Iodine is a chemical element with the symbol I and the atomic number 53. The heaviest of the halogen is stable, it exists as a shiny purple-black metal solid at standard sublime conditions ready to form gas violet. The elemental shape was invented by the French chemist Bernard Courtois in 1811. Named two years later by Joseph-Louis Gay-Lussac of this property, after the Greek ????? "purple".
Iodine occurs in many oxidizing states, including iodide (I - ), iodate ( IO -
3 ) , and various periodic anions. This is the least abundant halogen, being the most abundant of the sixty-one elements. Even less than the so-called rare earth. This is the toughest essential mineral nutrient. Iodine is essential in the synthesis of thyroid hormones. Iodine deficiency affects about two billion people and is a major cause of preventable intellectual disabilities.
The dominant iodine manufacturers today are Chile and Japan. Iodine and its compounds are mainly used in nutrients. Due to the high number of atoms and the ease of attachment to organic compounds, it also benefits as a non-toxic radiocontrast material. Due to the specificity of the absorption by the human body, radioactive isotope of iodine can also be used to treat thyroid cancer. Iodine is also used as a catalyst in industrial production of acetic acid and some polymers.
Video Iodine
History
In 1811, iodine was discovered by the French chemist Bernard Courtois, who was born from a producer of burp (an essential component of gunpowder). At the time of Napoleonic Wars, belchas are in great demand in France. Saltpeter produced from French nitre beds requires sodium carbonate, which can be isolated from seaweed collected on the beaches of Normandy and Brittany. To isolate sodium carbonate, the seaweed is burned and the ash is washed with water. The remaining waste is destroyed by adding sulfuric acid. Courtois once added excess sulfuric acid and purple vapor clouds rose. He notes that the vapor crystallizes on a cold surface, making the crystals dark. Courtois suspects that this material is a new element but lacks the funds to pursue it further.
Courtois gave an example to his friends, Charles Bernard Desormes (1777-1838) and Nicolas ClÃÆ'Â © ment (1779-1841), to continue his research. He also gave some chemicals Joseph Louis Gay-Lussac (1778-1850), and to the physicist AndrÃÆ' © -Marie AmpÃÆ'¨re (1775-1836). On 29 November 1813, Desormes and ClÃÆ'Â © ment made the discovery of Courtois public. They describe the substance to the meeting of the French Imperial Institute. On December 6, Gay-Lussac announced that the new substance is an element or an oxygen compound. It is Gay-Lussac who suggests the name "iode" , from the Greek ??????? ( ioeid? s ) for violet (due to the color of iodine vapor). AmpÃÆ'¨re has given some of his samples to the English chemist Humphry Davy (1778-1829), who experimented on the substance and noted the similarity to chlorine. Davy sent a letter dated December 10 to the Royal Society of London stating that he had identified a new element. Arguments erupted between Davy and Gay-Lussac about who identified the first iodine, but both scientists recognized Courtois as the first to isolate the elements.
Maps Iodine
Properties
Iodine is the fourth halogen, being a member of group 17 in the periodic table, under fluorine, chlorine, and bromine; it is the toughest stable member of his group. (The fifth rare and fugitive halogen, radioactive astatin, is not well studied because of its cost and inaccessibility in large quantities, but seems to exhibit a variety of unusual properties due to its relativistic effect.) Iodine has electron configurations [Kr] 4d 10 5s 2 5p 5 , with seven electrons on the fifth and outer shell is the valence electron. Like other halogens, it is one of the short electrons of a full octet and hence a strong oxidizing agent, reacting with many elements to resolve its outermost shell, albeit in keeping with the periodic trend, it is the weakest oxidizing agent amongst the stable. halogen: has the lowest electronegativity between them, only 2.66 on the Pauling scale (compare fluorine, chlorine, and bromine at 3.98, 3.16, and 2.96 respectively; astatin continues the trend with electronegativity 2.2). The iodine element therefore forms a diatomic molecule with the chemical formula I 2 , in which two iodine atoms share a pair of electrons for each to achieve a stable octet for themselves; at high temperatures, this diatomal molecule reversibly separates a pair of iodine atoms. Similarly, anion iodide, I - , is the strongest reducing agent amongst stable halogens, which is most readily oxidized back to diatomic I 2 . (Astatine goes further, becoming unstable as in - and easily oxidizes to At 0 or On , despite the presence of At 2 not completed.)
Halogen is dark when the group is lowered: fluorine is a very pale yellow gas, yellowish-green chlorine, and bromine is a reddish brown volatile liquid. Iodine in accordance with prevailing trends, becomes a shiny black crystalline solid that melts at 114 ° C and boils at 183 ° C to form violet gases. This tendency occurs because the visible wavelength of visible light absorbed by the halogen rises down the group (though astatine may not correspond to it, depending on how the metal turns out). In particular, the violet color of iodine gas results from the transition of electrons between the orbitals of the highest occupied orbital molecule ? g and the lowest empty antibody ? u molecular orbital.
