Fluorine
![]() Liquid fluorine (F2 at extremely low temperature) | |||||||||||||||||||||
Fluorine | |||||||||||||||||||||
---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
Pronunciation | |||||||||||||||||||||
Allotropes | alpha, beta (see Allotropes of fluorine) | ||||||||||||||||||||
Appearance | gas: very pale yellow liquid: bright yellow solid: alpha is opaque, beta is transparent | ||||||||||||||||||||
Standard atomic weight Ar°(F) | |||||||||||||||||||||
Fluorine in the periodic table | |||||||||||||||||||||
| |||||||||||||||||||||
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
| |||||||||||||||||||||
Atomic properties | |||||||||||||||||||||
Oxidation states | −1, 0[8] (oxidizes oxygen) | ||||||||||||||||||||
Electronegativity | Pauling scale: 3.98[3] | ||||||||||||||||||||
Ionization energies | |||||||||||||||||||||
Covalent radius | 64 pm[10] | ||||||||||||||||||||
Van der Waals radius | 135 pm[11] | ||||||||||||||||||||
![]() | |||||||||||||||||||||
Other properties | |||||||||||||||||||||
Natural occurrence | primordial | ||||||||||||||||||||
Crystal structure | cubic | ||||||||||||||||||||
Thermal conductivity | 0.02591 W/(m⋅K)[12] | ||||||||||||||||||||
Magnetic ordering | diamagnetic (−1.2×10−4)[13][14] | ||||||||||||||||||||
CAS Number | 7782-41-4[3] | ||||||||||||||||||||
History | |||||||||||||||||||||
Naming | after the mineral fluorite, itself named after Latin fluo (to flow, in smelting) | ||||||||||||||||||||
Discovery | André-Marie Ampère (1810) | ||||||||||||||||||||
First isolation | Henri Moissan[3] (June 26, 1886) | ||||||||||||||||||||
Named by | |||||||||||||||||||||
Isotopes of fluorine | |||||||||||||||||||||
| |||||||||||||||||||||
Fluorine is the chemical element with atomic number 9, represented by the symbol F. It is the lightest member of the halogen column of the periodic table and has a single stable isotope, fluorine-19. At standard pressure and temperature, fluorine is a pale yellow gas composed of diatomic molecules, F
2. In stars, fluorine is relatively rare compared to other light elements. On Earth, fluorine is more common; it is the 13th most abundant element in the crust.
Fluorine's most important mineral, fluorite, was first formally described in 1530, in the context of metal smelting. The mineral's name derives from the Latin verb fluo, which means "stream" or "flow," because fluorite was added to metal ores to lower their melting points. Suggested to be a chemical element in 1811, fluorine was named after the source mineral. It was not until 1886 that elemental fluorine was obtained by French chemist Henri Moissan, whose method of electrolysis remains the only industrial production method of fluorine gas. The main use of elemental fluorine, uranium enrichment, was developed during the Manhattan Project. The vast majority, however, of commercial fluorine is never reduced to the element.
Fluorine has the highest electron affinity of any element but chlorine, and for this reason it is a very strong oxidizing agent. Fluorine forms stable compounds, fluorides, with all elements, except for helium and neon, for which the reaction has been attempted. Hydrofluoric acid, in contrast to other haloacids such as hydrochloric acid, is only a weak acid in water, but it is nonetheless extremely corrosive. Fluorides of lighter metal elements are ionic compounds (salts), which are usually water soluble. Heavier metal elements such as uranium can form volatile coordination compounds (separate molecules with several fluorine atoms surrounding a metal atom). Organic fluorine compounds tend to have high chemical and thermal stability and water-repellent properties. Several have large-scale commercial application, such as the fluorinated plastic polytetrafluoroethylene ("Teflon") used in cookware. Another major application is fluorinated refrigerants. Here, traditional chlorofluorocarbons ("Freons"), which cause ozone depletion, have been largely replaced by hydrofluorocarbons.
Although it helps prevent tooth decay, fluorine is not an essential mineral for mammals. Some organofluorine compounds are synthesized in microorganisms and plants. Several fluorine compounds, as well as elemental fluorine itself, are dangerously toxic. Nevertheless, an increasing number of pharmaceuticals (about 10% of new drugs) contain fluorine.
Characteristics
Electronic structure
A fluorine atom has nine protons and thus nine electrons, arranged in electronic configuration [He]2s22p5, one fewer than neon.[16] Fluorine's outer electrons are relatively separate from each other, and do not shield each other from the nucleus. Therefore, they experience a relatively high effective nuclear charge of +7. Fluorine tightly holds its own electrons and has an attraction for one more electron to achieve the extremely stable neon arrangement.[16]
Fluorine's first ionization energy (energy required to remove an electron to form F+) is 1,681 kilojoules per mole, which is higher than for any other element except neon and helium.[17] The second and third ionization energies of fluorine are 3,374 and 6,147 kilojoules per mole, respectively.[17] Fluorine's electron affinity (energy released by adding an electron to form F–) is 328 kilojoules per mole, which is higher than that of any other element except chlorine.[18] Fluorine has a relatively small covalent radius, on average about 60 picometers, which slightly exceeds the numbers of neon but is surpassed by those of oxygen.[19]
Isotopes
Fluorine occurs naturally on Earth exclusively in the form of its only stable isotope, fluorine-19,[20] which makes the element both monoisotopic and mononuclidic. In total, at least 17 radioisotopes have been synthesized, ranging in mass number from 14 to 31.[21]
Fluorine-18 is the most stable radioisotope of fluorine, with a half-life of 109.77 minutes, and the lightest unstable nuclide with equal odd numbers of protons and neutrons.[22] All isotopes heavier than the stable fluorine-19 decay via beta minus decay (electron emission), for some isotopes possibly together with neutron emission.[21] Isotopes lighter than the stable fluorine-19 undergo electron capture, while fluorine-17 and fluorine-18 decay via beta plus decay (positron emission).[21]
To date, only one nuclear isomer, fluorine-18m, has been characterized.[23] Its half-life before gamma ray emission is approximately 160 nanoseconds, which is less than that of the ground states of the isotopes from fluorine-17 to fluorine-30, except for fluorine-28.[23]
Physical description
Fluorine atoms form diatomic molecules that are gaseous at room temperature. Though sometimes cited as yellow-green, fluorine gas is a very pale yellow; its color can only be observed in concentrated fluorine gas kept in long glass tubes.[24] The element has a "pungent" odor that is noticeable even in extremely low concentrations. Fluorine gas's density is 1.696 grams per liter at 100 kilopascal and 0 °C,[25] about 1.3 times as dense as air.[note 1]
Fluorine condenses to a bright yellow liquid at −188.1 °C (−306.6 °F),[27] a comparable temperature to the boiling points of oxygen and nitrogen. Fluorine solidifies at −219.6 °C (−363.3 °F)[27] into a cubic structure, called beta-fluorine. This phase is transparent and soft, with significant disorder of the molecules. At −227.5 °C (−377.5 °F) fluorine undergoes a solid–solid phase transition into a monoclinic structure called alpha-fluorine. This phase is opaque and hard with close-packed layers of molecules. The solid state phase change requires more energy than the melting point transition and can be violent, shattering samples and blowing out sample holder windows. In general, fluorine's solid state is more similar to oxygen (which has a monoclinic low-temperature structure and similar temperature alpha-beta transition) than to the other halogens.[28][29]
-
Three circles: left is same color as background (old paper somewhat tanned); middle is very light yellow; right is more yellow.
-
Alpha-fluorine-unit-cell-B-3D-vdW.png
-
Solid alpha-fluorine's crystal structure: fluorine molecules lie in shingled layers.
Chemical reactivity
X | XX | HX | BX3 | AlX3 | CX4 |
---|---|---|---|---|---|
F | 159 | 574 | 645 | 582 | 456 |
Cl | 243 | 428 | 444 | 427 | 327 |
Br | 193 | 363 | 368 | 360 | 272 |
I | 151 | 294 | 272 | 285 | 239 |
Fluorine's chemistry is dominated by its tendency to gain an electron. It is the most electronegative element.[31] The removal of an electron from a fluorine atom requires so much energy that no known oxidant can oxidize fluorine to any positive oxidation state.[32]
The high direct reactivity of fluorine gas results from the relative weakness of the fluorine–fluorine bond in elemental fluorine. The bond energy is similar to the easily cleaved oxygen–oxygen bonds of peroxides or nitrogen–nitrogen bonds of hydrazines and significantly weaker than those of dichlorine or dibromine molecules.[33] The covalent radius of fluorine in difluorine molecules, about 71 picometers, is significantly larger than that in other compounds, due to the weak bonding between fluorine atoms.[34]
Reactions between fluorine and other elements are often sudden or explosive.[25] Fluorine is so reactive that water,[25] halogens,[35] and most other substances, even generally nonreactive ones such as radon,[32] burn with a bright flame in a jet of fluorine gas. It can even oxidize elemental nitrogen, which is extremely nonreactive due to its triple bonds, to give nitrogen trifluoride, though this occurs only when activated by electric discharge.[36]
External videos | |
---|---|
![]() |
All metals react with fluorine to form fluorides, but different conditions are required for the reaction depending on the metal. Most frequently, the metals must be in powder forms, because many metals form layers of fluoride on their surfaces that resist further oxidation. Alkali metals react with fluorine violently and form fluorides with formula MF; alkaline earth metals react at room temperature as well, but such reactions are not so exothermic. The metals ruthenium, rhodium, palladium, platinum and gold react least readily with fluorine, and are oxidized by the halogen only in atmospheres of pure fluorine at temperatures of 300–450 °C (575–850 °F).[37] Fluorine reacts explosively with hydrogen in a manner similar to that of alkali metals.[38]
Fluorine is known to form compounds with all elements up to einsteinium, element 99,[39] except for helium, neon, astatine and francium; it is also known to form compounds with rutherfordium, element 104,[40] and seaborgium, element 106.[41] No attempt has been made to oxidize astatine, francium, four later actinides, dubnium or any elements above seaborgium with fluorine, due to the radioactive instability of these elements, though such oxidations are possible in theory.[42] Computational studies have suggested that helium could form a bond with fluorine,[43] and excited states containing neon–fluorine bonds have been observed in a mixture of neon and fluorine irradiated with electrons.[44] Argon reacts with hydrogen fluoride to form argon fluorohydride at low temperatures.[45]
Origin and occurrence
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 |
From the perspective of cosmology, fluorine is relatively rare with 400 ppb in the universe because of its tendency to undergo nuclear fusion with hydrogen to form oxygen and helium, or with helium to become neon and hydrogen, at solar core temperatures. Most fluorine forms either in Type II supernovae when a neutrino hits an atom of neon, in asymptotic giant branch stars, or in blue Wolf-Rayet stars with masses over 40 solar masses (in which stellar winds blow the fluorine out of the star before hydrogen or helium can destroy it).[47][48] Even though fluorine, due to its chemical activity, does not exist in its elementary state on Earth, it can be found in the interstellar medium,[49] and fluorine cations exist in stars and planetary nebulae.[50]
Fluorine is the thirteenth most common element in Earth's crust, making up between 600 and 700 ppm of the crust by mass.[51] Three minerals exist on earth that contain enough fluorine to be mined and used as industrial resources.[51] The most important is fluorite, which is used in smelting, construction, and the manufacture of hydrogen fluoride.[52] Fluorapatite is mined along with other apatites for its phosphate content, and is used mostly for production of phosphate fertilizers. The hexafluorosilicates produced as by-product phosphoric acid are mostly disposed of as waste. Cryolite is the least abundant of the three and is directly used for the production of aluminium. The latter two minerals originate from meteoric water; cryolite has also been found in magmatic water.[53] Fluorocarbon-containing chlorofluorocarbons and tetrafluoromethane have been reported in rocks, presumably having formed without action of living organisms. They are not a commercially or environmentally important source of fluorine.[48]
-
pink globular mass with crystal facets
-
Fluorapatite -
A white rock on a black background
History
![Black and white image of the head of a bearded man looking left](http://upload.wikimedia.org/wikipedia/commons/thumb/0/0b/Henri_Moissan.jpg/220px-Henri_Moissan.jpg)
"Fluorine" is a word that ultimately derives from the Latin verb fluo, meaning "flow." The mineral fluorite, a natural form of calcium fluoride, was first mentioned in 1529 by Georgius Agricola, who named it after its use as a "flux"—an additive that helps melt ores and slags during smelting.[54][55] Agricola first named the mineral "fluorspar" as a latinization of the German Flußspat.[56] Since then, the mineral has been renamed "fluorite," although "fluorspar" is still sometimes used.[57]
