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2: Naming Binary Compounds That Contain a Metal and a Nonmetal (Types I and II)

2: Naming Binary Compounds That Contain a Metal and a Nonmetal (Types I and II)

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466 Chapter 14 Liquids and Solids

d. Solid argon contains argon atoms, which cannot form covalent

bonds to each other. It is an atomic solid.

e. Sulfur contains S8 molecules, so it is a molecular solid.



Self-Check EXERCISE 14.2 Name the type of crystalline solid formed by each of the following substances:

a. sulfur trioxide

b. barium oxide

c. gold

See Problems 14.41 and 14.42. ■



C H A P T E R



14



REVIEW



F



Key Terms

normal boiling

point (14.1)

heating/cooling

curve (14.1)

normal freezing

point (14.1)

intramolecular

forces (14.2)

intermolecular

forces (14.2)

molar heat of

fusion (14.2)

molar heat of

vaporization (14.2)

dipole–dipole

attraction (14.3)



hydrogen bonding (14.3)

London dispersion

forces (14.3)

vaporization

(evaporation) (14.4)

condensation (14.4)

vapor pressure (14.4)

crystalline solid (14.5)

ionic solid (14.5)

molecular solid (14.5)

atomic solid (14.5)

electron sea model (14.6)

alloy (14.6)

substitutional

alloy (14.6)

interstitial alloy (14.6)



Summary

1. Liquids and solids exhibit some similarities and are

very different from the gaseous state.

2. The temperature at which a liquid changes its state to

a gas (at 1 atm pressure) is called the normal boiling

point of that liquid. Similarly, the temperature at

which a liquid freezes (at 1 atm pressure) is the normal freezing point. Changes of state are physical

changes, not chemical changes.



VP



directs you to the Chemistry in Focus feature in the chapter

indicates visual problems

interactive versions of these problems are assignable in OWL



3. To convert a substance from the solid to the liquid

and then to the gaseous state requires the addition of

energy. Forces among the molecules in a solid or a liquid must be overcome by the input of energy. The energy required to melt 1 mole of a substance is called

the molar heat of fusion, and the energy required to

change 1 mole of liquid to the gaseous state is called

the molar heat of vaporization.

4. There are several types of intermolecular forces.

Dipole–dipole interactions occur when molecules

with dipole moments attract each other. A particularly strong dipole–dipole interaction called hydrogen bonding occurs in molecules that contain hydrogen bonded to a very electronegative element such as

N, O, or F. London dispersion forces occur when instantaneous dipoles in atoms or nonpolar molecules

lead to relatively weak attractions.

5. The change of a liquid to its vapor is called vaporization or evaporation. The process whereby vapor molecules form a liquid is called condensation. In a

closed container, the pressure of the vapor over its

liquid reaches a constant value called the vapor pressure of the liquid.

6. Many solids are crystalline (contain highly regular

arrangements of their components). The three types



Chapter Review

of crystalline solids are ionic, molecular, and atomic

solids. In ionic solids, the ions are packed together in

a way that maximizes the attractions of oppositely

charged ions and minimizes the repulsions among

identically charged ions. Molecular solids are held together by dipole–dipole attractions if the molecules

are polar and by London dispersion forces if the molecules are nonpolar. Atomic solids are held together

by covalent bonding forces or London dispersion

forces, depending on the atoms present.



Active Learning Questions

These questions are designed to be considered by groups of

students in class. Often these questions work well for introducing a particular topic in class.

1. You seal a container half-filled with water. Which

best describes what occurs in the container?

a. Water evaporates until the air becomes saturated

with water vapor; at this point, no more water

evaporates.

b. Water evaporates until the air becomes overly saturated (supersaturated) with water, and most of

this water recondenses; this cycle continues until

a certain amount of water vapor is present, and

then the cycle ceases.

c. The water does not evaporate because the container is sealed.

d. Water evaporates, and then water evaporates and

recondenses simultaneously and continuously.

e. The water evaporates until it is eventually all in vapor form.

Justify your choice and for choices you did not pick,

explain what is wrong with them.

2. Explain the following: You add 100 mL of water to a

500-mL round-bottomed flask and heat the water until it is boiling. You remove the heat and stopper the

flask, and the boiling stops. You then run cool water

over the neck of the flask, and the boiling begins

again. It seems as though you are boiling water by

cooling it.

