1. What is the conversion from Celsius to Kelvin and why is this conversion useful when applying gas law calculations?
2. A room is heated from 15 °C to 30 °C. What are these temperatures in Kelvin?
3. Following question 2, what is the temperature change in Celsius? What is the temperature change in Kelvin?
4. What is the barometric pressure in your region? If necessary, convert from mm Hg to atmospheres (atm). Divide the measured pressure by 760 mm Hg. This will give you pressure, (P), in atmospheres.
5. Convert a volume of 564.3 mL of oxygen from mL to liters (L).
6. Rearrange the ideal gas law (PV = nRT) to algebraically solve for n, using the values you calculated above at a temperature of 30 °C.
experiment 1: ideal gas law – finding percent H2O2
Table 1: Temperature, Pressure, and Volume Data Temperature of Distilled H2O:
Room (or regional) Pressure (atm):
Initial Volume of Air (mL)
Final Volume of Air (after reaction) (mL)
Volume of O2 Collected (Final Volume – Initial Volume) 220C= 295K
43.7mL= 0.0437LTable 2: Reaction Time Data Time Reaction Started
Time Reaction Ended
Total Reaction Time 15: 26
1. Calculate the number of moles of O2 produced using the ideal gas law. Then, use this value to calculate the number of moles of hydrogen peroxide you began the experiment with. Hint: Use the balanced equation provided in the lab introduction.
2. Calculate the number of moles of hydrogen peroxide you would have if you used 5 mL of a pure hydrogen peroxide solution. Hint: The density of hydrogen peroxide is 1.02 g/mL.
3. Determine the percentage of hydrogen peroxide in your solution.
4. Was the calculated percentage of hydrogen peroxide close to the percentage on the label (3%)? Calculate percent error of your value.
5. 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?
6. What was going on in the graduated cylinder as the H2O was pushed out?
7. How would the number of moles (n) of O2 change if your atmospheric pressure was doubled and all other variables stayed the same?
experiment 2: charles’s law part 1
1. What did you observe as heat was added to the system over time? What did you observe as the system cooled down?
2. Consider the balloon and air inside the flask to be a closed system. Explain what happened to the balloon as heat was added by the environment.
experiment 3: charles’s law part 2
Table 3: Temperature vs. Volume of Gas Data Temperature Conditions
Volume (mL) Room Temperature
2.1 Hot Water
3.4 Ice Water
1. 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.
2. Create a graph of temperature and volume data. Place temperature (remember to use degrees Celsius) on the x-axis and volume (mL) on the y-axis. Leave room on the left side of your chart for temperature values below zero. Construct your graph on a computer program such as Microsoft Excel®. If you do not have a graphing program installed on your computer, you can access one on the internet via the following links: http://nces.ed.gov/nceskids/createagraph/ or http://www.onlinecharttool.com. 3. Draw a straight line of best fit through your data points, extrapolating the line until it intersects the (negative) x-axis. Why can you assume a linear relationship (a straight-lined slope)? Hint: To extrapolate the trend line, change the minimum values on the x-axis and y-axis to negative values until the line intersects with the x-axis.4. At what temperature does your line intersect the x-axis? What volume corresponds to this temperature?5. Would it be possible to cool a real gas down to zero volume? Why or why not? What do you think would happen before that volume was reached? 6. Is your measurement of absolute zero close to the actual value (-273 °C)? Calculate a percent error. How might you change the experiment to get closer to the actual value?