Chemistry A Molecular Approach

Nivaldo J. Tro

Chapter 10 Chemical Bonding ll: Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory - all with Video Answers

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Chapter Questions

01:27

Problem 1

Why is molecular geometry important? Cite some examples.

Kara Cecil

Numerade Educator

01:09

Problem 2

According to VSEPR theory, what determines the geometry of a molecule?

Eugene Schneider

University of Minnesota - Twin Cities

04:15

Problem 3

Name and sketch the five basic electron geometries, and state the number of electron groups corresponding to each. What constitutes an electron group?

Kara Cecil

Numerade Educator

01:24

Problem 4

Explain the difference between electron geometry and molecular geometry. Under what circumstances are they not the same?

Eugene Schneider

University of Minnesota - Twin Cities

11:44

Give the correct electron and molecular geometries that correspond to each set of electron groups around the central atom of a molecule.
\begin{equation}
\begin{array}{l}{\text { a. four electron groups overall; three bonding groups and one }} \\ {\text { lone pair }} \\ {\text { b. four electron groups overall; two bonding groups and two }} \\ {\text { lone pairs }} \\ {\text { c. five electron groups overall; four bonding groups and one }} \\ {\text { d. five electron groups overall; three bonding groups and two }} \\ {\text { lone pairs }}\\{\text { e. five electron groups overall; two bonding groups and three }} \\ {\text { lone pairs }} \\ {\text { f. } \text { six electron groups overall; five bonding groups and one }} \\ {\text { lone pair }} \\ {\text { g. six electron groups overall; four bonding groups and two }} \\ {\text { lone pairs }}\end{array}
\end{equation}

Kara Cecil

Numerade Educator

00:38

Problem 6

How do you apply VSEPR theory to predict the shape of a molecule with more than one interior atom?

Eugene Schneider

University of Minnesota - Twin Cities

02:39

Problem 7

How do you determine whether a molecule is polar? Why is polarity important?

Kara Cecil

Numerade Educator

00:54

Problem 8

What is a chemical bond according to valence bond theory?

Eugene Schneider

University of Minnesota - Twin Cities

01:31

Problem 9

In valence bond theory, what determines the geometry of a molecule?

Kara Cecil

Numerade Educator

01:10

Problem 10

In valence bond theory, the interaction energy between the electrons and nucleus of one atom with the electrons and nucleus of another atom is usually negative (stabilizing) when ___________.

Eugene Schneider

University of Minnesota - Twin Cities

01:52

Problem 11

What is hybridization? Why is hybridization necessary in valence bond theory?

Kara Cecil

Numerade Educator

01:05

Problem 12

How does hybridization of the atomicorbitals in the central atom of a molecule help lower the overall energy of the molecule?

Eugene Schneider

University of Minnesota - Twin Cities

01:42

Problem 13

How is the number of hybrid orbitals related to the number of standard atomic orbitals that are hybridized?

Kara Cecil

Numerade Educator

02:18

Problem 14

Sketch each set of hybrid orbitals.
\begin{equation}\begin{array}{lll}{\text { a. } s p} & {\text { b. } s p^{2}} & {\text { c. } s p^{3}} \\ {\text { d. } s p^{3} d} & {\text { e. }} {s p^{3} d^{2}}\end{array}\end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

01:38

Problem 15

In the Lewis model, the two bonds in a double bond look identical. However, valence bond theory shows that they are not. Describe a double bond according to valence bond theory. Explain why rotation is restricted about a double bond but not about a single bond.

Lottie Adams

Numerade Educator

02:04

Problem 16

Name the hybridization scheme that corresponds to each electron geometry.
\begin{equation}\begin{array}{lll}{\text { a. }} {\text { linear }} & {\text { b. }} & {\text { trigonal planar }} \\ {\text { c. tetrahedral }} & {\text { d. }} & {\text { trigonal bipyramidal }} \\ {\text { e. }} {\text { octahedral }}\end{array}\end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

00:23

Problem 17

What is a chemical bond according to molecular orbital theory?

Joshua Speer

Numerade Educator

01:13

Problem 18

Explain the difference between hybrid atomic orbitals in valence bond theory and LCAO molecular orbitals in molecular orbital theory.

Eugene Schneider

University of Minnesota - Twin Cities

00:43

Problem 19

What is a bonding molecular orbital?

Joshua Speer

Numerade Educator

00:47

Problem 20

What is an antibonding molecular orbital?

Eugene Schneider

University of Minnesota - Twin Cities

01:18

Problem 21

What is the role of wave interference in determining whether a molecular orbital is bonding or antibonding?

Joshua Speer

Numerade Educator

00:49

Problem 22

In molecular orbital theory, what is bond order? Why is it important?

Eugene Schneider

University of Minnesota - Twin Cities

01:27

Problem 23

How is the number of molecular orbitals approximated by a linear combination of atomic orbitals related to the number of atomic orbitals used in the approximation?

