Chapter 5

Define oxidation and reduction in terms of (a) electron transfer and (b) oxidation numbers.

Why must both oxidation and reduction occur simultaneously during a redox reaction? What is an oxidizing agent and what happens to it in a redox reaction? What is a reducing agent and what happens to it in a redox reaction?

In the compound \(\mathrm{As}_{4} \mathrm{O}_{6}\), arsenic has an oxidation \(\mathrm{num}\) ber of +3 . What is the oxidation state of arsenic in this compound?

If the oxidation number of nitrogen in a certain molecule changes from +3 to -2 during a reaction, is the nitrogen oxidized or reduced? How many electrons are gained or lost by the nitrogen atom?

When balancing redox reactions, which side of a halfreaction gets the electrons?

The following equation is not balanced. $$ \mathrm{Ag}+\mathrm{Fe}^{2+} \longrightarrow \mathrm{Ag}^{+}+\mathrm{Fe} $$ Why? Use the ion-electron method to balance it. Can you balance this reaction by inspection?

What are the possible products of the reduction of sulfuric acid?

What is a single replacement reaction?

When manganese reacts with silver ions, is manganese oxidized or reduced? Is it an oxidizing agent or a reducing agent?

Define combustion, rusting, and tarnishing.

What are the major products produced in the combustion of \(\mathrm{C}_{10} \mathrm{H}_{22}\) under the following conditions: (a) an excess of oxygen, (b) a slightly limited oxygen supply, (c) a very limited supply of oxygen, (d) the compound is burned in air?

How does a titration of a redox reaction differ from the titration of an acid with a base?

Assign oxidation numbers to the atoms in the following: (a) \(\mathrm{ClO}_{4}^{-},\) (b) \(\mathrm{Cl}^{-}\), (c) \(\mathrm{SF}_{6}\), and (d) \(\mathrm{Au}\left(\mathrm{NO}_{3}\right)_{3}\).

Assign oxidation numbers to the atoms in the following: (a) \(S^{2-}\), (b) \(\mathrm{SO}_{2}\), (c) \(\mathrm{P}_{4},\) and (d) \(\mathrm{PH}_{3}\)

Assign oxidation numbers to each atom in the following: (a) \(\mathrm{Na} \mathrm{OBr}\) (b) \(\mathrm{NaBrO}_{2}\), (c) \(\mathrm{NaBrO}_{3},\) and (d) \(\mathrm{NaBrO}_{4}\)

Assign oxidation numbers to the elements in the and following: (a) \(\mathrm{MnCl}_{2}\), (b) \(\mathrm{MnO}_{4}^{-}\) (c) \(\mathrm{MnO}_{4}^{2-},\) (d) \(\mathrm{MnO}_{2}\)

Assign oxidation numbers to the elements in the following: (a) \(\mathrm{Bi}_{2} \mathrm{~S}_{3},\) (b) \(\mathrm{CeCl}_{4}\) (c) \(\mathrm{CsO}_{2},\) and (d) \(\mathrm{O}_{2} \mathrm{~F}_{2}\).

Assign oxidation numbers to the elements in the following: (a) \(\operatorname{Sr}\left(\mathrm{BrO}_{3}\right)_{2}\) (b) \(\mathrm{Cr}_{2} \mathrm{~S}_{3}\) (c) \(\mathrm{OF}_{2},\) and (d) \(\mathrm{HOF}\)

Titanium burns in pure nitrogen to form TiN. What are the oxidation states of titanium and nitrogen in TiN?

Zirconia, which is \(\mathrm{Zr} \mathrm{O}_{2}\), is used to make ceramic knives. What are the oxidation states of zirconium and oxygen in zirconia?

Ozone, \(\mathrm{O}_{3},\) is an allotrope of oxygen and is one of the oxidants responsible for the haze over the Smoky Mountains. What is the oxidation number of the oxygen atoms in ozone?

The other major air pollutant is \(\mathrm{NO}_{2}\). What are the oxidation numbers of the atoms in \(\mathrm{NO}_{2} ?\)

