Chapter 10

Draw the Lewis structure for chloroform, CHCl \(_{3}\). What are its electron- pair and molecular geometries? What orbitals on \(\mathrm{C}, \mathrm{H},\) and $$\mathrm{Cl}$ overlap to form bonds involving these elements?

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?

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

Specify the electron-pair and molecular geometry for each of the following. Describe the hybrid orbital set used by the underlined atom in each molecule or ion. (a) \(\underline{\mathrm{CSe}}_{2}\) (b) \(\underline{\mathrm{SO}}_{2}\) (c) \(\underline{\mathrm{CH}}_{2} \mathrm{O}\) (d) \(\underline{\mathrm{NH}}_{4}^{+}\)

Draw the Lewis structure and then specify the electronpair and molecular geometries for each of the following molecules or ions. Identify the hybridization of the central atom. (a) \(\mathrm{SiF}_{6}^{2-}\) (b) \(\mathrm{SeF}_{4}\) (c) \(\mathrm{ICl}_{2}^{-}\) (d) \(\mathrm{XeF}_{4}\)

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

Draw the Lewis structures of the acid 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 the O atom in the acid.)

Draw the Lewis structures of HSO \(_{3} \mathrm{F}\) and \(\mathrm{SO}_{3} \mathrm{F}^{-}\). What is the molecular geometry and hybridization for the sulfur atom in each species? (H is bonded to the 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 sulfur atom in sulfuryl fluoride, \(\mathrm{SO}_{2} \mathrm{F}_{2} ?\)

The hydrogen molecular ion, \(\mathrm{H}_{2}^{+},\) can be detected spectroscopically. Write the electron configuration of the ion in molecular orbital terms. What is the bond order of the ion? Is the hydrogen-hydrogen bond stronger or weaker in \(\mathrm{H}_{2}^{+}\) than \(\operatorname{in} \mathrm{H}_{2} ?\)

Give the electron configurations for the ions \(\mathrm{Li}_{2}^{+}\) and \(\mathrm{Li}_{2}^{-}\) in molecular orbital terms. Compare the \(\mathrm{Li}-\mathrm{Li}\) bond order in these ions with the bond order in \(\mathrm{Li}_{2}\).

Oxygen, \(\mathrm{O}_{2},\) can acquire one or two electrons to give \(\mathrm{O}_{2}^{-}\) (superoxide ion) or \(\mathrm{O}_{2}^{2-}\) (peroxide ion). Write the electron configuration for the ions in molecular orbital terms, and then compare them with the \(\mathrm{O}_{2}\) molecule on the following bases. (a) magnetic character (b) net number of \(\sigma\) and \(\pi\) bonds (c) bond order (d) oxygen-oxygen bond length

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

Describe the \(\mathrm{O}-\mathrm{S}-\mathrm{O}\) angle and the hybrid orbital set used by sulfur in each of the following molecules or ions: (a) \(\mathrm{SO}_{2}\) (c) \(\mathrm{SO}_{3}^{2-}\) (b) \(\mathrm{SO}_{3}\) (d) \(\mathrm{SO}_{4}^{2-}\) Do all have the same value for the \(\mathrm{O}-\mathrm{S}-\mathrm{O}\) angle? Does the S atom in all these species use the same hybrid orbitals?

Sketch the Lewis structures of \(\mathrm{ClF}_{2}^{+}\) and \(\mathrm{ClF}_{2}^{-} .\) What are the electron-pair and molecular geometries of each ion? Do both have the same \(\mathrm{F}-\mathrm{Cl}-\mathrm{F}\) angle? What hybrid orbital set is used by Cl 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 \(\mathbf{O}-\mathbf{N}-\mathbf{O}\) bond angle, the average NO bond order, and the \(\mathrm{N}\) atom hybridization.

Sketch the resonance structures for the nitrate ion, \(\mathrm{NO}_{3}^{-}\). Is the hybridization of the N atom the same or different in each structure? Describe the orbitals involved in bond formation by the central \(\mathrm{N}\) atom.

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 \(\mathrm{O}-\mathrm{C}-\) O bond angles, the CO bond order, and the C atom hybridization.

The sulfamate ion, \(\mathrm{H}_{2} \mathrm{NSO}_{3}^{-},\) can be thought of as having been formed from the amide ion, \(\mathrm{NH}_{2}^{-},\) and sulfur trioxide, \(\mathrm{SO}_{3}\) (a) Sketch a structure for the sulfamate ion and estimate the bond angles. (b) What changes in hybridization do you expect for N and S in the course of the reaction \(\mathrm{NH}_{2}^{-}+\mathrm{SO}_{3} \longrightarrow \mathrm{H}_{2} \mathrm{N}-\mathrm{SO}_{3}^{-} ?\)

Iodine and oxygen form a complex series of ions, among them \(\mathrm{IO}_{4}^{-}\) and \(\mathrm{IO}_{5}^{3-} .\) Draw the Lewis structures for these ions, and specify their electron-pair geometries and the shapes of the ions. What is the hybridization of the I atom in these ions?

Antimony pentafluoride reacts with HF according to the equation $$2 \mathrm{HF}+\mathrm{SbF}_{5} \longrightarrow\left[\mathrm{H}_{2} \mathrm{F}\right]^{+}\left[\mathrm{SbF}_{6}\right]^{-}$$ (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. 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, \(\mathrm{O}_{2}^{2-}\) (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 \(\mathrm{O}_{2}^{2-} .\) 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 \(\mathrm{O}_{2}^{2-} ?\)

Nitrogen, \(\mathrm{N}_{2},\) can ionize to form \(\mathrm{N}_{2}^{+}\) 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 homonuclear, 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?

The elements of the second period from boron to oxygen form compounds of the type \(\mathrm{X}_{n} \mathrm{E}-\mathrm{EX}_{n},\) where \(\mathrm{X}\) can be H or a halogen. Sketch possible molecular 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 hybridizations of \(\mathrm{E}\) in each molecule and specify approximate \(\mathrm{X}-\mathrm{E}-\mathrm{E}\) bond angles.

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 \(\mathrm{BF}_{4}^{-}\). (b) Are \(\mathrm{SiF}_{4}\) and \(\mathrm{SF}_{4}\) 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 \(\left(\mathrm{C}_{2} \mathrm{H}_{6}\right),\) ethylene \(\left(\mathrm{C}_{2} \mathrm{H}_{4}\right)\) and acetylene \(\left(\mathrm{C}_{2} \mathrm{H}_{2}\right)\) as examples.

When is it desirable to use MO theory rather than valence bond theory?

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