The iodine element 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 through the formation of triiodide ions, among other polyodides. Nonpolar solvents such as hexane and carbon tetrachloride give higher solubility. Polar solutions, such as aqueous solutions, are brown, reflecting the role of this solvent as a Lewis base; on the other hand, a purple nonpolar solution, iodine vapor color. The charge transfer complex is formed when iodine is dissolved in a polar solvent, thus changing its color. Iodine is violet when dissolved in carbon tetrachloride and saturated hydrocarbons but dark brown in alcohols and amines, solvents that make charge-transfer addition.
The melting and boiling point of iodine is the highest among halogens, in keeping with the increasing trend in the group, since iodine has the largest electron cloud among those most easily polarized, so the molecule has the strongest van der Waals interaction. between halogens. Similarly, iodine is the most unstable of halogens. Because it has the largest atomic radius among the halogens, iodine has the first lowest ionization energy, the lowest electron affinity, the lowest electronegativity and the lowest reactivity of the halogen.
The interhalogenic bond in diiodine is the weakest of all halogens. Thus, 1% of the iodine gas sample at atmospheric pressure is separated into an iodine atom at 575 ° C. Temperatures greater than 750 ° C are required for fluorine, chlorine, and bromine to dissociate to the same extent. Most of the bonds to iodine are weaker than analogous bonds with lighter halogens. Iodine gas consists of I 2 molecules with I-I bond length 266.6Ã, pm. I-I bond is one of the longest known single bonds. This is even longer (271.5 μm) in solid orthorhombic iodine crystals, which have the same crystal structure as chlorine and bromine. (Records held by xenon iodine neighbors: Xe-Xe bond length is 308,71 pm). Thus, in iodine molecules, significant electronic interactions occur with the next two next neighbors of each atom, and this interaction gives rise, in large quantities of iodine, to a glossy appearance and semiconductor properties. Iodine is a two-dimensional semiconductor with band gap 1.3Ã, eV (125 kJ/mol): it is a semiconductor in the field of crystal layers and insulators in the direction perpendicular.
Isotope
Of the thirty-seven known iodine isotopes, only one occurs in nature, iodine-127. Others are radioactive and have half-lives too short to be primordial. Thus, iodine is monoisotopic and its atomic weight is known to be very precise, due to its constant nature.
The longest living iodine radioactive isotope is iodine-129, which has a half-life of 15.7 million years, decaying through beta decay into stable xenon-129. Some iodine-129 was formed together with iodine-127 before the formation of the Solar System, but it has now completely decomposed, making it an extinct radionuclide that is still useful in dating the history of the early Solar System or very old. groundwaters, because of their mobility in the environment. Her previous presence can be determined from the surplus of her daughter xenon-129. The iodine-129 trace still exists today, as it is also a cosmogenic nuclide, formed from cosmic ray spallation of atmospheric xenon: this trace forms 10 -14 to 10 -10 of all iodine terrestrial. This also happens from open nuclear tests, and is harmless 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 formed only about 10 -7 of all terrestrial iodine. The interesting conditions of iodine-127 and iodine-129 are often used in MÃÆ'¶ssbauer spectroscopy.
Another iodine radioisotope has a shorter half-life, no more than a day. 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 capturing electrons to tellurium-123, emitting gamma radiation; used in nuclear medicine imaging, including scanning emission computed tomography (SPECT) emissions and X-ray computed tomography (X-Ray CT) scans. Iodine-125 has a half-life of fifty-nine days, decaying by capturing electrons to tellurium-125 and emitting low-energy gamma radiation; the second longest living iodine radioisotope, has use in biological tests, imaging of nuclear medicine and in radiation therapy as brachytherapy to treat a number of conditions, including prostate cancer, uveal melanoma, and brain tumors. Finally, iodine-131, with a half-life of eight days, beta decay into a stable zinc state of xenon-131 which then converts to ground state by emitting gamma radiation. It is a common fission product and is thus present at a high level in radioactive fallout. These can then be absorbed through contaminated food, and will also accumulate in the thyroid. When it decays, it can cause damage to the thyroid. The main risk of exposure to high levels of iodine-131 is the possibility of radiogenic thyroid cancer later in life. Other risks include the possibility of non-cancerous growth and thyroiditis.
The usual protective measure against the negative effects of iodine-131 is to saturate the thyroid gland with stable iodine-127 in the form of potassium iodide tablets, taken daily for optimal prophylaxis. However, iodine-131 â € <â €
Chemicals and compounds
Although this is the most reactive of halogens, iodine is still one of the more reactive elements. For example, while chlorine gas will halogenate carbon monoxide, nitric oxide, and sulfur dioxide (for phosgene, nitrosil chloride, and sulfuryl chloride), iodine will not do it. Furthermore, metal iodination tends to produce lower oxidation than chlorination or bromination; for example, the rhenium metal reacts with chlorine to form rhenium hexachloride, but with bromine only forming rhenium pentabromide and iodine can only reach rhenium tetraiodide. In the same way, however, since iodine has the lowest ionisation energy among the halogens and the most oxidizable of them, it has more significant cationic chemistry and higher oxidation states are somewhat more stable than bromine and chlorine, since the examples in yeptafluoride iodine.
