Fluorine Uses

Fluorine (F) is a member of Group 17 in the periodic table. It is the lightest halogen of the group that includes Chlorine (Cl), Bromine (Br), Iodine (I), and Astatine (At).

The history of fluorine dates back to 1529, when it was described in the use of fluorspar, as a flux, by Georgius Agricola. The name fluorine is derived from the Latin word fluore for "flow" or "flux". After 141 years, in the 1670s, a German glassworker found that hydrofluoric acid could help in glass etching.

However, it was only in the year 1764 and 1771 that the conceptualization of fluoric acid, when heated in glass with sulfuric acid was recorded and identified by chemists, Andreas Sigismund Marggraf (Germany) and Carl Wilhelm Scheele (Sweden), respectively.

Andreas Sigismund Marggraf (Germany) and Carl Wilhelm Scheele (Sweden), respectively. Later in the year 1886, with nearly 74 years of continuous effort, the element fluorine was isolated from its compounds by a French chemist Ferdinand Frederick Henri Moissan -- for which, he later received the Nobel Prize for chemistry in the year 1906.

Uses
Fluorine is commonly used -- in its elemental form -- as rocket fuels. It has a working similar to oxygen and helps oxidizers in the rocket fuel to burn.

• Elemental fluorine is employed in the manufacturing and chemical laboratories. Isotopically fractionated uranium is created with the help of fluorinated compounds, which is an important step in uranium purification in power plants.

• Fluorspar, calcium, and fluoride are used in smelting metals to make them flow, and is also used to make lenses for focusing infrared radiation.
• Because of fluorine's reactive nature with glass, it is considered valuable in the production of computer chips, microelectronic sensors, and television screens.

• It is also found in toothpastes and mouthwashes, and hence, young children with developing teeth can use fluorinated toothpastes for stronger and healthier teeth. It is used to prevent tooth cavities and dental decay as well.

• Fluorinated compounds help in the production of polymers and plastics, some of which are specially designed to withstand high temperatures and large amount of stress.

• They are used to make polytetrafluoroethylene (teflon) that is used in manufacturing retractable roofs -- Teflon and Tefzel.

• Fluorinated compounds are also used to replace chloroflurocarbons or CFCs (that are found in refrigerators, air conditioners, and spray cans), which were considered to be very dangerous for the Earth's ozone layer.
• They are also used for glass etching and in other industries involving hydrofluoric acid.

• Pure aluminum can be obtained only by using fluorine compounds in the electrolysis process of aluminum.
• Fluorinated compounds (in the form of gas) are used in inhalational anesthetics.

• Hydrofluoric acid is used to mark light bulbs and highly sensitive glass pieces that are too thin for other etching methods.

• The other most common use of fluorinated compounds is in the purification of public water supplies. This has shown to reduce the incidence of dental caries and other such problems to a great extent.

Atomic Number of Fluorine (F)
9

Atomic Mass
18.998404 amu

Number of Protons/Electrons
9

Number of Neutrons
10

Melting Point
-219.62 °C

Boiling Point
-188.12 °C

Density @ 293K
1.696 g/cm3

Crystal Structure
Cubic

Some of the properties of fluorine can be listed as under:

• It is one of the most electronegative and reactive chemical elements. It can react with all organic and inorganic compounds.
• It is pale yellow in color and a corrosive gas.
• Glass, ceramics, carbon, finely divided metals, and even water burn in fluorine with a yellow flame.

• Fluorine has a pungent odor and can be easily detected in concentrations as low as 20 parts per billion.
• This is a highly reactive element and is never found in its 'free form'. The most common fluorine minerals are fluorspar (CaF2), apatite [Ca5(F,Cl,OH)(PO4)3] and cryolite (Na3AlF6).

Reactions

1) Fluorine can react with hydrogen in an explosive manner -- even in the dark and at low temperatures. It reacts violently with water to form hydrogen fluoride and releases oxygen that is highly charged with ozone.
2F2 + 2H2O ==> 4HF + O2

2) Fluorine also reacts with potassium chlorate as it is a powerful oxidizing agent. It oxidizes the potassium chlorate solution into potassium perchlorate.
F2 + 2KClO3 + H2O ==> 2HF + KClO4

3) It also reacts with sulfur, selenium, and tellurium to form halides. It can combine with chlorine, only when heated, by forming a gaseous product of chlorine fluoride and chlorine trifluoride. It can readily combine with bromine and iodine and form colorless liquids, like BrF3 (bromine trifluoride) and IF5 (iodine pentafluoride).

Some of the facts about fluorine, which you may like to know, have been listed as under.

• Fluorine is a highly explosive, corrosive, and poisonous halogen.
• It is found in its gaseous state at room temperature.
• It is obtained from its mineral fluorite forms from nature.

• It is a diatomic element and can react with all elements including noble gases.
• It has a higher oxidation potential than ozone.
• It can damage the soft tissues of the respiratory tract, skin, and eyes. Hence, extreme caution must be taken when you come in contact with pure fluorine, or fluorine compounds.

This highly reactive element is surely one of the most commercially used elements of the periodic table.