Bromine

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"Bromo" redirects here. For other uses, see Bromo (disambiguation).
35 selenium ← bromine → krypton
Cl
↑
Br
↓
I
General
Name, Symbol, Number bromine, Br, 35
Element category halogens
Group, Period, Block 17, 4, p
Appearance gas/liquid: red-brown
solid: metallic luster
Standard atomic weight 79.904(1)  g·mol−1
Electron configuration [Ar] 4s2 3d10 4p5
Electrons per shell 2, 8, 18, 7
Physical properties
Phase liquid
Density (near r.t.) (Br2, liquid) 3.1028  g·cm−3
Melting point 265.8 K
(-7.2 °C, 19 °F)
Boiling point 332.0 K
(58.8 °C, 137.8 °F)
Critical point 588 K, 10.34 MPa
Heat of fusion (Br2) 10.571  kJ·mol−1
Heat of vaporization (Br2) 29.96  kJ·mol−1
Specific heat capacity (25 °C) (Br2)
75.69  J·mol−1·K−1
Vapor pressure
P(Pa) 1 10 100 1 k 10 k 100 k
at T(K) 185 201 220 244 276 332
Atomic properties
Crystal structure orthorhombic
Oxidation states 7, 5, 4, 3, 1, -1
(strongly acidic oxide)
Electronegativity 2.96 (Pauling scale)
Ionization energies
(more)		 1st:  1139.9  kJ·mol−1
2nd:  2103  kJ·mol−1
3rd:  3470  kJ·mol−1
Atomic radius 120  pm
Covalent radius 120±3  pm
Van der Waals radius 185 pm
Miscellaneous
Magnetic ordering diamagnetic[1]
Electrical resistivity (20 °C) 7.8×1010  Ω·m
Thermal conductivity (300 K) 0.122  W·m−1·K−1
Speed of sound (20 °C) ? 206 m/s
CAS registry number 7726-95-6
Most-stable isotopes
Main article: Isotopes of bromine
iso NA half-life DM DE (MeV) DP
79Br 50.69% 79Br is stable with 44 neutrons
81Br 49.31% 81Br is stable with 46 neutrons
References

Bromine (pronounced /ˈbroʊmiːn/ or /ˈbroʊmɨn/), Greek: βρῶμος, brómos, meaning "stench (of he-goats)" [2]), 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 vapours are corrosive and toxic. Approximately 556,000 metric tonnes were produced in 2007. [3] The main applications for bromine are in fire retardants and fine chemicals.

Contents

[edit] History

Illustrative and secure bromine sample for teaching

Bromine was discovered independently by two chemists Antoine Balard[4] and Carl Jacob Löwig[5] in 1825 and 1826.[6]

Balard found bromide salts 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 found a new element and named it muride, derived from the Latin word muria for brine.[4]

Carl Jacob Löwig isolated bromine from mineral water spring from his home town 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.[5]

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 where presented at a lecture of the Académie des Sciences and published in Annales de Chimie et Physique[7]. 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 due to the characteristic smell of the vapors.[8] Bromine was not produced in quantity 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 previous used iodine vapour to create the light sensitive silver halide layer used for daguerreotypy.[9]

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.[10]

[edit] Characteristics

Bromine is the only liquid nonmetallic element at room temperature, and one of only two elements on the periodic table that are liquid at room temperature. The melting point of bromine is −7.2 °C and has the boiling point 58.8 °C. The pure chemical element has the physical form of a diatomic molecule, Br2. It is a dense, mobile, 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.

[edit] 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.[11] 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.[12]

[edit] Occurrence and production

See also Halide minerals.
World bromine production trend

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),[13] 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).[14][15]

View of salt evaporation pans on the Dead Sea, where Jordan (right) and Israel (left) produce salt and bromine

Approximately 556,000 metric tons (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.[16][17][18] 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.[19] 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 between 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 vapour pressure. Sodium thiosulphate is an excellent reagent for getting rid of bromine completely including the stains and odour.

[edit] Compounds

See also: Category:Bromine compounds

[edit] 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.

[edit] Inorganic chemistry

Oxidation states
of bromine
-1 HBr
+1 BrCl
+3 BrF3
+5 BrF5
+5 BrO3−
+7 BrO4−

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 quantitive. The resulting bromate is a strong oxidation reagent and very similar to the chlorate.

