Bromide

A bromide is a chemical compound containing a bromide ion or ligand. This is a bromine atom with an ionic charge of −1 (Br); for example, in caesium bromide, caesium cations (Cs+) are electrically attracted to bromide anions (Br) to form the electrically neutral ionic compound CsBr. The term "bromide" can also refer to a bromine atom with an oxidation number of −1 in covalent compounds such as sulfur dibromide (SBr2).

Bromide
Br-.svg
Bromide ion.svg
Names
Systematic IUPAC name
Bromide[1]
Identifiers
3D model (JSmol)
3587179
ChEBI
ChEMBL
ChemSpider
14908
KEGG
UNII
Properties
Br
Molar mass 79.904 g mol−1
Conjugate acid Hydrogen bromide
Thermochemistry
82 J·mol−1·K−1[2]
−121 kJ·mol−1[2]
Pharmacology
N05CM11 (WHO)
Pharmacokinetics:
12 d
Related compounds
Other anions
Fluoride

Chloride
Iodide

Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references


Description of the FamilyEdit

Prior to describing Bromide ion, a short description of the family to which it belongs must be given. Group 17 consists of six members, which are, in order of increasing number of protons (atomic number), Fluorine (9F), Chlorine (17Cl), Bromine (35Br), Iodine (53I), Astatine (85At), and Tennessine (117Ts). These elements are nonmetal – not ductile and not as good as metals in terms of conducting heat and electricity. Members of Group 17 are also called the “Halogen Family” or “Salt-Producers” due to their ability to bond with Sodium (Na) to produce a sodium salt [Schneider et al., 2010]. Pure halogens do not occur in nature. But chemists can produce them in laboratory. All Halogens in their pure form are hazardous. Pure Chloride (Cl2) and Pure Fluorine (F2) are both poisonous gasses that readily react with water to produce strong acids (HCl and HF respectively). Pure Bromine (Br2), as mentioned, is a toxic liquid. Pure Iodine (I2), at standard temperature and pressure, is a hazardous solid that can cause serious eye injuries. Pure Astatine (At2) is also solid. But it does not last long in sold phase due to being highly unstable or radioactive [Knight & Schlager, 2002]. Like At2, Pure Tennessine (Ts2) is another radioactive substance that spontaneously converts to other elements of lighter atomic masses. Due to being extremely unstable, scientists cannot assign a definitive physical phase for Pure Tennessine (Ts2). But they expect it to be solid as well [Knight & Schlager, 2002]. The compounds of four Halogens, which are F, Cl, Br, and I, occur in nature. Because At has no stable isotope, its compounds can hardly be found in nature. Many scientists consider At to be the most rarest of all natural elements because it is only present as a part of natural radioactive decomposition reactions [Knight & Schlager, 2002]. Unlike all other Halogens, Ts is completely man-made. It is an artificial element that has no naturally produced compounds whatsoever. The most important similarity that Halogens share is that they all can participate in oxidation-reduction reaction (or “redox” reaction for short). Simply defined, redox reaction refers to a type of chemical reaction by which at least one electron is simultaneously exchanged between two chemical species. That chemical species that gains an electron is said to be reduced, whereas the other that losses an electron is said to be oxidized. All halogens have a strong tendency to participate in redox reaction since their outermost electron shell, consisting of seven electrons, needs one more electron to become full or octet. In redox reactions taking place in nature, Halogens act as oxidizing agents by gaining an electron (i.e. becoming reduced). This results in the formation of Halogen Anions. The anions of three Halogens, F-, Cl-, and Br-, are all major constituents of seawater with Cl- being the most abundant and conservative constituent at the concentration of 19.162 ppt [Mackenzie et al., 2011]. Dissolved Iodine (I) is only present as a trace metal (or minor constituent) in the forms of I- and IO3- [Millero, 2013].

Natural occurrenceEdit

Bromide is present in typical seawater (35 PSU) with a concentration of around 65 mg/L, which is around 0.2% of all dissolved salts. Seafoods and deep sea plants generally have high levels of bromide, while foods derived from land have variable amounts. Bromargyryte - natural, crystalline silver bromide - is the most common bromide mineral currently known, but is still very rare. Beside silver, bromine is sometimes found in minerals combined with mercury and copper.[3]

The Formation of Dissolved Bromide IonEdit

Like all other halogens, Br2 cannot exit in its pure form due to being very reactive. Like gaseous Cl2 and F2, Br2 as a liquid can readily react with water upon being dissolved:

                                            Br2 (l) + H2O (l) —> HOBr (aq) + HBr (g)

This results in the formation of Hypobromous acid (HOBr), and Hydrogen Bromide (HBr), which is a colorless gas. The solution is called “Bromine Water” and its product HBr gas can immediately react with water to become hydrobromic acid (HBr(aq)). Like Hydrochloric acid (HCl), HBr is a strong acid that easily dissociates by giving off its proton

                                                     HBr(aq)—> H⁺(aq) + Br⁻(aq)

The result is the formation of the Bromide ion (Br-), which was originally derived from the breakdown of pure Bromide Br2. Though, it should be kept in mind that this was not how the present Br- concentration in seawater accumulated since, prior to its construction by man in 1826, Br2 was nonexistent. But the equation above shows one example of how anthropogenic substances containing Br2 and HBr can cause an increase in sweater Br-. The other important consideration is the acidification impact caused by dissolving Br2 in a water sample. As more protons (H+) become available, the pH of the water will get lower (i.e. acidified).


