Download Lab # 28: Calculating the Value of the Ideal Gas Constant “R”

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Lab # 28: Calculating the Value of the Ideal Gas Constant “R”
Chemistry 1
Disposable plastic lighters are filled with lighter fluid, mostly butane (C4H10), which liquefies under pressure. In
this lab, you will use the ideal gas law to calculate the value of the ideal gas constant “R”.
We will collect a sample of gas from a lighter by water displacement. See Figure 1 for the set up. Here are a
few “nuts and bolts” of the lab to keep in mind:
- The volume of gas, V, is simply the volume of water displaced.
- Since the gas is collected at atmospheric pressure, the total pressure of the gas, Ptotal, is the same as the
barometric (atmospheric) pressure.
- The mass of the lighter fluid released to displace the water can be calculated by subtracting the initial
mass of the lighter from the mass after the lab has been carried out.
- The number of moles of butane collected “n” can be calculated using this change in mass and the molar
mass of butane.
- Using the ideal gas law, and knowing P, V, n, and T, we can calculate R, the ideal gas constant.
You will need to make one correction to determine the amount of gas
collected. Whenever a gas is collected over water, the result is a mixture of
the collected gas and water vapor. Dalton’s law of partial pressure states that
the total pressure of any gas mixture is equal to the sum of the component
pressure of each of the gases. In this experiment,
Patmosphere = Pgas + Pwater vapor
Ptotal=Pgas+PH2O vapor
H2O level
inside=outside
The vapor pressure of water depends only on the temperature, and can be
obtained from a reference table. By difference, we can readily calculate the
vapor pressure due to the gas alone.
Hold lighter in
this position
Safety: Wear goggles. Butane is very flammable. DO NOT FLICK THE
LIGHTERS!
PROCEDURE
Figure 1
1. Record the barometric pressure.
2. Obtain a lighter. Dip in water, then dry as well as possible with a paper towel; do this very carefully. Mass
the lighter to the nearest 0.01 gram. This is the initial mass of the lighter.
3. Fill a trough or large beaker with room temperature water (adjust the water temperature so it is room
temperature ±3°C). Record the temperature.
4. Fill a 100-mL graduated cylinder to the top with room temperature water from the trough by submerging the
cylinder in the trough as demonstrated by your teacher. Flip the graduated cylinder upside down without
allowing air to enter the cylinder. There should be no air bubbles trapped in the cylinder.
5. As one partner holds the cylinder in place, the other should hold the lighter under the water, with the valve
in the position shown in Figure 1, just below the cylinder. Press the valve and release gas until you have
displaced a little less than 100 mL of water from the graduated cylinder. Adjust the level of the cylinder
until the water level is the same inside the graduated cylinder and in the trough.
6. Remove the lighter and dry very carefully with a paper towel. Determine the mass of the dried lighter and
record.
7. Remove the cylinder from the trough, cover, bring to the hood, and pour out the gas (it is heavier than air).
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Determining the Value of “R”
Data Table 1.
1. Initial mass of lighter
2. Final mass of lighter
3. Mass of butane collected
( # 1 - # 2)
4. Volume of butane gas collected
5. Atmospheric pressure (in Hg)
6. Atmospheric pressure (mm Hg)
7. Room temperature (K)
Calculations. SHOW YOUR WORK!
1. Use Table 2 to determine the vapor pressure of water at the temperature of your experiment.
2. Calculate the partial pressure of dry butane by correcting for the partial pressure of water at room
temperature. Remember the total pressure of wet butane is the same as the atmospheric pressure in mm Hg.
Patmosphere = Pbutane + Pwater vapor.
3. Convert the pressure of butane found in question number 3 to atm.
4. Calculate the molar mass of butane (C4H10).
Table 2. Vapor Pressure of Water
Temperature, °C
Pressure, mm Hg
18
15.5
19
16.5
20
17.5
21
18.6
22
19.8
23
21.1
24
22.4
25
23.8
5. Calculate the number of moles of butane gas collected by using
the mass of butane collected recorded in Table 1 and the molar mass of butane.
6. Using the values in Data Table 1 and calculations 3 and 5, calculate the value of “R”, the ideal gas
constant.) Hint: Think about the values you need to use the ideal gas law, PV=nRT. Show your work.
Include all units.
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Synthesis: SHOW YOUR WORK and Explain!
1. One group eliminated Step 2 of the Procedure and used the mass of the lighter straight from the stockroom
as the initial mass of the lighter. This introduces a potential source of error. Describe the error.
2. Why must the Pwater be subtracted from the total pressure?
3. A student rushed through the experiment and did not equalize the water levels inside and
outside the graduated cylinder. The water level in his graduated cylinder was 40 mm
higher than the water level in the beaker (see Figure 2).
a. Which pressure was greater, the pressure of the collected gas, or the atmospheric
pressure?
40 mm
Figure 2
b. With your answer to “a” in mind, what assumption did the student make that
introduced error in the calculation of the value of “R”?
4. Calculate the % Error of your experimental value of R. (The accepted value is 0.082 L atm/mol K)
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