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485VOLUME V / INSTRUMENTS

9.1.1 Introduction

In the hydrocarbon extraction and treatment industry, allclasses of materials are used: metallic, polymeric, ceramic,composite and cement materials. In terms of quantity andtype of use, metallic materials represent the maincomponent. Evidence of this are, for example, oil wells,refining columns, transportation pipelines, storage tanks andmany other important parts of facilities. Given theirimportance and strategic impact, this chapter will dealmainly with metallic materials and the main forms ofcorrosion affecting them.

Generally speaking, materials in contact with aggressiveenvironments undergo a form of chemical and physical decaywhich, as far as metallic materials in particular are concerned,is known as corrosion. Corrosion can be defined as an attackby atmospheric agents or other aggressive means onmaterials, especially metals, which leads to the slow butprogressive alteration of the characteristics – often not onlyon the surface – of the material. It can also be defined as thedestruction or deterioration of a material by reaction with theenvironment or as the tendency of a metallic material toreturn to its original state, as it is found in nature (Fontana,1986); as a consequence, it is also known as metallurgy inreverse. Corrosion processes tend to bring metallic materials

spontaneously to their most stable thermodynamic state, inwhich they combine with other elements, especially oxygenand sulphur. Starting from this state, metallic materials areobtained (‘extracted’) with metallurgy processes that entailthe supply of large amounts of energy (Fig. 1).

The economic impact of corrosionAn idea of the importance of corrosion in industrial

activities can be obtained from its economic impact. Inindustrialized countries, the cost of corrosion is around 3-4%of gross domestic product, calculated as the sum of the costsof direct damage (such as the cost of damaged materialswhich must be replaced, the cost of replacement operations)and of indirect damage (such as the cost of lost production,plant inactivity, the costs of pollution, poor image in thecommunity, the cost of environmental cleanup, damage topersons and things; Hoar, 1971; Eni-Agip, 1994). Indirectcosts, which are difficult to evaluate, generally exceed directcosts. It has been estimated (Hoar, 1971) that the cost ofcorrosion can be reduced by 15-20% by applying techniquesthat make use of a basic understanding of corrosion; thebest-known of these include: cathodic protection, selectingthe most resistant material, the use of corrosion inhibitors,improvements in design (for example, eliminating stagnantconditions).

9.1

General aspects of corrosion

iron metallurgyprocesses

aggressive environment(gas, process fluids,

etc.)

aggressive environment(soil, sea water,

etc.)

energy

energy

tubes

sheet

rust(oxides)ore

(oxides)

mine

energy

energy

Fig. 1. The corrosion process as metallurgy in reverse (Fontana, 1986).

Morphology of corrosion processesCorrosion phenomena may occur at the surface of

metallic materials in a generalized or localized way(Pedeferri, 2007). Generalized corrosion occurs when theattack affects the entire surface of the material exposed tothe environment, and uniform corrosion when thegeneralized attack takes place in a uniform way. Localizedcorrosion, by contrast, occurs when the attack affects onlysome parts of the surface of the material exposed to theenvironment, with a specific morphology, for example in theform of fissures or cracks, cavities, craters, ulcers. We speakof selective corrosion when specific constituents of thematerial are attacked, such as some phases present as grainsor around the grains.

Corrosion rateAny corrosion process, regardless of the morphology of

attack, entails a loss of mass. The corrosion rate can beexpressed as the mass loss per unit of surface area per unit oftime. Generally speaking, however, it is preferable toconsider the penetration rate of corrosion; in this case, thecorrosion rate is expressed as the decrease in thickness perunit of time. For some forms of corrosion, such as stresscorrosion cracking or corrosion fatigue which lead to theformation of cracks, mass loss and the decrease in thicknessare less important than the time to failure or the growth orpropagation rate of the cracks.

Uniform attackIn conditions of uniform corrosion, in other words an

attack distributed uniformly over the surface of the material,the rate of mass loss per unit of surface area exposed to theaggressive environment expresses the extent of damagecaused by the attack over time and can be calculated with theequation:

[1]

where Dm is the mass loss during time interval t, and A isthe exposed surface area. The most commonly used unit ofmeasurement for the corrosion rate in terms of mass loss isthe mg/dm2�day (mdd). Mass loss becomes important whenwe wish to know the quantity of dissolved metal, forexample to evaluate the pollution produced. An example istin poisoning caused by the tins used to preserve tomatoes.

Generally speaking, thinning is more important thanmass loss, so the uniform corrosion rate is expressed as theloss of thickness, given by:

[2]

where g is the density of the metal. In this case, the mostfrequently used unit of measurement for the corrosion rate ismm�yr. For the most widely used metals (iron, copper andzinc), which have densities between 7 and 8 t/m3, thefollowing rough equivalents can be obtained 1 mdd�5mm�yr; 1 mm�yr�220 mdd. The rate of generalizedcorrosion is usually classified according to the values shownin Table 1.

Localized attackIn conditions of localized corrosion, it is necessary to

distinguish between the rate of mass loss (which expresses a

mean velocity over the whole exposed surface) and the rateof penetration into the area attacked. In the presence oflocalized attack, the loss of efficiency is given, for example,by the perforation of the metal wall (as in the case of a tankor pipe) and not by the metal’s mass loss.

Types and mechanisms of corrosionThe corrosion of metallic materials takes two forms:

high temperature corrosion (or dry corrosion), typical ofmetallic materials operating at high temperatures in thepresence of hot gases, such as the fume side of boilers andgas turbines; wet corrosion, characteristic of materialsexposed to an electrolytic solution such as sea water, soil,concrete polluted by chlorides or carbonated concrete,process fluids. The distinction between wet and drycorrosion derives from the two different mechanismsgoverning the phenomenon: in the first case, anelectrochemical mechanism; in the second, a chemicalmechanism, typical of heterogeneous reactions.

9.1.2 High temperature corrosion

The corrosion of metals in contact with air at temperaturesabove 400°C and up to 1,300°C is known as high temperaturecorrosion. The presence of oxygen leads to the formation ofan oxide scale on the surface of the metal, whilst the presencein the hot gases of some chemical species such as sulphur,sodium and vanadium leads to the formation of salts with alow melting point that react with the metal.

The oxidation of metals and alloys at high temperature isknown and well-documented (ASM, 1987; Revie, 2000). Inorder to predict the formation of the scale and its growth, thethermodynamic conditions and the kinetics of the reactionsinvolved must be considered. The thermodynamic conditionsdetermine if the oxidation reaction proceeds spontaneouslyat the operating temperature, whilst the kinetics determinethe velocity at which the scale growth reaction takes place.

Decay processes at high temperature include: a) thinningdue to the formation of a non-protective scale; b) corrosionby molten salts with evaporation of the corrosion products;c) erosion-corrosion caused by solid particles in suspension;d ) localized attacks at the grain boundaries; e) embrittlementof the material.

Given the fairly extreme operating conditions, whichoften present the risk of catastrophic consequences should afailure occur, the choice of materials generally requiresgreater care than for low temperature applications.

V mAt

Vcorr

corr m= =∆γ γ

,

V mAtcorr m, = ∆

MATERIALS

486 ENCYCLOPAEDIA OF HYDROCARBONS

Table 1. Classification of the corrosion rate

Uniform corrosion rateCorrosion rate

mm/yr

Negligible �50

Low �50-100

Modest 100-500

Severe 500-1,000

Very severe �1,000

Thermodynamic conditionsWhereas wet corrosion processes are electrochemical in

nature, high temperature corrosion works in accordance withthe kinetics of chemical reactions in the gas phase;thermodynamic conditions and solid-phase diffusionprocesses in the products or corrosion scales are thereforeimportant.

Thermodynamic studies show that all metals oxidizespontaneously in the presence of oxygen or in air with theexception of gold and platinum. However, numerous metalscan be used for extremely long periods even at hightemperatures because the oxide growth kinetics aresufficiently slow. This fact lies behind the development ofalloys resistant to oxidation due to the formation of a scalethat acts as a barrier between the metal and the environment,characterized by a low growth rate.

In the absence of aqueous solutions, in other words indry oxygen or air, metals at ambient temperature form aprotective scale 1-10 nm thick which prevents furtheroxidation of the metal. As the temperature rises, thethickness of the film increases, leading to mechanicaldetachment on many metals due to the excessive volume ofthe oxide film.

From a thermodynamic point of view, the oxidation of ametal occurs only if the partial pressure of oxygen atoperating conditions is greater than the dissociation pressureof the metal, calculated using the Ellingham diagram (ASM,1987).

KineticsThe formation process of the oxide layer can be

interpreted with an electrochemical mechanism. Theoxidation reaction is as follows (for the sake of simplicity, adivalent metal has been considered):

[3] M �1�2O2��MO

which can be subdivided into the following twocomplementary reactions:

[4] M��M2��2e�

[5] 1�2O2�2e��O2�

the first of which takes place at the metal-oxide interface andthe second at the oxide-gas interface. The oxide layer mayincrease in thickness if two conditions are met: the electronsmigrate from the metal to the oxide-gas interface where theoxygen reduction reaction takes place and, simultaneously,due to diffusion phenomena, the metal ions move away fromthe metal-oxide interface or the O2�oxygen ions movetowards it, or both diffusion phenomena occur. Differencesin the diffusion rate determine where the film grows: at theoxide-gas interface if the diffusion of the metal ions is faster;at the metal-oxide interface if the diffusion of the oxygenions is faster; in all positions if the two diffusion velocitiesare comparable.

The kinetics of oxidation involve a series of stages whichcomprise: adsorption of oxygen onto the surface of themetal; the formation of an oxide nucleus extending over theentire surface; increase in the thickness of the film.

The nucleation of the oxide is favoured at high energysites such as surface defects (dislocations, grainboundaries, precipitates) and is influenced by surfacetreatments, temperature and the partial pressure of the

oxygen. Once the film has formed over the whole surface,its growth occurs through solid state diffusion processeswithin the scale.

The protective capacity of an oxide is evaluatedqualitatively with the Pilling-Bedworth ratio, defined as theratio of the volume of the oxide to that of the metal whichhas produced it. If this ratio is lower than 1 or higher than2.5, the oxide is not protective because in the first case, it isinsufficient to cover the metal, whereas in the second case,the oxide detaches due to the compressive stresses whichoriginate during growth.

If the oxide is not protective, the metal is continuallyexposed to the oxidizing atmosphere and the oxidation rateis constant. The growth of the oxide is linear and the metal’sloss of thickness is therefore given by an equation of thefollowing type:

[6] x �C1 t

where t is time and C1 is a constant of the metal.When the oxide layer is protective, its growth depends

on the diffusion processes of the O2�oxygen ions and themetal ions. The growth rate is therefore proportional to theslowest flow of the ions (Jion); this flow is given by Fick’slaw:

Assuming a constant concentration gradient, the loss ofthickness due to the formation of the oxide is given by anexpression of the following type:

[7]

where t is time and C2 and C3 are constants of the metal. Thegrowth of the oxide is parabolic.

The dependence on temperature of the diffusioncoefficient D of the diffusing species is given by:

[8]

where Do is a constant, Q is the activation energy, T thetemperature and R the gas constant. As such, the growth rateof an oxide controlled by diffusion presents an Arrheniuspattern according to temperature.

