Kelter P., Mosher M., Scott A. Chemistry. The Practical Science

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348 Chapter 8 Bonding Basics

equatorial position The position of a portion of a mole-

cule when it is arranged along the x axis or y axis of

the molecule. (p. 340)

expanded octet The existence of more than eight elec-

trons in the valence shell of an atom. Atoms of

atomic number greater than 12 are capable of this

feat. In the third row of the periodic table, it typi-

cally occurs in sulfur and phosphorus. (p. 330)

formal charge The difference between the number of

valence electrons on the free atom and the number

of electrons assigned to that atom in the molecule.

(p. 324)

ion pair Ions of opposite charges that exist in solution

as an ionically bonded pair. (p. 308)

ionic bond A strong electrostatic attraction between

ions of opposite charges. (p. 304)

isoelectronic Having the same electron conﬁguration.

(p. 305)

lattice energy The amount of energy released as a solid

ionic crystal is formed. (p. 314)

lattice enthalpy The amount of energy required to sep-

arate 1 mol of a solid ionic crystalline compound

into its gaseous ions. (p. 313)

Lewis dot structure A drawing convention used to

determine the arrangement of atoms in a molecule

or ion. (p. 305)

Lewis dot symbol A drawing convention used to deter-

mine the number of valence electrons on an atom

or monatomic ion. (p. 305)

lone pairs Pairs of electrons that are not involved in

bonding. (p. 322)

metallic bond A bond in which metal cations are

spaced throughout a sea of mobile electrons.

(p. 304)

molecular geometry The geometry of the atoms in a

molecule or ion. The molecular geometry comes

from the electron-group geometry but considers the

lone pairs “invisible.” (p. 337)

molecular models Models used by chemists to explain

the three-dimensional nature of molecules. Although

many can be made from plastic or wood, these

models can also be constructed using computers.

(p. 303)

nonpolar covalent bond A covalent bond with a very

small difference in electronegativity between the

bonded atoms. The result is minimal or no charge

separation in the bond. (p. 320)

nonpolar molecule A molecule in which there is no net

charge separation and no net dipole moment.

(p. 345)

octet rule The existence of eight electrons in the

valence shell of an atom, creating a stable atom or

ion. (p. 306)

pharmacognocist A person who studies natural drugs,

including their biological and chemical components,

botanical sources, and other characteristics. (p. 303)

polar covalent bond A covalent bond in which the atoms

differ in electronegativity, resulting in an unequal

sharing of electrons between two adjacent atoms.

(p. 320)

polar molecule A molecule in which there exists a

charge separation resulting in a net dipole moment.

(p. 345)

polarity The location of a compound on a scale from

polar to nonpolar. The dipole moment is a measure

of the polarity of a compound. (p. 345)

polarization A molecule or bond containing electrons

that are unequally distributed. (p. 319)

polarized A molecule or bond that contains an unequal

distribution of electrons. (p. 319)

radical A molecule that contains an unpaired non-

bonding electron. (p. 331)

resonance hybrid An equal or unequal (based on exper-

imental evidence) combination of all of the reso-

nance structures for a molecule. (p. 328)

resonance structure A model of a molecule in which the

positions of the electrons have changed, but the

positions of the atoms have remained ﬁxed.

(p. 328)

triple bond A covalent bond in which three electron

pairs (six electrons) are shared between adjacent

atoms. (p. 326)

valence shell electron-pair repulsion model A model that

proposes the three-dimensional shapes of molecules

on the basis of the number of electron groups

attached to a central atom. Also known as VSEPR.

(p. 337)

VSEPR The valence shell electron-pair repulsion model.

(p. 337)

zeolite A porous solid with a well-deﬁned structure,

typically made of aluminum and silicon. Zeolites are

capable of binding to ions of a speciﬁc radius based

on the relative size of its pores. (p. 310)

Focus Your Learning 349

The answers to the odd-numbered problems and some selected

problems appear at the back of the book, as represented by the

blue numbering.

Section 8.1 Modeling Bonds

Skill Review

1. Use Lewis electron dot structures to answer these questions

about aluminum and nitrogen.

a. Which neutral atom, Al or N, has the greater number of

unpaired valence electrons?

b. Which neutral atom, Al or N, has the greater number of

valence electrons?

c. Which neutral atom, Al or N, is more likely to gain

electrons to form an octet?

2. Use Lewis electron dot structures to answer these questions

about carbon and sulfur.

a. Which neutral atom, C or S, has the greater number of

unpaired valence electrons?

b. Which neutral atom, C or S, has the greater number of

valence electrons?

c. Which neutral atom, C or S, is more likely to gain electrons

to form an octet?

3. Of the following, which, if any, would not have the same

conﬁguration as a hydrogen atom with two electrons?

a. C

b. B

c. N

3–

d. H

4. Of the following, which, if any, would not have the same con-

ﬁguration as a ﬂuorine atom with eight valence electrons?

a. Ne b. Na

c. Br

−

d. S

2−

5. To which group does each of these ions belong?

a. Ion of an element with a –2 charge and a full octet

b. Ion of an element with a +2 charge and a full octet

c. Ion of an element with a –3 charge and a full octet

6. To which group does each of these ions belong?

a. Ion of an element with a +1 charge and a full octet

b. Ion of an element with a −1 charge and a full octet

c. Atom of an element with no charge and a full octet

7. The charges have been omitted from these Lewis structure

diagrams of ions. From your knowledge of the ground state

for atoms, determine the number of extra electrons that each

structure is showing, and provide the proper charge for each

ion.

a. b. c.

8. The charges have been omitted from these Lewis structure

diagrams of ions. From your knowledge of the ground state

for atoms, determine the number of extra electrons that each

structure is showing, and provide the proper charge for each

ion.

a. b. c.

9. Write Lewis electron dot diagrams for each of these ions.

a. Se

2–

b. I

–

c. Sr

d. Sc

e. Si

10. Write Lewis electron dot diagrams for each of these ions.

a. S

2–

b. Br

–

c. Ti

d. Ti

e. B

NBrO

ClPS

Focus Your Learning

11. Which of these ion(s) would have the same electron conﬁgu-

ration as the noble gas neon?

–

12. Which of these ion(s) would have the same electron conﬁgu-

ration as ﬂuoride (F

−

–

2−

13. Which of these, if any, are isoelectronic?

SArCl

–

14. Which of these, if any, are isoelectronic?

Na Be

Ar F

−

15. Which neutral atoms could be represented by the Lewis dot

symbol shown below?

16. Which cations with a +1 charge could be represented by the

Lewis dot symbol shown below?

Chemical Applications and Practices

17. Lithium, sodium, and potassium are all very reactive alkali

metals.

a. Diagram the Lewis electron dot structure for each neutral

atom.

b. Predict the formula of the oxide compound of each.

18. Calcium, magnesium, and barium are members of the alka-

line earth metals.

a. Diagram the Lewis electron dot structure for each neutral

atom.

b. Predict the formula of the oxide compound of each.

Section 8.2 Ionic Bonding

Skill Review

19. Consider each statement below and determine whether it is

true or false for compounds with ionic bonding.

a. They are typically composed of a nonmetal and a metal.

b. Electrons are shared among atoms in the compound.

c. The metal typically becomes a cation.

d. Like charges repel each other.

20. Consider each statement below and determine whether it is

true or false for compounds with ionic bonding.

a. They require only one metal and one nonmetal.

b. The metal is always written ﬁrst in the formula.

c. The nonmetal must have a –1 charge.

d. The metal and nonmetal must be in the same period.

21. Using the octet rule, explain why aluminum loses three elec-

trons in both aluminum oxide and aluminum chloride, yet

the ratio of aluminum to the nonmetal is 1:3 in aluminum

chloride and 2:3 in aluminum oxide.

22. Using the octet rule, explain why magnesium loses two elec-

trons in both magnesium oxide and magnesium chloride, yet

the ratio of magnesium to the nonmetal is 1:2 in magnesium

chloride and 1:1 in magnesium oxide.

350 Chapter 8 Bonding Basics

23. Within each of the series presented, rank the atoms or ions

from smallest to largest radius.

a. I

–

c. C

–

b. S

2–

d. Fe Fe

24. Within each of the series presented, rank the atoms or ions

from smallest to largest radius.

a. Cu

Cu c. N

−

3−

b. Cr

Cr d. Pd Pd

25. Using Coulomb’s law, arrange these alkali halides from

strongest to weakest attraction of anion and cation: KCl, NaCl,

CsCl, LiCl, RbCl.

26. Using Coulomb’s law, arrange these alkali halides from

strongest to weakest attraction of anion and cation: LiCl, LiF,

LiBr, LiI.

Chemical Applications and Practices

27. Potassium chloride (KCl) has been used as a sodium substi-

tute in some commercial salt (NaCl) products. Show the for-

mation of this important ionic compound using the Lewis

dot symbol model.

28. When some metals oxidize, the product formed may be

detrimental to the strength of the metal structure. This is the

case when iron corrodes. However, aluminum forms an oxide

that resists further reaction and adheres to the metal, form-

ing a protective covering. Show the formation of this ionic

compound using the Lewis dot symbol model.

29. When placed in water, the oxide compounds of many metals

form alkaline, or basic, solutions. This takes place when the

metal oxide reacts with water to produce hydroxide ions. Di-

agram the Lewis dot structure of sodium oxide, and diagram

the structure of the compound that forms when sodium

oxide reacts with water.

30. Calcium oxide is often called lime. Lime has a number of

uses, including reacting with acidic gases from coal plants.

Diagram the Lewis dot structure for this compound, and di-

agram the structure of the hydroxide ion that forms when

lime is placed in water.

31. If Mn

and Mn

were in a solution together, which would

ﬁt into a smaller zeolite hole?

32. Which of these copper ions would ﬁt into a smaller zeolite

hole, Cu

or Cu

33. The melting point of KBr is 734

C. The melting point of

CsBr is 640

C. Which of the two compounds do you predict

to have the higher lattice enthalpy? Explain your answer in

terms of Coulomb’s law.

34. Of NaCl, KCl, and MgCl

, which would you predict to have

the highest lattice enthalpy? Explain your answer in terms of

Coulomb’s law.

Section 8.3 Covalent Bonding

Skill Review

35. The terms electronegativity and electron afﬁnity are both used

to explain some bonding concepts. Compare and contrast

these two important terms.

36. The expression covalent bond is actually used in three differ-

ent contexts within the chapter. Explain the differences

among nonpolar covalent bonds, polar covalent bonds, and

coordinate covalent bonds. Provide an example of a com-

pound that illustrates each of these three types of covalent

bonds.

37. Diagram the outline of a periodic table. Using arrows to

indicate a direction of increase, show the general periodic

trendfor electronegativity withina group and withina period.

38. Diagram the outline of a periodic table. Using arrows to in-

dicate a direction of increase, show the general periodic trend

for atomic size within a group and within a period.

39. Diagram the Lewis structure for each of these neutral com-

pounds. Identify any that contain an odd number of (radical)

electrons

ONO CO NO

HCl PCl

NBr

40. Diagram the Lewis structure for each of these neutral com-

pounds. Identify any that contain an odd number of (radical)

electrons.

OPCl

SeH

41. Diagram the Lewis structure for each of these ions. Identify

any that contain an odd number of (radical) electrons.

–

−

3−

2−

BrO

−

42. Diagram the Lewis structure for each of these ions. Identify

any that contain an odd number of (radical) electrons.

–

−

3−

−

CrO

2−

43. Diagram the Lewis structure for each of these com-

pounds. Identify any that contain an odd number of (radical)

electrons.

44. Diagram the Lewis structure for each of these com-

pounds. Identify any that contain an odd number of (radical)

electrons.

45. It is possible to diagram three resonance structures for the ni-

trate (NO

–

) ion. Show, by giving formal charges, that each

contributes equally to the overall hybrid structure that

depicts bonds of equal strength between the three oxygen

atoms and one nitrogen atom.

46. It is possible to diagram many resonance structures for the

phosphate (PO

3−

) ion. Show, by giving formal charges, that

each contributes equally to the overall hybrid structure that

depicts bonds of equal strength between the four oxygen

atoms and one phosphorus atom.

47. Under the proper conditions, carbon atoms can bond to

form closed ring structures when the carbon atoms are

bonded to each other. These are called cyclic compounds.

Three four-carbon ring compounds (C

, and C

)

are known to exist. Diagram the Lewis dot structure of each

and predict which, if any, may have a resonance structure.

Diagram the resonance structures of any possible results.

Focus Your Learning 351

48. Three straight-chain four-carbon compounds that are not

cyclic (C

, and C

) are known to exist. Diagram

the Lewis dot structure of each and predict which, if any, may

have a resonance structure. Diagram the resonance struc-

tures of any possible results.

49. The compounds N

O, N

, and N

all contain nitrogen to

nitrogen bonds. Diagram the Lewis electron structure of

each. Then match these bond energies for the nitrogen-to-

nitrogen bond in each. (946 kJ/mol, 160 kJ/mol, 418 kJ/mol)

50. In which structure would you expect to ﬁnd the shorter

carbon-to-oxygen bond, carbon monoxide or carbon diox-

ide? Explain the basis of your choice. In which would you

expect to ﬁnd the higher value for bond energy? Explain the

basis of your choice.

