Bromine is a dark reddish-brown, volatile liquid halogen at standard conditions. It forms corrosive, toxic vapors and commonly occurs as diatomic Br₂.
Bromine has the electronic configuration \([Ar]3d^{10}\,4s^2\,4p^5\). It is in Group 17 (Halogens) and Period 4 of the periodic table. Like other halogens, it has seven valence electrons and tends to gain one electron to form the bromide ion \(\mathrm{Br^-}\).
Bromine’s intermediate atomic size and polarizability give it stronger van der Waals forces than chlorine (gas) but weaker than iodine (solid). Thus, it exists as a volatile liquid at room temperature (~20 °C).
Bromine exhibits oxidation states of −1, +1, +3, +5, and +7.
Bromine is obtained by oxidizing bromide ions (e.g., from sea water brines) with chlorine:
\(\mathrm{2\,Br^- + Cl_2 \rightarrow 2\,Cl^- + Br_2}\)
The bromine vapor is condensed and purified by distillation.
With hydrogen:
\(\mathrm{H_2 + Br_2 \xrightarrow{heat/light} 2\,HBr}\)
The reaction is slower than that of chlorine and reversible at high temperatures.
With metals: Bromine forms metal bromides such as \(\mathrm{2\,Fe + 3\,Br_2 \rightarrow 2\,FeBr_3}\).
Bromine reacts with water in a reversible disproportionation reaction:
\(\mathrm{Br_2 + H_2O \rightleftharpoons HBr + HBrO}\)
This forms hydrobromic acid (HBr) and hypobromous acid (HOBr). The latter is a weak oxidizing agent and used in disinfection.
Certain brominated compounds (like halons and PBDEs) are persistent pollutants that can deplete ozone or bioaccumulate. Modern use is controlled under environmental agreements to limit ozone-depleting substances.
Yes. Bromine and its vapors are highly corrosive and toxic, causing burns, respiratory irritation, and lung damage. Always handle bromine in a fume hood with gloves, goggles, and protective clothing. Contact with organic material can lead to violent reactions.