The oxides of some common metals, such as Al, Be, Znand Cr present growth kinetics of logarithmic type:

[9]

where t is time and C4 and C5 are constants of the metal. Thereasons for this deviation from parabolic behaviour arecomplex and, in the case of Al and Be, lie in the lowmobility of the electrons participating in the oxidationprocess, whereas in the case of Zn and Cr, they lie in thelower diffusion rate of the ions. In the case of Al and Cr, theoxide has a crystalline structure coherent with that of theunderlying metal and is therefore extremely adherent andprotective.

The growth of the oxide must be studied in the lightof the properties of semiconductors since the oxide whichforms during oxidation is non-stoichiometric; in otherwords, the metal to oxygen ratio is not precisely that ofthe formula even though the compound is electricallyneutral.

x C C t= +( )4 5 1ln

D D eQRT=

−

o

x C t C= +2 3

dxdt

J J Ddcdxion ion∝ = −

GENERAL ASPECTS OF CORROSION


Non-stoichiometric ionic compounds are classified assemiconductors and may be of n-type or p-type. The n-type(lack of oxygen) results from an oxygen vacancy MO1�x(Nb, Ta, Zr): ionic conduction takes place through thediffusion of oxygen vacancies, whilst the excess of electronsfalls into the band of electronic conduction. The p-type (lackof metal) results from a metal vacancy M1�xO (Fe, Ni, Co,Cu, Mn, Cr): ionic conduction takes place through thediffusion of metal vacancies whilst electron conductionoccurs through electron holes. The presence of differentatoms (or dopants) may give the same result (for example,the addition of Cr2O3 to NiO) with an increase in metalvacancies and a decrease in the concentration of electronholes.

Linear or surface defects (such as dislocations, grainboundaries, oxide-metal or oxide-gas interfaces, cracks inthe scale) form an easy route for the diffusion of defects.Some models of diffusion in polycrystalline materials arebased on equations in which the effective diffusioncoefficient depends on the diffusion coefficient in thecrystal lattice, on the diffusion coefficient in the grainboundary, on the thickness of the grain boundary and thesize of the grain, and explain the fact that diffusionincreases proportionally to the increase in the thicknessof the grain boundary and the decrease in grain size.Diffusion through defects is important at hightemperatures for Cr2O3 and Al2O3 whilst for NiO it issignificant at low temperatures. The type of oxide, p or n,determines where the film grows: at the oxide-gasinterface for p-type oxides such as NiO; at the metal-oxide interface for n-type oxides.

Multilayer oxide scalesIf a metal forms different oxides, the scale may

consist of a sequence of oxides with differentcompositions, for example M�M2O�MO. If the inner scalegrows due to the diffusion of metal ions towards theoutside, M2O forms at the M2O�MO interface through thedisplacement reaction M��e��MO��

��M2O. Variousfactors influence the oxidation rate: the purity of themetal, the composition of the gas, impurities in the gas,the pressure and temperature, the flow rate of the gas, theorientation of the crystal lattice, the surface finish of themetal, variations in temperature, the geometry andthickness of the metal.

Nickel. Forms a stable p-type oxide NiO (semiconductorwith a lack of metal). Ni1�xO has a value of x�10�4 at900°C and the partial oxygen pressure of 1 bar. The growthof the oxide occurs by migration of the metal ions with theformation of columnar oxide grains. The presence ofimpurities in the Ni leads to the formation of a fine-grainedporous oxide at the interface with the metal and a columnaroxide on the outside.

Iron. Forms three stable oxides: hematite Fe2O3,magnetite Fe3O4 and wüstite FeO. The latter is stable only attemperatures above 570°C, but lacks significant protectiveproperties. It is therefore necessary to distinguish betweenthe following situations:• if oxidation occurs at temperatures above 570°C, the

sequence of oxides is as follows:Fe�FeO�Fe3O4�Fe2O3�O2. The thickness ratio of theoxides is roughly 95:4:1 with FeO presenting significantthickness. Growth occurs by diffusion of the Fe2� ions

and reduction of Fe3O4. The mobility of Fe2� in FeO isextremely high, leading to a high oxidation rate;

• if oxidation occurs at temperatures below 570°C, noFeO forms but only Fe�Fe3O4�Fe2O3�O2 oxides. Sincethe diffusion of iron ions in magnetite is slow, steelspresent good resistance to oxidation at temperatures ofup to 550°C.Chromium. Forms the oxide Cr2O3 (spinel-corundum

structure) of p-type (though at low oxygen pressures, it seemsthat the oxide becomes n-type). Cr2�xO3 has a value ofx�9�10�5 at 1,100°C at the partial oxygen pressure of 1 bar.Since the oxide is relatively stoichiometric (low concentrationof defects), transportation within the scale is influenced bydiffusion at the grain boundary. At temperatures above 900°Cin oxygen-rich atmospheres, Cr2O3 oxidizes to volatile CrO3and loses its protective capabilities.

Aluminium. Forms the oxide Al2O3 which is very stableand protective because it is very stoichiometric. Some alloysare designed to form a film of Al2O3 which offers protectionup to 1,300°C.

Silicon. Like aluminium it forms an oxide, SiO2, whichis very stable and protective because it is verystoichiometric. New alloys are designed to form a film ofSiO2 which offers protection up to 1,200°C.

Titanium. The oxidation of Ti is complex due to theformation of numerous stable oxides (Ti2O, TiO, Ti2O3,Ti3O5, TiO2). At temperatures below 1,000°C and a partialoxygen pressure of 1 bar, only TiO2 is formed. Attemperatures above 600°C, the growth kinetics are parabolicand may become pseudolinear when exposure times arelong. At high temperatures, oxygen dissolves into the metalin significant quantities leading to the formation of cracksand the exfoliation of the metal.

Molybdenum. The oxidation of Mo leads to theformation of volatile oxides (for example, MoO3 melts at795°C). These oxides are not protective and oxidationfollows a catastrophic pattern.

Oxidation of alloysThe oxidation of alloys follows more complex

mechanisms since these consist of numerous metals whichpresent different affinities with oxygen and differentdiffusion rates. In addition, mixed oxides may form, withsolubility zones between the oxides and the elements in thealloy. For the sake of simplicity and assuming an alloy of twometals (A and B), three different situations may arise: theoxides AO and BO are completely miscible or completelyimmiscible or partially miscible. From the point of view ofcomposition, it may also be the case that if one of the twometals prevails, the other exerts a doping effect on the oxideof the prevalent metal, whilst if the composition isintermediate, the situation is more complex.

In general terms, it can be stated that resistance tooxidation is determined by the presence of reactive elementssuch as Ni, Cr, Al and Si in the alloy which form stable oxidefilms that adhere to the surfaces and provide effectiveprotection against the progress of oxidation. The strategiesfor protecting materials operating at high temperatures arenumerous and include: varying the conditions of the gas(composition, temperature, velocity); varying thecomposition of the material so as to form a protective scale;recourse to a protective coating; reducing residual stressesand loads.

MATERIALS


9.1.3 Wet corrosion

A generic corrosion reaction for a metallic material M canbe schematized as follows:

[10] M�aggressive environment��corrosion products of M

where M is a generic metallic material. When theenvironment is an electrolytic solution, the overall corrosionreaction [10] involves a metal oxidation processaccompanied by the reduction of the oxygen dissolved in thesolution, for example in the case of iron:

[11] iron�oxygen�water��corrosion products

or a second process, typical of acid solutions, in which thereduction of the hydrogen ion occurs according to thereaction, again in the case of iron:

[12] iron�acid solution��iron ions�hydrogen

These two reactions proceed according to anelectrochemical mechanism involving the electrons of themetallic material. The reaction is the sum of twocomplementary electrode processes: an anodic processinvolving the oxidation of the metallic material, makingelectrons available in the metallic phase; a cathodic processwhich consumes the electrons made available by the anodicprocess, through a reduction reaction (of molecular oxygenor the hydrogen ion or both).

Since electroneutrality must be maintained, the tworeactions must occur simultaneously and with the samevelocity.

Electrode processes

Anodic processesThe generic anodic process of a metal can be represented

by the oxidation reaction of a metal to its ion which passesinto solution:

[13] M��Mz��ze�

where z is the valence of the metal, e� indicates the electron,M the generic metallic material and Mz� its ion which passesinto solution. In cases where the metallic material tends toform hydroxides, the anodic reaction is of the followingtype:

[14] M �zH2O��M(OH)z�zH��ze�

Cathodic processesBy contrast, the cathodic processes which are of

practical interest for corrosion are limited in number. In thecase of corrosion in an acid solution, the cathodic process isthe reduction of the hydrogen ion and the production ofmolecular hydrogen, according to the reaction:

[15] 2H��2e��H2

where e� indicates the electron.In natural environments, by far the most important is the

oxygen reduction reaction, in a neutral or basic environment:

[16] O2�2H2O �4e��4OH–

or, in an acid environment:

[17] O2�4H��4e��2H2O

The oxygen which appears as a reagent is molecularoxygen dissolved in the water, whose concentration rangesfrom 0 to 12 mg per kg of water (ppm).

Stoichiometric aspectsSince the mechanism is electrochemical, it is possible to

apply Faraday’s laws which establish the relationshipbetween the masses and the circulating electrical current(number of electrons) using the electrochemical equivalents:

[18]

where q is the charge circulated (coulomb), eech is theelectrochemical equivalent, echem is the chemical equivalentand F is Faraday’s constant (96,500 coulomb/eq). Therelationship between the corrosion rate as mass loss and thecurrent exchanged by the metallic material is:

[19]

where ia is the anodic current or corrosion density.The equivalence between the corrosion rate

expressed in mA�m2 and mm�y for divalent metals with adensity of about 8 t�m3 (for example, Fe, Zn and Cu)is 1 mA�m2�1 mm/yr (for iron, the exact value is1 mA�m2�1.17 mm�yr).

Thermodynamic aspectsThe generic corrosion reaction [10] for a metallic

material M occurs if it is thermodynamically favoured, inother words if the variation of the free energy DGassociated with it is negative. If we consider a corrosionreaction and the two complementary anodic and cathodicreactions:

[20] M �(z�a)A��Mz��(z/a) Aa�

[21] M��Mz��ze� (anodic)

[22] (z�a)A�ze��(z/a) Aa� (cathodic)

the general thermodynamic condition outlined above can beapplied to all three reactions. Since these are electrochemicalreactions, the variation of free energy DG can be expressedas a variation of the electrical force associated with thereaction:

[23]

where DE is equivalent to the electromotive force of thereaction considered, and z and F have the usual meaning.Below DE is also referred to as the driving force or potentialdifference.

