Also found in: Dictionary, Thesaurus, Medical, Financial, Acronyms, Wikipedia.
fluorine(flo͞o`ərēn, –rĭn), gaseous chemical element; symbol F; at. no. 9; at. wt. 18.9984; m.p. −219.6°C;; b.p. −188.14°C;; density 1.696 grams per liter at STP; valence −1. Fluorine is a yellowish, poisonous, highly corrosive gas. It is the most chemically active nonmetallic element and is the most electronegative of all the elements. It is a member of Group 17 (the halogenshalogen
[Gr.,=salt-bearing], any of the chemically active elements found in Group 17 of the periodic table; the name applies especially to fluorine (symbol F), chlorine (Cl), bromine (Br), and iodine (I).
..... Click the link for more information. ) of the periodic tableperiodic table,
chart of the elements arranged according to the periodic law discovered by Dmitri I. Mendeleev and revised by Henry G. J. Moseley. In the periodic table the elements are arranged in columns and rows according to increasing atomic number (see the table entitled
..... Click the link for more information. .
Fluorine readily displaces the other halogens from their salts. It combines spontaneously with most other elements—exceptions are chlorine, nitrogen, oxygen, and the so-called inert gasesinert gas
or noble gas,
any of the elements in Group 18 of the periodic table. In order of increasing atomic number they are: helium, neon, argon, krypton, xenon, and radon. They are colorless, odorless, tasteless gases and were once believed to be entirely inert, i.e.
..... Click the link for more information. (helium, neon, argon, krypton, xenon, and radon), but it even combines with most of these when heated. Fluorine reacts with most inorganic and organic compounds. With hydrogen it forms hydrogen fluoridehydrogen fluoride,
chemical compound, HF, a colorless, fuming liquid or colorless gas that boils at 19.54°C;. It is miscible with water and is soluble in benzene, toluene, and concentrated sulfuric acid.
..... Click the link for more information. gas, whose water solution is called hydrofluoric acid.
Because of its extreme reactivity, fluorine does not occur uncombined in nature. Fluorine gas is produced commercially by electrolysis of a solution of hydrogen fluoride containing potassium hydrogen fluoride. The mineral fluoritefluorite
, mineral appearing in various colors, e.g., green, yellow-brown, rose, and red. Chemically, it is calcium fluoride, CaF2. Its crystals, commonly cubic, are transparent or translucent and under certain conditions exhibit fluorescence.
..... Click the link for more information. , or fluorspar (calcium fluoride), is the chief commercial source. Cryolitecryolite
[Gr.,=frost stone], mineral usually pure white or colorless but sometimes tinted in shades of pink, brown, or even black and having a luster like that of wax. Chemically, it is a double fluoride of sodium and aluminum, Na3AlF6.
..... Click the link for more information. and apatiteapatite
, mineral, a phosphate of calcium containing chlorine or fluorine, or both, that is transparent to opaque in shades of green, brown, yellow, white, red, and purple. Apatite is a minor constituent in igneous and metamorphic rocks.
..... Click the link for more information. are other important natural compounds.
The importance of fluorine lies largely in its compounds. Fluorite is used as a flux in refining iron; cryolite serves as the electrolyte in the production of aluminum. Compounds of fluorine are also used in the ceramic and glass industries; hydrofluoric acid is used to etch glass and in the manufacture of light bulbs. The addition of one part per million of soluble fluorides to public water supplies has reduced the incidence of tooth decay in many communities, but water with naturally occurring levels as low as four parts per million can damage teeth and bones. In even larger amounts fluorine and fluoride compounds are poisonous. Sodium fluoride is employed as an insecticide.
Halocarbons (compounds of carbon, fluorine, chlorine, and hydrogen) are used extensively in refrigeration and air-conditioning systems. They were widely used as aerosol propellants; but, since they cause depletion of the ozone layerozone layer
region of the stratosphere containing relatively high concentrations of ozone, located at altitudes of 12–30 mi (19–48 km) above the earth's surface.
..... Click the link for more information. , government restrictions have nearly abolished such use. The linking of fluorine and carbon has created some of the most chemically inert compounds known. Fluorocarbons such as TeflonTeflon,
trade name for a solid, chemically inert polymer of tetrafluoroethylene (C2F4), F2C=CF2. Stable up to temperatures around 572°F; (300°C;), Teflon is used in electrical insulation, gaskets, and in making low-adhesion
..... Click the link for more information. have found extensive use as lubricants and bearing materials because of their low friction. Because of their inertness and heat resistance they may be used, for example, as a coating on cooking ware. Because they are not wetted by water or oils, they are sometimes used to add antisoil properties to textiles.
