ELECTROLYSIS

Contents for this page	Related topics
Introduction Electrolysis of water Extraction of aluminium Electrorefining of copper Electroplating Faraday's laws of electrolysis Confusion with electrodes Additional questions	The Daniell (Cu-Zn) cell Standard electrode potentials How to balance redox reactions	Data Glossary
Learning Outcomes
After studying this section, you will (a) be familiar with the processes which take place during electrolysis, and (b) know the use of electrolysis in the refining of certain metals.

Introduction:

Electrolysis takes place in an electrolytic cell, the simplest form of which is shown here on the right: The components which make contact with the electrolyte are called ELECTRODES. The electrode which is attached to the negative pole of the battery, and which supplies electrons to the electrolyte, is called the CATHODE. Reduction takes place at the cathode. The electrode which is attached to the positive pole of the battery, and which accepts electrons from the electrolyte, is called the ANODE. Oxidation takes place at the anode.

When a direct electric current is passed through an ELECTROLYTE (such as a molten salt or an aqueous solution of a salt, acid or base), chemical reactions take place at the contacts between the circuit and the solution. This process is called ELECTROLYSIS. The combination of container, electrodes and electrolyte constitute the ELECTROLYTIC CELL. In such a cell, electrical energy from the current supply (battery or DC generator) is converted to chemical energy.

Various reactions take place at the electrodes during electrolysis. In general, reduction takes place at the cathode, and oxidation takes place at the anode.

Electrolysis is a hugely important process - many elements were discovered by electrolysis of their salts (the alkali metals, for example), and it is used industrially in a number of ways:

• The extraction of metals, notably aluminium, magnesium, and sodium.
• The preparation of halogens, for example chlorine.
• The refining of metals, such as copper and zinc.

At the electrodes, ions become discharged, or the anode may become oxidized and pass into solution. INERT ELECTRODES are electrodes which do not undergo any change during electrolysis. Platinum and carbon are frequently used when inert electrodes are required.

Let us consider what happens when a fairly concentrated solution of copper(II) chloride (CuCl2) is electrolyzed in the cell shown below, where the electrodes are made of carbon: When the current is flowing, the following is seen to happen: Metallic copper is deposited on the cathode Chlorine gas is evolved at the anode.

At the cathode, electrons are supplied to cations, which migrate to the cathode. The Cu²⁺ cation is discharged by accepting electrons. At the anode, electrons are supplied to the anions, which migrate to the anode. The half reactions are

Generally speaking, prediction of the nature of the products is complicated by (i) concentration effects and (ii) the rate at which the oxidation and reduction of different ions takes place. Standard electrode potentials can give a rough guide as to what may happen.

As a guide, in fairly concentrated aqueous solution, metals with a positive (reduction potential) will be formed at the cathode. Otherwise, hydrogen is formed by reduction of water or of the H⁺ ion. Halogens (chlorine, bromine and iodine) are formed at the anode when aqueous solutions of halides are electrolyzed. The sulphate ion is not discharged (oxidized) at the anode in aqueous solution, rather, oxygen is formed by oxidation of water or of the OH^- ion. Thus, electrolysis of sulphates produce oxygen at the anode.

Electrolysis of water:

Water may be electrolyzed in the apparatus shown below. Pure water is however a very poor conductor of electricity, and one has to add dilute sulphuric acid in order to have a significant current flow.

The electrodes consist of platinum foil. The electrolyte is dilute sulphuric acid. Hydrogen gas is evolved at the cathode, and oxygen at the anode. The ratio, by volume, of hydrogen to oxygen, is exactly 2:1. Remember that electron flow in the circuit is opposite to the conventional current flow. Thus, while the conventional current flows from the positive pole through the electrolyte and ends up at the negative pole, electrons flow from the negative pole in the reverse direction. The positive pole of a battery accepta electrons from the electrolyte by means of the anode of the electrolytic cell. The reaction at the cathode (tube A) is the reduction of protons:

Oxidation takes place at the anode (tube B). There are two anions competing to give up their electrons, namely sulphate (SO₄^2-), and hydroxide (OH^-) from the ionization of water. The oxidation of OH^- according to the reaction

has a standard electrode potential of -0.40V, compared to the oxidation of sulphate (-0.17V), and consequently, OH^- will be oxidized preferentially. The overall reaction is therefore

and the electrons are reurned to the battery, thus completing the circuit.

