Halogens and their compounds
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Chlorine
Uses
The first large-scale use of chlorine was in the manufacture of bleaching powder for use in making paper and cotton textiles. Bleaching powder was later replaced by liquid chlorine, which also came into widespread use as a germicide for public water supplies. Presently the principal use of chlorine is in making chemical compounds. Important inorganic chemicals made by direct action of chlorine on other substances include sulfur chloride, thionyl chloride, phosgene, aluminum chloride, iron(III) chloride, titanium(IV) chloride, tin(IV) chloride, and potassium chlorate.
Organic chemicals made directly from chlorine include derivatives of methane (methyl chloride, methylene chloride, chloroform, and carbon tetrachloride); chlorobenzene and ortho- and para-dichlorobenzenes; ethyl chloride; and ethylene chloride.
Commercial preparation
Of several processes that have been used for the manufacture of chlorine, the oldest employed the reaction of hydrochloric acid with manganese dioxide. The procedure was inefficient, and its commercial application was short-lived.
A process introduced about 1868 by the English chemist Henry Deacon was based on the reaction of atmospheric oxygen with hydrochloric acid, which was available as a by-product of the Leblanc process for making soda ash; when the Leblanc process became obsolete, the Deacon process fell into disuse.
The chlor-alkali industry—in which chlorine and caustic soda (sodium hydroxide) are produced simultaneously by electrolytic decomposition of salt (sodium chloride)—has become the principal source of chlorine during the 20th century. As noted earlier, in the two important versions of the electrolytic process, brine is the electrolyte (in which the passage of electric current occurs by the movement of charged particles called ions), and graphite rods are the anodes (positive terminals). The difference between the two processes derives from the distinct behaviour of iron and of mercury when those metals are used as cathodes (negative terminals).
In brine, the two substances susceptible to chemical reduction are positively charged sodium ions and neutral water molecules. At a reversible cathode, reduction of sodium ions requires a higher voltage than does the reduction of water molecules, and application of a voltage high enough to reduce sodium ions would effect reduction of a considerable amount of water but of a very small number of sodium ions. The reaction occurring at the surface of an iron cathode is represented by the following equation:
At a mercury cathode, on the other hand, appreciable reduction of water requires a much higher voltage than that needed at an iron cathode. This so-called overpotential is so great, in fact, that the electrode voltage can be raised to that needed for the reduction of sodium ions without affecting the water molecules.
Passage of a direct electric current through brine is attended by chemical changes at the surfaces where the electrodes come in contact with the electrolyte. At the graphite anode, chloride ions present in the dissolved salt are converted by oxidation to elemental chlorine, which is led away through a vent. At the iron cathode, reduction of water takes place, according to the equation shown above. The hydrogen gas is removed, while the hydroxide ions remain in the solution. The net result is that chloride ions and water are consumed and chlorine gas, hydrogen gas, and hydroxide ions are produced. Complete conversion of chloride to hydroxide is not practical, but as brine is continuously introduced at the top of the cell, a solution containing nearly equal amounts of salt and caustic soda is withdrawn at the bottom. Purification of the effluent liquor yields solid sodium hydroxide containing only a small amount of salt.
Successful production of chlorine and caustic soda in these cells requires that the two products be separated, because upon mixing they would react with one another. The chlorine is kept away from the caustic by interposing a diaphragm between the electrodes: such cells are commonly called diaphragm cells.
In the other main variant of the chlor-alkali process, the so-called mercury cell is employed. The cathode in such a cell is a shallow layer of mercury flowing across the bottom of the vessel; graphite anodes extend down into the brine electrolyte. A powerful direct current is caused to pass between the graphite rods and the mercury surface. At the anodes, chloride ions are converted to chlorine gas, as in the diaphragm cell; the reaction occurring at the mercury cathode, however, differs from that at an iron cathode. Positively charged sodium ions in the brine migrate to the mercury surface, where the voltage is high enough to reduce them to sodium metal without reducing the water because of the above-noted overpotential of mercury. The metallic sodium formed at the cathode dissolves in the mercury, and the solution (called an amalgam) flows out of the cell into another vessel, where it is brought into contact with water, which reacts with the sodium to form sodium hydroxide and hydrogen.
The overall result of operating a mercury cell is the same as that of operating a diaphragm cell: sodium chloride and water are changed into sodium hydroxide, chlorine, and hydrogen. Use of the mercury cell, however, makes it possible to generate the sodium hydroxide in the absence of salt, so that evaporation of the caustic liquor produces solid sodium hydroxide completely free of sodium chloride. The higher purity of the product makes it more desirable for certain applications, notably in the manufacture of rayon.
John V. KillhefferFluorine
Aluminum refining
The fluorine industry is intimately related to the production of aluminum. Alumina (aluminum oxide, Al2O3) can be reduced to metallic aluminum by electrolysis when fused with a flux consisting of sodium fluoroaluminate (Na3AlF6), usually called cryolite. After starting the process, the cryolite is not used up in massive quantities, but a small supply is needed to make up for inevitable losses. Cryolite is a rare mineral, however, found in commercial quantities only in Greenland. The supply is limited, and it has other uses in glass, in enamels, and as a filler for resin-bonded grinding wheels.
The supply problem was solved by the development of synthetic cryolite. For this synthetic, however, a source of fluorine was needed. Fluorine is actually somewhat more abundant in the Earth’s crust than chlorine, but most of it is distributed in various rocks in very small quantities. In a form available to the industrial chemist, it is much scarcer than chlorine. Until the 1960s almost the only source was fluorspar (CaF2), a mineral long known and used as a flux in various metallurgical operations. It is still so used, in quantities larger than before, because the processes that are coming into greatest use for making steel, the basic oxygen process and the electric furnace, use two to three times as much flux as the earlier open-hearth furnaces did. The mineral fluorspar is widely distributed, but the supplies of good quality ore are not large; it has been found necessary to utilize lower grade ores, making the processing more expensive. A very large reserve that can be tapped for fluorine is the 3 percent or so that is present in some phosphate rock. In the past this fluorine content was seldom recovered; the future will undoubtedly see a major reversal.