The use of fluorite as a flux was described in 1529 by Georgius Agricola. Many early chemists experimented with hydrogen fluoride gas, among them Scheele, Davy, Lavoisier, and Gay-Lussac. Fluorine gas was first prepared in 1886 by Henri Moissan after nearly three quarters of a century of effort. There was no commercial production of fluorine before World War II, when the use of the gas in a process for refining uranium ores prompted its manufacture.
F, a chemical element in Group VII of Mendeleev’s periodic system; a halogen. Atomic number, 9; atomic weight, 18.998403.
Under normal conditions (0°C and a pressure of 0.1 mega-newton/m2, or 1 kilogram-force/cm2), fluorine is a pale-yellow gas, with a pungent odor. It occurs in nature in the form of one stable isotope—19F. Five radioisotopes have been produced artificially: 16F (half-life T½ < 1 sec), 17F (T½ = 70 sec), 18F (T½ = 111 min), 20F(T½ = 11.4sec), and 21F(T½ = 5 sec).
History. The first fluorine compound—fluorite (fluorspar), CaF2—was described at the end of the 15th century under the name of fluor (from the Latin fluo, “flow”; this property of CaF2 imparts fluidity to the viscous slag of metallurgical products). K. Scheele obtained hydrofluoric acid in 1771. In 1886, H. Moissan isolated free fluorine by the electrolysis of liquid anhydrous hydrogen fluoride containing potassium bifluoride, KHF2.
The chemistry of fluorine began developing in the 1930’s and proceeded especially rapidly during and after World War II (1939–45), in connection with the needs of the atomic industry and rocket engineering. The term ftor (from the Greek phthoros, “destruction,” “loss”), proposed by A. Ampère in 1810, is used only in Russian; the term fluor is used in many countries.
Distribution in nature. The average content of fluorine in the earth’s crust (clarke) is 6.25 × 10–2 percent by weight; its content in acidic igneous rocks (granites) is 8 × 10–2 percent, in basic rocks, 3.7 × 10–2 percent, and in ultrabasic rocks, 1 × 10–2 percent. Fluorine is present in volcanic gases and thermal waters. Its most important compounds are fluorite, cryolite, and topaz (see). A total of 86 fluorine-containing minerals are known. Fluorine compounds are also found in apatites, phosphorites, and other minerals. Fluorine is an important biogenic element. Throughout the earth’s geological history, various products of volcanic eruption, such as gases, have been the source of fluorine penetration into the biosphere.
Physical and chemical properties. Fluorine gas has a density of 1.693 g/liter (0°C and a pressure of 0.1 meganewton/m2, or 1 kilogram-force/cm2), while liquid fluorine has a density of 1.5127 g/cm3 (at the boiling point). The melting point is – 219.61°C, and the boiling point, – 188.13°C. The fluorine molecule is composed of two atoms (F2); at 1000°C, 50 percent of the molecule dissociates, with a dissociation energy of approximately 155 ± 4 kilo-joules/mole (37 ± 1 kcal/mole). Fluorine is poorly soluble in liquid hydrogen fluoride, with a solubility of 2.5 × 10–3 g per 100 g of HF at –70°C and 0.4 × 10–3g at –20°C. In liquid form it is freely soluble in liquid oxygen and ozone.
The electron configuration of the outer shell of a fluorine atom is 2s22p5. In compounds fluorine exhibits an oxidation state of –1. The covalent atomic radius is 0.72 angstrom, and the ionic radius is 1.33 angstroms. The electron affinity is 3.62 eV, and the ionization energy (F→F+) is 17.418 eV. The high electronegativity of the fluorine atom, the highest of all the elements, is explained by the high values of electron affinity and ionization energy. The high reactivity of fluorine determines exothermic fluorination, which in turn is determined by the unusually low dissociation energy of the fluorine molecule and by the high bond energies between the fluorine atom and other atoms. Direct fluorination has a chain mechanism and can therefore easily transform into combustion or explosion.