For the following reactions, identify the substance oxidized, the substance reduced, the oxidizing agent, and the reducing agent. $$ \begin{array}{l} \text { (a) } 2 \mathrm{HNO}_{3}+3 \mathrm{H}_{3} \mathrm{AsO}_{3} \longrightarrow \\ 2 \mathrm{NO}+3 \mathrm{H}_{3} \mathrm{AsO}_{4}+\mathrm{H}_{2} \mathrm{O} \\ \text { (b) } \mathrm{NaI}+3 \mathrm{HOCl} \longrightarrow \mathrm{NaIO}_{3}+3 \mathrm{HCl} \end{array} $$ (c) \(2 \mathrm{KMnO}_{4}+5 \mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}+3 \mathrm{H}_{2} \mathrm{SO}_{4} \longrightarrow\) $$ 10 \mathrm{CO}_{2}+\mathrm{K}_{2} \mathrm{SO}_{4}+2 \mathrm{MnSO}_{4}+8 \mathrm{H}_{2} \mathrm{O} $$ (d) \(6 \mathrm{H}_{2} \mathrm{SO}_{4}+2 \mathrm{Al} \longrightarrow \mathrm{Al}_{2}\left(\mathrm{SO}_{4}\right)_{3}+3 \mathrm{SO}_{2}+6 \mathrm{H}_{2} \mathrm{O}\)

For the following reactions, identify the substance oxidized, the substance reduced, the oxidizing agent, and the reducing agent. (a) \(\mathrm{Cu}+2 \mathrm{H}_{2} \mathrm{SO}_{4} \longrightarrow \mathrm{CuSO}_{4}+\mathrm{SO}_{2}+2 \mathrm{H}_{2} \mathrm{O}\) (b) \(3 \mathrm{SO}_{2}+2 \mathrm{HNO}_{3}+2 \mathrm{H}_{2} \mathrm{O} \longrightarrow 3 \mathrm{H}_{2} \mathrm{SO}_{4}+2 \mathrm{NO}\) (c) \(5 \mathrm{H}_{2} \mathrm{SO}_{4}+4 \mathrm{Zn} \longrightarrow 4 \mathrm{ZnSO}_{4}+\mathrm{H}_{2} \mathrm{~S}+4 \mathrm{H}_{2} \mathrm{O}\) (d) \(\mathrm{I}_{2}+10 \mathrm{HNO}_{3} \longrightarrow 2 \mathrm{HIO}_{3}+10 \mathrm{NO}_{2}+4 \mathrm{H}_{2} \mathrm{O}\)

One pollutant in smog is nitrogen dioxide, \(\mathrm{NO}_{2}\). The gas has a reddish brown color and is responsible for the redbrown color associated with this type of air pollution. \(\mathrm{Ni}\) trogen dioxide is also a contributor to acid rain because when rain passes through air contaminated with \(\mathrm{NO}_{2}\), it dissolves and undergoes the following reaction: \(\mathrm3{NO}_{2}(g)+\mathrm{H}_{2} \mathrm{O} \longrightarrow \mathrm{NO}(g)+2 \mathrm{H}^{+}(a q)+2 \mathrm{NO}_{3}^{-}(a q)\) In this reaction, which element is reduced and which is oxidized? Which is the oxidizing agent and which is the reducing agent?

Balance the following equations for reactions occurring in an acidic solution. (a) \(\mathrm{S}_{2} \mathrm{O}_{3}^{2-}+\mathrm{OCl}^{-} \longrightarrow \mathrm{Cl}^{-}+\mathrm{S}_{4} \mathrm{O}_{6}^{2-}\) (b) \(\mathrm{NO}_{3}^{-}+\mathrm{Cu} \longrightarrow \mathrm{NO}_{2}+\mathrm{Cu}^{2+}\) (c) \(\mathrm{IO}_{3}^{-}+\mathrm{H}_{3} \mathrm{AsO}_{3} \longrightarrow \mathrm{I}^{-}+\mathrm{H}_{3} \mathrm{AsO}_{4}\) (d) \(\mathrm{SO}_{4}^{2-}+\mathrm{Zn} \longrightarrow \mathrm{Zn}^{2+}+\mathrm{SO}_{2}\)

Balance the following equations for reactions occurring in an acidic solution. (a) \(\mathrm{NO}_{3}^{-}+\mathrm{Zn} \longrightarrow \mathrm{NH}_{4}^{+}+\mathrm{Zn}^{2+}\) (b) \(\mathrm{Cr}^{3+}+\mathrm{BiO}_{3}^{-} \longrightarrow \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}+\mathrm{Bi}^{3+}\) (c) \(\mathrm{I}_{2}+\mathrm{OCl}^{-} \longrightarrow \mathrm{IO}_{3}^{-}+\mathrm{Cl}^{-}\) (d) \(\mathrm{Mn}^{2+}+\mathrm{BiO}_{3}^{-} \longrightarrow \mathrm{MnO}_{4}^{-}+\mathrm{Bi}^{3+}\)