3 OBr− → BrO3− + 2 Br−

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

OBr− + H2O + F2 → BrO4− + 2 HF

[edit] 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.[3]

Illustrative of the addition reaction[20] is the preparation of 1,2-Dibromoethane, the organobromine compound produced in the largest amounts:

C2H4 + Br2 → CH2BrCH2Br

[edit] Flame retardant

Tetrabromobisphenol A

Brominated flame retardants represent a commodity of growing importance. If the material burns the flame retardents 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.[21][22] The bromine containing compounds can be placed in the polymeres either during polymerisation if a small amount of brominated monomer is added or the bromine containing compound is added after polymerisation. Tetrabromobisphenol A can be added to produce polyesters or epoxy resins. An epoxy resigns used for printed circuit boards (PCB) are normally made from flame retardant resigns, 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 polymeres.[23]

[edit] 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.[24] Ethylene bromide is also used as a fumigant, but again this application is declining.[18]

[edit] 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.[25][26]

[edit] Other Use

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.[18][27]
  • Bromine is also used in for 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 emulsifers were developed.[28] By 1997 in the US still soft drinks are available containing brominated vegetable oil.[29]
Tralomethrin

[edit] Biological role

Tyrian purple

Bromine has no known role in human health, but organobromine compounds do occur naturally. Marine organisms are the main source of organobromine compounds. In 1999 over 1600 compounds were identified. The most abundant one is methyl bromide with a estimate amount of 56000 metric tons produced by marine algae.[30] The essential oil of the Hawaiian alga Asparagopsis taxiformis consists of 80% methyl bromide.[31] A famous example of a bromine-containing organic compound that has been used by humans for a long time is Tyrian purple.[30][32] 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.[33] Most organobromine compounds in nature arise via the action of vanadium bromoperoxidase.[34]

[edit] 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 arise, medical attention is needed.

