General Chemistry

Problem 12

What factors determine the strength of interaction between two atomic orbitals to form a molecular orbital?

Lottie Adams

Numerade Educator

02:15

Problem 13

Describe the formation of bonding and antibonding molecular orbitals resulting from the interaction of two $2 s$ orbitals.

Lottie Adams

Numerade Educator

01:47

Problem 14

Describe the formation of molecular orbitals resulting from the interaction of two $2 p$ orbitals.

Lottie Adams

Numerade Educator

02:51

Problem 15

How does molecular orbital theory describe the bonding in the HF molecule?

Lottie Adams

Numerade Educator

02:08

Describe the bonding in $\mathrm{O}_{3}$, using molecular orbital theory. Compare this with the resonance description. 10.17 According to the VSEPR model, the $\mathrm{H}-\mathrm{P}-\mathrm{H}$ bond angle in $\mathrm{PH}_{3}$ is
a. $120^{\circ}$
b. $109.5^{\circ}$
c. $90^{\circ}$
d. a little less than $120^{\circ}$
e. a little less than $109.5^{\circ}$

Lottie Adams

Numerade Educator

Problem 17

According to the VSEPR model, the $\mathrm{H}-\mathrm{P}-\mathrm{H}$ bond angle in $\mathrm{PH}_{3}$ is
a. $120^{\circ}$
b. $109.5^{\circ}$
c. $90^{\circ}$
d. a little less than $120^{\circ}$
e. a little less than $109.5^{\circ}$

Lottie Adams

Numerade Educator

01:40

Problem 18

Which of the following molecular geometries does the $\mathrm{XeF}_{5}^{+}$ cation exhibit?
a. tetradral
b. T-shaped
c. octahedral
d. square pyramidal
e. trigonal bipyramidal

Lottie Adams

Numerade Educator

02:47

Problem 19

Which of the following would be a polar molecule?
a. $\mathrm{CO}_{2}$
b. $\mathrm{H}_{2} \mathrm{~S}$
c. $\mathrm{CH}_{4}$
d. $\mathrm{SF}_{6}$
e. $\mathrm{BeF}_{2}$

Lottie Adams

Numerade Educator

01:42

Problem 20

What is the bond order of NO?
a. $\frac{1}{2}$
b. 1
c. $\frac{3}{2}$
d. 3
e. $\frac{5}{2}$

Lottie Adams

Numerade Educator

03:21

Problem 21

A student writes the Lewis electron-dot formula for the carbonate anion, $\mathrm{CO}_{3}^{2-}, \mathrm{as}$
a. Does this Lewis formula obey the octet rule? Explain. What are the formal charges on the atoms? Try describing the bonding for this formula in valence bond terms. Do you have any difficulty doing this?
b. Does this Lewis formula give a reasonable description of the electron structure, or is there a better one? If there is a better Lewis formula, write it down and explain why it is better.
C. The same student writes the following resonance description for $\mathrm{CO}_{2}:$
Is there something wrong with this description? (What would you predict as the geometries of these formulas?)
d. Is one or the other formula a better description? Could a value for the dipole moment help you decide?
e. Can you write a Lewis formula that gives an even better description of $\mathrm{CO}_{2}$ ? Explain your answer.

Lottie Adams

Numerade Educator

02:34

Problem 22

a. Shown here are several potential Lewis electron-dot formulas for the $\mathrm{SeO}_{2}$ molecule.
Are there any electron-dot formulas that you expect to give the best description? How would you describe the bonding in terms of electron-dot formulas?
b. Describe the bonding in $\mathrm{SeO}_{2}$ in valence bond terms. If there is delocalized bonding, note this in your description and explain how molecular orbital theory can be useful here.
C. Determine the arrangement of electron pairs about Se in the $\mathrm{SeO}_{2}$ molecule. What would you expect for the molecular geometry of the $\mathrm{SeO}_{2}$ molecule?
d. Would you expect the $\mathrm{O}-\mathrm{Se}-\mathrm{O}$ angle in the $\mathrm{SeO}_{2}$ molecule to be greater than, equal to, or less than $120^{\circ} .$ Explain your answer.
e. Draw an electron-dot formula for the $\mathrm{H}_{2}$ Se molecule and determine its molecular geometry.
f. Compare the $\mathrm{H}-\mathrm{Se}-\mathrm{H}$ bond angle to the $\mathrm{O}-\mathrm{Se}-\mathrm{O}$
bond angle. Is it larger or smaller? Explain.
g. Determine whether the $\mathrm{H}_{2} \mathrm{Se}$ and $\mathrm{SeO}_{2}$ molecules have dipole moments. Describe how you arrived at your answer.

