2. The Gas Laws The density of a gas decreases as its temperature increases.

3. Chemical Properties Produce Gases Chemists harness chemical properties to produce a desired gas through chemical reactions. Such as the reaction of zinc and hydrochloric acid.Chemists harness chemical properties to produce a desired gas through chemical reactions. Such as the reaction of zinc and hydrochloric acid.

4. Physical Properties of Gases Gases are compressible and that they assume the shape and volume of any container. Gases are all infinitely soluble in one another. Each of these characteristics can be explained by the distances between the molecules (or atoms) in a gaseous sample.

5. Physical Properties of Gases are affected by temperature and pressure States of matter simulationStates of matter simulation.

6. Collisions of Gas Particles

7. Kinetic Molecular Theory  explains why gases behave as they do  deals with “ideal” gas particles…

8. Kinetic Theory

9. Kinetic Molecular Theory Postulates of the Kinetic Molecular Theory of Gases 1.Gases consist of tiny particles (atoms or molecules) 2.These particles are so small, compared with the distances between them, that the volume (size) of the individual particles can be assumed to be negligible (zero). 3. The particles are in constant random motion, colliding with the walls of the container. These collisions with the walls cause the pressure exerted by the gas. 4. The particles are assumed not to attract or to repel each other. 5. The average kinetic energy of the gas particles is directly proportional to the Kelvin temperature of the gas.

10. Kinetic Molecular Theory Postulates Evidence 1. Gases are tiny molecules in mostly empty space. The compressibility of gases. 2. There are no attractive forces between molecules. Gases do not clump. 3. The molecules move in constant, rapid, random, straight-line motion. Gases mix rapidly. 4. The molecules collide classically with container walls and one another. Gases exert pressure that does not diminish over time. 5. The average kinetic energy of the molecules is proportional to the Kelvin temperature of the sample. Charles’ Law

11. Newton’s First Law of Motion (Law of Inertia) Object at rest tends to stay at rest, and object in motion tends to stay in motion at constant velocity unless object is acted upon by an unbalanced, external force.

12. Elastic vs. Inelastic Collisions 8 3

13. 8 8 v1v1 elastic collision inelastic collision POW v2v2 v3v3 v4v4

14. 8 Elastic Collision 8 v1v1 before v2v2 after

15. Model Gas Behavior All collisions must be elasticAll collisions must be elastic Take one step per beat of the metronomeTake one step per beat of the metronome ContainerContainer –Class stands outside tape box Higher temperatureHigher temperature –Faster beats of metronome Decreased volumeDecreased volume –Divide box in half More MolesMore Moles –More students are inside box  Mark area of container with tape on ground.  Add only a few molecules of inert gas  Increase temperature  Decrease volume  Add more gas  Effect of diffusion  Effect of effusion (opening size)

16. Kinetic Molecular Theory Particles in an ideal gas…Particles in an ideal gas… –have no volume. –have elastic collisions. –are in constant, random, straight-line motion. –don’t attract or repel each other. –have an avg. KE directly related to Kelvin temperature. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

17. Real Gases Particles in a REAL gas…Particles in a REAL gas… –have their own volume –attract each other Gas behavior is most ideal…Gas behavior is most ideal… –at low pressures –at high temperatures –in nonpolar atoms/molecules Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

18. Characteristics of Gases Gases expand to fill any container. –random motion, no attraction Gases are fluids (like liquids). –no attraction Gases have very low densities. –no volume = lots of empty space Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

19. Characteristics of Gases Gases can be compressed. –n–n–n–no volume = lots of empty space Gases undergo diffusion & effusion. –r–r–r–random motion Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

20. Pressure Pressure is defined as force divided by the area.Pressure is defined as force divided by the area.

21. Pressure The mercury in the inverted tube is pushed upward by the force of atmospheric pressure pushing down on the surface of the mercury in the dish. The height of the mercury in the tube changes with changing atmospheric pressure. Under conditions of standard atmospheric pressure, the height of the mercury in the tube is 760 mm. (1 atm = 760 mm Hg = 760 torr = 1.01325  kPa)The mercury in the inverted tube is pushed upward by the force of atmospheric pressure pushing down on the surface of the mercury in the dish. The height of the mercury in the tube changes with changing atmospheric pressure. Under conditions of standard atmospheric pressure, the height of the mercury in the tube is 760 mm. (1 atm = 760 mm Hg = 760 torr = 1.01325  kPa)

22. Collisions cause Pressure The pressure of a gas is caused by the collision of molecules against the sides of the container. The force of the collision against the container can be calculated by Newton’s Second Law of Motion: F=ma. The “F” = force, “m”=mass in kg and “a” is the acceleration in m/s 2.

23. Low Pressure vs. High Pressure inside a System The number of collisions of gas molecules against the wall of the container determines the pressure in the container. Notice the difference in the number of collisions. Figure (a) would have a lower pressure than Figure (b).

24. Pressure Is caused by the collisions of molecules with the walls of a container is equal to force/unit area SI units = Newton/meter 2 = 1 Pascal (Pa) 1 standard atmosphere = 101,325 Pa 1 standard atmosphere = 1 atm = 760 mm Hg = 760 torr

25. Pressure KEY UNITS AT SEA LEVEL 101.325 kPa (kilopascal) 1 atm 760 mm Hg 760 torr 14.7 psi 1 bar = 100 kPa Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Sea level

26. The first device for measuring atmospheric pressure was developed by Evangelista Torricelli during the 17 th century. The device was called a “barometer” Baro = weight Meter = measure Evangelista Torricelli, circa 1644 Measuring Pressure “We live submerged at the bottom of an ocean of air.”

27. Barometer Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 401 Empty space (a vacuum) Hg Weight of the mercury in the column Weight of the atmosphere (atmospheric pressure)

28. Barometer Mercury filledMercury filled 760 mm = 1 atm Water filledWater filled 10400 mm = 1 atm The barometer measures air pressure Water column (34.0 ft. high or 10.4 m) Atmospheric pressure Mercury column (30.0 in. high or 76 cm )

29. Barometers Mount Everest Sea level On top of Mount Everest Sea level fraction of 1 atm average altitude (m)(ft) 100 1/25,48618,000 1/38,37627,480 1/1016,13252,926 1/10030,901101,381 1/100048,467159,013 1/1000069,464227,899 1/10000096,282283,076

30. Pressure Practice Convert the following: 1.145 mm Hg into bars 2.450 psi into kPa 3.900 mm Hg into torrs 4.4580 Pa into kPa 5. 25 psi into atm 6. 150 atm into Pa 7.109 kPa into atm 8.76.9 mm Hg into bars 9.98.6 torr into kPa 10. 3 atm into kPa

31. Answers 1)0.19 bars 2)3102.4 kPa 3)900 torr 4)4.58 kPa 5)1.7 at5m 6)15199108.32 Pa 7)0.10 bars 8)13.14 kPa 9)303.98 kPa 10)1.08 atm

32. Temperature ºF ºC K -45932212 -2730100 0273373 K = ºC + 273 Always use absolute temperature (Kelvin) when working with gases. Always use absolute temperature (Kelvin) when working with gases. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

33. STP 0°C 0°C 1 atm 1 atm - OR - STP Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Standard Temperature & Pressure 273 K 101.325 kPa 760 mm Hg