Iodine

Iodine (from the Gr. Iodes, meaning "violet"), is a chemical element in the periodic table that has the symbol I and atomic number 53. It is required as a trace element for most living organisms. Chemically, iodine is the least reactive of the halogens, and the most electropositive metallic halogen. Iodine is primarily used in medicine, photography and in dyes.

1 Notable characteristics
2 Applications
3 History
4 Occurrence
5 Isotopes
6 Precautions
7 Compounds
8 References
9 External links
10 See also

Notable characteristics

Iodine is a bluish-black, lustrous solid that sublimes at standard temperatures into a blue-violet gas that has an irritating odor. This halogen also forms compounds with many elements, but is less active than the other members of its series and has some metallic-like properties. Iodine dissolves easily in chloroform, carbon tetrachloride, or carbon disulfide to form purple solutions (It is only slightly soluble in water). The deep blue color with starch solution is characteristic of the free element.

Applications

In areas where there is little iodine in the diet—typically remote inland areas and semi-arid equatorial climates where no marine foods are eaten—iodine deficiency gives rise to goitre, so called endemic goitre. In some such areas, this is now combated by the addition of small amounts of iodine to table salt in form of sodium iodide, potassium iodide, potassium iodate—this product is known as iodized salt. Iodine deficiency is the leading cause of mental retardation. Iodine deficency remains a serious problem that affects people around the globe. Other uses:

One of the halogens, it is an essential trace element; the thyroid hormones, thyroxine and triiodothyronine contain iodine.
Tincture of iodine (3% elemental iodine in water/ethanol base) is an essential component of any emergency survival kit, used both to disinfect wounds and to sanitize surface water for drinking (3 drops per liter, let stand for 30 minutes)
Iodine compounds are important in the field of organic chemistry and are very useful in medicine.
Iodides and thyroxine which contains iodine, are both used in internal medicine and, in combination with alcohol (as tincture of iodine) are used externally to disinfect wounds
Silver iodide is used in photography
Potassium iodide (KI) tablets can be given to people in a nuclear disaster area. KI prevents the body from absorbing the radioactive iodine and caesium (chemically related to potassium) produced at the disaster area; by flushing the body with an excess of a related less-harmful chemical, the body removes the toxic chemical along with the antidote.
Tungsten iodide is used to stabilize the filaments in light bulbs
Nitrogen triiodide is an explosive, too unstable to be used commercially, but is commonly used in college pranks
Iodine-131 is used as a tracer in medicine

History

Iodine (Gr. iodes meaning violet) was discovered by Barnard Courtois in 1811. He was the son of a manufacturer of saltpeter (potassium nitrate, a vital part of gunpowder). At the time France was at war and gunpowder was in great demand. Saltpeter was isolated from seaweed washed up on the coasts of Normandy and Brittany. To isolate the potassium nitrate, seaweed was burned and the ash then washed with water. The remaining waste was destroyed by adding sulfuric acid. One day Courtois added too much sulfuric acid and cloud of purple vapor rose. Courtois noted that the vapor crystallized on cold surfaces making dark crystals. Courtois suspected that this was a new element but lacked the money to pursue his observations.

However he gave samples to his friends, Charles Bernard Desormes (1777-1862) and Nicolas Clément (1779-1841) to continue research. He also gave some of the substance to Joseph Louis Gay-Lussac (1778–1850), a well-known chemist at that time, and to André-Marie Ampère (1775-1836). On November 29 1813 Dersormes and Clément made public Courtois’ discovery. They described the substance to a meeting of the Imperial Institute of France. On December 6 Gay-Lussac announced that the new substance was either an element or a compound of oxygen. Ampère had given some of his sample to Humphry Davy (1778-1829). Davy did some experiments on the substance and noted its similarity to chlorine. Davy sent a letter dated December 10 to the Royal Society of London stating that he had identified a new element. A large argument erupted between Davy and Gay-Lussac over who identified iodine first but both scientists acknowledged Bernard Courtois as the first to isolate the chemical element.

Occurrence

Iodine can be prepared in an ultrapure form through the reaction of potassium iodide with copper (II) sulfate. There are also several other methods of isolating this element. Although the element is actually quite rare, kelp and certain other plants have the ability to concentrate iodine, which helps introduce the element into the food chain as well as keeping its cost down.

Isotopes

There are 37 isotopes of iodine and only one, I-127, is stable. The artificial radioisotope I-131 (a beta emitter), also known as radioiodine which has a half-life of 8.0207 days, has been used in treating cancer and other pathologies of the thyroid glands. The most common compounds of iodine are the iodides of sodium and potassium (KI) and the iodates (KIO₃).

I-129 (half-life 15.7 million years) is a product of Xe-129 spallation in the atmosphere and uranium and plutonium fission, both in subsurface rocks and nuclear reactors. Nuclear processes, in particular nuclear fuel reprocessing and atmospheric nuclear weapons tests have now swamped the natural signal for this isotope. I-129 was used in rainwater studies following the Chernobyl accident. It also has been used as a ground-water tracer and as an indicator of nuclear waste dispersion into the natural environment.

In many ways, I-129 is similar to chlorine-36. It is a soluble halogen, fairly non-reactive, exists mainly as a non-sorbing anion, and is produced by cosmogenic, thermonuclear, and in-situ reactions. In hydrologic studies, I-129 concentrations are usually reported as the ratio of I-129 to total I (which is virtually all I-127). As is the case with Cl-36/Cl, I-129/I ratios in nature are quite small, 10^-14 to 10^-10 (peak thermonuclear I-129/I during the 1960s and 1970s reached about 10^-7). I-129 differs from Cl-36 in that its half-life is longer (15.7 vs. 0.301 million years), it is highly biophilic, and occurs in multiple ionic forms (commonly, I- and iodate) which have different chemical behaviors. This makes it fairly easy for I-129 to enter the biosphere as it becomes incorporated into vegetation, soil, milk, animal tissue, etc.

Excesses of Xe-129 in meteorites have been shown to result from decay of I-129. This was the first extinct radionuclide to be identified as present in the early solar system. Its decay is the basis of the I-Xe radiometric dating scheme, which covers the first 50 million years of solar system evolution.

Precautions

Direct contact with skin can cause lesions so care needs to be taken in handling iodine. Iodine vapor is very irritating to eyes and mucous membranes. The maximum allowable concentration of iodine in air should not exceed 1 mg/m³ (8-hour time-weighted average).

Compounds

Ammonium iodide (NH₄), Caesium iodide (CsI), Copper(I) iodide Hydroiodic acid (HI), Lithium iodide (LiI), Iodine cyanide (ICN), Iodine heptafluoride (IF₇), Iodine pentafluoride (IF₅), Iodic acid (HIO₃), Lead(II) iodide (PBI₂), Nitrogen triiodide (NI₃), Potassium iodide (KI), sodium iodide (NaI)

References

Los Alamos National Laboratory - Iodine

External links

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Contents

Notable characteristics

Applications

History

Occurrence

Isotopes

Precautions

Compounds

References

External links

See also

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