Mandatory Experiment 4.8
Determination of the percentage (w/v) of hypochlorite in
bleach
Student Material
Theory
Many commercial bleaches are simply solutions of hypochlorite salts such as sodium
hypochlorite (NaOCl) or calcium hypochlorite (Ca(OCl)2). Hypochlorite ion reacts with
excess iodide ion in the presence of acid to generate an iodine solution:
ClO- + 2I- + 2H+ → Cl- + I2 + H2O
The liberated iodine solution can be titrated against sodium thiosulfate solution using a
freshly prepared starch solution as indicator. The titration reaction may be represented by
the equation:
I2 + 2S2O32- → 2I- + S4O62Starch indicator is added during the titration when the colour of the solution in the
conical flask fades to a pale yellow colour. The solution becomes blue-black, and the
titration is continued until it goes colourless.
Chemicals and Apparatus
Bleach solution
Potassium iodide
i
Dilute sulfuric acid
i
0.1 M sodium thiosulfate solution
Starch solution
Deionised (or distilled) water
Safety glasses
Volumetric flask (250 cm3) and stopper
Dropping pipette
Pipette (25 cm3)
Pipette filler
Burette (50 cm3)
Filter funnel
Conical flask (250 cm3)
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White tile
White card
Graduated cylinders (100 cm3)
Retort stand
Boss-head
Clamp
Beakers (250 cm3)
Wash bottle
Procedure
NB: Wear your safety glasses.
1. The bleach solution must first be diluted to make a solution of suitable
concentration for the titration. Using a pipette, add 25 cm3 of bleach to a 250 cm3
volumetric flask, and make the solution up to the mark with deionised water. The
flask should be stoppered and inverted several times.
2. Wash the pipette, burette and conical flask with deionised water. Rinse the pipette
with the diluted bleach solution and the burette with the sodium thiosulfate
solution.
3. Using a pipette filler, fill the pipette with the diluted bleach solution and transfer
the contents of the pipette to the conical flask.
4. Add 1 g potassium iodide and 10 cm3 of dilute sulfuric acid to the conical flask –
this liberates iodine.
5. Using a funnel, fill the burette with sodium thiosulfate solution, making sure that
the part below the tap is filled before adjusting to zero.
6. With the conical flask standing on a white tile, add the solution from the burette to
the flask. Swirl the flask continuously and occasionally wash down the walls of
the flask with deionised water using a wash bottle.
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7. Add a few drops of the starch indicator solution just prior to the end-point, when
the colour of the solution fades to pale yellow. A blue-black colour appears. The
thiosulfate solution should now be added dropwise, with thorough swirling.
8. The end-point of the titration is detected when the blue-black colour changes to
colourless. Note the burette reading.
9. Repeat the procedure, adding the sodium thiosulfate solution dropwise
approaching the end-point, until two titres agree to within 0.1 cm3.
10. Calculate the concentration of the iodine solution.
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11. Calculate the concentration of hypochlorite in the bleach solution.
12. Calculate the percentage (w/v) of hypochlorite in the bleach solution.
Table of results
Copy this table into your practical report book.
Rough titre
Second titre
Third titre
Average of accurate titres
Volume of iodine solution used in each titration
Concentration of sodium thiosulfate solution
Concentration of iodine solution
Concentration of hypochlorite in the diluted bleach solution
Concentration of hypochlorite in the bleach
Percentage (w/v) of hypochlorite in the bleach
=
=
=
=
=
=
=
=
=
=
Questions relating to the experiment
1. Give one reason why, in making up the solution of diluted bleach, a volumetric
flask is preferable to a graduated cylinder.
2. A burette, a pipette and a conical flask were used in the titration. State the correct
washing procedures for each of these items before starting the titration.
3. Why could you not use hydrochloric acid when acidifying the bleach?
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Teacher Material

Milton solution is very suitable for use in this experiment, as is sodium
hypochlorite solution purchased from the usual school laboratory suppliers, and
Parazone thin bleach.