Andreas Sigismund Marggraf made the first recorded preparation of "fluoric acid" (hydrofluoric acid in modern nomenclature) in 1764, when he heated fluorite with sulfuric acid in glass, which was greatly corroded by the product.[58] In 1771, Swedish chemist Carl Wilhelm Scheele repeated this reaction.[58] In 1810, French physicist André-Marie Ampère suggested that the acid was a compound of hydrogen with an unknown element, analogous to chlorine;[59] Fluorite was then shown to be mostly composed of calcium fluoride.[60][61] Sir Humphry Davy originally suggested the name fluorine, taking the root from the name of "fluoric acid" and the -ine suffix, similarly to other halogens; this name, with modifications, came to most European languages. Greek, Russian, and several other languages use the ftor or deratives, which comes from Greek φθόριος, meaning "destructive." The new Latin name (fluorum) gave the element its current symbol F, although the symbol Fl is seen in pre-Moissan papers.[62]
Owing to its extreme reactivity, elemental fluorine was not isolated until many years after the characterization of fluorite. Progress in isolating elemental fluorine was slow because it could only be prepared electrolytically and even then under stringent conditions, since the gas reacts with most materials. The generation of elemental fluorine from hydrofluoric acid proved to be exceptionally dangerous, killing or blinding several people who attempted early experiments on this halogen. Jean Dussaud referred to these scientists as "fluorine martyrs," a term still used.[61] In 1886, French chemist Henri Moissan reported the isolation of elemental fluorine, after almost 74 years of effort by other chemists.[63] Moissan received the 1906 Nobel Prize in chemistry for the feat.[note 2]
The two most prominent developments of organofluorine compounds are chlorofluorocarbon refrigerants such as Freon-12, and Teflon. Both were associated with the DuPont company.[58][65] Chlorofluorocarbons are now being replaced by hydrofluorocarbons.[66]
![top level view of large industrial buildings](http://upload.wikimedia.org/wikipedia/commons/thumb/1/1b/K25_Aerial.jpg/240px-K25_Aerial.jpg)
Large-scale productions of elemental fluorine began during World War II. Germany used high-temperature electrolysis to produce tons of chlorine trifluoride, a compound planned to be used as an incendiary.[67] The Manhattan project in the United States produced even more fluorine for use in uranium separation. Gaseous uranium hexafluoride, was used to separate uranium-235, an important nuclear explosive, from the heavier uranium-238 in centrifuges and diffusion plants.[60] Because uranium hexafluoride releases small quantities of corrosive fluorine, the separation plants were built with special materials. All pipes were coated with nickel, which forms a protective fluoride layer on its surface after exposure to fluorine. Joints and flexible parts were fabricated from Teflon.[66][68]
Production
Industrially, fluorine is used either directly as the mined mineral fluorite or as hydrogen fluoride (obtained from the reaction of sulfuric acid with fluorite). Only a very small fraction of industrial fluorine is ever electrolyzed to molecular fluorine, F2. Most fluorine in synthesized organofluorines derives from hydrogen fluoride, not molecular fluorine.[69][70][71]
Electrolytic synthesis
![A machine room](http://upload.wikimedia.org/wikipedia/commons/thumb/0/04/Fluorine_cell_room.jpg/300px-Fluorine_cell_room.jpg)
Several thousand tons of elemental fluorine are produced annually by electrolysis of potassium bifluoride in hydrogen fluoride.[52] Potassium bifluoride forms spontaneously from potassium fluoride and the hydrogen fluoride:
- HF + KF → KHF2
A mixture with the approximate composition KF•2HF melts at 70 °C (158 °F) and is electrolyzed between 70 °C and 130 °C (158–266 °F).[71] Potassium bifluoride increases the electrical conductivity of the solution and provides the bifluoride anion, which is oxidized to form fluorine at the anode. When HF is electrolyzed, hydrogen forms at the cathode and the fluoride ions remain in solution. After electrolysis, potassium fluoride remains in solution.[69]
- 2 HF2– → H2↑ + F2↑ + 2 F–
Henri Moissan first pioneered this method of electrolysis. Moissan used platinum group metal electrodes and carved fluorite containers, but the modern process uses the steel cells that act as cathodes, while blocks of carbon are used as anodes (the Söderberg carbon electrodes are similar to those used in the electrolysis of aluminium). The voltage for the electrolysis varies between 8 and 12 volts.[72]
Pure fluorine gas may be stored in steel cylinders, where the inside surface becomes passivated by a metal fluoride layer that resists further attack.[69][71]
Chemical routes
In 1986, when preparing for a conference to celebrate the 100th anniversary of the discovery of fluorine, Karl Christe discovered a purely chemical preparation of fluorine gas. It involved the reaction of potassium hexafluoromanganate and antimony pentafluoride at 150 °C, in an atmosphere of hydrogen fluoride:
- 2 K
2MnF
6 + 4 SbF
5 → 4 KSbF
6 + 2 MnF
3 + F
2↑
This synthetic route is a rare chemical preparation of elemental fluorine, a reaction not previously thought possible.[73] The manganese(IV) fluoride has to be prepared by reaction with fluorine gas itself,[74] or with krypton difluoride,[75] which is synthesized by reaction with elemental fluorine. This reaction is therefore not an industrially viable way to produce fluorine. Fluorine can by synthesized by the reaction of hexafluoronickelate ion NiF2−
6[51] with the fluorides of the heavier noble gases (krypton[76] and xenon[77]), which can only be produced by a reaction of krypton or xenon with fluorine gas.
Economic aspects
With the exception of countries with planned economics, about 17,000 tonnes of fluorine are produced per year by 11 companies in G7 countries.[78] Fluorine is relatively inexpensive, costing about $5–8 per kilogram when sold as uranium hexafluoride or sulfur hexafluoride. Because of difficulties in storage and handling, the price of pure fluorine gas is much higher.[78]
Compounds
Fluorine exists in the −1 oxidation state in all compounds except for elemental fluorine, where the atoms are bonded to each other and thus at oxidation state 0. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. Most frequently, covalent bonds involving fluorine atoms are single bonds. Higher bonding can occur, for example boron monofluoride features a triple bond.[79] Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding.
Inorganic compounds
Inorganic acids
![Counter-chemical hood like arrangement. Lots of plastic and clean-looking surfaces of the equipment.](http://upload.wikimedia.org/wikipedia/commons/thumb/1/17/Wet_etching_tanks_at_LAAS_%286_inches%29_0468.jpg/250px-Wet_etching_tanks_at_LAAS_%286_inches%29_0468.jpg)
Unlike other hydrohalic acids, such as hydrochloric acid, hydrofluoric acid is only a weak acid in water solution, with acid dissociation constant (pKa) equal to 3.18.[80] The acid's weak acidity in water is due to the hydrogen bonding between the fluoride and hydronium ions and the decrease in entropy that occur when hydrofluoric acid dissociates.[81] When less basic solvents such as dry acetic acid are used, hydrofluoric acid is the strongest of the hydrohalic acids.[82][83] Despite its weakness as an acid in water, hydrogen fluoride is very corrosive, attacking glass. Due to the basicity of the fluoride ion, soluble fluorides give basic water solutions.
Perfluoroacids, which are acids that contain only hydrogen, fluorine and atoms of one other element in the center of the acid's anion, are generally very strong. Fluoroantimonic acid, one such acid, is a "superacid" and the strongest acid known.[84] It has an extremely low pKa of −31.3 and is 20 quintillion (2×1019) times stronger than pure sulfuric acid, which has pKa of −12.[84] This happens because fluorine atoms are univalent and thus cannot form strong chemical bonds to both the antimony atom and the hydron. By occupying all antimony's valence electrons, fluorine atoms prevent the hydron from bonding to it.[note 3]
Metal fluorides
Metal[note 4] fluorides have similarities to other metal halides and to metal oxides, but the ionic character is stronger in metal fluorides than in the corresponding chlorides or oxides. The solubity of ionic fluorides varies greatly, but tends to decrease as the number of fluorides increases. Alkali metal fluorides often resemble the chlorides in terms of structure (all having the sodium chloride structure) and solubilities.[85] Because the fluoride anion is highly basic, many alkali metal fluorides form bifluorides with the formula MHF2; this is a well-known process for sodium and potassium in chemical industry.[86] Among other monofluorides, only silver(I)[87] and thallium(I)[88] fluorides are well-characterized; both are very soluble, unlike other corresponding halides. Unlike the chlorides, alkaline earth metals (except for beryllium)[89] form fluorides that are only sparingly soluble.[62] Several other difluorides, such as those of copper(II) and nickel(II), are soluble.[62] No trifluoride is soluble in water, but several may be soluble in other solvents.[90]
-
white powder in a tube and on a spoon
-
Iron trifluoride -
pink powder in a tube and on a spoon
-
Nickel difluoride -
light blue powder in a tube and on a spoon
While metal tri- and lower fluorides are ionic solids, metal penta- and higher fluorides are molecular and volatile. Tetrafluorides are the borderline: for example, zirconium tetrafluoride is an ionic solid,[91] but germanium tetrafluoride is a molecular gas.[92] This property of the fluoride ion is partially caused by its small radius.[note 5] Only rhenium is known to form bonds with seven fluorides, forming rhenium heptafluoride, which holds the record for number of charged ligands for a metal compound. The compounds shares[93] the pentagonal bipyramid molecular geometry with iodine heptafluoride, the only well-studied analogous nonmetal compound.[note 6] Metal hexafluorides and higher fluorides are oxidants: for example, platinum hexafluoride was the first compound to oxidize molecular oxygen[95] and xenon[96] (see below). Metal fluorides can be volatile solids,[97] liquids,[98] or gases[99] at room temperature.
Nonmetal fluorides
The nonmetal[note 7] fluorides are volatile. Period 2 elements (with the exception of boron, which forms a trifluoride) form fluorides that follow the octet rule: carbon tetrafluoride,[100] nitrogen trifluoride,[101] and oxygen difluoride.[102][103] Lower periods may form fluorides that are hypervalent molecules, such as phosphorus pentafluoride.[104] The reactivity of such species varies greatly: sulfur hexafluoride is inert, while chlorine fluorides are oxidants.
Boron trifluoride is a Lewis acid.[105] Silicon tetrafluoride is a weaker acid and less thermally stable, but carbon tetrafluoride is relatively chemically stable.[100] Among pnictogens, reactivity and acidity of pentafluorides increases down the group; bismuth is an exception, because bismuth pentafluoride is not as acidic as antimony pentafluoride due to its polymeric structure[106] and its trifluoride is ionic.[107] Nitrogen is another special case because it is not known to form a pentafluoride, but tetrafluoroammonium ion, NF+
4, with nitrogen in the formal oxidation state of +5, is known.[108] Chalcogens show a similar trend: hexafluorides increase in acidity and reactivity down the group; oxygen is not known to be oxidized above difluoride. In halogens, unlike previous groups, not all elements form fluorides of their highest oxidation states; chlorine[109] and bromine[110] form pentafluorides, both strong fluorinators; iodine may be oxidized up to iodine heptafluoride.[111] Astatine is not well-studied, and astatine fluoride has not been produced,[112] even though this should be possible.[112]
Noble gas compounds
![black-and-white photo showing icy-looking crystals in a dish](http://upload.wikimedia.org/wikipedia/commons/thumb/b/b2/Xenon_tetrafluoride_crop.gif/220px-Xenon_tetrafluoride_crop.gif)
The noble gases are generally non-reactive because they all have fully filled electronic shells, which are extremely stable. The ability of fluorine-containing platinum hexafluoride to react with xenon was first reported by Neil Bartlett in 1962. He called the compound he prepared xenon hexafluoroplatinate, but since then the product has been revealed to be mixture of different chemicals.[note 8] Later that year, xenon was oxidized directly with fluorine, to form xenon difluoride. Today, only xenon and krypton have well-characterized binary noble gas–fluorine compounds, which include xenon difluoride, krypton difluoride, xenon tetrafluoride, krypton tetrafluoride, xenon hexafluoride and their deratives.[114][115][116] Several oxyfluorides and oxyfluoroxenates are known, including xenon oxytetrafluoride, XeOF4.[117][118]
Radon readily reacts with fluorine to form a solid compound, which is generally thought to be radon difluoride. However, it decomposes on attempted vaporization and its exact composition is uncertain.[119] Calculations indicate that radon difluoride can be ionic, unlike all other binary noble gas fluorides.[120]
Argon can react at extreme conditions with hydrogen fluoride, to form its only stable compound—argon fluoride hydride.[121] Helium can form an analogous helium fluoride hydride but it is metastable,[43] with a lifetime of at most 14 nanoseconds.[122] Argon forms binary argon monofluoride, ArF•, which, because of its metastability, is used in the argon fluoride laser.[123] Even though the reactivity of neon is the lowest of all the elements,[note 9] the element forms a metastable chemical compound, neon monofluoride, NeF•.[44]
Ununoctium, the last currently known group 18 element, is predicted to form ununoctium difluoride, UuoF
2, and ununoctium tetrafluoride, UuoF
4, which is likely to have the tetrahedral Td configuration.[125] However, only a few atoms of ununoctium have been synthesized,[126] and its chemical properties have not been examined yet (as of 2011).
