View source for Chemistry:Bromide

{{Short description|Chemical compound or ion}}
<!-- sometimes excessively wordy or uses weird phrasing -->
{{Chembox
|ImageFileL1 = Br-.svg
|ImageSizeL1 = 50px
|ImageFileR1 = Bromide ion.svg
|ImageSizeR1 = 60px
|Verifiedfields = changed
|Watchedfields = changed
|verifiedrevid = 458436846
|SystematicName = Bromide<ref>{{Cite web|title = Bromide – PubChem Public Chemical Database|url = https://pubchem.ncbi.nlm.nih.gov/summary/summary.cgi?cid=259|work = The PubChem Project|location = USA|publisher = National Center for Biotechnology Information|url-status = live|archive-url = https://web.archive.org/web/20121103072332/https://pubchem.ncbi.nlm.nih.gov/summary/summary.cgi?cid=259|archive-date = 2012-11-03}}</ref>
|Section1={{Chembox Identifiers
|CASNo_Ref = {{cascite|correct|CAS}}
|CASNo = 24959-67-9
|UNII_Ref = {{fdacite|correct|FDA}}
|UNII = 952902IX06
|PubChem = 259
|ChemSpiderID = 254
|ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
|KEGG = C01324
|KEGG_Ref = {{keggcite|changed|kegg}}
|ChEBI_Ref = {{ebicite|correct|EBI}}
|ChEBI = 15858
|ChEMBL = 11685
|ChEMBL_Ref = {{ebicite|correct|EBI}}
|Beilstein = 3587179
|Gmelin = 14908
|SMILES = [Br-]
|StdInChI = 1S/BrH/h1H/p-1
|StdInChI_Ref = {{stdinchicite|correct|chemspider}}
|StdInChIKey = CPELXLSAUQHCOX-UHFFFAOYSA-M
|StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
}}
|Section2={{Chembox Properties
|Formula = Br<sup>−</sup>
|MolarMass = 79.904&nbsp;g·mol<sup>−1</sup>
|ConjugateAcid = [[Chemistry:Hydrogen bromide|Hydrogen bromide]]
}}
|Section6={{Chembox Pharmacology
|ATCCode_prefix = N05
|ATCCode_suffix = CM11
|HalfLife = 12 d}}
|Section5={{Chembox Thermochemistry
|DeltaHf = −121&nbsp;kJ·mol<sup>−1</sup><ref name=b1>{{cite book| last= Zumdahl |first=Steven S.|title =Chemical Principles |edition=6th | publisher = Houghton Mifflin | year = 2009| isbn = 978-0-618-94690-7}}</ref>
|Entropy = 82&nbsp;J·mol<sup>−1</sup>·K<sup>−1</sup><ref name=b1/>
}}
|Section8={{Chembox Related
|OtherAnions = [[Chemistry:Fluoride|Fluoride]]<br />
[[Chemistry:Chloride|Chloride]]<br />
[[Chemistry:Iodide|Iodide]]
}}
}}

A '''bromide''' [[Physics:Ion|ion]] is the negatively charged form ('''Br<sup>−</sup>''') of the element [[Chemistry:Bromine|bromine]], a member of the [[Chemistry:Halogen|halogens group]] on the [[Chemistry:Periodic table|periodic table]]. Most bromides are colorless. Bromides have many practical roles, being found in anticonvulsants, flame-retardant materials, and cell stains.<ref>{{cite journal |doi=10.1038/nchem.1361|title=Ambiguous bromine|year=2012|last1=Rattley|first1=Matt|journal=Nature Chemistry|volume=4|issue=6|page=512|pmid=22614389|bibcode=2012NatCh...4..512R|doi-access=free}}</ref>  Although uncommon, chronic toxicity from bromide can result in [[Medicine:Bromism|bromism]], a syndrome with multiple neurological symptoms. Bromide toxicity can also cause a type of skin eruption, see [[Chemistry:Potassium bromide|potassium bromide]]. The bromide ion has an [[Chemistry:Ionic radius|ionic radius]] of 196&nbsp;pm.<ref>{{cite journal|title=Revised effective ionic radii and systematic studies of interatomic distances in halides and chalcogenides |journal=Acta Crystallographica A |date=1976 |volume=32 |pages=751–767 |doi=10.1107/s0567739476001551|last1=Shannon |first1=R. D. |issue=5 }}</ref>

