| Anions move toward the anode;
Cations move toward the cathode.
When something is oxidized, it
Think of iron rusting. The iron (zero charge) is oxidized to Fe3+.
Electrons are stripped away from the iron to give it a positive net
Reducing the rust
to iron involves giving back electrons
so Fe3+ can become neutral Fe. One way of looking at this
is that the magnitude of the positive
is being reduced.
Part II: Rational Design
Electrolysis of certain compounds could produce
or unstable products.
We can't control what you do or how you do it. To use this website or
can design simple electrolysis demonstrations starting with a table of
standard reduction potentials. These are available in almost any
chemistry reference, such as the CRC
of Chemistry and Physics. However, the reduction potentials
themselves don't entirely answer the complexity of what goes on in an
cell. Often there are side reactions and other variables which can
be so complex that it's difficult to predict them with theory.
Let's go over some things one should at least consider when designing
Suppose our reference book gives only the standard reduction potentials. Why not
too? It's not necessary. To get the oxidation potentials, just
switch the sign. For example, the reaction
Pb2+ + 2e- <----> Pb
has a reduction potential of -0.126 volt relative to the Standard
Electrode. The reverse reaction (in other words, the oxidation of Pb
to Pb2+) therefore has an oxidation potential of +0.126 volt
to the Standard Hydrogen Electrode. The positive value means the
of lead is actually favored. Metallic lead tends to oxidize to Pb2+
ions, although spontaneous transformation is slow.
/ reduction potentials indicate energetic favorability, but they
tell us nothing about how long a reaction could take.
A positive potential means the
over the reactants. In other words, the products have a lower energy
than the reactants.
A negative potential means reactants are energetically favored
products; in other words, energy would have to be added to the
in order to turn them into the product(s).
Suppose one reaction has a higher (i.e., less negative)
Let's say Half-Reaction A has a potential of -1.455 volts, while
B has a potential of -0.815 volts. If both half-reactions could
occur in solution, we might expect Half-Reaction B to occur
because its potential is less negative. Sometimes this doesn't
2H2O <---> O2 + 4H+ + 2e-
is energetically less difficult than
Pb2+ + 2H2O <---> PbO2 + 4H+
+ 2e- (-1.455 volt),
at least according to the oxidation potentials.
In real life, however, the formation of oxygen from water generally
an overvoltage in order to
the actual overvoltage depends on the type of anode material. The
is the excess voltage required to drive the reaction under real circumstances (that is, over
above the theoretical voltage). Oxygen overvoltages can be as much
as 1 volt, depending on the anode composition. There are also hydrogen
overvoltages, chlorine overvoltages, etc.; in any case, they depend
greatly on the electrode material.
is one more reason why much of electrolysis involves
with what works and what doesn't. Don't mistake this for random
trial-and-error; search the literature instead.
are given in terms of variation from the standard hydrogen potential,
which itself represents an arbitarily-assigned value (i.e.,
electrochemistry's answer to the Prime Meridian). One cannot expect
these values to translate readily to actual applied voltages.
we form PbO2 in an electrolytic cell. The dissociation of
PbO2 back into Pb2+ and water is much more
favorable than the formation of the lead dioxide. If this compound
could find its way back to the cathode, reduction would indeed occur;
our PbO2 would cease to exist. This doesn't happen,
and the reason is simple: PbO2 is not all that soluble;
it forms a heavy precipitate that either clings to the anode or falls
the bottom of the cell. Physically, it never makes it back over to the
another example. Let's imagine we form hydrogen from an aqueous
solution via the half-reaction:
2H+ + 2e- <---> H2,
which is actually the Standard Hydrogen Electrode half-reaction (E0
= 0.00000 volts).
of its oxidation or reduction potential, this one is no different
from any other electrolytic half-reaction; in other words, it could
go in either direction (depending on applied voltage). The reason this
doesn't happen in your electrolysis cell, however, is that H2
a gas with poor solubility in water. As soon as it forms, it vacates
the cell. The equilibrium essentially becomes a one-way reaction.
on the other hand, that we made sodium persulfate in an electrolytic
cell. S2O82- ions don't form an
precipitate with anything in our cell. They don't form a gas that
the cell. As a result, they stay in solution and can meander over to
the cathode where they are not going to last very long (reduction
of persulfate ion is around +2.0 volts relative to the H electrode,
higher in acidic conditions). Therefore, we'd better use a salt bridge
or some other way to keep S2O82- away
example where "logistics" come into play is the reduction of Fe3+
to Fe2+ in an electrolytic cell. Supposing the voltage and
current combination we're using are working smoothly, we still have to
that atmospheric oxygen will oxidize Fe2+ back to Fe3+.
This will happen most readily at the interface of electrolyte and air.
Obviously, you don't want to cap a cell that produces gas bubbles, since
buildup could burst it. If the gases form an explosive mixture, the
cap could make them build up to dangerous levels. However, a
cover can greatly reduce the amount of atmospheric oxygen that can act
4. Proven Systems
By now there
are many well-characterized systems. It is possible to
run them with some set of parameters to yield known products in known
Electrolysis of water, mentioned before, is probably the simplest
example. There are also others, especially ones that are or have been
used for industrial processes. One of the best-known (also mentioned
previously) is electrolytic production of chlorine bleach; indeed,
it seems there are few general chem textbooks that don't give it at
Back to Part I of
Electrolysis Apparatus - information and on-line ordering
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