Lab 20: Ideal Gas Law
Once the number of moles of O2 gas is calculated, the percent of H2O2 present in the solution can be determined. To do
this, you first need to calculate the theoretical number of moles of O2 there would be if the solution was 100% hydrogen
peroxide. This can be found by using the following equation:
For this experiment:
mL H2O2 used is the volume of H2O2 you actually use
(approximately 5 mL).
H2O2 density is 1.02 g/mL
1 mol H2O2 / 34.0 g H2O2 is the reciprocal (inverted
fraction). of the molar mass of H2O2 . The molar mass
of H2O2 is 34.0 g /mol, so this is equal to 1 mol H2O2 /
34.0 g H2O2.
1 mol O2 / 2 mol H2O2 is used since the decomposition
produces 1 mole O2 from 2 moles of H2O2 .
The units in the entire equation cancel to give moles
The percent hydrogen peroxide can now be found. To do this,
divide (n), the actual number of moles you calculated, by the
theoretical moles of O2 there would be if the hydrogen
peroxide were 100%. This number is then multiplied by 100%.
This value can now be compared to the 3% hydrogen peroxide shown on the label to see if any decomposition has
2 2 2
2 2 2 2 2
2 2 2 2
1 mol H O 1 mol O
Theoretical moles O = H O used × H O density × ×
34.0 g H O 2 mol H O
Actual moles O (n)
% H O = 100
Theoretical moles O
In this experiment, we will use yeast to accelerate the decomposition of the hydrogen peroxide into water and O2 gas.
Yeast contains the enzyme catalase, which is a catalyst for this reaction. You will add yeast activated in warm water to a
known amount of hydrogen peroxide and quickly seal off the system so that the O2 gas formed is collected in a graduated
cylinder. After measuring the total volume of gas produced, its temperature, and the atmospheric pressure, the ideal gas
law can then be used to calculate how many moles of O2 gas is formed. We can do this by solving the ideal gas law
equation for n.
Figure 3: Carbonated beverages contain dissolved CO2 at
high pressure. When the container is opened, this pressure
can create a powerful burst, such as with this sparkling
wine bottle, or when your soda “explodes.”
Lab 20: Ideal Gas Law
1. What is it in yeast that aids in the decomposition of hydrogen peroxide?
2. List the ideal gas law and define each term with units.
3. How many moles of O2 were produced in a decomposition reaction of H2O2 if the barometric pressure was
0.980 atm, the temperature was 298 K and the volume of O2 gas collected was 0.0500 L?
4. If you decomposed 10.00 mL of 100% H2O2, how many moles of O2 could you theoretically obtain?
Lab 20: Ideal Gas Law
Experiment: Finding Percent H2O2 with Yeast
1. Prepare the materials for the apparatus as shown in Figure 1. Insert the smaller rigid tubing into one end of
the larger, flexible tubing. Insert the free end of the rigid tubing securely into the rubber stopper hole.
2. Bend the free end of the flexible tubing into a U shape, and use a rubber band to hold this shape in place.
This will allow you to more easily insert this end of the flexible tubing into the inverted graduated cylinder.
Make sure the tubing is not pinched and that gas can flow freely through it.
2. Fill the 600 mL beaker with 400 mL distilled water.
3. Fill the 100 mL graduated cylinder with distilled water slightly over the 100 mL mark.
Figure 3: Gas Collection Apparatus (not to exact scale)
600 mL Beaker
Safety Equipment: Safety goggles, gloves
|Rubber band||Flexible tubing (18 in.)|
|2 Droppers (pipettes)||250 mL Beaker|
|Stir rod||600 mL Beaker|
|Warm water*||Ring stand*|
10 mL Hydrogen peroxide 10 and 100 mL Graduated cylinders Erlenmeyer flask Stopper with hole Rigid plastic tubing (3 in.) Large ring* Distilled water*
*You must provide *Optional Materials (not provided)
4. Take the temperature of the water in the 600 mL beaker, and record it in the Data section. Also, determine
the barometric pressure in the room, and record it in the Data section. HINT: The pressure in your region
may be found online—if necessary, convert this value to mm Hg.
5. Mix 100 mL of warm water (45°C) and 1 packet of baker’s yeast in a 250 mL beaker. This will activate the
yeast from the dormant (dry) state. Be sure to mix well with a stir rod until the yeast is completely dissolved.
6. Use a 10 mL graduated cylinder and pipette to measure out 5.00 mL of hydrogen peroxide. Pour this
hydrogen peroxide into the Erlenmeyer flask, and place the stopper with stopper tube over the top.
7. Clean the 10 mL graduated cylinder by rinsing it at least three times with distilled water. Dispose of the rinse
down the drain.
8. Cover the opening of the graduated cylinder with two or three fingers and quickly turn it upside down into
the 600 mL beaker already containing 400 mL of water. DO NOT remove your fingers from the opening until
the graduated cylinder is fully submerged under the water. If the amount of trapped air exceeds 10 mL, refill
the cylinder and try again.
