|Name, Symbol, Number|| fluorine, F, 9
|Group, Period, Block||17, 2, p|
|Appearance|| Yellowish brown gas |
|Atomic mass||18.9984032(5) g/mol|
|Electron configuration||1s2 2s2 2p5|
|Electrons per shell||2, 7|
|Density|| (0 °C, 101.325 kPa)|
|Melting point|| 53.53 K|
(-219.62 °C, -363.32 °F)
|Boiling point|| 85.03 K|
(-188.12 °C, -306.62 °F)
|Critical point||144.13 K, 5.172 MPa|
|Heat of fusion||(F2) 0.510 kJ·mol−1|
|Heat of vaporization||(F2) 6.62 kJ·mol−1|
|Heat capacity|| (25 °C) (F2)|
|Oxidation states|| −1|
(strongly acidic oxide)
|Electronegativity||3.98 (Pauling scale)|
| Ionization energies|
|1st: 1681.0 kJ·mol−1|
|2nd: 3374.2 kJ·mol−1|
|3rd: 6050.4 kJ·mol−1|
|Atomic radius||50 pm|
|Atomic radius (calc.)||42 pm|
|Covalent radius|| 71 pm|
(see covalent radius of fluorine)
|Van der Waals radius||147 pm|
|Thermal conductivity||(300 K) 27.7 mW·m−1·K−1|
|CAS registry number||7782-41-4|
Fluorine (IPA: /ˈflʊərɪːn, -ɔːrɪːn/, from Latin fluere, meaning "to flow"), is the chemical element with 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 elementally isolated (pure) form, fluorine 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.
Fluorine's relatively large electronegativity and small atomic radius give it interesting bonding characteristics, particularly in conjunction with carbon.
Pure fluorine (F2) is a corrosive pale yellow gas that is a powerful oxidizing agent. Since it is one electron away from 8 in its outer cloud (which is the magic number for stability), 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 farts and burps. 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 be kept in glass vessels. In moist air it reacts with water to form also-dangerous hydrofluoric acid.
In aqueous solution, fluorine commonly occurs as the fluoride ion F-, although 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.
Fluorine in the form of fluorspar (also called fluorite) (calcium fluoride) was described in 1530 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 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; fluorine is separated from its compounds only electrolytically, and then it immediately attacks any succeptable materials in the area. Finally, in 1886, elemental fluorine was isolated by Henri Moissan after almost 74 years of continuous effort by other chemists. 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 uranium hexafluoride (UF6) was needed as a gasseous carrier of uranium 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. In the Manhattan project, it was found that elemental fluorine was present whenever UF6 was, due to the spontaneous decomposition of this compound into UF4 and F2. The corrosion problem was eventually solved by making piping which was electrolytically coated with nickel metal, which resists fluorine's attack. Joints and flexible parts where made from Teflon, a newly discovered flourine-containing plastic at that time.
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. Its MAC-value is 1 1 µL/L. All equipment must be passivated before exposure to fluorine. For more information consult an MSDS.
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. This may be followed by cardiac arrest brought on by sudden chemical changes within the body, or large, difficult healing open wounds.
Elemental fluorine is a powerful oxidizer which can cause organic material, combustibles, or other flammable materials to ignite.
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 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: 2Template:PotassiumTemplate:ManganeseF6 + 4Template:AntimonyF5 → 4Template:PotassiumTemplate:AntimonyF6 + 2Template:ManganeseF3 + F2 This is not a practical synthesis, but demonstrates that electrolysis is not essential.
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.
For a list of fluorine compounds, see here.
- 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 ozone destruction and the ozone hole. Interestingly, since it is chlorine and bromine radicals which harm the ozone layer, not fluorine, compounds which do not have chlorine or bromine and contain only flourine, carbon and hydrogen (called hydrofluorocarbons), are not on the E.P.A. list of ozone-depleting substances , 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, and a member of a class of compounds that are potent greenhouse gases. rawr
- Many important agents for general anaesthesia such as sevoflurane, desflurane, and isoflurane are hydrofluorocarbon derivatives.
- Fluconazole is a triazole antifungal drug used in the treatment and prevention of superficial and systemic fungal infections.
- Fluoroquinolones are a family of broad-spectrum antibiotics.
- Sodium hexafluoroaluminate (cryolite), is used in the electrolysis of aluminium.
- Compounds of fluorine, including sodium fluoride (NaF), stannous fluoride (SnF2) and sodium MFP, 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 its half-life of 110 minutes is long by the standards of positron-emitters.
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 of course proved extremely toxic and corrosive.
- WebElements.com – Fluorine
- It's Elemental – Fluorine
- Picture of liquid fluorine – chemie-master.de
- Periodic Table of Elements
- Discovery of fluorine
|This page uses content from Wikipedia. The original article was at Fluorine. The list of authors can be seen in the page history. As with Chemistry, the text of Wikipedia is available under the GNU Free Documentation License.|