1. Gases Notes

2. 12.1 A. Physical Properties:
Gases have mass. The density is much smaller than solids or liquids, but they have mass. (A full balloon weighs more than an empty one.)
2. Gases can be compressed. It is very easy to reduce the volume of a gas.
Example: air in a bicycle tire

3. 3. Unlike liquids, gases completely fill their containers.

4. 4. Gases can move through each other
rapidly. Example: the diffusion of food
smells and perfume.

6. 5. Gases exert pressure.

7. 6. The pressure of a gas depends upon
temperature.
high temp. = high pressure,
low temp. = low pressure

8. B. Kinetic-Molecular Theory:
1. Gases are small particles that have mass. These particles are usually molecules, except for the Noble Gases.

9. 2. The particles in gases are separated by relatively large distances.

10. 3. The particles in gases are in constant rapid motion (random).
4. Gases exert pressure because their particles frequently collide with the walls of their container and each other.

11. 5. Collisions of gas particles are perfectly elastic.
Inelastic Collision
Elastic Collision

12. Gas particles do not slow down when hitting each other or the walls of their container… How do we know?

13. 6. Temperature of a gas is simply a measure of the average kinetic energy of the gas particles. High temp. = high KE, Low temp. = low KE

14. 7. Gas particles exert no force on one another
7. Gas particles exert no force on one another. Attractive forces are so weak between particles they are assumed to be zero.

15. C. Measuring Gases:
The following 4 variables will be used to do gas calculations:
n - amount of a gas, it is measured in moles
V - volume of a gas, it is measured in L (dm3) or mL (cm3)
“STP” = Standard Temperature & Pressure:
T- Standard Temperature: 0 oC = 273 K
Taken in oC converted to Kelvin (K) T(K) = °C + 273
Ex #1) 22°C = K °C = 373 K
-27.3°C = K °C = 0 K

16. P - Standard Pressure: can be measured using the following units:
(sea level pressure)
1 atm (atmospheres) 760 mm Hg
101,325 Pa (Pascals) 760 torr
kPa (kilopascals) 14.7 lb/in2 (psi)
Ex #2) Convert atm to kPa:
Ex #3) Convert 98,500 Pa to mm Hg:

17. 2. Atmospheric pressure is the pressure exerted by the air in the atmosphere. This pressure varies with altitude and water vapor content.

18. 3. Atmospheric pressure is measured with a barometer
3. Atmospheric pressure is measured with a barometer. This is a glass tube sealed at one end and filled with Hg.
Why toxic mercury and not water?

19. Atmospheric Pressure
Affects Weather:
Changes with Altitude:

20. Boiling Point Changes w/Altitude & Pressure

21. 12.2 The Gas Laws
A. Boyle’s Law: Pressure - Volume Relationship. The pressure & volume of a sample of gas at constant temperature are inversely proportional to each other.
Indirect
P1V1 = P2V2

24. Doubling Pressure = ½ Volume Quadrupling Pressure = ¼ Volume

26. Boyle’s Law

27. Ex #4) A gas has a volume of 300. mL under a pressure of 740
Ex #4) A gas has a volume of 300. mL under a pressure of 740. mm of mercury. If the temperature remains constant, calculate the volume when under a pressure of 750. mm Hg.
P1V1 = P2V2

28. B. Charles’ Law: Temperature - Volume Relationship
B. Charles’ Law: Temperature - Volume Relationship. At constant pressure the volume of a fixed amount of gas is directly proportional to its absolute temperature.
Direct
*Temperatures must be in Kelvin!
K = °C

30. Ballon in cool and cold water:

32. What is the Celsius temperature of
Ex #5)
What is the Celsius temperature of
68.20 mL of methane, if it occupies a volume of L at 50.0°C?
22.00 mL
22.00 mL

34. 1 mole gas = 22.4 L = 6.02 x 1023 particles at STP (273 K & 1 atm)
C. Avogadro’s Law: Amount - Volume Relationship. Equal volumes of gases at the same temperature and pressure contain an equal number of particles.
1 mole gas = 22.4 L = 6.02 x 1023 particles at STP (273 K & 1 atm)
constant
volume
22.4L
4 He
molar mass
222 Rn

35. He O2 Rn
Therefore because of Avogadro’s Law if these three gases have the same number of particles and are at the same temperature and pressure, they must take up the same volume.

