Fluorine is a non- metallic chemical element, chemical Fluorine Symbol F. Atomic number 9. Fluorine is one of the halogen elements, belonging to the periodic system VIIA, and is located in the second period in the periodic table. Fluorine Uses
The element of fluorine is F2, which is a pale yellow, highly toxic gas. Fluorine gas is very corrosive, chemically active, and is one of the most oxidizing substances. It can even react with some inert gases under certain conditions.
Fluorine is a key element in specialty plastics, rubber, and freezers (chlorofluorocarbons). Due to the special chemical nature of fluorine, fluorine chemistry plays an important role in the history of chemical development.
Physical Properties of Fluorine
Fluorine is a pale yellow gas under standard conditions and a yellow liquid when liquefied. It became a colorless liquid at -252 °C.
Due to the special nature of the fluoro-chemical, resulting in the difficulty of measuring the physical properties is large, is not very high accuracy of some data, the following data using the latest time reference data or similar data in the valid digital more digits By.
Atomic radius: 71 pm (FF), 64 pm (FC)
Ion radius: 133pm
Density: 1.696g / L (273.15K, 0 °C)
Melting point: -219.66 °C
Heat of fusion: 510.36 ± 2.1 J · mol-1.
Boiling point: -188.12 °C.
Gasification heat: 6543.69±12.55J·mol -1 (84.71K, 9.81kPa).
Critical temperature: 144K.
Critical pressure: 55atm.
Thermal conductivity: W / (m · K) 27.7.
Electronic layer layout: [He]2s 2 2p 5
Main oxidation state: F(-I), F(0)
Electronegativity: 3.98 (Pauling scale), 4.10 (Alai – Luo Zhou scale), 3.91 (Malken scale)
Chemical bond energy (kJ/mol): FF: 159; FH: 569, FO: 190; FN: 272; FC: 456, FB:644, F-Al:582.
Elemental dissociation energy: 157.7kJ·mol -1
F – Water and energy: -506.3kJ · mol -1
Standard entropy: F: 158.6 J·mol -1 · K -1 , F2 : 202.5 J·mol -1 · K -1
Standard electrode potential: E ∅ ( F2 /HF) = 3.053V, E (F2/F – ) = 2.87V.
Fluorine is the most non-metallic element of the known element, which makes it have no positive oxidation state.
The ground state valence electron layer structure of fluorine is 2s2 2p5, and fluorine has a very small atomic radius, and thus has a strong electron-promoting tendency and strong oxidizing property, and is one of the strongest oxidants known.
Fluorine intercalates of fluorine include ClF, ClF3, BrF3, IF6 and the like. Fluorine Uses
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Reaction with elemental
The chemical reaction between hydrogen and fluorine is extremely intense, and it can be explosively combined with hydrogen to form hydrogen fluoride even at a low temperature of -250 °C.
Not only hydrogen, but fluorine can react with all elements except O, N, He, Ne, Ar, and Kr to form the highest fluoride.
Except for the metal fluoride with the highest valence state and a few pure perfluoro organic compounds, almost all compounds can react with fluorine.
Even perfluorinated organic compounds can be burned in fluorine gas if they are contaminated with combustibles.
The reaction of most organic compounds with fluorine will explode, and the reaction of carbon or most hydrocarbons with excess fluorine will produce carbon tetrafluoride and a small amount of tetrafluoroethylene or hexafluoropropane.
Due to the strongly oxidizing properties of fluorine, fluorine can even be directly neutralized with hydrazine. The product may be XeF2, XeF4, XeF6 depending on the reaction conditions. Fluorine Uses
Generally, since nitrogen is inert to fluorine, it can be used as a diluent gas for gas-phase reactions. Nitrogen and fluorine can be combined into NF3 by glow discharge. When fluorine reacts with copper, nickel or magnesium, a dense fluoride protective film forms on the metal surface to prevent the reaction, so the fluorine gas can be stored in a container made of these materials
Reaction with a compound
The reaction between fluorine gas and water is complicated, and the main reaction is:
Generating hydrogen fluoride and oxygen, side reactions to produce small amounts of hydrogen peroxide, oxygen difluoride, and ozone.
