Lesson Explainer: Uses of Electrolytic Cells Chemistry • Third Year of Secondary School

In this explainer, we will learn how to describe conditions and applications for the electrolysis of molten salts and salt solutions.

Specifically, we will look at the role of electrolytic cells in the purification of copper and in electroplating.

An electrolytic cell is a type of electrochemical cell where an external power source drives a reaction. We can use electricity to break apart substances, turning them into pure elemental substances or other chemicals.

To sustain an electrolytic reaction, we need to have a complete circuit; we need to be able to continuously get power from the battery or power supply. To do this, we need our ions to be able to move.

In electrolytic cells, we use a direct current power source, meaning that the electrodes are always either positive or negative.

A substance or mixture that conducts electricity and can undergo electrolysis is called an electrolyte.

So, our electrolyte must be either a salt solution or a molten salt.

The names anode and cathode refer to either the electrode that electron flow comes from (anode) or goes to (cathode).

Anions in the electrolyte travel toward the anode and are oxidized. Cations in the electrolyte travel toward the cathode and are reduced.

This diagram shows the movement of cations toward the cathode and anions toward the anode.

Once cations reach the cathode, they can be reduced by gaining electrons.

Once anions reach the anode, they can be oxidized by losing electrons.

One of the most important uses of electrolysis is the production of pure copper. Copper is used for making circuits, wires, pipes, and cooling units because of its high electrical and thermal conductivity. It is also malleable and ductile.

There are many ways of extracting copper from its ores; however, we will focus on the final purification step. This purification step takes place after copper ore has been extracted from the ground and then converted into impure copper plates and copper(II) sulfate solution:

The impure copper plates produced from this process are not sufficiently pure for use in electronics and other critical applications. Purer copper plates can be produced by electrolysis.

The impure copper plates are placed in a bath of copper(II) sulfate solution and connected to the positive terminal of an external power supply. The negative end of the power supply is connected to a pure copper plate sitting in the same copper sulfate bath. We will not look at how this pure copper plate is originally made.

The key thing is that the power supply drives two reactions, one at the anode and one at the cathode. Combined, these two reactions result in copper leaving the anode and plating the cathode. The impurities are left behind.

This is the final result:

• The impure copper anode loses mass as copper dissolves in the electrolyte and impurities fall away.
• The pure copper cathode gains mass as copper from the electrolyte is plated onto it.
• Soluble impurities stay in the electrolyte.
• Insoluble impurities stay in the anode or fall to the bottom of the reaction vessel.

We can look at this in more detail. In this process, the anode is not inert—it is reactive. The solid copper in the impure copper anode undergoes the following oxidation:

Some impurities will be oxidized as well (e.g., zinc and iron) and pass into the electrolyte:

Less-reactive impurities, like silver and gold, will not be oxidized at all and will fall to the bottom of the cell.

At the cathode, the reverse happens:

Copper ions in the electrolyte are reduced to copper metal, plating the cathode.

While this process does not produce pure copper, it produces copper that is substantially more pure than what we start with.

This process of electroplating is not exclusive to copper, however.

Example 1: Identifying the Products of the Electrolysis of Copper(II) Sulfate Using Graphite or Copper Electrodes

A student conducts two electrolysis experiments using copper(II) sulfate as an electrolyte. He uses graphite electrodes in the first experiment and copper electrodes in the second experiment. Which line in the table correctly identifies what happens at the electrodes in the two experiments?

	Experiment 1: Graphite Electrodes		Experiment 2: Copper Electrodes
	Anode	Cathode	Anode	Cathode
Line 1	Oxygen is produced.	Hydrogen is formed.	The anode dissolved.	Copper is formed.
Line 2	The anode dissolved.	Copper is formed.	Oxygen is produced.	Copper is formed.
Line 3	Oxygen is produced.	Copper is formed.	The anode dissolved.	Copper is formed.
Line 4	The anode dissolved.	Copper is formed.	The anode dissolved.	Copper is formed.
Line 5	Copper is formed.	Oxygen is produced.	Copper is formed.	The anode dissolved.

Answer

These are the scenarios we have been given.

In experiment 1, we have inert electrodes; graphite will not react with copper sulfate solution under these conditions. So, what we expect is for the copper sulfate solution to be electrolyzed; copper is less reactive than hydrogen, so we expect copper ions in solution to be reduced at the graphite cathode, turning into copper metal.

At the anode, since the sulfate anion is not a halide, we expect oxygen to be produced via the oxidation of hydroxide ions from the water.

