R3.2.16 Electroplating (Higher Level Only)

Electroplating and Electrolytic Cells

1. Imagine holding a dull steel ring in your hand.
2. What if you could transform it into a shiny, corrosion-resistant piece of jewelry by coating it with copper?

This transformation is possible through electroplating, an intriguing application of electrolytic cells.

Understanding Electroplating

By passing an electric current through an electrolyte solution containing metal ions, we can deposit a layer of metal onto the object.

This setup allows copper ions from the solution to deposit onto the steel ring, while copper atoms from the anode replenish the ions in the solution.

The Role of Electrodes in Electroplating

In an electrolytic cell, the roles of the electrodes can be summarized as follows:

• Anode (Positive Electrode):
  • Oxidation occurs here, where metal atoms lose electrons and form ions that enter the solution.
• Cathode (Negative Electrode):
  • Reduction occurs here, where metal ions from the solution gain electrons and deposit as solid metal onto the object.

Writing Half-Equations for Electroplating

To understand the chemical changes, we write half-equations for the reactions at each electrode. Let’s use the example of copper electroplating:

At the Cathode (Reduction):

The steel ring acts as the cathode, where copper ions ${\text{Cu}}^{2+}$ in the solution are reduced to solid copper $\text{Cu}(s)$:

$${\text{Cu}}^{2+}(\text{aq})+2e^{-}\to \text{Cu}(s)$$

At the Anode (Oxidation):

The copper anode is oxidized, releasing copper ions into the solution:

$$\text{Cu}(s)\to {\text{Cu}}^{2+}(\text{aq})+2e^{-}$$

The Overall Reaction

When the two half-equations are combined, the electrons cancel out, resulting in the overall reaction:

$$\text{Cu}(s)\text{(anode)}\to \text{Cu}(s)\text{(cathode)}$$

Common Mistakes in Electroplating Calculations

Reflection