This document discusses electrochemistry and provides definitions and examples of key concepts. It defines oxidation as the loss of electrons and reduction as the gain of electrons. An electrochemical reaction is one where a chemical reaction produces an electric current or where an electric current causes a chemical reaction. There are two types of electrochemical cells - galvanic cells where a chemical reaction produces a current, and electrolytic cells which use electricity to drive a non-spontaneous reaction. The document also discusses standard electrode potentials and how they are used to calculate cell emf and predict reaction spontaneity. Industrial applications of electrochemical cells mentioned include electrolysis, chloralkali production, and metal extraction from ores.

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13. Electrochemistry
Chemistry
Grade 12

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Revision
Oxidation is the loss of electrons and reduction is the gain of electrons.
A redox reaction is one where there is always a change in the oxidation numbers of the
elements that are involved in the reaction.
It is possible to balance redox equations using the half-reactions that take place within the
overall reaction.

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Electrochemical reactions: definitions
An electrochemical reaction is one where either a chemical reaction produces an electric
current, or where an electric current causes a chemical reaction to take place.
An electrochemical cell is a device where electrochemical reactions take place. There are
two types of electrochemical cell: galvanic and electrolytic.
An electrode is an electrical conductor that connects the electrochemical species from its
solution to the external electrical circuit of the cell. In electrochemical cells the two electrodes
are referred to as the anode and the cathode.
An electrolyte is a solution that contains free ions, and which therefore behaves as a
conductor of charges (electrical conductor) in solution.
A salt bridge is a material which contains electrolytic solution and acts as a connection
between two half-cells (completes the circuit). It maintains electrical neutrality in and between
the electrolytes in the half-cell compartments.

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Galvanic cells
In a galvanic cell a chemical reaction produces a current in the external circuit. An example is
the zinc-copper cell.
In a galvanic cell each electrode is
placed in a separate container in an
electrolyte solution.
The two electrolytes are connected
by a salt bridge.
We can use standard notation to
represent a galvanic cell:
X (s) | X+
(aq) || Y+
(aq) | Y (s)

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Electrolytic cell
An electrolytic cell is an electrochemical cell that uses electricity to drive a non-spontaneous
reaction. In an electrolytic cell, electrolysis occurs, which is a process of separating elements
and compounds using an electric current.
In an electrolytic cell both electrodes are placed
in the same container in an electrolyte solution.

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Standard electrode potentials and cell EMF
Different metals have different reaction potentials. The reduction potential of metals (in other
words, their ability to ionise), is recorded in a table of standard electrode reduction potentials.
The more negative the value, the greater the tendency of the metal to be oxidised. The more
positive the value, the greater the tendency of the metal to be reduced.
The values on the table of standard electrode potentials are measured relative to the
standard hydrogen electrode.
The EMF of an electrochemical cell can be calculated using one of the following equations:
It is possible to predict whether a reaction is spontaneous or not, either by looking at the sign
of the cell EMF or by comparing the electrode potentials of the two half-cells. A negative EMF
indicates that the reaction will not occur spontaneously.
E°(cell)
= E° (reduction half-reaction) – E° (oxidation half-reaction)
E°(cell)
= E° (oxidising agent) – E° (reducing agent)
E°(cell)
= E° (cathode) – E° (anode)

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Applications
Industrial applications of cells include electrolysis (the electrowinning of copper), in the
chloralkali industry (mercury, diaphragm and membrane cells), as well as the extraction of
metals from ores (e.g. aluminium from bauxite).
Electrowinning of copper

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Applications – chloralkali
The chlorine-alkali (chloralkali) industry is an important part of the chemical industry, which
produces chlorine and sodium hydroxide through the electrolysis of the raw material brine.
Three types of cell are used: mercury cell, diaphragm cell and membrane cell.
Mercury cell
This method only produces a
fraction of the chlorine and
sodium hydroxide that is used
by industry. The chlorine
produced is very pure, the
cell requires a lot of electricity
to run and produces toxic
mercury waste.

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Diaphragm cell
A porous diaphragm divides
the electrolytic cell into an
anode compartment and a
cathode compartment. This
cell uses less energy than the
mercury cell but the products
are not very pure.
Membrane cell
An ion-selective membrane
divides the electrolytic cell into
an anode compartment and a
cathode compartment.
This cell is the cheapest to
operate and produces high
purity products.

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