In experiment 2, we have a reactive anode. The external power supply will draw electrons away from the anode, turning copper metal into copper ions, which go into solution.

On the cathode side, copper ions in solution will be reduced to form solid copper metal. The net result is that the copper anode will dissolve into solution and the copper cathode will be coated with more solid copper.

Therefore, in experiment 1, we expect oxygen to be formed at the anode and copper to be formed at the cathode. In experiment 2, we expect the anode to dissolve into solution and the cathode to be plated with copper.

The answer is line 3.

Some extremely useful materials are made by plating one metal onto a part (like a hubcap). Plating can be done by dipping a part in molten metal, by sputtering (bombarding a metal target with ions from a plasma so that bits fall onto the part), or by evaporating metal and letting it condense on the part. While all these are useful in their way, one of the most common methods is electroplating.

The process of electroplating is simple:

1. Immerse the part you want to electroplate in an electrolyte that contains the ions of the metal that you want to electroplate onto the cathode. For example, if you want to copper-plate something, you might use a solution of copper(II) sulfate as the electrolyte.
2. Connect the part to the negative terminal of a power supply. During the electrolysis, the part that will be electroplated will act as the cathode. The surface that is to be electroplated must be electrically conductive.
3. Introduce a piece of the metal that you want to electroplate with, for example, copper. This should be immersed so that a lot of surface area is in contact with the electrolyte.
4. Connect this piece of metal to the positive terminal of the power supply where it will act as the anode during the electrolysis.
5. Turn on the power supply.

If the metal is compatible with the part, the part will be plated with the metal the anode is made of. The longer the electroplating goes on, the thicker the coating will get.

There is one important consideration for electroplating: the metal that is being plated onto the part will generally be more reactive than the metals the part is made of. If we try to coat copper with silver by immersing copper in silver nitrate solution, the reaction will happen spontaneously (copper will react with silver ions in solution without any help from a power supply). There will still be a coating of silver on the copper that is left, but it will be very difficult to control the quality and thickness of such a coating.

If the reaction between the metal and the solution is slow enough, we can still do electroplating. This usually happens when the metals have very similar reactivities.

Here is a diagram of the basics of an electroplating setup.

If, for example, we wanted to coat a bell made of copper with zinc, we could use electroplating. Copper is less reactive than zinc, so it should work just fine.

If the power supply is set up properly, two things will happen when it is switched on.

At the anode, the zinc metal will be oxidized to form zinc ions, which will go into solution. The zinc anode will lose mass as the electroplating goes on:

At the cathode, the copper bell will be coated with zinc as zinc ions in solution are reduced to form zinc metal:

Example 2: Identifying the Setup Appropriate for Electroplating Iron with Copper

Which diagram shows the appropriate apparatus to electroplate the iron crown with copper?

Answer

In the four setups, we can see two blue solutions and two green solutions. The blue solutions contain copper(II) sulfate, and the green solutions contain iron(II) sulfate.

If we want to electroplate a metal onto the crown, ions of that metal must come from the solution. If we want to electroplate the iron crown with copper, we need a solution of copper ions. Therefore, we can dismiss any setup that uses iron(II) sulfate as the electrolyte.

In the setups with copper(II) sulfate solution as the electrolyte, the iron crown is connected to the negative terminal of the power source (making it the cathode). In the other, the iron crown is the anode.

When electroplating, positively charged metal ions from the solution will gain electrons at the cathode, turning into solid metal atoms. If we want to electroplate the crown, it must be the cathode.

Therefore, the right setup is answer C: the one with a copper anode (providing fresh copper ions to the electrolyte), an electrolyte of copper(II) sulfate solution, and the iron crown as the cathode.

We can then ask ourselves the following question: why would we want to coat one metal onto another?

A surface coating of a particular metal or mixture of metals can improve corrosion resistance, appearance, economic value, and mechanical function and can alter chemical and physical properties. Consider the following examples:

• Gold-coated silver has the same appearance as solid gold, but it is sold at a fraction of the cost. Jewelry can also be plated with platinum and other noble metals.
• Zinc-coated iron does not rust, and the layer of zinc oxide on the surface does not peel off like rust does. Instead, this layer forms a long-term barrier to corrosion.
• Aluminum can be plated onto metal parts in a thick layer that can be dyed to give brilliant colors. Parts treated in this way are said to have an anodized aluminum finish.
• Electrical contacts can be gold-plated to improve electrical conductivity and corrosion resistance.