Chapter 17

Make a drawing showing the principal parts of (a) a voltaic cell: show the anode, the cathode, the direction of electron movement outside the cell, and the direction of ion movement inside the cell. (b) a standard hydrogen electrode: describe the components of the electrode and explain how it works.

Explain how product-favored electrochemical reactions can be used to do useful work.

Explain how reactant-favored electrochemical reactions can be induced to make products.

Identify each statement as true or false. Rewrite each false statement to make it true. (a) Oxidation always occurs at the anode of an electrochemical cell. (b) The anode of a discharging voltaic cell is the site of reduction and is negative. (c) Standard-state conditions for electrochemical cells are a concentration of \(1.0 \mathrm{M}\) for dissolved species and a pressure of 1 bar for gases. (d) The potential of a voltaic cell does not change with temperature. (e) All product-favored oxidation-reduction reactions have a standard cell potential \(E_{\text {cell }}^{\circ}\), with a negative sign.

In this reaction, assign an oxidation number to each atom in reactants and products. Identify which substance is oxidized and which is reduced. Identify the oxidizing agent and the reducing agent. $$ \begin{aligned} 8 \mathrm{H}^{+}(\mathrm{aq})+\mathrm{MnO}_{4}^{-}(\mathrm{aq})+& 5 \mathrm{Fe}^{2+}(\mathrm{aq}) \longrightarrow \\ & 5 \mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{Mn}^{2+}(\mathrm{aq})+4 \mathrm{H}_{2} \mathrm{O}(\ell) \end{aligned} $$

In each of these reactions assign an oxidation number to each atom in reactants and products. Identify which substance is oxidized and which is reduced. Identify the oxidizing agent and the reducing agent. (a) \(\mathrm{Fe}(\mathrm{s})+\mathrm{Br}_{2}(\ell) \longrightarrow \mathrm{FeBr}_{2}(\mathrm{~s})\) (b) \(2 \mathrm{Al}(\mathrm{s})+3 \mathrm{Cl}_{2}(\mathrm{~g}) \longrightarrow 2 \mathrm{AlCl}_{3}(\mathrm{~s})\) (c) \(8 \mathrm{HI}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq}) \longrightarrow\) $$ \mathrm{H}_{2} \mathrm{~S}(\mathrm{aq})+4 \mathrm{I}_{2}(\mathrm{~s})+4 \mathrm{H}_{2} \mathrm{O}(\ell) $$

Write half-reactions for these changes: (a) Oxidation of cadmium to \(\mathrm{Cd}^{2+}\) ions (b) Reduction of \(\mathrm{Fe}^{3+}\) ions to Fe metal (c) Reduction of \(\mathrm{Sn}^{4+}\) ions to \(\mathrm{Sn}^{2+}\) ions (d) Reduction of chlorine to \(\mathrm{Cl}^{-}\) ions (e) Oxidation of sulfur dioxide to sulfate ions in acidic solution

Write half-reactions for these changes: (a) Reduction of \(\mathrm{MnO}_{4}^{-}\) ion to \(\mathrm{Mn}^{2+}\) ion in acid solution (b) Reduction of \(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}\) ion to \(\mathrm{Cr}^{3+}\) ion in acid solution (c) Oxidation of chlorine gas to \(\mathrm{ClO}^{-}\) ions (d) Reduction of hydrogen peroxide to water in acidic solution (e) Oxidation of nitrous acid to nitrate ions in acidic solution

