In aqueous solution, metal cations exist as hydrated complexes. When these solution complexes react with bases such as sodium hydroxide (NaOH) or ammonia (NH₃), they undergo acid-base reactions rather than ligand substitution. The base accepts protons from the water ligands, leading to the formation of neutral, insoluble metal hydroxide precipitates.
A precipitate is an insoluble solid that emerges from a liquid solution during a chemical reaction.
An amphoteric hydroxide is a metal hydroxide that can react chemically as both an acid (reacting with bases) and a base (reacting with acids) to form soluble species.
Acidity of Metal-Aqua Ions
In water, Fe²⁺ and Cu²⁺ ions form [M(H₂O)₆]²⁺ ions, and Al³⁺ and Fe³⁺ ions form [M(H₂O)₆]³⁺ ions. Each is octahedral, with six water ligands bonded to the metal ion.
The metal ion attracts electron density from its water ligands (it polarises them). This weakens the O–H bonds, so a water ligand can lose H⁺ to a water molecule in the solution:
[M(H₂O)₆]³⁺ is more acidic than [M(H₂O)₆]²⁺ because the 3+ ion has a higher charge/size ratio: it is smaller and more highly charged. It polarises the water ligands more, so H⁺ is released more readily and the equilibrium lies further to the right.
Key Principle: Use the Charge/Size Ratio
AQA asks you to explain why [M(H₂O)₆]³⁺ is more acidic than [M(H₂O)₆]²⁺ in terms of the charge/size ratio of the metal ion, so use that phrase. A full answer: the 3+ ion has a higher charge/size ratio, so it polarises its water ligands more, weakening the O–H bonds, so H⁺ is released more readily.
Note: the [Fe(H₂O)₆]³⁺ ion itself is pale purple, but iron(III) solutions usually look yellow-brown because some [Fe(H₂O)₅(OH)]²⁺ has formed.
Reactions with Sodium Hydroxide (NaOH)
Adding sodium hydroxide dropwise provides hydroxide ions (OH⁻) which act as a base. They remove hydrogen ions (protons) from the water ligands in the hexaaqua metal complexes, forming a neutral hydroxide precipitate. For example:
For convenience in the AQA syllabus, these products are commonly written as simple metal hydroxides, e.g. M(OH)₂(s) or M(OH)₃(s).
| Metal Ion | Aqua Complex Reactant | NaOH (Dropwise) Observation | Precipitate Formula | NaOH (Excess) Observation |
|---|---|---|---|---|
| Copper(II) | [Cu(H₂O)₆]²⁺ (pale blue) | Blue precipitate | Cu(OH)₂(s) | Insoluble (remains blue precipitate) |
| Iron(II) | [Fe(H₂O)₆]²⁺ (pale green) | Green precipitate | Fe(OH)₂(s) | Insoluble (darkens on standing in air) |
| Iron(III) | [Fe(H₂O)₆]³⁺ (yellow/brown) | Brown precipitate | Fe(OH)₃(s) | Insoluble (remains brown precipitate) |
| Manganese(II) Beyond the AQA specification |
[Mn(H₂O)₆]²⁺ (pale pink) | Cream/pale brown precipitate | Mn(OH)₂(s) | Insoluble (darkens in air to brown) |
| Chromium(III) Beyond the AQA specification |
[Cr(H₂O)₆]³⁺ (green) | Green precipitate | Cr(OH)₃(s) | Dissolves to a dark green solution: [Cr(OH)₆]³⁻(aq) |
| Aluminium(III) | [Al(H₂O)₆]³⁺ (colourless) | White precipitate | Al(OH)₃(s) | Dissolves to a colourless solution: [Al(OH)₄]⁻(aq) |
Key Principle: Amphoteric Cations (Al³⁺ and Cr³⁺)
Aluminium hydroxide and chromium(III) hydroxide are amphoteric (AQA 7405 only requires aluminium hydroxide; chromium(III) is beyond the AQA specification). When excess strong base (like NaOH) is added, they act as acids by donating a proton / reacting further with hydroxide ions to form soluble, charged hydroxo-complexes: \[ \text{Al}(\text{OH})_3(\text{s}) + \text{OH}^-(\text{aq}) \rightarrow [\text{Al}(\text{OH})_4]^-(\text{aq}) \quad \text{(colourless solution)} \] \[ \text{Cr}(\text{OH})_3(\text{s}) + 3\text{OH}^-(\text{aq}) \rightarrow [\text{Cr}(\text{OH})_6]^{3-}(\text{aq}) \quad \text{(dark green solution)} \] The other hydroxides do not display this behaviour and remain as insoluble precipitates.
Reactions with Ammonia (NH₃)
Ammonia acts as a weak base in water, creating a low concentration of hydroxide ions: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻. Addition of dropwise ammonia therefore yields the identical hydroxide precipitates as sodium hydroxide.
