Determination of the Molar Volume of a Gas
An Italian chemist, Amadeo Avogadro, declared in 1811 that "equal volumes of gases at the same temperature and pressure contain an equal number of molecules." Given that information, it did not take long for early scientists to realize that a standard volume of each different gas had a different mass. Since the number of molecules in that standard volume is a constant, the "mole concept" was soon born. The mass of that fixed number of molecules (Avogadro's number) is called the molecular mass. In this experiment you are going to determine the molecular mass of the gaseous hydrocarbon, butane, from the mass of a measured volume of gas.
Unfortunately, the experiment is not that simple. When an insoluble gas is collected by the downward displacement of water, the water evaporates into the butane until the butane becomes saturated with water molecules. The amount of water which evaporates is related to the force exerted by the water molecules as they push themselves between the butane molecules. This force is called the vapor pressure of water. To correct your gas volume for the added water vapor, simply subtract this vapor pressure from the atmospheric pressure. The difference, the partial pressure of the butane alone, is then substituted in the ideal gas equation when calculating its molecular mass.
Materials: water trough, thermometer, 10-mL graduated cylinder, Bic lighter
Procedure:
1. Fill the water trough about 3/4 full of water. Place a butane lighter under water for about 30 seconds. Remove the lighter from the water, shake it, and dry it with a paper towel to remove as much water as possible. Determine and record the mass of the lighter in the data table.
2. Submerge a 10-mL graduated cylinder completely under the water, allowing it to completely fill with water and invert it. If air bubbles appear in the cylinder, remove it from the water and repeat the procedure until there are no air bubbles. Stand the cylinder in the pail.
3. Keeping the mouth of the cylinder under water at all times, raise the cylinder off the bottom of the pail, and place the butane lighter under the mouth of the inverted cylinder.
4. Open the release valve on the lighter and allow butane bubbles to bubble up into the cylinder until the level of the water inside the cylinder is just below the level of the water outside the cylinder.
5. Remove the lighter from the pail, shake it to remove any excess water, dry exactly it as you did before, and find the mass of the butane lighter recording the result in the data table.
6. Raise up the graduated cylinder until the level of water inside and outside the cylinder are equal. (This essentially ensures that the air pressure on the inside of the cylinder is equal to atmospheric pressure outside the cylinder and allows us to calculate the partial pressure of butane, given atmospheric pressure and the partial pressure of water vapor at the temperature of the water.) Record this volume of gas in the data table.
7. Measure the temperature of the water in the pail. We will assume that this is also the temperature of the butane gas.
8. Record the barometric pressure in the room. Repeat for one more trial.
Vapor Pressure of Water at Selected Temperatures
ºC / torr / ºC / torr / ºC / torr0 / 4.6 / 19 / 16.5 / 26 / 25.2
5 / 6.5 / 20 / 17.5 / 27 / 26.7
10 / 9.2 / 21 / 18.7 / 28 / 28.3
15 / 12.8 / 22 / 19.8 / 29 / 30.0
16 / 13.6 / 23 / 21.1 / 30 / 31.8
17 / 14.5 / 24 / 22.4 / 40 / 55.3
18 / 15.5 / 25 / 23.8 / 50 / 92.5
DATA
Trial 1 / Trial 2Atmospheric pressure (kPa à atm)
Initial mass of the butane source (g)
Final mass of the butane source (g)
Mass of butane (g)
Adjusted volume of butane in the cylinder (L)
Water temperature (ºC)
Vapor pressure of water at @ water temperature (atm)
Partial pressure of butane gas (atm)
Moles of butane in the cylinder (mol)
Molar volume of butane @ water temperature (L/mol)
Average molar volume of butane @ water temperature (L/mol)
% Error
(molar volume)
Analysis (show work for Trial 1 only)
1. Convert the barometric pressure in the room into atmospheres.
2. Calculate the mass of butane released from the butane lighter.
3. Determine the vapor pressure of water in the cylinder from the table of the vapor pressures of water at various temperatures. Convert torr to atm.
4. Calculate the partial pressure of butane in the cylinder using Dalton's Law of Partial Pressures.
5. Calculate the molar mass of butane, C4H10.
6. Calculate the number of moles, n, of butane gas contained in the cylinder using the Ideal Gas Equation, the mass of butane released (Question 2) and the molar mass of butane (Question 5). Make sure to convert ºC into K.
7. Since the volume of a gas is directly proportional to the number of moles of the gas, we can calculate the actual molar volume (L/mol) of butane at the temperature of the water during the experiment using a proportion for your experimental data. For example, if 0.003 mol had a volume of 0.046 L at the temperature of the water, then you can calculate the volume of 1 mol of butane by setting up the following proportion:
0.046 L = x L
0.003 mol 1 mol
` Calculate the actual molar volume (L/mol) of butane using your experimental data.
8. Calculate the theoretical volume of 1 mole of butane at the temperature of the water using the Ideal Gas Equation.
9. Compare your actual and theoretical molar volumes of butane at the temperature of the water (Questions 10 and 11) by calculating the percent error of the molar volume of butane.
Discussion & Conclusion
Be sure to discuss possible sources of error and HOW those errors would have affected your molar volume determination.