Among the elements, fluorine ranks 24th in cosmic abundance and 13th in crustal abundance. Fluorite, the primary mineral source of fluorine, which gave the element its name, was first described in 1529; as it was added to metal ores to lower their melting points for smelting, the Latin verb fluo meaning 'to flow' gave the mineral its name. Proposed as an element in 1810, fluorine proved difficult and dangerous to separate from its compounds, and several early experimenters died or sustained injuries from their attempts. Only in 1886 did French chemist Henri Moissan isolate elemental fluorine using low-temperature electrolysis, a process still employed for modern production. Industrial production of fluorine gas for uranium enrichment, its largest application, began during the Manhattan Project in World War II.
Owing to the expense of refining pure fluorine, most commercial applications use fluorine compounds, with about half of mined fluorite used in steelmaking. The rest of the fluorite is converted into hydrogen fluoride en route to various organic fluorides, or into cryolite, which plays a key role in aluminium refining. The carbon–fluorine bond is usually very stable. Organofluorine compounds are widely used as refrigerants, electrical insulation, and PTFE (Teflon). Pharmaceuticals such as atorvastatin and fluoxetine contain C−F bonds. The fluoride ion from dissolved fluoride salts inhibits dental cavities and so finds use in toothpaste and water fluoridation. Global fluorochemical sales amount to more than US$15billion a year.
Fluorine atoms have nine electrons, one fewer than neon, and electron configuration 1s22s22p5: two electrons in a filled inner shell and seven in an outer shell requiring one more to be filled. The outer electrons are ineffective at nuclear shielding, and experience a high effective nuclear charge of 9− 2= 7; this affects the atom's physical properties.[3]
The bond energy of difluorine is much lower than that of either Cl 2 or Br 2 and similar to the easily cleaved peroxide bond; this, along with high electronegativity, accounts for fluorine's easy dissociation, high reactivity, and strong bonds to non-fluorine atoms.[21][22] Conversely, bonds to other atoms are very strong because of fluorine's high electronegativity. Unreactive substances like powdered steel, glass fragments, and asbestos fibers react quickly with cold fluorine gas; wood and water spontaneously combust under a fluorine jet.[5][23]
Reactions of elemental fluorine with metals require varying conditions. Alkali metals cause explosions and alkaline earth metals display vigorous activity in bulk; to prevent passivation from the formation of metal fluoride layers, most other metals such as aluminium and iron must be powdered,[21] and noble metals require pure fluorine gas at 300–450°C (572–842°F).[24] Some solid nonmetals (sulfur, phosphorus) react vigorously in liquid fluorine.[25]Hydrogen sulfide[25] and sulfur dioxide[26] combine readily with fluorine, the latter sometimes explosively; sulfuric acid exhibits much less activity, requiring elevated temperatures.[27]
Hydrogen, like some of the alkali metals, reacts explosively with fluorine.[28]Carbon, as lamp black, reacts at room temperature to yield tetrafluoromethane. Graphite combines with fluorine above 400°C (752°F) to produce non-stoichiometriccarbon monofluoride; higher temperatures generate gaseous fluorocarbons, sometimes with explosions.[29] Carbon dioxide and carbon monoxide react at or just above room temperature,[30] whereas paraffins and other organic chemicals generate strong reactions:[31] even completely substituted haloalkanes such as carbon tetrachloride, normally incombustible, may explode.[32] Although nitrogen trifluoride is stable, nitrogen requires an electric discharge at elevated temperatures for reaction with fluorine to occur, due to the very strong triple bond in elemental nitrogen;[33] ammonia may react explosively.[34][35]Oxygen does not combine with fluorine under ambient conditions, but can be made to react using electric discharge at low temperatures and pressures; the products tend to disintegrate into their constituent elements when heated.[36][37][38] Heavier halogens[39] react readily with fluorine as does the noble gas radon;[40] of the other noble gases, only xenon and krypton react, and only under special conditions.[41]Argon does not react with fluorine gas; however, it does form a compound with fluorine, argon fluorohydride.
Phases
Crystal structure of β-fluorine. Spheres indicate F 2 molecules that may assume any angle. Other molecules are constrained to planes.
Animation showing the crystal structure of beta-fluorine. Molecules on the faces of the unit cell have rotations constrained to a plane.
At room temperature, fluorine is a gas of diatomic molecules,[5] pale yellow when pure (sometimes described as yellow-green).[42] It has a characteristic halogen-like pungent and biting odor detectable at 20ppb.[43] Fluorine condenses into a bright yellow liquid at −188°C (−306.4°F), a transition temperature similar to those of oxygen and nitrogen.[44]
Fluorine has two solid forms, α- and β-fluorine. The latter crystallizes at −220°C (−364.0°F) and is transparent and soft, with the same disordered cubic structure of freshly crystallized solid oxygen,[44][note 3] unlike the orthorhombic systems of other solid halogens.[46][47] Further cooling to −228°C (−378.4°F) induces a phase transition into opaque and hard α-fluorine, which has a monoclinic structure with dense, angled layers of molecules. The transition from β- to α-fluorine is more exothermic than the condensation of fluorine, and can be violent.[46][47]
Among the lighter elements, fluorine's abundance value of 400ppb (parts per billion) – 24th among elements in the universe – is exceptionally low: other elements from carbon to magnesium are twenty or more times as common.[58] This is because stellar nucleosynthesis processes bypass fluorine, and any fluorine atoms otherwise created have high nuclear cross sections, allowing collisions with hydrogen or helium to generate oxygen or neon respectively.[58][59]
Beyond this transient existence, three explanations have been proposed for the presence of fluorine:[58][60]
Fluorine is the thirteenth most common element in Earth's crust at 600–700ppm (parts per million) by mass.[61] Though believed not to occur naturally, elemental fluorine has been shown to be present as an occlusion in antozonite, a variant of fluorite.[62] Most fluorine exists as fluoride-containing minerals. Fluorite, fluorapatite and cryolite are the most industrially significant.[61][63] Fluorite (CaF 2), also known as fluorspar, abundant worldwide, is the main source of fluoride, and hence fluorine. China and Mexico are the major suppliers.[63][64][65][66][67] Fluorapatite (Ca5(PO4)3F), which contains most of the world's fluoride, is an inadvertent source of fluoride as a byproduct of fertilizer production.[63] Cryolite (Na 3AlF 6), used in the production of aluminium, is the most fluorine-rich mineral. Economically viable natural sources of cryolite have been exhausted, and most is now synthesised commercially.[63]
Fluorite: Pink globular mass with crystal facets
Fluorapatite: Long, prismatic crystal, dull in lustre, protruding, at an angle, from matrix of aggregate-like rock
Cryolite: A parallelogram-shaped outline with diatomic molecules arranged in two layers
Other minerals such as topaz contain fluorine. Fluorides, unlike other halides, are insoluble and do not occur in commercially favorable concentrations in saline waters.[63] Trace quantities of organofluorines of uncertain origin have been detected in volcanic eruptions and geothermal springs.[68] The existence of gaseous fluorine in crystals, suggested by the smell of crushed antozonite, is contentious;[69][62] a 2012 study reported the presence of 0.04% F 2 by weight in antozonite, attributing these inclusions to radiation from the presence of tiny amounts of uranium.[62]
In 1529, Georgius Agricola described fluorite as an additive used to lower the melting point of metals during smelting.[70][71][note 5] He penned the Latin word fluorēs (fluor, flow) for fluorite rocks. The name later evolved into fluorspar (still commonly used) and then fluorite.[64][75][76] The composition of fluorite was later determined to be calcium difluoride.[77]
