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Chapter 18: Problem 108

What reaction will take place at the cathode and the anode when each of the following is electrolyzed? a. molten \(\mathrm{KF} \quad\) b. molten \(\mathrm{CuCl}_{2} \quad\) c. molten \(\mathrm{MgI}_{2}\)

Short Answer

Expert verified
During the electrolysis of the following molten salts: a. KF: Potassium metal forms at the cathode (\(K^+ + e^- \rightarrow K (s)\)) and fluoride gas forms at the anode (\(2F^- \rightarrow F_2(g) + 2e^-\)). b. CuCl2: Copper metal forms at the cathode (\(Cu^{2+} + 2e^- \rightarrow Cu (s)\)) and chloride gas forms at the anode (\(2Cl^- \rightarrow Cl_2(g) + 2e^-\)). c. MgI2: Magnesium metal forms at the cathode (\(Mg^{2+} + 2e^- \rightarrow Mg (s)\)) and iodine gas forms at the anode (\(2I^- \rightarrow I_2(g) + 2e^-\)).

Step by step solution

01

Electrolysis of molten KF

Cathode reaction (reduction): The cation in KF is K+ (potassium ion), which will migrate to the cathode and gain an electron to form potassium metal. \[K^+ + e^- \rightarrow K (s)\] Anode reaction (oxidation): The anion in KF is F- (fluoride ion), which will migrate to the anode, lose an electron, and form fluoride gas. \[2F^- \rightarrow F_2(g) + 2e^-\] So, during the electrolysis of molten KF, potassium metal will form at the cathode, and fluoride gas will form at the anode.
02

Electrolysis of molten CuCl2

Cathode reaction (reduction): The cation in CuCl2 is Cu2+ (copper (II) ion), which will migrate to the cathode and gain two electrons to form copper metal. \[Cu^{2+} + 2e^- \rightarrow Cu (s)\] Anode reaction (oxidation): The anion in CuCl2 is Cl- (chloride ion), which will migrate to the anode, lose an electron, and form chloride gas. \[2Cl^- \rightarrow Cl_2(g) + 2e^-\] So, during the electrolysis of molten CuCl2, copper metal will form at the cathode, and chloride gas will form at the anode.
03

Electrolysis of molten MgI2

Cathode reaction (reduction): The cation in MgI2 is Mg2+ (magnesium ion), which will migrate to the cathode and gain two electrons to form magnesium metal. \[Mg^{2+} + 2e^- \rightarrow Mg (s)\] Anode reaction (oxidation): The anion in MgI2 is I- (iodide ion), which will migrate to the anode, lose an electron, and form iodine gas. \[2I^- \rightarrow I_2(g) + 2e^-\] So, during the electrolysis of molten MgI2, magnesium metal will form at the cathode, and iodine gas will form at the anode.

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Most popular questions from this chapter

Three electrochemical cells were connected in series so that the same quantity of electrical current passes through all three cells. In the first cell, 1.15 g chromium metal was deposited from a chromium (III) nitrate solution. In the second cell, 3.15 \(\mathrm{g}\) osmium was deposited from a solution made of \(\mathrm{Os}^{n+}\) and nitrate ions. What is the name of the salt? In the third cell, the electrical charge passed through a solution containing \(\mathrm{X}^{2+}\) ions caused deposition of 2.11 \(\mathrm{g}\) metallic \(\mathrm{X}\) . What is the electron configuration of \(\mathrm{X} ?\)

The following standard reduction potentials have been determined for the aqueous chemistry of indium: $$\operatorname{In}^{3+}(a q)+2 \mathrm{e}^{-} \longrightarrow \operatorname{In}^{+}(a q) \quad \mathscr{E}^{\circ}=-0.444 \mathrm{V}$$ $$\operatorname{In}^{+}(a q)+\mathrm{e}^{-} \longrightarrow \operatorname{In}(s) \qquad \quad \mathscr{E}^{\circ}=-0.126 \mathrm{V}$$ a. What is the equilibrium constant for the disproportionation reaction, where a species is both oxidized and reduced, shown below? $$3 \ln ^{+}(a q) \longrightarrow 2 \operatorname{In}(s)+\operatorname{In}^{3+}(a q)$$ b. What is \(\Delta G_{i}^{\circ}\) for \(\operatorname{In}^{+}(a q)\) if $\Delta G_{f}^{\circ}=-97.9 \mathrm{kJ} / \mathrm{mol}\( for \)\operatorname{In}^{3+}(a q) ?$

If the cell potential is proportional to work and the standard reduction potential for the hydrogen ion is zero, does this mean that the reduction of the hydrogen ion requires no work?

Calculate \(\mathscr{E}^{\circ}\) values for the following cells. Which reactions are spontaneous as written (under standard conditions)? Balance the equations. Standard reduction potentials are found in Table 18.1. a. $\mathrm{MnO}_{4}^{-(a q)}+\mathrm{I}^{-}(a q) \longrightarrow \mathrm{I}_{2}(a q)+\mathrm{Mn}^{2+}(a q)$ b. $\mathrm{MnO}_{4}^{-}(a q)+\mathrm{F}^{-}(a q) \longrightarrow \mathrm{F}_{2}(g)+\mathrm{Mn}^{2+}(a q)$

The overall reaction in the lead storage battery is $\mathrm{Pb}(s)+\mathrm{PbO}_{2}(s)+2 \mathrm{H}^{+}(a q)+2 \mathrm{HSO}_{4}^{-}(a q) \longrightarrow$ $$\quad\quad\quad\quad\quad\quad\quad 2 \mathrm{PbSO}_{4}(s)+2 \mathrm{H}_{2} \mathrm{O}(l)$$ a. For the cell reaction \(\Delta H^{\circ}=-315.9 \mathrm{kJ}\) and $\Delta S^{\circ}=\( 263.5 \)\mathrm{J} / \mathrm{K}\( . Calculate \)\mathscr{E}^{\circ} \mathrm{at}-20 .^{\circ} \mathrm{C}\( . Assume \)\Delta H^{\circ}\( and \)\Delta S^{\circ}$ do not depend on temperature. b. Calculate \(\mathscr{E}\) at \(-20 .^{\circ} \mathrm{C}\) when \(\left[\mathrm{HSO}_{4}^{-}\right]=\left[\mathrm{H}^{+}\right]=4.5 \mathrm{M}\) . c. Consider your answer to Exercise 69. Why does it seem that batteries fail more often on cold days than on warm days?
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