Corrosion and Degradation of Materials

Interaction of the material with its environment.Changes in properties due to the interaction (mechanical – ductility or strength, physical properties, or appearance).Mechanisms of Deterioration :These depend on the type of material:Metals : Either materials loss by dissolution (corrosion) or formation fo a non-metallic scale (oxidation).Ceramic : Highly resistance to deterioration. Deterioration at high temperature or extreme environments (corrosion).Polymers : May dissolve or absorb and swell when exposed to liquid solvents, or deteriorate in the presence of UV and heat (degradation).

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

Cost of Corrosion

* 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).

Corrosion of MetalsDestructive electrochemical attack of a metal. Usually begins at the surface.It can be classified in two different types:• Dry corrosion – Oxidation (scaling). It takes place in gaseous atmospheres, normally air• Wet corrosion – Electrochemical corrosion. It takes place in aqueous solution, humid environments.

OxidationMechanisms

The process of an oxide layer formation is an electrochemical process. In the case of iron heated to high temperature the reaction is as follows:

2Fe + O2 2FeOOxidation Fe Fe2+ + 2e-

Reduction ½ O2 + 2e- O2-

For the oxide layer to increase is necessary that electrons be conducted to the scale gas interface at which reduction occurs.In addition, M2+ must diffuse away from the metal-scale interface, and/or O2-

ions must diffuse towards the metal-scale interface. The oxide scale layer serves as an electrolyte through which ions diffuse and as an electrical circuit for the passage of electrons.

The scale may protect the metal from rapid oxidation (barrier to ionic diffusion and/or electrical conduction).

Scale TypesWhere AO is the molecular or formula weight of the oxide; AM is the atomic weight of the metal, and ρO and ρM are the oxide and metal densities respectively.When P-B < 1 the oxide film tends to be porous and unprotective.When P-B > 2-3 the oxide film has compressive stresses that may crack and flake off, continuously exposing a new fresh and unprotected surface.

OM

MO

AA

ratioBPRatioBedworthPillingρρ

=−=−

If P-B = 1 is the ideal ratio for a protective film.

Pilling-Bedworth ratios for a number of metalsProtective Nonprotective

Al 1.28 K 0.45Pb 1.40 Na 0.57Ni 1.52 Ag 1.59Cu 1.68 Ti 1.95Fe 1.77 Mo 3.40Cr 1.99 W 3.40

Other factors also influence the oxidation resistance imparted by the film; these include the degree of adherence between film and metal, comparable coefficients of thermal expansion, good high temperature plasticity, high melting point.

Electrochemical Corrosion (wet corrosion)Metal atoms characteristically lose or give up electron (oxidation).The site at which oxidation takes place is called the anode. −+

−+

−+

−+

+→

+→

+→

+→

eAlAleFeFeeFeFe

neMM n

332

3

3

2

The electrons generated from each metal atom that is oxidized must be transferred to and become a part of another chemical species (reduction). Examples

( )( )

MneMMeM

OHeOHO

OHeHO

HeH

n

nn

→+

→+

→++

→++

→+

−+

+−−+

−−

−+

−+

122

22

2

442

244

22Corrosion in acid solution having a high concentration of hydrogen ions (H+)…………..

Acid solutions having dissolved oxygen……….

For neutral or basic aqueous solution …………

Metals ions present in the solution may also be reduced (if exists in different valence states).........

The metal may be totally reduced………………..

The location at which reduction - cathode

Example•Consider Zn metal immersed in an acid solution containing H+ ions.

)(2

)(222

22

2

2

gasHZnHZn

gasHeHeZnZn

+→+

→+

+→

++

−+

−+

•Rusting of iron in water which contains dissolved oxygen. This process occurs in two stages:

( )

( ) ( )3222

22

22

2212

221

OHFeOHOOHFe

OHFeOHFeOHOFe

→++

→+→++ −+ ….dissolved in water

…. it is the familiar rust

Electrode PotentialsConsider the following cells each one immersed in 1M concentration of its ions. (molarity = 1 mol per 1dm3).

