H_{2}O_{2}Decomposition Lab I(Reference: CH112 General Chemistry II Laboratory Manual, Department of Chemistry , Michigan Technological University, Wiley, 1998)

PREPARATION10 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

BACKGROUNDIn this experiment we want to determine how much hydrogen peroxide (H

_{2}O_{2}) is in a store-bought solution of H_{2}O_{2}, because in the SOLUTION there will be some H_{2}O_{2}and the rest is mostly water. In order to do this we will decompose the H_{2}O_{2}using bakers yeast. This decomposition will give off oxygen gas (O_{2}), 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 O_{2}given off you can then use temperature, pressure and the gas constant to calculate the number of moles of O_{2}produced. From the decomposition reaction you know the mole to mole ratios of O_{2}to H_{2}O_{2}, so you can find the moles and mass of H_{2}O_{2}. Once you have that you can calculate the mass percent of H_{2}O_{2}in the H_{2}O_{2}solution.PV=nRT 2H

_{2}O_{2}2H_{2}O + O_{2}Where “P” is the PRESSURE of O

_{2}gas, “V” is the VOLUME of O_{2}gas collected, “R” is the GAS CONSTANT (0.08206 L·atm/mol·K), “T” is the TEMPERATURE of the O_{2}gas, and “n” are the NUMBER OF MOLES of O_{2}gas in the volume that was collected

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

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

_{2}O_{2}yet2. 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 H

_{2}O_{2}solution, record exact volume4. Pour H

_{2}O_{2}solution into flask, add about 10 mL of distilled water and swirl gently5. 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

flask7. 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!

DATAMass of H

_{2}O_{2}solution, g _________________Water Temperature, °C _________________

H

_{2}O 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 O

_{2}collected, mL _________________Volume of O

_{2}collected, L _________________

CALCULATIONS1. 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 = PO

_{2}+ PH_{2}O

2. Calculate the number of moles of O

_{2}produced using the ideal gas law

3. From the balanced equation, now you can calculate the number of moles of H

_{2}O_{2}you started with

4. Next, calculate the number of grams of H

_{2}O_{2}

5. Using the mass of the solution you started with and the number of grams of H

_{2}O_{2}find the mass percent of H_{2}O_{2}in the solution

Practice Calculations for BEFORE the LabYou 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 H

_{2}O_{2}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 = 306.03g1. Calculate the pressure exerted by the collected O

_{2}at the water temperature. (Vapor pressure of H_{2}O at 24 °C = 22.4 mm Hg)

2. Covert this pressure to atmospheres

3. Calculate the volume of collected O

_{2}in liters

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

5. Calculate the number of moles of O

_{2}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 H

_{2}O_{2}in the sample

7. Calculate the number of grams of H

_{2}O_{2}solution you started with

8. Using the grams of H

_{2}O_{2}and the mass of the initial solution determine the percent mass of H_{2}O_{2}in the solution

9. If the solution was 3.0% mass percent H

_{2}O_{2}what would be your percent error?

ANSWERS to Practice CalculationsYou 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 H

_{2}O_{2}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 = 306.03g1. Calculate the pressure exerted by the collected O

_{2}at the water temperature. (Vapor pressure of H_{2}O at 24 °C = 22.4 mm Hg)745.2 mm Hg = PO

_{2}+ 22.4 mm Hg PO_{2}=722.8 mm Hg2. Covert this pressure to atmospheres

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

0.9511 atm3. Calculate the volume of collected O

_{2}in liters306.03g – 248.25g = 57.78g H

_{2}O 57.78 g H_{2}O = 0.05778 L H_{2}O =0.05778 L O_{2}4. Convert the water temperature, in Celsius, to Kelvin

273.15 K = 0 °C 273.15 + 24.0 =

297.15 K5. Calculate the number of moles of O

_{2}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 O26. From the balanced equation (in the lab BACKGROUND) determine the number of grams of H

_{2}O_{2}in the sample0.002254moles O

_{2}x (2mol H_{2}O_{2}/1mol O_{2}) x 34.00 g/mol O_{2}=0.1533g H_{2}O_{2}7. Calculate the number of grams of H

_{2}O_{2}solution you started with5.00 mL x 1.01 g/mL H

_{2}O_{2}solution =5.05g H_{2}O_{2}solution8. Using the grams of H

_{2}O_{2}and the mass of the initial solution determine the percent mass of H_{2}O_{2}in the solution(0.1533g H

_{2}O_{2 }/ 5.05g H_{2}O_{2 }solution) x 100 =3.03% mass H_{2}O_{2}9. If the solution was 3.0% mass percent H

_{2}O_{2}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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