Chapter 22

Give the formula of the stable fluoride formed by \(\mathrm{Li}\) \(\mathrm{Be}, \mathrm{B}, \mathrm{C}, \mathrm{N},\) and \(\mathrm{O} .\) For these fluorides, describe the variation in the bonding that occurs as we move from left to right across the period.

Fluorine is able to stabilize elements in very high oxidation states. For each of the elements \(\mathrm{Na}, \mathrm{Mg}, \mathrm{Al}, \mathrm{Si}\) P, S, and Cl, give the formula of the highest-order fluoride that is known to exist. Then, describe the variation in bonding that occurs as we move from left to right across the period.

The oxides of the phosphorus(III), antimony(III), and bismuth(III) are \(\mathrm{P}_{4} \mathrm{O}_{6}, \mathrm{Sb}_{4} \mathrm{O}_{6},\) and \(\mathrm{Bi}_{2} \mathrm{O}_{3} .\) Only one of these oxides is amphoteric. Which one? Which of these oxides is most acidic? Which is most basic?

The oxides of the selenium(IV) and tellurium(IV) are \(\mathrm{SeO}_{2}\) and \(\mathrm{TeO}_{2} .\) One of these oxides is amphoteric and one is acidic. Which is which?

A 55 L cylinder contains \(A r\) at 145 atm and \(26^{\circ}\) C. What minimum volume of air at STP must have been liquefied and distilled to produce this Ar? Air contains \(0.934 \%\) Ar, by volume.

Use VSEPR theory to predict the probable geometric structures of (a) \(\mathrm{XeO}_{3} ;\) (b) \(\mathrm{XeO}_{4} ;\) (c) \(\mathrm{XeF}_{5}^{+}\).

Use VSEPR theory to predict the probable geometric structures of the molecules (a) \(\mathrm{O}_{2} \mathrm{XeF}_{2} ;\) (b) \(\mathrm{O}_{3} \mathrm{XeF}_{2}\) (c) OXeF \(_{4}\).

Write a chemical equation for the hydrolysis of \(\mathrm{XeF}_{4}\) that yields \(\mathrm{XeO}_{3}, \mathrm{Xe}, \mathrm{O}_{2},\) and \(\mathrm{HF}\) as products.

Write a chemical equation for the hydrolysis in alkaline solution of \(\mathrm{XeF}_{6}\) that yields \(\mathrm{XeO}_{6}^{4-}, \mathrm{Xe}, \mathrm{O}_{2}, \mathrm{F}^{-}\) and \(\mathrm{H}_{2} \mathrm{O}\) as products.

Provide an explanation for the observation that helium, neon, and argon do not react directly with fluorine.

Provide an explanation for the inability of \(\mathrm{O}_{2}\) to react directly with xenon.

Freshly prepared solutions containing iodide ion are colorless, but over time they usually turn yellow. Describe a plausible chemical reaction (or reactions) to account for this observation.

Fluorine can be prepared by the reaction of hexafluoromanganate(IV) ion, MnF \(_{6}^{2-}\), with antimony pentafluoride to produce manganese(IV) fluoride and \(\mathrm{SbF}_{6}^{-}\) followed by the disproportionation of manganese(IV) fluoride to manganese(III) fluoride and \(\mathrm{F}_{2}(\mathrm{g}) .\) Write chemical equations for these two reactions.

Make a general prediction about which of the halogen elements, \(\mathrm{F}_{2}, \mathrm{Cl}_{2}, \mathrm{Br}_{2},\) or \(\mathrm{I}_{2},\) displaces other halogens from a solution of halide ions. Which of the halogens is able to displace \(\mathrm{O}_{2}(\mathrm{g})\) from water? Which is able to displace \(\mathrm{H}_{2}(\mathrm{g})\) from water?

The abundance of \(\mathrm{F}^{-}\) in seawater is \(1 \mathrm{g} \mathrm{F}^{-}\) per ton of seawater. Suppose that a commercially feasible method could be found to extract fluorine from seawater. (a) What mass of \(\mathrm{F}_{2}\) could be obtained from \(1 \mathrm{km}^{3}\) of seawater \(\left(d=1.03 \mathrm{g} \mathrm{cm}^{-3}\right) ?\) (b) Would the process resemble that for extracting bromine from seawater? Explain.

Show by calculation whether the disproportionation of chlorine gas to chlorate and chloride ions will occur under standard-state conditions in an acidic solution.

Show by calculation whether the reaction \(2 \mathrm{HOCl}(\mathrm{aq}) \longrightarrow \mathrm{HClO}_{2}(\mathrm{aq})+\mathrm{H}^{+}(\mathrm{aq})+\mathrm{Cl}^{-}(\mathrm{aq})\) will go essentially to completion as written for standardstate conditions.

