The document describes procedures for determining the available chlorine content in a bleaching powder sample using iodometric titration. Key steps include: 1. Standardizing a sodium thiosulfate solution by titrating it against a primary standard potassium dichromate solution in the presence of excess potassium iodide. 2. Dissolving a bleaching powder sample in water and liberating chlorine by adding acetic acid. 3. Reacting the liberated chlorine with excess potassium iodide to generate iodine, then titrating the iodine with the standardized sodium thiosulfate solution. 4. Calculating the available chlorine content based on the titration

1. 58
EXPT. 8 IODOMETRIC DETERMINATION
OF AVAILABLE CHLORINE IN A
SAMPLE OF BLEACHING POWDER
Structure
8.1 Introduction
Objectives
8.2 Principle
8.3 Requirements
8.4 Solutions Provided
8.5 Procedure
8.6 Observations and Calculations
8.7 Result
8.1 INTRODUCTION
In the previous experiment you learnt about and performed iodimetric determination
of ascorbic acid in a tablet of Vitamin C. In this process we recalled from Unit 10 of
the Basic Analytical Chemistry course (Section 10.7.1) that the I2/2I‒
couple is of
medium oxidising power.Molecular iodine is a weak oxidant, whereas the iodide ions
are relatively weak reductant. Accordingly, the I2/I–
redox couple can be used for the
determination of reductants as well as the oxidants. In the previous experiment you
used I2 (generated in situ from acidification of KIO3) as the oxidizing agent to
determine the reductant, ascorbic acid. In the present experiment you will be using
iodide ions (from KI) to determine available chlorine in bleaching powder. Such a
determination wherein iodide ions are used as a reducing agent is termed as
iodometric determination.
Bleaching powder, also known as chlorinated lime, is a yellowish-white powder
having a smell of chlorine and is readily soluble in water. It is prepared by passing
chlorine gas over slaked lime at a temperature of 35-450
and consists of a mixture of
calcium hypochlorite Ca(OCl)2 and calcium chloride CaCl2; in addition some amount
of free slaked lime i.e. Ca(OH)2.H2O is generally present. Of these, Ca(OCl2) is
responsible for the bleaching action of bleaching powder . On treatment with glacial
acetic acid, it liberates chlorine gas (Cl2) as per the following reaction.
( ) ( ) 222332 CIOHCOOCHCaCOOH2CHOCICa ++→+ ….(8.1)
The amount of chlorine liberated by the action of an acid on bleaching powder
(CaOCl2) is termed as available chlorine. The chlorine content of bleaching powder
varies from 35 – 40%. Besides bleaching action it has got strong germicidal and
disinfectant properties also. Accordingly, it finds application as a disinfectant for
drinking water or swimming pool water. Industrially, the bleaching powder finds
major use in chemical, paper, textile and oil industries. The bleaching, oxidizing or
disinfecting potential of a sample of bleaching powder depends on the percentage of
chlorine liberated on action of acid. We may define available chlorine to be the grams
of chlorine liberated from 100 g of the bleaching powder on treatment with dilute acid.
Due to its hygroscopic nature, bleaching powder absorbs moisture from atmosphere
and evolves chlorine as per the following reaction
( ) 222
2-
CIOHCaOHCaClOCl +→+++ +−
….(8.2)
Due to this deterioration, a sample of bleaching powder may always contain lesser
amount of chlorine than expected and therefore a sample of bleaching powder needs to
be analysed for its effective or available chlorine. In the next experiment you would

