Fluorine
Allotropes of fluorine) | |||||||||||||||||||||
Appearance | gas: very pale yellow liquid: bright yellow solid: alpha is opaque, beta is transparent | ||||||||||||||||||||
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Standard atomic weight Ar°(F) | |||||||||||||||||||||
Fluorine in the periodic table | |||||||||||||||||||||
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Atomic number (Z) | 9 | ||||||||||||||||||||
Group | group 17 (halogens) | ||||||||||||||||||||
Period | period 2 | ||||||||||||||||||||
Block | p-block | ||||||||||||||||||||
Electron configuration | [He] 2s2 2p5[3] | ||||||||||||||||||||
Electrons per shell | 2, 7 | ||||||||||||||||||||
Physical properties | |||||||||||||||||||||
Phase at STP | gas | ||||||||||||||||||||
Melting point | (F2) 53.48 K (−219.67 °C, −363.41 °F)[4] | ||||||||||||||||||||
Boiling point | (F2) 85.03 K (−188.11 °C, −306.60 °F)[4] | ||||||||||||||||||||
Density (at STP) | 1.696 g/L[5] | ||||||||||||||||||||
when liquid (at b.p.) | 1.505 g/cm3[6] | ||||||||||||||||||||
Triple point | 53.48 K, .252 kPa[7] | ||||||||||||||||||||
Critical point | 144.41 K, 5.1724 MPa[4] | ||||||||||||||||||||
Heat of vaporization | 6.51 kJ/mol[5] | ||||||||||||||||||||
Molar heat capacity | Cp: 31 J/(mol·K)[6] (at 21.1 °C) Cv: 23 J/(mol·K)[6] (at 21.1 °C) | ||||||||||||||||||||
Vapor pressure
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Atomic properties | |||||||||||||||||||||
Discovery | André-Marie Ampère (1810) | ||||||||||||||||||||
First isolation | Henri Moissan[3] (June 26, 1886) | ||||||||||||||||||||
Named by | |||||||||||||||||||||
Isotopes of fluorine | |||||||||||||||||||||
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Fluorine is a chemical element; it has symbol F and atomic number 9. It is the lightest halogen[note 1] and exists at standard conditions as a highly toxic, pale yellow diatomic gas. Fluorine is extremely reactive, as it reacts with all other elements except for the light inert gases.
Among the elements, fluorine ranks 24th in universal abundance and 13th in terrestrial 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 corrosive hydrogen fluoride en route to various organic fluorides, or into cryolite, which plays a key role in aluminium refining. Molecules containing a carbon–fluorine bond often have very high chemical and thermal stability; their major uses are as refrigerants, electrical insulation and cookware, 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$15 billion a year.
Fluorocarbon gases are generally greenhouse gases with global-warming potentials 100 to 23,500 times that of carbon dioxide, and SF6 has the highest global warming potential of any known substance. Organofluorine compounds often persist in the environment due to the strength of the carbon–fluorine bond. Fluorine has no known metabolic role in mammals; a few plants and marine sponges synthesize organofluorine poisons (most often monofluoroacetates) that help deter predation.[16]
Characteristics
Electron configuration
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]
Fluorine's
Reactivity
External videos | |
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Bright flames during fluorine reactions | |
Fluorine reacting with caesium |
The
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.[22][23] 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][24]
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,[22] and noble metals require pure fluorine gas at 300–450 °C (575–850 °F).[25] Some solid nonmetals (sulfur, phosphorus) react vigorously in liquid fluorine.[26] Hydrogen sulfide[26] and sulfur dioxide[27] combine readily with fluorine, the latter sometimes explosively; sulfuric acid exhibits much less activity, requiring elevated temperatures.[28]
Phases
At room temperature, fluorine is a gas of
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
Isotopes
Only one
F, with a half-life of 162(7) nanoseconds, and 26m
F, with a half-life of 2.2(1) milliseconds.[55]
Occurrence
Universe
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 ppb (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.[57] 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.[57][58]
Beyond this transient existence, three explanations have been proposed for the presence of fluorine:[57][59]
- during type II supernovae, bombardment of neon atoms by neutrinos could transmute them to fluorine;
- the solar wind of Wolf–Rayet stars could blow fluorine away from any hydrogen or helium atoms; or
- fluorine is borne out on convection currents arising from fusion in asymptotic giant branch stars.
