Content check — Chemistry 2e, Chapter 17: Electrochemistry

Every problem in the chapter beside the book’s answer, a chapter at a time as you scroll.

Chapter 17: Electrochemistry

34 problems · 3 figures

Example 17.1 §17.1 source ↗

The reaction above results when using relatively diluted nitric acid. If concentrated nitric acid is used, nitrogen dioxide is produced instead of nitrogen monoxide. Write a balanced equation for this reaction.

Answer

Cu (s) +

2 H^{+} (a q) +

2 {HNO}_{3} (a q) ⟶

{Cu}^{2+} (a q) +

2 {NO}_{2} (g) +

2 H_{2} O (l)

Example 17.2 §17.1 source ↗

Aqueous permanganate ion may also be reduced using aqueous bromide ion, Br⁻, the products of this reaction being solid manganese(IV) oxide and aqueous bromate ion, BrO₃⁻. Write the balanced equation for this reaction occurring in a basic medium.

Answer

H_{2} O (l) +

{2MnO}_{4}^{−} (a q) +

{Br}^{−} (a q) ⟶

2 {MnO}_{2} (s) +

{BrO}_{3}^{−} (a q) +

2 {OH}^{−} (a q)

Example 17.3 §17.2 source ↗

Omitting solute concentrations and spectator ion identities, write the schematic for a galvanic cell whose net cell reaction is shown below.

{Sn}^{4+} (a q) + Zn (s) ⟶ {Sn}^{2+} (a q) + {Zn}^{2+} (a q)

Answer

Zn (s) │ {Zn}^{2+} (a q) ║ {Sn}^{4+} (a q),

{Sn}^{2+} (a q) │ Pt (s)

Example 17.4 §17.3 source ↗

What is the standard cell potential expected if the silver cathode half-cell in Figure 17.3 is replaced with a lead half-cell:

{Pb}^{2+} (a q) +

2 e^{−} ⟶

Pb (s)

This figure contains a diagram of an electrochemical cell. Two beakers are shown. Each is just over half full. The beaker on the left contains a blue solution and is labeled below as “1 M solution of copper (II) nitrate ( C u ( N O subscript 3 ) subscript 2 ).” The beaker on the right contains a colorless solution and is labeled below as “1 M solution of silver nitrate ( A g N O subscript 3 ).” A glass tube in the shape of an inverted U connects the two beakers at the center of the diagram. The tube contents are colorless. The ends of the tubes are beneath the surface of the solutions in the beakers and a small gray plug is present at each end of the tube. The plug in the left beaker is labeled “Porous plug.” At the center of the diagram, the tube is labeled “Salt bridge ( N a N O subscript 3 ). Each beaker shows a metal strip partially submerged in the liquid. The beaker on the left has an orange-brown strip that is labeled “C u anode negative” at the top. The beaker on the right has a silver strip that is labeled “A g cathode positive” at the top. A wire extends from the top of each of these strips to a rectangle indicating “external circuit” that is labeled “flow of electrons” with an arrow pointing to the right following. A curved arrow extends from the C u strip into the surrounding solution. The tip of this arrow is labeled “C u superscript 2 plus.” A curved arrow extends from the salt bridge into the beaker on the left into the blue solution. The tip of this arrow is labeled “N O subscript 3 superscript negative.” A curved arrow extends from the solution in the beaker on the right to the A g strip. The base of this arrow is labeled “A g superscript plus.” A curved arrow extends from the colorless solution to salt bridge in the beaker on the right. The base of this arrow is labeled “N O subscript 3 superscript negative.” Just right of the salt bridge in the colorless solution is the label “N a superscript plus.” Just above this region of the tube appears the label “Flow of cations.” Just left of the salt bridge in the blue solution is the label “N O subscript 3 superscript negative.” Just above this region of the tube appears the label “Flow of anions.” — Figure 17.3 in the book ↗

Answer

−0. 47 V

Example 17.5 §17.3 source ↗

Use the data in Table 17.1 to predict the spontaneity of the oxidation of bromide ion by molecular iodine under standard state conditions, supporting the prediction by calculating the standard cell potential for the reaction. Repeat for the oxidation of iodide ion by molecular bromine.

