The dominant producers of iodine today areChile andJapan. Due to its high atomic number and ease of attachment toorganic compounds, it has also found favour as a non-toxicradiocontrast material. Because of the specificity of its uptake by the human body, radioactive isotopes of iodine can also be used to treatthyroid cancer. Iodine is also used as acatalyst in the industrial production ofacetic acid and somepolymers.
In 1811, iodine was discovered by French chemistBernard Courtois,[10][11] who was born to a family of manufacturers ofsaltpetre (an essential component ofgunpowder). At the time of theNapoleonic Wars, saltpetre was in great demand inFrance. Saltpetre produced from Frenchnitre beds requiredsodium carbonate, which could be isolated fromseaweed collected on the coasts ofNormandy andBrittany. To isolate the sodium carbonate, seaweed was burned and the ash washed with water. The remaining waste was destroyed by addingsulfuric 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) andNicolas Clément (1779–1841), to continue research. He also gave some of the substance to chemistJoseph Louis Gay-Lussac (1778–1850), and to physicistAndré-Marie Ampère (1775–1836). On 29 November 1813, Desormes and Clément made Courtois' discovery public by describing the substance to a meeting of the ImperialInstitut de France.[14] On 6 December 1813, Gay-Lussac found and announced that the new substance was either an element or a compound ofoxygen and he found that it is an element.[15][16][17] Gay-Lussac suggested the name "iode" (anglicised as "iodine"), from theAncient GreekΙώδης (iodēs, "violet"), because of the colour of iodine vapour.[10][15] Ampère had given some of his sample to British chemistHumphry Davy (1778–1829), who experimented on the substance and noted its similarity tochlorine and also found it as an element.[18] Davy sent a letter dated 10 December to theRoyal Society 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 researcherCasimir Davaine (1812–1882) discovered the antiseptic action of iodine.[20]Antonio Grossich (1849–1926), an Istrian-born surgeon, was among the first to usesterilisation 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 earlyperiodic tables, iodine was often given the symbolJ, forJod, its name inGerman; in German texts,J is still frequently used in place ofI.[22]
Iodine vapour in a flask, demonstrating its characteristic rich purple colour
Iodine is the fourthhalogen, being a member of group 17 in the periodic table, belowfluorine,chlorine, andbromine; sinceastatine andtennessine 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 itsvalence 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 withperiodic trends, it is the weakest oxidising agent among the stable halogens: it has the lowestelectronegativity 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 formsdiatomic 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 oftriiodide ions, among other polyiodides.[25] Nonpolar solvents such ashexane andcarbon tetrachloride provide a higher solubility.[26] Polar solutions, such as aqueous solutions, are brown, reflecting the role of these solvents asLewis 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 inalcohols andamines, 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 strongestVan 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 vapour.[23] Due to this property iodine is commonly used to demonstratesublimation directly fromsolid togas, which gives rise to a misconception that it does notmelt inatmospheric pressure.[28] Because it has the largestatomic radius among the halogens, iodine has the lowest firstionisation energy, lowestelectron affinity, lowestelectronegativity and lowest reactivity of the halogens.[23]
Structure of solid iodine
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 solidorthorhombic crystalline iodine, which has the same crystal structure as chlorine and bromine. (The record is held by iodine's neighbourxenon: the Xe–Xe bond length is 308.71 pm.)[29] 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 andsemiconducting properties.[23] Iodine is a two-dimensional semiconductor with aband 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]
