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Brief History

Fluorine was discovered in 1530 by Georgius Agricola.  He originally found it in the compound Fluorspar, which was used to promote the fusion of metals.  It was under this application until 1670, when Schwanhard discovered its usefulness in etching glass.  Pure Fluoride was not isolated until 1886 by Henri Moissan, burning and even killing many scientists along the way. (see Figure 1) It has many uses today, a particular one being used in the Manhattan project to help create the first nuclear bomb.

Figure 1: Image courtesy of Wikipedia

Electronegativity of Fluorine

Fluorine is the most electronegative element on the periodic table, which means that it is a very strong oxidizing agent and accepts other elements' electrons.  Fluorine's electron configuration is 1s²2s²2p⁵. (see Figure 2)

Figure 2: Electronic configuration of Sulfur

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

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 (F₂), or covalent bonds where it oxidizes other halogens (ClF, ClF₃, ClF₅). 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^-. 

3F₂+2H₂O → O₂+4HF

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. 

HF+H₂O → H₃O⁺+F^-

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, C₂F₆, an C₅F₁₂.

C(s) + F₂(g) → CF₄(g) + C₂F₆ + C₅F₁₂

Fluorine reacts with Oxygen to form OF₂ because Fluorine is more electronegative than Oxygen.  The reaction goes:

2F₂ + O₂ → 2OF₂

Fluorine is so electronegative that sometimes it will even form molecules with noble gases like Xenon, such as the the molecule Xenon Difluoride, XeF₂.

Xe + F₂ → XeF₂

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

F₂ + 2NaOH → O₂ + 2NaF +H₂

4F₂ + HCl + H₂O → 3HF + OF₂ + ClF₃

F₂ + 2HNO₃ → 2NO₃F + H₂

Applications of Fluorine

There are many applications of Fluorine in society, including:

• 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
• Toothpaste
• 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 (CaF₂) and Cryolite (Na₃AlF₆).

References

1. Newth, G. S. Inorganic Chemistry. Longmans, Green, and Co.:New York, 1903.
2. Latimer, Wendell M., Hildebrand, Joel H. Reference Book of Inorganic Chemistry. The Macmillan Company: New York, 1938.

Outside Links

• http://www.buzzle.com/articles/fluorine-uses.html
  
• http://www.chemicool.com/elements/fluorine.html

Problems

(Highlight to view answers)

1. Q. What is the electron configuration of Fluorine? of F^-?

     A. 1s²2s²2p⁵        

         1s²2s²2p⁶

2. Q. Is Fluorine usually oxidized or reduced? explain.

     A. Fluorine is usually reduced because it accepts an electron from other elements since it is so electronegative.

3. Q. What are some common uses of Fluorine?

    A. Toothpaste, plastics, rocket fuels, glass etching, etc.

4. Q. Does Fluorine form compounds with nonmetals? if so, give two examples, one of them being of an oxide.

     A.  OF₂,   ClF

5. Q. What group is Fluorine in? (include name of group and number)

     A. 17, Halogens

Contributors

• Rachel Feldman, University of California, Davis

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