IB Chemistry home > Syllabus 2016 > Redox processes > Reaction at the electrodes

Syllabus ref: 9.2 Syllabus ref: 19.1

Electrolytic cells break compounds apart using electricity. Here we look specific examples and the reactions occurring at the electrodes.

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.

An analogy

As an analogy, one could consider the processes occurring in a bank. Certain clients deposit money and other clients withdraw money. To the bank it appears as though there is a constant flow of money in and out. What happens to this money outside of the bank is immaterial. In the electrolytic cell certain ions deposit electrons onto an electrode and other ions withdraw electrons from an electrode. The electrodes are attached to the battery that performs the accountancy. The battery 'sees' electrons leaving from one terminal (the negative) and arriving at the other terminal (the positive) - it detects the flow of current in and out.


Electrolysis of molten lead(II) 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).


Electrolysis of molten sodium chloride

Sodium chloride contains Na+ ions and Cl- ions

The positive sodium ions are attracted by electrostatic forces to the cathode (negative electrode). Once they get there they encounter many avilable electrons. Each sodium ion can accept one electron and be reduced to a sodium atom:

Na+ + 1e Na

The negative chloride ions are attracted to the positive electrode by elecrostatic forces and migrate there. Once there, the negative chloride ions encounter an anode seriously lacking in electrons with positive 'holes'. The chloride ions drop off their electrons into the positive 'holes' and become chlorine atoms:

2Cl- - 2e Cl2

(The chlorine atoms produced then pair up to become chlorine molecules)

The overall effect for the circuit is that for each Na+ and Cl- reaction at the electrodes, one electron has been removed from the cathode and another electron has been deposited on the anode.

The battery detects electrons flowing around the circuit as an electric current and the bulb lights. Once again the cell equation is obtained by adding together the reactions occurring at the electrodes (having ensured that the number of electrons is the same):

2Na+ + 2e 2Na
2Cl- - 2e Cl2
2Na+ + 2Cl- 2Na + Cl2

Electrolysis of molten sodium chloride


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)