Bromine is achemical element; it hassymbolBr andatomic number 35. It is a volatile red-brownliquid at room temperature that evaporates readily to form a similarly coloured vapour. Its properties are intermediate between those ofchlorine andiodine. Isolated independently by two chemists,Carl Jacob Löwig (in 1825) andAntoine Jérôme Balard (in 1826), its name was derived from Ancient Greek βρῶμος (bromos)'stench', referring to its sharp and pungent smell.
Elemental bromine is very reactive and thus does not occur as afree element in nature. Instead, it can be isolated from colourless soluble crystalline mineral halidesalts analogous totable salt, a property it shares with the otherhalogens. While it is rather rare in the Earth's crust, the high solubility of thebromide ion (Br−) has caused itsaccumulation in the oceans. Commercially the element is easily extracted from brineevaporation ponds, mostly in theUnited States andIsrael. The mass of bromine in the oceans is about one three-hundredth that of chlorine.
Large amounts of bromide salts are toxic from the action of soluble bromide ions, causingbromism. However, bromine is beneficial for humaneosinophils,[10] and is an essential trace element forcollagen development in all animals.[11] Hundreds of known organobromine compounds are generated by terrestrial and marine plants and animals, and some serve important biological roles.[12] As apharmaceutical, the simple bromide ion (Br−) has inhibitory effects on the central nervous system, and bromidesalts were once a major medical sedative, before replacement by shorter-acting drugs. They retain niche uses asantiepileptics.
Löwig isolated bromine from a mineral water spring from his hometownBad Kreuznach in 1825. Löwig used a solution of the mineral salt saturated with chlorine and extracted the bromine withdiethyl 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 ofLeopold Gmelin inHeidelberg. The publication of the results was delayed and Balard published his results first.[17]
Balard found bromine chemicals in the ash ofseaweed from thesalt marshes ofMontpellier. 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 wasiodine monochloride (ICl), but after failing to do so he was sure that he had found a new element and named it muride, derived from theLatin wordmuria ("brine").[15][18][19]
After the French chemistsLouis Nicolas Vauquelin,Louis Jacques Thénard, andJoseph-Louis Gay-Lussac approved the experiments of the young pharmacist Balard, the results were presented at a lecture of theAcadémie des Sciences and published inAnnales de Chimie et Physique.[14] In his publication, Balard stated that he changed the name frommuride tobrôme on the proposal ofM. Anglada [fr]. The namebrôme (bromine) derives from theGreekβρῶμος (brômos, "stench").[14][20][18][21] Other sources claim that the French chemist and physicistJoseph-Louis Gay-Lussac suggested the namebrôme for the characteristic smell of the vapours.[22][23] Bromine was not produced in large quantities until 1858, when the discovery of salt deposits inStassfurt enabled its production as a by-product ofpotash.[24]
Apart from some minor medical applications, the first commercial use was thedaguerreotype. In 1840, bromine was discovered to have some advantages over the previously used iodine vapour to create the light sensitivesilver halide layer in daguerreotypy.[25]
By 1864, a 25% solution of liquid bromine in .75 molar aqueous potassium bromide[26] was widely used[27] to treatgangrene during the American Civil War, before the publications ofJoseph Lister andPasteur.[28]
Bromine is the thirdhalogen, being anonmetal in group 17 of the periodic table. Its properties are thus similar to those offluorine,chlorine, andiodine, and tend to be intermediate between those of chlorine and iodine, the two neighbouring halogens. Bromine has the electron configuration [Ar]4s23d104p5, with the seven electrons in the fourth and outermost shell acting as itsvalence 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.[31] Corresponding toperiodic trends, it is intermediate inelectronegativity between chlorine and iodine (F: 3.98, Cl: 3.16, Br: 2.96, I: 2.66), 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, thebromide ion is a weaker reducing agent than iodide, but a stronger one than chloride.[31] These similarities led to chlorine, bromine, and iodine together being classified as one of the original triads ofJohann Wolfgang Döbereiner, whose work foreshadowed theperiodic law for chemical elements.[32][33] It is intermediate inatomic 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 firstionisation energy,electron affinity, enthalpy of dissociation of the X2 molecule (X = Cl, Br, I), ionic radius, and X–X bond length.[31] The volatility of bromine accentuates its very penetrating, choking, and unpleasant odour.[34]
