metb113 emat 13 corrosion

8/11/2019 METB113 EMat 13 Corrosion

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METB113/1Corrosion and Protection Methods

METB113ENGINEERING MATERIALS

CORROSION AND PROTECTION METHODS
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Types and Its Mechanisms

Protection & Control MethodsStructure Analysis


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Corrosion: -- the destructive and unintentional electrochemical

attack of a materials.Rate of corrosiondepends upon temperature and concentration of

reactants and products.

Metals have free electrons that setup electrochemical cellswithin

their structure.

Metals have tendency to go back to low energystate by corroding.

Ceramics and polymers suffer corrosion by direct chemical attack.

-- Ex: Al Capone's

ship, Sapona,

off the coast

of Bimini.

Cost:-- 4 to 5% of the Gross National Product (GNP)*-- in the U.S. this amounts to just over $400 billion/yr**

Photos courtesy L.M. Maestas, SandiaNational Labs. Used with permission.

Introduction


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OxidationReduction Reactions

A metal (EgZn) placed in HCl undergoes corrosion.

Zn + 2HCl ZnCl2+ H2or

Zn + 2H+ Zn2+ + H2also

Zn Zn 2+ + 2e- (oxidation half cell reaction)

2H+ + 2e- H2 (Reduction half cell reaction)

Oxidation reaction: Metals form ions at local anode. Reduction reaction: Metal is reduced in local charge at

local cathode.

Oxidation and reduction takes place at same rate.

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5

Two reactions are necessary:-- oxidation half-cell reaction:-- reductionhalf-cell reaction:

Zn Zn2 2e

2H

2e H2(gas)

Other reductionreactions in solutions with dissolved oxygen:

-- acidic solution -- neutral or basic solution

O2 4H 4e 2H2O O2 2H2O 4e 4(OH)

ELECTROCHEMICAL CORROSION

Zinc

Oxidation reactionZn Zn2+

2e-Acidsolution

reduction reaction

H+H+

H2(gas)

H+

H+

H+

H+

H+

flow of e-in the metal

Ex: consider the corrosion of zinc in an acid solution


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Standard electrode Half-Cell Potential of Metals

Oxidation/Reduction half cell potentials are compared withstandard hydrogen electrode

half cell potential.

Voltage of metal (Ex: Zn) is

directly measured against

hydrogen half cell electrode.

Anodic to hydrogen More tendency to corrode

Examples:-Fe (-0.44), Na (-2.74)

Cathodic to hydrogen Less tendency to corrode

Examples:-Au (1.498), Cu (0.33)

Figure 12.

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7

Standard Hydrogen Electrode

Oxidation/Reduction half cell potentials are compared withstandard hydrogen electrodehalf cell potential.

Two outcome:

0o

metalV (relative to Pt)

Standard Electrode Potential

Adapted from Fig. 16.2,

Callister & Rethwisch 3e.

-- Electrodeposition

-- Metal is the cathode (+)

Mn+

ions

ne-

e- e-

25

C

1M Mn+soln1M H+ soln

Platinum

metal,M H

+

H+

2e-

0o

metalV (relative to Pt)

-- Corrosion

-- Metal is the anode (-)

Platinum

metal,M

Mn+

ions

ne- H2(gas)

25

C

1M Mn+soln1M H+ soln

2e

-

e-e-

H+

H+


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8

Standard EMF Series

metalo

Metal with smaller

V corrodes.

EMFseries

AuCuPbSn

NiCoCdFeCr

ZnAlMgNaK

+1.420 V+0.340- 0.126- 0.136

- 0.250- 0.277- 0.403- 0.440- 0.744

- 0.763- 1.662- 2.363- 2.714- 2.924

metalVmetal

o

Data based on Table 17.1,Callister 7e.

moreanodic

more

cathodic

V =0.153V

o

Adapted from Fig. 16.2,Callister & Rethwisch 3e.

-

1.0 M

Ni2+ solution

1.0 M

Cd2 + solution

+

25

C NiCd

Ex: Cd-Ni cell

V < V Cd corrodesCdo

Nio


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9

Effect of Solution Concentration and Temperature

Ex: Cd-Ni cell withstandard 1 M solutions

VNio

VCdo 0.153 V

-

Ni

1.0 M

Ni2+ solution

1.0M

Cd2 + solution

+

Cd 25

C

Ex: Cd-Ni cell with

non-standard solutions

Y

Xln

nF

RTVVVV

o

Cd

o

NiCdNi

n= #e-per unitoxid/redreaction(= 2 here)F=Faraday'sconstant= 96,500C/mol.

