1. The Behavior of Gases
Ch. 12

2. The Properties of Gases
12-1

3. Kinetic Theory (Revisited)
Kinetic Theory of Gases
Gases are composed of hard, spherical particles in constant, rapid, random, straight-line motion.
There are no attractive forces between particles. All collisions are perfectly elastic (no energy lost).
Kinetic Energy is directly proportional to Kelvin temperature.
A gas expands to fill the container. There is empty space between gas particles (explains compressibility). Gas particles occupy negligible volume.

4. Variables describing a gas
4 Variables are used to describe a gas:
Pressure (P) in kilopascals (kPa)
Volume (V) in liters (L)
Temperature (T) in Kelvin (K)
Number of moles (n) – amount of gas particles

5. Factors affecting Gas Pressure
12-2

6. 1) Amount of Gas
(Review) Pressure of a gas = force per unit area that gas particles exert on the walls of their container
Gas particles Pressure
Ex: Inflating a basketball adds more gas molecules which increases the pressure in the ball

7. 2) Volume Volume Pressure
Ex: piston in car engine, forces gas in cylinder to reduce in volume, creating more pressure.

8. 3) Temperature
Remember…as temperature increases kinetic energy increases…
temperature pressure
Ex: heat up a propane tank, gas molecules will move faster and faster… and finally blow up

9. Gas Laws
12-3

10. Boyle’s Law (Pressure vs. Volume)
Boyle’s Law = the pressure and volume of a gas are inversely proportional at a constant temperature.
Pressure Volume
Calculation:
P1V1 = P2V2
DRAW!

11. Calculating Boyle’s Law
Calculation:
P1V1 = P2V2 (P= pressure, V = volume)
Ex: The volume of a scuba tank is 10.0 L. It contains 290 atm of gas pressure in it. What would be the volume of gas at 2.40 atm?
P1 = 290 atm P2 = 2.40 atm
V1 = 10.0 L V2 = ?
290atm(10.0L) = 2.40atm(V2)
V2 = 1208 L

12. Charles’ Law (Temperature vs. Volume)
Charles’ Law = at constant pressure, the volume of a gas is directly proportional to temperature (K).
Volume Temperature
Ex: if you put a balloon in the freezer, it will shrink in size, take it out and it will expand!
Calculation:
V1 = V2
T T2
DRAW!

13. Calculating Charles’ Law
Calculation:
V1 = V2 V = volume, T = temperature (K)
T T2
Ex: A balloon is filled with 3.0 L of helium at 22°C and 760 mm Hg. It is then placed outdoors at 31°C, what will the new volume be? (convert °C to K first!)
3.0 L = V2
295 K 304K
V2 = 3.1 L

14. Gay-Lussac’s Law (Temperature vs. Pressure)
Gay-Lussac’s Law = the pressure of a gas is directly proportional to the temperature, in Kelvin, at a constant volume.
Pressure Temperature
Calculation:
P1 = P2
T1 T2
DRAW!

15. Calculating Gay-Lussac’s Law
Calculation:
P1 = P2 P = pressure, T = temperature (K)
T T2
Ex: If a can at a of pressure is 103kPa at 25°C is thrown into a fire, what will the resulting pressure be at 928°C?
103 kPa = P2
298 K K
P2 = 415 kPa

16. Combined Gas Law Combined Gas Law: P1V1 = P2V2 T1 T2
Ex: A 2.7 L sample of nitrogen is at 121kPa and 288K. If the pressure increases to 202kPa and the temp. rises to 303K, what is the new volume?
121kPa(2.7L) = 202kPa(V2)
288K 303K
V2 = 1.7L

17. Gas Law Overview Boyle Charles Gay-Lussac Combined Gas Law Proportion
Variable
Constant
Calculation
Boyle
Charles
Gay-Lussac
Combined

18. Gas Law Overview Boyle Inverse Pressure, volume Temp P1V1=P2V2 Charles
Proportion
Variable
Constant
Calculation
Boyle
Inverse
Pressure, volume
Temp
P1V1=P2V2
Charles
Direct
Volume, temp
Pressure
V1 = V2
T T2
Gay-Lussac
Pressure, temp
Volume
P1 = P2
T1 T2
Combined
Pressure, temp, volume
P1V1 = P2V T1 T2

19. Ideal Gases
12-4

20. Ideal Gas Law Calculation
Ideal Gas Constant (R) = 8.31 L•kPa/K•mol
Ideal Gas Law = PV = nRT; n = # of moles
Ex: You fill a rigid cylinder will a volume of 20.0L with N2 gas to a pressure of 2.00x104 kPa at 28°C. How many moles of N2 gas does the cylinder contain?
2.00x104 kPa x 20.0L = n x 8.31 L•kPa/K•mol x 301K
n = 160 mol N2

21. Ideal Gas Law + Kinetic Theory
Kinetic theory assumes all gases are ideal gases.
An ideal gas follows all of the gas laws under ALL conditions of pressure and temperature.
An ideal gas does NOT exist; however, the behavior of real gases is very similar.

22. Departures from the Ideal Gas Law
Main 2 differences between real + ideal behavior:
Gas molecules ARE attracted to each other (otherwise could not become a liquid)
Gas particles DO have volume (made up of particles, must have volume)

23. Gas Molecules: Mixtures + Movements
12-5

24. Avogadro’s Hypothesis
Avogadro’s Hypothesis = equal volumes of gases at the same temperature and pressure contain equal numbers of particles.
Thus, at STP, 1 mol = 6.02x1023 atoms, of any gas regardless of size, occupies a volume of 22.4 L

25. At STP!!

26. Dalton’s Law
Partial Pressure = the contribution each gas in a mixture adds to the total pressure.
Dalton’s Law of partial pressures = at constant volume and temperature, the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of the component gases.
Dalton’s Calculation:
Ptotal = P1 + P2 + P3 + …

28. Dalton’s Law Practice
Ex: Air contains O2, N2, CO2, and trace amounts of other gases. What is the partial pressure of O2 if the total pressure is kPa and the partial pressures of N2 = kPa, CO2 = kPa, and other gases = 0.94 kPa?
Ptotal = PN2 + PO2 + PCO2 + Pother
= PO
PO2 = kPa

29. Graham’s Law
Diffusion = the tendency of molecules to move from areas of high concentration to areas of low concentration until the concentration is uniform throughout.
Effusion = process in which a gas escapes through a tiny hole in a container.

30. Graham’s Law of Effusion/Diffusion = the rate of effusion of a gas is inversely proportional to the square root of a gas’s molar mass.
Graham’s Law Calculation:
Rate A = √molar massB
Rate B √molar massA
What?! Gases with lower molar mass move faster than gases with higher molar mass.
Ex: A balloon filled with helium
deflates faster than a balloon
filled with air –
which has a larger mass?

31. Calculating Graham’s Law
Ex: Compare the rates of effusion of nitrogen (B) (molar mass = 28.0g) to helium (A) (molar mass = 4.0g).
RateHe = √28.0g = 5.3g = 2.7
RateN2 √4.0g g
Helium effuses/diffuses 2.7 times faster than nitrogen at the same temperature!!