Copyright ©2008 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, Copyright Division, 1440 South creek Drive, Houston, Texas 777084. 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. Printed in the U.S.A.

Corrosion Rate Measurements in the Presence of Hydrogen Sulfide: Electrochemical Activity of Solution Species

David L. Miller

Faradaic Enterprises, Inc. 11201 Stormy Ridge Road Austin, Texas 78739 USA

DLMILLER@REALTIME.NET

K. Daniel Efird and Nathan Davis Efird Corrosion International, Inc.

9391 Grogan's Mill Road, Suite A-6 The Woodlands, Texas 77380 USA

Abdalla Ali

Chevron Energy Technology Co. 1600 Smith Street

Houston, TX 77002 USA

James Ward Chevron Energy Technology Co.

100 Chevron Way Richmond, California 94802 USA

ABSTRACT Electrochemical techniques are often used for measuring corrosion rates of metals such as carbon steel. They can provide rapid response and high sensitivity that make them useful for on-line corrosion monitoring. However, electrochemical methods are only accurate if the anodic current derives primarily from the corrosion process. The presence of electroactive species, either in solution or on the electrode surface, will interfere with the electrochemical measurements and if interpreted as corrosion current will lead the user to an erroneous conclusion of high corrosion rates. The present work evaluated the effects of electroactive species from a diethanolamine/water solution containing hydrogen sulfide. Carbon steel was evaluated in this fluid at H2S loading levels of 0.46 to 0.56 mole/mole and temperatures of 80, 120 and 140°C. Corrosion rates from coupon mass loss were compared to electrochemical techniques including linear polarization resistance, Tafel extrapolation (potentiodynamic polarization), electrochemical impedance spectroscopy and electrochemical noise.

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Paper No.

08414

The results indicate that bisulfide is electroactive on carbon steel and generates electrochemical currents as much as 20 times higher than the corrosion current. The effect was further evaluated by performing potentiodynamic sweeps with a graphite electrode and showed that bisulfide was readily oxidized and generated an anodic current. The oxidation of bisulfide is believed to form elemental sulfur on the electrode surface. The electroactivity of bisulfide interferes with all of the electrochemical techniques studied and strongly affects their use for corrosion monitoring of carbon steel. Key Words: Corrosion monitoring, linear polarization resistance, electrochemical, hydrogen sulfide, bisulfide sulfur redox electroactivity, amine, gas treating.

INTRODUCTION The fluids used for removal of hydrogen sulfide (H2S) and carbon dioxide (CO2) from natural gas become inherently corrosive to carbon steel(1). These acid gases lower the pH of the gas treating fluid thus affecting carbon steel both by moving away from the high pH region where iron oxides/hydroxides provide passivation(2) and by increasing the cathodic kinetics for reduction of hydrogen ions from solution. The use of aqueous amine solutions as gas treating fluids has a number of advantages, primarily a substantial increase in the acid gas absorption capacity and absorption kinetics relative to water. The amines also help with regard to corrosion, since they provide alkalinity, giving a high starting pH of near 12, and pH buffering such that, even for highly loaded solutions, the pH does not drop much below 8. This is considerably higher than would be achieved for water alone, where pH values can drop below 3 after absorption of the acid gases. Understanding the corrosion science of a gas treating unit becomes more complicated, however, by the amine chemistry and how it affects the corrosion mechanisms. For a gas treating amine containing H2S and CO2, the solution contains mixtures of protonated amine, protonated water (H3O+), sulfide (S2

-), bisulfide (HS-), undissociated H2S, carbonate (CO3

2-), bicarbonate (HCO3-) and, for primary and

secondary amines, also carbamates (RNCO2-, where RN is from the amine). The distribution of these

species depend on the acid gas loading, the pH, the alkaline strength of the amine (defined by its pKa) and on the temperature. Additionally, the chemical reactivity of the amine plays a role(1), as the electron density of the nitrogen and steric hindrance created by the ligands attached to the nitrogen affect both the thermodynamics and kinetics of the reaction between the amine with H2S and CO2. The formation and stability of passive films on carbon steel are also important factors affecting the corrosion rates. For H2S applications, iron sulfide films can provide good corrosion protection. However, the iron sulfide film can be degraded mechanically or chemically. Mechanical degradation occurs from erosion by high fluid flow and turbulence or by particulates carried in the fluid. Chemical degradation of iron sulfide films occurs if the solution concentration of sulfur is too low (very lean conditions), if the pH strays out of the stability range for iron sulfide(1,3-5), or by chelation from solution species(6-8) such as bicine (N,N-bis(hydroxyethyl)glycine). To monitor for corrosion of carbon steel in gas treating units, a number of techniques have been used. Each technique has advantages and disadvantages. Coupons provide direct mass loss which can be converted to corrosion rates but don't identify the time when the corrosion occurred. For applications involving the buildup of protective films, much of the mass loss can occur very early in the exposure time. Alternatively, a process disruption can create very corrosive conditions over a short time span. Coupons do not provide real time feedback during these events. Electrochemical techniques such as linear polarization resistance (LPR) make measurements very quickly, within a few minutes, that provide direct corrosion rate readings. The LPR technique involves application of a potential (generally less than 20 mV) to an electrode and measures the resulting

