Metals

Overview

Comprehensive guide covering the chemical and physical properties of metals, reactivity, extraction, and corrosion. This document covers both Core and Extended syllabus requirements.

Syllabus Coverage: Core & Extended

1. Properties and Structure

Physical Properties (Core)

  • Appearance: Lustrous (shiny) solids at room temperature (Exception: Mercury is liquid).
  • Mechanical Properties: Malleable (can be hammered into sheets) and ductile (can be drawn into wires).
  • Conductivity: High thermal and electrical conductivity due to a lattice of positive ions surrounded by a “sea” of delocalized electrons.
  • Density: Generally high density and high melting/boiling points.

Image: Diagram of metallic bonding showing positive ions and delocalized electrons

Chemical Properties (Core/Extended)

  • Electron Transfer: Metals typically lose valence electrons to form positive ions (cations).
  • Oxidation: The process of losing electrons.
    • General: $\text{Metal} \rightarrow \text{Metal}^+ + e^-$
  • Protective Oxide Layers (Supplement): Some metals, like aluminium, are highly reactive but appear unreactive because they form a thin, tough, impermeable layer of metal oxide ($\text{Al}_2\text{O}_3$) on the surface, which prevents further reaction.

2. The Reactivity Series (Core/Extended)

The reactivity series ranks metals by their ability to lose electrons.

The Series (Decreasing Reactivity):

  1. Potassium (K) - Most Reactive
  2. Sodium (Na)
  3. Calcium (Ca)
  4. Magnesium (Mg)
  5. Aluminium (Al)
  6. Carbon (C) - Non-metal reference
  7. Zinc (Zn)
  8. Iron (Fe)
  9. Tin (Sn)
  10. Lead (Pb)
  11. Copper (Cu)
  12. Silver (Ag)
  13. Gold (Au) - Least Reactive

3. Chemical Reactions (Core)

Reaction with Oxygen

Most metals react with oxygen to form metal oxides: $$\text{Metal} + \text{Oxygen} \rightarrow \text{Metal Oxide}$$

  • Highly reactive metals (K, Na) react vigorously; less reactive metals (Cu) require heating.

Reaction with Water

  • Very Reactive (K, Na): React violently with cold water to produce metal hydroxide and hydrogen gas.
    • $\text{2Na}(\text{s}) + \text{2H}_2\text{O}(\text{l}) \rightarrow \text{2NaOH}(\text{aq}) + \text{H}_2(\text{g})$
  • Medium Reactive (Mg, Al, Zn, Fe): React slowly with steam to produce metal oxide and hydrogen gas.
    • $\text{Mg}(\text{s}) + \text{H}_2\text{O}(\text{g}) \rightarrow \text{MgO}(\text{s}) + \text{H}_2(\text{g})$

Reaction with Acids

Reactive metals displace hydrogen from dilute acids to form a salt and hydrogen gas: $$\text{Metal} + \text{Acid} \rightarrow \text{Salt} + \text{Hydrogen}$$

  • Metals below Hydrogen in the reactivity series (Cu, Ag, Au) do not react with dilute acids.

4. Displacement Reactions (Extended)

A more reactive metal will displace a less reactive metal from its compound (salt solution).

  • Example: $\text{Zn}(\text{s}) + \text{CuSO}_4(\text{aq}) \rightarrow \text{ZnSO}_4(\text{aq}) + \text{Cu}(\text{s})$
  • Zinc is more reactive than copper; therefore, it displaces copper from the solution.

5. Extraction of Metals (Extended)

Extraction is based on the metal’s position in the reactivity series.

A. Metals below Carbon (Zn, Fe, Sn, Pb, Cu)

Extracted via reduction using carbon (coke).

Blast Furnace

  • The Blast Furnace (Iron Extraction):
    1. Fuel: Coke (C) burns to produce $\text{CO}_2$: $\text{C}(\text{s}) + \text{O}_2(\text{g}) \rightarrow \text{CO}_2(\text{g})$
    2. Reducing Agent Formation: $\text{CO}_2$ reacts with more coke to form Carbon Monoxide ($\text{CO}$): $\text{C}(\text{s}) + \text{CO}_2(\text{g}) \rightarrow 2\text{CO}(\text{g})$
    3. Reduction of Iron Ore: $\text{CO}$ removes oxygen from haematite ($\text{Fe}_2\text{O}_3$): $\text{Fe}_2\text{O}_3(\text{s}) + 3\text{CO}(\text{g}) \rightarrow \text{2Fe}(\text{l}) + 3\text{CO}_2(\text{g})$
    4. Flux and Slag: Limestone ($\text{CaCO}_3$) decomposes to calcium oxide ($\text{CaO}$), which removes silica ($\text{SiO}_2$) as slag: $\text{CaCO}_3(\text{s}) \rightarrow \text{CaO}(\text{s}) + \text{CO}_2(\text{g})$ $\text{CaO}(\text{s}) + \text{SiO}_2(\text{s}) \rightarrow \text{CaSiO}_3(\text{l})$

