effects of water and gas compositions on the internal...

CORROSION SCIENCE SECTION

221CORROSION–Vol. 57, No. 30010-9312/01/000045/$5.00+$0.50/0

© 2001, NACE International

Submitted for publication April 2000; in revised form, November2000.

* CNWRA, Southwest Research Institute, San Antonio, TX.** OLI Systems, Inc., Morris Plains, NJ.

*** Consultant, Houston, TX.

Effects of Water and Gas Compositionson the Internal Corrosion of Gas Pipelines—Modeling and Experimental Studies

N. Sridhar, D.S. Dunn,* A.M. Anderko, M.M. Lencka,** and H.U. Schutt***

ABSTRACT

Results of a laboratory study of internal corrosion in wet gaspipelines under conditions involving a stagnant or slowlyflowing aqueous phase indicate that the corrosion rate is notaffected by calcium and magnesium scale-forming tendencyof the test solutions, provided that they are buffered. It ispossible that the scales formed are sufficiently porous to per-mit electrolyte contact with steel. However, an increase in pHand decrease in oxygen resulted in significant corrosion ratereduction. Thermodynamic analyses indicate that the forma-tion of metastable iron sulfide (FeS) precipitates is promotedby higher pH and higher dissolved iron concentration, andhindered by the presence of carbon dioxide (CO2). The pres-ence of oxygen on corrosion may lead to the transformationof a metastable FeS phase, mackinawite (Fe1+xS), to anothermetastable phase, greigite (Fe3S4). Electrochemical polariza-tion data indicate that the steel behaves in an active mannersymptomatic of a nonprotective corrosion product film. Theeffect of oxygen in increasing the corrosion rate may be re-lated to the increase in corrosion potential. Surface analysisby Raman spectroscopy confirmed some of the thermody-namic predictions of stable phases. Additionally, Ramanspectroscopy indicated the predominance of akaganeite (β-)and repidocrocite (γ-FeOOH). Based on the results of thisstudy, a regression equation developed from nonscale-form-ing solutions may be adopted for predicting the internal cor-rosion of gas transmission and gathering lines caused bycondensed water and acid gases.

KEY WORDS: carbon dioxide, hydrogen sulfide, natural gas,oxygen, Raman spectroscopy, scaling, thermodynamicmodeling

INTRODUCTION

The ability to safely transport wet, untreated naturalgases through pipelines offshore or at other inacces-sible locations is an important factor in the develop-ment of new gas fields. The internal corrosion rate ofsteel pipelines varies in a complex way with the gascomposition—specifically carbon dioxide (CO2), O2,and hydrogen sulfide (H2S) partial pressures—andcondensed water chemistry. Estimating the corrosionrate of steel at inaccessible locations from the analy-sis of the gas and water composition will enable abetter determination of the need for corrosion inhibi-tors. Quantitative understanding of the corrosionrate of steel under these conditions will be key to anaccurate risk assessment of pipelines from internalcorrosion.

Corrosion in gas gathering and transmissionlines has been reported in the presence of H2S andCO2 in the gas and high chloride concentrations orsulfur/polysulfide sludge from the formation water.1-2

While O2-induced corrosion of gas transmission orgathering lines has not been reported, a wide rangeof O2 concentration limits can be found in the gasindustry. A 1988 survey of 44 natural gas transmis-sion pipeline companies indicated that the gasquality specifications allowed maximum O2 concen-trations ranging from < 10 vol% to 1.0 vol%.3 Theactual concentrations of O2 in the gas stream were


222 CORROSION–MARCH 2001

low, 0 vol% to 0.02 vol%. However, the typical O2

concentrations in the natural gas pipelines are be-lieved to be rising, particularly in lines transportinggases pumped from storage fields and from fieldsbeing produced under vacuum. At present, little isknown of the safe and economical limits of O2. Theobjectives of the research presented in this paper areto understand the effects of water chemistry and gascomposition on corrosion through thermodynamicand kinetic analyses and to determine the effect ofscale-forming species in the water on corrosion.These effects are examined through a combinationof experimental investigation and thermodynamicmodeling.

LITERATURE REVIEW

CO2 Corrosion — Dissolved CO2 forms carbonicacid (H2CO3), which then increases the cathodic reac-tion kinetics by dissociation to bicarbonate. Understagnant conditions, dissolved ferrous ions combinewith the H2CO3 to form ferrous carbonate (siderite;FeCO3). However, under flow conditions, parts of theFeCO3 scale may be removed, resulting in an in-creased corrosion rate of the steel attributable to theH2CO3. DeWaard and Milliams developed a semi-em-pirical correlation between corrosion rate under flow-ing conditions (≈ 1 m/s at the metal surface) and CO2

partial pressure.4 This relationship was later simpli-fied in the form of a nomogram by DeWaard, et al.:5

log . ., / . . logc r mm yT

pCO( ) = − + ( )5 81780

0 67 2 (1)

Equation (1) provides a conservative estimate of cor-rosion rate under flowing conditions because it doesnot account for the effect of nonideality of the gasphase and scale formation. The effect of total pres-sure of gas and FeCO3 scale formation on decreasingthe corrosion rate was given in terms of correctionfactors.5 One of the major limitations of this ap-proach is that the correction factors attributable toscale formation, pH, and total iron concentration arenot considered in terms of a consistent thermody-namic speciation model. Other correlations havebeen presented in a recent review of the literaturethat apply to flowing solutions containing a CO2/H2S

ratio > 200 although no rationale is provided for thiscutoff in the ratio.6 Pitting was observed when the pHwas allowed to attain its natural value (~ 4) whereasno pitting was seen when the pH was controlled at avalue of ~ 6 using bicarbonate.6

H2S Corrosion and Iron Sulfides — A small con-centration of H2S in the gas stream is generally be-lieved to be beneficial because of the formation of aprotective iron sulfide film. Shoesmith, et al., exam-ined the behavior of iron and steel in aqueous sulfidesolutions because of its importance to the Girdler-Sulfide process used to produce heavy water in theCanadian nuclear program.7-8 In strongly alkalineconditions, steel was passivated by an oxide film thatmight have contained deposited sulfur in the pores,but no sulfide was detected.7 In solutions of pH rang-ing from 9 to 12, mackinawite (Fe1+xS) formed afterinitial oxidation of steel. The total sulfide concentra-tion in these investigations ranged from 0.04 M to0.5 M. The mackinawite film was not completely pro-tective and tended to flake off as it thickened.8 Ironsulfides are known to increase the corrosion rate ofsteel. The iron sulfides and their crystal structuresare given in Table 1.

Smith and Miller reviewed the effects of variousiron sulfide compounds on the corrosion of ironin aqueous solutions.9 Their review revealed thefollowing:

— Although iron sulfide was thermodynamicallypredicted to form over a wider range of pH and po-tential than iron oxide, iron sulfide film was generallynot as protective as iron oxide and a more negativecathodic protection potential was required to protectsteel in the presence of mackinawite than was gener-ally considered adequate for steel.

— The protectiveness of sulfide scale dependedon the sulfide concentration in the aqueous solutionand pH. For example, protective pyrrhotite (Fe1–xS)scale was reported in the range of 15 ppm to1,700 ppm whereas nonprotective mackinawitewas observed above this concentration and in thepH range of 6.5 to 8.8.

— The corrosiveness of iron sulfide depended onthe aqueous solution chemistry. In distilled watersaturated with H2S, a nonprotective mackinawitescale was formed initially but transformed to a pro-tective pyrrhotite and pyrite (FeS2) scale. In the pres-

TABLE 1The Types of Iron Sulfides Generally Observed and their Crystal Structures

Sulfide Chemical Formula Stoichiometry Crystal Structure

Mackinawite Fe1+xS x = 0.057-0.064 TetragonalPyrrhotite Fe1–xS x = 0 to 0.14 VariableGreigite Fe3S4 — CubicSmythite Fe3+xS4 x = 0 to 0.25 HexagonalPyrite FeS2 S or Fe deficient CubicMarcasite FeS2 S deficient, unstable Orthorhombic


223CORROSION–Vol. 57, No. 3

ence of brine and CO2, only a nonprotectivemackinawite scale was formed. The corrosion ratecontinued to increase in the presence of brine.

