Hydrogen bromide is the inorganic compound with the formula HBr. It is a hydrogen halide consisting of hydrogen and bromine. A colorless gas, it dissolves in water, forming hydrobromic acid, which is saturated at 68.85% HBr by weight at room temperature. Aqueous solutions that are 47.6% HBr by mass form a constant-boiling azeotrope mixture that boils at 124.3 °C (255.7 °F). Boiling less concentrated solutions releases H2O until the constant-boiling mixture composition is reached.
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| Names | |||
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| IUPAC name
Hydrogen bromide
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| Preferred IUPAC name
Bromane[1] | |||
| Identifiers | |||
3D model (JSmol)
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| 3587158 | |||
| ChEBI | |||
| ChEMBL | |||
| ChemSpider | |||
| ECHA InfoCard | 100.030.090 | ||
| EC Number |
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| KEGG | |||
| MeSH | Hydrobromic Acid | ||
PubChem CID
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| RTECS number |
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| UNII | |||
| UN number | 1048 | ||
CompTox Dashboard (EPA)
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| Properties | |||
| HBr | |||
| Molar mass | 80.91 g/mol | ||
| Appearance | Colorless gas | ||
| Odor | Acrid | ||
| Density | 3.307 g/L (25 °C)[2] | ||
| 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 alcohol, organic solvents | ||
| Vapor pressure | 2.308 MPa (at 21 °C) | ||
| Acidity (pKa) | −8.8 (±0.8);[3] ~−9[4] | ||
| Basicity (pKb) | ~23 | ||
| Conjugate acid | Bromonium | ||
| Conjugate base | Bromide | ||
Refractive index (nD)
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1.325[citation needed] | ||
| Structure | |||
| Linear | |||
| 820 mD | |||
| Thermochemistry | |||
Heat capacity (C)
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350.7 mJ/(K·g) | ||
Std molar
entropy (S⦵298) |
198.696–198.704 J/(K·mol)[5] | ||
Std enthalpy of
formation (ΔfH⦵298) |
−36.45...−36.13 kJ/mol[5] | ||
| Hazards | |||
| Occupational safety and health (OHS/OSH): | |||
Main hazards
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Highly corrosive | ||
| GHS labelling: | |||
 
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| Danger | |||
| H314, H335 | |||
| P261, P280, P305 P351 P338, P310 | |||
| NFPA 704 (fire diamond) | |||
| Lethal dose or concentration (LD, LC): | |||
LC50 (median concentration)
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2858 ppm (rat, 1 h) 814 ppm (mouse, 1 h)[7] | ||
| NIOSH (US health exposure limits): | |||
PEL (Permissible)
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TWA 3 ppm (10 mg/m3)[6] | ||
REL (Recommended)
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TWA 3 ppm (10 mg/m3)[6] | ||
IDLH (Immediate danger)
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30 ppm[6] | ||
| Safety data sheet (SDS) | hazard.com | ||
| Related compounds | |||
Related compounds
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Hydrogen fluoride Hydrogen chloride Hydrogen iodide Hydrogen astatide | ||
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Hydrogen bromide, and its aqueous solution, hydrobromic acid, are commonly used reagents in the preparation of bromide compounds.
Reactions
editOrganic chemistry
editHydrogen bromide and hydrobromic acid are important reagents in the production of organobromine compounds.[8][9][10] In an electrophilic addition reaction, HBr adds to alkenes:
- RCH=CH2 HBr → R−CHBr−CH3
The resulting alkyl bromides are useful alkylating agents, e.g., as precursors to fatty amine derivatives. Related free radical additions to allyl chloride and styrene give 1-bromo-3-chloropropane and phenylethylbromide, respectively.
Hydrogen bromide reacts with dichloromethane to give bromochloromethane and dibromomethane, sequentially:
- HBr CH2Cl2 → HCl CH2BrCl
- HBr CH2BrCl → HCl CH2Br2
These metathesis reactions illustrate the consumption of the stronger acid (HBr) and release of the weaker acid (HCl).
Allyl bromide is prepared by treating allyl alcohol with HBr:
- CH2=CHCH2OH HBr → CH2=CHCH2Br H2O
HBr adds to alkynes to yield bromoalkenes. The stereochemistry of this type of addition is usually anti:
- RC≡CH HBr → RC(Br)=CH2
Also, HBr adds epoxides and lactones, resulting in ring-opening.
