Fundamentals of Electrochemistry, Corrosion and Corrosion Protection Chapter 13 - Christian D. Fernández-Solis, Ashokanand Vimalanandan.pdf

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Chapter 2
Fundamentals of Electrochemistry,
Corrosion and Corrosion Protection
Christian D. Fernández-Solis, Ashokanand Vimalanandan,
Abdulrahman Altin, Jesus S. Mondragón-Ochoa, Katharina Kreth,
Patrick Keil and Andreas Erbe
Abstract
This chapter introduces the basics of electrochemistry, with a focus on
electron transfer reactions. We will show that the electrode potential formed when a
metal is immersed in a solution is most of the time not an equilibrium potential, but
a mixed potential in a stationary state. This mixed potential formation is the basis of
corrosion of metals in aqueous solutions. Organic coatings are introduced as pro-
tecting agents, and several types of coatings are discussed: classical passive coat-
ings, and active coatings as modern developments. Three electrochemical
techniques, which are commonly used to asses the protecting properties of coatings,
are shortly introduced as well: linear polarisation measurements, electrochemical
impedance spectroscopy, and scanning Kelvin probe measurements.
2.1
Basics of Electrochemistry
Electron transfer reactions are wide-spread in nature, e.g. in the respiratory chain,
they are important technologically, e.g. in electrolysers and for metal plating, and
they contribute to the degradation of materials, in corrosion processes of metals. This
chapter shall serve as an introductory text into the basic concepts, with a special
focus on their importance in the
eld
of corrosion science. Electrochemistry is quite
an old science, and hence a number of good textbooks are available for more detailed
introductions of the fundamental concepts [1–3]. It is worth pointing out that there
are practically important corrosion mechanisms which are not at all based on elec-
trochemical reactions. Details are available in dedicated textbooks [4,
5].
C.D. Fernández-Solis
Á
A. Vimalanandan
Á
A. Altin
Á
J.S. Mondragón-Ochoa
Á
A. Erbe (&)
Max-Planck-Institut für Eisenforschung GmbH, Max-Planck-Str. 1,
40237 Düsseldorf, Germany
e-mail: a.erbe@mpie.de
K. Kreth
Á
P. Keil
BASF Coatings GmbH, 48165 Münster, Germany
©
Springer International Publishing Switzerland 2016
P.R. Lang and Y. Liu (eds.),
Soft Matter at Aqueous Interfaces,
Lecture Notes in Physics 917, DOI 10.1007/978-3-319-24502-7_2
29
30
C.D. Fernández-Solis et al.
2.1.1
Electrostatic Potentials at Interfaces
Interfaces, in particular aqueous interfaces, are almost always charged, e.g. by
dissociation of surface groups, or adsorption of ions from solution. On conductive
substrates, the charge state of an interface may be actively controlled, as will be
discussed below. Around charged interfaces, a static electric
eld
is present, which
affects the distribution of ions in solution around this interface. The rather complex
multibody-interactions between solvent, mobile charges on both sides of the phase
boundary, and stationary atoms leads to a complicated interfacial structure, which is
schematically shown in Fig.
2.1
[6]. This interface is often referred to by the
misleading term
“double
layer” (see also Chap.
4
by G. H. Findenegg).
Closest to the interface is a structured layer containing adsorbed molecules of the
solvent and other adsorbed species. The region where the electric charges of the
absorbed ions are allocated is called inner Helmholtz layer; this region is at a
distance
d
1
(Fig.
2.1).
Solvated ions can approach the interface up to a distance
d
2
.
The region where the electric charges of the solvated ions are located is called outer
Helmholtz layer. Due to thermal motion in the system, theses solvated ions are
distributed in a three dimensional region ranging from the outer Helmholtz layer to
the bulk. This region is often referred to as
“diffuse
layer”. Counter-ion distribution
Fig. 2.1
Schematic
representation of the layers at
a solid/liquid interface.
1
Inner Helmholtz layer,
2
Outer Helmholtz layer,
3
Solvated ions (cations),
4
Diffuse layer,
5
Electrolyte
solvent,
6
Specically
adsorbed ions. Drawing
inspired by [7]
1
2
3
6
4
5
d2 d1 0
2 Fundamentals of Electrochemistry, Corrosion
…
31
in the diffuse layer is important for some processes at a variety of interfaces (e.g.,
solid electrodes, biomolecules, etc.) and in technological applications as well (e.g.,
corrosion, paints, etc.). A simple quantitative approach to the description of the
diffuse layer, balancing entropy against mean-eld electrostatic attraction, is the
Poisson-Boltzmann equation, which is extensively discussed—including its limi-
tations—in [8]. Because deviations from the classic picture are found experimen-
tally (e.g. [9]), modern conceptual works discuss in more detail solvation effects,
uctuations
and ion correlation [10–14].
2.1.2
Electrochemical Potential
As an interface is charged, the work needed for a distribution of charges needs to be
considered when analysing the total free enthalpy or free energy of a system. For
this purpose, in charged systems, the electrochemical potential takes the role of the
chemical potential.
