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Introduction to Electrolysis
Think of electrolysis and electrolytic cells as the opposite of electrochemical cells:
Energy
conversion
Spontaneous
chemical
reaction?
Value of E°
Electrochemical
Cells
Electrolytic
Cells
Chemical →
Electrical
Electrical →
Chemical
Yes
No
Positive
Negative
Source: www.physchem.co.za/Redox/Electrolysis.htm
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 nonspontaneous 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  Alumina
(Al2O3)  Al)
Copper refining (blister copper  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.
Electroplating: an application of electrolysis.
Source: http://content.answers.com
Introduction to Electrolysis, R. Slider
Page 1
Figure A
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.
Figure B
Introduction to Electrolysis, R. Slider
Page 2
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°
2Na+(l) + 2e- → Na(s)
reduction
oxidation
Cl-(l) → Cl2 (g) + 2 e-
net voltage required
-2.71 V
-1.36V
- 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 molten sodium chloride
Source: http://www.saskschools.ca
Introduction to Electrolysis, R. Slider
Page 3
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
Electrolysis of a concentrated salt solution (brine)
Source: www.answers.com/topic/electrolysis
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.
Introduction to Electrolysis, R. Slider
Page 4
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.
Industrial production of chlorine gas and sodium hydroxide
Source: http://www.drbateman.net/asa2sums/sum2.5/image2.jpg
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)  Ag2+(aq) + 2eFe(s)  Fe2+(aq) + 2e2H2o(l)  O2(g) + 4H+(aq) + 4e-
Introduction to Electrolysis, R. Slider
E0 = -0.80V
E0 = +0.44V
E0 = -1.23V
Page 5
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.
Electrolysis of water
Source: http://www.saskschools.ca
Introduction to Electrolysis, R. Slider
Page 6
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.
Silver electroplating of flatware
Source: http://www.saskschools.ca
Introduction to Electrolysis, R. Slider
Page 7
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)  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
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