Transcript
Corrosion Mechanisms
Randhir Kumar Singh
Asst Professor
OPJIT
Chemical vs. Electrochemical Reactions
Chemical reactions are those in which elements are
added or removed from a chemical species.
Electrochemical reactions are chemical reactions in
which not only elements may be added or removed
from a chemical species but at least one of the
species undergoes a change in the number of
valence electron.
Corrosion processes are electrochemical in nature.
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Simplest Example: Dry Cell Battery
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Faraday’s laws of electrolysis, in chemistry,
quantitative laws used to express magnitudes of
electrolytic effects, first described by the English scientist
Michael Faraday in 1833.
The laws state that
(1) the amount of chemical change produced by current at
an electrode-electrolyte boundary is proportional to the
quantity of electricity used, and
(2) the amounts of chemical changes produced by the same
quantity of electricity in different substances are
proportional to their equivalent weights.
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Contd…
In electrolytic reactions, the equivalent weight of a
substance is the gram formula weight associated with a
unit gain or loss of electron. The quantity of electricity
that will cause a chemical change of one equivalent
weight unit has been designated a faraday. It is
equivalent to 9.6485309 × 104 coulombs of electricity.
Thus, in the electrolysis of fused magnesium chloride,
MgCl2, one faraday of electricity will deposit 24.312/2
grams of magnesium at the negative electrode and
liberate 35.453 grams of chlorine at the positive
electrode.
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Faraday’s Law
The mass of an element discharged at an
electrode is directly proportional to the
amount of electrical charge passed through
the electrode
weight of metal reacting = kIt
where I = Current Intensity
t = time of current passage
k = Constant
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What Happens if the Battery is Not in Use?
There will be some “local action current”
generated by “local action cells” because of
other metallic impurities in zinc
Shelf life of an ordinary zinc-carbon rod battery
is limited
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Local Action Cell
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Anode & Cathode
Anode
Loss of electron in oxidation
Oxidation always occurs at the anode
Cathode
Gain of electron in reduction
Reduction always occurs at the cathode
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Difference between Electrochemical and Electrolytic Cell
fes/fe2+
aq //sn2+aq/sns Both types of cells consist of two electrodes
connected to an electrolyte (an ionically conducting phase).
Electrode reactions then take place at the electrode-solution
surfaces. The change from electronic current to ionic current and
visa versa are always accompanied by oxidation/reduction
reactions.
An electrochemical cell is simply a device that converts chemical
energy into electrical energy when a chemical reaction is occurring
in a cell. An electrolytic cell converts electrical energy into chemical
energy.
In an electrochemical cell the reaction occurs spontaneously at the
electrodes, while an electrolytic cell reaction is not spontaneous at
the electrodes - the reaction has to be forced by applying an
external electrical current.
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In an electrochemical cell the cathode is positive and the anode is
negative. In an electrolytic cell the cathode is negative and the anode is
positive (does this mean that the electrons are going against their
gradient here?).
In a spontaneous chemical reaction electrons are passed directly from
one element to another. In an electrochemical cell these simultaneous
redox reactions are "spatially separated" - i.e. happen at different places.
The resultant ions then combine to form a new product. During this
process electrons are conducted from the anode to the cathode through
an outside electrical current which can be used. This action can be
reversed in a electrolytic cell.
Electrochemical cells are used usually as batteries, while electrolytic
cells are used for electroplating metals. Also, the recharging of a
rechargeable battery is an electrolytic reaction.
Corrosion Cells
Galvanic cell (Dissimilar electrode cell) – dissimilar metals
Salt concentration cell – difference in composition of aqueous environment
Differential aeration cell – difference in oxygen concentration
Differential temperature cell – difference in temperature distribution over the body of the metallic material
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Dissimilar Electrode Cell
When a cell is produced
due to two dissimilar metals
it is called dissimilar
electrode cell
Dry cell
Local action cell
A brass fitting connected to a
steel pipe
A bronze propeller in contact
with the steel hull of a ship
Zn anode
HCl Solution
Cu cathode
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Salt Concentration Cell
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Differential Aeration Cell
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Corrosion at the bottom of the electrical poles
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Differential Temperature Cell
This is the type of cell form when two identical electrodes are immersed in same electrolyte, but the electrodes are immersed into solution of two different temperatures
This type of cell formation takes place in the heat exchanger equipment where temperature difference exists at the same metal component exposed to same environment
For example for CuSO4 electrolyte & Cu electrode the electrode in contact with hot solution acts as cathode.
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Factors affecting choice of an engineering material
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Factors affecting corrosion resistance of a metal
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Corrosion Rate Expressions
mm/y – millimeters penetration per year
gmd – grams per square meter per day
ipy – inches penetration per year
mpy – mils penetration per year (1000 mil =
1 inch)
mcd – milligrams per square centimeter per
day
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Corrosion Rate Expressions
The most used expression for Corrosion Rate in the US is the mpy (Mils per
year).
