Chapter 8

Chemists can use the octet rule to predict the structures of new compounds to synthesize. Draw Lewis structures showing all resonance forms for the hypothetical compound ClSeNSO, where the atoms are connected in the order they are written.

Describe how formal charges are used to choose between possible molecular structures.

How do the electronegativities of elements influence the selection of which Lewis structure is favored?

In a molecule containing \(\mathrm{S}\) and \(\mathrm{O}\) atoms, is a structure with a negative formal charge on sulfur more likely to contribute to bonding than an alternative structure with a negative formal charge on oxygen? Explain.

In a cation containing \(\mathrm{N}\) and \(\mathrm{O},\) why do Lewis structures with a positive formal charge on nitrogen contribute more to the actual bonding in the molecule than do those structures with a positive formal charge on oxygen?

Hydrogen isocyanide (HNC) has the same elemental composition as hydrogen cyanide (HCN), but the H atom in HNC is bonded to the nitrogen atom. Draw a Lewis structure for HNC and assign formal charges to each atom. How do the formal charges on the atoms differ in the Lewis structures for HCN and HNC?

Molecules in Interstellar Space Hydrogen cyanide (HCN) and cyanoacetylene (HC \(_{3}\) N) have been detected in the interstellar regions of space and in comets close to Earth (Figure \(\mathrm{P} 8.86\) ). Draw Lewis structures for these molecules, and assign formal charges to each atom. The hydrogen atom is bonded to the carbon atom in both cases.

Origins of Life The discovery of polyatomic organic molecules such as cyanamide (H \(_{2} \mathrm{NCN}\) ) in interstellar space has led some scientists to believe that the molecules from which life began on Earth may have come from space. Draw Lewis structures for cyanamide, and select the preferred structure on the basis of formal charges.

Draw all resonance forms of the sulfur-nitrogen anion, \(\mathrm{S}_{4} \mathrm{N}^{-},\) and assign formal charges. The atoms are arranged as SSNSS.

Nitrogen is the central atom in molecules of nitrous oxide \(\left(\mathrm{N}_{2} \mathrm{O}\right) .\) Draw Lewis structures for another possible arrangement: \(\mathrm{N}-\mathrm{O}-\mathrm{N} .\) Assign formal charges and suggest a reason why this structure is not likely to be stable.

Use formal charges to determine which resonance form of each of the following ions is preferred: CNO"; NCO"; \(\mathrm{CON}^{-}\).

Are all odd-electron molecules exceptions to the octet rule?

Describe the factors that contribute to the stability of structures in which the central atoms have more than eight valence electrons.

Why do \(\mathrm{C}, \mathrm{N}, \mathrm{O},\) and \(\mathrm{F}\) atoms in covalently bonded molecules and ions have no more than eight valence electrons?

Do atoms with \(Z>12\) always expand their valence shell? Explain your answer.

In which of the following molecules does the sulfur atom have an expanded valence shell? (a) \(\mathrm{SF}_{6} ;\) (b) \(\mathrm{SF}_{5} ;\) (c) \(\mathrm{SF}_{4}\) (d) \(\mathrm{SF}_{2}\).

In which of the following molecules does the phosphorus atom have an expanded valence shell? (a) \(\mathrm{POCl}_{3} ;\) (b) \(\mathrm{PF}_{5}\) (c) \(\left.\mathrm{PF}_{3} ; \text { (d) } \mathrm{P}_{2} \mathrm{F}_{4} \text { (which has a } \mathrm{P}-\mathrm{P} \text { bond }\right)\).

How many electrons are there in the covalent bonds surrounding the central atom in the following species? (a) \(\left(\mathrm{CH}_{3}\right)_{3} \mathrm{Al} ;\) (b) \(\mathrm{B}_{2} \mathrm{Cl}_{4} ;\) (c) \(\mathrm{SO}_{3} ;\) (d) \(\mathrm{SF}_{5}^{-}\).

How many electrons are there in the covalent bonds surrounding the central atom in the following species? (a) \(\mathrm{POCl}_{3} ;(\mathrm{b}) \mathrm{InCl}_{5}^{2-} ;(\mathrm{c}) \mathrm{FBO} ;(\mathrm{d}) \mathrm{PF}_{4}^{-}\).

Draw Lewis structures for \(\mathrm{NOF}_{3}\) and \(\mathrm{POF}_{3}\) in which the group 15 element is the central atom and the other atoms are bonded to it. What differences are there in the types of bonding in these molecules?

The phosphate anion is common in minerals. The corresponding nitrogen- containing anion, \(\mathrm{NO}_{4}^{3-},\) is unstable but can be prepared by reacting sodium nitrate with sodium oxide at \(300^{\circ} \mathrm{C}\). Draw Lewis structures for each anion. What are the differences in bonding between these ions?