Lottie Adams

Numerade Educator

01:24

Problem 23

Match the following molecular substances with one of the molecular models (i) to (iv) that correctly depicts the geometry of the corresponding molecule.
a. $\mathrm{SeO}_{2}$
b. $\mathrm{BeCl}_{2}$
c. $\mathrm{PBr}_{3}$
d. $\mathrm{BCl}_{3}$

Lottie Adams

Numerade Educator

Problem 24

Which of the following molecular models correctly depicts the geometry of CICN?

Lottie Adams

Numerade Educator

01:27

Problem 25

Suppose that an ethane molecule, $\mathrm{CH}_{3} \mathrm{CH}_{3}$, is broken into two $\mathrm{CH}_{3}$ molecules in such a way that one $: \mathrm{CH}_{3}$ molecule retains the electron pair that was originally the one making up the $\mathrm{C}-\mathrm{C}$ bond. The other $\mathrm{CH}_{3}$ molecule has two fewer electrons. Imagine that momentarily this $\mathrm{CH}_{3}$ molecule has the geometry it had in the ethane molecule. Describe the electron repulsions present in this molecule, and indicate how they would be expected to rearrange its geometry.

Lottie Adams

Numerade Educator

02:09

Problem 26

Suppose that a $\mathrm{BF}_{3}$ molecule approaches the lone pair on the $\mathrm{N}$ atom of an $: \mathrm{NH}_{3}$ molecule, and that a bond forms between the $\mathrm{B}$ atom and the $\mathrm{N}$ atom. Consider the arrangement of electron pairs about the $\mathrm{B}$ atom at the moment of this bond formation, and describe the repulsions among the electron pairs and how they might be expected to change the geometry about the $\mathrm{B}$ atom.

Lottie Adams

Numerade Educator

Problem 27

Indicate what hybrid orbital depicted below is expected for the central atom of each of the following:
a. $\mathrm{BeF}_{2}$
b. $\mathrm{SiF}_{4}$
C. $\mathrm{SeF}_{4}$
d. $\bar{R} n F_{4}$

Lottie Adams

Numerade Educator

01:25

Problem 28

An atom in a molecule has two bonds to other atoms and one lone pair. What kind of hybrid orbitals do you expect for this atom? Describe how you arrived at your answer.

Lottie Adams

Numerade Educator

01:31

Problem 29

Two compounds have the same molecular formula, $\mathrm{C}_{2} \mathrm{H}_{2} \mathrm{Br}_{2}$. One has a dipole moment; the other does not. Both compounds react with bromine, $\mathrm{Br}_{2}$, to produce the same compound. This reaction is a generally accepted test for double bonds, and each bromine atom of $\mathrm{Br}_{2}$ adds to a different atom of the double bond. What is the identity of the original compounds? Describe the argument you use.

Lottie Adams

Numerade Educator

01:51

Problem 30

A neutral molecule is identified as a tetrafluoride, $\mathrm{XF}_{4,}$ where $\mathrm{X}$ is an unknown atom. If the molecule has a dipole moment of $0.63 \mathrm{D}$, can you give some possibilities for the identity of $\mathrm{X}$ ?

Lottie Adams

Numerade Educator

01:08

Problem 31

Acetic acid, the sour constituent of vinegar, has the following structure:
Indicate what geometry given below is expected to be found about each of the atoms labeled $\mathrm{a}, \mathrm{b}$, and $\mathrm{c}$.

Lottie Adams

Numerade Educator

Problem 32

What are the bond angles predicted by the VSEPR model about the carbon atom in the formate ion, $\mathrm{HCO}_{2}^{-}$ ? Considering that the bonds to this atom are not identical, would you expect the experimental values to agree precisely with the VSEPR values? How might they differ?

