Physical Chemistry 1 Laboratory Report
Virtual Experiment 01
Demonstration of Gas Laws
(Boyle’s Law, Charles’ Law, Avogadro’s Principle & Daltons Law of Partial Pressure)
Submitted by:
Carreon, Michael Angelo C.
Gocotano, Carl Lyster
Gonzales, Kristian Karl O.
Submitted to:
Fr. Antonio Basilio, SJ
December 22, 2021
Abstract:
The gas laws are a set of physical laws developed from experimental observations used to model
the behavior of gases. These fundamental laws describe the relationship of temperature, volume,
pressure, and the number of moles of a gas. In this experiment, two of the gas laws are presented and
discussed – namely, Charles’ law, Boyle’s law, Avogadro’s principle, and Dalton’s law of partial
pressure. Charles’ law is a statement relating temperature and volume at constant pressure. While
Boyle’s law is a statement relating pressure and volume at constant temperature. On the other hand ,
the Avogadro’s principle states the direct relationship between volume and the amount of gaseous
particles under constant temperature and pressure. And lastly, Dalton’s law of partial pressure states
that the total pressure of a gaseous mixture is equal to the sum of the partial pressures of the components.
Using the PhET Colorado’s simulation of gas properties, the stated laws are demonstrated. Results and
observations are analyzed and interpreted according to the principles of Kinetic Molecular Theory
(KMT).
Introduction:
Gases have three key characteristics- they are compressible, they expand to fill their containers,
and they occupy far more space than the liquids or solids from which they form. Brought by this unique
set of characteristics attributable to its relatively weak intermolecular forces, gases behave differently
compared to other states of matter. Gas behavior can be modelled using the experimental gas laws which
are established to assist in determining volumes, pressures, temperatures, and quantities when coming
to matters related to gas. Charles’ law describes how gases expand when subjected to heat when pressure
is held constant. Boyle’s law describes the inverse relationship between pressure and volume at constant
temperature. However, these models are only applicable when gases are said to be “ideal” as described
by the Kinetic Molecular Theory. The idea of ideal gases is just a theoretical conception wherein the
characteristics of a gas adhere to the following governing assumptions: (1) The gas particles have
negligible volume; (2) The gas particles are equally sized and do not have intermolecular forces
(attraction or repulsion) with other gas particles; (3) The gas particles move randomly in agreement
with Newton’s Laws of Motion; and (4) The gas particles have perfect elastic collisions with no energy
loss. In this experiment, the Charles’ law and Boyle’s law are demonstrated assuming that gases are at
their ideal state. Results and observations will be used to support the theoretical assumptions of the gas
laws and the principles of the Kinetic Molecular Theory.
Procedure
In this virtual experiment, the demonstration of the Gas Laws was performed using PhET
Interactive Simulations virtual laboratory.
For the virtual demonstration of Charles’ Law, the “Ideal” simulation in the PhET virtual
laboratory was used. (1) The experiment started by inserting 150 heavy particles controlled in the
“Particles” section of the virtual lab. After the addition of the particles, in the “Hold Constant” section,
the “Pressure ↕V” was selected, and the “Width” was checked to show the width of the container. For
this experiment the initial width was set to 5 nm, while the pressure and number of particles was set
constant at 29.2 atm and 150 particles, respectively.
(2) The temperature can be set by rising (for higher temperature) and lowering (for lower
temperature) the lever in the bucket located below the experimental set-up. The initial temperature was
250 K, and the corresponding initial width was recorded to calculate the volume. Procedure 2 was
repeated for 275 K, 300 K, 325 K, and 350 K, and the resulting width of each temperature was also
recorded. Each trial was then triplicated.
*For actual laboratory demonstration of Charles’ Law, the equipment needed are a hot plate,
1000 mL beaker, and a balloon. (1) Fill the beaker with 500 mL of water, and heat it using the hot plate.
(2) Fill the balloon with air up to the size to fit inside the beaker. (3) As the water starts to boil, place
the balloon in the beaker touching the water. (4) Observe the balloon going up the beaker as the
temperature rises. (This process is behind the rising of hot air balloons. As the hot air balloons heat
up, they increase in size, and the hot air wants to expand by moving up, taking the balloon with it) *
For the virtual demonstration of Boyle’s Law, the “Ideal” simulation in the PhET virtual
laboratory was used. (1) The experiment started by inserting 150 heavy particles controlled in the
“Particles” section of the virtual lab. After the addition of the particles, in the “Hold Constant” section,
the “Temperature” was selected, and the “Width” was checked to show the width of the container. The
movement of the particles can be paused and played by pressing the “Play” button beside the bucket
below the set-up. For this experiment, the temperature was set constant at 300 K.
