Chapter 18

Write a balanced chemical equation for the overall cell reaction represented as (a) \(\mathrm{Mg}\left|\mathrm{Mg}^{2+} \| \mathrm{Sc}^{3+}\right| \mathrm{Sc}\) (b) Sn \(\left|\mathrm{Sn}^{2+} \| \mathrm{Pb}^{2+}\right| \mathrm{Pb}\) (c) \(\mathrm{Pt}\left|\mathrm{Cl}^{-}\right| \mathrm{Cl}_{2} \| \mathrm{NO}_{3}^{-}|\mathrm{NO}| \mathrm{Pt}\)

Write a balanced chemical equation for the overall cell reaction represented as (a) \(\mathrm{Ag}\left|\mathrm{Ag}^{+} \| \mathrm{Sn}^{4+}, \mathrm{Sn}^{2+}\right| \mathrm{Pt}\) (b) \(\mathrm{Al}\left|\mathrm{Al}^{3+} \| \mathrm{Cu}^{2+}\right| \mathrm{Cu}\) (c) \(\mathrm{Pt}\left|\mathrm{Fe}^{2+}, \mathrm{Fe}^{3+} \| \mathrm{MnO}_{4}^{-}, \mathrm{Mn}^{2+}\right| \mathrm{Pt}\)

Draw a diagram for a salt bridge cell for each of the following reactions. Label the anode and cathode, and indicate the direction of current flow throughout the circuit. (a) \(\mathrm{Zn}(s)+\mathrm{Cd}^{2+}(a q) \longrightarrow \mathrm{Zn}^{2+}(a q)+\mathrm{Cd}(s)\) (b) \(2 \mathrm{AuCl}_{4}^{-}(a q)+3 \mathrm{Cu}(s) \longrightarrow 2 \mathrm{Au}(s)+8 \mathrm{Cl}^{-}(a q)+3 \mathrm{Cu}^{2+}(a q)\) (c) \(\mathrm{Fe}(s)+\mathrm{Cu}(\mathrm{OH})_{2}(s) \longrightarrow \mathrm{Cu}(s)+\mathrm{Fe}(\mathrm{OH})_{2}(s)\)

Consider a salt bridge voltaic cell represented by the following reaction: $$ \mathrm{Fe}(s)+2 \mathrm{Tl}^{+}(a q) \longrightarrow \mathrm{Fe}^{2+}(a q)+2 \mathrm{Tl}(s) $$ Choose the best answer from the choices in each part below: (a) What is the path of electron flow? Through the salt bridge, or through the external circuit? (b) To which half-cell do the negative ions in the salt bridge move? The anode, or the cathode? (c) Which metal is the electrode in the anode?

Consider a salt bridge voltaic cell represented by the following reaction: $$ \mathrm{Ca}(s)+2 \mathrm{H}^{+}(a q) \longrightarrow \mathrm{Ca}^{2+}(a q)+\mathrm{H}_{2}(g) $$ (a) What is the direction of the electrons in the external circuit? (b) What electrode can be used at the cathode? (c) What is the reaction occurring at the anode?

Consider a salt bridge cell in which the anode is a manganese rod immersed in an aqueous solution of manganese(II) sulfate. The cathode is a chromium strip immersed in an aqueous solution of chromium(III) sulfate. Sketch a diagram of the cell, indicating the flow of the current throughout. Write the half- equations for the electrode reactions, the overall equation, and the abbreviated notation for the cell.

Which species in each pair is the stronger oxidizing agent? (a) \(\mathrm{NO}_{3}^{-}\) or \(\mathrm{I}_{2}\) (b) \(\mathrm{Fe}(\mathrm{OH})_{3}\) or \(\mathrm{S}\) (c) \(\mathrm{Mn}^{2+}\) or \(\mathrm{MnO}_{2}\) (d) \(\mathrm{ClO}_{3}^{-}\) in acidic solution or \(\mathrm{ClO}_{3}^{-}\) in basic solution

Which species in each pair is the stronger reducing agent? (a) \(\mathrm{Cr}\) or \(\mathrm{Cd}\) (b) \(\mathrm{I}^{-}\) or \(\mathrm{Br}^{-}\) (c) \(\mathrm{OH}^{-}\) or \(\mathrm{NO}_{2}^{-}\) (d) NO in acidic solution or NO in basic solution

Consider the following species. $$ \begin{array}{llll} \mathrm{Cr}^{3+} & \mathrm{Hg}(l) & \mathrm{H}_{2} \text { (acidic) } & \mathrm{Sn}^{2+} \end{array} $$ \(\mathrm{Br}_{2}\) (acidic) Classify each species as oxidizing agent, reducing agent, or both. Arrange the oxidizing agents in order of increasing strength. Do the same for the reducing agents.

