EXPERIMENT 2 – REDOX TITRATION OF Fe2+ WITH MnO4BACKGROUND
Oxidation and reduction reactions are those chemical reactions that involve
the transfer of electrons. A substance is oxidized when it loses electrons, and a
substance is reduced when it gains electrons. An easy way to help remember this is
OIL RIG (Oxidation Is Loss, Reduction Is Gain). Since one species can gain an
electron only when another loses an electron, oxidation and reduction processes
always accompany one another and always occur to an equal extent. Thus, redox
reactions can serve as the basis of titrations and can provide an accurate and
sensitive method to determine quantitatively the amount of either species. If we
have a known concentration of a material that is easily reduced (which we call an
oxidizing agent) and carefully measure the volume required to react exactly with a
material that is easily oxidized (which we call a reducing agent), we can determine
the number of electrons gained by the oxidizing agent and lost by the reducing
agent. Knowledge of the stoichiometry of a reaction makes it possible to calculate
how many moles of the reducing agent are present in an unknown.
In this experiment, permanganate (MnO4-) will be used as the oxidizing
agent, and iron(II) will be the material oxidized. In acid solution, the two balanced
half-cell reactions are:
–
Reduction: 8 H+ + MnO4 + 5 e
–
Mn
2+
+ 4 H2O
Fe3+ + e–
Oxidation: Fe2+
When combining these half reactions, one must keep in mind that the number of
electrons lost and the number gained must be the same. Therefore, the overall
stoichiometry of the balanced reaction is:
(1)
+
2+
–
8 H (aq) + 5 Fe (aq) + MnO4 (aq)
2+
3+
Mn (aq) + 5 Fe (aq) + 4 H2O (l)
Each iron(II) ion loses 1 electron to become iron(III) while each manganese ion
gains 5 electrons. This knowledge of the stoichiometry is necessary in using the
reaction for a quantitative determination – e.g., a titration.
In this experiment, you will begin by preparing a solution of potassium
permanganate which has a molarity of about 0.02M. Then, you will determine its
exact molarity (i.e., “standardize” it) by titrating your permanganate solution
against pure iron(II) ammonium sulfate hexahydrate, Fe(NH4)2(SO4)2·6H2O. (An
older name for this compound is ferrous ammonium sulfate.) In this titration, the
–
permanganate ion will serve as its own indicator. The MnO4 ion is deep purple,
2+
but the Mn ion is colorless (as are both Fe2+ and Fe3+). Therefore, when all the
2+
–
Fe has been consumed, the added MnO4 will no longer react and will cause the
solution to turn pink. The first hint of a stable pink color signals the end point of
the titration. You will do three of these standardization titrations and use the
average as the concentration of the permanganate solution.
After having standardized your permanganate solution, you will be given an
iron(II) unknown in the form of a white crystalline powder. You will weigh out,
dissolve, and titrate three samples of this unknown using your standardized
permanganate solution and determine the average percent Fe present in the
unknown.
EXPERIMENTAL PROCEDURE
Preparation of approximate permanganate solution (~0.02 M MnO4-)
1. Weigh approximately 0.65 grams of KMnO4 onto weigh paper. Place about
75 mL of deionized water in a 600 mL beaker, add about a third of the
permanganate to the water, and stir well. Add another 75 mL of water to the beaker
and another one-third of the permanganate. Continue stirring. Finally add 50 mL of
water to the beaker along with the rest of the permanganate and stir well for several
minutes. Let the solution sit for about 5 minutes while you clean your buret
thoroughly with soap and water. Carefully decant the permanganate solution into a
400 mL beaker. Don't allow any undissolved or solid material to be transferred to
the 400 mL beaker (transferred solids will clog the buret!). The resulting solution
is approximately 0.02 M and will be used for all your titrations.
2. After rinsing the clean buret at least 3 times with water, rinse it 3 more times
with small amounts (about 5 mL) of the permanganate solution and then fill the
buret with the permanganate solution. This solution may be too deeply colored to
allow you to read the bottom of the meniscus; if so, you will need to read the top of
the liquid column instead.
Standardization of the permanganate solution
3. Weigh carefully three portions (approximately 1 gram each) of the iron(II)
ammonium sulfate (Fe-ASH) into 3 Erlenmeyer flasks. Record the weight
accurately to 0.0001 g. Dissolve the samples in about 40 mL of deionized water.
