4.8 - Determination of the Percentage of Hypochlorite in Bleach

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Experiment Summary

This experiment aims to determine the percentage of hypochlorite ( $ClO⁻$ ) in commercial bleach.

The hypochlorite ion reacts with excess iodide ions in the presence of acid to liberate iodine ( $I₂$ ).
The iodine is then titrated with sodium thiosulfate ( $Na₂S₂O₃$ ), with starch as an indicator.
The endpoint is marked by the disappearance of the blue-black colour formed by the starch-iodine complex.

Materials and Apparatus Required

Chemicals

Bleach solution (sodium hypochlorite, $NaOCl$ )
Potassium iodide ( $KI$ )
Dilute sulfuric acid ( $H₂SO₄$ )
Sodium thiosulfate solution ( $Na₂S₂O₃$ )
Starch indicator solution
Deionised water

Apparatus

Volumetric flask (250 cm³)
Pipette (25 cm³) and pipette filler
Burette (50 cm³) with retort stand and clamp
Conical flask (250 cm³)
Graduated cylinders (100 cm³)
White tile
Filter funnel
Wash bottle
Beakers (250 cm³)

Safety Precautions

Wear safety glasses throughout the experiment.
Sodium hypochlorite is corrosive and can irritate the skin and eyes. Avoid contact and handle with care.
Sulfuric acid is corrosive. Always dilute acid by adding it to water, never the other way around.
Potassium iodide and sodium thiosulfate are irritants; avoid inhaling dust or vapours.

Method

Use a pipette to transfer 25 cm³ of bleach solution into a 250 cm³ volumetric flask.
Fill the flask to the mark with deionised water and invert several times to mix.
Rinse the pipette with diluted bleach solution and the burette with sodium thiosulfate solution.
Use the pipette to transfer 25 cm³ of the diluted bleach solution into a clean conical flask.
Add about 1 g of potassium iodide and 10 cm³ of dilute sulfuric acid to the conical flask. This liberates iodine, turning the solution a brown colour.
Fill the burette with sodium thiosulfate solution and record the initial burette reading.
Titrate the iodine solution with sodium thiosulfate, swirling continuously.
When the solution turns pale yellow, add a few drops of starch indicator. The solution will turn blue-black.
Continue adding sodium thiosulfate dropwise until the solution turns colourless.
Record the final burette reading.
Repeat the titration until two concordant titres (within 0.1 cm³) are obtained.

Results

Measurement	Value
Rough titre	22.6 cm³
Second titre	22.4 cm³
Third titre	22.5 cm³
Average of accurate titres	22.45 cm³
Volume of diluted bleach solution	25.0 cm³
Concentration of sodium thiosulfate	0.1 M
Concentration of hypochlorite in bleach	0.449 M
Percentage of hypochlorite in bleach	3.35%

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Sample Calculation

Using the balanced equation:

\text{ClO}^- + 2\text{I}^- + 2\text{H}^+ \rightarrow \text{Cl}^- + \text{I}_2 + \text{H}_2\text{O}

Moles of thiosulfate used:

\frac{22.45 \times 0.1}{1000} = :highlight[0.002245 \, \text{moles}]

From the balanced equation:

\text{ClO}^- : \text{S}_2\text{O}_3^{2-} = :highlight[1 : 2]

Moles of hypochlorite:

0.002245 / 2 = :highlight[0.0011225 \, \text{moles}]

The concentration of hypochlorite in diluted bleach:

\frac{0.0011225}{25 / 1000} = :highlight[0.0449 \, \text{M}]

The concentration of hypochlorite in undiluted bleach:

0.0449 \times 10 = :highlight[0.449 \, \text{M}]

Percentage (w/v) of hypochlorite in bleach:

\frac{74.5 \times 0.449}{10} = :success[3.35\%]

Example Questions with Answers

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Q1: Why is a volumetric flask used to dilute the bleach instead of a graduated cylinder?

A volumetric flask is more accurate for making precise dilutions compared to a graduated cylinder.

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Q2: How should the burette, pipette, and conical flask be washed before the titration?

Rinse the burette with sodium thiosulfate, the pipette with a diluted bleach solution, and the conical flask with deionised water.

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Q3: Why can't hydrochloric acid be used to acidify the bleach solution?

Hydrochloric acid would react with hypochlorite, producing chlorine gas, which is hazardous.

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Q4: Why is starch added near the end of the titration?

Starch forms a blue-black complex with iodine, providing a clear endpoint when the colour disappears.

Adding it too early could result in an inaccurate endpoint.

- Determination of the Percentage of Hypochlorite in Bleach Simplified Revision Notes for Leaving Cert Chemistry

4.8 - Determination of the Percentage of Hypochlorite in Bleach

Experiment Summary

Materials and Apparatus Required

Chemicals

Apparatus

Safety Precautions

Method

Results

Sample Calculation

Example Questions with Answers

Q1: Why is a volumetric flask used to dilute the bleach instead of a graduated cylinder?

Q2: How should the burette, pipette, and conical flask be washed before the titration?

Q3: Why can't hydrochloric acid be used to acidify the bleach solution?

Q4: Why is starch added near the end of the titration?

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