Forms_of_Corrosion.pdf

8/14/2019 Forms_of_Corrosion.pdf

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Corrosion Engineering

N Balasubramanian

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CHE3166: Materials and Corrosion

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Corrosion:Corrosion: ?

Corrosion: Destruction of a material by reactionCorrosion: Destruction of a material by reaction

with its environment.with its environment.

Corrodere (Latin)Corrodere (Latin) -- To Eat AwayTo Eat Away

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The basic corrosion reaction

M M+ + e-

e a ecomes a ca on

Oxidation reaction

Electrochemical reaction

Need to have a corresponding reduction

reaction

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Categories

Infrastructure

Utilities

Cost of Corrosion

Transportation

Production & Manufacturing

Government

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Motor Vehicles 23.4

Ships 2.7

Aircraft 2.2

Oil & Gas Exploration & Production 1.4

Mining 0.1

Petroleum Refining 3.7

Chemical, Petrochemical, & Pharmaceutical 1.7

Cost of Corrosion

Railroad Cars 0.5

Hazardous Materials Transport 0.9

TOTAL: 29.7

Pulp & Paper 6.0

Agricultural Production 1.1

Food Processing 1.1

Electronics -

Home Appliances 1.5

TOTAL 17.6


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Cost of Corrosion

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The eight forms are:

forms of corrosion

5. intergranular corrosion. , ,

2. galvanic, or two-metal corrosion,

3. crevice corrosion,

4.pitting

6. selective leaching, or parting

7. erosion corrosion, and

8. stress corrosion.

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Forms of Corrosion

Uniform CorrosionUniform Corrosion

Galvanic CorrosionGalvanic Corrosion

Crevice CorrosionCrevice Corrosion

ErosionErosion--corrosioncorrosion

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Uniform Corrosion

Uniform attack is the most common form of corrosion. It is normally

characterized by a chemical or electrochemical reaction which

proceeds uniformly over the entire exposed surface or over a large

area. The metal becomes thinner and eventually fails. For example, a

piece of steel or zinc immersed in dilute sulfuric acid will normally

dissolve at a uniform rate over its entire surface. A sheet iron roof willshow essentiall the same de ree of rustin over its entire outside

surface.

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Forms Of Corrosion

Uniform CorrosionUniform Corrosion

Galvanic CorrosionGalvanic Corrosion

PittingPitting



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Dissimilar Metal Couples

GraphiteTitanium

Alloy C-276

Stainless Steel

OK?

Copper

Cast Iron

SteelAluminum

Zinc

OK

Can BeBad

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Galvanic Corrosion

A potential difference usually exists between two dissimilar metals

when they are immersed in a corrosive or conductive solution. If these

metals are placed in contact (or otherwise electrically connected), this

potential difference produces electron flow between them. Corrosion of

the less corrosion-resistant metal is usually increased and attack of the,

these metals when they are not in contact. The less resistant metal

becomes anodic and the more resistant metal cathodic. Usually the-

cathode or cathodic metal corrodes very little or not at all in this type of

couple. Because of the electric currents and dissimilar metals involved,

this form of corrosion is called galvanic, or two-metal, corrosion. It is

electrochemical corrosion, but we shall restrict the term galvanic to

dissimilar-metal effects for purposes of clarity.

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Galvanic Corrosion

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Forms Of Corrosion

Uniform CorrosionUniform Corrosion Galvanic CorrosionGalvanic Corrosion

PittingPitting

Crevice CorrosionCrevice Corrosion ros onros on--corros oncorros on

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Pitting Corrosion

Pitting is a form of extremely localized attack that results in holes in the

metal. These holes may be small or large in diameter, but in most cases they

are relatively small. Pits are sometimes isolated or so close together that

they look like a rough surface. Generally a pit may be described as a cavity

or hole with the surface diameter about the same as or less than the depth.

Pitting is one of the most destructive and insidious forms of corrosion. It

weight loss of the entire structure. It is often difficult to detect pits because

of their small size and because the pits are often covered with corrosion

products. In addition, it is difficult to measure quantitatively and compare

the extent of pitting because of the varying depths and numbers of pits that

may occur under identical conditions.