Hydrogen iodide
The simplest iodine compound is hydrogen iodide, HI. It is a colorless gas that reacts with oxygen to provide water and iodine. Although useful in laboratory iodination reactions, it has no large-scale industrial use, unlike other hydrogen halides. Commercially, it is usually made by reacting iodine with hydrogen sulphide or hydrazine:
- 2 2
4 H 2 sub> O ? 4 HI N 2
At room temperature, it is a colorless gas, like all hydrogen halides except hydrogen fluoride, because hydrogen can not form strong hydrogen bonds to large and few electronegative iodine atoms. It melts at -51.0 ° C and boils at -35.1 ° C. It is an exothermic endothermic compound dissociates at room temperature, although the process is very slow unless there is a catalyst: the reaction between hydrogen and iodine at room temperature for providing hydrogen iodide is not continued until complete. The dissociation energy of the H-I bond is also the smallest of the hydrogen halide, at 295 kJ/mol.
The aqueous hydrogen iodide is known as hydroiodic acid, which is a strong acid. Hydrogen iodide is very soluble in water: one liter of water will dissolve 425 liters of hydrogen iodide, and saturated solution has only four water molecules per molecule of hydrogen iodide. Commercial so-called "concentrated" hydroiodic acids usually contain 48-57% HI by mass; the solution forms an azeotrope with a boiling point of 126.7 Â ° C at 56.7 g HI per 100 g solution. Therefore hydroiodic acid can not be concentrated past this point by evaporation of water.
Unlike hydrogen fluoride, anhydrous liquid hydrogen iodide is difficult to work with as a solvent, since its boiling point is low, it has a small liquid range, its dielectric constant is low and does not dissociate considerably into H 2 I and HI -
- 2 ) due to the very weak hydrogen bonds between hydrogen and iodine, although its salts with very large and weak polarization cations such as Cs and
4
Other binary Iodides
Almost all elements in the periodic table form a binary iodide. The exception is clear in the minority and comes from every case of one of three causes: extreme reluctance and unwillingness to participate in chemical reactions (noble gases); extreme nuclear instability prevents chemical investigations before decay and transmutation (many of the toughest elements outside bismuth); and have higher electronegativity than iodine (oxygen, nitrogen, and the first three halogens), so the formally produced binary compound is not an iodide but oxide, nitride, or iodine halide. (Nonetheless, nitrogen triiodides are named as iodides because they are analogous to other nitrogen trihalides.)
Given the large size of iodide anion and the weak oxidation power of iodine, the high oxidation state is difficult to achieve in iodide binaries, known to be maximum in the niobium, tantalum, and protactinium pentaiodids. Iodides can be prepared by reaction of an element or oxide, hydroxide, or carbonate with hydroiodic acid, and then dehydrated by a somewhat high temperature combined with a low pressure or anhydrous hydrogen iodide gas. This method works best when the iodide product is stable against hydrolysis; if not, the possibility of including high temperature oxidative iodination of elements with iodine or hydrogen iodide, high temperature iodination of metal oxides or other halides by iodine, volatile metal halides, tetrazy carbon, or organic iodides. For example, molybdenum (IV) oxide reacts with aluminum (III) iodide at 230 ° C to give molybdenum (II) iodide. Examples involving a halogen exchange are given below, involving the reaction of tantalum (V) chloride with excess aluminum (III) iodide at 400 ° C to give tantalum (V) iodide:
Lower Iodide can be produced either by thermal decomposition or disproportionation, or by reducing higher iodides with hydrogen or metals, for example:
Most of the iodides of pre-transition metals (groups 1, 2, and 3, together with lanthanides and actinides in 2 and 3 oxidation numbers) are largely ionic, whereas non-metals tend to form covalent molecular iodides, as do metals in high oxidation states of 3 and above. Ionic ionics MI n tend to have the lowest melting and boiling points between the MX halides n of the same element, because of the electrostatic attraction between the cation and the weakest anion for the large anion iodide. In contrast, covalent iodides tend to have the highest melting and boiling points among halides of the same element, since iodine is the most polarized of halogens and, having the most electrons among them, can contribute most to van der Waals forces. Of course, exceptions are abundant in the intermediate iodide where one trend gives way to another. Similarly, water solubilities predominantly ionic iodide (eg potassium and calcium) are the largest among the ionic halides of the element, while those of the covalent iodide (eg silver) are the lowest of the element. Specifically, silver iodide is very insoluble in water and its formation is often used as a qualitative test for iodine.
