Electrolytic Cells

Electrolytic cells

In an electrolytic cell electrical potential energy is converted to chemical potential energy. An electrolytic cell uses an electric current to force a particular chemical reaction to occur, which would otherwise not take place.

Definition: Electrolytic cell

An electrolytic cell is an electrochemical cell that converts electrical potential energy to chemical potential energy by using electricity to drive a non-spontaneous chemical reaction.

Fact:

Sometimes galvanic cells are just called electrochemical cells. While they are electrochemical cells, electrolytic cells are also electrochemical cells. Electrolytic and galvanic cells are not the same however.

An electrolytic cell is activated by applying an electrical potential across the electrodes to force an internal chemical reaction between the electrodes and the ions that are in the electrolyte solution. This process is called electrolysis.

Definition: Electrolysis

Electrolysis is a method of driving chemical reactions by passing an electric current through an electrolyte.

A sketch of an electrolytic cell.

In an electrolytic cell (for example the cell shown in the figure above):

The electrolyte solution consists of the metal cations and spectator anions.
The oxidation and reduction reactions occur in the same container but are non-spontaneous.They require the electrodes to be connected to an external power source to proceed.
The electrodes in an electrolytic cell can be the same metal or different metals. The prinicple is the same. Let there be only one metal, and let that metal be Z.
An electrode is connected to the $\color{blue}{\textbf{positive terminal}}$ of the battery.
- To balance the charge at the $\color{blue}{\textbf{positive electrode}}$ metal atoms are $\color{blue}{\textbf{oxidised}}$ to form metal ions. The ions move into solution, leaving their electrons on the electrode.
- The following reaction takes place: $\color{blue}{\textbf{Z(s)} \to {\textbf{Z}}^{+}{\textbf{(aq)}} + {\textbf{e}}^{-}}$
- $\color{blue}{\textbf{Ox}}$idation is loss at the $\color{blue}{\textbf{an}}$ode, therefore this electrode is the $\color{blue}{\textbf{anode}}$.
An electrode is connected to the $\color{red}{\textbf{negative terminal}}$ of the battery.
- When positive ions come in contact with the $\color{red}{\textbf{negative electrode}}$ the ions gain electrons and are $\color{red}{\textbf{reduced}}$.
- The following reaction takes place: $\color{red}{\textbf{Z}^{+}{\textbf{(aq)}} + {\textbf{e}}^{-} \to {\textbf{Z(s)}}}$
- $\color{red}{\textbf{Red}}$uction is gain at the $\color{red}{\textbf{cat}}$hode, therefore this electrode is the $\color{red}{\textbf{cathode}}$.
This means that the overall reaction is: $\text{Z}(\text{s}) + \text{Z}^{+}(\text{aq})$ $\to$ $\text{Z}^{+}(\text{aq}) + \text{Z}(\text{s})$. While this might seem trivial this is an important technique to purify metals (see Section 13.7).

In the movement of coloured ions experiment, to make the ammonia and ammonium chloride buffer solution you need to combine equimolar amounts of ammonia and ammonium chloride. The volume doesn't matter, so long as there are the same number of moles of each compound in the solution.

Concentrated, strong bases can cause serious burns. Please remind the learners to be careful and wear the appropriate safety equipment when handling all chemicals, especially strong, concentrated bases. The safety equipment includes gloves, safety glasses, and protective clothing.

Optional Experiment: The movement of coloured ions under the effect of electrical charge

Aim

To demonstrate how ions migrate in solution towards oppositely charged electrodes.

Apparatus

Filter paper, glass slide, a 9V battery, two crocodile clips connected to wires, tape
Ammonia ($\text{NH}_{3}(\text{aq})$) and ammonium chloride ($\text{NH}_{4}\text{Cl}$) buffer solution, copper(II) chromate solution

Method

Connect the wire from one crocodile clip to one end of the battery and secure with tape. Repeat with the other crocodile clip wire and the other end of the battery.
Soak a piece of filter paper in the ammonia and ammonium chloride buffer solution and place it on the glass slide.
Connect the filter paper to the battery using one of the crocodile clips, keep the other one nearby.
Place a line of copper(II) chromate solution at the centre of the filter paper. The colour of this solution is initially green-brown.
Attach the other crocodile clip opposite the first one (as shown in the diagram) and leave the experiment to run for about $\text{20}$ $\text{minutes}$.

