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• • • Fluorine is a with symbol F and 9. It is the lightest and exists as a highly toxic pale yellow gas. As the most element, it is extremely reactive: almost all other elements, including some, form compounds with fluorine. Among the elements, fluorine ranks., 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, the Latin verb fluo meaning '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 isolate elemental fluorine using low-temperature, a process still employed for modern production. Industrial production of fluorine gas for, its largest application, began during the in. Owing to the expense of refining pure fluorine, most commercial applications use fluorine compounds, with about half of mined fluorite used in. The rest of the fluorite is converted into corrosive en route to various organic fluorides, or into which plays a key role in. Organic fluorides have very high chemical and thermal stability; their major uses are as, electrical insulation and cookware, the last as (Teflon).

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Pharmaceuticals such as and also contain fluorine, and the inhibits dental cavities, and so finds use in toothpaste and. Sales amount to more than 15 billion a year. Gases are generally with 100 to 20,000 times that of. Persist in the environment due to the strength of the. Fluorine has no known metabolic role in mammals; a few plants synthesize organofluorine poisons that deter herbivores.

Simplified structure of the fluorine atom Fluorine atoms have nine electrons, one fewer than, and 1s 22s 22p 5: 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, and experience a high of 9 − 2 = 7; this affects the atom's physical properties. Fluorine's is third-highest among all elements, behind helium and neon, which complicates the removal of electrons from neutral fluorine atoms.

It also has a high, second only to, and tends to capture an electron to become with the noble gas neon; it has the highest of any element. Fluorine atoms have a small of around 60, similar to those of its neighbors oxygen and neon.

Main article: The of is much lower than that of either Cl 2 or Br 2 and similar to the easily cleaved bond; this, along with high electronegativity, accounts for fluorine's easy dissociation, high reactivity, and strong bonds to non-fluorine atoms. Conversely, bonds to other atoms are very strong because of fluorine's high electronegativity. Unreactive substances like powdered steel, glass fragments, and fibers react quickly with cold fluorine gas; wood and water spontaneously combust under a fluorine jet. External video during fluorine reactions Fluorine Reactions of elemental fluorine with metals require varying conditions. Cause explosions and display vigorous activity in bulk; to prevent from the formation of metal fluoride layers, most other metals such as aluminium and iron must be powdered, and require pure fluorine gas at 300–450 °C (575–850 °F).

Some solid nonmetals (sulfur, phosphorus) react vigorously in liquid air temperature fluorine. And combine readily with fluorine, the latter sometimes explosively; exhibits much less activity, requiring elevated temperatures., like some of the alkali metals, reacts explosively with fluorine., as, reacts at room temperature to yield. Graphite combines with fluorine above 400 °C (750 °F) to produce; higher temperatures generate gaseous, sometimes with explosions.

Carbon dioxide and carbon monoxide react at or just above room temperature, whereas and other organic chemicals generate strong reactions: even fully substituted such as, normally incombustible, may explode. Although is stable, nitrogen requires an at elevated temperatures for reaction with fluorine to occur, due to the very strong in elemental nitrogen; ammonia may react explosively. 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. Heavier halogens react readily with fluorine as does the noble gas; of the other noble gases, only and react, and only under special conditions. 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, pale yellow when pure (sometimes described as yellow-green). It has a characteristic pungent odor detectable at 20. Fluorine condenses into a bright yellow liquid at −188 °C (−306 °F), a transition temperature similar to those of oxygen and nitrogen. Fluorine has two solid forms, α- and β-fluorine.

The latter crystallizes at −220 °C (−364 °F) and is transparent and soft, with the same disordered structure of freshly crystallized solid oxygen, unlike the systems of other solid halogens. Further cooling to −228 °C (−378 °F) induces a into opaque and hard α-fluorine, which has a structure with dense, angled layers of molecules. The transition from β- to α-fluorine is more than the condensation of fluorine, and can be violent. Isotopes [ ]. Main article: Only one of fluorine occurs naturally in abundance, the stable isotope 19 F.

It has a high and exceptional sensitivity to magnetic fields; because it is also, it is in. Seventeen with from 14 to 31 have been synthesized, of which is the most stable with a of 109.77 minutes. Other radioisotopes have half-lives less than 70 seconds; most decay in less than half a second.