Iodine halide
Halogen forms many binary, interalogenic compounds diamagnetic with stoichiometry XY, XY 3 , and XY 5 (where X is heavier than Y ), and iodine is no exception. Iodine forms the three possible diatomic interhalogen, trifluoride and trichloride, as well as pentafluoride and, strongly among halogens, a heptafluoride. Many cationic and anionic derivatives are also characterized, such as a bright red wine or orange compound from
2 and dark chocolate or purplish black compound from I 2 Cl . In addition to this, some pseudohalides are also known, such as cyanogen iodide (ICN), iodine thiocyanate (ISCN), and iodine azide (IN 3 ).
Iodine monofluoride (IF) is unstable at room temperature and disproportionately very easy and irreversible to iodine and iodine pentafluoride, and thus can not be obtained purely. It can be synthesized from the iodine reaction with fluorine gas in trichlorofluoromethane at -45 Â ° C, with iodine trifluoride in trichlorofluoromethane at -78 Â ° C, or with silver (I) fluoride at 0 Â ° C. Iodine monochloride (ICl) and iodine monobromide (IBr), on the other hand, is quite stable. The first, a volatile red-brown compound, was discovered separately by Joseph Louis Gay-Lussac and Humphry Davy in 1813-4 shortly after the discovery of chlorine and iodine, and he imitated medium-sized bromine halogens so well that Justus von Liebig was misled into bromine (which he found) for iodine monochloride. Iodine monochloride and iodine monobromide can be prepared only by reacting iodine with chlorine or bromine at room temperature and purified by fractional crystallization. Both are quite reactive and invade even platinum and gold, though not boron, carbon, cadmium, lead, zirconium, niobium, molybdenum, and tungsten. Their reaction with organic compounds depends on the conditions. Iodine chloride steam tends to chlorinate phenol and salicylic acid, because when iodine chloride undergoes homolytic dissociation, chlorine and iodine are produced and the former is more reactive. However, iodine chloride in tetrachloromethane solution produces iodination as the main reaction, because now the heterolytic cleavage of I-Cl bonds occurs and I attacks phenol as electrophil. However, iodine monobromides tend to brominate phenols even in tetrachloromethane solutions because they tend to dissociate into the elements in solution, and bromine is more reactive than iodine. When the liquid, iodine monochloride and iodine monobromide dissociate into I
2 X and IX -
2 anions (X = Cl, Br); thus they are significant electrical conductors and can be used as ionizing solvents.
Iodine trifluoride (IF 3 ) is an unstable yellow solid unraveling above -28 ° C. Thus little is known. Difficult to produce because fluorine gases will tend to oxidize iodine all the way to pentafluoride; reaction at low temperature with xenon difluoride is required. Iodine trichloride, present in solid state as planar dimer I 2 Cl 6 , is a bright yellow solid, synthesized by reacting iodine with liquid chlorine at -80 ° C; caution is required during purification because it easily dissociates with iodine monochloride and chlorine and can therefore act as a powerful chlorination agent. Liquid iodine trichloride conducts electricity, may show dissociation to ICl
2 and ICl -
4 ion.
Iodine pentafluoride (IF 5 ), a colorless, volatile liquid, is the most stable thermodynamic iodine fluoride, and can be prepared by reacting iodine with a fluorine gas at room temperature. It is a fluorinating agent, but light enough to be stored in glassware. Again, a little electrical conductivity is present in a liquid state because it dissociates into IF > 4 and IF -
6 . The pentagonal bipyramidal iodine heptafluoride (IF 7 ) is a very powerful fluorinating agent, behind only chlorine trifluoride, pentafluoride chlorine, and pentafluoride bromine between the interhalogens: it reacts with almost all elements even at low temperatures, fluorinate Pyrex glass to form iodine (VII) oxyfluoride (IOF 5 ), and burn carbon monoxide.
Iodine oxides and oxoacids
Iodine oxide is the most stable of all halogen oxides, because of the strong I-O bonds resulting from large electronegativity differences between iodine and oxygen, and they have been known for a long time. The stable, white, iodine iodine hygroscopic (I 2 5 ) has been known since its formation in 1813 by Gay-Lussac and Davy. This is most easily made by the dehydration of iodic acid (HIO 3 ), which is anhydride. It will rapidly oxidize carbon monoxide completely into carbon dioxide at room temperature, and thus is a useful reagent in determining the carbon monoxide concentration. It also oxidizes nitrogen oxide, ethylene, and hydrogen sulfide. Reacts with sulfur trioxide and peroxydisulfuryl difluoride (S 2 O 6 2 ) to form the salts of the iodyl cations, [IO 2 ] , and reduced by concentrated sulfuric acid to iodosil salts involving [IO] . It may be fluorinated by fluorine, bromine trifluoride, sulfur tetrafluoride, or chloryl fluoride, producing iodine pentafluoride, which also reacts with iodine pentoxide, giving iodine (V) oxyfluoride, IOF 3 . Some other less stable oxides are known, especially my 4 O 9 and I 2 O 4 ; the structure is undetermined, but a reasonable guess is that I III (I V O 3 ) 3 and [IO ] [IO 3 ] - respectively.