3. Is it possible for the dispersion forces in a particular

substance to be stronger than hydrogen-bonding

forces in another substance? Explain your answer.

4. Does the nature of intermolecular forces change

when a substance goes from a solid to a liquid, or

from a liquid to a gas? What causes a substance to undergo a phase change?

5. How does vapor pressure change with changing temperature? Explain.



467



8. How do the following physical properties depend on

the strength of intermolecular forces? Explain.

a. melting point

b. boiling point

c. vapor pressure

9. Look at Figure 14.2. Why doesn’t temperature increase continuously over time? That is, why does the

temperature stay constant for periods of time?

10. Which are stronger, intermolecular or intramolecular forces for a given molecule? What observation(s)

have you made that supports this position? Explain.

11. Why does water evaporate at all?

12. Sketch a microscopic picture of water and distinguish

between intramolecular bonds and intermolecular forces.

Which correspond to the bonds we draw in Lewis

structures?

13. Which has the stronger intermolecular forces: N2 or

H2O? Explain.

14. Which gas would behave more ideally at the same

conditions of pressure and temperature: CO or N2?

Why?

15. You have seen that the water molecule has a bent

shape and therefore is a polar molecule. This accounts for many of water’s interesting properties.

What if the water molecule were linear? How would

this affect the properties of water? How would life be

different?

16. True or false? Methane (CH4) is more likely to form

stronger hydrogen bonding than is water because

each methane molecule has twice as many hydrogen

atoms. Provide a concise explanation of hydrogen

bonding to go with your answer.

17. Why should it make sense that N2 exists as a gas?

Given your answer, how is it possible to make liquid

nitrogen? Explain why lowering the temperature

works.

18. White phosphorus and sulfur both are called molecular solids even though each is made of only phosphorus and sulfur, respectively. How can they be considered molecular solids? If this is true, why isn’t

diamond (which is made up only of carbon) a molecular solid?



6. What occurs when the vapor pressure of a liquid is

equal to atmospheric pressure? Explain.



19. Why is it incorrect to use the term “molecule of

NaCl” but correct to use the term “molecule of

H2O”? Is the term “molecule of diamond” correct?

Explain.



7. What is the vapor pressure of water at 100 °C? How

do you know?



20. Which would you predict should be larger for a given

substance: ⌬Hvap or ⌬Hfus? Explain why.



468 Chapter 14 Liquids and Solids

VP 21. In the diagram below, which lines represent the hydrogen bonds?



What is the empirical formula of this ionic compound?



H

H



Questions and Problems



O



H

O



O

H



H



H



H



H



O



O

H



H

O



14.1 Water and Its Phase Changes



H



H



H

H



O



a. The dotted lines between the hydrogen atoms of

one water molecule and the oxygen atoms of a different water molecule.

b. The solid lines between a hydrogen atom and oxygen atom in the same water molecule.

c. Both the solid lines and dotted lines represent the

hydrogen bonds.

d. There are no hydrogen bonds represented in the

diagram.



VP 22. Use the heating/cooling curve below to answer the

following questions.



1. Gases have (higher/lower) densities than liquids or

solids.

2. Liquids and solids are (more/less) compressible than

are gases.

3. What evidence do we have that the solid form of water is less dense than the liquid form of water at its

freezing/melting point?

4. The enthalpy (⌬H) of vaporization of water is about

seven times larger than water’s enthalpy of fusion

(41 kJ/mol versus 6 kJ/mol). What does this tell us

about the relative similarities among the solid, liquid,

and gaseous states of water?

5. Consider a sample of ice being heated from Ϫ5 °C to

ϩ5 °C. Describe on both a macroscopic and a microscopic basis what happens to the ice as the temperature reaches 0 °C.

6. Sketch a heating/cooling curve for water, starting out

at Ϫ20 °C and going up to 120 °C, applying heat to

the sample at a constant rate. Mark on your sketch

the portions of the curve that represent the melting

of the solid and the boiling of the liquid.



160

140



Temperature (°C)



QUESTIONS



120

100



14.2 Energy Requirements for the Changes of State



80



QUESTIONS



60

40



7. Are changes in state physical or chemical changes for

molecular solids? Why?