Hydrofluoric acid was used in glass etching from 1720 onward.[note 6]Andreas Sigismund Marggraf first characterized it in 1764 when he heated fluorite with sulfuric acid, and the resulting solution corroded its glass container.[79][80] Swedish chemist Carl Wilhelm Scheele repeated the experiment in 1771, and named the acidic product fluss-spats-syran (fluorspar acid).[80][81] In 1810, the French physicist André-Marie Ampère suggested that hydrogen and an element analogous to chlorine constituted hydrofluoric acid.[82] He also proposed in a letter to Sir Humphry Davy dated August 26, 1812 that this then-unknown substance may be named fluorine from fluoric acid and the -ine suffix of other halogens.[83][84] This word, often with modifications, is used in most European languages; however, Greek, Russian, and some others, following Ampère's later suggestion, use the name ftor or derivatives, from the Greek φθόριος (phthorios, destructive).[85] The New Latin name fluorum gave the element its current symbol F; Fl was used in early papers.[86][note 7]
Isolation
1887 drawing of Moissan's apparatus
Initial studies on fluorine were so dangerous that several 19th-century experimenters were deemed "fluorine martyrs" after misfortunes with hydrofluoric acid.[note 8] Isolation of elemental fluorine was hindered by the extreme corrosiveness of both elemental fluorine itself and hydrogen fluoride, as well as the lack of a simple and suitable electrolyte.[77][87]Edmond Frémy postulated that electrolysis of pure hydrogen fluoride to generate fluorine was feasible and devised a method to produce anhydrous samples from acidified potassium bifluoride; instead, he discovered that the resulting (dry) hydrogen fluoride did not conduct electricity.[77][87][88] Frémy's former student Henri Moissan persevered, and after much trial and error found that a mixture of potassium bifluoride and dry hydrogen fluoride was a conductor, enabling electrolysis. To prevent rapid corrosion of the platinum in his electrochemical cells, he cooled the reaction to extremely low temperatures in a special bath and forged cells from a more resistant mixture of platinum and iridium, and used fluorite stoppers.[87][89] In 1886, after 74 years of effort by many chemists, Moissan isolated elemental fluorine.[88][90]
[I]n recognition of the great services rendered by him in his investigation and isolation of the element fluorine... The whole world has admired the great experimental skill with which you have studied that savage beast among the elements.[note 9]
The Frigidaire division of General Motors (GM) experimented with chlorofluorocarbon refrigerants in the late 1920s, and Kinetic Chemicals was formed as a joint venture between GM and DuPont in 1930 hoping to market Freon-12 (CCl 2F 2) as one such refrigerant. It replaced earlier and more toxic compounds, increased demand for kitchen refrigerators, and became profitable; by 1949 DuPont had bought out Kinetic and marketed several other Freon compounds.[80][92][93][94]Polytetrafluoroethylene (Teflon) was serendipitously discovered in 1938 by Roy J. Plunkett while working on refrigerants at Kinetic, and its superlative chemical and thermal resistance lent it to accelerated commercialization and mass production by 1941.[80][92][93]
Large-scale production of elemental fluorine began during World War II. Germany used high-temperature electrolysis to make tons of the planned incendiary chlorine trifluoride[95] and the Manhattan Project used huge quantities to produce uranium hexafluoride for uranium enrichment. Since UF 6 is as corrosive as fluorine, gaseous diffusion plants required special materials: nickel for membranes, fluoropolymers for seals, and liquid fluorocarbons as coolants and lubricants. This burgeoning nuclear industry later drove post-war fluorochemical development.[96]
Fluorine has a rich chemistry, encompassing organic and inorganic domains. It combines with metals, nonmetals, metalloids, and most noble gases.[97] Fluorine's high electron affinity results in a preference for ionic bonding; when it forms covalent bonds, these are polar, and almost always single.[98][99][note 10]
Oxidation states
In compounds, fluorine almost exclusively assumes an oxidation state of −1. Fluorine in F 2 is defined to have oxidation state 0. The unstable species F− 2 and F− 3, which decompose at around 40K, have intermediate oxidation states;[100]F+ 4 and a few related species are predicted to be stable.[101]
Boiling points of hydrogen halides and chalcogenides, showing the unusually high values for hydrogen fluoride and water
Hydrogen and fluorine combine to yield hydrogen fluoride, in which discrete molecules form clusters by hydrogen bonding, resembling water more than hydrogen chloride.[126][127][128] It boils at a much higher temperature than heavier hydrogen halides and unlike them is miscible with water.[129] Hydrogen fluoride readily hydrates on contact with water to form aqueous hydrogen fluoride, also known as hydrofluoric acid. Unlike the other hydrohalic acids, which are strong, hydrofluoric acid is a weak acid at low concentrations.[130][131] However, it can attack glass, something the other acids cannot do.[132]
Other reactive nonmetals
Chlorine trifluoride, whose corrosive potential ignites asbestos, concrete, sand and other fire retardants
Binary fluorides of metalloids and p-block nonmetals are generally covalent and volatile, with varying reactivities. Period 3 and heavier nonmetals can form hypervalent fluorides.[134]
Boron trifluoride is planar and possesses an incomplete octet. It functions as a Lewis acid and combines with Lewis bases like ammonia to form adducts.[135]Carbon tetrafluoride is tetrahedral and inert;[note 13]its group analogues, silicon and germanium tetrafluoride, are also tetrahedral[136] but behave as Lewis acids.[137][138] The pnictogens form trifluorides that increase in reactivity and basicity with higher molecular weight, although nitrogen trifluoride resists hydrolysis and is not basic.[139] The pentafluorides of phosphorus, arsenic, and antimony are more reactive than their respective trifluorides, with antimony pentafluoride the strongest neutral Lewis acid known, only behind gold pentafluoride.[118][140][141]
Chalcogens have diverse fluorides: unstable difluorides have been reported for oxygen (the only known compound with oxygen in an oxidation state of +2), sulfur, and selenium; tetrafluorides and hexafluorides exist for sulfur, selenium, and tellurium. The latter are stabilized by more fluorine atoms and lighter central atoms, so sulfur hexafluoride is especially inert.[142][143] Chlorine, bromine, and iodine can each form mono-, tri-, and pentafluorides, but only iodine heptafluoride has been characterized among possible interhalogen heptafluorides.[144] Many of them are powerful sources of fluorine atoms, and industrial applications using chlorine trifluoride require precautions similar to those using fluorine.[145][146]
These xenon tetrafluoride crystals were photographed in 1962. The compound's synthesis, as with xenon hexafluoroplatinate, surprised many chemists.
Noble gases, having complete electron shells, defied reaction with other elements until 1962 when Neil Bartlett reported synthesis of xenon hexafluoroplatinate;[148]xenon difluoride, tetrafluoride, hexafluoride, and multiple oxyfluorides have been isolated since then.[149] Among other noble gases, krypton forms a difluoride,[150] and radon and fluorine generate a solid suspected to be radon difluoride.[151][152] Binary fluorides of lighter noble gases are exceptionally unstable: argon and hydrogen fluoride combine under extreme conditions to give argon fluorohydride.[41] Helium has no long-lived fluorides,[153] and no neon fluoride has ever been observed;[154] helium fluorohydride has been detected for milliseconds at high pressures and low temperatures.[153]
Organic compounds
Immiscible layers of colored water (top) and much denser perfluoroheptane (bottom) in a beaker; a goldfish and crab cannot penetrate the boundary; quarters rest at the bottom.