Joshua Speer

Numerade Educator

01:23

Problem 24

Sketch each molecular orbital.
\begin{equation}\begin{array}{llll}{\text { a. }} & {\sigma_{2 s}} & {\text { b. }} & {\sigma_{2 s}^{*}} & {\text { c. } \sigma_{2 p}} \\ {\text { d. }} & {\sigma_{2 p}^{*}} & {\text { e. }} & {\pi_{2 p}} & {\text { f. }} {\pi_{2 p}^{*}}\end{array}\end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

05:38

Problem 25

Draw an energy diagram for the molecular orbitals of period 2 diatomic molecules. Show the difference in ordering for $B_{2}, C_{2}$ and $N_{2}$ compared to $\mathrm{O}_{2}, \mathrm{F}_{2},$ and $\mathrm{Ne}_{2} .$

Lottie Adams

Numerade Educator

01:18

Problem 26

Why does the energy ordering of the molecular orbitals of the period 2 diatomic molecules change in going from $\mathrm{N}_{2}$ to $\mathrm{O}_{2} ?$

Eugene Schneider

University of Minnesota - Twin Cities

00:35

Problem 27

Explain the difference between a paramagnetic species and a diamagnetic one.

Joshua Speer

Numerade Educator

01:10

Problem 28

When applying molecular orbital theory to heteronuclear diatomic molecules, the atomic orbitals used may be of different energies. If two atomic orbitals of different energies make two molecular orbitals, how are the energies of the molecular orbitals related to the energies of the atomic orbitals? How is the shape of the resultant molecular orbitals related to the shape of the atomic orbitals?

Eugene Schneider

University of Minnesota - Twin Cities

00:36

Problem 29

In molecular orbital theory, what is a nonbonding orbital?

Joshua Speer

Numerade Educator

02:00

Problem 30

Write a short paragraph describing chemical bonding according to the Lewis model, valence bond theory, and molecular orbital theory. Indicate how the theories differ in their description of a chemical bond and indicate the strengths and weaknesses of each theory. Which theory is correct?

Eugene Schneider

University of Minnesota - Twin Cities

02:03

Problem 31

A molecule with the formula AB $_{3}$ has a trigonal pyramidal geometry. How many electron groups are on the central atom (A)?

Kara Cecil

Numerade Educator

00:26

Problem 32

A molecule with the formula $A B_{3}$ has a trigonal planar geometry. How many electron groups are on the central atom?

Eugene Schneider

University of Minnesota - Twin Cities

02:38

Problem 33

For each molecular geometry, list the number of total electron groups, the number of bonding groups, and the number of lone pairs on the central atom.

Kara Cecil

Numerade Educator

02:12

Problem 34

For each molecular geometry, list the number of total electron groups, the number of bonding groups, and the number of lone pairs on the central atom.

Eugene Schneider

University of Minnesota - Twin Cities

01:35

Problem 35

Determine the electron geometry, molecular geometry, and idealized bond angles for each molecule. In which cases do you expect deviations from the idealized bond angle?
\begin{equation}\text { a. } \mathrm{PF}_{3} \quad \text { b. } \mathrm{SBr}_{2} \quad \text { c. } \mathrm{CHCl}_{3} \quad \text { d. } \mathrm{CS}_{2}\end{equation}

Lottie Adams

Numerade Educator

02:59

Problem 36

Determine the electron geometry, molecular geometry, and idealized bond angles for each molecule. In which cases do you expect deviations from the idealized bond angle?
\begin{equation}\text { a. }\mathrm{CF}_{4} \quad \text { b. } \mathrm{NF}_{3} \quad \text { c. } \mathrm{OF}_{2} \quad \text { d. } \mathrm{H}_{2} \mathrm{S}\end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

00:39

Problem 37

Which species has the smaller bond angle, $\mathrm{H}_{3} \mathrm{O}^{+}$ or $\mathrm{H}_{2} \mathrm{O}$ ?Explain.

Lottie Adams

Numerade Educator

01:12

Problem 38

Which species has the smaller bond angle, $\mathrm{ClO}_{4}^{-}$ or $\mathrm{ClO}_{3}^{-2}$ ? Explain.

Eugene Schneider

University of Minnesota - Twin Cities

07:48

Problem 39

Determine the molecular geometry and sketch each molecule or ion using the bond conventions shown in "Representing Molecular Geometries on Paper" in Section 10.4.
\begin{equation}\text { a. }\mathrm{SF}_{4} \quad \text { b. } \mathrm{ClF}_{3} \quad \text { c. } \mathrm{IF}_{2}^{-} \quad \text { d. } \operatorname{IBr}_{4}^{-}\end{equation}

Tom Rutherford

Numerade Educator

02:00

Problem 40

Determine the molecular geometry and sketch each molecule or ion, using the bond conventions shown in "Representing Molecular Geometries on Paper" in Section 10.4.
\begin{equation}\text { a. } \mathrm{BrF}_{5} \quad \text { b. } \mathrm{SCl}_{6} \quad \text { c. } \mathrm{PF}_{5} \quad \text { d. } \mathrm{IF}_{4}^{+}\end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

