CHEM10007 Lecture Notes - Lecture 20: Galvanic Cell, Electrolytic Cell, Electrical Network
12 Jun 2018
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LECTURE 20 - ELECTROLYSIS
CHAPTER 12 - OXIDATION & REDUCTION
12.6 ELECTROLYSIS
•We know that spontaneous redox reactions can be used of generate electrical energy. Now we
will investigate how electrical energy can be used to force non spontaneous redox reactions to
occur.
•These are the kinds of reactions that occur when batteries (Daniell cells) are recharged.
WHAT IS ELECTROLYSIS
•Electrolysis is a chemical reaction that occurs when
electricity is passed through a molten ionic
compound or through a solution of an electrolyte.
•It converts electrical energy to chemical energy; it’s
the reverse of the spontaneous reaction occurring in
galvanic cells, using an external power source.
•An example of an electrolysis apparatus is an
electrolysis cell or electrolytic cell.
•A substance undergoing electrolysis must be molten
or in solution so that its ions can move freely and
conduction can occur.
•Inert electrodes, which will not react with the
molten substance, are dipped into the cell and then
connected to a source of direct current electricity.
•The direct current electricity source serves as an
electron pump, pulling electrons away from one
electrode and pushing them through the external
wiring onto the other electrode. Therefore, the
electrode from which electrons are removed becomes
positively charged, while the electrode which accepts
the electrons becomes negatively charged.
•At the positive electrode, oxidation occurs as the electrons are pulled away from the negatively
charged chloride ions. Therefore, the positive electrode becomes the anode, to which the anions
move.
•The direct current source pumps the electrons through the external electric circuit to the negative
electrode. Here, reduction takes place, as the electrons are forced onto positively charged sodium
ions. Therefore, the negative electrode becomes the cathode, to which cations move.
•In the above example of electrolysis of molten sodium chloride, these are the chemical changes
that occur at the electrodes;
COMPARISON OF ELECTROLYTIC CELLS & GALVANIC CELLS
•In a galvanic cell, the spontaneous reaction deposits electrons on the anode and removes them
from the cathode, resulting in the anode carrying a slight negative charge and the cathode a slight
positive charge.
•Galvanic cells produce energy, a spontaneous
reaction occurs and their chemical energy turns into
electrical energy.
•In an electrolytic cell, the situation is reversed.
Here, oxidation at the anode must be forced to
occur, which requires that the anode is positive so
that it can remove electrons from the reactant at
that electrode, and the cation is negative so that it
can force the reactant at the electrode to accept
electrons.
•Electrolytic cells consume energy, a non-spontaneous reaction occurs and their electrical energy
turns into chemical energy.
•The 4 key rules are:
1.
•An example of the reversal of a spontaneous reaction using an external power source (below)
•Note: The zinc electrode is now the cathode, and the copper electrode is now the anode.
•The polarity of the electrodes: Zn is still negative and Cu is still positive.
ELECTROLYSIS IN AQUEOUS SOLUTIONS
•When electrolysis is carried out in an aqueous solution, the electrode reactions can be more
complicated, as we have to consider the oxidation and reduction of the solute as well as oxidation
and reduction of the water.
‣Eg. The electrolysis of a solution of potassium sulfate gives hydrogen and oxygen.
‣Sulfate and nitrate ions will always be spectator ions.
‣At the cathode, water is
reduced, not K+:
‣At the anode, water is
oxidised, not the sulfate ion:
‣If we examine reduction potential data, at the cathode we have the following competing
reduction reactions:
‣Water has a much less negative reduction potential than K+, which means that H2O is
much easier to reduce than K+. During electrolysis, the more easily reduced substance is
reduced and H2 is formed at the
cathode. At the anode, the possible
oxidation half reactions are:
‣In
the
table,
they are written in the opposite direction:
‣The
E0
values
tell us that S2O82- is more easily reduced
than O2. So if S2O82- is more easily
reduced, then the product, SO42-, must
be less easily oxidised.
‣Basically, the half reaction with the smaller (less positive) reduction potential occurs more
easily as an oxidation.
‣So, during electrolysis, water is oxidised instead of SO42- and O2 is formed at the anode.
‣The overall cell reaction for the electrolysis of the K2SO4 solution is:
Document Summary
12. 6 electrolysis: we know that spontaneous redox reactions can be used of generate electrical energy. Now we will investigate how electrical energy can be used to force non spontaneous redox reactions to occur: these are the kinds of reactions that occur when batteries (daniell cells) are recharged. Therefore, the electrode from which electrons are removed becomes positively charged, while the electrode which accepts the electrons becomes negatively charged: at the positive electrode, oxidation occurs as the electrons are pulled away from the negatively charged chloride ions. Therefore, the positive electrode becomes the anode, to which the anions move: the direct current source pumps the electrons through the external electric circuit to the negative electrode. Here, reduction takes place, as the electrons are forced onto positively charged sodium ions. Therefore, the negative electrode becomes the cathode, to which cations move: in the above example of electrolysis of molten sodium chloride, these are the chemical changes that occur at the electrodes;
