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CW 6: Molar Mass of Butane
Experimental Background:
Ideal Gas Law
𝑃𝑉 = 𝑛𝑅𝑇
Common Gas Law Constant Values
𝑅=
8.31 𝐿 ∙ 𝑘𝑃𝑎
𝑚𝑜𝑙 ∙ 𝐾
𝑅=
0.0821 𝐿 ∙ 𝑎𝑡𝑚
𝑚𝑜𝑙 ∙ 𝐾
𝑅=
62.364 𝐿 ∙ 𝑚𝑚𝐻𝑔
𝑚𝑜𝑙 ∙ 𝐾
The ideal gas law may be rearranged to solve for moles:
𝑃𝑉
𝑛=
𝑅𝑇
Molar Mass
If you solve for the moles of a gas, and you know the mass of the gas, you can determine
molar mass of the gas.
𝑚𝑎𝑠𝑠 (𝑔𝑟𝑎𝑚𝑠)
𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 =
𝑛 (𝑚𝑜𝑙𝑒𝑠)
Empirical Formula
If you know the composition of the elements in the gas, you can find the empirical formula.
1. Convert the grams of each element into moles of each element using the correct molar
mass. NOTE: If you are given percent composition, assume that there are 100 grams
total of the sample, so that the percent is equal to the number of grams.
2. Choose the smallest number of moles you found in step 1. Divide each of your answers
from step 1 by that smallest number of moles.
3. Your answers from step 2 should all be whole numbers. If not, multiply each number by
some factor such that you get while numbers. These will tell you the ratio of each
element in the chemical formula.
Molecular Formula
If you know the empirical formula and the molar mass of the gas, you can find the molecular
formula.
1. Find the empirical formula.
2. Determine the molar mass of the empirical formula.
3. The problem will give you the molecular mass of the molecular formula. Divide this mass
by the molar mass of the empirical formula from step 2.
4. Your answer to step 3 should be a whole number. Multiply the subscripts in the
empirical formula by this number to get the molecular formula.
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Pre-lab Exercises:
A sample of gas is collected and found to have a mass of 0.15 g, at a temperature of 22°C, and
a pressure of 0.994 atm. The volume of the gas is 79.37 mL.
1. Using the ideal gas law, solve for the moles of the gas. (See Ideal Gas Law in the
Experimental Background section.)
2. Using the molar mass equation, solve for the molar mass of the gas. (See Molar Mass in
the Experimental Background section.)
3. The percent composition of the gas is 52.14% C, 13.13% H, and 34.73% O. Find the
empirical formula. Then find the molar mass of the empirical formula. (See Empirical
Formula in the Experimental Background section.)
4. Using the molar mass you found in question 2, and the molar mass of the empirical
formula in question 3, find the molecular formula of the gas. (See Molecular Formula in
the Experimental Background section.)
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Experimental Goals:
 Use the ideal gas law to determine the molar mass and molecular formula of a gas.
Materials:
 Goggles
 Electronic balance
 Barometer
 Large beaker
 Butane lighter
 Thermometer
 100 mL Graduated cylinder
 Rubber stopper (#5)
 Answer sheet
CAUTION: Butane is a flammable gas.
Procedure:
1. Fill the large beaker with room temperature water.
2. Submerge the disposable butane lighter in water, remove it, and then dry it off as
thoroughly as possible.
3. Record the temperature of the water, as well as the barometric pressure.
4. Weigh the disposable butane lighter to the nearest 0.01 g and record the mass.
5. Submerge the graduated cylinder in the beaker and fill it completely with water. There
should be no air bubbles in the cylinder.
6. Place the butane lighter underneath the opening of the graduated cylinder and fill the
cylinder with butane by holding down the trigger on the butane lighter. Do not touch
the striker or flint. Be careful not to let any of the gas escape around the graduated
cylinder.
7. Displace about 90 mL of water from the graduated cylinder. To collect the last 10 mL of
butane, adjust the height of the graduated cylinder so that the 100-mL mark lines up
with the level of water in the trough or beaker. Fill the graduated cylinder to the 100-mL
mark with butane. (This ensures that the pressure of gas inside the graduated cylinder
will be the same as the atmospheric pressure.)
8. Dry off the butane lighter and weigh it again. Record the mass.
9. See the Molar Mass of Butane Worksheet for calculations.
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Molar Mass of Butane Worksheet
Data:
Water bath temperature
Barometric pressure
Volume of gas collected
Initial mass of butane lighter
Final mass of butane lighter
Mass of butane collected
Analysis:
1. Consider how you collect the butane. What two gases are present in the graduated
cylinder?
2. What is the vapor pressure of water (PH2O) at the water bath temperature? (See unit
packet p. 27.)
3. Determine the partial pressure of butane: Pbut = Patm - PH2O.
4. Use the ideal gas law to solve for the moles of butane gas.
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5. Determine the experimental molar mass of the butane.
6. Butane is 82.66% C and 17.34% H. Find the empirical formula of butane.
7. Find the molar mass of the empirical formula.
8. Using the experimental molar mass and the molar mass of the empirical formula,
determine the molecular formula of butane.
Conclusion:
1. The molecular formula of butane is C4H10, with a molar mass of 58.12 g/mol. Determine
the percent error.
(𝐴𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑉𝑎𝑙𝑢𝑒 − 𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑉𝑎𝑙𝑢𝑒)
% 𝐸𝑟𝑟𝑜𝑟 = |
| × 100%
𝐴𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑉𝑎𝑙𝑢𝑒
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2. Before you massed the lighter, you got it wet and dried it off. Given that the lighter gets
wet when collecting the butane gas, why was this step necessary?
3. For each measurement you made, give a source of error, and explain how it affected the
experimental molar mass of butane.
Measurement Source of error
How it affected the experimental molar mass of butane
Water bath
temperature
Barometric
pressure
Volume of
gas collected
Initial mass of
butane lighter
Final mass of
butane lighter
Mass of
butane
collected