8.f^Chapter 8 Ending^319^325^,,^15659^16786%
SUMMARY

Different chemical bonds give substances different properties. In bonding, atoms tend to lose, gain, or share electrons to achieve a more stable electron configuration, often with eight valence electrons—the octet rule. The tendency is to fill the valence orbitals, so hydrogen usually has two electrons when bonded, and some elements have more than eight.

Ionic bonding involves the complete transfer of one or more electrons from a metal to a nonmetal, often giving both elements a noble-gas configuration. In such compounds, electrostatic forces hold the ions together in crystal lattices. Covalent bonds result from the sharing of a pair of electrons between two atoms. In some molecules, two atoms share two electron pairs to form a double bond or three electron pairs to form a triple bond. Some elements have a stronger tendency to attract electrons (measured by their electronegativity values), and the electrons in a covalent bond are not always shared equally. Such covalent bonds are polar.

Lewis formulas represent the bonding in molecules and ions. In some cases, the actual bonding is a composite of all the resonance structures, and two or more Lewis formulas are needed to show the arrangement of electrons in a covalent molecule or ion. Lewis formulas are also important for determining the shapes of molecules.

VSEPR theory, which assumes that electron pairs stay as far apart as possible, predicts the shapes of molecules. The total number of bonded atoms and of unshared electron pairs on the central atom determines the parent structure: linear, trigonal planar, or tetrahedral. The molecular shape is determined by the position of atoms around the central atom, which is affected by the presence of unshared electron pairs.

A nonpolar molecule has either nonpolar bonds or a symmetrical structure. A molecule is polar only if the bonds are polar and the molecule is not symmetrical.

KEY TERMS
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QUESTIONS AND PROBLEMS

The following questions and problems, except for those in the Additional Questions section, are paired. Questions in a pair focus on the same concept. Answers to the odd-numbered questions and problems are in Appendix D.

Matching Definitions with Key Terms
  • 8.1

    Match the key terms with the descriptions provided.

    1. a covalent bond that involves the sharing of one pair of electrons

    2. a hydrocarbon containing only carbon-carbon single bonds

    3. bonding between two atoms resulting from the sharing of electrons

    4. a solid structure in which the ions are arranged in a regular repeating pattern

    5. a rule stating that atoms tend to gain, lose, or share electrons to achieve an electron configuration with eight electrons in the valence shell

    6. a bond resulting from the unequal sharing of electron pairs

    7. a hydrocarbon containing carbon-carbon triple bonds

    8. a measure of the ability of an atom to attract electrons within a bond to itself

    9. a covalent bond that involves the sharing of three pairs of electrons

    10. the repeating pattern of atoms or ions in a crystal

    11. representation of a compound consisting of the symbols of the component elements, each surrounded by dots representing shared and unshared electrons

    12. force that holds atoms together in a molecule or compound

  • 8.2

    Match the key terms with the descriptions provided.

    1. bonding between cations and anions resulting from electrostatic attractions of opposite charges

    2. a covalent bond that involves the sharing of two pairs of electrons

    3. a group of atoms substituted for hydrogen in the formula of a hydrocarbon that gives the compound its characteristic properties

    4. representation of an atom consisting of the symbol for the element surrounded by a number of dots equal to the number of valence electrons

    5. the angle between the two lines defined by a central atom attached to two surrounding atoms

    6. a bond resulting from the equal sharing of electron pairs

    7. a hydrocarbon containing carbon-carbon double bonds

    8. the separation of electronic charge within a bond (or molecule)

    9. a class of chemical compounds containing only carbon and hydrogen

    10. an average or composite Lewis formula derived from two or more valid Lewis formulas that closely represents the bonding in a molecule

    11. a hydrocarbon derivative containing the –OH group

    12. the shapes of molecules result from the tendency for electron pairs to maximize the distance between them to minimize repulsions

Types of Bonds
  • 8.3

    What is a chemical bond?