Balance the following equations for reactions occurring in an acidic solution. (a) \(\mathrm{Sn}+\mathrm{NO}_{3}^{-} \longrightarrow \mathrm{SnO}_{2}+\mathrm{NO}\) (b) \(\mathrm{PbO}_{2}+\mathrm{Cl}^{-} \longrightarrow \mathrm{PbCl}_{2}+\mathrm{Cl}_{2}\) (c) \(\mathrm{Ag}+\mathrm{NO}_{3}^{-} \longrightarrow \mathrm{NO}_{2}+\mathrm{Ag}^{+}\) (d) \(\mathrm{Fe}^{3+}+\mathrm{NH}_{3} \mathrm{OH}^{+} \longrightarrow \mathrm{Fe}^{2+}+\mathrm{N}_{2} \mathrm{O}\)

Balance the following equations for reactions occurring in an acidic solution. (a) \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}+\mathrm{HNO}_{2} \longrightarrow \mathrm{CO}_{2}+\mathrm{NO}\) (b) \(\mathrm{HNO}_{2}+\mathrm{MnO}_{4}^{-} \longrightarrow \mathrm{Mn}^{2+}+\mathrm{NO}_{3}^{-}\) (c) \(\mathrm{H}_{3} \mathrm{PO}_{2}+\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-} \longrightarrow \mathrm{H}_{3} \mathrm{PO}_{4}+\mathrm{Cr}^{3+}\) (d) \(\mathrm{XeF}_{2}+\mathrm{Cl}^{-} \longrightarrow \mathrm{Xe}+\mathrm{F}^{-}+\mathrm{Cl}_{2}\)

Hydroiodic acid reduces chlorine to hydrochloric acid and iodine. Write a balanced net ionic equation for the reaction.

Calcium oxalate is one of the minerals found in kidney stones. If a strong acid is added to calcium oxalate, the compound will dissolve and the oxalate ion will be changed to oxalic acid (a weak acid). Oxalate ion is a moderately strong reducing agent. Write a balanced net ionic equation for the oxidation of \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\) by \(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) in an acidic solution. The reaction yields \(\mathrm{Cr}^{3+}\) and \(\mathrm{CO}_{2}\) among the products.

Laundry bleach such as Clorox is a dilute solution of sodium hypochlorite, \(\mathrm{NaOCl}\). Write a balanced net ionic equation for the reaction of \(\mathrm{NaOCl}\) with \(\mathrm{Na}_{2} \mathrm{~S}_{2} \mathrm{O}_{3}\). The \(\mathrm{OCl}^{-}\) is reduced to chloride ion and the \(\mathrm{S}_{2} \mathrm{O}_{3}{\underline{\phantom{xx}}}^{2-}\) is oxidized to sulfate ion.

Write balanced molecular, ionic, and net ionic equations for the reactions of the following metals with hydrochloric acid to give hydrogen plus the metal ion in solution. (a) Manganese (gives \(\mathrm{Mn}^{2+}\) ) (b) Cadmium (gives \(\mathrm{Cd}^{2+}\) ) (c) \(\operatorname{Tin}\) (gives \(\mathrm{Sn}^{2+}\) )

Write balanced molecular, ionic, and net ionic equations for the reactions of the following metals with hydrochloric acid to give hydrogen plus the metal ion in solution. (a) Cobalt (gives \(\mathrm{Co}^{2+}\) ) (b) Cesium (gives \(\left.\mathrm{Cs}^{+}\right)\) (c) Zinc (gives \(\mathrm{Zn}^{2+}\) )

Write balanced molecular, ionic, and net ionic equations for the reaction of each of the following metals with dilute sulfuric acid. (a) Nickel (gives \(\mathrm{Ni}^{2+}\) ) (b) Chromium (gives \(\mathrm{Cr}^{3+}\) )

Write balanced molecular, ionic, and net ionic equations for the reaction of each of the following metals with dilute sulfuric acid. (a) Barium (gives \(\mathrm{Ba}^{2+}\) ) (b) Aluminum (gives \(\mathrm{Al}^{3+}\) )

The following reactions occur spontaneously. $$ \begin{aligned} 2 \mathrm{Y}+3 \mathrm{Ni}^{2+} & \longrightarrow 2 \mathrm{Y}^{3+}+3 \mathrm{Ni} \\ 2 \mathrm{Mo}+3 \mathrm{Ni}^{2+} & \longrightarrow 2 \mathrm{Mo}^{3+}+3 \mathrm{Ni} \\ \mathrm{Y}^{3+}+\mathrm{Mo} \longrightarrow & \mathrm{Y}+\mathrm{Mo}^{3+} \end{aligned} $$ List the metals Y, \(\mathrm{Ni}\), and Mo in order of increasing ease of oxidation.