Predict the geometric structures of (a) \(\mathrm{BrF}_{3} ;\) (b) IF \(_{5}\); (c) \(\mathrm{Cl}_{3} \mathrm{IF}^{-}\). (Central atom underlined.).

Which of the following species has a linear structure: \(\mathrm{ClF}_{2}^{+}, \mathrm{IBrF}^{-}, \mathrm{OCl}_{2}, \mathrm{ClF}_{3},\) or \(\mathrm{SF}_{4} ?\) (Central atom underlined.) Do any two of these species have the same structure?

When iodine is added to an aqueous solution of iodide ion, the \(I_{3}^{-}\) ion is formed, according to the reaction below: $$\mathrm{I}_{2}(\mathrm{aq})+\mathrm{I}^{-}(\mathrm{aq}) \rightleftharpoons \mathrm{I}_{3}^{-}(\mathrm{aq})$$ The equilibrium constant for the reaction above is \(K=7.7 \times 10^{2}\) at \(25^{\circ} \mathrm{C}\) (a) What is \(E^{\circ}\) for the reaction above? (b) If a 0.0010 mol sample of \(I_{2}\) is added to 1.0 L of \(0.0050 \mathrm{M} \mathrm{NaI}(\mathrm{aq})\) at \(25^{\circ} \mathrm{C},\) then what fraction of the \(\mathrm{I}_{2}\) remains unreacted at equilibrium?

The trichloride ion, \(\mathrm{Cl}_{3}^{-1}\), is not very stable in aqueous solution. The equilibrium constant for the following dissociation reaction is 5.5 at \(25^{\circ} \mathrm{C}\) : $$\mathrm{Cl}_{3}^{-}(\mathrm{aq}) \rightleftharpoons \mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{Cl}_{2}(\mathrm{aq})$$ (a) Draw a Lewis structure for the \(\mathrm{Cl}_{3}^{-}\) ion and predict the geometry. (b) Calculate the equilibrium concentration of \(\mathrm{Cl}_{3}^{-}\) if 0.0010 moles each of \(\mathrm{KCl}\) and \(\mathrm{Cl}_{2}\) are dissolved in water at \(25^{\circ} \mathrm{C}\) to make \(1.0 \mathrm{L}\) of solution.

Each of the following compounds decomposes to produce \(\mathrm{O}_{2}(\mathrm{g})\) when heated: \((\mathrm{a}) \mathrm{HgO}(\mathrm{s}) ;\) (b) \(\mathrm{KClO}_{4}(\mathrm{s})\) Write plausible equations for these reactions.

\(\mathrm{O}_{3}(\mathrm{g})\) is a powerful oxidizing agent. Write equations to represent oxidation of \((a) I^{-}\) to \(I_{2}\) in acidic solution; (b) sulfur in the presence of moisture to sulfuric acid; (c) \(\left[\mathrm{Fe}(\mathrm{CN})_{6}\right]^{4-}\) to \(\left[\mathrm{Fe}(\mathrm{CN})_{6}\right]^{3-}\) in basic solution. In each case \(\mathrm{O}_{3}(\mathrm{g})\) is reduced to \(\mathrm{O}_{2}(\mathrm{g})\).

Without per frming detailed calculations, determine which of the following compounds has the greatest percent oxygen by mass: dinitrogen tetroxide, aluminum oxide, tetraphosphorus hexoxide, or carbon dioxide.

A typical concentration of \(\mathrm{O}_{3}\) in the ozone layer is \(5 \times 10^{12} \mathrm{O}_{3}\) molecules \(\mathrm{cm}^{-3} .\) What is the partial pressure of \(\mathrm{O}_{3},\) expressed in millimeters of mercury, in that layer? Assume a temperature of \(220 \mathrm{K}\).

Explain why the volumes of \(\mathrm{H}_{2}(\mathrm{g})\) and \(\mathrm{O}_{2}(\mathrm{g})\) obtained in the electrolysis of water are not the same.

In the electrolysis of a sample of water \(22.83 \mathrm{mL}\) of \(\mathrm{O}_{2}(\mathrm{g})\) was collected at \(25.0^{\circ} \mathrm{C}\) at an oxygen partial pressure of \(736.7 \mathrm{mmHg} .\) Determine the mass of water that was decomposed.

In water, \(\mathrm{O}^{2-}\) is a strong base. If \(50.0 \mathrm{mg}\) of \(\mathrm{Li}_{2} \mathrm{O}\) is dissolved in \(750.0 \mathrm{mL}\) of aqueous solution, what will be the pH of the solution?

Use Lewis structures and other information to explain the observation that (a) \(\mathrm{H}_{2} \mathrm{S}\) is a gas at room temperature, whereas \(\mathrm{H}_{2} \mathrm{O}\) is a liquid. (b) \(\mathrm{O}_{3}\) is diamagnetic.

