ALKALI METAL

any of the six chemical elements that make up Group I of the periodic tablenamely, lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). The alkali metals are so called because they form alkalies (i.e., strong bases capable of neutralizing acids) when they combine with other elements. The only members of the alkali metal family that are relatively abundant in the Earth's crust are potassium and sodium. These two were the first alkali metals to be isolated (1807). Alkali metals bear little resemblance to the more familiar metals such as iron and copper. They are silver-white in colour, malleable, and soft enough to cut with a knife. They are the most chemically active of all metals, readily forming ions with a single positive charge. This property is a result of their atoms having only a single, highly mobile electron in the outermost shell. Alkali metals react rapidly, sometimes violently, with both oxygen and water. Because of their reactivity, they always occur in nature in combination with other elements as simple and complex compounds. Pure alkali metals can be extracted from such compounds by the electrolysis of molten salts or hydroxides. Another process, known as thermal reduction, is also employed to obtain lithium and cesium. Figure 1: Modern version of the periodic table of the elements. To see more information about an any of the six chemical elements that make up Group I of the periodic tablenamely, lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). (See Figure.) The alkali metals are so called because they form alkalies (i.e., strong bases capable of neutralizing acids) when they combine with other elements. Sodium and potassium are the sixth and seventh most abundant of the elements, comprising, respectively, 2.6 and 2.4 percent of the Earth's crust. The other alkali metals are considerably more rare, with rubidium, lithium, and cesium, respectively, forming 0.03, 0.007, and 0.0007 percent of the Earth's crust. Francium, a natural radioactive isotope, is very rare and was not discovered until 1939. Because of their high reactivity, the alkali metals are never found as free metals in their natural state. They generally are found combined with other elements in the form of simple or complex compounds. The simpler compounds of the alkali metals are soluble in water and therefore are easily extracted and subjected to chemical operations for purposes of separation and purification. Minerals belonging to this class, such as halite (sodium chloride, NaCl), sylvite (potassium chloride, KCl), and carnallite (a potassium-magnesium chloride, KCl MgCl2 6H2O), although somewhat rare, are the most important commercial sources of the alkali metals; the more complex, water-soluble minerals are far more abundant in the Earth's crust. The alkali metals have all of the physical properties generally associated with metals, including silver-like lustre, high ductility, and excellent conductivity of electricity and heat. Lithium is the lightest metallic element. The alkali metals are low melting, ranging from a high of 179 C for lithium to a low of 28.5 C for cesium. Alloys of alkali metals exist that melt as low as -78 C. The alkali metals are extremely reactive and combine readily with most of the substances found in the atmosphere. (Only lithium, however, reacts with nitrogen.) The alkali metals all react vigorously, and often violently, with water, releasing hydrogen and forming strong caustic solutions. Most common nonmetallic substances such as the halogens, halogen acids, sulfur, and phosphorus react with the alkali metals. The alkali metals themselves react with many organic compounds, particularly those containing an active hydrogen atom (that is, one that can be replaced readily). Sodium is by far the most important alkali metal in terms of industrial use. It is employed in the reduction of organic compounds, in the preparation of sodium peroxide and sodium cyanide, and in the manufacture of tetraethyl lead. As a free metal, it is used as a heat transfer fluid in nuclear reactors. Its compounds, of the widest industrial importance and manufactured in hundreds of thousands of tons annually, include common salt (NaCl), baking soda (NaHCO3), soda ash (Na2CO3), and caustic soda (NaOH). Potassium has considerably less use than sodium as a free metal. Potassium salts, however, are consumed in considerable tonnages in the manufacture of fertilizers. Lithium metal is used in certain light-metal alloys, and it is employed as a reactant in organic syntheses. Lithium also has potential use as a battery anode in electrically propelled automobiles. Rubidium and cesium and their compounds have but limited use. Additional reading J.W. Mellor, A Comprehensive Treatise on Inorganic and Theoretical Chemistry, 16 vol. (192237), and supplements, is a comprehensive treatment devoted to the alkali metals and their compounds, including history, production methods, analysis, alloy systems, and physical, chemical, biological, and spectrochemical properties. Further information is available in J.W. Mausteller, F. Tepper, and S.J. Rodgers, Alkali Metal Handling System Operating Techniques (1968), on technology devoted chiefly to laboratory handling aspects, including a summary of physical properties, chemical properties of the pure metals, analytical techniques, handling, and safety procedures; Marshall Sittig, Sodium (1956), covering the manufacture, inorganic and organic reactions, physical properties, handling, and uses of sodium; The Alkali Metals (1967), 44 symposium papers covering most aspects of alkali metals, including preparation, chemical and physical properties, and compounds; and C.C. Addison, The Chemistry of the Liquid Alkali Metals (1984). The history of the alkali industry in Britain to 1926 is chronicled in Kenneth Warren, Chemical Foundations (1980). H.M. Emrich, J.B. Aldenhoff, and H. Dieter Lux, Basic Mechanisms in the Action of Lithium (1982), covers physiological effects. Jean-Paul Gabano (ed.), Lithium Batteries (1983), discusses an important commercial use. Frederick Tepper The Editors of the Encyclopdia Britannica