alkaline-earth metal

/al"keuh luyn'errth", -lin-/, Chem.

any of the group of bivalent metals including barium, radium, strontium, calcium, and, usually, magnesium, the hydroxides of which are alkalis but less soluble than those of the alkali metals.

[1900-05]

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▪ chemical element

Introduction

 any of the six chemical elements that comprise Group 2 (IIa) of the periodic (periodic law) 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 2 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 1 (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 the cation M²⁺, where M represents any Group 2 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 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 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.

History

      The earliest known alkaline earth was lime (Latin: calx), which is now known to be calcium oxide; it was used in ancient times in the composition of mortar. Magnesia (the name derives probably from the ancient district of Magnesia in Asia Minor), the oxide of magnesium, was shown to be an alkaline earth different from lime by the Scottish chemist Joseph Black (Black, Joseph) in 1755; he observed that magnesia gave rise to a soluble sulfate, whereas that derived from lime was known to be insoluble. In 1774 Carl Wilhelm Scheele (Scheele, Carl Wilhelm), the Swedish chemist who discovered oxygen, found that the mineral called heavy spar or barys (Greek: heavy) contained a new earth, which became known as baryta ( barium oxide). A further earth, strontia ( strontium oxide), was identified by the London chemists Adair Crawford and William Cruickshank in 1790 on examining a mineral (strontium carbonate) found in a lead mine at Strontian in Argyllshire, Scotland. Beryllia ( beryllium oxide) was extracted from the mineral beryl and recognized as an earth by the French analytical chemist Nicolas-Louis Vauquelin (Vauquelin, Nicolas-Louis) in 1798. Though at first confused with alumina (aluminum oxide) because both dissolve in alkali, beryllia was shown to be distinct; unlike alumina, it reprecipitated when the alkaline solution was boiled for some time. Beryllia was originally called glucina (Greek glykys, sweet) because of its sweet taste. (This etymological root is retained in France, where the element beryllium is also known as glucinium.)

      Magnesium, calcium, strontium, and barium—elements derived from alkaline earths—were isolated as impure metals by Sir Humphry Davy (Davy, Sir Humphry, Baronet) in 1808 by means of the electrolytic method he had previously used for isolating the alkali metals potassium and sodium. The alkaline-earth metals were later produced by reduction of their salts with free alkali metals, and it was in this way (the action of potassium on beryllium chloride) that beryllium was first isolated by the German chemist Friedrich Wöhler (Wöhler, Friedrich) and the French chemist Antoine Bussy independently in 1828. radium was discovered in 1898 by means of its radioactivity by Pierre and Marie Curie, who separated it from barium.

Physical and chemical behaviour

       Some properties of the alkaline metalsThe alkaline-earth elements are highly metallic and are good conductors of electricity. They have a gray-white lustre when freshly cut but tarnish readily in air, particularly the heavier members of the group. beryllium is sufficiently hard to scratch glass, but barium is only slightly harder than lead. The melting points (melting point) and boiling points (boiling point) of the group (see Table (Some properties of the alkaline metals)) are higher than those of the corresponding alkali metals; they vary in an irregular fashion, magnesium having the lowest (mp 650° C and bp 1,105° C) and beryllium the highest (mp 1,283° and bp about 2,500°). The elements crystallize in one or more of the three regular close-packed metallic crystal forms. Chemically, they are all strong reducing agents. The free metals are soluble in liquid ammonia—the dark-blue solutions of calcium, strontium, and barium arousing considerable interest because they are thought to contain metal ions and the most unusual species, solvated electrons, or electrons (electronegativity) resulting from the interaction of the metal and the solvent. Highly concentrated solutions of these elements have a metallic, copper-like appearance, and further evaporation yields residues containing ammonia, which correspond to the general formula M(NH₃)₆. The solutions are strong reducing agents and are useful in a number of chemical processes.

