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
Antoine Jérôme Balard (in 1826), its name was derived
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,
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.
Although large amounts are toxic and cause bromism, a clear biological
role for bromide and hypobromous acid has recently been elucidated,
and it now appears that bromine is an essential trace element. 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.
3 Chemistry and compounds
3.2 Other binary bromides
3.4 Polybromine compounds
Bromine oxides and oxoacids
3.6 Organobromine compounds
4 Occurrence and production
5.1 Flame retardants
5.2 Other uses
6 Biological role and toxicity
Antoine Balard, one of the discoverers of bromine
Bromine was discovered independently by two chemists, Carl Jacob
Löwig and Antoine Balard, in 1825 and 1826, respectively.
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
Leopold Gmelin in Heidelberg. The publication of the
results was delayed and Balard published his results first.
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
muria for brine.
After the French chemists Louis Nicolas Vauquelin, Louis Jacques
Joseph-Louis Gay-Lussac approved the experiments of the
young pharmacist Balard, the results were presented at a lecture of
Académie des Sciences
Académie des Sciences and published in Annales de Chimie et
Physique. 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). Other
sources claim that the French chemist and physicist Joseph-Louis
Gay-Lussac suggested the name brôme for the characteristic smell of
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.
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.
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. In the early years of the First World War, bromine
compounds such as xylyl bromide were used as poison gas.
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. 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. 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. 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. The volatility of bromine accentuates its very
penetrating, choking, and unpleasant odour.
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. 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.
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. The colour fades at low temperatures, so that solid
bromine at −195 °C is pale yellow.
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.
At a pressure of 55
GPa (roughly 540,000 times atmospheric
pressure) bromine undergoes an insulator-to-metal transition. At
GPa it changes to a face-centered orthorhombic structure. At
GPa it changes to a body centered orthorhombic monatomic
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. 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
Chemistry and compounds
Halogen bond energies (kJ/mol)
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
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:
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. 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
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
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.
Other binary bromides
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.
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. Another method is halogen exchange in the
presence of excess "halogenating reagent", for example:
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:
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
Silver bromide is very insoluble in water and is thus often
used as a qualitative test for bromine.
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
4, and BrF+
6. Apart from these, some pseudohalides are also known, such as
cyanogen bromide (BrCN), bromine thiocyanate (BrSCN), and bromine
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.
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. 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+.
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
Bromine trifluoride is a useful nonaqueous ionising solvent,
since it readily dissociates to form BrF+
2 and BrF−
4 and thus conducts electricity.
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.
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. The tribromide anion, Br−
3, has also been characterised; it is analogous to triiodide.
Bromine oxides and oxoacids
Standard reduction potentials for aqueous Br species
a(H+) = 1
a(OH−) = 1
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 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
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.
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
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
3 BrO− ⇌ 2 Br− + BrO−
K = 1015
Bromous acids and bromites are very unstable, although the strontium
and barium bromites are known. 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
Bromic acid is a strong acid. Bromides and bromates may
comproportionate to bromine as follows:
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 83SeO2−
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.
Main article: Organobromine compound
Structure of N-bromosuccinimide, a common brominating reagent in
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.
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.
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. The oceans are estimated to release 1–2
million tons of bromoform and 56,000 tons of bromomethane
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
Occurrence and production
View of salt evaporation pans on the Dead Sea, where Jordan (right)
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. It is from these sources that
bromine extraction is mostly economically feasible.
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.
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.
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
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.
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,
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).
Poisonous bromomethane was widely used as pesticide to fumigate soil
and to fumigate housing, by the tenting method. Ethylene bromide was
similarly used. 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
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.
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). Zinc–bromine batteries
are hybrid flow batteries used for stationary electrical power backup
and storage; from household scale to industrial scale.
