Fluorine is a pale yellow gas. It enters into a chemical reaction with almost all substances, including glass.
Fluorine: Fluorite - Ca-F2
Fluorine is used to make medicines. Tablets containing sodium fluoride are prescribed to people to prevent the development of caries. Sodium fluoride is also found in toothpastes.
Teflon is used to make non-stick cookware. This is amazing material. Clean solid slippery like ice. Teflon is very heavy, unlike most plastics, which tend to be lighter than water, Teflon is more than twice the density of water. Teflon is very useful because almost nothing sticks to it and it is impervious to most chemicals. The main value of Teflon is that it has a surprisingly high percentage of fluorine in a small space. By weight, Teflon (polytetrafluoroethylene) is composed of almost 76% fluorine, the remaining 24% is carbon. There are two fluorine atoms for every carbon atom, and each fluorine atom weighs more than a carbon atom.
Freon, or C-H-Cl-F2, is a refrigerant or substance that is used in refrigeration machines (refrigerators and air conditioners).
Fluorine, Properties and parameters of fluorine
Fluorine, Introduction |
|
| Symbol | F |
| Latin name | Fluorine |
| Substance type | simple chemical element |
| Discoverer | A. Moissan |
| Opening year | 1886 |
The main parameters of fluorine according to the periodic table |
|
| Atomic number Z | 9 |
| Atomic mass | 18.9984032 |
| Group | 17 |
| Period | 2 |
| Group affiliation | halogens |
Mechanical properties of fluorine |
|
| Density of gaseous substances (at 0°C and 760mmHg) | 1.696 (Kilogram / Meter 3) |
Thermodynamic properties of fluorine |
|
| Aggregate state under normal conditions | gas |
| Melting point in Kelvin | 53.55 (Kelvin) |
| Melting point Celsius | -219.6 (°C) |
| Boiling point Kelvin | 85.03 (Kelvin) |
| Boiling point Celsius | -188.12 (°C) |
Properties of the fluorine atom |
|
| e-cloud configuration | 1s 2 2s 2 2p 5 |
| Atom radius | 42 10 − 12 (Meter) |
| Number of protons p | 9 |
| Number of neutrons n | 10 |
| Number of electrons e | 9 |
| Mass number A | 19 |
 |
|
Chemical properties of fluorine |
|
| Valence | 1 |
Prevalence of fluoride |
|
| 0.00004% | |
| The sun is made up of fluorine | 0.00005% |
| The oceans are made up of fluorine | 0.00013% |
| The human body is made up of fluorine at | 0.0037% |
Universe |
|
| The universe is made up of fluorine | 0.00004% |
Fluorine
FLUORINE-a; m.[from Greek. phthoros - death, destruction] Chemical element (F), light yellow gas with a pungent odor. Add to drinking water f.
fluorine(lat. Fluorum), a chemical element of group VII of the periodic system, refers to halogens. Free fluorine consists of diatomic molecules (F 2); pale yellow gas with a pungent odor t pl –219.699°C, t bale –188.200°C, density 1.7 g/l. The most active non-metal: reacts with all elements except helium, neon and argon. The interaction of fluorine with many substances easily turns into combustion and explosion. Fluorine destroys many materials (hence the name: Greek phthóros - destruction). The main minerals are fluorite, cryolite, fluorapatite. Fluorine is used to obtain organofluorine compounds and fluorides; fluorine is part of the tissues of living organisms (bones, tooth enamel).
FLUORINEFLUORINE (lat. Fluorum), F (read "fluorine"), a chemical element with atomic number 9, atomic mass 18.998403. Natural fluorine consists of one stable nuclide (cm. NUCLIDE) 19 F. Outer electron layer configuration 2 s 2
p 5
. In compounds, it exhibits only the oxidation state –1 (valency I). Fluorine is located in the second period in group VIIA of the periodic system of elements of Mendeleev, refers to halogens (cm. HALOGENS).
The radius of the neutral fluorine atom is 0.064 nm, the radius of the F ion is 0.115 (2), 0.116 (3), 0.117 (4) and 0.119 (6) nm (the value of the coordination number is indicated in brackets). The successive ionization energies of a neutral fluorine atom are 17.422, 34.987, 62.66, 87.2, and 114.2 eV, respectively. Electron affinity 3.448 eV (the largest among atoms of all elements). According to the Pauling scale, the electronegativity of fluorine is 4 (the highest value among all elements). Fluorine is the most active non-metal.
In its free form, fluorine is a colorless gas with a pungent, suffocating odor.
