Chapter 19

Write Lewis structures for all the resonance forms of sulfuric acid.

Write Lewis structures for all the resonance forms of nitric acid.

Iodine trichloride, \(\mathrm{ICl}_{3},\) is an interhalogen compound. (a) Write the Lewis structure of \(\mathrm{ICl}_{3}\) (b) Does the central atom have more than an octet of valence electrons? (c) Using VSEPR theory, predict the molecular shape of \(\mathrm{ICl}_{3}\).

At some temperature, a gaseous mixture in a 1.00 - \(\mathrm{L}\) vessel originally contained \(1.00 \mathrm{~mol} \mathrm{SO}_{2}\) and \(5.00 \mathrm{~mol} \mathrm{O}_{2}\). When equilibrium was reached, \(77.8 \%\) of the \(\mathrm{SO}_{2}\) had been converted to \(\mathrm{SO}_{3}\). Calculate the equilibrium constant \(\left(K_{\mathrm{c}}\right)\) for this reaction at this temperature.

White phosphorus is soluble in carbon disulfide. At 10. \({ }^{\circ} \mathrm{C}, 800 . \mathrm{g} \mathrm{P}_{4}\) dissolves in \(100 . \mathrm{g} \mathrm{CS}_{2}\). The density of carbon disulfide at \(10 .{ }^{\circ} \mathrm{C}\) is \(1.26 \mathrm{~g} / \mathrm{mL}\) (a) Calculate the molarity of phosphorus in the solution. (b) Calculate the molality of phosphorus in the solution.

Drug stores sell \(3 \%\) aqueous hydrogen peroxide that is used as an antiseptic. Hydrogen peroxide, \(\mathrm{H}_{2} \mathrm{O}_{2},\) decomposes to water and oxygen. Calculate the volume of oxygen produced if \(250 . \mathrm{mL}\) of \(3 \%\) hydrogen peroxide decomposes fully at \(750 . \mathrm{mmHg}\) and \(22^{\circ} \mathrm{C}\).

Calculate the volume of concentrated \((98 \%)\) sulfuric acid that is needed to produce two tons of phosphoric acid from the reaction of sulfuric acid with sufficient phosphate-bearing rock. Density of conc. sulfuric acid = \(1.84 \mathrm{~g} / \mathrm{mL}\) \(\mathrm{Ca}_{3}\left(\mathrm{PO}_{4}\right)_{2}(\mathrm{~s})+3 \mathrm{H}_{2} \mathrm{SO}_{4}(\ell) \longrightarrow 2 \mathrm{H}_{3} \mathrm{PO}_{4}(\ell)+3 \mathrm{CaSO}_{4}(\mathrm{~s})\)

The \(K_{\mathrm{sp}}\) of \(\mathrm{Ca}(\mathrm{OH})_{2}\) is \(5.5 \times 10^{-6} ;\) that of \(\mathrm{Mg}(\mathrm{OH})_{2}\) is \(1.8 \times 10^{-11}\). Calculate the equilibrium constant for the reaction $$\mathrm{Ca}(\mathrm{OH})_{2}(\mathrm{~s})+\mathrm{Mg}^{2+}(\mathrm{aq}) \longrightarrow \mathrm{Ca}^{2+}(\mathrm{aq})+\mathrm{Mg}(\mathrm{OH})_{2}(\mathrm{~s})$$ Use it to explain why this reaction can be used commercially to extract magnesium from seawater.

Commercial concentrated nitric acid contains 69.5 mass percent \(\mathrm{HNO}_{3}\) and has a density of \(1.42 \mathrm{~g} / \mathrm{mL}\). (a) Calculate the molarity of this solution. (b) Calculate what volume of the concentrated acid must be used to prepare \(10.0 \mathrm{~L}\) of \(6.00-\mathrm{M} \mathrm{HNO}_{3}\).

The compound nitrosyl azide, \(\mathrm{N}_{4} \mathrm{O},\) is a covalent compound with an NNNNO atomic arrangement. Write a plausible Lewis structure for this compound.