The electrochemical potential
l
i
is dened as the mechanical work required to
bring 1 mol of ions with valency
z
and hence charge
ze
from a standard state to a
specied electrical potential and concentration. Thermodynamically, it is a measure
of the chemical potential that takes into account the electrostatic contributions; it is
expressed as energy per mole,
l
i
¼
l
i
þ
z
i
FU
ð2:1Þ
where
l
i
is the chemical potential of species i without considering the charge,
F
is
the Faraday constant and
U
the local electrostatic potential. The
rst
term of Eq.
2.1
takes into account the chemical potential of the species i (see Chap.
8
by R. Sigel
for a thorough discussion on the chemical potential in relation to the formation of
mixtures), while the second term includes the free enthalpy change brought by
altering the potential
U
of the phase in which the charged species is located.
U
is
also referred to as the inner potential or Galvani potential of the phase. This
approach is analogous to the addition of the interface term to the chemical potential,
described in detail in Chap.
4
by G. H. Findenegg.
Bringing two phases
a
and
b
with mobile charges into contact (consider metal/
salt solution as an example) will result in an equilibrium where the electrochemical
potentials of the two phases will equalise in a similar fashion as the chemical
potentials for uncharged systems,
l
a
¼
l
b
ð2:2Þ
For simplicity, let’s consider a single-valent system, for which we obtain, using
Eq.
2.1,
32
C.D. Fernández-Solis et al.
l
b
À
l
a
:
U
À
U
¼
F
a
b
ð2:3Þ
The result is the built-up of an electrostatic potential across the interface. The
electrostatic potential difference between the two phases is the electrode potential,
E
cell
¼
U
a
À
U
b
:
ð2:4Þ
We need a certain reference to quantify the potential, which is an arbitrarily
chosen system, for which different choices exist in the literature [15]. In analogy to
the standard term in the chemical potential, to which the concentration dependence
is added, we can
“dump”
the standard terms into a standard electrode potential
E
0
for each redox system. These potentials are published in the literature typically as
potential differences against the standard hydrogen electrode. They can be found
e.g. in the Handbook of Chemistry and Physics [16]. The potentials are usually
reported as reduction potentials, and represent the tendency of a certain species to
“obtain”
one or more additional electrons. The higher the potential, the higher is
this tendency. The standard electrode potential is therefore a quantity similar to the
free enthalpy of formation of a certain species.
The electrode potential
E
cell
is actually the potential of a single electrode. When
putting two of these half-cells into electrical contact, electrons can
ow
through an
external circuit and balance the electrode potential differences which are present. At
the same time, a chemical transformation occurs: one species is reduced and one is
oxidised. In one of the half cells, there will thus occur oxidation, i.e. electron
“loss”
of
the atoms or molecules, and on the other side, there will be reduction, i.e.
“gain”
in
electrons for a certain species. As other chemical reactions generate heat or light,
redox reaction can be used to generate an electrical current, and likewise, electrical
current can be used to
“drive”
redox reactions in a certain direction. The half cell, in
which oxidation occurs, is called the
“anode”,
and the half-cell, in which reduction
occurs is the
“cathode”.
Note that these denitions rely on the nature of the reaction
that is occurring, not on the direction of the current
ow
or the charge. The difference
in standard free enthalpy is related to the difference in standard electrode potentials to
DG
¼ ÀzFE
cell
:
ð2:5Þ
If
E
cell
[
0, the process is spontaneous, as e.g. in galvanic cells, batteries or
corrosion of metals. On the other hand, if
E
cell
\0,
the reverse reaction is sponta-
neous, as e.g. in electrolytic cells.
Moving away from the standard standard state, Eq.
2.5
keeps its validity. So, at
each concentration of involved species,
DG
¼ ÀzFE
cell
:
ð2:6Þ
2 Fundamentals of Electrochemistry, Corrosion
…
33
For a redox reaction of the type
aOx
þ
ze
À
!
cRed
ð2:7Þ
the Nernst equation determines the electrode potential for an individual half-cell as
RT
½Red
c
ln
;
E
¼
E
À
nF
½Ox
a
0
ð2:8Þ
where the activities may be approximated as concentrations in dilute solution. Here
we dropped the subscript
cell
for convenience. Equipped with the Nernst equation,
we can determine electrode potential differences, and hence free enthalpy differ-
ences from equilibrium, from tabulated standard electrode potentials, knowing
activities/concentrations of dissolved species [1].
2.1.3
Currents Are a Measure of Reaction Rates
Electric current is the
ow
rate of electric charge
q
through a system, where
t
denotes the time,
I
¼
dq
dt
ð2:9Þ
In interface science, normalising the current in Eq.
2.9
by the interface area
A
is
in general convenient, introducing current density
i
¼
I=A,
i.e. current per unit area
of electrode. The current through an interface is thus a convenient measure of the
rate of electron transfer. If there is only a single electron transfer going on, the
current density is directly related to the rate of the chemical reaction. It is possible
to relate to the current through Faraday’s law
I
¼
dq
nzF
¼
dt
t
ð2:10Þ
to the total amount
n
of transformed substance, e.g. in a reaction like in Eq.
2.7.
The
magnitude of the current
owing
at any potential depends on the kinetics of electron
transfer. (Alternatively, in practise it often depends on the kinetics of transport to
the interface. This case, is, however, not of interest in the understanding of the
mechanistic aspects of reactions, which is why it is not considered here.)
At any electrode potential, the measured current density is given as the sum of an
anodic and a cathodic partial current
i
¼
i
a
þ
i
c
:
ð2:11Þ

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