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Chapter 12

Chapter 12. The Alkaline Earth Metals. Alkaline Earth Properties. harder, denser, and less reactive than the alkali metals less dense and more reactive than a transition metal Beryllium acts as a semimetal Radium is radioactive. Group Trends. silvery and of fairly low density

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Chapter 12

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  1. Chapter 12 The Alkaline Earth Metals

  2. Alkaline Earth Properties • harder, denser, and less reactive than the alkali metals • less dense and more reactive than a transition metal • Beryllium acts as a semimetal • Radium is radioactive

  3. Group Trends • silvery and of fairly low density • density increases down the group

  4. Group Trends • greater enthalpies of atomization than the alkali metals • stronger metallic bonding • also harder and have higher melting points

  5. Group Trends • Less chemically reactive than the alkalis • calcium, strontium, and barium react with water Ba(s) + 2H2O(l)  Ba(OH)2(aq) + H2(g) • magnesium reacts with hot water • Also will react with nonmetal diatomics Ca(s) + Cl2(g)  CaCl2(s) 3Mg(s) + N2(g)  Mg3N2(s)

  6. Ionic Character • Oxidation number is always +2 • Compounds are stable, colorless, ionic solids • unless with a colored anion • Bonds tend to be mainly ionic • exceptions found in beryllium and magnesium

  7. Ion Hydration • Salts are almost always hydrated • As charge density decreases, so does the hydration number

  8. Solubility • Many compounds are insoluble in water • Mononegative anions tend to be soluble • hydroxides • insoluble  soluble • Dinegative and trinegative anions tend to be insoluble • sulfates • soluble  insoluble

  9. Solubility • Enthalpy considerations • electrostatic attractions are much more for a dipositive alkaline earth cation than the monopositive alkali metal cations • higher charge densities make the hydration enthalpy much more

  10. Solubility • Entropy considerations • lattice entropy increases more than alkalis • two vs. three gaseous ions produced • hydration entropy is more negative due to charge density

  11. Solubility • Free energy considerations • for mononegative anions, more soluble than the alkalis

  12. Solubility • Mononegative vs. polynegative anions • polynegative anions have much higher lattice energies • polynegative anions have fewer total ions making the hydration enthalpy less • combinations of these two give a more positive free energy, thus lower solubility

  13. Beryllium • Steel gray • Hard • High melting point • Low density • High resistance to corrosion • Nonmagnetic

  14. Uses • gyroscopes • windows of X-ray tubes • transparency

  15. Sources • Bertrandite • Be4Si2O7(OH)2 • Beryl • Be3Al2Si6O18 • aquamarine • emerald • contamination of Cr(III)

  16. Compound Properties • Sweet taste • extremely poisonous • Inhalation results in berylliosis

  17. Chemistry of Beryllium • Covalent bonds predominate • high charge density polarizes any anion causing overlap to occur • Simple ionic compounds are a mixture • BeCl2·4H2O • [Be(OH2)4]2+·2Cl-

  18. Beryllium • Metallic, but can form oxyanions • amphoteric • “weak” metal H2O(l) + BeO(s) + 2H3O+(aq)  [Be(OH2)4]2+(aq) H2O(l) + BeO(s) + 2OH-(aq)  [Be(OH)4]2-(aq)

  19. Magnesium • Found in many minerals in nature • carnallite • KMgCl3·6H2O • dolomite • MgCO3·CaCO3 • 3rd most common ion in seawater • 108 million tons

  20. Isolation of Magnesium • Precipitation reaction Ca(OH)2(s) + Mg2+(aq)  Ca2+(aq) + Mg(OH)2(s) • Neutralization reaction Mg(OH)2(s) + 2HCl(aq)  MgCl2(aq) + 2H2O(l) • Placed in a Downs cell Mg2+(MgCl2) + 2e- Mg(l) 2Cl-(MgCl2) Cl2(g) + 2e-

  21. Magnesium Reactivity • E° = -2.37V • not as reactive in air because of the formation of a protective coating 2Mg(s) + O2(g)  2MgO(s) 2Mg(s) + CO2(g)  2MgO(s) + C(s) • special fire exstinguishing measures

  22. Magnesium Uses • 4 x 105 tons produced annually • ½ is used in aluminum-magnesium alloys • very low density (1.74 g/ml) • used in aircraft, railroads, ships, etc…

  23. Magnesium Chemistry • Decomposition upon heating MgCl2·H2O(s) + heat  Mg(OH)Cl(s) + HCl(g) • Formation of covalent bonds • Grignard reagents C2H5Br(ether) + Mg(s)  C2H5MgBr(ether)

  24. Calcium and Barium • grayish metals • react with oxygen 2Ca(s) + O2(g)  2CaO(s) 2Ba(s) + O2(g)  2BaO(s) Ba(s) + O2(g)  BaO2(g)