All four stable halogens experience intermolecularvan 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 (leading to high volatility) thanks to their diatomic molecular structure.[31] 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. (Iodine is a shiny black solid.) This trend occurs because the wavelengths of visible light absorbed by the halogens increase down the group.[31] Specifically, the colour of a halogen, such as bromine, results from theelectron transition between thehighest occupied antibondingπg molecular orbital and the lowest vacant antibondingσu molecular orbital.[35] The colour fades at low temperatures so that solid bromine at −195 °C is pale yellow.[31]
Like solid chlorine and iodine, solid bromine crystallises in theorthorhombic crystal system, in a layered arrangement of Br2 molecules. The Br–Br distance is 227 pm (close to the gaseous Br–Br distance of 228 pm) and the Br···Br distance between molecules is 331 pm within a layer and 399 pm between layers (compare the van der Waals radius of bromine, 195 pm). This structure means that bromine is a very poor conductor of electricity, with a conductivity of around 5 × 10−13 Ω−1 cm−1 just below the melting point, although this is higher than the essentially undetectable conductivity of chlorine.[31]
At a pressure of 55 GPa (roughly 540,000 times atmospheric pressure) 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.[37]
Bromine has two stableisotopes,79Br and81Br. These are its only two natural isotopes, with79Br making up 51% of natural bromine and81Br making up the remaining 49%. Both have nuclear spin 3/2− and thus may be used fornuclear magnetic resonance, although81Br 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, withhalf-lives too short to occur in nature. Of these, the most important are80Br (t1/2 = 17.7 min),80mBr (t1/2 = 4.421 h), and82Br (t1/2 = 35.28 h), which may be produced from theneutron activation of natural bromine.[31] The most stable bromine radioisotope is77Br (t1/2 = 57.04 h). The primary decay mode of isotopes lighter than79Br iselectron capture to isotopes ofselenium; that of isotopes heavier than81Br isbeta decay to isotopes ofkrypton; and80Br may decay by either mode to stable80Se or80Kr. Br isotopes from87Br and heavier undergo beta decay with neutron emission and are of practical importance because they are fission products.[38]
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 thestandard electrode potentials of the X2/X− couples (F, +2.866 V; Cl, +1.395 V; Br, +1.087 V; I, +0.615 V; At, approximately +0.3 V). 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.[35]
Hydrogen bromide
The simplest compound of bromine ishydrogen bromide, HBr. It is mainly used in the production of inorganicbromides andalkyl bromides, and as a catalyst for many reactions in organic chemistry. Industrially, it is mainly produced by the reaction ofhydrogen gas with bromine gas at 200–400 °C with aplatinum catalyst. However, reduction of bromine withred phosphorus is a more practical way to produce hydrogen bromide in the laboratory:[39]
2 P + 6 H2O + 3 Br2 → 6 HBr + 2 H3PO3
H3PO3 + H2O + Br2 → 2 HBr + H3PO4
At room temperature, hydrogen bromide is a colourless gas, like all the hydrogen halides apart fromhydrogen fluoride, since hydrogen cannot form stronghydrogen bonds to the large and only mildly electronegative bromine atom; however, weak hydrogen bonding is present in solid crystalline hydrogen bromide at low temperatures, similar to the hydrogen fluoride structure, before disorder begins to prevail as the temperature is raised.[39] Aqueous hydrogen bromide is known ashydrobromic acid, which is a strong acid (pKa = −9) because the hydrogen bonds to bromine are too weak to inhibit dissociation. The HBr/H2O system also involves many hydrates HBr·nH2O forn = 1, 2, 3, 4, and 6, which are essentially salts of bromineanions andhydroniumcations. Hydrobromic acid forms anazeotrope with boiling point 124.3 °C at 47.63 g HBr per 100 g solution; thus hydrobromic acid cannot be concentrated beyond this point by distillation.[40]
