Laboratory Evaluation and Mechanistic Understanding of the Impact of

Jul 30, 2018 - However, the presence of iron species in oilfield produced water can considerably impair the performance of scale inhibitors. To date, ...
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Laboratory Evaluation and Mechanistic Understanding of the Impact of Ferric Species on Oilfield Scale Inhibitor Performance Zhang Zhang,†,‡,∥,⊥ Ping Zhang,*,§,⊥ Zhejun Li,§ Amy T. Kan,†,‡ and Mason B. Tomson†,‡ Department of Civil and Environmental Engineering and ‡Nanosystems Engineering Research Center for Nanotechnology-Enabled Water Treatment, Rice University, Houston, Texas 77005, United States § Department of Civil and Environmental Engineering, Faculty of Science and Technology, University of Macau, Taipa, Macau, China Downloaded via DURHAM UNIV on August 5, 2018 at 06:30:41 (UTC). See https://pubs.acs.org/sharingguidelines for options on how to legitimately share published articles.

†

S Supporting Information *

ABSTRACT: Scale inhibitor chemicals are widely used in oilfield operations for mineral scale control. However, the presence of iron species in oilfield produced water can considerably impair the performance of scale inhibitors. To date, few studies have been conducted to experimentally investigate the mechanism of iron effect on scale inhibitors. Although Fe(II) is the major form of iron species in oilfield produced water, Fe(III) can be formed in produced waters due to oxidation of Fe(II). In this study, Fe(III) effect on various scale inhibitors was evaluated by examining the inhibitor performance to control barium sulfate (barite) scale formation. This study finds that Fe(III) can significantly impair the performance of both phosphonate and polymeric inhibitors with an iron concentration below 1 mg L−1. Moreover, the mechanism of the influence of Fe(III) on scale inhibitors was studied by investigating the adsorption capacity of ferric hydroxide solid of phosphonate scale inhibitor and also examining the efficacy of the unadsorbed inhibitor in aqueous solution. It can be concluded that the Fe(III) impact on phosphonate inhibitor is due to the adsorption of inhibitor to the surface of ferric hydroxide solids. Furthermore, two common chelating chemicals (EDTA and citrate) were tested for their effects in reversing the adverse impact of Fe(III) on scale inhibitor. Experimental results suggest that citrate is more effective than EDTA in reversing the detrimental impact of Fe(III) despite the fact the EDTA is a stronger chelating agent. The mechanisms of these two chelating chemicals in terms of interacting with Fe(III) were discussed and compared. This study provides the theoretical basis and technical insights for oilfield iron control to minimize iron impairment on scale inhibitor performance.

1. INTRODUCTION Together with corrosion and gas hydrate, mineral scale deposition is among the top three water-associated production chemistry challenges for oilfield operations.1,2 Mineral scale deposition phenomenon stems from supersaturation of an aqueous solution with respect to an inorganic mineral at a given physicochemical condition. Mineral scale solid will precipitate from the supersaturated solution and subsequently deposit on the surface of the surrounding environment.3−5 Deposited scale solids can block the production tubing and flow line, resulting in a considerable reduction in the throughput of tubing or pipeline. Moreover, scale can deposit on the surface of heat exchanger changing the thermal conductivity of the equipment and also affect the efficiency of oil−water separator leading to discharge of hydrocarbons into the environment.6 One of the most commonly observed mineral scales in many industrial systems, especially oilfield operations, is barium sulfate (barite). Barite formation in oilfield is typically as a result of mixing of incompatible waters, such as seawater with formation water during water flooding operations. The high concentration (can be 3000 mg L−1 and above) of sulfate ion from seawater can precipitate with the barium ion in formation water at the near-wellbore formation of the injection well in the absence of scale prevention measures.7 Supersaturation level as well as scale threat can be mathematically predicted by calculating the saturation index (SI) of a solution with respect to a specific mineral.8 SI at a © XXXX American Chemical Society

certain temperature and pressure condition can be calculated via eq 1: ij IAP yz zz SIT , P = logjjjj z j K sp zz k {

(1)

where SIT,P represents the saturation index at a given temperature and pressure condition. IAP denotes the ion activity product, and Ksp corresponds to the solubility product of a mineral at a certain temperature and pressure condition. If the calculated SI is greater than zero, this suggests that the solution is supersaturated with respect to a mineral of interest. Furthermore, scaling tendency can be evaluated via scale deposition kinetics by measuring the induction time. Induction time is the time period between the attainment of supersaturation and the appearance of crystals.9,10 Experimentally, induction time can be measured as the elapsed time period between the onset of mixing cationic and anionic solutions and when the scale crystals start to appear in solution. The longer the induction time, the longer it takes to precipitate scale solids from aqueous solutions.8 In oilfield operations and other industries, scale threat control mainly relies on the delivery of scale inhibitor chemical into the production system to inhibit Received: May 26, 2018 Revised: July 24, 2018 Published: July 30, 2018 A