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Pitting Corrosion

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Forms of Corrosion

Uniform CorrosionUniform Corrosion Galvanic CorrosionGalvanic Corrosion

PittingPitting



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Crevice Corrosion

Intense localized corrosion frequently occurs within crevices and other shielded

areas on metal surfaces exposed to corrosives. This type of attack is usually

associated with small volumes of stagnant solution caused by holes, gasket

surfaces, lap joints, surface deposits, and crevices under bolt and rivet heads. As a

result, this form of corrosion is called crevice corrosion or, sometimes, deposit or

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Crevice Corrosion

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Forms Of Corrosion

Erosion corrosion is the acceleration or increase in rate of deterioration or

attack on a metal because of relative movement between a corrosive fluid and

the metal surface. Generally, this movement is quite rapid, and mechanical

wear effects or abrasion are involved. Metal is removed from the surface as

dissolved ions, or it forms solid corrosion products which are mechanically

swept from the metal surface. Sometimes, movement of the environment

Uniform CorrosionUniform Corrosion Galvanic CorrosionGalvanic Corrosion PittingPitting



,

conditions, but this is not erosion corrosion because deterioration is not

increased.

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Erosion-corrosion

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Forms of Corrosion

Selective leaching is the removal of one element from a solid alloy by corrosion

processes. The most common example is the selective removal of zinc in brass

alloys (dezincification). Similar processes occur in other alloy systems in which

aluminum; iron, cobalt, chromium, and other elements are removed. Selective

leaching is the general term to describe these processes, and its use precludes

the creation of terms such as dealuminumification, decobaltification, etc. Parting is

Selective LeachingSelective Leaching Intergranular Attack

Corrosion Fatigue

Environmental

Cracking

, .

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Selective Leaching

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Forms Of Corrosion

Grain boundary effects are of little or no consequence in most applications

or uses of metals. If a metal corrodes, uniform attack results since grain

boundaries are usually only slightly more reactive than the matrix.

However, under certain conditions, grain interfaces are very reactive and

intergranular corrosion results. Localized attack at and adjacent to grain

boundaries, with relatively little corrosion of the grains, is intergranular

corrosion. The alloy disintegrates (grains fall out) and/or loses its strength.

Selective Leaching

Intergranular Attack (IGA)Intergranular Attack (IGA)

Corrosion Fatigue

Environmental Cracking26


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Intergranular Corrosion


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Intergranular Corrosion

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Corrosion Fatigue

60

80

100

120

ksi

)

ksi

)

0

20

40

10,000 100,000 1,000,000 10,000,000 100,000,000

In AirIn Air In SeawaterIn Seawater

Stress

Stress

Cycles to FailureCycles to Failure

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Corrosion Fatigue Fracture

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Stress Corrosion Cracking

Alloy Susceptible

Tensile Stress

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Vertical Heat Exchanger

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Caustic Carryover In Steam

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Forms of Corrosion

Erosion corrosion is the acceleration or increase in rate of deterioration or

attack on a metal because of relative movement between a corrosive fluid

and the metal surface. Generally, this movement is quite rapid, and

mechanical wear effects or abrasion are involved. Metal is removed from the

surface as dissolved ions, or it forms solid corrosion products which are

mechanically swept from the metal surface. Sometimes, movement of the

Selective Leaching Intergranular Attack

Corrosion Fatigue

Environmental CrackingEnvironmental Cracking

,

under stagnant conditions, but this is not erosion corrosion because

deterioration is not increased.


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Factors Favoring

SCC Stainless Steels

Temperature > 55 C

Chlorides

Evaporative Conditions

Pit - Stress Raiser

Residual Or Applied Stress

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Mechanical And Physical Factors

Strength, Ductility

Formability

,

Thermal Conductivity

Thermal Expansion

Unique Properties

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Other Factors

Safety

Compatibility

Cost

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Economics

Managements Language Is EconomicsManagements Language Is Economics

Use Life Cycle Costs To Compare MaterialsUse Life Cycle Costs To Compare Materials

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Materials Selection

Define Environment

Define Required Equipment Define Test And Inspection

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Define Environment

Major Reactants

Minor Ingredients

pH, Contaminants, Aeration

Process Conditions

Temperature, Pressure, Excursions

Catalytic/inhibitive Effects

e.g., Metallic Ions

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Define Equipment

Pressure Vessels

Valves And Piping

Tankage es ys ems

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Test & Inspection

Fabrication Verify Materials, Workmanship, Dimensions,

Storage Conditions

Testing & Inspection

Develop Program

Consider Corrosion Test Locations

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Dont Expect Inspect

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CHE3166Corrosion: Thermodynamics

Net reaction is oxidation of zinc by copper(II) ions:

Zn(s) + Cu2+ Zn2+ + Cu(s)

Half reactions take place in separate locations:

Left electrode: Zn(s) Zn2+ + 2e oxidation

Right electrode: Cu2++ 2e Cu(s) reduction

Electrochemical cells allow measurement and control of a redox (reduction/oxidation)

reaction. Also, we can force the reaction to proceed in its non-spontaneous, or reversedirection by connecting a source of current to the two electrodes.

