Chapter 14 “The Behavior of Gases”

1. Chapter 14“The Behavior of Gases” Pre-AP Chemistry Charles Page High School Stephen L. Cotton

2. Section 14.1The Properties of Gases • OBJECTIVES: • Explain why gases are easier to compress than solids or liquids are. • OBJECTIVES: • Describe the three factors that affect gas pressure.

3. Compressibility • Gases can expand to fill its container, unlike solids or liquids • The reverse is also true: • They are easily compressed, or squeezed into a smaller volume • Compressibility is a measure of how much the volume of matter decreases under pressure

4. Variables that describe a Gas • The four variables and their common units: 1. pressure (P) in kilopascals 2. volume (V) in Liters 3. temperature (T) in Kelvin 4. amount (n) in moles • The amount of gas, volume, andtemperature are factors that affect gas pressure.

5. 1. Amount of Gas • When we inflate a balloon, we are adding gas molecules. • More molecules means more collisions (higher pressure) • Fewer molecules means fewer collisions (lower pressure)

6. Common use? • A practical application is Aerosol (spray) cans • gas moves from higher pressure to lower pressure • a propellant forces the product out • whipped cream, hair spray, paint

7. 2. Volume of Gas • In a smaller container, the molecules have less room to move. • The particles hit the sides of the container more often. • As volume decreases, pressure increases. (think of a syringe) • Thus, volume and pressure are inversely related to each other

8. 3. Temperature of Gas • Raising the temperature of a gas increases the pressure, if the volume is held constant. (Temp. and Pres. are directly related) • The higher the temperature, the more energy the molecules have. • The molecules hit the walls harder, and more frequently!

10. Section 14.2The Gas Laws • OBJECTIVES: • Describe the relationships among the temperature, pressure, and volume of a gas. • Use the combined gas law to solve problems.

11. #1 Boyle’s Law - 1662 Gas pressure is inversely proportional to the volume, when temperature is held constant. • Equation: P1V1 = P2V2

12. Graph of Boyle’s Law – page 418 Boyle’s Law says the pressure is inverse to the volume. Note that when the volume goes up, the pressure goes down

13. Example 1 A balloon contains 30.0 L of helium gas at 103 kPa. What is the volume of the helium when the pressure drops to 25.0 kPa?

14. #2 Charles’s Law - 1787 • The volume of a fixed mass of gas is directly proportional to the Kelvin temperature, when pressure is held constant.

15. Converting Celsius to Kelvin • Gas law problems involving temperature will always require that the temperature be in Kelvin. (Remember that no degree sign is shown with the kelvin scale.) Kelvin = C + 273 °C = Kelvin - 273 and

16. Example 2 A balloon inflated in a room at 24°C has a volume of 4.00 L. The balloon is then heated to a temperature of 58 °C. What is the new volume?

17. #3 Gay-Lussac’s Law - 1802 • The pressure and Kelvin temperature of a gas are directly proportional, provided that the volume remains constant. • How does a pressure cooker affect the time needed to cook food? (Note page 422)

18. Example 3 The gas in a used aerosol can is at a pressure of 103 kPa at 25°C. If the can is thrown onto a fire, what will the pressure be when the temperature reaches 928°C?

19. #4 The Combined Gas Law The combined gas law expresses the relationship between pressure, volume and temperature of a fixed amount of gas.

20. The combined gas law contains all the other gas laws! • If the temperature remains constant... P1 V1 P2 x V2 x = T1 T2 Boyle’s Law

21. The combined gas law contains all the other gas laws! • If the pressure remains constant... P1 V1 P2 x V2 x = T1 T2 Charles’s Law

22. The combined gas law contains all the other gas laws! • If the volume remains constant... P1 V1 P2 x V2 x = T1 T2 Gay-Lussac’s Law

23. Example 4 The volume of a gas-filled balloon is 30.0 L at 313 K and 153 kPa of pressure. What would the volume be at 273 K and 101.3 kPa of pressure?

24. Section 14.3Ideal Gases • OBJECTIVES: • Compute the value of an unknown using the ideal gas law. • Compare and contrast real an ideal gases.

25. #5 The Ideal Gas Law • Equation: PV = nRT • Pressure times Volume equals the number of moles (n) times the Ideal Gas Constant (R) times the Temperature in Kelvin.

26. The Ideal Gas Law • We now have a new way to count moles (the amount of matter), by measuring T, P, and V. We aren’t restricted to only STP conditions: P x V R x T n =

27. Ideal Gases • We are going to assume the gases behave “ideally”- in other words, they obey the Gas Laws under all conditions of temperature and pressure • An ideal gas does not really exist, but it makes the math easier and is a close approximation.

28. Ideal Gases don’t exist, because: • Molecules do take up space • There are attractive forces between particles - otherwise there would be no liquids formed

29. Section 14.4Gases: Mixtures and Movements • OBJECTIVES: • Relate the total pressure of a mixture of gases to the partial pressures of the component gases. • Explain how the molar mass of a gas affects the rate at which the gas diffuses and effuses.

30. #6 Dalton’s Law of Partial Pressures For a mixture of gases in a container, PTotal = P1 + P2 + P3 + . . . • P1 represents the “partial pressure”, or the contribution by that gas. • Dalton’s Law is particularly useful in calculating the pressure of gases collected over water.

31. If the first three containers are all put into the fourth, we can find the pressure in that container by adding up the pressure in the first 3: 2 atm + 1 atm = 6 atm + 3 atm 4 3 2 1

32. Diffusion is: • Molecules moving from areas of high concentration to low concentration. • Example: perfume molecules spreading across the room. • Effusion: Gas escaping through a tiny hole in a container. • Both of these depend on the molar mass of the particle, which determines the speed.

33. Diffusion:describes the mixing of gases. The rate of diffusion is the rate of gas mixing. • Molecules move from areas of high concentration to low concentration.

34. Effusion: a gas escapes through a tiny hole in its container -Think of a nail in your car tire… Diffusion and effusion are explained by Graham’s Law.

35. End of Chapter 14