Lottie Adams

Numerade Educator

01:37

Problem 33

Predict the shape or geometry of the following molecules, using the VSEPR model.
a. $\mathrm{SiF}_{4}$
b. $\mathrm{SF}_{2}$
c. $\mathrm{COF}_{2}$
d. $\mathrm{PCl}_{3}$

Lottie Adams

Numerade Educator

Problem 34

Use the electron-pair repulsion model to predict the geometry of the following molecules:
a. $\mathrm{GeCl}_{2}$
b. $\mathrm{NF}_{3}$
c. $\mathrm{SCl}_{2}$
d. $\mathrm{XeO}_{4}$

Lottie Adams

Numerade Educator

01:29

Problem 35

Predict the geometry of the following ions, using the electron-pair repulsion model.
a. $\mathrm{ClO}_{3}^{-}$
b. $\mathrm{PO}_{4}{ }^{3-}$
c. $\mathrm{SCN}^{-}$
d. $\mathrm{H}_{3} \mathrm{O}^{+}$

Lottie Adams

Numerade Educator

01:31

Problem 36

Use the VSEPR model to predict the geometry of the following ions:
a. $\mathrm{N}_{3}^{-}$
b. $\mathrm{BH}_{4}^{-}$
c. $\mathrm{SO}_{3}^{2-}$
d. $\mathrm{NO}_{2}^{-}$

Lottie Adams

Numerade Educator

02:10

Problem 37

For each of the following molecules, state the bond angle (or bond angles, as appropriate) that you would expect to see on the central atom based on the simple VSEPR model. Would you expect the actual bond angles to be greater or less than this?
a. $\mathrm{CCl}_{4}$
b. $\mathrm{SCl}_{2}$
c. $\mathrm{COCl}_{2}$
d. $\mathrm{AsH}_{3}$

Lottie Adams

Numerade Educator

01:44

Problem 38

For each of the following molecules, state the bond angle (or bond angles, as appropriate) that you would expect to see on the central atom based on the simple VSEPR model. Would you expect the actual bond angles to be greater or less than this?
a. $\mathrm{NCl}_{3}$
b. $\mathrm{CF}_{2} \mathrm{CF}_{2}$
c. $\mathrm{GeF}_{4}$
d. $\mathrm{OF}_{2}$

Lottie Adams

Numerade Educator

Problem 39

What geometry is expected for the following molecules, according to the VSEPR model?
a. $\mathrm{PF}_{5}$
b. $\mathrm{BrF}_{3}$
c. $\mathrm{BrF}_{5}$
d. $\mathrm{SCl}_{4}$

Lottie Adams

Numerade Educator

01:42

Problem 40

From the electron-pair repulsion model, predict the geometry of the following molecules:
a. $\mathrm{ClF}_{5}$
b. $\mathrm{SbF}_{5}$
c. $\mathrm{SeF}_{4}$
d. $\mathrm{TeF}_{6}$

Lottie Adams

Numerade Educator

01:29

Problem 41

Predict the geometries of the following ions, using the VSEPR model.
a. $\mathrm{SnCl}_{5}^{-}$
b. $\mathrm{PF}_{6}^{-}$
c. $\mathrm{ClF}_{2}^{-}$
d. $\mathrm{IF}_{4}^{-}$

Lottie Adams

Numerade Educator

Problem 42

Name the geometries expected for the following ions, according to the electron-pair repulsion model.
a. $\mathrm{ICl}_{2}^{-}$
b. $\mathrm{BrF}_{4}^{-}$
c. $\mathrm{ClF}_{6}^{+}$
d. $\mathrm{CIF}_{4}^{+}$

Lottie Adams

Numerade Educator

01:40

Problem 43

(a) The molecule $\mathrm{AsF}_{3}$ has a dipole moment of $2.59 \mathrm{D}$. Which of the following geometries are possible: trigonal planar, trigonal pyramidal, or T-shaped? (b) The molecule $\mathrm{H}_{2} \mathrm{~S}$ has a dipole moment of $0.97 \mathrm{D}$. Is the geometry linear or bent?

Lottie Adams

Numerade Educator

01:23

Problem 44

(a) The molecule $\mathrm{BrF}_{3}$ has a dipole moment of $1.19 \mathrm{D}$. Which of the following geometries are possible: trigonal planar, trigonal pyramidal, or T-shaped? (b) The molecule $\mathrm{TeCl}_{4}$ has a dipole moment of $2.54 \mathrm{D}$. Is the geometry tetrahedral, seesaw, or square planar?

Lottie Adams

Numerade Educator

02:00

Problem 45

Which of the following molecules would be expected to have zero dipole moment on the basis of their geometry?
a. $\mathrm{CS}_{2}$
b. $\mathrm{TeF}_{2}$
c. $\mathrm{SeCl}_{4}$
d. $\mathrm{XeF}_{4}$

Lottie Adams

Numerade Educator

Problem 46

Which of the following molecules would be expected to have a dipole moment of zero because of symmetry?
a. $\mathrm{SeF}_{2}$
b. $\mathrm{BeI}_{2}$
c. TeF $_{6}$
d. $\mathrm{TeF}_{4}$

Lottie Adams

Numerade Educator

Problem 47

What hybrid orbitals would be expected for the central atom in each of the following molecules or ions?

