Fluorine (pronounced /ˈflʊəriːn/, /ˈflʊərɨn/, or /ˈflɔr-/) is the chemical element A chemical element is a pure chemical substance consisting of one type of atom distinguished by its atomic number, which is the number of protons in its nucleus. The term is also used to refer to a pure chemical substance composed of atoms with the same number of protons. Common examples of elements are iron, copper, silver, gold, hydrogen, carbon, with atomic number In chemistry and physics, the atomic number is the number of protons found in the nucleus of an atom and therefore identical to the charge number of the nucleus. It is conventionally represented by the symbol Z. The atomic number uniquely identifies a chemical element. In an atom of neutral charge, the atomic number is also equal to the number of 9, represented by the symbol F. Fluorine forms a single bond with itself in elemental form, resulting in the diatomic F₂ molecule. F₂ is a supremely reactive, poisonous In the context of biology, poisons are substances that can cause disturbances to organisms, usually by chemical reaction or other activity on the molecular scale, when a sufficient quantity is absorbed by an organism. Legally and in hazardous chemical labeling, poisons are especially toxic substances; less toxic substances are labeled ", pale, yellowish brown gas. Elemental fluorine is the most chemically reactive and electronegative Electronegativity, symbol χ , is a chemical property that describes the ability of an atom (or, more rarely, a functional group) to attract electrons (or electron density) towards itself. An atom's electronegativity is affected by both its atomic weight and the distance that its valence electrons reside from the charged nucleus. The higher the of all the elements. For example, it will readily "burn" hydrocarbons In organic chemistry, a hydrocarbon is an organic compound consisting entirely of hydrogen and carbon. Hydrocarbons from which one hydrogen atom has been removed are functional groups, called hydrocarbyls. Aromatic hydrocarbons , alkanes, alkenes, cycloalkanes and alkyne-based compounds are different types of hydrocarbons at room temperature, in contrast to the combustion Combustion or burning is the sequence of exothermic chemical reactions between a fuel and an oxidant accompanied by the production of heat and conversion of chemical species. The release of heat can result in the production of light in the form of either glowing or a flame. Fuels of interest often include organic compounds in the gas, liquid or of hydrocarbons by oxygen Oxygen (pronounced /ˈɒksɨdʒɨn/, OK-si-jin, from the Greek roots ὀξύς (acid, literally "sharp", from the taste of acids) and -γενής (-genēs) (producer, literally begetter), is the element with atomic number 8 and represented by the symbol O. It is a member of the chalcogen group on the periodic table, and is a highly, which requires an input of energy In chemistry, activation energy is a term introduced in 1889 by the Swedish scientist Svante Arrhenius, that is defined as the energy that must be overcome in order for a chemical reaction to occur. Activation energy may also be defined as the minimum energy required to start a chemical reaction. The activation energy of a reaction is usually with a spark. Therefore, molecular fluorine is highly dangerous, more so than other halogens The halogens or halogen elements are a series of nonmetal elements from Group 17 IUPAC Style of the periodic table, comprising fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). The artificially created element 117, provisionally referred to by the systematic name ununseptium, may also be a halogen such as the poisonous chlorine Chlorine (pronounced /ˈklɔəriːn/ KLOR-een, from the Greek word 'χλωρóς' , is the chemical element with atomic number 17 and symbol Cl. It is a halogen, found in the periodic table in group 17 (formerly VII, VIIa, or VIIb). As the chloride ion, which is part of common salt and other compounds, it is abundant in nature and necessary to gas.