  • 8.4

    Describe the difference between ionic and covalent bonding.

  • 8.5

    Which type of elements are most likely to form compounds using ionic bonding?

  • 8.6

    Which type of elements are most likely to form compounds using covalent bonding?

  • 8.7

    Which of the following compounds are likely to have covalent bonds?

    1. HF

    2. NaF

    3. NCl3

    4. MgBr2

    5. CF4

  • 8.8

    Which of the following compounds are likely to have ionic bonds?

    1. CsCl

    2. CaCl2

    3. OCl2

    4. NBr3

    5. IF5

  • 8.9

    Identify the bonding in each of the following substances as ionic or covalent.

    1. CuCl2

    2. F2

    3. CO

    4. FeCl3

  • 8.10

    Identify the bonding in each of the following substances as ionic or covalent.

    1. NaCl

    2. H2

    3. CH4

    4. SiO2

  • 8.11

    Which of the following images represents an ionic compound?

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  • 8.12

    Which of the following images represents a covalent compound?

  • 8.13

    Which of the following compounds are likely to occur as gases at room temperature?

    1. CH4

    2. Ca3N2

    3. SF4

    4. KBr

    5. HCl

  • 8.14

    Which of the following compounds are likely to occur as solids at room temperature?

    1. BrCl

    2. CaCO3

    3. SO2

    4. CaO

    5. NO

  • 8.15

    Predict whether each of the following substances is likely to have a relatively high or low boiling point.

    1. TlCl3

    2. CsCl

    3. CO2

    4. CaO

    5. O2

  • 8.16

    Predict whether each of the following substances is likely to have a relatively high or low melting point.

    1. P4

    2. CaCl2

    3. CCl4

    4. AlN

    5. CaC2

  • 8.17

    Describe how electronegativity values change going down a group in the periodic table.

  • 8.18

    Compare the electronegativity of metallic and nonmetallic elements.

  • 8.19

    What kinds of bonds are always nonpolar?

  • 8.20

    Describe how to decide whether a bond is polar.

  • 8.21

    Using periodic trends, arrange the following atoms in order of increasing electronegativity.

    1. Br, Cl, F, N, O

    2. C, F, H, N, O

  • 8.22

    Using periodic trends, arrange the following atoms in order of increasing electronegativity.

    1. B, C, H, Se, Si

    2. C, Ca, Cl, Cs, Cu

  • 8.23

    Decide which bond in each pair is polar and explain your reasoning. Place δ+ and δ- superscripts on the element symbols for each polar bond to indicate the direction of the polarity.

    1. H—F and H—H

    2. Cl—I and Cl—Cl

    3. H—H and B—H

  • 8.24

    Decide which bond in each pair is polar and explain your reasoning. Place δ+ and δ- superscripts on the element symbols for each polar bond to indicate the direction of the polarity.

    1. F—F and C—F

    2. O—H and H—H

    3. O—O and N—O

  • 8.25

    Arrange the following bonds in order of increasing polarity.

    1. O—H, C—H, H—H, F—H

    2. O—Cl, C—Cl, H—Cl, F—Cl

  • 8.26

    Arrange the following bonds in order of increasing polarity.

    1. H—F, F—F, H—H, H—I

    2. B—F, O—F, C—F, H—F

Ionic Bonding
  • 8.27

    What information can be determined from a Lewis symbol?

  • 8.28

    What is the maximum number of valence electrons represented in the Lewis symbols of the main-group elements?

  • 8.29

    Draw Lewis symbols showing the valence electrons of the following atoms.

    1. C

    2. I

    3. Se

    4. Sr

    5. Cs

    6. Ar

  • 8.30

    Draw Lewis symbols showing the valence electrons of the following atoms.

    1. Kr

    2. Sb

    3. F

    4. In

    5. Ba

    6. N

  • 8.31

    Draw Lewis symbols showing the valence electrons of the following ions.

