• Home
  • Textbooks
  • Chemistry: The Central Science in SI Units, Global Edition
  • Electrochemistry

Chemistry: The Central Science in SI Units, Global Edition

Theodore L. Brown, Matthew W. Stoltzfus, Michael W. Lufaso

Chapter 20

Electrochemistry - all with Video Answers

Educators

Chapter Questions

01:17

Problem 1

In the Brønsted-Lowry concept of acids and bases, acidbase reactions are viewed as proton-transfer reactions. The stronger the acid, the weaker is its conjugate base. If we were to think of redox reactions in a similar way, what particle would be analogous to the proton? Would strong oxidizing agents be analogous to strong acids or strong bases?

Nicole Smina
Nicole Smina
Numerade Educator
01:02

Problem 2

You may have heard that "antioxidants" are good for your health. Is an "antioxidant" an oxidizing agent or a reducing agent?

Nicole Smina
Nicole Smina
Numerade Educator
01:03

Problem 3

The diagram that follows represents a molecular view of a process occurring at an electrode in a voltaic cell. (a) Does the process represent oxidation or reduction? (b) Is the electrode the anode or cathode? (c) Why are the atoms in the electrode represented by larger spheres than those in the solution?

Nicole Smina
Nicole Smina
Numerade Educator
02:21

Problem 4

Assume that you want to construct a voltaic cell that uses the following half-reactions:
$$
\begin{array}{ll}
\mathrm{A}^{2+}(a q)+2 \mathrm{e}^{-} \longrightarrow \mathrm{A}(s) & E_{\mathrm{red}}^{\circ}=-0.10 \mathrm{~V} \\
\mathrm{~B}^{2+}(a q)+2 \mathrm{e}^{-} \longrightarrow \mathrm{B}(s) & E_{\mathrm{red}}^{\mathrm{o}}=-1.10 \mathrm{~V}
\end{array}
$$
You begin with the incomplete cell pictured here in which the electrodes are immersed in water. (a) What additions must you make to the cell for it to generate a standard emf? (b) Which electrode functions as the cathode? (c) Which direction do electrons move through the external circuit? (d) What voltage will the cell generate under standard conditions?

Nicole Smina
Nicole Smina
Numerade Educator
01:35

Problem 5

For a spontaneous reaction $\mathrm{A}(a q)+\mathrm{B}(a q) \longrightarrow \mathrm{A}^{-}(a q)+$
$\mathrm{B}^{+}(a q),$ answer the following questions:
(a) If you made a voltaic cell out of this reaction, what halfreaction would be occurring at the cathode, and what half reaction would be occurring at the anode?
(b) Which half-reaction from (a) is higher in potential energy?
(c) What is the sign of $E_{\text {cell }}^{\circ}$ ?

Nicole Smina
Nicole Smina
Numerade Educator
02:41

Problem 6

Consider the following table of standard electrode potentials for a series of hypothetical reactions in aqueous solution:
$$
\begin{array}{lr}
\hline \text { Reduction Half-Reaction } & {E^{\circ}(\mathrm{V})} \\
\hline \mathrm{A}^{+}(a q)+\mathrm{e}^{-} \longrightarrow \mathrm{A}(s) & 1.33 \\
\mathrm{~B}^{2+}(a q)+2 \mathrm{e}^{-} \longrightarrow \mathrm{B}(s) & 0.87 \\
\mathrm{C}^{3+}(a q)+\mathrm{e}^{-} \longrightarrow \mathrm{C}^{2+}(a q) & -0.12 \\
\mathrm{D}^{3+}(a q)+3 \mathrm{e}^{-} \longrightarrow \mathrm{D}(s) & -1.59 \\
\hline
\end{array}
$$
(a) Which substance is the strongest oxidizing agent? Which is weakest?
(b) Which substance is the strongest reducing agent? Which is weakest?
(c) Which substance(s) can oxidize $\mathrm{C}^{2+} ?$

Nicole Smina
Nicole Smina
Numerade Educator
01:02

Problem 7

Consider a redox reaction for which $E^{\circ}$ is a negative number.
(a) What is the sign of $\Delta G^{\circ}$ for the reaction?
(b) Will the equilibrium constant for the reaction be larger or smaller than $1 ?$
(c) Can an electrochemical cell based on this reaction accomplish work on its surroundings?

Nicole Smina
Nicole Smina
Numerade Educator
13:11

Problem 8

Consider the following voltaic cell:
(a) Which electrode is the cathode?
(b) What is the standard emf generated by this cell?
(c) What is the change in the cell voltage when the ion concentrations in the cathode half-cell are increased by a factor of $10 ?$
(d) What is the change in the cell voltage when the ion concentrations in the anode half-cell are increased by a factor of $10 ?$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
05:31

Problem 9

Consider the half-reaction $\mathrm{Ag}^{+}(a q)+\mathrm{e}^{-} \longrightarrow \mathrm{Ag}(s)$ (a) Which of the lines in the following diagram indicates how the reduction potential varies with the concentration of $\mathrm{Ag}^{+}(a q) ?(\mathbf{b})$ What is the value of $E_{\text {red }}$ when $\log \left[\mathrm{Ag}^{+}\right]=0 ?$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
05:26

Problem 10

The electrodes in a silver oxide battery are silver oxide $\left(\mathrm{Ag}_{2} \mathrm{O}\right)$ and zinc. (a) Which electrode acts as the anode? (b) Which battery do you think has an energy density most similar to the silver oxide battery: a Li-ion battery, a nickelcadmium battery, or a lead-acid battery?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
01:57

Problem 11

Bars of iron are put into each of the three beakers as shown here. In which beaker-A, B, or C-would you expect the iron to show the most corrosion?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
View

Problem 12

Magnesium, the element, is produced commercially by electrolysis from a molten salt (the "electrolyte") using a cell similar to the one shown here. (a) What is the most common oxidation number for Mg when it is part of a salt? (b) Chlorine gas is evolved as voltage is applied in the cell. Knowing this, identify the electrolyte. (c) Recall that in an electrolytic cell the anode is given the + sign and the cathode is given the $-\operatorname{sign}$, which is the opposite of what we see in batteries. What half-reaction occurs at the anode in this electrolytic cell? (d) What half-reaction occurs at the cathode?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:41

Problem 13

(a) What is meant by the term oxidation? (b) On which side of an oxidation half-reaction do the electrons appear? (c) What is meant by the term oxidant? (d) What is meant by the term oxidizing agent?

Jennifer Landry
Jennifer Landry
Numerade Educator
02:14

Problem 14

(a) What is meant by the term reduction? (b) On which side of a reduction half-reaction do the electrons appear? (c) What is meant by the term reductant? (d) What is meant by the term reducing agent?

Keenan Mintz
Keenan Mintz
University of Miami
02:40

Problem 15

Indicate whether each of the following statements is true or false:
(a) If something is oxidized, it is formally losing electrons.
(b) For the reaction $\mathrm{Fe}^{3+}(a q)+\mathrm{Co}^{2+}(a q) \longrightarrow \mathrm{Fe}^{2+}(a q)+$ $\mathrm{Co}^{3+}(a q), \mathrm{Fe}^{3+}(a q)$ is the reducing agent and $\mathrm{Co}^{2+}(a q)$ is the oxidizing agent.
(c) If there are no changes in the oxidation state of the reactants or products of a particular reaction, that reaction is not a redox reaction.

Jennifer Landry
Jennifer Landry
Numerade Educator
01:48

Problem 16

Indicate whether each of the following statements is true or false:
(a) If something is reduced, it is formally losing electrons.
(b) A reducing agent gets oxidized as it reacts.
(c) An oxidizing agent is needed to convert $\mathrm{CO}$ into $\mathrm{CO}_{2}$.

Keenan Mintz
Keenan Mintz
University of Miami
01:17

Problem 17

For each of the following balanced oxidation-reduction reactions, (i) identify the oxidation numbers for all the elements in the reactants and products and (ii) state the total number of electrons transferred in each reaction.
(a) $14 \mathrm{H}^{+}(a q)+2 \mathrm{Mn}^{2+}(a q)+5 \mathrm{NaBiO}_{3}(s)$ $\quad \longrightarrow 7 \mathrm{H}_{2} \mathrm{O}(l)+2 \mathrm{MnO}_{4}^{-}+5 \mathrm{Bi}^{3+}(a q)+5 \mathrm{Na}^{+}(a q)$
(b) $2 \mathrm{KMnO}_{4}(a q)+3 \mathrm{Na}_{2} \mathrm{SO}_{3}(a q)+\mathrm{H}_{2} \mathrm{O}(l)$ $\quad \longrightarrow 2 \mathrm{MnO}_{2}(s)+3 \mathrm{Na}_{2} \mathrm{SO}_{4}(a q)+2 \mathrm{KOH}(a q)$
(c) $\mathrm{Cu}(s)+2 \mathrm{AgNO}_{3}(a q) \longrightarrow 2 \mathrm{Ag (s)+\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}(a q)$

Nicole Smina
Nicole Smina
Numerade Educator
01:35

Problem 18

For each of the following balanced oxidation-reduction reactions, (i) identify the oxidation numbers for all the elements in the reactants and products and (ii) state the total number of electrons transferred in each reaction.
(a) $\mathrm{H}_{2}(g)+\mathrm{F}_{2}(g) \longrightarrow 2 \mathrm{HF}(g)$
(b) $2 \mathrm{Fe}^{2+}(a q)+\mathrm{H}_{2} \mathrm{O}_{2}(a q)+2 \mathrm{H}^{+}(a q) \longrightarrow 2 \mathrm{Fe}^{3+}(a q)+\mathrm{H}_{2} \mathrm{O}(l)$
(c) $\mathrm{CH}_{4}(g)+2 \mathrm{O}_{2}(g) \longrightarrow \mathrm{CO}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(l)$

