Corrosion in Acid Gas Solutions

8/10/2019 Corrosion in Acid Gas Solutions

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AH Solid state diffusion kinetic constant for H

through mackinawite film,

AH 4:0 104 molm2 s1

ACO2 Solid state diffusion kinetic constant for CO2

through mackinawite film,

ACO2 2:0 106

molm

2

s

1

baFe Anodic Tafel slope for Fe oxidation (V)

bcH Cathodic Tafel slope for H+ ion reduction (V)

bcH2CO3 Cathodic Tafel slope for H2CO3

reduction (V)

bcH2O Cathodic Tafel slope for H2O reduction (V)

BFeCO3 Constant in the Arrhenius-type equation

forkrFeCO3 (kJ mol1)

cCO2 Bulk aqueous concentration of CO2 (kmol m3)

cCO23

Bulk aqueous concentration of CO23 ions

(kmol m3)

cFe2 Bulk aqueous concentration of Fe2 ions

(kmol m3)

cH Bulk aqueous concentration of H ions

(kmol m3)

csH Near-zero concentration of H underneath

the mackinawite film at the steel surface, set

to 1:0 107 (kmol m3)

cHCO3

Bulk aqueous concentration of HCO3 ions

(kmol m3)

cH2 CO3 Bulk aqueous concentration of H2CO3

(kmol m3)

cH2 SBulk aqueous concentration of H2S (kmol m3)

cHS Bulk aqueous concentration of HS ions(kmol m3)

ci Bulk aqueous concentration of a given aqueous

species (kmol m3)

ciH2 S Aqueous concentration of H2S at the inner

sulfide film/outer sulfide layer interface

(kmol m3)

cS2 Bulk aqueous concentration of S2 ions

(kmol m3)

csH2S Near-zero aqueous concentration of

H2S underneath the mackinawite film

at the steel surface, set to 1:0 107

(kmol m3)

coH2 S Aqueous concentration of H2S at the

outer sulfide layer/solution interface

(kmol m3)

csCO2 Aqueous concentration of CO2 underneath

the mackinawite film at the steel surface

dCharacteristic dimension for a given flow

geometry (m)

dp Diameter of a pipe (m)

dc Diameter of a rotating cylinder (m)

DDiffusion coefficient of a given species (m2 s1)

DH2CO3 Aqueous diffusion coefficient of H2CO3

(m2 s1)

DrefH2 CO3 Reference aqueous diffusion coefficient

of H2CO3,Dref;H2CO3 1.3 109 m2 s1 at

25C

DH

Aqueous diffusion coefficient for H

DrefH Reference aqueous diffusion coefficient for

H,DrefH 2.80 108 m2 s1 at 25 C

DH2S Aqueous diffusion coefficient for dissolved

H2S

DCO2 Aqueous diffusion coefficient for dissolved

CO2,DCO2 1.96 109, m2 s1

EPotential (V)

Ecorr Corrosion (open circuit) potential (V)

ErevFe Reversible potential of Fe oxidation,

ErevFe 0.488 V

ErevHReversible potential for H ion reduction (V)

ErevH2 CO3 Reversible potential for H2CO3

reduction (V)

ErevH2 O Reversible potential for H2O reduction

(A m2)

fH2CO3 Flow factor for the chemical reaction

boundary layer

FFaradays constant,F 96485 C mol1eFluxH2S Flux of H2S (kmol m

2 s1)

FluxH Flux of H ions (kmol m2 s1)

FluxCO2 Flux of CO2 (mol m2 s1)

HsolCO2 Henrys constant for dissolution of CO2

(bar kmol m3

)DHFe Activation enthalpy for Fe oxidation,

DHFe 50kJ mol1

DHH Activation enthalpy for H ion reduction,

DHH 30kJmol1

DHH2CO3 Activation enthalpy for H2CO3 reduction,

DHH2 CO3 57.5kJ mol1

DHH2O Activation enthalpy for H2O reduction,

DHH2 O 30kJmol1

iCurrent density (A m2)

icorr Corrosion current density (A m2)

iaFe

Anodic current density of iron oxidation

(A m2)

icH Cathodic current density for H ion reduction

(A m2)

icH2 CO3 Cathodic current density for H2CO3

reduction (A m2)

icH2 O Cathodic current density for H2O reduction

(A m2)

idlimH Mass transfer (diffusion) limiting current

density for H ion reduction (A m2)

irlimH2 CO3

Chemical reaction limiting current density

for H2CO3 reduction (A m2)

Corrosion in Acid Gas Solutions 1271


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ioFe Exchange current density of iron oxidation

(A m2)

ioHExchange current density for H ion reduction

(A m2)

ioH2 CO3 Exchange current density for H2CO3

reduction (A m

2

)ioH2 O Exchange current density for water

reduction (A m2)

irefoFe

Reference exchange current density of Fe

oxidation, irefoFe 1 A m

2

irefoH Reference exchange current density of H

oxidation, irefoH 0.03 A m

2 atTc;ref 25C

and pH 4

irefoH2 CO3

Reference exchange current density for

H2CO3 reduction,irefoH2 CO3

0.06 A m2 at

Tc;ref 25C, pH 5, andcH2 CO3;ref 10

4

kmol m3

irefoH2 O

Reference exchange current density for H2O

reduction,irefoH2O

3 105A m2 at

Tc;ref 20C

iaH Charge transfer current density for H ion

reduction (A m2)

iaH2 CO3 Charge transfer current density for H2CO3

reduction (A m2)

IIonic strength kmol m3

kbhyd Backward reaction rate of H2CO3 dehydration

reaction (1 s1),kbhyd=kfhyd=Khyd

kfhyd Forward reaction rate for the CO2 hydration

reaction (1 s1

)kmH Aqueous mass transfer coefficient for H

(A m2)

kmH2 CO3 Aqueous mass transfer coefficient for

H2CO3 (A m2)

kmH2 S Aqueous mass transfer coefficient for H2S

(A m2)

kmCO2 Aqueous mass transfer coefficient for CO2

(A m2)

krFeCO3 Kinetic constant in the ferrous carbonate

precipitation rate equation (1 mol1 s1)

Khyd Equilibrium hydration constant for CO2

,

Khyd kfhyd=kbhyd 2:58 10

3

Kbi Equilibrium constant for dissociation of HCO3

(kmol m3)

Kbs Equilibrium constant for dissociation HS

(kmol m3)

Kca Equilibrium constant for dissociation of H2CO3

(kmol m3)

Khs Equilibrium constant for dissociation H2S

(kmol m3)

KsolH2S Solubility constant for dissolution of H2S

(kmol m3 bar1)

KsolCO2 Solubility constant for dissolution of CO2

(kmol m3 bar1)

KspFeCO3 Solubility product constant for ferrous

carbonate (kmol m3 bar1)

KmackinspFeS Solubility product constant for

mackinawite (kmol m

3

bar

1

)mos Mass of the outer sulfide layer (kg)

MFe Molecular mass of iron (kg kmol1Fe)

MFeS Molecular mass of ferrous sulfide

(kgmol1FeS)

n Number of electrons used in reducing or oxidizing

a given species (kmole kmol1)

pCO2 Partial pressure of CO2 (bar)

pH2S Partial pressure of H2S (bar)

RElectrochemical reaction rate

(kmol m2 s1)

RFeCO3 Precipitation rate for iron carbonate

(kmol m3 s1)

RUniversal gas constant,R 8.314 J mol1 K1

ReReynolds number,Re vrH2 Od=mH2OScSchmidt number of a given species,

Sc mH2 O=rH2 OD

Shp Sherwood number of a given species

for a straight pipe flow geometry,

Shp kmdp=D

Shr Sherwood number of a given species

for a rotating cylinder flow geometry,

Shr kmdc=D

SSFeCO3 Supersaturation of iron carbonateSTScaling tendency

Tc Temperature (C)

Tc;ref Reference temperature, Tc;ref= 25C

Tf Temperature (F)

TkTemperature (K)

vWater characteristic velocity (m s1)

zi Species charge of various aqueous species

dmH2 CO3 Thickness of the mass transfer layer for

H2CO3 (m)

drH2CO3 Thickness of the chemical reaction layer

for H2

CO3

(m)

dos Thickness of the outer sulfide layer (m),

dos mos=rFeSA

DtTime interval (s)

mH2 O Water dynamic viscosity (Pa s

mH2 O;refReference water dynamic viscosity (Pa s) at

a reference temperature,

mH2 O;ref 1:002 104 Pas at 20 C

zH2CO3 Ratio of the mass transfer layer and

chemical reaction thicknesses for

H2CO3

eOuter sulfide layer porosity

1272 Liquid Corrosion Environments


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cOuter sulfide layer tortuosity factor

rH2 O Density of water (kg m3)

rFe Density of iron (kg m3)

rFeS Density of ferrous sulfide (kg m3)

2.25.1 Introduction

As oil and gas emerge from the geological formation,

they are always accompanied by some water and

varying amounts of acid gases: carbon dioxide,

CO2, and hydrogen sulfide, H2S. This is a corrosive

combination, which affects the integrity of mild steel.

This has been known for over 100 years; aqueous CO2and H2S corrosion of mild steel still represents a

significant problem for the oil and gas industry.1

Although corrosion resistant alloys that are able to

withstand this type of corrosion exist, mild steel is

often the most cost effective construction material

used in this industry for these applications. All the

pipelines, many wells, and much of the processing

equipment in the oil and gas industry are built out of

mild steel. The cost of equipment failure due to internal

CO2/H2S corrosion is enormous, both in terms of

direct costs such as repair costs and lost production,

as well as in indirect costs such as environmental cost,

impact on the downstream industries, etc.

