Question

What reactions take place at the cathode and the anode when each of the following is electrolyzed? (Assume standard conditions.) a. \(1.0 M \mathrm{NiBr}_{2}\) solution b. \(1.0 M \mathrm{AlF}_{3}\) solution c. \(1.0 M \mathrm{MnI}_{2}\) solution

Short answer

For the electrolysis of the given solutions, the most likely reactions to occur at the cathode and anode are as follows: a. \(1.0 M \mathrm{NiBr_2}\) - Cathode: \(\mathrm{Ni^{2+}} + 2e^- \rightarrow \mathrm{Ni}\) - Anode: \(2\mathrm{Br^-} \rightarrow \mathrm{Br_2} + 2e^-\) b. \(1.0 M \mathrm{AlF_3}\) - Cathode: \(\mathrm{Al^{3+}} + 3e^- \rightarrow \mathrm{Al}\) - Anode: \(2\mathrm{H_2O} \rightarrow \mathrm{O_2} + 4\mathrm{H^+}+ 4e^-\) c. \(1.0 M \mathrm{MnI_2}\) - Cathode: \(\mathrm{Mn^{2+}} + 2e^- \rightarrow \mathrm{Mn}\) - Anode: \(2\mathrm{I^-} \rightarrow \mathrm{I_2} + 2e^-\)

Step-by-step solution

Identify ions in solution

In a \(1.0 M \mathrm{NiBr}_2\) solution, the ions present are \(\mathrm{Ni^{2+}}\) and \(\mathrm{Br^-}\).

Determine possible redox reactions

At the cathode, possible reduction reactions are: - \(\mathrm{Ni^{2+}} + 2e^- \rightarrow \mathrm{Ni}\) - \(2\mathrm{H_2O} + 2e^- \rightarrow \mathrm{H_2} + 2\mathrm{OH^-}\) At the anode, possible oxidation reactions are: - \(2\mathrm{Br^-} \rightarrow \mathrm{Br_2} + 2e^-\) - \(2\mathrm{H_2O} \rightarrow \mathrm{O_2} + 4\mathrm{H^+}+ 4e^-\)

Identify most likely reactions

Using standard reduction potentials, we find that the most likely reactions to occur are: Cathode: \(\mathrm{Ni^{2+}} +2e^- \rightarrow \mathrm{Ni}\) Anode: \(2\mathrm{Br^-} \rightarrow \mathrm{Br_2} + 2e^-\) b. \(1.0 M \mathrm{AlF}_3\) solution

Identify ions in solution

In a \(1.0 M \mathrm{AlF}_3\) solution, the ions present are \(\mathrm{Al^{3+}}\) and \(\mathrm{F^-}\).

Determine possible redox reactions

At the cathode, possible reduction reactions are: - \(\mathrm{Al^{3+}} + 3e^- \rightarrow \mathrm{Al}\) - \(2\mathrm{H_2O} + 2e^- \rightarrow \mathrm{H_2} + 2\mathrm{OH^-}\) At the anode, possible oxidation reactions are: - \(2\mathrm{F^-} \rightarrow \mathrm{F_2} + 2e^-\) - \(2\mathrm{H_2O} \rightarrow \mathrm{O_2} + 4\mathrm{H^+}+ 4e^-\)

Identify most likely reactions

Using standard reduction potentials, we find that the most likely reactions to occur are: Cathode: \(\mathrm{Al^{3+}} + 3e^- \rightarrow \mathrm{Al}\) Anode: \(2\mathrm{H_2O} \rightarrow \mathrm{O_2} + 4\mathrm{H^+}+ 4e^-\) (since the oxidation of \(\mathrm{F^-}\) requires extremely high potentials and is unlikely to occur under standard conditions) c. \(1.0 M \mathrm{MnI}_2\) solution

Identify ions in solution

In a \(1.0 M \mathrm{MnI}_2\) solution, the ions present are \(\mathrm{Mn^{2+}}\) and \(\mathrm{I^-}\).

Determine possible redox reactions

At the cathode, possible reduction reactions are: - \(\mathrm{Mn^{2+}} + 2e^- \rightarrow \mathrm{Mn}\) - \(2\mathrm{H_2O} + 2e^- \rightarrow \mathrm{H_2} + 2\mathrm{OH^-}\) At the anode, possible oxidation reactions are: - \(2\mathrm{I^-} \rightarrow \mathrm{I_2} + 2e^-\) - \(2\mathrm{H_2O} \rightarrow \mathrm{O_2} + 4\mathrm{H^+}+ 4e^-\)

Identify most likely reactions

Using standard reduction potentials, we find that the most likely reactions to occur are: Cathode: \(\mathrm{Mn^{2+}} + 2e^- \rightarrow \mathrm{Mn}\) Anode: \(2\mathrm{I^-} \rightarrow \mathrm{I_2} + 2e^-\)

Key concepts

Standard Reduction Potentials

Understanding standard reduction potentials is essential when examining electrolysis reactions. These potentials, typically listed in volts, help predict which species will be reduced at the cathode during electrolysis. They are measured under standard conditions, which include a temperature of 298 K, a concentration of 1 M for aqueous species, and a pressure of 1 atm for gases.

The more positive the standard reduction potential, the greater the species' tendency to gain electrons and be reduced. Conversely, species with lower reduction potentials are more likely to lose electrons and be oxidized. In the exercise above, the selection of nickel, aluminum, and manganese for reduction at the cathode over the possible reduction of water was informed by comparing their standard reduction potentials.

Cathode and Anode Reactions

In electrolysis, the cathode acts as the site of reduction, while the anode is where oxidation occurs. At the cathode, cations are attracted and gain electrons, resulting in a reduction reaction. For example, as seen in the exercise, metal ions like \(\mathrm{Ni^{2+}}\), \(\mathrm{Al^{3+}}\), and \(\mathrm{Mn^{2+}}\) are accepted to gain electrons and transform into their metallic forms.

At the anode, anions lose electrons and are oxidized, leading to the generation of diatomic halogen gases in the given solutions. However, when the anions are difficult to oxidize, like fluoride ions, water is oxidized instead, releasing oxygen gas and hydrogen ions.

Redox Reactions in Electrolysis

Redox reactions are at the heart of electrolysis, involving the transfer of electrons between species. For a redox reaction to occur, there must be a reductant (electron donor) and an oxidant (electron acceptor). During electrolysis, the external power supply forces the electron transfer, driving non-spontaneous reactions.

This process splits compounds into their constituent elements, as indicated by the reduction of metal ions at the cathode and the generation of halogen gases or oxygen at the anode. The competition between water and other ions for reduction or oxidation is resolved by referencing standard reduction potentials.

Standard Conditions in Electrochemistry

Standard conditions refer to a set of specific parameters under which electrochemical reactions are measured to ensure consistency and comparability. These include a solute concentration of 1 M, gas pressures at 1 atm, and a temperature of 25°C (298 K). In electrolysis, these conditions are assumed unless otherwise stated.

Under standard conditions, the behavior of electrochemical cells can be predicted with the help of standard reduction potentials. It is important to note, however, that actual reaction conditions may vary, altering the expected outcomes. For instance, temperature and concentration changes can shift potentials and change which reactions are favorable.