Nicole Smina
Nicole Smina
Numerade Educator
01:22

Problem 19

Indicate whether the following balanced equations involve oxidation-reduction. If they do, identify the elements that undergo changes in oxidation number.
(a) $2 \mathrm{NO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}_{2}(a q) \longrightarrow 2 \mathrm{HNO}_{3}(a q)$
(b) $\mathrm{FeS}(s)+2 \mathrm{HCl}(a q) \longrightarrow \mathrm{FeCl}_{2}(a q)+\mathrm{H}_{2} \mathrm{~S}(g)$
(c) $\mathrm{Fe}(s)+2 \mathrm{HNO}_{3}(a q) \longrightarrow \mathrm{H}_{2} \mathrm{O}(l)+ 2 \mathrm{NO}_{2}(g)+\mathrm{FeO}(s)$

Nicole Smina
Nicole Smina
Numerade Educator
04:22

Problem 20

Indicate whether the following balanced equations involve oxidation-reduction. If they do, identify the elements that undergo changes in oxidation number.
(a) $2 \mathrm{AgNO}_{3}(a q)+\mathrm{CoCl}_{2}(a q) \longrightarrow 2 \mathrm{AgCl}(s)+ \mathrm{Co}\left(\mathrm{NO}_{3}\right)_{2}(a q)$
(b) $2 \mathrm{PbO}_{2}(s) \longrightarrow 2 \mathrm{PbO}(s)+\mathrm{O}_{2}(g)$
(c) $2 \mathrm{H}_{2} \mathrm{SO}_{4}(a q)+2 \mathrm{NaBr}(s) \longrightarrow \mathrm{Br}_{2}(l)+\mathrm{SO}_{2}(g)+ \mathrm{Na}_{2} \mathrm{SO}_{4}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l)$

Keenan Mintz
Keenan Mintz
University of Miami
01:39

Problem 21

The purification process of silicon involves the reaction of silicon tetrachloride vapor $\left(\mathrm{SiCl}_{4}(g)\right)$ with hydrogen to $1250^{\circ} \mathrm{C}$ to form solid silicon and hydrogen chloride. $(\mathbf{a})$ Write a balanced equation for this reaction. (b) What is being oxidized, and what is being reduced? (c) Which substance is the reductant, and which is the oxidant?

Nicole Smina
Nicole Smina
Numerade Educator
03:44

Problem 22

Hydrazine $\left(\mathrm{N}_{2} \mathrm{H}_{4}\right)$ and dinitrogen tetroxide $\left(\mathrm{N}_{2} \mathrm{O}_{4}\right)$ form a self-igniting mixture that has been used as a rocket propellant. The reaction products are $\mathrm{N}_{2}$ and $\mathrm{H}_{2} \mathrm{O}$. (a) Write a balanced chemical equation for this reaction. (b) What is being oxidized, and what is being reduced? (c) Which substance serves as the reducing agent and which as the oxidizing agent?

Keenan Mintz
Keenan Mintz
University of Miami
11:27

Problem 23

Complete and balance the following half-reactions. In each case, indicate whether the half-reaction is an oxidation or a reduction.
(a) $\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q) \longrightarrow \mathrm{Cr}^{3+}(a q)$ (acidic solution)
(b) $\mathrm{Mn}^{2+}(a q) \longrightarrow \mathrm{MnO}_{4}^{-}(a q)$ (acidic solution)
(c) $\mathrm{I}_{2}(s) \longrightarrow \mathrm{IO}_{3}^{-}(a q)$ (acidic solution)
(d) $\mathrm{S}(s)(a q) \longrightarrow \mathrm{H}_{2} \mathrm{~S}(g)$ (acidic solution)
(e) $\mathrm{NO}_{3}^{-}(a q) \longrightarrow \mathrm{NO}_{2}^{-}(a q)$ (basic solution)
(f) $\mathrm{H}_{2} \mathrm{O}_{2}(a q) \longrightarrow \mathrm{OH}^{-}(a q)$ (basic solution)

Susan Hallstrom
Susan Hallstrom
Numerade Educator
05:43

Problem 24

Complete and balance the following half-reactions. In each case indicate whether the half-reaction is an oxidation or a reduction.
(a) $\mathrm{Mo}^{3+}(a q) \longrightarrow \mathrm{Mo}(s)$ (acidic solution)
(b) $\mathrm{H}_{2} \mathrm{SO}_{3}(a q) \longrightarrow \mathrm{SO}_{4}^{2-}(a q)$ (acidic solution)
(c) $\mathrm{NO}_{3}^{-}(a q) \longrightarrow \mathrm{NO}(g)$ (acidic solution)
(d) $\mathrm{O}_{2}(g) \longrightarrow \mathrm{H}_{2} \mathrm{O}(l)$ (acidic solution)
(e) $\mathrm{O}_{2}(g) \longrightarrow \mathrm{H}_{2} \mathrm{O}(l)$ (basic solution)
(f) $\mathrm{Mn}^{2+}(a q) \longrightarrow \mathrm{MnO}_{2}(s)$ (basic solution)
(g) $\mathrm{Cr}(\mathrm{OH})_{3}(s) \longrightarrow \mathrm{CrO}_{4}^{2-}(a q)$ (basic solution)

Keenan Mintz
Keenan Mintz
University of Miami
07:17

Problem 25

Complete and balance the following equations, and identify the oxidizing and reducing agents:
(a) $\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{I}^{-}(a q) \longrightarrow \mathrm{Cr}^{3+}(a q)+\mathrm{IO}_{3}^{-}(a q)$ (acidic solution)
(b) $\mathrm{MnO}_{4}^{-}(a q)+\mathrm{CH}_{3} \mathrm{OH}(a q) \longrightarrow \mathrm{Mn}^{2+}(a q)+$
$\mathrm{HCOOH}(a q)$ (acidic solution)
(c) $\mathrm{I}_{2}(s)+\mathrm{OCl}^{-}(a q) \longrightarrow \mathrm{IO}_{3}^{-}(a q)+\mathrm{Cl}^{-}(a q)$ (acidic solution)
(d) $\mathrm{As}_{2} \mathrm{O}_{3}(s)+\mathrm{NO}_{3}^{-}(a q) \longrightarrow \mathrm{H}_{3} \mathrm{AsO}_{4}(a q)+\mathrm{N}_{2} \mathrm{O}_{3}(a q)$ (acidic solution)
(e) $\mathrm{MnO}_{4}^{-}(a q)+\mathrm{Br}^{-}(a q) \longrightarrow \mathrm{MnO}_{2}(s)+\operatorname{BrO}_{3}^{-}(a q)$ (basic solution)
(f) $\mathrm{Pb}(\mathrm{OH})_{4}^{2-}(a q)+\mathrm{ClO}^{-}(a q) \longrightarrow \mathrm{PbO}_{2}(s)+\mathrm{Cl}^{-}(a q)$ (basic solution)

Lottie Adams
Lottie Adams
Numerade Educator
25:08

Problem 26

Complete and balance the following equations, and identify the oxidizing and reducing agents. (Recall that the $\mathrm{O}$ atoms in hydrogen peroxide, $\mathrm{H}_{2} \mathrm{O}_{2}$, have an atypical oxidation state.)
(a) $\mathrm{NO}_{2}^{-}(a q)+\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q) \longrightarrow \mathrm{Cr}^{3+}(a q)+\mathrm{NO}_{3}^{-}(a q)$ (acidic solution)
(b) $\mathrm{S}(s)+\mathrm{HNO}_{3}(a q) \longrightarrow \mathrm{H}_{2} \mathrm{SO}_{3}(a q)+\mathrm{N}_{2} \mathrm{O}(g)$ (acidic solution)
(c) $\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{CH}_{3} \mathrm{OH}(a q) \longrightarrow \mathrm{HCOOH}(a q)+ \mathrm{Cr}^{3+}(a q)$ (acidic solution)
(d) $\mathrm{BrO}_{3}^{-}(a q)+\mathrm{N}_{2} \mathrm{H}_{4}(g) \longrightarrow \mathrm{Br}^{-}(a q)+\mathrm{N}_{2}(g)$ (acidic solution)
(e) $\mathrm{NO}_{2}^{-}(a q)+\mathrm{Al}(s) \longrightarrow \mathrm{NH}_{4}^{+}(a q)+\mathrm{AlO}_{2}^{-}(a q)$ (basic solution)
(f) $\mathrm{H}_{2} \mathrm{O}_{2}(a q)+\mathrm{ClO}_{2}(a q) \longrightarrow \mathrm{ClO}_{2}^{-}(a q)+\mathrm{O}_{2}(g)$ (basic solution)

Keenan Mintz
Keenan Mintz
University of Miami
02:05

Problem 27

Indicate whether each statement is true or false: (a) The cathode is the electrode at which oxidation takes place. (b) A galvanic cell is another name for a voltaic cell. (c) Electrons flow spontaneously from anode to cathode in a voltaic cell.

Jennifer Landry
Jennifer Landry
Numerade Educator
01:40

Problem 28

Indicate whether each statement is true or false: (a) The anode is the electrode at which oxidation takes place. (b) A voltaic cell always has a positive emf. (c) A salt bridge or permeable barrier is necessary to allow a voltaic cell to operate.

Keenan Mintz
Keenan Mintz
University of Miami
01:54

Problem 29

A voltaic cell similar to that shown in Figure 20.5 is constructed. One electrode half-cell consists of a magnesium strip placed in a solution of $\mathrm{MgCl}_{2}$, and the other has a nickel strip placed in a solution of $\mathrm{NiCl}_{2}$. The overall cell reaction is
$$
\mathrm{Mg}(s)+\mathrm{Ni}^{2+}(a q) \longrightarrow \mathrm{Ni}(s)+\mathrm{Mg}^{2+}(a q)
$$
(a) What is being oxidized, and what is being reduced? (b) Write the half-reactions that occur in the two half-cells. (c) Which electrode is the anode, and which is the cathode?(d) Indicate the signs of the electrodes. (e) Do electrons flow
from the magnesium electrode to the nickel electrode or from the nickel to the magnesium? (f) In which directions do the cations and anions migrate through the solution?