To sustain the cell reaction, the charge carried by the electrons through the external

circuit must be accompanied by a compensating transport of ions between the two cells.

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To sustain the cell reaction, the charge carried by

the electrons through the external circuit must be

accompanied by a compensating transport of ions

between the two cells.

Relative amounts of charge can be carried by

Galvanic cells and electrodes

negat ve or pos t ve ons epen s on t e r re at ve

mobilities) through the solution.

Salt bridge, consists of an intermediate

compartment filled with saturated salt solution and

fitted with porous barriers at each end, is used forprecise measurements. The purpose of salt bridge is

to minimize the natural potential difference

(junction potential).

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Consider zinc in a solution of hydrochloric acid

Thermodynamics

Zn + 2 ClZnCl2+H2

This reaction can be separated into two half-cell reactions

n n + e

2H+ + 2e H2

potential ea)

Reduction reaction (cathodic reaction withpotential ec)

has an associated free-energy change (G)

G= nFE= nF ea+ ec( )

n=number of electrons exchanged, F=96500 C/e-, E=potential

Nernst Equation

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Standard hydrogen electrode

Pt is inert and acts as

catalyst site

1 atm H2 gas is the standards a e

Solid Zn and 1N Zn2+ is the standard state

Measured potential difference V is 0.762V

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Standard hydrogen electrode

For a pure Cu electrode in a

1 N Cu 2+ solution, the

potential difference in this

cell is 1.104 V

Zn2+ + 2e Zn 0.762V( )

Cu2+

+ 2e

Cu +0.342V( )

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Electromotive Force Series

Zn + 2HClZnCl2+ 2

or

ZnZn2+ + 2e

2H+ + 2e H2

= ea + ec

E= +0.762 + 0

= +0.762V

G= nFE< 0

At standard state, reaction

will occur

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The half-cell reaction with the

more active (negative) half-

cell potential always proceeds

as an oxidation, and thereaction with the more noble

half-cell potential always

proceeds as a reduction in the

spontaneous reaction

produced by the pairShipboard corrosion

protection

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Question. Is copper (Cu)

oxidized (dissolved) by ferric

(Fe3+) ions?

Ferric reduction is noble relative

to copper so copper is oxidized

(copper becomes an ion and

gives off electrons)

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Reaction Direction

Written as reduction

reactions from left to right

y rogen re uc on s

chosen (arbitrarily) as the

reference and assigned

zero potential

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Standard conditions

Standard states:

For a solid, the solid is taken as unit activity

For a gas, 1 atm pressure is taken as unit activity

For dilute or strongly dissociated solutes typically found in

most instances of corrosion, activity is reasonably approximated

A( ) = CA

A=activity of atom A= activity coefficient (tabulated)

CA=concentration (g/l) of atom A

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Activities

Gases A=p (partial pressure in atm)

Pure li uids and solids A=1

Solutes A=M (molarity)

Liquids in solutions A=X (mole fraction)

Solvent in dilute solution A=1

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Non-standard conditions

Analyze a general half-cell equation between atom A and atom B in

an aqueous solution

aA +mH+ +ne = bB +dH2O

or reac an s an pro uc s n e s an ar s a e

G0 = bGB0 +dGH2O

0( ) aGA0 +mGH+0( )

For reactants and products in a non-standard state

G= bGB +dGH2O( ) aGA +mGH+( )59


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Non-standard conditions

The departure of free energy in the non-standard state from the

standard state is:

G G0 = b GB GB0( )+d GH2O GH2O

0( )[ ] a GA GA0( )+m GH+ GH+0( )[ ]

The corrected concentration A available for the reaction is known as

the activity of A and is given by:

a GA GA0( )=aRTln(A) =RTlnA( )

a

G G0 =RTln B( )

bH2O( )

d

A( )a

H+( )m

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Nernst Equation

G G0 =RTln B( )

bH2O( )

d

A( )a

H+( )m

= 0 RT B( )b

H2O( )d

G= nFE= nF ea + ec( )Recall

2.3 T = 0.059

At 25C

nF A( )a

H+( )m

e= e0 +2.3RT

nFlog

A( )a

H+( )m

B( )b

H2O( )d

e= e0 +0.059

nlog

A( )a

B( )b

m

n0.059pH

pH= logH+( )

