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= Bromine =
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Introduction
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Bromine is a chemical element; it has symbol Br and atomic number 35.
It is a volatile red-brown liquid at room temperature that evaporates
readily to form a similarly coloured vapour. Its properties are
intermediate between those of chlorine and iodine. Isolated
independently by two chemists, Carl Jacob Löwig (in 1825) and Antoine
Jérôme Balard (in 1826), its name was derived , referring to its sharp
and pungent smell.
Elemental bromine is very reactive and thus does not occur as a free
element in nature. Instead, it can be isolated from colourless soluble
crystalline mineral halide salts analogous to table salt, a property
it shares with the other halogens. 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 evaporation ponds, mostly in the United
States and Israel. The mass of bromine in the oceans is about one
three-hundredth that of chlorine.
At standard conditions for temperature and pressure it is a liquid;
the only other element that is liquid under these conditions is
mercury. 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. The same property
causes ultraviolet sunlight to dissociate volatile organobromine
compounds in the atmosphere to yield free bromine atoms, causing ozone
depletion. As a result, many organobromine 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.
Large amounts of bromide salts are toxic from the action of soluble
bromide ions, causing bromism. However, bromine is beneficial for
human eosinophils, and is an essential trace element for collagen
development in all animals. Hundreds of known organobromine compounds
are generated by terrestrial and marine plants and animals, and some
serve important biological roles. 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.
History
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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
laboratory of 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 Latin word
("brine").
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'. In his publication, Balard stated that he changed the name
from 'muride' to 'brôme' on the proposal of . The name 'brôme'
(bromine) derives from the Greek (, "stench"). Other sources claim
that the French chemist and physicist Joseph-Louis Gay-Lussac
suggested the name 'brôme' for the characteristic smell of the vapors.
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.
By 1864, a 25% solution of liquid bromine in .75 molar aqueous
potassium bromide was widely used to treat gangrene during the
American Civil War, before the publications of Joseph Lister and
Pasteur.
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.
Properties
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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
chlorine and iodine, the two neighbouring halogens. Bromine has the
electron configuration [Ar]4s(2)3d(10)4p(5), 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 X 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 the
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
freezes 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 'π' molecular
orbital and the lowest vacant antibonding 'σ' molecular orbital. The
colour fades at low temperatures so that solid bromine at −195 °C is
pale yellow.
Liquid bromine is infrared-transparent.
Like solid chlorine and iodine, solid bromine crystallises in the
orthorhombic crystal system, in a layered arrangement of Br 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 higher
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 75 GPa it
changes to a face-centered orthorhombic structure. At 100 GPa it
changes to a body centered orthorhombic monatomic form.
Isotopes
==========
Bromine has two stable isotopes, (79)Br and (81)Br. These are its only
two natural isotopes, with (79)Br making up 51% of natural bromine and
(81)Br making up the remaining 49%. Both have nuclear spin 3/2− and
thus may be used for nuclear magnetic resonance, although (81)Br is
more favourable. The relatively 1:1 distribution of the two isotopes
in nature is helpful in identification of bromine containing compounds
using mass spectroscopy. Other bromine isotopes are all radioactive,
with half-lives too short to occur in nature. Of these, the most
important are (80)Br ('t' = 17.7 min), (80m)Br ('t' = 4.421 h), and
(82)Br ('t' = 35.28 h), which may be produced from the neutron
activation of natural bromine. The most stable bromine radioisotope is
(77)Br ('t' = 57.04 h). The primary decay mode of isotopes lighter
than (79)Br is electron capture to isotopes of selenium; that of
isotopes heavier than (81)Br is beta decay to isotopes of krypton; and
(80)Br may decay by either mode to stable (80)Se or (80)Kr. Br
isotopes from 87Br and heavier undergo beta decay with neutron
emission and are of practical importance because they are fission
products.
Chemistry and compounds
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Halogen bond energies (kJ/mol)
X XX HX BX AlX CX
F 159 574 645 582 456
Cl |243 |428 |444 |427 |327
Br |193 |363 |368 |360 |272
I |151 |294 |272 |285 |239
Bromine is intermediate in reactivity between chlorine and iodine, and
is one of the most reactive elements. Bond energies to bromine tend to
be lower than those to chlorine but higher than those to iodine, and
bromine is a weaker oxidising agent than chlorine but a stronger one
than iodine. This can be seen from the standard electrode potentials
of the X/X(−) couples (F, +2.866 V; Cl, +1.395 V; Br, +1.087 V; I,
+0.615 V; At, approximately +0.3 V). Bromination often leads to higher
oxidation states than iodination but lower or equal oxidation states
to chlorination. Bromine tends to react with compounds including M-M,
M-H, or M-C bonds to form M-Br bonds.
Hydrogen bromide
==================
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 HO + 3 Br → 6 HBr + 2 HPO
: HPO + HO + Br → 2 HBr + HPO
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 (p'K' = −9) because
the hydrogen bonds to bromine are too weak to inhibit dissociation.
