Chapter 18

Describe oxidation and reduction. Compare the electron transfer in a redox reaction with the electron donation in a Lewis acid-base reaction.

List the halogens in order of increasing oxidizing power.

List four species that can oxidize \(\mathrm{Fe}^{2+}\) to \(\mathrm{Fe}^{3+}\).

In a "dead" battery, the chemical reaction has come to equilibrium. What are the values of \(\Delta G\) and \(E\) for a dead battery?

What is the difference between a battery and a fuel cell?

What are the differences between anodic and cathodic protection from corrosion?

Assign the oxidation numbers of all atoms in the following species. (a) \(\mathrm{ClO}_{3}^{-}\) (b) \(\mathrm{PF}_{3}\) (c) \(\mathrm{CO}\)

Assign the oxidation numbers of all atoms in the following species. (a) \(\mathrm{N}_{2}\) (b) \(\mathrm{B}(\mathrm{OH})_{3}\) (c) \(\mathrm{IF}_{4}^{-}\)

Assign the oxidation numbers of all atoms in the following ions. (a) \(\mathrm{NO}_{3}^{-}\) (b) \(\mathrm{NO}_{2}^{-}\) (c) \(\mathrm{NH}_{4}^{+}\)

Essign the oxidation numbers of all atoms in the following species. (a) \(\mathrm{Br}_{2}\) (b) \(\mathrm{CO}_{3}^{2-}\) (c) \(\mathrm{CO}_{2}\)

Assign the oxidation numbers of all atoms in the following compounds. (a) \(\mathrm{ZrO}_{2}\) (b) \(\mathrm{FeO}\) (c) \(\mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}\)

Assign the oxidation numbers of all atoms in the following species. (a) \(\mathrm{PF}_{5}\) (b) \(\mathrm{Na}_{2} \mathrm{CrO}_{4}\) (c) \(\mathrm{NO}_{2}^{-}\)

Assign the oxidation numbers of all atoms in the following species. (a) \(\mathrm{BaO}_{2}\) (b) \(\mathrm{F}_{2}\) (c) \(\mathrm{Sn}^{2+}\)

Assign the oxidation numbers of all atoms in the following species. (a) \(\mathrm{KMnO}_{4}\) (b) \(\mathrm{H}_{2} \mathrm{O}\) (c) \(\mathrm{Cl}_{2}\)

Assign the oxidation numbers of all atoms in the following species. (a) \(\mathrm{NO}_{2}\) (b) \(\mathrm{CrO}_{2}^{-}\) (c) \(\mathrm{Co}\left(\mathrm{NO}_{3}\right)_{3}\)

Assign the oxidation numbers of all atoms in the following species. (a) \(\mathrm{CaCO}_{3}\) (b) \(\mathrm{HBrO}_{4}\) (c) \(\mathrm{Fe}^{3+}\)

Assign the oxidation numbers of all atoms in the following compounds. (a) \(\mathrm{KHF}_{2}\) (b) \(\mathrm{H}_{2} \mathrm{Se}\) (c) \(\mathrm{NaO}_{2}\) (d) \(\mathrm{C}_{2} \mathrm{H}_{6}\)

Assign the oxidation numbers of all atoms in the following species. (a) \(\mathrm{NO}\) (b) \(\mathrm{BO}_{2}^{-}\) (c) \(\mathrm{Cr}\left(\mathrm{NO}_{3}\right)_{3}\) (d) \(\mathrm{CH}_{3} \mathrm{OH}\)

Balance the following reactions, and specify which species is oxidized and which is reduced. (a) \(\mathrm{Fe}_{2} \mathrm{O}_{3}+\mathrm{H}_{2} \rightarrow \mathrm{Fe}+\mathrm{H}_{2} \mathrm{O}\) (b) \(\mathrm{CuCl}_{2}+\mathrm{Na} \rightarrow \mathrm{NaCl}+\mathrm{Cu}\) (c) \(\mathrm{C}+\mathrm{O}_{2} \rightarrow \mathrm{CO}_{2}\)

Balance the following reactions, and specify which species is oxidized and which is reduced. (a) \(\mathrm{Na}+\mathrm{FeCl}_{3} \rightarrow \mathrm{Fe}+\mathrm{NaCl}\) (b) \(\mathrm{SnCl}_{2}+\mathrm{FeCl}_{3} \rightarrow \mathrm{SnCl}_{4}+\mathrm{FeCl}_{2}\) (c) \(\mathrm{CO}+\mathrm{Cr}_{2} \mathrm{O}_{3} \rightarrow \mathrm{Cr}+\mathrm{CO}_{2}\)

Balance the following reactions, and specify which species is oxidized and which is reduced. (a) \(\mathrm{Na}+\mathrm{Hg}_{2} \mathrm{Cl}_{2} \rightarrow \mathrm{NaCl}+\mathrm{Hg}\) (b) \(\mathrm{HCl}+\mathrm{Zn} \rightarrow \mathrm{ZnCl}_{2}+\mathrm{H}_{2}\) (c) \(\mathrm{H}_{2}+\mathrm{CO}_{2} \rightarrow \mathrm{CO}+\mathrm{H}_{2} \mathrm{O}\)

Complete and balance each half-reaction in acid solution, and identify it as an oxidation or a reduction. (a) \(\mathrm{Cr}^{3+}(\mathrm{aq}) \rightarrow \mathrm{Cr}(\mathrm{s})\) (b) \(\mathrm{I}^{-}(\mathrm{aq}) \rightarrow \mathrm{I}_{2}(\mathrm{aq})\) (c) \(\mathrm{NO}_{2}^{-}(\mathrm{aq}) \rightarrow \mathrm{NO}_{3}^{-}(\mathrm{aq})\)

Write balanced equations for the following half reactions. Specify whether each is an oxidation or reduction. (a) \(\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq}) \rightarrow \mathrm{O}_{2}(\mathrm{~g})\) (b) \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}(\mathrm{aq}) \rightarrow \mathrm{CO}_{2}(\mathrm{~g})\) (c) \(\mathrm{NO}_{3}^{-}(\mathrm{aq}) \rightarrow \mathrm{NO}(\mathrm{g})\)

Complete and balance each half-reaction in acid solution, and identify it as an oxidation or a reduction. (a) \(\mathrm{UO}_{2}^{2+}(\mathrm{aq}) \rightarrow \mathrm{U}^{4+}(\mathrm{aq})\) (b) \(\mathrm{Zn}(\mathrm{s}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})\) (c) \(\mathrm{IO}_{3}^{-}(\mathrm{aq}) \rightarrow \mathrm{I}^{-}(\mathrm{aq})\)

Complete and balance each half-reaction in acid solution, and identify it as an oxidation or a reduction. (a) \(\mathrm{N}_{2} \mathrm{O}_{4}(\mathrm{~g}) \rightarrow \mathrm{NO}_{3}^{-}(\mathrm{aq})\) (b) \(\mathrm{Mn}^{3+}(\mathrm{aq}) \rightarrow \mathrm{MnO}_{4}^{-}(\mathrm{aq})\) (c) \(\mathrm{HOCl}(\mathrm{aq}) \rightarrow \mathrm{ClO}_{3}^{-}(\mathrm{aq})\)

Balance each of the following redox reactions in acid solution. (a) \(\mathrm{Sn}(\mathrm{s})+\mathrm{Fe}^{3+}(\mathrm{aq}) \rightarrow \mathrm{Sn}^{2+}(\mathrm{aq})+\mathrm{Fe}^{2+}(\mathrm{aq})\) (b) \(\mathrm{HAsO}_{3}^{2-}(\mathrm{aq})+\mathrm{I}_{2}(\mathrm{aq}) \rightarrow \mathrm{H}_{2} \mathrm{AsO}_{4}^{-}(\mathrm{aq})+\mathrm{I}^{-}(\mathrm{aq})\) (c) \(\mathrm{Cu}(\mathrm{s})+\mathrm{Ag}^{+}(\mathrm{aq}) \rightarrow \mathrm{Cu}^{2+}(\mathrm{aq})+\mathrm{Ag}(\mathrm{s})\)

Balance each of the following redox reactions in acid solution. (a) \(\mathrm{MnO}_{4}^{-}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}(\mathrm{aq}) \rightarrow \mathrm{Mn}^{2+}(\mathrm{aq})+\mathrm{CO}_{2}(\mathrm{~g})\) (b) \(\mathrm{Cl}_{2}(\mathrm{~g})+\mathrm{Br}^{-}(\mathrm{aq}) \rightarrow \mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{Br}_{2}(\ell)\) (c) \(\mathrm{Cu}(\mathrm{s})+\mathrm{NO}_{3}^{-}(\mathrm{aq}) \rightarrow \mathrm{NO}(\mathrm{g})+\mathrm{Cu}^{2+}(\mathrm{aq})\)

