SODIUMTHIOSULPHATE-IODINE TITRATIONS

Objective

To determine the concentration of an iodine solution using standard sodium thiosulphate solution through redox titration with starch as an indicator.

Chemical Reaction

2S₂O₃^2- + I₂ → S₄O₆^2- + 2I^–

In this redox reaction:

Sodium thiosulphate (Na₂S₂O₃) acts as a reducing agent
Iodine (I₂) acts as an oxidizing agent
The endpoint is detected by the disappearance of the blue-black starch-iodine complex

Iodine Solution

I₂

Blue-black with starch

Sodium Thiosulphate

S₂O₃^2-

Colorless

→

Products

Colorless solution

Stoichiometry

The balanced chemical equation shows a 1:2 molar ratio between iodine and sodium thiosulphate:

1 mol I₂ : 2 mol Na₂S₂O₃

This ratio is crucial for calculating the concentration of the iodine solution from titration data.

Materials Required

Chemicals

Sodium thiosulphate (Na₂S₂O₃) solution (0.2M)
Iodine (I₂) solution (unknown concentration)
Starch indicator solution (1%)
Dilute sulphuric acid (H₂SO₄)
Potassium iodide (KI)
Distilled water

Apparatus

Burette (50 mL)
Pipette (20 mL)
Conical flask (250 mL)
White tile
Burette stand
Volumetric flask (1 L)
Analytical balance
Beakers and measuring cylinders

Experimental Procedure

Iodine + Starch (Blue-Black)

Starch-Iodine Complex

Endpoint (Colorless)

Rinse the burette with the 0.2M sodium thiosulphate solution and fill it with the same solution, ensuring no air bubbles are present.
Pipette 20 mL of the iodine solution into a clean conical flask.
Add 1-2 mL of starch indicator solution to the iodine solution. The solution will immediately turn blue-black due to the formation of starch-iodine complex.
Place the conical flask on a white tile to clearly observe the colour change during titration.
Slowly add sodium thiosulphate solution from the burette to the conical flask while swirling continuously.
The blue-black colour will gradually fade as iodine is reduced to colourless iodide ions.
As the endpoint approaches, the colour change becomes slower. Add the titrant dropwise at this stage.
The endpoint is reached when the blue-black colour completely disappears, leaving a colourless solution.
Record the initial and final readings of the burette.
Repeat the titration at least three times to obtain concordant results.

Observations

Sr. No.	Initial Reading (mL)	Final Reading (mL)	Volume of Na₂S₂O₃ Used (mL)
1	0.00	24.80	24.80
2	0.00	24.90	24.90
3	0.00	24.85	24.85
Mean Volume of Na₂S₂O₃			24.85 mL

Key Observations During Titration

Initial solution: Deep blue-black colour after adding starch indicator
During titration: Gradual fading of blue-black colour to pale blue
Near endpoint: Sudden change from pale blue to colourless
Endpoint: Colourless solution that remains colourless for at least 30 seconds

Preparing Iodine Solution for Titration

Iodine has limited solubility in water but dissolves readily in potassium iodide (KI) solution due to the formation of the triiodide ion (I₃^–).

I₂ + KI → KI₃ (Potassium Triiodide)

Step-by-Step Preparation of 0.1M Iodine Solution

Step 1: Calculate Required Amounts

For 1 L of 0.1M iodine solution:

Molar mass of I₂ = 253.8 g/mol
Mass required = 0.1 × 253.8 = 25.38 g
KI required (1.5 times molar ratio) = 1.5 × 0.1 × 166 = 24.9 g

Step 2: Dissolve Potassium Iodide

Weigh 24.9 g of KI
Dissolve in 200-300 mL distilled water in a beaker
Stir until completely dissolved

Step 3: Add Iodine Crystals

Weigh 25.38 g of iodine crystals
Add gradually to KI solution with continuous stirring
Iodine will dissolve slowly forming a dark brown solution

Step 4: Transfer and Dilute

Transfer solution to 1 L volumetric flask
Rinse beaker with distilled water and add to flask
Make up to mark with distilled water
Mix thoroughly by inverting the flask

Step 5: Storage

Store in amber glass bottle to protect from light
Label with concentration and preparation date
Standardize against primary standard if needed

Safety Precautions

Wear gloves, goggles, and lab coat
Work in well-ventilated area
Avoid direct contact with iodine
Clean spills immediately

Safety Measures

Personal Protection: Always wear gloves, safety goggles, and a lab coat to protect against chemical splashes.
Ventilation: Perform the experiment in a well-ventilated area or under a fume hood to avoid inhaling iodine vapors.
Chemical Handling: Iodine can stain skin and clothing. Handle with care using appropriate tools.
Spill Management: For iodine spills, use sodium thiosulphate solution to neutralize. For sodium thiosulphate spills, clean with plenty of water.
Disposal: Dispose of waste solutions according to institutional guidelines. Never pour chemicals down the drain without proper treatment.