Mils per year (mpy) = 534W/DAT
To convert corrosion rate (corrosion rate conversion) between the mpy and
the equivalent in metric unit mm/y (millimeter per year):
1 mpy = 0.0254 mm/y = 25.4 micron/y
To calculate the corrosion rate from metal loss:
mm /y = 87.6 x (W / DAT)
W = weight loss in mg
D = density of specimen material in g/cm3
A = area in cm2
T= exposure time in hours
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Classification of metallic materials
according to their rate of uniform attack
A. <0.005 ipy (<0.15 mm/y) – Metals in this
category have good corrosion resistance
and can be used for critical parts
B. 0.005 to 0.05 ipy (0.15 mm/y to 1.5 mm/y) –
Metals in this group are satisfactory if a
higher rate of corrosion can be tolerated
C. >0.05 ipy (>1.5 mm/y) – Usually not
satisfactory
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Electrochemical Aspects
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Electrochemical Reactions
Electrochemical Reactions
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The above concept is illustrated in the Fig.1
Fig.1 Electrochemical reactions
occurring during corrosion of zinc
in air-free hydrochloric acid
Electrochemical Reactions
25 Fig.2 Electrochemical reactions occurring during corrosion
of zinc in aerated hydrochloric acid
Let us look at a Zn-H Cell
(a)
(b)
The Zn electrode moves away from equilibrium by the removal of negative charges from the Zn plate and positive ions are released from the Zn plate to the liquid (a)
Zn is dissolved at the same rate as electrons are transported to the Pt plate, where they are consumed in the hydrogen reaction
The same cell process can be totally obtained on a Zn plate submerged in a solution containing hydrogen ions and Zn ions (b)
The reactions are accompanied by the same changes in free enthalpy and have the same equilibrium potentials as before
However, there is a higher resistance against the hydrogen reaction on the Zn plate than on Pt, and thus the reaction rate will be lower on the Zn surface
So We Also Need to Know …
Electrode kinetics to predict the corrosion
rates for the actual conditions
Single and mixed electrodes
Whenever only one electrode reaction takes place on a
metal surface in a given solution, that system is called a
single electrode.
This is the case for copper immersed into de-aerated and
slightly acidic copper sulfate solution:
Cu2+ + 2 e = Cu -- (1)
The open circuit potential (ocp) is the potential set up
spontaneously by an electrode in the absence of an
external current. For a single electrode, the open circuit
potential is equal to the equilibrium potential, Erev.
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Single and mixed electrodes Most often, several electrode reactions take place simultaneously at a metal
electrolyte interface. Such systems are referred to as mixed electrodes. If, in the
previous example, the copper sulfate solution is aerated, two electrode reactions
(called partial reactions) are observed at the open circuit potential; the oxidation of
copper
Cu→Cu2+ + 2 e -- (2)
and the reduction of oxygen:
½ O2 + 2H+ + 2 e → H2O -- (3)
The corresponding overall reaction is
Cu + ½ O2 + 2H+ → Cu 2+ H2O -- (4)
The copper thus corrodes without any external current. The open circuit potential of
a mixed electrode undergoing corrosion, is called the corrosion potential (in the
literature it is sometimes also called the free corrosion potential).
The corrosion potential has a value that lies in between the equilibrium potentials
of the partial electrode reactions. In contrast to the equilibrium potential, which is a
thermodynamic quantity, the corrosion potential is determined by kinetics; its value
depends on the rates of both the anodic and the cathodic partial reactions present.
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Single and mixed electrodes
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Figure A copper electrode in contact with itw own ions (single electrode)
and with an aerated solution (mixed electrode).
Polarization and Overpotential Polarization
Electrode reactions are assumed to induce deviations from equilibrium due to the passage of an electrical current through an electrochemical cell causing a change in the electrode potential. This electrochemical phenomenon is referred to as polarization.
The polarization ζ expresses the difference between the potential of a mixed
electrode subjected to anodic or cathodic polarization and its corrosion potential.
ζ = E – Ecor
A polarization of ζ > 0 indicates an anodic and a polarization of ζ < 0 a cathodic
current flow
Overpotential
The deviation from equilibrium causes an electrical potential difference between the polarized and the equilibrium (unpolarized) electrode potential known as overpotential
η = E – Erev
A positive overpotential indicates that an anodic current is crossing the interface; a
negative one means that the current is cathodic.
Polarization and Overpotential
Equilibrium potential for cathodic reaction = Eoc
Equilibrium potential for anodic reaction = Eoa
Real potential = E
Cathodic Overpotential ηc = E – Eoc < 0
anodic Overpotential ηa = E – Eoa > 0
The Polarized Cell
Exchange Current Density
At the equilibrium potential of a reaction, a reduction and an oxidation reaction occur, both at the same rate.