Lottie Adams

Numerade Educator

Problem 48

What hybrid orbitals would be expected for the central atom in each of the following molecules or ions?

Lottie Adams

Numerade Educator

Problem 49

What hybrid orbitals would be expected for the central atom in each of the following molecules or ions?
a. $\mathrm{SeCl}_{2}$
b. $\mathrm{NO}_{2}^{-}$
C. $\mathrm{CO}_{2}$
d. $\mathrm{COF}_{2}$

Lottie Adams

Numerade Educator

Problem 50

What hybrid orbitals would be expected for the central atom in each of the following molecules or ions?
a. $\mathrm{GeCl}_{4}$
b. $\mathrm{PBr}_{3}$
c. $\mathrm{BeF}_{2}$
d. $\mathrm{SO}_{2}$

Lottie Adams

Numerade Educator

01:35

Problem 51

(a) Mercury(II) chloride dissolves in water to give poorly conducting solutions, indicating that the compound is largely nonionized in solution - it dissolves as $\mathrm{HgCl}_{2}$ molecules. Describe the bonding of the $\mathrm{HgCl}_{2}$ molecule, using valence bond theory. (b) Phosphorus trichloride, $\mathrm{PCl}_{3}$, is a colorless liquid with a highly irritating vapor. Describe the bonding in the $\mathrm{PCl}_{3}$ molecule, using valence bond theory. Use hybrid orbitals.

Lottie Adams

Numerade Educator

01:24

Problem 52

(a) Nitrogen trifluoride, $\mathrm{NF}_{3}$, is a relatively unreactive, colorless gas. How would you describe the bonding in the $\mathrm{NF}_{3}$ molecule in terms of valence bond theory? Use hybrid orbitals.
(b) Silicon tetrafluoride, $\mathrm{SiF}_{4}$, is a colorless gas formed when hydrofluoric acid attacks silica $\left(\mathrm{SiO}_{2}\right)$ or glass. Describe the bonding in the $\mathrm{SiF}_{4}$ molecule, using valence bond theory.

Lottie Adams

Numerade Educator

01:17

Problem 53

What hybrid orbitals would be expected for the central atom in each of the following?
a. $\mathrm{XeF}_{2}$
b. $\mathrm{BrF}_{5}$
c. $\mathrm{PCl}_{5}$
d. $\mathrm{CIF}_{4}^{+}$

Lottie Adams

Numerade Educator

Problem 54

What hybrid orbitals would be expected for the central atom in each of the following?
a. $\mathrm{BrF}_{3}$
b. $\mathrm{TeF}_{4}$
c. $\mathrm{XeF}_{4}$
d. $\mathrm{IF}_{4}^{-}$

Lottie Adams

Numerade Educator

01:35

Problem 55

Phosphorus pentachloride is normally a white solid. It exists in this state as the ionic compound $\left[\mathrm{PCl}_{4}^{+}\right]\left[\mathrm{PCl}_{6}^{-}\right] .$

Lottie Adams

Numerade Educator

01:01

Problem 56

Iodine, $\mathrm{I}_{2}$, dissolves in an aqueous solution of iodide ion, $\mathrm{I}^{-}$, to give the triiodide ion, $\mathrm{I}_{3}^{-}$. This ion consists of one iodine atom bonded to two others. Describe the bonding of $I_{3}^{-}$ in terms of valence bond theory.

Lottie Adams

Numerade Educator

01:35

Problem 57

(a) Formaldehyde, $\mathrm{H}_{2} \mathrm{CO}$, is a colorless, pungent gas used to make plastics. Give the valence bond description of the formaldehyde molecule. (Both hydrogen atoms are attached to the carbon atom.) (b) Nitrogen, $\mathrm{N}_{2}$, makes up about $80 \%$ of the earth's atmosphere. Give the valence bond description of this molecule.

Lottie Adams

Numerade Educator

01:29

Problem 58

(a) The molecule HN=NH exists as a transient species in certain reactions. Give the valence bond description of this species. (b) Hydrogen cyanide, HCN, is a very poisonous gas or liquid with the odor of bitter almonds. Give the valence bond description of HCN. (Carbon is the central atom.)