Unlikehydrogen fluoride, anhydrous liquid hydrogen bromide is difficult to work with as a solvent, because its boiling point is low, it has a small liquid range, itsdielectric constant is low and it does not dissociate appreciably into H2Br+ andHBr− 2 ions – the latter, in any case, are much less stable than thebifluoride ions (HF− 2) due to the very weak hydrogen bonding between hydrogen and bromine, though its salts with very large and weakly polarising cations such asCs+ andNR+ 4 (R =Me,Et,Bun) may still be isolated. Anhydrous hydrogen bromide is a poor solvent, only able to dissolve small molecular compounds such asnitrosyl chloride andphenol, or salts with very lowlattice energies such as tetraalkylammonium halides.[40]
Nearly all elements in the periodic table form binary bromides. The exceptions are decidedly in the minority and stem in each case from one of three causes: extreme inertness and reluctance to participate in chemical reactions (thenoble gases, with the exception ofxenon in the very unstableXeBr2); extreme nuclear instability hampering chemical investigation before decay and transmutation (many of the heaviest elements beyondbismuth); and having an electronegativity higher than bromine's (oxygen,nitrogen,fluorine, andchlorine), so that the resultant binary compounds are formally not bromides but rather oxides, nitrides, fluorides, or chlorides of bromine. (Nonetheless,nitrogen tribromide is named as a bromide as it is analogous to the other nitrogen trihalides.)[41]
Bromination of metals with Br2 tends to yield lower oxidation states than chlorination with Cl2 when a variety of oxidation states is available. Bromides can be made by reaction of an element or its oxide, hydroxide, or carbonate with hydrobromic acid, and then dehydrated by mildly high temperatures combined with either low pressure or anhydrous hydrogen bromide gas. These methods work best when the bromide product is stable to hydrolysis; otherwise, the possibilities include high-temperature oxidative bromination of the element with bromine or hydrogen bromide, high-temperature bromination of a metal oxide or other halide by bromine, a volatile metal bromide,carbon tetrabromide, or an organic bromide. For example,niobium(V) oxide reacts with carbon tetrabromide at 370 °C to formniobium(V) bromide.[41] Another method is halogen exchange in the presence of excess "halogenating reagent", for example:[41]
FeCl3 + BBr3 (excess) → FeBr3 + BCl3
When a lower bromide is wanted, either a higher halide may be reduced using hydrogen or a metal as a reducing agent, or thermal decomposition ordisproportionation may be used, as follows:[41]
3 WBr5 + Althermal gradient→475 °C → 240 °C 3 WBr4 + AlBr3
EuBr3 +1/2 H2 → EuBr2 + HBr
2 TaBr4500 °C→ TaBr3 + TaBr5
Most metal bromides with the metal in low oxidation states (+1 to +3) are ionic. Nonmetals tend to form covalent molecular bromides, as do metals in high oxidation states from +3 and above. Both ionic and covalent bromides are known for metals in oxidation state +3 (e.g.scandium bromide is mostly ionic, butaluminium bromide is not).Silver bromide is very insoluble in water and is thus often used as a qualitative test for bromine.[41]
Bromine halides
The halogens form many binary,diamagneticinterhalogen compounds with stoichiometries XY, XY3, XY5, and XY7 (where X is heavier than Y), and bromine is no exception. Bromine forms a monofluoride and monochloride, as well as a trifluoride and pentafluoride. Some cationic and anionic derivatives are also characterised, such asBrF− 2,BrCl− 2,BrF+ 2,BrF+ 4, andBrF+ 6. Apart from these, somepseudohalides are also known, such ascyanogen bromide (BrCN), brominethiocyanate (BrSCN), and bromineazide (BrN3).[42]