Of the forty knownisotopes of iodine, only one occurs in nature,iodine-127. The others are radioactive and have half-lives too short to beprimordial. As such, iodine is bothmonoisotopic andmononuclidic 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 isiodine-129, which has a half-life of 15.7 million years, decaying viabeta decay to stablexenon-129.[30] Some iodine-129 was formed along with iodine-127 before the formation of theSolar System, but it has by now completely decayed away, making it anextinct radionuclide. Its former presence may be determined from an excess of itsdaughter xenon-129, but early attempts[31] 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.[32] Due to its mobility in the environment iodine-129 has been used to date very old groundwaters.[33][34] Traces of iodine-129 still exist today, as it is also acosmogenic nuclide, formed fromcosmic 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.[35] Excited states of iodine-127 and iodine-129 are often used inMössbauer spectroscopy.[23]
Protection usually used against the negative effects of iodine-131 is by saturating the thyroid gland with stable iodine-127 in the form ofpotassium iodide tablets, taken daily for optimal prophylaxis.[39] However, iodine-131 may also be used for medicinal purposes inradiation therapy for this very reason, when tissue destruction is desired after iodine uptake by the tissue.[40] Iodine-131 is also used as aradioactive tracer.[41][42][43][44]
Iodine is quite reactive, but it is less so than the lighter halogens, and it is a weaker oxidant. For example, it does nothalogenatecarbon monoxide,nitric oxide, andsulfur dioxide, whichchlorine does. Many metals react with iodine.[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 iniodine heptafluoride.[25]
I2•PPh3 charge-transfer complexes inCH2Cl2. From left to right: (1) I2 dissolved in dichloromethane – no CT complex. (2) A few seconds after excess PPh3 was added – CT complex is forming. (3) One minute later after excess PPh3 was added, the CT complex [Ph3PI]+I− has been formed. (4) Immediately after excess I2 was added, which contains [Ph3PI]+[I3]−.[45]
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 aπ* toσ* 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.[46]
The simplest compound of iodine ishydrogen iodide, HI. It is a colourless gas that reacts with oxygen to give water and iodine. Although it is useful iniodination 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 withhydrogen sulfide orhydrazine:[47]
2 I2 + N2H4H2O⟶ 4 HI + N2
At room temperature, it is a colourless gas, like all of the hydrogen halides excepthydrogen fluoride, since hydrogen cannot form stronghydrogen 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 anendothermic compound that can exothermically dissociate at room temperature, although the process is very slow unless acatalyst is present: the reaction between hydrogen and iodine at room temperature to give hydrogen iodide does not proceed to completion. The H–Ibond dissociation energy is likewise the smallest of the hydrogen halides, at 295 kJ/mol.[48]
Aqueous hydrogen iodide is known ashydroiodic 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.[49] Commercial so-called "concentrated" hydroiodic acid usually contains 48–57% HI by mass; the solution forms anazeotrope 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.[48] Unlike gaseous hydrogen iodide, hydroiodic acid has major industrial use in the manufacture ofacetic acid by theCativa process.[50][51]
With the exception of thenoble 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[52] was only known as an ammonia adduct. Ammonia-free NI3 was found to be isolable at –196 °C but spontaneously decomposes at 0 °C.[53] 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).[54] Sulfur-iodine and selenium-iodine polyatomic cations (e.g., [S2I42+][AsF6–]2 and [Se2I42+][Sb2F11–]2) have been prepared and characterised crystallographically.[55]
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 ofniobium,tantalum, andprotactinium. 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 withaluminium(III) iodide at 230 °C to givemolybdenum(II) iodide. An example involving halogen exchange is given below, involving the reaction oftantalum(V) chloride with excess aluminium(III) iodide at 400 °C to givetantalum(V) iodide:[56]
Lower iodides may be produced either through thermal decomposition or disproportionation, or by reducing the higher iodide with hydrogen or a metal, for example:[56]