The following section summarizes the degree of

understanding of the so-called sweet CO2corrosionand the so-called sour or H2S corrosion of mild steel

exposed to aqueous environments. It also casts the

knowledge in the form of mathematical equations

whenever possible. This should enable corrosion

engineers and scientists to build entry level corrosion

simulation and prediction models.

2.25.2 Aqueous CO2Corrosion ofMild Steel

Aqueous CO2corrosion of carbon steel is an electro-

chemical process involving the anodic dissolution of

iron and the cathodic evolution of hydrogen. The

overall reaction is

Fe CO2 H2O ! FeCO3 H2 1

CO2corrosion of mild steel is reasonably well under-

stood. A number of chemical, electrochemical, and

transport processes occur simultaneously. They are

briefly described below.

2.25.2.1 Chemistry of CO2Saturated

Aqueous Solutions Equilibrium

Considerations

CO2gas is soluble in water:

CO2g ,

Ksol

CO2 2For ideal gases and ideal solutions in equilibrium,

Henrys law can be used to calculate the aqueous

concentration of dissolved CO2, cCO2 , given that the

respective concentration in the gas phase (often

expressed in terms of partial pressure,pCO2 ) is known:

HsolCO2 1

KsolCO2

pCO2cCO2

3

The CO2solubility constant,KsolCO2, is a function of

temperature, Tf, and ionic strength, I2:

KsolCO2

14:5

1:00258 102:275:6510

3Tf8:06106T2

f0:075I

4

Ionic strength,I, can be calculated as

I 1

2

Xi

ciz2i

1

2c1z

21 c2z

22 5

The concentration of CO2in the aqueous phase is of

the same order of magnitude as the one in the gasphase. For example, at pCO2 1 bar, at 25C, the gas-

eous CO2 concentration is 4 mol l1 (kmol1 m3)

while in the water it is about 3 mol l1. Since the

solubility of CO2 decreases with temperature, at

100 C, the respective concentrations are 3.3 mol l1

in the gas and 1.1 mol l1 in water.

A rather small fraction (about 1 in 500) of the

dissolved CO2 molecules hydrates to make a weakcarbonic acid, H2CO3:

CO2 H2O ,Khyd

H2CO3 6

due to a relatively slow forward (hydration) rate.

Assuming that the concentration of water remains

unchanged, the equilibrium concentration cH2CO3 is

determined by:

Khyd cH2CO3

cCO27

The equilibrium hydration/dehydration constant,

Khyd 2:58 103, does not change much across

the typical temperature range of interest (20100 C).3

Carbonic acid is considered to be weak because

it only partially dissociates in water to produce



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hydronium, H ions and bicarbonate ions, HCO3:

H2CO3 ,Kca

H HCO3 8

The HCO3 dissociates further to give some more H

and carbonate ion, CO32:

HCO3 ,Kbi

H CO23 9

The respective equilibrium relations can be written as

KcacHcHCO3

cH2CO310

KbicHcCO23

cHCO311

The equilibrium constants can be calculated as func-

tions of temperatureTf, and ionic strength,Ias2

Kca 387:6 106:411:59410

3Tf8:5210

6

T2f3:07105p14:70:4772I0:5 0:118I

12

Kbi 1010:614:97103Tf1:33110

5 T2f2:624105

p14:71:166I0:50:3466I

13

One can use the equations above to calculate the pH

for a pure aqueous CO2 saturated system. Assuming

that the concentration of CO2 (or partial pressure,

pCO2 ) in the gas phase is known, one can calculate

the concentration of aqueous/dissolved CO2 cCO2 ,

via eqn [3]. Then the concentration cH2CO3 can bedetermined via eqn [7]. However, in the remaining

two eqns [12] and [13], there are three unknowns:

cH , cHCO3 , andcCO23 , and therefore one more equa-

tion is needed to close the system: a constraint that

describes charge conservation, that is, electroneutral-

ity of the solution. Clearly, chemical reactions [8] and

[9], which involve ions, always remain balanced with

respect to charge and therefore one can write

cH cHCO3 2cCO23 14

Now, the system of equations is closed and concentra-

tions of all the aqueous species can be determined,

including thecH and the corresponding pH. The pH

of pure water as a function ofpCO2 at room tempera-

ture is shown inFigure 1.

If there are other ions in the aqueous solution,

such as for example Fe2 produced by corrosion of

steel, theneqn [14]is extended to read

2cFe2 cH cHCO3 2cCO23 15

By inspecting the equations above, one can see that,

as iron dissolution causes an increase in cFe2 , it isaccompanied by a decrease ofcH due to the cathodic

reaction and a corresponding increase in pH. Other

cations and anions as well as other chemical reactions

can be introduced into the mix in a similar way.

An example of a CO2aqueous species distribution as

afunctionofpHforanopensystemisgiveninFigure 2.

It is worth noting that this simple water chemistry

calculation procedure is valid only for the case whenthe concentration of gaseous CO2, i.e., the partial

3

4

5

6

0.001

pCO2(bar)

pH

3 wt% NaCl

Pure H2O

0.01 0.1 101 100

Figure 1 Calculated pH of a pure aqueous solution saturated with CO2as a function of partial pressure of CO2;T 25C,

1 wt% NaCl.



6/29

pressure, pCO2 is known, constant, and independent

from what is happening in the aqueous phase. This is

often referred to as an open system. It is relevant to

field situations where there is an overwhelming

amount of CO2 in the gas phase (such as seen in

wet gas lines, multiphase pipelines, gas/liquid sep-

arators, etc.). In the lab setting, this condition iseasily achieved by continuous purge of a vessel with

gaseous CO2.

In contrast, there are many systems where there

is a limited amount of CO2 in the gas phase com-

pared to the amount in the liquid phase, such as in

oil well tubing, oil transportation lines, liquid/liquid

separators, etc. In the lab, aqueous systems with a

limited gas phase are frequently found in high-

pressure autoclaves and flow loops. Consequently

they are often referred to as closed systems, and in

principle can have varying gas/liquid volume ratios.

Anopensystem can be seen as a closedsystem with an

infinitely large gas/liquid volume ratio. In closedsystems, the concentration of gaseous CO2, that is,

the partial pressure, pCO2 , is not known explicitly

and typically depends on the aqueous chemistry. In

mathematical terms, this means that there is onemore unknown: pCO2 , and therefore one needs one

more equation to be able to solve for species con-

centrations. The extra equation comes from the

additional constraint: in a closed system, the total

amount of carbonic species remains constant; they

are just redistributed between the gas and aqueous

phases as conditions change. When one accounts for

this, an extra equation is obtained:

nCO2g nCO2aq nH2CO3aq nHCO3 aq

nCO23 aq Const: 16

wherendenotes the number of moles of a particularspecies in a gaseous or aqueous phase of a closedsystem.

The dissociation steps [8] and [9] are very fastcompared to all other processes occurring simulta-

neously in corrosion of mild steel, thus preserving

chemical equilibrium. However, the CO2dissolution

reaction [2] and the hydration reaction [6] are much

slower. When such chemical reactions proceed

slowly, other faster processes (such as electrochemi-

cal reactions or diffusion) can lead to local nonequi-

librium in the solution.

Either way, the occurrence of chemical reactions

can significantly alter the rate of electrochemical pro-

cesses at the surface and the rate of corrosion. This is

particularly true when, due to high local concentra-

tions of species, the solubility limit of salts is exceeded

and precipitation of a surface layer occurs. In a precip-

itation process, heterogeneous nucleation occurs first

on the surface of the metal or within the pores of anexisting layer since homogenous nucleation in the

bulk requires a much higher concentration of species.

Nucleation is followed by crystalline layer growth.

1.E07

1.E06

1.E05

1.E04

1.E03

1.E02

1.E01

1.E+00

2 3 4 5 6 7

pH

Speciesconcentration

(moll1)

HCO3

H2CO3

CO2

CO2

CO2(g)

3

Ferrous carbonate

Mild steel

Figure 2 Calculated carbonic species concentrations as a function of pH for a CO2saturated aqueous solution;pCO2 1 bar,25 C, 1 wt% NaCl.



7/29

Under certain conditions, surface layers can become

very protective and reduce the rate of corrosion.

In CO2 corrosion, when the concentrations of

Fe2 and CO32 ions exceed the solubility limit,

they form solid ferrous carbonate according to

Fe2 CO23 ,KspFeCO3 FeCO3s 17

where the solubility product constant for ferrous

carbonateKspFeCO3 is4

KspFeCO3 1059:34980:041377Tk2:1963=Tk

24:5724 log Tk2:518I0:50:657I 18

Actually ferrous and carbonate ions are frequently

found in the aqueous solution at concentrations much

higher than predicted by the equilibrium KspFeCO3.

This is termedsupersaturationand is a necessary con-

dition before any substantial precipitation can occur.The ferrous carbonate supersaturation, SSFeCO3, is

defined as:

SSFeCO3 cFe2cCO23KspFeCO3

19

The precipitation process can be seen as the process

of the solution returning to equilibrium and is driven

by the magnitude of supersaturation. The rate of the

precipitation (


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barium sulfate, strontium sulfate, etc. The presence

of calcium carbonate, in particular, can have a bene-

ficial effect upon corrosion and upon the stability of

the FeCO3 scale. Finally, in the presence of H2S,

various types of sulfides form as discussed in a sepa-

rate section below.

