2.3 Chemistry- Metals and their extraction

• Ore


  Rock found on the Earth's crust which contains metal compounds
• Metal extraction


  • Different methods can be used to extract the pure metal from ores
  • The method used depends on the metal's position in the reactivity series
  • The more reactive a metal is, the more stable its metal compound is so the harder it is to extract pure metal
• Reactivity of metals


  • Metals more reactive than carbon are usually extracted by electrolysis
  • Metals less reactive than carbon can be extracted by reduction with carbon
  • Metals less reactive than hydrogen tend to be so unreactive they are often found pure in their native form
• Reactivity series


  Arrangement of metals in order of decreasing reactivity, with the most reactive metal at the top and the least reactive metal at the bottom
• Displacement reaction


  1. A more reactive metal displaces a less reactive metal from a compound

  2. Can be used to investigate a metal's relative reactivity

  3. Some displacement reactions can be observed, like the reaction between copper sulfate and magnesium

  4. The thermite reaction is a displacement reaction used to produce pure iron
• Oxidation


  Gain of oxygen, loss of electrons, loss of hydrogen
• Reduction


  Loss of oxygen, gain of electrons, gain of hydrogen
• A reaction in which reduction and oxidation both occur is called a redox reaction
• The species that is reduced is known as the oxidising agent
• The species that is oxidised is known as the reducing agent
• Extraction of iron in the blast furnace

  1. Iron ore (iron oxide compounds) is reduced to obtain pure iron

  2. Limestone undergoes thermal decomposition to form calcium oxide, which reacts with impurities

  3. Coke reacts with oxygen to form carbon monoxide, which then reduces the iron ore

  4. Oxygen levels are controlled so carbon monoxide forms and not carbon dioxide
• Coke reacts with oxygen


  Forms carbon dioxide, which is an exothermic reaction that heats up the furnace
• Carbon dioxide reacts with more coke

  Forms carbon monoxide, which reduces the iron oxide
• Iron oxide reacts with coke or carbon monoxide

  Is reduced to form molten iron
• Removing impurities with limestone

  1. Limestone undergoes thermal decomposition into calcium oxide

  2. Calcium oxide reacts with silicon dioxide (impurity) to form solid calcium silicate, which can be removed from the furnace
• Electrolysis


  Process of using electricity to break down a molten or dissolved compound into its elements
• Metals more reactive than carbon are usually extracted by electrolysis of molten compounds
• Metals less reactive than carbon can also be extracted by electrolysis
• Electrolysis requires large amounts of energy, making it an expensive method compared to reduction with carbon
• Electrolysis setup


  1. Positively charged ions move to the negative electrode (cathode)

  2. Negatively charged ions move to the positive electrode (anode)

  3. Ions are discharged at the electrodes, producing elements
• Half equations


  Represent the oxidation and reduction reactions occurring at the electrodes during electrolysis
• Electrolysis of molten ionic compounds

  1. Identify the ions present in the compound

  2. Positive ions move to the cathode, negative ions move to the anode

  3. Reactions at the electrodes produce the elements
• Industrial electrolysis of aluminium

  1. Aluminium oxide is melted by dissolving it in cryolite to lower the melting point

  2. Graphite electrodes are used, with aluminium metal forming at the negative electrode and oxygen forming at the positive electrode
• Electrolysis can be used to separate water into hydrogen and oxygen
• Aluminium oxide


  Melted so electricity can pass through it
• Aluminium oxide


  • Melting point is very high which makes it expensive to melt
  • Dissolved in cryolite which lowers the melting point
  • Use of molten cryolite as a solvent reduces some of the energy costs involved in extracting aluminium
• Negative electrode (cathode)


  Made of graphite, a form of carbon
• Positive electrode (anode)


  Made of graphite, a form of carbon
• Aluminium metal formation


  1. Forms at the negative electrode

  2. Sinks to the bottom of the tank and is tapped off here
• Oxygen formation


  1. Forms at the positive electrode

  2. Reacts with the carbon forming carbon dioxide, which bubbles out of the tank
• As the electrode itself reacts with oxygen the positive electrode gradually burns away
• The positive electrode has to be replaced often, adding to the cost of the process
• Electrolysis of water


  Used to separate water into hydrogen gas and oxygen gas
• Electrolysis of water


  1. Positive hydrogen ions move to the negative electrode (cathode) and gain electrons in a reduction reaction

  2. Negative OH- ions move to the positive electrode (anode) and lose electrons to form water and oxygen in an oxidation reaction
• The overall equation for the electrolysis of water is: 2H2O (l) → 2H2 (g) + O2 (g)
• Ionic solution


  Contains the ions that make up the ionic compound, and the ions in water (OH- and H+)
• Electrolysis of aqueous solutions

  1. At the cathode (-): Hydrogen (from H+ in water) is produced UNLESS the + ions in the ionic compound are from a metal less reactive than hydrogen

  2. At the anode (+): Oxygen (from OH- in water) will be produced UNLESS the ionic compound contains halide ions (Cl-, Br-, I-)
• Copper chloride solution


  • Cu+ ions go to cathode, Cu (s) is produced (Cu is less reactive than hydrogen)
  • Cl- ions go to anode, Cl2 (g) is produced (Cl- are halide ions)
• Sodium chloride solution


  • H+ ions go to cathode, H2 (g) is produced (Na is more reactive than hydrogen)
  • Cl- ions go to anode, Cl2 (g) is produced (Cl- are halide ions)
• Sodium sulfate solution


  • H+ ions go to cathode, H2 (g) is produced (Na is more reactive than hydrogen)
  • OH- ions go to anode, O2 (g) is produced (SO4^2- ions are not halide ions)