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electrochemical current. The slope at the origin provides the polarization resistance value, Rp. The corrosion current density (icorr) is then determined mathematically from the Rp value using the Tafel slopes for the anodic and cathodic reactions(9,10). This relationship is shown as: Rp = βa βc/2.3(βa + βc)icorr (1) Where βa and βc are the anodic and cathodic Tafel slopes, respectively. The corrosion rate is then calculated from the corrosion current density by: Corrosion Rate = icorr x M/nFD (2) Where: M is the atomic mass of the metal or weighted average for an alloy n is the number of electrons transferred in the corrosion reaction F is the Faraday constant D is the density of the metal However, an assumption is made that essentially all of the current flow comes from the corrosion process. If electroactive species are present in solution, then electrochemical current can flow to those species and be interpreted as corrosion(11). These electroactive species, or redox couples, need to be accounted for. Other electrochemical techniques, such as electrochemical impedance spectroscopy (EIS) and electrochemical noise (ECN), have been used for measuring corrosion rates but these can also be affected by redox couples that contribute to the measured currents. It is desired to be able to use electrochemical techniques for real time monitoring of corrosion in a gas treating unit, but only if the data are accurate. The literature shows very high corrosion rates in H2S solutions measured by LPR. High corrosion rates have been measured by LPR in gas treating amines as the H2S content is increased. Cummings, Veatch and Keller(7) used LPR to measure corrosion rates in diethanolamine (DEA) and methyl-diethanolamine (MDEA) with low H2S loadings. For a refinery sample of 30% MDEA, addition of H2S to a loading of 0.015 m/m produced a corrosion rate of around 20 mpy (0.5 mm/y) at 250°F. When the H2S loading was increased to 0.06 m/m, the corrosion rate increased to 160 mpy (4.1 mm/y), but this fluid also contained bicine. Cummings, Veatch and Keller also tested 30% DEA starting with an H2S loading of 0.015 m/m, which produced a corrosion rate of around 60 mpy (1.5 mm/y) at 245°F. Increasing the total acid gas loading to 0.06 m/m (0.03 m/m H2S plus 0.03 m/m CO2) increased the LPR corrosion rate to around 190 mpy (4.8 mm/y), although this fluid also contained partially neutralized formic acid. A number of authors have evaluated carbon steel in aqueous H2S / CO2 but without amines. Agrawal, Durr and Koch(12) performed EIS on carbon steel in aqueous sodium chloride solution with H2S and H2S / CO2 and found increasing measured corrosion rates as the temperature increased from 27°C to 79°C. They report that the EIS corrosion rates were generally much higher than determined using coupon mass loss in the same environment.

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Kvarekval, Nyborg and Choi(13) performed LPR and coupon testing on carbon steel in brine solution containing CO2 and H2S / CO2. In their case, the B value used for determining icorr from the LPR data was adjusted to correlate with their mass loss data. While this may allow use of LPR for real time corrosion readings that are closer to those generated with mass loss, it deviates from the standard methods by circumventing the connection between icorr and the electrochemical kinetics. Also, known values of steady state corrosion rates from mass loss coupons are needed in order to do this. Other papers report electrochemical methods to evaluate the buildup of iron sulfide films from aqueous solutions (no amine) containing H2S. Ramanarayanan and Smith(14) performed electrochemical potentiodynamic polarization scans on carbon steel in aqueous sodium chloride solution with H2S and determined that an iron sulfide scale forms and produces a non-linear corrosion rate during the first two days. This was further explored by Vedage, Ramanarayanan, Mumford and Smith(15) using EIS, showing that the polarization resistance increases with time but decreases with temperature. Huang, Tsai and Lee(16) performed EIS on carbon steel in NACE TM-0177 solution (aqueous 5% sodium chloride plus 0.5% acetic acid) saturated with H2S at a pH of 2.8 and found a non-linear increase in the polarization resistance during the first 10 hours as the iron sulfide film developed. Overall, the reported corrosion rates were found to be non-linear due to the buildup of iron sulfide films. This makes use of coupons difficult, since the mass loss that occurs during initial exposure should be excluded to determine the long term corrosion rates. With some effort, this can be accomplished in laboratory testing but may not be possible for process monitoring. However, the electrochemical methods were found to produce higher measured corrosion rates compared to coupon mass loss. The work reported here evaluates this phenomenon and the usefulness of electrochemical techniques, LPR in particular, for monitoring corrosion rates in amine gas treating systems containing both high (rich) and low (lean) loadings of H2S.

EXPERIMENTAL PROCEDURE Test environment. The test solution consisted of 35% DEA by weight in distilled water for all tests. The electrochemical and coupon tests performed in one liter 316SS autoclaves contained a total solution volume of 971 mL. The DEA was obtained from Huntsman Corp. The acid gas loading for each test was based upon the molar ratio of H2S and/or CO2 to DEA. The tests containing CO2 (supercritical fluid grade) were performed at a 3:1 H2S/CO2 molar ratio. Test samples. 1018 carbon steel samples (UNS G10180), purchased from the Metal Samples Company, were used in the electrochemical and coupon tests. The cylindrical electrochemical samples were 3/8� in diameter and 1/2� long. The coupons were 2� long, 1� wide and 1/16� thick. A photograph of the two sample types and the elemental analysis provided for the 1018 carbon steel coupons are shown in Table 1. Test protocols. The basic test procedures and autoclave setup for the electrochemical and coupon tests were similar. All test solutions were deaerated overnight with ultra high purity nitrogen prior to the acid gas loading and heating of the autoclaves. The electrochemical samples and coupons were electrically isolated from their support rods with Teflon spacers. A schematic of the general autoclave setup and electrochemical cap setup are shown in Figure 1. Pure H2S was condensed in a deaerated and chilled sample bomb to efficiently and accurately measure the quantity add and meet the targeted H2S/DEA molar ratios. When used, CO2 was loaded

