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 summarizes the procedures, equations, and colours needed to identify these ions.
Core Specification Link
This practical links directly to topic 3.2.5 Transition metals, covering the properties, complexes, precipitation, and amphoteric behavior of transition metal ions.
Aim
To carry out simple test-tube reactions to observe the reactions and identify the transition metal ions in aqueous solution (specifically \(\text{Cu}^{2+}\), \(\text{Fe}^{2+}\), \(\text{Fe}^{3+}\), \(\text{Al}^{3+}\) / \(\text{Cr}^{3+}\)).
Equipment List
- Test tubes and test tube rack
- Dropping pipettes
- Aqueous transition metal ion solutions: copper(II) sulfate (\(\text{CuSO}_4\)), iron(II) sulfate (\(\text{FeSO}_4\)), iron(III) chloride (\(\text{FeCl}_3\)), aluminum(III) sulfate (\(\text{Al}_2(\text{SO}_4)_3\)) or chromium(III) chloride (\(\text{CrCl}_3\))
- Sodium hydroxide solution (\(\text{NaOH}\), 1.0 mol dm⁻³)
- Ammonia solution (\(\text{NH}_3\), 1.0 mol dm⁻³)
- Sodium carbonate solution (\(\text{Na}_2\text{CO}_3\), 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 (\(\text{NaOH}\))
- Add 10 drops of the hexaaqua metal ion solution into a clean test tube.
- Add dilute sodium hydroxide solution dropwise . Record the initial colour of any precipitate formed.
- 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 (\(\text{NH}_3\))
- Add 10 drops of the hexaaqua metal ion solution into a clean test tube.
- Add dilute ammonia solution dropwise . Record the initial precipitate colour.
- Continue to add ammonia until in excess . Shake the tube and observe whether the precipitate dissolves.
Summary of Precipitates, Excess Reactions, and Equations
| Metal Cation | Starting Hexaaqua Ion | Add drops of base (\(\text{OH}^-\) or \(\text{NH}_3\)) | Adding excess \(\text{NaOH}\) | Adding excess \(\text{NH}_3\) |
|---|---|---|---|---|
| \(\text{Cu}^{2+}\) | \([\text{Cu(H}_2\text{O)}_6]^{2+}(\text{aq})\) (blue) |
\(\text{Cu(OH)}_2(\text{H}_2\text{O)}_4(\text{s})\) Blue precipitate |
Insoluble. Remains blue precipitate. |
Dissolves to form a deep blue solution. \([\text{Cu(NH}_3)_4(\text{H}_2\text{O)}_2]^{2+}\) |
| \(\text{Fe}^{2+}\) | \([\text{Fe(H}_2\text{O)}_6]^{2+}(\text{aq})\) (pale green) |
\(\text{Fe(OH)}_2(\text{H}_2\text{O)}_4(\text{s})\) Green precipitate |
Insoluble. Precipitate slowly darkens to brown at surface in air. |
Insoluble. Precipitate slowly darkens to brown. |
| \(\text{Fe}^{3+}\) | \([\text{Fe(H}_2\text{O)}_6]^{3+}(\text{aq})\) (yellow-brown) |
\(\text{Fe(OH)}_3(\text{H}_2\text{O)}_3(\text{s})\) Rust-brown precipitate |
Insoluble. Remains brown precipitate. |
Insoluble. Remains brown precipitate. |
| \(\text{Al}^{3+}\) | \([\text{Al(H}_2\text{O)}_6]^{3+}(\text{aq})\) (colourless) |
\(\text{Al(OH)}_3(\text{H}_2\text{O)}_3(\text{s})\) White precipitate |
Dissolves to form a colourless solution. \([\text{Al(OH)}_4]^-\)(aq) (amphoteric) |
Insoluble. Remains white precipitate. |
| \(\text{Cr}^{3+}\) | \([\text{Cr(H}_2\text{O)}_6]^{3+}(\text{aq})\) (violet-green) |
\(\text{Cr(OH)}_3(\text{H}_2\text{O)}_3(\text{s})\) Grey-green precipitate |
Dissolves to form a dark green solution. \([\text{Cr(OH)}_6]^{3-}\) (amphoteric) |
Insoluble. Remains grey-green precipitate. |
