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"Bromo" redirects here. For other uses, see Bromo (disambiguation).
Bromine, 35Br
Liquid and gas bromine inside transparent cube
Bromine
Pronunciation/ˈbroʊmiːn, -mɪn, -maɪn/ ​(BROH-meen, -⁠min, -⁠myne)
Appearancereddish-brown
Standard atomic weight Ar°(Br)
Bromine in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
Cl
↑
Br
↓
 I 
selenium ← bromine → krypton
Atomic number (Z)35
Groupgroup 17 (halogens)
Periodperiod 4
Block  p-block
Electron configuration[Ar] 3d10 4s2 4p5
Electrons per shell2, 8, 18, 7
Physical properties
Phase at STPliquid
Melting point(Br2) 265.8 K ​(−7.2 °C, ​19 °F)
Boiling point(Br2) 332.0 K ​(58.8 °C, ​137.8 °F)
Density (near r.t.)Br2, liquid: 3.1028 g/cm3
Triple point265.90 K, ​5.8 kPa[3]
Critical point588 K, 10.34 MPa[3]
Heat of fusion(Br2) 10.571 kJ/mol
Heat of vaporization(Br2) 29.96 kJ/mol
Molar heat capacity(Br2) 75.69 J/(mol·K)
Vapor pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 185 201 220 244 276 332
Atomic properties
Oxidation statescommon: −1, +1, +3, +5
+2,[4] +4,[5] +7[5]
ElectronegativityPauling scale: 2.96
Ionization energies
  • 1st: 1139.9 kJ/mol
  • 2nd: 2103 kJ/mol
  • 3rd: 3470 kJ/mol
Atomic radiusempirical: 120 pm
Covalent radius120±3 pm
Van der Waals radius185 pm
Color lines in a spectral range
Spectral lines of bromine
Other properties
Natural occurrenceprimordial
Crystal structure ​orthorhombic (oS8)
Lattice constants
Orthorhombic crystal structure for bromine
a = 674.30 pm
b = 466.85 pm
c = 870.02 pm (at triple point: 269.60 K)[6]
Thermal conductivity0.122 W/(m⋅K)
Electrical resistivity7.8×1010 Ω⋅m (at 20 °C)
Magnetic orderingdiamagnetic[7]
Molar magnetic susceptibility−56.4×10−6 cm3/mol[8]
Speed of sound206 m/s (at 20 °C)
CAS Number7726-95-6
History
Discovery and first isolationAntoine Jérôme Balard and Carl Jacob Löwig (1825)
Isotopes of bromine
Main isotopes[9] Decay
abun­dance half-life (t1/2) mode pro­duct
75Br synth 96.7 min β+ 75Se
76Br synth 16.2 h β+ 76Se
77Br synth 57.04 h β+ 77Se
79Br 50.6% stable
80Br synth 17.68 min β− 80Kr
80mBr synth 4.4205 h IT 80Br
81Br 49.4% stable
82Br synth 35.282 h β− 82Kr
 Category: Bromine
| references

Bromine (Template:Pron-en BROH-meen or /ˈbroʊmɨn/ BROH-min, from [βρῶμος, brómos] Error: {{Lang-xx}}: text has italic markup (help), meaning "stench (of he-goats)"),[10] is a chemical element with the symbol Br and atomic number 35. A halogen element, bromine is a reddish-brown volatile liquid at standard room temperature that is intermediate in reactivity between chlorine and iodine. Bromine vapors are corrosive and toxic. Approximately 556,000 metric tons were produced in 2007.[11] The main applications for bromine are in fire retardants and fine chemicals.

History

Illustrative and secure bromine sample for teaching

Bromine was discovered independently by two chemists Antoine Balard[12] and Carl Jacob Löwig[13] in 1825 and 1826.[14]

Balard found bromide chemicals in the ash of sea weed from the salt marshes of Montpellier in 1826. The sea weed 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 resembled that of an intermediate of chlorine and iodine; with those results 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.[12]

Carl Jacob 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 diethylether. After evaporation of the ether a brown liquid remained. With this liquid as a sample for 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.[13]

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.[15] In his publication Balard states that he changed the name from muride to brôme on the proposal of M. Anglada. Other sources claim that the French chemist and physicist Joseph-Louis Gay-Lussac suggested the name brôme for the characteristic smell of the vapors.[16] Bromine was not produced in large quantities until 1860.

