Transcript Lecture 8

Electrochemistry MAE-212
Dr. Marc Madou, UCI, Winter 2016Class VIII
Corrosion
Table of Content
Definition
Why study corrosion?
Thermodynamics of Corrosion
Corrosion Illustrated
Protection Mechanisms
Evans Diagram
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Definition
 Corrosion is the deterioration of materials by chemical
interaction with their environment. The term corrosion is
sometimes also applied to the degradation of plastics,
concrete and wood, but generally refers to metals. The most
widely used metal is iron (usually as steel) and the following
discussion is mainly related to metal corrosion.
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Why Study Corrosion?
 Materials are precious resources
 Engineering design is incomplete without knowledge of corrosion
 Corrosion contaminates products such as pharmaceutical, food and dairy







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products
Corrosion products are a threat to the environment
Ensuring maximum life of new equipment
Preservation of existing equipment
Protecting or improving the quality of a product in order to maintain or
improve a competitive position.
Avoiding costly interruptions of production.
Reducing or eliminating losses of valuable products by spillage or leaks.
Reducing hazards to life and property that might be associated with
corrosion:
 Explosions of pressure vessels or piping systems
 release of poisonous or explosive gases or vapors
 Artificial implants for the human body!!!
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Thermodynamics of Corrosion: Gibbs Free
Energy
 Thermodynamic considerations allow the determination of whether a
reaction can occur spontaneously
 If metal dissolution is unfavorable thermodynamically in a given set of
circumstances – the job of the corrosion engineer is done
 Example: Copper in pure deoxygenated water
 Free Energy: Driving Force of a Chemical Reaction
 The larger the value of E for any cell – the higher is the
tendency for the overall cell reaction to proceed:
G  EnF
Ecell = Ecathode - Eanode
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Thermodynamics of Corrosion: Gibbs Free
Energy
Spontaneous
Spontaneous
Not Spontaneous
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Thermodynamics of Corrosion: The
Nernst Equation
General Reaction for a Galvanic Cell
lL  mM        qQ  rR       
Nernst Equation:
q
r
RT aQ .aR ...
EE 
ln l m
nF aL .aM ...
0
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Thermodynamics of Corrosion: Oxygen
Electrodes
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Thermodynamics of Corrosion: Oxygen
Electrode and Differential Aeration Cell
 Consider two O2 electrodes:
 one in contact with O2 at 1 atm
 other in contact with O2 at 0.2 atm
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Thermodynamics of Corrosion: Oxygen Electrode and
Differential Aeration Cell
 The reaction is not thermodynamically possible as written
 Thus, the electrode 1 is cathode electrode 2 the anode.
 In a differential aeration cell, the electrode in lower O2 pressure
acts as the anode and the one in higher O2 pressure acts as the
cathode
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Thermodynamics of Corrosion:
Galvanic Series
 Galvanic series is an arrangement of both metals and alloys according
to their actual measured potentials in a particular environment. There
is a Galvanic series for each environment
 Metals and alloys showing active-passive behavior are listed in both
active and passive states.
In the 1910s, before the First World War, people were concerned about how easily gun barrels made of steel would
corrode. A scientist called Henry Brearly found that adding about 10% chromium to the steel made an alloy which
was very resistant to corrosion. He called this new alloy ‘rustless steel’.
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Thermodynamics of Corrosion:
Galvanic Scale of Metals in Seawater
 Stainless steel is used to make cutlery,
saucepans, surgical instruments and
food transporters.
 Stainless steel is an alloy of iron which
contains around 12% chromium and
smaller amounts of nickel and carbon
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Thermodynamics of Corrosion:
Pourbaix Diagram
 The Nernst Equation allows us to compute lines for equilibrium
reactions of interest …
 Electrochemical reactions of pure charge transfer are horizontal
lines since there is no H+ or OH- dependence;
 Pure acid-base reactions are vertical lines since no electron
transfer occurs, thus there is no dependence on potential;
 Electrochemical reactions involving charge transfer and H+(OH-)
are represented as sloping lines.
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Thermodynamics of Corrosion:
Pourbaix Diagram
 Marcel Pourbaix developed potential-pH diagrams to show the




