Class 10 Science Notes Chapter 3 (Metals and non metals) – Science

Okay, let's focus on the key concepts from NCERT Class 10 Science, Chapter 3: Metals and Non-metals, tailored for government exam preparation.

Chapter 3: Metals and Non-metals - Detailed Notes

1. Classification of Elements:
Elements are broadly classified into Metals and Non-metals based on their properties. Metalloids (like Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium) exhibit properties intermediate between metals and non-metals.

2. Physical Properties:

3. Chemical Properties of Metals:

• Reaction with Air (Oxygen):

  • Metal + Oxygen → Metal Oxide
  • Example: 2Cu + O₂ → 2CuO (Black copper(II) oxide)
  • Example: 4Al + 3O₂ → 2Al₂O₃ (Aluminium oxide)
  • Reactivity Variation: Na and K react vigorously, catching fire in open air (kept under kerosene). Mg, Al, Zn, Pb form a protective oxide layer preventing further oxidation. Fe doesn't burn but iron filings burn vigorously. Cu doesn't burn but forms a black oxide layer. Ag and Au don't react with oxygen.
  • Amphoteric Oxides: Oxides that react with both acids and bases to produce salt and water.
    • Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O (Acting as a base)
    • Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O (Sodium aluminate) (Acting as an acid)
    • ZnO + 2HCl → ZnCl₂ + H₂O (Acting as a base)
    • ZnO + 2NaOH → Na₂ZnO₂ + H₂O (Sodium zincate) (Acting as an acid)

• Reaction with Water:

  • Metal + Water → Metal Oxide / Metal Hydroxide + Hydrogen Gas
  • Reactivity Variation:
    • Na, K, Ca react with cold water:
      • 2Na + 2H₂O → 2NaOH + H₂ + Heat (Violent, H₂ catches fire)
      • Ca + 2H₂O → Ca(OH)₂ + H₂ (Less violent, Ca floats as H₂ bubbles stick)
    • Mg reacts with hot water:
      • Mg + 2H₂O(hot) → Mg(OH)₂ + H₂ (Mg also floats)
    • Al, Zn, Fe react with steam:
      • 2Al + 3H₂O(g) → Al₂O₃ + 3H₂
      • 3Fe + 4H₂O(g) → Fe₃O₄ (Iron(II,III) oxide) + 4H₂
    • Pb, Cu, Ag, Au do not react with water/steam.

• Reaction with Dilute Acids (HCl, H₂SO₄):

  • Metal + Dilute Acid → Salt + Hydrogen Gas
  • Example: Fe + 2HCl → FeCl₂ + H₂
  • Example: Zn + H₂SO₄ → ZnSO₄ + H₂
  • Reactivity Variation: Reactivity decreases down the series (Mg > Al > Zn > Fe). Cu, Ag, Au do not react with dilute HCl or H₂SO₄ as they are less reactive than Hydrogen.
  • Reaction with Nitric Acid (HNO₃): HNO₃ is a strong oxidizing agent. It oxidizes the H₂ produced to water and itself gets reduced to nitrogen oxides (N₂O, NO, NO₂).
    • Cu + 4HNO₃(conc) → Cu(NO₃)₂ + 2NO₂ + 2H₂O
    • Exception: Very dilute HNO₃ reacts with Magnesium (Mg) and Manganese (Mn) to evolve H₂ gas.
      • Mg + 2HNO₃(very dilute) → Mg(NO₃)₂ + H₂

• Reaction with Solutions of Other Metal Salts (Displacement Reaction):

  • A more reactive metal displaces a less reactive metal from its salt solution.
  • Metal A + Salt Solution of Metal B → Salt Solution of Metal A + Metal B (If A is more reactive than B)
  • Example: Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s) (Blue solution turns green, reddish-brown Cu deposits)
  • Example: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
  • Example: Cu(s) + AgNO₃(aq) → Cu(NO₃)₂(aq) + 2Ag(s)
  • This forms the basis of the Reactivity Series.

