Chapter 10

Identify the hybrid orbitals on the oxygen atoms that form the \(\sigma\) bonds in the following species. (a) \(\mathrm{H}_{3} \mathrm{O}^{+}\) (b) \(\mathrm{H}_{3} \mathrm{COH}\) (c) \(\mathrm{Cl}_{2} \mathrm{O}\)

Identify the hybrid orbitals on the nitrogen atoms that form the \(\sigma\) bonds in.the following species. (a) \(\mathrm{HNCl}_{2}\) (b) \(\mathrm{NO}_{3}^{-}\) (c) \(\mathrm{N}_{2} \mathrm{H}_{2}\)

What orbitals on selenium and fluorine form the bonds in \(\mathrm{SeF}_{4} ?\) What orbital holds the lone pair on selenium?

Nitrous acid has the skeleton structure HONO. What are the hybrid orbitals on the nitrogen atom and the central oxygen atom?

If the \(z\) axis is defined as the bond axis, draw a picture that shows the overlap of each of the following pairs of orbitals; then indicate whether a \(\sigma\) or \(\pi\) bond forms. (a) \(p_{z}, p_{z}\) (b) \(p_{y}, p_{y}\) (c) sp hybrid formed from \(p_{z}\) and \(s\) orbitals, \(p_{z}\)

Identify the orbitals on each of the atoms that form the bonds in \(\mathrm{H}_{3} \mathrm{CCN}\). How many \(\sigma\) bonds and \(\pi\) bonds form?

Give the hybridization of each central atom in the following molecules. (a) cyclohexene (b) phosgene, \(\mathrm{Cl}_{2} \mathrm{CO}\) (c) glycine, \(\mathrm{H}_{2} \mathrm{NC}_{(1)} \mathrm{H}_{2} \mathrm{C}_{(2)} \mathrm{OOH}\) (Note: Numbers in parentheses label each carbon atom.)

Two resonance structures can be written for \(\mathrm{NO}_{2}^{-}\). Indicate the hybridization on the central atom for each resonance form.

Orlon is produced from acrylonitrile, \(\mathrm{H}_{2} \mathrm{CCHCN}\). Draw the Lewis structure of acrylonitrile, and indicate the hybridization of each central atom.

Tetrafluoroethylene, \(\mathrm{C}_{2} \mathrm{~F}_{4},\) is used to produce Teflon. Draw the Lewis structure of tetrafluoroethylene, and indicate the hybridization of each carbon atom.

Draw the molecular orbital diagram, including the electrons, and write the electron configuration of \(\mathrm{H}_{2}^{-}\). Give the bond order and the number of unpaired electrons, if any. Is this a stable species?

Draw the molecular orbital diagram, including the electrons, and write the electron configuration of \(\mathrm{Li}_{2}\). Give the bond order and the number of unpaired electrons, if any. Is this a stable species?

Draw the molecular orbital diagram, including the electrons, and write the electron configuration of \(\mathrm{C}_{2}\). Give the bond order and the number of unpaired electrons, if any. Is this a stable species?

Write the molecular orbital electron configuration and determine the bond order and number of unpaired electrons for the following ions. (a) \(\mathrm{C}_{2}^{+}\) (b) \(\mathrm{N}_{2}^{-}\) (c) \(\mathrm{Be}_{2}^{-}\)

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

Which species, \(\mathrm{N}_{2}\) or \(\mathrm{N}_{2}^{-}\), has the higher bond order? Explain your answer.

Identify two homonuclear diatomic molecules or ions with each of the following molecular orbital electron configurations. Are these species stable? (a) \(\left(\sigma_{2 s}\right)^{2}\left(\sigma_{2 s}^{*}\right)^{2}\left(\pi_{2 p}\right)^{4}\left(\sigma_{2 p}\right)^{2}\left(\pi_{2 p}^{*}\right)^{3}\) (b) \(\left(\sigma_{2 s}\right)^{2}\left(\sigma_{2 s}^{*}\right)^{2}\left(\pi_{2 p}\right)^{4}\left(\sigma_{2 p}\right)^{2}\) (c) \(\left(\sigma_{2 s}\right)^{2}\left(\sigma_{2 s}^{*}\right)^{2}\)

Identify two homonuclear diatomic molecules or ions with each of the following molecular orbital electron configurations. Are these species stable? (a) \(\left(\sigma_{2 s}\right)^{2}\left(\sigma_{2 s}^{*}\right)^{2}\left(\pi_{2 p}\right)^{4}\left(\sigma_{2 p}\right)^{1}\) (b) \(\left(\sigma_{2 s}\right)^{2}\left(\sigma_{2 s}^{*}\right)^{2}\left(\pi_{2 p}\right)^{4}\) (c) \(\left(\sigma_{2 s}\right)^{2}\left(\sigma_{2 s}^{*}\right)^{1}\)

The nitrosyl ion, \(\mathrm{NO}^{+}\), has an interesting chemistry. (a) Is \(\mathrm{NO}^{+}\) diamagnetic or paramagnetic? If paramagnetic, how many unpaired electrons does it have? (b) Assume the molecular orbital diagram for a homonuclear diatomic molecule applies to \(\mathrm{NO}^{+} .\) What is the highest-energy molecular orbital occupied by electrons? (c) What is the nitrogen-oxygen bond order? (d) Is the \(\mathrm{N}-\mathrm{O}\) bond in \(\mathrm{NO}^{+}\) stronger or weaker than the bond in \(\mathrm{NO}\) ?