DOI: 10.1021/acs.energyfuels.8b01837 Energy Fuels XXXX, XXX, XXX−XXX

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Energy & Fuels scale deposition. Scale inhibitors are a group of specialty chemicals capable of delaying scale deposition kinetics. Scale inhibitors fall into the category of threshold inhibitors. Different from chelating agents forming complexes with metals stoichiometrically, scale inhibitors can effectively control scale threat at a very low inhibitor concentration of several milligrams per liter or lower.2,6−8 The presence of scale inhibitors can substantially increase the induction time, allowing the produced fluid to travel from reservoir to topside processing facilities without scale threat. Common oilfield scale inhibitors include phosphonate and polymeric inhibitors. Phosphonate inhibitors typically function as scale crystal growth inhibitors by shielding the active sites of crystal surface from exposure to crystal-forming ions. Polymeric inhibitors prevent scale formation by disturbing scale crystal nucleation process.2,6,8 A number of previous studies have found that the performance of scale inhibitors can be affected by many factors, such as system temperature, solution pH, and the presence of foreign species, such as iron.11 As the fourth most abundant element in the earth’s crust, the element of iron exists extensively not only in the natural environment but also in various industrial systems, such as oilfield produced waters.1,2,6,8 Depending on the geochemistry of the producing formation, the type of hydrocarbons produced, and the characteristics of the production wells, ion species concentration in oilfield produced water can vary from several milligrams per liter to hundreds of milligrams per liter.1,2 Guerra et al. reported a total iron concentration of up to 1100 mg L−1 in conventional produced waters and of up to 258 mg L−1 in unconventional produced waters by reviewing numerous studies and the U.S. Geological Survery (USGS) produced water database.12 Iron species in oilfield produced waters have two major sources. One is from the prolonged contact with iron-bearing minerals, such as siderite (FeCO3), pyrite (FeS2), and ankerite (Ca(Mg,Fe)(CO3)2) in reservoir formation.13 During acid stimulation and CO2 flooding treatment, reservoir fluid can become acidic, which accelerates the dissolution of these minerals from formation, resulting in an elevated concentration of iron dissolved in the produced waters. The other source is from corrosion of the pipelines. Under anaerobic condition, iron can be oxidized to ferrous or Fe(II) state through reduction of water molecules, which generates hydrogen.14 Typically during the hydrocarbon production process, the production system including production tubing and flow lines remains an anaerobic condition before the fluid reaching the surface. Thus, the majority of the iron species in oilfield produced water should remain in the form of Fe(II). Although Fe(II) is the major form of Fe in produced water, magnetite (Fe3O4) can be occasionally found in the producing wells, suggesting that ferric or Fe(III) can form in a downhole reservoir.1 The sources of Fe(III) in formation waters include the following: 1. The first source is attack of Fe(II) by aqueous cation (hydronium, H3O+) in formation water before it is in contact with the atmosphere.14,15 The responsible chemical reaction is

Fe 2 + +

1 1 O2 + 2OH− + H 2O → Fe(OH)3 4 2

Fe(III) is a transition metal with an extremely low solubility in water and can be easily hydrolyzed into ferric hydroxide solids at near-neutral solution pHs.14 Ferric hydroxide minerals have a large surface area and a considerable adsorption capability for organic anions including scale inhibitor functional groups.14,17 The existence of iron species in solution can have multiple effects on produced water chemistry. When pH rises at downhole conditions, such as when acid in the treatment fluid becomes spent, iron will precipitate and cause formation damage and reduce the effectiveness of acidizing operation.6 The release of Fe(III) can cause asphaltic sludging and formation damage. Asphaltic sludging is formed by the reaction of the iron with polar groups in asphaltenes.2,18 Furthermore, Fe(III) has a strong interaction with common oilfield scale inhibitor functional groups. Several studies on Fe(III) effect on scale inhibitors have found a detrimental effect of Fe(III) on scale inhibitor performance. However, the mechanism of this effect has not been thoroughly investigated. Stoppelenburg et al. found that Fe(III) significantly impaired phosphonate inhibitor performance, and the authors proposed that the reason could be the adsorption of phosphonate inhibitor onto ferric hydroxide particles.19 Kelland found Fe(III) has a detrimental effect on three common scale inhibitors. The author speculated that the reason for the detrimental effect was the complexation between Fe(III) with inhibitor function groups.20 Amjad evaluated the antagonistic effect of iron oxide rust on the performance of calcium phosphate inhibitor and proposed several possible mechanisms responsible for this effect.21 In this study, the mechanism of Fe(III) effect on scale inhibitor performance is systematically investigated. In order to control the adverse effect of iron on scale inhibitors, one of the most common iron control methods in oilfields is to employ chelating agents to form complexes with iron. Ethylenediaminetetraacetic acid (EDTA), citrate, and nitrilotriacetic acid (NTA) are among the commonly used iron complexing agents.2 A number of studies have been carried out using EDTA, NTA, citrate, or their blend for iron control in acid stimulation conditions.22−24 EDTA has a much higher stability constant with iron than NTA or citrate, but has a low solubility in acidic conditions. Another problem with EDTA is that it is not readily biodegradable. NTA is acid soluble and biodegradable but has a lower stability constant than EDTA. Citrate is acid-soluble and environmentally friendly, but it might form calcium citrate precipitate at high Ca 2+ concentrations. Shen et al. adopted citrate to sequestrate iron and found that the presence of citrate could significantly improve inhibitor performance. These authors attributed the reason to be the prevention of the formation of iron carbonate seed particles, which can promote scale crystal growth in the brine.25 Although the presence of iron can have a substantial impact on scale inhibitor performance, little research has been conducted so far to experimentally investigate the mechanism of iron effect on scale inhibitors. In this paper, the effect of Fe(III) on both phosphonate and polymeric inhibitors was systematically evaluated by measuring the induction time of barite at different experimental conditions. Moreover, the mechanism of Fe(III) on phosphonate inhibitor was studied by evaluating the adsorption capacity of ferric hydroxide of

Fe2 + + 3H3O+ + e− → Fe3 + + 3OH− + 3H 2↑

2. Once the formation waters are in contact with atmospheric oxygen as a result of pump leakage and/or when formation waters flow close to the surface, Fe(II) species can be oxidized by molecular oxygen via the following reaction:16 B

DOI: 10.1021/acs.energyfuels.8b01837 Energy Fuels XXXX, XXX, XXX−XXX

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Energy & Fuels phosphonate inhibitor. Furthermore, EDTA and citrate were tested and compared for their abilities to reverse the detrimental effect of Fe(III) on scale inhibitors. This study provides the theoretical basis and technical insights for oilfield iron control to minimize iron impairment on scale inhibitor performance.

Figure 1. Schematic of the apparatus to measure barite induction time.

