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Introduction to Electrolysis

Think of electrolysis and electrolytic cells as the opposite of electrochemical cells:

Electrochemical
Cells / Electrolytic
Cells
Energy conversion / Chemical → Electrical / Electrical → Chemical
Spontaneous chemical reaction? / Yes / No
Value of E° / Positive / Negative


Recall that in an electrochemical cell, a spontaneous redox reaction is used to create an electric current.

In an electrolytic cell the reverse will occur - an electric current is used to cause a non-spontaneous chemical reaction to occur.

An electrochemical cell spontaneously ‘pushes’ electrons from the anode (-) to the cathode (+). However, in an electrolytic cell, an external voltage or EMF is applied to make the electrons flow in the opposite direction in a single electrolyte solution or pure liquid.

Note that this application of voltage will reverse the sites of oxidation and reduction leading to the occurrence of non-spontaneous reactions. Electrons produced at the negative terminal of a battery flow to the cathode where reduction takes place. Electrons are released at the anode which flow to the positive terminal of a battery. This means, in an electrolytic cell:

·  The cathode is negative (-)

·  The anode is positive (+)

Uses of electrolysis

Electrolysis is used in many industrial processes. Some of these applications are:

·  Extraction of Aluminium from its ore (Bauxite ore à Alumina (Al2O3) à Al)

·  Copper refining (blister copper à pure copper metal)

·  Electrolysis of molten NaCl to produce Na metal and Cl2 gas

·  Electrolysis of water to produce H2 and O2 gases

·  Electroplating of a thin layer of one metal on to another.

Fig.A: The diagram above shows the setup for the electrolysis of molten sodium chloride. The cut-away diagram below shows an industrial setup. Fig.B: Chlorine gas rises from the molten salt and molten sodium metal is less dense than water and floats up to a holding tank as it is produced at the cathode.


Electrolysis of Molten Sodium Chloride

The following equation represents the breaking apart of NaCl(l):

2NaCl(l) → 2Na(l) + Cl2 (g)

The half-reactions involved in this process are:

E°
reduction / 2Na+(l) + 2e- → Na(s) / -2.71 V
oxidation / Cl-(l) → Cl2 (g) + 2 e- / -1.36V
net voltage required / - 4.07V

The negative voltage (-4.07V) that results when we add up the half-reactions indicates that the overall reaction will not be spontaneous. An EMF of more than 4.07 volts will be required for this reaction to occur.

The electrolytic cell will have:

·  Inert electrodes (e.g. carbon)

·  Molten NaCl as the electrolyte

·  two half-reactions that are not separated by a salt bridge

·  an electrochemical cell (or other source of electric current)

Other important items to note:

§  At the anode – oxidation (+)

§  At the cathode - reduction (-)

The process:

1.  Electrons are "produced" in the battery at the anode, the site of oxidation.

2.  The electrons leave the electrochemical cell through the external circuit.

3.  These negative electrons create a negative electrode in the electrolytic cell which attracts the positive Na+ ions in the electrolyte. Na+ ions combine with the free electrons and become reduced (2Na+ + 2e- → Na)

4.  Meanwhile the negative Cl- ions in solution become attracted to the positive electrode of the electrolytic cell. At this electrode chlorine is oxidized, releasing electrons (Cl-→ Cl2 + 2 e-)

5.  These electrons travel through the external circuit, returning to the electrochemical cell.


Electrolysis of Sodium Chloride Solution – Competing reactions

The electrolysis of a pure molten substance like NaCl means there is only one chemical present at each electrode. However, there are sometimes several chemicals present at each electrode and thus, more than one possible reaction. So, the nature of the electrolyte can affect the chemical reactions that take place during electrolysis.

When there are competing reactions, we can look at a table of reduction potentials to decide which reaction is most likely to occur. The more likely of two reactions that will occur can be predicted by:

·  Reduction: the most + E0 on the reduction potential table (E0red most +)

·  Oxidation: the most - E0 on the reduction potential table (E0oxid most +)

In the electrolysis of concentrated salt water, there is more than one reaction possible at each electrode.

At the anode the possible reactions are:

Cl-(l) → Cl2 (g) + 2 e- E0oxid = -1.36V

2H2O(l) → O2 (g) +4H+(aq) + 4e- E0oxid = -1.23V

At the cathode the possible reactions are:

2H2O(l) + 2e-→ +H2(g) + 2OH- (aq) E0red = -0.83V

Na+(l) + e- → Na(s) E0red = -2.71V

Using the standard table of reduction potentials, the predicted reactions are the production of oxygen gas at the anode and the production of hydrogen gas at the cathode. This is, in fact, what occurs unless there is a high concentration of chloride ions, leading to the production of chlorine gas instead. This means, the concentration of the electrolyte can affect the products in an electrolytic cell.


