Bromine (Br) is a chemical element (atomic number 35) in group 17 (the halogens) of the periodic table. It is a dark reddish-brown, volatile liquid at standard conditions – in fact it is the only nonmetal element that is liquid at room temperature (the only other element liquid under these conditions is mercury) Bromine’s properties are intermediate between those of chlorine and iodine (its neighboring halogens). As a halogen it has seven valence electrons (electron configuration [Ar] 3d¹⁰4s²4p⁵) so it commonly forms the bromide ion (Br⁻, oxidation state –1) but can also exist in positive oxidation states (+1, +3, +5, +7) in highly oxidizing compounds. At room temperature bromine appears as a dense liquid (density ≈ 3.12 g/cm³) with a pungent, disagreeable odor. It freezes to a red-black crystalline solid at –7.2 °C and boils at 58.8 °C Because elemental bromine is very reactive, it never occurs freely in nature; it is typically obtained from bromide salts in brines. Bromine is used especially in organobromine flame retardants (over half of all bromine produced each year is used to make fire-retardant materials) as well as in drilling fluids, chemical synthesis, photography (silver bromide), and other applications.

Atomic Structure and Electron Configuration

Bromine
Atomic number	35
Symbol	Br
Group	17 (halogens)
Boiling point	58.8 °C
Electronegativity	2.96 (Pauling)
Electron configuration	[Ar] 3d10 4s2 4p5
Melting point	−7.2 °C
Main isotopes	79Br, 81Br
Block	p
Phase STP	Liquid
Oxidation states	−1, +1, +5
Wikidata	Q879

Bromine has 35 protons and, in its neutral atom form, 35 electrons. Its full electron configuration is [Ar] 3d¹⁰4s²4p⁵ The outermost shell (n=4) has seven electrons (4s²4p⁵), so bromine is one electron short of a filled shell. This configuration explains its chemistry: bromine readily accepts one electron to become Br⁻ (like other halides) or can share electrons in covalent bonds. In the periodic table, bromine sits in the fourth period and belongs to the p-block of nonmetals.

In periodic trends, bromine’s properties lie between those of chlorine (above it) and iodine (below it). For example, bromine’s atomic (covalent) radius is about 120 pm – larger than chlorine’s (≈100 pm) but smaller than iodine’s (≈133 pm). Its electronegativity (Pauling scale) is about 2.96 lower than chlorine (~3.16) and higher than iodine (~2.66). Correspondingly, the first ionization energy of bromine is roughly 1139 kJ/mol (Cl: ≈1250 kJ/mol; I: ≈1008 kJ/mol). These intermediate values reflect bromine’s intermediate reactivity among the halogens: it is a strong but not extreme oxidizing agent.

Isotopes and Nuclear Properties

Bromine has two stable isotopes: ⁷⁹Br and ⁸¹Br, present in almost equal abundance (about 51% ⁷⁹Br and 49% ⁸¹Br in nature) Both have nuclear spin 3/2⁻. The roughly 1:1 ratio of these two isotopes leads to the characteristic pattern of two nearly equal peaks, two mass units apart, in mass spectrometry of brominated compounds. Their nuclear quadrupole moments mean bromine is not ideal for high-resolution NMR, but such NMR is nevertheless possible with special techniques.

All other bromine isotopes are radioactive and very short-lived, so they are not found naturally. The most commonly studied radioisotopes are ⁸²Br (half-life ~35 hours), ⁸⁰Br (17.7 minutes), and ⁷⁷Br (57 hours) These decay by beta emission or electron capture to selenium or krypton. For example, ⁸²Br decays to ⁸²Kr by beta emission. These radioisotopes have niche uses in medical or tracer research but no major practical applications and occur only in laboratory-produced samples. No long-lived fissionable isotopes of bromine exist, so bromine has no direct role in nuclear power or dating.