Biological role and toxicity
2-Octyl 4-bromo-3-oxobutanoate, an organobromine compound found in
mammalian cerebrospinal fluid
A 2014 study suggests that bromine (in the form of bromide ion) is a
necessary cofactor in the biosynthesis of collagen IV, making the
element essential to basement membrane architecture and tissue
development in animals. Nevertheless, no clear deprivation
symptoms or syndromes have been documented. In other biological
functions, bromine may be non-essential but still beneficial when it
takes the place of chlorine. For example, in the presence of hydrogen
peroxide, H2O2, formed by 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
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. Although α-haloesters are generally thought of as
highly reactive, and therefore, toxic intermediates in organic
synthesis, mammals, including humans, cats, and rats, appear to
biosynthesize traces of an α-bromoester, 2-octyl
4-bromo-3-oxobutanoate, which is found in their cerebrospinal fluid
and appears to play a yet unclarified role in inducing REM sleep.
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. The essential oil of the Hawaiian alga Asparagopsis
taxiformis consists of 80% bromoform. Most of such organobromine
compounds in the sea are made by the action of a unique algal enzyme,
The bromide anion is not very toxic: a normal daily intake is 2 to
8 milligrams. 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.
caused by a neurotoxic effect on the brain which results in
somnolence, psychosis, seizures and delirium.
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. Caution is required when transporting
bromine; it is commonly carried in steel tanks lined with lead,
supported by strong metal frames. 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
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
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.
^ a b Haynes, William M., ed. (2011). CRC Handbook of Chemistry and
Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.121.
^ Lide, D. R., ed. (2005). "
Magnetic susceptibility of the elements
and inorganic compounds".
CRC Handbook of Chemistry and Physics
CRC Handbook of Chemistry and Physics (PDF)
(86th ed.). Boca Raton (FL): CRC Press. ISBN 0-8493-0486-5.
^ Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca
Raton, Florida: Chemical Rubber Company Publishing. pp. E110.
^ Löwig, Carl Jacob (1829). "Das Brom und seine chemischen
Bromine and its chemical relationships] (in German).
Heidelberg: Carl Winter.
^ a b c Balard, A. J. (1826). "Mémoire sur une substance
particulière contenue dans l'eau de la mer" [Memoir on a peculiar
substance contained in sea water]. Annales de Chimie et de Physique.
2nd series (in French). 32: 337–381.
^ a b Balard, Antoine (1826). "Memoir on a peculiar Substance
contained in Sea Water". Annals of Philosophy. 28: 381–387 and
^ Weeks, Mary Elvira (1932). "The discovery of the elements: XVII. The
halogen family". Journal of Chemical Education. 9 (11): 1915.
^ Landolt, Hans Heinrich (1890). "Nekrolog: Carl Löwig". Berichte der
deutschen chemischen Gesellschaft. 23 (3): 905–909.
^ Vauquelin, L. N.; Thenard, L.J.; Gay-Lussac, J.L. (1826). "Rapport
sur la Mémoire de M. Balard relatif à une nouvelle Substance"
[Report on a memoir by Mr. Balard regarding a new substance]. Annales
de Chimie et de Physique. 2nd series (in French). 32: 382–384.
^ On page 341 of his article, A. J. Balard (1826) "Mémoire sur une
substance particulière contenue dans l'eau de la mer" [Memoir on a
peculiar substance contained in sea water], Annales de Chimie et de
Physique, 2nd series, vol. 32, pp. 337–381, Balard states that Mr.
Anglada persuaded him to name his new element brôme. However, on page
382 of the same journal – "Rapport sur la Mémoire de M. Balard
relatif à une nouvelle Substance" [Report on a memoir by Mr. Balard
regarding a new substance], Annales de Chimie et de Physique, series
2, vol. 32, pp. 382–384. – a committee of the French Academy of
Sciences claimed that they had renamed the new element brôme.
^ Wisniak, Jaime (2004). "Antoine-Jerôme Balard. The discoverer of
bromine" (PDF). Revista CENIC Ciencias Químicas. 35 (1):
^ Greenwood and Earnshaw, p. 790
^ Barger, M. Susan; White, William Blaine (2000). "Technological
Practice of Daguerreotypy". The Daguerreotype: Nineteenth-century
Technology and Modern Science. JHU Press. pp. 31–35.
