1. Gas and Pressure

2. Review Kinetic Vs. Potential Energy

3. Kinetic Molecular Theory
Explains the behavior of gases
Real gases– do not follow all the assumptions of this theory.
Theoretical, applies to an ideal gas
Gas following all the assumptions of the theory.

4. Kinetic Molecular Theory
Gases are made up of TONS of particles, constantly moving, and spread out.
Gas particles drive straight until they hit/collide with something (ex. Wall, particles).
Small particles, HUGE space for them to roam! Gas volume mainly empty space

5. Kinetic Molecular Theory (cont.)
4) No force of attraction! –gas particles randomly move around
5) When particles collide with each other or container wall, elastic collisions are the result.
-elastic collisions—no kinetic energy is lost when gas particles collide
6) Temperature determines average kinetic energy of gas particles.
-not all gas particles are moving at the same kinetic energy

7. Properties of Gases 1) Expandable 2) Fluid 3) Low Density No shape
No defined volume
Fills whatever container is available to it.
Move quickly and no attraction
2) Fluid
Gases move similar to liquids
No attraction to worry about
3) Low Density
Least dense state of matter for most substances
Due to distance between particles

8. Properties of Gases (cont.)
4) Diffusion
Gas particles mix with each other and disperse
“Spontaneous mixing of gas particles from 2 substances due to random motion” (ex. Perfume)
5) Effusion
Describes gas movement through a small opening. (ex. Tire puncture)
Related to how fast gas particles can move.
6) Compressibility
Gas particles are able to be packed close together.
Decreased volume

9. Pressure “force per unit area on a surface”
Gas particles colliding against container and creating force
Force a gas exerts on its surroundings
Unit = Newton (N)
Amount of pressure dependent on Volume, Temperature, and particle/molecule number.
Ex. Atmospheric pressure

10. How do we measure pressure?
Barometer– atmospheric pressure
Evangelista Torricelli (1600s)
Mercury falls to 760 mm
Air pressure holds 760 mm mercury column.

11. Barometer Vacuum 760 mm Hg 1 atm Pressure
The pressure of the atmosphere at sea level will hold a column of mercury 760 mm Hg.
1 atm = 760 mm Hg
760 mm Hg
1 atm Pressure

12. Pressure Units Common: mmHg (millimeters of Hg)
Measurements done in “atmosphere units”
1 atmosphere of pressure (atm) = (sea level, 0° C)
760 mmHg
760 torr
x 105 Pa
kPa

13. Example 1:
1.75 atm of pressure to mmHg

14. Example 2:
570 torr of pressure to atmospheres.
kPa?

15. Gas Laws
Describes the relationship between variables associated with gases
Volume (V)
Temperature (T)
Pressure (P)
Concentration/amount of gas (n)
**Two variables change in relation to each other while the remaining variables are held constant.

16. Gas Laws Boyle’s Law Charle’s Law Gay-Lussac Law Avagadro’s Law
Combined Gas Law
Ideal Gas Law

17. Boyle’s Law
Pressure is INVERSELY proportional to the volume of the gas.
Temperature and particle amount constant
P , V 
P , V 
Pressure and volume relationship
P1V1 = P2V2
Remember to keep units the same.

18. Examples
What would happen to the volume of a balloon filled with L of H2 gas collected at mmHg if the atmospheric pressure increased to mmHg? (temperature is constant)
Calculate the volume of a balloon that could be filled at 1.00 atm with helium in a 2.50L compressed gas cylinder in which the pressure is 200 atm at 25°C.
13.3 L in scuba tank, pressure goes down while volume goes up
2) 0.349L H2, pressure goes up while volume goes down

19. Charle’s Law T1V2 = T2V1 Temperature and volume relationship
The volume of a gas is DIRECTLY proportional to the temperature
T , V 
T , V 
V1 = V2 SOOO T1 = T2
T1V2 = T2V1
Remember to keep units the same. Temperature MUST be in Kelvin

20. Examples
A sample of O2 gas with a volume of 0.357L was collected at 21°C. Calculate the volume of the gas when it is cooled to 0°C if the pressure remains constant.
How hot will a 2.3L balloon have to get to expand to a volume of 400L? Assume the initial balloon temperature is 25°C.
0.3315L Temperature decreases, volume decreases
51, 826°K, volume increases, temperature increases

21. Gay-Lussac’s Law P1T2 = P2T1 Pressure and temperature relationship
Pressure results from molecular collisions
Pressure of gas is DIRECTLY proportional to temperature.
P , T 
P , T 
P1 = P2 SOOO T1 = T2
P1T2 = P2T1
Remember to keep units the same. Temperature MUST be in Kelvin