Seawater BromideEdit

In terms of its behavior in seawater, the ion of Bromine, which is Bromide (Br-), falls in the conservative category. That is, its ratio to other major ions remains nearly constant in every part of the global ocean. Like other conservative elements, Br- concentration shows very minor changes with respect to depth and is not affect by any processes other than precipitation and evaporation. Despite being a major ion, not much research has been done on Br-. One can even argue that Br- has been the most neglected major ion in the history of Chemical Oceanography. For example, in their famous article regarding how major ions get removed from the global ocean on long time scales, Fred T. Mackenzie and Robert M. Garrels, who were the two eminent marine chemists of their time, never even touched upon Br- [1966]. Neither did John P. Riley and his colleagues in their great multivolume work entitled Chemical Oceanography mentioned any details about the processes that act as sources and sinks (i.e. removal) of Br- in seawater [1975]. All that was known about seawater Br- at that time was that it was conservative. Our current understanding of Br- has slightly improved compared to the past. But even today, there exists not even one encyclopedic article about seawater Br- in English language. Both editions of Encyclopedia of Ocean Sciences, the most authoritative encyclopedia on marine science, lack an article on sweater Br- ion [Steele et al., 2001] [Steele et al., 2009]. It is assumed that Br- has reached its present seawater concentration due to bromine salts present in the Earth’s crust, which haven continuously supplied Br- to seawater over a long geological timescale. The most common bromide salts are potassium bromide (KBr) and sodium bromide (NaBr). Once Br- is introduced to seawater, it will behave conservatively by not reacting with other elements. As a result, its concentration is always greater in sweater compared to Earth’s crust (which continuously loses bromide salts).

Extraction of bromine from seawaterEdit

In order to extract bromine from seawater, the same method utilized by Balard and Löwig can be used. First a seawater sample is tested for the presence of bromide compounds such as potassium bromide (KBr) and sodium bromide (NaBr). Next, that same seawater sample is titrated with aqueous chlorine (Cl2 (aq)) to produce pure bromine (Br2). The extracted Br2 will then be quantified by measuring its mass or volume. The chemical reaction can be equationally written as: Cl2 (aq) + 2Br2 (aq) —> 2Cl2 + Br2 (l) The science behind this chemical reaction is due to the fact Cl is a stronger halogen or oxidizing agent than bromine. In other words, when the two elements participate in the same chemical reaction, it is Cl that will act as the oxidizing agent of that reaction by oxidizing Br and becoming reduced (i.e. gaining an electron) [Magazinovic, 2004].


ChemistryEdit

One can test for a bromide ion by adding excess dilute HNO3 followed by dilute aqueous AgNO3 solution. The formation of creamy silver bromide precipitate confirms the existence of bromides.

Medical usesEdit

Bromide compounds, especially potassium bromide, were frequently used as sedatives in the 19th and early 20th century. Their use in over-the-counter sedatives and headache remedies (such as Bromo-Seltzer) in the United States extended to 1975, when bromides were withdrawn as ingredients, due to chronic toxicity.[4]

This use gave the word "bromide" its colloquial connotation of a boring cliché, a bit of conventional wisdom overused as a calming phrase, or verbal sedative.[5]

The bromide ion is antiepileptic, and bromide salts are still used as such, particularly in veterinary medicine. Bromide ion is excreted by the kidneys. The half-life of bromide in the human body (12 days) is long compared with many pharmaceuticals, making dosing difficult to adjust (a new dose may require several months to reach equilibrium). Bromide ion concentrations in the cerebrospinal fluid are about 30% of those in blood, and are strongly influenced by the body's chloride intake and metabolism.[6]

Since bromide is still used in veterinary medicine (particularly to treat seizures in dogs) 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 U.S., since there are no FDA-approved uses for bromide, and (as noted) it is no longer available in over-the-counter sedatives. Therapeutic bromide levels are measured in European countries like Germany, where bromide is still used therapeutically in human epilepsy.

Chronic toxicity from bromide can result in bromism, a syndrome with multiple neurological symptoms. Bromide toxicity can also cause a type of skin eruption. See potassium bromide.

Lithium bromide was used as a sedative beginning in the early 1900s, but it fell into disfavor in the 1940s, possibly 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 lithium chloride).[7] Like lithium carbonate and lithium chloride it was used as treatment for bipolar disorder.

