Bromine


Bromine is a chemical element; it has symbol Br and atomic number 35. It is a volatile red-brown liquid at room temperature that evaporates readily to form a similarly coloured vapour. Its properties are intermediate between those of chlorine and iodine. Isolated independently by two chemists, Carl Jacob Löwig and Antoine Jérôme Balard, its name was derived, referring to its sharp and pungent smell.
Elemental bromine is very reactive and thus does not occur as a free element in nature. Instead, it can be isolated from colourless soluble crystalline mineral halide salts analogous to table salt, a property it shares with the other halogens. While it is rather rare in the Earth's crust, the high solubility of the bromide ion has caused its accumulation in the oceans. Commercially the element is easily extracted from brine evaporation ponds, mostly in the United States and Israel. The mass of bromine in the oceans is about one three-hundredth that of chlorine.
At standard conditions for temperature and pressure it is a liquid; the only other element that is liquid under these conditions is mercury. At high temperatures, organobromine compounds readily dissociate to yield free bromine atoms, a process that stops free radical chemical chain reactions. This effect makes organobromine compounds useful as fire retardants, with more than half the bromine produced worldwide each year put to this purpose. The same property causes ultraviolet sunlight to dissociate volatile organobromine compounds in the atmosphere to yield free bromine atoms, causing ozone depletion. As a result, some organobromine compounds—such as the pesticide methyl bromide—are no longer used. Bromine compounds are still used in well drilling fluids, in photographic film, and as an intermediate in the manufacture of organic chemicals.
Large amounts of bromide salts are toxic from the action of soluble bromide ions, causing bromism. However, bromine is beneficial for human eosinophils, and is an essential trace element for collagen development in all animals. Hundreds of known organobromine compounds are generated by terrestrial and marine plants and animals, and some serve important biological roles. As a pharmaceutical, the simple bromide ion has inhibitory effects on the central nervous system, and bromide salts were once a major medical sedative, before replacement by shorter-acting drugs. They retain niche uses as antiepileptics.

History

Bromine was discovered independently by two chemists, Carl Jacob Löwig and Antoine Balard, in 1825 and 1826, respectively.
Justus von Liebig discovered Bromine in 1825, however, he did not recognize that he was looking at an unknown element and mistook it for iodine chloride.
Löwig isolated bromine from a mineral water spring from his hometown Bad Kreuznach in 1825. Löwig used a solution of the mineral salt saturated with chlorine and extracted the bromine with diethyl ether. After evaporation of the ether, a brown liquid remained. With this liquid as a sample of his work he applied for a position in the laboratory of Leopold Gmelin in Heidelberg. The publication of the results was delayed and Balard published his results first.
Balard found bromine chemicals in the ash of seaweed from the salt marshes of Montpellier. The seaweed was used to produce iodine, but also contained bromine. Balard distilled the bromine from a solution of seaweed ash saturated with chlorine. The properties of the resulting substance were intermediate between those of chlorine and iodine; thus he tried to prove that the substance was iodine monochloride, but after failing to do so he was sure that he had found a new element and named it muride, derived from the Latin word muria.
After the French chemists Louis Nicolas Vauquelin, Louis Jacques Thénard, and Joseph-Louis Gay-Lussac approved the experiments of the young pharmacist Balard, the results were presented at a lecture of the Académie des Sciences and published in Annales de Chimie et Physique. In his publication, Balard stated that he changed the name from muride to brôme on the proposal of. The name brôme derives from the Greek βρῶμος. Other sources claim that the French chemist and physicist Joseph-Louis Gay-Lussac suggested the name brôme for the characteristic smell of the vapours. Bromine was not produced in large quantities until 1858, when the discovery of salt deposits in Stassfurt enabled its production as a by-product of potash.
Apart from some minor medical applications, the first commercial use was the daguerreotype. In 1840, bromine was discovered to have some advantages over the previously used iodine vapour to create the light sensitive silver halide layer in daguerreotypy.
By 1864, a 25% solution of liquid bromine in.75 molar aqueous potassium bromide was widely used to treat gangrene during the American Civil War, before the publications of Joseph Lister and Pasteur.
Potassium bromide and sodium bromide were used as anticonvulsants and sedatives in the late 19th and early 20th centuries, but were gradually superseded by chloral hydrate and then by the barbiturates. In the early years of the First World War, bromine compounds such as xylyl bromide were used as poison gas.

