Chemistry 162
Lab 1: Behavior of gases
Part 1 Absolute Zero
This part is done as a class demonstration, but write it up as a lab.
Objective: Extrapolate some pressure-temperature data to determine the Celsius degree
value of absolute zero.
Materials and methods: Obtain five different temperature baths, large enough to
accommodate the bulb of the pressure-measuring device. The five baths should be:
• A cylindrical Dewar flask, half-filled with liquid nitrogen (LN2);
temperature = –195.8°C
• A cylindrical Dewar flask, half-filled with a slurry made from acetone (CH3COCH3)
and dry ice; temperature = –78.0°C
• A beaker, half-filled with ice water; temperature = 0.0°C
• A beaker, half-filled with room temperature water; temperature = TBD
• A beaker, half-filled with boiling water; temperature = 100.0°C
If you really are getting into how cold various cooling solutions are, check out
http://www2.uni-siegen.de/~pci/versuche/english/v105-2.html which has pictures
as well.
Procedure: Prepare the baths as described. Using the computer’s LoggerPro software
and the datalogger attached to the pressure-measuring device, monitor the pressure as
the bulb of the device is set sequentially into each of the five baths. When the pressure
has reached an equilibrium value, record it in your notebook in a table similar to the
one below:
Bath
Liquid nitrogen
Temperature (°C)
–195.8
Dry ice/acetone slurry
–78.0
Ice water
0.0
Room temp. water
Pressure
Boiling water
100.0
What are the units given for the pressure measurement?
Do remember to record the temperature of the room temp water bath.
Analysis: Using Excel, plot temperature (x-axis) versus pressure (y-axis). Make sure
there is enough room in the graph to be able to extrapolate the graph back at least a
hundred degrees. In other words, make the x-axis range from –300°C to 100°C. Usual
rules apply; give the graph a good title, and axes labels (with units).
Use the “Add Trendline” function to draw a best-fit line using a linear regression
algorithm, and give the equation for the line.
Determine the correlation coefficient for the best-fit line. In other words, find r2 for the
line. Then complete the following sentence that explains why the correlation coefficient
is useful:
“________ percent of the variation in the pressure is explained by the variation in the
temperature.”
Important statistical note: Most of the time when you plot a graph, you put the
independent variable (the one you control) on the x-axis and the dependent variable
(the one you can measure but have no direct control over) on the y-axis. When you use
the correlation coefficient and best-fit lines, you are initially assuming that the two
variables are independent of each other. Only after fitting the line and determining r2
can you say that the two variables depend on each other.
Explain how you will fulfill the objective for the first part of this exercise using the
graph you just made. Hint: use the word “extrapolate”.
So what is the value of absolute zero in °C?
Whose law does the linearity of your graph confirm?
The gas inside the ball of the pressure-measuring device is air. What assumption did
you make about the behavior of air as you extrapolated the data to absolute zero? Why
did you have to make this assumption? How good of an assumption is it?
Part 2 Molar Mass of CO2
This part of the experiment is done in pairs.
Objective: Calculate the molar mass of carbon dioxide gas using the ideal gas law.
Introduction: The ideal gas law gives the relationship between the characteristic
properties of a gas: PV = nRT where P is the gas pressure, V is the gas volume, n is the
amount (in this case, moles) of gas and T is the gas temperature.
Of course, the number of moles of a gas is equal to the mass of a given amount of gas
divided by the gas’s molar mass. To put it another way, the molar mass (MM) of a gas is
equal to the mass of a given amount of gas divided by the number of moles in that
amount of gas.
Symbolically, MM = m/n, where m is the mass of a given amount of gas.
The ideal gas law can be algebraically rearranged so that it is solved for the number of
moles of gas: n = PV/RT.
Combining the previous two equations:
MM = mRT/PV
In other words, you can find the molar mass of a pure gaseous substance by measuring
the mass of gas in a certain space, the temperature and pressure of the gas and the
volume of the space. Recall that R is the ideal gas constant, whose value can be found
in nearly any reference.
In this experiment, you will fill a flask full of carbon dioxide, weigh it and determine the
volume, pressure and temperature of the carbon dioxide and calculate the molar mass
of carbon dioxide using the equation above.
