3. 10.1 C HARACTERISTICS OF G ASES  Air behaves physically as one gaseous material N 2 (78%), O 2 (21%) and Ar (0.9%)  Only a few elements exist as gases under standard conditions H 2, N 2, O 2, F 2, and Cl 2, the noble gases (He, Ne, Ar, Kr, Xe)  Gas molecules are relatively far apart Each molecule behaves largely as though the others are not present Readily compressible and expansible Forms homogeneous mixtures with other gases Low density

4. 10.1 C HARACTERISTICS OF G ASES

5. 10.2 P RESSURE  Pressure is defined as:  A TMOSPHERIC P RESSURE AND THE B AROMETER

6. 10.2 P RESSURE  In the 17 th century, people believed that the atmosphere had no weight  Torricelli’s experiment Proved the atmosphere has weight  Pascal’s experiment Measured the height of the mercury column at two different places Supported Torricelli’s explanation  Standard atmospheric pressure  A TMOSPHERIC P RESSURE AND THE B AROMETER Figure 10.2 A mercury barometer invented by Torricelli

7.  Manometer This device is used to measure the difference in pressure between atmospheric pressure and that of a gas in a vessel.  A TMOSPHERIC P RESSURE AND THE B AROMETER 10.2 P RESSURE

8. 10.3 T HE G AS L AW  Hypertension is abnormally high blood pressure. The usual criterion is a blood pressure greater than 140/90. mercury manometer or related device closed, air- filled cuff stethoscope  Blood pressure is reported by two values Systolic pressure: maximum pressure (pumping) Diastolic pressure: resting pressure

9. 10.3 T HE G AS L AW  Pressure-volume relationship  The volume of a fixed quantity of gas at constant temperature is inversely proportional to the pressure  B OYLE ’ S L AW

10. 10.3 T HE G AS L AW  For a fixed quantity of gas at constant temperature, the volume of the gas is inversely proportional to its pressure

11.  C HARLES ’ S L AW  Temperature-volume relationship  The volume of a fixed amount of gas at constant pressure is directly proportional to its absolute temperature. 10.3 T HE G AS L AW

12.  A VOGADRO ’ S L AW  Quantity-volume relationship  Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules  The volume of a gas at constant temperature and pressure is directly proportional to the number of moles of the gas 10.3 T HE G AS L AW

13.  A VOGADRO ’ S L AW 10.3 T HE G AS L AW  At the same volume, pressure and temperature, samples of different gases have the same number of molecules but different masses

14. 10.4 T HE I DEAL -G AS E QUATION  The term R is called the gas constant R = 0.08206 L-atm/mol-K = 8.314 J/mol-K  Molar volume: the volume occupied by one mole of ideal gas at STP (273.15K and 1 atm), 22.41 L

15.  One mole of an ideal gas at STP occupies a volume of 22.41 L. One mole of various real gases at STP occupies close to this ideal volume 10.4 T HE I DEAL -G AS E QUATION ▲ Figure 10.11 Comparison of molar volumes at STP

16. Sample Exercise 10.4 Using the Ideal-Gas equation Calcium carbonate, CaCO 3 (s), decomposes upon heating to give CaO(s) and CO 2 (g). A sample of CaCO 3 is decomposed, and the carbon dioxide is collected in a 250-mL flask. After the decomposition is complete, the gas has a pressure of 1.3 atm at a temperature of 31 °C. How many moles of CO 2 gas were generated?

17. Sample Exercise 10.5 The gas pressure in an aerosol can is 1.5 atm at 25 °C. Assuming that the gas inside obeys the ideal-gas equation, what would the pressure be if the can were heated to 450 °C?

19.  G AS D ENSITIES AND M OLAR M ASS 10.5 F URTHER A PPLICATIONS

20.  G AS D ENSITIES AND M OLAR M ASS 10.5 F URTHER A PPLICATIONS ▲ Figure 10.12 Carbon dioxide gas flows downhill because it is denser than air.

21.  G AS D ENSITIES AND M OLAR M ASS 10.5 F URTHER A PPLICATIONS

22.  V OLUMES OF G ASES IN C HEMICAL R EACTIONS 10.5 F URTHER A PPLICATIONS

23. 10.6 G AS M IXTURES AND P ARTIAL P RESSURES  The total pressure of a mixture of gases equals the sum of the pressures that each would exert if it were present alone. - Dalton’s law of partial pressure  Partial pressure The pressure exerted by a particular component of a mixture of gases

24. Sample Exercise 10.10 Applying Dalton’s Law to Partial Pressures A gaseous mixture made from 6.00 g O 2 and 9.00 g CH 4 is placed in a 15.0-L vessel at 0 °C. What is the partial pressure of each gas, and what is the total pressure in the vessel?

