Fluorine

Distinguished from fluorene and fluorone.

Fluorine, ₉F

Liquid fluorine (F2 at extremely low temperature)
Fluorine
Pronunciation	/ˈflʊəriːn/ /ˈflɔːriːn/ (FLOR-een)
Allotropes	alpha, beta (see Allotropes of fluorine)
Appearance	gas: very pale yellow liquid: bright yellow solid: alpha is opaque, beta is transparent
Standard atomic weight Ar°(F)
	18.998403162±0.000000005 18.998±0.001 (abridged)
Fluorine in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson – ↑ F ↓ Cl oxygen ← fluorine → neon
Atomic number (Z)	9
Group	group 17 (halogens)
Period	period 2
Block	p-block
Electron configuration	[He] 2s2 2p5
Electrons per shell	2, 7
Physical properties
Phase at STP	gas
Melting point	(F2) 53.48 K ​(−219.67 °C, ​−363.41 °F)
Boiling point	(F2) 85.03 K ​(−188.11 °C, ​−306.60 °F)
Density (at STP)	1.696 g/L
when liquid (at b.p.)	1.505 g/cm3
Triple point	53.48 K, ​.252 kPa
Critical point	144.41 K, 5.1724 MPa
Heat of vaporization	6.51 kJ/mol
Molar heat capacity	Cp: 31 J/(mol·K) (at 21.1 °C) Cv: 23 J/(mol·K) (at 21.1 °C)
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, 0 (oxidizes oxygen)
Electronegativity	Pauling scale: 3.98
Ionization energies	1st: 1681 kJ/mol 2nd: 3374 kJ/mol 3rd: 6147 kJ/mol (more)
Covalent radius	64 pm
Van der Waals radius	135 pm
Spectral lines of fluorine
Other properties
Natural occurrence	primordial
Crystal structure	​cubic
Thermal conductivity	0.02591 W/(m⋅K)
Magnetic ordering	diamagnetic (−1.2×10−4)
CAS Number	7782-41-4
History
Naming	after the mineral fluorite, itself named after Latin fluo (to flow, in smelting)
Discovery	André-Marie Ampère (1810)
First isolation	Henri Moissan (June 26, 1886)
Named by	Humphry Davy André-Marie Ampère
Isotopes of fluorine
Main isotopes Decay abun­dance half-life (t1/2) mode pro­duct 18F trace 109.734 min β+ 18O 19F 100% stable

Fluorine (IPA: /ˈflʊərɪːn, -ɔːrɪːn/, 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, yellowish brown gas, with chemical formula F₂. Like other halogens, molecular fluorine is highly dangerous; it causes severe chemical burns on contact with skin.

Fluorine's large electronegativity and small atomic radius gives it interesting bonding characteristics, particularly in conjunction with carbon. See covalent radius of fluorine.

Notable characteristics

Pure fluorine (F₂) is a corrosive pale yellow or brown 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. Its oxidation number is a constant, at -1. Fluorine even combines with the noble gases, krypton, xenon, and radon. Even in dark, cool conditions, fluorine reacts explosively with hydrogen. 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. It is far too reactive to be found in elemental form. 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 highly diluted 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. The carbon-fluoride bond is covalent and stable, so that organofluorines are inert, in contrast to other organohalogens.

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 , 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. 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 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.

Safety

Main article: fluoride poisoning

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.

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 hydrofluoric acid molecule is capable of rapidly migrating through lipid layers of cells which would ordinarily stop an ionized acid, and the burns are typically deep. HF may react with calcium, permanently damaging the bone. More seriously, reaction with the body's calcium can cause cardiac arrhythmias, followed by cardiac arrest brought on by sudden chemical changes within the body. 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. If the patient survives, hydrofluoric acid burns typically produce open wounds of an especially slow-healing nature.

Elemental fluorine is a powerful oxidizer which can cause organic material, combustibles, or other flammable materials to ignite.

Fluorocarbons are generally inert and nontoxic; the electronegativity of fluorine means that a nearby fluorine atom makes a carboxylic acid group very much more reactive. For example, trifluoroacetic acid is 100,000 times stronger than acetic acid.

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, when 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:

Template:PotassiumTemplate:ManganeseF₆ + 2Template:AntimonyF₅ → 2Template:PotassiumTemplate:AntimonyF₆ + Template:ManganeseF₃ + ½F₂

This is not a practical synthesis, but demonstrates that electrolysis is not essential.

Compounds

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 prevent toxication.

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.

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.

Fluorite (CaF₂) crystals

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

For a list of fluorine compounds, see here.

Applications

Chemical uses:

• Atomic fluorine and molecular fluorine are used for plasma etching in semiconductor manufacturing, flat panel display production and MEMS (microelectromechanical systems) fabrication. 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.
• 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 fluorine, 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, very useful as an insulator in high-voltage electrical equipment. It doesn't occur in nature so 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.
• Polytetrafluoroethylene, also known as the non-stick Teflon surface in baking pans.
• 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:

• 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 number of health concerns has sometimes led to controversy.
• Many important agents for general anesthesia 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.
• 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.

See also

• Fluorocarbon
• Isotopes of fluorine
• Halide minerals

References

• Los Alamos National Laboratory – Fluorine

External links

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