14.00 Chapter Summary

1. Classification of Elements

• Metals:
  • Abundance: Most elements in the Periodic Table.
  • Properties:
    • Conduct heat and electricity due to delocalised electrons.
    • Malleable (can be hammered into shapes) and ductile (can be drawn into wires).
    • Typically have high melting and boiling points.
    • Exhibit metallic bonding with strong electrostatic attraction between positive ions and delocalised electrons.
• Non-Metals:
  • Abundance: Fewer elements compared to metals.
  • Properties:
    • Do not conduct heat and electricity (except graphite).
    • Brittle when solid; not malleable or ductile.
    • Generally have low melting and boiling points.
    • Many are gases at room temperature due to weak intermolecular forces.
    • Exceptions: Diamond and silicon(IV) dioxide have high melting and boiling points.
• Metalloids (Semi-Metals):
  • Location: Along the zig-zag line separating metals and non-metals.
  • Properties: Exhibit characteristics of both metals and non-metals.

2. Physical Properties of Metals vs. Non-Metals

3. Chemical Properties of Metals

• Reactivity Series:
  • Purpose: Order metals based on their reactivity from most to least reactive.
  • Reactivity Indicators:
    • Reactions with Water:
      • Highly Reactive Metals: React violently with cold water to form metal hydroxides and hydrogen gas.
      • Moderately Reactive Metals: React with steam to form metal oxides and hydrogen gas.
      • Unreactive Metals: Do not react with water.
    • Reactions with Acids:
      • Above Hydrogen: React with dilute acids to produce salts and hydrogen gas.
      • Below Hydrogen: Do not react with dilute acids.
    • Reactions with Oxygen:
      • Highly Reactive Metals: React easily with oxygen to form metal oxides.
      • Less Reactive Metals: React slowly or not at all.
• Reactivity Series of Metals:

• Mnemonic to Remember the Reactivity Series:
  • “Please Send Cats, Monkeys And Cute Zebras Into Hot Countries Signed Gordon”
    • P: Potassium
    • S: Sodium
    • C: Calcium
    • M: Magnesium
    • A: Aluminium
    • C: Carbon
    • Z: Zinc
    • I: Iron
    • H: Hydrogen
    • C: Copper
    • S: Silver
    • G: Gold

4. Reactions of Metals

• With Water:
  • General Equation:
    • Highly Reactive Metals: Metal + Cold Water → Metal Hydroxide + Hydrogen Gas
      • Example:
        • Calcium: Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)
        • Potassium: 2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g)
  • With Steam:
    • Less Reactive Metals: Metal + Steam → Metal Oxide + Hydrogen Gas
      • Example:
        • Zinc: Zn(s) + H₂O(g) → ZnO(s) + H₂(g)
        • Magnesium: Mg(s) + H₂O(g) → MgO(s) + H₂(g)
• With Acids:
  • General Equation:
    • Metal + Acid → Salt + Hydrogen Gas
      • Example:
        • Iron: Fe(s) + 2HCl(aq) → FeCl₂(aq) + H₂(g)
        • Zinc: Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)
  • Observations:
    • Highly Reactive Metals: Vigorous or explosive reactions.
    • Less Reactive Metals: Slower or no reaction.
• With Oxygen:
  • General Equation:
    • Metal + Oxygen → Metal Oxide
      • Example:
        • Copper: 2Cu(s) + O₂(g) → 2CuO(s)
  • Reactivity:
    • Highly Reactive Metals: React easily with oxygen.
    • Unreactive Metals: Do not react with oxygen (e.g., Gold, Platinum).

