Reactions of Aqueous Metal Ions

This lesson covers the reactions of aqueous transition metal ions with sodium hydroxide and ammonia, the acidity of hexaaqua ions, and reactions with carbonate ions (AQA 3.2.5). These reactions are frequently tested in practical and written papers.

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Hexaaqua Ions and Their Acidity

When transition metal ions dissolve in water, they form hexaaqua complex ions [M(H₂O)₆]^(n+). These ions are acidic because the highly charged metal cation polarises the O–H bonds in the coordinated water molecules, making it easier for a proton to be released.

Acidity Increases with Charge Density

Ion	Charge	Ionic radius / pm	Charge density	pH of solution
Na⁺	+1	102	Low	≈7 (neutral)
Mg²⁺	+2	72	Moderate	≈6
Al³⁺	+3	53.5	High	≈3
Fe³⁺	+3	64.5	High	≈2–3
Fe²⁺	+2	78	Moderate	≈5

Explanation: The metal cation attracts the electron density of the coordinated water molecules towards itself. This weakens the O–H bond and makes it easier for a proton (H⁺) to be lost to a free water molecule:

[M(H₂O)₆]^(n+)(aq) + H₂O(l) ⇌ [M(OH)(H₂O)₅]^((n-1)+)(aq) + H₃O⁺(aq)

The higher the charge and the smaller the ion, the greater the polarisation and the more acidic the solution.

Exam Tip: M³⁺ ions produce more acidic solutions than M²⁺ ions because they have higher charge density and polarise the water ligands more. This is why Fe³⁺ solutions are more acidic than Fe²⁺ solutions.

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Precipitation Reactions with NaOH

When sodium hydroxide is added to solutions of transition metal ions, insoluble metal hydroxides precipitate.

Table of Observations

Ion	Formula of aqua ion	Colour of aqua ion	With dilute NaOH	Precipitate colour	With excess NaOH
Fe²⁺	[Fe(H₂O)₆]²⁺	Green	Green precipitate	Green	No change (insoluble)
Fe³⁺	[Fe(H₂O)₆]³⁺	Yellow/brown	Brown precipitate	Brown (rust-coloured)	No change (insoluble)
Cu²⁺	[Cu(H₂O)₆]²⁺	Blue	Pale blue precipitate	Pale blue	No change (insoluble)
Cr³⁺	[Cr(H₂O)₆]³⁺	Green	Green precipitate	Green	Dissolves → dark green solution
Al³⁺	[Al(H₂O)₆]³⁺	Colourless	White precipitate	White	Dissolves → colourless solution
Mn²⁺	[Mn(H₂O)₆]²⁺	Very pale pink	Cream/white precipitate	Off-white (darkens in air)	No change (insoluble)
Co²⁺	[Co(H₂O)₆]²⁺	Pink	Blue-green precipitate	Blue-green	No change (insoluble)

Equations

Iron(II): [Fe(H₂O)₆]²⁺(aq) + 2OH⁻(aq) → Fe(OH)₂(s) + 6H₂O(l) — green precipitate

Fe(OH)₂ is oxidised in air to Fe(OH)₃: 4Fe(OH)₂(s) + O₂(g) + 2H₂O(l) → 4Fe(OH)₃(s) — turns brown

Iron(III): [Fe(H₂O)₆]³⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s) + 6H₂O(l) — brown precipitate

Copper(II): [Cu(H₂O)₆]²⁺(aq) + 2OH⁻(aq) → Cu(OH)₂(s) + 6H₂O(l) — pale blue precipitate

Chromium(III): [Cr(H₂O)₆]³⁺(aq) + 3OH⁻(aq) → Cr(OH)₃(s) + 6H₂O(l) — green precipitate

With excess NaOH (amphoteric behaviour): Cr(OH)₃(s) + 3OH⁻(aq) → [Cr(OH)₆]³⁻(aq) — dark green solution

Aluminium(III): [Al(H₂O)₆]³⁺(aq) + 3OH⁻(aq) → Al(OH)₃(s) + 6H₂O(l) — white precipitate

With excess NaOH (amphoteric behaviour): Al(OH)₃(s) + OH⁻(aq) → [Al(OH)₄]⁻(aq) — colourless solution

Exam Tip: Only Cr(OH)₃ and Al(OH)₃ dissolve in excess NaOH (amphoteric). Fe(OH)₂, Fe(OH)₃, and Cu(OH)₂ do NOT dissolve in excess NaOH.

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Precipitation Reactions with Ammonia (NH₃)

When ammonia solution is added, it first acts as a base (forming the hydroxide precipitate), then in excess it may act as a ligand (dissolving the precipitate by forming an ammine complex).

Table of Observations