Fluorine is a chemical element with the symbol F and the atomic number 9. This is the lightest halogen and exists as pure yellow diatomic gas that is highly toxic under standard conditions. As the most electronegative element, it is highly reactive: almost all other elements, including some noble gases, form compounds with fluorine.
Among the elements, fluor is ranked 24th in universal abundance and 13 in terrestrial abundance. Fluorite, the main fluorine mineral source that gives the element its name, was first described in 1529; such as those added to metal ores to lower their melting point for smelting, Latin verb fluo meaning "flow" giving mineral name. Proposed as an element in 1810, fluorine proved difficult and dangerous to separate from its compounds, and some early investigators died or injured suffered from their efforts. New in 1886 the French chemist Henri Moissan isolated the fluorine element using low-temperature electrolysis, a process still used for modern production. The production of the fluorine gas industry for uranium enrichment, its largest application, began during the Manhattan Project in World War II.
Because of the cost of purifying pure fluorine, most commercial applications use fluorine compounds, with about half of the mined fluorite used in steelmaking. The rest of the fluorite is converted into corrosive hydrogen fluoride on the way to various organic fluorides, or becomes a cryolite that plays a key role in aluminum refining. Organic fluoride has very high chemical and thermal stability; Their main uses are as cooling, electrical insulation and cooking utensils, the latter being PTFE (Teflon). Drugs such as atorvastatin and fluoxetine also contain fluorine, and fluoride ions inhibit cavities, and thus find use in toothpaste and water fluoridation. The number of global fluorochemical sales reaches more than US $ 15 billion per year.
Fluorocarbon gas is generally a greenhouse gas with a global warming potential of 100 to 20,000 times that of carbon dioxide. Organofluorin compounds survive in the environment due to the strength of carbon-fluorine bonds. Fluor has no known metabolic role in mammals; some plants synthesize organofluorine toxins that block herbivores.
Video Fluorine
Characteristics
Electron configuration
Fluorine atoms have nine electrons, one less than neon, and electron configuration 1 2 2 2 2p 5 : two electrons in the inner skin are filled and seven in the outer shell that need one more to fill. The outer electrons are ineffective on the nuclear shield, and experience a high effective nuclear charge of 9-2 = 7; this affects the physical properties of atoms.
The first ionisation energy of fluorine is the third highest among all elements, behind helium and neon, which complicates the removal of electrons from neutral fluorine atoms. It also has a high electron affinity, second only to chlorine, and tends to capture electrons to be isoelectronic with a noble fluorescent gas; it has the highest electronegativity of each element. The fluorine atom has a small covalent radius of about 60 pixometers, similar to its neighbors with oxygen and neon.
Reactivity
Difluorine bond energy is much lower than Cl
2 or Br
2 and similar to peroxide bonds that are easy to split; This, along with high electronegativity, contributes to fluorine easy dissociation, high reactivity, and strong bonds for non-fluorine atoms. In contrast, bonds to other atoms are very strong because of the high electronegativity of fluorine. Non-reactive substances such as powdered steel, broken glass, and asbestos fibers react quickly with cold fluorine gas; wood and water spontaneously burning under a fluorine jet.
The reaction of fluorine element with metal requires various conditions. Alkali metals cause explosions and alkaline earth metals exhibit vigorous activity in large quantities; to prevent passivation from formation of metal fluoride coatings, most other metals such as aluminum and iron should be powdered, and precious metals require pure fluorine gas at 300-450 ° C (575-850 ° F). Some solid nonmetals (sulfur, phosphorus) react strongly in fluorine liquid air temperatures. Hydrogen sulfide and sulfur dioxide easily combine with fluorine, the latter sometimes explosive; sulfuric acid exhibits much less activity, requiring high temperatures.
Hydrogen, like some alkali metals, reacts explosively with fluorine. Carbon, as a black lamp, reacts at room temperature to produce fluoromethane. Graphite combines fluorides above 400 ° C (750 ° F) to produce non-stoichiometric carbon monofluoride; higher temperatures produce gas fluorocarbons, sometimes with explosions. Carbon dioxide and carbon monoxide react at or just above room temperature, whereas paraffin and other organic chemicals produce a strong reaction: even completely substituted haloalkanes such as carbon tetrachloride, usually unburnt, can explode. Although nitrogen trifluoride is stable, nitrogen requires the release of electricity at high temperatures for the reaction with fluorine to occur, due to the very strong triple bond in the nitrogen element; ammonia can react explosively. Oxygen does not combine with fluorine under ambient conditions, but can be made to react using electricity release at low temperature and pressure; the product tends to crumble into its constituent elements when heated. Heavier halogens react with fluorine as well as noble radon gas; other noble gases, only xenon and krypton react, and only under special conditions.