Earth
Fluorine is the thirteenth most common element in Earth's crust at 600–700 ppm (parts per million) by mass.[60] Though believed not to occur naturally, elemental fluorine has been shown to be present as an occlusion in antozonite, a variant of fluorite.[61] Most fluorine exists as fluoride-containing minerals. Fluorite, fluorapatite and cryolite are the most industrially significant.[60][62] 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.[62][63][64][65][66] Fluorapatite (Ca5(PO4)3F), which contains most of the world's fluoride, is an inadvertent source of fluoride as a byproduct of fertilizer production.[62] 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.[62]
-
Fluorite: Pink globular mass with crystal facets
-
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.[62] Trace quantities of organofluorines of uncertain origin have been detected in volcanic eruptions and geothermal springs.[67] The existence of gaseous fluorine in crystals, suggested by the smell of crushed antozonite, is contentious;[68][61] 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.[61]
History
Early discoveries
In 1529,
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.[78][79] Swedish chemist Carl Wilhelm Scheele repeated the experiment in 1771, and named the acidic product fluss-spats-syran (fluorspar acid).[79][80] In 1810, the French physicist André-Marie Ampère suggested that hydrogen and an element analogous to chlorine constituted hydrofluoric acid.[81] 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.[82][83] 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).[84] The New Latin name fluorum gave the element its current symbol F; Fl was used in early papers.[85][note 7]
Isolation
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.[76][86] 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.[76][86][87] 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.[86][88] In 1886, after 74 years of effort by many chemists, Moissan isolated elemental fluorine.[87][89]
In 1906, two months before his death, Moissan received the Nobel Prize in Chemistry,[90] with the following citation:[86]
[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]
Later uses
The
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.[79][91][92][93] 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.[79][91][92]
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[94] 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.[95]
Compounds
Fluorine has a rich chemistry, encompassing organic and inorganic domains. It combines with metals, nonmetals, metalloids, and most noble gases,[96] and almost exclusively assumes an oxidation state of −1.[note 10] Fluorine's high electron affinity results in a preference for ionic bonding; when it forms covalent bonds, these are polar, and almost always single.[99][100][note 11]
Metals
Alkali metals form ionic and highly soluble
Covalent bonding first comes to prominence in the
6 and ReF
6, and gaseous WF
6.[120][121][122] Rhenium heptafluoride, the only characterized metal heptafluoride, is a low-melting molecular solid with pentagonal bipyramidal molecular geometry.[123] Metal fluorides with more fluorine atoms are particularly reactive.[124]
Structural progression of metal fluorides | ||
Sodium fluoride, ionic | Bismuth pentafluoride, polymeric | Rhenium heptafluoride, molecular |
Hydrogen
Hydrogen and fluorine combine to yield hydrogen fluoride, in which discrete molecules form clusters by hydrogen bonding, resembling water more than
Other reactive nonmetals
Binary fluorides of metalloids and p-block nonmetals are generally covalent and volatile, with varying reactivities.
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. [141][142] Chlorine, bromine, and iodine can each form mono-, tri-, and pentafluorides, but only iodine heptafluoride has been characterized among possible interhalogen heptafluorides.[143] Many of them are powerful sources of fluorine atoms, and industrial applications using chlorine trifluoride require precautions similar to those using fluorine.[144][145]
Noble gases
Noble gases, having complete electron shells, defied reaction with other elements until 1962 when Neil Bartlett reported synthesis of xenon hexafluoroplatinate;[147] xenon difluoride, tetrafluoride, hexafluoride, and multiple oxyfluorides have been isolated since then.[148] Among other noble gases, krypton forms a difluoride,[149] and radon and fluorine generate a solid suspected to be radon difluoride.[150][151] Binary fluorides of lighter noble gases are exceptionally unstable: argon and hydrogen fluoride combine under extreme conditions to give argon fluorohydride.[42] Helium has no long-lived fluorides,[152] and no neon fluoride has ever been observed;[153] helium fluorohydride has been detected for milliseconds at high pressures and low temperatures.[152]
Organic compounds
The carbon–fluorine bond is organic chemistry's strongest,[155] and gives stability to organofluorines.[156] It is almost non-existent in nature, but is used in artificial compounds. Research in this area is usually driven by commercial applications;[157] the compounds involved are diverse and reflect the complexity inherent in organic chemistry.[91]
Discrete molecules
The substitution of hydrogen atoms in an
The term
Polymers
Polymers exhibit the same stability increases afforded by fluorine substitution (for hydrogen) in discrete molecules; their melting points generally increase too.
