Chapter 19

Write balanced equations for the following halfreactions. Specify whether each is an oxidation or reduction. (a) \(\mathrm{Cr}(\mathrm{s}) \rightarrow \mathrm{Cr}^{3+}(\text { aq })\) (in acid) (b) \(\mathrm{AsH}_{3}(\mathrm{g}) \rightarrow \mathrm{As}(\mathrm{s})\) (in acid) (c) \(\mathrm{VO}_{3}^{-}(\mathrm{aq}) \rightarrow \mathrm{V}^{2+}(\mathrm{aq})\) (in acid) (d) \(\mathrm{Ag}(\mathrm{s}) \rightarrow \mathrm{Ag}_{2} \mathrm{O}(\mathrm{s})\) (in base)

Write balanced equations for the following halfreactions. Specify whether each is an oxidation or reduction. (a) \(\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq}) \rightarrow \mathrm{O}_{2}(\mathrm{g})\) (in acid) (b) \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}(\mathrm{aq}) \rightarrow \mathrm{CO}_{2}(\mathrm{g})\) (in acid) (c) \(\mathrm{NO}_{3}^{-}(\mathrm{aq}) \rightarrow \mathrm{NO}(\mathrm{g})\) (in acid) (d) \(\mathrm{MnO}_{4}^{-}(\mathrm{aq}) \rightarrow \mathrm{MnO}_{2}(\mathrm{s})\) (in base)

Balance the following redox equations. All occur in acid solution. (a) \(\mathrm{Ag}(\mathrm{s})+\mathrm{NO}_{3}^{-}(\mathrm{aq}) \rightarrow \mathrm{NO}_{2}(\mathrm{g})+\mathrm{Ag}^{+}(\mathrm{aq})\) (b) \(\mathrm{MnO}_{4}^{-}(\mathrm{aq})+\mathrm{HSO}_{3}^{-}(\mathrm{aq}) \rightarrow\) \(\mathrm{Mn}^{2+}(\mathrm{aq})+\mathrm{SO}_{4}^{2-}(\mathrm{aq}\) (c) \(\mathrm{Zn}(\mathrm{s})+\mathrm{NO}_{3}^{-}(\mathrm{aq}) \rightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+\mathrm{N}_{2} \mathrm{O}(\mathrm{g})\) (d) \(\mathrm{Cr}(\mathrm{s})+\mathrm{NO}_{3}^{-}(\mathrm{aq}) \rightarrow \mathrm{Cr}^{3+}(\mathrm{aq})+\mathrm{NO}(\mathrm{g})\)

Balance the following redox equations. All occur in acid solution. (a) \(\operatorname{sn}(s)+H^{+}(a q) \rightarrow S n^{2+}(a q)+H_{2}(g)\) (b) \(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(\mathrm{aq})+\mathrm{Fe}^{2+}(\mathrm{aq}) \rightarrow\) \(\mathrm{Cr}^{3+}(\mathrm{aq})+\mathrm{Fe}^{3+}(\mathrm{aq})\) (c) \(\mathrm{MnO}_{2}(\mathrm{s})+\mathrm{Cl}^{-}(\mathrm{aq}) \rightarrow \mathrm{Mn}^{2+}(\mathrm{aq})+\mathrm{Cl}_{2}(\mathrm{g})\) (d) \(\mathrm{CH}_{2} \mathrm{O}(\mathrm{aq})+\mathrm{Ag}^{+}(\mathrm{aq}) \rightarrow \mathrm{HCO}_{2} \mathrm{H}(\mathrm{aq})+\mathrm{Ag}(\mathrm{s})\)

Balance the following redox equations. All occur in basic solution. (a) \(\mathrm{Al}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{Al}(\mathrm{OH})_{4}-(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{g})\) (b) \(\mathrm{CrO}_{4}^{2-}(\mathrm{aq})+\mathrm{SO}_{3}^{2-}(\mathrm{aq}) \rightarrow\) \(\mathrm{Cr}(\mathrm{OH})_{3}(\mathrm{s})+\mathrm{SO}_{4}^{2-}(\mathrm{aq})\) (c) \(\mathrm{Zn}(\mathrm{s})+\mathrm{Cu}(\mathrm{OH})_{2}(\mathrm{s}) \rightarrow\) \(\left[\mathrm{Zn}(\mathrm{OH})_{4}\right]^{2-}(\mathrm{aq})+\mathrm{Cu}(\mathrm{s})\) (d) \(\mathrm{HS}^{-}(\mathrm{aq})+\mathrm{ClO}_{3}^{-}(\mathrm{aq}) \rightarrow \mathrm{S}(\mathrm{s})+\mathrm{Cl}^{-}(\mathrm{aq})\)

