Bromine (Br): Properties & Uses


Bromine is the third lightest halogen positioned between chlorine and iodine in group 17 of the periodic table. The property that makes bromine exclusively special than other halogens is its existence as a red-brown liquid at room temperature. In fact, bromine is one of the only two elements that exist in a liquid state at room temperature, the other one being mercury. The red-brown bromine liquid is viscous and toxic in nature and has a suffocating odor. Like other halogen elements, bromine is also highly reactive, and therefore, does not exist in its elemental form on earth’s crust naturally. However, it is fairly abundant in the ocean beds in compound forms. Just like chlorine, bromine also forms mineral salts that are soluble in water and organic solvents. An approximate abundance of bromine compounds in the oceans is almost one-third of that of chlorine compounds found in oceans. On earth’s crust, it is the 44th most abundant element with an approximate density of 2.4 parts per million.


Discovery and Naming



Antonie Balard

Antoine Jérôme Balard in 1870s

The first person to officially announce the discovery of bromine was Antonie Balard. However, he was not the first one to discover it. In 1825, Carl Lowig, a research scholar at the University of Heidelberg, Zurich, prepared a liquid from salt-spring water by mixing it with diethyl-ether and passing chlorine through it. After distilling off the ether from the solution, he was left with a red-brown liquid, which he brought to his professor, Leopold Gmelin, in the laboratory. Gmelin asked him to prepare that solution in large quantities so that they can study its properties. In the meantime, another professor of chemistry at College De France, Antonie Balard, announced the discovery of Bromine in his paper “Annales de Chimie et Physique.” He was conducting experiments on seaweed ash to produce iodine, but to his astonishment, he also found bromine in the distilled solution of seaweed ash when he saturated it with chlorine.


Carl Jacob Loewig

The credit for the naming of bromine is also as complicated as its discovery. Some of the sources suggest that Antonie Ballard changed the name from “muride” to “brome,” in reference to a Greek word for the stench in one of his following papers on bromine, whereas others suggest that the French chemist and physicist Joseph-Louis Gay-Lusac proposed the name “brome” for its pungent odor. Moreover, the large-scale production of bromine was not done until 1858 when the discovery of a large repository of salts in Stassfurt enabled it.


Louis Joseph Gay Lussac

Isotopes of Bromine


Bromine exist in 30 isotopes ranging from _{}^{68}{Br} to _{}^{97}{Br}. It can be found dominantly existing in two of its naturally occurring stable isotopes, _{}^{79}{Br} and _{}^{81}{Br}. Nevertheless, _{}^{79}{Br} makes up 51%of naturally occurring bromine, whereas _{}^{81}{Br} constitutes the remaining 49%. Bromine also exists in 30 other radioactive isotopes with _{}^{77}{Br} having the longest half-life of 57.036 hours. Among other radioisotopes of bromine only _{}^{80}{Br},_{}^{80m}{Br}, and _{}^{81}{Br} have the significant half-lives of  17.7 min, 4.421 hours and 35.28 hours, respectively. These radioactive isotopes can be produced in a lab by the process of electron capture. While the radioisotopes with mass no. less than 79 prefer to decay into selenium by the process of electron capture, those with the mass no. greater than 81 prefer to decay into krypton by the process of beta decay.

Properties of Bromine

Bromine belongs to the 17th group of the periodic table, otherwise known as the halogen family. Like other members of the halogen family, such as fluorine, chlorine, iodine, and astatine, bromine is also one electron less than its optimal electronic configuration due to which it shows many resemblances in its properties with other elements of the 17th group.


Physical Properties


Bromine is an element in the periodic table with atomic no. 35 and atomic mass 79.904 u. It is a dark reddish-brown, fuming, highly corrosive, and viscous liquid with a strong pungent odor at room temperature. However, the red color of bromine varies as the temperature changes. At the temperature of 20 K, it is yellow-orange in color, which changes to black as the temperature reaches the melting point, i.e., 266 K, and it changes from a red-brown liquid at room temperature to orange-red fumes as the temperature reaches the boiling point (332 K). The density of bromine at room temperature is 3.1028 grams per cubic centimeters. Bromine is soluble in water and several organic compounds.

Chemical Properties

Atomic structure

Bromine has an electronic configuration of  [Ar]{3d}^{10}{4s}^{2}{4p}^{5}. Like other halogen elements, bromine either needs to gain 1 electron or lose 5 electrons to achieve a stable electronic configuration. Therefore, the most common oxidizing states of bromine is -1 and +5. However, the oxidation state of +3, +1, and +7 can also be observed in some cases. The electronegativity of bromine is less than that of chlorine and fluorine, therefore, it is moderately oxidizing in comparison to the latter. Conversely, it is a better reducing agent than chlorine. Let’s understand these properties with the help of the following reactions.

1. Reaction of Bromine with Water

Bromine reduces to bromide when treated with water. The reducing agent can be either the water or bromine itself. When water acts as a reducing agent, oxygen is formed.

