Qualitative analysis involves performing simple test-tube reactions to observe precipitation, gas evolution, or characteristic flame emissions. This guide details how to identify key metal cations and non-metal anions following the AQA specification.
Core Specification Link
This practical supports topic 3.2.3 Group 2, 3.2.4 Group 7 (Halogens), and 3.2.6 Transition Metals of inorganic chemistry.
Aim
To carry out simple test-tube reactions to identify the cations and anions in a set of unknown solutions (including Group 2 metals, ammonium ions, halides, carbonates, sulfates, and hydroxides).
Equipment List
- Nichrome wire loop
- Concentrated hydrochloric acid (\(\text{HCl}\))
- Unknown solution samples (containing \(\text{Group 2}\) or ammonium cations, and halide or other anions)
- Bunsen burner
- Test tubes and test tube rack
- Dropping pipettes
- Sodium hydroxide solution (\(\text{NaOH}\), 1.0 mol dm⁻³)
- Dilute nitric acid (\(\text{HNO}_3\), 1.0 mol dm⁻³)
- Silver nitrate solution (\(\text{AgNO}_3\), 0.05 mol dm⁻³)
- Dilute and concentrated ammonia solutions (\(\text{NH}_3\))
- Barium chloride solution (\(\text{BaCl}_2\), 0.1 mol dm⁻³)
- Red litmus paper or pH indicator paper
- Water bath (beaker of hot water or electronic water bath)
Part 1: Identifying Cations (Flame & NaOH Tests)
1. Flame Tests
Metal ions exhibit unique emissions when heated. Electrons absorb heat and are promoted to higher energy levels. When they relax back to their ground state, they emit light of specific wavelengths within the visible spectrum.
- Clean the nichrome wire loop by dipping it in concentrated hydrochloric acid and holding it in the hottest part of the blue Bunsen burner flame . Repeat until the wire produces no colour in the flame.
- Dip the clean wire into the sample solution (or touch it to a tiny amount of the solid sample).
- Hold the wire loop in the edge of the blue Bunsen flame and record the colour observed.
2. Sodium Hydroxide Tests for Metal Aquajoin Cations
Transition metals and group 3 elements form metal hydroxide precipitates in the presence of hydroxide ions (\(\text{OH}^-\)). Some, like aluminium, are amphoteric and redissolve in excess alkali.
- Place 1 cm³ of the sample solution into a clean test tube.
- Add dilute sodium hydroxide dropwise, observing and recording any precipitate that forms.
- Continue adding sodium hydroxide solution until it is in excess, and observe if the precipitate redissolves.
| Cation | Equation for Precipitate Formation | Precipitate Colour | With Excess NaOH |
|---|---|---|---|
| \(\text{Cu}^{2+}(\text{aq})\) | \(\text{Cu}^{2+}(\text{aq}) + 2\text{OH}^-(\text{aq}) \rightarrow \text{Cu(OH)}_2(\text{s})\) | Blue | Insoluble (remains blue precipitate) |
| \(\text{Fe}^{2+}(\text{aq})\) | \(\text{Fe}^{2+}(\text{aq}) + 2\text{OH}^-(\text{aq}) \rightarrow \text{Fe(OH)}_2(\text{s})\) | Green (darkens on air contact) | Insoluble (remains green precipitate) |
| \(\text{Fe}^{3+}(\text{aq})\) | \(\text{Fe}^{3+}(\text{aq}) + 3\text{OH}^-(\text{aq}) \rightarrow \text{Fe(OH)}_3(\text{s})\) | Brown / Rust | Insoluble (remains brown precipitate) |
| \(\text{Al}^{3+}(\text{aq})\) | \(\text{Al}^{3+}(\text{aq}) + 3\text{OH}^-(\text{aq}) \rightarrow \text{Al(OH)}_3(\text{s})\) | White | Redissolves to form a colourless solution: \(\text{Al(OH)}_3(\text{s}) + \text{OH}^-(\text{aq}) \rightarrow [\text{Al(OH)}_4]^-(\text{aq})\) |
| \(\text{Ca}^{2+}(\text{aq})\) | \(\text{Ca}^{2+}(\text{aq}) + 2\text{OH}^-(\text{aq}) \rightarrow \text{Ca(OH)}_2(\text{s})\) | White | Insoluble (remains white precipitate) |
| \(\text{Mg}^{2+}(\text{aq})\) | \(\text{Mg}^{2+}(\text{aq}) + 2\text{OH}^-(\text{aq}) \rightarrow \text{Mg(OH)}_2(\text{s})\) | White | Insoluble (remains white precipitate) |
3. Test for Ammonium Ion (\(\text{NH}_4^+\))
Ammonium ions react with warm hydroxide ions to release volatile ammonia gas (\(\text{NH}_3\)):
\[ \text{NH}_4^+(\text{aq}) + \text{OH}^-(\text{aq}) \rightarrow \text{NH}_3(\text{g}) + \text{H}_2\text{O}(\text{l}) \]- Add dilute sodium hydroxide solution to the sample in a test tube.