20

0



Heat added



a. What is the freezing point of the liquid?

b. What is the boiling point of the liquid?

c. Which is greater: the head of fusion or the heat of

vaporization? Explain.



VP 23. Assume the two-dimensional structure of an ionic

compound, MxAy, is



8. Describe in detail the microscopic processes that take

place when a solid melts and when a liquid boils.

What kind of forces must be overcome? Are any

chemical bonds broken during these processes?

9. Explain the difference between intramolecular and

intermolecular forces.

10. The forces that connect two hydrogen atoms to an

oxygen atom in a water molecule are (intermolecular/intramolecular), but the forces that hold water molecules close together in an ice cube are (intermolecular/intramolecular).

11. Discuss the similarities and differences between the

arrangements of molecules and the forces between

molecules in liquid water versus steam, and in liquid

water versus ice.

12. What does the molar heat of fusion of a substance represent?



All even-numbered Questions and Problems have answers in the back of this book and solutions in the Solutions Guide.



Chapter Review

PROBLEMS

13. The following data have been collected for substance

X. Construct a heating curve for substance X. (The

drawing does not need to be absolutely to scale, but

it should clearly show relative differences.)

normal melting point



Ϫ15 °C



molar heat of fusion



2.5 kJ/mol



normal boiling point



134 °C



molar heat of vaporization



55.3 kJ/mol



14. The molar heat of fusion of aluminum metal is

10.79 kJ/mol, whereas its heat of vaporization is

293.4 kJ/mol.

a. Why is the heat of fusion of aluminum so much

smaller than the heat of vaporization?

b. What quantity of heat would be required to vaporize 1.00 g of aluminum at its normal boiling point?

c. What quantity of heat would be evolved if 5.00 g

of liquid aluminum freezes at its normal freezing

point?

d. What quantity of heat would be required to melt

0.105 mole of aluminum at its normal melting

point?

15. The molar heat of fusion of benzene is 9.92 kJ/mol.

Its molar heat of vaporization is 30.7 kJ/mol. Calculate the heat required to melt 8.25 g of benzene at its

normal melting point. Calculate the heat required to

vaporize 8.25 g of benzene at its normal boiling

point. Why is the heat of vaporization more than

three times the heat of fusion?

16. The molar heats of fusion and vaporization for silver

are 11.3 kJ/mol and 250. kJ/mol, respectively. Silver’s

normal melting point is 962 °C, and its normal boiling point is 2212 °C. What quantity of heat is required to melt 12.5 g of silver at 962 °C? What quantity of heat is liberated when 4.59 g of silver vapor

condenses at 2212 °C?

17. Given that the specific heat capacities of ice and

steam are 2.06 J/g °C and 2.03 J/g °C, respectively,

and considering the information about water given

in Exercise 16, calculate the total quantity of heat

evolved when 10.0 g of steam at 200. °C is condensed, cooled, and frozen to ice at Ϫ50. °C.

18. It requires 113 J to melt 1.00 g of sodium metal at its

normal melting point of 98 °C. Calculate the molar

heat of fusion of sodium.



14.3 Intermolecular Forces

QUESTIONS

19. Consider the iodine monochloride molecule, ICl. Because chlorine is more electronegative than iodine,

this molecule is a dipole. How would you expect iodine monochloride molecules in the gaseous state to

orient themselves with respect to each other as the



469



sample is cooled and the molecules begin to aggregate? Sketch the orientation you would expect.

20. Dipole–dipole forces become

between the dipoles increases.



as the distance



21. The text implies that hydrogen bonding is a special

case of very strong dipole–dipole interactions possible among only certain atoms. What atoms in addition to hydrogen are necessary for hydrogen bonding? How does the small size of the hydrogen atom

contribute to the unusual strength of the dipole–

dipole forces involved in hydrogen bonding?

22. The normal boiling point of water is unusually high,

compared to the boiling points of H2S, H2Se, and

H2Te. Explain this observation in terms of the hydrogen bonding that exists in water, but that does not exist in the other compounds.

23. Why are the dipole–dipole interactions between polar molecules not important in the vapor phase?

24. What are London dispersion forces, and how do they

arise?