— In an aqueous solution with iron sulfide sus-pensions and on an equal molar basis, pyrite wasfound to be the most corrosive, followed by smythite,greigite (Fe3S4), mackinawite, and pyrrhotite. Thecorrosivity was found to be proportional to the sul-fur-to-iron ratio of these compounds (Table 1). It wasalso found that sulfides produced a strong depolar-ization of the metal surface, which counteracted theprotectiveness of the film. Thus, any defect in thefilm is likely to enhance the corrosion rate of thesteel considerably.

Shoesmith, et al., using galvanostatic tests inaqueous solutions saturated with H2S, showed thatbetween pH of 4 and 7 the corrosion product consistsof mostly mackinawite, and that this layer is protec-tive at first but becomes more porous with time.10 Itwas observed that, in these solutions, the potentialinitially increases corresponding to film formation,then decreases corresponding to disruption, andthen stays constant. Huang, et al., observed thatsaturation of a 4 wt% sodium chloride (NaCl) solutionby H2S resulted in an overall decrease in corrosionrate of A516 (UNS K01880)(1) steel, attributable tosulfide precipitation, but increased the corrosion ofthe ferrite phase because of preferential cathodic re-actions on the iron sulfide.11 A recent investigationhas shown that, in the absence of iron sulfide forma-tion, small additions of H2S (up to 30 mmol) acceler-ate the dissolution rate of steel.12 Thus, from theliterature, it appears that iron sulfide scales protectcorrosion of steel only under a limited set of condi-tions and that any defect in the film is likely to exac-erbate the corrosion considerably.

CO2 ± H2S Corrosion — The earliest studies of thecombined effects of CO2 and H2S were carried out byGreco and Wright13 and Sardisco, et al.,14 who foundthat a protective sulfide film formed at concentra-tions of H2S < 1,700 ppm corresponding to gas pres-sure of < 0.1 psia. At higher concentrations, anonprotective sulfide film was reported. Videm andKvarekval also observed that small concentrations ofH2S (0.02 mmol or 0.0065 psi) decreased the corro-sion rate of steel at 70°C and 80°C in a 1 M NaClsolution with 10 psi CO2.15 However, pitting occurredin these solutions, possibly caused by selective disso-lution of the ferrite phase. At higher H2S concentra-tions (0.002 psi to 0.0082 psi), considerable scatterin corrosion rate was observed, with the corrosionrate generally increasing with H2S concentration. Atthe natural pH of the H2S and CO2 solutions (~ 4),the film formed was not visible and no evidence ofsulfides was found. At a pH of 6.9, a thick corrosion

product was observed, which consisted of a mixtureof FeS, FeCO3, and pyrrhotite. Ho-Chung-Qui andWilliamson reported that, in an environment contain-ing H2S/CO2 at the ratio of ~ 4, chloride concentra-tion > 10,000 ppm caused severe localized corrosion.1

The corrosion was associated with the presence offerrous chloride (FeCl2), which formed as a layerbetween iron sulfide and the metal.

O2 Effects — Very few studies of the effect of O2

have been reported in the literature. Durr andBeavers examined the effect of various O2 concentra-tions in a 1,200-psi gas mixture consisting of 1% CO2

and 3.76 ppm H2S above a stagnant solution (flowinggas mixture) of 1 wt% NaCl.16 The corrosion rate washighest at the vapor/liquid interface. The corrosionrate was significant (0.086 mm/y overall corrosionrate and 0.356 mm/y at the vapor/liquid interface)even at the lowest O2 concentration studied (10 ppm),and there was a gradual increase in penetration withO2 concentration.

Lyle and Schutt showed that, while formation ofiron sulfide films decreased corrosion rates, the sul-fide film was not completely protective, as evidencedby localized corrosion.17 Buffering the solutions to apH of 6 resulted in a marked decrease in uniformand localized corrosion rate. Chloride increased thelocalized corrosion rate. Slowly flowing solution in-creased general corrosion rates and reduced the pit-ting susceptibility of carbon steel relative to stagnantconditions. Under slowly flowing conditions, thepresence of O2 increased the corrosion rate signifi-cantly and enhanced localized corrosion. Specimenswere more susceptible to pitting corrosion on por-tions of partially immersed specimens exposed to thevapor phase. Pitting occurred on fully immersedspecimens in only 1 of 11 test conditions, while onthe vapor phase portions of partially immersed speci-mens pitting occurred in 7 of the 11 test conditions.Severe pitting occurred on steel surfaces exposedto the vapor phase when O2 at a concentration of100 ppm by volume (ppmv) or 1,000 ppmv waspresent as the only constituent in the gas being usedin the tests. When 0.5 psi H2S was combined with100 ppmv or 1,000 ppmv of O2, severe corrosion ofsteel surfaces exposed in the vapor phase resulted.In the absence of O2, 0.5 psi H2S had little effect onsteel corrosion rates. Increasing CO2 partial pres-sures tended to increase general corrosion rates ofcarbon steel and made it more susceptible to pitting,particularly on portions of specimens exposed to thevapor phase. The experimental data of fully im-mersed specimens yielded a statistical regressionequation for general corrosion:

CR O O

pH CO H S

CO O H S O O pH

= + × ( ) − × ( ) −

( ) + × ( )( ) − ×

( )( ) − × ( )( ) − × ( )( )

− −

− −

− −

8 7 9 86 10 1 48 10

1 31 4 93 10 4 82 10

2 37 10 1 11 10

32

72

2

22 2

5

2 23

2 23

2

. . .

. . .

. .(2)(1) UNS numbers are listed in Metals and Alloys in the Unified

Numbering System, published by the Society of AutomotiveEngineers (SAE) and cosponsored by ASTM.



The adjusted regression coefficient (R2) value forEquation (2) was 0.6235. Equation (2) indicates thatO2 is the most important accelerator of corrosion,that there is a synergistic action between H2S andCO2 in increasing corrosion rate, and that increasingpH results in a decrease in corrosion rate. Unlike theDeWaard, et al., equation (Equation [1]), no directeffect of CO2 partial pressure was found under theconditions examined. Lyle and Schutt conductedone test for a longer period of time (1 month vs the14 days for all other tests) and found that the generalcorrosion rates increased with time whereas pittingrate (pit depth divided by total test time) decreased.17

This would suggest that the corrosion occurredmostly under a nonprotective film and any “localized”corrosion was caused by the nonuniform nature ofthe corrosion product formed.

Effect of Scale-Forming Species — Some of thecorrosion rates measured by Lyle and Schutt werehigher than have been experienced by some pipelineoperators.17 This was thought to be caused by theinfluence of scale-forming species in the pipelineenvironment. There have been no reported studies ofthe effects of scale-forming species such as calciumand magnesium on the internal corrosion of pipe-lines. Although such data are available from fieldcoupon studies, the environmental conditions towhich these coupons are exposed are quite variableand complex. The method used to predict precipita-tion of calcium sulfate (CaSO4) and calcium carbon-ate (CaCO3) in oilfield waters is based on theexperimental approach used by Stiff and Davis.18-19

For CaSO4, empirical solubility equations were estab-lished that modified the solubility of CaSO4 in dis-tilled water through a number of multiplicationfactors dependent on the concentration of other ionssuch as sodium and magnesium. For CaCO3, an em-pirical stability index was established similar to theLangelier index, which corrected for ionic strength,temperature, and alkalinity. These approaches lack asound theoretical justification and cannot be appliedto systems outside the range in which experimentalsolubility data are available. The experimental stud-ies presented in this paper examine the role of scale-forming species on corrosion in the presence of H2Sand CO2. A thermodynamic analysis of the system isalso presented to better understand the range of con-ditions possible in these systems.