    1. Cl−

    2. Sc3+

    3. S2−

    4. Ba2+

    5. B3+

  • 8.32

    Draw Lewis symbols showing the valence electrons of the following ions.

    1. N3−

    2. C4−

    3. Br−

    4. Mg2+

    5. Al3+

  • 8.33

    Write a formula for each of the following ionic salts using the Lewis symbols of the ions.

    1. LiCl

    2. BaCl2

    3. BaS

  • 8.34

    Write a formula for each of the following ionic salts using the Lewis symbols of the ions.

    1. Na2O

    2. LiOH

    3. CsF

  • 8.35

    Explain why K+ is known in nature, but K2+ is not.

  • 8.36

    Explain why Cl− is known in nature, but Cl2− is not.

  • 8.37

    Explain why sodium fluoride has the composition represented by the formula NaF.

  • 8.38

    Explain why calcium oxide has the composition represented by the formula CaO.

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  • 8.39

    What is the difference between an ionic crystal and a crystal lattice?

  • 8.40

    What holds ions together in a crystal lattice?

  • 8.41

    Describe the sodium chloride structure shown in Figure 8.8. How many chloride ions surround each sodium ion?

  • 8.42

    Describe the cesium chloride structure shown in Figure 8.9. How many chloride ions surround each cesium ion?

  • 8.43

    Why does CaF2 have a different crystal structure than NaCl?

  • 8.44

    Would you expect LiCl to have the same crystal structure as NaCl or CsCl? Explain your selection.

Covalent Bonding
  • 8.45

    Draw the Lewis formulas for O2 and F2.

    1. How many valence electrons are around each atom in each of these formulas?

    2. Describe the bonding in each as a single, double, or triple covalent bond.

  • 8.46

    Draw the Lewis formulas for I2 and N2.

    1. How many valence electrons are around each atom in each of these formulas?

    2. Describe the bonding in each as a single, double, or triple covalent bond.

  • 8.47

    Why does hydrogen exist as a diatomic molecule?

  • 8.48

    How many electrons does each hydrogen have in the H2 molecule?

  • 8.49

    How many single bonds are typically formed by the following atoms?

    1. H

    2. N

    3. F

    4. Ne

  • 8.50

    How many single bonds are typically formed by the following atoms?

    1. B

    2. O

    3. I

    4. C

  • 8.51

    Identify what main-group element (X) could form each of the following compounds.

  • 8.52

    Identify what main-group element (X) could form each of the following compounds.

  • 8.53

    Write a Lewis formula for each of the following:

    1. HCN

    2. H3CCN

    3. C2H2

    4. C2H4

    5. C2H6

  • 8.54

    Write a Lewis formula for each of the following:

    1. NH2OH

    2. CCl4

    3. C2H3Cl

    4. C2Br2

    5. HOCl

  • 8.55

    Write a Lewis formula for each of the following:

    1. NO3−

    2. SO42−

    3. SO32−

    4. NO2−

    5. NO+

  • 8.56

    Write a Lewis formula for each of the following:

    1. NH4+

    2. H2CO

    3. (CH3)2CO

    4. OH−

    5. CH3NO2

  • 8.57

    Determine whether each of the following has the same number of bonding electrons as the cyanide ion, CN−.

    1. O2

    2. NO+

    3. CO

    4. N2

    5. NH3

  • 8.58

    Determine whether each of the following has the same number of bonding electrons as the sulfur dioxide molecule, SO2.

    1. NO2−

    2. OF2

    3. CO2

    4. O3

    5. N2

  • 8.59

    When is it necessary to use the concept of resonance?

  • 8.60

    How is the concept of resonance consistent with the octet rule?

  • 8.61

    Indicate whether or not each of the following molecules or ions exhibits resonance.

    1. O2

    2. H2O

    3. SO2

    4. NO2

    5. SO32−

  • 8.62

    Indicate whether or not each of the following molecules or ions exhibits resonance.

