Chapter 6

Which of these molecules have Lewis structures that involve exceptions to the octet rule? Classify each exception. (a) \(\mathrm{SF}_{6}\) (b) \(\mathrm{NH}_{4} \mathrm{NO}_{3}\) (c) \(\mathrm{SeF}_{2}\) (d) \(\mathrm{ClO}_{2}\)

The \(\mathrm{NO}^{2-}\) radical dianion of nitrogen oxide has been formed, as well as \(\mathrm{N}_{2}^{3-},\) the radical trianion of diatomic nitrogen, electronically analogous to superoxide ion, \(\mathrm{O}_{2}^{-}\). Write Lewis structures for these two nitrogen radical anions.

All carbon-to-carbon bond lengths are identical in benzene. Does this argue for or against the presence of \(\mathrm{C}=\mathrm{C}\) bonds in benzene? Explain.

Carbon-to-carbon double bonds, \(\mathrm{C}=\mathrm{C}\), react by addition. Cite experimental evidence that benzene does not have \(\mathrm{C}=\mathrm{C}\) bonds.

Three dibromobenzenes are known. Write the Lewis structure and name for each compound.

Write the structural formula for 1,2 -diiodobenzene (also known as ortho- diiodobenzene). Write the structural formulas for the meta and para isomers as well.

Use MO theory to predict the MO diagram, the number of bonds, and the number of unpaired electrons in (a) peroxide ion, \(\mathrm{O}_{2}^{2-}\) (b) \(\mathrm{B}_{2}^{+}\) (c) \(\mathrm{Li}_{2}^{+}\) (d) \(\mathrm{O}_{2}^{+}\)

Use MO theory to predict the number of electrons in each of the molecular orbitals, the number of bonds, ar the number of unpaired electrons in (a) \(\mathrm{CO}\) (b) \(\mathrm{F}_{2}^{-}\) (c) \(\mathrm{NO}^{-}\)

Use molecular orbital theory to predict the arrangement of electrons in MOs, the bond order, and the number of unpaired electrons in (a) \(\mathrm{BN}\) (b) cyanide ion, \(\mathrm{CN}^{-}\)

Both polyatomic ions and uncharged molecules can be detected using spectroscopic measurements. Two examples of polyatomic ions are \(\mathrm{He}_{2}^{2+}\) and \(\mathrm{HHe}^{+}\). Predict the arrangement of electrons in MOs and the bond order for each ion.

Using just a periodic table (not a table of electronegativities), decide which of these is likely to be the most polar bond. Explain your answer. (a) \(\mathrm{C}-\mathrm{F}\) (b) \(S-F\) (c) \(\mathrm{Si}=\mathrm{F}\) (d) \(O-F\)

Your friend says, "Elements that are close together in the periodic table form covalent bonds, whereas elements that are far apart form ionic bonds." Is your friend correct? Why or why not?

In nitryl chloride, \(\mathrm{NO}_{2} \mathrm{Cl}\), there is no oxygen-oxygen bond. Write a Lewis structure for the molecule. Write any resonance structures for this molecule.

Write Lewis structures for (a) \(\mathrm{SCl}_{2}\) (b) \(\mathrm{Cl}_{3}^{+}\) (c) \(\mathrm{SOCl}_{2}\) (d) \(\mathrm{ClOClO}_{3}(\) contains \(\mathrm{Cl}-\mathrm{O}-\mathrm{Cl}\) bond \()\)

Write a Lewis structure for each molecule. Which formulas represent alkanes? Which are probably aromatic? Which fall into neither category? (a) \(\mathrm{C}_{8} \mathrm{H}_{10}\) (b) \(\mathrm{C}_{10} \mathrm{H}_{8}\) (c) \(\mathrm{C}_{6} \mathrm{H}_{12}\) (d) \(\mathrm{C}_{6} \mathrm{H}_{14}\) (e) \(\mathrm{C}_{8} \mathrm{H}_{18}\) (f) \(\mathrm{C}_{6} \mathrm{H}_{10}\)

A student drew this incorrect Lewis structure for \(\mathrm{ClO}_{3}^{-}\). What errors were made when determining the number of valence electrons?

This Lewis structure for \(\mathrm{SF}_{5}^{+}\) is drawn incorrectly. What error was made when determining the number of valence electrons?