Phase
At room temperature, fluorine gas is a diatomic molecule, pale yellow when pure (sometimes described as yellow-green). It has a distinctive sharp odor detected at 20 ppb. Fluorides condense into a bright yellow liquid at -188 ° C (-306 ° F), a transition temperature similar to oxygen and nitrogen.
Fluor has two solid forms ,? - and? -fluorine. The latter crystallizes at -220 ° C (-364 ° F) and is transparent and soft, with the irregular cube structure of newly crystallized solid oxygen, unlike the orthorhombic system of other solid halogens. Further cooling up to -228 Â ° C (-378 Â ° F) induces phase transitions to be blurry and loud? -fluorine, which has a monoclinic structure with solid and angular molecular layers. Transition from? - to? -fluorine is more exothermic than fluorine condensation, and can be violent.
Isotope
Only one fluorine isotope occurs naturally in abundance, stable isotope 19 > F . It has a high magnetogical ratio and exceptional sensitivity to the magnetic field; because it is also the only stable isotope, used in magnetic resonance imaging. Seventeen radioisotopes with mass numbers from 14 to 31 have been synthesized, which is 18 The F is the most stable with a half-life of 109.77 minutes. Other radioisotopes have a half-life of less than 70 seconds; most of the decomposition in less than half a second. Isotopes 17 F and < span> 18 F live? electron decay and arrest, lighter isotope decay by proton emission, and heavier than 19 F experience? - decay (which is the heaviest with delayed neutron emissions). Two measurable fluorine isomers are known, 18m
F , with a half-life of 162 (7) nanoseconds, and 26 m
F , with a half-life of 2.2 (1) milliseconds.
Maps Fluorine
Genesis
Universe
Among the light elements, the fluorine abundance value of 400 ppb (parts per billion) - 24 among the elements in the universe - is very low: the other elements of carbon into magnesium are twenty times or more as normal. This is because the nucleosynthetic process of the star passes fluorine, and every fluorine atom formed has a high nuclear cross section, allowing further fusion with hydrogen or helium to produce oxygen or neon respectively.
Beyond this temporary existence, three explanations have been put forward for the presence of fluorine:
- during a Type II supernova, neon bombardment of a neon atom by a neutrino can transmute it to fluorine;
- the solar wind from Wolf-Rayet stars can blow away fluorine from hydrogen or helium atoms; or fluorine is borne by convection currents arising from fusion of giant asymptotic star branches.
Earth
Fluor is the thirteenth most common element in the Earth's crust at 600-700 ppm (parts per million) by mass. The fluorine element in the Earth's atmosphere will readily react with water vapor in the atmosphere, blocking its natural occurrence; it is found only in the form of combined minerals, in which fluorite, fluorapatite and cryolite are the most important in the industry. Fluorite or fluorspar ( CaF
2 ), colorful and abundant throughout the world, is a major source of fluorine; China and Mexico are the main suppliers. US-led extraction at the beginning of the 20th century but stopped mining in 1995. Despite the fluorapatite (Ca 5 (PO 4 ) 3 F) contains most of the world's fluorine, a low mass fraction of 3.5% means it is mostly used as a phosphate. In the US, small amounts of fluorine compounds are obtained through fluorosilicic acid, a by-product of the phosphate industry. Cryolite ( Na
3 AlF > 6 ), after being used directly in aluminum production, is the rarest and concentrated from these three minerals. The main commercial mine on the west coast of Greenland was closed in 1987, and most of the cryolites are now synthesized.
Other minerals such as topazs contain fluorine. Fluoride, unlike other halides, is insoluble and does not occur at commercially advantageous concentrations in saline waters. The number of traces of organofluorin uncertain origin has been detected in volcanic eruptions and geothermal springs. The presence of fluorine gas in crystals, suggested by a crumbling, controversial antozonite odor; a 2012 study reported a 0.04% F
2 weight in antozonite, connecting this inclusion with radiation from the presence of a small amount of uranium.
History
Initial discovery
In 1529, Georgius Agricola described fluorite as an additive used to lower the metal melting point during melting. He wrote Latin fluorÃÆ'Â © s ( fluor, flow) for fluorite rocks. It later evolved into fluorspar (still commonly used) and then fluorite . The fluorite composition is then determined to be calcium difluoride.
