Chapter 5

The Lewis structure for \(\mathrm{N}_{4} \mathrm{O},\) with the skeletal structure O-N-N-N-N, contains one \(\mathrm{N}-\mathrm{N}\) single bond, one \(\mathrm{N}=\mathrm{N}\) double bond, and one \(\mathrm{N} \equiv \mathrm{N}\) triple bond. Is the hybridization of all the nitrogen atoms the same?

The trifluorosulfate anion was isolated in 1999 as the tetramethylammonium salt \(\left(\mathrm{CH}_{3}\right)_{4} \mathrm{NSOF}_{3}\). Determine the geometry around sulfur in the anion and describe the bonding according to valence bond theory.

Can molecules with more than one central atom have resonance forms?

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

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

Can hybrid orbitals be associated with more than one atom?

Which of the following objects are chiral? (a) a golf club; (b) a spoon; (c) a glove; (d) a shoe

Which of the following objects are chiral? (a) a key; (b) a screwdriver; (c) a fluorescent coil lightbulb; (d) a baseball

Could an sp hybridized carbon atom be a chiral center? Explain your answer.

Two compounds have the same Lewis structure and the same optical activity. Are they enantiomers or the same compound?

Are racemic mixtures homogeneous or heterogeneous?

Can a mixture of enantiomers rotate plane-polarized light? Explain your answer.

Which better explains the visible emission spectra of molecular substances: valence bond theory or molecular orbital theory?

Which better explains the magnetic properties of molecular substances: valence bond theory or molecular orbital theory?

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

Are s atomic orbitals with different principal quantum numbers \((n)\) as likely to overlap and form MOs as s atomic orbitals with the same value of \(n ?\)

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

Make a sketch showing how two \(2 p_{y}\) orbitals overlap "sideways" to form a \(\pi_{2 p}\) bonding molecular orbital and a \(\pi_{2 p}^{*}\) antibonding molecular orbital.

Use MO theory to predict the bond orders of the following molecular ions: \(\mathrm{N}_{2}^{+}, \mathrm{O}_{2}^{+}, \mathrm{C}_{2}^{+},\) and \(\mathrm{Br}_{2}^{2-} .\) Do you expect any of the species to exist?

Diatomic noble gas molecules, such as \(\mathrm{He}_{2}\) and \(\mathrm{Ne}_{2},\) do not exist. Would removing an electron create molecular ions, such as \(\mathrm{He}_{2}^{+}\) and \(\mathrm{Ne}_{2}^{+},\) that are more stable than \(\mathrm{He}_{2}\) and \(\mathrm{Ne}_{2} ?\)

Which of the following molecular ions are paramagnetic? (a) \(\mathrm{N}_{2}^{+} ;\) (b) \(\mathrm{O}_{2}^{+} ;\) (c) \(\mathrm{C}_{2}^{2+} ;\) (d) \(\mathrm{Br}_{2}^{2-}\)

Which of the following molecular ions are paramagnetic? (a) \(\mathrm{N}_{2}^{+} ;\) (b) \(\mathrm{O}_{2}^{+} ;\) (c) \(\mathrm{C}_{2}^{2+} ;\) (d) \(\mathrm{Br}_{2}^{2-}\)

Which of the following molecular anions have electrons in \(\pi\) antibonding orbitals? (a) \(\mathrm{C}_{2}^{2-} ;\) (b) \(\mathrm{N}_{2}^{2-} ;\) (c) \(\mathrm{O}_{2}^{2-} ;\) (d) \(\mathrm{Br}_{2}^{2-}\)

Which of the following molecular cations have electrons in \(\pi\) antibonding orbitals? (a) \(\mathrm{N}_{2}^{+} ;\) (b) \(\mathrm{O}_{2}^{+} ;\) (c) \(\mathrm{C}_{2}^{2+} ;\) (d) \(\mathrm{Br}_{2}^{2+}\)

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

For which of the following diatomic molecules does the bond order increase with the loss of two electrons, forming the corresponding \(2+\) cation? (a) \(\mathrm{B}_{2} ;\) (b) \(\mathrm{C}_{2} ;\) (c) \(\mathrm{N}_{2} ;\) (d) \(\mathrm{O}_{2}.\)

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 \(1-\) and \(2-\) anions.

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

By December \(31,2003,\) concerns over arsenic contamination had prompted the manufacturers of pressure-treated lumber to voluntarily cease producing lumber treated with CCA (chromated copper arsenate) for residential use. CCA-treated lumber has a light greenish color and was widely used to build decks, sand boxes, and playground structures. Draw the Lewis structure for the arsenate ion \(\left(\mathrm{As} \mathrm{O}_{4}^{3-}\right)\) that yields the most favorable formal charges. Predict the angles between the arsenic-oxygen bonds in the arsenate anion.