Most metal iodides with the metal in low oxidation states (+1 to +3) are ionic. Nonmetals tend to form covalent molecular iodides, as do metals in high oxidation states from +3 and above. Both ionic and covalent iodides are known for metals in oxidation state +3 (e.g.scandium iodide is mostly ionic, butaluminium iodide is not). Ionic iodides MIn tend to have the lowest melting and boiling points among the halides MXn of the same element, because the electrostatic forces of attraction between the cations and anions are weakest for the large iodide anion. In contrast, covalent iodides tend to instead have the highest melting and boiling points among the halides of the same element, since iodine is the most polarisable of the halogens and, having the most electrons among them, can contribute the most to van der Waals forces. Naturally, exceptions abound in intermediate iodides where one trend gives way to the other. Similarly, solubilities in water of predominantly ionic iodides (e.g.potassium andcalcium) are the greatest among ionic halides of that element, while those of covalent iodides (e.g.silver) are the lowest of that element. In particular,silver iodide is very insoluble in water and its formation is often used as a qualitative test for iodine.[56]
The halogens form many binary,diamagneticinterhalogen compounds with stoichiometries XY, XY3, XY5, and XY7 (where X is heavier than Y), and iodine is no exception. Iodine forms all three possible diatomic interhalogens, a trifluoride and trichloride, as well as a pentafluoride and, exceptionally among the halogens, a heptafluoride. Numerous cationic and anionic derivatives are also characterised, such as the wine-red or bright orange compounds ofICl+ 2 and the dark brown or purplish black compounds of I2Cl+. Apart from these, somepseudohalides are also known, such ascyanogen iodide (ICN), iodinethiocyanate (ISCN), and iodineazide (IN3).[57]
Iodine monochloride
Iodine monofluoride (IF) is unstable at room temperature and disproportionates very readily and irreversibly to iodine andiodine pentafluoride, and thus cannot be obtained pure. It can be synthesised from the reaction of iodine with fluorine gas intrichlorofluoromethane at −45 °C, withiodine trifluoride in trichlorofluoromethane at −78 °C, or withsilver(I) fluoride at 0 °C.[57]Iodine monochloride (ICl) andiodine monobromide (IBr), on the other hand, are moderately stable. The former, a volatile red-brown compound, was discovered independently byJoseph Louis Gay-Lussac andHumphry Davy in 1813–1814 not long after the discoveries of chlorine and iodine, and it mimics the intermediate halogen bromine so well thatJustus von Liebig was misled into mistaking bromine (which he had found) for iodine monochloride. Iodine monochloride and iodine monobromide may be prepared simply by reacting iodine with chlorine or bromine at room temperature and purified byfractional crystallisation. Both are quite reactive and attack evenplatinum andgold, though notboron,carbon,cadmium,lead,zirconium,niobium,molybdenum, andtungsten. Their reaction with organic compounds depends on conditions. Iodine chloride vapour tends to chlorinatephenol andsalicylic acid, since when iodine chloride undergoeshomolytic fission, chlorine and iodine are produced and the former is more reactive. However, iodine chloride incarbon tetrachloride solution results in iodination being the main reaction, since nowheterolytic fission of the I–Cl bond occurs and I+ attacks phenol as an electrophile. However, iodine monobromide tends to brominate phenol even in carbon tetrachloride solution because it tends to dissociate into its elements in solution, and bromine is more reactive than iodine.[57] When liquid, iodine monochloride and iodine monobromide dissociate intoI 2X+ andIX− 2 ions (X = Cl, Br); thus they are significant conductors of electricity and can be used as ionising solvents.[57]
Iodine trifluoride (IF3) is an unstable yellow solid that decomposes above −28 °C. It is thus little-known. It is difficult to produce because fluorine gas would tend to oxidise iodine all the way to the pentafluoride; reaction at low temperature withxenon difluoride is necessary.Iodine trichloride, which exists in the solid state as the planar dimer I2Cl6, is a bright yellow solid, synthesised by reacting iodine with liquid chlorine at −80 °C; caution is necessary during purification because it easily dissociates to iodine monochloride and chlorine and hence can act as a strong chlorinating agent. Liquid iodine trichloride conducts electricity, possibly indicating dissociation toICl+ 2 andICl− 4 ions.[58]
Iodine pentafluoride (IF5), a colourless, volatile liquid, is the most thermodynamically stable iodine fluoride, and can be made by reacting iodine with fluorine gas at room temperature. It is a fluorinating agent, but is mild enough to store in glass apparatus. Again, slight electrical conductivity is present in the liquid state because of dissociation toIF+ 4 andIF− 6. Thepentagonal bipyramidaliodine heptafluoride (IF7) is an extremely powerful fluorinating agent, behind onlychlorine trifluoride,chlorine pentafluoride, andbromine pentafluoride among the interhalogens: it reacts with almost all the elements even at low temperatures, fluorinatesPyrex glass to form iodine(VII) oxyfluoride (IOF5), and setscarbon monoxide on fire.[59]
