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Bromine
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Bromine, 35Br
General properties
Pronunciation
/ˈbroʊmiːn/ or /ˈbroʊmɪn/
BROH-meen or BROH-min
Appearance
reddish-brown
Standard atomic weight (Ar)
[7001799010000000000♠79.901, 7001799070000000000♠79.907] conventional: 7001799040000000000♠79.904[1]
Bromine in the periodic table
Cl
↑
Br
↓
I
selenium ← bromine → krypton
Atomic number (Z)
35
Group, period
group 17 (halogens), period 4
Block
p-block
Element category
diatomic nonmetal
Electron configuration
[Ar] 3d10 4s2 4p5
Electrons per shell
2, 8, 18, 7
Physical properties
Spectral lines

Phase (at STP)
liquid
Melting point
265.8 K (−7.2 °C, 19 °F)
Boiling point
332.0 K (58.8 °C, 137.8 °F)
Density (near r.t.)
Br2, liquid: 3.1028 g/cm3
Triple point
265.90 K, 5.8 kPa[2]
Critical point
588 K, 10.34 MPa[2]
Heat of fusion
(Br2) 10.571 kJ/mol
Heat of vaporisation
(Br2) 29.96 kJ/mol
Molar heat capacity
(Br2) 75.69 J/(mol·K)
Vapour pressure
P (Pa)
1
10
100
1 k
10 k
100 k
at T (K)
185 201 220 244 276 332
Atomic properties
Oxidation states
7, 5, 4, 3, 1, −1 (a strongly acidic oxide)
Electronegativity
Pauling scale: 2.96
Ionisation energies
1st: 1139.9 kJ/mol
2nd: 2103 kJ/mol
3rd: 3470 kJ/mol
Atomic radius
empirical: 120 pm
Covalent radius
120±3 pm
Van der Waals radius
185 pm
Miscellanea
Crystal structure
orthorhombic
Speed of sound
206 m/s (at 20 °C)
Thermal conductivity
0.122 W/(m·K)
Electrical resistivity
7.8×1010 Ω·m (at 20 °C)
Magnetic ordering
diamagnetic[3]
Magnetic susceptibility
−56.4·10−6 cm3/mol[4]
CAS Number
7726-95-6
History
Discovery and first isolation
Antoine Jérôme Balard and Carl Jacob Löwig (1825)
Main isotopes of bromine
Isotope
Abundance
Half-life (t1/2)
Decay mode
Product
79Br
51% stable
81Br
49% stable
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Bromine is a chemical element with symbol Br and atomic number 35. It is the third-lightest halogen, and is a fuming red-brown liquid at room temperature that evaporates readily to form a similarly coloured gas. Its properties are thus intermediate between those of chlorine and iodine. Isolated independently by two chemists, Carl Jacob Löwig (in 1825) and Antoine Jérôme Balard (in 1826), its name was derived from the Ancient Greek βρῶμος "stench", referencing its sharp and disagreeable smell.
Elemental bromine is very reactive and thus does not occur free in nature, but in colourless soluble crystalline mineral halide salts, analogous to table salt. While it is rather rare in the Earth's crust, the high solubility of the bromide ion (Br−) has caused its accumulation in the oceans. Commercially the element is easily extracted from brine pools, mostly in the United States, Israel and China. The mass of bromine in the oceans is about one three-hundredth of that of chlorine.
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, and more than half the bromine produced worldwide each year is put to this purpose. Unfortunately, the same property causes sunlight to dissociate volatile organobromine compounds in the atmosphere to yield free bromine atoms, causing ozone depletion. As a result, many organobromide 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.
Bromine has sometimes been considered to be possibly essential in humans, but this is supported by only limited circumstantial evidence, and no clear biological role. As a pharmaceutical, the simple bromide ion (Br−) 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.
Contents
[hide]
1 History
2 Properties
2.1 Isotopes
3 Chemistry and compounds
3.1 Hydrogen bromide
3.2 Other binary bromides
3.3 Bromine halides
3.4 Polybromine compounds
3.5 Bromine oxides and oxoacids
3.6 Organobromine compounds
4 Occurrence and production
5 Applications
5.1 Flame retardants
5.2 Other uses
6 Biological role and toxicity
7 References
8 Bibliography
History[edit]
Antoine Balard, one of the discoverers of bromine
Bromine was discovered independently by two chemists, Carl Jacob Löwig[5] and Antoine Balard,[6][7] in 1825 and 1826, respectively.[8]
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.[9]
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 (ICl), 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 for brine.[7]
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.[6] In his publication, Balard states that he changed the name from muride to brôme on the proposal of M. Anglada. Brôme (bromine) derives from the Greek βρωμος (stench).[6][10] Other sources claim that the French chemist and physicist Joseph-Louis Gay-Lussac suggested the name brôme for the characteristic smell of the vapors.[11][12] 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.[13]
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 vapor to create the light sensitive silver halide layer in daguerreotypy.[14]
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.[15] In the early years of the First World War, bromine compounds such as xylyl bromide were used as poison gas.[16]
Properties[edit]
