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Gases. I. Physical Properties. Real vs. Ideal Gases:. Ideal gas = an imaginary gas that conforms perfectly to all the assumptions of the kinetic-molecular theory We will assume that the gases used for the gas law problems are ideal gases . Real Gases vs. Ideal Gases:.

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I. Physical Properties


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    1. Gases I. Physical Properties

    2. Real vs. Ideal Gases: • Ideal gas = an imaginary gas that conforms perfectly to all the assumptions of the kinetic-molecular theory • We will assume that the gases used for the gas law problems are ideal gases.

    3. Real Gases vs. Ideal Gases: • Real gas = a gas that does not behave completely according to the assumptions of the kinetic-molecular theory. •  All real gases deviate to some degree from ideal gas behavior. However, most real gases behave nearly ideally when their particles are sufficiently far apart and have sufficiently high kinetic energy.

    4. Causes of non-ideal behavior: • The kinetic-molecular theory is more likely to hold true for gases whose particles have NO attraction for each other.

    5. Causes of non-ideal behavior: 1. High pressure (low volume): • Space taken up by gas particles becomes significant •  Intermolecular forces are more significant between gas particles that are closer together

    6. Causes of non-ideal behavior: 2. Low temperature: • Gas particles move slower so intermolecular forces become more important

    7. KMT (Kinetic Molecular Theory) • Kinetic Molecular Theory (KMT) = the idea that particles of matter are always in motion and that this motion has consequences. • theory developed in the late 19th century to account for the behavior of the atoms and molecules that make up matter

    8. KMT (Kinetic Molecular Theory) • based on the idea that particles in all forms of matter are always in motion and that this motion has consequences LIQUID GAS SOLID

    9. KMT (Kinetic Molecular Theory) • can be used to explain the properties of solids, liquids, and gases in terms of the energy of particles and the forces that act between them

    10. B. Kinetic Molecular Theory – Ideal gas • KMT describing particles in an IDEAL gas: • have no volume. • have elastic collisions. • are in constant, random motion. • don’t attract or repel each other. • have an avg. KE directly related to Kelvin temperature.

    11. C. Kinetic Molecular Theory - Real Gases • KMT describing particles in a REAL gas: • have their own volume • attract each other • Gas behavior is most ideal… • at low pressures • at high temperatures • in nonpolar atoms/molecules

    12. D. Characteristics of Gases- using KMT • Gases expand to fill any container. • KMT - gas particles move rapidly in all directions without significant attraction or repulsion between particles

    13. D. Characteristics of Gases- using KMT • Gases are fluids (like liquids). • KMT - No significant attraction or repulsion between gas particles; glide past each other

    14. D. Characteristics of Gases- using KMT • Gases have very low densities. • KMT - particles are so much farther apart in the gas state

    15. D. Characteristics of Gases- using KMT • Gases can be compressed. • KMT - gas particles are far apart from one anther with room to be “squished” together

    16. D. Characteristics of Gases- using KMT • Gases undergo diffusion & effusion. • KMT – gas particles move in continuous, rapid, random motion Effusion

    17. E. Temperature K = ºC + 273 ºF -459 32 212 ºC -273 0 100 K 0 273 373 • Always use Kelvin temperature when working with gases.

    18. F. Pressure Which shoes create the most pressure?

    19. F. Pressure • Why do gases exert pressure? • Gas particles exert a pressure on any surface with which they collide! • More collisions = increase in pressure!

    20. F. Pressure Mercury Barometer • Barometer= measures atmospheric pressure • The height of the Hg in the tube depends on the pressure • The pressure of the atmosphere is proportional to the height of the Hg column, so the height of the Hg can be used to measure atmospheric pressure!

    21. F. Pressure • Pressure UNITS 101.3 kPa (kilopascal) 1 atm 760 mm Hg 760 torr

    22. G. STP STP Standard Temperature & Pressure 0°C (exact) 1 atm (exact) 273 K101.3 kPa 760 mm Hg (exact) 760 torr (exact)

    23. V T P Gases II. The Gas Laws

    24. GAS LAW PROBLEMS- MUST USE KELVIN K = ºC + 273

    25. A. Boyle’s Law P V PV = k

    26. A. Boyle’s Law P V • The pressure and volume of a gas are INVERSELY related • at constant mass & temp P1V1 = P2V2

    27. A. Boyle’s Law • Real life application:When you breathe, your diaphragm moves downward, increasing the volume of the lungs. This causes the pressure inside the lungs to be less than the outside pressure so air rushes in.