2. MATERIALS AND METHODS 2.1. Chemicals. Commercial grade diethylenetriamine pentakis(methylenephosphonic acid) (DTPMP) with 50% activity (w/w) and poly(vinylsulfonate) (PVS) with 25% activity (w/w) were purchased from Sigma-Aldrich and used as scale inhibitors. Phosphino poly(carboxylic acid) (PPCA) with 50% activity (w/w) was purchased from BWA Water Additives and used as scale inhibitor. The molecular structures of DTPMP, PPCA, and PVS are shown in Figure S1 in the Supporting Information. Chemicals such as ferric chloride, calcium chloride, sodium chloride, nitric acid, hydrochloric acid, sodium hydroxide, sodium sulfate, barium chloride, sodium citrate dihydrate, ethylenediaminetetraacetic acid (EDTA), and piperazine-1,4-bis(2-ethanesulfonic acid) sodium salt (PIPES) were reagent grade and purchased from Fisher Scientific. PIPES was used as a chemical buffer to control the solution pH. The presence of PIPES buffer is not expected to materially impact barite scale deposition, as illustrated in Figure S2 in the Supporting Information. Citrate and EDTA were used as chelating chemical agents to form complexes with ferric ion. Deionized water (DI water) was prepared by reverse osmosis followed by a four-stage ion exchange water purification process, consisting of a high capacity cation/anion column, two ultrapure ion exchange columns, and an organics removal column (Barnstead Internationals, Dubuque, IA). 2.2. Solution Preparation. A Na2SO4 stock brine solution was prepared by dissolving a calculated amount of sodium sulfate, calcium chloride, and PIPES solids into a 1 M NaCl brine solution. The concentrations of Ca2+ ion and PIPES in the Na2SO4 stock brine were 400 mg L−1 and 10 mM, respectively. For the tests involving scale inhibitor, a known volume of DTPMP, PPCA, or PVS stock solution was added into the Na2SO4 stock brine. Similarly, for the tests involving EDTA or citrate, a known amount of EDTA or sodium citrate dihydrate was added into the Na2SO4 stock solution. A BaCl2 stock brine was prepared by dissolving a calculated amount of barium chloride, calcium chloride, and PIPES solids into a 1 M NaCl brine solution. The concentrations of Ca2+ ion and PIPES in the BaCl2 stock brine were 400 mg L−1 and 10 mM, respectively. The solution pH of both BaCl2 and Na2SO4 solutions was adjusted to 6.74. Barite induction time measurement was initiated by mixing the Na2SO4 stock brine with the BaCl2 stock brine with a 1:1 volume ratio inside a sample vial at various testing conditions. The sample vial was constantly stirred by a magnetic stirrer. The stock solution of ferric chloride (Fe(III) stock solution) was prepared by adding a known amount of FeCl3 solid into a 1% HNO3 solution. Unless otherwise specified, in each experiment evaluating the impact of Fe(III) on scale inhibitor, Fe(III) stock solution was gradually delivered into the BaCl2 stock brine before mixing with the Na2SO4 stock solution to initiate the induction time measurement. Due to the presence of PIPES buffer, the Na2SO4−BaCl2 mixture brine pH was not significantly influenced by the addition of acidic Fe(III) stock solution. All experiments were performed at an ambient pressure of 1 atm. 2.3. Experimental Apparatus and Measurement of Induction Time. The apparatus to measure barite induction time adopted in this study is based on a laser nucleation detection method developed previously.26 In the previous study, this apparatus has been applied to test the inhibition efficiency of several scale inhibitors at different pHs, temperatures, and saturation indexes. Figure 1 illustrates the schematic of the experimental apparatus. The apparatus includes a laser source, a photodetector, a sample vial with a stirrer inside, and a stir plate. The Na2SO4 stock brine and the BaCl2 stock brine were mixed at a 1:1 volume ratio inside the sample vial. In each experiment, both Na2SO4 brine and BaCl2 brine were delivered into the sample vial by use of glass syringes. The error of the amount of

solution taken by the syringe was less than 1%. The temperature of the solution inside the sample vial was carefully controlled by a digital water bath. The laser component for scale crystal detection includes a green (532 nm) laser diode, lenses, a photodetector, and the associated mounting accessories. A photodetector made by a silicon photodiode was used to detect the laser signal and was connected to a multimeter which transforms the laser signal to electronic current in milliamps. When barite scale crystals form due to nucleation, the laser light will be scattered by the formed scale particles suspended in solution, causing a decrease of the laser signal intensity and hence a reduction in the electric current reading. The moment when the electric current reading starts to drop noticeably indicates the appearance of crystals. In this study, induction time is the time duration from the mixing of Na2SO4 brine with BaCl2 brine until the appearance of crystals, i.e., the moment when the electric current reading started to drop noticeably. A representative data set of electric current versus time elapsed post brine mixing is shown in Figure S3 in the Supporting Information. For all experiments, solution pH values before and after the barite induction time measurement changed insignificantly, due to the presence of PIPES buffer. 2.4. Separation of Ferric Hydroxide Particles. Due to the fine particle size of ferric hydroxide solid, an analytical ultracentrifuge machine (Optima L-80XP, Beckman Coulter) was adopted to separate the ferric hydroxide particle from the bulk solution. After centrifuging the solution at 45 000 rpm for 40 min at 25 °C, the majority of ferric hydroxide particles appeared to have been separated from the bulk solution with the supernatant solution being either colorless or a slightly yellowish color. Ultracentrifugation were conducted at room temperature due to the operational limit of the ultracentrifuge machine. Following the ultracentrifugation treatment, a known volume of the supernatant was removed by pipet. The remaining liquid and solid in the ultracentrifuge tube was acidified by 1 N HCl acid to dissolve the ferric hydroxide particles. Immediately after adding the HCl acid, a white precipitate was observed, which was the acidified PIPES buffer solid. The acidified solution with the white precipitate was filtered through a 0.45 μm filter paper. The retentate was dissolved into 0.1 M NaOH solution. No detectable Fe(III) or DTPMP was found in the retentate. 2.5. Analytical Methods. Concentrations of cationic species, sulfur, and phosphorus of the samples were measured by an inductively coupled plasma optical emission spectrometer (ICPOES; Optima 4300, PerkinElmer). DTPMP inhibitor concentration was measured based on the phosphorus measurement. PVS concentration in the stock solution was measured based on the sulfur measurement. PPCA concentration in the stock solution was measured via total dissolved organic carbon measurement (TOCVCSH, Shimadzu Corp.).

