A general les-grizzlys-catalans.orgistryles-grizzlys-catalans.orgTextmaporganized around the textbookles-grizzlys-catalans.orgistry: Principles, Patterns, and Applicationsby Bruce A. Averill
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Learning ObjectivesTo describe how to isolate the alkaline earth metals. To be familiar with the reactions, compounds, and complexes of the alkaline earth metals.
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Like the alkali metals, the alkaline earth metals are so reactive that they are never found in elemental form in nature. Because they form +2 ions that have very negative reduction potentials, large amounts of energy are needed to isolate them from their ores. Four of the six group 2 elements—magnesium (Mg), calcium (Ca), strontium (Sr), and barium (Ba)—were first isolated in the early 19th century by Sir Humphry Davy, using a technique similar to the one he used to obtain the first alkali metals. In contrast to the alkali metals, however, compounds of the alkaline earth metals had been recognized as unique for many centuries. In fact, the name alkali comes from the Arabic al-qili, meaning “ashes,” which were known to neutralize acids. Medieval alles-grizzlys-catalans.orgists found that a portion of the ashes would melt on heating, and these substances were later identified as the carbonates of sodium and potassium (\(M_2CO_3\)). The ashes that did not melt (but did dissolve in acid), originally called alkaline earths, were subsequently identified as the alkaline earth oxides (MO). In 1808, Davy was able to obtain pure samples of Mg, Ca, Sr, and Ba by electrolysis of their chlorides or oxides.
Beryllium (Be), the lightest alkaline earth metal, was first obtained in 1828 by Friedrich Wöhler in Germany and simultaneously by Antoine Bussy in France. The method used by both men was reduction of the chloride by the potent “new” reductant, potassium:
Radium was discovered in 1898 by Pierre and Marie Curie, who processed tons of residue from uranium mines to obtain about 120 mg of almost pure \(RaCl_2\). Marie Curie was awarded the Nobel Prize in les-grizzlys-catalans.orgistry in 1911 for its discovery. Because of its low abundance and high radioactivity however, radium has few uses.
Preparation of the Alkaline Earth Metals
The alkaline earth metals are produced for industrial use by electrolytic reduction of their molten chlorides, as indicated in this equation for calcium:
The group 2 metal chlorides are obtained from a variety of sources. For example, \(BeCl_2\) is produced by reacting \(HCl\) with beryllia (\(BeO\)), which is obtained from the semiprecious stone beryl \(
les-grizzlys-catalans.orgical reductants can also be used to obtain the group 2 elements. For example, magnesium is produced on a large scale by heating a form of limestone called dolomite (CaCO3·MgCO3) with an inexpensive iron/silicon alloy at 1150°C. Initially \(CO_2\) is released, leaving behind a mixture of \(CaO\) and MgO; Mg2+ is then reduced:
\<2CaO·MgO_(s) + Fe/Si_(s) \rightarrow 2Mg(l) + Ca_2SiO_4\;(s) + Fe(s) \labelEq3\>
An early source of magnesium was an ore called magnesite (\(MgCO_3\)) from the district of northern Greece called Magnesia. Strontium was obtained from strontianite (\(SrCO_3\)) found in a lead mine in the town of Strontian in Scotland. The alkaline earth metals are somewhat easier to isolate from their ores, as compared to the alkali metals, because their carbonate and some sulfate and hydroxide salts are insoluble.
General Properties of the Alkaline Earth Metals
Several important properties of the alkaline earth metals are summarized in Table \(\PageIndex1\). Although many of these properties are similar to those of the alkali metals (Table \(\PageIndex1\)), certain key differences are attributable to the differences in the valence electron configurations of the two groups (ns2 for the alkaline earth metals versus ns1 for the alkali metals).
|valence electron configuration||2s2||3s2||4s2||5s2||6s2||7s2|
|melting point/boiling point (°C)||1287/2471||650/1090||842/1484||777/1382||727/1897||700/—|
|density (g/cm3) at 25°C||1.85||1.74||1.54||2.64||3.62||~5|
|atomic radius (pm)||112||145||194||219||253||—|
|first ionization energy (kJ/mol)||900||738||590||549||503||—|
|most common oxidation state||+2||+2||+2||+2||+2||+2|
|ionic radius (pm)*||45||72||100||118||135||—|
|electron affinity (kJ/mol)||≥ 0||≥ 0||−2||−5||−14||—|
|standard electrode potential (E°, V)||−1.85||−2.37||−2.87||−2.90||−2.91||−2.8|
|product of reaction with O2||BeO||MgO||CaO||SrO||BaO2||—|
|type of oxide||amphoteric||weakly basic||basic||basic||basic||—|
|product of reaction with N2||none||Mg3N2||Ca3N2||Sr3N2||Ba3N2||—|
|product of reaction with X2||BeX2||MgX2||CaX2||SrX2||BaX2||—|
|product of reaction with H2||none||MgH2||CaH2||SrH2||BaH2||—|
As with the alkali metals, the atomic and ionic radii of the alkaline earth metals increase smoothly from Be to Ba, and the ionization energies decrease. As we would expect, the first ionization energy of an alkaline earth metal, with an ns2 valence electron configuration, is always significantly greater than that of the alkali metal immediately preceding it. The group 2 elements do exhibit some anomalies, however. For example, the density of Ca is less than that of Be and Mg, the two lightest members of the group, and Mg has the lowest melting and boiling points. In contrast to the alkali metals, the heaviest alkaline earth metal (Ba) is the strongest reductant, and the lightest (Be) is the weakest. The standard electrode potentials of Ca and Sr are not very different from that of Ba, indicating that the opposing trends in ionization energies and hydration energies are of roughly equal importance.
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One major difference between the group 1 and group 2 elements is their electron affinities. With their half-filled ns orbitals, the alkali metals have a significant affinity for an additional electron. In contrast, the alkaline earth metals generally have little or no tendency to accept an additional electron because their ns valence orbitals are already full; an added electron would have to occupy one of the vacant np orbitals, which are much higher in energy.