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
4. Group Trends
• greater enthalpies of atomization than the
alkali metals
– stronger metallic bonding
• also harder and have higher melting points
Element Melting Point (°C) Hatm (kJ/mol)
Mg 649 149
Ca 839 177
Sr 768 164
Ba 727 175
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
Element MCl2 M(NO3)2 MSO4
Mg 6 6 7
Ca 6 4 2
Sr 6 4 0
Ba 2 0 0
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
Compound Lattice Energy
(kJ/mol)
Hydration Enthalpy
(kJ/mol)
Net Enthalpy Change
(kJ/mol)
MgCl2 +2526 -2659 -133
NaCl +788 -784 +4
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
Compound Lattice Entropy
(kJ/mol)
Hydration Entropy
(kJ/mol)
Net Entropy Change
(kJ/mol)
MgCl2 +109 -143 -34
NaCl +68 -55 +13
11. Solubility
• Free energy considerations
– for mononegative anions, more soluble than the
alkalis
Compound Enthalpy Change
(kJ/mol)
Entropy Change
(kJ/mol)
Free Energy Change
(kJ/mol)
MgCl2 -133 -34 -99
NaCl +4 +13 -9
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
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-
Be
OH2
H2O OH2
OH2
2+
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
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…
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
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
Hydroxide Solubility (g/L)
Mg 0.0001
Ca 1.2
Sr 10
Ba 47
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/cont
ent/visualizations/es1405/es1405page01.cfm?chapter_
no=visualization
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
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,
C2
2-
– adopts the sodium chloride crystal structure
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
