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The Alkaline Earths As Metals

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The Alkaline Earths As Metals CHAPTER 1 The Alkaline Earths as Metals OUTLINE 1.1. General Properties 1 1.2.3. Calcium 12 1.2.4. Strontium 15 1.2. Properties of the Alkaline Earth 1.2.5. Barium 18 Metals 4 1.2.6. Radium 19 1.2.1. Beryllium 4 1.2.2. Magnesium 8 The alkaline earth metals comprise Group 2 of the 1.1. GENERAL PROPERTIES periodic table and include the elements Be, Mg, Ca, Sr, Ba and Ra. These metals form cations with a formal Like other groups, the members of this family show charge of þ2 in solution and are the second most electro- specific patterns in their electron configuration, espe- positive metals of all of the elements (the alkali metals cially the outermost shells, that results in trends in are the most electropositive). The name of this specific chemical behavior (Table 1.1). group in the periodic table stems from the fact that their Another way to depict the electronic structure of oxides produce basic alkaline solutions and that these these elements is shown in Table 1.2. elements melt at such high temperatures that they All of the alkaline earth metals have two electrons in remain solid (earths) in fires. The alkaline earth metals their outer valence shell, so the energetically preferred provide a good example of group trends in chemical state of achieving a filled electron shell is to lose two þ properties within the periodic table, with well-character- electrons to form doubly charged cations, M2 . The alka- ized homologous behavior as one goes down the group. line earth metals are silver-colored, soft metals that react With the exception of Be and Mg, the metals have readily with halogens to form ionic salts. They also react a distinguishable flame color,ELSEVIER orange-red for Ca, with water, though not as rapidly as the alkali metals, to magenta-red for Sr, green for Ba and crimson-red for Ra. form strongly alkaline (basic) hydroxides. For example, TABLE 1.1 TABLE 1.2 Z Element No. of electrons/shell Element Symbol Electronic configuration 4 Beryllium 2, 2 Beryllium Be [He]2s2 12 Magnesium 2, 8, 2 Magnesium Mg [Ne]3s2 20 Calcium 2, 8, 8, 2 Calcium Ca [Ar]4s2 38 Strontium 2, 8, 18, 8, 2 Strontium Sr [Kr]5s2 56 Barium 2, 8, 18, 18, 8, 2 Barium Ba [Xe]6s2 88 Radium 2, 8, 18, 32, 18, 8, 2 Radium Ra [Rn]7s2 Encyclopedia of the Alkaline Earth Compounds http://dx.doi.org/10.1016/B978-0-444-59550-8.00001-6 1 Copyright Ó 2013 Elsevier B.V. All rights reserved. 2 1. THE ALKALINE EARTHS AS METALS whereas Na and K react with water at room tempera- These elements are all found in the Earth’s crust, but ture, Mg reacts only with steam and Ca with hot water: not in the elemental form because they are so reactive. Instead, they are widely distributed in rock structures. MgðsolidÞþ2HOðgasÞ 2 The main minerals in which magnesium is found are 0 MgðOHÞ ðsolidÞþH ðgasÞ 2 2 “Carnellite”, “Magnesite” and “Dolomite”. Calcium is Be is an exception. It does not react with water or found in “Chalk”, “Limestone”, “Gypsum” and “Anhy- steam, and its halides are covalent. drite”. Magnesium is the eighth most abundant element The alkaline earth metals are named after their oxides, in the Earth’s crust, and calcium is the fifth. the alkaline earths, whose old-fashioned names were Some of the physical properties of the alkaline earth Beryllia, Magnesia, Lime, Strontia and Baryta. “Earth” metals are shown in Table 1.3. is the old term applied by early chemists to nonmetallic The metals of Group 2 are harder and denser than substances that were insoluble in water and resistant to sodium and potassium, and have higher melting points. heating, properties shared by these oxides. The realiza- These properties are due largely to the presence of two tion that these earths were not elements but compounds valence electrons on each atom, which leads to stronger is attributed to the chemist Antoine Lavoisier. In his metallic bonding than occurs in Group 1. “Traite´ E´ lementaire de Chemie” (Elements of Chemistry) Three of these elements give characteristic colors of 1789, he called them “salt-forming” earth elements. when heated in a flame: Later, he suggested that the alkaline earths might be metal oxides, but admitted that this was mere conjecture. Mg ¼ brilliant white Ca ¼ brick À red In 1808, acting on Lavoisier’s idea, Humphrey Davy Sr ¼ crimson became the first to obtain samples of the metals by elec- trolysis of their molten “earths”. In all their compounds, these metals have