Corrosion and Environmental Degradation of Materials-1.pdf
1. Corrosion and Environmental Degradation of Materials
(Subject Code: ME3103, No. of Credits: 4)
Syllabus:
Review of electrochemical Principles. –Faradays laws-Electrode potentials –Cathodic and anodic reactions- polarization over voltage. Current
efficiency, throwing power and its evolution, electro plating of Cu, Ni, Cr, Zn and alloy Plating. Structure and properties of electrodeposits,
Testing methods of electro deposits Electrochemical aspects of Corrosion, Corrosion Cells/Electrochemical Cells, Concentration Cells,
Temperature Cells; Determination of electrode potential, Thermodynamic aspects, Nernest equation Introduction, Principles of Corrosion,
Exchange Potential Theory, Pourbauix Diagram, Passivity diagrams, Polarization resistance, Linear Polarization Technique, Introduction,
classification, forms of corrosion. Uniform corrosion, galvanic corrosion, and galvanic series. Beneficial applications of galvanic corrosion, Pitting
corrosion, season cracking, dezincification, Crevice corrosion, stress corrosion cracking, Intergranullar corrosion, weld decay, Knife-line attack,
Erosion corrosion, frettling corrosion, Corrosion protection methods, selection of materials for corrosion services, selection of environment-use of
inhibitors, surface protection methods including painting, metallic coating. Cathodic protection, sacrificial anode. High Temperature Corrosion,
Oxidation, Pilling Bed-worth Ratio; practical examples of high temperature oxidation
TEXT BOOKS;
1. Principles of Electroplating and Electroforming - William Blum
2. Corrosion Engineering-Mars G. Fontana
REFERENCES;
1. Material science- Van Vlack
2. Electroplating Basic principles and practice - Kanan. N (Elsevier) 2004
3. Elements of Physical Metallurgy – a. guy
4. Corrosion and Protection – Einaravrdet (Spinger) 2004
2. Corrosion and Environmental Degradation of Materials
(Subject Code: ME3103, No. of Credits: 4)
Detailed Syllabus:
Unit-I: Introduction and basics: Importance of corrosion, concept of classification of corrosion ,concept of Degradation process – Mechanical
and chemical process, concept of electrochemical cells ,concept of corrosion cell formation ,concept of free energy and cell potential and EMF ,
Nernst equation , single electrode potential , concept of reference electrodes, Galvanic series and EMF series-applications in corrosion.
.Unit-II: Thermodynamic Aspects: Pourbaix diagrams (EH-PH diagram) and fundamental aspects, construction of EH-PH diagrams for different
systems,Concept of corrosion rate expressions, Faradays law, cathodic - anodic reactions, concept of polarization or over potential and
classification,Concept of exchange current density, Butlor – Volmer equation.
Unit-III: HEADING: Kinetic Aspects & corrosion rate determination: Concept of Evans diagrams(Mixed potentials) and basics ,
construction of Evans diagrams of corroding metals in different solutions and Bimetallic couples , Mixed potential theory- applications ,concept
of Activaation and diffusion control process , Corrosion rate measurement methods ,Linear polarization and Tafel methods.
Unit-IV: Passivity & Forms of corrosion: Faradays passivity experiment, passivity and influencing paramaters and definitions, passivity
diagrams. Theories of Passivation, factors and conditions for passivity ,Design of corrosion resistant alloys , Critical current density,Concept of
Galvanic corrosion and beneficial applications , Different forms of corrosion and mechanisms , crevise corrosion ,stress corrosion cracking ,pitting
and inter granular corrosion concept of Hydrogen attracts and hydrogen embrittlement and blistering.
Unit-V: Prevention of corrosion: Corrosion prevention principles , concept of Material selection , concept of inhibitors and classification ,
concept of cathodic protection - principles, classification, cathodic protection - influencing factors and monitoring , Sacrificial anode, High
Temperature Corrosion, Oxidation, Pilling Bed-worth Ratio, practical examples of high temperature oxidation.
Unit-VI: High temperature corrosion, conversion coating and chemical degradation of non metallic materials: Oxide defect structures, semi
conducting oxides, wagner – Hauffe valence approach , Doping of oxides ,catastropic and internal oxidation ,High temperature alloy design
,alloying elements and major defects , concept of conversion coating ,chemical degradation of non metallic materials like ceramics, plastics and
rubbers, Bond rupture ,swelling and weathering.
References:
1) M. G. Fontana, Corrosion Engineering
2) R. W. Revie & H. H. Uhlig, Corrosion and corrosion control,
3) Philip A. Schweitze, Fundamentals of Corrosion
4) E.E. Stansbury & R.A. Buchanan, Fundamentals of Electrochemical Corrosion
5) Zaki Ahmad, Principles of Corrosion Engineering and Corrosion Control
3. Corrosion and Environmental Degradation of Materials (ME3103)
Instructor Name: Dr P Justin
Units of
syllabus
Study Material
(by Dr. P. Justin)
Video Lectures
(by
Prof.A.S.KHANNA,
IIT, Bombay
Assignment
&
Problems
Additional
Assignments &
Problems
Examination details
Unit-I
&
Unit-II
Lecture1:
Technological
importance of corrosion
and study
(15 Pages)
Video Lecture 1
Assignment &
Problems-1
worked-out-
problems
Assignments-1
Assignments-2
Mid 1
(20 Marks
Descriptive)
Syllabus
covered within
first 25 days of
instruction
period
EST
(20 Marks
Objective &
40 Marks
Descriptive)
Syllabus
covered in
total 80 days
of instruction
period in a
semester
Lecture2:
Electrochemical
principles of corrosion
(19 Pages)
Video Lecture 2
Assignment &
Problems-2
Lecture 3:
Concept of free energy
(21 Pages)
Video Lecture 2
Assignment &
Problems-3
Lecture 4:
Concept of single
electrode potential (14
Pages)
Video Lecture 2
Assignment &
Problems-4
Lecture 5:
EMF and Galvanic series
(14 Pages)
Video Lecture 2
Assignment &
Problems-5
Lecture 6:
EH-PH diagram-
Fundamental aspects (12
Pages)
Video Lecture 2
Assignment &
Problems-6
Lecture 7: Construction
of EH-PH diagram (12
Pages)
Assignment &
Problems-7
4. Lecture 8: Corrosion rate
Expression
(8 Pages)
Video Lecture2, 7
Assignment &
Problems-8
Lecture 9:
Electrode- solution
interface
(24 Pages)
Video Lecture 4
Assignment &
Problems-9
Lecture 10:
Exchange current density
(16 Pages)
Video Lecture 4
Assignment &
Problems-10
Unit-III and
Unit-IV
Lecture 11:
Mixed potentials –
concepts and basics
(12 Pages)
Video Lecture 4
Assignment &
Problems-11
Assignments-3
Assignments-4
Assignments-5
Mid 2
(20 Marks
Descriptive)
Syllabus
covered within
second 25 days
of instruction
period
Lecture 12:
Mixed potentials –
Bimetallic couples
(8 Pages)
Video Lecture 4
Assignment &
Problems-12
Lecture 13:
Mixed potentials-
Activation and diffusion
controlled process
(9 Pages)
Assignment &
Problems-13
Lecture 14:
Corrosion rate
measurements
(20 Pages)
Video Lecture 2,7
Assignment &
Problems-14
Lecture 15:
Passivity and influencing
parameters
(13 Pages)
Video Lecture 7
Assignment &
Problems-15
5. Lecture 16:
Passivity :Design of
corrosion resistant alloys
(18 Pages)
Assignment &
Problems-16
Lecture 17:
Different forms of
corrosion
(21 Pages)
Video Lecture 3,7
Assignment &
Problems-17
Lecture 18:
Pitting and intergranular
corrosion
(16 Pages)
Video Lecture 3,7
Assignment &
Problems-18
Lecture 19:
Stress corrosion cracking
(18 Pages)
Video Lecture 3
Assignment &
Problems-19
Lecture 20:
Hydrogen attack
(17 Pages)
Video Lecture 3 Assignment &
Problems-20
Part of Unit-
V and Unit-
VI
Lecture 21:
Principles of corrosion
prevention-Material
selection
(34 Pages)
Assignment &
Problems-21
Mid 3
(20 Marks
Descriptive)
Syllabus
covered within
third 25 days of
instruction
period
Lecture 22: Principles of
corrosion prevention-
Inhibitors
(22 Pages) Video Lecture 5
Assignment &
Problems-22
Lecture 23:
Cathodic protection-
principles and
classification
(6 Pages)
Video Lecture
5,10,11
Assignment &
Problems-23
Lecture 24:
Cathodic protection- Video Lecture 5,10
Assignment &
Problems-24
6. Influencing Factors and
Monitoring
(6 Pages)
Assignments-6
Assignments-7
Assignments-8
Lecture 25:
High temperature
corrosion (7 Pages) Video Lecture 8,12
Assignment &
Problems-25
Lecture 26:
Oxide defect
structure
(9 Pages)
Assignment &
Problems-26
Lecture 27:
Catastropic oxidation and
internal oxidation
(7 Pages)
Assignment &
Problems-27
Lecture 28:
High temperature
Alloy design
(7 Pages)
Assignment &
Problems-28
Lecture 29:
Conversion coating-1
(Phosphate, Chromate &
Oxide)
(15 Pages)
Assignment &
Problems-29
Lecture 30:
Conversion coating-2
(Anodizing)
(32 Pages)
Assignment &
Problems-30
Lecture 31:
Chemical degradation of
non metallic materials
(8 Pages)
Assignment &
Problems-31
8. Corrosion is the degradation of useful properties of a metal by
chemical or electrochemical reaction (mostly) with an environment,
in which an anodic and cathodic reaction takes place at equal rate but
at different site
2
Except noble metals like Au, Pd and Pt, all
other metals undergo corrosion
9. In nature, metals are not found in free state due to their reactivity..
But metals exist as its oxides, sulphides, sulphates, or carbonates,
which are thermodynamically stable states of the metals. The metals
are extracted from these minerals by spending huge amount of
energy. So, they develop a natural tendency to revert to their most
stable forms.