Balance the following redox equations. All occur in basic solution. (a) \(\mathrm{Fe}(\mathrm{OH})_{3}(\mathrm{s})+\mathrm{Cr}(\mathrm{s}) \rightarrow\) \(\mathrm{Cr}(\mathrm{OH})_{3}(\mathrm{s})+\mathrm{Fe}(\mathrm{OH})_{2}(\mathrm{s})\) (b) \(\mathrm{NiO}_{2}(\mathrm{s})+\mathrm{Zn}(\mathrm{s}) \rightarrow\) \(\mathrm{Ni}(\mathrm{OH})_{2}(\mathrm{s})+\mathrm{Zn}(\mathrm{OH})_{2}(\mathrm{s})\) (c) \(\mathrm{Fe}(\mathrm{OH})_{2}(\mathrm{s})+\mathrm{CrO}_{4}^{2-}(\mathrm{aq}) \rightarrow\) \(\mathrm{Fe}(\mathrm{OH})_{3}(\mathrm{s})+\left[\mathrm{Cr}(\mathrm{OH})_{4}\right]^{-}(\mathrm{aq})\) (d) \(\mathrm{N}_{2} \mathrm{H}_{4}(\mathrm{aq})+\mathrm{Ag}_{2} \mathrm{O}(\mathrm{s}) \rightarrow \mathrm{N}_{2}(\mathrm{g})+\mathrm{Ag}(\mathrm{s})\)

A voltaic cell is constructed using the reaction \(\mathrm{Mg}(\mathrm{s})+2 \mathrm{H}^{+}(\mathrm{aq}) \rightarrow \mathrm{Mg}^{2+}(\mathrm{aq})+\mathrm{H}_{2}(\mathrm{g})\) (a) Write equations for the oxidation and reduction half-reactions. (b) Which half-reaction occurs in the anode compartment, and which occurs in the cathode compartment? (c) Complete the following sentences: Electrons in the external circuit flow from the electrode to the \(\quad\) electrode. Negative ions move in the salt bridge from the \- half-cell to the \(\quad\) half-cell. The half-reaction at the anode is \(=\) and that at the cathode is

The half-cells \(\mathrm{Fe}^{2+}(\text { aq }) | \mathrm{Fe}(\mathrm{s})\) and \(\mathrm{O}_{2}(\mathrm{g}) | \mathrm{H}_{2} \mathrm{O}\) (in acid solution) are linked to create a voltaic cell. (a) Write equations for the oxidation and reduction half-reactions and for the overall (cell) reaction. (b) Which half-reaction occurs in the anode compartment, and which occurs in the cathode compartment? (c) Complete the following sentences: Electrons in the external circuit flow from the electrode to the electrode. Negative ions move in the salt bridge from the \(-\) half-cell to the \(\quad\) half-cell.

The half-cells \(\operatorname{Sn}^{2+}(\text { aq }) | \operatorname{Sn}(s)\) and \(C l_{2}(g) | C l^{-}(a q)\) are linked to create a voltaic cell. (a) Write equations for the oxidation and reduction half-reactions and for the overall (cell) reaction. (b) Which half-reaction occurs in the anode compartment, and which occurs in the cathode compartment? (c) Complete the following sentences: Electrons in the external circuit flow from the electrode to the electrode. Negative ions move in the salt bridge from the \(-\) half-cell to the half-cell.

For each of the following electrochemical cells, write equations for the oxidation and reduction half-reactions and for the overall reaction. (a) \(\mathrm{Pb}(\mathrm{s})\left|\mathrm{Pb}^{2+}(\mathrm{aq}) \| \mathrm{Sn}^{4+}(\mathrm{aq}), \mathrm{Sn}^{2+}(\mathrm{aq})\right| \mathrm{C}(\mathrm{s})\) (b) \(\mathrm{Hg}(\ell)\left|\mathrm{Hg}_{2} \mathrm{Cl}_{2}(\mathrm{s})\right| \mathrm{Cl}^{-}(\mathrm{aq}) \| \mathrm{Ag}^{+}(\mathrm{aq}) | \mathrm{Ag}(\mathrm{s})\)

