RUSTING
OF
IRON
Rust is an iron oxide, a usually red oxide formed by the redox reaction
of iron and oxygen in the presence of water or air moisture. Several forms of rust are
distinguishable both visually and by spectroscopy, and form under different
circumstances.[1] Rust consists of hydratediron(III) oxides Fe2O3·nH2O and iron(III) oxide-
hydroxide (FeO(OH), Fe(OH)3).
Given sufficient time, oxygen, and water, any iron mass will eventually convert entirely
to rust and disintegrate. Surface rust is flaky andfriable, and it provides no protection to
the underlying iron, unlike the formation of patina on copper surfaces. Rusting is the
common term for corrosion of iron and its alloys, such as steel. Many
other metalsundergo similar corrosion, but the resulting oxides are not commonly called
rust.[2]
Other forms of rust exist, like the result of reactions between iron and chloridein an
environment deprived of oxygen. Rebar used in underwater concretepillars, which
generates green rust, is an example. Although rusting is generally a negative aspect of
iron, a particular form of rusting, known as "stable rust," causes the object to have a thin
coating of rust over the top, and if kept in low relative humidity, makes the "stable" layer
protective to the iron below, but not to the extent of other oxides, such as aluminum.[3]
Chemical reactions
Rust is another name for iron oxide,[4] which occurs when iron or an alloy that contains
iron, like steel, is exposed to oxygen and moisture for a long period of time. Over time,
the oxygen combines with the metal at an atomic level, forming a new compound called
an oxide and weakening the bonds of the metal itself. Although some people refer to
rust generally as "oxidation", that term is much more general and describes a vast
number of processes involving the loss of electrons or increased oxidation state, as part
of a reaction. The best-known of these reactions involve oxygen, hence the name
"oxidation". The terms "rust" and "rusting" only mean oxidation of iron and its resulting
products. Many other oxidation reactionsexist which do not involve iron or produce rust.
But only iron or alloys that contain iron can rust. However, other metals can corrode in
similar ways.
The main catalyst for the rusting process is water. Iron or steel structures might appear
to be solid, but water molecules can penetrate the microscopic pits and cracks in any
exposed metal. The hydrogen atoms present in water molecules can combine with other
elements to form acids, which will eventually cause more metal to be exposed. If
chloride ions are present, as is the case with saltwater, the corrosion is likely to occur
more quickly. Meanwhile, the oxygen atoms combine with metallic atoms to form the
destructive oxide compound. As the atoms combine, they weaken the metal, making the
structure brittle and crumbly.
Oxidation of iron
When impure (cast) iron is in contact with water, oxygen, other strong oxidants, or
acids, it rusts. If salt is present, for example in seawater or salt spray, the iron tends to
rust more quickly, as a result of electrochemical reactions. Iron metal is relatively
unaffected by pure water or by dry oxygen. As with other metals, like aluminium, a
tightly adhering oxide coating, a passivation layer, protects the bulk iron from further
oxidation. The conversion of the passivating ferrous oxide layer to rust results from the
combined action of two agents, usually oxygen and water.
Other degrading solutions are sulfur dioxide in water and carbon dioxide in water. Under
these corrosive conditions, iron hydroxide species are formed. Unlike ferrous oxides,
the hydroxides do not adhere to the bulk metal. As they form and flake off from the
surface, fresh iron is exposed, and the corrosion process continues until either all of the
iron is consumed or all of the oxygen, water, carbon dioxide, or sulfur dioxide in the
system are removed or consumed.[5]
When iron rusts, the oxides take up more volume than the original metal; this expansion
can generate enormous forces, damaging structures made with iron. See economic
effect for more details.
Associated reactions
The rusting of iron is an electrochemical process that begins with the transfer
of electrons from iron to oxygen.[6] The iron is the reducing agent (gives up electrons)
while the oxygen is the oxidising agent (gains electrons). The rate of corrosion is
affected by water and accelerated by electrolytes, as illustrated by the effects of road
salt on the corrosion of automobiles. The key reaction is the reduction of oxygen:
O2 + 4 e − + 2 H
2O → 4 OH−
Because it forms hydroxide ions, this process is strongly affected by the presence of
acid. Indeed, the corrosion of most metals by oxygen is accelerated at low pH.
