Chapter 9

Can molecules with more than one central atom have resonance forms? Explain your answer.

Can hybrid orbitals be associated with more than one atom? Explain your answer.

Are resonance structures examples of electron delocalization? Explain your answer.

Can \(s p^{2}\) and \(s p\) hybridized carbon atoms be chiral centers? Explain your answer.

Which of the following objects are chiral? (a) a baseball bat with no lettering on it; (b) a pair of scissors; (c) a boot; (d) a fork

Why is it difficult to assign a single geometry to a molecule with more than one central atom?

Do all \(\sigma\) molecular orbitals result from the overlap of \(s\) atomic orbitals? Explain your answer.

Do all \(\pi\) molecular orbitals result from the overlap of \(p\) atomic orbitals? Explain your answer.

Which atomic orbitals are more likely to mix to form a set of molecular orbitals-a \(2 s\) and a \(3 p\) orbital or a \(4 s\) and a \(5 p\) orbital?

Why might some molecules with even numbers of valence electrons be paramagnetic?

How does the molecular orbital diagram for a homonuclear diatomic species differ from that of a heteronuclear diatomic species?

How does the sea-of-electrons model (Chapter 8 ) explain the high electrical conductivity of gold? How does band theory explain this?

Some scientists believe that the solid hydrogen that forms at very low temperatures and high pressures may conduct electricity. Is this hypothesis supported by band theory?

Describe in general terms the differences in composition and conduction between n-type and p-type semiconductors.

How might doping of silicon with germanium affect the conductivity of silicon?

Make a sketch showing how two 1 s orbitals overlap to form a \(\sigma_{1s}\) bonding molecular orbital and a \(\sigma_{1 s}^{*}\) antibonding molecular orbital.

Consider the following molecular ions: \(\mathrm{N}_{2}^{+}, \mathrm{O}_{2}^{+}, \mathrm{C}_{2}^{+},\) and \(\mathrm{Br}_{2}^{2-} .\) Using MO theory, (a) write their orbital electron configuration; (b) predict their bond orders; (c) state whether you expect any of these species to exist.

Diatomic noble gas molecules, such as \(\mathrm{He}_{2}\) and \(\mathrm{Ne}_{2},\) do not exist. a. Write their orbital electron configurations. b. Does removing one electron from each of these molecules create molecular ions $$\left(\mathrm{He}_{2}^{+} \text {and } \mathrm{Ne}_{2}^{+}\right)$$ that are more stable than \(\mathrm{He}_{2}\) and \(\mathrm{Ne}_{2} ?\)

Which of the following molecular ions is expected to have one or more unpaired electrons? (a) \(\mathrm{N}_{2}^{+} ;\) (b) \(\mathrm{O}_{2}^{+} ;\) (c) \(\mathrm{C}_{2}^{2+}\) (d) \(\mathrm{Br}_{2}^{2-} ;\) (e) \(\mathrm{O}_{2}^{-} ;\) (f) \(\mathrm{O}_{2}^{2-} ;\) (g) \(\mathrm{N}_{2}^{2-} ;\) (h) \(\mathrm{F}_{2}^{+}\)

Which of the following molecular ions have electrons in \(\pi\) antibonding orbitals? (a) \(\mathrm{O}_{2}^{-} ;\) (b) \(\mathrm{O}_{2}^{2-} ;\) (c) \(\mathrm{N}_{2}^{2-} ;\) (d) \(\mathrm{F}_{2}^{+}\) (e) \(\mathrm{N}_{2}^{+} ;\) (f) \(\mathrm{O}_{2}^{+} ;(\mathrm{g}) \mathrm{C}_{2}^{2+} ;\) (h) \(\mathrm{Br}_{2}^{2+}\)

The odd-electron molecule ClO affects the atmospheric chemistry of chlorofluorocarbons as illustrated by the reaction (where the \(^{*}\) indicates an excited-state oxygen atom): $$\mathrm{CF}_{2} \mathrm{Cl}_{2}+\mathrm{O}^{*} \rightarrow \mathrm{ClO}+\mathrm{CF}_{2} \mathrm{Cl}$$ Draw a molecular orbital diagram for ClO. Is the odd electron in a bonding or antibonding orbital?

The elusive molecule boron monoxide, \(\mathrm{BO},\) can be stabilized by bonding to platinum. Draw a molecular orbital diagram for BO. Is the odd electron in a bonding or antibonding orbital?