PROBLEMS

25. What type of intermolecular forces is active in the liquid state of each of the following substances?

a.

b.

c.

d.



Ne

CO

CH3OH

Cl2



26. Discuss the types of intermolecular forces acting in

the liquid state of each of the following substances.

a.

b.

c.

d.



Xe

NH3

F2

ICl



27. The boiling points of the noble gas elements are listed

below. Comment on the trend in the boiling points.

Why do the boiling points vary in this manner?

He



Ϫ272 °C



Kr



Ϫ152.3 °C



Ne



Ϫ245.9 °C



Xe Ϫ107.1 °C



Ar



Ϫ185.7 °C



Rn



Ϫ61.8 °C



28. The heats of fusion of three substances are listed below. Explain the trend this list reflects.

HI



2.87 kJ/mol



HBr



2.41 kJ/mol



HCl



1.99 kJ/mol



29. When dry ammonia gas (NH3) is bubbled into a

125-mL sample of water, the volume of the sample

(initially, at least) decreases slightly. Suggest a reason

for this.



All even-numbered Questions and Problems have answers in the back of this book and solutions in the Solutions Guide.



470 Chapter 14 Liquids and Solids

30. When 50 mL of liquid water at 25 °C is added to

50 mL of ethanol (ethyl alcohol), also at 25 °C, the

combined volume of the mixture is considerably less

than 100 mL. Give a possible explanation.



14.4 Evaporation and Vapor Pressure

QUESTIONS

31. What is evaporation? What is condensation? Which

of these processes is endothermic and which is

exothermic?

32. If you’ve ever opened a bottle of rubbing alcohol or

other solvent on a warm day, you may have heard a

little “whoosh” as the vapor that had built up above

the liquid escapes. Describe on a microscopic basis

how a vapor pressure builds up in a closed container

above a liquid. What processes in the container give

rise to this phenomenon?

33. What do we mean by a dynamic equilibrium? Describe

how the development of a vapor pressure above a liquid represents such an equilibrium.

34. Consider Figure 14.10. Imagine you are talking to a

friend who has not taken any science courses, and explain how the figure demonstrates the concept of vapor pressure and enables it to be measured.

PROBLEMS

35. Which substance in each pair would be expected to

have a lower boiling point? Explain your reasoning.

a. CH3OH or CH3CH2CH2OH

b. CH3CH3 or CH3CH2OH

c. H2O or CH4

36. Which substance in each pair would be expected to

show the largest vapor pressure at a given temperature? Explain your reasoning.

a. H2O(l) or HF(l)

b. CH3OCH3(l) or CH3CH2OH(l)

c. CH3OH(l) or CH3SH(l)

37. Although water and ammonia differ in molar mass

by only one unit, the boiling point of water is over

100 °C higher than that of ammonia. What forces in

liquid water that do not exist in liquid ammonia

could account for this observation?

38. Two molecules that contain the same number of each

kind of atom but that have different molecular structures are said to be isomers of each other. For example,

both ethyl alcohol and dimethyl ether (shown below) have the formula C2H6O and are isomers. Based

on considerations of intermolecular forces, which

substance would you expect to be more volatile?

Which would you expect to have the higher boiling

point? Explain.

dimethyl ether

CH3OOOCH3



ethyl alcohol

CH3OCH2OOH



14.5 The Solid State: Types of Solids

QUESTIONS

39. What are crystalline solids? What kind of microscopic structure do such solids have? How is this microscopic structure reflected in the macroscopic appearance of such solids?

40. On the basis of the smaller units that make up the

crystals, cite three types of crystalline solids. For each

type of crystalline solid, give an example of a substance that forms that type of solid.



14.6 Bonding in Solids

QUESTIONS

41. How do ionic solids differ in structure from molecular

solids? What are the fundamental particles in each?

Give two examples of each type of solid and indicate

the individual particles that make up the solids in

each of your examples.

42. A common prank on college campuses is to switch

the salt and sugar on dining hall tables, which is usually easy because the substances look so much alike.

Yet, despite the similarity in their appearance, these

two substances differ greatly in their properties, since

one is a molecular solid and the other is an ionic

solid. How do the properties differ and why?