[edit] References

  1. ^ Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics 81th edition, CRC press.
  2. ^ Gemoll W, Vretska K: Griechisch-Deutsches Schul- und Handwörterbuch ("Greek-German dictionary"), 9th ed., published by öbvhpt, ISBN 3-209-00108-1
  3. ^ a b Jack F. Mills "Bromine" in Ullmann's Encyclopedia of Chemical Technology Wiley-VCH Verlag; Weinheim, 2002. DOI: 10.1002/14356007.a04_391
  4. ^ a b Balard, Antoine (1826). "Memoire of a peculire Substance contained in Sea Water". Annals of Philosophy: 387– and 411–. http://books.google.de/books?id=A-M4AAAAMAAJ. 
  5. ^ a b Landolt, Hans Heinrich (1890). "Nekrolog: Carl Löwig". Berichte der deutschen chemischen Gesellschaft 23 (3): 905–909. doi:10.1002/cber.18900230395. http://gallica.bnf.fr/ark:/12148/bpt6k907222/f920.chemindefer. 
  6. ^ Weeks, Mary Elvira (1932). "The discovery of the elements: XVII. The halogen family.". Journal of Chemical Education 9: 1915. 
  7. ^ Balard, A.J. Annales de Chimie et Physique (1826). 32. pp. 337–382. 
  8. ^ Wisniak, Jaime (2004). "Antoine-Jerôme Balard. The discoverer of bromine". Revista CENIC Ciencias Químicas 35 (1). http://revistas.mes.edu.cu:9900/EDUNIV/03-Revistas-Cientificas/Rev.CENIC-Ciencias-Quimicas/2004/1/09204109.pdf. 
  9. ^ Barger, M. Susan; White, William Blaine (2000). "Technological Practice of Daguerreotypy". The Daguerreotype: Nineteenth-century Technology and Modern Science. JHU Press. pp. 31–35. ISBN 9780801864582. 
  10. ^ Shorter, Edward (1997). A History of Psychiatry: From the Era of the Asylum to the Age of Prozac. John Wiley and Sons. pp. 200–202. ISBN 9780471245315. 
  11. ^ GE Nuclear Energy (1989). Chart of the Nuclides, 14th Edition. 
  12. ^ Audi, Georges (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A (Atomic Mass Data Center) 729: 3–128. doi:10.1016/j.nuclphysa.2003.11.001. 
  13. ^ Tallmadge, John A.; Butt, John B.; Solomon Herman J.. "Minerals From Sea Salt". Ind. Eng. Chem. 56 (7): 44–65. doi:10.1021/ie50655a008. 
  14. ^ Oumeish, Oumeish Youssef (1996). "Climatotherapy at the Dead Sea in Jordan". Clinics in Dermatology 14 (6): 659–664. doi:10.1016/S0738-081X(96)00101-0. 
  15. ^ Al-Weshah, Radwan A. (2008). "The water balance of the Dead Sea: an integrated approach". Hydrological Processes 14 (1): 145–154. 
  16. ^ Emsley, John (2001). "Bromine". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. pp. 69–73. ISBN 0198503407. 
  17. ^ Lyday, Phyllis A.. "Comodity Report 2007: Bromine". United States Geological Survey. http://minerals.usgs.gov/minerals/pubs/commodity/bromine/bromimcs07.pdf. Retrieved on 2008-09-03. 
  18. ^ a b c d Lyday, Phyllis A.. "Mineral Yearbook 2007: Bromine". United States Geological Survey. http://minerals.usgs.gov/minerals/pubs/commodity/bromine/myb1-2006-bromi.pdf. Retrieved on 2008-09-03. 
  19. ^ Bromine:An Important Arkansas Industry, Butler Center for Arkansas Studies
  20. ^ N. A. Khan, F. E. Deatherage, and J. B. Brown (1963). "Stearolic Acid". Org. Synth.; Coll. Vol. 4: 851. 
  21. ^ Green, Joseph (1996). "Mechanisms for Flame Retardancy and Smoke suppression -A Review". Journal of Fire Sciences 14 (6): 426–442. doi:10.1177/073490419601400602. 
  22. ^ Kaspersma, Jelle; Doumena, Cindy; Munrob Sheilaand; Prinsa, Anne-Marie (2002). "Fire retardant mechanism of aliphatic bromine compounds in polystyrene and polypropylene". Polymer Degradation and Stability 77 (2): 325–331. doi:10.1016/S0141-3910(02)00067-8. 
  23. ^ Weil, Edward D.; Levchik, Sergei (2004). "A Review of Current Flame Retardant Systems for Epoxy Resins". Journal of Fire Sciences 22 (1): 25–40. doi:10.1177/0734904104038107. 
  24. ^ Alaeea, Mehran; Ariasb, Pedro; Sjödinc, Andreas; Bergman, Åke (2003). "An overview of commercially used brominated flame retardants, their applications, their use patterns in different countries/regions and possible modes of release". Environment International 29 (6): 683–689. doi:10.1016/S0160-4120(03)00121-1. 
  25. ^ Messenger, Belinda; Braun, Adolf (2000). "Alternatives to Methyl Bromide for the Control of Soil-Borne Diseases and Pests in California". Pest Management Analysis and Planning Program. http://www.cdpr.ca.gov/docs/emon/methbrom/alt-anal/sept2000.pdf. Retrieved on 2008-11-17. 
  26. ^ Decanio, Stephen J.; Norman, Catherine S. (2008). "Economics of the "Critical Use" of Methyl bromide under the Montreal Protocol". Contemporary Economic Policy 23 (3): 376–393. doi:10.1093/cep/byi028. 
  27. ^ Darley, H. C. H.; Gray, George Robert (1988). Composition and Properties of Drilling and Completion Fluids. Gulf Professional Publishing. pp. 61–62. ISBN 9780872011472. 
  28. ^ Kaufman, Vered R.; Garti, Nissim (1984). "Effect of cloudy agents on the stability and opacity of cloudy emulsions for soft drinks". International Journal of Food Science & Technology 19 (2): 255–261. doi:10.1111/j.1365-2621.1984.tb00348.x. 
  29. ^ Horowitz, B. Zane (1997). Bromism from Excessive Cola Consumption',Clinical Toxicology. 35. pp. 315–320. doi:10.3109/15563659709001219. 
  30. ^ a b Gordon W. Gribble (1999). "The diversity of naturally occurring organobromine compounds". Chemical Society Reviews 28 (5): 335. doi:10.1039/a900201d. 
  31. ^ "Volatile halogen compounds in the alga Asparagopsis taxiformis (Rhodophyta)". Journal of Agricultural snd Food Chemistry 24 (4): 856–861. 1976. doi:10.1021/jf60206a040. 
  32. ^ Gordon W. Gribble (1998). "Naturally Occurring Organohalogen Compounds". Acc. Chem. Res. 31 (3): 141–152. doi:10.1021/ar9701777. 
  33. ^ Friedländer, P.. "Über den Farbstoff des antiken Purpurs aus murex brandaris". Berichte der deutschen chemischen Gesellschaft 42 (1): 765–770. 
  34. ^ Butler, Alison; Carter-Franklin, Jayme N. (2004). "The role of vanadium bromoperoxidase in the biosynthesis of halogenated marine natural products". Natural Product Reports 21: 180–188. doi:10.1039/b302337k. 

[edit] External links

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v • d • e
Diatomic chemical elements

Hydrogen H2 | Nitrogen N2 | Oxygen O2 | Fluorine F2 | Chlorine Cl2 | Bromine Br2 | Iodine I2 | Astatine At2 |
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