Fluorine (F) is the first element in the Halogen group (group 17) in the periodic table. Its atomic number is 9 and its atomic weight is 19, and it"s a gas at room temperature. It is the most electronegative element, given that it is the top element in the Halogen Group, and therefore is very reactive. It is a nonmetal, and is one of the few elements that can form diatomic molecules (F2). It has 5 valence electrons in the 2p level. Its electron configuration is 1s22s22p5. It will usually form the anion F- since it is extremely electronegative and a strong oxidizing agent. Fluorine is a Lewis acid in weak acid, which means that it accepts electrons when reacting. Fluorine has many isotopes, but the only stable one found in nature is F-19.

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Quick Reference Table

Symbol Atomic Number Group Electron Configuration Atomic Weight Density Melting Point Boiling Point Critical Point Oxidation States Electronegativity Stable Isotopes
F
9
17 (Halogens)
1s22s22p5
18.998 g
1.7 g/L
-219.62oC
-188.12oC
144.13K, 5.172 MPa
-1
3.98
F-19

Brief History

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 ubraintv-jp.comical 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.

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Figure 2: Electronic configuration of Fluorine

Fluorine is the most electronegative element because it has 5 electrons in it"s 2P shell. The optimal electron configuration of the 2P orbital contains 6 electrons, so since Fluorine is so close to ideal electron configuration, the electrons are held very tightly to the nucleus. The high electronegativity of fluorine explains its small radius because the positive protons have a very strong attraction to the negative electrons, holding them closer to the nucleus than the bigger and less electronegative elements.


Reactions of Fluorine

Because of its reactivity, elemental fluorine is never found in nature and no other ubraintv-jp.comical element can displace fluorine from its compounds. Fluorine bonds with almost any element, both metals and nonmetals, because it is a very strong oxidizing agent. It is very unstable and reactive since it is so close to its ideal electron configuration. It forms covalent bonds with nonmetals, and since it is the most electronegative element, is always going to be the element that is reduced. It can also form a diatomic element with itself ((F_2)), or covalent bonds where it oxidizes other halogens ((ClF), (ClF_3), (ClF_5)). It will react explosively with many elements and compounds such as Hydrogen and water. Elemental Fluorine is slightly basic, which means that when it reacts with water it forms (OH^-).

<3F_2+2H_2O ightarrow O_2+4HF ag1>

When combined with Hydrogen, Fluorine forms Hydrofluoric acid ((HF)), which is a weak acid. This acid is very dangerous and when dissociated can cause severe damage to the body because while it may not be painful initially, it passes through tissues quickly and can cause deep burns that interfere with nerve function.

There are also some organic compounds made of Fluorine, ranging from nontoxic to highly toxic. Fluorine forms covalent bonds with Carbon, which sometimes form into stable aromatic rings. When Carbon reacts with Fluorine the reaction is complex and forms a mixture of (CF_4), (C_2F_6), an (C_5F_12).

Fluorine reacts with Oxygen to form (OF_2) because Fluorine is more electronegative than Oxygen. The reaction goes:

<2F_2 + O_2 ightarrow 2OF_2 ag4>

Fluorine is so electronegative that sometimes it will even form molecules with noble gases like Xenon, such as the the molecule Xenon Difluoride, (XeF_2).

Fluorine also forms strong ionic compounds with metals. Some common ionic reactions of Fluorine are:

<4F_2 + HCl + H_2O ightarrow 3HF + OF_2 + ClF_3 ag7>


Applications of Fluorine

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. There are many applications of fluorine:

Rocket fuels Polymer and plastics production teflon and tefzel production When combined with Oxygen, used as a refrigerator cooler Hydrofluoric acid used for glass etching Purify public water supplies Uranium production Air conditioning

Sources

Fluorine can either be found in nature or produced in a lab. To make it in a lab, compounds like Potassium Fluoride are put through electrolysis with Hydrofluoric acid to create pure Fluorine and other compounds. It can be carried out with a variety of compounds, usually ionic ones involving Fluorine and a metal. Fluorine can also be found in nature in various minerals and compounds. The two main compounds it can be found in are Fluorspar ((CaF_2)) and Cryolite ((Na_3AlF_6)).

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References

Newth, G. S. Inorganic ubraintv-jp.comistry. Longmans, Green, and Co.:New York, 1903. Latimer, Wendell M., Hildebrand, Joel H. Reference Book of Inorganic ubraintv-jp.comistry. The Macmillan Company: New York, 1938.