Chapter 10

Draw a plausible Lewis structure for the following series of molecules and ions: \((a) \operatorname{SiF}_{6}^{2-} ;\) (b) \(\mathrm{PF}_{5} ;\) (c) \(\mathrm{SF}_{4}\); (d) \(\mathrm{XeF}_{4}\). Describe the electron group geometry and molecular structure of these species.

Sketch the propyne molecule, \(\mathrm{CH}_{3} \mathrm{C} \equiv \mathrm{CH}\). Indicate the bond angles in this molecule. What is the maximum number of atoms that can be in the same plane?

Sketch the propene molecule, \(\mathrm{CH}_{3} \mathrm{CH}=\mathrm{CH}_{2}\). Indicate the bond angles in this molecule. What is the maximum number of atoms that can be in the same plane?

Levulinic acid has the formula \(\mathrm{CH}_{3}(\mathrm{CO}) \mathrm{CH}_{2} \mathrm{CH}_{2}\) COOH. Sketch the levulinic acid molecule, and indicate the various bond angles.

One of the isomers of chloromethanol has the formula \(\mathrm{ClCH}_{2} \mathrm{OH} .\) Sketch, by using the dash and wedge symbolism, this isomer of chloromethanol, and indicate the various bond angles.

Predict the shapes of the following molecules, and then predict which would have resultant dipolemoments: (a) \(\mathrm{SO}_{2} ;\) (b) \(\mathrm{NH}_{3} ;\) (c) \(\mathrm{H}_{2} \mathrm{S} ;\) (d) \(\mathrm{C}_{2} \mathrm{H}_{4} ;\) (e) \(\mathrm{SF}_{6}\); (f) \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\).

Which of the following molecules would you expect to be polar: (a) \(\mathrm{HCN} ;\) (b) \(\mathrm{SO}_{3} ;\) (c) \(\mathrm{CS}_{2} ;\) (d) OCS; (e) \(\operatorname{SOCl}_{2} ;\) (f) \(\operatorname{SiF}_{4} ;\) (g) \(\operatorname{POF}_{3}\) ? Give reasons for your conclusions.

The molecule \(\mathrm{H}_{2} \mathrm{O}_{2}\) has a resultant dipole moment of 2.2 D. Can this molecule be linear? If not, describe a shape that might account for this dipole moment.

Refer to the Integrative Example. A compound related to nitryl fluoride is nitrosyl fluoride, FNO. For this molecule, indicate (a) a plausible Lewis structure and (b) the geometric shape. (c) Explain why the measured resultant dipole moment for FNO is larger than the value for \(\mathrm{FNO}_{2}\).

Without referring to tables in the text, indicate which of the following bonds you would expect to have the greatest bond length, and give your reasons. (a) \(\mathrm{O}_{2}\); (b) \(\mathrm{N}_{2} ;\) (c) \(\mathrm{Br}_{2} ;\) (d) \(\mathrm{BrCl}\).

Estimate the lengths of the following bonds and indicate whether your estimate is likely to be too high or too low: (a) I \(-\mathrm{Cl} ;\) (b) \(\mathrm{C}-\mathrm{F}\).

In which of the following molecules would you expect the oxygen-to-oxygen bond to be the shortest: (a) \(\mathrm{H}_{2} \mathrm{O}_{2},\) (b) \(\mathrm{O}_{2},\) (c) \(\mathrm{O}_{3} ?\) Explain.

One reaction involved in the sequence of reactions leading to the destruction of ozone is $$\mathrm{NO}_{2}(\mathrm{g})+\mathrm{O}(\mathrm{g}) \longrightarrow \mathrm{NO}(\mathrm{g})+\mathrm{O}_{2}(\mathrm{g})$$ Calculate \(\Delta H^{\circ}\) for this reaction by using the thermodynamic data in Appendix D. Use your \(\Delta H^{\circ}\) value, plus data from Table \(10.3,\) to estimate the nitrogenoxygen bond energy in \(\mathrm{NO}_{2}\). [Hint: The structure of nitrogen dioxide, \(\mathrm{NO}_{2}\), is best represented as a resonance hybrid of two equivalent Lewis structures.]

The following statements are not made as carefully as they might be. Criticize each one. (a) Lewis structures with formal charges are incorrect. (b) Triatomic molecules have a planar shape. (c) Molecules in which there is an electronegativity difference between the bonded atoms are polar.

A compound consists of \(47.5 \%\) S and \(52.5 \%\) Cl, by mass. Write a Lewis structure based on the empirical formula of this compound, and comment on its deficiencies. Write a more plausible structure with the same ratio of \(\mathrm{S}\) to \(\mathrm{Cl}\).

A \(0.325 \mathrm{g}\) sample of a gaseous hydrocarbon occupies a volume of \(193 \mathrm{mL}\) at \(749 \mathrm{mmHg}\) and \(26.1^{\circ} \mathrm{C}\). Determine the molecular mass, and write a plausible condensed structural formula for this hydrocarbon.