5. Uses of Metals

• Aluminium (Al):
  • Properties:
    • High strength-to-weight ratio, low density.
    • Forms a protective layer of aluminium oxide, making it appear unreactive.
  • Uses:
    • Aeroplane Bodies: Due to high strength-to-weight ratio.
    • Overhead Power Cables: Good electrical conductor, low density.
    • Saucepans: Excellent thermal conductor.
    • Food Cans: Non-toxic, corrosion-resistant.
• Copper (Cu):
  • Properties:
    • Excellent electrical and thermal conductivity.
    • Ductile and malleable.
    • Antibacterial properties.
  • Uses:
    • Electrical Wiring: High conductivity and ductility.
    • Pots and Pans: Good heat conductor, unreactive, malleable.
    • Water Pipes: Non-toxic, unreactive, malleable.
    • Surfaces in Hospitals: Antibacterial properties.
• Other Metals:
  • Iron (Fe): Used in construction, machinery, and transportation.
  • Zinc (Zn): Used in galvanizing to prevent rusting.
  • Silver (Ag) & Gold (Au): Used in jewelry, electronics, and as monetary standards.

6. Alloys: Properties and Uses

• Definition: Mixtures of a metal with other elements (usually metals) to enhance properties.
• Common Alloys:
  • Brass: Copper + Zinc
    • Properties: Stronger than either metal alone.
    • Uses: Musical instruments, ornaments, door knobs.
  • Stainless Steel: Iron + Chromium + Nickel + Carbon
    • Properties: Hard, corrosion-resistant.
    • Uses: Cutlery, medical instruments, kitchen appliances.
  • Other Examples:
    • Iron + Tungsten: Extremely hard, high-temperature resistance.
    • Aluminium + Copper/Manganese/Silicon: Stronger with low density, ideal for aircraft bodies.
• Properties of Alloys:
  • Enhanced Strength: Often stronger and harder than pure metals.
  • Corrosion Resistance: More resistant to rusting and corrosion.
  • Temperature Resistance: Can withstand extreme temperatures.
  • Tailored Properties: Can be designed for specific applications based on desired properties.
• Structure of Alloys:
  • Irregular Arrangement: Different-sized atoms distort the regular metal lattice.
  • Effect: Makes it more difficult for layers to slide over each other, increasing hardness and strength.
• Key Point: Alloys are not chemically combined (they are mixtures), which distinguishes them from compounds.

7. Rusting of Iron

• What is Rust?
  • Definition: Hydrated iron(III) oxide formed from the reaction of iron with water and oxygen.
  • Reaction:
    • Overall Equation: Fe + H₂O + O₂ → Fe₂O₃·nH₂O (rust)
  • Oxidation:
    • Iron is oxidised during rusting.
• Requirements for Rusting:
  • Presence of oxygen and water.
• Investigating Rusting:
  • Experiment Setup:
    • Test Tube 1: Open to air and exposed to water → Rusts.
    • Test Tube 2: Coated with oil (prevents air contact) → Does not rust.
    • Test Tube 3: Dry with calcium chloride (removes moisture) → Does not rust.
  • Observation: Only the iron nail in Test Tube 1 shows rusting.
• Rust Prevention Methods:
  • Barrier Methods: Coating iron with substances to prevent exposure to water and oxygen.
    • Examples:
      • Grease
      • Oil
      • Paint
      • Plastic
  • Sacrificial Protection: Using a more reactive metal to protect iron.
    • Example: Galvanising with Zinc.
  • Note: Barrier coatings must remain intact; if scratched or worn, rusting can resume.

8. Galvanising & Sacrificial Protection

• Sacrificial Protection:
  • Concept: A more reactive metal (e.g., Zinc) is attached to a less reactive metal (e.g., Iron).
  • Process:
    • Reaction: Zn → Zn²⁺ + 2e⁻
    • Protection: Zinc oxidises preferentially, protecting iron from rusting.
  • Maintenance: Replace sacrificial metal before it fully corrodes.
• Galvanising:
  • Process: Coating iron with a layer of zinc.
    • Methods: Electroplating or dipping in molten zinc.
  • Advantages:
    • Barrier Protection: Zinc layer prevents contact with oxygen and water.
    • Sacrificial Protection: If the zinc layer is damaged, zinc continues to protect iron.
  • Chemical Reaction:
    • Formation of Protective Layer: ZnCO₃ forms when zinc reacts with oxygen and carbon dioxide, enhancing protection.
• Examiner Tips:
  • Suitability: Metals higher in the reactivity series than iron (e.g., Zinc) are suitable for sacrificial protection.
  • Identification: Alloys and protective coatings often involve more reactive metals to prevent corrosion.

