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ENT 487
ENVIRONMENTALLY ASSISTED
CRACKING IN METALS
DR. HAFTIRMAN
LECTURE 12
WED, 8 OCTOBER 2008
ENVIRONMENTALLY ASSISTED CRACKING IN
METALS
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Environmentally assisted cracking (EAC) is a common
problem in a variety of industries.
In the petroleum industry, for example, EAC is pervasive
(merembes). Offshore platforms are susceptible to
corrosion-assisted fatigue.
Equipment in refineries and petrochemical plants are
exposed to a myriad of aggressive environments that
lead to stress corrosion cracking and hydrogen
embrittlement.
Similar problems exist in other settings, including fossil
and nuclear power plants, pulp and paper plants, ships,
bridges, and aircrafts.
Environmental cracking can occur even when there are
no visible signs of corrosions.
CORROSION PRINCIPLES
1.
ELECTROCHEMICAL
REACTIONS
All corrosion processes
involve
electrochemical
reactions. Figure illustrates
a simple electrochemical
cell. The anode and
cathode are physically
connected to one another
and are immersed in a
conductive medium called
an electrolyte. Atom from
the anode material give up
electrons, resulting in ions
being released into the
electrolyte and electrons
flowing to the cathode.
ELECTROCHEMICAL REACTIONS
Note that the corrosion cell forms an
electrical circuit. There is a voltage
drop,ΔE, between the anode and
cathode. Over time, the anode is
consumed( corrodes), as it releases ions
into the electrolyte.
ELECTROCHEMICAL REACTIONS
In cases where the two electrodes in an
electrochemical cell are different metals,
the anode is the metal that has a higher
propensity to give up electrodes
(oxidize).
 For example, in an electrochemical cell
with gold and iron electrodes, iron would
be the anode because it oxidizes more
readily than gold.
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ELECTROCHEMICAL REACTIONS
An electrochemical cell need not include
a bond between dissimilar metals. A
single metal in contact with an electrolyte
may be sufficient to form a corrosion cell,
depending on the respective chemical
compositions
of
the
metal
and
electrolyte.
 For example, consider a coupon of iron
immersed in hydrochloric acid (HCl).
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ELECTROCHEMICAL REACTIONS
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The chemical reaction is Fe  2HCl  FeCl2  H2
The iron is consumed by this reaction and hydrogen
gas (H2) is generated. If we consider only the
interaction between iron and hydrogen, the above
reaction can be written in the following forms:
Fe  2H   Fe2  2H  Fe2   H 2
Therefore, iron reacts with hydrogen ions to form iron
ions, atomic hydrogen, and hydrogen gas. This
reaction can be divided into two parts:
Fe  Fe2  2e
2 H   2e  2 H  H 2
ELECTROCHEMICAL REACTIONS
Iron is oxidized to iron ions and hydrogen
ions are reduced to H atoms that can be
either be absorbed by the electrode or
recombined and evolve into electrolyte
as hydrogen gas.
 The former is an anodic reaction and the
latter is a cathodic reaction.
 An oxidizing indicates a reducing or
cathodic reaction.
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ANODIC AND CATHODIC REACTIONS
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Figure schematically
illustrates the anodic
and cathodic
reactions occur at
the same physical
location.
 Every corrosion
process consist of an
anodic and cathodic
reaction.
ANODIC AND CATHODIC REACTIONS
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The anodic reaction normally involves the
oxidation of a metal to its ion. The general
form for the anodic reaction is given by
M M
n
 ne
Where n is number of electrons produced, which
equals the valence of the iron.
Most metallic corrosion processes involve one
or more of the cathodic (reduction) reactions.
CATHODIC REACTIONS
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Hydrogen evolution:
2 H  2e  H 2
 Oxygen reduction (acid solution):
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O2  4 H   4e  2 H 2O
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Oxygen reduction (neutral or basic solutions):
O2  2H 2O  4e  4OH 
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Metal ion reduction:
Metal deposition:
M
M
3
n
e M
2
 ne  M
CATHODIC REACTIONS
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The overall reaction is
2Fe  2H 2O  O2  2Fe2   4OH   2Fe(OH )2
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Ferrous hydroxide, which is the product of the
above reaction, is unstable in oxygenated
water. It oxidizes to ferric hydroxide, which is
known to the layperson as rust:
1
2 Fe(OH ) 2  H 2O  O2  2 Fe(OH )3
2
Note
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Both water and oxygen are required to corrode
steel.
