Fluorine

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9 oxygen ← fluorine → neon - ↑ F ↓ Cl Periodic table - Extended periodic table
General
Name, symbol, number	fluorine, F, 9
Element category	halogens
Group, period, block	17, 2, p
Appearance	Yellowish brown gas
Standard atomic weight	18.9984032(5) g·mol−1
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)
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
Specific heat capacity	(25 °C) (F2) 31.304 J·mol−1·K−1
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−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
Miscellaneous
Magnetic ordering	nonmagnetic
Thermal conductivity	(300 K) 27.7 m W·m−1·K−1
CAS registry number	7782-41-4
Selected isotopes
Main article: Isotopes of fluorine iso NA half-life DM DE (MeV) DP 18F syn 109.77 min ε 1.656 18O 19F 100% 19F is stable with 10 neutrons
References
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Fluorine (from Latin: fluorum, meaning "to flow") is the chemical element with the symbol F and atomic number 9. Fluorine forms a single bond with itself in elemental form, resulting in the diatomic F₂ molecule. F₂ is a supremely reactive, poisonous, pale, yellowish brown gas. Elemental fluorine is the most chemically reactive and electronegative of all the elements. For example, it will readily 'burn' hydrocarbons at room temperature, in contrast to the combustion of hydrocarbons by oxygen, which requires an input of energy with a spark. Therefore, molecular fluorine is highly dangerous, more so than other halogens such as the poisonous chlorine gas.

Fluorine's highest electronegativity and small atomic radius give unique properties to many of its compounds. For example, the enrichment of ²³⁵U, the principal nuclear fuel, relies on the volatility of UF₆. Also, the carbon–fluorine bond is considered as the strongest bond in organic chemistry. This contributes to the stability and persistence of fluorocarbon-based organofluorine compounds, such as PTFE and PFOS. The carbon–fluorine bond also sharply increases the efficacy of many pharmaceuticals and results in the strength of many superacids.

Contents 1 Characteristics 2 Applications 2.1 Dental and medical uses 3 Chemistry of fluorine 4 Production 5 History 6 Biological role 7 Precautions 7.1 Elemental fluorine 7.2 Fluoride ion 7.3 Hydrogen fluoride and hydrofluoric acid 7.4 Organofluorines 8 See also 9 Notes 10 References 11 External links

[edit] Characteristics

F₂ is a corrosive pale yellow or brown^[1] gas that is a powerful oxidizing agent. It is the most reactive and most electronegative of all the elements (4.0), and readily forms compounds with most other elements. It has an oxidation number -1, except when bonded to another fluorine in F₂ which gives it an oxidation number of 0. Fluorine even combines with the noble gases argon, krypton, xenon, and radon. Even in dark, cool conditions, fluorine reacts explosively with hydrogen. The reaction with hydrogen occurs even at extremely low temperatures, using liquid hydrogen and solid fluorine. It is so reactive that metals, and even water, as well as other substances, burn with a bright flame in a jet of fluorine gas. In moist air it reacts with water to form also-dangerous hydrofluoric acid.

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.

Hydrogen fluoride is a weak acid when dissolved in water. Consequently, fluorides of alkali metals produce basic solutions. For example, a 1 M solution of NaF in water has a pH of 8.59 compared to a 1 M solution of NaOH, a strong base, which has a pH of 14.00.^[2]

[edit] Applications

Industrial use:

• Atomic fluorine and molecular fluorine are used for plasma etching in semiconductor manufacturing, flat panel display production and MEMS (microelectromechanical systems) fabrication.^[3] Xenon difluoride is also used for this last purpose.
• Hydrofluoric acid (chemical formula HF) is used to etch glass in light bulbs and other products.
• Tetrafluoroethylene and perfluorooctanoic acid (PFOA) are directly used in the production of low friction plastics such as Teflon (or polytetrafluoroethylene).
• Fluorine is used indirectly the production of 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 contain chlorine or bromine but contain only fluorine, carbon and hydrogen (called hydrofluorocarbons) are not on the United States Environmental Protection Agency list of ozone-depleting substances,^[4] 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.
• Sodium hexafluoroaluminate (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.
• Compounds of fluorine such as fluoropolymers, potassium fluoride and cryolite are utilized in applications such as anti-reflective coatings and dichroic mirrors on account of their unusually low refractive index.

