Iodine

53 tellurium ← iodine → xenon Br ↑ I ↓ At 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)
Vapor pressure (rhombic) P/Pa 1 10 100 1 k 10 k 100 k at T/K 260 282 309 342 381 457
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
Main article: Isotopes of iodine iso NA half-life DM DE (MeV) DP 127I 100% I is stable with 74 neutrons 129I syn 1.57×107y Beta- 0.194 129Xe 131I syn 8.02070 d Beta- 0.971 131Xe

Iodine (chemical symbol I, atomic number 53) is a nonmetal that belongs to a group of chemical elements known as halogens. At ordinary temperatures and pressures, iodine is a dark-gray/purple-black solid that readily sublimes (goes directly from the solid phase to the gas phase) to form a purple-pink gas with an irritating odor. Chemically, it is the least reactive of the halogens and the most electropositive halogen after astatine.

• It is required as a trace element for most living organisms.
• 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.

• from the Greek word iodes, meaning "violet."

Occurrence

Iodine occurs in nature in the form of iodide ions, chiefly in solution in seawater but also in some minerals and soils. Although the element is quite rare, it is concentrated in kelp and some other plants, which help introduce the element into the food chain and keep its cost down.

Discovery

Iodine was discovered in 1811 by the Frenchman Bernard Courtois, when he was working with his father to manufacture saltpeter (potassium nitrate). At the time, Napoleon's army was engaged in war and saltpeter, a key component of gunpowder, was in great demand. The process of producing saltpeter from French niter beds required sodium carbonate, which could be isolated from the ashes of burned seaweed. The remaining waste was destroyed by adding sulfuric acid. One day, Courtois added too much sulfuric acid and a cloud of purple vapor arose. The vapor condensed on cold surfaces to produce dark crystals.

Courtois performed a few experiments that led him to suspect that this was a new element, but he lacked the funds to pursue his observations. He therefore gave samples of the material to his friends, Charles Bernard Désormes (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, Désormes and Clément made public Courtois’ discovery, describing 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). When Davy experimented with the substance, he 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 major argument erupted between Davy and Gay-Lussac over who identified iodine first, but both scientists acknowledged Courtois as the first to isolate the substance.

Isolation

The element may be prepared in an ultrapure form by reacting potassium iodide with copper(II) sulfate. There are also several other methods of isolating this element.

Notable characteristics

In the periodic table, iodine is located in group 17 (former group 7A), the halogen family, between bromine and astatine. In addition, it lies in period 5, between tellurium and xenon. The molecular formula of iodine is I₂.

Iodine forms compounds with many elements, but it is less active than the other halogens and has some metallic properties. It is only slightly soluble in water, giving a yellow solution, but it readily dissolves in chloroform, carbon tetrachloride, or carbon disulphide to form purple solutions. The free element forms a deep blue complex with starch.

• When mixed with ammonia, iodine can form nitrogen triiodide, which is extremely sensitive and can explode unexpectedly.

• The most common compounds of iodine are the iodides of sodium and potassium (KI) and the iodates (KIO₃).

Isotopes

There are 37 isotopes of iodine, of which 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.

Excess quantities of stable ¹²⁹Xe in meteorites appear to have resulted from the decay of "primordial" ¹²⁹I produced newly by the supernovas that 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.

Compounds

A list of notable inorganic compounds of iodine is given below, listed in alphabetical order.

• Ammonium iodide (NH₄I)
• Cesium 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 iodate (KIO₃)
• Potassium iodide (KI)
• Sodium iodate (NaIO₃)
• Sodium iodide (NaI)

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.

Uses

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

Radioiodine and biology

Radioiodine and the thyroid

The artificial radioisotope ¹³¹I (which emits beta particles) has a half-life of 8.0207 days and has been used in treating cancer and other diseases 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).

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).

References

ISBN links support NWE through referral fees

• Los Alamos National Laboratory - Iodine
• WebElements.com - Iodine
• The History of Iodine
• 21 CFR 101.9 (c)(8)(iv) (Text PDF) — Food and Drug Administration nutritional facts label information for vitamins and minerals

External links

• ChemicalElements.com - Iodine
• who.int - WHO Global Database on Iodine Deficiency
• Network for Sustained Elimination of Iodine Deficiency
• Oxidizing Agents > Iodine

See also

• Iodised salt
• Chemical Oxygen Iodine Laser

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