Allotropes and Typical Compounds

Elemental bromine exists only as diatomic molecules (Br₂) – there are no polymeric or crystalline allotropes beyond ordinary solid bromine. Solid Bromine below –7.2 °C forms dark red-brown crystals with an orthorhombic structure It vaporizes easily, giving an orange-red gas without changing chemical composition.

Bromine forms many compounds, akin to those of chlorine. It typically makes single bonds to other atoms (no multiple bonding except as part of ions). Salient classes of bromine compounds include:

Hydrogen bromide (HBr) – a gas that dissolves in water to form hydrobromic acid, a very strong acid.
Bromide salts (Br⁻) – Ionic salts analogous to chlorides. Examples include sodium bromide (NaBr, used in photography) and potassium bromide (KBr). Silver bromide (AgBr) is light-sensitive and was widely used in photographic film.
Bromine oxoanions – such as hypobromite (BrO⁻), bromite (BrO₂⁻), bromate (BrO₃⁻), and perbromate (BrO₄⁻). For instance, potassium bromate (KBrO₃) is a strong oxidizer (occasionally used in baking, though it is carcinogenic). Hypobromous acid (HOBr) is a weak acid and disinfectant (used in swimming pools). Higher oxyacids (like HBrO₂, HBrO₄) are unstable or only known as salts.
Interhalogens – compounds with bromine bonded to other halogens. Examples include bromine monochloride (BrCl, a yellow gas), bromine trifluoride (BrF₃, a powerful fluorinating agent), bromine pentafluoride (BrF₅), and iodine monobromide (IBr).
Organobromine compounds – any molecules with C–Br bonds. These are widely used in industry and nature. Important examples: methyl bromide (CH₃Br, formerly a major fumigant, now largely phased out for environmental reasons), bromoform (CHBr₃, a solvent), brominated phenols and aromatics in flame retardants (e.g., decabromodiphenyl ether), and pharmaceuticals (many drugs feature bromine). In nature, hundreds of organobromine compounds are produced by marine organisms (like algae), some with biological functions

These compounds illustrate bromine’s chemistry: as a halogen it readily forms bromides (Br⁻) in ionic compounds and shares electrons in covalent bonds with nonmetals. Unlike chlorine, bromine is less likely to form stable organic radicals or diatomic chains; it stays as Br₂ molecules or becomes Br⁻ in salts.

 Bromine is moderately soluble in water, forming a reddish-brown solution known as bromine water. This solution is often used as a chemical test: many unsaturated organic compounds (like alkenes) react with and decolorize bromine water, indicating a double bond. Bromine is much more soluble in organic solvents (e.g. carbon disulfide), where it forms deep red solutions. In solution it can participate in similar chemistry as the gas or liquid form (e.g. oxidizing or brominating organic substrates).

Physical Properties

Bromine’s most striking physical feature is its color and phase. At room temperature it is a dense (3.12 g/cm³) red-brown liquid with a sharp smell. The liquid has moderate viscosity and evaporates quickly on standing. When frozen (below –7.2 °C) it forms a dark red crystalline solid (orthorhombic Br₂ molecules) The vapor above liquid bromine is orange-red. Bromine boils at 58.8 °C and at higher temperatures dissociates to atomic bromine under strong heating.

Thermally, liquid bromine has a low thermal conductivity (about 0.12 W/m·K) and a specific heat of ~0.47 J/(g·K) typical of heavy molecular liquids. It is a poor electrical conductor (nonmetal); like chlorine gas, bromine gas and liquid are diamagnetic (all electrons paired). Spectroscopically, bromine absorbs in the ultraviolet and visible range: for example, Br₂ shows broad absorption bands in the visible that give it its distinctive color. In flame tests, bromide salts impart a red color (similar to carmine) to a flame; free bromine vapor will extinguish a flame if enough is present.