Write balanced chemical equations for the complete compustion (in the presence of excess oxygen) of the following: (a) \(\mathrm{C}_{6} \mathrm{H}_{6}\) (benzene, an important industrial chemical and solvent), (b) \(\mathrm{C}_{4} \mathrm{H}_{10}\) (butane, a fuel used in cigarette ighters), (c) \(\mathrm{C}_{21} \mathrm{H}_{44}\) (a component of paraffin wax used in candles).

Write balanced chemical equations for the complete combustion (in the presence of excess oxygen) of the following: (a) \(\mathrm{C}_{12} \mathrm{H}_{26}\) (a component of kerosene), (b) \(\mathrm{C}_{18} \mathrm{H}_{36}\) (a component of diesel fuel), (c) \(\mathrm{C}_{7} \mathrm{H}_{8}\) (toluene, a raw material in the production of the explosive TNT).

The metabolism of carbohydrates such as glucose, \(\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6},\) produces the same products as combustion in excess oxygen. Write a chemical equation representing the metabolism (combustion) of glucose.

Methanol, \(\mathrm{CH}_{3} \mathrm{OH},\) has been suggested as an alternative to gasoline as an automotive fuel. Write a balanced chemical equation for its complete combustion.

Write the balanced equation for the combustion of dimethylsulfide, \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{~S},\) in an abundant supply of oxygen.

Thiophene, \(\mathrm{C}_{4} \mathrm{H}_{4} \mathrm{~S},\) is an impurity in crude oil and is a source of pollution if not removed. Write an equation for the combustion of thiophene.

In an acidic solution, permanganate ion reacts with tin(II) ion to give manganese(II) ion and tin(IV) ion. (a) Write a balanced net ionic equation for the reaction. (b) How many milliliters of \(0.230 M\) potassium permanganate solution are needed to react completely with \(40.0 \mathrm{~mL}\) of \(0.250 \mathrm{M}\) tin(II) chloride solution?

Iodate ion reacts with sulfite ion to give sulfate ion and iodide ion. (a) Write a balanced net ionic equation for the reaction. (b) How many grams of sodium sulfite are needed to react with \(5.00 \mathrm{~g}\) of sodium iodate?

Sulfites are used worldwide in the wine industry as antioxidant and antimicrobial agents. However, sulfites have also been identified as causing certain allergic reactions suffered by asthmatics, and the FDA mandates that sulfites be identified on the label if they are present at levels of 10 ppm (parts per million) or higher. The analysis of sulfites in wine uses the "Ripper method" in which a standard iodine solution, prepared by the reaction of iodate and iodide ions, is used to titrate a sample of the wine. The iodine is formed in the reaction $$ \mathrm{IO}_{3}^{-}+5 \mathrm{I}^{-}+6 \mathrm{H}^{+} \longrightarrow 3 \mathrm{I}_{2}+3 \mathrm{H}_{2} \mathrm{O} $$ The iodine is held in solution by adding an excess of \(\mathrm{I}^{-}\), which combines with \(\mathrm{I}_{2}\) to give \(\mathrm{I}_{3}^{-}\). In the titration, the \(\mathrm{SO}_{3}^{2-}\) is converted to \(\mathrm{SO}_{2}\) by acidification, and the reaction during the titration is $$ \mathrm{SO}_{2}+\mathrm{I}_{3}^{-}+2 \mathrm{H}_{2} \mathrm{O} \longrightarrow \mathrm{SO}_{4}^{2-}+3 \mathrm{I}^{-}+4 \mathrm{H}^{+} $$ Starch is added to the wine sample to detect the end point, which is signaled by the formation of a dark blue color when excess iodine binds to the starch molecules. In a certain analysis, \(0.0421 \mathrm{~g}\) of \(\mathrm{NaIO}_{3}\) was dissolved in dilute acid and excess NaI was added to the solution, which was then diluted to a total volume of \(100.0 \mathrm{~mL}\) A \(50.0 \mathrm{~mL}\) sample of wine was then acidified and titrated with the iodine- containing solution. The volume of iodine solution required was \(2.47 \mathrm{~mL}\). (a) What was the molarity of the iodine (actually, \(\left.\mathrm{I}_{3}^{-}\right)\) in the standard solution? (b) How many grams of \(\mathrm{SO}_{2}\) were in the wine sample? (c) If the density of the wine was \(0.96 \mathrm{~g} / \mathrm{mL}\), what was the percentage of \(\mathrm{SO}_{2}\) in the wine? (d) Parts per million (ppm) is calculated in a manner similar to percent (which is equivalent to parts per hundred). $$ \mathrm{ppm}=\frac{\text { grams of component }}{\text { grams of sample }} \times 10^{6} \mathrm{ppm} $$ What was the concentration of sulfite in the wine, expressed as parts per million \(\mathrm{SO}_{2} ?\)

