1. CHAPTER 14
GAS LAWS

2. Boyles Law: Pressure (kPa) Volume (L)
-for a given mass of gas at constant temperature, the volume of the gas varies inversely with pressure.
Pressure increases, Volume decreases
Inversely proportional
Smaller volume- more pressure b/c particles collide more often
EX: syringe
P1V1=P2V2

3. P1V1=P2V2
where P1 and V1 are the pressure & volume before the gas expands
and P2 and V2 are the pressure & volume after the gas expands.

5. Syringe— Piston can slide in (towards water) or out (towards hand) No cap—Vol. Inc while Pr. Dec Piston slides out Capped—Vol. Dec while Pr. Inc Piston slides in

6. Boyles Law: Pressure (kPa) Volume (L)
Practice: If you had a gas that exerted 202 kPa of pressure and took up a space of 3.00 liters, & you decided to expand the space to 7.00 liters, what would be the new pressure? (temperature remains constant)
Answer: P1V1 = P2V2
So, P1 = 202 kPa, V1 = 3.00L, V2 = 7.00L, and you need to solve for P2, the new pressure. Plug the numbers into the equation, & you have:
(P1V1)/V2 = P2
(202 kPa) x (3.00 L) = (P2) x (7.00 L). P2 = 86.6 kPa.

7. Charles’ Law: Temp. (K) Volume (L)
-the volume of a fixed mass of gas is directly proportional to its Kelvin temperature if the pressure is kept constant.
Temperature increases, Volume increases
Directly proportional
Higher temperature, gas particles speed up and move farther away from one another
EX: heating a sealed container
V1 = V2
T T2

8. V1 = V2
T T2
where V1 and T1 are the volume & temperature before the gas expands
and V2 and T2 are the volume & temperature after the gas expands.

10. Charles’ Law: Temp. (K) Volume (L)
Practice: If you took a balloon outside that was at 20oC at 2L in volume, & it heated up to 29oC, what would its volume be? Assume constant pressure.
Answer: V1 / T1 = V2 / T2
V1 = 2.0L, T1 = 20oC, T2 = 29oC, you must solve for V2.
Wait!! You have to convert the temperatures to Kelvin So: T1 = 20oC = 20oC = 293 K T2 = 29oC = 29oC = 302 K
Now, you can plug in the numbers and solve for V2.
2.0L = V2
293 K K and V2 = 2.1L

11. Gay-Lussac's Law : Temp. (K) Pressure (kPa)
-the pressure of a fixed mass of gas is directly proportional to its temperature if the volume is kept constant.
Temperature increases, Pressure increases
Directly proportional
P1 = P2
T T2

13. -combined the three gas laws into a single expression.
Combined Gas Law:
-combined the three gas laws into a single expression.
V1P1 = V2P2
T T2
The Vice President is higher than the
Treasurer
BOYLE’S LAW
CHARLES’ LAW
GAY-LUSSAC’S LAW

14. COMBINED GAS LAW DOESN‘T:
ACCOUNT FOR THE AMOUNT OF GAS
IT JUST LOOKS AT PRESSURE, VOLUME, & TEMPERATURE

15. Ideal Gas Law:
-allows you to solve for the number of moles of a gas(n).
PV = nRT
P=pressure in kPa or atm
V=volume in L
T=temperature in K
n=#of moles
R=Ideal gas constant (R) is 8.31 (L x kPa)
(K x mol)
(L x atm)

16. Ideal Gas Law: PV = nRT Practice:
A sample of O2 gas has a volume of 4.52L at a temp. of 10oC and a pressure of kPa. Calculate the number of moles of O2 gas present in this sample.
Answer: Rearrange to solve for n (number of moles):
n = PV/RT.
P = kPa,
V = 4.52L,
T = 10oC = 283K, and
R = 8.31 L x kPa/K x mol.
n= (110.5 kPa)(4.52L)/(8.31L kPa/K mol)(283 K) = .212 moles.

17. Do Now
If 4.50 g of methane gas (CH4) is in a 2.00-L container at 35°C, what is the pressure inside the container?
Solution: PV=nRT - Ideal Gas Law Equation.
P = ?
V = 2.00 L
R = 8.31 L * kPa
T = 308 K
n = 4.50 g CH4 x
1 mol / 16.0 g CH4 = mol
mol * K

18. Do Now
If 4.50 g of methane gas (CH4) is in a 2.00-L container at 35°C, what is the pressure inside the container?
Solution: PV=nRT - Ideal Gas Law Equation.
P = 3.60 x 102 kPa.
V = 2.00 L
R = 8.31 L * kPa
T = 308 K
n = 4.50 g CH4 x
1 mol / 16.0 g CH4 = mol
mol * K

19. Mathematical Relationship
Each slice weighs 0.5 lbs. Eight slices make up a pie.
How much does this pie weigh? What can I ignore in my calculation?