For which of the following diatomic molecules does the bond order increase with the gain of two electrons, forming the corresponding anion with a \(2-\) charge? a. \(\mathrm{B}_{2}+2 \mathrm{e}^{-} \rightarrow \mathrm{B}_{2}^{2-}\) b. \(C_{2}+2 e^{-} \rightarrow C_{2}^{2-}\) c. \(\mathrm{N}_{2}+2 \mathrm{e}^{-} \rightarrow \mathrm{N}_{2}^{2-}\) d. \(\mathrm{O}_{2}+2 \mathrm{e}^{-} \rightarrow \mathrm{O}_{2}^{2-}\)

For which of the following diatomic molecules does the bond order increase with the loss of two electrons, forming the corresponding cation with a \(2+\) charge? a. \(\quad B_{2} \rightarrow B_{2}^{2+}+2 e^{-}\) b. \(C_{2} \rightarrow C_{2}^{2+}+2 e^{-}\) c. \(\mathrm{N}_{2} \rightarrow \mathrm{N}_{2}^{2+}+2 \mathrm{e}^{-}\) d. \(\mathrm{O}_{2} \rightarrow \mathrm{O}_{2}^{2+}+2 \mathrm{e}^{-}\)

Do the \(1+\) cations of homonuclear diatomic molecules of the second-row elements always have shorter bond lengths than the corresponding neutral molecules?

Do any of the anions of the homonuclear diatomic molecules formed by \(\mathrm{B}, \mathrm{C}, \mathrm{N}, \mathrm{O},\) and \(\mathrm{F}\) have shorter bond lengths than those of the corresponding neutral molecules? Consider only the anions with \(1-\) or \(2-\) charge.

Thin films of doped diamond hold promise as semiconductor materials. Trace amounts of nitrogen impart a yellow color to otherwise colorless pure diamonds. a. Are nitrogen-doped diamonds examples of semiconductors that are \(p\) -type or \(n\) -type? b. Draw a picture of the band structure of diamond to indicate the difference between pure diamond and N-doped (nitrogen-doped) diamond. "c. N-doped diamonds absorb violet light at about \(425 \mathrm{nm}\) What is the magnitude of \(E_{\mathrm{g}}\) that corresponds to this wavelength?

Hope Diamond Trace amounts of boron give diamonds (including the Smithsonian's Hope Diamond) a blue color (Figure \(\mathrm{P} 9.114).\) a. Are boron-doped diamonds examples of semiconductors that are \(p\) -type or n-type? b. Draw a picture of the band structure of diamond to indicate the difference between pure diamond and B-doped diamond. c. What is the band gap in energy if blue diamonds absorb red-orange light with a wavelength of 675 nm?

Draw the Lewis structure for the two ions in ammonium perchlorate \(\left(\mathrm{NH}_{4} \mathrm{ClO}_{4}\right),\) which is used as a propellant in solid fuel rockets, and determine the molecular geometries of the two polyatomic ions.

The molecule trinitramide, \(\mathrm{N}\left(\mathrm{NO}_{2}\right)_{3},\) was first prepared in late 2010 by chemists in Sweden. Draw Lewis structures for trinitramide and predict the geometry about the central nitrogen atom. Do all resonance forms of \(\mathrm{N}\left(\mathrm{NO}_{2}\right)_{3}\) have the same geometry?

The fluoroaluminate anions \(\mathrm{AlF}_{4}^{-}\) and \(\mathrm{AlF}_{6}^{3-}\) have been known for more than a century, but the structure of the pentafluoroaluminate ion, \(\mathrm{AlF}_{5}^{2-},\) was not determined until 2003. Draw the Lewis structures for AlF \(_{3},\) AlF \(_{4}^{-}\), \(\mathrm{AlF}_{5}^{2-},\) and \(\mathrm{AlF}_{6}^{3-} .\) Determine the molecular geometry of each molecule or ion. Describe the bonding in \(\mathrm{AlF}_{3}\) \(\mathrm{AlF}_{4}^{-}, \mathrm{AlF}_{5}^{2-},\) and \(\mathrm{AlF}_{6}^{3-}\) by using valence bond theory.

Thermally unstable compounds can sometimes be synthesized using matrix isolation methods in which the compounds are isolated in a nonreactive medium such as frozen argon. The reaction of boron with carbon monoxidn produces compounds with these skeletal structures: \(\mathrm{B}-\mathrm{B}-\mathrm{C}-\mathrm{O}\) and \(\mathrm{O}-\mathrm{C}-\mathrm{B}-\mathrm{B}-\mathrm{C}-\mathrm{O} .\) For each of these compounds, draw the Lewis structure that minimizes formal charges. Do any of your structures contain atoms with incomplete octets? Predict the molecular geometries of BBCO and OCBBCO.