Cause of corrosion:
3
Refining-corrosion cycle:
10. Mechanical degradation:
Corrosion reduces the thickness of the metal, which causes
in loss of mechanical strength and failure of the structure.
Efficiency of the machine is reduced due to corrosion.
Because of corrosion, pipes are blocked and pumps are
difficult to operate. It also damages boilers.
Corrosion seriously shortens the predicted design life.
Chemical Impact:
4
Chemical Impact:
Buildings and other historic monuments are damaged due to
corrosion (e.g.Taj Mahal).
Corrosion products could make sanitizing of equipment more
difficult, eg. milk and dairy product plant.
Corrosion of steel, leads to Cr3+ or Cr6+ ions releasing into
drinking water.
11. Uniform or general corrosion:
It is the uniform thinning of a metal without any localized attack.
Corrosion does not penetrate very deep inside. Uniform corrosion is more likely
to take place in acidic environments (more severe in the case of SOx).
rusting of iron or structural steels in open air or tarnishing
of silver or electrical contacts or “Fogging ” of nickel
Corrosion of offshore drilling platforms or galvanized steel
stairways or underground pipes
high - temperature oxidation of metals
5
Local corrosion:
In localized type of attack, rate of corrosion being greater at some areas than
at others. Here the anode is very small and fixed area of corroding part of metal. So
remaining area of the metal acting as cathode. Local corrosion is more likely to take
place in alkaline, neutral, or slightly acidic environments.
localized corrosion such as pitting, crevice corrosion, galvanic corrosion,
or stress corrosion cracking are more dangerous than uniform corrosion
SCC of austenitic stainless steels in acid chloride media, leading to crack
12. 1. Dry or chemical corrosion
2. Wet or Electrochemical corrosion
1. Dry or chemical corrosion:
Corrosion is due to direct attack of chemicals in dry condition.
Corrosion products are formed at the place of corrosion.
Corrosion is slow, depending on the nature of the chemical reaction.
6
Corrosion is slow, depending on the nature of the chemical reaction.
a) Oxidation Corrosion b) Corrosion by Gases:
It is due to direct attack of
oxygen on metals
Carbon dioxide, Sulphur dioxide, NOx,
Hydrogen Sulphide, Flourine, Chlorine, etc
Oxygen molecules are attracted to
the surface by Vander Wall Force
S- more corrosion
C- less corrosion
13. 2. Wet or Electrochemical corrosion
Corrosion is due to the formation of large number of galvanic cells in wet condition.
Corrosion occurs at anode whereas corrosion products are accumulated at cathode.
Corrosion is fast, due to electrochemical cells.
7
Important condition:
O2 and moisture
rise in temperature
rise in concentration of H+ ions
increases with O2 and moisture
Corrosion increases with
14. Sl.No. Wet or Electrochemical corrosion Dry or chemical corrosion:
1 Corrosion is due to the formation of large
number of galvanic cells in wet condition
Corrosion is due to direct attack
chemicals in dry condition
2 Corrosion occurs at anode whereas corrosion
products are accumulated at cathode
Corrosion products are formed at
the place of corrosion
3 Corrosion is fast, due to electrochemical Corrosion is slow, depending on the
8
3 Corrosion is fast, due to electrochemical
cells
Corrosion is slow, depending on the
nature of the chemical reaction
4 Explained by electrochemical reactions Explained by adsorption
mechanism
5 Corrosion takes place at active centre only
i.e., not uniform
Corrosion is uniform
16. direct replacement of corroded equipment, components, and
structures
labor and material cost to control the corrosion rate of equipment
and structures
Losses sustained by industry and by governments ~$ 276 billion in the
United States, or 3.1% of the GDP
Economic losses can be divided into direct and indirect losses
Direct losses:
10
and structures
additional costs incurred by the use of corrosion-resistant metals or
alloys instead of less-expensive carbon steel
Plant shutdowns: due to unexpected corrosion failures of equipment
lead to loss of production and consequently loss of profit
Contamination: Pb pipes were used to transport water until it was
found that the lead pick-up in the water caused Pb poisoning in humans
Indirect losses:
17. Loss of product: leakage of gasoline or uranium compound due to
corrosion pipes or storage of tanks
Loss of efficiency: overdesign (additional thicknesses of vessel shells)
and the corrosion products (scales) decrease the heat-transfer rate in heat
exchangers
Environmental damage: leakage of dangerous chemicals
Sudden failure can cause explosions and fire, release of toxic products
11
and collapse of structures.
Safety is a critical consideration in the design of equipment for nuclear
power plants and for disposal of nuclear wastes.
Medicals metals used for hip joints, screws, heart valves, etc – high
reliability is of paramount importance.
Corrosion products could make sanitizing of equipment more
difficult,eg. milk and dairy product plant.
18. Loss of metal by corrosion is a waste not only of the metal, but also of
the energy, the water, and the human effort that was used to produce and
fabricate the metal structures
Corrosion destroys the aesthetic appeal of the product and damages the
product image
12
Corrosion seriously shortens the predicted design life
Corrosion Engineering, M. G. Fontana
Corrosion and corrosion control, R. W. Revie H. H. Uhlig
Principles of Corrosion Engineering and Corrosion Control, Zaki Ahmad
Fundamentals of Corrosion, Philip A. Schweitzer
References
19. 1. Corrosion of metals involves
(a) Physical reactions (b) Chemical reactions (c) Both (d) None
2. The following factors play vital role in corrosion process
(a) Temperature (b) Solute concentration (c) Both (d) None
3. Following equation is related to corrosion rate
(a) Nernst equation (b) Faraday’s equation (c) Either (d) Neither
4.The following influences deterioration of polymers
(a) Weather (b) Radiation (c) Temparature (d) All
5. Passivity is not reason for inertness of the following
(a) Au (b) Al (c) Ti (d) Ni
13
(a) Au (b) Al (c) Ti (d) Ni
6. When Pt and Co are electrically connected, which one gets corroded
(a) Pt (b) Co (c) None (d) Can’t decide
7. Following is not the main form of polymer deterioration
(a) Corrosion (b) Swelling and Dissolution (c) Weathering (d)
Scission
8. Main form of ceramic degradation
(a) Corrosion (b) Weathering (c) Dissolution (d) Swelling
9. Corrosion is a process of
(a) reduction (b) oxidation (c) ozonolysis (d) electrolysis
21. 1. Discuss the five important consequences of corrosion.
2. Discuss the economic aspect of corrosion
3. Describe with an example how corroded structures can lead to
environmental pollution.
4. What is the relationship between depletion of natural resources
and corrosion?
5. Explain how corrosion can be considered as extractive
metallurgy in reverse.
6. Name three cities in Southeast Asia and the Middle East which
15
6. Name three cities in Southeast Asia and the Middle East which
have the most corrosive environment.
7. State two important corrosion websites.
8. What is corrosion? Why do metals corrode?
9. Why most of the metals are found in the ore form and not in the
pure form? Explain.
www.intercorrosion.com
www.learncorrosion.com
www.nace.org
www.iom3.org
Websites
23. In a discharging battery or a galvanic cell, the anode is the
negative terminal because it is where the current flows into
the device . Here oxidation takes place.
This inward current is carried externally by electrons moving
outwards, negative charge moving one way constituting positive
current flowing the other way..
Anode:
In a discharging battery or a galvanic cell the cathode is
the positive terminal since that is where the current flows
Cathode:
2
the positive terminal since that is where the current flows
out of the device. Here reduction takes place.
This outward current is carried internally by positive ions
moving from the electrolyte to the positive cathode
(chemical energy is responsible for this uphill motion).
Electrolyte: KCl, NaCl, Nafion, etc
Salt bridge: to complete circuit (KCl in agar-agar)
25. Example 2: Rusting of iron
4
Anodic reaction
)
(
e -
2
Fe 2
Fe Oxidation
+
+
→
26. All metals above hydrogen in electrochemical series can show this type of corrosion
Case 1: Evolution of H2
The hydrogen ions (H+) are formed due to the acidic environment and the following
reaction occurs in the absence of oxygen.
Cathodic reaction: Two possibilities
5
All metals above hydrogen in electrochemical series can show this type of corrosion
In hydrogen evolution type of corrosion, anodic area is large as compared to its cathodic area
27. Case 2: Absorption of O2
This type of corrosion takes place in neutral or basic medium in the presence of
oxygen.
The small scratch on the surface creates small anodic area and rest of the surface
acts as cathodic area.
6
If oxygen is limited, anhydrous magnetite,Fe3O4 [a mixture of
(FeO + Fe2O3)], is formed, which is brown-black in colour.
In O2 absorption type of corrosion, anodic area is small as
compared to its cathodic area.
28. Corrosion occurs at anode, but rust is deposited at or near cathode.
This is because of the rapid diffusion of Fe2+ as compared to OH.-
7
29. 1. Dissimilar electrode cells or Galvanic cells
2. Concentration cells:
a. Salt concentration cell
b. Differential aeration cell
3. Differential temperature cells
1. Dissimilar electrode cells or Galvanic cells:
A dissimilar electrode cell or galvanic cell is formed when two
8
A dissimilar electrode cell or galvanic cell is formed when two
dissimilar metals differing in potential are joined together or
due to the heterogeneity of the same metal surface.
a copper pipe connected to a steel pipe
a bronze propeller in contact with the steel hull of a ship
a steel plate with copper rivets
Examples:
30. Dissimilar electrode cell formation by joining of an old pipe with a new pipe
Some more examples:
9
Old pipe has a more positive potential (cathode)
than new pipe due to films formation on the old pipe
31. 2. Concentration cells:
Concentration cells are formed when the electrodes are identical but
are in contact with different concentration of solutions or differential
aerations or different soil resistivity or different stress.