EXPERIMENTAL PROCEDURES

Cylindrical specimens of Type 1018 carbon steel(UNS G10180) (Table 2) were exposed to stagnantaqueous solutions saturated with gas mixtures con-taining 0.5 psia H2S, 10 psia CO2, and 100 ppmv O2,in addition to high-purity nitrogen in an amount suf-ficient to maintain a total pressure of 500 psi. High-purity nitrogen was used instead of methane (CH4) toreduce the risk of exceeding the flammability limit forO2-CH4 gas mixtures. Tests were conducted in stag-nant solutions for several reasons: In many gatheringand transmission lines, stagnant or slow flow condi-tions cause high corrosion rates, stagnant conditionstend to maximize the scale-forming tendency, andoperational difficulties with the plugging of transferlines in the laboratory system by scale were encoun-tered. The solution chemistries employed corre-sponded to the range of Ca2+, Mg2+, Cl–, and SO4

2–

concentrations observed in the field. The test matrixis shown in Table 3. Solution 1 was designed to pro-duce a significant amount of calcium hydroxide(Ca[OH]2) scale by the addition of sodium hydroxide(NaOH) to calcium chloride (CaCl2). In contrast, Solu-tion 2 was slightly undersaturated with respect toCa(OH)2 and CaCO3. Solution 3 was designed to pre-cipitate magnesium hydroxide (Mg[OH]2) from magne-sium chloride (MgCl2). Solution 4 was designed toprecipitate CaSO4. Solutions 5, 8, 9, and 10 werecontrol tests to reproduce previous test results aswell as further examine the effects of H2S and O2.Finally, Solutions 6 and 7 simulated the condensatecompositions observed by one gas transmission com-pany. Solution 7 differs from Solution 6 only in termsof a slightly lower pH. The pH of test solutions werecalculated after addition of gases containing H2S andCO2. The calculations were performed using theExpress Calculate module of OLI Systems Environ-mental Simulation Program (ESP),† Versions 5.1 and6.0. NaOH, in amounts indicated by ESP, was usedto make up the pH. As a check on these calculations,the pH was measured prior to the addition of the acidgases and immediately after the test. The measuredpH values after the tests are reported in Table 4.There was reasonable agreement between the mea-sured and calculated values.

The test facilities were described fully in a previ-ous paper.17 Test exposures were conducted simulta-neously in 3 L of solution in Type 316L (UNS S31603)stainless steel autoclaves. The temperature withinthe cell containing the autoclaves was maintained at

TABLE 2Chemical Composition of 1018 Carbon Steel used in Corrosion Tests

Test Material C Mn P S Si Cr Cu Ni Al0.18 0.70 0.005 0.007 0.20 0.01 0.05 0.02 < 0.004

Composition (wt%)

† Trade name.



15.5 ± 1°C (60 ± 2°F). Premixed gases containing re-quired concentrations of H2S, CO2, and O2 were con-tinuously flowed through liquids in the autoclavesat 10 mL/min. To initiate a test, 189 L (50 gal) ofdeionized water and the required amounts of NaOHor sodium bicarbonate (NaHCO3) and NaCl weremixed in a glass storage vessel, and the solution wasdeaerated by evacuating gas from the vessel, followedby sparging high-purity nitrogen through the solu-tion. While the solution in the storage tank was beingdeaerated, specimens were placed in an autoclave,and the autoclave and the transfer lines between theautoclave, storage tank, and gas bottles were alter-nately evacuated and back-flushed with high-puritynitrogen to remove air from the system. The solutionwas then introduced into the autoclave by pressuriz-ing the glass vessel. The system was then pressur-ized to 500 psi with a premixed gas mixturecontaining the required partial pressures of H2S,CO2, and oxygen in high-purity nitrogen. In eachsolution, two fully immersed and three partially im-mersed specimens were exposed. For Solutions 6 and7, an additional specimen was placed at the bottomof the autoclave to increase the degree of scalingsince the precipitates tend to settle to the bottom.All tests were conducted for 14 days. The cylindricalspecimens had a surface area of ≈ 9.1 cm2. Prior totesting, the specimens were degreased and cleanedultrasonically in a detergent solution, rinsed in alco-hol, dried in air, and weighed on an analytical bal-ance accurate to ± 0.1 mg. Average corrosion rates oftest specimens were determined from specimenweight-change measurements. The total area of thespecimens was used for all cases, although some ofthe partially immersed specimens showed corrosiononly in the immersed portion. For the two specimensplaced at the bottom of the autoclave, the area of thethreaded holes was also considered.

Cyclic potentiodynamic polarization curves for acompletely immersed specimen were generated usinga tungsten/tungsten oxide reference electrode, andthe walls of the three autoclaves served as the auxil-iary electrodes. The W/WO3 reference electrode wascalibrated against a saturated calomel reference elec-trode (SCE) prior to the tests in H2S environments.As shown in Equation (3), the W/WO3 electrode po-tential has linear dependence on pH:

E W WO vs SCE pH/ . .3 0 0271 0 039 ( ) = ⋅ − (3)

However, since most of the solutions were bufferedwith bicarbonate, the W/WO3 electrode can be usedas a reference electrode. The suitability of the refer-ence electrode was verified by conducting ASTMG-5(2) (potentiodynamic polarization test on Type 430stainless steel [UNS S43000] in deaerated 1.0 N sul-furic acid [H2SO4] at 30°C) and G-61(2) (cyclicpotentiodynamic polarization test on Type 304L[UNS S30403] stainless steel in 3.56 wt% NaCl solu-tion at 25°C) tests prior to the tests in CO2 + H2Senvironments.

At the completion of a test and prior to cleaningthe specimen, Raman spectroscopy and x-ray diffrac-tion (XRD) analyses were performed on selectedspecimens to characterize the chemical compositionsof deposits formed on specimen surfaces. Specimenswere then cathodically cleaned to remove surfacedeposits in an alkaline cleaning compound (Endox214†). Tests on unexposed steel specimens haveshown that the cleaning procedure causes no mea-surable change in the weight of uncorroded speci-mens, and, therefore, has no measurable effect oncorrosion rates. After cathodic cleaning, specimenswere weighed and examined under a stereoscope forlocalized corrosion.

Raman spectra were obtained with a KaiserHoloProbe† system. A 532-nm laser, used as the exci-tation source, was connected with a fiber optic cable

TABLE 3Compositions of Solutions Used in the Experimental Program and their Calculated Scaling Tendencies(A)

SolutionNumber Solution Composition Final pH Scaling Tendency

1 166 g CaCl2, 58.4 g NaOH, 10 psi CO2, 0.5 psi H2S, 100 ppm O2 11.50 CaCO3: 1.0 Ca(OH)2: 1.02 166 g CaCl2, 0.821 g NaOH, 10 psi CO2, 0.5 psi H2S, 100 ppm O2 5.40 CaCO3: 0.1463 235 g MgCl2, 0.78 g NaOH, 10 psi CO2, 0.5 psi H2S 100 ppm O2 5.10 MgCO3: 1.3·10–4

4 15 g CaSO4, 0.372 g NaOH, 10 psi CO2, 0.5 psi H2S, 100 ppm O2 5.60 CaSO4·2H2O: 1.05 30 g NaCl, 1.289 g NaHCO3, 10 psi CO2, 0.5 psi H2S, 100 ppm O2 5.64 No scaling6 285 g CaCl2, 102 g MgCI2, 7.7 g KCL, 351 g NaCI, 0.639 g NaOH, 5.50 Dolomite:1.0

10 psi CO2, 0.5 psi HV, 100 ppm O2

7 285 g CaCI2, 102 g MgCl2, 7.7 g KCI, 351 g NaCI, 0.209 g NaOH, 4.80 No scaling10 psi CO2, 0.5 psi HV, 100 ppm O2

8 30 g NaCl, 0.346 g NaOH, 10 psi CO2, 0.5 psi HV, zero O2 5.20 No scaling9 30 g NaCl, 0.346 g NaOH, 10 psi CO2, zero O2 5.12 No scaling

10 0.283 g NaOH, 10 psi CO2, 0.5 psi HV, zero O2 4.95 No scaling

(A) All solutions assumed 3,000 mL of solution and the volume of the gas phase was 0.792 L.

