133
Discovering The Gas Laws
PRE-LAB ASSIGNMENTS:
To be assigned by your lab instructor.
STUDENT LEARNING OUTCOMES:





Students should have a qualitative empirical sense of what happens to
o Volume of a gas when
 amount (# of particles) of gas changes
 temperature changes
 pressure changes
o Pressure of a gas when
 amount (# of particles) of gas changes
 temperature changes
 volume of gas changes
Students will understand the principle of absolute temperature in terms of the volumetemperature behavior of gases and be able to work with the absolute temperature scale.
Students will understand that volume varies proportionally with temperature and that
proportionality constant varies with the amount of starting gas.
Students will be able solve Charles’ Law problems.
Students will be able to uses graphs to determine mathematical relationships between two
variables
ORIENTATION:
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 this lab you will explore the behavior of gas in response to these kinds
of changes and use your data to define four gas laws and learn how to use those laws to predict
changes that occur in a gas sample when one or more to the above properties are changed.
134
Discovering The Gas Laws
Part 1 – The temperature-volume relationship
(Rediscovering Charles’ Law)
EXPERIMENTAL GOALS:
In this experiment, you will explore quantitatively how the volume of a sample of gas is affected
by temperature and use your observations to develop a mathematical approach for determining
how one can predict how the volume of a gas will change with a change in temperature.
Process skill goal: students will be able to analyze data to detect proportional and linear
relationships. They will be able to modify any linear relationship to obtain a proportional
relationship and use proportional relationships to predict how one parameter changes in response
to change in another.
PROCEDURE:
Safety Information
 Be careful with the hot water; use some folded paper towels to move the beaker of hot
water onto/off of the ring stand.
 When a liquid is to be heated to or near to the boiling point, boiling chips or crushed
pieces of porcelain should be placed in the liquid before heating is started. This will
prevent the liquid from becoming superheated (heated to a temperature above the boiling
point): superheated water can boil suddenly and vigorously and presents a serious burn
hazard.
Students will work in pairs. Each pair should team up with one other pair. One pair in the team
should use a starting syringe volume of 25-30 mL and the other team should start with 40-43 mL.
Each pair in a group should use a different starting volume.
As soon as possible at the start of the lab period, each pair of students should start heating 600
mL of water to 90°C in a 800 mL beaker using a Bunsen burner (see step 1 below).
The Charles’ Law apparatus illustrated below will be used in this experiment.
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Figure 2. Charles’ Law apparatus
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 prevent superheating of the water. Heat the water to 90° using a Meker burner. Do
not place the Charles’ Law apparatus in the water while it is being heated!
2. Draw the starting amount of air that you selected above (25-30 or 40-43 mL) into the syringe
and firmly place a cap on it. Make sure that each pair of students in your group is using a
different starting amount. Compress the syringe approximately 5 cc’s and allow 20 seconds
for it settle into it final position. Record the exact amount used to as many significant figures
as possible.
3. Place the plunger handle of the syringe through the large hole (bottom to top) in the syringethermometer holder block (see Figure 2). Place the thermometer into the small hole of the
block (see Figure 2). A small slice of rubber tubing can be placed on the thermometer to
hold it off the bottom of the beaker.) Place the block assembly 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 20 seconds for it settle into it final
position. This compress and release procedure provides the most precise results.
Record the volume in the syringe as well as the temperature. Repeat the compress, release,
wait and read procedure one more time to obtain a second measurement at this temperature
recording both the volume and the temperature.
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6. Make several additional readings at progressively lower temperatures by following these
steps.
 Add 6-10 pieces of ice to the beaker. Stir gently and monitor the temperature until
the ice melts and the temperature stabilizes. The temperature should drop 15-20 oC.
You can adjust the number of pieces of ice needed to get in this range.
Be careful to not get water into the syringe barrel when adding ice or stirring.
If the water level in the beaker gets too high, use the specially marked wash bottle
to draw water out of the beaker.
 Allow two minutes for the temperature of the syringe and gas to equilibrate with the
water.
 Read the temperature and volume two times using the compress, release, wait and
record procedure described above.
The final temperature should be between 10-20 oC. Using crushed ice for the last reading
will speed up the melting of the ice at this lower temperature.
