Fluorine

Fluorine (symbol F, atomic weight 19), a chemical element of the halogen group. It is never found in the uncombined condition, but in combination with calcium as fluor-spar CaF2 it is widely distributed; it is also found in cryolite Na3AlF6, in fluor-apatite, CaF2·3Ca3P2O8, and in minute traces in sea-water, in some mineral springs, and as a constituent of the enamel of the teeth. It was first isolated by H. Moissan in 1886 by the electrolysis of pure anhydrous hydrofluoric acid containing dissolved potassium fluoride. The U-shaped electrolytic vessel and the electrodes are made of an alloy of platinum-iridium, the limbs of the tube being closed by stoppers made of fluor-spar, and fitted with two lateral exit tubes for carrying off the gases evolved. Whilst the electrolysis is proceeding, the apparatus is kept at a constant temperature of −23° C. by means of liquid methyl chloride. The fluorine, which is liberated as a gas at the anode, is passed through a well cooled platinum vessel, in order to free it from any acid fumes that may be carried over, and finally through two platinum tubes containing sodium fluoride to remove the last traces of hydrofluoric acid; it is then collected in a platinum tube closed with fluor-spar plates. B. Brauner (Jour. Chem. Soc., 1894, 65, p. 393) obtained fluorine by heating potassium fluorplumbate 3KF·HF·PbF4. At 200° C. this salt decomposes, giving off hydrofluoric acid, and between 230-250° C. fluorine is liberated.

Fluorine is a pale greenish-yellow gas with a very sharp smell; its specific gravity is 1.265 (H. Moissan); it has been liquefied, the liquid also being of a yellow colour and boiling at -187° C. It is the most active of all the chemical elements; in contact with hydrogen combination takes place between the two gases with explosive violence, even in the dark, and at as low a temperature as −210° C; finely divided carbon burns in the gas, forming carbon tetrafluoride; water is decomposed even at ordinary temperatures, with the formation of hydrofluoric acid and “ozonised” oxygen; iodine, sulphur and phosphorus melt and then inflame in the gas; it liberates chlorine from chlorides, and combines with most metals instantaneously to form fluorides; it does not, however, combine with oxygen. Organic compounds are rapidly attacked by the gas.

Only one compound of hydrogen and fluorine is known, namely hydrofluoric acid, HF or H2F2, which was first obtained by C. Scheele in 1771 by decomposing fluor-spar with concentrated sulphuric acid, a method still used for the commercial preparation of the aqueous solution of the acid, the mixture being distilled from leaden retorts and the acid stored in leaden or gutta-percha bottles. The perfectly anhydrous acid is a very volatile colourless liquid and is best obtained, according to G. Gore (Phil. Trans., 1869, p. 173) by decomposing the double fluoride of hydrogen and potassium, at a red heat in a platinum retort fitted with a platinum condenser surrounded by a freezing mixture, and having a platinum receiver luted on. It can also be prepared in the anhydrous condition by passing a current of hydrogen over dry silver fluoride. The pure acid thus obtained is a most dangerous substance to handle, its vapour even when highly diluted with air having an exceedingly injurious action on the respiratory organs, whilst inhalation of the pure vapour is followed by death. The anhydrous acid boils at 19°.5 C. (H. Moissan), and on cooling, sets to a solid mass at −102°.5 C, which melts at −92°.3 C. (K. Olszewski, Monats. für Chemie, 1886, 7, p. 371). Potassium and sodium readily dissolve in the anhydrous acid with evolution of hydrogen and formation of fluorides. The aqueous solution is strongly acid to litmus and dissolves most metals directly. Its most important property is that it rapidly attacks glass, reacting with the silica of the glass to form gaseous silicon fluoride, and consequently it is used for etching. T.E. Thorpe (Jour. Chem. Soc., 1889, 55, p. 163) determined the vapour density of hydrofluoric acid at different temperatures, and showed that there is no approach to a definite value below about 88° C. where it reaches the value 10.29 corresponding to the molecular formula HF; at temperatures below 88° C. the value increases rapidly, showing that the molecule is more complex in its structure. (For references see J.N. Friend, The Theory of Valency (1909), p. 111.) The aqueous solution behaves on concentration similarly to the other halogen acids; E. Deussen (Zeit. anorg. Chem., 1905, 44, pp. 300, 408; 1906, 49, p. 297) found the solution of constant boiling point to contain 43.2% HF and to boil at 110° (750 mm.).

The salts of hydrofluoric acid are known as fluorides and are easily obtained by the action of the acid on metals or their oxides, hydroxides or carbonates. The fluorides of the alkali metals, of silver, and of most of the heavy metals are soluble in water; those of the alkaline earths are insoluble. A characteristic property of the alkaline fluorides is their power of combining with a molecule of hydrofluoric acid and with the fluorides of the more electro-negative elements to form double fluorides, a behaviour not shown by other metallic halides. Fluorides can be readily detected by their power of etching glass when warmed with sulphuric acid; or by warming them in a glass tube with concentrated sulphuric acid and holding a moistened glass rod in the mouth of the tube, the water apparently gelatinizes owing to the decomposition of the silicon fluoride formed. The atomic weight of fluorine has been determined by the conversion of calcium, sodium and potassium fluorides into the corresponding sulphates. J. Berzelius, by converting silver fluoride into silver chloride, obtained the value 19.44, and by analysing calcium fluoride the value 19.16; the more recent work of H. Moissan gives the value 19.05.

See H. Moissan, Le Fluor et ses composes (Paris, 1900).

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