Comparison between the highest oxidation states of oxides and fluorides
For groups 1—6 and 13—16 the highest oxidation states of oxides and fluorides are always equal, and differences are only seen in groups 7—11, mercury, halogens, and the noble gases. The general trend is fluorination allows to achieve relatively low[note 10] but hardly achievable oxidation states; for example, no binary oxide is known for krypton, but krypton difluoride is well-studied.[116] However, very high oxidation states of several elements are known for oxygen only; for example, none has shown that existence of ruthenium octafluoride is possible yet, while ruthenium tetroxide is well-studied.[127]
Later transition metals
![central gray sphere with 4 others attached by spokes in a plane, 90 degrees apart](http://upload.wikimedia.org/wikipedia/commons/thumb/0/09/Mercury-tetrafluoride-3D-balls.png/250px-Mercury-tetrafluoride-3D-balls.png)
With the exceptions of the +7 and +8 oxidation states, fluorine is the key in achieving many rare high oxidation states of the transition metals. For instance, direct reaction of the respective metals with fluorine gives rise to palladium(VI)[128] and platinum(VI).[129] The only occurrence of mercury(IV)[130] is binary mercury(IV) fluoride, synthesized at temperatures close to absolute zero. Gold(V)[131] is only known in the hexafluoroaurate(V) ion, which can be synthesized indirectly under extreme conditions, and the gold(V) fluoride, which is obtained during hexafluoroaurate(V) decomposition. The high oxidizing potential of fluorine has led to the claim of the gold(VII) existence in gold heptafluoride,[132] but current calculations show that the claimed AuF7 molecule was AuF5·F2.[133] The great oxidizing power of fluorine is also illustrated by the fluorine-containing complexes of copper(IV),[134] silver(IV),[135] nickel(IV),[136] iridium(VI),[137] and others. It is possible that the element 113, ununtrium, will be the first element in boron group to form a species in the +5 oxidation state, the fluorine-based hexafluoroununtrate(V), UutF−
6;[138] the possibility of a +5 oxygen-based species is not known to be calculated.
Fluorine is a generally stronger oxidizer than oxygen; however, this strength does not apply for every case. Dinitrogen pentoxide, with nitrogen in the oxidation state of +5, is known; but creating nitrogen pentafluoride would need to squeeze five fluorine atoms attached to the central atom. This is hard to perform, as a nitrogen atom is smaller than most other atoms. It is not known whether the molecule is possible to produce or not, and if possible, whether it is stable or not.[139] Similarly, the highest oxidation states of several late transition metals may be achieved in oxides only: for example, even though only gold(V) is known now, and only in the form of a fluoride, calculations provided by Dementyev et al. in 1997 show that the element may be oxidized up to gold(IX), in the form of the tetroxoaurlyl(IX) ion, [AuO4]+, but not as a fluorine-based compound or ion.[140] Similarly, this and another (Rother et al., 1969) calculations revealed that oxygen-based complexes that contain iridium(IX),[141] platinum(X),[140] and mercury(VIII)[140] might be possible. However, these species were denied by the University of Würzburg in a 2006 paper; it expects platinum(VI), gold(V), and mercury(IV), known in binary fluorides, to be the highest for the elements.[94] It has been shown osmium and iridium may form heptafluorides;[94] for osmium, even an octafluoride may be possible.[94]
Halogens and noble gases
Among halogens, chlorine and bromine form perchlorates[142] and perbromates,[143] both oxygen-based and with the representative halogen in +7 state; however, chlorine, unlike bromine, forms a binary heptoxide.[144] Out of their stable fluorinated species, pentafluorides are the species in the highest oxidation state achieved; however, bromine hexafluoride, BrF6•, is known as well.[145] Iodine shows the reverse picture: no heptoxide is known, unlike heptafluoride, a well-known stable compound; however, periodic acid, containing iodine(VII), is known as well.[146]
Noble gases do not show a trend as well: as noted above, krypton has no known binary oxides, but has a well-studied difluoride. Xenon forms a tetroxide of oxygen-based species,[147] but only a hexafluoride of fluorine-based ones. Neutral xenon octafluoride is not known nor expected to be stable,[148] but octafluoroxenate(VI), XeF2−
8, has been synthesized.[149] Contradictory data is known about fluorides and especially oxides of radon; no binary fluoride or oxide of lighter noble gases are known.
Organic compounds
Organofluorine compounds are chemical compounds that contain a carbon–fluorine chemical bond. This bond is the strongest covalent bond in organic chemistry and is very stable.[150] Fluorine replaces hydrogen in hydrocarbons even at room temperature; after the reaction, the molecular size is not changed significantly. The range of organofluorine compounds is thus diverse; consequently, the research in the area and its uses are driven by commercial value of such compounds in materials science and pharmaceutical chemistry.[66] Organofluorine compounds are synthesized via both direct reaction with fluorine gas, which can be dangerously reactive, or reaction with fluorinating reagents such as sulfur tetrafluoride.[66]
The most industrially important compounds of fluorine include Polytetrafluoroethylene (also called PTFE or Teflon) and hydrofluorocarbons, the main properties of which are affected by the carbon–fluorine bonds in them. The slippery nature of PTFE is the result of chemical stability and repulsion of highly charged fluorine atoms in polymeric chains. Its resistance to van der Waals forces makes PTFE the only known surface to which a gecko cannot stick.[151] Properties of the chlorofluorocarbons and hydrochlorofluorocarbons depend on the number and identity of the halogen atoms. The volatility of these compounds is lower than in most organic compounds because of the strength of the carbon–fluorine bond and carbon–chlorine bond, as well as the molecular polarity induced by the halides and the polarity of halides themselves, which cause intermolecular interactions. The large difference between chlorine and fluorine atomic radii makes chlorofluorocarbons asymmetric, which increases the polarity in the molecules; these effects lead to high solubility potential and higher boiling points of chlorofluorocarbons compared to those of parent hydrocarbons.[152] Chlorofluorocarbons are far less flammable than methane, in part because they contain fewer carbon–hydrogen bonds and in part because the released halides quench the free radicals that sustain flames.[152]
The large inductive effect of the trifluoromethyl group results in the high acid strength of many fluorinated organic acids, which may be comparable to mineral acids. In these compounds, the cation's affinity for the acid proton is decreased by the cation's fluorine content, which increases its affinity for the extra electron left when the acidic proton leaves. For example, acetic acid is a weak acid, with pKa equal to 4.76, while its fluorinated derivative, trifluoroacetic acid has pKa of −0.23, giving it 33,000 times greater formal acidic potential.[153]
![several space-filling (globular) molecular models](http://upload.wikimedia.org/wikipedia/commons/thumb/2/21/Fluorocarbon-montage_rotated-2.png/800px-Fluorocarbon-montage_rotated-2.png)
Applications
Approximately half of mined fluorite is used to help molten metal flow, especially in iron smelting. The other half is converted to hydrofluoric acid, which is mostly used to produce organofluorides or synthetic cryolite.[52]
Uses of fluorine gas
Elemental fluorine is occasionally used as a fluorinating agent in industrial processes. The largest application for elemental fluorine is preparation of uranium hexafluoride, used in the production of nuclear fuels. To obtain the compound, uranium dioxide is treated with hydrofluoric acid, to produce uranium tetrafluoride, which is oxidized by fluorine to give uranium hexafluoride.[78] The second largest application for fluorine gas is sulfur hexafluoride, which is used as an inert dielectric medium in high voltage switching. Sulfur hexafluoride may be produced without using fluorine gas, but the reaction between pure sulfur and pure fluorine gas is the most commonly used in industry.[154]
Elemental fluorine is used for production of tetrafluoromethane,[155] which is utilized for plasma etching in semiconductor manufacturing,[156] flat panel display production, and microelectromechanical systems fabrication.[157][158] These and other require up to 2,000 tonnes annually.[78]
United States and Soviet space scientists in the early 1960s studied elemental fluorine as a possible rocket propellant, due to its exceptionally high specific impulse when used as an oxidizer. The experiments failed because fluorine proved difficult to handle, and its combustion product (typically hydrogen fluoride) was extremely toxic and corrosive.[159][160][161]
Isotope applications
![A rotating, transparent image of a human figure with targeted organs highlighted](http://upload.wikimedia.org/wikipedia/commons/thumb/3/3d/PET-MIPS-anim.gif/170px-PET-MIPS-anim.gif)
Natural fluorine is monoisotopic, consisting solely of fluorine-19. Fluorine compounds are highly amenable to nuclear magnetic resonance, because fluorine-19 has a nuclear spin of ½, a high nuclear magnetic moment, and a high magnetogyric ratio, which allows it to make measurements quickly, comparable with a similar effect based on hydrogen-1.[162] Although it is not one of basic NMR active nuclei used in science and medicine, fluorine-19 is still commonly used in nuclear magnetic resonance, especially in the study of protein structures and conformational changes.[163] Natural fluorine's monoisotopic occurrence makes it useful in uranium enrichment, because uranium hexafluoride molecules differ in mass only due to mass differences between uranium-235 and uranium-238. These mass differences are used to separate uranium-235 and uranium-238 via diffusion and gas centrifugation.
Compounds containing fluorine-18, a radioactive isotope that emits positrons, are often used in PET scanning, because its half-life of about 110 minutes is long by the standards of positron-emitters. One such species is 2-deoxy-2-(18F)fluoro-D-glucose, commonly abbreviated as 18F-FDG.[164][165] In PET imaging, 18F-FDG can be used for the assessment of glucose metabolism in the brain and for imaging tumors in oncology. This radiopharmaceutical is retained by cells and is taken up by tissues with a high need for glucose, such as the brain and most types of malignant tumors.[166] Tomography can thus be used for diagnosis, staging, and monitoring treatment of cancers, particularly in Hodgkin's disease, lung cancer, breast cancer, and many others.
Uses of compounds
Inorganic fluorides and organofluorine compounds find use in a variety of materials and chemicals, including important pharmaceuticals, agrochemicals, lubricants, and textiles.
Hydrofluoric acid and certain fluoride-containing salts are useful for etching glass, light bulbs.[167] Laboratory-produced sodium hexafluoroaluminate, better known as synthetic cryolite, the mineral composed mostly out of this chemical, is used in the electrolysis of aluminium and its purification metallurgy to lower the melting point of aluminium oxide; the compounds also act as a powerful flux for glass.[168]
Perfluorooctanoic acid and tetrafluoroethylene are used in water resistant coatings and in the production of low friction plastics such as PTFE, or Teflon. The low van der Waals forces in solid Teflon give it unusual antiadhesive properties.[169] Nafion, a strongly acidic fluorinated polymer, is a component of fuel cells.[170]
Other fluorine-based compounds were once heavily used in the production of haloalkanes such as chlorofluorocarbons, which are used extensively in air conditioning and in refrigeration. CFCs have been banned for these applications because they contribute to ozone destruction, and have therefore been partly replaced with hydrofluorocarbons that contain no halogen other than fluorine, such as 1,1,1,2-tetrafluoroethane and 2,3,3,3-tetrafluoropropene.[171] Another haloalkane, bromotrifluoromethane ("Halon") is still widely used in ship and aircraft gaseous fire suppression systems. Because Halon production has been banned since 1994, systems are dependent on the pre-ban stores and on recycling.[172]
-
Teflon tape, being wrapped on a pipe thread.
-
A water droplet DWR-coated surface2 edit1.jpg
-
Fluorine-containing durable water repellent makes a fabric water-resistant.
