An electrolytic cell is an electrochemical
cell that drives a non-spontaneous redox reaction
through the application of electrical energy.
They are often used to decompose chemical
compounds, in a process called electrolysis—the
Greek word lysis means to break up.
Important examples of electrolysis are the
decomposition of water into hydrogen and oxygen,
and bauxite into aluminium and other chemicals.
Electroplating (e.g. of copper, silver, nickel
or chromium) is done using an electrolytic
cell.
Electrolysis is a technique that uses a direct
electric current (DC).
An electrolytic cell has three component parts:
an electrolyte and two electrodes (a cathode
and an anode).
The electrolyte is usually a solution of water
or other solvents in which ions are dissolved.
Molten salts such as sodium chloride are also
electrolytes.
When driven by an external voltage applied
to the electrodes, the ions in the electrolyte
are attracted to an electrode with the opposite
charge, where charge-transferring (also called
faradaic or redox) reactions can take place.
Only with an external electrical potential
(i.e. voltage) of correct polarity and sufficient
magnitude can an electrolytic cell decompose
a normally stable, or inert chemical compound
in the solution.
The electrical energy provided can produce
a chemical reaction which would not occur
spontaneously otherwise.
== Galvanic cells compared to electrolytic
cells ==
A galvanic cell can be considered an electrolytic
cell acting in reverse.
While electrolytic cells convert electrical
energy into chemical energy, galvanic cells
convert chemical energy into electrical energy.
Galvanic cells are often used in batteries.
== Anode and cathode definitions depend on
charge and discharge ==
Michael Faraday defined the cathode of a cell
as the electrode to which cations (positively
charged ions, like silver ions Ag+) flow within
the cell, to be reduced by reacting with electrons
(negatively charged) from that electrode.
Likewise he defined the anode as the electrode
to which anions (negatively charged ions,
like chloride ions Cl−) flow within the
cell, to be oxidized by depositing electrons
on the electrode.
To an external wire connected to the electrodes
of a Galvanic cell (or battery), forming an
electric circuit, the cathode is positive
and the anode is negative.
Thus positive electric current flows from
the cathode to the anode through the external
circuit in the case of a Galvanic cell.
Consider two voltaic cells of unequal voltage.
Mark the positive and negative electrodes
of each one as P and N, respectively.
Place them in a circuit with P near P and
N near N, so the cells will tend to drive
current in opposite directions.
The cell with the larger voltage is discharged,
making it a galvanic cell, so P is the cathode
and N is the anode as described above.
But, the cell with the smaller voltage charges,
making it an electrolytic cell.
In the electrolytic cell, negative ions are
driven towards P and positive ions towards
N. Thus, the P electrode of the electrolytic
cell meets the definition of anode while the
electrolytic cell is being charged.
Similarly, the N electrode of the electrolytic
cell is the cathode while the electrolytic
cell is being charged.
== Uses ==
As already noted, water, particularly when
ions are added (salt water or acidic water),
can be electrolyzed (subject to electrolysis).
When driven by an external source of voltage,
H+ ions flow to the cathode to combine with
electrons to produce hydrogen gas in a reduction
reaction.
Likewise, OH− ions flow to the anode to
release electrons and an H+ ion to produce
oxygen gas in an oxidation reaction.
In molten sodium chloride, when a current
is passed through the salt the anode oxidizes
chloride ions (Cl−) to chlorine gas, releasing
electrons to the anode.
Likewise the cathode reduces sodium ions (Na+),
which accept electrons from the cathode and
deposits on the cathode as sodium metal.
NaCl dissolved in water can also be electrolyzed.
The anode oxidizes chloride ions (Cl−),
and Cl2 gas is produced.
However, at the cathode, instead of sodium
ions being reduced to sodium metal, water
molecules are reduced to hydroxide ions (OH−)
and hydrogen gas (H2).
The overall result of the electrolysis is
the production of chlorine gas and aqueous
sodium hydroxide (NaOH) solution.
Commercially, electrolytic cells are used
in electrorefining and electrowinning of several
non-ferrous metals.
Almost all high-purity aluminium, copper,
zinc and lead is produced industrially in
electrolytic cells.
== Cell types ==
Concentration cell
Electrochemical cell
Galvanic cell
