Chapter 3

For each of the following groups, place the atoms and/or ions in order of decreasing size. a. \(\mathrm{V}, \mathrm{V}^{2+}, \mathrm{V}^{3+}, \mathrm{V}^{5+}\) b. \(\mathrm{Na}^{+}, \mathrm{K}^{+}, \mathrm{Rb}^{+}, \mathrm{Cs}^{+}\) c. \(\mathrm{Te}^{2-}, \mathrm{I}^{-}, \mathrm{Cs}^{+}, \mathrm{Ba}^{2+}\) d. \(P, P, P^{2}, P^{3}\) e. \(\mathrm{O}^{2-}, \mathrm{S}^{2-}, \mathrm{Se}^{2-}, \mathrm{Te}^{2-}\)

Which compound in each of the following pairs of ionic substances has the most negative lattice energy? Justify your answers. a. NaCl, KCl b. LiF, LiCl c. \(\mathrm{Mg}(\mathrm{OH})_{2}, \mathrm{MgO}\) d. \(\mathrm{Fe}(\mathrm{OH})_{2}, \mathrm{Fe}(\mathrm{OH})_{3}\) e. \(\mathrm{NaCl}, \mathrm{Na}_{2} \mathrm{O}\) f. \(\mathrm{MgO}, \mathrm{BaS}\)

Which compound in each of the following pairs of ionic substances has the most negative lattice energy? Justify your answers. a. LiF, CsF b. NaBr, NaI c. \(\mathrm{BaCl}_{2}, \mathrm{BaO}\) d. \(\mathrm{Na}_{2} \mathrm{SO}_{4}, \mathrm{CaSO}_{4}\) e. \(\mathrm{KF}, \mathrm{K}_{2} \mathrm{O}\) f. \(\mathrm{Li}_{2} \mathrm{O}, \mathrm{Na}_{2} \mathrm{S}\)

Use the following data for potassium chloride to estimate \(\Delta E\) for the reaction: $$\mathrm{K}(s)+\frac{1}{2} \mathrm{Cl}_{2}(g) \longrightarrow \mathrm{KCl}(s) \quad \Delta E=?$$ Lattice energy Ionization energy for \(\mathbf{K}\) Electron affinity of Cl Bond energy of \(\mathrm{Cl}_{2}\) Energy of sublimation for \(\mathrm{K}\) \(-2913 \mathrm{kJ} / \mathrm{mol}\) \(735 \mathrm{kJ} / \mathrm{mol}\) \(1445 \mathrm{kJ} / \mathrm{mol}\) -328 kJ/mol \(154 \mathrm{kJ} / \mathrm{mol}\) 150\. kJ/mol

Use the following data formagnesium fluoride to estimate \(\Delta E\) for the reaction: $$\mathrm{Mg}(s)+\mathrm{F}_{2}(g) \longrightarrow \mathrm{MgF}_{2}(s) \quad \Delta E=?$$ Lattice energy First ionization energy of \(\mathrm{Mg}\) Second ionization energy of \(\mathbf{M g}\) Electron affinity of \(\mathbf{F}\) Bond energy of \(\mathrm{F}_{2}\) Energy of sublimation for \(\mathrm{Mg}\) \(-2913 \mathrm{kJ} / \mathrm{mol}\) \(735 \mathrm{kJ} / \mathrm{mol}\) \(1445 \mathrm{kJ} / \mathrm{mol}\) \(-328 \mathrm{kJ} / \mathrm{mol}\) \(154 \mathrm{kJ} / \mathrm{mol}\) 150\. kJ/mol

Consider the following energy changes: Magnesium oxide exists as \(\mathrm{Mg}^{2+} \mathrm{O}^{2-}\) and not as \(\mathrm{Mg}^{+} \mathrm{O}^{-}\) Explain.

Compare the electron affinity of fluorine to the ionization energy of sodium. Does the process of an electron being "pulled" from the sodium atom to the fluorine atom have a negative or a positive \(\Delta E ?\) Why is NaF a stable compound? Does the overall formation of NaF have a negative or a positive \(\Delta E ?\) How can this be?