The pale-brownbromine monofluoride (BrF) is unstable at room temperature, disproportionating quickly and irreversibly into bromine, bromine trifluoride, and bromine pentafluoride. It thus cannot be obtained pure. It may be synthesised by the direct reaction of the elements, or by the comproportionation of bromine and bromine trifluoride at high temperatures.[42]Bromine monochloride (BrCl), a red-brown gas, quite readily dissociates reversibly into bromine and chlorine at room temperature and thus also cannot be obtained pure, though it can be made by the reversible direct reaction of its elements in the gas phase or incarbon tetrachloride.[41] Bromine monofluoride inethanol readily leads to the monobromination of thearomatic compounds PhX (para-bromination occurs for X = Me, But, OMe, Br;meta-bromination occurs for the deactivating X = –CO2Et, –CHO, –NO2); this is due to heterolytic fission of the Br–F bond, leading to rapid electrophilic bromination by Br+.[41]
At room temperature,bromine trifluoride (BrF3) is a straw-coloured liquid. It may be formed by directly fluorinating bromine at room temperature and is purified through distillation. It reacts violently with water and explodes on contact with flammable materials, but is a less powerful fluorinating reagent thanchlorine trifluoride. It reacts vigorously withboron,carbon,silicon,arsenic,antimony, iodine, andsulfur to give fluorides, and will also convert most metals and many metal compounds to fluorides; as such, it is used to oxidiseuranium touranium hexafluoride in the nuclear power industry. Refractory oxides tend to be only partially fluorinated, but here the derivatives KBrF4 and BrF2SbF6 remain reactive. Bromine trifluoride is a useful nonaqueous ionising solvent, since it readily dissociates to formBrF+ 2 andBrF− 4 and thus conducts electricity.[43]
Bromine pentafluoride (BrF5) was first synthesised in 1930. It is produced on a large scale by direct reaction of bromine with excess fluorine at temperatures higher than 150 °C, and on a small scale by the fluorination ofpotassium bromide at 25 °C. It also reacts violently with water and is a very strong fluorinating agent, although chlorine trifluoride is still stronger.[44]
Polybromine compounds
Although dibromine is a strong oxidising agent with a high first ionisation energy, very strong oxidisers such asperoxydisulfuryl fluoride (S2O6F2) can oxidise it to form the cherry-redBr+ 2 cation. A few other bromine cations are known, namely the brownBr+ 3 and dark brownBr+ 5.[45] The tribromide anion,Br− 3, has also been characterised; it is analogous totriiodide.[42]
Bromine oxides and oxoacids
Standard reduction potentials for aqueous Br species[46]
So-called "bromine dioxide", a pale yellow crystalline solid, may be better formulated as bromineperbromate, BrOBrO3. It is thermally unstable above −40 °C, violently decomposing to its elements at 0 °C.Dibromine trioxide,syn-BrOBrO2, is also known; it is the anhydride ofhypobromous acid andbromic acid. It is an orange crystalline solid which decomposes above −40 °C; if heated too rapidly, it explodes around 0 °C. A few other unstable radical oxides are also known, as are some poorly characterised oxides, such asdibromine pentoxide,tribromine octoxide, and bromine trioxide.[48]
The fouroxoacids,hypobromous acid (HOBr),bromous acid (HOBrO),bromic acid (HOBrO2), andperbromic acid (HOBrO3), are better studied due to their greater stability, though they are only so in aqueous solution. When bromine dissolves in aqueous solution, the following reactions occur:[46]
Hypobromous acid is unstable to disproportionation. Thehypobromite ions thus formed disproportionate readily to give bromide and bromate:[46]
3 BrO− ⇌ 2 Br− +BrO− 3
K = 1015
Bromous acids andbromites are very unstable, although thestrontium andbarium bromites are known.[49] More important are thebromates, which are prepared on a small scale by oxidation of bromide by aqueoushypochlorite, and are strong oxidising agents. Unlike chlorates, which very slowly disproportionate to chloride and perchlorate, the bromate anion is stable to disproportionation in both acidic and aqueous solutions. Bromic acid is a strong acid. Bromides and bromates may comproportionate to bromine as follows:[49]
BrO− 3 + 5 Br− + 6 H+ → 3 Br2 + 3 H2O