9. Extraction of Metals

• Metal Ores:
  • Definition: Naturally occurring minerals containing a metal combined with other elements.
  • Examples:
    • Hematite (Fe₂O₃): Iron ore.
    • Bauxite (Al₂O₃·H₂O): Aluminium ore.
    • Native Metals: Gold, Platinum (found in pure form).
• Extraction Methods Based on Reactivity Series:
  • Highly Reactive Metals (Above Carbon):
    • Method: Electrolysis.
    • Reason: Cannot be reduced by carbon; require direct reduction using electricity.
    • Examples: Potassium, Sodium, Aluminium.
  • Less Reactive Metals (Below Carbon):
    • Method: Reduction with carbon (coke) or carbon monoxide in a blast furnace.
    • Reason: Can be reduced by carbon as a reducing agent.
    • Examples: Iron, Zinc.
  • Unreactive Metals (Below Hydrogen):
    • Method: Mined directly as native metals.
    • Reason: Do not react easily, found in pure form.
    • Examples: Gold, Silver.

10. Extraction of Iron from Hematite (Fe₂O₃)

• Blast Furnace Process:
  • Raw Materials:
    • Iron Ore (Hematite)
    • Coke (Carbon)
    • Limestone (Calcium Carbonate, CaCO₃)
  • Stages:
    1. Zone 1:
       • Reaction: C(s) + O₂(g) → CO₂(g)
       • Purpose: Burns coke to produce carbon dioxide, releasing heat.
    2. Zone 2:
       • Reaction: CO₂(g) + C(s) → 2CO(g)
       • Purpose: Reduces carbon dioxide to carbon monoxide, a reducing agent.
    3. Zone 3:
       • Reaction: Fe₂O₃(s) + 3CO(g) → 2Fe(l) + 3CO₂(g)
       • Purpose: Reduces iron(III) oxide to molten iron.
    4. Formation of Slag:
       • Reaction: CaCO₃(s) → CaO(s) + CO₂(g)
       • Reaction: CaO(s) + SiO₂(s) → CaSiO₃(l)
       • Purpose: Removes impurities (e.g., silicon dioxide) by forming slag (calcium silicate).
• Products:
  • Molten Iron: Collected at the bottom.
  • Slag: Floats on top and is removed.
• Examiner Tips:
  • Symbol Equations: Core students do not need to memorize symbol equations for each zone.
  • Key Concepts: Understand the role of each raw material and the purpose of each reaction stage.

11. Extraction of Aluminium from Bauxite (Al₂O₃)

• Electrolysis Process:
  • Ore: Bauxite (Al₂O₃·H₂O).
  • Purification:
    • Bayer Process: Purifies bauxite to produce aluminium oxide (Al₂O₃).
  • Preparation for Electrolysis:
    • Dissolution: Aluminium oxide is dissolved in molten cryolite (Na₃AlF₆) to lower its melting point.
  • Electrolysis Cell:
    • Components:
      • Cathode (Negative Electrode): Aluminium ions gain electrons (reduction) to form aluminium metal.
        • Reaction: Al³⁺ + 3e⁻ → Al(s)
      • Anode (Positive Electrode): Oxide ions lose electrons (oxidation) to form oxygen gas.
        • Reaction: 2O²⁻ → O₂(g) + 4e⁻
  • Overall Reaction:
    • 2Al₂O₃(l) → 4Al(l) + 3O₂(g)
  • Additional Reaction:
    • Carbon in graphite anodes reacts with oxygen to form carbon dioxide.
      • Reaction: C(s) + O₂(g) → CO₂(g)
    • Effect: Anodes wear away and need regular replacement.
• Energy Consideration:
  • Electricity Requirement: High amount of electricity makes aluminium extraction energy-intensive and expensive.
• Examiner Tips:
  • Core Students: Do not need to explain the electrolysis process in detail.
  • Extended Students: Must understand the electrolysis steps and the overall chemical reactions involved.