Fluorine's highest electronegativity and small atomic radius give unique properties to many of its compounds. For example, the enrichment of ²³⁵U Uranium is a silvery-white metallic chemical element in the actinide series of the periodic table with atomic number 92. It is assigned the chemical symbol U. A uranium atom has 92 protons and 92 electrons, in which 6 of the electrons are valence electrons. The uranium nucleus binds between 141 and 146 neutrons, establishing six isotopes, the most, the principal nuclear fuel, relies on the volatility of UF₆ Uranium hexafluoride , referred to as "hex" in the nuclear industry, is a compound used in the uranium enrichment process that produces fuel for nuclear reactors and nuclear weapons. It forms solid grey crystals at standard temperature and pressure (STP), is highly toxic, reacts violently with water and is corrosive to most metals. It. Also, the carbon–fluorine bond is one of the strongest bonds in organic chemistry. This contributes to the stability and persistence of fluoroalkane Fluorocarbons, sometimes referred to as perfluorocarbons, are organofluorine compounds that contain only carbon and fluorine bonded together in strong carbon–fluorine bonds. Fluoroalkanes, that contain only single bonds, are more chemically and thermally stable than alkanes. However, fluorocarbons with double bonds and especially triple bonds ( based organofluorine compounds Organofluorine compounds are organic chemical compounds that contain carbon and fluorine bonded in the polarized and remarkably strong carbon–fluorine bond. Organofluorine compounds are diverse, they can be fluorocarbons, perfluorinated, or biologically synthesized mono-fluorinated compounds, among other possibilities. These compounds have a, such as PTFE In chemistry, polytetrafluoroethylene is a synthetic fluoropolymer of tetrafluoroethylene that finds numerous applications. PTFE is most well known by the DuPont brand name Teflon/(Teflon) and PFOS Perfluorooctanesulfonic acid , or perfluorooctane sulfonate, is a man-made fluorosurfactant and global pollutant. It was added to Annex B of the Stockholm Convention on Persistent Organic Pollutants in May 2009. PFOS can form from the degradation of precursors in addition to industrial production. The PFOS levels that have been detected in. The carbon–fluorine bond's inductive effects result in the strength of many fluorinated acids, such as triflic acid Trifluoromethanesulfonic acid, also known as triflic acid, HOTf or TfOH, is a sulfonic acid with the chemical formula CF3SO3H . It is often regarded as one of the strongest acids, and is one of a number of so-called "superacids". Triflic acid is widely used especially as a catalyst and a precursor in organic chemistry and trifluoroacetic acid Trifluoroacetic acid is the simplest perfluorinated carboxylic acid chemical compound with the formula CF3CO2H. It is a strong carboxylic acid due to the influence of the electronegative trifluoromethyl group. TFA is almost 100,000-fold more acidic than acetic acid. TFA is widely used in organic chemistry. Drugs are often fluorinated at biologically reactive positions, to prevent their metabolism and prolong their half-lives.