There were many failed attempts to obtain perbromates and perbromic acid, leading to some rationalisations as to why they should not exist, until 1968 when the anion was first synthesised from the radioactivebeta decay of unstable83 SeO2− 4. Today, perbromates are produced by the oxidation of alkaline bromate solutions by fluorine gas. Excess bromate and fluoride are precipitated assilver bromate andcalcium fluoride, and the perbromic acid solution may be purified. The perbromate ion is fairly inert at room temperature but is thermodynamically extremely oxidising, with extremely strong oxidising agents needed to produce it, such as fluorine orxenon difluoride. The Br–O bond inBrO− 4 is fairly weak, which corresponds to the general reluctance of the 4p elementsarsenic,selenium, and bromine to attain their group oxidation state, as they come after thescandide contraction characterised by the poor shielding afforded by the radial-nodeless 3d orbitals.[50]
Structure ofN-bromosuccinimide, a common brominating reagent in organic chemistry
Like the other carbon–halogen bonds, the C–Br bond is a common functional group that forms part of coreorganic chemistry. Formally, compounds with this functional group may be considered organic derivatives of the bromide anion. Due to the difference of electronegativity between bromine (2.96) and carbon (2.55), the carbon atom in a C–Br bond is electron-deficient and thuselectrophilic. The reactivity of organobromine compounds resembles but is intermediate between the reactivity oforganochlorine andorganoiodine compounds. For many applications, organobromides represent a compromise of reactivity and cost.[51]
Organobromides are the most common organohalides in nature, even though the concentration of bromide is only 0.3% of that for chloride in sea water, because of the easy oxidation of bromide to the equivalent of Br+, a potent electrophile. The enzymebromoperoxidase catalyses this reaction.[52] The oceans are estimated to release 1–2 million tons ofbromoform and 56,000 tons ofbromomethane annually.[12]
Bromine addition to alkene reaction mechanism
An old qualitative test for the presence of thealkene functional group is that alkenes turn brown aqueous bromine solutions colourless, forming abromohydrin with some of the dibromoalkane also produced. The reaction passes through a short-lived strongly electrophilicbromonium intermediate. This is an example of ahalogen addition reaction.[53]
Occurrence and production
View of salt evaporation pans on the Dead Sea, whereJordan (right) and Israel (left) produce salt and bromine
Bromine is significantly less abundant in the crust than fluorine or chlorine, comprising only 2.5 parts per million of the Earth's crustal rocks, and then only as bromide salts. It is significantly more abundant in the oceans, resulting from long-termleaching. There, it makes up 65 parts per million, corresponding to a ratio of about one bromine atom for every 660 chlorine atoms. Salt lakes and brine wells may have higher bromine concentrations: for example, theDead Sea contains 0.4% bromide ions.[54] It is from these sources that bromine extraction is mostly economically feasible.[55][56][57] Bromine is the tenth most abundant element in seawater.[58]
The main sources of bromine production areIsrael andJordan.[59] The element is liberated by halogen exchange, using chlorine gas to oxidise Br− to Br2. This is then removed with a blast of steam or air, and is then condensed and purified.[60] Today, bromine is transported in large-capacity metal drums or lead-lined tanks that can hold hundreds of kilograms or even tonnes of bromine. The bromine industry is about one-hundredth the size of the chlorine industry. Laboratory production is unnecessary because bromine is commercially available and has a long shelf life.[61]
Applications
A wide variety of organobromine compounds are used inindustry. Some are prepared from bromine and others are prepared fromhydrogen bromide, which is obtained by burninghydrogen in bromine.[62]
Flame retardants
Tetrabromobisphenol A
Brominated flame retardants represent a commodity of growing importance, and make up the largest commercial use of bromine. When the brominated material burns, the flame retardant produceshydrobromic acid which interferes in the radicalchain reaction of theoxidation reaction of the fire. The mechanism is that the highly reactive hydrogen radicals, oxygen radicals, andhydroxyl radicals react with hydrobromic acid to form less reactive bromine radicals (i.e., free bromine atoms). Bromine atoms may also react directly with other radicals to help terminate the free radical chain-reactions that characterise combustion.[63][64]