The first commercial use, besides some minor medical applications, was the use of bromine for the daguerreotype. In 1840 it was discovered that bromine had some advantages over the previously used iodine vapor to create the light sensitive silver halide layer used for daguerreotypy.[17]

Potassium bromide and sodium bromide were used as anticonvulsants and sedatives in the late 19th and early 20th centuries, until they were gradually superseded by chloral hydrate and then the barbiturates.[18]

Characteristics

Bromine is the only nonmetallic element that is a liquid at room temperature, and one of only two elements on the periodic table that are liquids at room temperature (mercury is the other). The melting point of bromine is −7.2 °C and the boiling point 58.8 °C (138 °F). The pure chemical element has the physical form of a diatomic molecule, Br2. It is a dense, mobile, slightly transparent reddish-brown liquid, that evaporates easily at standard temperature and pressures to give a red vapor (its color resembles nitrogen dioxide) that has a strong disagreeable odor resembling that of chlorine. Bromine is a halogen, and is less reactive than chlorine and more reactive than iodine. Bromine is slightly soluble in water, and highly soluble in carbon disulfide, aliphatic alcohols (such as methanol), and acetic acid. It bonds easily with many elements and has a strong bleaching action. Bromine, like chlorine, is also used in maintenance of swimming pools.

Certain bromine-related compounds have been evaluated to have an ozone depletion potential or bioaccumulate in living organisms. As a result many industrial bromine compounds are no longer manufactured, are being restricted, or scheduled for phasing out. The Montreal Protocol mentions several organobromine compounds for this phase out.

Bromine is a powerful oxidizing agent. It reacts vigorously with metals, especially in the presence of water, as well as most organic compounds, especially upon illumination.

Isotopes

Main article: Isotopes of bromine

Bromine has 2 stable isotopes: 79Br (50.69 %) and 81Br (49.31%). At least another 23 radioisotopes are known to exist.[19] Many of the bromine isotopes are fission products. Several of the heavier bromine isotopes from fission are delayed neutron emitters. All of the radioactive bromine isotopes are relatively short lived. The longest half life is the neutron deficient 77Br at 2.376 days. The longest half life on the neutron rich side is 82Br at 1.471 days. A number of the bromine isotopes exhibit metastable isomers. Stable 79Br exhibits a radioactive isomer, with a half life of 4.86 seconds. It decays by isomeric transition to the stable ground state.[20]

Allotropes

At a pressure of 55 GPa bromine converts to a metal. At 75 GPa it converts to a face centered orthorhombic structure. At 100 GPa it converts to a body centered orthorhombic monoatomic form.[21]

Occurrence and production

See also: Category:Halide minerals and Halide mineral
World bromine production trend
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

The diatomic element Br2 does not occur naturally. Instead, bromine exists exclusively as bromide salts in diffuse amounts in crustal rock. Due to leaching, bromide salts have accumulated in sea water (65 ppm),[22] but at a lower concentration than chloride. Bromine may be economically recovered from bromide-rich brine wells and from the Dead Sea waters (up to 50000 ppm).[23][24]

Approximately 556,000 metric tonnes (worth around US$2.5 billion) of bromine are produced per year (2007) worldwide with the United States, Israel, and China being the primary producers.[25][26][27] Bromine production has increased sixfold since the 1960s. The largest bromine reserve in the United States is located in Columbia and Union County, Arkansas, U.S.[28] China's bromine reserves are located in the Shandong Province and Israel's bromine reserves are contained in the waters of the Dead Sea. The bromide-rich brines are treated with chlorine gas, flushing through with air. In this treatment, bromide anions are oxidized to bromine by the chlorine gas.

2 Br− + Cl2 → 2 Cl− + Br2

Because of its commercial availability and long shelf-life, bromine is not typically prepared. Small amounts of bromine can however be generated through the reaction of solid sodium bromide with concentrated sulfuric acid (H2SO4). The first stage is formation of hydrogen bromide (HBr), which is a gas, but under the reaction conditions some of the HBr is oxidized further by the sulfuric acid to form bromine (Br2) and sulfur dioxide (SO2).

NaBr (s) + H2SO4 (aq) → HBr (aq) + NaHSO4 (aq)
2 HBr (aq) + H2SO4 (aq) → Br2 (g) + SO2 (g) + 2 H2O (l)

Similar alternatives, such as the use of dilute hydrochloric acid with sodium hypochlorite, are also available. The most important thing is that the anion of the acid (in the above examples, sulfate and chloride, respectively) be more electronegative than bromine, allowing the substitution reaction to occur.