It has been said that during World War I, British soldiers were given bromide to curb their sexual urges,[8] although this is not well supported by documentation, and has been disputed as an urban myth, because the sedative effects of bromide would have hampered military performance. Lord Dunsany mentions a soldier being given bromide as a sedative for nervous exhaustion and overwork in his play Fame and the Poet (1919).[9]

There are more substantiated reports that bromide was used in the food served at some concentration camps during the Holocaust. This was apparently done in an effort to both chemically restrain the interned and prevent menstruation in women.[10]

In biologyEdit

According to one study, bromine (as bromide) is an essential cofactor in the peroxidasin catalysis of sulfilimine crosslinks in collagen IV. This post-translational modification occurs in all animals, and bromine is an essential trace element for humans.[11]

Bromide is needed by eosinophils (white blood cells of the granulocyte class, specialized for dealing with multi-cellular parasites), which use it to generate antiparasitic brominating compounds such as hypobromite, by the action of eosinophil peroxidase, a haloperoxidase enzyme which is able to use chloride, but preferentially uses bromide when available.[12] Other than its role in collagen IV production and its facultative use in eosinophils by the body, bromide is not known in other cases necessary for animal life, as its functions may generally be replaced (though in some cases not as well) by chloride. Land plants do not use bromide.

Bromide salts are also sometimes used in hot tubs and spas as mild germicidal agents, using the action of an added oxidizing agent to generate in situ hypobromite, in a similar fashion to the peroxidase in eosinophils.

Bromide is perhaps a minor necessary nutrient for collagen IV-producing animals in the sea. However, a few sea animals, such as Murex snails, use bromide to make organic compounds. Bromide ion is also heavily concentrated by some species of ocean algae, which construct methyl bromide and a great number of bromoorganic compounds with it, using the unusual enzymes called vanadium bromoperoxidases to do these reactions.

The average concentration of bromide in human blood in Queensland, Australia is 5.3±1.4 mg/L and varies with age and gender.[13] Much higher levels may indicate exposure to brominated chemicals (e.g. methyl bromide). However, since bromide occurs in relatively high concentration in seawater and many types of seafood, bromide concentrations in the blood are heavily influenced by seafood contributions to the diet.

Reference and Further ReadingEdit

Encyclopedic Articles and Books

Christe, K., and S. Schneider (2020), Bromine, Encyclopedia 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):

Rattley, M. (2012), Ambiguous bromine, Nature Chemistry, 4(6), 512–512, doi:10.1038/nchem.1361.

Wisniak, J. (2002), The history of bromine from discovery to commodity, NOPR.

Peer-Reviewed Journal Articles for Bromide (Br-)

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.


ReferencesEdit

  1. ^ "Bromide – PubChem Public Chemical Database". The PubChem Project. USA: National Center for Biotechnology Information. Archived from the original on 2012-11-03.
  2. ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. ISBN 978-0-618-94690-7.
  3. ^ "Mindat.org - Mines, Minerals and More". www.mindat.org. Archived from the original on 2 March 2001. Retrieved 29 April 2018.
  4. ^ Adams, Samuel Hopkins (1905). The Great American fraud. Press of the American Medical Association..
  5. ^ "the definition of bromide". Dictionary.com. Archived from the original on 24 December 2016. Retrieved 21 December 2016.
  6. ^ Goodman, L. S. and Gilman, A. (eds.) (1970) "Hypnotics and Sedatives", p. 121 in Chapter 10 in The Biological Basis of Therapeutics, Fourth Edition, The MacMillan Co., London.
  7. ^ Bipolar disorder. webmd.com
  8. ^ Tanaka, Yuki (2002) Japan's Comfort Women: Sexual slavery and prostitution during World War II and the US Occupation, Routledge, p. 175. ISBN 0415194008.
  9. ^ Lord Dunsany (August 1919). "Fame and the Poet". The Atlantic Monthly: 175–183. Archived from the original on 2015-12-22.
  10. ^ Jackson, “The Coming of Age” in Women and the Holocaust, eds Rittter & Roth, p. 80.
  11. ^ McCall AS, Cummings CF, Bhave G, Vanacore R, Page-McCaw A, Hudson BG (2014). "Bromine Is an Essential Trace Element for Assembly of Collagen IV Scaffolds in Tissue Development and Architecture". Cell. 157 (6): 1380–92. doi:10.1016/j.cell.2014.05.009. PMC 4144415. PMID 24906154.
  12. ^ Mayeno, AN; Curran, AJ; Roberts, RL; Foote, CS (1989). "Eosinophils preferentially use bromide to generate halogenating agents". The Journal of Biological Chemistry. 264 (10): 5660–8. PMID 2538427.
  13. ^ Olszowy, HA; Rossiter, J; Hegarty, J; Geoghegan, P (1998). "Background levels of bromide in human blood". Journal of Analytical Toxicology. 22 (3): 225–30. doi:10.1093/jat/22.3.225. PMID 9602940.