lectrochemistry: Electrolysis
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Goal: to understand and explore the process of electrolysis
Working Definitions:
Electrolysis is a electrochemical (redox) reaction
brought about by the application of a direct current.
Introduction:
Electrolytic cells use an external source of direct current (DC) to drive reactions that would not otherwise be spontaneous. Electrolytic cells do not produce a cell potential but require a potential to operate. Electrolytic reactions are used purify metals and to plate metals on many types of substrates.
Using the electrolytic process to purify a metal (refining):
Because impurities can dramatically decrease the conductivity of copper wires, impure copper must be purified. One method of purifying copper is by electrolysis.
When a strip of impure metallic copper is used as the anode in the electrolysis of an aqueous preparation of copper(II) sulfate, copper is oxidized. The oxidation of copper is more facile than the oxidation of water (see the standard oxidation potentials below) so metallic copper dissolves into solution as copper(II) ions, leaving behind many of the impurities (less active metals):
Cu(s) 
Cu2+(aq) + 2 e–
(anode)
E°ox(Cu) = –0.34 V vs. E°ox(H2O) = –1.23 V
The copper(II) ions formed at the anode migrate to the cathode where they are more easily reduced than water and metallic copper “plates” on the cathode where it can be collected:
2 e– +
Cu2+(aq) 
Cu(s)
(cathode)
E°red(Cu2+) = +0.34 V vs. E°red(H2O) = –0.80 V
We had to pass sufficient current between the electrodes to cause the otherwise non-spontaneous reaction to occur! By carefully regulating the electrical potential, the metallic impurities that are active enough to be oxidized with copper at the anode are not reduced at the cathode and copper is selectively deposited.
| Note: Not all metals are more easily reduced and/or oxidized than H2O. If this is the case, the electrochemical reaction that requires the least potential will occur first! For example, if we were to use Al electrodes (both anode and cathode), metallic Al would be oxidized at the anode: E°ox(Al) = +1.66 V vs. E°ox(H2O ) = –1.23 V but H2O would be reduced at the cathode and aluminum ions would remain in solution! E°red(Al3+) = –1.66 V vs. E°red(H2O) = –0.83 V |
Ready for a recap?
| In aqueous solutions at 25°C | |
| 2e– + Hg2+ (aq) Hg(l) |
E° = 0.86 V |
| 1e– + Ag+ (aq) Ag (s) |
E° = 0.80 V |
| 2e– + 2 H+ (aq) H2 (g) |
E° = 0 V |
| 4e– +
Zr4+ (aq)
Zr (s) |
E° = –1.53 V |
| 2e– + Mg+2 (aq) Mg
(s) |
E° = –2.37 V |
| Oxidation | Water |
| 6 H2O(l)
O2(g) + 4
H3O+(aq) + 4
e– |
E° = –1.23 V |
| Reduction | |
| 2e– + 2 H2O(l) H2(g) + 2
OH– (aq) |
E° = –0.83 V |
Use the previous table of standard reduction potentials to answer the following questions. There may be more than one correct answer!
Which
of these metal ions can be selectively reduced in an aqueous solution?
Is
the zirconium ion easier to reduce than water? Try again.
Well
done! Silver ion is much easier to reduce than water... A nice way
to silverplate flatware!
Is
the magnesium ion easier to reduce than water? Try again.
Which
of these metals can be refined using the experiment described for copper
in the preceeding section?
Though
zirconium ion is easier to oxidize than water, it is more difficult
to reduce. You would be unable to efficiently deposit zirconium metal
at the cathode. Try again.
Great
job! Silver ion is much easier to reduce that is water and is sufficiently
easier to oxidize.
Right!
Mercury ion is both easier to reduce and oxidize than is water.
Before moving to the next page, you should understand what is meant by electrolysis and how to determine if oxidation or reduction of a chosen species can selectively occur in the presence of water.
Electrolysis 