Tribromide, \(\mathrm{Br}_{3}^{-},\) and triiodide \(, \mathrm{I}_{3}^{-},\) ions are often found in aqueous solutions, but trifluoride ion, \(\mathrm{F}_{3}^{-},\) is so rare that its bond strength was only measured in \(2000 .\) Explain.

Explain why nonmetal atoms in Period 3 and beyond can accommodate greater than an octet of electrons and those in Period 2 cannot do so.

In another universe, elements try to achieve a nonet (nine valence electrons) instead of an octet when forming chemical bonds. As a result, covalent bonds form when a trio of electrons is shared between two atoms. Draw Lewis structures for the compounds that would form between (a) hydrogen and oxygen, and (b) hydrogen and fluorine.

The elements As, Br, Cl, S, and Se have electronegativity values of \(2.6,2.3,2.1,2.7,\) and \(2.4,\) but not in that order. Using the periodic trend for electronegativity, assign the values to the elements. Which assignments are you certain about? Which are you not?

Which of these molecules is least likely to exist: \(\mathrm{NF}_{5}\), \(\mathrm{PF}_{5}, \mathrm{SbF}_{5},\) or \(\mathrm{IF}_{5}\) ? Explain why.

Write the Lewis structure for nitrosyl fluoride, FNO. Using only a periodic table, identify (a) which is the longer bond. (b) which is the stronger bond. (c) which is the more polar bond.

Write the Lewis structure for nitrosyl chloride, CINO. Using only a periodic table, identify (a) which is the longer bond. (b) which is the stronger bond. (c) which is the more polar bond.

For many people, tears flow when they cut raw onions. The tear-inducing agent is thiopropionaldehyde-S-oxide, \(\mathrm{C}_{3} \mathrm{H}_{6} \mathrm{OS}\). Write the Lewis structure for this molecule.

Phosphorus and sulfur form a series of compounds, one of which is tetraphosphorus trisulfide. Write the Lewis structure for this molecule.

Tetrasulfur tetranitride reacts with disulfur dichloride to form \(\mathrm{S}_{4} \mathrm{~N}_{3} \mathrm{Cl}\), a salt. $$ 3 \mathrm{~S}_{4} \mathrm{~N}_{4}+2 \mathrm{~S}_{2} \mathrm{Cl}_{2} \longrightarrow 4 \mathrm{~S}_{4} \mathrm{~N}_{3} \mathrm{Cl} $$ Write a plausible Lewis structure for the two reactants and the cation of the salt.

Elemental phosphorus has the formula \(\mathrm{P}_{4}\). Propose a Lewis structure for this molecule. [Hints: (1) Each phosphorus atom is bonded to three other phosphorus atoms. (2) Visualize the structure three-dimensionally, not flat on a page.]

