Fluorine

Fluorine, ₀₀F
	Liquid fluorine (at extremely low temperatures)
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)
	18.998403163(6)
Fluorine in the periodic table
	oxygen ← fluorine → neon
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	53.48 K (−219.67 °C, −363.41 °F)
Boiling point	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, 90 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)
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
Isotopes of fluorine v; e;
	Category: Fluorine; view; talk; edit; \| references

Fluorine (symbol F) is a chemical element that is very poisonous. Its atomic number (which is the number of protons in it) is 9, and its atomic mass is 19. It is part of the Group 7 (halogens) on the periodic table of elements.

Properties

Fluorine is a light yellow diatomic gas. It is very reactive gas, which exists as diatomic molecules. It is the most reactive element. Fluorine has a very high attraction for electrons because it is missing one. This makes it the most powerful oxidizing agent. It can rip electrons from water (making oxygen) and ignite propane on contact. It does not need a spark. Metals can catch on fire when placed in a stream of fluorine. After it is reduced by reacting with other things, it forms the stable fluoride ion. Fluorine is very poisonous. Fluorine bonds very strongly with carbon. It can react with the unreactive noble gases. It explodes when mixed with hydrogen. The melting point of fluorine is -363.33°F (-219.62°C), the boiling point is -306.62°F (-188.12°C).

Chemical compounds

Chemical compounds containing fluorine ions are called fluorides. Fluorine only exists in one oxidation state: -1.

Aluminium fluoride
Antimony trifluoride
Antimony pentafluoride
Arsenic trifluoride
Arsenic pentafluoride
Bismuth(III) fluoride
Bismuth(V) fluoride
Bromine trifluoride
Bromine pentafluoride
Chlorine monofluoride
Chlorine trifluoride
Cobalt(II) fluoride
Cobalt(III) fluoride
Disulfur decafluoride
Hydrofluoric acid, a solution of hydrogen fluoride in water
Hydrogen fluoride
Iodine trifluoride
Iodine pentafluoride
Iodine heptafluoride
Manganese(II) fluoride
Manganese(III) fluoride
Manganese(IV) fluoride
Potassium fluoride
Selenium tetrafluoride
Selenium hexafluoride
Silver(I) fluoride, brown-yellow
Silver(II) fluoride, highly reactive, white or gray
Sodium aluminium fluoride, cryolite
Sodium fluoride
Sulfur hexafluoride
Sulfur tetrafluoride
Tellurium(IV) fluoride
Tellurium(VI) fluoride
Thallium(I) fluoride
Thallium(III) fluoride
Tin(II) fluoride
Tin(IV) fluoride
Zinc fluoride

Occurrence

Fluorine is not found as an element on the earth becase it is too reactive. Several fluorides are found in the earth, though. When calcium phosphate is reacted with sulfuric acid to make phosphoric acid, some hydrofluoric acid is produced. Also, fluorite can be reacted with sulfuric acid to make hydrofluoric acid. Fluorite naturally occurs on the earths' crust in rocks, coal and clay.

Preparation

Fluorine is normally made by electrolysis. Hydrogen fluoride is dissolved in potassium fluoride. This mixture is melted and an electric current is passed through it. This is electrolysis. Hydrogen is produced at one side and fluorine at the other side. If the sides are not separated, the cell may explode.

Someone made fluorine in 1986 without using electrolysis. They produced manganese(IV) fluoride by using various chemical compounds, which released fluorine gas.

Uses

Fluorine is used to enrich uranium for nuclear weapons. It is also used to make sulfur hexafluoride. Sulfur hexafluoride is used to propel stuff out of an aerosol can. It is also used to make integrated circuits. Fluorine compounds have many uses. Fluoride ions are in fluorine compounds. Fluoride ions can be in toothpaste. Some are used in nonstick coatings. Freons contain fluorine.

Safety

Fluorine as an element is extremely reactive and toxic. It can react with almost everything, even glass. Fluorine is also poisonous.

Fluoride ions are somewhat toxic. If too much toothpaste containing fluoride is eaten then fluoride poisoning may occur. Fluoride is not reactive, though.

Related pages

Sources

↑ "Standard Atomic Weights: Fluorine". CIAAW. 2013.
↑ ^2.0 ^2.1 ^2.2 ^2.3 ^2.4 Jaccaud et al. 2000, p. 381.
↑ ^3.0 ^3.1 ^3.2 ^3.3 Haynes 2011, p. 4.121.
↑ ^4.0 ^4.1 Jaccaud et al. 2000, p. 382.
↑ ^5.0 ^5.1 ^5.2 Compressed Gas Association 1999, p. 365.
↑ Himmel, D.; Riedel, S. (2007). "After 20 Years, Theoretical Evidence That 'AuF₇' Is Actually AuF₅·F₂". Inorganic Chemistry. 46 (13). 5338–5342. doi:10.1021/ic700431s.
↑ Dean 1999, p. 4.6.
↑ Dean 1999, p. 4.35.
↑ Matsui 2006, p. 257.
↑ Yaws & Braker 2001, p. 385.
↑ Mackay, Mackay & Henderson 2002, p. 72.
↑ Cheng et al. 1999.
↑ ^13.0 ^13.1 Chisté & Bé 2011.

[1] "Standard Atomic Weights: Fluorine". CIAAW. 2013.

[FOOTNOTEJaccaud_et_al.2000381-2] 2.0 ^2.1 ^2.2 ^2.3 ^2.4 Jaccaud et al. 2000, p. 381.

[FOOTNOTEHaynes20114.121-3] 3.0 ^3.1 ^3.2 ^3.3 Haynes 2011, p. 4.121.

[FOOTNOTEJaccaud_et_al.2000382-4] 4.0 ^4.1 Jaccaud et al. 2000, p. 382.

[FOOTNOTECompressed_Gas_Association1999365-5] 5.0 ^5.1 ^5.2 Compressed Gas Association 1999, p. 365.

[6] Himmel, D.; Riedel, S. (2007). "After 20 Years, Theoretical Evidence That 'AuF₇' Is Actually AuF₅·F₂". Inorganic Chemistry. 46 (13). 5338–5342. doi:10.1021/ic700431s.

[FOOTNOTEDean19994.6-7] Dean 1999, p. 4.6.

[FOOTNOTEDean19994.35-8] Dean 1999, p. 4.35.

[FOOTNOTEMatsui2006257-9] Matsui 2006, p. 257.

[FOOTNOTEYawsBraker2001385-10] Yaws & Braker 2001, p. 385.

[FOOTNOTEMackayMackayHenderson200272-11] Mackay, Mackay & Henderson 2002, p. 72.

[FOOTNOTECheng_et_al.1999-12] Cheng et al. 1999.

[FOOTNOTEChistéBé2011-13] 13.0 ^13.1 Chisté & Bé 2011.

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