Iodine

Iodine is a chemical element. Its atomic number (which is the number of protons in it) is 53, and its atomic mass (the number of protons plus neutrons) is 127. It is part of the Group 7 (halogens) on the periodic table of elements and its electronic configuration is 2,8,8,18,17. It is a nonmetal. 

Properties
Physical properties
Iodine is a gray-violet solid. It is also a nonmetal and a halogen. It sublimes when heated to make a bright purple vapor. This vapor is very irritating, although colorful. Iodine does not dissolve in water, but it does dissolve in a solution of an iodide in water. It dissolves easily in organic solvents.

Chemical properties
Iodine is at the bottom of the halogens and is the least reactive of the halogens. Iodine is an oxidizing agent, though not as strong as fluorine or bromine or chlorine. It can become the colorless iodide ion if another atom or molecule gives an iodine atom one of its electrons. This is called reduction. Iodine reacts with hydrogen sulfide or hydrazine to make hydriodic acid. Iodine dissolves in oleum to make a bright blue solution of I2+ and sulfur dioxide. This cation is also made when iodine reacts with antimony pentafluoride. 

Iodine reacts with bases to make hypoiodite and iodide. The hypoiodite disproportionates to iodate and iodide. Iodine reacts with nitric acid and chlorates to make iodates, too. Iodine reacts violently with some metals like aluminium to make large amounts of heat and toxic vapors.

Chemical compounds

Iodine forms chemical compounds in several oxidation states: -1, +1, +3, +5,  +7. This is typical of the halogens except for fluorine. -1 compounds as iodides are the most common. Most organic iodine compounds are in the -1 oxidation state. Most iodides are colorless or reddish-yellow and are weak reducing agents. +1 and +3 compounds are quite rare. They are both oxidizing agents. +1 compounds are found as iodine monochloride and hypoiodites. +3 compounds are found as iodine trichloride or iodine trifluoride. +5 compounds are found as iodates and iodic acid. Iodine pentafluoride and iodine pentoxide are iodine compounds in the +5 state. +5 and +7 compounds are all strong oxidizing agents. +7 compounds are normally found as periodates and periodic acid. Iodine heptafluoride is also in the +7 oxidation state.

Potassium iodide is one of the most common iodides. Iodides are normally colorless but turn yellow after being in air because the iodide is oxidized to iodine. Iodides react with acidified manganese dioxide or chlorine to make iodine. Iodine monochloride is a red or brown liquid. Iodine trichloride, as well as iodine trifluoride, are yellow solids. Iodates like potassium iodate are colorless solids similar to chlorates. Iodic acid is the only stable oxidizing halogen acid. Chloric acid, by comparison, is very unstable and dangerous. Iodine pentoxide is a white solid, and iodine pentafluoride is a yellowish liquid. Periodates, like sodium periodate, are colorless solids similar to perchlorates, although weaker oxidizing agents. They decompose to iodates when they are heated. Periodic acid is not as stable as iodic acid. Iodine heptafluoride is the only iodine(VII) compound.

History
It was discovered by Barnard Courtois in 1811. He burned seaweed to get sodium carbonate and reacted the residue with sulfuric acid. One time, he poured too much sulfuric acid on the seaweed and purple vapor came out which made dark blue-black crystals on a cold surface. He did not have enough money to see whether it was a new element or not. He gave samples to several chemists, who identified it as an element in 1813.

Occurrence
Iodine is too reactive to be in nature as an element. Iodine compounds are not common in the ground. There is a mineral called caliche that is found in dry deserts. It has iodate in it. Iodide is found in seawater. Kelp has quite large amounts of iodine in it. Iodine is one of the heaviest elements needed by living things. If people do not get enough of it, they get a disease known as goitre. Most salt has iodine in it, in the form of iodide or iodate. People can also take iodine pills to make sure they get enough.

Preparation
Kelp was burned and the iodine extracted from the ashes in the 18th and 19th century, but now it is not good enough for common use. Most iodine now is extracted from seawater. Chlorine is reacted with seawater to change the iodide to iodine. Air is blown over the water to evaporate the iodine. Then the iodine is reduced with sulfur dioxide to hydriodic acid. Then the hydriodic acid is oxidized with chlorine again to make iodine.  This process is repeated until the iodine is pure. 

A very pure form of iodine is made by reacting copper sulfate and potassium iodide. This makes copper(II) iodide which is not stable. It breaks down into copper(I) iodide and iodine, which is pure.

Uses
Iodine is used as a catalyst for making acetic acid. Animal feed also has iodine in it for nutrition. Iodine is used for a disinfectant. Tincture of iodine is a common iodine disinfectant. Iodine can be used to prevent thyroid cancer after a nuclear explosion. After a nuclear explosion, radioactive iodine fills the air. The radioactive iodine easily goes into the thyroid and causes cancer. When people take iodine tablets, it fills up the thyroid with non-radioactive iodine which prevents any radioactive iodine from coming in.

Iodine compounds can be used in X-rays because X-rays cannot go through them easily. Some metals are purified using their iodides. Silver iodide is used in photographic film and cloud seeding. Erythrosine is a food coloring that has iodine in it. Iodine can be used to make illegal drugs so it is restricted in some areas. Iodine can be used as a test for reducing agents. Iodine turns colorless when mixed with a reducing agent. Iodine also turns black when mixed with starch. Iodides become colored when mixed with oxidizing agents.

Safety
Iodine can irritate skin. Its vapors are irritating to the lungs. It is less toxic than the other halogens, though. 2-3 grams of iodine can kill. Iodides are only toxic in large amounts.

References

Category:Halogens