3a. Extraction of Aluminium

The current method for extracting aluminium is expensive because it involves several stages and uses large amounts of costly electrical energy. It is much more expensive than using carbon reduction to make iron in a blast furnace.

Aluminium is very abundant in the Earth's crust but it is always found as very stable compounds in many sources e.g. bauxite (mainly aluminium oxide) and alumino-silicate minerals in many rocks. Bauxite has the highest concentration of aluminium in these sources and is mined extensively around the world.

Reminders: Electrolysis (of aluminium oxide) is a way of splitting up (decomposition) of the compound (aluminium oxide) using electrical energy. The electrical energy comes from a d.c. (direct current) power pack supply. A conducting liquid, containing ions, called the electrolyte (molten aluminium oxide), must contain the compound (aluminium oxide) that is being broken down. The electricity must flow through electrodes dipped into the electrolyte to complete the electrical circuit with the battery. Electrolysis can only happen when the circuit is complete, and a d.c. electrical current (electricity) is flowing, then the products of electrolysing molten aluminium oxide are released on the electrode surfaces where they can be collected. Electrolysis always involves a flow of electrons in the external wires and electrodes and a flow of ions in the electrolyte and there is always a reduction at the negative cathode electrode (which attracts positive ions, cations) and an oxidation at the positive anode electrode (which attracts negative ions, anions).

The basic design of the industrial electrolysis cell used in the extraction of aluminium from molten purified aluminium oxide extracted from bauxite ore.

• The process of electrolysis uses of large amounts of energy in the extraction of a reactive metals and makes aluminium expensive to produce.

• Aluminium is a very useful metal but expensive to produce.

• Aluminium is theoretically a very reactive metal, so, because its position in the reactivity series of metals, aluminium cannot be extracted using carbon because it is above carbon in the reactivity series ie more reactive than carbon in the series.
  • Carbon is not reactive enough to displace aluminium from its compounds such as aluminium oxide.
• So, if aluminium is too reactive to be obtained by carbon reduction of its oxide another method must be employed which is called electrolysis.
• Aluminium is obtained from mining the mineral bauxite which is mainly aluminium oxide (Al₂O₃) and bauxite must be purified prior to electrolysis, adding to the manufacturing costs.
• The purified bauxite ore of aluminium oxide is continuously fed in. The mineral cryolite is added to lower the melting point and dissolve the ore. So the electrolyte is a mixture of molten aluminium oxide and cryolite minerals.
  • The addition of cryolite brings the melting point of aluminium oxide down to ~900^oC.
  • The melting point of aluminium oxide is over 2000^oC and it would require a lot of extra energy to keep purified bauxite ore molten for the electrolysis to take place - remember the ions (Al³⁺ and O^2–) must be free to move to electrodes for the electrolysis to work.
• The ore–compound containing the aluminium must be molten so the ions are free to move to the electrodes. The conducting melt is called the electrolyte, so extracting aluminium this way involves the electrolysis of molten aluminium oxide. See the electrolysis cell diagram on the left
• Ions must be free to move to the electrodes called the cathode (–, negative), attracting positive ions e.g. Al³⁺, and the anode (+, positive) which attracts negative ions e.g. O^2–.
• When the d.c. current is passed through aluminium forms at the negative cathode (metal*) and sinks to the bottom of the tank where it can tapped off, collected and run into moulds to cool down before transportation to it will be used to make things.
• The waste gases; At the positive anode, oxygen gas is formed (non–metal*). This is quite a problem. At the high temperature of the electrolysis cell it burns and oxidises away the carbon electrodes to form toxic carbon monoxide or carbon dioxide. So the carbon–graphite electrode is regularly replaced and the waste gases dealt with! 
• It is a costly process (6x more than Fe!) due to the large quantities of expensive electrical energy needed for the process.
• * Two general rules:
  • Metals and hydrogen (from positive ions), form at the negative cathode electrode.
  • Non–metals (from negative ions), form at the positive anode electrode.

Raw materials for the electrolysis process:

• Bauxite ore of impure aluminium oxide [Al₂O₃ made up of Al³⁺ and O^2– ions]

• Carbon (graphite) for the electrodes.

• Cryolite reduces the melting point of the ore and saves energy, because the ions must be free to move to carry the current and less energy is needed to melt the aluminium oxide obtained from the bauxite ore.

• Electrolysis means using d.c. electrical energy to bring about chemical changes e.g. decomposition of a compound to form metal deposits or release gases. The electrical energy splits the compound!

• At the electrolyte connections called the anode electrode (+, attracts – ions) and the cathode electrode (–, attracts + ions). An electrolyte is a conducting melt or solution of freely moving ions which carry the charge of the electric current.