Discovery history
The history of the discovery of fluorine is associated with the mineral fluorite (cm. FLUORITE), or fluorspar. The composition of this mineral, as now known, corresponds to the formula CaF 2 , and it is the first substance containing fluorine that began to be used by man. In ancient times, it was noted that if fluorite is added to ore during the smelting of metal, the melting temperature of ore and slag decreases, which greatly facilitates the process (hence the name of the mineral - from Latin fluo - flow).
In 1771, by treating fluorite with sulfuric acid, the Swedish chemist K. Scheele (cm. SCHEELE Karl Wilhelm) prepared acid, which he called hydrofluoric acid. French scientist A. Lavoisier (cm. Lavoisier Antoine Laurent) suggested that this acid included a new chemical element, which he proposed to call "fluorine" (Lavoisier believed that hydrofluoric acid is a compound of fluorium with oxygen, because, according to Lavoisier, all acids must contain oxygen). However, he could not select a new element.
The new element was given the name "fluor", which is also reflected in its Latin name. But long-term attempts to isolate this element in a free form were not successful. Many scientists who tried to get it in a free form died during such experiments or became disabled. These are the English chemists brothers T. and G. Knox, and the French J.-L. Gay Lussac (cm. GAY LUSSAC Joseph Louis) and L. J. Tenard (cm. TENAR Louis Jacques), and many others. Sam G. Davy (cm. DEVI Humphrey), who was the first to receive sodium, potassium, calcium and other elements in a free form, as a result of experiments on the production of fluorine by electrolysis, he was poisoned and became seriously ill. Probably, under the impression of all these failures, in 1816, a name similar in sound, but completely different in meaning, was proposed for the new element - fluorine (from the Greek phtoros - destruction, death). This name of the element is accepted only in Russian, the French and Germans continue to call fluorine “fluor”, the British - “fluorine”.
Even such an outstanding scientist as M. Faraday could not obtain free fluorine (cm. FARADEUS Michael). Only in 1886 the French chemist A. Moissan (cm. Moissan Henri), using the electrolysis of liquid hydrogen fluoride HF, cooled to a temperature of -23 ° C (the liquid should contain a little potassium fluoride KF, which ensures its electrical conductivity), was able to obtain the first portion of a new, extremely reactive gas at the anode. In the first experiments, Moissan used a very expensive electrolyzer made of platinum and iridium to obtain fluorine. At the same time, each gram of the resulting fluorine "ate" up to 6 g of platinum. Later, Moissan began to use a much cheaper copper electrolyzer. Fluorine reacts with copper, but during the reaction a very thin film of fluoride is formed, which prevents further destruction of the metal.
Being in nature
The content of fluorine in the earth's crust is quite high and amounts to 0.095% by weight (significantly more than the closest analogue of fluorine in the group - chlorine (cm. CHLORINE)). Due to the high chemical activity of fluorine in the free form, of course, is not found. The most important fluorine minerals are fluorite (fluorspar), as well as fluorapatite 3Ca 3 (PO 4) 2 CaF 2 and cryolite (cm. CRYOLITE) Na 3 AlF 6 . Fluorine as an impurity is part of many minerals and is found in groundwater; in sea water 1.3 10 -4% fluorine.
Receipt
At the first stage of obtaining fluorine, hydrogen fluoride HF is isolated. Preparation of hydrogen fluoride and hydrofluoric acid (cm. HYDROFLUORIC ACID)(hydrofluoric) acid occurs, as a rule, along with the processing of fluorapatite into phosphate fertilizers. The gaseous hydrogen fluoride formed during the sulfuric acid treatment of fluorapatite is then collected, liquefied and used for electrolysis. Electrolysis can be subjected to both a liquid mixture of HF and KF (the process is carried out at a temperature of 15-20°C), and a KH 2 F 3 melt (at a temperature of 70-120°C) or a KHF 2 melt (at a temperature of 245-310°C) .
In the laboratory, to prepare small amounts of free fluorine, one can use either heating MnF 4, during which fluorine is eliminated, or heating a mixture of K 2 MnF 6 and SbF 5:
2K 2 MnF 6 + 4SbF 5 = 4KSbF 6 + 2MnF 3 + F 2 .
Physical and chemical properties
Under normal conditions, fluorine is a gas (density 1.693 kg / m 3) with a pungent odor. Boiling point -188.14°C, melting point -219.62°C. In the solid state, it forms two modifications: the a-form, which exists from the melting point to –227.60°C, and the b-form, which is stable at temperatures lower than –227.60°C.
Like other halogens, fluorine exists as diatomic molecules F 2 . The internuclear distance in the molecule is 0.14165 nm. The F 2 molecule is characterized by an anomalously low energy of dissociation into atoms (158 kJ/mol), which, in particular, determines the high reactivity of fluorine.
