1. Chapter 10: Characteristics of Gases
The Kinetic-molecular theory-particles of matter are always in motion
The kinetic-molecular theory has 5 assumptions:
1. Gases consist of large numbers of tiny particles that are far apart relative to their size (gases are tiny particles that are far apart)

2. 2. Collisions between gas particles and between particles and container walls are elastic collisions (the particles bounce off each other and the walls with no loss of energy)
3. Gas particles are in continuous, rapid, random motion. They therefore possess kinetic energy, which is energy of motion (gases move in random directions)

3. 4. There are no forces of attraction or repulsion between gas particles (gas particles not attracted to each other and bounce apart when they hit)
5. The average kinetic energy of gas particles depends on the temperature of the gas (energy of gases depends on their speed, speed will increase with increased temperature)

4. Kinetic energy: KE = ½ mv2
Ideal gas=imaginary gas that fits all 5 assumptions
Ideal gases do not really exist; real gases will behave almost like ideal gases at low pressure and high temperature

5. Physical Properties of gases
1. Expansion-gases can completely fill a container
2. Fluidity-gases can slide past one another easily, similar to how liquids flow. Both are called fluids.
3. Low density-particles are far apart so they are less dense than solid and liquid

6. 4. Compressibility-gases can be squeezed together in smaller volumes under pressure
5. Diffusion-gases spread out and mix with other gases because of their random motion
Effusion-process by which gases move through a tiny opening.

7. Real Gases
Real gases-does not behave according to the 5 assumptions of the kinetic-molecular theory
2 factors van der Waals proposed to explain why real gases are different than ideal gases:
Gases occupy space
Gases exert attractive force on each other

8. Conditions that real gases will act similar to ideal gases:
1. High temperature
Low pressure

9. Gases that act similar to ideal gases:
Nonpolar
Diatomic
Noble gases
Ex: H2 N2 He
Gases that act different from ideal gases:
Polar molecules

10. 10-2 Pressure
Pressure-the force per unit area on a surface, expressed in Newtons (N)
P = F A
Consider the ballerina: Which has more pressure- flat footed, 2 feet tippy toes, 1 foot tippy toed

11. Each kilogram exerts 9.8N of force due to gravity
If the ballerina has a mass of 51 kg, then she exerts a force of 500 N (51 X 9.8)

12. Atmospheric Pressure
The atmosphere is a blanket of gases surrounding Earth. These gases exert a pressure downward, atmospheric pressure.
The atmospheric pressure exerts a pressure on everything so why doesn’t it crush us?
Crushing can demo

13. Measuring Atmospheric Pressure
Torricelli discovered the first barometer
Barometer-measures atmospheric pressure

14. Units of Pressure mm Hg (millimeters of mercury) Torr
atm (atmospheres)
Pascal (Pa)
Kilopascals (kPa)

15. Atmospheric pressure =
760 mm Hg = 1 atm = X 105 Pa = kPa = 760 torr

16. Standard temperature and pressure
STP = 1 atm and 0ºC

17. Converting Units of Pressure
Sample Problem 10-1 pg. 312
The average atmospheric pressure in Denver, Colorado, is atm. Express this pressure in mm Hg and kPa.

18. 1. Convert a pressure of 1.75 atm to kPa to mm Hg.
2. Convert a pressure of 570 torr to atmospheres and to kPa.

19. 10-3 Gas Laws
Simple mathematical relationships between the volume, temperature, pressure, and amount in a gas.

20. Boyle’s Law: Pressure-Volume Relationship
Boyle discovered that doubling the pressure on a gas will reduce the volume by one-half.
If you reduce the volume of a container of gas, there are still the same number of gas particles that are colliding, increasing the pressure

21. Boyle’s Law-the volume of a fixed mass of gas varies inversely with the pressure at constant temperature
PV = k
Use Boyle’s Law to compare changing conditions for a gas
P1V1 = P2V2

22. Sample problem 10-2
A sample of oxygen gas has a volume of 150 mL when its pressure is atm. What will the volume of the gas be at a pressure of atm if the temperature remains constant?

23. A gas has a pressure of 1. 26 atm and occupies a volume of 7. 40 L
A gas has a pressure of 1.26 atm and occupies a volume of 7.40 L. If the gas is compressed to a volume of 2.93 L, what will its pressure be, assuming constant temperature?

24. Charles’s Law: Volume-Temperature Relationship
Jacques Charles discovered that as temperature of a gas increases, volume also increases
Charles’s law states that the volume of a fixed mass of gas as constant temperature varies directly with the Kelvin temperature.

25. Converting Temperature
Kelvin (K) = ºC
Charles’s Law:
V = k V1 = V2
T T T2

26. Sample problem
A sample of neon gas occupies a volume of 752 mL at 25ºC. What volume will the gas occupy at 50ºC if the pressure remains constant?

27. Charles’s Law practice
A gas at 65ºC occupies 4.22L. At what Celsius temperature will the volume be 3.87L, assuming the same pressure?

28. Gay-Lussac’s Law-Pressure/Temperature
Joseph Gay-Lussac discovered
Gay-Lussac’s Law-the pressure of a fixed mass of gas at constant volume varies directly with the Kelvin temperature
P = k P1 = P2
T T1 T2

29. Sample Problem
The gas in an aerosol can is at a pressure of 3.00 atm at 25ºC. Directions on the can warn the user not to keep the can in a place where the temperature exceeds 52ºC. What would the gas pressure in the can be at 52ºC?

30. Gay-Lussac’s Practice
A sample of helium gas has a pressure of 1.20 atm at 22ºC. At what Celsius temperature will the helium reach a pressure of 2.00 atm?

31. Combined Gas Law
Expresses the relationship between pressure, volume, and temperature of a fixed amount of gas.
PV = k P1V1 = P2V2
T T T2

32. Sample Problem
A helium-filled balloon has a volume of 50 L at 25ºC and 1.08 atm. What volume will it have at atm and 10ºC?

33. Combined Gas Law Practice
A 700 mL gas sample at STP is compressed to a volume of 200 mL, and the temperature is increased to 30ºC. What is the new pressure of the gas in Pa?

34. Dalton’s Law of Partial Pressures
The pressure of each gas in a mixture is called the partial pressure of that gas.
Dalton’s Law of Partial Pressures states that the total pressure of a mixture of gases is equal to the sum of the partial pressures of the component gases.

35. Dalton’s Partial Pressure Formula
PT = P1 + P2 + P3 +…
Patm = Pgas + PH2O

36. Dalton’s Law Sample
A sample of nitrogen gas is collected over water at a temperature of 23ºC. What is the pressure of the nitrogen gas if atmospheric pressure is 785 mm Hg?
Use Table A-8