Illustrative and secure bromine sample for teaching
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 the two neighbouring halogens, chlorine and iodine. Bromine has the electron configuration [Ar]3d104s24p5, 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.[17] Corresponding to periodic trends, it is intermediate in electronegativity 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, the bromide ion is a weaker reducing agent than iodide, but a stronger one than chloride.[17] 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.[18][19] 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 X2 molecule (X = Cl, Br, I), ionic radius, and X–X bond length.[17] The volatility of bromine accentuates its very penetrating, choking, and unpleasant odour.[20]
All four stable halogens experience intermolecular van der Waals forces of attraction, and their strength increases together with 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.[17] 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 melts 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.[17] Specifically, the colour of a halogen, such as bromine, results from the electron transition between the highest occupied antibonding πg molecular orbital and the lowest vacant antibonding σu molecular orbital.[21] The colour fades at low temperatures, so that solid bromine at −195 °C is pale yellow.[17]
Like solid chlorine and iodine, solid bromine crystallises in the orthorhombic crystal system, in a layered lattice 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 better than the essentially undetectable conductivity of chlorine.[17]
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.[22]
Isotopes[edit]
Main article: Isotopes of bromine
Bromine has two stable isotopes, 79Br and 81Br. These are its only two natural isotopes, with 79Br making up 51% of natural bromine and 81Br making up the remaining 49%. Both have nuclear spin 3/2− and thus may be used for nuclear magnetic resonance, although 81Br is more favourable. The other bromine isotopes are all radioactive, with half-lives too short to occur in nature. Of these, the most important are 80Br (t1/2 = 17.7 min), 80mBr (t1/2 = 4.421 h), and 82Br (t1/2 = 35.28 h), which may be produced from the neutron activation of natural bromine.[17] The most stable bromine radioisotope is 77Br (t1/2 = 57.04 h). The primary decay mode of isotopes lighter than 79Br is electron capture to isotopes of selenium; that of isotopes heavier than 81Br is beta decay to isotopes of krypton; and 80Br may decay by either mode to stable 80Se or 80Kr.[23]
Chemistry and compounds[edit]
Halogen bond energies (kJ/mol)[21]
X
XX
HX
BX3
AlX3
CX4
F
159 574 645 582 456
Cl
243 428 444 427 327
Br
193 363 368 360 272
I
151 294 272 285 239
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 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.[21]
Hydrogen bromide[edit]
The simplest compound of bromine is hydrogen bromide, HBr. It is mainly used in the production of inorganic bromides and alkyl bromides, and as a catalyst for many reactions in organic chemistry. Industrially, it is mainly produced by the reaction of hydrogen gas with bromine gas at 200–400 °C with a platinum catalyst. However, reduction of bromine with red phosphorus is a more practical way to produce hydrogen bromide in the laboratory:[24]
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 from hydrogen fluoride, since hydrogen cannot form strong hydrogen 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.[24] Aqueous hydrogen bromide is known as hydrobromic 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 for n = 1, 2, 3, 4, and 6, which are essentially salts of bromine anions and hydronium cations. Hydrobromic acid forms an azeotrope 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.[25]
Unlike hydrogen 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, its dielectric constant is low and it does not dissociate appreciably into H2Br+ and HBr−
2 ions – the latter, in any case, are much less stable than the bifluoride 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 as Cs+ and NR+
4 (R = Me, Et, Bun) may still be isolated. Anhydrous hydrogen bromide is a poor solvent, only able to dissolve small molecular compounds such as nitrosyl chloride and phenol, or salts with very low lattice energies such as tetraalkylammonium halides.[25]
Other binary bromides[edit]
Silver bromide (AgBr)
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 (the noble gases, with the exception of xenon in the very unstable XeBr2); extreme nuclear instability hampering chemical investigation before decay and transmutation (many of the heaviest elements beyond bismuth); and having an electronegativity higher than bromine's (oxygen, fluorine, and chlorine), so that the resultant binary compounds are formally not bromides but rather oxides, fluorides, or chlorides of bromine.[26]
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 form niobium(V) bromide.[26] Another method is halogen exchange in the presence of excess "halogenating reagent", for example:[26]
