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Chapter 17: Problem 109

What reactions take place at the cathode and the anode when each of the following is electrolyzed? a. molten \(\mathrm{NiBr}_{2}\) b. molten \(\mathrm{AlF}_{3}\) c. molten \(\mathrm{MnI}_{2}\)

Short Answer

Expert verified
For each compound, the electrolysis reactions at the cathode and anode are: a. \(\mathrm{NiBr}_{2}\): Cathode: \(\mathrm{Ni^{2+}} + 2e^{-} \rightarrow \mathrm{Ni}\) Anode: \(2\: \mathrm{Br^{-}} \rightarrow \mathrm{Br_{2}} + 2e^{-}\) b. \(\mathrm{AlF}_{3}\): Cathode: \(\mathrm{Al^{3+}} + 3e^{-} \rightarrow \mathrm{Al}\) Anode: \(2\: \mathrm{F^{-}} \rightarrow \mathrm{F_{2}} + 2e^{-}\) c. \(\mathrm{MnI}_{2}\): Cathode: \(\mathrm{Mn^{2+}} + 2e^{-} \rightarrow \mathrm{Mn}\) Anode: \(2\: \mathrm{I^{-}} \rightarrow \mathrm{I_{2}} + 2e^{-}\)

Step by step solution

01

Understand the molten compounds

The molten compounds provided are: a. \(\mathrm{NiBr}_{2}\: (Nickel\,bromide)\) b. \(\mathrm{AlF}_{3}\: (Aluminum\,fluoride)\) c. \(\mathrm{MnI}_{2}\: (Manganese\,iodide)\) Next, identify the ions present in the compounds.
02

Identify the ions present in each compound

The ions present in the compounds are: a. \(\mathrm{NiBr}_{2}\): \(\mathrm{Ni^{2+}}\) and \(\mathrm{Br^{-}}\) b. \(\mathrm{AlF}_{3}\): \(\mathrm{Al^{3+}}\) and \(\mathrm{F^{-}}\) c. \(\mathrm{MnI}_{2}\): \(\mathrm{Mn^{2+}}\) and \(\mathrm{I^{-}}\)
03

Determine the cathode reaction

During electrolysis, metal cations are reduced at the cathode. Reduction occurs by gaining electrons. The cathode reactions for the molten compounds are as follows: a. \(\mathrm{NiBr}_{2}\): \(\mathrm{Ni^{2+}} + 2e^{-} \rightarrow \mathrm{Ni}\) b. \(\mathrm{AlF}_{3}\): \(\mathrm{Al^{3+}} + 3e^{-} \rightarrow \mathrm{Al}\) c. \(\mathrm{MnI}_{2}\): \(\mathrm{Mn^{2+}} + 2e^{-} \rightarrow \mathrm{Mn}\)
04

Determine the anode reaction

During electrolysis, non-metal anions are oxidized at the anode. Oxidation occurs by losing electrons to form elemental substances. The anode reactions for the molten compounds are: a. \(\mathrm{NiBr}_{2}\): \(2\: \mathrm{Br^{-}} \rightarrow \mathrm{Br_{2}} + 2e^{-}\) b. \(\mathrm{AlF}_{3}\): \(2\: \mathrm{F^{-}} \rightarrow \mathrm{F_{2}} + 2e^{-}\) c. \(\mathrm{MnI}_{2}\): \(2\: \mathrm{I^{-}} \rightarrow \mathrm{I_{2}} + 2e^{-}\)
05

Combine the cathode and anode reactions

For each compound, combine the cathode and anode reactions: a. \(\mathrm{NiBr}_{2}\) electrolysis reactions: Cathode: \(\mathrm{Ni^{2+}} + 2e^{-} \rightarrow \mathrm{Ni}\) Anode: \(2\: \mathrm{Br^{-}} \rightarrow \mathrm{Br_{2}} + 2e^{-}\) b. \(\mathrm{AlF}_{3}\) electrolysis reactions: Cathode: \(\mathrm{Al^{3+}} + 3e^{-} \rightarrow \mathrm{Al}\) Anode: \(2\: \mathrm{F^{-}} \rightarrow \mathrm{F_{2}} + 2e^{-}\) c. \(\mathrm{MnI}_{2}\) electrolysis reactions: Cathode: \(\mathrm{Mn^{2+}} + 2e^{-} \rightarrow \mathrm{Mn}\) Anode: \(2\: \mathrm{I^{-}} \rightarrow \mathrm{I_{2}} + 2e^{-}\)