Hydrofluoric acid is used in etching glass from 1720 onwards. Andreas Sigismund Marggraf was first characterized in 1764 when he heated up fluorite with sulfuric acid, and the resulting solution rusted his glass container. The Swedish chemist Carl Wilhelm Scheele repeated the experiment in 1771, and named the acidic product [fluss-spats-syran] (fluorspar acid). In 1810, the French physicist AndrÃÆ' © -Marie AmpÃÆ'¨re stated that hydrogen and an element analogous to chlorine form hydrofluoric acid. Sir Humphry Davy proposes that this currently unknown substance is named fluorine from the fluorine acid and the suffix -in of the other halogen. This word, with modification, is used in most European languages; Greek, Russian, and a few others (following Ampè¨re suggestions) using the name ftor or a derivative, from the Greek ??????? ( phthorios , damaging). New Latin Name fluorum gives its current symbol element F ; Fl is used in the initial paper.
Isolation
Preliminary studies of fluorine are so dangerous that some nineteenth-century experiments are considered "fluorine martyrs" after adversity with hydrofluoric acid. Isolation of the fluorine element is blocked by the extreme corrosion of both the fluorine element itself and hydrogen fluoride, as well as the lack of a simple and appropriate electrolyte. Edmond FrÃÆ' Â © I postulated that pure hydrogen fluoride electrolysis to produce fluorine was feasible and found a method for producing an anhydride sample of acidified potassium bifluoride; instead, he found that the fluoride fluoride (dried) produced did not conduct electricity. My former student Henri Moissan survived, and after much trial and error found that the mixture of potassium bromluoride and fluoride dried fluoride were conductors, which allowed electrolysis. To prevent rapid corrosion of platinum in electrochemical cells, it cools the reaction to very low temperatures in special baths and forged cells of a more resilient platinum and iridium mixture, and uses a fluorite plug. In 1886, after 74 years of effort by many chemists, Moissan isolated the fluorine element.
In 1906, two months before his death, Moissan received the Nobel Prize in Chemistry, with the following quote:
[I] recognize the great service it provides in its investigation and fluorine elemental isolation... The whole world has admired the great experimental skills that you have learned about the beast among the elements.
Later using
The Frigidaire General Motors (GM) division experimented with refrigerant chlorofluorocarbons in the late 1920s, and Kinetic Chemicals was formed as a joint venture between GM and DuPont in 1930 in the hope of marketing Freon-12 ( CCl
2 F
2 ) as one of the refrigerants. It replaces earlier and more toxic compounds, increases in demand for kitchen refrigerators, and becomes profitable; in 1949 DuPont had purchased Kinetik and marketed several other Freon compounds. Polytetrafluoroethylene (Teflon) was coincidentally discovered in 1938 by Roy J. Plunkett while working on refrigerant in Kinetic, and its chemical resistance and superlative heat gave it to accelerated commercialization and mass production in 1941.
The large production of fluorine elements began during World War II. Germany uses high-temperature electrolysis to make tons of chloride trifluoride burned and the Manhattan Project uses large quantities to produce uranium hexafluoride for uranium enrichment. Because UF
6 is as corrosive as fluorine, gas diffusion plants require special materials: nickel for membranes, fluoropolymers for seals, and liquid fluorocarbons as coolants and lubricants. This burgeoning nuclear industry then promotes the development of postwar fluorochemistry.
​​Compound
Fluor has a rich chemical, including organic and inorganic domains. It combines with metal, not metal, metaloid, and the most noble gases, and usually assumes a -1 oxidation state. The high electron affinity of fluorine produces a preference for ionic bonding; when forming a covalent bond, it is polar, and almost always singular.
Metal
Alkali metals form ionic and highly soluble monofluorides; It has a cubic arrangement of sodium chloride and an analog chloride. Alkaline earth diffluorides have strong ionic bonds but are not soluble in water, with the exception of beryllium difluoride, which also exhibits some covalent characters and has quartz-like structures. Rare earth elements and many other metals form most of the ionic trifluorides.
Covalent bonds first appear in tetrafluoride: zirconium, hafnium, and some ionic actinics with high melting point, while titanium, vanadium, and niobium are polymeric, melting or decomposing at no more than 350 ° C (660 ° C). Â ° F). Pentafluoride continues this trend with the linear polymer and its oligomer complex. Thirteen metal hexafluorides are known, all octahedral, and most are volatile solids but for liquids MoF
6 and ReF
6 , and WF
6 . Rhenium heptafluoride, the only characterized metal heptafluoride, is a low molecular solid melt with pentagonal bipyramidal molecular geometry. Metal fluoride with more fluorine atoms is highly reactive.
Hydrogen
Hydrogen and fluorine combine to produce hydrogen fluoride, in which separate molecules form clusters by hydrogen bonds, resembling water over hydrogen chloride. It boils at much higher temperatures than the heavier hydrogen halides and unlike them completely dissolved with water. Hydrogen fluoride easily hydrates contact with water to form an aqueous hydrogen fluoride, also known as hydrofluoric acid. Unlike other hydrohalic acids, which are strong, fluoride acids are weak acids at low concentrations. However, it can strike glass, something no other acid can do.
