Surfaces and interfaces
Corrosion of metals
11
Corrosion and durability of metals
11.1
Introduction
“Rusting” is probably the most familiar, but by far not the only form of corrosion. In contrast
to mechanical damage, metal corrosion is a reaction of the metal with its environment,
starting from the surface of the metal. The actual corrosion reactions take place in a few
nanometers thick metal/electrolyte interface, which does not correspond to the bulk phases on
either the metallic or the electrolyte side. Furthermore, “corrosion products” may be present
as a thin, well-adhering oxidic surface film, which protects the underlying metal from further
corrosion (passive film). Studies of corrosion and passivation processes are thus closely
related to surface analysis.
11.2
Basic concepts of corrosion
11.2.1.
Corrosion as a short-circuited galvanic element
The term corrosion stands for the reaction of a material with its environment, which leads to a
measurable alteration of the material (properties, behavior), and may cause functional
impairment (damage) of a component part or the whole system. In the case of metallic
materials, the reaction is mostly electrochemical in nature.
Figure 11.1: Schematic of a corroding metal electrode. Formal breakdown into two half-cells of a galvanic
element (oxidation reaction at the anode, reduction reaction at the cathode).
Corrosion processes on metallic materials are, apart from few exceptions, always electrochemical processes (Redox processes). The total reaction (Fig. 11.1) can be formally split up
into two partial reactions:
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a) Oxidation reaction. This is the actual corrosion process, i.e. the metal dissolution
(conversion of iron atoms from the metallic into the ionic state) the oxidation reaction
takes place at the anode:
Fe Fe2+ + 2 eb) Reduction reaction. Due to the electro-neutrality principle, the electrons released during
the anodic reaction must be taken up by a part of the environment adjacent to the metal,
which is then reduced. This process is taking place at the cathode. If the corrosive agent is
an acidic solution, protons are reduced forming hydrogen gas:
2 H+ + 2 e- H2 (gas)
In contrast, if oxygen, dissolved in (neutral or alkaline) electrolytes, interacts with the
metal, oxygen is the oxidizing agent, i.e. it will be reduced:
O2 + 2 H2O + 4 e- 4 OHDue to the electro-neutrality (the electrons released from the iron atom need to be taken up by
the oxidizing agent), the total corrosion process is composed of at least one oxidation and one
reduction process, which must take place simultaneously (Fig. 11.1).
An anodic (positive) current corresponds to the iron dissolution, a cathodic (negative) current
corresponds to the reduction reaction. Since metals (iron) are electrical conductors and the
electrolyte is in general well electrolytically conductive, both the anodic and cathodic reaction
constitute a formally short-circuited galvanic element – a current I (corrosion current) is
flowing:
I = U / (Ra + Rc + Re)
The intensity of the corrosion current is determined by the voltage difference U of the
galvanic element and the resistance of the anode Ra, the cathode Rc and the electrolyte Re.
Thermodynamic and kinetic basic principles of the corrosion reaction allow the prediction of
whether a corrosion reaction is possible or not (thermodynamics) and how fast it proceeds
(kinetics). Both thermodynamic and kinetic considerations have to take both the metal
as well as its environment into account.
11.2.2 Thermodynamics
From the thermodynamic laws one can derive whether a corrosion reaction can take place or
not.
G < 0 :
G > 0 :
the reaction takes place
the reaction does not take place
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For electrochemical reactions, G is replaced by the cell potential U (G = nFU), which can
be calculated from the equilibrium potentials Ea und Ec of the anodic and cathodic partial
reactions, respectively:
U = Ea – Ec
The equilibrium potentials can be calculated with the aid of the standard potentials (tabulated)
and Nernst’s law:
Ea = E0 + 2.3 RT/nF ln (cMez+)
Letters stand for:
Ea
E0
c Mez+
n
normal potential
standard potential
concentration of metal ions in solution
number of transmitted electrons
Using the common logarithm, the equation can be written as:
Ea = E0 + 0.059/n * log (cMez+)
Electrochemical series of metals
The standard potentials E0 are tabulated for standard conditions (metal immersed in a solution
of its metal ions with the concentration cMez+ = 1 Mol/l):
Partial reaction
E0Me/Me+ (Volt)
Au Au3+ + 3eAg Ag+ + eCu Cu2+ + 2eH2 2H+ + 2ePb Pb2+ + 2eNi Ni2+ + 2eFe Fe2+ + 2eZn Zn2+ + 2eAl Al3+ + 3eMg Mg2+ + 2e-
+ 1.50
+ 0.80
+ 0.34
0.00
- 0.13
- 0.25
- 0.44
- 0.76
- 1.66
- 2.37
E'Me/Me+ (Volt)
with cMez+ = 10-6 mol/l
+ 1.38
+ 0.44
+ 0.16
- 0.30
- 0.42
- 0.61
- 0.94
- 1.78
- 2.54
- "Precious" metals possess standard potentials E0 > 0. The cell potential U is positive in
combination with a hydrogen electrode, i.e. no corrosion takes place in non-oxidizing
(oxygen free) acids.
- "Non-noble" metals possess standard potentials E0 < 0. The cell potential U is negative in
combination with a hydrogen electrode, i.e. corrosion in non-oxidizing acids is possible.
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The standard conditions cj = cj 0 are usually not fulfilled in the case of corrosion processes, i.e.
the corrosive agent contains only traces of metal ions, e.g. cMe+ = 10-6 mol/l. It is better for
practical calculations to use the calculated value of the potential Ea (see column 3 in table
above).
Electrochemical series of nonmetals
The application of Nernst’s equation to cathodic partial reactions results in the normal
potentials of the partial reactions.