Iodine oxides are the most stable of all the halogen oxides, because of the strong I–O bonds resulting from the large electronegativity difference between iodine and oxygen, and they have been known for the longest time.[27] The stable, white,hygroscopiciodine pentoxide (I2O5) has been known since its formation in 1813 by Gay-Lussac and Davy. It is most easily made by the dehydration ofiodic acid (HIO3), of which it is the anhydride. It will quickly oxidise carbon monoxide completely tocarbon dioxide at room temperature, and is thus a useful reagent in determining carbon monoxide concentration. It also oxidisesnitrogen oxide,ethylene, andhydrogen sulfide. It reacts withsulfur trioxide and peroxydisulfuryl difluoride (S2O6F2) to form salts of the iodyl cation, [IO2]+, and is reduced by concentratedsulfuric acid to iodosyl salts involving [IO]+. It may be fluorinated byfluorine,bromine trifluoride,sulfur tetrafluoride, orchloryl fluoride, resultingiodine pentafluoride, which also reacts withiodine pentoxide, giving iodine(V) oxyfluoride, IOF3. A few other less stable oxides are known, notably I4O9 and I2O4; their structures have not been determined, but reasonable guesses are IIII(IVO3)3 and [IO]+[IO3]− respectively.[60]
Standard reduction potentials for aqueous I species[61]
Hypoiodous acid is unstable to disproportionation. The hypoiodite ions thus formed disproportionate immediately to give iodide and iodate:[61]
3 IO− ⇌ 2 I− +IO− 3K = 1020
Iodous acid and iodite are even less stable and exist only as a fleeting intermediate in the oxidation of iodide to iodate, if at all.[61] Iodates are by far the most important of these compounds, which can be made by oxidisingalkali metal iodides with oxygen at 600 °C and high pressure, or by oxidising iodine withchlorates. Unlike chlorates, which disproportionate very slowly to form chloride and perchlorate, iodates are stable to disproportionation in both acidic and alkaline solutions. From these, salts of most metals can be obtained. Iodic acid is most easily made by oxidation of an aqueous iodine suspension byelectrolysis or fumingnitric acid. Iodate has the weakest oxidising power of the halates, but reacts the quickest.[62]
Many periodates are known, including not only the expected tetrahedralIO− 4, but also square-pyramidalIO3− 5, octahedral orthoperiodateIO5− 6, [IO3(OH)3]2−, [I2O8(OH2)]4−, andI 2O4− 9. They are usually made by oxidising alkalinesodium iodate electrochemically (withlead(IV) oxide as the anode) or by chlorine gas:[63]
IO− 3 + 6 OH− →IO5− 6 + 3 H2O + 2 e−
IO− 3 + 6 OH− + Cl2 →IO5− 6 + 2 Cl− + 3 H2O
They are thermodymically and kinetically powerful oxidising agents, quickly oxidising Mn2+ toMnO− 4, and cleavingglycols, α-diketones, α-ketols, α-aminoalcohols, and α-diamines.[63] Orthoperiodate especially stabilises high oxidation states among metals because of its very high negative charge of −5.Orthoperiodic acid, H5IO6, is stable, and dehydrates at 100 °C in a vacuum toMetaperiodic acid, HIO4. Attempting to go further does not result in the nonexistent iodine heptoxide (I2O7), but rather iodine pentoxide and oxygen. Periodic acid may be protonated bysulfuric acid to give theI(OH)+ 6 cation, isoelectronic to Te(OH)6 andSb(OH)− 6, and giving salts with bisulfate and sulfate.[27]
When iodine dissolves in strong acids, such as fuming sulfuric acid, a bright blueparamagnetic solution includingI+ 2 cations is formed. A solid salt of the diiodine cation may be obtained by oxidising iodine withantimony pentafluoride:[27]
2 I2 + 5 SbF5SO2⟶20 °C 2 I2Sb2F11 + SbF3
The salt I2Sb2F11 is dark blue, and the bluetantalum analogue I2Ta2F11 is also known. Whereas the I–I bond length in I2 is 267 pm, that inI+ 2 is only 256 pm as the missing electron in the latter has been removed from an antibonding orbital, making the bond stronger and hence shorter. Influorosulfuric acid solution, deep-blueI+ 2 reversibly dimerises below −60 °C, forming red rectangular diamagneticI2+ 4. Other polyiodine cations are not as well-characterised, including bent dark-brown or blackI+ 3 and centrosymmetricC2h green or blackI+ 5, known in theAsF− 6 andAlCl− 4 salts among others.[27][64]
The only important polyiodide anion in aqueous solution is lineartriiodide,I− 3. Its formation explains why the solubility of iodine in water may be increased by the addition of potassium iodide solution:[27]
I2 + I− ⇌I− 3 (Keq = c. 700 at 20 °C)
Many other polyiodides may be found when solutions containing iodine and iodide crystallise, such asI− 5,I− 9,I2− 4, andI2− 8, whose salts with large, weakly polarising cations such asCs+ may be isolated.[27][65]