      The atoms of the alkaline-earth elements all have similar electronic structures (electronic configuration), consisting of a pair of electrons (designated s electrons) in an outermost orbital, within which is a stable electronic configuration corresponding to that of a noble gas. The noble gas elements—helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn)—have generally complete electron shells. Strontium has the configuration 1s²2s²2p⁶3s² 3p⁶3d¹⁰4s²4p⁶5s², which may be written as (krypton core) 5s², or simply [Kr] 5s². Similarly, Be may be designated as [He] 2s², Mg as [Ne] 3s², Ca as [Ar] 4s², Ba as [Xe] 6s², and Ra as [Rn] 7s². The prominent lines in the atomic spectra of the elements, obtained when the elements are heated under certain conditions, arise from states of the atom in which one of the two s electrons has been promoted to a higher energy orbital. The s electrons are relatively easily ionized (removed from the atom), and this ionization is the characteristic feature of alkaline-earth chemistry. The ionization energy (the energy required to strip an electron from the atom) falls continuously in the series from beryllium (9.32 electron volts [eV]) to barium (5.21 eV); radium, the heaviest in the group, has a slightly higher ionization energy (5.28 eV). The small irregularities observed in the otherwise smooth change as one proceeds down the group as it appears in the periodic table are explained by the uneven filling of electron shells in the successive rows of the table. The s electrons may also be promoted to p orbitals of the same principal quantum number (within the same shell) by energies similar to those required to form chemical bonds; the atoms are, therefore, able to form stable covalently bonded structures, unlike helium, which has the otherwise analogous electronic configuration of 1s².

      Zinc (zinc group element), cadmium, and mercury, the Group 12 (IIb) elements, are often compared with the alkaline-earth elements calcium, strontium, and barium. Cadmium, for example, has the electronic configuration [Kr] 4d¹⁰5s², with the ten 4d electrons taking virtually no part in chemical bonding. The 5s² electrons, however, are much less readily ionized in cadmium than they are in strontium, for the 4d electrons act as an ineffective shield for the corresponding increased charge on the cadmium nucleus. The chemistry of the IIb metals, therefore, is markedly less ionic than the chemistry of the alkaline-earth metals.

Ionic (ionic bond) character

      The chemistry of the alkaline-earth metals, like that of the alkali metals, is for the most part reasonably interpreted in terms of an ionic model for the compounds formed. This model is less satisfactory for the chemistry of beryllium and magnesium than for the heavier alkaline-earth metals. In fact, most beryllium compounds are molecular (covalent (valence)) rather than ionic. Although there is some evidence for the transient existence of singly charged alkaline-earth ions, in most cases the chemistry of these elements is dominated by the formation and properties of the doubly charged M²⁺ ions, in which the outermost s electrons have been stripped from the metal atom. The resulting ion is stabilized by electrostatic interaction with a solvent, like water, which has a high dielectric constant and a great ability to absorb electrical charge, or by combination with ions of opposite charge in an ionic lattice such as is found in salts. The extra energy required to remove the second s electron (the second ionization energy being approximately twice the first) is more than compensated for by the extra binding energy present in the doubly charged ion. The removal of a third electron from an alkaline-earth atom, however, would require an expenditure of energy greater than could be recouped from any chemical environment. As a result, therefore, the alkaline-earth metals show a constant oxidation state of +2 in their compounds.

      The sizes of the ions of the alkaline-earth elements increase steadily from Be²⁺, which has a radius of 0.31Å or 31 × 10⁻¹⁰ cm, to Ra²⁺ with a radius of 1.40Å. The ionic radius of the europium ion Eu²⁺ (1.12Å) is very close to that of strontium Sr²⁺ (1.13Å); this means that Eu²⁺ ions can sometimes be used as a “probe” for the alkaline-earth metals, substituting for strontium ions in situations where advantage can be taken of the spectroscopic and magnetic properties that make Eu²⁺ readily identifiable. The ionic radius of the cadmium ion Cd²⁺ (0.97Å) is very similar to that of calcium Ca²⁺ (0.99Å). A quantitative comparison of cadmium and calcium chemistry, therefore, clearly shows up the less ionic character of cadmium chemistry without complications due to differences in ionic size. A related comparison may be made between mercury and strontium, because of the similar ionic radii of Hg²⁺ (1.10Å) and Sr²⁺ (1.13Å).

      The chemistry of radium is less well investigated than that of the other alkaline-earth metals. As expected, however, it is in general an extrapolation of the chemistry of calcium, strontium, and, especially, barium.

      The Group 2 ions are readily hydrated, with the strength of bonding to the water molecules increasing with decreasing ionic radius. Although the number of water molecules directly attached to the metal ion may be greater with the larger ions for purely steric (geometrical) reasons, the total number of water molecules associated with the metal ion nevertheless increases inversely with the size of the ion itself (as shown by migration experiments conducted in aqueous solution). Large anions, such as sulfate, tend to form weak ion-pair complexes more readily with the larger metal ions of the family, but weak-acid anions, such as acetate, tend to form stronger complexes with the smaller metal ions, particularly those of magnesium and beryllium. That many of these complexes are molecular rather than ionic is shown by their ready extraction from aqueous solution (which preferentially dissolves ionic substances) into organic solvents (which dissolve molecular ones).

Courtenay Stanley Goss Phillips

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. (1965–66).Courtenay Stanley Goss Phillips Ed.

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