Chemistry Lab Manual 4th Edition
Experiment 9: Understanding Molar Volume via Sublimation of Carbon
Dioxide
The objective of the lab is to use the sublimation process of solid carbon dioxide (dry ice) to
calculate the standard molar volume (volume of one mol of a gas at STP). This experiment applies
the gas laws covered in the lecture portion of the course to accomplish this.
Background Information
The relationships of the properties of gases should be understood in order to better interpret the
observed phenomena. One of the first relationships that students learn in general chemistry is
Boyle’s Law, which describes the inverse relationship of pressure and volume, while the amount
(# of gas particles) and temperature remain constant. Another relationship is known as Charles’s
Law. This law describes the direct relationship between volume and temperature, while the gas
pressure and amount remain constant. While Boyle’s and Charles’s laws are important they do
not encompass all of the necessary relationships and cannot solely be used to solve all gas
problems. Avogadro’s law is important because it explains the notion that equal volumes of ideal
gases under the same conditions of temperature and pressure contain equal numbers of gas
particles. For example, one mol of one type of ideal gas occupies the same volume as one mol of
another type of ideal gas at the same temperature and pressure. The volume occupied by one mol
of a gas is the molar volume of that gas under the given conditions of temperature and pressure.
At standard conditions, one atmosphere (1 atm) pressure and 0°C (273.15 K), the molar volume
of any ideal gas is approximately 22.4 liters.
The ideal gas law in the equation below expresses the relation between pressure, P, volume, V,
number of mol of gas, n; and the absolute temperature, T, of the gas.
PV =nRT
R is the gas law constant and has a value of 0.08206 Lꞏatm/molꞏK for all gases independent of
temperature, volume, pressure or amount (# of mol).
It is well known that carbon dioxide (CO2) exhibits ideal gas behavior, however, we will perform
this experiment in non-ideal conditions. By employing the combined gas law (combination of
Boyle’s, Charles’s, and Gay-Lussac’s law) we can convert/correct the volume collected in nonideal conditions to reflect the volume that would have been collected at “standard conditions”
(STP). The combined gas law describes the relationship between pressure (P), volume (V) and
temperature (T) of a gas under two sets of conditions for the same sample. The equation,
=
,
is the mathematical representation of the combined gas law, whereby the amount of substance (#
of mol) is constant. We can identify conditions “initial” and “final” as laboratory and standard
conditions (STP), respectively. Ultimately, our goal will be to use the volume of CO2 sample at
standard conditions, which can be calculated using the following equation
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Chemistry Lab Manual 4th Edition
VSTP =
.
A plan of action to solve for the volume of the CO2 gas is initiated by first finding the mass of a
known volume of the gas at the present laboratory temperature and pressure. This information
will be disseminated at the beginning of your lab section. The mass of the CO2 gas is determined
by first weighing a flask filled with air and re-weighing the flask filled with carbon dioxide gas.
The flask is filled with carbon dioxide gas by allowing a piece of solid CO2 (dry ice) to sublime in
the flask since the atmospheric conditions are favorable for this process to occur. Sublimation of
the CO2 gas forces the initial air out of the flask and subsequently filled with CO2 gas at laboratory
pressure and temperature conditions. The volume of the gas at lab conditions is simply the volume
of the flask that it is contained in.
Mathematically, the procedure can be explained by the following series of equations.
First, the initial weighing of the seemingly “empty” flask is really the combined weight of the
Erlenmeyer Flask and the mass of the air filling the volume of the flask (with the stopper). Mass 1
(m1) includes the mass of the flask, rubber stopper, and the air contained within. The following
equation shows the breakdown of the aforementioned factors included in m1:
m1 = mflask w/stopper + mair
After placing the solid CO2 sample in the same Erlenmeyer Flask and allowing it to sublime, the
initial air is displaced and the second weighing is
m2 = mflask w/stopper + 𝑚
.
Subtraction of m1 from m2 is equal to the mass of CO2 minus the mass of the air. The 𝑚
be derived by adding the mair value to both sides of the equation.
m2 - m1 = 𝑚
𝑚
can
- mair
= m2 - m1 + mair
In order to solve for molar volume, your 𝑚 value is divided by the molar mass of CO2 (you are
expected to know this value!) to solve for the actual number of mol (n) present in your sample.
This step is followed by dividing n into the VSTP value
𝑛
.
