3c. Metal Displacement Reactions (metal + salt solution)

• A displacement reaction is where a more reactive element displaces a less reactive element.

• Displacement reactions involve a simultaneous oxidation and reduction, which is why they are called redox reactions.

  • In this experiment it is important to understand the 'electron' definitions of oxidation and reduction.

  • Oxidation is electron loss - the atom/ion/molecule losing one or more electrons is said to be oxidised.

  • Reduction is electron gain - the atom/ion/molecule gaining one or more electrons is said to be reduced.

  • You must learn these definitions!

  • There are lots more examples on my Introducing REDOX reactions page.

• All you need is a selection of metals, salt solutions of the same metals (eg. chloride, nitrate or sulphate) and rack of test tubes.

• The physical state of the metal in terms of granule size or area of sheet doesn't matter in this experiment.

• You just pop the bits of metals into the solutions and carefully observe what happens on the surface.

• See also metal extraction experiments in section 3(d)

3c. Results: Table of possible observations

• No reaction = no observed change, in some cases where theoretically there should be a reaction, you might not see any change (see Al note below).

• Some of the best observations are with the more reactive metals than copper.

• The picture on the right shows what happens if you put a strip of magnesium in copper sulfate solution.

  • The solution starts of a deep blue of the copper sulfate solution.

  • Then the blue colour fades as the brown copper deposit forms on the magnesium.

  • Eventually most of the magnesium dissolves and the solution becomes colourless as magnesium sulfate is formed (which is colourless!).

• Aluminium, again gives problems with the observations because of the oxide layer inhibiting the reaction with the salt solution of a less reactive metal.

• The displaced copper can display a variety of colours depending on how the precipitate-coating forms, and how much of it - I'm afraid that's the way it is!

• In some cases the metal crystals formed by the displacement reaction are very small and scatter the light so that they can look quite dark - almost black, rather than a silver-shiny precipitate.

3c. Conclusions and comments

• Any colour change indicates a displacement reaction has occurred,

• AND the rule is that

• a more reactive metal will displace a less reactive metal from its salt solution (irrespective of which salt)

• The way that you deduce the order is quite simple and logical.

  • e.g. zinc cannot displace magnesium, but it can displace copper, and magnesium displaces both zinc and copper,

  • therefore zinc is less reactive than magnesium and zinc is more reactive than copper,

  • therefore you can deduce the reactivity order for these three metals is Mg > Zn > Cu.

• Of the series of metals tested magnesium appears to be the most reactive because it displaces all the other metals being investigated here.

• Copper appears to be the least reactive in this limited series because it does not displace any of the other metals.

• Therefore by considering what will displace what from the results table, you can quite legitimately deduce that in terms of reactivity order

  • magnesium > aluminium > zinc > iron > lead > copper

  • BUT the observations for aluminium and lead may not show up clearly.

  • What you can definitely deduce from the observations in a more restricted experiment is the reactivity series order

  • magnesium > zinc > iron > copper

  • There can be problems in observations of the reactivity of aluminium and lead due to a protective oxide layer.

• Examples of displacement equations

• The first metal on the left is the most reactive, and the last metal on the right, is the displaced less reactive metal.

  • magnesium + copper sulphate ==> magnesium sulphate + copper
    • Mg(s) + CuSO_4(aq) ==> MgSO_4(aq) + Cu_(s)
    • The blue of the copper sulphate solution fades as colourless magnesium sulfate is formed and the brown deposit of copper increases on the surface of the magnesium.
    • -
  • zinc + copper sulphate ==> zinc sulphate + copper
    • Zn(s) + CuSO_4(aq) ==> ZnSO_4(aq) + Cu_(s)
    • _-
  • aluminium + copper(II) sulphate ==> aluminium sulphate + copper
    • 2Al(s) + 3CuSO₄(aq) ==> Al₂(SO₄)₃(aq) + 3Cu(s)
    • -

    • See also 'The Reactivity Series of Metals'

    • -

• Again, these reactions can be seen as redox reactions i.e. an oxidation-reduction reaction in terms of electron loss and gain.

  • Oxidation is electron loss and the neutral metal atom loses electrons to form a positive ion (cation).

    • eg.  Mg  ==>  Mg²⁺  +  2e^-,   Zn  ==>  Zn²⁺  +  2e^-   and  Al  ==>  Al³⁺  +  3e^-

    • Reminder: The more easily the electrons are lost, the more reactive the metal.

  • Reduction is electron gain and the hydrogen ions gain electrons to form a neutral hydrogen molecule.

    • Cu²⁺  +  2e^-  ==>  Cu

  • You can then combine these half reaction to give the full ionic equation that excludes spectator ions like the chloride ion, which simply remain in solution.

