Variable Oxidation States

Full Notes: Variable Oxidation States

Why Do Transition Metals Show Variable Oxidation States?

Transition metals have a partially filled d-subshell.

The energy difference (ΔE) between successive oxidation states is small. This allows transition metals to lose different numbers of electrons and form multiple oxidation states that can be interchanged relatively easily.

Reduction of Vanadium (V) Ions in Acidic Solution

Vanadium (V) exists as VO₂⁺ (vanadate(V) ion) in acidic solution.

Zinc (Zn) in acidic solution (H⁺) reduces vanadate(V) stepwise to vanadium(IV), vanadium(III), and vanadium(II) compounds. The species formed have different colours.

Reduction reactions with Zn and HCl:

VO₂⁺ (yellow) → VO²⁺ (blue)
VO₂⁺ + 2H⁺ + e⁻ → VO²⁺ + H₂O

VO²⁺ (blue) → V³⁺ (green)
VO²⁺ + 2H⁺ + e⁻ → V³⁺ + H₂O

V³⁺ (green) → V²⁺ (purple)
V³⁺ + e⁻ → V²⁺

Vanadium(II) is the most reduced form and is the strongest reducing agent.

Factors Affecting Redox Potential of Transition Metals

See electrode potentials for more information.

The redox potential (E°) of a transition metal depends on:

• pH:
  In acidic conditions, oxidation is more likely.
  In alkaline conditions, reduction is more likely.
• Ligands bonded to the metal ion:
  

  Different ligands affect the stability of oxidation states.

  For Example
  [Fe(H₂O)₆]³⁺ has a different redox potential than [Fe(CN)₆]³⁻ due to ligand effects.

Tollens’ Reagent and the Silver Mirror Test

Tollens’ reagent is [Ag(NH₃)₂]⁺ and it is used to distinguish between aldehydes and ketones (see here for more).

Aldehydes reduce Ag⁺ to metallic silver (silver mirror forms).

Equation:
[Ag(NH₃)₂]⁺ + e⁻ → Ag (s) + 2NH₃

Ketones do not react (no silver mirror).

Redox Titrations: Fe²⁺ and C₂O₄²⁻ with MnO₄⁻

Manganate(VII) (MnO₄⁻) is a strong oxidising agent often used in redox titrations.

The titrations are considered self-indicating because the solution changes colour from purple to colourless when the manganese has been reduced (to Mn²⁺).

Matt’s exam tip

Make sure you understand the half equations behind each of the redox examples given below and how the overall equation is constructed. It’s really important you can recognise the species being reduced and oxidised rather than just remember the equations.

Reaction of MnO₄⁻ with Fe²⁺ in acidic solution

Equation:
MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O

Half-equations
Reduction: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O
Oxidation: Fe²⁺ → Fe³⁺ + e⁻

Colour change:
Purple MnO₄⁻ → Colourless Mn²⁺

Reaction of MnO₄⁻ with C₂O₄²⁻ (ethanedioate ion)

Equation:
2MnO₄⁻ + 16H⁺ + 5C₂O₄²⁻ → 2Mn²⁺ + 10CO₂ + 8H₂O

Half-equations
Reduction: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O
Oxidation: C₂O₄²⁻ → CO₂ + 2e⁻

Colour change:
Purple MnO₄⁻ → Colourless Mn²⁺
Effervescence due to CO₂ gas.

Summary Table: Variable Oxidation States