The isotopes 17 F and 18 F undergo and, lighter isotopes decay by, and those heavier than 19 F undergo (the heaviest ones - with delayed ). Two of fluorine are known, 18m F, with a half-life of 162(7) nanoseconds, and 26m F, with a half-life of 2.2(1) milliseconds. Occurrence [ ]. Main article: Universe [ ] Solar System abundances Atomic number Element Relative amount 6 Carbon 4,800 7 Nitrogen 1,500 8 Oxygen 8,800 9 Fluorine 1 10 Neon 1,400 11 Sodium 24 12 Magnesium 430 Among the lighter elements, fluorine's abundance value of 400 (parts per billion) – 24th among elements in the universe – is exceptional: other elements from carbon to magnesium are twenty or more times as common.

This is because processes bypass fluorine, and any fluorine atoms otherwise created have high, allowing further fusion with hydrogen or helium to generate oxygen or neon respectively. Beyond this transient existence, three explanations have been proposed for the presence of fluorine: • during, bombardment of neon atoms by could transmute them to fluorine; • the solar wind of could blow fluorine away from any hydrogen or helium atoms; or • fluorine is borne out on convection currents arising from fusion in stars.

See also: Fluorine is the thirteenth most at 600–700 ppm (parts per million) by mass. Elemental fluorine in Earth's atmosphere would easily react with atmospheric, precluding its natural occurrence; it is found only in combined mineral forms, of which, and are the most industrially significant. Fluorite or fluorspar ( CaF 2), colorful and abundant worldwide, is fluorine's main source; China and Mexico are the major suppliers. Led extraction in the early 20th century but ceased mining in 1995. Although fluorapatite (Ca 5(PO 4) 3F) contains most of the world's fluorine, its low of 3.5% means that most of it is used as a phosphate.

Small quantities of fluorine compounds are obtained via, a phosphate industry byproduct. Cryolite ( Na 3AlF 6), once used directly in aluminium production, is the rarest and most concentrated of these three minerals.

The main commercial mine on Greenland's west coast closed in 1987, and most cryolite is now synthesized. Major fluorine-containing minerals Fluorite Fluorapatite Cryolite Other minerals such as contain fluorine. Fluorides, unlike other halides, are insoluble and do not occur in commercially favorable concentrations in saline waters. Trace quantities of organofluorines of uncertain origin have been detected in volcanic eruptions and geothermal springs. The existence of gaseous fluorine in crystals, suggested by the smell of crushed, is contentious; a 2012 study reported the presence of 0.04% F 2 by weight in antozonite, attributing these to radiation from the presence of tiny amounts of. An ampoule of or hex The division of (GM) experimented with chlorofluorocarbon refrigerants in the late 1920s, and was formed as a joint venture between GM and in 1930 hoping to market Freon-12 ( ) as one such.

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 compounds. (Teflon) was serendipitously discovered in 1938 by while working on refrigerants at Kinetic, and its superlative chemical and thermal resistance lent it to accelerated commercialization and mass production by 1941. 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 and the used huge quantities to produce for uranium enrichment. Since UF 6 is as corrosive as fluorine, 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. Compounds [ ].

See also: Alkali metals form ionic and highly soluble; these have the and analogous chlorides. Alkaline earth possess strong ionic bonds but are insoluble in water, with the exception of, which also exhibits some covalent character and has a -like structure. And many other metals form mostly ionic.

Covalent bonding first comes to prominence in the: those of zirconium, hafnium and several actinides are ionic with high melting points, while those of titanium, vanadium, and niobium are polymeric, melting or decomposing at no more than 350 °C (660 °F). Continue this trend with their linear polymers and complexes. Thirteen metal are known, all octahedral, and are mostly volatile solids but for liquid and, and gaseous., the only characterized metal, is a low-melting molecular solid with. Metal fluorides with more fluorine atoms are particularly reactive.

Structural progression of metal fluorides Sodium fluoride, ionic Bismuth pentafluoride, polymeric Rhenium heptafluoride, molecular Hydrogen [ ]. 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. It boils at a much higher temperature than heavier hydrogen halides and unlike them is fully with water. 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, hydrofluoric acid is a at low concentrations. However, it can attack glass, something the other acids cannot do.

Other reactive nonmetals [ ] Metalloids are included in this section. 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. And heavier nonmetals can form fluorides. Is planar and possesses an incomplete octet. It functions as a and combines with Lewis bases like ammonia to form.