The chemical activity of fluorine is extremely high. Of all the elements with fluorine, only three light inert gases do not form fluorides - helium, neon and argon. In all compounds, fluorine exhibits only one oxidation state -1.
Fluorine reacts directly with many simple and complex substances. So, upon contact with water, fluorine reacts with it (it is often said that “water burns in fluorine”):
2F 2 + 2H 2 O \u003d 4HF + O 2.
Fluorine reacts explosively on simple contact with hydrogen:
H 2 + F 2 \u003d 2HF.
In this case, hydrogen fluoride gas HF is formed, which is unlimitedly soluble in water with the formation of a relatively weak hydrofluoric acid.
Fluorine interacts with most non-metals. So, in the reaction of fluorine with graphite, compounds of the general formula CF x are formed, in the reaction of fluorine with silicon, SiF 4 fluoride, and with boron, BF 3 trifluoride. When fluorine interacts with sulfur, compounds SF 6 and SF 4 are formed, etc. (see Fluorides (cm. FLUORIDE)).
A large number of fluorine compounds with other halogens are known, for example, BrF 3, IF 7, ClF, ClF 3 and others, moreover, bromine and iodine ignite in a fluorine atmosphere at ordinary temperature, and chlorine interacts with fluorine when heated to 200-250 ° C.
Do not react directly with fluorine, in addition to the indicated inert gases, also nitrogen, oxygen, diamond, carbon dioxide and carbon monoxide.
Nitrogen trifluoride NF 3 and oxygen fluorides О 2 F 2 and OF 2 were obtained indirectly, in which oxygen has unusual oxidation states +1 and +2.
When fluorine interacts with hydrocarbons, their destruction occurs, accompanied by the production of fluorocarbons of various compositions.
With slight heating (100-250°C), fluorine reacts with silver, vanadium, rhenium and osmium. With gold, titanium, niobium, chromium and some other metals, the reaction involving fluorine begins to proceed at temperatures above 300-350°C. With those metals whose fluorides are nonvolatile (aluminum, iron, copper, etc.), fluorine reacts with a noticeable rate at temperatures above 400-500°C.
Some higher metal fluorides, such as uranium hexafluoride UF 6 , are obtained by acting with fluorine or a fluorinating agent such as BrF 3 on lower halides, for example:
UF 4 + F 2 = UF 6
It should be noted that not only medium fluorides of the NaF or CaF 2 type, but also acidic fluorides - hydrofluorides of the NaHF 2 and KHF 2 types, correspond to the already mentioned hydrofluoric acid HF.
A large number of different organofluorine compounds have also been synthesized. (cm. organofluorine compounds), including the famous Teflon (cm. TEFLON)- material, which is a polymer of tetrafluoroethylene (cm. TETRAFLUOROETHYLENE) .
Application
Fluorine is widely used as a fluorinating agent in the production of various fluorides (SF 6 , BF 3 , WF 6 and others), including compounds of inert gases (cm. NOBLE GASES) xenon and krypton (see Fluorination (cm. FLUORINATION)). Uranium hexafluoride UF 6 is used to separate uranium isotopes. Fluorine is used in the production of Teflon and other fluoroplastics. (cm. Fluoroplastics), fluororubber (cm. fluororubbers), fluorine-containing organic substances and materials that are widely used in engineering, especially in cases where resistance to aggressive media, high temperatures, etc. is required.
Biological role
As a trace element (cm. MICROELEMENTS) Fluoride is found in all organisms. In animals and humans, fluorine is present in bone tissue (in humans, 0.2–1.2%) and, especially, in dentin and tooth enamel. The body of an average person (body weight 70 kg) contains 2.6 g of fluorine; the daily requirement is 2-3 mg and is met mainly with drinking water. A lack of fluoride leads to dental caries. Therefore, fluorine compounds are added to toothpastes, sometimes introduced into drinking water. Excess fluoride in water, however, is also harmful to health. It leads to fluorosis (cm. FLUOROSIS)- changes in the structure of enamel and bone tissue, bone deformation. MPC for the content of fluoride ions in water is 0.7 mg/l. Maximum concentration limit for gaseous fluorine in the air is 0.03 mg/m 3 . The role of fluorine in plants is unclear.
encyclopedic Dictionary. 2009 .
