Chapter 18

An element that is a good reducing agent is also ______ a. easily oxidized b. a good oxidizing agent c. easily reduced d. a noble gas

Regarding the porous separator between the two halves of an electrochemical cell: a. Describe how it allows electrical charge to flow between the two half- cells. b. Explain why a piece of wire could not perform the same function.

In the redox reaction below, how many moles of electrons are transferred for each mole of chlorine gas consumed? $$2 \mathrm{Fe}^{2+}(a q)+\mathrm{Cl}_{2}(g) \rightarrow 2 \mathrm{Fe}^{3+}(a q)+2 \mathrm{Cl}^{-}(a q)$$

In the redox reaction below, how many electrons are transferred for each molecule of \(\mathrm{H}_{2} \mathrm{O}_{2}\) consumed? \(2 \mathrm{MnO}_{4}^{-}(a q)+3 \mathrm{H}_{2} \mathrm{O}_{2}(a q) \rightarrow\) $$2 \mathrm{MnO}_{2}(s)+3 \mathrm{O}_{2}(g)+2 \mathrm{OH}^{-}(a q)+2 \mathrm{H}_{2} \mathrm{O}(\ell)$$

Complete and balance the partial chemical equation below, using the appropriate half-reactions in Table \(A 6.1\) for acidic solutions. $$\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{Fe}^{2+}(a q) \rightarrow \mathrm{Cr}^{3+}(a q)+\mathrm{Fe}^{3+}(a q)$$

Select the appropriate half-reactions from Appendix 6 to write net ionic equations describing the reaction between: a. tin and \(A g^{+}\) ions in solution that produces dissolved \(\mathrm{Sn}^{2+}\) ions and silver metal. b. copper and \(\mathrm{O}_{2}\) in an acidic solution that produces \(\mathrm{Cu}^{2+}\) ions. c. solid \(\mathrm{Cr}(\mathrm{OH})_{3}\) and \(\mathrm{O}_{2}\) in a basic solution that produces \(\mathrm{CrO}_{4}^{2-}\) ions.

An electrochemical cell with an aqueous electrolyte is based on the reaction between \(\mathrm{Ni}^{2+}(a q)\) and \(\mathrm{Cd}(s)\) producing \(\mathrm{Ni}(s)\) and \(\mathrm{Cd}^{2+}(a q)\) a. Write half-reactions for the anode and cathode. b. Write a balanced net ionic equation describing the cell reaction. c. Draw the cell diagram.

A voltaic cell is based on the reaction between \(\mathrm{Cu}^{2+}(a q)\) and \(\mathrm{Ni}(s),\) producing \(\mathrm{Cu}(s)\) and \(\mathrm{Ni}^{2+}(a q)\) a. Write the anode and cathode half-reactions. b. Write a balanced cell reaction. c. Draw the cell diagram.

A voltaic cell with a basic aqueous background electrolyte is based on the oxidation of \(\mathrm{Cd}(s)\) to \(\mathrm{Cd}(\mathrm{OH})_{2}(s)\) and the reduction of \(\mathrm{MnO}_{4}^{-}(a q)\) to \(\mathrm{MnO}_{2}(s)\) a. Write half-reactions for the cell's anode and cathode. b. Write a balanced net ionic equation describing the cell reaction. c. Draw the cell diagram.

A voltaic cell is based on the reduction of \(\mathrm{Ag}^{+}(a q)\) to \(\mathrm{Ag}(s)\) and the oxidation of \(\operatorname{Sn}(s)\) to \(\mathrm{Sn}^{2+}(a q)\) a. Write half-reactions for the cell's anode and cathode. b. Write a balanced cell reaction. c. Draw the cell diagram.

Super Iron Batteries In \(1999,\) scientists in Israel developed a battery based on the following cell reaction with iron(VI), nicknamed "super iron": \(\beth \mathrm{K}_{2} \mathrm{FeO}_{4}(a q)+3 \mathrm{Zn}(s) \rightarrow \mathrm{Fe}_{2} \mathrm{O}_{3}(s)+\mathrm{ZnO}(s)+2 \mathrm{K}_{2} \mathrm{ZnO}_{2}(a q)\) a. Determine the number of electrons transferred in the cell reaction. b. What are the oxidation states of the transition metals in the reaction? c. Draw the cell diagram.

