chlorinous, adj.
/klawr"een, -in, klohr"-/, n.
a halogen element, a heavy, greenish-yellow, incombustible, water-soluble, poisonous gas that is highly irritating to the respiratory organs, obtained chiefly by electrolysis of sodium chloride brine: used for water purification, in the making of bleaching powder, and in the manufacture both of chemicals that do not contain chlorine, as ethylene glycol, and of those that do. Symbol: Cl; at. wt.: 35.453; at. no.: 17.
[1800-10; CHLOR-1 + -INE2]

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Nonmetallic chemical element, chemical symbol Cl, atomic number 17.

It is a toxic, corrosive, greenish yellow gas (as the diatomic molecule Cl2) that severely irritates the eyes and respiratory system (and was used for that purpose as a chemical-warfare agent in World War I). As the chloride ion and in the hypochlorite ion, it has valence 1; in the chlorite, chlorate, and perchlorate ions, it has higher valences. Chlorine and its compounds are important industrial materials with myriad uses in the manufacture of other chlorinated compounds (e.g., PVC, hydrochloric acid, ethylene dichloride, trichloroethylene, PCBs), in water purification (municipal systems, swimming pools), in textile industries, in flame retardants, in special batteries, and in food processing. Sodium chloride (table salt) is by far the most familiar of its compounds. See also bleach.

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 chemical element, second lightest member of the halogen elements (halogen element), or Group 17 (Group VIIa) of the periodic table (periodic law). Chlorine is a toxic, corrosive, greenish-yellow gas that is irritating to the eyes and to the respiratory system (respiration, human).

      Rock salt (common salt, or sodium chloride) has been known for several thousand years; it is the main constituent of the salts dissolved in seawater, from which it was obtained in ancient Egypt by evaporation. In Roman times, soldiers were partially paid in salt (salarium, the root of the modern word salary). In 1648 the German chemist Johann Rudolf Glauber (Glauber, Johann Rudolf) obtained a strong acid, which he called spirit of salt, by heating moist salt in a charcoal furnace and condensing the fumes in a receiver. Later he obtained the same product, now known to be hydrochloric acid, by heating salt with sulfuric acid.

      In 1774 the Swedish chemist Carl Wilhelm Scheele (Scheele, Carl Wilhelm) treated powdered black oxide of manganese with hydrochloric acid and obtained a greenish-yellowish gas, which he failed to recognize as an element. The true nature of the gas as an element was recognized in 1810 by English chemist Humphry Davy (Davy, Sir Humphry, Baronet), who later named it chlorine (from the Greek chloros, meaning “yellowish-green”) and provided an explanation for its bleaching action.

Occurrence and distribution
  Apart from very small amounts of free chlorine (Cl) in volcanic gases, chlorine is usually found only in the form of chemical compounds (chemical compound). It constitutes 0.017 percent of Earth's (Earth) crust. Natural chlorine is a mixture of two stable isotopes (isotope): chlorine-35 (75.53 percent) and chlorine-37 (24.47 percent). The most common compound of chlorine is sodium chloride, which is found in nature as crystalline rock salt, often discoloured by impurities. Sodium chloride is also present in seawater, which has an average concentration of about 2 percent of that salt. Certain landlocked seas, such as the Caspian Sea, the Dead Sea, and the Great Salt Lake of Utah, contain up to 33 percent dissolved salt. Small quantities of sodium chloride are present in blood and in milk. Other chlorine-containing minerals are sylvite ( potassium chloride [KCl]), bischofite (MgCl2 ∙6H2O), carnallite (KCl∙MgCl2 ∙6H2O), and kainite (KCl∙MgSO4 ∙3H2O). It is found in evaporite minerals such as chlorapatite and sodalite. Free hydrochloric acid is present in the stomach.

      Present-day salt deposits must have been formed by evaporation of prehistoric seas, the salts with the least solubility in water crystallizing first, followed by those with greater solubility. Because potassium chloride is more soluble in water than sodium chloride, certain rock salt deposits—such as those at Stassfurt, Ger.—were covered by a layer of potassium chloride. In order to gain access to the sodium chloride, the potassium salt, important as a fertilizer, is removed first.

Physical and chemical properties
      Chlorine is a greenish-yellow gas at room temperature and atmospheric pressure. It is two and a half times heavier than air. It becomes a liquid at −34 °C (−29 °F). It has a choking smell, and inhalation causes suffocation, constriction of the chest, tightness in the throat, and—after severe exposure— edema (filling with fluid) of the lungs. As little as one part per thousand in air causes death within a few minutes, but less than one part per million may be tolerated. Chlorine was the first gas used in chemical warfare in World War I. The gas is easily liquefied by cooling or by pressures of a few atmospheres at ordinary temperature.