Example:
Hydrogen electrode
H2 + 2H2O 2H3O+ + 2e-
(anodic notation)
For a hydrogen partial pressure pH2 = 1 and room temperature, the hydrogen electrode
potential is only dependent on the pH-value:
EH2/H+ = 0.059logcH+ = -0.059pH
Example:
Oxygen electrode
4OH- O2 + 2H2O + 4e-
(anodic notation)
For an oxygen partial pressure pO2 = 1 bar and room temperature, the oxygen electrode
potential is only dependent on the pH-value:
EOH-/O2 = + 1.27 - 0.059pH
The following table gives an overview of the normal potentials of some electrode reactions,
which are important as cathodic partial reactions in corrosion processes electrochemical
series of nonmetals:
Partial reaction
2Cl- Cl2 + 2e-
E0 (Volts)
+ 1.36
2Cr3+ Cr2O72- + 7H2O + 14H+ + 6e2Br- Br2 + 2e-
+ 1.33
NO + 2H2O NO3- + 4H+ + 3eFe2+ Fe3+ + e4OH- O2 + 2H2= + 4e-
+ 0.96
+ 0.77
+ 0.44
Cu Cu2+ + 2eH2 2H+ + 2e-
+ 0.34
0.00
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+ 1.07
Surfaces and interfaces
Corrosion of metals
Potential pH diagram (Pourbaix diagram)
The thermodynamic data of a corrosion system are graphically displayed in Pourbaix
diagrams. They provide information about possible equilibria between metal, solution with
dissolved metal ions, and stable oxygen compounds, as a function of the pH-value and the
electrode potential.
Figure 11.2: Pourbaix diagrams for different metals. Example: Aluminum corrodes in the acidic and alkaline
range; it is passive in the neutral range (pH 4-8).
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Corrosion of iron in non-oxidizing acids
Gross reaction:
Fe + 2HCl FeCl2 + H2
Partial reactions:
anodic (oxidation)
Fe Fe2+ + 2e-
cathodic (reduction)
2H+ + 2e- H2
Figure 11.3 schematically shows the different steps of a corrosion reaction at an electrode:
Figure 11.3: Schematic of corrosion processes of an iron bar in hydrochloric acid
The anodic as well as the cathodic partial reaction take place at the same electrode. This
means that both partial reactions cannot be separated, thus no external voltage can be
measured. The electrochemical cell is short-circuited. Nevertheless, a (theoretical) cell
potential can be defined, which is calculated from the potentials of the anodic Ea and cathodic
Ec partial reactions.
Practical series of elements
As already mentioned, thermodynamics can show which reactions are in principle possible.
When comparing the “electrochemical” and the “practical” series of elements (experimental
potential values), one can observe considerable differences in many cases. By definition,
thermodynamics does not allow us to draw conclusions for the reaction rate. Even if the
reaction is possible according to thermodynamics, it can proceed so slowly that there is no
practical corrosion risk; this is the case for the formation of protective oxide layers that
dissolve only very slowly (passivity). These lead to much more positive potentials than
predicted for the metal itself.
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Figure 11.4: Comparison between electrochemical and practical series of elements. In particular, note the
metals Al, Cr and Ni; their behavior is more “noble” in practice than expected in theory.
Thermodynamic calculations show whether and in which direction a reaction can
proceed. However, they do not allow conclusions to be drawn about reaction rates.
11.2.3 Kinetics of electrochemical reactions
During the contact of a metallic material with an aqueous environment (immersed or as a
liquid film through condensation of moisture), a potential difference is established over
the electric double layer. For mere adsorption of solvent molecules and ions, this potential
difference is determined by the condition that the charge has to be identical on both sides of
the double layer. Such adsorption phenomena can be observed for mercury, for example.
If charge transfer reactions are possible besides the adsorption processes (which is always the
case for corrosion processes), the potential difference can be determined kinetically.
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The two fundamental concepts of potential dependency and the additivity of partial reactions
can be combined as follows:
- Potential dependency
The rate of a charge-transfer reaction at an electrode depends, in addition to the
chemical factors (rate constant k , concentration c), exponentially on the electric field,
i.e. on the potential difference in the double layer:
[
i = k • c • exp (1 ) • F • RT
]
This exponential dependency of the current density on the potential forms the basis of
electrode kinetics.
- Additivity of the partial reactions
According to Wagner and Traud, anodic and cathodic partial reactions proceed
independently of each other at the metal surface. The corresponding partial current
densities ia and ic can be algebraically added up to the total current density is (Fig.
11.5):
is () = ia () + ik ()
This results in the total current density/potential curve, which can be measured
experimentally.
Figure 11.5: Total current/voltage diagram of a metal electrode corroding with H2 formation.
___total current density
----- partial current density
At the corrosion potential corr the electro-neutrality condition implies is = 0, so that
ia ( korr ) = ik ( korr )
i.e. both partial current densities have to be equal.
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This means that the corrosion potential corr – unlike the thermodynamic normal potential
EMe/Mez+ – is always a mixed potential, which is determined by the kinetics of the anodic and
cathodic partial reaction(s). Typical examples are the change of corrosion potential of zinc
with the pH-value of the solution (Figure 11.6a) or the influence of the oxygen content on the
corrosion potential of iron (Figure 11.6b).
Figure 11.6: a) Influence of the solution’s pH-value on the corrosion potential of zinc (schematically). pH 3 < 2
< 1. b) Influence of the oxygen content on the corrosion potential of iron (schematically). O2 content 1 < 2 <3
For uniform dissolution of a homogeneous alloy, anodic and cathodic partial reactions are
statistically distributed over the whole metal surface. Thus, for the whole metal, the same
potential is measured independently of the position of the reference electrode.
Charge transfer reactions
Current densities/potential curves (see Fig. 11.5) allow us to obtain information about the
corrosion mechanism and the corrosion rate. The additivity principle of Wagner and Traud
and the exponential potential dependence of electrochemical reactions lead to the
mathematical equation for the total current density/potential curve
korr
korr exp is = ikorr exp
bMe
bH
This equation has two limiting cases, which are very important for electrochemical
measurements:
1.