[edit] Dental and medical uses

• Compounds of fluorine, including sodium fluoride (NaF), stannous fluoride (SnF₂) and sodium MFP, in addition to a fluorine ion are used in toothpaste to prevent dental cavities. These compounds are also added to municipal water supplies, a process called water fluoridation, though a number of health concerns has sometimes led to controversy.
• Many important agents for general anesthesia such as sevoflurane, desflurane, and isoflurane are hydrofluorocarbon derivatives.
• The fluorinated antiinflammatories dexamethasone and triamcinolone are among the most potent of the synthetic corticosteroids class of drugs.^[5]
• Fludrocortisone ("Florinef") is one of the most common mineralocorticoids, a class of drugs which mimics the actions of aldosterone.
• 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.
• SSRI antidepressants, except in a few instances, are fluorinated molecules. These include citalopram,escitalopram oxalate,fluoxetine,fluvoxamine maleate, and paroxetine. A notable exception is 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 positrons, is often used in positron emission tomography, because its half-life of 110 minutes is long by the standards of positron-emitters.

[edit] Chemistry of fluorine

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 electronegativity. For example, hydrofluoric acid is extremely dangerous, while in synthetic drugs incorporating an aromatic ring (e.g. flumazenil), fluorine is used to help prevent toxication or to delay metabolism.

The fluoride ion is basic, therefore hydrofluoric acid is a weak acid in water solution. However, water is not an inert solvent in this case: when less basic solvents such as anhydrous acetic acid are used, hydrofluoric acid is the strongest of the hydrohalogenic acids. Also, owing to the basicity of the fluoride ion, soluble fluorides give basic water solutions. The fluoride ion is a Lewis base, and has a high affinity to certain elements such as calcium and silicon. For example, deprotection of silicon 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 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.

The carbon-fluoride bond is covalent and very stable. The use of a fluorocarbon polymer, poly(tetrafluoroethene) or Teflon, is an example: it is thermostable and waterproof enough to be used in frying pans. Organofluorines may be safely used in applications such as drugs, without the risk of release of toxic fluoride. In synthetic drugs, toxication can be prevented. For example, an aromatic ring is useful but presents a safety problem: enzymes in the body metabolize some of them into poisonous epoxides. When the para position is substituted with fluorine, the aromatic ring is protected and epoxide is no longer produced.

The substitution of hydrogen for fluorine in organic compounds offers a very large number of compounds. An estimated fifth of pharmaceutical compounds and 30% of agrochemical compounds contain fluorine.^[6] The -CF₃ and -OCF₃ moieties provide further variation, and more recently the -SF₅ group.^[7]

Flourine might form a molecule with unbihexium (Ubh), the theorized island of stability. Specifically, the stable monofluoride UbhF may result from a bonding interaction between the 5g orbital on Ubh and the 2p orbital on fluorine.^[8]

This element is recovered from fluorite, cryolite, and fluorapatite.

For a list of fluorine compounds, see here.

[edit] Production

Industrial fluorine production starts with fluorspar (CaF₂), which is heated with sulfuric acid (H₂SO₄) to produce anhydrous hydrogen fluoride (HF). The hydrogen fluoride is added to potassium fluoride (KF) to make potassium bifluoride (KHF₂). Electrolysis of potassium bifluoride produces fluorine gas at the anode, and hydrogen gas at the cathode. This is essentially the same method employed by Moissan in 1886; the use of potassium bifluoride rather than hydrogen fluoride itself aids electrolysis by greatly increasing the conductivity.

CaF₂ + H₂SO₄ → 2 HF + CaSO₄

HF + KF → KHF₂

2 KHF₂ → 2 KF + H₂ + F₂

In 1986, when preparing for a conference to celebrate the 100th anniversary of the discovery of fluorine, Karl Christe discovered a purely chemical preparation involving the reaction of solutions in anhydrous HF, K₂MnF₆, and SbF₅ at 150 °C:^{[citation needed]}

K₂MnF₆ + 2SbF₅ → 2KSbF₆ + MnF₃ + ½F₂

Though not a practical synthesis on the large scale, it demonstrates that electrolysis is not essential.

[edit] 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,^[9] 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. Carl Wilhelm Scheele and many later researchers, including Humphry Davy, Caroline Menard, 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 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.^[10] 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".^{[citation needed]} 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 gaseous 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 due to the F₂ was eventually solved by electrolytically coating all UF₆ carrying piping with nickel metal, which resists fluorine's attack. Joints and flexible parts were made from teflon, then a very recently discovered fluorocarbon plastic which was not attacked by F₂.