Lottie Adams

Numerade Educator

Problem 59

The hyponitrite ion, $^{-} \mathrm{O}-\mathrm{N}=\mathrm{N}-\mathrm{O}^{-}$, exists in solid compounds as the trans isomer. Using valence bond theory, explain why cis-trans isomers might be expected for this ion. Draw structural formulas of the cis-trans isomers.

Lottie Adams

Numerade Educator

01:25

Problem 60

Fumaric acid, $\mathrm{C}_{4} \mathrm{H}_{4} \mathrm{O}_{4}$, occurs in the metabolism of glucose in the cells of plants and animals. It is used commercially in beverages. The structural formula of fumaric acid is Maleic acid is the cis isomer of fumaric acid. Using valence bond theory, explain why these isomers are possible.

Lottie Adams

Numerade Educator

01:55

Problem 61

Describe the electronic structure of each of the following, using molecular orbital theory. Calculate the bond order of each and decide whether it should be stable. For each, state whether the substance is diamagnetic or paramagnetic.
a. $\mathrm{B}_{2}$
b. $\mathrm{B}_{2}^{+}$
c. $\mathrm{O}_{2}^{-}$

Lottie Adams

Numerade Educator

01:34

Problem 62

Use molecular orbital theory to describe the bonding in the following. For each one, find the bond order and decide whether it is stable. Is the substance diamagnetic or paramagnetic?
a. $\mathrm{Be}_{2}^{+}$
b. $\mathrm{Ne}_{2}$
c. $\mathrm{B}_{2}^{-}$

Lottie Adams

Numerade Educator

01:14

Problem 63

Assume that the cyanide ion, $\mathrm{CN}^{-}$, has molecular orbitals similar to those of a homonuclear diatomic molecule. Write the configuration and bond order of $\mathrm{CN}^{-}$. Is a substance of the ion diamagnetic or paramagnetic?

Lottie Adams

Numerade Educator

01:08

Problem 64

Write the molecular orbital configuration of the diatomic molecule BN. What is the bond order of BN? Is the substance diamagnetic or paramagnetic? Use the order of energies that was given for homonuclear diatomic molecules.

Lottie Adams

Numerade Educator

01:18

Problem 65

Predict the molecular geometry of the following:
a. $\mathrm{SnCl}_{2}$
b. $\mathrm{COBr}_{2}$
c. $\mathrm{ICl}_{2}$
d. $\mathrm{PCl}_{6}^{-}$

Lottie Adams

Numerade Educator

Problem 66

Predict the molecular geometry of the following:
a. HOF
b. $\mathrm{NH}_{4}^{+}$
c. $\mathrm{PF}_{5}$
d. $\mathrm{ClF}_{4}^{+}$

Lottie Adams

Numerade Educator

01:16

Problem 67

Which of the following molecules or ions are linear?
a. $\mathrm{SeF}_{2}$
b. $\mathrm{CHCl}_{3}$
c. $\mathrm{SeF}_{4}$
d. $\mathrm{SnF}_{6}^{2-}$

Lottie Adams

Numerade Educator

01:16

Problem 67

Which of the following molecules or ions are linear?
a. $\mathrm{SeF}_{2}$
b. $\mathrm{CHCl}_{3}$
c. $\mathrm{SeF}_{4}$
d. $\mathrm{SnF}_{6}^{2-}$

Lottie Adams

Numerade Educator

01:11

Problem 68

Which of the following molecules or ions are trigonal planar?
a. $\mathrm{BBr}_{3}$
b. $\mathrm{AsH}_{3}$
c. $\mathrm{ClF}_{4}^{-}$
d. $\mathrm{TeF}_{5}^{-}$

Lottie Adams

Numerade Educator

01:48

Problem 69

Describe the hybrid orbitals used by each carbon atom in the following molecules:

Lottie Adams

Numerade Educator

01:19

Problem 70

Describe the hybrid orbitals used by each nitrogen atom in the following molecules:

Lottie Adams

Numerade Educator

01:14

Problem 71

Explain how the dipole moment could be used to distinguish between the cis and trans isomers of 1,2 -dibromoethene:

Lottie Adams

Numerade Educator

Problem 72

Two compounds have the formula $\mathrm{Pt}\left(\mathrm{NH}_{3}\right)_{2} \mathrm{Cl}_{2}$. (Compound $\mathrm{B}$ is cisplatin, mentioned in the opening to Chapter $1 .)$ They have square planar structures. One is expected to have a dipole moment; the other is not. Which one would have a dipole moment?