CuFeCuFeCueCu

eFeFe

+→+

→+

+→

++

−+

−+

22

2

2

22

Cu2+ ions deposits (electrodeposit) as metallic copper on the copper electrode (cathode)Fe dissolves (corrodes) on the other side of the cell and goes into solution as Fe2+ (anode).Galvanic couple – two metals electrically connected in a liquid electrolyte.An electrical potential or voltage will exists between the two cell halves (0.78V).

++ +→+ 22 ZnFeZnFeThe electrode potential associated with this reaction is 0.323V.Thus, various electrode pairs have different voltages; the magnitude of such a voltage may be though of as representing the driving force for the electrochemical oxidation-reduction reaction.Then, metallic materials may be rated as to their tendency to experience oxidation when coupled to other metals in solutions of their respective ions.

A half cell of a pure metal electrode immersed in a 1M solution of its ions at 25oC is termed a standard half cell.

Standard EMF SeriesIt is convenient to establish a reference point, or reference cell, to which other cell half may be compared.The reference cell, arbitrarily chosen, is the standard hydrogen electrode.

It consists of an inert platinum electrode in a 1M solution of H+ ions, saturated with hydrogen gas that is bubble through the solution at a pressure of 1 atm and a temperature of 25oC.

The platinum electrode does not take part in the electrochemical reaction

4

• Two outcomes:

--Metal sample mass --Metal sample mass

Pla

tin

um

me

tal,

M

Mn+ ions

ne- H2(gas)

25°C 1M Mn+ sol’n 1M H+ sol’n

2e-

e- e-

H+

H+

--Metal is the anode (-) --Metal is the cathode (+)

Vmetalo < 0 (relative to Pt) Vmetal

o > 0 (relative to Pt)

Standard Electrode Potential

STANDARD HYDROGEN (EMF) TEST

Mn+ ions

ne-

e- e-

25°C 1M Mn+ sol’n 1M H+ sol’n

Pla

tin

um

me

tal,

M

H+ H+

2e-

The electromotive force (emf) series is generated by coupling to the standard hydrogen electrode, standard half cells for various metals and ranking them according to measured voltage.It represents the corrosion tendency of various metals.Those at the top, are noble, or chemically inert.Those at the bottom are highly active.

5

• EMF series • Metal with smaller VOMETAL

corrodes.• Ex: Cd-Ni cell

-

Ni

1.0 M

Ni2+ solution

1.0 M

Cd2+ solution

+

Cd 25°C

mo

re a

no

dic

mo

re c

ath

od

ic AuCuPbSnNiCoCdFeCrZnAlMgNaK

+1.420 V+0.340- 0.126- 0.136- 0.250- 0.277- 0.403- 0.440- 0.744- 0.763- 1.662- 2.262- 2.714- 2.924

metal Vmetalo

ΔV = 0.153V

o

Data based on Table 17.1, Callister 6e.

STANDARD EMF SERIES

Consider the generalized reactions

01

022121

0222

0111

VVVMMMM

VMneM

VneMM

Onn

n

n

−+=Δ+→+

+→+

−+→

++

−+

−+

For the reaction to occur spontaneously, ΔVO must be positive.

If it is negative the reaction is reverse.

Influence of Concentration and Temperature on Cell potential.

( ) [ ][ ]

( ) [ ][ ]+

+

+

+

−−=Δ

−−=Δ

n

nOO

n

nOO

MM

nVVV

MM

nFRTVVV

2

112

2

112

log0592.0

lnNerst equation: R =gas constant; n = number of electrons participating; F is the Faraday constant (96,500C/mol;

[M1n+] and [M2

n+] are the molar concentrations of M1 and M2.

In the copper electrode:

2H+ + 2e- 2H H2

in the iron electrode:

Fe2+ + 2(OH)- Fe(OH)2

Soluble in water

4Fe(OH)2 + 2H2O + O2 4Fe(OH)3

(rust)

Example:One half of an electrochemical cell consists of a pure nickel electrode in a

solution of Ni2+ ions; the other half is a cadmium electrode immersed in a Cd2+

solution.

(a) If the cell is a standard one write the spontaneous overall reaction and calculate the voltage generated.

