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1812, from borax + ending abstracted unetymologically from carbon (it resembles carbon). Originally called boracium by Sir Humphrey Davy because it was drawn from boracic acid. Related: Boric.
Description of the substance
Introduction to Boron
Boron is one of the simplest of atoms. The only simpler ones are hydrogen, helium, lithium and beryllium. Boron has chemical symbol B, atomic number 5, and occurs naturally as 80% B11 and 20% B10. The latter isotope has a high cross section for thermal neutron absorption, 3800 barns. Thermal neutron counters are often filled with BF3 gas. The gamma ray from the neutron capture reaction B10(n,-)B11* followed by decay of the B11* to an a plus Li7 produces ionization which is then detected. Boron is also used in reactor control rods. This is a nuclear property of boron, and has nothing at all to do with its chemistry. The atomic weight of boron is 10.81.
Boron is found in a variety of similar minerals all related to borax, sodium tetraborate, Na2B4O7·10H2O. The name comes from the Arabic buraq, "white." Borax is the same in French and German as in English, but the element is bor. In Spanish, the words are bóraxo and boro. It is a relatively rare element in the earth's crust, representing only 0.001%. In the United States, borax is found in large amounts in California, in Searles Lake brines and in the Mojave desert. It is also found in Turkey, South America and other places. The natural deposits are dried-up lake beds. Molten borax reacts with metal oxides to form borates that dissolve in the melt, so it is a useful as a welding and soldering flux, and in colored enamels for iron. In fact, this was the earliest use of borax, as a pottery glaze. This same property is used for borax bead tests in chemistry, where the characteristic colors produced in a transparent borax drop melted on a loop of platinum wire in a bunsen burner flame are observed. Blue, for example, is the color of cobalt; green, of chromium. The color can differ in oxidizing (blue) and reducing (yellow) flames.
The electron configuration of boron is 1s22s22p. It has only three electrons to work with, so the ion is unpolarizable, and does not hydrate. For this reason, boron is not eager to donate electrons in an electrovalent bond, and can also not accept them easily. Therefore, most of its bonds are covalent, and even forms half-bonds in which only one electron is shared covalently, not the usual two. This gives boron an apparent valence of +6 that we shall see in some interesting compounds. The first ionization potential is 8.30 V, which is not unusually high.
The other commonly-used boron compounds are orthoboric acid, or simply boric or boracic acid, H3BO3, and boron trioxide, B2O3, its anhydride. Note that the formula for boric acid can be written B(OH)3, as boron hydroxide. If boron were a normal metal, the hydroxyl ions would separate in water, creating the trivalent boron ion B+++. This, however, does not happen to the smallest degree, and boron does not form ionic bonds. Moreover, boric acid is not gelatinous like aluminium hydroxide, but crystallizes nicely. Boron trifluoride, BF3 is not an ionic compound like NaF, which has no ions in its crystals. Instead, the hydrogens are lost long before the oxygens. In aqueous solution, boric acid is a very weak acid, weaker even than carbonic. Its first ionization product is 6.4 x 10-10. Boric acid solution (solubility 6.35% at room temperature, 27.6% at 100°C) is used as an antiseptic, especially as an eyewash. Boric acid can be produced from borax by treatment with H2SO4, and boron trioxide by heating boric acid. These are the initial processes in the synthesis of all boron compounds.
Elemental boron was discovered by Davy, Gay-Lussac and Thénard in 1808. They produced metallic potassium by electrolysis, and then used it to reduce borates to impure boron. Davy called the element boracium, the Frenchmen bore. It can also be obtained in impure form by reduction of the oxide B2O3 by magnesium, or in pure form by the reduction of BCl3 by hydrogen on hot filaments. The first pure boron was produced by Weintraub in 1909. Ordinary boron is a brown-black amorphous powder. Pure boron can be made as extremely hard yellow monoclinic crystals that are a semiconductor resembling silicon. The band gap is 1.50 or 1.56 eV, which is higher than that of either silicon or germanium, and very similar to that of diamond. Crystalline boron is an insulator at low temperatures, but becomes a conductor at elevated temperatures, as would be expected as carriers are thermally excited into the conduction band. Fabrication difficulties have so far prevented the use of boron as a semiconductor. The density of crystalline boron is 2.34 g/cc, of amorphous boron, 2.37. It melts at 2300°C and boils at 2550°C (some sources say 2040°C and 4100°C), so it is a very refractory substance. Boron fibers have been used in composite materials because of their great strength.
Boron gives a blue-green flame, and the brown amorphous form is often used in pyrotechical devices for this purpose. The color can be distinguished from the emerald green of copper. Boron is also used in borosilicate glasses, which are 12%-15% B2O3, 80% SiO2, and 2% Al2O3. Sodium, potassium, magnesium and calcium oxides are kept to minimal amounts, since it is their exclusion that gives the glass its desirable properties. "Pyrex" is a common trade name for a borosilicate glass. This glass is chemically resistant, and has a small coefficient of thermal expansion. Boron carbide, B4C, harder than SiC, is formed by decomposing B2O3 with carbon in the electric furnace: 2B2O3 + 7C ? B4C + 6CO. It was first suggested for commercial use in 1934, and is an excellent abrasive. Boron nitride, BN, is another very hard compound, used in cutting tools. There are numerous other borides with complex structures. Boron is also used in porcelain enamels for iron, and for tiles and sanitary ware.