Is tetrahedral and inert; analogues, silicon and germanium tetrafluoride, are also tetrahedral but behave as Lewis acids. The form trifluorides that increase in reactivity and basicity with higher molecular weight, although resists hydrolysis and is not basic. The pentafluorides of phosphorus, arsenic, and antimony are more reactive than their respective trifluorides, with the strongest neutral Lewis acid known. 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 is especially inert. Chlorine, bromine, and iodine can each form mono-, tri-, and pentafluorides, but only has been characterized among possible heptafluorides.

Many of them are powerful sources of fluorine atoms, and industrial applications using chlorine trifluoride require precautions similar to those using fluorine. Noble gases [ ]. These xenon tetrafluoride crystals were photographed in 1962. The compound's synthesis, as with xenon hexafluoroplatinate, surprised many chemists., having complete electron shells, defied reaction with other elements until 1962 when reported synthesis of;,,, and multiple oxyfluorides have been isolated since then. Among other noble gases, krypton forms a, and radon and fluorine generate a solid suspected to be. Binary fluorides of lighter noble gases are exceptionally unstable: argon and hydrogen fluoride combine under extreme conditions to give.

Helium and neon have no long-lived fluorides, and no neon fluoride has ever been observed; helium fluorohydride has been detected for milliseconds at high pressures and low temperatures. Install Block Erupter Windows 7. Organic compounds [ ]. Main articles: and The substitution of hydrogen atoms in an 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, in which all hydrogen atoms are substituted, are insoluble in most organic solvents, reacting at ambient conditions only with sodium in liquid ammonia.

The term is used for what would otherwise be a perfluorocarbon if not for the presence of a, often a. These compounds share many properties with perfluorocarbons such as stability and, while the functional group augments their reactivity, enabling them to adhere to surfaces or act as;, in particular, can lower the of water more than their hydrocarbon-based analogues., which have some unfluorinated carbon atoms near the functional group, are also regarded as perfluorinated. Polymers [ ] Polymers exhibit the same stability increases afforded by fluorine substitution (for hydrogen) in discrete molecules; their melting points generally increase too. (PTFE), the simplest fluoropolymer and perfluoro analogue of with – CF 2–, demonstrates this change as expected, but its very high melting point makes it difficult to mold.

Various PTFE derivatives are less temperature-tolerant but easier to mold: replaces some fluorine atoms with groups, do the same with groups, and contains perfluoroether side chains capped with groups. Other fluoropolymers retain some hydrogen atoms; has half the fluorine atoms of PTFE and has a quarter, but both behave much like perfluorinated polymers. Production [ ] Industrial [ ]. Industrial fluorine cells at Moissan's method is used to produce industrial quantities of fluorine, via the electrolysis of a potassium fluoride/hydrogen fluoride mixture: hydrogen and fluoride ions are reduced and oxidized at a steel container and a carbon block, under 8–12 volts, to generate hydrogen and fluorine gas respectively.

Temperatures are elevated, KF•2HF melting at 70 °C (158 °F) and being electrolyzed at 70–130 °C (158–266 °F). KF, which acts as catalyst, is essential since pure HF cannot be electrolyzed.

Fluorine can be stored in steel cylinders that have passivated interiors, at temperatures below 200 °C (392 °F); otherwise nickel can be used. Regulator valves and pipework are made of nickel, the latter possibly using instead.

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; some sources instead recommend nickel-Monel-PTFE systems. Chemical [ ] While preparing for a 1986 conference to celebrate the centennial of Moissan's achievement, 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: 2 KMnO 4 + 2 KF + 10 HF + 3 H 2O 2 → 2 K 2MnF 6 + 8 H 2O + 3 O 2↑ 2 K 2MnF 6 + 4 SbF 5 → 4 KSbF 6 + 2 MnF 3 + F 2↑ Christe later commented that the reactants 'had been known for more than 100 years and even Moissan could have come up with this scheme.'

As late as 2008, some references still asserted that fluorine was too reactive for any chemical isolation. Industrial applications [ ]. Main article: Fluorite mining, which supplies most global fluorine, peaked in 1989 when 5.6 million of ore were extracted.

Chlorofluorocarbon restrictions lowered this to 3.6 million tons in 1994; production has since been increasing. Around 4.5 million tons of ore and revenue of 550 million were generated in 2003; later reports estimated 2011 global fluorochemical sales at $15 billion and predicted 2016–18 production figures of 3.5 to 5.9 million tons, and revenue of at least $20 billion. 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 hydrogen fluoride. See also: 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 because of handling challenges. Most processes using free fluorine in large amounts employ in situ generation under. The largest application of fluorine gas, consuming up to 7,000 metric tons annually, is in the preparation of UF 6 for the.