thermodynamic state of most metals in dilute aqueous solutions
With pH as abscissa and potential as ordinate, these diagrams have
curves representing chemical and electrochemical equilibria
between metal and aqueous environment
These diagrams ultimately show the conditions for immunity,
corrosion or passivation. A Pourbaix Diagram is a thermodynamic
map of corrosion, passivity and nobility of a particular metal as a
function of system pH and potential.
Operating environments are then superimposed on the diagram to
determine the optimum operating conditions for a particular metal.
Pourbaix diagrams may also be used to show regions of different
forms of corrosion:
 General corrosion;
 Pitting corrosion;
 Passivation.
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Thermodynamics of Corrosion:
Simplified Pourbaix Diagram for Iron
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Thermodynamics of Corrosion:
Simplified Pourbaix Diagram for Iron
 It may also be
observed that the
form of corrosion
is highly sensitive
to the level of
dissolved oxygen
in the
environment:
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Thermodynamics of Corrosion: Benefits
of Pourbaix Diagram
 Pourbaix diagrams offer a large volume of thermodynamic information in a
very efficient and compact format.
 The information in the diagrams can be beneficially used to control
corrosion of pure metals in the aqueous environment
 By altering the pH and potential to the regions of immunity and passivation,
corrosion can be controlled. For example, on increasing the pH of
environment in moving to slightly alkaline regions, the corrosion of iron can
be controlled
 Changing the potential of iron to more negative values eliminate corrosion,
this technique is called cathodic protection.
 Raising the potentials to more positive values reduces the corrosion by
formation of stable films of oxides on the surface of transition metals
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Corrosion Illustrated: Metal Loss
Through Voltaic Cells
In moist air, exposed iron can be oxidized to Fe2+ … these areas are referred to
as anodic areas
Other regions of the iron serve as cathodic areas
Electrons from the anodic areas reduce atmospheric oxygen to the OH– ion
Iron (II) ions migrate from the anodic areas to the cathodic areas, where they combine
with the hydroxide ions and are further oxidized to iron(III) hydroxide by atmospheric
oxygen … Fe2O3·xH2O is
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common rust
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Corrosion Illustrated: : Metal Loss
Through Voltaic Cells
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Corrosion Protection
 The simplest line of defense
against the corrosion of iron
is to paint it to exclude
oxygen from the surface
 Another approach is to coat the
iron with a thin layer of a less
active metal
 An entirely different approach is
to protect iron with a more
active metal, as in the zinc-clad
iron known as galvanized iron or
in stainless steel.
air
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 Stainless Steel:
 Chromium is more reactive than
iron. It reacts quickly with oxygen
from the air to form a very thin
layer of chromium oxide on the
surface of the steel.
 This protects the iron atoms from
reacting with the oxygen in the air
and prevents rust forming.
 If the steel is scratched or cut, more
chromium atoms quickly form a
new protective layer.
protective
Cr2O3 layer
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Corrosion Protection
Cathodic protection: the iron object to be
protected is connected to a chunk of an active
metal and the iron serves as the reduction halfcell
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Example Pourbaix Diagram:
 Consider a fairly simple system … aluminum corroding in
water:
 the following species are considered to be predominant in the
system:
 Al (s);
 Al3+ (aq);
 Al2O3.H2O (s);
 AlO2- (aq).
 We can consider several “equilibrium” reactions describing
the thermodynamic stability of these aluminum-containing
species …
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Example Pourbaix Diagram:

Pure charge transfer reactions:
1.

Al3+(aq) +3 e- n Al (s)
Pure acid/base reactions:
Al2O3.H2O + 6 H+ n 2 Al3+ + 4 H2O
3. 2 AlO2- + 2 H+
n
Al2O3.H2O
2.

