Chapter 9

Why is the shape of a molecule determined by repulsions between electron pairs and not by repulsions between nuclei?

Do all resonance forms of a molecule have the same molecular geometry? Explain your answer.

How can \(\mathrm{SO}_{3}\) and \(\mathrm{BF}_{3}\) have different numbers of bonds but the same trigonal planar geometry?

Account for the range of bond angles from less than \(100^{\circ}\) to \(180^{\circ}\) in triatomic molecules.

In a molecule of ammonia, why is the repulsion between the lone pair and a bonding pair of electrons on nitrogen greater than the repulsion between two N-H bonding pairs?

Why is it important to draw a correct Lewis structure for a molecule before predicting its geometry?

Why does the seesaw structure have lower energy than a trigonal pyramidal structure derived by removing an axial atom from a trigonal bipyramidal \(\mathrm{AB}_{5}\) molecule?

Which geometry do you predict will have lower energy: a square pyramid or a trigonal bipyramid? Why?

Arrange the following molecular geometries in order of increasing bond angle: (a) trigonal planar; (b) octahedral; (c) tetrahedral.

Arrange the following molecular geometries in order of increasing bond angle: (a) square planar; (b) tetrahedral; (c) square pyramidal.

Which of the molecular geometries discussed in this chapter have more than one characteristic bond angle?

Which molecular geometries for molecules of the general formula \(\mathrm{AB}_{x}(x=2 \text { to } 6\) ) discussed in this chapter have the same bond angles when lone pairs replace one or more atoms?

Which of the following molecular geometries does not lead to linear triatomic molecules after the removal of one or more atoms? (a) tetrahedral; (b) octahedral; (c) T-shaped

Which of the following molecular geometries does not lead to linear triatomic molecules after the removal of one or more atoms? (a) trigonal bipyramidal; (b) seesaw; (c) trigonal planar

Determine the molecular geometries of the following molecules: (a) \(\mathrm{GeH}_{4}\); (b) \(\mathrm{PH}_{3} ;\) (c) \(\mathrm{H}_{2} \mathrm{S} ;\) (d) \(\mathrm{CHCl}_{3}\)

Determine the molecular geometries of the following molecules and ions: (a) \(\mathrm{NO}_{3}^{-} ;\) (b) \(\mathrm{NO}_{4}^{3-}\) (c) \(\mathrm{S}_{2} \mathrm{O} ;\) (d) \(\mathrm{NF}_{3}\)

Determine the bond angles in the following ions: (a) \(\mathrm{NH}_{4}^{+}\) (b) \(\mathrm{SO}_{3}^{2-} ;\) (c) \(\mathrm{NO}_{2}^{-} ;\) (d) \(\mathrm{XeF}_{5}^{+}\)

Determine the bond angles in the following ions: (a) \(\mathrm{SCN}^{-}\) (b) \(\mathrm{BF}_{2}^{+} ;\) (c) \(\mathrm{ICl}_{2}^{-} ;\) (d) \(\mathrm{PO}_{3}^{3-}\)

Determine the geometries of the following ions: (a) \(\mathrm{S}_{2} \mathrm{O}_{3}^{2-}\) (b) \(\mathrm{PO}_{4}^{3-} ;\) (c) \(\mathrm{NO}_{3}^{-} ;\) (d) \(\mathrm{NCO}^{-}\)

Determine the geometries of the following molecules: (a) \(\mathrm{ClO}_{2} ;\) (b) \(\mathrm{ClO}_{3} ;(\mathrm{c}) \mathrm{IF}_{3} ;(\mathrm{d}) \mathrm{SF}_{4}\)

The anion \(\mathrm{C}(\mathrm{CN})_{3}\). has a trigonal planar geometry about the central carbon atom. Draw Lewis structures for \(\mathrm{C}(\mathrm{CN})_{3}^{-},\) including resonance forms, and determine which structure contributes the most to the bonding.

The anion \(\mathrm{C}\left(\mathrm{NO}_{2}\right)_{3}\) - has a trigonal planar geometry about the carbon atom. Draw Lewis structures for \(\mathrm{C}\left(\mathrm{NO}_{2}\right)_{3}^{-}\) including resonance forms, and determine which structure contributes the most to the bonding.

