Fluorine

Notable characteristics

Pure fluorine (F₂) is a corrosive pale yellow or brown^[1] hydrofluoric acid.

In aqueous solution, fluorine commonly occurs as the fluoride ion F⁻, although highly diluted HF is such a weak acid that substantial amounts of it are present in any water solution of fluoride at near neutral pH. Other forms are fluoro-complexes, such as [FeF₄]⁻, or H₂F⁺.

Fluorides are compounds that combine fluorine with some positively charged counterpart. They often consist of crystalline ionic salts. Fluorine compounds with metals are among the most stable of salts.

Applications

Chemical uses:

• Xenon difluoride is also used for this last purpose.
• HF) is used to etch glass in light bulbs and other products.
• Fluorine is indirectly used in the production of low friction freon.
• Along with some of its compounds, fluorine is used in the production of pure uranium hexafluoride and in the synthesis of numerous commercial fluorochemicals, including vitally important pharmaceuticals, agrochemical compounds, lubricants, and textiles.
• Fluorochlorohydrocarbons are used extensively in air conditioning and in hydrofluorocarbons) are not on the EPA list of ozone-depleting substances,^[3] and have been widely used as replacements for the chlorine- and bromine-containing fluorocarbons. Hydrofluorocarbons do have a greenhouse effect, but a small one compared with carbon dioxide and methane.
• Sulfur hexafluoride is an extremely inert and nontoxic gas, very useful as an insulator in high-voltage electrical equipment. It does not occur in nature, so it is a useful tracer gas, though as an exceptionally potent greenhouse gas its use in unenclosed systems is inadvisable.
• cryolite), is used in the electrolysis of aluminium.
• In much higher concentrations, sodium fluoride has been used as an insecticide, especially against cockroaches.
• Fluorides have been used in the past to help molten metal flow, hence the name.
• Some researchers including US space scientists in the early 1960s have studied elemental fluorine gas as a possible rocket propellant due to its exceptionally high specific impulse. The experiments failed because fluorine proved difficult to handle, and its combustion products proved extremely toxic and corrosive.
• Teflon surface in baking pans.
• Compounds of fluorine such as fluoropolymers, potassium fluoride and refractive index.

Dental and medical uses:

• Compounds of fluorine, including water fluoridation, though a number of health concerns has sometimes led to controversy.
• Many important agents for general anesthesia such as isoflurane are hydrofluorocarbon derivatives.
• The fluorinated antiinflammatories corticosteroids class of drugs. ^[4]
• aldosterone.
• Fluconazole is a triazole antifungal drug used in the treatment and prevention of superficial and systemic fungal infections.
• broad-spectrum antibiotics.
• sertraline. Because of the difficulty of biological systems in dealing with metabolism of fluorinated molecules, fluorinated antibiotics and antidepressants are among the major fluorinated organics found in treated city sewage and wastewater.

• ¹⁸F, a radioactive isotope that emits positron emission tomography, because its half-life of 110 minutes is long by the standards of positron-emitters.

Compounds

Fluorine forms a variety of very different compounds, owing to its small atomic size and covalent behavior, and on the other hand, its oxidizing ability and extreme flumazenil), fluorine is used to prevent toxication or to delay metabolism.

The fluoride ion is basic, therefore protecting groups is achieved with a fluoride. The fluoride ion is poisonous.

Fluorine as a freely reacting oxidant gives the strongest oxidants known. Chlorine trifluoride, for example, can burn water and sand, both compounds of a weaker oxidant, oxygen.

Fluorine compounds involving noble gases were first synthesised by argon fluorohydride has been prepared, although it is only stable at cryogenic temperatures.

The carbon-fluoride bond is epoxides. When the para position is substituted with fluorine, the aromatic ring is protected and epoxide is no longer produced.

Fluorine can often be substituted for compounds.

 

This element is recovered from fluorapatite.

For a list of fluorine compounds, see here.

History

Fluorine in the form of fluorspar) with concentrated sulfuric acid.

