Practical: Electroplating a Metal Object

The Electroplating Cell Setup

Setting up the cell correctly is like wiring a plug; getting the positive and negative sides mixed up completely ruins the result. The object you want to plate (such as an iron spoon or a brass key) must be connected to the negative terminal to become the cathode. You can remember this with the rule: "Coat the Cathode".

The pure metal you want to use for the plating (like pure silver or pure copper) is connected to the positive terminal, acting as the anode. Both electrodes are submerged in an electrolyte. This cannot just be pure water or any random liquid; it must be an aqueous solution of a soluble salt that contains the exact same metal ions as the plating metal (for example, using copper(II) sulfate when copper plating).

Step-by-Step Experimental Procedure

To get full marks when asked to describe this practical, you must outline the steps in a logical sequence. First, thoroughly clean the metal object (the cathode) using emery paper or steel wool to strip away any resistive oxide layer. Then, degrease it using propanone or detergent to ensure the new metal layer will adhere smoothly to the surface.

Next, suspend both the cleaned object and the pure plating metal into your chosen electrolyte. Connect the object to the negative terminal and the plating metal to the positive terminal of a low-voltage d.c. power supply (around 2 to 6 volts).

Turn on the power supply and leave a constant current running, gently rotating the object to ensure an even coating. Finally, remove the plated object and wash it with distilled water before rinsing it again with propanone. Because propanone is a highly volatile solvent, it evaporates extremely quickly, leaving the object completely dry so you can measure its final mass accurately.

The Chemistry of Plating: Migration and Redox

While the battery runs, an invisible exchange of atoms and ions happens beneath the surface of the solution. Positively charged metal ions, known as cations, exist freely in the electrolyte and are physically attracted to the negatively charged cathode.

When these cations reach the cathode, they gain electrons to become neutral solid metal atoms. This process is called reduction, and it is what creates the thin metal coating on your object. At the exact same time, neutral metal atoms in the pure anode lose electrons to become positive ions in a process called oxidation.

These newly formed ions dissolve into the electrolyte, replacing the ones that were just lost at the cathode. Because the rate of oxidation at the anode matches the rate of reduction at the cathode, the concentration of metal ions in the electrolyte remains perfectly constant. You will observe the anode decreasing in mass while the cathode increases in mass.

Half-Equations for Electroplating (Higher Tier)

Chemists represent this precise exchange of electrons using ionic half-equations. If we are silver plating an object, the pure silver anode dissolves while solid silver forms on the cathode.

At the anode (oxidation happens here as silver atoms lose electrons):

$$Ag(s) \rightarrow Ag^+(aq) + e^-$$

At the cathode (reduction happens here as silver ions gain electrons):

$$Ag^+(aq) + e^- \rightarrow Ag(s)$$

This same pattern applies to any generic metal, $M$, with an ion charge of $n+$. The anode will always show $M(s) \rightarrow M^{n+}(aq) + ne^-$, and the cathode will always show $M^{n+}(aq) + ne^- \rightarrow M(s)$.