    1. N2

    2. F2O

    3. ClO2−

    4. CO2

    5. SO3

  • 8.63

    Write a Lewis formula, including the resonance forms, for each of the following molecules or ions.

    1. NO2−

    2. SO3

    3. CO2

    4. CO32−

    5. HNO3−

  • 8.64

    Write a Lewis formula, including the resonance forms, for each of the following molecules or ions.

    1. CS2

    2. NCO−

    3. HNO2

    4. SO42−

    5. SO32−

  • 8.65

    In HF, the hydrogen atom shares two electrons with the fluorine atom, but has no unshared electron pairs. Discuss why this observation is consistent with the principle behind the octet rule.

  • 8.66

    Describe the bonding in S2Cl2. The atoms are connected in the order, Cl S S Cl. Why does this molecule not exist as SCl?

  • 8.67

    Decide whether the indicated atom obeys the octet rule. If not, indicate how the octet rule is broken.

    1. O in H2O

    2. S in SF4

    3. F in SF4

    4. S in SF2

  • 8.68

    Decide whether the indicated atom obeys the octet rule. If not, indicate how the octet rule is broken.

    1. B in BCl3

    2. Cl in ClF2− ion

    3. F in F2

    4. S in SF6

  • 8.69

    An atom in each of the following molecules does not obey the octet rule. Decide which atom violates the rule and explain the nature of the violation.

    1. SF4

    2. BH3

    3. XeF6

    4. ClO2

  • 8.70

    An atom in each of the following molecules does not obey the octet rule. Decide which atom violates the rule and explain the nature of the violation.

    1. NO2

    2. XeF2

    3. BeCl2

    4. ICl3

  • 8.71

    What are the different classes of hydrocarbons?

  • 8.72

    How many bonds does carbon normally form? What are the different ways it can form this many bonds?

  • 8.73

    Draw the Lewis formula of benzene, C6H6, a cyclic compound.

  • 8.74

    Draw the Lewis formula of ethylene, C2H4.

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  • 8.75

    Identify the class of organic substance for each of the following molecules.

    1. CH3—OH

    2. H3C—CH3

    3. H3C—O—CH3

    4. H3C—CH==CH2

  • 8.76

    Identify the class of organic substance for each of the following molecules.

    1. H3C—NH2

    2. H2C==CH2

    5. C6H6

  • 8.77

    Identify the class of organic substance for each of the following molecules.

    1. CH3CH2OH

    2. CH3CCH

    3. CH3CH2CHO

    4. CH3CH2OCH2CH3

  • 8.78

    Identify the class of organic substance for each of the following molecules.

    1. (CH3)2NH

    2. C2H2

    3. CH3CH2CO2H

    4. CH3CH2COCH3

  • 8.79

    Identify the class of organic substance for each of the following molecules.

  • 8.80

    Identify the class of organic substance for each of the following molecules.

  • 8.81

    Draw an aldehyde that has the molecular formula C4H8O.

  • 8.82

    Draw a ketone that has the molecular formula C4H8O.

Shapes of Molecules
  • 8.83

    How can VSEPR theory be used to predict molecular shapes?

  • 8.84

    Why are unshared pairs of electrons on a central atom not considered to be part of the molecular shape?

  • 8.85

    Why is it important to draw Lewis structures before predicting molecular shapes?

  • 8.86

    Explain how nonbonding pairs of electrons influence molecular shape.

  • 8.87

    Draw each of the following geometric arrangements.

    1. tetrahedral

    2. trigonal planar

    3. bent

    4. trigonal pyramidal

  • 8.88

    In which of the following molecular shapes would you expect to find one or more unshared pairs of electrons on the central atom?

    1. bent

    2. tetrahedral

    3. trigonal planar

    4. trigonal pyramidal

  • 8.89

    Predict the parent structures of the following molecules.

    1. BeCl2

    2. PH3

    3. SCl2

    4. SO2

    5. H2Te

    6. SiH4

    7. BBr3

    8. H2O

  • 8.90

    Predict the parent structures of the following molecules.

