Chapter Questions
Name the following species.(a) $\mathrm{HIO}_{4}$(b) $\mathrm{BrO}_{2}^{-}$(c) $\mathrm{HIO}$(d) $\mathrm{NaClO}_{3}$
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) chloric acid(b) periodic acid(c) hypobromous acid(d) hydriodic acid
Write the formula for each of the following compounds.(a) potassium bromite(b) calcium bromide(c) sodium periodate(d) magnesium hypochlorite
Write the formula of a compound of each of the following elements that cannot act as an oxidizing agent.
Write the formula of an oxoanion of each of the following elements that cannot act as a reducing agent.
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.
Write the formulas of the following compounds.(a) ammonia(b) laughing gas(c) hydrogen peroxide(d) sulfur trioxide
Write the formula of a compound of hydrogen with(a) nitrogen, which is a gas at $25^{\circ} \mathrm{C}$ and $1 \mathrm{~atm}$.(b) phosphorus, which is a liquid at $25^{\circ} \mathrm{C}$ and $1 \mathrm{~atm}$.(c) oxygen, which contains an $\mathrm{O}-\mathrm{O}$ bond.
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 $\mathrm{S}$ has an oxidation number of -2 .(b) two anions in which $S$ 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) $\mathrm{F}_{2}$ from HF.(b) $\mathrm{Br}_{2}$ from NaBr.(c) $\mathrm{NH}_{4}^{+}$ from $\mathrm{NH}_{3}$
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 $\mathrm{SO}_{4}{ }^{2-}$ ion is reduced to $\mathrm{SO}_{2}$.
Give the Lewis structure of(a) $\mathrm{NO}_{2}$(b) 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) $\mathrm{N}_{2}$
Which of the molecules in Question 27 are polar?
Which of the molecules in Question 28 are polar?
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 structures of the conjugate bases of the species in Question 31 .
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 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 seawater.
A 425-gal tank is filled with water containing $175 \mathrm{~g}$ of sodium iodide. How many liters of chlorine gas at $758 \mathrm{~mm} \mathrm{Hg}$ and $25^{\circ} \mathrm{C}$ will be required to oxidize all the iodide to iodine?
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 $\mathrm{HBr}$ dissolves with no volume change, what is the molarity of the hydrobromic acid solution produced?
When ammonium nitrate explodes, nitrogen, steam, and oxygen gas are produced. If the explosion is carried out by heating one kilogram of ammo-nium nitrate sealed in a rigid bomb with a volume of one liter, what is the total pressure produced by the gases before the bomb ruptures? Assume that the reaction goes to completion and that the final temperature is $500^{\circ} \mathrm{C}$ (3 significant figures).
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 -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 $\mathrm{pH}$ 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 $2.80 \times 10^{-3} \mathrm{~atm}$ for both hydrogen and bromine gases. What is $\mathrm{K}$ for the formation of two moles of $\mathrm{HBr}$ from $\mathrm{H}_{2}$ and $\mathrm{Br}_{2}$ ?
Given$$\begin{aligned}\mathrm{HF}(a q) & \rightleftharpoons \mathrm{H}^{+}(a q)+\mathrm{F}^{-}(a q) & & K_{\mathrm{a}}=6.9 \times 10^{-4} \\\mathrm{HF}(a q)+\mathrm{F}^{-}(a q) & \rightleftharpoons \mathrm{HF}_{2}^{-}(a q) & & K=2.7\end{aligned}$$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 $\mathrm{BaF}_{2}, K_{\mathrm{sp}}=1.8 \times 10^{-7} ?$
Calculate the solubility in grams per $100 \mathrm{~mL}$ of $\mathrm{BaF}_{2}$ in a $0.10 \mathrm{M} \mathrm{BaCl}_{2}$ solution.