Some bleaches, such as Domestos, are less suitable –when Domestos is diluted, it
forms suds, which makes it very difficult to accurately make up the diluted
bleach solution.

Though sodium thiosulfate is available in a high state of purity, it is efflorescent,
which creates some doubts over the water content in crystals from opened
containers.

Furthermore, solutions of thiosulfate are broken down by reaction with even trace
quantities of acid resulting from absorption of atmospheric carbon dioxide by
distilled water. Decomposition can also be catalysed by bacterial action if
solutions are left stand for long periods.

The stability of sodium thiosulfate solutions is enhanced if they are made up in
boiled deionised water (to remove carbon dioxide) with sodium carbonate (0.1
g/l) added. Larger quantities of sodium carbonate cause too high an increase in
pH and accelerate air oxidation of the thiosulfate ion.

Solutions that can act as primary standards can be made up from anhydrous
sodium thiosulfate. Anhydrous sodium thiosulfate can be prepared by refluxing
sodium thiosulfate pentahydrate crystals in methanol for about 30 minutes. The
anhydrous solid is isolated by filtration and dried in an oven at 70 oC.

Alternatively, sodium thiosulfate solution can be standardised using a standard
iodine solution, as in mandatory experiment 4.7.

Starch solutions are widely used in the detection of the end-point of iodine thiosulfate titrations. The great advantage of starch lies in the fact that it gives a
very definite colour change at the end-point. Without the starch indicator, the
colour of the iodine solution in the conical flask near the end-point fades slowly
from pale yellow to colourless. With the starch indicator added, the colour of the
solution in the conical flask at the end-point changes suddenly from blue black to
colourless.
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
There are however a number of disadvantages in the use of starch as an indicator.
Solid starch has a limited shelf life. Other disadvantages include the insolubility
of starch in cold water and the fact that starch forms an insoluble complex with
iodine. The latter problem precludes the addition of the indicator at the start of
the titration and this is why the indicator is to be added just prior to the end point
when the colour of the solution has changed to pale yellow. In addition to this
there is sometimes a 'drift' in the end point, particularly when solutions are dilute.

Starch solution needs to be freshly prepared, as it deteriorates fairly quickly.
However, if 0.5 g of salicylic acid is added to 50 cm3 of the solution immediately
after it is made up, the solution will last for a few months.

About 0.5 to 0.75 cm3 of starch solution is needed per titration, assuming a total
volume of 50 to 75 cm3 in the titration flask at the end-point.

A fuller description of titration procedure is to be found in the Student Material
relating to Mandatory Experiment 4.2.
Preparation of reagents
Sodium thiosulfate solution: To prepare a 0.l M solution of sodium thiosulfate, dissolve
15.8 g of anhydrous sodium thiosulfate in boiled deionised water and make the solution
up to 1 litre using a volumetric flask. To increase the stability of the solution, add 0.1 g of
sodium carbonate.
Starch solution:
Pour with stirring a paste containing 1 g starch and a little cold water into 100 cm3 of
boiling water. Boil for two minutes, and allow to cool. The solution should be stored in
stoppered bottles.
The dilute (1.5 M) sulfuric acid solution is prepared as follows:
(Always dilute sulfuric acid by adding the acid to water and not the other way round.) 85
cm3 of the concentrated acid is added slowly to about 700 cm3 of deionised water
containing about 20 ice cubes. The mixture is stirred and made up to 1 litre in a
volumetric flask with deionised water. The flask is stoppered and inverted a number of
times.
Quantities per working group
50 cm3 bleach solution
5 g potassium iodide
100 cm3 dilute sulfuric acid
200 cm3 sodium thiosulfate (0.1 M)
3 cm3 starch indicator solution
Deionised water
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Safety considerations

