H2O2 Decomposition Lab I
(Reference:  CH112 General Chemistry II Laboratory Manual, Department of Chemistry , Michigan Technological University, Wiley, 1998)
 

PREPARATION

10 g bakers yeast (each group needs the tip of a spatula full)
5 or 6 little caps to put the yeast in – 1 per group
1 250 mL Erlenmeyer for each group
1 gas collection apparatus for each group
1 ring stand for each group
1 clamp for each group
30 mL 3.0 % hydrogen peroxide (~ 5 mL for each group)
Distilled water for each group
 

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 decompose the H2O2 using bakers yeast.  This decomposition will give off oxygen gas (O2), see the reaction below, and then we can collect this gas and find out how much is given off.  Once you know the volume of  O2 given off you can then use temperature, pressure and the gas constant to calculate the number of moles of O2 produced.  From the decomposition reaction you know the mole to mole ratios of  O2 to H2O2, so you can find the moles and mass of H2O2.  Once you have that you can calculate the mass percent of H2O2 in the H2O2 solution.

         PV=nRT   2H2O2        2H2O  +  O2

Where “P” is the PRESSURE of O2 gas, “V” is the VOLUME of  O2 gas collected, “R” is the GAS CONSTANT (0.08206 L·atm/mol·K), “T” is the TEMPERATURE of the O2 gas, and “n” are the NUMBER OF MOLES of  O2 gas in the volume that was collected

PROCEDURE

Below is a diagram of the apparatus we will be using.


 

1. Set up your apparatus as shown, but DO NOT add the yeast or H2O2 yet

2. When you fill the gas collection bottle make sure there are no bubbles inside.  You may need to use a watchglass to cover the opening of the bottle while turning it upside down into the pan of water

3. Mass out about 5.0 mL H2O2 solution, record exact volume

4. Pour H2O2 solution into flask, add about 10 mL of distilled water and swirl gently

5. Obtain small red cap, fill with a small scoop of yeast

6. One of you will need to hold the gas collection bottle and tube, the other should then quickly drop the cap of yeast into the flask and IMMEDIATELY stopper the
flask

7. Gas flow will begin quickly through the tube to the gas collection bottle

8. One of you should gently swirl the flask to allow a complete reaction while the other person continues to hold the tube in the collection bottle

9. Allow reaction to proceed for 10 minutes

10. Take the temperature of the water in the pan when reaction is complete, record temperature

11. When complete, remove the gas tube, lift the collection bottle up so that the level of the water left inside the bottle is even with the water level in the pan

12. Slide the watchglass over the bottle opening, and carefully turn the bottle right-side-up without spilling any of the water inside the bottle

13. Wipe all water off the sides of the bottle, cover it with the watchglass and then mass it, record mass

14. Fill the bottle up with water, slide the watchglass over the mouth, wipe away all excess water, and mass that, record mass

15. Clean up!
 
 

DATA

Mass of H2O2 solution, g   _________________

Water Temperature, °C   _________________

H2O vapor pressure, mm Hg   _________________

Barometric pressure, mm Hg   _________________

Mass of bottle + water, part full  _________________

Mass of bottle, full    _________________

Mass of water displaced   _________________

Volume of water displaced, mL   _________________

Volume of O2 collected, mL   _________________

Volume of O2 collected, L   _________________
 
 

CALCULATIONS

1. Determine the pressure of collected oxygen, using Dalton’s Law of Partial Pressures.  (Use the Lange’s handbook to determine the vapor pressure of water.)

Ptotal = PO2  +  PH2O
 
 
 
 

2. Calculate the number of moles of O2 produced using the ideal gas law
 
 
 
 
 
 

3. From the balanced equation, now you can calculate the number of moles of H2O2 you started with
 
 
 
 
 
 

4. Next, calculate the number of grams of H2O2
 
 
 
 
 
 
 

5. Using the mass of the solution you started with and the number of grams of H2O2 find the mass percent of H2O2 in the solution
 
 









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, using 5.00 mL of an aqueous H2O2 solution, with a density of 1.01 g/mL.  The water temperature was 24 °C, and the barometric pressure was 745.2 mm Hg.  The masses of the gas collections bottles and water were:  partially full = 248.25 g and full = 306.03g

1. Calculate the pressure exerted by the collected O2 at the water temperature.  (Vapor pressure of H2O at 24 °C = 22.4 mm Hg)
 
 

2. Covert this pressure to atmospheres
 
 

3. Calculate the volume of collected O2 in liters
 
 

4. Convert the water temperature, in Celsius, to Kelvin
 
 

5. Calculate the number of moles of O2 produced, using the above values and the ideal gas law
 
 

6. From the balanced equation (in the lab BACKGROUND) determine the number of grams of H2O2 in the sample
 
 

7. Calculate the number of grams of H2O2 solution you started with
 
 

8. Using the grams of H2O2 and the mass of the initial solution determine the percent mass of H2O2 in the solution
 
 

9. If the solution was 3.0% mass percent H2O2 what would be your percent error?
 
 

ANSWERS to Practice Calculations

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, using 5.00 mL of an aqueous H2O2 solution, with a density of 1.01 g/mL.  The water temperature was 24.0 °C, and the barometric pressure was 745.2 mm Hg.  The masses of the gas collections bottles and water were:  partially full = 248.25 g and full = 306.03g

1. Calculate the pressure exerted by the collected O2 at the water temperature.  (Vapor pressure of H2O at 24 °C = 22.4 mm Hg)

745.2 mm Hg = PO2  +  22.4 mm Hg  PO2 = 722.8 mm Hg

2. Covert this pressure to atmospheres

722.8 mm Hg x (1 atm / 760 mm Hg) = 0.9511 atm

3. Calculate the volume of collected O2 in liters

306.03g – 248.25g = 57.78g H2O 57.78 g H2O = 0.05778 L H2O = 0.05778 L O2

4. Convert the water temperature, in Celsius, to Kelvin

273.15 K = 0 °C  273.15 + 24.0 = 297.15 K

5. Calculate the number of moles of O2 produced, using the above values and the ideal gas law

(0.9511atm x 0.05778 L) / (0.08206 L·atm/mol·K x 297.15 K) = 0.002254 moles O2

6. From the balanced equation (in the lab BACKGROUND) determine the number of grams of H2O2 in the sample

0.002254moles O2 x (2mol H2O2/1mol O2) x 34.00 g/mol O2 = 0.1533g H2O2

7. Calculate the number of grams of H2O2 solution you started with

5.00 mL x 1.01 g/mL H2O2 solution = 5.05g H2O2 solution

8. Using the grams of H2O2 and the mass of the initial solution determine the percent mass of H2O2 in the solution

(0.1533g H2O2 / 5.05g H2O2 solution) x 100 = 3.03% mass H2O2

9. If the solution was 3.0% mass percent H2O2 what would be your percent error?

((3.0 – 3.03) / 3.0) x 100 = 1.2% error
 
 

DISPOSAL
 · 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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