Characteristics

F₂ is a corrosive pale yellow or brown^[1] gas Gas is one of three classical states of matter. Near absolute zero, a substance exists as a solid. As heat is added to this substance it melts into a liquid at its melting point , boils into a gas at its boiling point, and if heated high enough would enter a plasma state in which the electrons are so energized that they leave their parent atoms that is a powerful oxidizing Redox describes all chemical reactions in which atoms have their oxidation number (oxidation state) changed. This can be either a simple redox process, such as the oxidation of carbon to yield carbon dioxide (CO2) or the reduction of carbon by hydrogen to yield methane (CH4), or a complex process such as the oxidation of sugar(C6H12O6) in the agent. It is the most reactive and most electronegative of all the elements on the classic Pauling scale Electronegativity, symbol χ, is a chemical property that describes the ability of an atom to attract electrons (or electron density) towards itself in a covalent bond. An atom's electronegativity is affected by both its atomic weight and the distance that its valence electrons reside from the charged nucleus. The higher the associated (4.0), and readily forms compounds with most other elements. It is found in the -1 oxidation state In chemistry, the oxidation state is an indicator of the degree of oxidation of an atom in a chemical compound. The formal oxidation state is the hypothetical charge that an atom would have if all bonds to atoms of different elements were 100% ionic. Oxidation states are typically represented by integers, which can be positive, negative, or zero, except when bonded to another fluorine in F₂ which gives it an oxidation number of 0. Fluorine combines with the noble gases The noble gases are a group of chemical elements with very similar properties: under standard conditions, they are all odorless, colorless, monatomic gases, with very low chemical reactivity. The six noble gases that occur naturally are helium , neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and the radioactive radon (Rn) argon Argon is a chemical element represented by the symbol Ar. Argon has atomic number 18 and is the third element in group 18 of the periodic table (noble gases). Argon is the third most common gas in the Earth's atmosphere, at 0.93%, making it more common than carbon dioxide. The complete octet (eight electrons) in the outer atomic shell makes it, krypton Krypton is a chemical element with the symbol Kr and atomic number 36. It is a member of Group 18 and Period 4 elements. A colorless, odorless, tasteless noble gas, krypton occurs in trace amounts in the atmosphere, is isolated by fractionally distilling liquified air, and is often used with other rare gases in fluorescent lamps. Krypton is inert, xenon Xenon is a chemical element represented by the symbol Xe. The element name is pronounced /ˈzɛnɒn/, ZEN-on or /ˈziːnɒn/, ZEE-non. Its atomic number is 54. A colorless, heavy, odorless noble gas, xenon occurs in the Earth's atmosphere in trace amounts. Although generally unreactive, xenon can undergo a few chemical reactions such as the, and radon Radon is formed as part of the normal radioactive decay chain of uranium. Uranium has been around since the earth was formed and its most common isotope has a very long half-life , which is the amount of time required for one-half of uranium to break down. Uranium, radium, and thus radon, will continue to occur for millions of years at about the. Even in dark, cool conditions, fluorine reacts explosively with hydrogen Hydrogen is the chemical element with atomic number 1. It is represented by the symbol H. With an average atomic weight of 1.00794 u (1.007825 u for Hydrogen-1), hydrogen is the lightest and most abundant chemical element, constituting roughly 75 % of the Universe's elemental mass. Stars in the main sequence are mainly composed of hydrogen in its. The reaction with hydrogen can occur at extremely low temperatures, using liquid hydrogen and solid fluorine. It is so reactive that metals A metal is a chemical element that is a good conductor of both electricity and heat and forms cations and ionic bonds with non-metals. In chemistry, a metal is an element, compound, or alloy characterized by high electrical conductivity. In a metal, atoms readily lose electrons to form positive ions (cations). Those ions are surrounded by, water Water is a chemical substance with the chemical formula H2O. Its molecule contains one oxygen and two hydrogen atoms connected by covalent bonds. Water is a liquid at ambient conditions, but it often co-exists on Earth with its solid state, ice, and gaseous state, water vapor or steam, as well as most other substances In chemistry, a chemical substance is a material with a specific chemical composition, burn with a bright flame in a jet of fluorine gas. In moist air, it reacts with water to form the also dangerous hydrofluoric acid Hydrofluoric acid is a solution of hydrogen fluoride in water. While it is extremely corrosive and difficult to handle, it is technically a weak acid. Hydrogen fluoride, often in the aqueous form as hydrofluoric acid, is a valued source of fluorine, being the precursor to numerous pharmaceuticals such as fluoxetine (Prozac), diverse polymers such.

Fluorides Fluoride is the anion F−, the reduced form of fluorine. Both organic and inorganic compounds containing the element fluorine are sometimes called fluorides. Fluoride, like other halides, is a monovalent ion . Its compounds often have properties that are distinct relative to other halides. Structurally, and to some extent chemically, the fluoride are compounds that combine fluorine with some positively charged counterpart. They often consist of crystalline ionic salts. Fluorine compounds with metals are among the most stable of salts.

Hydrogen fluoride Hydrogen fluoride is a chemical compound with the formula HF. It is the principal industrial source of fluorine, often in the aqueous form as hydrofluoric acid, and thus is the precursor to many important compounds including pharmaceuticals and polymers . HF is widely used in the petrochemical industry and a component of many superacids. HF boils is a weak acid A weak acid is an acid that dissociates incompletely. It does not release all of its hydrogens in a solution, donating only a partial amount of its protons to the solution. These acids have higher pKa than strong acids, which release all of their hydrogen atoms when dissolved in water when dissolved in water, but is still very corrosive and attacks glass. Consequently, fluorides of alkali metals The alkali metals are a series of chemical elements forming Group 1 of the periodic table: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). (Hydrogen, although nominally also a member of Group 1, very rarely exhibits behavior comparable to the alkali metals). The alkali metals provide one of the best produce basic solutions. For example, a 1 M solution of NaF in water has a pH of 8.59 compared to a 1 M solution of NaOH, a strong base, which has a pH of 14.00.^[2]