01:52

Problem 41

Determine the molecular geometry about each interior atom and sketch each molecule.
\begin{equation}\begin{array}{l}{\text { a. } \mathrm{C}_{2} \mathrm{H}_{2} \text { (skeletal structure HCCH) }} \\ {\text { b. } \mathrm{C}_{2} \mathrm{H}_{4} \text { (skeletal structure } \mathrm{H}_{2} \mathrm{CCH}_{2} )} \\ {\text { c. } \mathrm{C}_{2} \mathrm{H}_{6} \text { (skeletal structure } \mathrm{H}_{3} \mathrm{CCH}_{3} )}\end{array}\end{equation}

Lottie Adams

Numerade Educator

01:20

Problem 42

Determine the molecular geometry about each interior atom and sketch each molecule.
\begin{equation}\begin{array}{l}{\text { a. } \mathrm{N}_{2} \quad \text { b. } \mathrm{N}_{2} \mathrm{H}_{2}(\text { skeletal structure HNNN })} \\ {\text { c. } \mathrm{N}_{2} \mathrm{H}_{4}\left(\text { skeletal structure } \mathrm{H}_{2} \mathrm{NNH}_{2}\right)}\end{array}\end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

01:13

Problem 43

Each ball-and-stick model shows the electron and molecular geometry of a generic molecule. Explain what is wrong with each molecular geometry and provide the correct molecular geometry, given the number of lone pairs and bonding groups on the central atom.

Lottie Adams

Numerade Educator

02:00

Problem 44

Eugene Schneider

University of Minnesota - Twin Cities

01:48

Problem 45

Lottie Adams

Numerade Educator

03:40

Problem 46

Determine the geometry about each interior atom in each molecule and sketch the molecule. (Skeletal structure is indicated in parentheses.)
\begin{equation}\begin{array}{l}{\text { a. } \mathrm{CH}_{3} \mathrm{NH}_{2}\left(\mathrm{H}_{3} \mathrm{CNH}_{2}\right)} \\ {\text { b. } \mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{CH}_{3}\left(\mathrm{H}_{3} \mathrm{CCOOCH}_{3} \text { One } \mathrm{O} \text { atom attached to } 2 \mathrm{nd} \mathrm{C}\right.} \\ {\text { atom; the other O atom is bonded to the } 2 \mathrm{nd} \text { and } 3 \mathrm{rd} \text { Catom } )} \\ {\text { c. } \mathrm{NH}_{2} \mathrm{CO}_{2} \mathrm{H}\left(\mathrm{H}_{2} \mathrm{NCOOH} \text { both } \mathrm{O} \text { atoms attached to } \mathrm{C}\right)}\end{array}\end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

01:16

Problem 47

Explain why $\mathrm{CO}_{2}$ and $\mathrm{CCl}_{4}$ are both nonpolar even though they contain polar bonds.

Kara Cecil

Numerade Educator

01:22

Problem 48

$\mathrm{CH}_{3} \mathrm{F}$ is a polar molecule, even though the tetrahedral geometry often leads to nonpolar molecules. Explain.

Eugene Schneider

University of Minnesota - Twin Cities

03:27

Problem 49

Determine whether each molecule in Exercise 35 is polar or nonpolar.

Lottie Adams

Numerade Educator

01:26

Problem 50

Determine whether each molecule in Exercise 36 is polar or nonpolar.

Eugene Schneider

University of Minnesota - Twin Cities

02:30

Problem 51

Determine whether each molecule is polar or nonpolar.
\begin{equation}\text { a. }\mathrm{SCl}_{2} \quad \text { b. } \mathrm{SCl}_{4} \quad \text { c. } \mathrm{BrCl}_{5}\end{equation}

Lottie Adams

Numerade Educator

02:19

Problem 52

Determine whether each molecule is polar or nonpolar.
\begin{equation}\text { a. }\mathrm{SCl}_{4} \quad \text { b. } \mathrm{CF}_{2} \mathrm{Cl}_{2} \quad \text { c. } \mathrm{SeF}_{6} \quad \text { d. } \mathrm{IF}_{5}\end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

01:13

Problem 53

The valence electron configurations of several atoms are shown here. How many bonds can each atom make without hybridization?
\begin{equation}\text { a. } \mathrm{Be} \ 2 s^{2} \quad \text { b. } \mathrm{P} 3 s^{2} 3 p^{3} \quad \text { c. } \mathrm{F} \ 2 \mathrm{s}^{2} 2 p^{5}\end{equation}

Lottie Adams

Numerade Educator

01:06

Problem 54

The valence electron configurations of several atoms are shown here. How many bonds can each atom make without hybridization?
\begin{equation}\text { a. }\mathrm{B} \ 2 \mathrm{s}^{2} 2 p^{1} \quad \text { b. } \mathrm{N} \ 2 \mathrm{s}^{2} 2 \mathrm{p}^{3} \quad \text { c. } \mathrm{O} \ 2 \mathrm{s}^{2} 2 \mathrm{p}^{4}\end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