Chemical structure of Nafion, a fluoropolymer used in fuel cells and many other applications
The carbon–fluorine bond is organic chemistry's strongest,[156] and gives stability to organofluorines.[157] It is almost non-existent in nature, but is used in artificial compounds. Research in this area is usually driven by commercial applications;[158] the compounds involved are diverse and reflect the complexity inherent in organic chemistry.[92]
The substitution of hydrogen atoms in an alkane by progressively more fluorine atoms gradually alters several properties: melting and boiling points are lowered, density increases, solubility in hydrocarbons decreases and overall stability increases. Perfluorocarbons,[note 14] in which all hydrogen atoms are substituted, are insoluble in most organic solvents, reacting at ambient conditions only with sodium in liquid ammonia.[159]
The term perfluorinated compound is used for what would otherwise be a perfluorocarbon if not for the presence of a functional group,[160][note 15] often a carboxylic acid. These compounds share many properties with perfluorocarbons such as stability and hydrophobicity,[162] while the functional group augments their reactivity, enabling them to adhere to surfaces or act as surfactants.[163]Fluorosurfactants, in particular, can lower the surface tension of water more than their hydrocarbon-based analogues. Fluorotelomers, which have some unfluorinated carbon atoms near the functional group, are also regarded as perfluorinated.[162]
Polymers
Polymers exhibit the same stability increases afforded by fluorine substitution (for hydrogen) in discrete molecules; their melting points generally increase too.[164]Polytetrafluoroethylene (PTFE), the simplest fluoropolymer and perfluoro analogue of polyethylene with structural unit –CF 2–, demonstrates this change as expected, but its very high melting point makes it difficult to mold.[165] Various PTFE derivatives are less temperature-tolerant but easier to mold: fluorinated ethylene propylene replaces some fluorine atoms with trifluoromethyl groups, perfluoroalkoxy alkanes do the same with trifluoromethoxy groups,[165] and Nafion contains perfluoroether side chains capped with sulfonic acid groups.[166][167] Other fluoropolymers retain some hydrogen atoms; polyvinylidene fluoride has half the fluorine atoms of PTFE and polyvinyl fluoride has a quarter, but both behave much like perfluorinated polymers.[168]
The gaseous HF can then be absorbed in water or liquefied.[170]
About 20% of manufactured HF is a byproduct of fertilizer production, which produces hexafluorosilicic acid (H2SiF6), which can be degraded to release HF thermally and by hydrolysis:
Moissan's method is used to produce industrial quantities of fluorine, via the electrolysis of a potassium bifluoride/hydrogen fluoride mixture: hydrogen ions are reduced at a steel container cathode and fluoride ions are oxidized at a carbon block anode, under 8–12volts, to generate hydrogen and fluorine gas respectively.[65][171] Temperatures are elevated, KF•2HF melting at 70°C (158°F) and being electrolyzed at 70–130°C (158–266°F). KF, which acts to provide electrical conductivity, is essential since pure HF cannot be electrolyzed because it is virtually non-conductive.[80][172][173] Fluorine can be stored in steel cylinders that have passivated interiors, at temperatures below 200°C (392°F); otherwise nickel can be used.[80][174] Regulator valves and pipework are made of nickel, the latter possibly using Monel instead.[175] Frequent passivation, along with the strict exclusion of water and greases, must be undertaken. In the laboratory, glassware may carry fluorine gas under low pressure and anhydrous conditions;[175] some sources instead recommend nickel-Monel-PTFE systems.[176]
Laboratory routes
While preparing for a 1986 conference to celebrate the centennial of Moissan's achievement, Karl O. Christe reasoned that chemical fluorine generation should be feasible since some metal fluoride anions have no stable neutral counterparts; their acidification potentially triggers oxidation instead. He devised a method which evolves fluorine at high yield and atmospheric pressure:[177]
Christe later commented that the reactants "had been known for more than 100 years and even Moissan could have come up with this scheme."[178] As late as 2008, some references still asserted that fluorine was too reactive for any chemical isolation.[179]
Fluorite mining, which supplies most global fluorine, peaked in 1989 when 5.6million metric tons of ore were extracted. Chlorofluorocarbon restrictions lowered this to 3.6million tons in 1994; production has since been increasing. Around 4.5million tons of ore and revenue of US$550million were generated in 2003; later reports estimated 2011 global fluorochemical sales at $15billion and predicted 2016–18 production figures of 3.5 to 5.9million tons, and revenue of at least $20billion.[80][180][181][182][183]Froth flotation separates mined fluorite into two main metallurgical grades of equal proportion: 60–85% pure metspar is almost all used in iron smelting whereas 97%+ pure acidspar is mainly converted to the key industrial intermediate hydrogen fluoride.[65][80][184]
Clickable diagram of the fluorochemical industry according to mass flowsSF 6 current transformers at a Russian railway
At least 17,000metrictons 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 because of handling challenges. Most processes using free fluorine in large amounts employ in situ generation under vertical integration.[185]
The largest application of fluorine gas, consuming up to 7,000metrictons annually, is in the preparation of UF 6 for the nuclear fuel cycle. Fluorine is used to fluorinate uranium tetrafluoride, itself formed from uranium dioxide and hydrofluoric acid.[185] Fluorine is monoisotopic, so any mass differences between UF 6 molecules are due to the presence of 235 U or 238 U, enabling uranium enrichment via gaseous diffusion or gas centrifuge.[5][65] About 6,000metrictons per year go into producing the inert dielectricSF 6 for high-voltage transformers and circuit breakers, eliminating the need for hazardous polychlorinated biphenyls associated with oil-filled devices.[186] Several fluorine compounds are used in electronics: rhenium and tungsten hexafluoride in chemical vapor deposition, tetrafluoromethane in plasma etching[187][188][189] and nitrogen trifluoride in cleaning equipment.[65] Fluorine is also used in the synthesis of organic fluorides, but its reactivity often necessitates conversion first to the gentler ClF 3, BrF 3, or IF 5, which together allow calibrated fluorination. Fluorinated pharmaceuticals use sulfur tetrafluoride instead.[65]
Inorganic fluorides
Aluminium extraction depends critically on cryolite
As with other iron alloys, around 3kg (6.6lb) metspar is added to each metric ton of steel; the fluoride ions lower its melting point and viscosity.[65][190] Alongside its role as an additive in materials like enamels and welding rod coats, most acidspar is reacted with sulfuric acid to form hydrofluoric acid, which is used in steel pickling, glass etching and alkane cracking.[65] One-third of HF goes into synthesizing cryolite and aluminium trifluoride, both fluxes in the Hall–Héroult process for aluminium extraction; replenishment is necessitated by their occasional reactions with the smelting apparatus. Each metric ton of aluminium requires about 23kg (51lb) of flux.[65][191] Fluorosilicates consume the second largest portion, with sodium fluorosilicate used in water fluoridation and laundry effluent treatment, and as an intermediate en route to cryolite and silicon tetrafluoride.[192] Other important inorganic fluorides include those of cobalt, nickel, and ammonium.[65][103][193]
Organic fluorides
Organofluorides consume over 20% of mined fluorite and over 40% of hydrofluoric acid, with refrigerant gases dominating and fluoropolymers increasing their market share.[65][194]Surfactants are a minor application but generate over $1billion in annual revenue.[195] Due to the danger from direct hydrocarbon–fluorine reactions above −150°C (−238°F), industrial fluorocarbon production is indirect, mostly through halogen exchange reactions such as Swarts fluorination, in which chlorocarbon chlorines are substituted for fluorines by hydrogen fluoride under catalysts. Electrochemical fluorination subjects hydrocarbons to electrolysis in hydrogen fluoride, and the Fowler process treats them with solid fluorine carriers like cobalt trifluoride.[92][196]
Halogenated refrigerants, termed Freons in informal contexts,[note 16] are identified by R-numbers that denote the amount of fluorine, chlorine, carbon, and hydrogen present.[65][197]Chlorofluorocarbons (CFCs) like R-11, R-12, and R-114 once dominated organofluorines, peaking in production in the 1980s. Used for air conditioning systems, propellants and solvents, their production was below one-tenth of this peak by the early 2000s, after widespread international prohibition.[65] Hydrochlorofluorocarbons (HCFCs) and hydrofluorocarbons (HFCs) were designed as replacements; their synthesis consumes more than 90% of the fluorine in the organic industry. Important HCFCs include R-22, chlorodifluoromethane, and R-141b. The main HFC is R-134a[65] with a new type of molecule HFO-1234yf, a Hydrofluoroolefin (HFO) coming to prominence owing to its global warming potential of less than 1% that of HFC-134a.[198]
Polymers
Fluorosurfactant-treated fabrics are often hydrophobic.