Hydrazoic acid, \(\mathrm{HN}_{3}\), is very explosive in its pure state but can be studied in aqueous solution. The acid is prepared by the reaction of hydrazine with nitrous acid. \(\mathrm{N}_{2} \mathrm{H}_{4}(\ell)+\mathrm{HNO}_{2}(\mathrm{aq}) \longrightarrow \mathrm{HN}_{3}(\mathrm{aq})+2 \mathrm{H}_{2} \mathrm{O}(\ell)\) (a) Determine the oxidation states of nitrogen in the compounds in this reaction. (b) What is the oxidizing agent in this reaction? (c) The \(K_{\mathrm{a}}\) of hydrazoic acid is \(1.0 \times 10^{-5}\) at \(25^{\circ} \mathrm{C}\). Calculate the \(\mathrm{pH}\) of a \(0.010-\mathrm{M}\) solution of \(\mathrm{HN}_{3}\).

Given the reaction $$\mathrm{Cl}_{2}(\mathrm{~g})+\mathrm{H}_{2} \mathrm{O}(\ell) \rightleftharpoons \mathrm{H}^{+}(\mathrm{aq})+\mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{HOCl}(\mathrm{aq})$$ (a) identify the oxidizing agent and the reducing agent. (b) write the equilibrium constant expression for the reaction. (c) calculate the concentration of \(\mathrm{HOCl}\) in equilibrium with $$\mathrm{Cl}_{2}(\mathrm{~g}) \text { at } 1.0 \mathrm{~atm} . K_{\mathrm{c}}=2.7 \times 10^{-5}$$.

Dinitrogen trioxide, \(\mathrm{N}_{2} \mathrm{O}_{3}\), is a blue liquid formed by the reaction of \(\mathrm{NO}_{2}\) and \(\mathrm{NO}\). (a) Write a balanced chemical equation for the formation of \(\mathrm{N}_{2} \mathrm{O}_{3}\) (b) Write the Lewis structure of \(\mathrm{N}_{2} \mathrm{O}_{3}\) and any plausible resonance forms. (c) Predict the \(\mathrm{O}-\mathrm{N}-\mathrm{O}\) and the \(\mathrm{N}-\mathrm{N}-\mathrm{O}\) bond angles.

(a) Write the resonance forms of \(\mathrm{SO}_{3}\). (b) Predict the molecular shape of \(\mathrm{SO}_{3}\) and the \(\mathrm{O}-\mathrm{S}-\mathrm{O}\) bond angle.

Iodine can be produced by the oxidation of iodide ion with permanganate ion. $$\begin{aligned} \mathrm{MnO}_{4}^{-}(\mathrm{aq})+2 \mathrm{I}^{-}(\mathrm{aq})+8 \mathrm{H}^{+}(\mathrm{aq}) & \longrightarrow \\ \mathrm{I}_{2}(\mathrm{~s})+\mathrm{Mn}^{2+}(\mathrm{aq})+4 \mathrm{H}_{2} \mathrm{O}(\ell)\end{aligned}$$ Excess HI is added to \(0.200 \mathrm{~g} \mathrm{MnO}_{4}^{-}\). Assuming \(100 \%\) yield, calculate the mass (g) of iodine produced.

In a Downs cell, molten \(\mathrm{NaCl}\) is electrolyzed to sodium metal and chlorine gas. $$2 \mathrm{NaCl}(\ell) \longrightarrow 2 \mathrm{Na}(\ell)+\mathrm{Cl}_{2}(\mathrm{~g})$$ \(\Delta_{\mathrm{r}} H^{\circ}\) and \(\Delta_{\mathrm{r}} S^{\circ}\) for the reaction are \(+820 \mathrm{~kJ} / \mathrm{mol}\) and \(+180 \mathrm{~J} \mathrm{~K}^{-1} \mathrm{~mol}^{-1},\) respectively. (a) Calculate \(\Delta_{\mathrm{r}} G^{\circ}\) at \(600 .{ }^{\circ} \mathrm{C},\) the electrolysis temperature. (b) Calculate the voltage required for the electrolysis.