01:12

Problem 55

Write orbital diagrams (boxes with arrows in them) to represent the electron configurations-without hybridization-for all the atoms in $\mathrm{PH}_{3}$ . Circle the electrons involved in bonding. Draw a three-dimensional sketch of the molecule and show orbital overlap. What bond angle do you expect from the unhybridized orbitals? How well does valence bond theory agree with the experimentally measured bond angle of $93.3^{\circ} ?$

Lottie Adams

Numerade Educator

02:56

Problem 56

Write orbital diagrams (boxes with arrows in them) to represent the electron configurations-without hybridization-for all the atoms in SF. Circle the electrons involved in bonding. Draw a three dimensional sketch of the molecule and show orbital overlap. What bond angle do you expect from the unhybridized orbitals? How well does valence bond theory agree with the experimentally measured bond angle of $98.2^{\circ} ?$

Eugene Schneider

University of Minnesota - Twin Cities

07:14

Problem 57

Write orbital diagrams (boxes with arrows in them) to represent the electron configuration of carbon before and after $s p^{3}$ hybridization.

Connor Siggins

Numerade Educator

00:58

Problem 58

Write orbital diagrams (boxes with arrows in them) to represent the electron configurations of carbon before and after $s p$ hybridization.

Eugene Schneider

University of Minnesota - Twin Cities

01:05

Problem 59

Which hybridization scheme allows the formation of at least one $\pi$ bond?
$$s p^{3}, s p^{2}, s p^{3} d^{2}$$

Lottie Adams

Numerade Educator

00:42

Problem 60

Which hybridization scheme allows the central atom to form more than four bonds?
$$s p^{3}, s p^{3} d, s p^{2}$$

Eugene Schneider

University of Minnesota - Twin Cities

03:36

Problem 61

Write a hybridization and bonding scheme for each molecule. Sketch the molecule, including overlapping orbitals, and label all bonds using the notation shown in Examples 10.6 and 10.7.
$$\text { a. }\mathrm{CCl}_{4} \quad \text { b. } \mathrm{NH}_{3} \quad \text { c. } \mathrm{OF}_{2} \quad \text { d. } \mathrm{CO}_{2}$$

Lottie Adams

Numerade Educator

04:55

Problem 62

Write a hybridization and bonding scheme for each molecule. Sketch the molecule, including overlapping orbitals, and label all bonds using the notation shown in Examples 10.6 and 10.7.
\begin{equation}\text { a. }\mathrm{CH}_{2} \mathrm{Br}_{2} \quad \text { b. } \mathrm{SO}_{2} \quad \text { c. }\mathrm{NF}_{3} \quad \text { d. } \mathrm{BF}_{3}\end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

02:00

Problem 63

Write a hybridization and bonding scheme for each molecule or ion. Sketch the structure, including overlapping orbitals, and label all bonds using the notation shown in Examples 10.6 and 10.7.
\begin{equation}\begin{array}{l}{\text { a. } \mathrm{COCl}_{2}(\text { carbon is the central atom })} \\ {\text { b. } \mathrm{BrF}_{5}} \\ {\text { c. } \mathrm{XeF}_{2}} \\ {\text { d. } \mathrm{I}_{3}^{-}}\end{array}\end{equation}

Lottie Adams

Numerade Educator

04:07

Problem 64

Write a hybridization and bonding scheme for each molecule or ion. Sketch the structure, including overlapping orbitals, and label all bonds using the notation shown in Examples 10.6 and 10.7.
\begin{equation}\text { a. }\mathrm{SO}_{3}^ {2-} \quad \text { b. } \mathrm{PF}_{6}^{-} \quad \text { c. } \mathrm{BrF}_{3} \quad \text { d. } \mathrm{HCN}\end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

02:28

Problem 65

Write a hybridization and bonding scheme for each molecule that contains more than one interior atom. Indicate the hybridization about each interior atom. Sketch the structure, including overlapping orbitals, and label all bonds using the notation shown in Examples 10.6 and 10.7.
\begin{equation}\begin{array}{l}{\text { a. } \mathrm{N}_{2} \mathrm{H}_{2} \text { (skeletal structure HNNH) }} \\ {\text { b. } \mathrm{N}_{2} \mathrm{H}_{4} \text { (skeletal structure } \mathrm{H}_{2} \mathrm{NNH}_{2} )} \\ {\text { c. } \mathrm{CH}_{3} \mathrm{NH}_{2} \text { (skeletal structure } \mathrm{H}_{3} \mathrm{CNH}_{2} )}\end{array}\end{equation}

Lottie Adams

Numerade Educator

03:15

Problem 66

Write a hybridization and bonding scheme for each molecule that contains more than one interior atom. Indicate the hybridization about each interior atom. Sketch the structure, including overlapping orbitals, and label all bonds using the notation shown in Examples 10.6 and 10.7.
\begin{equation}\begin{array}{l}{\text { a. } \mathrm{C}_{2} \mathrm{H}_{2} \text { (skeletal structure HCCH) }} \\ {\text { b. } \mathrm{C}_{2} \mathrm{H}_{4} \text { (skeletal structure } \mathrm{H}_{2} \mathrm{CCH}_{2} )} \\ {\text { c. } \mathrm{C}_{2} \mathrm{H}_{6} \text { (skeletal structure } \mathrm{H}_{3} \mathrm{CCH}_{3} )}\end{array}\end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

01:17

Problem 67

Consider the structure of the amino acid alanine. Indicate the hybridization about each interior atom.

