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The Gas Laws
PRE-LAB ASSIGNMENTS:
To be assigned by your lab instructor.
STUDENT LEARNING OUTCOMES:


Boyle’s Law, Charles’s Law, and Avogadro’s Law.
Using graphs to determine mathematical relationships between two variables
EXPERIMENTAL GOALS:
The purpose of this lab is to verify Boyle’s, Charles’s, and Avogadro’s Law and the gas law
constant, R.
INTRODUCTION:
Have you ever wondered why it gets harder to push on a partially full soda bottle as you
squeeze on it and reduce the volume inside? Similarly, why does your plastic bottle of soda get
rigid when you leave it in your car during the summer? The answer to both questions deals with
what we call gas laws. In the first case, we know that pressure is inversely proportional to
volume and as you reduce the volume of gas in the bottle, more force is required to continue
reducing the volume due to the increased pressure. In the second case, we know that gases
expand to exert more pressure when the temperature increases. We are going to discuss four gas
laws as an introduction to this lab.
Avogadro’s Law
An Italian scientist named Amedeo Avogadro (1776-1856) examined the relationship
between the amount of gas particles or molecules (now understood as the number of moles) of a
gas and its volume. This law is probably the one that makes the most sense to everybody. If you
have more of something, chances are that it will occupy more space and have a greater volume.
If temperature and pressure are kept constant, the volume that a gas occupies is directly
proportional to the number of moles of the gas. Avogadro’s Law may be expressed
mathematically as
Volume  Number of moles
(at constant pressure and temperatu re)
where the mathematical symbol  means ‘is proportional to.’ The volume and the number of
moles of gas are directly proportional to each other — that is, when one variable increases, the
other one also increases.
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Boyle’s Law
The British scientist Robert Boyle (1627-1691), discovered the relationship between the
volume that a gas occupies and the pressure that is exerted upon it. He used a glass, J-shaped
device that resembles an older style of barometer, that is appropriately called a J-tube. The tube
was sealed at one end and mercury was added to the open end.
mercury added
gas sample
As more mercury was added, thereby exerting more pressure on the gas, Boyle discovered that
the volume the gas occupied decreased. He realized that if you keep the temperature and number
of moles of gas constant, the volume occupied by the gas was inversely proportional to the
pressure exerted upon it. In other words, as the pressure that is exerted upon a gas increases (at
constant temperature and moles of gas) the volume that it occupies decreases and vice versa. In
honor of his discovery, we now call this relationship Boyle’s law. Boyle’s law may be expressed
mathematically as
Pressure 
1
Volume
(at constant t emperature and number of moles)
When one variable is proportional to the reciprocal of another variable, the two variables are said
to be inversely proportional to each other — that is, when one variable increases, the other
decreases.
Charles’ Law
In the late 1700s, a French mathematician and physicist named Jacques Alexandre César
Charles (1746-1823) was interested in the relationship between the volume that a gas occupied
and its temperature. He was also one of the first people to use a hydrogen-filled balloon to lift
himself off of the ground. Charles noticed that as you heat a gas, the volume that it occupies
increases, if the pressure and amount of gas is held constant. Similarly, as temperature is
decreased, the volume that the gas occupies decreases. He correctly concluded that the volume
that a gas occupies is directly proportional to the temperature, provided that the pressure and
number of moles of gas are kept constant. Charles’ law may be expressed mathematically as
Volume  Temperatur e
(at constant pressure and number of moles)
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The Ideal Gas Law
The Ideal Gas Law describes the general behavior of an ideal gas. Boyle’s, Charles’ and
Avogadro’s Laws are specific gas laws that may be combined to form this general gas law.
Similarly, these three named laws can easily be derived from the Ideal Gas Law. The Ideal Gas
Law can be written as
PV = nRT
where P is the pressure of the gas, V is the volume of the gas, n is the number of moles of the gas
and T is the temperature. The last term, R, is the ideal gas constant, whose value depends on the
units in which it is expressed. For example, in non-SI units, the value of R is 0.08206 L atm K-1
mol-1. It is important to emphasize that P, V, and T must have the same units as the gas constant
(think dimensional analysis). The Ideal Gas Law does make assumptions about gases behaving
in an “ideal” manner so there are some limitations to using it. What constitutes an ideal gas has
to do with intermolecular forces, etc., but we can approximate most gases as ideal at normal
pressures and temperatures. Hence, this Law is very useful for understanding the behavior of a
gas or gases in terms of the amount of gas (n) pressure (P), volume (V) and temperature (T). In
summary, this Law can be used to derive the other gas laws, and it allows chemists and engineers
to determine one of the variables provided that the other three are known.
PROCEDURE:
Safety Information
 In the Avogadro’s Law experiment, make sure that you carefully examine your ignition
tube for any defects, since any cracks can cause the tube to explode when heating. Get
another tube if you are concerned about the quality of your ignition tube and make your
instructor aware of the faulty tube.
