Chapter 21

Name the following compounds. (a) \(\mathrm{HBrO}_{3}\) (b) KIO (c) \(\mathrm{NaClO}_{2}\) (d) \(\mathrm{NaBrO}_{4}\)

Write the formula for each of the following compounds. (a) potassium bromite (b) calcium bromide (c) sodium periodate (d) magnesium hypochlorite

Give the formula for the acidic oxide of (a) \(\mathrm{HNO}_{3}\) (b) \(\mathrm{HNO}_{2}\) (c) \(\mathrm{H}_{2} \mathrm{SO}_{4}\)

Write the formula of the acid formed when each of these acidic oxides reacts with water. (a) \(\mathrm{SO}_{2}\) (b) \(\mathrm{Cl}_{2} \mathrm{O}\) (c) \(\mathrm{P}_{4} \mathrm{O}_{6}\)

Write the formulas of the following compounds. (a) ammonia (b) laughing gas (c) hydrogen peroxide (d) sulfur trioxide

Write the formula for the following compounds. (a) sodium azide (b) sulfurous acid (c) hydrazine (d) sodium dihydrogen phosphate

Write the formula of a compound of hydrogen with (a) sulfur. (b) nitrogen, which is a liquid at \(25^{\circ} \mathrm{C}\) and \(1 \mathrm{~atm}\). (c) phosphorus, which is a poisonous gas at \(25^{\circ} \mathrm{C}\) and \(1 \mathrm{~atm}\).

Give the formula of (a) an anion in which \(S\) has an oxidation number of \(-2\). (b) two anions in which \(\mathrm{S}\) has an oxidation number of \(+4\). (c) two different acids of sulfur.

Give the formula of a compound of nitrogen that is (a) a weak base. (b) a strong acid. (c) a weak acid. (d) capable of oxidizing copper.

Write a balanced net ionic equation for (a) the electrolytic decomposition of hydrogen fluoride. (b) the oxidation of iodide ion to iodine by hydrogen peroxide in acidic solution. Hydrogen peroxide is reduced to water.

Write a balanced net ionic equation for (a) the oxidation of iodide to iodine by sulfate ion in acidic solution. Sulfur dioxide gas is also produced. (b) The preparation of iodine from an iodide salt and chlorine gas.

Write a balanced net ionic equation for the disproportionation reaction (a) of iodine to give iodate and iodide ions in basic solution. (b) of chlorine gas to chloride and perchlorate ions in basic solution.

Write a balanced net ionic equation for the disproportionation reaction of (a) hypochlorous acid to chlorine gas and chlorous acid in acidic solution. (b) chlorate ion to perchlorate and chlorite ions.

Complete and balance the following equations. If no reaction occurs, write NR. (a) \(\mathrm{Cl}_{2}(g)+\mathrm{I}^{-}(a q) \longrightarrow\) (b) \(\mathrm{F}_{2}(g)+\mathrm{Br}^{-}(a q) \longrightarrow\) (c) \(\mathrm{I}_{2}(s)+\mathrm{Cl}^{-}(a q) \longrightarrow\) (d) \(\mathrm{Br}_{2}(l)+\mathrm{I}^{-}(a q) \longrightarrow\)

Complete and balance the following equations. If no reaction occurs, write NR. (a) \(\mathrm{Cl}_{2}(g)+\mathrm{Br}^{-}(a q) \longrightarrow\) (b) \(\mathrm{I}_{2}(s)+\mathrm{Cl}^{-}(a q) \longrightarrow\) (c) \(\mathrm{I}_{2}(s)+\mathrm{Br}^{-}(a q) \longrightarrow\) (d) \(\mathrm{Br}_{2}(l)+\mathrm{Cl}^{-}(a q) \longrightarrow\)

Write a balanced equation for the preparation of (a) \(\mathrm{N}_{2}\) from \(\mathrm{Pb}\left(\mathrm{N}_{3}\right)_{2}\) (b) \(\mathrm{O}_{2}\) from \(\mathrm{O}_{3}\) (c) \(\mathrm{S}\) from \(\mathrm{H}_{2} \mathrm{~S}\)

