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When a transition-metal salt dissolves in water, the cation does not float around as a bare M^n+; it sits at the centre of a hexa-aqua complex [M(H₂O)₆]^n+, with six water ligands arranged octahedrally. The chemistry of that aqua complex — its colour, its acidity, and how it responds to added base — is what colours every test-tube reaction you will perform in this course. In this lesson we systematically tabulate the reactions of the AQA-mandated cations (Cu²⁺, Co²⁺, Fe²⁺, Fe³⁺, Al³⁺, Cr³⁺) with strong alkali (NaOH), with weak alkali (NH₃, in small amount and in excess), and with sodium carbonate (Na₂CO₃). We explain why 3+ aqua ions are spectacularly more acidic than 2+ aqua ions (~four orders of magnitude in Ka), why some hydroxides redissolve in excess base and others do not, and why a 3+ aqua ion fizzes with Na₂CO₃ whereas a 2+ aqua ion merely precipitates. The lesson anchors Required Practical 11 (qualitative tests for cations and anions) and finishes with a diagnostic flow-scheme suitable for an unknown-ion paper question.
Spec mapping (AQA 7405): This lesson maps to §3.2.6 ("Reactions of ions in aqueous solution"), and provides the assessment anchor for Required Practical 11 ("Carry out simple test-tube reactions to identify transition metal ions in aqueous solution"). It builds on lesson 5 (formation of complex ions), lesson 6 (ligand-substitution reactions), and lesson 7 (origin of colour and d–d transitions). The systematic NaOH / NH₃ / Na₂CO₃ scheme developed here is the same scheme used for cation identification in the multi-stage qualitative analysis common to Paper 3. Refer to the official AQA specification document for exact wording.
Assessment objectives: AO1 recall the colour of each aqua ion, the colour and identity of the hydroxide precipitate formed with NaOH, the behaviour of the precipitate in excess NaOH and in excess NH₃, and the products with Na₂CO₃. AO2 write balanced ionic equations for partial deprotonation (hydroxide precipitation), for amphoteric dissolution in excess OH⁻, for ammine complex formation in excess NH₃, and for the metal-carbonate vs metal-hydroxide-plus-CO₂ outcomes with Na₂CO₃. AO3 rationalise the acidity difference between 2+ and 3+ aqua ions in terms of charge density and ligand-water polarisation, predict gas-vs-precipitate behaviour with carbonate from first principles, and design a logical qualitative test scheme to distinguish white precipitates (Al(OH)₃, Mg(OH)₂, Zn(OH)₂).
When a soluble transition-metal salt such as CuSO₄ or FeCl₃ dissolves in water, the cation immediately coordinates six water molecules through lone pairs on oxygen, forming an octahedral hexa-aqua complex [M(H₂O)₆]^n+. The water molecules are tightly bound — exchange of a coordinated water with bulk solvent is slow on the laboratory timescale for ions such as Cr³⁺ (water-exchange half-life ~10⁵ s) but much faster for ions such as Cu²⁺ (sub-microsecond). For all six AQA target ions, the species you should write in equations is [M(H₂O)₆]^n+, not the bare M^n+.
| Aqua ion | Colour in solution | d-electron count |
|---|---|---|
| [Fe(H₂O)₆]²⁺ | Pale green | d⁶ |
| [Fe(H₂O)₆]³⁺ | Pale violet (often appears yellow-brown due to partial hydrolysis) | d⁵ |
| [Cu(H₂O)₆]²⁺ | Pale blue | d⁹ |
| [Co(H₂O)₆]²⁺ | Pink | d⁷ |
| [Cr(H₂O)₆]³⁺ | Violet/ruby (often appears dark green from chloride ligand substitution) | d³ |
| [Al(H₂O)₆]³⁺ | Colourless | d⁰ (not a transition metal — included for the practical) |
Key Point: Aluminium is not a transition metal (no partly-filled d sub-shell), but [Al(H₂O)₆]³⁺ undergoes the same acid–base chemistry as the genuine 3+ aqua ions and is tested alongside them in qualitative-analysis practicals.
