Reactions of Transition Metal Elements

Full Notes

Colours and Redox of Vanadium Ions

Vanadium exists in multiple oxidation states with characteristic colours:

• VO₂⁺ (+5) – yellow
• VO²⁺ (+4) – blue
• V³⁺ (+3) – green
• V²⁺ (+2) – purple

Vanadium Reduction with Zinc in Acid

You can reduce vanadium stepwise using zinc in acidic conditions:

1. VO₂⁺ + 2H⁺ + e⁻ → VO²⁺ + H₂O
2. VO²⁺ + 2H⁺ + e⁻ → V³⁺ + H₂O
3. V³⁺ + e⁻ → V²⁺

Each step involves electron transfer and results in distinct colour changes:

VO₂⁺ (Yellow) → VO²⁺ (Blue) → V³⁺ (Green) → V²⁺ (Violet)

Explaining with E° Values (E_cell)

We can explain these reductions in terms of standard electrode potentials:

• VO₂⁺ + 2H⁺ + e⁻ → VO²⁺ + H₂O  E° = +1.00 V
• VO²⁺ + 2H⁺ + e⁻ → V³⁺ + H₂O  E° = +0.34 V
• V³⁺ + e⁻ → V²⁺  E° = −0.26 V
• Zn²⁺ + 2e⁻ → Zn  E° = −0.76 V
• V²⁺ + 2e⁻ → V(s) E° = −1.18 V

Since zinc's E° is more negative than all the vanadium half-cells, it can act as a reducing agent for all three steps shown previously.

However, V²⁺ cannot be reduced further by zinc, as zinc (E° = −0.76 V) does not have a low enough potential to drive this final reduction. Therefore, reaction stops at V²⁺.

Chromium Redox Chemistry

These are the reactions and chemistry of chromium (Cr) that you need to know.

Reduction using zinc in acidic conditions

Cr₂O₇²⁻ can be reduced to Cr³⁺ and Cr²⁺ using zinc in acidic conditions

Dichromate(VI) to Chromium(III)
Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O

Chromium(III) to Chromium(II)
With excess zinc, Cr³⁺ can be further reduced:
Cr³⁺ + e⁻ → Cr²⁺ (blue)

Explanation

• Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O E° = +1.33 V (Orange → Green)
• Cr³⁺ + e⁻ → Cr²⁺ E° = −0.41 V (Green → Blue)
• Zn²⁺ + 2e⁻ → Zn  E° = −0.76 V

Zinc has a more negative E° than both half-cells, so it can reduce:

• Cr₂O₇²⁻ to Cr³⁺
• And further: Cr³⁺ to Cr²⁺

Zinc stops at Cr²⁺, as further reduction (to Cr metal, E° = −0.74 V) is not feasible due to the similar potential to Zinc (−0.76 V).

Oxidation using H₂O₂

Cr³⁺ can be oxidised to Cr₂O₇²⁻ using H₂O₂ in alkaline conditions.

CrO4- is formed before then being acidified.

2Cr³⁺ + 10OH⁻ + 3H₂O₂ → 2CrO₄²⁻ + 8H₂O

On acidification:
2CrO₄²⁻ + 2H⁺ → Cr₂O₇²⁻ + H₂O (Orange)

So overall:
Cr³⁺ is oxidised to Cr₂O₇²⁻ via CrO₄²⁻, using H₂O₂ in alkali, followed by acidification.

Hydroxide Precipitation and Ligand Reactions

Transition metal ions form precipitates with OH⁻ and NH₃:

Amphoteric behaviour: Cr(OH)₃ dissolves in both acid and excess base.

Ligand Exchange and Stability

Ligands in a complex ion can sometimes be substituted for different ligands in what are called ligand substitution or ligand exchange reactions.

Water NH₃ and H₂O are similar in size and uncharged, and usually six molecules of each can fit around a central ion in complex, getting close enough to form co-ordinate bonds to it. Giving the complex a co-ordination number of 6 and an octahedral shape.

However, chloride ions, Cl⁻ are larger than H₂O and NH₃, and only four Cl⁻ ligands can fit around a central ion, giving the complex a co-ordination number of 4 and a tetrahedral shape.

Substitution of H₂O by NH₃:

Only six H₂O or NH₃ ligands can get close enough to the metal ion to co-ordinately bond to it.

As a result, if ligand substitution occurs and H₂O ligands are substituted or exchanged for NH₃ ligands, there is no change in co-ordination number (6 → 6).

Example [Cu(H₂O)₆]²⁺ with ammonia

[Cu(H₂O)₆]²⁺ + 4NH₃ ⇌ [Cu(NH₃)₄(H₂O)₂]²⁺ + 4H₂O

• Colour change: Blue → Deep Blue.
• Substitution is incomplete (only 4 NH₃ molecules replace H₂O).