A 425-gal tank of water contains \(175 \mathrm{~g}\) NaI. Calculate the volume (L) of chlorine gas at \(758 \mathrm{mmHg}\) and \(25^{\circ} \mathrm{C}\) required to convert all the iodide to iodine.

In the laboratory, small quantities of bromine can be produced by the reaction of hydrobromic acid with \(\mathrm{MnO}_{2}\). The unbalanced equation for the reaction is $$\mathrm{MnO}_{2}(\mathrm{~s})+\mathrm{Br}^{-}(\mathrm{aq}) \longrightarrow \mathrm{Mn}^{2+}(\mathrm{aq})+\mathrm{Br}_{2}(\ell)$$ (a) Balance the equation. (b) If the reaction has \(100 \%\) yield, calculate the amount of bromide ions that reacts to produce \(6.50 \mathrm{~g}\) bromine. (c) Calculate the mass (g) of \(\mathrm{MnO}_{2}\) required in part (b).

Lapis lazuli is an aluminum silicate whose brilliant blue color is due to the presence of \(\mathrm{S}_{3}^{-}\) ions. Write a plausible Lewis structure for this ion.

Magnesium can be extracted from dolomite, a mineral that contains \(13.2 \% \mathrm{Mg}, 21.7 \% \mathrm{Ca}, 13.0 \% \mathrm{C},\) and \(52.1 \% \mathrm{O} .\) Determine the simplest formula for this ionic compound.

A natural brine found in Arkansas has a bromide ion concentration of \(5.00 \times 10^{-3} \mathrm{M}\). If \(210 . \mathrm{g} \mathrm{Cl}_{2}\) were added to \(1.00 \times 10^{3} \mathrm{~L}\) of the brine, (a) determine the limiting reactant. (b) calculate the theoretical yield of \(\mathrm{Br}_{2}(d=3.12 \mathrm{~g} / \mathrm{mL})\).

Chlorine gas was first prepared by Carl Wilhelm Scheele in 1774 by the reaction of sodium chloride, manganese(IV) oxide, and sulfuric acid. In addition to chlorine, the reaction produces water, sodium sulfate, and manganese(II) sulfate. Write the balanced equation for this reaction.

Mercury(II) azide, \(\mathrm{Hg}\left(\mathrm{N}_{3}\right)_{2},\) is an unstable compound used as a detonator in blasting caps. Calculate the volume (L) of nitrogen produced at \(1 \mathrm{~atm}\) and \(25^{\circ} \mathrm{C}\) when \(2.50 \mathrm{~g}\) mercury azide decomposes to liquid mercury and nitrogen.

At \(20 .{ }^{\circ} \mathrm{C}\) the vapor pressure of white phosphorus is \(0.0254 \mathrm{mmHg}\); at \(40 .{ }^{\circ} \mathrm{C}\) it is \(0.133 \mathrm{mmHg}\). Use the Clausius-Clapeyron equation to estimate the heat of sublimation (J/mol) of white phosphorus.

The density of sulfur vapor at \(700 .{ }^{\circ} \mathrm{C}\) and \(1.00 \mathrm{~atm}\) is \(0.8012 \mathrm{~g} / \mathrm{L}\). Determine the molecular formula of sulfur in the vapor.

Assume that the radius of Earth is \(6400 \mathrm{~km},\) the crust is \(50 . \mathrm{km}\) thick, the density of the crust is \(3.5 \mathrm{~g} / \mathrm{cm}^{3},\) and \(25.7 \%\) of the crust is silicon by mass. Calculate the total mass of silicon in the crust of Earth.

A \(5.00-\mathrm{g}\) sample of white phosphorus is burned in excess oxygen and the product is dissolved in sufficient water to form \(250 . \mathrm{mL}\) of solution. (a) Write the balanced chemical equation for the burning of phosphorus in excess oxygen. (b) Calculate the \(\mathrm{pH}\) of the resulting solution. (c) An excess of aqueous calcium nitrate is added to the solution causing a white precipitate to form. Write a balanced chemical equation for this reaction and calculate the mass of precipitate formed. (d) An excess of zinc is added to the remaining solution. The reaction generates a colorless gas. Identify the gas and calculate its volume at STP.

