Summary — What This Chapter Covers
- Vanadium chemistry: four oxidation states (+2 to +5), each with its own colour, reduced step-by-step by zinc.
- Chromium chemistry: oxidation states +2, +3, +6; the famous yellow ⇌ orange chromate/dichromate equilibrium.
- Ions in aqueous solution: how metal-aqua ions react with hydroxide and ammonia — precipitates, ligand substitution, and amphoteric behaviour.
- Catalysts: heterogeneous (different phase, surface adsorption theory) vs homogeneous (same phase, intermediate species).
- The Contact Process: V₂O₅ as a heterogeneous catalyst that changes oxidation state mid-reaction.
- Catalytic converters: Pt/Rh catalysts converting car exhaust pollutants.
- Homogeneous catalysis: Fe²⁺/Fe³⁺ catalysing the iodide–peroxodisulfate reaction.
- Autocatalysis: Mn²⁺ speeding up its own formation in the manganate(VII)–oxalate reaction.
1. Vanadium Chemistry
Vanadium is unusual because it comfortably exists in four different oxidation states, and — brilliantly for
students trying to remember it — every single one has a different, vivid colour. Think of it like a traffic light
system, except it has four "lights" instead of three, and they cycle from yellow all the way down to purple as
vanadium keeps gaining electrons.
| Oxidation State | Formula | Name | Colour |
|---|
| +5 | VO₂⁺ | Dioxovanadium(V) | Yellow |
| +4 | VO²⁺ | Oxovanadium(IV) | Blue |
| +3 | V³⁺ | Vanadium(III) | Green |
| +2 | V²⁺ | Vanadium(II) | Purple |
Memory trick
"Yellow Bus Goes Purple" — VO₂⁺ (Yellow) → VO²⁺ (Blue) → V³⁺ (Green) → V²⁺ (Purple), as you add a reducing agent
like zinc and travel down the oxidation states from +5 to +2.
Why does zinc do this?
Zinc metal is a strong reducing agent — it's very willing to lose its own electrons (Zn → Zn²⁺ + 2e⁻),
and those electrons have to go somewhere. In acidic conditions, vanadium(V) grabs them one pair at a time, stepping
down through +4, +3, and finally +2. Each step is its own separate redox reaction, and each one has a colour change
you could watch happen in a test tube — genuinely one of the most visually striking reactions in A Level chemistry.
Using standard electrode potentials (E°) to predict reactions
This is the technique the exam is really testing — not memorising vanadium chemistry, but knowing how to use
E° values to work out which reaction happens. Here's the golden rule, stated as simply as possible:
The half-equation with the more positive E° value happens as written (this is the reduction).
The half-equation with the less positive (more negative) E° value gets reversed (this becomes the oxidation).
Then balance electrons between the two half-equations and add them together — electrons cancel out completely in the final equation.
Worked example — reducing V⁵⁺ to V⁴⁺ using zinc:
- Zn²⁺(aq) + 2e⁻ ⇌ Zn(s) E° = −0.76 V
- VO₂⁺(aq) + 2H⁺(aq) + e⁻ ⇌ VO²⁺(aq) + H₂O(l) E° = +1.00 V
2. Chromium Chemistry
redox
reactions
acid-base
equilibrium