Chapter 13 Metals and Reactivity (notes)

Alloy: A mixture of metals.
Purpose of Alloying:
- To make the metal harder.
- To improve resistance to corrosion.
Characteristics of Pure Metals:
- Soft due to neat arrangement of atoms.
- When force is applied, atoms slide over each other, making metals ductile and malleable.
Why Alloys are Harder:
- Presence of different-sized foreign atoms disrupts the neat arrangement of metal atoms.
- This disruption prevents layers of atoms from sliding over each other.
Examples of Alloys:
- Brass:
  - Props: Stronger yet malleable.
  - Components: Copper + Zinc.
  - Uses: Musical instruments.
- Bronze:
  - Props: Very hard.
  - Components: Copper + Tin (Sn).
  - Uses: Medals, statues.
- Stainless Steel:
  - Props: Does not rust.
  - Components: Iron + Chromium + Nickel.
  - Uses: Surgical instruments, cutlery.
- Duralumin:
  - Props: Light.
  - Components: Aluminium alloy.
  - Uses: Aircraft body.
- Solder:
  - Props: Low melting point.
  - Components: Lead + Tin.
  - Uses: Soldering (joining metal wires).

Reactivity Series of Metals:
1. Potassium (K)
2. Sodium (Na)
3. Calcium (Ca)
4. Magnesium (Mg)
5. Aluminium (Al)
6. Carbon (C)
7. Zinc (Zn)
8. Iron (Fe)
9. Tin (Sn)
10. Lead (Pb)
11. Hydrogen (H)
12. Copper (Cu)
13. Mercury (Hg)
14. Silver (Ag)
15. Gold (Au)
16. Platinum (Pt)
Mnemonic for Order: "Please carry my apple, Zainal is the loving handsome charming man, silver, gold and platinum."
Reactions of Metals:
- Metal + O₂ → Metal Oxide
- Metal + H₂O → Metal Hydroxide + H₂
- Metal + Acid → Salt + H₂
- Reactivity observations:
  - Potassium/Sodium react with water producing many bubbles.
  - Magnesium and Aluminium react less vigorously.
  - Metals like Iron, Tin, and Lead show little to no reaction.
Displacement Reactions:
- More reactive metals displace less reactive metals from compounds.

Definition: More reactive metal displaces less reactive metal from a compound.
Example Reaction:
- Zn + CuSO₄ → ZnSO₄ + Cu
- Observation:
  - Zn dissolves, blue solution turns colorless, brown precipitate of copper forms.
Understanding the Reaction:
- Zinc is more reactive than copper, resulting in copper being displaced.
Reactivity Insights:
- More reactive metals lose outer shell electrons more easily.
- Displacement involves Redox Reactions:
  - Oxidation Half Equation: Zn → Zn²⁺ + 2e⁻
  - Reduction Half Equation: Cu²⁺ + 2e⁻ → Cu
Displacement of Metals by Hydrogen:
- H₂ + CuO → Cu + H₂O
- Observation: Black CuO turns brown.

Ores:
- Most metals found in ores as ionic compounds.
- Common Ores:
  - Bauxite - Aluminium oxide
  - Cassiterite - Tin oxide
  - Haematite - Iron (III) oxide
  - Zinc blende - Zinc sulfide.
Extraction Reactions:
- Metals can react with more reactive metals or carbon to extract.
- Example with iron displacing copper:
  - Fe + CuO → FeO + Cu
  - Iron displaces copper due to higher reactivity.
Carbon Extraction Details:
- Carbon reduces metal oxides by removing oxygen.
- Only metals less reactive than carbon can be extracted this way; others require electrolysis.

Aluminium Behavior:
- Reactive but does not readily react with water or acids due to oxide layers formed.
- Freshly prepared Aluminium will react with acids but oxidizes immediately upon exposure to air, forming Al₂O₃.
Properties of Oxide Layer:
- Tough and adheres strongly to the surface, preventing further reactions.

Metal Hydroxides:
- Decompose into Metal Oxides and Water.
- Example: Zn(OH)₂ → ZnO + H₂O
- Exceptions: Group I hydroxides (Li can decompose).
Metal Nitrates:
- Decompose into metal nitrites/nitrogen oxides and oxygen.
- Example: 2KNO₃ → 2KNO₂ + O₂
Thermal Stability:
- More reactive metals = more stable compounds = require more heat for decomposition.
- Work with NO₂ must be conducted in fume cupboards due to toxicity.