A | B | C | D | E | F | G | H | CH | I | J | K | L | M | N | O | P | Q | R | S | T | U | V | W | X | Y | Z | 0 | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9
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| Iodine | ||||||||||||||||||||||||||||||||||||||||||||||
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| Pronunciation | /ˈaɪədaɪn, -dɪn, -diːn/ | |||||||||||||||||||||||||||||||||||||||||||||
| Appearance | lustrous metallic gray solid, black/violet liquid, violet gas | |||||||||||||||||||||||||||||||||||||||||||||
| Standard atomic weight Ar°(I) | ||||||||||||||||||||||||||||||||||||||||||||||
| Iodine in the periodic table | ||||||||||||||||||||||||||||||||||||||||||||||
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| Atomic number (Z) | 53 | |||||||||||||||||||||||||||||||||||||||||||||
| Group | group 17 (halogens) | |||||||||||||||||||||||||||||||||||||||||||||
| Period | period 5 | |||||||||||||||||||||||||||||||||||||||||||||
| Block | p-block | |||||||||||||||||||||||||||||||||||||||||||||
| Electron configuration | [Kr] 4d10 5s2 5p5 | |||||||||||||||||||||||||||||||||||||||||||||
| Electrons per shell | 2, 8, 18, 18, 7 | |||||||||||||||||||||||||||||||||||||||||||||
| Physical properties | ||||||||||||||||||||||||||||||||||||||||||||||
| Phase at STP | solid | |||||||||||||||||||||||||||||||||||||||||||||
| Melting point | (I2) 386.85 K (113.7 °C, 236.66 °F) | |||||||||||||||||||||||||||||||||||||||||||||
| Boiling point | (I2) 457.4 K (184.3 °C, 363.7 °F) | |||||||||||||||||||||||||||||||||||||||||||||
| Density (at 20° C) | 4.944 g/cm3[3] | |||||||||||||||||||||||||||||||||||||||||||||
| Triple point | 386.65 K, 12.1 kPa | |||||||||||||||||||||||||||||||||||||||||||||
| Critical point | 819 K, 11.7 MPa | |||||||||||||||||||||||||||||||||||||||||||||
| Heat of fusion | (I2) 15.52 kJ/mol | |||||||||||||||||||||||||||||||||||||||||||||
| Heat of vaporisation | (I2) 41.57 kJ/mol | |||||||||||||||||||||||||||||||||||||||||||||
| Molar heat capacity | (I2) 54.44 J/(mol·K) | |||||||||||||||||||||||||||||||||||||||||||||
Vapour pressure (rhombic)
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| Atomic properties | ||||||||||||||||||||||||||||||||||||||||||||||
| Oxidation states | −1, 0, +1, +2,[4] +3, +4, +5, +6, +7 (a strongly acidic oxide) | |||||||||||||||||||||||||||||||||||||||||||||
| Electronegativity | Pauling scale: 2.66 | |||||||||||||||||||||||||||||||||||||||||||||
| Ionisation energies |
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| Atomic radius | empirical: 140 pm | |||||||||||||||||||||||||||||||||||||||||||||
| Covalent radius | 139±3 pm | |||||||||||||||||||||||||||||||||||||||||||||
| Van der Waals radius | 198 pm | |||||||||||||||||||||||||||||||||||||||||||||
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| Other properties | ||||||||||||||||||||||||||||||||||||||||||||||
| Natural occurrence | primordial | |||||||||||||||||||||||||||||||||||||||||||||
| Crystal structure | base-centered orthorhombic (oS8) | |||||||||||||||||||||||||||||||||||||||||||||
| Lattice constants |  b = 478.28 pm c = 982.38 pm (at 20 °C)[3] | |||||||||||||||||||||||||||||||||||||||||||||
| Thermal expansion | 74.9×10−6/K (at 20 °C)[a] | |||||||||||||||||||||||||||||||||||||||||||||
| Thermal conductivity | 0.449 W/(m⋅K) | |||||||||||||||||||||||||||||||||||||||||||||
| Electrical resistivity | 1.3×107 Ω⋅m (at 0 °C) | |||||||||||||||||||||||||||||||||||||||||||||
| Magnetic ordering | diamagnetic[5] | |||||||||||||||||||||||||||||||||||||||||||||
| Molar magnetic susceptibility | −88.7×10−6 cm3/mol (298 K)[6] | |||||||||||||||||||||||||||||||||||||||||||||
| Bulk modulus | 7.7 GPa | |||||||||||||||||||||||||||||||||||||||||||||
| CAS Number | 7553-56-2 | |||||||||||||||||||||||||||||||||||||||||||||
| History | ||||||||||||||||||||||||||||||||||||||||||||||
| Discovery and first isolation | Bernard Courtois (1811) | |||||||||||||||||||||||||||||||||||||||||||||
| Isotopes of iodine | ||||||||||||||||||||||||||||||||||||||||||||||
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Iodine is a chemical element; it has symbol I and atomic number 53. The heaviest of the stable halogens, it exists at standard conditions as a semi-lustrous, non-metallic solid that melts to form a deep violet liquid at 114 °C (237 °F), and boils to a violet gas at 184 °C (363 °F). The element was discovered by the French chemist Bernard Courtois in 1811 and was named two years later by Joseph Louis Gay-Lussac, after the Ancient Greek Ιώδης, meaning 'violet'.
