Bromine

• 1. Introduction

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  Bromine, element number 35, pours like water, yet it gives off a vapour with the same red-brown colour as the liquid itself. At room temperature only one other element, mercury, stays liquid. The vapour is penetrating and unpleasant to breathe, and the element reacts with almost everything it touches.

  Still, this difficult substance has calmed patients, shaped early photography, and slowed the spread of fire. It also sits in the sea in enormous quantity, and a few traces of it turn up inside living bodies.

  Who found it first, and how close did a famous chemist come to holding it without knowing? Why does it sit so neatly between two better-known elements? And what does an element this hostile actually do inside animals?

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• 2. A Race Nobody Knew About

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  In 1825, in his hometown of Bad Kreuznach, the chemist Carl Jacob Löwig took water from a mineral spring and saturated it with chlorine. He shook the mixture with diethyl ether. When the ether evaporated, a brown liquid remained. Löwig offered that liquid as a sample of his skill when he applied for a place in the Heidelberg laboratory of Leopold Gmelin. His publication was delayed.

  Meanwhile, in Montpellier, a young pharmacist named Antoine Jérôme Balard was studying the ash of seaweed gathered from the local salt marshes. The seaweed was already used to make iodine. Balard treated a solution of the ash with chlorine and distilled off a substance whose behaviour sat between that of chlorine and iodine. He suspected iodine monochloride and tried to prove it. He failed, and concluded that he held a new element. His results appeared in 1826, ahead of Löwig's.

  Justus von Liebig had the unluckiest part in the story. He had handled the same material in 1825, but he took it for iodine chloride and moved on.

  Balard first called his discovery muride, from the Latin word muria, meaning brine. Three established French chemists, Louis Nicolas Vauquelin, Louis Jacques Thénard, and Joseph-Louis Gay-Lussac, approved the work. It was then presented at a lecture of the Académie des Sciences and printed in the Annales de Chimie et Physique. In that paper Balard explained that he had changed the name to brôme on the suggestion of a certain M. Anglada. The word comes from the Greek brômos, which means stench. Some sources instead credit Gay-Lussac with choosing the name, inspired by the smell of the vapours.

  A name and a lecture did not make the element common. Bromine was not made in quantity until 1858, when salt deposits found at Stassfurt allowed it to be collected as a by-product of potash.

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• 3. Between Chlorine and Iodine

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  Johann Wolfgang Döbereiner grouped chlorine, bromine, and iodine into one of his original triads, an early step toward the periodic law. Bromine is the middle member of that set, and nearly everything about it lands in between its neighbours. It is the third halogen, in group 17, with seven electrons in its outer shell. One more would complete the shell, so it is a strong oxidising agent that grabs electrons from many other elements.

  The numbers show the pattern. On the electronegativity scale, fluorine scores 3.98, chlorine 3.16, bromine 2.96, and iodine 2.66. Bromine is less reactive than chlorine and more reactive than iodine. Its melting point, boiling point, density, and bond lengths follow the same middle path.

  Colour tells the same story. Fluorine is a very pale yellow gas, chlorine is greenish-yellow, and iodine is a shiny black solid. Bromine is a reddish-brown volatile liquid that freezes at -7.2 °C and boils at 58.8 °C. The darkening happens because the halogens absorb longer wavelengths of visible light as one moves down the group. Cool bromine far enough, to -195 °C, and the colour fades to pale yellow.

  Solid bromine forms layers of Br2 molecules in an orthorhombic crystal, and it conducts electricity very poorly. Natural bromine is a nearly even mix of two stable isotopes, 79Br at 51% and 81Br at 49%. That roughly 1:1 split leaves a recognisable signature in mass spectroscopy, so chemists can spot bromine-containing compounds quickly.

  Squeeze the element hard enough and its character changes. At 55 GPa, roughly 540,000 times atmospheric pressure, bromine turns from an insulator into a metal. At 100 GPa it settles into a body-centred orthorhombic structure made of single atoms.

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• 4. Brine and Bitter Seas

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  The Earth's crust holds only 2.5 parts per million of bromine, always locked in bromide salts. The oceans hold 65 parts per million, gathered over a very long time as water leached the rocks. That works out to about one bromine atom for every 660 chlorine atoms, and it makes bromine the tenth most abundant element in seawater.