Calculations for Iodine Determination

Given Data:

Molarity of Na₂S₂O₃ solution, M₁ = 0.2 M

Volume of Na₂S₂O₃ used, V₁ = 24.85 mL

Volume of I₂ solution used, V₂ = 20 mL

From the balanced equation: 1 mole I₂ reacts with 2 moles Na₂S₂O₃

Therefore, n₁ (for Na₂S₂O₃) = 2, n₂ (for I₂) = 1

Using the formula:

M₁V₁/n₁ = M₂V₂/n₂

Calculation:

(0.2 × 24.85) / 2 = (M₂ × 20) / 1

4.97 / 2 = 20M₂

2.485 = 20M₂

M₂ = 2.485 / 20 = 0.12425 M ≈ 0.124 M

Amount of I₂ per dm³ of solution:

= Molarity × Molar mass of I₂

= 0.124 mol/L × 253.8 g/mol

= 31.47 g/L ≈ 31.5 g/L

Result

The concentration of the given iodine solution is 0.124 M.

The amount of I₂ dissolved per dm³ of solution is 31.5 g.

Post-Lab Discussion

Result Analysis

The titration results show good precision with less than 0.5% variation between trials.
Compare the experimental value with the theoretical or expected concentration if known.
Calculate percentage error if a standard value is available.

Sources of Error

Systematic Errors: Incorrect concentration of standard solution, improper calibration of glassware.
Random Errors: Parallax error in reading burette, difficulty in detecting endpoint colour change.
Experimental Errors: Loss of iodine due to volatilization, incomplete transfer of solutions, starch indicator added too early or too late.

Conclusion

The experiment successfully demonstrated the determination of iodine concentration using sodium thiosulphate titration. The starch indicator provided a clear and sharp endpoint, and the results were consistent across multiple trials. This method is reliable for quantitative analysis of oxidizing agents that react with iodide ions.

15 Short Questions (1 mark each)

1. What is the purpose of adding starch indicator in iodine titration?

To form a blue-black complex with iodine, providing a clear visual endpoint.

2. Why is potassium iodide added when preparing iodine solution?

To increase iodine’s solubility in water by forming the soluble triiodide ion (I₃^–).

3. What is the molar ratio between iodine and sodium thiosulphate in the titration reaction?

1:2 (1 mole I₂ reacts with 2 moles Na₂S₂O₃).

4. What colour change is observed at the endpoint of this titration?

Blue-black to colourless.

5. Why should starch indicator be added near the endpoint rather than at the beginning?

To prevent the starch-iodine complex from becoming too stable and causing a sluggish endpoint.

6. What is the oxidation product of thiosulphate in this reaction?

Tetrathionate ion (S₄O₆^2-).

7. What is the reduction product of iodine in this reaction?

Iodide ions (I^–).

8. Why is the iodine solution stored in amber or brown bottles?

To protect it from light, which can cause decomposition of iodine.

9. What safety precaution is essential when handling iodine?

Wear gloves and goggles as iodine can cause stains and is toxic.

10. What is the name of the starch-iodine complex?

Amylose-iodine complex or starch-iodine inclusion complex.

11. Why is a white tile used during titration?

To provide a contrasting background for better observation of colour changes.

12. What is the molar mass of iodine (I₂)?

253.8 g/mol.

13. What would happen if you overshoot the endpoint in this titration?

You would get an inaccurate high volume reading, leading to incorrect concentration calculation.

14. Why is sulphuric acid sometimes added in iodine titrations?

To provide an acidic medium that prevents side reactions and ensures complete reaction.

15. What practical application uses sodium thiosulphate to neutralize chlorine?

Water treatment to remove excess chlorine from drinking water.

Interactive Quiz: Iodine-Thiosulphate Titration

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Concept Assessment Exercise

Design an activity to determine the strength of given solution of iodine using 0.2M Na₂S₂O₃ solution.

Experimental Design:

Preparation: Clean and rinse all glassware. Fill burette with 0.2M Na₂S₂O₃ solution.
Measurement: Pipette 25 mL of the given iodine solution into a clean conical flask.
Indicator Addition: Add 1-2 mL of freshly prepared starch indicator solution. The solution turns blue-black.
Titration: Titrate with Na₂S₂O₃ solution while swirling continuously.
Endpoint Detection: Continue until the blue-black colour disappears completely, leaving a colourless solution.
Recording: Record the volume of Na₂S₂O₃ used. Repeat for concordant values.
Calculation: Use the formula M₁V₁/n₁ = M₂V₂/n₂ where:
- M₁ = 0.2 M (Na₂S₂O₃)
- V₁ = Volume used (from burette)
- n₁ = 2 (stoichiometric coefficient for Na₂S₂O₃)
- M₂ = Concentration of I₂ (to be determined)
- V₂ = 25 mL (volume of I₂ solution)
- n₂ = 1 (stoichiometric coefficient for I₂)
Result: Calculate and report the concentration of iodine solution in mol/L and g/L.

Sample Calculation: If 28.5 mL of 0.2M Na₂S₂O₃ is used:

(0.2 × 28.5) / 2 = (M₂ × 25) / 1

M₂ = (0.2 × 28.5) / (2 × 25) = 0.114 M

Concentration in g/L = 0.114 × 253.8 = 28.93 g/L

Daily Life Applications

Water Treatment

Sodium thiosulphate neutralizes excess chlorine in drinking water and wastewater treatment plants.

Medical Uses

Used as an antidote for cyanide poisoning and in some topical medications.

Photography

Known as “hypo,” it fixes photographic films by removing unreacted silver halide.

Analytical Chemistry

Widely used in iodometric titrations to determine concentrations of various oxidants.

Dechlorination

Used in aquaculture and aquarium maintenance to remove chlorine from tap water.

Leather Industry

Employed in the tanning process to remove excess chromium from leather.