A coordinated water molecule is significantly more acidic than free water. Free water has Ka ≈ 10⁻¹⁴; coordinated waters in [Fe(H₂O)₆]³⁺ have Ka ≈ 6 × 10⁻³ — roughly the acidity of formic acid. The reason is electron-density withdrawal: the highly charged metal cation pulls electron density from the lone pair on oxygen towards itself, weakening the O–H bond. A proton can then be transferred to bulk water:
[M(H₂O)₆]^n+(aq) + H₂O(l) ⇌ [M(H₂O)₅(OH)]^(n−1)+(aq) + H₃O⁺(aq)
Further deprotonations are possible (loss of a second and third proton), but each successive Ka is much smaller because the complex becomes less positively charged with each proton loss.
| Aqua ion | Approximate Ka (first deprotonation) | pH of 0.1 M solution |
|---|---|---|
| [Na(H₂O)ₙ]⁺ | ~10⁻¹⁴ | 7 (effectively unchanged) |
| [Mg(H₂O)₆]²⁺ | ~10⁻¹² | ~6.5 |
| [Cu(H₂O)₆]²⁺ | ~10⁻⁸ | ~4.5 |
| [Fe(H₂O)₆]²⁺ | ~10⁻¹⁰ | ~5 |
| [Co(H₂O)₆]²⁺ | ~10⁻⁹ | ~5 |
| [Al(H₂O)₆]³⁺ | ~10⁻⁵ | ~3 |
| [Cr(H₂O)₆]³⁺ | ~10⁻⁴ | ~3 |
| [Fe(H₂O)₆]³⁺ | ~6 × 10⁻³ | ~2 |
The crucial generalisation: 3+ hexa-aqua ions are roughly 10⁴ times more acidic than 2+ hexa-aqua ions of comparable size. This single ratio explains essentially every qualitative-test observation in this lesson.
Charge density (charge ÷ ionic radius) measures how strongly the cation polarises its ligands. For [Fe(H₂O)₆]³⁺ (ionic radius 64.5 pm, charge 3+), the charge density is roughly 47 charge units / nm; for [Fe(H₂O)₆]²⁺ (radius 78 pm, charge 2+), it is roughly 26 charge units / nm. The 3+ ion polarises the O–H bonds nearly twice as effectively, lowering the activation barrier for proton transfer and shifting the hydrolysis equilibrium far to the right.
Synoptic link: The same logic — small, highly-charged cation polarising a soft anion — explains the increased covalent character of LiH vs NaH, and the increased thermal stability of MgCO₃ vs Na₂CO₃. Charge density is one of the most transferable concepts in inorganic chemistry.
Sodium hydroxide is a strong base. Added to a solution of [M(H₂O)₆]^n+, hydroxide ions deprotonate coordinated waters until the complex becomes electrically neutral; at that point it falls out of solution as an insoluble hydroxide precipitate. The general half-equation for a 2+ aqua ion is two deprotonations:
[M(H₂O)₆]²⁺(aq) + 2OH⁻(aq) → M(H₂O)₄(OH)₂ + 2H₂O(l)
For a 3+ aqua ion, three deprotonations are needed:
[M(H₂O)₆]³⁺(aq) + 3OH⁻(aq) → M(H₂O)₃(OH)₃ + 3H₂O(l)
The product is conventionally written without the coordinated waters as M(OH)₂ or M(OH)₃, but in formal equations the hydrated form [M(H₂O)_n(OH)_m] is preferred — AQA mark schemes accept either, but the hydrated form scores depth marks.
| Ion | Formula of precipitate | Colour of precipitate | Notes |
|---|---|---|---|
| [Cu(H₂O)₆]²⁺ | [Cu(H₂O)₄(OH)₂] | Light blue | "Robin's egg" blue gelatinous solid |
| [Co(H₂O)₆]²⁺ | [Co(H₂O)₄(OH)₂] | Blue, turning pink | Pink form is more thermodynamically stable; blue is kinetic |
| [Fe(H₂O)₆]²⁺ | [Fe(H₂O)₄(OH)₂] | Dark green (sometimes "dirty green") | Turns brown on standing as Fe(OH)₂ is air-oxidised to Fe(OH)₃ |
| [Fe(H₂O)₆]³⁺ | [Fe(H₂O)₃(OH)₃] | Rust-brown / red-brown | Often described as "ferric hydroxide" |
| [Al(H₂O)₆]³⁺ | [Al(H₂O)₃(OH)₃] | White gelatinous | Voluminous, jelly-like — characteristic of Al |
| [Cr(H₂O)₆]³⁺ | [Cr(H₂O)₃(OH)₃] | Grey-green to green | Colour depends on counter-anion |
Two of these hydroxides — [Al(H₂O)₃(OH)₃] and [Cr(H₂O)₃(OH)₃] — are amphoteric. In a large excess of OH⁻, further deprotonation removes additional protons from coordinated waters, producing soluble anionic complexes:
Al(H₂O)₃(OH)₃ + OH⁻(aq) → [Al(OH)₄]⁻(aq) + 3H₂O(l) (colourless)
Cr(H₂O)₃(OH)₃ + 3OH⁻(aq) → [Cr(OH)₆]³⁻(aq) + 3H₂O(l) (dark green)
The other hydroxides (Cu(OH)₂, Co(OH)₂, Fe(OH)₂, Fe(OH)₃) are not amphoteric under A-Level conditions — no observable change with excess NaOH; the precipitate persists.
