Iodine, ₅₃I

Iodine
Pronunciation	/ˈaɪədaɪn, -dɪn, -diːn/ ​(EYE-ə-dyne, -⁠din, -⁠deen)
Appearance	lustrous metallic gray solid, black/violet liquid, violet gas
Standard atomic weight Ar°(I)
	126.90447±0.00003[1] 126.90±0.01 (abridged)[2]
Iodine in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson Br ↑ I ↓ At tellurium ← iodine → xenon
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) P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 260 282 309 342 381 457
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	1st: 1008.4 kJ/mol 2nd: 1845.9 kJ/mol 3rd: 3180 kJ/mol
Atomic radius	empirical: 140 pm
Covalent radius	139±3 pm
Van der Waals radius	198 pm
Spectral lines of iodine
Other properties
Natural occurrence	primordial
Crystal structure	​base-centered orthorhombic (oS8)
Lattice constants	a = 725.79 pm 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 v e
Main isotopes Decay abun­dance half-life (t1/2) mode pro­duct 123I synth 13 h β+100% 123Te 124I synth 4.176 d ε 124Te 125I synth 59.40 d ε 125Te 127I 100% stable 129I trace 1.57×107 y β− 129Xe 131I synth 8.02070 d β−100% 131Xe 135I synth 6.57 h β− 135Xe
Category: Iodine view talk edit | references

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. They described 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" (Englished 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]}

Chemistry and compounds

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.^[23] 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.^[25]

Charge-transfer complexes

The iodine molecule, I₂, dissolves in CCl₄ and aliphatic hydrocarbons to give bright violet solutions. In these solvents the absorption band maximum occurs in the 520 – 540 nm region and is assigned to a π^* to σ^* transition. When I₂ reacts with Lewis bases in these solvents a blue shift in I₂ peak is seen and the new peak (230 – 330 nm) arises that is due to the formation of adducts, which are referred to as charge-transfer complexes.^[47]

Hydrogen iodide

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:^[48]

2 I₂ + N₂H₄ H₂O⟶ 4 HI + N₂

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 295 kJ/mol.^[49]

Aqueous hydrogen iodide is known as hydroiodic acid, which is a strong acid. Hydrogen iodide is exceptionally soluble in water: one litre of water will dissolve 425 litres of hydrogen iodide, and the saturated solution has only four water molecules per molecule of hydrogen iodide.^[50] Commercial so-called "concentrated" hydroiodic acid usually contains 48–57% HI by mass; the solution forms an azeotrope with boiling point 126.7 °C at 56.7 g HI per 100 g solution. Hence hydroiodic acid cannot be concentrated past this point by evaporation of water.^[49] Unlike gaseous hydrogen iodide, hydroiodic acid has major industrial use in the manufacture of acetic acid by the Cativa process.^[51][52]

Unlike hydrogen fluoride, anhydrous liquid hydrogen iodide is difficult to work with as a solvent, because its boiling point is low, it has a small liquid range, its permittivity is low and it does not dissociate appreciably into H₂I⁺ and HI⁻
₂ ions – the latter, in any case, are much less stable than the bifluoride ions (HF⁻
₂) 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⁺
₄ (R = Me, Et, Buⁿ) 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.^[49]

Other binary iodine compounds

With the exception of the noble gases, nearly all elements on the periodic table up to einsteinium (EsI₃ is known) are known to form binary compounds with iodine. Until 1990, nitrogen triiodide^[53] was only known as an ammonia adduct. Ammonia-free NI₃ was found to be isolable at –196 °C but spontaneously decomposes at 0 °C.^[54] For thermodynamic reasons related to electronegativity of the elements, neutral sulfur and selenium iodides that are stable at room temperature are also nonexistent, although S₂I₂ and SI₂ are stable up to 183 and 9 K, respectively. As of 2022, no neutral binary selenium iodide has been unambiguously identified (at any temperature).^[55] Sulfur- and selenium-iodine polyatomic cations (e.g., ₂ and ₂) have been prepared and characterized crystallographically.^[56]

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:^[57]

${\displaystyle 3{\text{TaCl}}_5^{}+\underset{(\text{excess})}{5{\text{AlI}}_3^{}}⟶<!--\; ⟶\; -->3}$Zdroj:https://en.wikipedia.org?pojem=Iodine
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