H2O2 Decomposition Lab II
(Reference:  Chen, C.,Ehrenkranz, D., Sconzo, P., Chemistry in Microscale: Advanced Laboratories, 11th Biennial Conference on Chemical Education, Georgia Tech University, 1990.)

PREPARATION

Polyethelene pipets – 2 per group
6 inch test tubes – 1 per group
Balances – 1 per group
Fresh bottle of hydrogen peroxide (3%)
50 mL 6 M H2SO4 (at reagent bench or 10 mL per group)
KMnO4 solution moles solute/ gram solution – NOT MOLARITY or MOLALITY
· Mass 150 mL beaker
· Add ~ 0.7 g KMnO4
· Mass beaker again
· Record concentration as moles of solute/ gram of solution

BACKGROUND

In this experiment we want to determine how much hydrogen peroxide (H2O2) is in a store-bought solution of H2O2, because in the SOLUTION there will be some H2O2 and the rest is mostly water.  In order to do this we will oxidize the H2O2 using potassium permanganate (KMnO4)- this is an example of an oxidation/reduction reaction, or REDOX reaction.  You will be able to determine when the decomposition is complete because the solution in the test tube will turn from clear and colorless to clear and pink (presence of the Mn2+ ions that you see on the PRODUCTS side of the reaction give the pink color).  You will have the mass of both of the reactants, H2O2 and KMnO4, and you know the concentration of the KMnO4 solution so you can calculate the mass of H2O2 reacted.  Once you have that you can calculate the mass percent of H2O2 in the H2O2 solution.

2MnO4- + 6H+ + 5H2O2  8H2O + 5O2 + 2Mn2+

The concentration units are like MOLARITY, except the units are number of moles of solute per kilogram of solution (mol/kg).  Remember that molarity units are number of moles of solute per liter of solution. These units are convenient because you to prepare a certain concentration of solution by mass, and since you are keeping track of the masses of the reactants during your experiment these units for the reactants will be more convenient for calculations.

PROCEDURE

1. Label, and then fill a pipet with H2O2 solution

2. Label, and fill a second pipet with KMnO4 solution

3. Mass both pipets and record

4. Add 7 drops of the H2O2 solution to a small test tube

5. Add 3 drops of 6M H2SO4 (CAUTION!!)

6. Add KMnO4 solution dropwise, mixing after each addition, until a faint pink color persists

7. Re-mass each pipet and record

8. Using same pipets do a second trial

9. Clean up!

DATA

 o TRIAL I TRIAL II Concentration of KMnO4, mol/kg 0 0 Concentration of KMnO4, mol/g 0 0 Initial Mass of H2O2 pipet, g 0 0 Final Mass of H2O2 pipet, g 0 0 Initial Mass of KMnO4, g 0 0 Final Mass of KMnO4, g 0 0
CALCULATIONS

1. Determine the number of MOLES of KMnO4 reacted

2. What is the MOLE to MOLE relationship between H2O2 and KMnO4

3. Determine the number of grams of H2O2 reacted

4. Determine the MASS PERCENT of H2O2 in the H2O2 solution

5. If the mass percent of commercial H2O2 is 3.0%, determine you PERCENT ERROR

Practice Calculations for BEFORE the Lab

You should read the lab before you do these calculations so that you will understand the procedure for determining each value.

Suppose you performed the experiment as described, you used 25.0g of H2O2 solution and 2.18g KMnO4 solution.  The concentration of your KMnO4 solution is 4.0 x 10–8 mol solute/kg solution.  Complete the following calculations to find MASS PERCENT of the H2O2 solution.

2MnO4- + 6H+ + 5H2O2 8H2O + 5O2 + 2Mn2+

1. Find mol/g from mol/kg.

2. Determine the number of moles of KMnO4 that reacted.

3. From the balanced reaction, determine the mole to mole ratio of H2O2 to KMnO4.

4. Determine the number of grams of H2O2 consumed in the reaction. (Hint: H2O2 pure, not grams of the solution, you already know that from the question).

5. Now determine the grams of H2O2 per gram of H2O2 solution then multiply by 100, this is the mass percent of H2O2 in the solution.

6. Determine the percent error in this experiment if the mass percent on the bottle of H2O2 solution read 3.1%.  For percent error you subtract the actual from the experimental and then divide by the actual, then multiply by 100.

You should read the lab before you do these calculations so that you will understand the procedure for determining each value.

Suppose you performed the experiment as described, you used 25.0g of H2O2 solution and 2.18g KMnO4 solution.  The concentration of your KMnO4 solution is 4.0 x 10–8 mol solute/kg solution.  Perform the following calculations to find MASS PERCENT of the H2O2 solution.

2MnO4- + 6H+ + 5H2O2 8H2O + 5O2 + 2Mn2+

1. Find mol/g from mol/kg.

(4.0 x 10–8 mol/kg) x (1000 g/kg) =   4.0 x 10–5 mol/g KMnO4

2. Determine the number of moles of KMnO4 that reacted.

(4.0 x 10–5 mol/g KMnO4) x (2.18 g KMnO4) =  8.72 x 10-5 mol KMnO4

3. From the balanced reaction, determine the mole to mole ratio of H2O2 to KMnO4.

5 mol H2O2 to 2 mol KMnO4

4. Determine the number of grams of H2O2 consumed in the reaction. (Hint: H2O2 pure, not grams of the solution, you already know that from the question).

(8.72 x 10-5 mol KMnO4) x (5 mol H2O2 / 2 mol KMnO4) x (34 g/mol H2O2) =  7.4 x 10-3 g H2O2

5. Now determine the grams of H2O2 per gram of H2O2 solution then multiply by 100, this is the mass percent of H2O2 in the solution.

(7.4 x 10-3 g H2O2) / (25.0 g H2O2 solution) x 100 =   3.0 % by mass H2O2

6. Determine the percent error in this experiment if the mass percent on the bottle of H2O2 solution read 3.1%.  For percent error you subtract the actual from the experimental and then divide by the actual, then multiply by 100.

(3.1 –3.0) / (3.1) x 100 =   ~3 % error

DISPOSAL

· Fresh KMnO4 should be made each year, old can be reacted with hydrogen peroxide
· Extra 6 M sulfuric acid can be saved
· Extra hydrogen peroxide should be tightly capped and kept in the dark
· Fresh yeast should be purchased each year
· Student solutions can all go down the drain

Ginger Chateauneuf, 2000.
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