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Fluorine is the chemical element with atomic number 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, pale, yellowish brown gas. Elemental fluorine is the most chemically reactive and electronegative of all the elements. For example, it will readily "burn" hydrocarbons at room temperature, in contrast to the combustion of hydrocarbons by oxygen, which requires an input of energy with a spark. Therefore, molecular fluorine is highly dangerous, more so than other halogens such as the poisonous chlorine gas. Fluorine's highest electronegativity and small atomic radius give unique properties to many of its compounds. For example, the enrichment of 235U, the principal nuclear fuel, relies on the volatility of UF6. Also, the carbon–fluorine bond is one of the strongest bonds in organic chemistry. This contributes to the stability and persistence of fluoroalkane based organofluorine compounds, such as PTFE/(Teflon) and PFOS. The carbon–fluorine bond's inductive effects result in the strength of many fluorinated acids, such as triflic acid and trifluoroacetic acid. 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 that is a powerful oxidizing agent. It is the most reactive and most electronegative of all the elements on the classic Pauling scale (4.0), and readily forms compounds with most other elements. It is found in the -1 oxidation state, except when bonded to another fluorine in F2 which gives it an oxidation number of 0. Fluorine combines with the noble gases argon, krypton, xenon, and radon. Even in dark, cool conditions, fluorine reacts explosively with hydrogen. The reaction with hydrogen can occur at extremely low temperatures, using liquid hydrogen and solid fluorine. It is so reactive that metals, water, as well as most other substances, 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. Fluorides 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 is a weak acid when dissolved in water, but is still very corrosive and attacks glass. Consequently, fluorides of alkali metals 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] Although fluorine (F) has multiple isotopes, only one of these isotopes (F-19) is stable, and the others have short half-lives and are not found in nature. Fluorine is thus a mononuclidic element. The nuclide 18F is the radionuclide of fluorine with the longest half life (about 110 minutes), and commercially is an important source of positrons, finding its major use in positron emission tomography scanning. Elemental fluorine, F2, is mainly used for the production of two compounds of commercial interest, uranium hexafluoride and sulfur hexafluoride.[3] * Atomic fluorine and molecular fluorine are used for plasma etching in semiconductor manufacturing, flat panel display production and MEMS (microelectromechanical systems) fabrication.[4] Xenon difluoride is also used for this last purpose. 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. * 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.[8] The -CF3 and -OCF3 moieties provide further variation, and more recently the -SF5 group.[9]
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 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:[10] 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. The mineral fluorspar (also called fluorite), consisting mainly of calcium fluoride, was described in 1530 by Georgius Agricola for its use as a flux.[11] 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.[12] The generation of elemental fluorine from hydrofluoric acid is exceptionally dangerous, killing or blinding several scientists who attempted early experiments on this halogen. These individuals came to be referred to as "fluorine martyrs".[13] For Moissan, it earned him the 1906 Nobel Prize in chemistry.[14] 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. 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.[15] 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.[16] The effect is predominantly topical, although prior to 1981 it was considered primarily systemic (occurring through ingestion).[17] 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 ions are toxic: the lethal dose of sodium fluoride for a 70 kg human is estimated at 5–10 g.[18] 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.[19] 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. Fluorine chemistry at the millennium: fascinated by fluorine, Ronald Eric Banks Fluorine and Health: Molecular Imaging, Biomedical Materials and Pharmaceuticals, Günter Haufe Fluorine Chemistry Research Advances, Ira V. Gardiner
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. External links
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