Table 17.1 in the book ↗

Half-Reaction	E° (V)
$F_{2} (g) +$ ${2e}^{−} ⟶$ ${2F}^{−} (a q)$	+2.866
${PbO}_{2} (s) +$ ${SO}_{4}^{2−} (a q) +$ ${4H}^{+} (a q) +$ ${2e}^{−} ⟶$ ${PbSO}_{4} (s) +$ ${2H}_{2} O (l)$	+1.69
${MnO}_{4}^{−} (a q) +$ ${8H}^{+} (a q) +$ ${5e}^{−} ⟶$ ${Mn}^{2+} (a q) +$ ${4H}_{2} O (l)$	+1.507
${Au}^{3+} (a q) +$ ${3e}^{−} ⟶$ $Au (s)$	+1.498
${Cl}_{2} (g) +$ ${2e}^{−} ⟶$ ${2Cl}^{−} (a q)$	+1.35827
$O_{2} (g) +$ ${4H}^{+} (a q) +$ ${4e}^{−} ⟶$ ${2H}_{2} O (l)$	+1.229
${Pt}^{2+} (a q) +$ ${2e}^{−} ⟶$ $Pt (s)$	+1.20
${Br}_{2} (a q) +$ ${2e}^{−} ⟶$ ${2Br}^{−} (a q)$	+1.0873
${Ag}^{+} (a q) +$ $e^{−} ⟶$ $Ag (s)$	+0.7996
${Hg}_{2}^{2+} (a q) +$ ${2e}^{−} ⟶$ $2Hg (l)$	+0.7973
${Fe}^{3+} (a q) +$ $e^{−} ⟶$ ${Fe}^{2+} (a q)$	+0.771
${MnO}_{4}^{−} (a q) +$ ${2H}_{2} O (l) +$ ${3e}^{−} ⟶$ ${MnO}_{2} (s) +$ ${4OH}^{−} (a q)$	+0.558
$I_{2} (s) +$ ${2e}^{−} ⟶$ ${2I}^{−} (a q)$	+0.5355
${NiO}_{2} (s) +$ ${2H}_{2} O (l) +$ ${2e}^{−} ⟶$ ${Ni(OH)}_{2} (s) +$ ${2OH}^{−} (a q)$	+0.49
${Cu}^{2+} (a q) +$ ${2e}^{−} ⟶$ $Cu (s)$	+0.34
${Hg}_{2} {Cl}_{2} (s) +$ ${2e}^{−} ⟶$ $2Hg (l) +$ ${2Cl}^{−} (a q)$	+0.26808
$AgCl (s) +$ $e^{−} ⟶$ $Ag (s) +$ ${Cl}^{−} (a q)$	+0.22233
${Sn}^{4+} (a q) +$ ${2e}^{−} ⟶$ ${Sn}^{2+} (a q)$	+0.151
${2H}^{+} (a q) +$ ${2e}^{−} ⟶$ $H_{2} (g)$	0.00
${Pb}^{2+} (a q) +$ ${2e}^{−} ⟶$ $Pb (s)$	−0.1262
${Sn}^{2+} (a q) +$ ${2e}^{−} ⟶$ $Sn (s)$	−0.1375
${Ni}^{2+} (a q) +$ ${2e}^{−} ⟶$ $Ni (s)$	−0.257
${Co}^{2+} (a q) +$ ${2e}^{−} ⟶$ $Co (s)$	−0.28
${PbSO}_{4} (s) +$ ${2e}^{−} ⟶$ $Pb (s) +$ ${SO}_{4}^{2−} (a q)$	−0.3505
${Cd}^{2+} (a q) +$ ${2e}^{−} ⟶$ $Cd (s)$	−0.4030
${Fe}^{2+} (a q) +$ ${2e}^{−} ⟶$ $Fe (s)$	−0.447
${Cr}^{3+} (a q) +$ ${3e}^{−} ⟶$ $Cr (s)$	−0.744
${Mn}^{2+} (a q) +$ ${2e}^{−} ⟶$ $Mn (s)$	−1.185
${Zn(OH)}_{2} (s) +$ ${2e}^{−} ⟶$ $Zn (s) +$ ${2OH}^{−} (a q)$	−1.245
${Zn}^{2+} (a q) +$ ${2e}^{−} ⟶$ $Zn (s)$	−0.7618
${Al}^{3+} (a q) +$ ${3e}^{−} ⟶$ $Al (s)$	−1.662
${Mg}^{2} (a q) +$ ${2e}^{−} ⟶$ $Mg (s)$	−2.372
${Na}^{+} (a q) +$ $e^{−} ⟶$ $Na (s)$	−2.71
${Ca}^{2+} (a q) +$ ${2e}^{−} ⟶$ $Ca (s)$	−2.868
${Ba}^{2+} (a q) +$ ${2e}^{−} ⟶$ $Ba (s)$	−2.912
$K^{+} (a q) +$ $e^{−} ⟶$ $K (s)$	−2.931
${Li}^{+} (a q) +$ $e^{−} ⟶$ $Li (s)$	−3.04