Consider the following: $$\mathrm{Li}(s)+\frac{1}{2} \mathrm{I}_{2}(s) \longrightarrow \mathrm{LiI}(s) \quad \Delta E=-272 \mathrm{kJ} / \mathrm{mol}$$ LiI(s) has a lattice energy of -753 kJ/mol. The ionization energy of \(\mathrm{Li}(g)\) is \(520 . \mathrm{kJ} / \mathrm{mol},\) the bond energy of \(\mathrm{I}_{2}(g)\) is \(151 \mathrm{kJ} /\) mol, and the electron affinity of \(\mathrm{I}(g)\) is \(-295 \mathrm{kJ} / \mathrm{mol} .\) Use these data to determine the energy of sublimation of Li(s).

Rationalize the following lattice energy values: $$\begin{array}{|lc|} \hline & \text { Lattice Energy } \\ \text { Compound } & \text { (kj/mol) } \\ \hline \text { CaSe } & -2862 \\ \text { Na }_{2} \text { Se } & -2130 \\ \text { CaTe } & -2721 \\ \text { Na }_{2} \text { Te } & -2095 \\ \hline \end{array}$$

The lattice energies of \(\mathrm{FeCl}_{3}, \mathrm{FeCl}_{2},\) and \(\mathrm{Fe}_{2} \mathrm{O}_{3}\) are (in no particular order) \(-2631,-5359,\) and -14,774 kJ/mol. Match the appropriate formula to each lattice energy. Explain.

The major industrial source of hydrogen gas is by the following reaction: $$ \mathrm{CH}_{4}(g)+\mathrm{H}_{2} \mathrm{O}(g) \longrightarrow \mathrm{CO}(g)+3 \mathrm{H}_{2}(g) $$ Use bond energies to predict r this reaction.

Consider the following reaction: $$ \mathrm{A}_{2}+\mathrm{B}_{2} \longrightarrow 2 \mathrm{AB} \quad \Delta E=-285 \mathrm{kJ} $$ The bond energy for \(A_{2}\) is one-half the amount of the AB bond energy. The bond energy of \(\mathbf{B}_{2}=432 \mathrm{kJ} / \mathrm{mol} .\) What is the bond energy of \(\mathrm{A}_{2} ?\)

Write Lewis structures that obey the octet rule (duet rule for H) for each of the following molecules. Carbon is the central atom in \(\mathrm{CH}_{4}\), nitrogen is the central atom in \(\mathrm{NH}_{3}\), and oxygen is the central atom in \(\mathrm{H}_{2} \mathrm{O}\). a. \(\mathrm{F}_{2}\) \(\mathbf{b} . \mathbf{O}_{2}\) c. CO d. \(\overline{\mathrm{CH}_{4}}\) \(\mathbf{e} . \mathrm{NH}_{3}\) \(\mathbf{f .} \quad \mathbf{H}_{2} \mathbf{O}\) g. IIF

Write Lewis structures that obey the octet rule (duet rule for H) for each of the following molecules. a. \(\mathrm{H}_{2} \mathrm{CO}\) b. \(\mathrm{CO}_{2}\) c. HCN Carbon is the central atom in all of these molecules.

Write Lewis structures that obey the octet rule for each of the following molecules. a. \(\mathrm{CCl}_{4}\) b. \(\mathrm{NCl}_{3}\) c. \(\operatorname{Sec} 1_{2}\) d. ICl In each case, the atom listed first is the central atom.

Write Lewis structures that obey the octet rule for each of the following molecules and ions. (In each case the first atom listed is the central atom.) a. \(\mathrm{POCl}_{3}, \mathrm{SO}_{4}^{2-}, \mathrm{XeO}_{4}, \mathrm{PO}_{4}^{3-}, \mathrm{ClO}_{4}^{-}\) b. \(\mathrm{NF}_{3}, \mathrm{SO}_{3}^{2-}, \mathrm{PO}_{3}^{3-}, \mathrm{ClO}_{3}^{-}\) c. \(\mathrm{ClO}_{2}^{-}, \mathrm{SCl}_{2}, \mathrm{PCl}_{2}^{-}\) d. Considering your answers to parts a, b, and c, what conclusions can you draw concerning the structures of species containing the same number of atoms and the same number of valence electrons?

One type of exception to the octet rule are compounds with central atoms having fewer than eight electrons around them. \(\mathrm{BeH}_{2}\) and \(\mathrm{BH}_{3}\) are examples of this type of exception. Draw the Lewis structures for \(\mathrm{BeH}_{2}\) and \(\mathrm{BH}_{3}\)

Lewis structures can be used to understand why some molecules react in certain ways. Write the Lewis structures for the reactants and products in the reactions described below. a. Nitrogen dioxide dimerizes to produce dinitrogen tetroxide. b. Boron trihydride accepts a pair of electrons from ammonia, forming \(\mathrm{BH}_{3} \mathrm{NH}_{3}\) Give a possible explanation for why these two reactions occur.