Determine whether the following redox reaction is spontaneous at $25^{\circ} \mathrm{C}$ and 1 atm:$$2 \mathrm{KIO}_{3}(s)+\mathrm{Cl}_{2}(g) \longrightarrow 2 \mathrm{KClO}_{3}(s)+\mathrm{I}_{2}(s)$$Use data in Appendix 1 and the following information: $\Delta H_{\mathrm{f}}^{\circ} \mathrm{KIO}_{3}(s)=$ $-501.4 \mathrm{~kJ} / \mathrm{mol}, \mathrm{S}^{\circ} \mathrm{KIO}_{3}(s)=151.5 \mathrm{~J} / \mathrm{mol} \cdot \mathrm{K} .$ What is the lowest temperatureat which the reaction is spontaneous?
Follow the directions for Problem 49 for the reaction$$2 \mathrm{KBrO}_{3}(s)+\mathrm{Cl}_{2}(g) \longrightarrow 2 \mathrm{KClO}_{3}(s)+\mathrm{Br}_{2}(l)$$The following thermodynamic data may be useful:$$\Delta H_{\mathrm{f}}^{\circ} \mathrm{KBrO}_{3}=-360.2 \mathrm{~kJ} / \mathrm{mol} ; S^{\circ} \mathrm{KBrO}_{3}=149.2 \mathrm{~J} / \mathrm{mol} \cdot \mathrm{K}$$
Consider the equilibrium system$$\begin{aligned}& \mathrm{HF}(a q) \rightleftharpoons \mathrm{H}^{+}(a q)+\mathrm{F}^{-}(a q) \\\text { Given } \Delta H_{\mathrm{f}}^{\circ} \mathrm{HF}(a q)=&-320.1 \mathrm{k} \mathrm{J} / \mathrm{mol} \\\Delta H_{\mathrm{f}}^{\circ} \mathrm{F}^{-}(a q)=&-332.6 \mathrm{~kJ} / \mathrm{mol} ; S^{\circ} \mathrm{F}^{-}(a q)=-13.8 \mathrm{~J} / \mathrm{mol} \cdot \mathrm{K} ; \\K_{\mathrm{a}} \mathrm{HF}=& 6.9 \times 10^{-4} \text {at } 25^{\circ} \mathrm{C}\end{aligned}$$calculate $S^{\circ}$ for $\mathrm{HF}(a q)$.
Applying the Tables in Appendix 1 to$$4 \mathrm{HCl}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{Cl}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(l)$$determine(a) whether the reaction is spontaneous at $25^{\circ} \mathrm{C}$ and $1 \mathrm{~atm}$.(b) $K$ for the reaction at $25^{\circ} \mathrm{C}$.
Consider the reaction$$4 \mathrm{NH}_{3}(g)+5 \mathrm{O}_{2}(g) \longrightarrow 4 \mathrm{NO}(g)+6 \mathrm{H}_{2} \mathrm{O}(g)$$(a) Calculate $\Delta H^{\circ}$ for this reaction. Is it exothermic or endothermic?(b) Would you expect $\Delta S^{\circ}$ to be positive or negative? Calculate $\Delta S^{\circ}$.(c) Is the reaction spontaneous at $25^{\circ} \mathrm{C}$ and $1 \mathrm{~atm} ?$(d) At what temperature, if any, is the reaction at equilibrium at 1 atmpressure?
Data are given in Appendix 1 for white phosphorus, $\mathrm{P}_{4}(s) . \mathrm{P}_{4}(g)$ has the following thermodynamic values: $\Delta H_{\mathrm{f}}^{\circ}=58.9 \mathrm{~kJ} / \mathrm{mol}, S^{\circ}=280.0 \mathrm{~J} / \mathrm{K} \cdot \mathrm{mol}$. What is the temperature at which white phosphorus sublimes at 1 atm pressure?
In the electrolysis of a KI solution, using $5.00 \mathrm{~V}$, how much electrical energy in kilojoules is consumed when one mole of $\mathrm{I}_{2}$ 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)?
