Chapter 9

Draw the Lewis structure for \(\mathrm{NF}_{3}\). What are its electronpair and molecular geometries? What is the hybridization of the nitrogen atom? What orbitals on \(\mathrm{N}\) and \(\mathrm{F}\) overlap to form bonds between these elements?

Draw the Lewis structure for hydroxylamine, \(\mathbf{H}_{2} \mathrm{NOH}\). What is the hybridization for nitrogen and oxygen in this molecule? What orbitals overlap to form the bond between nitrogen and oxygen?

Draw the Lewis structure for 1,1 -dimethylhydrazine \(\left[\left(\mathrm{CH}_{3}\right)_{2} \mathrm{NNH}_{2},\) a compound used as a rocket fuel. \right. What is the hybridization for the two nitrogen atoms in this molecule? What orbitals overlap to form the bond between the nitrogen atoms?

Draw the Lewis structure for carbonyl fluoride, \(\mathrm{COF}_{2}\) What are its electron-pair geometry and molecular geometry? What is the hybridization of the carbon atom? What orbitals overlap to form the \(\sigma\) and \(\pi\) bonds between carbon and oxygen?

Draw the Lewis structure for acetamide, \(\mathrm{CH}_{3} \mathrm{CONH}_{2}\) What are its electron-pair geometry and molecular geometry? What is the hybridization of the central \(\mathrm{C}\) atom? What orbitals overlap to form the \(\sigma\) and \(\pi\) bonds between carbon and oxygen?

Specify the electron-pair and molecular geometry for each underlined atom in the following list. Describe the hybrid orbital set used by this atom in each molecule or ion. (a) \(\underline{\mathrm{BBr}}_{\mathrm{s}}\) (b) \(\underline{\mathrm{CO}}_{2}\) (c) \(\underline{\mathrm{CH}}_{2} \mathrm{Cl}_{2} \quad\) (d) \(\underline{\mathrm{CO}}_{3}^{2-}\)

Draw the Lewis structure, and then specify the electron-pair and molecular geometries for each of the following molecules or ions. Identify the hybridiration of the central atom. (a) \(\operatorname{SiF}_{6}^{2-}\) (b) \(\mathrm{SeF}_{4}\) (c) \(1 \mathrm{Cl}_{2}^{-}\) (d) \(\mathrm{XeF}_{4}\)

Draw the Lewis structure, and then specify the electron-pair and molecular geometries for each of the following molecules or ions. Identify the hybridization of the central atom. (a) \(\mathrm{XeOF}_{4}\) (c) central \(\mathrm{S}\) in \(\mathrm{SOF}_{4}\) (b) \(\mathrm{BrF}_{5}\) (d) central Br in \(\mathrm{Br}_{3}^{-}\)

Draw the Lewis structures of the acid \(\mathrm{HPO}_{2} \mathrm{F}_{2}\) and its anion \(\mathrm{PO}_{2} \mathrm{F}_{2}^{-}\). What is the molecular geometry and hybridization for the phosphorus atom in each species? (H is bonded to an O atom in the acid.)

Draw the Lewis structures of HSOs \(\mathrm{F}\) and \(\mathrm{SO}_{3} \mathrm{F}^{-}\). What is the molecular geometry and hybridization for the sulfur atom in each species? ( \(\mathrm{H}\) is bonded to an \(\mathrm{O}\) atom in the acid.)

What is the hybridization of the carbon atom in phosgene, \(\mathrm{Cl}_{2}\) CO? Give a complete description of the \(\sigma\) and \pi bonding in this molecule.

What is the hybridization of the carbon atoms in benzene, \(\overline{\mathrm{C}_{6} \mathrm{H}_{6}}\) ? Describe the \(\sigma\) and \(\pi\) bonding in this compound.

What is the electron-pair and molecular geometry around the central \(S\) atom in thionyl chloride, \(\operatorname{SOCl}_{2} ?\) What is the hybridization of sulfur in this compound?

What is the electron-pair and molecular geometry around the central \(\mathrm{S}\) atom in sulfury \(\mathrm{chloride}, \mathrm{SO}_{2} \mathrm{Cl}_{2} ?\) What is the hybridization of sulfur in this compound?

Platinum hexafluoride is an extremely strong oxidizing agent. It can even oxidize oxygen, its reaction with \(\mathrm{O}_{2}\) giving \(\mathrm{O}_{2}^{+} \mathrm{PtF}_{6}^{-}\). Sketch the molecular orbital energy Level diagram for the \(\mathrm{O}_{2}^{+}\) ion. How many net \(\sigma\) and \(\pi\) bonds does the ion have? What is the oxygen-oxygen bond order? How has the bond order changed on taking away electrons from \(\mathbf{O}_{2}\) to obtain \(\mathbf{O}_{2}^{+}\) ? Is the \(\mathbf{O}_{2}^{+}\) ion paramagnetic?