Use Lewis structures and other information to explain the observation that (a) the oxygen-to-oxygen bond lengths in \(\mathrm{O}_{2}, \mathrm{O}_{3}\) and \(\mathrm{H}_{2} \mathrm{O}_{2}\) are \(121,128,\) and \(148 \mathrm{pm},\) respectively. (b) the oxygen-to-oxygen bond length of \(\mathrm{O}_{2}\) is \(121 \mathrm{pm}\) and for \(\mathrm{O}_{2}^{+}\) is \(112 \mathrm{pm}\). Why is the bond length for \(\mathrm{O}_{2}^{+}\) so much shorter than for \(\mathrm{O}_{2} ?\)

Which of the following reactions are likely to go to completion or very nearly so? (a) \(\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq})+2 \mathrm{I}^{-}(\mathrm{aq})+2 \mathrm{H}^{+}(\mathrm{aq}) \longrightarrow\) \(\mathrm{I}_{2}(\mathrm{s})+2 \mathrm{H}_{2} \mathrm{O}(1)\) (b) \(\mathrm{O}_{2}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(1)+4 \mathrm{Cl}^{-}(\mathrm{aq}) \stackrel{-}{\longrightarrow}\) \(2 \mathrm{Cl}_{2}(\mathrm{g})+4 \mathrm{OH}^{-}(\mathrm{aq})\) (c) \(\mathrm{O}_{3}(\mathrm{g})+\mathrm{Pb}^{2+}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \longrightarrow\) \(\mathrm{PbO}_{2}(\mathrm{s})+2 \mathrm{H}^{+}(\mathrm{aq})+\mathrm{O}_{2}(\mathrm{g})\) (d) \(\mathrm{HO}_{2}^{-}(\mathrm{aq})+2 \mathrm{Br}^{-}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \longrightarrow\) \(3 \mathrm{OH}^{-}(\mathrm{aq})+\mathrm{Br}_{2}(1)\)

Each of the following compounds produces \(\mathrm{O}_{2}(\mathrm{g})\) when strongly heated: (a) \(\mathrm{HgO}(\mathrm{s}) ;\) (b) \(\mathrm{KClO}_{4}(\mathrm{s})\) (c) \(\mathrm{Hg}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{s}) ;\) (d) \(\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq}) .\) Write a plausible equation for the reaction that occurs in each instance.

Joseph Priestley, a British chemist, was credited with the discovering oxygen in \(1774 .\) In his experiments, he generated oxygen gas by heating \(\mathrm{HgO}(\mathrm{s}) .\) The other product of the decomposition reaction is \(\mathrm{Hg}(1) .\) What volume of wet \(\mathrm{O}_{2}(\mathrm{g})\) is obtained from the decomposition of \(1.0 \mathrm{g} \mathrm{HgO}(\mathrm{s}),\) if the gas is collected over water at \(25^{\circ} \mathrm{C}\) and a barometric pressure of \(756 \mathrm{mmHg} ?\) The vapor pressure of water is \(23.76 \mathrm{mmHg}\) at \(25^{\circ} \mathrm{C}\).

Give an appropriate name to each of the following compounds: (a) \(\mathrm{ZnS} ;\) (b) \(\mathrm{KHSO}_{3} ;\) (c) \(\mathrm{K}_{2} \mathrm{S}_{2} \mathrm{O}_{3} ;\) (d) \(\mathrm{SF}_{4}\).

Give an appropriate formula for each of the following compounds: (a) calcium sulfate dihydrate; (b) hydrosulfuric acid; (c) sodium hydrogen sulfate; (d) disulfuric acid.

Give a specific example of a chemical equation that illustrates the (a) reaction of a metal sulfide with \(\mathrm{HCl}(\mathrm{aq})\) (b) action of a nonoxidizing acid on a metal sulfite; (c) oxidation of \(\mathrm{SO}_{2}(\mathrm{aq})\) to \(\mathrm{SO}_{4}^{2-}(\mathrm{aq})\) by \(\mathrm{MnO}_{2}(\mathrm{s})\) in acidic solution; (d) disproportionation of \(S_{2} \mathrm{O}_{3}^{2-}\) in acidic solution.

Show how you would use elemental sulfur, chlorine gas, metallic sodium, water, and air to produce aqueous solutions containing (a) \(\mathrm{Na}_{2} \mathrm{SO}_{3} ;\) (b) \(\mathrm{Na}_{2} \mathrm{SO}_{4}\) (c) \(\mathrm{Na}_{2} \mathrm{S}_{2} \mathrm{O}_{3} .\) [Hint: You will have to use information from other chapters as well as this one.].

