Metals and Non-metals

Introduction

• Elements are classified as metals or non-metals based on their properties.
• Key considerations:
  • Uses of metals and non-metals in daily life.
  • Properties used for classification.
  • Relationship between properties and uses.

Physical Properties

Metals

• Metallic Lustre: Metals have a shiny surface in their pure state (Activity 3.1).
• Hardness: Most metals are hard, though hardness varies (Activity 3.2).
• Malleability: Metals can be beaten into thin sheets (e.g., gold and silver are highly malleable) (Activity 3.3).
• Ductility: Metals can be drawn into thin wires (gold is the most ductile; 1g can form a 2 km wire) (Activity 3.4).
• Conductivity:
  • Good conductors of heat (silver and copper are best; lead and mercury are poor) (Activity 3.5).
  • Good conductors of electricity (Activity 3.6).
• Sonority: Metals produce sound when struck, making them suitable for bells.
• Exceptions:
  • Mercury is liquid at room temperature.
  • Gallium and caesium have low melting points.
  • Alkali metals (lithium, sodium, potassium) are soft, low density, and low melting point.

Non-metals

• Exist as solids or gases (except bromine, a liquid).
• Examples: carbon, sulphur, iodine, oxygen, hydrogen (Activity 3.7).
• Properties (opposite to metals):
  • Non-lustrous (except iodine).
  • Not malleable or ductile.
  • Poor conductors of heat and electricity (except graphite).
  • Not sonorous.
• Exceptions:
  • Diamond (carbon allotrope) is the hardest natural substance with high melting/boiling points.
  • Graphite conducts electricity.

Chemical Properties

Metals

• Reaction with Oxygen (Activity 3.9):
  • Form metal oxides (basic or amphoteric).
  • Examples:
    • 2Cu + O₂ → 2CuO (Copper(II) oxide)
    • 4Al + 3O₂ → 2Al₂O₃ (Aluminium oxide)
  • Amphoteric oxides (e.g., Al₂O₃, ZnO) react with both acids and bases.
  • Reactivity order: Sodium > Magnesium > Zinc, Iron, Copper, Lead.
  • Some oxides (Na₂O, K₂O) dissolve in water to form alkalis.
• Reaction with Water (Activity 3.10):
  • Form metal oxides/hydroxides and hydrogen.
  • Examples:
    • 2K + 2H₂O → 2KOH + H₂ (violent, exothermic)
    • Ca + 2H₂O → Ca(OH)₂ + H₂ (less violent)
    • 2Al + 3H₂O(g) → Al₂O₃ + 3H₂ (with steam)
  • Reactivity: Potassium, Sodium > Calcium > Magnesium > Aluminium, Iron, Zinc.
  • Lead, copper, silver, gold do not react with water.
• Reaction with Acids (Activity 3.11):
  • Form salt and hydrogen gas.
  • Examples:
    • Mg + 2HCl → MgCl₂ + H₂
    • Zn + H₂SO₄ → ZnSO₄ + H₂
  • Reactivity: Mg > Al > Zn > Fe; copper does not react with dilute HCl.
  • Nitric acid does not produce H₂ (except with Mg, Mn in very dilute form).
  • Aqua regia (3:1 HCl:HNO₃) dissolves gold and platinum.
• Reaction with Metal Salts (Activity 3.12):
  • More reactive metals displace less reactive ones from their salts.
  • Example: Fe + CuSO₄ → FeSO₄ + Cu (iron displaces copper).
• Reactivity Series:

  Metal	Symbol	Reactivity
  Potassium	K	Most reactive
  Sodium	Na
  Calcium	Ca
  Magnesium	Mg
  Aluminium	Al
  Zinc	Zn	Reactivity decreases
  Iron	Fe
  Lead	Pb
  [Hydrogen]	[H]
  Copper	Cu
  Mercury	Hg
  Silver	Ag
  Gold	Au	Least reactive

Non-metals

• Reaction with Oxygen (Activity 3.8):
  • Form acidic or neutral oxides.
  • Example: S + O₂ → SO₂ (acidic).
• Do not displace hydrogen from acids.
• Form hydrides with hydrogen.

Reactions of Metals and Non-metals

• Metals lose electrons to form positive ions (cations).
• Non-metals gain electrons to form negative ions (anions).
• Example: Na + Cl → NaCl (ionic compound).
• Ionic compounds (e.g., MgCl₂):
  • Solid, hard, brittle.
  • High melting/boiling points.
  • Soluble in water, insoluble in petrol/kerosene.
  • Conduct electricity in molten state or aqueous solution (Activity 3.13).

Occurrence of Metals

• Found in earth’s crust as minerals or ores.
• Some metals (e.g., gold, silver) occur in free state; others as compounds.

Extraction of Metals

Low Reactivity Metals

• Extracted by heating oxides alone.
• Example: HgS → HgO → Hg.

Medium Reactivity Metals

• Converted to oxides via roasting (sulphides) or calcination (carbonates).
• Reduced using carbon or displacement reactions.
• Example: ZnO + C → Zn + CO.

High Reactivity Metals

• Extracted by electrolytic reduction.
• Example: NaCl electrolysis → Na (cathode), Cl₂ (anode).

Refining

• Electrolytic refining: Impure metal (anode), pure metal (cathode).
• Example: Copper refining using CuSO₄ solution.

Corrosion

• Metals corrode in moist air:
  • Silver → silver sulphide (black).
  • Copper → basic copper carbonate (green).
  • Iron → rust (Fe₂O₃·nH₂O).
• Conditions for rusting: Presence of air and water (Activity 3.14).
• Prevention:
  • Painting, oiling, greasing.
  • Galvanising (zinc coating).
  • Alloying (e.g., stainless steel with nickel and chromium).
  • Anodising (thick aluminium oxide layer).

Key Questions

1. Sodium is stored in kerosene to prevent reaction with oxygen (Q1).
2. Reactions (Q2):
   • Fe + H₂O(g) → Fe₂O₃ + H₂
   • Ca + 2H₂O → Ca(OH)₂ + H₂
   • 2K + 2H₂O → 2KOH + H₂
3. Displacement reactions (Q3):
   • B is most reactive.
   • B displaces Cu from CuSO₄.
   • Order: B > A > C > D.
4. Hydrogen gas is produced with dilute acids; Fe + H₂SO₄ → FeSO₄ + H₂ (Q4).
5. Zinc displaces iron: Zn + FeSO₄ → ZnSO₄ + Fe (Q5).