For example, on the Zn electrode, Zn ions are released from the metal and discharged on the metal at the same rate
The reaction rate in each direction can also be expressed by the transport rate of electric charges, i.e. by current or current density, called, respectively, exchange current, Io, and (more frequently used) exchange current density, io.
The net reaction rate and net current density are zero
How Polarization is Measured
Causes of Polarization
Depending on the type of resistance that
limits the reaction rate, we are talking about
three different kinds of polarization activation polarization
concentration polarization and
resistance (ohmic) polarization or IR Drop
Activation Polarization
When current flows through the anode and the cathode electrodes, their shift in potential is partly because of activation polarization
An electrochemical process that is controlled by reaction sequence at
the metal-electrolyte interface.
This is easily illustrated by considering hydrogen evolution reaction on
zinc during corrosion in acid solution.
An electrochemical reaction may consist of several steps
The slowest step determines the rate of the reaction which requires activation energy to proceed
Subsequent shift in potential or polarization is termed activation polarization
Activation polarization usually is the controlling factor during corrosion
in media containing a high concentration of active species(e.g.
concentrated acids).
Most important example is that of hydrogen ion reduction at a cathode, H+ + e- → ½ H2, the polarization is termed as hydrogen overpotential
Activation Polarization
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Fig. Hydrogen-reduction reaction under activation control(simplified).
Hydrogen Overpotential
Hydrogen evaluation at
a platinum electrode:
H+ + e- → Hads
2Hads → H2
Step 2 is rate limiting
step and its rate
determines the value of
hydrogen overpotential
on platinum
Tafel Equation
Activation polarization (η) increases with
current density in accord with Tafel equation:
The Tafel constant is given by:
oi
ilog
nF
RTβ
α
3.2
Overpotential
Values
Concentration Polarization
It refers to electrochemical reactions that are controlled
by the diffusion in electrolyte.
Sometimes the mass transport within the solution may
be rate determining – in such cases we have
concentration polarization
Concentration polarization implies either there is a
shortage of reactants at the electrode or that an
accumulation of reaction product occurs
Concentration polarization generally predominates when
the concentration of the reducible species is small(e.g.
dilute acids, aerated salt solutions).
Concentration Polarization
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Fig.3 Concentration polarization during hydrogen reduction.
Concentration Polarization: reduction of
oxygen O2H4e4HO 22
IR Drop
When polarization is measured with a potentiometer and a reference electrode-Luggin probe combination, the measured potential includes the potential drop due to the electrolyte resistance and possible film formation on the electrode surface
The drop in potential between the electrode and the tip of Luggin probe equals iR.
If l is the length of the electrode path of cross sectional area s, k is the specific conductivity, and i is the current density then resistance
iR drop in volts = k
lR
k
il
k
il
Combined Polarization
Total polarization of an electrode is the sum
of the individual contributions,
If neglect IR drop or resistance polarization is
neglected then:
rcaT ηηηη
caT ηηη
Combined Polarization
Effect of temperature, concentration and velocity of the aqueous
environment on combined polarization is shown in the figure
Passivity
Passivity refers to the loss of chemical
reactivity experienced by certain metals and
alloys under particular environmental
conditions.
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Fig. Corrosion rate of a metal as a
function of solution oxidizing
power (electrode potential).
Passivity
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Fig. Corrosion characteristics of an active-passive metal as a function
of solution oxidizing power (electrode potential).
Passivity in Iron Chromium Alloys
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Dissolution of an active metal (active dissolution), involves a charge transfer at
the metal-electrolyte interface. Soluble ions, either hydrated or complexed, are
formed and dissolve into the electrolyte, while the liberated electrons either flow
to the cathode or are taken up by an oxidizing agent.
When a passive metal dissolves (passive dissolution), cations are formed by a
charge transfer reaction at the metal-film interface. They migrate across the
passive film to the film-electrolyte interface, where they dissolve into solution as
hydrated or complexed ions.
Cont… Because of the presence of an oxide film, the dissolution rate of a passive
metal at a given potential is much lower than that of an active metal. It
depends mostly on the properties of the passive film and its solubility in the
electrolyte.
During passivation, which is a term used to describe the transition from the
active to the passive state, the rate of dissolution therefore decreases
abruptly.
The polarization curve of a stainless steel in sulfuric acid, given in Figure,
illustrates this phenomenon. In this electrolyte, the corrosion potential of the
alloy is close to –0.3 V.
Anodic polarization leads to active dissolution up to about –0.15 V, where the
current density reaches a maximum. Beyond this point, the current density,
and hence the dissolution rate, drops sharply.
It then shows little further variation with potential up to about 1.1 V. Above
that value the current density increases again because transpassive
dissolution and oxidation of water to oxygen becomes possible.
