Iodine

Vapor pressure (rhombic)
53	tellurium ← iodine → xenon
	periodic table
General
Name, Symbol, Number	iodine, I, 53
Chemical series	halogens
Group, Period, Block	17, 5, p
Appearance	violet-dark gray, lustrous ;
Atomic mass	126.90447(3) g/mol
Electron configuration	[Kr] 4d10 5s2 5p5
Electrons per shell	2, 8, 18, 18, 7
Physical properties
Phase	solid
Density (near r.t.)	4.933 g/cm³
Melting point	386.85 K; (113.7 °C, 236.66 °F)
Boiling point	457.4 K; (184.3 °C, 363.7 °F)
Critical point	819 K, 11.7 MPa
Heat of fusion	(I2) 15.52 kJ/mol
Heat of vaporization	(I2) 41.57 kJ/mol
Heat capacity	(25 °C) (I2) 54.44 J/(mol·K)
Atomic properties
Crystal structure	orthorhombic
Oxidation states	±1, 5, 7; (strongly acidic oxide)
Electronegativity	2.66 (Pauling scale)
Ionization energies	1st: 1008.4 kJ/mol
	2nd: 1845.9 kJ/mol
	3rd: 3180 kJ/mol
Atomic radius	140 pm
Atomic radius (calc.)	115 pm
Covalent radius	133 pm
Van der Waals radius	198 pm
Miscellaneous
Magnetic ordering	nonmagnetic
Electrical resistivity	(0 °C) 1.3×107 Ω·m
Thermal conductivity	(300 K) 0.449 W/(m·K)
Bulk modulus	7.7 GPa
CAS registry number	7553-56-2
Notable isotopes

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Iodine (from the Greek word 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 halogen after astatine. Iodine is primarily used in medicine, photography and in dyes.

As with all other halogens (members of Group VII in the Periodic Table), iodine forms diatomic molecules, and hence, has the molecular formula of I₂.

Occurrence on earth

Iodine occurs in the environment chiefly as dissolved iodide in seawater, though it is also found in some minerals and soils. The element may 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.

Uses

Iodine is used in pharmaceuticals, antiseptics, medicine, food supplements, dyes, catalysts and photography.

Isotopes

There are 37 isotopes of iodine and only one, ¹²⁷I, is stable.

In many ways, ¹²⁹I is similar to ³⁶Cl. 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, ¹²⁹I concentrations are usually reported as the ratio of ¹²⁹I to total I (which is virtually all ¹²⁷I). As is the case with ³⁶Cl/Cl, ¹²⁹I/I ratios in nature are quite small, 10⁻¹⁴ to 10⁻¹⁰ (peak thermonuclear ¹²⁹I/I during the 1960s and 1970s reached about 10⁻⁷). ¹²⁹I differs from ³⁶Cl 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 IO₃⁻) which have different chemical behaviors. This makes it fairly easy for ¹²⁹I to enter the biosphere as it becomes incorporated into vegetation, soil, milk, animal tissue, etc.

Excesses of stable ¹²⁹Xe in meteorites have been shown to result from decay of "primordial" ¹²⁹I produced newly by the supernovas which created the dust and gas from which the solar system formed. ¹²⁹I 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.

Effects of various radioiodine isotopes in biology are discussed below.

Notable characteristics

Iodine is a dark-gray/purple-black solid that sublimes at standard temperatures into a purple-pink gas that has an irritating odor. This halogen forms compounds with many elements, but is less active than the other members of its Group VII (halogens) and has some metallic-like properties. Iodine dissolves easily in chloroform, carbon tetrachloride, or carbon disulphide to form purple solutions (It is only slightly soluble in water, giving a yellow solution). The deep blue color of starch-iodine complexes is produced only by the free element.

History

Iodine was discovered by Barnard Courtois in 1811. He was born to a manufacturer of saltpeter (potassium nitrate, a vital part of gunpowder). At the time France was at war, saltpeter, a component of gunpowder, was in great demand. Saltpeter produced from French niter beds required sodium carbonate, which could be isolated from seaweed washed up on the coasts of Normandy and Brittany. To isolate the sodium carbonate, 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 a 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 29 November 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 Barnard Courtois as the first to isolate the chemical element.

Notable inorganic iodine compounds

Ammonium iodide (NH₄I)
Caesium iodide (CsI)
Copper(I) iodide (CuI)
Hydroiodic acid (HI)
Iodic acid (HIO₃)
Iodine cyanide (ICN)
Iodine heptafluoride (IF₇)
Iodine pentafluoride (IF₅)
Lead(II) iodide (PbI₂)
Lithium iodide (LiI)
Nitrogen triiodide (NI₃)
Potassium iodide (KI)
Sodium iodide (NaI)

See also iodine compounds

Stable iodine in biology

One of the halogens, iodine is an essential trace element; the thyroid hormones, thyroxine and triiodothyronine contain iodine.

