and synthetic compounds.

Related to the oxides are the alkoxides and boronic acids with the formula B(OR)₃ and R₂BOH, respectively. Boron forms a wide variety of such metal-organic compounds, some of which are used in the synthesis of pharmaceuticals. These developments, especially the Suzuki reaction, was recognized with the 2010 Nobel Prize in Chemistry to Akira Suzuki.

Hydrides

Boranes and borohydrides are neutral and anionic compounds of boron and hydrogen, respectively. Sodium borohydride is the progenitor of the boranes. Sodium borohydride is obtained by hydrogenation of trimethylborate:
:
Sodium borohydride is a white, fairly air-stable salt.

Sodium borohydride converts to diborane by treatment with boron trifluoride:
:
Diborane is the dimer of the elusive parent called borane, BH₃. Having a formula akin to ethane's (C₂H₆), diborane adopts a very different structure, featuring a pair of bridging H atoms. This unusual structure, which was deduced only in the 1940s, was an early indication of the many surprises provided by boron chemistry.
class=skin-invert-image, left, 220px, Structure of diborane

Pyrolysis of diborane gives boron hydride clusters, such as pentaborane(9) and decaborane . A large number of anionic boron hydrides are also known, e.g. ₁₂H₁₂sup>2−. In these cluster compounds, boron has a coordination number greater than four. The analysis of the bonding in these polyhedra clusters earned William N. Lipscomb the 1976 Nobel Prize in Chemistry for "studies on the structure of boranes illuminating problems of chemical bonding". Not only are their structures unusual, many of the boranes are extremely reactive. For example, a widely used procedure for pentaborane

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