The melting and boiling points of iodine are the highest among the halogens, conforming to the increasing trend down the group, since iodine has the largest electron cloud among them that is the most easily polarised, resulting in its molecules having the strongest van der Waals interactions among the halogens. Similarly, iodine is the least volatile of the halogens, though the solid still can be observed to give off purple vapor. [21] Due to this property iodine is commonly used to demonstrate sublimation directly from solid to gas, which gives rise to a misconception that it does not melt in atmospheric pressure. [27] Because it has the largest atomic radius among the halogens, iodine has the lowest first ionisation energy, lowest electron affinity, lowest electronegativity and lowest reactivity of the halogens. [21] Structure of solid iodine At room temperature, it is a colourless gas, like all of the hydrogen halides except hydrogen fluoride, since hydrogen cannot form strong hydrogen bonds to the large and only mildly electronegative iodine atom. It melts at −51.0°C and boils at −35.1°C. It is an endothermic compound that can exothermically dissociate at room temperature, although the process is very slow unless a catalyst is present: the reaction between hydrogen and iodine at room temperature to give hydrogen iodide does not proceed to completion. The H–I bond dissociation energy is likewise the smallest of the hydrogen halides, at 295kJ/mol. [47]

Iodine - British Dietetic Association (BDA)

due to the very weak hydrogen bonding between hydrogen and iodine, though its salts with very large and weakly polarising cations such as Cs + and NR + R = Me, Et, Bu n) may still be isolated. Anhydrous hydrogen iodide is a poor solvent, able to dissolve only small molecular compounds such as nitrosyl chloride and phenol, or salts with very low lattice energies such as tetraalkylammonium halides. [47] Other binary iodine compounds [ edit ] The halogens darken in colour as the group is descended: fluorine is a very pale yellow, chlorine is greenish-yellow, bromine is reddish-brown, and iodine is violet. In early periodic tables, iodine was often given the symbol J, for Jod, its name in German. [20] Properties [ edit ] Iodine vapour in a flask.monosaccharides – such as glucose close glucose A simple sugar used by cells for respiration. and fructose close fructose A monosaccharide which joins with glucose to make sucrose. The simplest compound of iodine is hydrogen iodide, HI. It is a colourless gas that reacts with oxygen to give water and iodine. Although it is useful in iodination reactions in the laboratory, it does not have large-scale industrial uses, unlike the other hydrogen halides. Commercially, it is usually made by reacting iodine with hydrogen sulfide or hydrazine: [46] 2 I 2 + N 2H 4 H 2O ⟶ 4 HI + N 2 The dominant producers of iodine today are Chile and Japan. 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. Given the large size of the iodide anion and iodine's weak oxidising power, high oxidation states are difficult to achieve in binary iodides, the maximum known being in the pentaiodides of niobium, tantalum, and protactinium. Iodides can be made by reaction of an element or its oxide, hydroxide, or carbonate with hydroiodic acid, and then dehydrated by mildly high temperatures combined with either low pressure or anhydrous hydrogen iodide gas. These methods work best when the iodide product is stable to hydrolysis. Other syntheses include high-temperature oxidative iodination of the element with iodine or hydrogen iodide, high-temperature iodination of a metal oxide or other halide by iodine, a volatile metal halide, carbon tetraiodide, or an organic iodide. For example, molybdenum(IV) oxide reacts with aluminium(III) iodide at 230°C to give molybdenum(II) iodide. An example involving halogen exchange is given below, involving the reaction of tantalum(V) chloride with excess aluminium(III) iodide at 400°C to give tantalum(V) iodide: [55] 3 TaCl 5 + 5 AlI 3 ( excess ) ⟶ 3 TaI 5 + 5 AlCl 3 {\displaystyle {\ce {3TaCl5 + {\underset {(excess)}{5AlI3}}-> 3TaI5 + 5AlCl3}}} Elemental iodine is slightly soluble in water, with one gram dissolving in 3450mL at 20°C and 1280mL at 50°C; potassium iodide may be added to increase solubility via formation of triiodide ions, among other polyiodides. [23] Nonpolar solvents such as hexane and carbon tetrachloride provide a higher solubility. [24] Polar solutions, such as aqueous solutions, are brown, reflecting the role of these solvents as Lewis bases; on the other hand, nonpolar solutions are violet, the color of iodine vapour. [23] Charge-transfer complexes form when iodine is dissolved in polar solvents, hence changing the colour. Iodine is violet when dissolved in carbon tetrachloride and saturated hydrocarbons but deep brown in alcohols and amines, solvents that form charge-transfer adducts. [25] I 2• PPh 3 charge-transfer complexes in CH 2Cl 2. From left to right: (1) I 2 dissolved in dichloromethane – no CT complex. (2) A few seconds after excess PPh 3 was added – CT complex is forming. (3) One minute later after excess PPh 3 was added, the CT complex [Ph 3PI] +I − has been formed. (4) Immediately after excess I 2 was added, which contains [Ph 3PI] +[I 3] −. [26]

Unusual uses for white iodine - Ask Dr Gott MD Unusual uses for white iodine - Ask Dr Gott MD

Iodine is quite reactive, but it is much less reactive than the other halogens. 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. [21] 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. [23] Charge-transfer complexes [ edit ]

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Taking high doses of iodine for long periods of time could change the way your thyroid gland works. Lower iodides may be produced either through thermal decomposition or disproportionation, or by reducing the higher iodide with hydrogen or a metal, for example: [55] TaI 5 + Ta → 630 ∘ C ⟶ 575 ∘ C thermal gradient Ta 6 I 14 {\displaystyle {\ce {TaI5{}+Ta->[{\text{thermal gradient}}][{\ce {630