Introduction to Variable Oxidation States Revision Notes | A Level AQA Chemistry

1. Why Transition Metals Behave Differently

Cast your mind back to the electronic configurations you learned at AS level. For a main-group element such as sodium, removing electrons beyond the first ionisation energy costs so much energy that only one oxidation state is chemically accessible. Transition metals play by different rules — and the reason is buried in their electronic structure.

Transition metals are defined as elements that form at least one stable ion with a partially filled 3d subshell. The critical point is that the 3d and 4s subshells are very close in energy. Because of this proximity, different numbers of electrons can be removed without an enormous energy penalty, giving rise to a range of stable oxidation states — a property called variable oxidation states.

Consider manganese: its ground-state configuration is [Ar] 3d$^5$ 4s$^2$. Manganese forms stable species in oxidation states ranging from +2 all the way to +7 — from $\text{Mn}^{2+}$ in pale pink solution to the deep purple $\text{MnO}_4^-$ permanganate ion. Each oxidation state corresponds to a different number of electrons removed from the 3d and 4s subshells. This is the hallmark of transition metal chemistry.

Contrast this with a non-transition metal such as zinc. Zinc has the configuration [Ar] 3d$^{10}$ 4s$^2$. Once the 4s electrons are removed to give $\text{Zn}^{2+}$, the 3d subshell is full and stable — removing further electrons requires far too much energy. Zinc therefore shows only one oxidation state (+2) in its chemistry, which is why it does not qualify as a transition metal by the strict AQA definition.

2. Vanadium: A Case Study in Variable Oxidation States

Vanadium is the AQA specification's showcase example of variable oxidation states, and for good reason — it displays four distinct, coloured oxidation states that you can observe in a single experiment. The reaction of vanadate(V) ions with zinc in acidic solution is a stepwise reduction, and each step produces a visually different species.

The starting material in acidic solution is the dioxovanadium(V) ion, $\text{VO}_2^+$, which is yellow. Zinc acts as the reducing agent throughout, supplying electrons. As reduction proceeds, the vanadium passes through oxidation states +4, +3, and finally +2, each with a characteristic colour. Recognising these colours and the corresponding species is a firm AQA requirement.

The reduction proceeds in stages because each step has its own standard electrode potential. Zinc is a sufficiently strong reducing agent to drive all three steps under the acidic conditions used, but the reactions occur sequentially — you see the colour change from yellow through blue and green to violet as the reduction progresses.

Worked example: Stepwise Reduction of Vanadate(V) by Zinc in Acidic Solution

Write the half-equations for each step in the reduction of $\text{VO}_2^+$ to $\text{V}^{2+}$ by zinc in acidic solution, and identify the colour change at each stage.

1. $\text{VO}_2^+(\text{aq}) + 2\text{H}^+(\text{aq}) + e^- \rightarrow \text{VO}^{2+}(\text{aq}) + \text{H}_2\text{O}(\text{l})$ — One electron is gained; two H$^+$ ions are needed to remove an oxygen as water. Colour changes from yellow to blue.
2. $\text{VO}^{2+}(\text{aq}) + 2\text{H}^+(\text{aq}) + e^- \rightarrow \text{V}^{3+}(\text{aq}) + \text{H}_2\text{O}(\text{l})$ — Another electron is gained; the remaining oxygen is removed as water using two more H$^+$ ions. Colour changes from blue to green.
3. $\text{V}^{3+}(\text{aq}) + e^- \rightarrow \text{V}^{2+}(\text{aq})$ — A simple one-electron reduction; no oxygen atoms remain to remove, so no H$^+$ is needed. Colour changes from green to violet.
4. $\text{Zn}(\text{s}) \rightarrow \text{Zn}^{2+}(\text{aq}) + 2e^-$ — Zinc is oxidised, supplying the electrons for each of the three reduction steps. Each zinc atom provides two electrons, so it can drive two single-electron reductions.

Answer: The four species observed in sequence are: $\text{VO}_2^+$ (yellow) $\rightarrow$ $\text{VO}^{2+}$ (blue) $\rightarrow$ $\text{V}^{3+}$ (green) $\rightarrow$ $\text{V}^{2+}$ (violet). Zinc is oxidised to $\text{Zn}^{2+}$ throughout.

3. Redox Potentials: pH and Ligand Effects

Now that you have seen how vanadium steps through its oxidation states, there is a deeper question worth addressing: what determines whether a particular reduction is feasible? The answer lies in the standard electrode potential, $E^\circ$, for each half-reaction — but at A-level, you also need to appreciate that this value is not fixed in stone. Two factors can shift it significantly: the pH of the solution and the identity of the ligand surrounding the metal ion.

Look back at the half-equation for the V(+5) to V(+4) step: $\text{VO}_2^+(\text{aq}) + 2\text{H}^+(\text{aq}) + e^- \rightarrow \text{VO}^{2+}(\text{aq}) + \text{H}_2\text{O}(\text{l})$. The H$^+$ ions appear on the left-hand side of this equilibrium. By Le Chatelier's principle, increasing the concentration of H$^+$ (lowering the pH) shifts the equilibrium to the right, making the reduction more favourable. In other words, a lower pH makes the $E^\circ$ for this step more positive, so the reduction becomes easier to drive.