  25. Calcium and Barium • strong absorbers of x-rays • bones • dyes CaCl2·2H2O(s) + heat  CaCl2(s) + 2H2O(g) BaCl2·2H2O(s) + heat  BaCl2(s) + 2H2O(g)

  26. Oxides • Formed upon reaction with air or heating of the carbonate 2Mg(s) + O2(g) 2MgO(s) CaCO3(s) + heat  CaO(s) + CO2(g)

  27. Oxides • MgO • very high melting point • used as a refractory compound in furnaces • crystalline MgO is a good conductor of heat, but not electricity

  28. Oxides • CaO • called quicklime • undergoes thermoluminescence

  29. Oxides • CaO • reacts with water to form slaked lime CaO(s) + H2O(l)  Ca(OH)2(s) • used as a neutralizing agent in gardening, along with CaCO3 Ca(OH)2(s) + 2H+(aq)  Ca2+(aq) + 2H2O(l) CaCO3(s) + 2H+(aq)  Ca2+(aq) + H2O(l)

  30. Hydroxides • Solubility of hydroxides increase down the group

  31. Hydroxides • Limewater • a solution of calcium hydroxide Ca(OH)2(aq) + CO2(g)  CaCO3(s) + H2O(l) CaCO3(s) + H2O(l) + CO2(g)  Ca2+(aq) + 2HCO3-(aq) • leads to deterioration of marble

  32. Calcium Carbonate • Two naturally occurring crystalline forms • calcite • Iceland spar • two refractive indices • aragonite

  33. Calcium Carbonate • Chemistry involved in the formation of caves, stalagmites, and stalactites CaCO3(s) + CO2(aq) + H2O(l)  Ca2+(aq) + 2HCO3-(aq) Ca(HCO3)2(aq)  CaCO3(s) + CO2(g) + H2O(l) See animation at: http://www.classzone.com/books/earth_science/terc/content/visualizations/es1405/es1405page01.cfm?chapter_no=visualization

  34. Calcium Carbonate • Biological importance • dietary supplement • helps reduce osteoporosis • antacid • constipative

  35. Cement • Originally a paste of calcium hydroxide and sand • process perfected by the Romans • One of the largest chemical industries • 700 million tons annually

  36. Production of Cement • Grind limestone and shales together and heat to 1500°C • calcium carbonate and aluminosilicates • carbon dioxide is released and the melt is called “clinker” • The clinker is ground with a small amount of calcium sulfate • Portland cement

  37. Portland cement • 26% dicalcium silicate • Ca2SiO4 • 51% tricalcium silicate • Ca3SiO4 • 11% tricalcium aluminate • Ca3Al2O6 2Ca2SiO4(s) + 4H2O(l)  Ca3Si2O7·3H2O(s) + Ca(OH)2(s)

  38. Calcium Chloride • White, deliquescent solid • Formation of CaCl2·6H2O is very exothermic (H = -82 kJ/mol) • used in hot packs

  39. Calcium Chloride • Concentrated solutions • used to melt ice • lowers m.p. to -55°C • used to coat unpaved roads • minimizes dust • used to fill tires of earth-moving equipment • better traction

  40. Calcium Sulfate • Found naturally as a dihydrate • CaSO4·2H2O • gypsum or alabaster • Upon heating, forms the hemihydrate • CaSO4·1/2H2O • Plaster of Paris

  41. Calcium Sulfate • Uses • fire-resistant wallboard • endothermic process to form the hemihydrate • releases water

  42. Calcium Carbide • CaC2 • contains the acetylide ion or dicarbide(2-) ion, C22- • adopts the sodium chloride crystal structure

  43. Calcium Carbide • Preparation CaO(s) + 3C(s) + heat  CaC2(s) + CO(g) • 5 million tons produced annually • China is the main producer

  44. Calcium Carbide • Uses • production of acetylene CaC2(s) + 2H2O(l)  Ca(OH)2(s) + C2H2(g) 2C2H2(g) + 5O2(g)  4CO2(g) + 2H2O(g) • production of cyanamide ion CaC2(s) + N2(g) + heat  CaCN2(s) + C(s) CaCN2(s) + 3H2O(l)  CaCO3(s) + 2NH3(aq)

  45. Biological Aspects • Photosynthesis • magnesium is contained in chlorophyll and keeps the molecule in a specific configuration 6CO2(g) + 6H2O(l)  C6H12O6(aq) + 6O2(g)

  46. Biological Aspects • Magnesium is concentrated inside cells • triggers the relaxation of muscles • Calcium is concentrated outside cells • important in blot-clotting • trigger the contractions of muscles

  47. Magnesium Reaction Flowchart

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