an oxidation þ If the alkaline earths are compared to the alkalis, number of 2 and, with few exceptions, their compounds many similarities are apparent. The main difference is are ionic in nature. The reason for this can be seen by the electron configuration, which is ns2 for alkaline earth examination of the electron configuration, which always metals and ns1 for alkali metals. But for the alkaline has two electrons in an outer quantum level. These elec- earth metals, the nucleus also contains an additional trons are relatively easy to remove, but removing the positive charge. Also, the elements of Group 2 (alkaline third electron is much more difficult, as it is close to the nucleus and in a filled quantum shell. This results in earths) have much higher melting points and boiling 2þ points compared to those of Group 1 (alkali metals). the formation of M . The ionization energies reflect The alkalis are softer and more lightweight than the this electron arrangement. The first two ionization ener- alkaline earth metals that are much harder and denser. gies are relatively low, and the third very much higher. The second valence electron is very important when it In general, the chemical properties of Group 2 comes to comparing chemical properties of the alkaline elements are dominated by the strong reducing power earth and the alkali metals. The second valence electron of the metals. The elements become increasingly electro- is in the same “sublevel” as the first valence electron. positive as one descends within the Group. In direct contact with oxygen or chlorine gas, little or no reaction Therefore, the Zeff is much greater. This means that the elements of Group 2 have a smaller atomic radius and occurs. However, once started, the reactions with much higher ionization energy than those of Group 1. oxygen and chlorine are vigorous: Even though the Group 2 contains a much higher ioniza- ð Þþ ð Þ 0 ð Þþ tion energy, they still form ionic compoundsELSEVIER containing 2Mg solid O2 g 2MgO solid heat þ ð Þþ ð Þ 0 ð Þþ 2 cations. Beryllium, however, behaves differently. Ca solid Cl2 gas CaCl2 solid heat This is due to the fact that in order to remove two elec- trons from this particular atom, significantly more All the metals except beryllium form oxide layers in þ energy is required. It never forms the Be2 cation and air at room temperature that dulls the surface of the its bonds are polar covalent. metal. Barium is so reactive that it is stored under oil. Atomic and ionic radii increase smoothly down the All of the metals except beryllium reduce water and Group. The ionic radii are all much smaller than the cor- dilute acids to hydrogen: responding atomic radii. This arises because the atom ð Þþ þð Þ 0 ð Þþ ð Þ contains two electrons in an s level relatively far from Mg solid 2H aq Mg aq H2 gas the nucleus. It is these electrons that are removed to form the ion. Remaining electrons are thus in levels Magnesium reacts only slowly with water unless the closer to the nucleus, and in addition the increased effec- water is boiling, but calcium reacts rapidly even at room tive nuclear charge attracts the electrons toward the temperature, and forms a cloudy white suspension of nucleus and decreases the size of the ion. sparingly soluble calcium hydroxide. 1.1. GENERAL PROPERTIES 3 TABLE 1.3 Element Atomic number Relative atomic mass Melting point, C Density in kgm/m3 Be 4 9.012 1551 1847.7 Mg 12 24.31 922 1738 Ca 20 40.08 1112 1550 Sr 38 87.62 1042 2540 Ba 56 137.33 1002 3594 Ionization energies in kJ/mol 1st 2nd 3rd Be 899.4 1757.1 14,848 Mg 737.7 1450.7 7732.6 Ca 589.7 1145 4910 Sr 549.5 1064.2 4210 Ba 502.8 965.1 3600 Standard electrode D Atomic radius/A˚ Ionic radius/A˚ (M2 ) potentials/V Be 1.13 0.34 À1.85 Mg 1.60 0.78 2.36 Ca 1.97 1.06 À2.87 Sr 2.15 1.27 À2.89 Ba 2.17 1.43 À2.90 Calcium, strontium and barium can reduce hydrogen Calcium hydroxide is known as “slaked lime”. It is gas when heated, forming the hydride: sparingly soluble in water and the resulting mildly alka- line solution is known as “limewater” which is used to ð Þþ ð Þ 0 ð Þ Ca solid H2 gas CaH2 solid test for the acidic gas, carbon dioxide. The Group 2 halides are normally found in the The hot metals are also sufficiently strong reducing hydrated form. They are all ionic except beryllium chlo- agents to reduce nitrogen gas and form nitrides: ride. Anhydrous calcium chloride has such a strong ð Þþ ð Þ 0 ð Þ 3Mg solid N2 gas Mg3N2 solid affinity for water that it is used as a drying agent.
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