When sodium and oxygen react, one of the products obtained is sodium peroxide, \(\mathrm{Na}_{2} \mathrm{O}_{2}\). The anion in this compound is the peroxide ion, \(\mathbf{O}_{2}^{2-} .\) Write the electron configuration for this ion in molecular orbital terms, and then compare it with the electron configuration of the \(\mathbf{O}_{2}\) molecule with respect to the following criteria: (a) magnetic character (b) net number of \(\sigma\) and \(\pi\) bonds (c) bond order (d) oxygen-oxygen bond length

Consider the following list of small molecules and ions: \(\mathrm{C}_{2}, \mathrm{O}_{2}^{-}, \mathrm{CN}^{-}, \mathrm{O}_{2}, \mathrm{CO}, \mathrm{NO}, \mathrm{NO}^{+}, \mathrm{C}_{2}^{2-}, \mathrm{OF}^{-} .\) Identify (a) all species that have a bond order of 3 (b) all species that are paramagnetic (c) species that have a fractional bond order

Sketch the Lewis structures of \(\mathrm{C} 1 \mathrm{F}_{2}^{+}\) and \(\mathrm{C} \mathrm{AF}_{2}^{-} .\) What are the electron-pair and molecular geometrics of each ion? Do both have the same \(\mathbf{F}-\mathbf{C} \mathbf{7}-\mathbf{F}\) angle? What hybrid orbital set is used by \(\mathrm{G}\) in each ion?

Sketch the resonance structures for the nitrite ion, \(\mathrm{NO}_{2}^{-} .\) Describe the electron-pair and molecular geometries of the ion. From these geometries, decide on the O-N-O bond angle, the average NO bond order, and the \(\mathrm{N}\) atom hybridization.

Sketch the resonance structures for the \(\mathrm{N}_{2} \mathrm{O}\) molecule. Is the hybridization of the N atoms the same or different in each structure? Describe the orbitals involved in bond formation by the central \(\mathrm{N}\) atom.

Compare the structure and bonding in \(\mathrm{CO}_{2}\) and \(\mathrm{CO}_{3}^{2-}\) with regard to the \(\mathrm{O}-\mathrm{C}-\mathrm{O}\) bond angles, the \(\mathrm{CO}\) bond order, and the \(\mathrm{C}\) atom hybridization.

Antimony pentafluoride reacts with HF according to the equation $$2\mathrm{HF}+\mathrm{SbF}_{3}\rightarrow\left[\mathrm{H}_{2}\mathrm{F}\right]^{+}\left[\mathrm{SbF}_{6}\right]^{-}$$ (a) What is the hybridization of the Sb atom in the reactant and product? (b) Draw a Lewis structure for \(\mathrm{H}_{2} \mathrm{F}^{+}\). What is the geometry of \(\mathrm{H}_{2} \mathrm{F}^{+}\) ? What is the hybridization of \(\mathrm{F}\) in \(\mathrm{H}_{2} \mathrm{F}^{+} ?\)

Xenon forms well-characterized compounds (4 page 400 ). Two xenon-oxygen compounds are \(\mathrm{XeO}_{3}\) and \(\mathrm{XeO}_{4} .\) Draw the Lewis structures of these compounds, and give their electron-pair and molecular geometries. What are the hybrid orbital sets used by xenon in these two oxides?

The simple valence bond picture of \(\mathrm{O}_{2}\) does not agree with the molecular orbital view. Compare these two theories with regard to the peroxide ion, \(\mathbf{O}_{2}^{2-}\) (a) Draw an electron dot structure for \(\mathrm{O}_{2}^{2-}\). What is the bond order of the ion? (b) Write the molecular orbital electron configuration for O \(_{2}^{2-} .\) What is the bond order based on this approach? (c) Do the two theories of bonding lead to the same magnetic character and bond order for \(\mathbf{O}_{2}^{2-} ?\)

Nitrogen, \(\mathbf{N}_{2}\), can ionize to form \(\mathbf{N}_{2}^{+}\) or add an electron to give \(\mathrm{N}_{2}\). . Using molecular orbital theory, compare these species with regard to (a) their magnetic character, (b) net number of \(\pi\) bonds, (c) bond order, (d) bond length, and (e) bond strength.

Which of the homo-nuclear, diatomic molecules of the second-period elements (from \(\mathrm{Li}_{2}\) to \(\mathrm{Ne}_{2}\) ) are paramagnetic? Which have a bond order of \(1 ?\) Which have a bond order of \(2 ?\) Which diatomic molecule has the highest bond order?