The electrode potential lies far from the corrosion potential. Then the cathodic (or
anodic) partial current can be neglected (see Fig. 11.5) and the experimentally
measured total current density equals exactly the anodic (or cathodic) partial current
density:
korr is = ikorr • exp bMe
korr is = ikorr • exp bH
anodic
cathodic
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Surfaces and interfaces
Corrosion of metals
Using this, one can study the kinetics of the anodic and cathodic partial reactions. In
the logarithmic scale, the equations result in straight lines (the so-called Tafel lines)
(-korr) = a + b * ln i
natural logarithm
(-korr) = a + b' * log i
common logarithm
From the slope, which equals the Tafel constant b
RT • n
= bMe =
,
lni
F • (1 )
bMe' = 2.3 • bMe
conclusions can be drawn about the mechanism of the electrochemical reaction.
2.
The electrode potential is very close to the corrosion potential. Thus ( - corr) 0
and the exponential functions can be approximated as
exp x (x0) = x
The equation for the total current density/potential curve results in
1
1
is ( ) = ikorr • ( korr ) • +
bMe bH This results in a straight line with the slope
1
1
dis
= Rp 1
= ikorr • +
bMe bH d
that determines the polarization resistance Rp of the system, from which the corrosion
rate for uniform corrosion can be calculated (Fig. 11.7). The Tafel constants of the
anodic partial reaction can be assumed to be 60 mV, the one for the cathodic partial
reaction can be assumed to be 120 mV.
Therefore, current density/potential curves can be used to determine the most important
values of a corrosion system, the corrosion rate icorr, and the mechanism of the ongoing
reaction.
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Surfaces and interfaces
Corrosion of metals
Figure 11.7: Polarization resistance as a slope of the total current density/potential curve at the corrosion
potential korr
Diffusion-controlled reactions
For corrosion reactions in which oxygen is the most important oxidizing agent (all natural
environments with neutral to alkaline pH), the oxygen reduction rate is only charge-transfer
controlled close to the equilibrium potential. For larger polarization (i.e. around the corrosion
potential), the rate is limited by the transport of O2 to the electrode. A concentration gradient
is therefore formed at the interface (oxygen depletion at the metal surface). This gradient
leads to O2 diffusion towards the electrode surface.
The diffusion current density i in the diffusion layer can be approximated with the help of
Fick’s first law:
i = - nFD dc
dx x=0
In this equation, n stands for the number of transferred electrons, F for the Faraday constant,
D for the diffusion constant and c for the concentration of the diffusing gas. When observing
corrosion processes it is sufficient to assume a linear concentration gradient in the diffusion
layer. Fick’s law becomes
i = - nFD c0 - c
with c0 as the concentration in the electrolyte, c as the concentration at the metal surface and
as the thickness of the diffusion layer.
The diffusion layer thickness is determined by the hydrodynamic conditions at the surface.
In the case of strongly stirred solutions, equals about 0.001 cm, while in stationary solutions
(natural convection) layer thicknesses of up to 0.05 cm can be reached (Fig. 11.8).
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Surfaces and interfaces
Corrosion of metals
If at the metal surface the solution is fully depleted (c = 0), a limiting current is present, which
can be calculated with the following equation:
c
igr = - nFD 0
Figure 11.8: Concentration profile in the diffusion interface according to Nernst. The diffusion layer depends
on the hydrodynamic conditions (flow rate).
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Surfaces and interfaces
Corrosion of metals
11.2.4 Faraday’s law
The corrosion current flowing is proportional to the amount of metal dissolved, i.e. the corrosion rate as given by Faraday’s law:
G= M I t
z F
Meaning of the symbols
G:
M:
z:
F:
I:
t:
transferred mass
atomic mass
valency of the metal ion
Faraday’s constant
electric current
time
g
g/mol
A.s/mol
A
s
With Faraday’s law the units used to express corrosion rates can be converted to each other
by consideration of metal density :
vcorr
dcorr
weight loss per time and area
thickness abrasion per time
icorr
vR
current density of metal dissolution
crack propagation rate
g/m2day
mm/year
A/cm2
m/s
For various, frequently used metals these rates are listed in the following table.
Reaction
Cu --> Cu2+
M = 63.57
= 8,92
z=2
Fe --> Fe2+
M = 55.85
= 7.86
z=2
Zn --> Zn2+
M = 65.38
= 7.13
z=2
Al --> Al3+
M = 26.97
= 2.70
z=3
i
(mA/cm2)
0.001
0.01
0.1
1.0
10.0
0.001
0.01
0.1
1.0
10.0
0.001
0.01
0.1
1.0
10.0
0.001
0.01
0.1
1.0
10.0
vcorr
(g/m2day)
0.285
2.845
28.454
284.54
2845.4
0.250
2.500
24.998
249.98
2499.8
0.293
2.926
29.264
292.64
2926.4
0.081
0.805
8.048
80.48
804.8
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dcorr
(mm/year)
0.012
0.116
1.164
11.64
116.4
0.012
0.116
1.160
11.60
116.0
0.015
0.150
1.498
14.98
149.8
0.011
0.109
1.088
10.88
108.8
vR
(m/s)
3.7 . 10-13
3.7 . 10-12
3.7 . 10-11
3.7 . 10-10
3.7 . 10-9
3.7 . 10-13
3.7 . 10-12
3.7 . 10-11
3.7 . 10-10
3.7 . 10-9
4.76 . 10-13
4.76 . 10-12
4.75 . 10-11
4.75 . 10-10
4.75 . 10-9
3.48 . 10-13
3.46 . 10-12
3.45 . 10-11
3.45 . 10-10
3.45 . 10-9
Surfaces and interfaces
Corrosion of metals
11.2.5 Corrosion rate - pH - diagrams
In practice, one often observes, e.g. for iron, a pH-dependent corrosion or mass-loss rate (Fig.