Answer

\begin{matrix} I_{2} (s) + 2 {Br}^{−} (a q) ⟶ 2 I^{−} (a q) + {Br}_{2} (l) & E^{°}_{cell} = –0.5518 V (nonspontaneous) \\ {Br}_{2} (s) + 2 I^{−} (a q) ⟶ 2 {Br}^{−} (a q) + I_{2} (s) & E^{°}_{cell} = +0.5518 V (spontaneous) \end{matrix}

Example 17.6 §17.4 source ↗

What is the standard free energy change and the equilibrium constant for the following reaction at room temperature? Is the reaction spontaneous?

Sn (s) + 2 {Cu}^{2+} (a q) ⇌ {Sn}^{2+} (a q) + 2 {Cu}^{+} (a q)

Answer

Spontaneous; n = 2;

E_{cell}^{°} =

+0.291 V;

Δ G ° =

−56.2 \frac{kJ}{mol};

K = 6.8

\times

10⁹.

Example 17.7 §17.4 source ↗

For the cell schematic below, identify values for n and Q, and calculate the cell potential, E_cell.

Al (s) │ {Al}^{3+} (a q, 0.15 M) ║ {Cu}^{2+} (a q, 0.025 M) │ Cu (s)

Answer

n = 6; Q = 1440; E_cell = +1.97 V, spontaneous.

Example 17.8 §17.4 source ↗

The concentration cell above was allowed to operate until the cell reaction reached equilibrium. What are the cell potential and the concentrations of zinc(II) in each half-cell for the cell now?

Answer

E_cell = 0.000 V; [Zn²⁺]_cathode = [Zn²⁺]_anode = 0.30 M

Example 17.9 §17.7 source ↗

Aluminum metal can be made from aluminum(III) ions by electrolysis. What is the half-reaction at the cathode? What mass of aluminum metal would be recovered if a current of 25.0 A passed through the solution for 15.0 minutes?

Answer

{Al}^{3+} (a q) +

3 e^{−} ⟶

Al (s);

0.0777 mol Al = 2.10 g Al.

Example 17.10 §17.7 source ↗

What mass of zinc is required to galvanize the top of a 3.00 m

\times

5.50 m sheet of iron to a thickness of 0.100 mm of zinc? If the zinc comes from a solution of Zn(NO₃)₂ and the current is 25.5 A, how long will it take to galvanize the top of the iron? The density of zinc is 7.140 g/cm³.

Answer

11.8 kg Zn requires 379 hours.

17-1 §17.1 source ↗

Identify each half-reaction below as either oxidation or reduction.

(a) ${Fe}^{3+} + {3e}^{−} ⟶ Fe$

(b) $Cr ⟶ {Cr}^{3+} + {3e}^{−}$

(d) ${Li}^{+} + e^{−} ⟶ Li$

Answer

(a) reduction; (b) oxidation; (c) oxidation; (d) reduction

17-3 §17.1 source ↗

Assuming each pair of half-reactions below takes place in an acidic solution, write a balanced equation for the overall reaction.

(a) $Ca ⟶$ ${Ca}^{2+} +$ ${2e}^{−},$ $F_{2} + {2e}^{−} ⟶ {2F}^{−}$

(b) $Li ⟶ {Li}^{+} + e^{−},$ ${Cl}_{2} +$ ${2e}^{−} ⟶$ ${2Cl}^{−}$

(d) $Ag ⟶ {Ag}^{+} + e^{−},$ ${MnO}_{4}^{-} +$ ${4H}^{+} +$ ${3e}^{−} ⟶$ ${MnO}_{2} +$ ${2H}_{2} O$

Answer

(a)

F_{2} +

Ca ⟶

{2F}^{−} +

{Ca}^{2+};

(b)

{Cl}_{2} +

2Li ⟶

{2Li}^{+} +

{2Cl}^{−};

(c)

{3Br}_{2} +

2Fe ⟶

{2Fe}^{3+} +

{6Br}^{−};

(d)

{MnO}_{4}^{—} +

{4H}^{+} +

3Ag ⟶

{3Ag}^{+} +

{MnO}_{2} +

{2H}_{2} O

17-5 §17.1 source ↗

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

Answer

Oxidized: (a) Sn²⁺; (b) Hg; (c) Al; reduced: (a) H₂O₂; (b) PbO₂; (c)

