The reactivity series

• • Metal atoms form positive ions by loss of electrons when they react with other substances
  • The tendency of a metal to lose electrons is a measure of how reactive the metal is
• • A metal that is high up on the series loses electrons easily and is thus more reactive than one which is lower down on the series
• • Note that although carbon and hydrogen are non-metals, they are included in the series as they are useful in extracting metals from their oxides by reduction processes

  
• Metal reacting with water

  • Some metals react with water
  • Metals above hydrogen in the reactivity series will react with water
  • For some metals such as iron, the reaction may be very slow
  • For other metals such as the alkali metals, the reaction may be quick and potentially hazardous because of their reactivity
• • Metals that react with cold water form a metal hydroxide and hydrogen gas:

  metal + water → metal hydroxide + hydrogen

  • For example, calcium:

  Ca       +    2H2O     → Ca(OH)2      +      H2

  calcium + water → calcium hydroxide + hydrogen
• • Magnesium reacts very slowly with cold water when finely divided
  • Magnesium reacts with gaseous water to form a metal oxide and hydrogen gas:

  Mg       +    H2O     → MgO      +      H2 

  magnesium + water → magnesium oxide + hydrogen
• Reaction with water

  • Iron- slow rusting
• • Most metals react with dilute acids such as HCl
  • Only the ones below hydrogen in the reactivity series will not react with acids
• • When acids and metals react, the hydrogen atom in the acid is replaced by the metal atom to produce a salt and hydrogen gas:

  metal + acid → metal salt + hydrogen

  • For example iron:

  Fe    +     2HCI    →    FeCl2    +    H2

  iron + hydrochloric acid   →  iron(II)chloride + hydrogen

  • In both these types of reactions (water and acids) the metals are becoming positive ions
• • The reactivity of the metals is related to their tendency to become an ion
  • The more reactive the metal the more easily it becomes an ion (by losing electrons)
• • Aluminium is high in the reactivity series, but it does not react with water and the reaction with dilute acids can be quite slow.
  • This is because it has a protective oxide layer that prevents reaction with these reagents.
• • Carbon is a cheap reducing agent which can be used to remove oxygen from metal oxide ores
  • Placing carbon in the reactivity series allows us to see whether a metal oxide can be reduced or not by carbon
• • From the reactions with water and acids we have seen that whether a reaction takes place depends on the position of the metal in the reactivity series relative to hydrogen
  • A reaction takes place if the metal is able to displace hydrogen from water or acids
• • Metals below carbon can be extracted by heating the metal oxide with carbon
  • Metals higher than carbon have to be extracted by other methods, such as electrolysis
• • The reactivity of metals decreases going down the reactivity series.
  • This means that a more reactive metal will displace a less reactive metal from its compounds
  • Two examples are:
  • Reacting a metal with a metal oxide (by heating the metal will reduce the other metal from its oxide)
  • Reacting a metal with an aqueous solution of a metal compound
  • For example, it is possible to reduce copper(II) oxide by heating it with zinc.
  • The reducing agent in the reaction is zinc:

  Zn    +     CuO    →    ZnO    +    Cu

  zinc + copper oxide → zinc oxide + copper
• • The thermite reaction is a highly exothermic reaction that occurs when iron(III) oxide is reacted with aluminium metal
  • As aluminium is more reactive than iron, the iron in iron(III) oxide is displaced 

  iron(III) oxide + aluminium → iron + aluminium oxide

  Fe2O3 + 2Al → 2Fe + Al2O3

  • The reaction releases so much heat that the iron formed is molten so this process is used in welding and incendiary devices
• • The reactivity between two metals can be compared using displacement reactions in salt solutions of one of the metals
  • This is easily seen as the more reactive metal slowly disappears from the solution, displacing the less reactive metal
  • For example, magnesium is a reactive metal and can displace copper from a copper sulfate solution:

  Mg + CuSO4  →  MgSO4 + Cu

  • The blue colour of the CuSO4 solution fades as the copper ions are reduced and colourless magnesium sulfate solution is formed.
  • Copper coats the surface of the magnesium and also forms solid metal which falls to the bottom of the beaker