    • These are the full ionic equations (redox equations) for metal displacement reactions, and I've included state symbols too.

      • Mg(s)  +  Cu²⁺(aq)  ==>  Mg²⁺(aq)  +  Cu(s)

      • 

      • -

      • Zn(s)  +  Cu²⁺(aq)  ==>  Zn²⁺(aq)  +  Cu(s)

      • 

      • 2Al(s)  +  2Cu²⁺(aq)  ==>  2Al³⁺(aq)  + 3Cu(s)  (a bit tricky!)

      • Note the electrons are not shown because electrons lost = electrons gained, so cancel out.

• and for more on redox theory behind displacement reactions

3d. Simple Metal Extraction Experiments  (with reducing agents)

• Metals of lower reactivity (below aluminium) can be displaced-extracted using carbon as the reducing agent.
• Heating oxides or carbonates with powdered charcoal (mainly carbon)
  • If you strongly heat copper(II) carbonate with finely powdered charcoal (mainly carbon) you can reddish-brown specks of copper in resulting mixture. The dark green copper(II) carbonate turns black initially as copper(II) oxide is formed, but this is then reduced to copper by the charcoal.
    • initially a thermal decomposition: CuCO₃ ==> CuO + CO₂
    • then the reduction reaction (O loss): 2CuO + C ==> 2Cu + CO₂
    • You can of course start the experiment with copper(II) oxide, but copper carbonate is closer to the sort of naturally occurring copper ore that is mined.
  • You can do a similar experiment by heating lead(II) oxide with powdered carbon and you can get silvery lead formed BUT in a fume cupboard please, since lead fumes are very poisonous!
    • 2PbO + C  ==> 2Pb + CO₂
    • Lead is a sufficiently unreactive metal for carbon to displace it.
  • If you heat iron oxides with powdered carbon nothing happens however strongly you heat the mixture in a test tube.
    • Although carbon is reactive enough to displace iron, the temperature in the test tube isn't high enough - you need a blast furnace!
    • If you heat the white powder of aluminium oxide with carbon powder, nothing happens because carbon isn't 'reactive' enough to displace aluminium ie aluminium is too reactive forming very stable compounds.
    • For more see notes on Metal Extraction Fe, Cu, Al etc.
• Metal displacement reactions with solid oxides plus reactive metal as the reducing agent
  • Here you can use a more reactive metal to displace a less reactive one from its oxide.
  • Your teacher can demonstrate two metal extraction displacement reactions.
  • Neither are used commercially, but they do illustrate how chromium and titanium are extracted.
  • Both these reactions are very exothermic - lots of heat released
  • The Thermit Reaction
    • You ignite a mixture of brown iron(III) oxide and silvery grey aluminium powder using a magnesium fuse.
    • The mixture goes off like a firework in a shower of sparks - a VERY exothermic reaction!
    • On examining the cold residue you find a lump of iron and specks of white aluminium oxide.
    • 2Al + Fe₂O₃ ==> Al₂O₃ + 2Fe
    • Aluminium is more reactive than iron and so will displace iron from iron compounds.
  • Heating a mixture of magnesium powder and copper(II) oxide
    • Silvery grey magnesium powder is mixed with black copper oxide powder and heated strongly in a test tube.
    • The mixture glows red hot and on examining the cold mixture you see specks of white magnesium oxide and reddish-brown specks of copper.
    • Mg + CuO ==> MgO + Cu
    • Magnesium is more reactive than copper and so will displace copper from copper compounds.
• See also Metal Displacement Reactions (metal + salt solution) section 3c

• and 'The Reactivity Series of Metals'

4. Action of heat on carbonates

Often, but not always, the more reactive a metal, the more thermally stable is the metal carbonate i.e.

reactivity trend with water and acids: sodium > calcium > copper,

thermal stability of carbonate: Na₂CO₃ > CaCO₃ > CuCO₃

BUT many carbonates do not fit into such a sequence, so you can't regards this as a general rule

See notes on the Limestone: Detailed section on the thermal decompositions of carbonates page

and advanced level chemistry notes on the thermal decomposition of carbonates

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OTHER ASSOCIATED PAGE LINKS

SEE ALSO 2. RUSTING & Introducing REDOX reactions

and 3. Metal Reactivity Series Experiments-Observations

and GCSE/IGCSE m/c QUIZZES on metal reactivity

Foundation-tier Level (easier) multiple choice quiz on the Reactivity Series of Metals

or Higher-tier Level (harder) multiple choice quiz on the Reactivity Series of Metals

and GCSE/IGCSE reactivity gap-fill worksheet or Rusting word-fill worksheet

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