Properties

Bromine is the third halogen, being a nonmetal in group 17 of the periodic table. Its properties are thus similar to those of fluorine, chlorine, and iodine, and tend to be intermediate between those of chlorine and iodine, the two neighbouring halogens. Bromine has the electron configuration 4s3d4p, with the seven electrons in the fourth and outermost shell acting as its valence electrons. Like all halogens, it is thus one electron short of a full octet, and is hence a strong oxidising agent, reacting with many elements in order to complete its outer shell. Corresponding to periodic trends, it is intermediate in electronegativity between chlorine and iodine, and is less reactive than chlorine and more reactive than iodine. It is also a weaker oxidising agent than chlorine, but a stronger one than iodine. Conversely, the bromide ion is a weaker reducing agent than iodide, but a stronger one than chloride. These similarities led to chlorine, bromine, and iodine together being classified as one of the original triads of Johann Wolfgang Döbereiner, whose work foreshadowed the periodic law for chemical elements. It is intermediate in atomic radius between chlorine and iodine, and this leads to many of its atomic properties being similarly intermediate in value between chlorine and iodine, such as first ionisation energy, electron affinity, enthalpy of dissociation of the X molecule, ionic radius, and X–X bond length. The volatility of bromine accentuates its very penetrating, choking, and unpleasant odour.
All four stable halogens experience intermolecular van der Waals forces of attraction, and their strength increases together with the number of electrons among all homonuclear diatomic halogen molecules. Thus, the melting and boiling points of bromine are intermediate between those of chlorine and iodine. As a result of the increasing molecular weight of the halogens down the group, the density and heats of fusion and vaporisation of bromine are again intermediate between those of chlorine and iodine, although all their heats of vaporisation are fairly low thanks to their diatomic molecular structure. The halogens darken in colour as the group is descended: fluorine is a very pale yellow gas, chlorine is greenish-yellow, and bromine is a reddish-brown volatile liquid that freezes at −7.2 °C and boils at 58.8 °C. This trend occurs because the wavelengths of visible light absorbed by the halogens increase down the group. Specifically, the colour of a halogen, such as bromine, results from the electron transition between the highest occupied antibonding π molecular orbital and the lowest vacant antibonding σ molecular orbital. The colour fades at low temperatures so that solid bromine at −195 °C is pale yellow.
Liquid bromine is infrared-transparent.
Like solid chlorine and iodine, solid bromine crystallises in the orthorhombic crystal system, in a layered arrangement of Br molecules. The Br–Br distance is 227 pm and the Br···Br distance between molecules is 331 pm within a layer and 399 pm between layers. This structure means that bromine is a very poor conductor of electricity, with a conductivity of around 5 × 10 Ω cm just below the melting point, although this is higher than the essentially undetectable conductivity of chlorine.
At a pressure of 55 GPa bromine undergoes an insulator-to-metal transition. At 75 GPa it changes to a face-centered orthorhombic structure. At 100 GPa it changes to a body centered orthorhombic monatomic form.

Isotopes

Bromine has two stable isotopes, Br and Br. These are its only two natural isotopes, with Br making up 51% of natural bromine and Br making up the remaining 49%. Both have nuclear spin 3/2− and thus may be used for nuclear magnetic resonance, although Br is more favourable. The relatively 1:1 distribution of the two isotopes in nature is helpful in identification of bromine containing compounds using mass spectroscopy. Other bromine isotopes are all radioactive, with half-lives too short to occur in nature. Of these, the most important are Br, Br, and Br, which may be produced from the neutron activation of natural bromine. The most stable bromine radioisotope is Br. The primary decay mode of isotopes lighter than Br is electron capture to isotopes of selenium; that of isotopes heavier than Br is beta decay to isotopes of krypton; and Br may decay by either mode to stable Se or Kr. Br isotopes from 87Br and heavier undergo beta decay with neutron emission and are of practical importance because they are fission products.

Chemistry and compounds

XXXHXBXAlXCX
F159574645582456
Cl243428444427327
Br193363368360272
I151294272285239

Bromine is intermediate in reactivity between chlorine and iodine, and is one of the most reactive elements. Bond energies to bromine tend to be lower than those to chlorine but higher than those to iodine, and bromine is a weaker oxidising agent than chlorine but a stronger one than iodine. This can be seen from the standard electrode potentials of the X/X couples. Bromination often leads to higher oxidation states than iodination but lower or equal oxidation states to chlorination. Bromine tends to react with compounds including M–M, M–H, or M–C bonds to form M–Br bonds.