Thecarbon–iodine bond is a common functional group that forms part of coreorganic chemistry; formally, these compounds may be thought of as organic derivatives of theiodide anion. The simplestorganoiodine compounds,alkyl iodides, may be synthesised by the reaction ofalcohols withphosphorus triiodide; these may then be used innucleophilic substitution reactions, or for preparingGrignard reagents. The C–I bond is the weakest of all the carbon–halogen bonds due to the minuscule difference in electronegativity between carbon (2.55) and iodine (2.66). As such, iodide is the bestleaving group among the halogens, to such an extent that many organoiodine compounds turn yellow when stored over time due to decomposition into elemental iodine; as such, they are commonly used inorganic synthesis, because of the easy formation and cleavage of the C–I bond.[70] They are also significantly denser than the other organohalogen compounds thanks to the high atomic weight of iodine.[71] A few organic oxidising agents like theiodanes contain iodine in a higher oxidation state than −1, such as2-iodoxybenzoic acid, a common reagent for the oxidation of alcohols toaldehydes,[72] andiodobenzene dichloride (PhICl2), used for the selective chlorination ofalkenes andalkynes.[73] One of the more well-known uses of organoiodine compounds is the so-callediodoform test, whereiodoform (CHI3) is produced by the exhaustive iodination of amethyl ketone (or another compound capable of being oxidised to a methyl ketone), as follows:[74]
Some drawbacks of using organoiodine compounds as compared to organochlorine or organobromine compounds is the greater expense and toxicity of the iodine derivatives, since iodine is expensive and organoiodine compounds are stronger alkylating agents.[75] For example,iodoacetamide andiodoacetic acid denature proteins by irreversibly alkylatingcysteine residues and preventing the reformation ofdisulfide linkages.[76]
Halogen exchange to produce iodoalkanes by theFinkelstein reaction is slightly complicated by the fact that iodide is a better leaving group than chloride or bromide. The difference is nevertheless small enough that the reaction can be driven to completion by exploiting the differential solubility of halide salts, or by using a large excess of the halide salt.[74] In the classic Finkelstein reaction, analkyl chloride or analkyl bromide is converted to analkyl iodide by treatment with a solution ofsodium iodide inacetone. Sodium iodide is soluble in acetone andsodium chloride andsodium bromide are not.[77] The reaction is driven toward products bymass action due to the precipitation of the insoluble salt.[78][79]
Iodine is the least abundant of the stable halogens, comprising only 0.46 parts per million of Earth's crustal rocks (compare:fluorine: 544 ppm,chlorine: 126 ppm,bromine: 2.5 ppm) making it the 60th most abundant element.[80] Iodide minerals are rare, and most deposits that are concentrated enough for economical extraction are iodate minerals instead. Examples includelautarite, Ca(IO3)2, and dietzeite, 7Ca(IO3)2·8CaCrO4.[80] These are the minerals that occur as trace impurities in thecaliche, found inChile, whose main product issodium nitrate. In total, they can contain at least 0.02% and at most 1% iodine by mass.[81]Sodium iodate is extracted from the caliche and reduced to iodide bysodium bisulfite. This solution is then reacted with freshly extracted iodate, resulting incomproportionation to iodine, which may be filtered off.[23]
The caliche was the main source of iodine in the 19th century and continues to be important today, replacingkelp (which is no longer an economically viable source),[82] but in the late 20th centurybrines emerged as a comparable source. The JapaneseMinami Kantō gas field east ofTokyo and the AmericanAnadarko Basin gas field in northwestOklahoma are the two largest such sources. The brine is hotter than 60 °C from the depth of the source. The brine is firstpurified and acidified usingsulfuric acid, then the iodide present is oxidised to iodine with chlorine. An iodine solution is produced, but is dilute and must be concentrated. Air is blown into the solution toevaporate the iodine, which is passed into an absorbing tower, wheresulfur dioxide reduces the iodine. Thehydrogen iodide (HI) is reacted with chlorine to precipitate the iodine. After filtering and purification the iodine is packed.[81][83]
2 HI + Cl2 → I2↑ + 2 HCl
I2 + 2 H2O + SO2 → 2 HI + H2SO4
2 HI + Cl2 → I2↓ + 2 HCl
These sources ensure that Chile and Japan are the largest producers of iodine today.[80] Alternatively, the brine may be treated withsilver nitrate to precipitate out iodine assilver iodide, which is then decomposed by reaction with iron to form metallic silver and a solution ofiron(II) iodide. The iodine is then liberated by displacement withchlorine.[84]
