Coaching Centre
Identifying cations
How to identify Ba²⁺, Ca²⁺, Mg²⁺, Pb²⁺, Ag⁺, Cu²⁺, Fe²⁺ and Fe³⁺ in solution using flame tests, precipitation reactions and complexation reactions, with net ionic equations and a flowchart that separates them in as few tests as possible.
A pathology lab gets a sample of bright blue liquid from a farm: a calf died after drinking from a trough. Copper sulfate is used against foot rot in livestock, so the blue colour is a clue, but a clue isn't evidence. Before the instruments are switched on, a chemist can prove which metal ion is there in a few minutes with a dropper bottle, a test tube and a Bunsen burner.
How qualitative tests work
A qualitative test tells you what is present, not how much. Each test relies on a property that only some ions have:
- a characteristic flame colour;
- forming an insoluble precipitate with a particular anion (from your solubility rules in Module 5);
- forming a coloured complex ion.
A good test is specific (only one ion responds) and easy to see. Few tests are specific on their own, so you combine them in an order that narrows the choices step by step.
Flame tests
Dip a clean nichrome wire (or a wooden splint soaked overnight) in the sample and hold it in the hottest part of a blue Bunsen flame. The heat excites electrons in the metal atoms to higher energy levels; when they return to lower levels they emit light with energy $E = hf$ equal to the difference between the levels. Each element has a unique set of energy levels, so a unique emission spectrum, and some give a strong visible colour.
| Ion | Flame colour |
|---|---|
| Ba²⁺ | pale (apple) green |
| Ca²⁺ | brick red (orange-red) |
| Cu²⁺ | blue-green |
| Na⁺ / K⁺ / Sr²⁺ | yellow-orange / lilac / crimson (useful to know, not on the syllabus list) |
| Mg²⁺, Ag⁺, Fe²⁺, Fe³⁺ | no useful colour |
| Pb²⁺ | faint blue-white (unreliable) |
Flame tests have limits. Only some metals emit in the visible range, sodium contamination (from fingers or glassware) swamps other colours, and colours are judged by eye. In a mixture the strongest colour hides the others. A spectroscope splits the light into lines, which removes the guesswork. You met line spectra in Year 9; at this level, connect them to energy levels and to atomic absorption spectroscopy (8E), which uses exactly the same energy gaps in reverse.
Precipitation tests
Adding a reagent that forms an insoluble salt with only some cations sorts them into groups. Write each result as a net ionic equation, leaving out spectator ions:
| Reagent | Result | Net ionic equation |
|---|---|---|
| dilute HCl (Cl⁻) | white ppt with Ag⁺ and Pb²⁺ only. PbCl₂ dissolves in hot water; AgCl does not, and darkens in light | $\text{Ag}^+(aq) + \text{Cl}^-(aq) \rightarrow \text{AgCl}(s)$ $\text{Pb}^{2+}(aq) + 2\text{Cl}^-(aq) \rightarrow \text{PbCl}_2(s)$ |
| dilute H₂SO₄ (SO₄²⁻) | white ppt with Ba²⁺ and Pb²⁺ (Ca²⁺ only if concentrated: CaSO₄ is slightly soluble) | $\text{Ba}^{2+}(aq) + \text{SO}_4^{2-}(aq) \rightarrow \text{BaSO}_4(s)$ |
| NaOH (OH⁻), a few drops | Cu²⁺ blue; Fe²⁺ green, turning brown at the surface; Fe³⁺ red-brown; Mg²⁺ white; Pb²⁺ white (dissolves in excess); Ag⁺ brown (Ag₂O); Ca²⁺ faint white | $\text{Cu}^{2+}(aq) + 2\text{OH}^-(aq) \rightarrow \text{Cu(OH)}_2(s)$ $\text{Fe}^{3+}(aq) + 3\text{OH}^-(aq) \rightarrow \text{Fe(OH)}_3(s)$ |
| KI (I⁻) | bright yellow PbI₂; pale yellow AgI | $\text{Pb}^{2+}(aq) + 2\text{I}^-(aq) \rightarrow \text{PbI}_2(s)$ |
Fe(OH)₂ turns brown because oxygen in the air oxidises Fe²⁺ to Fe³⁺. Pb(OH)₂ dissolves in excess hydroxide because it is amphoteric: it reacts with a base to form the soluble plumbate ion.