Nicole Smina
Nicole Smina
Numerade Educator
05:47

Problem 30

A voltaic cell similar to that shown in Figure 20.5 is constructed. One half-cell consists of an iron strip placed in a solution of $\mathrm{FeSO}_{4}$, and the other has an aluminum strip placed in a solution of $\mathrm{Al}_{2}\left(\mathrm{SO}_{4}\right)_{3} .$ The overall cell reaction is
$$
2 \mathrm{Al}(s)+3 \mathrm{Fe}^{2+}(a q) \longrightarrow 3 \mathrm{Fe}(s)+2 \mathrm{Al}^{3+}(a q)
$$
(a) What is being oxidized, and what is being reduced? (b) Write the half-reactions that occur in the two half-cells. (c) Which electrode is the anode, and which is the cathode? (d) Indicate the signs of the electrodes. (e) Do electrons flow from the aluminum electrode to the iron electrode or from the iron to the aluminum? (f) In which directions do the cations and anions migrate through the solution? Assume the $\mathrm{Al}$ is not coated with its oxide.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:27

Problem 31

(a) What is the definition of the volt? (b) Do all voltaic cells produce a positive cell potential?

Jennifer Landry
Jennifer Landry
Numerade Educator
01:34

Problem 32

(a) Which electrode of a voltaic cell, the cathode or the anode, corresponds to the higher potential energy for the electrons? (b) What are the units for electrical potential? How does this unit relate to energy expressed in joules?

Keenan Mintz
Keenan Mintz
University of Miami
04:18

Problem 33

(a) Write the half-reaction that occurs at an oxygen electrode in acidic aqueous solution when it serves as the cathode of a voltaic cell. (b) Write the half-reaction that occurs at an oxygen electrode in acidic aqueous solution when it serves as the anode of a voltaic cell. (c) What is standard about the standard oxygen electrode?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:23

Problem 34

(a) What conditions must be met for a reduction potential to be a standard reduction potential? (b) What is the standard reduction potential of a standard hydrogen electrode? (c) Why is it impossible to measure the standard reduction potential of a single half-reaction?

Keenan Mintz
Keenan Mintz
University of Miami
06:46

Problem 35

A voltaic cell that uses the reaction
$$
\mathrm{Tl}^{3+}(a q)+2 \mathrm{Cr}^{2+}(a q) \longrightarrow \mathrm{Tl}^{+}(a q)+2 \mathrm{Cr}^{3+}(a q)
$$
has a measured standard cell potential of $+1.19 \mathrm{~V} .(\mathbf{a})$ Write the two half-cell reactions. (b) By using data from Appendix E, determine $E_{\text {red }}^{\circ}$ for the reduction of $\mathrm{Tl}^{3+}(a q)$ to $\mathrm{Tl}^{+}(a q) .$ (c) Sketch the voltaic cell, label the anode and cathode, and indicate the direction of electron flow.

Shubham Kumar
Shubham Kumar
Numerade Educator
05:20

Problem 36

A voltaic cell that uses the reaction
$$
\mathrm{PdCl}_{4}^{2-}(a q)+\mathrm{Cd}(s) \longrightarrow \mathrm{Pd}(s)+4 \mathrm{Cl}^{-}(a q)+\mathrm{Cd}^{2+}(a q)
$$
has a measured standard cell potential of $+1.03 \mathrm{~V} .(\mathbf{a})$ Write the two half-cell reactions. $(\mathbf{b})$ By using data from Appendix E, determine $E_{\text {red }}^{\circ}$ for the reaction involving Pd. (c) Sketch the voltaic cell, label the anode and cathode, and indicate the direction of electron flow.

Keenan Mintz
Keenan Mintz
University of Miami
04:44

Problem 37

Using standard reduction potentials (Appendix E), calculate the standard emf for each of the following reactions:
(a) $\mathrm{Cl}_{2}(g)+2 \mathrm{I}^{-}(a q) \longrightarrow 2 \mathrm{Cl}^{-}(a q)+\mathrm{I}_{2}(s)$
(b) $\mathrm{Ni}(s)+2 \mathrm{Ce}^{4+}(a q) \longrightarrow \mathrm{Ni}^{2+}(a q)+2 \mathrm{Ce}^{3+}(a q)$
(c) $\mathrm{Fe}(s)+2 \mathrm{Fe}^{3+}(a q) \longrightarrow 3 \mathrm{Fe}^{2+}(a q)$
(d) $2 \mathrm{NO}_{3}^{-}(a q)+8 \mathrm{H}^{+}(a q)+3 \mathrm{Cu}(s) \longrightarrow 2 \mathrm{NO}(g)+ 4 \mathrm{H}_{2} \mathrm{O}(l)+3 \mathrm{Cu}^{2+}(a q)$

Jennifer Landry
Jennifer Landry
Numerade Educator
08:31

Problem 38

Using data in Appendix E, calculate the standard emf for each of the following reactions:
(a) $\mathrm{H}_{2}(g)+\mathrm{F}_{2}(g) \longrightarrow 2 \mathrm{H}^{+}(a q)+2 \mathrm{~F}^{-}(a q)$
(b) $\mathrm{Cu}^{2+}(a q)+\mathrm{Ca}(s) \longrightarrow \mathrm{Cu}(s)+\mathrm{Ca}^{2+}(a q)$
(c) $3 \mathrm{Fe}^{2+}(a q) \longrightarrow \mathrm{Fe}(s)+2 \mathrm{Fe}^{3+}(a q)$
(d) $2 \mathrm{ClO}_{3}^{-}(a q)+10 \mathrm{Br}^{-}(a q)+12 \mathrm{H}^{+}(a q) \longrightarrow \mathrm{Cl}_{2}(g)+ 5 \mathrm{Br}_{2}(l)+6 \mathrm{H}_{2} \mathrm{O}(l)$

Keenan Mintz
Keenan Mintz
University of Miami
03:29

Problem 39

The standard reduction potentials of the following half-reactions are given in Appendix E:
$$
\begin{array}{l}
\mathrm{Fe}^{2+}(a q)+2 \mathrm{e}^{-} \longrightarrow \mathrm{Fe}(s) \\
\mathrm{Cd}^{2+}(a q)+2 \mathrm{e}^{-} \longrightarrow \operatorname{Cd}(s) \\
\mathrm{Sn}^{2+}(a q)+2 \mathrm{e}^{-} \longrightarrow \operatorname{Sn}(s) \\
\mathrm{Ag}^{+}(a q)+\mathrm{e}^{-} \longrightarrow \operatorname{Ag}(s)
\end{array}
$$
(a) Determine which combination of these half-cell reactions leads to the cell reaction with the largest positive cell potential and calculate the value. (b) Determine which combination of these half-cell reactions leads to the cell reaction with the smallest positive cell potential and calculate the value.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
11:20

Problem 40

Given the following half-reactions and associated standard reduction potentials:
$$
\begin{aligned}
\mathrm{AuBr}_{4}^{-}(a q)+3 \mathrm{e}^{-} \longrightarrow \mathrm{Au}(s)+4 \mathrm{Br}^{-}(a q) & \\
E_{\mathrm{red}}^{\circ}=-0.86 \mathrm{~V} \\
\mathrm{Eu}^{3+}(a q)+\mathrm{e}^{-} \longrightarrow \mathrm{Eu}^{2+}(a q) & \\
E_{\mathrm{red}}^{\circ}=-0.43 \mathrm{~V} \\
\mathrm{IO}^{-}(a q)+\mathrm{H}_{2} \mathrm{O}(l)+2 \mathrm{e}^{-} \longrightarrow \mathrm{I}^{-}(a q)+2 \mathrm{OH}^{-}(a q) & \\
E_{\mathrm{red}}^{\circ}=&+0.49 \mathrm{~V}
\end{aligned}
$$
(a) Write the equation for the combination of these halfcell reactions that leads to the largest positive emf and calculate the value. (b) Write the equation for the combination of half-cell reactions that leads to the smallest positive emf and calculate that value.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
05:57

Problem 41

A $1 \mathrm{M}$ solution of $\mathrm{AgNO}_{3}$ is placed in a beaker with a strip of Ag metal. A $1 M$ solution of $\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}$ is placed in a second beaker with a strip of Cu metal. A salt bridge connects the two beakers, and wires to a voltmeter link the two metal electrodes. (a) Which electrode serves as the anode, and which as the cathode? (b) Which electrode gains mass, and which loses mass as the cell reaction proceeds? (c) Write the equation for the overall cell reaction. (d) What is the emf generated by the cell under standard conditions?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
03:52

Problem 42

A voltaic cell consists of a strip of cadmium metal in a solution of $\mathrm{Cd}\left(\mathrm{NO}_{3}\right)_{2}$ in one beaker, and in the other beaker a platinum electrode is immersed in a NaCl solution, with $\mathrm{Cl}_{2}$ gas bubbled around the electrode. A salt bridge connects the two beakers. (a) Which electrode serves as the anode, and which as the cathode? (b) Does the Cd electrode gain or lose mass as the cell reaction proceeds? (c) Write the equation for the overall cell reaction. (d) What is the emf generated by the cell under standard conditions?

Keenan Mintz
Keenan Mintz
University of Miami
04:18

Problem 43

From each of the following pairs of substances, use data in Appendix $\mathrm{E}$ to choose the one that is the stronger reducing agent:
(a) $\mathrm{Al}(s)$ or $\mathrm{Mg}(s)$
(b) $\mathrm{Fe}(s)$ or $\mathrm{Ni}(s)$
(c) $\mathrm{H}_{2}(g$, acidic solution) or $\operatorname{Sn}(s)$
(d) $\mathrm{I}^{-}(a q)$ or $\mathrm{Br}^{-}(a q)$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:24

Problem 44

From each of the following pairs of substances, use data in Appendix $\mathrm{E}$ to choose the one that is the stronger oxidizing agent:
(a) $\mathrm{Cl}_{2}(g)$ or $\mathrm{Br}_{2}(l)$
(b) $\mathrm{Zn}^{2+}(a q)$ or $\mathrm{Cd}^{2+}(a q)$
(c) $\mathrm{Cl}^{-}(a q)$ or $\mathrm{ClO}_{3}^{-}(a q)$
(d) $\mathrm{H}_{2} \mathrm{O}_{2}(a q)$ or $\mathrm{O}_{3}(g)$

Keenan Mintz
Keenan Mintz
University of Miami
03:19

Problem 45

By using the data in Appendix E, determine whether each of the following substances is likely to serve as an oxidant or a reductant: (a) $\mathrm{H}_{2}(g,$ basic $)$, (b) $\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q,$ acidic $)$ (c) $\mathrm{F}_{2}(g),(\mathbf{d}) \mathrm{Li}(s)$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
04:20