2.25.2.2 Electrochemistry of Mild Steel

Corrosion in CO2Saturated Aqueous

Solutions

The electrochemical dissolution of iron in a water

solution:

Fe ! Fe2 2e 23

is the dominant anodic reaction in CO2 corrosion.The reaction is pH dependent in acidic solutions

with a reaction order with respect to OH between

1 and 2, decreasing toward 1 and 0 at pH > 4, which isthe typical range for CO2corrosion. Measured Tafel

slopes are typically 3080 mV. This subject, which is

still somewhat controversial with respect to the

mechanism, has been reviewed for acidic corrosion6,7

and CO2solutions.8

The presence of CO2increases the rate of corro-

sion of mild steel in aqueous solutions primarily by

increasing the rate of the hydrogen evolution reac-

tion. It is well known that in strongacids, which are

fully dissociated, the rate of hydrogen evolution

occurs according to

2H 2e ! H2 24

and is, for the case of mild steel corrosion, limited by

the rate at which H ions are transported from the

bulk solution to the steel surface (mass transfer limi-

tation). In CO2solutions, where typically pH > 4, thislimiting flux would be small, and therefore it is the

presence of H2CO3 which enables hydrogen evolu-

tion at a much higher rate. Thus, for pH > 4, thepresence of CO2 leads to a much higher corrosion

rate than would be found in a solution of astrongacidat the same pH.

This can be readily explained by considering that

the homogenous dissociation of H2CO3, as given by

reaction [8], serves as an additional source of H ions,

which are subsequently adsorbed at the steel surface

and reduced according to reaction [24].1 A different

pathway is also possible, where the H2CO3 first

adsorbs at the steel surface followed by heterogeneous

dissociation and reduction of the H

ion. This is

often referred to as direct reduction of carbonic

acid911 and is written as

2H2CO3 2e ! H2 2HCO

3 25

Clearly, the addition of the reactions [8] and [24] gives

the reaction [25] proving that the overall reaction isthe same and the distinction is only in the pathway,

that is, in the sequence of reactions. The rate of

reaction [25] is limited primarily by the slow hydra-

tion step [6]11,12 and in some cases by the slow CO2dissolution reaction [2].

It can be conceived that in CO2 solutions at

pH > 5 the direct reduction of the bicarbonate ionbecomes important13:

2HCO3 2e ! H2 2CO

23 26

which seems plausible, as the concentration of HCO3

increases with pH and can exceed that of H2CO3as seen in Figure 2. However, it is difficult to dis-

tinguish experimentally the effect of this particular

reaction pathway for hydrogen evolution from the

two previously discussed (eqns [8] and [25]). In

addition, evidence exists that suggests that the rate

of this reaction is comparatively low and can be

neglected. For example, as the pH increases, the

amount of HCO3 increases as well (see Figure 2),

suggesting that the corrosion rate should follow the

same trend, if one is to believe that the direct reduc-tion of the bicarbonate ion [26] is a significant

cathodic reaction. Experimental evidence does not

support this scenario and shows the opposite trend:

the corrosion rate actually decreases with an increas-

ing pH, even if no protective ferrous carbonate layer

forms.

Hydrogen evolution by direct reduction of water:

2H2O 2e ! H2 2OH

27

is always possible, but is comparatively very slow and

is important only atpCO2 0:1 bar and pH> 6.14,15

Therefore, this reaction is rarely a factor in practical

CO2corrosion situations.

The various electrochemical processes described

above can be quantified using the well established

electrochemical theory. The rate of the electrochem-

ical reactions, < in kmol m2 s1, can be readily exp-ressed in terms of current density, iin A m2, since

the two are directly related: for example, during

hydrogen evolution [24] for every kmol of H

1 kmol of electrons is used (n 1 kmole kmol1),

while for every kmol of iron dissolved [23] two



9/29

kmoles of electrons are used (n 2 kmole kmol1).

Therefore, one can write

i nF< 28

2.25.2.2.1 Oxidation of iron

In the corrosion of mild steel, the oxidation (dissolu-tion) of iron [23] is the dominant anodic reaction.

The anodic dissolution of iron at the corrosion

potential (and up to 200 mV above) is under charge

transfer control. Thus, pure Tafel behavior can be

assumed close to the corrosion potential:

ia Fe io Fe 10EcorrErev Fe =ba Fe 29

The exchange current density of iron oxidation is a

function of temperature:

io Fe irefo Fe

exp

DHFe

R

1

Tc 273:15

1

Tc;ref 273:15

30

The Tafel slope of this reaction is given by

baFe 2:303R Tc 273:15

1:5F 31

2.25.2.2.2 Reduction of hydronium ion

In general, the H

ion reduction reaction [24] can beeither under charge transfer or mass transfer (diffu-

sion) control, therefore, one can write:

1

ic H

1

ia H

1

idlim H

32

The charge transfer current density can be calculated

by

ia H io H 10EcorrErev H =bc H 33

The exchange current densityio H is a function of

pH and temperature. The pH dependence is

@log io H @pH

0:5 34

The temperature dependence of the exchange cur-

rent density can be calculated via an Arrhenius-type

relation:

io H

irefo H

exp

DHH

R

1

Tc 273:15

1

Tc;ref 273:15 35

The reversible potential for H+ reductionErev H is a

function of temperature and pH:

Erev H 2:303R Tc 273:15

F pH 36

The cathodic Tafel slope bc H is calculated as

bcH2:303R Tc 273:15

0:5F 37

The limiting mass transfer current density idlim H

is

related to the rate of transport of H+ ions from the

bulk of the solution through the boundary layer to

the steel surface:

idlim H kmHFcH 38

where the mass transfer coefficient, kmH can be

calculated from a correlation of the Sherwood, Rey-

nolds, and Schmidt numbers as explained in thefollowing section.

2.25.2.2.3 Reduction of carbonic acid

The carbonic acid reduction reaction [25] can be

under charge transfer control or limited by the

slow chemical reactionhydration step [6], preceding

it.11,12 The rate of this reaction in terms of current

density is

1

ic H2CO3

1

ia H2CO3

1

irlim H2CO3

39

The charge transfer current density ia H 2CO3 is cal-

culated as

ia H2CO3 io H2CO3 10EcorrErev H2CO3

=bc H2CO3 40

The exchange current density io H2CO3 depends on

pH, H2CO3concentration, and temperature:

@logio H2CO3 @pH

0:5 41

@logio H2CO3

@cH2CO3 1 42io H2CO3

irefo H2CO3 exp

DHH2CO3

R

1

Tc 273:15

1

Tc;ref 273:15

43

The cathodic Tafel slope bc H2CO3 is

bc H2CO3 2:303R Tc 273:15

0:5F 44

Since the reductions of H2CO3and H+ are equivalent



10/29

thermodynamically, the reversible potential for

H2CO3reductionErev H2CO3 is calculated as

Erev H2CO3 2:303R Tc 273:15

F pH 45

The chemical reaction limiting current densityirlim H2CO3 can be calculated from

16:

irlim H2CO3 FcCO2fH2CO3

ffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi ffiffiDH2CO3 Khydk

fhyd

q 46

The diffusion coefficient for carbonic acid DH2CO3 as

a function of temperature can be calculated using

Einsteins relation:

D DrefTc 273:15

Tc;ref 273:15 mH2O;ref

mH2O 47

where T is temperature and m is dynamic viscosity.

The forward reaction rate for the CO2 hydration

reactionkfhyd is calculated as

kfhyd 10169:253:0 log Tc273:15 11715=Tc273:15 48

The flow factorfH2CO3 is

fH2CO3 coth zH2CO3 49

where

zH2CO3 dmH2CO3

drH2CO350

and

dmH2CO3 DH2CO3kmH2CO3

51

drH2CO3

ffiffiffiffiffiffiffiffiffiffiffiffiffiffiDH2CO3

kbhyd

s 52

The carbonic acid mass transfer coefficient kmH2CO3is discussed inSection 2.25.2.3.

2.25.2.2.4 Reduction of water

Unless water is mixed with methanol or glycol to

prevent hydrate formation or somehow diluted oth-

erwise, it can be assumed that water molecules are

present in virtually unlimited quantities at the steel

surface, and the reduction rate of H2O is controlled

by the charge-transfer process and, hence, pure Tafel

behavior:

ic H2O io H2O 10Ecorr Erev H2O

=bc H2O 53

Since the reduction of H2O and H are equivalent

thermodynamically, they have the same reversible

potential at a given pH:

Erev H2O 2:303R Tc 273:15

F pH 54

The exchange current density for water reduction

io H2O depends on temperature:

io H2O

irefo H2O exp

DHH2O

R

1

Tc 273:15

1

Tc;ref 273:15 55The Tafel slope for H2O reduction was found to bethe same as that for H reduction:

bcH2O 2:303R Tc 273:15

0:5F 56

2.25.2.3 Transport Processes in CO2Corrosion of Mild Steel

From the description of the electrochemical processes

above, it is clear that certain species in the solution are

produced at the metal surface (e.g., Fe2+) while others

are depleted (e.g., H). The established concentration

gradients lead to molecular diffusion of the species

toward and away from the surface. In cases when the

diffusion processes are much faster than the electro-chemical processes, the concentration change at the

metal surface is small. In contrast, when the diffusion is

unable to keep up with the rate of the electrochemi-

cal reactions, the concentration of species at the metal

surface can become very different from that in the

bulk solution. The rate of the electrochemical pro-cesses depends on the concentration of the reactants

at the surface. Therefore, there exists a two-way cou-

pling between the electrochemical processes at the

metal surface (corrosion) and processes in the adjacent

solution layer (i.e., diffusion in the boundary layer).

The same is true for chemical reactions, which interact

with both the transport and electrochemical processes

in a complex way.

In most practical systems, the water solution

moves with respect to the metal surface. Therefore,

the effect of convection on transport processes cannot



11/29

be ignored. Turbulent eddies can penetrate deep into

the hydrodynamic boundary layer and significantly

alter the rate of species transport to and from the

surface. Very close to the surface no turbulence can

exist and the species are transported solely by diffu-

sion. The effect of turbulent flow is captured most

easily by using the concept of mass transfer coeffi-

cient, described below.