For each of the following electrochemical cells, write equations for the oxidation and reduction half-reactions and for the overall reaction. (a) \(\mathrm{Pb}(\mathrm{s})\left|\mathrm{Pb}^{2+}(\mathrm{aq}) \| \mathrm{Sn}^{4+}(\mathrm{aq}), \mathrm{Sn}^{2+}(\mathrm{aq})\right| \mathrm{C}(\mathrm{s})\) (b) \(\mathrm{Hg}(\ell)\left|\mathrm{Hg}_{2} \mathrm{Cl}_{2}(\mathrm{s})\right| \mathrm{Cl}^{-}(\mathrm{aq}) \| \mathrm{Ag}^{+}(\mathrm{aq}) | \mathrm{Ag}(\mathrm{s})\)

Use cell notation to depict an electrochemical cell based upon the following reaction that is productfavored at equilibrium. \(\mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{Ag}(\mathrm{s})+\mathrm{Cl}^{-}(\mathrm{aq}) \rightarrow \mathrm{Fe}^{2+}(\mathrm{aq})+\mathrm{AgCl}(\mathrm{s})\)

What are the similarities and differences between dry cells, alkaline batteries, and Ni-cad batteries?

What reactions occur when a lead storage battery is recharged?

Balance each of the following unbalanced equations; then calculate the standard potential, \(E^{\circ}\) and decide whether each is product-favored at equilibrium as written. (All reactions are carried out in acid solution. (a) \(\mathrm{I}_{2}(\mathrm{s})+\mathrm{Br}^{-}(\mathrm{aq}) \rightarrow \mathrm{I}^{-}(\mathrm{aq})+\mathrm{Br}_{2}(\ell)\) (b) \(\mathrm{Fe}^{2+}(\mathrm{aq})+\mathrm{Cu}^{2+}(\mathrm{aq}) \rightarrow \mathrm{Cu}(\mathrm{s})+\mathrm{Fe}^{3+}(\mathrm{aq})\) (c) \(\mathrm{Fe}^{2+}(\mathrm{aq})+\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(\mathrm{aq}) \rightarrow\) \(\mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{Cr}^{3+}(\mathrm{aq})\) (d) \(\mathrm{MnO}_{4}-(\mathrm{aq})+\mathrm{HNO}_{2}(\mathrm{aq}) \rightarrow\) \(\mathrm{Mn}^{2+}(\mathrm{aq})+\mathrm{NO}_{3}^{-}(\mathrm{aq})\)

Which of the following elements is the best reducing agent under standard conditions? (a) Cu (b) Zn (c) Fe (d) \(\mathrm{Ag}\) (e) \(\mathrm{Cr}\)

From the following list, identify those elements that are easier to oxidize than \(\mathrm{H}_{2}(\mathrm{g})\) (a) Cu (b) Zn (c) Fe (d) \(\mathrm{Ag}\) (e) \(\mathrm{Cr}\)

One half-cell in a voltaic cell is constructed from a copper wire electrode in a \(4.8 \times 10^{-3} \mathrm{M}\) solution of \(\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}\). The other half-cell consists of a zinc electrode in a \(0.40 \mathrm{M}\) solution of \(\mathrm{Zn}\left(\mathrm{NO}_{3}\right)_{2}\). Calculate the cell potential.

One half-cell in a voltaic cell is constructed from an iron electrode in an \(\mathrm{Fe}\left(\mathrm{NO}_{3}\right)_{2}\) solution of unknown concentration. The other half-cell is a standard hydrogen electrode. A potential of 0.49 V is measured for this cell. Use this information to calculate the concentration of \(\mathrm{Fe}^{2+}(\text { aq })\)

Which product, \(\mathrm{O}_{2}\) or \(\mathrm{F}_{2,}\) is more likely to form at the anode in the electrolysis of an aqueous solution of KF? Explain your reasoning.

Which product, Ca or \(\mathrm{H}_{2}\), is more likely to form at the cathode in the electrolysis of \(\mathrm{CaCl}_{2}\) ? Explain your reasoning.

An aqueous solution of KBr is placed in a beaker with two inert platinum electrodes. When the cell is attached to an external source of electrical energy, electrolysis occurs. (a) Hydrogen gas and hydroxide ion form at the cathode Write an equation for the halfreaction that occurs at this electrode. (b) Bromine is the primary product at the anode. Write an equation for its formation.