{2Br}_{2}{+}{2H}_{2}{O}{4H}^{+}{+}{4Br}^{-} + {O}_{2}

Whereas, when bromine itself acts as a reducing agent, bromide or hypobromite is formed.

{Br}_{2}{+}{H}_{2}{O}{H}^{+}{+}{Br}^{-} + {HOBr}

2. Reaction of Bromine with Air

Bromine does not react with an oxygen molecule, {O}_{2}, and nitrogen molecule, {N}_{2}, in air. However, it does react with ozone, at -78 °C, forming bromine(IV) oxide, {BrO}_{2}

{Br}_{2}(l) + 2{O}_{3}(g) → 2{BrO}_{2}(s) + {O}_{2}(g)

Bromine reacts with carbon monoxide, CO, forming {COBr}_{2}.

{Br}_{2}(l) + {CO}(g) →{COBr}_{2}(l)

3. Reaction of Bromine with Hydrogen

Hydrogen reacts with {Br}_{2} forming hydrogen bromide. The reaction is slow at room temperature and and becomes fast with increasing temperatures.

{H}_{2}(g) + {Br}_{2}(g) → 2 {HBr}(g)

4. Reaction of Bromine with Halogens

Bromine, {Br}_{2}, reacts with fluorine, {F}_{2}, in the gas phase, forming BrF. The product is difficult to obtain pure as BrF reacts with itself, forming {Br}_{2}, {BrF}_{3}, and {BrF}_{5}.

{Br}_{2}(g) + {F}_{2}(g) → 2{BrF}(g)

3{BrF}(g) → {Br}_{2}(l) + {BrF}_{3}(l)

5{BrF}(g) → 2{Br}_{2}(l) + {BrF}_{5}(l)

Using excess fluorine at 150 °C, bromine will react with fluorine forming {BrF}_{5}.

{Br}_{2}(l) + 5{F}_{2}(g) → 2{BrF}_{5}(l)

Of the bromine fluoride compounds (BrF, BrF3, BrF5), BrF3 is of greatest practical importance because of its use in fluorinating organic substances.

Chlorine, {Cl}_{2}, reacts with bromine, {Br}_{2}, in gas phase, forming the unstable bromine(I) chloride, {ClBr}.

{Cl}_{2}(g) + {Br}_{2}(g) → 2{ClBr}(g)

Bromine, {Br}_{2}, reacts with iodine, {I}_{2}, at room temperature, forming bromine(I) iodide, {BrI}.

{Br}_{2}(l) + {I}_{2}(s) → 2{BrI}(s)

5. Reaction of Bromine with Organic Compounds

Bromine adds readily to unsaturated compounds. Such reactions are usually run at low temperature to avoid substitution side reactions. Although a catalyst is usually not required, ultraviolet radiation or high temperature may be used to accelerate the reaction. Electrophilic aromatic substitution reaction is by far the most important type of aromatic bromination. In the presence of a catalyst, bromine reacts with aromatic compounds to give aryl bromides and hydrogen bromide.


Bromine is efficiently used in aromatic substitution reactions to produce hydrochloric acid with the help of chlorine.


Bromination of saturated hydrocarbons and alkyl side chains of aromatic compounds occurs by a free-radical chain reaction.


Dissociation of bromine can be achieved thermally, photolytically, by gamma rays, and by the use of peroxide initiators.



6. Reaction of Bromine with Metals

Bromine reacts with many metals to form bromides. Sodium is stable in dry bromine, but sodium vapor reacts vigorously. Potassium and cesium react violently with bromine. Bromine is also highly reactive with aluminum and titanium. Aluminum reacts with the emission of light. Magnesium, silver, nickel, and lead become coated with their bromides, which prevents further reaction.

Mn(II)-ions are readily oxidized to MnO2 by brome under alkaline conditions

{Mn}^{2+}(aq) + {Br}_{2}(aq) + 2 {OH}_{-}(aq) → {MnO}_{2}(s) [brown-black] + 2{HBr}(aq)

Manganese with oxidation steps greater than 2 will be reduced to Mn(II) by {Br}^{-} under acidic conditions under the formation of {Br}_{2}, e.g.

{MnO}_{2}(s) + 2 {Br}^{-}(aq) + 4 {H}^{+}(aq){Mn}^{2+}(aq) + {Br}_{2}(aq) + 2 {H}_{2}{O}(l)

Uses of Bromine

1. Gasoline Additives


The largest industrial application of bromine from the early 1920s to late 1980s was in the fuel industry as 1,2-dibromoethane ({C}_{2}{H}_{4}{Br}_{2}). Ethylene Dibromide or EDB was used on large scale as a lead scavenger additive with tetralkyls to reduce a destructive phenomenon known as engine knocking. However, after the discovery of carcinogenic properties of EDB, its use as a gasoline additive was replaced by another less harmful alternative, methyl tert-butyl ether.