Ronald Prasad

Numerade Educator

01:21

Problem 68

Consider the structure of the amino acid aspartic acid. Indicate the hybridization about each interior atom.

Eugene Schneider

University of Minnesota - Twin Cities

01:22

Problem 69

Sketch the bonding molecular orbital that results from the linear combination of two 1$s$ orbitals. Indicate the region where interference occurs and state the kind of interference (constructive or destructive).

Joshua Speer

Numerade Educator

00:52

Problem 70

Sketch the antibonding molecular orbital that results from the linear combination of two 1$s$ orbitals. Indicate the region where interference occurs and state the kind of interference (constructive or destructive).

Eugene Schneider

University of Minnesota - Twin Cities

03:05

Problem 71

Draw an MO energy diagram and predict the bond order of $\mathrm{Be}_{2}^{+}$ and $\mathrm{Be}_{2}^{-} .$ Do you expect these molecules to exist in the gas phase?

Lottie Adams

Numerade Educator

02:05

Problem 72

Draw an MO energy diagram and predict the bond order of $\mathrm{I} \mathrm{i}_{2}^{+}$ and $\mathrm{Li}_{2}^{-} .$ Do you expect these molecules to exist in the gas phase?

Eugene Schneider

University of Minnesota - Twin Cities

06:41

Problem 73

Sketch the bonding and antibonding molecular orbitals that result from linear combinations of the 2$p_{x}$ atomic orbitals in a homonuclear diatomic molecule. (The 2$p_{x}$ orbitals are those whose
lobes are oriented along the bonding axis.)

Wan Deng

Numerade Educator

01:44

Problem 74

Sketch the bonding and antibonding molecular orbitals that result from linear combinations of the 2$p_{z}$ atomic orbitals in a homonuclear diatomic molecule. (The 2$p_{z}$ orbitals are those whose
lobes are oriented perpendicular to the bonding axis. How do these molecular orbitals differ from those obtained from linear combinations of the 2$p_{y}$ atomic orbitals? (The 2$p_{y}$ orbitals are also oriented perpendicular to the bonding axis, but also perpendicular to the 2$p_{z}$ orbitals.)

Eugene Schneider

University of Minnesota - Twin Cities

02:56

Problem 75

Using the molecular orbital energy ordering for second-row homonuclear diatomic molecules in which the $\pi_{2 p \text { orbitals lie at }}$ lower energy than the $\sigma_{2 p},$ draw MO energy diagrams and predict the bond order in a molecule or ion with each number of total valence electrons. Will the molecule or ion be diamagnetic or paramagnetic?
\begin{equation}\text { a. }4 \quad \text { b. } 6 \quad \text { c. } 8 \quad \text { d. } 9\end{equation}

Lottie Adams

Numerade Educator

07:32

Problem 76

Using the molecular orbital energy ordering for second-row homonuclear diatomic molecules in which the $\pi_{2 p}$ orbitals lie at higher energy than the $\sigma_{2 p}$ , draw MO energy diagrams and predict the bond order in a molecule or ion with each number of total valence electrons. Will the molecule or ion be diamagnetic or paramagnetic?
\begin{equation}\text { a. }10 \quad \text { b. } 12 \quad \text { c. } 13 \quad \text { d. } 14\end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

03:02

Problem 77

Use molecular orbital theory to predict if each molecule or ion exists in a relatively stable form.
\begin{equation}\text { a. }\mathrm{H}_{2}^{2-} \quad \text { b. } \mathrm{Ne}_{2} \quad \text { c. } \mathrm{He}_{2}^{2+} \quad \text { d. } \mathrm{F}_{2}^{2-}\end{equation}

Lottie Adams

Numerade Educator

03:05

Problem 78

Use molecular orbital theory to predict if each molecule or ion exists in a relatively stable form.
\begin{equation}\text { a. }\mathrm{C}_{2}^{2+} \quad \text { b. } \mathrm{Li}_{2} \quad \text { c. } \mathrm{Be}_{2}^{2+} \quad \text { d. } \mathrm{Li}_{2}^{2-}\end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

00:50

Problem 79

According to MO theory, which molecule or ion has the highest bond order? Highest bond energy? Shortest bond length?
$$\mathrm{C}_{2}, \mathrm{C}_{2}^{+}, \mathrm{C}_{2}^{-}$$

Lottie Adams

Numerade Educator

01:51

Problem 80

According to MO theory, which molecule or ion has the highest bond order? Highest bond energy? Shortest bond length?
$$\mathrm{O}_{2}, \mathrm{O}_{2}^{-}, \mathrm{O}_{2}^{2-}$$

Eugene Schneider

University of Minnesota - Twin Cities

01:27

Problem 81

Draw an MO energy diagram for CO. (Use the energy ordering of $\mathrm{O}_{2} .$ Predict the bond order and make a sketch of the lowest energy bonding molecular orbital.