Providing the electrons for the above reaction is the oxidation of iron that may be
described as follows:
Fe → Fe2+ + 2 e−
The following redox reaction also occurs in the presence of water and is crucial
to the formation of rust:
4 Fe2+ + O2 → 4 Fe3+ + 2 O2−
In addition, the following multistep acid–base reactions affect the course of
rust formation:
Fe2+ + 2 H2O ⇌ Fe(OH)2 + 2 H+
Fe3+ + 3 H2O ⇌ Fe(OH)3 + 3 H+
as do the following dehydration equilibria:
Fe(OH)2 ⇌ FeO + H
2O
Fe(OH)3 ⇌ FeO(OH) + H
2O
2 FeO(OH) ⇌ Fe2O3 + H
2O
From the above equations, it is also seen that the
corrosion products are dictated by the availability of water
and oxygen. With limited dissolved oxygen, iron(II)-
containing materials are favoured, including FeO and
black lodestone or magnetite(Fe3O4). High oxygen
concentrations favour ferric materials with the nominal
formulae Fe(OH)3−xOx⁄2. The nature of rust changes with
time, reflecting the slow rates of the reactions of solids.
Furthermore, these complex processes are affected by the
presence of other ions, such as Ca2+, which serve as
electrolytes which accelerate rust formation, or combine
with the hydroxides and oxides of iron to precipitate a
variety of Ca, Fe, O, OH species.
Onset of rusting can also be detected in laboratory with the
use of ferroxyl indicator solution. The solution detects both
Fe2+ions and hydroxyl ions. Formation of Fe2+ ions and
hydroxyl ions are indicated by blue and pink patches
respectively.
Prevention
Because of the widespread use and importance of iron and
steel products, the prevention or slowing of rust is the
basis of major economic activities in a number of
specialized technologies. A brief overview of methods is
presented here; for detailed coverage, see the cross-
referenced articles.
Rust is permeable to air and water, therefore the interior
metallic iron beneath a rust layer continues to corrode.
Rust prevention thus requires coatings that preclude rust
formation.
Rust-resistant alloys
See also: Stainless steel and Weathering steel
Stainless steel forms a passivation layer of chromium(III)
oxide.[7][8] Similar passivation behavior occurs
with magnesium, titanium, zinc, zinc
oxides, aluminium, polyaniline, and other electroactive
conductive polymers.[citation needed]
Special "weathering steel" alloys such as Cor-Ten rust at a
much slower rate than normal, because the rust adheres to
the surface of the metal in a protective layer. Designs
using this material must include measures that avoid
worst-case exposures, since the material still continues to
rust slowly even under near-ideal conditions.[citation needed]
Galvanization
Main article: Galvanization
Galvanization consists of an application on the object to be
protected of a layer of metallic zinc by either hot-dip
galvanizing or electroplating. Zinc is traditionally used
because it is cheap, adheres well to steel, and
provides cathodic protection to the steel surface in case of
damage of the zinc layer. In more corrosive environments
(such as salt water), cadmium plating is preferred.
Galvanization often fails at seams, holes, and joints where
there are gaps in the coating. In these cases, the coating
still provides some partial cathodic protection to iron, by
acting as a galvanic anode and corroding itself instead of
the underlying protected metal. The protective zinc layer is
consumed by this action, and thus galvanization provides
protection only for a limited period of time.
More modern coatings add aluminium to the coating
as zinc-alume; aluminium will migrate to cover scratches
and thus provide protection for a longer period. These
approaches rely on the aluminium and zinc oxides
reprotecting a once-scratched surface, rather than
oxidizing as asacrificial anode as in traditional galvanized
coatings. In some cases, such as very aggressive
environments or long design life, both zinc and
a coating are applied to provide enhanced corrosion
protection.
Typical galvanization of steel products which are to subject
to normal day to day weathering in an outside environment
consists of a hot dipped 85 µm zinc coating. Under normal
weather conditions, this will deteriorate at a rate of 1 µm
per year, giving approximately 85 years of protection.[citation
needed]
Cathodic protection
Main article: Cathodic protection
Cathodic protection is a technique used to inhibit corrosion
on buried or immersed structures by supplying an electrical
charge that suppresses the electrochemical reaction. If
correctly applied, corrosion can be stopped completely. In
its simplest form, it is achieved by attaching a sacrificial
anode, thereby making the iron or steel the cathode in the
cell formed. The sacrificial anode must be made from
something with a more negative electrode potential than
the iron or steel, commonly zinc, aluminium, or
magnesium. The sacrificial anode will eventually corrode
away, ceasing its protective action unless it is replaced in a
timely manner.
Cathodic protection can also be provided by using a
special-purpose electrical device to appropriately induce
an electric charge.[9]
