CHEM 12032

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CHEM 12032
ELECTROCHEMISTRY
Part ii
Dr.K.A.S Pathiratne Head of the department of
Chemistry University of Kelaniya Sri Lanka
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CHARGE DISTRYBUTION AT THE ELECTRODE /
ELECTROLYTE INTERFACE
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Even when an electrode not connected to an
external power supply it can have either ()ve or
(-)ve excess charges as shown above on its
surface. The oppositely charged counter ions from
the solution gather towards the surface of the
electrode due to electrostatic attractions. The
region of the solution in between the surface of
the electrode and the surface of the solution
next to the electrode is called the INTERFACE.
The interface also includes the two surfaces.
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The oxidation and reduction reactions all occurs
on the electrode surface. Electrons exist on the
electrode surface. They are donated to ions
coming from the solutions towards the electrode
or species present on the electrode during
reduction reactions. Electrons are donated to
electrodes at the electrode surface during
oxidation by ions coming from the solution or
species present on the electrode during oxidation
reaction.
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Oxidation process at a metal electrode
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Oxidation process at a redox electrode
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This charge transfer process is controlled by the
charge distribution at the interface
Helmoltz -Perrin Theory
The parallel plate condenses model The charge on
the electrode is present on the two dimensional
surface of the electrode. The counter ions
gathered from the solution also are found very
near to electrode surface on a 2D surface. Note
The electrode was considered to have a planar
surface here. If it is a spherical the counter
ions will also be found on a spherical shell.
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A planer electrode is considered for the present
discussion The charge distribution at the
interface resembles a parallel plates capacitor
found in electrical circuits.
d interface
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Note The Interface is very thin. d is very
small and is of atomic dimension (order of an
angstrom).
The properties of circuitry element-Capacitor
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Q a V
Q CV C Q/V
charge on either of the plate per unit
potential difference (called capacitance).
Further, it can be shown that, Where,
? permittivity of the space between the plate.
?0 permittivity of free space
d the distance between the two metal plates.
C ? ?0 / d
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This relationship shows that if a charge
distribution similar to what is present on
parallel plates capacitors exist at the
electrode/electrolyte interface, the capacitance
of it should be a constant that is independent on
the potential applied to the electrode. Experiment
shows that the capacitance at the
electrode/electrolyte interface does depends on
the potential applied to the electrode. (see fig
7.68)
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e.g.
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The above results indicate that the charge
distribution at the interface is different from
what is described by the Helmoltz-Perrin model .
The Gouy-Chapman Diffuse-charge model for
electrode/ electrolyte surface
The excess charge on the electrode surface are
distributed on a 2 dimensional plane as described
by Helmoltz Perin model The counter ions in the
solution side however are not limited to a 2
dimensional plane, due to the random thermal
movement of ions.
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The ions in solution possess thermal energy. As a
result they move within the solution volume
randomly. These movement increases the disorder
in their distribution. The ions at the same time
are attracted electrostatically by the oppositely
charged electrode. These electrostatic attraction
between the electrode and the ions forces the
ions to be present on a more orderly arrangement.
Under the existing situation each an every ion is
subjected to two opposing forces. One that forces
them to be in a more organized arrangement (e.g.
distribution on a 2D plane near the electrode
surface)
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The other that forces them to be distributed
randomly in a 3D Region. Under the influence of
these two opposing forces, the ions reaches to an
equilibrium resulting a distribution in a 3
dimensional region known as diffuse layer
(region) of few angstroms thick. The density of
the counter ions in this diffusion layer
gradually decreases with increasing the distances
away for metal the electrode. The thickness of
the layer increases with decreasing concentrates
of ions and decreasing the potential applied to
the electrode.
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If is very important to note that the density of
the diffused layer however decreases with
increasing its thickness under this
circumstances. On the other hand, when the
potential of the electrode is increased and the
concentration of the solution increased the
thickness of the diffuse layer, reduces. The
decrease in thickness increases its density.
Using basic principles of chemistry and physics
and statistics, Gouy Chapman derived
mathematical expressions, for (a) the charge
density on the metal surface or the diffuse layer
(b) capacitance at the interface.
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The total charge, QM on the electrode, equals in
magnitude and opposite in the sign to the total
charge, QS in the diffuse layer.
i.e. QM (-) QS Further, the charge
density, sM, on the metal surface equals in
magnitude and opposite in sign to the charge
density, sS, in the diffuse layer.
i.e. sM -sS
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Also, the charge density in the solution side sS
decreases exponentially with the distance away
from the electrode towards the bulk of the
solution. See the figure 7.66.