Bromine, \(\mathrm{Br}_{2}\), reacts vigorously with hydrogen, \(\mathrm{H}_{2}\), to form hydrogen bromide. At STP, \(100 .\) mL hydrogen gas reacts with a stoichiometric amount of bromine. The resulting hydrogen bromide is dissolved in sufficient water to form \(250 . \mathrm{mL}\) solution. Calculate the \(\mathrm{pH}\) of the solution.

At 1 atm and approximately \(1800^{\circ} \mathrm{C}, 50 \%\) of \(\mathrm{P}_{4}\) is dissociated into \(2 \mathrm{P}_{2}\). If equilibrium is established under these conditions, calculate the equilibrium constant.

The Gibbs free energy of formation, \(\Delta_{f} G^{\circ}\), of HI is \(+1.70 \mathrm{~kJ} / \mathrm{mol}\) at \(25^{\circ} \mathrm{C}\). Calculate the equilibrium constant for the reaction \(\mathrm{HI}(\mathrm{g}) \Longrightarrow \frac{1}{2} \mathrm{H}_{2}(\mathrm{~g})+\frac{1}{2} \mathrm{~J}_{2}(\mathrm{~g})\).

Predict whether the formation of ions of the alkaline earth elements requires less energy or more energy than formation of ions of the alkali metals. Explain your answer.

Use a Born-Haber cycle ( Sec. \(5-13\) ) to calculate the lattice energy of \(\mathrm{MgF}_{2}\) using these thermodynamic data. $$\begin{aligned} \Delta_{\text {sub }} H \mathrm{Mg}(\mathrm{s}) &=+146 \mathrm{~kJ} / \mathrm{mol} \\ \text { B.E. } \mathrm{F}_{2}(\mathrm{~g}) &=+158 \mathrm{~kJ} / \mathrm{mol} \\\ \text { I.E. } \mathrm{Mg}(\mathrm{g}) &=+738 \mathrm{~kJ} / \mathrm{mol} \\ \text { I.E. }_{2} \mathrm{Mg}^{+}(\mathrm{g}) &=+1450 \mathrm{~kJ} / \mathrm{mol} ; \\ \text { E.A. } \mathrm{F}(\mathrm{g}) &=-328 \mathrm{~kJ} / \mathrm{mol} \\ \Delta_{\mathrm{f}} H^{\circ} \mathrm{MgF}_{2}(\mathrm{~s}) &=-1124 \mathrm{~kJ} / \mathrm{mol} . \end{aligned}$$ Compare this lattice energy with that of \(\mathrm{SrF}_{2}\), \(-2496 \mathrm{~kJ} / \mathrm{mol}\). Explain the difference in the values in structural terms.

Elemental analysis of a borane indicates this composition: \(84.2 \% \mathrm{~B}\) and \(15.7 \% \mathrm{H}\). The compound has a molar mass of \(76.7 \mathrm{~g} / \mathrm{mol}\). Determine the molecular formula of the borane.

The reaction of iodine with excess liquid chlorine produces the interhalogen compound \(\mathrm{I}_{2} \mathrm{Cl}_{6}\). (a) Write the Lewis structure of this molecule. (b) Use VSEPR theory to predict the structure of the molecule.

Hydroxyapatite is the important compound in tooth enamel. It dissociates according to the equation \(\mathrm{Ca}_{5}\left(\mathrm{PO}_{4}\right)_{3} \mathrm{OH}(\mathrm{s})\) \(5 \mathrm{Ca}^{2+}(\mathrm{aq})+3 \mathrm{PO}_{4}^{3+}(\mathrm{aq})+\mathrm{OH}^{-}(\mathrm{aq})\) Children drink milk to obtain calcium, but the fermentation of the milk produces lactic acid, which remains on the teeth. (a) Use the dissociation equation to explain how drinking milk helps babies to produce "strong" teeth. (b) Explain why the lactic acid inhibits formation of strong teeth.