The most common exceptions to the octet rule are compounds or ions with central atoms having more than eight electrons around them. \(\mathrm{PF}_{5}, \mathrm{SF}_{4}, \mathrm{ClF}_{3},\) and \(\mathrm{Br}_{3}^{-}\) are examples of this type of exception. Draw the Lewis structure for these compounds or ions. Which elements, when they have to, can have more than eight electrons around them? How is this rationalized?

\(\mathrm{SF}_{6}, \mathrm{ClF}_{5,}\) and \(\mathrm{XeF}_{4}\) are three compounds whose central atoms do not follow the octet rule. Draw Lewis structures for these compounds.

Write Lewis structures for the following. Show all resonancestructures where applicable.a. \(\mathrm{NO}_{2}^{-}, \mathrm{NO}_{3}^{-}, \mathrm{N}_{2} \mathrm{O}_{4}\left(\mathrm{N}_{2} \mathrm{O}_{4} \text { exists as } \mathrm{O}_{2} \mathrm{N}-\mathrm{NO}_{2} .\right)\) b. \(\mathrm{OCN}^{-}, \mathrm{SCN}^{-}, \mathrm{N}_{3}^{-}\) (Carbon is the central atom in \(\mathrm{OCN}^{-}\) and \(\mathrm{SCN}^{-} .\) )

Some of the important pollutants in the atmosphere are ozone \(\left(\mathrm{O}_{3}\right),\) sulfur dioxide, and sulfur trioxide. Write Lewis structures for these three molecules. Show all resonance structures where applicable.

Benzene \(\left(\mathrm{C}_{6} \mathrm{H}_{6}\right)\) consists of a six- membered ring of carbon atoms with one hydrogen bonded to each carbon. Write Lewis structures for benzene, including resonance structures.

Borazine \(\left(B_{3} N_{3} H_{6}\right)\) has often been called "inorganic" benzene. Write Lewis structures for borazine. Borazine contains a sixmembered ring of alternating boron and nitrogen atoms with one hydrogen bonded to each boron and nitrogen.

An important observation supporting the concept of resonance in the localized electron model was that there are only three different structures of dichlorobenzene \(\left(\mathrm{C}_{6} \mathrm{H}_{4} \mathrm{Cl}_{2}\right) .\) How does this fact support the concept of resonance (see Exercise 89)?

Consider the following bond lengths: \(\begin{array}{cccccc}\mathbf{C}-\mathbf{O} & 143 \mathrm{pm} & \mathbf{C}=\mathbf{O} & 123 \mathrm{pm} & \mathbf{C} \equiv \mathbf{O} & 109 \mathrm{pm}\end{array}\) In the \(\mathrm{CO}_{3}^{2-}\) ion, all three \(\mathrm{C}-\mathrm{O}\) bonds have identical bond lengths of \(136 \mathrm{pm}\). Why?

A toxic cloud covered Bhopal, India, in December 1984 when water leaked into a tank of methyl isocyanate, and the product escaped into the atmosphere. Methyl isocyanate is used in the production of many pesticides. Draw the Lewis structures for methyl isocyanate, \(\mathrm{CH}_{3} \mathrm{NCO}\), including resonance forms. The skeletal structure is

Order the following species with respect to carbon-oxygen bond length (longest to shortest). $$\mathrm{CO}, \quad \mathrm{CO}_{2}, \quad \mathrm{CO}_{3}^{2-}, \quad \mathrm{CH}_{3} \mathrm{OH}$$ What is the order from the weakest to the strongest carbonoxygen bond? \(\left(\mathrm{CH}_{3} \mathrm{OH} \text { exists as } \mathrm{H}_{3} \mathrm{C}-\mathrm{OH} .\right)\)

Place the species below in order of the shortest to the longest nitrogen- oxygen bond. $$\mathrm{H}_{2} \mathrm{NOH}, \quad \mathrm{N}_{2} \mathrm{O}, \quad \mathrm{NO}^{+}, \quad \mathrm{NO}_{2}^{-}, \quad \mathrm{NO}_{3}^{-}$$ $$\left(\mathrm{H}_{2} \mathrm{NOH} \text { exists as } \mathrm{H}_{2} \mathrm{N}-\mathrm{OH.}\right)$$

Use the formal charge arguments to rationalize why \(\mathrm{BF}_{3}\) would not follow the octet rule.