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 or disproportionation may be used, as follows:[26]
3 WBr5 + Al thermal gradient→475°C → 240°C 3 WBr4 + AlBr3
EuBr3 + 1/2 H2 → EuBr2 + HBr
2 TaBr4 500°C→ TaBr3 + TaBr5
Most of the bromides of the pre-transition metals (groups 1, 2, and 3, along with the lanthanides and actinides in the +2 and +3 oxidation states) are mostly ionic, while nonmetals tend to form covalent molecular bromides, as do metals in high oxidation states from +3 and above. Silver bromide is very insoluble in water and is thus often used as a qualitative test for bromine.[26]
Bromine halides[edit]
The halogens form many binary, diamagnetic interhalogen 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 as BrF−
2, BrCl−
2, BrF+
2, BrF+
4, and BrF+
6. Apart from these, some pseudohalides are also known, such as cyanogen bromide (BrCN), bromine thiocyanate (BrSCN), and bromine azide (BrN3).[27]
The pale-brown bromine 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.[27] 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 in carbon tetrachloride.[26] Bromine monofluoride in ethanol readily leads to the monobromination of the aromatic 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+.[26]
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 explosively with water and hydrocarbons, but is a less violent fluorinating reagent than chlorine trifluoride. It reacts vigorously with boron, carbon, silicon, arsenic, antimony, iodine, and sulfur to give fluorides, and also reacts with most metals and their oxides: as such, it is used to oxidise uranium to uranium hexafluoride in the nuclear 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 form BrF+
2 and BrF−
4 and thus conducts electricity.[28]
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 of potassium bromide at 25 °C. It is a very vigorous fluorinating agent, although chlorine trifluoride is still more violent. Bromine pentafluoride explodes on reaction with water and fluorinates silicates at 450 °C.[29]
Polybromine compounds[edit]
Although dibromine is a strong oxidising agent with a high first ionisation energy, very strong oxidisers such as peroxydisulfuryl fluoride (S2O6F2) can oxidise it to form the cherry-red Br+
2 cation. A few other bromine cations are known, namely the brown Br+
3 and dark brown Br+
5.[30] The tribromide anion, Br−
3, has also been characterised; it is analogous to triiodide.[27]
Bromine oxides and oxoacids[edit]
Standard reduction potentials for aqueous Br species[31]
E°(couple)
a(H+) = 1
(acid)
E°(couple)
a(OH−) = 1
(base)
Br2/Br− +1.052 Br2/Br− +1.065
HOBr/Br− +1.341 BrO−/Br− +0.760
BrO−
3/Br− +1.399 BrO−
3/Br− +0.584
HOBr/Br2 +1.604 BrO−/Br2 +0.455
BrO−
3/Br2 +1.478 BrO−
3/Br2 +0.485
BrO−
3/HOBr +1.447 BrO−
3/BrO− +0.492
BrO−
4/BrO−
3 +1.853 BrO−
4/BrO−
3 +1.025
Bromine oxides are not as well-characterised as chlorine oxides or iodine oxides, as they are all fairly unstable: it was once thought that they could not exist at all. Dibromine monoxide is a dark-brown solid which, while reasonably stable at −60 °C, decomposes at its melting point of −17.5 °C; it is useful in bromination reactions[32] and may be made from the low-temperature decomposition of bromine dioxide in a vacuum. It oxidises iodine to iodine pentoxide and benzene to 1,4-benzoquinone; in alkaline solutions, it gives the hypobromite anion.[33]
So-called "bromine dioxide", a pale yellow crystalline solid, may be better formulated as bromine perbromate, 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 of hypobromous acid and bromic 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 as dibromine pentoxide, tribromine octoxide, and bromine trioxide.[33]
The four oxoacids, hypobromous acid (HOBr), bromous acid (HOBrO), bromic acid (HOBrO2), and perbromic 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:[31]
Br2 + H2O ⇌ HOBr + H+ + Br− Kac = 7.2 × 10−9 mol2 l−2
Br2 + 2 OH− ⇌ OBr− + H2O + Br− Kalk = 2 × 108 mol−1 l
Hypobromous acid is unstable to disproportionation. The hypobromite ions thus formed disproportionate readily to give bromide and bromate:[31]
3 BrO− ⇌ 2 Br− + BrO−
3 K = 1015
Bromous acids and bromites are very unstable, although the strontium and barium bromites are known.[34] More important are the bromates, which are prepared on a small scale by oxidation of bromide by aqueous hypochlorite, and are strong oxidising agents. Unlike chlorates, which very slowly disproportionate to chloride and perclorate, 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:[34]
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 radioactive beta decay of unstable 83
SeO2−
4. Today, perbromates are produced by the oxidation of alkaline bromate solutions by fluorine gas. Excess bromate and fluoride are precipitated as silver bromate and calcium 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 or xenon difluoride. The Br–O bond in BrO−
