Untitled - Kelly Walsh High School

Electrochemistry 277
Be Careful!
The same rules apply to the anode. From these rules, we see that the anode
half-reaction must be:
2 H 2O(l) l 4 H (aq) O 2(g) 4 e E 1.23 V
There is a phenomenon known as overvoltage, which leads to variations in these
rules. The water/oxygen half-reaction, shown here, is often subject to this complication.
If overvoltage is present, the oxygen value may become 1.40 V
instead of 1.23 V. Your instructor should let you know if you need to consider
overvoltage.
Once we know the two half-reactions that will occur, we can determine the cell
reaction using the final steps for balancing redox equations.
2 (2 H2O(l) 2 e l H2(g) 2 OH (aq))
2 H2O(l) l 4 H (aq) O2(g) 4 e 6 H 2O(l) 4 e l 2 H 2(g) 4 OH (aq) 4 H (aq) O 2(g) 4 e
The 4 H (aq) and the 4 OH (aq) become 4 H 2O(l):
6 H 2O(l) 4 e l 2 H 2(g) 4 H 2O(l) O 2(g) 4 e
We can now cancel to get the final balanced equation:
2 H 2O(l) l 2 H 2(g) O 2(g)
In an electrolysis process, such as this one, the potassium ions and the fluoride
ions are spectator ions. They must be present for the procedure to work, but
they will remain unchanged.
A redox reaction is a simultaneous reduction (gain of electrons) and oxidation
(loss of electrons). The oxidizing agent is the reactant in the reduction halfreaction,
while the reducing agent is the reactant in the oxidation half-reaction.
Balancing redox reactions involves balancing all the atoms in addition to the
number of electrons lost and gained. Galvanic (voltaic) cells use a redox reaction
to produce electricity, while electrolytic cells use electricity to produce a
desired redox reaction. The anode is the electrode at which oxidation takes