25.  P ARTIAL P RESSURE AND M OLE F RACTIONS  Each gas in a mixture behaves independently  We can relate the amount of a given gas in a mixture to its partial pressure 10.6 G AS M IXTURES AND P ARTIAL P RESSURES

26. Sample Exercise 10.11 Relating Mole Fractions to Partial Pressures A study of the effects of certain gases on plant growth requires a synthetic atmosphere composed of 1.5 mol% CO 2, 18.0 mol% O 2, and 80.5 mol% Ar. (a) Calculate the partial pressure of O 2 in the mixture if the total pressure of the atmosphere is to be 745 torr. (b) If this atmosphere is to be held in a 121-L space at 295 K, how many moles of O 2 are needed?

27.  C OLLECTING G ASES OVER W ATER  How to measure the amount of gases generated from a chemical reaction 10.6 G AS M IXTURES AND P ARTIAL P RESSURES ▲ Figure 10.15 Collecting water-insoluble gas over water.

28.  C OLLECTING G ASES OVER W ATER 10.6 G AS M IXTURES AND P ARTIAL P RESSURES

29. 122.6 g/mol

30. 1.Gases consist of large numbers of molecules that are in continuous, random motion 2.The combined volume of all the molecules of the gas is negligible relative to the total volume in which the gas is contained 3.Attractive and repulsive forces between gas molecules are negligible 10.7 K INETIC - MOLECULAR T HEORY  This is a model that aids in our understanding of what happens to gas particles as environmental conditions change.  Summaries of the theory

31. 4.Energy can be transferred between molecules during collisions, but the average kinetic energy of the molecules does not change with time, as long as the temperature of the gas remains constant 5.The average kinetic energy of the molecules is proportional to the absolute temperature  The pressure of a gas is caused by collisions of the molecules with the walls of the container 10.7 K INETIC - MOLECULAR T HEORY  This is a model that aids in our understanding of what happens to gas particles as environmental conditions change.  Summaries of the theory

32.  Although the molecules in a sample of gas have an average kinetic energy and hence an average speed, the Individual molecules move at varying speeds  D ISTRIBUTIONS OF M OLECULAR S PEED 10.7 K INETIC - MOLECULAR T HEORY mp: most probable speed av: average speed rms: root-mean-square speed

33.  Effect of a volume increase at constant temperature If the volume is increased, the molecules must move a longer distance between collisions → pressure decreases  Effect of a temperature increase at constant volume An increase in T means an increase the average kinetic E of the molecule and thus increase in u If there is no change in volume, there will be more collisions with the walls per unit time → pressure increases  A PPLICATIONS TO THE G AS L AWS 10.7 K INETIC - MOLECULAR T HEORY

34. (a) Constant (b) Constant (c) Increase (d) Increase

35. 10.8 M OLECULAR E FFUSION AND D IFFUSION ▲ Figure 10.19 The effect of molecular mass on molecular speeds.

36.  G RAHAM ’ S L AW OF E FFUSION  Effusion ( 유출 ) is the escape of gas molecules through a tiny hole into an evacuated space. 10.8 M OLECULAR E FFUSION AND D IFFUSION ▲ Figure 10.19 Effusion. Gas molecules in top half effuse through pinhole only when they happen to hit pinhole

37.  G RAHAM ’ S L AW OF E FFUSION 10.8 M OLECULAR E FFUSION AND D IFFUSION

38.  D IFFUSION AND M EAN F REE P ATH  Diffusion ( 확산 ) is the spread of one substance throughout a space or throughout a second substance 10.8 M OLECULAR E FFUSION AND D IFFUSION  The diffusion of gases is much slower than molecular speeds because of molecular collisions  The mean free path of a molecule is the average distance traveled by the molecule between collisions  The mean free path for air molecule 60 nm at sea level 10 cm at 100 km in altitude

39. 10.9 R EAL G ASES  Although the ideal-gas equation is a very useful description of gases, all real gases fail to obey the relationship to some degree ▲ Figure 10.24 Gases behave more ideally at low pressure than at high pressure. The volume of gas molecules is not negligible at high pressure. ▲ Figure 10.25 In any real gas, attractive intermolecular forces reduce pressure to values lower than in an ideal gas.

40. 10.9 R EAL G ASES At high P, gas volumes are not negligible Attractive forces between molecules reduce the pressure ▲ Figure 10.22 The effect of pressure on the behavior of several real gases at constant T. The deviations increases with increasing P.

41. 10.9 R EAL G ASES  Cooling a gas increase the chance for molecules to interact with each other ▲ Figure 10.23 The effect of temperature and pressure on the behavior of nitrogen gas. The deviations increase with decreasing T.

42.  T HE VAN DER W AALS E QUATION  The ideal-gas equation can be adjusted to take the deviations from ideal behavior into account  The van der Waals Equation 10.9 R EAL G ASES = P ideal = V ideal

43.  T HE VAN DER W AALS E QUATION  a and b values increase with mass of the molecule and the complexity of its structure 10.9 R EAL G ASES

46. Homework Practice Exercises p397, 399, 402, and 412 Due on 06-13 (Thur)