43. Ionic solids are generally considerably harder than

most molecular solids. Explain.

44. Although crystals of table salt (sodium chloride) and

table sugar (sucrose) look very similar to the naked

eye, the melting point of sucrose (186 °C) is several

hundred degrees less than the melting point of

sodium chloride (801 °C). Explain.

45. The forces holding together a molecular solid are

much (stronger/weaker) than the forces between particles in an ionic solid.

46. Explain the overall trend in melting points given below in terms of the forces among particles in the

solids indicated.

Hydrogen, H2



Ϫ259 °C



Ethyl alcohol, C2H5OH



Ϫ114 °C



Water, H2O



0 °C



Sucrose, C12H22O11



186 °C



Calcium chloride, CaCl2



772 °C



47. What is a network solid? Give an example of a network

solid and describe the bonding in such a solid. How

does a network solid differ from a molecular solid?

48. Ionic solids do not conduct electricity in the solid

state, but are strong conductors in the liquid state

and when dissolved in water. Explain.



All even-numbered Questions and Problems have answers in the back of this book and solutions in the Solutions Guide.



Chapter Review

49. What is an alloy? Explain the differences in structure

between substitutional and interstitial alloys. Give an

example of each type.

F 50. The “Chemistry in Focus” segment Metal with a Mem-



ory discusses Nitinol, an alloy that “remembers” a

shape originally impressed in it. Which elements

compose Nitinol, and why is it classified as an alloy?



Additional Problems

M AT C H I N G

For Exercises 51–60 choose one of the following terms to

match the definition or description given.

a.

b.

c.

d.

e.

f.

g.

h.

i.

j.

k.

l.

m.

n.



alloy

specific heat

crystalline solid

dipole–dipole attraction

equilibrium vapor pressure

intermolecular

intramolecular

ionic solids

London dispersion forces

molar heat of fusion

molar heat of vaporization

molecular solids

normal boiling point

semiconductor



51. boiling point at pressure of 1 atm

52. energy required to melt 1 mole of a substance

53. forces between atoms in a molecule

54. forces between molecules in a solid

55. instantaneous dipole forces for nonpolar molecules

56. lining up of opposite charges on adjacent polar molecules

57. maximum pressure of vapor that builds up in a closed

container

58. mixture of elements having metallic properties overall

59. repeating arrangement of component species in a solid

60. solids that melt at relatively low temperatures

61. Given the densities and conditions of ice, liquid water,

and steam listed in Table 14.1, calculate the volume of

1.0 g of water under each of these circumstances.

62. In carbon compounds a given group of atoms can often be arranged in more than one way. This means

that more than one structure may be possible for the

same atoms. For example, both the molecules diethyl

ether and 1-butanol have the same number of each

type of atom, but they have different structures and

are said to be isomers of one another.

diethyl ether

1-butanol



CH3OCH2OOOCH2OCH3

CH3OCH2OCH2OCH2OOH



471



Which substance would you expect to have the larger

vapor pressure? Why?

63. Which of the substances in each of the following sets

would be expected to have the highest boiling point?

Explain why.

a. Ga, KBr, O2

b. Hg, NaCl, He

c. H2, O2, H2O

64. Which of the substances in each of the following sets

would be expected to have the lowest melting point?

Explain why.

a. H2, N2, O2

b. Xe, NaCl, C (diamond)

c. Cl2, Br2, I2

65. When a person has a severe fever, one therapy to reduce the fever is an “alcohol rub.” Explain how the

evaporation of alcohol from the person’s skin removes heat energy from the body.

66. What is steel?

67. Some properties of potassium metal are summarized

in the following table:

Normal melting point



63.5 °C



Normal boiling point



765.7 °C



Molar heat of fusion



2.334 kJ/mol



Molar heat of vaporization



79.87 kJ/mol



Specific heat of the solid



0.75 J/g °C



a. Calculate the quantity of heat required to heat

5.00 g of potassium from 25.3 °C to 45.2 °C.

b. Calculate the quantity of heat required to melt

1.35 moles of potassium at its normal melting point.

c. Calculate the quantity of heat required to vaporize

2.25 g of potassium at its normal boiling point.

68. What are some important uses of water, both in nature

and in industry? What is the liquid range for water?

