   #copyright

Fluorine

2007 Schools Wikipedia Selection. Related subjects: Chemical elements

          Distinguish from fluorene and fluorone.


                 9                oxygen ← fluorine → neon
                 -
                ↑
                F
                ↓
                Cl

                                  Periodic Table - Extended Periodic Table

                                                                   General
                                       Name, Symbol, Number fluorine, F, 9
                                                  Chemical series halogens
                                             Group, Period, Block 17, 2, p
                                            Appearance Yellowish brown gas
                                          Atomic mass 18.9984032 (5) g/mol
                                     Electron configuration 1s^2 2s^2 2p^5
                                                  Electrons per shell 2, 7
                                                       Physical properties
                                                                 Phase gas
                                              Density (0 °C, 101.325 kPa)
                                                                   1.7 g/L
                                                    Melting point 53.53  K
                                              (-219.62 ° C, -363.32 ° F)
                                                     Boiling point 85.03 K
                                              (-188.12 ° C, -306.62 ° F)
                                        Critical point 144.13 K, 5.172 MPa
                                  Heat of fusion (F[2]) 0.510 kJ·mol^−1
                             Heat of vaporization (F[2]) 6.62 kJ·mol^−1
                                             Heat capacity (25 °C) (F[2])
                                                31.304 J·mol^−1·K^−1

   CAPTION: Vapor pressure

                                            P/Pa  1  10 100 1 k 10 k 100 k
                                           at T/K 38 44 50  58   69   85

                                                         Atomic properties
                                                   Crystal structure cubic
                                                     Oxidation states −1
                                                   (strongly acidic oxide)
                                    Electronegativity 3.98 (Pauling scale)
                                                       Ionization energies
                                          ( more) 1st: 1681.0 kJ·mol^−1
                                                  2nd: 3374.2 kJ·mol^−1
                                                  3rd: 6050.4 kJ·mol^−1
                                                       Atomic radius 50 pm
                                               Atomic radius (calc.) 42 pm
                                                     Covalent radius 71 pm
                                         (see covalent radius of fluorine)
                                               Van der Waals radius 147 pm
                                                             Miscellaneous
                                             Magnetic ordering nonmagnetic
                      Thermal conductivity (300 K) 27.7 mW·m^−1·K^−1
                                             CAS registry number 7782-41-4
                                                         Selected isotopes

                 CAPTION: Main article: Isotopes of fluorine

                                    iso   NA  half-life  DM DE ( MeV)  DP
                                    ^18F syn  109.77 min ε  1.656     ^18O
                                    ^19F 100% F is stable with 10 neutrons

                                                                References

   Fluorine ( IPA: /ˈflʊərɪːn, -ɔːrɪːn/, Latin: fluere, meaning "to
   flow"), is the chemical element with the symbol F and atomic number 9.
   Atomic fluorine is univalent and is the most chemically reactive and
   electronegative of all the elements. In its elementally isolated (pure)
   form, fluorine is a poisonous, pale, yellow- green gas, with chemical
   formula F[2]. Like other halogens, molecular fluorine is highly
   dangerous; it causes severe chemical burns on contact with skin.

   Fluorine's relatively large electronegativity and small atomic radius
   gives it interesting bonding characteristics, particularly in
   conjunction with carbon. See covalent radius of fluorine.

Notable characteristics

   Pure fluorine (F[2]) 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.

   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 [FeF[4]]^-, or H[2]F^+.

   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.

History

   Fluorine in the form of fluorspar (also called fluorite) ( calcium
   fluoride) 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. Karl 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.

   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 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).

   The first large-scale production of fluorine was needed for the atomic
   bomb Manhattan project in World War II where the compound uranium
   hexafluoride (UF[6]) was needed as a gaseous carrier of uranium to
   separate the ^235U and ^238U isotopes of uranium. Today both the
   gaseous diffusion process and the gas centrifuge process use gaseous
   UF[6] to produce enriched uranium for nuclear power applications. In
   the Manhattan Project, it was found that elemental fluorine was present
   whenever UF[6] was, due to the spontaneous decomposition of this
   compound into UF[4] and F[2]. The corrosion problem due to the F[2] was
   eventually solved by electrolytically coating all UF[6] 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 F[2].