Balance these redox reactions, and identify the oxidizing agent and the reducing agent. (a) \(\mathrm{CO}(\mathrm{g})+\mathrm{O}_{3}(\mathrm{~g}) \longrightarrow \mathrm{CO}_{2}(\mathrm{~g})\) (b) \(\mathrm{H}_{2}(\mathrm{~g})+\mathrm{Cl}_{2}(\mathrm{~g}) \longrightarrow \mathrm{HCl}(\mathrm{g})\) (c) \(\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq})+\mathrm{Ti}^{2+}(\mathrm{aq}) \longrightarrow \mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{TiO}_{2}(\mathrm{~s})\) in acidic solution (d) \(\mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{MnO}_{4}^{-}(\mathrm{aq}) \longrightarrow \mathrm{Cl}_{2}(\mathrm{~g})+\mathrm{MnO}_{2}(\mathrm{~s})\) in acidic solution (e) \(\mathrm{FeS}_{2}(\mathrm{~s})+\mathrm{O}_{2}(\mathrm{~g}) \longrightarrow \mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{~s})+\mathrm{SO}_{2}(\mathrm{~g})\) (f) \(\mathrm{O}_{3}(\mathrm{~g})+\mathrm{NO}(\mathrm{g}) \longrightarrow \mathrm{O}_{2}(\mathrm{~g})+\mathrm{NO}_{2}(\mathrm{~g})\) (g) \(\mathrm{Zn}(\mathrm{s})+\mathrm{HgO}(\mathrm{s}) \longrightarrow \mathrm{Zn}(\mathrm{OH})_{2}(\mathrm{~s})+\operatorname{Hg}(\ell)\) in basic solution

Balance these redox reactions, and identify the oxidizing agent and the reducing agent. (a) \(\mathrm{FeO}(\mathrm{s})+\mathrm{O}_{3}(\mathrm{~g}) \longrightarrow \mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{~s})\) (b) \(\mathrm{P}_{4}(\mathrm{~s})+\mathrm{Br}_{2}(\ell) \longrightarrow \mathrm{PBr}_{5}(\ell)\) (c) \(\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq})+\mathrm{Co}^{2+}(\mathrm{aq}) \longrightarrow \mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{Co}^{3+}(\mathrm{aq})\) in acidic solution (d) \(\mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(\mathrm{aq}) \longrightarrow \mathrm{Cl}_{2}(\mathrm{~g})+\mathrm{Cr}^{3+}(\mathrm{aq})\) in acidic solution (e) \(\mathrm{MnO}_{4}^{-}(\mathrm{aq})+\mathrm{Zn}(\mathrm{s}) \longrightarrow \mathrm{MnO}_{2}(\mathrm{~s})+\mathrm{Zn}(\mathrm{OH})_{2}(\mathrm{~s})\) in basic solution (f) \(\mathrm{H}_{2} \mathrm{CO}(\mathrm{g})+\mathrm{O}_{2}(\mathrm{~g}) \longrightarrow \mathrm{CO}_{2}(\mathrm{~g})+\mathrm{H}_{2} \mathrm{O}(\ell)\) (g) \(\mathrm{C}_{3} \mathrm{H}_{8}(\mathrm{~g})+\mathrm{O}_{2}(\mathrm{~g}) \longrightarrow \mathrm{CO}_{2}(\mathrm{~g})+\mathrm{H}_{2} \mathrm{O}(\ell)\)

For the reaction \(\mathrm{Cu}^{2+}(\mathrm{aq})+\mathrm{Zn}(\mathrm{s}) \longrightarrow \mathrm{Cu}(\mathrm{s})+\mathrm{Zn}^{2+}(\mathrm{aq})\) why can't you generate electric current by placing a piece of copper metal and a piece of zinc metal in a solution containing \(\mathrm{CuCl}_{2}(\mathrm{aq})\) and \(\mathrm{ZnCl}_{2}(\mathrm{aq})\) ?

Explain the function of a salt bridge in a voltaic cell.

Tell whether this statement is true or false. If false, rewrite it to make it a correct statement: The value of an electrode potential changes when the half-reaction is multiplied by a factor. That is, \(E^{\circ}\) for \(\mathrm{Li}^{+}+\mathrm{e}^{-} \longrightarrow \mathrm{Li}\) is different from that for \(2 \mathrm{Li}^{+}+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{Li}\).

A voltaic cell is assembled with \(\mathrm{Sn}(\mathrm{s})\) and \(\mathrm{Sn}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})\) in one compartment and \(\mathrm{Ag}(\mathrm{s})\) and \(\mathrm{AgNO}_{3}(\mathrm{aq})\) in the other. An external wire connects the two electrodes, and a salt bridge containing \(\mathrm{KNO}_{3}(\) aq) connects the two solutions. (a) In the product-favored reaction, \(\mathrm{Ag}^{+}\) is reduced to silver metal. Write a balanced net ionic equation for this reaction. (b) Which half-reaction occurs at each electrode? Which is the anode and which is the cathode? (c) Draw a diagram of the cell, indicating the direction of electron movement outside the cell and of ion movement within the cell.