22. Examples
A coke can has 5.00atm of gas at 21°C. Calculate the pressure inside the can when it is found in a warehouse during the summer at 38°C.
The pressure of my tires before a road trip to Wyoming was 1.5atm at 25°C. After returning to North Carolina, my tire pressure is 1.7atm. What is the temperature (in °C) outside?
5.29atm, increase temperature, increase pressure
65°C°, increase temperature, increase pressure

23. Gas Worksheet

24. Avogadro’s Law
Volume of a gas is DIRECTLY proportional to # of gas particles (moles of gas)
Temperature and Pressure are held constant
V1 = V n1 = n2
 # gas particles,  volume
 # gas particles,  volume
Ex. Blowing up a balloon

25. Ideal Gas Law Describes the general relationship among the variables:
Temperature
Pressure
Volume
Number of moles of gas
Enables us to determine the value of a variable if the other three variables are known

26. Ideal Gas Law (cont.) PV = nRT P = pressure (atmospheres)
V = volume (liters)
T = temperature (Kelvin)
n = moles of the gas
R = Latm/molK (ideal gas constant)

27. Examples
Many gases are available for use in the laboratory among compressed gas cylinders stored at high pressures. Calculate the mass of O2 (in grams) that could be stored at 21°C and 170atm in a cylinder with a volume of 60.0L.
Calculate the molecular weight of butane if g of the gas fills a 250.0ml flask at a temperature of 24.4°C and a pressure of mmHg
13529g oxygen gas
58.13g/mol butane

28. Examples (cont.)
3) Calculate the density in grams per liter of O2 gas at 0°C and 1.00 atm
3) 1.43 g/L oxygen gas

29. Combined Gas Law Do variables remain constant for gases???
Temperature, pressure, and volume are CONSTANTLY changing for a gas based on the conditions
Gas amount (n) is constant

30. Combined Gas Law (cont.)
Combination of all three laws into one equation (Boyle’s, Charle’s, and Gay-Lussac’s Laws)
Describes the relationship between pressure, volume, and temperature
Focus on initial and final conditions
Rearrange ideal gas law with the gas constant (R) remaining the same

31. Combined Gas Law
P1V1 = P2V2
T T2
Temperature—Kelvin

32. Examples
A gas has a volume of 80.0ml at 27°C and atm. What volume will the gas have at standard conditions?
A gas has a volume of 60.0ml at standard conditions. This volume is reduced to 10.0ml at 25.0°C. What is the necessary pressure for this volume reduction?
14.56ml
2) 6.55 atm

33. Dalton’s Law of Partial Pressures
Pressure of each gas DIRECTLY proportional to amount of moles of a gas
Increase gas particles, increase pressure
Decrease gas particles, decrease pressure
Partial pressure—
Pressure of one gas that contributes to the total pressure in a mixture of gases
Total mixture pressure----
The sum of the individual gas pressures in a mixture

34. Dalton’s Law of Partial Pressures
Total Pressure (PT) of gas mixture =
Sum of partial pressures of each gas in the mixture
PT = P1 + P2 + P3

35. Example
Calculate the partial pressure (in mmHg) exerted by the 4 main gases in air at 760 mmHg: nitrogen, oxygen, argon, and carbon dioxide. Their abundance by volume is 78.08%, 20.95%, 0.934%, and 0.035% respectively.
Nitrogen gas = 593.4mmHg
Oxygen gas = mmHg
Argon gas = 7.10 mmHg
Carbon dioxide gas = 0.27mmHg

36. Water Displacement with Dalton’s Law
How do we collect and measure gases? Water displacement
Gas displaces water but the gas is mixed with water vapor
Application of Dalton’s Law allows the adjustment for the amount of water vapor to be made so just the amount of gas collected can be measured.

37. Water Displacement with Dalton’s Law (cont.)
Water vapor is mixed in with gas of interest so need to separate.
PT = Pgas + Pwater
look up vapor pressure of water at different temperatures (p. 196, Table 15. 4)

38. Examples
A sample of nitrogen gas is collected over water at a temperature of 23.0°C. What is the pressure of the nitrogen gas if atmospheric pressure is 785 mmHg?
A student has stored ml of neon gas over water on a day when the temperature was 27.0°C. If the barometer in the room reads mmHg, what is the pressure of the neon gas in the container?
764 mmHg of nitrogen gas
717 mmHg neon gas

39. Homework
Gas Worksheet #2