11.9). In highly acidic media, i.e. pH < 4, the dissolution rate strongly increases with
decreasing pH. In the pH ranges between 4 and 10.5, the corrosion rate effectively does not
depend on pH (diffusion control of oxygen reduction). In strongly alkaline media, for pH >
12, the corrosion rate is strongly decreasing.
Figure 11.9 Influence of the pH-value on the corrosion rate of steel
At low pH values the chemical composition and the structure of the iron lattice (annealed,
deformed) have a great influence on the corrosion rate, while in the neutral pH-range the
oxygen concentration, i.e. the flow rate of the liquid, and the oxygen access influence the
corrosion rate. In the alkaline range, iron and steel are spontaneously passivated and the
corrosion rate is practically zero.
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11.3
Corrosion of metals
Passivity
Although thermodynamically non-noble, many of these metals and their alloys are in practice
resistant to corrosion (see comparison of theoretical and practical electrochemical series of
metals, Fig. 11.4). This behavior is due to a spontaneous formation of a thin protecting oxide
layer, the passive layer, as a product of the anodic metal dissolution. The passive layer is a
kinetic barrier at the metal/electrolyte interface and inhibits further metal dissolution. But this
purely kinetic barrier is also the weak spot of passive metals: if the passive layer is destroyed
(e.g. by exposure to chloride), high corrosion rates can occur locally (pitting, stress-corrosion
cracking).
Typical metals and alloys that owe their corrosion resistance to passivation are the highalloyed steels, aluminum alloys, titanium and also normal steel in alkaline media (steel in
concrete). The alloys based on iron form very thin passive layers (nm range), the passive
layers on aluminum and titanium are in the μm range, their thickness can be increased by
anodization.
11.3.1 Formation of the passive layer
The passivation process is always associated with a (minimal) anodic dissolution, but the
resulting corrosion products are not dissolved in water or loosely present on the surface (rust);
they form a compact and adhesive oxide layer.
Figure 11.10: Current density-potential curve of a metal that can be passivated and the influence of Cr and Mo
on the anodic current density-potential curve of high-alloy steel.
The anodic current density-potential curve exponentially increases from the equilibrium
potential (charge-transfer controlled) and reaches a current-density maximum icrit at the
passivation potential p. The subsequent drop in current density, which can be several orders
of magnitude, is due to the passivation of the electrode. In the passive range itself, the metal is
only dissolving at very low rates (ip < 10-6 Acm-2) (passivity, anodic protection). The current
density of metal dissolution in the passive range is constant over a large potential range. At
very positive potential it increases again for some metals (Fe, Ni, Cr) due to transpassive
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dissolution of the passive film (e.g. chromium oxide Cr2O3 --> CrO42- (aq.)). The oxygen
formation (2 H2O --> O2 + 4H+ + 4e-) mostly takes place in the same potential range.
The following conditions have to be fulfilled for spontaneous formation of the protecting
passive layer (Fig. 11.11):
1. Ec > p, the normal potential of the cathodic partial reaction is more positive than the
passivation potential
2. |ic (p) | > icrit , the cathodic partial current density at the passivation potential is larger than
the critical current density
As seen in Fig. 11.11 and under the above conditions, both the metal (passivation potential
icrit) and the environment (cathodic partial current curve Ec, ic) play a role in passivation. A
certain alloy can thus be passive or not depending on the environment (Fig. 11.11).
Figure 11.11: Different partial reactions and their influence on passivation. 1 to 3 represent increasing oxygen
content in solution.
Spontaneous and stable passivation is only attainable in case 3 (Fig. 11.11). In this case, the
cathodic partial reaction (here O2 reduction) possesses a sufficiently high limiting current
density such that the critical current density icrit is exceeded. The corrosion potential corr then
lies in the passive range, the corrosion current density being icorr = ip.
The values p, icrit und ip depend on the type and composition of the alloy and on the
environment. Passivation is thus a typical system behavior; an alloy can be passive in one
environment while in another (e.g. oxygen-depleted) it may not be passive and corrode
actively (Fig. 11.11).
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Influence of the alloy composition: By alloying with chromium or molybdenum, the critical
current density icrit of iron-based alloys decreases (Fig. 11.12) and the passivation potential
moves to more cathodic (negative) values. Chromium also decreases the passive current
density, i.e. the dissolution rate in the passive state is lower. Molybdenum lowers the critical
current density icrit for the passivation (Fig. 11.12). The passive current density is only
marginally influenced by the Mo-content.
Figure 11.12: Influence of the alloy elements chromium and molybdenum on the critical current density for the
passivation (medium 0.1 N H2SO4).
Influence of the environment: For the environment, the pH-value and the O2 content are
crucial. The pH-value of the environment is especially important for iron, steel and iron-based
alloys, an increasing pH-value lowers the critical current density for the passivation (2 pHunits for a factor of 10). This is the reason for the high corrosion resistance of normal steel in
concrete.
Figure 11.13: Current density / Potential curves of an active (unalloyed steel) in comparison to a passive metal
(high-alloyed steel). Note the high difference of the corrosion potentials.
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11.3.2 Electrochemical investigation of passivation
The passivation behavior of an alloy/electrolyte system can be investigated with electrochemical experiments (potential measurement, current density/potential curves, potentiostatic
passivation). The formation of the protecting oxide layer manifests itself in an asymptotic
increase of the corrosion potential or in strongly decreasing current densities during
potentiostatic passivation (Fig 11.14).
Figure 11.14: Passivation of high-alloyed steels in alkaline 0.1 N NaOH solution. a) corrosion potential versus
time, b) passivation current density vs time ( = - 0.1 V SCE) D. Addari, PhD Thesis University Cagliari (2005)
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11.3.3 Composition of passive films
Passive films are oxide/hydroxide layers. For high-alloyed steels, it is mainly the chromium
content (in addition the molybdenum content) that facilitates passivation. Surface-analytical
methods such as XPS (X-ray photoelectron spectroscopy) allow the thickness and the
chemical composition of passive films to be determined.
