Transition Metals - Examples of Redox Reactions in Transition Metals (A-Level Chemistry)

Examples of Redox Reactions in Transition Metals

Reduction of Vanadate (V) ions

Metallic zinc is a good reducing agent.

It reacts with dilute acid to form ZnĀ²āŗ ions and releases electrons for reduction.

Zn(s) ā†’ ZnĀ²āŗ(aq) + 2eā»

Addition of zinc to vanadium(V) ions from a solution of ammonium vanadate (V) (NHā‚„VOā‚ƒ) in acidic solution, will reduce the vanadium through each successive oxidation state. Each oxidation state produces a different colour.

Vanadium (V) ions are first reduced to vanadium (IV) ions:

2[VOā‚‚(Hā‚‚O)ā‚„]āŗ(aq) + Zn(s) + 4Hāŗ(aq) ā†’ 2[VO(Hā‚‚O)ā‚…]Ā²āŗ(aq) + ZnĀ²āŗ(aq)
Yellow                                                                Blue

Vanadium (IV) ions are then reduced to vanadium (III) ions:

2[VO(Hā‚‚O)ā‚…]Ā²āŗ(aq) + Zn(s) +4Hāŗ(aq) ā†’ 2[V(Hā‚‚O)ā‚†]Ā³āŗ+(aq) + ZnĀ²āŗ(aq)
Blue                                                                           Green

Vanadium (III) ions are finally reduced to vanadium (II) ions:

Until now, the reactions have been feasible because their redox potentials have been positive.

Examples of Redox in Transition Metals
Examples of Redox in Transition Metals

Vanadium (II) ions will however not be further reduced by zinc to vanadium metal because the redox potential for the reaction is negative, which means that the reaction is not feasible under standard conditions.

2[V(Hā‚‚O)ā‚†]Ā²āŗ (aq) + Zn(s) ā†’ V(s) + ZnĀ²āŗ(aq) + 6Hā‚‚O(l)                             Eā¦µ = -0.42V

Reduction of Silver Diammine [Ag(NHā‚ƒ)ā‚‚]āŗ

Silver diammine [Ag(NHā‚ƒ)ā‚‚]āŗ is used in Tollenā€™s reagent to distinguish between aldehydes and ketones.

Aldehydes are strong enough reducing agents to reduce the silver (I) to metallic silver.

2[Ag(NHā‚ƒ)ā‚‚]āŗ + CHā‚ƒCHO + Hā‚‚O ā†’ 2Ag(s) + 4NHā‚ƒ + CHā‚ƒCOOH + 2Hāŗ

Ketones are unable to do this and have no reaction.

Reduction and Oxidation of Chromium Ions

The large number of oxidation states that can be taken by chromium, allow its to part-take in a wide variety of redox reactions.

When added to an acidic solution of zinc metal, dichromate (VI) ions are reduced to chromium (III) ions (Eā¦µ = +2.09V):

Crā‚‚Oā‚‡Ā²ā»(aq) + 3Zn(s) + 14Hāŗ(aq) ā†’ 2CrĀ³āŗ(aq) + 3ZnĀ²āŗ(aq) + 7Hā‚‚O(l)
Orange                                                         Green

Zinc will further reduce CrĀ³āŗ to CrĀ²āŗ (Eā¦µ = +0.35V):

2CrĀ³āŗ(aq) + 3Zn(s) ā†’ 2CrĀ²āŗ(aq) + ZnĀ²āŗ(aq)
Green                               Yellow

When added to an alkaline solution of hydrogen peroxide, chromium (III) ions are oxidised to chromate (IV) ions (Eā¦µ = +1.08V):

2CrĀ³āŗ(aq) + 3Hā‚‚Oā‚‚(s) + 10OHā» (aq) ā†’ CrOā‚„Ā²ā»(aq) + 8Hā‚‚O(l)
Green                                                                Yellow

If we then add some dilute sulphuric acid to the solution, Crā‚‚Oā‚‡Ā²ā» ions form turning the solution orange.

CrOā‚„Ā²ā»(aq) + 2Hāŗ(aq) ā‡Œ Crā‚‚Oā‚‡Ā²ā»(aq) + Hā‚‚O(l)
Yellow                                  Orange

This reaction is reversible and an equilibrium between CrOā‚„Ā²ā» and Crā‚‚Oā‚‡Ā²ā» ions sets up.

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