Dissolving NaF in selenium tetrafluoride (SeF \(_{4}\) ) produces NaSeF \(_{5 .}\) Draw Lewis structures for \(\operatorname{SeF}_{4}\) and \(\operatorname{SeF}_{5}\). In which structure does Se have more than eight valence electrons?

Reaction between \(\mathrm{NF}_{3}, \mathrm{F}_{2},\) and \(\mathrm{SbF}_{3}\) at \(200^{\circ} \mathrm{C}\) and 100 atm pressure gives the ionic compound \(\mathrm{NF}_{4} \mathrm{SbF}_{6}\) $$ \mathrm{NF}_{3}(g)+2 \mathrm{F}_{2}(g)+\mathrm{SbF}_{3}(g) \rightarrow \mathrm{NF}_{4} \mathrm{SbF}_{6}(s) $$ Draw Lewis structures for the ions in this product.

Trimethylaluminum reacts with dimethylamine, \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{NH},\) forming methane and \(\mathrm{Al}_{2}\left(\mathrm{CH}_{3}\right)_{4}\left[\mathrm{N}\left(\mathrm{CH}_{3}\right)_{2}\right]_{2}\) by the balanced chemical equation: \(2\left(\mathrm{CH}_{3}\right)_{3} \mathrm{Al}+2 \mathrm{HN}\left(\mathrm{CH}_{3}\right)_{2} \rightarrow\) $$ 2 \mathrm{CH}_{4}+\mathrm{Al}_{2}\left(\mathrm{CH}_{3}\right)_{4}\left[\mathrm{N}\left(\mathrm{CH}_{3}\right)_{2}\right]_{2} $$ Draw the Lewis structures for the reactants and products. Must any of these structures contain atoms with fewer than eight valence electrons?

Which of the following chlorine oxides are odd-electron molecules? (a) \(\mathrm{Cl}_{2} \mathrm{O}_{7} ;\) (b) \(\mathrm{Cl}_{2} \mathrm{O}_{6} ;\) (c) \(\mathrm{ClO}_{4} ;\) (d) \(\mathrm{ClO}_{3}\) (e) \(\mathrm{ClO}_{2}\).

Which of the following nitrogen oxides are odd-electron molecules? (a) \(\mathrm{NO} ;\) (b) \(\mathrm{NO}_{2} ;\) (c) \(\mathrm{NO}_{3} ;\) (d) \(\mathrm{N}_{2} \mathrm{O}_{4} ;\) (e) \(\mathrm{N}_{2} \mathrm{O}_{5}\).

Using Lewis structures, explain why dimethylaluminum chloride is more likely to exist as \(\left(\mathrm{CH}_{3}\right)_{4} \mathrm{Al}_{2} \mathrm{Cl}_{2}\) than as \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{AlCl}\).

When left at room temperature, mixtures of \(\mathrm{BF}_{3}\) and \(\mathrm{BCl}_{3}\) are found to contain significant amounts of \(\mathrm{BF}_{2} \mathrm{Cl}\) and BFCl \(_{2} .\) Using Lewis structures, explain the origin of the latter two compounds..

Some have argued that \(\mathrm{SF}_{6}\) has ionic resonance forms that do not require an expanded octet for S. Draw a resonance structure consistent with this hypothesis and assign formal charges to each atom. Is this resonance form better than or the same as the one with an expanded octet?

The synthesis of an extremely unusual molecule, \([\mathrm{NO}]_{2}+\left[\mathrm{XeF}_{8}\right]^{2-},\) was reported close to 50 years ago. The structure of \(\left[\mathrm{XeF}_{8}\right]^{2-}\) shows that all eight fluorine atoms are within bonding distance of the Xe. Draw a Lewis structure for the \(\left[\mathrm{XeF}_{8}\right]^{2-}\) anion.

Do you expect the nitrogen-oxygen bond length in the nitrate ion to be the same as in the nitrite ion? Explain.

Why is the oxygen-oxygen bond length in \(\mathrm{O}_{3}\) not the same as in \(\mathrm{O}_{2} ?\)

Explain why the nitrogen-oxygen bond lengths in \(\mathrm{N}_{2} \mathrm{O}_{4}\) (which has a nitrogen-nitrogen bond) and \(\mathrm{N}_{2} \mathrm{O}\) are nearly identical ( 118 and 119 pm, respectively).

Do you expect the sulfur-oxygen bond lengths in \(\mathrm{SO}_{3}^{2-}\) and \(\mathrm{SO}_{4}^{2-}\) ions to be about the same? Why?

Rank the following ions in order of (a) increasing nitrogenoxygen bond lengths and (b) increasing bond energies: \(\mathrm{NO}_{2}^{-} ; \mathrm{NO}^{+} ; \mathrm{NO}_{3}^{-}\).

Rank the following compounds and ions in order of (a) increasing carbon-oxygen bond lengths and (b) increasing bond energies: \(\mathrm{CO} ; \mathrm{CO}_{2} ; \mathrm{CO}_{3}^{2-}\).

