Bromine

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Bromine,  35Br
Bromine 25ml (transparent).png
Bromine
Pronunciation /ˈbrmn,-mɪn,-mn/ (BROH-meen, -min, -myn)
Appearancereddish-brown
Standard atomic weight Ar, std(Br)[79.901, 79.907]conventional: 79.904
Bromine in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
Cl

Br

 I 
seleniumbrominekrypton
Atomic number (Z)35
Group group 17 (halogens)
Period period 4
Block p-block
Element category   Reactive nonmetal
Electron configuration [ Ar ] 3d10 4s2 4p5
Electrons per shell
2, 8, 18, 7
Physical properties
Phase at  STP liquid
Melting point (Br2) 265.8  K (−7.2 °C,19 °F)
Boiling point (Br2) 332.0 K(58.8 °C,137.8 °F)
Density (near r.t.)Br2, liquid: 3.1028 g/cm3
Triple point 265.90 K,5.8 kPa [1]
Critical point 588 K, 10.34 MPa [1]
Heat of fusion (Br2) 10.571  kJ/mol
Heat of vaporisation (Br2) 29.96 kJ/mol
Molar heat capacity (Br2) 75.69 J/(mol·K)
Vapour pressure
P (Pa)1101001 k10 k100 k
at T (K)185201220244276332
Atomic properties
Oxidation states −1, +1, +3, +4, +5, +7 (a strongly acidic oxide)
Electronegativity Pauling scale: 2.96
Ionisation energies
  • 1st: 1139.9 kJ/mol
  • 2nd: 2103 kJ/mol
  • 3rd: 3470 kJ/mol
Atomic radius empirical:120  pm
Covalent radius 120±3 pm
Van der Waals radius 185 pm
Color lines in a spectral range Bromine spectrum visible.png
Color lines in a spectral range
Spectral lines of bromine
Other properties
Natural occurrence primordial
Crystal structure orthorhombic
Orthorhombic.svg
Speed of sound 206 m/s (at 20 °C)
Thermal conductivity 0.122 W/(m·K)
Electrical resistivity 7.8×1010 Ω·m(at 20 °C)
Magnetic ordering diamagnetic [2]
Magnetic susceptibility 56.4·10−6 cm3/mol [3]
CAS Number 7726-95-6
History
Discovery and first isolation Antoine Jérôme Balard and Carl Jacob Löwig (1825)
Main isotopes of bromine
Iso­tope Abun­dance Half-life (t1/2) Decay mode Pro­duct
79Br51% stable
81Br49%stable
| references
Bromine 3D molecule, Br2 Bromine-3D-Balls.png
Bromine 3D molecule, Br2

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 (in 1826), its name was derived from the Ancient Greek βρῶμος ("stench"), referencing its sharp and disagreeable smell.

Chemical element a species of atoms having the same number of protons in the atomic nucleus

A chemical element is a species of atom having the same number of protons in their atomic nuclei. For example, the atomic number of oxygen is 8, so the element oxygen consists of all atoms which have 8 protons.

Atomic number number of protons found in the nucleus of an atom

The atomic number or proton number of a chemical element is the number of protons found in the nucleus of every atom of that element. The atomic number uniquely identifies a chemical element. It is identical to the charge number of the nucleus. In an uncharged atom, the atomic number is also equal to the number of electrons.

Halogen group of elements that tend to be very reactive and form salts with metals and form acids when bonded to hydrogen, consisting of fluorine, chlorine, bromine, iodine, astatine, and possibly tennessine

The halogens are a group in the periodic table consisting of five chemically related elements: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), and Astatine (At). The artificially created element 117 may also be a halogen. In the modern IUPAC nomenclature, this group is known as group 17.

Contents

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, Israel and China. The mass of bromine in the oceans is about one three-hundredth that of chlorine.

In chemistry, a free element is a chemical element that is not combined with or chemically bonded to other elements. Examples of elements which can occur as free elements include the oxygen molecule (O2) and carbon. All atoms of free elements have an oxidation number of 0. They hardly ever bond with other atoms. Other examples of free elements include the noble metals gold and platinum.

Brine pool An area of high density brine collected in a depression on the ocean floor

A brine pool is a large area of brine on the ocean basin. These pools are bodies of water that have a salinity three to eight times greater than the surrounding ocean. For deep-sea brine pools, the source of the salt is the dissolution of large salt deposits through salt tectonics. The brine often contains high concentrations of methane, providing energy to chemosynthetic animals that live near the pool. These creatures are often extremophiles. Brine pools are also known to exist on the Antarctic Shelf where the source of brine is salt excluded during formation of sea ice. Deep-sea and Antarctic brine pools can be toxic to marine animals.