Other reactive non-metals
- Metalloids are included in this section
The binary fluoride of the metalloid and the nonmetallic p-block are usually covalent and volatile, with various reactivity. The heavier 3 and non-metallic periods may form a hypervalent fluoride.
Boron trifluoride is planar and has an incomplete octet. It serves as Lewis acid and combines with Lewis bases such as ammonia to form adducts. Tetrafluoride carbon is tetrahedral and inert; its group analogs, silicon and germanium tetrafluoride, also tetrahedral but behave like Lewis acid. Pnictogen forms trifluoride which increases reactivity and alkalinity with higher molecular weights, although nitrogen trifluoride counteracts hydrolysis and is not fundamental. Pentaphluoride phosphorus, arsenic, and antimony are more reactive than the trifluoride respectively, with the strongest known neutral Lewis pentafluoride antimony.
Chalcogens have various fluorides: unstable difluorides have been reported for oxygen (the only known compounds with oxygen under oxidation state 2), sulfur, and selenium; tetrafluoride and hexafluoride present for sulfur, selenium, and tellurium. The latter is stabilized by fluorine atoms and lighter central atoms, so the sulfur hexafluoride is very inert. Chlorine, bromine and iodine can each form mono-, tri-, and pentafluoride, but only iodine-heptafluoride has been characterized among possible interhalogenous heptafluorides. Many of them are powerful sources of fluorine atoms, and industrial applications using trifluoride chlorine require precautions similar to those that use fluorine.
Precious Gas
The noble gases, having complete electron shells, opposed reactions with other elements until 1962 when Neil Bartlett reported the synthesis of xenon hexafluoroplatinate; xenon difluoride, tetrafluoride, hexafluoride, and some oxybluorides have been isolated since then. Among other noble gases, krypton forms difluoride, and radon and fluorine produce a suspected solid of radon difluoride. Binary fluoride from lighter gaseous gases is highly unstable: argon and hydrogen fluoride join in extreme conditions to give argon fluorohydride. Helium and neon have no long-lived fluoride, and no fluoride fluid has ever been observed; helium fluorohydride has been detected for milliseconds at high pressure and low temperature.
Organic compound
The carbon-fluorine bond is the strongest organic chemical, and gives stability to the organofluorine. It is virtually non-existent in nature, but is used in artificial compounds. Research in this area is usually driven by commercial applications; The compounds involved vary and reflect the inherent complexity of organic chemistry.
Discrete molecule
Substitution of a hydrogen atom in an alkane by more fluorine atoms gradually changes some properties: melting and boiling points are lowered, density increases, solubility in hydrocarbons decreases and overall stability increases. Perfluorocarbons, in which all hydrogen atoms are replaced, insoluble in most organic solvents, react to ambient conditions only with sodium in liquid ammonia.
The term perfluorinated compounds is used for what should be perfluorocarbons if not for the presence of functional groups, often a carboxylic acid. These compounds share many properties with perfluorocarbons such as stability and hydrophobicity, whereas functional groups add to their reactivity, allowing them to stick to surfaces or act as surfactants; Fluorosurfactants, in particular, can lower water surface tension over their hydrocarbon analogs. Fluorotelomers, which have several unfluorinated carbon atoms near the functional group, are also considered perfluorinated.
Polymers
The polymer exhibits the same stability enhancement given by fluorine substitution (for hydrogen) in discrete molecules; Their melting points generally increase as well. Polytetrafluoroethylene (PTFE), the simplest fluoropolymer and perfluoro analogue of polyethylene with structural units - CF
2 -, shows this change as expected, but its very high melting point makes it difficult to form. PTFE derivatives are less temperature tolerant but easier to form: fluorinated propylene ethylene replaces some fluorine atoms with trifluoromethyl groups, the perfluoroalkoxy alkane does the same with the trifluorometoxy group, and Nafion contains perfluoroethane side chains covered with sulfonic acid groups. Other fluoropolymers retain some of the hydrogen atoms; polyvinylidene fluoride has half the fluorine atoms of PTFE and polyvinyl fluoride has a quarter, but both behave like perfluorinated polymers.
Production
Industrial
The Moissan method is used to produce fluorine in industrial quantities, through electrolysis of a mixture of potassium fluoride/hydrogen fluoride: hydrogen and fluoride ions are reduced and oxidized to the cathode of steel containers and carbon block anodes, below 8-12 volts, to produce hydrogen and fluorine gases respectively, respectively. Temperature increases, KFo2HF melts at 70 ° C (158 ° F) and is electrolysed at 70 ° -30 ° C (158 ° -26 ° F). KF, which acts as a catalyst, is essential because pure HF can not be electrolyzed. Fluor can be stored in a steel cylinder that has a passive interior, at temperatures below 200 ° C (392 ° F); otherwise, nickel can be used. The regulator valve and pipework are made of nickel, the latter probably using Monel instead. Frequent passivation, along with water exceptions and tight grease, should be done. In the laboratory, glassware can carry fluorine gas under low pressure and anhydrous conditions; some sources recommend the nickel-monel-PTFE system.