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into a separate sample bomb and introduced into the autoclave after the H2S had fully dissolved into the test solution. All sample bombs were weighed, post autoclave loading and prior to heat up, to determine if any acid gas remained un-dissolved in the test solution. Sample bombs that did not return to their initial weight were left attached and open to the autoclave during the entire test. The autoclaves were heated to their respective test temperatures once the acid gas saturation was complete. As much as 48 hours was allowed for those tests with rich H2S+CO2 loadings at 80°C. Most tests containing only H2S were heated to the appropriate test temperature within 8 hours of the start of the acid gas saturation. Electrochemical tests The electrochemical tests were performed using a Gamry electrochemical system1. The electrochemical tests were performed in a specific order to minimize differences between the tests with respect to time of electrode exposure, film development and polarization effects. The test order was as follows:

1) Corrosion potential scan 2) LPR scan 3) Corrosion potential scan 4) EIS scan 5) Corrosion potential scan 6) ECN scan 7) Corrosion potential scan 8) Anodic and cathodic potentiodynamic scans 9) Test terminated.

Electrode mass loss data was used to verify the accuracy of the electrochemical corrosion rate measurements. Due to the high pressures and temperatures, a Hastelloy® C2000 (UNS N060200) electrode was used as a reference electrode in the electrochemical autoclave tests. The potential difference between the Hastelloy C2000 and 1018 carbon steel electrodes and a saturated calomel electrode (SCE) was determined at a 0.46 H2S/DEA mole/mole (m/m) ratio at 23°C. All potential measurements reported have been adjusted to reflect the potential difference vs. SCE at the test temperature (80 or 140°C). The SCE potentials versus SHE (standard hydrogen electrode) used to make the temperature adjustments were: SCE at 23°C was 246 mV, SCE at 80°C was 208 mV and SCE at 140°C was 176 mV. The potentiodynamic scans were performed at a rate of 1 mV/sec except where noted. Several electrochemical tests were performed in glass reaction kettles to investigate secondary redox reactions suspected of distorting the results of the electrochemical autoclave tests. Those tests were performed at 80°C with sulfide loadings of 1000 ppm, which is subsequently converted to bisulfide at the test pH of 9.75. Sodium sulfide was used instead of H2S in these tests because the glass reaction kettles could only maintain a small positive pressure and require a constant nitrogen purge to maintain an oxygen free environment. The test solution�s pH was adjusted to 9.75 with hydrochloric acid. An SCE reference electrode was used in all glass reaction kettle tests. A schematic of the glass reaction kettle test cell is shown in Figure 2. Coupon tests The coupon tests were allowed to run for a minimum of two weeks. Upon removal from the autoclave, the coupons were photographed and then cleaned in an inhibited HCl bath with hexamethylene 1 Gamry Instruments Framework Version 4.20 (copyright 2003).

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tetramine following ISO 9226. Each coupon was examined for localized attack in the form of pitting and crevice corrosion, photographed under a stereoscope, and weighed to 0.1 mg. Corrosion rate calculations were made as specified in ASTM G1. When appropriate, some electrochemical samples were treated in a similar manner.

RESULTS Electrochemical tests The corrosion rates from LPR obtained at loadings of 0.46 to 0.56 m/m are shown in Figure 3. The LPR corrosion rates were based on Tafel slopes determined for the respective conditions, discussed in the next section. As shown in Figure 3, the LPR corrosion rates are strongly dependent on temperature. At 140°C, steady corrosion rates on the order of 2,000 mpy (50 mm/y) were obtained. At the lower temperature of 80°C, the corrosion rates were on the order of 10 to 30 mpy (0.25 to 0.76 mm/y) but trended upward over time. The change in H2S loading had relatively little effect; with the higher loadings of 0.55 to 0.56 m/m showing slightly lower corrosion rates than did the lower loadings of 0.46 to 0.47 m/m for either temperature. In addition to LPR, corrosion rates were also obtained using EIS and ECN electrochemical techniques. Table 2 shows the corrosion rates obtained from all three methods. The rates at 80°C range from 31 to 266 mpy (0.79 to 6.76 mm/y). The corrosion rates at 140°C are much higher yet, with rates ranging from 740 to 10,521 mpy (19 to 267 mm/y). These rates are much higher than expected for DEA loaded with H2S at these temperatures, especially for the 80°C conditions. Measurement of the mass loss of the electrodes and converting to corrosion rates, meanwhile, produced far lower rates than did any of the electrochemical methods (Table 2). Based on mass loss, the 80°C electrode showed only 1.65 mpy (0.0419 mm/y) vs. 31 to 160 mpy (0.79 to 4.1 mm/y) for the electrochemical techniques for the 0.56 m/m solution. Corrosion rates of around 89 mpy (2.3 mm/y) were obtained at 140°C based on electrode mass loss; far lower than predicted from the electrochemical methods, which showed as much as 10,521 mpy (267 mm/y). Clearly, there is a substantial divergence in the results such that the electrochemical methods are producing very high corrosion rates. This indicates there is interference by one or more electroactive redox species contributing to Faradaic current and is being interpreted as corrosion. Within the dataset of the electrochemical methods, the corrosion rates are somewhat inconsistent. The EIS method produced higher rates compared to LPR at 140°C but similar rates at 80°C. Meanwhile, the ECN technique produced higher rates at 80°C but lower rates at 140°C. Potentiodynamic results Four potentiodynamic tests were performed to determine the Tafel coefficients and evaluate the anodic and cathodic reactions involved in the corrosion process. Figures 8 and 9 show the scans for the 0.46 to 0.47 m/m and 0.55 to 0.56 m/m ratios, respectively. Both figures compare the results obtained at 80°C and at 140°C. The anodic and cathodic scans were run separately using duplicate electrodes in each test environment with the respective scans starting from the zero current potential. Thus, any films that had formed on the electrodes, passive or otherwise, were not destroyed prior to the respective scan. The Tafel values generated from the four potentiodynamic scans are shown in Table 3.

