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Experiment 11 – Exploring the Properties of Gases
I. Purpose
Experimentally test Boyle’s law, Charles’s law and the ideal gas law constant. Graphical
analysis of data, to derive mathematical relationships that relate the different variables
used when studying gases, i.e. pressure, volume, amount of substance n and temperature.
II. Introduction
Intuition is a useful attribute. When you apply intuition to the behavior of gases, you can
arrive at the same relationships discovered 200 years ago by Robert Boyle, Jacques
Charles, Amedeo Avogadro, Joseph Gay-Lussac and Guillaume Amontons. Consider a
gas confined in a cylinder with a volume, V1, at temperature T1, and pressure, P1.
Boyle’s Law – Boyle studied what would happen to the volume of a gas if mass is added
to the piston (increasing the pressure on the gas) when temperature is held constant. He
found out that the volume will decrease or as the pressure increases, the volume
decreases. The simplest algebraic equation that can be written that predicts this behavior
is an inverse proportionality, V = k/P or PV = k where k is a constant. Other relationships
such a V = k /P2 are also consistent with a volume decrease as the pressure increases.
The question is which equation not only agrees with the direction of the changes but also
quantitatively agrees with the behavior. It turns out as Boyle discovered that the simple
inverse proportionality correctly predicts the behavior of gases and we call PV = k,
Boyle’s law.
Charles’s Law – He studied the effect on volume when the temperature is increased and
the pressure held constant. He observed that to maintain constant pressure the volume
has to increase and drive the piston up. The simplest algebraic expression of this
relationship that is consistent with this observation is the direct proportionality of
Charles’s law, V= k T or V/T = k. Inspection of this equation reveals that a temperature
of zero predicts a volume of zero. Now you know that the volume of a gas (e.g., the
interior of a balloon) will not go to zero at 0 ◦C (put a balloon in your freezer). A
temperature scale where the volume theoretically would be zero needs to be devised. The
Kelvin scale conforms to this criteria where 0 K = -273.15 ◦C. *Remember to use Kelvin
for all gas law calculations.
Gay-Lussac’s Law – Amontons and Gay-Lussac explored the relationship when the
piston is held in fixed position and the pressure is increased. They found that the
temperature has to increase to keep the piston in the same position. The direct
proportionality P = k T, or P/T = k , accurately predicts experimental behaviors and is
called either Amontons’s or Gay-Lussac’s law.
Avogadro’s Law – Consider the effect on the volume of adding more gas to the cylinder
while maintaining constant temperature and pressure. The relationship is the simplest
one with the volume proportional to the number of moles, V = k n (n equals the number
of moles).
Ideal Gas Law – Combining all these relationships yields V = nRT/P where R is the
proportionality constant and is appropriately called the gas constant. In the Part II of this
experiment, you will determine the value of R and compare it to the accepted literature
value of 0.08206 L atm/mol K. The equation is usually written in the form PV = nRT
and called the ideal gas law. All of the previously discussed gas laws can be easily
derived from this one equation. For example, if n and T are held constant, PV = constant
which is Boyle’s law. An in teresting relationship is that at STP (0◦ C and 1 atm), 1 mole
of any gas occupies 22.4 L. The unit conversion 22.4 L/mol has very limited use as it
applies only at STP whereas PV = nRT applies whenever gas behavior approximates that
of an ideal gas.
Absolute Zero – This is the theoretical value for the coldest temperature that matter can
attain. This temperature can be extrapolated using Charles’s Law, and plotting volume as
a function of temperature. To get the numerical value, extend (extrapolate) the graph line
to the axis where the volume of the gas would decrease to 0 mL.
III. Materials
25 mL 6 M HCl
1 g Zn metal (high surface area)
IV. Safety Precautions
When doing Part II, make sure your hydrogen is produced away from any flames.
V. Experimental Procedure
Part 1. Studying Boyle’s and Charle’s Laws Using Vernier
Review each section of this part before starting your work. You will need to decide the
best way to conduct the testing, so it is wise to make some plans before you begin. You
may wish to conduct a test run without collecting data, in order to observe how the
experiment will proceed. In each section of part A, you will investigate the relationship
between two of the four possible variables, the other two being constant.
A. Boyle’s Law: The Relationship between Pressure and Volume
1. Position the piston of a plastic 20 mL syringe so that there will be a measured volume
of air trapped in the barrel of the syringe.
2. Attach the syringe to the valve of the Gas Pressure sensor, as shown in figure 1. A
gentle half turn should connect the syringe to the sensor securely. Note: read the volume
at the front edge of the inside black ring on the piston of the syringe, as indicated by the
arrow in Figure 1.
Figure 11.1. Gas Syringe connected to Vernier Pressure Sensor
3. Connect the Gas Pressure sensor to Channel 1 of the Vernier computer interface.
Connect the interface to the computer using the proper cable.
4. Start the Logger Pro program on your computer. Open the file “30a Gases” from the
Advanced Chemistry with Vernier folder. This file allows you to collect pressure data
from the Gas Pressure sensor, using Events with Entry mode. For each pressure reading
you take with a “keep” button mode, enter a volume value. Collect at least six data
5. Measure the pressure of the air in the syringe at various volumes. Print a copy of the
graph and data.
B. Charles’s Law: The Relationship between Volume and Absolute Temperature
Using the apparatus shown in Figure 2, you will place an Erlenmeyer flask containing an
air sample in a water bath and you will vary the temperature of the water bath.
Figure 11.2. Set-up for Collecting Pressure at Changing Temperatures
Keep some of these factors in mind as you plan your procedure:
– If you are starting with a cold-water bath, set the piston at the 0 mL mark on the
syringe. This will allow the gas volume to be increased in warmer water baths.