The thermodynamic condition of spontaneity of thecorrosion process thus becomes:

[24]

Since the variation of free energy is expressed by theequation:

[25]

introducing the potential leads to Nernst’s law:

[26] E E RTzF

M Mz

z+

= ° + [ ][ ]

+/ ln

MM

∆ ∆ ΠΠ

G G RTprod

react= ° + ln

∆ ∆G E< >0 0ovvero

∆ ∆G zF E= −

VV m

Ate qAt

e icorrcorr m ech

ech a= = = =,

γ γ γ γ∆ 1

∆m e qeF

qechchem= =



or

[27]

or

This makes it easy to obtain the variation of free energyfor the overall reaction as the sum of the variations of the twopartial reactions (Hess’ law). Using the potentials we obtain:

[28]

Condition [24] thus becomes:

[29] E E E Ec a c a− > >0 or

∆E E Ec a= −

E E RTzFc c

a= − ° −ln A

E E RTzF

A Aa

a/ ln−

= ° −

−A

A

E E RTzFa a

z= + ° +ln M

MATERIALS


Table 2. Electrochemical series of standard potentials

Electrode reactionsE

(V vs. SHE)Electrode reactions

E(V vs. SHE)

F2�2H��2e��2HF �3.03 2H��2e��

��H2 0

O3�2H��2e��O2�H2O �2.07 2D��2e��

��D2 �0.0034

Co3��3e��Co �1.842 Fe3��3e��

��Fe �0.036

Au��e��Au �1.68 Pb2��2e��

��Pb �0.1263

Au3��3e��Au �1.50 Sn2��2e��

��Sn �0.1364

MnO4��8H��5e��

��Mn2��4H2O �1.491 Ge4��4e��Ge �0.15

PbO2�4H��2e��Pb2��2H2O �1.467 Mo3��3e��

��Mo �0.20

Cl2�2e��2Cl� �1.3583 Ni2��2e��

��Ni �0,25

Cr2O72��14H��6e��

��2Cr3��7H2O �1.33 Co2��2e��Co �0.28

O2�4H��4e��2H2O �1.23 Mn3��3e��

��Mn �0.283

CrO42��8H��3e��

��Cr3��4H2O �1.195 In3��3e��In �0.342

Pt2��2e��Pt �1.19 Cd2��2e��

��Cd �0.40

Br2�2e��2Br� �1.087 Cr3��e��

��Cr2+ �0.41

HNO3�3H��3e��NO�2H2O �0.96 Fe2��2e��

��Fe �0.44

2Hg2��2e��Hg2

2� �0.92 Cr3��3e��Cr �0.74

Hg2��2e��Hg �0.851 Zn2��2e��

��Zn �0.76

Ag��e��Ag �0.7996 V3��3e��

��V �0.876

Hg22��2e��

��2Hg �0.7961 Cr2��2e��Cr �0.913

Fe3��e��Fe2+ �0.770 Nb3��3e��

��Nb �1.10

O2�2H��2e��H2O2 �0.682 Mn2��2e��

��Mn �1.18

Hg2SO4�2e��2Hg�SO4

2� �0.62 V2��2e��V �1.18

MnO4��2H2O�3e��

��MnO2�4OH� �0.588 Ti3��3e��Ti �1.21

I2�2e��2I� �0.534 Zr4��4e��

��Zr �1.53

Cu��e��Cu �0.522 Ti2��2e��

��Ti �1.63

Cu2��2e��Cu �0.34 Al3��3e��

��Al �1.66

AgCl�e��Ag�Cl� �0.22 Mg2��2e��

��Mg �2.36

Cu2��e��Cu� �0.158 Na��e��

��Na �2.71

Sn4��2e��Sn2� �0.15 Ca2��2e��

��Ca �2.86

2H��2e��H2 0 Li��e��

��Li �3.05

in other words, the corrosion reaction is spontaneous if thepotential of the cathodic reaction is more noble than thepotential of the anodic reaction. Only in this case is apositive driving force available (DE�0) which makes thereaction possible.

Considering the anodic reaction alone [21], thethermodynamic condition for it to proceed in the anodicdirection (corrosion) emerges when its potential E is morenoble than the equilibrium potential, indicated by Eeq, givenby

[30]

where E° is the standard potential of the metal and [Mz�]the concentration of its ions in the electrolyte in contactwith its surface. Table 2 shows the standard potentials –expressed in volt with respect to the hydrogen electrode(SHE, Standard Hydrogen Electrode) – of the mostcommon metals and electrode reactions. Thethermodynamic condition for the reaction to proceed in theanodic direction (corrosion) is thus E�Eeq, correspondingto a negative variation of free energy (DG�0). If, however,E�Eeq, we have a positive variation of free energy(DG�0), and the reaction proceeds in the cathodicdirection (the oxidation of the metal cannot occur, in otherwords a condition of thermodynamic immunity is created).If a reference concentration is established above which it isconsidered that the metal undergoes corrosion, for examplea concentration of 10�6 mol/L, as suggested by M.Pourbaix (1973), using [30] we can calculate the immunitypotential of the metal used in practical applications, forexample to determine the conditions for cathodicprotection (see below).

Measuring potential and reference electrodesConventionally, potentials are referred to the hydrogen

electrode, taken as the zero reference point. In practice,potential is measured by connecting a voltmeter to the metal(or metal structure) and to a reference electrode which hasthe property of keeping its own potential stable. One of themost commonly used electrodes in the laboratory is theSaturated Calomel Electrode (SCE), which has a potential of+0.24 V compared to SHE. Potential measurements on realstructures make use of a copper-saturated copper sulphateelectrode CSE (Copper Sulphate Electrode) in soils (whichhas a potential of �0.3 V compared to SHE), zinc (which

has a potential of �0.8 V compared to SHE) andsilver-silver chloride (Ag/AgCl) in sea water (which has apotential of �0.25 V compared to SHE). Table 3 shows thereference electrodes used in the laboratory and in the field tomeasure potential.

Pourbaix diagramsIn 1946 Marcel Pourbaix introduced the potential-pH

diagrams which supply the equilibrium potentials as the pHvaries for metals in contact with electrolytes (Pourbaix,1973). The diagrams show the cathodic reactions ofhydrogen release [15] and oxygen reduction [16],represented by two parallel lines with an angular coefficientof �0.059 and at a distance of 1.23 V, obtained from theNernst equations:

[31]

[32]

The dissolution reaction of a metal M is represented bythe equation:

EeqO H O2 2/

. .= −1 229 0 0059pH

EeqH H+

= −/ .0 0059pH

E E RTzFeq

z= + ° +ln M



E (

V v

s. S

HE

)

2

�2

�1

0

1

pH

Mimmunity

M(OH)zpassivity

Mz�

activity

0 7 14

0�3�6

0 �3 �6

b

a

Fig. 2. E-pH diagram for a generic metal M which formshydroxides. Line a represents the hydrogen release reaction and line b the oxygen reduction reaction.

Table 3. Reference electrodes

Reference electrode Description SemireactionE

(V vs. SHE)

Standard hydrogen electrode H2(1 atm)�H�(a�1) 2H��2e��H2

0

Calomel Hg�Hg2Cl2, KCl (sat) Hg2Cl2�2e��2Hg �2Cl� �0.244

Silver/silver cloride (0.1M) Ag�AgCl, KCl (0.1M) AgCl �e��Ag �Cl� �0.288

Silver/silver cloride/sea water Ag�AgCl, sea water AgCl �e��Ag �Cl� �0.250

Copper/saturated copper sulphate Cu�CuSO4 (sat) Cu2��2e��Cu �0.318

Zinc/sea water Zn�sea water corrosion reaction ��0.80

[33]

where EoMz��M is the standard potential of the dissolution

reaction of metal M. In the E-pH diagram, the equilibriumpotential Eeq

Mz��M does not depend on pH and is thereforerepresented by a series of straight lines parallel to the x-axis(Fig. 2) where each line corresponds to a value of theparameter log aMz�. The line characterized by a value of theparameter corresponding to the concentration of 10�6 mol/Ldivides the plane into two regions: the upper corrosionregion for concentrations above this value and the lowerthermodynamic stability region of metal M for lowerconcentrations, also known as the immunity zone.

In the simplest case, if the dissolution reaction of metalM leads to the formation of hydroxides, especially in aneutral or basic environment, according to the reaction

[34] M �zH2O��M(OH)z�zH��ze�

the equilibrium condition is given by the equation

[35]

By substituting the unitary activities of the metal and thehydroxide and introducing the value of the constants (R is

the universal gas constant; F is Faraday’s constant) weobtain:

[36]

showing that the equilibrium potential varies with pHaccording to a line with the same slope as lines a and b ofthe cathodic processes. In Fig. 2, the generic line of equation[36] separates the immunity zone from that of hydroxideformation, also known as the passivity zone. Fig. 3 shows thesimplified Pourbaix diagram for iron.

Pourbaix diagrams, constructed on the basis ofthermodynamic equilibrium data for the electrochemicalreactions involving the metal as pH varies, show the stabilityzones of the chemical species involved (immunity zone ofthe metal; the stability zone for oxides, which may give riseto passivation phenomena; the stability zone of the metalions and complexed forms in strongly alkalineenvironments) but they cannot provide information (exceptin a qualitative sense) on the kinetics of the processes, inother words on the corrosion rate. It should also be notedthat in conditions of non-equilibrium, as is the case inpractice, the stability zones are different, with a significantimpact on the kinetics. In the case of iron, for example, thediagram obtained experimentally in agitated water presents afar larger passivity zone than that predicted theoretically.

Kinetic aspectsThe availability of a driving force represents a necessary

but insufficient condition for the corrosion reaction to takeplace: the intervention of reaction resistances (or generalizedattrition) conditions the corrosion rate and may even stopcorrosion. In any case, knowing the driving force alone doesnot make it possible to predict the evolution of the corrosionprocess over time. In other words, as in many chemical andphysical phenomena, kinetic factors intervene which mayradically alter the behaviour of a material subject tocorrosion as deduced from thermodynamic considerations.This is the case, for example, of titanium, which accordingto thermodynamics is more reactive than iron since itsstandard potential is more negative by over 1 V. In practice,however, titanium behaves like a noble metal and does notundergo corrosion in environments where iron is corroded,such as sea water. This behaviour is due to the interventionof passivation phenomena (see below) with the formation ofprotective films. In the corrosion process, precisely becauseit is of electrochemical nature, generalized attrition islocalized on the surface of the material, where the electronic

E EeqM M OH M M OHz z/ ( ) / ( )

.= −o

pH0 059

E E RTzF

a a

aeqM M OH M M OH M OH H

z

M

z z z/ ( ) / ( ) ( )ln= +

+

o

E Ez

aeqM M M M

M

z z

z

+ +

+= −/ / . logo0 0059

MATERIALS


E (

V v

s. S

HE

)2.0

�1.2

�0.4

0.4

1.2

1.6

�0.8

0

0.8

pH0 7 14

0�3

�6�10

b

a

Fe3�

Fe2� Fe2O3

HFeO2�

Fe3O4

Fe

Fig. 3. Simplified Pourbaix diagram for iron.

Ia � Ienv � Ic � Im � Icurr

anodicprocess

current flowin metal

cathodicprocess

current flowin environment

Ia

Im

Ienv

Ic

Fig. 4. Schematic illustrationof the electrochemicalmechanism (Pedeferri, 2007).

reactions take place, and in the electrolyte, as shown in thediagram in Fig. 4, where the four partial processes arecomplementary; in other words, they take place at the samevelocity. These processes are:• the iron oxidation reaction (anodic process) which makes

electrons available in the metallic phase (Ia);• the oxygen reduction reaction (cathodic process) which

consumes these electrons and produces alkalinity:O2�2H2O�4e��

��4OH� (Ic);• the flow of the electrons inside the metal from the anodic

to the cathodic regions (Im);• the current flow into the electrolyte (Ienv).

