Practical: Identifying Metal Ions in Solution Revision Notes | A Level AQA Chemistry

1. Why Identifying Metal Ions Matters

Picture this: you have a colourless or pale solution and you need to determine which transition metal ion it contains. You cannot see inside the solution at the atomic level — but you can provoke it into revealing itself through a series of simple test-tube reactions.

This is the heart of Required Practical 11. The reactions you carry out are not arbitrary — each one exploits a specific chemical property of transition metal ions, and the colour and behaviour of the precipitate you observe is your evidence. Getting this right in an exam means knowing not just what you see, but why you see it.

2. The Chemistry Behind the Colours

Transition metal ions in aqueous solution exist as aqua complexes — the metal ion is surrounded by six water molecules acting as ligands, giving the general form $[M(H_2O)_6]^{n+}$. The colour of these aqua complexes arises because d-electrons absorb specific wavelengths of visible light when they are promoted between split d-orbitals.

When you add sodium hydroxide solution ($NaOH$) or aqueous ammonia ($NH_3$) to these complexes, the $OH^-$ or $NH_3$ ligands displace the water molecules. The result is a precipitate or a new soluble complex — and the colour change is your diagnostic signal.

The key insight is that each transition metal ion produces a distinctive colour of precipitate with $OH^-$ ions, and some of those precipitates then behave differently when you add excess reagent. That two-stage behaviour — initial precipitate, then response to excess — is what allows you to distinguish between ions that might otherwise look similar.

3. Reactions with Sodium Hydroxide Solution

Adding a few drops of dilute sodium hydroxide solution to a transition metal ion solution introduces $OH^-$ ions. These react with the aqua complex, replacing water ligands and causing the metal hydroxide to precipitate out of solution. The general equation follows this pattern:

The behaviour of $Cr^{3+}$ with excess $NaOH$ is particularly important. Chromium(III) hydroxide is amphoteric — it dissolves in excess alkali to form the tetrahydroxochromate(III) ion, $[Cr(OH)_4]^-(aq)$, which is dark green. This distinguishes $Cr^{3+}$ from all other common transition metal ions in this test.

4. Reactions with Aqueous Ammonia

Aqueous ammonia ($NH_3(aq)$) acts as both a weak base and a ligand. As a weak base, it partially ionises in water to provide $OH^-$ ions — but at a much lower concentration than sodium hydroxide. These $OH^-$ ions are sufficient to precipitate the metal hydroxide, so when you add a small amount of ammonia solution, you still get the same coloured hydroxide precipitate as with $NaOH$. The diagnostic difference comes when you add excess ammonia.

Ammonia is also a ligand. In excess, it can displace water molecules from the metal complex through ligand substitution, forming a new soluble ammine complex. This dissolves the precipitate and produces a characteristic colour change — and that colour change is your key distinguishing observation.

Notice that $Cu^{2+}$ is uniquely identified by the deep blue tetraamminediaquacopper(II) complex formed in excess ammonia. No other common transition metal ion produces this intense deep blue colour, making it one of the most reliable identification tests at A-level.

5. Distinguishing Between Similar Ions

Several pairs of ions can look deceptively similar at first glance, and examiners know this. The table below highlights the most commonly confused pairs and the observations that separate them.

The $Fe^{2+}$/$Fe^{3+}$ distinction is one of the most reliable in this practical — the green versus reddish-brown precipitate colours are visually distinct and do not require any further testing. For $Cu^{2+}$ versus $Co^{2+}$, however, the initial precipitate colours can overlap, so always proceed to the excess ammonia test.

6. Required Practical 11 — Procedure and Evaluation

Required Practical 11 asks you to carry out systematic test-tube reactions to identify transition metal ions in aqueous solution. The procedure follows a logical sequence: first test with dilute NaOH, observe the precipitate colour, then add excess NaOH and record any change, then repeat the sequence with aqueous ammonia on a fresh sample.

One practical consideration worth noting: $Fe^{2+}$ precipitates can oxidise in air over time, turning from green $Fe(OH)_2$ towards reddish-brown $Fe(OH)_3$. If you observe a green precipitate that darkens on standing, this is consistent with $Fe^{2+}$ — and it is worth mentioning this in an exam evaluation question as a source of ambiguity.

Worked example: Systematic Identification of an Unknown Transition Metal Ion

An unknown solution produces a blue-green precipitate when dilute NaOH is added. The precipitate does not dissolve in excess NaOH. When the test is repeated with aqueous ammonia, the same blue-green precipitate forms initially, then dissolves in excess ammonia to give a straw-yellow/brown solution. Identify the ion and write the ionic equation for the formation of the precipitate.

1. Blue-green precipitate with NaOH, does not dissolve in excess. — This rules out $Cr^{3+}$ (which dissolves in excess NaOH to give dark green). The blue-green colour is consistent with $Co^{2+}$ or possibly $Cu^{2+}$, though $Cu^{2+}$ typically gives a more distinctly blue precipitate.
2. Precipitate dissolves in excess $NH_3$ to give straw-yellow/brown solution. — This is the key distinguishing observation. $Cu^{2+}$ would give a deep blue solution in excess ammonia. A straw-yellow/brown solution is characteristic of $[Co(NH_3)_6]^{2+}$, confirming the ion is $Co^{2+}$.
3. Ion = $Co^{2+}$ — The combination of blue-green precipitate with NaOH (no dissolution in excess) and dissolution in excess ammonia to give straw-yellow/brown solution is unique to cobalt(II).
4. $[Co(H_2O)_6]^{2+}(aq) + 2OH^-(aq) \rightarrow Co(OH)_2(s) + 6H_2O(l)$ — The aqua complex loses two water ligands and gains two $OH^-$ ions, forming the insoluble cobalt(II) hydroxide precipitate. State symbols are required for full marks.

Answer: The ion is $Co^{2+}$. Ionic equation: $[Co(H_2O)_6]^{2+}(aq) + 2OH^-(aq) \rightarrow Co(OH)_2(s) + 6H_2O(l)$

7. Exam Technique and Common Mistakes

AQA exam questions on this practical tend to fall into two categories: recall of specific observations, and application to an unknown ion scenario. Both require precise language — vague answers like "a coloured precipitate forms" will not score marks. You need to name the colour and state whether it dissolves in excess reagent.

When writing ionic equations for these reactions, always include state symbols and write the metal ion as its aqua complex $[M(H_2O)_6]^{n+}$ if the question asks for the full ionic equation. If the question asks for a simplified ionic equation, you may omit the water ligands — but check the command word carefully.

Required Practical 11 tests your ability to carry out systematic test-tube reactions and interpret observations precisely — both skills that appear directly in AQA exam questions, whether as short 'describe' tasks or extended 'explain and identify' scenarios. Mastering the colour of each precipitate, the behaviour with excess reagent, and the ionic equations that underpin each observation gives you complete coverage of this specification requirement.