For the following half-reactions, answer the questions below: $$ \begin{array}{cc} \mathrm{Ce}^{4+}(a q)+e^{-} \longrightarrow \mathrm{Ce}^{3+}(a q) & E^{\circ}=+1.61 \mathrm{~V} \\ \mathrm{Ag}^{+}(a q)+e^{-} \longrightarrow \mathrm{Ag}(s) & E^{\circ}=+0.80 \mathrm{~V} \\ \mathrm{Hg}_{2}^{2+}(a q)+2 e^{-} \longrightarrow 2 \mathrm{Hg}(l) & E^{\circ}=+0.80 \mathrm{~V} \\ \mathrm{Sn}^{2+}(a q)+2 e^{-} \longrightarrow \mathrm{Sn}(s) & E^{\circ}=-0.14 \mathrm{~V} \\ \mathrm{Ni}^{2+}(a q)+2 e^{-} \longrightarrow \mathrm{Ni}(s) & E^{\circ}=-0.24 \mathrm{~V} \\ \mathrm{Al}^{3+}(a q)+3 e^{-} \longrightarrow \mathrm{Al}(s) & E^{o}=-1.68 \mathrm{~V} \end{array} $$ (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 \(\mathrm{Sn}(s)\) reduce \(\mathrm{Ag}^{+}(\mathrm{aq})\) to \(\mathrm{Ag}(s) ?\) (f) Will \(\mathrm{Hg}(l)\) reduce \(\mathrm{Sn}^{2+}(a q)\) to \(\mathrm{Sn}(s) ?\) (g) Which ion(s) can be reduced by \(\operatorname{Sn}(s)\) ? (h) Which metal(s) can be oxidized by \(\mathrm{Ag}^{+}(a q)\) ?

For the following half-reactions, answer the questions below. $$ \begin{array}{cc} \mathrm{Co}^{3+}(a q)+e^{-} \longrightarrow \mathrm{Co}^{2+}(a q) & E^{\circ}=+1.953 \mathrm{~V} \\ \mathrm{Fe}^{3+}(a q)+e^{-} \longrightarrow \mathrm{Fe}^{2+}(a q) & E^{\circ}=+0.769 \mathrm{~V} \\ \mathrm{I}_{2}(a q)+2 e^{-} \longrightarrow 2 \mathrm{I}^{-}(a q) & E^{o}=+0.534 \mathrm{~V} \\ \mathrm{~Pb}^{2+}(a q)+2 e^{-} \longrightarrow \mathrm{Pb}(s) & E^{\circ}=-0.127 \mathrm{~V} \\ \mathrm{Cd}^{2+}(a q)+2 e^{-} \longrightarrow \mathrm{Cd}(s) & E^{\circ}=-0.402 \mathrm{~V} \\ \mathrm{Mn}^{2+}(a q)+2 e^{-} \longrightarrow \mathrm{Mn}(s) & E^{\circ}=-1.182 \mathrm{~V} \end{array} $$ (a) Which is the weakest reducing agent? (b) Which is the strongest reducing agent? (c) Which is the strongest oxidizing agent? (d) Which is the weakest oxidizing agent? (e) Will \(\mathrm{Pb}(s)\) reduce \(\mathrm{Fe}^{3+}(a q)\) to \(\mathrm{Fe}^{2+}(a q) ?\) (f) Will \(\mathrm{I}^{-}(a q)\) reduce \(\mathrm{Pb}^{2+}(a q)\) to \(\mathrm{Pb}(s) ?\) (g) Which ion(s) can be reduced by \(\mathrm{Pb}(s)\) ? (h) Which if any metal(s) can be oxidized by \(\mathrm{Fe}^{3+}(a q)\) ?

Calculate \(E^{\circ}\) for the following voltaic cells: (a) \(\mathrm{MnO}_{2}(s)+4 \mathrm{H}^{+}(a q)+2 \mathrm{I}^{-}(a q) \longrightarrow\) \(\mathrm{Mn}^{2+}(a q)+2 \mathrm{H}_{2} \mathrm{O}+\mathrm{I}_{2}(s)\) (b) \(\mathrm{H}_{2}(g)+2 \mathrm{OH}^{-}(a q)+\mathrm{S}(s) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}+\mathrm{S}^{2-}(a q)\) (c) an \(\mathrm{Ag}-\mathrm{Ag}^{+}\) half-cell and an \(\mathrm{Au}-\mathrm{AuCl}_{4}^{-}\) half-cell.

Calculate \(E^{\circ}\) for the following voltaic cells: (a) \(\mathrm{Pb}(s)+2 \mathrm{Ag}^{+}(a q) \longrightarrow \mathrm{Pb}^{2+}(a q)+2 \mathrm{Ag}(s)\) (b) \(\mathrm{O}_{2}(\mathrm{~g})+4 \mathrm{Fe}^{2+}(a q)+4 \mathrm{H}^{+}(a q) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}+4 \mathrm{Fe}^{3+}(a q)\) (c) a Cd-Cd \(^{2+}\) half-cell and a \(\mathrm{Zn}-\mathrm{Zn}^{2+}\) half-cell

Calculate \(E^{\circ}\) for the following cells: (a) \(\mathrm{Mn}\left|\mathrm{Mn}^{2+} \| \mathrm{H}^{+}\right| \mathrm{H}_{2} \mid \mathrm{Pt}\) (b) \(\mathrm{Au}\left|\mathrm{AuCl}_{4}^{-} \| \mathrm{Co}^{3+}, \mathrm{Co}^{2+}\right| \mathrm{Pt}\) (c) \(\mathrm{Pt}\left|\mathrm{S}^{2-}\right| \mathrm{S} \| \mathrm{NO}_{3}^{-}|\mathrm{NO}| \mathrm{Pt} \quad\) (basic medium)

Calculate \(E^{\circ}\) for the following cells: (a) \(\mathrm{Pb}\left|\mathrm{PbSO}_{4} \| \mathrm{Pb}^{2+}\right| \mathrm{Pb}\) (b) \(\mathrm{Pt}\left|\mathrm{Cl}_{2}\right| \mathrm{ClO}_{3}^{-} \| \mathrm{O}_{2}\left|\mathrm{H}_{2} \mathrm{O}\right| \mathrm{Pt}\) (c) \(\mathrm{Pt}\left|\mathrm{OH}^{-}\right| \mathrm{O}_{2} \| \mathrm{ClO}_{3}^{-}, \mathrm{Cl}^{-} \mid \mathrm{Pt} \quad\) (basic medium)

I Suppose \(E_{\text {red }}^{\circ}\) for \(\mathrm{H}^{+} \longrightarrow \mathrm{H}_{2}\) were taken to be \(0.300 \mathrm{~V}\) instead of \(0.000 \mathrm{~V}\). What would be (a) \(E_{\text {ox }}^{o}\) for \(\mathrm{H}_{2} \longrightarrow \mathrm{H}^{+}\) ? (b) \(E_{\text {red }}^{\circ}\) for \(\mathrm{Br}_{2} \longrightarrow \mathrm{Br}^{-} ?\) (c) \(E^{\circ}\) for the cell in \(24(\mathrm{c})\) ? Compare your answer with that obtained in \(24(\mathrm{c})\).

Which of the following reactions is (are) spontaneous at standard conditions? (a) \(2 \mathrm{NO}_{3}^{-}(a q)+8 \mathrm{H}^{+}(a q)+6 \mathrm{Cl}^{-}(a q) \longrightarrow\) \(2 \mathrm{NO}(g)+4 \mathrm{H}_{2} \mathrm{O}+3 \mathrm{Cl}_{2}(g)\) (b) \(\mathrm{O}_{2}(g)+4 \mathrm{H}^{+}(a q)+4 \mathrm{Cl}^{-}(a q) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}+2 \mathrm{Cl}_{2}(g)\) (c) \(3 \mathrm{Fe}(s)+2 \mathrm{AuCl}_{4}^{-}(a q) \longrightarrow 2 \mathrm{Au}(s)+8 \mathrm{Cl}^{-}(a q)+3 \mathrm{Fe}^{2+}(a q)\)

Which of the following reactions is (are) spontaneous at standard conditions? (a) \(\mathrm{Zn}(s)+2 \mathrm{Fe}^{3+}(a q) \longrightarrow \mathrm{Zn}^{2+}(a q)+2 \mathrm{Fe}^{2+}(a q)\) (b) \(\mathrm{Cu}(s)+2 \mathrm{H}^{+}(a q) \longrightarrow \mathrm{Cu}^{2+}(a q)+\mathrm{H}_{2}(g)\) (c) \(2 \mathrm{Br}^{-}(a q)+\mathrm{I}_{2}(s) \longrightarrow \mathrm{Br}_{2}(l)+2 \mathrm{I}^{-}(a q)\)

Use the following half-equations to write three spontaneous reactions. Justify your answers by calculating \(E^{\circ}\) for the cells. (1) \(\mathrm{MnO}_{4}^{-}(a q)+8 \mathrm{H}^{+}(a q)+5 e^{-} \longrightarrow \mathrm{Mn}^{2+}(a q)+4 \mathrm{H}_{2} \mathrm{O}\) \(E^{\circ}=+1.512 \mathrm{~V}\) (2) \(\mathrm{O}_{2}(\mathrm{~g})+4 \mathrm{H}^{+}(\mathrm{aq})+4 e^{-} \longrightarrow 2 \mathrm{H}_{2} \mathrm{O} \quad E^{\circ}=+1.229 \mathrm{~V}\) (3) \(\mathrm{Co}^{2+}(a q)+2 e^{-} \longrightarrow \mathrm{Co}(s) \quad E^{\circ}=-0.282 \mathrm{~V}\)

Write the equation for the reaction, if any, that occurs when each of the following experiments is performed under standard conditions. (a) Crystals of iodine are added to an aqueous solution of potassium bromide. (b) Liquid bromine is added to an aqueous solution of sodium chloride. (c) A chromium wire is dipped into a solution of nickel(II) chloride.

Write the equation for the reaction, if any, that occurs when each of the following experiments is performed under standard conditions. (a) Sulfur is added to mercury. (b) Manganese dioxide in acidic solution is added to liquid mercury. (c) Aluminum metal is added to a solution of potassium ions.

Which of the following species will react with \(1 \mathrm{M} \mathrm{HNO}_{3}\) ? (a) \(\mathrm{I}^{-}\) (b) Fe (c) \(\mathrm{Ag}\) (d) \(\mathrm{Pb}\)

Which of the following species will be oxidized by \(1 \mathrm{MHCl}\) ? (a) \(\mathrm{Au}\) (b) \(\mathrm{Mg}\) (c) Cu (d) \(\mathrm{F}^{-}\)

Given the following standard reduction potentials $$ \begin{gathered} \mathrm{Ag}^{+}(a q)+e^{-} \longrightarrow \mathrm{Ag}(s) \quad E^{\circ}=0.799 \mathrm{~V} \\ \mathrm{Ag}(\mathrm{CN})_{2}^{-}+e^{-} \longrightarrow \mathrm{Ag}(s)+2 \mathrm{CN}^{-}(a q) \quad E^{\circ}=-0.31 \mathrm{~V} \end{gathered} $$ find \(K_{f}\) for \(\mathrm{Ag}(\mathrm{CN})_{2}^{-}(a q)\) at \(25^{\circ} \mathrm{C}\).

What is \(E^{\circ}\) at \(25^{\circ} \mathrm{C}\) for the following reaction? $$ \mathrm{Ba}^{2+}(a q)+\mathrm{SO}_{4}^{2-}(a q) \longrightarrow \mathrm{BaSO}_{4}(s) $$ \(K_{s p}\) for \(\mathrm{BaSO}_{4}\) is \(1.1 \times 10^{-10}\)

Consider a voltaic cell at \(25^{\circ} \mathrm{C}\) in which the following reaction takes place. $$ 3 \mathrm{H}_{2} \mathrm{O}_{2}(a q)+6 \mathrm{H}^{+}(a q)+2 \mathrm{Au}(s) \longrightarrow 2 \mathrm{Au}^{3+}(a q)+6 \mathrm{H}_{2} \mathrm{O} $$ (a) Calculate \(E\). (b) Write the Nernst equation for the cell. (c) Calculate \(E\) when \(\left[\mathrm{Au}^{3+}\right]=0.250 \mathrm{M},\left[\mathrm{H}^{+}\right]=1.25 \mathrm{M},\left[\mathrm{H}_{2} \mathrm{O}_{2}\right]=\) \(1.50 M\)

Consider a voltaic cell at \(25^{\circ} \mathrm{C}\) in which the following reaction takes place. $$ 3 \mathrm{O}_{2}(g)+4 \mathrm{NO}(g)+2 \mathrm{H}_{2} \mathrm{O} \longrightarrow 4 \mathrm{NO}_{3}^{-}(a q)+4 \mathrm{H}^{+}(a q) $$ (a) Calculate \(E^{\circ}\) (b) Write the Nernst equation for the cell. (c) Calculate \(E\) under the following conditions: \(\left[\mathrm{NO}_{3}{\underline{\phantom{xx}}}^{-}\right]=0.750 M\), \(P_{\mathrm{NO}}=0.993 \mathrm{~atm}, P_{\mathrm{O}_{2}}=0.515 \mathrm{~atm}, \mathrm{pH}=2.85\)

Consider a voltaic cell in which the following reaction takes place. $$ 2 \mathrm{NO}_{3}^{-}(a q)+3 \mathrm{H}_{2}(g) \longrightarrow 2 \mathrm{NO}(g)+2 \mathrm{OH}^{-}(a q)+2 \mathrm{H}_{2} \mathrm{O} $$ (a) Calculate \(E^{\circ}\). (b) Write the Nernst equation for the cell. (c) Calculate \(E\) under the following conditions: \(\left[\mathrm{NO}_{3}^{-}\right]=0.0315 M\), \(P_{\mathrm{NO}}=0.922 \mathrm{~atm}, P_{\mathrm{H}_{2}}=0.437 \mathrm{~atm}, \mathrm{pH}=11.50 .\)

Calculate voltages of the following cells at \(25^{\circ} \mathrm{C}\) and under the following conditions. (a) \(\mathrm{Fe}\left|\mathrm{Fe}^{2+}(0.010 \mathrm{M}) \| \mathrm{Cu}^{2+}(0.10 \mathrm{M})\right| \mathrm{Cu}\) (b) \(\mathrm{Pt}\left|\mathrm{Sn}^{2+}(0.10 \mathrm{M}), \mathrm{Sn}^{4+}(0.010 \mathrm{M}) \| \mathrm{Co}^{2+}(0.10 \mathrm{M})\right| \mathrm{Co}\)

Consider the reaction $$ \mathrm{S}(s)+2 \mathrm{H}^{+}(a q)+2 \mathrm{Ag}(s)+2 \mathrm{Br}^{-}(a q) \longrightarrow 2 \mathrm{AgBr}(s)+\mathrm{H}_{2} \mathrm{~S}(a q) $$ At what \(\mathrm{pH}\) is the voltage zero if all other species are at standard concentrations?

Consider the reaction below at \(25^{\circ} \mathrm{C}\) : $$ 3 \mathrm{SO}_{4}{\underline{\phantom{xx}}}^{2-}(a q)+12 \mathrm{H}^{+}(a q)+2 \mathrm{Cr}(s) \longrightarrow 3 \mathrm{SO}_{2}(g)+2 \mathrm{Cr}^{3+}(a q)+6 \mathrm{H}_{2} \mathrm{O} $$ Use Table \(18.1\) to answer the following questions. Support your answers with calculations. (a) Is the reaction spontaneous at standard conditions? (b) Is the reaction spontaneous at a \(\mathrm{pH}\) of \(3.00\) with all other ionic species at \(0.100 \mathrm{M}\) and gases at \(1.00\) atm? (c) Is the reaction spontaneous at a pH of \(8.00\) with all other ionic species at \(0.100 \mathrm{M}\) and gases at \(1.00 \mathrm{~atm}\) ? (d) At what \(\mathrm{pH}\) is the reaction at equilibrium with all other ionic species at \(0.100 \mathrm{M}\) and gases at \(1.00 \mathrm{~atm}\) ?