Write a balanced equation for the reaction of ammonia with (a) \(\mathrm{Cu}^{2+}\) (b) \(\mathrm{H}^{+}\) (c) \(\mathrm{Al}^{3+}\)

Write a balanced equation for the reaction of hydrogen sulfide with (a) \(\mathrm{Cd}^{2+}\) (b) \(\mathrm{OH}^{-}\) (c) \(\mathrm{O}_{2}(g)\)

Write a balanced net ionic equation for the reaction of nitric acid with (a) a solution of \(\mathrm{Ca}(\mathrm{OH})_{2}\). (b) \(\mathrm{Ag}(s)\); assume the nitrate ion is reduced to \(\mathrm{NO}_{2}(g)\) (c) \(\mathrm{Cd}(s)\); assume the nitrate ion is reduced to \(\mathrm{N}_{2}(g)\).

Write a balanced net ionic equation for the reaction of sulfuric acid with (a) \(\mathrm{CaCO}_{3}(s)\) (b) a solution of \(\mathrm{NaOH}\). (c) Cu; assume the \(\mathrm{SO}_{4}{\underline{\phantom{xx}}}^{2-}\) ion is reduced to \(\mathrm{SO}_{2}\).

Give the Lewis structure of (a) \(\mathrm{NO}_{2}\) (b) \(\mathrm{NO}\) (c) \(\mathrm{SO}_{2}\) (d) \(\mathrm{SO}_{3}\)

Give the Lewis structure of (a) \(\mathrm{Cl}_{2} \mathrm{O}\) (b) \(\mathrm{N}_{2} \mathrm{O}\) (c) \(\mathrm{P}_{4}\) (d) \(\mathbf{N}_{2}\)

Give the Lewis structure of (a) \(\mathrm{HNO}_{3}\) (b) \(\mathrm{H}_{2} \mathrm{SO}_{4}\) (c) \(\mathrm{H}_{3} \mathrm{PO}_{4}\)

Give the Lewis structure of (a) the strongest oxoacid of bromine. (b) a hydrogen compound of nitrogen in which there is an \(-\mathrm{N}-\mathrm{N}-\) bond. (c) an acid added to cola drinks.

Give the Lewis structure of (a) an oxide of nitrogen in the \(+5\) state. (b) the strongest oxoacid of nitrogen. (c) a tetrahedral oxoanion of sulfur.

The average concentration of bromine (as bromide) in seawater is \(65 \mathrm{ppm} .\) Calculate (a) the volume of seawater \(\left(d=64.0 \mathrm{lb} / \mathrm{ft}^{3}\right)\) in cubic feet required to produce one kilogram of liquid bromine. (b) the volume of chlorine gas in liters, measured at \(20^{\circ} \mathrm{C}\) and \(762 \mathrm{~mm}\) \(\mathrm{Hg}\), required to react with this volume of sea water.

37\. Iodine can be prepared by allowing an aqueous solution of hydrogen iodide to react with manganese dioxide, \(\mathrm{MnO}_{2}\). The reaction is $$2 \mathrm{I}^{-}(a q)+4 \mathrm{H}^{+}(a q)+\mathrm{MnO}_{2}(s) \longrightarrow \mathrm{Mn}^{2+}(a q)+2 \mathrm{H}_{2} \mathrm{O}+\mathrm{I}_{2}(s)$$ If an excess of hydrogen iodide is added to \(0.200 \mathrm{~g}\) of \(\mathrm{MnO}_{2}\), how many grams of iodine are obtained, assuming \(100 \%\) yield?

When a solution of hydrogen bromide is prepared, \(1.283 \mathrm{~L}\) of \(\mathrm{HBr}\) gas at \(25^{\circ} \mathrm{C}\) and \(0.974\) atm is bubbled into \(250.0 \mathrm{~mL}\) of water. Assuming all the HBr dissolves with no volume change, what is the molarity of the hydrobromic acid solution produced?