B. Metals above Carbon (K, Na, Ca, Mg, Al)

Extracted via electrolysis of their molten oxides (since carbon is not strong enough to reduce them). Image: Diagram of the Hall-Héroult cell for Aluminium extraction

  • Aluminium Extraction (Hall-Héroult Process):
    1. Ore: Bauxite ($\text{Al}_2\text{O}_3$).
    2. Cryolite: $\text{Al}_2\text{O}_3$ has a very high melting point. It is dissolved in molten cryolite ($\text{Na}_3\text{AlF}_6$) to lower the melting point, reducing energy costs.
    3. Process: Molten $\text{Al}_2\text{O}_3$ is electrolyzed.
    • Cathode (Reduction): $\text{Al}^{3+} + 3\text{e}^- \rightarrow \text{Al}$
    • Anode (Oxidation): $2\text{O}^{2-} \rightarrow \text{O}_2 + 4\text{e}^-$
    1. Anodes: Graphite anodes are used but are slowly consumed by oxygen produced at the anode, forming $\text{CO}_2$.

6. Alloys and Corrosion (Core/Extended)

Alloys (Extended)

An alloy is a mixture of a metal with another element.

  • Structure: Pure metals have regular layers of atoms that slide easily (making them soft). Alloys introduce atoms of different sizes, which disrupts the regular lattice and prevents layers from sliding.
  • Result: Alloys are typically harder and stronger than pure metals.

Alloys

Corrosion (Core)

  • Rusting: Specifically refers to the corrosion of iron. Requires both water and oxygen.
    • $\text{Fe} \rightarrow \text{Fe}^{2+} + 2e^-$ (Oxidation)
  • Prevention:
    1. Barrier Methods: Painting, oiling, or plastic coating.
    2. Galvanizing: Coating iron with Zinc. Zinc is more reactive than iron; it loses electrons more readily and is oxidized first, providing sacrificial protection.
    • Sacrificial Protection (Supplement): A more reactive metal (e.g., Mg or Zn) is attached to the iron. The more reactive metal is oxidized (loses electrons) instead of the iron, preventing rust.
    1. Alloying: Adding Chromium to iron to create Stainless Steel, which forms a protective oxide layer.

7. Amphoteric Oxides (Extended)

Some metal oxides can react with both acids and bases to form salts. These are called amphoteric.

  • Examples: Aluminium oxide ($\text{Al}_2\text{O}_3$) and Zinc oxide ($\text{ZnO}$).
  • $\text{Al}_2\text{O}_3 + 6\text{HCl} \rightarrow \text{2AlCl}_3 + 3\text{H}_2\text{O}$ (Acting as base)
  • $\text{Al}_2\text{O}_3 + 2\text{NaOH} + 3\text{H}_2\text{O} \rightarrow \text{2Na[Al(OH)}_4]$ (Acting as acid)

8. Redox Principles (Extended)

  • Oxidation: Loss of electrons / Increase in oxidation state / Gain of oxygen.
  • Reduction: Gain of electrons / Decrease in oxidation state / Loss of oxygen.
  • Redox: A reaction where both oxidation and reduction occur simultaneously.

Quick Summary (Simplified)

  • What are metals? Shiny, conductive solids that lose electrons to become positive.
  • Reactivity: Top of the list = very reactive = need electricity (electrolysis) to extract. Bottom of the list = less reactive = can use carbon/heat to extract.
  • Displacement: Stronger (more reactive) metal kicks out the weaker one from a compound.
  • Rust: Iron + Water + Oxygen = Rust. Stop it with paint or Zinc (galvanizing).
  • Alloys: Mix metals $\rightarrow$ disrupt layers $\rightarrow$ make it harder.
  • Redox: Oxidation = lose electrons; Reduction = gain electrons.

Attempt the quiz: C9 - Metals Quiz