3. RESULTS AND DISCUSSION 3.1. Mechanistic Understanding of Fe(III) Effect on Scale Inhibitors. According to a recent USGS report, in the United States among all the produced waters and formation waters with iron species present, about half of these waters contain a total iron concentration below 10 mg L−1 and ca. 30% of waters contain between 10 and 100 mg L−1.27 Thus, in this study a Fe(III) concentration range of 0.4−2.0 mg L−1 was adopted, assuming less than 10% of total iron in formation waters is in the form of Fe(III). The effect of Fe(III) species C

DOI: 10.1021/acs.energyfuels.8b01837 Energy Fuels XXXX, XXX, XXX−XXX

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Energy & Fuels Table 1. Mixture Brine Solution Compositions and Experimental Conditions brine solution

NaCl (M)

Ba2+ (mg L−1)

SO42− (mg L−1)

Ca2+ (mg L−1)

PIPES (mM)

pHa

SIb

tempc (°C)

MBS-I MBS-II MBS-III ref brine-18 ref brine-228 ref brine-329

1 1 1 0.7 2 0.6

204 67 0 550 10.3 187

143 47 0 10 2000 533

400 400 400 6500 445 629

10 10 10 NA NA NA

6.38 6.38 6.74 6.80 6.50 6.23

2.0 1.0 NA 1.0 0.2 2.6

70 70 25 70 65 70

a

pH value is reported at the respective temperature. pH values in this study (MBS-I, -II, and -III) are controlled by buffer. bSI is calculated at the respective temperature for each brine solution. cTemperature is reported as the experimental temperatures for each mixture brine solution adopted in this study or the temperature conditions of the reference brines reported in respective studies.

exacerbated when Fe(III) concentration was elevated to 1.0 and 2.0 mg L−1 with the corresponding induction time being further shortened to 64 and 22 s, respectively (nucleation-4 and nucleation-5). At a higher DTPMP concentration of 1.9 mg L−1, the detrimental effect of Fe(III) was confirmed again by reducing the induction time from over 23 000 s without Fe(III) (nucleation-6) to only 403 s in the presence of 1 mg L−1 Fe(III) (nucleation-7). A similar impairment of scale inhibitor performance was also observed for polymeric inhibitors such as PPCA and PVS (nucleation-8 to nucleation-11). The presence of 1 mg L−1 Fe(III) impaired the performance of both PPCA and PVS by shortening the induction time from a few hundred seconds to less than 70 s. Obviously, it can be confirmed in this study that the presence of Fe(III) species has a considerable detrimental effect on the inhibitory performance of DTPMP, PPCA, and PVS inhibitors. Ferric hydroxide is sparingly soluble in aqueous solution at near-neutral solution pH.14 Table 3 lists the calculated

on scale inhibitors was evaluated by studying the detrimental effect of Fe(III) on scale inhibitors by measuring barite scale induction time. Three common oilfield scale inhibitors were employed in this study, i.e., DTPMP, PPCA, and PVS. In these experiments, mixing the BaCl2 stock brine and the Na2SO4 stock brine via a 1:1 volume ratio formed the mixture solution of mixture brine solution I (MBS-I) with brine compositions and experimental conditions shown in Table 1. Ca2+ was added into the brine solution due to the abundancy of Ca2+ ion in oilfield produced waters. Table 1 also presents the compositions of oilfield brines from three different fields as references.8,28,29 This is to demonstrate that the selected ionic species concentrations in this study are within the range of ionic species in oilfield brines. The calculated SI value of barite, i.e., SI(barite), for MBS-I at 70 °C was 2.0, indicative of the solution being supersaturated with respect to barite. The effect of Fe(III) on scale inhibitor is tabulated in Table 2 and is also Table 2. Barite Induction Time under Different Experimental Conditions

Table 3. Saturation Index (SI) of Various Iron Oxide Minerals at the Experimental Condition of pH 6.38 and 70 °C

a

expt

DTPMP (mg L−1)

PPCA (mg L−1)

PVS (mg L−1)

Fe(III) (mg L−1)

tind (s)

nucleation-1 nucleation-2 nucleation-3 nucleation-4 nucleation-5 nucleation-6 nucleation-7 nucleation-8 nucleation-9 nucleation-10 nucleation-11

0 1.2 1.2 1.2 1.2 1.9 1.9 0 0 0 0

0 0 0 0 0 0 0 1.2 1.2 0 0

0 0 0 0 0 0 0 0 0 1.2 1.2

0 0 0.4 1.0 2.0 0 1.0 0 1.0 0 1.0

20 1751 241 64 22 23698 403 987 42 600 67

mineral name Fe(OH)2.7Cl0.3 ferrihydrite ferrihydrite (aged) goethite hematite lepidocrocite maghemite a

a

ion activity product −2.7

−0.3

aFe aOH aCl aFe3+aOH−3 aFe3+aOH−3 aFe3+aOH−3 aFe3+aOH−3 aFe3+aOH−3 aFe3+aOH−3 3+

log Ksp

calcd SIa

−40.84 −38.80 −39.31 −41.51 −42.71 −40.63 −38.81

7.54 5.68 6.19 7.48 8.77 5.20 3.38

Calculation of SI is achieved by Visual MINTEQ (version 3.1).

solubility product (Ksp) and SI values of various forms of ferric hydroxide solids. Obviously, MBS-I is significantly supersaturated with respect to each type of ferric hydroxide solid listed in Table 3. Therefore, Fe(III) species should exist in the solution as ferric hydroxide solid particles at the experimental condition. It should be noted that the exact form and molecular structure of ferric hydroxide solids depend on brine chemistry, solution physicochemical conditions, and the stability of the ferric hydroxide solid.14 Mineral nucleation rate and the corresponding induction time can be affected by the presence of a trace amount of dissolved foreign species in the aqueous solution.28 In some cases, these species can serve as inhibitors to delay the nucleation process, while in other cases these species can behave as accelerators to expedite the nucleation process. The function of foreign species in terms of their impact on the nucleation process cannot be easily predicted based upon the composition of the aqueous solution

As detailed in the text, the majority of the added Fe(III) species should be in the form of ferric hydroxide solid. The listed Fe(III) concentration is the equivalent iron concentration of the ferric hydroxide solid.

illustrated in Figure S4 in the Supporting Information. According to Table 2, the barite induction time (tind) is approximately 20 s in the absence of DTPMP and Fe(III) (nucleation-1). In the presence of 1.2 mg L−1 DTPMP inhibitor (nucleation-2), the barite induction time was considerably extended to beyond 1700 s as a result of the inhibitory effect of DTPMP. However, the addition of Fe(III) species can adversely impact the performance of DTPMP in barite nucleation inhibition. Nucleation-3 suggests that the presence of as low as 0.4 mg L−1 Fe(III) led to a substantial reduction of the induction time to less than 250 s. The impairment of DTPMP performance by Fe(III) was D

