Chapter 9

Draw the lewis structure for \(\mathrm{NF}_{3}\). What are its electron-pair 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, \(\mathrm{H}_{2} \mathrm{NOH} .\) 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 \([\left(\mathrm{CH}_{3}\right)_{2} \mathrm{NNH}_{2}, \mathrm{a}\) compound used as a rocket fuel]. 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, COF \(_{2}\). What are the electron-pair geometry and molecular geometry around the central atom? 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 the electron-pair geometry and molecular geometry around the two C atoms? What is the hybridization of each of the \(\mathrm{C}\) atoms? 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) \(\underline{\mathrm{BBr}}_{3}\) (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 underlined atom in the following list. Describe the hybrid orbital set used by this atom in each molecule or ion. $$\begin{aligned} &\text { (a) } \underline{\mathrm{CSe}_{2}}\\\ &\text { (b) } \underline{\mathbf{S O}_{2}} \end{aligned}$$ (c) \(\underline{\mathrm{CH}_{2} \mathrm{O}}\) (d) \(\underline{\mathrm{NH}}_{4}^{+}\)

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

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

What is the hybridization of the carbon atoms in benzene, \(\mathrm{C}_{6} \mathrm{H}_{6}\) ? 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, \(\mathrm{SOCl}_{2} ?\) What is the hybridization of sulfur in this compound?

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

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 in \(\mathrm{H}_{2} ?\)

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

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 \(\mathrm{O}_{2}\) to obtain \(\mathrm{O}_{2}^{+} ?\) Is the \(\mathrm{O}_{2}^{+}\) 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, \(\mathrm{O}_{2}^{2-}\) Write the electron configuration for this ion in molecular orbital terms, and draw the electron dot structure. (a) Compare the ion with the \(\mathrm{O}_{2}\) molecule with respect to the following: magnetic character, net number of \(\sigma\) and \(\pi\) bonds, bond order, and oxygen-oxygen bond length. (b) Compare the valence bond and MO pictures with regard to the number of \(\sigma\) and \(\pi\) bonds and the bond order.

Among the following, which has the shortest bond and which has the longest: \(\mathrm{Li}_{2}, \mathrm{B}_{2}, \mathrm{C}_{2}, \mathrm{N}_{2}, \mathrm{O}_{2} ?\)

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

Draw the Lewis structure for \(\mathrm{AlF}_{4} .\) What are its electron-pair and molecular geometries? What orbitals on Al and F overlap to form bonds between these elements? What are the formal charges on the atoms? Is this a reasonable charge distribution?

What is 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) \(\mathrm{SO}_{2}\) (b) \(\mathrm{SO}_{3}\) (c) \(\mathrm{SO}_{3}^{2-}\) (d) \(\mathrm{SO}_{4}^{2-}\) Do all have the same value for the \(\mathrm{O}-\mathrm{S}-\mathrm{O}\) angle? Does the \(\mathrm{S}\) atom in all these species use the same hybrid orbitals?

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 N atom hybridization.

Sketch the resonance structures for the nitrate ion, \(\mathrm{NO}_{3}^{-} .\) Is the hybridization of the \(\mathrm{N}\) 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 \(\mathrm{O}-\mathrm{C}-\mathrm{O}\) bond angles, the CO bond order, and the C atom hybridization.

Acrolein, a component of photochemical smog, has a pungent odor and irritates eyes and mucous membranes. (a) What are the hybridizations of carbon atoms 1 and \(2 ?\) (b) What are the approximate values of angles \(A\) \(B,\) and \(C ?\) (c) Is cis-trans isomerism possible here?

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 \(\mathrm{X}_{n} \mathrm{E}-\mathrm{EX}_{m}\) 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 hybridizations of \(E\) in each molecule and specify approximate \(X-E-E\) bond angles.

Suppose you carry out the following reaction of ammonia and boron trifluoride in the laboratory. (a) What is the geometry of the boron atom in \(\mathrm{BF}_{3} ? \mathrm{In} \mathrm{H}_{3} \mathrm{~N} \rightarrow \mathrm{BF}_{3} ?\) (b) What is the hybridization of the boron atom in the two compounds? (c) Considering the structures and bonding of \(\mathrm{NH}_{3}\) and \(\mathrm{BF}_{3}\), why do you expect the nitrogen on \(\mathrm{NH}_{3}\) to donate an electron pair to the \(\mathrm{B}\) atom of \(\mathrm{BF}_{3} ?\) (d) \(\mathrm{BF}_{3}\) also reacts readily with water. Based on the ammonia reaction above, speculate on how water can interact with \(\mathrm{BF}_{3}\).

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 geometry 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 \(S-O\) bonds in this compound?

Draw a Lewis structure for diimide, H-N=N-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 \(\mathrm{BF}_{4}\) (b) Are \(\operatorname{sir}_{4}\) and \(\mathrm{SF}_{4}\) isoelectronic? (c) What is the hybridization of the central atom in \(\mathrm{BF}_{4}^{-}\) and \(\mathrm{SiF}_{4} ?\)

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.

Show how valence bond theory and molecular orbital theory rationalize the O-O bond order of 1.5 in 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?

Melamine is an important industrial chemical, used to make fertilizers and plastics. (a) The carbon-nitrogen bond lengths in the ring are all the same length (about \(140 \mathrm{pm}\) ). Explain. (b) Melamine is made by the decomposition of urea, \(\left(\mathrm{H}_{2} \mathrm{N}\right)_{2} \mathrm{CO}\) \(6\left(\mathrm{H}_{2} \mathrm{N}\right)_{2} \mathrm{CO}(\mathrm{s}) \rightarrow \mathrm{C}_{3} \mathrm{H}_{6} \mathrm{N}_{6}(\mathrm{s})+6 \mathrm{NH}_{3}(\mathrm{g})+3 \mathrm{CO}_{2}(\mathrm{g})\) Calculate the enthalpy change for this reaction. Is it endo- or exothermic? \(\left[\Delta_{f} H^{\circ} \text { for melamine(s) }=\right.\) \(-66.1 \mathrm{kJ} / \mathrm{mol}\) and for urea(s) \(=-333.1 \mathrm{kJ} / \mathrm{moll}\)

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 9.16 applies to BrO, write its electron configuration (where Br uses \(4 s\) and \(4 p\) orbitals). What is the highest occupied molecular orbital (HOMO) for the molecule?

The following problem is taken from the Theoretical Examination of the 44 th annual International Chemistry Olympiad in \(2012,\) a competition attended by four secondary school students from each of about 70 countries. (Used with permission. Graphene is a sheet of carbon atoms arranged in a two-dimensional honeycomb pattern. It can be considered as an extreme case of a polyaromatic hydrocarbon with essentially infinite length in two dimensions. Graphene has remarkable strength, flexibility, and electrical properties. The Nobel Prize for Physics was awarded in 2010 to Andre Geim and Konstantin Novoselov for groundbreaking experiments on graphene. A section of the graphene sheet is shown below. The area of one hexagonal 6 -carbon unit is \(-52400 \mathrm{pm}^{2} .\) Calculate the number of \(\pi\) electrons in a tiny \(25 \mathrm{nm} \times 25 \mathrm{nm}\) sheet of graphene. For this problem you can ignore edge electrons (i.e., those outside the full hexagons in the picture). (IMAGE CAN'T COPY)