About 180,000metrictons of fluoropolymers were produced in 2006 and 2007, generating over $3.5billion revenue per year.[199] The global market was estimated at just under $6billion in 2011.[200] Fluoropolymers can only be formed by polymerizing free radicals.[164]
Polytetrafluoroethylene (PTFE), sometimes called by its DuPont name Teflon,[201] represents 60–80% by mass of the world's fluoropolymer production.[199] The largest application is in electrical insulation since PTFE is an excellent dielectric. It is also used in the chemical industry where corrosion resistance is needed, in coating pipes, tubing, and gaskets. Another major use is in PFTE-coated fiberglass cloth for stadium roofs. The major consumer application is for non-stick cookware.[201] Jerked PTFE film becomes expanded PTFE (ePTFE), a fine-pored membrane sometimes referred to by the brand name Gore-Tex and used for rainwear, protective apparel, and filters; ePTFE fibers may be made into seals and dust filters.[201] Other fluoropolymers, including fluorinated ethylene propylene, mimic PTFE's properties and can substitute for it; they are more moldable, but also more costly and have lower thermal stability. Films from two different fluoropolymers replace glass in solar cells.[201][202]
The chemically resistant (but expensive) fluorinated ionomers are used as electrochemical cell membranes, of which the first and most prominent example is Nafion. Developed in the 1960s, it was initially deployed as fuel cell material in spacecraft and then replaced mercury-based chloralkali process cells. Recently, the fuel cell application has reemerged with efforts to install proton exchange membrane fuel cells into automobiles.[203][204][205]Fluoroelastomers such as Viton are crosslinked fluoropolymer mixtures mainly used in O-rings;[201]perfluorobutane (C4F10) is used as a fire-extinguishing agent.[206]
Fluorosurfactants are small organofluorine molecules used for repelling water and stains. Although expensive (comparable to pharmaceuticals at $200–2000 per kilogram), they yielded over $1billion in annual revenues by 2006; Scotchgard alone generated over $300million in 2000.[195][207][208] Fluorosurfactants are a minority in the overall surfactant market, most of which is taken up by much cheaper hydrocarbon-based products. Applications in paints are burdened by compounding costs; this use was valued at only $100million in 2006.[195]
Agrichemicals
About 30% of agrichemicals contain fluorine,[209] most of them herbicides and fungicides with a few crop regulators. Fluorine substitution, usually of a single atom or at most a trifluoromethyl group, is a robust modification with effects analogous to fluorinated pharmaceuticals: increased biological stay time, membrane crossing, and altering of molecular recognition.[210]Trifluralin is a prominent example, with large-scale use in the U.S. as a weedkiller,[210][211] but it is a suspected carcinogen and has been banned in many European countries.[212]Sodium monofluoroacetate (1080) is a mammalian poison in which one sodium acetate hydrogen is replaced with fluorine; 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 in the early 20th century, and was later deployed in its current use. New Zealand, the largest consumer of 1080, uses it to protect kiwis from the invasive Australian common brushtail possum.[213] Europe and the U.S. have banned 1080.[214][215][note 17]
Population studies from the mid-20th century onwards show topical fluoride reduces dental caries. This was first attributed to the conversion of tooth enamel hydroxyapatite into the more durable fluorapatite, but studies on pre-fluoridated teeth refuted this hypothesis, and current theories involve fluoride aiding enamel growth in small caries.[216] After studies of children in areas where fluoride was naturally present in drinking water, controlled public water supply fluoridation to fight tooth decay[217] began in the 1940s and is now applied to water supplying 6 percent of the global population, including two-thirds of Americans.[218][219] Reviews of the scholarly literature in 2000 and 2007 associated water fluoridation with a significant reduction of tooth decay in children.[220] Despite such endorsements and evidence of no adverse effects other than mostly benign dental fluorosis,[221]opposition still exists on ethical and safety grounds.[219][222] The benefits of fluoridation have lessened, possibly due to other fluoride sources, but are still measurable in low-income groups.[223]Sodium monofluorophosphate and sometimes sodium or tin(II) fluoride are often found in fluoride toothpastes, first introduced in the U.S. in 1955 and now ubiquitous in developed countries, alongside fluoridated mouthwashes, gels, foams, and varnishes.[223][224]
Twenty percent of modern pharmaceuticals contain fluorine.[225] One of these, the cholesterol-reducer atorvastatin (Lipitor), made more revenue than any other drug until it became generic in 2011.[226] The combination asthma prescription Seretide, a top-ten revenue drug in the mid-2000s, contains two active ingredients, one of which – fluticasone – is fluorinated.[227] Many drugs are fluorinated to delay inactivation and lengthen dosage periods because the carbon–fluorine bond is very stable.[228] Fluorination also increases lipophilicity because the bond is more hydrophobic than the carbon–hydrogen bond, and this often helps in cell membrane penetration and hence bioavailability.[227]
A full-body F PET scan with glucose tagged with radioactive fluorine-18. The normal brain and kidneys take up enough glucose to be imaged. A malignant tumor is seen in the upper abdomen. Radioactive fluorine is seen in urine in the bladder.
Fluorine-18 is often found in radioactive tracers for positron emission tomography, as its half-life of almost two hours is long enough to allow for its transport from production facilities to imaging centers.[239] The most common tracer is fluorodeoxyglucose[239] which, after intravenous injection, is taken up by glucose-requiring tissues such as the brain and most malignant tumors;[240]computer-assisted tomography can then be used for detailed imaging.[241]
Liquid fluorocarbons can hold large volumes of oxygen or carbon dioxide, more so than blood, and have attracted attention for their possible uses in artificial blood and in liquid breathing.[242] Because fluorocarbons do not normally mix with water, they must be mixed into emulsions (small droplets of perfluorocarbon suspended in water) to be used as blood.[243][244] One such product, Oxycyte, has been through initial clinical trials.[245] These substances can aid endurance athletes and are banned from sports; one cyclist's near death in 1998 prompted an investigation into their abuse.[246][247] Applications of pure perfluorocarbon liquid breathing (which uses pure perfluorocarbon liquid, not a water emulsion) include assisting burn victims and premature babies with deficient lungs. Partial and complete lung filling have been considered, though only the former has had any significant tests in humans.[248] An Alliance Pharmaceuticals effort reached clinical trials but was abandoned because the results were not better than normal therapies.[249]
Fluorine is not essential for humans and other mammals, but small amounts are known to be beneficial for the strengthening of dental enamel (where the formation of fluorapatite makes the enamel more resistant to attack, from acids produced by bacterial fermentation of sugars). Small amounts of fluorine may be beneficial for bone strength, but the latter has not been definitively established.[250] Both the WHO and the Institute of Medicine of the US National Academies publish recommended daily allowance (RDA) and upper tolerated intake of fluorine, which varies with age and gender.[251][252]
Elemental fluorine is highly toxic to living organisms. Its effects in humans start at concentrations lower than hydrogen cyanide's 50ppm[256] and are similar to those of chlorine:[257] significant irritation of the eyes and respiratory system as well as liver and kidney damage occur above 25ppm, which is the immediately dangerous to life and health value for fluorine.[258] The eyes and nose are seriously damaged at 100ppm,[258] and inhalation of 1,000ppm fluorine will cause death in minutes,[259] compared to 270ppm for hydrogen cyanide.[260]
Hydrofluoric acid is the weakest of the hydrohalic acids, having a pKa of 3.2 at 25°C.[263] Pure hydrogen fluoride is a volatile liquid due to the presence of hydrogen bonding, while the other hydrogen halides are gases. It is able to attack glass, concrete, metals, and organic matter.[264]
Hydrofluoric acid is a contact poison with greater hazards than many strong acids like sulfuric acid even though it is weak: it remains neutral in aqueous solution and thus penetrates tissue faster, whether through inhalation, ingestion or the skin, and at least nine U.S. workers died in such accidents from 1984 to 1994. It reacts with calcium and magnesium in the blood leading to hypocalcemia and possible death through cardiac arrhythmia.[265] Insoluble calcium fluoride formation triggers strong pain[266] and burns larger than 160cm2 (25in2) can cause serious systemic toxicity.[267]