(2) The width of the container can be adjusted by dragging the handle located at its left side.
The initial width of the container was set to 5 nm, the movement of the particles was paused, and the
corresponding pressure was recorded; the procedure was done twice more, and the collected pressure
results were averaged. Procedure 2 was repeated for 7 nm, 9 nm, 11 nm, and 13 nm, and the resulting
pressure for each width was also recorded.
*For an actual laboratory demonstration of Boyle’s Law, the equipment needed is two small
balloons, a large plastic syringe, and a scissor. (1) Fill the balloon with air up to the size where it can
fit inside the syringe. (2) Place the balloon inside the syringe. (3) For the first test, place the syringe's
plunger at the highest point of the syringe; cover the bottom opening with a finger. (4) Press the plunger
down; observe the balloon shrink inside the syringe. (5) Remove the finger at the bottom of the syringe
to release the pressure. (6) For the second part, place the plunger at the lowest point it can be without
touching the balloon inside; cover the bottom opening with a finger. (7) Pull the plunger upwards;
observe the balloon inside increase in size. (8) Remove the finger at the bottom of the syringe to release
the pressure. (Part one and two of this experiment exhibits the inversely proportional relationship
between the Volume and Pressure of an ideal gas.)*
For the virtual demonstration of Avogadro’s Principle, the “ideal” simulation in PhET virtual
laboratory was used. (1) the experiment started by inserting 80 heavy particles controlled in the
“Particles” section of the virtual lab. After the addition of the particles, in the “Hold Constant” section,
the “Pressure ↕V” was selected, and the “Width” was checked to show the width of the container. For
this experiment the initial width was set to 5 nm, while the pressure and the temperature was set constant
at 18.7 atm and 300 K, respectively.
(2) The initial volume was taken note of, and more heavy particles were added until there were
90 heavy particles. The final volume was taken note of, and the procedure was repeated twice more.
The average final volume of the three trials was considered as the final volume for its corresponding
number of heavy particles. The process was repeated for 100, 110, and 120 number of particles.
*For an actual laboratory demonstration of Avogadro’s Principle, simply blow up a balloon.
To do this, (1) breathe into the balloon to add molecules of gas within causing the balloon to increase
in volume. (2) Notice that when air is released out of the balloon, that is when the number of gaseous
particles in the system decreases at constant temperature and pressure, and consequently the volume
is also decreased. Another principle acquired from this law is that if two gas samples have the same
volume, temperature and pressure, then they will all have the same number of molecules within them
For the virtual demonstration of Dalton’s Partial Pressure, the “ideal” simulation in PhET
virtual laboratory was used. (1) the experiment started by inserting 100 heavy particles controlled in the
“Particles” section of the virtual lab. After the addition of the particles, in the “Hold Constant” section,
the “Temperature” was selected, and the “Width” was checked to show the width of the container. For
this experiment the initial number of heavy particles was set to 100, while the width and the temperature
was set constant at 5 nm and 300 K, respectively.
(2) The initial number of heavy particles was taken note of, the movement of the particles was
paused, and the corresponding pressure was recorded; the procedure was done twice more, and the
collected pressure results were averaged. Procedure 2 was then repeated for 110, 120, 130, and 140
varying numbers of heavy particles.
(3) The heavy particles were replaced with 100 lighter particles and the same procedure 2 were
executed.
(4) The light particles were removed and replaced with a mixture of 100 heavy particles and
100 light balls. The entire procedure 2 was then repeated.
*An actual laboratory demonstration of Dalton’s Partial Pressure can be done by combining
methane, oxygen, and hydrogen gas in a gas mixture. To compute the total pressure, the pressure of
these gases must first be measured as a mixture. Their pressure can then be measured independently,
as though each gas exerted force on its own in a container with a defined volume. If the sum of each
lone gas's partial pressures equals or nearly equals the total pressure of the mixture, they are said to
obey Dalton's Law.