(2) ASTM, 100 Barr Harbor Drive, West Conshohocken, PA 19428.



to a probe head with a 63.5-mm focal length objec-tive. The laser power was limited to 2.5 mW in orderto minimize fluorescence and prevent oxidation of thecorrosion products. Exposure time for each spectrumaccumulation varied but was limited to a maximumof 240 s. Typically, at least two accumulations wereused to obtain the Raman spectra of the specimens.Test specimens were placed in a V-block mounted onan X-Y µm stage with a 50-mm travel and 0.01-mm

resolution, and Raman spectra were obtained on atleast two locations along the length of the specimens.Additional spectra were collected for specimenswhere multiple corrosion products appeared to bepresent. The spectra from standard compounds werecompared in order to identify the corrosion productsformed. The spectra of standard compounds areshown in Appendix A. Selected specimens were alsoanalyzed by XRD using a Siemens D-500† x-ray

TABLE 4Results of Corrosion Tests and Surface Analyses using Raman Spectroscopy

Specimen Test Final Gas Fully Partially Observed PredictedNumber Solution pH (psi) Immersed Immersed Solids Solids

1, 2, 3, 4, 5 1 11.5 CO2 = 10 0.068, 0.06 0.11, 0.12, CaCO3 CaCO3

H2S = 0.5 0.10 FeCO3 Fe1+xSN2 = BAL β-FeOOHO2 (ppmv) = 100 Fe1+xS

6, 7, 8, 9, 10 2 5.4 CO2 = 10 0.19, 0.088 0.067, 0.40, FeCO3 Fe1+xSH2S = 0.5 0.44 β-FeOOH Very slight tendencyN2 = BAL γ-FeOOH to scale by CaCO3

O2 (ppmv) = 100 MgCO3

11, 12, 14 3 5.1 CO2 = 10 0.058, 0.035 0.1, 0.045, Fe1+xSH2S = 0.05 0.04 Unassigned Extremely slightN2 = BAL Peaks tendency to scaleO2 (ppmv) = 100 by MgCO3

16–20 4 5.6 CO2 = 10 0.147, 0.04 0.046, 0.53, CaSO4? Fe1+xSH2S = 0.5 0.056 CaCO2? CaSO4

N2 = BALO2 (ppmv) = 100

22, 24 5 5.6 CO2 = 10 0.17, 0.033 0.56, 0.57, MCO3 Fe1+xSH2S = 0.5 0.33 MHCO3

N2 = BAL γ-FeOOHO2 (ppmv) = 100 Greigite?

26, 31 6 4.9 CO2 = 10 0.107, 0.15, 0.067, β-FeOOH Fe1+xSH2S = 0.5 0.062, 1.3* 0.054 γ-FeOOH DolomiteN2 = BAL γ-Fe2O2

O2 (ppmv) = 100 Dolomite ?

32, 33, 37 7 4.8 CO2 = 10 0.13, 0.053, 0.144, 0.13, Fe1+xS Fe1+xSH2S = 0.5 1 47* 0.11 β-FeOOHN2 = BAL γ-FeOOHO2 (ppmv) = 100

38, 42 8 5.2 CO2 = 10 0.21, 0.05 0.081, Fe1+xS Fe1+xSH2S = 0.5 0.053, 0.068 γ-FeOOHN2 = BALO2 (ppmv) = 0

45, 46 9 5.1 CO2 = 10 0.35, 0.35 0.019, FeCO3 FeCO2

H2S = 0.0 0.043, 0.064 γ-FeOOHN2 = BALO2 (ppmv) = 0

49, 52 10 Final CO2 = 10 0.176, 0.12 0.26, 0.11, FeS1+x FeS1+x

4.9 H2S = 0.5 0.26 Other peaksN2 = BAL unassignedO2 (ppmv) = 0

Corrosion Rate (mm/y)



diffractometer with Cu K-alpha radiation. Thediffractometer was operated at 40 kV with a currentof 30 mA. XRD analyses of the corrosion productsagreed with the Raman analyses.

THERMODYNAMIC ANALYSIS

The Corrosion Simulation Program (CSP), devel-oped by OLI Systems† was used to find the equilib-rium speciation of complex fluids and, thus, predictthe occurrence of scale-forming compounds and cor-rosion products. The equilibrium computations arebased on a comprehensive thermodynamic model,which is applicable to aqueous systems of virtuallyany complexity.20,24-25 The most common sparinglysoluble compounds that may deposit on the metalsurface in scale-forming solutions containing Ca2+

are calcium carbonate and calcium sulfate. However,various calcium, magnesium carbonates may alsoprecipitate in multicomponent aqueous solutions(e.g., CaMg[CO3]2 and CaMg3[CO3]4 in the presence ofmagnesium ions). Stiff and Davis developed a tech-nique for predicting the solubility of CaCO3 andCaSO4 in natural brines.18-19 However, the method isnot based on rigorous thermodynamic analysis,employs arbitrary assumptions (e.g. the independentcharacter of the common ion, sodium, and magne-sium factors), and is not designed for systems withmultiple ions and neutral species. To predict theequilibrium solid and dissolved species from corro-sion reactions, two kinds of stability diagrams areavailable in CSP: Electrochemical diagrams (i.e.,those utilizing the potential [E] and pH or concentra-tions of aqueous species as independent variables)are extensions of the classical Pourbaix diagrams toreal (e.g., concentrated) solutions and offer a flexiblechoice of independent variables.24-25 These diagramsmake it possible to analyze various combinations ofredox processes on metal surfaces that may lead tothe formation of potentially protective films or to thedissolution of metals. In the second type of diagram,chemical diagrams (i.e., those utilizing concentra-tions or various species and/or temperature asindependent variables), the potential is not an inde-pendent variable, but results from the solution ofchemical and phase equilibria in the system of inter-est. These diagrams show the most stable species inthe system and, thus, represent the global equilib-rium of the system. These diagrams have been de-scribed by Lencka, et al.26

In all cases, it was assumed that the amount ofthe gas phase is 100 moles/kg of solution. This as-sumption is somewhat arbitrary and is intended tosatisfy the condition that the gas phase be suffi-ciently large so that the composition of the gas is notchanged as a result of the dissolution of CO2 and H2Sin the aqueous phase and the subsequent reactionwith iron. Thus, 0.5 psi H2S in a total gas pressure of

500 psi corresponds to 0.1 moles/kg of solution. Thevolume of gas phase in the experiments at any timeis 0.79 L. Assuming ideal gas behavior, the numberof moles of gas introduced during the course of thetest (336 h) at a gas flow rate of 10 mL/min is~ 290 moles for 3 L of solution. Therefore, the totalamount of gas may be considered to be 100 moles/kgof solution. In pipelines, the gas volume-to-solutionratio may be considerably larger than in laboratoryautoclave tests.

RESULTS

Prediction of Scaling Tendencyand Corrosion Products

Scale in this paper is distinguished from corro-sion products as resulting mainly from the chemicalreactions occurring in the solution without the pres-ence of cationic species from the metal dissolution.The detailed stability diagrams are presented else-where21 and the results are as follows:

— The formation of calcium- and magnesium-containing scales is predicted to be independent ofthe formation of iron sulfides.