However, adding excess ammonia produces different results because ammonia can also act as a ligand, leading to ligand substitution for certain metal ions.
| Metal Ion | NH₃ (Dropwise) Observation | NH₃ (Excess) Observation | Soluble Species / Cause |
|---|---|---|---|
| Copper(II) | Blue precipitate, Cu(OH)₂(s) | Dissolves to deep blue solution | [Cu(NH₃)₄(H₂O)₂]²⁺(aq) (ligand substitution) |
| Iron(II) | Green precipitate, Fe(OH)₂(s) | Insoluble | Remains Fe(OH)₂(s) |
| Iron(III) | Brown precipitate, Fe(OH)₃(s) | Insoluble | Remains Fe(OH)₃(s) |
| Manganese(II) Beyond the AQA specification |
Cream precipitate, Mn(OH)₂(s) | Insoluble | Remains Mn(OH)₂(s) |
| Chromium(III) Beyond the AQA specification |
Green precipitate, Cr(OH)₃(s) | Dissolves to purple solution | [Cr(NH₃)₆]³⁺(aq) (ligand substitution) |
| Aluminium(III) | White precipitate, Al(OH)₃(s) | Insoluble | Remains Al(OH)₃(s) (ammonia is a weak base, cannot redissolve Al(OH)₃) |
Key Principle: Distinguishing Aluminium from Chromium and Copper
The distinction between excess NaOH and excess NH₃ is critical for identifying unknown ions:
- Al³⁺: forms a white precipitate with NaOH that dissolves in excess, but the white precipitate with NH₃ does not dissolve in excess.
- Cr³⁺ (beyond the AQA specification): forms a green precipitate that dissolves in excess NaOH (dark green solution) and also dissolves in excess NH₃ (purple solution).
- Cu²⁺: forms a blue precipitate that does not dissolve in excess NaOH but does dissolve in excess NH₃ (deep blue solution).
You must write ionic equations with correct state symbols. When hydroxide precipitates form, the aqueous ions react to form a solid. For 2+ metal ions: \[ \text{M}^{2+}(\text{aq}) + 2\text{OH}^-(\text{aq}) \rightarrow \text{M}(\text{OH})_2(\text{s}) \] For 3+ metal ions: \[ \text{M}^{3+}(\text{aq}) + 3\text{OH}^-(\text{aq}) \rightarrow \text{M}(\text{OH})_3(\text{s}) \] These equations apply whether the source of hydroxide is NaOH or NH₃.
Reactions with Sodium Carbonate (Na₂CO₃)
Carbonate ions are a base, but how they react depends on how acidic the metal-aqua ion is.
2+ ions: [M(H₂O)₆]²⁺ ions are not acidic enough to react with carbonate ions as an acid, so the metal carbonate precipitates:
Iron(II) gives a green precipitate of FeCO₃ and copper(II) gives a blue-green precipitate of CuCO₃. No gas is given off.
3+ ions: [M(H₂O)₆]³⁺ ions are acidic enough to react with carbonate ions, so carbon dioxide is given off (effervescence) and the hydroxide precipitates:
Iron(III) gives a brown precipitate and bubbles, and aluminium gives a white precipitate and bubbles. No metal(III) carbonate forms. The difference comes from the greater acidity of the 3+ ions, explained at the top of this page.
Sodium carbonate tells iron(II) from iron(III). Iron(II) gives a green precipitate with no bubbles; iron(III) gives a brown precipitate and bubbles of carbon dioxide.
Step 1: Identify the metal ion
Of the ions on this page, the only one that forms a green hydroxide precipitate dissolving in excess sodium hydroxide is the chromium(III) ion, Cr³⁺. Chromium(III) is beyond the AQA specification, so treat this example as extension practice. (Iron(II) also forms a green precipitate, but it is insoluble in excess NaOH.)
Step 2: Write the equation for the precipitation
The green precipitate is chromium(III) hydroxide:
\[ \text{Cr}^{3+}(\text{aq}) + 3\text{OH}^-(\text{aq}) \rightarrow \text{Cr}(\text{OH})_3(\text{s}) \]
Step 3: Write the equation for dissolving in excess NaOH
Chromium(III) hydroxide is amphoteric and reacts with excess hydroxide ions to form a soluble hexahydroxochromate(III) complex:
\[ \text{Cr}(\text{OH})_3(\text{s}) + 3\text{OH}^-(\text{aq}) \rightarrow [\text{Cr}(\text{OH})_6]^{3-}(\text{aq}) \]
Step 1: Dropwise addition
Observation: A light blue precipitate forms.
Equation: \[ \text{Cu}^{2+}(\text{aq}) + 2\text{OH}^-(\text{aq}) \rightarrow \text{Cu}(\text{OH})_2(\text{s}) \]
(The hydroxide ions are supplied by the reaction of ammonia with water).
Step 2: Excess addition
Observation: The blue precipitate dissolves to form a deep blue solution.
Equation: \[ \text{Cu}(\text{OH})_2(\text{s}) + 4\text{NH}_3(\text{aq}) + 2\text{H}_2\text{O}(\text{l}) \rightarrow [\text{Cu}(\text{NH}_3)_4(\text{H}_2\text{O})_2]^{2+}(\text{aq}) + 2\text{OH}^-(\text{aq}) \]
This is a ligand substitution reaction where ammonia molecules displace hydroxide and water ligands.
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