Reduce VNi- VCdby-- increasingX-- decreasing Y-- increasing T

- +

Ni

YM

Ni2+ solution

XM

Cd2 + solution

Cd T


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Microscopic galvanic cell can exist in metals or alloys due to

difference in structure, composition and stress concentration.

When single electrodeis immersed in an air-free electrolyte,

microscopiccathodes and anodes are formed on the surface due

to difference in structure and composition.

Oxidation reaction occurs at local anodeand reduction reaction

at local cathode.

If iron is immersed in

oxygenated water,

2Fe + 2H2O + O2 2Fe2++ 4OH- 2Fe(OH)2

Microscopic Galvanic Cell Corrosion of Single Electrode

Fe Fe2++ 2e-

O2+ 2H2O + 4e- 4OH-

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Grain-grain-boundary Electrochemicalcells

Grain boundaries are more anodicand hence get corroded

by electrochemical attack.

Grain boundaries are at higher energy.

Impurities migrateto grain boundaries.

Solute segregation might cause grain boundaries to

become more cathodic.

Cartridge Brass

Grain

BoundaryGrain boundary

(anode)

Grain boundary

(cathode)

anode

Figure 12.913-11
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Multiple Phase Electrochemical Cells In multiple alloys, one phase is more anodicthan another.

Corrosion rates are higherin multiphase alloys.

Example:In pearlite gray cast iron, graphite flake is

cathodic than surrounding pearlite matrix.

Anodic pearlite corrodes

Steel, in martensiticcondition

(single phase) after quenching

from austenitic condition, has

better corrosion resistance. Impurities in metals leads toprecipitationof intermetallic

phases and hence forms anodic and cathodic regions

leading to corrosion.Figure 12.10

After Metals handbook, vol. 7, 8thed., American Society for Metals, 1972, p.83.13-12


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Rate of Uniform Corrosion Faradays equation:

W = weight of metal (g), corroded or electroplated in an

aqueous solution in time t, seconds.

I = Current flow A, i = current density A/cm2

M = atomic mass of metal g/mol

n = number of atoms/electron produced or consumedF = Faradays Constant, A = area, cm2

Corrosion rateis expressed as weight loss per unit areaper

unit time or loss in depth per unit time.

nF

iAtM

nF

ItMW

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Activation and Concentration Polarization

Activation polarization:Electrochemical reactions that are controlled by

the slowest stepin the reaction sequence.

There is a critical activation energyto surmount energy barrier associated

with slowest step.

Figure 12.15After M. G. Fontana and N. D. Greene, Corrosion Engineering, 2nded., McGraw-Hill, 1978, p.15.13-15

Polarization controls the corrosion rate.

1. Adsorption of H+from solutiononto zinc surface.

2. Electron transfer from zinc tohydrogen atom.

3. Reduction of hydrogen.4. Combining two atom hydrogen

together.

5. Coalescence of many hydrogenmolecules to form bubbles.

The slowest of these steps controlthe overall reaction.


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Concentration polarization:

exists when the reaction rate is limited by the diffusion of ions in the

solution. Example:Reduction rate of H+ions at surface is controlled by diffusion

of H+ions onto metal surface.

Activation and Concentration Polarization


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Passivation

Passivation is loss of chemical reactivityin

presence of an environmental condition.

Formation of surface layerof reaction products that

inhibit further reaction. Oxide film theory: Passive film is always a

diffusion barrier of reaction products.

Adsorption theory: Passive metals are covered

by chemisorbed films of oxygen.

Examples:-Stainless steel, nickel alloys,

titanium and aluminum alloys.

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METB113/19C


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Uniform and Galvanic Types of Corrosion

Uniform or general attack corrosion:Reaction

proceeds uniformly on the entire surface.

Controlled by protective coatings, inhibitors and

cathodic protection.

Galvanic or two metal corrosion:

Electrochemical reaction leads to corrosion of on

metal.

Zinc coatings on steel protects steel as zinc is anodictosteel and corrodes.

Large cathode area to small anode area should be

avoided.

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METB113/20C i d P t ti M th d


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Pitting Corrosion

Pitting:Localized corrosive attacks that

produces holes or pits in a metal.

Results in sudden unexpected failure as pits go

undetected(covered by corrosion products). Pitting requires an initiation

periodand grows in

direction of gravity. Pits initiate at structural

and compositional

heterogeneities. Pitting of stainless steel

Figure 12.20

Courtesy of LaQue Center for Corrosion Technology, Inc.13-20



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Growth of Pit

Growth of pit involves dissolution of metalin pitmaintaining high acidity at the bottom.