Stellar nucleosynthesis

Fluorine is relatively rare because the solar temperatures needed to make it also enable it to quickly fuse with hydrogen Hydrogen is the chemical element with atomic number 1. It is represented by the symbol H. With an average atomic weight of 1.00794 u (1.007825 u for Hydrogen-1), hydrogen is the lightest and most abundant chemical element, constituting roughly 75 % of the Universe's elemental mass. Stars in the main sequence are mainly composed of hydrogen in its to form oxygen Oxygen (pronounced /ˈɒksɨdʒɨn/, OK-si-jin, from the Greek roots ὀξύς (acid, literally "sharp", from the taste of acids) and -γενής (-genēs) (producer, literally begetter), is the element with atomic number 8 and represented by the symbol O. It is a member of the chalcogen group on the periodic table, and is a highly and helium Helium is the chemical element with atomic number 2 and an atomic weight of 4.002602, which is represented by the symbol He. It is a colorless, odorless, tasteless, non-toxic, inert monatomic gas that heads the noble gas group in the periodic table. Its boiling and melting points are the lowest among the elements and it exists only as a gas except , or with helium to become neon Neon is the chemical element that has the symbol Ne and an atomic number of 10. Although a very common element in the universe, it is rare on Earth. A colorless, inert noble gas under standard conditions, neon gives a distinct reddish-orange glow when used in discharge tubes and neon lamps and advertising signs. It is commercially extracted from. Most fluorine is created either during a supernova when a neutrino A neutrino is an elementary particle that usually travels close to the speed of light, is electrically neutral, and is able to pass through ordinary matter almost undisturbed. This makes neutrinos extremely difficult to detect. Neutrinos have a very small, but nonzero mass. They are denoted by the Greek letter ν (nu) hits an atom of neon, or when a blue Wolf-Rayet star massing over 40 solar masses has a stellar wind blowing the fluorine out of the star before hydrogen or helium can destroy it.^[3]

Isotopes

Although fluorine has multiple isotopes Isotopes are different types of atoms of the same chemical element, each having a different number of neutrons. In a corresponding manner, isotopes differ in mass number (or number of nucleons) but never in atomic number. The number of protons (the atomic number) is the same because that is what characterizes a chemical element. For example,, only one of these isotopes (¹⁹F) is stable, and the others have short half-lives and are not found in nature. Fluorine is thus a mononuclidic element.

The nuclide ¹⁸F Fluorine-18 is a fluorine radioisotope which is an important source of positrons. It has a mass of 18.0009380 u and its half-life is 109.771(20) minutes is the radionuclide of fluorine with the longest half life (about 110 minutes), and commercially is an important source of positrons The positron or antielectron is the antiparticle or the antimatter counterpart of the electron. The positron has an electric charge of +1e, a spin of 1⁄2, and the same mass as an electron. When a low-energy positron collides with a low-energy electron, annihilation occurs, resulting in the production of two or more gamma ray photons, finding its major use in positron emission tomography Positron emission tomography is a nuclear medicine imaging technique which produces a three-dimensional image or picture of functional processes in the body. The system detects pairs of gamma rays emitted indirectly by a positron-emitting radionuclide (tracer), which is introduced into the body on a biologically active molecule. Images of tracer scanning.

Applications

Elemental fluorine, F₂, is mainly used for the production of two compounds of commercial interest, uranium hexafluoride Uranium hexafluoride , referred to as "hex" in the nuclear industry, is a compound used in the uranium enrichment process that produces fuel for nuclear reactors and nuclear weapons. It forms solid grey crystals at standard temperature and pressure (STP), is highly toxic, reacts violently with water and is corrosive to most metals. It and sulfur hexafluoride Sulfur hexafluoride is an inorganic, colorless, odorless, non-toxic and non-flammable gas (under standard conditions). SF6 has an octahedral geometry, consisting of six fluorine atoms attached to a central sulfur atom. It is a hypervalent molecule. Typical for a nonpolar gas, it is poorly soluble in water but soluble in nonpolar organic solvents.^[4]