— The stability of the potentially protectivescales is primarily determined by pH. Only CaSO4 ×2H2O can be stable under acidic conditions. Carbon-ates become stable for pH values ranging from 5.5 to8 for realistic CO2 concentrations. The value of pH forinitiating the precipitation of carbonates depends onthe concentration of CO2.

— There is a synergy between calcium and mag-nesium ions in the precipitation of carbonates. Themixed calcium/magnesium carbonate can be stableat somewhat lower pH values than the carbonatesthat contain only calcium or magnesium.

— To facilitate the analysis of data in a widevariety of environmental conditions, it is convenientto use solid scaling tendencies because they indicatehow close a given environmental condition is tothe precipitation of a given solid. If a separate solidphase “B” precipitates according to the followingreaction:

B Aii

= ∑ (4)

its scaling tendency is defined as:

sa

K

ii

B

=∏

(5)

where ai is the activity of the Ai species and KB is theequilibrium constant for the reaction shown in Equa-tion (4). When the solid B is in equilibrium with theaqueous phase, the scaling tendency is equal to 1. Ifone of the Ai species is absent in the solution, the



scaling tendency with respect to that species is equalto 0. In an unsaturated solution, the scaling ten-dency varies from 0 to 1. If the scaling tendency ishigh (e.g., > 0.8 or even 0.5), a relatively small varia-tion in the solution composition may cause the pre-cipitation of the solid. For scaling tendencies close to1, the corrosion rates should be smaller providedthat the scale is protective. A comparison of the an-ticipated precipitation and the precipitates observedby Raman spectroscopy and XRD are shown in Table4. Generally, there was agreement between predictedand observed scale and corrosion products.

Four phases—β-FeOOH, FeCO3, mackinawite,and CaCO3—were detected on Specimens 1 to 5, andmost of the spectra were consistent with β-FeOOH.Identification of FeCO3 and CaCO3 was made bythe location of the 280 cm–1 peak (CaCO3) and the296 cm–1 and 508 cm–1 peaks (FeCO3). The small2,171 cm–1 peak indicated that a minor amount ofmackinawite was present in the corrosion products.These results agree with the predicted scaling ten-dency of 1 with respect to CaCO3, but Ca(OH)2 wasnot observed in the Raman spectra. It is possible thatCa(OH)2 is kinetically hindered or not easily distin-guishable from the O-H stretch peak of water. Fortest Specimens 6 to 10 tested in 0.50 M CaCl2·2H2O+ 0.007 M NaOH, β-FeOOH was the most prevalentphase found in the corrosion products. Someindication of γ-FeOOH was also found based on the248 cm–1 peak. While no mackinawite was found, asignificant amount of FeCO3 was observed in thespectra of corrosion products from these specimens.No calcium scale was found. This corresponds rea-sonably well with the low tendency predicted for thescaling of CaCO3 for this solution. Raman spectra forSpecimens 11 to 15 tested in 0.82 M MgCl2·6H2O +0.007 M NaOH showed a large 3,420 cm–1 peak thatcan be attributed to water on the specimen surfacefrom the test solution. All of the spectra have a verystrong peak at 289 cm–1 to 293 cm–1, which cannot beattributed to either water or any of the possible cor-rosion products or mineral phases that would be ex-pected to precipitate from the solution. The veryslight scale-forming tendency predicted for this solu-tion is in agreement with the lack of observable scalein this solution.

The Raman spectra of corrosion products fromSpecimens 16 to 20, tested in 0.035 M CaSO4·2H2O +0.003 M NaOH, clearly indicated the presence of amixture of CaSO4 and CaCO3, in agreement with pre-dicted scaling of CaSO4. Specimens tested in 0.17 MNaCl + 0.005 M NaHCO3 were found to have mainlyγ-FeOOH and lesser amounts of β-FeOOH. Anotherphase was also present that had Raman peaks at298 cm–1 and 352 cm–1 and a broad peak at 675 cm–1.These peaks are not consistent with the presence ofNaHCO3 or any of the possible iron corrosion prod-ucts. These peaks may correspond to Fe3S4. For

example, the peak at 352 cm–1 is similar to the347 cm–1 peak found for other iron sulfides. Nakamotocites several investigations of sulfate-reducing bacte-rial proteins which have terminal Fe-S stretchingfrequencies in the 352 cm–1 region.28 Corrosion prod-ucts consisting of β-FeOOH and γ-FeOOH were alsofound on Specimens 26 and 31 tested in Solution 6,containing 0.86 M CaCl2·2H2O, 0.36 M MgCl2·6H2O,0.034 M KCl, 2.00 M NaCl, and 0.005 NaOH.Although the peak at ~ 300 cm–1 may arise fromdolomite (CaMg[CO3]2), other characteristic Ramanpeaks for CaMg(CO3)2 were absent. The presence ofCaMg(CO3)2 would be in general agreement with thepredicted formation of CaMg(CO3)2 in this solution.When the concentration of NaOH was reduced to0.0016 M (Solution 7), mackinawite was found in thecorrosion product layers, in addition to β-FeOOH andγ-FeOOH. No scaling was noted, in agreement withthermodynamic prediction.

Raman spectra for specimens tested in 0.17 MNaCl + 0.003 M NaOH indicate only the presence ofmackinawite and β-FeOOH. The relatively weak spec-tra for mackinawite likely indicates that only a smallamount of mackinawite was formed in the corrosionproduct layers. The Raman spectra of specimensfrom a subsequent test performed without H2S in thesystem (Solution 9) indicates that the corrosionproducts contain both γ-FeOOH and FeCO3. Theunique peak for FeCO3, located at 508 cm–1, is quiteweak in this spectra, especially compared to thesame peak in Solution 2 containing CaCl2. Becauseof the strong 1,080 cm–1 carbonate peak, this mayindicate that other carbonate species are forming inthe corrosion product layer. Based on the solutioncomposition, sodium carbonate (Na2CO3) would bethe only other carbonate species to form in this test,but it is unlikely given the near-neutral pH. Ramanspectra for specimens tested in 0.0023 M NaOH(Solution 10) show strong peaks at 300 cm–1 and356 cm–1. These peaks cannot be assigned to any ofthe corrosion products or mineral phases examinedpreviously. Other minor peaks located at 2,171,2,820, 3,414, and 3,692 cm–1 can be assigned tomackinawite.

The corrosion rates are shown in Table 4. Inanalyzing the corrosion behavior, results from a pre-vious program were also considered.17 The weight-loss corrosion rate is plotted against scalingtendencies of the laboratory test solutions (Table 3)in Figures 1 and 2 for fully and partially immersedsamples, respectively. There appears to be no corre-lation between the calculated scaling tendency andthe corrosion rate.

Effect of pHAs shown in Equation (2), increasing the pH

resulted in a decrease in corrosion. This is illustratedin Figure 3, which shows that the corrosion rate in



chloride-free solutions under a variety of gas mix-tures was generally lower when the pH was increasedto ~ 6. The test conducted at pH 11.5 had a solutionwith 3.53 wt% chloride. The large scatter in the datamakes an unequivocal conclusion of the role of pH oncorrosion difficult and may well be related to the dualrole of sulfide films in protecting and enhancing thecorrosion of steel.

Effect of Chloride ConcentrationThe corrosion rate in various chloride solutions,

under a gas mixture of CO2 + H2S + O2, increasedwith small additions of chloride but was insensitiveto larger concentrations. This is illustrated further inFigure 4 where the average corrosion rates of fullyimmersed and partially immersed specimens areplotted as a function of chloride concentration.However, at concentrations > 5%, localized corrosionwas observed.

