AP Chemistry Laboratory #4: Determination of Concentration by Redox Titration
Lab days: Wednesday and Thursday, September 25-26, 2013
Lab due: Friday, September 27, 2013
Goal (list in your lab book):
The goal of this lab is to determine the actual amount of iron in an iron pill by titrating it with a
standard potassium permanganate solution.
Introduction/Background (DON’T write in your lab book):
A common task in analytical chemistry is the determination of the amount of a substance present in
a sample or product. If the product contains a substance that can be oxidized, then it is possible to
determine the number of moles of that substance by titrating the sample with a solution of a strong
oxidizing agent. In this lab, a solution of an oxidizing agent will first be standardized and then used
to determine the amount of iron in an iron pill.
During the standardization part of the lab (Day 1), ferrous ammonium sulfate will be oxidized to iron
(III) ions in acidic solution according to the following unbalanced equation:
H+ (aq) + MnO4- (aq) + Fe2+ (aq) -----> Mn2+(aq) + Fe3+ (aq) + H2O (R)
[Eqn 1]
For this titration, the equivalence point occurs when the exact number of moles of Fe2+ ions has been
added to react completely with all the permanganate ions in the solution.
Instead of phenolphthalein or some other acid-base indicator, the indicator for this titration is the
permanganate ion itself. The permanganate ion is purple (violet) in solution. At the endpoint of the
titration, the solution changes from light pink to colorless.
In the second part of the lab (Day 2), the amount of iron in an iron pill will be determined by titration
with the standardized potassium permanganate solution from Day 1.
An animation of a similar reaction is available at
http://group.chem.iastate.edu/Greenbowe/sections/projectfolder/flashfiles/redoxNew/redox.html
NOTE: This lab will take two days.
Research questions (answer in your lab book in complete sentences; don’t write the questions):
1) List the names and formulas for all the chemicals that will be used on Day 1.
2) What is oxidation? What is reduction?
3) (a) Write and balance the half-reactions (and full reaction) of Eqn 1 from the
introduction/background above.
(b) Indicate what is oxidized and what is reduced.
(c) What is the ratio of moles of iron (II) ions to moles of permanganate ions?
4) How many moles of Fe2+ ions will be oxidized by 0.043 moles of permanganate ions?
5) Iron can be analyzed by dissolving it in acid to convert Fe0 to Fe2+ ions and then titrating the Fe2+
ions with potassium permanganate. First 1.630 g of impure iron ore are dissolved in acidic
solution. This solution is titrated to a pink endpoint with 27.15 mL of a 0.020 M potassium
permanganate solution.
(a) How many moles of permanganate ion were consumed?
(b) Using the ratio you determined in research question 3c and your answer to 5a,
how many moles of Fe2+ were in the iron ore sample?
(c) How many grams of Fe2+ (NOT impure iron) were in the iron ore sample?
(d) Using your answer to 5c and the original mass, what is the percent of iron in the iron ore
sample?
6) What is a standard solution? Which solution will we standardize during this lab?
7) Summarize what will happen and/or the goal of Day 1.
8) Summarize what will happen and/or the goal of Day 2.
Materials - Day 1 (DON’T list in your lab book):
50 mL 0.100 M ferrous ammonium
sulfate
30 mL 6 M H2SO4
1 50 mL buret
1 ringstand
1 10 mL graduated cylinder
1 test tube brush
1 250 mL beaker (for cleaning and waste)
1 10 mL volumetric pipet
1 30 mL beaker
Materials - Day 2:
1 balance
(The rest of the materials are up to you!)
100 mL approx. 0.02 M KMnO4
3 100 mL beakers
1 double buret clamp
3 250 mL Erlenmeyer flasks
1 permanent marker
1 buret funnel
60 mL 3% hydrogen peroxide solution
1 pipet pump
iron pills
Hazards (list in your lab book):
(include the safety contract and the hazards of the potassium permanganate, ferrous ammonium
sulfate, hydrogen peroxide, and sulfuric acid)
Procedure - Day 1 (DON’T list in your lab book):
1. Clean beakers, flasks, and the graduated cylinder
2. Label glassware
a. label one 100 mL beaker KMnO4 and initials
b. label another 100 mL beaker Fe(NH4)2(SO4)2 and initials
c. label the buret KMnO4
d. label another 100 mL beaker H2O2 and initials
e. label the 30 mL beaker H2SO4 and initials
f. label the 250 mL beaker waste
3. Obtain solutions
a. obtain approximately 80 mL of potassium permanganate solution in a 100 mL
beaker
b. obtain approximately 50 mL of ferrous ammonium sulfate solution in a 100 mL
beaker
c. obtain approximately 10 mL of sulfuric acid in a 30 mL beaker
3. Set up the 50 mL buret
4. Prepare the buret
a. use the 250 mL beaker for waste
b. rinse the burets with about 10 mL of distilled water
c. rinse the buret twice with about 5 mL of potassium permanganate solution
5. Fill the buret to the zero mark with potassium permanganate solution
6. Using the pipet, transfer 10 mL of ferrous ammonium sulfate solution into the first
Erlenmeyer flask
7. Acidify the solution
a. measure out 10 mL of 6 M H2SO4 into the graduated cylinder
b. add it to the Erlenmeyer flask
c. swirl to mix
8. Titrate the ferrous ammonium sulfate
a. position the Erlenmeyer flask under the buret of potassium permanganate so that
the tip of the buret is in the flask but at least 2 cm above the surface of the liquid
b. slowly add potassium permanganate solution until the first trace of pink color
persists for 30 seconds
c. while adding the permanganate, swirl the flask and wash down the walls of the
flask with distilled water
d. record the final buret reading
9. Repeat steps 6-8 twice more for a total of three trials
10. Clean up
a. empty the buret into a 250 mL beaker
b. obtain approximately 60 mL of hydrogen peroxide in the third 100 mL beaker
c. add the hydrogen peroxide to the buret
d. drain into waste beaker
e. rinse the buret with distilled water again
f. draw tap water up into the pipet twice to rinse it out
Procedure - Day 2 (You will need to write your own procedure!):
· On this day, you will use enough iron pills to make up at least a 1 g sample, grind them into a
powder, make them into a solution, acidify the solution with sulfuric acid, and use your
knowledge from Day 1 to determine how much iron is actually in the pills.
Post-lab calculations:
1) Use the day 1 data and your answer to research question #3 to calculate the molarity of the
potassium permanganate solution:
(a) Determine how many moles of Fe2+ were used from the molarity and milliliters used in each
trial. (Note: It was the same for each trial!)
(b) Use the ratio determined in research question 3c to calculate the number of moles of KMnO4
used in each trial. (Note: It was the same for each trial!)
(c) Convert the milliliters of KMnO4 used in each trial to liters.
(d) Divide the moles of KMnO4 used in each trial by the liters used in each trial to calculate the
exact molarity of the KMnO4 used. Watch your sig figs!
(e) Average your three trials.
2) Calculate the grams of iron in each iron pill you used on Day 2 using the molarity you obtained
on Day 1 and the process used in research question 5.
3) Compare the actual amount of iron you calculated to the amount listed on the bottle to obtain a
percent error (assuming the amount on the bottle is correct).
Lab handout based on the experiment “Oxidation-Reduction Titrations” in Laboratory Experiments for
Advanced Placement Chemistry (Second Edition) by S.A. Vonderbrink (Flinn Scientific, 2006),
Investigation #8 of AP Chemistry Guided-Inquiry Experiments: Applying the Science Practices (College
Board, 2013), and the AP Chemistry Summer Institute materials of consultant Mark Case.