7.f^Chapter 7 Ending. CHAPTER HIGHLIGHTS^222^227^,,^13920^14729%
CHAPTER HIGHLIGHTS
KEY TERMS
KEY CONCEPTS
  1. What is pressure and what units are used to measure it? (7.2)

    • Pressure is the force per unit area. The pressure of a gas is the force exerted when gas particles strike a surface. Pressure is measured by a barometer and recorded in atmospheres (atm), millimeters of mercury (mm Hg), or pounds per square inch (psi).

    • 1 atm = 760 mm Hg = 14.7 psi.

  2. What are gas laws and how are they used to describe the relationship between the pressure, volume, and temperature of a gas? (7.3)

    • Because gas particles are far apart and behave independently, a set of gas laws describes the behavior of all gases regardless of their identity. Three gas laws—Boyle's law, Charles's law, and Gay–Lussac's law—describe the relationship between the pressure, volume, and temperature of a gas. These gas laws are summarized in “Key Equations—The Gas Laws” on page 223.

    • For a constant amount of gas, the following relationships exist.

      • The pressure and volume of a gas are inversely related, so increasing the pressure decreases the volume at constant temperature.

      • The volume of a gas is proportional to its Kelvin temperature, so increasing the temperature increases the volume at constant pressure.

      • The pressure of a gas is proportional to its Kelvin temperature, so increasing the temperature increases the pressure at constant volume.

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  4. Describe the relationship between the volume and number of moles of a gas. (7.4)

    • Avogadro's law states that when temperature and pressure are held constant, the volume of a gas is proportional to its number of moles.

    • One mole of any gas has the same volume, the standard molar volume of 22.4 L, at 1 atm and 273 K (STP).

  5. What is the ideal gas law? (7.5)

    • The ideal gas law is an equation that relates the pressure (P), volume (V), temperature (T), and number of moles (n) of a gas; PV = nRT, where R is the universal gas constant. The ideal gas law can be used to calculate any one of the four variables, as long as the other three variables are known.

  6. What is Dalton's law and how is it used to relate partial pressures and the total pressure of a gas mixture? (7.6)

    • Dalton's law states that the total pressure of a gas mixture is the sum of the partial pressures of its component gases. The partial pressure is the pressure exerted by each component of a mixture.

  7. What types of intermolecular forces exist and how do they determine a compound's boiling point and melting point? (7.7)

    • Intermolecular forces are the forces of attraction between molecules. Three types of intermolecular forces exist in covalent compounds. London dispersion forces are due to momentary changes in electron density in a molecule. Dipole–dipole interactions are due to permanent dipoles. Hydrogen bonding, the strongest intermolecular force, results when a H atom bonded to an O, N, or F, is attracted to an O, N, or F atom in another molecule.

    • The stronger the intermolecular forces, the higher the boiling point and melting point of a compound.

  8. Describe three features of the liquid state—vapor pressure, viscosity, and surface tension. (7.8)

    • Vapor pressure is the pressure exerted by gas molecules in equilibrium with the liquid phase. Vapor pressure increases with increasing temperature. The higher the vapor pressure at a given temperature, the lower the boiling point of a compound.

    • Viscosity measures a liquid's resistance to flow. More viscous compounds tend to have stronger intermolecular forces or they have high molecular weights.

    • Surface tension measures a liquid's resistance to spreading out. The stronger the intermolecular forces, the higher the surface tension.

  9. Describe the features of different types of solids. (7.9)

    • Solids can be amorphous or crystalline. An amorphous solid has no regular arrangement of particles. A crystalline solid has a regular arrangement of particles in a repeating pattern. There are four types of crystalline solids. Ionic solids are composed of ions. Molecular solids are composed of individual molecules. Network solids are composed of vast repeating arrays of covalently bonded atoms in a regular three-dimensional arrangement. Metallic solids are composed of metal cations with a cloud of delocalized electrons.

  10. Describe the energy changes that accompany changes of state. (7.10)

    • A phase change converts one state to another. Energy is absorbed when a more organized state is converted to a less organized state. Thus, energy is absorbed when a solid melts to form a liquid, or when a liquid vaporizes to form a gas.