Recall

Activity of water=1 for

aqueous solutions

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Electrochemical potential

e= e0 +0.059

nlog

A( )a

B( )b

m

n0.059pH

Examine specific half-cell reactions

+ e = 2

eH+ /H2

= eH+ /H2

0 0.059pH

,

so activity=1

As the activity of any dissolved oxidizer (e.g., H+) increases,

more noble or positive potentials are routinely measured

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Electrochemical potential

The above reaction occurs in an acid solution. An

equivalent reaction in neutral or alkaline solutions is:

2H+ + 2e =H2 eH+ /H2 = eH+ /H20

0.059pH

0

The electrochemical evolution of hydrogen is the

decomposition of water

2 + e = 2+ eH+ /H2= e

H+ /H2 . p

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Potential/pH (Pourbaix) Diagram

eH+ /H2

= eH+ /H2

0 0.059pH

oxygen evolution

and acidification

For hydrogen evolution at a fixedpotential

2H+

+ 2e

=H2

+ hydrogen evolution

and alkalization

decreases), we are less likely to

form H2 gas

Note that, consistent with the half-cell potential definitions,

the hydrogen line goes through zero potential at zero pH

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oxygen evolution

and acidification

At higher potential

O2+ 4H+

+ 4e

2H2O

is feasible

eO2 /H2O = eO2 /H2O 0.059pHhydrogen evolutionand alkalization

Note that, consistent with the half-cell potential definitions,the oxygen line goes through 1.229V at zero pH and 0.401V atpH=14.

This reaction is the basis for water electrolysis65


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Corrosion-soluble ions of the

metal are stable

Passivation- oxides are stable

Immunity-reduced form of the

metal is stable

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Reaction 1: Metal oxidizes to

aqueous cations

Pourbaix diagram for Al

M=Mn +

+ne

l= l3+ + 3e

eAl/Al3+

= e0 +0.059

nlog

A( )a

B( )b

m

n0.059pH

eAl/Al3+

= 1.662 +0.059

3logAl3+( )

Independent of pH since

no H+ is involved.

Only depends on Al3+

activity

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Reaction 2: Metal reacts to metal

hydroxide or oxide


M+nH2O =M OH( )n+nH+ +ne

+

2 = 2 3 e

eAl/Al2O3 = 1.55 0.059pH

At higher pH Al2O3 is formed. At lower pH Al2O3 dissolves to Al3+

Intersection depends depends on Al3+ activity (dashed lines areportions of the reactions with no significance)

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The reaction rate constant for


2Al3+

+ 3H2O =Al2O3+ 6H+

K=H+( )

6

Al3+( )2

= 1011.4

For (Al3+)=10-6,

logK= 6logH+( ) 2log Al3+( )= 6pH 2log Al3+( )

pH= 3.9 Independent of potential

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Reaction 3: Metal reacts to form

soluble aqueous anions


= n2m + m

Al+ 2H2O =AlO2 + 4H+ + 3e

eAl/AlO2

= eAl/AlO2

0 +0.059

3logAlO2

( )

4

30.059pH

eAl/AlO2

= 1.262 + 0.020log AlO2( ) 0.079pH

At higher pH, Al2O3 dissolves to AlO2-

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P b i Di F


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Pourbaix Diagram Fe

Example 2: stability diagram for iron (little

bit complicated).

Horizontal lines: reactions areinvolved with electrons (potential),

but independent of pH (H(+a) or OH(-a)

ions):

Fe2+=Fe3++e

Ee=0.771+0.0591lg(aFe3+/aFe

2+)

Fe=Fe2++2e

Ee=0.771+0.0591lg(aFe3+/aFe

2+)

Vertical lines: reactions are

involved with pH, but independent of

electrons:

Fe3++H2O=FeOH2++H+

lg(aFeOH2+/aFe

3+)=-2.22+pH

Fe2++2H2O=Fe(OH)2+2H+

lg(aFe2+)=13.37-2pH

Diagonal lines: reactions are

involved both pH and electrons:

Fe2++H2O=FeOH2++H++eEe=0.877-0.0591pH+0.0591lg(aFe(OH)

2+/aFe2+)

Fe2++3H2O=Fe(OH)3+3H++e

Ee=0.748-0.1773pH-0.0591lgaFe2+

Fe+2H2O=Fe(OH)2+2H++2e

Ee=-0.045-0.0591pH

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Standard Hydrogen Electrode

Platinum Foil suspended in

sulfuric acid with unit activity of

H+. Platinum is noble and does

not participate (acts as catalyst

Connected to another half-cell

through a solution bridge that

allows for charge but not masstransfer

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Ohter Electrodes

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Questions?

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