In aqueous solution, metal cations exist as hydrated complexes. When these solution complexes react with bases such as sodium hydroxide (\(\text{NaOH}\)) or ammonia (\(\text{NH}_3\)), 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.
Reactions with Sodium Hydroxide (\(\text{NaOH}\))
Adding sodium hydroxide dropwise provides hydroxide ions (\(\text{OH}^-\)) 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. \(\text{M}(\text{OH})_2(\text{s})\) or \(\text{M}(\text{OH})_3(\text{s})\).
| Metal Ion | Aqua Complex Reactant | NaOH (Dropwise) Observation | Precipitate Formula | NaOH (Excess) Observation |
|---|---|---|---|---|
| Copper(II) | \([\text{Cu}(\text{H}_2\text{O})_6]^{2+}\) (pale blue) | Blue precipitate | \(\text{Cu}(\text{OH})_2(\text{s})\) | Insoluble (remains blue precipitate) |
| Iron(II) | \([\text{Fe}(\text{H}_2\text{O})_6]^{2+}\) (pale green) | Green precipitate | \(\text{Fe}(\text{OH})_2(\text{s})\) | Insoluble (darkens on standing in air) |
| Iron(III) | \([\text{Fe}(\text{H}_2\text{O})_6]^{3+}\) (yellow/brown) | Brown precipitate | \(\text{Fe}(\text{OH})_3(\text{s})\) | Insoluble (remains brown precipitate) |
| Manganese(II) | \([\text{Mn}(\text{H}_2\text{O})_6]^{2+}\) (pale pink) | Cream/pale brown precipitate | \(\text{Mn}(\text{OH})_2(\text{s})\) | Insoluble (darkens in air to brown) |
| Chromium(III) | \([\text{Cr}(\text{H}_2\text{O})_6]^{3+}\) (green) | Green precipitate | \(\text{Cr}(\text{OH})_3(\text{s})\) | Dissolves to a dark green solution: \([\text{Cr}(\text{OH})_6]^{3-}(\text{aq})\) |
| Aluminium(III) | \([\text{Al}(\text{H}_2\text{O})_6]^{3+}\) (colourless) | White precipitate | \(\text{Al}(\text{OH})_3(\text{s})\) | Dissolves to a colourless solution: \([\text{Al}(\text{OH})_4]^-(\text{aq})\) |
🔑 Key Principle: Amphoteric Cations (Al³⁺ and Cr³⁺)
Aluminium hydroxide and chromium(III) hydroxide are amphoteric. When excess strong base (like \(\text{NaOH}\)) 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 (\(\text{NH}_3\))
Ammonia acts as a weak base in water, creating a low concentration of hydroxide ions: \(\text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^-\). 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, \(\text{Cu}(\text{OH})_2(\text{s})\) | Dissolves to deep blue solution | \([\text{Cu}(\text{NH}_3)_4(\text{H}_2\text{O})_2]^{2+}(\text{aq})\) (ligand substitution) |
| Iron(II) | Green precipitate, \(\text{Fe}(\text{OH})_2(\text{s})\) | Insoluble | Remains \(\text{Fe}(\text{OH})_2(\text{s})\) |
| Iron(III) | Brown precipitate, \(\text{Fe}(\text{OH})_3(\text{s})\) | Insoluble | Remains \(\text{Fe}(\text{OH})_3(\text{s})\) |
| Manganese(II) | Cream precipitate, \(\text{Mn}(\text{OH})_2(\text{s})\) | Insoluble | Remains \(\text{Mn}(\text{OH})_2(\text{s})\) |
| Chromium(III) | Green precipitate, \(\text{Cr}(\text{OH})_3(\text{s})\) | Dissolves to purple solution | \([\text{Cr}(\text{NH}_3)_6]^{3+}(\text{aq})\) (ligand substitution) |
| Aluminium(III) | White precipitate, \(\text{Al}(\text{OH})_3(\text{s})\) | Insoluble | Remains \(\text{Al}(\text{OH})_3(\text{s})\) (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⁺: 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 \(\text{NaOH}\) or \(\text{NH}_3\).
Step 1: Identify the metal ion
The only metal ion in the syllabus that forms a green hydroxide precipitate dissolving in excess sodium hydroxide is the chromium(III) ion, \(\text{Cr}^{3+}\). (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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