Two types of concentration cells:
a. Salt concentration cell
b. Differential aeration cell
10
2.a. Salt concentration cell:
Salt concentration cells are formed when the electrodes are identical
but are in contact with different concentration of solutions.
33. 2.b. Differential aeration cell:
Differential aeration cells are formed when the electrodes are identical
but are in contact with differential aerations i.e., different oxygen
concentrations.
12
36. Differential aeration cells can also cause pitting damage under rust
and at the water line — that is, at the water – air interface
pitting damage under rust
15
Water - line corrosion
37. 3. Differential Temperature Cells:
Differential temperature cells are formed when electrodes of the same
metal, each of which is at a different temperature, are immersed in an
electrolyte of the same initial composition.
Such a situation may arise in practice in components of heat
exchangers, boilers, and similar heat transfer equipment.
16
Polarity developed in an electrode varies from system to system. For
a copper electrode in a copper sulfate solution, the electrode at
the higher temperature is the cathode; but for lead, the situation is
just the reverse. For iron immersed in dilute aerated sodium chloride
solutions, the hot electrode is initially anodic to the colder metal, but
the polarity may reverse with the progress of corrosion.
42. -the driving force of a any chemical reaction
We know that for any reversible reaction,
At standard state,
2
A chemical reaction at constant temperature and pressure will takes place only if
there is an decrease in the over all free energy of the system during the reaction.
43. The tendency for any chemical reaction to go, including the reaction of a metal
with its environment, is measured by the Gibbs free - energy change, ΔG .
3
The large negative value of ΔG indicates a pronounced
tendency for magnesium to react with water and oxygen.
The reaction tendency of cu is less when compared to Mg, i.e., the
corrosion tendency of copper in aerated water is not as pronounced as
that of magnesium.
44. The free energy is positive, indicating that the reaction
has no tendency to go at all; and gold, correspondingly,
does not corrode in aqueous media to form Au(OH)3
It should be emphasized that the tendency to corrode is not a measure of
reaction rate.
4
If ΔG is negative, the corrosion rate may be
rapid or slow, depending on various factors
If ΔG is positive, corrosion will not go at all
under the particular conditions described
It should be emphasized that the tendency to corrode is not a measure of
reaction rate.
45. Electrochemical cells generate an electrical energy due to electrochemical reactions.
In any electrochemical reaction, there exists an integer correspondence between the
moles of chemical species reacting and number of moles of electrons (n) transferred. To
5
moles of chemical species reacting and number of moles of electrons (n) transferred. To
convert this molar quantity of electrons to a total charge (Q), we must multiply the
number of electrons (n) with Avogadro’s number (NA=6.022 × 1023 electron mol-1) and the
charge per atom (q =1.602 × 10-19 C electron-1).
From combining equation 3 and 4,
46. Here, F, is the Faraday’s constant and it is really the product of NA × q (6.022 × 1023
electron mol-1 × 1.602 × 10-19 C electron-1 = 96500 C mol-1). Since is F large, a little
chemistry produces lot of electricity.
This maximum electrical work can be done only through a decrease in Gibbs free
energy at constant temperature and pressure.
6
47. Nernst equation express the emf of a cell in terms of activities of
products and reactants of the cell.
Let us consider a reversible cell reaction of type
When ‘n’ faraday of electricity is passed through the cell, the decrease in free
energy is given by
7
energy is given by
Where k is called equilibrium constant and at any arbitrary condition is given by
48. Now, at any arbitrary condition is
arbitrary condition
standard state
Divide equation 5 by nF
8
Divide equation 5 by nF
The equation 6 and 7 is called as Nernst equation
49. To construct Pourbaix diagram, which represents
conditions of thermodynamic equilibrium for some reaction
To calculate theoretical open – circuit potential
To correlate polarization with potential
To calculate cell potential
Application on corrosion:
Calculation of cell potential:
Reversible cell:
9
Net cell reaction:
Reversible cell:
52. A cell is said to be reversible if the following two conditions are fulfilled.
1. The chemical reaction of the cell stops, when an exactly equal amount of
opposing emf is applied. In Daniell cell, the following reaction will stop, if
exactly1.1 V emf is applied from external source.
2. The chemical reaction of the cell is reversed and the current flows in opposite
12
2. The chemical reaction of the cell is reversed and the current flows in opposite
direction, when slightly higher amount of opposing emf is applied. In Daniell cell,
the above reaction reversed and the current flows in opposite direction, if slightly
higher than 1.1 V emf is applied from external source.
Examples:
Daniell cell, Li-ion battery, Lead – storage battery
used in automobiles. Edison cell; Ni-Cd cells used in
calculators and flash lamps.
53. A cell is said to be irreversible, if the two above conditions of
reversible cells are not fulfilled.
Examples:
Leclanche cell (ordinary flash light battery), Zn-Hg cell used in
cameras, clocks, hearing aids and watches. A cell consisting of zinc
and copper (or Ag) electrodes dipped into the solution of sulphuric
acid is irreversible.
13
acid is irreversible.
63. When a metal is immersed in an electrolyte, a dynamic equilibrium is established across
the interface with a potential difference between the metal and an electrolyte.
2
64. The metal is left with a negative charge on its surface and
its positively charged metal ions, Mn+
(aq), in the electrolyte
Now, metal ions in the electrolyte are attracted back
towards the metal surface
Thus, a potential difference and dynamic equilibrium
between the metal and the solution is established
The tendency of a metal electrode to loose or gain an electrons
when it contact with its own ions is called electrode potential.
3
The tendency of a metal electrode to loose or gain an electrons
when it contact with its own ions at standard state (Concentration =
1M, Pressure = 1 atm, Temperature = 25oC )is called standard
electrode potential.
The tendency of a metal electrode to loose an electrons when it
contact with its own ions at standard state (Concentration = 1M,
Pressure = 1 atm, Temperature = 25oC )is called standard oxidation
potential.
65. The tendency of a metal electrode to loose an electrons when it
contact with its own ions at standard state (Concentration = 1M,
Pressure = 1 atm, Temperature = 25oC )is called standard reduction
potential.
In the measurement of emf, interest is usually focused on the
reaction that occurs at only one electrode. Measurements of this kind are
made by using an electrode, called a reference electrode.
4
Important requirements: It should have a stable electrochemical potential, regardless of
the environment in which it is used.
It should have low temperature coefficient.
The potential should be stable for long time.
Precautions:
The reference electrode should be placed as close as possible to
the protected structure to minimize the internal resistance (IR).
There should be no gas bubble on the electrolyte path in order
to ensure zero impedance.
66. The standard hydrogen electrode consists of a platinized platinum foil
immersed in a solution containing hydrogen ions and saturated with hydrogen gas as
shown below.
5
Platinum foil should be is immersed in arsenic free acid, and H2 gas
should be free from O2 and CO.
Slowly, the air is displaced by hydrogen and the reversible potential
is achieved.
67. Electrode representation:
Electrode reaction:
Nernst Equation:
6
Limitations
It cannot be used in oxidizing media (reducible materials like nitrates, magnates, permanganate,
chlorates and perchlorates).
It cannot be used in solutions containing metals ions like Ag+, Pb2+, Hg22+ and Cd2+, since metals
may be displaced by H+ ions.
If a current is withdrawn from the electrode, the electrode acts as an anode because of the
ionization of gas molecules.
The electrode is fragile and delicate to handle.
68. The silver – silver chloride, can be prepared by chloridizing silver wire in
dilute hydrochloric acid. The chloridized silver wire is sealed into a glass
tube and immersed in a solution containing chloride ions as shown below.
.
7
69. Electrode representation:
Electrode reaction:
Nernst Equation:
8
At 25O C, the potential of silver – silver chloride electrodes depend only on
the concentration of Cl- ions as shown below.
Concentration of KCl Potential* vs. SHE
0.1 M 0.288 V
3.5 M 0.202 V
Saturated (~4 M) 0.199 V
*The values cited neglect the liquid - junction potential (LJP) at the KCl
boundary. For example, in the case of strong acids, the LJP increases to
an average of several mV.
70. The calomel electrode basically consists of a platinum wire dipped into pure
mercury and covered with a paste of mercurous chloride (Hg2Cl2) and mercury as
shown below. The paste is in contact with a solution of potassium chloride, which
acts as a salt bridge to the other half of the cell.
.
9
71. Electrode representation:
Electrode reaction:
Nernst Equation:
10
At 25O C, the potential of calomel electrodes depend only on the
concentration of Cl- ions as shown below.
Concentration of KCl Potential* vs. SHE
0.1 M 0.334 V
3.5 M 0.280 V
Saturated (~4 M) 0.242 V
*The values cited neglect the liquid - junction potential (LJP) at
the KCl boundary. For example, in the case of strong acids, the LJP
increases to an average of several mV.
72. The copper–copper sulfate electrode basically consists of a copper plate immersed
in a solution containing copper sulfate and copper sulfate crystals placed in a non-
conducting holder with a porous plug as shown below.
11
73. Electrode reaction:
A saturated solution of 1.47 M CuS04 at 25°C is used and the potential
value of copper–copper sulfate electrode is 0.316 V vs. SHE.
Nernst Equation:
Advantage:
12
This is a reference electrode robust, stable and easy to construct. It is used
mainly in cathodic protection measurements, such as the measurement of pipe-to-
soil potential.
