CH110 Chapter 6: Gases

1. CH110 Chapter 6: Gases Kinetic Molecular Theory Pressure Gas Laws

2. Density Shape Compressibility Fast moving, Low density, Expands to fill container Vapor Large compressibility, Expands w/ heat Moderate movement, Dense, Takes shape of container Small compressibility, Small heat expansion Liquid Small compressibility, Very small heat expansion Slow moving, dense, Fixed shape Solid

3. Kinetic molecular theory of Gases Model to explain behavior of gases Vapor • 1. All gases are made up of tiny particlesmoving in • straight lines • in all directions • at various speeds.

4. Kinetic molecular theory 2. Particles far apart have no effect oneach other. (Don’t attract or repel) • 3. V of a gas = • V of container • V of a gas is mostly empty space.

5. KE T Kinetic molecular theory 4. Theave KE as theT (K.E. a T) • The average KE is the same for all gases atthe same T.

6. Kinetic molecular theory 5. Gas molecules exert pressure as they collide with container walls The > the # of collisions (per unit time), the > the pressure

7. Force Area Pressure = Force per unit of Area. P = Force Area In the atmosphere, molecules of air (N2, O2, Ar, H2O, etc..) are constantly bouncing off us.

8. Atmospheric Pressure We live at the bottom of an ocean of air Atmosphere: A sea of colorless, odorless gases surrounding the earth

9. Pressure At higher elevations, there is less air so the P is less.

10. Boiling Point = Temp where molecules overcome atmospheric Pressure H2O 270 torr Mt. Everest(20,000’) = 73 oC 467 torr Mt. Evans,CO(14,000’) = 87 oC 630 torr Denver (5280’) = 95 oC 760 torr Sea Level = 100 oC

11. Measuring Pressure Attempts to pump water out of flooded mines often failed because H2O can’t be lifted more than 34 feet.

12. Measuring Pressure Torricelli believed reason was that P of atmosphere could not hold anything heavier than a 34’ column of water.

13. Like drinking from a straw. Atmospheric Pressure What causes the liquid to move up the straw to your mouth ?

14. Measuring Pressure The atmosphere would support a column of H2O > 34 feet high. 1 Atm 34’ column of water

15. Torricelli Barometer Pressure of the atmosphere supports a column of Hg 760 mm high. 1 atm = 760 mm Hg 760 torr 29.92 in Hg 14.7 lb/in2 101,325 Pa vacuum 1 atm Mercury used because it’s so dense.

16. Blood pressure (systolic over diastolic): most often in mm Hg. (ex. 120/80) Meteorologists refer to pressure systems in mm or inches of Hg. ex. 30.01 in

17. STP Standard Temperature & Pressure 1 atm = 760 mm Hg 760 torr 29.92 in Hg 14.7 lb/in2 101,325 Pa 0oC 273K 1 atm

18. Gas laws Laws that show relationships between volume and properties of gases Boyle’s Law Charles’ Law Gay-Lussac’s Law Combined Gas Law Avogadro’s Law Dalton’s Law

19. P V P V P V Boyle’s law: V vs P V is inversely proportional to P when T is constant. If P goes down V goes up

20. Boyle’s law: V vs P P1V1= P2V2 2 L V2 = P2 = 0.5 Atm P1V1 = V2 P2 1atm (1L) = 0.5 atm 2 L P1 = 1 Atm 1 L V1 =

21. Boyle’s law: V vs P 2 L • Drive to top of mountain - ears start popping. • Breathing at high altitudes is more difficult because the pressure of O2 is less. 1 L

22. Boyle’s law It all “Boyle’s” down to Breathing in and out.

23. P T V Charles’s law: V vs T The volume of a gas is directly proportional to the absolute temperature (K). If T goes up V goes up

24. Charles’s law: V vs T V1= V2 T1 T2 T2V1= V2 T1 V1 = 125 mL (546K)125 mL = 273 K 250 mL V2= T2 = 546 K T1 = 273 K

25. V P T Gay-Lussac’s Law (PT) Pressure of a gas is directly proportional to Absolute Temp (K) when Volume is constant P1= P2 T1 T2 If P goes up T goes up

26. V T P P T V T P V Gas Laws P1V1 = P2V2 V1= V2 T1 T2 P1= P2 T1 T2 Boyle’s Charles’ Gay-Lussac’s

27. V T P P T V P1V1 T1 P2V2 T2 = T P V Gas Laws Boyle’s Combined Gas Law Charles’ Gay-Lussac’s

28. P1V1 T1 P2V2 T2 = T2P1V1 P2 T1 = V2 (240 K)(740 mm)(10 m3 ) (370 mm) (300 K) V2 = Combined Gas Law A 10 m3 balloon contains helium on the ground where the temperature is 27ºC and the pressure is 740 torr. Find the volume at an altitude of 5300 m if pressure is 370 mm Hg and temperature is -33 ºC. P1 = 740 mm P2 = 370 mm T1 = 27 + 273 = 300 K T2 = -33 + 273 = 240 K V1 = 10 m3 V2 = ? = 16 m3

29. Avogadro’s law The volume of a gas is directly proportional to the number of molecules V1= V2 n1 n2 More moles of a gas, takes up more space.

