Chemical & Chemical Engineering News (80th Anniversary Issue), Vol. 81, No. 36, 2003, Sept. Edited by X. Lu Introduction

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.

IODINE AT A GLANCE

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

ASTATINE

DALE R. CORSON, CORNELL UNIVERSITY

In 1938-40, I served two years as a postdoctoral fellow in the Radiation Laboratory of the University of California, Berkeley, helping construct a new 60-inch cyclotron that produced 32-MeV -particles.

With construction nearing completion, the laboratory staff met in a special session one evening in the fall of 1939 to discuss experiments designed to exploit the newly available high-energy particles. In the course of the discussion, Emilio Segrè pointed out the obvious fact that once one looked at the Periodic Table of the Elements, adding an -particle to a bismuth nucleus of atomic number 83 could produce a nucleus of element 85, which had been until then missing in the periodic table.

The next day, Robert Cornog, a graduate student colleague who possessed a small piece of bismuth, and I bombarded the shiny metal for a short time with the 32-MeV -particles. I had a nearby laboratory with instrumentation I had built, anticipating the new research opportunities. At the conclusion of the short bombardment, we carried the bismuth to my laboratory and placed it in front of an ionization chamber connected to a linear amplifier. The oscilloscope screen recording the output of the amplifier was alive with large pulses, characteristic of -particles. We clearly had something of interest.

After some preliminary experiments seeking to sort out the radioactivities we had produced, it was clear that I was going to need help if I were to pursue the effort with vigor. Cornog was already occupied with the discovery of tritium, working with Luis Alvarez. Kenneth MacKenzie was a graduate student ready for a dissertation project, and I invited him to help me. He proved to be an ingenious and energetic colleague.

We identified the many radiations we had produced: two groups of a-particles, -rays, X-rays, low-energy electrons, and some energetic positrons. Curiously, all these radiations possessed the same 7.5-hour half-life, except for the positrons. We later showed that these came from copper contamination in the bismuth.

For a reason I do not know, our 7.5-hour half-life is now recorded in the literature as 7.2 hours. The precision of our measurement did not permit this large an experimental error. It probably has to do with the volatility of the radioactive product we were seeking to identify.

We applied all the ingenuity and imagination we could muster, seeking to determine the origins of all the observed radiations. We finally arrived at the correct radioactive decay scheme after a lengthy series of experiments designed to check one possibility after another. Fortunately, bismuth has only a single stable isotope (209). Through an (-2n) process--that is, an -particle enters the bismuth nucleus and two neutrons leave it--the element 85 nucleus with atomic mass number 211 is produced, with a half-life of 7.5 hours. The ²¹¹85 nucleus decays in two ways: 40% of the time, it decays by capturing a K-shell electron to go to ²¹¹Po, which then decays with a very short half-life, emitting an a-particle, to go to stable ²⁰⁷Pb; 60% of the time, ²¹¹85 decays with emission of an -particle followed by a process to go to ²⁰⁷Pb. We failed to identify the process.

Identification of the nucleus emitting the X-rays following the K-capture process was vital in nailing down the disintegration scheme. We employed the critical absorption edges in tungsten and platinum absorbers to show that the X-rays are polonium X-rays, as demanded by our decay scheme.

RINGING IN A 1939 photograph of the 60-inch cyclotron group at the University of California, Berkeley. (From left) Don Cooksey, Corson, Ernest O. Lawrence, Bob Thornton, John Backus, W. W. Salisbury, (above) Alvarez, and Edwin McMillan.

Segrè's help was essential in determining the general chemical properties of 85. Some of the properties are similar to those of iodine, its lower homolog. It also exhibits metallic properties, more like its metallic neighbors Po and Bi.

There was some investigation of element 85's behavior in guinea pigs. I collaborated with J. G. Hamilton, an M.D., in one experiment. In this study, element 85 was injected in the animal, and a few hours later tissue samples were examined for 85 activity. The element was concentrated in the tiny thyroid gland in much the same way iodine would be, establishing the physiological similarity to iodine.

The literature now records some 20 different isotopes of 85. Small amounts of some naturally occurring 85 isotopes have been found, but probably no more than a few grams total in the entire Earth's crust. I have appeared in the "Guinness Book of World Records" for having discovered the rarest substance on Earth.

In a note to Nature in 1947, MacKenzie, Segrè, and I proposed the name astatine (from the Greek word meaning unstable) for this element, in keeping with the naming of the other halogens where the name relates to some property of the substance.

Dale R. Corson is a professor of physics emeritus and president emeritus of Cornell University. He is a Public Welfare Medalist of the National Academy of Sciences and a Bueche Medalist of the National Academy of Engineering.

ASTATINE AT A GLANCE

Name: From the Greek astatos, unstable.

Atomic mass: (210)

History: First produced in 1940 by Dale R. Corson, K. R. MacKenzie, and Emilio Segrè at the University of California, Berkeley, by bombarding a bismuth isotope with particles.

Occurrence: The rarest of the 92 naturally occurring elements; less than 30 g exists on Earth at any one time as a natural by-product of uranium and thorium decay.

Appearance: Solid nonmetal of unknown color at room temperature.

Behavior: Highly radioactive; decays very quickly.

THE NOBLE GASES