4 is fairly weak, which corresponds to the general reluctance of the 4p elements (especially arsenic, selenium, and bromine) to attain their maximum possible oxidation state, as they come after the scandide contraction characterised by the poor shielding afforded by the radial-nodeless 3d orbitals.[35]
Organobromine compounds[edit]
Main article: Organobromine compound
Structure of N-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 core organic 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 in a C–Br bond is electron-deficient and thus electrophilic. The reactivity of organobromine compounds resembles but is intermediate between the reactivity of organochlorine and organoiodine compounds. For many applications, organobromides represent a compromise of reactivity and cost.[36]
Organobromides are typically produced by additive or substitutive bromination of other organic precursors. Bromine itself can be used, but due to its toxicity and volatility safer brominating reagents are normally used, such as N-bromosuccinimide. The principal reactions for organobromides include dehydrobromination, Grignard reactions, reductive coupling, and nucleophilic substitution.[36]
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 enzyme bromoperoxidase catalyzes this reaction.[37] The oceans are estimated to release 1–2 million tons of bromoform and 56,000 tons of bromomethane annually.[38]
Bromine addition to alkene reaction mechanism
An old qualitative test for the presence of the alkene functional group is that alkenes turn brown aqueous bromine solutions colourless, forming a bromohydrin with some of the dibromoalkane also produced. The reaction passes through a short-lived strongly electrophilic bromonium intermediate. This is an example of a halogen addition reaction.[39]
Occurrence and production[edit]
View of salt evaporation pans on the Dead Sea, where Jordan (right) and Israel (left) produce salt and bromine 31°9′0″N 35°27′0″E / 31.15000°N 35.45000°E / 31.15000; 35.45000
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 the forty-sixth most abundant element in Earth's crust. It is significantly more abundant in the oceans, resulting from long-term leaching. 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, the Dead Sea contains 0.4% bromide ions.[40] It is from these sources that bromine extraction is mostly economically feasible.[41][42][43]
The main sources of bromine are in the United States and Israel. 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. 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.[44]
Applications[edit]
A wide variety of organobromine compounds are used in industry. Some are prepared from bromine and others are prepared from hydrogen bromide, which is obtained by burning hydrogen in bromine.[45]
Flame retardants[edit]
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 produces hydrobromic acid which interferes in the radical chain reaction of the oxidation reaction of the fire. The mechanism is that the highly reactive hydrogen radicals, oxygen radicals, and hydroxy 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.[46][47]
To make brominated polymers and plastics, bromine-containing compounds can be incorporated into the polymer during polymerisation. 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 of polyethylene, polyvinylchloride or polypropylene. Specific highly brominated molecules can also be added that participate in the polymerisation process For example, tetrabromobisphenol A can be added to polyesters or epoxy resins, where it becomes part of the polymer. Epoxys used in printed circuit boards are normally made from such flame retardant resins, indicated by the FR in the abbreviation of the products (FR-4 and FR-2). In some cases the bromine containing compound may be added after polymerisation. For example, decabromodiphenyl ether can be added to the final polymers.[48]
A number of gaseous or highly volatile brominated halomethane compounds are non-toxic and make superior fire suppressant agents by this same mechanism, and are particular 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 include bromochloromethane (Halon 1011, CH2BrCl), bromochlorodifluoromethane (Halon 1211, CBrClF2), and bromotrifluoromethane (Halon 1301, CBrF3).[49]
Other uses[edit]
Ethylene bromide was an additive 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).[50]
Poisonous bromomethane was widely used as pesticide to fumigate soil and to fumigate housing, by the tenting method. Ethylene bromide was similarly used.[51] These volatile organobromine compounds are all now regulated as ozone depletion agents. The Montreal Protocol on Substances that Deplete the Ozone Layer scheduled the phase out for the ozone depleting chemical by 2005, and organobromide pesticides are no longer used (in housing fumigation they have been replaced by such compounds as sulfuryl 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 control nematodes, fungi, weeds and other soil-borne diseases.[52][53]
