Fluorine

Distinguish from fluorene and fluorone.

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Fluorine (from L. fluere, meaning "to flow"), is the chemical element in the periodic table that has the symbol F and atomic number 9. Atomic fluorine is univalent and is the most chemically reactive and electronegative of all the elements. In its pure form, it is a poisonous, pale, yellow-green gas, with chemical formula F₂. Like other halogens, molecular fluorine is highly dangerous; it causes severe chemical burns on contact with skin.

Contents

1 Notable characteristics
2 Applications
3 History
4 Precautions
5 Preparation
6 Compounds
7 See also
8 References
9 External links

Notable characteristics

Pure fluorine (F₂, since fluorine is diatomic) is a corrosive pale yellow gas that is a powerful oxidizing agent. It is the most reactive and electronegative of all the elements, and readily forms compounds with most other elements. Fluorine even combines with the noble gases, krypton, xenon, and radon. Even in dark, cool conditions, fluorine reacts explosively with hydrogen. It is so reactive that glass, metals, and even water, as well as other substances, burn with a bright flame in a jet of fluorine gas. It is far too reactive to be found in elemental form and has such an affinity for most elements, including silicon, that it can neither be prepared nor should be kept in glass vessels. In moist air it reacts with water to form the equally dangerous hydrofluoric acid .

In aqueous solution, fluorine commonly occurs as the fluoride ion F^-. Other forms are fluoro-complexes, such as [FeF₄]^-, or H₂F⁺.

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

Applications

Atomic fluorine and molecular fluorine are used for plasma etching in semiconductor manufacturing, flat panel display production and MEMS fabrication. Other uses:

Hydrofluoric acid (chemical formula HF) is used to etch glass in light bulbs and other products.
Fluorine is indirectly used in the production of low friction plastics such as Teflon, and in halons such as Freon.
Along with some of its compounds, fluorine is used in the production of pure uranium from 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 refrigeration. Chlorofluorocarbons have been banned for these applications because they contribute to the ozone hole.
Sulfur hexafluoride is an extremely inert and nontoxic gas, and a member of a class of compounds that are potent greenhouse gases.
Many important agents for general anaesthesia such as sevoflurane, desflurane, and isoflurane are fluorohydrocarbon derivatives.
Sodium hexafluoroaluminate (cryolite), is used in the electrolysis of aluminium.
Compounds of fluorine, including sodium fluoride, are used in toothpaste to prevent dental cavities. These compounds are also added to municipal water supplies, a process called water fluoridation, though a combination of health concerns and urban legends has sometimes led to controversy.
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.
¹⁸F, a radioactive isotope that emits positrons, is often used in positron emission tomography because of its half-life of 110 minutes.

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 was so hard to handle.

History

Fluorine in the form of fluorspar (also called fluorite) (calcium fluoride) was described in 1530 by Georgius Agricola for its use as a flux Fluoride History [Discovery of fluorine] , which is a substance that is used to promote the fusion of metals or minerals. In 1670 Schwanhard found that glass was etched when it was exposed to fluorspar that was treated with acid. Karl Scheele and many later researchers, including Humphry Davy, Gay-Lussac, Antoine Lavoisier, and Louis Thenard all would experiment with hydrofluoric acid, easily obtained by treating calcium fluoride (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 - it is separated from its compounds only with difficulty and then it immediately attacks the remaining materials of the compound. Finally, in 1886, fluorine was isolated by Henri Moissan after almost 74 years of continuous effort. It was an effort which cost several researchers their health or even their lives, and for Moissan, it earned him the 1906 Nobel Prize in chemistry.

The first large scale production of fluorine was needed for the atomic bomb Manhattan project in World War II where the compound uranium hexafluoride (UF₆) was used to separate the ²³⁵U and ²³⁸U isotopes of uranium. Today both the gaseous diffusion process and the gas centrifuge process use gaseous (UF₆) to produce enriched uranium for nuclear power applications.

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."

Precautions

Both fluorine and HF must be handled with great care and any contact with skin and eyes should be strictly avoided. All equipment must be passivated before exposure to fluorine.

Contact with exposed skin may result in the HF molecule rapidly migrating through the skin and flesh into the bone where it reacts with calcium permanently damaging the bone, followed by cardiac arrest brought on by sudden chemical changes within the body.

Both elemental fluorine and fluoride ions are highly toxic. When it is a free element, fluorine has a characteristic pungent odor that is detectable in concentrations as low as 20 nL/L. It is recommended that the maximum allowable concentration for a daily 8-hour time-weighted exposure is 1 µL/L (part per million by volume) (lower than, for example, hydrogen cyanide).

Fluorine is a powerful oxidizer which can cause organic material, combustibles, or other flammable materials to ignite. However, safe handling procedures enable the transport of liquid fluorine by the ton.

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, preparing for a conference to celebrate the 100th aniversary 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. This is not a practical synthesis, but demonstrates that electrolysis is not essential.

Compounds

Fluorite (CaF₂) crystals

Fluorine can often be substituted for hydrogen when it occurs in organic compounds. Through this mechanism, fluorine can have a very large number of compounds. Fluorine compounds involving noble gases were first synthesised by Neil Bartlett in 1962 - xenon hexafluoroplatinate, XePtF₆, being the first. Fluorides of krypton and radon have also been prepared. Also argon fluorohydride has been prepared, although it is only stable at cryogenic temperatures. This element is recovered from fluorite, cryolite, and fluorapatite.

References

[Los Alamos National Laboratory – Fluorine]

External links

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