Alkaline-earth Metal | Properties, List, & Reactivity - Britannica

Occurrence, properties, and uses

Prior to the 19th century, substances that were nonmetallic, insoluble in water, and unchanged by fire were known as earths. Those earths, such as lime (calcium oxide), 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 2 (IIa) of the periodic table ever since Russian chemist Dmitry Mendeleyev proposed his first table in 1869.

As with the alkali metals of Group 1 (Ia), the atoms of the alkaline-earth metals easily lose electrons to become positive ions (cations). Most of their typical compounds are therefore ionic: salts in which the metal occurs as the cation M2+, where M represents any Group 2 atom. The salts are colourless unless they include a coloured anion (negative ion). The formulas of typical alkaline-earth compounds, such as calcium chloride (CaCl2) and calcium oxide (CaO), may be contrasted with the corresponding compounds of the alkali metals (which contain M+ ions), 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 actually amphoteric, rather than basic, whereas barium oxide is 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.

1 of 2

2 of 2

With the exception of radium, 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 (they are among the six most common elements on Earth) and play significant roles in geological and biological processes. Radium is a rare element, and all its isotopes are radioactive. There has never been commercial production of the metal, and, although its compounds were frequently used in the first half of the 20th century for cancer treatment, they have largely been superseded by less expensive alternatives.

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