 Steel that is completely submerged in water
normally corrodes very slowly because the
cathodic reaction is starved for oxygen.
 Steel corrodes most quickly when there is an
ample supply of both moisture and oxygen,
such as in a climate with high relative humidity
and frequent rain showers.
 The corrosion rate is also accelerated if steel
is coupled galvanically to a more noble metal.
Note
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Consider a steel structure in a seawater
environment, such as an offshore platform.
The most aggressive environments occur just
above and below the water surface.
 In the splash zone above the surface, both
oxygen and water are plentiful. Within the first
few feet below surface, the water is oxygen
rich because wave motion traps air bubbles.
 The relatively simple situation is complicated
by tight crevice geometries, the presence of
additional dissolved ions in the electrolyte, and
the imposed cathodic protection.
Corrosion Current an Polarization
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Since corrosion is an electrochemical process, the magnitude of
the electric current in the corrosion cell is a fundamental measure
of the corrosion rate. The corrosion current can be reduced by
inhibiting either reaction, or by reducing the conductivity of the
electrolyte.
When an electrochemical reaction is retarded by one or more
environmental factors, it is said to be polarized. There are three
types of polarization: activation polarization, concentration
polarization, and resistance polarization.
Activation polarization refers to processes that are controlled by
the rate of the reaction at the metal-electrolyte interface.
Concentration polarization occurs when the rate-limiting step is
diffusion of ions in the electrolyte.
Corrosion Current an Polarization
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Resistance polarization is a consequence of the
resistivity of the electrolyte. A reaction can also be
polarized by an externally applied current (galvanic
polarization) or potential (potentiostatic polarization).
Resistance polarization is a major factor in the
corrosiveness of seawater compared to tap water and
de-ionized water.
Seawater is very conductive because there is a ample
supply of sodium and chlorides ions, while de-ionized
water has relatively low electrical conductivity.
Normal tap water falls somewhere between the
extremes.
ELECTRODE POTENTIAL AND PASSIVITY
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A key factor that controls the corrosion current is the electrode
potential. The simple corrosion cell in previous Figure, which
showed an electric potential drop (ΔE) between the anode and
cathode. The elctrode potential refers to the half-cell potential of
the electrode. It is define as the potential difference between the
electrode of interest and a reference electrode, such as a
standard hydrogen electrode (SHE). The magnitude of the
electrode potential is a function of the chemical composition of
the electrode and the oxidizing power of the electrolyte.
The oxidizing power is a function of the reagents that are present
as well as their concentration.
Normally, the corrosion current increases exponentially with
increasing electrode potential. However, many technologically
important materials ( steel, aluminum, and titanium alloys) exhibit
a more complex behavior call passivity.
POLARIZATION DIAGRAM OF A METAL THAT
EXHIBITS PASSIVITY EFFECTS
POLARIZATION DIAGRAM OF A METAL THAT
EXHIBITS PASSIVITY EFFECTS
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Figure illustrates the typical behavior of a metal that exhibits
passivity effects. There are three distinct regimes: active,
passivity, and transpassive. In the active region, a small increase
in the electrode potential causes a large increase in corrosion
rate. A plot of electrode potential vs the logarithm of current
density is a straight line in the active region. As electrode
potential is increased further by any of the polarization processes,
the current density exhibits a sudden decrease at the beginning
of the passive region. The corrosion rate in the passive region is
typically 3 to 6 orders of magnitude slower than one would predict
by extrapolating the trend in the active region.
In the passive region, a surface film that acts as a protective
barrier forms on the surface. This surface film remains stable over
a wide range of electrode potential. The surface film breaks down
in the transpassive region due to the presence of very powerful
oxidizers. The highly protective surface films are very thin,
perhaps tens of nanometers. Such films are easily damaged by
mechanical means, but quickly reform to prtect the metal from
corrosion.