Safety glasses must be worn.
Chemical hazard notes
Sodium thiosulfate: Irritant to eyes.
Methanol
: Toxic by ingestion or inhalation. Much more poisonous than
ethanol. Highly flammable.
Potassium iodide may be harmful by ingestion. Eye irritant.
Starch: Starch powder is explosive when dry. Dust may irritate eyes and lungs.
Concentrated sulfuric acid
is very corrosive to eyes and skin. Due to its very
considerable heat of reaction with water, it is essential that the acid be added to water
when it is being diluted.
Dilute sulfuric acid
Bleach
i:
i
is harmful to eyes and an irritant to skin.
Hypochlorite solutions are corrosive, harmful and irritant.
Disposal of wastes
Dilute with water, neutralise with anhydrous sodium carbonate and flush to foul water
drain with excess water.
Specimen Results (for sodium hypochlorite solution)
Rough titre
Second titre
Third titre
Average of accurate titres
Volume of diluted bleach used to release iodine
Concentration of sodium thiosulfate solution
= 22.6cm3
= 22.4 cm3
= 22.5 cm3
= 22.45 cm3
= 25.0 cm3
= 0.1 M
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Specimen Calculations
(a) First principles method
Volume of thiosulfate solution used
Moles of thiosulfate used
= 22.45 cm3
= 22.45 x 0.1 /1000
= 0.002245 moles
Balanced equations:
ClO- + 2I- + 2H+  Cl- + I2 + H2O
1 mole 2 moles 2 moles 1 mole 1 mole 1 mole
I2 + 2S2O32-  2I- + S4O621 mole 2moles
2 moles 1 mole
Thus ClO- : S2O32- = 1 : 2
Moles of hypochlorite solution used
Moles/cm3 of hypochlorite solution
Moles/litre of hypochlorite solution
Concentration of hypochlorite in diluted bleach solution
Concentration of hypochlorite in undiluted bleach solution
Percentage (w/v) of hypochlorite (NaOCl) in bleach
= 0.002245/2
= 0.0011225 moles
= 0.0011225 / 25
= 0.0000449
= 0.0449
= 0.0449 M
= 0.0449 x 10 M
= 0.449 M
= 74.5 x 0.449 / 10
= 3.35%
(b) Formula method
VA x MA x nB = VB x MB x nA
25.0 x MA x 2 = 22.45 x 0.1 x 1
MA = 22.45 x 0.1 x 1 / (25.0 x 2)
= 0.0449
Concentration of hypochlorite in diluted bleach solution
Concentration of hypochlorite in undiluted bleach solution
Percentage (w/v) of hypochlorite in bleach
= 0.0449 M
= 0.0449 x 10 M
= 0.449 M
= 74.5 x 0.449 / 10
= 3.35%
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Specimen Results (for Parazone thin bleach)
Rough titre
Second titre
Third titre
Average of accurate titres
Volume of diluted bleach used to release iodine
Concentration of sodium thiosulfate solution
= 29.1 cm3
= 28.9 cm3
= 28.9 cm3
= 28.9 cm3
= 25.0 cm3
= 0.1 M
Specimen Calculations
Using either the first principles method or the formula method,
Concentration of hypochlorite in diluted bleach solution
= 0.0578 M
Therefore, concentration of hypochlorite in undiluted bleach solution = 0.0578 x 10 M
= 0.578 M
Percentage (w/v) of hypochlorite (NaOCl) in bleach
= 74.5 x 0.578 / 10
= 4.31%
Solutions to student questions
1. Give one reason why, in making up the solution of diluted bleach, a volumetric flask is
preferable to a graduated cylinder.
A volumetric flask is quite an accurate measuring vessel, whereas a graduated cylinder is
not.
2. A burette, a pipette and a conical flask were used in the titration. State the correct
washing procedures for each of these items before starting the titration.
Rinse the burette, pipette and conical flask respectively with deionised water. Rinse the
burette with sodium thiosulfate solution, and rinse the pipette with diluted bleach
solution.
3. Why could you not use hydrochloric acid when acidifying the bleach?
Hydrochloric acid is not suitable, as it will react with hypochlorite to liberate chlorine
gas.
Extension Work
Analysis of a different brand of bleach may be done, using the same method, and the
results compared.
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