Do you expect the boron-fluorine bond energy to be the same in \(\mathrm{BF}_{3}\) and \(\mathrm{F}_{3} \mathrm{BNH}_{3} ?\)

The boron-oxygen distances in the \(\mathrm{BO}_{2}^{+}\) cation are equal. Does this mean the bond order of the \(\mathrm{B}\) -O bond is two? Explain.

Why must the stoichiometry of a reaction be known in order to estimate the enthalpy change from bond energies?

Why must the structures of the reactants and products be known in order to estimate the enthalpy change of a reaction from bond energies?

When calculating the enthalpy change for a chemical reaction by using bond energies, why is it important to know the phase (solid, liquid, or gaseous) for every compound in the reaction?

If the energy needed to break 2 moles of \(\mathrm{C}=\mathrm{O}\) bonds is greater than the sum of the energies needed to break the \(\mathrm{O}=\mathrm{O}\) bonds in 1 mole of \(\mathrm{O}_{2}\) and vaporize 1 mole of carbon, why does the combustion of pure carbon release heat?

Use average bond energies to estimate the enthalpy changes of the following reactions: a. \(\mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \rightarrow 2 \mathrm{NH}_{3}(g)\) b. \(\mathrm{N}_{2}(g)+2 \mathrm{H}_{2}(g) \rightarrow \mathrm{H}_{2} \mathrm{NNH}_{2}(g)\) c. \(2 \mathrm{N}_{2}(g)+\mathrm{O}_{2}(g) \rightarrow 2 \mathrm{N}_{2} \mathrm{O}(g)\)

Use average bond energies to estimate the enthalpy changes of the following reactions: a. \(\mathrm{CO}_{2}(g)+\mathrm{H}_{2}(g) \rightarrow \mathrm{H}_{2} \mathrm{O}(g)+\mathrm{CO}(g)\) b. \(\mathrm{N}_{2}(g)+\mathrm{O}_{2}(g) \rightarrow 2 \mathrm{NO}(g)\) c.\(\mathrm{C}(s)+\mathrm{CO}_{2}(g) \rightarrow 2 \mathrm{CO}(g)\)

Carbon and oxygen form three oxides: \(\mathrm{CO}, \mathrm{CO}_{2},\) and carbon suboxide \(\left(\mathrm{C}_{3} \mathrm{O}_{2}\right) .\) Draw a Lewis structure for \(\mathrm{C}_{3} \mathrm{O}_{2}\) in which the three carbon atoms are bonded to each other, and predict whether the carbon-oxygen bond lengths in the carbon suboxide molecule are equal.

Spectroscopic analysis of the linear molecule \(\mathrm{N}_{4} \mathrm{O}\) reveals that the nitrogen-oxygen bond length is \(135 \mathrm{pm}\) and that there are three nitrogen-nitrogen bond lengths: 148,127 and \(115 \mathrm{pm} .\) Draw the Lewis structure for \(\mathrm{N}_{4} \mathrm{O}\) consistent with these observations.

Use formal charges to predict whether the atoms in carbon disulfide are arranged CSS or SCS.

The following is a family of weak acids: \(\mathrm{HClO}, \mathrm{HClO}_{2}\) and HClO \(_{3} .\) Draw their Lewis structures, using formal charges to predict the best arrangement of atoms. Show any resonance forms of the molecules.

Chemical Weapons Phosgene is a poisonous gas first used in chemical warfare during World War I. It has the formula \(\mathrm{COCl}_{2}(\mathrm{C}\) is the central atom). a. Draw its Lewis structure. b. Phosgene kills because it reacts with water in nasal passages, in the lungs, and on the skin to produce carbon dioxide and hydrogen chloride. Write a balanced chemical equation for this process, showing the Lewis structures for reactants and products.

Silver cyanate (AgOCN) is a source of the cyanate ion \(\left(\mathrm{OCN}^{-}\right),\) which reacts with several small molecules. Under certain conditions the species OCN is an anion with a charge of \(1-;\) under other conditions it is a neutral, odd-electron molecule. a. Two molecules of OCN combine to form OCNNCO. Draw the Lewis structures for this molecule, including all resonance forms. b. The OCN" ion reacts with BrNO, forming the unstable molecule OCNNO. Draw the Lewis structures for BrNO and OCNNO, including all resonance forms. c. The OCN- ion reacts with \(\mathrm{Br}_{2}\) and \(\mathrm{NO}_{2}\) to produce \(\mathrm{N}_{2} \mathrm{O}, \mathrm{CO}_{2}, \mathrm{BrNCO},\) and \(\mathrm{OCN}(\mathrm{CO}) \mathrm{NCO} .\) Draw three of the resonance forms of \(\mathrm{OCN}(\mathrm{CO}) \mathrm{NCO},\) which has the arrangement of atoms shown in Figure P8.145.