51. Each of these compounds contains both ionic and covalent

bonding. Using their Lewis dot structures, indicate where

each type of bonding occurs.

a. Na

b. CaCO

c. Fe(NO

)

52. Each of these compounds contains both ionic and covalent

bonding. Using their Lewis dot structures, indicate where

each type of bonding occurs.

a. NaOH b. NaHCO

c. K

CrO

53. This list gives bonds that could form in compounds. Arrange

them in order from least polar bond to most polar bond.

CaONHOCCaOH

COCNOO

54. This list gives bonds that could form in compounds. Arrange

them in order from least polar bond to most polar bond.

HOFOOPLiOF

SOFNOS

55. a. Calculate the formal charges on the carbon atoms in

ethane (C

b. If one of the hydrogen atoms is replaced with an OH

group, the compound becomes ethanol. What if any

changes in the formal charge of the carbon atoms takes

place when this conversion is made?

c. What is the formal charge on the oxygen atom in ethanol?

d. If the hydrogen atom were then removed from the oxygen

atom, with no other atom taking its place, what would you

calculate as the formal charge for oxygen?

56. Diagram the Lewis dot structure of the sulfate ion (SO

2−

)

without an expanded octet and with an expanded octet.

According to the reasoning about formal charges suggested

in the chapter, which of the two structures is more likely to

represent the actual structure of this important ion?

Chemical Applications and Practices

57. Ammonium chloride (NH

Cl) is used in the manufacture of

some dry cell batteries. Show the structure of this important

ionic battery component using the Lewis dot symbol model.

58. One of the major manufactured compounds containing

iodine is hydrogen iodide. It is used to produce other com-

pounds that contain the iodide anion. Diagram the Lewis dot

structure of this compound.

59. Diagram the Lewis structure for the air pollutant nitrogen

dioxide (NO

). Determine the formal charges on each atom

in the structure. Circle and label any electrons that would be

considered bonding, lone-pair, or radical.

60. Many states are blending ethanol with gasoline to produce

“gasohol” products to use in automobiles. This mixture al-

lows us to extend our dwindling gasoline supplies and im-

prove the octane rating of the gasoline. Diagram the Lewis

structure for this renewable energy extender (C

OH). De-

termine the formal charges on each atom in the structure.

Circle and label any electrons that would be considered

bonding, lone-pair, or radical.

61. The cyanide ion (CN

–

) plays a critical role in reacting metals

from ores in the process of producing pure metals. Diagram

the Lewis structure of this important ion, and determine the

formal charges on both the carbon and nitrogen atoms.

When the ion reacts with H

, hydrocyanic acid is formed. On

the basis of your calculations of formal charges, would it be

more likely for the H

to attach to the carbon side or the

nitrogen side of the ion? Explain your choice.

62. Hypochlorous acid is the acid that can be used to make the

salt sodium hypochlorite that is found in most commercial

bleach preparations. Hypochlorous acid consists of one atom

each of hydrogen, chlorine, and oxygen. Diagram three dif-

ferent arrangements of the atoms in the molecule, and use

formal charge considerations to predict which is most likely.

63. Sulfurous acid (H

) is one of the molecules that con-

tributes to acid rain. The molecule has three oxygen atoms

attached to the central sulfur atom. Diagram the Lewis struc-

ture for the acid, and predict the formal charge on the sulfur

atom. Is the formal charge the same as the oxidation number?

Explain any differences.

64. Dimethyl sulfoxide is a solvent used in some veterinary ap-

plications. The molecule consists of a central sulfur atom

bonded to an oxygen atom and two carbon atoms. Each car-

bon atom has three bonded hydrogen atoms attached. Each

atom follows the Lewis electron dot rules. Diagram the Lewis

structure and determine the formal charge on the sulfur

atom.

65. The concentrations of both nitrite ion (NO

–

) and nitrate ion

(NO

–

) must be monitored in well water. Both can be quite

harmful for humans. Diagram the Lewis dot structures of

both (the nitrogen atom is in the center of both, and there are

no oxygen-to-oxygen bonds), and assign a formal charge to

the nitrogen atom in each.