Problem 46

Is each of the following substances likely to serve as an oxidant or a reductant: $(\mathbf{a}) \mathrm{Ce}^{3+}(a q),(\mathbf{b}) \mathrm{Ca}(s),(\mathbf{c}) \mathrm{ClO}_{3}^{-}(a q),$ (d) $\mathrm{N}_{2} \mathrm{O}_{5}(g) ?$

Shubham Kumar
Shubham Kumar
Numerade Educator
07:09

Problem 47

(a) Assuming standard conditions, arrange the following in order of increasing strength as oxidizing agents in acidic solution: $\mathrm{MnO}_{4}^{-}(a q), \mathrm{O}_{3}(g), \mathrm{HSO}_{4}^{-}(a q), \mathrm{O}_{2}(g), \mathrm{HClO}(a q)$ (b) Arrange the following in order of increasing strength as reducing agents in basic solution: $\mathrm{Cr}(\mathrm{OH})_{3}(s), \mathrm{Fe}(s), \mathrm{Ca}(s),$ $\mathrm{H}_{2}(g), \mathrm{Mn}(s)$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
07:31

Problem 48

Based on the data in Appendix E, (a) which of the following is the strongest oxidizing agent, and which is the weakest in acidic solution: $\mathrm{Br}_{2}, \mathrm{H}_{2} \mathrm{O}_{2}, \mathrm{Zn}, \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-} ?(\mathbf{b})$ Which of the following is the strongest reducing agent, and which is the weakest in acidic solution: $\mathrm{F}^{-}, \mathrm{Zn}, \mathrm{N}_{2} \mathrm{H}_{5}^{+}, \mathrm{I}_{2}, \mathrm{NO}$ ?

Shubham Kumar
Shubham Kumar
Numerade Educator
03:22

Problem 49

The standard reduction potential of $\mathrm{Eu}^{2+}(a q)$ is $-0.43 \mathrm{~V}$. Using Appendix E, which of the following substances is capable of reducing $\mathrm{Eu}^{3+}(a q)$ to $\mathrm{Eu}^{2+}(a q)$ under standard conditions: $\mathrm{Al}, \mathrm{Co}, \mathrm{H}_{2} \mathrm{O}_{2}, \mathrm{~N}_{2} \mathrm{H}_{5}^{+}, \mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4} ?$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:49

Problem 50

The standard reduction potential for the reduction of $\mathrm{RuO}_{4}^{-}(a q)$ to $\mathrm{RuO}_{4}^{2-}(a q)$ is $+0.59 \mathrm{~V}$. By using Appendix $\mathrm{E}$ which of the following substances can oxidize $\mathrm{RuO}_{4}^{2-}(a q)$ to $\mathrm{RuO}_{4}^{-}(a q)$ under standard conditions: $\mathrm{Br}_{2}(l), \mathrm{BrO}_{3}^{-}(a q),$ $\mathrm{Mn}^{2+}(a q), \mathrm{O}_{2}(g), \mathrm{Sn}^{2+}(a q) ?$

Keenan Mintz
Keenan Mintz
University of Miami
11:00

Problem 51

Given the following reduction half-reactions:
$$
\begin{aligned}
\mathrm{Fe}^{3+}(a q)+\mathrm{e}^{-} \longrightarrow \mathrm{Fe}^{2+}(a q) & E_{\mathrm{red}}^{\circ}=+0.77 \mathrm{~V} \\
\mathrm{~S}_{2} \mathrm{O}_{6}^{2-}(a q)+4 \mathrm{H}^{+}(a q)+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{H}_{2} \mathrm{SO}_{3}(a q) & E_{\mathrm{red}}^{\circ}=+0.60 \mathrm{~V} \\
\mathrm{~N}_{2} \mathrm{O}(g)+2 \mathrm{H}^{+}(a q)+2 \mathrm{e}^{-} \longrightarrow \mathrm{N}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) & E_{\mathrm{red}}^{\circ}=-1.77 \mathrm{~V} \\
\mathrm{VO}_{2}^{+}(a q)+2 \mathrm{H}^{+}(a q)+\mathrm{e}^{-} \longrightarrow \mathrm{VO}^{2+}+\mathrm{H}_{2} \mathrm{O}(l) & E_{\mathrm{red}}^{\circ}=+1.00 \mathrm{~V}
\end{aligned}
$$
(a) Write balanced chemical equations for the oxidation of $\mathrm{Fe}^{2+}(a q)$ by $\mathrm{S}_{2} \mathrm{O}_{6}^{2-}(a q),$ by $\mathrm{N}_{2} \mathrm{O}(a q),$ and by $\mathrm{VO}_{2}^{+}(a q)$ (b) Calculate $\Delta G^{\circ}$ for each reaction at $298 \mathrm{~K}$. (c) Calculate the equilibrium constant $K$ for each reaction at $298 \mathrm{~K}$.

Jennifer Landry
Jennifer Landry
Numerade Educator
18:18

Problem 52

For each of the following reactions, write a balanced equation, calculate the standard emf, calculate $\Delta G^{\circ}$ at $298 \mathrm{~K},$ and calculate the equilibrium constant $K$ at $298 \mathrm{~K}$. (a) Aqueous iodide ion is oxidized to $\mathrm{I}_{2}(s)$ by $\mathrm{Hg}_{2}^{2+}(a q) .$ (b) In acidic solution, copper(I) ion is oxidized to copper(II) ion by nitrate ion. (c) In basic solution, $\mathrm{Cr}(\mathrm{OH})_{3}(s)$ is oxidized to $\mathrm{CrO}_{4}^{2-}(a q)$ by $\mathrm{ClO}^{-}(a q)$

Shubham Kumar
Shubham Kumar
Numerade Educator
03:35

Problem 53

If the equilibrium constant for a one-electron redox reaction at $298 \mathrm{~K}$ is $2.2 \times 10^{-5},$ calculate the corresponding $\Delta G^{\circ}$ and $E^{\circ}$.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:29

Problem 54

If the equilibrium constant for a two-electron redox reaction at $298 \mathrm{~K}$ is $2.2 \times 10^{5},$ calculate the corresponding $\Delta G^{\circ}$ and $E^{\circ}$.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
08:12

Problem 55

Using the standard reduction potentials listed in Appendix E, calculate the equilibrium constant for each of the following reactions at $298 \mathrm{~K}$ :
$$
\begin{array}{l}
\text { (a) } \mathrm{Fe}(s)+\mathrm{Ni}^{2+}(a q) \longrightarrow \mathrm{Fe}^{2+}(a q)+\mathrm{Ni}(s) \\
\text { (b) } \mathrm{Co}(s)+2 \mathrm{H}^{+}(a q) \longrightarrow \mathrm{Co}^{2+}(a q)+\mathrm{H}_{2}(g) \\
\text { (c) } 10 \mathrm{Br}^{-}(a q)+2 \mathrm{MnO}_{4}^{-}(a q)+16 \mathrm{H}^{+}(a q) \longrightarrow
2 \mathrm{Mn}^{2+}(a q)+8 \mathrm{H}_{2} \mathrm{O}(l)+5 \mathrm{Br}_{2}(l)
\end{array}
$$

Jennifer Landry
Jennifer Landry
Numerade Educator
07:59

Problem 56

Using the standard reduction potentials listed in Appendix E, calculate the equilibrium constant for each of the following reactions at $298 \mathrm{~K}$ :
(a) $\mathrm{Cu}(s)+2 \mathrm{Ag}^{+}(a q) \longrightarrow \mathrm{Cu}^{2+}(a q)+2 \mathrm{Ag}(s)$
(b) $3 \mathrm{Ce}^{4+}(a q)+\mathrm{Bi}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 3 \mathrm{Ce}^{3+}(a q)+ \mathrm{BiO}^{+}(a q)+2 \mathrm{H}^{+}(a q)$
(c) $\mathrm{N}_{2} \mathrm{H}_{5}^{+}(a q)+4 \mathrm{Fe}(\mathrm{CN})_{6}^{3-}(a q) \longrightarrow \mathrm{N}_{2}(g)+ 5 \mathrm{H}^{+}(a q)+4 \mathrm{Fe}(\mathrm{CN})_{6}^{4-}(a q)$

Keenan Mintz
Keenan Mintz
University of Miami
04:33

Problem 57

A cell has a standard cell potential of $+0.257 \mathrm{~V}$ at $298 \mathrm{~K}$. What is the value of the equilibrium constant for the reaction $(\mathbf{a})$ if $n=1 ?(\mathbf{b})$ if $n=2 ?(\mathbf{c})$ if $n=3 ?$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
01:41

Problem 58

At $298 \mathrm{~K}$ a cell reaction has a standard cell potential of $+0.63 \mathrm{~V}$. The equilibrium constant for the reaction is $3.8 \times 10^{10}$. What is the value of $n$ for the reaction?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
03:55

Problem 59

A voltaic cell is based on the reaction
$$
\operatorname{Sn}(s)+\mathrm{I}_{2}(s) \longrightarrow \mathrm{Sn}^{2+}(a q)+2 \mathrm{I}^{-}(a q)
$$
Under standard conditions, what is the maximum electrical work, in joules, that the cell can accomplish if $75.0 \mathrm{~g}$ of Sn is consumed?

Jennifer Landry
Jennifer Landry
Numerade Educator
03:10

Problem 60

Consider the voltaic cell illustrated in Figure $20.5,$ which is based on the cell reaction
$$
\mathrm{Zn}(s)+\mathrm{Cu}^{2+}(a q) \longrightarrow \mathrm{Zn}^{2+}(a q)+\mathrm{Cu}(s)
$$
Under standard conditions, what is the maximum electrical work, in joules, that the cell can accomplish if $50.0 \mathrm{~g}$ of copper is formed?

Keenan Mintz
Keenan Mintz
University of Miami
02:10

Problem 61

(a) In the Nernst equation, what is the numerical value of the reaction quotient, $Q,$ under standard conditions? (b) Can the Nernst equation be used at temperatures other than room temperature?

Jennifer Landry
Jennifer Landry
Numerade Educator
03:20

Problem 62

A voltaic cell is constructed with all reactants and products in their standard states. Will the concentration of the reactants increase, decrease, or remain the same as the cell operates?