Balance each of the following redox reactions in acid solution. (a) \(\mathrm{Fe}(\mathrm{s})+\mathrm{Ag}^{+}(\mathrm{aq}) \rightarrow \mathrm{Ag}(\mathrm{s})+\mathrm{Fe}^{2+}(\mathrm{aq})\) (b) \(\mathrm{I}_{2}(\mathrm{aq})+\mathrm{S}_{2} \mathrm{O}_{3}^{2-}(\mathrm{aq}) \rightarrow \mathrm{I}^{-}(\mathrm{aq})+\mathrm{S}_{4} \mathrm{O}_{6}^{2-}(\mathrm{aq})\) (c) \(\mathrm{MnO}_{4}^{-}(\mathrm{aq})+\mathrm{Fe}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{Mn}^{2+}(\mathrm{aq})\)

Balance each of the following redox reactions in acid solution. (a) \(\mathrm{Zn}(\mathrm{s})+\mathrm{NO}_{3}^{-}(\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+\mathrm{N}_{2}(\mathrm{~g})\) (b) \(\mathrm{IO}_{3}^{-}(\mathrm{aq})+\mathrm{I}^{-}(\mathrm{aq}) \rightarrow \mathrm{I}_{2}(\mathrm{aq})\) (c) \(\mathrm{Ce}^{4+}(\mathrm{aq})+\mathrm{Cl}^{-}(\mathrm{aq}) \rightarrow \mathrm{Cl}_{2}(\mathrm{aq})+\mathrm{Ce}^{3+}(\mathrm{aq})\)

Balance each of the following redox reactions in basic solution. (a) \(\mathrm{Al}(\mathrm{s})+\mathrm{ClO}^{-}(\mathrm{aq}) \rightarrow \mathrm{Al}(\mathrm{OH})_{4}^{-}(\mathrm{aq})+\mathrm{Cl}^{-}(\mathrm{aq})\) (b) \(\mathrm{MnO}_{4}^{-}(\mathrm{aq})+\mathrm{SO}_{3}^{2-}(\mathrm{aq}) \rightarrow \mathrm{MnO}_{2}(\mathrm{~s})+\mathrm{SO}_{4}^{2-}(\mathrm{aq})\) (c) \(\mathrm{Zn}(\mathrm{s})+\mathrm{NO}_{3}(\mathrm{aq}) \rightarrow \mathrm{Zn}(\mathrm{OH})_{4}^{2-}(\mathrm{aq})+\mathrm{NH}_{3}(\mathrm{aq})\)

Balance each of the following redox reactions in basic solution. (a) \(\mathrm{ClO}^{-}(\mathrm{aq})+\mathrm{CrO}_{2}^{-}(\mathrm{aq}) \rightarrow \mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{CrO}_{4}^{2-}(\mathrm{aq})\) (b) \(\mathrm{Br}_{2}(\mathrm{aq}) \rightarrow \mathrm{Br}^{-}(\mathrm{aq})+\mathrm{BrO}_{3}^{-}(\mathrm{aq})\) (c) \(\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq})+\mathrm{N}_{2} \mathrm{H}_{4}(\mathrm{aq}) \rightarrow \mathrm{N}_{2}(\mathrm{~g})+\mathrm{H}_{2} \mathrm{O}(\ell)\)

Balance each of the following redox reactions in basic solution. (a) \(\mathrm{Cl}_{2}(\mathrm{aq}) \rightarrow \mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{ClO}_{3}^{-}(\mathrm{aq})\) (b) \(\cdot \mathrm{MnO}_{4}^{-}(\mathrm{aq})+\mathrm{I}^{-}(\mathrm{aq}) \rightarrow \mathrm{IO}_{3}^{-}(\mathrm{aq})+\mathrm{MnO}_{2}(\mathrm{~s})\) (c) \(\mathrm{ClO}_{3}^{-}(\mathrm{aq})+\mathrm{CN}^{-}(\mathrm{aq}) \rightarrow \mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{CNO}^{-}(\mathrm{aq})\)

Balance each of the following redox reactions in basic solution. (a) \(\mathrm{PH}_{3}(\mathrm{~g})+\mathrm{CrO}_{4}^{2-}(\mathrm{aq}) \rightarrow \mathrm{CrO}_{2}^{-}(\mathrm{aq})+\mathrm{P}_{4}(\mathrm{~s})\) (b) \(\mathrm{F}_{2}(\mathrm{~g})+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{F}^{-}(\mathrm{aq})+\mathrm{O}_{2}(\mathrm{~g})\) (c) \(\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq})+\mathrm{Cr}(\mathrm{OH})_{3}(\mathrm{~s}) \rightarrow \mathrm{CrO}_{4}^{2-}(\mathrm{aq})\)

