foundations of materials science and engineering third edition · 01/01/2016 · 1/1/2016 5...

1/1/2016

1

CHAPTER

8

Corrosion

1

Foundations of Materials Science and Engineering, 5th Edn. Smith and Hashemi

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display

• Announcement:

• Please fill in Class Evaluation Survey (compulsory)

• http://goo.gl/forms/i8vCyteO0Y

• Access from www.afendirojan.wordpress.com

2



3

ISSUES TO ADDRESS...

• How does corrosion occur?

• Which metals are most likely to corrode?

• What environmental parameters affect

corrosion rate?

• How do we prevent or control corrosion?



4

• Corrosion: -- the destructive electrochemical attack of a material.

-- 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**

* H.H. Uhlig and W.R. Revie, Corrosion and Corrosion Control: An Introduction

to Corrosion Science and Engineering, 3rd ed., John Wiley and Sons, Inc.,

1985.

**Economic Report of the President (1998).

Photos courtesy L.M. Maestas, Sandia

National Labs. Used with permission.

THE COST OF CORROSION



5

CORROSION IN A GRAPEFRUIT

Zn 2+

2e - oxidation reaction

Acid

H + H +

H +

H +

H +

H +

H + - +

Zn (anode) Cu (cathode)

O2 4H 4e 2H2O

2H 2e H2(gas)

reduction reactions

Zn Zn2+ 2e



Introduction

• Corrosion: Deterioration of a metal

resulting from chemical attack by its

environment.

• Rate of corrosion depends upon

temperature and concentration of

reactants and products.

• Metals have free electrons that setup

electrochemical cells within their

structure.

• Metals have tendency to go back to

low energy state by corroding.

• Ceramics and polymers suffer

corrosion by direct chemical attack.

6

http://goo.gl/forms/i8vCyteO0Y

http://goo.gl/forms/i8vCyteO0Y

1/1/2016

2



Oxidation – Reduction Reactions

• A metal (E.g. – Zn) placed in HCL undergoes corrosion.

Zn + 2HCL ZnCl2 + H2

or

Zn + 2H+ Zn2+ + H2

also

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.

7 Foundations of Materials Science and Engineering, 5th Edn. Smith and Hashemi


8

• Two reactions are necessary: -- oxidation reaction:

-- reduction reaction:

Zn Zn2 2e

2H 2e H2(gas)

• Other reduction reactions in solutions with dissolved oxygen:

-- acidic solution -- neutral or basic solution

O2 4H 4e 2H2O

O2 2H2O 4e 4(OH)

Adapted from Fig. 17.1,

Callister & Rethwisch 8e.

(Fig. 17.1 is from M.G.

Fontana, Corrosion

Engineering, 3rd ed., McGraw-

Hill Book Company, 1986.)

ELECTROCHEMICAL CORROSION

Zinc

Oxidation reaction Zn Zn 2+

2e - Acid solution

reduction reaction

H + H +

H 2 (gas)

H +

H +

H +

H +

H +

flow of e- in the metal

Ex: consider the corrosion of zinc in an acid solution



Standard electrode Half-Cell Potential of Metals

• Oxidation/Reduction half cell potentials are compared with standard hydrogen ion

half cell potential.

• Voltage of metal (E.g.-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)



10

STANDARD HYDROGEN ELECTRODE

• Two outcomes:

0o

metal V (relative to Pt)

Standard Electrode Potential Adapted from Fig. 17.2,


-- Corrosion

-- Metal is the anode (-)

Pla

tinum

meta

l, M

M n+ ions

ne - H2(gas)

25ºC 1M M n+ sol’n 1M H + sol’n

2e -

e - e -

H +

H +

-- Electrodeposition

-- Metal is the cathode (+)

M n+ ions

ne -

e - e -

25ºC 1M M n+ sol’n 1M H + sol’n

Pla

tinum

meta

l, M

H +

H + 2e -

0o

metal V (relative to Pt)

H2(gas)



11

STANDARD EMF SERIES

metal o

• Metal with smaller

V corrodes. • EMF series

Au

Cu

Pb

Sn

Ni

Co

Cd

Fe

Cr

Zn

Al

Mg

Na

K

+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

metal V metal o

Data based on Table 17.1,

Callister 8e.

mo

re a

no

dic

m

ore

ca

tho

dic

DV =

0.153V

o

Adapted from Fig. 17.2,


-

1.0 M

Ni 2+ solution

1.0 M

Cd 2 + solution

+

25ºC Ni Cd

Cd o

Ni o

• Ex: Cd-Ni cell

V < V Cd corrodes



Macroscopic Galvanic Cells with 1M Electrolyte

• Two dissimilar metal electrodes immersed in solution of their own ions.

