# 4.1 The electrolytic cell

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## The electrolytic cell

In "The Galvanic Cell" , we saw that a chemical reaction that involves a transfer of electrons, can be used to produce an electric current. In this section, we are going to see whether the 'reverse' process applies. In other words, is it possible to use an electric current to force a particular chemical reaction to occur, which would otherwise not take place? The answer is 'yes', and the type of cell that is used to do this, is called an electrolytic cell .

Electrolytic cell

An electrolytic cell is a type of cell that uses electricity to drive a non-spontaneous reaction.

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

Electrolysis

In chemistry and manufacturing, electrolysis is a method of separating bonded elements and compounds by passing an electric current through them.

## Demonstration : the movement of coloured ions

A piece of filter paper is soaked in an ammonia-ammonium chloride solution and placed on a microscope slide. The filter paper is then connected to a supply of electric current using crocodile clips and connecting wire as shown in the diagram below. A line of copper chromate solution is placed in the centre of the filter paper. The colour of this solution is initially green-brown.

The current is then switched on and allowed to run for about 20 minutes. After this time, 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 chromate.

Explanation:

• 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 (CrO ${}_{4}^{2-}$ ) ions in the copper chromate solution are attracted to the positive electrode, while the Cu ${}^{2+}$ ions are attracted to the negative electrode.

Conclusion:

The movement of ions occurs because the electric current in the outer circuit sets up a potential difference between the two electrodes.

Similar principles apply in the electrolytic cell, where substances that are made of ions can be broken down into simpler substances through electrolysis.

## The electrolysis of copper sulphate

There are a number of examples of electrolysis. The electrolysis of copper sulphate is just one.

## Demonstration : the electrolysis of copper sulphate

Two copper electrodes are placed in a solution of blue copper sulphate and are connected to a source of electrical current as shown in the diagram below. The current is turned on and the reaction is left for a period of time.

Observations:

• The initial blue colour of the solution remains unchanged.
• It appears that copper has been deposited on one of the electrodes but dissolved from the other.

Explanation:

• At the negative cathode, positively charged Cu ${}^{2+}$ ions are attracted to the negatively charged electrode. These ions gain electrons and are reduced to form copper metal, which is deposited on the electrode. The half-reaction that takes place is as follows: $C{u}^{2+}\left(aq\right)+2{e}^{-}\to Cu\left(s\right)$ (reduction half reaction)
• At the positive anode, copper metal is oxidised to form Cu ${}^{2+}$ ions. This is why it appears that some of the copper has dissolved from the electrode. The half-reaction that takes place is as follows: $Cu\left(s\right)\to C{u}^{2+}\left(aq\right)+2{e}^{-}$ (oxidation half reaction)
• 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 Cu ${}^{2+}$ ions in the solution therefore remains almost the same and the blue colour of the solution is unchanged.

Conclusion:

In this demonstration, an electric current was used to split CuSO ${}_{4}$ into its component ions, $\mathrm{Cu}{}^{2+}$ and SO ${}_{4}^{2-}$ . This process is called electrolysis .

## The electrolysis of water

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

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

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 (normally platinum metal), submerged in an electrolyte and connected to a source of electric current.

The reduction half-reaction that takes place at the cathode is as follows:

$2{H}_{2}O\left(l\right)+2{e}^{-}\to {H}_{2}\left(g\right)+2O{H}^{-}\left(aq\right)$

The oxidation half-reaction that takes place at the anode is as follows:

$2{H}_{2}O\left(l\right)\to {O}_{2}\left(g\right)+4{H}^{+}\left(aq\right)+4{e}^{-}$

## A comparison of galvanic and electrolytic cells

It should be much clearer now that there are a number of differences between a galvanic and an electrolytic cell. Some of these differences have been summarised in [link] .

 Item Galvanic cell Electrolytic cell Metals used for electrode Two metals with different reaction potentials are used as electrodes The same metal can be used for both the cathode and the anode Charge of the anode negative positive Charge of the cathode positive negative The electrolyte solution/s The electrolyte solutions are kept separate from one another, and are connected only by a salt bridge The cathode and anode are in the same electrolyte Energy changes Chemical potential energy from chemical reactions is converted to electrical energy An external supply of electrical energy causes a chemical reaction to occur Applications Run batteries, electroplating Electrolysis e.g. of water, NaCl

## Electrolyis

1. An electrolytic cell consists of two electrodes in a silver chloride (AgCl) solution, connected to a source of current. A current is passed through the solution and Ag ${}^{+}$ ions are reduced to a silver metal deposit on one of the electrodes.
1. Give the equation for the reduction half-reaction.
2. Give the equation for the oxidation half-reacion.
2. Electrolysis takes place in a solution of molten lead bromide (PbBr) to produce lead atoms.
1. Draw a simple diagram of the electrolytic cell.
2. Give equations for the half-reactions that take place at the anode and cathode, and include these in the diagram.
3. On your diagram, show the direction in which current flows.

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