Aluminum-Air Batteries In recent years engineers have been working on an aluminum-air battery as an alternative energy source for electric vehicles. The battery consists of an aluminum anode, which is oxidized to solid aluminum hydroxide, immersed in an electrolyte of aqueous KOH. At the cathode oxygen from the air is reduced to hydroxide ions on an inert metal surface. Write the two half-reactions for the battery and diagram the cell. Use the generic \(\mathrm{M}(s)\) symbol for the metallic cathode material.

Of the group 1 elements \(\mathrm{Li}, \mathrm{K},\) and \(\mathrm{Na},\) which is the strongest reducing agent?

If a piece of silver is placed in a solution in which \(\left[\mathrm{Ag}^{+}\right]=\left[\mathrm{Cu}^{2+}\right]=1.00 \mathrm{M},\) will the following reaction proceed spontaneously? $$2 \mathrm{Ag}(s)+\mathrm{Cu}^{2+}(a q) \rightarrow 2 \mathrm{Ag}^{+}(a q)+\mathrm{Cu}(s)$$

Sometimes the anode half-reaction in the zinc-air battery (Figure 18.7 ) is written with the zincate ion, \(\mathrm{Zn}(\mathrm{OH})_{4}^{2-}\) as the product. Write a balanced equation for the cell reaction based on this product.

Sometimes the cell reaction of nickel-cadmium batteries is written with Cd metal as the anode and solid \(\mathrm{NiO}_{2}\) as the cathode. Assuming that the products of the reactions are a solid hydroxide of cadmium(II) at the anode and a solid hydroxide of nickel(II) at the cathode, write balanced equations for the cathode and anode half-reactions and the overall cell reaction.

Voltaic cells based on the following pairs of half-reactions are constructed. For each pair, write a balanced equation for the cell reaction, and identify which half-reaction takes place at each anode and cathode. a. \(\mathrm{Cd}^{2+}(a q)+2 \mathrm{e}^{-} \rightarrow \mathrm{Cd}(s)\) \(\mathrm{Ag}^{+}(a q)+\mathrm{e}^{-} \rightarrow \mathrm{Ag}(s)\) b. \(\mathrm{AgBr}(s)+\mathrm{e}^{-} \rightarrow \mathrm{Ag}(s)+\mathrm{Br}^{-}(a q)\) \(\mathrm{MnO}_{2}(s)+4 \mathrm{H}^{+}(a q)+2 \mathrm{e}^{-} \rightarrow \mathrm{Mn}^{2+}(a q)+2 \mathrm{H}_{2} \mathrm{O}(\ell)\) c. \(\mathrm{PtCl}_{4}^{2-}(a q)+2 \mathrm{e}^{-} \rightarrow \mathrm{Pt}(s)+4 \mathrm{Cl}^{-}(a q)\) \(\mathrm{AgCl}(s)+\mathrm{e}^{-} \rightarrow \mathrm{Ag}(s)+\mathrm{Cl}^{-}(a q)\)

Which of the following reductions will occur in the presence of \(\mathrm{H}_{2}\) gas under standard conditions? a. \(A g^{+}\) to \(A g\) b. \(\mathrm{Mg}^{2+}\) to \(\mathrm{Mg}\) c. \(\mathrm{Cu}^{2+}\) to \(\mathrm{Cu}\) d. \(\mathrm{Cd}^{2+}\) to \(\mathrm{Cd}\)

Which of the following oxidations will occur in the presence of \(\mathrm{H}_{2}\) gas under standard conditions? a. \(\mathrm{Zn}^{2+}\) to \(\mathrm{Zn}\) b. \(\mathrm{Fe}^{2+}\) to \(\mathrm{Fe}^{3+}\) c. \(\mathrm{Cr}(\mathrm{OH})_{3}\) to \(\mathrm{CrO}_{4}^{2-}\) d. Ni to Ni \(^{2+}\)

The half-reactions and standard potentials for a nickelmetal hydride battery with a titanium-zirconium anode are as follows: Cathode: \(\quad \mathrm{NiO}(\mathrm{OH})(s)+\mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{e}^{-} \rightarrow \mathrm{Ni}(\mathrm{OH})_{2}(s)+\mathrm{OH}^{-}(a q)\) \(E^{*}=0.52 \mathrm{V}\) Anode: \(\quad \mathrm{TiZr}_{2} \mathrm{H}(s)+\mathrm{OH}^{-}(a q) \rightarrow \mathrm{TiZr}_{2}(s)+\mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{e}^{-}\) \(E^{\circ}=0.00 \mathrm{V}\) a. Write the overall cell reaction for this battery. b. Calculate the standard cell potential.