66. Phosphoric acid (H

) is used in the production of phos-

phate fertilizers and is often found in soft drinks (check the

labels). In the molecule, phosphorus is at the center, and the

hydrogen atoms are attached to the oxygen atoms. Phospho-

rous acid (H

) differs slightly in that one of the hydrogen

atoms is attached to the central phosphorus atom. Diagram

the Lewis electron dot structure of both, and compare the

formal charges on the phosphorus atoms.

67. Of the approximately 20 naturally occurring amino acids,

glycine, the principal component in silk, has the simplest

structure. Show the Lewis dot structure of glycine. (It is com-

posed of two atoms of carbon, one of which is bonded to two

hydrogen atoms, and a nitrogen, which is also bonded to

two hydrogen atoms. The other carbon atom is bonded

to two oxygen atoms, one of which is bonded to an atom of

hydrogen. Examine the structure. Is there a possible reso-

nance structure to draw for the one you have shown? If so,

diagram the other resonance form.

352 Chapter 8 Bonding Basics

68. Oxalic acid (C

) is a toxin found in rhubarb leaves. It

can also be used in general chemistry labs, in its pure form, as

a standard acid with which to react unknown base solutions.

In the structure, the two carbons are connected by a single

bond. Each carbon is connected to two oxygens (one with a

double bond and one with a single bond). The hydrogens are

on opposite sides of the structure. Diagram the structure and

determine whether a possible resonance structure exists for

the compound. If so, draw all of the possible resonance

structures.

69. The three simplest two-carbon hydrocarbon compounds

are the heating fuel ethane (C

), the plant hormone ethene

), and the welding gas ethyne (C

). Diagram the

Lewis structure of each, and predict which would have

the highest bond energy. Which would have the longest

carbon-to-carbon bond?

70. Absorbed light can be sufﬁcient to break chemical bonds.

Compare the bonding of Cl

and O

. In which would the

light absorbed have to be greater in energy? Use Planck’s con-

stant to calculate the frequency of the energy needed to break

apart the molecule you selected. What would the energy need

to be in order to break 1 mol of those bonds?

71. Butane (C

) is the fuel typically used in small, hand-held

lighters to provide heat for other reactions through combus-

tion. Use bond energies to determine the heat of the reaction

between butane and oxygen to produce water and carbon

dioxide. When another determination for the heat of this

reaction was made, using a calorimeter rather than tabular

values, the answer was slightly different. Explain why the two

values may not agree, even though they were obtained for the

same combustion reaction.

72. One way to produce ethanol (CH

OH) is by the reaction

of ethene (C

) and water. If the heat of this reaction was

–37 kJ/mol, what would you calculate as the bond energy for

the carbon double bond in ethene? How does this compare

with the tabular value for the CPC?

73. Using hydrogen gas and oxygen gas, calculate the heat of re-

action both for water and for the bleaching agent hydrogen

peroxide (H

). Judging on the basis of your calculation,

which of the two should be more stable? Explain the basis of

your choice.

74. Small portable tanks often contain propane (C

) as the

fuel to provide heat for camping trips. Use bond energies and

the balanced chemical equation for the combustion of

propane to determine the heat of the reaction between oxy-

gen and propane to produce carbon dioxide and water. On a

per-gram basis, which fuel provides more heat, propane or

butane (C

Section 8.4 VSEPR Theory—A Better Model

Skill Review

75. For those compounds listed in Problem 39 that contain more

than two atoms, determine the electron group and molecular

geometry of the central atoms.

76. For those compounds listed in Problem 40 that contain more

than two atoms, determine the electron group and molecular

geometry of the central atoms.

77. For those compounds listed in Problem 41 that contain more

than two atoms, determine the geometry of the central

atoms.

78. For those compounds listed in Problem 42 that contain more

than two atoms, determine the geometry of the central

atoms.

79. For those compounds listed in Problem 43, determine the

geometry of each carbon atom.

80. For those compounds listed in Problem 44, determine the

geometry of each carbon atom.

81. Use Lewis dot structures and the VSEPR theory to explain the

bond angles in ClO

82. Use Lewis dot structures and the VSEPR theory to explain the

bond angles in SF

83. Use VSEPR modeling to arrange this list in order from largest

to smallest Cl-to-element-to-Cl bond angle:

BeCl

AlCl

CCl

XeCl

NCl

84. Use VSEPR modeling to arrange this list in order from largest

to smallest F-to-element-to-F bond angle:

MgF

Chemical Applications and Practices

85. Carbon tetrachloride follows the “octet rule” and has tetrahe-

dral geometry. However, another tetrachloride compound,

SeCl

, has a different geometry and does not follow the octet

rule. Predict the shape of the second compound, and

diagram its Lewis dot structure.

