Name: Fluorine
Symbol: F
Atomic Number: 9
Atomic Weight: 18.998403
Family:  Halogens
CAS RN: 7782-41-4
Description: A pale yellow gas.
State (25C): Gas
Oxidation states: -1

Molar Volume: 17.1 cm3/mole
Valence Electrons: 2p5

Boiling Point:  85.1K, 188.05C, 306.49F
Melting Point:
53.63K, 219.52C, 363.14F
Electrons Energy Level: 2, 7
Isotopes: 12 + 1 stable
Heat of Vaporization: 3.2698 kJ/mol
Heat of Fusion: 0.2552 kJ/mol
Density: 1.696 g/L @ 273K & 1atm
Specific Heat: 0.82 J/gK
Atomic Radius: 0.57
Ionic Radius: 1.33
Electronegativity: 3.98 (Pauling); 4.1 (Allrod Rochow)
Fluorine (from the Latin fluere, for "flow") was isolated by Henri Moissan in 1886.  It is a highly toxic and reactive greenish-yellow gas at room temperature.  Because of its reactivity, elemental Fluorine is never found in nature and no other chemical element can displace Fluorine from its compounds.

In the late 1600's minerals which we now know contain Fluorine were used in etching glass.   The discovery of the element was prompted by the search for the chemical substance which was able to attack glass (it is HF, a weak acid). The early history of the isolation and work with Fluorine and Hydrogen Fluoride is filled with accidents since both are extremely dangerous. Eventually, electrolysis of a mixture of KF and HF (carefully ensuring that the resulting Hydrogen and Fluorine would not come in contact) in a platinum apparatus yielded the element.

Compounds of Fluorine are present in fluoridated toothpaste and in many municipal water systems where they help to prevent tooth decay. And, of course, fluorocarbons such as Teflon have made a major impact on life in the 20th. century.






Pure Fluorine (F2) is a corrosive pale yellow or brown gas that is a powerful oxidizing agent.  It is the most reactive and electronegative of all the elements (4.0), and readily forms compounds with most other elements. 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 glass, 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 and has such an affinity for most elements, including Silicon, that it can neither be prepared nor be kept in ordinary glass vessels. Instead, it must be kept in specialized quartz tubes lined with a very thin layer of fluorocarbons. In moist air it reacts with water to form also-dangerous Hydrofluoric Acid, HF.

In aqueous solution, Fluorine commonly occurs as the Fluoride Ion F-, although 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.


Fluorine is the most reactive of all elements and no chemical substance is capable of freeing Fluorine from any of its compounds.  For this reason, Fluorine does not occur free in nature and was extremely difficult for scientists to isolate.  The first recorded use of a Fluorine compound dates to around 1670 to a set of instructions for etching glass that called for Bohemian Emerald (CaF2).  Chemists attempted to identify the material that was capable of etching glass and George Gore was able to produce a small amount of fluorine through an electrolytic process in 1869.  Unknown to Gore, Fluorine gas, F2, explosively combines with Hydrogen gas, H2.  That is exactly what happened in Gore's experiment when the Fluorine gas that formed on one electrode combined with the Hydrogen gas that formed on the other electrode.  Ferdinand Frederic Henri Moissan, a French chemist, was the first to successfully isolate fluorine in 1886.   He did this through the electrolysis of Potassium Fluoride (KF) and Hydrofluoric Acid (HF).  He also completely isolated the fluorine gas from the Hydrogen gas and he built his electrolysis device completely from Platinum.  His work was so impressive that he was awarded the Nobel Prize for chemistry in 1906.  Today, Fluorine is still produced through the electrolysis of Potassium Fluoride and Hydrofluoric Acid as well as through the electrolysis of molten Potassium acid Fluoride (KHF2).

Fluorine in the form of fluorspar (also called fluorite) (Calcium Fluoride, CaF2) 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, Gay-Lussac, Antoine Lavoisier, and Louis Thenard all would experiment with Hydrofluoric Acid, easily obtained by treating Calcium Fluoride ( fluorspar) with concentrated Sulfuric Acid, H2SO4.

1s2 2s2p5

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, HF, 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).

2s2 2p5

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 isotope 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 Fluorine-containing plastic which was not attacked by F2.


40px-Skull_and_crossbones.svg.jpg (1420 bytes) 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 L/L.  All equipment must be passivated before exposure to Fluorine.

Contact of exposed skin with HF solutions posses one of the most extreme and insidious industrial threats—one which is exacerbated by the fact that HF damages nerves in such a way as to make such burns initially painless.  The HF 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.  HF spills over just 2.5% of the body's surface area, despite copious immediate washing, have been fatal (this corresponds with an area of about 9 in2 or 23 cm2).  If the patient survives, HF 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.


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, 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:

2K2MnF6 + 4SbF5 rarrow.gif (63 bytes) 4KSbF6 + MnF2 + F2

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


Fluorine is added to city water supplies in the proportion of about one part per million to help prevent tooth decay.  Sodium Fluoride (NaF), Stannous(II) Fluoride (SnF2) and Sodium monofluorophosphate (Na2PO3F) are all Fluorine compounds added to toothpaste, also to help prevent tooth decay. Hydrofluoric Acid (HF) is used to etch glass, including most of the glass used in light bulbs.  Uranium Hexafluoride (UF6) is used to separate isotopes of Uranium.  Crystals of Calcium Fluoride (CaF2), also known as Fluorite and Fluorspar, are used to make lenses to focus infrared light.  Fluorine joins with Carbon to form a class of compounds known as fluorocarbons. Some of these compounds, such as Dichlorodifluoromethane (CF2Cl2), were widely used in air conditioning and refrigeration systems and in aerosol spray cans, but have been phased out due to the damage they were causing to the earth's Ozone layer.