An aqueous solution of \(\mathrm{Na}_{2} \mathrm{S}\) is placed in a beaker with two inert platinum electrodes. When the cell is attached to an external battery, electrolysis OCCurs. (a) Hydrogen gas and hydroxide ion form at the cathode Write an equation for the halfreaction that occurs at this electrode. (b) Sulfur is the primary product at the anode. Write an equation for its formation.

In the electrolysis of a solution containing \(\mathrm{Ni}^{2+}(\mathrm{aq}),\) metallic \(\mathrm{Ni}(\mathrm{s})\) deposits on the cathode. Using a current of 0.150 A for 12.2 minutes, what mass of nickel will form?

In the electrolysis of a solution containing \(\mathrm{Ag}^{+}(\mathrm{aq}),\) metallic \(\mathrm{Ag}(\mathrm{s})\) deposits on the cathode Using a current of 1.12 A for 2.40 hours, what mass of silver forms?

Electrolysis of a solution of \(\mathrm{CuSO}_{4}\) (aq) to give copper metal is carried out using a current of 0.66 A. How long should electrolysis continue to produce \(0.50 \mathrm{g}\) of copper?

Electrolysis of a solution of \(\mathrm{Zn}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})\) to give zinc metal is carried out using a current of \(2.12 \mathrm{A}\) How long should electrolysis continue in order to prepare \(2.5 \mathrm{g}\) of zinc?

A voltaic cell can be built using the reaction between Al metal and \(\mathrm{O}_{2}\) from the air. If the Al anode of this cell consists of 84 g of aluminum, how many hours can the cell produce \(1.0 \mathrm{A}\) of electricity, assuming an unlimited supply of \(\mathbf{O}_{2} ?\)

Use \(E^{\circ}\) values to predict which of the following metals, if coated on iron, will provide cathodic protection against corrosion to iron. (a) Cu (b) \(\mathrm{Mg}\) (c) Ni (d) Sn

Use \(E^{\circ}\) values to predict which of the following metals, if coated on nickel, will provide cathodic protection against corrosion to nickel. (a) Cu (b) \(\mathrm{Mg}\) (c) Zn (d) \(\mathrm{Cr}\)

In the presence of oxgyen and acid, two half. reactions responsible for the corrosion of iron are $$ \begin{array}{c} \mathrm{Fe}(\mathrm{s}) \rightarrow \mathrm{Fe}^{2+}(\mathrm{aq})+2 e^{-} \\ \mathrm{O}_{2}(\mathrm{g})+4 \mathrm{H}^{+}(\mathrm{aq})+4 e^{-} \rightarrow 2 \mathrm{H}_{2} \mathrm{O}(\ell) \end{array} $$ Calculate the the standard potential, \(E^{\circ},\) and decide whether the reaction is product-favored at equilibrium. Will decreasing the pH make the reaction less thermodynamically product-favored at equilibrium?

Write balanced equations for the following half-reactions. (a) \(\mathrm{UO}_{2}^{+}(\mathrm{aq}) \rightarrow \mathrm{U}^{4+}(\mathrm{aq})\) (acid solution) (b) \(\mathrm{ClO}_{3}^{-}(\mathrm{aq}) \rightarrow \mathrm{Cl}^{-}(\mathrm{aq})\) (acid solution) (c) \(\mathrm{N}_{2} \mathrm{H}_{4}(\mathrm{aq}) \rightarrow \mathrm{N}_{2}(\mathrm{g})\) (basic solution) (d) \(\mathrm{ClO}^{-}(\mathrm{aq}) \rightarrow \mathrm{Cl}^{-}(\mathrm{aq})\) (basic solution)

S. Balance the following equations. (a) \(\mathrm{Zn}(\mathrm{s})+\mathrm{VO}^{2+}(\mathrm{aq}) \rightarrow\) \(\mathrm{Zn}^{2+}(\mathrm{aq})+\mathrm{V}^{3+}(\mathrm{aq}) \quad\) (acid solution) (b) \(\mathrm{Zn}(\mathrm{s})+\mathrm{VO}_{3}^{-}(\mathrm{aq}) \rightarrow\) \(\mathrm{V}^{2+}(\mathrm{aq})+\mathrm{Zn}^{2+}(\mathrm{aq}) \quad\) (acid solution) (c) \(\mathrm{Zn}(\mathrm{s})+\mathrm{ClO}^{-}(\mathrm{aq}) \rightarrow\) \(\mathrm{Zn}(\mathrm{OH})_{2}(\mathrm{s})+\mathrm{Cl}^{-}(\mathrm{aq}) \quad\) (basic solution) (d) CIO-(aq) \(+\left[\mathrm{Cr}(\mathrm{OH})_{4}\right]^{-}(\mathrm{aq}) \rightarrow\) \(\mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{CrO}_{4}^{2-}(\mathrm{aq})\) (basic solution)