      Chlorine has a high electronegativity and a high electron affinity, the latter being even slightly higher than that of fluorine. The affinity of chlorine for hydrogen is so great that the reaction proceeds with explosive violence in light, as in the following equation (where hν is light):

      In the presence of charcoal, the combination of chlorine and hydrogen takes place rapidly (but without explosion) in the dark. A jet of hydrogen will burn in chlorine with a silvery flame. Its high affinity for hydrogen allows chlorine to react with many compounds containing hydrogen. Chlorine reacts with hydrocarbons (hydrocarbon), for example, substituting chlorine atoms for the hydrogen atoms successively. If the hydrocarbon is unsaturated, however, chlorine atoms readily add to the double or triple bond.

      Chlorine molecules (molecule) are composed of two atoms (atom) (Cl2). Chlorine combines with almost all the elements, except for the lighter noble gases (noble gas), to give chlorides; those of most metals are ionic crystals (crystal), whereas those of the semimetals and nonmetals are predominantly molecular.

      The products of reaction with chlorine usually are chlorides with high oxidation numbers (oxidation number), such as iron trichloride (FeCl3), tin tetrachloride (SnCl4), or antimony pentachloride (SbCl5), but it should be noted that the chloride of highest oxidation number of a particular element is frequently in a lower oxidation state than the fluoride with the highest oxidation number. Thus, vanadium forms a pentafluoride, whereas the pentachloride is unknown, and sulfur gives a hexafluoride but no hexachloride. With sulfur, even the tetrachloride is unstable.

      Aside from the −1 oxidation states of some chlorides, chlorine exhibits +1, +3, +5, and +7 oxidation states, respectively, in the following ions: hypochlorite (ClO1.5PT), chlorite (ClO2), chlorate (ClO3), and perchlorate (ClO4). Five oxides—chlorine monoxide (Cl2O), chlorine dioxide (ClO2), chlorine perchlorate (Cl2O4), dichlorine hexoxide (Cl2O6), and dichlorine heptoxide (Cl2O7)—all highly reactive and unstable, have been indirectly synthesized. Chlorine can undergo addition or substitution reactions with organic compounds.

      Chlorine displaces the heavier, less electronegative halogens (halogen element), bromine and iodine, from compounds. The displacement of bromides, for example, occurs according to the following equation:

      Furthermore, it converts several oxides into chlorides. An example is the conversion of iron trioxide to the corresponding chloride:

      Chlorine is moderately soluble in water, yielding chlorine water, and from this solution a solid hydrate of ideal composition, Cl2∙7.66H2O, is obtained. This hydrate is characterized by a structure that is more open than that of ice; the unit cell contains 46 molecules of water and 6 cavities suitable for the chlorine molecules. When the hydrate stands, disproportionation takes place; that is, one chlorine atom in the molecule is oxidized, and the other is reduced. At the same time, the solution becomes acidic, as shown in the following equation:

      in which the oxidation numbers are written above the atomic symbols. Chlorine water loses its efficiency as an oxidizing agent on standing, because hypochlorous acid gradually decomposes. The reaction of chlorine with alkaline solutions yields salts of oxyacids.

      The first ionization energy of chlorine is high. Although ions in positive oxidation states are not very stable, high oxidation numbers are stabilized by coordination, mainly with oxygen and fluorine. In such compounds bonding is predominantly covalent, and chlorine is capable of exhibiting the oxidation numbers +1, +3, +4, +5, +6, and +7.

Production and use
      Rock salt deposits are usually mined; occasionally water is pumped down, and brine, containing about 25 percent sodium chloride, is brought to the surface. When the brine is evaporated, impurities separate first and can be removed. In warm climates salt is obtained by evaporation of shallow seawater by the Sun, producing bay salt.

      Chlorine is produced on a large scale by any of a number of different methods:
      By electrolysis of a concentrated solution of sodium chloride in water. Hydrogen is generated at the cathode and chlorine at the anode. At the same time, sodium hydroxide is produced in the electrolyte; hence, this process is often referred to as chlorine-alkali electrolysis.

      The chemical reactions that take place at each electrode and the overall cell process are given in the following equations:

      in which the symbol e represents a single electron. In the reaction vessel, free chlorine and hydroxide ions (ion) must not come in contact with each other, because chlorine would be consumed according to the reaction

      To accomplish the separation of chlorine gas and the hydroxide ion, a porous wall is inserted between the electrodes (electrode) (diaphragm process), or the iron cathode is replaced by a cathode consisting of liquid mercury (mercury cathode process), which avoids the production of hydroxide ions at the electrode. Instead, free sodium is discharged at the cathode, and this metal is readily dissolved in the mercury, forming an amalgam, as follows:

      The amalgam is allowed to react with water outside the cell:

      The overall process is equivalent to the cell process given above.

      By electrolysis of fused sodium chloride, which also produces metallic sodium; chlorine is again evolved at the anode.

      By electrolysis of fused magnesium chloride, in which chlorine is formed as a by-product in the manufacture of metallic magnesium.