Draw a diagram of each cell. Label the anode, the cathode, the species in each half-cell solution, the direction of electron movement in an external circuit, and the direction of movement of ions within the cell. (a) \(\mathrm{Cu}(\mathrm{s})\left|\mathrm{Cu}^{2+}(\mathrm{aq}) \| \mathrm{Fe}^{2+}(\mathrm{aq})\right| \mathrm{Fe}(\mathrm{s})\) (b) \(\mathrm{Pt}(\mathrm{s})\left|\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq}), \mathrm{H}^{+}(\mathrm{aq}) \| \mathrm{Fe}^{2+}(\mathrm{aq}), \mathrm{Fe}^{3+}(\mathrm{aq})\right| \mathrm{Pt}(\mathrm{s})\)

You light a \(25-W\) light bulb with the current from a \(12-V\) lead-acid storage battery. Calculate how much energy the light bulb utilized after \(1.0 \mathrm{~h}\) of operation. Calculate how many coulombs passed through the bulb. Assume \(100 \%\) efficiency. \((1 \mathrm{~W}=1 \mathrm{~J} / \mathrm{s} .)\)

Copper can reduce silver ion to metallic silver, a reaction that could, in principle, be used in a battery. $$ \mathrm{Cu}(\mathrm{s})+2 \mathrm{Ag}^{+}(\mathrm{aq}) \longrightarrow \mathrm{Cu}^{2+}(\mathrm{aq})+2 \mathrm{Ag}(\mathrm{s}) $$ (a) Write equations for the half-reactions involved. (b) Which half-reaction is an oxidation and which is a reduction? Which half- reaction occurs in the anode compartment and which takes place in the cathode compartment?

Chlorine gas can oxidize zinc metal in a reaction that has been suggested as the basis of a battery. Write the half reactions involved. Label which is the oxidation half reaction and which is the reduction half-reaction.

Consider these half-reactions: $$ \begin{array}{lr} \hline \text { Half-reaction } & E^{\circ}(\mathrm{V}) \\ \hline \mathrm{Au}^{3+}(\mathrm{aq})+3 \mathrm{e}^{-} \longrightarrow \mathrm{Au}(\mathrm{s}) & 1.52 \\ \mathrm{Pt}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \longrightarrow \mathrm{Pt}(\mathrm{s}) & 1.118 \\ \mathrm{Co}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \longrightarrow \mathrm{Co}(\mathrm{s}) & -0.277 \\ \mathrm{Mn}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \longrightarrow \mathrm{Mn}(\mathrm{s}) & -1.18 \\ \hline \end{array} $$ (a) Which is the weakest oxidizing agent? (b) Which is the strongest oxidizing agent? (c) Which is the strongest reducing agent? (d) Which is the weakest reducing agent? (e) Will \(\mathrm{Co}(\mathrm{s})\) reduce \(\mathrm{Pt}^{2+}(\mathrm{aq})\) to \(\mathrm{Pt}(\mathrm{s})\) ? (f) Will \(\mathrm{Pt}(\mathrm{s})\) reduce \(\mathrm{Co}^{2+}(\mathrm{aq})\) to \(\mathrm{Co}(\mathrm{s})\) ? (g) Which ions can be reduced by Co(s)?