The geometry about nitrogen in \(\mathrm{N}\left(\mathrm{CF}_{3}\right)_{3}\) and \(\mathrm{N}\left(\mathrm{SCF}_{3}\right)_{3}\) is trigonal planar for both complexes, as shown in Figure P9.36. Draw Lewis structures for each that are consistent with the observed geometry. (Hint: For \(\mathrm{N}\left(\mathrm{CF}_{3}\right)_{3}\) consider an ionic form \(\left[\left(\mathrm{CF}_{3}\right)_{2} \mathrm{NCF}_{2}\right]^{+}[\mathrm{F}]^{-}\) )

For many years, it was believed that the noble gases could not form covalently bonded compounds. However, xenon reacts with fluorine and oxygen. Reaction between xenon tetrafluoride and fluoride ions produces the pentafluoroxenate anion: $$\mathrm{XeF}_{4}+\mathrm{F}^{-} \rightarrow \mathrm{XeF}_{5}^{-}$$ Draw Lewis structures for \(\mathrm{XeF}_{4}\) and \(\mathrm{XeF}_{5}^{-}\), and predict the geometry around xenon in \(\mathrm{XeF}_{4} .\) The crystal structure of \(\mathrm{XeF}_{5}^{-}\) compounds indicates a pentagonal bipyramidal orientation of valence pairs around Xe. Sketch the structure for \(\mathrm{XeF}_{5}^{-}.\)

The first compound containing a xenon-sulfur bond was isolated in \(1998 .\) Draw a Lewis structure for HXeSH and determine its molecular geometry at Xe.

The Cl-O distances in \(\mathrm{ClO}_{2}^{+}, \mathrm{ClO}_{2},\) and \(\mathrm{ClO}_{2}^{-}\) are found to be \(131 \mathrm{pm}, 147 \mathrm{pm},\) and \(156 \mathrm{pm},\) respectively. The corresponding \(\mathrm{O}-\mathrm{Cl}-\mathrm{O}\) bond angles are \(122^{\circ}, 118^{\circ},\) and \(110^{\circ} .\) Draw Lewis structures consistent with these data.

Explain the difference between a polar bond and a polar molecule.

Must a polar molecule contain polar covalent bonds? Why or why not?

Can a nonpolar molecule contain polar covalent bonds?

Consider the following molecules: (a) \(\mathrm{CCl}_{4} ;\) (b) \(\mathrm{CHCl}_{3}\) (c) \(\mathrm{CO}_{2} ;\) (d) \(\mathrm{H}_{2} \mathrm{S} ;\) (e) \(\mathrm{SO}_{2}\) a. Which of them contain polar bonds? b. Which are polar molecules? c. Which are nonpolar molecules?

Freon Ban Compounds containing carbon, chlorine, and fluorine are known as Freons or chlorofluorocarbons (CFCs). Widespread use of these substances was banned because of their effect on the ozone layer in the upper atmosphere. Which of the following CFCs are polar and which are nonpolar? (a) Freon \(11\left(\mathrm{CFCl}_{3}\right) ;\) (b) Freon 12 \(\left(\mathrm{CF}_{2} \mathrm{Cl}_{2}\right) ;(\mathrm{c})\) Freon \(113\left(\mathrm{Cl}_{2} \mathrm{FCCF}_{2} \mathrm{Cl}\right).\)

Which of the following chlorofluorocarbons (CFCs) are polar and which are nonpolar? (a) Freon \(\mathrm{C} 318\) (C \(_{4} \mathrm{F}_{8}\), cyclic structure); (b) Freon \(1113\left(\mathrm{C}_{2} \mathrm{ClF}_{3}\right) ;\) (c) \(\mathrm{Cl}_{2} \mathrm{HCCClF}_{2}.\)

Which molecule in each of the following pairs has the larger dipole moment? (a) \(\mathrm{BF}_{3}\) or \(\mathrm{BCl}_{3} ;\) (b) \(\mathrm{BCl}_{2} \mathrm{F}\) or \(\mathrm{BClF}_{2}\)

Cleaning Silicon Chips Nitrogen trifluoride, \(\mathrm{NF}_{3}\), is used in the electronics industry to clean surfaces. NF \(_{3}\) is also a potent greenhouse gas. a. Draw the Lewis structure of \(\mathrm{NF}_{3}\) and determine its molecular geometry. b. \(\mathrm{BF}_{3}\) and \(\mathrm{NF}_{3}\) both have three covalently bonded fluorine atoms around a central atom. Do they have the same dipole moment? c. Could BF \(_{3}\) also behave as a greenhouse gas?