Results:
Table 1 shows the results of varying temperature on volume at constant pressure and number of
particles:
Volume (nm3)
Ave. Volume (nm3)
Trial
Temperature (K)
1
250
125
125
125
125.00
2
275
166.375
166.375
166.375
166.38
3
300
216
216
216
216.00
4
325
274.625
274.625
274.625
274.63
5
350
343
343
343
343.00
Table 1: Results of the Charles’ Law demonstration
Graph 1 shows the direct relationship between temperature and volume obeying Charles’ law of perfect
gases:
Temperature vs Volume
400
Volume (nm3)
350
300
250
200
150
100
225
250
275
300
325
Temperature K
Graph 1: Temperature versus volume
350
375
Table 2 shows the results of varying volume on pressure at constant temperature and number of
particles:
Trial
Volume (nm3)
1
125
35.2
35.0
34.9
35.03
2
343
25.0
24.7
25.2
24.97
3
729
19.1
19.9
19.6
19.53
4
1331
15.8
16.0
16.3
16.03
5
2197
13.8
13.9
13.2
13.63
Pressure (atm)
Final Pressure (atm)
Table 2: Results of the Boyle’s Law demonstration
Graph 2 shows the inverse relationship between pressure and volume obeying Boyle’s law of perfect
gases:
Volume vs Pressure
40
Pressure (atm)
35
30
25
20
15
10
0
250
500
750
1000
Volume
1250
1500
(nm3)
Graph 2: Volume versus pressure
1750
2000
2250
Table 3 shows the results of varying number of heavy particles on volume at constant temperature and
pressure:
Volume (nm3)
Ave. Volume (nm3)
Trial
No. of Particles
1
80
125
125
125
125.00
2
90
175.616
175.616
175.616
175.62
3
100
250.047
250.047
250.047
250.05
4
110
328.509
328.509
328.509
328.51
5
120
421.875
421.875
421.875
421.88
Table 3: Results of the Avogadro’s Law demonstration
Graph 3 shows the direct relationship between number of heavy particles and volume obeying
Avogadro’s law of perfect gases:
No. of Particles vs Volume
450
Volume (nm3)
400
350
300
250
200
150
100
75
80
85
90
95
100
105
110
115
No. of Particles
Graph 3: Number of particles versus volume
120
125
Table 4.1 shows the results of varying number of heavy particles on pressure at constant temperature
and volume:
Trial
No. of Heavy Particles
Pressure (atm)
Ave Pressure (atm)
1
100
23.2
23.6
23.5
23.43
2
110
25.9
26.0
25.6
25.83
3
120
27.7
28.0
28.1
27.93
4
130
30.7
30.1
30.3
30.37
5
140
32.7
32.9
33.1
23.43
Table 4.1: Pressure at different number of heavy particles
Table 4.2 shows the results of varying number of light particles on pressure at constant temperature
and volume:
Trial
No. of light Particles
Pressure (atm)
Ave Pressure (atm)
1
100
23.7
23.6
23.3
23.53
2
110
25.5
25.6
25.0
25.37
3
120
28.1
27.7
27.6
27.80
4
130
30.5
30.2
30.0
30.23
5
140
32.7
33.0
32.9
32.87
Table 4.2: Pressure at different number of light particles
Table 4.3 shows the results of varying mixture of light and heavy particles on pressure at constant
temperature and volume:
Trial
No. of light
Particles
No. of heavy
particles
Ave Pressure
(atm)
1
100
100
46.5
46.9
46.6
46.67
2
110
110
51.4
51.1
51.2
51.23
3
120
120
56.1
56.0
55.7
55.93
4
130
130
60.6
60.8
60.4
60.60
5
140
140
65.1
65.7
65.4
65.40
Pressure (atm)
Table 4.3: Average pressure of mixture
Table 4.4 shows the difference between the sum of average partial pressures of light and heavy
particles, and the average pressure of the mixture:
Trial
Sum of Ave. Partial Pressure
(atm)
Ave. Pressure of Mixture
(atm)
Difference of Ave. Pressures
(atm)
1
46.97
46.67
0.30
2
51.20
51.23
-0.03
3
55.73
55.93
-0.20
4
60.60
60.60
0.00
5
65.77
65.40
0.30
Table 4.4: Results of the Dalton’s Partial Pressure demonstration
Graph 4 the difference of the total pressure of the system when measured as the sum of the partial
pressures and when measured as a whole are close to zero obeying Dalton’s law of partial pressure:
Average Pressures
Sum of Ave. Partial Pressure
Ave. Pressure of Mixture
Pressure (atm)
70
65
60
55
50
45
1
1.5
2
2.5
3
3.5
4
4.5
5
Trial
Graph 4: Sum of partial pressures and average pressures of mixture
Discussion
According to Charles’ law, states that the volume of a given mass of gas varies directly with
the absolute temperature of the gas when pressure is kept constant (Libretext.com, 2021). In terms of
KMT, when the gas heats up, the average kinetic energy of the particles increases, resulting in more
particle collisions with the walls of the containing and a higher pressure. The flexible container may
expand and increase its volume to lessen the number of collisions per square unit area and maintain
constant pressure. When a gas is cooled, the average kinetic energy of the particles lowers, resulting in
fewer collisions and a lower pressure. As a result, the container can contract to increase the number of
collisions per square unit area in order to maintain constant pressure. The results in Table 1 shows the
variation of volume corresponding to the change in temperature in accordance with Charles’ law. As it
can be seen in Graph 1, there is a positive and almost linear proportionality between the temperature
and volume of a gaseous system.
