8.2.8 Identifying Ions

Test for Group 2 Ions Using Sodium Hydroxide

Aim:

To carry out simple test-tube reactions to identify Group 2 cations using sodium hydroxide and observe the formation of precipitates for different Group 2 ions ($Ba²⁺$, $Ca²⁺$, $Mg²⁺$, $Sr²⁺$).

Introduction:

Group 2 metal ions ($Ba²⁺$, $Ca²⁺$, $Mg²⁺$, $Sr²⁺$) can be identified using their reaction with sodium hydroxide ($NaOH$). Upon the addition of NaOH, these cations react to form characteristic precipitates. By carefully observing the precipitates formed at different stages of the reaction, we can determine the identity of the metal ion present in the solution.

Materials and Equipment:

• 0.1 mol dm⁻³ barium chloride solution
• 0.1 mol dm⁻³ calcium bromide solution
• 0.1 mol dm⁻³ magnesium chloride solution
• 0.1 mol dm⁻³ strontium chloride solution
• 0.6 mol dm⁻³ sodium hydroxide solution
• Test tubes
• Test tube rack
• Dropping pipette

Method:

Step 1: Initial Preparation

1. Clean the test tubes thoroughly to avoid contamination, which could interfere with the results.
2. Label the test tubes for each solution to prevent any confusion when observing results.

Step 2: Adding Group 2 Ions

1. Place 10 drops of 0.1 mol dm⁻³ barium chloride solution into a clean test tube. This serves as the test for $Ba²⁺$ ions.
2. Repeat the procedure for:

• Calcium bromide ($Ca²⁺$)
• Magnesium chloride ($Mg²⁺$)
• Strontium chloride ($Sr²⁺$)

Step 3: Addition of Sodium Hydroxide

1. Add 10 drops of 0.6 mol dm⁻³ sodium hydroxide solution to each test tube. Mix well by swirling and record any immediate observations.

• Note: The formation of a precipitate indicates a reaction between the sodium hydroxide and the Group 2 ion.

2. Continue to add sodium hydroxide dropwise to each test tube, shaking gently until in excess. This allows you to observe if more precipitate forms or if any precipitate dissolves.

Expected Results:

Group 2 Ion	Initial (Before NaOH)	After 10 Drops of NaOH	Excess NaOH
$Ba²⁺$	Colourless solution	Colourless solution	Colourless solution
$Ca²⁺$	Colourless solution	Slight white precipitate	Slight white precipitate
$Mg²⁺$	Colourless solution	Slight white precipitate	White precipitate
$Sr²⁺$	Colourless solution	Slight white precipitate	Slight white precipitate

Explanation of Observations:

• Barium ions ($Ba²⁺$): No precipitate is formed when sodium hydroxide is added to barium chloride. The solution remains clear, indicating that barium hydroxide is soluble in water.
• Calcium ions $(Ca²⁺$): Upon the addition of $NaOH$, a slight white precipitate of calcium hydroxide forms. This precipitate remains even in excess sodium hydroxide, indicating limited solubility.
• Magnesium ions ($Mg²⁺$): A white precipitate of magnesium hydroxide forms upon the addition of sodium hydroxide. The precipitate becomes more pronounced with excess $NaOH$, indicating that magnesium hydroxide is sparingly soluble in water.
• Strontium ions ($Sr²⁺$): Like calcium, a slight white precipitate of strontium hydroxide forms, which remains in excess sodium hydroxide. This suggests that strontium hydroxide has a similar solubility to calcium hydroxide.

Test for Group 2 Ions Using Dilute Sulfuric Acid

Aim:

To carry out test-tube reactions to identify Group 2 cations using dilute sulfuric acid ($H₂SO₄$) and observe the formation of precipitates for different Group 2 ions ($Ba²⁺$, $Ca²⁺$, $Mg²⁺$, $Sr²⁺$).

Introduction:

Group 2 metal ions $(Ba²⁺, Ca²⁺, Mg²⁺, Sr²⁺$) react differently with dilute sulfuric acid, forming varying degrees of precipitates. These reactions can be used to distinguish between the cations of Group 2 elements based on the solubility of their sulfates.

• Barium sulfate ($BaSO₄$) and strontium sulfate ($SrSO₄$) are insoluble and form white precipitates.
• Calcium sulfate ($CaSO₄$) forms a slight white precipitate due to partial solubility.
• Magnesium sulfate ($MgSO₄$) is soluble, producing a colourless solution.