9. Insert the U shaped side of the flexible tubing into the beaker, and carefully snake it into the submerged
opening of the graduated cylinder. You want as little air as possible to be in the graduated cylinder.
10. Secure the graduated cylinder to the ring stand by sliding a ring under the submerged cylinder, then attach‐
ing the ring to the stand.
If your kit does not include a ring stand, you will hold the graduated cylinder in
place while gas is collected. Make sure to keep the open end of the cylinder completely submerged to pre‐
vent additional gas from entering. Rest the graduated cylinder against the side of the beaker during experi‐
11. With the cylinder vertical, record the volume of air inside (the line at which the water reaches in the
cylinder) in the Data section in Table 1.
12. Using the pipette, measure out 5.00 mL of yeast solution into the rinsed 10 mL graduated cylinder. NOTE: Do
not immediately pour the yeast solution into the Erlenmeyer flask.
13. Prepare to place the stopper (still connected to the hose) on the Erlenmeyer flask. Reset the stopwatch.
14. Quickly pour the 5.0 mL of yeast solution into the Erlenmeyer flask. Immediately place the stopper securely
in the opening of the Erlenmeyer flask by twisting it down into the flask gently.
15. Start timing the reaction with the stopwatch.
16. Swirl the Erlenmeyer flask to mix the two solutions together.
17. You will begin to see bubbles coming up into the 100 mL graduated cylinder. HINT: If gas bubbles are not
immediately visible, make sure the stopper is on tight enough and the tubing is not leaking. You will need to
start over after correcting any problems.
18. Continue to swirl the Erlenmeyer flask and let the reaction run until no more bubbles form to assure the
reaction has gone to completion. This should take approximately 6‐10 minutes. HINT: Catalase works best
around the temperature of the human body. You can speed the reaction up by warming the Erlenmeyer flask
with your hands.
19. Record the time when the reaction is finished in Table 2 of the Data section, along with the final volume of
air in Table 1. Remember to read it at eye‐level and measure from the bottom of the meniscus.
20. Pour all other liquids down the drain and clean the labware.
Lab 20: Ideal Gas Law
Water temperature: ⁰C
Barometric Pressure: mm Hg
|Initial volume of air (mL)||Final volume of air after reac‐|
|Volume of O2 collected|
(Final volume ‐ initial volume)
Table 1: Volume data
|Time reaction started||Time reaction ended||Reaction time (s)|
Table 2: Reaction time data
The goal is to find the percentage of hydrogen peroxide in the solution! This can be found by working through the
1. Convert the temperature of the water from ⁰C to Kelvin (K). Use the equation K = ⁰C + 273. This will be your
value for absolute T or the temperature in Kelvin.
2. If necessary, convert the barometric pressure in the room from mm Hg to atmospheres (atm).
Divide the measured pressure from the Data section by 760 mm Hg. This will give you pressure
(P) in atmospheres.
T = ⁰C + 273 = K
P = mm Hg * = atm
760 mm Hg
3. Convert the volume of oxygen from mL to liters (L).
4. Rearrange the ideal gas law to solve for n.
5. You are now ready to solve for the number of moles of O2. Be sure the units cancel so that you end up with
only the moles of O2 left. Use the value for the constant R given:
Actual number of moles of O2 (n) = moles
V = mL * = L
R = 0.0821 L∙atm/mol∙K
6. Calculate the theoretical number of moles of O2 there would be if the hydrogen peroxide were 100%, and not
an aqueous solution.
To use the above equation, calculate the following:
— H2O2 volume is the volume (mL) of hydrogen peroxide used: Volume = mL H2O2
— H2O2 density is known: Density = 1.02 g/mL
Molar mass of H2 O2 = g H2O2/1 mol H2O2
Molar mass of H2 O2 reciprocal =
Now you have all of the information needed to solve the equation for the theoretical moles of O2. All you need
to do is fill in the blanks and do the calculations.
Theoretical moles of O2 =
Theoretical moles of O2 = mol
is the reciprocal of the molar mass of H2O2. First write the molar mass of H2O2 then
find the reciprocal.
* * *
1 mol O2
Theoretical moles of O2 = H2O2 volume * H2O2 density * * 2 mol H2O2
7. Find the percent hydrogen peroxide.
% H2O2 = * 100% = %
8. You can also easily determine the reaction rate. To do this, divide the total volume of oxygen collected by the
total time of the reaction.
Reaction rate = = mL/sec
Actual moles O2
Theoretical moles O2
Volume O2 (mL)
Reaction time (s)
1. Was the calculated percentage of hydrogen peroxide close to the same as the percentage on the label?
2. Considering that catalysts are not consumed in a reaction, how do you think increasing the amount of catalyst
would affect the reaction rate for the decomposition of hydrogen peroxide?
Lab 20: Ideal Gas Law