With triphenylphosphine, HBr gives triphenylphosphonium bromide, a solid "source" of HBr.[11]
- P(C6H5)3 HBr → [HP(C6H5)3] Br−
Inorganic chemistry
editVanadium(III) bromide and molybdenum(IV) bromide were prepared by treatment of the higher chlorides with HBr. These reactions proceed via redox reactions:[12]
- 2 VCl4 8 HBr → 2 VBr3 8 HCl Br2
Industrial preparation
editHydrogen bromide (along with hydrobromic acid) is produced by combining hydrogen and bromine at temperatures between 200 and 400 °C. The reaction is typically catalyzed by platinum or asbestos.[9][13]
Laboratory synthesis
editHBr can be prepared by distillation of a solution of sodium bromide or potassium bromide with phosphoric acid or sulfuric acid:[14]
- KBr H2SO4 → KHSO4 HBr
Concentrated sulfuric acid is less effective because it oxidizes HBr to bromine:
- 2 HBr H2SO4 → Br2 SO2 2 H2O
The acid may be prepared by:
- reaction of bromine with water and sulfur:[14]
- 2 Br2 S 2 H2O → 4 HBr SO2
- bromination of tetralin:[14]
- C10H12 4 Br2 → C10H8Br4 4 HBr
- reduction of bromine with phosphorous acid:[9]
- Br2 H3PO3 H2O → H3PO4 2 HBr
Anhydrous hydrogen bromide can also be produced on a small scale by thermolysis of triphenylphosphonium bromide in refluxing xylene.[11]
Hydrogen bromide prepared by the above methods can be contaminated with Br2, which can be removed by passing the gas through a solution of phenol at room temperature in tetrachloromethane or other suitable solvent (producing 2,4,6-tribromophenol and generating more HBr in the process) or through copper turnings or copper gauze at high temperature.[13]
Safety
editHBr is highly corrosive and, if inhaled, can cause lung damage.[15]
References
edit- ^ Favre, Henri A.; Powell, Warren H., eds. (2014). Nomenclature of Organic Chemistry: IUPAC Recommendations and Preferred Names 2013. Cambridge: The Royal Society of Chemistry. p. 131. ISBN 9781849733069.
- ^ Lide, David R., ed. (2006). CRC Handbook of Chemistry and Physics (87th ed.). Boca Raton, Florida: CRC Press. ISBN 0-8493-0487-3.
- ^ Trummal, Aleksander; Lipping, Lauri; Kaljurand, Ivari; Koppel, Ilmar A; Leito, Ivo (2016). "Acidity of Strong Acids in Water and Dimethyl Sulfoxide". The Journal of Physical Chemistry A. 120 (20): 3663–9. Bibcode:2016JPCA..120.3663T. doi:10.1021/acs.jpca.6b02253. PMID 27115918. S2CID 29697201.
- ^ Perrin, D. D. Dissociation constants of inorganic acids and bases in aqueous solution. Butterworths, London, 1969.
- ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. ISBN 978-0-618-94690-7.
- ^ a b c NIOSH Pocket Guide to Chemical Hazards. "#0331". National Institute for Occupational Safety and Health (NIOSH).
- ^ "Hydrogen bromide". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
- ^ Dagani, M. J.; Barda, H. J.; Benya, T. J.; Sanders, D. C. "Bromine Compounds". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a04_405. ISBN 978-3527306732.
{{cite encyclopedia}}: CS1 maint: multiple names: authors list (link) - ^ a b c Greenwood, N. N.; Earnshaw, A. Chemistry of the Elements; Butterworth-Heineman: Oxford, Great Britain; 1997; pp. 809–812.
- ^ Vollhardt, K. P. C.; Schore, N. E. Organic Chemistry: Structure and Function; 4th Ed.; W. H. Freeman and Company: New York, NY; 2003.
- ^ a b Hercouet, A.; LeCorre, M. (1988) Triphenylphosphonium bromide: A convenient and quantitative source of gaseous hydrogen bromide. Synthesis, 157–158.
- ^ Calderazzo, Fausto; Maichle-Mössmer, Cäcilie; Pampaloni, Guido; Strähle, Joachim (1993). "Low-Temperature Syntheses of Vanadium(III) and Molybdenum(IV) Bromides by Halide Exchange". J. Chem. Soc., Dalton Trans. (5): 655–658. doi:10.1039/DT9930000655.
- ^ a b Ruhoff, J. R.; Burnett, R. E.; Reid, E. E. "Hydrogen Bromide (Anhydrous)" Organic Syntheses, Vol. 15, p. 35 (Coll. Vol. 2, p. 338).
- ^ a b c 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.
- ^ "Hydrobromic Acid". PubChem. Springer Nature. 2024-02-03. Retrieved 2024-02-09.