Fluorine is used to fluorinate, itself formed from uranium dioxide and hydrofluoric acid. 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. About 6,000 metric tons per year go into producing the inert SF 6 for high-voltage transformers and circuit breakers, eliminating the need for hazardous associated with oil-filled devices. Several fluorine compounds are used in electronics: rhenium and tungsten hexafluoride in, in and in cleaning equipment.

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 instead. Inorganic fluorides [ ]. Aluminium extraction depends critically on cryolite As with other iron alloys, around 3 kg (6.5 lb) metspar is added to each metric ton of steel; the fluoride ions lower its melting point and.

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, glass etching and. One-third of HF goes into synthesizing and, both fluxes in the for aluminium extraction; replenishment is necessitated by their occasional reactions with the smelting apparatus. Each metric ton of aluminium requires about 23 kg (51 lb) of flux. Fluorosilicates consume the second largest portion, with used in water fluoridation and laundry effluent treatment, and as an intermediate en route to cryolite and silicon tetrafluoride. Other important inorganic fluorides include those of,, and.

Organic fluorides [ ] Organofluorides consume over 20% of mined fluorite and over 40% of hydrofluoric acid, with dominating and increasing their market share. Are a minor application but generate over $1 billion in annual revenue. Due to the danger from direct hydrocarbon–fluorine reactions above −150 °C (−240 °F), industrial fluorocarbon production is indirect, mostly through such as, in which chlorocarbon chlorines are substituted for fluorines by hydrogen fluoride under catalysts. Subjects hydrocarbons to electrolysis in hydrogen fluoride, and the treats them with solid fluorine carriers like. Refrigerant gases [ ]. See also: Halogenated refrigerants, termed Freons in informal contexts, are identified by that denote the amount of fluorine, chlorine, carbon, and hydrogen present. (CFCs) like,, and 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. 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,, and. The main HFC is with a new type of molecule, a (HFO) coming to prominence owing to its of less than 1% that of HFC-134a. Polymers [ ]. Main article: About 180,000 metric tons of fluoropolymers were produced in 2006 and 2007, generating over $3.5 billion revenue per year. The global market was estimated at just under $6 billion in 2011 and was predicted to grow by 6.5% per year up to 2016.

Fluoropolymers can only be formed by free radicals. Polytetrafluoroethylene (PTFE), sometimes called by its DuPont name Teflon, represents 60–80% by mass of the world's fluoropolymer production. The largest application is in since PTFE is an excellent. 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 for stadium roofs. The major consumer application is for. Jerked PTFE film becomes expanded PTFE (ePTFE), a fine-pored sometimes referred to by the brand name and used for rainwear,, and; may be made into and. Other fluoropolymers, including, 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. The chemically resistant (but expensive) fluorinated are used as electrochemical cell membranes, of which the first and most prominent example is.

Developed in the 1960s, it was initially deployed as fuel cell material in spacecraft and then replaced mercury-based cells. Recently, the fuel cell application has reemerged with efforts to install fuel cells into automobiles. Such as are fluoropolymer mixtures mainly used in; (C 4F 10) is used as a fire-extinguishing agent. Surfactants [ ]. Main articles: and Fluorosurfactants are small organofluorine molecules used for repelling water and stains.

Although expensive (comparable to pharmaceuticals at $200–2000 per kilogram), they yielded over $1 billion in annual revenues by 2006; alone generated over $300 million in 2000. Fluorosurfactants are a minority in the overall surfactant market, most of which is taken up by much cheaper hydrocarbon-based products. Applications in are burdened by costs; this use was valued at only $100 million in 2006. Agrichemicals [ ] About 30% of contain fluorine, most of them and with a few. Fluorine substitution, usually of a single atom or at most a group, is a robust modification with effects analogous to fluorinated pharmaceuticals: increased biological stay time, membrane crossing, and altering of molecular recognition. Is a prominent example, with large-scale use in the U.S.

As a weedkiller, but it is a suspected carcinogen and has been banned in many European countries. (1080) is a mammalian poison in which two hydrogens are replaced with fluorine and sodium; it disrupts cell metabolism by replacing acetate in the. First synthesized in the late 19th century, it was recognized as an insecticide in the early 20th, and was later deployed in its current use. New Zealand, the largest consumer of 1080, uses it to protect from the invasive Australian. Europe and the U.S. Have banned 1080. Medicinal applications [ ] Dental care [ ].

Main articles:,, and Population studies from the mid-20th century onwards show fluoride reduces. This was first attributed to the conversion of tooth enamel 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.