Synonyms:See what "fluorine" is in other dictionaries:
fluorine- fluorine, and ... Russian spelling dictionary
fluorine- fluorine/… Morphemic spelling dictionary
- (lat. Fluorum) F, a chemical element of group VII of the periodic system of Mendeleev, atomic number 9, atomic mass 18.998403, belongs to halogens. Pale yellow gas with a pungent odor, mp? 219.699 .C, tbp? 188.200 .C, density 1.70 g / cm & sup3. ... ... Big Encyclopedic Dictionary
F (from Greek phthoros death, destruction, lat. Fluorum * a. fluorine; n. Fluor; f. fluor; and. fluor), chem. element of group VII periodic. system of Mendeleev, refers to halogens, at. n. 9, at. m. 18.998403. In nature, 1 stable isotope 19F ... Geological Encyclopedia
- (Fluorum), F, chemical element of group VII of the periodic system, atomic number 9, atomic mass 18.9984; refers to halogens; gas, boiling point 188.2shC. Fluorine is used in the production of uranium, freons, medicines and others, as well as in ... ... Modern Encyclopedia
The most reactive element in the Periodic Table is Fluorine. Despite the explosive properties of fluorine, it is a vital element for humans and animals, found in drinking water and toothpaste.
just the facts
- Atomic number (number of protons in the nucleus) 9
- Atomic symbol (in the Periodic Table of the Elements) F
- Atomic weight (average mass of an atom) 18.998
- Density 0.001696 g/cm3
- At room temperature - gas
- Melting point minus 363.32 degrees Fahrenheit (-219.62°C)
- Boiling point minus 306.62 degrees F (-188.12°C)
- Number of isotopes (atoms of the same element with different numbers of neutrons) 18
- Most common F-19 isotopes (100% natural abundance)
fluorite crystal
Chemists have been trying for years to free the element fluorine from various fluorides. However, fluorine does not have a free nature: no chemical substance is able to release fluorine from its compounds, due to its reactive nature.
For centuries, the mineral fluorspar has been used to recycle metals. Calcium fluoride (CaF 2 ) has been used to separate pure metal from unwanted minerals in the ore. "Fluer" (from the Latin word "fluere") means "to flow": the fluid property of fluorspar made it possible to make metals. The mineral was also called Czech emerald because it was used in glass etching.
For many years, fluorine salts or fluorides have been used for welding and for glazed glass. For example, hydrofluoric acid has been used to etch the glass of light bulbs.
Experimenting with fluorspar, scientists have studied its properties and composition for decades. Chemists often produced fluoric acid (hydrofluoric acid, HF), an incredibly reactive and dangerous acid. Even small splashes of this acid on the skin could be fatal. Many scientists were injured, blinded, poisoned or died during the experiments.
- In the early 19th century, André-Marie Ampère of France and Humphry Davy of England announced the discovery of a new element in 1813 and named it fluorine, at the suggestion of Ampère.
- Henry Moisan, a French chemist, finally isolated fluorine in 1886 by electrolysis of dry potassium fluoride (KHF 2) and dry hydrofluoric acid, for which he was awarded the Nobel Prize in 1906.
From now on, fluorine is a vital element in nuclear energy. It is used to produce uranium hexafluoride, which is essential for the separation of uranium isotopes. Sulfur hexafluoride is a gas used to insulate high power transformers.
Chlorofluorocarbons (CFCs) were once used in aerosols, refrigerators, air conditioners, foam packaging and fire extinguishers. These uses have been banned since 1996 because they contribute to ozone depletion. Until 2009, CFCs were used in asthma inhalers, but these types of inhalers were also banned in 2013.
Fluorine is used in many fluorine-containing substances, including solvents and high-temperature plastics such as Teflon (poly-tetrafluoroethene, PTFE). Teflon is well known for its non-stick properties and is used in pans. Fluorine is also used to insulate cables, for plumber's tape, and as the basis of waterproof boots and clothing.
According to the Jefferson Lab, fluoride is added to city water supplies at a rate of one part per million to prevent tooth decay. Several fluoride compounds are added to toothpaste, also to prevent tooth decay.
Although all humans and animals are exposed to and need fluorine, the element fluorine in large enough doses is extremely toxic and dangerous. Fluorine can naturally enter water, air, and vegetation as well as animal hosts in small amounts. Large amounts of fluoride are found in some foods such as tea and shellfish.
Although fluoride is essential for maintaining the strength of our bones and teeth, too much of it can have the opposite effect, causing osteoporosis and tooth decay, and it can also damage the kidneys, nerves, and muscles.
In its gaseous form, fluorine is incredibly dangerous. Small amounts of fluorinated gas are irritating to the eyes and nose, and large amounts can be fatal. Hydrofluoric acid is also fatal, even in small skin contact.
Fluorine, the 13th most abundant element in the earth's crust; it usually settles in the soil and mixes easily with sand, pebbles, coal and clay. Plants can absorb fluorine from the soil, although high concentrations result in plant death. For example, corn and apricot are among the plants most susceptible to damage when exposed to elevated fluorine concentrations.