Use formal charge arguments to explain why CO has a less polar bond than expected on the basis of electronegativity.

Write Lewis structures that obey the octet rule for the following species. Assign the formal charge for each central atom. a. \(\mathrm{POCl}_{3}\) b. \(\mathrm{SO}_{4}^{2-}\) c. \(\mathrm{ClO}_{4}\) d. \(\mathrm{PO}_{4}^{3-}\) e. \(\mathrm{SO}_{2} \mathrm{Cl}_{2}\) f. \(\quad X \in O_{4}\) g. \(\mathrm{ClO}_{3}\) h. \(\mathrm{NO}_{4}^{3-}\)

Oxidation of the cyanide ion produces the stable cyanate ion, OCN \(^{-}\). The fulminate ion, \(\mathrm{CNO}^{-}\), on the other hand, is very unstable. Fulminate salts explode when struck; \(\mathrm{Hg}(\mathrm{CNO})_{2}\) is used in blasting caps. Write the Lewis structures and assign formal charges for the cyanate and fulminate ions. Why is the fulminate ion so unstable? (C is the central atom in OCN - and \(\mathrm{N}\) is the central atom in \(\mathrm{CNO}^{-} .\) )

Nitrous oxide \(\left(\mathrm{N}_{2} \mathrm{O}\right)\) has three possible Lewis structures: $$\therefore N=N=O^{\cdot} \leftrightarrow: N \equiv N-\vec{O}: \longleftrightarrow: N-N \equiv 0$$ Given the following bond lengths, $$\begin{aligned} &\mathrm{N}-\mathrm{N} \quad 167 \mathrm{pm} \quad \mathrm{N}=\mathrm{O} \quad 115 \mathrm{pm}\\\ &\mathrm{N}=\mathrm{N} \quad 120 \mathrm{pm} \quad \mathrm{N}-\mathrm{O} \quad 147 \mathrm{pm}\\\ &\mathrm{N} \equiv \mathrm{N} \quad 110 \mathrm{pm} \end{aligned}$$ rationalize the observations that the \(\mathrm{N}-\mathrm{N}\) bond length in \(\mathrm{N}_{2} \mathrm{O}\) is \(112 \mathrm{pm}\) and that the \(\mathrm{N}-\mathrm{O}\) bond length is \(119 \mathrm{pm}\). Assign formal charges to the resonance structures for \(\mathrm{N}_{2} \mathrm{O}\). Can you eliminate any of the resonance structures on the basis of formal charges? Is this consistent with observation?

Name the compounds in parts a-d and write the formulas for the compounds in parts e-h. a. \(\mathrm{NaBr}\) b. \(\mathrm{Rb}_{2} \mathrm{O}\) c. CaS d. \(\mathrm{AlI}_{3}\) e. strontium fluoride f. aluminum selenide g. potassium nitride h. magnesium phosphide

Name the compounds in parts a-d and write the formulas for the compounds in parts e-h. a. \(\mathrm{Hg}_{2} \mathrm{O}\) b. \(\operatorname{FeBr}_{3}\) c. CoS d. \(\mathrm{TiCl}_{4}\) e. \(\operatorname{tin}(\text { II })\) nitride f. cobalt(III) iodide g. mercury(II) oxide h. chromium(VI) sulfide

Name each of the following compounds: a. CsF b. \(\mathrm{Li}_{3} \mathrm{N}\) c. \(A g_{2} S\) d. \(\mathrm{MnO}_{2}\) e. \(\mathrm{TiO}_{2}\) f. \(\mathrm{Sr}_{3} \mathrm{P}_{2}\)

Write the formula for each of the following compounds: a. zinc chloride b. \(\operatorname{tin}(1 \mathrm{V})\) fluoride c. calcium nitride d. aluminum sulfide e. mercury(I) selenide f. silver iodide

Name each of the following compounds: a. \(\mathrm{BaSO}_{3}\) b. \(\mathrm{NaNO}_{2}\) c. \(\mathrm{KMnO}_{4}\) \(\mathbf{d .} \mathbf{K}_{2} \mathrm{Cr}_{2} \mathbf{O}_{7}\)

Write the formula for each of the following compounds: a. chromium(III) hydroxide b. magnesium cyanide c. lead(IV) carbonate d. ammonium acetate

Write the formula for each of the following compounds: a. diboron trioxide b. arsenic pentafluoride c. dinitrogen monoxide d. sulfur hexachloride