    1. OCS

    2. FNO

    3. FCN

    4. HN3

    5. PF3

    6. SF2

    7. NO2

  • 8.91

    Predict the shapes and give approximate bond angles for the following molecules.

    1. BeCl2

    2. PH3

    3. SCl2

    4. SO2

    5. H2Te

    6. SiH4

    7. BBr3

    8. H2O

  • 8.92

    Predict the shapes and give approximate bond angles for the following molecules.

    1. OCS

    2. FNO

    3. FCN

    4. HN3

    5. PF3

    6. SF2

    7. NO2

  • 8.93

    Predict the bond angles in the following molecules.

    1. NH3

    2. H2O

    3. HCl

    4. HCN

    5. BF3

    6. H2CO

    7. PCl3

  • 8.94

    Predict the bond angles in the following molecules.

    1. NF3

    2. SO2Cl2

    3. CBr4

    4. F2CO

    5. PH3

    6. HOCl

  • 8.95

    Give an example of a molecule or polyatomic ion that has the following features.

    1. three bonded atoms, no unshared electrons on the central atom

    2. three bonded atoms, one unshared pair of electrons on the central atom

    3. two bonded atoms, two unshared pairs of electrons on the central atom

  • 8.96

    Give an example of a molecule or polyatomic ion that has the following features.

    1. four bonded atoms, no unshared electrons on the central atom

    2. two bonded atoms, no unshared electrons on the central atom

    3. two bonded atoms, one unshared pair of electrons on the central atom

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  • 8.97

    Identify a molecule or ion that could have the following structures.

  • 8.98

    Identify a molecule or ion that could have the following structures.

  • 8.99

    Is this the structure of NO3− or ClO3−?

  • 8.100

    Is this the structure of SCl2 or BeCl2?

  • 8.101

    Which of the following molecules or ions have unshared electron pairs on the central atom?

  • 8.102

    Which of the following molecules or ions have unshared electron pairs on the central atom?

  • 8.103

    Hydrazine, N2H4, is a colorless, oily liquid that fumes in air and has an odor much like that of ammonia. It is used as a rocket fuel. The order of its atoms is H2NNH2. How many unshared pairs of electrons are on each nitrogen atom?

  • 8.104

    Oxalic acid, H2C2O4, a poisonous colorless solid, is found in some vegetables such as spinach and rhubarb. It is present in concentrations well below the toxic limit, so you can't use this as a reason to refuse a helping of spinach. The order of atoms in a molecule of oxalic acid is HO2CCO2H. How many unshared pairs of electrons are on the carbon and oxygen atoms?

  • 8.105

    Chloropicrin, Cl3CNO2, is an insecticide that has been used against insects that attack cereals and grains. It is a liquid with an intense odor. Predict the Cl—C—Cl, Cl—C—N, C—N—O, and O—N—O bond angles in a molecule of chloropicrin.

  • 8.106

    Fuel cells are used in many areas, such as the aerospace industry, where energy efficiency is more important than high power output. One fuel cell uses methanol, CH3OH, as a fuel. Predict the H—C—H, H—C—O, and C—O—H bond angles.

  • 8.107

    Distinguish between bond polarity and molecular polarity.

  • 8.108

    Why does molecular polarity depend not only on bond polarity but also on the geometry of the molecule?

  • 8.109

    Explain how carbon tetrachloride can have polar bonds but still be a nonpolar molecule.

  • 8.110

    Explain why hydrocarbons are all essentially nonpolar substances.

  • 8.111

    Which of the following molecules are polar?

    1. HI

    2. CHF3

    3. SO2Cl2

    4. PF3

  • 8.112

    Which of the following molecules are polar?

    1. BF3

    3. CH2Cl2

  • 8.113

    For each pair of molecules decide which molecule is polar and explain why it is polar while the other is not.