After studies of children in areas where fluoride was naturally present in drinking water, controlled fluoridation to fight tooth decay began in the 1940s and is now applied to water supplying 6 percent of the global population, including two-thirds of Americans. Reviews of the scholarly literature in 2000 and 2007 associated water fluoridation with a significant reduction of tooth decay in children.

Despite such endorsements and evidence of no adverse effects other than mostly benign, still exists on ethical and safety grounds. The benefits of fluoridation have lessened, possibly due to other fluoride sources, but are still measurable in low-income groups. And sometimes sodium or are often found in fluoride, first introduced in the U.S. In 1955 and now ubiquitous in developed countries, alongside fluoridated mouthwashes, gels, foams, and varnishes. Pharmaceuticals [ ].

Capsules Twenty percent of modern pharmaceuticals contain fluorine. One of these, the cholesterol-reducer (Lipitor), made more revenue than any other drug until it became generic in 2011.

The combination asthma prescription, a top-ten revenue drug in the mid-2000s, contains two active ingredients, one of which – – is fluorinated. Many drugs are fluorinated to delay inactivation and lengthen dosage periods because the carbon–fluorine bond is very stable. Fluorination also increases because the bond is more hydrophobic than the, and this often helps in cell membrane penetration and hence. And other pre-1980s had several side effects due to their non-selective interference with other than the target; the fluorinated was selective and one of the first to avoid this problem.

Many current antidepressants receive this same treatment, including the, its isomer, and and. Are artificial that are often fluorinated to enhance their effects. These include and. Fluorine also finds use in steroids: is a blood pressure-raising, and and are strong. The majority of inhaled are heavily fluorinated; the prototype is much more inert and potent than its contemporaries. Later compounds such as the fluorinated and are better than halothane and are almost insoluble in blood, allowing faster waking times. PET scanning [ ].

See also: and 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. 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. One such product,, has been through initial clinical trials.

These substances can aid endurance athletes and are banned from sports; one cyclist's near death in 1998 prompted an investigation into their abuse. 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. An Alliance Pharmaceuticals effort reached clinical trials but was abandoned because the results were not better than normal therapies. Biological role [ ]. The is one of the few organofluorine-synthesizing organisms Fluorine is for humans or other mammals; small amounts may be beneficial for bone strength, but this has not been definitively established. As there are many environmental sources of trace fluorine, the possibility of a could apply only to artificial diets.

Natural organofluorines have been found in microorganisms and plants but not animals. The most common is, which is used as a by at least 40 plants in Africa, Australia and Brazil. Other examples include terminally fluorinated,, and 2-fluorocitrate. An enzyme that binds fluorine to carbon – – was discovered in bacteria in 2002.

Toxicity [ ]. Hazard signs for commercially transported fluorine Elemental fluorine is highly toxic to living organisms. Its effects in humans start at concentrations lower than 's 50 ppm and are similar to those of chlorine: significant irritation of the eyes and respiratory system as well as liver and kidney damage occur above 25 ppm, which is the value for fluorine. Eyes and noses are seriously damaged at 100 ppm, and inhalation of 1,000 ppm fluorine will cause death in minutes, compared to 270 ppm for hydrogen cyanide. See also: 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 and possible death through.

Insoluble calcium fluoride formation triggers strong pain and burns larger than 160 cm 2 (25 in 2) can cause serious systemic toxicity. 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. 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.

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. Fluoride ion [ ]. See also: Soluble fluorides are moderately toxic: 5–10 g sodium fluoride, or 32–64 mg fluoride ions per kilogram of body mass, represents a lethal dose for adults. One-fifth of the lethal dose can cause adverse health effects, and chronic excess consumption may lead to, which affects millions in Asia and Africa. 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. Exposure limits are determined by urine testing of the body's ability to clear fluoride ions.

Historically, most cases of fluoride poisoning have been caused by accidental ingestion of insecticides containing inorganic fluorides. Most current calls to poison control centers for possible fluoride poisoning come from the ingestion of fluoride-containing toothpaste. Malfunctioning water fluoridation equipment is another cause: one incident in Alaska affected almost 300 people and killed one person. Dangers from toothpaste are aggravated for small children, and the recommends supervising children below six brushing their teeth so that they do not swallow toothpaste. 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. A larger study across the U.S. Had similar findings: 80% of cases involved children under six, and there were few serious cases. Environmental concerns [ ] Atmosphere [ ]. See also: and The, signed in 1987, set strict regulations on chlorofluorocarbons (CFCs) and 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. Even with the ban, and early indications of its efficacy, predictions warned that several generations would pass before full recovery.