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Also shown in Table 3 are the corrosion current densities and corrosion rates. As with the other electrochemical techniques (Table 3), the corrosion rates are much higher at 140°C than at 80°C. All are high relative to the electrode mass loss. Of interest from Figures 8 and 9 is that the corrosion potentials2 are very similar irrespective of loading and there is nearly no temperature effect. At 80°C, the corrosion potential was -559 mV vs. SCE for loadings of 0.46 to 0.47 m/m and -562 mV for loadings of 0.55 to 0.56 m/m. At 140°C, potentials of -527 mV vs. SCE were observed for both loading levels. Distinct slope changes were observed in all the cathodic scans shown in Figures 8 and 9. A reasonably smooth curve would be expected for a single cathodic reaction, e.g. hydrogen reduction. The changes in slope and current during the cathodic scans indicate that one or more reduction reactions are occurring on the surface of the 1018 carbon steel electrodes. Coupon results Coupon testing was performed to evaluate rich acid gas loadings with H2S at 80°C for comparison to the electrochemical results. Higher temperature testing at 120°C and 140°C was performed with lean loadings to correlate with conditions in the Stripper reboiler. The results are shown for all three temperatures in Figures 10, 11 and 13 for both H2S and a combined acid gas composition of H2S+CO2. Coupon tests at 80°C. The coupon results at 80°C (Figure 10) show corrosion rates increasing from 0.37 to 3.33 mpy (0.0049 to 0.0846 mm/y) as the H2S loading is increased from 0.50 m/m to 1.00 m/m when tested without CO2. These corrosion rates are much lower than observed by any of the electrochemical methods (Tables 2 and 3). Addition of CO2 increased the corrosion rates moderately compared to the fluids containing H2S alone. The corrosion rates are seen to correlate reasonably well with the total acid gas loadings. This implies that the total acid gas loading determines the corrosivity when adding CO2 to DEA/H2S at 80ºC. Coupon tests at 120°C and 140°C. The coupon tests performed at 120°C and 140°C simulated lean acid gas loadings in the stripper-reboiler areas with H2S only and with H2S+CO2. The acid gas loadings ranged from 0.05 to a 0.40 m/m ratio. As shown in Figure 11, increasing the H2S loading at 120ºC shows a more substantial increase in corrosion rates than observed at 80ºC. The decrease in corrosion rates at the highest loading levels at 120°C may be due to increasing protectiveness of iron sulfide film on the coupons. The adherence of the surface film appears more intact and protective for the 0.32 m/m test compared to the wrinkled, loose films observed on coupons exposed to 0.19 and 0.25 m/m acid gas loadings (Figure 12). At 140°C, the corrosion rates also increase with acid gas loading but not as rapidly as observed at 120°C. Also, the corrosion rate did not fall off at the highest loading level at 140°C as it did at 120°C. Comparing the corrosion rates from the coupon study to those from the electrochemical tests show that for H2S loadings of 0.5 to 0.6 m/m at 80°C, the coupons exhibit 0.37 to 0.61 mpy (0.0094 to 0.015 mm/y) while the electrochemical tests showed rates of 22 to 166 mpy (0.56 to 4.22 mm/y). The loading levels tested at 140°C don't overlap, but for 0.30 to 0.32 m/m H2S loading, the coupon results show 8.31 to 17.75 mpy (0.211 to 0.451 mm/y) at 120°C and 140°C, respectively. These corrosion rates are much lower than the rates of 740 to 10,521 mpy (19 to 267 mm/y) observed by the electrochemical methods at loadings of 0.47 to 0.55 m/m H2S at 140°C.