- Warm the mixture gently in a water bath.
- Hold a piece of damp red litmus paper over the mouth of the test tube.
- Observation: The litmus paper turns blue due to alkaline ammonia gas dissolving in the dampness. Ammonia also has a characteristic sharp, pungent odour.
Part 2: Identifying Anions (Carbonate, Sulfate, & Halides)
| Anion | Reagents & Procedure | Observation | Ionic Equation |
|---|---|---|---|
| Carbonate (\(\text{CO}_3^{2-}\)) | Add dilute hydrochloric acid. Bubble any gas evolved through limewater. | Effervescence (fizzing). Limewater turns cloudy/milky. | \(\text{CO}_3^{2-}(\text{aq}) + 2\text{H}^+(\text{aq}) \rightarrow \text{CO}_2(\text{g}) + \text{H}_2\text{O}(\text{l})\) |
| Sulfate (\(\text{SO}_4^{2-}\)) | Add dilute hydrochloric acid (to clear carbonates), then add barium chloride solution (\(\text{BaCl}_2\)). | Thick white precipitate. | \(\text{Ba}^{2+}(\text{aq}) + \text{SO}_4^{2-}(\text{aq}) \rightarrow \text{BaSO}_4(\text{s})\) |
| Chloride (\(\text{Cl}^-\)) | Acidify with dilute nitric acid (\(\text{HNO}_3\)), then add silver nitrate solution (\(\text{AgNO}_3\)). Add dilute ammonia. | White precipitate. Redissolves in dilute ammonia. | \(\text{Ag}^+(\text{aq}) + \text{Cl}^-(\text{aq}) \rightarrow \text{AgCl}(\text{s})\) |
| Bromide (\(\text{Br}^-\)) | Acidify with dilute nitric acid, then add silver nitrate solution. Add concentrated ammonia. | Cream precipitate. Redissolves only in concentrated ammonia. | \(\text{Ag}^+(\text{aq}) + \text{Br}^-(\text{aq}) \rightarrow \text{AgBr}(\text{s})\) |
| Iodide (\(\text{I}^-\)) | Acidify with dilute nitric acid, then add silver nitrate solution. Add concentrated ammonia. | Yellow precipitate. Insoluble in both dilute and concentrated ammonia. | \(\text{Ag}^+(\text{aq}) + \text{I}^-(\text{aq}) \rightarrow \text{AgI}(\text{s})\) |
Step 1: Test for the ammonium ion using sodium hydroxide solution and heat
Add dilute sodium hydroxide solution to a sample of each solution and warm gently in a water bath. Hold damp red litmus paper over the mouth of each tube.
- The solution containing \(\text{NH}_4\text{Cl}\) will release a gas (ammonia) that turns the damp red litmus paper blue: \[ \text{NH}_4^+(\text{aq}) + \text{OH}^-( ext{aq}) \rightarrow \text{NH}_3(\text{g}) + \text{H}_2\text{O}(\text{l}) \]
- The other three solutions (\(\text{BaCl}_2\), \(\text{NaCl}\), and \(\text{Al}_2(\text{SO}_4)_3\)) will produce no alkaline gas.
Step 2: Add excess sodium hydroxide solution to the remaining three solutions
Add sodium hydroxide dropwise until in excess to fresh samples of the remaining three solutions:
- The solution containing \(\text{Al}_2(\text{SO}_4)_3\) will form a white precipitate of aluminium hydroxide, which dissolves in excess sodium hydroxide to form a colourless solution: \[ \text{Al}^{3+}(\text{aq}) + 3\text{OH}^-(\text{aq}) \rightarrow \text{Al(OH)}_3(\text{s}) \quad \text{(white ppt)} \] \[ \text{Al(OH)}_3(\text{s}) + \text{OH}^-(\text{aq}) \rightarrow [\text{Al(OH)}_4]^-(\text{aq}) \quad \text{(colourless solution)} \]
- The solutions containing \(\text{BaCl}_2\) and \(\text{NaCl}\) will show no change (sodium does not precipitate, and barium hydroxide is soluble at this concentration).