==Natural occurrence==
Bromide is present in typical [[Earth:Seawater|seawater]] (35&nbsp;[[Earth:Salinity|PSU]]) with a concentration of around 65&nbsp;mg/L, which is about 0.2% of all dissolved [[Chemistry:Salt|salt]]s. Seafood and deep sea plants generally have higher levels than land-derived foods. [[Chemistry:Bromargyrite|Bromargyrite]]—natural, crystalline silver bromide—is the most common bromide mineral known but is still very rare. In addition to silver, bromine is also in minerals combined with mercury and copper.<ref>{{cite web|url=https://www.mindat.org|title=Mindat.org - Mines, Minerals and More|website=www.mindat.org|access-date=29 April 2018|url-status=live|archive-url=https://web.archive.org/web/20010302042623/https://www.mindat.org/|archive-date=2 March 2001}}</ref>

==Formation and reactions of bromide==
===Dissociation of bromide salts===
Bromide salts of [[Chemistry:Alkali metal|alkali metal]], [[Chemistry:Alkaline earth metal|alkaline earth metal]]s, and many other metals dissolve in water (and even some alcohols and a few ethers) to give bromide ions.  The classic case is sodium bromide, which fully dissociates in water:
:NaBr  →  Na<sup>+</sup>  +  Br<sup>−</sup>

Hydrogen bromide, which is a [[Chemistry:Diatomic molecule|diatomic molecule]], takes on salt-like properties upon contact with water to give an ionic solution called [[Chemistry:Hydrobromic acid|hydrobromic acid]]. The process is often described simplistically as involving formation of the hydronium salt of bromide:
:HBr + H<sub>2</sub>O → H<sub>3</sub>O<sup>+</sup> + Br<sup>−</sup>

===Hydrolysis of bromine===
Bromine readily reacts with water, i.e. it undergoes hydrolysis:			
:Br<sub>2</sub> + H<sub>2</sub>O → HOBr + HBr

This forms [[Chemistry:Hypobromous acid|hypobromous acid]] (HOBr), and [[Chemistry:Hydrobromic acid|hydrobromic acid]] (HBr in water). The solution is called "[[Chemistry:Bromine water|bromine water]]".  The hydrolysis of bromine is more favorable in the presence of base, for example [[Chemistry:Sodium hydroxide|sodium hydroxide]]:

:Br<sub>2</sub> + NaOH → NaOBr  + NaBr
This reaction is analogous to the production of [[Chemistry:Bleach|bleach]], where chlorine is dissolved in the presence of sodium hydroxide.<ref>Chemistry of the Elements, N. N. Greenwood, A. Earnshaw, Elsevier, 2012, pp 789</ref>

===Oxidation of bromide===
One can test for a bromide ion by adding an oxidizer. One method uses dilute [[Chemistry:Nitric acid|HNO<sub>3</sub>]].

Balard and Löwig's method can be used to extract bromine from seawater and certain brines. For samples testing for sufficient bromide concentration, addition of chlorine produces bromine (Br<sub>2</sub>):<ref>{{Cite journal |last1 = Magazinovic |first1 = Rodney S. |last2 = Nicholson |first2 = Brenton C. |last3 = Mulcahy |first3 = Dennis E. |last4 = Davey |first4 = David E. |date = 2004 |title = Bromide levels in natural waters: its relationship to levels of both chloride and total dissolved solids and the implications for water treatment |url = https://linkinghub.elsevier.com/retrieve/pii/S0045653504003935 |journal = Chemosphere |language = en |volume = 57 |issue = 4 |pages = 329–335 |doi = 10.1016/j.chemosphere.2004.04.056 |pmid = 15312731 |bibcode = 2004Chmsp..57..329M |access-date = 2021-03-07 |archive-date = 2021-05-25 |archive-url =https://web.archive.org/web/20210525180633/https://linkinghub.elsevier.com/retrieve/pii/S0045653504003935 |url-status = live}}</ref>
:Cl<sub>2</sub> + 2&nbsp;Br<sup>−</sup>  → 2&nbsp;Cl<sup>−</sup> + Br<sub>2</sub>