  Some waters are far richer. The Dead Sea contains 0.4% bromide ions, and other salt lakes and brine wells are also concentrated. Almost all economical extraction starts from sources like these.

  The sea also makes its own bromine chemistry. Marine plants and animals build organobromine compounds, and more than 1600 of them had been identified by 1999. Bromide is scarce next to chloride in seawater, yet organobromides are the most common organohalides in nature. The reason is that bromide oxidises easily into a potent electrophile, and the enzyme bromoperoxidase speeds that step. Algae rely on a distinctive vanadium bromoperoxidase. The essential oil of the Hawaiian alga Asparagopsis taxiformis is 80% bromoform. Marine algae make an estimated 56,000 tonnes of methyl bromide each year, and the oceans are estimated to release 1-2 million tons of bromoform annually.

  Industry works with the brine instead. The main producers are Israel and Jordan. Chlorine gas oxidises bromide to bromine, a blast of steam or air carries the bromine out, and it is then condensed and purified. The whole bromine industry is about one-hundredth the size of the chlorine industry.

  The finished liquid travels in large metal drums or lead-lined tanks that hold hundreds of kilograms or even tonnes. Because commercial supplies are easy to buy and keep well on the shelf, laboratories have no need to repeat Balard's distillation themselves.

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• 5. Strange Bonds, Unstable Oxides

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  Hydrogen bromide, HBr, is the simplest bromine compound, and it is a colourless gas. Dissolved in water it becomes hydrobromic acid, a strong acid with a pKa of -9. Distillation cannot concentrate that acid beyond a boiling mixture at 124.3 °C containing 47.63 g of HBr per 100 g of solution.

  Nearly every element forms a binary bromide. The exceptions fall into three small groups: the extremely inert noble gases, the heaviest elements that decay before chemists can study them, and the elements more electronegative than bromine, such as oxygen and fluorine. Silver bromide barely dissolves in water, so it serves as a classic test for the presence of bromine. Another old test uses brown bromine water, which turns colourless when it meets an alkene.

  Bromine also combines with other halogens, and some of the results are dramatic. Bromine trifluoride is a straw-coloured liquid that reacts violently with water and explodes on contact with flammable materials. The nuclear power industry uses it to oxidise uranium to uranium hexafluoride. Bromine pentafluoride, first made in 1930, is another strong fluorinating agent that reacts violently with water.

  The oxides are the most elusive family. Bromine oxides are all fairly unstable, and it was once thought they could not exist at all. Dibromine monoxide, a dark-brown solid, is reasonably steady at -60 °C but decomposes at its melting point of -17.5 °C. Another so-called oxide, once called bromine dioxide, is better described as bromine perbromate, and it breaks down violently at 0 °C.

  Perbromates held out longest. Many attempts to make them failed, and some chemists built arguments for why they should not exist. In 1968 the perbromate anion finally appeared, formed from the radioactive beta decay of an unstable selenium compound. Today it is made by oxidising alkaline bromate solutions with fluorine gas.

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• 6. Fire, Film, and Fuel

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  In 1840, a small improvement in photography gave bromine its first commercial job. Daguerreotype makers had been using iodine vapour to create a light-sensitive layer of silver halide, and bromine turned out to have advantages. Silver bromide, alone or mixed with silver chloride and silver iodide, is still the light-sensitive part of photographic emulsions.

  Today most of bromine's output goes to fire prevention. More than half the bromine produced worldwide each year is used for this purpose, and brominated flame retardants are the largest commercial use. The trick is chemical. When the treated material burns, it releases hydrobromic acid. Highly reactive hydrogen, oxygen, and hydroxyl radicals react with it to form far less reactive bromine atoms, and the chain reaction that sustains the flame stalls.

  Manufacturers build the retardant into plastics in several ways. A small amount of brominated monomer such as vinyl bromide can join polyethylene, polyvinyl chloride, or polypropylene during polymerisation. Tetrabromobisphenol A can be added to polyesters and epoxy resins. The epoxies in printed circuit boards usually come from flame-retardant resins, marked by the letters FR in the names FR-4 and FR-2.

  Some gaseous bromine compounds make superior fire suppressants. Halon 1301, Halon 1211, and Halon 1011 are non-toxic and work especially well in enclosed spaces such as submarines, airplanes, and spacecraft. They are expensive, and they harm the ozone layer. They no longer appear in routine fire extinguishers, though aerospace and military systems still use them.