• Electrode that has more

negative oxidation potential

will be oxidized.

Zn Zn2+, Cu2+ Cu

Half cell reactions are

Zn Zn 2+ + 2e- E0 = -0.763 V

Cu Cu2+ + 2e- E0 = + 0.337 V

Or Cu2+ + 2e- Cu E0 = -0.337 V (negative sign)

Adding two reactions,

Zn + Cu2+ zn2+ + Cu E0cell = -1.1V

Oxidized Reduced

12

1/1/2016

3



Galvanic Cells With Electrolytes not 1M

• If the concentration of electrolyte surrounding anode is not

I molar, driving force for corrosion is greater.

• There will be more negative emf half cell reaction

M Mn+ + ne-

• Nernst Equation:

ionCn

EE log0592.0

0

E = Net efm of half cell

E0 = Standard emf of half cell

N = Number of electrons transferred

Cion = Molar concentration of ions.



Galvanic Cells With Acid or Alkaline Electrolytes

• Consider iron and copper electrodes in acidic electrolyte.

• Since standard electrode potential of Fe to oxidize is –

0.44 , compared to 0.337 of copper,

Fe Fe2+ + 2e-

• Since there are no copper ions to reduce

2H+ + 2e- H2

• If electrolyte contains oxidizing agent

O2 + 4H4+ + 4e- 2H2O

• If electrolyte is neutral or basic,

O2 + 2H2O + 4e- 4OH-

14



Microscopic Galvanic Cell Corrosion of Single Electrode

• When single electrode is immersed in an electrolyte, microscopic cathodes and anodes are formed due to structural irregularities.

• Oxidation reaction occurs at local anode and reduction reaction at local cathode.

• If iron is immersed in

oxygenated water,

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

Fe Fe2+ + 2e-

O2 + 2H2O + 4e- 4OH-



Concentration of Galvanic Cells

• If a concentration cell is created by immersing 2 electrodes in electrolytes of different concentrations of same ion, electrode in dilute electrolyte will be anode.

• Example:- 2 Fe electrodes

immersed in electrolytes

of 0.001M and 0.01 M

Fe 2+ electrolyte.

E fe2+ = E0 + 0.0296 log Cion

for 0.001M E Fe2+ = -0.44V + 0.0296 log 0.001 = -0.529 V

for 0.01M E Fe2+ = -0.44V + 0.0296 log 0.01 = -0.499 V

• Since –0.529 V is more negative than –0.499 V, electrode in 0.001 M solution is anode and gets corroded.

16



17

EFFECT OF SOLUTION CONCENTRATION AND

TEMPERATURE

• Ex: Cd-Ni cell with

standard 1 M solutions

VNi

o VCd

o 0.153 V

-

Ni

1.0 M

Ni 2+ solution

1.0 M

Cd 2 + solution

+

Cd 25ºC

• Ex: Cd-Ni cell with

non-standard solutions

Y

Xln

nF

RTVVVV o

Cd

o

NiCdNi

n = #e- per unit

oxid/red

reaction

(= 2 here) F =

Faraday's

constant

= 96,500

C/mol. • Reduce VNi - VCd by

-- increasing X

-- decreasing Y

-- increasing T

- +

Ni

Y M

Ni 2+ solution

X M

Cd 2 + solution

Cd T



Oxygen Concentration Cells

• If two electrodes are immersed in electrolytes of different

oxygen concentrations, electrode in low-oxygen content

electrolyte is anode.

• Example: Two iron electrodes, one in low oxygen

concentration water and another in high oxygen

concentration water.

Anode reaction : Fe Fe2+ + 2e-

Cathode reaction: O2 + 2H2O + 4e- 4OH-

• Since cathode reaction requires O2 and electrons, high

concentration oxygen is cathode.

18

1/1/2016

4



Grain – Grain boundary Electrochemical cells

• Grain boundaries are more anodic and hence get

corroded by electrochemical attack.

• Grain boundaries are at higher energy.

• Impurities migrate to grain boundaries.

• Solute segregation might cause grain boundaries to

become more cathodic.

Cartridge Brass

Grain

Boundary

Grain boundary

(anode)

Grain boundary

(cathode)

anode



Multiple Phase Electrochemical Cells

• In multiple alloys, one phase is more anodic than another.

• Corrosion rates are higher in multiphase alloys.

• Example: In pearlite gray cast iron, graphite flake is

cathodic than surrounding pearlite matrix.

Anodic pearlite corrodes

• Steel, in martensitic condition

(single phase) after quenching

from austenitic condition, has

better corrosion resistance.

• Impurities in metals leads to precipitation of intermetallic

phases and hence forms anodic and cathodic regions

leading to corrosion.

Figure 12.10

20



Rate of Uniform Corrosion

• Faraday’s 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 consumed

F = Faradays Constant, A = area Cm2

• Corrosion rate is expressed as weight loss per unit area per

unit time or loss in depth per unit time.

nF

iAtM

nF

ItMW



Corrosion Reaction and Polarization

• When a metal corrodes, the potentials of local cathode and

anode are not at equilibrium.

• Polarization: Displacement of

electrode potential from their

equilibrium values to some

intermediate value and cre-

ation of net current flow.

• Point A : equilibrium potential and current density of Zn

• Point B : equilibrium potential and current density of H

• Point C : Intermediate point

Zn in acid

solution

22



Activation and Concentration Polarization

• Activation polarization:

Electrochemical reactions that

are controlled by a slow step in

a reaction sequence.

• There is a critical activation

energy to surmount energy

barrier associated with slowest

step.

• Concentration polarization:

Associate with electrochemical

reaction and controlled by diffusion

of ions.

• Example: Reduction rate of H+ ions

at surface is controlled by diffusion

of H+ ions into metal surface.



Passivation

• Passivation is loss of chemical reactivity in presence of a

environmental condition.

Formation of surface layer of 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.

24

1/1/2016

5



Polarization Curve

• Polarization curve shows how the potential of a metal

varies with current density.

• As the electrode potential

is made more positive, the

metal behaves as an active

metal.

• When potential reaches Epp

(primary passive potential)

current density decreases an

hence the corrosion rate.

• Further increase in potential makes metal active again.



The Galvanic Series.

• Many metals do not

behave as galvanic cells

due to passive films.

• Galvanic series gives the

cathodic, anodic

relationship between the

metals.

• In flowing seawater,

Zinc is more active than

aluminum.

• Series is determined

experimentally for every

corrosive environment.

26



27

GALVANIC SERIES

• Ranking of the reactivity of metals/alloys in seawater

Based on Table 17.2, Callister &

Rethwisch 8e. (Source of Table

17.2 is M.G. Fontana, Corrosion

Engineering, 3rd ed., McGraw-

Hill Book Company, 1986.)

Platinum

Gold

Graphite

Titanium

Silver

316 Stainless Steel (passive)

Nickel (passive)

Copper

Nickel (active)

Tin

Lead

316 Stainless Steel (active)

Iron/Steel

Aluminum Alloys

Cadmium

Zinc

Magnesium

mo

re a

no

dic

(a

ctive

) m

ore

ca

tho

dic

(i

ne

rt)



28

• Uniform Attack Oxidation & reduction

reactions occur uniformly

over surfaces.

• Selective Leaching Preferred corrosion of

one element/constituent

[e.g., Zn from brass (Cu-Zn)].

• Stress corrosion Corrosion at crack tips

when a tensile stress

is present.

• Galvanic Dissimilar metals are

physically joined in the

presence of an

electrolyte. The

more anodic metal

corrodes.

• Erosion-corrosion Combined chemical attack and

mechanical wear (e.g., pipe

elbows).

FORMS OF CORROSION

Forms of

corrosion

• Crevice Narrow and

confined spaces.

Fig. 17.15, Callister & Rethwisch 8e. (Fig. 17.15

is courtesy LaQue Center for Corrosion

Technology, Inc.)

Rivet holes

• Intergranular Corrosion along

grain boundaries,

often where precip.

particles form.

Fig. 17.18, Callister &

Rethwisch 8e.

attacked

zones

g.b.

prec.

• Pitting Downward propagation

of small pits and holes.

Fig. 17.17, Callister &

Rethwisch 8e. (Fig. 17.17

from M.G. Fontana,

Corrosion Engineering,

3rd ed., McGraw-Hill Book

Company, 1986.)



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

anodic to steel and corrodes.

Large cathode area to small anode area should be

avoided.



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

period and grows in

direction of gravity.

• Pits initiate at structural

and compositional

heterogeneities.

Pitting of stainless steel 30

1/1/2016

6



Growth of Pit

• 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 bottom, metal chloride + water Metal hydroxide +

free acid.

• Some metals (stainless steel) have better resistance than

others (titanium).



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

crevice to balance positive charge and form metal hydroxide and free acid that causes corrosion.

32



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 (500-8000C) chromium

carbide precipitate along grain boundaries.