^ Shorter, Edward (1997). A History of Psychiatry: From the Era of the
Asylum to the Age of Prozac. John Wiley and Sons. p. 200.
^ Corey J Hilmas, Jeffery K Smart, Benjamin A Hill (2008). "Chapter 2:
History of Chemical Warfare (pdf)". Medical Aspects of Chemical
Warfare (PDF). Borden Institute. pp. 12–14. CS1 maint:
Uses authors parameter (link)
^ a b c d e f g h Greenwood and Earnshaw, pp. 800–4
^ "Johann Wolfgang Dobereiner". Purdue University. Archived from the
original on 2014-11-14. Retrieved 2008-03-08.
^ "A Historic Overview: Mendeleev and the Periodic Table" (PDF). NASA.
^ Greenwood and Earnshaw, p. 793–4
^ a b c Greenwood and Earnshaw, pp. 804–9
^ Duan, Defang; et al. (2007-09-26). "Ab initio studies of solid
bromine under high pressure". Physical Review B. 76 (10): 104113.
^ Audi, G.; Bersillon, O.; Blachot, J.; Wapstra, A. H. (2003). "The
NUBASE evaluation of nuclear and decay properties" (PDF). Nuclear
Physics A. 729: 3–128. Bibcode:2003NuPhA.729....3A.
doi:10.1016/j.nuclphysa.2003.11.001. Archived from the original (PDF)
^ a b Greenwood and Earnshaw, pp. 809–12
^ a b Greenwood and Earnshaw, pp. 812–6
^ a b c d e f g Greenwood and Earnshaw, pp. 821–4
^ a b c Greenwood and Earnshaw, pp. 824–8
^ Greenwood and Earnshaw, pp. 828–31
^ Greenwood and Earnshaw, pp. 832–5
^ Greenwood and Earnshaw, pp. 842–4
^ a b c Greenwood and Earnshaw, pp. 853–9
^ Perry, Dale L.; Phillips, Sidney L. (1995), Handbook of Inorganic
Compounds, CRC Press, p. 74, ISBN 0-8493-8671-3, retrieved
25 August 2015
^ a b Greenwood and Earnshaw, pp. 850–1
^ a b Greenwood and Earnshaw, pp. 862–5
^ Greenwood and Earnshaw, pp. 871–2
^ a b Ioffe, David and Kampf, Arieh (2002) "Bromine, Organic
Compounds" in Kirk-Othmer Encyclopedia of Chemical Technology. John
Wiley & Sons. doi: 10.1002/0471238961.0218151325150606.a01.
^ Carter-Franklin, Jayme N.; Butler, Alison (2004). "Vanadium
Bromoperoxidase-Catalyzed Biosynthesis of Halogenated Marine Natural
Products". Journal of the American Chemical Society. 126 (46):
15060–6. doi:10.1021/ja047925p. PMID 15548002.
^ Gribble, Gordon W. (1999). "The diversity of naturally occurring
organobromine compounds". Chemical Society Reviews. 28 (5): 335–346.
^ Clayden, Jonathan; Greeves, Nick; Warren, Stuart (2012). Organic
Chemistry (2nd ed.). Oxford University Press. pp. 427–9.
^ Greenwood and Earnshaw, pp. 795–6
^ Tallmadge, John A.; Butt, John B.; Solomon Herman J. (1964).
"Minerals From Sea Salt". Ind. Eng. Chem. 56 (7): 44–65.
^ Oumeish, Oumeish Youssef (1996). "Climatotherapy at the
Dead Sea in
Jordan". Clinics in Dermatology. 14 (6): 659–664.
^ Al-Weshah, Radwan A. (2008). "The water balance of the Dead Sea: an
integrated approach". Hydrological Processes. 14: 145–154.
^ a b c Greenwood and Earnshaw, pp. 798–9
^ Mills, Jack F. (2002). Bromine: in Ullmann's Encyclopedia of
Chemical Technology. Weinheim: Wiley-VCH Verlag.