Who knew? Interesting facts about fluoride
- Sodium fluoride is rat poison.
- Fluorine is the most chemically reactive element on our planet; it can explode on contact with any element except oxygen, helium, neon, and krypton.
- Fluorine is also the most electronegative element; it attracts electrons more easily than any other element.
- The average amount of fluoride in the human body is three milligrams.
- Fluorine is mainly mined in China, Mongolia, Russia, Mexico and South Africa.
- Fluorine is formed in solar stars at the end of their lives (Astrophysical Journal in Letters, 2014). The element forms at the highest pressures and temperatures inside a star as it expands to become a red giant. As the outer layers of a star are shed off, creating a planetary nebula, fluorine moves along with other gases into the interstellar medium, eventually forming new stars and planets.
- About 25% of drugs and medications, including those for cancer, the central nervous system, and the cardiovascular system, contain some form of fluoride.
According to a study (report in the Journal of Fluorine Chemistry) in drug active ingredients, replacing carbon-hydrogen or carbon-oxygen bonds with carbon-fluorine bonds usually shows an improvement in drug efficacy, including increased metabolic stability, increased binding to molecules- targets and improve membrane permeability.
According to this study, a new generation of anti-cancer drugs, as well as fluoride probes for drug delivery, have been tested against cancer stem cells and show promise in fighting cancer cells. The researchers found that the drugs that included fluoride were several times more potent and showed better stability than traditional anti-cancer drugs.
Fluorine(lat. fluorum), f, a chemical element of group vii of the periodic system of Mendeleev, refers to halogens, atomic number 9, atomic mass 18.998403; under normal conditions (0 °С; 0.1 MN/m 2, or 1 kgf / cm 2) is a pale yellow gas with a pungent odor.
Natural F. consists of one stable isotope 19 f. Artificially obtained five radioactive isotopes: 16 f with a half-life T 1/2 < 1 sec, 17 f ( t 1/2 = 70 sec), 18 f ( t 1/2 = 111 min), 20 f ( t 1/2 = 11,4 sec), 21 f ( t1/2 = 5 sec).
History reference. The first F. compound, fluorite (fluorspar) caf 2, was described at the end of the 15th century. under the name "fluor" (from lat. fluo - flow, by the property of cafa 2 to make viscous slags of metallurgical industries fluid-flowing). In 1771 K. Scheele received hydrofluoric acid. Free F. singled out A. Moissan in 1886 by electrolysis of liquid anhydrous hydrogen fluoride containing an admixture of acidic potassium fluoride khf 3 .
Phytochemical chemistry began to develop in the 1930s, especially rapidly during World War II (1939–45) and after it in connection with the needs of the nuclear industry and rocket technology. The name "F." (from the Greek phth o ros - destruction, death), proposed by A. ampere in 1810, used only in Russian. language in many countries the name "fluor" is accepted.
distribution in nature. The average content of F. in the earth's crust (clarke) is 6.25 10 -2% by weight; in acid igneous rocks (granites) it is 8 10 -2%, in basic - 3.7 10 -2%, in ultrabasic - 1 10 -2%. F. is present in volcanic gases and thermal waters. The most important compounds F. - fluorite, cryolite and topaz (cf. Natural fluorides). A total of 86 fluorine-containing minerals are known. F.'s connections are also in apatite, phosphorites etc. F. - important biogenic element. In the history of the Earth, the products of volcanic eruptions (gases, etc.) have been the source of F.'s entry into the biosphere.
Physical and chemical properties . Gaseous F. has a density of 1.693 g/l(0°С and 0.1 MN/m 2, or 1 kgf / cm 2), liquid - 1.5127 g/cm 3(at boiling point); t pl - 219.61°C; t kip - 188.13°C. The F. molecule consists of two atoms (f 2); at 1000°C, 50% of the molecules dissociate, the dissociation energy is about 155 ± 4 kJ/mol(37 ± 1 kcal/mol). F. is poorly soluble in liquid hydrogen fluoride; solubility 2.5 10 -3 G at 100 G hf at -70°C and 0.4 10 -3 at -20°C; in liquid form, it is infinitely soluble in liquid oxygen and ozone. The configuration of the outer electrons of the atom F. 2 s2 2 p2. In compounds, it exhibits an oxidation state of - 1. The covalent radius of the atom is 0.72 a, the ionic radius is 1.33 a. Electron affinity 3.62 ev, ionization energy (f ® f +) 17.418 ev. The high values of electron affinity and ionization energy explain the strong electronegativity of the Ph. atom, the highest among all other elements. The high reactivity of F. causes the exothermicity of fluorination, which, in turn, is determined by the anomalously low dissociation energy of the F. molecule and the large values of the bonding energy of the F. atom with other atoms. Direct fluoridation has a chain mechanism and can easily turn into combustion and explosion. F. reacts with all elements except helium, neon, and argon. It interacts with oxygen in a glow discharge, forming at low temperatures oxygen fluorides o 2 f 2, o 3 f 2, etc. F. reactions with other halogens are exothermic, resulting in the formation interhalogen compounds. Chlorine reacts with F. when heated to 200-250°C, giving chlorine monofluoride cif and chlorine trifluoride clf 3 . Also known cif 5 obtained by fluorination of clf 3 at high temperature and pressure 25 MN/m 2 (250 kgf / cm 2). Bromine and iodine ignite in an F. atmosphere at ordinary temperatures, and brf 3, brf 5, if 5, if 7 can be obtained. F. directly reacts with krypton, xenon and radon, forming the corresponding fluorides (for example, xef 4, xef 6, krf 2). Xenon oxyfluorides are also known.
