Bromine Br

a chemical element in Group VII of the Mendeleev periodic system. It is a member of the halogen family. Atomic number, 35; atomic weight, 79.904. It is a reddish-brown liquid with a strong, unpleasant smell. It was discovered in 1826 by the French chemist A. J. Balard while doing research on Mediterranean brine deposits. The name comes from the Greek bromos (“stench”). Bromine in nature consists of two stable isotopes, ⁷⁹Br (50.54 percent) and ⁸¹Br (49.46 percent). The most interesting of the artificially produced radioactive isotopes of bromine is ⁸⁰Br. I. V. Kurchatov discovered and studied the phenomenon of isomerism of atomic nuclei of ⁸⁰Br.

Natural occurrence. The earth’s crust is estimated to contain 10¹⁵ to 10¹⁶ tons of bromine (1.6 × 10⁻⁴ percent by weight). For the most part bromine occurs dispersed in magmatic rock, as well as in widely distributed halides. Bromine invariably occurs with chlorine. Bromine salts (NaBr, KBr, and MgBr₂) are found in deposits of chlorides (up to 0.03 percent Br in common salt and up to 0.3 percent in potassium salts—sylvine and carnallite), seawater (0.065 percent), the natural brine of salt lakes (up to 0.2 percent), and underground brines, usually associated with salt and petroleum deposits (up to 0.1 percent). Because of the high water solubility of bromine salts, they accumulate in the residual brines of sea and lake basins. Bromine migrates as easily soluble compounds that very rarely yield solid mineral forms—for example, bromyrite, AgBr; embolite, Ag(Cl,Br); and iodembolite, Ag(Cl, Br, I). The formation of minerals takes place in oxidation zones of deposits containing silver sulfide that form in arid desert regions.

Physical and chemical properties. Liquid bromine solidifies at −7.2° C to form reddish-brown acicular crystals with a slight metallic luster. Bromine vapor is yellowish-brown; its boiling point is 58.78° C. Liquid bromine has a density (at 20° C) of 3.1 g/cm³. It has limited solubility in water, but it is more soluble than other halogens (3.58 g of bromine in 100 g of water at 20° C). Garnet-red crystals of Br₂ · 8H₂O are precipitated from water at temperatures below 5.84° C. Bromine is very readily soluble in many organic solvents; therefore, they are used to extract it from aqueous solutions. The bromine molecule is diatomic in the solid, liquid, and gaseous states. Appreciable dissociation into atoms starts at about 800° C; dissociation is also observed under the influence of light.

The electron configuration of bromine is 4s²4p⁵. The valence of bromine in compounds is variable, and the degree of oxidation is −1 (in bromides—for example, KBr), +1 (in hypobromites, NaBrO), +3 (in biomites, NaBrO₂), +5 (in bromates, KBrO₃), or +7 (in perbromates, NaBrO₄). Chemically, bromine is very reactive; its reactivity is between that of chlorine and iodine. The reaction of bromine with sulfur, selenium, tellurium, arsenic, and antimony is accompanied by large heat evolution and sometimes even an appearance of flame. Bromine reacts just as vigorously with some metals—for example, potassium and aluminum. However, many metals react with difficulty with dry bromine, owing to the formation of a protective bromide film, which is insoluble in bromine. The metals that are most resistant to the action of bromine, even at high temperatures and in the presence of moisture, are silver, lead, platinum, and tantalum. (Unlike platinum, gold reacts vigorously with bromine.) Even at high temperatures, bromine does not combine directly with oxygen, nitrogen, and carbon. Indirect methods are used to obtain compounds of bromine with these elements, such as the extremely unstable oxides Br₂O, BrO₂, and Br₃O₈ (the latter is obtained, for example, by the action of ozone on bromine at 80° C). Bromine reacts directly with halogens, forming BrF₃, BrF₅, BrCl, IBr, and so on (interhalogen compounds).

