The Alkaline Earth Metals

CHEM 251 LECTURE 3 & 4: The Alkaline Earth Metals

GROUP II A ELEMENTS

  • The term "alkaline earth metals" originates from early chemists who referred to any water-insoluble substance that remained unchanged by fire as "earth." The names for the respective "earths" of this group include lime (CaO), magnesia (MgO), strontia (SrO), and baryta (BaO), as these compounds displayed alkaline properties when dissolved in water.
  • Valence Shell Configuration: All alkaline earth metals possess a valence shell electron configuration of ns^2, indicating they have two electrons outside the noble gas core.
    • Elements and Electronic Configuration:
      • Beryllium (Be): [He] 2s²
      • Magnesium (Mg): [Ne] 3s²
      • Calcium (Ca): [Ar] 4s²
      • Strontium (Sr): [Kr] 5s²
      • Barium (Ba): [Xe] 6s²
      • Radium (Ra): [Rn] 7s²

GROUP II A ELEMENTS - SIMILARITIES AND TRENDS

  • While there are notable similarities in both the chemical and physical properties of alkaline earth metals, these similarities are less pronounced than among alkali metals (Group I A).
  • Example: Lithium has properties more akin to cesium than beryllium relates to barium.
  • A general phenomenon across the periodic table indicates that as one moves towards the center from either side, there is an increasing disparity between the properties of upper and lower elements in a group.

GROUP II A ELEMENTS - OCCURRENCE

  • Reactivity: Group IIA elements are reactive and do not exist in a free state in nature.
  • Abundance in Nature:
    • Magnesium (Mg) and Calcium (Ca) are abundant.
    • Beryllium (Be) is not abundant.
    • Strontium (Sr) and Barium (Ba) are relatively less abundant.
  • Common Occurrences: These metals primarily occur in nature as carbonates, sulfates, and silicates.

PHYSICAL PROPERTIES OF GROUP II A ELEMENTS

  • | Element | Melting Point (°C) | Boiling Point (°C) | Ionization Enthalpies (kJ/mol) | Reduction Potential (V) | Ionic Radius (Å) | Radius of M+ Ion (Å) |
  • |---|---|---|---|---|---|---|
  • | Be | 1278 | 2500 | 1st: 899, 2nd: 1757 | -1.85 | 0.31 | 0.30 |
  • | Mg | 651 | 1105 | 1st: 737, 2nd: 1450 | -2.37 | 0.78 | 0.65 |
  • | Ca | 843 | 1494 | 1st: 590, 2nd: 1146 | -2.76 | 1.06 | 0.99 |
  • | Sr | 769 | 1381 | 1st: 549, 2nd: 1064 | -2.89 | 1.27 | 1.13 |
  • | Ba | 725 | 1850 | 1st: 503, 2nd: 965 | -2.90 | 1.43 | 1.35 |
  • | Ra | 700 | 1700 | 1st: 509, 2nd: 979 | -2.92 | 1.37 | 1.40 |

Additional Physical Properties

  • Size: Atomic radii are smaller than those of Group I A elements due to an increased nuclear charge, leading to tighter electron binding. Each atom loses a loosely held outer electron, resulting in doubly charged cations that are bound more tightly.
  • Metallic Bonding: Alkaline earth metals involve two electrons in metallic bonding, making them harder, heavier, and with higher melting points compared to alkali metals.
  • Density: Alkaline earth metals are generally lighter than alkali metals but denser due to the greater nuclear charge. Density trends decrease from Be to Mg and Ca, then increase from Ca to Ra.
  • Hardness: Hardness results from stronger bonding due to the presence of two valence electrons; decreases down the group.
  • Melting and Boiling Points: High melting points correlate with hardness and ionic radius.
  • Ionization Energies: Ionization energies decrease down the group, with second ionization energy being roughly double the first but significantly higher than Group I elements' energies.
  • Flame Colors: Calcium, Strontium, and Barium produce characteristic flame colors (brick red, crimson, and apple green, respectively). Beryllium and Magnesium do not contribute color due to stronger electron binding.

Electronegativities and Reducing Properties

  • The electronegativity trends reveal low values (except for Be), decreasing down the group. Alkaline earth metals act as reducing agents, with reducing potential increasing negatively as one descends from Be to Ra. Hydration energy notably decreases down the group.

