Chapter 17, Electrochemistry Video Solutions, Chemistry

Chapter Questions

Problem 1

Identify each half-reaction below as either oxidation or reduction.
(a) $\mathrm{Fe}^{3+}+3 \mathrm{e}^{-} \longrightarrow \mathrm{Fe}$
(b) $\mathrm{Cr} \longrightarrow \mathrm{Cr}^{3+}+3 \mathrm{e}^{-}$
(c) $\mathrm{MnO}_{4}^{2-} \longrightarrow \mathrm{MnO}_{4}^{-}+\mathrm{e}^{-}$
(d) $\mathrm{Li}^{+}+\mathrm{e}^{-} \longrightarrow \mathrm{Li}$

Nicole S.

Numerade Educator

00:51

Problem 2

Identify each half-reaction below as either oxidation or reduction.
(a) $\mathrm{Cl}^{-} \longrightarrow \mathrm{Cl}_{2}$
(b) $\mathrm{Mn}^{2+} \longrightarrow \mathrm{MnO}_{2}$
(c) $\mathrm{H}_{2} \longrightarrow \mathrm{H}^{+}$
(d) $\mathrm{NO}_{3}^{-} \longrightarrow \mathrm{NO}$

Nicole S.

Numerade Educator

03:01

Problem 3

Assuming each pair of half-reactions below takes place in an acidic solution, write a balanced equation for the overall reaction.
(a) $\mathrm{Ca} \longrightarrow \mathrm{Ca}^{2+}+2 \mathrm{e}^{-}, \quad \mathrm{F}_{2}+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{~F}^{-}$
(b) $\mathrm{Li} \longrightarrow \mathrm{Li}^{+}+\mathrm{e}^{-}, \quad \mathrm{Cl}_{2}+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{Cl}^{-}$
(c) $\mathrm{Fe} \longrightarrow \mathrm{Fe}^{3+}+3 \mathrm{e}^{-}, \mathrm{Br}_{2}+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{Br}^{-}$
(d) $\mathrm{Ag} \longrightarrow \mathrm{Ag}^{+}+\mathrm{e}^{-}, \quad \mathrm{MnO}_{4}^{-}+4 \mathrm{H}^{+}+3 \mathrm{e}^{-} \longrightarrow \mathrm{MnO}_{2}+2 \mathrm{H}_{2} \mathrm{O}$

Nicole S.

Numerade Educator

14:13

Problem 4

Balance the equations below assuming they occur in an acidic solution.
(a) $\mathrm{H}_{2} \mathrm{O}_{2}+\mathrm{Sn}^{2+} \longrightarrow \mathrm{H}_{2} \mathrm{O}+\mathrm{Sn}^{4+}$
(b) $\mathrm{PbO}_{2}+\mathrm{Hg} \longrightarrow \mathrm{Hg}_{2}^{2+}+\mathrm{Pb}^{2+}$
(c) $\mathrm{Al}+\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-} \longrightarrow \mathrm{Al}^{3+}+\mathrm{Cr}^{3+}$

Vishal S.

Numerade Educator

Problem 5

Identify the oxidant and reductant of each reaction of the previous exercise.

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Problem 6

Balance the equations below assuming they occur in a basic solution.
(a) $\mathrm{SO}_{3}^{2-}(a q)+\mathrm{Cu}(\mathrm{OH})_{2}(s) \longrightarrow \mathrm{SO}_{4}^{2-}(a q)+\mathrm{Cu}(\mathrm{OH})(s)$
(b) $\mathrm{O}_{2}(g)+\mathrm{Mn}(\mathrm{OH})_{2}(s) \longrightarrow \mathrm{MnO}_{2}(s)$
(c) $\mathrm{NO}_{3}^{-}(a q)+\mathrm{H}_{2}(g) \longrightarrow \mathrm{NO}(g)$
(d) $\mathrm{Al}(s)+\mathrm{CrO}_{4}^{2-}(a q) \longrightarrow \mathrm{Al}(\mathrm{OH})_{3}(s)+\mathrm{Cr}(\mathrm{OH})_{4}-(a q)$

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Problem 7

Identify the oxidant and reductant of each reaction of the previous exercise.

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01:40

Problem 8

Why don't hydroxide ions appear in equations for half-reactions occurring in acidic solution?