Bromide compounds, especially potassium 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, and it was removed from over-the-counter sedative products like Bromo-Seltzer, in 1975.[54]
Other uses of organobromine compounds include high-density drilling fluids, dyes (such as Tyrian purple and the indicator bromothymol 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).[44]
Biological role and toxicity[edit]
Unlike its neighbours in the halogen column (chlorine and iodine), bromine is not known to have a confirmed biological role in humans. A single 2014 study suggests that it may be a necessary catalyst to make collagen IV, making it essential to basement membrane architecture and tissue development, but these findings have not been confirmed and the precise role posited for bromine is not certain.[55] Aside from this, only circumstantial evidence in the form of hypothesised deprivation symptoms exists for a trace role for bromine.[56] Nevertheless, even if bromine is not an essential element, it can be pharmacologically beneficial. For example, in the presence of hydrogen peroxide, H2O2, formed by the eosinophil, and either chloride or bromide ions, eosinophil peroxidase provides a potent mechanism by which eosinophils kill multicellular parasites (such as, for example, the nematode worms involved in filariasis) and some bacteria (such as tuberculosis bacteria). Eosinophil peroxidase is a haloperoxidase that preferentially uses bromide over chloride for this purpose, generating hypobromite (hypobromous acid), although the use of chloride is possible.[57]
Marine organisms are the main source of organobromine compounds, and it is in these organisms that the essentiality of bromine is on much firmer ground. More than 1600 such organobromine compounds were identified by 1999. The most abundant is methyl bromide (CH3Br), of which an estimated 56,000 tonnes is produced by marine algae each year.[58] The essential oil of the Hawaiian alga Asparagopsis taxiformis consists of 80% bromoform.[59] Most of such organobromine compounds in the sea are made by the action of a unique algal enzyme, vanadium bromoperoxidase.[60]
The bromide anion is not very toxic: a normal daily intake is 2 to 8 milligrams.[56] However, high levels of bromide chronically impair the membrane of neurons, which progressively impairs neuronal transmission, leading to toxicity, known as bromism. Bromide has an elimination 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.[61] Bromism is caused by a neurotoxic effect on the brain which results in somnolence, psychosis, seizures and delirium.[62]
Elemental bromine is toxic and causes chemical burns on human flesh. Inhaling bromine gas results in similar irritation of the respiratory tract, causing coughing, choking, and 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.[63] Caution is required when transporting bromine; it is commonly carried in steel tanks lined with lead, supported by strong metal frames.[44] The Occupational Safety and Health Administration (OSHA) of the United States has set a permissible exposure limit (PEL) for bromine at a time-weighted average (TWA) of 0.1 ppm. The National Institute for Occupational Safety and Health (NIOSH) has set a recommended exposure limit (REL) of TWA 0.1 ppm and a short-term limit of 0.3 ppm. The exposure to bromine immediately dangerous to life and health (IDLH) is 3 ppm.[64] Bromine is classified as an extremely 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.[65]
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Bibliography[edit]
Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9.
[show]
v
t
e
Bromine compounds
AgBr
AlBr3
AmBr3
AsBr3
AuBr3
BBr3
B2Br4
BaBr2
BeBr2
BiBr3
BrCl
CsBr
BrF3
BrF5
HBr
IBr
KBr
LiBr
BrN3
NaBr
BrO2
RbBr
TlBr
CaBr2
CdBr2
CoBr2
CuBr2
EuBr2
FeBr2
GeBr2
HgBr2
Hg2Br2
MgBr2
MnBr2
NiBr2
Br2O
Br2O3
Br2O5
PbBr2
PdBr2
PtBr2
S2Br2
Se2Br2
SnBr2
SrBr2
TeBr2
YbBr2
ZnBr2
CeBr3
CrBr3
CsBr3
DyBr3
ErBr3
EuBr3
FeBr3
GaBr3
GdBr3
HoBr3
InBr
InBr3
IrBr3
LaBr3
LuBr3
MoBr3
NBr3
NdBr3
PBr3
PrBr3
ReBr3
SbBr3
SmBr3
TbBr3
TmBr3
UBr3
VBr3
YBr3
YbBr3
GeBr4
HfBr4
PbBr4
SeBr4
SiBr4
SnBr4
TeBr4
ThBr4
TiBr4
UBr4
WBr4
ZrBr4
NbBr5
PBr5
TaBr5
UBr5
WBr5
CBr4
CrBr2
DyBr2
HfBr2
HfBr3
Chemical formulas
[show]
v
t
e
Diatomic chemical elements
Normal
H
2
N
2
O
2
F
2
Cl
2
Br
2
I
2
Others
C
2
P
2
S
2
Li
2
[hide]
v
t
e
Periodic table
(Large cells)
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
1
H He
2
Li Be B C N O F Ne
3
Na Mg Al Si P S Cl Ar
4
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
5
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
6
Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
7
Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
Alkali metal Alkaline earth metal Lanthanide Actinide Transition metal Post-transition metal Metalloid Polyatomic nonmetal Diatomic nonmetal Noble gas Unknown
chemical
properties
Authority control
LCCN: sh85017053
GND: 4146635-4
BNF: cb12166058m (data)
NDL: 00572452
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/wiki/Diatomic_nonmetal
mw-redirect
Diatomic nonmetal

