The thirty-fifth element of the periodic table has properties that are immediately evident if you have a quality periodic table at hand. Bromine (pronounced as an angry younger sister might declare while pointing: “Bro Mean!”) is both a member of the reactive halogens family (Group 17) and one of only two naturally occurring elements that are liquids at room temperature. (The other is mercury.) Besides its basic litany of facts, you should be aware of some of the more common reactions and applications of bromine in today’s world.
Readers seeking a simple table of basic data might want some of these values for bromine, taken directly from the 82nd edition of the CRC Handbook:
Formula: Br2
Formula Mass: 159.808 g/mol (amu) – divide by two for a Br atom: 79.904 g/mol
Appearance: red liquid
Melting Point: -7.2ºC
Boiling Point: 58.8ºC
Density: 3.1028 g/mL
Solubility: slightly soluble in water
Isotopes: bromine-79 (50.69%) and bromine-81 (49.31%)
As you may have seen from its formula, bromine is one of only seven elements which are commonly found in a diatomic (two atoms bound together) state. This is common behavior among the halogens, which all have seven electrons in their outermost energy level. A shared bond gives them access to the last electron they need to satisfy the “octet rule.” Although this bond completes an octet for each, bromine is a rather electronegative element. Often listed with a value of 3.0 (on a four-point scale), bromine likes to draw electron density toward itself, making it strongly reactive. This also makes it a hazardous chemical, as it reacts with skin or mucous membranes, causing irritation or painful burns.
One reaction of bromine is well known to any student of organic chemistry. Bromine will add to the carbons at either end of a double bond in what is commonly called “bromination of an alkene.” Electron-hungry bromine readily hooks up with the electron-rich double bond between two carbons. If the molecule has a rigid structure, this reaction results in two “mirror-image” (enantiomer) products. The carbon-bromine bonds are also polar, since carbon has a lower electronegativity. Bromination by various methods is possible, and is important to modern-day consumers since brominated materials often serve as fire-retardants.
Bromine reacts with water to produce the hypobromite ion (BrO-). This is analogous to its cousin, the hypochlorite ion, commonly found in household bleach. With similar sterilizing properties, bromine is sometimes used in water as a sanitizer. Sodium hypobromite also finds fame with chemistry students in its incarnation as a “Chemistry Cat” joke.
Alkali metals (Group 1 on the periodic table) react violently with bromine to form stable salts (sodium bromide, potassium bromide, etc.). Hydrogen also reacts with bromine, making hydrogen bromide (HBr). In water, HBr behaves as a strong acid, dissociating completely into hydronium ions and bromide ions. In its aqueous (dissolved in water) form, HBr is referred to as hydrobromic acid. (Again, bromine has a more familiar chlorine cousin–hydrochloric acid–found in stomach acid and drain cleaners.)
Bromine undergoes a variety of other interesting reactions, and finds other applications, including pharmaceuticals and agriculture. Though not the most familiar element in daily speech, chances are that it plays a part in your daily life. And now, Br-illiant reader, if this Br-ief article has stimulated your Br-ain, Br-avo!