Industrial use of fluorine-containing compounds

• Atomic The name Atom applies to a pair of related standards. The Atom Syndication Format is an XML language used for web feeds, while the Atom Publishing Protocol is a simple HTTP-based protocol for creating and updating web resources fluorine and molecular fluorine are used for plasma etching Plasma etching is a form of plasma processing used to fabricate integrated circuits. It involves a high-speed stream of glow discharge of an appropriate gas mixture being shot (in pulses) at a sample. The plasma source, known as etch species, can be either charged (ions) or neutral (atoms and radicals). During the process, the plasma will generate in semiconductor manufacturing, flat panel display production and MEMS (microelectromechanical systems) fabrication.^[5] Xenon difluoride is also used for this last purpose.
• Hydrofluoric acid (chemical formula HF) is used to etch glass in light bulbs and other products.
• Tetrafluoroethylene and perfluorooctanoic acid (PFOA) are directly used in the production of low friction plastics such as Teflon (or polytetrafluoroethylene).
• Fluorine is used indirectly in the production of halons such as freon.
• Along with some of its compounds, fluorine is used in the production of pure uranium from uranium hexafluoride and in the synthesis of numerous commercial fluorochemicals, including vitally important pharmaceuticals, agrochemical compounds, lubricants, and textiles.
• Chlorofluorocarbons (CFCs) are used extensively in air conditioning and in refrigeration. CFCs have been banned for these applications because they contribute to ozone destruction and the ozone hole. It is interesting to note that, since it is chlorine and bromine radicals that harm the ozone layer, not fluorine, compounds that do not contain chlorine or bromine but contain only fluorine, carbon, and hydrogen (called hydrofluorocarbons) are not on the United States Environmental Protection Agency list of ozone-depleting substances,^[6] and have been widely used as replacements for the chlorine- and bromine-containing fluorocarbons. Hydrofluorocarbons do have a greenhouse effect, but a small one compared with carbon dioxide and methane.
• Sodium hexafluoroaluminate (cryolite), is used in the electrolysis of aluminium.
• In much higher concentrations, sodium fluoride has been used as an insecticide, especially against cockroaches.
• Fluorides have been used in the past to help molten metal flow, hence the name, which derives from Latin verb fluere, meaning to flow.^[7]
• Some researchers including US space scientists in the early 1960s have studied elemental fluorine gas as a possible rocket propellant due to its exceptionally high specific impulse. The experiments failed because fluorine proved difficult to handle, and its combustion products proved extremely toxic and corrosive.
• Compounds of fluorine such as fluoropolymers, potassium fluoride, and cryolite are utilized in applications such as anti-reflective coatings and dichroic mirrors on account of their unusually low refractive index.

Dental and medical uses

• Inorganic compounds of fluoride, including sodium fluoride (NaF), stannous fluoride (SnF₂) and sodium MFP, are used in toothpaste to prevent dental cavities. These or related compounds are also added to some municipal water supplies, a process called water fluoridation, although the practice has remained controversial since its beginnings in 1945.
• Many important agents for general anesthesia such as sevoflurane, desflurane, and isoflurane are hydrofluorocarbon derivatives.
• The fluorinated antiinflammatories dexamethasone and triamcinolone are among the most potent of the synthetic corticosteroids class of drugs.^[8]
• Fludrocortisone ("Florinef") is one of the most common mineralocorticoids, a class of drugs that mimics the actions of aldosterone.
• Fluconazole is a triazole antifungal drug used in the treatment and prevention of superficial and systemic fungal infections.
• Fluoroquinolones are a family of broad-spectrum antibiotics.
• SSRI antidepressants, except in a few instances, are fluorinated molecules. These include citalopram, escitalopram, fluoxetine, fluvoxamine, and paroxetine. A notable exception is sertraline. Because of the difficulty of biological systems in dealing with metabolism of fluorinated molecules, fluorinated antibiotics and antidepressants are among the major fluorinated organics found in treated city sewage and wastewater.

• Compounds containing ¹⁸F, a radioactive isotope that emits positrons, are often used in positron emission tomography, because its half-life of 110 minutes is long by the standards of positron-emitters. One such species is fluorodeoxyglucose.