Mixed charge transfer and acid/base reactions:
4.
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Al2O3.H2O + 6 H+ + 6 e- n 2 Al (s) + 4 H2O
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Example Pourbaix Diagram:
 Lines of pure charge transfer are calculated directly through the
Nernst Equation …
Al3+
+
3
e
« Al (s)
(aq)
E Al 3+ /Al
æa ö
2.303RT
o
Al ÷
ç
= E Al
log
3+
/Al
ça ÷
nF
è Al 3+ ø
 Eo is calculated through the ∆Grxno or looked up in a table of
standard potentials; activity of solid Al = 1 …
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Example Pourbaix Diagram:
 Thus;
(
E Al 3+ /Al = -1.662 + 0.0197log aAl 3+
)
 We see that the potential is dependent upon the concentration of Al3+
in the system.
 In general, when constructing Pourbaix diagrams we choose ionic
concentrations of …
100, 10-2, 10-4 and 10-6 mol/L
(more correctly this is mol/kg but at room temperature they are moreor-less equal)
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Example Pourbaix Diagram:
Potential (E) vs SHE
-1.4
Al3+
-1.6
100
10-2
10-4
10-6
-1.8
Al
-2
0
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2
4
6
8
pH
10
12
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Example Pourbaix Diagram:
 Lines for pure acid/base reactions (no electron transfer) are
calculated through the equilibrium constant or reaction quotient
…
Al 2O3.H 2O + 6 H+ ® 2 Al3+ + 4 H 2O
( )
o
-DG rxn
= RT ln K eq
2
K eq =
aAl 3+ a
aAl O .H O aH6 +
2 3
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4
H 2O
2
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Example Pourbaix Diagram:
 Accounting for Al2O3.H2O as a solid and that water will have an
activity of 1.0 gives:
aH6 + =
 Taking logs …
6logaH +
aA2 l 3+
K eq
æ a 2 3+ ö
= log çç Al ÷÷
è K eq ø
( )
( )
pH = 0.166log ( K ) - 0.333log (a )
-pH =1/ 3log aAl 3+ -1/ 6log K eq
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eq
Al 3+
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Example Pourbaix Diagram:
 Calculating Keq from the Gibbs energies of formation tabulations:
Gf ------
Al3+
H20
Al2O3.H2O
H+
-485.4
-237.0
-1825.5
0.0
kJ/mol
kJ/mol
kJ/mol
kJ/mol
æ -é2(-485.4) + 4(-237)-(-1825.5)ù ´1000
ë
û
K eq = exp ç
ç
8.314 298
è
Thus;
(
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)( )
(
pH = 2.732 - 0.333log aAl 3+
)
ö
÷ = 2.882 ´1016
÷
ø
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Example Pourbaix Diagram:
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Example Pourbaix Diagram:
 For the other acid/base reaction:
2 AlO-2 + 2 H+ « Al 2O3.H 2O
K eq =
aAl O .H O
2 3
2
(aAlO - )2 (aH + )2
2
( )
( )
pH = 0.5log K eq + log aAlO 31
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Example Pourbaix Diagram:
 Gibbs energies of formation …
Gf ------
AlO2Al2O3.H2O
H+
-838.968
-1825.5
0.0
kJ/mol
kJ/mol
kJ/mol
Keq = 7.354 x 1025
( )
pH = 12.93+ log a AlO-
2
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Example Pourbaix Diagram:
2.5
10-6
Potential (E) vs SHE
2
10-4
10-2
100
1.5
1
0.5
AlO2-
0
Al2O3.H2O
-0.5
-1
-1.5
-2
-2.5
4
6
8
10
12
14
16
pH
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Example Pourbaix Diagram:
 Finally, for mixed reactions involving both charge transfer and
acid/base reactions …
Al2O3.H2O + 6 H+ + 6 e- → 2 Al (s) + 4 H2O
 We calculate these directly through the Nernst equation:
E Al O /Al
2 3
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4ö
2
æ
ç aAl aH 2O ÷
2.303RT
o
= E Al O /Al log ç
2 ÷
2 3
nF
ç aAl O a + ÷
è 23 H
ø
( )(
( )
)
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Example Pourbaix Diagram:
 Eo is calculated through the Gibbs Free Energy of reaction …
Gf ------
H 20
Al2O3.H2O
Al
H+
(
-237.0
-1825.5
0.0
0.0
) (
kJ/mol
kJ/mol
kJ/mol
kJ/mol
) (
é4 ´ -237.0 - -1825.5 ù ´ 1000J / kJ
o
-DG
ë
û
o
rxn
E =
=
nF
6 ´ 96485
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E o = -1.516 Volts
)
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Example Pourbaix Diagram:
 So …
( )
E Al O /Al = -1.516 + 0.059log aH +
2 3
E Al O /Al = -1.516 - 0.059 pH
2 3
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Example Pourbaix Diagram:
Potential (E) vs SHE
-1.4
Al2O3.H2O
-1.6
Al
-1.8
-2
0
2
4
6
8
10
12
14
pH
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Example Pourbaix Diagram:
 Putting all our lines together … we get the “Pourbaix”
Diagram for Aluminum …
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Example Pourbaix Diagram:
 Note the stability lines (a) & (b) that are indicated on the
diagram.
 These denote the “region of stability” for water based on the
following reactions …
Line a:
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2H + + 2e - « H 2
Line b:
O2 + 4H + + 4 e - « H 2O
or equivalently:
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Example Pourbaix Diagram:
 Again, from the Nernst equation …
E H + /H = - 0.059 pH
2
- 0.0295log p H
2
 We usually consider the case where pH2 = 1 atm …
Line a:
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E H + /H = - 0.059 pH
2
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Example Pourbaix Diagram:
 For line b:
EO = E - 0.0148log
2
o
O2
aH O
2
pO aH4 +
2
 EO2o = 1.230 V
(as calculated through Gibbs of formation
data), thus:
EO =1.230 - 0.0591 pH
2
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Evans Diagrams: Butler-Volmer Equation
where:
I = electrode current, Amps
Io= exchange current density, Amp/m2
E = electrode potential, V
Eeq= equilibrium potential, V
 A = electrode active surface area, m2
 T = absolute temperature, K
 n = number of electrons involved in the electrode reaction
 F = Faraday constant
 R = universal gas constant
 α = so-called symmetry factor or charge transfer coefficient dimensionless
The equation is named after chemists John Alfred Valentine Butler and Max Volmer





