C4 - Chemical Changes - Revision Cards in GCSE Science

Separating Metals from Metal Oxides

Common metals react with oxygen to form oxides - oxidation - gain of oxygen. Oxides often ores metals need to be extracted from.
Reaction that separates metal from its oxides - reduction - loss of oxygen.
Some metals can be extracted from ores chemically by reduction using carbon. Ore is reduced and carbon is oxidised.
Position of metal in reactivity series determines whether it can be extracted by carbon.
- Metals higher than carbon - extracted using electrolysis - expensive.
- Metals below carbon - extracted by reduction using carbon - iron oxide reduced in blast furnace to make iron. Carbon can only take oxygen away from metals less reactive than itself.
Some metals are so unreactive they are in earth as metal itself - gold mined as its elemental form.

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Redox reaction if electrons are transferred.
Loss of electrons - oxidation. Gain - reduction - OIL RIG - Oxidation Is Loss, Reduction Is Gain.
REDuction & OXidation happen at same time - hence term 'REDOX'
Displacement reactions - Redox. Involves one metal kicking out another from compound - MORE REACTIVE metal will displace LESS REACTIVE from its compound.
If you put reactive metal into solution of a dissolved metal compound, reactive metal will replace less reactive one in compound.
In displacement reaction, metal ion always gains electrons and is reduced. Metal atom always loses electrons and is oxidised.
In ionic equation, only particles that react and products they form are shown - Mg(s) + Zn'2+(aq) --> Mg'2+(aq) + Zn(s). Shows displacement of zinc ions by magnesium. Full equation of above reaction would be this if you started with zinc chloride - Mg(s) + ZnCl2(aq) --> MgCl2(aq) + Zn(s).
Ionic equation concentrates on substances oxidised and reduced.

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Means 'splitting up with electricity'. Electric current is passed through electrolyte (molten or dissolved ionic compound that conducts electricity). Ions move towards electrodes (solid that conducts electricity submered in electrolyte), where they react, and compund decomposes.
+ ions - cathode - gain electrons - reduced. -ve ions - anode - lose electrons - oxidised. As ions gain/lose electrons, they form uncharged element and are discharged from electrolyte
Lead Bromide - + metal ion reduced to element at cathode --> Pb2+ + 2e- --> Pb. -ve non metal ions oxidised to element at anode --> 2Br- --> Br2 + 2e-
Electrolysis used if metal is too reactive to be reduced by carbon - expensive - lots of enegry required to melt ore and produce required current.
Aluminium extracted from ore bauxite - containes aluminium oxide. Has very high melting point so mixed with cryolite to lower it.
+ Al3+ ions attracted to cathode - gain 3 electrons each and turn into aluminium atoms. -ve O2- ions attracted to anode - lose 2 elelectrons each. Neutral O atoms combine to form O2 molecules. 2Al2O3 --> 4Al + 3O2
Anode made of carbon - needs regular replacement - reacts with O to produce CO2

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In aqueous solutions, there'll also be H+ ions and OH- ions from water - H2O H+ + OH-
Cathode - if H+ ions and metal ions are present, H2 gas will be produced if metal ions form elemental metal more reactive than hydrogen. If elemental metal is less reactive than H2, solid layer of pure metal will be produced. Anode - OH- ions and halide ions present, molecules of halide will be formed. No halide ions - OH- ions discharged and O2 formed.
CuSO4 - Cu2+, SO42-, H+ and OH-. Copper less reactive than hydrogen. Cathode - Cu produced - Cu2+ + 2e- --> Cu. No halide ions. Anode - O2 & H2O produced - 4OH- --> O2 + 2H2O + 4e-
NaCl - Na+, Cl-, OH-, H+. Sodium more reactive than H2. Cathode - H2 - 2H + 2e- --> H2. Chlorine ions present - Anode - Cl - 2Cl- --> Cl2 + 2e-
Once experiment is finished you can test any gaseous products to work out what's been produced.
- Chlorine bleaches damp litmus paper white
- Hydrogen - squeaky pop with lighted splint
- Oxygen relights glowing splint.

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