{Cr}_{2} O_{7}^{2−};

oxidizing agent: (a) H₂O₂; (b) PbO₂; (c)

{Cr}_{2} O_{7}^{2−};

reducing agent: (a) Sn²⁺; (b) Hg; (c) Al

17-7 §17.1 source ↗

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

Answer

Oxidized = reducing agent: (a)

{SO}_{3}^{2−};

(b) Mn(OH)₂; (c) H₂; (d) Al; reduced = oxidizing agent: (a) Cu(OH)₂; (b) O₂; (c)

{NO}_{3}^{−};

(d)

{CrO}_{4}^{2−}

17-9 §17.1 source ↗

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

Answer

In basic solution, [OH⁻] > 1

\times

10⁻⁷ M > [H⁺]. Hydrogen ion cannot appear as a reactant because its concentration is essentially zero. If it were produced, it would instantly react with the excess hydroxide ion to produce water. Thus, hydrogen ion should not appear as a reactant or product in basic solution.

17-11 §17.2 source ↗

Write cell schematics for the following cell reactions, using platinum as an inert electrode as needed.

(a) $Mg (s) +$ ${Ni}^{2+} (a q) ⟶$ ${Mg}^{2+} (a q) +$ $Ni (s)$

(b) $2 {Ag}^{+} (a q) +$ $Cu (s) ⟶$ ${Cu}^{2+} (a q) +$ $2Ag (s)$

(d) $3 {CuNO}_{3} (a q) +$ ${Au(NO}_{3})_{3} (a q) ⟶$ ${3Cu(NO}_{3})_{2} (a q) +$ $Au (s)$

Answer

(a)

Mg (s) │ {Mg}^{2+} (a q) ║ {Ni}^{2+} (a q) │ Ni (s);

(b)

Cu (s) │ {Cu}^{2+} (a q) ║ {Ag}^{+} (a q) │ Ag (s);

(c)

Mn (s) │ {Mn}^{2+} (a q) ║ {Sn}^{2+} (a q) │ Sn (s);

(d)

Pt (s) │ {Cu}^{+} (a q {), Cu}^{2+} (a q) ║ {Au}^{3+} (a q) │ Au (s)

17-13 §17.2 source ↗

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

Answer

(a)

Mg (s) +

{Cu}^{2+} (a q) ⟶

{Mg}^{2+} (a q) +

Cu (s);

(b)

2 {Ag}^{+} (a q) +

Ni (s) ⟶

{Ni}^{2+} (a q) +

2Ag (s)

17-15 §17.2 source ↗

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

Answer

Species oxidized = reducing agent: (a) Al(s); (b) NO(g); (c) Mg(s); and (d) MnO₂(s); Species reduced = oxidizing agent: (a) Zr⁴⁺(aq); (b) Ag⁺(aq); (c)

{SiO}_{3}^{2−} (a q)

; and (d)

{ClO}_{3}^{−} (a q)

17-17 §17.2 source ↗

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

Answer

Without the salt bridge, the circuit would be open (or broken) and no current could flow. With a salt bridge, each half-cell remains electrically neutral and current can flow through the circuit.

17-19 §17.2 source ↗

An active (metal) electrode was found to lose mass as the oxidation-reduction reaction was allowed to proceed. Was the electrode an anode or a cathode? Explain.

Answer

Active electrodes participate in the oxidation-reduction reaction. Since metals form cations, the electrode would lose mass if metal atoms in the electrode were to oxidize and go into solution. Oxidation occurs at the anode.

17-21 §17.3 source ↗

Calculate the standard cell potential for each reaction below, and note whether the reaction is spontaneous under standard state conditions.

(a) $Mg (s) +$ ${Ni}^{2+} (a q) ⟶$ ${Mg}^{2+} (a q) +$ $Ni (s)$

(b) $2 {Ag}^{+} (a q) +$ $Cu (s) ⟶$ ${Cu}^{2+} (a q) +$ $2Ag (s)$

(d) $3 {Fe(NO}_{3})_{2} (a q) +$ ${Au(NO}_{3})_{3} (a q) ⟶$ ${3Fe(NO}_{3})_{3} (a q) +$ $Au (s)$

Answer

(a) +2.115 V (spontaneous); (b) +0.4626 V (spontaneous); (c) +1.048 V (spontaneous); (d) +0.727 V (spontaneous)

17-23 §17.3 source ↗

Write the balanced cell reaction for the cell schematic below, calculate the standard cell potential, and note whether the reaction is spontaneous under standard state conditions.