DOI: 10.1021/acs.energyfuels.8b01837 Energy Fuels XXXX, XXX, XXX−XXX

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Energy & Fuels and the nature of these species.28 In the presence of a suitable combination of foreign species and crystal surfaces, the nucleation process can be facilitated. Therefore, one hypothesis to elucidate the observed effect of Fe(III) is that the precipitated ferric hydroxide particles provide the necessary nucleation sites to promote barite nucleation. Additional nucleation experiments were designed and conducted in order to examine this proposed mechanism of Fe(III) on barite nucleation. In these experiments, the barite induction times at different Fe(III) concentration levels were compared. Based upon the aforementioned hypothesis, it can be speculated that, at the same experimental conditions (SI, temperature, pH, etc.), a higher Fe(III) concentration should produce more ferric hydroxide particles generating more nucleation sites to promote barite nucleation, resulting in a reduced induction time. Experiments were carried out by adopting a mixture brine solution II (MBS-II) at pH 6.38 and 70 °C, with brine compositions shown in Table 1. The only difference between MBS-I and MBS-II is that SI(barite) of MBS-II was lowered from 2.0 to 1.0 to slow down the nucleation process with a prolonged induction time, allowing a more reliable and accurate observation of the impact of Fe(III). It has been reported by several scholars that the deposition kinetics of scale particles is affected by a number of factors including SI.8−10,30−32 As long as the solution is still supersaturated with respect to the scale of interest, a reduced SI will lead to a lower scale deposition kinetics. Figure 2

hydroxide particles to promote barite nucleation might not be responsible for the observed Fe(III) effect on barite induction time. DTPMP inhibitor was chosen to obtain mechanistic understanding of the Fe(III) impact of scale inhibitor. This is because DTPMP has an extensive application in industrial scale control6 and also, compared with polymeric inhibitors, the aqueous concentration of DTPMP can be determined in a relatively easy manner using ICP-OES. Ferric hydroxide particles are known to strongly adsorb anionic species in aqueous solution, such as carboxylic, oxalate, sulfonate, and lactate.33 Once formed in aqueous solution, the surface of ferric hydroxide solid will be covered with active hydroxyl functional groups. It is typically assumed that the freshly precipitated ferric hydroxide solids have a large surface area.17 Therefore, ferric hydroxide particles could adsorb inhibitor molecules via the formation of chemical bonds between the surface hydroxyl groups and functional groups of the inhibitors, leading to a reduction of the functionality of the inhibitors. Phosphonates are known to be able to strongly adsorb to the surface of metal hydroxides.34 Therefore, another plausible explanation of the detrimental effect of Fe(III) on scale inhibitors is the adsorption of scale inhibitors onto the surface of ferric hydroxide particles. When Fe(III) stock solution is added into the BaCl2 brine solution, due to the presence of PIPES buffer, the BaCl2−Fe(III) mixture solution pH is about 6.0. At a near-neutral pH, Fe(III) will be immediately hydrolyzed forming ferric hydroxide colloidal particles. One can conceive that adsorption of scale inhibitors onto the ferric hydroxide surfaces will occur subsequent to the mixing of Na2SO4−inhibitor brine with BaCl2−Fe(III) brine, resulting in a reduction in inhibitory efficiency. In addition, it can be inferred that any remaining unabsorbed freely mobile scale inhibitor in the brine solution would still be available to inhibit barite nucleation. To examine the hypothesis that the detrimental effect of Fe(III) is due to the precipitation of ferric hydroxide solid and the subsequent adsorption of scale inhibitor to ferric hydroxide solid surface, the following three arguments can be inferred accordingly: 1. Adsorption of scale inhibitor by ferric hydroxide will lead to a reduction of aqueous phase inhibitor concentration. 2. Scale inhibitors being removed from the aqueous phase due to adsorption are transferred to the surface of ferric hydroxide based upon the adsorption capacity of the ferric hydroxide solid. 3. The remaining unadsorbed mobile scale inhibitor after achieving adsorption equilibrium should be able to effectively inhibit barite scale formation. The efficacy of the remainder mobile inhibitor should match with the efficacy of scale inhibitor at the same testing conditions in an aqueous solution of otherwise the same composition except in the absence of

Figure 2. Illustration of induction time with Fe(III) concentration.

illustrates the measured induction time at different Fe(III) concentrations. Statistical analysis suggests that the difference in the barite induction time at different Fe(III) concentrations was statistically insignificant with a p-value of 0.85. This suggests that the increase in Fe(III) concentration did not promote barite nucleation. Therefore, the proposed mechanism of the formation of more nucleation sites by ferric

Table 4. Details of Adsorption Experiments of DTPMP to Surface of Ferric Hydroxide Particles before ultracentrifugation

after ultracentrifugation

expt

init DTPMP concn (mg L−1)

init Fe(III) concn (mg L−1)

ICP measd DTPMP concn in supernatant (mg L−1)

ICP measd Fe(III) concn in supernatant (mg L−1)

adsorption capacitya (mg mg−1)

calcd free DTPMP concnb (mg L−1)

measd 0 tinh ind/tind

calcd DTPMP concnc (mg L−1)

adsorption-1 adsorption-2

2.29 3.00

1.00 1.00

1.27 2.03

0.17 0.22

1.23 1.24

1.06 1.76

8.23 33.45

1.17 1.95

a The unit of adsorption capacity is mg of DTPMP per mg of Fe(III). bFreely mobile DTPMP concentration was calculated by subtracting the adsorbed DTPMP concentration from the ICP measured (total) concentration of DTPMP in solution after centrifugation. cThis column lists the calculated DTPMP concentration based upon the measured induction time via eq 2.

E

DOI: 10.1021/acs.energyfuels.8b01837 Energy Fuels XXXX, XXX, XXX−XXX

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Energy & Fuels

1.24 mg mg−1 in this study, surface coverages of ferrihydrite by DTPMP can be calculated to be ca. 1.07 and 1.08 layers, respectively. This suggests that, in this adsorption study, DTPMP molecules achieved an approximately monolayer surface coverage on ferrihydrite surface. Thus, far, the first two statements can be confirmed by the adsorption experiments. The last statement can be verified via laboratory measurement of barite induction time and the relationship between induction time and inhibitor concentration. One can speculate that once the adsorption experiment reached an equilibrium, the unadsorbed freely mobile DTPMP molecules should be able to independently function as scale inhibitors to inhibit barite nucleation and the efficacy of the unadsorbed DTPMP should be the same as those at the same testing conditions in an aqueous solution of otherwise the same composition except in the absence of iron. Therefore, the measured induction time of barite at these two aqueous solutions (one with iron and one without iron) should be the same as long as SI(barite), free DTPMP concentration, and other physical conditions, such as temperature and ionic strength, are the same. Based upon the relationship of barite induction time and DTPMP concentration, the measured induction time of barite can be used to calculate the effective DTPMP concentration in an iron-free aqueous solution. A number of previous studies have investigated the relationship between DTPMP aqueous concentration and the induction time.36−38 More recently, such a relationship has been shown to be a consequence of classical nucleation theory.39 The relationship between DTPMP inhibitor and barite induction time can be characterized by the following expression (eq 2):