• Electrolysis reminders – the negative electrode (–) is called the cathode and attracts positive ions or cations e.g. Al³⁺, and the positive electrode (+) is called the anode and attracts negative ions or anions e.g. O^2–.
• Read the following in conjunction with the 'concept diagram' for the electrolysis of molten aluminium oxide, after the electrode equations.
• The negative cathode electrode attracts positive ions, the aluminium ion.
• At the negative (–) cathode, reduction occurs (electron gain) when the positive aluminium ions are attracted to it. They gain three electrons to change to neutral Al atoms.

  • half equation: aluminium ion (3+) + 3 electrons (–) ==> neutral and free aluminium atoms

  • Al³⁺ + 3e^– ==> Al
  • This is a reduction - electron gain. It might look a bit different from an iron ore reduction equation, but this is the equation that precisely expresses the reduction of an aluminium ore to aluminium metal.
• The positive anode attracts negative ions, the oxide ion.
• At the positive (+) anode, oxidation takes place (electron loss) when the negative oxide ions are attracted to it. They lose two electrons forming neutral atoms, which combine to form oxygen molecules.
  • half equation
  • 2O^2– ==> O₂ + 4e^– 
  • or
  • 2O^2– – 4e^– ==> O₂ 
  • This is an oxidation - electron loss.
• Note: Reduction and Oxidation always go together!
• The overall electrolytic decomposition is ...
  • aluminium oxide ==> aluminium + oxygen
  • 2Al₂O₃ ==> 4Al + 3O₂
  • 2Al₂O₃(l) ==> 4Al(l) + 3O₂(g)  (equation with state symbols)
  • and is a very endothermic process, lots of electrical energy input!
  • Note that the aluminium oxide loses its oxygen, therefore in this electrolytic process the compound aluminium oxide is reduced to the metal aluminium.

Concept diagram for the electrolysis of molten aluminium oxide

• GENERAL NOTE ON ELECTROLYSIS:
  • Any molten or dissolved material in which the liquid contains free moving ions is called the electrolyte.
  • Ions are charged particles e.g. Na⁺ sodium ion, or Cl^– chloride ion, and their movement or flow constitutes an electric current, because a current is moving charged particles.
  • What does the complete electrical circuit consist of?
    • There are two ion currents in the electrolyte flowing in opposite directions:
      • positive cations e.g. Al³⁺ attracted to the negative cathode electrode,
      • and negative anions e.g. O^2– attracted to the positive anode electrode,
      • BUT remember no electrons flow in the electrolyte, only in the graphite or metal wiring!
    • The circuit of 'charge flow' is completed by the electrons moving around the external circuit e.g. copper wire or graphite electrode, from the positive to the negative electrode
    • This e^– flow from the +ve to the –ve electrode perhaps doesn't make sense until you look at the electrode reactions, electrons released at the +ve anode move round the external circuit to produce the electron rich negative cathode electrode.
  • Electron balancing: In the above process it takes the removal of four electrons from two oxide ions to form one oxygen molecule and the gain of three electrons by each aluminium ion to form one aluminium atom.
    • Therefore for every 12 electrons you get 3 oxygen molecules and 4 aluminium atoms formed.
    • This means you can do mole ratio product calculations.
      • See electrolysis calculations in section 13. of the Chemical Calculations pages
• NOTE on RECYCLING Aluminium
  • About 39% of the aluminium in foil, car components etc. is recycled aluminium.
  • This makes good economics because recycling saves on costs AND allows a mineral resource like aluminium's bauxite ore to last a lot longer – slower depletion of the Earth's mineral ore resources will make it last longer.
    • Transport costs may be less (ie within UK now), but much more importantly
    • mining costs are omitted – energy/machinery involved in digging out the ore, crushing it, transporting the ore,
    • and the cost of actually extracting the metal from its finite ore resource – electrolysis plant, expensive electrical energy used
  • So, scrap metal merchants are doing a roaring trade at the moment.
  • The savings are partly reduced by the cost off collecting waste/scrap metal and purifying for further use.
    • It is estimated that recycling aluminium only uses 5% of the energy required to extracted the same mass of aluminium from its ore – the original aluminium extraction uses very expensive electrical energy for the electrolysis.

• The social, economic and environmental impacts of exploiting metal ores including RECYCLING are further discussed on a separate page.

• USES of ALUMINIUM

• Aluminium is very useful metal and used as a lightweight construction material eg greenhouse frames.

• Aluminium is a reactive metal but it is resistant to corrosion. This is because aluminium reacts in air to form a layer of aluminium oxide which then protects the aluminium from further attack.

  • This is why it appears to be less reactive than its position in the reactivity series of metals would predict.

• Aluminium can be made more resistant to corrosion by a process called anodising.

• For some uses of aluminium it is desirable to increase artificially the thickness of the protective oxide layer in a process is called anodising.

  • This involves removing the oxide layer by treating the aluminium sheet with sodium hydroxide solution.