Polarization Behavior of SteelA typical cyclic potentiodynamic polarization

curve is shown in Figure 5, indicating a completelyactive behavior. In the H2S environment, the increasein polarization at high current densities is essentiallythe result of transport-limited dissolution ratherthan protective film formation or a large ohmic com-ponent. The potential-current density relationship isnot linear, and there is a slight hysteresis in the po-larization behavior. The transport-limiting processwas not determined in this study, but could be thediffusion of dissolved ferrous hydrogen sulfide(Fe[HS]+) complexes away from the electrode. Thelack of passivity is evident not only in the shape ofthe polarization curves, but also in the lack of signifi-cant hysteresis between the forward and backwardscans. Typically, if passive films exist and localizedcorrosion is observed, significant hysteresis in the

polarization curves will occur. The lack of passivebehavior indicates that the scale or corrosion prod-ucts formed on the steel may have provided minimalprotection. The presence of oxygen in the gas phaseincreases the corrosion potential significantly, asshown in Figure 6. This indicates that O2 promotescorrosion through cathodic depolarization.

DISCUSSION

The lack of observed correlation between mea-sured corrosion rates and predicted scaling tendencymay be attributable to the scale being nonprotectiveor only partially covering the surface. It is possiblethat most of the precipitation occurred on the bottomof the autoclave rather than on the specimens. How-ever, the specimens on the floor of the autoclave in

FIGURE 1. Effect of scaling, as predicted by scaling tendency oncorrosion rates of fully immersed specimens.

FIGURE 2. Effect of scaling on measured corrosion rates of partiallyimmersed specimens.

FIGURE 3. Effect of pH on general corrosion rate of completelyimmersed specimens.



Solutions 6 and 7 did not exhibit lower corrosion ratethan the fully or partially immersed specimens.Indeed, these specimens showed considerablyhigher corrosion rates (1.3 mm/y in Solution 6 and1.47 mm/y in Solution 7 for floor specimens vs0.085 mm/y and 0.09 mm/y for fully immersedspecimens in the same solutions). It is possible thatthe floor specimens generated greater amounts ofiron sulfides locally, which acted as cathodic depo-larizers and increased the corrosion rates. The testtime also may not have been sufficiently long to dif-ferentiate the corrosion rates between scaling andnonscaling solutions. The range of corrosion ratesunder scale-forming conditions was generally nar-rower, with the exception of one partially immersed

specimen. This would suggest that a larger samplesize and variations in geometry of exposure may benecessary to determine the effect of scaling on corro-sion more accurately.

Empirical models developed through limitedtests are not valid beyond the range of factors usedto develop the models. Thus, the DeWaard, et al.,model (Equation [1]) cannot be used in environmentscontaining H2S in addition to CO2. By the sametoken, the empirical model developed by Lyle andSchutt (Equation [2]) cannot be used in an environ-ment without H2S and O2 because the model does notindicate a dependence on CO2 alone. The objective ofthis paper is not to develop a comprehensive modelcontaining any or all three gaseous species, but toprovide an understanding of the factors that couldinfluence the corrosion behavior in the presence ofthese gases.

While thermodynamic models cannot explain thecorrosion kinetics, the environmental conditionsunder which corrosion products can precipitate maybe examined. Whether these corrosion products areprotective or not can be examined by kinetic mea-surements and models. FeS2 and FeS (troilite orstoichiometric pyrrhotite) are the most stable ironsulfide forms. However, it is known that they do notform as the initial corrosion products. For example,FeS2 is stable under mildly acidic conditions, but itsprecipitation is kinetically inhibited.22-23 It has beenshown24-25 that stability diagrams can be used tostudy the formation of metastable species in additionto stable ones. The E-pH diagram in Figure 7 hasbeen generated by considering the metastable ironsulfide species. The dashed lines indicate the stabil-ity fields of various aqueous complexes, while thesolid lines indicate the stability fields of the solid

FIGURE 4. Effect of chloride concentration on general corrosion rate.FIGURE 5. Cyclic potentiodynamic polarization curves for steel invarious environments. Scan rate: 0.167 mV/s.

FIGURE 6. Cyclic potentiodynamic polarization curves for steel vsoxygen concentration in the gas. Scan rate: 0.167 mV/s.



phases. The dashed vertical line at ≈ pH 4.5 indicatesthe equilibrium pH predicted for this environment.Mackinawite is the metastable solid that is likely toform as a corrosion product. However, the stabilityfield of mackinawite starts at pH > ≈ 8. Thus, thesystem is predicted to be in the range of activedissolution as shown by the vertical dotted line, ata pH of ~ 4.5.

The stability fields of metastable Fe-S species areaffected by the partial pressure of H2S and theamount of iron dissolved in the aqueous phase. Sincethe experiments were performed in a stagnant solu-tion, a buildup of corrosion products could be ex-pected in the vessel. This should have increased theamount of dissolved iron that was available for theformation of mackinawite. Figure 8 shows the sameE-pH diagram for an assumed dissolved iron concen-tration of 10–4 M. It can be seen that the mackinawitestability region has shifted to lower pH values (≈ 6).The addition of 10 psi CO2 results in the appearanceof a small region of FeCO3, as shown in Figure 9.

Because the stability of mackinawite is influ-enced by both the H2S and dissolved iron concentra-tions, a different kind of stability diagram has beenconstructed to investigate the effects of these speciesin an explicit way. Figure 10 shows a stability dia-gram with the molality of iron in the aqueous phaseas an x-axis variable. The y-axis shows the amountof H2S added to the solution per 1 kg of water on theassumption that the vapor phase contains 100 molesof neutral gases. The latter assumption was arbitraryand was consistent with the requirement that theamount of the vapor phase be large so that its com-position does not change. An important result,observable in Figure 10, is the dependence of thestability field of mackinawite on the amounts of both

H2S and dissolved iron. If a sufficient gas phase ispresent, the amount of dissolved iron required toform mackinawite may be quite significant. For theexperimental condition of 0.5 psia H2S (10–1 moles/kg)mackinawite starts to appear when the amount ofdissolved iron is roughly ~ 10–4.1 (or 7.9 × 10–5) M. Inthe absence of H2S, iron oxides are predicted to form(Figure 10). In the laboratory experiments, ironoxyhydroxides, FeOOH, form prior to the formationof magnetite (Fe3O4). The amount of iron required toform mackinawite is greater when CO2 is present inthe system (Figure 11). Another interesting observa-tion is the boundary between FeCO3 and mackinawite.This boundary shows how much H2S is necessary toform mackinawite in the presence of CO2. The limit-

FIGURE 7. E-pH diagram of a system containing 0.5 psi H2S and10–6 M dissolved iron, considering only metastable iron sulfidespecies.

FIGURE 8. E-pH diagram of a system containing 0.5 psi H2S and10–4 M dissolved iron.

FIGURE 9. Effect of 10 psi CO2 on the stability fields of iron inwater and 0.5 psi H2S at 15.5°C. A dissolved concentration of10–6 moles/kg iron is assumed.



ing partial pressure of H2S for mackinawite in a CO2-containing system is ≈ 0.1 psia (10–1.67 moles/kg),which is fairly low. Since the literature indicates thatlarge amounts of mackinawite precipitation is notlikely to be protective,9 prediction of mackinawiteformation alone is not sufficient to determine thecorrosion rate.

