w a t e r r e s e a r c h 5 3 ( 2 0 1 4 ) 3 1 0 e3 2 1
Available online at www.sciencedirect.com
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Reactive iron sulfide (FeS)-supported ultrafiltration for removal of mercury (Hg(II)) from water Dong Suk Han a,*, Maria Orillano a, Ahmed Khodary a, Yuhang Duan b, Bill Batchelor b, Ahmed Abdel-Wahab a a b
Chemical Engineering Program, Texas A&M University at Qatar, Education City, Doha 23874, Qatar Zachry Department of Civil Engineering, Texas A&M University, College Station, TX 77843-3136, USA
article info
abstract
Article history:
This study investigated removal of Hg(II) from water using FeS(s) with batch and contin-
Received 28 September 2013
uous contact filtration systems. For the batch system, kinetic experiments showed that
Received in revised form
removal of Hg(II) by FeS(s) was rapid at lower concentration (500 mM), but at higher con-
15 January 2014
centration (1000 and 1250 mM), more time was required to achieve greater than 99%
Accepted 18 January 2014
removal. The concentration of iron released to the solution remained relatively low,
Available online 28 January 2014
typically below 3 mM. This would theoretically present less than 1% of the Hg(II) removed. Thus, a simple exchange of Hg(II) for Fe(II) in the solid (FeS(s)) does not explain the results,
Keywords:
but if the Fe(II) released could react to form another solids, low concentrations of Fe do not
FeS
preclude a mechanism in which Hg(II) reacts to form HgS and release Fe(II). A continuous
Mackinawite
contact dead-end ultrafiltration (DE/UF) system was developed to treat water containing
Mercury
Hg(II) by applying a FeS(s) suspension with stirred or non-stirred modes. A major reason for
Ultrafiltration
applying stirring to the system was to investigate the role of “shear” flow in rejection of
Adsorption
Hg(II)-contacted FeS(s) by a UF membrane and the stability of Hg on the FeS(s). The Hg(II)-
Desorption
contacted FeS(s) was completely rejected by the DE/UF system and mercury was strongly retained on the FeS(s) particles. Almost no release of Hg(II) (z0 mM) from the FeS(s) solids was observed when they were contacted with 0.1M-thiosulfate, regardless of whether the system was operated in stirred or non-stirred mode. However, rapid oxidation of FeS(s) was observed in the stirred system but not in the non-stirred system. Determining the mechanism of oxidation requires further study, but it is important because oxidation reduces the ability of the solids to remove additional Hg(II). ª 2014 Elsevier Ltd. All rights reserved.
1.
Introduction
Mercury has been considered to be a global contaminant of significant concern for centuries due to its high toxicity and potential for bioaccumulation via the aquatic food chain, which seriously affects natural ecosystems and the health of humans. Exposure to high levels of mercury can cause
inhibition of enzyme activity, cell damage, impairment of pulmonary function and kidney performance, chest pain, and damage to the central nervous system (Jeong et al., 2007; Kaneko, 1988; Zahir et al., 2005). Mercury contamination occurs due to release from anthropogenic and natural sources, including chloro-alkali plants, mining and smelting activities, coal-fired power plants, electrical and electronic
* Corresponding author. Tel.: þ974 6629 1478; fax: þ974 4423 0065. E-mail address: [email protected] (D.S. Han). 0043-1354/$ e see front matter ª 2014 Elsevier Ltd. All rights reserved. http://dx.doi.org/10.1016/j.watres.2014.01.033
w a t e r r e s e a r c h 5 3 ( 2 0 1 4 ) 3 1 0 e3 2 1
manufacturing plants, and a variety of incinerator facilities (Vieira and Beppu, 2006; Weisener et al., 2005). Among them, coal-fired power plants and incinerator facilities are the major sources with emissions of 120 ton per year which constitutes 77% of anthropogenic Hg emissions in the United States (Liu et al., 1998). Removal of inorganic Hg from emissions is a crucial process because in aquatic ecosystems, most inorganic Hg is transformed to methylmercury (MeHg), which leads to highly elevated concentrations in aquatic fish and wildlife (Hsu-Kim et al., 2013; Skyllberg et al., 2006). Hg(II) is a soft Lewis acid that preferentially bonds with soft Lewis bases. Since the thiol functional group is a soft base, sulfur-containing chemicals have been widely used to remove mercury (Biester and Zimmer, 1998; Chen et al., 2004; Gibson et al., 2011; Wolfenden et al., 2005). For the same reason, mercury forms a very insoluble solid with sulfide (Brown et al., 1979; Ebadian et al., 2001; Nriagu, 1979; Sadiq, 1992). However, there are problems associated with controlling the sulfide dose in treatment processes. If excess sulfide is added, then soluble HgeS complexes are formed and the effluent concentrations of soluble mercury increases (U.S. Environmental Protection Agency, 1997). In general, the solubility of Hgsulfide solid phases depends on pH and sulfide concentration. At low pH and low sulfide concentration, formation of insoluble mercuric sulfide solid phase (HgS) is preferred, whereas soluble HgeS complexes occur at high pH and high sulfide concentrations (Ravichandran et al., 1999). A way of avoiding the problems that