Key physical values: solid Br (–78 °C) density ~3.41 g/cm³; liquid at 20 °C density 3.12 g/cm³; melting point –7.2 °C; boiling point 58.8 °C; standard atomic weight 79.904; molar mass ~79.9 g/mol

Chemical Reactivity and Trends

Bromine is chemically very reactive. It is a fairly strong oxidizing agent – capable of accepting electrons – but not as powerful as chlorine. This means bromine will oxidize iodide ions (I⁻) to iodine, but bromide ions (Br⁻) can be oxidized by chlorine (Cl₂ + 2Br⁻ → Br₂ + 2Cl⁻). In aqueous solution, bromine undergoes a rapid disproportionation: ½ Br₂ + H₂O → HOBr + HBr (equivalently Br₂ + 2H₂O → BrO⁻ + Br⁻ + 2H⁺). This reaction produces hypobromous acid (HOBr), a mild oxidant.

Some typical behaviors and trends:

Reactivity with hydrogen and metals: Bromine reacts directly with hydrogen gas to form hydrogen bromide (H₂ + Br₂ → 2 HBr). It also reacts exothermically with many metals. Alkali metals and alkaline earth metals give ionic bromides (e.g., 2 Na + Br₂ → 2 NaBr). Transition metals form metal-bromine complexes (e.g., FeBr₃, CuBr₂, AlBr₃) analogous to chlorides.
Halogen displacement: In the halogen series F₂ > Cl₂ > Br₂ > I₂ in oxidizing power. Thus fluorine can oxidize bromide to bromine, chlorine can oxidize bromide, bromine can oxidize iodide, and bromine will not oxidize chloride.
Acid/Base behavior: Hydrobromic acid (aqueous HBr) is a very strong acid (strength similar to HCl); it dissociates completely to H⁺ and Br⁻. Hypobromous acid (HOBr) is a weak acid (pKa ~8.7) and a mild disinfectant. Bromine with alkali hydroxide reacts to form bromide and bromate: in the cold it produces hypobromite (Br⁻ + HOBr), while hot or concentrated conditions typically yield bromide plus bromate (Br₂ + 6 OH⁻ → 5 Br⁻ + BrO₃⁻ + 3 H₂O).
Organic reactivity: Bromine adds readily to carbon–carbon double bonds, yielding vicinal dibromides (R–CH=CH–R + Br₂ → R–CHBr–CHBr–R). Aromatic rings can be brominated (replacing an H) in the presence of catalyst (e.g. FeBr₃). Because of this reactivity, a drop of bromine water in a hydrocarbon solution quickly loses its color if unsaturated compounds are present.
Complex formation: Bromide (Br⁻) can form complexes with certain metals (for example, excess bromide converts AgBr to the soluble [AgBr₂]⁻ complex).
Corrosion and passivity: Elemental bromine and hydrobromic acid are corrosive to metals (similar to chlorine); many metals are attacked by hot bromine, though some, like stainless steel and certain alloys, offer resistance. In general, bromine solutions must be handled in inert materials (glass, certain plastics, or alloy-lined containers) to avoid corrosion.

Overall, bromine’s chemistry is dominated by its role as an oxidizer and halogenator. It will take part in redox reactions to reach the –1 oxidation state and can form acids (HBr, HOBr) and oxides under suitable conditions.

Occurrence and Production

Bromine is relatively rare in Earth’s crust (~2–3 ppm) but quite abundant in seawater (about 65 mg of bromide per kilogram of seawater, roughly 1/300 the abundance of chloride) It mainly occurs as bromide ion (Br⁻) dissolved in salty water. Very few bromine minerals exist – one example is bromargyrite (AgBr, a silver bromide mineral in some silver ore deposits).

Commercially, bromine is extracted from bromide-rich brines and salt lake deposits. Major bromine-producing sources include:

Dead Sea (Israel/Jordan) and the Great Salt Lake (USA), where natural evaporation concentrates bromide;
Oilfield brine wells (e.g. Michigan and Arkansas in the USA) containing bromide;
Evaporation ponds in other saline lakes or manufactured brines.

