Measurement of Fe2+ in Iron Pills
Redox Titration
Purpose:
You will prepare and standardize a potassium permanganate solution and use it to
determine the amount of iron, specifically Fe2+, in iron pills.
Overview:
Titration is the volumetric measurement of a solution of known concentration when it
reacts completely with a measured volume or mass of another substance. In the first part
of this experiment, the potassium permanganate solution, the titrant, is standardized, by
quantitatively reacting it with iron (II) ammonium sulfate hexahydrate,
Fe(NH4)2(SO4)2•6H2O (abbreviated FAS). The second part of the lab uses the
standardized potassium permanganate to determine the amount of the analyte, also iron
(II), in iron supplement pills.
In this experiment, a purple-colored solution of potassium permanganate, KMnO4, with
an approximate concentration of 0.02 M will be added to solutions containing Fe2+ ions.
In redox titration the chemical reaction is an oxidation-reduction reaction with electrons
transferred from one species to another. The total number of electrons lost in the
oxidation half-reaction is equal to the total number of electrons gained in the reduction
half-reaction. The permanganate ion, MnO4-, is a strong oxidizing agent which causes
the Fe2+ to be oxidized to Fe3+ ions. The manganese gains five electrons and is reduced
from a 7+ oxidation state in the permanganate ion to form colorless Mn2+ ions. The Fe2+
loses one electron as it is oxidized to Fe3+. The end point is indicated when the reaction
is complete at the point when all of the Fe2+ ions in solution are oxidized and the
colorless mixture retains the purple tint of unreacted permanganate. Note that the color
may be somewhat orange in appearance, depending upon the concentration of the
Fe3+(aq), which has an orange-yellow-green tint. A few drops of concentrated
phosphoric acid, H3PO4, will be added to form a complex with the iron to minimize this
color. Sulfuric acid, H2SO4, is also added to increase the concentration of hydrogen ions
in the solution.
The net ionic equation for the reaction is:
+
5Fe2+ + MnO4- + 8H  5Fe3+ + Mn2+ + 4H2O
Safety Precautions:
Safety goggles and aprons must be worn in at all times. Potassium permanganate,
sulfuric acid, and phosphoric acid can cause chemical burns to skin and eyes. The
potassium permanganate is also an oxidizer and also will stain skin and clothing. Avoid
skin contact with these chemicals and wash any contaminated area thoroughly with cold
water. Report any spills.
Procedure:
Preparation and Standardization of KMnO4 Solution
1)
Prepare 250 mL of an approximately 0.02M potassium permanganate solution
by adding the appropriate amount of potassium permanganate to a 250mL
volumetric flask and diluting to the line using deionized water.
2)
Rinse a clean buret several times with a few mL of deionized water. Rinse the
buret with a few mL of the potassium permanganate and discard the washings
into a waste beaker. Fill the buret with the KMnO4 solution.
3)
Clean three Erlenmeyer flasks and label the flasks "1", "2", and “blank”.
Place approximately 2 g of iron (II) ammonium sulfate hexahydrate,
Fe(NH4)2(SO4)2•6H2O (abbreviated FAS), in a weighing boat and mass with
an analytical balance to ±0.0001 g. Transfer half of the FAS to one of the
flasks, then remass the boat. The difference is the mass of FAS in sample 1.
Transfer the remaining FAS to the second flask and remass the boat.
Determine mass of sample 2. Add 25 mL of deionized water, 15 mL of 3M
H2SO4, and a few drops of concentrated H3PO4 to each flask and swirl to
dissolve the FAS.
4)
Place about 50 mL of deionized water in the flask labeled “blank” and add 1
drop of the permanganate solution. This is the color standard for the reaction.
When the end point is reached, the color intensity of the mixture should match
this standard.
5)
Record the initial volume reading in the buret to ± 0.01 mL. While swirling
continuously, add the permanganate solution to the FAS solution in flask 1
until the end point is reached. Record the final volume reading in the buret
and determine the exact volume of permanganate solution used.
6)
Repeat the titration process using the FAS in flask 2.
7)
Wash the mixtures in the flasks down the drain, then rinse the flasks with
deionized water.
8)
From the mass of each FAS sample, calculate the moles of Fe2+ ions present.
Use the balanced equation to determine the moles of KMnO4 needed to reach
the end point. Using the numbers of moles and the volumes of KMnO4
solution used to titrate the samples, calculate the molarity of your
permanganate solution.
Analysis of the Iron Pill
1)
Obtain 5 iron supplement pills and place the pills in a weighing boat. Mass
with an analytical balance to ± 0.0001 g.
2)
Grind the 5 iron pills in a clean, dry mortar and pestle. Place the ground pills
in a weighing boat and mass with an analytical balance to ± 0.0001 g.
Transfer half of the powder to a clean Erlenmeyer flask, then remass the boat.
The difference is the exact mass of iron pill in sample 1. Transfer the
remaining powder to the second flask and remass the boat. Determine the
exact mass of iron pill in sample 2. Add 25 mL of deionized water, 15 mL of
3M H2SO4, and a few drops of concentrated H3PO4 to each flask and swirl to
dissolve the iron pill.
3)
Record the initial volume reading in the buret to ± 0.01 mL. While swirling
continuously, add the permanganate solution to the iron pill solution in flask 1
until the end point is reached. Record the final volume reading in the buret
and determine the volume of permanganate solution used.
4)
Repeat the titration process using the iron pill in flask 2.
5)
Wash the mixtures in the flasks down the drain, then rinse the flasks with
deionized water. Clean the buret, rinse several times with tap water, and then
with deionized water.
6)
From the volume of KMnO4used to titrate the samples, determine the moles of
KMnO4 used. Calculate the moles of Fe2+ present in the pill samples based on
the balanced equation. Determine the mass of Fe2+ equivalent to this number
of moles, then calculate the mass percentage of Fe2+in the iron pill samples
and the average % composition. Finally calculate the mass of iron in each
pill.
Adapted from an East Penn School District laboratory experiment.
Questions:
1) What does the manufacturer claim about the iron content of the tablets? How do
your results compare?
2) In what ratio do the MnO4- and Fe2+ react?
3) What is the purpose of the phosphoric acid?
4) What is the purpose of the sulfuric acid?
5) Potassium permanganate titration is also used to measure concentrations of
compounds such as hydrogen peroxide and oxalic acid. What information would
be needed to successfully determine the concentrations of these compounds?
6) If the iron in the pills was Fe3+ instead of Fe2+, could you still use a potassium
permanganate titration to determine the iron concentration?
7) What would have happened if you had prepared a potassium permanganate
solution with a concentration of 0.01M instead of the approximately 0.02M?