In turbulent flow of dilute ideal solutions, a mass

transfer coefficient km for a given species (H ions,

H2CO3etc.) can be calculated from a correlation, suchas the straight pipe correlation of Berger and Hau25:

Shp 0:0165Re0:86Sc0:33 57

or the rotating cylinder correlation of Eisenberg et al.26:

Shr 0:0791Re0:7Sc0:356 58

or anyother similar correlation for the flow geometryathand. It should be noted that most of the mass transfer

correlations found in the literature (including the two

listed above) are suited only for single-phase flow.

Therefore, extension of this approach to multiphase

flow situations needs to be done with careful

consideration.Overall, CO2 corrosion of mild steel is not very

sensitive to flow, at least not so when compared to

mild steel corrosion in strong acids. This is due to the

fact that the main corrosive species in CO2corrosion

is H2CO3which can easily be depleted due to a slowchemical step which precedes it: the hydration reac-

tion [6]. Therefore, the limiting rate of CO2corrosion

is primarily affected by the rate of this chemical

reaction [46], which is a function of temperature and

CO2partial pressure and not very sensitive to flow.

2.25.2.4 Calculation of Mild Steel CO2Corrosion Rate

Leading to this point, the main processes underpin-

ning CO2

corrosion were defined: the speciation of

the aqueous CO2solution using the thermodynamic

approach outlined in Section 2.25.2.1, the electro-

chemical theory described in Section 2.25.2.2, and

the transport processes as covered in Section 2.25.2.3.

Using this information, the corrosion rate of mild steel

can now be calculated. The unknown corrosion

potential Ecorr in [33], [40], [53], and [29] can be

found from the current (charge) balance equation at

the steel surface:

ic H ic H2CO3 ic H2O ia Fe 59

which expresses the simple fact that at steady state all

the electrons generated by the oxidation processes are

consumed by the sum of the reduction processes. By

substituting the expressions for the various currents

given byeqns [33], [40], [53], and [29]intoeqn [59]a

single nonlinear equation is now obtained withEcorras

the only unknown, which can be easily solved. When

the calculated value ofEcorr is now returned to eqns

[33], [40], [53], and [29], the rate of each individual

reaction can be explicitly computed. This also

includes the corrosion current density obtained fromeqn [29]:

icorr ia Fe 60

Finally, the CO2corrosion rate is recovered by using

Faradays law:

CR icorrMFerFenF

61

whereMis the molecular mass and ris the density. If

the unit amperes per square meters is used for the

corrosion current density icorr, then conveniently

the corrosion rate for iron and steel expressed in

millimeter per year takes almost the same numerical

value, precisely, CR 1:155icorr .

2.25.2.5 Successes and Limitations ofModeling of Aqueous CO2Corrosion of

Mild Steel

Evidence that our basic understanding of the pro-

cesses underlying CO2 corrosion of mild steel is

reasonably sound can be found by comparing the

predictions made by the mechanistic model outlined

above with experimental values. InFigure 4, below,

one can see the comparison of a potentiodynamic

sweep obtained in the experiments and the one pre-

dicted by the model. Many other comparisons of the

predicted and measured corrosion rates are given in

the following section, where the effect of key factors

in CO2corrosion of mild steel is discussed.

Despite the relative progress we have made in

understanding and modeling of aqueous CO2corro-

sion of mild steel, many questions persist. One is theissue of localized CO2corrosion, which is still a topic

of intense ongoing research. Effect of other factors

such as steel metallurgy, organic acids, oxygen, mul-

tiphase flow, and inhibitors are challenges that need

further effort. Some of those are discussed in the

following sections.



12/29

2.25.2.6 Key Factors Affecting Aqueous

CO2Corrosion of Mild Steel

Armed with the understanding and the ability to

calculate CO2 corrosion rates, as described in thesections above, in this section, the effect of key factors

which affect the rate of CO2corrosion are discussed,

and the predictions made by the model are compared

to empirical results.

2.25.2.6.1 Effect of pH

The pH has a significant influence on the CO2corro-

sion rate. Lower pH leads to higher corrosion rates and

vice versa, just like in many other acidic solutions.

Typical pH in CO2 saturated condensed water is

about pH 4 while in buffered brines, one frequentlyencounters 5< pH< 7. At pH 4 or below, direct reduc-tion ofH ions, reaction [24], is important, particularly

at lower partial pressures of CO2, when direct reduc-

tion of carbonic acid, reaction [25], can be ignored. In

that case, the pH has a direct effect on the corrosion

rate. Another important effect of pH is indirect and

relates to how pH changes conditions for the formation

of ferrous carbonate layers. Higher pH (5 < pH< 7)results in a decreased solubility of ferrous carbonate

and leads to an increased precipitation rate and a

higher scaling tendency. The effect of various pH and

supersaturations are shown in Figure 5. At lower

supersaturations obtained at the lower pH of 6, shown

inFigure 5, the corrosion rate does not change much

with time, even if some ferrous carbonate precipitationoccurs, reflecting the fact that a relatively porous,

detached and unprotective layer is formed (low scaling

tendency ST). The higher pH of 6.6 results in higher

supersaturation, faster precipitation, and formation ofmore protective ferrous carbonate, reflected by a rapid

decrease of the corrosion rate with time. There are

other indirect effects of pH, and by almost all accounts,

higher pH leads to a reduction of the corrosion rate,

making the pH stabilization (meaning: pH increase)

technique an attractive way of managing CO2 corro-

sion. The drawback of this technique is that it can lead

to excessive scaling and can rarely be used with forma-tion water systems.

2.25.2.6.2 Effect of CO2partial pressure

In the case of scale-free CO2corrosion, an increase of

pCO2 typically leads to an increase in the corrosion

rate. The commonly accepted explanation is that

with increasing pCO2 the concentration of H2CO3increases and accelerates the cathodic reaction, eqn

[25], and ultimately the corrosion rate. The detri-

mental effect ofpCO2 at a constant pH is illustrated in

Figure 6. The model described above reasonably

1

0.9

0.8

0.7

0.6

0.5

0.4

0.3

0.2

0.1

0

1010.1

i(A m2)

E

vs.

SHE(V)

H+reductionH2CO3reduction

Total cathodic

Total anodic

(Fe dissolution)

H2O reduction

Model sweep

Experimental

sweep

icorr

Ecorr

Figure 4 Potentiodynamic sweep, experimental (points) vs. model (lines); 20 C,pCO2 1bar, pH 4, 2ms1.



13/29

captures well this trend up to approximately

pCO2 10 bar. However, when other conditions arefavorable for the formation of ferrous carbonate

layers, increased pCO2 can have a beneficial effect.

At a high pH, higher pCO2 leads to an increase in

bicarbonate and carbonate ion concentration and a

higher supersaturation which accelerates precipita-

tion and protective layer formation. The effect of

pCO2 on the corrosion rate in the presence of ferrous

carbonate precipitation is illustrated in Figure 7

where in stratified wet gas flow, corrosion rate is

reduced both at top and bottom of the pipe with the

increase partial pressure of CO2.

2.25.2.6.3 Effect of temperature

Temperature accelerates all the processes involved incorrosion: electrochemical, chemical, transport, etc.

One would expect then that the corrosion rate

steadily increases with temperature, and this is the

case at low pH when precipitation of ferrous carbon-ate or other protective layers does not occur. An

example is shown Figure 8. The situation changes

markedly when solubility of ferrous carbonate is

exceeded, typically at a higher pH. In that case,

increased temperature rapidly accelerates the kinet-

ics of precipitation and protective layer formation,

decreasing the corrosion rate. The peak in the

3.00

2.50

2.00

1.50

1.00

0.50

0.00

Corrosionrate(mmy

ear

1)

SS=150

SS=9

SS=7

SS=37SS=30

0 5 10 15 20 25 30 35 40 45 50 55 60 65 70

Time (h)

Figure 5 Effect of ferrous carbonate supersaturation SSFeCO3 on corrosion rate obtained at a range of pH 6.06.6, for

5ppm


14/29

corrosion rate is usually seen between 60 and 80 C

depending on water chemistry and flow conditions as

shown inFigure 8(dotted line).

2.25.2.6.4 Effect of flow

There are two main ways in which flow may affectCO2corrosion, which can be distinguished based on

whether or not other conditions are conducive to

protective layer formation or not.

In the case of corrosion where protective layers do

not form (typically at low pH as found in condensed

water and in the absence of inhibitors), the main role

of turbulent flow is to enhance transport of species

toward and away from the metal surface. This may

lead to an increase in the corrosion rate as illustrated

inFigure 9. At lower pH 4, the effect is much more

pronounced as the dominant cathodic reaction is

direct H ion reduction [24], which is under mass

transfer control (seeeqn [38]).

When protective ferrous carbonate layers form

(typically at higher pH in produced water) or when

inhibitor films are present on the steel surface, theabove-mentioned effect of flow becomes insignificant

as the main resistance to corrosion is now in the surface

layer or inhibitor film. In this case, the effect of flow is

to interfere with the formation of protective surface

layers or to remove them once they are in place, often

leading to an increased risk of localized attack.

The two flow accelerated corrosion effects dis-

cussed above are frequently aggravated by flow dis-

turbances such as valves, constrictions, expansions,

bends, etc. where local increases of near-wall turbu-

lence and wall-shear stress are seen. However, flow

can lead to onset of localized attack only when giventhe right set of circumstances as discussed in a

separate heading below.

