1. Corrosion and Degradation of Materials
Chapter 17 Part 2
Corrosion and Degradation of Materials

2. Electrochemical Cells
Electrochemical cells can be set up under different situations and can lead to the corrosion of a metal
Grain/Grain Boundary
Generally the grain boundary is a region of disarray, and has higher energy, and atoms can be “pulled out” of the solid more easily
This phenomenon is used in metallography, when polished metal surfaces are etched with acids to reveal grain boundaries
Sometimes, segregation of solutes to grain boundaries may make them more “noble” than the surrounding grain, resulting in corrosion in regions adjacent to grain boundaries
When certain stainless steels are cooled slowly, chromium precipitates as chromium carbide along grain boundaries, robbing the surrounding grain of the protective chromium oxide. The grain boundary is cathodic compared to surrounding grain

3. Electrochemical Cells
In multiphase materials, one phase may be anodic with respect to another
Corrosion rates are higher in multiphase materials.
For example
In pearlitic gray cast iron, graphite is more noble than pearlite, leading to corrosion of pearlitic regions
Martensite (single phase) is more corrosion resistant than austenite that has been slow cooled to form pearlite (two phase)
Tempered martensite corrodes more easily than martensite
Low temperature tempering results in finer Fe3C particles and more corrosion sites
Higher temperature tempering results in coarser Fe3C particles and fewer corrosion sites

4. Corrosion Kinetics
Up to this point, we have dealt with the thermodynamics of corrosion, i.e. which combinations of conditions results in anodic and cathodic regions under equilibrium
Corrosion does not occur under equilibrium conditions
Of interest is the corrosion kinetics, i.e. the rate at which a metal corrodes
For each atom of a metal that participates in the oxidation reaction, n electrons need to get transported away
The weight of a metal that is lost due to corrosion is given by Faraday’s law
M  Mn+ + ne-
w = weight loss during corrosion (or weight gain during electroplating)
I = current in amps = iA
i = current density
A = area of corroding surface
t = time in seconds
M = atomic mass g/mole
n = number of electrons involved in the corrosion reaction
F = Faraday’s constant
= 96,500 C/mol

5. Corrosion Penetration Rate (CPR)
If a metal exposed to a corrosive environment, dissolves uniformly, the thickness removed can be calculated
If a sample with a surface area of A (in2) is exposed for t (hours) and a weight loss of w (mg) is measured then
CPR is the rate of thickness loss
CPR<20 mil/yr or about 0.5 mm/yr is acceptable

6. Polarization
In an electrochemical half cell the metal atoms are in a state of equilibrium with its ions in solution
There is an equilibrium exchange current density i0 associated with the transfer of electrons at the standard emf potential E0 (or V0) of the half cell
There is an i0 and E0 associated with the anodic and cathodic reactions
However, potential differences cannot be maintained in a conductive metal, such as Zn
There is a displacement of the electrode potentials and currents from points A and B to C
This displacement of electrode potentials is called polarization
Point A: V, 10-7A/cm2
Point B: 0V, 10-10A/cm2
Point C: ~-0.5V, 10-4A/cm2
icorr = 10-4A/cm2 is used in
CPR calculations

7. Activation and Concentration Polarization
Corrosion reaction may result in a build up or depletion of the ions or atoms that are required for a corrosion reaction.
Activation polarization: In a multistep electrochemical reaction the rate is controlled by the slowest step.

8. Passivation
Passivation is the loss of chemical reactivity in presence of a environmental condition.
The formation of surface layer of reaction products that inhibit further reaction
Oxide film theory: A passive film of reaction products acts as a diffusion barrier.
Adsorption theory: Passive metals are covered by chemisorbed films of oxygen.
Examples: Stainless steel, nickel alloys, titanium and aluminum alloys passivate in certain environments

9. Polarization Curve for Passivation
Initially, the potential of the metal increases with current density, i.e. the metal undergoes active corrosion
When potential reaches Epp the primary passive potential, current density decreases, i.e., the corrosion rate also decreases
In order to make the metal active again, there may need to be an externally applied potential

10. Galvanic Series
The Standard emf series gives the relative oxidation or reduction behavior under standard conditions.
The Galvanic series ranks materials on the basis of corrosion behavior in sea water