Lottie Adams

Numerade Educator

02:00

Problem 82

Draw an energy diagram for HCl. Predict the bond order and make a sketch of the lowest energy bonding molecular orbital.

Eugene Schneider

University of Minnesota - Twin Cities

17:26

Problem 83

For each compound, draw the Lewis structure, determine the geometry using VSEPR theory, determine whether the molecule is polar, identify the hybridization of all interior atoms, and make a sketch of the molecule, according to valence bond theory, showing orbital overlap.
\begin{equation}\begin{array}{l}{\text { a. } \mathrm{COF}_{2}(\text { carbon is the central atom })} \\ {\text { b. } \mathrm{S}_{2} \mathrm{Cl}_{2}(\mathrm{ClSSCl})} \\ {\text { c. } \mathrm{SF}_{4}}\end{array}\end{equation}

America Guerrero

Numerade Educator

04:29

Problem 84

For each compound, draw the Lewis structure, determine the geometry using VSEPR theory, determine whether the molecule is polar, identify the hybridization of all interior atoms, and make
a sketch of the molecule, according to valence bond theory, showing orbital overlap.
\begin{equation}\text { a. } \mathrm{IF}_{5} \quad \text { b. } \mathrm{CH}_{2} \mathrm{CHCH}_{3} \quad \text { c. } \mathrm{CH}_{3} \mathrm{SH}\end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

03:49

Problem 85

Amino acids are biological compounds that link together to form proteins, the workhorse molecules in living organisms. The skeletal structures of several simple amino acids are shown here. For each skeletal structure, complete the Lewis structure, determine the geometry and hybridization about each interior ventions of Section $10.4 .$

Lottie Adams

Numerade Educator

06:18

Problem 86

The genetic code is based on four different bases with the structures shown here. Assign a geometry and hybridization to each interior atom in these four bases.
\begin{equation}\text { a. cytosine }\quad \text { b. adenine } \quad \text { c. thymine } \quad \text { d. guanine } \end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

06:02

Problem 87

The structure of caffeine, present in coffee and many soft drinks, is shown here. How many pi bonds are present in caffeine? How many sigma bonds? Insert the lone pairs in the molecule. What kinds of orbitals do the lone pairs occupy?

Ronald Prasad

Numerade Educator

04:11

Problem 88

The structure of acetylsalicylic acid (aspirin) is shown here. How many pi bonds are present in acetylsalicylic acid? How many sigma bonds? What parts of the molecule are free to rotate? What parts are rigid?

Ronald Prasad

Numerade Educator

05:59

Problem 89

Most vitamins can be classified as either fat soluble, which results in their tendency to accumulate in
the body (so that taking too much can be harmful), or water soluble, which results in their tendency to
be quickly eliminated from the body in urine. Examine the structural formulas and space-filling models
of these vitamins and determine whether each one is fat soluble (mostly nonpolar) or water soluble
(mostly polar).

Keenan Mintz

University of Miami

04:48

Problem 90

Water does not easily remove grease from dishes or hands because grease is nonpolar and water is polar. The addition of soap to water, however, allows the grease to dissolve. Study the structure of sodium stearate (a soap) and describe how it works.

Ronald Prasad

Numerade Educator

04:27

Problem 91

Draw a molecular orbital energy diagram for ClF. (Assume that the $\sigma_{p}$ orbitals are lower in energy than the $\pi$ orbitals.) What is the bond order in ClF?

Anchita Batra

Numerade Educator

04:36

Problem 92

Draw Lewis structures and MO diagrams for $\mathrm{CN}^{+}, \mathrm{CN},$ and $\mathrm{CN}^{-} .$ According to the Lewis model, which species is most stable? According to MO theory, which species is most stable? Do the two theories agree?

Eugene Schneider

University of Minnesota - Twin Cities

05:11

Problem 93

Bromine can form compounds or ions with any number of fluorine atoms from one to five. Write the formulas of all five of these species, assign a hybridization, and describe their electron and molecular geometry.

Joshua Speer

Numerade Educator

01:20

Problem 94

The compound $\mathrm{C}_{3} \mathrm{H}_{4}$ has two double bonds. Describe its bonding and geometry, using a valence bond approach.

Eugene Schneider

University of Minnesota - Twin Cities

00:51

Problem 95

Draw the structure of a molecule with the formula $\mathrm{C}_{4} \mathrm{H}_{6} \mathrm{Cl}_{2}$ that has a dipole moment of $0 .$

Lottie Adams

Numerade Educator

01:18

Problem 96

Draw the structures of two compounds that have the composition $\mathrm{CH}_{3} \mathrm{NO}_{2}$ and have all three $\mathrm{H}$ atoms bonded to the C. Predict which compound has the larger ONO bond angle.