[37] Ia�Ic�Im�Ienv�Icurr

The total value of these currents (Icurr) measures, inelectrochemical units, the velocity of the overall corrosionprocess. The corrosion rate is determined by the slowest ofthe four partial processes; in practice, since the electricalresistance of the metal is always negligible, the current flowin the metal is never a slow process and therefore does nothelp to reduce the corrosion rate. Each of the other threeprocesses, on the other hand, under specific environmentalconditions, may take place at an extremely slow rate andtherefore represent the kinetically controlling process. Atypical instance is the presence of passivity, which reducesthe velocity of the anodic process to negligible values; or thereduction of the velocity of the cathodic process due to theintervention of overpotentials (for example, the use ofcorrosion inhibitors in acid solutions). In the first instance,we speak of anodic control, in the latter of cathodic control.

Potential-current diagramsTo describe corrosion systems, it is opportune to use the

diagrams proposed in 1945 by Ulick Richardson Evans, wherethe two axes show the potential (E) and the current density(log j), which show the characteristic curves of the anodic andcathodic reactions obtained experimentally starting from thereaction’s equilibrium potential. As the current increases, in ananodic or cathodic direction, the reaction potential shifts fromthe equilibrium potential by a value known as theoverpotential, indicated by h, which represents the dissipativeterm expended as the reaction proceeds and which increaseswith the reaction velocity following a logarithmic pattern. Theoverpotentials of anodic reactions have a positive sign andthose of cathodic reactions a negative sign. If the characteristiccurves of both the anodic and cathodic processes are shownon the same diagram, the corrosion potential and thecorrosion current given by the point where the two curvesmeet can be identified (Fig. 5). The corrosion potential (Ecorr)can be measured directly. Usually, Evans diagrams are ofsemilogarithmic type, with the current density represented ona logarithmic scale to render the characteristics linear.

The driving force DE made available by the anodicprocess is in this case given by the difference between thecorrosion potential Ecorr and the equilibrium potential Eeq, sothat DE�Ecorr�Eeq. The potential Ecorr, to which the systembrings itself, is higher than the equilibrium potential of themetal but lower than the equilibrium potential of thecathodic process.

Electrode overpotentialsIn corrosion processes, we can distinguish

between two types of overpotential: the activation

or charge transfer overpotential and the concentrationoverpotential.

The activation overpotential is associated with allelectrode reactions in which a charge transfer takes place:

[38] Ox �ze��Red

and becomes the activation energy of the reaction. It can bededuced that the velocity constant follows Arrhenius’ equation:

[39]

where Ea is the activation energy referred to one mole, Tthe thermodynamic temperature, R the gas constant and Z aconstant.

The general expression of current density as a functionof overpotential is given by the Butler-Volmer equation:

[40]

where h is the overpotential given by h��E�Eeq�, io is theexchange current density, b is the symmetry factor betweenthe anodic and cathodic branches (usually equal to 0.5), F isFaraday’s constant and z the equivalence of the reaction(number of electrons in the electrode reaction underconsideration). The Butler-Volmer equation establishes thatcurrent exchange at the surface of an electrode occurs only ifan activation energy is exceeded, in other words bydissipating part of the available driving force.

If the activation overpotential of an electrode reaction isrelatively large, due also to low values of density currentexchanged i, in the order of �h��50 mV, the Butler-Volmerequation [40] is reduced, assuming b�0.5, to the well-known Tafel equation:

[41] h �a �b logi

or, in its equivalent form

[42] η = ±b ii

logo

i i e ezF

RTzFRT= −

−( )−

o

1 β βη η

k T ZeERTa

( ) =−



cathodic process

anodic process

E

Eeq,c

Eeq,a

Ecorr

logjlogjcorr

Fig. 5. Diagram of potential-logarithmic of the current density of a corrosion process (anodic and cathodic property).

where the � sign is used for anodic processes in whichthe h are positive and the � sign is used for cathodicprocesses in which the h are negative; a is a (positive)constant which depends on the exchange current densityio, b is a (positive) constant equivalent to the slope of lineh-i in a semilogarithmic diagram, known as the Tafelslope. Introducing the decimal logarithms, constant btakes a value of 59 mV/decade for bivalent reactions and118 mV/decade for monovalent reactions (such ashydrogen release).

Overpotentials of metal dissolution. Experimental resultsshow (Piontelli, 1961) that the characteristics of the anodicmetal dissolution reactions and the opposing cathodicdeposition reactions are almost symmetrical. As far as thevalues of the overpotentials are concerned, three classes ofbehaviour have been identified, which depend on the atomicand crystalline properties of metals:• Normal metals, which have very low overpotentials,

below 10 mV, even at high current densities (both anodicand cathodic). This class comprises metals which have alow melting point (below 600°C), in other wordsrelatively weak interatomic bonds: Cd, Hg, Sn, Pb, Mg,Al; ZN on the anodic side only.

• Inert metals, which present overpotentials above 100 mVeven at low current densities. In contrast to the precedingclass, this class comprises metals with a high meltingpoint (above 1,400°C), in other words strongerinteratomic bonds: Fe, Co, Ni, Cr, Mo, Ti, metals of thePt group and transition metals.

• Intermediate metals, which present intermediatebehaviour compared to the two preceding classes. Thisgroup comprises metals with an intermediate meltingpoint, around 1,000°C, such as Cu, Ag, Au.The correlation between the overpotential and melting

point of metals also extends to other physical properties,again correlated with the nature of the interatomic bonds; forexample, it is easy to predict that normal metals (lowoverpotentials) also have low hardness and mechanicalresistance and high interatomic distances, whilst the exactopposite is true of inert metals. The variation ofoverpotentials with current density is practically linear fornormal metals whilst it follows Tafel’s law for intermediateand inert metals.

Passivity. Iron, and carbon and low alloy steels are foundin natural environments, such as soils and waters, and ingeneral in acid solutions, under so-called active conditions;in other words, with an anodic reaction resistance which ispractically nil. In these cases, the anodic process cannotcontribute to reducing the corrosion rate. Numerous metalsand their alloys with a high affinity for oxygen have thecharacteristic of covering themselves with a protective layerof oxide which prevents their corrosion in corrosiveenvironments. These conditions are described as passivityconditions and determine the state of ‘inertia’ of the metal,which behaves like a noble metal. A well-known example isstainless steel, which owes its passivity to the formation onits surface of a very protective chromium oxide layer whichmakes it, as the name suggests, resistant to manyenvironments where iron or carbon and low alloy steels maysuffer even very severe corrosion. Another example ofpassive behaviour, which often goes unnoticed, is that of thereinforcement bars of reinforced concrete which areperfectly passivated by the alkalinity of the hydrated cement

mixture. The iron in concrete (or very alkaline solutions)behaves like stainless steel in fresh water. Aluminium andtitanium are resistant to corrosion thanks to their ability topassivate.

From an electrochemical point of view, active andpassive behaviour are characterized by two different patternsin the anodic property: in active materials, it is a straight linewith a slight slope; in materials with active-passivebehaviour, it is a vertical line which follows the typicalpattern shown in Fig. 6.

Overpotential of hydrogen releaseThe attritions (overpotentials) of the hydrogen release

process, represented by reaction [15], depend on the natureof the metallic material M on which the reaction takes place,according to an anticorrelation with the overpotential ofmetal dissolution. A distinction should therefore be madebetween:• Normal metals characterized by very low dissolution

overpotentials, below 10 mV, and high hydrogen releaseoverpotentials, corresponding to low values of theexchange current density, between 10�3 and 10�6

mA�m2 (Hg, Sn, Pb, Mg, Al, Zn).• Inert metals which present overpotentials above 100 mV

and very low hydrogen release overpotentials,corresponding to high values of the exchange currentdensity, between 10 and 105 mA�m2 (Fe, Co, Ni, Cr, Mo,Ti, Pt).As seen above, the relation linking the overpotential

hH to the current density is given by Tafel’s law [41]. Theslope b of Tafel’s line in a semilogarithmic diagram is intheory equal to 118 mV/decade and in practice fallswithin the range 120 to 150 mV/decade for all metals,whilst the exchange current density jo,H varies by asmuch as 11 orders of magnitude passing from mercuryto platinum.

MATERIALS


transpassivity

passivity

activity

E

Et

Eeq,a

Ep

logjlogjp

Fig. 6. Diagram of potential-logarithmic of the current density of a metal with active-passive behaviour.Ep, passivation potential; Et, transpassivity potential; jp, passivity current density.

Overpotential of oxygen reductionThe oxygen reduction reaction is the main cathodic

process in reactions which occur in natural environments andin neutral or weakly alkaline solutions. As already stated, theoxygen participating in wet corrosion processes is thatdissolved in water, in equilibrium with the oxygen present asa gas in the atmosphere. The solubility of oxygen in waterdecreases as temperature increases (becoming practically nilabove 60°C) and as the content of dissolved salts in thewater increases. In pure water at 0°C, the solubility ofoxygen is 10 mL/L, falling to 5.28 mL/L at 30°C. In seawater, with a salinity of 36 g/L, the solubility of oxygen is 8 mL/L at 0°C and 4.33 mL/L at 30°C. In closed systemsoperating at temperatures above 60°C, such as heatingcircuits, all the oxygen present participates in the cathodicreduction reaction at high velocity given the hightemperature, despite its extremely low solubility, lower than1 mL/L (1 mL/L of oxygen in water corresponds to about 1.2 ppm or mg/kg).

In neutral or alkaline solutions, in the absence of oxygen,corrosion is negligible, as in the case of the closed circuitsused for heating or cooling in plants: since these are closedsystems, oxygen is initially present when the plant is startedup but is rapidly consumed by corrosion; subsequently, thewater is deaerated and non-corrosive. As such, in neutral oralkaline solutions, the corrosion rate is equal to the oxygenreduction rate on the metal, determined by its supply bydiffusion since oxygen is a neutral species not influenced bycurrent flow. Under stationary conditions, the oxygen supplyis given by Fick’s first law, according to which it is directlyproportional to the concentration gradient and the diffusioncoefficient, and inversely proportional to the limit layer ofdiffusion. In electrochemical units, we obtain:

[43]

where D is the diffusion coefficient, x is the coordinate inthe direction of oxygen flow, d the thickness of the limitlayer of diffusion, CO2

and CM,O2the concentration in the

solution and on the surface of the metal. Equation [43]admits a maximum when, chemico-physical conditionsbeing equal, the concentration of oxygen on the surfaceof the metal falls to zero: CM,O2

�0, in other words, whenall the oxygen reaching the surface of the metal bydiffusion has been consumed by the corrosion process.The maximum value of the oxygen flow reaching thesurface of the metal from the solution, corresponding tothe limit diffusion current of oxygen jl, is therefore givenby the equation:

[44]

The value of jl depends on three factors: the diffusioncoefficient D which increases as temperature increases, theconcentration of oxygen CO2

which decreases as temperatureincreases, the limit diffusion layer d which is highest understagnant conditions and decreases as turbulence increases. Insea water, for example, since the oxygen concentration doesnot exceed 11 ppm, the limit diffusion layer d ranges inpractice from 0.1 to 3 mm and the diffusion coefficient Dranges from 1.3 to 2.5 10�9 m2�s�1, for temperaturesbetween 10 and 30°C, the limit current density of oxygendiffusion ranges from about 50 mA/m2 to about 2 A/m2.

The characteristic curve of the cathodic oxygenreduction reaction follows the pattern shown in Fig. 7 and ischaracterized by a vertical line which represents the limitcurrent density of diffusion jl. The importance of jl lies in thefact that for common metals (such as Fe, Zn), the corrosionrate coincides with the limit current.