2. Flame Retardants


Flame retardants (FR) are added or applied to a material to increase the fire resistance of that product. Most of the daily life products, such as clothing, furniture, electronics, vehicles, and computers are petroleum-based polymeric materials, and hence, they are flammable. To meet fire safety regulations, Brominated Flame Retardants (BFRs) are commonly applied to these materials to increase their fire resistance. BFRs are divided into three subgroups depending on the mode of incorporation of these compounds into the polymers: brominated monomers, reactive and additive. A brominated monomer such as brominated styrene or brominated butadiene is added before the polymerization, whereas reactive flame retardants, such as tetrabromobisphenol A (TBBPA), and additive flame retardants, such as polybrominated diphenyl ethers (PBDEs) and hexabromocyclododecane (HBCDD) are simply blended during the process.

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3. Pesticides


1,2-dibromoethane ({C}_{2}{H}_{4}{Br}_{2}) was first used in 1952 to control pests in grain storage units. Its efficacy has made it employable as a pesticide against many soil pests, such as Nematodes, soil fungi, wild weed, parasitic plants, and several soil insects. However, the use of EBD as a pesticide went under subsequent decline in many countries due to the awareness for a clean environment.

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The need for a pesticide for preventing the degradation of crops and making the land arable was met by another bromine compound known as bromomethane (commonly known as methyl bromide, {CH}_{3}{Br}). Due to its low boiling point, it is an active gas even at relatively low temperatures that can be diffused into the soil to a considerable depth.

4. Used in Petroleum and Natural Gas Drilling Plant


Natural Gas Drilling Plant

Oil and gas are hydrocarbons that are found in porous stones. Heavy drilling machinery is used to extract these oils from underneath. The drilling requires a specific gravity needed to compensate for the pressure to avoid closing the pores whilst maintaining permeability. Because of their higher densities than freshwater, brines (saline liquids) are used when penetrating a pay zone (a reservoir or portion of a reservoir that contains economically producible hydrocarbons). Calcium bromide, sodium bromide, and zinc bromide are collectively referred to as clear brine fluids. They are used in the oil and gas well drilling industry for high-density solid-free completion, packer, and workover fluids to reduce the likelihood of damage to the wellbore and productive zone.

5. Photographic Chemical


Silver bromide (AgBr) is used in photography as a component of an emulsion that helps in the development of a photographic image. Silver bromide is sensitive to light, and when suspended in gelatin, silver bromide’s grains create a photographic emulsion. When exposed to light, silver bromide decomposes and as a result, it preserves a photographic image. After silver bromide creates a photographic image, the image needs to be developed. Grains of silver bromide, which have reacted to light, become metallic silver, whereas those unaffected by light do not change. These remaining grains are washed away in a fixing solution.

6. Dyes

Purple Dye

Ancient Romans in purple-dyed clothes

Perhaps the earliest mentioned use of bromine goes back to the biblical era when Romans used a reddish-purple natural dye named Tyrian Purple to dye their clothes. They used the sea-snails in the production of this dye. It was a very high valued dye, as it required thousands of snails and great labor to just produce a few grams of it; also, it does not fade easily. In fact, the expensive nature of this dye became a status symbol among the royalty of the twentieth century. It was later discovered that the chemical responsible for this royal shade was 6,6-dibromoindigo. Many bromine compounds are still used in the textile industry to produce elegant shades on clothes.

7. Pharmaceuticals


Although elemental bromine is toxic in nature, many over-the-counter life-saving drugs have bromine compounds as a major ingredient. Many organobromines are used as a sedative because bromide ions have the ability to decrease the sensitivity of the central nervous system. Bromine compounds are used in manufacturing sedatives, analgesics, and antihistamines medicines. They are also used in various drugs for treating pneumonia, cocaine addiction; moreover, they are also used as a catalyst to speed up the drug-reaction inside the body. Apart from these uses, bromine-based medicine is also supplied to those suffering from heart problems, thyroid hyperactivity, and hysteria. Some organic bromides are also used in the manufacturing of designer drugs that have the potential to cure many terminal diseases. For instance, 5-bromouracil is used as an artificial mutagen, which helps scientists modify genetic materials such as DNA.

Biological Aspects


From a physiological point of view, in its vapor state, bromine is very similar to chlorine. Due to its oxidizing action, inhalation of 10 ppm and higher concentrations of bromine may cause high irritation to the entire respiratory tract, the mucous membranes, and the eyes; producing symptoms such as coughing, nose bleed, feeling of oppression, dizziness, headache, and possibly delayed abdominal pain and diarrhea. Pneumonia can also be a late complication of severe exposure.


Liquid bromine produces a mild cooling sensation on the first contact with the skin, which is followed by a sensation of heat. It can leave severe blisters and burns on the skin if not removed immediately by flooding water on the area of contact. Contact with concentrated vapor can also cause burns and blisters. For treating tiny areas of contact in the laboratory, a 10% solution of sodium thiosulfate in water can neutralize bromine, and such a solution should be available when working with bromine.


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