Cr2p3/2
Fe2p3/2
1.4301
15 min 0.1 N NaOH
Intensity (a.u.)
Intensity (a.u.)
Fe15Cr
15 min 0.1 N NaOH
718
716
714
712
710
708
706
704
702
582
580
Binding Energy (eV)
• signal of metallic iron
• signal of Fe(II) in Fe3O4
satellite • signal of Fe(III) in Fe2O3
• signal of Fe(III) in FeOOH
578
576
574
572
Binding Energy (eV)
706.7 eV
708.8 eV
714.0 eV
710.3 eV
712.1 eV
• signal of metallic chromium
• signal of Cr(III) in Cr2O3
• signal of Cr(III) in Cr(OH)3
574.0 eV
576.6 eV
578.1 eV
Figure 11.15: High-resolution XPS spectra of iron Fe2p3/2 and chromium Cr2p3/2 and a Fe15Cr alloy after 15
min in 0.1 N NaOH pH 13. D. Addari, PhD Thesis University of Cagliari (2005)
Investigations on different high-alloyed steels (Tab. 11.2) showed that the chromium (III)
content in the passive film is determined by the pH-value of the environment; there is no big
difference between high-alloyed steel with different chromium content (Fig, 11.16).
Table 11.2 Cr, Ni and Mo content of various high-alloyed steels
Steel DIN
Cr %
Ni %
Mo %
Fe15Cr
15
-
-
1.4301
17.3
8.6
-
1.4462
20.2
4.2
2
1.4529
20.8
24.9
6.4
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100
100
composition (weight %)
80
60
40
20
0
Fe
Cr
Ni
Mo
80
composition (weight %)
Feox
Crox
Niox
Moox
60
40
20
C-steel
Fe15Cr
1.4301
1.4462
0
C-steel
Fe15Cr
1.4301
1.4462
Figure 11.16: Composition of the passive film (left) and the interface below the passive film (right) for different
iron-based alloys. Passivation = -0.1 V SCE, 0.1 N NaOH, 24 h. D. Addari, PhD Thesis University of Cagliari
(2005)
Potential dependency
XPS analyses of high-alloyed steel after potentiostatic passivation are used to investigate the
potential and pH dependency of passive film compositions in acidic solutions (0.1 N sulfuric
acid). With increasing potential the films become richer in iron and the thickness of the
passive films increases (Fig. 11.17), This trend could also be observed in less acidic and
neutral media.
Figure 11.17: Potential dependency of passive film composition (left) and passive film thickness (right) for
1.4301 CrNi steel in 0.1 N H2SO4. B. Beccu, PhD Thesis University of Cagliari (1997)
The pH value strongly influences the passive-film composition: The highest chromium
enrichment in passive films can be found in acidic media, an increasing pH-value leads to a
decrease of the chromium content in the film (Fig. 11.18).
Prof. B. Elsener, Institute for Building Materials, ETH Zürich
264
Surfaces and interfaces
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Figure 11.18: Passive film composition of high-alloyed steel DIN 1.4301 in 0.1 N sulfate solution for different
pH-values, potential 0.2 V SCE. B. Beccu, PhD Thesis University of Cagliari (1997)
As in alkaline media, pronounced nickel enrichment at the passive film/alloy interface can
also be found in acidic environments. This nickel enrichment is especially strong for the
highly durable and corrosion resistant CrNiMo steel (DIN 1.4529); the interface contains
(independently of the potential) more than 50% nickel (Fig. 11.20).
Figure 11.19: Comparison of alloy compositions (bulk) with the passive film (film) and the composition below
the passive film (interface). The amount of nickel in the interface is twice as much as in the alloy.
Prof. B. Elsener, Institute for Building Materials, ETH Zürich
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11.4 Localized corrosion
For passive metals and alloys, uniform corrosion is rare or only possible in a very aggressive
(acidic) environment. In the presence of chloride ions, microscopic local corrosion attacks
can take place under certain conditions.
11.4.1 Morphology
Localized corrosion can manifest itself as crevice corrosion, pitting, or intercrystalline
corrosion (Fig. 11.20).
Figure 11.20: Manifestations of local corrosion: crevice corrosion (left), pitting (middle), intercrystalline
corrosion (right)
In the case of pitting corrosion, local corrosion attacks form at the passive surface. Crevice
corrosion is to a large extent a geometrical problem (more intense attack in the crevices
arising from construction), and intercrystalline corrosion is a material problem (sensitization,
“non-noble” grain boundaries, which are preferentially attacked).
11.4.2 Pitting corrosion
Pitting only occurs on passive metals. For pitting to occur, some material and environmental
conditions have to be fulfilled (similar to passivation).
Pitting potential
To evaluate the pitting susceptibility of an alloy, the electrochemically determined pitting
potential pit is often used (Fig. 11.21). Above the pitting potential, the measured dissolution
current densities are very high and the sample displays pits. The pitting potential is dependent
on the chloride concentration and the alloy composition (Fig. 11.22, Tab. 11.3).
Prof. B. Elsener, Institute for Building Materials, ETH Zürich
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Fig. 11.21: Current density/potential curve of metal that can be passivated in the absence and in the presence of
pitting inducing anions. For explanation see text
Table 11.3: Pitting potentials for different metals and alloys
Metal
Aluminum
Nickel
18/8 CrNi steel
12% Cr steel
30% Cr steel
Titanium
eL (V NWE, 0.1 N NaCl
-0.37
0.28
0.26
0.20
0.62
> 1.0 (1 N NaCl)
Figure 11.22: Pitting potentials of different metals depending on the Cl- concentrations
Higher chloride concentrations (more aggressive conditions) shift the pitting potential to more
negative values (Fig. 11.22).