12. Displacement Reactions in the Reactivity Series

• Concept: A more reactive metal can displace a less reactive metal from its compound.
• Mechanism:
  • More Reactive Metal: Loses electrons easily, forming ions.
  • Less Reactive Metal: Gains electrons, forming elemental metal.
• Example Reactions:
  • Magnesium Displacing Copper:
    • Reaction: Mg(s) + CuSO₄(aq) → MgSO₄(aq) + Cu(s)
    • Observation: Blue CuSO₄ solution fades; copper precipitates as a solid.
  • Zinc Displacing Iron:
    • Reaction: Zn(s) + FeCl₂(aq) → ZnCl₂(aq) + Fe(s)
    • Observation: Iron metal is deposited.
  • Iron Displacing Silver:
    • Reaction: Fe(s) + 2AgNO₃(aq) → Fe(NO₃)₂(aq) + 2Ag(s)
    • Observation: Silver metal is deposited.
• Displacement Table Summary:

• Reactivity Order Deduced:
  • Mg > Zn > Fe > Cu > Ag
• Examiner Tips:
  • Reactivity Order: Use displacement reactions to establish or confirm the order of reactivity.
  • Observations: Look for color changes, precipitation of metals, or fading of solutions.

13. Explaining Reactivity

• Electron Loss:
  • Tendency to Lose Electrons: High reactivity metals lose electrons easily, forming positive ions.
• Reaction Tendency:
  • High in Reactivity Series: Metals lose electrons readily, making them good reducing agents.
  • Low in Reactivity Series: Metals hold onto electrons tightly, making them less reactive.
• Displacement Reactions:
  • More Reactive Metal: Displaces less reactive metal from its compounds.
  • Less Reactive Metal: Gets reduced to its elemental form.
• Key Concept: OIL RIG (Oxidation Is Loss, Reduction Is Gain) to remember electron transfer.

14. Rusting Prevention Techniques

• Barrier Methods:
  • Coatings: Apply grease, oil, paint, or plastic to prevent exposure to water and oxygen.
  • Effectiveness: Only effective if coatings remain intact.
• Galvanising:
  • Method: Coat iron with a layer of zinc.
  • Benefit: Zinc acts as a sacrificial metal, oxidising before iron.
• Sacrificial Protection:
  • Method: Attach a more reactive metal (e.g., Zinc) to protect iron.
  • Benefit: More reactive metal corrodes first, shielding iron.
• Examiner Tips:
  • Suitability: Use metals higher in the reactivity series for sacrificial protection.
  • Identification: Recognize methods based on the type of protection provided.

15. Summary Tables

Properties Comparison: Metals vs. Non-Metals

Reactivity Series Mnemonic

• “Please Send Cats, Monkeys And Cute Zebras Into Hot Countries Signed Gordon”
  • P: Potassium
  • S: Sodium
  • C: Calcium
  • M: Magnesium
  • A: Aluminium
  • C: Carbon
  • Z: Zinc
  • I: Iron
  • H: Hydrogen
  • C: Copper
  • S: Silver
  • G: Gold

Group 1 (Alkali Metals) Reactivity with Water

Group 7 (Halogens) Physical States at Room Temperature

Group 7 (Halogens) Displacement Reactions

Transition Metal Ion Colors

16. Key Takeaways

• Element Classification: Distinguish between metals, non-metals, and metalloids based on their physical and chemical properties.
• Reactivity Series: Understand the order of reactivity and its implications for displacement reactions and extraction methods.
• Reactions of Metals: Familiarize yourself with how different metals react with water, acids, and oxygen.
• Alloys: Recognize the properties and uses of common alloys and understand their structural differences from pure metals.
• Rusting and Prevention: Comprehend the process of rusting, its requirements, and methods to prevent it using barrier techniques and sacrificial protection.
• Extraction Methods: Know the extraction methods for various metals based on their position in the reactivity series, including the blast furnace for iron and electrolysis for aluminium.
• Exam Preparation: Utilize summary tables, mnemonics, and understand examiner tips to effectively recall and apply knowledge during IGCSE exams.