Exposure may not be evident for eight hours for 50% HF, rising to 24 hours for lower concentrations, and a burn may initially be painless as hydrogen fluoride affects nerve function. If skin has been exposed to HF, damage can be reduced by rinsing it under a jet of water for 10–15 minutes and removing contaminated clothing.[268]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 rinsing solutions.[269][270][271] Hydrofluoric acid absorption requires further medical treatment; calcium gluconate may be injected or administered intravenously. Using calcium chloride – a common laboratory reagent – in lieu of calcium gluconate is contraindicated, and may lead to severe complications. Excision or amputation of affected parts may be required.[267][272]
Soluble fluorides are moderately toxic: 5–10g sodium fluoride, or 32–64mg fluoride ions per kilogram of body mass, represents a lethal dose for adults.[273] One-fifth of the lethal dose can cause adverse health effects,[274] and chronic excess consumption may lead to skeletal fluorosis, which affects millions in Asia and Africa, and, in children, to reduced intelligence.[274][275] Ingested fluoride forms hydrofluoric acid in the stomach which is easily absorbed by the intestines, where it crosses cell membranes, binds with calcium and interferes with various enzymes, before urinary excretion. Exposure limits are determined by urine testing of the body's ability to clear fluoride ions.[274][276]
Historically, most cases of fluoride poisoning have been caused by accidental ingestion of insecticides containing inorganic fluorides.[277] Most current calls to poison control centers for possible fluoride poisoning come from the ingestion of fluoride-containing toothpaste.[274] Malfunctioning water fluoridation equipment is another cause: one incident in Alaska affected almost 300 people and killed one person.[278] Dangers from toothpaste are aggravated for small children, and the Centers for Disease Control and Prevention recommends supervising children below six brushing their teeth so that they do not swallow toothpaste.[279] One regional study examined a year of pre-teen fluoride poisoning reports totaling 87 cases, including one death from ingesting insecticide. Most had no symptoms, but about 30% had stomach pains.[277] A larger study across the U.S. had similar findings: 80% of cases involved children under six, and there were few serious cases.[280]
Environmental concerns
Atmosphere
NASA projection of stratospheric ozone over North America without the Montreal Protocol
The Montreal Protocol, signed in 1987, set strict regulations on chlorofluorocarbons (CFCs) and bromofluorocarbons due to their ozone damaging potential (ODP). The high stability which suited them to their original applications also meant that they were not decomposing until they reached higher altitudes, where liberated chlorine and bromine atoms attacked ozone molecules.[282] Even with the ban, and early indications of its efficacy, predictions warned that several generations would pass before full recovery.[283][284] With one-tenth the ODP of CFCs, hydrochlorofluorocarbons (HCFCs) are the current replacements,[285] and are themselves scheduled for substitution by 2030–2040 by hydrofluorocarbons (HFCs) with no chlorine and zero ODP.[286] In 2007 this date was brought forward to 2020 for developed countries;[287] the Environmental Protection Agency had already prohibited one HCFC's production and capped those of two others in 2003.[286] Fluorocarbon gases are generally greenhouse gases with global-warming potentials (GWPs) of about 100 to 10,000; sulfur hexafluoride has a value of around 20,000.[288] An outlier is HFO-1234yf which is a new type of refrigerant called a Hydrofluoroolefin (HFO) and has attracted global demand due to its GWP of less than 1 compared to 1,430 for the current refrigerant standard HFC-134a.[198]
Organofluorines exhibit biopersistence due to the strength of the carbon–fluorine bond. Perfluoroalkyl acids (PFAAs), which are sparingly water-soluble owing to their acidic functional groups, are noted persistent organic pollutants;[290]perfluorooctanesulfonic acid (PFOS) and perfluorooctanoic acid (PFOA) are most often researched.[291][292][293] PFAAs have been found in trace quantities worldwide from polar bears to humans, with PFOS and PFOA known to reside in breast milk and the blood of newborn babies. A 2013 review showed a slight correlation between groundwater and soil PFAA levels and human activity; there was no clear pattern of one chemical dominating, and higher amounts of PFOS were correlated to higher amounts of PFOA.[291][292][294] In the body, PFAAs bind to proteins such as serum albumin; they tend to concentrate within humans in the liver and blood before excretion through the kidneys. Dwell time in the body varies greatly by species, with half-lives of days in rodents, and years in humans.[291][292][295] High doses of PFOS and PFOA cause cancer and death in newborn rodents but human studies have not established an effect at current exposure levels.[291][292][295]
↑ Assuming that hydrogen is not considered a halogen.
↑ Sources disagree on the radii of oxygen, fluorine, and neon atoms. Precise comparison is thus impossible.
↑ α-Fluorine has a regular pattern of molecules and is a crystalline solid, but its molecules do not have a specific orientation. β-Fluorine's molecules have fixed locations and minimal rotational uncertainty.[45]
↑ The ratio of the angular momentum to magnetic moment is called the gyromagnetic ratio. "Certain nuclei can for many purposes be thought of as spinning round an axis like the Earth or like a top. In general the spin endows them with angular momentum and with a magnetic moment; the first because of their mass, the second because all or part of their electric charge may be rotating with the mass."[49]
↑ Basilius Valentinus supposedly described fluorite in the late 15th century, but because his writings were uncovered 200 years later, this work's veracity is doubtful.[72][73][74]
↑ Or perhaps from as early as 1670 onwards; Partington[78] and Weeks[77] give differing accounts.
↑ ZrF 4 melts at 932°C (1,710°F),[112]HfF 4 sublimes at 968°C (1,774°F),[109] and UF 4 melts at 1,036°C (1,897°F).[113]
↑ These thirteen are those of molybdenum, technetium, ruthenium, rhodium, tungsten, rhenium, osmium, iridium, platinum, polonium, uranium, neptunium, and plutonium.
↑ Carbon tetrafluoride is formally organic, but is included here rather than in the organofluorine chemistry section – where more complex carbon-fluorine compounds are discussed – for comparison with SiF 4 and GeF 4.
↑ Perfluorocarbon and fluorocarbon are IUPAC synonyms for molecules containing carbon and fluorine only, but in colloquial and commercial contexts the latter term may refer to any carbon- and fluorine-containing molecule, possibly with other elements.
↑ This terminology is imprecise, and perfluorinated substance is also used.[161]
↑ This DuPont trademark is sometimes further misused for CFCs, HFCs, or HCFCs.
↑ American sheep and cattle collars may use 1080 against predators like coyotes.
Alavi, Abbas; Huang, Steve S. (2007). "Positron Emission Tomography in Medicine: An Overview". In Hayat, M. A. (ed.). Cancer Imaging, Volume 1: Lung and Breast Carcinomas. Burlington: Academic Press. pp.39–44. ISBN978-0-12-370468-9.
Arana, L. R.; Mas, N.; Schmidt, R.; Franz, A. J.; Schmidt, M. A.; Jensen, K. F. (2007). "Isotropic Etching of Silicon in Fluorine Gas for MEMS Micromachining". Journal of Micromechanics and Microengineering. 17 (2): 384–392. Bibcode:2007JMiMi..17..384A. doi:10.1088/0960-1317/17/2/026. S2CID135708022.
Babel, Dietrich; Tressaud, Alain (1985). "Crystal Chemistry of Fluorides". In Hagenmuller, Paul (ed.). Inorganic Solid Fluorides: Chemistry And Physics. Orlando: Academic Press. pp.78–203. ISBN978-0-12-412490-5.
Baelum, Vibeke; Sheiham, Aubrey; Burt, Brian (2008). "Caries Control for Populations". In Fejerskov, Ole; Kidd, Edwina (eds.). Dental Caries: The Disease and Its Clinical Management (2nded.). Oxford: Blackwell Munksgaard. pp.505–526. ISBN978-1-4051-3889-5.
Barbee, K.; McCormack, K.; Vartanian, V. (2000). "EHS Concerns with Ozonated Water Spray Processing". In Mendicino, L. (ed.). Environmental Issues in the Electronics and Semiconductor Industries. Pennington, NJ: The Electrochemical Society. pp.108–121. ISBN978-1-56677-230-3.
Barry, Patrick L.; Phillips, Tony (26 May 2006). "Good News and a Puzzle". National Aeronautics and Space Administration. Archived from the original on 27 May 2010. Retrieved 6 January 2012.
Becker, S.; Müller, B. G. (1990). "Vanadium Tetrafluoride". Angewandte Chemie International Edition in English. 29 (4): 406–407. doi:10.1002/anie.199004061.
Bégué, Jean-Pierre; Bonnet-Delpon, Danièle (2008). Bioorganic and Medicinal Chemistry of Fluorine. Hoboken: John Wiley & Sons. ISBN978-0-470-27830-7.
Bihary, Z.; Chaban, G. M.; Gerber, R. B. (2002). "Stability of a Chemically Bound Helium Compound in High-pressure Solid Helium". The Journal of Chemical Physics. 117 (11): 5105–5108. Bibcode:2002JChPh.117.5105B. doi:10.1063/1.1506150.
Brantley, L. R. (1949). Squires, Roy; Clarke, Arthur C. (eds.). "Fluorine". Pacific Rockets: Journal of the Pacific Rocket Society. 3 (1). South Pasadena: Sawyer Publishing/Pacific Rocket Society Historical Library: 11–18. ISBN978-0-9794418-5-1.