Eugene Schneider

University of Minnesota - Twin Cities

03:55

Problem 97

How many types of hybrid orbitals do we use to describe each molecule?
\begin{equation}\begin{array}{l}{\text { a. } \mathrm{N}_{2} \mathrm{O}_{5}} \\ {\text { b. } \mathrm{C}_{2} \mathrm{H}_{5} \mathrm{NO} \text { (four } \mathrm{C}-\mathrm{H} \text { bonds and one } \mathrm{O}-\mathrm{H} \text { bond } )} \\ {\text { c. } \mathrm{BrCN}(\text { no formal charges) }}\end{array}\end{equation}

Chang Qu

Numerade Educator

01:17

Problem 98

Indicate which orbitals overlap to form the $\sigma$ bonds in each molecule.
\begin{equation}\text { a. }\mathrm{BeBr}_{2} \quad \text { b. } \mathrm{HgCl}_{2} \quad \text { c. }\mathrm{ICN}\end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

00:58

Problem 99

In VSEPR theory, which uses the Lewis model to determine molecular geometry, the trend of decreasing bond angles in $\mathrm{C} \mathrm{CH}_{4}, \mathrm{NH}_{3},$ and $\mathrm{H}_{2} \mathrm{O}$ is accounted for by the greater repulsion of lone pair electrons compared to bonding pair electrons. How would this trend be accounted for in valence bond theory?

Lottie Adams

Numerade Educator

02:34

Problem 100

The results of a molecular orbital calculation for $\mathrm{H}_{2} \mathrm{O}$ are shown here. Examine each of the orbitals and classify them as bonding, antibonding, or nonbonding. Assign the correct number of electrons to the energy diagram. According to this energy diagram, is $\mathrm{H}_{2} \mathrm{O}$ stable? Explain.

Eugene Schneider

University of Minnesota - Twin Cities

00:54

Problem 101

The results of a molecular orbital calculation for $\mathrm{NH}_{3}$ are shown here. Examine each of the orbitals and classify them as bonding, antibonding, or nonbonding. Assign the correct number of
electrons to the energy diagram. According to this energy diagram, is $\mathrm{NH}_{3}$ stable? Explain.

Lottie Adams

Numerade Educator

02:28

Problem 102

cis-2-Butene isomerizes to trans-2-butene via the reaction shown here.
\begin{equation}\begin{array}{l}{\text { a. If isomerization requires breaking the } \pi \text { bond, what minimum }} \\ {\text { energy is required for isomerization in } \mathrm{J} / \mathrm{mol} ? \text { In }} \\ {\text { J/molecule? }} \\ {\text { b. If the energy for isomerization came from light, what }} \\ {\text { minimum frequency of light would be required? In what }} \\ {\text { portion of the electromagnetic spectrum does this }} \\ {\text { frequency lie? }}\end{array}\end{equation}

Eugene Schneider

University of Minnesota - Twin Cities

00:46

Problem 103

The species $\mathrm{NO}_{2}, \mathrm{NO}_{2}^{+},$ and $\mathrm{NO}_{2}^{-}$ in which $\mathrm{N}$ is the central atom have very different bond angles. Predict what these bond angles might be with respect to the ideal angles and justify your prediction.

Lottie Adams

Numerade Educator

01:03

Problem 104

The bond angles increase steadily in the series $\mathrm{PF}_{3}, \mathrm{PCl}_{3}, \mathrm{PBr}_{3}$ and $\mathrm{PI}_{3} .$ After consulting the data on atomic radii in Chapter $8,$
provide an explanation for this observation.

Eugene Schneider

University of Minnesota - Twin Cities

01:08

Problem 105

The ion $\mathrm{CH}_{5}^{+}$ can form under very special high-energy conditions in the vapor phase in a mass spectrometer. Propose a hybridization for the carbon atom and predict the geometry.

Lottie Adams

Numerade Educator

01:02

Problem 106

Neither the VSEPR model nor the hybridization model is able to account for the experimental observation that the $\mathrm{F}-\mathrm{Ba}-\mathrm{F}$ bond angle in gaseous BaF $_{2}$ is $108^{\circ}$ rather than the predicted $180^{\circ} .$ Suggest some possible explanations for this observation.

Eugene Schneider

University of Minnesota - Twin Cities

16:43

Problem 107

Draw the Lewis structure for acetamide $\left(\mathrm{CH}_{3} \mathrm{CONH}_{2}\right),$ an organic compound, and determine the geometry about each interior atom. Experiments show that the geometry about the nitrogen atom in acetamide is nearly planar. What resonance structure can account for the planar geometry about the nitrogen atom?