-
a sailer in coveralls wipes down red gas bottles
Biological roles
Living organisms
![](http://upload.wikimedia.org/wikipedia/commons/thumb/7/76/Gifblaar.jpg/200px-Gifblaar.jpg)
Fluoride is not considered an essential mineral element for mammals and humans, though its role in prevention of tooth decay is well-established.[173] Sodium fluoride, tin(II) fluoride, and, most commonly, sodium monofluorophosphate, are used in toothpaste. These or related compounds, such as fluorosilicates, are added to many municipal water supplies, a process called water fluoridation, which has been controversial since its inception in 1945.[173][174] Small amounts of fluoride may be beneficial for bone strength, but this is an issue only in the formulation of artificial diets.[175]
Biologically synthesized organofluorines have been found in microorganisms and plants,[176] but not in animals.[48] The most common example is fluoroacetate, which is used as a defense against herbivores by at least 40 plants in Australia, Brazil and Africa.[177] Other biologically synthesized organofluorines include ω-fluoro fatty acids, fluoroacetone, and 2-fluorocitrate, all of which are believed to be biosynthesized from the intermediate fluoroacetaldehyde.[48] The enzyme adenosyl-fluoride synthase is capable of biologically synthesizing the carbon–fluorine bond.[178]
Pharmaceuticals, agricultural chemicals, and poisons
![](http://upload.wikimedia.org/wikipedia/commons/thumb/4/46/Diflunisal_structure.svg/180px-Diflunisal_structure.svg.png)
Several important pharmaceuticals contain fluorine.[180] Because of the considerable stability of the carbon-fluorine bond, many drugs are fluorinated to prevent their metabolism and prolong their half-lives, allowing for longer times between dosing and activation. For example, an aromatic ring may add to prevent the metabolism of a drug, but this presents a safety problem, because enzymes in the body metabolize some aromatic compounds into poisonous epoxides. Substituting a fluorine into a para position, however, protects the aromatic ring and prevents the epoxide from being produced.[181] Adding fluorine to biologically active organics increases their lipophilicity, because the carbon–fluorine bond is even more hydrophobic than the carbon–hydrogen bond. This effect often increases a drug's bioavailability due to increased cell membrane penetration.[182] Since the carbon–fluorine bond is strong, organofluorides are generally very stable, although the potential of the fluorine to be released as a fluoride leaving group is heavily dependent on the its position in the molecule.
Of drugs that have been commercialized in the past 50 years, 5–15% contain fluorine, and the percentage of currently available fluorine-containing drugs is increasing.[183] For example, fludrocortisone is one of the most common mineralocorticoids, a class of drugs that mimics the actions of aldosterone. The anti-inflammatories dexamethasone and triamcinolone, which are among the most potent of the synthetic corticosteroids class of drugs, contain fluorine.[184] Several inhaled general anesthetic agents, including the most commonly used inhaled agents, also contain fluorine. Examples include sevoflurane, desflurane, and isoflurane, which are hydrofluorocarbon derivatives.[183]
Many SSRI antidepressants are fluorinated organics,[185] including citalopram, escitalopram, fluoxetine, fluvoxamine, and paroxetine. Fluoroquinolones are a commonly used family of broad-spectrum antibiotics.[186] Because biological systems do not metabolize fluorinated molecules easily, fluorinated pharmaceuticals (often antibiotics and antidepressants) are among the major fluorinated organics found in treated city sewage and wastewater.[187]
In addition to pharmaceuticals, an estimated 30% of agrochemical compounds contain fluorine.[188] Because of these, water from agricultural sites contaminates rivers with runoff organofluorines. Synthetic sodium fluoroacetate has been used as an insecticide, especially against cockroaches, and is effective as a bait-poison against mammalian pests.[189] Several other insecticides contain sodium fluoride, which is much less toxic than fluoroacetate.[190]
Chronology
Because groundwater contains fluorine ions, organic items such as bone that are buried in soil will absorb those ions over time. As such, it is possible to determine the relative age of an object by comparing the amount of fluoride with another object found in the same area. It is important as a separation technique in intra-site chronological analysis and inter-site comparisons.[191]
However, if no actual age of any object is known, the ages can only be expressed in terms of one of the objects being older or younger than the other. The fluctuating amount of fluoride found in groundwater means the objects being compared must be in the same local area in order for the comparisons to be accurate. This technique is not always reliable, given that not all objects absorb fluorine at the same rates.[192]
Environmental concerns
![An animation showing colored representation of ozone distribution by year, above North America, through 6 steps. It starts with a lot of ozone especially over Alaska and by 2060 is almost all gone from north to south.](http://upload.wikimedia.org/wikipedia/commons/thumb/4/46/Future_ozone_layer_concentrations.gif/300px-Future_ozone_layer_concentrations.gif)
Chlorofluorocarbons and bromofluorocarbons have come under strict environmental regulation due to their long residence times in the atmosphere, and their contribution to ozone depletion. Since it is specifically 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,[193] and have been widely used as replacements for halocarbons containing chlorine and bromine. Hydrofluorocarbons and perfluorocarbons are greenhouse gases about 4,000 to 10,000 times as potent as carbon dioxide.[194] Sulfur hexafluoride exhibits an even stronger effect, having 20,000 times the global warming potential of carbon dioxide.[194]
Because of the strength of the carbon–fluorine bond, many synthetic fluorocarbons and fluorocarbon-based compounds are persistent in the environment. The fluorosurfactants perfluorooctanesulfonic acid (PFOS) and perfluorooctanoic acid (PFOA), used in waterproofing sprays, and other related chemicals, are persistent global contaminants. PFOS is a persistent organic pollutant and may be harming the health of wildlife.[195] The potential health effects of PFOA to humans are unclear; its tissue distribution in humans is unknown, but studies in rats suggest it is likely to be present primarily in the liver, kidney, and blood, as it is absorbed easily via the gastrointestinal tract in rats. PFOA has been shown not to metabolize in the body, and, unlike chlorinated hydrocarbons, it is neither genotoxic nor lipophilic. It binds to serum albumin and is excreted primarily from the kidney.[196]
Precautions
Elementary state, fluoride ion, and fluoroacetate
![4 diagonal placards with warnings, poison, corrosive, inhalant, oxidant](http://upload.wikimedia.org/wikipedia/commons/thumb/c/c4/Fluorine_DOT_signs.jpg/340px-Fluorine_DOT_signs.jpg)
Elemental fluorine is a highly toxic, corrosive oxidant, and is extremely reactive to organic material (except for perfluorinated substances) even at very low concentrations and can cause ignition at higher ones.[198] Fluorine gas has a pungent characteristic odor that is noticeable in concentrations as low as 20 ppb.[199] Significant irritation to humans can be caused by concentration of fluorine of 25 ppm; at this and higher concentrations fluorine attacks the eyes, respiratory tract, lungs, liver and kidneys. At a concentration of 100 ppm, human eyes and noses are irritated and seriously damaged.[200]
Soluble fluorides are moderately toxic. In the case of the simple salt sodium fluoride, the lethal dose for most adult humans is estimated at 5 to 10 g, which is equivalent to 32 to 64 mg/kg elemental fluoride/kg body weight.[201][202][203] A toxic dose that may lead to adverse health effects is estimated at 3 to 5 mg/kg of fluoride.[204] The fluoride ion is somewhat toxic, in part because of its ability to form, by equilibration, small amounts of hydrogen fluoride in water. This mobile uncharged species diffuses across cell membranes to attack intracellular calcium. The fluoride ion is readily absorbed by the stomach, intestines and excreted through urine. Urine tests have been used to ascertain rates of excretion in order to set upper limits in exposure to fluoride compounds and associated detrimental health effects.[205] Ingested fluoride initially acts locally on the intestinal mucosa, where it forms hydrofluoric acid in the stomach.[206] Thereafter it binds with calcium and interferes with various enzymes.[206] Excess fluoride consumption can lead to skeletal fluorosis, which currently affects millions of people.[207]
Historically, most cases of fluoride poisoning have been caused by accidental ingestion of insecticides containing inorganic fluoride,[208] or (more rarely) rodenticides containing sodium fluoroacetate ("Compound 1080").[209] Currently, most fluoride poisonings are due to the ingestion of fluoride-containing toothpaste.[206] Malfunction of water fluoridation equipment has occurred several times, including a notable incident in Alaska, which affected nearly 300 people and killed one person.[210]
Hydrofluoric acid
![A tube of gel and the box it came in, marked calcium gloconate](http://upload.wikimedia.org/wikipedia/commons/thumb/7/72/Calcium_gluconate_gel_crop.jpg/240px-Calcium_gluconate_gel_crop.jpg)
Hydrofluoric acid is a contact poison, and must be handled with extreme care far beyond that accorded to other mineral acids, even the analogous hydrochloric acid, HCl. Owing to its lesser chemical dissociation in water (remaining a neutral molecule), hydrogen fluoride penetrates tissue more quickly than typical acids. Poisoning can occur readily through exposure of skin or eyes, or when inhaled or swallowed. Symptoms of exposure to hydrofluoric acid may not be immediately evident. Hydrogen fluoride interferes with nerve function, meaning that burns may not initially be painful. Accidental exposures can go unnoticed, delaying treatment and increasing the extent and seriousness of the injury.[211]
Once absorbed into blood through the skin, hydrogen fluoride reacts with blood calcium and may cause cardiac arrest.[212] Formation of insoluble calcium fluoride possibly causes both a fall in calcium serum and the strong pain associated with tissue toxicity.[213] In some cases, exposures can lead to hypocalcemia. Burns with areas larger than 160 cm2 (25 in2) can cause serious systemic toxicity from interference with blood and tissue calcium levels.[214]
Hydrofluoric acid exposure is often treated with calcium gluconate, a source of Ca2+ that binds with the fluoride ions. Hydrogen fluoride chemical burns to the skin can be treated with a water wash and 2.5% calcium gluconate gel[215][216] or special rinsing solutions.[217][218] However, because it is absorbed, medical treatment is necessary; in some cases, amputation may be required.[214]
See also
- Halogens—the periodic group to which fluorine belongs
- Caesium—the least electronegative element, fluorine's "opposite"
Notes
- ^ Density of air at 100 kilopascal and 0 °C is 1.2724 grams per liter.[26]
- ^ In addition to his fluorine isolation, Moissan was honored for inventing the electric arc furnace.[64]
- ^ Hydron is bonded to fluorine by very weak dative bonds, but the low strength of the bond makes the hexafluoroantimonate anion non-coordinating and causes the extreme acidity of the compound.[84]
- ^ In this article, metalloids are not treated separately from metals and nonmetals, but among elements they are closer to. For example, germanium is treated as a metal, and silicon as a nonmetal. Antimony, bismuth and polonium are included for comparison among nonmetals, even though they are closer to metals chemically than to nonmetals.
- ^ For example, larger oxide ions, which are weaker oxidants and are more likely to form covalent bonds, form covalent molecules with only four metals in neutral binary compounds (manganese heptoxide, technetium heptoxide, ruthenium tetroxide and osmium tetroxide). Fluorine forms covalent bonds to twelve metals; see fluoride volatility.
- ^ Calculation shows that currently unknown (possible but not produced) osmium heptafluoride and iridium heptafluoride have this structure.[94]
- ^ The article treats noble gases separately from nonmetals and hydrogen is discussed in the Inorganic acids section. P-block period 7 elements have not been studied and thus are not included.
- ^ Bartlett probably synthesized a mixture of monofluoroxenyl(II) hexafluoroplatinate, [XeF]+[PtF6]–, monofluoroxenyl(II) undecafluorodiplatinate, [XeF]+[Pt2F11]–, and trifluorodixenyl(II) hexafluoroplatinate, [Xe2F3]+[PtF6]–.[113]
- ^ Helium and neon are the only elements that have no known "stable" compounds; i. e., the compounds that do not decay on specific conditions. Calculations show that helium, unlike neon, may form some compounds that will not decay over time in specific conditions; additionally, clathrates are known for every noble gas but neon.[124]
- ^ There is no general line where oxidation states are "relatively low" or "relatively high," they rely on specific elements (and defined only for elements that have highest oxides and fluorides are in different oxidation states); in general, +7 and +8 are high, while +4 and below are low. States +5 and +6 rely on element properties, like atomic radius; for a small nitrogen atom, +5 is what is called "high" here, but for larger palladium and platinum +6 is still "low."
References
- ^ "Standard Atomic Weights: Fluorine". CIAAW. 2021.