More important are the four oxoacids: hypoiodous acid (HIO), iodous acid (HIO 2 ), iodic acid (HIO 3 ), and periodic acid (HIO 4 or H 5 IO 6 ). When iodine dissolves in an aqueous solution, the following reactions occur:
Hypoiodous acid is unstable against disproportion. Hypoiodit ions thus form disproportionately immediately to provide iodide and iodate:
Iodic and ioditic acids are even less stable and exist only as intermediates at a glance in oxidation of iodide to iodate, if at all. Iodates are the most important of these compounds, which can be prepared by oxidizing alkali metal iodides with oxygen at 600 ° C and high pressure, or by oxidizing iodine with chlorate. Unlike chlorates, which are highly disproportionate to form chlorides and perchlorates, iodates are stable for disproportionation in acid and base solutions. From this, salts of most metals can be obtained. The best iodic acid is prepared by oxidation of an aqueous iodine suspension by electrolysis or fuming nitric acid. Iodate has the weakest oxidizing strength of halat, but it reacts most quickly.
Many known periods, including not only the expected tetrahedral IO - 4 , but also square-pyramidal IO 3 -
5 , orthoperiodate oktahedral
6 , [IO 3
2 O 4 -
- 3 6 OH - -> IO 5 - 6 3 H 2 O 2 e -
6 2 Cl - 3 H 2 O
They are powerful thermodymically and kinetic oxidizing agents, rapidly oxidize Mn 2 to MnO - 4 , and splitting glycols ,? -diketon ,? -ketol ,? -aminoalcohol, and? -diamina. Orthoperiodate primarily stabilizes the high oxidation state between metals because of its very high negative charge -5. Orthoperiodic acid, H 5 IO 6 , stable, and dehydrated at 100 ° C in a vacuum for metaperiodic acid, HIO 4 . Trying to go further does not produce the missing iodine heptoxide (I 2 O 7 ), but preferably iodine pentoxide and oxygen. Periodic acid can be protonated by sulfuric acid to give I (OH) 6 cation, isoelectronics for Te (OH) 6 and Sb (OH) - 6 , and salt with bisulfate and sulfate.
Polyodin compounds
When iodine dissolves in strong acids, such as angry sulfuric acid, a bright blue paramagnetic solution includes 2 cation is formed. The solid salt of the diiodin cation can be obtained by oxidizing iodine with an antimony pentafluoride:
- 2 I 2 5 SbF 5 SO 2 ? 20Ã, Â ° C 2 I 2 Sb 2 F 11 SbF 3
The salt I 2 Sb 2 F 11 is dark blue, and the blue tantalum analog I 2 Ta 2 F 11 is also known. Whereas the length of the bond II in I 2 is 267 pm, which is in font-size: inherit; line-height: inherit; vertical-align: baseline " >
2 only 256 pm as electron lost on the latter has been removed from the anti-bonding orbital, making the bond stronger and therefore shorter. In fluorosulfuric acid solution, deep-blue I 2 dimerises reversible below -60 ° C, form a diamagnetic red rectangle I am 2
4 . Other polyiodine cations do not have good characteristics, including dark-brown or black bends
3 and centrosymmetric C > 2 h green or black
5 , known in AsF -
6 and AlCl -
4 salt among others.
The only important polyiodide anion in aqueous solution is a linear triiodide, -
3 . Its formation explains why iodine solubility in water can be increased by the addition of a potassium iodide solution:
- I 2 I am - ? I - 3 ( K eq = ~ 700 at 20 Â ° C)
Many other polyiodides can be found when solutions containing iodine and iodide crystallize, such as I -
5 , I < sup style = "font-size: inherit; line-height: inherit; vertical-align: baseline"> -
9 , 2 -
4 , and I 2 -
8 , whose salts with large width polarization cations such as Cs can be isolated.
Organoiodin compound
Organoiodin compounds have been fundamental in the development of organic synthesis, as in the removal of Hofmann amines, Williamson ether synthesis, Wurtz partner reactions, and in Grignard reagents.