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Evans Diagrams: Butler-Volmer Equation – High
Field Strength
 1   nF a 
ia `  i0 exp 

RT


at high anodic overpotent ial
 nF c 
ic `  i0 exp 

RT


at high cathodic overpotent ial
ia and ic are the
exhange current
densities for the
anodic and
cathodic
reactions
These equations can be rearranged to give the Tafel
equation which was obtained experimentally
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Evans Diagrams: Butler Volmer Equation - Tafel
Equation
c 
RT
RT
ln i0 
ln ic
 c nF
 c nF
c 
0.059
0.059
log i0 
log ic at 25 0 C cathodic reaction
cn
cn
a 
0.059
0.059
log i0 
log ia at 25 0 C anodic reaction
1  c n
1  c n
The equation is the well known Tafel equation
  a  b log i
a
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0.059
0.059
ln io or a 
ln io
1  n
n
0.059
0.059
b
or
1  n
n
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Evans Diagrams: Current Voltage Curves for
Electrode Reactions
Without concentration and
therefore mass transport
effects to complicate the
electrolysis it is possible to
establish the effects of voltage
on the current flowing. In this
situation the quantity E - Ee
reflects the activation energy
required to force current i to
flow. Plotted below are three
curves for differing values of
io with α = 0.5.
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Evans Diagrams: Current Voltage Curves for Single
Electrode Reactions
The iE curves from the
previous slide have been
rotated.
Voltage
Electrochemical reactions of
different i0 or degrees of
reversibility
Current
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Evans Diagrams: Single Chemical Reaction
Only at appreciable overpotentials does the reverse
reaction become negligible
At Ee the forward and
reverse currents are
equal
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Electrochemical reaction which has a large exchange current
density, i0, This means that a small applied voltage results in an
appreciable increase in current.
Electrode reactions which have a
high exchange current density are
not easily polarised. Examples are
the hydrogen evolution reaction on
Pt and AgCl + e ↔ Ag + ClThe H+/H2(Pt) and Ag/AgCl make
good reference electrodes because
they are not easily polarised
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Electrochemical reaction in which the i0 value is very low. This
means that it takes an appreciable over-potential to produce a
significant current.
This electrode is easily
polarisable since a
small current would
result in a significant
change in voltage
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At low overpotential the Butler Volmer equation is linear (Stern Geary
equation)
nF
i  io

RT
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So far we have looked mainly at single
electrochemical reactions
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KINETICS OF AQUEOUS CORROSION
Anodic and cathodic reactions are coupled at a corroding metal surface
Schematics of two distinct corrosion processes.
(a) The corrosion process M + O  Mn+ + R showing the separation of anodic and
cathodic sites. (b) The corrosion process involving two cathodic reactions.
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Butler Volmer graphs for two electrochemical reactions
Wagner Traud Method
The cathodic and
anodic reactions are
drawn together on the
same graph to show
how the currents are
equal at the corrosion
potential
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 Note in the previous diagram that:


ia = ic = icorr at the corrosion potential Ecorr
Ecorr is a mixed potential which lies between (Ee)c and (Ee)a. In this case it is
closer to (Ee)a because the i0 and the kinetics of the anodic reaction is faster.
 The metal dissolution is driven by the anodic activation overpotential
Ecorr - (Ee)a
 The cathodic reaction is driven by the cathodic activation overpotential
Ecorr - (Ee)c
ηa =
ηc =
 The thermodynamic driving force

