AP CHEMISTRY
REDOX
GUIDED INQUIRY LAB
Name: ______________________
Partner #1: _______________________
Partner #2: _______________________
LAB: QUANTITATIVE OXIDATION/REDUCTION
TITRATION
INTRODUCTION
Oxidation-reduction reactions form a large category of important chemical reactions. In this lab you investigate the
reaction between iron (II) sulfate and one of the most powerful oxidizing agents found in the lab, namely, potassium
permanganate. This oxidation-reduction reaction between potassium permanganate and iron (II) sulfate in an acidic
environment will be carried out quantitatively and will also introduce you to the oxidation states of manganese..
CONTEXT:
Iron (Fe2+) is a common component of well water in many parts of the country. Many Fe 2+ compounds are soluble.
But Fe2+ iis easily oxidized to Fe3+ when it comes through a faucet and comes in contact with oxygen in the
atmosphere. This oxidation-reduction reaction leads to the formation of Fe3+ compounds, which tend to be insoluble.
This is the source of red deposits often seen on bathroom fixtures. High levels of iron in the water can lead to
plumbing problems and issues with household water quality.
Environmental scientists use titration techniques and oxidation-reduction reactions to quantitatively determinE the
level of iron in household and industrial site water. In this lab, you will determine the concentration of iron in a
sample of iron (II) sulfate using titration and the powerful oxidizing agent, potassium permanganate.
MATERIALS
100 mL of 0.020 M KMnO4
15 mL 3.0 M H2SO4
FeSO4 concentration unknown and freshly prepared.
2 burets
Graduated cylinders
Beakers and flasks
PRELAB QUESTIONS
1. What measuring devices are used to obtain precise measurements of volumes? What measuring devices are
used to obtain approximate measurements of volumes?
2. Write a balanced half-reaction for the reduction of permanganate ions in acidic solution. What are the oxidation
states of manganese in this reaction?
3. Write a balanced half-reaction for the oxidation in the iron (II) sulfate. What are the oxidation states of iron in
this reaction?
4. Write the balanced reaction for the net oxidation-reduction reaction. How many electrons are transferred in this
reaction?
5. If this reaction is carried out in neutral solution, what is the most likely oxidation state for the manganese
product?
Due Date:
______________________
AP CHEMISTRY
Name: ______________________
REDOX
GUIDED INQUIRY LAB
Partner #1: _______________________
Partner #2: _______________________
6. If this reaction is carried out in strongly basic solution, what is the most likely oxidation state for the manganese
product?
7.
A sample of oxalic acid, H2C2O4, was analyzed using a standardized solution of KMnO4. 25.0 mL of oxalic acid
is titrated after heating. 12.30 mL of a standardized 0.226 M KMnO 4 solution was added to the sample when a
faint pink color was observed. The balanced equation for the oxidation-reduction reaction is shown below:
6H+(aq)
+ 2MnO4- (aq) + 5H2C2O4 (aq) -> 10 CO2 (g) + 8H2O (l) + 2Mn2+ (aq)
a. What is the mole ratio of MnO4- ions to H2C2O4 molecules in this reaction?
b. How many moles of MnO4- ions reacted with the given amount of oxalic acid solutuion?
c. How many moles of H2C2O4 were present
d. What is the molarity of the oxalic acid solution?
e. If the density of the oxalic acid solution was 1.00 g/mL, what was the percentage by mass of oxalic acid
in the solution?
f. What does the faint pink color indicate about the reaction?
8. Consider your answer to question 7f. Why is this important to the oxidation-reduction reaction you are about to
carry out?
PROCEDURE
1. Before coming to class, develop a procedure that
will allow you to determine the concentration of
the iron (II) sulfate solution.. The procedure
should collect qualitative and quantitative data
that will allow you to determine the concentration
of the iron (II) sulfate solution so that you can
calculate the iron (II) sulfate concentration. You
must have approval from your teacher to begin.
2. You will accomplish these tasks by first obtaining
about 100 mLof the standardized KMnO4 solution
(.020 M) from the stock bottle. (CAUTION:
KMnO4 stains if it comes in contact with skin)
3. Obtain about the same amount of FeSO4 solution.
4. Place about 50 mL of distilled water in a container
and add 1 drop of the KMnO4 solution to the
water. Observe the color of the solution. This
will be the color standard for the reaction.
5. Measure 10.0. mL samples of the FeSO4 solution
for this experiment and add approximately 5.0 mL
of 3.0 M H2SO4 to increase the hydrogen ion
concentration.
6. Any equipment listed is available for you to
accomplish your task.
DATA
Due Date:
______________________
PREPARE A DATA TABLE TO RECORD ALL APPROPRIATE DATA
CALCULATIONS AND ANALYSIS
1.
Calculate the moles of 0.020 M KMnO4 solution needed to react with all of the FeSO4 in this reaction.
2.
Calculate the molarity of the FeSO4 solution for each trial.
3.
Calculate the average molarity of the FeSO4 solution.
4.
Calculate the normality of the FeSO4 solution.
5.
If the density of the FeSO4 solution is 1.10 5/mL, calculate the percentage of FeSO4 in solution for each
trial.
6.
If the FeSO4 solution was left in an open beaker overnight before the experiment was performed, would
you expect your molarity determination to yield the same result you obtained? Higher value? Lower
value? Explain.
7.
If you did not rinse your buret with the titrant after cleaning it with soap and water, what impact would this
have on your result. Would the molarity be higher, lower, or the same? Explain.
8.
What was the purpose of the drop of KMnO4 in a beaker of water?
9.
How would the concentrations of KMnO4 and FeSO4 solutions have been affected if the following
observations were made about the Erlenmeyer flask? Justify your answers.
a. When the titration was completed, the flask was colorless.
b. When the titration was completed, the flask was dark red or purple.