BACKGROUND:
“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, C4H10. Why use butane to study gases 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 gases 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 gases can be obtained as well).
LAB: MOLAR MASS OF A GAS DETERMINED EXPERIMENTALLY
In this lab, we are going to determine the molar mass of butane, C4H10, experimentally. A simple calculation using a periodic table would give us the correct answer for the molar mass of butane, however, you are going to conduct an experiment in the lab to see how close you can come to the accepted value.
According to Avogadro’s law, the volume of 1 mol of any gas should be the same as the volume of 1 mol of any other gas, if temperature and pressure are the same. Molar volume is expressed in L/mol. By combining Avogadro’s law and the gas laws, we can calculate pressure, temperature, volume, or amount of gas. The Ideal Gas Law is summarized as:
PV = nRT
In this lab we will measure the mass of the butane released from a lighter (pressurized container), and we will measure the volume of the gas collected. You can use the method of water displacement at room conditions and then substitute the measurements of volume, temperature, and pressure into the ideal gas law equation in order to find n, the number of moles of butane. Once we know the mass, and the number of moles, we can calculate the molar mass (g/mol).
WARNING: Butane is a flammable gas, and at NO TIME during this lab should there be any use of an open flame or other heat source!
PURPOSE: Determine, experimentally, the molar mass of butane. Determine the % error for your lab.
HYPOTHESIS: If we can capture a sample of gas and determine the mass of the sample and the number of moles, then we can calculate the molar mass using the following equation:
Molar Mass = mass of sample in grams / number of moles
PROCEDURE:
1. Determine the initial mass for the butane lighter provided by your instructor ( +/- .01 gram ).
2. Set up the water basin (fill 2/3 full of water) for collecting a gas in an inverted graduated cylinder by water displacement. Use a 100 mL graduated cylinder.
3. Release butane from the pressurized container and collect a sample of gas with a volume of approximately 70.0 mL to 90.0 mL. Collect every bubble that leaves the lighter. If you miss a bubble, you will have to repeat the entire procedure beginning with re-weighing the clean and dry lighter. Don’t tilt the cylinder!
4. Adjust the cylinder up or down so that the water level inside the graduated cylinder is the same as the level outside the cylinder. THIS IS VERY IMPORTANT! If you miss this simple step, the pressures inside and outside the cylinder will not be equal, and you will have an inaccurate estimation of the pressure inside the cylinder where the gas is collected.
5. Record the volume of butane that was collected in the Observations and Data section. Put your hand over the mouth of the cylinder, keeping it inverted, bring it to the fumehood and release the gas by turning it right side up.
6. Measure and record the temperature of the water in the basin ( +/- 0.1 °C ). This will be used to determine the vapor pressure of water H2O (g) in the graduated cylinder. (a table of water vapor pressure at a given temperature is on the back)
7. Read the barometric pressure in the lab (inches of Hg ) and convert this to mm Hg (using the conversion factor 1 inch Hg = 25.4 mm Hg) and then to kPa.
8. Make sure the lighter is dry then, measure and record the mass of it ( +/- .01 gram).
9. Clean and return all materials to the starting position for the next class.
10. Report your measurements and calculations on the Observations and Analysis sheet below.
OBSERVATIONS AND ANALYSIS
Name ______Group members______
1. Beginning mass of butane container (minitial ) ______grams
2. Final mass of butane container (mfinal ) ______grams
mbutane = Mass of butane collected (line 1 – line 2) ______grams
3. Volume of the gas collected ______mL
Vbutane = Volume of butane, converted to liters ______L
4. Temperature of water basin in °C ______°C
Tbutane = Temperature of water in K ______K (assume the water T = butane T)
8. Barometric pressure in the room Proom ______mm Hg = Ptotal ______mm Hg
9. Vapor pressure of water at the temperature in line 6 (see Table 1 attached)
Pwater ______mm Hg
Note: The gas collected in the cylinder is a mixture of C4H10 (g) and H2O (g). Use Dalton’s Law of Partial Pressure to calculate the pressure of butane in the mixture.
Ptotal = P butane + P water vapor
10. Pressure of dry butane P butane ______mm Hg
11. Convert P butane from mm Hg into kPa using the unit factor method. SHOW your conversion below.
______= Pbutane ______kPa
To Calculate Moles
1. Calculate the number of moles of butane collected using PV = nRT
R = 8.31 kPa · L / mol · K
P= ______
V= ______
n =
Τ= ______
2. Calculate the experimental molar mass of butane by dividing the mass of butane released by n, the number of moles.
M butane experimental = ______
3. Calculate the accepted value for the molar mass of butane using a periodic table.
M C4H10 from periodic table =
4. Calculate the percentage error in your result.
% error = (theoretical value – lab value) x 100 %
theoretical value
QUESTIONS:
1. How many grams of butane did you collect in this experiment?
2. What volume would the butane have occupied at a higher atmospheric pressure?
3. What volume would the same number of grams of ethane (C2H6) have occupied at the same conditions of temperature and pressure?
Table 1: Vapor Pressure of Water at Various TemperaturesTemperature (°C) / Pressure (atm)
18 / .0204
19 / .0217
20 / .0230
21 / .0245
22 / .0261
23 / .0276
24 / .0295
25 / .0312
26 / .0332
27 / .0351
28 / .0372
29 / .0395