The Butane Lab
One of the important uses of the Ideal Gas Law is to determine the molar mass of a gas. To
measure the molar mass of a gas we determine the mass of a given sample of the gas under known
conditions of temperature and pressure. If the pressure is in atmospheres, the volume in Liters
and the temperature in K, then the gas constant, R, is 0.0821 Liter-atm/mol -K. The number
of moles can be expressed as (grams of gas/Molar Mass of gas). PV= nRT is then rearranged to:
MM = g R T
PV
Purpose: To determine the molecular mass of butane gas
Materials: 250 ml Erlenmeyer flask, glass plate, butane lighter, trough, thermometer
Procedure:
1. Be sure the lighter is clean and dry and will turn up high. Weight the lighter.
2. Fill a 250 ml Erlenmeyer flask full of water. Cover the top with a glass plate. Fill
the trough half full of water. Holding the glass plate in place, turn the flask upside
down and place the mouth under the water in the trough. There should be no air
bubbles in the flask. If there is, redo this part. If not remove the glass plate.
3. Hold the butane lighter under the water beneath the mouth of the inverted flask.
Press the release lever and be sure all the gas enters the flask. Hold the lever down
and collect approximately 200 ml of the gas. Then remove the lighter and let it dry
completely. Weigh it only when you are sure it is completely dry.
4. Raise or lower the flask until the level of water inside the flask is the same as the
water level outside the flask. When you have done this, the pressure inside the flask
equals the atmospheric pressure. Mark the water level on the flask with a wax pencil.
5. Remove the flask, fill it with water to the mark and pour the water into a graduated
cylinder to accurately determine the volume of gas that was in the flask.
6. Record the temperature of the water in the trough and the barometric pressure.
Data
Mass of lighter before the experiment
_______________
Mass of lighter after the experiment
_______________
Mass of butane used in the experiment
_______________
Volume of Gas collected in ml
_______________
Volume of gas collected in Liters
_______________
Temperature of the water in Celsius
_______________
Temperature of water in Kelvin
_______________
Atmospheric pressure (barometer reading)
_______________
Calculations
1. Using Dalton's Law of Partial Pressure, calculate the pressure of the dry butane
Pressure of dry butane = Barometric Pressure - Vapor Pressure of water (at
temperature of water)
Temperature (oC )
18
19
20
21
22
Vapor Pressure of Water at Various Temperatures
Pressure (mm Hg)
Temperature (oC )
Pressure (mm Hg)
15.6
23
21.0
16.5
24
22.4
17.5
25
23.7
18.6
26
25.2
19.8
27
26.7
2. Convert the pressure of the dry butane from mm Hg to atmospheres.
3. Using the formula above, find your experimental molecular mass of butane.
V= Volume of gas collected in liters
P = pressure of dry butane in atmospheres
g = mass of butane used in the experiment R = 0.0821 L-atm/mol-Kelvin
T = Temperature of water in Kelvin
4. Now find the theoretical molecular mass of butane from its formula, C4 H10.
5. Compare your experimentally determined molecular mass of butane to the
theoretically determined molecular mass by determining your percent accuracy error.
Conclusion:
Source of Error: Give this part some thought. Do not give as a source of error
anything you are supposed to know at this point such as; calculations may be wrong or
the masses may not be correct. Instead evaluate the procedure to see if there is any
area where even with the greatest of care error may have occurred.