Handout: Determining the Molar Mass of Butane
Name ___________________________ Class ________
Teacher__________________________
Introduction
“Mr. Zippo”, also known as George G. Blaisdell, invented the Zippo™ lighter in 1932. Like
many others, he was looking for financial stability during the Depression. Originally, he
patterned it after an Austrian lighter, improving the appearance, but these didn’t sell. He
tried again, this time making it smaller, adding a hinged lid, using what’s called a “ wind
hook” around the wick, and marketing the Zippo™ with the first lifetime guarantee. It sold
for $1.95. Since then, these lighters have become extremely popular. Perhaps it is because
of their resilience and utility. (One story is told of a Zippo™ that lit on the first try—after
being removed from a fish.) Zippo’s TM are especially known for their utility in war. Soldiers
have carried them since World War II, using them for everything from signaling helicopters
to storing salt that would replenish what was lost sweating*.
Zippo™ lighters are different from everyday, plastic lighters because they contain lighter
fluid, not butane. A plastic lighter like the ones you will be using today contains only
butane, C4 H10. Why use butane to study gasses and gas laws? First of all, butane is easily
collected, as we will show today. Most important, though, is that butane is close to “ideal” at
standard temperature and pressure. Ideal gasses are described by the ideal
gas law, which states that the product of the pressure and volume of a gas is proportional
to the product of the number of moles and the Kelvin temperature. Emil Clapeyron first
wrote this in 1834, and we’ll write it again here.
R is the gas constant. The value depends on the units used. When pressure is reported in
atmospheres (atm), volume in liters (L), and temperature in Kelvin (K), the gas constant
has a value of:
You will use the ideal gas law today to find the molar mass of butane. Although butane can
be described by the ideal gas law, it is important to remember that it is not ideal. Later on
you will use corrections to the ideal gas law to see how butane’s behavior deviates from
ideal gas behavior.
To find the molar mass of butane, you will collect butane gas by releasing it from a lighter
and collecting it over water in a graduated cylinder. You will be able to find the volume of
gas released in this way. In addition, you will make measurements of both temperature and
pressure. With the volume, temperature, and pressure of butane, you can use the ideal gas
law to find the moles of butane released from the lighter.
While measuring volume and temperature is accomplished easily, measuring the pressure
of the butane gas in the graduated cylinder is more complicated. The relevant gas law
which will help you to do this is Dalton’s Law of Partial Pressures. It states that a gas exerts
a certain pressure regardless of the presence of other gasses. This means that calculating
the pressure of each gas in a mixture independently and summing these individual
pressures determines the total pressure.
Or in our case, since we are collecting the gas over water,
Water vapor pressure (Pwater) will depend on temperature. Since water is a factor in so
many experiments, charts that provide pressure values at common temperatures are
readily available (though charts for other gasses can be obtained as well). Such a chart is
provided here.
Handout : Determining the Molar Mass of Butane
Teachers Guide
Purpose
To determine the molar mass of the butane in a disposable cigarette lighter.
Materials
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disposable lighter (Bic®, Cricket® or other brands) with striking machinery removed
100 mL graduated cylinder
balance
1L beaker or tub
water
thermometer
barometer (in the absence of a barometer you can also find the barometric pressure at
your location in weatherchannel.com)
small beaker (100 mL)
15 cm x 15 cm square plastic wrap
pliers
paper towels
hair dryer (optional)
Lab Hints
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Corrections for water vapor pressure and for the difference in internal and external
pressure by equalizing levels are required.
This experiment can be performed by students if balances with 0.1 mg sensitivity are
available. Collect about 100 mL of gas in a 100-mL graduated cylinder.
Use pliers to remove the guard, striker, spring and flint from a butane lighter. Make
weighings to the nearest 0.01 g.
In the calculation of Pbutane using Ptotal (the same as barometric pressure) – Pwater vapor and
in the calculation of molar mass of butane using wRT/PV take into consideration the
following unit conversions:
If the barometric pressure is given in inches of mercury (in Hg):
1 in Hg = 25.40 torr
1 atmosphere (atm) = 760 torr
1 atmosphere (atm) = 29.92 in Hg
Time
Teacher preparation: 30 minutes
Presentation: 30 minutes (Several drying steps)
Class Time: 45 minutes
Hazards
Butane is flammable.
Precautions
Remove flames or other ignition sources from the laboratory. Students should not be allowed to
play with the lighters.
Procedure
1. Immerse the lighter in water; remove. Shake the lighter to remove any excess of water
and dry it thoroughly. A hair dryer may also be used to dry the lighter.
2. Weigh the thoroughly dried lighter. Record the mass.
3. Fill a 1 L beaker with about 200 mL of tap water.
4. Fill the graduate to the top with water. Cover the cylinder with plastic wrap.
5. Invert and immerse under the water surface in a large beaker or other appropriate
vessel; remove plastic. Make sure no bubbles get in the cylinder; if you see bubbles
inside repeat the steps 4-5 until you don’t get bubbles inside the cylinder.
6. Position the cylinder near the wall of the beaker. Position the lip 1-2 cm below the water
surface.
7. Hold the modified lighter under the opening of the cylinder; press the release valve to
release gas. Collect about 100 mL of gas. Avoid losing bubbles coming from the
lighter; all butane bubbles coming from the lighter should be collected inside the
cylinder. If you miss some butane bubbles you will need to repeat the steps from
the beginning.
8. Remove the lighter; shake off excess water; pat dry with a paper towel.
9. Adjust the internal and external water levels using a 100 mL beaker until they are the
same; record the internal level.
10. Measure and record the temperature of the water.
11. Dry the lighter thoroughly. A hair dryer may also be used to dry the lighter. Weigh the
thoroughly dried lighter. Record the mass.
12. Record the barometric pressure.
Disposal
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Release the butane collected by displacement in a well ventilated area. Flush the water
down the sink. Place used towels and plastic wrap with ordinary solid trash.
Release any butane remaining in a lighter out of doors before disposing of that lighter
with ordinary solid trash.
Set Questions:
1. Find the pressure of a sample of oxygen that is collected over water at 27 °C if the total
pressure of the sample is 778.2 torr.
2. Find the molar mass of an unknown gas if a 0.16 g sample of the gas has a dry pressure
of 755 torr and a volume of 90.0 mL at a temperature of 27 °C .
Sample Data
Initial mass of lighter = 107.51 g
Final mass of lighter = 106.74 g
Mass of gas collected = 0.77 g
Volume of gas collected = 325 mL
Barometric Pressure = 744.7 torr
Water temperature = 24.0 °C
Vapor press of water = 22.4 torr at 24.0 °C
Pressure of butane = 722.3 torr
Molar mass of butane = wRT/PV = 61 g/mol
Closure Questions:
1. Identify at least 4 possible sources of experimental error
2. The accepted molar mass of butane is 58.0 g/mole. Determine percentage error for your
data.
3. Use your data, assume a molar mass of 58.0 g/mol, and derive the value of R, the gas
constant.
4. The molar mass of the gas is 58.0 g/mol. Determine the volume of one mole of butane at
STP based on your data. What is the percentage error between your value and 22.4 L
(the accepted value).
5. Butane contains only carbon and hydrogen. There are 4 carbon atoms per molecule.
Determine the empirical formula of butane based on your data.
References:
http://www.chem.uiuc.edu/chem103/molar_mass/introduction.htm
http://chemmovies.unl.edu/chemistry/dochem/DoChem077.html