$Cu (s) │ {Cu}^{2+} (a q) ║ {Au}^{3+} (a q) │ Au (s)$

Answer

3 Cu (s) +

{2Au}^{3+} (a q) ⟶

{3Cu}^{2+} (a q) +

2Au (s);

+1.16 V; spontaneous

17-25 §17.3 source ↗

Determine the cell reaction and standard cell potential at 25 °C for a cell made from an anode half-cell containing a cadmium electrode in 1 M cadmium nitrate and a cathode half-cell consisting of an aluminum electrode in 1 M aluminum nitrate solution. Is the reaction spontaneous at standard conditions?

Answer

3 Cd (s) +

{2Al}^{3+} (a q) ⟶

{3Cd}^{2+} (a q) +

2Al (s);

−1.259 V; nonspontaneous

17-27 §17.4 source ↗

For each pair of standard cell potential and electron stoichiometry values below, calculate a corresponding standard free energy change (kJ).

(a) 0.000 V, n = 2

(b) +0.434 V, n = 2

Answer

(a) 0 kJ/mol; (b) −83.7 kJ/mol; (c) +235.3 kJ/mol

17-29 §17.4 source ↗

Determine the standard cell potential and the cell potential under the stated conditions for the electrochemical reactions described here. State whether each is spontaneous or nonspontaneous under each set of conditions at 298.15 K.

(a) $Hg (l) +$ $S^{2−} (a q, 0.10 M) +$ $2 {Ag}^{+} (a q, 0.25 M) ⟶$ $2 Ag (s) +$ $HgS (s)$

(b) The cell made from an anode half-cell consisting of an aluminum electrode in 0.015 M aluminum nitrate solution and a cathode half-cell consisting of a nickel electrode in 0.25 M nickel(II) nitrate solution.

(c) The cell comprised of a half-cell in which aqueous bromide ion (1.0 M) is being oxidized to aqueous bromine (0.11 M) and a half-cell in which Al³⁺ (0.023 M) is being reduced to aluminum metal.

Answer

(a) standard cell potential: 1.50 V, spontaneous; cell potential under stated conditions: 1.43 V, spontaneous; (b) standard cell potential: 1.405 V, spontaneous; cell potential under stated conditions: 1.423 V, spontaneous; (c) standard cell potential: −2.749 V, nonspontaneous; cell potential under stated conditions: −2.733 V, nonspontaneous

17-31 §17.4 source ↗

Use the data in Appendix L to calculate equilibrium constants for the following reactions. Assume 298.15 K if no temperature is given.

(a) $AgCl (s) ⇌$ ${Ag}^{+} (a q) +$ ${Cl}^{−} (a q)$

(b) $CdS (s) ⇌$ ${Cd}^{2+} (a q) +$ $S^{2−} (a q)$ $at 377 K$

(d) $H_{2} O (l) ⇌$ $H^{+} (a q) +$ ${OH}^{−} (a q)$ $at 25 °C$

Answer

(a) 1.7

\times

10⁻¹⁰; (b) 2.6

\times

10⁻²¹; (c) 4.693

\times

10²¹; (d) 1.0

\times

10⁻¹⁴

17-33 §17.5 source ↗

Consider a battery with the overall reaction: $Cu (s) +$ $2 {Ag}^{+} (a q) ⟶$ $2Ag (s) +$ ${Cu}^{2+} (a q) .$

(a) What is the reaction at the anode and cathode?

(b) A battery is “dead” when its cell potential is zero. What is the value of Q when this battery is dead?

Answer

(a)

\begin{array}{l} anode: Cu (s) ⟶ {Cu}^{2+} (a q) + {2e}^{−} E_{anode}^{°} = 0.34 V \\ cathode: 2 \times ({Ag}^{+} (a q) + e^{−} ⟶ Ag (s)) E_{cathode}^{°} = 0.7996 V \end{array};

(b) 3.5

\times

10¹⁵; (c) 5.6

\times

10⁻⁹ M

17-34 §17.5 source ↗

Why do batteries go dead, but fuel cells do not?

Answer

Batteries are self-contained and have a limited supply of reagents to expend before going dead. Alternatively, battery reaction byproducts accumulate and interfere with the reaction. Because a fuel cell is constantly resupplied with reactants and products are expelled, it can continue to function as long as reagents are supplied.