The effect of multiphase flow on CO2corrosion is

complicated by the different flow patterns that exist,

the most common being stratified, slug, and annular-

mistflow. In the liquid phase, water and oil can flow

separated or mixed with either phase being continu-

ous with the other flowing as a dispersed phase.

Different flow patterns lead to a variety of steel

surface wetting mechanisms: stable water wetting,

stable oil wetting, intermittent wetting, etc., which

15

0.2 0.2

0.06

BottomTop

100

10

1

0.1

0.01

Corrosionrate(mm

year

1)

P= 3.8 bar P=10.6barCO2partial pressure

Figure 7 Experimental measurements of the corrosion rate at the top and bottom of the pipe in stratified gasliquidflow showing the effect of CO2partial pressure,pCO2 on formation of ferrouscarbonate layer. Test conditions: 90

C, pH 6,

100mm ID,Vsg 1 0 m s1,Vsl 0.1ms

1. Data taken from Sun and Nesic.18

25

20

15

10

5

00 20 40 60 80 100 120

Corrosionrate(mmy

ear1)

Temperature (C)

Figure 8 The effect of temperature on CO2corrosion rate

of mild steel; pH 4, pCO2 1 bar, 100 mm ID single phasepipe flow. Points are experimental values and the solid line

is the model. The dotted line is a model simulation of the

same conditions at pH 6.6 accounting for protective ferrous

carbonate film formation.



15/29

greatly affect corrosion. In annular mist flow, the

liquid droplets move at high velocity and can lead

to protective layer damage at points of impact such as

bends, valves, tees, constrictions/expansions, and

other pipe fittings. Slug flow can lead to significant

short-lived fluctuations in the wall-shear stress,

which can help remove a protective surface layer of

ferrous carbonate or possibly affect an inhibitor film.

2.25.2.6.5 Effect of corrosion inhibition

The two most common sources of corrosion inhibi-

tion need to be considered:

(a) inhibition by addition of corrosion inhibitors and

(b) inhibition by components present in the crude oil.

Corrosion inhibitors

Describing the effect of corrosion inhibitors is not a

straightforward task due to the enormous complexity

of the subject. Quantifying them and predicting their

behavior are even harder. There is a plethora of

approaches in the open literature, varying from the

use of simple inhibitor factorsand inhibition efficiencies

to the application of complicated molecular modelingtechniques to describe inhibitor interactions with the

steel surface and ferrous carbonate layer. A middle-

of-the-road approach is based on the assumption that

corrosion protection is achieved by surface coverage,

that is, that the inhibitor adsorbs onto the steel sur-

face and slows down one or more electrochemical

reactions by blocking. The degree of protection is

0

1

2

3

4

CR(m

my

ear

1)

pH=4

0

1

2

3

4

CR

(mmy

ear

1)

pH=5

0

1

2

3

4

0 2 4 6 8 10 12 14

Velocity (m s1)

CR(mmy

ear

1)

pH=6

Figure 9 Predicted (line) and experimentally measured corrosion rates (points) showing the effect of velocity in the

absence of ferrous carbonate layers. Test conditions: 20 C,pCO2 1 bar, 15 mm ID single-phase pipe flow. Experimentaldata taken from Nesicet al.19



16/29

assumed to be directly proportional to the fraction of

the steel surface blocked by the inhibitor. In this type

of model, one needs to establish a relationship

between the surface coverage y and the inhibitor

concentration in the solutioncinh. This is most com-

monly done by the use of adsorption isotherms.

Corrosion inhibition by crude oil

It has been known for a while that CO2 corrosion

rates seen in the field in the presence of crude oil are

much lower than those obtained in laboratory condi-

tions where crude oil was not used or synthetic crude

oil was used. One can identify two main effects of

crude oil on the CO2corrosion rate.

The first is a wettability effect and relates to a

hydrodynamic condition where crude oil entrains

the water and prevents it from wetting the steelsurface (continuously or intermittently).

The second effect is corrosion inhibitionby compo-

nents of the crude oil that reach the steel surface

either by direct contact or by first partitioning into

the water phase. Various surface active organic com-pounds found in crude oil (typically oxygen, sulfur

and nitrogen containing molecules) have been iden-

tified to directly inhibit corrosion of mild steel in

CO2solutions.

2.25.2.6.6 Effect of organic acidsThe low molecular weight organic acids are primarily

soluble in water and can lead to corrosion of mild

steel. Higher molecular weight organic acids are not

water soluble, but are typically soluble in the oil phase

and pose a corrosion threat at higher temperatures in

the refineries. Acetic acid CH3COOH (denoted as

HAc in the text below) is the most prevalent low

molecular weight organic acid found in brines.

Other acids typically found in the brine are propionic,

formic, etc.; however, their behavior and corrosiveness

is very similar to that of HAc and therefore HAc can

be used as a surrogate for all the organic acids found

in the brine. HAc is a weak acid; however, it is stronger

than H2CO3(pKa4.76 vs. 6.35 at 25C), and it is the

main source of H ions when the two acid concentra-

tions are similar. The effect of HAc is particularly

pronounced at higher temperatures and low pH

when the abundance of undissociated HAc can

increase the CO2corrosion rate dramatically as seen

in Figure 10. Solid iron acetate does not precipitate in

the pH range of interest since its solubility is much

higher than that of ferrous carbonate. There are some

indications that the presence of organic acids impairs

the protectiveness of ferrous carbonate layers; how-

ever, the mechanism is still not clear.

2.25.2.6.7 Effect of glycol/methanol

Glycol and methanol are often added to flowing

systems in order to prevent hydrates from forming.

The quantities are often significant (50% of total

liquid phase is not unusual). In the very few studies

available, it has been assumed that the main inhibi-

tive effect of glycol/methanol on corrosion comes

from dilution of the water phase, which leads to

a decreased activity of water. However, there are

many unanswered questions such as the changes inmechanisms of CO2 corrosion in water/glycol

mixtures which are yet to be discovered.

2.25.2.6.8 Effect of condensation in

wet gas flow

When transporting humid natural gas, due to the cool-

ing of the stream, condensation of water vapor occurs

on the internal pipe wall. The condensed water is pure

and, due to dissolved CO2, typically has a pH < 4.This leads to the so-called top-of-the-line corrosion(TLC) scenario. If the rate of condensation is high,

plenty of acidic water flows down the internal pipewalls leading to a very corrosive situation. If the con-

densation rate is low, the water film is not renewed and

flows down very slowly and the corrosion process can

release enough Fe2+ to raise the local pH and saturate

the solution, leading to the formation of protective

ferrous carbonate layer. The layer is often protective;

however, incidents of localized attack in TLC were

reported.21 Either way, the stratified or stratified-wavy

flow regime, typical for TLC, does not lead to a good

opportunity for inhibitors to reach the upper portion

of the internal pipe wall and protect it. A very limited

0

10

20

30

40

50

60

1000100101Undissociated aqueous HAc concentration (ppm)

C

R(mmy

ear

1)

Figure 10 Predicted (line) and experimentally measured

data (points) showing the effect of the concentration of

undissociated acetic acid (HAc) on the CO2corrosion rate,

60 C,pCO2 0.8 bar, pH 4, 12 mm OD rotating cylinderflowat 1000 rpm. Experimental data taken from Sunet al.20



17/29

range of corrosion management options for TLC

exists. To qualitatively and quantitatively describe the

phenomenon of corrosion occurring at the top of

the line, a deep insight into the combined effect

of the chemistry, hydrodynamics, thermodynamics,

and heat and mass transfer in the condensed water is

needed. A full description exceeds the scope of this

review, and the interested reader is directed to some

recent articles published on this topic.21,22

2.25.2.6.9 Nonideal solutions and gases

In many cases produced, water has very high dissolved

solids content (>10 wt%). At such high concentra-tions, the infinite dilution theory used above does not

hold, and corrections need to be made to account for

solution nonideality. A simple way to account for the

effect on nonideal homogenous water chemistry is

to correct the equilibrium constants by using the con-cept of ionic strength as indicated above. This

approach seems to work well only for moderately

concentrated solution (up to a few weight percentage

of dissolved solids). For more concentrated solutions, a

more accurate way is to use activity coefficients as

described by Anderko et al.23 The effect of concen-

trated solutions on heterogeneous reactions such as

precipitation of ferrous carbonate and other layers is

still largely unknown. Furthermore, it is unclear howthe highly concentrated solutions affect surface elec-

trochemistry. Some experience suggests that corrosionrates can be dramatically reduced in very concentrated

brines; nevertheless a more systematic study is needed.

At very high total pressure, the gasliquid equili-

bria cannot be accounted for by Henrys law. A simple

correction can be made by using a fugacity coeffi-

cient, which accounts for nonideality of the CO2/

natural gas mixture24 and can be obtained by solving

the equation of state for the gas mixture. Those cases,

in which critical point for CO2 is approached or

exceeded, warrant a separate analysis and are not

covered by the considerations discussed above.

2.25.2.7 Localized CO2Corrosion of Mild

Steel in Aqueous Solutions

As illustrated above, significant progress has been

achieved in understanding uniform CO2 corrosion,

without or with protective layers, and hence success-

ful uniform corrosion models can be built. However,

much less is known about localized CO2corrosion. It

is thought that one of the main factors that triggers

localized attack is flow, tempered by other environ-

mental variables such as pH, temperature, partial

pressure of CO2, etc. It seems that localized attack

occurs when the conditions are such that partially

protective ferrous carbonate layers form. It is well

known that when fully protective ferrous carbonate

forms, low general corrosion rates are obtained and

vice versa: when no protective layers form, a high rate

of general corrosion is seen. It is when the corrosive

environment is in between, in the so-called gray

zone, that localized attack can be initiated most

often by some extreme flow conditions. There are

many combinations of environmental and metallur-gical parameters that define the grey zone, making

this sound like a difficult proposal. However, there is

a single parameter which is easy to calculate: ferrous

carbonate supersaturation, SSFeCO3 (see eqn [19]

above), which can be successfully used as a good

delineator for the gray zone and as such as a predictor

for the probability for localized attack. When bulkferrous carbonate supersaturation is in the range

0.5 < SSFeCO3


18/29

Therefore, the mechanism of H2S corrosion remains

much less understood when compared to that of CO2corrosion. This uncertainty makes it more difficult to

develop a model to predict the corrosion rate of mild

steel in H2S saturated aqueous solution.