Chemistry
While preparing for the 1986 conference to celebrate the hundredth anniversary of the achievement of Moissan, Karl O. Christe argued that the generation of fluorine chemistry should be viable because some metal fluoride anions do not have stable neutral partners; their acidification has the potential to trigger oxidation. He invented a method that developed fluorine on high yield and atmospheric pressure:
- 2 KMnO 4 2 KF 10 HF 3 H 2 O 2 -> 2 K 2 MnF 6 8 H 2 O 3 O 2 ?
- 2 K 2 MnF 6 4 SbF 5 -> 4 KSbF 6 2 MnF 3 F 2 ?
Christe later commented that the reactants "have been known for over 100 years and even Moissan could come up with this scheme." Until late 2008, some references still confirm that fluorine is too reactive for chemical isolation.
Industrial applications
Fluorite mining, which supplies most of the global fluorine, peaked in 1989 when 5.6 million metric tons of ore were extracted. The limitation of chlorofluorocarbons decreased this to 3.6 million tons in 1994; production has increased. Approximately 4.5 million tonnes of ore and revenues of US $ 550 million were generated in 2003; the report then estimates 2011 global fluorochemical sales of $ 15 billion and forecasts for 2016-18 production figures of 3.5 to 5.9 million tons, and revenues of at least $ 20 billion. Flotation froth separates the fluorite mined into two main metallurgical classes with the same proportion: 60-85% pure metspar is almost all used in iron smelting whereas 97% pure acids are usually converted to hydrogen intermediates of the hydrogen industry.
At least 17,000 metric tons of fluorine are produced each year. It costs only $ 5-8 per kilogram as uranium or sulfur hexafluoride, but many times more as an element as it handles challenges. Most processes use large amounts of free fluoride using in situ generation under vertical integration.
The largest application of fluorine gas, consuming up to 7,000 metric tons per year, is in preparation of UF
6 for the nuclear fuel cycle. Fluor is used to fluorize uranium tetrafluoride, which is formed from uranium dioxide and hydrofluoric acid. Fluor is monoisotopic, so there is a mass difference between UF
6 molecule is caused by the presence of 235
U or 238
U , enabling uranium enrichment through gas diffusion or gas centrifuge. Approximately 6,000 metric tons per year is used to produce dielectric inertrics SF
6 for high voltage transformers and circuit breakers, eliminating the need for dangerous polychlorinated biphenyls associated with oil-filled devices. Some fluorine compounds are used in electronics: rhenium and tungsten hexafluoride in chemical vapor deposition, tetrafluoromethane in plasma etching and nitrogen trifluoride in cleaning apparatus. Fluor is also used in organic fluoride synthesis, but its reactivity often necessitates the first conversion to a softer ClF
3 , Brf
3 , or IF
5 , which together allow for calibrated fluorination. Fluorated drugs use sulfur tetrafluoride instead. Inorganic Fluor
Like other iron alloys, about 3 kg (6.5 lb) metspar is added to every metric ton of steel; fluoride ions decrease the melting point and its viscosity. In addition to its role as an additive in materials such as enamel and welded rod layers, most acidpar is reacted with sulfuric acid to form hydrofluoric acid, used in steel preservation, glass etching, and alkane cracking. One-third of HF goes into the synthesis of cryolite and aluminum trifluoride, both of which are flux in the Hall-HÃÆ' Â © roult process for aluminum extraction; replenishment is required by their occasional reactions with the melting apparatus. Each metric ton of aluminum requires about 23 kg (51 pounds) of flux. Fluorosilicate consumes the second largest part, with sodium fluorosilicate used in water fluoridation and laundry waste treatment, and as an intermediate on the way to cryolite and silicon tetrafluoride. Other important inorganic fluorides include cobalt, nickel, and ammonium.
Organic Fluoride
Organofluoride consumes more than 20% of the mined fluorite and more than 40% hydrofluoric acid, with cooling gases dominating and fluoropolymers increasing their market share. Surfactants are small applications but generate more than $ 1 billion in annual revenue. Due to the danger of direct hydrocarbon-fluorine reactions above -150Ã, Â ° C (-240Ã, Â ° F), indirect industrial fluorocarbon production, mostly through halogen exchange reactions such as Swarts fluorination, in which chlorine carbon is substituted for fluorine with hydrogen. fluoride under the catalyst. Substitution of electrochemical fluoroxy hydrocarbons for electrolysis in hydrogen fluoride, and the Fowler process treats them with dense fluorine carriers such as cobalt trifluoride.
