1. Gases Dr. Chin Chu River Dell Regional High School

2. Particles in a Solid, Liquid and Gas

3. Random Motion of Gas Particles

4. II. Gas Pressure A. Pressure is force per unit area
1. result of particle collisions
2. measured by a barometer
3. influenced by temperature, gas
volume, and the number of gas
particles
a. as the number of particle
collisions increases the pressure
increases

5. Kinetic Theory A. Assumptions 1. gas particles do not attract each
other
2. gas particles are very small
3. particles are very far apart
4. constant, random motion
5. elastic collisions
6. kinetic energy varies with temperature

6. B. Properties of Gases
1. low density (grams/liter)
2. can expand and can be compressed
3. can diffuse and effuse
a. rate related to molar mass
b. diffusion is the movement of
particles from an area of greater
concentration to an area of lesser
concentration
c. effusion is the movement of gas
particles through a small opening

7. B. Units of Pressure
1. pascal - 1Pa = 1N(newton)/m2
(meter squared)
2. psi (pounds/square inch)
3. mm Hg or torr
4. atm (atmospheres)
* 1 atm = 760 mm Hg = 760 torr = KPa = 14.7 psi (pounds per square inch)

8. Measuring Atmospheric Pressure

9. Aneroid Baramoter Mercury Barometer

10. II. The Gas Laws
A.Boyle’s Law (P1V1 = P2V2 )inverse relationship
1.As the volume of a gas increases the pressure decreases (temperature remains constant)
2. Example A sample of gas in a balloon is
compressed from 7.00 L to 3.50 L. The pressure
at 7.00L is 125 KPa. What will the pressure be at
2.50L if the temperature remains constant?
P1 = 125 KPa P2 = X V1 = 7.00L V2 = 3.50L
(125)(7.00) = (X) (3.50)
X = 250.KPa

11. Boyle’s Law

12. As volume increases the pressure decreases when temperature remains constant

13. Boyles Law

14. B. Charles’ Law V1 = V2 must use kelvin
T1 T temperature
As the temperature of a gas increases the volume increases (direct relationship)
2. Example A gas sample at 20.0 C occupies
a volume of 3.00 L. If the temperature is raised to
50.0 C, what will the volume be if the pressure
remains constant?
V1 = 3.00L V2 = X T1 = 293K T2 = 323K
3.00 = X X = (3)(323) X = (3)(323)
X = 3.31 L

15. Charles’ Law – Temperature increases – volume increases

16. Charles’ Law

17. 1. as the temperature increases the pressure
C. Gay Lussac’s Law P1 = P2
T1 T2
1. as the temperature increases the pressure
increases when the volume remains constant
2. Example The pressure of a gas in a tank is
4.00 atm at 200.0C. If the temperature rises
rises to 280.0C, what will be the pressure of the
gas in the tank?
P1 = 4.00 atm P2 = X T1 = 473K T2 = 553K
4.00 = X X = (4)(553) X = (4)(553)
X = 4.68 atm

18. D. Combined Gas Law P1 V1 = P2 V2 T1 T2
1. Combines Boyle’s, Charle’s and Gay Lussac’s
2. Example A gas at 70.0KPa and 10.0C fills a
flexible container with an initial volume of 4.00L
If the temperature is raised to 60C and the pressure
is raised to 80.0 KPa, what is the new volume?
P1 = 70.0 KPa P2 = 80.0 KPa V1 = 4.00L V2 = X
T1 = 283K T2 = 333K
(70.0)(4.00) = (80.0)(X)
X = (333)(70.0)(4.00)
(283)(80.0) X = 4.12L

19. E. Dalton’s Law of Partial Pressures Ptotal = P1 + P2 + P3 + .....Pn
The total pressure of a mixture of gases is equal to
the sum of the pressures of all the gases in the
mixture
1. Example Find the total pressure for a mixture that contains four gases with partial pressures of 5.00 kPa, 4.56 kPa, 3.02 kPa and 1.20kPa.

20. Dalton’s Law Partial Pressures

21. Dalton’s Law of Partial Pressures

22. 2. Suppose two gases in a container have a total
pressure of 1.20 atm. What is the pressure of gas
B if the partial pressure of gas A is 0.75 atm?
3. What is the partial pressure of hydrogen gas in
a mixture of hydrogen and helium if the total
pressure is 600.0mmHg and the partial pressure of
helium is 439 mmHg?

23. III. Avogadro’s Principle
A. Equal volumes of gases at the same temperature and pressure have the same number of particles
B. Molar Volume (22.4 L at STP)
1. volume of one mole of gas particles at
STP(standard temperature and pressure)
0C and 1.00 atm (760mm Hg)
* 1 mole of any gas at STP = L
2. conversion factors 1 mol L
22.4 L mol

24. Avogadro’s Principle

25. Equal volumes of gases at the same temperature and pressure contain the same number of particles

26. C. Sample Problems
1. Calculate the volume occupied by .250 mol of
oxygen gas at STP.
2. Calculate the number of moles of methane gas
in a 11.2 L flask at STP.

27. 3. Calculate the volume of 88.0 g of CO2 at STP.
4. How many grams of He are found in a 5.60L balloon at STP?

28. 5. Calculate the density of H2 at STP.
D = molar mass
molar volume
6. Calculate the molar mass of a gas that has a
density of 3.2 g/L.

29. IV. Ideal vs Real Gases
A. Ideal compared to Real Gases
1. ideal gas
a) particles do not have volume
b) there are no intermolecular
attractions
c) all particle collisions are elastic
d) obey all kinetic theory
assumptions

30. 2. real gases behave like ideal gases except when
a) pressure is very high
b) temperatures are low
c) molecules are very large
d) spaces between particles is small
(small volume)

31. B. Ideal Gas Law - PV = nRT
1. pressure ( atm,mm Hg, KPa)
2. volume (liters)
3. temperature (kelvin)
4. number of moles (n)
5. R = constant (L) (pressure unit*)
(mol) (K)

32. unit for pressure determines which
constant must be used in the Ideal Gas Law
PV = nRT
a) R = (pressure is mm Hg)
b) R = (pressure is atm)
c) R = (pressure is KPa)

33. Ideal Gas Law PV = nRT

34. What Principles and/or Laws are Ilustrated?

35. C. Application Problems (PV = nRT)
1. How many moles of O2 are in a 2.00L container
at 2.00 atm pressure and 200K?
2. Calculate the volume occupied by 2.00 mol of N2 at 300K and .800 atm pressure.

36. 3. What is the pressure in mm Hg of .200 moles of
gas in a 5.00 L container at 27C?
4. Calculate the number of grams of oxygen in a 4.00 L sample of gas at 1.00 atm and 27 C.