Bromine is a strong oxidizing agent. Thus, it oxidizes sulfites and thiosulfates in aqueous solutions to sulfates, nitrites to nitrates, and ammonia to free nitrogen (3Br₂ + 8NH₃ = N₂ + 6NH₄Br). It displaces iodine from its compounds, but is itself displaced by chlorine and fluorine. Free bromine is also liberated from aqueous acid solutions of bromides by strong oxidizing agents (KMnO₄, K₂Cr₂O₇). When dissolved in water, part of the bromine reacts (Br₂ + H₂O ⇆ HBr + HBrO) to give hydrobromic acid and unstable hypobromous acid, HBrO. A solution of bromine in water is called bromine water. When bromine is dissolved in alkali solutions in the cold, bromide and hypobromite are formed (2NaOH + Br₂ = NaBr + NaBrO + H₂O); at high temperatures (about 100° C), bromide and bromate are obtained (6NaOH + 3 Br₂ = 5 NaBr + NaBrO₃ + 3H₂O). The most characteristic reactions of bromine with organic compounds are its addition to double bonds, C==C, and its substitution for hydrogen (usually under the action of catalysts or light).

Production and use. Seawater, lake and underground brines, and the liquors from caustic potash production, all of which contain bromine in the form of the bromide ion Br⁻ (from 65 g/m³ for seawater to 3-4 kg/m³ in liquors from caustic potash production), serve as raw materials for the production of bromine. Bromine is usually isolated by the action of chlorine (2Br⁻ + Cl₂ = Br₂ 2Cl) and is distilled with water vapor or air. Steam distillation is effected in columns of granite, ceramic, or other bromine-resistant material. Hot brine is fed in at the top of the column; chlorine and steam are fed in at the bottom. The bromine vapors coming out of the column are condensed in ceramic condensers. The bromine is separated from the water, and traces of chlorine are removed by distillation. Aeration permits the use of brines of low bromine content, for which the use of steam would be impractical because of high steam consumption. The bromine is removed from the resulting bromine-air mixture using chemical absorbents. For this purpose, use is made of either ferrous bromide solution (2FeBr₂ + Br₂ = 2FeBr₃), which is made by reducing FeBr₃ with iron filings; sodium hydroxide or carbonate solutions; or gaseous sulfur dioxide, which reacts with bromine in the presence of water vapor to give hydrobromic and sulfuric acids (Br₂ + SO₂ + 2H₂O = 2HBr + H₂SO₄). Bromine is obtained from the resultant intermediates either by the action of chlorine (in the cases of FeBr₃ and HBr) or acids (5NaBr + NaBrO₃ + 3H₂SO₄ = 3Br₂ + 3Na₂SO₄ + 3H₂O). According to demand, the intermediates are converted to bromine compounds without the liberation of elemental bromine.

The inhalation of bromine vapors at an atmospheric content of 1 mg/m³ or more causes coughing, catarrh, nosebleed, dizziness, and headache; higher concentrations cause asthma, bronchitis, and sometimes death. The maximum safe concentration of bromine vapor in air is 2 mg/m³. Liquid bromine causes skin burns that do not heal readily. Work with bromine must be done under an exhaust hood. In cases of poisoning by bromine vapor, inhalation of ammonia is recommended, using a very dilute aqueous or alcohol solution. Throat pain resulting from inhalation of bromine vapor can be cured by drinking hot milk. Bromine on the skin should be washed off with plenty of water or blown off with a powerful air jet. Burns should be treated with lanolin.

Bromine is used rather extensively. It is a starting material for making a number of bromine salts and organic bromine derivatives. Large amounts of bromine are used to make ethyl bromide and dibromethane, which are components of the ethyl fluid added to gasoline to improve its antiknock qualities. Bromine compounds are used in photography and in making a number of dyes; methyl bromide and various other bromine compounds are used as insecticides. Certain organic bromine compounds are efficient fire extinguishers. Bromine and bromine water are used in chemical analysis to determine many substances. Sodium, potassium, and ammonium bromides, as well as organic bromine compounds, are used in medicine to treat neuroses, hysteria, hypertension, insomnia, hypertonic maladies, epilepsy, and chorea.