values:
/wiki/File:Bromine_spectrum_visible.png
image
Color lines in a spectral range

values: Standard conditions for temperature and pressure

values: STP: standard temperature and pressure: 0 °C and 101.325 kPa

values:
/wiki/Kilojoule_per_mole
mw-redirect
Kilojoule per mole

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/wiki/Molar_heat_capacity
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Molar heat capacity

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values: Orthorhombic crystal structure for bromine

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Stable isotope

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Antoine Jerome Balard

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Free radical

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Well drilling fluids

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Joseph-Louis Gay-Lussac

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Académie des Sciences

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First World War

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Periodic trend

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Periodic law

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Ionisation energy

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/wiki/Alkyl_bromide
mw-redirect
Alkyl bromide

values:
/wiki/Red_phosphorus
mw-redirect
Red phosphorus

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/wiki/Dielectric_constant
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Dielectric constant

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/wiki/Carbon_tetrabromide
mw-redirect
Carbon tetrabromide

values:
/wiki/Niobium(V)_oxide
mw-redirect
Niobium(V) oxide

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values: Peroxydisulfuryl fluoride (page does not exist)

values: Edit section: Bromine oxides and oxoacids

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/wiki/Scandide_contraction
mw-redirect
Scandide contraction

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/wiki/Organochlorine_compound
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Organochlorine compound

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mw-redirect
Parts per million

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Leaching (chemical science)

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Gasoline additive

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/wiki/Ozone_depleting
mw-redirect
Ozone depleting

values: Edit section: Biological role and toxicity

values:
/wiki/Collagen_IV
mw-redirect
Collagen IV

values:
/wiki/Methyl_bromide
mw-redirect
Methyl bromide

values:
/wiki/Elimination_half-life
mw-redirect
Elimination half-life

values:
/wiki/Chemical_burns
mw-redirect
Chemical burns

values: Occupational Safety and Health Administration

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values: Immediately dangerous to life and health

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values:
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external text
//doi.org/10.1515%2Fpac-2015-0305

values:
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external text
//doi.org/10.1021%2Fed009p1915

values: Berichte der deutschen chemischen Gesellschaft

values:
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external text
//doi.org/10.1002%2Fcber.18900230395

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external text
//doi.org/10.1103%2FPhysRevB.76.104113

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/wiki/Nuclear_Physics_A
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Nuclear Physics A

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PubMed Identifier

values:
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external text
//www.ncbi.nlm.nih.gov/pubmed/15548002

values:
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external text
//doi.org/10.1021%2Fja047925p

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external text
//doi.org/10.1039%2Fa900201d

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external text
//doi.org/10.1021%2Fie50655a008

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external text
//doi.org/10.1002%2F14356007.a04_391

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external text
//www.ncbi.nlm.nih.gov/pubmed/12850087

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external text
//doi.org/10.1093%2Fcep%2Fbyi028

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external text
//www.ncbi.nlm.nih.gov/pubmed/24906154

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external text
//doi.org/10.1016%2Fj.cell.2014.05.009

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external text
http://handle.nal.usda.gov/10113/45649

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values: Journal of Agricultural and Food Chemistry

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values: United States Government Publishing Office

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values: Template talk:Bromine compounds (page does not exist)

values: Tetrabromodiborane (page does not exist)

values:
/wiki/Europium_dibromide
mw-redirect
Europium dibromide

values: Germanium(II) bromide (page does not exist)

values: Ytterbium(II) bromide (page does not exist)

values: Caesium(III) bromide (page does not exist)

values: Dysprosium(III) bromide (page does not exist)

values:
/wiki/Erbium_tribromide
mw-redirect
Erbium tribromide

values:
/wiki/Europium_tribromide
mw-redirect
Europium tribromide

values: Iridium(III) bromide (page does not exist)

values: Molybdenum bromide (page does not exist)

values:
/wiki/Neodymium_bromide
mw-redirect
Neodymium bromide

values: Rhenium(III) bromide (page does not exist)

values:
/wiki/Terbium_bromide
mw-redirect
Terbium bromide

values: Uranium tribromide (page does not exist)

values: Germanium tetrabromide (page does not exist)

values: Selenium tetrabromide (page does not exist)

values: Tellurium tetrabromine (page does not exist)

values: Thorium(IV) bromide (page does not exist)

values: Tungsten(IV) bromide (page does not exist)

values: Dysprosium(II) bromide (page does not exist)

values: Hafnium(II) bromide (page does not exist)

values: Hafnium(III) bromide (page does not exist)

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values: Template:Periodic table (32 columns, compact)

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values: Meitnerium, 109 (unknown chemical properties)

values: Darmstadtium, 110 (unknown chemical properties)

values: Roentgenium, 111 (unknown chemical properties)

values: Copernicium, 112 (post-transition metal)

values: Nihonium, 113 (unknown chemical properties)

values: Flerovium, 114 (unknown chemical properties)

values: Moscovium, 115 (unknown chemical properties)

values: Livermorium, 116 (unknown chemical properties)

values: Tennessine, 117 (unknown chemical properties)

values: Oganesson, 118 (unknown chemical properties)

values:
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periodic-table-legend

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/wiki/Polyatomic_nonmetal
mw-redirect
Polyatomic nonmetal