Safety

   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 odour that is detectable in concentrations as
   low as 20 nL/L. Its MAC-value is 1 1 µL/L. All equipment must be
   passivated before exposure to fluorine. For more information consult an
   MSDS.

   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 in^2 or 23 cm ^2). 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.

Preparation

   Elemental fluorine is prepared industrially by Moissan's original
   process: electrolysis of anhydrous HF in which KHF[2] 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
   K[2]MnF[6] and of SbF[5]. The reaction is: 2K[2]MnF[6] + 4SbF[5] →
   4KSbF[6] + MnF[2] + F[2] This is not a practical synthesis, but
   demonstrates that electrolysis is not essential.

Compounds

   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 synthesised by Neil Bartlett in 1962 - xenon
   hexafluoroplatinate, XePtF[6], 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.
   Fluorite (CaF2) crystals
   Enlarge
   Fluorite (CaF[2]) crystals

   This element is recovered from fluorite, cryolite, and fluorapatite.

   For a list of fluorine compounds, see here.

Applications

   Atomic fluorine and molecular fluorine are used for plasma etching in
   semiconductor manufacturing, flat panel display production and MEMS
   fabrication. Other uses:
     * Hydrofluoric acid (chemical formula HF) is used to etch glass in
       light bulbs and other products.
     * Fluorine is indirectly used in the production of low friction
       plastics such as Teflon, and in halons such as Freon.
     * Along with some of its compounds, fluorine is used in the
       production of pure uranium from uranium hexafluoride and in the
       synthesis of numerous commercial fluorochemicals, including vitally
       important pharmaceuticals, agrochemical compounds, lubricants, and
       textiles.
     * Fluorochlorohydrocarbons are used extensively in air conditioning
       and in refrigeration. Chlorofluorocarbons have been banned for
       these applications because they contribute to ozone destruction and
       the ozone hole. Interestingly, since it is chlorine and bromine
       radicals which harm the ozone layer, not fluorine, compounds which
       do not have chlorine or bromine and contain only fluorine, carbon
       and hydrogen (called hydrofluorocarbons), are not on the E.P.A.
       list of ozone-depleting substances , and have been widely used as
       replacements for the chlorine and bromine containing fluorocarbons.
       Hydrofluorocarbons do have a greenhouse effect, but a small one
       compared with carbon dioxide and methane.
     * Sulfur hexafluoride is an extremely inert and nontoxic gas, and a
       member of a class of compounds that are potent greenhouse gases.
     * Many important agents for general anesthesia such as sevoflurane,
       desflurane, and isoflurane are hydrofluorocarbon derivatives.
     * Fluconazole is a triazole antifungal drug used in the treatment and
       prevention of superficial and systemic fungal infections.
     * Fluoroquinolones are a family of broad-spectrum antibiotics.
     * Sodium hexafluoroaluminate ( cryolite), is used in the electrolysis
       of aluminium.
     * Compounds of fluorine, including sodium fluoride (NaF), stannous
       fluoride (SnF[2]) and sodium MFP, are used in toothpaste to prevent
       dental cavities. These compounds are also added to municipal water
       supplies, a process called water fluoridation, though a combination
       of health concerns and urban legends has sometimes led to
       controversy.
     * In much higher concentrations, sodium fluoride has been used as an
       insecticide, especially against cockroaches.
     * Fluorides have been used in the past to help molten metal flow,
       hence the name.
     * ^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.

   Some researchers including US space scientists in the early 1960s have
   studied elemental fluorine gas as a possible rocket propellant due to
   its exceptionally high specific impulse. The experiments failed because
   fluorine proved difficult to handle, and its combustion products of
   course proved extremely toxic and corrosive.

   Retrieved from " http://en.wikipedia.org/wiki/Fluorine"
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   with only minor checks and changes (see www.wikipedia.org for details
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