The interaction of sulfur with sulfur is accompanied by the release of heat and leads to the formation of numerous sulfur fluorides. Selenium and tellurium form the higher fluorides sef 6 tef 6 . F. with hydrogen react with ignition; this creates hydrogen fluoride. This is a chain branching radical reaction: hf* + h 2 = hf + h 2 *; h 2 * + f 2 \u003d hf + H + f (where hf * and h 2 * are molecules in a vibrationally excited state); the reaction is used in chemical lasers. F. reacts with nitrogen only in an electric discharge. Charcoal, when interacting with F., ignites at ordinary temperatures; graphite reacts with it when heated strongly, and solid graphite fluoride (cf) x or gaseous perfluorocarbons cf 4 , c 2 f 6 , etc., can form. F. interacts with boron, silicon, phosphorus, and arsenic in the cold, forming the corresponding fluorides. F. vigorously combines with most metals; alkali and alkaline earth metals ignite in the atmosphere of F. in the cold, bi, sn, ti, mo, w - with slight heating, hg, pb, u, v react with F. at room temperature, pt - at a temperature of dark red heat. When metals interact with F., as a rule, higher fluorides are formed, for example, uf 6 , mof 6 , hgf 2 . Some metals (fe, cu, al, ni, mg, zn) react with F. to form a protective film of fluorides that prevents further reaction.
When F. interacts with metal oxides in the cold, metal fluorides and oxygen are formed; the formation of metal oxyfluorides is also possible (eg moo 2 f 2). Non-metal oxides either add F., for example, so 2 + f 2 \u003d so 2 f 2, or oxygen in them is replaced by F., for example, sio 2 + 2f 2 \u003d sif 4 + o 2. Glass reacts very slowly with F.; in the presence of water, the reaction proceeds rapidly. Water interacts with F.: 2h 2 o + 2f 2 = 4hf + o 2; in this case, of 2 and hydrogen peroxide h 2 o 2 are also formed. Nitrogen oxides no and no 2 easily attach F. with the formation of nitrosyl fluoride fno and nitrile fluoride fno 2, respectively. Carbon monoxide adds F. when heated to form carbonyl fluoride: co + f 2 = cof 2.
Metal hydroxides react with F., forming metal fluoride and oxygen, for example, 2ba (oh) 2 + 2f 2 \u003d 2baf 2 + 2h 2 o + o 2. Aqueous solutions of naoh and koh react with F. at 0°C to form of 2 .
Halides of metals or nonmetals interact with F. in the cold, and F. replaces all halogens. Sulfides, nitrides, and carbides are easily fluorinated. Metal hydrides form metal fluoride and hf with F. in the cold; ammonia (in vapor) - n 2 and hf. F. replaces hydrogen in acids or metals in their salts, for example hno 3 (or nano 3) + f 2 ® fno 3 + hf (or naf); under more severe conditions, F. displaces oxygen from these compounds, forming sulfuryl fluoride, for example, na 2 so 4 + 2f 2 \u003d 2naf + so 2 f 2 + o 2. Alkali and alkaline-earth metal carbonates react with F. at ordinary temperatures; this yields the corresponding fluoride, co 2 and o 2 .
F. vigorously reacts with organic substances.
Receipt. The source for the production of F. is hydrogen fluoride, which is obtained mainly either by the action of sulfuric acid h 2 so 4 on caf 2 fluorite, or by processing apatites and phosphorites. F. is produced by electrolysis of a melt of acidic potassium fluoride kf · (1.8–2.0) hf, which is formed when the melt kf · hf is saturated with hydrogen fluoride to a content of 40–41% hf. The material for the electrolyzer is usually steel; electrodes - carbon anode and steel cathode. Electrolysis is carried out at 95-100°C and voltage 9-11 in; F.'s current output reaches 90-95%. The resulting F. contains up to 5% hf, which is removed by freezing followed by absorption with sodium fluoride. F. is stored in a gaseous state (under pressure) and in liquid form (when cooled with liquid nitrogen) in devices made of nickel and alloys based on it ( monel metal), from copper, aluminum and its alloys, brass, stainless steel.
