Student Name: Click here to enter text.
Access Code (located on the lid of your lab kit):
Pre-Lab Questions:
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1. Write the balanced equation for the decomposition of hydrogen peroxide.
2. According to Charles’s law, what is the relationship between temperature and pressure?
3. Atmospheric pressure depends on the altitude (or height) of your location.
How should the air pressure change if you were in Denver, Colorado, which is 1.5 kilometers (1 mile) above sea level?
4. 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?
Experiment 1: Charles’s Law
Table 1 Temperature vs. Volume of Gas Data
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Temperature Conditions
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Temperature (˚C)
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Volume (mL)
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Room Temperature
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Hot Water
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Ice Water
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Post-Lab Questions
1. Use a pencil and graph paper to create a graph of temperature and volume data. Place temperature on the x-axis (in Kelvin) and volume (mL) on the y-axis. Leave room on the left side of your chart for temperature values below zero. You can also use a graphing program to create your graph. Using a ruler, draw a straight line of best fit through your data points, extrapolating the line until it intersects the (negative) x-axis.
2. Why can you assume a linear relationship (a straight-lined slope)?
3. What happened to the volume of gas when the syringe was exposed to various temperature conditions? Using the concepts explored in the Introduction, describe why this occurred, keeping in mind the definition of temperature.
4. At what temperature does your line intersect the x-axis? What volume corresponds to this temperature?
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Experiment 2: Using the Ideal Gas Law
Table 2 Temperature, Pressure and Volume Data
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Temperature of Distilled H2O:
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Room (or regional) Pressure (atm):
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Initial Volume
of Air (mL)
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Final Volume of Air
(after reaction) (mL)
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Volume of O2 Collected
(Final Volume – Initial Volume)
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Table 3 Reaction Time Data
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Time Reaction Started
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Time Reaction Ended
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Total Reaction Time
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Post-Lab Questions
1. What would happen if you added more than five mL of yeast to the H2O2?
2. What would happen if you added more than 5 mL H2O2 to the 5 mL of yeast?
3. What was going on in the graduated cylinder as the H2O was pushed out?
4. How would the number of moles (n) of O2 change if your atmosphere was doubled and all other variables stayed the same?
5. How would the number of moles (n) of O2 change if your temperature was doubled and all other variables stayed the same?
6. In this experiment, the temperature of the gas evolved is equal to the temperature of the water in the beaker, which ideally should be the same as the air temperature. Explain how the volume of oxygen evolved would change if you used ice water instead of room temperature water. How would it change if you used boiling water?
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