17-36 §17.5 source ↗

Automobile batteries are often rated with regard to “cold-cranking amps,” a measure of the ability of the battery to perform in low-temperature weather. Suggest a reason to explain why a battery’s ability to deliver current is diminished at low temperatures. (Hint: Consider the kinetics of chemical reactions.)

Answer

According to kinetic theory, chemical reactions occur more slowly at lower temperatures due to the decreased rate of collisions between reactant molecules. The rates of the redox reactions taking place within batteries decrease with temperature, resulting in a decreased production of electrical current.

17-38 §17.6 source ↗

Consider the following metals: Ag, Au, Mg, Ni, and Zn. Which of these metals could be used as a sacrificial anode in the cathodic protection of an underground steel storage tank? Steel is an alloy composed mostly of iron, so use −0.447 V as the standard reduction potential for steel.

Answer

Mg and Zn

17-40 §17.6 source ↗

If a sample of iron and a sample of zinc come into contact, the zinc corrodes but the iron does not. If a sample of iron comes into contact with a sample of copper, the iron corrodes but the copper does not. Explain this phenomenon.

Answer

Both examples involve cathodic protection. The (sacrificial) anode is the metal that corrodes (oxidizes or reacts). In the case of iron (−0.447 V) and zinc (−0.7618 V), zinc has a more negative standard reduction potential and so serves as the anode. In the case of iron and copper (0.34 V), iron has the smaller standard reduction potential and so corrodes (serves as the anode).

17-42 §17.6 source ↗

Why would a sacrificial anode made of lithium metal be a bad choice

Answer

While the reduction potential of lithium would make it capable of protecting the other metals, this high potential is also indicative of how reactive lithium is; it would have a spontaneous reaction with most substances. This means that the lithium would react quickly with other substances, even those that would not oxidize the metal it is attempting to protect. Reactivity like this means the sacrificial anode would be depleted rapidly and need to be replaced frequently. (Optional additional reason: fire hazard in the presence of water.)

17-46 §17.7 source ↗

What mass of each product is produced in each of the electrolytic cells of the previous problem if a total charge of 3.33

\times

10⁵ C passes through each cell?

Answer

(a)

\begin{array}{l} mass Ca = 69.1 g \\ {mass Cl}_{2} = 122 g \end{array};

(b)

\begin{array}{l} mass Li = 23.9 g \\ {mass H}_{2} = 3.48 g \end{array};

(c)

\begin{array}{l} mass Al = 31.0 g \\ {mass Cl}_{2} = 122 g \end{array};

(d)

\begin{array}{l} mass Cr = 59.8 g \\ {mass Br}_{2} = 276 g \end{array}

17-48 §17.7 source ↗

A current of 2.345 A passes through the cell shown in Figure 17.19 for 45 minutes. What is the volume of the hydrogen collected at room temperature if the pressure is exactly 1 atm? (Hint: Is hydrogen the only gas present above the water?)

This figure shows an apparatus used for electrolysis. A central chamber with an open top has a vertical column extending below that is nearly full of a clear, colorless liquid, which is labeled “H subscript 2 O plus H subscript 2 S O subscript 4.” A horizontal tube in the apparatus connects the central region to vertical columns to the left and right, each of which has a valve or stopcock at the top and a stoppered bottom. On the left, the stopper at the bottom has a small brown square connected just above it in the liquid. The square is labeled “Anode plus.” A black wire extends from the stopper at the left to a rectangle which is labeled “Voltage source” on to the stopper at the right. The left side of the rectangle is labeled with a plus symbol and the right side is labeled with a negative sign. The stopper on the right also has a brown square connected to it which is in the liquid in the apparatus. This square is labeled “Cathode negative.” The level of the solution on the left arm or tube of the apparatus is significantly higher than the level of the right arm. Bubbles are present near the surface of the liquid on each side of the apparatus, with the bubbles labeled as “O subscript 2 ( g )” on the left and “H subscript 2 ( g )” on the right. — Figure 17.19 in the book ↗

Answer

0.79 L

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Problems and answers are adapted from Chemistry 2e by Paul Flowers, Klaus Theopold, Richard Langley and William R. Robinson, OpenStax, and are used under the CC BY-NC-SA 4.0 licence. Access for free at https://openstax.org/books/chemistry-2e/pages/1-introduction. Illustrations: Copyright Rice University, OpenStax, under CC BY-NC-SA 4.0 license. This adaptation is released under the same licence, is non-commercial, and is neither affiliated with nor endorsed by OpenStax or Rice University.