Prof. B. Elsener, Institute for Building Materials, ETH Zürich
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Surfaces and interfaces
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For high-alloyed steel, the resistance against pitting increases with higher chromium and
molybdenum content. Figure 11.23 shows an example of the pitting potential changing with
the alloy composition. Three steel types with different compositions (Tab. 11.2) were
potentiodynamically investigated in 1N HCl. Figure 11.23 shows that steel with the highest
Ni and Mo content possesses the highest pitting potential and is the easiest to passivate.
10 0
1.4439, 1.4529
i
10-4
[mA/cm 2 ]
-2
10
i
[mA/cm 2 ]
1.4301
1.4401
10
2
10
4
10
0
10
-2
1.4301
1.4439
1.4529
0.1 M NaCl
6.0 M LiCl
-6
-4
10
-0.4
-0.2
0
0.2
0.4
0.6
0.8
10
1
-0.6
Potential [VSCE]
-0.4
-0.2
0
0.2
0.4
Potential [VSCE]
0.6
0.8
1
702.1
Figure 11.23 Current density/potential curves of different high-alloyed steel in 1N HCl
When does pitting occur?
The knowledge of the pitting potential and its influencing variables (chloride concentrations,
alloy composition) alone does not allow one to establish whether pitting occurs in a given
situation or not. Similar to the passivation process, the occurrence of pitting is also a system
behavior, i.e. the metal as well as the environment influence this system. The condition for
the occurrence of pitting is
corr > pit
meaning that pitting occurs when the corrosion potential of the originally passive metal
exceeds the pitting potential (for the respective chloride concentration). To be sure that pitting
will not occur, alloys with a high pitting potential have to be used (Fig. 11.22, Tab. 11.3)
and/or the corrosion potential has to be as negative as possible. For high-alloyed steel, the
corrosion potential is a function of the pH-value and, for passive surfaces, practically
independent of the alloy composition (Fig. 11.24).
The resistance of stainless steels against pitting and crevice corrosion is positively influenced
by the elements chromium (Cr), molybdenum (Mo), and nitrogen (N) (Tab. 1.8), and
negatively influenced by the elements sulfur (S), manganese (Mn, together with sulfur), and
carbon (C). The elements that have a positive effect are summarized in the PRE (Pitting
Resistance Equivalent). To calculate the PRE, the following equation is used:
PRE = a %Cr + b %Mo + c %N.
For the constants, a = 1, b = 3.3 und c = 0 – 30 are used. The PRE has no strict scientific
meaning. Nevertheless, it is a suitable guideline to a comparative evaluation of the resistance
of solution annealed / quenched stainless steel (and nickel alloys). Impurities, inclusions,
degree of deformation, previous heat treatments and the surface condition of the steel are not
considered in the PRE.
Prof. B. Elsener, Institute for Building Materials, ETH Zürich
268
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Regarding their resistance to pitting and crevice corrosion, stainless steels can be roughly
divided into groups (Tab. 11.4).
Table 11.4
Some steel and nickel qualities sorted in groups with increasing pitting and crevice corrosion resistance;
qualities in brackets: can be present in different groups
group
I
II
III
IV
V
VI
materials
1.4301, 1.4303, 1.4306, 1.4541, 1.4543
1.4401, 1.4404, 1.4406, 1.4429, 1.4435, 1.4436, 1.4438, 1.4571
1.4439, 1.4462, 1.4503, (1.4539)
(1.4539), 1.4563
1.4529, 1.4565, Avesta 254 SMO, (2.4856)
2.4602, 2.4610, 2.4819, (2.4856)
This division into groups of provides only a general rule for corrosion resistance. In specific
cases, the material impurities (Mn and S content) and especially the surface quality (ground,
brushed, polished, honed) play an important role. Lower classified steel with better surface
quality behaves in a more resistant way than higher classified steel with bad surface (rough,
grooves, scratches).
Total
O-
OH-
S-
Figure 11.25: ToF-SIMS images of a 1.4305 CrNi steel surface with 0.25% sulfur. The MnS dispersion is
associated with less OH-, i.e. the passive film is less resistant. A Rossi, B. Elsener, G. Hähner, M. Textor and
N.D. Spencer, Surf. Int. Anal. 29 (2000) 460
Prof. B. Elsener, Institute for Building Materials, ETH Zürich
269
Surfaces and interfaces
Corrosion of metals
Why does pitting occur?
In the literature, different models are proposed to explain the start of pitting (nucleation). It is
generally assumed that chemical or structural inhomogeneities, weak spots, in the passive
film are the starting points. With the high lateral resolution that became possible in recent
years in micro-electrochemical as well as in surface-analytical measurement techniques, the
model of MnS inclusions has been confirmed to a large extent.
ToF-SIMS investigations on sulfur containing high-alloyed steel (DIN 1.4305) showed that
the film on combined MnS/oxide inclusion has less OH- above the MnS particles than above
the undisturbed passive film. This can be interpreted as a weak spot of the passive film and
thus as a nucleus for pit formation. Latest analyses show that chromium depletion occurs at
the border of dispersions, which then leads to a microscopic thin area, where the passive film
is no longer corrosion resistant.
11.4.3 Intercrystalline corrosion
Intercrystalline corrosion, the intergranular attack accompanied by a drastic loss of strength,
can occur for Al-Cu alloys and high-alloyed steel due to an erroneous heat treatment
(sensitization). For austenitic CrNi steels the critical temperature range is between 500 and
700 °C, for temperatures around 750°C a few minutes, for 500°C a few hours are enough for
sensitization. In practice, slow cooling rates or welding operations are the causes of
sensitization. Fast cooling prevents intercrystalline corrosion.
During sensitization, chromium carbides form at the grain boundaries (energetically favorable
areas). The necessary chromium diffuses from the bulk to the grain boundaries and chromium
depletion occurs. The chromium-depleted areas cannot be passivated anymore under the same
environmental conditions (pH) and start to corrode (Fig. 11.26).
% Cr
90
18
12
Figure 11.26 Intercrystalline corrosion (micro photography left) through chromium carbide formation at the
grain boundaries (middle scheme) leads to chromium depletion at the grain boundaries (right).