2 More accurately called zero current potentials or redox potentials for the case of solution redox species.

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All of the electrochemical methods are producing much higher rates than observed via coupons, indicating that an interfering electron transfer reaction is contributing to the currents measured by the electrochemical techniques. Mechanism investigation and discussion Several electrochemical tests were performed to elucidate the redox reaction that interferes with the corrosion rate measurements. Using a glass cell, a potentiodynamic scan was performed on 1018 carbon steel in DEA not containing any acid gas. The pH of the DEA fluid was adjusted to 9.75 (same pH as measured for the H2S loaded solutions) using hydrochloric acid. This test was performed at 80°C and the result is compared to the scans obtained with 0.46 m/m H2S loading at 80°C and at 140°C (Figure 14). All three scans in Figure 14 have essentially the same corrosion potential. However, addition of H2S at 80°C produces significantly higher anodic and cathodic currents with the resultant increase in the corrosion current density. The effect is quite substantial, with a three order increase in observed corrosion current density. Increasing the temperature to 140°C increases the corrosion current density by yet another order of magnitude, primarily due to increased cathodic reaction kinetics. As presented in Table 3, the corrosion current density of 0.337 mA/cm2 obtained from the potentiodynamic scan for 0.46 m/m H2S loading at 80°C (Figure 14) corresponds to a corrosion rate of 154 mpy (3.91 mm/y). However, the coupon results (Figure 10, 0.50 m/m H2S loading) show that the real corrosion rate under similar conditions is only 0.37 mpy (0.0094 mm/y). Considering the relatively low corrosion rates obtained by the mass loss tests, the currents measured in the potentiodynamic scans with the addition of H2S must be primarily from electron transfer reactions other than corrosion. Specifically, the dissolution of iron contributes only a small portion to the anodic currents measured in the presence of H2S in Figure 14. The corrosion potentials observed in the potentiodynamic tests, discussed above for Figure 8, were -559 mV vs. SCE for loadings of 0.46 to 0.47 m/m and -562 mV for loadings of 0.55 to 0.56 m/m at 80°C. This is higher than usually observed for corrosion carbon steel, which tends to be around -850 mV depending on conditions. Instead, the observed corrosion potentials point to the sulfur/bisulfide redox couple (Equation 3). S + H2O + 2e- → HS- + OH- (3) The standard reduction potential for the sulfur/bisulfide redox couple is -478 mV vs. NHE(17) or -719 mV vs. SCE. Using the Nernst equation(18,19), the reversible potential for the sulfur/bisulfide redox couple can be corrected for temperature and concentrations. E = Eo + RT/nF ln(1/[HS-][OH-]) (4) The concentration of bisulfide in the test solutions is approximately 1.66 M. This is based on an H2S /DEA ratio of 0.5 m/m and the expectation(6) that all of the H2S exists as bisulfide in the DEA solution at pH = 9.75. Solving the Nernst equation for 80°C (353 K), [OH-] = (5.62 x 10-5 M) for pH = 9.75, and [HS-] = 1.66 M gives a redox potential of -578 mV vs. SCE. This is within 20 mV of the observed potentials for carbon steel in the DEA/ H2S test solutions at 80°C.

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Based on the evidence that the electrochemical reactions are strongly affected by a solution based redox couple and the proximity of the sulfur/bisulfide redox potential to the observed test potentials, it is postulated here that the interfering electrochemical reaction involves the oxidation of bisulfide to elemental sulfur (reverse of the reaction shown in Equation 3). The oxidation reaction is likely converting bisulfide from solution to elemental sulfur or iron sulfide on the electrode surface. Additional tests were performed to investigate the effects of the bisulfide oxidation on the electrochemistry. Potentiodynamic scans were run on 1018 carbon steel in DEA containing relatively small concentrations of bisulfide. In order to do this in a non-pressurized glass cell, sodium sulfide (Na2S) was added to the DEA and the pH adjusted to 9.75 with hydrochloric acid, thus converting the sulfide to bisulfide and achieving the same pH as the DEA fluids loaded with 0.55 m/m H2S. Figure 15 shows the results obtained using separate 1018 carbon steel electrodes for the anodic and cathodic potentiodynamic scans, each initiated from the corrosion potential. The addition of bisulfide to the test solution had a significant effect on the anodic scan, increasing rapidly with relatively little potential change. This indicates that very fast oxidation kinetics are associated with the presence of bisulfide. The 1018 carbon steel electrode used for the anodic scan with 1,000 ppm bisulfide produced substantial scale buildup, as shown in Figure 16. Apparently, the current peak in the anodic scan from -407 to -332 mV involves the formation of an iron sulfide film. This was not observed on electrodes used in the autoclave experiments in Figures 3, 8, 9 and 14. The cathodic scans shown in Figure 15 are very similar. This indicates essentially no affect of the bisulfide on the cathodic reaction, at least at this low concentration of 1,000 ppm. The corrosion current density also shows little change with the addition of bisulfide. Compared to the higher concentrations of bisulfide (Figures 8 and 9), the scans at 1,000 ppm in Figure 15 show that the cathodic reaction is limiting the observed corrosion current density. This is probably due to the low concentration of bisulfide (1,000 ppm) compared to that present in the 0.46 and 0.56 H2S/DEA m/m tests. The presence of bisulfide affects the anodic reaction more strongly, at lower concentrations, than it affects the cathodic reaction. Additional tests were performed under similar conditions, with 1,000 ppm bisulfide, except this time the potentiodynamic scan was initiated at the negative extreme (-1.0 V vs. Ecorr) instead of starting from Ecorr. Also, the scan was carried through Ecorr into the anodic region until reaching +1.0 V vs. Ecorr (Figure 17). At that point, the scan was reversed and the potential scanned back to -1.0 V vs. Ecorr. The reverse scan showed a much higher corrosion current density due to both increased anodic current and, especially, increased cathodic current. This indicates that the forward scan changed the electrode surface to make it more electroactive on the reverse scan. It is believed that bisulfide from solution was oxidized to elemental sulfur during the anodic portion of the forward scan. The resultant elemental sulfur deposited onto the electrode surface and added additional anodic current during the reverse scan, noticeable over the potential range from about -0.3 V vs. SCE to Ecorr. After passing through Ecorr, the elemental sulfur may have been reduced back to bisulfide during the cathodic portion of the scan. There are alternative possibilities, however. One such possibility is that the electrode surface is now more catalytically active toward reduction of solution species, namely protonated DEA to form hydrogen gas, due to the new surface film. In order to eliminate the contribution of iron oxidation to the anodic current scans, additional tests were performed using a graphite electrode. Figure 18 shows the results from potentiodynamic scans on graphite starting from the negative potential extreme. As with carbon steel, the solutions were DEA with and without added sodium sulfide; pH adjusted to 9.75.

