Oxidation Reduction Titration

Title: Oxidation-Reduction Titration Iodimetry
Activity No. 5
Date: August 22, 2015
Names: Ma. Theresa M. Llasos and Jessica
Magayanes
ABSTRACT
There are several types of redox titrations and
two of which are iodimetry and iodometry.
Iodimetry is a redox titration which uses iodine
as the titrant, an oxidizing agent that reacts with
the analyte, a reducing agent. Iodometry is a
redox reaction in which the liberated iodine
produced in the analyte of the sample is titrated
with a standard solution of thiosulfate, Na 2S2O3.
The Na2S2O3 solution is the reducing agent while
the liberated iodine in sample is the oxidizing
agent in iodometry. In this experiment, iodometry
was performed with starch as the indicator. Its
reaction is given by:
I2 + S2O32- 2I- + S4O62The objectives of this experiment were to
prepare a standard solution of sodium
thiosulfate (Na2S2O3) and determine the strength
of a bleaching agent by oxidation-reduction
titration.
The amount of Na2S2O3
5H2O crystals
needed to prepare a 500 mL of 0.1 M Na 2S2O3

solution was 12. 41 g. The average molarity of
Na2S2O3 from the four trials performed was
0.3085 g/mol. The % Cl present in the sample
was 0.826 % given 15.1 mL of Na 2S2O3 and 10.0
g sample.
In this experiment, the %Cl content of
the unknown bleach was determined. Possible
sources of error include oxidation and loss of
iodine by vapor, insufficient acidity of standard
iodine solution which can cause incomplete
reduction of dichromate by iodide.
INTRODUCTION
There are several types of redox titrations and

two of which are iodimetry and iodometry. Both
of these titrations involve the use of iodine.
Iodine (I2) is an oxidizing agent that can be used
to titrate fairly strong reducing agents. On the
other hand, iodide ion (I -) is a mild reducing
agent and serves as the basis for determining
strong oxidizing agents. 1
In iodimetry, iodine is a moderately strong
oxidizing agent and can be used to titrate
reducing agents. Titrations with I 2 are called
iodimetric methods. These methods are usually
performed in neutral or mildly alkaline (pH 8) to
weakly acid solutions. If the pH is too alkaline, I 2
will disproportionate to hypoiodate and iodide:2
I2 + 2OH- = IO- + I- + H2O
Iodine has a low solubility in water but the
complex, I3-, is very soluble. So iodine solutions
are prepared by dissolving I2 in a concentrated
solution of KI:3
I2 + I- I3I3- is therefore the actual species used in
titration.
In iodometry, iodide ion is a weak reducing agent
and will reduce oxidizing agents. It is not used,
however, as a titrant because of lack of
convenient visual indicator system, as well as
other factors such as speed of reaction. 4
When an excess of iodide is added to a solution
of an oxidizing agent, I2 is produced in an
amount equivalent to the oxidizing agent. This I 2
can therefore be titrated with a reducing agent
and the result will be the same as if the oxidizing
agent were titrated directly. The titrating agent
used is sodium thiosulfate. 5 Standard solution of
sodium thiosulfate is one of the few reducing
agents that is stable toward air oxidation.
The end point for iodometric titrations is
detected with starch. The disappearance of the
blue starch-I2 color indicates the end of the
titration. The starch is not added at the
beginning of the titration when the iodine
concentration is high. Instead, it is added just
before the endpoint when dilute iodine color

becomes pale yellow.
An example of this procedure is the
determination of hypochlorite in bleaches. The
reactions are
2S2O32- + I2 2I- + S4O622H+ + ClO- + 2I- I2+ Cl- + H2O
In this experiment, sodium thiosulfate solution
was standardized iodometrically against a pure
oxidizing agent, the K2Cr2O7 and the strength of
a bleaching agent was determined by oxidationreduction titration.
METHODOLOGY
A. Preparation of Starch Indicator
A 0.5 g of starch was weighed and was
dissolved with 5 mL of distilled water. The
starch solution was added to a 100 mL
boiling water and was boiled for another 2
minutes.
B. Preparation and Standardization of 0.1
M Na2S2O3 solution
The weight of Na2S2O3
5H2O crystals
needed to prepare a 500 mL of 0.1 M

Na2S2O3 solution was calculated. The
crystals were dissolved in a beaker with 100
mL distilled water (pre-boiled). It was diluted
to make a 500 mL solution. A 0.2 g of
Na2CO3 was added to the Na2S2O3 solution.
A 0.10 0.05 g of pure, dry K 2Cr2O7 was
weighed into each of three 250 mL
Erlenmeyer flasks. It was dissolved by
adding 50 mL distilled water (pre-boiled). A 4
mL of 1:2 H2SO4 was added to the solution.
A 5.0g of KI was weighed and was dissolved
by adding water. A 5 mL of the KI solution
was added to each Erlenmeyer flask and
was covered with watch glass. The analyte
was allowed to stand by for 3 minutes. The
analyte should have an initial brown color.
After 3 minutes, the analyte was diluted with