Chemistry of fluorine

Fluorine forms a variety of very different compounds, owing to its small atomic size and covalent behavior. Elemental fluorine is a dangerously powerful oxidant, reflecting the extreme electronegativity of fluorine. Hydrofluoric acid is extremely dangerous, whereas, in synthetic drugs incorporating an aromatic ring (e.g., flumazenil), fluorine is used to help prevent toxication or to delay metabolism^{[citation needed]}.

The fluoride ion is basic, therefore hydrofluoric acid is a weak acid in water solution. However, water is not an inert solvent in this case: When less basic solvents such as anhydrous acetic acid are used, hydrofluoric acid is the strongest of the hydrohalogenic acids. Also, owing to the basicity of the fluoride ion, soluble fluorides give basic water solutions. The fluoride ion is a Lewis base, and has a high affinity to certain elements such as calcium and silicon. For example, deprotection of silicon protecting groups is achieved with a fluoride. The fluoride ion is poisonous.

Fluorine as a freely reacting oxidant gives the strongest oxidants known.

The reactivity of fluorine toward the noble gas xenon was first reported by Neil Bartlett in 1962. Fluorides of krypton and radon have also been prepared. Argon fluorohydride has been observed at cryogenic temperatures.

The carbon-fluoride bond is covalent and very stable. The use of a fluorocarbon polymer, poly(tetrafluoroethene) or Teflon, is an example: It is thermostable and waterproof enough to be used in frying pans. Organofluorines may be safely used in applications such as drugs, without the risk of release of toxic fluoride. In synthetic drugs, toxication can be prevented. For example, an aromatic ring is useful but presents a safety problem: enzymes in the body metabolize some of them into poisonous epoxides. When the para position is substituted with fluorine, the aromatic ring is protected and epoxide is no longer produced.

The substitution of fluorine for hydrogen in organic compounds offers a very large number of compounds. An estimated fifth of pharmaceutical compounds and 30% of agrochemical compounds contain fluorine.^[9] The -CF₃ and -OCF₃ moieties provide further variation, and more recently the -SF₅ group.^[10]

Production

Industrial production of fluorine entails the electrolysis of hydrogen fluoride in the presence of potassium fluoride. This method is based on the pioneering studies by Moissan (see below). Fluorine gas forms at the anode, and hydrogen gas at the cathode. Under these conditions, the potassium fluoride (KF) converts to potassium bifluoride (KHF₂), which is the actual electrolyte. This potassium bifluoride aids electrolysis by greatly increasing the electrical conductivity of the solution.

HF + KF → KHF₂

2 KHF₂ → 2 KF + H₂ + F₂

The HF required for the electrolysis is obtained as a byproduct of the production of phosphoric acid. Phosphate-containing minerals contain significant amounts of calcium fluorides, such as fluorite. Upon treatment with sulfuric acid, these minerals release hydrogen fluoride:

CaF₂ + H₂SO₄ → 2 HF + CaSO₄

In 1986, when preparing for a conference to celebrate the 100th anniversary of the discovery of fluorine, Karl Christe discovered a purely chemical preparation involving the reaction of solutions in anhydrous HF, K₂MnF₆, and SbF₅ at 150 °C:^[11]

2 K₂MnF₆ + 4 SbF₅ → 4 KSbF₆ + 2 MnF₃ + F₂

Though not a practical synthesis on the large scale, this report demonstrates that electrolysis is not the sole route to the element.

History

The mineral fluorspar (also called fluorite), consisting mainly of calcium fluoride, was described in 1530 by Georgius Agricola for its use as a flux.^[12] Fluxes are used to promote the fusion of metals or minerals. The etymology of the element's name reflects its history: Fluorine is pronounced /ˈflʊəriːn/, /ˈflʊərɨn/, or commonly /ˈflɔr-/; from Latin: fluere, meaning "to flow". In 1670 Schwanhard found that glass was etched when it was exposed to fluorspar that had been treated with acid. Carl Wilhelm Scheele and many later researchers, including Humphry Davy, Caroline Menard, Gay-Lussac, Antoine Lavoisier, and Louis Thenard all would experiment with hydrofluoric acid, easily obtained by treating fluorite with concentrated sulfuric acid.