Which of the following molecules or ions are paramagnetic? What is the highest occupied molecular orbital (HOMO) in each one? Assume the molecular orbital diagram in Figure 9.18 applies to all of them. (a) NO (c) \(\mathrm{O}_{2}^{2-}\) (b) \(\mathrm{OF}^{-}\) (d) \(\mathrm{Ne}_{2}^{+}\) (e) CN

The elements of the second period from boron to oxygen form compounds of the type \(\mathrm{X}_{n} \mathrm{E}-\mathrm{EX}_{m}\), where X can be H or a halogen. Sketch possible Lewis structures for \(\mathrm{B}_{2} \mathrm{F}_{4}, \mathrm{C}_{2} \mathrm{H}_{4}, \mathrm{N}_{2} \mathrm{H}_{4},\) and \(\mathrm{O}_{2} \mathrm{H}_{2},\) Give the hybrid-izations of \(E\) in each molecule and specify approximate \(\mathbf{X}-\mathbf{E}-\mathbf{E}\) bond angles.

Draw the two resonance structures that describe the bonding in the acetate ion. What is the hybridization of the carbon atom of the \(-\mathrm{CO}_{2}^{-}\) group? Select one of the two resonance structures and identify the orbitals that overlap to form the bonds between carbon and the three elements attached to it.

Carbon dioxide \(\left(\mathrm{CO}_{2}\right),\) dinitrogen monoxide \(\left(\mathrm{N}_{2} \mathrm{O}\right)\) the azide ion \(\left(\mathrm{N}_{3}^{-}\right),\) and the cyanate ion (OCN^-) have the same arrangement of atoms and the same number of valence shell electrons. However, there are significant differences in their electronic structures. (a) What hybridization is assigned to the central atom in each species? Which orbitals overlap to form the bonds between atoms in each structure. (b) Evaluate the resonance structures of these four species. Which most closely describe the bonding in these species? Comment on the differences in bond lengths and bond orders that you expect to see based on the resonance structures.

Draw the two resonance structures that describe the bonding in \(\mathrm{SO}_{2}\). Then describe the bonding in this compound using MO theory. How does MO theory rationalize the bond order of 1.5 for the two \(\mathrm{S}-\mathrm{O}\) bonds in this compound?

Draw a Lewis structure for diimide, \(\mathbf{H}-\mathbf{N}=\mathbf{N}-\mathbf{H}\) Then, using valence bond theory, describe the bonding in this compound. What orbitals overlap to form the bond between nitrogen atoms in this compound?

What is the maximum number of hybrid orbitals that a carbon atom may form? What is the minimum number? Explain briefly.

Consider the three fluorides \(\mathrm{BF}_{4}^{-}, \mathrm{SiF}_{4},\) and \(\mathrm{SF}_{4}\) (a) Identify a molecule that is isoelectronic with \(\mathbf{B F}_{4}^{-}\) (b) Are SiF, and SF, isoelectronic? (c) What is the hybridization of the central atom in each of these species?

What is the connection between bond order, bond length, and bond energy? Use ethane \(\left(C_{2} \mathrm{H}_{6}\right),\) ethylene \(\left(\mathrm{C}_{2} \mathrm{H}_{4}\right),\) and acetylene \(\left(\mathrm{C}_{2} \mathrm{H}_{2}\right)\) as examples.

How do valence bond theory and molecular orbital theory differ in their explanation of the bond order of 1.5 for ozone?

Let's look more closely at the process of hybridization. (a) What is the relationship between the number of hybrid orbitals produced and the number of atomic orbitals used to create them? (b) Do hybrid atomic orbitals form between different p orbitals without involving s orbitals? (c) What is the relationship between the energy of hybrid atomic orbitals and the atomic orbitals from which they are formed?

Bromine and fluorine react at temperatures higher than \(150^{\circ} \mathrm{C}\) to give a compound that is \(45.69 \%\) Br and \(54.31 \% \mathrm{F}\) (a) What is the empirical formula of the compound? (b) Assuming the molecular formula of the compound is the same as its empirical formula, suggest a structure for the molecule. What is the Br atom hybrid. ization in the molecule? (c) The molecule has a small dipole moment. Does this agree with your structural proposal in (b) above? Why or why not? If it does not agree, can you propose an alternative structure?

Bromine forms a number of oxides of varying stability. (a) One oxide has \(90.90 \%\) Br and \(9.10 \%\) O. Assuming its empirical and molecular formulas are the same, draw a Lewis structure of the molecule and specify the hybridization of the central atom (O). (b) Another oxide is unstable BrO. Assuming the molecular orbital diagram in Figure 9.18 applies to BrO, write its electron configuration (where Br uses \(4 s\) and \(4 p\) orbitals). What is the HOMO for the molecule?