Erin Wagner

Numerade Educator

01:20

Problem 108

Use VSEPR theory to predict the geometry (including bond angles) about each interior atom of methyl azide $\left(\mathrm{CH}_{3} \mathrm{N}_{3}\right),$ and make a sketch of the molecule. Would you expect the bond angle between the two interior nitrogen atoms to be the same or different? Would you expect the two nitrogen-nitrogen bond lengths to be the same or different?

Eugene Schneider

University of Minnesota - Twin Cities

01:21

Problem 109

Which statement best captures the fundamental idea behind VSEPR theory? Explain what is wrong with each of the other statements.
\begin{equation}\begin{array}{l}{\text { a. The angle between two or more bonds is determined primarily }} \\ {\text { by the repulsions between the electrons within those }} \\ {\text { bonds and other (lone pair) electrons on the central atom of }} \\ {\text { a molecule. Each of these electron groups (bonding electrons }}\\{\text { or lone pair electrons) will lower its potential energy by maximizing }} \\ {\text { its separation from other electron groups, thus determining }} \\ {\text { the geometry of the molecule. }}\\{\text { b. The angle between two or more bonds is determined primarily }} \\ {\text { by the repulsions between the electrons within those }} \\ {\text { bonds. Each of these bonding electrons will lower its potential }} \\ {\text { energy by maximizing its separation from other electron }} \\ {\text { groups, thus determining the geometry of the molecule. }}\\{\text { c. The geometry of a molecule is determined by the shapes of }} \\ {\text { the overlapping orbitals that form the chemical bonds. Therefore, }} \\ {\text { to determine the geometry of a molecule, you must }} \\ {\text { determine the shapes of the orbitals involved in bonding. }}\end{array}\end{equation}

Lottie Adams

Numerade Educator

00:59

Problem 110

Suppose that a molecule has four bonding groups and one lone pair on the central atom. Suppose further that the molecule is confined to two dimensions (this is a purely hypothetical assumption for the sake of understanding the principles behind VSEPR theory). Make a sketch of the molecule and estimate the bond angles.

Eugene Schneider

University of Minnesota - Twin Cities

03:18

Problem 111

How does each of the three major bonding theories (the Lewis model, valence bond theory, and molecular orbital theory) define a single chemical bond? A double bond? A triple bond? How are these definitions similar? How are they different?

Joshua Speer

Numerade Educator

01:34

Problem 112

The most stable forms of the nonmetals in groups $4 \mathrm{A}, 5 \mathrm{A},$ and 6 $\mathrm{A}$ of the second period are molecules with multiple bonds. Beginning with the third period, the most stable forms of the non-metals of these groups are molecules without multiple bonds. Propose an explanation for this observation based on valence bond theory.

Eugene Schneider

University of Minnesota - Twin Cities

01:05

Problem 113

In complete sentences, describe why someone might expect the bond angles in methane $\left(\mathrm{CH}_{4}\right)$ to be $90^{\circ}$ even though the bond angles are actually $109.5^{\circ} .$

Lottie Adams

Numerade Educator

01:20

Problem 114

At least two different numbers of electron groups can result in a linear molecule. What are they? What are the numbers of bonding groups and lone pairs in each case? Provide an example of a linear molecule in each case.

Eugene Schneider

University of Minnesota - Twin Cities

02:16

Problem 115

Have each member of your group select one of the molecules shown below and complete steps a-d. Each member should then present his or her results to the rest of the group, explaining the reasoning used to determine the answers.
$$\mathrm{CS}_{2} \quad \mathrm{NCl}_{3} \quad \mathrm{CF}_{4} \quad \mathrm{CH}_{2} \mathrm{F}_{2}$$
\begin{equation}\begin{array}{l}{\text { a. Draw the Lewis dot structure. }} \\ {\text { b. Determine the molecular geometry and draw it accurately. }} \\ {\text { c. Indicate the polarity of any polar bonds within the structure. }} \\ {\text { d. Classify the molecule as polar or nonpolar. }}\end{array}\end{equation}

Lottie Adams

Numerade Educator

00:45

Problem 116

How many atomic orbitals form a set of $s p^{3}$ hybrid orbitals? A set of $s p^{2}$ hybrid orbitals? A set of sp hybrid orbitals? What is the relationship between these numbers and the number of electron groups around the central atom?

Eugene Schneider

University of Minnesota - Twin Cities

01:21

Problem 117

Use molecular orbital theory to explain in detail why $\mathrm{N}_{2}^{+}$ and $\mathrm{N}_{2}^{-}$ have similar bond strengths and both are very different from neutral $\mathrm{N}_{2}$ .

Lottie Adams

Numerade Educator

03:04

Problem 118

The distance between two atoms that are attached to a common central atom is approximately equal to the sum of their "target distance" atomic radii. For example, consider the structures of $\mathrm{CF}_{4}$ and $\mathrm{OCF}_{2}$ in Figure a $\nabla$ . Given that the target distance radius of fluorine when it is bonded to carbon is $108 \mathrm{pm},$ the target distance between the two fluorine atoms

Eugene Schneider

University of Minnesota - Twin Cities