occur in precipitation processes that use soluble reagents is to dose with large particles of iron sulfide, which gradually dissolve to produce the needed sulfide (Nriagu, 1979). If smaller particles are added, the removal mechanism will include sorption and surface reaction. Nano-scale iron sulfide particles have been produced microbially or abiotically and they have been applied to removal of mercury (Ito et al., 2004; Watson et al., 2001, 1995). Particle sizes around 2e5 nm and specific surface areas of 280e500 m2/g have been reported (Watson et al., 2001). Evidence for the production of solid phase products via surface reaction is found in the decrease in soluble mercury concentration with time (1e24 h) that becomes more pronounced as the surface concentration is increased (Watson et al., 1995). Other metals besides Hg were shown to be removed by biogenic FeS(s) (Watson et al., 2001, 1995). Chemical synthesis of nano-sized FeS has also been reported (Butler and Hayes, 1998; Martellaro et al., 2001). Others have also reported that iron sulfides are good reagents for removing mercury from solution (Behra et al., 2001; Bower et al., 2008; Brown et al., 1979; Ehrhardt et al., 2000; Gong et al., 2012; Han et al., 2012; Jean and Bancroft, 1986; Jeong et al., 2007, 2010; Liu et al., 2008; Skyllberg and Drott, 2010). For example, Liu et al. (2008) investigated interactions between aqueous Hg(II) and FeS(s) in batch sorption experiments and found that the maximum removal capacity at lower pH was approximately 0.75 mol Hg(II)/mol FeS. They also used X-ray powder diffraction (XRPD) to identify the major products as metacinnabar, cinnabar, and mercury iron sulfide with less solubility (Ksp ¼ 2E-53) (Barnett et al., 2001; Liu et al., 2008). In this respect, iron sulfide particle has been widely used as immobilizer or methylation inhibitor for Hg(II) in anoxic environment such as estuaries or groundwater (Liu et al., 2009), but
311
less in industrial wastewater released from gas or coal-fired power plant requiring appropriate separation and disposal of final residual solids. However, if appropriate tool or process to safely treat the final solids is provided, use of iron sulfide may have a great merit in such environment. The purpose of this work is to remove Hg(II) in water using reactive FeS(s) with an ultrafiltration system to produce purified water and stable final residual solids that can be greatly rejected by the ultrafiltration (UF) membrane. Hg would be stabilized on the solids by surface reactions after removal by adsorption. In order to meet this goal, we performed a batch kinetic test to evaluate removal of Hg(II) and its sorption behavior by synthetic mackinawite (FeS(s)) as affected by mercury concentration and contact time. Then, lab-scale UF filtration was performed to produce purified water by separating Hg(II)-contacted FeS(s) from the synthetic feed water. In addition, the solid phases deposited on the membrane were evaluated to better understand the chemical changes that occur and how they improve stability of Hg(II) on the FeS(s).
2.
Materials and method
2.1.
Materials
All materials and chemicals applied in this study were used as received. The chemicals used here were: FeCl2 (SigmaeAldrich, 99.9%), Na2S (SigmaeAldrich, 99.9%), HgCl2 (SigmaeAldrich, 99.99%), and Na2S2O3 (SigmaeAldrich, 98%). Components of the UF system were a 300-mL UF stirred cell (Millipore) and 30 kDa regenerated cellulose (RC) ultrafiltration membrane with surface area of 31.7 cm2 (Millipore) (Scheme 1). The water used in this study was prepared through several steps. First, distilled water was collected from a Barnstead Mega-pure distillation device (Thermo Co.) and deionized by a Labconco purifier system. Then, the deionized water was deoxygenated for 2 h by purging with nitrogen gas (99.99% of purity). The deionized/deoxygenated water (DDW) obtained in this way was used for the FeS(s) synthesis experiments and the Hg(II) removal experiments. All experiment in this study was conducted at anaerobic chamber (99.9% N2) unless otherwise noticed.
2.2.
Synthesis of FeS
FeS(s) was synthesized using the procedure reported in previous studies (Han et al., 2011, 2013). Briefly, the required doses of chemicals corresponding to stoichiometric equivalent molar amounts of FeCl2 and Na2S were weighted and then mixed in a 1-L polyethylene bottle to obtain 0.05-mol/L FeS and the suspension was aged for 3 days in an anaerobic chamber. However, a slight excess of sulfur element was observed in finally obtained FeS(s), as reported by previous study (Han et al., 2013). Thereafter, the aged amorphous FeS(s) (mackinawite) suspension was transferred into 30-mL centrifuge tubes and centrifuged for 30 min at 10,000 rpm (6700 G-force). The supernatant solution in the tubes was decanted and the remaining solids were collected in one bottle and washed with DDW and centrifuged again using the same procedure described above. This washing procedure was
312
w a t e r r e s e a r c h 5 3 ( 2 0 1 4 ) 3 1 0 e3 2 1
Scheme 1 e Schematic representation of FeS-supported dead-end ultrafiltration system for removal of Hg(II) (modified from (Millipore, 2004)) and flowchart of experimental procedures.
