EXPERIMENT 3.4
HYDROGEN PEROXIDE TITRATION
INTRODUCTION
In your study of inorganic chemistry, you should have encountered some discussion
of the peroxides. Sodium and to some extent barium naturally form peroxides when
combined with oxygen under ordinary conditions. The molecular compound
hydrogen peroxide is useful as both an oxidizing and reducing agent. This dual
nature comes from the instability of oxygen in peroxides (at oxidation number -1)
relative to neutral oxygen (0) and normal oxides (-2).
In this experiment, you will analyze an unknown solution containing H2O2 as a
solute by titrating it with a known solution of the oxidizing agent KMnO4. In an
acidic solution, the purple colored MnO4- ion oxidizes H2O2 to form O2 while it is
reduced to the colorless Mn2+ ion. The color change allows KMnO4 to act as its own
indicator. As a solution of KMnO4 is added to an H2O2 sample, the added drops lose
color until the H2O2 is used up. The first drop of KMnO4 which is in excess retains
its purple color and marks the stopping point of the titration.
Using the given concentration of the KMnO4 solution and measurements of volume
from a buret, you will be able to calculate the moles of KMnO4 used. You will then
be required to balance the redox reaction equation to relate moles of KMnO4 to the
moles of H2O2 in a sample. You will convert this to a mass in grams to allow you to
calculate the percent by mass of the sample that was H2O2. This concentration will
be typical for commercial solutions of antiseptic H2O2.
MATERIALS AND SAFETY
Although the concentrations will not be high, oxidizing and reducing agents can be
corrosive an irritating. Wear safety goggles at all times as usual. The KMnO4
solution will stain skin with its purple color unless treated with a “sacrificial”
reducing agent. For this purpose, a flask of sodium bisulfite will be available. A
very small amount of the bisulfite can be placed on stained areas of skin to take out
the color. Handle the KMnO4 solution with care and clean up spills promptly. All
solutions used and produced in the experiment can be disposed of in the sinks when
finished.
PROCEDURE
Set up a buret with a buret clamp. Use distilled water to verify that the stopcock
functions properly with no leaks or air bubbles.
Use a small beaker to transfer KMnO4 solution from the reservoir to the buret.
Record the KMnO4 concentration as data. Weigh a clean, dry 250 mL Erlenmeyer
flask on the balance to the nearest 0.001 g. Remove it from the balance and add 10
mL of the unknown H2O2 solution using one of the 10 mL graduated cylinders
provided. (DO NOT ADD LIQUIDS TO CONTAINERS WHILE THEY ARE ON
THE BALANCES.) Immediately reweigh the flask with sample. Add about 50 mL
of distilled water and add 15 mL of dilute sulfuric acid from a storage bottle. Swirl
to mix the contents of the flask.
Record the initial level of KMnO4 solution in the buret. Read all buret volumes to
the nearest 0.01 mL. Titrate the sample by adding the KMnO4 solution to the flask
and swirling until a light pink/purple color begins to persist. Then add slowly until
a drop or fraction of a drop produces a pink color that lasts for one minute or more.
It may be helpful to have a white paper background under the flask. Record the
final buret level.
Repeat the process on a second and third sample of the H2O2 solution. Each group
member should perform at least one titration.
When finished, dispose of all solutions in the sink and rinse all glassware.
Carefully return the burets and pipets to the cart along with the buret clamps.
PRELAB QUESTIONS
1. Balance the redox reaction below. Show your work.
MnO4 - + H2O2  Mn2+ + O2 + H2O
2. At warm temperatures, H2O2 slowly decomposes according to the reaction
2H2O2  2H2O + O2
If you waited a long time between preparing and titrating your samples, state
and explain the specific effect on the calculated results (i.e., include whether
they will likely be high or low and trace the cause through the calculation steps).
3. If a sample solution is 5.00% H2O2 by mass and has a density close to 1.00 g/mL,
find the approximate molar concentration (molarity) of the H2O2 in the solution.
Show work.
DATA TABLES (write into lab notebook as part of prelab)
a) Molarity of standard KMnO4 solution_____________
TRIAL
b) Mass of empty flask
c) Mass of flask with sample
d) Initial buret level
e) Final buret level
f) Qualitative comments
1
___________
___________
___________
___________
2
___________
___________
___________
___________
3
___________
___________
___________
___________
POSTLAB CALCULATIONS
TRIAL
1
g) Sample solution mass
___________
h) Volume of KMnO4 used
___________
i) Moles of KMnO4 used
___________
j) Moles of H2O2 in sample
___________
k) Grams of H2O2 in sample
___________
l) Percent by mass of H2O2
___________
m) Average percent by mass of H2O2___________
n) Deviation from average
___________
o) Average deviation___________
p) Discussion
q) Problems and Errors
2
___________
___________
___________
___________
___________
___________
3
___________
___________
___________
___________
___________
___________
___________ ___________