My watch list  


35 seleniumbrominekrypton


Name, Symbol, Number bromine, Br, 35
Chemical series 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
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 5, 4,[1] 3,[2] 1, -1
(strongly acidic oxide)
Electronegativity 2.96 (Pauling scale)
Ionization energies
1st:  1139.9  kJ·mol−1
2nd:  2103  kJ·mol−1
3rd:  3470  kJ·mol−1
Atomic radius 115  pm
Atomic radius (calc.) 94  pm
Covalent radius 114  pm
Van der Waals radius 185 pm
Magnetic ordering nonmagnetic
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
Selected isotopes
Main article: Isotopes of bromine
iso NA half-life DM DE (MeV) DP
79Br 50.69% Br is stable with 44 neutrons
81Br 49.31% Br is stable with 46 neutrons

Bromine (pronounced /ˈbroʊmiːn/, /ˈbroʊmaɪn/, /ˈbroʊmɪn/, Greek: βρῶμος, brómos, meaning "stench (of he-goats)" [3]), is a chemical element with the symbol Br and atomic number 35. A halogen element, bromine is a red volatile liquid at standard room temperature that is intermediate in reactivity between chlorine and iodine. Bromine vapours are corrosive toxic. Approximately 730,000,000 kg were produced in 1993.[4] The main applications for bromine are in fire retardants and fine chemicals.



Bromine was discovered by Antoine Balard at the salt marshes of Montpellier in 1826, but was not produced in quantity until 1860. The French chemist and physicist Joseph-Louis Gay-Lussac suggested the name bromine due to the characteristic smell of the vapors. Some also suggest that it may have been discovered by Bernard Courtois, the man who discovered iodine.


Main article: Isotopes of bromine

Bromine has 2 stable isotopes: Br-79 (50.69%) and Br-81 (49.31%). At least another 23[5] isotopes are known to exist. 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 Br-77 at 2.376 days. The longest half life on the neutron rich side is Br-82 at 1.471 days. A number of the bromine isotopes exhibit metastable isomers. Stable Br-79 exhibits a radioactive isomer, with a half life of 4.86 seconds. It decays by isomeric transition to the stable ground state.

Notable characteristics

Bromine is the only liquid nonmetallic element at room temperature and one of only six elements on the periodic table that are liquid at or close to room temperature. 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.

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.

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.

Bromine has no known role in human health. Organobromine compounds do occur naturally, a famous example being Tyrian purple. Most organobromine compounds in nature arise via the action of vanadium bromoperoxidase.

Occurrence and production

See also Halide minerals.

  The diatomic compound 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 (85 ppm), 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).

Approximately 500,000 metric tons (worth around US$350 million) of bromine are produced per year (2001) worldwide with the United States and Israel being the primary producers. 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.[6] 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 anion is 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 by further 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.


See also: Category:Bromine compounds

Organic chemistry

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, the corresponding bromohydrin is formed instead. So reliable is the reactivity or bromine that bromine water is employed as a reagent to test for the presence 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 reagent]]s. Thus, Grignard and organolithium compound are most often generated from the corresponding bromides.

Inorganic chemistry

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 metaloids 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.


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.[4]

Illustrative of the addition reaction[7] is the preparation of 1,2-Dibromoethane, the organobromine compound produced in the largest amounts:

C2H4 + Br2 → CH2BrCH2Br

Ethylene bromide is a 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 has declined since the 1970's due to environmental regulations. Ethylene bromide is also used as a fumigant, but again this application is declining.

Brominated flame retardants represent a commodity of growing importance. Specific compound used produced for this purpose include tetrabromobisphenol A, decabromodiphenyl ether, and vinyl bromide.

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.

Miscellaneous uses:


Elemental bromine is toxic and causes burns. As an oxidizing agent, it is incompatible with most organic and inorganic compounds.

See also


  1. ^ Bromine: bromine(IV) oxide compound data. Retrieved on 2007-12-10.
  2. ^ Bromine: bromine(III) fluoride compound data. Retrieved on 2007-12-10.
  3. ^ Gemoll W, Vretska K: Griechisch-Deutsches Schul- und Handwörterbuch ("Greek-German dictionary"), 9th ed., published by öbvhpt, ISBN 3-209-00108-1
  4. ^ a b Jack F. Mills "Bromine" in Ullmann's Encyclopedia of Chemical Technology Wiley-VCH Verlag; Weinheim, 2002. DOI: 10.1002/14356007.a04_391
  5. ^ GE Nuclear Energy (1989). Chart of the Nuclides, 14th Edition. 
  6. ^ Bromine:An Important Arkansas Industry, Butler Center for Arkansas Studies
  7. ^ N. A. Khan, F. E. Deatherage, and J. B. Brown (1963). "Stearolic Acid". Org. Synth.; Coll. Vol. 4: 851. 
  • Los Alamos National Laboratory – Bromine
This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Bromine". A list of authors is available in Wikipedia.
Your browser is not current. Microsoft Internet Explorer 6.0 does not support some functions on Chemie.DE