fluorine

Oxygen ← Fluorine → Neon - ↑ F ↓ Cl 9F Periodic table
Appearance
Yellowish brown gas
General properties
Name, symbol, number	Fluorine, F, 9
Element category	halogen
Group, period, block	17, 2, p
Standard atomic weight	18.9984032(5) g·mol−1
Electron configuration	1s2 2s2 2p5
Electrons per shell	2, 7 (Image)
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
Oxidation states	−1 (Weaklyacidic 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
Covalent radius	57±3 pm (see covalent radius of fluorine)
Van der Waals radius	147 pm
Miscellanea
Crystal structure	cubic
Magnetic ordering	nonmagnetic
Thermal conductivity	(300 K) 27.7 m W·m−1·K−1
CAS registry number	7782-41-4
Most stable isotopes
Main article: Isotopes of Fluorine
iso NA half-life DM DE (MeV) DP 18F syn 109.77 min β+ (97%) 0.64 18O ε (3%) 1.656 18O 19F 100% 19F is stable with 10 neutrons
This box: view • talk • edit

Fluorine is the chemical element with atomic number 9, represented by the symbol F. 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 one of the strongest bonds in organic chemistry. This contributes to the stability and persistence of fluoroalkane based organofluorine compounds, such as PTFE/(Teflon) and PFOS. The carbon–fluorine bond's inductive effects result in the strength of many fluorinated acids, such as triflic acid and trifluoroacetic acid. Drugs are often fluorinated at biologically reactive positions, to prevent their metabolism and prolong their half-lives.

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 on the classic Pauling scale (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 the 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, but is still very corrosive and attacks glass. 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]

Isotopes

Although fluorine (F) has multiple isotopes, only one of these isotopes (F-19) is stable, and the others have short half-lives and are not found in nature. Fluorine is thus a mononuclidic element.

The nuclide ¹⁸F is the radionuclide of fluorine with the longest half life (about 110 minutes), and commercially is an important source of positrons, finding its major use in positron emission tomography scanning.

Applications

Elemental fluorine, F₂, is mainly used for the production of two compounds of commercial interest, uranium hexafluoride and sulfur hexafluoride.^[3]

Industrial use of fluorine-containing compounds

Fluorite (CaF₂) crystals

• Atomic fluorine and molecular fluorine are used for plasma etching in semiconductor manufacturing, flat panel display production and MEMS (microelectromechanical systems) fabrication.^[4] 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 in 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,^[5] 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.
• 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, which derives from Latin verb fluere, meaning to flow.^{[clarification needed]}
• 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.

Dental and medical uses

• Inorganic compounds of fluoride, including sodium fluoride (NaF), stannous fluoride (SnF₂) and sodium MFP, are used in toothpaste to prevent dental cavities. These or related compounds are also added to some municipal water supplies, a process called water fluoridation, although the practice has remained controversial since its beginnings in 1945.
• 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.^[6]
• 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.

• Compounds containing ¹⁸F, a radioactive isotope that emits positrons, are often used in positron emission tomography, because its half-life of 110 minutes is long by the standards of positron-emitters. One such species is fluorodeoxyglucose.

Chemistry of fluorine

Fluorine forms a variety of very different compounds, owing to its small atomic size and covalent behavior. Elemental fluorine is a dangerously powerful oxidant, reflecting the extreme electronegativity of fluorine. Hydrofluoric acid is extremely dangerous, whereas in synthetic drugs incorporating an aromatic ring (e.g. flumazenil), fluorine is used to help prevent toxication or to delay metabolism^{[citation needed]}.

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.

The reactivity of fluorine toward the noble gas xenon was first reported by Neil Bartlett in 1962. Fluorides of krypton and radon have also been prepared. Argon fluorohydride has been observed 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 fluorine for hydrogen in organic compounds offers a very large number of compounds. An estimated fifth of pharmaceutical compounds and 30% of agrochemical compounds contain fluorine.^[7] The -CF₃ and -OCF₃ moieties provide further variation, and more recently the -SF₅ group.^[8]

Production

Fluorine cell room at F2 Chemicals Ltd, Preston, UK

Industrial production of fluorine entails the electrolysis of hydrogen fluoride in the presence of potassium fluoride. This method is based on the pioneering studies by Moissan (see below). Fluorine gas forms at the anode, and hydrogen gas at the cathode. Under these conditions, the potassium fluoride (KF) converts to potassium bifluoride (KHF₂), which is the actual electrolyte. This potassium bifluoride aids electrolysis by greatly increasing the electrical conductivity of the solution.