Materials and Equipment:

• 0.1 mol dm⁻³ barium chloride solution ($Ba²⁺$)
• 0.1 mol dm⁻³ calcium bromide solution ($Ca²⁺$)
• 0.1 mol dm⁻³ magnesium chloride solution ($Mg²⁺$)
• 0.1 mol dm⁻³ strontium chloride solution ($Sr²⁺$)
• 1.0 mol dm⁻³ dilute sulfuric acid ($H₂SO₄$)
• Test tubes
• Test tube rack
• Dropping pipette
• Safety gloves (since barium salts are toxic)

Method:

Step 1: Initial Preparation

1. Clean the test tubes thoroughly to avoid contamination, which could interfere with the results.
2. Label the test tubes for each solution to prevent any confusion when observing results.

Step 2: Adding Group 2 Ions

1. Place 10 drops of 0.1 mol dm⁻³ barium chloride solution into a clean test tube.
2. Repeat this procedure for:

• Calcium bromide ($Ca²⁺$)
• Magnesium chloride ($Mg²⁺$)
• Strontium chloride ($Sr²⁺$)

Step 3: Addition of Sulfuric Acid

1. Add 10 drops of 1.0 mol dm⁻³ sulfuric acid to each test tube. Mix well by swirling the solution and observe the formation of any precipitate.
2. Continue adding sulfuric acid dropwise to each test tube with gentle shaking until the solution is in excess. Record any further changes in the precipitate or solution.

Step 4: Observation

1. After the reactions are complete, dispose of the contents of each test tube into a bowl of water for safe disposal.
2. Repeat the experiment for all Group 2 ions ($Ba²⁺, Ca²⁺, Mg²⁺, Sr²⁺$).

Expected Results:

Group 2 Ion	10 drops of 1.0 mol dm⁻³ $H₂SO₄$	Excess $H₂SO₄$
$Ba²⁺$(Barium)	White precipitate ($BaSO₄$)	White precipitate remains
$Ca²⁺ ($Calcium)	Slight white precipitate ($CaSO₄$)	Slight white precipitate remains
$Mg²⁺$(Magnesium)	Slight white precipitate	Colourless solution (soluble $MgSO₄$)
$Sr²⁺$(Strontium)	White precipitate ($SrSO₄$)	White precipitate remains

Explanation of Results:

• Barium ions ($Ba²⁺$: A white precipitate of barium sulfate forms immediately, which is insoluble in excess sulfuric acid.
• Calcium ions ($Ca²⁺$): A slight white precipitate of calcium sulfate forms, indicating its partial solubility. The precipitate remains in excess sulfuric acid.
• Magnesium ions ($Mg²⁺$): Initially, a slight precipitate may form, but upon adding excess sulfuric acid, the solution turns colourless because magnesium sulfate is soluble in water.
• Strontium ions ($Sr²⁺$): A white precipitate of strontium sulfate forms and remains even with excess sulfuric acid, indicating low solubility similar to barium sulfate.

Overall Group 2 Results Summary:

Group 2 Ion	Ammonium Solution	Excess NaOH	Excess H₂SO₄
$Mg²⁺$(Magnesium)	White precipitate of $Mg(OH)₂$	White precipitate of $Mg(OH)₂$	Colourless solution (soluble $MgSO₄$)
$Ca²⁺$(Calcium)	No change	White precipitate of $Ca(OH)₂$	Slight white precipitate of $CaSO₄$
$Sr²⁺$(Strontium)	No change	Slight white precipitate of $Sr(OH)₂$	White precipitate of $SrSO₄$
$Ba²⁺$(Barium)	No change	No change (no precipitate)	White precipitate of $BaSO₄$

Safety Precautions:

• Barium salts are toxic, so wear gloves while handling them and take care to avoid direct contact with your skin.
• Always dispose of chemicals safely, ensuring that test tubes are rinsed and cleaned after the reactions.

Test for Ammonium Ions (NH₄⁺)

Method:

3. Place 10 drops of ammonium chloride into a clean test tube, then add 10 drops of sodium hydroxide. Shake to mix.
4. Gently warm the test tube in a water bath to release any ammonia gas.
5. Test the fumes by holding damp red litmus paper at the mouth of the test tube. Record observations.
6. Dispose of the test tube contents.
7. Record that the damp red litmus paper turns blue. Expected Results: The litmus paper will turn blue, indicating the presence of ammonia gas.

Explanation: Ammonium ions react with sodium hydroxide, releasing ammonia gas, which is alkaline and turns red litmus paper blue.