Why is the following balanced reaction not a proper redox reaction? $$ \mathrm{Fe}^{2+}(\mathrm{aq})+2 \mathrm{Br}^{-}(\mathrm{aq}) \rightarrow \mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{Br}_{2}(\ell) $$

A voltaic cell is based on the reaction $$ \mathrm{Pb}(\mathrm{s})+2 \mathrm{Ag}^{+}(\mathrm{aq}) \rightarrow \mathrm{Pb}^{2+}(\mathrm{aq})+2 \mathrm{Ag}(\mathrm{s}) $$ Voltage measurements show that the Ag electrode is positive. Sketch the cell, and label the anode and cathode, the positive and negative electrodes, the direction of electron flow in the external circuit, and the direction of flow of cations and anions through the salt bridge. Write the halfreaction that occurs at each electrode.

A voltaic cell is based on the reaction $$ \mathrm{Zn}(\mathrm{s})+\mathrm{Ni}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+\mathrm{Ni}(\mathrm{s}) $$ Voltage measurements show that the Ni electrode is positive. Sketch the cell, and label the anode and cathode, the positive and negative electrodes, the direction of electron flow in the external circuit, and the direction of flow of cations and anions through the salt bridge. Write the halfreaction that occurs at each electrode.

A platinum wire is in contact with a mixture of mercury and solid mercury(I) chloride \(\left(\mathrm{Hg}_{2} \mathrm{Cl}_{2}\right)\) in a beaker containing \(1 M\) KCl solution. A salt bridge connects this halfcell to a beaker that contains a copper electrode immersed in \(1 \mathrm{M} \mathrm{CuSO}_{4}\) solution. Voltage measurements show that the copper electrode is positive. (a) Write balanced half-reactions for the two electrodes. (b) Write the equation for the spontaneous cell reaction. (c) In which direction do electrons flow in the external circuit? (d) Would direct reaction occur if both the \(\mathrm{Hg} / \mathrm{Hg}_{2} \mathrm{Cl}_{2}\) and copper electrodes were placed in a container holding an aqueous solution that is \(1 \mathrm{M} \mathrm{CuSO}_{4}\) and \(1 \mathrm{M} \mathrm{KCl}^{2}\)

Two electrodes are immersed in a \(1 M \mathrm{HBr}\) solution. One of the electrodes is a silver wire coated with a deposit of \(\mathrm{AgBr}(\mathrm{s})\). The second electrode is a platinum wire in electrical contact with a mixture of metallic mercury and \(\mathrm{Hg}_{2} \mathrm{Br}_{2}(\mathrm{~s}) .\) Voltage measurements show that the \(\mathrm{Pt}\) electrode is positive. (a) Write balanced half-reactions for the two electrodes. (b) Write the equation for the spontaneous cell reaction. (c) In which direction do electrons flow in the external circuit? (d) Why is a salt bridge unnecessary in this cell?

A half-cell that consists of a silver wire in a \(1.00 M\) \(\mathrm{AgNO}_{3}\) solution is connected by a salt bridge to a \(1.00 \mathrm{M}\) thallium(I) acetate solution that contains a metallic T1 electrode. The voltage of the cell is \(1.136 \mathrm{~V}\), with the silver as the positive electrode. (a) Write the half-reactions and the overall chemical equation for the spontaneous reaction. (b) Use the standard potential of the silver half-reaction, with the voltage of the cell, to calculate the standard reduction potential for the thallium half-reaction.

The standard potential of the half-reaction $$ 2 \mathrm{D}^{+}(\mathrm{aq})+2 \mathrm{e}^{-} \rightarrow \mathrm{D}_{2}(\mathrm{~g}) $$ (where \(\mathrm{D}=\) deuterium, or \(\left.{ }^{2} \mathrm{H}\right)\) is \(-0.013 \mathrm{~V}\). Determine \(\Delta G^{\circ}\) and \(K_{\mathrm{eq}}\) for the reaction $$ \mathrm{H}_{2}(\mathrm{~g})+2 \mathrm{D}^{+}(\mathrm{aq}) \rightarrow 2 \mathrm{H}^{+}(\mathrm{aq})+\mathrm{D}_{2}(\mathrm{~g}) $$ In a mixture of hydrogen and deuterium, which isotope more favors its elemental form under standard conditions?