It has a lower accuracy than other electrodes used for laboratory work.
Limitations:
74. 0.0 V recorded vs SCE should be reported as 0.244 V vs. SHE
13
0.0 V recorded vs SCE should be reported as 0.244 V vs. SHE
(To covert SCE to SHE, you have to add 0.241 V to SCE potential)
0.244 V recorded vs SHE should be reported as 0.0 V vs SCE
(To covert SHE to SCE, you have to subtract 0.241 V from SHE potential)
The standard electrode potential of the Mercury/Mercury Oxide half-cell is
+0.098 volts vs. NHE in 20% (4.24 KOH).
77. In EMF series, the metals or elements are arranged in their increased
order of electrochemical activities based on the equilibrium potential of
a metal, which is in contact with its own ions of unit activity.
Applications:
The metal with a more negative potential is generally the anode, and the metal
with a less negative potential, the cathode.
2
with a less negative potential, the cathode.
The metals with high positive potentials are recognized as metals with good
corrosion resistance.
The metal or nonmetal's ions, which possess high positive reduction potentials
have the tendency to accept electrons readily, i.e., readily reduced
Metal ions above hydrogen are more readily reduced than the hydrogen ions.
An oxidizing agent takes electrons from another species, and thus it is reduced
and the reducing agent gives electrons to another species, and thus it is oxidized.
A less electropositive metal would displace a more electropositive metal from one
of its salts in aqueous solution. Eg., Zn dissolves and cu would be deposited.
79. Limitations:
EMF series lists only metals (has little engineering application). Alloys
not included .
Electrode potentials listed are calculated from thermodynamic principles
(corrosion potentials are more relevant).
Prediction about galvanic coupling can only be made when the two metals
forming the galvanic couple have their ionic activities at unity.
Predicts only tendency to corrode (Role of passive films and oxidation
kinetics not predicted).
Effect of environment not predicted (Eg: Sn – Fe couple as in Tin cans)*.
4
Effect of environment not predicted (Eg: Sn – Fe couple as in Tin cans)*.
*Internally tinned (tin-coated) steel cans are used to preserve
vegetable and fruit juices. Tin (Sn) is nobler to iron (Fe) in the EMF
series. Such a cathodic protection of iron by tin is however only
limited. Because many food constituents such as organic acids can
combine with Sn2+ to form soluble tin complexes, resulting in lowering the
activity of stannous ions. The polarity of Fe– Sn couple can reverse
under these conditions.
80. 5
or, the ratio (Sn2+/Fe2+) must be less than 5 ×1 0−11 for tin to
become active to iron.
81. In Galvanic series, the metals and alloys are arranged in their increased
order of electrochemical activities based on the practical measurement
of corrosion potential at equilibrium in seawater. It enables the
corrosion engineer to predict the corrosivity of metals and whether the
coupling of two metals will be compatible or not.
6
84. The Galvanic Series is an arrangement of metals and alloys in accord
with their actual measured potentials in a given environment
Practically measured potentials vs reference electrode. It includes
steady - state values in addition to truly reversible values.
The galvanic series indicates that alloys can be coupled without being
corroded.
Some metals occupy two positions in the Galvanic Series, depending on
whether they are active or passive*.
*Only in active state, there is a true equilibrium attained between metal
and its own ions. But in passive state, there is non-equilibrium state
Merits:
9
*Only in active state, there is a true equilibrium attained between metal
and its own ions. But in passive state, there is non-equilibrium state
between metal and its own ions because of surface films.
The role of passive films and oxidation kinetics are considered.
*Eg 1: Aluminium exhibits higher corrosion resistance due to Al2O3
layer present on surface.
Eg 2: Chromium exhibits stable Cr2O3 layer and is used as alloying
element for corrosion resistance in stainless steels.
85. 1. Which of the following pairs of metals would show the highest rate of corrosion in
seawater?
(a) Copper and steel (b) Copper and aluminum (c) Copper and brass (d) Copper and
zinc
2. Referring to the galvanic series of some commercial metals and alloys in seawater,
mark the condition which would lead to minimum corrosion by galvanic coupling.
(a) Magnesium and aluminum plates (b) Silver and copper (c) 70-30 brass with pure
copper (d) Coupling of 18-8 steel active with chromium stainless steel, 13% Cr
(active)
3. If the free energy of a reversible process is negative, it implies that
(a) the cell reaction is spontaneous (b) the cell reaction is not spontaneous (c)
the cell reaction proceeds from right to left (d) no reaction takes place at all
10
the cell reaction proceeds from right to left (d) no reaction takes place at all
4. A galvanic series is
(a) a list of alloys arranged according to their corrosion potentials in a given
environment (b) a list of metals and alloys according to their corrosion potentials
in a given environment (c) a grouping of metals and alloys based on their ability to
get oxidized in a stated environment
(d) a list of standard electrode potentials of alloys or metals arranged in order of their
values
5. The metal at the top of electrochemical series is
(a) most stable (b) ) more noble (c) less active (d) more active
6. Food stuff containers should not be
(a)Galvanized (b) tinned(c) electroplated (d) cladded
86. 1. In the galvanic series why active steel is placed far away from the
passive steel and not bracketed together?
2. In the following couples, which one would form the anode and
which one the cathode?
(a) Active and passive steel (b) Zinc and aluminum (c) Copper
and iron (d) Stainless steel and brass
3. State the limitations of the emf series and the advantages of
galvanic series for an engineer.
4. Distinguish between emf series and galvanic series.
11
4. Distinguish between emf series and galvanic series.
91. Potential-pH diagrams are also called Pourbaix diagrams after
the name of their originator, Pourbaix (1963), a Belgium
electrochemist and corrosion scientist.
Characteristic features:
Pourbaix diagrams is a graphical representations of the stability of a metal and its
corrosion products as a function of the potential and pH (acidity or alkalinity) of the
aqueous solution.
2
aqueous solution.
The potential is shown on the vertical axis (y-axis) and the pH on the horizontal axis (x-
axis). The hydrogen and oxygen lines are indicated in Pourbaix diagrams by dotted line.
These diagrams are constructed from calculations based on the Nernst equation and the
solubility data for various metal compounds. The concentration of all metal ions is assumed
to be 10-6 mol per liter of solution.
By controlling potential (e.g., by cathodic protection) and/or by adjusting the pH in
specific domains identified using Pourbaix diagrams, it may be possible to prevent
corrosion from taking place.
92. The potential–pH diagram shows three clear-cut zones::
Immunity zone: Under these conditions of potential and pH, iron
remains in metallic form.
Corrosion zone: Under these conditions of potential and pH, iron
corrodes, forming Fe2+ or Fe3+ or HFeO2-
Passive zone: Under these conditions of potential and pH, protective
layers of Fe (OH)3 form on iron and further corrosion of iron does not
take place.
3
With reference to Pourbaix diagrams, corrosion prevention can be
achieved by lowering the electrode potential down to the zone of immunity,
raising the electrode potential up to the region of passivity, or raising the
pH or alkalinity of the solution so that a passive film is formed.
93. Advantages:
Predicting the spontaneous direction of reactions.
Estimating the stability and composition of corrosion products.
Predicting environmental changes that will prevent or reduce corrosion*.
Limitations:
It represent equilibrium conditions and hence cannot be used for predicting the
rate of a reaction.
4
rate of a reaction.
The corrosion products (oxides, hydroxides, etc.) are assumed to be protective
(due to passivity ) which is not always true .
The possibility of precipitation of other ions such as chlorides, sulfates, and
phosphates has been ignored.
The activity of species is arbitrarily selected as 10-6 g mol-1 which is not
realistic.
It deal only with metals (has little engineering application).
The pH at the metal surface may vary drastically because of side reactions, and
a prediction of corrosion based on the bulk pH of the solution may be misleading.
103. Identify the possible components of the system. (There can be hundreds.)
Get K , acid and base dissociation constants, redox potentials, etc.
2
Get Ksp, acid and base dissociation constants, redox potentials, etc.
Decide on concentrations, temperature, pressure, etc.
Identify the possible reaction of the system.
Apply the Nernst equation to the electron transfer reactions and fix
horizontal line, which is dependent solely on potential, but independent of pH.
Apply the equilibrium constant to the non-electron transfer reactions and
fix vertical line, which is dependent solely on pH, but independent of potential.
Connect both horizontal line and vertical line by a sloping line, which depends
on both EH and pH.
105. Reaction dependent only on EH, but independent of pH)
Apply Nernst equation,
Line 1: Horizontal line 1
The activity of the species is arbitrarily selected as 10-6 g mol-1 all the
species other than H+ and OH-. The activity of water is taken as 1 and
temperature is at 25oC.
4
The corresponding EH value is indicated and drawn as horizontal line on the Y-axis in
the Pourbaix diagram.
Below this line (i.e., at EH 0.77 V), ferric ions in solution is the stable phase and
above this line (i.e., at EH 0.77 V), ferrous ions in solution are stable. So,
corrosion is expected to take place without any protection afforded by a surface
oxide film in both the phases.
106. Apply Nernst equation,
Reaction dependent only on EH, but independent of pH)
Line 2: Horizontal line 2
5
The corresponding EH value is indicated and drawn as horizontal line on the Y-axis in
the Pourbaix diagram.
Below this line (i.e., at EH -0.617 V), Fe metal is the stable phase due to immunity
condition. Above this line (i.e., at EH -0.617 V), ferric ions in solution are stable, and
corrosion is expected to take place without any protection afforded by immunity
condition.
107. Solubility product, Ksp of Fe(OH)2 is
Reaction dependent only on pH, but independent of EH)
Line 3: vertical line 1
6
The corresponding pH value is indicated and drawn as vertical line on the x-axis in
the Pourbaix diagram.