Owing to its extreme reactivity, elemental fluorine was not isolated until many years after the characterization of fluorite. Progress in isolating elemental fluorine was slowed because it could be prepared only electrolytically and even then under stringent conditions, since the gas attacks many materials. In 1886, the isolation of elemental fluorine was reported by Henri Moissan after almost 74 years of effort by other chemists.^[13] The generation of elemental fluorine from hydrofluoric acid proved to be exceptionally dangerous, killing or blinding several scientists who attempted early experiments on this halogen. These individuals came to be referred to as "fluorine martyrs".^[14] For Moissan, it earned him the 1906 Nobel Prize in chemistry.^[15]

The first large-scale production of fluorine was undertaken in support of the Manhattan project, where the compound uranium hexafluoride (UF₆) had been selected as the form of uranium that would allow separation of its ²³⁵U and ²³⁸U isotopes. Today both the gaseous diffusion process and the gas centrifuge process use gaseous UF₆ to produce enriched uranium for nuclear power applications. In the Manhattan Project, it was found that UF₆ decomposed into UF₄ and F₂. The corrosion problem due to the F₂ was eventually solved by electrolytically coating all UF₆ carrying piping with nickel metal, which forms a nickel difluoride that is not attacked by fluorine. Joints and flexible parts were made from teflon, then a very recently discovered fluorocarbon plastic that is also not attacked by F₂.

Biological role

Though F₂ is too reactive to have any natural biological role, fluorine is incorporated into compounds with biological activity. Naturally occurring organofluorine compounds are rare, the most notable example is fluoroacetate, which functions as a plant defence against herbivores in at least 40 plants in Australia, Brazil, and Africa.^[16] The enzyme adenosyl-fluoride synthase catalyzes the formation of 5'-deoxy-5'-fluoroadenosine. Fluorine is not an essential nutrient, but its importance in preventing tooth decay is well-recognized.^[17] The effect is predominantly topical, although prior to 1981 it was considered primarily systemic (occurring through ingestion).^[18]

Precautions

Elemental fluorine

Elemental fluorine (fluorine gas) is a highly toxic, corrosive oxidant, which can cause ignition of organic material. Fluorine gas has a characteristic pungent odor that is detectable in concentrations as low as 20 ppb. As it is so reactive, all materials of construction must be carefully selected and metal surfaces must be passivated.

Fluoride ion

Fluoride ions are toxic: the lethal dose of sodium fluoride for a 70 kg human is estimated at 5–10 g.^[19]

Hydrogen fluoride and hydrofluoric acid

Hydrogen fluoride and hydrofluoric acid are dangerous, far more so than the related hydrochloric acid, because undissociated molecular HF penetrates the skin and biological membranes, causing deep and painless burns. The free fluoride, once released from HF in dissociation, also is capable of chelating calcium ion to the point of causing death by cardiac dysrhythmia. Burns with areas larger than 25 square inches (160 cm²) have the potential to cause serious systemic toxicity.^[20]

Organofluorines

Organofluorines are naturally rare compounds. They can be nontoxic (perflubron and perfluorodecalin) or highly toxic (perfluoroisobutylene and fluoroacetic acid). Many pharmacueticals are organofluorines, such as the anti-cancer fluorouracil. Perfluorooctanesulfonic acid (PFOS) is a persistent organic pollutant.