Electrolysis of Saturated Sodium Chloride Solution

Sodium hydroxide is also produced using this type of electrolytic cell. Using concentrated NaCl solution, a semi-permeable membrane is inserted between the electrodes which allows only Na+ ions to migrate towards the cathode. Here they join with OH- ions that are produced in the reduction of water. Without this type of membrane, sodium hydroxide contains sodium chloride as a contaminant which is removed in subsequent processes.

Inert vs. Active Electrodes

Inert electrodes (carbon and platinum are the most common) generally do not take part in the reactions at the surface of the electrode. They simply provide a surface for the electron transfer reactions to take place.

If we were to substitute the carbon anode for a more active substance such as iron or silver, these electrodes may be oxidized themselves. This type of electrode is known as an active electrode as they actively take part in the reactions of the cell. Notice how much more positive (i.e. more likely) these oxidations are in comparison to water.

Ag(s) à Ag2+(aq) + 2e- E0 = -0.80V

Fe(s) à Fe2+(aq) + 2e- E0 = +0.44V

2H2o(l) à O2(g) + 4H+(aq) + 4e- E0 = -1.23V
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Electrolysis of Water

The following equation represents the breaking apart of H2O(l):

2H2O(l) → 2H2(g) + O2 (g)

The half-reactions for this process are:

E°
reduction / 2H2O(l) + 2e- → H2 + 2 OH- / -0.83 V
oxidation / 2H2O(l) → O2 + 4H+ + 4e- / -1.23V
net voltage required / - 2.06V

Again, the negative value indicates the need for more than 2.06V. The electrolytic cell is similar to the salt cell with a few minor alterations. These are:

·  Water has a low charge carrying capacity, so an H+ electrolyte is added to the water.

·  Hydrogen and oxygen gases are collected in inverted test tubes

At the anode (+) water will undergo oxidation:

2H2O(l) → O2 (g) + 4H+ (aq) + 4e-

At the cathode (-) water will be reduced:

2H2O(l) + 2e- → H2(g) + 2 OH-(aq)

Notice that there is two times as much H2 gas as there is O2 gas – the correct stoichiometric amount.


Electroplating

Electroplating is a technique in which a thin layer of a desired metal is used to coat (or "plate") another object. This process is often used to protect objects against corrosion or to improve their appearance.

For example, silver plating of flatware is a common use of this technique as it provides “silver” flatware at a fraction of the cost to pure silver.

For silver plating, the following is used:

§  An electrolyte solution of AgNO3 which will provide Ag+ ions,

§  a source of current (an electrochemical cell - a battery), and

§  two electrodes. One of the electrodes will be the object to be coated (flatware), while the other must be the plating metal (a bar of silver).

The half-reactions are simple and involve the reverse of the same reaction:

cathode / reduction / Ag+(aq)+ e- → Ag(s)
anode / oxidation / Ag(s) → Ag+(aq)+ e-

The process:

§  The flatware electrode acts as the cathode, or site of reduction.

§  Positive Ag+ ions from the solution will be deposited on the cathode (-) as solid silver.

§  As the Ag+ from the solution get used up, they will need to be replaced.

§  The silver bar acts as the anode and will be oxidised to Ag+ ions in solution.


Refining of Copper

Copper ore (generally copper pyrites) is roasted in air to produce an impure copper product known as “blister copper”. This product is generally about 98% pure and further refining is accomplished by electrolysis.

2CuFeS2 (s) + 5O2 (g) à 2Cu(s) + 2FeO(s) + 4SO2 (g)

The half-reactions are:

cathode / reduction / Cu2+(aq)+ 2e- → Cu(s)
anode / oxidation / Cu(s) → Cu2+(aq)+ 2e-

The process:

·  The blister copper is the anode and gradually dissolves

·  Pure copper forms on the cathode

·  Impurities such as Au an Ag are less readily oxidised than copper and fall to the bottom of the cell as a sludge

·  Other impurities such as Ni, Fe and Zn readily dissolve, but are not as likely to react at the cathode to form solid metals.

Introduction to Electrolysis, R. Slider Page 5