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.  Fluorine compounds involving noble gases were first synthesized 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.  Fluorine is also recovered from Fluorite, Cryolite, and Fluorapatite.


Atomic Fluorine and molecular Fluorine are used for plasma etching in semiconductor manufacturing, flat panel display production and MEMS fabrication. Other uses:

Dental and Medical Uses:


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.

atom.gif (700 bytes)

Isotope Atomic Mass Half-Life
F14 14.036  
F15 15.018 1 MeV
F16 16.0115 40 keV
F17 17.0021 64.49 seconds
F18 18.0009 109.77 minutes
F19 18.9984 Stable
F20 20. 11 seconds
F21 21. 4.158 seconds
F22 22.003 4.23 seconds
F23 23.0036 2.23 seconds
F24 24.0081 0.34 seconds
F25 25.0121 59 ms
F26 26.02  
F27 27.027 >200 ns
F28 28.036  
F29 29.043 >200 ns

(L. and F. fluere, flow or flux) In 1529, Georigius Agricola described the use of fluorspar as a flux, and as early as 1670 Schwandhard found that glass was etched when exposed to fluorspar treated with acid. Scheele and many later investigators, including Davy, Gay-Lussac, Lavoisier, and Thenard, experimented with hydrofluoric acid, some experiments ending in tragedy. The element was finally isolated in 186 by Moissan after nearly 74 years of continuous effort. Fluorine is the most electronegative and reactive of all elements. It is a pale yellow, corrosive gas, which reacts with practically all organic and inorganic substances. Finely divided metals, glass, ceramics, carbon, and even water burn in fluorine with a bright flame. Until World War II, there was no commercial production of elemental fluorine. The nuclear bomb project and nuclear energy applications, however, made it necessary to produce large quantities. Safe handling techniques have now been developed and it is possible at present to transport liquid fluorine by the ton. Fluorine and its compounds are used in producing uranium (from the hexafluoride) and more than 100 commercial fluorochemicals, including many well-known high-temperature plastics. Hydrofluoric acid is extensively used for etching the glass of light bulbs, etc. Fluorochlorohydrocarbons are extensively used in air conditioning and refrigeration. It has been suggested that fluorine can be substituted for hydrogen wherever it occurs in organic compounds, which could lead to an astronomical number of new fluorine compounds. The presence of fluorine as a soluble fluoride in drinking water to the extent of 2 ppm may cause mottled enamel in teeth, when used by children acquiring permanent teeth; in smaller amounts, however, fluorides are said to be beneficial and used in water supplies to prevent dental cavities. elemental fluorine has been studied as a rocket propellant as it has an exceptionally high specific impulse value. Compounds of fluorine with rare gases have now been confirmed. Fluorides of xenon, radon, and krypton are among those known. Elemental fluorine and the fluoride ion are highly toxic. The free element has a characteristic pungent odor, detectable in concentrations as low as 20 ppb, which is below the safe working level. The recommended maximum allowable concentration for a daily 8-hour time-weighted exposure is 1 ppm.

Source: CRC Handbook of Chemistry and Physics, 1913-1995. David R. Lide, Editor in Chief. Author: C.R. Hammond

atom.gif (700 bytes)

Fluorine Data

Atomic Radius (): 0.57
Atomic Volume cm3/mol : 17.1cm3/mol
Covalent Radius: 0.72
Crystal Structure: Cubic
Ionic Radius: 1.33

Chemical Properties

Electrochemical Equivalents: 0.70883 g/amp-hr
Electron Work Function: unknown
Electronegativity: 3.98 (Pauling); 4.1 (Allrod Rochow)
Heat of Fusion: 0.2552 kJ/mol
Incompatibilities: Water, nitric acid, oxidizers, organic compounds
First Ionization Potential: 17.422
Second Ionization Potential: 34.97
Third Ionization Potential: 62.707
Valence Electron Potential: -10.1
Ionization Energy (eV): 17.423 eV

Physical Properties

Atomic Mass Average: 18.9984
Boiling Point: 85.1K, 188.05C, 306.49F
Melting Point: 53.63K, 219.52C, 363.14F
Heat of Vaporization: 3.2698 kJ/mol
Coefficient of Lineal Thermal Expansion/K-1: N/A
Electrical Conductivity: unknown
Thermal Conductivity: 0.000279 W/cmK
Density: 1.696 g/L @ 273K & 1atm
Enthalpy of Atomization: 79.08 kJ/mole @ 25C
Enthalpy of Fusion: 0.26 kJ/mole
Enthalpy of Vaporization: 3.31 kJ/mole
Molar Volume: 17.1 cm3/mole
Optical Refractive Index: 1.000195
Relative Gas Density (Air=1): 1.31 
Specific Heat: 0.82 J/gK
Vapor Pressure: unknown
Estimated Crustal Abundance: 5.85102 milligrams per kilogram
Estimated Oceanic Abundance: 1.3 milligrams per liter