When we estimate \(\Delta_{\mathrm{r}} H^{\circ}\) from bond enthalpies we assume that all bonds of the same type (single, double, triple) between the same two atoms have the same energy, regardless of the molecule in which they occur. The purpose of this problem is to show you that this is only an approximation. You will need these standard enthalpies of formation: $$ \begin{array}{ll} \mathrm{C}(\mathrm{g}) & \Delta_{\mathrm{f}} H^{\circ}=716.7 \mathrm{~kJ} / \mathrm{mol} \\ \mathrm{CH}(\mathrm{g}) & \Delta_{\mathrm{f}} H^{\circ}=596.3 \mathrm{~kJ} / \mathrm{mol} \\ \mathrm{CH}_{2}(\mathrm{~g}) & \Delta_{\mathrm{f}} H^{\circ}=392.5 \mathrm{~kJ} / \mathrm{mol} \\ \mathrm{CH}_{3}(\mathrm{~g}) & \Delta_{\mathrm{f}} H^{\circ}=146.0 \mathrm{~kJ} / \mathrm{mol} \\ \mathrm{H}(\mathrm{g}) & \Delta_{\mathrm{f}} H^{\circ}=218.0 \mathrm{~kJ} / \mathrm{mol} \end{array} $$ (a) What is the average \(\mathrm{C}-\mathrm{H}\) bond energy in methane, \(\mathrm{CH}_{4} ?\) (b) Using bond enthalpies, estimate \(\Delta_{1} H^{\circ}\) for the reaction $$ \mathrm{CH}_{4}(\mathrm{~g}) \longrightarrow \mathrm{C}(\mathrm{g})+2 \mathrm{H}_{2}(\mathrm{~g}) $$ (c) By heating \(\mathrm{CH}_{4}\) in a flame it is possible to produce the reactive gaseous species \(\mathrm{CH}_{3}, \mathrm{CH}_{2}, \mathrm{CH},\) and even carbon atoms, C. Experiments give these values of \(\Delta_{r} H^{\circ}\) for the reactions shown: $$ \begin{aligned} \mathrm{CH}_{3}(\mathrm{~g}) \longrightarrow \mathrm{C}(\mathrm{g})+\mathrm{H}_{2}(\mathrm{~g})+\mathrm{H}(\mathrm{g}) & & \Delta_{\mathrm{r}} H^{\circ} &=788.7 \mathrm{~kJ} \\ \mathrm{CH}_{2}(\mathrm{~g}) \longrightarrow \mathrm{C}(\mathrm{g})+\mathrm{H}_{2}(\mathrm{~g}) & & \Delta_{\mathrm{r}} H^{\circ} &=324.2 \mathrm{~kJ} \\ \mathrm{CH}(\mathrm{g}) \longrightarrow \mathrm{C}(\mathrm{g})+\mathrm{H}(\mathrm{g}) & & \Delta_{\mathrm{r}} H^{\circ} &=338.3 \mathrm{~kJ} \end{aligned} $$ For each of the reactions in part (c), draw a diagram similar to Figure 6.6 . Then, calculate the average \(\mathrm{C}-\mathrm{H}\) bond energy in \(\mathrm{CH}_{3}, \mathrm{CH}_{2}\), and \(\mathrm{CH}\). Comment on any trends you see.

Nitrosyl azide, \(\mathrm{N}_{4} \mathrm{O},\) is a pale yellow solid first synthesized in 1993 . Write the Lewis structure for nitrosyl azide.

Write the Lewis structures for (a) \(\left(\mathrm{Cl}_{2} \mathrm{PN}\right)_{3}\) (b) \(\left(\mathrm{Cl}_{2} \mathrm{PN}\right)_{4}\)

Nitrous oxide, \(\mathrm{N}_{2} \mathrm{O}\), is a linear molecule that has the two nitrogen atoms adjacent to each other. Using concepts from this chapter, explain why the structure is not NON with the nitrogen atoms each bonded to the oxygen.

The azide ion, \(\mathrm{N}_{3}^{-}\), has three resonance hybrid structures. (a) Write the Lewis structure of each. (b) Use formal charges to determine which is the most favorable resonance structure.

Hydrazoic acid, \(\mathrm{HN}_{3}\), has three resonance hybrid structures. (a) Write the Lewis structure of each. (b) Use formal charges to determine which is the least favorable resonance structure.

Write the Lewis structures for (a) \(\left(\mathrm{SO}_{3}\right)_{3}\) (b) \(\mathrm{FXeN}\left(\mathrm{SO}_{2} \mathrm{~F}\right)_{2}\)

Experimental evidence indicates the existence of \(\mathrm{HC}_{3} \mathrm{~N}\) molecules in interstellar clouds. Write a plausible Lewis structure for this molecule.

Molecules of \(\mathrm{SiC}_{3}\) have been discovered in interstellar clouds. Write a plausible Lewis structure for this molecule.

One of the structural isomers of \(\mathrm{C}_{3} \mathrm{H}_{6} \mathrm{OS}\) is the compound that makes you cry when you slice onions. Write Lewis structures for two isomers of this molecule.

Sulfur and oxygen form a series of \(2-\) anions including sulfite, \(\mathrm{SO}_{3}^{2-},\) and sulfate, \(\mathrm{SO}_{4}^{2-} .\) In addition to these, there are three other more complex anions - dithionite, \(\mathrm{S}_{2} \mathrm{O}_{4}^{2-},\) dithionate, \(\mathrm{S}_{2} \mathrm{O}_{6}^{2-},\) and tetrathionate, \(\mathrm{S}_{4} \mathrm{O}_{6}^{2-} .\) Write correct Lewis structures for dithionite, dithionate, and tetrathionate ions.