    • Energy is released when a less organized state is converted to a more organized state. Thus, energy is released when a gas condenses to form a liquid, or a liquid freezes to form a solid.

    • The heat of fusion is the energy needed to melt 1 g of a substance, while the heat of vaporization is the energy needed to vaporize 1 g of a substance.

KEY EQUATIONS—THE GAS LAWS

Name

Equation

Variables Related

Constant Terms

Boyle's law

P1V1 = P2V2

P, V

T, n

Charles's law

V, T

P, n

Gay–Lussac's law

P, T

V, n

Combined gas law

P, V, T

n

Avogadro's law

V, n

P, T

Ideal gas law

PV = nRT

P, V, T, n

R

Page 224
PROBLEMS

Selected in-chapter and end-of-chapter problems have brief answers provided in Appendix B.

Pressure
  • 7.43

    What is the relationship between the units mm Hg and atm?

  • 7.44

    What is the relationship between the units mm Hg and psi?

  • 7.45

    The highest atmospheric pressure ever measured is 814.3 mm Hg, recorded in Mongolia in December, 2001. Convert this value to atmospheres.

  • 7.46

    The lowest atmospheric pressure ever measured is 652.5 mm Hg, recorded during Typhoon Tip on October 12, 1979. Convert this value to atmospheres.

  • 7.47

    Convert each quantity to the indicated unit.

    1. 2.8 atm to psi

    2. 520 mm Hg to atm

    3. 20.0 atm to torr

    4. 100. mm Hg to Pa

  • 7.48

    The compressed air tank of a scuba diver reads 3,200 psi at the beginning of a dive and 825 psi at the end of a dive. Convert each of these values to atm and mm Hg.

Boyle's Law
  • 7.49

    Assuming a fixed amount of gas at constant temperature, complete the following table.

      P1 V1 P2 V2
    a. 2.0 atm 3.0 L 8.0 atm ?
    b. 55 mm Hg 0.35 L 18 mm Hg ?
    c. 705 mm Hg 215 mL ? 1.52 L
  • 7.50

    Assuming a fixed amount of gas at constant temperature, complete the following table.

      P1 V1 P2 V2
    a. 2.5 atm 1.5 L 3.8 atm ?
    b. 2.0 atm 350 mL 750 mm Hg ?
    c. 75 mm Hg 9.1 mL ? 890 mL
  • 7.51

    If a scuba diver releases a 10.-mL air bubble below the surface where the pressure is 3.5 atm, what is the volume of the bubble when it rises to the surface and the pressure is 1.0 atm?

  • 7.52

    If someone takes a breath and the lungs expand from 4.5 L to 5.6 L in volume, and the initial pressure was 756 mm Hg, what is the pressure inside the lungs before any additional air is pulled in?

Charles's Law
  • 7.53

    Assuming a fixed amount of gas at constant pressure, complete the following table.

      V1 T1 V2 T2
    a. 5.0 L 310 K ? 250 K
    b. 150 mL 45 K ? 45 °C
    c. 60.0 L 0.0 °C 180 L ?
  • 7.54

    Assuming a fixed amount of gas at constant pressure, complete the following table.

      V1 T1 V2 T2
    a. 10.0 mL 210 K ? 450 K
    b. 255 mL 55 °C ? 150 K
    c. 13 L −150 °C 52 L ?
  • 7.55

    If a balloon containing 2.2 L of gas at 25 °C is cooled to −78 °C, what is its new volume?

  • 7.56

    How hot must the air in a balloon be heated if initially it has a volume of 750. L at 20 °C and the final volume must be 1,000. L?

Gay–Lussac's Law
  • 7.57

    Assuming a fixed amount of gas at constant volume, complete the following table.

      P1 T1 P2 T2
    a. 3.25 atm 298 K ? 398 K
    b. 550 mm Hg 273 K ? −100. °C
    c. 0.50 atm 250 °C 955 mm Hg ?
  • 7.58

    Assuming a fixed amount of gas at constant volume, complete the following table.