See also

• Halide minerals

References

1. ^ Theodore Gray. "Real visible fluorine". The Wooden Periodic Table. http://theodoregray.com/PeriodicTable/Samples/009.5/index.s12.html.
2. ^ "pKa's of Inorganic and Oxo-Acids". Evans Group. http://www2.lsdiv.harvard.edu/labs/evans/pdf/evans_pKa_table.pdf. Retrieved 2008-11-29.
3. ^ Ken Croswell (September 2003). "Fluorine: An Element-ary Mystery". Sky and Telescope. "http://kencroswell.com/fluorine.html"
4. ^ M. Jaccaud, R. Faron, D. Devilliers, R. Romano (2005). Fluorine, in Ullmann’s Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. ISBN 3527310975.
5. ^ Leonel R Arana, Nuria de Mas, Raymond Schmidt, Aleksander J Franz, Martin A Schmidt and Klavs F Jensen (2007). "Isotropic etching of silicon in fluorine gas for MEMS micromachining". J. Micromech. Microeng. 17: 384. doi:10.1088/0960-1317/17/2/026.
6. ^ "Class I Ozone-Depleting Substances". Ozone Depletion. U.S. Environmental Protection Agency. http://www.epa.gov/ozone/ods.html.
7. ^ compiled by Alexander Senning. (2007). Elsevier's dictionary of chemoetymology : the whies and whences of chemical nomenclature and terminology. Amsterdam: Elsevier. p. 149. ISBN 9780444522399. http://books.google.com/?id=Fl4sdCYrq3cC&pg=PT158.
8. ^ Steve S Lim. "eMedicine - Corticosteroid-Induced Myopathy". http://www.emedicine.com/pmr/topic35.htm.
9. ^ "Fluorine's treasure trove". ICIS news. 2006-10-02. http://www.icis.com/Articles/2006/09/30/2016413/fluorines-treasure-trove.html. Retrieved 2008-11-29.
10. ^ Bernhard Stump, Christian Eberle, W. Bernd Schweizer, Marcel Kaiser, Reto Brun, R. Luise Krauth-Siegel, Dieter Lentz, François Diederich (2009). "Pentafluorosulfanyl as a Novel Building Block for Enzyme Inhibitors: Trypanothione Reductase Inhibition and Antiprotozoal Activities of Diarylamines". ChemBioChem 10 (1): 79. doi:10.1002/cbic.200800565. PMID 19058274.
11. ^ K. Christe (1986). "Chemical synthesis of elemental fluorine". Inorg. Chem. 25: 3721–3724. doi:10.1021/ic00241a001.
12. ^ "Discovery of fluorine". Fluoride History. http://www.fluoride-history.de/fluorine.htm.
13. ^ H. Moissan (1886). "Action d'un courant électrique sur l'acide fluorhydrique anhydre". Comptes rendus hebdomadaires des séances de l'Académie des sciences 102: 1543–1544. http://gallica.bnf.fr/ark:/12148/bpt6k3058f/f1541.chemindefer.
14. ^ Richard D. Duncan. (2008). Elements of faith : faith facts and learning lessons from the periodic table. Green Forest, Ark.: Master Books. p. 22. ISBN 9780890515471. http://books.google.com/?id=kgVAlzGXx6oC.
15. ^ "The Nobel Prize in Chemistry 1906". Nobelprize.org. http://nobelprize.org/nobel_prizes/chemistry/laureates/1906/. Retrieved 2009-07-07.
16. ^ Proudfoot AT, Bradberry SM, Vale JA (2006). "Sodium fluoroacetate poisoning". Toxicol Rev 25 (4): 213–9. doi:10.2165/00139709-200625040-00002. PMID 17288493.
17. ^ Olivares M and Uauy R (2004). "Essential nutrients in drinking-water (Draft)". WHO. http://www.who.int/water_sanitation_health/dwq/en/nutoverview.pdf. Retrieved 2008-12-30.
18. ^ Pizzo G, Piscopo MR, Pizzo I, Giuliana G (September 2007). "Community water fluoridation and caries prevention: a critical review". Clin Oral Investig 11 (3): 189–93. doi:10.1007/s00784-007-0111-6. PMID 17333303.
19. ^ Aigueperse, Jean; Paul Mollard, Didier Devilliers, Marius Chemla, Robert Faron, Renée Romano, Jean Pierre Cuer (2005). "Fluorine Compounds, Inorganic". in Ullmann. Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH.
20. ^ "Recommended Medical Treatment for Hydrofluoric Acid Exposure" (PDF). Honeywell Specialty Materials. http://www51.honeywell.com/sm/hfacid/common/documents/HF_medical_book.pdf. Retrieved 2009-05-06.

External links

• The Periodic Table of Videos video of Fluorine at YouTube
• WebElements.com – Fluorine
• It's Elemental – Fluorine
• Picture of liquid fluorine – chemie-master.de
• Chemsoc.org
• Fluorine: An Element-ary Mystery by Ken Croswell, published in Sky and Telescope September 2003

Categories: Fluorine | Biology and pharmacology of chemical elements | Chemical elements | Fluorinating agents | Halogens | Oxidizing agents

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