Oxides and Oxoanions of Transition Elements

Full Notes

Transition metals form a variety of oxides and oxoanions by reacting with oxygen, particularly at high temperatures. Their oxidation states, chemical nature, and solubility change based on the group and metal. These oxides play a critical role in redox chemistry and industrial processes.

Formation and Nature of Oxides

• All d-block metals except scandium form monoxides (MO) which are predominantly ionic.
• The highest oxidation number matches the group number:
  • Sc₂O₃ (Sc³⁺)
  • Mn₂O₇ (Mn⁷⁺)
• Beyond Mn (Group 7), no higher oxides than Fe₂O₃ are observed.

Ionic to Covalent Nature

As oxidation number increases, ionic character decreases and covalent character increases.

For Example: Mn₂O₇ is a covalent green oil and highly acidic whereas CrO₃ and V₂O₅ have low melting points and are acidic oxides.

Acidic, Amphoteric, and Basic Nature

• Acidic: Mn₂O₇, CrO₃, V₂O₅
  • Mn₂O₇ forms HMnO₄ (permanganic acid).
  • CrO₃ forms H₂CrO₄, H₂Cr₂O₇.
• Amphoteric: V₂O₅, Cr₂O₃
  • V₂O₅ reacts with acids to form VO₂⁺ and with bases to form VO₃⁻.
  • Cr₂O₃ dissolves in both acids and bases.
• Basic: CrO

Oxocations

Stabilised in aqueous medium: VO⁺ (V⁵⁺), VO₂⁺ (V⁴⁺), TiO₂⁺ (Ti⁴⁺).

Chromium Compounds

Preparation of Chromates and Dichromates

From chromite ore (FeCr₂O₄) by fusion with Na₂CO₃ in air:

4FeCr₂O₄ + 8Na₂CO₃ + 7O₂ → 8Na₂CrO₄ + 2Fe₂O₃ + 8CO₂

Acidify sodium chromate to get sodium dichromate:

2Na₂CrO₄ + 2H⁺ → Na₂Cr₂O₇ + 2Na⁺ + H₂O

Convert sodium dichromate to K₂Cr₂O₇ (orange crystals):

Na₂Cr₂O₇ + 2KCl → K₂Cr₂O₇ + 2NaCl

Equilibrium Between Chromate and Dichromate

• Acidic: 2CrO₄²⁻ + 2H⁺ ⇌ Cr₂O₇²⁻ + H₂O
• Basic: Cr₂O₇²⁻ + 2OH⁻ ⇌ 2CrO₄²⁻ + H₂O

Structure

CrO₄²⁻: Tetrahedral

Cr₂O₇²⁻: Two tetrahedra sharing one corner; Cr–O–Cr bond angle = 126°

Oxidising Properties of K₂Cr₂O₇

Reaction in acid:

Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O   E° = +1.33 V

Oxidation Examples:

• 6Fe²⁺ → 6Fe³⁺ + 6e⁻
• 3Sn²⁺ → 3Sn⁴⁺ + 6e⁻
• 3H₂S → 3S + 6H⁺ + 6e⁻
• 6I⁻ → 3I₂ + 6e⁻

Example: Combined ionic reaction between Cr₂O₇²⁻ and Fe²⁺:

Cr₂O₇²⁻ + 14H⁺ + 6Fe²⁺ → 2Cr³⁺ + 6Fe³⁺ + 7H₂O

Manganese Compounds

Preparation of Potassium Permanganate

• Fusion of MnO₂ with KOH and O₂: 2MnO₂ + 4KOH + O₂ → 2K₂MnO₄ + 2H₂O
• Disproportionation in acidic solution: 3MnO₄²⁻ + 4H⁺ → 2MnO₄⁻ + MnO₂ + 2H₂O
• Electrolytic oxidation (following fusion of MnO₂ with KOH and O₂): MnO₄²⁻ → MnO₄⁻
• Lab preparation: 2Mn²⁺ + 5S₂O₈²⁻ + 8H₂O → 2MnO₄⁻ + 10SO₄²⁻ + 16H⁺

Properties of KMnO₄

• Purple/black crystals
• Not very soluble (6.4 g/100 g at 293 K)
• Decomposes on heating: 2KMnO₄ → K₂MnO₄ + MnO₂ + O₂
• Molecular Orbital Theory explains intense colour and weak paramagnetism.
• Manganate and permanganate ions are tetrahedral, with π bonding occurring between overlap of p-orbitals of oxygen and d-orbitals of manganese.

Manganate, MnO₄²⁻ – green and paramagnetic (one unpaired electron).

Permanganate, MnO₄⁻ – purple and diamagnetic (no unpaired electrons).

Oxidising Properties of KMnO₄

KMnO₄ is a powerful oxidising agent, readily gaining electrons and oxidising another species.

In Acidic Medium:

MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O   E° = +1.52 V

Hydrogen ion concentration ([H⁺]) plays a crucial role in determining which product is formed when KMnO₄ reacts (MnO₄²⁻, MnO₂, or Mn²⁺).

Half-Reactions and Corresponding Standard Electrode Potentials:

• MnO₄⁻ + e⁻ → MnO₄²⁻   E° = +0.56 V   (mildly basic)
• MnO₄⁻ + 4H⁺ + 3e⁻ → MnO₂ + 2H₂O   E° = +1.69 V   (weakly acidic)
• MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O   E° = +1.52 V   (strongly acidic)

Example reactions (acidic):

• 10I⁻ + 2MnO₄⁻ + 16H⁺ → 2Mn²⁺ + 5I₂ + 8H₂O
• 5Fe²⁺ + MnO₄⁻ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O
• 5C₂O₄²⁻ + 2MnO₄⁻ + 16H⁺ → 2Mn²⁺ + 10CO₂ + 8H₂O
• 5S²⁻ + 2MnO₄⁻ + 16H⁺ → 2Mn²⁺ + 5S + 8H₂O
• 5SO₃²⁻ + 2MnO₄⁻ + 6H⁺ → 2Mn²⁺ + 5SO₄²⁻ + 3H₂O
• 5NO₂⁻ + 2MnO₄⁻ + 6H⁺ → 5NO₃⁻ + 2Mn²⁺ + 3H₂O

In Neutral/Alkaline Medium:

• I⁻ + MnO₄⁻ + H₂O → IO₃⁻ + MnO₂ + OH⁻
• 3S₂O₃²⁻ + 8MnO₄⁻ + H₂O → 6SO₄²⁻ + 8MnO₂ + 2OH⁻
• 3Mn²⁺ + 2MnO₄⁻ + 2H₂O → 5MnO₂ + 4H⁺

Note: KMnO₄ should not be used in HCl, as Cl⁻ gets oxidised to Cl₂.

Uses of KMnO₄

• Strong oxidising agent in organic synthesis
• Used for bleaching and decolourisation
• Acts as a volumetric titrant in redox titrations

Summary

• Lower oxides are more ionic and basic and higher oxides become covalent and acidic.
• Chromate–dichromate interconvert with pH and both species show strong oxidising behaviour in acid.
• KMnO₄ is a powerful oxidant with products that depend on medium and acidity.
• MnO₄⁻ is purple and diamagnetic and MnO₄²⁻ is green and paramagnetic.