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 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.

Organobromine compounds, also called organobromides, are organic compounds that contain carbon bonded to bromine. The most pervasive is the naturally produced bromomethane.

A chain reaction is a sequence of reactions where a reactive product or by-product causes additional reactions to take place. In a chain reaction, positive feedback leads to a self-amplifying chain of events.

A fire retardant is a substance that is used to slow or stop the spread of fire or reduce its intensity. This is commonly accomplished by chemical reactions that reduce the flammability of fuels or delay their combustion. Fire retardants may also cool the fuel through physical action or endothermic chemical reactions. Fire retardants are available as powder, to be mixed with water, as fire-fighting foams and fire-retardant gels. Fire retardants are also available as coatings or sprays to be applied to an object.

Large amounts of bromide salts are toxic from the action of soluble bromide ion, causing bromism. However, a clear biological role for bromide ion and hypobromous acid has recently been elucidated, and it now appears that bromine is an essential trace element in humans. The role of biological organobromine compounds in sea life such as algae has been known for much longer. 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.

Bromism is the syndrome which results from the long-term consumption of bromine, usually through bromide-based sedatives such as potassium bromide and lithium bromide. Bromism was once a very common disorder, being responsible for 5 to 10% of psychiatric hospital admissions, but is now uncommon since bromide was withdrawn from clinical use in many countries and was severely restricted in others.

Salt (chemistry) Ionic compound consisting of cations and anions

In chemistry, a salt is a solid chemical compound consisting of an ionic assembly of cations and anions. Salts are composed of related numbers of cations and anions so that the product is electrically neutral. These component ions can be inorganic, such as chloride (Cl), or organic, such as acetate ; and can be monatomic, such as fluoride (F), or polyatomic, such as sulfate.

History

Antoine Balard, one of the discoverers of bromine Antoine Jerome Balard 1870s.jpg
Antoine Balard, one of the discoverers of bromine

Bromine was discovered independently by two chemists, Carl Jacob Löwig [4] and Antoine Balard, [5] [6] in 1825 and 1826, respectively. [7]

Carl Jacob Löwig German chemist

Carl Jacob Löwig was a German chemist and discovered bromine independently of Antoine Jérôme Balard.

Antoine Jérôme Balard French chemist

Antoine Jérôme Balard (1802–1876) was a French chemist and one of the discoverers of bromine.

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. [8]

Bad Kreuznach Place in Rhineland-Palatinate, Germany

Bad Kreuznach is a town in the Bad Kreuznach district in Rhineland-Palatinate, Germany. It is a spa town most well known for its medieval bridge dating from around 1300, the Alte Nahebrücke, which is one of the few remaining bridges in the world with buildings on it.

Diethyl ether chemical compound

Diethyl ether, or simply ether, is an organic compound in the ether class with the formula (C
2
H
5
)
2
O
, sometimes abbreviated as Et
2
O
. It is a colorless, highly volatile flammable liquid. It is commonly used as a solvent in laboratories and as a starting fluid for some engines. It was formerly used as a general anesthetic, until non-flammable drugs were developed, such as halothane. It has been used as a recreational drug to cause intoxication. It is an isomer of butanol.

Leopold Gmelin German chemist

Leopold Gmelin was a German chemist. Gmelin was professor at the University of Heidelberg among other things, he worked on the red prussiate and created Gmelin's test.

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 muria for brine. [6]

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. [5] 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). [5] [9] Other sources claim that the French chemist and physicist Joseph-Louis Gay-Lussac suggested the name brôme for the characteristic smell of the vapors. [10] [11] 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. [12]

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. [13]

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. [14] In the early years of the First World War, bromine compounds such as xylyl bromide were used as poison gas. [15]

Properties

Illustrative and secure bromine sample for teaching Bromine vial in acrylic cube.jpg
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. [16] 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. [16] 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. [17] [18] 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. [16] The volatility of bromine accentuates its very penetrating, choking, and unpleasant odour. [19]

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. [16] 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. [16] 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. [20] The colour fades at low temperatures, so that solid bromine at −195 °C is pale yellow. [16]

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. [16]

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. [21]

Isotopes

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 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 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. [16] 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 80Kr. [22]

