Redox Reactions of Transistion Elements (OCR A Level Chemistry)

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Iron(II) & Iron(III) Ions

Oxidation of Fe2+ to Fe3+

  • A redox titration involves an oxidising agent being titrated against a reducing agent
  • Electrons are transferred from one species to another
  • In acid-base titrations indicators are used to show the endpoint of a reaction; however redox titrations using transition metal ions naturally change colour when changing oxidation state, so indicators are not always necessary
  • They are said to be 'self-indicating'
  • The redox reaction between iron(II) ions and manganate(VII) ions in acidic conditions is used as a basis for a redox titration
  • In this reaction:
    • Fe2+ is oxidised to Fe3+
    • MnO4- is reduced to Mn2+
  • Potassium manganate(VII) is commonly used which is an oxidising agent and is a deep purple colour
  • In acidic solutions it is reduced to the almost colourless manganese(II) ion by the Fe2+(aq) 
  • The equation for the reaction is:

   MnO4- (aq) + 8H(aq)  + 5Fe2+   →   Mn2+ (aq) + 5Fe3+  + 4H2O (aq)

                      purple                                                    colourless

Reduction of Fe3+ to Fe2+

  • An orange-brown solution of Fe3+(aq) ions can be reduced to pale green Fe2+(aq) ions by various reducing agents.
  • A potassium iodide solution is commonly used
  • The colour change can be masked by the formation of iodine which has a brown colour
  • In this reaction
    • Fe3+ is reduced to Fe2+
    • I- is oxidised to I2
  • The equation for this reaction is:

   2Fe3+ (aq) + 2I- (aq)     →   2Fe2+  + I2 (aq)

      orange-brown                    pale-green   brown

Chromium(III) & Dichromate Ions

Reduction of Cr2O72- to Cr3+

  • Aqueous dichromate (VI) ions, Cr2O72-, have an orange colour whilst aqueous chromium (III) ions, Cr3+(aq), have a green colour
  • Acidified Cr2O72- ions can be reduced to Cr3+(aq) ions by the addition of zinc
  • Zinc is a strong reducing agent and is capable of reducing both  Cr2O72- to Cr3+ and Cr3+ to Cr2+
  • The equation for this reaction is:

   Cr2O72- (aq) + 14H(aq) + 3Zn (s) → 2Cr3+ (aq) + 7H2O (l) + 3Zn2+ (aq)

                orange                                                           green

  • With an excess of zinc, chromium(III) ions are reduced further to chromium(II), which is a pale blue colour

   Zn (s) + 2Cr3+ (aq)  → Zn2+ (aq) + 2Cr2+ (aq)

                                                 green                                pale blue

Examiner Tip

Fe2+(aq) is a weaker reducing agent than zinc and will only reduce the dichromate to Cr3+

Standard electrode potentials can be used to compare the strength of reducing agents and you should be able to use E to explain why certain redox reactions take place

Oxidation of Cr3+ to CrO42-

  • When transition metals in low oxidation states are in an alkaline solution, they are more easily oxidised than when in acidic solution
  • Hot alkaline hydrogen peroxide, H2O2, is a powerful oxidising agent which can be used to oxidise chromium(III) to chromium(VI), CrO42-
  • The equation for this reaction is

3H2O2 (aq) + 2Cr3+ (aq) + 10OH- (aq) → 2CrO42- (aq) + 8H2O (l)

                                         dark green                                 yellow     

    • Chromium is oxidised from +3 in Cr3+ to +6 in CrO42-
    • Oxygen is reduced from -1 in H2O2 to -2 in CrO42-

Oxidation of  CrO42- to Cr2O72-

  • Dilute sulfuric acid can be added to chromate (VI), CrO42- (aq), solution to produce a dichromate (VI), Cr2O72- (aq), solution
  • The equation for this reaction is

2CrO42- (aq) + 2H+ (aq) →  Cr2O72- (aq) + H2O (l)

                                  yellow                                      orange

Examiner Tip

Don't worry - you won't be expected to memorise and reproduce these full equations in exams. 

You could be asked to construct and interpret redox equations from half equations and oxidation numbers though - so make sure you understand how to do it.

Reduction & Disproportionation of Copper Ions

Reduction of Cu2+ to Cu+

  • A pale blue solution of Cu2+ can be reduced to Cu+ by various reducing agents
  • A potassium iodide solution is commonly used
  • When excess iodide ions are present the following reaction occurs:

2Cu2+ (aq) + 4I(aq) → 2CuI (s) + I(aq)

pale blue            white precipitate   brown

    • I- is oxidised to brown iodine, I2
    • Cu2+ is reduced to Cu+, forming a white precipitate

Disproportionation of copper(I) ions

  • When solid copper(I) oxide, Cu2O, reacts with hot dilute sulfuric acid, a brown precipitate of copper is formed together with a blue solution of copper(II) sulfate
  • In this reaction copper(I) ions, Cu+, have been simultaneously oxidised and reduced
  • As the same element has been reduced and oxidised, this reaction is disproportionation

Cu2O (s) + H2SO(aq) → Cu (s) + CuSO4 (aq)+ H2O (l)   

    • Copper has been reduced from +1 in Cu2O to O in Cu
    • Copper has been oxidised from +1 in Cu2O to +2 in CuSO4

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Sonny

Author: Sonny

Expertise: Chemistry

Sonny graduated from Imperial College London with a first-class degree in Biomedical Engineering. Turning from engineering to education, he has now been a science tutor working in the UK for several years. Sonny enjoys sharing his passion for science and producing engaging educational materials that help students reach their goals.