Lab 4.3 Quantitative Titration Involving Redox Reactions
Name _______________
Acid-base titration is used to determine the quantitative relationship between two reacting solutions when no precipitate
is formed. Titration methods can also be used to determine the concentration of an ion solution if the substance is easily
oxidized or reduced. This process is called oxidation-reduction or redox titration.
The analytical process for a redox titration is similar to that of an acid-base titration. In both cases, a balanced ionic
equation must be written in order to determine the molar ratio of the reactants. In this experiment you will titrate
standardized KMnO4 with FeSO4 of unknown concentration. You will then use the volume data for the reacting solutions
and the concentration of the standard to determine the concentration of the unknown solution.
PROCEDURE
1. Obtain 2 pipets, verify that they are labeled KMnO4 and FeSO4.
2. Clean the two beakers, the test tube, and the flask.
3. Fill your smallest beaker with about 20 mL of the standardized KMnO4 solution (0.010M) from the correct flask.
Fill your middle size beaker with 20 mL of FeSO4 solution. CAUTION: KMnO4 is a strong irritant to skin and eyes.
Keep away from heat and flames.
4. Place about 50 mL of distilled water in an Erlenmeyer flask and add 1 drop of the KMnO 4 solution. Observe the color
of the solution. This solution will be the color standard for the reaction. Leave the solution in the flask.
5. Pipet 20 drops of FeSO4 solution in the test tube. Add 5 drops of H2SO4 in the test tube. Add enough ditilled water
to the test tube so that it is roughly ¾ - filled. CAUTION: H2SO4 causes severe burns, avoid skin contact.
6. Begin titrating with the KMnO4 solution. Count the number of drops of KMnO4 solution until the end point is reached.
Record the number of drops in the data table. Discard the contents of the test tube.
Rinse the test tube with distilled water. Repeat the titration two more times, then average.
Data Table
Trial 1
Trial 2
Trial 3
Average
# drops of KMnO4
ANALYSIS
The following reaction is the redox reaction in this lab, involving KMnO4, FeSO4, and H2SO4.
10 FeSO4 + 2 KMnO4 + 8 H2SO4  5 Fe2(SO4)3 + 2 MnSO4 + K2SO4 + 8 H2O
The net ionic equation:
Fe+2 + MnO4-1  Fe+3 + Mn+2
1. Determine the charge of the manganese in the permanganate ion, MnO4-1 .
Charge: _______________
2. How many electrons did each manganese gain or lose in this reaction?
_______________
3. Was manganese oxidized or reduced?
_______________
4. How many electrons did iron gain or lose in this reaction?
_______________
5. Was iron oxidized or reduced?
_______________
6. You may notice that the number of electrons gained and lost were not the same.
Please multiply the following half-reactions correctly so that the number of electrons transferred balances.
1 Mn+7 + ___ e-  1 Mn+2
1 Fe+2  1 Fe+3 + ___ e7. Please combine the half-reactions into one, balanced redox reaction.
___ Mn+7 +
___ Fe+2  ___ Mn+2 ___ Fe+3
8. Now the titration. This is worth 10 points. You must show your work. Complete the conversion table at the bottom.
We will pretend that each drop of solution is 1 mL. Using your average number of drops of KMnO4 as the number
of mLs, use the data to calculate the molarity of the FeSO4.
___ Mn+7 + ___ Fe+2  ___ Mn+2 ___ Fe+3
___ mLs
20 mLs
0.01 M
?M
= __________________ M