Moreover, Boyle’s law of ideal gases states that the pressure of a given quantity of gas varies
inversely with its volume at constant temperature (Britannica.com, n.d.). In terms of KMT, if the
temperature is constant, so is the average kinetic energy of the particles. However, as the container
shrinks, the lesser the mean free path will be, and the particles is closer to each other resulting to more
collisions occurring against the walls of the container. With this, for the same amount of kinetic energy
there will be an increase in the overall pressure of the system. In contrast, as the container expands,
particles tend to be distant from each other resulting to fewer collisions for the mount of kinetic energy,
causing the pressure to drop. As the results in Table 2 shows, the increase in the volume of the system
leads to a decrease in pressure at constant temperature. Thus, Graph 2 statistically proves that there is
an inverse relationship between the volume and the pressure of the system.
Furthermore, Avogadro’s Principle of ideal gases states that under the same pressure and
temperature conditions, equal volumes of all gases contain the same number of molecules (Lumen,
n.d.). Thus, the volume of a gaseous system is directly dependent to the number of particles contained.
In terms of KMT, more particles contained in the system mean more collisions will occur, which
consequently results in an increase in pressure. The container can expand to reduce the number of
collisions per square unit area while maintaining constant pressure. Conversely, fewer particles indicate
fewer collisions. This results in a reduced pressure. The container can then contract to increase the
number of collisions per square unit area in order to maintain constant pressure. The results in Table 3
shows the increase of volume corresponding to the increase in the number of gaseous particles added
in the container, following Avogadro’s principle. Together with Charles’ law, the Avogadro’s principle
also explains a direct positive relationship between volume and the number of particles, as it can be
seen in Graph 3.
And lastly, according to Dalton’s law of partial pressure, the total pressure of a mixture of gases
is equal to the sum of the partial pressures of the component gases (Khan Academy, n.d.). In terms of
KMT, non-reacting gases do not interact with each other outside elastic collisions, since it is assumed
that there are no attractive or repulsive forces present between particles and ideal gases do not undergo
chemical reaction. Because of this, there is no other reason for a pressure shift when two or more nonreacting gases are mixed together except due to an increase in the number of particles. Thus, simply
combining the individual pressures of the gases in different containers yields the overall pressure of the
mixture these gases in one container. The experimental results in Tables 4.4 demonstrate a small margin
of difference between the total pressure obtained by adding the partial pressures of the particles together
and the pressure of the mixture as a whole. In this case Dalton’s law of partial pressure is in effect.
Conclusion:
The Kinetic Molecular Theory (KMT) provides better understanding about the different
macroscopic properties of gaseous systems in relation to its microscopic activities. Charles' Law,
Boyle's Law, Avogadro's Principle, and Dalton's Law of Partial Pressures are all supported by KMT.
The positive relationship between volume and pressure, as well as the number of particles and volume,
were demonstrated using an ideal gas simulation, as were the negative relationship between pressure
and volume, and the equivalence of the pressure of a mixture of non-reacting gases when measured as
a whole and when measured its components separately. It was obtained by adding the partial pressures
of its constituent gases. It can be concluded that the experiment was a success since the results
theoretically observed the stated gas laws. However, some variations from the expected results can be
blamed to the occurrence of some technical errors at the conduct of the virtual laboratory experiment.
Nevertheless, this experiment provided the experimenters new and sufficient knowledge about perfect
gases which is helpful and relevant to their course.
References:
Britannica.com. (n.d.). Boyle's Law. Retrieved from Britannica:
https://www.britannica.com/science/Boyles-law
Gregerson, E. (n.d.). Gas Laws. Retrieved from Britannica.com:
https://www.britannica.com/science/kinetic-theory-of-gases
Khan Academy. (n.d.). Dalton's law of partial pressure. Retrieved from KhanAcademy.org:
https://www.khanacademy.org/science/chemistry/gases-and-kinetic-molecular-theory/idealgas-laws/a/daltons-law-of-partial-pressure
Letran, D. (2020). The Ideal Gas Laws. Retrieved from Chemistry Libretext:
https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Ma
ps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Physical_Properties_of_
Matter/States_of_Matter/Properties_of_Gases/Gas_Laws/The_Ideal_Gas_Law
Libretext.com. (2021). Charles Law. Retrieved from Chemistry Libretext:
https://chem.libretexts.org/Bookshelves/Introductory_Chemistry/Book%3A_Introductory_Ch
emistry_(CK-12)/14%3A_The_Behavior_of_Gases/14.04%3A_Charles's_Law
Lumen. (n.d.). Avogadro’s Law: Volume and Amount. Retrieved from Lumen Learning:
https://courses.lumenlearning.com/introchem/chapter/avogadros-law-volume-and-amount/