The carbon-iodine bond is a general functional group that forms part of the core organic chemical; formally, this compound can be considered as an organic derivative of anion iodide. The simplest organoiodin compounds, alkyl iodides, can be synthesized by alcohol reactions with phosphorus triiodide; this can then be used in nucleophilic substitution reactions, or to prepare Grignard reagents. The C-I bond is the weakest of all carbon-halogen bonds because of the very small difference in electronegativity between carbon (2.55) and iodine (2.66). Thus, iodide is the best abandoned group among halogens, such that many organoiodin compounds turn yellow when stored over time because of decomposition into iodine elements; thus, they are commonly used in organic synthesis, due to the easy formation and cleavage of C-I bonds. They are also significantly denser than other organohalogen compounds due to the high iodine atoms. Some organic oxidizing agents such as iodan contain iodine in the oxidation state higher than -1, such as 2-iodoxybenzoic acid, general reagents for the oxidation of alcohols to aldehydes, and iodobenzene dichloride (PhICl 2 ), used for chlorination selective alkene and alkene. One of the more well-known uses of organoiodin compounds is the so-called iodoform test, in which iodoform (CHI 3 ) is produced by complete iodination of methyl ketone (or other oxidizable compound to methyl ketone) as follows:
Some of the disadvantages of using organoiodin compounds as compared to organochlorine compounds or organobromins are the greater cost and toxicity of the iodine derivatives, because expensive iodine and organoiodin compounds are stronger alkylation agents. For example, iodoacetamide and iodoacetic acid denaturation proteins with irreversible alkyl cysteine ​​residues and prevent disulfide relation reform.
The exchange of halogens to produce iodoalkanes by Finkelstein's reaction is somewhat complicated by the fact that iodides are a leaving group better than chloride or bromide. The difference is small enough that the reaction can be moved to completion by exploiting the differential solubility of the halide salt, or by using a large excess of the halide salt. In classical Finkelstein's reaction, the alkyl chloride or alkyl bromide is converted to alkyl iodide by treatment with sodium iodide solution in acetone. Sodium iodide dissolves in acetone and sodium chloride and no sodium bromide. The reaction is driven toward the product by mass action due to the deposition of insoluble salts.
Genesis and production
Iodine is the least abundant of stable halogens, comprising only 0.46 parts per million of rocks of the earth's crust (compare: fluorine 544 ppm, chlorine 126 ppm, 2.5 ppm bromine). Among the eighty-four elements that occur in significant quantities (elements 1-42, 44-60, 62-83, and 90-92), it ranks sixty-one in abundance. Rare mineral iodides, and most of the concentrated deposits for economical extraction are mineral iodates instead. Examples include seaarite, Ca (IO 3 ) 2 , and dietzeit, 7Ca (IO 3 ) 2 Ã, Â · 8CaCrO 4 . This is a mineral that occurs as a trace of dirt in caliche, found in Chile, whose main product is sodium nitrate. In total, they can contain at least 0.02% and at most 1% iodine. Sodium iodate is extracted from caliche and reduced to iodide by sodium bisulfite. This solution is then reacted with a new extracted iodate, resulting in a compression of iodine, which can be filtered.
The caliche was the main source of iodine in the 19th century and continues to be important today, replacing seaweed (which is no longer an economically viable source), but by the end of the 20th century salt water emerged as a comparable source. Japan's Minami Kanto gas field east of Tokyo and the American gas field Anadarko Basin in northwestern Oklahoma are the two largest sources. Salt water is hotter than 60 Â ° C from source depth. The salt water is first purified and acidified using sulfuric acid, the existing iodide is oxidized to iodine with chlorine. Iodine solution is produced, but it is dilute and must be concentrated. Air is exhaled into the solution to evaporate iodine, which is passed to an absorbent tower where sulfur dioxide reduces iodine. Hydrogen iodide (HI) is reacted with chlorine to precipitate iodine. After filtration and purification, iodine is packed.
- 2 HI Cl 2 -> I 2 ? 2 HCl
- I 2 2 H 2 O SO 2 -> 2 HI H 2 SO 4
- 2 HI Cl 2 -> I 2 ? 2 HCl
These sources ensure that Chile and Japan are the largest iodine producers today. Alternatively, salt water can be treated with silver nitrate to precipitate iodine as silver iodide, which is then described by reaction with iron to form metallic silver and iron (II) iodide solution. Iodine can then be liberated by displacement with chlorine.
Apps
Unlike chlorine and bromine, which has one major significant use dwarfs everything else, iodine is used in many applications of varying importance. About half of all iodine is produced into various organoiodin compounds; The remaining 15% as a pure element, another 15% is used to form potassium iodide, and another 15% for other inorganic iodine compounds. 5% left for small usage. Among the main uses of iodine compounds are catalysts, animal feed supplements, stabilizers, dyes, dyes and pigments, pharmaceuticals, sanitation (from iodine tincture), and photography; minor uses include smoke haze inhibition, cloud seeding, and various uses in analytical chemistry.
Analysis
Potassium tetraiodomercurate (II), K 2 HgI 4 , also known as Nessler reagent. It is often used as a sensitive spot test for ammonia. Similarly, Cu 2 HgI 4 is used as the originator to test the alkaloids. Anion iodide and iodate are often used for quantitative volumetric analysis, for example in iodometry and iodine clock reaction (where iodine also serves as a test for starch, forming dark blue complexes), and an aqueous iodoidal alkaline solution is used in the iodoform test for methyl ketone. Iodine test for flour is still used to detect counterfeit banknotes printed on paper containing starch.