| \(\text{Mn}^{2+}\) | \([\text{Mn(H}_2\text{O)}_6]^{2+}(\text{aq})\) (very pale pink) |
\(\text{Mn(OH)}_2(\text{H}_2\text{O)}_4(\text{s})\) Buff/cream precipitate |
Insoluble. Rapidly darkens to brown in air. |
Insoluble. Rapidly darkens to brown in air. |
| \(\text{Co}^{2+}\) | \([\text{Co(H}_2\text{O)}_6]^{2+}(\text{aq})\) (pink) |
\(\text{Co(OH)}_2(\text{H}_2\text{O)}_4(\text{s})\) Blue-green precipitate |
Insoluble. Precipitate darkens/turns pink. |
Dissolves to form a straw-yellow solution. \([\text{Co(NH}_3)_6]^{2+}\) |
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 center. Because neutral complexes carry no charge, they precipitate out of aqueous solution:
\[ [\text{Cu(H}_2\text{O)}_6]^{2+}(\text{aq}) + 2\text{OH}^-(\text{aq}) \rightarrow \text{Cu(OH)}_2(\text{H}_2\text{O)}_4(\text{s}) + 2\text{H}_2\text{O}(\text{l}) \] \[ [\text{Fe(H}_2\text{O)}_6]^{3+}(\text{aq}) + 3\text{NH}_3(\text{aq}) \rightarrow \text{Fe(OH)}_3(\text{H}_2\text{O)}_3(\text{s}) + 3\text{NH}_4^+(\text{aq}) \]2. Amphoteric Behavior of Chromium(III)
Chromium(III) hydroxide is amphoteric. In excess sodium hydroxide (a strong base), the neutral precipitate undergoes further deprotonation to form a soluble negatively charged hexahydroxochromate(III) complex:
\[ \text{Cr(OH)}_3(\text{H}_2\text{O)}_3(\text{s}) + 3\text{OH}^-(\text{aq}) \rightarrow [\text{Cr(OH)}_6]^{3-}(\text{aq}) + 3\text{H}_2\text{O}(\text{l}) \]3. Ligand Exchange with Excess Ammonia
Ammonia is a nucleophile and can displace water ligands. With excess ammonia, the copper(II) and cobalt(II) precipitates dissolve via a ligand exchange reaction:
\[ \text{Cu(OH)}_2(\text{H}_2\text{O)}_4(\text{s}) + 4\text{NH}_3(\text{aq}) \rightarrow [\text{Cu(NH}_3)_4(\text{H}_2\text{O)}_2]^{2+}(\text{aq}) + 2\text{OH}^-(\text{aq}) + 2\text{H}_2\text{O}(\text{l}) \]4. Reactions with Sodium Carbonate (\(\text{Na}_2\text{CO}_3\))
Aqueous metal ions react differently with sodium carbonate depending on their charge density:
- Metal(II) ions (\(\text{M}^{2+}\)): The lower charge density of \(2+\) ions means they are not sufficiently acidic to react with carbonate bases. Instead, they undergo simple precipitation to form insoluble metal carbonates: \[ [\text{M(H}_2\text{O)}_6]^{2+}(\text{aq}) + \text{CO}_3^{2-}(\text{aq}) \rightarrow \text{MCO}_3(\text{s}) + 6\text{H}_2\text{O}(\text{l}) \] For example, copper(II) forms a blue-green precipitate of copper(II) carbonate (\(\text{CuCO}_3\)).
- Metal(III) ions (\(\text{M}^{3+}\)): Due to their high charge density, \(3+\) ions polarise water ligands strongly, releasing \(H^+\) ions. They are sufficiently acidic to react with the carbonate base, undergoing an acid-base reaction that yields a neutral metal hydroxide precipitate and carbon dioxide gas (bubbles): \[ 2[\text{M(H}_2\text{O)}_6]^{3+}(\text{aq}) + 3\text{CO}_3^{2-}(\text{aq}) \rightarrow 2\text{M(OH)}_3(\text{H}_2\text{O)}_3(\text{s}) + 3\text{CO}_2(\text{g}) + 3\text{H}_2\text{O}(\text{l}) \] For example, iron(III) forms a brown precipitate of \(\text{Fe(OH)}_3(\text{H}_2\text{O)}_3\) and bubbles of \(\text{CO}_2\). Also, aluminium(III) forms a white precipitate of \(\text{Al(OH)}_3(\text{H}_2\text{O)}_3\) and bubbles of \(\text{CO}_2\).
