Variable Oxidation States, Redox Titrations & Catalysis (AQA A Level Chemistry): Video

Hi, I'm Eleanor with 3 years of experience teaching A-Level Chemistry, and this video is about variable oxidation states, redox titrations and catalysis of the transition metals.

Transition metals switch between oxidation states easily, and that's exactly what lets them act as catalysts by cycling between two states. That same easy, colour-changing switch is what makes their titrations self-indicating, with no external indicator needed.

Transition metals can exist in different oxidation states and switch between them relatively easily. This makes them useful in redox reactions and catalysis, because a transition metal can be oxidised and then reduced again, allowing it to cycle between two oxidation states without being used up. Their different oxidation states often have different colours, which is also useful in redox titrations.

So this video will cover: why oxidation states vary so easily, how that shows up in redox titrations with manganate(VII), and how it lets transition metals catalyse reactions — heterogeneously, homogeneously, and through autocatalysis.

Transition metals like titanium through to copper have several stable oxidation states. That's because their first few ionisation energies are relatively small and close together, so losing a small number of electrons costs little extra energy. Vanadium shows this vividly: reducing it with zinc in acid takes it from yellow, through blue and green, to violet as it drops from an oxidation state of plus 5 to plus 2. The redox potential for these changes also shifts depending on the pH of the solution and which ligand is attached.

A redox titration is an oxidising agent titrated against a reducing agent. With manganate(VII), the colour itself does the work. The deep purple MnO 4 minus ion is reduced to almost colourless Mn 2 plus, so no indicator is needed — the reaction is self-indicating. The manganate(VII) goes in the burette, and the endpoint is the first drop that leaves a permanent pink in the flask.

A heterogeneous catalyst is in a different phase from the reactants and the reaction takes place at active sites on its surface. Vanadium(V) oxide catalysing the Contact process is a classic example, being reduced and then re-oxidised as it does its job. A homogeneous catalyst is in the same phase as the reactants: Fe 2 plus and Fe 3 plus can each act in turn as oxidising and reducing agents, cycling back and forth as they catalyse a reaction.

Autocatalysis is when a reaction is sped up by one of its own products. In the reaction between manganate(VII) and oxalate ions, the Mn 2 plus produced catalyses the reaction itself, cycling between the plus 2 and plus 3 oxidation states. That's why a concentration-time graph for this reaction looks unusual as it speeds up before it slows down, instead of just slowing down from the start.

One thing worth remembering for redox titration calculations: always show your working. Marks are awarded for correct steps along the way, even if the final answer ends up wrong.

Oxidation states vary easily because electrons can come and go with little energy cost. The colour of an ion varies with the different oxidation states which is why titrations with manganate(VII) are self-indicating, and it's also why transition metals make such effective catalysts — whether heterogeneous, homogeneous, or auto.

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