 In the Charles’ Law experiment, be careful with the hot water; use some folded paper
towels to move the beaker of hot water onto the ring stand.
Students will work in groups of four. Data will be shared between the four students. Two
students will verify Boyle’s and Charles’s laws, while the other two verify Avogadro’s law.
Each pair should pay attention to the work that the other pair is doing, so that everyone
understands what’s going on in the experiments.
The group verifying Boyle’s and Charles’ laws should start heating 600 mL of water in a 800 mL
beaker to 90°C (use a Bunsen burner) and, while they are waiting, begin verifying Boyle’s law.
A. Boyle’s Law Verification.
Determining The Pressure Inside the Syringe
It is easy to confuse the properties of force and pressure. In fact, many of the ways we use the
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term pressure in the English language make it synonymous with force. For example, one of the
definitions of pressure in the online version of The American Heritage Dictionary is stated as:
“Pressure: The application of continuous force by one body on another that it is
touching; compression.”1
As with many terms we use in science, the meanings that scientists give to force and pressure are
more precisely defined than those applied in everyday usage. The scientific definition of
pressure in the same online dictionary is given as:
“Pressure: (Abbr. P) Physics. Force applied uniformly over a surface, measured
as force per unit of area.”2
If someone were to place their hand against your back and apply a force of 10 lbs3, it would feel
about the same as if you were lying on your stomach and a 10 lb weight were placed on your
back. However, if someone were to place their hand on your back and apply a pressure of 2 psi
(pounds per square inch, lbs/in2, or lbs in-2), it would feel the same if you were on your stomach
and 50 lbs of weight was placed on your back because a hand has an area of approximately 25
in2 (25 in2  2 lbs/in2 = 50 lbs).
In the Boyle’s Law experiment, you will apply different forces to the syringe plunger (in
lbs) by placing various weights on it. Inside the syringe, that force is distributed across the
internal face of the plunger. The resulting pressure increase inside the syringe due to the
addition of the weights can be determined using the formula
Padded 
Force applied to plunger
Area of plunger face
Assuming the weight of the top is negligible, the initial pressure inside the syringe (with no
weights added) is equal to atmospheric pressure. The barometric (atmospheric) pressure (Patm)
will be measured by your instructor and written on the blackboard in the lab. The pressure inside
the syringe (Psyringe) is the sum of these two components (Psyringe = Patm + Padded). Make sure both
pressures are converted to atm before adding them together.
Pressure-Volume Measurements
1. Slide the wooden block marked “TOP” onto the syringe end (see Figure 1a).
2. Draw at least 35 cc of air into the syringe, insert the syringe into the hole in the wooden
block marked “BOTTOM,” and firmly place the cap on the syringe through the hole on the
bottom of the wooden block (Figure 1b).
1
http://www.answers.com/pressure&r=67, accessed on 8/4/2008.
http://www.answers.com/pressure&r=67, accessed on 8/4/2008.
3
Pound is a measure of the force exerted by the gravitational pull of the Earth on a mass. Since the gravitational
force is essentially the same anywhere on the face of the earth, that force is proportional to mass and can be used as
a measure of mass on the Earth.
2
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3. Stand the apparatus upright (Figure 1c). While your partner is steadying the syringe, place a
2.5 lb weight on the top of the apparatus (Figure 1d). Record the resulting volume and the
amount of weight added. Repeat this with various increasing amounts of weight in 2.5 lb
increments until you have made five measurements of weights and volume. IMPORTANT:
DO NOT use more than 15 lbs on the apparatus.
4. Remove all of the weights. Reset the syringe by pulling it up to its original volume (you may
need to remove the cap end to do this).
5. Repeat the experiment so that you have two volume measurements for each weight.
6. Record the air temperature using a thermometer. We will assume that the air temperature is
the same as the temperature of the gas (air) in the syringe.
(a)
(b)
(c)
(d)
Figure 1. Boyle’s Law apparatus
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B. Charles’s Law Verification.
1. Place an 800 mL beaker containing 600 mL of water on a tripod stand, with a piece of wire
gauze under the beaker. Place some boiling chips or crushed pieces of porcelain into the
flask to provide a surface for bubbles to form on. Heat the water to 90° using a Bunsen
burner. Do not place the Charles’ Law apparatus in the water while it is being heated!
2. Draw at least 30 cc of air into the syringe and firmly place a cap on it.
3. Place the syringe into the large hole and a thermometer into the small hole of the long block
of wood (see Figure 2). Place the block on top of the beaker of hot water. Do not allow
water to enter the top of the syringe. The gas chamber in the syringe should remain
submerged in the water throughout the experiment, but the portion of the syringe above the
gas chamber does not need to be submerged. One partner can steady the syringe while the
other makes temperature and volume readings.
4. Allow two minutes for the syringe to equilibrate with the water.
5. Compress the syringe approximately 5 cc’s and allow it to return to position. Record the
volume in the syringe as well as the temperature.