- ^ Prohaska, Thomas; Irrgeher, Johanna; Benefield, Jacqueline; Böhlke, John K.; Chesson, Lesley A.; Coplen, Tyler B.; Ding, Tiping; Dunn, Philip J. H.; Gröning, Manfred; Holden, Norman E.; Meijer, Harro A. J. (4 May 2022). "Standard atomic weights of the elements 2021 (IUPAC Technical Report)". Pure and Applied Chemistry. doi:10.1515/pac-2019-0603. ISSN 1365-3075.
- ^ a b c d Jaccaud et al. 2000, p. 381.
- ^ a b c Haynes 2011, p. 4.121.
- ^ a b Jaccaud et al. 2000, p. 382.
- ^ a b c Compressed Gas Association 1999, p. 365.
- ^ "Triple Point | The Elements Handbook at KnowledgeDoor". KnowledgeDoor.
- ^ Himmel, D.; Riedel, S. (2007). "After 20 Years, Theoretical Evidence That 'AuF7' Is Actually AuF5·F2". Inorganic Chemistry. 46 (13). 5338–5342. doi:10.1021/ic700431s.
- ^ Dean 1999, p. 4.6.
- ^ Dean 1999, p. 4.35.
- ^ Matsui 2006, p. 257.
- ^ Yaws & Braker 2001, p. 385.
- ^ Mackay, Mackay & Henderson 2002, p. 72.
- ^ Cheng et al. 1999.
- ^ Chisté & Bé 2011.
- ^ a b Aigueperse et al. 2005, "Fluorine," p. 1.
- ^ a b Dean 1999, p. 564.
- ^ Blondel, C.; Delsart, C.; Goldfarb, F. (2001). "Electron spectrometry at the μeV level and the electron affinities of Si and F". Journal of Physics B: Atomic, Molecular and Optical Physics (34): L281–L288. doi:10.1088/0953-4075/34/9/101.
{{cite journal}}
: Unknown parameter|pii=
ignored (help)CS1 maint: multiple names: authors list (link) - ^ Cordero, Beatriz; Gómez, Verónica; Platero-Prats, Ana E.; Revés, Marc; Echeverría, Jorge; Cremades, Eduard; Barragán, Flavia; Alvarez, Santiago (2008). "Covalent radii revisited". Dalton Trans. (21): 2832–2838. doi:10.1039/b801115j.
{{cite journal}}
: CS1 maint: multiple names: authors list (link) - ^ National Nuclear Data Center. "NuDat 2.1 Database – Fluorine-19". Brookhaven National Laboratory. Retrieved 2005-09-31.
{{cite web}}
: Check date values in:|accessdate=
(help) - ^ a b c National Nuclear Data Center. "NuDat 2.1 Database". Brookhaven National Laboratory. Retrieved 2005-09-31.
{{cite web}}
: Check date values in:|accessdate=
(help) - ^ National Nuclear Data Center. "NuDat 2.1 Database – Fluorine-18". Brookhaven National Laboratory. Retrieved 2005-09-31.
{{cite web}}
: Check date values in:|accessdate=
(help) - ^ a b Holden, N. E. (2004). "Table of the Isotopes". In D. R. Lide (ed.). CRC Handbook of Chemistry and Physics (85th ed.). CRC Press. Section 11. ISBN 978-0849304859.
{{cite book}}
: Unknown parameter|nopp=
ignored (|no-pp=
suggested) (help) - ^ Burdon, J.; Emson, B.; Edwards, A. J. (1987). "Is Fluorine Really Yellow?". Journal of Fluorine Chemistry. 34: 471–474. doi:10.1016/S0022-1139(00)85188-X.
- ^ a b c Aigueperse et al. 2005, "Fluorine," p. 2.
- ^ Shelquist, Richard (2010). "An Introduction to Air Density and Density Altitude Calculations". Shelquist Engineering. Retrieved 29 April 2011.
- ^ a b Dean 1999, p. 523.
- ^ Meyer, Lothar (1969). "Phase transitions in Van der Waal's lattices". In Prigogine, I. (ed.). Advances in chemical physics, volume XVI. John Wiley & Sons, Inc. pp. 364–375. doi:10.1002/9780470143612.ch7.
- ^ Attention: This template ({{cite doi}}) is deprecated. To cite the publication identified by doi:10.1063/1.1711946, please use {{cite journal}} (if it was published in a bona fide academic journal, otherwise {{cite report}} with
|doi=10.1063/1.1711946
instead. - ^ Greenwood & Earnshaw 1998, p. 804.
- ^ Cite error: The named reference
Allred
was invoked but never defined (see the help page). - ^ a b Moore, John W.; Stanitski, Conrad L.; Jurs, Peter C. (2010). Principles of Chemistry: The Molecular Science. Cengage Learning. p. 156. ISBN 9780495390794. Retrieved 7 May 2011. Cite error: The named reference "rrr" was defined multiple times with different content (see the help page).
- ^ Macomber, Roger S. (1996). Organic chemistry, Volume 1. University Science Books. p. 230. ISBN 0935702903. Retrieved 26 July 2011.
- ^ Benson, S. W. (1965). "Bond Energies". Journal Chemical Education. 42 (9): 502–519. Bibcode:1965JChEd..42..502B. doi:10.1021/ed042p502.
- ^ Aigueperse et al. 2005, "Fluorine," p. 3.
- ^ Lidin, Molochko & Andreeva 2000, p. 252.
- ^ Lidin, Molochko & Andreeva 2000, pp. 442–455.
- ^ Greenwood & Earnshaw 1998, p. 844.
- ^ Young, J. P.; Haire, R. G.; Peterson, J. R.; Ensor, D. D.; Fellow, R. L. (1981). "Chemical Consequences of Radioactive Decay. 2. Spectrophotometric Study of the Ingrowth of Berkelium-249 and Californium-249 Into Halides of Einsteinium-253". Inorganic Chemistry. 20 (11): 3979–3983. doi:10.1021/ic50225a076.
- ^ Kratz, J. V. (2003). "Critical Evaluation of the Chemical Properties of the Transactinide Elements (IUPAC Technical Report)" (PDF). Pure and Applied Chemistry. 75 (1): 103. doi:10.1351/pac200375010103. Retrieved 4 May 2011.
- ^ Kahn, Bernd (2007). Radioanalytical Chemistry. ISBN 9780387341224. Retrieved 4 May 2011.
- ^ Even, J.; Kratz, W.; Ballof, J.; Buda, R. A.; Eberhardt, K.; Gromm, E.; Hild, D.; Liebe, D.; Mendel, M.; Reichert, P.; Thörle-Pospiech, P.; Wiehl, N.; Wunderlich, T.; Brüchle, W.; Düllmann, Ch.E.; Jäger, E.; Krier, J.; Schädel, M.; Schausten, B.; Semchenkov, A.; Nayak, D.; Toyoshima, A.; Türler, A.; Yakushev, A. (2008). First Transactinide Chemistry Behind TASCA (PDF) (Report). Gesellschaft für Schwerionenforschung. p. 1. Retrieved 29 April 2011.
- ^ a b Lewars 2008, pp. 70–71.
- ^ a b Steigerwald, F.; Walter, W.; Langhoff, H.; Hammer, W. (1988). "Emission Spectrum and Formation Kinetics of Neon Fluoride". Zeitschrift für Physik D. 7 (4): 379–382. Bibcode:1988ZPhyD...7..379S. doi:10.1007/BF01439807. Cite error: The named reference "NeF" was defined multiple times with different content (see the help page).
- ^ Khriachtchev, Leonid; Pettersson, Mika; Runeberg, Nino; Lundell, Jan; Räsänen, Markku (24 August 2000). "A Stable Argon Compound". Nature. 406 (6798): 874–876. doi:10.1038/35022551. PMID 10972285. Retrieved 29 April 2011.
- ^ Cameron, A. G. W. (1973). "Abundance of the elements in the Solar System" (PDF). Space Science Review. 15: 121–146. Bibcode:1973SSRv...15..121C. doi:10.1007/BF00172440.
- ^ Croswell, Ken (2003). "Fluorine: An Element–ary Mystery". Sky and Telescope. Retrieved 3 May 2011.
{{cite journal}}
: Unknown parameter|month=
ignored (help) - ^ a b c d Renda, Agostino; Fenner, Yeshe; Gibson, Brad K.; Karakas, Amanda I.; Lattanzio, John C.; Campbell, Simon; Chieffi, Alessandro; Cunha, Katia; Smith, Verne V. (2004). "On the origin of fluorine in the Milky Way". Monthly Notices of the Royal Astronomical Society. 354 (2): 575–580. arXiv:astro-ph/0410580. Bibcode:2004MNRAS.354..575R. doi:10.1111/j.1365-2966.2004.08215.x. Cite error: The named reference "Murphy2003" was defined multiple times with different content (see the help page).
- ^ Snow, T. P., Jr.; York, D. G. (1981). "The detection of interstellar fluorine in the line of sight toward Delta Scorpii". Astrophysical Journal. 247: L39. Bibcode:1981ApJ...247L..39S. doi:10.1086/183585.
{{cite journal}}
: CS1 maint: multiple names: authors list (link) - ^ Zhang, Y.; Liu, X.-W. (2005). "Fluorine abundances in planetary nebulae" (PDF). The Astrophysical Journal. 631 (1): L61–L63. arXiv:astro-ph/0508339. Bibcode:2005ApJ...631L..61Z. doi:10.1086/497113. Retrieved 3 May 2011.
- ^ a b c Aigueperse et al. 2005, "Fluorine," p. 4.
- ^ a b c Villalba, Gara; Ayres, Robert U.; Schroder, Hans (2008). "Accounting for Fluorine: Production, Use, and Loss". Journal of Industrial Ecology. 11: 85–101. doi:10.1162/jiec.2007.1075.
- ^ Morrison, Roger B. (1935). "The Occurence and Origine of Celestite and Fluorite at Clay Center, Ohio" (PDF). American Mineralogist. 20: 787–789. Retrieved 3 May 2011.
- ^ Greenwood & Earnshaw 1998, p. 790.
- ^ Senning, Alexander (2007). Elsevier's dictionary of chemoetymology: the whies and whences of chemical nomenclature and terminology. Elsevier. p. 149. ISBN 9780444522399. Retrieved 7 May 2011.
- ^ Agricola, Georgius (2003) [1912]. De Re Metallica. trans. Hoover, Herbert Clark; Hoover, Henry Lou. p. 381. ISBN 9780766131972.
- ^ Greenwood & Earnshaw 1998, p. 109.
- ^ a b c Kirsch, Peer (2004). Modern fluoroorganic chemistry: synthesis, reactivity, applications. Wiley-VCH. p. 3. ISBN 978-3527306916.
- ^ Template:Fr iconAmpere, André-Marie (1816). "Suite d'une Classification naturelle pour les Corps simples". Annales de chimie et de physique (in French): 1–5. Retrieved 7 May 2011.
- ^ a b Peter, Meiers (2010). "Discovery of fluorine". Fluoride History. Retrieved 17 January 2011.
- ^ a b Weeks, Mary Elvira (1932). "The discovery of the elements. XVII. The halogen family". Journal of Chemical Education. 9 (11): 1915–1939. Bibcode:1932JChEd...9.1915W. doi:10.1021/ed009p1915.
- ^ a b c Storer, Frank Humphreys (1864). First outlines of a dictionary of solubilities of chemical substances. Cambridge. pp. 278–280. ISBN 978-1176622562.
- ^ Moissan, Henry (1886). "Action d'un courant électrique sur l'acide fluorhydrique anhydre". Comptes rendus hebdomadaires des séances de l'Académie des sciences (in French). 102: 1543–1544. Retrieved 7 May 2011.
- ^ "The Nobel Prize in Chemistry 1906". Nobelprize.org. Retrieved 7 July 2009.
- ^ Hounshell, David A. (1988). Science and corporate strategy: DuPont R&D, 1902–1980. Cambridge University Press. pp. 156–157. ISBN 0521327679.
{{cite book}}
: Unknown parameter|coauthors=
ignored (|author=
suggested) (help) - ^ a b c d Okazoe, Takashi (2009). "Overview on the history of organofluorine chemistry from the viewpoint of material industry". Proceedings of the Japan Academy, Series B. 85 (8): 276–289. Bibcode:2009PJAB...85..276O. doi:10.2183/pjab.85.276.
- ^ Leech, H. R. (1949). "Laboratory and technical production of fluorine and its compounds". Quarterly Reviews, Chemical Society. 3: 22–35. doi:10.1039/QR9490300022.