2.25.3.1 Chemistry of H2S Saturated

Aqueous Solutions Equilibrium

Considerations

Similar to CO2 discussed above, H2S gas is also

soluble in water:

H2S g ,KH2 S

H2S 62

where KH2S is the solubility constant of H2S in

mol l1 bar1:

KsolH2S cH2S

pH2S63

and can be found from34

KsolH2S 10634:270:2709TK0:1113210

3T2K16719=TK261:9logTK 64

As shown inFigure 11, the solubility of H2S decreases

with temperature, as it is observed for CO2. However,

for the same partial pressure and temperature, theconcentration of dissolved H2S actually exceeds that

in the gas phase as shown inFigure 12.

Aqueous H2S is another weak acid which partly

dissociates in two steps:

H2S ,Khs

H HS 65

HS ,Kbs

H S2 66

whereKhs is the dissociation constant of H2S:

Khs cH cHS

cH2S67

and can be calculated as35

Khs 10782:439450:361261TK1:672210

4T2K20565:7315=TK142:741722lnTK 68

andKbs is the dissociation constant of HS:

0.00

0.05

0.10

0.15

0.20

0 20 40 60 80 100T(C)

Speciesconcentration(moll1)

H2S

CO2

Figure 11 Calculated solubility of H2S and CO2as afunction of temperature; 25 C,pH2 S 1 bar,pCO2 1 bar.

1.E07

1.E06

1.E05

1.E04

1.E03

1.E02

1.E01

1.E+00

2 3 4 5 6 7pH

Speciesconcentra

tion(moll1)

HS

H2S

H2S(g)

Figure 12 Calculated sulfide species concentrations as a function of pH for an H2S saturated aqueous solution at

pH2 S 1 mbar, 25 C, 1 wt% NaCl.



19/29

Kbs cH cS2

cHS69

There is a very large discrepancy in the reported

values for Kbs, varying from 1:0 1019 to

1:1 1012

kmol m3

at room temperature (sevenorders of magnitude). In addition, these values are

very small compared with other equilibrium con-

stants, all suggesting that using Kbs to calculate the

concentration of sulfide species, cS2 and further to

predict the solubility product constants for ferrous

sulfides should be avoided.

Given the same gaseous concentrations of

H2S and CO2, one obtains a similar aqueous concen-

tration of dissolved H2S and CO2(seeFigure 11) and

the resulting pH is within 0.1 pH unit, therefore,

values shown in Figure 1 for CO2 can be used for

H2S as the first approximation. The equilibrium dis-

tribution of sulfide species as a function of pH for an

open system is shown inFigure 12. The concentra-

tion of bisulfide ion, cHS , becomes significant only

above pH 4, while the concentration of the sulfideion, cS2 , is not even shown as it is very low and

unreliable to calculate.

Many types of iron sulfides, such as amorphous

ferrous sulfide (FeS), mackinawite (Fe1xS), cubic

ferrous sulfide (FeS), troilite (FeS), pyrrhotite

(Fe1xS or FeS1x), smythite (Fe3xS4), greigite

(Fe3S4), and pyrite (FeS2) occur. Studies have sug-gested that some of these are stoichiometric such

as cubic ferrous sulfide, troilite, greigite, and

pyrite, while others such as mackinawite, pyrrho-

tite, and smythite are not. Some are electrically

nonconductive, others apparently behave as semi-

conductors. However, there is no consensus on

these issues and the interested reader is directed

to the vast literature on iron sulfides for a more

in-depth treatment. The thermodynamics of thesesystems is very complicated; depending on envi-

ronmental conditions and time, transformationfrom one type of ferrous sulfide into the other

occurs. Limited information exists on aqueous

solubility of the various sulfides. Avoiding the

usage of the sulfide ion concentration, cS2 , one

can write a general equation for precipitation of

ferrous sulfide as

Fe2 H2S ,KspFeS

FeSs 2H 70

where the solubility constant for one type of

ferrous sulfide mackinawite is known as a func-

tion of temperature36,37

KmackinspFeS 102848:779=Tk6:347 71

For other ferrous sulfides, only the values at room

temperature are known, as listed in Table 1 below.It is convenient to show various ferrous sulfide solu-

bilities in terms of an equilibrium concentration of

the Fe2+ as a function of pH at a given H2S partial

pressure (concentration). An example is presented in

Figure 13 where it can be seen that the much less

soluble pyrrhotite and troilite are thermodynami-

cally more stable forms compared to mackinawite

and amorphous ferrous sulfide. For a typical ferrous

ion concentration of cFe2 1 ppm, the saturationwith respect to troilite and pyrrhotite is reached

already at pH 5.4, while for mackinawite it is pH 6and for amorphous ferrous sulfide pH 6.7. Keeping in

mind that the concentration of Fe2 at a corroding

steel surface can easily be much higher than in the

bulk (e.g., 10 ppm or even higher) and that the pH is

also higher at the surface than in the bulk (typicallyabove pH 6), usingFigure 13one can expect a whole

range of different ferrous sulfides to form on a cor-

roding steel surface at this H2S concentration at

different points in time.

SEM images of a ferrous sulfide surface layer

formed on mild steel after a week long exposure areshown in Figure 14. The layered structure of the

sulfide is prominent, and it can be identified as mack-

inawite. In longer exposures, the ferrous sulfide layer

thickens and eventually becomes more protective. An

image of a ferrous sulfide layer after a month long

exposure is shown inFigure 15. The composition of

the layer is a mixture of mackinawite and pyrrhotite.

Another layered structure composed of a mixture of

ferrous carbonate and ferrous sulfide is shown in

Figure 16.

Table 1 Solubility product constants for various ferrous

sulfides at 25 C38

Type of ferrous sulfide log Ksp(FeS)

Amorphous (FeS) 2.95

Mackinawite (Fe1xS) 3.6

Pyrrhotite (Fe1xS or FeS1x) 5.19

Troilite (FeS) 5.31



20/29

2.25.3.2 Mild Steel Corrosion in H2S andMixed H2S/CO2Saturated Aqueous

Solutions

As aqueous H2S is another weak acid, it can be seen

as an additional reservoir of H ions according to

reaction [65], similar to H2CO3. Therefore, stimula-

tion of the hydrogen evolution reaction could also be

expected in the presence of H2S. Using the analogywith CO2 corrosion, one must also allow the possi-

bility of direct reduction of H2S, that is, that the

H2S molecule can be adsorbed at the steel surface,

1.E+00

1.E+01

1.E+02

1.E+03

1.E+04

3

pH

Fe

2+concentration(ppm)

Mackinaw

ite

Pyrrhotite AmorphorusTroilite

3.5 4 4.5 5 5.5 6 6.5 7

Figure 13 Calculated solubility of various iron sulfides as a function of pH shown in terms of the equilibrium concentration

of Fe2

,pH2 S 1 mbar, 25

C, 1 wt% NaCl.

Mild steel

Ferrous sulfide

Acc.V

20.0kV

Spot

5.0

Magn

100x

Det

SE

WD

10.3

200 m

Acc.V

20.0kV

Spot

5.0

Magn

100x

Det

SE

WD

10.3

200m

Figure 14 SEM images showing a cross-section

and a top view of a ferrous sulfide layer formed on mild

steel; 60 C, pH 6,pCO2 7.7bar,pH2S 0.25mbar, 1ms1

single phase flow in a 100 mm ID pipe, 7 days exposure.

Mild steel

Ferrous sulfide

Acc.V

15.0 kV

Spot

3.0

Magn

250x

WD

11.9100m

Figure 15 SEM images showing a cross-section view of a

ferrous sulfide layer formed on mild steel; 60 C, pH 6,

pCO2 7.7 bar,pH2S 0.25mbar, 1 m s1 single phase flow

in a 100 mm ID pipe, 30 day exposure.



21/29

followed by a reduction of the H

and oxidation ofiron in the steel. One can write the overall corrosion

reaction as

Fes H2S ! FeSs H2 72

As solid ferrous sulfide (mackinawite) is always found

on the corroding steel surface in the presence of H2S,

even below the solubility limit, this can been referred

to as a direct solid state reaction pathway as both the

initial and final state of Fe are solid(s).39

Experimental evidence suggests that corrosion ofmild steel by H2S initially proceeds by adsorption of

H2S to the steel surface followed by a very fast redox

reaction at the steel surface to form an adherent

mackinawite film (much like a tarnish). This initial

mackinawite film is very thin (1 mm) but appar-ently rather dense and acts as a solid state diffusion

barrier for the species involved in the corrosion reac-

tion. Therefore, this thin mackinawite film is one of

the most important factors governing the corrosionrate in H2S corrosion. It also impedes the mobility of

other species in reaching the steel surface and there-

fore corrosion rates due to CO2 are affected even if

very small amounts of H2S are present in the gas

phase (as little as 105 bar).

