Electrolysis

Electrolysis is the very first electrochemical technique ever used. The word is derived from the Greek words "electron" and "lysis" the first in its present day meaning refering to electricity, the second to decomposition, so: decomposition by the means of an electricity.

Example
a Daniell cell consists of a zinc anode (an electron collector), is oxidized as it dissolves into a zinc sulfate solution, the dissolving zinc leaving behind its electrons in the electrode according to the oxidation reaction (s = solid electrode; aq = aqueous solution):


 * $$\mathrm{Zn_{(s)} \rightarrow Zn^{2+}_{(aq)} + 2 e ^- \ } $$

The zinc sulfate is the electrolyte in that half cell. It is a solution which contains zinc cations $$\mathrm{Zn}_{} ^{2+}$$, and sulfate anions $$\mathrm{SO}_4^{2-}\ $$ with charges that balance to zero.

In the other half cell, the copper cations in a copper sulfate electrolyte are drawn to the copper cathode to which they attach themselves as they adopt electrons from the copper electrode by the reduction reaction:


 * $$ \mathrm{Cu^{2+}_{(aq)} + 2 e^- \rightarrow Cu_{(s)}\ } $$

Prerequisites
A basic understanding of the following topics is assumed:
 * Atom
 * Reduction and oxidation or redox reaction
 * Electrical potential

Short History
Electrolysis dates back to the turn of the 18th into the 19th century. At that time Sir Humphry Davy applied the then newly discovered voltaic pile to molten potassium hydroxide, releasing the metal out of its compound. Although the source of the electrical current has changed over time, principals have not.

Principals
At its base during electrolysis a redox reaction is performed. The difference between the plain redox reaction between two chemical compounds and electrolysis is found in the route an electron uses to move from one atom to another. In ordinairy redox reactions the electron is directly transferred between the atoms. In electrolysis the electrons travel along the electrical circuitry.

Techniques
The basic apparatus used in electrolysis is shown on the right. It conists of a voltage source, leads, two electrodes and a container with the substance to be electrolysed, in this case: a beaker glass with water. The beaker glass with its content is referred to as: electrochemical cel. The part of the leads immerged in the solution are called "electrodes". In the electrolysis of water these electrodes are made of platinum. The voltage source in general is adjusted to a few volts.

The names of the electrodes are: The naming of the two electrodes does not depend on the use of the electrochemical cell.
 * Anode, the electrode at which the oxidation reaction takes place.
 * Cathode, the electrode at which the reduction reaction occurs.

Reactions
The reactions taking place are treated separately, as they are spatially separated too.

Anodic Or Oxidation Reaction
At the anode an oxidation reaction is taking place. The anode is made positive by its connection to the positive terminal of the voltage source. It has a great oxidizing power. Electrons do "want" to go there. As in this example water is the only component in the cell, it is to be oxidized according to: 2 H2O O2 + 4 H+ + 4 e&minus; The formed oxygen ascends as bubbles and might be caught in a test tube and tested for, otherwise it escapes into the air. The ions of hydrogen simply dissolve in the remaining water. The electrons are taken in by the anode.

The anode, however, is connected to the positive terminal of the voltage source. The drop in potential due to the arrival of the electrons on the anode leads to the transfer of these electrons to the negative terminal: the potential of the anode does not change.

Cathodic Or Reduction Reaction
At the cathode a reduction reaction is taking place. The cathode is very rich in electrons, as it is connected to the negative terminal of the voltage source. It has a great reducing power. As in this example water is the only component in the cell, it is to be reduced according to: 2 H2O + 2 e&minus; H2 + 2 OH&minus; The formed hydrogen ascends as bubbles and might be caught in a test tube and tested for, otherwise it escapes into the air. The ions of hydroxide simply dissolve in the remaining water and of course will react with the hydrogen ions formed at the anode. The electrons are taken in by the anode.

The anode, however, is connected to the positive terminal of the voltage source. The drop in potential due to the arrival of the electrons on the anode leads to the transfer of these electrons to the negative terminal: the potential of the anode does not change.

Complete Reaction
As is seen in the reaction equations, hydrogen and oxygen are formed. Although the equations are balanced in themselves, the over all reaction needs some adjustment, as four electrons are taken in at the anode, and only two are released at the cathode. To balance the over all reaction, the cathodic one should be read twice, so: 2 H2O 2 H2 + O2

More Complex Electrolysis
The example of water being electrolysed is straight forward. A more complicated situation arises when a solute is added. Now at both electrodes several reactions might be possible, but only one will proceed. As an example a solution of zinc iodide is used. To predict the reactions at each electrode that will occur the following procedure will help:

Determine the components present at the anode. The components and their respective redox potentials for oxydation are:

Determine the components present at the cathode. The components and their respective redox potentials for reduction are:

Reference

 * Electrochemical cell