Question

(a) What is electrolysis? (b) Are electrolysis reactions thermodynamically spontaneous? Explain. (c) What process occurs at the anode in the electrolysis of molten NaCl? (d) Why is sodium metal not obtained when an aqueous solution of \(\mathrm{NaCl}\) undergoes electrolysis?

Short answer

(a) Electrolysis is a process where electrical energy is used to drive a non-spontaneous chemical reaction, causing redox reactions at the electrodes. (b) Electrolysis reactions are generally not thermodynamically spontaneous as they have positive Gibbs free energy change (\(ΔG > 0\)) and require external energy input. (c) The oxidation of chloride ions (\(Cl^-\)) into chlorine gas (\(Cl_2\)) occurs at the anode during the electrolysis of molten NaCl. (d) Sodium metal is not obtained in the electrolysis of aqueous NaCl because the reduction of water, producing hydrogen gas and hydroxide ions, is more favorable than the reduction of sodium ions in this scenario.

Step-by-step solution

(a) Definition of Electrolysis

Electrolysis is a process in which electrical energy is used to drive a non-spontaneous chemical reaction. During electrolysis, an electric current is passed through an electrolyte, causing the ions in the electrolyte to move towards their respective electrodes. This results in redox reactions at the electrodes, where the reduction occurs at the cathode (negative electrode) and oxidation occurs at the anode (positive electrode).

(b) Thermodynamic Spontaneity of Electrolysis Reactions

Electrolysis reactions are generally not thermodynamically spontaneous. Spontaneous reactions have a negative Gibbs free energy change (\(ΔG < 0\)), meaning they occur without the need for external energy input. On the other hand, electrolysis reactions usually have a positive Gibbs free energy change (\(ΔG > 0\)), indicating that they are non-spontaneous and require an external energy input, such as an electrical current, to occur.

(c) Anode Process in Electrolysis of Molten NaCl

During the electrolysis of molten NaCl, oxidation occurs at the anode. In this case, chloride ions (\(Cl^-\)) lose electrons to become chlorine gas (\(Cl_2\)). The half-reaction that takes place at the anode can be written as: \[2Cl^-(l)\rightarrow Cl_{2(g)} + 2e^-\]

(d) Lack of Sodium Metal Formation in Aqueous NaCl Electrolysis

When an aqueous solution of NaCl undergoes electrolysis, sodium metal is not produced due to the presence of water. In aqueous solutions, water molecules can also undergo reduction and oxidation reactions at the electrodes. At the cathode, water is more easily reduced than sodium ions, producing hydrogen gas and hydroxide ions: \[2H_2O(l) + 2e^-\rightarrow H_{2(g)} + 2OH^-(aq)\] At the anode, the oxidation of the chloride ions still takes place, producing chlorine gas as previously mentioned. Since the reduction of water is more favorable in this scenario, the production of sodium metal is not observed.

Key concepts

Thermodynamic Spontaneity

When discussing chemical reactions, thermodynamic spontaneity is an essential concept to grasp. It determines whether a reaction can take place without the addition of external energy. For a process to be considered spontaneous, it must have a negative Gibbs free energy change, represented by the symbol \(ΔG < 0\). This negative value indicates that the reaction is energetically favorable and can occur on its own.

Electricity is needed for electrolysis because it's a non-spontaneous process, which implies that \(ΔG > 0\), meaning the reaction requires an input of energy to proceed. This is counterintuitive because many students of chemistry are familiar with spontaneous reactions that happen readily, such as a piece of sodium reacting vigorously with water. In electrolysis, however, we are forcing a reaction to occur that wouldn't happen under normal circumstances.

Oxidation and Reduction Reactions

The reactions that happen during electrolysis are fundamentally redox (reduction-oxidation) processes. Understanding these is vital for dissecting what occurs in electrolysis. An oxidation reaction is where a substance loses electrons, while reduction is the gain of electrons by a substance.

In our study of electrolysis, these redox processes occur at two locations within the cell: the anode, where oxidation takes place, and the cathode, where reduction happens. Electrons are transferred from the substance being oxidized to the substance being reduced, facilitated by the flow of electric current. For example, in a basic scenario, a metal ion in solution might gain electrons at the cathode to form a solid metal, while a non-metal ion might lose electrons at the anode to form a non-metal element.

Molten NaCl Electrolysis

The electrolysis of molten sodium chloride (NaCl) serves as an instructive example. When NaCl is melted, it separates into its constituent ions, sodium (\(Na^+\)) and chloride (\(Cl^-\)). Upon the application of an electric current, sodium ions migrate towards the cathode where they gain electrons (reduced) to form molten sodium metal, while chloride ions travel to the anode where they give up electrons (oxidized) and form chlorine gas.

To visualize, consider the two half-reactions that take place: at the cathode, \[Na^+ + e^-\rightarrow Na(l)\], and at the anode, \[2Cl^-\rightarrow Cl_{2(g)} + 2e^-\]. This process is used commercially to produce both elemental sodium and chlorine gas. The separation of ions in the molten state is key here, as it enables each ion to participate in the redox reactions without competition from other species, such as water in an aqueous solution.

Aqueous Solution Electrolysis

Electrolysis becomes a bit more complex when performed in aqueous solutions, as water itself can participate in redox reactions. Why can't we obtain sodium metal from saline water, despite sodium being present? The answer lies in the preferential discharge of ions.

Water molecules can be both oxidized and reduced more easily compared to certain ions in the solution. During the electrolysis of aqueous \(NaCl\), at the cathode, water molecules instead of sodium ions are reduced to form hydrogen gas and hydroxide ions with the half-reaction \[2H_2O(l) + 2e^-\rightarrow H_{2(g)} + 2OH^-\]. At the anode, chloride ions are oxidized to form chlorine gas, but because the reduction of water happens more readily than that of sodium ions, no sodium metal is produced.

These preferences of ion reduction and oxidation dramatically affect the outcome of electrolysis in aqueous solutions, showcasing the importance of understanding the chemical context in which electrolysis is conducted.