Fluoroaluminate anions \(\mathrm{AIF}_{4}^{-}\) and \(\mathrm{AIF}_{6}^{3-}\) have been known for over a century, but the structure of the pentafluoroaluminate ion, AlFs \(_{5}^{2-},\) was not determined until 2003. Draw the Lewis structures for AlF \(_{3},\) AlF \(_{4}^{-},\) AlF \(_{5}^{2-}\), and 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 \(\mathrm{AlF}_{6}^{3-}\) using valence bond theory.

Rocket Fuel Hydrazine, \(\mathrm{NH}_{2} \mathrm{NH}_{2},\) fueled the spacecraft that delivered the Mars rover Curiosity to the planet's surface in 2012 a. Draw the Lewis structure of hydrazine. b. What is the hybridization of its \(\mathrm{N}\) atoms? c. Is hydrazine a polar compound? Explain why or why not.

Two compounds formed by the reaction of boron with carbon monoxide have the following skeletal structures: \(\mathrm{B}-\mathrm{B}-\mathrm{C}-\mathrm{O}\) and \(\mathrm{O}-\mathrm{C}-\mathrm{B}-\mathrm{B}-\mathrm{C}-\mathrm{O}\) a. Draw the Lewis structures of both compounds that minimize formal charges. b. What are the B-B-C bond angles in the molecules?

Boron reacts with NO, forming a compound with the formula BNO. a. Draw the Lewis structure for BNO, including any resonance forms. b. Assign formal charges and predict which structure provides the best description of the bonding in this molecule. c. Predict the molecular geometry of BNO.

Borazine, \(\mathrm{B}_{3} \mathrm{N}_{3} \mathrm{H}_{6}\) (a cyclic compound with alternating B and \(N\) atoms in the ring), is isoelectronic with benzene \(\left(\mathrm{C}_{6} \mathrm{H}_{6}\right) .\) Are there delocalized \(\pi\) electrons in borazine?

Compounds That 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, \(\mathrm{CH}_{3} \mathrm{NCS}\). a. Draw the Lewis structure for methyl isothiocyanate, including all resonance forms. Hint: The nitrogen atom is bonded to the methyl \(\left(\mathrm{CH}_{3} \longrightarrow\right)\) group. b. Assign formal charges and determine which structure is likely to contribute the most to bonding. c. Predict the molecular geometry of the molecule at both carbon atoms.

Toxic to Insects and People Methyl thiocyanate (CH \(_{3} \mathrm{SCN}\) ) is used as an agricultural pesticide and fumigant. It is slightly water soluble and is readily absorbed through the skin, but it is highly toxic if ingested. Its toxicity stems in part from its metabolism to cyanide ion. a. Draw the Lewis structure for methyl thiocyanate, including all resonance forms. b. Assign formal charges and predict which structure contributes the most to bonding. c. Predict the molecular geometry of the molecule at both carbon atoms.

Some chemists think HArF consists of \(\mathrm{H}^{+}\) ions and \(\mathrm{ArF}^{-}\) ions. Using an appropriate MO diagram, determine the bond order of the Ar-F bond in \(\mathrm{ArF}^{-}\)

To model the bonding in \(\mathrm{SF}_{6}\) gas, some chemists assume the existence of \(\mathrm{SF}_{4}^{2+}\) cations surrounded by two \(\mathrm{F}^{-}\) ions. a. Draw the Lewis structure of a \(\mathrm{SF}_{4}^{2+}\) ion. b. What are the formal charges on \(S\) and \(F\) in the structure you drew? c. What is the shape of the ion? d. Does the S atom have an expanded octet?

Which of the 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? \(\left(\mathrm{N}_{2} \mathrm{O}_{2} \text { and } \mathrm{N}_{2} \mathrm{O}_{3} \text { have } \mathrm{N}-\mathrm{N}\right.\) bonds; \(\left.\mathrm{N}_{2} \mathrm{O}_{5} \text { does not. }\right)\)

Hydrogen atoms have one electron. Does this mean that hydrogen gas is paramagnetic? Why or why not?

Draw the molecular orbital diagram of the valence shell of a \(\mathrm{F}_{2}+\) ion, and use it to determine the bond order in the ion.

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 the bond order values consistent with those predicted from Lewis structures?

Trimethylamine, \(\left(\mathrm{CH}_{3}\right)_{3} \mathrm{N},\) has a trigonal pyramidal structure, while trisilylamine, \(\left(\mathrm{SiH}_{3}\right)_{3} \mathrm{N},\) has a trigonal planar geometry. Draw Lewis structures for both compounds consistent with the observed geometries and explain your reasoning.

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

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?

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