With one-tenth the ODP of CFCs, hydrochlorofluorocarbons (HCFCs) are the current replacements, and are themselves scheduled for substitution by 2030–2040 by hydrofluorocarbons (HFCs) with no chlorine and zero ODP. In 2007 this date was brought forward to 2020 for developed countries; the had already prohibited one HCFC's production and capped those of two others in 2003.

Fluorocarbon gases are generally with (GWPs) of about 100 to 10,000; sulfur hexafluoride has a value of around 20,000. An outlier is which is a new type of refrigerant called a (HFO) and has attracted global demand due to its GWP of 4 compared to 1,430 for the current refrigerant standard. Biopersistence [ ]. Main article: Organofluorines exhibit biopersistence due to the strength of the carbon–fluorine bond. (PFAAs), which are sparingly water-soluble owing to their acidic functional groups, are noted; (PFOS) and (PFOA) are most often researched. 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. In the body, PFAAs bind to proteins such as; 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. 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. See also [ ] •, which measures fluoride concentration • • •, a process used to separate reagents from • and • and fluorination Notes [ ]. • 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. For further detail on α-fluorine, see the 1970 structure by Pauling.

For further detail on the concept of disorder in crystals, see the referenced general reviews. • A loud click is heard.

Samples may shatter and sample windows blow out. • 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.' • supposedly described fluorite in the late 15th century, but because his writings were uncovered 200 years later, this work's veracity is doubtful. • Or perhaps from as early as 1670 onwards; Partington and Weeks give differing accounts.

• Fl, since 2012, is used for. •,,, and the Irish chemists Thomas and George Knox were injured. Belgian chemist and French chemist died. Moissan also experienced serious hydrogen fluoride poisoning. • Also honored was his invention of the. • Fluorine in F 2 is defined to have oxidation state 0.

The unstable species F − 2 and F − 3, which decompose at around 40 K, have intermediate oxidation states; F + 4 and a few related species are predicted to be stable. • The metastable and have higher-order fluorine bonds, and some use it as a. Is another possibility. • ZrF 4 melts at 932 °C (1710 °F), HfF 4 sublimes at 968 °C (1774 °F), and UF 4 melts at 1036 °C (1897 °F). • These thirteen are those of molybdenum, technetium, ruthenium, rhodium, tungsten, rhenium, osmium, iridium, platinum, polonium, uranium, neptunium, and plutonium. • See also the explanation by Clark.

• Carbon tetrafluoride is formally organic, but is included here rather than in the – where more complex carbon-fluorine compounds are discussed – for comparison with SiF 4 and GeF 4. • Perfluorocarbon and fluorocarbon are 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. • This DuPont trademark is sometimes further misused for CFCs, HFCs, or HCFCs. • American sheep and cattle collars may use 1080 against predators like coyotes. Sources [ ] Citations [ ]. • Meija, J.; et al.

88 (3): 265–91.. • ^, p. 4.121. •, pp. 10.137–10.138. •, pp. 442–455. •, pp. 76, 804.

•, pp. 64–78. •, pp. 752, 754. •, p. 030001-27. •, pp. 030001–24. •, pp. 140, 145.

•, footnotes and commentary, pp. Xxx, 38, 409, 430, 461, 608.

•, preface, pp.. • ^, pp. 3–10. •, pp. 789–791. • ^, pp. 156–157.

• ^, pp. 420–422. •, pp. 89–97. •, pp. 4.60, 4.76, 4.92, 4.96. •, pp. 4.72, 4.91, 4.93.

• ^, pp. 561–563. •, pp. 256–277. •, pp. 355–356. •, (various pages, by metal in respective chapters). •, pp. 4.71, 4.78, 4.92. •, pp. 184–185.

•, pp. 812–816. •, pp. 328–329. •, pp. 638–640, 683–689, 767–778. •, pp. 435–436.

•, pp. 828–830. •, pp. 392–393. •, p. 395–397, 400. •, pp. 451–452. • ^, pp. 3–4.

•, pp. 384–285. •, pp. 796–797. •, pp. 384–385. • ^, pp. 390–391. •, pp. 391–392. •, pp. 41, 50.

•, pp. 538, 543–547. •; see for a summary. •, pp. 4–6, 41, 46–47. •, pp. 212–213.

Indexed references [ ].