4. The Reactivity Series (Activity Series):
A list of metals arranged in order of their decreasing chemical reactivity.

• K (Potassium) - Most Reactive
• Na (Sodium)
• Ca (Calcium)
• Mg (Magnesium)
• Al (Aluminium)
• Zn (Zinc)
• Fe (Iron)
• Pb (Lead)
• (H) (Hydrogen) - Reference point
• Cu (Copper)
• Hg (Mercury)
• Ag (Silver)
• Au (Gold) - Least Reactive

Uses: Predicts displacement reactions, reaction with water/acids, method of extraction.

5. How do Metals and Non-metals React? (Formation of Ionic Compounds)

• Metals tend to lose valence electrons to form positive ions (cations).

  • Na → Na⁺ + e⁻ (Configuration: 2,8,1 → 2,8)

• Non-metals tend to gain electrons in their valence shell to form negative ions (anions).

  • Cl + e⁻ → Cl⁻ (Configuration: 2,8,7 → 2,8,8)

• The cation and anion are held together by strong electrostatic forces of attraction, forming an Ionic Bond or Electrovalent Bond.

• The resulting compounds are called Ionic Compounds or Electrovalent Compounds.

  • Example: Formation of NaCl (Sodium Chloride)
    • Na → Na⁺ + e⁻
    • Cl + e⁻ → Cl⁻
    • Na⁺ + Cl⁻ → NaCl
  • Example: Formation of MgCl₂ (Magnesium Chloride)
    • Mg → Mg²⁺ + 2e⁻
    • Cl + e⁻ → Cl⁻ (Needs 2 Cl atoms)
    • Mg²⁺ + 2Cl⁻ → MgCl₂

• Properties of Ionic Compounds:

  • Physical Nature: Solid, hard, generally brittle (due to strong inter-ionic attraction).
  • Melting & Boiling Points: High MP & BP (significant energy needed to break strong ionic bonds).
  • Solubility: Generally soluble in water, insoluble in organic solvents like kerosene, petrol.
  • Electrical Conductivity: Do not conduct electricity in the solid state (ions not mobile). Conduct electricity in the molten state or in aqueous solution (ions become free to move).

6. Occurrence and Extraction of Metals (Metallurgy):

• Minerals: Elements or compounds occurring naturally in the Earth's crust.
• Ores: Minerals from which metals can be extracted profitably and conveniently. (All ores are minerals, but not all minerals are ores).
• Gangue: Earthly impurities (sand, soil, etc.) associated with the ore.
• Metallurgy: The entire scientific and technological process used for isolation of the metal from its ore. Steps involved:

  1. Concentration of Ore (Enrichment/Dressing): Removal of gangue. Methods depend on physical/chemical property differences between ore and gangue (e.g., hydraulic washing, magnetic separation, froth flotation, chemical leaching).

  2. Extraction of Metal from Concentrated Ore: Depends on the metal's position in the reactivity series.

     • Metals Low in Reactivity (Ag, Au, Cu, Hg): Found often in free state (Au, Ag) or as oxides/sulphides. Oxides reduced by heating alone. Sulphides converted to oxides by Roasting, then reduced by heating.
       • 2HgS (Cinnabar) + 3O₂ (Heat - Roasting) → 2HgO + 2SO₂
       • 2HgO (Heat) → 2Hg + O₂
       • 2Cu₂S + 3O₂ (Heat - Roasting) → 2Cu₂O + 2SO₂
       • 2Cu₂O + Cu₂S (Heat) → 6Cu + SO₂ (Auto-reduction)
     • Metals of Medium Reactivity (Zn, Fe, Pb, etc.): Usually present as sulphides or carbonates.
       • Convert to Oxides:
         • Roasting: Heating sulphide ores strongly in the presence of excess air.
           • 2ZnS + 3O₂ (Heat) → 2ZnO + 2SO₂
         • Calcination: Heating carbonate ores strongly in limited or no air.
           • ZnCO₃ (Heat) → ZnO + CO₂
       • Reduction of Oxide to Metal:
         • Using Carbon (Smelting): ZnO + C → Zn + CO
         • Using Reactive Metals (Displacement/Thermite Reaction): More reactive metals like Na, Ca, Al used if carbon reduction is not feasible or economical.
           • Thermite Reaction: Fe₂O₃(s) + 2Al(s) → 2Fe(l) + Al₂O₃(s) + Heat (Highly exothermic, used to weld railway tracks).
     • Metals High in Reactivity (K, Na, Ca, Mg, Al): Cannot be reduced by carbon (metals have higher affinity for oxygen than carbon). Extracted by Electrolysis of their molten chlorides or oxides.
       • Metal deposited at the Cathode (negative electrode).
       • Non-metal (e.g., Chlorine) liberated at the Anode (positive electrode).
       • Example (Molten NaCl):
         • At Cathode: Na⁺ + e⁻ → Na
         • At Anode: 2Cl⁻ → Cl₂ + 2e⁻
       • Example (Alumina - Al₂O₃ in molten Cryolite Na₃AlF₆ + Fluorspar CaF₂ to lower MP and increase conductivity):
         • At Cathode: Al³⁺ + 3e⁻ → Al
         • At Anode: 2O²⁻ → O₂ + 4e⁻ (Oxygen reacts with carbon anode: C + O₂ → CO₂)