Lottie Adams

Numerade Educator

01:11

Problem 73

Explain in terms of bonding theory why all four hydrogen atoms of allene, $\mathrm{H}_{2} \mathrm{C}=\mathrm{C}=\mathrm{CH}_{2}$, cannot lie in the same plane.

Lottie Adams

Numerade Educator

01:06

Problem 74

Explain in terms of bonding theory why all atoms of $\mathrm{H}_{2} \mathrm{C}=\mathrm{C}=\mathrm{C}=\mathrm{CH}_{2}$ must lie in the same plane.

Lottie Adams

Numerade Educator

01:05

Problem 75

What is the molecular orbital configuration of $\mathrm{HeH}^{+} ?$ Do you expect the ion to be stable?

Lottie Adams

Numerade Educator

01:02

Problem 76

What is the molecular orbital configuration of $\mathrm{He}_{2}^{+} ?$ Do you expect the ion to be stable?

Lottie Adams

Numerade Educator

01:19

Problem 77

Calcium carbide, $\mathrm{CaC}_{2}$, consists of $\mathrm{Ca}^{2+}$ and $\mathrm{C}_{2}^{2-}$ (acetylide) ions. Write the molecular orbital configuration and bond order of the acetylide ion, $\mathrm{C}_{2}{ }^{2-}$.

Lottie Adams

Numerade Educator

01:24

Problem 78

Sodium peroxide, $\mathrm{Na}_{2} \mathrm{O}_{2}$, consists of $\mathrm{Na}^{+}$ and $\mathrm{O}_{2}^{2-}$ (peroxide) ions. Write the molecular orbital configuration and bond order of the peroxide ion, $\mathrm{O}_{2}{ }^{2-}$.

Lottie Adams

Numerade Educator

02:06

Problem 79

The oxygen-oxygen bond in $\mathrm{O}_{2}^{+}$ is $112 \mathrm{pm}$ and in $\mathrm{O}_{2}$ is $121 \mathrm{pm}$. Explain why the bond length in $\mathrm{O}_{2}{ }^{+}$ is shorter than in $\mathrm{O}_{2}$. Would you expect the bond length in $\mathrm{O}_{2}^{-}$ to be longer or shorter than that in $\mathrm{O}_{2}$ ? Why?

Lottie Adams

Numerade Educator

01:43

Problem 80

The nitrogen-nitrogen bond distance in $\mathrm{N}_{2}$ is $109 \mathrm{pm} .$ On the basis of bond orders, would you expect the bond distance in $\mathrm{N}_{2}{ }^{+}$ to be less than or greater than $109 \mathrm{pm}$ ? Answer the same question for $\mathrm{N}_{2}^{-}$.

Lottie Adams

Numerade Educator

01:30

Problem 81

Using molecular orbital theory, determine the electronic structure of the first excited electronic state of $\mathrm{N}_{2}$. What differences are expected in the properties of the excited state of $\mathrm{N}_{2}$ compared with the same properties of the ground state?

Lottie Adams

Numerade Educator

01:55

Problem 82

The ionization energy of $\mathrm{O}_{2}$ is smaller than the ionization energy of atomic $\mathrm{O} ;$ the opposite is true for the ionization energies of $\mathrm{N}_{2}$ and atomic $\mathrm{N}$. Explain this behavior in terms of the molecular orbital energy diagrams of $\mathrm{O}_{2}$ and $\mathrm{N}_{2}$.

Lottie Adams

Numerade Educator

01:18

Problem 83

Describe the property of "handedness." What is it about the bonding of atoms in molecules that can give rise to handedness in molecules?

Lottie Adams

Numerade Educator

Problem 84

Biological molecules frequently display handedness. Describe the difference between the molecule responsible for the flavor of spearmint and the molecule responsible for the flavor of caraway seeds. In which way are the molecules similar and in which way do they differ?

Lottie Adams

Numerade Educator

01:53

Problem 85

Color vision results from the absorption of light by the cone cells of the retina. Briefly describe how different types of cone cells give rise to color vision. What is the chemical substance primarily responsible for human vision? How can its color absorption be changed?

Lottie Adams

Numerade Educator

01:59

Problem 86

Describe the cis-trans conversion that occurs in the ultraviolet absorption of cis $-1,2$ -dichloroethene.

Lottie Adams

Numerade Educator

01:18

Problem 87

What is the biological importance of stratospheric ozone? Explain.