The delocalized bonding that describes \(\mathrm{O}_{3}\) also applies to \(\mathrm{NO}_{2}^{-}\). Draw the delocalized \(\pi\) molecular orbital for \(\mathrm{NO}_{2}^{-}\)

Draw the delocalized \(\pi\) orbital for benzene. Clearly indicate the atomic orbitals that form the molecular orbital.

Write one Lewis structure of \(\mathrm{N}_{2} \mathrm{O}_{5}\left(\mathrm{O}_{2} \mathrm{NONO}_{2}\right.\) skeleton structure). What are the bond angles around the central oxygen atom and the two nitrogen atoms? What is the hybridization of each?

The ions \(\mathrm{ClF}_{2}^{-}\) and \(\mathrm{ClF}_{2}^{+}\) have both been observed. Use the VSEPR model to predict the F-Cl-F bond angle in each.

A Recently, the structure of an amine compound, \(\mathrm{NR}_{3}\) \((\mathrm{R}=\) large organic group), has been determined to have C-N-C bond angles of 119.2 degrees. It is believed that the bond angles of about 109 degrees expected from the VSEPR model are not observed because of the large substituents bonded to the nitrogen atom. Given this large bond angle, what type of orbital on the nitrogen atom makes the \(\mathrm{N}-\mathrm{C} \sigma\) bonds, and in what type of orbital is the lone pair located?

Phosgene, \(\mathrm{COCl}_{2}\), is a highly toxic gas that was used in combat during World War I. It is an important intermediate in the preparation of a number of organic compounds but must be handled with extreme care. Given that carbon is the central atom in phosgene, determine the Lewis structure, the bonded-atom lone-pair arrangement, the hybridization of the carbon atom, and the polarity of the molecule.

Calcium cyanamide, CaNCN, is used both to kill weeds and as a fertilizer. Give the Lewis structure of the \(\mathrm{NCN}^{2-}\) ion and the bonded-atom lone- pair arrangement and hybridization of the carbon atom.

A Formamide, \(\mathrm{HC}(\mathrm{O}) \mathrm{NH}_{2},\) is prepared at high pressures from carbon monoxide and ammonia, and serves as an industrial solvent (the parentheses around the \(\mathrm{O}\) indicate that it is bonded only to the carbon atom and that the carbon atom is also bonded to the \(\mathrm{H}\) and the \(\mathrm{N}\) atoms). Two resonance forms (one with formal charges) can be written for formamide. Write both resonance structures, and predict the bond angles about the carbon and nitrogen atoms for each resonance form. Are they the same? Describe how the experimental determination of the \(\mathrm{H}-\mathrm{N}-\mathrm{H}\) bond angle could be used to indicate which resonance form is more important.

Draw the molecular orbital diagrams for \(\mathrm{NO}^{-}\) and \(\mathrm{NO}^{+}\). Compare the bond orders in these two ions.

Ionization energies can be determined for molecules and atoms. Draw the molecular orbital diagrams for \(\mathrm{NO}\) and \(\mathrm{CO}\), and predict which compound has the lower ionization energy.

A compound is analyzed and found to contain \(54.53 \%\) carbon, \(9.15 \%\) hydrogen, and \(36.32 \%\) oxygen by mass. A mass spectrometry experiment shows that the molar mass is \(44 \mathrm{~g} / \mathrm{mol}\). What is the molecular formula? There are two reasonable ways to draw noncyclic skeleton structures of this molecule. Draw the Lewis structure for each, indicating the bond angles and hybridization of each central atom.

The reaction of sulfur, \(\mathrm{S}_{8}\), with fluorine, \(\mathrm{F}_{2}\), yields a product with the general formula \(\mathrm{SF}_{x}\). If \(4.01 \mathrm{~g} \mathrm{~S}_{8}\) reacts with \(4.76 \mathrm{~g} \mathrm{~F}_{2}\) to yield only \(\mathrm{SF}_{x},\) what is the value of \(x ?\) Draw the Lewis structure of this compound, indicating the \(\mathrm{F}-\mathrm{S}-\mathrm{F}\) bond angles and the hybrid orbitals on sulfur.

Two compounds have the formula \(\mathrm{S}_{2} \mathrm{~F}_{2}\). Disulfur difluoride has the skeleton structure \(\mathrm{F}-\mathrm{S}-\mathrm{S}-\mathrm{F}\), whereas thiothionyl fluoride has the skeletal structure Determine Lewis structures for each compound.

Recently, the compound \(\mathrm{CF}_{3} \mathrm{SF}_{5}\) was discovered in the atmosphere and identified as a potential greenhouse gas. Assume the carbon and sulfur atoms are both central atoms, and draw the Lewis structure for this compound. What is the hybridization of each central atom and the bond angles with the surrounding atoms?

A 1.30-g sample of \(\mathrm{C}_{2} \mathrm{H}_{2}\) reacts with exactly \(1.22 \mathrm{~L} \mathrm{H}_{2}\) gas at \(27^{\circ} \mathrm{C}\) and 1.01 atm of pressure to yield a compound with the formula \(\mathrm{C}_{2} \mathrm{H}_{x}\). What is the value of \(x\), and what are the orbitals on the carbon atoms that form the \(\mathrm{C}-\mathrm{C}\) bond \((\mathrm{s}) ?\)