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
Chemical element with atomic number 53 (I) Br
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I
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At tellurium ← iodine → xenon Atomic number (Z)53Groupgroup 17 (halogens)Periodperiod 5Block p-blockElectron configuration[Kr] 4d10 5s2 5p5Electrons per shell2, 8, 18, 18, 7Physical propertiesPhase at STPsolidMelting 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 point386.65 K, 12.1 kPa Critical point819 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 Atomic propertiesOxidation statescommon: −1, +1, +3, +5, +7
+2,[4] +4,[5] +6[6]ElectronegativityPauling scale: 2.66 Ionisation energies
- 1st: 1008.4 kJ/mol
- 2nd: 1845.9 kJ/mol
- 3rd: 3180 kJ/mol
b = 478.28 pm
c = 982.38 pm (at 20 °C)[3]Thermal expansion74.9×10−6/K (at 20 °C)[a]Thermal conductivity0.449 W/(m⋅K) Electrical resistivity1.3×107 Ω⋅m (at 0 °C) Magnetic orderingdiamagnetic[7] Molar magnetic susceptibility−88.7×10−6 cm3/mol (298 K)[8]Bulk modulus7.7 GPa CAS Number7553-56-2 HistoryNamingfrom the Ancient Greek ιώδης, "violet", for the color of its vaporDiscovery and first isolationBernard 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.[9] Iodine deficiency affects about two billion people and is the leading preventable cause of intellectual disabilities.[10]
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.[11]
It is on the World Health Organization's List of Essential Medicines.[12]
History
In 1811, iodine was discovered by French chemist Bernard Courtois,[13][14] 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 ashes washed with water. While investigating the cause of corrosion to the copper vessels used in the process,[15] Courtois added an excess of sulfuric acid to the waste remaining and a cloud of violet vapour arose. He noted that the vapour crystallised on cold surfaces, forming dark crystals.[16] Courtois suspected that this material was a new element but lacked funding to pursue it further.[17]
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 Institut de France.[18] 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.[19][20][21] Gay-Lussac suggested the name "iode" (anglicised as "iodine"), from the Ancient Greek Ιώδης (iodēs, "violet"), because of the colour of iodine vapour.[13][19] 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.[22] Davy sent a letter dated 10 December to the Royal Society stating that he had identified a new element called iodine.[23] Arguments erupted between Davy and Gay-Lussac over who identified iodine first, but they both ultimately recognized Courtois as the discoverer of the element.[17]
In 1873, the French medical researcher Casimir Davaine (1812–1882) discovered the antiseptic action of iodine.[24] 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.[25]
In early periodic tables, iodine was often given the symbol J, for Jod, its name in German; in German texts, J is still frequently used in place of I.[26]
Properties
Iodine is the fourth halogen, being a member of group 17 in the periodic table, below fluorine, chlorine, and bromine; since astatine and tennessine are radioactive, iodine is the heaviest stable halogen. Iodine has an electron configuration of [Kr]5s24d105p5, 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, while that of tennessine is unknown). 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.[27] (Astatine goes further, being indeed unstable as At− and readily oxidised to At0 or At+.)[28]
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.[29] Nonpolar solvents such as hexane and carbon tetrachloride provide a higher solubility.[30] 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 colour of iodine vapour.[29] 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.[31]
The density, and the melting and boiling points of iodine follow the trend across all of the halogens of increasing with atomic number.[32]: 405 Among the stable halogens, 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 stable halogens. Similarly, iodine is the least volatile of the stable halogens, though the solid still can be observed to give off purple vapour.[27] 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.[33] Because it has the largest atomic radius among the stable halogens, iodine has the lowest first ionisation energy, lowest electron affinity, lowest electronegativity and lowest reactivity of the stable halogens.[27]
The interhalogen bond in diiodine is the weakest of all the stable 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.[27] 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.)[34] 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.[27] 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.[27]
Isotopes
Main article: Isotopes of iodineNaturally occurring iodine consists of one stable isotope, 127I, and is a mononuclidic element for atomic weight, which is thus a constant of nature determined by that isotope.[1] Radioisotopes are known from 108I to 147I. As other isotopes have half-lives too short to be primordial, it is also monoisotopic.
The longest-lived of the radioactive isotopes of iodine is iodine-129, which has a half-life of 16.1 million years, decaying via beta decay to stable xenon-129.[35] 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. Its former presence may be determined from an excess of its daughter xenon-129, but early attempts[36] to use this characteristic to date the supernova source for elements in the Solar System are made difficult by alternative nuclear processes giving iodine-129 and by iodine's volatility at higher temperatures.[37] Due to its mobility in the environment iodine-129 has been used to date very old groundwaters.[38][39]
The vast majority of iodine-129 on Earth today derives from human nuclear activity. Iodine-129 increased 3-8 orders of magnitude after nuclear activity began.[40] A small amount of naturally occurring iodine-129 forms from cosmic ray spallation of atmospheric xenon and as a fission product; the ratio 129I/127I is about 10−12.[41]
Excited states of iodine-127 and iodine-129 are often used in Mössbauer spectroscopy.[27]
The other iodine radioisotopes have much shorter half-lives, less than 60 days.[35] Some of them have medical applications involving the thyroid gland, where the iodine that enters the body is stored and concentrated. Iodine-123 (half-life 13.223 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.[42] Iodine-125 (half-life 59.392 days) is similar, 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.[43] Finally, iodine-131 (half-life 8.0249 days) beta-decays to xenon-131 and also emits gamma radiation. It is also be used for medicinal purposes in radiation therapy to the thyroid, when tissue destruction is desired after iodine uptake by the tissue.[44]
Iodine-131 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 and damage it through its radiation. 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.[45] Protection against the negative effects of iodine-131 upon a release is effected by saturating the thyroid gland with stable iodine-127 in the form of potassium iodide tablets, taken daily for optimal prophylaxis.[46]
Iodine-131 has also been used as a radioactive tracer.[47][48][49][50]
Chemistry and compounds
Main article: Iodine compoundsIodine is quite reactive, but it is less so than the lighter halogens, and it is a weaker oxidant. For example, it does not halogenate carbon monoxide, nitric oxide, and sulfur dioxide, which chlorine does. Many metals react with iodine.[27] For the same reason, 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.[29]
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.[52]
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:[53]
2 I2 + N2H4 H2O⟶ 4 HI + N2At 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 (−59.8 °F) and boils at −35.1 °C (−31.2 °F). 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.[54]
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.[55] Commercial so-called "concentrated" hydroiodic acid usually contains 48–57% HI by mass; the solution forms an azeotrope with boiling point 126.7 °C (260.1 °F) at 56.7 g HI per 100 g solution. Hence hydroiodic acid cannot be concentrated past this point by evaporation of water.[54] Unlike gaseous hydrogen iodide, hydroiodic acid has major industrial use in the manufacture of acetic acid by the Cativa process.[56][57]