Exam Tip: Only Al(OH)₃ and Cr(OH)₃ dissolve in excess NaOH. This is one of the two most common diagnostic discriminators on Paper 3.
Ammonia plays two roles in qualitative tests. Acting as a Brønsted base in low concentration, it deprotonates coordinated waters exactly as NaOH does — producing the same hydroxide precipitate with the same colour. Acting as a Lewis base (a ligand) in excess, however, it can substitute coordinated waters or hydroxides to form an ammine complex, dissolving the precipitate.
Same as with dilute NaOH — the hydroxide precipitate forms with the same colour and equation.
| Precipitate | Excess NH₃ result | Final complex (if soluble) | Colour |
|---|---|---|---|
| Cu(OH)₂ | Dissolves | [Cu(NH₃)₄(H₂O)₂]²⁺ | Deep royal blue |
| Co(OH)₂ | Dissolves; further air-oxidation to Co(III) | [Co(NH₃)₆]²⁺ → [Co(NH₃)₆]³⁺ | Straw-yellow → brown |
| Fe(OH)₂ | No change — precipitate remains | — | Green (or brown after air-oxidation) |
| Fe(OH)₃ | No change — precipitate remains | — | Rust-brown |
| Al(OH)₃ | No change — precipitate remains | — | White |
| Cr(OH)₃ | Slowly dissolves (very large excess) | [Cr(NH₃)₆]³⁺ | Purple-violet |
The Cr(OH)₃ + NH₃ reaction is slow at room temperature; the AQA mark scheme typically accepts either "purple solution forms" or "no observable change in normal lab timeframe" provided reasoning is shown.
Copper(II) — the canonical example:
[Cu(H₂O)₆]²⁺(aq) + 4NH₃(aq) → [Cu(NH₃)₄(H₂O)₂]²⁺(aq) + 4H₂O(l)
Or equivalently, dissolving the freshly-formed hydroxide precipitate:
Cu(H₂O)₄(OH)₂ + 4NH₃(aq) → [Cu(NH₃)₄(H₂O)₂]²⁺(aq) + 2OH⁻(aq) + 2H₂O(l)
The intense, characteristic deep-blue colour is a definitive test for Cu²⁺.
Cobalt(II):
[Co(H₂O)₆]²⁺(aq) + 6NH₃(aq) → [Co(NH₃)₆]²⁺(aq) + 6H₂O(l) (straw-yellow)
On standing in air, the Co(II) complex is oxidised to the more stable Co(III) ammine:
4[Co(NH₃)₆]²⁺(aq) + O₂(g) + 2H₂O(l) → 4[Co(NH₃)₆]³⁺(aq) + 4OH⁻(aq) (brown/dark-yellow)
Why don't Al³⁺, Fe³⁺ redissolve? Ammonia is too weak a ligand (and too weakly basic) to disrupt the strong Al–O and Fe–O bonds in the hydroxide lattice. Cu²⁺ and Co²⁺ form especially stable ammine complexes (large stability constants — K_stab ≈ 10¹² for [Cu(NH₃)₄]²⁺, ~10⁵ for [Co(NH₃)₆]²⁺), so for them ligand substitution is thermodynamically favoured. For Al³⁺/Fe³⁺/Cr³⁺ the hydroxide is the kinetic and thermodynamic sink.
Common Misconception: Students assume "ammonia base = hydroxide base." For partial-deprotonation purposes they behave identically, but in excess the two diverge sharply: excess OH⁻ dissolves the amphoteric Al(OH)₃ and Cr(OH)₃; excess NH₃ dissolves the soft-acid Cu(OH)₂ and Co(OH)₂. The two reagents therefore probe complementary chemistry.
Carbonate is a weak base (pKb of CO₃²⁻/HCO₃⁻ ≈ 3.7) and a precipitating anion (CO₃²⁻ forms insoluble salts with most divalent metals). When Na₂CO₃ is added to an aqueous metal ion, the outcome depends entirely on whether the aqua ion is acidic enough to protonate the carbonate.