The typical process is oxidation of bromide (Br⁻) to bromine (Br₂). For example, chlorine gas or sodium hypochlorite is added to bromide-containing brine:

<code>2 Br⁻ + Cl₂ → Br₂ + 2 Cl⁻
</code>

This liberates bromine, which is then distilled from the brine (often under vacuum) to produce pure bromine liquid. Other methods include electrolysis of bromide-rich solutions.

World production of bromine is on the order of a few hundred thousand tonnes per year. As of recent data, leading producers are the United States (notably Arkansas, Michigan, and Louisiana brines), Israel and Jordan (Dead Sea), China and Japan (often from seawater), as well as smaller outputs from Russia and Europe The United States Geological Survey notes that a significant portion of bromine in the atmosphere today comes from human activities, largely due to bromine use in industry

Applications and Technology

Bromine’s principal uses leverage its chemical reactivity and the properties of its compounds. The largest application is flame retardants: many polymers (plastics, foams, textiles, electronics) are treated with brominated flame-retardant additives to inhibit combustion. Over half of all bromine produced goes into brominated compounds such as polybrominated diphenyl ethers (PBDEs) and tetrabromobisphenol A, which trap free radicals in flames and slow fire spread

Other major applications include:

Oil and gas drilling: Dense brine solutions of bromide salts (e.g. zinc bromide ZnBr₂, calcium bromide CaBr₂) are used in well drilling and completion to control pressure. These heavy, transparent fluids help prevent blowouts in deep wells.
Chemical synthesis: Bromine is an intermediate in making many chemicals. Methyl bromide (CH₃Br) was widely used as a pesticide fumigant and as a chemical building block (now largely phased out due to ozone effects). Bromine reacts with olefins or aromatic compounds to produce alkyl bromides or aryl bromides, used in fine chemicals and pharmaceuticals. For example, bromoform (CHBr₃) and bromobenzene are prepared from bromine.
Photography and imaging: Silver bromide (AgBr) is light-sensitive and was the basis for traditional black-and-white photographic film and paper. (Digital imaging has largely supplanted this, but AgBr emulsions were major consumers of bromide.)
Water treatment and disinfection: Bromine (often in the form of sodium bromide plus chlorine) generates hypobromous acid (HOBr), an effective biocide. It is used to disinfect swimming pools, hot tubs, and industrial cooling systems as an alternative to chlorine. Bromides also serve as corrosion inhibitors in recirculating water systems.
Medicine: In the 19th century, bromide salts (especially potassium bromide) were used as sedatives and anticonvulsants, exploiting bromide’s depressing effect on the nervous system. Modern medications rarely use bromides, but bromine still appears in some drugs (e.g. respiratory anesthetics or antihistamines).
Analytical chemistry: Bromine water (aqueous Br₂) is a classical reagent for testing unsaturation in organic molecules (alkenes decolorize it). N-bromosuccinimide (NBS) is a mild brominating agent in organic synthesis.
Materials & electronics: Bromine compounds are used in printing circuit boards (flame-retardant laminates) and in some specialty semiconductors or electrolytes (e.g. zinc-bromine flow batteries for energy storage).

Overall, bromine acts as a crucial intermediate in industry: its ability to add or remove bromine atoms makes it versatile for making complex molecules, while its flame-retardant properties meet fire-safety needs.

Biology, Environment, and Safety

Bromine and its compounds have notable biological and environmental effects. Bromide (Br⁻) is present in sea water and in the human body at trace levels, and it is now recognized as beneficial (possibly essential) for certain biological functions. For example, marine algae and some bacteria produce organobromine compounds (like bromo-phenols and bromo-alkanes) that can serve as chemical signals or defenses. In humans, bromide is utilized by eosinophil white blood cells to produce hypobromous acid (HOBr) as an antimicrobial oxidant. Recent studies suggest bromide is needed for proper formation of collagen tissues (making it a trace element). However, excess bromide (from salts like KBr) can cause neurological symptoms (“bromism”) revealing its toxicity at high doses

