Variable oxidation states and redox titrationsAQA A-Level Chemistry: Revision notes

Transition elements show variable oxidation states because the 4s and 3d electrons are close in energy, so different numbers of electrons can be lost. Iron forms Fe²⁺ and Fe³⁺, copper forms Cu⁺ and Cu²⁺, manganese forms +2, +4 and +7, and vanadium forms +2 to +5.

A change in oxidation state is often accompanied by a colour change, which makes redox changes easy to follow.

Vanadate(V) ions in acidic solution (VO₂⁺, yellow) are reduced by zinc through the oxidation states IV, III and II:

• +5 VO₂⁺, yellow
• +4 VO²⁺, blue
• +3 V³⁺, green
• +2 V²⁺, violet

Each step is a reduction, for example VO₂⁺ + 2H⁺ + e⁻ → VO²⁺ + H₂O and V³⁺ + e⁻ → V²⁺. Zinc (E° = −0.76 V) reduces V³⁺ to V²⁺ (−0.26 V) but cannot reduce V²⁺ to the metal (−1.18 V): the reaction has a positive E°cell only if the metal's potential is more negative than the half-cell being reduced.

The redox potential for a transition metal ion going from a higher to a lower oxidation state is influenced by pH and by the ligand.

pH: if H⁺ appears in the half-equation, such as MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O, a higher pH (lower [H⁺]) makes E less positive, so manganate(VII) is a weaker oxidising agent.

Ligand: E° for Co³⁺/Co²⁺ is +1.82 V with water ligands but +0.11 V with ammonia ligands. Ammonia stabilises Co(III), so [Co(NH₃)₆]²⁺ is oxidised by air but [Co(H₂O)₆]²⁺ is not.

Tollens' reagent contains the complex ion [Ag(NH₃)₂]⁺. It is reduced to metallic silver, [Ag(NH₃)₂]⁺ + e⁻ → Ag + 2NH₃, which forms a silver mirror.

Aldehydes are oxidised to carboxylic acids and give the silver mirror. Ketones cannot be oxidised easily and give no reaction. This distinguishes the two.

Overall: CH₃CHO + 2[Ag(NH₃)₂]⁺ + H₂O → CH₃COOH + 2Ag + 4NH₃ + 2H⁺.

In acidic solution the purple MnO₄⁻ ion is reduced to almost colourless Mn²⁺, so it acts as its own indicator. The end-point is the first permanent pale pink colour. Use dilute sulfuric acid (not hydrochloric acid, which MnO₄⁻ oxidises).

With Fe²⁺: MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O With ethanedioate: 2MnO₄⁻ + 16H⁺ + 5C₂O₄²⁻ → 2Mn²⁺ + 8H₂O + 10CO₂ (warm to about 60 °C as it is slow at first; Mn²⁺ then catalyses it)

Worked example: 25.0 cm³ of Fe²⁺ solution needs 21.20 cm³ of 0.0200 mol dm⁻³ KMnO₄. Moles MnO₄⁻ = 4.24 × 10⁻⁴; Fe²⁺ = 5 × 4.24 × 10⁻⁴ = 2.12 × 10⁻³ mol; concentration = 2.12 × 10⁻³ / 0.0250 = 0.0848 mol dm⁻³.