RP11: Reactions of Transition Metal Ions in Aqueous Solution

Transition metal ions form hydrated hexaaqua complex ions in aqueous solution. Because of their differing oxidation states and electron structures, they exhibit characteristic precipitation and ligand exchange reactions when treated with bases. This guide summarises the procedures, equations, and colours needed to identify these ions.

Aim

To carry out simple test-tube reactions to observe the reactions and identify the transition metal ions in aqueous solution (specifically Cu²⁺, Fe²⁺, Fe³⁺ and Al³⁺; Cr³⁺ is beyond the AQA specification).

Equipment List

• Test tubes and test tube rack
• Dropping pipettes
• Aqueous transition metal ion solutions: copper(II) sulfate (CuSO₄), iron(II) sulfate (FeSO₄), iron(III) chloride (FeCl₃), aluminium(III) sulfate (Al₂(SO₄)₃), and chromium(III) chloride (CrCl₃) for the tests beyond the AQA specification
• Sodium hydroxide solution (NaOH, 1.0 mol dm⁻³)
• Ammonia solution (NH₃, 1.0 mol dm⁻³)
• Sodium carbonate solution (Na₂CO₃, 0.5 mol dm⁻³)
• White tile (for observing colour changes clearly)

Precipitate Colours Reference

The diagram below displays the characteristic colours of precipitates formed when a few drops of sodium hydroxide or ammonia are added to the hexaaqua metal complexes:

Experimental Procedures

Part A: Reactions with Sodium Hydroxide (NaOH)

1. Add 10 drops of the hexaaqua metal ion solution into a clean test tube.
2. Add dilute sodium hydroxide solution dropwise. Record the initial colour of any precipitate formed.
3. Continue to add sodium hydroxide until in excess. Shake the tube and observe whether the precipitate dissolves or remains insoluble.

Part B: Reactions with Ammonia (NH₃)

1. Add 10 drops of the hexaaqua metal ion solution into a clean test tube.
2. Add dilute ammonia solution dropwise. Record the initial precipitate colour.
3. Continue to add ammonia until in excess. Shake the tube and observe whether the precipitate dissolves.

Summary of Precipitates, Excess Reactions, and Equations

Important Chemical Mechanisms

1. Deprotonation (Precipitation)

Both sodium hydroxide and ammonia act as bases (proton acceptors) to deprotonate the acidic water ligands attached to the metal centre. Because neutral complexes carry no charge, they precipitate out of aqueous solution:

2. Amphoteric Behaviour of Aluminium Hydroxide

Aluminium hydroxide is amphoteric: it dissolves in both acids and bases. In excess sodium hydroxide (a strong base), the neutral precipitate is deprotonated further to form a soluble, negatively charged complex. In acid, the water ligands are restored and the hexaaqua ion re-forms:

The complex formed in excess sodium hydroxide is often written more simply as [Al(OH)₄]⁻, as in the table above. Ammonia is too weak a base to remove the extra proton, so aluminium hydroxide does not dissolve in excess ammonia.

Beyond the AQA specification: chromium(III) hydroxide is also amphoteric. In excess sodium hydroxide it forms a soluble, dark green hexahydroxochromate(III) complex:

3. Ligand Exchange with Excess Ammonia

Ammonia is a nucleophile and can displace water ligands. With excess ammonia, the copper(II) hydroxide precipitate dissolves via a ligand exchange reaction to give a deep blue solution. (Cobalt(II) hydroxide also dissolves in excess ammonia, but cobalt(II) is beyond the AQA specification for this practical.)

4. Reactions with Sodium Carbonate (Na₂CO₃)

Aqueous metal ions react differently with sodium carbonate depending on the charge/size ratio of the metal ion:

• Metal(II) ions (M²⁺): A 2+ ion has a lower charge/size ratio, so [M(H₂O)₆]²⁺ is only weakly acidic. It is not acidic enough to react with carbonate ions as an acid, so the metal carbonate precipitates instead: \[ [\text{M(H}_2\text{O)}_6]^{2+}(\text{aq}) + \text{CO}_3^{2-}(\text{aq}) \rightarrow \text{MCO}_3(\text{s}) + 6\text{H}_2\text{O}(\text{l}) \] Copper(II) forms a blue-green precipitate of copper(II) carbonate (CuCO₃), and iron(II) forms a green precipitate of iron(II) carbonate (FeCO₃).
• Metal(III) ions (M³⁺): A 3+ ion has a higher charge/size ratio, so it polarises its water ligands more strongly and weakens their O–H bonds. [M(H₂O)₆]³⁺ is therefore acidic enough to react with carbonate ions: carbon dioxide is given off (effervescence) and the metal hydroxide precipitates: \[ 2[\text{M(H}_2\text{O)}_6]^{3+}(\text{aq}) + 3\text{CO}_3^{2-}(\text{aq}) \rightarrow 2\text{M(OH)}_3(\text{H}_2\text{O)}_3(\text{s}) + 3\text{CO}_2(\text{g}) + 3\text{H}_2\text{O}(\text{l}) \] For example, iron(III) forms a brown precipitate of Fe(OH)₃(H₂O)₃ and bubbles of CO₂. Also, aluminium(III) forms a white precipitate of Al(OH)₃(H₂O)₃ and bubbles of CO₂.

Why 3+ aqua ions are more acidic than 2+ aqua ions is explained in full in the metal hydroxide precipitates notes.

Safety & Risk Assessment

Sources of Error & Improvements

• Oxidation of Fe(II) and Mn(II) precipitates: Iron(II) and manganese(II) hydroxides (manganese(II) is beyond the AQA specification) react with dissolved oxygen in water and air, turning brown rapidly. This can confuse students identifying unknown samples. Improvement: Use freshly prepared solutions and record colours immediately after precipitation.
• Contaminated glassware: Trace metal contaminants in test tubes can yield wrong colours or trigger false precipitates. Improvement: Thoroughly clean test tubes with dilute nitric acid and rinse with deionised water before starting.

CPAC Skills Assessed

• CPAC 1: Follows multi-step written protocols to carry out sequential qualitative analysis.
• CPAC 3: Safely handles corrosive alkalis and heavy metal ion solutions.
• CPAC 4: Records detailed descriptions of precipitate colours and solubility in excess.