Redox Reactions Involving Transition Metal Ions

This lesson covers: 

1. How transition metals can exist in various oxidation states
2. Using redox potentials to determine the ease of reduction
3. Redox reactions of transition metals

Transition metals have variable oxidation states

Transition metals can exist in various positive oxidation states in compounds and complexes. This occurs due to the similar energies of the outer d orbital electrons, allowing multiple oxidation states to be stable.

For example, vanadium exhibits four oxidation states from +2 to +5, each with different coloured ions:

Redox potentials indicate ease of reduction

The redox potential, E^⦵, indicates how easily an ion is reduced. Ions with higher, more positive E^⦵ values are more easily reduced to lower oxidation states.

For example, Cu²⁺ has a higher redox potential than Cr³⁺ so is more easily reduced.

However, E^⦵ values are measured under standard conditions, so the actual redox potential can vary depending on:

• Ligands - E° assumes metal ions are surrounded by water ligands in aqueous solution. Other ligands that form stronger bonds with the metal ion can raise or lower the potential by stabilising a particular oxidation state.
• pH - Acidic conditions provide excess H⁺ ions needed for reduction of some metal ions. For example: 

VO₂⁺_(aq) + 2H⁺_(aq) + e^- ➔ VO²⁺_(aq) + H₂O_(l)

Alkaline conditions favour oxidation reactions that consume OH^- ions instead. For example:

Cr(OH)_3(s) + 5OH^-_(aq) ➔ CrO₄^2-_(aq) + 4H₂O_(l) + 3e^-

In general, more acidic solutions give higher, more positive redox potentials.

Interconversion between common transition ions

Transition elements can change oxidation state by gaining or losing electrons in redox reactions.

For complex ions in solution, this change in oxidation state is often accompanied by a distinctive colour change.

You need to know about the redox reactions of iron, chromium, and copper.

Interconversion between Fe²⁺ and Fe³⁺

Pale-green Fe²⁺_(aq) ions are oxidised to yellow Fe³⁺_(aq) ions by acidified potassium manganate(VII) solution, KMnO_4(aq), in acidic conditions:

MnO₄^- + 8H⁺ + 5Fe²⁺ ➔ Mn²⁺ + 4H₂O + 5Fe³⁺

• Iron is oxidised from +2 to +3.
• Manganese is reduced from +7 to +2.

Fe³⁺ is reduced back to Fe²⁺ by iodide ions (I^-)_(aq):

2I^- + 2Fe³⁺ ➔ 2Fe²⁺ + I₂

• Iodine is oxidised from -1 to 0.
• Iron is reduced from +3 to +2.

Interconversion between Cr³⁺ and Cr₂O₇^2-

Dark-green Cr³⁺_(aq) ions are oxidised to yellow CrO₄^2-_(aq) ions by warming with hydrogen peroxide solution, H₂O_2(aq), in alkaline conditions:

3H₂O₂ + 2[Cr(OH)₆]^3- ➔ 2OH^- + 2CrO₄^2- + 8H₂O

• Chromium is oxidised from +3 to +6.
• Oxygen is reduced from -1 to -2.

Orange Cr₂O₇^2- ions are formed by adding sulfuric acid:

2CrO₄^2- + 2H⁺ ➔ Cr₂O₇^2- + H₂O

Cr₂O₇^2- is reduced back to Cr³⁺ by zinc in acidic conditions:

Cr₂O₇^2- + 14H⁺ + 3Zn ➔ 2Cr³⁺ + 7H₂O + 3Zn²⁺

• Zinc is oxidised from 0 to +2.
• Chromium is reduced from +6 to +3.

Reduction of Cu²⁺ and disproportionation of Cu⁺

Pale blue Cu²⁺_(aq) ions are reduced to the white precipitate copper(I) iodide, CuI, by iodide ions (I^-).

The dissolved iodine turns the solution brown:

2Cu²⁺_(aq) + 4I^-_(aq) ➔ 2CuI_(s) + I_2(aq)

• Iodide is oxidised from -1 to 0.
• Copper is reduced from +2 to +1.

The colourless Cu⁺ ion is unstable in aqueous solution and readily disproportionates into Cu²⁺_(aq) and Cu_(s):

2Cu⁺_(aq) ➔ Cu²⁺_(aq) + Cu_(s)

• Cu⁺ is both oxidised (from +1 to +2) and reduced (from +1 to 0).