Sulfur dioxide can be removed from the smokestack emissions of power plants by reacting it with hydrogen sulfide, producing sulfur and water. What volume of hydrogen sulfide at \(27^{\circ} \mathrm{C}\) and \(755 \mathrm{~mm} \mathrm{Hg}\) is required to remove the sulfur dioxide produced by a power plant that burns one metric ton of coal containing \(5.0 \%\) sulfur by mass? How many grams of sulfur are produced by the reaction of \(\mathrm{H}_{2} \mathrm{~S}\) with \(\mathrm{SO}_{2} ?\)

A \(1.500-\mathrm{g}\) sample containing sodium nitrate was heated to form \(\mathrm{NaNO}_{2}\) and \(\mathrm{O}_{2}\). The oxygen evolved was collected over water at \(23^{\circ} \mathrm{C}\) and \(752 \mathrm{~mm} \mathrm{Hg}\); its volume was \(125.0 \mathrm{~mL}\). Calculate the percentage of \(\mathrm{NaNO}_{3}\) in the sample. The vapor pressure of water at \(23^{\circ} \mathrm{C}\) is \(21.07 \mathrm{~mm} \mathrm{Hg}\).

Chlorine can remove the foul smell of \(\mathrm{H}_{2} \mathrm{~S}\) in water. The reaction is $$\mathrm{H}_{2} \mathrm{~S}(a q)+\mathrm{Cl}_{2}(a q) \longrightarrow 2 \mathrm{H}^{+}(a q)+2 \mathrm{Cl}^{-}(a q)+\mathrm{S}(s)$$ If the contaminated water has \(5.0\) ppm hydrogen sulfide by mass, what volume of chlorine gas at STP is required to remove all the \(\mathrm{H}_{2} \mathrm{~S}\) from \(1.00 \times 10^{3}\) gallons of water \((d=1.00 \mathrm{~g} / \mathrm{mL}) ?\) What is the \(\mathrm{pH}\) of the solution after treatment with chlorine?

The equilibrium constant at \(25^{\circ} \mathrm{C}\) for the reaction $$\mathrm{Br}_{2}(l)+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{H}^{+}(a q)+\mathrm{Br}^{-}(a q)+\mathrm{HBrO}(a q)$$ is \(1.2 \times 10^{-9} .\) This is the system present in a bottle of "bromine water." Assuming that HBrO does not ionize appreciably, what is the pH of the bromine water?

Calculate the \(\mathrm{pH}\) and the equilibrium concentration of \(\mathrm{HClO}\) in a \(0.10 M\) solution of hypochlorous acid. \(K_{\mathrm{a}} \mathrm{HClO}=2.8 \times 10^{-8}\)

At equilibrium, a gas mixture has a partial pressure of \(0.7324\) atm for \(\mathrm{HBr}\) and \(2.80 \times 10^{-3}\) atm for both hydrogen and bromine gases. What is \(K\) for the formation of two moles of HBr from \(\mathrm{H}_{2}\) and \(\mathrm{Br}_{2} ?\)

Given $$\begin{array}{ll} \mathrm{HF}(a q) \rightleftharpoons \mathrm{H}^{+}(a q)+\mathrm{F}^{-}(a q) & K_{\mathrm{a}}=6.9 \times 10^{-1} \\ \mathrm{HF}(a q)+\mathrm{F}^{-}(a q) \rightleftharpoons \mathrm{HF}_{2}-(a q) & K=2.7 \end{array}$$ calculate \(K\) for the reaction $$ 2 \mathrm{HF}(a q) \rightleftharpoons \mathrm{H}^{+}(a q)+\mathrm{HF}_{2}^{-}(a q) $$

What is the concentration of fluoride ion in a water solution saturated with \(\mathrm{BaF}_{2}, K_{s \mathrm{p}}=1.8 \times 10^{-7}\) ?

Consider the equilibrium system $$\mathrm{HF}(a q) \rightleftharpoons \mathrm{H}^{+}(a q)+\mathrm{F}^{-}(a q)$$ Given \(\Delta H_{\mathrm{f}}^{\circ} \mathrm{HF}(a q)=-320.1 \mathrm{~kJ} / \mathrm{mol}\), $$\begin{aligned} \Delta H_{\mathrm{f}}^{\circ} \mathrm{F}^{-}(a q) &=-332.6 \mathrm{~kJ} / \mathrm{mol} ; S^{\circ} \mathrm{F}^{-}(a q)=-13.8 \mathrm{~J} / \mathrm{mol} \cdot \mathrm{K} \\ K_{\mathrm{a}} \mathrm{HF} &=6.9 \times 10^{-4} \text { at } 25^{\circ} \mathrm{C} \end{aligned}$$ calculate \(S^{\circ}\) for \(\mathrm{HF}(a q)\).

