Electrolytic Cells: Electrolysis

Electroplating

Electroplating allows the production of metal coatings of such desirable commodities as silver and gold. People make fortunes gold or silver plating junk metal (usually aluminum) because they can sell gold plated necklaces for a comparable price to the real thing (or even pass them off as being solid gold). That's how electrochemistry can be used to rip you off! In our discussion of electroplating, we will discuss how you can set up a cell for electroplating, how you can calculate the amount of precious material consumed, and various other calculations you can perform with electroplating. In terms of the variety of electrochemistry problems possible to ask, this section, perhaps rivaled by Thermodynamics, is the richest.

The setup for electroplating is quite simple and the entire cell is usually conducted in a single solution as shown in .

The gold from the anode is oxidized and dissolves in solution as Au³⁺. The electrons arriving at the aluminum glasses frame cathode reduce the Au³⁺ in solution to Au (s) on the surface of the frame cathode. We can calculate how long we should have our glasses frame in solution if we want a certain amount of gold to be plated.

Let's assume it takes 1.0 g of gold to provide an adequate coating for our glasses and also assume that we are using an emf sufficient to produce 10 amperes (A) of current (1 A = 1 coulomb per second). how long it will take to plate that 1.0 g of gold.

As you can see from the , such a problem only involves the use of unit cancellation. To calculate the time needed to deposit a certain amount of material, you need to start with the amount, converted to moles. Then, multiply by the number of electrons consumed in the reduction (in this case 3). Using the definition of a faraday, 96500 C per mole of electrons, you can convert between moles and charge. Finally, by using the definition of an ampere, 1 C per second, you can convert the amount of charge required to deposit the material into a time in seconds. There are various ways of phrasing this same problem such as "how much gold is deposited in 146 seconds at 10 A" or "what current is required to deposit 1.0 g of gold in 146 seconds." Don't be fooled by those permutations of the same problem, they all boil down to simple unit cancellation which you have been doing since you learned how to do stoichiometry. Also note that in these problems, you do not need to know the cell potential. Students often try, incorrectly, to use the cell potential somewhere in that calculation. Furthermore, you need only know the number of electrons transferred--you could solve the same problem without even knowing what material was being plated (as long as you know its molar mass).

Previous page Electrolysis page 1 Next section