The thin mackinawite film continuously goes

through a cyclic process of growth, internal stress

growth, cracking, and delamination that generates

an outer sulfide layer, which thickens over time

(typically 1mm) and forms an additional diffusionbarrier. However, this outer sulfide layer is very

porous and rather loosely attached to the steel sur-

face. Over time it cracks, peels, and spalls, a

process accelerated by turbulent flow. If the pH

of the solution is below saturation level, the outer

sulfide layer will undergo a process of chemical

dissolution. Conversely, when the saturation is

exceeded, ferrous sulfide precipitation from the

bulk is possible. Eventually, the amount and protec-

tiveness of the outer sulfide layer is determined by

the balance of the various formation and removal

processes.39

The transformation of mackinawite into other

forms of less soluble and more stable ferrous sulfide

(pyrrhotite and troilite, see Figure 13) may happenover time. Among the various ferrous sulfides, mack-

inawite is the prevalent ferrous sulfide that forms

in the corrosion of mild steel at low H2S concentra-

tion and low temperature. At increased levels of

H2S, mackinawite is less prevalent and pyrrhotite

is the main corrosion product. At very high H2S con-

centrations, pyrite and elemental sulfur appear. Whilethermodynamics of ferrous sulfides may favor other

types of sulfide over mackinawite as the corrosion

product, the rapid kinetics of mackinawite formation

favors it as the initial corrosion product seen in most

situations. Overall, however, there is currently no

clearly defined relationship between the nature of

the sulfide layer and the underlying corrosion process.

It is generally thought that all types of ferrous sulfide

layers offer some degree of corrosion protection formild steel.

At very high H2S concentrations, elemental sulfurcan appear and lead to severe localized corrosion.

Large amounts of elemental sulfur can precipitate

out of the gas stream and can even block the line,

due to the changes in pressure and temperature.

Alternatively, when there is O2 present, the most

likely pathways for formation of elemental sulfur

are as follows:

ferrous sulfide reacts with O2and converts to ironoxide forming elemental sulfur probably via:

3FeS 2O2 ! Fe3O4 3S 73

at very high H2S concentration, the followingreaction can occur to yield elemental sulfur:

2H2S O2 ! 2H2O 2S 74

At very high temperatures, an alternative pathway is

H2S ! H2 S 75

Localized corrosion by elemental sulfur occurs via a

reaction with the iron in the steel, represented by the

Mild steel

Ferrous sulfide

Ferrous carbonate

Figure 16 SEM images showing a cross-section view of a

mixed ferrous carbonate and ferrous sulfide layer formed on

mild steel; 60 C, pH 6,pCO2 7.7 bar,pH2S 1.2 mbar,1 m s1 single phase flow in a 100 mm ID pipe, 25 day

exposure.



22/29

overall reaction

Fe S ! FeS 76

It is not very clear at this stage what the detailed

mechanism of this reaction is. It appears that rapid

attack is seen only when direct contact of sulfur with

the steel is achieved in the presence of water. A more

in-depth discussion about the corrosion mechanisms

of mild steel involving elemental sulfur exceeds the

scope of this review.

2.25.3.3 Calculation of Mild Steel

H2S Corrosion Rate

Due to the complexity of the underlying processes

and a lack of mechanistic understanding, predictive

models of H2S corrosion were not readily availableuntil recently. One approach40 which has the capabil-

ity to address a few simple H2S corrosion scenarios is

presented below. A pure H2S corrosion environment

is described first followed by a mixed H2S/CO2corrosion scenario.

2.25.3.3.1 Pure H2S aqueous environment

Due to the presence of the inner mackinawite film

and the outer porous sulfide layer, it is assumed thatthe corrosion rate of steel in H2S solutions is always

under mass transfer control. One can then write theflux of H2S due to:

convective diffusion through the mass transferboundary layer as

FluxH2S kmH2S cH2S coH2S

77

molecular diffusion through the liquid in theporous outer sulfide layer as

FluxH2S DH2Sec

doscoH2S ciH2S 78

solid state diffusion through the inner mackinawitefilm as

FluxH2S AH2S exp BH2S

RTk

ln

ciH2S

csH2S

! 79

In a steady state, the three fluxes are equal to each

other and are equivalent to the corrosion rate as

CRH2S FluxH2SMFe=rFe 80

further corrected for appropriate corrosion rate unit.

By eliminating the unknown interfacial concen-

trationscoH2S andciH2S fromeqns [77] to [79], the

following equation is obtained for the flux (corrosion

rate) due to H2S:

FluxH2S AH2SlncH2S FluxH2S

dos

DH2Sec 1

kmH2S !csH2S

81

This is an algebraic nonlinear equation with respect

to FluxH2S, which does not have an explicit solution

but can be solved by using a simple numerical algo-

rithm such as the interval halving method or similar

methods. These are available as ready-made routines

in spreadsheet applications or in any common com-

puter programming language. The prediction forFluxH2S depends on a number of constants used in

the model which can be either found in handbooks

(such asDH2S), calculated from the established theory

(e.g., kmH2S) or are determined from experiments

(e.g., AH2S; csH2S). The unknown thickness of theouter sulfide layer change with time and need to be

calculated as described below.

It is assumed that the amount of layer retained on

the metal surface at any point in time depends on the

balance of:

layer formation kinetics (as the layer is generated byspalling of the thin mackinawite film underneath it

and by the precipitation from the solution), and layer damage kinetics (as the layer is damaged by

intrinsic or hydrodynamic stresses and/or by

chemical dissolution):

gSRR

Sulfide layerretension rate

gSFR

Sulfide layerformation rate

gSDR

Sulfide layerdamage rate

82

where all the terms are expressed in kmol m2 s1. In

order to simplify the calculations, it can be assumed

that in the typical range of application (4 < pH < 7),precipitation and dissolution of ferrous sulfide layer

do not play a significant role and so it can be written

SRR CR SDRm 83

Some experiments involving mackinawite have shown

that even in stagnant conditions about half of the outer

sulfide layer that forms is lost from the steel surface

due to intrinsic growth stresses by internal cracking

and spalling, that is, SDRm 0:5CR, so one obtains:

SRR 0:5CR 84



23/29

that is, about half of the iron corroded is found on the

steel surface in the form of mackinawite. It is not

known if and how this ratio is different when other

types of ferrous sulfide layers form, for example, the

more adherent and protective pyrrhotite. Moreover,

additional experimentation is required to determine

how the mechanical layer damage is affected by hydro-

dynamic forces.

Once the layer retention rate SRR is known, the

change in mass of the outer sulfide layer can be easily

calculated as

Dmos SRRMFeSADt 85

The porosity of the outer sulfide layer was deter-

mined to be very high (e 0:9) by comparing theweight of the layer with the cross-sectional SEM

images showing its thickness. On the other hand,

this layer has proven to be rather protective (i.e.,

impermeable to diffusion) which can only be

explained by its low tortuosity arising from its lay-

ered structure. By comparing the measured and

calculated corrosion rates in the presence of the

outer sulfide layer, the tortuosity factor was calcu-

lated to be c 0:003.A time-marching explicit solution procedure

could now be established where

1. the corrosion rate FluxH2S in the absence of outer

sulfide layer can be calculated by using eqn [81],and assumingdos 0;

2. the amount of sulfide layer dmos formed over a

time interval Dtis calculated by usingeqn [85];

3. the new corrosion rate FluxH2S in the presence of

sulfide layer can be recalculated by usingeqn [81];

4. a new time interval Dt is set and steps 2 and 3

repeated.

At very low H2S gas concentrations (ppmwrange),

there is very little dissolved H2S and the corrosion

rate is directly affected by pH. A mackinawite layer

still forms and controls the corrosion rate; however,the corrosion process is largely driven by the reduc-

tion of H ions, rather than of H2S. By analogy with

the approach laid out above, the following expression

is obtained for the flux of H ions controlled by the

presence of the ferrous sulfide layers:

FluxH AH ln

cH FluxH

dos

DHec

1

kmH

csH

86

The flux FluxH is directly related to the corrosion

rate by H ions:CRH

FluxH

2

MFe

rFe87

further adjusted for the appropriate corrosion

rate unit.

By solvingeqns [81] and [86]sequentially in time,the total corrosion rate in mixed pure H2S aqueous

environments can be calculated as

CR CRH2S CRH 88

2.25.3.3.2 Mixed CO2/H2S environments

For mild steel corrosion in mixed CO2/

H2S containing environments, one can account for

the effect of CO2 by assuming that the rate

controlling step in this additional process is the dif-fusion of CO2 through the ferrous sulfide layers.

Then a similar expression can be obtained for the

corrosion rate due to CO2:

FluxCO2 ACO2 ln

cCO2 FluxCO2

dos

DCO2ec

1

kmCO2

csCO2

89

The flux FluxCO2 is equivalent to the corrosion rate

by CO2:

CRCO2 FluxCO2

2

MFe

rFe90

further adjusted for appropriate corrosion rate unit.

By solving eqns [81], [86], and [89], the total

corrosion rate in mixed CO2/H2S environments can

be calculated as

CR CRH2S CRH CRCO2 91

2.25.3.4 Limitations of Modeling ofAqueous H2S Corrosion of Mild Steel

The calculation model presented above covers

uniform H2S and CO2/H2S corrosion. There are

numerous limitations:

It does not predict localized corrosion in eitherenvironment.

While it covers a very broad range of H2S partialpressures, it is not recommended to use this model

below pH2S 0.01 mbar or above pH2S 10 bar.

Similar limits apply to the CO2 partial pressure.



24/29

This leaves a very broad area of applicability for

the present model. This H2S model does not account for any precipi-

tation of ferrous sulfide, ferrous carbonate, or any

other scale; therefore, in cases where this is

deemed important for corrosion, the model shouldbe used with caution. The model also does not

account for various transformations of sulfide

layer from one type to another which are known

to happen over time. The present model does not account for dissolu-

tion of the sulfide layer that may occur at very low

pH. Therefore, the use of this model at pH < 3 isnot recommended. Similarly, the model should be

used with caution for pH > 7 where it has not beentested.