The stability diagram for a system containing0.5 psia H2S as a function of the amount of dissolved

iron and O2 added to the vapor phase is shown inFigure 12. For the gas phase containing 100 moles ofneutral gas, 10–1 moles/kg O2 on the vertical axiscorresponds to 1,000 ppmv O2, 10–2 moles/kg corre-sponds to 100 ppmv, etc. As shown in Figure 12, theincrease in the amount of oxygen results first in theconversion of mackinawite into Fe3S4. This starts at1 ppmv O2 (10–4 moles/kg on the vertical scale). Theboundary between mackinawite and Fe3S4 dependson the amount of dissolved iron. However, mackinawitedisappears for O2 amounts > ≈ 5,000 ppmv for rea-sonable amounts of dissolved iron. A further increasein the amount of O2 results in the dissolution ofFe3S4 and a shift of the state of the system to theactive dissolution range. Similar to Figures 10and 11, addition of CO2 resulted in iron sulfidestability fields shifting to a higher dissolved ironconcentration.

Lyle and Shutt reported the effect of slowly flow-ing solution on the corrosion rate in O2-containingsystems.17 If the solution is slowly flowing, buildup ofcorrosion products in the vessels is minimized.Therefore, simulations were performed by assumingthat the concentration of the aqueous phase remainsvirtually unchanged. In particular, the pH of thesolution can be assumed to be constant rather thanto vary with the amount of dissolved iron. Thus, thestability diagrams have been regenerated using pHand O2 concentration as independent variables andthe amount of dissolved iron as a parameter, asexemplified by Figure 13. In this case, the effect ofpH and O2 in a system containing 10–4 m Fe, 0.5 psiH2S, and 10 psi CO2 is shown. The Fe3S4 andmackinawite incipient precipitation lines are ob-served at a pH slightly > 6.

The thermodynamic simulation can assist ex-perimental studies in identifying regions where pre-cipitation of specific corrosion products can occur.However, a thermodynamic model alone cannot pre-dict changes in corrosion rates. The polarizationcurves indicate that the dissolution behavior of thesteel in these environments does not exhibit any pas-sivity. Thus any reduction in corrosion rate observedin weight-loss experiments is essentially attributableto a barrier layer of precipitate blocking access of theelectrolyte to the metal. However, if this barrier layeris porous, the presence of iron sulfides, such asmackinawite and Fe3S4, can increase the corrosionrate of steel by providing sites for enhanced cathodicreduction reaction. The corrosion potential increasesby ~ 200 mV in the presence of H2S (Figure 5). Sincethe pH of the solution was buffered by the presenceof bicarbonate (Table 4), the increase in potential isnot likely to have been caused by the change in thereference potential of the W/WO3 electrode (Equation[3]). The increase in corrosion potential may be ex-plained by the presence of iron sulfides increasingthe cathodic reaction kinetics or the effect of H2S on

FIGURE 10. Stability diagram as a function of molality of dissolvediron and the number of moles of H2S per 1 kg of H2O and 100 mol ofneutral gas phase.

FIGURE 11. Stability diagram for a system containing 10 psi CO2.Other conditions are as in Figure 10.



the hydrogen evolution reaction. The addition of CO2

and H2S increases the anodic reaction kinetics evenfurther without a significant shift in the corrosionpotential. This is consistent with the thermodynamicspeciation model which suggests that addition of CO2

results in a shift of the mackinawite stability field tohigher iron concentrations, thus reducing the block-ing effect of iron sulfides. However, since the corro-sion potential is increased even in the presence ofCO2, the mechanism for the increase in corrosionpotential is not likely to be the catalytic effect of ironsulfides. The observed increase in anodic kinetics inthe combined presence of H2S and CO2 (Figure 5) isalso consistent with the regression equation (Equa-tion [2]) derived from weight-loss measurements.

As seen in Figure 6, the presence of O2 results inan increase in corrosion potential. Raman spectros-copy did not reveal the presence of any elementalsulfur, and hence the increase in corrosion potentialis essentially attributable to the additional cathodicreduction reaction of O2. The higher corrosion in thepartially immersed specimen (Table 4) points to thegreater availability of O2 in the vapor/liquid interface.It has also been speculated that the change in crystalstructure from mackinawite to Fe3S4 promotesdisruption of the protective sulfide and increasedlocalized corrosion.27

CONCLUSIONS

❖ The internal corrosion behavior of steel pipeline incondensate environment was studied using thermo-dynamic modeling and experimental techniques. Theimportant conclusions of this study can be summa-rized as follows:❖ No significant correlation between calcium andmagnesium scaling tendency and corrosion rate inthe presence of Cl–, CO2, H2S, and O2 was observed.This may be because the scales formed did not pro-vide adequate protection or did not form in sufficientamounts. A similar conclusion was reached byanalyzing the results of field corrosion coupons.Thermodynamic modeling provides a convenient toolto characterize the water chemistry in terms of itsscaling tendency.❖ Iron sulfide offers poor protection to the steel andcan enhance corrosion by acting as a cathodic depo-larizer. This finding was supported by electrochemi-cal polarization tests, which showed an essentiallyactive corrosion, without any passivity afforded bythe corrosion product films and an increase in corro-sion potential in H2S environments.❖ The corrosion rate measurements were also consis-tent with the anodic polarization behavior of steel,which showed a significant increase in corrosion po-tential in the presence of H2S in the gas phase andan increase in anodic currents in the combined pres-ence of CO2 and H2S.

❖ The corrosion rate of partially immersed specimenswas higher than that of fully immersed specimens.This may be related to the easier access of O2 to thesurface or the effect of corrosion products such asgreigite in enhancing cathodic reaction kinetics.❖ The comprehensive thermodynamic modeling ap-plied to the internal corrosion environment providesa useful tool to examine the corrosion behavior in acomplex environment in terms of the stability of vari-ous precipitated compounds. The model predictionswere in general agreement with surface analytical re-sults by Raman spectroscopy. Further developmentof standard Raman spectra may help confirm thepresence of certain compounds predicted by themodel and suspected to be present.

FIGURE 12. Stability diagram for a system containing 0.5 psi H2Sas a function of iron and O2 concentrations.

FIGURE 13. Stability diagram for a system containing 10–4 moles/kgdissolved iron and 10 psi CO2 as a function of pH and H2Sconcentration.



❖ In addition to the thermodynamic predictions ofthe regions of phase stability, the morphology, elec-trochemical kinetics of anodic and cathodic reac-tions, and possible deleterious effects of precipitatedcompounds such as mackinawite and Fe3S4 need tobe considered in evaluating the effects of environ-ment on corrosion.

ACKNOWLEDGMENTS

The authors acknowledge the contributions andsupport of various Southwest Research Institute staffmembers in carrying out the activities described inthis report: V.D. Aaron, J.L. Sievert, and S.T. Clay,who were responsible for the day-to-day conduct oflaboratory corrosion exposures; M. Bogart and J.F.Spencer, who performed the Raman spectroscopyand x-ray diffraction studies, respectively; and G.Cragnolino, O. Moghissi, and F.F. Lyle, who providedvaluable advice during the project. This work wasfunded by PRCI under the project number PR-15-9712. The opinions, findings, and conclusions ex-pressed are those of the authors and not necessarilythose of the PRCI. Mention of company or productnames is not to be considered an endorsement byPRCI or SwRI.

REFERENCES

1. D.F. Ho-Chung-Qui, A.I. Williamson, “Pipeline Corrosion—H2Sand O2 Corrosion of Pipelines,” Corrosion Information Compila-tion Series (Houston, TX: NACE International, 1999), p. 41.

2. N.N. Bich, K. Goerz, Pipeline Gas J. (1997): p. 46-48.3. F.F. Lyle, H.C. Burghard, Jr., E.P. George, “Effect of Natural

Gas Quality on Corrosion of CNG Storage Cylinders—Phase I,”Report No. 90-13, Project No. 730-FFES-FUC-85 (Albany, NY:New York State Energy Development Authority, 1990).

4. C. de Waard, D.E. Milliams. 1975. “Prediction of Carbonic AcidCorrosion in Natural Gas Pipelines,” in CO2 Corrosion in Oiland Gas Production (Houston, TX: NACE, 1984), p. 506-510.