2. 59
learn about precipitation titrations and perform a precipitation titration for the
determination of chloride ions in a solution.
8.2 OBJECTIVES
After studying and performing the experiment, you should be able to:
• define available chlorine,
• define iodometric titrations,
• state and explain the principle of iodometric titrations with reference to the
determination of available chlorine in a sample of bleaching powder,
• state the reasons for using acetic acid to liberate chlorine from bleaching
powder,
• prepare a standard solution of potassium dichromate and use it to standradise a
solution of sodium thiosulphate,
• write the chemical equations involved in the titration of a solution of bleaching
powder solution with sodium thiosulphate,
• prepare a solution of bleaching powder from the given sample,
• perform the determination of available chlorine in the solution of bleaching
powder, and
• calculate the amount of available chlorine in the solution of bleaching powder.
8.3 PRINCIPLE
As mentioned in the introduction, a sample of bleaching powder liberates chlorine gas
(Cl2) on treatment with glacial acetic acid, as per the following reaction.
( ) ( ) 222332 CIOHCOOCHCaCOOH2CHOCICa ++→+ …(8.1)
The amount of chlorine so liberated is termed as available chlorine. The liberated
chlorine can be used to oxidize KI (taken in excess) in presence of acid and liberate
out an equivalent amount of iodine as per the following equation:
22 I2KCl2KICl +→+ …(8.3)
This iodine can then be determined by titrating against a standardised solution of
sodium thiosulphate using freshly prepared starch solution as an indicator. The
chemical reactions involved can be given as follows
2NaIOSNaIOS2Na 6422322 +→+ …(8.4)
The overall reaction between the chlorine liberated from the bleaching powder and
sodium thiosulphate mediated by potassium iodide can be obtained by adding eq. 8.3
and eq. 8.4 can be written as follows,
22 I2KC2KICl +→+ l …(8.3)
2NaIOSNaIOS2Na 6422322 +→+ …(8.4)
ll CK22NaIOSNa2KICOS2Na 6422322 ++→++ …(8.5)
The sodium thiosulphate can be standardized by titrating against a primary standard
solution of potassium dichromate. (alternatively, you can use potassium iodate for the
purpose as described in the previous experiment)
Sodium thiosulphate,
Na2S203.5H20, can be obtained
chemically pure. However, a
standard solution of
thiosulphate cannot be made
by exact weighing as it reacts
with atmospheric O2 and also
the CO2 dissolved in water.
More so, even some
microorganisms can
decompose thiosulphate

3. 60
The standradisation of sodium thiosulphate by using potassium dichromate is also
based on iodimetric titrations. The reaction between potassium dichromate and sodium
thiosulphate is mediated through the participation of iodide ions provided by
potassium iodide as explained below.
In acidic medium, Cr2O7
2-
ions gets reduced to Cr (III) ions as per the following
equation
O7H2Cr6e14HOCr 2
3-2
72 +→++ ++−
…(8.6)
The iodide ions from KI on the other hand can get oxidised to I2 as follows:
−−
+→ 2eI2I 2 …(8.7)
In order to maintain the electron balance, we can multiply equation 8.7 by 3 to get the
following,
−−
+→ 6eI36I 2 … (8.8)
Adding equations 8.6 and 8.8, we get the overall ionic equation for the reaction
between potassium dichromate and potassium iodide as :
O7H3I2Cr6I14HOCr 22
32
72 ++→++ +−+−
… (8.9)
Thus, from Eq. 8.9, one mole of potassium dichromate reacts with 6 moles of
potassium iodide and liberates 3 moles of iodine in the process.
The liberated iodine, in turn, reacts with sodium thiosulphate solution as per the
following equation
−−−
+=+ 2
64
2
322 O3S6IO6S3I …(8.10 )
The iodide ions are regenerated back. The net chemical reaction involving a titration
of potassium dichromate and sodium thiosulphate in the presence of excess potassium
iodide can be written by combining Eq. 8.8 and Eq. 8.9, as shown below,
O7H3I2Cr6I14HOCr 22
32
72 ++→++ +−+−
... (8.9)
−−−
+=+ 2
64
2
322 O3S6IO6S3I ... (8.10)
O7HO3S2CrO6S14HOCr 2
2
64
32
32
2
72 ++→++ −+−+−
… (8.11)
We see from Eq. 8.11 that one mole of potassium dichromate is equivalent to 6 moles
of sodium thiosulphate. Therefore, the molarities are related by the following
relationship.
6 MDichromateVDichromate = MThiosulphateVThiosulphate … (8.12)
You would be using this equation to compute the molarity of the given solution of
sodium thiosulphate.
8.4 REQUIREMENTS
Apparatus Chemicals