iron. In other words, the unadsorbed mobile DTPMP molecules can inhibit barite scale formation independently. To test these three statements, a new set of inhibitor adsorption experiments were conducted by adopting a baritefree mixture brine solution III (MBS-III) as the background brine at 25 °C and pH 6.74, with brine compositions shown in Table 1. A DTPMP-containing brine and another Fe(III)containing brine were prepared by adding DTPMP stock solution and Fe(III) stock solution into two MBS-III brines. Adsorption of DTPMP by Fe(III) was initiated by mixing these two prepared brines. Adsorption experiments were carried out at an ambient temperature of 25 °C due to the temperature limitation of the ultracentrifuge machine. Table 4 lists the experimental details regarding the initial concentrations of DTPMP and Fe(III) of two adsorption experiments (adsorption-1 and adsorption-2). Both adsorption experiments adopted an initial concentration of Fe(III) of 1 mg L−1. In this study, it is of key importance to be able to physically separate the formed ferric hydroxide particles from the aqueous solution so that the unadsorbed DTPMP inhibitor freely mobile in the aqueous solution can be separated from the adsorbed inhibitor on ferric hydroxide surface. Otherwise, it is not feasible to distinguish and quantify the adsorbed inhibitor from the unadsorbed ones. The formed ferric hydroxide colloidal particles were so fine that they can pass through 0.02 μm membrane filters. Therefore, the separation of ferric hydroxide solids was achieved via a high-speed ultracentrifugation approach. After centrifuging the suspension at 45 000 rpm for 40 min, visual inspection suggested that most of the ferric hydroxide particles appeared to have been separated from the solution, with the supernatant being either colorless or a slightly yellowish color. Aqueous iron concentrations measured by ICP-OES after the ultracentrifuge for adsorption-1 and adsorption-2 were 0.17 and 0.22 mg L−1, respectively. This suggests that the majority of the added Fe(III) has been removed during the centrifugation and about 20% of the initially added Fe(III) was left in the solution after centrifugation. This amount of Fe(III) probably corresponds to the removal capacity of the centrifugation machine considering the centrifuge speed, centrifuge time, ferric hydroxide particle size, and fluid dynamic properties. As detailed in Table 4, the reductions of aqueous DTPMP and Fe(III) concentrations after centrifugation treatment in adsorption-1 were 1.02 and 0.83 mg L−1, respectively, suggesting a sorption capacity of ferric hydroxide as 1.23 mg of DTPMP/mg of Fe(III). Since as much as 0.17 mg L−1 Fe(III) was left in the supernatant after centrifugation, the amount of DTPMP adsorbed by these leftover ferric hydroxide solid can be easily calculated as 0.21 mg L−1. This suggests that the freely mobile DTPMP in the supernatant should be approximately 1.06 mg L−1, as shown in Table 4. Experimental data obtained from adsorption-2 suggest that the sorption capacity of ferric hydroxide solid is 1.24 mg L−1, which is very close to that obtained from the first experiment. Thus, the freely mobile DTPMP concentration in adsorption-2 is calculated to be ca. 1.76 mg L−1 (Table 4). Moreover, it has been reported that ferrihydrite particles ((Fe3+)2O3·0.5H2O) have a surface area of ca. 600 m2 g−1,17 and 1 mg of Fe(III) corresponds to 1.91 mg of ferrihydrite solid. As for the DTPMP molecule, a previous study indicated that 1 mg of DTPMP molecules can cover a surface area of 1 m2.35 Assuming the solid formed from ferric species is ferrihydrite, based upon the calculated adsorption capacities of 1.23 and

ji t zyz z = binh logjjj ind j t 0 zz k ind {

ij L yz jj z jj mg zzz Cinh k {

ij mg yz zz jj k L {

(2)

where tind (s) stands for the induction time of barite in the presence of DTPMP inhibitor; Cinh (mg L−1) is the DTPMP concentration; t0ind (s) represents the induction time of barite in the absence of inhibitor in solution; binh is a constant characterizing the inhibition efficiency. Hence, the term of tind 0 t ind

corresponds to the prolonged induction time due to the presence of inhibitor. In this study, experiments measuring barite induction time using the supernatant solutions have been performed. First, the supernatant solutions from adsorption-1 and adsorption-2 were employed as the background brine solutions to conduct barite nucleation experiments by mixing barium-containing solution with sulfatecontaining solution at 25 °C with an SI(barite) of 2.4. As shown in Table 4, the measured tind by using the supernatant 0 t ind

solutions of adsorption-1 and adsorption-2 are 8.22 and 33.45, respectively. In a separate study, barite induction time has been evaluated at different DTPMP concentrations using MBS-III with an SI(barite) of 2.4 at pH 6.74 and 25 °C. The objective is to obtain the relationship between the measured induction time and DTPMP concentration. Figure 3 provides the linear

( ) and

relation between log

t ind 0 t ind

DTPMP concentration at

testing conditions. Based upon the obtained relationship shown in Figure 3, the corresponding aqueous DTPMP concentrations can be calculated as 1.17 and 1.95 mg L−1 for adsorption-1 and adsorption-2, respectively. These results are very close to the DTPMP concentrations calculated based F

