Gases Notes

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 volumeof a gas. Example: air in a bicycle tire

3. 3. Unlike liquids, gases completelyfill 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 collidewith the walls of their container and each other.

11. Inelastic Collision 5. Collisions of gas particles are perfectly elastic. 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. 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: 0oC = 273K Taken in oC converted to Kelvin (K) T(K) = °C + 273 Ex #1) 22°C = 295 K 100.°C = 373 K -27.3°C = 245.7 K -273°C = 0 K

16. P- Standard Pressure: can be measured using the following units: (sea level pressure) 1 atm (atmospheres) 760mm Hg 101,325 Pa (Pascals) 760torr 101.325kPa (kilopascals) 14.7lb/in2 (psi) Ex #2) Convert 1.026 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. 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 VolumeQuadrupling Pressure = ¼ Volume

26. Boyle’s Law

27. 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. At constant pressurethe volume of a fixed amount of gas is directly proportional to its absolute temperature. Direct *Temperatures must be in Kelvin! K = °C + 273

30. Ballon in cool and cold water:

32. Ex #5) What is the Celsius temperature of 68.20 mL of methane, if it occupies a volume of 0.02200 L at 50.0°C? 22.00 mL 22.00 mL

34. C. Avogadro’s Law: Amount - Volume Relationship. Equal volumesof 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 222 Rn molar mass

35. 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. He O2 Rn

36. Molar Mass does not affect volume of a gas

37. D. Dalton’s Law of Partial Pressures: The sum of the partial pressuresof all the components in a gas mixture is equal to the total pressureof the gas mixture. Ex #6) A flask contains a mixture of oxygen, argon, and carbon dioxide with partial pressures of 745 torr, 0.278 atm, and 391 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. 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. Convert these to standard conditions. What is the new volume? P1 = 198 kPa P2 = 101.325 kPa V1 = 72.0 mL V2 = ? T1 = 298 K T2 = 273 K (273 K) (101.325 kPa)

44. 12.3 A. Ideal Gas Law Although no “ideal gas” exists, this law can be used to explain the behaviorof real gases under ordinary conditions. P = pressure (atm) V = volume (L or dm3) n = number of moles R = 0.0821 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 temperaturesand high pressuresdo NOT behave ideally. As you decrease the volume of a gas, the volumeof the particles themselves becomes significant. The Kinectic-Molecular Theory & Ideal Gas Law assumes the gas particles have novolume of their own. 3. Second, the attractiveforces which are very small when the particles are moving fast, become larger as they slowdown.