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The potentiodynamic scans on graphite (Figure 18) show very little anodic activity without the presence of bisulfide. However, the addition of 1,000 ppm bisulfide produces significant current after the potential passes through Ecorr. This shows the electroactivity of bisulfide toward oxidation and indicates that the current observed in Figure 17 during application of anodic potential oxidized bisulfide, most likely to elemental sulfur. Additional tests were performed to investigate the behavior of the cathodic scans on carbon steel in DEA containing 0.55 m/m H2S. Those scans (Figure 8) showed specific regions where the current change implied that the reaction mechanisms were changing. To explore this, a potentiodynamic test was performed at a 0.55 m/m H2S/DEA at 80°C using a slow scan rate (0.1 mV/sec instead of 1 mV/sec). The carbon steel electrode was first scanned from the zero current potential to -2.10 V and then from -1.56 V back to the zero current potential (Figure 19). Both scans show a similar current at potentials more negative than -1.1 V versus SCE, which can be interpreted as water reduction. Between -560 mV to -1.1 V there are mixed cathodic reactions believed to include current contributions from the reduction of hydrogen, bisulfide, elemental sulfur or other sulfur species such as polysulfides. During the return scan, a higher current density was observed above -1.2 V compared to the forward scan. The cathodic current is seen to hold steady over a 100 mV potential range near -1.0 V. The current density is very high in this region (>10 mA/cm2) and may represent the diffusion limited current density for bisulfide reduction at this H2S loading and temperature. The elevated current seen near -1.0 V on the return scan was further explored by performing a potentiostatic test on the electrode in the same system. The potential was held at -400 mV vs. Ecorr and the current was observed to maintain -60 to -65 mA/cm2 for 24 hours. Such a high current maintained for this duration of time cannot be due to reduction of a thin surface film. Instead, it is likely due to reduction of solution species such as bisulfide or protonated DEA.

CONCLUSIONS

• All of the electrochemical methods tested, including LPR, EIS and ECN, produced overly high corrosion rates for carbon steel in the presence of H2S at loadings of 0.46 to 0.56 m/m in DEA. The corrosion rates determined from coupon mass loss and electrode mass loss were much lower.

• An interfering redox couple, namely bisulfide/elemental sulfur, is believed to be responsible for

the high currents measured by the electrochemical techniques and interpreted as corrosion.

• The corrosion potential of the bisulfide/elemental sulfur redox couple is close to that of corroding iron. This makes it less obvious that an interfering reaction is taking place.

• Bisulfide was observed to be electroactive. This was determined using a graphite electrode.

• The oxidation of bisulfide is believed to form elemental sulfur on the electrode surface. This

process appears to increase the electroactivity of the surface, as observed for carbon steel in 1,000 ppm bisulfide. The product formed during the anodic scan produces significantly higher currents during the return scan from peak anodic potential.

• When LPR scans are performed on an electrode, the bisulfide to sulfur reaction is believed to

take place during application of anodic potential, even for overpotentials as little as 10 mV.

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• Sulfur and iron sulfide can exist in many different stoichiometries and oxidations states. The overall chemistry of the surface species could be quite complex.

• It may still be possible to discern the polarization resistance, and therefore the corrosion rate,

from the EIS scans. Otherwise, caution should be used when performing corrosion monitoring in H2S systems using methods based on current flow, which produce misleadingly high corrosion rates. Mass or thickness loss techniques, such as electrical resistance probes or inductive resistance probes should be considered although the buildup of conductive iron sulfide films can affect those on-line techniques.

REFERENCES

1. A. Kohl and R. Nielsen, Gas Purification, Fifth Ed., Gulf Publishing, 1997; pp. 40-56, 187-224. 2. M. Pourbaix, Atlas of Electrochemical Equilibria in Aqueous Solutions, Second English Ed.,

NACE and CEBELCOR, 1974; pp. 307-321. 3. Foroulis, Corrosion Prevention & Control, p. 84, Aug. 1993. 4. Nielsen, Lewis, Hansen, and McCullough, CORROSION/95, paper #571. 5. Liu, Dean and Bosen, CORROSION/95, paper #572. 6. Rooney, Bacon and DuPart, �Effect of heat stable salts on solution corrosivity of MDEA-based

alkanolamine plants. Part III� Laurance Reid conference March 1997. 7. Cummings, Veatch and Keller, CORROSION/97, paper #341. 8. Cummings, Veatch and Keller, Materials Performance, Jan. 1998, p. 42. 9. D. Jones, Principles and Prevention of Corrosion, Macmillan Publ., New York, 1992; p. 147. 10. N. Thompson and J. Payer, DC Electrochemical Test Methods, Vol. 6 of "Corrosion Testing

Made Easy" NACE Intl. 1998; pp. 6 and 66. 11. Scully, CORROSION/98, paper #304. 12. Agrawal, Durr and Koch, CORROSION/04, paper #04383. 13. Kvarekval, Nyborg and Choi, CORROSION/03, paper #03339. 14. Ramanarayanan and Smith, Corrosion, 46 (1), 66 (1990). 15. Vedage, Ramanarayanan, Mumford and Smith, Corrosion, 49 (2), 114 (1993). 16. Huang, Tsai and Lee, Corrosion, 52 (9), 708 (1996). 17. Handbook of Chemistry and Physics, 61st Ed. CRC Press, p. D-155. 18. J. Bockris and A. Reddy, Modern Electrochemistry, Plenum Press, New York, 1977; p. 901. 19. A. Bard and L. Faulkner, Electrochemical Methods: Fundamentals and Applications, John Wiley

& Sons, New York, 1980; p. 51.