Anodic reactionat the

bottom and cathodic

reactionat the metal

surface.

At bottom, metal chloride + water Metal

hydroxide + free acid.

Some metals (stainless steel) have better resistance

than others (titanium).

Figure 12.21

After M. G. Fontana and N. D. Greene, Corrosion Engineering, 2nded., McGraw-Hill, 197813-21



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Figure 12.21

After M. G. Fontana and N. D. Greene, Corrosion Engineering, 2nded., McGraw-Hill, 197813-21

Crevice Corrosion

Localized electrochemical corrosion in crevicesand undershielded surfaces where stagnant solutions can exist.

Occurs under valve gaskets, rivets and bolts in alloy systemslike 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 tocrevice to balance positive charge and form metal hydroxideand free acidthat causes corrosion. Figure 12

nd


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Stress Corrosion Stress corrosion cracking (SCC):Cracking caused by

combined effect of tensile stressand 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 propagatesperpendicular

to stress Crack growth stops if either stress or corrosive environment

is removed.

Figure 12.27

After R. W. Staehle.13-24



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Erosion Corrosion and Cavitation Damage

Erosion corrosion: Acceleration in rate of corrosion due torelative motionbetween corrosive fluid and surface.

Pits, grooves, valleys appear on surface in directionof

flow.

Corrosion is due to abrasive actionand removal ofprotective film.

Cavitation damage:Caused by collapseof 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.

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Mechanisms of Oxidation

Oxidation partial reaction: M M 2++ 2e-

Reduction partial reaction: O2+ 2e- O2-

Oxidation starts by lateral expansionof discrete oxide nuclei.

Metal diffusesas electrons or cations across oxide films.

Sometimes O2-ions diffuse to oxide metal interface and

electrons diffuse to oxide gas interface.

Figure 12.30

After L.L. Shreir (ed.) Corrosion, vol.1, 2nded., Newnes-Butterwirth, 1976, p. 1:242.13-28



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Oxidation Rates Oxidation rate is expressed

as weight gainedper unit area. Linearoxidation behavior

W = KLt

If ion diffusionis controlling the step (EgFe, Cu)

W2 = Kpt+C Kp= Parabolic rateconstant, C = constant

Some metals follow logarithmic ratelawW = KeLog(Ct + A) C, A = constants, Ke= logarithmic

rate constant

Examples:-Al, Cu, Fe (at slightly elevated temperature)

W=weight gained

per unit area

KL = linear rate

constant.T = time

Figure 12.31

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Corrosion ControlMaterial 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 polymersin presence of strong inorganic

acids.

Ceramics have better corrosion resistance but arebrittle.

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Coatings

Metallic Coatings:Used to protect metal by

separatingfrom corrosive environment and

serving asanode.

Coating applied through electroplatingor rollbonding.

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.

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Corrosion and Protection Methods

Design

General design rules:Provide allowancefor corrosion in thickness.

Weldrather than rivet to avoid crevice corrosion.

Avoid dissimilar metalsthat can cause galvanic

corrosion.

Avoid excessive stress and stress concentration.

Avoid sharp bendsin pipes to prevent erosioncorrosion.

Design tanks and containers for early draining.design so that parts can be easily replaced.

Design heating systems so that hot spotsdo notoccur.

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Alteration Environment

Lower the temperature Reduces reaction

rate.

Decrease velocityof fluids Reduces

erosion corrosion.

Removing oxygen from liquids reducescorrosion.

Reducing ion concentration decreases

corrosion rate. Adding inhibitors inhibitors are retarding

catalystsand hence reduce corrosion.

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Cathodic Protection

Electrons are suppliedto the metal structure tobe protected.

Example:Fe in acid

Fe Fe2+

+ 2e-

2H++ 2e- H2Corrosion of Fe will be

prevented if electrons

are supplied to steel

structure.

Electrons can be supplied by external DCsupply

or galvanic couplingwith more anodic metal.

Figure 12.33

After M. G. Fontana and N. D. Greene, Corrosion Engineering, 2nded., McGraw-Hill, 1978, p.207.13-34



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Anodic Protection

Externally impressed anodic currents form

protective passive filmson metal and alloy

surfaces.

Anodic currents are applied bypotentiostat toprotect metals that passivate.

Current makes them more passive and decreases

the corrosion rate.

Figure

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metb113 emat 13 corrosion

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