Magnesium metal is oxidized, and silver ions are reduced in a voltaic cell using \(\mathrm{Mg}^{2+}(\mathrm{aq}, 1 \mathrm{M}) | \mathrm{Mg}\) and \(\mathrm{Ag}^{+}(\text {aq, } 1 \mathrm{M}) |\) Ag half-cells. (a) Label each part of the cell. (b) Write equations for the half-reactions occurring at the anode and the cathode, and write an equation for the net reaction in the cell. (c) Trace the movement of electrons in the exter. nal circuit. Assuming the salt bridge contains NaNO_, trace the movement of the Nat and \(\mathrm{NO}_{3}^{-}\) ions in the salt bridge that occurs when a voltaic cell produces current. Why is a salt bridge required in a cell?

You want to set up a series of voltaic cells with specific cell potentials. \(A\) Zn \(^{2+}(\text { aq, } 1.0 \mathrm{M}) | \mathrm{Zn}(\mathrm{s})\) half- cell is in one compartment. Identify several half-cells that you could use so that the cell potential will be close to (a) \(1.1 \mathrm{V}\) and (b) \(0.50 \mathrm{V}\). Consider cells in which the zinc cell can be either the cathode or the anode.

The reaction occurring in the cell in which \(\mathrm{Al}_{2} \mathrm{O}_{3}\) and aluminum salts are electrolyzed is \(\mathrm{Al}^{3+}(\mathrm{aq})+\) \(3 e^{-} \rightarrow \mathrm{Al}(\mathrm{s}) .\) If the electrolysis cell operates at \(5.0 \mathrm{V}\) and \(1.0 \times 10^{5} \mathrm{A},\) what mass of aluminum metal can be produced in a 24 -hour day?

A A cell is constructed using the following half-reactions: $$ \begin{array}{c} \mathrm{Ag}^{+}(\mathrm{aq})+e^{-} \rightarrow \mathrm{Ag}(\mathrm{s}) \\ \mathrm{Ag}_{2} \mathrm{SO}_{4}(\mathrm{s})+2 e^{-} \rightarrow 2 \mathrm{Ag}(\mathrm{s})+\mathrm{SO}_{4}^{2-}(\mathrm{aq}) \end{array} $$ \(E^{\bullet}=0.653 \mathrm{V}\) (a) What reactions should be observed at the anode and cathode? (b) Calculate the solubility product constant, \(K_{\mathrm{zp}}\) for \(\mathrm{Ag}_{2} \mathrm{SO}_{4}\)

A A potential of 0.142 V is recorded (under standard conditions) for a voltaic cell constructed using the following half reactions: Cathode: \(\mathrm{Pb}^{2+}(\mathrm{aq})+2 e^{-} \rightarrow \mathrm{Pb}(\mathrm{s})\) Anode: \(\quad \mathrm{PbCl}_{2}(\mathrm{s})+2 e^{-} \rightarrow \mathrm{Pb}(\mathrm{s})+2 \mathrm{Cl}^{-}(\mathrm{aq})\) Net: \(\quad \mathrm{Pb}^{2+}(\mathrm{aq})+2 \mathrm{Cl}^{-}(\mathrm{aq}) \rightarrow \mathrm{PbCl}_{2}(\mathrm{s})\) (a) What is the standard reduction potential for the anode reaction? (b) Calculate the solubility product, \(K_{\mathrm{rp}}\) for \(\mathrm{PbCl}_{2}\)

The standard potential, \(E^{\circ},\) for the reaction of \(\mathrm{Zn}(\mathrm{s})\) and \(\mathrm{C}_{2}(\mathrm{g})\) is \(+2.12 \mathrm{V}\). What is the standard free energy change, \(\Delta_{i} G^{\circ},\) for the reaction?