ΔE = (Ee)c - (Ee)a

ΔE is usually large enough to put Ecorr
in the Tafel region for both reactions, i.e.
the reverse reaction is negligible.
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Evans Diagrams
It is convenient to represent the
linear plots of i and E as log i/E
plots with the negative cathodic
current plotted positively, i.e.
both the anodic and cathodic
current appear in the positive
quadrant.
The linear region gives us the
Tafel slopes
The i0 for the individual
reactions can be obtained by
extrapolating back to (Ee)a and
(Ee)c if these values are known.
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Evans Diagrams
In this case the
cathodic reaction
with the higher
oxidation potential is
controlling the
reaction
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Evans Diagrams
In this example because
of the faster kinetics.
the cathodic reaction
taking place at the
lower oxidation (+ve)
Potential is influencing
the corrosion rate
more,
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Evans Diagrams
 The situation in the previous example often occurs for a metal
corroding in acid, compared with the metal corroding in dissolved
oxygen.
 Despite the thermodynamic driving force, Ee, being greater for oxygen
than H2/H+, the acid corrosion is faster.
 In some cases the oxygen and acid have a synergistic effect. For
example in the case of Ni corrosion. The reaction is quite slow in
sulphuric acid (0.5 M) and it is also slow in water saturated with air at
pH 7. In the latter case a passive protective oxide film is formed.
However, in the presence of sulphuric acid and air. The corrosion rate
is relatively rapid. The acid dissolves the protective oxide film allowing
oxygen to corrode the metal.
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Evans Diagrams
• The relative corrosion rates of metals depends on the i0 and mass
transfer.
• With acid corrosion: 2H+ + e → H2
• i0 can vary from 10-3 – 10-12 A cm-2
• The Tafel slope  120 mV/decade
• For oxygen corrosion O2 + H2 O + 4e → 4OH• I0 is difficult to difficult to determine because it is very low, but it is of
the order of <10-10 A cm-2
• The Tafel slope >120 mV/decade
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Exchange Current Densities in 1 Molal H2SO4
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Electrode Material
-log10(A/cm2
Palladium
3.0
Platinum
3.1
Rhodium
3.6
Nickel
5.2
Gold
5.4
Tungsten
5.9
Niobium
6.8
Titantium
8.2
Cadmium
10.8
Manganese
10.9
Lead
12
Mercury
12.3
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α Values for Some Reactions
Metal
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System
α
Pt
Pt
Hg
Hg
Fe3+
Ce4+
Ti4+
2H+
+ e ↔ Fe2+
+ e ↔ Ce3+
+ e ↔ Ti3+
+ 2e ↔ H2
0.58
0.75
0.42
0.50
Ni
Ag
2H+ + 2e ↔ H2
Ag+ + e ↔ Ag
0.58
0.55
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Evans Diagrams
• The slowest reaction controls
the rate of corrosion.
• Normally this is the cathodic
reaction.
• In this example:
• A small changes in kinetics of
cathode have a large effect on
corrosion rate.
• A small changes in kinetics of
anode have small effect on
corrosion
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Mass Transfer Control
• If the cathodic reagent at the corrosion site (e.g., dissolved O2 in the
O2 reduction) is in short supply, mass transfer of the reagent can
become rate limiting.
• The cathodic charge-transfer reaction at the metal/solution interface
is fast enough to reduce the concentration of the reagent at interface
(cathodic sites) to a value less than that in the bulk solution.
• This sets up a concentration gradient and the reaction becomes
diffusion controlled.
ic 
i 
nFDCb  C s 
c Lim
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

nFDCb 

 icorr max
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Mass Transfer Control
• When the corrosion rate is limited by mass transfer it can be
increased by:
• By altering the bulk concentration
• By stirring and reducing the thickness of the Nernst diffusion layer
ic 
i 
nFDCb  C s 
c Lim
65


nFDCb 

 icorr max
Where :
ic  the cathodic current
n  the number of electrons
F  the Faraday constant
Cb  the bulk concentrat ion
C s  the surface concentrat ion
  the Nernst diffusion layer
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Mass Transfer Control
Activation Controlled
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Diffusion or Mass Transfer
Controlled
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Mass Transfer Control
Increase in corrosion potential, Ecorr, and the corrosion current, icorr, due to an
increase in mass transfer caused by stirring.
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Mixed Transfer Control
The cathodic Tafel plot often shows deviation from ideal Tafel behavior
Polarization curve
for the cathodic
process showing:
1. Activation
polarization
2. Joint activationconcentration
polarization
3. Mass transportlimited corrosion
control
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Evans Diagrams
Anodic Control
Cathodic
Control
Mixed
Control
69
8/7/2016
Galvanic Corrosion – Influence of i0
70
8/7/2016