Describe a chemical test you could use to determine whether a white solid is \(\mathrm{Na}_{2} \mathrm{SO}_{4}\) or \(\mathrm{Na}_{2} \mathrm{S}_{2} \mathrm{O}_{3} .\) Explain the basis of this test using a chemical equation or equations.

Explain why sulfur can occur naturally as sulfates, but not as sulfites.

What mass of \(\mathrm{Na}_{2} \mathrm{SO}_{3}\) was present in a sample that required \(26.50 \mathrm{mL}\) of \(0.0510 \mathrm{M} \mathrm{KMnO}_{4}\) for its oxidation to \(\mathrm{Na}_{2} \mathrm{SO}_{4}\) in an acidic solution? \(\mathrm{MnO}_{4}^{-}\) is reduced to \(\mathrm{Mn}^{2+}\).

A \(1.100 \mathrm{g}\) sample of copper ore is dissolved, and the \(\mathrm{Cu}^{2+}(\mathrm{aq})\) is treated with excess KI. The liberated \(\mathrm{I}_{3}^{-}\) requires \(12.12 \mathrm{mL}\) of \(0.1000 \mathrm{M} \mathrm{Na}_{2} \mathrm{S}_{2} \mathrm{O}_{3}\) for its titration. What is the mass percent copper in the ore?

A 25.0 L sample of a natural gas, measured at \(25^{\circ} \mathrm{C}\). and 740.0 Torr, is bubbled through \(\mathrm{Pb}^{2+}(\mathrm{aq}),\) yielding \(0.535 \mathrm{g}\) of \(\mathrm{PbS}(\mathrm{s}) .\) What mass of sulfur can be recovered per cubic meter of this natural gas?

What is the oxidation state of sulfur in the following compounds? (a) \(\mathrm{SF}_{4} ;\) (b) \(\mathrm{S}_{2} \mathrm{F}_{10} ;\) (c) \(\mathrm{H}_{2} \mathrm{S} ;\) (d) \(\mathrm{CaSO}_{3}\).

What is the oxidation state of sulfur in the following compounds? (a) \(\mathrm{S}_{2} \mathrm{Br}_{2}\) (b) \(\mathrm{SCl}_{2}\) (c) \(\mathrm{Na}_{2} \mathrm{S}_{2} \mathrm{O}_{3}\) (d) \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{S}_{4} \mathrm{O}_{6}\).

Write balanced equations for the following important commercial reactions involving nitrogen and its compounds. (a) the principal artificial method of fixing atmospheric \(\mathrm{N}_{2}\) (b) oxidation of ammonia to \(\mathrm{NO}\) (c) preparation of nitric acid from \(\mathrm{NO}\).

When heated, each of the following substances decomposes to the products indicated. Write balanced equations for these reactions. (a) \(\mathrm{NH}_{4} \mathrm{NO}_{3}(\mathrm{s})\) to \(\mathrm{N}_{2}(\mathrm{g}), \mathrm{O}_{2}(\mathrm{g}),\) and \(\mathrm{H}_{2} \mathrm{O}(\mathrm{g})\) (b) \(\mathrm{NaNO}_{3}(\mathrm{s})\) to sodium nitrite and oxygen gas (c) \(\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{s})\) to lead(II) oxide, nitrogen dioxide, and oxygen.

Sodium nitrite can be made by passing oxygen and nitrogen monoxide gases into an aqueous solution of sodium carbonate. Write a balanced equation for this reaction.

Concentrated \(\mathrm{HNO}_{3}(\text { aq })\) used in laboratories is usually \(15 \mathrm{M} \mathrm{HNO}_{3}\) and has a density of \(1.41 \mathrm{g} \mathrm{mL}^{-1}\) What is the percent by mass of \(\mathrm{HNO}_{3}\) in this concentrated acid?

In \(1968,\) before pollution controls were introduced, over 75 billion gallons of gasoline were used in the United States as a motor fuel. Assume an emission of oxides of nitrogen of 5 grams per vehicle mile and an average mileage of 15 miles per gallon of gasoline. How many kilograms of nitrogen oxides were released into the atmosphere in the United States in \(1968 ?\)

Use information from this chapter and previous chapters to write chemical equations to represent the following: (a) equilibrium between nitrogen dioxide and dinitrogen tetroxide in the gaseous state (b) the reduction of nitrous acid by \(\mathrm{N}_{2} \mathrm{H}_{5}^{+}\) forming hydrazoic acid, followed by the reduction of additional nitrous acid by the hydrazoic acid, yielding nitrogen and dinitrogen monoxide (c) the neutralization of \(\mathrm{H}_{3} \mathrm{PO}_{4}(\mathrm{aq})\) to the second equivalence point by \(\mathrm{NH}_{3}(\mathrm{aq})\).