Explain the apparent unreactivity of aluminium in terms of its oxide layer

Metals – Reactivity Series (Cambridge IGCSE 0620)

Learning objective

Explain why aluminium, although relatively high in the reactivity series, often appears unreactive. Emphasise the role of the surface aluminium‑oxide layer and connect the explanation to all relevant parts of the IGCSE Chemistry syllabus.

1. Quick syllabus refresher (core + supplementary points)

2. The reactivity series (full excerpt)

Why does bulk aluminium often look “inactive”?

Although aluminium lies above zinc and iron in the series, a self‑protecting aluminium‑oxide layer forms instantly on exposure to air, masking its true reactivity.

3. Formation & nature of the aluminium‑oxide layer

• Rapid oxidation (seconds on exposure to air)

  
  \$4\text{Al (s)} + 3\text{O}2(g) \rightarrow 2\text{Al}2\text{O}_3(s)\$

• Properties of the Al₂O₃ film

  • Thickness: 3–5 nm (≈ 5–10 atomic layers).
  • Hard, dense, colourless, adheres tightly to the metal.
  • Electrically insulating.
  • Self‑healing: any breach is re‑oxidised within seconds.

4. How the oxide layer masks reactivity

1. Physical barrier – prevents acids, water or steam from contacting the underlying metal.
2. Chemical passivation – the dense Al₂O₃ stops further oxidation of the metal.
3. Electrical insulation – blocks the flow of electrons required for redox processes.

Consequences:

• No visible bubbling when a clean aluminium strip is placed in dilute HCl.
• No reaction with steam unless the film is removed (e.g., by scratching or by using a strong base).

5. Demonstrations & practical work

• Acid test – untreated aluminium

  • Place a clean piece of aluminium in ~1 M HCl. No observable H₂ evolution.

• Acid test – oxide removed

  • Scratch the surface with fine sandpaper (or dip briefly in liquid mercury). Add to the same acid. Rapid bubbling of H₂.

• Base test – concentrated NaOH (≈ 5 M)

  
  \$2\text{Al} + 2\text{NaOH} + 6\text{H}2\text{O} \rightarrow 2\text{Na[Al(OH)}4] + 3\text{H}_2\uparrow\$

  • Vigorous bubbling; solution becomes clear, colourless sodium tetrahydroxoaluminate.

• Thermite‑type demonstration

  • Mix aluminium powder with Fe₂O₃ and ignite. The reaction proceeds violently, showing the true reducing power of aluminium when the oxide barrier is absent.

6. Stoichiometry worked example

Question: How many grams of aluminium are required to produce 22.4 L of H₂ (collected over water at STP) when reacted with excess dilute HCl?

1. Balanced equation:
   \$\text{Al} + 3\text{HCl} \rightarrow \text{AlCl}3 + \tfrac{3}{2}\text{H}2\$
2. From the equation, 1 mol Al → 1.5 mol H₂.
3. At STP, 1 mol gas = 22.4 L, so 22.4 L H₂ = 1 mol H₂.
4. Moles of Al needed = 1 mol ÷ 1.5 = 0.667 mol.
5. Mass of Al = 0.667 mol × 26.98 g mol⁻¹ ≈ 18 g.

This links the metal’s reactivity to the quantitative part of the syllabus (moles, molar mass, gas‑volume relationships).

7. Redox perspective

Aluminium’s reaction with steam:

Balancing electrons (multiply the reduction half‑reaction by 3 and the oxidation by 2) gives the overall equation:

\$2\text{Al} + 3\text{H}2\text{O} \rightarrow \text{Al}2\text{O}3 + 3\text{H}2\$

This satisfies the IGCSE requirement to identify oxidation numbers, write half‑equations, and combine them.

8. Chemical energetics

• ΔH (2 Al + 3 H₂O → Al₂O₃ + 3 H₂) ≈ ‑ 850 kJ mol⁻¹ (exothermic).
• The large heat release explains why the reaction, once initiated, is self‑sustaining and why aluminium is a key component of thermite mixtures.

9. Periodic‑table context

• Group 13 (B, Al, Ga, In, Tl). The +3 oxidation state follows the group‑number rule.
• Metallic character increases down the group, so aluminium is moderately metallic – consistent with its position in the reactivity series.

10. Industrial extraction of aluminium

1. Bauxite mining → Bayer process – bauxite is refined to pure alumina (Al₂O₃).
2. Hall‑Héroult electrolytic reduction (the only commercial method):

   
   At the cathode: Al³⁺ + 3e⁻ → Al(l)

   
   At the anode (carbon): C + O²⁻ → CO₂ + 4e⁻

   
   Overall: Al₂O₃(l) + 3C → 2Al(l) + 3CO₂(g)

   • Molten Al₂O₃ is dissolved in cryolite (Na₃AlF₆) to lower the melting point.
   • Carbon anodes are consumed, producing CO₂.

Link to the syllabus: this demonstrates electrolysis (core) and connects aluminium’s reactivity to its industrial production.

11. Alloys & modification of the oxide layer

• Al‑Cu (e.g., Duralumin) – copper disrupts the continuity of the oxide film, increasing strength for aircraft.
• Al‑Mg (marine alloys) – magnesium improves corrosion resistance in salty environments.
• Al‑Si (casting alloys) – silicon refines grain structure and can make the oxide layer more porous, aiding casting.
• Alloying can either strengthen the protective film (e.g., small Si additions) or make it more permeable, influencing apparent reactivity.

12. Comparison of protective oxide layers

13. Summary checklist

• Aluminium is a moderately reactive metal (group 13, +3 oxidation state).
• Exposure to air instantly produces a thin (3–5 nm) Al₂O₃ layer.
• The oxide layer acts as a physical, chemical and electrical barrier → “apparent unreactivity”.
• Removing or bypassing the layer (scratching, strong base, alloying) reveals the true reactivity.
• Key reactions (once the film is absent):

  • With dilute HCl: Al + 3HCl → AlCl₃ + 3/2 H₂
  • With steam: 2Al + 3H₂O → Al₂O₃ + 3H₂ ΔH ≈ –850 kJ
  • With concentrated NaOH: 2Al + 2NaOH + 6H₂O → 2Na[Al(OH)₄] + 3H₂

• Industrial relevance: Hall‑Héroult electrolytic production of aluminium.
• Alloying and corrosion considerations are part of the broader “metals” sub‑topic.
• Common uses: aircraft, beverage cans, foil, heat‑sinks, electrical conductors, marine fittings.

14. Sample exam‑style questions

1. Explain why a clean piece of aluminium does not visibly react with dilute hydrochloric acid, but a piece that has been scratched does. Include the relevant chemical equations.
2. Write the balanced overall equation for the reaction of aluminium with steam. Then:

   • Identify the oxidation and reduction half‑reactions.
   • State whether the reaction is exothermic or endothermic and give an approximate ΔH value.

3. Compare the protective oxide layers formed on aluminium and iron. Discuss how each influences long‑term corrosion and the practical implications for the use of these metals.
4. Calculate the mass of aluminium required to produce 5.0 L of hydrogen gas (collected over water at 25 °C and 1 atm) when reacted with excess dilute HCl. (Molar volume at 25 °C = 24.5 L mol⁻¹.)
5. Describe the Hall‑Héroult process for extracting aluminium from alumina, including the role of electricity and the by‑products formed.