repeated at least three times to remove the remnant iron or sulfur chemicals from the suspension. Concentrations of Fe and sulfide in the washed FeS(s) suspension were measured using ICP/OES and a sulfide ion selective electrode (Orion, Thermo Scientific) with precision and accuracy of 0.55% and 98%. It was shown that there were no concentrations of Fe and sulfide in the supernatant of the centrifuged FeS(s) suspension. All experiment procedures described above were conducted at anaerobic chamber except for short period when the solids were centrifuged. The method of calculating the mass of solids obtained was described in detail in a previous study (Han et al., 2013). The FeS(s) suspension was stored in an anaerobic chamber until use.
2.3.
Kinetic experiment
A standard stock solution of mercury was prepared using HgCl2. To avoid formation of HgO(s), the mercury stock solution did not exceed a concentration of 1.3 mM. A kinetic test was initiated by mixing a FeS(s) slurry with the mercury stock solution in 20-mL vials to achieve concentrations of 0.05 g/L-FeS(s) and 500, 1000, and 1250 mM of Hg. Acid (0.5 M HNO3) or base (0.5 M NaOH) were added to adjust the pH 8. The reaction vessels were mixed with an end-over-end rotator until the specified sampling time (10, 30, 60, 150, 210, 330, 510, 750, 1440 min). Approximately 10 mL of suspension were removed and filtered using 0.02-mm anodisc membrane filters (Whatman). All samples were stored in an anaerobic chamber filled with 99.99% N2 prior to atomic absorption spectrometry (AAS) analysis in order to avoid oxidation of mercury and changes in pH.
2.4. Removal of Hg(II) using FeS-enhanced DE/UF system Experiments to evaluate continuous removal of Hg(II) were performed using a low pressure dead-end (DE) ultrafiltration
(UF) device that could be operated in stirred or non-stirred mode under N2. The setup of the DE/UF system is schematically represented in Scheme 1. This system included a reservoir of feed water, a UF reactor with a 30 kDa RC UF membrane, and an adapter box for control of gas and water flows. A series of experiments was conducted in order to evaluate sorption capacity of FeS(s) for Hg(II) during continuous filtration (Scheme 1). These experiments were conducted with the following steps: Step I) Hg(II) and FeS(s) are allowed to react for 15 min; Step II) separate Hg(II)-contacted FeS(s) from solution using DE/UF; Step III) contact Hg(II)/FeS(s)-deposited membrane with 0.1M thiosulfate solution (S2O2 3 ) to evaluate stability of Hg(II) on FeS(s); Step IV) apply additional feed solution to determine additional capacity of Hg(II)/FeS(s)deposited on the membrane for removal of Hg(II). For all steps except step I, permeate water was collected over time in order to measure flux, concentrations of Hg and Fe, and pH. In step I, 250 mL of stock solution containing 2.0 g/L of FeS(s) and the equivalent volume of stock solution containing 10 mM Hg(II) were prepared in volumetric flasks and their pH was adjusted to 8.0 0.5. The stock solutions were purged with N2 for 30 min and placed into an anaerobic chamber. To initiate the experiment, the two stock solutions were added to the UF reservoir and mixed to produce a final feed solution of 1.0 g/LFeS(s) and 5 mM-Hg(II). During adding two stock solutions to the UF reservoir, the solutions can be shortly exposed to air, but the inflow air in the reservoir was immediately evacuated several times under 1 bar of nitrogen gas by using a pressure control plug. However, it cannot be secured that the UF system is perfect anoxic condition at the time when the prepared solution was transferred to the UF reservoir. The feed solution was allowed to react in the UF reservoir for 15 min and then 300 mL of the feed solution was transferred into DE/UF cell by applying a pressure of 1 bar. This is the start of step II, when permeate water was collected over time until all feed solution was consumed. Then step III began by feeding a 0.1 M
w a t e r r e s e a r c h 5 3 ( 2 0 1 4 ) 3 1 0 e3 2 1
thiosulfate (S2O2 3 ) solution to the UF stirred cell in order to investigate the extent of desorption of Hg(II) from FeS(s) retained on the membrane. Finally, step IV was performed to evaluate additional sorption capacity of FeS(s) on the membrane for Hg(II) by applying a feed solution that contained 5 mM-Hg(II), but no FeS(s). Permeate water was collected and analyzed for Hg(II). When the filtration experiment was finished, the Hg/FeS(s)-deposited membrane was collected and stored in the anaerobic chamber until SEM/EDS analysis. The surface analysis of solids collected on the membrane was conducted using scanning emission microscopy (SEM)/ energy dispersive X-ray spectroscopy (EDS) in order to characterize surface morphology, element composition, and extent of cake-layer formation on the membrane surface.