11

TABLE 1

Elemental analysis of the 1018 carbon steel coupons and a photograph of an electrochemical sample and a coupon. Balance of the composition is Fe.

Element Weight % Element Weight %

Al 0.030 V 0.003 Cu 0.080 C 0.190 N 0.007 Nb 0.004 S 0.001 Si 0.020

Ca 0.001 Ni 0.040 Mn 0.670 Sn 0.008 Cr 0.030 Mo 0.020 P 0.003 Ti 0.004

TABLE 2

Summary of the measured corrosion rates, mpy (mm/y), from the LPR, EIS, ECN and 1018 carbon steel electrode mass loss (exposures of 3 to 4 days).

Test Method 80°C

0.46 H2S/DEA 80°C

0.56 H2S/DEA 140°C

0.47 H2S/DEA 140°C

0.55 H2S/DEA LPR 33 (0.84) 31 (0.79) 2,083 (52.91) 1,959 (49.76) EIS 37 (0.94) 33 (0.84) 10,521 (267.23) 4,419 (112.2) ECN 266 (6.76) 160 (4.1) 740 (19) 1,093 (27.76)

Electrode Mass Loss N/A 1.65 (0.0419) 83.5 (2.12) 93.9 (2.39)

TABLE 3 Summary of the Tafel coefficients, corrosion current and corrosion rates, mpy (mm/y), extrapolated

from the potentiodynamic scans in the four electrochemical autoclave test environments.

H2S/DEA Mole Ratio

Anodic Tafel (βa)

Cathodic Tafel (βc)

icorr mA/cm2 Corrosion Rate

0.46 (80°C) 78 mV/Decade 122 mV/Decade 0.337 154 (3.91) 0.56 (80°C) 27 mV/Decade 124 mV/Decade 0.0456 21.8 (0.554) 0.47 (140°C) 88 mV/Decade 91 mV/Decade 3.42 1560 (39.6) 0.55 (140°C) 95 mV/Decade 91 mV/Decade 3.58 1630 (41.4)

12

Gas In

Thermocouple

Heater

Gas Out

140 C

Temperature Controller

Magnetic Stirrer

Test Coupons

Sample Bomb

FIGURE 1 - A schematic of the autoclave setup used for the coupon testing and a picture of the autoclave cap, with three test electrodes, used in the electrochemical testing.

Thermocouple

Nitrogen In

Gas Out

Magnetic Stirrer

Working Electrode

Luggin Probe - SCE

Counter Electrode (Graphite)

80 C

Heating Tape

FIGURE 2 - A schematic of the glass reaction kettle test cell.

13

0 5 10 15 20 25Time (hours)

1

10

100

1000

10000

Cor

rosi

on R

ate

(mpy

)

0.46 H2S/DEA mole ratio at 80C0.56 H2S/DEA mole ratio at 80C

0.47 H2S/DEA mole ratio at 140C0.55 H2S/DEA mole ratio at 140C

FIGURE 3 - Comparison of the LPR corrosion rates measured in the four electrochemical autoclave

test environments.

FIGURE 4 - Nyquist plot for the 0.46 H2S/DEA m/m ratio test at 80°C.

0 10 20 30 40 50 60Real Impedance (ohms)

0

5

10

15

20

25

Imag

inar

y Im

peda

nce

(ohm

s)

14

0 5 10 15 20 25 30 35Real Impedance (ohms)

0

2

4

6

8

10

12

14

Imag

inar

y Im

peda

nce

(ohm

s)

FIGURE 5 - Nyquist plot for the 0.56 H2S/DEA m/m ratio test at 80°C.

0.75 0.8 0.85 0.9 0.95 1Real Impedance (ohms)

0

0.01

0.02

0.03

0.04

0.05

Imag

inar

y Im

peda

nce

(ohm

s)

FIGURE 6 - Nyquist plot for the 0.47 H2S/DEA m/m ratio test at 140°C.

0.6 0.7 0.8 0.9 1 1.1 1.2 1.3Real Impedance (ohms)

0

0.05

0.1

0.15

0.2

Imag

inar

y Im

peda

nce

(ohm

s)

FIGURE 7 - Nyquist plot for the 0.55 H2S/DEA m/m ratio test at 140°C.

15

1E-6 1E-5 0.0001 0.001 0.01 0.1 1Current Density (amps/cm^2)

-2

-1.5

-1

-0.5

0

0.5

1

Volts

vs

SCE

0.46 H2S/DEA mole ratio at 80C 0.47 H2S/DEA mole ratio at 140C

FIGURE 8 - Potentiodynamic scans generated for 1018 carbon steel at 80 and 140°C at a 0.46 and 0.47 H2S/DEA m/m ratio, respectively.

1E-7 1E-6 1E-5 0.0001 0.001 0.01 0.1 1Current Density (amps/cm^2)

-2

-1.5

-1

-0.5

0

0.5

1

Volts

vs

SCE

0.56 H2S/DEA mole ratio at 80C 0.55 H2S/DEA mole ratio at 140C

FIGURE 9 - Potentiodynamic scans generated for 1018 carbon steel at 80 and 140°C at a 0.55 and 0.56 H2S/DEA m/m ratio, respectively.