Consider a cell in which the reaction is $$ 2 \mathrm{Ag}(s)+\mathrm{Cu}^{2+}(a q) \longrightarrow 2 \mathrm{Ag}^{+}(a q)+\mathrm{Cu}(s) $$ (a) Calculate \(E^{\circ}\) for this cell. (b) Chloride ions are added to the \(\mathrm{Ag} \mid \mathrm{Ag}^{+}\) half- cell to precipitate \(\mathrm{AgCl}\). The measured voltage is \(+0.060 \mathrm{~V}\). Taking \(\left[\mathrm{Cu}^{2+}\right]=1.0 \mathrm{M}\), calculate \(\left[\mathrm{Ag}^{+}\right]\). (c) Taking \(\left[\mathrm{Cl}^{-}\right]\) in \((\mathrm{b})\) to be \(0.10 M\), calculate \(K_{\text {sp }}\) of \(\mathrm{AgCl}\).

An electrolytic cell produces aluminum from \(\mathrm{Al}_{2} \mathrm{O}_{3}\) at the rate of ten kilograms a day. Assuming a yield of \(100 \%\), (a) how many moles of electrons must pass through the cell in one day? (b) how many amperes are passing through the cell? (c) how many moles of oxygen \(\left(\mathrm{O}_{2}\right)\) are being produced simultaneously?

The electrolysis of an aqueous solution of \(\mathrm{NaCl}\) has the overall equation $$ 2 \mathrm{H}_{2} \mathrm{O}+2 \mathrm{Cl}^{-}(a q) \longrightarrow \mathrm{H}_{2}(g)+\mathrm{Cl}_{2}(g)+2 \mathrm{OH}^{-}(a q) $$ During the electrolysis, \(0.228\) mol of electrons passes through the cell. (a) How many electrons does this represent? (b) How many coulombs does this represent? (c) Assuming \(100 \%\) yield, what masses of \(\mathrm{H}_{2}\) and \(\mathrm{Cl}_{2}\) are produced?

A solution containing a metal ion \(\left(\mathrm{M}^{3+}(a q)\right)\) is electrolyzed by a current of \(5.0 \mathrm{~A}\). After \(10.0\) minutes, \(1.19 \mathrm{~g}\) of the metal is plated out. (a) How many coulombs are supplied by the battery? (b) What is the metal? (Assume \(100 \%\) efficiency.)

A solution containing a metal ion \(\left(M^{2+}(a q)\right)\) is electrolyzed by a current of \(7.8\) A. After \(15.5\) minutes, \(2.39 \mathrm{~g}\) of the metal is plated out. (a) How many coulombs are supplied by the battery? (b) What is the metal? (Assume \(100 \%\) efficiency.)

A baby's spoon with an area of \(6.25 \mathrm{~cm}^{2}\) is plated with silver from \(\mathrm{AgNO}_{3}\) using a current of \(2.00 \mathrm{~A}\) for two hours and 25 minutes. (a) If the current efficiency is \(82.0 \%\), how many grams of silver are plated? (b) What is the thickness of the silver plate formed \(\left(d=10.5 \mathrm{~g} / \mathrm{cm}^{3}\right)\) ?

Given the following data: $$ \begin{array}{cc} \mathrm{PtCl}_{4}{\underline{\phantom{xx}}}^{2-}(a q)+2 e^{-} \longrightarrow \mathrm{Pt}(s)+4 \mathrm{Cl}^{-}(a q) & E^{o}=+0.73 \mathrm{~V} \\ \mathrm{Pt}^{2+}(a q)+4 \mathrm{Cl}^{-}(a q) \longrightarrow \mathrm{PtCl}_{4}^{2-}(a q) & K_{f}=1 \times 10^{16} \end{array} $$ find \(E^{\circ}\) for the half-cell $$ \mathrm{Pt}^{2+}(a q)+2 e^{-} \longrightarrow \operatorname{Pt}(s) $$

In a nickel-cadmium battery (Nicad), cadmium is oxidized to \(\mathrm{Cd}(\mathrm{OH})_{2}\) at the anode, while \(\mathrm{Ni}_{2} \mathrm{O}_{3}\) is reduced to \(\mathrm{Ni}(\mathrm{OH})_{2}\) at the cathode. A portable CD player uses \(0.175\) amp of current. How many grams of \(\mathrm{Cd}\) and \(\mathrm{Ni}_{2} \mathrm{O}_{3}\) are consumed when the CD player is used for an hour and a half?