Chemistry and compounds

Halogen bond energies (kJ/mol) [20]
XXXHXBX3AlX3CX4
F159574645582456
Cl243428444427327
Br193363368360272
I151294272285239

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 M–Br bonds. [20]

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: [23]

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. [23] 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 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. [24]

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. [24]

Other binary bromides

Silver bromide (AgBr) Bromid stribrny.PNG
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, 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.) [25]

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. [25] Another method is halogen exchange in the presence of excess "halogenating reagent", for example: [25]

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: [25]

3 WBr5 + Al thermal gradient475°C → 240°C 3 WBr4 + AlBr3
EuBr3 + 1/2 H2 → EuBr2 + HBr
2 TaBr4500°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 above. Silver bromide is very insoluble in water and is thus often used as a qualitative test for bromine. [25]

Bromine halides

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 BrF
2
, BrCl
2
, BrF+
2
, BrF+
4
, and BrF+
6
. Apart from these, some pseudohalides are also known, such as cyanogen bromide (BrCN), bromine thiocyanate (BrSCN), and bromine azide (BrN3). [26]

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. [26] 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. [25] 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+. [25]

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 reactive. Bromine trifluoride is a useful nonaqueous ionising solvent, since it readily dissociates to form BrF+
2
and BrF
4
and thus conducts electricity. [27]

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. [28]

Polybromine compounds

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
. [29] The tribromide anion, Br
3
, has also been characterised; it is analogous to triiodide. [26]

Bromine oxides and oxoacids

Standard reduction potentials for aqueous Br species [30]
E°(couple)a(H+) = 1
(acid)
E°(couple)a(OH) = 1
(base)
Br2/Br+1.052Br2/Br+1.065
HOBr/Br+1.341BrO/Br+0.760
BrO
3
/Br
+1.399BrO
3
/Br
+0.584
HOBr/Br2+1.604BrO/Br2+0.455
BrO
3
/Br2
+1.478BrO
3
/Br2
+0.485
BrO
3
/HOBr
+1.447BrO
3
/BrO
+0.492
BrO
4
/BrO
3
+1.853BrO
4
/BrO
3
+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 [31] 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. [32]

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. [32]

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 reactions occur: [30]

Br2 + H2O HOBr + H+ + BrKac = 7.2 × 10−9 mol2 l−2
Br2 + 2 OH OBr + H2O + BrKalk = 2 × 108 mol−1 l

Hypobromous acid is unstable to disproportionation. The hypobromite ions thus formed disproportionate readily to give bromide and bromate: [30]

3 BrO 2 Br + BrO
3
K = 1015

Bromous acids and bromites are very unstable, although the strontium and barium bromites are known. [33] 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: [33]

BrO
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 83
SeO2−
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 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. [34]

Organobromine compounds

Structure of N-bromosuccinimide, a common brominating reagent in organic chemistry N-Bromosuccinimide.svg
Structure of N-bromosuccinimide, a common brominating reagent in organic chemistry

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. [35]

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. [35]

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. [36] The oceans are estimated to release 1–2 million tons of bromoform and 56,000 tons of bromomethane annually. [37]

Bromine addition to alkene reaction mechanism Alkene-bromine-addition-2D-skeletal.png
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. [38]

Occurrence and production

View of salt evaporation pans on the Dead Sea, where Jordan (right) and Israel (left) produce salt and bromine 31deg9'0''N 35deg27'0''E / 31.15000degN 35.45000degE / 31.15000; 35.45000 STS028-96-65.jpg
View of salt evaporation pans on the Dead Sea, where Jordan (right) and Israel (left) produce salt and bromine 31°9′0″N35°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. [39] It is from these sources that bromine extraction is mostly economically feasible. [40] [41] [42]

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. [43]

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. [44]

Flame retardants

Tetrabromobisphenol A Tetrabromobisphenol A.svg
Tetrabromobisphenol A

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 characterise combustion. [45] [46]

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. [47]

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, CBrF3). [48]

Other uses

Baltimore's Emerson Bromo-Seltzer Tower, originally part of the headquarters of Emerson Drug Company, which made Bromo-Seltzer Bromo-Seltzer Tower MD2.jpg
Baltimore's Emerson Bromo-Seltzer Tower, originally part of the headquarters of Emerson Drug Company, which made Bromo-Seltzer

Silver bromide is used, either alone or in combination with silver chloride and silver iodide, as the light sensitive constituent of photographic emulsions. [43]

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). [49]