(b) Compute the cell potential at 25oC if the Cd2+ and Ni2+ concentrations are 0.5 and 10-3M respectively. Is the spontaneous reaction direction still the same as for the standard one

Solution

(a)

(b)

( ) VVVVCdNiCdNi

V .- V Ni eNi

V. VeCd CdO

Ni

OCd

153.0403.0205.0

20502

40302

22

2

2

+=−−−=Δ+→+

=→+

−=+→

++

−+

−+

( ) VV 073.010

50.0log2

0592.0153.0 3 =⎟⎠⎞

⎜⎝⎛−+=Δ −

- +

Ni

Y M

Ni2+ solution

X M

Cd2+ solution

Cd T

• Ex: Cd-Ni cell withstandard 1M solutions

VNio − VCd

o = 0.153

• Ex: Cd-Ni cell withnon-standard solutions

VNi − VCd = VNi

o −VCdo −

RTnF

lnXY

-

Ni

1.0 M

Ni2+ solution

1.0 M

Cd2+ solution

+

Cd 25°C

n = #e-

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

• Reduce VNi - VCd by--increasing X--decreasing Y

Galvanic Series• Ranks the reactivity of metals/alloys in seawaterm

ore

an

od

ic

(ac

tive

)m

ore

ca

tho

dic

(i

ne

rt)

PlatinumGoldGraphiteTitaniumSilver316 Stainless SteelNickel (passive)CopperNickel (active)TinLead316 Stainless SteelIron/SteelAluminum AlloysCadmiumZincMagnesium Based on Table 17.2, Callister 6e. (Source of Table 17.2 is M.G.

Fontana, Corrosion Engineering, 3rd ed., McGraw-Hill Book Company, 1986.)

More realistic and practical ranking.

Alloys are the top are cathodicand unreactive, whereas those at the bottom are most anodic.

No voltages are provided.

Forms of CorrosionThe classification according to the manner in which it is manifest.

Uniform Attack - Oxidation & reduction occur uniformly over surface.Galvanic Corrosion - Dissimilar metals are physically joined. The

more anodic one corrodes. (see emf table)Crevice Corrosion - Between two pieces of the same metal.Pitting - Downward propagation of small pits & holes.Intergranular Corrosion - Corrosion along grain boundaries, often

where special phases exist.Selective Leaching - Preferred corrosion of one element/constituent

(e.g., Zn from brass (Cu-Zn)).Erosion-Corrosion - Break down of passivating layer by erosion (pipe

elbows).Stress Corrosion - Stress & corrosion work together at crack tips.Localized Cold Work – Due to the presence of parts in the metal that

are heavily cold workedHydrogen Embrittlement – Produces by hydrogen atoms dissolved in

the metal.

Galvanic Corrosion:

It occurs when two metals or alloys having different compositions are electrically coupled while exposed to an electrolyte.

The more noble metal in that environment will act as cathode andwill be protected by the less noble metal (anode).

Example: (a) copper and steel tube joined in a domestic water heater.

Pitting CorrosionPitting is a form of a very localized surface attack in which small pits or holes form.

Intergranular CorrosionIt occurs preferentially along grain boundaries for some alloys and in specific environments. The net result is that a macroscopic specimen disintegrates along its grain boundaries.Example: Stainless steels heated at temperatures between 500 and 800oC for long periods (weld decay)

Localized Cold WorkDue to the stored energy difference between areas of cold work and annealed material.

Crevice CorrosionIt occurs as a consequence of concentration differences or ions or dissolved gases in the electrolyte solution, and between two regions of the same metal piece.After oxygen has been depleted within the crevice, oxidation of the metal occurs at this position.

Corrosion PreventionSelection of materials: Use of the standard corrosion references to select the appropriate material for the environment.Changing the character of the environment: Lowering temperature, velocity or concentration of the fluids.

•Using inhibitors: These are substances that when added in low concentrations to the environment decreases its corrosiveness. The inhibitor depends on both the alloy and the corrosive environment. They are used mainly in closed systems: e.g. car radiators, steam boilers.•Using physical barriers: Films and coatings applied to the surface. There is a large diversity of metallic and non-metallic coatings with a high degree of surface adhesion. •Cathodic Protection: Very effective. It involves supplying from an external source electrons to the metal to be protected, making it a cathode.

Galvanizing : a layer of zinc is applied to the surface of the steel by hot dipping.

Tin plating: a layer of tin is applied to the surface of the steel.