- ^ Hendricks, James O. (1953). "Industrial Fluorochemicals". Industrial & Engineering Chemistry. 45: 99–105. doi:10.1021/ie50517a034.
- ^ a b c Aigueperse et al. 2005, "Fluorine," pp. 4–5.
- ^ Noyes, Robert (1993). "Hydrofluoric acid". Pollution prevention technology handbook. pp. 149–152. ISBN 9780815513117. Retrieved 7 May 2011.
- ^ a b c Kirsch, Peer (2004). "Fluorine". Modern fluoroorganic chemistry: synthesis, reactivity, applications. pp. 5–10. ISBN 9783527306916. Retrieved 7 May 2011.
- ^ Aigueperse et al. 2005, "Fluorine," p. 6.
- ^ Christe, K. (1986). "Chemical synthesis of elemental fluorine". Inorganic Chemistry. 25 (21): 3721–3724. doi:10.1021/ic00241a001.
{{cite journal}}
: Unknown parameter|years=
ignored (help) - ^ Mazej, Z. (2002). "Room temperature syntheses of MnF3, MnF4 and hexafluoromanganete(IV) salts of alkali cations". Journal of Fluorine Chemistry. 114 (1): 75–80. doi:10.1016/S0022-1139(01)00566-8.
- ^ Lutar, Karel; Jesih, Adolf; Žemva, Boris (1988). "KrF2/MnF4 adducts from KrF2/MnF2 interaction in HF as a route to high purity MnF4". Polyhedron. 7 (13): 1217–19. doi:10.1016/S0277-5387(00)81212-7.
- ^ Lehmann, J. F.; Mercier, H. L. N. P. A.; Schrobilgen, G. J. (2002). "The chemistry of krypton". Coordination Chemistry Reviews. 233–234: 1–39. doi:10.1016/S0010-8545(02)00202-3.
- ^ Tius, M. A. (1995). "Xenon difluoride in synthesis". Tetrahedron. 51 (24): 6605–6634. doi:10.1016/0040-4020(95)00362-C.
- ^ a b c d Aigueperse et al. 2005, "Fluorine," p. 12.
- ^ Allan McQuarrie, Donald; Simon, John Douglas (1997). Physical chemistry: a molecular approach. University Science Books. p. 365. ISBN 0-935702-99-7.
{{cite book}}
: CS1 maint: multiple names: authors list (link) - ^ Herring, F. Geoffrey; Petrucci, Ralph H.; Harwood, William (2002). General chemistry: principles and modern applications (8th ed.). Prentice Hall. p. 678. ISBN 0-13-014329-4.
{{cite book}}
: CS1 maint: multiple names: authors list (link) - ^ Clark, Jim. "The Acidity of the Hydrogen Halides". Retrieved 4 September 2011.
- ^ Jolly, W. L. (1984). Modern Inorganic Chemistry. McGraw-Hill. p. 203. ISBN 0-07-032760-2.
- ^ Cotton, F. A.; Wilkinson, G. (1988). Advanced Inorganic Chemistry. John Wiley and Sons. p. 109. ISBN 0-471-84997-9.
- ^ a b c Olah, George A. (2005). "Crossing Conventional Boundaries in Half a Century of Research". Journal of Organic Chemistry. 70 (7): 2413–2429. doi:10.1021/jo040285o. PMID 15787527.
- ^ Aigueperse et al. 2005, "Fluorine Compounds, Inorganic," pp. 25–27.
- ^ Aigueperse et al. 2005, "Fluorine Compounds, Inorganic," pp. 26–27.
- ^ Milne, George W. A. (2005). Gardner's commercially important chemicals. John Wiley and Sons. p. 553. ISBN 978-0-471-73518-2.
- ^ Arora, M. G. (2003). P-block Elements. Anmol Publications. p. 35. ISBN 8174885633.
- ^ U.S. Department of Health and Human Services (1998). Barry Leonard (ed.). Eighth Annual Report on Carcinogens: 1998 Summary. Institut d'Estudis Catalans. p. 63. ISBN 9780788183966.
- ^ Sobolev, Boris Petrovich (2001). The Rare Earth Trifluorides: Introduction to materials science of multicomponent metal fluoride crystals. Institut d'Estudis Catalans. p. 51. ISBN 847283610X.
- ^ Brown, Paul L.; Mompean, Federico J.; Perrone, Jane; Illemassène, Myriam (2005). Chemical thermodynamics of zirconium. Gulf Professional Publishing. p. 144. ISBN 0-444-51803-7.
{{cite book}}
: CS1 maint: multiple names: authors list (link) - ^ Yaws & Braker 2001, p. 397.
- ^ Vogt, T.; Fitch, A. N.; Cockcroft, J. K. (1994). "Crystal and Molecular Structures of Rhenium Heptafluoride". Science. 263 (5151): 1265–7. Bibcode:1994Sci...263.1265V. doi:10.1126/science.263.5151.1265. PMID 17817431.
{{cite journal}}
: CS1 maint: multiple names: authors list (link) - ^ a b c d Bayerische Julius-Maximilians-Universität Würzburg (2006). The Highest Oxidation States of the 5d Transition Metals: a Quantum-Chemical Study (Report). Chemical Society. pp. 95–105.
{{cite report}}
:|access-date=
requires|url=
(help) - ^ Bartlett, Neil; Lohmann, D. H. (1962). "Dioxygenyl hexafluoroplatinate (V), O2+[PtF6]−". Proceedings of the Chemical Society (3). Chemical Society: 115. doi:10.1039/PS9620000097.
{{cite journal}}
: CS1 maint: multiple names: authors list (link) - ^ Bartlett, N. (1962). "Xenon hexafluoroplatinate (V) Xe+[PtF6]−". Proceedings of the Chemical Society (6). Chemical Society: 218. doi:10.1039/PS9620000197.
- ^ Lide, David R. (2009). Handbook of Chemistry and Physics (90 ed.). CRC Press. pp. 4–81. ISBN 978-1-4200-9084-0.
- ^ Bernhardt, N. A.; Bishop, H. W.; Brusie, J. P.; U.S. Atomic Energy Commission (1947). The preparation of molybdenum hexafluoride (Report). pp. 1–2.
{{cite report}}
: CS1 maint: multiple names: authors list (link) - ^ Neikov, Oleg D.; Naboychenko, Stanislav; Gopienko, Victor G.; Frishberg, Irina V. (2008). Handbook of Non-Ferrous Metal Powders: Technologies and Applications. Elsevier. p. 440. ISBN 9781856174220.
{{cite book}}
: CS1 maint: multiple names: authors list (link) - ^ a b Lemal, D. M. (2004). "Perspective on Fluorocarbon Chemistry". Journal of Organic Chemistry. 69 (1): 1–11. doi:10.1021/jo0302556. PMID 14703372.
- ^ Template:De iconOtto Ruff, Joseph Fischer, Fritz Luft (1928). "Das Stickstoff-3-fluorid". Zeitschrift für anorganische und allgemeine Chemie. 172 (1): 417–425. doi:10.1002/zaac.19281720132.
{{cite journal}}
: CS1 maint: multiple names: authors list (link) - ^ Yaws, Carl L.; Braker, William (2001). Matheson gas data book, Book 2001. McGraw-Hill Professional. p. 651. ISBN 0071358544.
{{cite book}}
: CS1 maint: multiple names: authors list (link) - ^ Lewars 2008, p. 57.
- ^ Musher, J. I. (1969). "The Chemistry of Hypervalent Molecules". Angewandte Chemie International Edition. 8: 54–68. doi:10.1002/anie.196900541.
- ^ Brinck, T.; Murray, J. S.; Politzer, P. (1993). "A computational analysis of the bonding in boron trifluoride and boron trichloride and their complexes with ammonia". Inorganic Chemistry. 32 (12): 2622–2625. doi:10.1021/ic00064a008.
- ^ Hebecker, C. (1971). "Zur Kristallstruktur von Wismutpentafluorid". Z. anorg. allg. Chem. (in German). 384 (2): 111–114. doi:10.1002/zaac.19713840204.
- ^ Wells, A. F. (1984). Structural Inorganic Chemistry (Fifth ed.). ISBN 0198553706.
{{cite book}}
: Unknown parameter|publications=
ignored (help) - ^ Christe, K. O.; Wilson, W. W. (1986). "Synthesis and characterization of tetrafluoroammonium(1+) tetrafluorobromate(1-) and tetrafluoroammonium(1+) tetrafluorooxobromate(1-)". Inorganic Chemistry. 25 (11): 1904–1906. doi:10.1021/ic00231a038.
- ^ Smith, D. F. (1963). "Chlorine Pentafluoride". Science. 141 (3585): 1039–1040. Bibcode:1963Sci...141.1039S. doi:10.1126/science.141.3585.1039. PMID 17739492.
- ^ Hyde, G. A.; Boudakian, M. M. (1968). "Synthesis routes to chlorine and bromine pentafluorides". Inorganic Chemistry. 7 (12): 2648–2649. doi:10.1021/ic50070a039.
- ^ Schumb, W. C.; Lynch, M. A., Jr. (1950). "Iodine Heptafluoride". Industrial & Engineering Chemistry. 42 (7): 1383–1386. doi:10.1021/ie50487a035.
{{cite journal}}
: CS1 maint: multiple names: authors list (link) - ^ a b Zuckerman, J. J.; Hagen, A. P. (1989). Inorganic Reactions and Methods, the Formation of Bonds to Halogens. John Wiley and Sons. p. 31. ISBN 0471186562.
{{cite book}}
: CS1 maint: multiple names: authors list (link) - ^ Wiberg, Wiberg & Holleman 2001, pp. 392–393.
- ^ Wiberg, Wiberg & Holleman 2001, p. 438.
- ^ Template:De iconRiedel, Erwin; Janiak, Christoph (2007). Anorganische Chemie (7th ed.). Walter de Gruyter. p. 393. ISBN 3110189038.
- ^ a b Lehmann, J. F.; Dixon, D. A.; Schrobilgen, G. J. (2001). "X-ray Crystal Structures of α-KrF
2, [KrF][MF
6] (M = As, Sb, Bi), [Kr
2F
3][SbF
6]·KrF
2, [Kr
2F
3]
2[SbF
6]
2·KrF
2, and [Kr
2F
3][AsF
6]·[KrF][AsF
6]; Synthesis and Characterization of [Kr
2F
3][PF
6]·nKrF
2; and Theoretical Studies of KrF
2, KrF+
, Kr
2F+
3, and the [KrF][MF
6] (M = P, As, Sb, Bi) Ion Pairs". Inorganic Chemistry. 40 (13). American Chemical Society: 3002–3017. doi:10.1021/ic001167w. PMID 11399167. - ^ Mackay, Kenneth Malcolm; Mackay, Rosemary Ann; Henderson, W. (2002). Introduction to modern inorganic chemistry (6th ed.). CRC Press. pp. 497–501. ISBN 0748764208. Retrieved 3 May 2011.
{{cite book}}
: CS1 maint: multiple names: authors list (link) - ^ Christe, K. O.; et al. (1995). "On the Structure of the [XeOF5]– Anion and of Heptacoordinated Complex Fluorides Containing One or Two Highly Repulsive Ligands or Sterically Active Free Valence Electron Pairs". Inorganic Chemistry. 34 (7): 1868–1874. doi:10.1021/ic00111a039.
{{cite journal}}
: Explicit use of et al. in:|author=
(help) - ^ Stein, L. (1970). "Ionic Radon Solution". Science. 168 (3929): 362–364. Bibcode:1970Sci...168..362S. doi:10.1126/science.168.3929.362. PMID 17809133.
- ^ Pitzer, Kenneth S. (1975). "Fluorides of radon and element 118". Journal of the Chemical Society, Chemical Communications (18): 760b–761. doi:10.1039/C3975000760b.
- ^ Stump, B.; Eberle, C.; Schweizer, W. B.; Kaiser, M.; Brun, R.; Krauth-Siegel, R. L.; Lentz, D.; Diederich, F. (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.
- ^ Chaban, Galina M.; Lundell, Jan; Gerber, R. Benny (2001). "Lifetime and decomposition pathways of a chemically bound helium compound". The Journal of Chemical Physics. 115 (16): 7341–7344. Bibcode:2001JChPh.115.7341C. doi:10.1063/1.1412467.
- ^ Basting, D.; Marowsky, G. (2005). "Historical Review of the Excimer Laser Development". Excimer Laser Technology. Springer. pp. 8–20. ISBN 9783540200567.
- ^ Lewars 2008, pp. 70–78.