Substitution of H₂O by Cl⁻:

As Cl⁻ ligands are larger than H₂O, the co-ordination number decreases (6 → 4) if ligand substitution occurs.

Ligand substitution examples you need to know

[Cu(H₂O)₆]²⁺ + 4NH₃ ⇌ [Cu(NH₃)₄(H₂O)₂]²⁺ + 4H₂O

• Colour change: Blue → Deep Blue.

[Cu(H₂O)₆]²⁺ + 4Cl⁻ ⇌ [CuCl₄]²⁻ + 6H₂O

• Colour change: Blue → Yellow.

[Co(H₂O)₆]²⁺ + 4Cl⁻ ⇌ [CoCl₄]²⁻ + 6H₂O

• Colour change: Pink → Blue.

Chelation and the Chelate Effect

There is a tendency for bidentate or multidentate ligands to replace monodentate ligands in ligands.

This is driven by an increase in entropy and is called the Chelate Effect.

Example Reaction of [Cu(H₂O)₆]²⁺ with C₂O₄²⁻ ions

[Cu(H₂O)₆]²⁺ + 3C₂O₄²⁻ → [Cu(C₂O₄)₃]⁴⁻ + 6H₂O

• Entropy increases (+ΔS) as the number of free particles increases (4 reactant particles compared to 7 product particles).
• ΔG is more negative, making the reaction more feasible.

Catalysis by Transition Metals

Transition metals are commonly used as catalysts due their variable oxidation states. There are two types of catalyst - heterogeneous and homogeneous.

Heterogeneous Catalysts

Heterogeneous catalysts are in a different phase than the reactants.

Example Vanadium(V) oxide (V₂O₅) in the Contact Process

• Reaction: SO₂ + ½O₂ → SO₃ (used to make H₂SO₄).

Catalytic Cycle:

1. SO₂ is oxidised to SO₃ via vanadium(V) oxide.
2. V₂O₅ is reduced to V₂O₄.
   V₂O₅ + SO₂ → V₂O₄ + SO₃
3. V₂O₄ is re-oxidised by oxygen.
   V₂O₄ + ½O₂ → V₂O₅

Catalyst remains unchanged overall.

Catalytic converters:

Use platinum or rhodium to remove harmful gases from car exhaust fumes.

CO and NO adsorb to surface → bonds weaken → reaction occurs → CO₂ and N₂ desorb.

Homogeneous Catalysts

Homogeneous catalysts are in the same phase as the reactants.

• The reaction proceeds via an intermediate species.
• Transition metals are effective due to variable oxidation states.

Example Fe²⁺ Catalysing the Reaction Between I⁻ and S₂O₈²⁻

Reaction:
S₂O₈²⁻ + 2I⁻ → 2SO₄²⁻ + I₂

This reaction is slow because both reactants are negatively charged.

Fe²⁺ speeds up the reaction by forming an intermediate:

• Fe²⁺ reduces S₂O₈²⁻:
  S₂O₈²⁻ + 2Fe²⁺ → 2SO₄²⁻ + 2Fe³⁺
• Fe³⁺ oxidises I⁻ to I₂:
  2Fe³⁺ + 2I⁻ → 2Fe²⁺ + I₂

Fe²⁺ is regenerated, so it remains a catalyst.

Autocatalysis

Autocatalysis occurs when a reaction produces its own catalyst.

Example Mn²⁺ catalysing the reaction between C₂O₄²⁻ and MnO₄⁻

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

Without Mn²⁺, the reaction is slow.

As Mn²⁺ is produced, it catalyses the reaction by forming intermediates:

• Mn²⁺ reacts with MnO₄⁻ to form Mn³⁺:
  4Mn²⁺ + MnO₄⁻ + 8H⁺ → 5Mn³⁺ + 4H₂O
• Mn³⁺ reacts with C₂O₄²⁻, regenerating Mn²⁺:
  2Mn³⁺ + C₂O₄²⁻ → 2Mn²⁺ + 2CO₂

The reaction speeds up as more Mn²⁺ is produced.

Summary

• Vanadium and chromium show multiple oxidation states with characteristic colour changes that can be predicted using E° values.
• Chromate and dichromate interconvert with pH change and H₂O₂ oxidises Cr³⁺ in alkaline conditions.
• Ligand substitution causes colour and coordination number changes; chloride often gives tetrahedral complexes.
• Chelation is favoured by entropy which makes multidentate complexes especially stable.
• Transition metals catalyse reactions heterogeneously on surfaces and homogeneously via intermediates with notable cases like V₂O₅ and Fe²⁺.
• Autocatalysis occurs when a product such as Mn²⁺ accelerates its own formation.