IODINE

Iodine is a chemical element with symbol I and atomic number 53. The heaviest of the stable halogens, it exists as a lustrous, purple-black non-metallic solid at standard conditions that melts to form a deep violet liquid at 114 degrees Celsius, and boils to a violet gas at 184 degrees Celsius. The element was discovered by the French chemist Bernard Courtois in 1811. It was named two years later by Joseph Louis Gay-Lussac from this property.
Iodine occurs in many oxidation states, including iodide , iodate, and the various periodate anions. It is the least abundant of the stable halogens, being the sixty-first most abundant element. It is the heaviest essential mineral nutrient. Iodine is essential in the synthesis of thyroid hormones.[4] Iodine deficiency affects about two billion people and is the leading preventable cause of intellectual disabilities.
The dominant producers of iodine today are Chile and Japan. Iodine and its compounds are primarily used in nutrition. Due to its high atomic number and ease of attachment to organic compounds, it has also found favour as a non-toxic radiocontrast material. Because of the specificity of its uptake by the human body, radioactive isotopes of iodine can also be used to treat thyroid cancer. Iodine is also used as a catalyst in the industrial production of acetic acid and some polymers.

• Appearance: lustrous metallic gray, violet as a gas
• Standard Atomic Weight: 126.904
• Atomic Number(Z): 53
• Group: group 17 (halogens)
• Period: period 5
• Block: p block
• Element Category: reactive non-metal
• PHASE AT STP : solid
• Melting Point : 386.85 K [113.7 ° C]
• Boiling Point : 457.4 K [184.3 ° C]
• Density : 4.933 g/ml
• Heat of Vaporization : 41.57 kj/mol
• Covalent Radius : 139+-3pm
• Crystal Structure : orthorhombic
• Discovery :  Bernard Courtosis (1811)

                   HISTORY
In 1811, iodine was discovered by French chemist Bernard Courtois who was born to a manufacturer of saltpetre (an essential component of gunpowder). At the time of the Napoleonic Wars, saltpetre was in great demand in France. Saltpetre produced from French nitre beds required sodium carbonate, which could be isolated from seaweed collected on the coasts of Normandy and Brittany. To isolate the sodium carbonate, seaweed was burned and the ash washed with water. The remaining waste was destroyed by adding sulfuric acid. Courtois once added excessive sulfuric acid and a cloud of purple vapour rose. He noted that the vapour crystallised on cold surfaces, making dark crystals Courtois suspected that this material was a new element but lacked funding to pursue it further.
                     PROPERTIES
Iodine is the fourth halogen, being a member of group 17 in the periodic table, below fluorine, chlorine, and bromine; it is the heaviest stable member of its group. Like the other halogens, it is one electron short of a full octet and is hence a strong oxidising agent, reacting with many elements in order to complete its outer shell, although in keeping with periodic trends, it is the weakest oxidising agent among the stable halogens: it has the lowest electronegativity among them, just 2.66 on the Pauling scale (compare fluorine, chlorine, and bromine at 3.98, 3.16, and 2.96 respectively; astatine continues the trend with an electronegativity of 2.2). Elemental iodine hence forms diatomic molecules with chemical formula I2, where two iodine atoms share a pair of electrons in order to each achieve a stable octet for themselves; at high temperatures, these diatomic molecules reversibly dissociate a pair of iodine atoms. Similarly, the iodide anion, I−, is the strongest reducing agent among the stable halogens, being the most easily oxidised back to diatomic I2. The halogens darken in colour as the group is descended: fluorine is a very pale yellow gas, chlorine is greenish-yellow, and bromine is a reddish-brown volatile liquid. Iodine conforms to the prevailing trend, being a shiny black crystalline solid that melts at 114 °C and boils at 183 °C to form a violet gas. This trend occurs because the wavelengths of visible light absorbed by the halogens increase down the group (though astatine may not conform to it, depending on how metallic it turns out to be). Specifically, the violet colour of iodine gas results from the electron transition between the highest occupied antibonding πg molecular orbital and the lowest vacant antibonding σu molecular orbital.

                                 REACTIONS OF IODINE



                        CHEMISTRY AND COMPOUNDS
Though it is the least reactive of the halogens, iodine is still one of the more reactive elements. For example, while chlorine gas will halogenate carbon monoxide, nitric oxide, and sulfur dioxide (to phosgene, nitrosyl chloride, and sulfuryl chloride respectively), iodine will not do so. Furthermore, iodination of metals tends to result in lower oxidation states than chlorination or bromination; for example, rhenium metal reacts with chlorine to form rhenium hexachloride, but with bromine it forms only rhenium pentabromide and iodine can achieve only rhenium tetraiodide. By the same token, however, since iodine has the lowest ionisation energy among the halogens and is the most easily oxidised of them, it has a more significant cationic chemistry and its higher oxidation states are rather more stable than those of bromine and chlorine, for example in iodine heptafluoride.

                  APPLICATIONS
Unlike chlorine and bromine, which have one significant main use dwarfing all others, iodine is used in many applications of varying importance. About half of all produced iodine goes into various organoiodine compounds; another 15% remains as the pure element, another 15% is used to form potassium iodide, and another 15% for other inorganic iodine compounds. The remaining 5% is for minor uses. Among the major uses of iodine compounds are catalysts, animal feed supplements, stabilisers, dyes, colourants and pigments, pharmaceutical, sanitation (from tincture of iodine), and photography; minor uses include smog inhibition, cloud seeding, and various uses in analytical chemistry.