Lottie Adams

Numerade Educator

01:54

Problem 88

In 1995, Sherwood Rowland and Mario Molina received the Nobel Prize in chemistry. Explain in some detail the work they did for which they were awarded this prize.

Lottie Adams

Numerade Educator

Problem 89

Calcium carbide, $\mathrm{CaC}_{2}$, has an ionic structure with ions $\mathrm{Ca}^{2+}$ and $\mathrm{C}_{2}^{2-} .$ Give the valence bond description of the bonding in the $\mathrm{C}_{2}^{2-}$ ion. Now write the MO configuration of this ion. What is the bond order according to this configuration? Compare the valence bond and molecular orbital descriptions.

Lottie Adams

Numerade Educator

02:21

Problem 90

Write Lewis formulas for the BF molecule (two with a single $\mathrm{B}-\mathrm{F}$ bond, two with a double $\mathrm{B}-\mathrm{F}$ bond, and one with a triple $\mathrm{B}-\mathrm{F}$ bond in which the octet rule is satisfied for at least one of the atoms. Obtain the formal charges of the atoms. Based on the information you have, is there one formula that you think best describes the molecule? Explain. For this best Lewis formula, give the valence bond description of the bonding.

Lottie Adams

Numerade Educator

01:09

Problem 91

Boron trifluoride, $\mathrm{BF}_{3}$, reacts with ammonia, $\mathrm{NH}_{3}$, to form an addition compound, $\mathrm{BF}_{3} \mathrm{NH}_{3}$. Describe the geometries about the $\mathrm{B}$ and the $\mathrm{F}$ atoms in this compound. Describe the hybridization on these two atoms. Now describe the bonding between the $\mathrm{B}$ and $\mathrm{F}$ atoms, using valence bond theory. Compare the geometries and hybridization of these atoms in this addition compound with that in the reactant molecules $\mathrm{BF}_{3}$ and $\mathrm{NH}_{3}$.

Lottie Adams

Numerade Educator

02:06

Problem 92

Acetylsalicylic acid (aspirin) has the structure shown by the following molecular model: Describe the geometry and bonding about the carbon atom shown at the top of the model, in which the carbon atom (in black) is bonded to two oxygen atoms (red) and to another carbon atom. One of the oxygen atoms is bonded to a hydrogen atom (blue).

Lottie Adams

Numerade Educator

01:17

Problem 93

A molecule $\mathrm{XF}_{6}$ (having no lone pairs) has a dipole moment of zero. (X denotes an unidentified element.) When two atoms of fluorine have been taken away, you get the molecule $\mathrm{XF}_{4}$, which has dipole moment of $0.632 \mathrm{D}$.
$$\mathrm{XF}_{6} \longrightarrow \mathrm{XF}_{4}+2 \mathrm{~F}$$
Describe the molecular geometry and bonding for each molecule. What are some possibilities for $\mathrm{X} ?$

Lottie Adams

Numerade Educator

01:30

Problem 94

Three different compounds have the same molecular formula, $\mathrm{C}_{2} \mathrm{H}_{2} \mathrm{~F}_{2} ;$ call them $\mathrm{A}, \mathrm{B}$, and $\mathrm{C}$. Compound $\mathrm{A}$ has a dipole moment of $2.42 \mathrm{D}$, compound $\mathrm{B}$ has a dipole moment of $1.38 \mathrm{D}$, and compound $\mathrm{C}$ has a dipole moment of $0 \mathrm{D} .$ All three compounds react with hydrogen, $\mathrm{H}_{2}$ :
$$\mathrm{C}_{2} \mathrm{H}_{2} \mathrm{~F}_{2}+\mathrm{H}_{2} \longrightarrow \mathrm{C}_{2}\mathrm{H}_{4} \mathrm{~F}_{2}$$
Compound A gives product D (molecular formula $\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{~F}_{2}$ ), compound B gives product $E$ (same molecular formula, $\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{~F}_{2}$, but different from product $\mathrm{D}$ ), and compound $\mathrm{C}$ gives product D. Identify the structural formulas of compounds $\mathrm{A}$, $\mathrm{B}$, and $\mathrm{C}$.

Lottie Adams

Numerade Educator

01:14

Problem 95

Obtain a Lewis formula of $\mathrm{H}_{2} \mathrm{~N}_{2}$. This molecule exhibits isomerism (it has two isomers). Sketch the electron-dot formulas of these isomers. Describe the bonding in terms of valence bond theory, and explain this isomerism in terms of this bonding description.