The products of the reaction between boron and NO can be trapped in solid argon matrices. Among the products is BNO. Draw the Lewis structure for BNO, including any resonance forms. Assign formal charges and predict which structure provides the best description of the bonding in this molecule. Do any of your structures contain atoms without complete octets? Predict the molecular geometry of BNO.

Compounds May Help Prevent Cancer Broccoli, cabbage, and kale contain compounds that break down in the human body to form isothiocyanates, whose presence may reduce the risk of certain types of cancer. The simplest isothiocyanate is methyl isothiocyanate, \(\mathrm{CH}_{3} \mathrm{NCS}\). Draw the Lewis structure for \(\mathrm{CH}_{3} \mathrm{NCS}\), including all resonance forms. Assign formal charges and determine which structure is likely to contribute the most to bonding. Predict the molecular geometry of the molecule at both carbon atoms.

Toxic to Insects and People Methyl thiocyanate \(\left(\mathrm{CH}_{3} \mathrm{SCN}\right)\) is used as an agricultural pesticide and fumigant. It is slightly water soluble and is readily absorbed through the skin; it is highly toxic if ingested. Its toxicity stems in part from its metabolism to cyanide ion. Draw three resonance structures for methyl thiocyanate. Assign formal charges and predict which structure would be the most stable. Predict the molecular geometry of the molecule at both carbon atoms.

Skunks The pungent smell of skunk spray is detected by receptors in the nose when the skunk secretes butanethiol, \(\mathrm{CH}_{3}\left(\mathrm{CH}_{2}\right)_{3} \mathrm{SH} .\) Draw a carbon- skeleton structure for butanethiol.

Unlike \(\mathrm{O}_{2},\) sulfur monoxide (SO) is highly unstable, decomposing to a mixture of \(\mathrm{S}_{2} \mathrm{O}\) and \(\mathrm{O}_{2}\) in less than 1 second. Using the \(\mathrm{O}_{2 \mathrm{s}}, \mathrm{O}_{2 p}, \mathrm{S}_{3, \mathrm{s}}\) and \(\mathrm{S}_{3 p}\) atomic orbitals, construct an approximate molecular orbital diagram for SO. Is SO diamagnetic or paramagnetic?

Which of the following unstable nitrogen oxides, \(\mathrm{N}_{2} \mathrm{O}_{2}\) \(\mathrm{N}_{2} \mathrm{O}_{5},\) and \(\mathrm{N}_{2} \mathrm{O}_{3},\) are polar molecules? \(\left(\mathrm{N}_{2} \mathrm{O}_{2} \text { and } \mathrm{N}_{2} \mathrm{O}_{3}\right.\) have \(\left.\mathrm{N}-\mathrm{N} \text { bonds; } \mathrm{N}_{2} \mathrm{O}_{5} \text { does not. }\right)\)

Explain why \(\mathrm{O}_{2}\) is paramagnetic.

Using an appropriate molecular orbital diagram, show that the bond order in the disulfide anion \(\mathrm{S}_{2}^{2-}\) is equal to \(1 .\) Is \(\mathrm{S}_{2}^{2-}\) diamagnetic or paramagnetic?

Use molecular orbital diagrams to determine the bond order of the peroxide \(\left(\mathrm{O}_{2}^{2-}\right)\) and superoxide \(\left(\mathrm{O}_{2}^{-}\right)\) ions. Are these bond order values consistent with those predicted from Lewis structures?

Elemental sulfur has several allotropic forms, including cyclic \(\mathrm{S}_{8}\) molecules. What is the orbital hybridization of sulfur atoms in this allotrope? The bond angles are about \(108^{\circ}.\)

Ozone \(\left(\mathrm{O}_{3}\right)\) has a dipole moment \((0.54 \mathrm{D}) .\) How can a molecule with only one kind of atom have a dipole moment?

The bond angle in \(\mathrm{H}_{2} \mathrm{O}\) is 104.5 \(^\circ\); the bond angles in \(\mathrm{H}_{2} \mathrm{S}\) \(\mathrm{H}_{2} \mathrm{Se},\) and \(\mathrm{H}_{2} \mathrm{Te}\) are very close to \(90^{\circ} .\) Which theory would you apply to describe the geometry in \(\mathrm{H}_{2} \mathrm{S}, \mathrm{H}_{2} \mathrm{Se},\) and \(\mathrm{H}_{2} \mathrm{Te}:\) VSEPR? Valence bond without invoking hybrid orbitals? Valence bond theory, using hybrid orbitals? Why?