Name __________________________________________________ Date ________________________
NORTH ALLEGHENY SENIOR HIGH SCHOOL
AP Chemistry
DETERMINATION OF IRON BY REACTION WITH PERMANGANATE: A
REDOX TITRATION
INTRODUCTION/THEORY
Potassium permanganate is widely used as an oxidizing agent in volumetric analysis. In acid
solutions the permanganate ion undergoes reduction to manganous (2+) ion. Since the permanganate ion is
violet and the manganous ion is colorless, the end point in titrations using potassium permanganate as the
titrant can be taken as the first permanent pink color that appears in solution.
Potassium permanganate will be used in this experiment to determine the percentage of iron in an
unknown mixture of iron(II) ammonium sulfate, Fe(NH4)2(SO4)2.6H2O, and ammonium sulfate. The
titration, which involves the oxidation of ferrous ions to ferric ions by the permanganate ion, is carried out
in sulfuric acid solution to prevent the air oxidation of the ferrous ion. The end point of the titration is
sharpened markedly if phosphoric acid is present. The reason being the ferric ion by itself has a yellow
color that can partly mask the pink endpoint, but the addition of phosphoric acid forms a colorless complex
with the Fe(III) ion.
Using potassium permanganate, unfortunately, does have its drawbacks. It is such a powerful
oxidizer that it reacts practically with anything that can be oxidized. This makes it difficult to store without
decomposition or change in concentration. Because of this, it is common to prepare, standardize, then use
potassium permanganate all in the same day. Potassium permanganate solutions are prepared to be an
approximate concentration, and then they are standardized against a known primary standard of the same
substance that is to be analyzed in the unknown sample.
The number of moles of potassium permanganate used in the titration is equal to the product of the
molarity of the potassium permanganate solution and the volume used. The number of moles of iron
present in the sample is obtained from the balanced equation for the reaction and the amount of
permanganate ion reacted. The percent by weight of iron in the solid sample follows directly using correct
equations.
PRE-LAB QUESTIONS
1. Write the balanced net ionic equation for the reaction between the permanganate ion and the
ferrous ion in acid solution.
2. What is the number of moles of ferrous ion that could be oxidized by 1.2 x 10-2 moles of
permanganate ion?
3. What is the percent Fe in pure iron(II) ammonium sulfate hexahydrate?
PROCEDURE
Part 1: Standardization of the KMnO4 solution
We are going to standardize the KMnO4 solution with two trials(its concentration is approximately
0.02 M). If your numbers are appreciably different between the two trials, you should complete a third trial,
but due to time constraints two trials in part 1 should be enough.
1. Label two 250-mL Erlenmeyer flasks as samples 1 and 2. Place 1 gram of iron(II) ammonium
sulfate, Fe(NH4)2(SO4)2.6H2O, into each flask. Record the mass of the iron(II) ammonium
sulfate to at least the nearest milligram.
2. Add 25 mL of distilled water to each sample and swirl the flasks to dissolve the iron(II)
ammonium sulfate. Then add 15 mL of 3 M sulfuric acid and a pipet full (about 2 mL) of
concentrated phosphoric acid.
3. Clean out a buret by rinsing it with several portions of tap water, followed by rinsings with
distilled water. Finally, rinse the buret with the KMnO4 solution you are going to standardize.
4. Allow a few mL of KMnO4 to run out of the tip to remove air bubbles.
5. Place sample 1 under the buret and begin addition of KMnO4 a few milliliters at a time, swirling
the flask after each addition. As the permanganate is added, red streaks may be visible in the
sample until the KMnO4 has a chance to mix and react with the iron(II).
6. When the red streaks become more persistent, add the KMnO4 1 drop at a time. The end point
is the first appearance of a permanent pale pink color.
7. Record the final volume of the KMnO4 solution used.
8. From the mass of each iron(II) ammonium sulfate and the volume of KMnO4 used calculate the
molarity of each KMnO4 solution, as well as the average molarity (You will need the balanced
redox reaction to do this).
Part 2: Analysis of the Unknown
1. Prepare 2 samples of the iron unknown just as you prepared the samples of iron(II) ammonium
sulfate. The only difference being you should use approximately 0.5 g of the iron unknown
instead of 1 gram.
2. Titrate the iron unknown just as you titrated the iron(II) ammonium sulfate samples.
3. If you have time, feel free to complete third trials for either part 1 or part 2. More data means
less error.
4. From the volume of you standardized KMnO4, determine the number of moles of permanganate,
the munber of moles of iron, the mass of iron and the mass % of iron in each sample. Then
calculate the average mass %.
DATA TABLE (You may use these if you wish)
Part 1:
Trial 1
Trial 2
Trial 1
Trial 2
Mass of iron(II) ammonium
sulfate
(*) Moles of iron(II) ammonium
sulfate in sample
(Moles of Fe in sample)
Initial buret reading (KMnO4)
Final buret reading (KMnO4)
(*) Volume KMnO4 used
(*) Moles MnO4- ion required
(*) Molarity of KMnO4
Part 2:
Mass of unknown sample
Molarity of KMnO4
Initial buret reading
Final buret reading
(*) Volume KMnO4 used
(*) Moles MnO4- ion required
(*) Moles of Fe in sample
(*) Mass of Fe in sample
(*) % Fe in sample
CALCULATIONS
For the starred (*) items above, please show detailed calculations on how the value was obtained.
You only have to show calculations for one trial in both parts 1 and 2.
CONCLUSION
1. Ask your instructor for the actual % of iron in your unknown. Determine the percentage error in
your titration. Propose a source of error that prevented you from obtaining a smaller percent
error.
2. What mass of Fe(NH4)2(SO4)2.6H2O was present in your unknown sample?
3. A potential source of error for this lab could arise if not enough acid was added to the iron
solution before the titration was started. This would allow much of the permanganate to go to
MnO2(s) rather than to Mn2+. How would this affect the % of iron found at the end of the
experiment?