Experiment No. ___________________
Date ___________________
OXIDATION-REDUCTION TITRATIONS-Permanganometry
INTRODUCTION
Potassium permanganate, KMnO4, is probably the most widely used of all volumetric oxidizing
agents. It is a powerful oxidant and readily available at modest cost. The intense color of the
permanganate ion, MnO4-, is sufficient to detect the end point in most titrations.
Depending upon reaction conditions permanganate ion is reduced to manganese in the 2+, 3+,
4+ or 6+ state.
In solutions that are 0.1 M or greater in mineral acid the common reduction product is
manganese (II) ion
MnO4- + 8H+ + 5e- ↔ Mn2+ + 4H2O
E0 = 1.51 V
This is the most widely used of the permanganate reactions.
In solutions that are weakly acidic (above pH 4) neutral, or weakly alkaline manganese dioxide
is the most common reduction product
MnO4- + 4H+ + 3e- ↔ MnO2(s) + 2H2O
E0 = 1.70 V
Titration in which manganese dioxide is the product suffer from the disadvantage that the
slightly soluble brown oxide obscures the end point; time must be allowed for the solid to settle
before an excess of the permanganate can be detected.
Some important volumetric analyses based on permanganate involve reduction to manganese
ion according to the half reaction given below;
MnO4- + e- ↔ MnO42-
E0 = 0.56 V
This stoichiometry tends to predominate in solutions that are greater than 1 M in sodium
hydroxide. Alkaline oxidations with permanganate have proved to be most useful in the
determination of organic compounds.
REAGENTS AND APPARATUS
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FeSO4.7H2O, Ca(NO3)2.4H2O, H2O2 for unknowns
Potassium permanganate, KMnO4: 1.0 L of 0.020 M (for 4 students)
Sodium oxalate, Na2C2O4 (primary standard in desiccator)
Ammonium oxalate, (NH4)2C2O4.H2O (solid in the balance room)
Methyl orange (in droppers)
Stannous chloride, SnCl2 (already prepared)
Mercury (II) Chloride, HgCl2 (already prepared)
Preventive solution (or Zimmermann-Reinhart reagent) (already prepared)
Sulfuric acid, H2SO4, 500 mL of 0.75 M (for 2 students)
Hydrochloric acid, HCl, 25 mL of 6.0 M (for each student)
Ammonia, NH3, 25 mL of 6.0 M (for each student)
Sulfuric acid, H2SO4, 250 mL of 3.0 M (for 2 students)
0.1 M AgNO3 (in droppers)
Blue band filter paper (2 for each student)
Glass wool
buret
Watch glass
250 mL conical flasks x2 ( for each student)
100 mL graduated cylinder
PROCEDURE
A) Preparation and Standardization of KMnO4 Solution
1)
2)
Prepare 1.0 L of 0.020 M KMnO4 in distilled water for 4 students. Pour into a 1.0 L beaker.
Keep the solution at a gentle boil for about 1 hour.

Distilled water may contain organic matter which will reduce MnO4-2 ion. The solution is
heated in order to hasten the oxidation of this material and coagulate the colloidal
precipitate of MnO2 which forms as a reduction product:
MnO 2 -  reducing agent  MnO  oxidation product (in neutral solution)
4
2
3)
4)
Cover and let stand overnight.
Remove MnO2 by filtration using glass wool.

5)
Particles of MnO2 should be removed since these particles catalyze further
decomposition of the solution. Filter paper can not be used for filtering because
permanganate ion reacts with it to form additional manganese dioxide.
Store the solution in a clean, glass-stoppered amber bottle and keep in the dark when not in
use.
B) Standardization of Permanganate Solution with Sodium Oxalate
1)
2)
3)
Weigh 0.2 to 0.3 g (0.1 mg) dry primary standard Na2C2O4 into 250 mL Erlenmeyer flasks.
Dissolve the sodium oxalate in 75.0 mL of 0.75 M H2SO4.
Heat the solution to 80 to 90 C and titrate with KMnO4, do not boil the solution. End point is
permanent pink color (~30 s). The net reaction in the titration can be written as follows:
2MnO 42-  5C 2 O42  16H   2Mn 2  10CO2  8H2O

Promptly wash any KMnO4 that spatters on the walls of the beaker into the bulk of the
liquid with a stream of water.