Medicine
Elemental iodine
The iodine element is used as a disinfectant either as an element, or as a water-soluble anion triiodide I 3 - is generated in situ by adding iodide to elemental iodine which is poorly soluble in water (the opposite chemical reaction makes some free iodine elements available for antisepsis). The iodine element may also be used to treat iodine deficiency.
Alternatively, iodine can be produced from iodophores, containing iodine complexed with a dissolving agent (iodide ions can be considered loose as an iodophor in a triiodide water solution). Examples of such preparations include:
- Iodine tincture: iodine in ethanol, or iodine and sodium iodide in a mixture of ethanol and water.
- Lugol's iodine: iodine and iodide in water alone, forming most of the triiodide. Unlike iodine tincture, Lugol iodine has a smaller amount of free iodine component (I 2 ).
- Povidone iodine (an iodophor).
The action of iodine antimicrobial is fast and works at low concentrations, and thus is used in the operating room. The specific action mode is unknown. Penetrating to microorganisms and attacking certain amino acids (such as cysteine ​​and methionine), nucleotides, and fatty acids, eventually result in cell death. It also has antiviral action, but nonlipid and parvovirus viruses are less sensitive than those covered by lipids. Iodine may attack surface proteins from wrapped viruses, and this can also damage the stability of fatty acids by reacting with unsaturated carbon bonds.
Other formats
In the medical world, saturated potassium iodide solution is used to treat acute thyrotoxicosis. It is also used to block iodine-131 uptake in the thyroid gland (see isotope section above), when this isotope is used as part of a radiopharmaceutical (such as iobenguane) that is not targeted to the thyroid or thyroid tissue.
Iodine-131 (usually an iodide) is a component of nuclear fallout, and is very dangerous because of the tendency of the thyroid gland to concentrate iodine to swallow and retain it for longer periods of radiological half-life of this isotope eight days. For this reason, people at risk of exposure to radioactive iodine environment (iodine-131) in the fallout may be instructed to take non-radioactive potassium iodide tablets. The usual adult dose is one tablet of 130 mg per 24 hours, providing 100% iodine ion (100,000 micrograms). (The daily dose of iodine common to normal health is to order 100 micrograms; see "Intake Diet" below.) This large dosage of non-radioactive iodine minimizes the absorption of radioactive iodine by the thyroid gland.
As an element with high electron density and atomic number, iodine absorbs X-rays weaker than 33.3 keV because of the deepest electronelectric photoelectric effect. Organoiodin compounds are used with intravenous injection as X-ray radiocontrast agents. These applications are often associated with advanced X-ray techniques such as angiography and CT scans. Currently, all radiocontrast agents that are water soluble depend on iodine.
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Production of ethylenediamine dihydroiodide, provided as a nutritional supplement for livestock, consumes most of the available iodine. Another significant use is the catalyst for the production of acetic acid by the Monsanto and Cativa processes. In this technology, which supports world demand for acetic acid, hydroiodic acid converts methanol feedstock into methyl iodide, which undergoes carbonylation. The resulting hydrolysis of acetyl iodide regenerates the hydroiodic acid and gives acetic acid.
Inorganic iodine finds special uses. Titanium, zirconium, hafnium, and thorium are purified by the Arkel van process, which involves the formation of a reversible tetraiodide of these elements. Silver iodide is the main ingredient for traditional photographic films. Thousands of kilograms of silver iodide are used every year for cloud seeding to induce rain.
The organoiodin compound of erythrocyte is an important food coloring agent. Perfluoroalkyl iodide is a precursor for important surfactants, such as perfluorooctanesulfonic acid.
The role of biology
Iodine is an essential element for life and is the toughest element normally required by living organisms. (Lanthanum and other lanthanides, as well as tungsten, are used by some microorganisms.) It is necessary for the synthesis of thyroid hormones that regulate the growth of thyroxine and triiodothyronine (T 4 and T 3 respectively, named after their number of iodine atoms). Iodine deficiency leads to a decrease in the production of T 3 and T 4 and the concurrent enlargement of the thyroid tissue in an attempt to obtain more iodine, leading to a disease known as simple goitre. The main form of thyroid hormone in the blood is thyroxine (T 4 ), which has a longer half-life than T 3 . In humans, the ratio of T 4 to T 3 is released into the blood is between 14: 1 and 20: 1. T 4 is converted to T 3 is active (three to four times stronger than T 4 ) inside the cell by deiodinases (5'-iodinase). This is further processed by decarboxylation and deiodination to produce iodothyronamine (T 1 a) and thyronamine (T 0 a '). The third isoform deiodinase is an enzyme containing selenium; so the selenium diet is very important for the production of T 3 .
