Hypobromite
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| Names | |
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| IUPAC name
hypobromite
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| Identifiers | |
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3D model (JSmol)
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| EC Number | Page Template:Plainlist/styles.css has no content. |
| 1040 | |
PubChem CID
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| Properties | |
| Page Module:Chem2/styles.css has no content.BrO− | |
| Molar mass | 95.903 g·mol−1 |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Hypobromite, also called alkaline bromine water, is an anion with the chemical formula Page Module:Chem2/styles.css has no content.BrO−. Bromine is in the +1 oxidation state. The Page Module:Chem2/styles.css has no content.Br−O bond length is 1.82 Å.[1] Hypobromite is the bromine compound analogous to hypochlorites found in common bleaches, and in immune cells. In many ways, hypobromite functions in the same manner as hypochlorite, and is also used as a germicide and antiparasitic in both industrial applications, and in the immune system.
Preparation
Hypobromite salts form upon treating bromine with aqueous alkali, such as sodium or potassium hydroxide. At Template:Cvt the reaction is rapid.[2]
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In this reaction the bromine disproportionates (some undergoes reduction and some oxidation) from oxidation state 0 (Page Module:Chem2/styles.css has no content.Br2) to oxidation state −1 (Page Module:Chem2/styles.css has no content.Br−) and oxidation state +1 (Page Module:Chem2/styles.css has no content.BrO−). Sodium hypobromite can be isolated as an orange solid.
A secondary reaction, where hypobromite spontaneously disproportionates to bromide (bromine oxidation state −1) and bromate (bromine oxidation state +5) takes place rapidly at Template:Cvt and slowly at Template:Cvt.[citation needed]
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Hence, in reaction 2, the formation and proportions of the −1, +1 and +5 bromine oxidation state products can be controlled by temperature. Hypobromite is not thermodynamically stable at any pH (see Pourbaix diagram for bromine[3]), but it is kinetically locked toward a further disproportion into bromate and bromide above the pKa of HBrO.
These reactions of bromine are analogous to those of chlorine forming hypochlorite and chlorate. The corresponding chlorine reaction 1 (to form Page Module:Chem2/styles.css has no content.ClO−) is fast at Template:Cvt and reaction 2 (to form Page Module:Chem2/styles.css has no content.ClO−3) is slow at Template:Cvt and fast at Template:Cvt.
In nature and industry
Bromide from the diet, naturally present in the blood, is used by eosinophils, white blood cells of the granulocyte class, specialized for dealing with multi-cellular parasites. These cells react the bromide with peroxide to generate hypobromite by the action of eosinophil peroxidase, a haloperoxidase enzyme which preferentially uses bromide over chloride for this purpose.[4]
Simple bromide salts (such as sodium bromide) are also sometimes used in hot tubs and spas as mild germicidal agents, using the action of an added oxidizing agent (such as hydrogen peroxide) to generate hypobromite in situ, in a similar fashion to the action of peroxidase on bromide in eosinophils.[citation needed]
Hypobromite has been proposed to be a reactive intermediate in the Hofmann rearrangement.[citation needed]
See also
Other bromine anions:
| Bromine oxidation state | −1 | +1 | +3 | +5 | +7 |
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| Name | bromide | hypobromite | bromite | bromate | perbromate |
| Formula | Page Module:Chem2/styles.css has no content.Br− | Page Module:Chem2/styles.css has no content.BrO− | Page Module:Chem2/styles.css has no content.BrO−2 | Page Module:Chem2/styles.css has no content.BrO−3 | Page Module:Chem2/styles.css has no content.BrO−4 |
| Structure | The bromide ion | The hypobromite ion | The bromite ion | The bromate ion | The perbromate ion |
References
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- ^ Page Module:Citation/CS1/styles.css has no content.Topić, Filip; Marrett, Joseph M.; Borchers, Tristan H.; Titi, Hatem M.; Barrett, Christopher J.; Friščić, Tomislav (2021). "After 200 Years: The Structure of Bleach and Characterization of Hypohalite Ions by Single-Crystal X-Ray Diffraction". Angew. Chem. Int. Ed. 60 (46): 24400–24405. doi:10.1002/anie.202108843.
- ^ Page Module:Citation/CS1/styles.css has no content.Kneen; Rogers; Simpson (1972). "The Halogens". Chemistry. Facts, Patterns and Principles. Addison-Wesley. ISBN 0-201-03779-3.
- ^ Page Module:Citation/CS1/styles.css has no content.Atlas of Eh-pH diagrams, Intercomparison of thermodynamic databases (PDF). eosremediation.com (Report). Geological Survey of Japan Open File Report. Vol. 419. National Institute of Advanced Industrial Science and Technology - Research Center for Deep Geological Environments. May 2005. p. 49. Archived from the original (PDF) on 30 October 2014.
- ^ Page Module:Citation/CS1/styles.css has no content.Mayeno, A N; Curran, A J; Roberts, R L; Foote, C S (5 April 1989), "Eosinophils preferentially use bromide to generate halogenating agents", Journal of Biological Chemistry, 264 (10): 5660–8, doi:10.1016/S0021-9258(18)83599-2, PMID 2538427, archived from the original on 26 May 2009, retrieved 25 October 2009