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values:
https://af.wikipedia.org/wiki/Broom
Broom – Afrikaans
af
af
interlanguage-link-target

values:
https://an.wikipedia.org/wiki/Bromo
Bromo – Aragonese
an
an
interlanguage-link-target

values:
https://ast.wikipedia.org/wiki/Bromu
Bromu – Asturian
ast
ast
interlanguage-link-target

values:
https://az.wikipedia.org/wiki/Brom
Brom – Azerbaijani
az
az
interlanguage-link-target

values: ব্রোমিন – Bangla

values: Бром – Belarusian (Taraškievica orthography)

values: ब्रोमिन – Bhojpuri

values: རྦོར་མིན། – Tibetan

values:
https://bs.wikipedia.org/wiki/Brom
Brom – Bosnian
bs
bs
interlanguage-link-target

values:
https://br.wikipedia.org/wiki/Brom
Brom – Breton
br
br
interlanguage-link-target

values:
https://ca.wikipedia.org/wiki/Brom
Brom – Catalan
ca
ca
interlanguage-link-target

values:
https://ceb.wikipedia.org/wiki/Bromo
Bromo – Cebuano
ceb
ceb
interlanguage-link-target

values:
https://cs.wikipedia.org/wiki/Brom
Brom – Czech
cs
cs
interlanguage-link-target

values:
https://co.wikipedia.org/wiki/Bromu
Bromu – Corsican
co
co
interlanguage-link-target

values:
https://cy.wikipedia.org/wiki/Bromin
Bromin – Welsh
cy
cy
interlanguage-link-target

values:
https://da.wikipedia.org/wiki/Brom
Brom – Danish
da
da
interlanguage-link-target

values:
https://de.wikipedia.org/wiki/Brom
Brom – German
de
de
interlanguage-link-target

values:
https://et.wikipedia.org/wiki/Broom
Broom – Estonian
et
et
interlanguage-link-target

values: Βρώμιο – Greek

values: Бром – Erzya

values:
https://es.wikipedia.org/wiki/Bromo
Bromo – Spanish
es
es
interlanguage-link-target

values:
https://eo.wikipedia.org/wiki/Bromo
Bromo – Esperanto
eo
eo
interlanguage-link-target

values:
https://eu.wikipedia.org/wiki/Bromo
Bromo – Basque
eu
eu
interlanguage-link-target

values:
https://hif.wikipedia.org/wiki/Bromine
Bromine – Fiji Hindi
hif
hif
interlanguage-link-target

values:
https://fo.wikipedia.org/wiki/Brom
Brom – Faroese
fo
fo
interlanguage-link-target

values:
https://fr.wikipedia.org/wiki/Brome
Brome – French
fr
fr
interlanguage-link-target

values:
https://fur.wikipedia.org/wiki/Brom
Brom – Friulian
fur
fur
interlanguage-link-target

values:
https://gv.wikipedia.org/wiki/Bromeen
Bromeen – Manx
gv
gv
interlanguage-link-target

values:
https://gl.wikipedia.org/wiki/Bromo
Bromo – Galician
gl
gl
interlanguage-link-target

values: બ્રોમિન – Gujarati

values:
https://hak.wikipedia.org/wiki/Chhiu
Chhiu – Hakka Chinese
hak
hak
interlanguage-link-target

values: ब्रोमीन – Hindi

values:
https://hr.wikipedia.org/wiki/Brom
Brom – Croatian
hr
hr
interlanguage-link-target

values:
https://io.wikipedia.org/wiki/Bromo
Bromo – Ido
io
io
interlanguage-link-target

values:
https://id.wikipedia.org/wiki/Bromin
Bromin – Indonesian
id
id
interlanguage-link-target

values:
https://ia.wikipedia.org/wiki/Bromo
Bromo – Interlingua
ia
ia
interlanguage-link-target

values:
https://is.wikipedia.org/wiki/Br%C3%B3m
Bróm – Icelandic
is
is
interlanguage-link-target

values:
https://it.wikipedia.org/wiki/Bromo
Bromo – Italian
it
it
interlanguage-link-target

values:
https://jv.wikipedia.org/wiki/Brom
Brom – Javanese
jv
jv
interlanguage-link-target

values: ಬ್ರೋಮಿನ್ – Kannada

values:
https://sw.wikipedia.org/wiki/Bromi
Bromi – Swahili
sw
sw
interlanguage-link-target

values:
https://ht.wikipedia.org/wiki/Bw%C3%B2m
Bwòm – Haitian Creole
ht
ht
interlanguage-link-target

values:
https://la.wikipedia.org/wiki/Bromium
Bromium – Latin
la
la
interlanguage-link-target

values:
https://lv.wikipedia.org/wiki/Broms
Broms – Latvian
lv
lv
interlanguage-link-target

values:
https://lb.wikipedia.org/wiki/Brom
Brom – Luxembourgish
lb
lb
interlanguage-link-target

values:
https://lt.wikipedia.org/wiki/Bromas
Bromas – Lithuanian
lt
lt
interlanguage-link-target

