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Organofluorine compounds are prepared by numerous routes, depending on the degree and regiochemistry of fluorination sought and the nature of the precursors. The direct fluorination of hydrocarbons with F 2, often diluted with N 2, is useful for highly fluorinated compounds: R 3 CH + F 2 → R 3 CF + HF
Radical fluorination with the pure element is difficult to control and highly exothermic; care must be taken to prevent an explosion or a runaway reaction. With chlorine the reaction is moderate to fast; with bromine, slow and requires intense UV irradiation; and with iodine, it is practically nonexistent and thermodynamically unfavored.
Chlorine pentafluoride (ClF 5) is made on a large scale by direct fluorination of chlorine with excess fluorine gas at 350 °C and 250 atm, and on a small scale by reacting metal chlorides with fluorine gas at 100–300 °C. It melts at −103 °C and boils at −13.1 °C.
Such reagents include cobalt trifluoride, chlorine trifluoride, and iodine pentafluoride. [4] The method electrochemical fluorination is used commercially for the production of perfluorinated compounds. It generates small amounts of elemental fluorine in situ from hydrogen fluoride. The method avoids the hazards of handling fluorine gas.
Chlorine and oxygen can bond in a number of ways: chlorine monoxide radical, ClO•, chlorine (II) oxide radical; chloroperoxyl radical, ClOO•, chlorine (II) peroxide radical; chlorine dioxide, ClO 2, chlorine (IV) oxide; chlorine trioxide radical, ClO 3 •, chlorine (VI) oxide radical; chlorine tetroxide radical, ClO 4 •, chlorine (VII ...
ClO 2 F was first reported by Schmitz and Schumacher in 1942, who prepared it by the fluorination of ClO 2. [2] The compound is more conveniently prepared by reaction of sodium chlorate and chlorine trifluoride [3] and purified by vacuum fractionation, i.e. selectively condensing this species separately from other products. This species is a ...
When chlorine gas (Cl 2) is irradiated, some of the molecules are split into two chlorine radicals (Cl•), whose free electrons are strongly nucleophilic. One of them breaks a C–H covalent bond in CH 4 and grabs the hydrogen atom to form the electrically neutral HCl.
By far most chlorine is manufactured from table salt (NaCl) by electrolysis in the chlor-alkali process. The resulting gas at atmospheric pressures is liquified at high pressure. The liquefied gas is transported and used as such. [citation needed] As a strong oxidizing agent, chlorine kills via the oxidation of organic molecules. [16]