  Ozone damage is bromine's great environmental cost. Ultraviolet sunlight splits volatile organobromine compounds in the atmosphere into free bromine atoms, which destroy ozone.

  Gasoline once consumed a surprising share of the supply. Ethylene bromide was added to leaded fuel to scavenge lead as volatile lead bromide, and that use accounted for 77% of United States bromine consumption in 1966. Environmental regulations have cut it since the 1970s.

  The pesticide bromomethane was widely used to fumigate soil and houses. Before the Montreal Protocol, an estimated 35,000 tonnes a year went to control nematodes, fungi, weeds, and other soil-borne diseases. The protocol scheduled its phase-out by 2005, and sulfuryl fluoride, which contains neither chlorine nor bromine, now replaces it in housing fumigation.

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• 7. Soft Drinks, Pools, and Batteries

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  Brominated vegetable oil is a mixture of plant-derived triglycerides that carry bromine atoms, and it keeps citrus-flavoured soft drinks from separating while they are shipped. It is one of the quieter places bromine reaches ordinary life.

  Industry has other uses. Organobromine compounds go into high-density drilling fluids, and they give colour to dyes such as Tyrian purple and the indicator bromothymol blue. Water treatment uses bromine and some of its compounds as well.

  Swimming pools show how a choice can be driven by history. Bromine came into pool use in the United States during World War II, when a shortage of chlorine was expected. It disinfects in much the same way as chlorine. It is rarely chosen outdoors, because it costs more and lacks a stabiliser that shields it from sunlight. Indoors it makes sense. It works across a wider pH range and holds up better in a heated pool or hot tub.

  Cooling towers use bromine instead of chlorine to control bacteria, algae, fungi, and zebra mussels.

  One newer role is energy storage. Zinc-bromine batteries are hybrid flow batteries that provide backup power and storage for stationary systems, from household scale all the way up to industrial scale.

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• 8. Bromide in the Body

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  By 1864, during the American Civil War, a 25% solution of liquid bromine in 0.75 molar aqueous potassium bromide was widely used to treat gangrene. That was before the publications of Joseph Lister and Pasteur changed ideas about infection.

  A gentler form, the simple bromide ion, had a longer medical career. In the late 19th and early 20th centuries, potassium bromide and sodium bromide were major sedatives and anticonvulsants. Chloral hydrate replaced them gradually, and then the barbiturates did. Sodium bromide was removed from over-the-counter sedative products such as Bromo-Seltzer in 1975. Bromide salts are still used as antiepileptics in veterinary and human medicine, though the United States Food and Drug Administration does not approve bromide for any disease.

  Bromism explains why. Bromide leaves the body slowly, with an elimination half-life of 9 to 12 days, so it builds up. Normal daily intake is 2 to 8 milligrams, and 0.5 to 1 gram per day can lead to bromism. The old therapeutic dose was about 3 to 5 grams, which is why chronic poisoning was once so common. It chronically impairs the membranes of neurons and produces drowsiness, psychosis, seizures, and delirium. Death from it is rare.

  Bromine compounds also had a darker career. In the early years of the First World War, xylyl bromide was used as a poison gas.

  Still, the body does use bromine. A 2014 study suggested that bromide is a necessary cofactor in building collagen IV, which makes it important for the architecture of basement membranes and for tissue development in animals. No clear deprivation syndrome has been documented in mammals.

  Eosinophils, white blood cells that fight parasites, use hydrogen peroxide with an enzyme called eosinophil peroxidase. The enzyme prefers bromide over chloride and produces hypobromous acid. The mechanism helps kill multicellular parasites, including the nematode worms behind filariasis, and some bacteria such as those that cause tuberculosis.

  A stranger compound turns up in the cerebrospinal fluid of humans, cats, and rats. It is 2-octyl 4-bromo-3-oxobutanoate, and it appears to help induce REM sleep, although its exact role is unclarified. A different enzyme, neutrophil myeloperoxidase, can brominate deoxycytidine, which could result in DNA mutations.

  Elemental bromine itself remains dangerous. It burns human flesh, and inhaled gas causes coughing, choking, and shortness of breath. Large enough doses are fatal. The United States Occupational Safety and Health Administration sets its permissible exposure limit at a time-weighted average of 0.1 ppm, and air containing 3 ppm is considered immediately dangerous to life and health.

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