10.2 Alkali and Alkaline earth Metals

Chemistry – Class 11

10.2.1 Alkali Metals : General characteristics of alkali metals ; Sodium [extraction from Down’s process, properties (action with Oxygen, water, acids nonmetals and ammonia) and uses] ; Properties (precipitation reaction and action with carbon monooxide) and uses of sodium hydroxide ; Properties (action with CO2, SO2, water, precipitation reactions) and uses of sodium carbonate 10.2.2 Alkaline Earth Metals : General characteristics of alkaline earth metals ; Molecular formula and uses of (quick lime, bleaching powder, magnesia, plaster of paris and epsom salt); Solubility of hydroxides, carbonates and sulphates of alkaline earth metals (general trend with explanation) ; Stability of carbonate and nitrate of alkaline earth metals (general trend with explanation).

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Class 11 Chemistry Nepal: Alkali and Alkaline Earth Metals Notes

Unit 10.2: Alkali and Alkaline Earth Metals (5 Teaching Hours)

  1. Alkali Metals

    • General Characteristics of Alkali Metals
      • Group 1 elements (Li, Na, K, Rb, Cs, Fr) with one valence electron (ns¹).
      • Soft, low melting points, highly reactive due to low ionization energy; reactivity increases down the group.
      • Form basic oxides (e.g., Na₂O) and hydroxides (e.g., NaOH).
      • Stored in oil to prevent reaction with air/moisture.
      • Figure 1: Periodic Trends in Alkali Metals (Diagram showing trends in ionization energy and reactivity).
    • Sodium
      • Extraction from Down’s Process
        • Electrolysis of molten NaCl with CaCl₂ (lowers melting point) in Down’s cell; Na forms at cathode, Cl₂ at anode.
        • Figure 2: Down’s Process (Diagram showing Down’s cell setup).
      • Properties
        • Action with Oxygen: Forms sodium peroxide: 2Na + O₂ → Na₂O₂.
        • Action with Water: Forms NaOH and H₂: 2Na + 2H₂O → 2NaOH + H₂.
        • Action with Acids: Reacts to form salts and H₂: 2Na + 2HCl → 2NaCl + H₂.
        • Action with Non-Metals: Forms compounds like NaCl, Na₂S (e.g., 2Na + Cl₂ → 2NaCl).
        • Action with Ammonia: Forms sodamide: 2Na + 2NH₃ → 2NaNH₂ + H₂.
      • Uses
        • Coolant in nuclear reactors.
        • Synthesis of chemicals (e.g., NaOH, Na₂CO₃).
        • Sodium vapor lamps.
    • Properties and Uses of Sodium Hydroxide (NaOH)
      • Precipitation Reaction: Forms insoluble hydroxides with metal ions (e.g., FeCl₃ + 3NaOH → Fe(OH)₃↓ + 3NaCl).
      • Action with Carbon Monoxide: Forms sodium formate at high pressure: NaOH + CO → HCOONa.
      • Uses: Soap and detergent production, paper industry, petroleum refining, pH regulation.
      • Figure 3: Precipitation Reaction of NaOH (Diagram showing Fe(OH)₃ precipitation).
    • Properties and Uses of Sodium Carbonate (Na₂CO₃)
      • Action with CO₂: Forms sodium bicarbonate: Na₂CO₃ + CO₂ + H₂O → 2NaHCO₃.
      • Action with SO₂: Forms sodium sulphite: Na₂CO₃ + SO₂ → Na₂SO₃ + CO₂.
      • Action with Water: Hydrolyses to form alkaline solution: Na₂CO₃ + H₂O ⇌ NaHCO₃ + NaOH.
      • Precipitation Reactions: Forms insoluble carbonates (e.g., CaCl₂ + Na₂CO₃ → CaCO₃↓ + 2NaCl).
      • Uses: Glass manufacturing, water softening, detergent production, pH buffer.

  2. Alkaline Earth Metals

    • General Characteristics of Alkaline Earth Metals
      • Group 2 elements (Be, Mg, Ca, Sr, Ba, Ra) with two valence electrons (ns²).
      • Less reactive than alkali metals; reactivity increases down the group.
      • Form basic oxides (e.g., CaO) and hydroxides (e.g., Ca(OH)₂).
      • Figure 4: Periodic Trends in Alkaline Earth Metals (Diagram showing trends in ionization energy and reactivity).
    • Molecular Formula and Uses of Compounds
      • Quick Lime (CaO): Used in cement, mortar, and neutralizing acidic soils.
      • Bleaching Powder (Ca(OCl)Cl): Used as disinfectant, bleaching agent in textiles/paper.
      • Magnesia (MgO): Used as refractory material, antacid in medicine.
      • Plaster of Paris (CaSO₄·½H₂O): Used in construction, medical casts, sculptures.
      • Epsom Salt (MgSO₄·7H₂O): Used as laxative, fertilizer, bath salts.
    • Solubility of Hydroxides, Carbonates, and Sulphates
      • Hydroxides: Solubility increases down the group (Be(OH)₂ insoluble, Ba(OH)₂ more soluble) due to decreasing lattice energy.
      • Carbonates: Solubility decreases down the group (MgCO₃ soluble, BaCO₃ insoluble) due to increasing lattice stability.
      • Sulphates: Solubility decreases down the group (MgSO₄ soluble, BaSO₄ insoluble) due to increasing lattice energy.
      • Figure 5: Solubility Trends of Alkaline Earth Compounds (Diagram showing solubility trends).
    • Stability of Carbonates and Nitrates
      • Carbonates: Thermal stability increases down the group (BeCO₃ least stable, BaCO₃ most stable) due to increasing ionic radius stabilizing lattice.
      • Nitrates: Thermal stability increases down the group (Be(NO₃)₂ decomposes easily, Ba(NO₃)₂ more stable) due to larger cations stabilizing anion.
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