Fluorine

From Freepedia

9 oxygen ← fluorine → neon
-
↑
F
↓
Cl
Image:F-TableImage.png
General
Name, Symbol, Number fluorine, F, 9
Chemical series halogens
Group, Period, Block 17, 2, p
Appearance pale greenish-yellow gas
Image:F,9.jpg
Atomic mass 18.9984032(5) g/mol
Electron configuration 1s2 2s2 2p5
Electrons per shell 2, 7
Physical properties
Phase gas
Density (0 °C, 101.325 kPa)
1.7 g/L
Melting point 53.53 K
(-219.62 °C, -363.32 °F)
Boiling point 85.03 K
(-188.12 °C, -306.62 °F)
Heat of fusion (F2) 0.510 kJ/mol
Heat of vaporization (F2) 6.62 kJ/mol
Heat capacity (25 °C) (F2)
31.304 J/(mol·K)
Vapor pressure
P/Pa 1 10 100 1 k 10 k 100 k
at T/K 38 44 50 58 69 85
Atomic properties
Crystal structure cubic
Oxidation states −1
(strongly acidic oxide)
Electronegativity 3.98 (Pauling scale)
Ionization energies
(more) 		1st: 1681.0 kJ/mol
2nd: 3374.2 kJ/mol
3rd: 6050.4 kJ/mol
Atomic radius 50 pm
Atomic radius (calc.) 42 pm
Covalent radius 71 pm
Van der Waals radius 147 pm
Miscellaneous
Magnetic ordering nonmagnetic
Thermal conductivity (300 K) 27.7 mW/(m·K)
CAS registry number 7782-41-4
Notable isotopes
Main article: Isotopes of fluorine
iso NA half-life DM DE (MeV) DP
19F 100% F is stable with 10 neutrons
References

Fluorine (from L. fluere, meaning "to flow"), is the chemical element in the periodic table that has the symbol F and atomic number 9. It is a poisonous pale yellow-green, univalent gaseous halogen that is the most chemically reactive and electronegative of all the elements. In its pure form, it is highly dangerous, causing severe chemical burns on contact with skin.

Contents

Notable characteristics

Pure fluorine 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. In a jet of fluorine gas, glass, metals, water and other substances burn with a bright flame. 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

Fluorine is used in the production of low friction plastics such as Teflon, and in halons such as Freon. Other uses:

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. Experiments failed since fluorine was so hard to handle.

History

Fluorine in the form of fluorspar (calcium fluoride) was described in 1529 by Georgius Agricola for its use as a flux, which is a substance that is used to promote the fusion of metals or minerals. In 1670 Schwandhard 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 U-235 and U-238 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.

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.

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 & micro;L/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 150C solutions in anhydrous HF of K2MnF6 and of SbF5. This is not a practical synthesis, but demonstrates that electrolysis is not essential.


Compounds

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, althought it is only stable at cryogenic temperatures. This element is recovered from fluorite, cryolite, and fluorapatite. See also: Fluorocarbon

References

External links

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