Reaction involving a strong oxidizing agent, such as potassium permanganate, on bromide ions in the presence of an acid also gives bromine. An acidic solution of bromate ions and bromide ions will also disproportionate slowly to give bromine.

Bromine is only slightly soluble in water. But the solubility can be increased by the presence of bromide ions. However, concentrated solutions of bromine are rarely prepared in the lab as they will continually give off toxic red-brown bromine gas due to its very high vapor pressure. Sodium thiosulphate is an excellent reagent for getting rid of bromine completely including the stains and odor.

Compounds

See also: Category:Bromine compounds

Organic chemistry

N-Bromosuccinimide

Organic compounds are brominated by either addition or substitution reactions. Bromine undergoes electrophilic addition to the double-bonds of alkenes, via a cyclic bromonium intermediate. In non-aqueous solvents such as carbon disulfide, this affords the di-bromo product. For example, reaction with ethylene will produce 1,2-dibromoethane. Bromine also undergoes electrophilic addition to phenols and anilines. When used as bromine water, a small amount of the corresponding bromohydrin is formed as well as the dibromo compound. So reliable is the reactivity of bromine that bromine water is employed as a reagent to test for the presence of alkenes, phenols, and anilines. Like the other halogens, bromine participates in free radical reactions. For example hydrocarbons are brominated upon treatment with bromine in the presence of light.

Bromine, sometimes with a catalytic amount of phosphorus, easily brominates carboxylic acids at the α-position. This method, the Hell-Volhard-Zelinsky reaction, is the basis of the commercial route to bromoacetic acid. N-Bromosuccinimide is commonly used as a substitute for elemental bromine, being easier to handle, and reacting more mildly and thus more selectively. Organic bromides are often preferable relative to the less reactive chlorides and more expensive iodide-containing reagents. Thus, Grignard and organolithium compound are most often generated from the corresponding bromides.

Inorganic chemistry

Oxidation states
of bromine
−1 HBr
0 Br
2
+1 BrCl
+3 BrF
3
+5 BrF
5
+5 BrO−
3
+7 BrO−
4

Bromine is an oxidizer, and it will oxidize iodide ions to iodine, being itself reduced to bromide:

Br2 + 2 I− → 2 Br− + I2

Bromine will also oxidize metals and metalloids to the corresponding bromides. Anhydrous bromine is less reactive toward many metals than hydrated bromine, however. Dry bromine reacts vigorously with aluminium, titanium, mercury as well as alkaline earths and alkali metals.

If bromine is dissolved in hydroxide containing water not only bromide (Br−) is formed, but also the hypobromite (OBr−). This hypobromite is responsible for the bleaching abilities of bromide solutions. In warm solutions the disproportion reaction of the hypobromite is quantitative. The resulting bromate is a strong oxidising agent and very similar to the chlorate.

3 BrO−
 → BrO−
3 + 2 Br−

The perbromates are not accessible through electrolysis like the perchlorates, but only by reacting bromate solutions with fluorine or ozone.

BrO3− + H2O + F2 → BrO−
4 + 2 HF
BrO3− + O3 → BrO−
4 + O2

Applications

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.[11]

Illustrative of the addition reaction[29] is the preparation of 1,2-dibromoethane, the organobromine compound produced in the largest amounts:

C2H4 + Br2 → CH2BrCH2Br

Flame retardant

Tetrabromobisphenol A

Brominated flame retardants represent a commodity of growing importance. If the material burns the flame retardants produce hydrobromic acid which interferes in the radical chain reaction of the oxidation reaction of the fire. The highly reactive hydrogen oxygen and hydroxy radicals react with hydrobromic acid and form less reactive bromine radicals.[30][31] The bromine-containing compounds can be placed in the polymers either during polymerization if a small amount of brominated monomer is added or the bromine containing compound is added after polymerization. Tetrabromobisphenol A can be added to produce polyesters or epoxy resins. Epoxy used in printed circuit boards (PCB) are normally made from flame retardant resins, indicated by the FR in the abbreviation of the products (FR-4 and FR-2. Vinyl bromide can be used in the production of polyethylene, polyvinylchloride or polypropylene. Decabromodiphenyl ether can be added to the final polymers.[32]

Gasoline additive

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 uses in 1966 in the US. This application has declined since the 1970s due to environmental regulations.[33] Ethylene bromide is also used as a fumigant, but again this application is declining.[27]

Pesticide

Methyl bromide (bromomethane)

Methyl bromide was widely used as pesticide to fumigate soil. The Montreal Protocol on Substances that Deplete the Ozone scheduled the phase out for the ozone depleting chemical until 2005. In 1991, an estimated 35,000 metric tonnes of the chemical were used to control nematodes, fungi, weeds and other soil-borne diseases.[34][35]

Medical and veterinary

  • Bromide compounds, especially potassium bromide, were frequently used as 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.