Test for Hydroxide Ions (OH⁻): Aqueous Solution

Method:

8. Add 1 cm depth of the test solution to a test tube.
9. Use red litmus paper or universal indicator paper to test the solution.
10. Record any observations, then dispose of the contents. Expected Result:

• Hydroxide ions will turn damp red litmus paper blue, indicating an alkaline solution.

Test for Hydroxide Ions (OH⁻): Ammonia Solution

Method:

11. Place 5 drops of 1.0 mol dm⁻³ ammonia solution onto filter paper and place inside a petri dish with a lid.
12. On the other side of the petri dish, place damp red litmus paper.
13. Close the petri dish and observe over a few minutes. Expected Result:

• Ammonia vapours will turn the damp red litmus paper blue, indicating that hydroxide ions are present due to ammonia's basic nature.

Test for Carbonate Ions (CO₃²⁻): Aqueous Solution

Method:

14. Add a small amount of dilute hydrochloric acid to the sodium carbonate solution in a test tube.
15. Use a delivery tube to transfer the gas produced into a second test tube containing a small amount of calcium hydroxide solution (limewater).
16. Put a stopper in the test tube containing the calcium hydroxide solution and shake it gently from side to side. Expected Result:

• Limewater turns cloudy if carbonate ions are present due to the formation of carbon dioxide ($CO₂$). Equations:

• $Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂$

• $Ca(OH)₂ + CO₂ → CaCO₃ + H₂O$

Test for Halide Ions: Solid Salts

Method:

17. Place a small amount of solid potassium chloride in a dry test tube.
18. Add a few drops of concentrated sulfuric acid and observe.
19. Test any evolved gas using moist blue litmus paper or universal indicator.
20. Repeat the test with solid potassium bromide and solid potassium iodide. Expected Results:

• KCl: White, steamy fumes; litmus paper turns red.
• KBr: Orange fumes; litmus paper turns red.
• KI: Purple fumes and solid; litmus paper turns red.

Safety Precautions:

• Barium chloride is harmful; wear gloves and eye protection.
• Concentrated ammonia is corrosive and should be handled in a fume cupboard.
• Dilute nitric acid and sulfuric acid are irritants, so proper lab attire, including gloves and goggles, should be worn at all times.

Test for Sulfate Ions ($SO₄²⁻$) in Aqueous Solution

Objective:

To identify the presence of sulfate ions in a solution using barium chloride ($BaCl₂$) and dilute hydrochloric acid ($HC$l).

Method

1. Preparation:

• Add an equal volume of dilute hydrochloric acid (HCl) to the test solution containing the suspected sulfate ions in a test tube.

2. Barium Chloride Addition:

• Add an equal volume of barium chloride ($BaCl₂$) solution to the test tube.
• Observe any immediate formation of a precipitate.

3. Observation:

• If a white precipitate forms, it indicates the presence of barium sulfate ($BaSO₄$), confirming the presence of sulfate ions in the solution.

4. Verification:

• To confirm the sulfate ions, add a small volume of dilute $HCl$ to the mixture.
• Observe whether the precipitate dissolves. If the precipitate remains insoluble, it is likely barium sulfate, confirming the presence of sulfate ions ($SO₄²⁻$).

Expected Results:

• Formation of Barium Sulfate: A white precipitate forms if sulfate ions are present. This precipitate is barium sulfate ($BaSO₄$), an insoluble salt.

Equations:

Formation of Barium Sulfate:

$$\text{BaCl}_2 + \text{SO}_4^{2-} \rightarrow \text{BaSO}_4(s) + 2\text{Cl}^-$$

Complete Reaction Example:

$$\text{MgSO}_4(aq) + \text{BaCl}_2(aq) \rightarrow \text{BaSO}_4(s) + \text{MgCl}_2(aq)$$

Accuracy:

• Use of HCl: The addition of dilute hydrochloric acid ensures that any carbonate ions ($CO₃²-$), which could also produce a white precipitate, are removed to avoid false positive results. This enhances the specificity of the test for sulfate ions.

Explanation:

• Insolubility of Barium Sulfate: Barium sulfate is an insoluble salt, and its formation in the presence of sulfate ions provides a clear indication of sulfate in the solution.

Safety Precautions:

• Barium Chloride Hazard: Barium chloride is harmful, so gloves and a lab coat should be worn during the experiment.

Test for Halide Ions: Aqueous Solution

Method:

1. Prepare the Solution:

• Add a small volume of dilute nitric acid to the solution of potassium chloride.
• Accuracy Tip: This step removes any ions that might form a different precipitate.

2. Addition of Silver Nitrate:

• Add 2 cm³ of silver nitrate to the solution. Record any observations.

