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1. The Behavior of Gases Chapter 14

2. Section 1 Properties of Gases

3. Compressibility • Compressibility – a measure of how much the volume of matter decreases under pressure. • Gases are easily compressed because of the space between the particles.

4. Factors Affecting Gas Pressure • The amount of gas, volume, and temperature are factors that affect gas pressure. • Pressure (P) in kilopascals (kPa) • Volume (V) in liters (L) • Temperature (T) in Kelvin (K) • Number of moles (n) in mole (mol)

5. Amount of Gas • You can use kinetic theory to predict and explain how gases will respond to a change of conditions. • As you add more gas particles the pressure increases.

6. Amount of Gas • Once the pressure exceeds the strength of the container the container will burst.

7. Aerosol cans depend on the movement of gas from a region of high pressure to a region of low pressure. • Pushing the spray button creates an opening between the inside of the can and the outside.

8. The gas flows through the opening to the lower pressure region outside. • The movement of the gas propels the paint out of the can until the gas can no longer propel paint out.

9. Volume • You can raise the pressure exerted by a contained gas by reducing its volume. • The more a gas is compressed the greater the pressure.

10. Temperature • As a gas is heated, the temperature increases and the average kinetic energy also increases.

11. When the volume of a container is held constant and the temperature increases and the pressure increases.

12. Section 2 The Gas Laws

13. Boyle’s Law: Pressure and Volume • If the temperature is constant, as the pressure of a gas increases, the volume decreases.

14. Boyle’s law– states that for a given mass of gas at a constant temperature, the volume of the gas varies inversely with pressure.

15. Charles’s Law: Temperature and Volume • As the temperature of an enclosed gas increases, the volume increases, if the pressure is constant.

16. Charles’s law– states that the volume of a fixed mass of gas is directly proportional to its Kelvin temperature if the pressure is kept constant.

17. How can you tell from the picture that there is a fixed amount of gas in the cylinder? Describe what is happening in the cylinder as it’s being heated.

18. Gay-Lussac’s Law: Pressure and Temperature • As the temperature of an enclosed gas increases, the pressure increases, if the volume is constant.

19. Gay-Lussac’s law– states that the pressure of a gas is directly proportional to the Kelvin temperature if the volume remains constant.

20. P V Boyle’s Law Charles’s Law Gay-Lussac’s Law T

21. The Combined Gas Law • Combined gas law – describes the relationships among the pressure, temperature, and volume of an enclosed gas.

22. The combined gas law allows you to do calculations for situations in which only the amount of gas is constant.

23. Boyle’s Law P V T Charles’s Law Gay-Lussac’s Law

24. Section 3 Ideal Gases

25. Ideal Gas Law • The combined gas law is good when the amount of gas does not change – this does not always stay constant though.

26. To calculate the number of moles of a contained gas requires an expression that contains the variable n. • The number of moles is directly proportional to the number of particles and can be introduced into the combined gas law by dividing each side by n.

27. Ideal gas constant – (R) has the value of 8.31 (L•kPa)/(K•mol). • Ideal gas law – includes the variables of P, V, T, and n. • P is the pressure (units of kPa) • V is the volume (units of L) • T is the temperature (units of K) • n is the number of moles (units of mol)

28. PV=nRT Song

29. Ideal Gases and Real Gases • An ideal gas is one that follows the gas laws under all conditions of temperature and pressure. • Real gases differ most from an ideal gas at low temperatures and high pressures.

30. Section 4 Gases: Mixtures and Movements

31. Dalton’s Law • Partial pressure – the contribution of each gas in a mixture makes to the total pressure. • In a mixture of gases, the total pressure is the sum of the partial pressures of the gases.

32. Dalton’s law of partial pressures – states that at constant volume and temperature, the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of the component gases.

33. Example: • Determine the total pressure of a gas mixture that contains oxygen, nitrogen, and helium. The partial pressures are: PO2 = 20kpa, PN2 = 46.7kPa; and PHe = 26.7kPa.

34. Graham’s Law • Diffusion – tendency of molecules to move toward areas of lower concentration until the concentration is uniform throughout.

35. Bromine gas is put in a cylinder and after several hours you can see how the gas has diffused.

36. Effusion – a gas escapes through a tiny hole in its container. • Gases of lower molar mass diffuse and effuse faster than gases of higher molar mass.

37. Thomas Graham’s Contribution • Scottish chemist Thomas Graham studied rates of effusion in the 1840’s • Relates to KE = ½ mv2. • Kinetic energy of the particles (KE) is related to the mass (m) and their velocity (v).

38. Graham’s law of effusion – states that the rate of effusion of a gas is inversely proportional to the square root of the gas’s molar mass.

39. Example: • Determine the rate of effusion for helium compared to nitrogen. • This result tells me that the helium effuses/diffuses faster than the nitrogen at the same temperature.