Keenan Mintz
Keenan Mintz
University of Miami
06:12

Problem 63

What is the effect on the emf of the cell shown in Figure $20.5,$ which has the overall reaction $\mathrm{Zn}(s)+\mathrm{Cu}^{2+}(a q)$ $\longrightarrow \mathrm{Cu}(s)+\mathrm{Zn}^{2+}(a q),$ for each of the following changes? (a) Zinc chloride is added to the anode halfcell. (b) Copper chloride is added to the anode half-cell. (c) The surface area of the anode is cut to half. (d) The cell operates for one hour.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
05:30

Problem 64

A voltaic cell utilizes the following reaction:
$$
\mathrm{Al}(s)+3 \mathrm{Ag}^{+}(a q) \longrightarrow \mathrm{Al}^{3+}(a q)+3 \mathrm{Ag}(s)
$$
What is the effect on the cell emf of each of the following changes? (a) Water is added to the anode half-cell, diluting the solution. (b) The size of the aluminum electrode is increased. (c) A solution of $\mathrm{AgNO}_{3}$ is added to the cathode half-cell, increasing the quantity of $\mathrm{Ag}^{+}$ but not changing its concentration. (d) HCl is added to the $\mathrm{AgNO}_{3}$ solution, precipitating some of the $\mathrm{Ag}^{+}$ as $\mathrm{AgCl}$.

Shubham Kumar
Shubham Kumar
Numerade Educator
06:15

Problem 65

A voltaic cell is constructed that uses the following reaction and operates at $298 \mathrm{~K}$ :
$$
\mathrm{Zn}(s)+\mathrm{Ni}^{2+}(a q) \longrightarrow \mathrm{Zn}^{2+}(a q)+\mathrm{Ni}(s)
$$
(a) What is the emf of this cell under standard conditions?
(b) What is the emf of this cell when $\left[\mathrm{Ni}^{2+}\right]=3.00 \mathrm{M}$ and $\left[\mathrm{Zn}^{2+}\right]=0.100 \mathrm{M} ?(\mathbf{c})$ What is the emf of the cell when $\left[\mathrm{Ni}^{2+}\right]=0.200 M$ and $\left[\mathrm{Zn}^{2+}\right]=0.900 \mathrm{M} ?$

Jennifer Landry
Jennifer Landry
Numerade Educator
03:50

Problem 66

A voltaic cell utilizes the following reaction and operates at 298 K:
$$
3 \mathrm{Ce}^{4+}(a q)+\mathrm{Cr}(s) \longrightarrow 3 \mathrm{Ce}^{3+}(a q)+\mathrm{Cr}^{3+}(a q)
$$
(a) What is the emf of this cell under standard conditions? (b) What is the emf of this cell when $\left[\mathrm{Ce}^{4+}\right]=3.0 \mathrm{M},$ $\left[\mathrm{Ce}^{3+}\right]=0.10 \mathrm{M},$ and $\left[\mathrm{Cr}^{3+}\right]=0.010 \mathrm{M} ?(\mathbf{c})$ What is the emf of the cell when $\left[\mathrm{Ce}^{4^{+}}\right]=0.010 \mathrm{M},\left[\mathrm{Ce}^{3+}\right]=2.0 \mathrm{M}$ and $\left[\mathrm{Cr}^{3+}\right]=1.5 \mathrm{M} ?$

Keenan Mintz
Keenan Mintz
University of Miami
08:39

Problem 67

A voltaic cell utilizes the following reaction:
$4 \mathrm{Fe}^{2+}(a q)+\mathrm{O}_{2}(g)+4 \mathrm{H}^{+}(a q) \longrightarrow 4 \mathrm{Fe}^{3+}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l)$
(a) What is the emf of this cell under standard conditions? (b) What is the emf of this cell when $\left[\mathrm{Fe}^{2+}\right]=1.3 \mathrm{M},\left[\mathrm{Fe}^{3+}\right]=$ $0.010 \mathrm{M}, P_{\mathrm{O}_{2}}=50.7 \mathrm{kPa},$ and the $\mathrm{pH}$ of the solution in the cathode half-cell is $3.50 ?$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
07:08

Problem 68

A voltaic cell utilizes the following reaction:
$$
2 \mathrm{Fe}^{3+}(a q)+\mathrm{H}_{2}(g) \longrightarrow 2 \mathrm{Fe}^{2+}(a q)+2 \mathrm{H}^{+}(a q)
$$
(a) What is the emf of this cell under standard conditions? (b) What is the emf for this cell when $\left[\mathrm{Fe}^{3+}\right]=3.50 \mathrm{M}, P_{\mathrm{H}_{2}}=$ $96.3 \mathrm{kPa},\left[\mathrm{Fe}^{2+}\right]=0.0010 \mathrm{M},$ and the $\mathrm{pH}$ in both half-cells is $4.00 ?$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
06:52

Problem 69

A voltaic cell is constructed with two $\mathrm{Cu}^{2+}-\mathrm{Cu}$ electrodes. The two half-cells have $\left[\mathrm{Cu}^{2+}\right]=0.100 \mathrm{M}$ and $\left[\mathrm{Cu}^{2+}\right]=1.00 \times 10^{-4} \mathrm{M}$, respectively. (a) Which electrode is the cathode of the cell? (b) What is the standard emf of the cell? (c) What is the cell emf for the concentrations given? (d) For each electrode, predict whether $\left[\mathrm{Cu}^{2+}\right]$ will increase, decrease, or stay the same as the cell operates.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
04:27

Problem 70

A voltaic cell is constructed with two silver-silver chloride electrodes, each of which is based on the following half-reaction:
$$
\operatorname{AgCl}(s)+\mathrm{e}^{-} \longrightarrow \mathrm{Ag}(s)+\mathrm{Cl}^{-}(a q)
$$
The two half-cells have $\left[\mathrm{Cl}^{-}\right]=0.0150 \mathrm{M}$ and $\left[\mathrm{Cl}^{-}\right]=$ $2.55 M,$ respectively. (a) Which electrode is the cathode of the cell? (b) What is the standard emf of the cell? (c) What is the cell emf for the concentrations given? (d) For each electrode, predict whether $\left[\mathrm{Cl}^{-}\right]$ will increase, decrease, or stay the same as the cell operates.

Keenan Mintz
Keenan Mintz
University of Miami
10:53

Problem 71

The cell in Figure 20.9 could be used to provide a measure of the $\mathrm{pH}$ in the cathode half-cell. Calculate the pH of the cathode half-cell solution if the cell emf at $298 \mathrm{~K}$ is measured to be $+0.663 \mathrm{~V}$ when $\left[\mathrm{Zn}^{2+}\right]=0.10 \mathrm{M}$ and $P_{\mathrm{H}_{2}}=101.3 \mathrm{kPa}$.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
03:18

Problem 72

A voltaic cell is constructed that is based on the following reaction:
$$
\mathrm{Sn}^{2+}(a q)+\mathrm{Pb}(s) \longrightarrow \mathrm{Sn}(s)+\mathrm{Pb}^{2+}(a q)
$$
(a) If the concentration of $\mathrm{Sn}^{2+}$ in the cathode half-cell is $1.00 M$ and the cell generates an emf of $+0.22 \mathrm{~V},$ what is the concentration of $\mathrm{Pb}^{2+}$ in the anode half-cell? $(\mathbf{b})$ If the anode half-cell contains $\left[\mathrm{SO}_{4}^{2-}\right]=1.00 M$ in equilibrium with $\mathrm{PbSO}_{4}(s),$ what is the $K_{s p}$ of $\mathrm{PbSO}_{4} ?$

Keenan Mintz
Keenan Mintz
University of Miami
06:11

Problem 73

During a period of discharge of a lead-acid battery, $300 \mathrm{~g}$ of $\mathrm{PbO}_{2}(s)$ from the cathode is converted into $\mathrm{PbSO}_{4}(s)$. (a) What mass of $\mathrm{Pb}(s)$ is oxidized at the anode during this same period? (b) How many coulombs of electrical charge are transferred from $\mathrm{Pb}$ to $\mathrm{PbO}_{2}$ ?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
04:16

Problem 74

During the discharge of an alkaline battery, $4.50 \mathrm{~g}$ of $\mathrm{Zn}$ is consumed at the anode of the battery. (a) What mass of $\mathrm{MnO}_{2}$ is reduced at the cathode during this discharge? (b) How many coulombs of electrical charge are transferred from $\mathrm{Zn}$ to $\mathrm{MnO}_{2} ?$

Keenan Mintz
Keenan Mintz
University of Miami
04:36

Problem 75

Heart pacemakers are often powered by lithium-silver chromate "button" batteries. The overall cell reaction is
$$
2 \mathrm{Li}(s)+\mathrm{Ag}_{2} \mathrm{CrO}_{4}(s) \longrightarrow \mathrm{Li}_{2} \mathrm{CrO}_{4}(s)+2 \mathrm{Ag}(s)
$$
(a) Lithium metal is the reactant at one of the electrodes of the battery. Is it the anode or the cathode? (b) Choose the two half-reactions from Appendix $\mathrm{E}$ that most closely approximate the reactions that occur in the battery. What standard emf would be generated by a voltaic cell based on these half-reactions? (c) The battery generates an emf of $+3.5 \mathrm{~V}$. How close is this value to the one calculated in part (b)? (d) Calculate the emf that would be generated at body temperature, $37^{\circ} \mathrm{C}$. How does this compare to the emf you calculated in part (b)?

Jennifer Landry
Jennifer Landry
Numerade Educator
04:07

Problem 76

Mercuric oxide dry-cell batteries are often used where a flat discharge voltage and long life are required, such as in watches and cameras. The two half-cell reactions that occur in the battery are
$$
\begin{array}{l}
\mathrm{HgO}(s)+\mathrm{H}_{2} \mathrm{O}(l)+2 \mathrm{e}^{-} \longrightarrow \mathrm{Hg}(l)+2 \mathrm{OH}^{-}(a q) \\
\mathrm{Zn}(s)+2 \mathrm{OH}^{-}(a q) \longrightarrow \mathrm{ZnO}(s)+\mathrm{H}_{2} \mathrm{O}(l)+2 \mathrm{e}^{-}
\end{array}
$$
(a) Write the overall cell reaction. (b) The value of $E_{\text {red }}^{\circ}$ for the cathode reaction is $+0.098 \mathrm{~V}$. The overall cell potential is $+1.35 \mathrm{~V}$. Assuming that both half-cells operate under standard conditions, what is the standard reduction potential for the anode reaction? (c) Why is the potential of the anode reaction different than would be expected if the reaction occurred in an acidic medium?

