Swain
What’s special about bromine?
What’s special about
bromine?
Pat Swain
This article focuses on aspects of the chemistry of bromine as a
typical halogen and as a special element with many useful
applications
Extracting the element from
the sea
Much of the bromine produced world-wide is
extracted from bromine-rich brine deposits from
below the Earth’s surface or from very saline waters
such as the Dead Sea. Some, however, is extracted
from sea water. There was a great demand for bromine
when a plant to manufacture it from coastal sea water
at Amlwch on the island of Anglesey commenced
operation in 1953 (Figure 1). This is the world’s largest
plant of its type, extracting over 30 000 tonnes each
year.
The raw materials used in the extraction process
are sea water, chlorine and SO2 gas made by burning
sulfur in air. The operation of the plant involves four
main stages. First, the sea water, containing an average
0.0065% by mass of bromide ions, is pumped in at a
rate of one third of a million gallons per minute. The
pH is then adjusted to about 3.5 to prevent hydrolysis
occurring at the second stage, when the Br– ions are
ABSTRACT
The article starts with information about an
industrial plant in North Wales which has been
extracting bromine from coastal sea water since
the 1950s. Then follows an account of the
discovery of bromine by Balard in 1826 and how
this event assisted scientists with theories about
arranging chemical elements, many years before
the first Periodic Table of the elements was
published. The article goes on to provide details
about the terrestrial occurrence and properties of
bromine, considered with the other halogens,
and the special uses of bromine in organic
chemistry. Other important applications for this
element past and present and its large-scale
manufacture are discussed.
oxidised to molecular bromine by injecting chlorine
gas into the water:
2Br–(aq) + Cl2(aq) → Br2(aq) + 2Cl–(aq)
Hydrolysis of molecular bromine and chlorine is
minimised at this acidity since the position of
equilibrium is well over to the left in the following
equations:
Cl2(aq) + H2O(l)
Br2(aq) + H2O(l)
HClO(aq) + H+(aq) + Cl–(aq)
HBrO(aq) + H+(aq) + Br–(aq)
The solution is passed down a large blowing-out tower
where bromine is stripped by an upward current of
air. Sea-water effluent flows from the bottom of the
towers and is made much less acidic by dilution before
passing into the sea.
At the third stage, bromine in the moist air leaving
the towers at the top is reduced to hydrogen bromide
with SO2 gas:
Br2(aq) + SO2(g) + 2H2O(l) →
H2SO4(aq) + 2HBr(aq)
Fresh water is added to the acidic vapour produced
before it goes through a special absorber packing layer
to remove air. The acid solution at this point contains
dilute sulfuric acid and about 13% bromine by mass
as hydrobromic acid .
At the fourth stage more chlorine is added to this
solution to displace bromine from the bromide ions.
Finally the solution passes to a distillation tower where
the required end-product, molecular bromine, is
separated by steam distillation.
Safety
Liquid bromine is classified as a poisonous material,
and hazardous brominated compounds are also made
on the site. They have to be transported and distributed
in accordance with national and international
School Science Review, September 2003, 85(310)
75
What’s special about bromine?
regulations, using special containers for road vehicles
(Figure 2). All aspects of the work at the Amlwch
plant involve handling hazardous materials so safety
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receives a great deal of attention. To retain its accreditation certificate as a ‘Firm of Assessed Capability’
there is a twice-yearly audit by safety inspectors.
Figure 1 A view of the works at Amlwch.
Figure 2 Transporting hazardous chemicals at the Amlwch plant.
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School Science Review, September 2003, 85(310)
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The discovery by Balard
This unpleasant reddish-brown, fuming element,
discovered by the French chemist Antoine-Jérôme
Balard (1802–1876) and published in 1826, is the only
liquid non-metal at room temperature. It attacks cork
and rubber so is usually kept in an all-glass container
(Figure 3). When he made this discovery Balard was
a young chemistry assistant at the Montpellier School
of Pharmacy in France who was interested in
Mediterranean salt-marsh flora and chemicals from
the sea.
Figure 3 A glass-stoppered bottle containing
bromine.
Balard isolated the element by passing chlorine
through the aqueous residues left over at the salt works
after manufacture of sea salt. Distillation of the
residues after addition of manganese dioxide and
sulfuric acid yielded a dark-red vapour which
condensed to a dark-red liquid. The analogy with the
What’s special about bromine?
preparation of chlorine aided his identification of a
new element. It was named ‘bromine’ (Greek bromos,
for ‘stench’) by the French Academy when Balard’s
proposed name muride was considered unacceptable.
In 1825, Carl Löwig, a new young student of
chemistry at Heidelberg, had shown his professor a
dark-red liquid he obtained by passing chlorine
through liquor from a salt spring near his home. He
was asked to prepare some more and study its
properties. Also, Justus von Liebig (1803–1873),
professor of chemistry at the university of Giessen,
had earlier isolated a new substance on passing
chlorine into mother liquor from salt works, followed
by distillation. He had decided it was a liquid chloride
of iodine and duly labelled the sample, which he kept
on a shelf in his laboratory.
It was therefore fortunate that Balard was quick
to publish his results, for on reading these both Löwig
and Liebig realised that they had also obtained the
new element bromine. They both later published
works on bromine and Liebig was able to use his
wrongly labelled sample for his studies.
In the same year that bromine was discovered,
the Swedish chemist Jons Jacob Berzelius (1779–
1848) published his painstaking and surprisingly
accurate measurements of atomic weights for the 50
known elements, which included chlorine and iodine
(see Table 1). Since it had rapidly been established
that bromine had chemical properties similar to those
of chlorine and iodine, the German chemist Johann
Wolfgang Döbereiner (1780–1849) was able to
include these three halogens (Greek hal for ‘salt’; gen
‘to produce’) as a triad for his theory of triads ( see
Box 1). He predicted from the arithmetic mean of
Berzelius’ atomic weights for chlorine and iodine that
the atomic weight for bromine would be about 80,
Table 1 Three halogen elements are discovered.
Element
Discoverer and brief details
Recognised as a new
element by:
Named by:
Chlorine
C. W. Scheele in 1774.
Reaction of manganese
dioxide with hydrochloric acid.
B. Courtois in 1811. Reaction
of sulfuric acid with seaweed
(Fucus) ashes. Announced in
1813.
A. J. Balard in 1826. Reaction
of chlorine with aqueous
residues from salt works.
Humphry Davy in 1810
Humphry Davy. From a Greek
word meaning ‘green’.
J. L. Gay-Lussac in 1813.
Humphry Davy in 1813.
J. L. Gay-Lussac. From a
Greek word meaning ‘violet’.
A. J. Balard in 1826 when
reporting his discovery.
The French academy. From a
Greek word meaning ‘stench’.
Iodine
Bromine
School Science Review, September 2003, 85(310)
77
What’s special about bromine?
Box 1
■
■
■
■
■
Döbereiner’s triads
In 1817 J. W. Döbereiner made the first attempt to group elements with similar chemical properties,
according to atomic weights.
He noticed that there were some triads (groups of three) of elements with similar chemical properties.
Elements of a triad could have atomic weights almost the same, or when arranged in order of atomic
weight the second element was approximately the arithmetic mean of the other two. This latter was
stated as his ‘Law of Triads’ in 1827.
His triads were:
Metals
Alkali metals
Alkaline earths Halogens
iron
lithium
calcium
chlorine (35.47)
cobalt
sodium
strontium
bromine
nickel
potassium
barium
iodine (126.47)
In 1827 he predicted from his Law that the atomic weight of bromine would be about 80.
From 1827 to 1858, other scientists extended Döbereiner’s relationships as new elements were
added to triad groups, and new families of elements were discovered:
fluorine added to the halogens;
magnesium added to the alkaline earths;
oxygen, sulfur, selenium, tellurium classified as a family of elements;
nitrogen, phosphorus, arsenic, antimony, bismuth, classified as a family of elements.
Table 2 Abundances of halogen elements.
Occurrence Concentration (ppm)
Fluorine Chlorine Bromine Iodine
Terrestrial
1108
390
(total)
Earth’s crust 360
140
Sea water
1.3 19 400
1.96
2.4
66
0.122
0.1
0.06
although the actual mean value of 80.97 is rather high.
In 1858 a pamphlet by the Italian chemist Stanislao
Cannizzaro (1826–1910), whose atomic weights were
used for the first periodic classifications, gave the
atomic weight of bromine as 80, determined from
vapour density measurements.
Terrestrial abundance and
occurrence in nature
Since Mendeléev published his Periodic Table in
1871, the halogens have formed the well recognised
group of non-metals, with bromine coming third in
order of reactivity on descending the group. Due to
their high chemical reactivity these elements are only
found in a combined state in nature. Bromine is also
third in order of terrestrial abundance (Table 2). It is
the 46th most abundant element in the Earth’s crust
78
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School Science Review, September 2003, 85(310)
and the seventh in ordinary sea water. There is
considerable enrichment of chlorine and bromine in
oceans and salt lakes due to the solubility of their salts.
Bromine chiefly occurs as bromide ions, alongside
the much more abundant chloride ions. It is found in
mineral deposits of rock salt (NaCl) and carnallite
(KCl.MgCl2.6H2O), such as the underground salt beds
near Stassfurt in Germany. The bromide ion
concentration can be as high as 800–2800 ppm in
some brines and mineral springs, as in the underground brines in Michigan and Arkansas in the USA
and inland lakes in California and in the Crimea. In
the Dead Sea, which contains about 35% total salts
compared to 3.5% in ordinary sea water, concentrations can be as high as 4000 ppm. Some silver
bromide is found in halide minerals such as cerargyrite
(horn silver) associated with native silver and other
mineral deposits in Germany, Chile and western USA.
In New Zealand, natural geothermal waters contain
magnesium and potassium bromides. The associated
mineral-containing mud has, reputedly, therapeutic
properties and is used in special soap preparations.
Many marine plants and animals have been shown
to contain bromine. The ancient Tyrian Purple dye,
found to be a dibromoindigo, was made from an
eastern Mediterranean sea snail Murex brandaris. The
sponges, phylum Porifera, contain bromine (or
iodine) in spongins – scleroproteins similar to the
keratin in skin, claws and hair. These marine species,
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including certain seaweeds and shellfish, take in and
concentrate bromine from their surroundings.
Many hundreds of organobromine compounds are
now known to be produced naturally by biological
processes in the environment. A host of sea creatures
such as slugs, sponges, corals and even mammals
along with bacteria, plants, seaweeds, fungi, lichen,
and algae, are known to be responsible. These compounds include ordinary molecules such as methyl
bromide, bromoform and bromophenols, as well as
new biochemicals with properties currently of special
interest.
Bromine is not an essential element for human
life. Small amounts present in the body as bromide
ions (Br–) are not concentrated in tissues but excreted
in the urine like chloride and iodide ions. In contrast
chlorine is a main mineral element occurring as
chloride ions in body fluids and iodine is an important
trace element found in the thyroid gland.
What’s special about bromine?
Table 3 Physical data for bromine.
Relative atomic mass
Atomic number
Electronic structure
79.904
35
1s2,2s2,2p6,3s2,
3p6,3d10,4s2,4p5
Melting point
–7.3 oC
Boiling point
59 oC
o
3.1
Specific gravity at 20 C
Stable isotopes in normal Br 50.52% 79Br;
49.48% 81Br
Radioactive isotopes
Seventeen from
72
Br to 92Br
I
I–
Br
Br–
Cl
Cl –
Physical characteristics and
typical chemistry
Physical chemistry data for bromine are shown in
Table 3. It is the fifth most electronegative element in
the Periodic Table, with a Pauling electronegativity
value of 2.8, surpassed only by nitrogen (3.0), oxygen
(3.5), fluorine (4.0) and chlorine (3.0) but not by
iodine (2.5). The size of its atom and ion are
intermediate between chlorine and iodine, consistent
with its central position in the group VII halogens
(Figure 4).
Like the other halogen atoms, bromine acquires a
stable octet of electrons (a noble gas structure) in the
outer electronic shell by gaining one electron when it
forms a negative bromide ion in ionic (electrovalent)
compounds, or by sharing an electron in covalent
compounds. The element normally exists as diatomic
molecules by forming a covalent bond between the
atoms, giving each a stable outer shell of 8 electrons:
x x
x
x
x
x
x
Br x Br xx
x x
x
Bromine molecule
(x = outer shell electron)
x
It is less reactive than chlorine since its outer valency
electrons are further from the positive nucleus, but
conversely it is more reactive than iodine. Some bond
enthalpy (∆Hθ/kJ mol–1) values illustrating this trend
are 431 (HCl), 366 (HBr) and 299 (HI). There is a
similar trend in the melting points (Tm/°C ) of halide
salts containing ionic bonds, e.g. 808 (NaCl), 750
0.2 nm
Figure 4 Sizes of the atoms and ions from covalent
and ionic radii.
(NaBr), 662 (NaI) and 772 (CaCl2), 765 (CaBr2), 740
(CaI2). Some of the typical inorganic reactions of
halogens as exhibited by bromine are shown in Box
2 (overleaf).
Bromide salts produce bromine when heated with
concentrated sulfuric acid and manganese dioxide,
since the hydrogen bromide initially formed is
oxidised. The equation for the usual preparation of
bromine by the reaction of potassium bromide is:
2KBr(aq) +3H2SO4(aq) + MnO2(s) →
2KHSO4(aq) + MnSO4(aq) + 2H2O(l)+ Br2(g)
In the 1950s, when glass retorts in school laboratories
were normally fitted with rubber bungs or corks, a
special glass stopper was needed which would not be
attacked by the bromine vapour (Figure 5).
School Science Review, September 2003, 85(310)
79
What’s special about bromine?
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glass stoppered retort
Box 2
Inorganic reactions of
bromine
Metals
Bromine reacts directly with most metals to form
bromides:
Sodium:
2Na(s) + Br2(l) → 2NaBr(s)
Potassium: 2K(s) + Br2(l) → 2KBr(s)
Magnesium: Mg(s) + Br2(l) → MgBr2(s)
Calcium:
Ca(s) + Br2(l) → CaBr2(s)
Copper:
Cu(s) + Br2(l) → CuBr2(s)
Zinc:
Zn(s) + Br2(l) → ZnBr2(s)
Aluminium: 2Al(s) + 3Br2(l) → 2AlBr3(s)
Iron:
4Fe(s) + 6Br2(l) → 2Fe2Br6(s)
Non-metals
Bromine reacts directly with many non-metals
forming bromides; e.g.
Phosphorus: 2P(s) + 3Br2(l) → 2PBr3(l);
2P(s) + 5Br2(l) → 2PBr5(s)
Silicon:
Si(s) + 2Br2(l) → SiBr4(l)
Hydrogen: H2(g) + Br2(l) → 2HBr(g)
It reacts with water:
It forms bromine water containing (bromic(I)) acid
and hydrobromic acid:
H+(aq) +Br–(aq) + HBrO(aq)
H2O(l) + Br2(l)
It reacts with alkalis:
A solution of a bromate(I) and a bromide is
formed when bromine reacts with excess cold
dilute alkali; e.g.
2NaOH(aq)+ Br2(l) →
NaOBr(aq) + NaBr(aq) + H2O(l)
A solution of a bromate(V) and a bromide is
formed when an excess of bromine reacts with
hot concentrated alkali; e.g.
6KOH(aq) + 3Br2(l) →
KBrO3(aq) + 5KBr(aq) + 3H2O(l)
Some other oxidation reactions of bromine
Ammonia is oxidised to nitrogen:
2NH3(g) + 3Br2(l) → N2(g) + 6HBr(g)
Hydrogen sulfide is oxidised to sulfur:
H2S(g) + Br2(l) → S(s) + 2HBr(g)
Sulfuric(IV) acid is oxidised to sulfuric(VI) acid:
H2SO3(aq) + Br2(l) + H2O(l) →
H2SO4 (aq) + 2HBr(aq)
Iodides are oxidised to iodine; e.g.
2KI(aq) + Br2(l) → I2(s) + 2KBr(aq)
CaI2(aq) + Br2(l) → I2(s) + CaBr2(aq)
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cold water jet
heat
liquid bromine
KBr + MnO2 + conc. H 2SO4
Figure 5 All-glass apparatus in the laboratory
preparation of bromine.