  • The aluminium is then placed in dilute sulphuric acid and is made the positive electrode (anode) used in the electrolysis of the acid.

  • Oxygen forms on the surface of the aluminium and reacts with the aluminium metal to form a thicker protective oxide layer. 

• Aluminium can be alloyed to make 'Duralumin' by adding copper (and smaller amounts of magnesium, silicon and iron), to make a stronger alloy used in aircraft components (low density = 'lighter'!), greenhouse and window frames (good anti–corrosion properties), overhead power lines (quite a good conductor and 'light'), but steel strands are included to make the 'line' stronger and poorly electrical conducting ceramic materials are used to insulate the wires from the pylons and the ground.

  • There is a note about the bonding and structure of alloys on another page.

3b. Electrolytic extraction of the very reactive metal like sodium

The process of electrolysis uses of large amounts of energy in the extraction of a reactive metals like sodium, potassium, magnesium and calcium etc. and makes them expensive to produce.

Because its position in the reactivity series of metals, sodium cannot be extracted using carbon, sodium is above carbon and cannot be displaced by it. So, sodium is too reactive to be obtained by carbon reduction of its oxide and another method must be employed which is called electrolysis.

Sodium, like many of the most reactive metals, can be extracted by electrolysis of its molten chloride. This can be done in the 'Down's Cell' shown in the diagram.

Electrolysis reminders – the negative electrode (–) is called the cathode and attracts positive ions or cations e.g. Na⁺, and the positive electrode (+) is called the anode and attracts negative ions or anions e.g. Cl^–. The ore–compound containing the sodium (or other metal) must be molten so the ions are free to move to the electrodes. The conducting melt is called the electrolyte.

In the molten salt the positive sodium ions migrate to the negative cathode electrode and are reduced by electron gain to form liquid sodium atoms.

The negative cathode electrode attracts positive ions, eg the sodium ion.

At the (–) cathode half equation: Na⁺ + e^– ==> Na

This is electron gain, a reduction, so reduction of a sodium compound to give the free metal sodium.

Equally mobile in the molten chloride salt are the negative chloride ions, which migrate to the positive anode electrode and get oxidised by electron loss to form green chlorine gas molecules. Initially two chlorine atoms are formed and these rapidly combine to give chlorine molecules.

The positive anode attracts negative ions, eg the chloride ion.

At the (+) anode half equation overall: 2Cl^– ==> Cl₂ + 2e^–

or 2Cl^– ==> 2Cl + 2e^– and then 2Cl ==> Cl₂

This is an oxidation change - electron gain to free the chlorine from the sodium chloride

Overall chemical change: 2NaCl ==> 2Na + Cl₂

Other very reactive metals like lithium, potassium and calcium can be extracted in the same way by electrolysing their molten salts. As you can see from the diagram on the right, all these metals are above carbon in the reactivity series and cannot be displaced by carbon.

Some general notes on electrolysis AND

ELECTROCHEMISTRY INDEX:  1. INTRODUCTION to electrolysis – electrolytes, non–electrolytes, electrode equations   2. Electrolysis of acidified water (dilute sulfuric acid)   3. Electrolysis of sodium chloride solution (brine)   4. Electrolysis of copper(II) sulfate solution and electroplating   5. Electrolysis of molten lead(II) bromide (and other molten compounds)   6. Electrolysis of copper(II) chloride solution   7. Electrolysis of hydrochloric acid   8. Summary of electrode equations and products   9. Summary of electrolysis products from various electrolytes   10. Simple cells (batteries)   11. Fuel Cells   12. The extraction of aluminium from purified molten bauxite ore   13. Anodising aluminium to thicken and strengthen the protective oxide layer   14. The extraction of sodium from molten sodium chloride using the 'Down's Cell'   15. The purification of copper by electrolysis   16. The purification of zinc by electrolysis   17. Electroplating coating conducting surfaces with a metal layer   18. Electrolysis of brine (NaCl) for the production of chlorine, hydrogen & sodium hydroxide 19. Electrolysis calculations

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Other associated KS4 Science GCSE/IGCSE chemistry web pages on this site

• GCSE/IGCSE Foundation–tier (easier) Multiple choice QUIZ on METAL EXTRACTION

• GCSE/IGCSE Higher–tier (harder) Multiple choice QUIZ on METAL EXTRACTION

• gap–word–fill exercise on METAL EXTRACTION

• Reactivity of metals : all the reactions of metals with oxygen (air), water and acids, displacement reactions

• Oxidation–reduction explained (redox theory)

• Group 1 Alkali Metals: chemical and physical properties

• Extra Industrial chemistry : properties and uses of metals including titanium and aluminium

• Transition metals physical and chemical properties and uses

• Metallic Bonding – structure and properties of metals

• Extra Electrochemistry–electrolysis theory, experiment designs and links to all electrolysis sections

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