11. Types of Corrosion
Though the basic corrosion mechanism is the same, i.e. the oxidation of a metal due to transfer of electrons, conditions under which the process occurs can be vastly different, and has lead to the classification of corrosion into the following types:
Uniform
Galvanic
Crevice
Pitting
Inter-granular
Selective leaching
Erosion/cavitation
Stress corrosion
Fretting

12. Types of Corrosion Uniform or general corrosion
Most common in terms of weight loss and destruction, but also the easiest to control
Galvanic or two metal corrosion
When dissimilar metals are in contact
Very sensitive to the relative areas of the anodic and cathodic regions
If the anodic region is small, then corrosion can be rapid and deep
Tin coating of steel in “tin” cans
If the coating is scratched, the underlying steel rusts quickly
Zinc coating of steel by galvanizing
If the coating is scratched, the galvanic cell results in corrosion of the coating, but due to large surface area of Zn, the CPR of the Zn coating is relatively low

13. Types of Corrosion Crevice Corrosion
Localized electrochemical corrosion in crevices and under shielded surfaces where stagnant solutions can exist.
Occurs under valve gaskets, rivets and bolts in alloy systems like steel, titanium and copper alloys.
Anode: M  M+ + e-
Cathode:O2 + 2H2O + 4e-  4OH-
As the solution is stagnant, oxygen is used up and not replaced.
Chloride ions migrate to the crevice to balance positive charge and form metal hydroxide and free acid that causes corrosion

14. Types of Corrosion Pitting
Initiates at non-uniformities in composition
Growth of pit involves dissolution of metal in pit maintaining high acidity at the bottom.
Anodic reaction at the bottom and cathodic reaction at the metal surface.
At the bottom of the pit
MCl + H2O  M(OH) + H+ + Cl-
The presence of H+ at the bottom of the pit pulls more M into solution
Some metals (stainless steel) have better resistance than others (titanium).

15. Types of Corrosion Intergranular corrosion
Localized corrosion at and/or adjacent to highly reactive grain boundaries resulting in disintegration.
When stainless steels are heated to or cooled through sensitizing temperature range ( C) chromium carbide precipitate along grain boundaries.
When exposed to corrosive environment, the region next to grain boundaries become anodic and corrode.

16. Types of Corrosion Stress Corrosion Cracking (SCC)
Cracking caused by combined effect of tensile stress and corrosive environment
Stress might be residual and applied
Only certain combination of alloy and environment causes SCC
Crack initiates at pit or other discontinuity
Crack propagates perpendicular to stress
Crack growth stops if either stress or corrosive environment is removed.

17. Types of Corrosion Erosion corrosion
Acceleration in rate of corrosion due to relative motion between corrosive fluid and surface.
Pits, grooves, valleys appear on surface in direction of flow.
Corrosion is due to abrasive action and removal of protective film.
Cavitation damage
Caused by collapse of air bubbles or vapor filled cavities in a liquid near metal surface.
Rapidly collapsing air bubbles produce very high pressure (60,000 PSI) and damage the surface.
Occurs at metal surface when high velocity flow and pressure are present.

18. Types of Corrosion Selective leaching:
Selective removal of one element of alloy by corrosion
Example: Dezincification or selective removal of zinc from copper and brasses
Weakens the alloy as single metal might not have same strength as the alloy
Fretting corrosion
Occurs at interface between materials under load subjected to vibration and sliding
Metal fragments get oxidized and act as abrasives between the surfaces

19. Oxidation of Metals Metals oxidize when exposed to air
Oxidation partial reaction: M  M2+ + 2e-
Reduction partial reaction: ½ O2 + 2e-  O2-
Oxidation starts by lateral expansion of discrete oxide nuclei.
The oxide layer is the medium through which:
Metal ion diffuse to the surface.
O2- ions diffuse to oxide-metal interface
Electrons diffuse to oxide gas interface
The oxide therefore serves as both the electrolyte for the electrochemical process as well as the electrically conductive medium for the flow of electrons
If the oxide is electrically non-conductive or does not allow the diffusion of ions, then a protective oxide layer can form

20. Oxidation of Metals
Whether an oxide films form depends on following factors:
Pilling-Bedworth Ratio
PB ratio should be close to 1
Good adherence
High melting point of the film
Low oxide partial pressure
Should not evaporate
Coefficient of expansion equal to that of metal
High temperature plasticity
Low conductivity and diffusion coefficients of metal ions and oxygen.
Where
AO = molecular wt. of oxide
rO = density of oxide
AM = molecular wt. of metal
rM = density of metal