5. C. de Waard, U. Lotz, A. Dugstad, “Influence of Liquid FlowVelocity on CO2 Corrosion: A Semi-Empirical Model,” CORRO-SION/95, paper no. 128 (Houston, TX: NACE, 1995).

6. C.A. Palacios, Y. Hernandez, “Application of SimulationTechniques for Internal Corrosion Prediction,” CORROSION/97,paper no. 2 (Houston, TX: NACE, 1997).

7. D.W. Shoesmith, P. Taylor, M.G. Bailey, B. Ikeda, Electrochem.Acta 23 (1978a): p. 903-916.

8. D.W. Shoesmith, M.G. Bailey, B. Ikeda, Electrochem. Acta 23(1978b): p. 1,329-1,339.

9. J.S. Smith, J.D.A. Miller, Brit. Corros. J. 10, 3 (1975): p. 136-143.

10. D.W. Shoesmith, P. Taylor, M.G. Bailey, D. Owen, J.Electrochem. Soc. 127 (1980): p. 1,007-1,015.

11. H.H. Huang, W.T. Tsai, J.T. Lee, Corrosion 52, 9 (1996): p. 708-713.

12. X.L. Cheng, H.Y. Ma, J.P. Zhang, X. Chen, S.H. Chen, H.Q.Yang, Corrosion 54, 5 (1998): p. 369-376.

13. E.C. Greco, W.B. Wright, Corrosion 18 (1962): p. 119t-124t.14. J.B. Sardisco, W.B. Wright, E.C. Greco, Corrosion 19 (1963):

p. 354t.15. K. Videm, I. Kvarekval, Corrosion 51, 4 (1995): p. 260-269.16. C.L. Durr, J.A. Beavers, “Effect of Oxygen on the Internal

Corrosion of Natural Gas Pipelines,” CORROSION/96, paper no.612 (Houston, TX: NACE, 1996).

17. F.F. Lyle, H.U. Schutt, “CO2/H2S Corrosion Under Wet GasPipeline Conditions in the Presence of Bicarbonate, Chloride,and Oxygen,” CORROSION/98, paper no. 11 (Houston, TX:NACE, 1998).

18. H. Stiff, L.E. Davis, “A Method for Predicting the Tendency ofOilfield Waters to Deposit Calcium Sulfate,” PetroleumTransactions. American Institute of Mechanical Engineering,195 (1952a): p. 25-28.

19. H. Stiff, L.E. Davis, “A Method for Predicting the Tendency ofOilfield Waters to Deposit Calcium Carbonate,” PetroleumTransactions. American Institute of Mechanical Engineering,195 (1952a): p. 213-216.

20. M. Rafal, J.W. Berthold, N.C. Scrivner, S.L. Grise, “Models forElectrolyte Solutions,” in Models for Thermodynamic and PhaseEquilibria Calculations, ed. S.I. Sandler (New York, NY: MarcelDekker, 1995), p. 601-607.

21. N. Sridhar, D.S. Dunn, A. Anderko, M. Lencka, “The Effects ofWater Chemistry on Internal Corrosion of Steel Pipelines,” FinalReport PR-15-9712 (Arlington, VA: Pipeline Research Commit-tee International, 1998).

22. M.A.A. Schoonen, H.L. Barnes, Cosmochim. Acta 55, 6 (1991):p. 1,495-1,504.

23. M.A.A. Schoonen, H.L. Barnes, Geochim. Cosmochim. Acta 55,6 (1991): p. 1,505-1,514.

24. A. Anderko, P.J. Shuler, Comput. Geosci. 23 (1997): p. 647-658.

25. A. Anderko, S.J. Sanders, R.D. Young, Corrosion 53, 1 (1997):p. 43-53.

26. M.M. Lencka, E. Nielson, A. Anderko, R.E. Rieman, Chemistryof Materials 9 (1997): p. 1,116-1,125.

27. M.B. McNeil, B.J. Little, Corrosion 46, 7 (1990): p. 599-600.28. K. Nakamoto, Infrared and Raman Spectra of Inorganic and

Coordination Compounds (New York, NY: John Wiley, 1978).29. D.S. Dunn, M.B. Bogart, C.S. Brossia, G.A. Cragnolino,

“Corrosion of Iron Under Alternating Wet and Dry Conditions,”CORROSION/99, paper no. 223 (Houston, TX: NACE, 1999).

APPENDIX A

Raman Spectra of Standard CompoundsThe Raman spectra for iron oxides and hydrox-

ides have been reported elsewhere.29 Corrosion prod-ucts formed on iron exposed to aqueous chloridesolutions typically contain γ-FeOOH, β-FeOOH,Fe3O4, and γ-Fe2O3. For most of these compounds,the Raman spectra can be used to identify thephases present. It is often difficult to identify the for-mation of Fe3O4 when γ-FeOOH, β-FeOOH, or γ-Fe2O3

are present due overlapping peaks hiding the rela-tively weak 667 cm–1 peak for Fe3O4. For γ-FeOOH,the strong peaks at 248 cm –1 and 1,311 cm–1 are themost distinguishing features. Sharp asymmetricpeaks at 719 to 725 and 3,516 cm–1 with a broadpeak at 1,389 can be used to determine the presenceof β-FeOOH. The 309 cm–1 and 388 cm–1 peaks arealso useful in determining the presence of β-FeOOH,however, these peaks are somewhat broad and aresimilar to the Raman peaks for other oxides andoxyhydroxides. The Raman spectra for Fe3O4 has onlythree peaks and of these only the 667 peak is consis-tently distinguishable. It should be noted that theRaman spectra for Fe3O4 is much less intense thanthe spectra of most of the other oxides and oxyhy-droxides. The Raman spectra for γ-Fe2O3 has someminor peaks in the 260 cm–1 to 510 cm–1 range. Themost intense peaks are broad peaks centered at717 cm–1 and 1,412 cm–1. Several other peaks arealso present in the 2,580 cm–1 to 3,910 cm–1 range.

Raman spectra of iron sulfides used as stan-dards for identifying the corrosion products formedon the test specimens are shown in Figures A-1. The



spectra for Fe2S, which is not expected to form underthe conditions used in this study, has very welldefined peaks at 344 cm–1 and 379 cm–1. A group of5 peaks is located in the range of 1,022 cm–1 to1,126 cm–1. The spectra for FeS (troilite or pyrrhotite),on the other hand, does not have such strong sharppeaks and is quite similar to the spectra obtained forγ-FeOOH. The spectrum for mackinawite has manysharp peaks. The strongest peaks are located at 347,2,171, 2,846, 3,527, and 3,691 cm–1. The spectrumof mackinawite shown in Figure A-1 is also notice-ably noisier than the other iron sulfide spectra. The

FIGURE A-2. Raman spectra of NaHCO3 and Na2CO2 spectraobtained with a 532-nm system using a laser power of 2 mW. Intensityscales are arbitrary. A strong peak located around 1,080 cm–1 iscommon for most carbonate minerals.

FIGURE A-1. Raman spectra of iron sulfides. Spectra obtained witha 532-nm system using a laser power of 2 mW. Intensity scales arearbitrary.

reference spectrum for Fe3S4 was not obtained inthis study.

The spectra shown in Figure A-2 are of sodiumcarbonate (Na2CO3) and sodium bicarbonate(NaHCO3). One of the test solutions used in this in-vestigation contained NaHCO3. The spectrum forNa2CO3 is quite similar to most other spectra for car-bonate minerals including FeCO3 and CaCO3. All ofthese carbonate minerals have a strong sharp peakat 1,080 cm–1 to 1,090 cm–1. The spectrum for CaCO3

has an additional peak at 280 cm–1 while FeCO3 hasadditional peaks at 290 cm–1 and 508 cm–1.

effects of water and gas compositions on the internal...

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