# Group IIA: Alkaline Earths

Submitted by ChemPRIME Staff on Thu, 12/16/2010 - 14:33

GroupThose elements that comprise a single column of the periodic table. Also called family. IIA also known as the alkaline-earth metals, include beryllium, magnesium, calcium, strontium, barium, and radium. The last member of the group, Ra, is radioactiveDescribes a substance that gives off radiation‐alpha particles, beta particles, or gamma rays‐by the disintegration of its nucleus. and will not be considered here. All alkaline earths are silvery-gray metals which are ductileAble to be drawn into a wire; usually applied to metals, of which ductility is a characteristic property. and relatively soft. However, the following table shows that they are much denser than the group IA metals, and their melting points are significantly higher. They are also harder than the alkali metals. This may be attributed to the general valence electronIn a neutral atom, any of the electrons found in the highest occupied shell as well as any electrons in incompletely filled subshells of lower shells. configurationThe 3-D arrangement of atoms about a chiral center in a molecule. ns2 for the alkaline earths, which involves two electrons per metal atom in metallic bonding (instead of just one as in an alkali metal).

Properties of the Group IIA Alkaline-Earth Metals

 Element Symbol Electron Configuration Usual Oxidation State Radius/pm Atomic Ionic (M2+) Beryllium Be [He]2s2 +2 89 31 Magnesium Mg [Ne]3s2 +2 136 65 Calcium Ca [Ar]4s2 +2 174 99 Strontium Sr [Kr]5s2 +2 191 113 Barium Ba [Xe]6s2 +2 198 135

 Symbol Ionization Energy/MJ mol–1 Density/ g cm–3 Electro- negativity Melting Point (in °C) First Second Third Be 0.906 1.763 14.86 1.86 1.5 1278 Mg 0.744 1.467 7.739 1.74 1.2 651 Ca 0.596 1.152 4.918 1.54 1.0 839 Sr 0.556 1.071 4.21 2.60 1.0 769 Ba 0.509 0.972 3.43 3.51 0.9 725

First and second ionizationA process in which an atom, molecule, or negative ion loses an electron; a process in which a covalent molecule reacts with a solvent to form positive and negative ions; for example, a weak acid reacting with water to form its conjugate base (an anion) and a hydrogen (hydronium) ion. energies for the alkaline earths (corresponding to removal of the first and second valence electrons) are relatively small, but the disruption of an octetA stable set of eight electrons in the valence shell of an atom. Each noble-gas atom has an octet. by removal of a third electron is far more difficult. Like the alkali metals, the alkaline-earth atomsThe smallest particle of an element that can be involved in chemical combination with another element; an atom consists of protons and neutrons in a tiny, very dense nucleus, surrounded by electrons, which occupy most of its volume. lose electrons easily, and so they are good reducing agents. Other trends among the data in the table are what we would expect. Ionization energies and electronegativities decrease from top to bottom of the group, and atomic and ionic radii increase. The radii of +2 alkaline-earth ions are much smaller than the +1 alkali-metal ions of the same periodThose elements from a single row of the periodic table., because the greater nuclear charge holds the inner shells more tightly. This effect is sufficiently large that an alkaline earth below and to the right of a given alkali metal in the periodic tableA chart showing the symbols of the elements arranged in order by atomic number and having chemically related elements appearing in columns. often has nearly the same ionic radiusAn estimate of the size of an ion in an ionic compound; found from the internuclear distance between ions in a crystal lattice.. Thus Na+ (95 pm), can fit into exactly the same type of crystal latticeAn orderly, repeating arrangement of points in 3-D space in which each p;oint has surroundings identical to every other point. A crystal's constituent atoms, molecules, and ions are arranged about each lattice point. as Ca2+ (99 pm), and these two elements are often found in the same minerals. The same is true of K+ and Ba2+. Below is the table for alkali metals, to compare with the table of alkaline earth metals.