21. Oxidation of Metals

22. Oxidation of Metals Linear: Catastrophic oxidation
Oxidation rate is expressed as weight gained per unit area.
Parabolic oxidation behavior
W2 = K1t+K2
K1 = Parabolic rate constant
K2 = constant
Ion diffusion is controlling the step (Eg – Fe, Cu, Co)
Linear oxidation behavior
W = K3t
W=weight gained
per unit area
K3 = linear rate constant.
t = time
Porous or flakey oxide layer allows continuous supply of oxygen to the metal surface. Ex. Na, K, Ta
Logarithmic oxidation behavior
W = K4 Log(K5t + K6)
K4, K5, K6 = constants
Ex. Al, Cu, Fe (at slightly elevated temperature)
Linear: Catastrophic oxidation
Parabolic or logarithmic: a protective oxide may form

23. Corrosion Control

24. Corrosion Control – Material Selection
Metallic Metals:
Use proper metal for particular environment
For reducing conditions, use nickel and copper alloys
For oxidizing conditions, use chromium based alloys
Nonmetallic Metals:
Limit use of polymers in presence of strong inorganic acids
Ceramics have better corrosion resistance but are brittle

25. Coatings
Metallic Coatings: Used to protect metal by separating from corrosive environment and serving as anode.
Coating applied through electroplating, roll bonding or hot dipping
Might have several layers
Inorganic coatings: Coating with steel and glass.
Steel is coated with porcelain and lined with glass.
Organic coatings: Organic polymers (paints and varnishes) are used for coatings.
Serve as barrier but should be applied carefully.

26. Design General design rules:
Provide allowance for corrosion in thickness.
Weld rather than rivet to avoid crevice corrosion.
Avoid dissimilar metals that can cause galvanic corrosion.
Avoid excessive stress and stress concentration.
Avoid sharp bends in pipes to prevent erosion corrosion.
Design tanks and containers for easy draining.
Design so that parts can be easily replaced.
Design heating systems so that hot spots do not occur.

27. Alteration Environment
Lower the temperature reduces reaction rate.
Decrease velocity of fluids reduces erosion corrosion.
Removing oxygen from liquids reduces corrosion.
Reducing ion concentration decreases corrosion rate.
Adding inhibitors which are retarding catalysts reduce corrosion

28. Cathodic Protection
Electrons are supplied to the metal structure to be protected.
Example: Fe in acid
Fe  Fe2+ + 2e-
2H+ + 2e-  H2
Corrosion of Fe will be prevented if electrons are supplied to steel structure.
Electrons can be supplied by:
External DC supply
Externally impressed anodic currents form protective passive films on metal and alloy surfaces.
Galvanic coupling with more anodic metal
A more easily corroded material is used as a sacrificial anode

29. Degradation of Polymers
Polymers degrade by a physico-chemical process rather than an electrochemical process
Different forms of degradation include
Swelling and Dissolution:
A diffusion of liquid or solute into the polymer
Small solvent molecules enter into the space between the macromolecules of the polymer causing chains to be forced apart.
Reduced intermolecular secondary bonding
Increased ductility and lower strength
Polymers can dissolve in solvents, especially when the solvent has a similar chemical structure
Example: Hydrocarbon polymers in gasoline

30. Degradation of Polymers

31. Degradation of Polymers
Bond Rupture or Scission by exposure to Radiation
a, b, g, electron beam, X-rays, UV can all cause bonds to break, i.e. an electron that makes a covalent bond may be removed
Subsequently, the bond may result in the breaking of the carbon chain (scission) or result in cross-linking, changing the structure of the polymer
Exposure to radiation can sometimes be used in constructive processes, such as Rapid Prototyping, in which cross-linking is enhanced by exposure to UV radiation
Bond Rupture by exposure to chemicals
Oxygen O2, Ozone O3,and other substances can cause chain scission, especially if there are double bonds, such as those in un-vulcanized or partially vulcanized rubber

32. Degradation of Polymers

33. Degradation of Polymers
Thermal Effects
High temperature breaks covalent bonds
C-F bond in PTFE (Teflon) is stronger than the C-H bond in most polymers and the C-Cl bond in polymers like PVC
Teflon use temperature is higher than other thermoplastics
Weathering
A general term referring to the degradation of polymers due to exposure to outdoor conditions that may include a combination of several of the above effects