2.5.
Analysis of aqueous-phase and solid-phase samples
Aqueous-phase samples of permeate water collected during filtration were analyzed for Hg, Fe, and pH as well as their volume used to calculate flux. The Hg concentration was measured using cold-vapor atomic absorbance spectrometry (CV-AAS). The average recovery (accuracy) and the relative standard deviation (precision) of the Hg measurement were 101.9% and 2.67%, respectively. A method detection limit (MDL) was calculated as 7.7 mg/L. Total Fe concentrations (solid þ solution) were measured by inductively coupled plasma optical emission spectrometry (ICP-OES) in filtered samples obtained during batch kinetic experiments and in the permeate water from UF experiments. The average recovery and the relative standard deviation of the Fe measurement were 98.8% and 2.85%, and the MDL was calculated as 11.3 mg/ L, respectively. The pH was measured using Thermo Triode pH meter calibrated with 3 pH buffers (4, 7, 10).
3.
Results and discussion
3.1.
Kinetics of Hg(II) removal
Fig. 1 shows the percentage of Hg(II) removed by FeS(s) as a function of time. The experiments were conducted with a 100
10 500 mM Hg(II) 1000 mM Hg(II) 1250 mM Hg(II)
% Hg(II) sorbed
96 94
8
6
92 4 90 88
2
Total Fe released (µ mol Fe/L)
98
86 0
200
400
600
800
1000
1200
1400
0 1600
Time, min
Fig. 1 e Percentage removal of Hg(II) and concentration of total Fe released as a function of time at pH 8 for three initial Hg(II) concentrations.
313
solids concentration of 0.05 g/L and different initial concentrations of Hg(II) (500, 1000, 1250 mM). Irrespective of reasonable experimental conditions that can reveal real environments contaminated with mercury, the reason for choosing high ratio of Hg(II) to FeS(s) is due to differentiating sorption behavior (i.e., fast and complete removal, slow and gradual removal) of Hg(II) over time because of the fact that even small amount of FeS(s) can quickly remove high concentration of mercury from water (Jeong et al., 2007, 2010; Liu et al., 2008). The experiment conducted with an initial concentration of 500 mM Hg(II) was observed to reach complete removal most rapidly and did so within 10 min. Experiments conducted with other initial concentrations required longer reaction times to reach more than 99% removal. In addition, the data in Fig. 1 shows biphasic removal, where removal is very fast initially, but slows later. The experiment with the lowest initial concentration (500 mM) shows that nearly complete removal was observed at the first sampling point. Experiments at higher initial concentrations showed removals near 94% at the first sampling point, followed by a slower approach to near complete removal. The biphasic sorption behavior was probably due to chemical interactions between mercury and the FeS(s) surface that are slower than the initial transport of mercury to the surface. The surface interactions could lead to formation of stronger bonds such as found in surface precipitates (Eq. (1)), discrete precipitates (Eq. (2)) or surface complexes (Eq. (3)), which have been found when the molar ratio of Hg(II) to FeS(s) is less than 1 or Hg(II) sorbs onto partly oxidized FeS(s) (Jeong et al., 2007; Skyllberg and Drott, 2010; Wolfenden et al., 2005). Herein, a surface complex formation by Eq. (3) was evident at molar ratio of Hg(II) to FeS less than 0.05, while precipitation of HgS(s) by Eq. (2) was shown to predominate at much higher molar ratio of Hg(II) to FeS(s), based on previous results (Jeong et al., 2007, 2010). However, one study shows that there was no surface complex formation between Hg(II) and FeS at even lower molar ratio ([Hg(II)]/[FeS]) of 0.002e0.012, rather postulating possible formation of precipitates (Skyllberg and Drott, 2010). The different observation can be caused by different experimental conditions such as aging time of FeS and ionic media (Cl-) as well as different loading ratio (Jeong et al., 2010; Skyllberg and Drott, 2010). Specifically, different ion media such as chloride has shown to slow down the surface reaction leading to formation of surface precipitate (HgS(s)) (Morse and Luther 1999; Skyllberg and Drott, 2010). Thus, many studies are still needed to clarify the formation of precipitates at various solution or solid conditions. (1) FeSðsÞ þ xHgðIIÞ/ Fe1x ; Hgx SðsÞ þ xFeðIIÞ FeSðsÞ þ HgðIIÞ/HgSðsÞ þ FeðIIÞ
(2)
hFeS þ HgðIIÞ/hFeS HgðIIÞ
(3)
In Eq. (3), FeS can be hydrolyzed over pH, resulting in formation of surface charge as similarly observed in metal oxides such as iron or aluminum oxide (Amirbahman et al., 2013; Butler and Hayes, 1998; Wolthers et al., 2005). In this study, the symbol of “hFeS” encompasses the charged FeS and also “Hg(II)” includes all types of divalent mercury complexed with inorganic anions such as Cl- or sulfur-containing oxyanions, if
314
w a t e r r e s e a r c h 5 3 ( 2 0 1 4 ) 3 1 0 e3 2 1
Since Hg(II) was shown to be successfully removed from solution by FeS(s) in batch kinetic tests, experiments were conducted to evaluate ultrafiltration as a means of removing the Hg-loaded FeS(s) particles from the solution to produce clean water. In the kinetic test, a wide range of molar ratio of Hg/ FeS(s) was applied by varying initial Hg concentration at given loading of FeS(s) in order to differentiate sorption behaviors. However, the DE/UF system chose the initial concentration of mercury with concentration level (
Available online at www.sciencedirect.com
ScienceDirect journal homepage: www.elsevier.com/locate/watres
Reactive iron sulfide (FeS)-supported ultrafiltration for removal of mercury (Hg(II)) from water Dong Suk Han a,*, Maria Orillano a, Ahmed Khodary a, Yuhang Duan b, Bill Batchelor b, Ahmed Abdel-Wahab a a b
Chemical Engineering Program, Texas A&M University at Qatar, Education City, Doha 23874, Qatar Zachry Department of Civil Engineering, Texas A&M University, College Station, TX 77843-3136, USA
article info
abstract
Article history:
This study investigated removal of Hg(II) from water using FeS(s) with batch and contin-
Received 28 September 2013
uous contact filtration systems. For the batch system, kinetic experiments showed that
Received in revised form
removal of Hg(II) by FeS(s) was rapid at lower concentration (500 mM), but at higher con-
15 January 2014
centration (1000 and 1250 mM), more time was required to achieve greater than 99%
Accepted 18 January 2014
removal. The concentration of iron released to the solution remained relatively low,
Available online 28 January 2014
typically below 3 mM. This would theoretically present less than 1% of the Hg(II) removed. Thus, a simple exchange of Hg(II) for Fe(II) in the solid (FeS(s)) does not explain the results,
Keywords:
but if the Fe(II) released could react to form another solids, low concentrations of Fe do not
FeS
preclude a mechanism in which Hg(II) reacts to form HgS and release Fe(II). A continuous
Mackinawite
contact dead-end ultrafiltration (DE/UF) system was developed to treat water containing
Mercury
Hg(II) by applying a FeS(s) suspension with stirred or non-stirred modes. A major reason for
Ultrafiltration
applying stirring to the system was to investigate the role of “shear” flow in rejection of
Adsorption
Hg(II)-contacted FeS(s) by a UF membrane and the stability of Hg on the FeS(s). The Hg(II)-
Desorption
contacted FeS(s) was completely rejected by the DE/UF system and mercury was strongly retained on the FeS(s) particles. Almost no release of Hg(II) (z0 mM) from the FeS(s) solids was observed when they were contacted with 0.1M-thiosulfate, regardless of whether the system was operated in stirred or non-stirred mode. However, rapid oxidation of FeS(s) was observed in the stirred system but not in the non-stirred system. Determining the mechanism of oxidation requires further study, but it is important because oxidation reduces the ability of the solids to remove additional Hg(II). ª 2014 Elsevier Ltd. All rights reserved.
1.
Introduction
Mercury has been considered to be a global contaminant of significant concern for centuries due to its high toxicity and potential for bioaccumulation via the aquatic food chain, which seriously affects natural ecosystems and the health of humans. Exposure to high levels of mercury can cause
inhibition of enzyme activity, cell damage, impairment of pulmonary function and kidney performance, chest pain, and damage to the central nervous system (Jeong et al., 2007; Kaneko, 1988; Zahir et al., 2005). Mercury contamination occurs due to release from anthropogenic and natural sources, including chloro-alkali plants, mining and smelting activities, coal-fired power plants, electrical and electronic
* Corresponding author. Tel.: þ974 6629 1478; fax: þ974 4423 0065. E-mail address: [email protected] (D.S. Han). 0043-1354/$ e see front matter ª 2014 Elsevier Ltd. All rights reserved. http://dx.doi.org/10.1016/j.watres.2014.01.033
w a t e r r e s e a r c h 5 3 ( 2 0 1 4 ) 3 1 0 e3 2 1
manufacturing plants, and a variety of incinerator facilities (Vieira and Beppu, 2006; Weisener et al., 2005). Among them, coal-fired power plants and incinerator facilities are the major sources with emissions of 120 ton per year which constitutes 77% of anthropogenic Hg emissions in the United States (Liu et al., 1998). Removal of inorganic Hg from emissions is a crucial process because in aquatic ecosystems, most inorganic Hg is transformed to methylmercury (MeHg), which leads to highly elevated concentrations in aquatic fish and wildlife (Hsu-Kim et al., 2013; Skyllberg et al., 2006). Hg(II) is a soft Lewis acid that preferentially bonds with soft Lewis bases. Since the thiol functional group is a soft base, sulfur-containing chemicals have been widely used to remove mercury (Biester and Zimmer, 1998; Chen et al., 2004; Gibson