Under stagnant conditions, the application of Fick’slaw is easy since the thickness of the limit layer d can beestimated from the experimental graphs reported inmanuals; an empirical rule has been developed accordingto which, at ambient temperature, the corrosion rate inmm/yr is 10 times the concentration of oxygen in ppm.Under non-stagnant conditions, the calculation of d isfairly complex, so the limit current of oxygen diffusion jl,which depends on fluid dynamic conditions, is obtainedfrom Sherwood’s number (or Nusselt’s number) expressedby the equation:

[45] Sh �/�d � jl/�(4FD CO2)

where / is a characteristic dimension, for example thediameter of a pipeline. The adimensional Sherwoodnumber is calculated as a function of the adimensionalReynolds (Re) and Schmidt numbers (Sc; Shreir et al.,1994; Lazzari and Pedeferri, 2006). For a rough but often acceptable estimate of the limit currentunder turbulent conditions, the value obtained under stagnant conditions is multiplied by themultiplication coefficient 1�(u)1�2, where u is themean velocity in m/s.

9.1.4 Types of corrosion

Generalized corrosionGeneralized corrosion affects the entire surface of a

metal or a large part of it. A distinction is made betweenuniform and non-uniform generalized corrosion: in the firstinstance the loss of thickness is uniform over the whole

j FDC

lO2= 4

δ

j

FD dC

dxD

C CO O M O2 2 2

4= − =

−,

δ



diffusioncontrol

activationcontrol

E

logjlogjl

Fig. 7. Pattern of the characteristic cathodic curve of oxygen reduction.

surface, whereas in the latter case it takes a more or lessregular form.

As far as the mechanism is concerned, generalizedcorrosion indicates the substantial coincidence of the anodicand cathodic areas, affecting metal-environment systems in anactive state, as is the case of carbon steel in acid solutions andnatural environments such as soils, waters and the atmosphere.When corrosion is controlled by the oxygen reductionreaction, the maximum velocity of uniform penetrationcoincides with the limit current of oxygen diffusion.

In acids, the cathodic reaction is the reduction ofhydrogen ions to gaseous molecular hydrogen. Thecontrolling parameter is pH: for iron and steels in general,acid corrosion becomes significant when the pH is lowerthan 4 and increases exponentially at lower pH values; whenthe pH is alkaline, by contrast, where the presence of OH�

ions in solution prevails, the corrosion rate of iron becomesnegligible due to the formation of a protective oxide film.

The prevention of this type of corrosion, less insidiousthan other types because its initiation and the mean velocityof thickness loss can be predicted, generally involves addingan extra thickness of metal, sized by multiplying the uniformcorrosion rate by the project life and, more generally, usingtraditional methods such as cathodic protection, coatings(organic, inorganic and metallic), paints and corrosioninhibitors.

Localized corrosionWhen the anodic and cathodic reactions take place on

different surfaces, localized corrosion occurs, affecting onlya limited part of the surface exposed to the environment. Theseparation of the areas leads to the circulation of a current,known as a galvanic current, with the circulation of electronsin the metal from the anodic to the cathodic area and thecirculation of ions in the solution: positive ions migrate fromthe anodic to the cathodic area and negative ones in theopposite direction. The causes leading to the establishmentof a galvanic current are numerous: the differing nobility ofthe metals or alloys in electrical contact, which causescorrosion by galvanic contact; the different availability ofoxygen which leads to the separation of the anodic andcathodic areas, for example on steels; the local failure of thepassivity film in active-passive materials.

Corrosion by galvanic contactThis is also known as galvanic or bimetallic corrosion

and occurs on a metal when it is in electrical contact withanother, more noble, metal (or a non-metallic material withelectron conductivity, such as graphite, or surface films,such as magnetite or iron sulphide) and both are exposed toan aggressive environment. Under these conditions, thecorrosion rate of the less noble metal undergoes anacceleration which depends on the ratio of the area of themore noble metal (cathodic area) to that of the less noblemetal which is corroded (anodic area). A typical example isattack on low alloy carbon steels in sea water or aeratedsolutions when they are coupled with more noble materialssuch as copper alloys, or materials whose practical nobilityis greater, such as stainless steels and titanium. For galvaniccorrosion to occur, three conditions are necessary: the twometals must be in electrical contact (often a mechanicalmetal-metal contact is sufficient); they must be of differentnobility (the less noble metal is corroded, whilst the more

noble metal is the main site of the cathodic process withreduction or even suppression of corrosion); they must beexposed to a corrosive environment where a cathodicprocess is possible (such as sea water where the cathodicprocess is oxygen reduction). It is worth stressing that inneutral and deaerated solutions, corrosion by galvaniccontact is negligible since the cathodic process is absent ortakes place at an extremely low velocity. Severe corrosionconditions due to galvanic contact occur in sea water (lowresistivity, high driving force and availability of oxygen asthe cathodic process) and in acid solutions (low resistivity,low driving force but high velocity of the cathodic process ofhydrogen release), especially when the area ratio isunfavourable (cathodic area�anodic area).

MATERIALS


Table 4. Series of practical potentials of materialsin sea water (LaQue, 1975)

Most noble

GraphiteTitanium Stainless steels with a high Cr and Mo(passive)Stainless steel 18-8-3, type AISI 316(passive)Stainless steel 18-8, type AISI 304(passive)Stainless steel 13%Cr, type AISI 410(passive)Nickel (passive)Silver for weldingBronze MBronze GCupronickel 70-30Cupronickel 90-10BronzeCopperRed brassBronze-AlAdmiralty brassYellow brassNickel (active)Naval brassBronze-MnMuntz metal

Hydrogen

Less noble

TinLeadStainless steel 18-8-3, type AISI 316(active)Stainless steel 18-8, type AISI 304(active)Stainless steel 13%Cr, type AISI 410(active)Cast ironLamination steelMild steelAluminium series 2024CadmiumAlcladAluminium series 6053Galvanized steelZincMagnesium alloysMagnesium

The areas of the two metals involved in the galvaniccorrosion process are determined by the resistivity of theelectrolyte and the driving force of the corrosion process(given by the difference between the equilibrium potentials ofthe cathodic and anodic processes, subtracting theoverpotentials of the two processes from this value). Given anidentical driving force, in resistive environments the attack islocalized and restricted to the areas where the two metals areclose to one another; by contrast, in environments with highconductivity, such as sea water, attack is more widespreadand affects surfaces distant from one another. For an accuratecalculation, mathematical models are needed, with finiteelements to solve the electrical field equation. If the areas ofthe two anodic and cathodic surfaces are known, thecorrosion rate due to galvanic contact is given by the generalexpression jcat (1�Sc�Sa), where jcat is the velocity of thecathodic process, and Sa and Sc are the anodic and cathodicareas respectively. In waters, the velocity of the cathodicprocess is the limit current density of oxygen diffusion jl.

To establish the presence of a galvanic contact, it is firstnecessary to compare the potentials of the two metals in theaggressive environment to which they are exposed.Generally, this is done using the scale of practical potentialsin sea water shown in Table 4 which also shows differentenvironments such as soils and waters (LaQue, 1975). Table 5 shows the practical potentials of metallic materialsused in oil wells (Wellmate R2, 1999).

Corrosion by galvanic contact is prevented by: a) avoiding coupling materials with different nobility (for

example, by using insulating flanges); b) using large anodicareas and small cathodic areas; c) avoiding electrolytes withlow resistivity; d ) applying cathodic protection. Painting isan effective remedy when applied to cathodic areas, in otherwords the more noble metal; if paint (or the insulatingcoating) is only applied to the less noble metal, it becomesdangerous since it increases the corrosion rate whereverthere are defects in the coating. Secondary effects ofcorrosion by galvanic contact may occur, such as the releaseof hydrogen on the more noble metal with potentialhydrogen embrittlement on susceptible stainless alloys andtitanium.

PittingPitting manifests itself as attacks, known as pits or

cavities which are extremely penetrating but affect only avery small part of the metallic surface compared to theexposed surface area. Linear dimensions vary from a fewtens of µm up to a few mm, and the shape may vary fromtrough pits to sideways pits.

This type of corrosion is typical of materials with active-passive behaviour. Among commonly used metals, pittingcorrosion affects stainless steels, copper and aluminium. Themechanism has two distinct stages: initiation andpropagation.

The initiation stage lasts sufficient time for the localrupture of the passivity film by specific chemical speciespresent in the corrosive environment, such as chloride ions,Cl�, in the case of stainless steels and aluminium alloys,



Table 5. Series of practical potentials of materials in contact with hydrocarbons (Wellmate R2, 1999)

Potentials of free corrosion of metallic materialsused in the oil industry

(V vs. SCE)

Metallic material Sweet environment Sour environment Aerated sea water

Carbon and low alloy steels �0.6 to �0.7 �0.65 to �0.75 �0.6

Martensitic stainless steel 9Cr 1Mo �0.5�0.5 (T�120°C)�0.4 (T�120°C)

Martensitic stainless steel (type 410) �0.55 �0.6

Martensitic stainless steel (type 420)�0.45 (T�120°C)�0.4 (T�120°C)

Austenitic stainless steel (type 304) �0.25 to �0.35 �0.3 to 0.45 �0.12 to �0.04

Austenitic stainless steel (type 316) �0.2 to �0.25 �0.35 to �0.4 �0.12 to �0.04

Austenitic-ferritic stainless steel(duplex)

�0.3 to �0.4 �0.35 to �0.45 �0.10 to �0.02

Austenitic-ferritic stainless steel(duplex) (type 2205)

�0.4 (T�120°C)�0.35 (T�120°C)

Nickel alloys type C-276, type 625,type 825

�0.4 (T�120°C)�0.35 (T�120°C)

Nickel alloys type G3, type 718

�0.35 (T�120°C)�0.33 (T�120°C)

Titanium �0.12 to �0.04

when their concentration exceeds a threshold value whichdepends on composition (for example the Cr and Mo contentof stainless steels) and environmental parameters (such astemperature and turbulence: stagnant conditions reduce theinitiation time and generally make pitting more severe;temperature always encourages initiation and increases thepropagation velocity). The presence of inclusions, theformation of precipitates and the metal’s degree of workhardening encourage the initiation of pitting.

Fig. 8 shows the effect of chloride ions on thecharacteristic curve of a generic stainless steel in an aqueousenvironment, with a reduction of the passivity interval, thegreater the higher the concentration of chlorides. A potential,known as the pitting or rupture potential Er, is shown on thecurves, above which initiation occurs. The pitting potential,which can be measured experimentally, is used to determinethe relative resistance to pitting corrosion of metallicmaterials.

During the propagation stage, a galvanic current isestablished between areas where the oxide has been damagedwith dissolution of the metal and passive areas with cathodicbehaviour. The resulting corrosion rate is extremely high, inmany cases in the order of several mm per year, given theunfavourable ratio of a small anodic area to a far largercathodic area. During this stage, the solution inside the pitbecomes increasingly aggressive due to the hydrolysisreactions, of the type Mz��zH2O��

��M(OH)z�zH�, of themetals passing into solution (especially Cr in stainlesssteels). These lead both to a progressive acidification of thesolution to a pH below 2 and an increase in the concentrationof chlorides drawn into the pit by the galvanic current(autocatalytic pitting mechanism).