137
LAB REPORT
The Gas Laws
Name ________________________________
Date _________
Partners ________________________________
Section _________
Report Grade ______
________________________________
B. Effect of Gas Temperature on Volume Verification.
Starting Room Temperature Volume _____________ mL
Temperature 1
(°C)
Volume 1
(mL)
Temperature 2
(°C)
Data from Other Pair in Team
Partners in Other Pair
___________________________
___________________________
___________________________
Starting Room Temp. Volume
Volume 2
(mL)
Average
Temperature
Average
Temperature
Avg. Volume
Avg. Volume
138
Analyzing the data
Average the two volume and the two temperature measurements for each temperature interval.
Plot the average volume (y-axis) versus average temperature in oC (x-axis) data of both pairs of
students in your team in one Excel spreadsheet. Include the trendline slope and intercept and R 2
values for each line. You should have two lines each representing a different initial volume
(sample size) of gas.
Thinking about the data
Examine the trends in the resulting lines from both pairs of students.
1. Does this represent a proportional relationship between temperature and pressure? Why or
why not?
2. Do the lines tend to run parallel or do they tend to converge in the same direction. If they are
converging, will they converge at higher or lower temperatures than those studied?
3. If the gas continued to behave in this fashion, eventually the volume would go to zero and
below! Of course, this does not happen.
a. Physically something else must happen first: what happens to a gas before its volume goes to
zero? (Hint: think about taking a sample of water vapor at 150 0C and cooling it. What happens
at 100 oC?)
139
b. To explore this idea of the volume eventually decreasing to zero further, we will try to
estimate the temperature at which your gas samples would reach a volume of zero. We are
assuming it continues to decrease linearly until that point. In reality, it will convert to liquid at
some point and the volume-temperature behavior will change drastically.
For each of the sets of data, right click on the trend line to get the “Format Trendline“ box and
click on the Options tab. Use the “Forecast” feature to extend the lines (units used are x-axis
units) until the volume goes below zero and all of the lines have intersected.
For each of your groups data sets (lines) determine the temperature at which the line crosses
zero. This can be estimated from the graph or from the best-fit line equation for the data set. To
use the best fit equation, calculate the temperature (x) when the volume (y) is zero.
Are these values reasonably close together or very different.?
c. How close are the points of intersection of the lines to each other and how close are they to the
place on the graph where the “volumes become zero”?
Does it appear that it is possible the points of intersection of all these lines might be
experimentally be the same (are they really close to each other or very far apart)?
Note: remember that the data region in this portion of the graph is being extrapolated quite a distance from the
experimental data. Extrapolation tends to exaggerate experimental error 1and so even modest similarities may
indicate the possibility of agreement.
d. Average temperatures obtained in 3b. What physical constant related to temperature is this
value near?
PRINT your final graph to turn in. Use appropriate axis labels, chart title, include trendline &
eqn, Each person must do their own graph.
Homework Assignment
Do the assignment titled “Obtaining Ideal Gas Data” which can be found in Blackboard at
. . . > ASSIGNMENTS > PRELAB WORK FOR LAB WEEK 10: NOV 3-9
Include the results with your lab report.
1
For this reason, scientists analyzing data avoid using extrapolation as much as possible, if it does become
necessary, as in the case of your V vs. T data, extra precautions should be taken to get very accurate and precise
data. In addition to error exaggeration, we cannot always be sure that the trends observed in one region of the graph
will hold true in regions out side of the data collection range. This is true for the data we are examining. The
behavior of the sample changes drastically as it approaches and drops below the boiling point of the sample.
140
Discovering The Gas Laws
Part 2 – The pressure-volume relationship
(Rediscovering Boyle’s Law)
EXPERIMENTAL GOALS:
In this experiment, you will explore quantitatively how the volume of a sample of gas is affected
by pressure and use your observations to develop a mathematical approach for determining how
one can predict how the volume of a gas will change with a change in pressure.
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
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.”2
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.”3
If someone were to place their hand against your back and apply a force of 10 lbs 4, 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
2
http://www.answers.com/pressure&r=67, accessed on 8/4/2008.
http://www.answers.com/pressure&r=67, accessed on 8/4/2008.
4
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.
3
141
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
Force
applied
to
plunger
P

added
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.
Procedure: Pressure-Volume Measurements
Students will work in pairs. Each pair should team up with one other pair. One pair in the team
should use a starting syringe volume of ca. 35 mL and the other team should start 45-50 mL.
Each pair in a group should use a different starting volume.
1. Slide the wooden block marked “TOP” onto the syringe end (see Figure 1a).
2. With the syringe cap off adjust the plunger so that the syringe holds the desired starting
volume (35 mL or 45 mL) of air. Try to get the plunger position to within 0.2 mL of the
desired volume. 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).