As an element with highelectron density and atomic number, iodine efficiently absorbs X-rays. X-rayradiocontrast agents is the top application for iodine.[85] In this application, Organoiodine compounds are injected intravenously. This application is often in conjunction with advanced X-ray techniques such asangiography andCT scanning. At present, all water-soluble radiocontrast agents rely oniodine-containing compounds.
Iodine absorbs X-rays with energies less than 33.3 keV due to thephotoelectric effect of the innermost electrons.[86]
Use of iodine as a biocide represents a major application of the element, ranked 2nd by weight.[85] Elemental iodine (I2) is used as anantiseptic in medicine.[87] A number of water-soluble compounds, fromtriiodide (I3−, generatedin situ by addingiodide to poorly water-soluble elemental iodine) to variousiodophors, slowly decompose to release I2 when applied.[88]
Thin-film-transistor liquid crystal displays rely on polarisation. The liquid crystal transistor is sandwiched between two polarising films and illuminated from behind. The two films prevent light transmission unless the transistor in the middle of the sandwich rotates the light.[89] Iodine-impregnated polymer films are used in polarising optical components with the highest transmission and degree of polarisation.[90]
Another significant use of iodine is as a cocatalyst for the production ofacetic acid by theMonsanto andCativa processes. In these technologies,hydroiodic acid converts themethanol feedstock into methyl iodide, which undergoescarbonylation. Hydrolysis of the resulting acetyl iodide regenerates hydroiodic acid and gives acetic acid. The majority of acetic acid is produced by these approaches.[91][92]
Salts of iodide and iodate are used extensively in human and animal nutrition. This application reflects the status of iodide as anessential element, being required for two hormones. The production ofethylenediamine dihydroiodide, provided as anutritional supplement for livestock, consumes a large portion of available iodine.[85] Iodine is a component ofiodised salt.
A saturated solution ofpotassium iodide is used to treat acutethyrotoxicosis. It is also used to block uptake ofiodine-131 in the thyroid gland (see isotopes section above), when this isotope is used as part of radiopharmaceuticals (such asiobenguane) that are not targeted to the thyroid or thyroid-type tissues.[93][94]
Inorganic iodides find specialised uses.Titanium,zirconium,hafnium, andthorium are purified by theVan Arkel–de Boer process, which involves the reversible formation of the tetraiodides of these elements. Silver iodide is a major ingredient to traditional photographic film. Thousands of kilograms of silver iodide are used annually forcloud seeding to induce rain.[85]
The organoiodine compounderythrosine is an important food colouring agent. Perfluoroalkyl iodides are precursors to important surfactants, such asperfluorooctanesulfonic acid.[85]
An iodine based thermochemical cycle has been evaluated for hydrogen production using energy from nuclear power.[96] The cycle has three steps. At 120 °C (248 °F), iodine reacts withsulfur dioxide and water to give hydrogen iodide andsulfuric acid:
I2 + SO2 + 2 H2O → 2 HI + H2SO4
After a separation stage, at 830–850 °C (1,530–1,560 °F) sulfuric acid splits in sulfur dioxide and oxygen:
2 H2SO4 → SO2 + 2 H2O + O2
Hydrogen iodide, at 300–320 °C (572–608 °F), gives hydrogen and the initial element, iodine:
2 HI → I2 + H2
The yield of the cycle (ratio between lower heating value of the produced hydrogen and the consumed energy for its production, is approximately 38%. As of 2020[update], the cycle is not a competitive means of producing hydrogen.[96]
The spectrum of the iodine molecule, I2, consists of (not exclusively) tens of thousands of sharp spectral lines in the wavelength range 500–700 nm. It is therefore a commonly used wavelength reference (secondary standard). By measuring with aspectroscopic Doppler-free technique while focusing on one of these lines, thehyperfine structure of the iodine molecule reveals itself. A line is now resolved such that either 15 components (from even rotational quantum numbers,Jeven), or 21 components (from odd rotational quantum numbers,Jodd) are measurable.[97]
Caesium iodide and thallium-doped sodium iodide are used in crystalscintillators for the detection of gamma rays. The efficiency is high and energy dispersive spectroscopy is possible, but the resolution is rather poor.