Application. Gaseous F. serves for fluorination uf 4, in uf 6, used for isotope separation uranium, as well as for the production of chlorine trifluoride clf 3 (fluorinating agent), sulfur hexafluoride sf 6 (gaseous insulator in the electrical industry), metal fluorides (for example, w and v). Liquid F. is an oxidizing agent for rocket fuels.
Numerous F compounds have been widely used. - hydrogen fluoride, aluminum fluoride, silicofluorides, fluorosulfonic acid (solvent, catalyst, reagent for obtaining organic compounds containing a group - so 2 f), bf 3 (catalyst), organofluorine compounds, etc.
Safety . F. is toxic, its maximum permissible concentration in the air is approximately 2 10 -4 mg/l, and the maximum permissible concentration during exposure is not more than 1 h is 1.5 10 -3 mg/l.
A. V. Pankratov.
Fluoride in the body. F. is constantly a part of animal and plant tissues; trace element. In the form of inorganic compounds, it is found mainly in the bones of animals and humans - 100-300 mg/kg; especially a lot of F. in the teeth. The bones of marine animals are richer in F. than those of terrestrial ones. It enters the body of animals and humans mainly with drinking water, the optimal content of F. in which is 1-1.5 mg/l. With a lack of F., a person develops dental caries, with increased intake - fluorosis. High concentrations of F. ions are dangerous due to their ability to inhibit a number of enzymatic reactions, as well as to bind biologically important elements (P, ca, mg, etc.), which disrupts their balance in the body. Organic derivatives of F. are found only in some plants (for example, in the South African dichapetalum cymosum). The main ones are derivatives of fluoroacetic acid, which are toxic to both other plants and animals. The biological role of F. is studied insufficiently. A relationship has been established between F.'s metabolism and the formation of bone tissue of the skeleton and, especially, teeth. The need for F. for plants has not been proven.
V. R. Polishchuk.
F. poisoning is possible for workers in the chemical industry, in the synthesis of fluorine-containing compounds and in the production of phosphate fertilizers. F. irritates the respiratory tract, causes skin burns. In acute poisoning, irritation of the mucous membranes of the larynx and bronchi, eyes, salivation, nosebleeds occur; in severe cases - pulmonary edema, damage to the central nervous system, etc .; in chronic - conjunctivitis, bronchitis, pneumonia, pneumosclerosis, fluorosis. Characterized by skin lesions such as eczema. First aid: washing the eyes with water, for skin burns - irrigation with 70% alcohol; with inhalation poisoning - inhalation of oxygen. Prevention: compliance with safety regulations, wearing special clothing, regular medical examinations, inclusion of calcium and vitamins in the diet. Preparations containing F. are used in medical practice as antitumor drugs (5-fluorouracil, ftorafur, fluorobenzotef), neuroleptics (trifluperidol, or trisedil, fluorophenazine, triftazine, etc.), antidepressants (fluorocyzine), narcotic (halothane), etc. funds.
Lit.: Ryss I. G., Chemistry of fluorine and its inorganic compounds, M., 1956; Fluorine and its compounds, trans. from English, vol. 1-2, M., 1953-56; Occupational diseases, 3rd ed., M., 1973.
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(first electron)
(according to Pauling)
| F | 9 |
| 18,9984 | |
| 2s 2 2p 5 | |
| Fluorine | |
Chemical properties
The most active non-metal, it violently interacts with almost all substances (rare exceptions are fluoroplasts), and with most of them - with combustion and explosion. The contact of fluorine with hydrogen leads to ignition and explosion even at very low temperatures (up to −252°C). Even water and platinum: uranium for the nuclear industry burn in a fluorine atmosphere.
chlorine trifluoride ClF 3 - a fluorinating agent and a powerful rocket fuel oxidizer
sulfur hexafluoride SF 6 - gaseous insulator in the electrical industry
metal fluorides (such as W and V), which have some beneficial properties
freons are good refrigerants
teflon - chemically inert polymers
sodium hexafluoroaluminate - for the subsequent production of aluminum by electrolysis
various fluorine compounds
Missile technology
Fluorine compounds are widely used in rocket technology as a propellant oxidizer.Application in medicine
Fluorine compounds are widely used in medicine as blood substitutes.