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examined with a DTPMP concentration of 1.2 mg L−1. The experimental results are illustrated in Figure 4, parts a and b, for the impact of EDTA and citrate, respectively. As discussed earlier, a barite induction time of 20 s was measured at the testing condition in the absence of Fe(III), DTPMP, or chelating agent (nucleation-1). The presence of 1.2 mg L−1 DTPMP can substantially increase the induction time to beyond 1700 s (nucleation-2). As expected, such an inhibition effect of DTPMP was considerably impaired with an induction time of only 241 s when 0.4 mg L−1 Fe(III) was added into the aqueous system (nucleation-3). One might expect that the addition of EDTA into the system would markedly reverse the impairment of Fe(III) on DTPMP with a considerably prolonged induction time. Experimental results suggest that the reversal effect of EDTA on Fe(III) is not significant when the EDTA to Fe(III) molar ratio is below 30:1 (Figure 4a). By increasing the EDTA concentration with a molar ratio of EDTA to Fe(III) from 1:1 to 30:1 (EDTA-1 to EDTA-6), the induction time marginally increased from 254 to 502 s. Only when the EDTA to Fe(III) molar ratio increased to 50:1 (EDTA-7), the induction time was extended beyond 1100 s, suggesting a pronounced reversal of impairment of Fe(III) impact due to the chelation effect of EDTA at such a concentration level. A similar experiment was conducted by evaluating the effect of citrate on Fe(III) by use of MBS-I at pH 6.38 and 70 °C. According to Table 5 and Figure 4b, with the increase of citrate concentration from a citrate to Fe(III) molar ratio of 1:1 to 10:1 (citrate-1 to citrate-4), the barite induction time was substantially increased from 315 s to beyond 1000 s. When the citrate to Fe(III) molar ratio reached 15:1 (citrate-5), the barite induction time was further extended to 1625 s, which is close to the induction time without Fe(III) (nucleation-2). The experimental data suggest that, under the testing condition, citrate is more effective than EDTA in reversing the impairment of Fe(III) to DTPMP. For instance, at the same ligand to Fe(III) molar ratio of 10:1, the measured values of induction time were 1019 and 399 s in the presence of citrate and EDTA, respectively. However, according to water chemistry principles, EDTA is known to be a stronger chelating agent for Fe(III) than citrate. The base-10 logarithms of stability constants of Fe(III) with EDTA and citrate are reported to be 25.7 and 11.9, respectively.40 In other words, Fe(III)−EDTA complex is 10 orders of magnitude more stable

Figure 3. Linear relationship between DTPMP inhibitor concentration and log tind/t0 without presence of Fe(III).

upon Fe(III) adsorption capacity, i.e, 1.06 and 1.76 mg L−1. This can support the argument that the efficacy of the remainder mobile DTPMP inhibitor after centrifugation matches well with the efficacy of scale inhibitor at the same conditions in an iron-free solution. In summary, based on the above investigations and discussions, the hypothesis that the detrimental effect of Fe(III) on scale inhibitors is due to the adsorption of scale inhibitors onto the surface of ferric hydroxide solid particles can be verified. 3.2. Use of Organic Chelating Agents To Reverse the Effect of Fe(III) on Scale Inhibitors. To prevent the detrimental impact of Fe(III) on phosphonate scale inhibitors, a number of chelating chemical agents are commonly used in the industry and especially oilfield for iron control.1,2 These chelating agents (ligands) can sequester iron species forming metal−ligand complexes, effectively reducing the concentration of free iron species and subsequently ferric hydroxide amount in the aqueous solution. With the reduction of iron concentration in the solution, more scale inhibitors become available for scale control and the negative impact of Fe(III) on the inhibitors can be reversed. Since both EDTA and citrate are among the most commonly used chelating agents, EDTA and citrate were selected in this study to evaluate their abilities to reverse Fe(III) impairment on scale inhibitors. The impact of EDTA and citrate was investigated by measuring the barite induction time at different scenarios with and without EDTA and citrate presence (Table 5). MBS-I was employed for this study with an SI(barite) of 2.0 at pH 6.38 and 70 °C. A series of ligand to Fe(III) molar ratios from 1:1 to 15:1 have been

Table 5. Experimental Details of the Impact of EDTA and Citrate on Fe(III) expt

DTPMP (mg L−1)

Fe(III) (mg L−1)

EDTA (mg L−1)

citrate (mg L−1)

ligand:Fe(III) molar ratio

tind (s)

nucleation-1 nucleation-2 nucleation-3 EDTA-1 EDTA-2 EDTA-3 EDTA-4 EDTA-5 EDTA-6 EDTA-7 citrate-1 citrate-2 citrate-3 citrate-4 citrate-5