Testing a seed for starch with a solution of iodine
The iodide and iodate anions can be used for quantitative volumetric analysis, for example iniodometry. Iodine and starch form a blue complex, and this reaction is often used to test for either starch or iodine and as anindicator in iodometry. The iodine test for starch is still used to detectcounterfeit banknotes printed on starch-containing paper.[98]
Theiodine value is the mass of iodine in grams that is consumed by 100 grams of achemical substance typically fats or oils. Iodine numbers are often used to determine the amount of unsaturation infatty acids. This unsaturation is in the form ofdouble bonds, which react with iodine compounds.
Thethyroid system of the thyroid hormonesT3 andT4Comparison of the iodine content in urine inFrance (in microgramme/day), for some regions and departments (average levels of urine iodine, measured in micrograms per litre at the end of the twentieth century (1980 to 2000))[100]
Iodine is anessential element for life and, at atomic numberZ = 53, is the heaviest element commonly needed by living organisms. (Lanthanum and the otherlanthanides, as well astungsten withZ = 74 anduranium withZ = 92, are used by a few microorganisms.[101][102][103]) It is required for the synthesis of the growth-regulating thyroid hormonestetraiodothyronine andtriiodothyronine (T4 and T3 respectively, named after their number of iodine atoms). A deficiency of iodine leads to decreased production of T3 and T4 and a concomitant enlargement of thethyroid tissue in an attempt to obtain more iodine, causing the diseasegoitre. The major form of thyroid hormone in the blood is tetraiodothyronine (T4), which has a longer life than triiodothyronine (T3). In humans, the ratio of T4 to T3 released into the blood is between 14:1 and 20:1. T4 is converted to the active T3 (three to four times more potent than T4) withincells bydeiodinases (5'-iodinase). These are further processed bydecarboxylation and deiodination to produceiodothyronamine (T1a) andthyronamine (T0a'). All three isoforms of the deiodinases areselenium-containing enzymes; thus metallic selenium is needed for triiodothyronine and tetraiodothyronine production.[104]
Iodine accounts for 65% of the molecular weight of T4 and 59% of T3. Fifteen to 20 mg of iodine is concentrated in thyroid tissue and hormones, but 70% of all iodine in the body is found in other tissues, including mammary glands,eyes, gastric mucosa, thymus,cerebrospinal fluid, choroid plexus, arteries,cervix, salivary glands. During pregnancy, theplacenta is able to store and accumulate iodine.[105][106] In the cells of those tissues, iodine enters directly bysodium-iodide symporter (NIS). The action of iodine in mammal tissues is related to fetal and neonatal development, and in the other tissues, it is known.[107]
The daily levels of intake recommended by theUnited StatesNational Academy of Medicine are between 110 and 130μg for infants up to 12 months, 90 μg for children up to eight years, 130 μg for children up to 13 years, 150 μg for adults, 220 μg for pregnant women and 290 μg for lactating women.[7][108] The Tolerable Upper Intake Level (TUIL) for adults is 1,100 μg/day.[109] This upper limit was assessed by analysing the effect of supplementation onthyroid-stimulating hormone.[107]
TheEuropean Food Safety Authority (EFSA) refers to the collective set of information as Dietary Reference Values, with Population Reference Intake (PRI) instead of RDA, and Average Requirement instead of EAR; AI and UL are defined the same as in the United States. For women and men ages 18 and older, the PRI for iodine is set at 150 μg/day; the PRI during pregnancy and lactation is 200 μg/day. For children aged 1–17 years, the PRI increases with age from 90 to 130 μg/day. These PRIs are comparable to the U.S. RDAs with the exception of that for lactation.[110]