The ligand effect is equally important. When a ligand coordinates to a transition metal ion, it alters the electron density around the metal centre. A ligand that donates electron density strongly to the metal stabilises the higher oxidation state less than a weakly donating ligand would — and this shifts the equilibrium between the two oxidation states, changing the $E^\circ$ value. A classic AQA example is the iron(III)/iron(II) couple: in the presence of water as ligand, $E^\circ(\text{Fe}^{3+}/\text{Fe}^{2+}) = +0.77$ V, but when cyanide ligands replace water, the $E^\circ$ shifts to approximately $+0.36$ V, making Fe$^{3+}$ a weaker oxidising agent in that environment.

4. Tollens' Reagent: Silver Mirror and Aldehyde Detection

The variable oxidation states of transition metals are not confined to vanadium or iron — silver provides one of the most visually striking and practically useful examples. Tollens' reagent exploits the ability of silver(I) to be reduced to silver(0), and in doing so it distinguishes between two functional groups that look deceptively similar on paper: aldehydes and ketones.

Tollens' reagent is an alkaline solution of the diamminesilver(I) complex ion, $[\text{Ag}(\text{NH}_3)_2]^+$. The silver is in the +1 oxidation state, coordinated to two ammonia ligands. When an aldehyde is added and the mixture is warmed gently, the aldehyde is oxidised to a carboxylate ion, and the silver(I) is reduced to metallic silver(0), which deposits as a bright, reflective silver mirror on the inside of the test tube — hence the name 'silver mirror test'.

The key half-equations are worth knowing precisely. The silver(I) complex is reduced: $[\text{Ag}(\text{NH}_3)_2]^+(\text{aq}) + e^- \rightarrow \text{Ag}(\text{s}) + 2\text{NH}_3(\text{aq})$. The aldehyde (represented as RCHO) is oxidised: $\text{RCHO}(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightarrow \text{RCOOH}(\text{aq}) + 2\text{H}^+(\text{aq}) + 2e^-$. Combining these gives the overall redox reaction in which the aldehyde reduces the silver complex to metallic silver.

Ketones cannot be oxidised under these mild conditions. A ketone has no hydrogen atom attached to the carbonyl carbon — the C=O group is flanked by two carbon groups — so there is no equivalent oxidation pathway available. When Tollens' reagent is warmed with a ketone, no silver mirror forms and the solution remains colourless. This difference in reactivity is the basis of the test.

5. Connecting the Concepts: Redox, Ligands, and Oxidation States

Earlier in this lesson, you saw how the ligand surrounding a metal ion can shift its electrode potential. Tollens' reagent is a perfect illustration of this principle in action. Free silver ions ($\text{Ag}^+$) in aqueous solution are a moderately strong oxidising agent. However, when silver(I) is coordinated by two ammonia ligands to form $[\text{Ag}(\text{NH}_3)_2]^+$, the electron density around the silver centre increases, stabilising the +1 oxidation state and making the Ag$^+$/Ag couple less positive. This is precisely why Tollens' reagent is selective: it is mild enough to oxidise aldehydes but not strong enough to oxidise ketones or other less reactive organic compounds.

The same logic applies to the vanadium system you studied earlier. The pH of the acidic solution influences each step of the reduction because H$^+$ ions are consumed as the oxo-ligands are removed as water. Change the pH, and you change the feasibility of each step. Change the ligand environment, and you change the electrode potential. These are not isolated facts — they are the same underlying principle applied to different systems.

6. Exam Technique and Common Pitfalls

Questions on variable oxidation states tend to fall into three categories at A-level: recall of species and colours (especially vanadium), half-equation writing and balancing, and extended 'explain' or 'evaluate' questions on pH/ligand effects. Each demands a different approach, and the marks are often lost in predictable ways.

For half-equation questions, the most reliable method is to balance in this order: (1) balance the element changing oxidation state, (2) balance oxygen using water molecules, (3) balance hydrogen using H$^+$ ions, (4) balance charge using electrons. Always include state symbols — AQA mark schemes frequently award a mark specifically for correct state symbols in ionic equations.

For Tollens' reagent questions, the command word matters enormously. 'Describe' requires you to state what you observe (silver mirror forms on the inside of the test tube) and what this indicates (aldehyde present). 'Explain' requires the chemistry: the aldehyde is oxidised to a carboxylate ion; silver(I) in $[\text{Ag}(\text{NH}_3)_2]^+$ is reduced to silver(0). Giving only observations when asked to explain will cost you marks.

Across this lesson, every concept connects back to one central idea: the 3d and 4s subshells in transition metals are close enough in energy that multiple oxidation states are accessible. From the stepwise reduction of vanadium through four coloured species, to the pH-dependence of electrode potentials, to the selective chemistry of Tollens' reagent — variable oxidation states underpin all of it. Recognise that pattern, and the individual facts become much easier to retain and apply under exam conditions.