- ^ Han, Young-Kyu; Lee, Yoon Sup (1999). "Structures of RgFn (Rg = Xe, Rn, and Element 118. n = 2, 4.) Calculated by Two-component Spin-Orbit Methods. A Spin-Orbit Induced Isomer of (118)F4". Journal of Physical Chemistry A. 103 (8): 1104–1108. doi:10.1021/jp983665k.
- ^ Barber, Robert C.; Karol, Paul J.; Nakahara, Hiromichi; Vardaci, Emanuele; Vogt, Erich W. (2011). "Discovery of the elements with atomic numbers greater than or equal to 113 (IUPAC Technical Report)". Pure and Applied Chemistry: 1. doi:10.1351/PAC-REP-10-05-01.
- ^ Singh, B.; Srivastava, S. (1991). "Kinetics and Mechanism of Ruthenium tetroxide Catalysed Oxidation of Cyclic Alcohols by Bromate in a Base". Transition Metal Chemistry. 16: 466–468. doi:10.1007/BF01129466.
{{cite journal}}
: CS1 maint: multiple names: authors list (link) - ^ Aullón, G.; Alvarez, S. (2007). "On the Existence of Molecular Palladium(VI) Compounds: Palladium Hexafluoride". Inorganic Chemistry. 46 (7): 2700–2703. doi:10.1021/ic0623819. PMID 17326630.
- ^ Weinstock, B.; Claassen, H. H.; Malm, J. G. (1957). "Platinum Hexafluoride". Journal of the American Chemical Society. 79 (21): 5822–5832. doi:10.1021/ja01578a073.
- ^ Wang, Xuefang; Andrews, Lester; Riedel, Sebastian; Kaupp, Martin (2007). "Mercury is a Transition Metal: The First Experimental Evidence for HgF4". Angewandte Chemie. 119 (44): 8523–8527. doi:10.1002/ange.200703710.
- ^ Emeléus, H. J.; Sharpe, A. G. (1983). Advances in Inorganic Chemistry and Radiochemistry. Academic Press. p. 83. ISBN 0120236273.
- ^ Template:Ru icon Timakov, A. A.; Prusakov, V. N.; Drobyshevskii, Y. V. (1986). Doklad akademii nauk SSSR. USSR Academy of Sciences. p. 125.
{{cite book}}
: Unknown parameter|unused_data=
ignored (help) - ^ Riedel, S.; Kaupp, M. (2006). "Has AuF7 Been Made?". Inorganic Chemistry. 45 (3): 1228–1234. doi:10.1021/ic051944y. PMID 16441134.
- ^ Popova, T. V.; Aksenova, N. V. (2003). "Complexes of Copper in Unstable Oxidation States". Russian Journal of Coordinated Chemistry. 29 (11): 743–765. doi:10.1023/B:RUCO.0000003432.39025.cc.
- ^ Template:Ru iconPopov, N. V.; Kiselev, Y. M.; Sukhoverkhov, V. F.; Spitsyn, V. I. (1987). Doklad akademii nauk SSSR. USSR Academy of Sciences. p. 628.
- ^ Taylor, J. C.; Wilson, P. W. (1973). "The Structures of Fluorides—IV : A Neutron Diffraction Study of K2NiF6". Journal of Inorganic and Nuclear Chemistry. 36 (7): 1561–1563. doi:10.1016/0022-1902(74)80623-8.
- ^ Drews, T; Supeł, J; Hagenbach, A; Seppelt, K. (2006). "Solid State Molecular Structures of Transition Metal Hexafluorides". Inorganic Chemistry. 45 (9): 3782–3788. doi:10.1021/ic052029f. PMID 16634614.
{{cite journal}}
: Unknown parameter|month=
ignored (help)CS1 maint: multiple names: authors list (link) - ^ Haire, Richard G. (2006). "Transactinides and the future elements". In Morss; Edelstein, Norman M.; Fuger, Jean (eds.). The Chemistry of the Actinide and Transactinide Elements (3rd ed.). Springer Science+Business Media. p. 1723. ISBN 1-4020-3555-1.
{{cite book}}
: CS1 maint: ref duplicates default (link) - ^ Lewars 2008, pp. 63–64.
- ^ a b c Template:Ru iconDementyev, A. I.; Kuznetsov, M. L.; Kiselev, Y. M. (1997). "On extremal oxidation states of heavy 5d elements". Zhurnal Neorganicheskoy Khimiyi (42): 1167.
{{cite journal}}
: CS1 maint: multiple names: authors list (link) - ^ Template:De iconRother, P.; Wagner, F.; Zahn, U. (1969). "Chemical consequences of the 193Os(β-)193Ir decay in osmium". Radiochimica Acta (11): 203.
{{cite journal}}
: CS1 maint: multiple names: authors list (link) - ^ Wolff, J. (1998). "Perchlorate and the Thyroid Gland". Pharmacological Reviews. 50 (1): 89–105. PMID 9549759.
- ^ Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic chemistry. Academic Press. p. 439. ISBN 0123526515.
- ^ Byrns, A. C.; Rollefson, G. K. (1934). "The formation of chlorine heptoxideon illuminations of mixtures of chlorine and ozone". Journal of the American Chemical Society. 56 (5): 1250–1251. doi:10.1021/ja01320a506.
- ^ Emeléus, Harry Julius; Sharpe, A. G. (1978). Advances in inorganic chemistry and radiochemistry. Vol. 21. Academic Press. p. 248. ISBN 0120236214.
{{cite book}}
: CS1 maint: multiple names: authors list (link) - ^ Alan Isaacs, John Daintith, Elizabeth Martin, ed. (1984). Concise Science Dictionary. Oxford University Press. p. 356. ISBN 0192115936.
{{cite book}}
: CS1 maint: multiple names: editors list (link) - ^ Gundersen, G.; Hedberg, K.; Huston, J. L. (1970). "Molecular Structure of Xenon Tetroxide, XeO4". J. Chem. Phys. 52: 812–815. Bibcode:1970JChPh..52..812G. doi:10.1063/1.1673060.
- ^ Seppelt, Konrad (1979). "Recent developments in the Chemistry of Some Electronegative Elements". Accounts of Chemical Research. 12 (6): 211–216. doi:10.1021/ar50138a004.
- ^ Chandra, Sulekh (2004). Comprehensive Inorganic Chemistry. New Age International. p. 308. ISBN 8122415121.
- ^ O'Hagan, D. (2008). "Understanding organofluorine chemistry. An introduction to the C–F bond". Chemical Society Reviews. 37 (2): 308–19. doi:10.1039/b711844a. PMID 18197347.
- ^ University of California, Berkeley. "Research into Gecko Adhesion". Archived from the original on 14 October 2007. Retrieved 29 April 2011.
- ^ a b Sukornick, B. (1989). "Potentially acceptable substitutes for the chlorofluorocarbons". International Journal of Thermophysics. 10 (3): 553–561. Bibcode:1989IJT....10..553S. doi:10.1007/BF00507978.
- ^ McMurry, John (2009). Organic Chemistry: With Biological Applications. Mary Finch. p. 618. ISBN 9780495391456. Retrieved 9 April 2011.
- ^ Aigueperse et al. 2005, "Fluorine Compounds, Inorganic," p. 35.
- ^ Aigueperse et al. 2005, "Fluorine," p. 11.
- ^ El-Kareh, Badih (1994). "Fluorine –Based Plasmas". Fundamentals of semiconductor processing technology. p. 317. ISBN 9780792395348. Retrieved 7 May 2011.
- ^ Arana, Leonel R.; de Mas, Nuria; Schmidt, Aleksander J.; Franz, Martin A.; Jensen, Schmidt F.; Jensen, Klaus F. (2007). "Isotropic etching of silicon in fluorine gas for MEMS micromachining". Journal Micromechanical Microenergy. 17 (2): 384. Bibcode:2007JMiMi..17..384A. doi:10.1088/0960-1317/17/2/026.
- ^ Bahreyni, Behraad (2008). "Vapour Phase Etching". Fabrication and design of resonant microdevices. pp. 27–28. ISBN 9780815515777. Retrieved 7 May 2011.
- ^ Krieger, F. J. (1960). The Russian Literature on Rocket Propellant (PDF) (Report). The Rand Corporation. p. 17. Retrieved 2011-02-31.
{{cite report}}
: Check date values in:|accessdate=
(help) - ^ Sutton, Oscar; Biblarz (2010). "Liquid Oxidizers". Rocket Propulsion Elements. p. 256. ISBN 9780470080245. Retrieved 7 May 2011.
- ^ Tormey, John (1957). "Liquid Rocket Propellants—Is There an Energy Limit?". Industrial & Engineering Chemistry. 49 (9): 1339–1343. doi:10.1021/ie50573a021.
- ^ Nelson, J. H. (2003). Nuclear Magnetic Resonance Spectroscopy. Prentice Hall. pp. 129–139. ISBN 0-13-033451-0. Retrieved 7 May 2011.
- ^ Danielson, Mark A.; Falke, Joseph J. (1996). "Use of 19F NMR to probe protein structure and conformational changes". Annual Review of Biophysics and Biomolecular Structure. 25: 163–195. doi:10.1146/annurev.bb.25.060196.001115. PMC 2899692. PMID 8800468.
- ^ Georges, Audi (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A. 729. Atomic Mass Data Center: 3–128. Bibcode:2003NuPhA.729....3A. doi:10.1016/j.nuclphysa.2003.11.001.
- ^ Schmitz, A.; Kälicke, T.; Willkomm, P.; Grünwald, F.; Kandyba, J.; Schmitt, O. (2000). "Use of fluorine-18 fluoro-2-deoxy-D-glucose positron emission tomography in assessing the process of tuberculous spondylitis". Journal of spinal disorders. 13 (6): 541–4. doi:10.1097/00002517-200012000-00016. PMID 11132989.
- ^ Bustamante, Ernesto; Pedersen, Peter L. (1977). "High Aerobic Glycolysis of Rat Hepatoma Cells in Culture: Role of Mitochondrial Hexokinase". Proceedings of the National Academy of Sciences. 74 (9): 3735–3739. Bibcode:1977PNAS...74.3735B. doi:10.1073/pnas.74.9.3735. PMC 431708. PMID 198801. Retrieved 7 May 2011.
- ^ Pulker, Hans K. (1999). Coatings on glass. ISBN 9780444501035.
- ^ Harper, Charles A. (2001). Handbook of ceramics, glasses, and diamonds. p. A26. ISBN 9780070267121. Retrieved 7 May 2011.
- ^ Autumn, Kellar (2006). "Properties, Principles, and Parameters of the Gecko Adhesive System". Biomedical and life sciences: 225–256. doi:10.1007/978-3-540-31049-5_12.
- ^ Mauritz, Kenneth A.; Moore, Robert B. (2004). "State of Understanding of Nafion". Chemical Reviews. 104 (10): 4535–4586. doi:10.1021/cr0207123. PMID 15669162.
- ^ Smart, Bruce E.; Tatlow, J. C. (1994). "Aliphatic Fluorides Go Commercial". Organofluorine chemistry: principles and commercial applications. pp. 12–15. ISBN 9780306446108.
- ^ "Executive summary". Washington, D.C.: National Academies Press. 1997. p. 1. ISBN 9780309074926 http://www.nap.edu/openbook.php?record_id=5744&page=1.
{{cite book}}
: Missing or empty|title=
(help); Unknown parameter|Title=
ignored (|title=
suggested) (help) - ^ a b Olivares, M.; Uauy, R. (2004). "Essential nutrients in drinking-water (Draft)" (PDF). WHO. Retrieved 30 December 2008.
- ^ Pizzo, G.; Piscopo, M. R.; Pizzo, I.; Giuliana, G. (2007). "Community water fluoridation and caries prevention: a critical review". Clinical Oral Investigations. 11 (3): 189–193. doi:10.1007/s00784-007-0111-6. PMID 17333303.
- ^ Nielsen, Forrest H. (2009). "Micronutrients in Parenteral Nutrition: Boron, Silicon, and Fluoride". Gastroenterology. 137 (5 Suppl): S55–S60. doi:10.1053/j.gastro.2009.07.072. PMID 19874950.
- ^ Gribble, Gordon W. (2002). "Naturally Occurring Organofluorines". The Handbook of Environmental Chemistry. 3N: 121–136. doi:10.1007/10721878_5.
- ^ Proudfoot, A. T.; Bradberry, S. M.; Vale, J. A. (2006). "Sodium fluoroacetate poisoning". Toxicological Reviews. 25 (4): 213–219. doi:10.2165/00139709-200625040-00002. PMID 17288493.