The HBr/HO system also involves many hydrates HBr·'n'HO for 'n' = 1,
2, 3, 4, and 6, which are essentially salts of bromine anions and
hydronium cations. Hydrobromic acid forms an azeotrope with boiling
point 124.3 °C at 47.63 g HBr per 100 g solution; thus hydrobromic
acid cannot be concentrated beyond this point by distillation.
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 HBr(+) and ions - the latter, in
any case, are much less stable than the bifluoride ions () 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
quaternary ammonium cation (R = Me, Et, Bu('n')) 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
=======================
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 XeBr); 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, nitrogen, fluorine, and chlorine), so
that the resultant binary compounds are formally not bromides but
rather oxides, nitrides, fluorides, or chlorides of bromine.
(Nonetheless, nitrogen tribromide is named as a bromide as it is
analogous to the other nitrogen trihalides.)
Bromination of metals with Br tends to yield lower oxidation states
than chlorination with Cl 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:
:FeCl + BBr (excess) → FeBr + BCl
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 WBr + Al 3 WBr + AlBr
: EuBr + H → EuBr + HBr
: 2 TaBr TaBr + TaBr
Most metal bromides with the metal in low oxidation states (+1 to +3)
are ionic. Nonmetals tend to form covalent molecular bromides, as do
metals in high oxidation states from +3 and above. Both ionic and
covalent bromides are known for metals in oxidation state +3 (e.g.
scandium bromide is mostly ionic, but aluminium bromide is not).
Silver bromide is very insoluble in water and is thus often used as a
qualitative test for bromine.
Bromine halides
=================
The halogens form many binary, diamagnetic interhalogen compounds with
stoichiometries XY, XY, XY, and XY (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 , , ,
, and . Apart from these, some pseudohalides are also known, such as
cyanogen bromide (BrCN), bromine thiocyanate (BrSCN), and bromine
azide (BrN).
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, Bu('t'),
OMe, Br; 'meta'-bromination occurs for the deactivating X = -COEt,
-CHO, -NO); this is due to heterolytic fission of the Br-F bond,
leading to rapid electrophilic bromination by Br(+).
At room temperature, bromine trifluoride (BrF) is a straw-coloured
liquid. It may be formed by directly fluorinating bromine at room
temperature and is purified through distillation. It reacts violently
with water and explodes on contact with flammable materials, but is a
less powerful fluorinating reagent than chlorine trifluoride. It
reacts vigorously with boron, carbon, silicon, arsenic, antimony,
iodine, and sulfur to give fluorides, and will also convert most
metals and many metal compounds to fluorides; as such, it is used to
oxidise uranium to uranium hexafluoride in the nuclear power industry.
Refractory oxides tend to be only partially fluorinated, but here the
derivatives KBrF and BrFSbF remain reactive. Bromine trifluoride is a
useful nonaqueous ionising solvent, since it readily dissociates to
form and and thus conducts electricity.
Bromine pentafluoride (BrF) 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 also reacts
violently with water and is a very strong fluorinating agent, although
chlorine trifluoride is still stronger.
Polybromine compounds
=======================
Although dibromine is a strong oxidising agent with a high first
ionisation energy, very strong oxidisers such as peroxydisulfuryl
fluoride (SOF) can oxidise it to form the cherry-red cation. A few
other bromine cations are known, namely the brown and dark brown .
The tribromide anion, , has also been characterised; it is analogous
to triiodide.
Bromine oxides and oxoacids
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Standard reduction potentials for aqueous Br species
!! (acid)!!!! (base)
|Br/Br(−) +1.052 |Br/Br(−) +1.065
|HOBr/Br(−) +1.341 BrO(−)/Br(−) +0.760
|/Br(−) +1.399 /Br(−) +0.584
|HOBr/Br +1.604 BrO(−)/Br +0.455
|/Br +1.478 /Br +0.485
|/HOBr +1.447 /BrO(−) +0.492
|/ +1.853 / +1.025
Bromine oxides are not as well-characterised as chlorine oxides or
iodine oxides, as they are all fairly unstable: it was once thought
that they could not exist at all. Dibromine monoxide is a dark-brown
solid which, while reasonably stable at −60 °C, decomposes at its
melting point of −17.5 °C; it is useful in bromination reactions and
may be made from the low-temperature decomposition of bromine dioxide
in a vacuum. It oxidises iodine to iodine pentoxide and benzene to
1,4-benzoquinone; in alkaline solutions, it gives the hypobromite
anion.
So-called "bromine dioxide", a pale yellow crystalline solid, may be
better formulated as bromine perbromate, BrOBrO. It is thermally
unstable above −40 °C, violently decomposing to its elements at 0 °C.