To make brominated polymers and plastics, bromine-containing compounds can be incorporated into the polymer duringpolymerisation. One method is to include a relatively small amount of brominated monomer during the polymerisation process. For example,vinyl bromide can be used in the production ofpolyethylene,polyvinyl chloride orpolypropylene. Specific highly brominated molecules can also be added that participate in the polymerisation process. For example,tetrabromobisphenol A can be added topolyesters or epoxy resins, where it becomes part of the polymer. Epoxies used inprinted circuit boards are normally made from such flame retardantresins, indicated by the FR in the abbreviation of the products (FR-4 andFR-2). In some cases, the bromine-containing compound may be added after polymerisation. For example,decabromodiphenyl ether can be added to the final polymers.[65]
A number of gaseous or highly volatile brominatedhalomethane compounds are non-toxic and make superior fire suppressant agents by this same mechanism, and are particularly effective in enclosed spaces such as submarines, airplanes, and spacecraft. However, they are expensive and their production and use has been greatly curtailed due to their effect as ozone-depleting agents. They are no longer used in routine fire extinguishers, but retain niche uses in aerospace and military automatic fire suppression applications. They includebromochloromethane (Halon 1011, CH2BrCl),bromochlorodifluoromethane (Halon 1211, CBrClF2), andbromotrifluoromethane (Halon 1301, CBrF3).[66]
Ethylene bromide was anadditive in gasolines containing lead anti-engine knocking agents. It scavenges lead by forming volatile lead bromide, which is exhausted from the engine. This application accounted for 77% of the bromine use in 1966 in the US. This application has declined since the 1970s due to environmental regulations (see below).[67]
Brominated vegetable oil (BVO), a complex mixture of plant-derived triglycerides that have been reacted to contain atoms of the element bromine bonded to the molecules, is used primarily to help emulsify citrus-flavored soft drinks, preventing them from separating during distribution.
Poisonousbromomethane was widely used aspesticide tofumigate soil and to fumigate housing, by the tenting method. Ethylene bromide was similarly used.[68] These volatile organobromine compounds are all now regulated asozone depletion agents. TheMontreal Protocol on Substances that Deplete the Ozone Layer scheduled the phase out for theozone depleting chemical by 2005, and organobromide pesticides are no longer used (in housing fumigation they have been replaced by such compounds assulfuryl fluoride, which contain neither the chlorine or bromine organics which harm ozone). Before the Montreal protocol in 1991 (for example) an estimated 35,000 tonnes of the chemical were used to controlnematodes,fungi,weeds and other soil-borne diseases.[69][70]
In pharmacology, inorganicbromide compounds, especiallypotassium bromide, were frequently used as general sedatives in the 19th and early 20th century. Bromides in the form of simple salts are still used as anticonvulsants in both veterinary and human medicine, although the latter use varies from country to country. For example, the U.S.Food and Drug Administration (FDA) does not approve bromide for the treatment of any disease, andsodium bromide was removed from over-the-counter sedative products likeBromo-Seltzer, in 1975.[71] Commercially available organobromine pharmaceuticals include the vasodilatornicergoline, the sedativebrotizolam, the anticancer agentpipobroman, and the antisepticmerbromin. Otherwise, organobromine compounds are rarely pharmaceutically useful, in contrast to the situation fororganofluorine compounds. Several drugs are produced as the bromide (or equivalents, hydrobromide) salts, but in such cases bromide serves as an innocuous counterion of no biological significance.[51]
Other uses of organobromine compounds include high-density drilling fluids, dyes (such asTyrian purple and the indicatorbromothymol blue), and pharmaceuticals. Bromine itself, as well as some of its compounds, are used in water treatment, and is the precursor of a variety of inorganic compounds with an enormous number of applications (e.g.silver bromide for photography).[61]Zinc–bromine batteries are hybridflow batteries used for stationary electrical power backup and storage; from household scale to industrial scale.