Iodine accounts for 65% of the molecular weight of T 4 and 59% T 3 . Fifteen to 20 mg of iodine is concentrated in the thyroid and hormonal tissues, but 70% of all iodine in the body is found in other tissues, including mammary glands, eyes, gastric mucosa, fetal thymus, cerebro-spinal fluid and choroid plexus, cervix, and salivary glands. In these tissue cells, the iodide enters directly by the sodium-iodide symporer (NIS). The act of iodine in the mammary tissue is related to fetal and neonatal development, but on other tissues, it is (at least) partially unknown.
food intake
Recommendations by the United States Institute of Medicine are between 110 and 130 Ã,Âμg for infants up to 12 months, 90 Ã,Âμg for children up to eight years, 130 Ã,Âμg for children up to 13 years, 150 Ã,Âμg for adults, 220Ã,Âμg for pregnant women and 290Ã,Âμg for lactation. Intake Tolerable (UL) level for adults is 1,100 Âμg/day. This upper limit is assessed by analyzing the effects of supplementation on the thyroid stimulating hormone.
The thyroid gland requires no more than 70 μg/day to synthesize the daily amount of T4 and T3 required. Higher recommended daily iodine removal levels are required for optimal functioning of a number of body systems, including breast-feeding, gastric mucosa, salivary glands, brain cells, choroid plexus, thymus, and artery walls.
Natural sources of dietary iodine include seafood, such as fish, seaweed (such as seaweed) and shellfish, dairy products and eggs as long as animals receive enough iodine, and plants grow on iodine-rich soils. The iodized salt is enriched with iodine in the form of sodium iodide.
In 2000, the median intake of iodine from food in the United States was 240 to 300 g/day for men and 190 to 210 g/day for women. The general US population has sufficient iodine nutrition, with women of childbearing age and pregnant women having a possible risk of mild deficiency. In Japan, consumption is considered much higher, ranging from 5,280? G/day up to 13,800 g/day of seaweed or kombu seaweed, often in the form of Kombu Umami extract for soup and potato chips. However, the new study shows that Japanese consumption is closer to 1,000-3,000 g/day. The adult UL in Japan was last revised to 3,000 Ã,Âμg/day by 2015.
After iodineation fortification programs such as iodisation of salt have been implemented, some cases of iodine induced hyperthyroidism have been observed (called Jod-Basedow phenomenon). This condition appears to occur mainly in people over forty, and the risk appears higher when severe iodine deficiency and early increase in intake of high iodine.
Disadvantages
In areas where there is a small amount of iodine in food, typically remote rural areas and semi-arid equatorial climate where no seafood is eaten, iodine deficiency leads to hypothyroidism, symptoms of extreme fatigue, mumps, mental retardation, depression, gain weight, and low basal body temperature. Iodine deficiency is the leading cause of intellectual disability that can be prevented, the result that occurs especially when infants or small children are given hypothyroidism by a lack of elements. The addition of iodine to table salt has largely eliminated this problem in rich countries, but iodine deficiency remains a serious public health problem today in developing countries. Iodine deficiency is also a problem in certain areas of Europe. Information processing, fine motor skills, and visual problem solving are enhanced by iodine filling in children with iodine deficiency.
Toxicity
The iodine element (I 2 ) is toxic if taken orally. The lethal dose for adult humans is 30 mg/kg, which is about 2.1-2.4 grams for humans weighing 70 to 80 kg (even if experiments in mice indicate that these animals can survive after eating 14,000 mg/dose kg) Excess iodine can be more cytotoxic in the presence of selenium deficiency. Iodine supplementation in the selenium deficiency population, in theory, is problematic, partly for this reason. Toxicity comes from its oxidizing properties, which weaken proteins (including enzymes).
The iodine element is also a skin irritant, and direct skin contact can cause damage and the solid iodine crystals must be handled with care. Solutions with high concentrations of iodine concentrations, such as iodine and Lugol tincture solutions, can cause tissue damage if used in prolonged cleaning or antisepsis; likewise, Povidone-iodine (Betadine) fluid trapped on the skin causes chemical burns in some cases reported.
Work exposure
People can be exposed to iodine at work through inhalation, consumption, skin contact, and eye contact. Occupational Safety and Health Administration (OSHA) has set a legal limit (Exposure Limit allowed) for iodine exposure at work at 0.1 ppm (1 mg/m 3 ) for 8 hours. The National Institute for Occupational Safety and Health (NIOSH) has set a Recommended Exposure limit (REL) of 0.1Ã, ppm (1 mg/m 3 ) for 8 hours. At a rate of 2 ppm, iodine is harmful to life and health.
Allergic reactions
Some people develop hypersensitivity to iodine-containing products and foods. Application of iodine or Betadine tincture can cause a rash, sometimes severe. The parenteral use of iodine contrast agents (see above) may cause reactions ranging from mild rash to fatal anaphylaxis. Such reactions have caused misunderstanding (widely held, even in between
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