A \(1.24 \mathrm{g}\) sample of a hydrocarbon, when completely burned in an excess of \(\mathrm{O}_{2}(\mathrm{g}),\) yields \(4.04 \mathrm{g} \mathrm{CO}_{2}\) and \(1.24 \mathrm{g} \mathrm{H}_{2} \mathrm{O} .\) Draw a plausible structural formula for the hydrocarbon molecule. [Hint: There is more than one possible arrangement of the C and H atoms.]

Draw Lewis structures for two different molecules with the formula \(\mathrm{C}_{3} \mathrm{H}_{4}\). Is either of these molecules linear? Explain.

Sodium azide, \(\mathrm{NaN}_{3}\), is the nitrogen gas-forming substance used in automobile air-bag systems. It is an ionic compound containing the azide ion, \(\mathrm{N}_{3}^{-}\). In this ion, the two nitrogen-to-nitrogen bond lengths are \(116 \mathrm{pm} .\) Describe the resonance hybrid Lewis structure of this ion.

Hydrogen azide, \(\mathrm{HN}_{3}\), is a liquid that explodes violently when subjected to physical shock. In the \(\mathrm{HN}_{3}\) molecule, one nitrogen- to-nitrogen bond length is \(113 \mathrm{pm},\) and the other is \(124 \mathrm{pm} .\) The \(\mathrm{H}-\mathrm{N}-\mathrm{N}\) bond angle is \(112^{\circ} .\) Draw Lewis structures and a sketch of the molecule consistent with these facts.

A few years ago the synthesis of a salt containing the \(\mathrm{N}_{5}^{+}\) ion was reported. What is the likely shape of this ion-linear, bent, zigzag, tetrahedral, seesaw, or square-planar? Explain your choice.

Carbon suboxide has the formula \(\mathrm{C}_{3} \mathrm{O}_{2} .\) The carbon- to-carbon bond lengths are \(130 \mathrm{pm}\) and carbon-to-oxygen, \(120 \mathrm{pm} .\) Propose a plausible Lewis structure to account for these bond lengths, and predict the shape of the molecule.

In certain polar solvents, \(\mathrm{PCl}_{5}\) undergoes an ionization reaction in which a \(\mathrm{Cl}^{-}\) ion leaves one \(\mathrm{PCl}_{5}\) molecule and attaches itself to another. The products of the ionization are \(\mathrm{PCl}_{4}^{+}\) and \(\mathrm{PCl}_{6}^{-}\). Draw a sketch showing the changes in geometric shapes that occur in this ionization (that is, give the shapes of \(\mathrm{PCl}_{5}\), \(\mathrm{PCl}_{4}^{+},\) and \(\mathrm{PCl}_{6}^{-}\) ). $$2 \mathrm{PCl}_{5} \rightleftharpoons \mathrm{PCl}_{4}^{+}+\mathrm{PCl}_{6}^{-}$$

Use the VSEPR theory to predict a probable shape of the molecule \(\mathrm{F}_{4} \mathrm{SCH}_{2}\), and explain the source of any ambiguities in your prediction.

One possibility for the electron-group geometry for seven electron groups is pentagonal-bipyramidal, as found in the IF \(_{7}\) molecule. Write the VSEPR notation for this molecule. Sketch the structure of the molecule, labeling all the bond angles.

Hydrogen azide, \(\mathrm{HN}_{3}\), can exist in two forms. One form has the three nitrogen atoms connected in a line; and the nitrogen atoms form a triangle in the other. Construct Lewis structures for these isomers and describe their shapes. Other interesting derivatives are nitrosyl azide ( \(\mathrm{N}_{4} \mathrm{O}\) ) and trifluoromethyl azide \(\left(\mathrm{CF}_{3} \mathrm{N}_{3}\right) .\) Describe the shapes of these molecules based on a line of nitrogen atoms.

A pair of isoelectronic species for \(C\) and \(N\) exist with the formula \(\mathrm{X}_{2} \mathrm{O}_{4}\) in which there is an \(\mathrm{X}-\mathrm{X}\) bond. \(\mathrm{A}\) corresponding fluoride of boron also exists. Draw Lewis structures for these species and describe their shapes.

Acetone \(\left(\mathrm{CH}_{3}\right) \mathrm{C}=\mathrm{O},\) a ketone, will react with a strong base \((\mathrm{A})\) to produce the enolate anion, \(\mathrm{CH}_{3}(\mathrm{C}=\mathrm{O}) \mathrm{CH}_{2}^{-} .\) Draw the Lewis structure of the enolate anion, and describe the relative contributions of any resonance structures.

The species \(\mathrm{PBr}_{4}^{-}\) has been synthesized and has been described as a tetrahedral anion. Comment on this description.

