Iodine: The Element

There is one experiment which I always like to try, because it proves something whichever way it goes. A solution of iodine in water is shaken with bone-black, filtered and tested with starch paste. If the colorless solution does not turn the starch blue, the experiment shows how completely charcoal extracts iodine from aqueous solution. If the starch turns blue, the experiment shows that the solution, though apparently colorless, still contains iodine which can be detected by means of a sensitive starch test.

These were the words of Wilder Dwight Bancroft, regarding the element iodine. The element does not occur in a free state in nature and when freed from its compounds, it forms diatomic molecules (I₂). Here are some of the properties and uses of iodine.

History
The name of the element originates from the Greek word Iodes, which means 'violet'. The element iodine was discovered by Bernard Courtois in 1811. His father was a manufacturer of saltpeter, which was a vital part of gunpowder. At the time of the Napoleonic Wars, France was at war and saltpeter was in great demand. Production of saltpeter required sodium carbonate, which was obtained by burning seaweed and extracting it from the ash washed in water. The remaining waste was usually destroyed with sulfuric acid. One day Courtois added too much sulfuric acid to the waste to be destroyed. He saw a cloud of purple vapor, which crystallized on cold surfaces, making dark crystals. As he was short of finances, he gave the samples to Charles Bernard Desormes and Nicolas Clément, who continued the research. On November 29, 1813, Dersormes and Clément published the discovery of Courtois.

Physical Properties
Iodine is a bluish-black, lustrous solid, which volatilizes at ordinary temperatures, into a blue-violet gas. Its melting point is 113.5°C, boiling point is 184.35°C, and specific gravity is 4.93, for its solid state at 20°C. Its gas density is 11.27 g/l and atomic number is 53. This halogen element forms compounds with many elements, but is less reactive than the other members of its Group VII (halogens) and has some metallic light reflectance. The element does not occur in its free state in nature and is slightly soluble in water, forming the tri-iodide anion I₃^-. Iodine also readily dissolves in carbon tetrachloride, chloroform, and carbon disulfide, forming purple solutions. The direct contact of iodine with skin can cause lesions and its vapor is highly irritating to the eyes and mucous membranes. Iodine forms various compounds with many elements, but as compared to other halogens, it is less reactive. This property is shown by the displacement of iodine for iodides, by other halogens. There are 37 isotopes of iodine, out of which, only ¹²⁷I is stable.

Uses
The chronic deficiency of iodine may cause thyroid gland dysfunction and various neurological, gastrointestinal and skin abnormalities. Its compounds are basically used in medicine, photography and dyes. Iodine also forms starch, which is an intense blue complex compound. Starch may be used for various applications, like, iodometry and iodine clock reaction. Being a heavy element, iodine is quite radio-opaque and thus can be used in medicine as the X-ray radio-contrast agent, for intravenous injection. Some of the radioactive isotopes can also be used for the treatment of thyroid cancer, as it can selectively damage the growing thyroid cancer cells.

Iodine forms many compounds, like, potassium iodide, sodium iodide, iodic acid and periodic acid. Potassium iodide is of commercial use and sodium iodide is especially useful in the Finkelstein reaction. Iodic acid and its salts are strong oxidizers, whereas periodic acid cleaves vicinal diols along the C-C bond to give aldehyde fragments. Iodine should always be handled with care because of its adverse effects on the skin.