"Corrosion: A CASE OF ENVIRONMENTAL
ELECTROCHEMISTRY"
You may be thinking that spontaneous electrochemical processes
are always beneficial, but consider the problem of corrosion, the natural
redox process that oxidizes metals to their oxides and sulfides. In
chemical terms, corrosion is the reverse of isolating a metal from its
oxide or sulfide ore; in electrochemical terms, the process shares many
similarities with voltaic cells. Although we focus here on the corrosion of
iron, many other metals, such as copper and silver, also corrode.
The most common and economically destructive form of corrosion
is the rusting of iron. Rust is not a direct product of the reaction
between iron and oxygen but arises through a complex electrochemical
process. Let's look at the facts of iron corrosion and then use the
features of a voltaic cell to explain them:
1- Iron does not rust in dry air: moisture must be present.
2- Iron does not rust in air-free water: oxygen must be present.
3- The loss of iron and the depositing of rust often occur at different
places on the same object.
4- Iron rusts more quickly at low pH (high [H+]).
5- Iron rusts more quickly in contact with ionic solutions.
6- Iron rusts more quickly in contact with a less active metal (such as Cu)
and more slowly in contact with a more active metal (such as Zn).
Strains, ridges and dents in contact with water are typically the
sites of iron loss (fact 1). these sites are called anodic regions because
the following half-reaction occurs there:
( )
(
)
[
]
hence the iron atoms lose electrons, the damage to the object gas been
done, and a pit forms where the iron is lost.
The freed electrons move through the external circuit until they
reach a region of relatively high O2 concentration (fact 2), near the
surface of a surrounding water droplet, for instance. At this cathodic
region, the electrons released from the iron atoms reduce O2 molecules:
( )
(
)
()
[
]
Notice that this overall redox process is complete, and thus the iron loss
has occurred, without any rust forming:
( )
( )
(
)
(
)
()
Rust forms through another redox reaction in which the reactants
contact each other directly. The Fe2+ ions formed originally at the anodic
region disperse through the surrounding water and react with O2, often
at some distance from the surface(fact 3). The overall reaction for this
step is
(
)
( )
(
)
( )
()
(
)
[The inexact coefficient n for H2O in the above equation appears
because rust, Fe2O3.nH2O, is a form of iron (III) oxide with variable
numbers of waters. The rust deposit is really incidental to the damage
caused by loss of iron—a chemical insult added to the original injury.
Adding the previous two equations together shows the overall
equation for …. Rusting of iron:
( )
( )
()
( )
(
(
)
)
we've shown the canceled H+ ions to emphasize that they act as a
catalyst; that they are used up in one step of the overall reaction and
created in another. As a result of this action, rusting is faster at low pH
(high [H+]) (fact 4). Ionic solutions speed rusting by improving the
conductivity of the aqueous medium near the anodic and cathodic
regions (fact 5).