The prevention of pitting is mainly carried out during theinitiation stage since once pitting has set in, it is moredifficult to stop its propagation; as such, it is based on thechoice of materials resistant to initiation under the operatingconditions. For austenitic and austenitic-ferritic stainless

steels and nickel alloys, an index is used known as thePitting Resistance Equivalent Number (PREN), calculatedon the basis of the alloy’s Cr, Mo, W and N contentaccording to the following formula:

[46] PREN �Cr% �3.3(Mo% �0.5W%) �16N%

For PREN values above 40, resistance to pitting is high,whilst for values below 35 it is low. Cathodic protection iseffective during both the initiation and propagation phases.

Crevice and deposit corrosion The presence of crevices, which are often extremely

small, measuring less than a µm, or parts of the surfacewhich are not freely exposed to the environment oftenrepresent an aggravating factor for the corrosion of a metal.Examples of crevices are: coupling using flanges, spotwelding, the presence of deposits of various types.

Crevices can make corrosion more severe throughseveral different mechanisms. For metals in an active state,such as carbon and low alloy steels in water or soils, thereduced oxygen supply inside the crevice creates adifferential aeration cell, with the anodic area located insidethe crevice, where the oxygen supply is low or non-existent,and the cathodic area outside the crevice, where the oxygensupply is high. The corrosion rate is given by the product ofthe limit current density of oxygen diffusion on the cathodicarea and the ratio of the cathodic area to the anodic areaunder a deposit. In many practical situations, this ratio mayrange from 10 to 20 depending on the conductivity of theenvironment and the geometry.

The crevice corrosion mechanism of stainless steels insolutions containing chloride ions is more complex. As forpitting, there are two distinct stages: initiation andpropagation. During the initiation phase, a series ofvariations takes place in the composition of the solutionwithin the crevice, until passivity is destroyed by theconsumption of all the oxygen present in the crevice by thepassivity current (if this is higher than the velocity of oxygensupply inside the crevice) and, subsequently, by theestablishment of a galvanic current which, as in the case ofpitting, gives rise to the propagation stage, the accumulationof chlorides and the decrease in pH.

In the case of the deposit corrosion of active materials,preventive measures mainly involve the elimination ofenvironmental dishomogeneities and secondarily the use ofpaints and the application of cathodic protection. To preventcrevice corrosion in stainless steels, solutions withoutcrevices must be adopted during the design stage; in the caseof joints, it is preferable to use full penetration welds ratherthan superimposition or other mechanical methods such asflanges which do not eliminate crevices. As with pitting,cathodic protection is effective during both the initiation andpropagation stages.

Intergranular corrosionThis type of corrosion appears as a localized attack at

the grain boundary due to the presence of precipitates. Infact, the grain boundary is a preferential site for processesinvolving the segregation and precipitation of compounds(such as carbides, sulphides, intermetallic compounds). As ageneral rule, all alloys in which precipitates are present atthe grain boundaries are vulnerable to intergranularcorrosion. In practice, the alloys which have shown the

MATERIALS


chloride-free solution

solutions withincreasingchloride content

E

logj

Er

Fig. 8. Active-passive behaviour of a stainless steel in an aqueous solution, with and without chloride ions(increasing concentrations).

greatest vulnerability are stainless steels and some nickelalloys, in other words materials thought to be resistant tocorrosion.

For the sake of simplicity, the mechanism leading to thistype of attack can be subdivided into the following phases,typical of the stainless steel AISI 304 (18Cr-8Ni): the steel issupplied in the form described as annealed, obtained with asolution heat treatment which involves bringing the steel to atemperature of 1,050°C to solubilize the chromium carbides,followed by rapid cooling to prevent their precipitation; ifafter start-up the steel is brought to the critical temperaturerange 600-850°C, for example during welding operations ina thermally altered zone, the chromium carbides (mostprobable formula Fe23Cr23C6) precipitate at the grainboundaries, lowering the chromium content to below 12% inthe zones adjacent to the grain boundaries themselves; thesezones lose their stainless property and undergo corrosion ofgalvanic type, in which the cathodic (stainless) areas are thegrains; given the unfavourable ratio of (very small) anodicareas to (large) cathodic areas, the corrosion rate is generallyvery high even in modestly corrosive environments.

To prevent this form of corrosion, it is necessary toprevent the precipitation of the carbides by repeating thesolution heat treatment on the whole welded article. When itis not possible to carry out the solution heat treatment, theprecipitation of chromium carbides is prevented byfollowing one of two routes: lowering the carbon content tobelow a threshold value which is generally lower than 0.03%to make the formation of carbides kinetically more difficult(grade L steels, such as AISI 304L); or by adding alloyelements which have a greater affinity with carbon thanchromium, such as titanium or niobium, so that the carbon isalready alloyed and therefore unavailable to the chromium(AISI 321 and AISI 347 stabilized steels). Steels with a lowcarbon content (grade L) are thought to be safest; however,these have inferior mechanical properties given their lowercarbon content. Stabilized steels have better mechanicalproperties, but may give rise to an unusual form ofintergranular corrosion, known as knifeline attack.

Selective corrosionThis type of corrosion affects metal alloys and involves

the preferential dissolution of the less noble metal due toexposure to environments of medium corrosiveness. Themost important examples are: the dezincification of brass(copper-zinc alloys), where the zinc passes into solution andthe copper remains as a reddish metallic residue; thegraphitization of cast iron (mainly affecting grey cast iron),in which the iron matrix undergoes selective corrosion withthe formation of a more or less superficial residue layer ofgraphite.

Stress corrosion crackingThe term stress corrosion cracking (SCC) is used to

describe phenomena involving the formation andpropagation of cracks in a metal subject to the combinedaction of tensile (but not compressive) mechanical stressesand a corrosive environment. This is a complex phenomenonwhich has numerous aspects: for example, initiationconditions are characteristic of highly specific combinationsof a metallic material and an environment; the phenomenonis triggered only above a given mechanical stress threshold;it mainly affects alloys; the propagation rate of the cracks,

though high, is lower than that of cracks of purelymechanical nature. Tensile stresses may result from externalloads applied to the metallic component but also frominternal stress states, such as residual stresses frommechanical working, welding operations and heattreatments. The cracks may be of intergranular type withpropagation along the grain boundaries or of transgranulartype when propagation takes place through the metal grains.The cracks may be individual or branched.

In simple terms, which do not fully describe thiscomplex phenomenon, the stress corrosion crackingmechanisms can be subdivided into: a dissolutionmechanism, where the propagation and growth of the cracksoccurs due to the anodic dissolution of the tip, and amechanism known as embrittlement, where the cracks growdue to the mechanical fracture of the zone near the crack tip.The latter involves the entry of atomic hydrogen, producedby the cathodic reaction, into the crystal lattice of the metal,its diffusion to the crack tip which is in a state of strongplastic deformation and interaction with the metal, leading toembrittlement. The first mechanism is determined by theanodic process of metal dissolution, whilst the second isdetermined by the cathodic hydrogen release process. Thesimplest and clearest demonstration of the two mechanismsis the effect of anodic and cathodic polarization: anodicpolarization increases the dissolution velocity of the firstmechanism and inhibits the second, whilst cathodicpolarization blocks the first mechanism and accelerates thesecond. Table 6 shows the most common instances of SCC.

Stress corrosion cracking by chlorides. This mainlyaffects austenitic stainless steels of the 300 series such asAISI 304 and 316 and, to a lesser extent, austenitic-ferriticstainless steels. Ferritic stainless steels with a Ni contentbelow 4% are practically immune. For this condition tooccur, the following is required: a temperature above 60°C(the initiation temperature for this phenomenon rises thelower the concentration of chloride ions in solution); aconcentration of chlorides in solution above a thresholdwhich depends on the composition of the steel and in anycase above 10 ppm; an applied or residual tensile stresshigher than about 30% of the breaking load. Austeniticstainless steels with a high nickel content (above 15%) andnickel alloys offer increasing resistance to stress corrosioncracking as the content of alloying elements rises; in anycase, their resistance is significantly higher than steels oftype 304 and 316.

Hydrogen embrittlement. The hydrogen ion reductionreaction (2H��2e��

��H2), which takes place in acidenvironments, involves a first stage (H��e��

��Hads) withthe formation of atomic hydrogen adsorbed onto the metallicsurface (Hads). This atomic hydrogen may form a hydrogenmolecule H2 (Hads�Hads

��H2), or penetrate inside the

metal’s crystal lattice (Habs) if specific poisons are presenton its surface which inhibit the formation of molecularhydrogen. These poisons include metal sulphides andhydrogen sulphide (H2S), often present in environmentsassociated with hydrocarbons. The presence of H2S givesrise to the form of stress corrosion cracking due to hydrogenembrittlement known as sulphide stress cracking (SSC).

The steels susceptible to SSC are those with a high yieldlimit (breaking load above 700 MPa) whose microstructureis particularly sensitive to the effects of hydrogen. Theruptures are characterized by transgranular cracking.



The prevention of stress corrosion cracking takes placeabove all during the design phase by choosing materialswhich are not vulnerable in the environment specified, andby putting in place measures able to remove states ofresidual stress (for example, a heat extension treatment afterthe welding of joints). In some cases, recourse is made toenvironmental conditioning, for example by modifying thepH (this is the case with the treatment of drilling muds foroil wells with a pH above 10 when H2S is present). Cathodicprotection is only effective for stress corrosion cracking withan anodic dissolution mechanism.

Corrosion fatigueThe simultaneous presence of variable mechanical loads,

in other words fatigue stresses, and a corrosion process givesrise to a form of failure known as corrosion fatigue (CF). Theaggressive environment influences fatigue behaviour only dueto low frequency variations in the stresses, whilst at highfrequencies the mechanical aspects of fatigue prevail. In thepresence of CF, the fatigue limit for steels is eliminated andthe S-N curves (intensity of stress applied-number of cycles)take on a pattern similar to that of non-ferrous materials. Inpractice, as in the case of the design of oil platform nodes, thefatigue limit is reduced by up to 50% compared to fatigue inair. CF phenomena are present in marine environments insubsea pipelines and, above all, the nodes of oil platforms(Design […], 1990).

Two approaches are adopted to determine the corrosionfatigue life of a structure: the fatigue limit curves as a

function of the number of cycles and the fracture mechanicsexpressed by the crack growth velocity curves per cycle as afunction of the DK variation of the stress intensity factor.The use of fracture mechanics to predict the corrosionfatigue life of structures is fairly complex but helpful ingaining an understanding of the influence of the variousfactors, including cathodic protection in offshore structures.Except for materials susceptible to hydrogen embrittlement,such as steels with a high yield limit, cathodic protection isoften the ideal method for eliminating the negative effect ofthe environment on fatigue phenomena, bringing fatigueresistance to levels typical of inert environments such as air.On the other hand, even the well known fact according towhich the fatigue limit of steel improves when it is coveredwith anodic coatings (such as zinc) bringing it to close to thevalues obtained in air, confirms the beneficial effect ofcathodic protection. Nonetheless, the choice of the mostappropriate protection potential for structures operating atsea remains fairly complex because the effects of potentialdiffer in the two stages characterizing corrosion fatiguephenomena (the initiation and propagation periods) and theextent of the variation in loads (low or high values of DK).Generally speaking, there is agreement that the applicationof cathodic protection has a beneficial effect because theinitiation period is lengthened, even though the growth ratemay be accelerated.