==Applications==
Bromide's main commercial value is its use in producing [[Chemistry:Organobromine compound|organobromine compound]]s, which themselves are rather specialized. Organobromine compounds are commonly used as brominated flame retardants.<ref>Michael J. Dagani, Henry J. Barda, Theodore J. Benya, David C. Sanders: ''Bromine Compounds'', Ullmann's Encyclopedia of Industrial Chemistry 2002, Wiley-VCH, Weinheim. {{doi|10.1002/14356007.a04_405}}</ref> Some brominated flame retardants were identified as persistent, bioaccumulative, and toxic to both humans and the environment and were suspected of causing neurobehavioral effects and endocrine disruption.<ref>{{cite web |url=http://www.epa.gov/oppt/existingchemicals/pubs/actionplans/pbde.html |title=Polybrominated Diphenyl Ethers (PBDEs) Action Plan Summary &#124; Existing Chemicals &#124; OPPT &#124; US EPA |access-date=2012-12-03 |url-status=live |archive-url=https://web.archive.org/web/20150901200633/http://www.epa.gov/oppt/existingchemicals/pubs/actionplans/pbde.html |archive-date=2015-09-01 }}</ref><ref>{{cite web |url=http://www.cerc.usgs.gov/pubs/center/pdfdocs/pbde.pdf |title=Archived copy |access-date=2012-12-03 |url-status=live |archive-url=https://web.archive.org/web/20160508170026/http://www.cerc.usgs.gov/pubs/center/pdfDocs/PBDE.pdf |archive-date=2016-05-08 }}</ref>

Many metal bromides are produced commercially, including [[Chemistry:Lithium bromide|LiBr]], [[Chemistry:Sodium bromide|NaBr]], [[Chemistry:Ammonium bromide|NH<sub>4</sub>Br]], [[Chemistry:Copper(I) bromide|CuBr]], [[Chemistry:Zinc(II) bromide|ZnBr<sub>2</sub>]] and AlBr<sub>3</sub>. [[Chemistry:Silver bromide|AgBr]] is used for the largely obsolete photographic gelatin silver process.<ref name=Weaver>{{cite web|last1=Weaver|first1=Gawain|title=A Guide to Fiber-Base Gelatin Silver Print Condition and Deterioration|url=http://gawainweaver.com/images/uploads/Weaver_Guide_to_Gelatin_Silver.pdf|publisher=George Eastman House, International Museum of Photography and Film|access-date=30 October 2009|date=2008}}</ref>

==Medicinal and veterinary uses==
[[File:Fox Plate XIV.jpg|thumb|[[Medicine:Dermatitis|Dermatitis]] reactions to bromide, all except lower right]]
===Folk and passé medicine===
[[Chemistry:Lithium bromide|Lithium bromide]] was used as a [[Chemistry:Sedative|sedative]] beginning in the early 1900s. However, it fell into disfavour in the 1940s due to the rising popularity of safer and more efficient sedatives (specifically, barbiturates) and when some heart patients died after using a salt substitute (see [[Chemistry:Lithium chloride|lithium chloride]]).<ref>[https://www.webmd.com/content/Article/87/99356.htm?pagenumber=1 Bipolar disorder] {{Webarchive|url=https://web.archive.org/web/20220224085437/https://www.webmd.com/bipolar-disorder/guide/bipolar-disorder-treatment-care?pagenumber=1 |date=2022-02-24}}. webmd.com</ref> Like [[Chemistry:Lithium carbonate|lithium carbonate]] and [[Chemistry:Lithium chloride|lithium chloride]], it was used as a treatment for [[Medicine:Bipolar disorder|bipolar disorder]].

From 1954 - 1977, the Australian [[Biochemist|biochemist]] Shirley Andrews was researching safe ways to use [[Chemistry:Lithium (medication)|lithium]] for the treatment of [[Medicine:Bipolar disorder|manic depressive illnesses]] while working at the Royal Park Psychiatric Hospital in Victoria. While conducting this research she discovered that bromide caused symptoms of mental illness, leading to a major reduction in its usage.<ref>{{Cite web |title=Papers of Shirley Andrews |url=https://nla.gov.au/nla.obj-233797980 |access-date=2022-10-26 |website=Trove |language=en}}</ref>
{{Further|Chemistry:Potassium bromide}}

Bromide compounds, especially [[Chemistry:Potassium bromide|potassium bromide]], were frequently used as sedatives in the 19th and early 20th centuries. Their use in over-the-counter sedatives and headache remedies (such as [[Chemistry:Bromo-Seltzer|Bromo-Seltzer]]) in the United States extended to 1975 when bromides were withdrawn as ingredients due to [[Chemistry:Chronic toxicity|chronic toxicity]].<ref>{{Cite book | url =https://archive.org/details/greatamericanfr03adamgoog | title = The Great American fraud | publisher =Press of the American Medical Association | last1 =Adams | first1 =Samuel Hopkins | year =1905}}.</ref> This use gave the word "bromide" its colloquial connotation of a comforting cliché.<ref name="Dictionary">{{cite web|title=the definition of bromide|url=https://www.dictionary.com/browse/bromide|website=Dictionary.com|access-date=21 December 2016|url-status=live|archive-url=https://web.archive.org/web/20161224192445/https://www.dictionary.com/browse/bromide|archive-date=24 December 2016}}</ref>

