HOW CAN WE DETERMINE THE ACTUAL PERCENTAGE OF H2O2 IN A
DRUGSTORE BOTTLE OF HYDROGEN PEROXIDE
CENTRAL CHALLENGE
This lab has two major tasks. The first task is to standardize the concentration of a KMnO4
solution. This task is necessary in order to complete the second task, which is to evaluate how
close commercial H2O2 solutions are to their labeled concentrations. Different groups of students
will work with different brands.
BACKGROUND
Hydrogen peroxide, H2O2, is easily oxidized. It is used in commercial bleaching processes and in wastewater treatment
plants as an environmentally alternative to chlorine. Dilute solutions of H2O2 are used to bleach hair and to clean wounds.
It readily decomposes in the presence of light, heat, or metallic catalysts into water and oxygen. It is important to know
the actual concentration of a solution of H2O2 as its effectiveness can decrease with smaller concentrations.
Oxidation–reduction (redox) reactions involve a transfer of electrons between the species being oxidized and the species
being reduced. The reactions are often balanced by separating the reaction components into two half-reactions: oxidation
(loss of electrons) and reduction (gain of electrons). In a redox reaction, the number of electrons lost by the species being
oxidized is always equal to the number of electrons gained by the species being reduced. In the reaction being studied in
this lab, solutions of hydrogen peroxide, H2O2, and potassium permanganate, KMnO4, will be combined in acidic solution.
Deep purple in solution, the Mn in KMnO4 undergoes reduction very easily. In acidic solution, permanganate ions
(MnO4 -) from KMnO4 reduce to nearly co colorless Mn2+. In the presence of permanganate ions in acidic solution, an
aqueous solution of H2O2 will undergo oxidation to make oxygen gas and hydrogen ions.
Solutions of KMnO4 are not easily standardized solely by preparation, as solid KMnO4 often contains impurities such as
chlorides, sulfates, and nitrates. KMnO4 can be standardized in acidic solutions with a known concentration of
Iron (II) ions. In this oxidation-reduction reaction, manganese (II) ions and iron (III) ions are formed.
Pre- Lab Questions
1. What measuring devices are used to obtain precise measurements of volumes? What measuring devices are
used to obtain approximate measurements of volume?
2. Write a balanced half- reaction for the reduction of permanganate ions in acidic solution. What the oxidation states
of manganese in this reaction?
3. Write a balanced half-reaction for the oxidation of hydrogen peroxide. What are the oxidation states of oxygen in
this reaction?
4. Write the overall balanced redox reaction for permanganate ions reacting with hydrogen peroxide in an acidic
solution? How many electrons are transferred in this reaction?
5. Write a balanced half-reaction for the oxidation of iron (II) ions.
6. Write the balanced redox reation for permanganate ions reacting with iron (II) ions in an acidic solution.
How many electrons are transferred in this reaction?
7. A sample of oxalic acid, H2C2O4, was analyzed using a standardized solution of KMnO4, 25.0 mL of oxalic acid
is titrated after heating, 12.30 mL of a 0.0226 M KMnO4 was added to the sample when a faint pink color was
observed.
6H+ (aq) + 2 MnO4- (aq) + 5 H2C2O4 (aq)  10 CO2 (g) + 8 H2O (l) + 2Mn2+
(a)
(b)
(c)
(d)
How many moles of MnO4- reacted with the given amount of oxalic acid.
How many moles of oxalic acid were consumed in the reaction.
What was the molarity of the oxalic acid solution?
If the density of oxalic acid solution was 1.00 g/mL, what was the percentage by mass of oxalic acid in the
solution? (hint: % mass = mass of solute/mass of solution)
Before coming to lab, please go www.chem.yorku.ca/courses/chem1000/equipment/pipette.html to read how to properly
use a pipette and propipette.
In the first part of the lab (Part A), you will standardize a MnO4- solution by using a solid standard known as
FAS – ferrous ammonium sulfate – Fe(NH4)2(SO4)2 . 6 H2O. The titration involves the oxidation of Fe2+ to Fe3+ and the
reduction of MnO4- to Mn2+. The solution is acidified by adding a few drops of sulfuric acid. The end point in this
titration is taken as the first permanent pink color that appears in the solution as a result of a slight excess of permanganate
ion. To “sharpen” the end point (make it easier to see by getting rid of the slight yellow color of the Fe 3+ ion and any
cloudiness that forms), we will use a few drops of phosphoric acid.
PROCEDURE PART A
1. Determine the mass of an empty weighing boat.
2. Add approximately 0.30 grams to 0.40 grams of FAS sample and determine mass of weighing boat and FAS
3. Transfer your FAS sample to a clean – but not necessarily dry 125 mL Erlenmeyer flask.
4. Add enough distilled water (25mL or so) to dissolve the FAS.
5. Add 5 – 10 drops of sulfuric acid to the FAS solution to ACIDIFY it.
6. Check to see that your buret is filled with permanganate solution so you can titrate. Refill if necessary
7. WITHOUT DELAY, titrate this solution with the MnO4-1 solution. After adding 5 drops or so of permanganate
solution, add 5 drops or so of phosphoric acid. Swirl the flask to mix.
8. Continue the titration until you reach the first pink color that persists for 15 – 30 seconds. You should be
constantly swirling your flask during the titration and occasionally rinsing the inside walls of the flask with
distilled water.
9. Repeat the titration (no more than two trials)
PROCEDURE PART B
1. Pipette 1.0 mL of the hydrogen peroxide solution into a clean – but not necessarily dry Erlenmeyer flask.
2. Add around 25 mL of distilled water and several drops of sulfuric acid, swirl to mix.
3. Titrate with the permanganate solution using the same procedure as part A. Perform 2 trials.
CALCULATIONS
Calculate the moles of KMnO4 solution needed to react with all the 0.100 M FAS solution for each trial.
1. Calculate the molarity of KMnO4 solution for each trial.
2. Calculate the average molarity of the KMnO4 solution.
3. Calculate the moles and mass of H2O2 in solution in each trial.
4. If the density of the H2O2 solution titrated was 1.00 g/mL, calculate the percent by mass of H2O2 in solution in
each trial. (Use 1.0 mL as the volume of your H2O2 solution)
5. Calculate the average percentage of the H2O2 solution.
POSTLAB QUESTIONS
1. What is a titrant? What is an analyte?
2. In this lab, did a substance serve as both a titrant and an analyte? If so, what is this substance? Explain
your answer.
3. How would the concentration of the KMnO4 in part A be affected if the following were observed about
the Erlenmeyer flask? Justify your answers
(a) When the titration was complete, the flask was colorless.
(b) When the titration was complete, the flask was very dark pink or purple.