A major concern is “brominated” persistent pollutants. Many industrial compounds, especially brominated flame retardants (like PBDEs), are very stable and bioaccumulate, leading to ecological and health worries (endocrine disruption, etc.). Volatile organobromines like methyl bromide also impact the atmosphere: methyl bromide was found to deplete stratospheric ozone much more effectively than chlorine and has been nearly banned by the Montreal Protocol. Even natural bromine cycles matter: sea-spray and oceanic emissions of bromine-containing gases contribute to atmospheric chemistry and can deplete ozone locally (bromine radicals destroy ozone 40–100 times more efficiently than chlorine radicals

Safety: Elemental bromine is highly hazardous. Its vapors (even a few ppm) irritate and burn the respiratory tract, eyes, and skin. Inhalation can cause coughing, chemical burns in the lungs, or pulmonary edema. Liquid bromine causes severe chemical burns on contact. The permissible exposure limit is very low (typically around 0.1–0.5 ppm); adequate ventilation or respirators are required when handling it. Bromine is also corrosive to many materials; it should be stored in glass or compatible metal containers (not ordinary plastics). Hydrobromic acid (HBr) is similarly corrosive and toxic.

In environmental terms, bromide ions in drinking water can form brominated byproducts during chlorination (such as bromate or bromoform), which are regulated due to potential carcinogenicity. Bromine’s strong oxidative chemistry demands careful handling: spills on organics can cause fires or toxic fumes. Overall, bromine must be handled with caution (gloves, fume hood, eye protection), and waste streams containing bromides are often treated before release.

History and Etymology

Bromine was discovered relatively late among the elements. In 1826 the French chemist Antoine-Jérôme Balard identified it by treating seaweed ash washings near Montpellier and observed a reddish vapor. Independently, Carl Löwig had obtained a similar gas in 1825. Balard published first, naming the element “bromine” from the Ancient Greek word brômos, meaning “stench,” referring to its sharp smell Initially it was isolated by adding chlorine to brine to free the reddish gas, which was then condensed into the liquid known today.

After its discovery, bromine found various uses. In the late 19th century, potassium bromide became a common sedative in Europe and the United States (sold as a calming medication). The light-sensitive nature of silver bromide was crucial to the invention of photography (Daguerreotype and later film). The compound bis(2,4,6-tribromophenyl) (a brominated phenol) was one of the first modern flame retardants. In World War I, an organic bromine compound called "xylyl bromide" was used covertly as a lachrymator (tear gas). The ozone-depleting effects of methyl bromide led to its phase-out around 2005.

Bromine’s history is thus tied both to traditional chemistry (naming the halogens) and to public policy (e.g., Montreal Protocol). Its name is one of a few elements directly reflecting a sensory property (here, odor). Over time, bromine has evolved from a laboratory curiosity to an industrial staple with important cultural and environmental impacts (for example, the rise and regulation of brominated flame retardants in electronics and furnishings).

Data Table

Property	Value
Atomic number (Z)	35
Symbol	Br
Standard atomic weight	79.904 u
Electron configuration	[Ar] 3d¹⁰ 4s² 4p⁵
Valence electrons	7
Group / Period / Block	17 (halogen) / 4 / p-block
Stable isotopes	⁷⁹Br (~51%), ⁸¹Br (~49%)
Common oxidation states	–1; +1, +3, +5, +7
State at STP	Liquid (reddish-brown)
Density (liquid, 20 °C)	3.12 g/cm³
Melting point	–7.2 °C
Boiling point	58.8 °C
Electronegativity	2.96 (Pauling)
First ionization energy	1139 kJ/mol
Electron affinity	324.6 kJ/mol
Covalent radius	~120 pm
Ionic radius (Br⁻)	~196 pm
Thermal conductivity	0.122 W·m⁻¹·K⁻¹ (liquid)
Heat capacity (liquid)	~0.47 J·g⁻¹·K⁻¹
Appearance	Reddish-brown liquid