Paper No.
11072
2011
Effect of FeCO3 Supersaturation and Carbide Exposure on the
CO2 Corrosion Rate of Carbon Steel
Marion Seiersten
Institute for Energy Technology
P.O. Box 40
N-2007 Kjeller
Norway
Tonje Berntsen
Statoil ASA
P.O. Box 308
N-5501 Haugesund
Norway
Tor Hemmingsen
University of Stavanger
N-4036 Stavanger
Norway
ABSTRACT
The pH stabilization technique is a widely used corrosion protection method for multiphase gas
pipelines with glycol as hydrate inhibitor. It implies to increase the pH by addition of HCO 3 in order to
enhance the formation of protective iron carbonate films. The protection mechanism at ~20°C is of
concern because the conditions for precipitating protective corrosion film are less favorable compared
to higher temperatures due to the increasing solubility of FeCO3 with decreasing temperature. The
scope of the ongoing work is to investigate whether corrosion mitigation of pipelines at ~20°C relies on
the formation of protective corrosion films or if the corrosion rate is sufficiently lowered by the elevated
pH. This paper discusses the corrosion rate and corrosion potential observed on carbon steel exposed
to varying concentrations of HCO 3 and Fe2+ at 20°C in a 1wt% NaCl and 50wt% glycol solution purged
with CO2 at 1 atm partial pressure. The objective was to promote protective FeCO3 films by high iron
and bicarbonate concentrations and study the effect of supersaturation and variations in iron and
bicarbonate concentration. Protective films did not form despite high supersaturation and long exposure
times. The reason for this is discussed in light of exposed iron carbide (Fe3C) and prerequisites for iron
carbonate growth.
Key words:
CO2 corrosion, pH stabilization, FeCO3 supersaturation, film formation, precipitation.
1
©2011 by NACE International. Requests for permission to publish this manuscript in any form, in part or in whole, must be in writing to NACE
International, Publications Division, 1440 South Creek Drive, Houston, Texas 77084. The material presented and the views expressed in this paper are
solely those of the author(s) and are not necessarily endorsed by the Association.
1
INTRODUCTION
Pipelines for oil and gas are often designed with a maximum allowable corrosion tolerance of 0.1 mm/y
to allow for process upsets or unpredictable incidents despite corrosion mitigation actions. Previous
experiments [1] have shown that the unmitigated corrosion rate at ~20°C, 1 atm CO2 and 1 wt% NaCl is
about 1 mm/y. It decreased to values between 0.1 mm/y (this work) and 0.2 mm/y [1] under pH
stabilized conditions without a protective FeCO3 film. Once a truly protective FeCO3 film had
precipitated on the steel surface, the corrosion rate decreased by one order of magnitude to less than
0.01 mm/y. As the corrosion rate reached this low level, the corrosion potential increased rapidly to
-0.45 V, and then stabilized around -0.5 V.
When a protective film covers the surface, it is film properties like porosity, thickness and composition
which control the transport of reactants and corrosion products through the film that governs the
corrosion rate. These properties depend on the FeCO3 precipitation process [2]. The precipitation is
facilitated by increased CO2 partial pressure, pH, bicarbonate concentration, temperature and Fe2+
concentration; consequently increased supersaturation, and all other measures which can reduce the
transport of reactants and corrosion products to and from the steel surface [3].
Dugstad and Drønen [2] studied precipitation of FeCO3 film on one steel with carbon content of 0.057%
and another with carbon content of 0.080% at low temperature; i.e. 20°C (pH 6.5, 0.6 MPa CO2 partial
pressure). A protective corrosion film formed on the 0.08% C steel giving a corrosion rate well below
0.1 mm/y. Only small FeCO3 crystallites had formed on the 0.057% C steel after ~4 months exposure,
and the corrosion rate was still well above 0.1 mm/y; i.e. a non-protective film.
It is a prerequisite for initiating growth of FeCO3 film that the solution must be supersaturated with
regards to iron carbonate, implying that the saturation ratio (SR) of FeCO3 must be >1. The saturation
ratio is defined as
SR 
aFe2  a CO2
3
K sp
where aFe2 is the activity of Fe2+, a CO 2  is the activity of CO32 and Ksp is the solubility product of FeCO3
3
[3]. The concentration-temperature curve for the solubility of FeCO3 is inverse compared to most salts,
meaning the solubility increases with decreasing temperature. This means that the driving force for
FeCO3 precipitation, consequently SR, decreases with falling temperature.
The precipitation process involves both nucleation and particle growth. According to Sun [4], the
nucleation rate is primarily important in homogeneous crystallization processes. In the case of
crystallization onto a metal surface, the crystallization process is classified as heterogeneous and the
overall process kinetics is dominated by crystal growth. According to Johnson and Tomson [5], FeCO3
has extremely slow precipitation kinetics at temperatures below 75°C. They claim that increased SR,
i.e. high Fe2+ and CO 32 concentrations and high pH, might improve the adherence of such a film [5].
This is in agreement with studies of the induction time for precipitation of FeCO3 that have been
performed at T> 60°C. At constant MEG and Fe2+ concentrations, the induction time for measurable
precipitation decreased as a function of increasing HCO 3 concentrations; consequently increasing SR
[6]. The growth of an FeCO3 layer on steel is strongly affected by the corrosion rate at low
supersaturation; consequently the rate at which Fe2+ is released from the surface. At high
supersaturation, the corrosion rate has less of an effect on the corrosion layer accumulation rate [4].
The main difference between protective and non-protective corrosion layer morphologies is the
absence or presence of empty Fe3C (i.e. not filled with FeCO3) in contact with the steel surface,
2
2
respectively. In the absence of Fe3C (freshly polished specimen), the precipitation of FeCO3 can only
occur at the steel surface where the concentrations of Fe2+ from the corrosion and HCO 3 from the
cathodic reduction of CO2 are at its maximum. This would lead to protective FeCO3 film when the
saturation limit is exceeded. Due to galvanic coupling between the Fe3C and the steel in a direction
perpendicular to the surface, preferential corrosion of the steel matrix will result in an empty Fe3C layer
in contact with the steel surface. The electron conductive Fe3C layer provides additional surface area to
the steel surface for the cathodic reduction of CO2 to HCO 3 . This leads to local alkalinization at a
certain distance from the steel surface. If the saturation limit is exceeded, precipitation of FeCO3 can
take place inside, or more likely, on the surface of the Fe3C. This gives a non-protective layer that even
enormous iron supersaturations cannot subsequently render protective. Only layers of FeCO3 that are
directly in contact with the steel can be protective [7].
Formation of protective FeCO3 film at 20°C
Experiments performed at similar conditions to the experiments presented in this paper are discussed
in a previous publication [1]. The conclusion from the previous work was that a protective FeCO3 film
(corrosion rate lower than 0.01 mm/y) formed on X-65 steel under pH stabilized conditions. Figure 1
shows the corrosion rate and SR variations for a long term experiment performed at room temperature.
A summary of the experimental conditions is provided in Table 3 for comparison with the experiments
discussed here. A SR above 300 was required for the FeCO3 film to precipitate, refer to Figure 1.
400
1
670 hrs
CR p.a grade MEG
Bicarb added 22->67 mmol/l
SR
300
-0.50
200
Ecorr / V
250
SR
CR / (mm/y)
1000 hrs
0.1
-0.45
350
150
0.01
-0.55
-0.60
100
-0.65
1650 hrs
0.001
Ecorr
-0.70
0
0
30
60
Bicarb added 22->67 mmol/l
50
90
Time / days
120
0
150
a)
30
60
90
Time / days
120
150
b)
Figure 1: a) Corrosion rate and SR and b) corrosion potential as a function of time in ~50 wt%
MEG at room temperature, 1wt% NaCl and 1 atm CO2. No pre-corrosion. Concentration of HCO 3
was 22-88 mmol/l.
The precipitation and film growth process was monitored by measuring the Fe2+ and HCO 3
concentrations and calculating the corresponding SR. From the onset of precipitation, the decrease in
the corrosion rate from ~0.1 mm/y down to a steady level below 0.01 mm/y was slow and took ~30
days. The film growth process on the metal surface was followed by retrieving samples at ~monthly
intervals (refer to the original paper for pictures covering the whole time span of the experiment [1]).
Figure 2 shows Scanning Electron Microscopy (SEM) pictures from a specimen that was immersed for
84 days, consequently after the corrosion rate had decreased to below 0.01 mm/y.
The SEM images in Figure 2 a) and c), in combination with Energy Dispersive Spectroscopy (EDS)
analysis of the phases, clearly show that the surface was covered with a 10-15 m thick FeCO3 (d) film
of very small (1-3 m) cubic crystals (b). The Fe3C structure is visible as a light grey network integrated
in the iron carbonate in Figure 2 d). Fe3C was exposed on the surface at an early stage when the SR
was still low and the corrosion rate quite high [1].
3
3
b)
a)
Fe3C
FeCO3
d)
c)
Figure 2: The surface of a specimen from a previous experiment [1]. The specimen was exposed
to a solution with 1 wt% NaCl, 22-88 mmol/l HCO 3 , ~50wt% MEG and 1 atm CO2 for 84 days at
room temperature. Magnification 100x, 1000x, 500x and 2000x (top left-lower right).
This long term experiment showed that it is possible to form a protective FeCO3 film under pH stabilized
conditions at temperatures close to 20°C since the corrosion rate decreased to less than 0.01 mm/y
and a dense film covered the surface. The other experiments discussed in [1] indicated that
precipitation of FeCO3 film occurred more easily on a pre-corroded surface with some exposed Fe3C
compared to a freshly ground surface. In addition to an elevated pH, a high iron concentration was
needed to promote precipitation. These results formed the background for the conditions chosen for the
new experiments presented in this paper; pre-corrosion, 20°C, 1wt% NaCl, ~50wt% MEG, target SR
300 and SR 500 with 70 and 100 mmol/l HCO 3 , and corresponding iron concentrations. The objective
of the experiments presented here was to reproduce the film formation process observed in the long
term experiment under controlled temperature and bicarbonate concentration, searching to map the
effects of SR and iron concentration on the precipitation and growth processes of FeCO3.
4
4
EXPERIMENTAL PROCEDURE
Apparatus and corrosion measurement method
The tests were performed at atmospheric pressure of CO2 in 3 liter jacketed glass cells as shown in
Figure 3. The cells were thermostatically controlled at 20°C and equipped with magnetic stirring. A
standard three-electrode setup was used for all the electrochemical measurements with the cylindrical
steel specimen as the working electrode. The counter electrode was a titanium ring. A saturated
Ag/AgCl reference electrode was connected to the cell via a Luggin capillary. A potentiostat with an
eight channel multiplexer was used for the electrochemical corrosion rate measurements.
Figure 3: Glass cell with lid equipped for corrosion testing.
The corrosion rate was monitored by linear polarization resistance (LPR) measurements and
electrochemical impedance (EIS) measurements were performed to assess the IR drop in the solution,
the experimental settings are listed in Table 1. The weight loss was measured and the corresponding
corrosion rate calculated. The reported LPR corrosion rates in this work are corrected for the measured
IR drop and weight loss.
Table 1: Experimental settings for the electrochemical measurements.
LPR
Potential ramp
Scan rate
EIS
Initial frequency
Final frequency
AC voltage
± 5 mV vs. Ecorr
0.1 mV/s
10 000 Hz
0.001 / 0.01 Hz
± 10 mV vs. Ecorr
5
5
Specimen material and preparation
The chemical composition of the X65 subsea pipeline steel used in the specimens is given in Table 2.
C
0.08
Si
0.25
Table 2: Alloying elements in the X65 steel used for the corrosion tests.
Mn
S
P
Cr
Ni
V
Mo
Cu
Al
Sn
1.54 0.001 0.019 0.04 0.03 0.045 0.01 0.02 0.038 0.001
Nb
0.043
The working electrode was a cylindrical-shaped specimen with dimensions 1 cm x 1 cm Ø, and the
exposed surface area was 3.14 cm2. The specimens were washed in soap, ground with 500 and 1000
grit silicon carbide abrasive paper, and rinsed in an ultrasonic bath with acetone for 10 minutes. It was
weighed and mounted on a steel bar covered with PEEK and sealed in both ends by Teflon rings. A
PEEK piece covered the end of the steel bar and isolated it electrically from the solution. The specimen
was rinsed with ethanol just prior to immersion.
Chemicals and monitoring of test solution
The test electrolyte consisted of distilled water, 1 wt% NaCl, ~50 wt% mono ethylene glycol purged with
CO2 gas at 1 atm for at least 24 hours. The mono ethylene glycol (MEG) used was routinely checked to
make sure it was not corrosion inhibiting. NaHCO3 was used to increase the pH. In order to increase
the concentration of dissolved Fe2+, small amounts of a 3 mol/l FeCl2·4H2O purged with N2 were added
to the cells by the use of a titrator, or iron powder (>99 wt%) was dissolved in the cells.
The water/MEG content was monitored by regular Karl Fischer titration analysis of solution samples.
The evaporated water was not replaced. The solution pH was measured at regular intervals with a pH
electrode. The pH meter was calibrated in aqueous buffer solutions, and the pH value was corrected for
the MEG content in solution by the equation given by Sandengen [8].
2+
The Fe concentration in the test solution was measured at regular intervals. According to an in-house
2+
procedure, a small volume was withdrawn from the cell and the Fe concentration in the sample was
determined spectrophotometrically.
Surface examination of specimens
The specimen was rinsed immediately in ethanol upon removal from the cell. It was dried, weighed and
stored in a dry atmosphere before the surface was examined using a Scanning Electron Microscope
(SEM). Elemental analysis was performed with Energy Dispersive Spectroscopy (EDS).
Detailed description of the experiments
Experiments 1-2: Target SR 300; 70 and 100 mmol/l bicarbonate concentration
The specimens were pre-corroded in a 1 wt% NaCl and distilled water solution at 20°C according to the
exposure time given in Table 3. Then the NaHCO3 salt, MEG and NaCl solutions were added to the
cells, maintaining the NaCl concentration at 1 wt% and reaching a MEG concentration of 50 wt% and
bicarbonate concentrations of 70 and 100 mmol/l in the two cells. The MEG/NaCl solution was
transferred using a N2 gas lift arrangement. The Fe2+ concentration was increased during the
experiment by adding known amounts of 3 mol/l FeCl2 solution to reach SR 300. The bicarbonate
concentration (total alkalinity) was measured by titration shortly after the NaHCO3 salt was dissolved,
and at regular intervals throughout the experiment.
Experiments 3-4: Target SR 500; 70 and 100 mmol/l bicarbonate concentration
The procedure was changed slightly compared to experiment 1-2. Iron powder was considered to be a
better choice as a source for Fe2+ since some of the FeCl2·4H2O salt might have oxidized to trivalent
iron during storage. Prior to immersing the steel specimen, iron powder was dissolved in the 1 wt%
NaCl solution purged with CO2 in order to obtain the desired SR by adding bicarbonate only after the
pre-corrosion period. The concentration of Fe2+ was monitored until all the iron powder was completely
dissolved. The specimens were then pre-corroded in the 1 wt% NaCl solution with dissolved iron
6
6
according to the exposure times given in Table 3. The amount of iron powder was calculated so as to
give the Fe2+ concentration required to reach SR 500. The rest of the procedure was the same as for
experiments 1-2.
In addition to the experiments primarily presented here, Table 3 lists the experimental conditions for the
previous long term experiment mentioned in the introduction for comparison [1].
Table 3: Experimental conditions for experiments 1- 4 a previous experiment [1].
All the experiments had X-65 steel, ~50 wt% MEG and 1 atm CO2.
Experiment
Previous
1
2
3
4
exp.
SRtarget
300
300
500
500

70
70
100
70
100
[ HCO ]target [mmol/l]
3
2+
[Fe ]target [mg/l]
Iron source
Corroding
specimens
Temperature [°C]
room
Pre-corrosion time [h]
* 1 wt% NaCl with ~200 mg/l Fe2+
** 1 wt% NaCl with ~80 mg/l Fe2+
64
FeCl2
21
FeCl2
105
Iron powder
34
Iron powder
20
24
20
24
20
44*
20
41**
RESULTS AND DISCUSSION
Effect of saturation ratio on the corrosion rate at 20°C
Four corrosion experiments were conducted under the experimental conditions described in the
previous section. A survey of corrosion rates, corrosion potentials and other key data is provided in
Table 4. A summary of the results from the previous long term experiment mentioned in the introduction
is included for comparison.
Table 4: Overview of corrosion rates based on weight loss, weight loss calibrated LPR
corrosion rates, corrosion potentials, exposure times, measured concentrations, max SR,
temperature and average pH for experiments 1-4 and a previous experiment [1].
Experiment
Previous
1
2
3
4
exp.
CRcal,pre-corr [mm/y]
0.81
0.98
0.74
0.60
CRcal,final 100 h [mm/y]
0.006
0.06
0.13
0.15
0.07
CRcal,average [mm/y]
0.04
0.09
0.12
0.14
0.07
CRWL [mm/y]
0.04
0.13
0.16
0.16
0.11
Ecorr, last 100 h [mVAg/AgCl]
-486
-676
-685
-671
-686
Total duration [days]
162
48
48
84
85
2+
[Fe ]max measured [mg/l]
~100
93
32
97
41
SRmax
~343
385
434
498
530
79 ± 7*
68.1 ± 1.3
98.6 ± 0.3
71.7 ± 0.6
98.8 ± 1.5
[ HCO 3 ] [mmol/l]
MEG conc. [wt%]
53 ± 11
50.8 ± 0.3
50.5 ± 0.5
51.9 ± 1.2
Temperature [°C]
22 ± 1
20.4 ± 0.2
20.4 ± 0.1
20.8 ± 0.7
pH
7.0 ± 0.2*
6.89 ± 0.01
7.05 ± 0.01
6.92 ± 0.02
*After increasing the bicarbonate concentration to the desired level, see Figure 1.
In the following, the various data from each experiment will be discussed.
7
7
51.3 ± 0.6
20.8 ± 0.7
7.05 ± 0.01
Pre-corrosion period
The corrosion rates for the pre-corrosion period for experiments 1-4 are listed in Table 4 and shown in
Figure 4, 6, 8 and 10. In experiments 3-4, the corrosion rate was slightly lower than in experiments 1-2.
This is probably due to the presence of the already-dissolved iron, which gives a lower driving force for
corrosion.
Corrosion period at high SR
Experiment 1:
FeCl2 was added successively in small portions, but because the true Fe2+ concentration of the FeCl2
solution was lower than intended, it took ~ one week before enough Fe2+ was added to reach SR 300;
refer to Figure 4 a). From that point on, the Fe2+ concentration stabilized around 83 mg/l and the SR
value remained at ~330.
The corrosion rate dropped to ~0.1 mm/y when the MEG and bicarbonate solution was added, and then
decreased slowly and reached 0.06 mm/y at the end of experiment 1, as shown in Table 4 and Figure 4
a). The corrosion potential shifted in the positive direction when Fe2+ was added to the solution; see
Figure 4 b). It then decreased slightly as the corrosion rate decreased, and was –676 mV at the end of
the experiment.
10
600
-0.63
500
-0.64
CR Exp 1
Ecorr Exp 1
SR Exp 1
-0.65
200
0.1
0.01
0
a)
SR
CR / (mm/y)
300
Ecorr / VAg/AgCl
400
1
30
Time / days
-0.66
-0.67
-0.68
100
-0.69
0
-0.70
0
60
30
Time / days
60
b)
Figure 4: a) Corrosion rate, calculated SR and b) corrosion potential of experiment 1 as a
function of time. Test conditions are given in Table 3.
The SEM pictures in Figure 5 show the corroded surface of the specimen from experiment 1. The
surface consists of mostly bare steel, but has small areas covered by clusters of small cubic (1-2 m)
and larger “barrel-shaped” FeCO3 crystals (5-10 m). There was practically no Fe3C film on the surface,
but the FeCO3 present had precipitated on carbide structures, see Figure 5 c) and d). The composition
of the two phases on the steel surface was analyzed by EDS, confirming that the crystals were FeCO3
and the “network” in between was iron carbide; see Figure 5 d).
The decreasing corrosion rate and the fact that FeCO3 had precipitated on small areas on the surface,
shows that the crystallization process had started, but FeCO3 did not grow to a continuous layer despite
of the high SR.
The corrosion rate was low, but not as low as expected if a protective film had formed. The change in
corrosion potential was also marginal compared to when a protective film forms, refer to Figure 1.
8
8
a)
b)
Fe3C
c)
FeCO3
d)
Figure 5: Surface, a) and b), and cross-section, c) and d) of the specimen from experiment 1.
Test conditions are given in Table 3.
There was less carbide on the surface than expected from the amount of removed iron. The corrosion
rate in the pre-corrosion period corresponds to about 2 m removed steel, and by the end of the
experiment about 12 m. Since the Fe3C phase is cathodic with respect to the steel, this should
theoretically leave a Fe3C layer of 12 m on the surface. The SEM images illustrate the lack of such a
layer, and this indicates that the carbide structure was fragile and had fallen off during the experiment,
or when it was taken out of the solution and rinsed in ethanol.
The lack of Fe3C and the decreasing corrosion rate suggests that the corrosion rate was not influenced
by galvanic coupling between Fe3C and the bare steel, which is expected to increase the corrosion
rate.
Experiment 2:
As in experiment 1, the FeCl2 solution was added in several steps to reach and exceed SR 300, refer to
Figure 4 a). After about 10 days, the Fe2+ concentration more or less stabilized at about 28 mg/l and the
SR value remained at ~370.
9
9
10
600
-0.63
500
-0.64
CR Exp 2
Ecorr Exp 2
SR Exp 2
200
0.1
0.01
0
a)
SR
CR / (mm/y)
300
30
Time / days
Ecorr / VAg/AgCl
-0.65
400
1
-0.66
-0.67
-0.68
100
-0.69
0
-0.70
0
60
30
Time / days
60
b)
Figure 6: a) Corrosion rate, calculated SR and b) corrosion potential of experiment 2 as a
function of time. Test conditions are given in Table 3.
The corrosion rate dropped to ~0.1 mm/y when the MEG and bicarbonate solution was added and
increased slightly toward the end of the exposure, becoming 0.13 mm/y in the final 100 hours of
exposure. This is significantly higher than in experiment 1. The corrosion potential was shifted in the
positive direction when Fe2+ was added to the solution; see Figure 4 b) at ~5-8 days. It then increased
slightly as the corrosion rate increased, and was –685 mV at the end of the experiment.
The SEM images in Figure 7 show the corroded surface of the specimen from experiment 2. Figure 7 c)
indicate that more carbide was present on the surface in experiment 2 compared to experiment 1. This
is expected since the corrosion rate was higher. The iron carbide and carbonate phases were identified
using EDS. The surface has a rough appearance (a-b) with a thin (~6-8 m), but continuous film of
Fe3C (c) and scattered, “barrel-shaped” particles of FeCO3 (5-10 m) (b) precipitated on top of it. A few
clusters (~20 m) of cubic crystals (~5 m) are also observed; see Figure 7 b).
The corrosion potential responded in the same way to the added FeCl2 in this experiment as in
experiment 1. As the corrosion rate increased toward the end, the corrosion potential also increased as
expected from the Nernst equation.
The pre-corrosion treatment should have removed about 3 m of metal, and 16 m should have been
removed by the end of exposure. The SEM images show that the Fe3C film was much thinner than 16
m, implying that some of the Fe3C had fallen off at the time of SEM and EDS analysis. Still, the
existence of a thin, continuous layer of Fe3C might have contributed to the increasing corrosion rate
toward the end of exposure, either by the increase in area for the cathodic reduction or by an
accelerated corrosion rate caused by galvanic coupling of Fe3C and the steel.
The SR levelled out in the same range as in experiment 1, but as intended, the bulk Fe2+ concentration
was considerably lower in experiment 2. The corrosion rate was higher in experiment 2; the surface
concentration of Fe2+ and local SR should be higher. The existence of a Fe3C layer would in experiment
2 represent a physical separation between the steel surface, where the Fe2+ concentration is the
highest, and the cathodic sites on the carbide. Only a few small crystallites of FeCO3 had precipitated in
the very outer parts of the Fe3C layer, which is in accordance with the theory of local alkalinization in
the Fe3C layer leading to precipitation if the local SR is high enough. Evidently, the local SR was not
high enough to promote FeCO3 film formation in this case. This may indicate that a higher Fe2+
concentration in the bulk is needed in order to form a FeCO3 film.
10
10
a)
b)
FeCO3
Fe3C
c)
d)
Figure 7: Surface, a) and b), and cross-section, c) and d) of the specimen from experiment 2.
Test conditions are given in Table 3.
Experiment 3:
The SR was raised to ~500 by adding bicarbonate together with MEG immediately after the precorrosion period; refer to Figure 8 a). The iron concentration was already 90 mg/l when the specimen
was immersed due to the addition of iron powder as described above. The SR decreased steadily
between 10 and 40 days of exposure, to below 100. The solution was observed to be cloudy around 28
days. This implies that precipitation of FeCO3 occurred, which was confirmed by the SEM images in
Figure 9, and EDS analysis.
The corrosion rate dropped to ~0.1 mm/y when the MEG and bicarbonate solution was added and
remained stable for ~30 days, refer to Figure 8 a). It then increased to ~0.14 mm/y in the course of
about 10 days, coincident with the observed decrease in SR. The corrosion rate stabilized, and the final
corrosion rate was 0.15 mm/y.
The corrosion potential was ~30 mV higher after the MEG and bicarbonate was added compared to
experiments 1-2, and decreased to a minimum at around -675 mV after ~35 days of exposure. This
coincided with the rise in corrosion rate and the decrease in SR. The corrosion potential continued to
11
11
increase toward the end of the experiment, and reached a final value of -671 mV for the last 100 hours
of exposure.
10
600
-0.63
500
-0.64
CR Exp 3
Ecorr Exp 3
SR Exp 3
0.1
200
SR
CR / (mm/y)
300
Ecorr / VAg/AgCl
-0.65
400
1
-0.66
-0.67
-0.68
0.01
0
a)
30
Time / days
60
100
-0.69
0
-0.70
0
90
30
Time / days
60
b)
Figure 8: a) Corrosion rate, calculated SR and b) corrosion potential of experiment 3 as a
function of time. Test conditions are given in Table 3.
The SEM images in Figure 9 show the surface of the corroded specimen from experiment 3. The
corrosion film was loosely adhered and came off easily when the specimen was rinsed with ethanol.
The film consisted of a discontinuous Fe3C layer 4-20 m thick with regions of 5-10 m thick FeCO3
film sitting on top, i.e. the FeCO3 was not integrated into the carbide structure. Figure 9 b) shows that
the carbonate crystals were “barrel-shaped” (5-10 m) with small cubic crystals (1-2 m) on top of the
“barrels”. The iron carbide and carbonate phases were identified using EDS.
The somewhat higher initial corrosion potential in this experiment compared to the two previous ones is
attributed to the high initial Fe2+ concentration (~90 mg/l) in the same way as before. The corrosion
potential followed the decrease in the SR; i.e. the falling Fe2+ concentration in the bulk, until day 35.
The pre-corrosion period should have produced a 4 m Fe3C layer, and after 35 days, it should have
been about 13 m thick, provided it was intact on the steel surface. By the end of exposure, 32 m of
material was removed. The theoretical Fe3C thickness at 35 days corresponds well to the thickness of
the Fe3C layer observed underneath the FeCO3 film in the SEM images, and together with the SR
decrease, this indicates that the FeCO3 film precipitated quite early in the process. The total amount of
removed material corresponds well to the distance between the steel surface and the outer edge of the
FeCO3 film.
The average corrosion rate in experiment 3 was 0.14 mm/y. This is the highest corrosion rate seen
among the experiments presented here, see Table 4. This should give a higher surface concentration
of Fe2+ compared to the other experiments, and better conditions for precipitation of FeCO3. The fact
that the FeCO3 had precipitated on top of the carbide, suggest that a local alkalinization occurred there
as a result of the cathodic reduction reaction occurring on the Fe3C, and that a critical SR for
precipitation was reached locally.
The sharp increase in the corrosion rate to ~0.14 mm/y at ~30 days, is coincident with the decrease in
SR and a decrease in the corrosion potential. The decrease in SR is due to precipitation which removes
2+
a considerable amount of the Fe from the bulk solution while the HCO 3 concentration and pH
change is minor. Some of the FeCO3 precipitated in the outer part of the Fe3C layer.
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12
90
FeCO3
Fe3C
a)
b)
FeCO3
Fe3C
c)
d)
Figure 9: Surface, a) and b), and cross-section, c) and d) of the specimen from experiment 3.
Test conditions are given in Table 3.
Beyond day 35, the Fe2+ concentration and SR continued to decrease as FeCO3 precipitated, gradually
covering the Fe3C and possibly trapping the Fe2+ released by corrosion inside the solution volume close
to the surface. Simultaneously, the corrosion potential started to increase as the corrosion rate
stabilized. The observed increase in corrosion potential could be ascribed to the possible accumulation
of Fe2+ on the steel surface, but this is not fully understood.
Experiment 4:
As in experiment 3, iron powder was dissolved prior to immersing the specimen. The first SR reading
was ~300 and the corresponding Fe2+ concentration ~25 mg/l; refer to Figure 10 a). The SR was ~500
after 8-9 days of exposure and remained high for ~45 days before it decreased at approximately the
same rate as in experiment 3. It was below 150 when the experiment was terminated. This implies that
precipitation of FeCO3 occurred toward the end of the experiment. The presence of FeCO3 crystals on
the surface was confirmed by the SEM images shown in Figure 11, and EDS analysis.
The corrosion rate dropped to well below 0.1 mm/y when the MEG and bicarbonate solution was
added; refer to Figure 10 a). It remained stable for about 10 days, and then it decreased slightly down
to a minimum of 0.05 mm/y after 45 days.
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13
10
600
-0.63
500
-0.64
CR Exp 4
Ecorr Exp 4
SR Exp 4
0.1
200
SR
CR / (mm/y)
300
Ecorr / VAg/AgCl
-0.65
400
1
-0.66
-0.67
-0.68
0.01
0
a)
30
Time / days
60
100
-0.69
0
-0.70
0
90
30
Time / days
60
b)
Figure 10: a) Corrosion rate, calculated SR and b) corrosion potential of experiment 4 as a
function of time. Test conditions are given in Table 3.
The corrosion potential was significantly lower and more stable after the MEG and bicarbonate was
added compared to experiment 3. The corrosion rate increased slowly after 45 days and the final
corrosion rate was 0.07 mm/y, see Table 4. This period of increasing corrosion rate coincided with the
decrease in SR from ~500 down to ~150 and the corrosion potential decreased, analogous to
experiment 3, only it occurred later.
The SEM images in Figure 11 show the surface of the corroded specimen from experiment 4. The
corrosion film was physically more adherent compared to the specimen from experiment 3. Figure 11 ab) show that there are scattered, 10-20 m large “barrel-shaped” FeCO3 crystals on the practically bare
steel surface. The steel has a “fuzzy” appearance where the fine Fe3C structure protrudes, with some
shallow grooves where preferential corrosion has taken place. Figure 11 c-d) illustrates how the FeCO3
crystals are clearly integrated into the thin Fe3C structure. This “film” is 10-20 m thick. The iron carbide
and carbonate phases were identified using EDS.
The corrosion potential initially in experiment 4 was lower compared to experiment 3, as expected from
the much lower initial iron concentration (~25 mg/l). At a later stage, the corrosion potential decreased
coincident with the decreasing SR, caused by decreasing Fe2+ concentration on the steel due to
precipitation of FeCO3. A slight increase in the corrosion rate at the same time, similar to the period
between 15-45 days in experiment 3, but the corrosion rate increase was less.