3. Stand the apparatus upright (Figure 1c). One partner will steady the apparatus and take
volume readings. The other partner will add weights and record data. Before adding weight,
depress the plunger ca. 5 mL, release it and allow one minute for the plunger position to
stabilize. Read the volume to the nearest 0.1 mL. While one 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. Wait one minute for the plunger position to stabilize. Record the
position of the plunger to the nearest 0.1 mL.
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. Place thermometer in a split rubber stopper
and clamp to a ring stand with a test tube clamp and wait for ten minutes to ensure an
accurate reading. We will assume that the air temperature is the same as the temperature of
the gas (air) in the syringe.
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(a)
(b)
(c)
(d)
Figure 1. Boyle’s Law apparatus
143
LAB REPORT
The Gas Laws
Name ________________________________
Date _________
Partners ________________________________
Section _________
Report Grade ______
________________________________
A. Boyle’s Law Verification.
Air temperature:
Weight
__________
Volume 1
Atmospheric pressure:
Volume 2
__________
Avg. Volume
Pressure (atm)
0.0 lbs (initial)
2.5 lbs
5.0 lbs
7.5 lbs
10.0 lbs
12.5 lbs
15.0 lbs
0.0 lbs (final)
Data from Other Pair in Team
Partners in Other Pair
___________________________
___________________________
___________________________
Starting Room Temp. Volume
Average
Pressure
Avg. Volume
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Study Questions
PRINT your final graphs to turn in. Use appropriate axis labels, chart title, include trendline &
eqn, Each person must do their own graph.
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). Include the trendline and R2 value.
4. Plot 1/(average volume) vs. pressure using Excel. Include the trendline and R2 value.
5. Which relationship is more linear. From the plot of your data do you think the more linear
relationship is proportional?
Show your work:
145
Discovering The Gas Laws
Part 3 – The mole-volume relationship
(Rediscovering Avogadro’s Law)
EXPERIMENTAL GOALS:
The goal of this experiment is to measure the volume of water displaced during the
decomposition of potassium chlorate to find the volume of oxygen gas generated. The data
collected for several different samples of O2 at similar temperatures and pressures will be used to
understand the relationship between moles of gas and gas volume. The data collected will be
sufficient to determine the ideal gas law constant in a future lab.
INTRODUCTION:
In this experiment oxygen gas is formed by the decomposition of potassium chlorate:


2KClO3(s) 
2KCl(s) + 3O2(g)
The potassium chlorate, with the catalyst manganese(IV) oxide, MnO2, is placed in a test tube,
and the weight is determined accurately. The test tube is then heated, and oxygen gas is evolved
as the potassium chlorate decomposes. (Potassium chlorate decomposes rather slowly when
heated. The MnO2 acts as a catalyst for the reaction, speeding the reaction up without itself
being consumed.) The entire weight loss of the test tube can be assumed to be due to the mass of
the oxygen gas which is formed which can be used to determine the moles of gas produced. This
oxygen gas is passed into a bottle containing water and displaces a volume of water equal to the
volume of oxygen gas produced; the water is forced out of the bottle through a hose and is
collected in a beaker, where its volume can be measured.
The temperature and pressure of the gas will also be measured and used for analysis in
next week's lab. The apparatus is balanced against atmospheric pressure so that the pressure of
the gas which is collected is the same as that day’s atmospheric pressure. The gas which is
collected is actually a mixture of oxygen and water vapor, since gases bubbled through water
pick up small amounts of moisture. The total pressure of the gas is the sum of the pressure of the
dry oxygen, PO 2 , and the vapor pressure of water, Pw (from Dalton’s Law of Partial Pressures),
P
P
P
total
O
w
2
The pressure of the dry oxygen is determined by subtracting the vapor pressure of water from the
total pressure.
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PROCEDURE:
1. Obtain the apparatus shown in Figure 1 (tubing assembly, 2.5 L bottle, 800 mL beaker, and
pinch clamp). Obtain from the stockroom a heavy-walled Pyrex test tube (ignition tube)
containing the KClO3/MnO2 mixture.
2. Fill the large (2.5 L) bottle with tap water. Insert the stopper/delivery tubes into the bottle.
Do not attach the test tube yet.
3. Fill the 800 mL beaker about one-third full of tap water. Open the pinch clamp and use a
squirt bottle labeled “MW of O2 Lab” to blow through the tubing at point A until water runs
into the beaker. Close the pinch clamp, keeping it toward the bottom of the rubber tubing.
4. Weigh the Pyrex test tube and KClO3/MnO2 mixture (without stopper). Spread the mixture
out along the lower portion of the test tube for a couple of inches.
Figure 1. Experimental setup.
5.
To make the pressure in the 2.5 L bottle equal to the room air pressure, attach the test
tube to the assembly. Keeping the end of the glass tube extension under water in the beaker,
open the pinch clamp and raise the beaker until the water level in the beaker is the same as the
water level in the 2.5 L bottle. Hold the beaker in this position for 10-15 seconds, then close the
pinch clamp and empty the beaker.
6. To test for leaks, open the pinch clamp. Only a very small amount of water, if any, should
dribble out of the glass extension tube; then it should stop. If water continues to flow, the
assembly has an air leak. Consult your instructor before proceeding.
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7. If your apparatus has no air leaks (water stops running), you are ready to begin heating the
test tube. Leave the pinch clamp off and do not discard the small amount of water collected
in the beaker. Have your instructor check your apparatus before proceeding.
8. Heat the upper part of the KClO3/MnO2 mixture, slowly at first, then increase the heating
until a steady stream of water is flowing into the beaker. If necessary, move the heating zone
toward the lower end of the mixture. Continue heating. On your first run quit heating your
sample when 200-300 mL of water has been collected in the beaker. On your first run quit
heating your sample when 400-500 mL of water has been collected in the beaker. Turn off
the burner when you quit heating the sample. The goal is to produce two substantially
different size oxygen samples. Let the entire assembly stand undisturbed with the glass tube
extension under water and pinch clamp open for ten minutes. Do NOT attach the pinch
clamp yet!
NOTE: Do not let the beaker overflow! If it starts to do so, catch any overflow in another
beaker. Also do not let the water in the 2.5 L bottle drop below the delivery tube. If it does
so, stop the experiment and start again.
9. Keeping the end of the glass tube extension under water, raise the beaker (or bottle) until the
levels of water are the same in both. Close the pinch clamp. Do NOT disconnect the
ignition tube until the other end of the apparatus is closed by the pinch clamp.
10. Measure the volume of water collected in a large graduated cylinder.
11. Remove the test tube (without the stopper) from the assembly and weigh it.
12. Record the barometric pressure of the room, which is written on the blackboard. Measure the
temperature of the water in the beaker.
13. Repeat the experiment, beginning at step 5 above. Since you were given enough
KClO3/MnO2 for two runs, use the final weight of the test tube and contents from the first run
as the initial weight for the second run.
14. After completing both runs, put your used sample test tubes in the beaker designated for that
purpose.
15. Dismantle the equipment and return it to its proper location.
148
LAB REPORT
The Mole-Volume Relationship
Name ________________________________
Date _________
Partner ________________________________
Section _________
Report Grade ______
First
Determination
Second
Determination
1.
Weight of test tube + KClO3/MnO2
______________ ______________
2.
Weight of test tube after heating
______________ ______________
3.
Weight of oxygen evolved
______________ ______________
4.
Volume of water (= volume of oxygen evolved) (V)
______________ ______________
5.
Temperature of water (= temp. of oxygen) (T)
______________ ______________
6.
Barometric pressure (Ptotal)
______________ ______________
7.
Vapor pressure of H2O at T (from Appendix IIIB)
______________ ______________
8.
Partial pressure of oxygen (show calculations)
______________ ______________
9.
Partial pressure of oxygen in atm (PO2) (show
calculations)
______________ ______________
10. Moles of oxygen calculated from mass O2
______________ ______________
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Obtain 4 additional data points (moles of O2 and volume) from two other pairs of students in
your class and combine it with your data in the table below.
Names
Moles
Volume
Your Data Points
from previous page
Data from
=>
Write names of team members =>
Data from
=>
Write names of team members =>
Plot this data on a graph with volume on the Y-axis and mole on the X-axis.
If any points seem like they are in error, you may want to exclude them from the graph. PRINT
your final graph to turn in. Use appropriate axis labels, chart title, include trendline & eqn,
Each person must do their own graph.
1. Does this graph indicate a proportional relationship between moles and volume?
Preparation for Next Weeks Lab
All data sheets from this Experiment (Parts 1-3) along with well labeled, high quality copies of
the graphs requested should be turned in at the start of the next lab. You should also make a
copy of your lab report to have on hand in the lab next week. In addition, your spreadsheet files
should be available either on some type of portable media such as a jump-drive or your ASU
network P-Drive.
Lab next week will meet in a computer lab. Lab locations will be announced when that
information becomes available.