86. Noble gases already have an octet of valence electrons. There-

fore, any combination with another atom is likely to produce

an exception to the octet rule. Predict the shape of the com-

pound XeF

, and diagram its Lewis dot structure.

87. The compound N

has a double bond between the two

nitrogen atoms. Use the VSEPR model to predict the

HONON bond angle in the molecule. If N

had a

triple bond between the two nitrogen atoms, how would

that affect the same bond angle? Explain the basis of your

prediction.

88. The bicarbonate ion (HCO

–

) is essential as a buffer

(discussed in Chapter 18) in your blood.

a. Diagram the Lewis dot structure for this important ion.

b. Is there a possibility for resonance in the structure? If so,

show the example(s).

c. In VSEPR modeling, what shape would this ion have?

Section 8.5 Properties of

Ionic Compounds and Molecules

Skill Review

89. Examine each of the bonds depicted as dashes between atoms

here. Rank them from least polar to most polar. In addition,

rewrite the bond, showing an arrow in the direction of the

more negative atom.

COCl CONCOO

COHCOBCOMg

Focus Your Learning 353

that would allow this variety in bonding? Diagram the

structure of the N

molecule.

99. When you inhale a breath of air, you are taking in oxygen

and nitrogen molecules. Examine the Lewis dot structure of

an N

molecule (see Problem 98). Suggest a reason why the

that you inhale is unaffected and is soon exhaled as the

same N

that you inhaled.

100. Prepare a diagram, using circles with charges inside, show-

ing a relative size comparison between magnesium oxide

(used as an abrasive) and calcium oxide (used in some plas-

ter formulations). On the basis of your representations, ex-

plain how Coulomb’s potential could be used to predict

which compound contains the larger lattice energy.

101. These circles represent cations and anions. From the choices

provided, select the cation and anion combination that

would have the lowest lattice energy. Explain your choice.

Electrical

Solubility In Conductivity As

Melting

Compound Point Water Hexane Solid Molten

A Very high Yes No No Yes

B Very low No Yes No No

C Medium Yes No No No

94. Refer to the clues in the previous problem and match these

substances as possible identities with A, B, and C.

KCl

citric acid, C

O(COOH)

95. The disinfectant hydrogen peroxide (H

) has a dipole mo-

ment greater than zero. Explain why this information sup-

ports suggesting that the molecule is not linear. If the molecule

were linear, what would you predict about its dipole moment?

96. The formula C

can be drawn many different ways. Pro-

vide the Lewis dot structure that would give a molecule with

the largest dipole moment. Provide the Lewis dot structure

that would give a molecule that did not have a dipole moment.

Comprehensive Problems

97. These diagrams represent a hydrogen atom with different

numbers of electrons. Determine which diagram represents

hydride (H

−

), which hydrogen (H), and which a proton (H

+1 +3 –2

102. a. In your own words explain the lattice energy formula.

b. Lithium forms many useful salts. For example, lithium

ﬂuoride is used extensively in industry to assist in

purifying aluminum. Lithium bromide has applications

in the air–conditioning industry. Explain which of the

two compounds has the stronger lattice energy.

103. On the basis of your arrangement of the ionic compounds

in Problem 101, which compound would you predict to

have the highest lattice enthalpy?

104. a. What is the qualitative relationship between lattice

enthalpy and melting point?

b. On the basis of that relationship, which would you

predict to have a higher melting point, MgF

(lattice

enthalpy = 2913 kJ/mol) or NaCl (lattice enthalpy =

787 kJ/mol)? Use another reference such as the Merck

Index or the Chemical Rubber Company’s Handbook of

Chemistry and Physics to obtain the melting points of

the two compounds to check your assumptions.

105. Most ionic compounds have very high melting points.

a. Use Coulomb’s law to explain this property.

b. Using Table 8.2, select the ionic combination that would

be likely to have the highest melting point and the one

that would be likely to have the lowest melting point.

Explain the basis of both selections.

106. Magnesium metal will burn so vigorously in air that it will

even combine with nitrogen.

a. What is the formula of magnesium nitride?

b. Would the compound be ionic or covalent?

c. Diagram the Lewis dot structure of the compound.

107. a. Explain why the bond energy listed in tables is only an

approximation, based on averages, of the bond energy

for a bond within a molecule.

90. Examine each of the bonds depicted as dashes between atoms

here. Rank the bonds from least polar to most polar. In

addition, rewrite the bond, using the δ+ and δ– symbols to

indicate the polarity of the bond.

HOOHOCHOF

HONHONa HOH

91. Diagram the Lewis structures of both SF

and SF

. Use the

VSEPR model to predict the shape and the resulting polarity

of each.

92. Diagram the Lewis structures of both PBr

and PBr

. Use the

VSEPR model to predict the shape and the resulting polarity

of each.

Chemical Applications and Practices

93. This table illustrates some properties of three unknown sub-

stances. On the basis of the clues provided, decide whether

each compound listed in the table is an ionic compound, a

nonpolar molecule, or a polar molecule.

98. In various compounds, nitrogen can be shown to form sin-

gle, double, or even triple bonds. Diagram the Lewis dot

structure of a nitrogen atom. What arrangement do you see

354 Chapter 8 Bonding Basics

b. In which case would the carbon-to-carbon bond be

expected to deviate more from the average, H

COCH

or H

COCF

? Explain the basis for your reasoning.

108. The environment of an electron in a bond varies consider-

ably depending on the type of bonding arrangement in

which the electron ﬁnds itself. Electrons are attracted to the

positive nucleus of an atom or ion. Using the examples of

covalent, polar covalent, ionic, and metallic bonds, describe

the degree of attraction of an electron to its “original” atom

once it has become part of a bond with another atom.

109. Iron metal can react with oxygen gas to make iron(III)

oxide.

a. Write the balanced equation for this reaction.

b. How many unpaired electrons are in an isolated atom of

iron(III)?

c. What is the mass percent of iron in this compound?

110. People on low-sodium diets sometimes use a salt substitute

for their meals. Suppose the salt substitute contains potas-

sium chloride.

a. What is the electron conﬁguration of potassium ions?

b. Would you expect potassium chloride to taste similar to

sodium chloride? Explain.

c. Which is larger, the potassium cation or the sodium

cation?

d. How many potassium cations are there in 1.00 g of

potassium chloride?

111. Calcium carbonate is a very common mineral found in

nature.

a. Is this compound soluble in water?

b. What is the mass percent carbon in this compound?

c. When hydrochloric acid is added to this compound,

calcium chloride and water are formed.What is the Lewis

dot structure of the other product?

112. A fertilizer salesperson remarks that potassium nitrate pro-

vides more nitrogen per pound than calcium nitrate.

a. Based on the mass percent nitrogen, is the salesperson

correct?

b. How many grams of nitrogen would be present in 1.00 lb

of the potassium nitrate fertilizer?

c. Write the Lewis dot structure for the nitrate ion.

113. The structure of nifedipine, Figure 8.18, contains four

different types of atoms.

a. What is the molar mass of this useful drug?

b. Draw at least one other resonance structure for this

molecule.

c. When this molecule is treated with an acid, a proton

) is added to one of the nitrogen atoms in the

structure. To which nitrogen does that proton bond?

Thinking Beyond the Calculation

114. Someone proposes the use of dichloroethane (C

) as a

solvent to remove small wax (nonpolar hydrocarbon)

buildups.

a. Diagram a way to depict the structure that possesses an

essentially zero dipole moment.

b. Diagram another structure that would yield a dipole

moment.

c. Which of the compounds, that in part a or that in part b,

would be best suited to dissolve wax?

d. Dichloroethane can be prepared by adding chlorine (Cl

)

to ethene (C

). Draw a Lewis dot structure for ethene.

e. Would you predict a nonzero dipole moment for ethene?

f. If 2.0 g of dichloroethane was prepared using the method

in part d, how many grams of ethene was consumed in

the reaction?

355

Contents and Selected Applications

Chemical Encounters: Molecular Structure and Eyesight

9.1 Valence Bond Theory

9.2 Hybridization

9.3 Molecular Orbital Theory

9.4 Putting It All Together

Chemical Encounters: Benzene, Stability, and MO Theory

9.5 Molecular Models in the Chemist’s Toolbox

Advanced

Models of

Bonding

At the back of the human eye are mil-

lions of photoreceptor cells known as

rods and cones. These cells are home

to rhodopsin, a molecule that uses

11-cis-retinal to catch light and initiate a

nerve impulse. The double bonds in

11-cis-retinal enable it to interact with

light in the visible region of the electro-

magnetic spectrum. This chapter will

focus on the construction of models

that help explain this property.