The value of \(E_{\text {cell for the reaction below is } 0.500 \mathrm{V} . \text { What is }}\) the value of \(\Delta G_{\text {cell }}^{\text {? }}\) $$\mathrm{Mn}^{3+}+2 \mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{Mn}^{2+}+\mathrm{MnO}_{2}+4 \mathrm{H}^{+}$$

What is the value of \(\Delta G_{\text {cell }}^{\circ}\) for an electrochemical cell based on a cell reaction described by the following net ionic equation? $$\mathrm{Mg}+2 \mathrm{Cu}^{+} \rightarrow \mathrm{Mg}^{2+}+2 \mathrm{Cu}$$

A cell in a lead-acid battery delivers exactly \(2.00 \mathrm{V}\) of cell potential in accordance with the following cell reaction: \(\mathrm{Pb}(s)+\mathrm{PbO}_{2}(s)+2 \mathrm{H}_{2} \mathrm{SO}_{4}(a q) \rightarrow 2 \mathrm{PbSO}_{4}(s)+2 \mathrm{H}_{2} \mathrm{O}(\ell)\) What is the value of \(\Delta G_{\text {cell }} ?\)

What is the function of platinum in the standard hydrogen electrode?

Platinum is very expensive, so why is it used in standard hydrogen standard electrodes where the half-reaction is \(2 \mathrm{H}^{+}(a q)+2 \mathrm{e}^{-} \rightarrow \mathrm{H}_{2}(\mathrm{g}) ?\)

The potential of the standard hydrogen electrode (SHE) is the reference against which other half-reaction potentials are expressed. Why, then, is the SHE not widely used as a reference electrode in electrochemical cells?

Suggest a replacement metal for platinum in the standard hydrogen electrode. Explain why you selected the metal you did.

An electrochemical cell consists of a standard hydrogen electrode and a second half-cell in which a magnesium electrode is immersed in a \(1.00 M\) solution of \(\mathrm{Mg}^{2+}\) ions. a. What is the value of \(E_{\text {cell }} ?\) b. Which electrode is the anode? c. Which is a product of the cell reaction: \(\mathrm{H}^{+}\) ions or \(\mathrm{H}_{2}\) gas?

An electrochemical cell consists of a standard hydrogen electrode and a second half-cell in which a cadmium electrode is immersed in a \(1.00 M\) solution of \(\mathrm{Cd}^{2+}\) ions. a. What is the value of \(E_{\text {cell }} ?\) b. Which electrode is the anode? c. Which is a product of the cell reaction: \(\mathrm{Cd}^{2+}\) ions or \(\mathrm{Cd}\) metal?

Why does the operating cell potential of most batteries change little until the battery is nearly discharged?

The standard potential of the Cu-Zn cell reaction, $$\mathrm{Zn}(s)+\mathrm{Cu}^{2+}(a q) \rightarrow \mathrm{Zn}^{2+}(a q)+\mathrm{Cu}(s)$$ is \(1.10 \mathrm{V} .\) Would the potential of the Cu-Zn cell differ from \(1.10 \mathrm{V}\) if the concentrations of both \(\mathrm{Cu}^{2+}\) and \(\mathrm{Zn}^{2+}\) were \(0.25 M ?\)

If the potential of a hydrogen electrode based on the half-reaction $$2 \mathrm{H}^{+}(a q)+2 \mathrm{e}^{-} \rightarrow \mathrm{H}_{2}(g)$$ is \(0.000 \mathrm{V}\) at \(\mathrm{pH}=0.00,\) what is the potential of the same electrode at \(\mathrm{pH}=7.00 ?\)