To the right of this line (i.e., at pH 9.65), Fe3O4 is the stable phase; and this oxide, as
a protective fi lm, would be expected to provide some protection against corrosion. To
the left of this line (i.e., at pH 9.65), ferric ions in solution are stable, and corrosion is
expected to take place without any protection afforded by a surface oxide film.
108. Reaction dependent on both EH and pH
Line 4: Sloping line 1
This sloping line separating Fe 2+ from Fe2O3 represents the
equilibrium.
7
A slope of —0.777 is obtained, which is indicated in the figure. The slope shows a pH
and potential dependence.
To the right of this line, Fe2O3 is a stable phase that is expected to form a surface
oxide fi lm that protects the underlying metal from corrosion. To the left of this
line, Fe 2+ is a stable species in solution.
109. The hydrogen line represents the equilibrium as follows:
Hydrogen and oxygen electrode lines:
The hydrogen and oxygen are also shown in the diagram by the dotted lines.
(a) Hydrogen electrode line:
8
These two reactions are equivalent and their pH dependence of SHE is represented by:
110. The oxygen line represents the equilibrium as follows:
(b) Oxygen electrode line:
9
These two reactions are equivalent and their pH dependence of SHE is represented by:
Water is stable in the area designated by these two lines. Below the hydrogen line it is
reduced to hydrogen gas, and above the oxygen line it is oxidized to oxygen.
111. Some more discussions:
The horizontal line at − 0.617 V means that iron will not corrode below this value to
form a solution of concentration 10-6 M of Fe2+ ions. If a value other that 10-6 M is
used, the lines separating the phases are shifted.
The fields marked Fe2O3 and Fe3O4 are sometimes labeled “ passivation ” on the
assumption that iron reacts in these regions to form protective oxide films. This is
correct only if the passivity is accounted by a diffusion - barrier oxide layer with
highly adherent in nature.
10
Metals like aluminum and steel are known to resist corrosion because of development
of oxide films in the air.
112. The potential–pH diagram shows three clear-cut zones::
Immunity zone: Under these conditions of potential and pH, iron
remains in metallic form.
Corrosion zone: Under these conditions of potential and pH, iron
corrodes, forming Fe2+ or Fe3+ or HFeO2-
Passive zone: Under these conditions of potential and pH, protective
layers of Fe (OH)3 form on iron and further corrosion of iron does not
take place.
11
With reference to Pourbaix diagrams, corrosion prevention can be
achieved by lowering the electrode potential down to the zone of immunity,
raising the electrode potential up to the region of passivity, or raising the
pH or alkalinity of the solution so that a passive film is formed.
113. Advantages:
Predicting the spontaneous direction of reactions.
Estimating the stability and composition of corrosion products.
Predicting environmental changes that will prevent or reduce corrosion*.
Limitations:
It represent equilibrium conditions and hence cannot be used for predicting the
rate of a reaction.
The corrosion products (oxides, hydroxides, etc.) are assumed to be protective
(due to passivity ) which is not always true .
The possibility of precipitation of other ions such as chlorides, sulfates, and
12
The possibility of precipitation of other ions such as chlorides, sulfates, and
phosphates has been ignored.
The activity of species is arbitrarily selected as 10-6 g mol-1 which is not
realistic.
It deal only with metals (has little engineering application).
The pH at the metal surface may vary drastically because of side reactions, and
a prediction of corrosion based on the bulk pH of the solution may be misleading.
115. Corrosion involves dissolution of metal, as a result of which the
metallic part loses its mass (or weight) and becomes thinner. Corrosion
rate expressions are therefore based on either weight loss (Faraday’s law ) or
penetration into the metal.
The weight loss due to corrosion can be converted to average corrosion rate
(mpy) using Faraday’s law. According to the Faraday’s first law of electrolysis “the
mass of substances liberated/deposited at an electrode by electrolysis is directly
2
mass of substances liberated/deposited at an electrode by electrolysis is directly
proportional to the quantity of electricity passed”.
1
......
ZIt
=
m
electrochemical equivalent = mass of
substances produced by 1 ampere-second
of a current (1 coulomb)
According to the Faraday’s second law of electrolysis “when the same
amount of electricity is passed through different electrolytes, the mass of
substance liberated/deposited at the electrodes is directly proportional to their
equivalent weights”.
I = Current, A
t = time, s
m = mass of substances, g
Z = electrochemical equivalent
116. 2
......
1
1
1
1
n
m Z
M ∝
∝
3
......
2
2
2
2
n
m Z
M ∝
∝
Substituting for Z, from equation (2) into (1)
m1, m2 = masses of substances, g
M1, M2 = Molar masses, g mol-1
n1, n2 = number of electrons
Z1, Z2 = electrochemical equivalent
4
......
t
k
m I
n
M
=
5
......
t
F
1
m I
n
M
×
=
3
F n
Dividing equation (5) by the exposed area of the metal
6
......
nFA
MI
At
=
w
density)
(current
A
I
i
But =
7
......
nF
M
At
i
w
=
Penetration per unit time can be obtained by dividing equation (7)
by density of the metal or alloy
117. 8
......
rate,
orrosion
ρ
n
Mi
C
r
C ×
=
M = Atomic weight, g mol-1
i = current density, µA cm–2
n = number of electrons
ρ = density, g cm–3
C= constant = 0.129, if corrosion rate is in mpy
Conversion of corrosion current of iron (=1µA cm–2) into corrosion rate (mpy)
4
118. Density
mdd
.00144
0 ×
=
ipy
4
.
5
2
ipy×
=
y
mm
One mil is one
thousandth of an inch
5
DAT
W
534
=
(mpy)
year
per
n
penetratio
Mils
W = weight loss, mg
D = density of specimen, g/cm3
A = area of specimen, (in.2)
T = exposure time, hr
The corrosion rates of resistant materials
generally range between 1 and 200 mpy
123. 2
When a metal is immersed in an electrolyte, its positively charged metal ions, Mn+
(aq), goes in
to the electrolyte and there by the metal is left with a negative charges on its surface.
As more and more +ions are released, the metal surface becomes increasingly negatively
charged.
Now, metal ions in the electrolyte are attracted back towards the metal surface due to the
excess negative charges on its surface.
Thereby forming a two oppositely charged layers looks like a capacitor in an interface and is
called Helmholtz double layer.
Thus, a dynamic equilibrium is established across the interface and a potential difference is
created between the metal and an electrolyte.
124. Helmholtz double layer model:
In the absence of external current, the electrode has a
charged double layer, i.e., the capacitor is charged.
3
125. If the metal have large -ive charge, the
H2O molecules would be oriented with the
+ive ends (hydrogen) towards the metal, and -
ive ends (oxygen) towards the large +ive
charge . The water molecules sometimes
contain the specifically adsorbed anions.
The line drawn through the center of
these molecules is called the inner Helmholtz
plane (φ1). In inner Helmholtz plane, potential
changes linearly with the distance comprises
the adsorbed water molecules and the
specifically adsorbed anions (X1).
4
The Helmholtz double layer model is only applicable to a concentrated solution. In dilute
solution, Guoy and Chapman observed that the net charge in the compact double layer does not
balance the charge on the metal surface.
Thus an outer diffuse layer (Guoy-Chapman layer) contains excess cations or anions
distributed in a diffuse layer extending up to 1 μm from the outer Helmholtz plane in the bulk
solution. In this layer the potential varies exponentially with the distance (X).
Like that, the H2O molecules would be oriented with the -ive ends (oxygen) towards the
metal ions (cations) in the electrolyte and the locus of the electrical centers of the positive
charges is called the outer Helmholtz plane (φ2). In outer Helmholtz plane, potential changes
linearly with the distance comprises the hydrated (solvated) cations (X2).
126. When a metal electrode is in equilibrium with its electrolyte, a partial current of forward
reaction (if) and partial current of reverse reaction (ir) are precisely equal and opposite in
direction. There is no net current flow.
5
127. In the presence of external superimposing emf (a net current flows into
or from the metal surface), there is a potential drop across the double
layer.
Now, there is a deviation from the equilibrium condition and the electrode is said to be
polarized. The extent of polarization is measured by the change in the potential drop (∆E)
across the double layer.
The shift of potentials from their equilibrium value on application of an external current is
called ‘polarization’. The magnitude of the deviation is termed ‘overvoltage’ which is directly
proportional to the magnitude of the external current density and its direction.
6
128. The direction of potential always changes from equilibrium and opposes the flow of
current, whether the current is impressed externally or is of galvanic origin.
7
When current flows in a galvanic cell, for example, the anode is always more cathodic
in potential and the cathode always becomes more anodic. The potential difference
between the anode and cathode becomes smaller as current is increased.
129. In anodic polarization, there is a shift of the potential of an electrode in a
positive direction (noble) by an external current.
In cathodic polarization, there is a shift of the potential of an electrode in a
negative direction (active) by an external current.
8
negative direction (active) by an external current.
The variation of η with i is linear. The departure from the equilibrium is shown by the
over-potential (η).
130. 1. Activation polarization
2. Concentration polarization
3. Ohmic polarization (IR drop)
1. Activation polarization:
Activation polarization is caused by a slow electrode reaction. The reaction
at the electrode requires an activation energy in order to proceed. An
9
at the electrode requires an activation energy in order to proceed. An
activation energy in the form of potential is required for the reaction to
proceed.
The most important example is that of
hydrogen ion reduction at a cathode.
For this reaction, the polarization is called hydrogen over-potential.
Hydrogen evolution occurs in four major steps:
131. 10
Either the electron transfer step (step
2) or the formation of hydrogen molecules
(step 3) is deemed the slowest step in
the reaction sequence.