It has been said that during [[Social:World War I|World War I]], British soldiers were given bromide to curb their sexual urges.<ref>Tanaka, Yuki (2002) ''Japan's Comfort Women: Sexual slavery and prostitution during World War II and the US Occupation'', Routledge, p. 175. {{ISBN|0415194008}}.</ref> Lord Dunsany mentions a soldier being given bromide as a sedative for nervous exhaustion and overwork in his play ''Fame and the Poet'' (1919).<ref>{{cite journal|author=Lord Dunsany|title=Fame and the Poet|journal=The Atlantic Monthly|date=August 1919|pages=175–183}}</ref>

Bromide salts are used in [[Engineering:Hot tub|hot tub]]s as mild germicidal agents to generate ''in situ'' [[Chemistry:Hypobromite|hypobromite]]

The bromide ion is antiepileptic and as bromide salt, is used in veterinary medicine in the US. The kidneys excrete bromide ions. The half-life of bromide in the human body (12 days) is long compared with many pharmaceuticals, making dosing challenging to adjust. (A new dose may require several months to reach equilibrium.) Bromide ion concentrations in the [[Biology:Cerebrospinal fluid|cerebrospinal fluid]] are about 30% of those in blood and are strongly influenced by the body's chloride intake and metabolism.<ref>{{cite book|editor1-last=Goodman |editor1-first=L. S. |editor2-last=Gilman |editor2-first=A. |date=1970 |chapter=10. Hypnotics and Sedatives |page=121 |title=The Biological Basis of Therapeutics |edition=4th |publisher=Macmillan |location=London}}</ref>
 
Since bromide is still used in veterinary medicine in the United States, veterinary diagnostic labs can routinely measure blood bromide levels. However, this is not a conventional test in human medicine in the US since there are no FDA-approved uses for the bromide. Therapeutic bromide levels are measured in European countries like {{wipe|Germany}}, where bromide is still used therapeutically in human epilepsy.

==Biochemistry==
Bromide is rarely mentioned in the biochemical context.  Some enzymes use bromide as substrate or as a cofactor. 

===Substrate===
[[Biology:Bromoperoxidase|Bromoperoxidase]] enzymes use bromide (typically in seawater) to generate electrophilic brominating agents Hundreds of [[Chemistry:Organobromine compound|organobromine compound]]s are generated by this process. Notable examples are bromoform, thousands of tons of which are produced annually in this way. The historical dye [[Tyrian purple]] is produced by similar enzymatic reactions.<ref>{{cite journal|first=Gordon W. |last=Gribble |title=The diversity of naturally occurring organobromine compounds |journal=Chemical Society Reviews |date=1999 |volume=28 |issue=5 |pages=335–346 |doi=10.1039/a900201d}}</ref>

===Cofactor===
In one specialized report, bromide is an essential cofactor in the peroxidising catalysis of sulfonimine crosslinks in collagen IV. This [[Biology:Post-translational modification|post-translational modification]] occurs in all animals and bromine is an essential trace element for humans.<ref name="pmid24906154">{{Cite journal |last1 = McCall |first1 = A. Scott |last2 = Cummings |first2 = Christopher F. |last3 = Bhave |first3 = Gautam |last4 = Vanacore |first4 = Roberto |last5 = Page-McCaw |first5 = Andrea |last6 = Hudson |first6 = Billy G. |title = Bromine Is an Essential Trace Element for Assembly of Collagen IV Scaffolds in Tissue Development and Architecture |url= |journal = Cell |language = en |volume = 157 |issue = 6 |pages = 1380–1392 |year = 2014 |doi = 10.1016/j.cell.2014.05.009 |pmc = 4144415 |pmid = 24906154}}</ref>

Eosinophils need bromide for fighting multicellular parasites. [[Chemistry:Hypobromite|Hypobromite]] is produced via [[Biology:Eosinophil peroxidase|eosinophil peroxidase]], an enzyme that can use chloride but preferentially uses bromide.<ref>
{{Cite journal |last1 = Mayeno |first1 = Arthur N. |last2 = Curran |first2 = A. Jane |last3 = Roberts |first3 = Robert L. |last4 = Foote |first4 = Christopher S. |date = 1989-04-05 |title = Eosinophils Preferentially Use Bromide to Generate Halogenating Agents |journal = Journal of Biological Chemistry |volume = 264 |issue = 10 |pages = 5660–5668 |doi = 10.1016/s0021-9258(18)83599-2 |issn = 0021-9258 |pmid = 2538427|doi-access = free}}</ref>