Sodium hypochlorite is produced by the electrolysis of cold sodium chloride solution. How long must a cell operate to produce $1.500 \times 10^{3} \mathrm{~L}$ of $5.00 \%$ NaClO by mass if the cell current is $2.00 \times 10^{3} \mathrm{~A} ?$ Assume that the density of the solution is $1.00 \mathrm{~g} / \mathrm{cm}^{3}$.
Sodium perchlorate is produced by the electrolysis of sodium chlorate. If a current of $1.50 \times 10^{3} \mathrm{~A}$ passes through an electrolytic cell, how many kilograms of sodium perchlorate are produced in an eight-hour run?
Taking $E_{\text {ox }}^{\circ} \mathrm{H}_{2} \mathrm{O}_{2}=-0.695 \mathrm{~V},$ determine which of the following species will be reduced by hydrogen peroxide (use Table 17.1 to find $E_{\mathrm{red}}^{\circ}$ values).(a) $\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}$(b) $\mathrm{Fe}^{2+}$(c) $\mathrm{I}_{2}$(d) $\mathrm{Br}_{2}$
Taking $E_{\text {red }}^{\circ} \mathrm{H}_{2} \mathrm{O}_{2}=+1.763 \mathrm{~V}$, determine which of the following species will be oxidized by hydrogen peroxide (use Table 17.1 to find $E_{\alpha x}^{\circ}$ values).(a) $\mathrm{Co}^{2+}$(b) $\mathrm{Cl}^{-}$(c) $\mathrm{Fe}^{2+}$(d) $\mathrm{Sn}^{2+}$
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 $\left(E_{\mathrm{red}}^{\circ}=0.155 \mathrm{~V}\right) .$ 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 $0.100 \mathrm{M},$ and the $\mathrm{pH}$ is 4.30 .
For the reaction in Problem 61 if gas pressures are at 1.00 atm and ionic species are at $0.100 M$ (except $\mathrm{H}^{+}$ ), at what pH will the voltage be $1.000 \mathrm{~V} ?$
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) $\mathrm{NH}_{4}{ }^{+}$
The density of sulfur vapor at one atmosphere pressure and $973 \mathrm{~K}$ is $0.8012 \mathrm{~g} / \mathrm{L} .$ What is the molecular formula of the vapor?
Give the formula of a substance discussed in this chapter that is used(a) to disinfect water.(b) in safety matches.(c) to prepare hydrazine.(d) to etch glass.
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 nonmetal 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.
Wuppose you wish to calculate the mass of sulfuric acid that can be obtained from an underground deposit of sulfur $1.00 \mathrm{~km}^{2}$ in area. What additional information do you need to make this calculation?
The reaction$$4 \mathrm{HF}(a q)+\mathrm{SiO}_{2}(a q) \longrightarrow \mathrm{SiF}_{4}(a q)+2 \mathrm{H}_{2} \mathrm{O}$$can be used to release gold that is distributed in certain quartz $\left(\mathrm{SiO}_{2}\right)$ veins of hydrothermal origin. If the quartz contains $1.0 \times 10^{-3} \%$ Au by weight and the gold has a market value of $\$ 425$ per troy ounce, would the process be economically feasible if commercial HF (50\% by weight, $d=1.17 \mathrm{~g} / \mathrm{cm}^{3}$ ) costs $75 c$ a liter? ( 1 troy ounce $=31.1 \mathrm{~g} .$ )
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 toreduce all the iodine produced by the bleach back to iodide, what is the mass percent of $\mathrm{NaClO}$ in the bleach?
What is the minimum amount of sodium azide, $\mathrm{NaN}_{3}$, that can be added to an automobile airbag to give a volume of $20.0 \mathrm{~L}$ of $\mathrm{N}_{2}(g)$ on inflation? Make any reasonable assumptions required to obtain an answer, but state what these assumptions are.