A zinc anode is used in this experiment because zinc is less reactive than copper. This difference in reactivity is crucial for the electroplating process. The anode must be oxidized to provide the copper ions needed for reduction at the cathode (the steel nail). Zinc is more readily oxidized than copper, meaning it will lose electrons more easily. This ensures that copper ions from the solution are preferentially reduced to copper metal on the steel nail, resulting in a copper coating.

The reaction occurring at the anode is oxidation of zinc. The half-equation for this reaction is:

Electroplating is a process where a thin layer of metal is deposited onto a conductive object using electrolysis. The key components and their roles are:

• Electrolyte Solution: This is an aqueous solution containing ions of the metal to be deposited. For example, a solution of copper sulfate (CuSO₄) would be used to electroplate with copper ions.
• Anode: This is the object to be coated with the metal. It is connected to the positive terminal of a DC power supply. The anode dissolves, releasing metal ions into the electrolyte.
• Cathode: This is a conductive object (usually made of the metal being deposited) immersed in the electrolyte solution. It is connected to the negative terminal of the DC power supply. Metal ions from the electrolyte are attracted to the cathode and gain electrons, depositing as a thin layer of metal.
• DC Power Supply: Provides the electrical potential difference to drive the electrolytic reaction.

During electroplating, the anode dissolves, releasing metal ions into the electrolyte. These ions migrate towards the cathode, where they gain electrons and deposit as a metallic coating on the cathode's surface. The process continues until a sufficient thickness of metal has been deposited.

The key difference lies in the mobility of the ions. In aqueous solutions, ions are surrounded by solvent molecules (water), which allows them to move and participate in the electrolytic reactions. This mobility enables the formation of gases at the electrodes. In contrast, molten ionic compounds have ions that are free to move throughout the molten state. Therefore, the ions can easily migrate to the electrodes and undergo the required redox reactions without the need for a gaseous product to be formed.

Electrolysis of aqueous solution (gas production): An example is the electrolysis of aqueous sodium chloride (NaCl) as described in the previous question. Chlorine gas is produced at the anode and hydrogen gas is produced at the cathode.

Electrolysis of molten ionic compound (no gas production): An example is the electrolysis of molten aluminium chloride (AlCl₃). At the cathode, aluminium ions (Al³⁺) are reduced to molten aluminium (Al), which then solidifies. At the anode, chloride ions (Cl^-) are oxidized to chlorine gas. However, the chlorine gas produced is immediately attracted to the cathode and reacts with the molten aluminium to form molten aluminium chloride, preventing it from escaping as a gas. The overall reaction is: