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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)
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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
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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.
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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.
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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
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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
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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
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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.
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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.
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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
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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.
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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
1284 Liquid Corrosion Environments
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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
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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
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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.
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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
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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.
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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.
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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
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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.
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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
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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
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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
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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
Figure 21 The corrosion rate vs. time; experimental data (exp.) shown as points, model predictions (mod.) shown as lines;
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 = 2.76 bar (exp.)
pH2S = 3.45 bar (exp.)
pH2S = 4.14 bar (exp.)
pH2S = 1.38 bar (mod.)
pH2S = 2.76 bar (mod.)
pH2S = 3.45 bar (mod.)
pH2S = 4.14 bar (mod.)
Figure 22 The corrosion rate vs. time; experimental data (exp.) shown as points, model predictions (mod.) shown as lines;
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
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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