Cooling Gas
Halogenated refrigerants, called Freons in an informal context, are identified by R-numbers indicating the amount of fluorine, chlorine, carbon, and hydrogen present. Chlorofluorocarbons (CFCs) such as R-11, R-12, and R-114 once dominated organofluorine, culminating in production in the 1980s. Used for air conditioning systems, propellants and solvents, their production was below one tenth of this peak in early 2000, following widespread international bans. Hydrochlorofluorocarbons (HCFCs) and hydrofluorocarbons (HFCs) are designed as substitutes; Their synthesis consumes more than 90% of fluorine in the organic industry. Important HCFCs include R-22, chlorodifluoromethane, and R-141b. The main HFC is R-134a with a new molecular type HFO-1234yf, Hydrofluoroolefin (HFO) which is notable for its global warming potential of less than 1% of HFC-134a.
Polymers
About 180,000 metric tons of fluoropolymers were produced in 2006 and 2007, generating over $ 3.5 billion in revenues per year. The global market is expected to be just under $ 6 billion in 2011 and is expected to grow by 6.5% annually through 2016. Fluoropolymers can only be formed with free radical polymerization.
Polytetrafluoroethylene (PTFE), sometimes referred to as DuPont Teflon, represents 60-80% of world fluoropolymer mass production. The biggest application in electrical insulation because PTFE is a very good dielectric. It is also used in the chemical industry where corrosion resistance is required, in pipe coating, tubing, and gaskets. Another major use is the PFTE-coated fiberglass fabric for the roof of the stadium. The main consumer app is for nonstick cookware. PTFE films snapped into an expanded PTFE (ePTFE), a fine porous membrane sometimes called the Gore-Tex brand name and used for raincoats, protective clothing, and filters; EPTFE fibers can be made into seals and dust filters. Other fluoropolymers, including fluorinated ethylene propylene, mimic PTFE properties and may replace them; they are more moldable, but also more expensive and have lower thermal stability. Films of two different fluoropolymers replace glass in a solar cell.
Chemically resistant (but expensive) fluorinated ionic is used as an electrochemical cell membrane, where the first and most prominent example is Nafion. Developed in the 1960s, it was originally used as a fuel cell material in a spacecraft and subsequently replaced mercury-based chloralkali cell processes. Recently, fuel cell applications have reappeared with efforts to install proton fuel cell membrane exchange into the car. Fluoroelastomers such as Viton is a crosslinked fluoropolymer mixture mainly used in O-ring; perfluorobutane (C 4 F 10 ) is used as a fire extinguisher.
Surfactant
Fluorosurfactants are small organofluorin molecules used to hold water and stains. While expensive (comparable to drugs at $ 200-2000 per kilogram), they generated over $ 1 billion in annual revenues in 2006; Scotchgard alone generated more than $ 300 million in 2000. Fluorosurfactants are a minority in the entire surfactant market, which is mostly taken up by cheaper hydrocarbon-based products. Application in the paint is burdened by the cost of compounding; this usage is only worth $ 100 million in 2006.
Agrichemicals
About 30% of agrichemicals contain fluorine, most of them herbicides and fungicides with some plant regulators. Fluorine substitution, usually of a single atom or at most trifluoromethyl groups, is a strong modification with effects analogous to fluorinated drugs: increased biological residence time, membrane crossings, and changes in molecular recognition. Trifluralin is a prominent example, with large-scale use in the US as a weedkiller, but it is a suspected carcinogen and has been banned in many European countries. Sodium monofluoroacetate (1080) is a mammalian toxin in which two acetic acid hydrogen are replaced by fluorine and sodium; it disrupts cell metabolism by replacing acetate in the citric acid cycle. First synthesized in the late 19th century, it was recognized as an insecticide at the beginning of 20, and then deployed in its current use. New Zealand, the largest consumer of 1080, uses it to protect kiwi from an invasive Australian toothbrush bag. Europe and the US have banned 1080.