Describe in your own words the differences between sigma and pi bonds.

Why aren't the orbitals on isolated atoms hybridized?

The bond angles in \(\mathrm{PF}_{3}\) are \(97.8^{\circ}.\) a. Explain the size of this angle by using hybrid orbitals. b. Explain the size of this angle without using hybrid orbitals.

What combination of \(s, p,\) and \(d\) orbitals would we need to form four \(\sigma\) and two \(\pi\) bonds to a sulfur atom?

What is the hybridization of nitrogen in each of the following ions and molecules? (a) \(\mathrm{NO}_{2}^{+} ;\) (b) \(\mathrm{NO}_{2}^{-}\) (c) \(\mathrm{N}_{2} \mathrm{O} ;\) (d) \(\mathrm{N}_{2} \mathrm{O}_{5} ;\) (e) \(\mathrm{N}_{2} \mathrm{O}_{3}\)

Airbags Azides such as sodium azide, NaN \(_{3},\) are used in automobile airbags as a source of nitrogen gas. Another compound with three nitrogen atoms bonded together is \(\mathrm{N}_{3} \mathrm{F} .\) What differences are there in the arrangement of the electrons around the nitrogen atoms in the azide ion \(\left(\overline{\mathrm{N}}_{3}^{-}\right)\) and \(\mathrm{N}_{3} \mathrm{F}\) ? Is there a difference in the hybridization of the central nitrogen atom?

How does the hybridization of the sulfur atom change in the series \(\mathrm{SF}_{2}, \mathrm{SF}_{4},\) and \(\mathrm{SF}_{6} ?\)

How does the hybridization of the central atom change in the series \(\mathrm{CO}_{2}, \mathrm{NO}_{2}, \mathrm{O}_{3},\) and \(\mathrm{ClO}_{2} ?\)

Draw the Lewis structure of the chlorite ion, \(\mathrm{ClO}_{2}^{-}\) which is used as a bleaching agent. Include all resonance structures in which formal charges are closest to zero. What is the shape of the ion? Suggest a hybridization scheme for the central chlorine atom that accounts for the structures you have drawn.

Perchlorate lon and Human Health Perchlorate ion adversely affects human health by interfering with the uptake of iodine in the thyroid gland, but because of this behavior, it also provides a useful medical treatment for hyperthyroidism, or overactive thyroid. Draw the Lewis structure of the perchlorate ion, \(\mathrm{ClO}_{4}^{-} .\) Include all resonance structures in which formal charges are closest to zero. What is the shape of the ion? Suggest a hybridization scheme for the central chlorine atom that accounts for this shape.

Draw a Lewis structure for \(\mathrm{CF}_{3} \mathrm{PCF}_{2}\) where the fluorine atoms are all bonded to carbon atoms. Determine its molecular geometry at \(P\) and the hybridization of the phosphorus atom.

Synthesis of the first compound of argon was reported in 2000. HArF was made by reacting Ar with HF. Draw a Lewis structure for HArF, and determine the hybridization of Ar in this molecule.

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

The trifluorosulfate anion was isolated in 1999 as the tetramethylammonium salt \(\left[\left(\mathrm{CH}_{3}\right)_{4} \mathrm{N}\right]^{+}\left[\mathrm{SO}_{2} \mathrm{F}_{3}\right]^{-}.\) a. Determine the geometry around the nitrogen atom in the cation and describe the \(C-N\) bonding according to valence bond theory. b. The \(S-O\) bond lengths in the anion are both \(143 \mathrm{pm}\). Draw the Lewis structure that is consistent with this bond length. c. What is the molecular geometry of the anion?

What features do Lewis structures, Kekulé structures, condensed structures, and carbon-skeleton structures share in common? What features differentiate these four kinds of structures?

Explain why alkanes don't have optical isomers.