[edit] Biological role

While F₂ is too reactive to have any natural biological role, the fluorine atom is incorporated into compounds with biological roles. However, organofluorine compounds are rare in nature. The most notable example is fluoroacetate, which functions as a plant defence against herbivores in at least 40 plants in Australia, Brazil and Africa.^[11] The enzyme adenosyl-fluoride synthase catalyzes the formation of 5'-deoxy-5'-fluoroadenosine. Additionally, fluoride might have a natural role in preventing tooth decay.^[12]

[edit] Precautions

[edit] Elemental fluorine

Elemental fluorine (fluorine gas) is a highly toxic, corrosive oxidant, which can cause organic material, combustibles, or other flammable materials to ignite. It must be handled with great care and any contact with skin and eyes should be strictly avoided. Fluorine gas has a characteristic pungent odor that is detectable in concentrations as low as 20 ppb. As it is so reactive, all materials of construction must be carefully selected. All metal surfaces must be passivated before exposure to fluorine.

[edit] Fluoride ion

Main article: fluoride poisoning

Fluoride ions are also toxic and must also be handled with great care and any contact with skin and eyes should be strictly avoided.

[edit] Hydrogen fluoride and hydrofluoric acid

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 HF molecule is a weaker acid which is significantly non-dissociated in water, and the intact molecule is capable of rapidly migrating through lipid layers of cells which would ordinarily stop an ion or partly ionized acid, and the burns it produces are typically deep. HF may react with calcium, permanently damaging the bone . More seriously, HF reaction with the body's calcium inside cells can cause cardiac arrhythmias, followed by cardiac arrest brought on by sudden chemical changes within the body (hypocalcaemia). 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.^[13] If the patient survives, hydrofluoric acid burns typically produce open wounds of an especially slow-healing nature.

Anhydrous hydrogen fluoride will rapidly form hydrofluoric acid on contact with moisture; its physiological effects are then the same.

[edit] Organofluorines

Organofluorines are naturally rare compounds. They can be nontoxic (perflubron and perfluorodecalin) or highly toxic (perfluoroisobutylene and fluoroacetic acid). Many pharmacuticals are organofluorines, such as the anti-cancer fluorouracil. Perfluorooctanesulfonic acid (PFOS) is a proposed persistent organic pollutant.

[edit] See also

• Fluorocarbon
• Isotopes of fluorine
• Halide minerals
• Water fluoridation

[edit] Notes

1. ^ FG33R Theodore Gray. "Real visible fluorine". The Wooden Periodic Table.
2. ^ "pKa's of Inorganic and Oxo-Acids". Evans Group. Retrieved on 2008-11-29.
3. ^ 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.
4. ^ "Class I Ozone-Depleting Substances". Ozone Depletion. U.S. Environmental Protection Agency.
5. ^ eMedicine - Corticosteroid-Induced Myopathy : Article by Steve S Lim, MD
6. ^ "Fluorine's treasure trove". ICIS news (2006-10-02). Retrieved on 2008-11-29.
7. ^ "Selective Direct Fluorination". F2 Chemicals Limited. Retrieved on 2008-11-29.
8. ^ Jacoby, Mitch (2006). "As-yet-unsynthesized superheavy atom should form a stable diatomic molecule with fluorine". Chemical & Engineering News 84 (10): 19, http://pubs.acs.org/cen/news/84/i10/8410notw9.html. Retrieved on 14 January 2008. 
9. ^ Fluoride History Discovery of fluorine
10. ^ 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, http://gallica.bnf.fr/ark:/12148/bpt6k3058f/f1541.chemindefer. 
11. ^ Proudfoot AT, Bradberry SM, Vale JA (2006). "Sodium fluoroacetate poisoning". Toxicol Rev 25 (4): 213–9. doi:10.2165/00139709-200625040-00002. PMID 17288493. 
12. ^ Yeung CA (2008). "A systematic review of the efficacy and safety of fluoridation". Evid Based Dent 9 (2): 39–43. doi:10.1038/sj.ebd.6400578. PMID 18584000, http://dx.doi.org/10.1038/sj.ebd.6400578. 
13. ^ Hydrogen fluoride (PIM 268)

[edit] References

• Los Alamos National Laboratory – Fluorine

[edit] External links

Wikimedia Commons has media related to: Fluorine

• WebElements.com – Fluorine
• It's Elemental – Fluorine
• Picture of liquid fluorine – chemie-master.de
• Chemsoc.org

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