Results

After $\text{20}$ $\text{minutes}$ you should see that the central coloured band disappears and is replaced by two bands, one yellow and the other blue, which seem to have separated out from the first band of copper(II) chromate.
The cell that is used to supply an electric current sets up a potential difference across the circuit, so that one of the electrodes is positive and the other is negative.
The chromate ($\text{CrO}_{4}^{2-}$) ions in the copper(II) chromate solution are attracted to the positive electrode, this creates a yellow band. The $\text{Cu}^{2+}$ ions are attracted to the negative electrode, this creates a blue band.

Conclusion

The movement of ions occurs because the electric current in the external circuit provides a potential difference between the two electrodes.

Optional Experiment: An electrolytic cell

Aim

To investigate the reactions that take place in an electrolytic cell.

Apparatus

Two copper plates (of equal size and mass), copper(II) sulfate ($\text{CuSO}_{4}$) solution ($\text{1}$ $\text{mol.dm$^{-3}$}$)
A $\text{9}$ $\text{V}$ battery, two connecting wires, a beaker.

Method

Half fill the beaker with copper(II) sulfate solution. What colour is the solution?
Weigh each copper electrode carefully and record the weight.
Place the two copper electrodes (of known mass) in the solution and make sure they are not touching each other.
Connect the electrodes to the battery as shown below and leave the experiment for a day. What colour is the solution after a day?

Discussion

What colour was the copper(II) sulfate solution before the experiment?
What colour was the copper(II) sulfate solution after the experiment?
Examine the two electrodes, what do you observe?
What is the charge on each electrode?
Which electrode is the anode and which is the cathode?

Observations

The initial blue colour of the solution remains unchanged throughout the experiment.
It appears that copper has been deposited on one of the electrodes (it increased in mass) but dissolved from the other (it decreased in mass).
The electrode connected to the negative terminal of the battery will have a negative charge. The electrode connected to the positive terminal of the battery will have a positive charge.
When positively charged $\text{Cu}^{2+}$ ions encounter the negatively charged electrode they gain electrons and are reduced to form copper metal. This metal is deposited on the electrode. The half-reaction that takes place is as follows:

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

Reduction occurs at the cathode. Therefore, the electrode which increased in mass is the cathode.
At the positive electrode, copper metal is oxidised to form $\text{Cu}^{2+}$ ions, leaving electrons on the electrode. The half-reaction that takes place is as follows:

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

Oxidation occurs at the anode. Therefore, the electrode which decreased in mass is the anode.
The amount of copper that is deposited at one electrode is approximately the same as the amount of copper that is dissolved from the other. The number of $\text{Cu}^{2+}$ ions in the solution therefore remains almost the same, and the blue colour of the solution is unchanged.

Conclusion

In this demonstration, the container held aqueous $\text{CuSO}_{4}$ ($\text{Cu}^{2+}(\text{aq})$ and $\text{SO}_{4}^{2-}(\text{aq})$). The copper atoms of the electrode connected to the positive terminal (the anode) were oxidised and formed $\text{Cu}^{2+}(\text{aq})$ ions, causing a decrease in mass. The copper atoms of the electrode connected to the negative terminal (the cathode) were reduced to form solid copper, causing an increase in mass. This process is called electrolysis, and is very useful in the purification of metals.

Note that the cathode is negative and the anode is positive. Reduction still occurs at the cathode (Red Cat), and oxidation still occurs at the anode (An Ox).

Tip:

The electrolysis of water

Fact:

August Wilhelm von Hofmann was a German chemist who invented the Hofmann cell, which uses a current to form $\text{H}_{2}(\text{g})$ and $\text{O}_{2}(\text{g})$ from water through electrolysis.

Water can undergo electrolysis to form hydrogen gas and oxygen gas according to the following reaction:

$2\text{H}_{2}\text{O}(\text{l})$ $\to$ $2\text{H}_{2}(\text{g}) + \text{O}_{2}(\text{g})$

Hofmann apparatus for the electrolysis of water.

This reaction is very important because hydrogen gas has the potential to be used as an energy source. The electrolytic cell for this reaction consists of two electrodes, submerged in an electrolyte and connected to a source of electric current (see figure above).

The $\color{blue}{\textbf{oxidation half-reaction}}$ is as follows: The $\color{red}{\textbf{reduction half-reaction}}$ is as follows:

$\color{blue}{\text{2H}_{2}\text{O(l)} \to \text{O}_{2}\text{(g) + 4H}^{+}\text{(aq) + 4e}^{-}} \color{red}{\text{ 2H}^{+}\text{(aq) + 2e}^{-} \to \text{H}_{2}\text{(g)}}$

Electrolytic Cells