Brown, Paul L.; Mompean, Federico J.; Perrone, Jane; Illemassène, Myriam (2005). Chemical Thermodynamics of Zirconium. Amsterdam: Elsevier B. V. ISBN978-0-444-51803-3.
Burney, H. (1999). "Past, Present and Future of the Chlor-Alkali Industry". In Burney, H. S.; Furuya, N.; Hine, F.; Ota, K.-I. (eds.). Chlor-Alkali and Chlorate Technology: R. B. MacMullin Memorial Symposium. Pennington: The Electrochemical Society. pp.105–126. ISBN1-56677-244-3.
Buznik, V. M. (2009). "Fluoropolymer Chemistry in Russia: Current Situation and Prospects". Russian Journal of General Chemistry. 79 (3): 520–526. doi:10.1134/S1070363209030335. S2CID97518401.
Cheng, H.; Fowler, D. E.; Henderson, P. B.; Hobbs, J. P.; Pascolini, M. R. (1999). "On the Magnetic Susceptibility of Fluorine". The Journal of Physical Chemistry A. 103 (15): 2861–2866. Bibcode:1999JPCA..103.2861C. doi:10.1021/jp9844720.
Chisté, V.; Bé, M. M. (2011). "F-18"(PDF). In Bé, M. M.; Coursol, N.; Duchemin, B.; Lagoutine, F.; etal. (eds.). Table de radionucléides (Report). CEA (Commissariat à l'énergie atomique et aux énergies alternatives), LIST, LNE-LNHB (Laboratoire National Henri Becquerel/Commissariat à l'Energie Atomique). Archived from the original(PDF) on 11 August 2020. Retrieved 15 June 2011.
Christe, Karl O. (1986). "Chemical Synthesis of Elemental Fluorine". Inorganic Chemistry. 25 (21): 3721–3722. doi:10.1021/ic00241a001.
Eaton, Charles (1997). "Figure hfl". E-Hand.com: The Electronic Textbook of Hand Surgery. The Hand Center (former practice of Dr. Eaton). Retrieved 28 September 2013.
Edwards, Philip Neil (1994). "Use of Fluorine in Chemotherapy". In Banks, R. E.; Smart, B. E.; Tatlow, J. C. (eds.). Organofluorine Chemistry: Principles and Commercial Applications. New York: Plenum Press. pp.501–542. ISBN978-0-306-44610-8.
Einstein, F. W. B.; Rao, P. R.; Trotter, J.; Bartlett, N. (1967). "The Crystal Structure of Gold Trifluoride". Journal of the Chemical Society A: Inorganic, Physical, Theoretical. 4: 478–482. doi:10.1039/J19670000478.
El-Kareh, Badih (1994). Fundamentals of Semiconductor Processing Technology. Norwell and Dordrecht: Kluwer Academic Publishers. ISBN978-0-7923-9534-8.
El Saadi, M. S.; Hall, A. H.; Hall, P. K.; Riggs, B. S.; Augenstein, W. L.; Rumack, B. H. (1989). "Hydrofluoric Acid Dermal Exposure". Veterinary and Human Toxicology. 31 (3): 243–247. PMID2741315.
Emeléus, H. J.; Sharpe, A. G. (1974). Advances in Inorganic Chemistry and Radiochemistry. Vol.16. New York: Academic Press. ISBN978-0-08-057865-1.
Emeléus, H. J.; Sharpe, A. G. (1983). Advances in Inorganic Chemistry and Radiochemistry. Vol.27. Academic Press. ISBN0-12-023627-3.
Fischman, Michael L. (2001). "Semiconductor Manufacturing Hazards". In Sullivan, John B.; Krieger, Gary R. (eds.). Clinical Environmental Health and Toxic Exposures (2nded.). Philadelphia: Lippincott Williams & Wilkins. pp.431–465. ISBN978-0-683-08027-8.
Forster, P.; Ramaswamy, V.; Artaxo, P.; Berntsen, T.; Betts, R.; Fahey, D. W.; Haywood, J.; Lean, J.; Lowe, D. C.; Myhr e, G.; Nganga, J.; Prinn, R.; Raga, G.; Schulz, M.; Van Dorland, R. (2007). "Changes in Atmospheric Constituents and in Radiative Forcing". In Solomon, S.; Manning, M.; Chen, Z.; Marquis, M.; Averyt, K. B.; Tignor, M.; Miller, H. L. (eds.). Climate Change 2007: The Physical Science Basis. Contribution of Working Group I to the Fourth Assessment Report of the Intergovernmental Panel on Climate Change. Cambridge: Cambridge University. pp.129–234. ISBN978-0-521-70596-7.
Fulton, Robert B.; Miller, M. Michael (2006). "Fluorspar". In Kogel, Jessica Elzea; Trivedi, Nikhil C.; Barker, James M.; Krukowski, Stanley T. (eds.). Industrial Minerals & Rocks: Commodities, Markets, and Uses. Littleton: Society for Mining, Metallurgy, and Exploration (U.S.). pp.461–473. ISBN978-0-87335-233-8.
Gabriel, J. L.; Miller, T. F. Jr.; Wolfson, M. R.; Shaffer, T. H. (1996). "Quantitative Structure-Activity Relationships of Perfluorinated Hetero-Hydrocarbons as Potential Respiratory Media". ASAIO Journal. 42 (6): 968–973. doi:10.1097/00002480-199642060-00009. PMID8959271. S2CID31161098.
Godfrey, S. M.; McAuliffe, C. A.; Mackie, A. G.; Pritchard, R. G. (1998). "Inorganic Derivatives of the Elements". In Norman, Nicholas C. (ed.). Chemistry of Arsenic, Antimony and Bismuth. London: Blackie Academic & Professional. pp.67–158. ISBN978-0-7514-0389-3.
Green, S. W.; Slinn, D. S. L.; Simpson, R. N. F.; Woytek, A. J. (1994). "Perfluorocarbon Fluids". In Banks, R. E.; Smart, B. E.; Tatlow, J. C. (eds.). Organofluorine Chemistry: Principles and Applications. New York: Plenum Press. pp.89–119. ISBN978-0-306-44610-8.
Greenwood, N. N.; Earnshaw, A. (1998). Chemistry of the Elements (2nded.). Oxford: Butterworth Heinemann. ISBN0-7506-3365-4.
Gribble, G. W. (2002). "Naturally Occurring Organofluorines". In Neison, A. H. (ed.). Organofluorines. The Handbook of Environmental Chemistry. Vol.3N. Berlin: Springer. pp.121–136. doi:10.1007/10721878_5. ISBN3-540-42064-9.
Grot, Walter (2011). Fluorinated Ionomers (2nded.). Oxford and Waltham: Elsevier. ISBN978-1-4377-4457-6.
Hagmann, W. K. (2008). "The Many Roles for Fluorine in Medicinal Chemistry". Journal of Medicinal Chemistry. 51 (15): 4359–4369. doi:10.1021/jm800219f. PMID18570365.
Harbison, G. S. (2002). "The Electric Dipole Polarity of the Ground and Low-lying Metastable Excited States of NF". Journal of the American Chemical Society. 124 (3): 366–367. doi:10.1021/ja0159261. PMID11792193.
Hoogers, G. (2002). "Fuel Cell Components and Their Impact on Performance". In Hoogers, G. (ed.). Fuel Cell Technology Handbook. Boca Raton: CRC Press. pp.4-1–4-27. ISBN0-8493-0877-1.
Kacmarek, Robert M.; Wiedemann, Herbert P.; Lavin, Philip T.; Wedel, Mark K.; Tütüncü, Ahmet S.; Slutsky, Arthur S. (2006). "Partial Liquid Ventilation in Adult Patients with Acute Respiratory Distress Syndrome". American Journal of Respiratory and Critical Care Medicine. 173 (8): 882–9. doi:10.1164/rccm.200508-1196OC. PMID16254269.
Katakuse, Itsuo; Ichihara, Toshio; Ito, Hiroyuki; Sakurai, Tohru; Matsuo, Takekiyo (1999). "SIMS Experiment". In Arai, T.; Mihama, K.; Yamamoto, K.; Sugano, S. (eds.). Mesoscopic Materials and Clusters: Their Physical and Chemical Properties. Tokyo: Kodansha. pp.259–273. ISBN4-06-208635-2.
Keplinger, M. L.; Suissa, L. W. (1968). "Toxicity of Fluorine Short-Term Inhalation". American Industrial Hygiene Association Journal. 29 (1): 10–18. doi:10.1080/00028896809342975. PMID5667185.
Kissa, Erik (2001). Fluorinated Surfactants and Repellents (2nded.). New York: Marcel Dekker. ISBN978-0-8247-0472-8.
Kuriakose, A. K.; Margrave, J. L. (1965). "Kinetics of the Reactions of Elemental Fluorine. IV. Fluorination of Graphite". Journal of Physical Chemistry. 69 (8): 2772–2775. doi:10.1021/j100892a049.
Lidin, R.; Molochko, V. A.; Andreeva, L. L. (2000). Химические свойства неорганических веществ[Chemical Properties of Inorganic Substances] (in Russian). Moscow: Khimiya. ISBN5-7245-1163-0.
Liteplo, R.; Gomes, R.; Howe, P.; Malcolm, H. (2002). Environmental Health Criteria 227 (Fluoride). Geneva: United Nations Environment Programme; International Labour Organization; World Health Organization. ISBN92-4-157227-2. Retrieved 14 October 2013.
Mackay, Kenneth Malcolm; Mackay, Rosemary Ann; Henderson, W. (2002). Introduction to Modern Inorganic Chemistry (6thed.). Cheltenham: Nelson Thornes. ISBN0-7487-6420-8.
Macomber, Roger (1996). Organic chemistry. Vol.1. Sausalito: University Science Books. ISBN978-0-935702-90-3.
McCoy, M. (2007). "SURVEY Market Challenges Dim the Confidence of the World's Chemical CEOs". Chemical & Engineering News. 85 (23): 11. doi:10.1021/cen-v085n023.p011a.
Moore, John W.; Stanitski, Conrad L.; Jurs, Peter C. (2010). Principles of Chemistry: The Molecular Science. Belmont: Brooks/Cole. ISBN978-0-495-39079-4.
Morrow, S. I.; Perry, D. D.; Cohen, M. S. (1959). "The Formation of Dinitrogen Tetrafluoride in the Reaction of Fluorine and Ammonia". Journal of the American Chemical Society. 81 (23): 6338–6339. doi:10.1021/ja01532a066.
The National Institute for Occupational Safety and Health (1994). "Fluorine". Documentation for Immediately Dangerous To Life or Health Concentrations (IDLHs). Retrieved 15 January 2014.
The National Institute for Occupational Safety and Health (1994). "Chlorine". Documentation for Immediately Dangerous To Life or Health Concentrations (IDLHs). Retrieved 13 July 2014.
O'Hagan, D. (2008). "Understanding Organofluorine Chemistry. An Introduction to the C–F Bond". Chemical Society Reviews. 37 (2): 308–319. doi:10.1039/b711844a. PMID18197347.
Okada, T.; Xie, G.; Gorseth, O.; Kjelstrup, S.; Nakamura, N.; Arimura, T. (1998). "Ion and Water Transport Characteristics of Nafion Membranes as Electrolytes". Electrochimica Acta. 43 (24): 3741–3747. doi:10.1016/S0013-4686(98)00132-7.
Parente, Luca (2001). "The Development of Synthetic Glucocorticoids". In Goulding, Nicolas J.; Flower, Rod J. (eds.). Glucocorticoids. Basel: Birkhäuser. pp.35–53. ISBN978-3-7643-6059-7.
Partington, J. R. (1923). "The early history of hydrofluoric acid". Memoirs and Proceedings of the Manchester Literary and Philosophical Society. 67 (6): 73–87.
Patnaik, Pradyot (2007). A Comprehensive Guide to the Hazardous Properties of Chemical Substances (3rded.). Hoboken: John Wiley & Sons. ISBN978-0-471-71458-3.
Posner, Stefan (2011). "Perfluorinated Compounds: Occurrence and Uses in Products". In Knepper, Thomas P.; Large, Frank T. (eds.). Polyfluorinated Chemicals and Transformation Products. Heidelberg: Springer Science+Business Media. pp.25–40. ISBN978-3-642-21871-2.
Preskorn, Sheldon H. (1996). Clinical Pharmacology of Selective Serotonin Reuptake Inhibitors. Caddo: Professional Communications. ISBN978-1-884735-08-0.
Pyykkö, Pekka; Atsumi, Michiko (2009). "Molecular Double-Bond Covalent Radii for Elements Li–E112". Chemistry: A European Journal. 15 (46): 12770–9. doi:10.1002/chem.200901472. PMID19856342.
Raghavan, P. S. (1998). Concepts and Problems in Inorganic Chemistry. Delhi: Discovery Publishing House. ISBN978-81-7141-418-5.
Raj, P. Prithvi; Erdine, Serdar (2012). Pain-Relieving Procedures: The Illustrated Guide. Chichester: John Wiley & Sons. ISBN978-0-470-67038-5.
Ramkumar, Jayshree (2012). "Nafion Perfluorosulphonate Membrane: Unique Properties and Various Applications". In Banerjee, S.; Tyagi, A. K. (eds.). Functional Materials: Preparation, Processing and Applications. London and Waltham: Elsevier. pp.549–578. ISBN978-0-12-385142-0.
Rhoades, David Walter (2008). Broadband Dielectric Spectroscopy Studies of Nafion (PhD dissertation). Ann Arbor: University of Southern Mississippi, MS. ISBN978-0-549-78540-8.
Richter, M.; Hahn, O.; Fuchs, R. (2001). "Purple Fluorite: A Little Known Artists' Pigment and Its Use in Late Gothic and Early Renaissance Painting in Northern Europe". Studies in Conservation. 46 (1): 1–13. doi:10.1179/sic.2001.46.1.1. JSTOR1506878. S2CID191611885.
Riedel, Sebastian; Kaupp, Martin (2009). "The highest oxidation states of the transition metal elements". Coordination Chemistry Reviews. 253 (5–6): 606–624. doi:10.1016/j.ccr.2008.07.014.
Roblin, I.; Urban, M.; Flicoteau, D.; Martin, C.; Pradeau, D. (2006). "Topical Treatment of Experimental Hydrofluoric Acid Skin Burns by 2.5% Calcium Gluconate". Journal of Burn Care & Research. 27 (6): 889–894. doi:10.1097/01.BCR.0000245767.54278.09. PMID17091088. S2CID3691306.
Salager, Jean-Louis (2002). Surfactants: Types and Uses(PDF). FIRP Booklet # 300-A. Laboratory of Formulation, Interfaces, Rheology, and Processes, Universidad de los Andes. Retrieved 13 October 2013.
Schimmeyer, S. (2002). "The Search for a Blood Substitute". Illumin. 15 (1). Columbia: University of Southern Carolina. Archived from the original on 2 October 2011. Retrieved 15 October 2013.
Schmedt Auf Der Günne, Jörn; Mangstl, Martin; Kraus, Florian (2012). "Occurrence of Difluorine F2 in Nature-In Situ Proof and Quantification by NMR Spectroscopy". Angewandte Chemie International Edition. 51 (31): 7847–7849. doi:10.1002/anie.201203515. ISSN1521-3773. PMID22763992.
Senning, A. (2007). Elsevier's Dictionary of Chemoetymology: The Whies and Whences of Chemical Nomenclature and Terminology. Amsterdam and Oxford: Elsevier. ISBN978-0-444-52239-9.
Shulman, J. D.; Wells, L. M. (1997). "Acute Fluoride Toxicity from Ingesting Home-use Dental Products in Children, Birth to 6 Years of Age". Journal of Public Health Dentistry. 57 (3): 150–158. doi:10.1111/j.1752-7325.1997.tb02966.x. PMID9383753.
Theodoridis, George (2006). "Fluorine-Containing Agrochemicals: An Overview of Recent Developments". In Tressaud, Alain (ed.). Fluorine and the Environment: Agrochemicals, Archaeology, Green Chemistry & Water. Amsterdam and Oxford: Elsevier. pp.121–176. ISBN978-0-444-52672-4.
Villalba, Gara; Ayres, Robert U.; Schroder, Hans (2008). "Accounting for Fluorine: Production, Use, and Loss". Journal of Industrial Ecology. 11: 85–101. doi:10.1162/jiec.2007.1075. S2CID153740615.