Poisonous bromomethane was widely used as pesticide to fumigate soil and to fumigate housing, by the tenting method. Ethylene bromide was similarly used. [50] 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. [51] [52]

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 it was removed from over-the-counter sedative products like Bromo-Seltzer, in 1975. [53] 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. [35]

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). [43] 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

Bromine
Hazards
GHS pictograms GHS-pictogram-acid.svg GHS-pictogram-skull.svg GHS-pictogram-pollu.svg
GHS Signal word Danger
H314, H330, H400
P260, P273, P280, P284, P305+351+338, P310 [54]
NFPA 704
Flammability code 0: Will not burn. E.g. waterHealth code 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasReactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no codeBromine
0
3
0
2-Octyl 4-bromo-3-oxobutanoate, an organobromine compound found in mammalian cerebrospinal fluid 2-octyl 4-bromo-3-oxobutanoate.png
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. [56] Nevertheless, no clear deprivation symptoms or syndromes have been documented. [57] 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 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. [58]

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. [37] Neutrophil myeloperoxidase can use H2O2 and Br- to brominate deoxycytidine, which could result in DNA mutations. [59] 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. [37] The essential oil of the Hawaiian alga Asparagopsis taxiformis consists of 80% bromoform. [60] Most of such organobromine compounds in the sea are made by the action of a unique algal enzyme, vanadium bromoperoxidase. [61]

The bromide anion is not very toxic: a normal daily intake is 2 to 8 milligrams. [57] 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. [62] Bromism is caused by a neurotoxic effect on the brain which results in somnolence, psychosis, seizures and delirium. [63]

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. [64] Caution is required when transporting bromine; it is commonly carried in steel tanks lined with lead, supported by strong metal frames. [43] 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. [65] 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. [66]

Related Research Articles

Chlorine Chemical element with atomic number 17

Chlorine is a chemical element with the symbol Cl and atomic number 17. The second-lightest of the halogens, it appears between fluorine and bromine in the periodic table and its properties are mostly intermediate between them. Chlorine is a yellow-green gas at room temperature. It is an extremely reactive element and a strong oxidising agent: among the elements, it has the highest electron affinity and the third-highest electronegativity on the Pauling scale, behind only oxygen and fluorine.

Iodine Chemical element with atomic number 53

Iodine is a chemical element with the symbol I and atomic number 53. The heaviest of the stable halogens, it exists as a lustrous, purple-black non-metallic solid at standard conditions that melts to form a deep violet liquid at 114 degrees Celsius, and boils to a violet gas at 184 degrees Celsius. The element was discovered by the French chemist Bernard Courtois in 1811. It was named two years later by Joseph Louis Gay-Lussac from this property, after the Greek ἰώδης "violet-coloured".

Haloalkane

The haloalkanes are a group of chemical compounds derived from alkanes containing one or more halogens. They are a subset of the general class of halocarbons, although the distinction is not often made. Haloalkanes are widely used commercially and, consequently, are known under many chemical and commercial names. They are used as flame retardants, fire extinguishants, refrigerants, propellants, solvents, and pharmaceuticals. Subsequent to the widespread use in commerce, many halocarbons have also been shown to be serious pollutants and toxins. For example, the chlorofluorocarbons have been shown to lead to ozone depletion. Methyl bromide is a controversial fumigant. Only haloalkanes which contain chlorine, bromine, and iodine are a threat to the ozone layer, but fluorinated volatile haloalkanes in theory may have activity as greenhouse gases. Methyl iodide, a naturally occurring substance, however, does not have ozone-depleting properties and the United States Environmental Protection Agency has designated the compound a non-ozone layer depleter. For more information, see Halomethane. Haloalkane or alkyl halides are the compounds which have the general formula "RX" where R is an alkyl or substituted alkyl group and X is a halogen.

Nonmetal term to classify chemical elements

In chemistry, a nonmetal is a chemical element that mostly lacks the characteristics of a metal. Physically, a nonmetal tends to have a relatively low melting point, boiling point, and density. A nonmetal is typically brittle when solid and usually has poor thermal conductivity and electrical conductivity. Chemically, nonmetals tend to have relatively high ionization energy, electron affinity, and electronegativity. They gain or share electrons when they react with other elements and chemical compounds. Seventeen elements are generally classified as nonmetals: most are gases ; one is a liquid (bromine); and a few are solids. Metalloids such as boron, silicon, and germanium are sometimes counted as nonmetals.