Keenan Mintz
Keenan Mintz
University of Miami
06:16

Problem 77

(a) Suppose that an alkaline battery was manufactured using cadmium metal rather than zinc. What effect would this have on the cell emf? (b) What environmental advantage is provided by the use of nickel-metal hydride batteries over nickel-cadmium batteries?

Jennifer Landry
Jennifer Landry
Numerade Educator
05:26

Problem 78

In some applications nickel-cadmium batteries have been replaced by nickel-zinc batteries. The overall cell reaction for this relatively new battery is:
$$
\begin{aligned}
2 \mathrm{H}_{2} \mathrm{O}(l)+2 \mathrm{NiO}(\mathrm{OH})(s) &+\mathrm{Zn}(s) \\
& \longrightarrow 2 \mathrm{Ni}(\mathrm{OH})_{2}(s)+\mathrm{Zn}(\mathrm{OH})_{2}(s)
\end{aligned}
$$
(a)What is the cathode half-reaction? (b) What is the anode half-reaction? (c) A single nickel-cadmium cell has a voltage of $1.30 \mathrm{~V}$. Based on the difference in the standard reduction potentials of $\mathrm{Cd}^{2+}$ and $\mathrm{Zn}^{2+}$, what voltage would you estimate a nickel-zinc battery will produce? (d) Would you expect the specific energy density of a nickel-zinc battery to be higher or lower than that of a nickel-cadmium battery?

Keenan Mintz
Keenan Mintz
University of Miami
06:11

Problem 79

In a Li-ion battery the composition of the cathode is $\mathrm{LiCoO}_{2}$ when completely discharged. On charging, approximately $50 \%$ of the $\mathrm{Li}^{+}$ ions can be extracted from the cathode and transported to the graphite anode where they intercalate between the layers. (a) What is the composition of the cathode when the battery is fully charged? (b) If the $\mathrm{LiCoO}_{2}$ cathode has a mass of $10 \mathrm{~g}$ (when fully discharged), how many coulombs of electricity can be delivered on completely discharging a fully charged battery?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
04:35

Problem 80

Li-ion batteries used in automobiles typically use a $\operatorname{LiMn}_{2} \mathrm{O}_{4}$ cathode in place of the $\mathrm{LiCoO}_{2}$ cathode found in most Li-ion batteries. (a) Calculate the mass percent lithium in each electrode material. (b) Which material has a higher percentage of lithium? Does this help to explain why batteries made with $\operatorname{LiMn}_{2} \mathrm{O}_{4}$ cathodes deliver less power on discharging? (c) In a battery that uses a $\mathrm{LiCoO}_{2}$ cathode, approximately $50 \%$ of the lithium migrates from the cathode to the anode on charging. In a battery that uses a $\operatorname{LiMn}_{2} \mathrm{O}_{4}$ cathode, what fraction of the lithium in $\mathrm{LiMn}_{2} \mathrm{O}_{4}$ would need to migrate out of the cathode to deliver the same amount of lithium to the graphite anode?

Keenan Mintz
Keenan Mintz
University of Miami
02:59

Problem 81

(a) Which reaction is spontaneous in the hydrogen fuel cell: hydrogen gas plus oxygen gas makes water, or water makes hydrogen gas plus oxygen gas? (b) Using the standard reduction potentials in Appendix E, calculate the standard voltage generated by the hydrogen fuel cell in acidic solution.

Jennifer Landry
Jennifer Landry
Numerade Educator
01:58

Problem 82

(a) What is the difference between a battery and a fuel cell? (b) Can the "fuel" of a fuel cell be a solid?

Keenan Mintz
Keenan Mintz
University of Miami
04:14

Problem 83

(a) Write the anode and cathode reactions that cause the corrosion of iron metal to aqueous iron(II). $(\mathbf{b})$ Write the balanced half-reactions involved in the air oxidation of $\mathrm{Fe}^{2+}(a q)$ to $\mathrm{Fe}_{2} \mathrm{O}_{3} \cdot 3 \mathrm{H}_{2} \mathrm{O}(s)$.

Jennifer Landry
Jennifer Landry
Numerade Educator
02:21

Problem 84

(a) Based on standard reduction potentials, would you expect copper metal to oxidize under standard conditions in the presence of oxygen and hydrogen ions? (b) When the Statue of Liberty was refurbished, Teflon spacers were placed between the iron skeleton and the copper metal on the surface of the statue. What role do these spacers play?

Keenan Mintz
Keenan Mintz
University of Miami
03:29

Problem 85

(a) Aluminum metal is used as a sacrificial anode to protect offshore pipelines in salt water from corrosion. Why is the aluminum referred to as a "sacrificial anode"? (b) Looking in Appendix E, suggest what metal the pipelines could be made from in order for aluminum to be successful as a sacrificial anode.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
01:06

Problem 86

An iron object is plated with a coating of tin (Sn) to protect against corrosion. Does the tin protect iron by cathodic protection?

John Nicolle
John Nicolle
Numerade Educator
04:13

Problem 87

Iron corrodes to produce rust, $\mathrm{Fe}_{2} \mathrm{O}_{3},$ but other corrosion products that can form are $\mathrm{Fe}(\mathrm{O})(\mathrm{OH})$, iron oxyhydroxide, and magnetite, $\mathrm{Fe}_{3} \mathrm{O}_{4} .$ (a) What is the oxidation number of Fe in iron oxyhydroxide, assuming oxygen's oxidation number is $-2 ?(\mathbf{b})$ The oxidation number for Fe in magnetite was controversial for a long time. If we assume that oxygen's oxidation number is $-2,$ and Fe has a unique oxidation number, what is the oxidation number for Fe in magnetite? (c) It turns out that there are two different kinds of Fe in magnetite that have different oxidation numbers. Suggest what these oxidation numbers are and what their relative stoichiometry must be, assuming oxygen's oxidation number is -2 .

Jennifer Landry
Jennifer Landry
Numerade Educator
02:19

Problem 88

Copper corrodes to cuprous oxide, $\mathrm{Cu}_{2} \mathrm{O},$ or cupric oxide, $\mathrm{CuO},$ depending on environmental conditions. (a) What is the oxidation state of copper in cuprous oxide? (b) What is the oxidation state of copper in cupric oxide? (c) Copper peroxide is another oxidation product of elemental copper. Suggest a formula for copper peroxide based on its name. (d) Copper(III) oxide is another unusual oxidation product of elemental copper. Suggest a chemical formula for copper(III) oxide.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
04:11

Problem 89

(a) What is electrolysis? (b) Are electrolysis reactions thermodynamically spontaneous? (c) What process occurs at the anode in the electrolysis of molten $\mathrm{NaCl}$ ? (d) Why is sodium metal not obtained when an aqueous solution of $\mathrm{NaCl}$ undergoes electrolysis?

Jennifer Landry
Jennifer Landry
Numerade Educator
05:33

Problem 90

(a) What is an electrolytic cell? (b) The negative terminal of a voltage source is connected to an electrode of an electrolytic cell. Is the electrode the anode or the cathode of the cell? Explain. (c) The electrolysis of water is often done with a small amount of sulfuric acid added to the water. What is the role of the sulfuric acid? (d) Why are active metals such as Al obtained by electrolysis using molten salts rather than aqueous solutions?

Keenan Mintz
Keenan Mintz
University of Miami
04:43

Problem 91

(a) A $\mathrm{Cr}^{3+}(a q)$ solution is electrolyzed, using a current of $7.60 \mathrm{~A}$. What mass of $\mathrm{Cr}(s)$ is plated out after 2.00 days? (b) What amperage is required to plate out $0.250 \mathrm{~mol} \mathrm{Cr}$ from a $\mathrm{Cr}^{3+}$ solution in a period of $8.00 \mathrm{~h}$ ?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
06:18

Problem 92

Metallic magnesium can be made by the electrolysis of molten $\mathrm{MgCl}_{2}$ (a) What mass of $\mathrm{Mg}$ is formed by passing a current of 4.55 A through molten $\mathrm{MgCl}_{2}$, for 4.50 days? (b) How many minutes are needed to plate out $25.00 \mathrm{~g} \mathrm{Mg}$ from molten $\mathrm{MgCl}_{2}$ using $3.50 \mathrm{~A}$ of current?

Shubham Kumar
Shubham Kumar
Numerade Educator
06:29

Problem 93

(a) Calculate the mass of Li formed by electrolysis of molten LiCl by a current of $7.5 \times 10^{4}$ A flowing for a period of 24 h. Assume the electrolytic cell is $85 \%$ efficient. (b) What is the minimum voltage required to drive the reaction?

Jennifer Landry
Jennifer Landry
Numerade Educator
04:27

Problem 94

Elemental calcium is produced by the electrolysis of molten $\mathrm{CaCl}_{2}$. (a) What mass of calcium can be produced by this process if a current of $7.5 \times 10^{3} \mathrm{~A}$ is applied for $48 \mathrm{~h}$ ? Assume that the electrolytic cell is $68 \%$ efficient. (b) What is the minimum voltage needed to cause the electrolysis?

Shubham Kumar
Shubham Kumar
Numerade Educator
01:42

Problem 95

Metallic gold is collected from below the anode when a mixture of copper and gold metals is refined by electrolysis. Explain this behavior.

Lottie Adams
Lottie Adams
Numerade Educator
02:25

Problem 96

A mixture of copper and gold metals that is subjected to electrorefining contains tellurium as an impurity. The standard reduction potential between tellurium and its lowest common oxidation state, $\mathrm{Te}^{4+}$, is
$$
\mathrm{Te}^{4+}(a q)+4 \mathrm{e}^{-} \longrightarrow \mathrm{Te}(s) \quad E_{\mathrm{red}}^{\circ}=0.57 \mathrm{~V}
$$
Given this information, describe the probable fate of tellurium impurities during electrorefining. Do the impurities fall to the bottom of the refining bath, unchanged, as copper is oxidized, or do they go into solution as ions? If they go into solution, do they plate out on the cathode?