Iodine has a single known role in biology: it is an essential trace element since the thyroid hormones, thyroxine (T4) and triiodothyronine (T3) contain iodine. These are made from addition condensation products of the amino acid tyrosine, and are stored prior to release in a protein-like molecule called thryroglobulin. T4 and T3 contain four and three atoms of iodine per molecule, respectively. The thyroid actively absorbs elemental iodine from the blood to make and release these hormones into the blood, actions which are regulated by a second hormone TSH from the pituitary. Thyroid hormones are phylogenetically very old molecules which are sythesized by most multicellular organisms, and which even have some effect on unicellular organisms.

Thyroid hormones play a very basic role in biology, acting on mitochondria to regulate metabolism. T4 acts largely as a precursor to T3, which is (with some minor exceptions) the biologically active hormone.

Dietary intake

The United States Food and Drug Administration recommends (21 CFR 101.9 (c)(8)(iv)) 150 micrograms of iodine per day for both men and women. This is necessary for proper production of thyroid hormone. Natural sources of iodine include seaweed, such as kelp and seafood. [1] Salt for human consumption is often enriched with iodine and is referred to as iodized salt.

Iodine deficiency

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 goiter, so called endemic goiter. In some such areas, this is now combatted 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 compounds have also been added to other foodstuffs, such as flour, in areas of deficiency. Iodine deficiency is the leading cause of preventable mental retardation. This is caused by lack of thyroid hormone in the infant. Iodine deficiency remains a serious problem that affects people around the globe.

Toxicity of Iodine

Excess iodine has symptoms similar to those of iodine deficiency. Commonly encountered symptoms are abnormal growth of the thyroid gland and disorders in functioning and growth of the organism as a whole.

Elemental iodine, I₂, is deadly poison if taken in larger amounts; if 2-3 grams of it is comsumed, it is fatal to humans.

Iodides are similar in toxicity to bromides.

Radioiodine and biology

Radioiodine and the thyroid

The artificial radioisotope ¹³¹I (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. ¹²³I is the radioisotope most often used in nuclear imaging of the kidney and thyroid as well as thyroid uptake scans (used for the evaluation of Grave's disease). The most common compounds of iodine are the iodides of sodium and potassium (KI) and the iodates (KIO₃).

¹²⁹I (half-life 15.7 million years) is a product of ¹³⁰Xe 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. ¹²⁹I 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.

If humans are exposed to radioactive iodine, the thyroid gland will absorb it as if it were non-radioactive iodine, leading to elevated chances of thyroid cancer. Isotopes with shorter half-lifes such as ¹³¹I present a greater risk than those with longer half-lives since they generate more radiation per unit of time. Taking large amounts of regular iodine will saturate the thyroid and prevent uptake. Iodine pills are sometimes distributed to persons living close to nuclear establishments, for use in case of accidents that could lead to releases of radioactive iodine.

Iodine-123 and iodine-125 are used in medicine as tracers for imaging and evaluating the function of the thyroid.
Iodine-131 is used in medicine for treatment of thyroid cancer and Grave's disease.
Uncombined (elemental) iodine is mildly toxic to all living things.
Potassium iodide (KI tablets, or "SSKI" = "Super-Saturated KI" liquid drops) can be given to people in a nuclear disaster area when fission has taken place, to flush out the radioactive iodine-131 fission product. The half-life of iodine-131 is only eight days, so the treatment would need to continue only a couple of weeks. In cases of leakage of certain nuclear materials without fission, or certain types of dirty bomb made with other than radioiodine, this precaution would be of no avail.

Non-hormone-related applications of 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). Alcohol-free iodine solutions such as Lugol's iodine, as well as other free iodine-providing antiseptics iodophors, are also available as effective elemental iodine sources for this purpose.
Iodine compounds are important in the field of organic chemistry and are very useful in medicine.
Silver iodide is used in photography.
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.

Precautions for stable iodine

Direct contact with skin can cause lesions, so it should be handled with care. Iodine vapor is very irritating to the eye and to mucous membranes. Concentration of iodine in the air should not exceed 1 mg/m³ (eight-hour time-weighted average). When mixed with ammonia, it can form nitrogen triiodide which is extremely sensitive and can explode unexpectedly.

References
ISBN links support NWE through referral fees

Los Alamos National Laboratory - Iodine
WebElements.com - Iodine
21 CFR 101.9 (c)(8)(iv) (Text PDF) — FDA nutritional facts label information for vitamins and minerals

External links

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