Part 4: Variable oxidation states and redox reactions

Track vanadium reduction, the effects of ligands and pH on redox behaviour, and the silver chemistry of Tollens’ reagent.

Balance electron transfers, not colour names

The similar energies of 3d and 4s electrons help make several oxidation states accessible for transition metals. Their relative stabilities depend on the element, ligands and reaction conditions. A colour change may reflect a ligand change, a redox reaction or both; establish the formulae before assigning oxidation states.

To balance an acidic half-equation, balance non-H/O atoms first, oxygen using H₂O, hydrogen using H⁺ and charge using electrons. A reduction has electrons on the left. Multiply half-equations so the electrons cancel in the overall equation; never carry uncancelled electrons into a complete ionic reaction.

The vanadium reduction sequence

In acidic solution, zinc can reduce V(V) through V(IV) and V(III) to V(II). The principal sequence is yellow → blue → green → violet. A mixture of yellow V(V) and blue V(IV) can appear green during the first stage; this is not by itself proof that V(III) has formed. Air can reoxidise low-oxidation-state products.

One vanadium needs three electrons for +5 to +2

Add the three vanadium half-equations: VO₂⁺ + 4H⁺ + 3e⁻ → V²⁺ + 2H₂O. Two vanadium ions therefore accept the six electrons released by three zinc atoms. In the full equation, the left-side charge is +10 and the right-side charge is also +10.

Original worked example: reducing 0.00600 mol V(V) completely to V(II) requires 0.0180 mol electrons, hence a theoretical minimum of 0.00900 mol Zn. With Aᵣ(Zn) = 65.4, this is 0.589 g. An actual experiment can consume additional zinc through its competing reaction with acid to release hydrogen.

pH and ligand change the reduction tendency

A redox potential applies to a specified pair of species under stated conditions. If H⁺ is consumed in the reduction half-equation, lowering [H⁺] generally makes that reduction less favourable relative to the same species at higher acidity. The actual products can also change: manganate(VII) gives Mn²⁺ in sufficiently acidic solution but may form brown MnO₂ when acidity is inadequate.

Ligands stabilise different oxidation states to different extents, changing the potential of a metal redox pair. Cobalt(III) is stabilised by ammonia relative to its simple aqua chemistry, which helps explain why cobalt(II) ammine species can be oxidised by oxygen in air. Do not reuse an aqua-ion potential unchanged for an ammine complex.

When reduction potentials are supplied for the relevant conditions, calculate Ecell = E(reduction) − E(other reduction couple). A positive value supports thermodynamic feasibility in that direction, not a guarantee of a measurable rate. As a constructed illustration, couples at +0.90 V and +0.30 V give +0.60 V when the first is reduced and the second reversed for oxidation.

Tollens’ reagent couples silver reduction to aldehyde oxidation

Tollens’ reagent contains linear [Ag(NH₃)₂]⁺ in alkaline solution. An aldehyde is oxidised to a carboxylate while Ag(I) gains electrons to form silver metal, observed as a silver mirror or grey/silver deposit under suitable conditions. A typical simple ketone gives no reaction in this test. The conclusion is limited by other reducing compounds which may also react.

For ethanal, use the alkaline oxidation half-equation below. Adding twice the silver half-equation cancels two electrons. Since the medium is alkaline, the organic product is CH₃COO⁻, not mainly free CH₃COOH. Freshly prepared reagent is used in supervised laboratory work and disposed of promptly according to the laboratory procedure; it must not be stored or allowed to dry.

Quick checks

Original Finesse questions. Reveal the indicative worked solutions after attempting each question; these are not official AQA mark allocations.

Sources