The Foundations of Chemistry

Now that we know how to use standard reduction potentials, let us use them to explain
the reaction that occurs in the electrolysis of aqueous NaCl. The first two electrolytic
cells we considered involved moltenNaCl and aqueousNaCl (see Sections 21-3 and 21-4).
There was no doubt that in molten NaCl, metallic Na would be produced by reduction
of Na, and gaseous Cl 2 would be produced by oxidation of Cl. But we found that in
aqueous NaCl, H 2 O, rather than Na, was reduced. This is consistent with the less nega-
tive reduction potential of H 2 O, compared with Na.

E^0

2H 2 O 2 e88nH 2 2OH 0.828 V

Nae88nNa 2.714 V

The more easily reduced species, H 2 O, is reduced.
Electrode potentials measure only the relative thermodynamiclikelihood for various half-
reactions. In practice kinetic factors can complicate matters. For instance, sometimes the
electrode process is limited by the rate of diffusion of dissolved species to or from the
electrode surface. At some cathodes, the rate of electron transfer from the electrode to a
reactant is the rate-limiting step, and a higher voltage (called overvoltage) must be applied
to accomplish the reduction. As a result of these factors, a half-reaction that is thermody-
namicallymore favorable than some other process still might not occur at a significant
rate. In the electrolysis of NaCl(aq), Clis oxidized to Cl 2 gas (1.360 V), instead of
H 2 O being oxidized to form O 2 gas (1.229 V), because of the overvoltage of O 2 on Pt,
the inert electrode.

CORROSION

Ordinary corrosionis the redox process by which metals are oxidized by oxygen, O 2 , in
the presence of moisture. There are other kinds, but this is the most common. The
problem of corrosion and its prevention are of both theoretical and practical interest.
Corrosion is responsible for the loss of billions of dollars annually in metal products. The
mechanism of corrosion has been studied extensively. It is now known that the oxidation
of metals occurs most readily at points of strain (where the metals are most “active”).
Thus, a steel nail, which is mostly iron (Section 22-7), first corrodes at the tip and head
(Figure 21-11). A bent nail corrodes most readily at the bend.

21-17

21-17 Corrosion 873

See the Saunders Interactive

General Chemistry CD-ROM,

Screen 21.9, Corrosion.

Figure 21-11 (a) A bent nail

corrodes at points of strain and

“active” metal atoms. (b) Two nails

were placed in an agar gel that

contained phenolphthalein and

potassium ferricyanide, K 3 [Fe(CN) 6 ].

As the nails corroded they produced

Fe^2 ions at each end and at the

bend. Fe^2 ions react with

[Fe(CN) 6 ]^3 ions to form

Fe 3 [Fe(CN) 6 ] 2 , an intensely blue-

colored compound. The rest of each

nail is the cathode, at which water is

reduced to H 2 and OHions. The

OHions turn phenolphthalein

pink.

Rust

Pits

Rust

(a) (b)