4. 61
Volumetric flask (100 cm3
) – 1
Burette (50 cm3
) – 1
Pipette (10 cm3
) – 1
Weighing bottle – 1
Burette stand with clamp – 1
Conical flasks (100 cm3
) – 2
Funnel – 1
Beakers (250 cm3
) – 2
Bleaching powder
Potassium dichromate
Potassium iodide
Sodium thiosulphate
Sodium carbonate
Sulphuric acid
Sodium bicarbonate
Acetic acid
Starch
8.5 SOLUTIONS PROVIDED
1. ~0.05 M Sodiun thiosulphate: It is prepared by dissolving about 7.9 g of
sodium thiosulfate pentahydrate in about 200 cm3
of distilled water taken in a
1.0 dm3
volumetric flask and adding about 0.1 g of sodium carbonate to it. The
solution is then diluted to the mark with distilled water.
2. 0.5% Starch indicator solution: It is prepared by mixing 0.25g of soluble
starch with 50 cm3
of distilled water taken in a 100 cm3
conical flask or beaker
and heating it with stirring at about 80o
C for about 5 minutes. The solution is
then allowed to cool to room temperature.
3. 10% Potassium iodide solution: It is prepared by dissolving 100g of KI in
about 200 cm3
of distilled water taken in a 1 dm3
beaker or conical flask and
stirring well to dissolve it. It is followed by making up the volume to 1 dm3
by
adding more distilled water.
8.6 PROCEDURE
The determination of available chlorine in a sample of bleaching powder using
iodometric titration consists of the following steps:
a) Preparation of potassium dichromate primary standard
b) Standradisation of sodium thiosulphate solution
c) Preparation of the solution of the bleaching powder sample
d) Determination of available chlorine in the above solution by iodimetric
titration
Follow the instructions given below in sequential manner
a) Preparation of potassium dichromate primary standard
• Accurately weigh about 0.3g of potassium dichromate in a clean dry
weighing bottle, and transfer the same to a clean conical flask of 100 cm3
capacity through a glass funnel.
• Add about 20 cm3
of distilled water and swirl the contents of the flask
until all the potassium dichromate is dissolved.
• Make the volume upto the mark by adding more distilled water.
b) Standradisation of sodium thiosulphate solution
Sodium carbonate is added to
adjust the pH to
approximately 9.3 so as to
inhibit the formation of
thiosulfonic acid.

5. 62
• Pipette out 10 cm3
of potassium dichromate solution in a 100 cm3
conical
flask, add 10 cm3
of dilute sulphuric acid and 1 g sodium hydrogen
carbonate with gentle swirling to liberate carbon dioxide.
• Add 10 cm3
of 10% KI solution, swirl, cover the flask with watch glass
and allow the solution to stand for about 5 minutes in a dark place.
• Titrate the liberated iodine against the sodium thiosulphate solution taken
in the burette until the solution acquires a light pale yellow colour.
• Add about 2 cm3
of starch solution and continue adding sodium
thiosulphate dropwise until the violet colour of the starch iodine
complex just disappears.
• Repeat the standardization procedure at least three times and record your
observations in Observation Table 8.1.
c) Preparation of the solution of the bleaching powder sample
• Accurately weigh about 3-4 g of the bleaching powder and put it into a
clean glass mortar. Add a little water, and rub the mixture to a smooth
paste.
• Add a little more water, triturate with the pestle and allow the mixture to
settle.
• Pour off the milky liquid into a 500-cm3
volumetric flask.
• Grind the residue with a little more water, and repeat the operation until
the whole of the sample has been transferred to the flask either in solution
or in a state of very fine suspension, and the mortar washed quite clean.
• Make the volume upto the mark by adding more distilled water.
d) Determination of available chlorine in the above solution by iodometric
titration
• Wash the burette with distilled water and rinse with standard solution of
sodium thiosulphate and then fill the burette with the same.
• Carefully pipette out 10 cm3
of homogenous solution of bleaching powder
and transfer into a 100 cm3
conical flask.
• Add about 10 cm3
of 10% potassium iodide (KI) solution and about half
test tube of glacial acetic acid to the flask.
• Keep the flask in a dark place for about 5 minutes
• Titrate the liberated iodine against the sodium thiosulphate solution taken
in the burette until the solution acquires a light pale yellow colour.
• Add about 2 cm3
of starch solution and continue adding sodium
thiosulphate dropwise until the violet colour of the starch iodine
complex just disappears.
• Repeat the standardization procedure at least three times and record your
observations in Observation Table 8.2.
8.7 OBSERVATIONS AND CALCULATIONS
Sodiu
carbon
an atm
CO2 i
which
air an
oxidat
from a