- ^ O'Hagan, D.; Schaffrath, C.; Cobb, S. L.; Hamilton, J. T.; Murphy, C. D. (2002). "Biochemistry: biosynthesis of an organofluorine molecule". Nature. 416 (6878): 279. doi:10.1038/416279a. PMID 11907567.
- ^ Seth, S. D.; Seth, Vimlesh. Textbook Of Pharmacology (3rd ed.). Elsevier India. p. VE-40. ISBN 9788131211588.
{{cite book}}
: Text "year-2009" ignored (help)CS1 maint: multiple names: authors list (link) - ^ National Research Council (U.S.). Committee on Fluoride in Drinking Water (2006). Fluoride in drinking water: a scientific review of EPA's standards. National Academies Press. p. 49. ISBN 978-0309101288.
- ^ Rentmeister, A.; Arnold, F. H.; Fasan, R. (2009). "Chemo-enzymatic fluorination of unactivated organic compounds". Nature Chemical Biology. 5 (1): 26–8. doi:10.1038/nchembio.128. PMC 2713661. PMID 19011638.
{{cite journal}}
: Unknown parameter|month=
ignored (help) - ^ "Fluorine-A vital element in the medicine chest" (PDF). Pharmaceutical Chemistry. 2005. pp. 1–4. Retrieved 26 August 2010.
- ^ a b Filler, R.; Saha, R. (2009). "Fluorine in medicinal chemistry: a century of progress and a 60-year retrospective of selected highlights". Future Medicinal Chemistry. 1 (5): 777–91. doi:10.4155/fmc.09.65. PMID 21426080.
{{cite journal}}
: Unknown parameter|month=
ignored (help)CS1 maint: multiple names: authors list (link) - ^ Hagmann, W.K. (2008). "The many roles for fluorine in medicinal chemistry". Journal of Medicinal Chemistry. 51 (15): 4359–69. doi:10.1021/jm800219f. PMID 18570365.
{{cite journal}}
: Unknown parameter|month=
ignored (help) - ^ Mitchell, Siobhan E.; Triggle, D. J. (2004). Antidepressants. Infobase Publishing. pp. 37–39. ISBN 9781438101927.
{{cite book}}
: CS1 maint: multiple names: authors list (link) - ^ Nelson, J.M.; Chiller, T. M.; Powers, J. H.; Angulo, F. J. (2007). "Fluoroquinolone-resistant Campylobacter species and the withdrawal of fluoroquinolones from use in poultry: a public health success story". Clinical Infectious Diseases. 44 (7): 977–80. doi:10.1086/512369. PMID 17342653.
{{cite journal}}
: CS1 maint: multiple names: authors list (link) - ^ Lietz, A. C.; Meyer, Michael T. (2006). Evaluation of Emerging Contaminants of Concern at the South District Wastewater Treatment Plant Based on Seasonal Sampling Events, Miami-Dade Country, Florida, 2004 (Report). U.S. Geological Survey Scientific Investigations. pp. 7–8.
{{cite report}}
:|access-date=
requires|url=
(help)CS1 maint: multiple names: authors list (link) - ^ "Fluorine's treasure trove". ICIS news. 2006. Retrieved 20 February 2011.
- ^ Eisler, Ronald (February 1995). "Biological Report 27: Sodium monofluoroacetate (1080) Hazards to Fish, Wildlife and Invertebrates: A Synoptic Review" (PDF). Patuxent Environmental Science Center (U.S. National Biological Service). Retrieved 5 June 2011.
- ^ "Class I Ozone-Depleting Substances". Sodium fluoride – pesticidal uses. Scorecard. Retrieved 20 February 2011.
- ^ Göksu Murphy, H. Y.; Oberhofer, Martin; Regulla, D. F. (1991). Scientific dating methods. Kluver Academic Publishers. p. 267. ISBN 978-0-7923-1461-5. Retrieved 7 May 2011.
- ^ Reiche, I. (2006). "Chapter 8 Fluorine and Its Relevance for Archaeological Studies". 2: 253–283. doi:10.1016/S1872-0358(06)02008-2.
{{cite journal}}
: Cite journal requires|journal=
(help) - ^ "Class I Ozone-Depleting Substances". Ozone Depletion. U.S. Environmental Protection Agency. 19 August 2010. Retrieved 20 February 2011.
- ^ a b Intergovernmental Panel on Climate Change. "Changes in Atmospheric Constituents and in Radiative Forcing.". Climate Change 2007: The Physical Science Basis. Contribution of Working Group I to the Fourth Assessment Report of the Intergovernmental Panel on Climate (PDF). Cambridge University. p. Table 2.14, Chap. 2, p. 212. ISBN 9780521705967.
- ^ Betts, Kellyn S. (1 May 2007). "Perfluoroalkyl Acids: What Is the Evidence Telling Us?". Environmental Health Perspectives. 115 (5): A250–A256. doi:10.1289/ehp.115-a250. Retrieved 5 September 2011.
- ^ Steenland, Kyle; Fletcher, Tony; Savitz, David A. (2010). "Epidemiologic Evidence on the Health Effects of Perfluorooctanoic Acid (PFOA)". Environmental Health Perspectives. 118 (8): 1100. doi:10.1289/ehp.0901827. PMC 2920088. PMID 20423814. Retrieved 3 May 2011.
- ^ "UN/NA 1045 (United Nations/North America fluorine data sheet)". Cameo Chemicals. National Oceanic and Atmospheric Administration. Retrieved 17 June 2011.
- ^ Aigueperse et al. 2005, "Fluorine Compounds, Organic," pp. 3–4.
- ^ Banks, Ronald Eric; Sharp, D. W. A; Tatlow, J. C. (1986). Fluorine: the first hundred years (1886–1986). p. 4. ISBN 9780444750396. Retrieved 2 May 2011.
- ^ Keplinger, M. L.; Suissa, L. W. (1968). "Toxicity of Fluorine Short-Term inhalation". American Industrial Hygiene Association Journal. 29 (1): 10–18. doi:10.1080/00028896809342975. PMID 5667185.
- ^ Gosselin, R. E. (1984). Clinical toxicology of commercial products. Williams & Wilkins. pp. III–185–93. ISBN 0683036327.
{{cite book}}
: Unknown parameter|coauthors=
ignored (|author=
suggested) (help) - ^ Baselt, R.C. (2008). Disposition of toxic drugs and chemicals in man. Biomedical Publications. pp. 636–40. ISBN 0962652370.
{{cite book}}
: Check|isbn=
value: checksum (help) - ^ IPCS (2002). Environmental health criteria 227 (Fluoride). International Programme on Chemical Safety, World Health Organization. p. 100. ISBN 9241572272.
- ^ Nochimson, Geofrey (11 May 2011). "Fluoride Toxicity". emedicine. Retrieved 19 May 2011.
- ^ Baez, Ramon J.; Baez, Martha X.; Marthaler, Thomas M. (2000). "Urinary fluoride excretion by children 4–6 years old in a south Texas community". Revista Panamericana de Salud Pública. 7 (4): 242–248. doi:10.1590/S1020-49892000000400005.
- ^ a b c Nochimson, G. (2008). "Toxicity, Fluoride". eMedicine. Retrieved 28 December 2008.
- ^ Reddy, D. R. (2009). "Neurology of endemic skeletal fluorosis". Neurology India. 57 (1): 7–12. doi:10.4103/0028-3886.48793. PMID 19305069.
{{cite journal}}
: CS1 maint: unflagged free DOI (link) - ^ Augenstein, W. L.; Spoerke, D. G.; Kulig, K. W.; Hall, A. H.; Hall, P. K.; Riggs, B. S.; El Saadi, M.; Rumack, B. H. (1991). "Fluoride ingestion in children: a review of 87 cases". Pediatrics. 88 (5): 907–912. PMID 1945630. Retrieved 2 May 2011.
- ^ Green, W. (July 2004). "The use of 1080 for pest control" (PDF). The Animal Health Board and The Department of Conservation. pp. 10–55. Retrieved 16 December 2008.
- ^ Gessner, Bradford D.; Beller, Michael; Middaugh, John P.; Whitford, Gary M. (13 January 1994). "Acute fluoride poisoning from a public water system". New England Journal of Medicine. 330 (2): 95–99. doi:10.1056/NEJM199401133300203. PMID 8259189. Retrieved 2 May 2011.
- ^ Yamashita, Masatomo; Yamashita, Mamoru; Suzuki, Mikio; Hirai, Hiroyasu; Kajigaya, Hiroshi (2001). "Ionophoretic delivery of calcium for experimental hydrofluoric acid burns". Critical Care Medicine. 29 (8): 1575–1578. doi:10.1097/00003246-200108000-00013. PMID 11505130.
- ^ Blodgett, David W.; Suruda, Anthony J.; Crouch, Barbara Insley (2001). "Fatal unintentional occupational poisonings by hydrofluoric acid in the U.S". American Journal of Industrial Medicine. 40 (2): 215–220. doi:10.1002/ajim.1090. PMID 11494350.
- ^ Goldfrank's Manual of Toxicologic Emergencies. New York: McGraw-Hill Professional. 2007. p. 1333. ISBN 978-0071443104. Retrieved 3 May 2011.
- ^ a b "Recommended Medical Treatment for Hydrofluoric Acid Exposure" (PDF). Honeywell Specialty Materials. pp. 2–12. Retrieved 6 May 2009.[dead link]
- ^ El Saadi, M. S.; Hall, A. H.; Hall, P. K.; Riggs, B. S.; Augenstein, W. L.; Rumack, B. H. (1989). "Hydrofluoric acid dermal exposure". Veterinary and human toxicology. 31 (3): 243–247. PMID 2741315.
- ^ Roblin, Isabelle; Urban, Martine; Flicoteau, Domitille; Martin, Chantel; Pradeau, Dominique (2006). "Topical treatment of experimental hydrofluoric acid skin burns by 2.5% calcium gluconate". Journal of Burn Care & Research. 27 (6): 889–894. doi:10.1097/01.BCR.0000245767.54278.09. PMID 17091088.
- ^ Hultén, Peter; Höjer, J.; Ludwigs, U.; Janson, A. (2004). "Hexafluorine vs. standard decontamination to reduce systemic toxicity after dermal exposure to hydrofluoric acid". Journal of Toxicology — Clinical Toxicology. 42 (4): 355–361. doi:10.1081/CLT-120039541. PMID 15461243.
- ^ "News & Views". Chemical Health and Safety. 12 (5): 35–37. 2005. doi:10.1016/j.chs.2005.07.007.
Bibliography
- Aigueperse, Jean; Mollard, Paul; Devilliers, Didier; Chemla, Marius; Faron, Robert; Romano, Renée; Cuer, Jean Pierre (2005). Ullmann (ed.). Encyclopedia of Industrial Chemistry. Wiley-VCH. p. 35. doi:10.1002/14356007. ISBN 9783527306732.
- Dean, John A. (1999). Lange's handbook of chemistry (Fifteenth edition). McGraw-Hill, Inc. ISBN 0070161909.
{{cite book}}
: Invalid|ref=harv
(help) - Greenwood, N. N.; Earnshaw, A. (1998). Chemistry of the Elements (second edition). Butterworth Heinemann. ISBN 0750633654.
{{cite book}}
: Invalid|ref=harv
(help) - Lewars, Errol G. (2008). Modelling Marvels. Springer. ISBN 1402069723. Retrieved 7 May 2011.
{{cite book}}
: Invalid|ref=harv
(help) - Lide, David R. (1992). Handbook of Chemistry and Physics (73rd edition, 1992–1993, special student edition). CRC Press. ISBN 0849305667.
{{cite book}}
: Invalid|ref=harv
(help) - Template:Ru icon Lidin, P. A.; Molochko, V. A.; Andreeva, L. L. (2000). Knimicheskiye svoystva neorganicheskikh veshchestv (Chemical properties of inorganic substances). Khimiya. ISBN 5-7245-1163-0.
{{cite book}}
: Invalid|ref=harv
(help) - Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic chemistry. Academic Press. ISBN 9780123526519. Retrieved 3 March 2011.
{{cite book}}
: Invalid|ref=harv
(help) - Yaws, Carl L.; Braker, William (2001). "Fluorine". Matheson gas data book, Book 2001. McGraw-Hill Professional. ISBN 9780071358545.
{{cite book}}
: Invalid|ref=harv
(help)
External links