Step 3: Distinguish between the two remaining chloride solutions
To distinguish between the remaining \(\text{BaCl}_2\) and \(\text{NaCl}\) solutions:
- Add acidified silver nitrate to both. Both will form a white precipitate of silver chloride: \[ \text{Ag}^+(\text{aq}) + \text{Cl}^-(\text{aq}) \rightarrow \text{AgCl}(\text{s}) \quad \text{(white ppt)} \]
- To confirm \(\text{BaCl}_2\) from \(\text{NaCl}\), perform a flame test or add dilute sulfuric acid to a fresh sample. Barium ions (\(\text{Ba}^{2+}\)) will yield an apple-green flame and a white precipitate of barium sulfate: \[ \text{Ba}^{2+}(\text{aq}) + \text{SO}_4^{2-}(\text{aq}) \rightarrow \text{BaSO}_4(\text{s}) \quad \text{(white ppt)} \] Sodium ions (\(\text{Na}^+\)) will yield a persistent bright yellow flame and no precipitate with sulfuric acid.
Safety & Risk Assessment
| Hazard | Risk | Precaution |
|---|---|---|
| Barium chloride (\(\text{BaCl}_2\)) | Highly toxic by ingestion and acts as a skin irritant. | Wear protective nitrile gloves. Wash hands immediately after use. Dispose of waste in a designated heavy metal residue bottle. |
| Silver nitrate (\(\text{AgNO}_3\)) | Corrosive and causes permanent brown-black stains on skin and clothing. | Wear gloves and a lab coat. Handle with care to prevent spills. |
| Concentrated Hydrochloric acid | Corrosive; releases toxic, choking hydrogen chloride gas. | Clean the wire in a fume cupboard. Wear safety goggles and gloves. |
Sources of Error & Improvements
- Contamination in flame tests: Sodium traces are extremely common and emit an intense yellow light that easily masks other ions. Improvement: Dip the wire loop in concentrated acid and burn off repeatedly until zero background colour remains.
- Halide false positives: If the sample contains carbonate (\(\text{CO}_3^{2-}\)) or sulfate (\(\text{SO}_4^{2-}\)) ions, adding silver nitrate without acidifying will yield white precipitates of \(\text{Ag}_2\text{CO}_3\) or \(\text{Ag}_2\text{SO}_4\). Improvement: Always acidify the sample with dilute nitric acid (\(\text{HNO}_3\)) first. The nitric acid destroys carbonate ions by converting them to carbon dioxide gas. Never use hydrochloric acid (\(\text{HCl}\)) to acidify, as the chloride ions would react with silver ions and yield a false positive white precipitate.
Common Exam Questions
1. State why sulfuric acid cannot be used to acidify a sample prior to testing for chloride ions using silver nitrate.
Show Answer
Sulfuric acid contains sulfate ions (\(\text{SO}_4^{2-}\)). Barium ions in the sample or silver ions in the reagent would combine with sulfate to form a white precipitate of silver sulfate, interfering with the chloride observation.
2. Explain why aluminium hydroxide dissolves in excess sodium hydroxide, but magnesium hydroxide does not.
Show Answer
Aluminium hydroxide is amphoteric. It reacts with excess hydroxide ions to form a soluble tetrahydroxoaluminate complex, \([\text{Al(OH)}_4]^-\). Magnesium hydroxide is basic and does not react with excess hydroxide ions, so it remains insoluble.
3. Describe how you would distinguish between aqueous solutions of potassium chloride and potassium iodide using chemical tests.
Show Answer
Add dilute nitric acid to both samples, followed by silver nitrate solution. Potassium chloride will yield a white precipitate (\(\text{AgCl}\)) that dissolves upon adding dilute ammonia. Potassium iodide will yield a yellow precipitate (\(\text{AgI}\)) that does not dissolve in dilute or concentrated ammonia.
CPAC Skills Assessed
- CPAC 1: Correctly follows written procedures to identify multiple unknown ions.
- CPAC 3: Safely handles toxic barium and staining silver nitrate solutions.
- CPAC 4: Accurately observes and records precipitate colour changes and gas evolution.
When asked to describe the halide test, you must state that you acidify with nitric acid (\(\text{HNO}_3\)). Simply stating "acidify" or using "HCl" or "H₂SO₄" will lose marks. Also, distinguish halide precipitate colours carefully: white (Cl⁻), cream (Br⁻), yellow (I⁻).