Glucose Metabolism The standard potentials for the reduction of nicotinamide adenine dinucleotide (NAD') and oxaloacetate (reactants in the multistep metabolism of glucose) are as follows: \(\mathrm{NAD}^{+}(a q)+2 \mathrm{H}^{+}(a q)+2 \mathrm{e}^{-} \rightarrow \mathrm{NADH}(a q)+\mathrm{H}^{+}(a q)\) $$ E^{*}=-0.320 \mathrm{V} $$ Oxaloacetate \(^{2-}(a q)+2 \mathrm{H}^{+}(a q)+2 \mathrm{e}^{-} \rightarrow\) malate \(^{2-}(a q)\) $$ E^{\circ}=-0.166 \mathrm{V} $$ a. Calculate the standard potential for the following reaction: Oxaloacetate \(^{2-}(a q)+\mathrm{NADH}(a q)+\mathrm{H}^{+}(a q) \rightarrow\) $$ \text { malate }^{-}(a q)+\mathrm{NAD}^{+}(a q) $$ b. Calculate the equilibrium constant for the reaction at \(25^{\circ} \mathrm{C}\)

A concentration cell can be constructed by using the same half-reaction for both the cathode and anode. What is the value of \(E_{\text {cell }}\) for a concentration cell that combines copper electrodes in contact with \(0.35 M\) copper (II) nitrate and \(0.00075 M\) copper \((\mathrm{II})\) nitrate solutions?

Chlorine dioxide \(\left(\mathrm{ClO}_{2}\right)\) is produced by the following reaction of chlorate \(\left(\mathrm{ClO}_{3}^{-}\right)\) with \(\mathrm{Cl}^{-}\) in acid solution: \(\begin{aligned} 2 \mathrm{ClO}_{3}^{-}(a q)+2 \mathrm{Cl}^{-}(a q)+4 \mathrm{H}^{+}(a q) \rightarrow & \\ 2 \mathrm{ClO}_{2}(g)+\mathrm{Cl}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(\ell) \end{aligned}\) a. Determine \(E^{\circ}\) for the reaction. b. The reaction produces a mixture of gases in the reaction vessel in which \(P_{\mathrm{C} 10_{2}}=2.0 \mathrm{atm} ; P_{\mathrm{C}_{2}}=1.00 \mathrm{atm}\) Calculate \(\left[\mathrm{ClO}_{3}^{-}\right]\) if, at equilibrium \(\left(T=25^{\circ} \mathrm{C}\right),\left[\mathrm{H}^{+}\right]=\) \(\left[\mathrm{Cl}^{-}\right]=10.0 \mathrm{M}\)

One 12 -volt lead-acid battery has a higher ampere-hour rating than another. Which of the following parameters are likely to be different for the two batteries? a. individual cell potentials b. anode half-reactions c. total masses of electrode materials d. number of cells e. electrolyte composition f. combined surface areas of their electrodes

In a voltaic cell based on the Cu-Zn cell reaction $$ \mathrm{Zn}(s)+\mathrm{Cu}^{2+}(a q) \rightarrow \mathrm{Cu}(s)+\mathrm{Zn}^{2+}(a q) $$ there is exactly 1 mole of each reactant and product. A second cell based on the Cd-Cu cell reaction $$ \mathrm{Cd}(s)+\mathrm{Cu}^{2+}(a q) \rightarrow \mathrm{Cu}(s)+\mathrm{Cd}^{2+}(a q) $$ also has exactly 1 mole of each reactant and product. Which of the following statements about these two cells is true? a. Their cell potentials are the same. b. The masses of their electrodes are the same. c. The quantities of electrical charge that they can produce are the same. d. The quantities of electrical energy that they can produce are the same.

Which of the following voltaic cells will produce the greater quantity of electrical charge per gram of anode material? \(\mathrm{Cd}(s)+2 \mathrm{NiO}(\mathrm{OH})(s)+2 \mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow 2 \mathrm{Ni}(\mathrm{OH})_{2}(s)+\mathrm{Cd}(\mathrm{OH})_{2}(s)\) \(\quad\) or \(\quad 4 \mathrm{Al}(s)+3 \mathrm{O}_{2}(g)+6 \mathrm{H}_{2} \mathrm{O}(\ell)+4 \mathrm{OH}^{-}(a q) \rightarrow 4 \mathrm{Al}(\mathrm{OH})_{4}^{-}(a q)\)