Intercrystalline corrosion on high-alloyed steel can be avoided when steel with low carbon
content or with additional titanium in the alloy is used. Low carbon content makes carbide
formation impossible, while alloying with titanium (Ti = 5 * C) leads to formation of (non
interfering) titanium carbides, and the chromium content remains unchanged.
Prof. B. Elsener, Institute for Building Materials, ETH Zürich
270
Surfaces and interfaces
Corrosion of metals
11.5 Nickel-phosphorus coatings
Nickel-phosphorus alloys, used in the past chiefly as corrosion protective coatings, constitute
the earliest industrial application of X-ray amorphous or nano-crystalline metals, dating back
to 1946. Ni-P alloys with ca. 20 at. % P (close to the eutectic composition) exhibit distinctly
better corrosion resistance than does pure Ni, showing a suppression of anodic dissolution in
the potential range where pure nickel dissolves actively in acids. This holds for melt spun,
electroless or electro-deposited Ni-P alloys. It is also generally accepted that only X-ray
amorphous or nano-crystalline alloys – irrespective of mode of production – exhibit this
superior corrosion resistance whereas crystalline Ni-P alloys with low P content are
susceptible to corrosion. Today there is a growing interest in the production of ternary Ni-P
alloys and for co-deposition with diamond or PTFE particles in order to produce tailor-made
functionalized surfaces. (The reason for the high corrosion resistance of these Ni-P alloys
containing more than 18% P was the topic of a recent bachelor thesis and research at IfB and
LSST combining electrochemical and surface-analytical techniques.)
11.5.1 Coating characterization
Ni-P coatings were produced on mild steel substrates in a commercial autocatalytic nickel
hypophosphite bath (Ronamax SR) by electroless deposition with a thickness of 10, 15 and
20 m. The pH of the bath was 4.6 – 5, temperature 85 – 90 °C, plating time ca. 2 h. The
surface morphology shown in Figure 11.27 shows the nodular growth of NiP. The structure of
the deposits was found to be x-ray amorphous. Chemical composition determined by EDX
showed phosphorus content of 18.4 – 19 at%—close to the eutectic composition.
Ibeob. (counts)
15000
10000
5000
0
20
40
60
Position (°2Theta)
80
Figure 11.27: AFM characterization of the surface of electroless Ni19P alloy and XRD measurement
showing the broad peak of nickel indicating an x-ray amorphous structure
11.5.2 Electrochemical and corrosion behaviour
The corrosion behaviour was studied in acidic (0.1 M H2SO4) and neutral (0.1 N NaCl)
solutions with potentiodynamic polarization curves comparing pure nickel with the Ni19P
alloy (Figure 11.28). Pure nickel shows the typical active/passive transition and a broad
passive range whereas the NiP alloy shows low current densities and a current plateau in the
range where nickel is actively corroding. From figure 11.28 it can thus already be concluded
that the high corrosion resistance of the Ni-P alloys is not due to a normal, oxide-type
passivity.
Prof. B. Elsener, Institute for Building Materials, ETH Zürich
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Surfaces and interfaces
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Ni und Ni-P in 0.1M H2SO4
1.00E+00
1.00E-01
i [A/cm2]
1.00E-02
1.00E-03
Ni
Ni-P 20.7
Ni-P 14.8
Ni-P 9.64
1.00E-04
1.00E-05
1.00E-06
1.00E-07
1.00E-08
-0.4
-0.2
0
0.2
0.4
0.6
0.8
1
1.2
1.4
E [V]
Figure 11.28: Current density/potential curves recorded in 0.1 M H2SO4 of pure nickel compared to
Ni19P alloy with coating thicknesses ranging from 10 to 20 μm.
Figure 11.29: Corrosion rate (in μA/cm2) determined from polarization resistance measurements of
the Ni19P alloy (three different coating thickness) in aerated 0.1 N NaCl over time (7 days).
The corrosion rate in chloride-containing (0.1 N NaCl) solution (determined from
polarization resistance measurements) is low and remains constant with time of immersion
(Figure 11.29). Potentiostatic passivation experiments were performed in the range of current
arrest for the NiP alloy (-0.1 V SCE) and in the passive range for pure nickel (figure 11.30).
The results show that nickel shows a -1 slope in the log i vs log t diagram, indicating the
formation of an oxide film similar to that of stainless steels (see figure 11.14). The Ni19P
alloy instead shows a slope of -0.5, indicating a diffusion-controlled process of alloy
dissolution.
Prof. B. Elsener, Institute for Building Materials, ETH Zürich
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Surfaces and interfaces
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Ni und NiP in 0.1M NaCl , polarisiert
1.00E-02
9.64um
1.00E-03
14.8um
i [A/cm2]
20.7um
1.00E-04
Ni
-0.5
Potenziell (14.8um)
Potenziell (Ni)
1.00E-05
1.00E-06
-1
1.00E-07
10
100
1000
10000
t [s]
Figure 11.30: Potentiostatic passivation of pure nickel and nickel-phosphorus alloy
11.5.3 XPS surface analysis
In the survey spectra of sputter-cleaned Ni-P alloy (Figure 11.31) only the signals from
photoelectron and X-ray induced Auger peaks of nickel and phosphorus were found. On the
as-received samples carbon C1s and oxygen O1s signals were also detected. In the following,
the detailed spectra of phosphorus (P2p and PKLL) and nickel Ni2p of a Ni-P sample
polarized for 1 h in 0.1 M Na2SO4 solution are shown as an example.