36. Molar Mass does not affect volume of a gas

37. D. Dalton’s Law of Partial Pressures: The sum of the partial pressures of all the components in a gas mixture is equal to the total pressure of the gas mixture.
Ex #6) A flask contains a mixture of oxygen, argon, and carbon dioxide with partial pressures of 745 torr, atm, and torr respectively. What is the total pressure in the flask?
PT = Pa + Pb Pc …

39. Dalton’s Law of Partial Pressures

40. Ex #7) The total pressure of a mixture of helium and neon is 498 mm Hg
Ex #7) The total pressure of a mixture of helium and neon is 498 mm Hg. If helium is 20.0 % of the mixture, what is the partial pressure of helium in mm Hg?

42. E. The Combined Gas Law: “Choyles” This law can be used to determine how changing two variables at a time affects a third variable.

43. Ex #8) A gas occupies 72. 0 mL at 25 °C and 198 kPa
Ex #8) A gas occupies 72.0 mL at 25 °C and 198 kPa. Convert these to standard conditions. What is the new volume?
P1 = 198 kPa P2 = kPa
V1 = mL V2 = ?
T1 = 298 K T2 = 273 K
(273 K)
( kPa)

44. 12.3 A. Ideal Gas Law
Although no “ideal gas” exists, this law can be used to explain the behavior of real gases under ordinary conditions.
P = pressure (atm)
V = volume (L or dm3)
n = number of moles
R = L•atm/mol•K
universal gas constant
T = Kelvin temperature
PV = nRT

45. B. Ideal Gas Law & The Kinetic – Molecular Theory:
Under normal conditions (temperature and pressure) gases behave ideally.

46. n ____ P ____ (more gas, ____________)
more collisions

47. T ____ P ____ (moves faster, ______________)
more collisions

49. V ____P ____ (smaller volume, ____________)
more collisions

50. 2. Gases at low temperatures and high pressures do NOT behave ideally
2. Gases at low temperatures and high pressures do NOT behave ideally. As you decrease the volume of a gas, the volume of the particles themselves becomes significant. The Kinectic-Molecular Theory & Ideal Gas Law assumes the gas particles have no volume of their own.
3. Second, the attractive forces which are very small when the particles are moving fast, become larger as they slow down.

51. Ex #9) How many grams of carbon dioxide occupy a volume of 36
Ex #9) How many grams of carbon dioxide occupy a volume of 36.9 mL at 158 kPa and 72 oC?
PV = nRT
1st convert units to “R” units
( L•atm/mol•K) (345 K)
( L•atm/mol•K) (345 K)
n = mol CO2

52. C. Lifting Power of Gases:
1. Uses a gas “lighter” than air (smaller molar mass.)
Ex) The Hindenburg used hydrogen
Ex) Today’s blimps use helium
2. Hot Air balloons heat air to lower its density.
3. Effusion is the movement of gas atoms or molecules through a hole so tiny they move one particle at a time. Smaller particles effuse faster than larger particles.
Ex) → → H2
fastest He
rare/expensive N2
(air) plentiful/ cheap/slowest

54. Effusion

55. Effusion

58. The Hindenburg was a Dirigible (a balloon that has engines to control its horizontal movement)

63. Passenger Quarters had 25 two-berth cabins

64. Titanic - 882 feet Hindenberg – 804 feet

66. May 6, 1937 Lakehurst, NJ

68. Memorial in Lakehurst, NJ