      P1 T1 P2 T2
    a. 1.74 atm 120 °C ? 20. °C
    b. 220 mm Hg 150 °C ? 300. K
    c. 0.75 atm 198 °C 220 mm Hg ?
  • 7.59

    An autoclave is a pressurized container used to sterilize medical equipment by heating it to a high temperature under pressure. If an autoclave containing steam at 100. °C and 1.0 atm pressure is then heated to 150. °C, what is the pressure inside it?

  • 7.60

    If a plastic container at 1.0 °C and 750. mm Hg is heated in a microwave oven to 80. °C, what is the pressure inside the container?

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Combined Gas Law
  • 7.61

    Assuming a fixed amount of gas, complete the following table.

      P1 V1 T1 P2 V2 T2
    a. 0.90 atm 4.0 L 265 K ? 3.0 L 310 K
    b. 1.2 atm 75 L 5.0 °C 700. mm Hg ? 50 °C
    c. 200. mm Hg 125 mL 298 K 100. mm Hg 0.62 L ?
  • 7.62

    Assuming a fixed amount of gas, complete the following table.

      P1 V1 T1 P2 V2 T2
    a. 0.55 atm 1.1 L 340 K ? 3.0 L 298 K
    b. 735 mm Hg 1.2 L 298 K 1.1 atm ? 0.0 °C
    c. 7.5 atm 230 mL −120 °C 15 atm 0.45 L ?
  • 7.63

    If a compressed air cylinder for scuba diving contains 6.0 L of gas at 18 °C and 200. atm pressure, what volume does the gas occupy at 1.0 atm and 25 °C?

  • 7.64

    What happens to the pressure of a sample with each of the following changes?

    1. Double the volume and halve the Kelvin temperature.

    2. Double the volume and double the Kelvin temperature.

    3. Halve the volume and double the Kelvin temperature.

Avogadro's Law
  • 7.65

    What is the difference between STP and standard molar volume?

  • 7.66

    Given the same number of moles of two gases at STP conditions, how do the volumes of two gases compare? How do the masses of the two gas samples compare?

  • 7.67

    How many moles of helium are contained in each volume at STP: (a) 5.0 L; (b) 11.2 L; (c) 50.0 mL?

  • 7.68

    How many moles of argon are contained in each volume at STP: (a) 4.0 L; (b) 31.2 L; (c) 120 mL?

  • 7.69

    Calculate the volume of each substance at STP.

    1. 4.2 mol Ar

    2. 3.5 g CO2

    3. 2.1 g N2

  • 7.70

    Calculate the volume of each substance at STP.

    1. 4.2 mol N2

    2. 6.5 g He

    3. 22.0 g CH4

  • 7.71

    What volume does 3.01 × 1021 molecules of N2 occupy at STP?

  • 7.72

    What volume does 1.50 × 1024 molecules of CO2 occupy at STP?

Ideal Gas Law
  • 7.73

    How many moles of gas are contained in a human breath that occupies 0.45 L and has a pressure of 747 mm Hg at 37 °C?

  • 7.74

    How many moles of gas are contained in a compressed air tank for scuba diving that has a volume of 7.0 L and a pressure of 210 atm at 25 °C?

  • 7.75

    How many moles of air are present in the lungs if they occupy a volume of 5.0 L at 37 °C and 740 mm Hg? How many molecules of air does this correspond to?

  • 7.76

    If a cylinder contains 10.0 g of CO2 in 10.0 L at 325 K, what is the pressure?

  • 7.77

    Which sample contains more moles: 2.0 L of O2 at 273 K and 500 mm Hg, or 1.5 L of N2 at 298 K and 650 mm Hg? Which sample weighs more?

  • 7.78

    An unknown amount of gas occupies 30.0 L at 2.1 atm and 298 K. How many moles does the sample contain? What is the mass if the gas is helium? What is the mass if the gas is argon?