^ Green, Joseph (1996). "Mechanisms for Flame Retardancy and Smoke
suppression – A Review". Journal of Fire Sciences. 14 (6):
^ 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.
^ Weil, Edward D.; Levchik, Sergei (2004). "A Review of Current Flame
Retardant Systems for Epoxy Resins". Journal of Fire Sciences. 22:
^ Günter Siegemund, Werner Schwertfeger, Andrew Feiring, Bruce Smart,
Fred Behr, Herward Vogel, Blaine McKusick "
Organic" Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH,
Weinheim, 2002. doi:10.1002/14356007.a11_349
^ 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):
683–9. doi:10.1016/S0160-4120(03)00121-1. PMID 12850087.
^ Lyday, Phyllis A. "Mineral Yearbook 2007: Bromine" (PDF). United
States Geological Survey. Retrieved 2008-09-03.
^ Messenger, Belinda; Braun, Adolf (2000). "Alternatives to Methyl
Bromide for the Control of Soil-Borne Diseases and Pests in
California" (PDF). Pest Management Analysis and Planning Program.
^ 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.
^ Samuel Hopkins Adams (1905). The Great American fraud. Press of the
American Medical Association. Retrieved 2011-06-25.
^ McCall AS; Cummings CF; Bhave G; Vanacore R; Page-McCaw A; et al.
Bromine Is an Essential Trace Element for Assembly of
Collagen IV Scaffolds in Tissue Development and Architecture". Cell.
157 (6): 1380–92. doi:10.1016/j.cell.2014.05.009.
PMC 4144415 . PMID 24906154.
^ a b Nielsen, Forrest H. (2000). "Possibly Essential Trace Elements".
Clinical Nutrition of the Essential Trace Elements and Minerals:
^ Mayeno AN; Curran AJ; Roberts RL; Foote CS (1989). "Eosinophils
preferentially use bromide to generate halogenating agents". J. Biol.
Chem. 264 (10): 5660–8. PMID 2538427.
^ Gribble, Gordon W. (1999-01-01). "The diversity of naturally
occurring organobromine compounds". Chemical Society Reviews. 28 (5).
doi:10.1039/A900201D. ISSN 1460-4744.
^ Gribble, Gordon W. (1999). "The diversity of naturally occurring
organobromine compounds". Chemical Society Reviews. 28 (5): 335–346.
^ Burreson, B. Jay; Moore, Richard E.; Roller, Peter P. (1976).
"Volatile halogen compounds in the alga Asparagopsis taxiformis
(Rhodophyta)". Journal of Agricultural and Food Chemistry. 24 (4):
^ Butler, Alison; Carter-Franklin, Jayme N. (2004). "The role of
vanadium bromoperoxidase in the biosynthesis of halogenated marine
natural products". Natural Product Reports. 21 (1): 180–8.
doi:10.1039/b302337k. PMID 15039842.
^ Olson, Kent R. (1 November 2003). Poisoning & drug overdose (4th
ed.). Appleton & Lange. pp. 140–141.
^ Galanter, Marc; Kleber, Herbert D. (1 July 2008). The American
Psychiatric Publishing Textbook of Substance Abuse Treatment (4th
United States of America: American Psychiatric Publishing Inc.
p. 217. ISBN 978-1-58562-276-4.
^ Science Lab.com. "Material Safety Data Sheet:
sciencelab.com. Retrieved 27 October 2016.
^ "NIOSH Pocket Guide to Chemical Hazards #0064". National Institute
for Occupational Safety and Health (NIOSH).
^ "40 C.F.R.: Appendix A to Part 355—The List of Extremely Hazardous
Substances and Their Threshold Planning Quantities" (PDF) (July 1,
2008 ed.). Government Printing Office. Retrieved October 29,
Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements
(2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9.
Diatomic chemical elements
Periodic table (Large cells)
Alkaline earth metal
BNF: cb12166058m (data)