Properties of the Group IA Alkali Metals

 Element Symbol Electron Configuration Usual Oxidation State Radius/pm Atomic Ionic (M+) Lithium Li [He]2s1 +1 122 60 Sodium Na [Ne]3s1 +1 157 95 Potassium K [Ar]4s1 +1 202 133 Rubidium Rb [Kr]5s1 +1 216 148 Cesium Cs [Xe]6s1 +1 235 169

 Symbol Ionization Energy/MJ mol–1 Density/ g cm–3 Electro- negativity Melting Point (in °C) First Second Li 0.526 7.305 0.534 1.0 179 Na 0.502 4.569 0.97 0.9 98 K 0.425 3.058 0.86 0.8 64 Rb 0.409 2.638 1.52 0.8 39 Cs 0.382 2.430 1.87 0.7 28

Similarity of ionic radii also leads to related properties for Li and Mg. Since these two elements are adjacent along a diagonal line from the upper left to the lower right in the periodic table, their similarity is called a diagonal relationshipIn the periodic table, similarities in chemical or physical properties of two elements (or their compounds); the second element is either above and to the left or below and to the right of the first element (on a diagonal in the periodic table).. Diagonal relationships are mainly evident in the second and third periods: Be is similar to Al, and B is like Si in many ways.

Farther toward the right-hand side of the table such relationships are less pronounced. The most striking similarity between Li and Mg is their ability to form covalent bonds with elements of average electronegativityThe tendency of an atom (nucleus and core electrons) within a molecule to attract electrons in bonds., such as C, while forming fairly ionic compounds with more electronegative elements, such as O or F. Two examples of covalent compounds are ethyllithium, CH3CH2Li, and diethylmagnesium, (CH3CH2)2Mg. Such compounds are likely in the case of Li and Mg but not the alkali or alkaline earths below them, because Li+ and Mg2+ are small enough to be strongly polarizing and thus form bonds with considerable covalent character.

Chemical Reactions and Compounds

The alkaline earth metals react directly with most nonmetallic elements. forming Except for beryllium, the alkaline earths react directly with hydrogen gasA state of matter in which a substance occupies the full volume of its container and changes shape to match the shape of the container. In a gas the distance between particles is much greater than the diameters of the particles themselves; hence the distances between particles can change as necessary so that the matter uniformly occupies its container. to form hydrides, MH2; M = Mg, Ca, Sr, Ba, or Ra. Beryllium hydride, BeH2 can also be prepared, but not directly from the elements. Alkaline-earth metals combine readily with oxygen from the air to form oxides, MO. This follows the general reaction:

2M(s) + O2(g) → 2MO2(s)      M = Be, Mg, Ca, Sr, Ba, or Ra      (1)

The following video shows the reaction of magnesium with oxygen:

In the video, magnesium is burned in air, and emits a bright white flame. A white powder of MgO remains after the reaction described by the equation:

2Mg(s) + O2(g) → 2MgO2(s)

It should also be noted that while MgO is the main productA substance produced by a chemical reaction., nitrogen is also present in the air, and so some magnesium nitride is also produced according to the chemical equationA representation of a chemical reaction in which chemical symbols represent reactants on the left side and products on the right side.:

3Mg(s) + N2(g) → Mg3N2(s)

These oxides will coat the surface of the metal and prevent other substances from contacting and reacting with it. A good example of the effect of such an oxide coating is the reaction of alkaline-earth metals with water. Beryllium and magnesium react much more slowly than the others because their oxides are insolubleUnable to dissolve appreciably in a solvent. and prevent water from contacting the metal.

Alkaline-earth metals react directly with halogens to form salts:

M(s) + Cl2(g) → MCl2(s)      M = Be, Mg, Ca, Sr, Ba, or Ra      (2)

Salt obtained by evaporating seawater (sea salt) contains a good deal of magnesium chloride and calcium chloride as well as sodium chloride. It also has small traces of iodide salts, accounting for the absence of simple goiter in communities which obtain their salt from the oceans. Simple goiter is an enlargement of the thyroid gland caused by iodine deficiency.

Alkaline earths also form sulfides: MS. In all these compounds the alkaline-earth elements occur as dipositive ions, Mg2+, Ca2+, Sr2+, or Ba2+.

Similar compounds of Be can be formed by roundabout means, but not by direct combination of the elements. Moreover, the Be compounds are more covalent than ionic. The Be2+ ion has a very small radius (31 pm) and is therefore capable of distorting (polarizing) the electron cloud of an anionA negatively charged ion. An ion that is attracted toward the anode in an electrolytic cell. in its vicinity. Therefore all bonds involving Be have considerable covalent character, and the chemistry of Be is significantly different from that of the other members of group IIA.