The reaction between fluorine and ammonia is:
However, if the ammonia gas is excessive, in addition to the formation of NF3, N2 F4, HNF2 and N2 F2 may be formed. If the above reaction is too intense, only nitrogen may be obtained:
Fluorine reacts with nitric oxide to form fluorinated nitrosyl :
Fluoro anhydrous sodium nitrite under heating conditions, or so dioxide combustion fluorine can be obtained fluorinated nitroxyl:
Fluorine reacts with dilute hydrogen azide to form azide fluoride:
In general, oxygen does not react with fluorine, but there are two known oxyfluorides, namely OF2 and O2F2. By introducing fluorine into a 2% sodium hydroxide solution, OF2 can be obtained:
Hydrofluoric acid (HOF) can be obtained by passing fluorine gas through the surface of ice water.
Usually, the reaction of fluorine with organic matter is too intense to obtain a simple organic fluoride, but if a certain proportion of fluorine is diluted, an organic addition reaction or an organic substitution reaction similar to chlorine and bromine may occur.
Silica can be burned in fluorine gas to produce oxygen:
Fluorine can also replace other halogens from many halogen-containing compounds.
(X is other halogens).
Some special properties of F can be explained in the following aspects:
- F has the largest electronegativity.
- Standard electrode potential F. 2/ F.
- F has a small atomic radius, so the repulsive force of the lone pair of electrons in the fluorine molecule is quite large, and fluorine has no available d orbital, so the d-p π bond cannot be formed so that the FF bond can be very small.
- In fluoride, the chemical bond formed by fluorine and other elements is very strong. In the ionic halide, the fluoride energy U is the largest in the covalent halide, the general fluoride Δ f Gm is the most negative
Some fluorine-containing compounds have extremely strong Lewis acidity, such as BF3, SbF5, etc., and SbF5 is dissolved in liquid hydrogen fluoride to obtain fluoroantimonic acid, which is a superacid.
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- UF6(g) can be obtained by the strong oxidizing property of fluorine. Using 238 the UF. 6 and235 the UF.6 different diffusion rates, to separate the isotopes of uranium.
- Used to synthesize coolants such as Freon. Fluorine Uses
- Used in the manufacture of fluorinating reagents (fluorene difluoride, etc.) and fluxing agents (ice crystals, etc.) in metal smelting, etc.
- ClF 3and BrF3 can be used as oxidants for rocket fuels.
- Used in the manufacture of insecticides and fire extinguishing agents.
- Fluorocarbons can be used as temporary substitutes for blood.
- Fluoride glass (containing ZrF4, BaF2, NaF) is 100 times more transparent than conventional oxide glass and does not darken even under strong radiation, optical fiber made of fluoride glass fiber is more effective than SiO2optical fiber. The effect is hundreds of times.
- Fluorine-containing plastics and fluoroelastomers have exceptionally good properties For oxyfluoride blowing and manufacturing of various fluorides.
- Fluorine element is also added to the toothpaste as fluoride toothpaste, fluorideand teeth of dibasic calcium phosphate react harder and the less soluble fluorapatite.
Fluorine is one of the most widely distributed elements in nature. The stock of fluorine in the earth’s crust is 6.5 × 10 -2% The number of sorts of the amount of existence is 13. Fluorine in nature is mainly composed of fluorite (CaF2), cryolite (Na3[AlF6]) and fluorapatite (Ca10(PO4)6F2).
German scientists have for the first time confirmed the presence of fluorine in nature, Florian Kraus of the Technical University of Munich, Germany, and other collaborators, such as Jorn Schmedt Auf Der Günne of Munich Ludwig-Maximilian University, the first in-situ confirmation that fluorine gas is vomiting The culprit of the stench of the stone.
They collected a large sample of vomit in the bean near the area that caused nausea and vomiting of the miners and then analyzed them with a solid nuclear magnetic resonance spectrometer. This technology can detect the fluorine gas in situ without breaking the sample.