Complexation reactions
A complex ion forms when molecules or ions with lone pairs (ligands) bond to a central metal ion. Complexes often have intense colours, which makes them excellent confirmatory tests:
- Fe³⁺ with thiocyanate gives a blood-red solution: $\text{Fe}^{3+}(aq) + \text{SCN}^-(aq) \rightleftharpoons [\text{FeSCN}]^{2+}(aq)$. You used this equilibrium in Module 5.
- Cu²⁺ with excess ammonia: the pale blue Cu(OH)₂ precipitate dissolves to give a deep blue solution, $\text{Cu}^{2+}(aq) + 4\text{NH}_3(aq) \rightarrow [\text{Cu(NH}_3)_4]^{2+}(aq)$.
- Ag⁺: a brown precipitate with a little ammonia dissolves in excess as colourless $[\text{Ag(NH}_3)_2]^+$. The same complex is why AgCl dissolves in ammonia.
- Fe²⁺ decolourises acidified permanganate (it is oxidised to Fe³⁺), which Fe³⁺ cannot do.
Putting the tests in order
The order matters when a solution might contain more than one cation. Chloride removes Ag⁺ and Pb²⁺ first; sulfate then removes Ba²⁺ (and any remaining Pb²⁺); hydroxide then separates the rest by colour. After each precipitation, filter or centrifuge and test the filtrate, so that an ion removed in an earlier step can't confuse a later one.
- Identify three unknowns. Try to do each one in three tests or fewer.
- Find an unknown that gives no flame colour and no precipitate with HCl or H₂SO₄. Which test separates the three ions that are left?
- Compare Cu²⁺ with a few drops of NaOH and with excess NH₃. Write a net ionic equation for each.
- Find two ions that both give a white precipitate with HCl. How does KI tell them apart?
Test yourself
A solution gives no precipitate with HCl, a white precipitate with H₂SO₄ and a pale green flame. The cation is
Pb²⁺ would also precipitate with HCl. Ba²⁺ gives a green flame.
Which reagent distinguishes Fe²⁺ from Fe³⁺ most clearly?
Fe³⁺ gives blood-red [FeSCN]²⁺; Fe²⁺ does not. NaOH would also work (green vs red-brown).
The net ionic equation for adding ammonia solution to a pale blue copper(II) hydroxide precipitate is
Ammonia ligands replace the hydroxide to form the deep blue tetraamminecopper(II) complex.
Why does each element give a different flame colour?
Emitted photon energies equal the gaps between energy levels.
Name each reagent or test in the flowchart.
Sort the cations. Some could fit more than one group: choose the most useful test.
Calcium ions give a flame. Iron(III) ions form a precipitate with hydroxide. Silver chloride dissolves in ammonia because it forms a ion. Lead(II) iodide is a bright precipitate.
A technician has four unlabelled bottles, known to contain solutions of barium nitrate, calcium nitrate, lead(II) nitrate and copper(II) nitrate. Describe a sequence of tests to identify each solution. Include expected observations and net ionic equations.
Step 1: the blue solution is likely copper(II) nitrate. Confirm by adding a few drops of NaOH: a pale blue precipitate forms, $\text{Cu}^{2+}(aq) + 2\text{OH}^-(aq) \rightarrow \text{Cu(OH)}_2(s)$, which dissolves in excess ammonia to a deep blue solution.
Step 2: add dilute HCl to fresh samples of the three colourless solutions. Only lead(II) nitrate gives a white precipitate (which dissolves on heating): $\text{Pb}^{2+}(aq) + 2\text{Cl}^-(aq) \rightarrow \text{PbCl}_2(s)$. Adding KI to a fresh sample confirms it with a bright yellow precipitate of PbI₂.
Step 3: add dilute H₂SO₄ to the remaining two. Barium nitrate gives a dense white precipitate, $\text{Ba}^{2+}(aq) + \text{SO}_4^{2-}(aq) \rightarrow \text{BaSO}_4(s)$; calcium nitrate gives little or none because CaSO₄ is slightly soluble. Confirm with flame tests: barium gives a pale green flame and calcium a brick-red flame.