Fluorine in Periodic Table
Fluorine (F), chemical formula F2, atomic number 9 is the most electronegative and most chemically reactive member of the halogen family or Group 17 of the periodic table element. Due to high reactivity and the small size of atoms, fluorine combines directly with metal, almost all nonmetal, and noble gases. The reactions with many elements are vigorous and sometimes explosive. The exceptionally high reactivity of fluorine was assigned due to low F-F bond energy ( 159 kJ mol-1) which gradually decreases in other members of the group (chlorine, bromine, and iodine).
Isolation and History
The name of the chemical element fluorine (chemical symbol F) origin from the mineral fluorspar (CaF2) suggested by French physicist André-Marie Ampère in a letter to Davy (1812). In 1771 the Swedish chemist Carl Wilhelm Scheele discovered corrosive compounds hydrofluoric acid (HF) by heating fluorspar with concentrated sulfuric acid in a glass retort but it could not be used for isolation of fluorine due to several difficulties. Anhydrous hydrofluoric acid is a nonconductor of electricity while an aqueous solution produces ozonized oxygen on electrolysis. No chemical oxidizing agent was able to oxidize a fluoride compound or hydrofluoric acid. The decomposition of fluoride compounds gives fluorine gas but the process is unlikely from the thermodynamic point of view since the compounds have high lattice or bond energy.
Finally, French chemist Henri Moissan in 1886 preparing fluorine by electrolysis reaction of potassium fluoride (KF) in anhydrous hydrofluoric acid in a U-tube of Platinum with a platinum-iridium electrode. He received the Noble prize for chemistry in 1906 for isolating chemical elements but due to toxic characteristics, the progressing of fluorine chemistry is very slow. Fluorine is now prepared by electrolysis molten mixture of potassium fluoride and hydrofluoric acid in a mild cell acting as the cathode and anode is the central rod of nongraphitic carbon separated from cathode by the porous diaphragm. For more than 150 years, all the chemical preparation methods had failed to produce fluorine, success of discovery has come in 1986 by American chemist Karl O. Christe through the reaction of K2MnF6 and antimony pentafluoride (SbF5). Both the chemical compounds are obtained by using hydrogen fluoride.
Abundance and Occurrence
None of the halogen family-like fluorine, chlorine, bromine, iodine occurs free in nature owing to their high reactivity. Fluorine occurs in the earth’s crust to the extent of 0.065 percent (544 ppm; thirteenth in abundance, which is higher than that of chlorine (0.055 percent, 126 ppm, twentieth in abundance). The most common minerals found in the earth’s environment concentrated with fluorine element are fluorite or fluorspar (CaF2), cryolite (Na3AlF6), and fluorapatite[3Ca3(PO4)2, CaF2]. The principal fluorine-containing minerals (fluorspar) are obtained in different parts of the world like India, Illinois, Kentucky, Derbyshire, southern Germany, the south of France, and Russia and Greenland is the chief source of mineral-like cryolite (Na3AlF6). In India, fluorite occurs in Chhattisgarh in the district of Drug, Raipur, and Jabalpur, in Rajasthan in Jaipur, Ajmer, and Kishangarh.
Physical and Chemical Properties
Fluorine (atomic number = 9, atomic weight or mass = 18.998403163) is a faint yellow greenish gas at ordinary or room temperature that turns into yellow liquid after cooling with only one stable isotope 19F or F-19. In the vapour state, fluorine normally exists as a diatomic molecule with a melting point -219 °C and boiling point -188 °C held by weak Van der Waals attraction. In the solid phase, it forms layers of the crystal lattice. The most electronegative chemical element in the periodic table, fluorine always assigns -1 oxidation number introduces polarity in hydrogen fluoride molecule which associates via intermolecular hydrogen bonding. The chemistry (physical and chemical properties) and periodic table facts of fluorine can be described by its electron configuration which is one electron short of the next noble gas.
Some of the interesting facts of higher halogens can be explained only by considering the participation of d-orbitals in reactions and chemical bonding. The noble gas electronic structure of the fluorine atom is achieved in two ways, by gaining one electron from reactive surroundings to form an ion or by forming a shared pair in covalent bonding. The tendency to accept one electron is so great due to high electron affinity and ionization energy. It brings about the maximum coordination number of other elements due to the small size of the atom and extremely high electronegativity. Both these facts permit a maximum number of small fluorine atoms without accumulation of excessive charge density on central atom like SF6, OsF6, [SiF6]2-, IF7, etc.
Uses of Fluorine
At least 10,000 tonnes of fluorine (F2) produces annually over the world, used in the manufacture of UF6 in nuclear power stations, in making SF6 dielectrics and fluorinating agents like ClF3, BrF3, etc. Tungsten (W) and Rhenium (Re) fluorinated to volatile WF6 and ReF6 requires in vapour deposition of metals on machine components, the chemical compounds of fluorine like hydrogen fluoride (HF), chlorofluorocarbon (CFC), boron trifluoride (BF3) and synthetic cryolite has very important industrial application, the chemical compound fluorspar uses as a flux in metallurgy.