For [M(H₂O)₆]²⁺ (Ka ≈ 10⁻⁸ to 10⁻¹⁰), the aqua complex is not acidic enough to protonate carbonate. Carbonate simply behaves as a counter-anion and precipitates the metal as MCO₃:
[M(H₂O)₆]²⁺(aq) + CO₃²⁻(aq) → MCO₃(s) + 6H₂O(l)
| Ion | Carbonate formed | Colour |
|---|---|---|
| Cu²⁺ | CuCO₃ | Green-blue |
| Co²⁺ | CoCO₃ | Pink |
| Fe²⁺ | FeCO₃ | Pale (dirty) green |
| Mn²⁺ | MnCO₃ | Off-white |
No gas evolution — this is diagnostic of a 2+ aqua ion.
For [M(H₂O)₆]³⁺ (Ka ≈ 10⁻³ to 10⁻⁵), the aqua complex is acidic enough to protonate carbonate. Carbonate behaves as a Brønsted base, abstracting protons from coordinated water; the metal precipitates as the neutral hydroxide, and CO₂ gas is released:
2[M(H₂O)₆]³⁺(aq) + 3CO₃²⁻(aq) → 2M(H₂O)₃(OH)₃ + 3CO₂(g) + 3H₂O(l)
This works for all three 3+ ions:
| Ion | Hydroxide formed | Colour | Gas? |
|---|---|---|---|
| Fe³⁺ | Fe(OH)₃ | Rust-brown | Yes — CO₂ effervescence |
| Al³⁺ | Al(OH)₃ | White gelatinous | Yes — CO₂ effervescence |
| Cr³⁺ | Cr(OH)₃ | Grey-green | Yes — CO₂ effervescence |
Effervescence (turning limewater milky) is diagnostic of a 3+ aqua ion.
Key Point: The carbonate test discriminates 2+ from 3+ aqua ions on the basis of solution acidity. M²⁺ gives a metal-carbonate precipitate with no gas; M³⁺ gives a metal-hydroxide precipitate plus CO₂ effervescence. This is the cleanest charge-discriminating test in qualitative inorganic chemistry.
Q: A pale-blue solution is treated with aqueous Na₂CO₃ and a green-blue precipitate forms without effervescence. A separate yellow-brown solution gives a rust-brown precipitate with vigorous effervescence. Identify the two solutions and write ionic equations for both reactions.
A:
This table consolidates everything you need to memorise for Paper 3 cation identification.
| Ion | OH⁻ (1 eq) | OH⁻ (excess) | NH₃ (1 eq) | NH₃ (excess) | Na₂CO₃ |
|---|---|---|---|---|---|
| Cu²⁺ (blue) | Light blue ppt [Cu(H₂O)₄(OH)₂] | No change | Light blue ppt | Dissolves → deep blue [Cu(NH₃)₄(H₂O)₂]²⁺ | Green-blue ppt CuCO₃, no gas |
| Co²⁺ (pink) | Blue ppt → pink | No change | Blue ppt → pink | Dissolves → straw [Co(NH₃)₆]²⁺ (browns in air) | Pink ppt CoCO₃, no gas |
| Fe²⁺ (green) | Green ppt (browns in air) | No change | Green ppt | No change | Pale-green ppt FeCO₃, no gas |
| Fe³⁺ (yellow-brown) | Rust-brown ppt [Fe(H₂O)₃(OH)₃] | No change | Rust-brown ppt | No change | Rust-brown ppt + CO₂ effervescence |
| Al³⁺ (colourless) | White gelatinous ppt | Dissolves → [Al(OH)₄]⁻ colourless | White ppt | No change | White ppt + CO₂ effervescence |
| Cr³⁺ (violet/green) | Grey-green ppt | Dissolves → [Cr(OH)₆]³⁻ dark green | Grey-green ppt | Slow dissolution → [Cr(NH₃)₆]³⁺ purple | Grey-green ppt + CO₂ effervescence |
The five discriminating features:
RP11 asks you to identify an unknown salt by carrying out a sequence of test-tube reactions on the cation and the anion. The protocol below is the version used at most centres.
Q: An unknown solid X dissolves in water to give a violet solution. Treatment with dilute NaOH gives a grey-green precipitate that dissolves on further addition of NaOH to give a dark-green solution. Treatment with dilute NH₃ gives a grey-green precipitate that does not dissolve readily in excess. Treatment with Na₂CO₃ gives a grey-green precipitate with effervescence. The acidified solution gives a white precipitate with BaCl₂. Identify X.
A:
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