A \(1.362 \mathrm{~g}\) sample of an iron ore that contained \(\mathrm{Fe}_{3} \mathrm{O}_{4}\) was dissolved in acid and all of the iron was reduced to \(\mathrm{Fe}^{2+} .\) The solution was then acidified with \(\mathrm{H}_{2} \mathrm{SO}_{4}\) and titrated with \(39.42 \mathrm{~mL}\) of \(0.0281 \mathrm{M} \mathrm{KMnO}_{4}\), which oxidized the iron to \(\mathrm{Fe}^{3+}\). The net ionic equation for the reaction is \(5 \mathrm{Fe}^{2+}+\mathrm{MnO}_{4}^{-}+8 \mathrm{H}^{+} \longrightarrow 5 \mathrm{Fe}^{3+}+\mathrm{Mn}^{2+}+4 \mathrm{H}_{2} \mathrm{O}\) (a) What was the percentage by mass of iron in the ore? (b) What was the percentage by mass of \(\mathrm{Fe}_{3} \mathrm{O}_{4}\) in the ore?

Both calcium chloride and sodium chloride are used to melt ice and snow on roads in the winter. A certain company was marketing a mixture of these two compounds for this purpose. A chemist, wanting to analyze the mixture, dissolved \(2.463 \mathrm{~g}\) of it in water and precipitated calcium oxalate by adding sodium oxalate, \(\mathrm{Na}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\) The calcium oxalate was carefully filtered from the solution, dissolved in sulfuric acid, and titrated with 0.1000 \(M \mathrm{KMnO}_{4}\) solution. The reaction that occurred was \(6 \mathrm{H}^{+}+5 \mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}+2 \mathrm{MnO}_{4}^{-} \longrightarrow\) $$ 10 \mathrm{CO}_{2}+2 \mathrm{Mn}^{2+}+8 \mathrm{H}_{2} \mathrm{O} $$ The titration required \(21.62 \mathrm{~mL}\) of the \(\mathrm{KMnO}_{4}\) solution. (a) How many moles of \(\mathrm{C}_{2} \mathrm{O}_{4}^{2-}\) were present in the calcium oxalate precipitate? (b) How many grams of calcium chloride were in the original \(2.463 \mathrm{~g}\) sample? (c) What was the percentage by mass of calcium chloride in the sample?

One way to analyze a sample for nitrite ion is to acidify a solution containing \(\mathrm{NO}_{2}^{-}\) and then allow the \(\mathrm{HNO}_{2}\) that is formed to react with iodide ion in the presence of excess \(\mathrm{I}^{-}\). The reaction is $$ 2 \mathrm{HNO}_{2}+2 \mathrm{H}^{+}+3 \mathrm{I}^{-} \longrightarrow 2 \mathrm{NO}+2 \mathrm{H}_{2} \mathrm{O}+\mathrm{I}_{3}- $$ Then the \(\mathrm{I}_{3}^{-}\) is titrated with \(\mathrm{Na}_{2} \mathrm{~S}_{2} \mathrm{O}_{3}\) solution using starch as an indicator. $$ \mathrm{I}_{3}^{-}+2 \mathrm{~S}_{2} \mathrm{O}_{3}^{2-} \longrightarrow 3 \mathrm{I}^{-}+\mathrm{S}_{4} \mathrm{O}_{6}^{2-} $$ In a typical analysis, a \(1.104 \mathrm{~g}\) sample that was known to contain \(\mathrm{NaNO}_{2}\) was treated as described above. The titration required \(29.25 \mathrm{~mL}\) of \(0.3000 \mathrm{M} \mathrm{Na}_{2} \mathrm{~S}_{2} \mathrm{O}_{3}\) solu- tion to reach the end point. (a) How many moles of \(\mathrm{I}_{3}^{-}\) had been produced in the first reaction? (b) How many moles of \(\mathrm{NO}_{2}^{-}\) had been in the original \(1.104 \mathrm{~g}\) sample? (c) What was the percentage by mass of \(\mathrm{NaNO}_{2}\) in the original sample?

Use oxidation numbers to show that the fermentation of glucose, \(\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6},\) to carbon dioxide and ethanol, \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH},\) is a redox reaction.