6. 63
a) Preparation of standard solution of potassium dichromate
Mass of weighing bottle + potassium dichromate = m1 g = ..............g
Mass of weighing bottle (after transferring potassium dichromate)
= m2 g = .............. g
Amount of potassium dichromate transferred = m1 – m2 = m g = ............. g
Molar mass (Mm) of potassium dichromate = 294.18 g mol−1
Volume of potassium dichromate prepared = 100 cm3
Molarity of standard potassium dichromate solution =
K Cr O2 2 7
1000 10
...........
100 294.2 294.2
m m
M M
×
= = =
×
b) Standardisation of sodium thiosulphate solution
Volume of standard K2Cr2O7 solution taken in conical flask, Vdichromate = …cm3
Volume of 10% KI added : 10 cm3
Volume of dilute sulphuric acid added : 10 cm3
Solution in the burette: Sodium thiosulphate
Indicator used: Starch
Observation Table 8.1: Standardisation of sodium thiosulphate
S.No. Volume of potassium
dichromate (in cm3
)
Burette reading
Initial Final
Titre value (in cm3
)
(Final-initial reading)
1
2
3
Concordant reading
The concentration of the given sodium thiosulphate solution can be determined as
follows.
The reactions involved:
O7HO3S2CrO6S14HOCr 2
2
64
32
32
2
72 ++→++ −+−+−
Molarity equation: 6MDichromateVDichromate = MThiosulphateVThiosulphate
Thiosulphate
6 Dichromate Dichromate
Thiosulphate
M V
M
V
=
Substituting the values, the molarity of thiosulphate =
The molarity of given thiosulphate solution is = ……….M
c) Preparation of the solution of the bleaching powder sample

7. 64
Mass of bleaching powder taken =
d) Determination of available chlorine in the given sample of bleaching
powder
Volume of bleaching powder solution taken in conical flask, = 10 cm3
Volume of acetic acid added : 10 cm3
Volume of 10% KI added : 10 cm3
Solution in the burette: Sodium thiosulphate
Indicator used: Starch
Observation Table 8.2: Determination of the amount of available chlorine in the
solution of given bleaching powder
S.No. Volume of bleaching
powder solution (in
cm3
)
Burette reading
Initial Final
Titre value (in cm3
)
(Final-initial reading)
1
2
3
Concordant reading
The molarity of the iodine liberated from KI solution (which in turn is equal to the
amount of chlorine liberated from the bleaching powder on the action of acetic acid)
can be determined as follows.
The reaction involved:
ll 2C2NaIOSNa2KICOS2Na 6422322 ++→++
Molarity equation: 2MChlorine VChlorine = M Thiosulphate V Thiosulphate
Thiosulphate Thiosulphate
2Chlorine
Chlorine
M V
M
V
=
Substituting the values, of the molarity and the volume of thiosulphate used, the
molarity of chlorine is found to be: ......M=
The mass of chlorine liberated = molarity × molar mass = ….M × 70.90 g mol−1
= ….×70.90 g = P g dm-3
The mass of bleaching powder dissolved per liter = 2 ×w g
(the solution was w g/ 500 cm3
)
Thus, the amount of available chlorine in 2 w g of bleaching powder = P g
• The amount of available chlorine in 100.0 g of bleaching powder = 50 P / w g
• The available chlorine (the grams of chlorine liberated from 100 grams of the
bleaching powder on treatment with dilute acid) = 50 P/w g
8.8 RESULTS
The available chlorine (the grams of chlorine liberated from 100 g of the bleaching
powder on treatment with dilute acid) = ….. g