et al., 2011; Wolfenden et al., 2005). For the same reason, mercury forms a very insoluble solid with sulfide (Brown et al., 1979; Ebadian et al., 2001; Nriagu, 1979; Sadiq, 1992). However, there are problems associated with controlling the sulfide dose in treatment processes. If excess sulfide is added, then soluble HgeS complexes are formed and the effluent concentrations of soluble mercury increases (U.S. Environmental Protection Agency, 1997). In general, the solubility of Hgsulfide solid phases depends on pH and sulfide concentration. At low pH and low sulfide concentration, formation of insoluble mercuric sulfide solid phase (HgS) is preferred, whereas soluble HgeS complexes occur at high pH and high sulfide concentrations (Ravichandran et al., 1999). A way of avoiding the problems that occur in precipitation processes that use soluble reagents is to dose with large particles of iron sulfide, which gradually dissolve to produce the needed sulfide (Nriagu, 1979). If smaller particles are added, the removal mechanism will include sorption and surface reaction. Nano-scale iron sulfide particles have been produced microbially or abiotically and they have been applied to removal of mercury (Ito et al., 2004; Watson et al., 2001, 1995). Particle sizes around 2e5 nm and specific surface areas of 280e500 m2/g have been reported (Watson et al., 2001). Evidence for the production of solid phase products via surface reaction is found in the decrease in soluble mercury concentration with time (1e24 h) that becomes more pronounced as the surface concentration is increased (Watson et al., 1995). Other metals besides Hg were shown to be removed by biogenic FeS(s) (Watson et al., 2001, 1995). Chemical synthesis of nano-sized FeS has also been reported (Butler and Hayes, 1998; Martellaro et al., 2001). Others have also reported that iron sulfides are good reagents for removing mercury from solution (Behra et al., 2001; Bower et al., 2008; Brown et al., 1979; Ehrhardt et al., 2000; Gong et al., 2012; Han et al., 2012; Jean and Bancroft, 1986; Jeong et al., 2007, 2010; Liu et al., 2008; Skyllberg and Drott, 2010). For example, Liu et al. (2008) investigated interactions between aqueous Hg(II) and FeS(s) in batch sorption experiments and found that the maximum removal capacity at lower pH was approximately 0.75 mol Hg(II)/mol FeS. They also used X-ray powder diffraction (XRPD) to identify the major products as metacinnabar, cinnabar, and mercury iron sulfide with less solubility (Ksp ¼ 2E-53) (Barnett et al., 2001; Liu et al., 2008). In this respect, iron sulfide particle has been widely used as immobilizer or methylation inhibitor for Hg(II) in anoxic environment such as estuaries or groundwater (Liu et al., 2009), but
311
less in industrial wastewater released from gas or coal-fired power plant requiring appropriate separation and disposal of final residual solids. However, if appropriate tool or process to safely treat the final solids is provided, use of iron sulfide may have a great merit in such environment. The purpose of this work is to remove Hg(II) in water using reactive FeS(s) with an ultrafiltration system to produce purified water and stable final residual solids that can be greatly rejected by the ultrafiltration (UF) membrane. Hg would be stabilized on the solids by surface reactions after removal by adsorption. In order to meet this goal, we performed a batch kinetic test to evaluate removal of Hg(II) and its sorption behavior by synthetic mackinawite (FeS(s)) as affected by mercury concentration and contact time. Then, lab-scale UF filtration was performed to produce purified water by separating Hg(II)-contacted FeS(s) from the synthetic feed water. In addition, the solid phases deposited on the membrane were evaluated to better understand the chemical changes that occur and how they improve stability of Hg(II) on the FeS(s).
2.
Materials and method
2.1.
Materials
All materials and chemicals applied in this study were used as received. The chemicals used here were: FeCl2 (SigmaeAldrich, 99.9%), Na2S (SigmaeAldrich, 99.9%), HgCl2 (SigmaeAldrich, 99.99%), and Na2S2O3 (SigmaeAldrich, 98%). Components of the UF system were a 300-mL UF stirred cell (Millipore) and 30 kDa regenerated cellulose (RC) ultrafiltration membrane with surface area of 31.7 cm2 (Millipore) (Scheme 1). The water used in this study was prepared through several steps. First, distilled water was collected from a Barnstead Mega-pure distillation device (Thermo Co.) and deionized by a Labconco purifier system. Then, the deionized water was deoxygenated for 2 h by purging with nitrogen gas (99.99% of purity). The deionized/deoxygenated water (DDW) obtained in this way was used for the FeS(s) synthesis experiments and the Hg(II) removal experiments. All experiment in this study was conducted at anaerobic chamber (99.9% N2) unless otherwise noticed.
2.2.