Which of the following voltaic cells will produce the greater quantity of electrical charge per gram of anode material? $$\mathrm{Zn}(s)+\mathrm{MnO}_{2}(s)+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{ZnO}(s)+\mathrm{Mn}(\mathrm{OH})_{2}(s)$$ or$$\mathrm{Li}(s)+\mathrm{MnO}_{2}(s) \rightarrow \mathrm{LiMnO}_{2}(s)$$

Which of the following voltaic cell reactions delivers more electrical energy per gram of anode material at \(25^{\circ} \mathrm{C} ?\) $$ \mathrm{Zn}(s)+\mathrm{Ni}(\mathrm{OH})_{2}(s) \rightarrow \mathrm{Zn}(\mathrm{OH})_{2}(s)+\mathrm{Ni}(s) \quad E_{\mathrm{cell}}^{*}=1.50 \mathrm{V}$$ or$$2 \mathrm{Zn}(s)+\mathrm{O}_{2}(g) \rightarrow 2 \mathrm{ZnO}(s) \quad E_{\mathrm{cell}}^{\circ}=2.08 \mathrm{V}$$

When the iron skeleton of the Statue of Liberty was replaced with stainless steel, the asbestos mats that had separated the skeleton from the copper exterior were replaced with Teflon spacers. Why was Teflon a good choice?

What does a sacrificial anode do to protect a metal structure, and why is the process called catbodic protection?

The positive terminal of a voltaic cell is the cathode. However, the cathode of an electrolytic cell is connected to the negative terminal of a power supply. Explain this difference in polarity.

The positive terminal of a voltaic cell is the cathode. However, the cathode of an electrolytic cell is connected to the negative terminal of a power supply. Explain this difference in polarity.

The anode in an electrochemical cell is defined as the electrode where oxidation takes place. Why is the anode in an electrolytic cell connected to the positive \((+)\) terminal of an external supply, whereas the anode in a voltaic cell battery is connected to the negative \((-)\) terminal?

The salts obtained from the evaporation of seawater can be a source of halogens, principally \(\mathrm{Cl}_{2}\) and \(\mathrm{Br}_{2},\) through the electrolysis of the molten alkali metal halides. As the potential of the anode in an electrolytic cell is increased, which of these two halogens forms first?

Quantitative Analysis Electrolysis can be used to determinethe concentration of \(\mathrm{Cu}^{2+}\) in a given volume of solution by electrolyzing the solution in a cell equipped with a platinum cathode. If all the \(\mathrm{Cu}^{2+}\) is reduced to Cu metal at the cathode, the increase in mass of the electrode provides a measure of the concentration of \(\mathrm{Cu}^{2+}\) in the original solution. To ensure the complete \((99.99 \%)\) removal of the \(\mathrm{Cu}^{2+}\) from a solution in which \(\left[\mathrm{Cu}^{2+}\right]\) is initially about 1.0 \(M,\) will the potential of the cathode (versus SHE) have to be more or less negative than \(0.34 \mathrm{V}\) (the standard potential for \(\left.\mathrm{Cu}^{2+}+2 \mathrm{e}^{-} \rightarrow \mathrm{Cu}\right) ?\)

How long does it take to deposit a coating of gold \(1.00 \mu \mathrm{m}\) thick on a disk-shaped medallion \(2.0 \mathrm{cm}\) in diameter and \(3.0 \mathrm{mm}\) thick at a constant current of \(45 \mathrm{A} ?\) The density of gold is \(18.3 \mathrm{g} / \mathrm{cm}^{3} .\) The gold solution contains gold( 111 ).

Oxygen Supply in Submarines Nuclear submarines can stay under water nearly indefinitely because they can produce their own oxygen by the electrolysis of water. a. How many liters of \(\mathrm{O}_{2}\) at \(25^{\circ} \mathrm{C}\) and 1.00 bar are produced in 1 hour in an electrolytic cell operating at a current of \(0.025 \mathrm{A} ?\) b. Could seawater be used as the source of oxygen in this electrolysis? Explain why or why not.

Describe two advantages of hybrid (gasoline engineelectric motor) power systems over all-electric systems based on fuel cells. Describe two disadvantages.