Hydration Energies

  • Hydration energies for group II ions are significantly higher than those for Group I, mostly due to smaller ionic size and increased charges. Hydration reactions are described as:
    M^{2+} + x H2O ightarrow [M(H2O)_x]^{2+} + ext{Energy}
  • Values for ionic radius and hydration energy:
    • Be²⁺: 0.31 Å, -2494 kJ/mol
    • Mg²⁺: 0.78 Å, -1921 kJ/mol
    • Ca²⁺: 1.06 Å, -1597 kJ/mol
    • Sr²⁺: 1.27 Å, -1443 kJ/mol
    • Ba²⁺: 1.43 Å, -1305 kJ/mol

LATTICE ENERGIES OF SOME COMPOUNDS OF GROUP II A ELEMENTS

  • The lattice energies for Group II A compounds are higher compared to Group I A due to increased ionic charge. Lattice energy decreases as the metal size increases:

U = -No A rac{Z+ Z- e^2 (1 - rac{1}{n})}{4 imes ext{π} ext{ε}o r_0}

Where:

  • $N_o$ = Avogadro's number
  • $A$ = Madelung constant
  • $Z+$ and $Z-$ = charges on cation and anion, respectively
  • $ ext{ε}_o$ = permittivity of free space
  • $r_0$ = internuclear distance

Trends in Solubility With Lattice and Hydration Energy

  • Lattice energies and hydration energies change in opposite directions as varying sizes of metal ions affect solubility:
    • Most compounds become less soluble as metals increase in size, while hydroxides and fluorides show increased solubility with decreasing lattice energy.

CHEMICAL PROPERTIES OF ALKALINE EARTH METALS

  • Alkaline earth metals act as good polarizers, but their oxides are more covalent compared to alkali metal oxides, and their hydroxides are generally less basic.
  • Common Reactions:
    • Reaction with Water:
      • M + 2H2O ightarrow M(OH)2 + H_2 (e.g., Ca)
      • Mg + 2H2O ightarrow Mg(OH)2 + H2 or Mg + H2O
        ightarrow MgO + H_2
    • Reaction with Acids:
      • M + 2HCl
        ightarrow MCl2 + H2
    • Amphoteric behavior of Beryllium noted in
      • Be + NaOH
        ightarrow Na2[Be(OH)4] + H_2

Summary of Chemical Reactions:

  • Formation of solid sulfides (M + S = MS), selenides (M + Se = MSe), nitrides (3M + N2 = M3N2), phosphides (3M + 2P = M3P2), and multiple halides (reacting with halogens).
  • Beryllium does not react with boiling water while magnesium reacts with hot water and produces varying results with cold water.

SOLUTION OF GROUP IIA METALS IN LIQUID AMMONIA

  • Group IIA metals dissolve in liquid ammonia, creating blue-colored solutions due to solvated electrons, leading over time to form amides and hydrogen evolution:
    2NH3 + 2e^- ightarrow 2NH2^- + H2 M(NH3)6 ightarrow M(NH2)2 + 4NH3 + H_2
  • Concentrated ammonia solutions turn bronze as metallic clusters form, demonstrating particular properties distinct to each alkaline earth metal.

OXIDES OF GROUP II ELEMENTS

  • All IIA elements react with O2 to form normal oxides (MO), increasing in reactivity down the group. Oxides such as BeO are high in melting point (2370 °C) and exhibit invariant properties.
  • The formation pathways and reactivity in air vary:
    • Example: 2M + O_2
      ightarrow 2MO [general oxide formation]
  • Experimental setups for these reactions generally involve heating metals in controlled environments to facilitate oxidation.

SOLUBILITY AND INERTNESS OF OXIDES

  • The oxides display specific characteristics useful in industrial applications, particularly in higher temperature applications like in furnace linings. They are conducive to heat conductivity, are chemically inert, and are utilized for insulating applications due to high melting points and low vapor pressure.

SULPHATES

  • Sulfates demonstrate a decreasing solubility trend down the group from Be > Mg >> Ca > Sr > Ba, suggesting varying hydrated states (Gypsum as CaSO4.2H2O).
  • Plaster of Paris (POP) involves reacting calcium sulfate derivatives and can achieve a quick set with water while forming stable structures suitable for various applications.
  • Order of Stability of the Sulphates:
    • BeSO4 < MgSO4 < CaSO4 < SrSO4 < BaSO_4
  • Hydration reactions: (e.g., CaSO4 undergoes transformations under heating leading to anhydrite).

WATER SOFTENING

  • Hard Water Issues: Presence of dissolved Ca²⁺ and Mg²⁺ leads to precipitates while reducing the function of soaps and detergents. This leads to issues such as deposition within plumbing systems.
  • Softening Agents: Calcium hydroxide is typically utilized to precipitate these ions, reverting carbonate precipitation in solutions.

Example Reaction:

Ca(HCO3)2 (aq) + Ca(OH)2 (aq) ightarrow 2CaCO3 (s) + 2H_2O (l)