69. Describe, on both a microscopic and a macroscopic

basis, what happens to a sample of water as it is

cooled from room temperature to 50 °C below its normal freezing point.

70. Cake mixes and other packaged foods that require

cooking often contain special directions for use at

high elevations. Typically these directions indicate

that the food should be cooked longer above 5000 ft.

Explain why it takes longer to cook something at

higher elevations.

71. Why is there no change in intramolecular forces

when a solid is melted? Are intramolecular forces

stronger or weaker than intermolecular forces?

72. What do we call the energies required, respectively, to

melt and to vaporize 1 mole of a substance? Which of

these energies is always larger for a given substance?

Why?



All even-numbered Questions and Problems have answers in the back of this book and solutions in the Solutions Guide.



472 Chapter 14 Liquids and Solids

73. The molar heat of vaporization of carbon disulfide,

CS2, is 28.4 kJ/mol at its normal boiling point of

46 °C. How much energy (heat) is required to vaporize 1.0 g of CS2 at 46 °C? How much heat is evolved

when 50. g of CS2 is condensed from the vapor to the

liquid form at 46 °C?

74. Which is stronger, a dipole–dipole attraction between

two molecules or a covalent bond between two atoms

within the same molecule? Explain.

75. For a liquid to boil, the intermolecular forces in the

liquid must be overcome. Based on the types of intermolecular forces present, arrange the expected

boiling points of the liquid states of the following

substances in order from lowest to highest: NaCl(l),

He(l), CO(l), H2O(l).

76. What are London dispersion forces and how do they

arise in a nonpolar molecule? Are London forces typically stronger or weaker than dipole–dipole attractions between polar molecules? Are London forces

stronger or weaker than covalent bonds? Explain.

77. Discuss the types of intermolecular forces acting in

the liquid state of each of the following substances.

a.

b.

c.

d.



N2

NH3

He

CO2 (linear, nonpolar)



81. Which type of solid is likely to have the highest melting point—an ionic solid, a molecular solid, or an

atomic solid? Explain.

82. What types of intermolecular forces exist in a crystal

of ice? How do these forces differ from the types of

intermolecular forces that exist in a crystal of solid

oxygen?

83. Discuss the electron sea model for metals. How does

this model account for the fact that metals are very

good conductors of electricity?

84. Water is unusual in that its solid form (ice) is less

dense than its liquid form. Discuss some implications

of this fact.

85. Describe in detail the microscopic processes that take

place when a liquid boils. What kind of forces must

be overcome? Are any chemical bonds broken during

these processes?

86. Water at 100 °C (its normal boiling point) could certainly give you a bad burn if it were spilled on the

skin, but steam at 100 °C could give you a much worse

burn. Explain.

87. What is a dipole–dipole attraction? Give three examples

of liquid substances in which you would expect

dipole–dipole attractions to be large.



78. Explain how the evaporation of water acts as a

coolant for the earth.



88. What is meant by hydrogen bonding? Give three examples of substances that would be expected to exhibit hydrogen bonding in the liquid state.



79. What do we mean when we say a liquid is volatile? Do

volatile liquids have large or small vapor pressures?

What types of intermolecular forces occur in highly

volatile liquids?



89. Although the noble gas elements are monatomic and

could not give rise to dipole–dipole forces or hydrogen bonding, these elements still can be liquefied and

solidified. Explain.



80. Although methane, CH4, and ammonia, NH3, differ

in molar mass by only one unit, the boiling point of

ammonia is over 100 °C higher than that of methane

(a nonpolar molecule). Explain.



90. Describe, on a microscopic basis, the processes of

evaporation and condensation. Which process requires

an input of energy? Why?



All even-numbered Questions and Problems have answers in the back of this book and solutions in the Solutions Guide.



This page intentionally left blank



15

15.1 Solubility

15.2 Solution Composition: An

Introduction

15.3 Solution Composition: Mass

Percent

15.4 Solution Composition: Molarity

15.5 Dilution

15.6 Stoichiometry of Solution

Reactions

15.7 Neutralization Reactions

15.8 Solution Composition:

Normality



Solutions



Seawater is an aqueous solution. (Georgette

Douwma/Getty Images)



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