30. 1 mol He 4 g He 22.4 L 1 mol N2 28 g N2 22.4 L 1 mol CO2 44 g CO2 22.4 L 1 mol He 4 g He 22.4 L 1 mol N2 28 g N2 22.4 L 1 mol CO2 44 g CO2 22.4 L Avogadro’s law At Standard Temperature & Pressure (STP) V of 1 mole of gas = 22.4 liters At T = 273 K (0ºC) P = 1 atm (760 mm) Equal volumes of gas (at same T and P) contain equal numbers of molecules.

31. STP Example: What volume will 66 grams of CO2 occupy at STP? 66 g CO2 1 mole CO2 44 g CO2 22.4 liters 1 mole CO2 = 33.6 L

32. Dalton’s law of Partial Pressures The total pressure of a gas mix = sum of the partial pressures of each gas. • PT = P1 + P2 + P3 + ..... Each gas acts independently of the others. Example: Air Pair = PN2 + PO2 + PAr + PCO2 + PH2O

33. Dalton’s law of Partial Pressures Pair = PN2 + PO2 + PCO2 + PH2O = 760 mm Typical values for Atmospheric air at 0 ºC (excluding argon): (594.0mm) +(160mm) +(0.3mm) +(5.7mm)= 760mm As T of air increases, more H2O enters mix. example: at 20 ºC, the PH2O = 18 mm Ptotal (760 mm) can’t change, so other gases get diluted to make room for the water.

34. Pair = PN2 + PO2 + PCO2 + PH2O = 760 mm Air moving over warm water has more water in it. Low pressure is often associated with this air. Typhoons and hurricanes are associated with very warm, moist air.

35. Blood Gases PBG = PO2 + PCO2 Normal PO2in the air =160 mm. If drops < 100 mm, can’t diffuse into the blood. Arterial Blood Gases (ABGs) PCO2 ~ 40 mm Hg PO2 ~ 100 mm Hg Venous Blood Gases (VBGs) PO2 ~ 40 mm Hg PCO2 ~ 46 mm Hg

36. We only use about 25% of the Oxygen we inhale. The rest is exhaled along with the Nitrogen and some carbon dioxide. THIS IS WHY CPR WORKS !!!

37. T P Sol HENRY’S LAW The solubility of a gas in a liquid is directly related to the pressure on the liquid. Gas solubility goes up (more gas will dissolve) If P goes up

38. T P Sol HENRY’S LAW If P goes down Gas solubility goes down (gases escape) Example: opening a soda. Soda under high pressure Soda under low pressure

39. The “Bends” Lower P Less dissolved gases Quick ascent Get bubbles in blood & joints  extreme pain High P Lots of dissolved N2

40. The “Bends” Lower P N2accumulates in brain, spinal cord, and peripheral nerves. Bubbles here can cause paralysis and convulsions. Effects often irreversible. Less dissolved gases High P Lots of dissolved gases

41. Nitrogen Narcosis • “Nitrogen Narcosis”, • = nitrogen euphoria or raptures of the deep. • (Effect somewhat like that observed • when alcohol levels rise in the blood.) So, Helium often substituted for N2 in divers air.

42. T S T S T S Temperature vs Solubility The solubility of a gas in a liquid is inversely related to the temperature . If T goes up Gas solubility goes down (gases escape) Gas Solubility

43. Temperature vs Solubility Cold H2O holds more gas than warm H2O If hot rivers lose too much dissolved O2 the fish can’t survive.

44. Temperature vs Solubility Carbonated beveragesbottled cold. Divers with bends often packed in ice for transport to hyperbaric chamber.

45. T Solubility P Solubility P T Gas Laws Henry’s T vs Sol

46. Bernoulli's Principle Faster moving gases exert less pressure than slow moving gases. Fast moving Gases Low P Slow moving Gases High P

47. Bernoulli's Principle Fast moving Gases Slow moving Gases Low P High P