One complication is that weighing the carbon dioxide gas is hard; you can’t simply tare
the balance and add carbon dioxide. You will need to figure out the mass of the air
inside the flask and subtract that off of the mass of the flask in order to get just what the
flask weighs.
Materials
• 250 mL Erlenmeyer flask
• Electronic balance
• Barometer (attached to classroom wall)
• Small piece of dry ice
Safety
• Rubber stopper
• Thermometer
• 100 mL graduated cylinder
Dry ice sublimes at –78.5°C; that is the temperature of the dry ice in this experiment. It will cause burns
upon prolonged or repeated exposure to bare skin, so do not pick up the dry ice with your bare hands; use
a paper towel. Carbon dioxide and water are not considered hazardous, so all fluids can be poured down
the sink.
Procedure
1. Obtain the Erlenmeyer flask and rubber stopper; make sure they are dry. Push the
rubber stopper firmly into the flask (not so firmly that the stopper won’t come back out).
Weigh this on the balance, and record the value in your data table.
2. Remove the rubber stopper and obtain a small piece of dry ice, about one-third the
size of the rubber stopper. Do not handle the dry ice with your bare hands; use a piece
of towel paper as shown. Put the dry ice into the flask. Do not put the stopper on the
flask just yet.
3. Allow the dry ice to sublime completely; do not move or shake the flask. After the last
bit of dry ice is gone, firmly stopper the flask and weigh it. Record this value in your data
table. Immediately place the thermometer inside the flask (you don’t need the rubber
stopper anymore) and, after it settles, note the temperature in your data table.
4. A 250 mL Erlenmeyer flask is never exactly 250 mL; you will have to determine its
volume. Using the flask, water and graduated cylinder, determine the volume of the
flask to the nearest tenth of a milliliter. Take into account the rubber stopper’s volume!
Write the volume in your data table. Record how you determined the value for the
volume.
5. Measure the atmospheric pressure in the room by reading the barometer on the wall.
If you are not sure how to set the barometer or to read a Vernier scale, please ask the
instructor. Note this value in your data table.
6. Look up the density of air in g/mL from any resource; write this value in your data
table and also cite the reference. Calculate the mass of the air in the flask from the
information you have, and enter this value on the calculation table.
Data: Recorded in a table.
Analysis: Prepare a calculation table. Show the work for each calculation for trial one.
7. Calculate the mass of the flask alone, then the mass of the carbon dioxide gas in the
flask.
8. Finally, calculate the molar mass of carbon dioxide.
9. Repeat this experiment twice more (clearly you will not need to look up or measure
as many things, since most properties will have remained the same).
10. Each pair of students will record their three calculated molar masses on the board,
along with the first initial and last name of each student in the pair.
11. Determine your average molar mass and the percent error of your average molar
mass, using the reference value of 44.01 g/mol for carbon dioxide.
Class Data:
12. After all of the groups have entered their molar masses on the board, copy the
information down and use Excel to determine the mean, standard deviation and percent
error for the class data. Also determine the one-sigma range and the two-sigma range.
Tape the Excel spreadsheet into your notebook.
Conclusion: Give your group’s results, and the class results. Rate the precision of your
group’s mean as excellent (within one sigma of the class mean), good (within two sigma of the
class mean), or poor (an outlier). Rate the accuracy of your group’s mean as excellent (less
than 1% error), good (less than 5% error) or poor greater than 5% error). Explain any poor
rankings.
Questions
1. a. Why couldn’t you simply tare the flask and weigh the carbon dioxide directly after it
sublimed? (Recalculate the molar mass of carbon dioxide if you had simply tared the
flask and let the dry ice sublime)
2. Why did you need to measure the temperature of the sublimed carbon dioxide
immediately after you weighed the flask? How would your calculated molar mass have
changed (higher or lower) if you’d waited for the gas to reach room temperature?
3. Was the class as a whole consistent with the reference (true) molar mass of carbon
dioxide? In other words, was the true molar mass within two-sigma of the class mean?
4. By the way, what are you assuming about the behavior of the carbon dioxide?