    1. SO2, CO2

    2. SO2, SO3

    3. SeCl2, BeCl2

    4. CH4, CH3I

  • 8.114

    Explain why the first molecule of each pair is polar and the second is not.

    1. CH2Cl2, CCl4

    2. PF3, BF3

    3. BF2Cl, BF3

    4. SO2, O3

  • 8.115

    Which of the following molecules would align with an electric field if placed in a capacitor? Assume the bonds are polar.

  • 8.116

    Which of the following molecules would align with an electric field if placed in a capacitor? Assume the bonds are polar.

  • 8.117

    Which molecule, CF4 or CCl2F2, is most likely to be soluble in water?

  • 8.118

    Which substance, SO2 or CO2, is likely to be more soluble in water?

  • 8.119

    Which of these molecules is polar? Assume the bonds are polar.

  • 8.120

    Which of these molecules is polar? Assume the bonds are polar.

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Additional Questions
  • 8.121

    Draw Lewis symbols showing the valence electrons of the following atoms.

    1. Br

    2. Pb

    3. S

    4. Ca

    5. Be

    6. Xe

  • 8.122

    Draw Lewis symbols showing the valence electrons of the following ions.

    1. P3−

    2. In3+

    3. Se2−

    4. Be2+

    5. C4−

  • 8.123

    Arrange the following atoms in order of decreasing electronegativity: Br, Cl, F, I.

  • 8.124

    Which of the following substances is likely to be ionic?

    1. H2

    2. Li2O

    3. BCl3

    4. ClBr

    5. SiO2

  • 8.125

    Classify each of the following substances according to the primary type of bonding.

    1. HCN

    2. AgCl

    3. S8

    4. CH4

    5. CoCl2

  • 8.126

    Write a Lewis formula for each of the following.

    1. HClO2

    2. HClO3

    3. HClO4

    4. BrF3

    5. ClO3−

  • 8.127

    Write a Lewis formula for each of the following molecules.

    1. N2H2

    2. CS2

    3. AsF3

    4. CO2

    5. CO

  • 8.128

    Write a Lewis formula, including the resonance forms, for each of the following molecules or ions.

    1. OCN−

    2. N3−

    3. ClO2−

    4. PO43−

    5. H2CO3

  • 8.129

    Draw the best Lewis formula for each of the following. Include any necessary resonance forms.

    1. ClO4−

    2. NO2−

    3. NCO−

    4. HCO2−

    5. BF3

  • 8.130

    Gaseous aluminum chloride exists as a dimer, Al2Cl6, with two chlorines bridging between the two aluminum atoms: Cl2AlCl2AlCl2. Draw a Lewis structure for Al2Cl6 and for AlCl3. Discuss a reason why AlCl3 forms Al2Cl6.

  • 8.131

    Describe the molecular shape of the following.

    1. SiCl4

    2. GaCl3

    3. NCl2+

    4. IO3−

    5. PCl4+

    6. OF2

    7. GeH4

    8. SOCl2

    9. Br2O

    10. ClO2−

  • 8.132

    Describe the structure and bonding in sulfuric acid, H2SO4, and in its two ions, HSO4− and SO42−.

  • 8.133

    Decide which of each pair of gaseous molecules is polar, and explain why that molecule is polar while the other is not.

    1. BeCl2, OCl2

    2. PH3, BH3

    3. BCl3, AsCl3

    4. SiH4, NH3

  • 8.134

    Which of the following are nonpolar molecules, even though they have polar bonds?

    1. CH2Cl2

    2. SiCl4

    3. SCl2

    4. ClO2

    5. PBr3

  • 8.135

    For each pair of molecules decide which molecule is most polar and explain your answer.

    1. CCl4, CH2Cl2

    2. CH3F, CH3Br

    3. NF3, NH3

    4. OF2, H2O

  • 8.136

    Based on predicted bond angles from VSEPR theory, why is the ozone molecule, O3, not a cyclic molecule?