Disproportionation is a type of redox reaction in which the same species is simultaneously oxidized and reduced. One species that undergoes disproportionation is \(\mathrm{Cu}^{+}(\) aq \()\). $$ 2 \mathrm{Cu}^{+}(\mathrm{aq}) \rightarrow \mathrm{Cu}(\mathrm{s})+\mathrm{Cu}^{2+}(\mathrm{aq}) $$ If the half-reactions are $$ \begin{array}{ll} \mathrm{Cu}^{2+}(\mathrm{aq})+\mathrm{e}^{-} \rightarrow \mathrm{Cu}^{+}(\text {aq }) & E^{\circ}=0.153 \mathrm{~V} \\ \mathrm{Cu}^{+}(\mathrm{aq})+\mathrm{e}^{-} \rightarrow \mathrm{Cu}(\mathrm{s}) & E^{\circ}=0.521 \mathrm{~V} \end{array} $$ what are \(E^{\circ}, \Delta G^{\circ},\) and \(K_{\mathrm{eq}}\) for the overall reaction?

A voltaic cell consists of a lead electrode and a reference electrode with a constant potential. This cell has a voltage of \(53 \mathrm{mV}\) when the lead electrode is placed in a \(0.100 \mathrm{M}\) \(\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}\) solution (the lead electrode is positive). What voltage is measured when the lead electrode is placed in a saturated lead chloride solution, in which \(\left[\mathrm{Pb}^{2+}\right]\) is \(1.6 \times 10^{-2} M ?\)

Calculate the value of the solubility product constant for \(\mathrm{Cd}(\mathrm{OH})_{2}\) from the half-cell potentials. $$ \begin{array}{lr} \mathrm{Cd}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \rightarrow \mathrm{Cd}(\mathrm{s}) & E^{\circ}=-0.403 \mathrm{~V} \\ \mathrm{Cd}(\mathrm{OH})_{2}(\mathrm{~s})+2 \mathrm{e}^{-} \rightarrow \mathrm{Cd}(\mathrm{s})+2 \mathrm{OH}^{-}(\mathrm{aq}) & \\ E^{\circ} & =-0.83 \mathrm{~V} \end{array} $$

Calculate the value of the solubility product constant for \(\mathrm{PbSO}_{4}\) from the half-cell potentials. $$ \begin{aligned} \mathrm{PbSO}_{4}(\mathrm{~s})+2 \mathrm{e}^{-} \rightarrow \mathrm{Pb}(\mathrm{s})+\mathrm{SO}_{4}^{2-}(\mathrm{aq}) & E^{\circ}=-0.356 \mathrm{~V} \\ \mathrm{~Pb}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \rightarrow \mathrm{Pb}(\mathrm{s}) & E^{\circ}=-0.126 \mathrm{~V} \end{aligned} $$

What is the voltage of a concentration cell of \(\mathrm{Fe}^{2+}\) ions where the concentrations are 0.0025 and \(0.750 M ?\) What is the spontaneous reaction?

What is the voltage of a concentration cell of \(\mathrm{Cl}^{-}\) ions where the concentrations are 1.045 and \(0.085 \mathrm{M}\) ? What is the spontaneous reaction?

Write the half-reactions and the balanced chemical equations for the reactions that occur in the electrolysis of (a) a zinc chloride aqueous solution, using zinc electrodes. (b) a calcium bromide solution, using inert electrodes. (c) a sodium iodide solution, using inert electrodes.

A solution contains the ions \(\mathrm{H}^{+}, \mathrm{Ag}^{+}, \mathrm{Pb}^{2+},\) and \(\mathrm{Ba}^{2+},\) each at a concentration of \(1.0 \mathrm{M}\). (a) Which of these ions would be reduced first at the cathode during an electrolysis? (b) After the first ion has been completely removed by electrolysis, which is the second ion to be reduced? (c) Which, if any, of these ions cannot be reduced by the electrolysis of the aqueous solution?

A solution contains the ions \(\mathrm{H}^{+}, \mathrm{Cu}^{2+}, \mathrm{Ca}^{2+},\) and \(\mathrm{Ni}^{2+},\) each at a concentration of \(1.0 \mathrm{M}\). (a) Which of these ions would be reduced first at the cathode during an electrolysis? (b) After the first ion has been completely removed by electrolysis, which is the second ion to be reduced? (c) Which, if any, of these ions cannot be reduced by the electrolysis of the aqueous solution?

The commercial production of magnesium is accomplished by electrolysis of molten \(\mathrm{MgCl}_{2}\) (a) Why is electrolysis of an aqueous solution of \(\mathrm{MgCl}_{2}\) not used in this process? (b) Write the anode and cathode half-reaction in the electrolysis of molten \(\mathrm{MgCl}_{2}\)