Refrigerants
These inorganic uses, as a flux and in the manufacture of aluminum, formerly constituted almost the whole of the fluorine industry. The organic fluorine industry, a separate branch, began in the late 1920s with the discovery by Thomas Midgley, Jr., of the United States, of the fluorine-containing refrigerants. A new refrigerant was needed for the domestic refrigerators that were just beginning to be produced on a large scale. Ammonia was unsuitable because even a minute leak would give an unpleasant smell, and breakdown would release poisonous quantities of the gas. Although many fluorine compounds were known to be poisonous, Midgley found some that were remarkably nontoxic. They also had the physical properties required for a refrigerant and were totally odourless.
The most used of these is Freon 12 (CCl2F2), dichlorodifluoromethane; also used is Freon 22 (CHClF2), chlorodifluoromethane. Several analogous compounds containing carbon, fluorine, chlorine, and sometimes hydrogen are available.
Isotope separation
The next advance in the fluorine industry was connected with the development of the atomic bomb during World War II. It was necessary to separate the small proportion of the fissionable isotope uranium-235 from other, nonfissionable uranium isotopes. This separation could be done by diffusion, working with uranium hexafluoride, a gas. Fluorine at that time was made only occasionally on a small laboratory scale, and it had a reputation for intense chemical reactivity and for being difficult to handle. The solution to the problem of large-scale preparation of elemental fluorine, which required the development and introduction of novel, fluorine-resistant materials of construction, made this important element generally available. Fluorine manufacture is now routine. Other uses have been developed: as a component in some rocket propellants, for the preparation of the extremely reactive interhalogen compounds such as chlorine trifluoride (ClF3), used for cutting steel, and for the preparation of sulfur hexafluoride, an extremely stable gas that has been employed as an insulator in electrical applications.
Other uses
Nonstick frying pans have been coated with a fluorocarbon resin, the best known of which is polytetrafluoroethylene. There are several other fluorocarbon and fluorinated hydrocarbon resins; some have highly specialized applications in the aerospace industry.
Fluorinated compounds are also used in textile treatments; some are soil-release agents that make fabric easy to wash. The salt sodium fluoroacetate is an extremely powerful rodenticide; it has been reported to give good control of rats, but it must be used with great care. Sodium bifluoride is used as a laundry sour; it also removes iron stains without weakening the fabric.
A minor but important use of fluorine in some countries is in the fluoridation of drinking water in the interest of dental health.
Bromine
The properties of bromine are significantly different from those of fluorine and chlorine, and it is far less abundant. Discovered in the early 19th century, in the form of its salts (bromides) in the bitterns remaining after evaporating seawater and extracting the sodium chloride, it was obtained later from Stassfurt, Germany, as a by-product in the production of potassium salts and from other salt deposits and salt lakes. Its main use was originally for bromides in medicine, still a minor use. Bromine first became of industrial importance with the development of the modern photographic process, in which the light-sensitive material is an emulsion of minute particles of silver bromide (together with silver chloride, or iodide, or both) in gelatin.
Tetraethyllead was another of Thomas Midgley’s discoveries in the 1920s. Long the only effective agent in preventing “knock” in gasoline engines, tetraethyllead is now supplemented by tetramethyl lead, a similar compound. Although the knock problem was solved, a method was needed to get all traces of lead out of the engine cylinder. This removal was achieved by the addition of small quantities of a scavenger, ethylene dibromide, often in a mixture with ethylene dichloride.
For a time the expanding world automobile industry threatened a scarcity of bromine, obtained from brines from the Great Lakes region and Searles Lake in the United States, and from the Dead Sea, which contains about 0.5 percent bromine. To meet the demand it was necessary to turn to seawater, which contains about 70 parts per million bromine.
To produce bromine from seawater, very large volumes of water must be processed. A preferable site for the operation is a neck of land projecting into the ocean so that water can be taken from one side and discharged to the other, avoiding the problem of processing the same water. The water is made acid with a little sulfuric acid and then treated with chlorine, which releases bromine from the bromides.
A current of air removes the bromine as a very dilute mixture of bromine with air. The bromine is absorbed in sodium carbonate, after which treatment with sulfuric acid releases the bromine again in a much more concentrated form.
By far the greater part of the bromine produced is converted to ethylene dibromide by treatment with ethylene. Most of the ethylene dibromide is used in gasoline as a scavenger for lead; but it is also used as a fumigant, as a solvent for certain gums, and for further syntheses. The next most important bromine compound is methyl bromide, which is used as a fumigant, sometimes as a fire extinguisher, and for further syntheses.
Iodine
Iodine enters the chemical industry on a smaller scale. The largest producer is Japan, where iodine is obtained from seaweed. Seawater contains only about 0.05 part per million iodine, but some species of seaweed are able to concentrate this iodine manyfold, so that commercial extraction of the iodine is possible.
The most important industrial use of iodine compounds is the small amount of silver iodide used with silver bromide in photography. Iodine is important also in medicine (although this is not a large-scale use) in the treatment of certain thyroid conditions, and it is added to common table salt to prevent such conditions. It is also used directly as a disinfectant. Iodine is a component of a few useful dyes. The laboratory chemist frequently makes use of iodine or iodine compounds in synthesis and also in analysis. Crystalline silver iodide is useful in cloud seeding.
Anthony Standen