50 mL of distilled water (pre-boiled) and was
titrated with thiosulfate solution until brown
color of iodine had disappeared. After the
first change in color, a 5 mL of starch
solution was added and the titration
continued until the last drop of the titrant
removes the blue color of the starch-iodine
complex gives a clear emerald green
solution. The molarity of the Na 2SO3 solution
was calculated. The average deviation
should be about 1-3ppt.
C. Analysis of the Unknown
A 2.0 mL of tap water was pipette into
each of the three Erlenmeyer flasks. A
50 mL of distilled water (pre-boiled), 3
mL of KI, 8 mL of 1:6 H2SO4 and 3 drops
of 3 % ammonium molybdate (optional)
were added. Each flask was covered
with a watch glass and was allowed to
stand for 3 minutes to allow the reaction
to be completed. The analyte was
titrated with standard thiosulfate solution
until brown color of iodine had
disappeared. After a change in color, a 5
mL of starch indicator was added and
the titration continued until the
disappearance of the blue color. The %
Cl was calculated in each product
assuming that the density of liquid is 1.0
g/mL.
RESULTS AND DISCUSSION
A. Preparation of Na2S2O3 solution
Amount of Na2S2O3
5H2O crystals needed:
12. 41 g
500 mL, 0.1 M Na2S2O3
(0.5 L)(0.1 mol/ 1L)(248.16 g Na2S2O3
5H2O) = 12.408 g = 12. 41 g
B. Standardization of Na2S2O3
Trial
Wt.
K2Cr2O7
(g)
Vol.
Na2S2O3
(mL)
Molarity
Na2S2O3
(g/mol)
0.1035
8.14
0.2593
0.1048
6.60
0.3238
0.1215
6.90
0.3591
4
Average
0.1010
7.06
0.2918
0.3085
Sample Computation:
Trial 1: M Na2S2O3 = 0.2593
(0.1035 g Cr2O72)(1 mol Cr2O72/ 294. 2 g Cr2O72)
(3 mol I2/ 1mol Cr2O72)(2 mol S2O32/ 1mol I2) =
0.2593 g/mol Na2S2O3
Trial 2: M Na2S2O3 = 0.3238
2
7
2
7
(0.1048 g Cr2O )(1 mol Cr2O7 / 294. 2 g Cr2O )

0.2593 g/mol Na2S2O3 = 0.3238 g/mol
Trial 3: M Na2S2O3 = 0.3591
(0.1215 g Cr2O72)(1 mol Cr2O72/ 294. 2 g Cr2O72)
0.2593 g/mol Na2S2O3 = 0.3591 g/mol
Trial 4: M Na2S2O3 = 0.2918
(0.1010 g Cr2O72)(1 mol Cr2O72/ 294. 2 g Cr2O72)
0.2593 g/mol Na2S2O3 = 0.2918 g/mol
Average Molarity of Na2S2O3 = 0.2593 +
0.3238 + 0.3591 + 0.2918/ 4 = 0. 3085 g/mol
C. Determination of Bleaching Power
Trial
Vol.
Na2S2O3
(mL)
15. 1
Wt.
sample
(g)
10.0 g
% Cl in
the Sx.
0.826 %
(0.0151 L)(0.3085 mol/LNa2S2O3)(1 mol I2/2 mol

S2O3)(1 mol Cl-/1 mol I2)(35.45 g/ 1 mol Cl -)
10.0 g sample 100 % = 0.826 % Cl
In the standardization, the molarity of the
thiosulfate solution was obtained by dividing the
weight of the dried dichromate by the molecular
mass of potassium dichromate. The quotient
was then stoichiometrically converted to moles
of S2O32-, and then divided by the volume of
sodium thiosulfate obtained through titration.
The molarity per trial exceeded the 0.1 M of the
standard sodium thiosulfate solution. The
volume of the titrant consumed is so small that it
produced such molarities. The possible sources
of this are from (1) the expired KI added to the
sample and (2) the additional H2SO4 to the
sample for it did not yield an initial brown color
before titration. The iodine solution was so
diluted that it did not yield the brown color.
In the determination of the bleaching power of
the unknown, the %Cl was obtained by
determining the moles of S2O3 and converting it
to g of Cl-. The answer was divided by 10.0g
sample and multiplied to 100 %. The % Cl in the
sample was 0.826 %.
CONCLUSION
In this experiment, the %Cl content of
the unknown bleach was determined. Possible
sources of error include oxidation and loss of
iodine by vapor, insufficient acidity of standard
iodine solution which can cause incomplete
reduction of dichromate by iodide.
REFERENCES
Sample Computation:
Skoog, et al. Fundamentals

Chemistry. 9th edition
of
Analytical
Gary Christian, Analytical Chemistry, 7th edition

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Title: Oxidation-Reduction Titration Iodimetry

needed to prepare a 500 mL of 0.1 M Na 2S2O3

There are several types of redox titrations and

before the endpoint when dilute iodine color

needed to prepare a 500 mL of 0.1 M

After 3 minutes, the analyte was diluted with

5H2O crystals needed:

(0.1048 g Cr2O )(1 mol Cr2O7 / 294. 2 g Cr2O )

(0.0151 L)(0.3085 mol/LNa2S2O3)(1 mol I2/2 mol

Skoog, et al. Fundamentals

Gary Christian, Analytical Chemistry, 7th edition

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