In soils, the influence of cathodic protection on thefatigue behaviour of carbon steels is more complex andbecomes negative in the presence of carbonate-bicarbonate

MATERIALS


Table 6. Examples of SCC (Pedeferri, 2007)

Metal or alloy Environment Notes

Copper alloysSolutions containing ammonia,ammonium salts, amines

Transgranular cracks; season crackingof brasses

Copper alloys Mercury Intergranular cracking

Stainless steels Chloride solutionsAffects stainless steels of the 300 series(AISI 304 and 316); occurs when T�60°C;intergranular, branched cracking

Stainless steelsSolutions containing H2S;politionic acids

–

TitaniumChlorides in anhydrous methanol or ethanol

–

Carbon steels Alkaline solutionsCaustic embrittlement: affects boilers;T�200°C; intergranular cracking.

Carbon steels Nitrate solutionsCalcium and ammonium nitrates(fertilizer industry); T�100°C

Carbon steels Solutions containing sulphidesSulphide stress corrosion cracking: occursin deaerated, neutral or acid environments,typical of the petroleum industry.

Carbon steelsSolutions of carbonates, phosphates,cyanides; liquid ammonia

–

Steels with a high yield limit Damp air, aqueous solutions –

solutions and very slow load variations: in these cases, theprotection potential for corrosion fatigue is �0.70 V CSE, inother words underprotective compared to the generalizedcorrosion of steel. The propagation velocity, by contrast, isaccelerated under overprotection conditions at �1.4V CSEwhere hydrogen has a negative influence that does notemerge under stable load conditions.

Erosion corrosionThis form of corrosion appears in the shape of attacks

when the corrosion process is associated with severe fluiddynamic conditions, such as a high fluid velocity whichleads to the removal of corrosion products. The presence ofsuspended solids accentuates the phenomenon. This synergybetween corrosion and the mechanical action of the fluidaffects most metallic materials when given conditions ofturbulence expressed by the critical erosion corrosion rateare exceeded; in the oil industry, this rate is calculated withthe following empirical formula (API, 1991):

[47]

where ue is the critical corrosion erosion rate (m/s), C is aconstant depending on the metal and rm is the density of thefluid under operating conditions (Mg/m3). Constant C takeson different values depending on the type of material, thetype of service (continuous or intermittent) and the presenceof a corrosion inhibitor. The values of constant C in theabsence of inhibitors are 40 for copper, 60 for copper-nickel70/30, 120 for carbon steels and 500 for stainless steels.Whenever aggressive conditions are particularly severe, itbecomes necessary to use hard and corrosion resistantcoatings such as stellites and ceramic coatings.

Microbiological corrosionIn waters and soils, a type of corrosion takes place

known as MIC (Microbiologically Induced Corrosion),which is associated with the presence and action of bacteria(Biological […], 1985). This type of corrosion involves bothaerobic bacteria which live and grow in the presence ofoxygen, usually creating acid conditions, and anaerobicbacteria which live and grow in the absence of oxygen. Thelatter, thought to be more dangerous, include sulphatereducing bacteria (SRB).

It can be observed that conditions where oxygen isabsent, ideal from the electrochemical point of view foreliminating corrosion processes, are actually those whichencourage the growth of these bacteria and which thusprovide the worst results from the point of view of corrosion.In the literature, corrosion rates above 1 mm/yr have beenreported. Anaerobic bacteria are extremely adaptable andable to resist temperatures of up to 80°C and pressures of400 bar. In aerobic environments they are inactive, but cansurvive, ready to grow if anaerobic microenvironments arecreated (for example, beneath deposits). The corrosionmechanism of sulphate reducing bacteria is fairly complexand has not yet been fully clarified. The most widelyaccepted modern theory suggests that the enzymehydrogenase first produces the hydrogen used by thesulphate reducing bacteria to reduce sulphates to sulphides;this is followed by cathodic depolarization due to theprecipitation of iron sulphide; finally, anodic stimulation bythe sulphide ion occurs with the passing into solution of the

iron and the formation of occluded cells. The anaerobicconditions suited to the growth of sulphate reducing bacteriamay be determined by measuring the redox potential with aplatinum electrode: potentials at pH 7 lower than 0.1 V SHEindicate the absence of oxygen, whilst potentials which aremore noble than 0.4 V SHE indicate oxygenated conditions.The presence of bacteria, or more accurately the so-calledbiofilm, is responsible for triggering pitting corrosion onstainless steels in waters and especially in sea water. Thecause of this behaviour is associated with the nobilization ofthe cathodic processes caused by the biofilm, bringing thecorrosion potential above steel’s pitting potential.

Stray current corrosionThis type of corrosion, also known as electrolytic

corrosion, is caused by electrical interference phenomena(Lazzari and Pedeferri, 2006). It appears most frequently aslocalized corrosion on underground pipelines and tanks nearthe tracks of railway lines, tram lines and underground linespowered by a direct electrical current or in the vicinity of thedirect current groundbeds used in cathodic protectionsystems. The mechanism is as follows: the currentcirculating in the earth (a return current in direct currentelectrical systems or a protection current in cathodicprotection systems) uses the buried metal structure as anelectricity conductor, causing corrosion attacks in areaswhere the current leaves the structure. The current entryzones, by contrast, are in conditions of cathodic protectionand do not undergo corrosion. Attacks are often severebecause the exit zones of the current are localized wherethere are defects in the structure’s insulating coating, nearelectrical substations or cathodically protected structures; itis sufficient to note that 1 mA released by 1 cm2 of metalsurface leads to a loss of thickness of 10 mm/yr.

To limit the damage caused by stray currents, these mustbe prevented from entering the buried structures; to this endrecourse is made to various strategies which include: theimprovement of the insulating coating, the electricalsectioning of the structures (for example, by placinginsulating joints on pipelines), electrical drainage andequipotential connections.

Alternating current corrosion. Alternating currrents mayalso cause damage to structures affected by electricalinterference, but the density of the current exchangedbetween the metal and the earth must be above 30 A/m2 – inother words, five orders of magnitude greater than theinterference currents from a direct current. Instances ofinterference by alternating currents occur when structuresrun parallel to high voltage power lines, above 125 kV andby dispersion into the earth from high speed railway lines.Corrosion attacks occur especially on structures equippedwith excellent insulating coatings, where there are defects inthe coating. The main preventive and control measure foralternating current corrosion is the correct positioning of thestructure in the ground.

9.1.5 Internal corrosion in oil wells

General corrosion conditionsSurfaces in contact with hydrocarbons may undergo

corrosion only if they are wetted by the water phase (waterwetting) in a permanent or intermittent way. The conditions

u Ce

m

m

=r



giving rise to water wetting vary depending on thecomposition of the hydrocarbon phase and the operatingconditions; according to an empirical rule which is oftenadopted, water wetting occurs in vertical flows (for example,inside the tubing of wells) when the percentage of theaqueous phase exceeds 20% in volume, whereas inhorizontal flows (for example, inside transportationpipelines) this percentage may fall to as low as 1%. Waterwetting is estimated not with reference to the total watercontent (or water cut) but to the free water phase, in otherwords water not emulsified in the hydrocarbon phase. In oilwells, the aqueous phase has the same composition asformation water, with variable but always high salinity,generally above 100 g/L. In gas and gas condensate wells,the aqueous phase may be present as entrained formationwater or as condensed water from the water vapourcontained within the gas or as a mixture of the two, so thesalinity is generally very low, ranging from that of distilledwater to a few g/L.

Water wetting conditions are influenced by the presenceof gas and vary depending on temperature and pressure; thepresence of gas is correlated with the bubble point pressureat the temperature considered: if the pressure is above thebubble point pressure, all the gas is dissolved in the liquidhydrocarbon phase.

Internal corrosion in oil wells may take various forms:some have already been examined (corrosion by galvaniccontact, stress corrosion cracking, erosion corrosion,microbiological corrosion), whereas others are specific tothe oil industry (corrosion by CO2, corrosion by H2O,sulphur corrosion).

Corrosion by CO2The presence of CO2 in the aqueous phase of

hydrocarbons leads to the formation of carbonic acid which,though it is a weak acid, is extremely aggressive compared tocarbon steel. For example, comparing the corrosion rate ofiron in carbonic acid and that in a strong acid such ashydrochloric acid, both with a pH of 4, the corrosion rate incarbonic acid is about ten times higher. The reason for this isthe different kinetics of the reaction of the hydrogen ionreduction reaction and the release of molecular hydrogen; inthe case of carbonic acid, the bicarbonate ion participatesdirectly in the reduction and hydrogen release reaction.

Corrosion by CO2 gives rise to: generalized corrosion attemperatures below 80°C, with a characteristic morphologyknown as mesa corrosion; generalized corrosion withsignificant localization in the temperature range 80-120°Cdue to the formation of partially protective Fe and Cacarbonates which create passivation conditions with anegligible corrosion rate at T�120°C. Generally speaking,the corrosion products are not protective at T�60°C and pCO2

�5 bar, whereas they are partially protective at60°C�T�100°C and protective at T�100°C. The additionof chromium to steels improves corrosion resistance by asmuch as 0.5% (especially resistance to mesa corrosion) attemperatures below 90°C.

As shown by the results of numerous laboratoryexperiments and field data, the corrosion rate of carbon steeldepends on three parameters: the partial pressure of the CO2(or fugacity when pressures are high), temperature and thepH of the aqueous phase (linked to salinity/ionic strengthand the concentration of hydrogencarbonates). The partial

pressure is given by the product of the mole fraction of CO2in the separated gas phase and the working pressure in thepresence of gas in equilibrium with the liquid phases or thebubble point pressure if the gas is dissolved in the liquidphase under operating conditions.

Calculating the corrosion rateIn the 1970s, C. De Waard and D.E. Milliams proposed

an equation developed from experimental laboratory resultsto calculate the corrosion rate in the presence of CO2 ingases/gas condensates, subsequently corrected with U. Lotz(De Waard and Milliams, 1975; De Waard et al., 1991):

[48]

where vCO2,dWLM is the corrosion rate in mm/yr and pCO2the

partial pressure of CO2 in the gas phase. Starting from thisbasic equation, correction factors have been introducedwhich take into account the fugacity for pressures above 100bar (system factor), the formation of deposits (scale factor)and the pH (pH factor).

Frequently, the pH is unknown because it is difficult tomeasure: recourse is made to calculations on the basis ofthe composition of the aqueous phase, the partial pressureof CO2 and H2S and the concentration ofhydrogencarbonates. The following empirical formulae areoften used (in the presence of bicarbonates and in theabsence of hydrogencarbonates, respectively):

[49] pH �4.4 �0.475ln(PCO2�PH2S) �

� 0.5ln[HCO3�] �0.00375T

[50] pH �3.8 �0.195ln(PCO2�PH2S) �0.00375T

where pCO2and pH2S are the partial pressures in the gas of

CO2 and H2S, T is the temperature (°C) and HCO3� is the

concentration of hydrogencarbonates in solution in meq/L. Ifthe concentration of HCO3

� is unknown, it is assumed that itis zero in the case of condensates alone (ionic strength m�0.5)and equal to 10 meq/L in the presence of formation water(ionic strength m�0.5). To calculate the pH, reference canalso be made to NorSoK M-506 Model (NorSoK, 1998). Alater model has been proposed (De Waard et al., 1995; EFC,1997a) which also takes into consideration the diffusion ofCO2. However, in order to apply this model, it is necessary toknow the fluid dynamic conditions of the aqueous phase.