Medical applications
Dental care
Population studies from the mid-20th century onwards showed topical fluoride reducing dental caries. This was first associated with conversion of tooth enamel hydroxyapatite into more durable fluorapatite, but research on pre-fluoride teeth denied this hypothesis, and current theory involves fluoride which helps the growth of enamel in small caries. Following the study of children in areas where fluoride is naturally present in drinking water, fluoride controlled public water supplies to fight tooth decay began in the 1940s and is now applied to water that supplies 6 percent of the global population, including two thirds of people America. A review of scientific literature in 2000 and 2007 relates to water fluoridation with significant tooth decline in children. Although support and evidence have no side effects other than most benign dental fluorosis, the opposition still exists on ethical and safety reasons. The benefits of fluoridation have been reduced, probably due to other sources of fluoride, but can still be measured in low-income groups. Sodium monofluorophosphate and sometimes sodium or tin (II) fluoride are often found in fluoride toothpaste, first introduced in the US in 1955 and now in developed countries, in addition to fluoride mouthwashes, gels, foams, and varnishes.
Pharmacy
Twenty percent of modern medicines contain fluorine. One of them, cholesterol-lowering atorvastatin (Lipitor), generates more income than other drugs until it becomes generic in 2011. The combination of Seretide asthma recipes, the top ten income drugs in the mid-2000s, contained two active ingredients, one of them - fluticasone - was fluorinated. Many fluorinated drugs to delay inactivation and prolong the dosing period because carbon-fluorine bonds are very stable. Fluorination also increases lipophilicity because bonding is more hydrophobic than carbon-hydrogen bonding, and this often helps in penetration of cell membranes and hence bioavailability.
Tricyclics and other antidepressants before 1980 had some side effects due to their non-selective disorder with neurotransmitters other than serotonin targets; fluorinated fluoxetine is selective and one of the first to avoid this problem. Many antidepressants currently receive this same treatment, including selective serotonin reuptake inhibitors: citalopram, escitalopram isomer, and fluvoxamine and paroxetine. Quinolone is a broad-spectrum antibiotic that is often fluorinated to increase its effect. These include ciprofloxacin and levofloxacin. Fluor also found use in steroids: fludrocortisone is a blood pressure boosting mineralocorticoid, and triamcinortone and dexamethasone are strong glucocorticoids. The majority of inhalation anesthetics are highly fluorinated; The prototype of halothane is much more inert and powerful than its contemporaries. Then compounds such as fluorinated fluorinated sevoflurane and desflurane are better than halothane and almost insoluble in the blood, allowing faster wake times.
PET scanning
Fluor-18 is often found in radioactive trackers for positron emission tomography, because the half-life of nearly two hours is long enough to allow transportation from the production facility to the imaging center. The most common tracker is fluorodeoxyglucose which, after intravenous injection, is taken up by tissues that require glucose such as the brain and most malignant tumors; Computer-assisted tomography can then be used for detailed imaging.
Operator oxygen
Liquid fluorocarbons can store large amounts of oxygen or carbon dioxide, more than blood, and have attracted attention to possible use in artificial blood and in respiratory fluids. Since fluorocarbons typically do not mix with water, they must be mixed into emulsions (tiny droplets of perfluorocarbons suspended in water) for use as blood. One such product, Oxycyte, has been through preliminary clinical trials. This substance can assist endurance athletes and is prohibited from sports; death near one of the cyclists in 1998 prompted an investigation of their abuse. Pure perfluorocarbon fluid breathing applications (which use pure perfluorocarbon fluids, not water emulsions) include helping burn victims and premature babies with deprived lungs. Partial and complete lung replenishment has been considered, although only the first one has significant tests in humans. The Alliance Pharmaceuticals effort achieved clinical trials but was abandoned because the results were no better than normal therapy.
The role of biology
Fluor is not important for humans or other mammals; Small amounts may be beneficial for bone strength, but this has not been determined for certain. Since there are many environmental sources of fluorine traces, the possibility of fluorine deficiency can only apply to artificial diets. Natural organofluorin has been found in microorganisms and plants but not animals. The most common is fluoroacetate, which is used as a defense against herbivores by at least 40 plants in Africa, Australia and Brazil. Other examples include limited fluorinated fatty acids, fluoroacetone, and 2-fluorocitrate. The enzyme that binds fluorine to carbon - adenosyl-fluoride synthase - was found in bacteria in 2002.
Toxicity
The fluorine element is highly toxic to living organisms. The effect on humans starts at lower concentrations of 50 ppm of hydrogen cyanide and is similar to chlorine: significant irritation of the eyes and respiratory system as well as liver and kidney damage occurs above 25 ppm, which is harmful to life and health. value for fluorine. Eyes and nose are seriously damaged at 100 ppm, and 1,000 ppm fluorine inhalation will cause death in minutes, compared to 270 ppm for hydrogen cyanide.
Hydrofluoric Acid
Hydrofluoric acid is a contact toxin with a greater danger than strong acids such as sulfuric acid although weak, it remains neutral in aqueous solutions and thus penetrates the tissue more quickly, either by inhalation, swallowing or skin, and at least nine US workers die. in an accident like 1984-1994. It reacts with calcium and magnesium in the blood causing hypocalcemia and possible death through cardiac arrhythmias. Insoluble formation of calcium fluoride triggers intense pain and burns greater than 160 cm 2 (25 in 2 ) can cause serious systemic poisoning.
