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|Section2={{Chembox Properties | Formula = | I=1 }} |Section3={{Chembox Thermochemistry | Entropy = 169.26 J K−1 mol−1 }} |Section4={{Chembox Related | OtherAnions = Fluoride Chloride Bromide }} }} An iodide ion is the ion I−. Compounds with iodine in formal oxidation state −1 are called iodides. This page is for the iodide ion and its salts, not organoiodine compounds. In everyday life, iodide is most commonly encountered as a component of iodized salt, which many governments mandate. Worldwide, iodine deficiency affects two billion people and is the leading preventable cause of intellectual disability.

Structure and characteristics of inorganic iodides

Iodide is one of the largest monatomic anions. It is assigned a radius of around 206 picometers. For comparison, the lighter halides are considerably smaller: bromide (196 pm), chloride (181 pm), and fluoride (133 pm). In part because of its size, iodide forms relatively weak bonds with most elements. Most iodide salts are soluble in water, but often less so than the related chlorides and bromides. Iodide, being large, is less hydrophilic compared to the smaller anions. One consequence of this is that sodium iodide is highly soluble in acetone, whereas sodium chloride is not. The low solubility of silver iodide and lead iodide reflects the covalent character of these metal iodides. A test for the presence of iodide ions is the formation of yellow precipitates of these compounds upon treatment of a solution of silver nitrate or lead(II) nitrate. Aqueous solutions of iodide salts dissolve iodine better than pure water. This effect is due to the formation of the triiodide ion, which is brown: I− + I2 ⇌

Redox, including antioxidant properties

Iodide salts are mild reducing agents and many react with oxygen to give iodine. A reducing agent is a chemical term for an antioxidant. Its antioxidant properties can be expressed quantitatively as a redox potential : I− ⇌  I2 + e− E° = −0.54 volts (versus SHE) Because iodide is easily oxidized, some enzymes readily convert it into electrophilic iodinating agents, as required for the biosynthesis of myriad iodide-containing natural products. Iodide can function as an antioxidant reducing species that can destroy reactive oxygen species such as hydrogen peroxide: 2 I− + peroxidase + H2O2 + tyrosine, histidine, lipid, etc. → iodo-compounds + H2O + 2 e− (antioxidants).

Representative iodides

Other oxyanions

Iodine can assume oxidation states of −1, +1, +3, +5, or +7. A number of neutral iodine oxides are also known.


External links

  • {{Cite web
| title = Seaweed use iodine as an antioxidant | work = Chemistry World blog | accessdate = 2008-05-15 | url = http://prospect.rsc.org/blogs/cw/?p=956 }}
  • {{Cite web
| title = Stressed seaweed contributes to cloudy coastal skies, study suggests | accessdate = 2008-05-15 | url = http://www.eurekalert.org/pub_releases/2008-05/uom-ssc050608.php }}
"green air" © 2007 - Ingo Malchow, Webdesign Neustrelitz
This article based upon the http://en.wikipedia.org/wiki/Iodide, the free encyclopaedia Wikipedia and is licensed under the GNU Free Documentation License.
Further informations available on the list of authors and history: http://en.wikipedia.org/w/index.php?title=Iodide&action=history
presented by: Ingo Malchow, Mirower Bogen 22, 17235 Neustrelitz, Germany