As in the case of the alkali metals, the most important and abundant alkaline earths, Mg and Ca, are in the third and fourth periods. Be is rare, although its strength and low density make it useful in certain special alloys. Sr and Ba occur naturally as the relatively insoluble sulfates SrSO4 (strontianite) and BaSO4 (barite), but these two elements are of minor commercial importance.

The most common ores of Mg and Ca are dolomite, MgCO3•CaCO3, after which an entire mountain range in Italy is named, and limestone, CaCO3, an important building material. Mg is also recovered from seawater on a wide scale. The oxides of the alkaline earths are commonly obtained by heating the carbonates. For example, lime, CaO, is obtained from limestone as follows:

CaCO3(s) $\xrightarrow{\Delta }$ CaO(s) + CO2(g)

Except for BeO, which is covalently bonded, alkaline-earth oxides contain O2– ions and are strongly basic. When treated with water (a process known as slaking), they are converted to hydroxides:

CaO(s) + H2O(l) → Ca(OH)2(s)

Ca(OH)2 (slaked lime) is an important strong baseA base that dissociates completely or ionizes completely in a particular solvent. for industrial applications, because it is cheaper than NaOH.

MgO has an extremely high melting point (2800°C) because of the close approach and large charges of its constituent Mg2+ and O2– ions in the crystalA solid with a regular polyhedral shape; for example, in sodium chloride (table salt) the crystal faces are all at 90° angles. A solid in which the atoms, molecules, or ions are arranged in a regular, repeating lattice structure. lattice. As a solidA state of matter having a specific shape and volume and in which the particles do not readily change their relative positions. it is a good electrical insulator, and so it is used to surround metal-resistance heating wires in electric ranges. MgO is also used to line high-temperatureA physical property that indicates whether one object can transfer thermal energy to another object. furnaces. When converted to the hydroxide, Mg finds a different use. Mg(OH)2 is quite insoluble in water, and so it does not produce a high enough concentrationA measure of the ratio of the quantity of a substance to the quantity of solvent, solution, or ore. Also, the process of making something more concentrated. of hydroxide ions to be caustic. It is basic, however, and gramOne thousandth of a kilogram. for gram can neutralize nearly twice the quantity of acidIn Arrhenius theory, a substance that produces hydrogen ions (hydronium ions) in aqueous solution. In Bronsted-Lowry theory, a hydrogen-ion (proton) donor. In Lewis theory, a species that accepts a pair of electrons to form a covalent bond. that NaOH can. Consequently a suspension of Mg(OH)2 in water (milk of magnesia) makes an excellent antacid, for those who can stand its taste.

Because the carbonate ion behaves as a Brönstedt-Lowry base, carbonate salts dissolve in acidic solutions. In nature, water often becomes acidic because the acidic oxide CO2 is present in the atmosphereA unit of pressure equal to 101.325 kPa or 760 mmHg; abbreviated atm. Also, the mixture of gases surrounding the earth.. When CO2 from the air dissolves in water, it can help dissolve limestone:

CO2(g) + H2O(l) + CaCO3(s) $\rightleftharpoons$ Ca2+(aq) + HCO3(aq)

This reaction often occurs underground as rainwater saturatedDescribes 1) a solution that contains the equilibrium concentration of a solute, or 2) an organic compound that contains no double or triple bonds (such as an alkane). with CO2 seeps through a layer of limestone. Caves from which the limestone has been dissolved are often prevalent in areas where there are large deposits of CaCO3. In addition, the groundwater and well water in such areas becomes hard. Hard waterWater containing high concentrations of cations having charge greater than +1; hardness can be removed by ion exchange. contains appreciable concentrations of Ca2+, Mg2+ , and certain other metal ions. These form insoluble compounds with soapA salt of a fatty acid produced by the saponification of fat., causing curdy, scummy precipitates. Hard water can be softened by adding Na2CO3, washing soda, which precipitates CaCO3, or by ion exchangeThe replacement of ions by other ions, usually on the surface of a resin designed as a reservoir for ions., a process in which the undesirable Ca2+ and Mg2+ ions are replaced in solution by Na+ ions, which do not precipitate soap. Most home water softeners work on the latter principle.