A An electrolysis cell for aluminum production operates at \(5.0 \mathrm{V}\) and a current of \(1.0 \times 10^{5} \mathrm{A}\) Calculate the number of kilowatt-hours of energy required to produce 1 metric ton \(\left(1.0 \times 10^{3} \mathrm{kg}\right)\) of aluminum. \(\left(1 \mathrm{kWh}=3.6 \times 10^{6} \mathrm{J} \text { and } 1 \mathrm{J}=\right.\) \(1 \mathbf{C} \cdot \mathbf{V}\)

A Electrolysis of molten NaCl is done in cells operating at \(7.0 \mathrm{V}\) and \(4.0 \times 10^{4} \mathrm{A}\). What mass of \(\mathrm{Na}(\mathrm{s})\) and \(\mathrm{Cl}_{2}(\mathrm{g})\) can be produced in 1 day in such a cell? What is the energy consumption in kilowatt-hours? ( \(1 \mathrm{kWh}=3.6 \times 10^{6} \mathrm{J}\) and \(1 \mathrm{J}=\) \(1 \mathrm{C} \cdot \mathrm{V}\)

Chlorine gas is obtained commercially by electrolysis of brine (a concentrated aqueous solution of NaCl). If the electrolysis cells operate at \(4.6 \mathrm{V}\) and \(3.0 \times 10^{5} \mathrm{A},\) what mass of chlorine can be produced in a 24 -hour day?

Write equations for the half-reactions that occur at the anode and cathode in the electrolysis of molten KBr. What are the products formed at the anode and cathode in the electrolysis of aqueous KBr?

The products formed in the electrolysis of aqueous \(\mathrm{CuSO}_{4}\) are \(\mathrm{Cu}(\mathrm{s})\) and \(\mathrm{O}_{2}(\mathrm{g}) .\) Write equations for the anode and cathode reactions.

Predict the products formed in the electrolysis of an aqueous solution of \(\mathrm{CdSO}_{4}\)

The metallurgy of aluminum involves electrolysis of \(\mathrm{Al}_{2} \mathrm{O}_{3}\) dissolved in molten cryolite \(\left(\mathrm{Na}_{3} \mathrm{AlF}_{6}\right)\) at about \(950^{\circ} \mathrm{C}\). Aluminum metal is produced at the cathode. Predict the anode product and write equations for the reactions occurring at both electrodes.

A voltaic cell set up utilizing the reaction $$ \mathrm{Cu}(\mathrm{s})+2 \mathrm{Ag}^{+}(\mathrm{aq}) \rightarrow \mathrm{Cu}^{2+}(\mathrm{aq})+2 \mathrm{Ag}(\mathrm{s}) $$ has a cell potential of 0.45 V at 298 K. Describe how the potential of this cell will change as the cell is discharged. At what point does the cell potential reach a constant value? Explain your answer.

Two \(\mathrm{Ag}^{+}(\mathrm{aq}) | \mathrm{Ag}(\mathrm{s})\) half-cells are constructed. The first has \(\left|\mathrm{Ag}^{+}\right|=1.0 \mathrm{M},\) the second has \(\left[\mathrm{Ag}^{+}\right]=\) \(1.0 \times 10^{-5} \mathrm{M} .\) When linked together with a salt bridge and external circuit, a cell potential is observed. (This kind of voltaic cell is referred to as a concentration cell.) (a) Draw a picture of this cell, labeling all components. Indicate the cathode and the anode, and indicate in which direction electrons flow in the external circuit. (b) Calculate the cell potential at \(298 \mathrm{K}\)

A Write balanced equations for the follow. ing reduction half-reactions involving organic compounds. (a) \(\mathrm{HCO}_{2} \mathrm{H} \rightarrow \mathrm{CH}_{2} \mathrm{O}\) (acid solution) (b) \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{CO}_{2} \mathrm{H} \rightarrow \mathrm{C}_{6} \mathrm{H}_{5} \mathrm{CH}_{3}\) (acid solution) (c) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CHO} \rightarrow\) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{OH}\) (acid solution) (d) \(\mathrm{CH}_{3} \mathrm{OH} \rightarrow \mathrm{CH}_{4}\) (acid solution)

A Balance the following equations involving organic compounds. (a) \(\mathrm{Ag}^{+}(\mathrm{aq})+\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{CHO}(\mathrm{aq}) \rightarrow\) \(\mathrm{Ag}(\mathrm{s})+\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{CO}_{2} \mathrm{H}(\mathrm{aq})\) (acid solution) (b) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}+\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(\mathrm{aq}) \rightarrow\) \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}(\mathrm{aq})+\mathrm{Cr}^{3+}(\mathrm{aq})\) (acid solution)