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Fig. Polarization curve of Fe-17Cr stainless steel in 0.5 M H2SO4. Sweep
rate is 0.02 V/min.
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Figure. Variation of partial anodic current density with potential for a
passivating metal (Evans Diagram).
Generally speaking, we can distinguish three potential regions in the polarization curve of a
passivating metal (Figure ):
• the active region;
• the passive region;
• the transpassive region.
In contrast to the active and passive regions, the surface state of the metal in the transpassive
region is not well defined and an oxide may or may not cover the surface.
The current density measured during a polarization experiment is the
sum of all anodic and cathodic partial current densities. Figure .
Schematically shows the variation of the anodic partial current density of
a passivating metal as a function of the potential. It allows us to define a
number of quantities that describe the polarization behavior of
passivating metals.
The passivation potential Ep separates the active from the passive
potential region.
The corresponding current density at the maximum is the passivation
current density, ip.
The passive current density ipp characterizes the dissolution behavior of
the metal in the passive potential region.
The transpassivation potential Eb marks the end of the passive
potential region and the transition from passive to transpassive behavior.
Beyond this point the anodic partial current density increases markedly
with increasing potential due one of the following processes: uniform
transpassive dissolution resulting from oxidation of the passive film,
dissolution by pitting resulting from local film breakdown, oxygen
evolution due to water oxidation.
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When dissolution by pitting is the dominating reaction, the potential Eb is
called pitting potential or critical pitting potential. Often, Eb is also
referred to as film breakdown potential, indicating that pitting is initiated
by passive film breakdown.
Depending on conditions, the value of Eb can be either above or below
the reversible potential of the oxygen electrode, Erev,O2. For sufficiently
stable passive films with good electronic onductivity, oxygen evolution
rather than transpassive dissolution may therefore account for the
observed current at high anodic potentials.
The pitting potential, plays an important role for the corrosion resistance
of passive metals and alloys. Generally speaking, to have a good
corrosion resistance an alloy should exhibit a low value of Ep and a high
value of Eb.
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Figure Variation of passivation potential for Fe-Cr alloys with pH in
sulfate solutions
Generally, the oxides of less noble metals exhibit a lower standard
potential of formation. These metals passivate
spontaneously in the presence of protons. Furthermore, many oxides
exhibit good chemical stability in acidic environments. This explains
the higher corrosion resistance of metals such as titanium, tantalum
and chromium.
According to the relation,
the value of Erev,oxide decreases by 59 mV per pH unit, regardless of
the stoichiometry of the oxide formed; this is because the number of
charges does not appear in the equation.
The condition Ep ≥ Erev,oxide, suggests that the passivation potential
also decreases with increasing pH.
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Figure confirms this behavior: When the passivation potentials of iron, chromium
and their alloys in sulfuric acid are plotted as a function of the pH, a straight line
is obtained with a slope of –59 mV/pH. Similar results have been found for other
metals.
In this figure, for zero pH (pH = 0; Erev,oxide = E°), the passivation potentials of
chromium and iron do not match the standard potentials of the oxides Cr2O3 and
Fe2O3 listed in Table 6.8.
One explanation is that kinetic limitations lead to a higher passivation potential
than predicted by thermodynamics. Another reason could be that the listed
standard potentials were measured on bulk samples. The extreme thinness of
passive films could influence their thermodynamic properties.
In addition, their composition does not always correspond to a simple
stoichiometry.
For example, chromium-iron alloys form passive films containing both iron and
chromium cations and their passivation lie between those of iron and chromium.
Then again, they exhibit the same pH dependence as the pure metals.
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Figure 6 Anodic polarization curves of Fe, Ni and Cr in 0.5 M H2SO4
In acidic media, the metals iron, nickel and chromium
have passivation current densities that increase in the
order Cr < Ni < Fe.
In Figure 6, the anodic polarization curves for the three
metals in 0.5 M sulfuric acid (25°C)are compared.
Chromium has lower values of both ip and Ep than the
other two metals.
By alloying increasing amounts of chromium to steel one
therefore improves the corrosion resistance.
Experience shows that above a chromium concentration
of 12 to 13%, a steel passivates spontaneously in
contact with aerated water. It becomes "stainless“,
meaning it does not rust easily
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Figure Corrosion potentials of Fe-Cr alloys in aerated 4% NaCl as a function of
their chromium content. The measured average rate of corrosion in salt spray
tests is also shown
Corrosion potentials of iron-chromium alloys in an
aerated solution of 4% NaCl together with the corrosion
rate measured in a salt spray test.
Spontaneous passivation above a chromium content of 8
to 12%, leads to a rise in corrosion potential and to a
drop of corrosion rate.
The magnitude of the passivation current density
depends on different factors:
the kinetics of active dissolution;
the mass transport of the dissolution products;
the pH of the electrolyte;
the water content of the electrolyte.
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