Dibromine trioxide, 'syn'-BrOBrO, 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 (HOBrO), and perbromic acid (HOBrO), are better studied
due to their greater stability, though they are only so in aqueous
solution. When bromine dissolves in aqueous solution, the following
reactions occur:
: Br + HO HOBr + H(+) + Br(−) 'K' = 7.2 × 10(−9) mol(2) l(−2)
Br + 2 OH(−) OBr(−) + HO + Br(−) 'K' = 2 × 10(8) mol(−1) l
Hypobromous acid is unstable to disproportionation. The hypobromite
ions thus formed disproportionate readily to give bromide and bromate:
: 3 BrO(−) 2 Br(−) + 'K' = 10(15)
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 perchlorate, the bromate
anion is stable to disproportionation in both acidic and aqueous
solutions. Bromic acid is a strong acid. Bromides and bromates may
comproportionate to bromine as follows:
: + 5 Br(−) + 6 H(+) → 3 Br + 3 HO
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 . 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 is fairly weak, which corresponds to the general
reluctance of the 4p elements arsenic, selenium, and bromine to attain
their group oxidation state, as they come after the scandide
contraction characterised by the poor shielding afforded by the
radial-nodeless 3d orbitals.
Organobromine compounds
=========================
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
atom 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 annually.
Bromine addition to alkene reaction mechanism
An old qualitative test for the presence of the alkene functional
group is that alkenes turn brown aqueous bromine solutions colourless,
forming a bromohydrin with some of the dibromoalkane also produced.
The reaction passes through a short-lived strongly electrophilic
bromonium intermediate. This is an example of a halogen addition
reaction.
Occurrence and production
======================================================================
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 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. Bromine is the tenth most
abundant element in seawater.
The main sources of bromine production are Israel and Jordan. The
element is liberated by halogen exchange, using chlorine gas to
oxidise Br(−) to Br. 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.
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.
Flame retardants
==================
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 hydroxyl radicals
react with hydrobromic acid to form less reactive bromine radicals
(i.e., free bromine atoms). Bromine atoms may also react directly with
other radicals to help terminate the free radical chain-reactions that
characterise combustion.
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, polyvinyl chloride 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. Epoxies 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 particularly 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, CHBrCl), bromochlorodifluoromethane
(Halon 1211, CBrClF), and bromotrifluoromethane (Halon 1301, CBrF).
Other uses
============
Silver bromide is used, either alone or in combination with silver
chloride and silver iodide, as the light sensitive constituent of
photographic emulsions.
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).
Brominated vegetable oil (BVO), a complex mixture of plant-derived
triglycerides that have been reacted to contain atoms of the element
bromine bonded to the molecules, is used primarily to help emulsify
citrus-flavored soft drinks, preventing them from separating during
distribution.
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 diseases.
In pharmacology, inorganic 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 sodium bromide was removed from
over-the-counter sedative products like Bromo-Seltzer, in 1975.
Commercially available organobromine pharmaceuticals include the
vasodilator nicergoline, the sedative brotizolam, the anticancer agent
pipobroman, and the antiseptic merbromin. Otherwise, organobromine
compounds are rarely pharmaceutically useful, in contrast to the
situation for organofluorine compounds. Several drugs are produced as
the bromide (or equivalents, hydrobromide) salts, but in such cases
bromide serves as an innocuous counterion of no biological
significance.
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.
Bromine is used in cooling towers (in place of chlorine) for
controlling bacteria, algae, fungi, and zebra mussels.
Because it has similar antiseptic qualities to chlorine, bromine can
be used in the same manner as chlorine as a disinfectant or
antimicrobial in applications such as swimming pools. Bromine came
into this use in the United States during World War II due to a
predicted shortage of chlorine. However, bromine is usually not used
outside for these applications due to it being relatively more
expensive than chlorine and the absence of a stabilizer to protect it
from the sun. For indoor pools, it can be a good option as it is
effective at a wider pH range. It is also more stable in a heated pool
or hot tub.
Biological role and toxicity
======================================================================
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 mammals. 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, HO, formed by the eosinophil, and either chloride, iodide,
thiocyanate, or bromide ions, eosinophil peroxidase provides a potent
mechanism by which eosinophils kill multicellular parasites (such as
the nematode worms involved in filariasis) and some bacteria (such as
tuberculosis bacteria). Eosinophil peroxidase is a haloperoxidase that
preferentially uses bromide over chloride for this purpose, generating
hypobromite (hypobromous acid), although the use of chloride is
possible.
α-Haloesters are generally thought of as highly reactive and
consequently toxic intermediates in organic synthesis. Nevertheless,
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. Neutrophil myeloperoxidase can
use HO and Br(−) to brominate deoxycytidine, which could result in DNA
mutations. Marine organisms are the main source of organobromine
compounds, and it is in these organisms that bromine is more firmly
shown to be essential. More than 1600 such organobromine compounds
were identified by 1999. The most abundant is methyl bromide (CHBr),
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, vanadium
bromoperoxidase.
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. Bromism is caused by a
neurotoxic effect on the brain which results in somnolence, psychosis,
seizures and delirium.
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Elemental bromine (Br) is toxic and causes chemical burns on human
flesh. Inhaling bromine gas results in similar irritation of the
respiratory tract, causing coughing, choking, 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 3 ppm. Bromine is classified as
an extremely hazardous substance in the United States as defined in
Section 302 of the U.S. Emergency Planning and Community Right-to-Know
Act (42 U.S.C. 11002), and is subject to strict reporting requirements
by facilities which produce, store, or use it in significant
quantities.
License
=========
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License URL:
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Original Article:
http://en.wikipedia.org/wiki/Bromine