20. Mathematical Relationship
Each lego block has dimensions 4 cm x 2 cm x 1 cm
How tall is the stack of lego blocks?

21. History
John Dalton ( ) was an English chemist and physicist born in Cumberland, England.
Early in life, influenced by meteorology.
Researched color-blindness a.k.a. “Daltonism.”
Most notably known for compiling fundamental ideas into a universal atomic theory.
His interest in gases and gas mixtures lead him to investigate humidity. This ultimately lead to Dalton’s Law.

22. Observe and Find a Pattern
Partial Pressures of Compressed Air
(Assuming air is 80% Nitrogen and 20% Oxygen)
Depth (meters)
Absolute Pressure
P O2
P N2
14.7
2.94
11.76 33
29.4
5.88
23.52 66
44.1
8.82
35.28 99
58.8
47.04
132
73.5
14.70
58.80
165
88.2
17.64
70.56
198
102.9
20.58
82.32
231
117.6
23.53
94.08
264
132.3
26.46
105.8
297
147.0
29.40

23. Explain
The Ideal Gas Law holds for virtually any gas, whether pure or a mixture, at ordinary conditions for two reasons:
Gases mix homogeneously in any proportions.
Each gas in a mixture behaves as if it were the only gas present.
Explain why we can use a gas mixture, such as air, to study the general behavior of an ideal gas under ordinary conditions.

24. Predict and Test
What is the relationship between the number of particles in containers A and C and the partial pressures of A and C.
Predict what the reading will be for container T.

25. Predict and Test
What is the relationship between the number of particles in containers A and C and the partial pressures of A and C.
Predict what the reading will be for container T.

26. Derive
Based on the outcome of the experiment relative to the prediction, make a judgment about whether the experiment disproved the hypothesis or not.
Write a mathematical expression that relates absolute (total) pressure to the individual gas pressures.
Ptotal = P1 + P2 + P3 + …this is known as Dalton’s Law of Partial Pressures. How can we qualitatively explain this mathematical relationship?

27. Mathematical Application
It was a dark and stormy night and your “friend” has locked you in the lab without a key. You remembered that the oldest window in the lab is always cracked open on a diagonal that won’t ever seem to budge completely. Someone calculated that you need a total pressure of 7.10 x 106 kPa to exert a strong enough force to bust open the window. Coincidentally, you have 2.00 moles of Ne, 4.00 moles of Xe, and 6.00 moles of Ar in a 5.00-L vessel at 27°C standing next to your lab station. Is the total pressure of the mixture enough for you to escape and make it to the club in time or will you be left alone in the lab synthesizing aspirin to relieve you of your torment?

28. Observe and Explain
KClO3
Gas produced is less dense than water, so it replaces the water in the test tube.
Gas collected is not pure because it contains vapor from the water.

29. Dalton’s Law of Partial Pressure:
-for a mixture of gases in a container, the total pressure exerted is the sum of the pressures that each gas would exert if it were alone.
Ptotal=P1 + P2 + P3 +…..

30. Dalton’s Law of Partial Pressure:
Practice:
Air contains O2, N2, CO2, and trace amounts of other gases. What is the partial pressure of O2 (PO2) at kPa of pressure if PN2 =79.10 kPa, PCO2 =0.040 kPa, and Pothers =0.94 kPa?
Answer: Ptotal = kPa
-Ptotal = PO2 + PN2 + PCO2 + Pothers
-PO2 = Ptotal - (PN2 + PCO2 + Pothers)
-PO2 = kPa - (79.10 kPa kPa kPa)
-PO2 = kPa

31. Ideal vs. Real Gases
An ideal gas is a hypothetical concept. No gas exactly follows the ideal gas law.

32. Ideal Gases Always a gas Not attracted to one another Have no volume
Follow the gas laws for all conditions of pressure and temperature
Ideal gases DO NOT exist!!!!!!!!

33. Real Gases
Can be liquified and sometimes solidified by cooling and applying pressure
(ideal gases can not)
Have a finite volume
Are attracted to one another, especially at low temperatures

34. Graham’s Law of Effusion:
the rate of effusion and diffusion of a gas is inversely proportional to the square root of its molar mass .
Lighter gas goes faster.
Heavier gas goes slower.
To figure out how much faster take the square root of the gfm of the heavier gas and divide by the square root of the gfm of the lighter gas.

36. Graham’s Law of Effusion:
Practice:
Does He effuse faster or slower than O2? What is the relative rate of diffusion of He compared to O2?
Answer: = 2.8, He is 2.8 times much faster