Figure 11.31: Survey spectrum of sputter-cleaned Ni19P alloy. AlK 300 W
Phosphorus: The detailed spectra of phosphorus P2p (Fig. 11.32) and corresponding X-ray
induced PKLL Auger peak (not shown) show three different peaks. The most intense P2p
signal appears at a binding energy of 129.7 eV and the corresponding PKLL at a kinetic
energy of 1858.3 eV. The modified Auger parameter ’ calculated as BE (2p) + KE(PKLL) is
equal to 1988.1 eV. The phosphorus signal at the binding energy of 133.6 eV and
corresponding PKLL signal at 1851.0 eV yields an Auger parameter of 1984.6 eV and can be
Prof. B. Elsener, Institute for Building Materials, ETH Zürich
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Surfaces and interfaces
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Intensity (a.u.)
assigned to phosphates. The third (intermediate) phosphorus signal shows a binding energy of
132.0 eV (Figure 11.32) and a kinetic energy of 1855.4 eV. The Auger parameter ’ is 1987.3
eV.
126
128
130
132
134
136
138
Binding Energy (eV)
Figure 11.32: High-resolution spectrum of phosphorus P2p. Ni19P alloy polarized for 1 h at -0.1 V
SCE in 0.1 N NaCl. AlK 300 W
Nickel: The detailed spectrum of nickel Ni2p3/2 is very similar to the typical Ni2p3/2 plus
satellite structure as known from pure metallic nickel. The binding energy of the Ni2p3/2
signal is found at 853.0 ± 0.1 eV. The distance of the satellite is found here to be 7.2 eV, thus
clearly higher than the 6 eV reported for pure nickel but in keeping with the concept of the
influence of phosphorus content on the electronic properties of Ni-P alloys. A very small
amount of nickel phosphate (856.4 ± 0.1 eV) is also found. No signals attributable to nickel
oxide have been detected.
Mechanism of corrosion resistance: Ni-P alloys in this work show a slope of -0.5 in the
logi/logt diagram (Figure 11.30). The electrochemical results can be interpreted in terms of
diffusion of a faster dissolving component of the alloy through the developing surface layer
enriched in the less soluble component. In the present case nickel dissolves preferentially
from the amorphous Ni-P alloys and the current decay curve describe the dissolution of nickel
limited by diffusion through the P enriched surface layer. The chemical state of this
(intermediate) enriched phosphorus species has been found with the Wagner chemical state
plot to be similar to elemental phosphorus, thus its concentration is so high as to show no
chemical bonds to nickel atoms. The outstanding corrosion resistance of electroless-deposited
Ni-19P alloys can be explained by the formation of a phosphorus-enriched layer at the alloy
surface. This layer accounts for the diffusion-controlled dissolution of the alloy. XPS/XAES
surface analysis data interpreted on the basis of the Auger parameter concept and the
chemical-state plot clearly show that phosphorus in this layer has a chemical state close to
that of elemental phosphorus.
Prof. B. Elsener, Institute for Building Materials, ETH Zürich
274
Surfaces and interfaces
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11. 6 Organic coatings
Metals coated with plastics (coating or paint) can only corrode at the pores of the coating.
Due to the high electric resistance of the coating, DC measurements do not yield a satisfying
result, and the properties of the coating cannot be detected. Impedance measurements allow
the properties and the changes of the layer to be detected, as well as the extent of the
corrosion of the metal below – and everything at the corrosion potential with a minimal
disturbance of the system.
Figure 11.33 Impedance measurements of the silicone polyester on steel system in 0.5 N NaCl,
deaerated. 1 at the beginning, 2 after 43 h. Thickness of the layer: 25 m
For the silicon polyester on steel system, impedance measurements at the corrosion potential
in 0.5 N NaCl, deaerated, were carried out at different times (Fig. 11.33). At the beginning of
the experiment (curve 1) a purely capacitive behavior of the layer is observed, while after 43
hours the resistance of the layer was already strongly decreased (pore formation) and
corrosion reactions have started. A detailed analysis of the system results in:
-
Curve 1, at the beginning of the experiment
purely capacitive behavior, slope dlogZ/dlog = -1.
From the value Z(log = 0) the capacity C of the coating can be calculated
C = 1/ Z(log = 0) = 0.28 nF/cm2
This very small value is typical for organic coatings, the value of the dielectric
permittivity is calculated from that to about 5.
-
Curve 2, after 43 h in solution
In the high frequency part, the impedance coincides with curve 1, i.e. the capacity of the
layer has not changed. In contrast, the resistance (local) strongly decreased, a value of R
= 290 kcm2 was measured. The impedance increases for low frequencies
Prof. B. Elsener, Institute for Building Materials, ETH Zürich
275
Surfaces and interfaces
Corrosion of metals
(simultaneous with a second maximum of the phase angle), which suggests that
corrosion processes occur in the pores (possibly diffusion controlled).
A visual observation of the probe shows that a pore has indeed formed on the surface. From
the diameter of the corroded area (F 0.0009 cm2) and the polarization resistance Rp of
approx. 1.0 M, a value for Rp of about 900 cm2 can be calculated, which is well in
accordance to the values for steel corrosion in deaerated NaCl.
Prof. B. Elsener, Institute for Building Materials, ETH Zürich
276
Surfaces and interfaces
Corrosion of metals
Literature
H. Kaesche, Die Korrosion der Metalle, Springer Verlag Berlin
H. Uhlig, Corrosion and Corrosion Control – an Introduction to Corrosion Science and
Engineering, John Wiley & Sons
Corrosion and Environmental Degradation, ed. M. Schütze, Wiley VCH (2000) in the Series
Materials Science and Technology, Two Volumes
L. Bertolini, B. Elsener, R. Polder, P. Pedeferri, Corrosion of Steel in Concrete, Wiley VCH
(2004)
Internet www.corrosion-doctors.org
Bernhard Elsener,
Institute for Building Materials, ETH Zürich
HIF E15, ETH Hönggerberg, 8093 Zürich
044 / 633 2791
elsener@ethz.ch
Professor for Materials Science, University of Cagliari,
belsener@unica.it, http//:dipcia.unica.it/superf/
Prof. B. Elsener, Institute for Building Materials, ETH Zürich
277