Synthesis of FeS
FeS(s) was synthesized using the procedure reported in previous studies (Han et al., 2011, 2013). Briefly, the required doses of chemicals corresponding to stoichiometric equivalent molar amounts of FeCl2 and Na2S were weighted and then mixed in a 1-L polyethylene bottle to obtain 0.05-mol/L FeS and the suspension was aged for 3 days in an anaerobic chamber. However, a slight excess of sulfur element was observed in finally obtained FeS(s), as reported by previous study (Han et al., 2013). Thereafter, the aged amorphous FeS(s) (mackinawite) suspension was transferred into 30-mL centrifuge tubes and centrifuged for 30 min at 10,000 rpm (6700 G-force). The supernatant solution in the tubes was decanted and the remaining solids were collected in one bottle and washed with DDW and centrifuged again using the same procedure described above. This washing procedure was
312
w a t e r r e s e a r c h 5 3 ( 2 0 1 4 ) 3 1 0 e3 2 1
Scheme 1 e Schematic representation of FeS-supported dead-end ultrafiltration system for removal of Hg(II) (modified from (Millipore, 2004)) and flowchart of experimental procedures.
repeated at least three times to remove the remnant iron or sulfur chemicals from the suspension. Concentrations of Fe and sulfide in the washed FeS(s) suspension were measured using ICP/OES and a sulfide ion selective electrode (Orion, Thermo Scientific) with precision and accuracy of 0.55% and 98%. It was shown that there were no concentrations of Fe and sulfide in the supernatant of the centrifuged FeS(s) suspension. All experiment procedures described above were conducted at anaerobic chamber except for short period when the solids were centrifuged. The method of calculating the mass of solids obtained was described in detail in a previous study (Han et al., 2013). The FeS(s) suspension was stored in an anaerobic chamber until use.
2.3.
Kinetic experiment
A standard stock solution of mercury was prepared using HgCl2. To avoid formation of HgO(s), the mercury stock solution did not exceed a concentration of 1.3 mM. A kinetic test was initiated by mixing a FeS(s) slurry with the mercury stock solution in 20-mL vials to achieve concentrations of 0.05 g/L-FeS(s) and 500, 1000, and 1250 mM of Hg. Acid (0.5 M HNO3) or base (0.5 M NaOH) were added to adjust the pH 8. The reaction vessels were mixed with an end-over-end rotator until the specified sampling time (10, 30, 60, 150, 210, 330, 510, 750, 1440 min). Approximately 10 mL of suspension were removed and filtered using 0.02-mm anodisc membrane filters (Whatman). All samples were stored in an anaerobic chamber filled with 99.99% N2 prior to atomic absorption spectrometry (AAS) analysis in order to avoid oxidation of mercury and changes in pH.
2.4. Removal of Hg(II) using FeS-enhanced DE/UF system Experiments to evaluate continuous removal of Hg(II) were performed using a low pressure dead-end (DE) ultrafiltration
(UF) device that could be operated in stirred or non-stirred mode under N2. The setup of the DE/UF system is schematically represented in Scheme 1. This system included a reservoir of feed water, a UF reactor with a 30 kDa RC UF membrane, and an adapter box for control of gas and water flows. A series of experiments was conducted in order to evaluate sorption capacity of FeS(s) for Hg(II) during continuous filtration (Scheme 1). These experiments were conducted with the following steps: Step I) Hg(II) and FeS(s) are allowed to react for 15 min; Step II) separate Hg(II)-contacted FeS(s) from solution using DE/UF; Step III) contact Hg(II)/FeS(s)-deposited membrane with 0.1M thiosulfate solution (S2O2 3 ) to evaluate stability of Hg(II) on FeS(s); Step IV) apply additional feed solution to determine additional capacity of Hg(II)/FeS(s)deposited on the membrane for removal of Hg(II). For all steps except step I, permeate water was collected over time in order to measure flux, concentrations of Hg and Fe, and pH. In step I, 250 mL of stock solution containing 2.0 g/L of FeS(s) and the equivalent volume of stock solution containing 10 mM Hg(II) were prepared in volumetric flasks and their pH was adjusted to 8.0 0.5. The stock solutions were purged with N2 for 30 min and placed into an anaerobic chamber. To initiate the experiment, the two stock solutions were added to the UF reservoir and mixed to produce a final feed solution of 1.0 g/LFeS(s) and 5 mM-Hg(II). During adding two stock solutions to the UF reservoir, the solutions can be shortly exposed to air, but the inflow air in the reservoir was immediately evacuated several times under 1 bar of nitrogen gas by using a pressure control plug. However, it cannot be secured that the UF system is perfect anoxic condition at the time when the prepared solution was transferred to the UF reservoir. The feed solution was allowed to react in the UF reservoir for 15 min and then 300 mL of the feed solution was transferred into DE/UF cell by applying a pressure of 1 bar. This is the start of step II, when permeate water was collected over time until all feed solution was consumed. Then step III began by feeding a 0.1 M
w a t e r r e s e a r c h 5 3 ( 2 0 1 4 ) 3 1 0 e3 2 1
thiosulfate (S2O2 3 ) solution to the UF stirred cell in order to investigate the extent of desorption of Hg(II) from FeS(s) retained on the membrane. Finally, step IV was performed to evaluate additional sorption capacity of FeS(s) on the membrane for Hg(II) by applying a feed solution that contained 5 mM-Hg(II), but no FeS(s). Permeate water was collected and analyzed for Hg(II). When the filtration experiment was finished, the Hg/FeS(s)-deposited membrane was collected and stored in the anaerobic chamber until SEM/EDS analysis. The surface analysis of solids collected on the membrane was conducted using scanning emission microscopy (SEM)/ energy dispersive X-ray spectroscopy (EDS) in order to characterize surface morphology, element composition, and extent of cake-layer formation on the membrane surface.
