| IUPAC name
| Other names
Anhydrous hydrobromic acid
|Molar mass||80.91 g/mol|
|Appearance|| Colorless gas|
Brownish gas (impure)
|Density||3.6452 kg/m3 (At 0 °C, 1013 mbar)|
|Melting point||−86.9 °C (−124.4 °F; 186.2 K)|
|Boiling point||−66.8 °C (−88.2 °F; 206.3 K)|
| 221 g/100 mL (0 °C)|
204 g/100 mL (15 °C)
193 g/100 mL (20 °C)
130 g/100 mL (100 °C)
|Solubility||Soluble in alcohols|
|Vapor pressure||2.308 MPa (at 21 °C)|
Std enthalpy of
|Safety data sheet||Praxair|
|Lethal dose or concentration (LD, LC):|
LC50 (Median concentration)
| 2858 ppm (rat, 1 hr)|
814 ppm (mouse, 1 hr)
| Hydrogen chloride|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Hydrogen bromide is a gaseous corrosive compound with the chemical formula HBr. When dissolved in water, the solution is known as hydrobromic acid.
Hydrogen bromide is a colorless gas, though in the presence of light it will slowly turn yellow-brownish, due to traces of bromine. It is extremely soluble in water, as well as alcohols and other organic solvents. Hydrogen bromide has an unpleasant acrid smell.
Anhydrous hydrogen bromide is only sold to industrial entities.
Reaction of diluted sulfuric acid with a bromide salt yields hydrobromic acid, which can be dehydrated through a variety of methods to hydrogen bromide, however it may decompose to form small amounts of elemental bromine.
A much better method involves the bromination of tetralin:
- C10H12 + 4 Br2 → C10H8Br4 + 4 HBr
All the bromide traces from the described processes can be removed by passing the impure hydrogen bromide though a layer of phenol or or through copper turnings at high temperature, in the absence of oxygen.
- Make anhydrous bromide salts
Hydrogen bromide is extremely corrosive and dangerous for lungs, mouth, nose and eyes if inhaled.
Anhydrous hydrogen bromide should be stored in special containers, in a separate location.
Gaseous hydrogen bromide can be neutralized with ammonia, though this will generate a mist of ammonium bromide. A better way is to react it with a base or sodium thiosulfate.
- M. Schmeisser "Chlorine, Bromine, Iodine" in Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 282.