5.3.2 Chromium Chemistry

• For chromium we need to consider the following standard electrode potential values

• The half equations are arranged from the most negative E^θ at the top to the most positive E^θ at the bottom
  • The strongest oxidising agent is the reactant from the electrode reaction with the most positive standard electrode potential
    • In this case, the most positive value is + 1.33 V, which means that Cr₂O₇^2– (aq)  is the strongest oxidising agent
  • The strongest reducing agent is the product from the electrode reaction with the most negative standard electrode potential
    • In this case, the most negative value is – 0.76 V, which means that Zn (s) is the strongest reducing agent 

Oxidation of Cr from +3 to +6 by H₂O₂ 

• The two half equations we need to consider are 3 and 4
• Chromium is being oxidised from an oxidation number of +6 to +3 in half equation 3
• So, the two half equations we need to consider are:
  • 3. CrO₄^2– (aq) + 4H₂O (l) + 3e^–  Cr(OH)₃ (aq) + 5OH^– (aq)     E^θ = – 0.13 V
  • 4. H₂O₂ (aq) + 2e^–  2OH^– (aq)     E^θ = + 1.24 V

• Half equation 4 has the most positive standard electrode potential, E^θ, value
  • Therefore, this is the reduction reaction 
  • H₂O₂ (aq) + 2e^–  2OH^– (aq)
• Half equation 3 has the least positive standard electrode potential, E^θ, value
  • Therefore, this is the oxidation reaction and needs to be reversed
  • Cr(OH)₃ (aq) + 5OH^– (aq) CrO₄^2– (aq) + 4H₂O (l) + 3e
• We can obtain the overall equation by combining the reduction equation (half equation 4) and the oxidation equation (the reversed half equation 3)
  • Remember: When combining half equations, they must have the same number of electrons in both equations
    

    3H2O2 (aq) + 6e–		6OH– (aq)
    2Cr(OH)3 (aq) + 10OH– (aq)		2CrO42– (aq) + 8H2O (l) + 6e–

  •  The equal number of electrons on both sides of the equation cancel out and 6OH^– (aq) cancel out on both sides to give the overall equation: 

 2Cr(OH)₃ (aq) + 4OH^- (aq) + H₂O₂ (aq)  CrO₄^2- (aq) + 8H₂O (l)

• This reaction is carried out in alkaline conditions due to the presence of OH^- ions in the equation

Reduction of Cr from +6 to +3 by Zn

• Chromium is being reduced from an oxidation number of +6 to +3 in half equation 5
• So, the two half equations we need to consider are:
  • 1. Zn²⁺ (aq) + 2e^– Zn (s)     E^θ = – 0.76 V
  • 5. Cr₂O₇^2– (aq) + 14H⁺ (aq) + 6e^–  2Cr³⁺ (aq) + 7H₂O (l)     E^θ = + 1.33 V
• Half equation 5 has the most positive standard electrode potential, E^θ, value
  • Therefore, this is the reduction reaction 
  • Cr₂O₇^2– (aq) + 14H⁺ (aq) + 6e^–  2Cr³⁺ (aq) + 7H₂O (l)
• Half equation 1 has the least positive / most negative standard electrode potential, E^θ, value
  • Therefore, this is the oxidation reaction and needs to be reversed
  • Zn (s) → Zn²⁺ (aq) + 2e^– 
• We can obtain the overall equation by combining the reduction equation (half equation 3) and the oxidation equation (the reversed half equation 1)
  • Remember: When combining half equations, they must have the same number of electrons in both equations
    

    Cr2O72– (aq) + 14H+ (aq) + 6e–		2Cr3+ (aq) + 7H2O (l)
    3Zn (s)		3Zn2+ (aq) + 6e–

  •  The equal number of electrons on both sides of the equation cancel out to give the overall equation: 

 Cr₂O₇^2- (aq) + 14H⁺ (aq) + 3Zn (s) 2Cr³⁺ (aq) + 7H₂O (l) + 3Zn²⁺ (aq)

• This reaction is carried out under acidic conditions due to the presence of H⁺ in the equation

Reduction of Cr from +3 to +2 by Zn

• Chromium is being reduced from an oxidation number of +3 to +2 in half equation 2
• So, the two half equations we need to consider are:
  • 1. Zn²⁺ (aq) + 2e^–  Zn (s)     E^θ = – 0.76 V
  • 2. Cr³⁺ (aq) + e^– Cr²⁺ (aq)     E^θ = – 0.41 V
• Half equation 2 has the most positive / least negative standard electrode potential, E^θ, value
  • Therefore, this is the reduction reaction 
  • Cr³⁺ (aq) + e^–  Cr²⁺ (aq)
• Half equation 2 has the least positive standard electrode potential, E^θ, value
  • Therefore, this is the oxidation reaction and needs to be reversed
  • Zn (s)  Zn²⁺ (aq) + 2e^– 
• We can obtain the overall equation by combining the reduction equation (half equation 2) and the oxidation equation (the reversed half equation 1)
  • Remember: When combining half equations, they must have the same number of electrons in both equations
    

    2Cr3+ (aq) + 2e–		2Cr2+ (aq)
    Zn (s)		Zn2+ (aq) + 2e–

  •  The equal number of electrons on both sides of the equation cancel out to give the overall equation: 

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

• As this reaction is a further step from the previous reduction this reaction is also carried out under acidic conditions

Colour changes of chromium 

• Another approach to transition metal chemistry is to look at the reactions of transition metal ions and complexes with various reagents
• This information is sometimes presented in the form of a reaction scheme:

Reaction scheme outlining the colour changes associated with the reactions of different chromium species 

 

• You would be expected to know:
  • The formula and corresponding state, (aq) or (s), of the transition metal species
  • The colour of the transition metal species
  • The reagents and conditions required to convert one transition metal species into another
  • How to write the equation for the conversion of one transition metal species into another
• For example:
  • A green solution of hexaaquachromium(III) ions, [Cr(H₂O)₆]³⁺ (aq), will react with dilute NaOH (aq) to form a grey-green precipitate of Cr(H₂O)₃(OH)₃ (s)
  • [Cr(H₂O)₆]³⁺ (aq) + 3OH^- (aq) → [Cr(H₂O)₃(OH)₃] (s) + 3H₂O (l) 

• The chromate CrO₄^2- and dichromate Cr₂O₇^2- ions can be converted from one to the other by the following equilibrium reaction

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

• Chromate(IV) ions are stable in alkaline solution, but in acidic conditions the dichromate(VI) ion is more stable
• Addition of acid will  push the equilibrium to the dichromate
  • This results in a colour change from yellow to orange
• Addition of alkali will remove the H⁺ ions and push the equilibrium to the chromate
• This is not a redox reaction as both the chromate and dichromate ions have an oxidation number of +6
  • This is an acid base reaction