In the electrolysis of a KI solution, using \(5.00 \mathrm{~V}\), how much electrical energy in kilojoules is consumed when one mole of \(I_{2}\) is formed?

If an electrolytic cell producing fluorine uses a current of \(7.00 \times 10^{3} \mathrm{~A}\) (at \(10.0 \mathrm{~V})\), how many grams of fluorine gas can be produced in two days (assuming that the cell operates continuously at \(95 \%\) efficiency)?

Consider the reduction of nitrate ion in acidic solution to nitrogen oxide \(\left(E_{\mathrm{red}}^{\circ}=0.964 \mathrm{~V}\right)\) by sulfur dioxide that is oxidized to sulfate ion \(\left(E_{\text {red }}^{o}=0.155 \mathrm{~V}\right)\). Calculate the voltage of a cell involving this reaction in which all the gases have pressures of \(1.00 \mathrm{~atm}\), all the ionic species (except \(\left.\mathrm{H}^{+}\right)\) are at \(0.100 M\), and the \(\mathrm{pH}\) is \(4.30 .\)

Choose the strongest acid from each group. (a) \(\mathrm{HClO}, \mathrm{HBrO}, \mathrm{HIO}\) (b) \(\mathrm{HIO}, \mathrm{HIO}_{3}, \mathrm{HIO}_{4}\) (c) \(\mathrm{HIO}, \mathrm{HBrO}_{2}, \mathrm{HBrO}_{4}\)

What intermolecular forces are present in the following? (a) \(\mathrm{Cl}_{2}\) (b) \(\mathrm{HBr}\) (c) \(\mathrm{HF}\) (d) \(\mathrm{HClO}_{4}\) (e) \(\mathrm{MgI}_{2}\)

Write a balanced equation for the reaction of hydrofluoric acid with \(\mathrm{SiO}_{2}\). What volume of \(2.0 \mathrm{M} \mathrm{HF}\) is required to react with one gram of silicon dioxide?

State the oxidation number of \(\mathrm{N}\) in (a) \(\mathrm{NO}_{2}^{-}\) (b) \(\mathrm{NO}_{2}\) (c) \(\mathrm{HNO}_{3}\) (d) \(\mathrm{NH}_{4}{\underline{\phantom{xx}}}^{+}\)

Why does concentrated nitric acid often have a yellow color even though pure \(\mathrm{HNO}_{3}\) is colorless?

Explain why (a) acid strength increases as the oxidation number of the central non- metal atom increases. (b) nitrogen dioxide is paramagnetic. (c) the oxidizing strength of an oxoanion is inversely related to \(\mathrm{pH}\). (d) sugar turns black when treated with concentrated sulfuric acid.

The amount of sodium hypochlorite in a bleach solution can be determined by using a given volume of bleach to oxidize excess iodide ion to iodine; \(\mathrm{ClO}^{-}\) is reduced to \(\mathrm{Cl}^{-}\). The amount of iodine produced by the redox reaction is determined by titration with sodium thiosulfate, \(\mathrm{Na}_{2} \mathrm{~S}_{2} \mathrm{O}_{3} ; \mathrm{I}_{2}\) is reduced to \(\mathrm{I}^{-}\). The sodium thiosulfate is oxidized to sodium tetrathionate, \(\mathrm{Na}_{2} \mathrm{~S}_{4} \mathrm{O}_{6}\). In this analysis, potassium iodide was added in excess to \(5.00 \mathrm{~mL}\) of bleach \(\left(d=1.00 \mathrm{~g} / \mathrm{cm}^{3}\right)\). If \(25.00 \mathrm{~mL}\) of \(0.0700 \mathrm{M} \mathrm{Na}_{2} \mathrm{~S}_{2} \mathrm{O}_{3}\) was required to reduce all the iodine produced by the bleach back to iodide, what is the mass percent of \(\mathrm{NaClO}\) in the bleach?