Gaseous molecules in the ground vibrational state have a vibrational energy given by the equation \(E=\frac{1}{2} h v\) where \(h\) is Planck's constant and \(v\) is the frequency of the vibration. Consider a water vapor molecule, \(\mathrm{HOH}\), in which one of the lighter isotope hydrogen atoms, \(\mathrm{H},\) is replaced by the heavier isotope deuterium atom, \(\mathrm{D},\) to form \(\mathrm{HOD}\). (a) Will an HOD molecule have a higher or lower ground vibrational state energy than an HOH molecule? Explain your answer. (b) Which bond is stronger: \(\mathrm{O}-\mathrm{H}\) or \(\mathrm{O}-\mathrm{D}\) ? Explain your answer.

Suppose in building up molecular orbitals, the \(\pi_{2 p}\) were placed above the \(\sigma_{2 p} .\) Prepare a diagram similar to Figure 6.11 based on these changes. For which species in Table 6.4 would this change in relative energies of the MOs affect the prediction of number of bonds and number of unpaired electrons?

In carbon suboxide, \(\mathrm{C}_{3} \mathrm{O}_{2}\), a linear molecule, the atoms are in the sequence \(\mathrm{OCCCO}\). The carbon-to-carbon bond distance is \(128 \mathrm{pm} ;\) the carbon-to-oxygen bond distance is \(119 \mathrm{pm}\). Describe the bonding in this molecule.

Compare and contrast the valence-bond and the molecular-orbital descriptions of bonding in the gaseous, diatomic \(\mathrm{S}_{2}\) molecule. The sulfur-to- sulfur bond distance is \(189 \mathrm{pm}\). A chemistry classmate says, "The valence bond description of this molecule is not correct." Explain whether the classmate's comment is accurate.

Write Lewis structures for these three anions: \(\mathrm{C}_{2}^{2-}\) (acetylide); \(\mathrm{N}_{3}^{-}\) (azide); and \(\mathrm{O}_{3}^{-}\) (ozonide). Write resonance structures if there are any. Use formal charges to determine which is the more likely resonance structure.

The bond length can be considered as approximately the sum of the atomic radii of the two bonded atoms. The atomic radii for single-bonded carbon and oxygen are \(77 \mathrm{pm}\) and \(74 \mathrm{pm}\), respectively. For double-bonded \(\mathrm{C}\) and \(\mathrm{O}\) the values are \(67 \mathrm{pm}\) and \(60 \mathrm{pm}\), respectively. Use these data to estimate the carbon-to-oxygen bond length in (a) methanol, \(\mathrm{CH}_{3} \mathrm{OH} ;\) (b) dimethyl ether, \(\mathrm{CH}_{3} \mathrm{OCH}_{3}\); (c) formaldehyde, \(\mathrm{H}_{2} \mathrm{CO}\). Explain any difference among the bond lengths (see Table 6.1 as a reference).

Consider the \(\mathrm{SC}\left[\mathrm{N}\left(\mathrm{CH}_{3}\right)_{2}\right]_{2}\) molecule in which \(\mathrm{S}\) is bonded to \(\mathrm{C}\) and each \(\mathrm{CH}_{3}\) group is bonded to \(\mathrm{N}\). Use Lewis structures and formal charges to write resonance structures and to determine which is the most plausible resonance structure.

Consider the reaction: \(\mathrm{P}_{4}(\mathrm{~g}) \longrightarrow 2 \mathrm{P}_{2}(\mathrm{~g})\) for which \(\Delta_{\mathrm{r}} H=229 \mathrm{~kJ} / \mathrm{mol}\). The bond energy of a \(\mathrm{P}-\mathrm{P}\) single bond is \(209 \mathrm{~kJ} / \mathrm{mol}\). (a) Calculate the bond energy of a phosphorus-tophosphorus triple bond. (b) Compare this calculated value with the bond energy of \(\mathrm{N}_{2}\) and propose an explanation for the difference in bond energies between \(\mathrm{P}_{2}\) and \(\mathrm{N}_{2}\).

A substance is analyzed and found to contain \(85.7 \%\) carbon and \(14.3 \%\) hydrogen by weight. A gaseous sample of the substance is found to have a density of \(1.87 \mathrm{~g} / \mathrm{L},\) and \(1 \mathrm{~mol}\) of it occupies a volume of \(22.4 \mathrm{~L}\). What are two possible Lewis structures for molecules of the compound? (Hint: First determine the empirical formula and molar mass of the substance.)