Go to college.hmco.com/pic/kelterMEE for online learning resources.

356

We begin our study of advanced models

of bonding by peering down into a most

remarkable organ, the human eye. Unlike our

hands, our eyes enable us to interact with our environ-

ment without direct contact. Light reﬂected from other

objects enters the eye through a small opening called the iris.

Once there, the photons hit the retina, where they are collected,

converted into electrical impulses, and sent to the brain for processing

(Figure 9.1). The retina contains two types of photoreceptor cells (rods and

cones) that collect the light. The number and type of rods and cones distinguish

the human eye from that of other species.

Over 120 million rods occupy the human retina. These cells are extremely sensi-

tive to light but do not distinguish color. It is these rods that allow humans some

limited nighttime vision. The 7 million cones, which are much less sensitive to

light than the rods, allow humans to perceive color.

How are rods and cones related to molecular structure and advanced bonding

models? It has to do with the way they absorb light. Located in each rod and cone

is a protein called opsin. The protein itself doesn’t absorb visible light. However,

when a molecule known as 11-cis-retinal (see Figure 9.2) is attached to the opsin,

the resulting protein, rhodopsin, becomes able to absorb wavelengths near 500 nm.

A photon that strikes the rhodopsin causes a reaction that converts 11-cis-retinal to

the more stable all-trans-retinal. This results in a change in the position of the atoms

in the retinal molecule, causing the rhodopsin to alter its shape. This change assists

in the generation of a nerve impulse and the release of trans-retinal from the protein.

After the impulse, the trans-retinal is converted back into 11-cis-retinal and attached

to another opsin—ready for the next photon of light to begin the process again (see

Figure 9.3).

What is it about rhodopsin that makes it capable of absorbing photons? Why

does 11-cis-retinal absorb light near 500 nm and not near, say, 850 nm?

Application

HEMICAL

ENCOUNTERS:

Molecular Structure

and Eyesight

Cornea

Pupil

Lens

Iris

Aqueous

humor

Suspensory

ligament

Ciliary

muscles

Vitreous

body

Retina

Fovea

Blind spot

Optic

nerve

Hyaloid

canal

Choroid

Retinal

artery

and vein

FIGURE 9.1

The human eye is a complex organ. Light enters the front

of the eye through the cornea and is focused by the lens

on the retina at the back of the eye.

Wavelength (nm)

Light fraction absorbed

400 500 600 700

There are three kinds of cones in the human eye (they are

known as S, M, and L). About 64% are sensitive to light in the

red region of the visible spectrum, 32% are green-sensitive,

and 2% respond best to blue light.

The answer to these questions largely lies in the structure of the retinal mole-

cule. However, this is best understood when we can examine more advanced

models of bonding. Although the VSEPR model and Lewis dot structures are

useful for determining the shape and connectivity of the atoms within a molecule,

they do not adequately predict many of its properties. For instance, the VSEPR

model of water does predict that the HOOOH bond angle should be less than

109.5°. However, Lewis dot structure and VSEPR models do not properly explain

why the COC bond in ethane (CH

OCH

) is longer than the COC bond in

butadiene (CH

PCHOCHPCH

). Moreover, VSEPR doesn’t explain how we can

end up with equally spaced bonds. For example, in methane (CH

), the bonding

electrons occupy s and p orbitals. VSEPR doesn’t even hint at why electrons in

p orbitals (90° apart) would, in reality, form 109.5° bond angles. In order to

explain the reasons behind the observed bond angles, and to more accurately

assess and predict the properties of molecules, we need to make better models.

Our goal in this chapter is to introduce you to some of the better models.

Advanced Models of Bonding 357

C N Opsin

N Opsin

500 nm

Nerve impulse

Rhodopsin

trans-Retinal

11-cis-Retinal

FIGURE 9.3

Transformation of 11-cis-retinal to trans-

retinal. After the absorption of light and

the reorganization of the double bond,

the rhodopsin molecule undergoes a confor-

mational change that brings about a nerve

impulse. The trans-retinal is recycled by con-

version back to 11-cis-retinal and reattached

to an opsin molecule.

11-cis-Retinal

FIGURE 9.2

11-cis-Retinal absorbs light with wavelengths

near 500 nm. The absorption of light causes an

electron in the double bonds to become excited,

allowing the molecule to rotate about the double

bond in red.