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Most popular questions from this chapter

A galvanic cell is based on the following half-reactions: $$\begin{array}{ll} \mathrm{Cu}^{2+}(a q)+2 \mathrm{e}^{-} \longrightarrow \mathrm{Cu}(s) & \mathscr{E}^{\circ}=0.34 \mathrm{V} \\ \mathrm{V}^{2+}(a q)+2 \mathrm{e}^{-} \longrightarrow \mathrm{V}(s) & \mathscr{E}^{\circ}=-1.20 \mathrm{V} \end{array}$$ In this cell, the copper compartment contains a copper electrode and \(\left[\mathrm{Cu}^{2+}\right]=1.00 \mathrm{M},\) and the vanadium compartment contains a vanadium electrode and \(V^{2+}\) at an unknown concentration. The compartment containing the vanadium \((1.00 \mathrm{L}\) of solution) was titrated with \(0.0800 M \space \mathrm{H}_{2} \mathrm{EDTA}^{2-},\) resulting in the reaction $$\mathrm{H}_{2} \mathrm{EDTA}^{2-}(a q)+\mathrm{V}^{2+}(a q) \rightleftharpoons \mathrm{VEDTA}^{2-}(a q)+2 \mathrm{H}^{+}(a q) \space \mathrm{K=?}$$ The potential of the cell was monitored to determine the stoichiometric point for the process, which occurred at a volume of \(500.0 \mathrm{mL} \space \mathrm{H}_{2} \mathrm{EDTA}^{2-}\) solution added. At the stoichiometric point, \(\mathscr{E}_{\text {cell }}\) was observed to be \(1.98 \mathrm{V}\). The solution was buffered at a pH of \(10.00 .\) a. Calculate \(\mathscr{E}_{\text {cell }}\) before the titration was carried out. b. Calculate the value of the equilibrium constant, \(K,\) for the titration reaction. c. Calculate \(\mathscr{E}_{\text {cell }}\) at the halfway point in the titration.

A factory wants to produce \(1.00 \times 10^{3} \mathrm{kg}\) barium from the electrolysis of molten barium chloride. What current must be applied for \(4.00 \mathrm{h}\) to accomplish this?

Consider the cell described below: $$ \mathrm{Zn}\left|\mathrm{Zn}^{2+}(1.00 M)\right|\left|\mathrm{Cu}^{2+}(1.00 M)\right| \mathrm{Cu} $$ Calculate the cell potential after the reaction has operated long enough for the \(\left[\mathrm{Zn}^{2+}\right]\) to have changed by \(0.20 \mathrm{mol} / \mathrm{L}\). (Assume \(T=25^{\circ} \mathrm{C} .\) )

Which of the following statement(s) is/are true? a. Copper metal can be oxidized by \(\mathrm{Ag}^{+}\) (at standard conditions). b. In a galvanic cell the oxidizing agent in the cell reaction is present at the anode. c. In a cell using the half reactions \(A l^{3+}+3 e^{-} \longrightarrow A l\) and \(\mathrm{Mg}^{2+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{Mg},\) aluminum functions as the anode. d. In a concentration cell electrons always flow from the compartment with the lower ion concentration to the compartment with the higher ion concentration. e. In a galvanic cell the negative ions in the salt bridge flow in the same direction as the electrons.

The equation \(\Delta G^{\circ}=-n F \mathscr{E}^{\circ}\) also can be applied to halfreactions. Use standard reduction potentials to estimate \(\Delta G_{\mathrm{r}}^{\circ}\) for \(\mathrm{Fe}^{2+}(a q)\) and \(\mathrm{Fe}^{3+}(a q) .\left(\Delta G_{\mathrm{f}}^{\circ} \text { for } \mathrm{e}^{-}=0 .\right)\)
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