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Fluorine

Oxygen - Fluorine - Neon
F Cl
	Full table

General

Name, Symbol, Number

Fluorine, F, 9

Series

Halogens

Group, Period, Block

17 (VIIA), 2 , p

Density, Hardness

1.696 kg/m³ (273 K), NA

Appearance

pale greenish-yellow gas

Atomic Properties

Atomic weight

18.9984 amu

Atomic radius (calc.)

50 (42) pm

Covalent radius

71 pm

van der Waals radius

147 pm

Electron configuration

[He]2s² 2p⁵

e^- 's per energy level

2, 7

Oxidation states (Oxide)

-1 (strong acid)

Crystal structure

cubic

Physical Properties

State of matter

Gas (nonmagnetic)

Melting point

53.53 K (-363.32 °F)

Boiling point

85.03 K (-306.62 °F)

Molar volume

11.20 ×10^-3 m³/mol

Heat of vaporization

3.2698 kJ/mol

Heat of fusion

0.2552 kJ/mol

Vapor pressure

no data

Speed of sound

no data

Miscellaneous

Electronegativity

3.98 (Pauling scale)

Specific heat capacity

824 J/(kg*K)

Electrical conductivity

no data

Thermal conductivity

0.0279 W/(m*K)

1^st ionization potential

1681.0 kJ/mol

2^nd ionization potential

3374.2 kJ/mol

3^rd ionization potential

6050.4 kJ/mol

4^th ionization potential

8407.7 kJ/mol

5^th ionization potential

11022.7 kJ/mol

6^th ionization potential

15164.1 kJ/mol

7^th ionization potential

17868 kJ/mol

8^th ionization potential

92038.1 kJ/mol

9^th ionization potential

106434.3 kJ/mol

Most Stable Isotopes

iso	NA	half-life	DM	DE M eV	DP
¹⁹F	100%	F is stable with 10 neutrons

SI units & STP are used except where noted.

Fluorine is a chemical element in the periodic table that has the symbol F and atomic number 9. It is a poisonous pale yellow, univalent gaseous halogen that is the most chemically reactive and electronegative of all the elements. In its pure form, it is highly dangerous, causing severe chemical burns[?] on contact with skin.

Table of contents

1 Notable Characteristics

3 History

Notable Characteristics

Pure fluorine is a corrosive pale yellow gas that is a powerful oxidizing agent. It is the most reactive and electronegative of all the elements, and forms compounds with most other elements, including the noble gases xenon and radon. Even in dark, cool conditions, fluorine reacts explosively with hydrogen. In a jet of fluorine gas, glass, metals, water and other substances burn with a bright flame. It always occurs combined and has such an affinity for most elements, especially silicon, that it can neither be prepared nor kept in glass vessels.

In solution, fluorine commonly occurs as the fluoride ion F-. Fluorides are compounds that combine this fluoride ion with some positively charged radical.

Applications

Fluorine is used in the production of low friction plastics such as Teflon, and in halons such as Freon. Other uses:

Hydrofluoric acid[?] (chemical formula HF) is used to etch glass in light bulbs and other products.
Monoatomic fluorine is used for plasma ashing in semiconductor manufacturing.
Along with its compounds, fluorine is used in the production of uranium (from the hexafluoride) and in more than 100 different commercial fluorochemicals, including many high-temperature plastics.
Fluorochlorohydrocarbons are used extensively in air conditioning and in refrigeration. Chlorofluorocarbons have been banned for these applications because they are suspected to contribute to the ozone hole. Both of these classes of compounds are potent greenhouse gases.
Sodium fluoride[?] has been used as an insecticide, especially against cockroaches.
Some other fluorides are often added to toothpaste and (somewhat controversially) to municipal water supplies to prevent dental cavities.

Some researchers have studied elemental fluorine gas a possible rocket propellant[?] due to its exceptionally high specific impulse.

History

Fluorine (L fluere meaning flow or flux) in the form of fluorspar was described in 1529 by Georigius Agricola[?] for its use as a flux, which is a substance that is used to promote the fusion of metals or minerals. In 1670 Schwandhard[?] 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 (some experiments would end in tragedy).

This element was not isolated for many years after this due to the fact that when it is separated from one of its compounds it immediately attacks the remaining materials of the compound. Finally in 1886 fluorine was isolated by Henri Moissan[?] after almost 74 years of continuous effort.

The first commercial production of fluorine was for the atomic bomb Manhattan project in World War II where the compound uranium hexafluoride[?] (UF₆) was used to separate isotopes of uranium. This process is still is use today in nuclear power applications.

Compounds

A certain hypothesis held by some says that fluorine can be substituted for hydrogen when it occurs in organic compounds. Through this mechanism it is thought that fluorine can have a very large number of compounds. Fluorine compounds involving rare gases have been confirmed with fluorides of krypton, radon, and xenon. This element is recovered from fluorite, cryolite, and fluorapatite[?].