HF + KF → KHF₂

2 KHF₂ → 2 KF + H₂ + F₂

The HF required for the electrolysis is obtained as a byproduct of the production of phosphoric acid. Phosphate-containing minerals contain significant amounts of calcium fluorides, such as fluorite. Upon treatment with sulfuric acid, these minerals release hydrogen fluoride:

CaF₂ + H₂SO₄ → 2 HF + CaSO₄

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:^[9]

2 K₂MnF₆ + 4 SbF₅ → 4 KSbF₆ + 2 MnF₃ + F₂

Though not a practical synthesis on the large scale, this report demonstrates that electrolysis is not the sole route to the element.

History

The mineral fluorspar (also called fluorite), consisting mainly of calcium fluoride, was described in 1530 by Georgius Agricola for its use as a flux.^[10] Fluxes are used to promote the fusion of metals or minerals. The etymology of the element's name reflects its history: Fluorine is pronounced /ˈflʊəriːn/, /ˈflʊərɨn/, or commonly /ˈflɔr-/; from Latin: fluere, meaning "to flow". In 1670 Schwanhard found that glass was etched when it was exposed to fluorspar that had been 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 fluorite with concentrated sulfuric acid.

Owing to its extreme reactivity, elemental fluorine was not isolated until many years after the characterization of fluorite. Progress in isolating elemental fluorine was slowed because it could only be prepared electrolytically and even then under stringent conditions since the gas attacks many materials. In 1886, the isolation of elemental fluorine was reported by Henri Moissan after almost 74 years of effort by other chemists.^[11] The generation of elemental fluorine from hydrofluoric acid is exceptionally dangerous, killing or blinding several scientists who attempted early experiments on this halogen. These individuals came to be referred to as "fluorine martyrs".^[12] For Moissan, it earned him the 1906 Nobel Prize in chemistry.^[13]

The first large-scale production of fluorine was undertaken in support of the Manhattan project, where the compound uranium hexafluoride (UF₆) had been selected as the form of uranium that would allow separation of its ²³⁵U and ²³⁸U isotopes. 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 UF₆ decomposed 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 forms a nickel difluoride that is not attacked by fluorine. Joints and flexible parts were made from teflon, then a very recently discovered fluorocarbon plastic which is also not attacked by F₂.

Biological role

Though F₂ is too reactive to have any natural biological role, fluorine is incorporated into compounds with biological activity. Naturally occurring organofluorine compounds are rare, the most notable example is fluoroacetate, which functions as a plant defence against herbivores in at least 40 plants in Australia, Brazil and Africa.^[14] The enzyme adenosyl-fluoride synthase catalyzes the formation of 5'-deoxy-5'-fluoroadenosine. Fluorine is not an essential nutrient, but its importance in preventing tooth decay is well-recognized.^[15] The effect is predominantly topical, although prior to 1981 it was considered primarily systemic (occurring through ingestion).^[16]

Precautions

Elemental fluorine

Elemental fluorine (fluorine gas) is a highly toxic, corrosive oxidant, which can cause ignition of organic material. 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 and metal surfaces must be passivated.

Fluoride ion

Fluoride ions are toxic: the lethal dose of sodium fluoride for a 70 kg human is estimated at 5–10 g.^[17]

Hydrogen fluoride and hydrofluoric acid

Hydrogen fluoride and hydrofluoric acid are dangerous, far more so than the related hydrochloric acid, because undissociated molecular HF penetrates the skin and biological membranes, causing deep and painless burns. The free fluoride, once released from HF in dissociation, also is capable of chelating calcium ion to the point of causing death by cardiac dysrhythmia. Burns with areas larger than 25 square inches (160 cm²) have the potential to cause serious systemic toxicity.^[18]

Organofluorines

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

See also

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

References

External links

Wikimedia Commons has media related to: Fluorine

Look up fluorine in Wiktionary, the free dictionary.

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