Name each of the following compounds: a. Cul b. \(\mathrm{CuI}_{2}\) c. \(\mathrm{CoI}_{2}\) d. \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) e. \(\mathrm{NaHCO}_{3}\) \(\mathbf{f .} \mathrm{S}_{4} \mathrm{N}_{4}\) g. \(\mathrm{SF}_{4}\) h. NaOCl i. \(\mathrm{BaCrO}_{4}\) J. \(\mathrm{NH}_{4} \mathrm{NO}_{3}\)

Name each of the following compounds. Assume the acids are dissolved in water. a. \(\mathrm{HC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\) b. \(\mathrm{NH}_{4} \mathrm{NO}_{2}\) c. \(\mathrm{Co}_{2} \mathrm{S}_{3}\) d. ICl e. \(\mathrm{Pb}_{3}\left(\mathrm{PO}_{4}\right)_{2}\) f. \(\mathrm{KClO}_{3}\) g. \(\mathrm{H}_{2} \mathrm{SO}_{4}\) \(\mathbf{h} . \mathrm{Sr}_{3} \mathrm{N}_{2}\) i. \(\quad \mathrm{Al}_{2}\left(\mathrm{SO}_{3}\right)_{3}\) J. \(\mathrm{SnO}_{2}\) \(\mathbf{k} . \mathrm{Na}_{2} \mathrm{CrO}_{4}\) I. \(\mathrm{HClO}\)

Elements in the same family often form oxyanions of the same general formula. The anions are named in a similar fashion. Give the names of the oxyanions of selenium and tellurium: \(\mathrm{ScO}_{4}^{2-}, \mathrm{SeO}_{3}^{2-}, \mathrm{TeO}_{4}^{2-}, \mathrm{TeO}_{3}^{2-}\)

Knowing the names of similar chlorine oxyanions and acids, deduce the names of the following: \(\mathrm{IO}^{-}, \mathrm{IO}_{2}^{-}, \mathrm{IO}_{3}^{-}, \mathrm{IO}_{4}^{-}\) \(\mathrm{HIO}, \mathrm{HIO}_{2}, \mathrm{HIO}_{3}, \mathrm{HIO}_{4}\)

Write the formula for each of the following compounds: a. sulfur difluoride b. sulfur hexafluoride c. sodium dihydrogen phosphate d. lithium nitride e. chromium(III) carbonate f. \(\operatorname{tin}(\text { II })\) fluoride g. ammonium acetate h. ammonium hydrogen sulfate i. cobalt(III) nitrate J. mercury(I) chloride k. potassium chlorate 1\. sodium hydride

Write the formula for each of the following compounds: a. chromium(VI) oxide b. disulfur dichloride c. nickel(II) fluoride d. potassium hydrogen phosphate e. aluminum nitride f. ammonia g. manganese(IV) sulfide h. sodium dichromate i. ammonium sulfite J. carbon tetraiodide

Write the formula for each of the following compounds: a. sodium oxide b. sodium peroxide c. potassium cyanide d. copper(II) nitrate e. selenium tetrabromide f. iodous acid g. lead(IV) sulfide h. copper(I) chloride I. gallium arsenide J. cadmium selenide k. zinc sulfide I. nitrous acid m. diphosphorus pentoxide

Write the formula for each of the following compounds: a. ammonium hydrogen phosphate b. mercury(I) sulfide c. silicon dioxide d. sodium sulfite e. aluminum hydrogen sulfate f. nitrogen trichloride g. hydrobromic acid h. bromous acid I. perbromic acid J. potassium hydrogen sulfide k. calcium iodide 1\. cesium perchlorate

Each of the following compounds is incorrectly named. What is wrong with each name, and what is the correct name for each compound? a. \(\mathrm{FeCl}_{3},\) iron chloride b. \(\mathrm{NO}_{2},\) nitrogen(IV) oxide c. \(\mathrm{CaO},\) calcium(II) monoxide d. \(\mathrm{Al}_{2} \mathrm{S}_{3},\) dialuminum trisulfide e. \(\mathrm{Mg}\left(\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\right)_{2},\) manganese diacetate f. \(\mathrm{FePO}_{4},\) iron(II) phosphide g. \(\mathrm{P}_{2} \mathrm{S}_{5},\) phosphorus sulfide h. \(\mathrm{Na}_{2} \mathrm{O}_{2},\) sodium oxide