ALKALINE-EARTH METAL

any of the six chemical elements that comprise Group IIa of the periodic table. The elements are beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). The designation earth for these metals derives from the Middle Ages when alchemists referred to substances that were insoluble in water and unchanged by fire as earths. Those earths, such as lime, that bore a resemblance to the alkalines (e.g., soda ash and potash) were called alkaline earths. By the early 1800s it became apparent that the earths, formerly regarded as elements, were in reality compounds of a metal and oxygeni.e., oxides. The metals whose oxides comprise the alkaline earths subsequently came to be known as the alkaline-earth metals. Magnesium, calcium, strontium, and barium were isolated as impure metals for the first time in 1808 by the English chemist Sir Humphry Davy. The alkaline-earth metals are basically grayish white in colour. They all are malleable but vary widely in hardness. Beryllium, for example, is hard enough to cut glass; barium, by contrast, is as soft as lead. Like other metals, the alkaline-earth elements are good conductors of electricity. Their melting points and boiling points, though varying in an irregular fashion, are higher than those of the corresponding alkali metals. The atoms of the alkaline-earth elements have a similar electronic structure, which consists of a pair of electrons in the outermost shell. These electrons are removed from their respective atoms relatively easily, and this ionization is the distinguishing chemical property of the alkaline-earth metals. These elements readily combine with most oxides and many nonmetals and thus never occur as pure metals in nature. Magnesium and calcium are the only abundant alkaline-earth elements in the crust of the Earth. They also are the most commercially important members of the family. Figure 1: Modern version of the periodic table of the elements. To see more information about an any of the six chemical elements that comprise Group IIa of the periodic table (see Figure). The elements are beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). Prior to the 19th century, substances that were nonmetallic, insoluble in water, and unchanged by fire were known as earths. Those earths, like lime, that resembled the alkalies (soda ash and potash) were designated alkaline earths. Alkaline earths were thus distinguished from the alkalies and from other earths, such as alumina and the rare earths. By the early 1800s it became clear that the earths, formerly considered to be elements, were in fact oxides, compounds of a metal and oxygen. The metals whose oxides make up the alkaline earths then came to be known as the alkaline-earth metals and have been classified in group II of the periodic table ever since Mendeleyev proposed his first table in 1869. The alkaline-earth metals are extremely electropositive; that is, like the alkali metals of group Ia, their atoms easily lose electrons to become positive ions (cations). Most of their typical compounds are therefore ionic: salts in which the metal occurs as uniformly divalent cations, M2+, where M represents any group IIa atom. The salts are colourless unless they include a coloured anion (negative ion). Typical alkaline-earth compounds, calcium chloride (CaCl2) and calcium oxide (CaO), may be contrasted with the compounds of the uniformly monovalent alkali metals, sodium chloride (NaCl) and sodium monoxide (Na2O). The oxides of the alkaline-earth metals are basic (i.e., alkaline, in contrast to acidic). A fairly steady increase in electropositive character is observed in passing from beryllium, the lightest member of the group, to radium, the heaviest; as a result of this trend, beryllium oxide is only weakly basic and even shows acidic properties, whereas barium and radium oxide are strongly basic. The metals themselves are highly reactive reducing agents; that is, they readily give up electrons to other substances that are, in the process, reduced. All the metals and their compounds find commercial application to some degree, especially magnesium alloys and a variety of calcium compounds. Magnesium and calcium, particularly the latter, are abundant in nature and play significant roles in geologic and biological processes. Radium is a rare element; all its isotopes are radioactive. Additional reading The occurrence, properties, and uses of the individual alkaline-earth elements and their more important compounds are given in Clifford A. Hampel (ed.), The Encyclopedia of the Chemical Elements (1968). Ulick R. Evans, Metals and Metallic Compounds, vol. 2 (1923), contains a detailed chapter on the Group IIa elements describing their physical properties, laboratory preparation, and common compounds. The general comparative and theoretical aspects of alkaline-earth chemistry are discussed in most modern textbooks of inorganic chemistry and particularly in C.S.G. Phillips and R.J.P. Williams, Inorganic Chemistry, 2 vol. (196566). Courtenay Stanley Goss Phillips The Editors of the Encyclopdia Britannica