Finely divided MnO2 will form along with Mn2+ if the KMnO4 is added too rapidly and will
cause the solution to acquire a faint brown discoloration. Precipitate formation is not a
serious problem as long as sufficient oxalate remain to reduce the MnO 2 to Mn2+; the
titration is simply discontinued until the brown color disappears. The solution must be
free of MnO2 at the end point.
4)
Repeat the titration with one additional sample.
5)
Determine a blank by titrating an equal volume of 0.75 M H2SO4. Correct the titration data
for the blank.
Your partners will do the standardization experiment twice. Record the results of all
titrations (totally 8 titrations) and calculate the molarity of the KMnO 4 solution for eigth
replicates and at the end calculate the average of molarity of KMnO4 solution.
6)
C) Determination of Iron
1)
Take your unknown sample into a 250 mL Erlenmeyer flask and place it in a steam bath
and evaporate to a volume of about 10 mL. The solution will probably show the yellow color
of ferric ion at this stage.

Estimate the volume comparison with 10 mL of water in a similar beaker. Reduction of
stannous chloride must be carried out in a relatively concentrated solution in order the
point at which reaction is complete can be determined by the change in color.
2)
Treat each unknown solution individually.
3)
While still hot, add SnCl2 solution a drop at a time until color changes to light green. The
first drop added may be sufficient to reduce all Fe3+ present. The reaction is given below:
2Fe3  Sn2  2Fe2  Sn 4
4)
5)
Add two drops of SnCl2 solution in excess after the color change is observed.
Cool the solution to room temperature and pour in rapidly 20.0 mL mercuric chloride, HgCl2
solution.

The excess reducing agent is eliminated by addition HgCl2. The slightly soluble mercury
(I) chloride (Hg2Cl2) does not reduce permanganate, nor does the excess mercury (II)
chloride (HgCl2) reoxidize iron (II).
Sn

6)
7)
 2HgCl2  Hg2Cl2 (s)  Sn
4
 2Cl

If the reagent is not added all at once, there may be a local excess of stannous ion
which will cause the reduction of HgCl2 to Hg rather than Hg2Cl2 mercurous chloride.
Elemental Hg(l) reacts with permanganate and causes the results of the analysis to be
high.
Sn

2
2
 HgCl2  Hg(l)  Sn
4
 2Cl

The precipitate should be small in amount and of a pure white color. A grayish
precipitate indicates reduction to mercury. This reacts slowly with permanganate and
will not give true titration values. If no precipitate appears, it means that not enough
stannous chloride has been added. If a small white precipitate is not obtained, the
sample is spoiled and should be discarded.
Allow to stand 2 minutes and then add 100 mL distilled water.
Add 20.0 mL preventive solution and titrate immediately with permanganate. The net
reaction in the titration can be written as follow:
5Fe 2  MnO   8H   5Fe 3  Mn 2  4H O
4
2
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8)
9)
The reaction of iron(II) with permanganate is smooth and rapid. The presence of iron (II)
in the reaction mixture, however, induces the oxidation of chloride ion by permanganate,
a reaction that does not ordinarily proceed rapidly enough to cause serious error. High
results are obtained if this parasitic reaction is not controlled. Its effects can be
eliminated through removal of the hydrochloric acid by evaporation with sulfuric acid or
by the addition of Zimmermann-Reinhart reagent, which contains manganese (II) in a
fairly concentrated mixture of sulfuric and phosphoric acid. The oxidation of chloride ion
during a titration is believed to involve a direct reaction between this species and the
manganese(II) ions that form as an intermediate in the reduction of permanganate ion by
iron (II). The presence of manganese(II) in the Zimmermann-Reinhart reagent is belived
to inhibit the formation of chlorine by decreasing the potential of the
manganese(III)/manganese(II) couple. Phosphate ion is believed to exert a similar effect
by forming stable manganese(III) complexes. Moreover, phosphate ions react with
iron(III) to form nearly colorless complexes so that the yellow color of the iron(II)/chloro
complexes doe not interfere with the end point.
Determine a blank by adding two drops of stannous chloride to 100 mL of distilled water in
a 250 mL Erlenmeyer flask; then proceed with addition of 20.0 mL of mercuric chloride,
20.0 mL of preventive solution and 100 mL of distilled water just in the titration of a sample.
Note the volume of permanganate needed to give the same color as the end point reached
in the titrations.
Subtract the blank from the total volume used to obtain the net volume for each portion of
unknown sample.
10) Report the result as mg iron.
D) Determination of Hydrogen Peroxide
1)
2)
Take your unknown sample into a 250 mL Erlenmeyer flask and add 100 mL distilled water.
Treat each unknown solution individually.
3)
4)
Add 7.0 mL of 3.0 M H2SO4.
Titrate with standard KMnO4 solution. End point is the first pink color that persists for 2
minutes.
Report the result as mg H2O2.
5)
The net reaction in the titration is
  6H   5O  2Mn 2  8H O
5H2O2  2MnO4
2
2
E) Determination of Calcium
1) Take your unknown sample into a 250 mL Erlenmeyer flask and add 100 mL distilled water
2) Treat each unknown solution individually.
3) Add 10.0 mL of 6.0 M HCl.
4) Heat to 60 C to 80 C and then add 3.0 g of (NH4)2C2O4.H2O.
5) Add 3 drops of methyl orange indicator, then introduce 6.0 M NH3 dropwise from pipette until
the color changes from red to yellow.
6) Allow the solution to stand half an hour.
7) Filter the solution through a blue band filter paper.
8) Wash the beaker and precipitate with 10 to 20 mL portions of distilled water until the
washings show a faint cloudiness when tested with an acidified AgNO3 solution. After
enough washings almost all Cl- ions can be removed.
9) Place the filter in which the precipitate was formed in a 250 mL Erlenmeyer flask. Add 50.0
mL of distilled water and 50.0 mL of 3.0 M H2SO4.
CaC O (s)  2H
2 4