Predictive modelsTo provide a rough calculation of the corrosiveness of

hydrocarbons containing CO2, the American PetroleumInstitute proposed an empirical rule (API, 1958) in the 1950sbased on the partial pressure of the CO2: if pCO2

�7 psi (0.5 bar), corrosion by CO2 is almost negligible, becomingsignificant when pCO2

�14 psi (1 bar). However, these rulesare considered insufficient to evaluate corrosion by CO2.

An empirical model derived from specific experiments(Crolet and Bonis, 1985) provides a classification ofcorrosion by CO2 as high, medium and low (where themedium class indicates a corrosion rate of 1.2 mm/yr) basedon: the partial pressure of CO2, pH, the content of organicacids (acetic and formic), the Ca2��HCO�

3 ratio. Forcondensed water, corrosiveness is low if pCO2

�0.05 bar or ifpCO2

�0.2 bar and acetic acid<0.1 meq�L; medium if0.2�pCO2

�5 bar and acetic acid�0.1 meq�L; in other cases,

log .,

. log,

uT

pCO dWLM CO2 2= −

++5 8

1 710

2730 67

MATERIALS


corrosiveness is high. For formation water, corrosiveness islow if pCO2

�0.05 bar or pH�5.6 or 0.05�pCO2�10 bar and

Ca2��HCO�3 �0.5 meq�meq; medium if acetic acid�0.1

meq�L or Ca2��HCO�3 �1,000 meq�meq; in other cases,

corrosiveness is high.

Corrosion by H2SThe presence of H2S leads to various forms of corrosion:

the generalized corrosion of carbon steels; the formation ofblisters and cracks, also known as HIC (Hydrogen InducedCracking); SSCC (Sulphide Stress Corrosion Cracking) ofmaterials susceptible to hydrogen embrittlement. Recently,other forms of corrosion resulting from a combination ofHIC and SSC have also been suggested, such as SOHIC(Stress Oriented Hydrogen Induced Cracking; EFC, 1997b;ISO, 2003).

Generalized corrosion by H2SH2S is a weak acid, but despite this it causes particularly

severe corrosion attacks because the formation of ironsulphide FeS, with an extremely low solubility product (inthe order of 10�24), leads to a decrease in the anodicpotential of iron and the consequent availability of a drivingforce even in neutral solutions. The sulphide is not alwayscompact and protective, and may lead to localized corrosionwith a galvanic contact mechanism due to the fact that it hasan electron conductivity and practical nobility higher thanthose of iron. The prediction of the corrosion rate is based onvarious empirical rules which can be summarized as follows:at T�60°C the corrosion rate is negligible for H2Sconcentrations in solution below 40 ppm but becomessignificant (0.5 mm/yr) at concentrations above 0.4%; atT�60°C the corrosion rate is 1 mm/yr at concentrationsabove 40 ppm.

Hydrogen induced cracking (HIC)In addition to hydrogen embrittlement, the entry of

atomic hydrogen into the metal may cause blistering,involving the formation of blisters and cracks as an effect ofthe recombination of the hydrogen atoms into molecularhydrogen, H2, where there are traps in the form of inclusions

or microvacuums in the metallic matrix. The hydrogenmolecules, too large to diffuse through the metal lattice,accumulate and generate extremely high internal pressures,sufficient to cause local plastic deformation of the metal andblistering. The inclusions with the greatest impact on thisphenomenon are those of manganese sulphide (MnS2 typeII), which during the hot lamination of the carbon steels usedfor pipelines and the sheets for pressurized containers aresqueezed and arranged parallel to the direction oflamination, thus forming an easy trap for the hydrogenatoms. The phenomenon is linked to the quantity ofhydrogen penetrating into the metal and time; if the partialpressure of H2S is above 0.1 bar (concentration in theaqueous phase of about 400 ppm), it emerges on vulnerablesteels in a time comparable to the mean project life of oilfacilities (15-20 years); at partial pressures below 0.1 bar, thephenomenon is nonetheless present on vulnerable steels andemerges over a longer time. When the concentration of H2Sis significant, for example pH2S�0.03 bar, two differentstrategies can be adopted (or both) to limit HIC: the use ofcorrosion inhibitors, which reduces the quantity of hydrogenproduced, and the use of steels which are not vulnerable tothis phenomenon. The latter may be of two types: steelstreated with rare earths, in which the sulphur contained inthe steel is combined with elements whose affinity is higherthan manganese, such as calcium and caesium which formhard sulphides that are not squeezed during hot lamination;steels with a low sulphur content (20-50 ppm) so that thequantity of Mn sulphide able to form is insufficient for thephenomenon to set in.

Sulphide Stress Corrosion Cracking (SSCC)The hydrogen embrittlement of vulnerable steels in the

presence of H2S is a phenomenon which depends onnumerous factors: the steel’s composition, heat treatments,microstructure, mechanical resistance, the pH, the partialpressure of the H2S, the load applied, temperature and time.

The onset of SSCC is evaluated based on the partialpressure of H2S in the gas phase, proportional to itsconcentration in the aqueous phase. The partial pressure ofH2S is calculated starting from the mole fraction in the gas



IG

IL

xH2S =15%

GOR � 890P = 18 bar

GOR � 890

sour service

1

10

100

1,000

0.0001 0.001 0.01 0.1 1 10 100

tota

l pre

ssur

e (b

ar)

P = 4.5 bar

pH2S= 0.7 bar

pH2S = 0.0035 bar

hydrogen sulphide molar fraction in gas phase, xH2S (%)

Fig. 9. Identification of conditions for the onset of SSCC(according to NACE, 2003).

and the working pressure (as for corrosion by CO2) takingaccount of whether the gas is free (pressure below thesaturation pressure of the gas or the bubble point pressure)or dissolved in the liquid phase (pressure above thesaturation pressure of the gas or the bubble point pressure).The calculation of partial pressure becomes complicatedwhen the mole fraction of H2S is not known with sufficientaccuracy; in these cases, the estimate of partial pressure isbased on the mass balance of the H2S according toprocedures described in specialist manuals.

Sour environment evaluation according to NACE(2003). The environment is described as sour when thepartial pressure of H2S and the concentration of H2S exceedgiven thresholds, as a function of the type of fluid and theoperating conditions. For gas wells, Fig. 9 is used; this graph,showing the mole fraction of H2S in the gas and the totalpressure on the x-axis, identifies the region in which theSSCC phenomenon emerges on vulnerable steels and alloys(P�4.5 bar and pH2S�0.0035 bar).

In the presence of a multiphase system (liquid phaseand gas phase), the following are considered conditions ofnon-onset: Gas Oil Ratio, GOR, �890 Nm3/m3; molefraction of H2S in the gas, x, �0.15 and pH2S�0.7 bar;working pressure P�18 bar. Under these conditions, giventhe low pressure, mechanical stresses are generallymoderate, whilst the prevalence of the oil phase (lowGOR) prevents contact with the water phase. Only in thepresence of a high water content should the definition ofsour environment be reconsidered, with the possible onsetof SSCC in the system.

For multiphase systems which do not fall into the aboveclassification, sour environment conditions can besummarized as follows: P�18 bar and pH2S�0.0035 bar;P18 bar, pH2S�0.0035 bar and GOR�890 Nm3/m3;4.5�P�18 bar, pH2S�0.7 bar and GOR�890 Nm3/m3;P�4.5 bar, xH2S�0,15 and GOR�890 Nm3/m3.

‘Sour environment’ evaluation according to ISO (2003).The criteria used to define a sour environment according tothe standard ISO 15156, which integrates the results of EFCdocuments (De Waard et al., 1991; NorSoK, 1998) and theNACE criteria, are based on three parameters: pH2S, pH ofthe aqueous phase (pHACT) and temperature. The pH can becalculated on the basis of the composition of the aqueousphase using appropriate calculation programmes.

Very severe sour conditions: pHACT3.5 or0.01�pH2S�1 bar and pHACT5.5�logpH2S or pH2S�1 barand pHACT5.5. Severe sour conditions: 0.001�pH2S�0.01bar and 3.5�pHACT6.5�logpH2S or 0.01�pH2S�1 bar and5.5�logpH2S�pHACT6.5�logpH2S or pH2S�1 bar and5.5�pHACT6.5. Moderate sour conditions: 0.001�pH2S�1bar and pHACT�6.5�logpH2S or pH2S�1 bar and pHACT�6.5.Non-sour (sweet) conditions: pH2S�0.001 bar andpHACT�3.5.

Corrosion by sulphurSulphur (melting point 113°C and boiling point 445°C)

is present in some petroleum reservoirs in association withH2S. The formation of sulphur may occur due to thereduction of sulphates to sulphur by methane(SO4

2��CH4��S�2H2O�CO2) or by catalytic

decomposition of hydrogen sulphide or the oxidation ofhydrogen sulphide by oxygen (due to contamination).Sulphur reacts with hydrogen sulphide to form polysulphides

which in turn release sulphur to form deposits (slugs) andentrainment in the tubing of oil wells.

Sulphur is an effective cathodic reagent, acting as anelectron acceptor through the catalytic action of thesulphide films which have electron conductivity, and givesrise to high corrosion rates when it comes into directcontact with steels; if it is present in a hydrocarbonsolution, the corrosion rate is controlled by diffusionprocesses. At high temperatures (above 120°C), the sulphurdismutes to H2S and sulphuric acid, causing the rupture ofthe passivation films and the onset of localized corrosion(crevice and SSC) on stainless alloys (Corrosion in the oil[…], 1998).

The sulphur corrosion rate on carbon and low alloysteels depends on temperature: when T�120°C and directcontact with sulphur, the corrosion rate is in the order ofseveral mm/yr and is only slightly influenced bytemperature. The sulphides present on the surface catalysethe reaction and lead to macrocell conditions similar tocorrosion by differential aeration. At temperatures above120°C, the corrosion rate increases due to the formation ofsulphuric acid, but only up to about 150°C; above thistemperature the corrosion rate begins to decrease due to theprotective action of the corrosion products. At 180°C thecorrosion rate is nonetheless so high (above 10 mm/yr) thatthese materials cannot be used.

The presence of elementary sulphur generally leads tothe localized corrosion of stainless steels: austenitic-ferritic (duplex) stainless steels undergo generalizedcorrosion even at ambient temperature like carbon steels,whereas austenitic stainless steels present low resistanceonly above 120°C. Martensitic stainless steels generallyhave low resistance. The presence of sulphur increasesvulnerability to SCC: when T �120°C, resistance tocorrosion increases with increasing Ni, Cr and Mocontent. Nickel alloys offering good resistance in severelyaggressive conditions (S, H2S, CO2, Cl) at hightemperatures must generally have a basic composition ofthe type Ni�55%, Mo�12%, Cr�15% (typically AlloyG-50 22Cr-52Ni-11Mo-0.7W-0.8Cu). Grade 2 titaniumgives rise to crevices when T �130°C; beta-C titaniumalloys present greater resistance.

Bibliography

Bennet L.H. et al. (1978) Economic effects of metallic corrosion inthe United States, Washington (D.C.), National Bureau of Standards,511, 1-3.

EFC (European Federation of Corrosion) (2002) Guidance on materialsrequirements for corrosion resistant alloys exposed to H2S in oiland gas production, London, The Institute of Materials, 17.

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Luciano Lazzari

Dipartimento di Chimica, Materiali eIngegneria chimica ‘Giulio Natta’

Politecnico di MilanoMilano, Italy



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