Iodine occurs in many oxidation states, including iodide (I−), iodate (IO−
3), and the various periodate anions. As the heaviest essential mineral nutrient, iodine is required for the synthesis of thyroid hormones.[7] Iodine deficiency affects about two billion people and is the leading preventable cause of intellectual disabilities.[8]
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.
It is on the World Health Organization's List of Essential Medicines.[9]
History
In 1811, iodine was discovered by French chemist Bernard Courtois,[10][11] who was born to a family of manufacturers 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 violet vapour rose. He noted that the vapour crystallised on cold surfaces, making dark black crystals.[12] Courtois suspected that this material was a new element but lacked funding to pursue it further.[13]
Courtois gave samples to his friends, Charles Bernard Desormes (1777–1838) and Nicolas Clément (1779–1841), to continue research. He also gave some of the substance to chemist Joseph Louis Gay-Lussac (1778–1850), and to physicist André-Marie Ampère (1775–1836). On 29 November 1813, Desormes and Clément made Courtois' discovery public by describing the substance to a meeting of the Imperial Institute of France.[14] On 6 December 1813, Gay-Lussac found and announced that the new substance was either an element or a compound of oxygen and he found that it is an element.[15][16][17] Gay-Lussac suggested the name "iode" (anglicized as "iodine"), from the Ancient Greek Ιώδης (iodēs, "violet"), because of the colour of iodine vapor.[10][15] Ampère had given some of his sample to British chemist Humphry Davy (1778–1829), who experimented on the substance and noted its similarity to chlorine and also found it as an element.[18] Davy sent a letter dated 10 December to the Royal Society of London stating that he had identified a new element called iodine.[19] Arguments erupted between Davy and Gay-Lussac over who identified iodine first, but both scientists found that both of them identified iodine first and also knew that Courtois is the first one to isolate the element.[13]
In 1873, the French medical researcher Casimir Davaine (1812–1882) discovered the antiseptic action of iodine.[20] Antonio Grossich (1849–1926), an Istrian-born surgeon, was among the first to use sterilisation of the operative field. In 1908, he introduced tincture of iodine as a way to rapidly sterilise the human skin in the surgical field.[21]
In early periodic tables, iodine was often given the symbol J, for Jod, its name in German.[22]
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. (The fifth and sixth halogens, the radioactive astatine and tennessine, are not well-studied due to their expense and inaccessibility in large quantities, but appear to show various unusual properties for the group due to relativistic effects.) Iodine has an electron configuration of 4d105s25p5, with the seven electrons in the fifth and outermost shell being its valence electrons. Like the other halogens, it is one electron short of a full octet and is hence an 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.[23] (Astatine goes further, being indeed unstable as At− and readily oxidised to At0 or At+.)[24]
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.
