Sections 13.1-13.4: Team Learning Worksheet
1. Why does air come out of a bicycle tire when you open the valve, even if the valve is pointing up?
The molecules inside the tube are crammed into that small area. There are so many they are creating a high pressure against the inside
of the tube and the valve. When the valve opens, there is much higher pressure than the atmospheric pressure pushing inward and the
tire deflates.
2. Why do aerosol cans come with the warning “Contents under pressure. Do not heat.”?
If the can is heated the pressure (or force/area) of the collisions will increase. The can is already at high pressure inside, if we increase
it, the can could burst.
3. Why is it insufficient to only know the volume of a gas when you are trying to determine the amount of gas
in a sample? If you know the volume, what other information would you need? Explain your answer.
If we want to know the amount of a gas it is important to know what the volume as well as the temperature and the pressure. For a
given volume container, I can cram in lots of molecules, but that would increase the pressure. I could also cram in more molecules if I
slow the particles down in that same volume and pressure. It is important to know all three to determine the amount. PV=nRT
4. Why is it incorrect to say that a sample of helium at 50 ºC is twice as hot as a sample of helium at 25 ºC?
Temperature must be measured in Kelvins. If the temperature is 50K and 25K then we can say one is twice as hot as another, but
323K is not twice as hot as 298K.
Sections 13.5-13.6: Team Learning Worksheet
5. Use the ideal gas law equation and derive a formula that relates pressure and number of moles of gas at
constant temperature and volume.
Divide both sides by P2V2
P1V1 n1R1T 1
P1V1
n1R1T 1
 Since P2V2 is equal to n2R2T2, Substitute into equation


P2V2
P2V2
P2V2 n 2 R 2T 2
P1 n1
Cancel out or get rid of any constants 

P2 n2
Start with P1V1=n1R1T1

6. Use the ideal gas law to derive Boyle’s law, Charles’s law, and the combined gas law.
Same as number 5 but crossing out different constants to give
P1V1 1

P2V2 1
or
V1 T 1

V2 T 2
or
P1V1 T 1

P2V2 T 2
7. Consider a mixture of equal masses of helium and hydrogen gases. Which gas exerts the greater partial
pressure? By what factor? Support your answer.
Helium will exert more pressure at a certain speed (temperature). It has twice the mass and will have twice the kinetic energy, thereby
exerting twice the pressure.
8. Consider a container with 3.40 g of N2 gas and an unknown mass of CO2 gas. If we know that the partial
pressure of the CO2 is twice that of the N2, determine the mass of the CO2 gas in the container.
If the partial pressure is twice that of the N2, then we have twice as much gas or twice as many moles. 3.40gN2
We can say there are 0.24mol CO2 and calculate the mass. 0.24mol CO2
44 gCO 2
= 10.56g CO2
1molCO 2
1molN 2
=0.12mol N2
28 gN 2
Sections 13.7-13.11: Team Learning Worksheet
9. Explain why a sealed balloon will contract if it is placed in a freezer.
The particles currently ramming into the inside wall of the balloon will slow down when placed in a cooler environment. The energy
that they hit the walls with will be less (KE=1/2mv2) and the pressure inside will be less. The outside (atmospheric) pressure will then
be able to crush the balloon along with the latex contracting.
10. At a given temperature and pressure, 2.0 L of N2 gas reacts completely with 6.0 L of H2 gas to form 4.0 L of
product. What is the formula of the product? Support your answer.
For every 22.4L at STP, ANY gas is 1mole. 2.0L = 0.09moles, 6.0L = 0.27moles, 4.0L = 0.18moles. This is a 1:3:2 ratio. For every
one N2 molecule we make 2 of our product. Each product will contain one N atom. For every 3 H2’s used we make 2 products. We
will have 3 H’s in each product. The formula will be NH 3.
11. Use the kinetic molecular theory to explain why we need to use the Kelvin scale when dealing with gases.
In Celsius temperature we have negative values, however, in Kelvin’s we do not. Using PV=nRT we would calculate a negative
pressure or a zero pressure when using negative or zero Celsius temperatures. KMT says that we cannot have a negative pressure
from a gas or a zero pressure unless the molecules in the gas are stopped. The molecules in a gas at a negative Celsius temperature are
sill moving and still causing pressure.
12. Draw a molecular level sketch of the interior of a sealed steel container of helium gas on a warm summer
day. Then draw a molecular level sketch of this same container on a cold winter day. Explain the drawings.
The pressure and frequency of collisions in the cold temperature container will be less. This is due to the lower speeds and, therefore,
the lower kinetic energies of the particles.
13. Draw a molecular level sketch of the interior of a sealed steel container of helium gas. Then draw a
molecular level sketch of this same container after half of the helium has been removed. Explain the drawings.
The pressure in the second container will be less due to the particles hitting the walls less frequently. The speeds will remain the same
along with the temperature and volume.
14. In what type of a set-up is the pressure constant for a gas in a closed system? Use the kinetic molecular
theory to explain why increasing the number of moles of gas at constant temperature and pressure causes an
increase in the volume.
The pressure will remain constant if the temp or the volume could change if the number of moles changed. If the number of moles is
increased in a container and the pressure and temperature remain constant, the only variable that can compensate would be the
volume. The volume would have to increase to accommodate all of the new molecules.