    28. A. Boyle’s Law • Ex: Halving the volume leads to twice the rate of collisions and a doubling of the pressure.

    29. A. Boyle’s Law • As the volume increases, the pressure decreases!

    30. B. Charles’ Law V T

    31. B. Charles’ Law V T • The volume and temperature (K) of a gas are DIRECTLY related • at constant mass & pressure V1_= V2__ T1 T2

    32. B. Charles’ Law • Real life application: Bread dough rises because yeast produces carbon dioxide. When placed in the oven, the heat causes the gas to expand, and the bread rises even further.

    33. B. Charles’ Law • As temperature decreases, the volume of the gas decreases • Liquid nitrogen’s temp. is about 63K or -210 ºC or -346 ºF!

    34. B. Charles’ Law • As temperature decreases, the volume of the gas decreases

    35. B. Charles’ Law • NUMBER OF PARTICLES & PRESSURE ARE CONSTANT!!!

    36. C. Guy-Lussac’sLaw P T

    37. C. Guy-Lussac’sLaw P T • The pressure and temperature (K) of a gas are DIRECTLY RELATED • at constant mass & volume _P1_= P2__ T1 T2

    38. C. Guy-Lussac’s Law • When the temp. of a gas increases (KE increases) and gas particles move faster and hit container walls more frequently and collisions are more forceful

    39. C. Guy-Lussac’s Law • Real life application: The air pressure inside a tire increases on a hot summer day.

    40. D. Combined Gas Law P1V1 T1 P2V2 T2 = V Charles’T P Guy Lussac’sT PV T BoylesPV = k

    41. E. EXAMPLE Problems • You need your calculators!

    42. Gas Law Problems 1.) A gas occupies 473 cm3 at 36°C. Find its volume at 94°C. CHARLES’ LAW V1_= V2__ T1 T2 WORK: GIVEN: V1= 473 cm3 T1 = 36°C = 309K V2 = ? T2 = 94°C = 367K V T V2 =V1T2 T1 V2= (473 cm3)(367 K) (309 K) V2 = 562 cm3 C. Johannesson

    43. Gas Law Problems 2.) A gas occupies 100. mL at 150. kPa. Find its volume at 200. kPa. BOYLE’S LAW P1V1 = P2V2 P WORK: V2 = P1V1 P2 V GIVEN: V1= 100. mL P1 = 150. kPa V2 = ? P2 = 200. kPa V2 = (150.kPa)(100.mL) 200.kPa V2 = 75.0 mL C. Johannesson

    44. Gas Law Problems 3.) A gas occupies 7.84 cm3 at 71.8 kPa & 25°C. Find its volume at STP. COMBINED GAS LAW P1V1 = __P2V2T1T2 P T V GIVEN: V1=7.84 cm3 P1=71.8 kPa T1=25°C = 298 K V2=? P2=101.3 kPa T2=273 K WORK: V2 = P1V1T2 P2T1 V2 = (71.8 kPa)(7.84 cm3)(273 K) (101.3 kPa)(298 K) V2= 5.09 cm3 C. Johannesson

    45. Gas Law Problems 4.) A gas’ pressure is 765 torr at 23°C. At what temperature will the pressure be 560. torr? GUY LUSSAC’S LAW P1_= P2__ T1 T2 GIVEN: P1= 765 torr T1 = 23°C = 296K P2 = 560. torr T2 = ? WORK: T2 = P2T1 P1 T P T2 = (560. torr)(296K) 765 torr T2 = 217 K C. Johannesson

    46. Gases III. Standard Molar Volume, Gas Densities & Molar Mass

    47. A. Standard Molar Volume • The volume occupied by one mole of a gas at STP • 22.41410 L/mol or about 22.4 L/mol • In other words, one mole of any ideal gas at STP will occupy 22.4 L

    48. B. Example Problems-standard molar volume 1.) A chemical reaction is expected to produce 0.0680 mol of oxygen gas. What volume of gas in L will be occupied by this gas sample at STP? • 0.068 mol O2 22.4 L O2 • 1 mol O2 = 1.52 L O2

    49. B. Example Problems-standard molar volume 1.) A chemical reaction is expected to produce 0.0680 mol of oxygen gas. What volume of gas in L will be occupied by this gas sample at STP? • 0.068 mol O2 22.4 L O2 • 1 mol O2 = 1.52 L O2

    50. B. Example Problems-standard molar volume 2.) A chemical reaction produced 98.0 mL of SO2 at STP. What mass (in grams) of the gas was produced? 64.064 g SO2 • 98.0 mL SO2 1 LSO21 mol O2 = 1000 mL SO2 22.4 L SO2 1 mol SO2 0.280 g SO2