The average concentration of bromide in human blood in Queensland, Australia, is {{val|5.3|1.4|u=mg/L}} and varies with age and gender.<ref>{{cite journal|pmid=9602940|year=1998|last1=Olszowy|first1=HA|last2=Rossiter|first2=J|last3=Hegarty|first3=J|last4=Geoghegan|first4=P|title=Background levels of bromide in human blood|volume=22|issue=3|pages=225–30|journal=Journal of Analytical Toxicology|doi=10.1093/jat/22.3.225|doi-access=free}}</ref> Much higher levels may indicate exposure to brominated chemicals. It is also found in seafood.

==Further reading==
===Encyclopedia articles and books===

* Christe, K., and S. Schneider (2020), Bromine, Encyclopædia Britannica. 
* Emerson, S., and J. Hedges (2011), Chemical Oceanography and the Marine Carbon Cycle, Cambridge University Press, Cambridge. 
* Glasow, R. von, and C. Hughes (2014), Biogeochemical Cycles: Bromine, Encyclopedia of Atmospheric Sciences (Second Edition). 
* Knight, J., and N. Schlager (2002), Real-life chemistry, Gale Group, Detroit, MI. 
* Millero, F. J. (2013), Chemical oceanography, Taylor & Francis, Boca Raton. 
* Newton D. E. (2010), Bromine (Revised), Chemical Elements: From Carbon to Krypton.    
* Riley, J. P., G. Skirrow, and R. Chester (1975), Chemical Oceanography, Academic Press, London 
* Ross, R. (2017), Facts About Bromine, LiveScience. 
* Steele, J. H., S. A. Thorpe, and K. K. Turekian (2001), Encyclopedia of Ocean Sciences, Academic Press, San Diego.
* Steele, J. H., S. A. Thorpe, and K. K. Turekian (2009), Encyclopedia of Ocean Sciences, Academic Press, Boston.	
* Watkins, T. (2011), Bromine, Environmental Encyclopedia.

===Peer-reviewed journal articles for bromine (Br)===
* Wisniak, J. (2002), The history of bromine from discovery to commodity, NOPR.

===Peer-reviewed journal articles for bromide (Br<sup>−</sup>)===
* Anbar, A. D., Y. L. Yung, and F. P. Chavez (1996), Methyl bromide: Ocean sources, ocean sinks, and climate sensitivity, AGU Journals. 
* Foti, S. C., and Naval Ordnance Lab White Oak Md (1972), Concentration of Bromide Ions in Seawater by Isotopic Exchange with Mercurous Bromide, DTIC. 
* Gribble, G. W. (2000), The natural production of organobromine compounds, Environmental Science and Pollution Research, 7(1), 37–49, {{doi|10.1065/espr199910.002}}. 
* Leri A. (2012), The Chemistry of Bromine in Terrestrial and Marine Environments, Science Highlight.
* Magazinovic, R. S., B. C. Nicholson, D. E. Mulcahy, and D. E. Davey (2004), Bromide levels in natural waters: its relationship to levels of both chloride and total dissolved solids and the implications for water treatment, Chemosphere, 57(4), 329–335, {{doi|10.1016/j.chemosphere.2004.04.056}}. 
* Pilinis, C., D. B. King, and E. S. Saltzman (1996), The oceans: A source or a sink of methyl bromide?, Geophysical Research Letters, 23(8), 817–820, {{doi|10.1029/96gl00424}}. 
* Stemmler, I., I. Hense, and B. Quack (2015), Marine sources of bromoform in the global open ocean – global patterns and emissions, Biogeosciences, 12(6), 1967–1981, {{doi|10.5194/bg-12-1967-2015}}. 
* Suzuki, A., Lim, L., Hiroi, T., & Takeuchi, T. (2006, March 20). Rapid determination of bromide in seawater samples by capillary ion chromatography using monolithic silica columns modified with cetyltrimethylammonium ion.

==References==
{{Reflist}}

{{GABAAR PAMs}}


[[Category:Anions]]
[[Category:Bromides| ]]
[[Category:GABAA receptor positive allosteric modulators]]

{{Sourceattribution|Bromide}}