Keenan Mintz
Keenan Mintz
University of Miami
15:03

Problem 97

A disproportionation reaction is an oxidation-reduction reaction in which the same substance is oxidized and reduced. Complete and balance the following disproportionation reactions:
(a) $\mathrm{Fe}^{2+}(a q) \longrightarrow \mathrm{Fe}(s)+\mathrm{Fe}^{3+}(a q)$
(b) $\mathrm{Br}_{2}(l) \longrightarrow \mathrm{Br}^{-}(a q)+\mathrm{BrO}_{3}^{-}(a q)$ (acidic solution)
(c) $\mathrm{Cr}^{3+}(a q) \longrightarrow \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{Cr}(s)$ (acidic solution)
(d) $\mathrm{NO}(g) \longrightarrow \mathrm{N}_{2}(g)+\mathrm{NO}_{3}^{-}(a q)$ (acidic solution)

Susan Hallstrom
Susan Hallstrom
Numerade Educator
10:09

Problem 98

A common shorthand way to represent a voltaic cell is
anode $\mid$ anode solution || cathode solution $\mid$ cathode
A double vertical line represents a salt bridge or a porous barrier. A single vertical line represents a change in phase, such as from solid to solution. (a) Write the half-reactions and overall cell reaction represented by $\mathrm{Fe}\left|\mathrm{Fe}^{2+}\right|\left|\mathrm{Ag}^{+}\right| \mathrm{Ag} ;$ calculate the standard cell emf using data in Appendix E. (b) Write the half-reactions and overall cell reaction represented by $\mathrm{Zn}\left|\mathrm{Zn}^{2+}\right|\left|\mathrm{H}^{+}\right| \mathrm{H}_{2} ;$ calculate the standard cell emf using data in Appendix E and use Pt for the hydrogen electrode. (c) Using the notation just described, represent a cell based on the following reaction:
$$
\begin{aligned}
\mathrm{ClO}_{3}^{-}(a q)+3 \mathrm{Cu}(s) &+6 \mathrm{H}^{+}(a q) \\
& \longrightarrow \mathrm{Cl}^{-}(a q)+3 \mathrm{Cu}^{2+}(a q)+3 \mathrm{H}_{2} \mathrm{O}(l)
\end{aligned}
$$
$\mathrm{Pt}$ is used as an inert electrode in contact with the $\mathrm{ClO}_{3}^{-}$ and $\mathrm{Cl}^{-}$. Calculate the standard cell emf given: $\mathrm{ClO}_{3}^{-}(a q)+$ $6 \mathrm{H}^{+}(a q)+6 \mathrm{e}^{-} \longrightarrow \mathrm{Cl}^{-}(a q)+3 \mathrm{H}_{2} \mathrm{O}(l) ; E^{\circ}=1.45 \mathrm{~V}$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
05:08

Problem 99

Predict whether the following reactions will be spontaneous in acidic solution under standard conditions: (a) oxidation of $\mathrm{Cu}$ to $\mathrm{Cu}^{2+}$ by $\mathrm{I}_{2}$ (to form $\mathrm{I}^{-}$ ), $(\mathbf{b})$ reduction of $\mathrm{Fe}^{2+}$ to $\mathrm{Fe}$ by $\mathrm{H}_{2}$ (to form $\mathrm{H}^{+}$ ), $\left(\mathbf{c}\right.$ ) reduction of $\mathrm{I}_{2}$ to $\mathrm{I}^{-}$ by $\mathrm{H}_{2} \mathrm{O}_{2},(\mathbf{d})$ reduction of $\mathrm{Ni}^{2+}$ to $\mathrm{Ni}$ by $\mathrm{Sn}^{2+}\left(\right.$ to form $\left.\mathrm{Sn}^{4+}\right)$.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
08:16

Problem 100

Gold exists in two common positive oxidation states, +1 and +3 . The standard reduction potentials for these oxidation states are
$$
\begin{array}{l}
\mathrm{Au}^{+}(a q)+\mathrm{e}^{-} \quad \longrightarrow \mathrm{Au}(s) \quad E_{\mathrm{red}}^{\circ}=+1.69 \mathrm{~V} \\
\mathrm{Au}^{3+}(a q)+3 \mathrm{e}^{-} \longrightarrow \mathrm{Au}(s) E_{\mathrm{red}}^{\circ}=+1.50 \mathrm{~V}
\end{array}
$$
(a) Can you use these data to explain why gold does not tarnish in the air? (b) Suggest several substances that should be strong enough oxidizing agents to oxidize gold metal. (c) Miners obtain gold by soaking gold-containing ores in an aqueous solution of sodium cyanide. A very soluble complex ion of gold forms in the aqueous solution because of the redox reaction
$$
\begin{aligned}
4 \mathrm{Au}(s)+8 \mathrm{NaCN}(a q) &+2 \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{O}_{2}(g) \\
\longrightarrow & 4 \mathrm{Na}\left[\mathrm{Au}(\mathrm{CN})_{2}\right](a q)+4 \mathrm{NaOH}(a q)
\end{aligned}
$$
What is being oxidized, and what is being reduced in this reaction? (d) Gold miners then react the basic aqueous product solution from part (c) with $\mathrm{Zn}$ dust to get gold metal. Write a balanced redox reaction for this process. What is being oxidized, and what is being reduced?

Keenan Mintz
Keenan Mintz
University of Miami
09:03

Problem 101

A voltaic cell is constructed from an $\mathrm{Cd}^{2+}(a q)-\operatorname{Cd}(s)$ half-cell and an $\mathrm{Ag}^{+}(a q)-\mathrm{Ag}(s)$ half-cell. The initial concentration of $\mathrm{Cd}^{2+}(a q)$ in the $\mathrm{Cd}^{2+}(a q)-\mathrm{Cd}(s)$ half-cell is $\left[\mathrm{Cd}^{2+}\right]=0.0200 \mathrm{M}$. The initial cell voltage is $+1.102 \mathrm{~V}$. (a) By using data in Appendix E, calculate the standard emf of this voltaic cell. (b) Will the concentration of $\mathrm{Cd}^{2+}(a q)$ increase or decrease as the cell operates? (c) What is the initial concentration of $\mathrm{Ag}^{+}(a q)$ in the $\mathrm{Ag}^{+}(a q)-\mathrm{Ag}(s)$ half-cell?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
09:13

Problem 102

A voltaic cell is constructed that uses the following half-cell reactions:
$$
\begin{array}{l}
\mathrm{Ag}^{+}(a q)+2 \mathrm{e}^{-} \longrightarrow \mathrm{Ag}(s) \\
\mathrm{I}_{2}(s)+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{I}^{-}(a q)
\end{array}
$$
The cell is operated at $298 \mathrm{~K}$ with $\left[\mathrm{Ag}^{+}\right]=0.15 \mathrm{M}$ and $\left[\mathrm{I}^{-}\right]=0.035 \mathrm{M}$. (a) Determine $E$ for the cell at these concentrations. (b) Which electrode is the anode of the cell? (c) Is the answer to part (b) the same as it would be if the cell were operated under standard conditions? (d) With $\left[\mathrm{Ag}^{+}\right]$ equal to $0.15 \mathrm{M}$, at what concentration of $\mathrm{I}^{-}$ would the cell have zero potential?

Henry He
Henry He
Numerade Educator
03:51

Problem 103

Using data from Appendix E, calculate the equilibrium constant for the disproportionation of the iron(II) ion at room temperature:
$$
3 \mathrm{Fe}^{2+}(a q) \longrightarrow \mathrm{Fe}(s)+2 \mathrm{Fe}^{3+}(a q)
$$

Henry He
Henry He
Numerade Educator
04:28

Problem 104

(a) Write the reactions for the discharge and charge of a nickel-cadmium (nicad) rechargeable battery. (b) Given the following reduction potentials, calculate the standard emf of the cell:
$$
\begin{aligned}
\mathrm{Cd}(\mathrm{OH})_{2}(s)+2 \mathrm{e}^{-} \longrightarrow \mathrm{Cd}(s)+2 \mathrm{OH}^{-}(a q) & \\
E_{\mathrm{red}}^{\circ} &=-0.76 \mathrm{~V} \\
\mathrm{NiO}(\mathrm{OH})(s)+\mathrm{H}_{2} \mathrm{O}(l)+\mathrm{e}^{-} \longrightarrow \mathrm{Ni}(\mathrm{OH})_{2}(s)+\mathrm{OH}^{-}(a q) \\
E_{\mathrm{red}}^{\circ} &=+0.49 \mathrm{~V}
\end{aligned}
$$
(c) A typical nicad voltaic cell generates an emf of $+1.30 \mathrm{~V}$. Why is there a difference between this value and the one you calculated in part (b)? (d) Calculate the equilibrium constant for the overall nicad reaction based on this typical emf value.

Keenan Mintz
Keenan Mintz
University of Miami
02:00

Problem 105

The capacity of batteries such as a lithium-ion battery is expressed in units of milliamp-hours (mAh). A typical battery of this type yields a nominal capacity of $2000 \mathrm{mAh}$. (a) What quantity of interest to the consumer is being expressed by the units of $\mathrm{mAh}$ ? (b) The starting voltage of a fresh lithium-ion battery is $3.60 \mathrm{~V}$. The voltage decreases during discharge and is $3.20 \mathrm{~V}$ when the battery has delivered its rated capacity. If we assume that the voltage declines linearly as current is withdrawn, estimate the total maximum electrical work the battery could perform during discharge.