Bromine is used in cooling towers (in place of chlorine) for controlling bacteria, algae, fungi, andzebra mussels.[72]
Because it has similar antiseptic qualities to chlorine, bromine can be used in the same manner as chlorine as a disinfectant or antimicrobial in applications such as swimming pools. Bromine came into this use in the United States duringWorld War II due to a predicted shortage of chlorine.[73] However, bromine is usually not used outside for these applications due to it being relatively more expensive than chlorine and the absence of a stabiliser to protect it from the sun. For indoor pools, it can be a good option as it is effective at a wider pH range. It is also more stable in a heated pool or hot tub.[74]
A 2014 study suggests that bromine (in the form of bromide ion) is a necessary cofactor in the biosynthesis ofcollagen IV, making the elementessential tobasement membrane architecture and tissue development in animals.[11] Nevertheless, no clear deprivation symptoms or syndromes have been documented in mammals.[75] In other biological functions, bromine may be non-essential but still beneficial when it takes the place of chlorine. For example, in the presence of hydrogen peroxide,H2O2, formed by theeosinophil, and either chloride, iodide, thiocyanate, or bromide ions,eosinophil peroxidase provides a potent mechanism by which eosinophils kill multicellularparasites (such as the nematode worms involved infilariasis) and somebacteria (such astuberculosis bacteria). Eosinophil peroxidase is ahaloperoxidase that preferentially uses bromide over chloride for this purpose, generatinghypobromite (hypobromous acid), although the use of chloride is possible.[10]
Octan-2-yl 4-bromo-3-oxobutanoate, an organobromine compound found in mammalian cerebrospinal fluid
α-Haloesters are generally thought of as highly reactive and consequently toxic intermediates in organic synthesis. Nevertheless, mammals, including humans, cats, and rats, appear to biosynthesise traces of an α-bromoester, 2-octyl 4-bromo-3-oxobutanoate, which is found in theircerebrospinal fluid and appears to play a yet unclarified role in inducing REM sleep.[12] Neutrophil myeloperoxidase can use H2O2 and Br− to brominate deoxycytidine, which could result in DNA mutations.[76] Marine organisms are the main source of organobromine compounds, and it is in these organisms that bromine is more firmly shown to be essential. More than 1600 such organobromine compounds were identified by 1999. The most abundant ismethyl bromide (CH3Br), of which an estimated 56,000 tonnes is produced by marine algae each year.[12] The essential oil of the Hawaiian algaAsparagopsis taxiformis consists of 80%bromoform.[77] Most of such organobromine compounds in the sea are made by the action of a unique algal enzyme,vanadium bromoperoxidase.[78]
The bromide anion is not very toxic: a normal daily intake is 2 to 8 milligrams.[75] However, high levels of bromide chronically impair the membrane of neurons, which progressively impairs neuronal transmission, leading to toxicity, known asbromism. Bromide has anelimination half-life of 9 to 12 days, which can lead to excessive accumulation. Doses of 0.5 to 1 gram per day of bromide can lead to bromism. Historically, the therapeutic dose of bromide is about 3 to 5 grams of bromide, thus explaining why chronic toxicity (bromism) was once so common. While significant and sometimes serious disturbances occur to neurologic, psychiatric, dermatological, and gastrointestinal functions, death from bromism is rare.[79] Bromism is caused by a neurotoxic effect on the brain which results insomnolence,psychosis,seizures anddelirium.[80]
Elemental bromine (Br2) is toxic and causeschemical burns on human flesh. Inhaling bromine gas results in similar irritation of the respiratory tract, causing coughing, choking, shortness of breath, and death if inhaled in large enough amounts. Chronic exposure may lead to frequent bronchial infections and a general deterioration of health. As a strong oxidising agent, bromine is incompatible with most organic and inorganic compounds.[83] Caution is required when transporting bromine; it is commonly carried in steel tanks lined with lead, supported by strong metal frames.[61] TheOccupational Safety and Health Administration (OSHA) of theUnited States has set apermissible exposure limit (PEL) for bromine at a time-weighted average (TWA) of 0.1 ppm. TheNational Institute for Occupational Safety and Health (NIOSH) has set arecommended exposure limit (REL) of TWA 0.1 ppm and a short-term limit of 0.3 ppm. The exposure to bromineimmediately dangerous to life and health (IDLH) is 3 ppm.[84] Bromine is classified as anextremely hazardous substance in the United States as defined in Section 302 of the U.S.Emergency Planning and Community Right-to-Know Act (42 U.S.C. 11002), and is subject to strict reporting requirements by facilities which produce, store, or use it in significant quantities.[85]
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