Human Body Distribution
Fluoride contains about 2 grams to 3 grams in normal adults, and the human body contains about 2.6 grams of fluorine, mainly distributed in bones and teeth About 90% of the fluoride is accumulated in the two, and the blood contains 0.04 μg to 0.4 μg per ml.
The fluorine required by the human body is mainly from drinking water. A daily intake of more than 4 mg of the human body can cause poisoning and damage health.
The isotope in which fluorine is abundant in nature is only 19F. There are 18 known fluorine isotopes, and only 19F is stable. 18F is a good source of positrons and is often used in the synthesis of positron emission tomography (PET) tracers.
At present, the most commonly used tracer in the clinic, fluorine-18 deoxyglucose (18F-FDG), is a tracer containing fluorine-18.
Safety of fluorine compounds
Fluorine compounds are harmful to the human body, and a small amount of fluorine (within 150 mg) can cause a series of illnesses. A large amount of fluoride enters the body and causes acute poisoning.
Due to the different amounts of inhalation, various conditions such as anorexia, nausea, abdominal pain, stomach ulcers, cramps, and even death can occur. Fluorine Uses
If the amount of poisoning is insufficient, the human body can quickly recover from fluorosis. Especially when intravenous or intramuscular calcium gluconate is used, about 90% of the fluorine can be quickly eliminated, and the remaining fluorine needs time to be removed.
Frequent exposure to fluoride can easily lead to hardening and embrittlement of bones, and the symptoms of brittle fracture and breakage. In some areas, too much fluoride in drinking water can easily lead to fluorosis.
Traces of fluoride are good for preventing dental caries. If the fluoride content in water is less than 0.5ppm, the incidence of dental caries will reach 70% to 90%.
However, if the amount of fluorine in drinking water exceeds 1 ppm, the teeth will gradually become spots and become brittle.
When the fluoride content in drinking water exceeds 4 ppm, people are susceptible to fluorosis and lead to bone marrow malformation. The way to reduce the fluoride content in drinking water is to boil drinking water.
Insoluble fluoride is low in toxicity and non-irritating to the skin, but if a large amount of dust is inhaled, it is easily absorbed by the body and chronically poisoned. Soluble fluoride can be quickly excreted after being absorbed.
Swallowing 5~10g at a time will cause gastrointestinal bleeding and death. Acidic fluorides, such as hydrofluoric acid and fluoroboric acid, can violently corrode the skin, causing redness and spread at the contact site, resulting in ulcers that are difficult to heal. Fluorine Uses
Mononuclear fluorine, hydrogen fluoride, and other gases are irritating to the human eye and nose. Excessive inhalation can cause severe bronchitis and pulmonary edema, leading to death.
The most serious and dangerous thing for people exposed to fluoride is the exposure of the face and skin to fluoride and fluoride.
Therefore, the use of fluorine and fluoride must comply with the operating procedures, and there are reliable safety measures, including operating equipment, rubber gloves, covered protective masks and gas masks with acid-proof gas.
The workplace should have good ventilation facilities, and explosion-proof devices should be provided for items with high reactivity.
When burning with hydrogen fluoride and other fluorides, treat the burned area in time with a large amount of water and then apply it with glycerol magnesium oxide. The best method is to immediately inject calcium gluconate into the affected area so that the fluorine is fixed as insoluble fluoride.
In addition, the use of hexafluoride rinse is a good method to deal with hydrofluoric acid accidents. The principle of hexafluoride is to reduce the number of hydrogen ions and fluoride ions in the human body through neutralization reaction and coordination.
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Some organic fluorides are very toxic. Among them, for the fluorine-containing carboxylic acid, the structural formula is F(CH2)nCOOH. If n is an odd number, the organic substance is extremely toxic, and when n is even, the toxicity is small or even non-toxic. Fluorine Uses
At the end of November 2013, there was an event of death caused by co-delivery of methyl fluoroacetate, in which the fluoro acetic acid portion of methyl fluoroacetate was n=1.
The toxicity of fluorine to insects is similar to that of hydrogen fluoride. For plants, the toxicity is similar to that of sulfur dioxide. Fluorine compounds cause white spots or brown spots on the leaves and veins of plants.