Nicole S.

Numerade Educator

01:00

Problem 9

Why don't hydrogen ions appear in equations for half-reactions occurring in basic solution?

Nicole S.

Numerade Educator

02:22

Problem 10

Why must the charge balance in oxidation-reduction reactions?

Emily P.

Numerade Educator

02:17

Problem 11

Write cell schematics for the following cell reactions, using platinum as an inert electrode as needed.
(a) $\mathrm{Mg}(s)+\mathrm{Ni}^{2+}(a q) \longrightarrow \mathrm{Mg}^{2+}(a q)+\mathrm{Ni}(s)$
(b) $2 \mathrm{Ag}^{+}(a q)+\mathrm{Cu}(s) \longrightarrow \mathrm{Cu}^{2+}(a q)+2 \mathrm{Ag}(s)$
(c) $\mathrm{Mn}(s)+\mathrm{Sn}\left(\mathrm{NO}_{3}\right)_{2}(a q) \longrightarrow \mathrm{Mn}\left(\mathrm{NO}_{3}\right)_{2}(a q)+\mathrm{Au}(s)$
(d) $3 \mathrm{CuNO}_{3}(a q)+\mathrm{Au}\left(\mathrm{NO}_{3}\right)_{3}(a q) \longrightarrow 3 \mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}(a q)+\mathrm{Au}(s)$

Nicole S.

Numerade Educator

01:03

Problem 12

Assuming the schematics below represent galvanic cells as written, identify the half-cell reactions occurring in each.
(a) $\mathrm{Mg}(s)\left|\mathrm{Mg}^{2+}(a q) \| \mathrm{Cu}^{2+}(a q)\right| \mathrm{Cu}(s)$
(b) $\mathrm{Ni}(s)\left|\mathrm{Ni}^{2+}(a q) \| \mathrm{Ag}^{+}(a q)\right| \mathrm{Ag}(s)$

Nicole S.

Numerade Educator

01:11

Problem 13

Write a balanced equation for the cell reaction of each cell in the previous exercise.

Nicole S.

Numerade Educator

Problem 14

Balance each reaction below, and write a cell schematic representing the reaction as it would occur in a galvanic cell.
(a) $\mathrm{Al}(s)+\mathrm{Zr}^{4+}(a q) \longrightarrow \mathrm{Al}^{3+}(a q)+\mathrm{Zr}(s)$
(b) $\mathrm{Ag}^{+}(a q)+\mathrm{NO}(g) \longrightarrow \mathrm{Ag}(s)+\mathrm{NO}_{3}-(a q)$
(acidic solution)
(c) $\mathrm{SiO}_{3}^{2-}(a q)+\mathrm{Mg}(s) \longrightarrow \mathrm{Si}(s)+\mathrm{Mg}(\mathrm{OH})_{2}(s) \quad$ (basic solution)
(d) $\mathrm{ClO}_{3}^{-}(a q)+\mathrm{MnO}_{2}(s) \longrightarrow \mathrm{Cl}^{-}(a q)+\mathrm{MnO}_{4}-(a q)$
(basic solution)

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Problem 15

Identify the oxidant and reductant in each reaction of the previous exercise.

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Problem 16

From the information provided, use cell notation to describe the following systems:
(a) In one half-cell, a solution of $\mathrm{Pt}\left(\mathrm{NO}_{3}\right)_{2}$ forms Pt metal, while in the other half-cell, Cu metal goes into a $\mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}$ solution with all solute concentrations $1 \mathrm{M}$
(b) The cathode consists of a gold electrode in a $0.55 \mathrm{M} \mathrm{Au}\left(\mathrm{NO}_{3}\right)_{3}$ solution and the anode is a magnesium electrode in $0.75 \mathrm{M} \mathrm{Mg}\left(\mathrm{NO}_{3}\right)_{2}$ solution.
(c) One half-cell consists of a silver electrode in a $1 \mathrm{M} \mathrm{AgNO}_{3}$ solution, and in the other half-cell, a copper electrode in $1 \mathrm{M} \mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}$ is oxidized.

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00:40

Problem 17

Why is a salt bridge necessary in galvanic cells like the one in Figure $17.3 ?$

Aadit S.

Numerade Educator