Consider these half-reactions: $$ \begin{array}{lr} \hline \text { Half-reaction } & E^{\circ}(\mathrm{V}) \\ \hline \mathrm{Ce}^{4+}(\mathrm{aq})+\mathrm{e}^{-} \longrightarrow \mathrm{Ce}^{3+}(\mathrm{aq}) & 1.72 \\ \mathrm{Ag}^{+}(\mathrm{aq})+\mathrm{e}^{-} \longrightarrow \mathrm{Ag}(\mathrm{s}) & 0.80 \\ \mathrm{Hg}_{2}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{Hg}(\ell) & 0.80 \\ \mathrm{Sn}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \longrightarrow \mathrm{Sn}(\mathrm{s}) & -0.14 \\ \mathrm{Ni}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \longrightarrow \mathrm{Ni}(\mathrm{s}) & -0.25 \\ \mathrm{Al}^{3+}(\mathrm{aq})+3 \mathrm{e}^{-} \longrightarrow \mathrm{Al}(\mathrm{s}) & -1.68 \\ \hline \end{array} $$ (a) Which is the weakest oxidizing agent? (b) Which is the strongest oxidizing agent? (c) Which is the strongest reducing agent? (d) Which is the weakest reducing agent? (e) Will \(\mathrm{Sn}(\mathrm{s})\) reduce \(\mathrm{Ag}^{+}(\mathrm{aq})\) to \(\mathrm{Ag}(\mathrm{s}) ?\) (f) Will \(\mathrm{Hg}(\ell)\) reduce \(\mathrm{Sn}^{2+}(\mathrm{aq})\) to \(\mathrm{Sn}(\mathrm{s}) ?\) (g) Name the ions that can be reduced by \(\operatorname{Sn}(\mathrm{s})\). (h) Which metals can be oxidized by \(\mathrm{Ag}^{+}(\mathrm{aq}) ?\)

In principle, a battery could be made from aluminum metal and chlorine gas. (a) Write a balanced equation for the reaction that would occur in a battery using \(\mathrm{Al}^{3+}(\mathrm{aq}) \mid \mathrm{Al}(\mathrm{s})\) and \(\mathrm{Cl}_{2}(\mathrm{~g}) \mid \mathrm{Cl}^{-}(\) aq \()\) half-cells. (b) Identify the half-reaction at the anode and at the cathode. Do electrons flow from the \(\mathrm{Al}\) electrode when the cell does work? Explain. (c) Calculate the standard potential, \(E_{\text {cell }}^{\circ}\), for the battery.

Choose the correct answers: In a product-favored chemical reaction, the standard cell potential, \(E_{\text {cell }}^{\circ}\), is (greater/less) than zero, and the Gibbs free energy change, \(\Delta_{\mathrm{r}} G^{\circ},\) is (greater/less) than zero.

Hydrazine, \(\mathrm{N}_{2} \mathrm{H}_{4},\) can be used as the reducing agent in a fuel cell. $$ \mathrm{N}_{2} \mathrm{H}_{4}(\mathrm{aq})+\mathrm{O}_{2}(\mathrm{aq}) \longrightarrow \mathrm{N}_{2}(\mathrm{~g})+2 \mathrm{H}_{2} \mathrm{O}(\ell) $$ (a) If \(\Delta_{t} G^{\circ}\) for the reaction is \(-598 \mathrm{~kJ},\) calculate the value of \(E^{\circ}\) expected for the reaction. (b) Suppose the equation is written with all coefficients doubled. Determine \(\Delta_{\mathrm{r}} G^{\circ}\) and \(E^{\circ}\) for this new reaction.

The standard cell potential for the oxidation of \(\mathrm{Mg}\) by $$ \begin{aligned} \mathrm{Br}_{2} \text { is } 3.42 \mathrm{~V} \\ & \mathrm{Br}_{2}(\ell)+\mathrm{Mg}(\mathrm{s}) \longrightarrow \mathrm{Mg}^{2+}(\mathrm{aq})+2 \mathrm{Br}^{-}(\mathrm{aq}) \end{aligned} $$ (a) Calculate \(\Delta_{\mathrm{r}} G^{\circ}\) for this reaction. (b) Suppose the equation is written with all coefficients doubled. Determine \(\Delta_{\mathrm{r}} G^{\circ}\) and \(E^{\circ}\) for this new equation.

The standard cell potential, \(E^{\circ},\) for the reaction of \(\mathrm{Zn}(\mathrm{s})\) and \(\mathrm{Cl}_{2}(\mathrm{~g})\) is \(2.12 \mathrm{~V}\). Write the chemical equation for the reaction of \(1 \mathrm{~mol}\) zinc. Calculate the standard Gibbs free energy change, \(\Delta_{t} G^{\circ},\) for this reaction.