An interesting test for boron and borates uses volatile (CH3)3BO3, which burns with a green flame. Combine the solution to be tested with methanol, then add concentrated sulphuric acid to make the methyl borate. Another test uses turmeric paper, which turns reddish brown when moistened with boric acid or borates. If then a drop of ammonium hydroxide or sodium hydroxide is put on the paper, it will turn dark green, blue or black. To make turmeric paper, wash some ground turmeric in water. Then digest the washed powder with alcohol, and filter. Soak unsized white paper in the filtrate, then dry and cut into test strips. The active ingredient in turmeric paper may be curcumin, which seems to act the same way. It is available in an alcoholic 0.1% solution for this purpose.
We have mentioned that boron would rather lose the H in B(OH)3 than the oxygen. Like silicon, which has a similar persuasion, boron can form chains like -B-O-B-O-B-, where each intermediate boron has a free valence to work with. This is seen in tetraboric acid, O=B-O-B(OH)-O-B(OH)-O-B=O, or H2B4O7, which can be derived from B(OH)3 by dehydration: 4H3(BO)3 ? H2B4O7 + 5H2O. Dehydration of boric acid also gives boron trioxide: 2B(OH)3 ? B2O3 + 3H2O. Boron trioxide can also be written O=B-O-B=O, where the boron-oxygen chain again appears.
Some compounds can be considered derivatives of the theoretical acid HBO2, called borates. This metaboric acid is boric acid less one water molecule. If it loses another water, it becomes B2O3, which we have already seen. Boron trioxide in water becomes boric acid. The use of borax as a flux depends on reactions like O=B-O-B(ONa)-O-B(ONa)-O-B=O (borax) + NiO ? 2NaBO2 + Ni(BO2)2.
Boron-oxygen chains also form linked planar structures that have the appearance of fence wire with hexagonal openings. Oxygen atoms project radially to make bonds with other atoms.
The icosahedron ("twenty-base") is one of the five Platonic solids, a regular convex polyhedron with 20 faces, 30 edges and 12 corners. In the absence of degeneracy, a symmetrical structure will have the lowest energy, and the regularity is possible because of the identity of the atoms. The tetrahedron, square, octahedron and dodecahedron also appear in molecular structures. It is remarkable that boron finds the icosahedron adapted for its purposes. Excellent drawings of the structures we discuss here are found in Pauling and Hayward. The icosahedral structures are exhibited by compounds known as boranes.
Just after world war II, the U. S. military was interested in developing an advanced aviation fuel for jets and ramjets that would replace the hydrocarbon-based JP-4 and similar fuels. Someone suggested that the then poorly-known boranes would give a better power-to-weight ratio, and Project ZIP was launched in 1952. Government money rained down on research projects dealing with boron, and for roughly a decade there was near hysteria in the scramble for research money, and many things were, incidentally, learned about boron. The boranes turned out to be not the powerhouses that were anticipated, were hard to manufacture and store, tended to start fires, and were also extremely toxic. Military enthusiasm cooled, and the research money went elsewhere. However, we found out the interesting structures of boranes, and a good deal of boron chemistry. Fifty years later, airplanes still burn JP-4.
A boron atom at an icosahedral vertex can make five equal bonds to its nearest neighbors, leaving one bond to stick out radially. Boron-boron bonds are half-bonds of length 0.180 nm. If we put one boron atom at each vertex, each with one proton, and assume that all bonds are half-bonds, then we have the neutral structure B12H12. However, the protons would like to have more electrons in their bonds, and overall 6 electrons are lacking for a full single bond. This would give far too large a charge to the ion, so the structure only attracts 2 extra electrons, which are spread over the 12 bonds to the protons. It happens that two extra electrons are ideal for resonance stablization of the ion. The proton bonds wind up at 0.120 nm. The substance potassium dodecaborohydride contains this ion in its crystals.
Boron will not form BH3. Instead, it forms larger borohydrides called boranes. In decaborane, B10H14, only 10 of the 12 vertices of the icosahedron are occupied. The extra 4 hydrogens are each bonded to two adjacent boron atoms, replacing the "missing" ones. The 10 radial hydrogens have bonds 0.118 nm long, while the bridging hydrogens have bonds 0.135 nm long. In tetraborane, B4H10, only four vertices are occupied, and there are 10 bridging hydrogens. Tetraborane has a foul smell. Diborane, B2H6 has no boron-boron bonds, but two bridging hydrogens out of plane in the planar molecule.
Davy might have made a borane when he treated boron with hydrochloric acid. The gas that was evolved burnt with a blue-green flame. Boranes react vigorously with oxygen, so they have been considered as candidates for rocket fuels. They are examples of electron-deficient compounds, since there are more bonding orbitals available than there are electrons to fill them. Carboranes are boranes where some of the borons are replaced by carbons. It is not easy to make boranes, and there was a considerable interval between Davy and the first recognized boranes.
One form of crystalline boron is tetragonal. The unit cell of this lattice consists of 4 icosahedrons of boron and two extra boron atoms, 50 in all. The radial bonds go to other icosahedrons or to the extra atoms, which show tetrahedral bonds. The bond length is 0.180 nm, as in the boranes.
The element is not found free in nature, but occurs as orthoboric acid usually in certain volcanic spring waters and as borates in boron and colemantie. Ulexite, another boron mineral, is interesting as it is nature's own version of "fiber optics." Important sources of boron are the ore rasorite (kernite) and tincal (borax ore).