Fluorine reacts with all the elements except helium, neon, and argon. At low temperatures it interacts with oxygen in a glow discharge, forming oxyfluorides, such as 02 F2 and O3F2. The reactions of fluorine with other halogens are exothermic and result in the formation of interhalogen compounds (seeINTERHALOGEN COMPOUNDS). Chlorine interacts with fluorine when heated to 200°–250°C, yielding chlorine fluoride, CIF, and chlorine trifluoride, CIF3. There is also CIF5, which is obtained by the fluorination of CIF3 at high temperatures and a pressure of 25 meganewtons/m2 (250 kilograms-force/cm2). Bromine and iodine ignite in a fluorine atmosphere at ordinary temperatures, sometimes yielding BrF3, BrF5, IF5, and IF7. Fluorine reacts directly with krypton, xenon, and radon to form the corresponding fluorides, such as XeF4, XeF6, and KrF2. Oxyfluorides of xenon are also known.
The interaction of fluorine with sulfur is accompanied by the liberation of heat and leads to the formation of many sulfur fluorides. Selenium and tellurium form the higher fluorides SeF6 and TeF6. Fluorine ignites on interaction with hydrogen, yielding hydrogen fluoride. This is a radical reaction with chain branching: HF* + H2= HF + H2*; H2* + F2= HF + H + F (where HF* and H2* are molecules in the vibrational-excited state); the reaction is used in chemical lasers. Fluorine reacts with nitrogen only in an electric discharge. Charcoal ignites at ordinary temperatures on interaction with fluorine, while graphite reacts with fluorine on strong heating, sometimes resulting in the formation of solid graphite fluoride (CF)x or gaseous carbon perfluorides, such as CF4 and C2F6. Fluorine interacts with boron, silicon, phosphorus, and arsenic at cold temperatures, forming the corresponding fluorides.
Fluorine combines vigorously with most metals; alkali and alkaline-earth metals ignite in a cold fluorine atmosphere, whereas Bi, Sn, Ti, Mo, and W ignite on slight heating. Hg, Pb, U, and V react with fluorine at room temperature, and Pt reacts with fluorine at red heat. The interaction of metals with fluorine generally yields higher fluorides, such as UF6, MoF6, and HgF2. Certain metals, for example, Fe, Cu, Al, Ni, Mg, and Zn, react with fluorine to form a protective fluoride film, which hinders further reaction.
The interaction of fluorine with metallic oxides at cold temperatures yields metallic fluorides and oxygen; metallic oxyfluorides, such as MoO2F2, may also be formed. Nonmetallic oxides either add fluorine, for example SO2 + F2 = SO2F2, or the oxygen in the oxides is replaced by fluorine, for example, SiO2 + 2F2 = SiF4 + O2. Glass reacts very slowly with fluorine, except in the presence of water, when the reaction proceeds rapidly. Water interacts with fluorine according to the reaction 2H2O + 2F2 = 4HF + O2, a process accompanied by the formation of OF2 and hydrogen peroxide, H2O2. The nitrogen oxides NO and NO2 combine readily with fluorine, forming nitrosyl fluoride, FNO, and nitroxyl fluoride, FNO2, respectively. Carbon monoxide combines with fluorine on heating to form carbonyl fluoride: CO + F2 = COF2.
Metallic hydroxides react with fluorine to form a metallic fluoride and oxygen; for example, 2Ba(OH)2 + 2F2 = 2BaF2 + 2H2O + O2. Aqueous solutions of NaOH and KOH interact with fluorine at 0°C to yield OF2.
The halides of metals or nonmetals interact with fluorine at cold temperatures, at which time all the halogens are replaced by fluorine.
Sulfides, nitrides, and carbides are readily fluorinated. Metallic hydrides react with fluorine at cold temperatures to form a metallic fluoride and HF; ammonia, in vapor form, interacts with fluorine to yield N2 and HF. Fluorine replaces hydrogen in acids or metals in the salts of acids; for example, HNO3 (or NaNO3) + F2 → FNO3+ HF (or NaF). Under more rigorous conditions, fluorine displaces oxygen in the aforementioned compounds to form sulfuryl fluoride; for example, Na2SO4 + 2F2 = 2NaF + SO2F2 + O2. Carbonates of alkali and alkaline-earth metals react with fluorine at ordinary temperatures to yield the corresponding fluoride, CO2, and O2.
Fluorine reacts vigorously with organic substances.