It was eventually realized that hydrofluoric acid contained a previously unknown element. This element was not isolated for many years after this, due to its extreme reactivity; fluorine can only be prepared from its compounds electrolytically, and then it immediately attacks any susceptible materials in the area. Finally, in 1886, elemental fluorine was isolated by Henri Moissan after almost 74 years of continuous effort by other chemists.^[6] It was an effort which cost several researchers their health or even their lives. The derivation of elemental fluorine from hydrofluoric acid is exceptionally dangerous, killing or blinding several scientists who attempted early experiments on this halogen. These men came to be referred to as "fluorine martyrs". For Moissan, it earned him the 1906 Nobel Prize in chemistry (Moissan himself lived to be 54, and it is not clear whether his fluorine work shortened his life).

The first large-scale production of fluorine was needed for the atomic bomb Manhattan project in World War II where the compound fluorocarbon plastic which was not attacked by F₂.

Preparation

Elemental fluorine is prepared industrially by Moissan's original process: electrolysis of anhydrous HF in which KHF₂ has been dissolved to provide enough ions for conduction to take place.

In 1986, when preparing for a conference to celebrate the 100th anniversary of the discovery of fluorine, Karl Christe discovered a purely chemical preparation by reacting together at 150 °C solutions in anhydrous HF of K₂MnF₆ and of SbF₅. The reaction is:

MnF₃ + ½F₂

This is not a practical synthesis, but demonstrates that electrolysis is not essential.

Safety

Main article: fluoride poisoning

Both elemental fluorine and fluoride ions are highly toxic and must be handled with great care and any contact with skin and eyes should be strictly avoided. When it is a free element, fluorine has a characteristic pungent odor that is detectable in concentrations as low as 20 nL/L. All equipment must be passivated before exposure to fluorine.

Contact of exposed skin with hydrofluoric acid solutions poses one of the most extreme and insidious industrial threats—one which is exacerbated by the fact that hydrofluoric acid damages nerves in such a way as to make such burns initially painless. The hydrofluoric acid molecule is capable of rapidly migrating through lipid layers of cells which would ordinarily stop an ionized acid, and the burns are typically deep. HF may react with calcium, permanently damaging the bone. More seriously, reaction with the body's calcium can cause cardiac arrhythmias, followed by cardiac arrest brought on by sudden chemical changes within the body. These cannot always be prevented with local or intravenous injection of calcium salts. Hydrofluoric acid spills over just 2.5% of the body's surface area (about 75 in² or 5 dm²), despite copious immediate washing, have been fatal.^[7] If the patient survives, hydrofluoric acid burns typically produce open wounds of an especially slow-healing nature.

Elemental fluorine is a powerful oxidizer which can cause organic material, combustibles, or other flammable materials to ignite.

Fluorocarbons are generally inert and nontoxic; the electronegativity of fluorine means that a nearby fluorine atom makes a carboxylic acid group very much more reactive. For example, acetic acid.

See also

• Fluorocarbon
• Isotopes of fluorine
• Halide minerals

References

• Los Alamos National Laboratory – Fluorine

1. ^ Theodore Gray. Real visible fluorine. The Wooden Periodic Table.
2. ^ Leonel R Arana, Nuria de Mas, Raymond Schmidt, Aleksander J Franz, Martin A Schmidt and Klavs F Jensen, Isotropic etching of silicon in fluorine gas for MEMS micromachining , J. Micromech. Microeng. 17 , 2007, pp. 384-392.
3. ^ Class I Ozone-Depleting Substances. Ozone Depletion. U.S. Environmental Protection Agency.
4. ^ http://www.emedicine.com/pmr/topic35.htm
5. ^ Fluoride History Discovery of fluorine
6. ^ H. Moissan (1886). "Action d'un courant électrique sur l'acide fluorhydrique anhydre". Comptes rendus hebdomadaires des séances de l'Académie des sciences 102: 1543-1544.
7. ^ [1]

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Hydrogen H2 | Nitrogen N2 | Oxygen O2 | Fluorine F2 | Chlorine Cl2 | Bromine Br2 | Iodine I2 | Astatine At2