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I first encountered iodine as a young child during the second World War, when my mother applied a mysterious purple-brown solution--with quite a sting--to disinfect a bruised, scraped knee. Little did I know then what iodine was or that nearly a half century later it would become one of my favorite elements.

My next encounter with iodine occurred as a teenager in Hungary, when I acquired it in the neighborhood pharmacy as part of my extensive home chemistry set and mixed it in small amounts with self-made smokeless gunpowder to impress my friends with its purple vapor.

Iodine, element 53 with a relative atomic mass of 126.90447, was first isolated by Bernard Courtois in 1811 from the ash of seaweed (by treating kelp with H2SO4). It was named by J. L. Gay Lussac in 1813, and its name derives from the Greek word iodes, meaning "violet-colored," reflecting the characteristic lustrous, deep purple color of resublimed crystalline iodine as well as the color of its vapor. Potassium iodide (KI) was used as a remedy for goiter (Derbyshire neck), an enlargement of the thyroid gland, as early as 1819. The thyroid is responsible for the production of thyroxine, a metabolism-regulating hormone. Iodine is an essential trace element for humans and plays an important role in many biological organisms. In modern times, KI is recommended for the treatment of radiation poisoning.

Early sources of iodine were the saltpeter deposits in Chile, whereas contemporary sources also include natural brines and salt wells. It is generally liberated from brine via chlorine gas. Annual production of I2 exceeds 10,000 tons, Japan being the dominant producer. It is the heaviest of the common halogens, and its isotopes range in mass from 117 to 139; the natural isotope, 127, occurs in 100% abundance. Radioactive isotopes include 124, 125, 128, 131, and 132 and can be used as radioactive tracer elements. It readily dissolves in such organic solvents as chloroform, carbon tetrachloride, ethanol, benzene, and ethyl ether to form beautiful purple solutions, but it is only slightly soluble in water.

Iodine readily forms compounds with most other elements in the periodic table. It occurs most commonly in monovalent form with an oxidation state of –1. It forms relatively weak bonds with first-row elements, including carbon, the typical C–I bond dissociation energy being only about 55 kcal per mole. Organoiodine compounds have been used since the mid-1800s, notably in Wurtz coupling reactions, the Williamson ether synthesis, and Hofmann's alkylation of amines.

Currently, the most important and common use of organoiodine compounds involves various metal-mediated cross-coupling reactions where they serve as premier electrophilic partners in Heck, Negishi, Suzuki, Sonogashira, Stille, and similar cross-coupling protocols. These metal-catalyzed cross-coupling reactions are extensively employed in preparative organic chemistry, the synthesis of complex natural products, and the manufacture of drugs, as well as in supramolecular and materials chemistry.

Because iodine is the largest, least electronegative, and most polarizable of the common halogens, it is also capable of forming stable polycoordinate high-valent (with a value of up to 7, IF7) compounds. The most common polyvalent organic iodine compounds are I(III) and I(V) species. The first stable polyvalent organic iodine compound, the trivalent PhICl2, was prepared by the German chemist C. H. C. Willgerodt in 1886.

In the 1980s, we and others developed alkynyliodonium salts, RCCI+PhX (X=OTs, OTf, BF4, etcetera), the newest member of the family of polyvalent organoiodine compounds, which may serve as electrophilic acetylene equivalents. This has engendered a renaissance in polyvalent organoiodine chemistry. Arguably, the most useful and widely employed contemporary polyvalent organoiodine compound is the I(V) Dess-Martin periodinane that has emerged as the reagent of choice for the oxidation of primary and secondary alcohols to aldehydes and ketones, respectively. Because of its ready availability; its convenience of use; its unique, selective oxidizing property; and, most importantly, its functional group tolerance, the Dess-Martin periodinane is widely employed in the synthesis of complex natural products of biological and medicinal interest.

Among the more common, everyday uses of iodine are the following: in halogen lamps, as a salt additive (to prevent goiter), and in ink pigments. Tincture of iodine is used as a topical antiseptic to kill bacteria. Silver iodide is used in the preparation of some photographic films.

Peter J. Stang is distinguished professor of chemistry and dean of the College of Science at the University of Utah. He is a member of the National Academy of Sciences and a fellow of the American Academy of Arts & Sciences. Since 2002, he has been the editor of the Journal of the American Chemical Society.


Chemical & Engineering News
Copyright © 2003 American Chemical Society

Name: From the Greek iodes, meaning violet-colored.
Atomic mass: 126.90.
History: Discovered in 1811 by French chemist Bernard Courtois.
Occurrence: Found in seaweed and brine wells. It is also found in Chilean saltpeter, caliche, old salt brines, and salt wells.
Appearance: Lustrous, violet-dark gray, nonmetal solid.
Behavior: Forms compounds with most elements. Volatilizes at ambient temperatures into a blue-violet gas with an irritating odor. It is only slightly soluble in water. Pure iodine is highly poisonous.
Uses: Essential to many species, including humans. It is part of thyroxine, a hormone produced by the thyroid gland; a lack of iodine causes goiter.

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