3. Swirling the Tubes:

• Swirl the tubes to ensure that the precipitates are evenly distributed.
• Divide the contents of each tube in half.
• Accuracy Tip: This allows for further identification tests as the initial precipitates can be difficult to distinguish.

4. Test with Aqueous Ammonia (Dilute):

• To one half of the contents, add an excess of dilute aqueous ammonia and observe any changes. Record observations.

5. Test with Concentrated Ammonia:

• To the other half, working in a fume cupboard, add excess concentrated ammonia solution. Record observations.

6. Repeat with Other Halides:

• Repeat the entire process with potassium bromide and potassium iodide in new, separate test tubes.

Results:

Substance	Silver Nitrate	Dilute Ammonia	Concentrated Ammonia
KCl	White precipitate	Colourless solution	Colourless solution
KBr	Cream precipitate	Cream precipitate	Colourless solution
KI	Yellow precipitate	Yellow precipitate	Yellow precipitate

Explanation:

• Potassium Chloride (KCl): Forms a white precipitate with silver nitrate. This precipitate dissolves in both dilute and concentrated ammonia, confirming the presence of chloride ions.
• Potassium Bromide (KBr): Forms a cream precipitate with silver nitrate, which only dissolves in concentrated ammonia, confirming bromide ions.
• Potassium Iodide (KI): Forms a yellow precipitate with silver nitrate, which does not dissolve in ammonia, indicating iodide ions.

Safety Precautions:

• Concentrated Ammonia is corrosive; wear chemical splash-proof eye protection and nitrile gloves.
• Always use a fume cupboard for working with ammonia.
• Dilute Nitric Acid is an irritant; handle with care.

Test for Group 7 Ions: Solid Salts

Method:

1. Preparation of Sample:

• Place a small spatula measure of solid potassium chloride into a clean, dry test tube in a test tube rack.

2. Addition of Concentrated Sulfuric Acid:

• In a fume cupboard, add a few drops of concentrated sulfuric acid to the test tube. Record any observations.

3. Testing of Evolved Gases:

• Test any gas that is produced using moist blue litmus paper or universal indicator paper. Record any colour changes observed.

4. Repeat the Experiment:

• Repeat the above procedure using solid potassium bromide and solid potassium iodide, and record any observations.

Observations and Results:

Substance	Concentrated Sulfuric Acid	Blue Litmus Paper
KCl	White, steamy fumes	Turns red
KBr	Orange fumes	Turns red
KI	Purple fumes, with purple/black solid	Turns red

Safety Precautions:

• Concentrated sulfuric acid is corrosive and must be handled carefully.
• The gases produced during this test are both toxic and corrosive. Ensure that the experiment is carried out in a fume cupboard.
• Wear appropriate personal protective equipment such as chemical splash-proof goggles and nitrile gloves to protect your skin and eyes from harmful chemicals.

Test for Halide Ions: Aqueous Solution

Method:

1. Preparation of Solution:

• Add a small volume of dilute nitric acid to the solution of potassium chloride.
• Reason: This step removes any ions that might interfere by forming a different precipitate.

2. Addition of Silver Nitrate:

• Add 2 cm³ of silver nitrate solution to the solution. Record your observations.

3. Swirl and Split:

• Swirl the tube to evenly distribute the precipitate, then divide the solution in half.

4. Test with Dilute Ammonia:

• To one half, add dilute aqueous ammonia and observe any changes. Record the results.

5. Test with Concentrated Ammonia:

• To the second half, add concentrated ammonia solution in a fume cupboard and record your observations.

6. Repeat for Other Halides:

• Repeat the test using solutions of potassium bromide and potassium iodide.

Observations and Results:

Substance	Silver Nitrate	Dilute Ammonia	Concentrated Ammonia
KCl	White precipitate	Colourless solution	Colourless solution
KBr	Cream precipitate	Cream precipitate	Colourless solution
KI	Yellow precipitate	Yellow precipitate	Yellow precipitate

Explanation:

• Potassium Chloride ($KCl$): The white precipitate formed by silver nitrate dissolves in both dilute and concentrated ammonia solutions, indicating the presence of chloride ions.
• Potassium Bromide ($KBr$): The cream precipitate only dissolves in concentrated ammonia, confirming bromide ions.
• Potassium Iodide ($KI$): The yellow precipitate does not dissolve in either dilute or concentrated ammonia, identifying iodide ions.

Safety:

• Concentrated ammonia is corrosive, so wear protective gear and use a fume cupboard.
• Dilute nitric acid is an irritant and should be handled carefully.