Consider the electrolysis of \(\mathrm{CuCl}_{2}\) to form \(\mathrm{Cu}(s)\) and \(\mathrm{Cl}_{2}(\mathrm{~g}) .\) Calculate the minimum voltage required to carry out this reaction at standard conditions. If a voltage of \(1.50 \mathrm{~V}\) is actually used, how many kilojoules of electrical energy are consumed in producing \(2.00 \mathrm{~g}\) of \(\mathrm{Cu} ?\)

An electrolytic cell consists of a \(100.0-\mathrm{g}\) strip of copper in \(0.200 \mathrm{M}\) \(\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}\) and a \(100.0-\mathrm{g}\) strip of \(\mathrm{Cr}\) in \(0.200 \mathrm{M} \mathrm{Cr}\left(\mathrm{NO}_{3}\right)_{3}\). The overall reac- tion is: $$ 3 \mathrm{Cu}(s)+2 \mathrm{Cr}^{3+}(a q) \longrightarrow 3 \mathrm{Cu}^{2+}(a q)+2 \mathrm{Cr}(s) \quad E^{\circ}=-1.083 \mathrm{~V} $$ An external battery provides 3 amperes for 70 minutes and 20 seconds with \(100 \%\) efficiency. What is the mass of the copper strip after the battery has been disconnected?

Consider the following reaction carried out at \(1000^{\circ} \mathrm{C}\). $$ \mathrm{CO}(g)+\frac{1}{2} \mathrm{O}_{2}(g) \longrightarrow \mathrm{CO}_{2}(g) $$ Assuming that all gases are at \(1.00 \mathrm{~atm}\), calculate the voltage produced at the given conditions. (Use Appendix 1 and assume that \(\Delta H^{\circ}\) and \(\Delta S^{\circ}\) do not change with an increase in temperature.)

Atomic masses can be determined by electrolysis. In one hour, a current of \(0.600\) A deposits \(2.42 \mathrm{~g}\) of a certain metal, \(\mathrm{M}\), which is present in solution as \(\mathrm{M}^{+}\) ions. What is the atomic mass of the metal?

Consider the following reaction at \(25^{\circ} \mathrm{C}\). $$ \mathrm{O}_{2}(g)+4 \mathrm{H}^{+}(a q)+4 \mathrm{Br}^{-}(a q) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}+2 \mathrm{Br}_{2}(l) $$ If \(\left[\mathrm{H}^{+}\right]\) is adjusted by adding a buffer that is \(0.100 M\) in sodium acetate and \(0.100 \mathrm{M}\) in acetic acid, the pressure of oxygen gas is \(1.00 \mathrm{~atm}\), and the bromide concentration is \(0.100 \mathrm{M}\), what is the calculated cell voltage? ( \(K_{\mathrm{a}}\) acetic acid \(\left.=1.8 \times 10^{-5} .\right)\).

For the cell $$ \mathrm{Zn}\left|\mathrm{Zn}^{2+}\right| \mathrm{Cu}^{2+} \mid \mathrm{Cu} $$ \(E^{\circ}\) is \(1.10 \mathrm{~V}\). A student prepared the same cell in the lab at standard conditions. Her experimental \(E^{\circ}\) was \(1.0 \mathrm{~V}\). A possible explanation for the difference is that (a) a larger volume of \(\mathrm{Zn}^{2+}\) than \(\mathrm{Cu}^{2+}\) was used. (b) the zinc electrode had twice the mass of the copper electrode. (c) \(\left[\mathrm{Zn}^{2+}\right]\) was smaller than \(1 M\). (d) \(\left[\mathrm{Cu}^{2+}\right]\) was smaller than \(1 M\). (e) the copper electrode had twice the surface area of the zinc electrode.