The metal bromides are ionic compounds,
generally very soluble in water. They can be identified
by adding silver nitrate solution. The pale yellow
precipitate of silver bromide formed if Br– ions are
present is only sparingly soluble in dilute ammonium
hydroxide but dissolves in a concentrated solution
because a soluble complex salt is formed:
AgBr(s) + 2NH4OH(aq) →
[Ag(NH3)2Br](aq) +2H2O(l)
Chlorides form white AgCl, soluble in dilute NH4OH,
and iodides form yellow AgI, only slightly soluble in
concentrated NH4OH.
The non-metal bromides are covalent compounds.
Hydrogen bromide is a colourless gas that dissolves
readily in water to give the strong monobasic hydrobromic acid:
HBr(g) + H2O(l) → H3O+(aq) + Br–(aq)
It is a stronger acid and reducing agent than HCl since
the H—Br bond is weaker than the H—Cl bond
(shown by bond enthalpies ∆Hθ/kJ mol–1 of 366 and
431 respectively). Concentrated sulfuric acid is
reduced to sulfur dioxide by hydrogen bromide:
2HBr(aq) + H2SO4(l) →
Br2(aq) + 2H2O(l) + SO2(g)
For this reason HBr cannot be prepared by the action
of concentrated sulfuric acid on a metal bromide since
it would be oxidised to bromine. The usual laboratory
method of preparation is drop-wise addition of liquid
bromine to a mixture of red phosphorus and water,
producing in situ phosphorus(III) and phosphorus(V)
bromides which immediately hydrolyse. Hydrogen
bromide is evolved:
PBr3 + 3H2O → 3HBr + H3PO3
and
PBr5 + 4H2O → 5HBr + H3PO4
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Bromine has three unstable acidic oxides, Br2O, BrO2
and BrO 3, which dissolve in water. The known
oxyacids, bromic(I), bromic(V) and bromic(VII) acids,
form salts with alkalis which are oxidising agents;
e.g. sodium bromate( I ) (NaOBr), potassium
bromate(v) (KBrO3), potassium bromate(VII) (KBrO4).
Bromine forms a number of stable binary
interhalogen compounds with fluorine, chlorine and
iodine, which can be made by direct combination.
Compounds known to exist, and their physical states
at 20 °C, are:
What’s special about bromine?
molecule, usually listed in common electrophiles
(electron-seeking), forms a complex with the π
electrons by becoming positively polarised at one end,
i.e. Brδ+ –Brδ–. This complex breaks down to form an
intermediate carbocation with a bromine atom
covalently bonded to one carbon of the double (or
triple) bond, rapidly followed by nucleophilic addition
of the residual Br– ion to the other carbon. An
unsymmetrical molecule such as hydrogen bromide
(Hδ+ –Brδ–) can react similarly with an alkene. In this
case the mono bromide is formed:
BrF(l); BrCl(g); IBr(s); BrF3(l); BrF5(l)
Since the interhalogen bond is weaker than that
between like halogens (fluorine is an exception), these
compounds are more reactive than the individual
halogens. Some, like bromine trifluoride, are able to
ionise in the molten state:
[BrF2]+ + [BrF4]–
2BrF3
A reagent in organic chemistry
Addition across the carbon-carbon double bond of
ethene, the simplest alkene, occurs when a few drops
of bromine water are added to the gas in a covered
gas jar. The bromine colour quickly disappears on
vigorous shaking.
C2H4 + Br2 → C2H4Br2
There is a similar addition to the carbon-carbon triple
bond in ethyne, the simplest alkyne, but it is slower
under the same conditions:
C2H2 + 2Br2 → C2H2Br4
This type of bromination – by addition to the double
and triple covalent bonds – is a general reaction of
the alkenes and alkynes. It takes place in all organic
compounds containing the functional groups of these
homologous series, e.g.
C
C
+ Br2
Br
Br
C
C
and has found important uses in organic chemistry.
The decolorisation of a solution of bromine is the
classic standard test for unsaturation in an organic
compound. A substance being tested is treated with
bromine water or a solution of bromine in chloroform
or tetrachloromethane. Reactions are usually rapid,
requiring no catalyst, and the disappearance of the
red bromine colour is easily seen.
