19.2 – Electrolysis
19.2.1 – Predict and explain the products of electrolysis of aqueous solutions
Electrolysis takes place in electrolytic cells, which means that the redox reaction takes place in reverse. Electricity is supplied to the cell to overcome the potential difference between the electrodes. The electrolyte contains ions that compete at the electrodes with the H+ and OH– ions, which may be oxidised or reduced instead of the dissolved salts.
Using the table of standard electrode potentials, we are able to predict whether a species can be oxidised or reduced in an aqueous solution.
At the cathode, only a species with less reducing power than H+ with be released into aqueous solution. Otherwise, the reaction is:
At the anode, only a species with more oxidising power than O2 will be released. Otherwise the reaction is:
Note that when the solution is highly concentrated, these rules may be overcome, causing the other species to be released instead.
There are three factors that will determine which ions are released at the electrodes:
Position of the Ion on the Electrochemical Series
The electrochemical series has the redox equilibrium expressed like in the example below:
In this form, they are expressed as reductions going from left to right. However, if this is simply reversed to show it from right to left, it would show the oxidation of the species.
In the form Cu2+ | Cu, the electrode potential is +0.34V, indicating that Cu2+ ions are more easily reduced than H+ ions. This means that if Cu2+ ions are competing with H+ ions, the Cu2+ ions are more able to accept electrons. We can hence see that species that are below hydrogen on the electrochemical series are better oxidisers, and will be preferentially deposited on the electrode.
Concentration of the Ion in the Solution
Concentration also has an effect when the ions have similar reactivity. The species with the higher concentration is more likely to be reduced.
The Nature of the Electrode
The electrode of an electrolytic cell is normally inert so that it will not affect the reaction. Their purpose is to complete the circuit and allow for the flow of electrons.
On the other hand, a metal anode may dissociate into its ions and participate in the reaction.
Example 1 – Electrolysis of Water
Water has a low ionisation and is not a good conductor, but its conductivity can be increased by adding ionic compounds. This allows electrolysis to take place.
For example, NaOH may be added to provide additional ions. Under these circumstances, the ions present in the solution are H+, OH– and Na+.
At the cathode, the H+ and Na+ ions compete. The H+ ions are reduced, forming H2 gas. The pH increases because H+ is discharged.
At the anode, the OH– ions accumulate and are oxidised, releasing electrons. There is no competition. As a result of this, O2 gas is released. The pH decreases because OH– is discharged.
Example 2 – Electrolysis of Sodium Chloride
The ions present in this solution are H+, OH–,Na+ and Cl–.
At the cathode, the sodium and hydrogen ions are competing. However, since H+ occurs lower down on the table of standard electrode potentials, it will be reduced. Bubbles of hydrogen gas will be produced.
On the other hand, at the anode, the hydroxide and chloride ions are competing. Since chlorine occurs lower down on the table, it will be oxidised, producing Cl2.
The remaining ions in the solution are Na+ and OH–, making sodium hydroxide.
Example 3 – Electrolysis of Copper (II) Sulfate
The ions present in the solution are Cu2+, SO42-, H+ and OH–.
At the cathode, the H+ and Cu2+ ions are competing. Since hydrogen is higher on the table, the copper will be reduced. This can also be determined by the fact that the Eo is positive.
The reaction of copper is:
At the anode, there is competition between the hydroxide ions and the sulfate ions, and the OH– ions are oxidised to release water and oxygen gas:
The remaining ions – H+ and SO42- – form sulfuric acid.
If a metal is more reactive than hydrogen, it will not be deposited during electrolysis. Instead, hydrogen gas is liberated at the cathode.
Halide ions are usually released. If these are not present, then the OH– ions will be released at the anode instead.
Unreactive metals can participate in the reaction at the anode.
19.2.2 – Determine the relative amounts of the products formed during electrolysis
Faraday’s law states that the mass produced is relative to the amount of charge passed in electrolysis.
One Faraday = 96500 Coulombs
Using the equation above, the number of Faradays used in the electrolysis is calculated. We then refer to the equation for the reaction to see how many electrons are required to produce on mole of product. We divide the number of Faradays used in the reaction by the number of electrons required to give the number of moles of product formed in the reaction. This process can be altered to determine the value of other variables if other factors are known.
Charge on the Ion
The charge on the ion being deposited will determine how many electrons it must pick up to become the solid metal. One Faraday of charge is required to release one mole of a single-charged ion. Likewise, if the charge on the ion is two or three, then two or three Faradays are required to release one mole.
Current
If more electrons are available in the circuit, then more ions can be deposited. The current tells how many electrons are being passed through the circuit.
Duration of Electrolysis
This can be used in combination with the current to calculate the total charge, which in turn shows how many moles of electrons have passed through the circuit..
19.2.3 – Describe the use of electrolysis in electroplating
Electroplating is the process of depositing a metal layer on another metal or conductive substance in an electrolytic cell. For this to occur, the electrolyte must contain the metal ions that are being deposited. The object that is being plated forms the cathode. The anode may also be made up of the metal that is coating the cathode to supply additional ions.
In this example, two copper electrodes are used. Remember that in electrolytic cells, the electrons flow towards the cathode and the copper ions will be deposited there in the form of a solid layer of copper.
As the electrolysis is allowed to continue, the anode will become smaller, and there will be a build-up of copper on the cathode. The mass of the cathode will increase as a result.
Electroplating is very common as it can be used for decoration, preventing corrosion or improving the function of certain items.
During the process, any neutral metal may be deposited on the cathode as the metal ions are taken out of solution in a reduction reaction. Unlike in the example above, both electrodes do not necessarily have to be the same substance. All that is important is that the cathode is the item to be plated, and the anode is made up of the metal to be plated over it. So, if you want to plate a spoon in silver, the spoon would be the cathode and the anode would be made of silver.
The process is also very effective because the metal always deposits in its pure form