Dalton's Law and Partial Pressure
  • 7.79

    Air pressure on the top of Mauna Loa, a 13,000-ft mountain in Hawaii, is 460 mm Hg. What are the partial pressures of O2 and N2, which compose 21% and 78% of the atmosphere, respectively?

  • 7.80

    If air contains 21% O2, what is the partial pressure of O2 in a cylinder of compressed air at 175 atm?

  • 7.81

    The partial pressure of N2 in the air is 593 mm Hg at 1 atm. What is the partial pressure of N2 in a bubble of air a scuba diver breathes when he is 66 ft below the surface of the water where the pressure is 3 atm?

  • 7.82

    If N2 is added to a balloon that contains O2 (partial pressure 450 mm Hg) and CO2 (partial pressure 150 mm Hg) to give a total pressure of 850 mm Hg, what is the partial pressure of each gas in the final mixture?

Intermolecular Forces
  • 7.83

    What is the difference between dipole–dipole interactions and London dispersion forces?

  • 7.84

    What is the difference between dipole–dipole interactions and hydrogen bonding?

  • 7.85

    Why is H2O a liquid at room temperature, but H2S, which has a higher molecular weight and a larger surface area, is a gas at room temperature?

  • 7.86

    Why is Cl2 a gas, Br2 a liquid, and I2 a solid at room temperature?

  • 7.87

    What types of intermolecular forces are exhibited by each compound? Chloroethane is a local anesthetic and cyclopropane is a general anesthetic.

  • Page 226
  • 7.88

    What types of intermolecular forces are exhibited by each compound? Acetaldehyde is formed when ethanol, the alcohol in alcoholic beverages, is metabolized, and acetic acid gives vinegar its biting odor and taste.

  • 7.89

    Which molecules are capable of intermolecular hydrogen bonding?

    1. H—C≡≡C—H

    2. CO2

    3. Br2

  • 7.90

    Which molecules are capable of intermolecular hydrogen bonding?

    1. N2

    2. HI

  • 7.91

    Can two molecules of formaldehyde (H2C==O) intermolecularly hydrogen bond to each other? Explain why or why not.

  • 7.92

    Why is the melting point of NaCl (801 °C) much higher than the melting point of water (0 °C)?

  • 7.93

    Ethylene and methanol have approximately the same molar mass.

    1. What types of intermolecular forces are present in each compound?

    2. Which compound has the higher boiling point?

    3. Which compound has the higher vapor pressure at a given temperature?

  • 7.94

    Ethanol and dimethyl ether have the same molecular formula.

    1. What types of intermolecular forces are present in each compound?

    2. Which compound has the higher boiling point?

    3. Which compound has the higher vapor pressure at a given temperature?

Liquids and Solids
  • 7.95

    What is the difference between vapor pressure and partial pressure?

  • 7.96

    What is the difference between viscosity and surface tension?

  • 7.97

    Given the following vapor pressures at 20 °C, arrange the compounds in order of increasing boiling point: butane, 1,650 mm Hg; acetaldehyde, 740 mm Hg, Freon-113, 284 mm Hg.

  • 7.98

    Using the given boiling points, predict which compound has the higher vapor pressure at a given temperature.

    1. ethanol (C2H6O, bp 78 °C) or 1-propanol (C3H8O, bp 97 °C)

    2. hexane (C6H14, bp 69 °C) or octane (C8H18, bp 125 °C)

  • 7.99

    Explain why glycerol is more viscous than water, but acetone is less viscous than water. Glycerol is a component of skin lotions and creams. Acetone is the main ingredient in nail polish remover.

  • 7.100

    Explain the following observation. When a needle is carefully placed on the surface of water, it floats, yet when its tip is pushed below the surface, it sinks to the bottom.

  • 7.101

    What is the difference between an ionic solid and a metallic solid?

  • 7.102

    What is the difference between a molecular solid and a network solid?

  • 7.103

    Classify each solid as amorphous, ionic, molecular, network, or metallic.

    1. KI

    2. CO2

    3. bronze, an alloy of Cu and Sn

    4. diamond

    5. the plastic polyethylene

  • 7.104

    Classify each solid as amorphous, ionic, molecular, network, or metallic.

    1. CaCO3

    2. CH3COOH (acetic acid)

    3. Ag

    4. graphite

    5. the plastic polypropylene

Energy and Phase Changes
  • 7.105

    What is the difference between evaporation and condensation?

  • 7.106

    What is the difference between vaporization and condensation?

  • Page 227
  • 7.107

    What is the difference between sublimation and deposition?

  • 7.108

    What is the difference between melting and freezing?

  • 7.109

    Indicate whether heat is absorbed or released in each process.

    1. melting 100 g of ice

    2. freezing 25 g of water

    3. condensing 20 g of steam

    4. vaporizing 30 g of water

  • 7.110

    What is the difference between the heat of fusion and the heat of vaporization?

  • 7.111

    Which process requires more energy, melting 250 g of ice or vaporizing 50.0 g of water? The heat of fusion of water is 79.7 cal/g and the heat of vaporization is 540 cal/g.

  • 7.112

    How much energy in kilocalories is needed to vaporize 255 g of water? The heat of vaporization of water is 540 cal/g.

General Problems
  • 7.113

    Explain the difference between Charles's law and Gay–Lussac's law, both of which deal with the temperature of gases.

  • 7.114

    Explain the difference between Charles's law and Avogadro's law, both of which deal with the volume of gases.

  • 7.115

    Explain the difference between Boyle's law and Gay–Lussac's law, both of which deal with the pressure of gases.

  • 7.116

    What is the difference between the combined gas law and the ideal gas law?

  • 7.117

    A balloon is filled with helium at sea level. What happens to the volume of the balloon in each instance? Explain each answer.

    1. The balloon floats to a higher altitude.

    2. The balloon is placed in a bath of liquid nitrogen at −196 °C.

    3. The balloon is placed inside a hyperbaric chamber at a pressure of 2.5 atm.

    4. The balloon is heated inside a microwave.

  • 7.118

    Suppose you have a fixed amount of gas in a container with a movable piston, as drawn. Re-draw the container and piston to illustrate what it looks like after each of the following changes takes place.

    1. The temperature is held constant and the pressure is doubled.

    2. The pressure is held constant and the Kelvin temperature is doubled.

    3. The pressure is halved and the Kelvin temperature is halved.

Applications
  • 7.119

    What is the difference between the systolic and diastolic blood pressure?

  • 7.120

    What is hypertension and what are some of its complications?

  • 7.121

    If you pack a bag of potato chips for a snack on a plane ride, the bag appears to have inflated when you take it out to open. Explain why this occurs. If the initial volume of air in the bag was 250 mL at 760 mm Hg, and the plane is pressurized at 650 mm Hg, what is the final volume of the bag?

  • 7.122

    Why does a bubble at the bottom of a glass of a soft drink get larger as it rises to the surface?

  • 7.123

    Explain why cooling a full glass water bottle to −10 °C causes the bottle to crack.

  • 7.124

    What happens to the density of a gas if the temperature is increased but the pressure is held constant? Use this information to explain how wind currents arise.

  • 7.125

    A common laboratory test for a patient is to measure blood gases—that is, the partial pressures of O2 and CO2 in oxygenated blood. Normal values are 100 mm Hg for O2 and 40 mm Hg for CO2. A high or low level of one or both readings has some underlying cause. Offer an explanation for each of the following situations.

    1. If a patient comes in agitated and hyperventilating—breathing very rapidly—the partial pressure of O2 is normal but the partial pressure of CO2 is 22 mm Hg.

    2. A patient with chronic lung disease has a partial pressure of O2 of 60 mm Hg and a partial pressure of CO2 of 60 mm Hg.

  • 7.126

    If a scuba diver inhales 0.50 L of air at a depth of 100. ft and 4.0 atm pressure, what volume does this air occupy at the surface of the water, assuming air pressure is 1.0 atm? When a scuba diver must make a rapid ascent to the surface, he is told to exhale slowly as he ascends. How does your result support this recommendation?

CHALLENGE QUESTIONS