values:
https://lij.wikipedia.org/wiki/Brommo
Brommo – Ligurian
lij
lij
interlanguage-link-target

values:
https://li.wikipedia.org/wiki/Broeam
Broeam – Limburgish
li
li
interlanguage-link-target

values:
https://jbo.wikipedia.org/wiki/xunkliru
xunkliru – Lojban
jbo
jbo
interlanguage-link-target

values:
https://lmo.wikipedia.org/wiki/Brom
Brom – Lombard
lmo
lmo
interlanguage-link-target

values:
https://hu.wikipedia.org/wiki/Br%C3%B3m
Bróm – Hungarian
hu
hu
interlanguage-link-target

values: ब्रोमीन – Marathi

values:
https://ms.wikipedia.org/wiki/Bromin
Bromin – Malay
ms
ms
interlanguage-link-target

values: ဘရိုမင်း – Burmese

values: ब्रोमिन – Nepali

values: ब्रोमिन – Newari

values:
https://frr.wikipedia.org/wiki/Brom
Brom – Northern Frisian
frr
frr
interlanguage-link-target

values:
https://no.wikipedia.org/wiki/Brom
Brom – Norwegian
no
no
interlanguage-link-target

values:
https://nn.wikipedia.org/wiki/Brom
Brom – Norwegian Nynorsk
nn
nn
interlanguage-link-target

values: ବ୍ରୋମିନ – Odia

values:
https://uz.wikipedia.org/wiki/Brom
Brom – Uzbek
uz
uz
interlanguage-link-target

values: ਬ੍ਰੋਮੀਨ – Punjabi

values: ब्रोमिन – Pali

values:
https://nds.wikipedia.org/wiki/Brom
Brom – Low German
nds
nds
interlanguage-link-target

values:
https://pl.wikipedia.org/wiki/Brom
Brom – Polish
pl
pl
interlanguage-link-target

values:
https://pt.wikipedia.org/wiki/Bromo
Bromo – Portuguese
pt
pt
interlanguage-link-target

values:
https://ro.wikipedia.org/wiki/Brom
Brom – Romanian
ro
ro
interlanguage-link-target

values:
https://qu.wikipedia.org/wiki/Bromu
Bromu – Quechua
qu
qu
interlanguage-link-target

values: ब्रोमिन – Sanskrit

values:
https://sco.wikipedia.org/wiki/Bromine
Bromine – Scots
sco
sco
interlanguage-link-target

values:
https://stq.wikipedia.org/wiki/Bromium
Bromium – Saterland Frisian
stq
stq
interlanguage-link-target

values:
https://sq.wikipedia.org/wiki/Bromi
Bromi – Albanian
sq
sq
interlanguage-link-target

values:
https://scn.wikipedia.org/wiki/Bromu
Bromu – Sicilian
scn
scn
interlanguage-link-target

values: බ්‍රෝමීන් – Sinhala

values:
https://sk.wikipedia.org/wiki/Br%C3%B3m
Bróm – Slovak
sk
sk
interlanguage-link-target

values:
https://sl.wikipedia.org/wiki/Brom
Brom – Slovenian
sl
sl
interlanguage-link-target

values:
https://so.wikipedia.org/wiki/Boromiin
Boromiin – Somali
so
so
interlanguage-link-target

values:
https://sh.wikipedia.org/wiki/Brom
Brom – Serbo-Croatian
sh
sh
interlanguage-link-target

values:
https://su.wikipedia.org/wiki/Bromin
Bromin – Sundanese
su
su
interlanguage-link-target

values:
https://fi.wikipedia.org/wiki/Bromi
Bromi – Finnish
fi
fi
interlanguage-link-target

values:
https://sv.wikipedia.org/wiki/Brom
Brom – Swedish
sv
sv
interlanguage-link-target

values:
https://tl.wikipedia.org/wiki/Bromina
Bromina – Tagalog
tl
tl
interlanguage-link-target

values: புரோமின் – Tamil

values: Бром – Tatar

values: బ్రోమిన్ – Telugu

values: โบรมีน – Thai

values: Бром – Tajik

values:
https://tr.wikipedia.org/wiki/Brom
Brom – Turkish
tr
tr
interlanguage-link-target

values: برومین – Urdu

values:
https://vep.wikipedia.org/wiki/Brom
Brom – Veps
vep
vep
interlanguage-link-target

values:
https://vi.wikipedia.org/wiki/Brom
Brom – Vietnamese
vi
vi
interlanguage-link-target

values:
https://war.wikipedia.org/wiki/Bromo
Bromo – Waray
war
war
interlanguage-link-target

values:
https://zh.wikipedia.org/wiki/%E6%BA%B4
溴 – Chinese
zh
zh
interlanguage-link-target

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