Consider a voltaic cell with the reaction given below. As the cell reaction proceeds, what happens to the values of \(E_{\text {cell }}, \Delta_{r} G,\) and \(K_{\mathrm{c}} ?\) Explain your answers. $$ \begin{array}{r} \mathrm{Cu}^{2+}(\mathrm{aq}, 1 \mathrm{M})+\mathrm{Zn}(\mathrm{s}) \longrightarrow \mathrm{Cu}(\mathrm{s})+\mathrm{Zn}^{2+}(\mathrm{aq}, 1 \mathrm{M}) \\ E_{\mathrm{cell}}^{\circ}=1.10 \mathrm{~V} \end{array} $$

Calculate the cell potential of a concentration cell that contains two hydrogen electrodes if the cathode contacts a solution with \(\mathrm{pH}=7.8\) and the anode contacts a solution with (conc. \(\left.\mathrm{H}^{+}\right)=0.05 \mathrm{M}\).

Calculate the potential of a cell with one electrode made from zinc metal immersed in a solution where (conc. \(\left.\mathrm{Zn}^{2+}\right)=0.010 \mathrm{M}\) and the other electrode is a standard hydrogen electrode.

Describe the advantages and disadvantages of lead-acid storage batteries.

Describe the advantages and disadvantages of Li-ion batteries.

How does a fuel cell differ from a battery?

Describe the principal parts of an \(\mathrm{H}_{2} \mid \mathrm{O}_{2}\) fuel cell. Write a balanced equation for the reaction at the cathode; at the anode. Give the formula of the product of the fuel cell reaction.

Hydrazine, \(\mathrm{N}_{2} \mathrm{H}_{4}\), has been proposed as the fuel in a fuel cell in which oxygen is the oxidizing agent. The reactions are $$ \begin{array}{r} \mathrm{N}_{2} \mathrm{H}_{4}(\mathrm{aq})+4 \mathrm{OH}^{-}(\mathrm{aq}) \longrightarrow \mathrm{N}_{2}(\mathrm{~g})+4 \mathrm{H}_{2} \mathrm{O}(\ell)+4 \mathrm{e}^{-} \\ \mathrm{O}_{2}(\mathrm{~g})+2 \mathrm{H}_{2} \mathrm{O}(\ell)+4 \mathrm{e}^{-} \longrightarrow 4 \mathrm{OH}^{-}(\mathrm{aq}) \end{array} $$ (a) Which reaction occurs at the anode and which at the cathode? (b) What is the overall cell reaction? (c) If the cell is to produce \(0.50 \mathrm{~A}\) of current for \(50.0 \mathrm{~h}, \mathrm{cal}-\) culate what mass in grams of hydrazine must be present. (d) Calculate what mass (g) of \(\mathrm{O}_{2}\) must be available to react with the mass of \(\mathrm{N}_{2} \mathrm{H}_{4}\) determined in part (c).

Consider the electrolysis of water in the presence of very dilute \(\mathrm{H}_{2} \mathrm{SO}_{4}\). What species is produced at the anode? At the cathode? What are the relative amounts of the species produced at the two electrodes?

Write chemical equations for the electrolysis of molten salts of three different alkali halides to produce the corresponding halogens and alkali metals.

Identify the products of the electrolysis of a 1 -M aqueous solution of NaBr. What species are present in the solution? What is formed at the cathode? What is formed at the anode?

For each of these solutions, tell what reactions take place at the anode and at the cathode during electrolysis. (a) \(\mathrm{NiBr}_{2}(\mathrm{aq})\) (b) \(\mathrm{NaI}(\mathrm{aq})\) (c) \(\mathrm{CdCl}_{2}(\mathrm{aq})\) (d) \(\mathrm{CuI}_{2}(\mathrm{aq})\) (e) \(\mathrm{MgF}_{2}(\mathrm{aq})\) (f) \(\mathrm{HNO}_{3}(\mathrm{aq})\)

A current of \(1.0 \mathrm{~mA}\) is passed through a solution containing \(\mathrm{Ag}^{+}(\mathrm{aq}) .\) Calculate the mass of silver in the solution if all the silver was deposited as Ag metal in \(14.5 \mathrm{~min}\).