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Further, the potential within the defuse layer
decreases exponentially with the distance from
the electrode as given by the figure,
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The potential ?x at any distance x, away from the
electrode according to the theory can be given by
the equation. ?x ?0
exp(-X.x) Where, X inverse of the Debye-Huckel
length i.e. X-1 the Debye Huckel reciprocal
length and it has dimension of length
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Using these equations they calculated the values
of charges and capacitances, for certain
electrode electrolyte interfaces under several
different electrode potentials electrolyte
concentrations. The capacitance according the
theory of Gouy-Chapman varies (increases)
asymptotically with increasing potential beyond
the potential of zero charges (pzc) to either
directions. The capacitance also increases with
the concentration. See the figure 12.3.5
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Gouy - Chapman Stern Model for the Electrified
Interface
The G.C. model was partly successful. It
predicted the variation of interfacial
capacitance with variation of potential applied
to the electrode and the concentration of the
electrolyte.
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However, variation of capacitance according to
the G-C model deviated at higher potentials from
that observed experimentally. An improvement to
the G-C model was proposed by Stern. The
resultant model was named Gouy -Chapman-Stern
model. According to the Stern modification ions
which are solvated in solution can approach the
electrode surface only up to a certain distance
of closest approach and that distance was labeled
a.
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All the ions approached the distance of a from
the electrode constitute a plane. This plane
passing through the centers of all these ions is
called Helmholtz Plane and no ions are present
between the metal surface and the Helmholtz
Plane. The potential decreases linearly from the
electrode towards the Helmholtz Plane from the
Helmholtz Plane towards bulk of the solution
solvated ions are present in a 3D region with
exponentially decaying charge density and
potential. See the figure 7.69
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Double Layer
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Next picture shows the formation of the fixed
Helmholtz layer and the mobile diffused layer in
animation
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A consequence of the stern modification An
electrified interface is equivalent to two
capacitors in series. Two types of potential
drops exist at the interface. One that is
linearly decreasing from the metal surface to the
Helmholtz Plane which is denoted by fM here and
the other that decreases exponentially from the
Helmholtz Plane towards the bulk to the zero
potential denoted by fG here.
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The charge separation between electrode surface
and the Helmholtz plane forms a capacitance
called Helmholtz capacitance and label here as
CH. The Helmholtz plane does not have enough
counter ions to balance all the charges present
on the electrode surface. The counter ions
aggregated in the defuse layer in a 3D region
balance the remaining charges on the metal plate.
The charges in the 3D region and the remaining
charges equivalent to be present at the Helmholtz
plane constitute a second capacitor known as
Gouy-Chapman capacitor.
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This capacitance here is indicated by CG. The
total capacitance represented by Cd is the recent
of two capacitance 1 connected in series and is
given by the expression.
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(a) A view of the differential capacitance in the
GCS model as a series network of Helmholtz layer
and diffuse layer capacitance. (b) Potential
profile through the solution side of the double
layer, according to GCS theory. Calculate from
(12.2.23) for 10-2 M 11 electrolyte in water 250C
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Further important to the GCS model specially
adsorbed ions on the surface of the
electrode Variation of experimentally observed
capacitance with variation of electrolyte
concentration was found to be different from what
predicted by GCS theory. GCS model was further
refined to accommodate the experimental
observation.
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Further Improvements
The electrode surface was considered to consist
of either specifically absorbed ions which are
partially desolvated or absorbed solvent
molecules with these dipole oriented to suit the
charge on the metal surface. The two types of
arrangements specifically adsorbed partially
desolvated ions called Type I arrangement and
adsorbed dipole oriented molecules Type o
arrangement are represented in the figure 7.72
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The distribution of ions at the metal/solution
interface and the variation of potential
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Corrosion of Metals
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Corrosion of metals Corrosion is an
electrochemical process. During corrosion a metal
undergoes oxidation. This process can be
represented by the following general equation.
M(s)
Mn(aq) ne The metal ions can undergo further
chemical reaction leading to a continuous loss of
metal. Corrosion is not limited to iron.
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Corrosion is a natural and though slow is a
spontaneous process. There are many different
types of corrosion process occurring under
different types of environmental conditions. In
all these corrosions another very important
process occurs near the corroding object or on
its surface. It is a reduction process where,
electrons produced during corrosion are consumed
by this cathodic or reduction reactions. This
reduction reaction (s) support the corrosion or
oxidation of the metal. Four such common
reduction reactions that occur in national
(aqueous) environment are given below.
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Wagner- Traud Model (Mechanism) for Corrosion
Process
A corroding system is equivalent (represents)
a short circuited electrochemical cell. Just as
in an electrochemical cell, in a corroding system
following 4 components can be seen 1. An
anode, (i.e. the corroding metal) 2. A
cathode (The site where the reduction reaction
occurs) 3. A medium for movement
of the oxidized and reduced species.
