Question

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

For each compound, the following half-reactions occur: a. molten NiBr2: Cathode (reduction): \(\mathrm{Ni^{2+}(l) + 2e^{-} \rightarrow Ni(s)}\) Anode (oxidation): \(\mathrm{2Br^{-}(l) \rightarrow Br_{2}(g) + 2e^{-}}\) b. molten AlF3: Cathode (reduction): \(\mathrm{Al^{3+}(l) + 3e^{-} \rightarrow Al(s)}\) Anode (oxidation): \(\mathrm{2F^{-}(l) \rightarrow F_{2}(g) + 2e^{-}}\) c. molten MnI2: Cathode (reduction): \(\mathrm{Mn^{2+}(l) + 2e^{-} \rightarrow Mn(s)}\) Anode (oxidation): \(\mathrm{2I^{-}(l) \rightarrow I_{2}(g) + 2e^{-}}\)

Step-by-step solution

(a. Analyzing molten NiBr2)

(In molten NiBr2, we have two ions: Ni(II) ions and Br- ions.)

(a. Determining reactions at the cathode and anode)

(For molten NiBr2, reduction occurs at the cathode and oxidation occurs at the anode. Reduction of Ni(II) ions produces nickel metal, and oxidation of Br- ions produces bromine gas.) Half-reactions for this process are: Cathode (reduction): \(\mathrm{Ni^{2+}(l) + 2e^{-} \rightarrow Ni(s)}\) Anode (oxidation): \(\mathrm{2Br^{-}(l) \rightarrow Br_{2}(g) + 2e^{-}}\)

(b. Analyzing molten AlF3)

(In molten AlF3, we have two ions: Al(III) ions and F- ions.)

(b. Determining reactions at the cathode and anode)

(For molten AlF3, reduction occurs at the cathode and oxidation occurs at the anode. Reduction of Al(III) ions produces aluminum metal, and oxidation of F- ions produces fluorine gas.) Half-reactions for this process are: Cathode (reduction): \(\mathrm{Al^{3+}(l) + 3e^{-} \rightarrow Al(s)}\) Anode (oxidation): \(\mathrm{2F^{-}(l) \rightarrow F_{2}(g) + 2e^{-}}\)

(c. Analyzing molten MnI2)

(In molten MnI2, we have two ions: Mn(II) ions and I- ions.)

(c. Determining reactions at the cathode and anode)

(For molten MnI2, reduction occurs at the cathode and oxidation occurs at the anode. Reduction of Mn(II) ions produces manganese metal, and oxidation of I- ions produces iodine gas.) Half-reactions for this process are: Cathode (reduction): \(\mathrm{Mn^{2+}(l) + 2e^{-} \rightarrow Mn(s)}\) Anode (oxidation): \(\mathrm{2I^{-}(l) \rightarrow I_{2}(g) + 2e^{-}}\)

Key concepts

Electrochemical Cells

Electrochemical cells are fundamental to the field of electrochemistry, providing physical sites for redox reactions where chemical energy is converted into electrical energy, or vice versa. These cells comprise two electrodes, an anode and a cathode, submerged in an electrolyte that allows ions to move freely.

Through a process known as electrolysis, an external voltage source forces a non-spontaneous reaction to occur. The key to understanding electrolysis is to recognize that it involves the movement of electrons through an external circuit, driven by the power supply, while ions migrate within the electrolyte to balance the charge.

For instance, when a molten ionic compound such as \(\mathrm{NiBr}_{2}\) is subjected to electrolysis, the Ni(II) ions are attracted to the cathode where they gain electrons, and Br- ions are attracted to the anode where they surrender electrons.

Reduction at the Cathode

At the cathode, the reduction reaction takes place. This is where cations (positive ions) are attracted and electrons from the external circuit are gained in a process governed by the principle 'Red Cat' - reduction occurs at the cathode. The gain of electrons by a substance is its reduction.

For example, in the electrolysis of molten \(\mathrm{NiBr}_{2}\), \(\mathrm{Ni^{2+}}\) ions are reduced to nickel metal by gaining two electrons. The half-reaction can be represented as:\[\mathrm{Ni^{2+}(l) + 2e^{-} \rightarrow Ni(s)}\].

Likewise, in the electrolysis of molten \(\mathrm{AlF}_{3}\) and \(\mathrm{MnI}_{2}\), \(\mathrm{Al^{3+}}\) and \(\mathrm{Mn^{2+}}\) ions undergo reduction at the cathode to form aluminum and manganese metal, respectively.

Oxidation at the Anode

The anode is the site of oxidation, the half of the redox reaction involving the loss of electrons. Since anions (negative ions) are attracted to the anode, this is where they give up their electrons, completing the circuit established by the external power source. The mnemonic 'An Ox' helps students remember that 'Anode Oxidation'.

Consider the anode reaction in the electrolysis of molten \(\mathrm{NiBr}_{2}\), the Br- ions oxidize, releasing their electrons and forming bromine gas, detailed by the half-reaction:\[\mathrm{2Br^{-}(l) \rightarrow Br_{2}(g) + 2e^{-}}\].

Similarly, when electrolyzing \(\mathrm{AlF}_{3}\) and \(\mathrm{MnI}_{2}\), fluoride and iodide ions respectively are oxidized at the anode to form fluorine and iodine gases.

Half-Reactions

Half-reactions are the individual processes that make up the overall redox reaction occurring in electrochemical cells. They are split into two types: reduction half-reactions, which occur at the cathode, and oxidation half-reactions, which occur at the anode.

By separating the two, it becomes easier to understand and balance the reactions. Students can analyze the movement of electrons and the changes in oxidation states. These half-reactions can be combined to yield the overall cell reaction.

For example, in the electrolysis of \(\mathrm{NiBr}_{2}\), you have the reduction of \(\mathrm{Ni^{2+}}\) and the oxidation of \(\mathrm{Br^{-}}\), which merge into a full reaction that reflects the conservation of charge and mass, showcasing the transformation from reactants to products.