6. Repeat step 5, recording this volume as your second measurement.
7. Allow the water to cool and continue to record volume measurements at approximately 20°C
increments. Use an ice bath to speed up the cooling process. After the temperature reaches
35°C, you can use a few ice chips to reduce the temperature further.
Figure 2. Charles’ Law apparatus
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C. Avogadro’s Law Verification.
For this portion of the lab, we will use the decomposition of KClO3 into KCl and O2 gas to
explore Avogadro’s Law according to the reaction below. The number of moles of KClO3 lost
upon heating is entirely due to the moles of oxygen that are produced, as measured by the
difference in weight before and after heating
2KClO3(s)  2KCl(s) + 3O2(g)
1. Obtain an ignition tube containing the KClO3/MnO2 mixture, stopper assembly, gas buret
with pinch clamp, and metal tub.
2. Attach a pinch clamp to the rubber tubing on the gas buret, fill the buret with tap water,
invert it, and place the lip of the buret under water. Hold the buret upright with a buret clamp
(see Figure 3a). Open the pinch clamp to allow water to run out of the buret until the water
level drops to a measurable level on the buret (below, but not exactly on the 50.00 mL mark)
(see Figure 3b). Read this volume level to the nearest 0.01 mL.
3. Place the L-shaped end of the stopper assembly into the gas buret (see Figure 3c).
4. Weigh the ignition tube containing the KClO3/MnO2 mixture.
5. Spread the KClO3/MnO2 mixture evenly along the length of the tube and clamp it in place.
(a)
(b)
Figure 3. Avogadro’s Law apparatus
(c)
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6. Attach the stopper to the ignition tube and attach this assembly to a ring stand. Gently heat
the tube until the mixture starts to bubble (1-3 minutes, depending on your Bunsen burner),
and remove the heat. (If the level of the gas drops below the zero mark on the buret, you
have produced too much gas and cannot use the data from that experiment.)
7. Allow the tube to cool, and record the final volume reading to the nearest 0.01 mL. Use the
initial and final volume readings to determine the volume of the gas produced. Weigh the
test tube to determine the mass of oxygen produced.
8. Reassemble the apparatus and make at least five successful measurements.
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LAB REPORT
The Gas Laws
Name ________________________________
Date _________
Partners ________________________________
Section _________
Report Grade ______
________________________________
________________________________
A. Boyle’s Law Verification.
Air temperature:
__________
Atmospheric pressure: __________
Weight
Volume 1
Volume 2
Avg. Volume
Pressure (atm)
2.5 lbs
5.0 lbs
7.5 lbs
10.0 lbs
12.5 lbs
15.0 lbs
1. Average the two volume measurements for each weight.
2. Calculate the pressure in atm exerted by each weight on the gas in the syringe. Pressure is
equal to force/area (P=F/A). For a circle, A= r2. The average diameter of the syringe you
will use in lab is 0.937 inches. Convert this number to atmospheres using the relationship
1 atm = 14.7 psi. You will need to add the atmospheric pressure for that day (in units of atm)
to this value in order to determine the total pressure exerted.
3. Plot average volume (y-axis) versus pressure (x-axis) using Excel. Include the trendline and
R2 value.
Show your work:
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B. Charles’ Law Verification.
Temperature
(°C)
Temperature
(K)
Volume 1
Volume 2
Avg. Volume
1. Average the two volume measurements for each temperature.
2. Plot the average volume (y-axis) versus temperature (x-axis), converted to Kelvin, using
Excel. Include the trendline and R2 value.
Show your work:
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C. Avogadro’s Law Verification.
Mass of test
tube before
heating
Mass of test
tube after
heating
Mass of O2
Moles of O2
Volume
1. Since oxygen is the only gas evolved, you can determine the moles of gas produced using the
molar mass of oxygen.
2. Plot the volume of gas produced (y-axis) versus moles of gas evolved (x-axis) using Excel.
Include the trendline and R2 value.
Show your work:
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D. Post Lab Questions
1. Based on your data, what is the relationship between pressure and volume in regard to gases?
Does this agree with Boyle’s Law? Explain your answer.
2. Based on your data, what is the relationship between temperature and volume in regard to
gases? Does this agree with Charles’s Law? Explain your answer.
3. Based on your data, what is the relationship between the number of moles of a gas and the
volume it occupies? Does this agree with Avogadro’s Law? Explain your answer.
4. You can use the data collected to verify the value of R, the ideal gas constant (R= 0.0821 L
atm K-1 mol-1). From the ideal gas law (PV = nRT), R is a ratio between the product of VP
and nT (R= VP/nT). Assume that the density of air is 1.20 g/L and that the molecular
weight of air is 28.8 g/mol, and use these values to convert the volume of air you began with,
35 mL, to moles of air. Use the data you collected in Boyle’s law to verify the gas constant
for each measurement you made. Make sure that you use proper units. Comment on the
accuracy of your calculations.