L11 - Electrolytic cells

objectives

Electrolytic cells:

• Electrolytic cells convert electrical energy to chemical energy, by bringing about non-spontaneous processes.

• Oxidation occurs at the anode (positive electrode) and reduction occurs at the cathode (negative electrode) in an electrolytic cell.

Applications and skills:

• Construction and annotation of both types of electrochemical cells.

• A redox reaction is used to explain how current is conducted in an electrolytic cell.

• Distinction between electron and ion flow in both electrochemical cells.

• Performance of laboratory experiments involving a typical voltaic cell using two metal/metal-ion half-cells.

• Deduction of the products of the electrolysis of a molten salt.

An Introduction to Electrolysis

• An electrolytic cell is the apparatus used to pass electricity through an electrolyte (a conducting liquid, either a molten ionic compound, or an ionic solution).

• The electrolytic cell has two (usually inert) electrodes to pass the electric current into the electrolyte.

• The negative electrode is called the cathode and the positive electrode is called the anode.

• The inert electrodes are usually made of graphite or platinium wire. Chemical reactions occur at the surface of the electrodes.

• The power supply is usually symbolised by a short fat line (negative) and a long thin line (positive)

Reactions at the Electrodes

• The power pack (battery) supplies electrons for the negative electrode (the cathode), which the positive ions in the electrolyte can pick up to become other species, usually elements.

• At the same time the positive potential at the positive electrode can remove electrons from the ions in solution to make new species, also usually elements.

• The consequence of these two processes is that electrons are removed from negative ions at the the anode (by the powerful positive potential there) and they are added to positive ions at the cathode (where there are many available electrons).

• The overall effect is that the anode gains electrons as the cathode loses them.

• As far as the battery is concerned there is a regular flow of electrons out of the negative terminal and into the positive terminal.

• Let's look at the reactions occurring at both electrodes in more detail.

Reaction at the Anode

• The anode is the positive electrode, connected directly to the positive terminal of the battery.

• This is strongly attractive to negative ions.

• Once a negative ion arrives at the anode its electron is removed by the strong positive potential and the ion becomes a neutral species.

• Hence the process that occurs is oxidation.

For example:

Cl-(aq) --> Cl + 1e

The chlorine atom is too reactive to exist for long on its own and pairs up with another chlorine atom to form a molecule.

2Cl --> Cl2(g)

Usually both steps are written as one:

2Cl-(aq) --> Cl2(g) + 2e

Reactions at the Cathode

• The cathode is supplied with electrons by the negative terminal of the battery.

• This attracts positive ions from the electrolyte, which pick up the electrons neutralising their positive charge.

• Thus the ions become atoms, or molecules depending on the type of ion.

For example, a sodium ion in a molten sodium chloride electrolyte will approach the cathode and pick up an electron:

Na+ + 1e --> Na

This results in the formation of metallic sodium. It is the way that sodium is manufactured in industry.

Conduction in an Electrolytic Cell

• Current passes around the external circuit to and from the battery in the normal way i.e. by the movement of electrons. However, in the cell itself there is a very different process occurring.

• Positive ions from the electrolyte pick up electrons at the cathode and use them to perform reduction of the ion (reduction = addition of electrons).

• At the same time negative ions migrate to the positive electrode (anode) to drop off electrons and get oxidised (oxidation = loss of electrons).

• Overall, there are ions picking up electrons from one electrode (the cathode) and DIFFERENT IONS dropping off different electrons at the other electrode (anode).

• The net effect is as if electrons are jumping from one electrode to the other.

• It should be stressed that at no time do electrons cross the electrolyte.

• The battery, however, cannot distinguish between electrons and, to all intents and purposes, as a current passes around the external circuit, it seems also to pass through the electrolyte.

Electrolysis of Molten Sodium Chloride

• In molten sodium chloride, the ions are free to migrate to the electrodes of an electrolytic cell.

• A simplified diagram of the cell commercially used to produce sodium metal and chlorine gas is shown in Figure 1.

• Sodium is a strong reducing agent and chlorine is used to purify water, and is used in antiseptics and in paper production. The reactions are

• The power supply (battery) must supply a minimum of 4 V, but, in practice, the applied voltages are typically higher because of inefficiencies in the process itself.

Electrolysis of Lead Bromide

• Lead bromide is an ionic solid and as such contains charged particles (ions).

• While in the solid state these ions are not free to move, but once the structure is melted the ions become free to move.

• If an electrical potential is applied across some molten lead(II) bromide, the lead 2+ ions can move to the cathode (negative electrode) and the bromide (negative) ions move to the anode (positive electrode).

• Once at their respective electrodes, the ions can undergo reaction.

At the anode (positive electrode)

The negative bromide ions arrive and drop off electrons to become bromine atoms:

Br- - 1e --> Br

These single atoms then pair up to become bromine molecules, which, at the temperature of the molten lead(II) bromide, is a gas:

2Br --> Br2(g)

Overall the process could be represented as:

2Br- - 2e --> Br2(g)

At the cathode (negative electrode)

The lead 2+ ions move to the cathode where they pick up electrons:

Pb2+ + 2e --> Pb(l)

At the temperature of the molten lead(II) bromide the lead formed falls to the bottom of the melt as a silvery liquid.

The overall electrolysis

The overall process occurring in the electrolytic cell can be obtained by adding together the two equations for the processes going on at the two electrodes (ensuring that the same number of electrons appear in each equation.

2Br- - 2e --> Br2(g)

Pb2+ + 2e --> Pb(l)

PbBr2 --> Pb(l) + Br2(g)

As we can see, the net result is that the ionic compound lead(II) bromide has been broken apart into its original elements. The term, 'electrolysis' means literally broken apart (lysis) by electricity (electro).

Industrial Extraction of Aluminium

• Aluminium is a reactive metal that can only be effectively extracted from its ores using electrolysis.

• However, there are a couple of difficulties.

• Firstly the chloride of aluminium is covalent at elevated temperatures, meaning that it cannot be electrolysed.

• The oxide of aluminium is ionic, but it has an extremely high melting temperature making it unsuitable for direct use.

• These two problems are overcome by using sodium aluminium fluoride, NaAlF4 (cryolite), which melts at a reasonably accessible temperature for industry.

• Alumnium oxide can then be dissolved in the molten cryolite and the mix electrolysed.

• Aluminium oxide can be continuously added to the molten mixture as it is used up.

• The whole continuous process is called the Hall process, or Heroult - Hall process after its developers.

At the anode

Oxide ions migrate to the anode where they drop off electrons and become oxygen gas. At the elevated temperature of the electrolysis cell this oxygen gas reacts with the carbon electrodes, gradually burning them away. This is one of the pollution problems of the Hall process, it produces a lot of carbon dioxide (a greenhouse effect gas).

2O2- - 4e --> O2(g)

C(s) + O2(g) --> CO2(g)

At the cathode

Aluminium ions migrate to the cathode and pick up electrons to form molten aluminium metal that sinks to the bottom of the cell and which can be tapped off via an appropriately located tap.

Al3+ + 3e --> Al(l)

The overall cell reaction:

2O2- - 4e --> O2(g)

Al3+ + 3e --> Al(l)

2Al2O3 --> 4Al(l) + 3O2(g)