– The temperature of the water bath cannot be increased by more than 30-40
degrees from your starting temperature.
– Even though you are not plotting pressure, it is important to monitor pressure in
the meter to ensure that it remains constant.
– It is important to know the total volume of air in the flask and syringe. The
volume of the flask, up to the bottom of rubber stopper, can be accurately
measured using a graduated cylinder. For the estimated volume of the tubing
(from the rubber stopper to the Gas Pressure sensor box), as well as in the valve
below the bottom of the syringe, use a value of ~4 mL.
1. Ensure that the Gas Pressure sensor is plugged into the Channel 1 and the
Temperature Probe is plugged into Channel 2 of the interface.
2. Assemble the apparatus as shown in figure 2. Be sure all fittings are airtight. Make
sure the rubber stopper and flask neck are dry, then twist and push hard on the rubber
stopper to ensure a tight fit. Be sure the water level is at least as high as the confined air
in the syringe.
3. Open the file “30c Gases” from the Advanced Chemistry with Vernier folder. This file
is set up to collect pressure and temperature data from the attached sensors, using Events
with Entry mode. This mode allows you to collect a data pair simultaneously from the
Gas Pressure sensor and Temperature Probe by clicking on the “keep” button and
entering a value for the volume. Even though the pressure reading will not be plotted on
the graph of volume versus temperature, it is important for pressure to be monitored so
that it can be kept constant.
4. Collect volume data at several different temperatures. Set up water baths in a largevolume container as you need them, ranging from ice water to hot water. Print a copy of
the graph and data.
Part 2. Determination of the Value of the Gas Constant, R
The determination of the gas constant R from the ideal gas law (R = PV/nT) requires the
values of P, V, n, and T. The values of P, V and T will be measured and the value of n
determined from a stoichiometric calculation. Hydrogen will be prepared by a single
replacement reaction:
Zn(s) + 2 HCl(aq) → ZnCl2(aq) + H2(g)
By utilizing a weighed amount of zinc, the number of moles of hydrogen that will be
produced can then be calculated.
1. Assemble the apparatus pictured in Figure 3. *Note:Be sure to keep all flames far
away from this apparatus once you start to generate hydrogen.
Figure 11.3. Set-up to Study Reaction (1)
2. Fill the 500 mL Florence flask (or it could be an Erlenmeyer flask) up to the neck with
water and add 25 mL of 6 M hydrochloric acid to the 250 mL Erlenmeyer flask.
3. While the Erlenmeyer flask has no stopper, apply compressed air (or use a rubber
bulb) to the tube inserted through the stopper. This will cause water to flow from the
Florence flask into the beaker. As soon as the flow begins, clamp the tube through which
the water is flowing while disconnecting the compressed air source.
4. Now that the delivery hose is filled with water, remove the stopper from the 500 mL
Florence flask, return the water you collected in the beaker to the Florence flask and
replace the stopper.
5. Weigh a little less than a gram of zinc metal to the nearest 0.001 g.
6. Holding the zinc sample with one hand and the stopper for the 250 mL Erlenmeyer in
the other, drop the sample into the HCl, quickly stopper the flask, and remove the clamp.
7. After the zinc has completely reacted, allow a few minutes for the gas in the flasks to
cool down to room temperature.
8. Equalize the pressure between the Florence flask and the beaker by raising the beaker
or flask (whichever liquid level is lower) so that the level of the water you have collected
in the beaker is the same as the level of water in the flask. Do not remove the stopper
from the flask until after you have measured the water.
9. Clamp the delivery hose when the water levels are equal.
10. Measure the amount of water you collected using a 500 mL graduated cylinder and
also measure the temperature of the water. Read the barometer to determine the current
atmospheric pressure.
Because you collected the hydrogen over water, part of the gas in the Florence flask was
water vapor. Consequently, to obtain the actual pressure of hydrogen that was produced,
you must subtract the vapor pressure of the water from the pressure at which the gas was
collected (atmospheric). The number of moles of hydrogen can be calculated from the
number of moles of zinc used and the stoichiometry of the reaction. The values of P, V,
and T can be determined from the measurements made during the experiment.
Table I. Vapor Pressure of Water as a Function of Temperature
Temperature (◦C )
Vapor Pressure (mm) Temperature (◦C )
Vapor Pressure (mm)
VI. Questions
1. For each section in Part I, write an equation using the two variables and a
proportionality constant, k. Also, graph the relationship between the two variables.
2. Assuming no elasticity forces in rubber, what volume should a 1.0 L ballon at 25 ◦C
and 1.00 atm be if immersed in liquid nitrogen (-196 ◦C, 1 atm)?
3. When this experiment is actually performed, the volume of the balloon approaches 0.0
L. How do you account for the discrepancy?
4. When a 0.134 g popcorn kernel pops, it loses about 13% of its mass. Assuming that
the volume of a typical kernel is 9.5 × 10-2 mL and that the mass loss is due to water lost,
use the ideal gas law to calculate the pressure of the water vapor in the kernel
immediately before it pops at 100 ◦C. Does the resulting calculated pressure account for
the popping of the popcorn?
VII. Discussion Questions
What other reactions could you sue to student gas phenomena? Based on today’s
findings what processes would you recommend to improve air quality in Los Angeles?
VIII. References
J. Randall, Advanced Chemistry with Vernier, Exploring properties of Gases,
(2007) 2nd Ed., Vernier Software & Technology, Beaverton, OR
Murov, S., Experiments in General Chemistry, Gas Law Studies, (2006) 5th Ed.,
Thomson Brooks/Cole

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