132. The relationship between reaction rate and change in potential
(overvoltage) is expressed by the Tafel equation:
where ηa is overvoltage polarization (in volts), and β is a constant, called the Tafel
constant (also expressed in volts), and is usually on the order of 0.1 V.
The exchange current density, io , represents the current density equivalent to the equal
forward and reverse reactions at the electrode at equilibrium.
11
forward and reverse reactions at the electrode at equilibrium.
The larger the value of io and the smaller the value of β , the smaller the
corresponding over potential.
Dissolution reactions (anodic) in corrosion are usually controlled by
activation polarization if the solution of ions is the probable rate-
controlling step. Hydrogen evolution reactions (cathodic reactions)
are controlled by activation polarization when the concentration of
hydrogen ions is high.
133. Factors Effect on activation polarization
Current density increased current density increases the ηa according to
Surface roughness increased surface roughness decreases the ηa
Temperature increased temperature decreases the ηa
Pressure increased pressure decreases the ηa
pH Over-voltage increase initially and decreases with increased
pH value.
12
Agitation no effect, because it is a charge transfer process involving
electrons and not a mass transfer
Adsorption of ions The hydrogen over-voltage is decreased by adsorption of
anions and increased by adsorption of cations.
134. 2. Concentration polarization (Diffusion or transport over-potential):
Concentration polarization is due to concentration changes near the
electrode/electrolyte interface caused by diffusion of ionic species in the
electrolyte. According to the Tafel equation
As long as there is no concentration built up, an E vs log I plot shows a
linearity and it is activation controlled.
13
The above assumption is often not correct especially for oxygen reduction because O2 takes
large time to diffuse in solution to the corroding interface, and metal ions also take a
definite time to cross the double layer. A mathematical expression for concentration
polarization (ηc) in volts involves limiting current density (il) is given by:
135. There is no question of concentration polarization when the supply of
reacting species is abundant. Hence, in metal dissolution reactions, its
effect is negligible as the supply of metal atoms for dissolution is
unlimited. On the other hand, for a hydrogen evolution reaction,
concentration polarization becomes significant in the solutions of low Η+
concentration.
More often, the reduction process is controlled by a combined polarization — that is,
activation polarization at lower reaction rates and concentration polarization at higher
reaction rates — as i approaches iL.
14
136. Factors Effect on concentration polarization
Agitation increased agitation decreases the ηc because the thickness
of the diffusion layer is decreases, and thereby rate of
diffusion of ions increases
Temperature increased temperature decreases the ηc because the
thickness of the diffusion layer is decreases
Velocity There is no concentration polarization at high velocity
15
(totally activation controlled).
Concentration of ionic
species.
The lower is the concentration of species, the greater would
be the concentration polarization (totally mass controlled).
Geometry The geometry of fluid flow and the design of the cell
(horizontal or vertical) affects concentration polarization.
137. 3. Ohmic polarization (IR drop or resistance polarization):
The effect of ohmic polarization is significant because the current flows
from the anode to the cathode through an electrolyte.
Painting the metal surface inserts a
high resistance into the corrosion
With an increasing resistance offered by the electrolytes, the magnitude of
the corrosion current decreases as shown by R1, R2 and R3 in the figure.
16
high resistance into the corrosion
circuit as illustrated by Fig
As polarization increases, corrosion decreases
147. At equilibrium, the rates of the anodic (oxidation) and cathodic (reduction)
processes are equal, and there is no net charge transfer. The magnitude of
current per area, where the rate of forward and reverse reactions are equal
and opposite is defined as the exchange current density (io).
At equilibrium, η = 0, E = Ee, inet = 0 (i.e., no net current)
But ia = -ic = io (i.e., the rate of forward and reverse reactions are equal and
opposite and the rate corresponds to the io)
2
Every reversible electrode reaction has its own exchange current density.
The io is a fundamental characteristic that can be defined as the rate of oxidation
or reduction of an electrode at equilibrium expressed in terms of current.
Exchange current density inversely proportional to polarization.
In corrosion studies, the exchange current density, io is referred to as the
corrosion current: io = icorr ∝ corrosion rate
148. Factors Effect on exchange current density
Metal composition It depends upon the composition of the metal or alloy and
the solution. (see the table). The exchange current density
of hydrogen evolution on platinum is approximately 10−2 A
cm−2 whereas on mercury and lead it is 10−13 A cm−2.
Surface roughness Large surface areas provide a high io. The io for H2 evolution
is 10−2 A cm−2 on platinized platinum whereas on smooth Pt
3
is 10 A cm on platinized platinum whereas on smooth Pt
it is 10−3 A cm−2.
Surface impurities The exchange current density is reduced by presence of
trace impurities, such as As, S, and Sb
150. Note that the value for the exchange current density of hydrogen evolution on platinum
is approximately 10−2 A cm−2 whereas on mercury and lead it is 10−13 A cm−2, eleven
orders of magnitude difference for the rate of this particular reaction, or one hundred
billion times easier on platinum than on mercury or lead !.
This is the reason why Hg is often added to popular alkaline primary cells to slowdown the
thermodynamically favored production of H2 gas. This is also why lead acid batteries
(car batteries) can provide power in a highly acidic environment in a relatively safe
manner unless excessive charging currents are used.
5
Magnitude of io will indicate as to whether the redox reaction is reversible or
irreversible. Lower io denotes higher overpotential, while higher io indicates lower
overpotential (which means the reaction tends towards reversibility).
151. The exchange current density can be determined experimentally from a plot
of η against log |i|, with the intercept yielding a value for i0 as shown in the
following figure.
6
152. The relationship between current density and potential of anodic and
cathodic electrode reactions under charge transfer control is given by the
Butler-Volmer equation:
−
−
−
= η
β
η
β
RT
nF
RT
nF
i
i o
]
1
[
exp
exp
where …
7
where …
R = gas constant;
T = absolute temperature;
n = no. charges transferred (≡ valency);
F = Faraday (96,500 coul/mol);
β = “symmetry coefficient” ( 0.5);
io = exchange current density (a constant for the system).
The first term in { } in B-V describes the forward (metal dissolution, anodic) reaction;
the second term in { } describes the backward (metal deposition, cathodic) reaction.
153. A plot of the B-V equation for the metal
dissolution/deposition reaction gives the polarization curve:
8
If the symmetry coefficient β = 0.5, the curve is symmetrical
about (i = 0, Ee) and the B-V equation has a sinh form.
154. At very high over potential or high-field approximation:
At very high η, the reaction is essentially in one direction i.e., one of the terms in the B-V-
E is negligible and can be dropped. Thus, for metal dissolution (anodic reaction):
= a
o
a
RT
nF
i
i η
β
exp
i
approximately 0.12 volt
9
o
a
a
i
i
b log
=
η
nF
RT
ba
β
303
.
2
=
an inverse relationship between ηa and io
The Tafel coefficient for metal dissolution (anodic)
155. For metal deposition (cathodic reaction):
an inverse relationship between ηc and io
10
The Tafel coefficient for metal deposition (cathodic)
nF
RT
bc
)
1
(
303
.
2
β
−
−
=
156. At very low over potential or low-field approximation:
In the narrow region of small over potentials, the relation becomes linear
11
η
RT
nF
i
i o
=
157. If a reaction has a large exchange current, io, the curve is shallow and a
large current is obtained for a small over potential.
12
i.e., the reaction is not easily polarized (approaching non-polarizable)
158. If a reaction has a small exchange current, io, the curve is steep and a large
over potential is needed for a small current .
13
i.e., the reaction is readily polarized
163. A polarization diagrams or mixed potential diagrams are also
called Evans diagrams after the name of their originator, U.
R. Evans, corrosion scientist University of Cambridge in
England.
Characteristic features:
Evans diagrams is a graphical representations of a corroding metals, as a function of the
potential and current (usually current density) in a corroding environment.
2
The potential (in volts) is shown on the vertical axis (y-axis, ordinate) and the current
density (in ampere per unit area) on the horizontal axis (x-axis, abscissa).
It combine thermodynamic factors (E values) with kinetics factors (i values). So it is
useful for predicting the corrosion rates (corrosion kinetics ).
The exchange current densities have been included in the polarization diagram by Stern,
and such diagrams are called Stern diagrams. Evans diagrams do not include io.
164. It was postulated by Wagner and Traud in 1938. It has two basic assumptions:
Electrochemical reactions are composed of two or more partial anodic and cathodic
reactions.
There cannot be any accumulation of charges (law of conservations of charges).
Any electrochemical reaction can be algebraically divided into separate
oxidation and reduction reactions with no net accumulation of electrical
3
oxidation and reduction reactions with no net accumulation of electrical
charge.
In the absence of an externally applied potential, the oxidation of the
metal and the reduction of some species in solution occur simultaneously at
the metal/electrolyte interface at equal rate. Under these circumstances the
net measurable current is zero and the corroding metal is charge neutral, with
no net accumulation of charge. i.e., anodic current = cathodic current
165. Identify the possible anodic and cathodic reaction of the system.
Get the Eo and i values of all the reactions.
Decide on activation or concentrations polarization controlled.
Get activation over potential for each process that is potentially involved.
Get any additional information that could be affected by ηc.
Advantages:
4
Predicting the corrosion rates.
It combines thermodynamic factors (E values) with kinetics factors (i values)
Predicting environmental changes that will affect corrosion rates.
167. In this system, the oxidation reaction may be the dissolution of metal
and the reduction reaction may be symbolized as
Two possibilities:
6
In an aerated neutral or basic aqueous
solution, the reduction reaction could be
in a de-aerated acid, the reduction
reaction could be
Two possibilities:
168. The anodic and cathodic half reaction has its own electrode potential and exchange current
density .