The reaction mechanism is an electrophilic
addition to the double or triple bond. A bromine
C
C
+ HBr
H
Br
C
C
Electrophilic addition of bromine to the benzene
nucleus in aromatic compounds does not take place.
Benzene itself reacts with bromine in the presence of
a halogen carrier, such as iron filings or AlCl3, which
assists polarisation of the electrophile. Substitution
occurs and bromobenzene is formed:
C6H6 + Br2 → C6H5Br + HBr
Electrophilic substitution occurs more readily in
compounds with an activating substituent group
present, such as in phenol and aniline:
C6H5OH + 3Br2 → C6H2(OH)Br3 + 3HBr
C6H5NH2 + 3Br2 → C6H2(NH2)Br3 + 3HBr
Alkyl bromides are very useful reactants in the
synthesis of organic compounds. They can be prepared
by the general methods for alkyl halides and also by
a special method where bromine is added to a
suspension of the silver salts of a carboxylic acid in
an organic solvent:
Br2 + RCOOAg → RBr + AgBr + CO2
In synthetic routes to required organic molecules they
can be used to produce intermediate molecules
containing the functional halogen atoms, which can
then easily be substituted by other groups.
The Hofmann reaction was discovered by the
German chemist A.W. von Hofmann in 1881, whereby
good yields of pure primary amines can be produced
by the reaction of alkaline bromine solution
(bromate(I)) on carboxylic acid amides:
RCONH2 + Br2 + 4KOH →
RNH2 + 2KBr + K2CO3 +2H2O
It is a method applicable to both aliphatic and aromatic
acid amides.
As with chlorine, ultraviolet light causes homolytic fission of the covalent bond in the bromine
molecule, generating highly reactive bromine radicals:
School Science Review, September 2003, 85(310)
81
What’s special about bromine?
Br–Br
UV
Br· + Br·
Consequently photochemical reactions occur with
bromine and organic compounds. The normally
unreactive alkanes produce mixtures of brominated
alkanes by substitution and benzene reacts by addition
to form benzene hexabromide. Similarly, free radical
reactions with alkenes can lead to substitution as well
as addition products.
Other uses
The bromide salts of potassium, sodium, calcium,
strontium, lithium and ammonium became useful in
medicine in the nineteenth century. The first mild
sedative prescriptions, in 1837, contained bromides.
This medicinal use continued into the twentieth
century until superseded by modern remedies.
Bromide ions are known to have a specific
tranquilizing effect on the central nervous system by
depressing the higher centres of the brain.
Silver bromide and the other silver halides, due
to their light sensitivity, became essential chemicals
in photography in the nineteenth century. Emulsions
of silver bromide were made by reacting silver nitrate
with KBr (usually), suspended in a suitable gelatinous
substance, to make photographic plates and paper. The
first ‘dry plate’ that could be mass produced was made
by R. L. Maddox in 1871, using silver bromide. The
British chemist and inventor Sir Joseph Swan patented
a bromide paper in 1879. This use of silver halides
increased from the early twentieth century with the
inventions of celluloid film and readily available roll
film for black-and-white photography, followed by
colour photography. Silver bromide continues to be
used today to manufacture photographic film and
paper.
Potassium bromate(v) is a powerful oxidising
agent in acid solution and is a useful reagent in
quantitative volumetric analysis (Vogel, 1953). An
accurate standard KBrO3 solution is easily made and
can be titrated with an acidic solution of the substance
to be determined by the oxidation-reduction reaction.
The KBrO3 is reduced to KBr in the titration:
BrO3– + 6H+ + 6e– → Br– + 3H2O
and the end-point, when all the reducing substance
has reacted, is detected with the appearance of free
bromine when a final drop of KBrO3 reacts with KBr:
5Br– + BrO3– + 6H+ → 3Br2 + 3H2O
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Bromine can be used in disinfectants and bleaches
and is preferable, for certain purposes, to the familiar
chlorate(I) (hypochlorite) products. Ordinary bromine
water – a saturated solution of bromine – contains
undissociated bromic(I) acid which has uses as a mild
oxidising agent:
Br2(aq) + H2O(l)
HBrO(aq) +H+(aq) + Br–(aq)
A solution of sodium bromate(I) is formed with caustic
soda solution:
2NaOH(aq) + Br2(l) →
NaOBr(aq) + NaBr(aq) + H2O(l)
This has been a minor use of bromine in the past but
now more is being used in bleach modifiers,
disinfectants for swimming pools and industrial water
treatment. There are a number of chlorine-free
swimming-pool disinfectants made from bromine,
some in the form of tablets to dissolve in the water.
The discovery in 1921 by the American engineer
and chemist Thomas Midgley Jr., that tetraethyl lead
was an effective antiknock agent for petrol, greatly
increased bromine consumption. Although the
tetraethyl lead additive solved one problem at that
time, it caused another due to the formation of lead
oxide in the internal combustion engine. This was
solved by also adding 1,2-dibromoethane to react and
form volatile lead bromide which could then escape
in exhaust gases. A large demand for bromine to make
1,2-dibromoethane was thus created. In the 1970s
however, environmental agencies became concerned
about high levels of airborne lead in urban areas. New
regulations were in place by the 1990s which made
unleaded petrol the standard, requiring no 1,2dibromoethane.
Bromine is used for the leaching of precious
metals. A lot more bromine is now needed to make
bromobutyl rubber for car tyres, brominated intermediate compounds for the pharmaceutical industries,
fire-retardant chemicals, plastics and dyes. It is also
used in Geiger counters to quench the electron
avalanche that occurs when nuclear radiation is
detected.
Manufacture
When elemental bromine began to be manufactured
on a large scale, the traditional sources in America
were the saline springs of west Pennsylvania and west
Virginia, where 125 000 lb were extracted in 1870. In
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Europe it was extracted from mother liquors from kelp
manufacture and bromine-rich brine waters and salt
deposits. The production from the Stassfurt salt
deposits was 10 000 lbs in 1873.
Large-scale manufacturing of 1,2-dibromoethane,
which began after 1928, greatly increased the demand
for bromine, causing a shortage. New commercial
processes to extract it from sea water were developed
and large quantities began to be extracted from the
Dead Sea, a virtually unlimited source.
In 1990 the total world production of bromine had
reached about 433 000 tonnes (Table 4). How it was
used in the USA, the largest producer, illustrates the
major applications in the 1990s (Figure 6). This
pattern has changed, mainly because 1,2-dibromoethane is no longer required as a petrol additive, but
Table 4 World bromine production in 1990
(Greenwood and Earnshaw, 1997).
Country
Tonnes/103
USA
Israel
Russia
UK
France
Japan
Total
177
135
60
28
18
14
433
What’s special about bromine?
the overall requirement for bromine has not fallen.
About 75% of world production is still from underground brines and lakes in the USA and from the Dead
Sea.
other
water
treatment
flame
retardants
17%
29%
5.5%
5.5%
11%
16%
16%
inorganic
bromides
drilling fluids
1,2-dibromoethane
agrochemicals
Figure 6 Applications for bromine in the USA in
1990 (Greenwood and Earnshaw, 1997).
Acknowledgement
I was pleased to be able to visit the Amlwch plant in 1997 on a very helpful and informative North West
Industry study tour for teachers sponsored by the Royal Society of Chemistry.
Sources
Encyclopaedia Britannica CD (2000) Geochemical
distribution of the elements: Terrestrial distribution: The
Earth’s crust, Figure 1. 1994–1999 Encyclopaedia
Britannica Inc. Oxford University Press.
Harrison, R. D. ed. (1972) Nuffield Advanced Science Book
of data, pp. 24–25. London: Longman.
Vogel, A. I. (1953) A textbook of quantitative inorganic
analysis. London: Longman.
Greenwood, N. N. and Earnshaw, A. (1997) Chemistry of the
elements. 2nd edn. Butterworth-Heinemann; Reed
Educational & Professional Publishing Ltd.
Pat Swain formerly taught chemistry and mathematics and is currently a GCSE examiner and an occasional
supply teacher. E-mail: [email protected]
School Science Review, September 2003, 85(310)
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