Elemental iodine is slightly soluble in water, with one gram dissolving in 3450 mL at 20 °C and 1280 mL at 50 °C; potassium iodide may be added to increase solubility via formation of triiodide ions, among other polyiodides.[25] Nonpolar solvents such as hexane and carbon tetrachloride provide a higher solubility.[26] 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.[25] 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.[27]
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.[23] 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.[29] 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.[23]
The interhalogen bond in diiodine is the weakest of all the halogens. As such, 1% of a sample of gaseous iodine at atmospheric pressure is dissociated into iodine atoms at 575 °C. Temperatures greater than 750 °C are required for fluorine, chlorine, and bromine to dissociate to a similar extent. Most bonds to iodine are weaker than the analogous bonds to the lighter halogens.[23] Gaseous iodine is composed of I2 molecules with an I–I bond length of 266.6 pm. The I–I bond is one of the longest single bonds known. It is even longer (271.5 pm) in solid orthorhombic crystalline iodine, which has the same crystal structure as chlorine and bromine. (The record is held by iodine's neighbour xenon: the Xe–Xe bond length is 308.71 pm.)[30] As such, within the iodine molecule, significant electronic interactions occur with the two next-nearest neighbours of each atom, and these interactions give rise, in bulk iodine, to a shiny appearance and semiconducting properties.[23] Iodine is a two-dimensional semiconductor with a band gap of 1.3 eV (125 kJ/mol): it is a semiconductor in the plane of its crystalline layers and an insulator in the perpendicular direction.[23]
Isotopes
Of the thirty-seven known isotopes of iodine, only one occurs in nature, iodine-127. The others are radioactive and have half-lives too short to be primordial. As such, iodine is both monoisotopic and mononuclidic and its atomic weight is known to great precision, as it is a constant of nature.[23]
The longest-lived of the radioactive isotopes of iodine is iodine-129, which has a half-life of 15.7 million years, decaying via beta decay to stable xenon-129.[31] Some iodine-129 was formed along with iodine-127 before the formation of the Solar System, but it has by now completely decayed away, making it an extinct radionuclide that is nevertheless still useful in dating the history of the early Solar System or very old groundwaters, due to its mobility in the environment. Its former presence may be determined from an excess of its daughter xenon-129.[32][33][34][35][36] Traces of iodine-129 still exist today, as it is also a cosmogenic nuclide, formed from cosmic ray spallation of atmospheric xenon: these traces make up 10−14 to 10−10 of all terrestrial iodine. It also occurs from open-air nuclear testing, and is not hazardous because of its very long half-life, the longest of all fission products. At the peak of thermonuclear testing in the 1960s and 1970s, iodine-129 still made up only about 10−7 of all terrestrial iodine.[37] Excited states of iodine-127 and iodine-129 are often used in Mössbauer spectroscopy.[23]
The other iodine radioisotopes have much shorter half-lives, no longer than days.[31] Some of them have medical applications involving the thyroid gland, where the iodine that enters the body is stored and concentrated. Iodine-123 has a half-life of thirteen hours and decays by electron capture to tellurium-123, emitting gamma radiation; it is used in nuclear medicine imaging, including single photon emission computed tomography (SPECT) and X-ray computed tomography (X-Ray CT) scans.[38] Iodine-125 has a half-life of fifty-nine days, decaying by electron capture to tellurium-125 and emitting low-energy gamma radiation; the second-longest-lived iodine radioisotope, it has uses in biological assays, nuclear medicine imaging and in radiation therapy as brachytherapy to treat a number of conditions, including prostate cancer, uveal melanomas, and brain tumours.[39] Finally, iodine-131, with a half-life of eight days, beta decays to an excited state of stable xenon-131 that then converts to the ground state by emitting gamma radiation. It is a common fission product and thus is present in high levels in radioactive fallout. It may then be absorbed through contaminated food, and will also accumulate in the thyroid. As it decays, it may cause damage to the thyroid. The primary risk from exposure to high levels of iodine-131 is the chance occurrence of radiogenic thyroid cancer in later life. Other risks include the possibility of non-cancerous growths and thyroiditis.[40]
Protection usually used against the negative effects of iodine-131 is by saturating the thyroid gland with stable iodine-127 in the form of potassium iodide tablets, taken daily for optimal prophylaxis.[41] However, iodine-131 may also be used for medicinal purposes in radiation therapy for this very reason, when tissue destruction is desired after iodine uptake by the tissue.[42] Iodine-131 is also used as a radioactive tracer.[43][44][45][46]