Finally, the molar volume is calculated by using the equation
𝑀𝑜𝑙𝑎𝑟 𝑉𝑜𝑙𝑢𝑚𝑒 𝑉
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Chemistry Lab Manual 4th Edition
Pre-Laboratory Assignment
Be sure to show your work.
1. What is the molar mass of CO2?
44.01 g/mol
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2. Calculate the number of mol present in a 0.376 g dry ice sample.
0.376 g/ 44.01g/mol =0.009 mol
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3. Define sublimation.
The process on which a solid directly evaporates instead of going to a liquid
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4. Calculate the theoretical volume of the CO2 sample in question 2 if it is sublimed at STP.
V = 22.71 * 0.009 = 0.204 L
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5. Use the combined gas law to solve for the pressure of the CO2 sample if the sample is
confined to 275 mL and the temperature is 19.3 C. Use your answer from question 4 at
STP as your initial conditions.
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Chemistry Lab Manual 4th Edition
Lab Safety Notes
Wear full PPE in lab as each week. Your lab instructor or a designee will disseminate pieces of
dry ice to your group, however, it is very important that you do not handle the dry ice with your
Procedure
1. By using the balance located at the end of your lab bench, carefully weigh a stoppered
Erlenmeyer flask to three decimal places. Record the mass of the flask and stopper on the
Data Sheet. **the flasks may have a volume of 50 mL, 125 mL, or 250 mL**
2. Retrieve a piece of solid CO2 (dry ice) for each size of Erlenmeyer flask. **Larger sized
pieces should be used for larger flasks**
CAUTION: Do not handle the Dry Ice with bare hands because of the danger of frostbite.
3. Place the solid carbon dioxide in the un-stoppered Erlenmeyer flask and allow to
sublime/vaporize.
4. Place the stopper onto the flask immediately after all the solid has just
disappeared/sublimed.
5. Weigh the flask, stopper, and carbon dioxide vapor to three decimal places (0.001). Record
the mass on the data sheet provided.
6. After the flask, stopper, and contents have been weighed, fill the flask to the brim with
water.
7. Stopper the flask tightly. (some water will spill outside of the flask!!!) Use a paper towel
to wipe the excess water outside of the flask. The outside of the flask should be dry before
proceeding to the next step.
8. Carefully pour the water from the flask into a 100-mL graduated cylinder and stop at the
100 mL mark. If the Erlenmeyer flask is not empty, pour the contents of the graduated
cylinder out until and refill with the remaining water up to the 100 mL mark (or lower).
Repeat if necessary to measure the volume of the contents of the Erlenmeyer flasks. Record
on a Data Sheet the volume of water.
9. Make certain that you obtain the information regarding the barometric pressure and the
temperature of the laboratory.
10. Do a second and third determination and, if time is available a third determination.
Upon completion of the laboratory experiment and BEFORE you leave the lab, please wash
your hands thoroughly with soap and water!!!
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Chemistry Lab Manual 4th Edition
Data Sheet
Trial 1
50 mL
Mass of flask, stopper and air
(m1) in g
Mass of flask, stopper and CO2
(m2) in g
Volume of flask (Vinitial) in mL
Volume of flask (Vinitial) in L
Laboratory barometric
pressure (Pinitial)
(inches Hg)
Conversion of barometric
pressure (Pinitial) to
atmospheres (atm)
Laboratory Temperature (°F)
(Tinitial)
Laboratory Temperature (K)
(Tinitial)
Density of air at Tinitial and
Pinitial (given by instructor)
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Trial 2
125 mL
Trial 3
250 mL
Chemistry Lab Manual 4th Edition
Post-Laboratory Assignment
1. Using the Combined Gas Law, determine the volume of the experimental CO2 gas sample
under standard conditions (0°C and 1 atm).
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2. Since density (g/L) = mass (g)/volume (L), Calculate the mass of air (mair) using the
density of air given by your instructor.
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3. Calculate the 𝑚
and number of mol contained in each trial.
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4. Calculate the experimentally determined molar volume of CO2 at STP using data from
each trial at 50 ml, 125 mL, and 250 mL.
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5. Explain why CO2 is denser than air. How was this determined experimentally? Could
this have been determined without experimentation?
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Chemistry Lab Manual 4th Edition
6. Describe the impact that using different volume Erlenmeyer flasks had on the
experiment.
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7. Plot the number of mol present versus the volume of the flasks in an electronic
spreadsheet program. Turn this plot in with your short lab report next week.
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