Preparation. The source used in the preparation of fluorine is hydrogen fluoride, which is generally obtained either by the action of sulfuric acid, H2SO4, on fluorite, CaF2, or by the processing of apatites and phosphorites. Fluorine is prepared by the electrolysis of molten potassium bifluoride KF · (1.8–2.0)HF, which is formed during the saturation of molten KF · HF with hydrogen fluoride up to a 40–41 percent content of HF. The electrolyzer is generally made of steel, the electrodes being a carbon anode and a steel cathode. Electrolysis is carried out at 95°–100°C at a voltage of 9–11 V; the fluorine yield over the current reaches 90–95 percent. The resulting fluorine contains up to 5 percent HF, which is removed by freezing with subsequent absorption by sodium fluoride. Fluorine is stored in the gaseous state, under pressure, or in liquid form, cooled by liquid nitrogen, in containers made of nickel and nickel-base alloys (Monel Metal), copper, aluminum and its alloys, brass, or stainless steel.
Uses. Fluorine gas is used for the fluorination of UF4 into UF6, which is used in uranium isotope separation, and for the preparation of chlorine trifluoride, CIF3(a fluorinating agent), sulfur hexafluoride, SF6 (a gaseous insulator in electrical engineering), and metallic fluorides (for example, fluorides of W and V). Liquid fluorine is used as an oxidizing agent in rocket fuels.
The following fluorine compounds are widely used: hydrogen fluoride, aluminum fluoride, silicon fluorides, fluosulfonic acid (a solvent, catalyst, reagent for the preparation of organic compounds containing the—SO2F group), BF3(a catalyst), and fluorocarbons.
Safety measures. Fluorine is toxic. Its maximum permissible concentration in the air is approximately 2 × 10–4 mg/liter; the maximum permissible concentration on exposure for not more than 1 hour is 1.5 × 10–3mg/liter.
A. V. PANKRATOV
Fluorine in the organism. A trace element, fluorine is a permanent component of animal and plant tissues. In inorganic compound form, it is found mainly in the bones of animals and humans, the content being 100–300 mg/kg. Particularly large amounts are present in teeth. The bones of marine animals are richer in fluorine than the bones of land animals. It enters the organism mainly with drinking water; the optimum fluorine content in water is 1–1.5 mg/liter. Fluorine deficiency in humans induces dental caries and in its excess induces fluorosis. High concentrations of fluorine ions are dangerous in view of their ability to inhibit a number of enzymic reactions and to bind biologically important elements (for example, P, Ca, and Mg), thus disturbing their balance in the body. Organic fluorine derivatives are found only in certain plants, such as the South African Dichapetalum cymosum. The principal ones are derivatives of fluoroacetic acid and are toxic to both plants and animals. The biological role of fluorine has not been sufficiently studied. A connection has been established between fluorine exchange and the formation of bone tissue in the skeleton and, in particular, the teeth.
The necessity of fluorine for plants has not been verified.
V. P. POLISHCHUK
Toxicity. Fluorine poisoning is possible for workers in the chemical industry involved in the synthesis of fluorine-containing compounds and the manufacture of phosphorus fertilizers. Fluorine irritates the respiratory tract and induces burns of the skin. Acute poisoning is characterized by irritation of the laryngeal and bronchial mucosa and the eyes, salivation, and the onset of nosebleeds. Severe cases are characterized by pulmonary edema and disturbances of the central nervous system, whereas chronic poisoning leads to conjunctivitis, bronchitis, pneumonia, pneumosclerosis, and fluorosis. An eczema-type skin condition is common. Immediate treatment involves rinsing the eyes with water and, in cases of skin burns, irrigation with 70-percent alcohol; immediate treatment for poisoning by the inhalation of fluorine is the inhalation of oxygen. Preventive measures include the observation of safety regulations, the wearing of protective clothing, regular medical examinations, and the inclusion of calcium and vitamins in the daily diet.
Preparations containing fluorine are used in medicine as antineoplastic agents (5-fluorouracil, florafur, ftorbenzotef), neuroleptics (trifluperidol, fluphenazine, triphthasine), antidepressants (ftorasizin), and narcotics (Fluothane).
Various fluorine compounds are discussed in separate articles (see Index).
REFERENCESRyss, I. G. Khimiia flora i ego neorganicheskikh soedinenii. Moscow, 1956.
Ftor i ego soedineniia, vols. 1–2. Moscow, 1953–56. (Translated from English.)
Professional’nye bolezni, 3rd ed. Moscow, 1973.