 Ca
2
H C O
2 2 4
10) Heat to 80 to 90 C to dissolve the precipitate. Titrate with standardized KMnO4 solution with
filter paper. The temperature of the solution should not be allowed to drop below 60 C.
Take as an end point the first pink color that persists for 15 to 20 sec. The net reaction in the
titration is given below:
2


2
5H C O
 2MnO  6H  2Mn
 10CO  8H O
2
2
4
2 2 4
11) Report the result as mg Ca.
PRE-LAB STUDIES
Read pages 510-531 from the textbook (9th Ed)
1) Why KMnO4 is preferred as an oxidizing agent in redox titrations?
2) What is the aim of heating KMnO4 during its preparation step? Explain and write the related
reaction equations.
3) What is the importance of filtering KMnO4 solution and why do we use glass-wool instead of
filter-paper?
4) Explain the importance of heating the solution to 80 to 90 C in the standardization of
KMnO4.
5) Why should the medium be acidic in KMnO4 standardization?
6) What is the aim of adding SnCl2 in Fe determination? Write the related reaction equations.
7) What are the function(s) of preventive solution? Explain.
8) Why an indirect method is applied for Ca2+ determination?
POST-LAB STUDIES
1) Why do we perform the titrations slowly?
2) During standardization of KMnO4 if you observe a brown precipitate, what does this
indicate? Do you continue the titration in this situation or stop and repeat it?
3) Why do we perform blank analysis in KMnO4 standardization?
4) Why do we add HgCl2 and why do we add it quickly? Write the related reaction equations.
5) If you add HgCl2 slowly what will you observe and how does this affect your result?
6) Why should the pH of the medium be about 4 in Ca2+ determination? How do we adjust this
pH?
7) Why do we use NH3 but not NaOH for pH adjustment in Ca2+ determination?
Name surname:
Section:
Date:
REPORT SHEET FOR PERMANGANOMETRY
B. Standardization of Permanganate Solution with Sodium Oxalate
Replicates
Group Members:
Name, Surname
Mass of
Na2C2O4, g
Blank corrected
volume of
KMnO4, mL
Concentration
of KMnO4, M
Concentration
of KMnO4, M
̅  𝐬)
(𝐗
1
2
3
4
5
6
7
8
C. Determination of Iron
Replicates
Blank corrected
volume of KMnO4,
mL
Mass of
iron, mg
Mean mass,
̅  𝐬)
mg (𝐗
True mass of
iron, mg
1
2
The following information (true values) will be sent to your e-mail address:
Concentration of iron in the unknown solution=
Volume of iron unknown, mL=
TA`s Name and Signature:
% Relative
Error
D. Determination of Hydrogen Peroxide
Replicates
Volume of KMnO4,
mL
Mass of
H2O2, mg
Mean mass,
̅  𝐬)
mg (𝐗
True mass of
H2O2, mg
% Relative
Error
1
2
The following information (true values) will be sent to your e-mail address:
Concentration of H2O2 in the unknown solution=
Volume of H2O2 unknown, mL=
E. Determination of Calcium
Replicates
Volume of KMnO4,
mL
Mass of
calcium,
mg
Mean mass,
̅  𝐬)
mg (𝐗
True mass of Ca,
mg
1
2
The following information (true values) will be sent to your e-mail address:
Concentration of calcium in the unknown solution=
Volume of calcium unknown, mL=
TA`s Name and Signature:
% Relative
Error