16

0.370.95

0.611.68

1.641.96

3.33

0 2 4Corrosion Rate (mpy)

H2S ONLYH2S + CO2

0.50 mole ratio0.68 mole ratio

0.60 mole ratio0.82 mole ratio

0.70 mole ratio0.94 mole ratio

1.00 mole ratioAcid

Gas

/ D

EA M

ole

Rat

io

FIGURE 10 - The average corrosion rates of the duplicate 1018 carbon steel coupons are shown for

each test environment at 80°C. The tests performed with CO2 were at a 3:1 H2S/CO2 m/m ratio.

0.840.6

3.955.46

15.7416.05

8.316.8

0 2 4 6 8 10 12 14 16 18Corrosion Rate (mpy)

H2S ONLYH2S + CO2

0.05 mole ratio0.07 mole ratio

0.12 mole ratio0.19 mole ratio

0.19 mole ratio0.25 mole ratio

0.32 mole ratioAci

d G

as /

DEA

Mol

e R

atio

0.40 mole ratio

FIGURE 11 - The average corrosion rates of the duplicate 1018 carbon steel coupons are shown for

each test environment at 120°C. The tests performed with CO2 were at a 3:1 H2S/CO2 m/m ratio.

17

FIGURE 12 - A comparison of the surface films formed at 120°C from tests at 0.19 H2S/DEA m/m (left), at 0.25 (H2S+CO2)/DEA m/m (center) and the 0.32 H2S/DEA m/m test (right).

0.720.96

1.842.97

10.474.1

17.7511.92

0 2 4 6 8 10 12 14 16 18 20Corrosion Rate (mpy)

H2S ONLYH2S + CO2

0.05 mole ratio0.08 mole ratio

0.13 mole ratio0.17 mole ratio

0.18 mole ratio0.25 mole ratio

0.30 mole ratioAci

d G

as /

DEA

Mol

e R

atio

0.39 mole ratio

FIGURE 13 - The average corrosion rates of the duplicate 1018 carbon steel coupons are shown for

each test environment at 140°C. The tests performed with CO2 were at a 3:1 H2S/CO2 m/m ratio.

18

1E-7 1E-6 1E-5 0.0001 0.001 0.01 0.1 1Current Density (amps/cm^2)

-2

-1.5

-1

-0.5

0

0.5

1

Volts

vs

SCE

0.46 DEA/H2S mole ratio at 80C0.47 DEA/H2S mole ratio at 140C

DEA - no bisulfide at 80C

FIGURE 14 - A comparison of the potentiodynamic scan performed in the DEA test solution without any sulfur species and two potentiodynamic scans performed with H2S loadings.

1E-7 1E-6 1E-5 0.0001 0.001 0.01 0.1 1Current Density (amps/cm^2)

-1.5

-1

-0.5

0

Volts

vs

SCE

No bisulfide at 80C 1000 ppm bisulfide at 80C

FIGURE 15 - A comparison of the potentiodynamic scans performed with and without 1000 ppm

bisulfide (added as Na2S) in the DEA test solution (pH 9.75) at 80°C.

19

FIGURE 16 - Electrode appearance upon removal from the DEA test solution containing 1000 ppm bisulfide at 80°C.

-2.00

-1.50

-1.00

-0.50

0.00

0.50

1.E-06 1.E-05 1.E-04 1.E-03 1.E-02 1.E-01 1.E+00

Amps

Pote

ntia

l -vs

- SC

E

DEA Scan (1018CS Blank, No Sulfide)

DEA Scan with sulfide -80C - (-1.0V to E apex)

DEA Scan with sulfide -80C -(return E apex to -1.0V)

FIGURE 17 - Potentiodynamic scans performed on 1018 carbon steel in DEA with and without 1000 ppm bisulfide (added as Na2S) at 80°C.

20

-1.6

-1.4

-1.2

-1

-0.8

-0.6

-0.4

-0.2

0

0.2

1.E-09 1.E-08 1.E-07 1.E-06 1.E-05 1.E-04 1.E-03 1.E-02 1.E-01 1.E+00

Current (amps)

Pote

ntia

l (vo

lts v

s SC

E)

Graphite- No Sulfide (-1.4V to 0.00V vs SCE)

Graphite -1000ppm S (-1.4V to 0.00V vs SCE)

FIGURE 18 - Potentiodynamic scans performed on graphite in DEA with and without 1000 ppm bisulfide (added as Na2S) at 80°C. For both tests, the scan was started at the negative potential extreme and scanned anodically to the positive potential extreme. Currents are not divided by the electrode surface area in this figure.

-2.50

-2.00

-1.50

-1.00

-0.50

0.00

1.000E-06 1.000E-05 1.000E-04 1.000E-03 1.000E-02 1.000E-01 1.000E+00Current Density (amps/cm^2)

Volts

vs

SCE

at 8

0C

0.55 H2S/DEA mole ratio at 80C - Scan from Ecorr to -2.10 V

0.55 H2S/DEA mole ratio at 80C - Scan from -1.56V to -0.51 V

FIGURE 19 - Cathodic polarization scans ran at a 0.55 H2S/DEA m/m ratio at 80°C using a single 1018

carbon steel electrode first from Ecorr to -2.10 V and then from -1.56 V back to Ecorr.

Start of Water Reduction

Hydrogen, Bisulfide and Sulfur Reduction

21