0 1.2 1.2 1.2 1.2 1.2 1.2 1.2 1.2 1.2 1.2 1.2 1.2 1.2 1.2

0 0 0.4 0.4 0.4 0.4 0.4 0.4 0.4 0.4 0.4 0.4 0.4 0.4 0.4

0 0 0 2.1 6.3 10.4 20.9 31.3 62.6 104.3 0 0 0 0 0

0 0 0 0 0 0 0 0 0 0 1.4 4.1 6.9 13.7 20.6

0 0 0 1:1 3:1 5:1 10:1 15:1 30:1 50:1 1:1 3:1 5:1 10:1 15:1

20 1751 241 254 308 311 399 428 502 1172 315 357 504 1019 1625

G

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hydrolyzed into ferric hydroxide particles, leaving limited Fe(III) free species to complex with EDTA. Since the chelation function of EDTA is to work directly with the ferric species rather than ferric hydroxide solids, EDTA was less effective in chelating with Fe(III) under the testing condition. Citrate, on the hand, follows a different chelating mechanism than EDTA in that citrate adsorbs to the surface of the ferric hydroxide particles.42 In citrate-1 to citrate-5, the Fe(III) species was hydrolyzed first forming ferric hydroxide particles as soon as being added into the BaCl2 brine solution. Subsequent to barium and sulfate solution mixing, the citrate adsorbed to the surface of ferric hydroxide solids, covering the active sorption sites on the particle surface. Once a considerable fraction of the active sorption sites of the formed ferric hydroxide particles were covered by citrate, ferric hydroxide particles could no longer be available to react with DTPMP to impair the inhibitor performance. Previous studies have shown that citrate is effective in covering the surface of ferric hydroxide via adsorption.43 A number of previous studies investigated the adsorption process of L-tartaric acid, mesotartaric acid, and citrate on the surface of amorphous ferric hydroxide.33,44 It was found that, among these three chelating agents, citrate has the strongest surface complex stability with ferric hydroxide. In order to confirm the aforementioned chelating mechanism of EDTA, another experiment was carried out by changing the way Fe(III) and EDTA were introduced into the mixture solution. Instead of adding the stock solution of Fe(III) into the BaCl2 solution, the acid Fe(III) stock solution was first mixed with another volume of EDTA solution and the final solution pH was ca. 2.2. The molar ratio of Fe(III) to EDTA was still maintained at 1:1. At solution pH 2.2, the hydrolysis reaction of Fe(III) is significantly suppressed and the majority of the Fe(III) molecules are in the form of freely mobile Fe(III). Subsequently, this Fe(III)−EDTA mixture solution was added into the Na2SO4 brine solution, followed by mixing with the BaCl2 brine solution. The final mixture solution had the same composition, SI(barite), and solution pH as in EDTA-1, except the addition sequence of the chemicals. The experimental temperature remained at 70 °C. The measured barite induction time in this experiment was as high as 1156 s, suggesting that EDTA effectively formed a complex with Fe(III) so that DTPMP was able to inhibit barite scale nucleation. The experimental result proved that if EDTA can effectively chelate Fe(III) species before Fe(III) is hydrolyzed into ferric hydroxide solids, EDTA can reverse the detrimental effect of Fe(III) on DTPMP inhibitor. In addition, EDTA has a higher stability constant with Fe(III) species than with other common cations present in oilfield waters, such as calcium, barium, and magnesium.40 This suggests that the presence of other commonly encountered cations in oilfield water will not be able to substantially compete with Fe(III) in terms of chelating with EDTA. Moreover, it should be noted that scale inhibitors as threshold inhibitors function differently from chelating chemicals, such as EDTA and citrate, which form chelates with metal ions based upon stoichiometric ratios. However, phosphonate inhibitors can chelate with metals due to their multidentate molecular structures. Additional studies are required to understand the impact of chelating agents on scale inhibitor performance when Fe(III) is present in oilfield brines.

Figure 4. Illustration of impact of chelating agents on Fe(III) to reverse the detrimental impact on DTPMP inhibitor. (a) Impact of EDTA. (b) Impact of citrate.

than Fe(III)−citrate complex. The discrepancy between the experimental results and the water chemistry principles in terms of chelating capability of EDTA and citrate with Fe(III) can be elucidated by the difference in the mechanism of these two ligands in forming complexes with Fe(III). As for EDTA, this ligand primarily works with freely mobile ferric species instead of ferric hydroxide particles.41 In EDTA-1 to EDTA-7, an acidic stock solution of Fe(III) was first added into a BaCl2 brine containing NaCl, BaCl2, and PIPES and the resultant solution pH was near 6 due to the presence of PIPES buffer. Immediately after the Fe(III) stock solution was added into the BaCl2 brine, the Fe(III) species was hydrolyzed, forming ferric hydroxide solids. Upon mixing the BaCl2−Fe(III) brine with Na2SO4−DTPMP−EDTA brine, ferric species has been H

DOI: 10.1021/acs.energyfuels.8b01837 Energy Fuels XXXX, XXX, XXX−XXX

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4. CONCLUSIONS

Article

ASSOCIATED CONTENT

S Supporting Information *

In this study, the detrimental impact of Fe(III) on scale inhibitors has been evaluated by measuring barite scale induction time in the presence of scale inhibitors. It shows that a sub milligram per liter concentration level of Fe(III) in brine solution can substantially impair the performance of both phosphonate and polymeric scale inhibitors with a considerable reduction in the induction time. Since Fe(III) has an extremely low solubility in aqueous solution at near-neutral pH, it is hypothesized that the functioning mechanism of Fe(III) to scale inhibitor is the adsorption of inhibitor molecules to the surface of the formed ferric hydroxide particles. Such a hypothesis can be verified by evaluating the adsorption capacity of the ferric hydroxide particles of DTPMP inhibitor and also the efficacy of the unadsorbed DTPMP in the aqueous phase. The results indicate an adsorption capacity of ca. 1.2 mg of DTPMP inhibitor per milligram of Fe(III) species. Considering the reported surface area of both ferric hydroxide and DTPMP, a monolayer surface coverage of ferric hydroxide by DTPMP can be calculated. Moreover, the unadsorbed freely mobile DTPMP inhibitor in the aqueous phase can inhibit barite scale with the same efficacy as the DTPMP molecules in an iron-free environment. This study also investigates the effect of chelating agents such as EDTA and citrate in reversing the detrimental effect of Fe(III) on DTPMP. The experimental results show that, at the testing condition, citrate demonstrates an enhanced performance compared with EDTA in reversing the Fe(III) effect, albeit there is a higher stability constant of Fe(III) with EDTA than with citrate. This phenomenon is due to the difference in the mechanism of these two ligands in forming complexes with Fe(III). EDTA primarily forms complexes with Fe(III) mobile species, while citrate adsorbs to the surface of ferric hydroxide particles. A significantly enhanced EDTA performance in reversing the Fe(III) effect can be achieved in this study by premixing the acidic Fe(III) solution with EDTA before Fe(III) is hydrolyzed. This study provides the theoretical basis and technical insights for oilfield iron control to minimize iron impairment on scale inhibitor performance.

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.energyfuels.8b01837. Molecular structures of DTPMP, PPCA, and PVS; turbidity measurement experiments to show no impact of PIPES buffer on barite scale deposition; representative data set of electric current versus time elapsed post brine mixing; illustration of the relationship of induction time with inhibitor concentration and Fe(III) concentration (PDF)

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Tel.: +853 8822 4917. ORCID

Ping Zhang: 0000-0003-0820-2056 Present Address ∥

Fulcrum Resources Inc., San Francisco, CA.

Author Contributions ⊥

Z.Z. and P.Z.: These authors contributed equally to the manuscript. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was financially supported by Brine Chemistry Consortium companies of Rice University, including Aegis, Apache, BHGE, BWA, Chevron, ConocoPhillips, Coastal Chemical, EOG Resources, ExxonMobil, Flotek Industries, Halliburton, Hess, Italmatch, JACAM, Kemira, Kinder Morgan, Nalco, Oasis, Occidental Oil and Gas, Range Resources, RSI, Saudi Aramco, Schlumberger, Shell, SNF, Statoil, Suez, Total, and the NSF Nanosystems Engineering Research Center for Nanotechnology-Enabled Water Treatment (ERC-1449500). The authors also appreciate the financial support of a Start-Up Research Grant provided by the University of Macau (SRG2018-00112-FST) and the sponsorship of Science and Technology Development Fund, Macao S.A.R. (FDCT) (0063/2018/A2).

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5. FUTURE WORK

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