4. Add approximately 10 mL of 2.0 M sulfuric acid to each Erlenmeyer flask just
before starting the titration. IMPORTANT! Don't forget to add the sulfuric acid!
To make the color change more easily seen, place a piece of white paper under the
Erlenmeyer flask. Begin the titration by adding the potassium permanganate
solution carefully to the flask, swirling constantly. As soon as the color of an
additional drop fades rather slowly, add the permanganate solution 1 drop at a
time. When the pink color persists for 30 seconds, stop the titration and read the
buret. Refill the buret and repeat the procedure with the second and third sample.
Report the concentration of your permanganate solution as an average of all trials,
including the 95% confidence interval.
Determination of the %Fe in the unknown
5. Using the balance, accurately weigh three portions (approximately 1 gram each)
of the iron(II) containing unknown into three Erlenmeyer flasks. Record the weight
accurately to 0.0001 g. Dissolve the samples in about 40 mL of deionized water.
Refill the buret with permanganate solution and proceed as in the standardization
procedure above. DON'T forget to add the sulfuric acid just before starting the
titration! Calculate the mass of Fe(II) present and the percent Fe(II) in the
unknown. Report your results as an average of all three trials, including the 95%
confidence interval.
CHEM 208
PRELAB QUESTIONS FOR EXPERIMENT 3:
REDOX TITRATION OF Fe2+ WITH MnO4Name________________________________________
Section_________
TA Name_____________________________________
1. (3 points) Write the balanced chemical equation that corresponds to the titration
of Fe2+ with MnO4-.
__________________________________________________________
2. (3 points) Suppose that 0.2553 moles of MnO4- were required to exactly titrate
2+
2+
the Fe in a sample. How many moles of Fe were in the sample?
3. (3 points) Suppose that 22.65 mL of a 0.02024 M permanganate solution was
2+
needed to react with the Fe in a sample that had a mass of 1.1243 g. What is the
2+
2+
mass of the Fe in the sample and what is the % of Fe in the sample?
4. (6 points) Balance the following oxidation-reduction (redox) reactions. SHOW
YOUR WORK.
(a)
C
+
H2SO4
CO2 +
(b)
Ag +
H2SO4
Ag2SO4 +
(c)
ClO3- +
Cl-
Cl2
SO2
+
SO2
ClO2-
CHEM 208
LAB WRITE-UP FOR EXPERIMENT 3:
REDOX TITRATION OF Fe2+ WITH MnO4Name____________________________________
Section______________
TA Name_________________________________
Date________________
Unknown #_______________________________
Standardization of permanganate solution (steps 3 and 4)
Initial buret reading (KMnO4)
Final buret reading (KMnO4)
Volume KMnO4 used
Trial
1
2
3
Mass Fe-ASH
n Fe2+
Trial 1
Trial 2
Trial 3
__________
__________
__________
__________
__________
__________
__________
__________
__________
n MnO4-
Vol KMnO4 used
M KMnO4
Note that "n" means “number of moles”
Sample calculation for n Fe2+, n MnO4-, and M KMnO4 for Trial 1:
Q test (90%) of M KMnO4 (show calculation):
Confidence Interval (95%) of M KMnO4 (show calculation):
Average M KMnO4 ± 95% Conf. Interval.:_______________________________
Determination of mass Fe2+ in the unknown (step 5)
Initial buret reading (KMnO4)
Final buret reading (KMnO4)
Volume KMnO4 used
Trial
Trial 1
Trial 2
Trial 3
__________
__________
__________
__________
__________
__________
__________
__________
__________
Average M Vol KMnO4
n MnO4- n Fe2+
KMnO4
used
Mass Fe2+
Weight of
Unknown
% Weight
Fe2+
1
2
3
Sample calculation of n MnO4-, n Fe2+, mass Fe2+, and % weight Fe2+ for Trial 1:
Q test (90%) of % Weight Fe2+ (show calculation):
Confidence Interval (95%) of % Weight Fe2+ (show calculation):
Final result:
Unknown number
Average % Weight Fe2+ ± 95% Conf. Interval
For TA Use
Accuracy_______________________
Grade__________________________