Fluorine
| General | |||||||||||||
|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
| Name, Symbol, Number | Fluorine, F, 9 | ||||||||||||
| Series | Halogens | ||||||||||||
| Group, Period, Block | 17 (VIIA), 2 , p | ||||||||||||
| Density, Hardness | 1.696 kg/m3 (273 K), NA | ||||||||||||
| Appearance | pale greenish-yellow gas | ||||||||||||
| Atomic Properties | |||||||||||||
| Atomic weight | 18.9984 amu | ||||||||||||
| Atomic radius (calc.) | 50 (42) pm | ||||||||||||
| Covalent radius | 71 pm | ||||||||||||
| van der Waals radius | 147 pm | ||||||||||||
| Electron configuration | [He]2s2 2p5 | ||||||||||||
| e- 's per energy level | 2, 7 | ||||||||||||
| Oxidation states (Oxide) | -1 (strong acid) | ||||||||||||
| Crystal structure | cubic | ||||||||||||
| Physical Properties | |||||||||||||
| State of matter | Gas (nonmagnetic) | ||||||||||||
| Melting point | 53.53 K (-363.32 °F) | ||||||||||||
| Boiling point | 85.03 K (-306.62 °F) | ||||||||||||
| Molar volume | 11.20 ×10-3 m3/mol | ||||||||||||
| Heat of vaporization | 3.2698 kJ/mol | ||||||||||||
| Heat of fusion | 0.2552 kJ/mol | ||||||||||||
| Vapor pressure | no data | ||||||||||||
| Velocity of sound | no data | ||||||||||||
| Miscellaneous | |||||||||||||
| Electronegativity | 3.98 (Pauling scale) | ||||||||||||
| Specific heat capacity | 824 J/(kg*K) | ||||||||||||
| Electrical conductivity | no data | ||||||||||||
| Thermal conductivity | 0.0279 W/(m*K) | ||||||||||||
| 1st ionization potential | 1681.0 kJ/mol | ||||||||||||
| 2nd ionization potential | 3374.2 kJ/mol | ||||||||||||
| 3rd ionization potential | 6050.4 kJ/mol | ||||||||||||
| 4th ionization potential | 8407.7 kJ/mol | ||||||||||||
| 5th ionization potential | 11022.7 kJ/mol | ||||||||||||
| 6th ionization potential | 15164.1 kJ/mol | ||||||||||||
| 7th ionization potential | 17868 kJ/mol | ||||||||||||
| 8th ionization potential | 92038.1 kJ/mol | ||||||||||||
| 9th ionization potential | 106434.3 kJ/mol | ||||||||||||
| Most Stable Isotopes | |||||||||||||
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| SI units & STP are used except where noted. | |||||||||||||
Fluorine is a chemical element in the periodic table that has the symbol F and atomic number 9. This is a poisonous pale yellow, 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. The fluorine compound fluoride is used to promote dental health.
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 forms compounds with most other elements, including the noble gases 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 always occurs combined and has such an affinity for most elements, especially silicon, that it can neither be prepared nor kept in glass vessels.
Applications
Fluorine is used in the production of low friction plastics such as Teflon, and in halons such as Freon. Other uses;
- Hydrofluoric acid (chemical formula HF) is used to etch glass in light bulbs and other products.
- Monoatomic fluorine is used for plasma ashing in semiconductor manufacturing.
- Along with its compounds, fluorine is used in the production of uranium (from the hexafluoride) and in more than 100 different commercial fluorochemicals, including many high-temperature plastics.
- Fluorochlorohydrocarbons are used extensively in air conditioning and in refrigeration.
- Fluoride has been sold in pill form as a very effective rat and cockroach poison.
- Fluoride is often added to toothpaste and municipal water supplies (see below).
Some researchers have studied elemental fluorine gas a possible rocket propellant due to its exceptionally high specific impulse.
History
Fluorine (L fluere meaning flow or flux) in the form of fluorspar was described in 1529 by Georigius 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 (some experiments would end in tragedy).
This element was not isolated for many years after this due to the fact that when it is separated from one of its compounds 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.
The first commercial production of fluorine was for the atomic bomb Manhattan project in World War II where the compound uranium hexafluoride (UF6) was used to separate isotopes of uranium. This process is still is use today in nuclear power applications.
Compounds
Hydrogen fluoride HF is used to etch glass. It dissolves silicates under formation of gaseous SiF4. Fluorides in drinking water may protect against dental cavities but are known to cause mottling of dental enamel in children; see below. Before World War I high concentrations of fluoride compounds were used as rat poisons, but did not work very well compared to other alternatives.
A certain hypothesis held by some says that fluorine can be substituted for hydrogen when it occurs in organic compounds. Through this mechanism it is thought that fluorine can have a very large number of compounds. Fluorine compounds involving rare gases have been confirmed with fluorides of krypton, radon, and xenon. This element is recovered from fluorite, cryolite, and fluorapatite.
See also: Fluorocarbon
Precautions
Fluorine and HF must be handled with great care and any contact with skin and eyes should be strictly avoided.
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 ppb. It is recommended that the maximum allowable concentration for a daily 8-hour time-weighted exposure is 1 ppm.
However, safe handling procedures enable the transport of liquid fluorine by the ton.
Fluorine and dental health
Fluoride compounds are naturally found in drinking water and some foods. Fluoride ions replace hydroxyl ions in hydroxyapatite in teeth, forming fluorapatite, which is more chemically stable and dissolves at a pH of 4.5, compared to 5.5 pH for hydroxyapatite. This is generally believed to lead to less cavities, since stronger acids are needed to attack the tooth enamel.
For this reason, fluorides are often added to toothpaste. A common view of many dentists and health organisations is that fluorides should also be added to municipal water supplies where the natural level is less than 0.7 ppm in water, to increase the concentration to between 0.7 and 1.2 ppm, with studies showing that incidence of dental cavities is much less where the water is fluoridated up to 1 ppm. Many cities worldwide add fluoride to their water supplies, citing its effectiveness, safety and low cost.
A dissenting view, supported by other scientific studies, suggests that the association of fluorine with dental health is a fallacy, and points to the dangers of developing dental fluorosis and weakened bones from high levels of fluorine intake, over 2 ppm. Some supporters of this viewpoint suggest that associating fluorine with health protects metal smelting companies from the massive lawsuits that began to be filed as much as a century ago for fluoride related damage to livestock and farms in agricultural areas.
The controversy over fluoridations effect on the public health is unlikely to end anytime soon. Currently there is no way to reduce the fluoride emissions from the metal smelting and processing industries to non-toxic levels without shutting them down entirely.