Biological and physiological role
Fluorine is a vital element for the body. In the human body, fluorine is mainly found in tooth enamel as part of fluorapatite - Ca 5 F (PO 4) 3 . With insufficient (less than 0.5 mg / liter of drinking water) or excessive (more than 1 mg / liter) fluoride intake by the body, dental diseases can develop: caries and fluorosis (mottled enamel) and osteosarcoma, respectively.
To prevent caries, it is recommended to use toothpastes with fluoride additives or drink fluoridated water (up to a concentration of 1 mg/l), or use local applications of 1-2% sodium fluoride or stannous fluoride solution. Such actions can reduce the likelihood of caries by 30-50%.
The maximum allowable concentration of bound fluorine in the air of industrial premises is 0.0005 mg/liter.
Additional Information
Fluorine, Fluorum, F(9)
Fluorine (Fluorine, French and German Fluor) was obtained in a free state in 1886, but its compounds have been known for a long time and were widely used in metallurgy and glass production. The first mention of fluorite (CaP,) under the name fluorspar (Fliisspat) dates back to the 16th century. One of the works attributed to the legendary Vasily Valentin mentions stones painted in various colors - fluxes (Fliisse from Latin fluere - flow, pour), which were used as fluxes in the smelting of metals. Agricola and Libavius write about the same. The latter introduces special names for this flux - fluorspar (Flusspat) and mineral melt. Many authors of chemical and technical writings of the 17th and 18th centuries. describe different types of fluorspar. In Russia, these stones were called plavik, spalt, spat; Lomonosov classified these stones as selenites and called them spar or flux (crystal flux). Russian masters, as well as collectors of mineral collections (for example, in the 18th century, Prince P.F. Golitsyn) knew that some types of spars glow in the dark when heated (for example, in hot water). However, even Leibniz in his history of phosphorus (1710) mentions in this connection thermophosphorus (Thermophosphorus).
Apparently, chemists and artisan chemists became acquainted with hydrofluoric acid no later than the 17th century. In 1670, the Nuremberg craftsman Schwanhard used fluorspar mixed with sulfuric acid to etch designs on glass goblets. However, at that time the nature of fluorspar and hydrofluoric acid was completely unknown. It was believed, for example, that silicic acid has an etching effect in the Schwanhard process. This erroneous opinion was eliminated by Scheele, proving that in the interaction of fluorspar with sulfuric acid, silicic acid is obtained as a result of the erosion of the glass retort by the resulting hydrofluoric acid. In addition, Scheele established (1771) that fluorspar is a combination of calcareous earth with a special acid, which was called "Swedish acid".
Lavoisier recognized the hydrofluoric acid radical (radical fluorique) as a simple body and included it in his table of simple bodies. More or less pure hydrofluoric acid was obtained in 1809. Gay-Lussac and Tenard by distilling fluorspar with sulfuric acid in a lead or silver retort. During this operation, both researchers were poisoned. The true nature of hydrofluoric acid was established in 1810 by Ampère. He rejected Lavoisier's opinion that hydrofluoric acid must contain oxygen, and proved the analogy of this acid with hydrochloric acid. Ampère reported his findings to Davy, who shortly before that had established the elemental nature of chlorine. Davy fully agreed with Ampere's arguments and spent a lot of effort on obtaining free fluorine by electrolysis of hydrofluoric acid and in other ways. Taking into account the strong corrosive effect of hydrofluoric acid on glass, as well as on plant and animal tissues, Ampere suggested calling the element contained in it fluorine (Greek - destruction, death, pestilence, plague, etc.). However, Davy did not accept this name and proposed another - fluorine (Fluorine), by analogy with the then name of chlorine - chlorine (Chlorine), both names are still used in English. In Russian, the name given by Ampere has been preserved.
Numerous attempts to isolate free fluorine in the 19th century did not lead to successful results. Only in 1886 did Moissan manage to do this and obtain free fluorine in the form of a yellow-green gas. Since fluorine is an unusually aggressive gas, Moissan had to overcome many difficulties before he found a material suitable for the apparatus in experiments with fluorine. The U-tube for electrolysis of hydrofluoric acid at 55°C (cooled with liquid methyl chloride) was made of platinum with fluorspar plugs. After the chemical and physical properties of free fluorine were investigated, it found wide application. Today, fluorine is one of the most important components in the synthesis of a wide range of organofluorine compounds. Russian literature of the early 19th century. fluorine was called differently: the base of hydrofluoric acid, fluorine (Dvigubsky, 1824), fluorine (Iovsky), fluor (Shcheglov, 1830), fluor, fluorine, fluorine. Hess from 1831 introduced the name fluorine.