Exposure may not be proven for eight hours for 50% HF, increased up to 24 hours for lower concentrations, and burns may initially not cause pain as hydrogen fluoride affects nerve function. If the skin has been affected by HF, the damage can be reduced by rinsing it under water for 10-15 minutes and removing contaminated clothing. Calcium gluconate is often applied next, providing calcium ions to bind with fluoride; Skin burns can be treated with 2.5% calcium gluconate gel or special flushing solution. Absorption of hydrofluoric acid requires further medical care; calcium gluconate can be injected or administered intravenously. Using calcium chloride - a common laboratory reagent - instead of calcium gluconate is contraindicated, and can cause severe complications. Excision or amputation of affected parts may be required.
Ion fluoride
Dissolved fluids are quite toxic: 5-10 g of sodium fluoride, or 32-64 mg of ionic fluoride per kilogram of body mass, is a lethal dose for adults. A fifth of the lethal doses can cause adverse health effects, and chronic excessive consumption can lead to bone fluorosis, affecting millions of people in Asia and Africa. Fluorate is submerged to form fluoride acid in the stomach that is easily absorbed by the intestine, where it crosses the cell membrane, binds with calcium and interferes with various enzymes, before urinary excretion. The limit of exposure is determined by the urine test of the body's ability to clean fluoride ions.
Historically, most cases of fluoride poisoning have been caused by accidental consumption of insecticides containing inorganic fluorine. Most current calls to the poison control center for possible fluoride intoxication come from the consumption of fluoride-containing toothpaste. Non-functioning water fluoridation equipment was another cause: an incident in Alaska affected nearly 300 people and killed one person. The dangers of toothpaste are aggravated for young children, and the Centers for Disease Control and Prevention recommends keeping an eye on children under six brushing their teeth so they do not swallow toothpaste. One regional study examined a year of pre-adolescent fluoride poisoning reports of 87 cases, including one death from ingestion of insecticide. Most have no symptoms, but about 30% have abdominal pain. A larger study across the US has similar findings: 80% of cases involve children under six years, and there are some serious cases.
Environmental issues
Atmosphere
The Montreal Protocol, signed in 1987, sets strict rules on chlorofluorocarbons (CFCs) and bromofluorocarbons due to the potential for ozone damage (ODP). High stability suitable for their native applications also means that they do not rot until they reach higher altitudes, where chlorine and bromine atoms are released to attack ozone molecules. Even with the ban, and an early indication of its efficacy, predictions warn that several generations will pass before full recovery. With one-tenth of ODP from CFCs, hydrochlorofluorocarbons (HCFCs) are the current substitutes, and themselves are scheduled for substitution by 2030-2040 by hydrofluorocarbons (HFC) without chlorine and zero ODP. In 2007 this date was brought forward to 2020 for developed countries; The Environmental Protection Agency has banned one HCFC production and closed two other products in 2003. Fluorocarbon gases are generally greenhouse gases with global warming potential (GWPs) of about 100 to 10,000; sulfur hexafluoride has a value of about 20,000. Outlier is HFO-1234yf which is a new type of refrigerant called Hydrofluoroolefin (HFO) and has attracted global demand due to its 4 GWP compared to 1430 for the current HFC-134a refrigerant standard.
Biopersistence
Organofluorin shows its biopersistence due to the strength of the carbon-fluorine bond. Perfluoroalkyl acid (PFA), which is slightly soluble in water due to their acid functional groups, is recorded persistent organic pollutants; perfluorooctanesulfonic acid (PFOS) and perfluorooctanoic acid (PFOA) are most frequently studied. PFAA has been found in small numbers worldwide from polar bears to humans, with PFOS and PFOA known to be in breast milk and newborn blood. The 2013 review shows little correlation between PFAA levels and human activity in ground water and humans; there is no clear pattern of one chemical that predominates, and higher PFOS numbers are correlated with higher PFOA numbers. In the body, PFAA binds proteins such as serum albumin; they tend to concentrate on humans in the liver and blood before excretion through the kidneys. The residence time in the body varies greatly by species, with live beings of rodents, and years in humans. High doses of PFOS and PFOA cause cancer and death in newborn mice but human studies have not determined the effect on current exposure levels.
See also
- Selective fluoride electrode, which measures fluoride concentration
- Determination of fluorine absorption
- 19 F NMR
- Fluorine chemistry, a process used to separate reagents from organic solvents
- Krypton and argon fluoride lasers
- Electrophilic and radical fluorilation
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Source of the article : Wikipedia
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