Lottie Adams

Numerade Educator

01:46

Problem 96

Dioxygen, $\mathrm{O}_{2}$, has sometimes been represented by the electron-dot formula
Note, however, that this formula implies that all electrons are paired. Give the molecular orbital description of the state of the $\mathrm{O}_{2}$ molecule of lowest energy in which all electrons are paired in this way. Compare this configuration with that of the ground state of $\mathrm{O}_{2} .$ In what way are they similar, and in what way are they different?

Lottie Adams

Numerade Educator

Problem 97

Solid sulfur normally consists of crystals of $\mathrm{S}_{8}$ molecules, but when heated strongly, the solid vaporizes to give $\mathrm{S}_{2}$ molecules (among other molecular species). Describe the bonding in $\mathrm{S}_{2}$ in molecular orbital terms, assuming the orbitals are analogous to those of the preceding period. What would you expect to happen to the sulfur-sulfur bond length if two electrons were added to give the $\mathrm{S}_{2}^{2-}$ ion? What would you expect to happen to the bond length if, instead, two electrons were taken away to give $\mathrm{S}_{2}{ }^{2+}$ ?

Lottie Adams

Numerade Educator

Problem 98

Formaldehyde has the formula $\mathrm{H}_{2} \mathrm{CO}$. Describe the bonding between the $\mathrm{C}$ and $\mathrm{O}$ atoms in VB terms. Formaldehyde absorbs UV light at about $270 \mathrm{~nm}$. The absorption is described as a $\pi$ to $\pi^{*}$ transition. Explain in molecular orbital terms what is involved.

Lottie Adams

Numerade Educator

01:39

Problem 99

A molecular compound is composed of $60.4 \%$ Xe, $22.1 \% \mathrm{O}$, and $17.5 \% \mathrm{~F}$, by mass. If the molecular weight is $217.3$ amu, what is the molecular formula? What is the Lewis formula? Predict the molecular geometry using the VSEPR model. Describe the bonding, using valence bond theory.

Lottie Adams

Numerade Educator

Problem 100

A molecular compound is composed of $58.8 \% \mathrm{Xe}, 7.2 \%$ $\mathrm{O}$, and $34.0 \% \mathrm{~F}$, by mass. If the molecular weight is $223 \mathrm{amu}$, what is the molecular formula? What is the Lewis formula? Predict the molecular geometry using the VSEPR model. Describe the bonding, using valence bond theory.

Lottie Adams

Numerade Educator

01:16

Problem 101

A compound of chlorine and fluorine, $\mathrm{CIF}_{n}$, reacts at about $75^{\circ} \mathrm{C}$ with uranium metal to produce uranium hexafluoride, UF $_{6}$, and chlorine monofluoride, $\mathrm{CIF}(g) .$ A quantity of uranium produced $3.53 \mathrm{~g} \mathrm{UF}_{6}$ and $343 \mathrm{~mL}$ CIF at $75^{\circ} \mathrm{C}$ and $2.50 \mathrm{~atm} .$ What is the formula $(n)$ of the compound? Describe the bonding in the molecule, using valence bond theory.

Lottie Adams

Numerade Educator

01:05

Problem 102

Excess fluorine, $\mathrm{F}_{2}(g)$, reacts at $150^{\circ} \mathrm{C}$ with bromine, $\mathrm{Br}_{2}(g)$, to give a compound $\mathrm{BrF}_{n}$. If $423 \mathrm{~mL} \mathrm{Br}_{2}(g)$ at $150^{\circ} \mathrm{C}$ and
$748 \mathrm{mmHg}$ produced $4.20 \mathrm{~g} \mathrm{BrF}_{n}$, what is $n ?$ Describe the bonding in the molecule, using valence bond theory.

Lottie Adams

Numerade Educator

01:43

Problem 103

Each of the following compounds has a nitrogen-nitrogen bond: $\mathrm{N}_{2}, \mathrm{~N}_{2} \mathrm{H}_{4}, \mathrm{~N}_{2} \mathrm{~F}_{2} .$ Match each compound with one of the
following bond lengths: $110 \mathrm{pm}, 122 \mathrm{pm}, 145 \mathrm{pm} .$ Describe the geometry about one of the $\mathrm{N}$ atoms in each compound. What hybrid orbitals are needed to describe the bonding in valence bond theory?

Lottie Adams

Numerade Educator