Other uses

Orange fluoresces of DNA Ethidium bromide intercalate
  • The bromides of calcium, sodium, and zinc account for a sizable part of the bromine market. These salts form dense solutions in water that are used as drilling fluids sometimes called clear brine fluids.[27][36]
  • Bromine is also used in the production of brominated vegetable oil, which is used as an emulsifier in many citrus-flavored soft drinks (e.g. Mountain Dew). After the introduction in the 1940s the compound was extensively used until the UK and the US limited its use in the mid 1970s and alternative emulsifiers were developed.[37]
    Soft drinks containing brominated vegetable oil are still sold in the US (2009).[38]
Tralomethrin
  • Several dyes, agrichemicals, and pharmaceuticals are organobromine compounds. 1-Bromo-3-chloropropane, 1-bromoethylbenzene, and 1-bromoalkanes are prepared by the antimarkovnikov addition of HBr to alkenes. Ethidium bromide, EtBr, is used as a DNA stain in gel electrophoresis.
  • High refractive index compounds
  • Water purification compounds, disinfectants and insecticides, such as tralomethrin (C22H19Br4NO3).[27]
  • Potassium bromide is used in some photographic developers to inhibit the formation of fog (undesired reduction of silver).
  • Vapor is used as the second step in sensitizing daguerreotype plates to be developed under Mercury (Hg) vapor. Bromine acts as an accelerator to the light sensitivity of the previously iodized plate.
  • Bromine is also used to reduce mercury pollution from coal-fired power plants. This can be achieved either by treating activated carbon with bromine or by injecting bromine compounds onto the coal prior to combustion.

Biological role

Tyrian purple

Bromine has no known essential role in human or mammalian health, but inorganic bromine and organobromine compounds do occur naturally, and some may be of use to higher organisms in dealing with parasites. For example, in the presence of 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 also certain bacteria (such as tuberculosis bacteria). Eosinophil peroxidase is a haloperoxidase that preferentially uses bromide over chloride for this purpose, generating hypobromite (hypobromous acid).[39]

Marine organisms are the main source of organobromine compounds. Over 1600 compounds were identified by 1999. The most abundant one is methyl bromide with estimated 56,000 metric tonnes produced by marine algae.[40] The essential oil of the Hawaiian alga Asparagopsis taxiformis consists of 80% methyl bromide.[41] A famous example of a bromine-containing organic compound that has been used by humans for a long time is Tyrian purple.[40][42] The brominated indigo is produced by a medium-sized predatory sea snail, the marine gastropod Murex brandaris. It took until 1909 before the organobromine nature of the compound was discovered by Paul Friedländer.[43] Most organobromine compounds in nature arise via the action of vanadium bromoperoxidase.[44]

Safety

See also: List of highly toxic gases

Elemental bromine is toxic and causes burns. As an oxidizing agent, it is incompatible with most organic and inorganic compounds. Care needs to be taken when transporting bromine; it is commonly carried in steel tanks lined with lead, supported by strong metal frames.

When certain ionic compounds containing bromine are mixed with potassium permanganate (KMnO4) and an acidic substance, they will form a pale brown cloud of bromine gas. This gas smells like bleach and is very irritating to the mucus membranes. Upon exposure, one should move to fresh air immediately. If symptoms of bromine poisoning arise, medical attention is needed.

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  43. ^ Friedländer, P. (1909). "Über den Farbstoff des antiken Purpurs aus murex brandaris". Berichte der deutschen chemischen Gesellschaft. 42: 765. doi:10.1002/cber.190904201122.
  44. ^ Butler, Alison (2004). "The role of vanadium bromoperoxidase in the biosynthesis of halogenated marine natural products". Natural Product Reports. 21: 180. doi:10.1039/b302337k. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)
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