The model in its present state does not cover the

effect of organic acids on mixed H2S and CO2/H2S corrosion, and therefore it should not be used

when organic acids are present in the system.

A practical threshold for the validity of the present

model is


25/29

total pressure (p 138 bar) and a high CO2 partialpressure (pCO2 13.8 bar). When comparing the pre-dictions with the experimental results, it can be seen

that the model underpredicts the observed rate

of steel corrosion by approximately a factor of 2.

However, when this is compared with a pure CO2(H2S-free) corrosion rate under the same conditions

(which is not reported but can be predicted to be

almost 20 mm year1), the accuracy of the model can

be considered as reasonable. At the highestpCO2 /pH2S

ratio of 3500 (pCO2 13.8 bar, pH2S 40 mbar), CO2accounts for 70% of the corrosion rate and 30%can be ascribed to H2S. At the lowestpCO2 /pH2S ratio

of 1180 (pCO2 13.8 bar, pH2S 116 mbar), CO2accounts for 57% of the corrosion rate and 43%can be ascribed to H2S.

Corrosion rates of mild steel at very high partial

pressures of H2S (pH2S 320 bar) and CO2(pCO2 312.8 bar) for exposures lasting up to 4 daysare shown inFigure 19.This is a situation where the

0

5

10

Corrosionrate(mm

year1)

15

20

0

H2S partial pressure (mbar)

0

H2S gas concentration (ppmm)

Mod.

Exp.Pure CO2corrosion rate

100 200 300 400 500 600 700

20 40 60 80 100 110 120

Figure 18 The corrosion rate vs. H2S partial pressure; experimental data (exp.) shown as points, model predictions

(mod.) shown as lines; conditions: total pressure p 137.9 bar,pCO2

13.8 bar,pH2

S 40120 mbar, T 50C, experiment

duration 3 days, pH 4.06.2, stagnant. Experimental data taken from Smith and Pachecoet al.31

0

5

10

15

20

25

30

35

40

0 20 40 60 80 100

Time (h)

Corrosionrate(mmy

ear1)

Test A and B mod.

Test C mod.

Test D mod.

Test E mod.

Test F mod.

Test A and B exp.

Test C exp.

Test D exp.

Test E exp.

Test F exp.

Figure 19 The corrosion rate vs. time; experimental data (exp.) shown as points, model predictions (mod.) shown as lines;

Test A and B:p 8.3 bar,pCO2 5.3 bar,pH2S =3 bar,T 60C, 71 h (a) and 91 h (b); Test C:p 24 bar,pCO2 4bar,

pH2 S20 bar,T 70C, 91 h; Test D:p 15.7bar,pCO2 3.5 bar,pH2 S12.2bar,T 65

C, 69 h; Test E:p 20.8bar,pCO2 12.8bar,pH2S8 bar,T 65

C, 91 h; Test F:p 7.2 bar,pCO2 3bar,pH2S4.2 bar,T 65C, 63 h; experimental data

taken from Bich and Goerz.43



26/29

H2S was the dominant corrosive species. At the high-

estpCO2 /pH2Sratio of 1.8 (pCO2 5.3 bar,pH2S 3 bar),H2S generated 86% of the corrosion rate. At thelowest pCO2 /pH2S ratio of 0.2 (pCO2 4 bar, pH2S 20bar), H2S generated 97% of the overall corrosion rate.

It is also noted that the model predictions show that

the corrosion rate in the first reaction hour is on

average 20 mm year1 with an initial corrosion rate

of 60 mmyear1 and a final corrosion rate of 10 mm

year1. The pitting corrosion rate was reported to be

30 mm year1 in a field case with similar conditions,which is related to the very high, H2S-driven corro-

sion seen at the beginning of experiments before a

thick protective ferrous sulfide film forms.

2.25.3.5.2 Effect of flow

The effect of flow velocity in H2S corrosion is shownin Figure 20 for the three long-term experiments

reported by Omar et al.44 Flow loop experiments

lasting 1521 days were conducted at severe condi-

tions: high partial pressure of H2S (pH2S1030 bar),high partial pressure of CO2 (pCO2 3.310 bar) andlow pH 2.93.2. No effect of velocity on the uniform

corrosion rate could be observed in these long-term

exposures, which is due to the build-up of a thick

protective sulfide layer. The model predictions also

shown in Figure 20 confirm this trend and show a

remarkable agreement with the experimental results

in the less extreme experiments 1 and 2 (pCO2 3.3bar;pH2S 10 bar) both at low (25

C) and high tem-

perature (80 C). In experiment 3 which was con-

ducted at the most extreme set of conditions

(pCO2 10 bar; pH2S 30 bar) and high temperature(80 C) the model overpredicts the corrosion rate by

a factor of2.5. In all three experiments reported by

Omaret al.,44 thepCO2 /pH2Sratio was about 0.3, that is,

the corrosion process and corrosion rate were

completely dominated by H2S, which contributed

95% of the corrosion rate.

2.25.3.5.3 Effect of time

A marked decrease of corrosion rate with time was

seen in autoclave tests as reported in Figure 19above;

the same was observed in stratified pipe flow experi-

ments where pure CO2corrosion rate decreased withtime due to the presence of H2S, as shown in Figure 21below. The latter is also a mixed CO2/H2S corrosion

scenario. At a pCO2 /pH2S ratio of 200 (pCO2 2 bar,pH2S 4 mbar), the CO2 contribution to the corro-sion rate is 75% with most of the balance providedby H2S. At the pCO2 /pH2S ratio of 28 (pCO2 2 bar,

pH2S70mbar), both CO2 and H2S account for50% of the overall corrosion rate.

Corrosion experiments at high temperature

(120 C), high partial pressures of CO2 (pCO2 6.9bar), and H2S (pH2S 1.384.14 bar) in exposures

0.0

1.0

2.0

3.0

4.0

5.0

0 1 2 3 4 5 6

Corrosion

rate(mmy

ear

1)

Exp. 1

Exp. 2

Exp. 3

Mod. 1

Mod. 2

Mod. 3

Velocity (m s1)

Figure 20 The corrosion rate vs. velocity; experimental data (exp.) shown as points, model predictions (mod.) shown as

lines; exp 1.: 19 days,p 40 bar,pCO2 3.3 bar,pH2 S 10 bar,T 80C, pH 3.1,v 15ms1; exp 2.: 21 days,p 40bar,

pCO2 3.3 bar,pH2S10bar,T 25C, pH 3.2,v 15ms1; exp 3.: 10 days,p 40bar,pCO2 10 bar,pH2S 30 bar,

T 80 C, pH 2.9,v 15ms

1; experimental data taken from Omar et al.44



27/29

lasting up to 16days are shown in Figure 22.

A steadily decreasing corrosion rate was observed

due to build-up of a protective ferrous sulfide layer.

The effect ofpH2Sincrease on corrosion rate was very

small and practically vanished over time. Both these

effects were readily captured by the model with very

good accuracy as seen in Figure 22. In this case, the

H2S is the dominant corrosive species. At the highest

pCO2 /pH2S ratio of 5 (pCO2 6.9 bar, pH2S 1.38 bar),

H2S generated 70% of the corrosion rate. At the

lowest pCO2 /pH2S ratio of 1.67 (pCO2 6.9 bar,pH2S 4.14 bar), H2S generated 82% of the overallcorrosion rate.

The longest H2S containing corrosion experi-

ments which are practically achievable in the lab are

of the order of a few weeks or at best a few months,

while predictions are meant to cover a period of

at least a decade, in order to be meaningful. With

this in mind, it is interesting to take the experimental

conditions above (pCO2 6.9 bar, pH2S 3.45 bar,

0

2

4

6

8

10

0

Time (h)

Corrosionrate(m

my

ear

1)

pH2S= 0 mbar (exp.)

pH2S= 4 mbar (exp.)

pH2S= 70mbar (exp.)pH2S= 4 mbar (mod.)

pH2S= 70mbar (mod.)

Pure CO2corrosion rate

100 200 300 400 500 600


conditions: total pressurep 3 bar,pCO2

2bar, pH2

S 370mbar,T 70C, experiment duration 221 days, pH 4.24.9,

liquid velocity 0.3m s1. Experimental data taken from Singer et al.42

0

1

10

100

0

Time (h)

Corrosionrate(mmy

ear

1)

Pure CO2corrosion rate

100 200 300 400 500 600

pH2S = 1.38 bar (exp.)




pH2S = 1.38 bar (mod.)





conditions: total pressurep 7 bar,pCO2 6.9 bar,pH2S 1.384.14bar,T 120C,experiment duration 116 days,

pH 3.954.96, liquid velocity 10 m s1. Experimental data taken from Kvarekval et al.45



28/29

T 120 C, pH 4,v 1 0 m s1) and extend the simu-lation to 25 years. The result is shown inFigure 23.

The corrosion rate was predicted to start out rather

high as observed in the experiments; however, it was

reduced to below 0.1 mm year1 after 2 years and

was as low as 0.03 mm year1 after 25 years. The aver-

age corrosion rate over this period was only 0.06 mmyear1, which amounts to a wall thickness loss of only

1.5 mm over the 25 years, an acceptable amount

by any practical account. Actually, most of the other

conditions simulated have shown that rather low

H2S uniform corrosion rates are obtained for very

long exposures, which agrees with general field expe-

rience as recently discussed by Bonis et al.33 Never-

theless, no quantitative long-term lab data are

currently available to back-up these long-term predic-

tion

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Corrosion in Acid Gas Solutions

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