The anodic electrode polarizes (shifts in potentials, ηa = Ecorr - (Ee)a) in cathodic directions
(η0) and cathodic electrode polarizes (ηc = Ecorr - (Ee)c) in anodic directions (η0) to an
intermediate value (between the two half-cell potentials). ΔE = (Ee)c - (Ee)a
7
intermediate value (between the two half-cell potentials). ΔE = (Ee)c - (Ee)a
This polarized potential is a mixture of the two half – cell potentials, so it is called as
MIXED POTENTIAL .The intersection of the two curves along the potential is called
corrosion potential (Ecorr) and the current density is called corrosion current (icorr).
At this point, rates of anodic and cathodic reactions are equal but in opposite directions.
The corrosion potential (Ecorr) is also called open circuit potential (EOCP). It corresponds to
potential with out any applied current.
171. Comparing the corrosion rates for zinc and iron (when present separately)
in dilute HCl solutions, zinc dissolution is expected to be higher than that of
iron from a thermodynamic view point (E°Zn
2+
/Zn = -0.76 V, E°Fe
2+
/Fe = —0.44 V).
Surprisingly, iron corrodes faster than zinc, when immersed in similar
concentrations of hydrochloric acid because of the lower io value of hydrogen
reduction on zinc (10-7A cm-2) compared to on iron (10-6A cm-2). So iron
corrodes faster than zinc as shown in Figure (next slide).
10
176. The rate of hydrogen evolution is decreased on zinc and increased on platinum because the
io on pt (10-3 A cm-2) is higher than Zn (10-10 A cm-2).
The rate of oxidation of zinc is increased significantly on coupling and zinc dissolves
vigorously. Nothing happens to platinum because Pt has more + ive reduction potential than Zn.
Platinum is an excellent catalyst for reduction of hydrogen and zinc is a poor catalyst.
If the area ratio of Pt to Zn (larger cathode in contact with smaller anode), is
increased, the effect of coupling with Pt on the corrosion rate of zinc will be magnified.
The corrosion rate of an active metal such as zinc or iron thus depends on
3
The corrosion rate of an active metal such as zinc or iron thus depends on
1. Cathodic metal they are coupled (i.e., exchange current density
for the reduction reaction is the controlling parameter)
For example, if Pb metal is coupled with Zn instead of
Pt in acid solution, the effect of coupling on anodic
oxidation of Zn would be negligible, since the io for
hydrogen reduction on Pb is very much lower than that
on Zn.
177. 2. pH and nature of the cathodic reactant
For example, if the Zn-Pt couple is exposed to
neutral pH solution where oxygen reduction is the
cathodic reaction (instead of H2 evolution), the
expected effect of noble metal (Pt) would be not
so significant since the io for oxygen reduction on
both the surfaces are nearly same.
4
179. According to the thermodynamic approach, the difference between the potential of zinc
(Eo = —0.76V) and gold (Eo = 1.50V) is higher than the difference between the potential of
the zinc (Eo = —0.76V) and platinum (Eo = +1.2V).
6
the zinc (Eo = —0.76V) and platinum (Eo = +1.2V).
So, the Zn-Au couple should corrode faster than Zn-Pt couple according to the
thermodynamic approach.
Surprisingly, Zn-Pt couple corrodes faster than Zn-Au couple, because the hydrogen
reduction rate is the highest on a platinum surface (io = 10-3A cm-2) compared to on gold (10-
6A cm-2). The reduction rate of hydrogen is very low on Zn surface (io = 10-10A cm-2).
180. Effect of change in cathodic surface area (anode-cathode area ratios):
7
181. The icorr of uncoupled zinc is lower than the icorr of zinc coupled either to Pt
(l cm2) or Pt (10 cm2).
The icorr of Zn coupled to Pt (10 cm2) is highest. The corrosion potential of coupled platinum
shifts to more noble values as the area of platinum surface is increased from l cm2 to 10 cm2.
The smaller the anode to cathode ratio as in the case of Zn coupled to Pt (10 cm2), the
larger is the magnitude of corrosion.
8
If the anodic surface areas are increased for a constant cathode area, anodic oxidation
rate can be lessened.
Avoid a small anode to cathode area ratio to minimize the risk of
serious galvanic corrosion.
183. 1. Effect of multiple reducible species (effect of oxidizer) on anodic corrosion:
Effect of Fe3+
2
184. The corrosion potential of active metal is shifted to a more noble direction.
The corrosion rate of the metal is increased.
The rate of hydrogen evolution is decreased.
Effect of added oxidant:
3
185. 2. Effect of oxygen (Effect of aeration and de-aeration):
The dissolved oxygen in aerated water is ~10 mg L-1 but in de-aerated water
is only ~0.01 mg L-1. So, cathodic reaction rate in aerated water is 103 higher
than de-aerated water.
4
186. Fe is oxidized to Fe2+.
Hydrogen is reduced.
Oxygen is reduced.
The rate of corrosion increases on aeration.
The rate of corrosion decreases on deaeration.
Effect of aeration :
5
For example, the Ni corrosion 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.
187. 1. Effect of velocity on limiting current density of O2 reduction:
6
Increased agitation decreases the ηc because the thickness of the
diffusion layer is decreases, and thereby rate of diffusion of ions
increases. So, corrosion current increases as the velocity increases.
188. Corrosion rate increases with solution velocity as long as the
cathodic process is under diffusion control and it becomes
independent of velocity at higher velocities when the cathodic
reaction is under activation control.
7
189. 1. The cathodic reaction with the higher oxidation potential is controlling the
reaction.
8
190. 2. The slowest reaction controls the rate of corrosion. Normally this is the
cathodic reaction.
• A small changes
in kinetics of
cathode have a
large effect on
corrosion rate.
• A small changes
in kinetics of
anode have small
9
anode have small
effect on
corrosion
192. Corrosion involves dissolution of metal, as a result of which the
metallic part loses its mass (or weight) and also it generates corrosion
current. Corrosion rate measurements are therefore based on either weight
loss or electrochemical measurements.
The simplest way of measuring the corrosion rate of a metal is to expose
(1) Weight loss measurements
(2) Linear polarization method
(3) Tafel Extrapolation Method
2
The simplest way of measuring the corrosion rate of a metal is to expose
the sample to the test medium (e.g. sea water) and measure the loss of weight of
the material as a function of time. The weight loss can be converted into
corrosion rate (mpy) by using following expression.
DAT
W
534
=
(mpy)
year
per
n
penetratio
Mils
W = weight loss, mg
D = density of specimen, g/cm3
A = area of specimen, (in.2)
T = exposure time, hr
1 Inch = 2.54 cm, 1 Feet = 30.48, 1 Feet = 12 Inch
193. Simple and inexpensive
Corrosion deposits can be observed and analyzed
Advantages
not applicable for localized corrosion like pitting
minimum exposure time should be 45 days
Disadvantages
3
A potentiostat is an electronic instrument that
controls the potential difference between a working
electrode (W) and a reference electrode (ref.).
It measures the current flow between the working
(W) and counter electrodes (C.E).
A galvanostat is an electronic instrument
that controls the current between the
working (W) and counter electrodes (C.E).
It measures the potential difference
between a working electrode (W) and a
reference electrode (ref.).
Potentiostat
Galvanostat
194. A potentiostat requires an electrochemical cell with three electrodes for
measuring corrosion potential and corrosion current.
4
195. (1)Three electrode setup:
Potentiostat
I V S
Vv
CA
Vi
Rm
C
C
CE
WE
R
E
I/E Converter
Control Amp
Electrometer
Cell Switch
5
The potential difference is controlled between the WE and the CE and measured between
the RE (kept at close proximity of the WE) and S, because the WE is connected with S.
This configuration allows the potential across the electrochemical interface at the WE to
be controlled with respect to the RE.
Normally, the potential between the WE and CE usually is not measured, because potential
is applied by the control amplifier (by the user)
This setup is typically used for corrosion rate measurement , electroanalytical experiments,
etc…
It is also used for the characterization of half cells of energy storage or conversion
devices like batteries, supercapacitor, fuel cells, photovoltaic panels etc…
196. (2)Two electrode setup:
In a two-electrode cell setup, CE and RE are shorted on one of the electrodes
while the WE and S are shorted on the opposite electrode.
The potential/current across the complete cell is measured.
This setup is typically used with energy storage or conversion devices like
batteries, supercapacitor, fuel cells, photovoltaic panels etc…
It is also used in measurements of ultrafast dynamics of electrode processes
or electrochemical impedance measurements at high frequencies ( 100 kHz).
6
(3) Four electrode setup:
This setup is used to measure the potential difference between RE
and S, due to the passage of a current across WE and CE)
It is used to calculate the resistance of the interface or the
membrane conductivity.
197. The corrosion rate can be determined from the polarization resistance
(Rp) using the Stearn-Geary equation, if the Tafel slopes are known.
In potentiostatic polarization method, you record a current versus voltage curve as the cell
voltage is swept over a small range of potential that is very near to Eoc (generally +10 mV for anodic
and -10 mV for cathodic ). Values of +5 mV or -5 mV and +20 mV or -20 mV are also commonly used.
7
Potentiodynamic measurements yield curves of log i vs. E and the reciprocal of the slope
of the curve (∆η/∆i) at the corrosion potential (Ecorr) is measured as shown below.
(Ecorr) = Eoc
198. Log i
8
As the potential is raised (i.e., anodically polarized), the current flow will increase and
behave like anodic polarization curve. Alternatively, if the potential were lowered (i.e.,
cathodically polarized) the current flow will decrease and behave like cathodic
polarization curve.
199. Here, the polarization resistance, Rp = (∆η/∆i) at ∆E = 0. By measuring this slope, the
rate of corrosion can be measured using Stearn-Geary equation
( )
β
β
β
β
η
c
a
corr
c
a
i
i +
=
∆
∆
303
.
