LAB 7:
DETERMINATION OF HYDROGEN PEROXIDE
CONCENTRATION BY REDOX TITRATION
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This lab does NOT require a lab report.
You will work with a partner for this lab.
I Introduction
Hydrogen peroxide, H2O2, is a topical antiseptic commonly used to treat cuts and
wounds to prevent infection. Typically, it is purchased off the shelf in solutions that
are 3% H2O2 by mass. Hydrogen peroxide, however, is relatively unstable and
decomposes to form water and oxygen gas:
H2O2(aq) → H2O(l) + O2(g)
This decomposition can be caused by heating, aging (thus the expiration date on
peroxide bottles), or absorption of light (thus the opaque, brown bottles peroxide is
commonly sold in). In reality, therefore, your hydrogen peroxide is probably not 3%.
In this lab, you will determine the percent hydrogen peroxide in a given sample
freshly purchased from a store shelf. You will then determine the concentration of
hydrogen peroxide in a solution that has been heated and exposed to light.
To determine the concentrations of your samples, you will use a redox titration. A
redox titration is similar to an acid-base titration (which we’ve done), but uses an
oxidizing or reducing agent to react with the analyte instead of an acid or base. In
this case, we will use the common oxidizing agent potassium permanganate, KMnO 4.
In acidic solutions, potassium permanganate reacts with hydrogen peroxide according
to the following unbalanced equation:
MnO4-(aq) + H2O2(l) → Mn2+(aq) + O2(g)
Several observations will assist you in determining when your titration is over. First,
as the reaction occurs, you can see bubbles of oxygen forming in the solution; when
you no longer see these bubbles forming, all of the peroxide has been consumed and
you have reached the end of your titration. A second (and more precise) indication is
a color change. In our recent acid-base titration, it was necessary to add an indicator
(such as phenolphthalein) to produce this color change. In the reaction you are
performing today, however, no indicator is necessary because MnO4- and Mn2+ are
different colors: the permanganate ion (MnO4-) is purple (pink at low concentrations)
and the manganese(II) ion (Mn2+) is pinkish (but relatively clear at these low
concentrations).Thus when the solution changes from pink to clear, your titration is
over.
Equipment and Reagents
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Buret and ringstand
250-mL Erlenmeyer flask
1-mL Mohr pipet and pipet bulb
Graduated cylinder
Stirring rod
Magnetic stirrer and stirbar
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Fresh hydrogen peroxide, H2O2
“Exposed and expired” H2O2
0.0200 M KMnO4, potassium permanganate
4.5 M H2SO4, sulfuric acid
Distilled water
! Warnings!
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Although the potassium permanganate, KMnO4, used in this experiment is dilute, it
is still a potent oxidizer and stains skin and clothing. Wear goggles and aprons and
clean up any spills immediately with wet paper towels.
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The sulfuric acid, H2SO4, used is significantly concentrated. Wear goggles and
aprons when working with it and immediately neutralize any spills with Acid Lock.
If you spill any acid on your skin, immediately rinse with water.
Procedure
Preparing the Titrant (0.0200 M KMnO4, acidified)
1) Prepare your lab notebook and don your goggles and aprons.
2) Rinse your buret with a few milliliters of the permanganate solution: pour some
in, rotate the buret around to coat all the surfaces, and then rinse your
stopcock by opening it and allowing the permanganate to drain into a waste
beaker (Don’t let it go down the sink!). Do NOT rinse your buret with water.
3) Fill your buret with the permanganate solution. (Make sure your stopcock is
closed before you start to pour!)
Your permanganate solution is now ready for titration.
Preparing the Analyte (3% H2O2)
4) Using a 1-mL Mohr pipet, transfer 1.00 mL of fresh H2O2 to an Erlenmeyer flask.
5) Dilute the peroxide with 25 mL distilled water and add 10 mL of 4.5 M sulfuric
acid.
6) Add a stir bar and begin stirring.
Your analyte is now ready for titration.
Performing the Titration
7) Perform this titration like you would any other: slow down when you start to
see a lingering color change. This one is easy to overshoot, so go slow towards
the end: drop by drop.
8) Record your observations (watch carefully) and the volume of titrant used.
9) Repeat the titration and calculations, this time using the “expired” solution that has
been heated and exposed to sunlight.
10) Answer the attached AP Free Response question.
Data and Calculations (to be recorded on your lab Data Sheet)
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Title and Purpose
Procedure used
Volumes of all reagents used
Observations of all reactions
Balanced equation describing titration reaction
Molarity of hydrogen peroxide solutions (old and new). Include: mols KMnO4
used, mols H2O2, and molarity
? Questions to Answer
1) Mass percent is a form of concentration, just as molarity is. The mass percent
of a solute is defined as follows:
Mass Percent of Solute = (Mass of Solute) / (Mass of Solution) * 100
Calculate the molarity of 3.0 % H2O2.
Hints:
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To simplify the math, assume the solution has a mass of 100 g.
Assume the mass of solution = mass of solute + mass of water.
Assume the density of water is 1.0 g/mL.
See Examples 11.1 and 11.2 on pages 514 and 517, respectively.
2) The reaction that occurs during the titration can be described as follows:
MnO4-(aq) + H2O2(l) → Mn2+(aq) + O2(g)
a) Identify the elements reduced and oxidized. Identify the reducing and
oxidizing agents.
b) Balance this equation. The reaction occurs in acidic solution.