One of the allotropes of sulfur is a ring of eight sulfur atoms. Draw the Lewis structure for the \(S_{8}\) ring. Is the ring likely to be planar? The \(S_{8}\) ring can be oxidized to produce \(S_{8}\) O. In \(S_{8} O,\) the oxygen atom is bonded to one of the S atoms and the Sg ring is still intact. Draw the Lewis structure for \(\mathrm{S}_{8} \mathrm{O}\).

Alternative strategies to the one used in this chapter have been proposed for applying the VSEPR theory to molecules or ions with a single central atom. In general, these strategies do not require writing Lewis structures. In one strategy, we write (1) the total number of electron pairs \(=[\) (number of valence electrons) \(\pm\) (electrons required for ionic charge) \(] / 2\) (2) the number of bonding electron pairs \(=\) (number of atoms) -1 (3) the number of electron pairs around central atom \(=\) total number of electron pairs \(-3 \times[\) number of terminal atoms (excluding \(\mathrm{H}\) )] (4) the number of lone-pair electrons = number of central atom pairs - number of bonding pairs After evaluating items \(2,3,\) and \(4,\) establish the VSEPR notation and determine the molecular shape. Use this method to predict the geometrical shapes of the following: (a) \(\mathrm{PCl}_{5} ;\) (b) \(\mathrm{NH}_{3} ;\) (c) \(\mathrm{ClF}_{3} ;\) (d) \(\mathrm{SO}_{2} ;\) (e) \(\mathrm{ClF}_{4}^{-}\); (f) \(\mathrm{PCl}_{4}^{+}\). Justify each of the steps in the strategy, and explain why it yields the same results as the VSEPR method based on Lewis structures. How does the strategy deal with multiple bonds?

In your own words, define the following terms: (a) valence electrons; (b) electronegativity; (c) bonddissociation energy; (d) double covalent bond;(e) coordinate covalent bond.

Briefly describe each of the following ideas: (a) formal charge; (b) resonance; (c) expanded valence shell; (d) bond energy.

Explain the important distinctions between (a) ionic and covalent bonds; (b) lone-pair and bond-pair electrons; (c) molecular geometry and electron-group geometry; (d) bond dipole and resultant dipole moment; (e) polar molecule and nonpolar molecule.

Of the following species, the one with a triple covalent bond is (a) \(\mathrm{NO}_{3}^{-} ;\) (b) \(\mathrm{CN}^{-} ;\) (c) \(\mathrm{CO}_{2} ;\) (d) \(\mathrm{AlCl}_{3}\).

The formal charges on the \(\mathrm{O}\) atoms in the ion \([\mathrm{ONO}]^{+}\) is \((\mathrm{a})-2 ;(\mathrm{b})-1 ;(\mathrm{c}) 0 ;(\mathrm{d})+1\).

Which molecule is nonlinear? (a) \(\mathrm{SO}_{2} ;\) (b) \(\mathrm{CO}_{2}\); (c) HCN; (d) NO.

Which molecule is nonpolar? (a) \(\mathrm{SO}_{3} ;\) (b) \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\); (c) \(\mathrm{NH}_{3} ;\) (d) FNO.

The highest bond-dissociation energy is found in (a) \(\mathrm{O}_{2} ;\) (b) \(\mathrm{N}_{2} ;\) (c) \(\mathrm{Cl}_{2} ;\) (d) \(\mathrm{I}_{2}\).

The greatest bond length is found in (a) \(\mathrm{O}_{2} ;\) (b) \(\mathrm{N}_{2}\) (c) \(\mathrm{Br}_{2} ;\) (d) BrCl.

Draw plausible Lewis structures for the following species; use expanded valence shells where necessary. (a) \(\mathrm{Cl}_{2} \mathrm{O} ;\) (b) \(\mathrm{PF}_{3} ;\) (c) \(\mathrm{CO}_{3}^{2-} ;\) (d) \(\mathrm{BrF}_{5}\).

Predict the shapes of the following sulfur-containing species. (a) \(\mathrm{SO}_{2} ;\) (b) \(\mathrm{SO}_{3}^{2-} ;\) (c) \(\mathrm{SO}_{4}^{2-}\).

Without referring to tables or figures in the text other than the periodic table, indicate which of the following atoms, \(\mathrm{Bi}, \mathrm{S}, \mathrm{Ba}, \mathrm{As},\) or \(\mathrm{Mg},\) has the intermediate value when they are arranged in order of increasing electronegativity.

What is the VSEPR theory? On what physical basis is the VSEPR theory founded?

Use the \(\mathrm{NH}_{3}\) molecule as an example to explain the difference between molecular geometry and electron-group geometry.

If you have four electron pairs around a central atom, under what circumstances can you have a pyramidal molecule? Similarly, how can you have a bent molecule? What are the expected bond angles in each case?

Draw three resonance structures for the sulfine molecule, \(\mathrm{H}_{2} \mathrm{CSO}\). Do not consider ring structures.

Construct a concept map illustrating the connections between Lewis dot structures, the shapes of molecules, and polarity.