Lottie Adams
Lottie Adams
Numerade Educator
02:42

Problem 106

Disulfides are compounds that have $S-S$ bonds, like peroxides have $\mathrm{O}-\mathrm{O}$ bonds. Thiols are organic compounds that have the general formula $\mathrm{R}-\mathrm{SH}$, where $\mathrm{R}$ is a generic hydrocarbon. The $\mathrm{SH}^{-}$ ion is the sulfur counterpart of hydroxide, $\mathrm{OH}^{-}$. Two thiols can react to make a disulfide, $\mathrm{R}-\mathrm{S}-\mathrm{S}-\mathrm{R} .$ (a) What is the oxidation state of sulfur in a thiol? (b) What is the oxidation state of sulfur in a disulfide? (c) If you react two thiols to make a disulfide, are you oxidizing or reducing the thiols? (d) If you wanted to convert a disulfide to two thiols, should you add a reducing agent or oxidizing agent to the solution? (e) Suggest what happens to the H's in the thiols when they form disulfides.

Henry He
Henry He
Numerade Educator
06:47

Problem 107

(a) How many coulombs are required to plate a layer of chromium metal $0.15 \mathrm{~mm}$ thick on an auto bumper with a total area of $0.40 \mathrm{~m}^{2}$ from a solution containing $\mathrm{CrO}_{4}^{2-}$ ? The density of chromium metal is $7.20 \mathrm{~g} / \mathrm{cm}^{3}$. (b) What current flow is required for this electroplating if the bumper is to be plated in $20.0 \mathrm{~s} ?(\mathbf{c})$ If the external source has an emf of $+5.5 \mathrm{~V}$ and the electrolytic cell is $60 \%$ efficient, how much electrical energy is expended to electroplate the bumper?

Henry He
Henry He
Numerade Educator
03:02

Problem 108

Magnesium is obtained by electrolysis of molten $\mathrm{MgCl}_{2}$. (a) Why is an aqueous solution of $\mathrm{MgCl}_{2}$ not used in the electrolysis? (b) Several cells are connected in parallel by very large copper bars that convey current to the cells. Assuming that the cells are $96 \%$ efficient in producing the desired products in electrolysis, what mass of $\mathrm{Mg}$ is formed by passing a current of 97,000 A for a period of 24 h?

Henry He
Henry He
Numerade Educator
03:35

Problem 109

Calculate the number of kilowatt-hours of electricity required to produce $500 \mathrm{~kg}$ of aluminum by electrolysis of $\mathrm{Al}^{3+}$ if the applied voltage is $5.00 \mathrm{~V}$ and the process is $50 \%$ efficient.

Henry He
Henry He
Numerade Educator
07:22

Problem 110

Some years ago a unique proposal was made to raise the Titanic. The plan involved placing pontoons within the ship using a surface-controlled submarine-type vessel. The pontoons would contain cathodes and would be filled with hydrogen gas formed by the electrolysis of water. It has been estimated that it would require about $7 \times 10^{8} \mathrm{~mol}$ of $\mathrm{H}_{2}$ to provide the buoyancy to lift the ship (J. Chem. Educ., 1973, Vol. 50, 61). (a) How many coulombs of electrical charge would be required? (b) What is the minimum voltage required to generate $\mathrm{H}_{2}$ and $\mathrm{O}_{2}$ if the pressure on the gases at the depth of the wreckage $(3 \mathrm{~km})$ is $30 \mathrm{MPa} ?(\mathbf{c})$ What is the minimum electrical energy required to raise the Titanic by electrolysis? (d) What is the minimum cost of the electrical energy required to generate the necessary $\mathrm{H}_{2}$ if the electricity costs 85 cents per kilowatt-hour to generate at the site?

Keenan Mintz
Keenan Mintz
University of Miami
05:31

Problem 111

The Haber process is the principal industrial route for converting nitrogen into ammonia:
$$
\mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \longrightarrow 2 \mathrm{NH}_{3}(g)
$$
(a) What is being oxidized, and what is being reduced? (b) Using the thermodynamic data in Appendix C, calculate the equilibrium constant for the process at room temperature. (c) Calculate the standard emf of the Haber process at room temperature.

Jennifer Landry
Jennifer Landry
Numerade Educator
04:30

Problem 112

In a galvanic cell the cathode is an $\mathrm{Ag}^{+}(1.00 \mathrm{M}) / \mathrm{Ag}(s)$ half-cell. The anode is a standard hydrogen electrode immersed in a buffer solution containing $0.10 \mathrm{M}$ benzoic acid $\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH}\right)$ and $0.050 \mathrm{M}$ sodium benzoate $\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COO}^{-} \mathrm{Na}^{+}\right)$. The measured cell voltage is $1.030 \mathrm{~V}$. What is the $\mathrm{p} K_{\mathrm{a}}$ of benzoic acid?

Keenan Mintz
Keenan Mintz
University of Miami
02:23

Problem 113

Aqueous solutions of ammonia $\left(\mathrm{NH}_{3}\right)$ and bleach (active ingredient $\mathrm{NaOCl}$ ) are sold as cleaning fluids, but bottles of both of them warn: "Never mix ammonia and bleach, as toxic gases may be produced." One of the toxic gases that can be produced is chloroamine, $\mathrm{NH}_{2} \mathrm{Cl}$. (a) What is the oxidation number of chlorine in bleach? (b) What is the oxidation number of chlorine in chloramine? (c) Is Cl oxidized, reduced, or neither, upon the conversion of bleach to chloramine? (d) Another toxic gas that can be produced is nitrogen trichloride, $\mathrm{NCl}_{3}$. What is the oxidation number of $\mathrm{N}$ in nitrogen trichloride? $(\mathbf{e})$ Is $\mathrm{N}$ oxidized, reduced, or neither, upon the conversion of ammonia to nitrogen trichloride?

Lottie Adams
Lottie Adams
Numerade Educator
07:11

Problem 114

A voltaic cell is based on $\mathrm{Cu}^{2+}(a q) / \mathrm{Cu}(s)$ and $\mathrm{Br}_{2}(l) /$ $\mathrm{Br}^{-}(a q)$ half-cells. (a) What is the standard emf of the cell? (b) Which reaction occurs at the cathode and which at the anode of the cell? (c) Use $S^{\circ}$ values in Appendix $\mathrm{C}$ and the relationship between cell potential and free-energy change to predict whether the standard cell potential increases or decreases when the temperature is raised above $25^{\circ} \mathrm{C}$. (Thestandard entropy of $\mathrm{Cu}^{2+}(a q)$ is $\left.S^{\circ}=-99.6 \mathrm{~J} / \mathrm{K}\right)$

Henry He
Henry He
Numerade Educator
02:33

Problem 115

Hydrogen gas has the potential for use as a clean fuel in reaction with oxygen. The relevant reaction is
$$
2 \mathrm{H}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}(l)
$$
Consider two possible ways of utilizing this reaction as an electrical energy source: (i) Hydrogen and oxygen gases are combusted and used to drive a generator, much as coal is currently used in the electric power industry; (ii) hydrogen and oxygen gases are used to generate electricity directly by using fuel cells that operate at $85^{\circ} \mathrm{C} .$ (a) Use data in Appendix $\mathrm{C}$ to calculate $\Delta H^{\circ}$ and $\Delta S^{\circ}$ for the reaction. We will assume that these values do not change appreciably with temperature. (b) Based on the values from part (a), what trend would you expect for the magnitude of $\Delta G$ for the reaction as the temperature increases? (c) What is the significance of the change in the magnitude of $\Delta G$ with temperature with respect to the utility of hydrogen as a fuel? (d) Based on the analysis here, would it be more efficient to use the combustion method or the fuel-cell method to generate electrical energy from hydrogen?

Lottie Adams
Lottie Adams
Numerade Educator
04:47

Problem 116

Cytochrome, a complicated molecule that we will represent as $\mathrm{CyFe}^{2+}$, reacts with the air we breathe to supply energy required to synthesize adenosine triphosphate (ATP). The body uses ATP as an energy source to drive other reactions (Section 19.7). At $\mathrm{pH} 7.0$ the following reduction potentials pertain to this oxidation of $\mathrm{CyFe}^{2+}$
$$
\begin{aligned}
\mathrm{O}_{2}(g)+4 \mathrm{H}^{+}(a q)+4 \mathrm{e}^{-} & \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}(l) & & E_{\mathrm{red}}^{\circ}=+0.82 \mathrm{~V} \\
\mathrm{CyFe}^{3+}(a q)+\mathrm{e}^{-} & \longrightarrow \mathrm{CyFe}^{2+}(a q) & E_{\mathrm{red}}^{\circ} &=+0.22 \mathrm{~V}
\end{aligned}
$$
(a) What is $\Delta G$ for the oxidation of $\mathrm{CyFe}^{2+}$ by air? $(\mathbf{b})$ If the synthesis of $1.00 \mathrm{~mol}$ of ATP from adenosine diphosphate (ADP) requires a $\Delta G$ of $37.7 \mathrm{~kJ},$ how many moles of ATP are synthesized per mole of $\mathrm{O}_{2} ?$

Keenan Mintz
Keenan Mintz
University of Miami
03:45

Problem 117

The standard potential for the reduction of $\operatorname{AgCl}(s)$ is $+0.222 \mathrm{~V}$.
$$
\mathrm{AgCl}(s)+\mathrm{e}^{-} \longrightarrow \mathrm{Ag}(s)+\mathrm{CI}^{-}(a q)
$$
Using this value and the electrode potential for $\mathrm{Ag}^{+}(a q)$, calculate the $K_{s p}$ for $\mathrm{AgCl}$.

Henry He
Henry He
Numerade Educator
04:49

Problem 118

The $K_{s p}$ value for $\operatorname{Agl}(s)$ is $8.3 \times 10^{-17}$. By using this value together with an electrode potential from Appendix E, determine the value of the standard reduction potential for the reaction
$$
\mathrm{Agl}(s)+\mathrm{e}^{-} \longrightarrow \mathrm{Ag}(s)+\mathrm{I}^{-}(a q)
$$

Keenan Mintz
Keenan Mintz
University of Miami
02:36

Problem 119

A student designs an ammeter (device that measures electrical current) that is based on the electrolysis of water into hydrogen and oxygen gases. When electrical current of unknown magnitude is run through the device for 90 min, $32.5 \mathrm{~mL}$ of water-saturated $\mathrm{H}_{2}(g)$ is collected. The temperature of the system is $20^{\circ} \mathrm{C},$ and the atmospheric pressure is $101.3 \mathrm{kPa}$. What is the magnitude of the average current in amperes?

Lottie Adams
Lottie Adams
Numerade Educator