4. At the end of the topic, the students must have:
a.) recognized the relationship between volume and
pressure at a constant temperature,
b.) practiced the volume-pressure relationship in
solving problems,
c.) valued the importance of Boyle’s Law in real-
life situations.
6. 1. Put the small balloon inside the syringe.
2. Cover the needle hub with your index
finger, pull the plunger and observe what
happens to the balloon.
3. Afterward, push the plunger and observe
what happens to the balloon.
7. 1.As you pull the plunger,
does the pressure inside
the balloon increase or
decrease?
2. As you push the
plunger, does the pressure
inside the balloon increase
or decrease?
1.If the pressure increases,
what happens to the size of
the balloon?
2.If the pressure
decreases, what happens to
the size of the balloon?
8. How would you state
the relationship
between volume and
pressure?
Note: Handle the laboratory
tools properly.
13. Boyle’s Law can be expressed
mathematically as:
Where:
V = Volume
P = Pressure
T = Temperature
n = amount of gas
𝑽 ∝
𝟏
𝑷
𝒂𝒕 𝒄𝒐𝒏𝒔𝒕𝒂𝒏𝒕 𝑻 𝒂𝒏𝒅 𝒏
14. 𝑽 ∝
𝟏
𝑷
𝒂𝒕 𝒄𝒐𝒏𝒔𝒕𝒂𝒏𝒕 𝒌 𝑜𝑟 𝑽 =
𝒌
𝑷
Thus, 𝒌 = 𝑽𝑷
𝒌 = 𝑽𝟏𝑷𝟏 𝒂𝒏𝒅 𝒌 = 𝑽𝟐𝑷𝟐
Whereas, 𝑽1 is the initial volume and 𝑷𝟏 is the
initial pressure.
𝑽𝟐 is the final volume and 𝑷𝟐 is the final
pressure.
16. The inflated balloon that slipped from the
hand of Renn has a volume of 0.50 L at sea level
(1.0 atm) and it reached a height of approximately
8 km where the atmospheric pressure is
approximately 0.33 atm. Assuming that the
temperature is constant, compute for the final
volume of the balloon.
18. Oxygen gas inside a 1.5 L gas tank has a
pressure of 0.95 atm. Provided that the
temperature remains constant, how much
pressure is needed to reduce its volume by ½?
21. A. Directions: Solve for V x P by multiplying volume
and pressure from the table.
Trial Volume (L) Pressure
(atm)
V x P
1 2.0 10.00
2 4.0 5.00
3 8.0 2.50
4 16.0 1.25
22. B. Directions: Plot the data from Table 1 in a graph
by placing the Volume in the X axis and Pressure in
the Y axis. Write a conclusion based from the results
of your graph.
23.
24. A. Directions: Solve for V x P by multiplying volume
and pressure from the table.
Trial Volume (L) Pressure
(atm)
V x P
1 2.0 10.00 20.00
2 4.0 5.00 20.00
3 8.0 2.50 20.00
4 16.0 1.25 20.00
26. 1. A scuba diver needs a diving tank in order to provide breathing
gas while he is underwater. How much pressure is needed for 6.00
liters of gas at 1.01 atmospheric pressure to be compressed in a
3.00 liter cylinder?
2. A sample of fluorine gas occupies a volume of 500 mL at 760 torr.
Given that the temperature remains the same, calculate the
pressure required to reduce its volume by 1/3.
Directions:
Solve the following problems and show your solution. Copy the
problems and write your answer in a one whole sheet of paper.
27. Directions: Answer the question given in a ½
crosswise sheet of paper.
1. Who introduced the Charles’ Law? (Give a
little background)
2. State the Charles’ Law and explain.
3. What is the mathematical expression of the
law?
The air that we breathe in and out is what causes our diaphragm to become bigger and smaller in its size.
As we inhale deeply, there is a force exerted so that the air goes inside our lungs which causes the size of our diaphragm to become smaller as it contracts. And as we exhale, there is lesser force being exerted which causes the size of our diaphragm to become bigger as it relaxes.
While doing this activity, you are going to answer the following questions based from your observation:
NEXT SLIDE!
I will give you 5 minutes only to do this activity. Each group must select one representative who will discuss or explain what happened to their activity and what are the results. Am I clear?
Do you have any questions or clarifications?
Then your time will start now!
To check your work and know more about the concept of the activity that you did a while ago, let us first discuss the relationship between the volume and pressure at constant temperature in Boyle’s Law. But before that, let me introduce to you the man who discovered this law.
He was an Anglo-Irish natural philosopher, chemist, physicist, alchemist and inventor. He was the first to state the relationship between the volume and pressure of gases at constant temperature during the 16th century. Boyle’s Law states that:
He performed an experiment wherein he trapped a fixed amount of air in the J-tube. He changed the pressure and controlled the temperature and then, he observed its effect to the volume of the air inside the J-tube.
He found out that as the volume is increased, the pressure decreases and if the volume is decreased, the pressure increases.
The volume of a gas is inversely proportional to its pressure, if temperature and amount of a gas are held constant.
Removing the proportionality symbol (∝) and using the equality sign (=) the equation will be as follows:
Thus we will have the final equation k = VP or it can be read as the product of Volume and Pressure is constant.
If you are going to consider the initial and final conditions, you will arrive at the following equations:
If the volume-pressure ratios are the same in the initial and final conditions, then we will arrive at this final equation:
Am I still clear class?
To illustrate the mathematical equations, let’s apply Boyle’s Law in solving problems related to volume-pressure relationship in gases.
You can see from the solution that the decrease in pressure from 1.0 atm to 0.33 atm has greatly affected the final volume. In this case, as the pressure on a gas decreases, the gas volume increases because the gas particles can now move farther apart. Am I clear class?
Now, let’s have another example.
That’s right! As you can see, the decrease in volume from 1.5 L to 0.75 L has affected the result of the final pressure. In this case, as the pressure on a gas increases, the volume of the gas decreases because the gas particles are forced closer together. Am I clear class?
Aside from that, understanding Boyle’s law will give us knowledge on how things with pressurized gases can affect its volume such as opening a can or a bottle of soda. Note that when you open the bottle of soda quickly, the gas rushes from everywhere in the form of foam, causing a mess. So, what is the cause of this mess?
This mess occurs because the soda is pumped into the soda bottle by passing the water on carbon dioxide. When you open the bottle, you are actually reducing the pressure on the gas, and the volume of the gas expands. If you remove the cap quickly, the gas pushes out of the bottle. Therefore, you should open the cap slowly and carefully until the gas comes out quietly. Am I understood?