Production
Elemental fluorine and virtually all fluorine compounds are produced from
- CaF2 + H2SO4 → 2 HF(g) + CaSO4
The gaseous HF can then be absorbed in water or liquefied.[169]
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:
- H2SiF6 → 2 HF + SiF4
- SiF4 + 2 H2O → 4 HF + SiO2
Industrial routes to F2
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 8–12 volts, to generate hydrogen and fluorine gas respectively.[64][170] 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.[79][171][172] Fluorine can be stored in steel cylinders that have passivated interiors, at temperatures below 200 °C (392 °F); otherwise nickel can be used.[79][173] Regulator valves and pipework are made of nickel, the latter possibly using Monel instead.[174] 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;[174] some sources instead recommend nickel-Monel-PTFE systems.[175]
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:[176]
Christe later commented that the reactants "had been known for more than 100 years and even Moissan could have come up with this scheme."[177] As late as 2008, some references still asserted that fluorine was too reactive for any chemical isolation.[178]
Industrial applications
Fluorite mining, which supplies most global fluorine, peaked in 1989 when 5.6 million
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 vertical integration.[184]
The largest application of fluorine gas, consuming up to 7,000 metric tons annually, is in the preparation of UF
6 for the
3, BrF
3, or IF
5, which together allow calibrated fluorination. Fluorinated pharmaceuticals use sulfur tetrafluoride instead.[64]
Inorganic fluorides
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
Organic fluorides
Organofluorides consume over 20% of mined fluorite and over 40% of hydrofluoric acid, with
Refrigerant gases
Halogenated refrigerants, termed Freons in informal contexts,
Polymers
About 180,000 metric tons of fluoropolymers were produced in 2006 and 2007, generating over $3.5 billion revenue per year.[198] The global market was estimated at just under $6 billion in 2011.[199] Fluoropolymers can only be formed by polymerizing free radicals.[163]
Polytetrafluoroethylene (PTFE), sometimes called by its DuPont name Teflon,
The chemically resistant (but expensive) fluorinated
Surfactants
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; Scotchgard alone generated over $300 million in 2000.[194][206][207] 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 $100 million in 2006.[194]
Agrichemicals
About 30% of
Medicinal applications
Dental care
Population studies from the mid-20th century onwards show
Pharmaceuticals
Twenty percent of modern pharmaceuticals contain fluorine.
PET scanning
Fluorine-18 is often found in
Oxygen carriers
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.
Biological role
Fluorine is not
Natural organofluorines have been found in microorganisms, plants
Toxicity
Elemental fluorine is highly toxic to living organisms. Its effects in humans start at concentrations lower than
Hydrofluoric acid
Hazards | |
---|---|
GHS labelling: | |
Danger | |
H270, H314, H330[260] | |
NFPA 704 (fire diamond) |
Hydrofluoric acid is the weakest of the
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
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.[267] 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.[268][269][270] 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.[266][271]
Fluoride ion
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.[272] One-fifth of the lethal dose can cause adverse health effects,[273] and chronic excess consumption may lead to skeletal fluorosis, which affects millions in Asia and Africa, and, in children, to reduced intelligence.[273][274] 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.[273][275]
Historically, most cases of fluoride poisoning have been caused by accidental ingestion of insecticides containing inorganic fluorides.[276] Most current calls to poison control centers for possible fluoride poisoning come from the ingestion of fluoride-containing toothpaste.[273] Malfunctioning water fluoridation equipment is another cause: one incident in Alaska affected almost 300 people and killed one person.[277] 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.[278] 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.[276] A larger study across the U.S. had similar findings: 80% of cases involved children under six, and there were few serious cases.[279]
Environmental concerns
Atmosphere
The
Biopersistence
Organofluorines exhibit biopersistence due to the strength of the carbon–fluorine bond.
See also
- Argon fluoride laser
- Electrophilic fluorination
- Fluoride selective electrode, which measures fluoride concentration
- Fluorine absorption dating
- Fluorous chemistry, a process used to separate reagents from organic solvents
- Krypton fluoride laser
- Radical fluorination
Notes
- ^ 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.[46]
- ^ 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."[50]
- ^ Or perhaps from as early as 1670 onwards; Partington[77] and Weeks[76] give differing accounts.
- ^ Fl, since 2012, is used for flerovium.
- ^ Davy, Gay-Lussac, Thénard, and the Irish chemists Thomas and George Knox were injured. Belgian chemist Paulin Louyet and French chemist Jérôme Nicklès died. Moissan also experienced serious hydrogen fluoride poisoning.[76][86]
- ^ Also honored was his invention of the electric arc furnace.
- ^ 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;[97] F+
4 and a few related species are predicted to be stable.[98] - Hydrogen bondingis another possibility.
- ^ ZrF
4 melts at 932 °C (1710 °F),[111] HfF
4 sublimes at 968 °C (1774 °F),[108] and UF
4 melts at 1036 °C (1897 °F).[112] - ^ 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. - IUPACsynonyms 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.[160]
- ^ This DuPont trademark is sometimes further misused for CFCs, HFCs, or HCFCs.
- ^ American sheep and cattle collars may use 1080 against predators like coyotes.
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External links
- Media related to Fluorine at Wikimedia Commons