4. A medium for flow of electrons.
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Consider the corrosion of a piece of Zn metal
immersed in an HCl solution (1 M).
Zn in 1M HCl will be corroded. Zn undergoes
oxidation as follows. Zn(s)
Zn2(aq) 2e
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H ions in the solution grab electrons released
from the Zn metal on the surface of the Zn metal
and is converted to it atoms producing H2 gas
from the Zinc surface can be seen in this case.
The reduction reaction is 2H(aq) 2e
H2(g)
This is the cathodic reaction that occurs during
the reaction. The surface of the metal is the
cathodic site.
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The electrons released from zinc atoms are
grabbed by H ions on the surface itself and
hence not available for circulation through an
external conductor. The solution itself supports
the movement of ions. The net reaction of the
corrosion is given below. Zn(s) 2H(aq)
It is possible to make the very same reaction
(reaction (4)) to occur in a different
environment as given below.
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The net reaction that occurs in the corrosion of
Zn in HCl occurs during the discharge of the
above cell with one very important difference.
i.e. The electron flow occurs through the
external wire from the Zinc electrode to Pt wire
of the H2 electrode. This shows that a corroding
cell is a short circuited cell of a conventional
electrochemical cell where the same reaction
occurs during the cell discharge.
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Thermodynamics of Corrosion
Corrosion is a natural, in other words a
spontaneous process. Therefore ?G corrosponding
to net corrosion reaction should be negative.
i.e. ?Gcorr lt 0
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Also, corroding system is an electrochemical
cell. The change in free energy ?G during the
cell discharge reaction can be proved to be
equaled to ?G -nFEcell Where, Ecell
the e.m.f. of the corroding short
circuited cell (n, and F have their
usual meanings)
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If corrosion occurs ?Gcorr lt 0 Hence according
to the above relationship
nFEcell gt 0 i.e. Fcell gt 0
Ecell Ecathode- EAnode For Ecell gt 0,
Ecathode gt EAnode
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In situation where, the reduction reaction
feasible in known this criterion can be used to
predict whether a particular metal is stable or
not against corrosion in the environmental
concerned. The previous example can be examined
to predict if a pure Zn vessel at 250C can be
used to store 1M HCl acid which has been
deaerated to expel dissolved O2 in it. The
corrosion reaction is
Zn(s) Zn2(aq) 2e The reduction reaction is
2H(aq) 2e
H2(g)
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(Assume concentration can be used in place of
activities.)
Note By convention, the concentration of 10-6M
is used for the concentration of ions produced
during a corrosion and when concentration exceeds
this limit, it is assumed that the metal has
undergone corrosion.
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Therefore ?G lt 0
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The Zn vessel dissolves (corrodes). Therefore
Zinc vessel cannot be used to store, in HCl (1 M)
in the above environment. For the corrosion to
occur the ?Gcorr must always be less than zero.
However, this criterion as it is can be used to
study corrosion of specific subjects in a
controlled environment. In natural environmental
there are many different competing cathodic
reactions and also wide range of temperatures and
purity of corrosions metals. As a result the
criterion should used cautiously.
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On the other hand, if a metal is found to be
thermodynamically unstable against corrosion it
may not under go corrosion. For example,
Aluminum metal it can be shown to be unstable
against corrosion in natural environment.
However, in natural environment it does not
corrode as kinetic of the reaction does not allow
the metal to corrode. The corrosion product of
Al, i.e. Al2O3 protect the metal against
corrosion, as it act as a hard non conducting
film the metal surface which covers the inner
part of the metal.
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Kinetics of corrosion Determination of
corrosion Potential (Ecorr) and Corrosion current
(icorr)
We can examine the previous example of corrosion
of pure Zinc in 1.0 M deaerated HCl solution at
250C. The corrosion reaction is Zn(s)
Zn2(aq) 2e And
E0cell (-) 0.777V
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A possible reduction reaction is 2H(aq) 2e
H2(g) E0H2
0.0 V We need to construct Butler-Volmer current
potential curves for the both the reactions.
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H2(g)?2H(aq) 2e
Zn2 2e Zn
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Evans Diagram
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Note On () ve section of the x axis () ve
potentials have been plotted. Under the given
condition
During a corrosion, the electron generated at the
corroding metal is consumed by the reduction
reaction at the same rate they are produced.
Therefore current of the Zn
? Zn2 2e Reaction is equal to the
cathodic reaction of the
2H 2e ? H2
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The potential of which these two currents became
equal can be found as indicated by the figure
above. This process gives a method to determine
corrosion current and corrosion potential. The
corrosion current can be used to calculate rate
of corrosion of metal in terms of mass loss per
unit time or penetration depth per unit time. The
values of Ecorr together with icorr can be used
to device methods for controlling of corrosion
rates.
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THE END