      By oxidation of hydrogen chloride, in which gaseous hydrogen chloride mixed with air or oxygen is passed over pumice in contact with cupric chloride as a catalyst, as shown in the following equation:

      The equilibrium constant for this reaction decreases with increase of temperature; i.e., the reaction proceeds less extensively at higher temperatures. In practice, however, a temperature of 400 °C (750 °F) is required to achieve a reasonable rate of conversion.

      Of historical interest is the process in which a mixture of almost any solid chloride and manganese dioxide (MnO2) yields chlorine when heated with concentrated sulfuric acid (H2SO4). The reaction occurs, as follows:

      In the laboratory chlorine is frequently prepared by the oxidation of concentrated hydrochloric acid with permanganate or dichromate salts:

      Most of the chlorine produced is used for chemical processes involving the introduction of chlorine into organic compounds, yielding carbon tetrachloride (used as a solvent, a fire extinguisher, and a dry-cleaning agent), glycols (glycol) (used as antifreeze), and other organic compounds for the manufacture of plastics (plastic) (polyvinyl chloride), dyes, and synthetic rubber. sulfur chloride, made by the action of chlorine on carbon disulfide or by combining sulfur and chlorine, is used in the vulcanization of rubber and as a chlorinating agent in organic synthesis. Sulfur dioxide combines with chlorine to give sulfuryl dioxide. Chlorine and carbon monoxide form carbonyl chloride (COCl2), or phosgene, which was employed as a chemical weapon in World War I and is used mainly in the preparation of isocyanates and polyurethanes (polyurethane) and in metallurgy to transform certain oxides into chlorides. The reactions that form phosgene and sulfuryl dioxide (SO2Cl2) are

      Much chlorine is used to sterilize water and wastes, and the substance is employed either directly or indirectly as a bleaching agent for paper or textiles (textile) and as “bleaching powder” (Ca[OCl]2∙CaCl2∙Ca[OH]2∙2H2O). Chlorine is applied in the manufacturing of high-purity hydrochloric acid, the extraction of titanium with the formation of titanium tetrachloride (TiCl4), and the removal of tin from old tinplate. Anhydrous aluminum chloride (AlCl3) is made by the reaction of chlorine with scrap aluminum or with aluminum oxide and carbon. Chlorine is also used to prepare silicon tetrachloride (SiCl4) and methyl chloride (CH3Cl), which are employed in the synthesis of silicon materials. Chlorine enters directly, or indirectly as an intermediate, into many organic syntheses of industrial importance.

      Free chlorine may be recognized by its smell, its colour, and its characteristic reaction with mercury to produce white mercury dichloride (HgCl2). Tests for chloride ions are:
      The formation of a white precipitate of silver chloride (AgCl) on addition of silver nitrate (AgNO3) in dilute nitric acid (HNO3). (This precipitate is soluble in the presence of ammonia.)

      The formation of chromyl chloride (CrO2Cl2), a red gas, by heating a solid sample with potassium dichromate (K2Cr2O7) and concentrated sulfuric acid. When chromyl chloride is passed into water, a yellow chromate solution forms (bromides and iodides do not form analogous compounds).

      The evolution of free chlorine by heating the sample with manganese dioxide (MnO2) and concentrated sulfuric acid.

      The following methods are available for the quantitative determination of free chlorine:
      The chlorine-containing gas is shaken with an aqueous solution of potassium iodide (KI), and the resulting iodine is determined by titration.

      Chlorine is reduced in alkaline solution by an alkali arsenite (e.g., NaAsO2). Back-titration of excess arsenite is carried out with potassium bromate (KBrO3).

      In the presence of an alkali hydroxide (e.g., NaOH), chlorine is reduced to the chloride ion by hydrogen peroxide (H2O2), and the excess alkali hydroxide is back-titrated with acid.

      With sulfur dioxide (SO2) or sodium thiosulfate (Na2S2O3), chlorine is reduced to chloride, and the latter is analyzed as silver chloride (see below).

      Colorimetric measurements are carried out in the presence of o-toluidine in hydrochloric acid.

      For the determination of chloride ions, one of the following methods may be recommended:
       gravimetric analysis (analysis by weight of a given product) as silver chloride.

      Titration of a neutral chloride solution with silver nitrate in the presence of potassium chromate.

      Potentiometric titration (measurement of voltage changes) with silver nitrate, a process that can be carried out in the presence of bromide and iodide ions.

      Most insoluble chlorides can be melted with soda (Na2CO3), and the resulting melt is then usually soluble in water. Organic compounds containing chlorine are heated with alkali peroxide, and the product is dissolved in water.

atomic number
atomic weight
melting point
−103 °C (−153 °F)
boiling point
−34 °C (−29 °F)
density (1 atm, 0 °C or 32 °F)
3.214 g/litre (0.429 ounce/gallon)
oxidation states
−1, +1, +3, +5, +7
electronic config.
Karl Christe Stefan Schneider

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Universalium. 2010.

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