5.6.6 Redox Reactions of Transistion Elements

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

Reduction of Cr₂O₇^2- to Cr³⁺

• Aqueous dichromate (VI) ions, Cr₂O₇^2-, have an orange colour whilst aqueous chromium (III) ions, Cr³⁺(aq), have a green colour
• Acidified Cr₂O₇^2- 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  Cr₂O₇^2- to Cr³⁺ and Cr³⁺ to Cr²⁺
• The equation for this reaction is:

   Cr₂O₇^2- (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

Oxidation of Cr³⁺ to CrO₄^2-

• 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), CrO₄^2-
• The equation for this reaction is

3H₂O₂ (aq) + 2Cr³⁺ (aq) + 10OH^- (aq) → 2CrO₄^2- (aq) + 8H₂O (l)

                                         dark green                                 yellow     

• • Chromium is oxidised from +3 in Cr³⁺ to +6 in CrO₄^2-
  • Oxygen is reduced from -1 in H₂O₂ to -2 in CrO₄^2-

Oxidation of  CrO₄^2- to Cr₂O₇^2-

• Dilute sulfuric acid can be added to chromate (VI), CrO₄^2- (aq), solution to produce a dichromate (VI), Cr₂O₇^2- (aq), solution
• The equation for this reaction is

2CrO₄^2- (aq) + 2H⁺ (aq) →  Cr₂O₇^2- (aq) + H₂O (l)

                                  yellow                                      orange

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) + I₂ (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) + H₂SO₄ (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