Extraction of aluminium:

Aluminium is obtained by the electrolytic reduction of its molten oxide, alumina (Al₂O₃). Because alumina has a very high melting point (2045 ºC), the mineral cryolite (Na₃AlF₆) is added to lower the melting point in order that the electrolysis may be carried out at about 950 ºC. The electrolytic cell has carbon anodes and a carbon cathode (which forms the lining of the tank in which the electrolysis takes place). Carbon dioxide is formed at the anodes, and aluminium at the cathode. It is heavier than the molten alumina/cryolite mixture, and sinks to the bottom of the cell, where it is tapped off. The procedure is known as the Hall-Héroult process.

Aluminium extraction is very demanding on electrical current (typically, 3-5 V and 100 000 A), and is economical only where power is cheap.

Electrorefining of copper:

When copper is first obtained by reduction of its ores, it is cast as impure slabs or ingots, called blister copper. In the electrorefining process, the blister ingots are used as anodes in an electrolytic cell, where an acid solution of copper (II) sulphate is used as electrolyte. Initially, the cathodes consist of thin sheets of pure copper.

During electrolysis, copper passes into solution from the anodes, (leaving the impurities, normally containing silver, gold and platinum) as an anode slime, which sinks to the bottom of the cell. The anode reaction is

At the cathode, copper (II) ions are discharged and the pure copper sheet becomes coated with an increasingly thick layer of very pure copper:

Electroplating:

Electroplating consists of depositing a thin layer of a metal on another, either for protection or for the sake of appearance. Typically, a brass or nickel object is coated with a layer of silver by making use of electrolysis of a silver solution, using the object to be coated as the cathode:

The anode consist of pure silver, and the cathode is the object to be plated. The electrolyte is a mixure of silver nitrate with potassium cyanide. The reactions are: At the anode: Ag → Ag+ + e- At the cathode: Ag+ + e- → Ag The cyanide ensures a low concentration of silver ions, a condition for providing the best plating results.

During the process, the concentration of silver in the electrolyte remains constant, as the rate of reduction at the cathode (which is the rate of deposition of silver on the object) is the same as the rate of reduction at the anode (which is the rate of rate of dissolution of the silver anode).

Faraday's laws of electrolysis:

Michael Faraday, a pioneer in the properties of electric currents, formulated two basic laws of electrolysis:

FARADAY'S FIRST LAW may be stated as follows:

The amount of electricity passing through the circuit in a given time is the number of moles of electrons passing through the circuit in that time, and the charge Q is related to the current I by

The charge on the electron is 1.602 x 10^-19 C, and Avogadro's number is 6.023 x 10²³. It follows that one mole of electrons has a charge of 9.65 x 10⁴ C. This quantity is known as the FARADAY or FARADAY'S CONSTANT (F).

FARADAY'S SECOND LAW may be stated as follows:

This means that z moles of electrons are needed to discharge an ion X^z+ or X^z-.

In the apparatus below, 1 Faraday will discharge 9 g Al (1/3 mole), 20 g Ca (1/2 mole) and 23 g Na (1 mole). The relevant half reactions are: Al3+ + 3e- Al Ca2+ +2e- Ca Na+ + e- Na

Confusion with electrode names:

It is very easy to be confused about the names CATHODE and ANODE and what their properties are, both with electrochemical cells and electrolytic cells. This, hopefully, will help you!

For BOTH electrochemical cells AND electrolytic cells,

Cathode is the site of reduCtion, or, if you prefer, CCC = Cathode Collects Cations.
Anode is the site of oxidAtion, or, AAA = Anode Attracts Anions.

This makes the table below quite simple:

	Electrochemical cells	Electrolytic cells
CATHODE	Accepts electrons from the external circuit (wire).	Accepts electrons from the external circuit (wire).
	Is the site of REDUCTION.	Is the site of REDUCTION.
	Is the conventional POSITIVE pole of the cell.	Is attached to the NEGATIVE pole of the DC source.
ANODE	Releases electrons to the external circuit (wire).	Releases electrons to the external circuit (wire).
	Is the site of OXIDATION.	Is the site of OXIDATION.
	Is the conventional NEGATIVE pole of the cell.	Is attached to the POSITIVE pole of the DC source.

Additional questions