Oxidizing agent Chemical compound used to oxidize another substance in a chemical reaction

In chemistry, an oxidizing agent is a substance that has the ability to oxidize other substances — in other words to accept their electrons. Common oxidizing agents are oxygen, hydrogen peroxide and the halogens.

Hydrobromic acid is a strong acid formed by dissolving the diatomic molecule hydrogen bromide (HBr) in water. "Constant boiling" hydrobromic acid is an aqueous solution that distills at 124.3 °C and contains 47.6% HBr by mass, which is 8.89 mol/L. Hydrobromic acid has a pKa of −9, making it a stronger acid than hydrochloric acid, but not as strong as hydroiodic acid. Hydrobromic acid is one of the strongest mineral acids known.

Halogenation is a chemical reaction that involves the addition of one or more halogens to a compound or material. The pathway and stoichiometry of halogenation depends on the structural features and functional groups of the organic substrate, as well as on the specific halogen. Inorganic compounds such as metals also undergo halogenation.

Hydrogen bromide chemical compound

Hydrogen bromide is the diatomic molecule with the formula HBr. It is a colorless compound and a hydrogen halide. Hydrobromic acid is a solution of HBr in water. Both the anhydrous and aqueous solutions of HBr are common reagents in the preparation of bromide compounds.

IHC-Inter Halogen Compounds.An interhalogen compound is a molecule which contains two or more different halogen atoms and no atoms of elements from any other group.

Hydrogen halide diatomic inorganic compound with the formula HX where X is one of the halogens: fluorine, chlorine, bromine, iodine, or astatine

Hydrogen halides are diatomic inorganic compounds with the formula HX where X is one of the halogens: fluorine, chlorine, bromine, iodine, or astatine. Hydrogen halides are gases that dissolve in water to give acids which are commonly known as hydrohalic acids.

Bromine pentafluoride chemical compound

Bromine pentafluoride, BrF5, is an interhalogen compound and a fluoride of bromine. It is a strong fluorination reagent.

Sodium bromate chemical compound

Sodium bromate, the inorganic compound with the chemical formula of NaBrO3, is the sodium salt of bromic acid. It is a strong oxidant.

Bromic acid chemical compound

Bromic acid, also known as hydrogen bromate, is an oxoacid with the molecular formula HBrO3. It only exists in aqueous solution. It is a colorless solution that turns yellow at room temperature as it decomposes to bromine. Bromic acid and bromates are powerful oxidizing agents and are common ingredients in Belousov-Zhabotinsky reactions. Belousov-Zhabotinsky reactions are a classic example of non-equilibrium thermodynamics.

Vanadium bromoperoxidase

Vanadium bromoperoxidases are a kind of enzymes called haloperoxidases. Its primary function is to remove hydrogen peroxide from in or around the cell which is produced during photosynthesis. By producing hypobromous acid (HOBr) a secondary reaction with dissolved organic matter, what results is the bromination of organic compounds that are associated with the defense of the organism. These enzymes produce the bulk of natural organobromine compounds in the world.

Hypobromite anion

The hypobromite ion, also called alkaline bromine water, is BrO. Bromine is in the +1 oxidation state. Hypobromite is the bromine compound analogous to hypochlorites found in common bleaches, and in immune cells. In many ways, hypobromite functions in the same manner as hypochlorite, and is also used as a germicide and antiparasitic in both industrial applications, and in the immune system.

Beryllium iodide chemical compound

Beryllium iodide is the chemical compound with the formula BeI2. It is very hygroscopic and reacts violently with water, forming hydroiodic acid.

Haloform reaction chemical reaction

The haloform reaction is a chemical reaction where a haloform (CHX3, where X is a halogen) is produced by the exhaustive halogenation of a methyl ketone (RCOCH3, where R can be either a hydrogen atom, an alkyl or an aryl group), in the presence of a base. The reaction can be used to transform acetyl groups into carboxyl groups or to produce chloroform (CHCl3), bromoform (CHBr3), or iodoform (CHI3).

Fluorine forms a great variety of chemical compounds, within which it always adopts an oxidation state of −1. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. Most frequently, covalent bonds involving fluorine atoms are single bonds, although at least two examples of a higher order bond exist. Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding. Fluorine's chemistry includes inorganic compounds formed with hydrogen, metals, nonmetals, and even noble gases; as well as a diverse set of organic compounds. For many elements the highest known oxidation state can be achieved in a fluoride. For some elements this is achieved exclusively in a fluoride, for others exclusively in an oxide; and for still others the highest oxidation states of oxides and fluorides are always equal.

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