2
=
)
(
Slope Rp
( )
β
β
β
β
c
a
p
c
a
corr
R
i +
=
303
.
2 ( )
β
β
β
β
c
a
c
a
B
+
=
303
.
2
9
β
β c
a
p
R
More accurate, instantaneous and fast
two-point measurement at potentials above
and below the OCP
Advantages
The extent of linearity of the potential –
current plot depends on βa and βc values
Disadvantages
( )
β
β c
a
R
i
p
corr
B
= Where B is Stearn-Geary constant
and β is Tafel constant
Tafel slopes are necessary to calculate icorr
Valid for corrosion under activation control
200. Usually, calculated corrosion rates are not wrong by more than a factor of
2–3 if the Tafel slopes are both assumed to be 100 mV/dec
10
201. This technique uses data obtained from cathodic and anodic polarization
measurements. Cathodic data are preferred, since these are easier to
measure experimentally.
In this technique, the polarization curves for the anodic and cathodic reactions are
obtained by applying potentials about 200 mV well away from the corrosion potential and
recording the current.
Plotting the logarithms of current (log I) vs potential and extrapolating the currents in the
11
Plotting the logarithms of current (log I) vs potential and extrapolating the currents in the
two Tafel regions gives the corrosion potential and the corrosion current icorr as shown in
next slide.
Knowing icorr the rate of corrosion can be calculated in desired units by using Faraday's law
as discussed in pdf 8.
Tafel constants must be calculated from both the anodic and cathodic portions of the
Tafel Plot. The units of the Tafel constants is V/decade. A decade of current is one order of
magnitude.
204. Advantage:
greater accuracy than conventional weight loss methods
it is possible to measure extremely low corrosion rates
rapid determination of corrosion rates
instantaneous corrosion rate determination
it provide a direct measure of the corrosion current
Disadvantage:
14
Disadvantage:
The specimen geometry requires a strict control to obtain a uniform current.
The Tafel region is often obscured by concentration polarization and by the existence
of more than one activation polarization process
206. Schematically show on overpotential (η) vs. current density (i) plot (η vs i) that if
exchange current density increases only for cathodic reaction, both the corrosion rate and
corrosion potential increase.
16
Show that current density is equivalent to corrosion rate.
Show the origin of exchange current density. What is its significance in corrosion rate?
Why is surface condition importance in deciding exchange current density?
207. Problem 1:
what is the corrosion rate for a steel coupon, 2 cm2 in
area, which has lost 0.03 g in 20 hrs.
Problem 2:
Problem 3:
17
Problem 3:
210. Problem 9:
Find out multiplication factor for the conversion of
corrosion rate from (a) mdd to mpy, (b) mpy to mmy-1
Problem 10:
Find out expression for corrosion rate from Faraday’s laws of
electrochemistry for (a) Uniform corrosion and (b) pitting
corrosion.
20
corrosion.
Problem 11:
Find out the corrosion rate of Fe in sea water if the current
density is 6.0 x10-6A/cm2 in (a) mdd, (b) mmy-1 and (c) mpy.
212. Iron will corrode in dilute nitric acid, but at higher
concentrations the corrosion rate of iron is very little or negligible.
2
Iron is resistant to corrosion in nitric acid at concentrations around 70%. Faraday
also conformed this by cell made up of passive iron coupled to platinum in concentrated
nitric acid produced little or no current.
Once passivated under these conditions, it can also exhibit low rates of corrosion as
the nitric acid is diluted.
However, if this passive film is disturbed, rapid corrosion will begin and re-passivation
will not be possible until the nitric acid concentration is raised to a sufficient level.
213. Definition 1. A metal or alloy become passive on increasing the electrode potential
towards more noble values (anodic polarization), at which the rate of
In the Eh – pH diagrams, resistance to metallic corrosion is indicated at stability regions
where either the metal remains thermodynamically stable (immunity) or the metal surface is
covered with an oxide/hydroxide layer (passivity).
Passivity is due to the formation of thin, impermeable and adherent surface films under
oxidizing conditions (e.g., iron in chromate or nitrite solutions) often associated with anodic
polarization (e.g., iron in H2SO4).
3
towards more noble values (anodic polarization), at which the rate of
anodic dissolution is less than the less noble potential in given
environment.
i.e., noble potential, low corrosion rate
Definition 2. A metal or alloy become passive on increasing the concentration of an
oxidizing agent in an adjacent solution or gas phase in absence of
external current, at which the rate of anodic dissolution is less than
the lower concentration of the oxidizing agent.
active potential, low corrosion rate
214. Examples of metals or alloys (active-passive ) that are passive under Definition 1 are Cr, Ni,
Mo, Ti, Zr, the stainless steel, 70%Ni – 30%Cu alloys (Monel), iron in dissolved chromates
(passive in passivator solutions) and several other metals and alloys.
Metals and alloys in this category show a marked tendency to polarize anodically and
corrosion potentials of this category approach the OCP of oxygen electrode (exhibit
potentials near those of the noble metals.)
Definition 1 usually conform as well as to Definition 2 based on low corrosion rates.
Definition 1
4
Definition 1 usually conform as well as to Definition 2 based on low corrosion rates.
Examples of metals that are passive under Definition 2 (passive metals) are Pb immersed in
sulfuric acid, or Mg in water, or iron in inhibited pickling acid or zinc based on lowcorrosion
rates, despite pronounced corrosion tendencies according to thermodynamic data (e.m.f. series).
Their corrosion potentials are relatively active, and polarization is not pronounced when they
are made the anode of a cell.
Definition 2
215. Electrochemical basis of active-passive behavior (anodic dissolution
behavior ) of a metal is illustrated in following figure.
5
216. As the potential increases towards more noble direction (anodic polarization) than EM/M
+, the
rate of dissolution of the metal also increases. At this point partially insulating films (probably
porous metal sulphate, nitrate or chromate) on metal surface is formed .
At this point the rate corrosion is maximum and maximum current density is called critical
current density (icritical). The potential corresponding to icritical is called the primary passive
potential (Epp) as it represents the transition of a metal from an active state to a passive state.
The potential at which the current becomes virtually independent of potential and remains
virtually stationary is called the flade potential (EF). At this point much thinner films (probably
MxOy or Fe (OH)2 or FeO) on metal surface is formed and metal becomes passive.
6
The minimum current density required to maintain the metal in a passive state is called
passive current density (ip). At ip, the metal dissolution occurs at a constant rate and the oxide
film begins to thicken. The dissolution rate in the passive region, therefore, remains constant.
On further increase in potential leads to an accelerated rate of corrosion due the breakdown
of passive films and is called transpassive potential (Etranspassive). Now the corrosion product is
Fe3+ and O2 evolution, which causes a sharp increase in the current.
The transpassive potential corresponding to the end of passive region, which corresponds to
the initial point of anodic evolution of oxygen . This may correspond either to the breakdown
(electrolysis) voltage of water, or, to the pitting potential.
217. Stability of passivity is related to EF. The lower the Eo
F, the easier it
becomes for passivation and higher the film stability. For Cr – Fe alloys, the
value ranges from 0.63 V to -0.10V with 25% chromium addition.
The passive films maybe as thin as 2-10 nm, and they offer a limited
electronic conductivity, and behave like semi-conductors with metallic
properties rather than the properties shown by bulk oxides.
7
218. In a variety of Fe-Cr alloys, Epp and icritical substantially reduced than iron
because of the formation of uniform protective films.
8
219. The transpassive region increase with increasing chromium content. As the
film dissolves, cation vacancies are created in the oxide surface and the
conductivity of the film is increased.
Metals, like Fe, Cr, Ni and Ti, show a strong active-passive behavior. The, cathodic reaction is
a deciding factor in the establishment of passivity. The rate of corrosion depends upon the
degree of polarization of the anode.
A metal not showing any passivity will exhibit a linear E vs log i relationship. On the other
hand, a metal exhibiting passivity would exhibit a non-linear anodic polarization.
9
hand, a metal exhibiting passivity would exhibit a non-linear anodic polarization.
Metals, like zinc, magnesium and aluminum, show a passive behavior in atmospheric corrosion.
The rate of corrosion depends on the degree of polarization of both the anode and the cathode.
220. Parameters Definitions of Parameters
Equilibrium potential
(Eeq or Eo or EM/Mz+)
The potential of an electrode in an electrolyte at which rate of
forward reaction is balanced by the rate of reverse reaction. At
equilibrium potential, the rates of the anodic (oxidation) and
cathodic (reduction) processes are equal, and there is no net
charge transfer.
10
charge transfer.
Passive potential
(Epassive)
The potential at which a metal surface changes from an active to a
passive state. Passive potential is defined as the potential below
which the metal surface remains active and above which the metal
surface remains passive.
Flade potential (EF). The potential at which a metal changes from a passive state to an
active state. Flade potential is defined as the potential below
which the metal surface remains passive and above which the
metal surface remains active.
221. Transpassive potential
(Etranspassive)
The potential corresponding to the end of passive region
which corresponds to the initial point of anodic evolution
of oxygen. This may correspond either to the breakdown
(electrolysis) voltage of water, or, to the pitting
potential.
Pitting potential (Ep). It is the potential at which there is a sudden increase in
the current density due to breakdown of passive film on
the metal surface in the anodic region.
11
Critical current density
(icritical)
The maximum current density observed in the active
region for a metal or alloy that exhibits an active-passive
behavior.
Passive current density
(ip)
The minimum current density required to maintain the
thickness of the film in the passive range.
223. 13
Show that when applied current density approaches the limiting current
density the over potential at the cathode increases rapidly where as over
potentialat the anode remains very small.