Which of the changes below will increase the voltage of the following cell? $$ \text { Co }\left|\mathrm{Co}^{2+}(0.010 M) \| \mathrm{H}^{+}(0.010 \mathrm{M})\right| \mathrm{H}_{2}(0.500 \mathrm{~atm}) \mid \mathrm{Pt} $$ (a) Increase the volume of \(\mathrm{CoCl}_{2}\) solution from \(100 \mathrm{~mL}\) to \(300 \mathrm{~mL}\). (b) Increase \(\left[\mathrm{H}^{+}\right]\) from \(0.010 \mathrm{M}\) to \(0.500 \mathrm{M}\) (c) Increase the pressure of \(\mathrm{H}_{2}\) from \(0.500 \mathrm{~atm}\) to \(1 \mathrm{~atm}\). (d) Increase the mass of the Co electrode from \(15 \mathrm{~g}\) to \(25 \mathrm{~g}\). (e) Increase \(\left[\mathrm{Co}^{2+}\right]\) from \(0.010 \mathrm{M}\) to \(0.500 \mathrm{M}\).

Consider the following standard reduction potentials: $$ \begin{array}{ll} \mathrm{Tl}^{+}(a q)+e^{-} \longrightarrow \mathrm{Tl}(s) & E_{\mathrm{red}}^{o}=-0.34 \mathrm{~V} \\ \mathrm{~T}^{3+}(a q)+3 e^{-} \longrightarrow \mathrm{T} 1(s) & E_{\mathrm{red}}^{\circ}=0.74 \mathrm{~V} \\ \mathrm{~T}^{3+}(a q)+2 e^{-} \longrightarrow \mathrm{T} 1^{+}(a q) & E_{\mathrm{red}}^{\circ}=1.28 \mathrm{~V} \end{array} $$ and the following abbreviated cell notations: $$ \text { (1) } \left.\mathrm{T} 1 \mid \mathrm{T}]^{+} \| \mathrm{T}^{3+}, \mathrm{T}\right]^{+} \mid \mathrm{Pt} $$ $$ \begin{aligned} &\text { (2) } \mathrm{Tl}\left|\mathrm{T}^{3+} \| \mathrm{T}^{3+}, \mathrm{Tl}^{+}\right| \mathrm{Pt} \\ &\text { (3) } \mathrm{Tl}\left|\mathrm{Tl}^{+} \| \mathrm{T}^{3+}\right| \mathrm{Tl} \end{aligned} $$ (a) Write the overall equation for each cell. (b) Calculate \(E^{\circ}\) for each cell. (c) Calculate \(\Delta G^{\circ}\) for each overall equation. (d) Comment on whether \(\Delta G^{\circ}\) and/or \(E^{\circ}\) are state properties. (Hint: \(\mathrm{A}\) state property is path-independent.)

Use Table \(18.1\) to answer the following questions. Use LT (for is less than), GT (for is greater than), EQ(for is equal to), or MI (for more information required). (a) For the half reaction (b) For the reaction (c) If the half reaction (d) For the reaction (e) For the reaction described in (d), the number of coulombs that passes through the cell is _____ \(9.648 \times 10^{4}\).

Consider three metals, \(\mathrm{X}, \mathrm{Y}\), and \(\mathrm{Z}\), and their salts, \(\mathrm{XA}, \mathrm{YA}\), and \(\mathrm{ZA}\). Three experiments take place with the following results: \- \(\mathrm{X}+\mathrm{hot} \mathrm{H}_{2} \mathrm{O} \longrightarrow \mathrm{H}_{2}\) bubbles \(\mathrm{X}+\mathrm{YA} \longrightarrow\) no reaction I \(\mathrm{X}+\mathrm{ZA} \longrightarrow \mathrm{X}\) discolored \(+\mathrm{Z}\) Rank metals \(\mathrm{X}, \mathrm{Y}\), and \(\mathrm{Z}\), in order of decreasing strength as reducing agents.

An alloy made up of tin and copper is prepared by simultaneously electroplating the two metals from a solution containing \(\mathrm{Sn}\left(\mathrm{NO}_{3}\right)_{2}\) and \(\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}\). If \(20.0 \%\) of the total current is used to plate tin, while \(80.0 \%\) is used to plate copper, what is the percent composition of the alloy?