- Adding a few drops of aqueous sodium hydroxide forms a green precipitate.
- Leaving the green precipitate to stand in air causes the surface to turn brown.
- Adding excess aqueous ammonia to a fresh sample of the starting solution also forms a green precipitate that does not dissolve.
Step 1: Identify the metal cation and starting complex
A green precipitate that is insoluble in excess sodium hydroxide points to iron(II) ions, \(\text{Fe}^{2+}\). Cobalt(II) forms a blue-green precipitate but dissolves in excess ammonia. Chromium(III) forms a grey-green precipitate but dissolves in excess sodium hydroxide. The starting hexaaqua complex formula is: \[ [\text{Fe(H}_2\text{O)}_6]^{2+}(\text{aq}) \]
Step 2: Write the ionic equation for the reaction with sodium hydroxide (Test 1)
Hydroxide ions deprotonate the hexaaqua complex to form the neutral, insoluble iron(II) hydroxide precipitate: \[ [\text{Fe(H}_2\text{O)}_6]^{2+}(\text{aq}) + 2\text{OH}^-(\text{aq}) \rightarrow \text{Fe(OH)}_2(\text{H}_2\text{O)}_4(\text{s}) + 2\text{H}_2\text{O}(\text{l}) \]
Step 3: Explain the change on standing in air (Test 2)
Oxygen in the air oxidises the iron(II) hydroxide (\(\text{Fe}^{2+}\)) to iron(III) hydroxide (\(\text{Fe}^{3+}\)), which has a rust-brown colour: \[ 4\text{Fe(OH)}_2(\text{H}_2\text{O)}_4(\text{s}) + \text{O}_2(\text{g}) \rightarrow 4\text{Fe(OH)}_3(\text{H}_2\text{O)}_3(\text{s}) + 2\text{H}_2\text{O}(\text{l}) \]
Safety & Risk Assessment
| Hazard | Risk | Precaution |
|---|---|---|
| 2.0 mol dm⁻³ Sodium Hydroxide | Corrosive; causes skin burns and severe eye damage. | Wear safety goggles and gloves. Wash splashes off skin immediately. |
| Ammonia solution | Corrosive and has a pungent, irritating vapour. | Handle in a fume cupboard or well-ventilated space. Keep sample bottles stoppered. |
| Cobalt(II) sulfate solution | May cause cancer if inhaled; toxic to aquatic life. | Wear gloves; handle minimum quantities; dispose of in the heavy metals waste container. |
Sources of Error & Improvements
- Oxidation of Fe(II) and Mn(II) precipitates: Iron(II) and manganese(II) hydroxides 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.
Common Exam Questions
1. Write a balanced equation for the reaction of chromium(III) hydroxide precipitate with excess sodium hydroxide.
Show Answer
\[ \text{Cr(OH)}_3(\text{H}_2\text{O)}_3(\text{s}) + 3\text{OH}^-(\text{aq}) \rightarrow [\text{Cr(OH)}_6]^{3-}(\text{aq}) + 3\text{H}_2\text{O}(\text{l}) \]
2. Explain why the green precipitate of Fe(II) hydroxide slowly turns brown at the top of the test tube when left standing in air.
Show Answer
The green precipitate of iron(II) hydroxide, \(\text{Fe(OH)}_2(\text{H}_2\text{O)}_4\), is oxidised by oxygen in the air to form the rust-brown iron(III) hydroxide, \(\text{Fe(OH)}_3(\text{H}_2\text{O)}_3\).
3. State the formula of the complex ion responsible for the deep blue solution formed when excess ammonia is added to a solution of copper(II) ions.
Show Answer
\[ [\text{Cu(NH}_3)_4(\text{H}_2\text{O)}_2]^{2+} \]
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 colors and solubility in excess.
Always distinguish clearly between precipitate colours. Copper(II) forms a blue precipitate, iron(II) forms a green precipitate, and chromium(III) forms a grey-green precipitate. Mixing up green and grey-green will lose marks.