A current of \(2.50 \mathrm{~A}\) is passed through a solution of \(\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}\) for \(2.00 \mathrm{~h} .\) Calculate the mass of copper deposited at the cathode.

A current of \(0.0125 \mathrm{~A}\) is passed through a solution of \(\mathrm{CuCl}_{2}\) for \(2.00 \mathrm{~h} .\) Calculate the mass of copper deposited at the cathode and the volume of \(\mathrm{Cl}_{2}\) gas (in \(\mathrm{mL}\) at STP) produced at the anode.

The major reduction half-reaction occurring in the cell in which molten \(\mathrm{Al}_{2} \mathrm{O}_{3}\) and molten aluminum salts are electrolyzed is \(\mathrm{Al}^{3+}(\mathrm{aq})+3 \mathrm{e}^{-} \longrightarrow \mathrm{Al}(\mathrm{s})\). The cell operates at \(5.0 \mathrm{~V}\) and \(1.0 \times 10^{5} \mathrm{~A} .\) Calculate the mass \((\mathrm{g})\) of aluminum metal produced in \(8.0 \mathrm{~h}\).

The vanadium(II) ion can be produced by electrolysis of a vanadium(III) salt in solution. Calculate how long you must carry out an electrolysis if you wish to convert completely \(0.125 \mathrm{~L}\) of \(0.0150-\mathrm{M} \mathrm{V}^{3+}(\mathrm{aq})\) to \(\mathrm{V}^{2+}(\mathrm{aq})\) using a current of \(0.268 \mathrm{~A}\).

Assume that the anode reaction for the lithium battery is $$ \mathrm{LiC}_{6}(\mathrm{~s}) \longrightarrow \mathrm{Li}^{+}(\text {electrolyte })+\mathrm{C}_{6}(\mathrm{~s})+\mathrm{e}^{-} $$ and the anode reaction for the lead-acid storage battery is $$ \mathrm{Pb}(\mathrm{s})+\mathrm{HSO}_{4}^{-}(\mathrm{aq}) \longrightarrow \mathrm{PbSO}_{4}(\mathrm{~s})+2 \mathrm{e}^{-}+\mathrm{H}^{+}(\mathrm{aq}) $$ Compare the masses of metals consumed when each of these batteries supplies a current of \(1.0 \mathrm{~A}\) for \(10 . \mathrm{min}\).

A hydrogen-oxygen fuel cell operates on the simple reaction $$ 2 \mathrm{H}_{2}(\mathrm{~g})+\mathrm{O}_{2}(\mathrm{~g}) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}(\ell) $$ If the cell is designed to produce \(1.5 \mathrm{~A}\) of current, determine how long it can operate if there is an excess of oxygen and only sufficient hydrogen to fill a \(1.0-\mathrm{L}\) tank at 200. bar pressure at \(25^{\circ} \mathrm{C}\).

Calculate how long it would take to electroplate a metal surface with \(0.500 \mathrm{~g}\) nickel metal from a solution of \(\mathrm{Ni}^{2+}\) with a current of \(4.00 \mathrm{~A}\).

Explain how rust is formed from iron materials by corrosion.

Why does iron corrode faster in salt water than in fresh water?

Name one common metal that does not corrode readily under normal conditions.

Explain how galvanizing iron stops corrosion of the underlying iron.

Three electrolytic cells are connected in series, so that the same current flows through all of them for \(20 .\) min. In cell A, 0.0234 g Ag plates out from a solution of \(\mathrm{AgNO}_{3}(\mathrm{aq}) ;\) cell \(\mathrm{B}\) contains \(\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq}) ;\) cell \(\mathrm{C}\) contains \(\mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3}(\) aq \() .\) Calculate what mass of \(\mathrm{Cu}\) will plate out in cell B. Calculate what mass of Al will plate out in cell C.