• When exposed to corrosive environment, the region next

to grain boundaries become anodic and corrode.



Stress 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.

34



Erosion Corrosion and Cavitation Damage

• 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.



Fretting Corrosion and Selective Leaching

• Fretting corrosion: Occurs at interface between materials

under load subjected to vibration and slip.

Metal fragments get oxidized and act as abrasives

between the surfaces.

• Selective leaching: Selective removal of one element of

alloy by corrosion.

Example: Dezincification Selective removal of

zinc from copper and brasses.

Weakens the alloy as single metal might not have

same strength as the alloy.

36

1/1/2016

7



Hydrogen Damage

• Load carrying capacity of a metallic component reduced

due to interaction with atomic/molecular hydrogen.

• Happens in low carbon and alloy steels, aluminum alloys

and titanium alloys.

• Cracking, blistering, hydride formation, reduced ductility

(hydrogen embrittlement).

• Caused due to the diffusion of hydrogen into metal.

• Bakeout is a process applied to the component to diffuse

the hydrogen out of the metal.



Oxidation- Protective Oxide Films

• Oxides form on metals due to reaction with air.

• Degree to which oxide films form depends on following factors.

Volume ratio of oxide to metal consumed after oxidation should be close to 1.

Good adherence.

High melting point of the film.

Low oxide pressure.

Coefficient of expansion equal to that of metal.

High temperature plasticity.

Low conductivity and diffusion coefficients of metal ions and oxygen.

38



Mechanisms of Oxidation

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

• Reduction partial reaction: ½ O2 + 2e- O2-

• Oxidation starts by lateral expansion of discrete oxide

nuclei.

• Metal diffuses as electrons or cations across oxide films.

• Sometimes O2- ions diffuse to oxide metal interface and

electrons diffuse to oxide gas interface.



Oxidation Rates

• Oxidation rate is expressed

as weight gained per unit area.

• Linear oxidation behavior

W = KLt

• If ion diffusion is controlling the step (Eg – Fe, Cu)

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

• Some metals follow logarithmic rate law

W = Ke Log(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

40



41

-- Use metals that passivate - These metals form a thin,

adhering oxide layer that

slows corrosion.

• Lower the temperature (reduces rates of oxidation and

reduction)

CORROSION PREVENTION (i)

Metal (e.g., Al, stainless steel)

Metal oxide

• Apply physical barriers -- e.g., films and coatings

• Materials Selection

-- Use metals that are relatively unreactive in the

corrosion environment -- e.g., Ni in basic solutions



42

• Add inhibitors (substances added to solution that decrease

its reactivity) -- Slow oxidation/reduction reactions by removing reactants

(e.g., remove O2 gas by reacting it w/an inhibitor).

-- Slow oxidation reaction by attaching species to

the surface.

CORROSION PREVENTION (ii)

Adapted

from Fig.

17.22(a),

Callister &

Rethwisch

8e.

Using a sacrificial anode

steel pipe

Mg anode

Cu wire e -

Earth

Mg 2+

• Cathodic (or sacrificial) protection -- Attach a more anodic material to the one to be protected.

Adapted

from Fig.

17.23,

Callister &

Rethwisch

8e. steel

zinc zinc

Zn 2+

2e - 2e -

e.g., zinc-coated nail

Galvanized Steel

e.g., Mg Anode

1/1/2016

8



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.



Coatings

• Metallic Coatings: Used to protect metal by separating

from corrosive environment and serving as anode.

Coating applied through electroplating or roll

bonding.

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.

44



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 early draining.

design so that parts can be easily replaced.

Design heating systems so that hot spots do not occur.



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 inhibitors are retarding

catalysts and hence reduce corrosion.

46



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 or galvanic coupling with more anodic metal.



Anodic Protection

• Externally impressed anodic currents form protective

passive films on metal and alloy surfaces.

• Anodic currents are applied by potentiostat to protect

metals that passivate.

• Current makes them more passive and decreases the

corrosion rate.

48

1/1/2016

9



49

• Metallic corrosion involves electrochemical reactions -- electrons are given up by metals in an oxidation reaction

-- these electrons are consumed in a reduction reaction

• Metals and alloys are ranked according to their

corrosiveness in standard emf and galvanic series.

• Temperature and solution composition affect corrosion

rates.

• Forms of corrosion are classified according to mechanism

• Corrosion may be prevented or controlled by: -- materials selection

-- reducing the temperature

-- applying physical barriers

-- adding inhibitors

-- cathodic protection

SUMMARY

foundations of materials science and engineering third edition · 01/01/2016 · 1/1/2016 5...

Documents