Walsh, Kenneth A. (2009). Beryllium Chemistry and Processing. Materials Park: ASM International. ISBN978-0-87170-721-5.
Zareitalabad, P.; Siemens, J.; Hamer, M.; Amelung, W. (2013). "Perfluorooctanoic acid (PFOA) and perfluorooctanesulfonic acid (PFOS) in surface waters, sediments, soils and wastewater – A review on concentrations and distribution coefficients". Chemosphere. 91 (6): 725–32. Bibcode:2013Chmsp..91..725Z. doi:10.1016/j.chemosphere.2013.02.024. PMID23498059.
Zorich, Robert (1991). Handbook of Quality Integrated Circuit Manufacturing. San Diego: Academic Press. ISBN978-0-323-14055-3.
The halogens are a group in the periodic table consisting of six chemically related elements: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and the radioactive elements astatine (At) and tennessine (Ts), though some authors would exclude tennessine as its chemistry is unknown and is theoretically expected to be more like that of gallium. In the modern IUPAC nomenclature, this group is known as group 17.
Fluoride is an inorganic, monatomic anion of fluorine, with the chemical formula F− , whose salts are typically white or colorless. Fluoride salts typically have distinctive bitter tastes, and are odorless. Its salts and minerals are important chemical reagents and industrial chemicals, mainly used in the production of hydrogen fluoride for fluorocarbons. Fluoride is classified as a weak base since it only partially associates in solution, but concentrated fluoride is corrosive and can attack the skin.
Hydrofluoric acid is a solution of hydrogen fluoride (HF) in water. Solutions of HF are colorless, acidic and highly corrosive. A common concentration is 49% (48-52%) but there are also stronger solutions and pure HF has a boiling point near room temperature. It is used to make most fluorine-containing compounds; examples include the commonly used pharmaceutical antidepressant medication fluoxetine (Prozac) and the material PTFE (Teflon). Elemental fluorine is produced from it. It is commonly used to etch glass and silicon wafers.
Calcium fluoride is the inorganic compound of the elements calcium and fluorine with the formula CaF2. It is a white solid that is practically insoluble in water. It occurs as the mineral fluorite (also called fluorspar), which is often deeply coloured owing to impurities.
Caesium fluoride is an inorganic compound with the formula CsF. A hygroscopic white salt, caesium fluoride is used in the synthesis of organic compounds as a source of the fluoride anion. The compound is noteworthy from the pedagogical perspective as caesium also has the highest electropositivity of all commonly available elements and fluorine has the highest electronegativity.
Silver(II) fluoride is a chemical compound with the formula AgF2. It is a rare example of a silver(II) compound - silver usually exists in its +1 oxidation state. It is used as a fluorinating agent.
Hydrogen fluoride (fluorane) is an inorganic compound with chemical formula HF. It is a very poisonous, colorless gas or liquid that dissolves in water to yield hydrofluoric acid. It is the principal industrial source of fluorine, often in the form of hydrofluoric acid, and is an important feedstock in the preparation of many important compounds including pharmaceuticals and polymers such as polytetrafluoroethylene (PTFE). HF is also widely used in the petrochemical industry as a component of superacids. Due to strong and extensive hydrogen bonding, it boils at near room temperature, which is much higher of a temperature than other hydrogen halides.
Xenon difluoride is a powerful fluorinating agent with the chemical formula XeF 2, and one of the most stable xenon compounds. Like most covalent inorganic fluorides it is moisture-sensitive. It decomposes on contact with water vapor, but is otherwise stable in storage. Xenon difluoride is a dense, colourless crystalline solid.
Sulfur tetrafluoride is a chemical compound with the formula SF4. It is a colorless corrosive gas that releases dangerous hydrogen fluoride gas upon exposure to water or moisture. Sulfur tetrafluoride is a useful reagent for the preparation of organofluorine compounds, some of which are important in the pharmaceutical and specialty chemical industries.
Aluminium fluoride is an inorganic compound with the formula AlF3. It forms hydrates AlF3·xH2O. Anhydrous AlF3 and its hydrates are all colorless solids. Anhydrous AlF3 is used in the production of aluminium. Several occur as minerals.
Antimony trifluoride is the inorganic compound with the formula SbF3. Sometimes called Swarts' reagent, it is one of two principal fluorides of antimony, the other being SbF5. It appears as a white solid. As well as some industrial applications, it is used as a reagent in inorganic and organofluorine chemistry.
Organofluorine chemistry describes the chemistry of organofluorine compounds, organic compounds that contain a carbon–fluorine bond. Organofluorine compounds find diverse applications ranging from oil and water repellents to pharmaceuticals, refrigerants, and reagents in catalysis. In addition to these applications, some organofluorine compounds are pollutants because of their contributions to ozone depletion, global warming, bioaccumulation, and toxicity. The area of organofluorine chemistry often requires special techniques associated with the handling of fluorinating agents.
Electrochemical fluorination (ECF), or electrofluorination, is a foundational organofluorine chemistry method for the preparation of fluorocarbon-based organofluorine compounds. The general approach represents an application of electrosynthesis. The fluorinated chemical compounds produced by ECF are useful because of their distinctive solvation properties and the relative inertness of carbon–fluorine bonds. Two ECF synthesis routes are commercialized and commonly applied: the Simons process and the Phillips Petroleum process. It is also possible to electrofluorinate in various organic media. Prior to the development of these methods, fluorination with fluorine, a dangerous oxidizing agent, was a dangerous and wasteful process. ECF can be cost-effective, but it may also result in low yields.
Fluorination by sulfur tetrafluoride produces organofluorine compounds from oxygen-containing organic functional groups using sulfur tetrafluoride. The reaction has broad scope, and SF4 is an inexpensive reagent. It is however hazardous gas whose handling requires specialized apparatus. Thus, for many laboratory scale fluorinations diethylaminosulfur trifluoride ("DAST") is used instead.
Vanadium(V) fluoride is the inorganic compound with the chemical formula VF5. It is a colorless volatile liquid that freezes near room temperature. It is a highly reactive compound, as indicated by its ability to fluorinate organic substances.
Fluorine forms a great variety of chemical compounds, within which it always adopts an oxidation state of −1. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. Most frequently, covalent bonds involving fluorine atoms are single bonds, although at least two examples of a higher order bond exist. Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding. Fluorine's chemistry includes inorganic compounds formed with hydrogen, metals, nonmetals, and even noble gases; as well as a diverse set of organic compounds. For many elements the highest known oxidation state can be achieved in a fluoride. For some elements this is achieved exclusively in a fluoride, for others exclusively in an oxide; and for still others the highest oxidation states of oxides and fluorides are always equal.
Fluorine may interact with biological systems in the form of fluorine-containing compounds. Though elemental fluorine (F2) is very rare in everyday life, fluorine-containing compounds such as fluorite occur naturally as minerals. Naturally occurring organofluorine compounds are extremely rare. Man-made fluoride compounds are common and are used in medicines, pesticides, and materials. Twenty percent of all commercialized pharmaceuticals contain fluorine, including Lipitor and Prozac. In many contexts, fluorine-containing compounds are harmless or even beneficial to living organisms; in others, they are toxic.
The global market for chemicals from fluorine was about US$16 billion per year as of 2006. The industry was predicted to reach 2.6 million metric tons per year by 2015. The largest market is the United States. Western Europe is the second largest. Asia Pacific is the fastest growing region of production. China in particular has experienced significant growth as a fluorochemical market and is becoming a producer of them as well. Fluorite mining was estimated in 2003 to be a $550 million industry, extracting 4.5 million tons per year.
Fluorine is a relatively new element in human applications. In ancient times, only minor uses of fluorine-containing minerals existed. The industrial use of fluorite, fluorine's source mineral, was first described by early scientist Georgius Agricola in the 16th century, in the context of smelting. The name "fluorite" derives from Agricola's invented Latin terminology. In the late 18th century, hydrofluoric acid was discovered. By the early 19th century, it was recognized that fluorine was a bound element within compounds, similar to chlorine. Fluorite was determined to be calcium fluoride.
Iodyl fluoride is an inorganic compound of iodine, fluorine, and oxygen with the chemical formula IO2F. The compound was initially synthesized in 1951.
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