What are alkaline earth metals?

Short Answer

Alkaline earth metals are the elements of Group 2 in the periodic table, including beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra).

  • They have two electrons in their outermost shell (ns²), making them less reactive than alkali metals but still highly reactive.
  • These metals are shiny, silvery, and good conductors, and they readily form compounds with non-metals, especially oxygen and halogens.
Detailed Explanation :

General Characteristics of Alkaline Earth Metals

Alkaline earth metals are located in Group 2 of the periodic table. Their general electronic configuration is ns², meaning there are two valence electrons in the outermost s-orbital. This configuration gives them distinct physical and chemical properties.

Physical Properties

  1. Shiny Metallic Luster:
    • Alkaline earth metals are silvery-white and reflect light when freshly cut.
  2. Hardness:
    • Harder than alkali metals because they have stronger metallic bonds due to two valence electrons.
  3. Density:
    • Density increases down the group, with beryllium and magnesium being lighter than calcium, strontium, and barium.
  4. Melting and Boiling Points:
    • Higher than alkali metals because of stronger metallic bonds.
    • Example: Beryllium → 1287°C, Magnesium → 650°C, Barium → 727°C
  5. Conductivity:
    • Good conductors of heat and electricity due to delocalized valence electrons.

Chemical Properties

  1. Reactivity:
    • Alkaline earth metals are reactive but less than alkali metals.
    • Reactivity increases down the group as atomic size increases and valence electrons are farther from the nucleus.
  2. Reaction with Water:
    • React with hot water (except beryllium and magnesium which react slowly) to form hydroxides and hydrogen gas.
    • Example: Ca + 2H₂O → Ca(OH)₂ + H₂
  3. Reaction with Oxygen:
    • Form oxides that are basic (alkaline).
    • Example: Mg + O₂ → 2MgO
  4. Reaction with Halogens:
    • Form ionic halides.
    • Example: Ca + Cl₂ → CaCl₂
  5. Reaction with Acids:
    • React vigorously to produce salts and hydrogen gas.
    • Example: Mg + 2HCl → MgCl₂ + H₂

Trends in Alkaline Earth Metals

  1. Atomic and Ionic Size:
    • Increases down the group due to addition of electron shells.
  2. Ionization Energy:
    • Decreases down the group → valence electrons are easier to remove as distance from nucleus increases.
  3. Electronegativity:
    • Decreases down the group → lower tendency to attract electrons.
  4. Reactivity:
    • Increases down the group → barium is more reactive than magnesium.

Occurrence and Uses

  1. Occurrence:
    • Highly reactive → not found free in nature.
    • Occur in compounds like CaCO₃ (limestone), MgSO₄ (Epsom salt), and BaSO₄.
  2. Uses:
    • Calcium: Cement, bones, and teeth.
    • Magnesium: Alloy production, fireworks.
    • Barium: X-ray contrast agents.
    • Strontium: Fireworks and flares.

Comparison with Alkali Metals

Feature Alkali Metals Alkaline Earth Metals
Group 1 2
Valence Electrons 1 2
Reactivity Very high High but less than Group 1
Density Lower Higher
Hardness Soft Harder
Conclusion

Alkaline earth metals are Group 2 elements with two valence electrons, giving them metallic, reactive, and basic properties. Reactivity increases down the group due to larger atomic size and lower ionization energy. Their physical and chemical properties make them important in industry, medicine, construction, and fireworks, and they are essential in understanding periodic trends in metals.