  3. Refining of Metals (Purification): Removal of impurities from the extracted metal.

     • Electrolytic Refining: Most common method. Used for Cu, Zn, Ni, Ag, Au, etc.
       • Anode: Impure metal slab.
       • Cathode: Thin strip of pure metal.
       • Electrolyte: Solution of a salt of the metal being refined (e.g., CuSO₄ for Copper refining).
       • Process: On passing current, pure metal from the anode dissolves into the electrolyte. An equivalent amount of pure metal from the electrolyte deposits onto the cathode. Impurities settle below the anode as Anode Mud.
       • At Anode: Cu (impure) → Cu²⁺(aq) + 2e⁻
       • At Cathode: Cu²⁺(aq) + 2e⁻ → Cu (pure)

7. Corrosion:

• The gradual deterioration of metals by the action of air, moisture, or chemicals (like acids) on their surface.
• Rusting of Iron: Major example. Hydrated Iron(III) oxide (Fe₂O₃.xH₂O) is formed. Requires both Oxygen and Water (or moisture). Rust is flaky and non-protective.
• Corrosion of Copper: Forms a green coating of basic copper carbonate (CuCO₃.Cu(OH)₂).
• Corrosion of Silver: Forms a black coating of silver sulphide (Ag₂S) by reacting with Sulphur compounds in the air.
• Prevention of Corrosion:
  • Barrier Protection: Painting, oiling, greasing, plastic coating.
  • Galvanizing: Coating iron or steel with a thin layer of Zinc. Zinc is more reactive than iron, so it corrodes preferentially (sacrificial protection).
  • Chromium Plating / Anodizing (for Aluminium): Depositing a layer of another metal (like Cr, Ni, Sn) or forming a thick, protective oxide layer (Anodizing).
  • Alloying: Mixing the metal with other metals or non-metals to improve properties, including corrosion resistance.

8. Alloys:

• A homogeneous mixture of two or more metals, or a metal and a non-metal.
• Prepared by melting the primary metal and then dissolving other elements in definite proportions.
• Purpose: To modify properties like hardness, strength, corrosion resistance, melting point, electrical conductivity.
• Examples:
  • Steel: Iron (Fe) + Carbon (C) (small amount) - Harder, stronger than pure iron.
  • Stainless Steel: Iron (Fe) + Nickel (Ni) + Chromium (Cr) - Corrosion resistant, hard.
  • Brass: Copper (Cu) + Zinc (Zn) - Used in utensils, screws.
  • Bronze: Copper (Cu) + Tin (Sn) - Used for statues, medals, bells.
  • Solder: Lead (Pb) + Tin (Sn) - Low melting point, used for welding electrical wires.
  • Amalgam: An alloy where one of the metals is Mercury (Hg). Used in dental fillings (with Ag, Sn, Zn).
• Note: The electrical conductivity and melting point of an alloy are generally lower than those of the constituent pure metals. Pure (24 carat) gold is very soft; hence, it is alloyed with Silver or Copper (e.g., 22 carat gold) to make it harder for jewellery.

This covers the essential points from Chapter 3 for exam preparation. Focus on definitions, key reactions (balanced equations), the reactivity series, extraction processes (roasting, calcination, electrolysis, smelting), properties of ionic compounds, corrosion prevention, and examples of alloys with their constituents and uses. Remember to understand the why behind the properties and reactions based on electronic configurations and reactivity.