The pre-corrosion period should have produced a 3 m Fe3C layer, and at the onset of precipitation (at
~45 days) it should have been about 10 m. By the end of exposure, 16 m of material was removed.
The SEM images show that the Fe3C film was discontinuous and thinner than 16 m, implying that
some of the Fe3C had fallen off at the time of SEM analysis.
The average corrosion rate was 0.07 mm/y, so it was considerably lower than in experiment 3 with
similar SR. The lower corrosion rate gives a lower surface concentration of Fe2+, and thus a lower
driving force for precipitation compared to experiment 3. Despite this, the SEM pictures reveal, some
precipitation occurred in the outer part of the apparently Fe3C structure. The crystallites were evenly
distributed over the surface and quite adherent.
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a)
b)
FeCO3
Fe3C
d)
c)
Figure 11: Surface, a) and b), and cross-section, c) and d) of the specimen from experiment 4.
Test conditions are given in Table 3.
Overall discussion
It was not possible to obtain a protective FeCO3 film under the conditions and time span of the four
2+
experiments discussed here. This was despite having a bulk solution with the same Fe and HCO 3
concentration as in the previous experiment with the same steel, where the corrosion rate was below
0.01 mm/y [1].
A high supersaturation (SR~500) induces bulk precipitation and precipitation of FeCO3 in the outer part
2+
of the Fe3C layer. That may lead to high local Fe concentration close to the metal, but insufficient
carbonate concentration to achieve the growth rate required for the formation of a protective film.
Whether this was due to local acidification as proposed by some [7] or was due to physical hindrance
remains an open question. The lower supersaturation (SR~300) leads to carbonate precipitation within
the carbide layer, but the time seemed to be insufficient in order to grow a protective carbonate film.
These experiments show that the kinetics of FeCO3 precipitation is very slow at the 20°C, in
accordance with Johnson and Tomson [5]. It seems that a high SR must be maintained for a certain
time before precipitation occurs and SR decreases, suggesting that there might exist some sort of
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“induction time” also for the heterogeneous crystallization process of FeCO3 as well as for the
homogeneous one [6].
A striking difference between the previous long term experiment [1] and all the four experiments
presented here is the size of the crystallites. In the case of the protective film, the crystals are solely
cubic and 1-3 m, whereas they are both cubic and “barrel-shaped” and generally larger in these
experiments. This point needs further studies.
Dugstad and Drønen did the same type of experiment with two different types of steel [2]. The 0.08% C
steel they studied had the same level of carbon as the X-65 steel used in the experiments presented in
this work, but it had somewhat higher levels of carbide-forming alloying elements, and hence it might
have contained more carbide. They obtained protective film on the 0.08% C steel, but not on the
0.057% C steel. This indicates that a key parameter for growth of FeCO3 film, apart from the
supersaturation, is the presence of carbide.
CONCLUSIONS
It is clear that it was not possible to obtain a protective FeCO3 film under the conditions and time span
of the four experiments discussed here. Neither of the experiments where the SR for FeCO3 was
2+
increased by adding Fe to the solution instead of letting it build up only due to corrosion, resulted in a
protective FeCO3 film and corrosion rates less than 0.01 mm/y at 20°C.
Apart from the supersaturation, presence of carbide is believed to be a key parameter for growth of
FeCO3 film. The effect of carbide content and the morphology of carbide film will be studied further by
corroding the matrix in a controlled manner, giving a reproducible amount of Fe3C on the steel surface.
ACKNOWLEDGEMENTS
This work has been carried out at the department of Materials and Corrosion Technology at Institute for
Energy Technology in Norway as part of my ongoing PhD project.
The financial support from Statoil ASA is highly appreciated.
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REFERENCES
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Event. No. 299, 7-11 September 2008), pp. 20.
2. A. Dugstad and P. E. Drønen, “Efficient Corrosion Control of Gas Condensate Pipelines by pHstabilisation,” CORROSION/99, paper no. 20 (Houston, TX: NACE, 1999), pp. 10.
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CORROSION/98, paper no. 31 (San Diego, CA: NACE, 1998), pp. 11.
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Corrosion,” CORROSION/ 91, paper no. 268 (Houston, TX: NACE, 1991), pp.15.
6. G. Watterud, “Precipitation of FeCO3 from MEG solutions”, Institute for Energy Technology, Kjeller,
Norway: IFE/KR/F–2009/087.
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Protectiveness of Corrosion Layers,” CORROSION/96, paper no. 4 (Houston, TX: NACE, 1996), pp.
14.
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