As of 2000, the median intake of iodine from food in the United States was 240 to 300 μg/day for men and 190 to 210 μg/day for women.[109] The general US population has adequate iodine nutrition,[118][119] with lactating women and pregnant women having a mild risk of deficiency.[119] In Japan, consumption was considered much higher, ranging between 5,280 μg/day to 13,800 μg/day fromwakame andkombu that are eaten,[107] both in the form of kombu and wakame and kombu and wakameumamiextracts forsoup stock andpotato chips. However, new studies suggest that Japan's consumption is closer to 1,000–3,000 μg/day.[120] The adult UL in Japan was last revised to 3,000 μg/day in 2015.[121]
After iodine fortification programs such as iodisation ofsalt have been done, some cases of iodine-inducedhyperthyroidism have been observed (so-calledJod-Basedow phenomenon). The condition occurs mainly in people above 40 years of age, and the risk is higher when iodine deficiency is high and the first rise in iodine consumption is high.[122]
In areas where there is little iodine in the diet,[123] which are remote inland areas and faraway mountainous areas where no iodine rich foods are eaten,iodine deficiency gives rise tohypothyroidism, symptoms of which areextreme fatigue,goitre,mental slowing,depression,low weight gain[clarification needed], andlow basal body temperatures.[124] Iodine deficiency is the leading cause of preventableintellectual disability, a result that occurs primarily when babies or small children are renderedhypothyroidic by no iodine. The addition of iodine to salt has largely destroyed this problem in wealthier areas, but iodine deficiency remains a serious public health problem in poorer areas today.[125] Iodine deficiency is also a problem in certain areas of all continents of the world. Information processing, fine motor skills, and visual problem solving are normalised by iodine repletion in iodine-deficient people.[126]
Elemental iodine (I2) istoxic if taken orally undiluted. The lethal dose for an adult human is 30 mg/kg, which is about 2.1–2.4 grams for a human weighing 70 to 80 kg (even when experiments on rats demonstrated that these animals could survive after eating a 14000 mg/kg dose and are still living after that). Excess iodine is morecytotoxic in the presence ofselenium deficiency.[129] Iodine supplementation in selenium-deficient populations is problematic for this reason.[107] The toxicity derives from its oxidising properties, through which it denaturates proteins (including enzymes).[130]
Elemental iodine is also a skin irritant. Solutions with high elemental iodine concentration, such astincture of iodine andLugol's solution, are capable of causingtissue damage if used in prolonged cleaning or antisepsis; similarly, liquidPovidone-iodine (Betadine) trapped against the skin resulted in chemical burns in some reported cases.[131]
Some people develop ahypersensitivity to products and foods containing iodine. Applications of tincture of iodine or Betadine can cause rashes, sometimes severe.[133]Parenteral use of iodine-based contrast agents (see above) can cause reactions ranging from a mild rash to fatalanaphylaxis. Such reactions have led to the misconception (widely held, even among physicians) that some people are allergic to iodine itself; even allergies to iodine-rich foods have been so construed.[134] In fact, there has never been a confirmed report of a true iodine allergy, as an allergy to iodine or iodine salts is biologically impossible. Hypersensitivity reactions to products and foods containing iodine are apparently related to their other molecular components;[135] thus, a person who has demonstrated an allergy to one food or product containing iodine may not have an allergic reaction to another. Patients with various food allergies (fishes, shellfishes, eggs, milk, seaweeds, kelp, meats, vegetables, kombu, wakame) do not have an increased risk for a contrast medium hypersensitivity.[136][135] The patient's allergy history is relevant.[137]
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