2.5.
Analysis of aqueous-phase and solid-phase samples
Aqueous-phase samples of permeate water collected during filtration were analyzed for Hg, Fe, and pH as well as their volume used to calculate flux. The Hg concentration was measured using cold-vapor atomic absorbance spectrometry (CV-AAS). The average recovery (accuracy) and the relative standard deviation (precision) of the Hg measurement were 101.9% and 2.67%, respectively. A method detection limit (MDL) was calculated as 7.7 mg/L. Total Fe concentrations (solid þ solution) were measured by inductively coupled plasma optical emission spectrometry (ICP-OES) in filtered samples obtained during batch kinetic experiments and in the permeate water from UF experiments. The average recovery and the relative standard deviation of the Fe measurement were 98.8% and 2.85%, and the MDL was calculated as 11.3 mg/ L, respectively. The pH was measured using Thermo Triode pH meter calibrated with 3 pH buffers (4, 7, 10).
3.
Results and discussion
3.1.
Kinetics of Hg(II) removal
Fig. 1 shows the percentage of Hg(II) removed by FeS(s) as a function of time. The experiments were conducted with a 100
10 500 mM Hg(II) 1000 mM Hg(II) 1250 mM Hg(II)
% Hg(II) sorbed
96 94
8
6
92 4 90 88
2
Total Fe released (µ mol Fe/L)
98
86 0
200
400
600
800
1000
1200
1400
0 1600
Time, min
Fig. 1 e Percentage removal of Hg(II) and concentration of total Fe released as a function of time at pH 8 for three initial Hg(II) concentrations.
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solids concentration of 0.05 g/L and different initial concentrations of Hg(II) (500, 1000, 1250 mM). Irrespective of reasonable experimental conditions that can reveal real environments contaminated with mercury, the reason for choosing high ratio of Hg(II) to FeS(s) is due to differentiating sorption behavior (i.e., fast and complete removal, slow and gradual removal) of Hg(II) over time because of the fact that even small amount of FeS(s) can quickly remove high concentration of mercury from water (Jeong et al., 2007, 2010; Liu et al., 2008). The experiment conducted with an initial concentration of 500 mM Hg(II) was observed to reach complete removal most rapidly and did so within 10 min. Experiments conducted with other initial concentrations required longer reaction times to reach more than 99% removal. In addition, the data in Fig. 1 shows biphasic removal, where removal is very fast initially, but slows later. The experiment with the lowest initial concentration (500 mM) shows that nearly complete removal was observed at the first sampling point. Experiments at higher initial concentrations showed removals near 94% at the first sampling point, followed by a slower approach to near complete removal. The biphasic sorption behavior was probably due to chemical interactions between mercury and the FeS(s) surface that are slower than the initial transport of mercury to the surface. The surface interactions could lead to formation of stronger bonds such as found in surface precipitates (Eq. (1)), discrete precipitates (Eq. (2)) or surface complexes (Eq. (3)), which have been found when the molar ratio of Hg(II) to FeS(s) is less than 1 or Hg(II) sorbs onto partly oxidized FeS(s) (Jeong et al., 2007; Skyllberg and Drott, 2010; Wolfenden et al., 2005). Herein, a surface complex formation by Eq. (3) was evident at molar ratio of Hg(II) to FeS less than 0.05, while precipitation of HgS(s) by Eq. (2) was shown to predominate at much higher molar ratio of Hg(II) to FeS(s), based on previous results (Jeong et al., 2007, 2010). However, one study shows that there was no surface complex formation between Hg(II) and FeS at even lower molar ratio ([Hg(II)]/[FeS]) of 0.002e0.012, rather postulating possible formation of precipitates (Skyllberg and Drott, 2010). The different observation can be caused by different experimental conditions such as aging time of FeS and ionic media (Cl-) as well as different loading ratio (Jeong et al., 2010; Skyllberg and Drott, 2010). Specifically, different ion media such as chloride has shown to slow down the surface reaction leading to formation of surface precipitate (HgS(s)) (Morse and Luther 1999; Skyllberg and Drott, 2010). Thus, many studies are still needed to clarify the formation of precipitates at various solution or solid conditions. (1) FeSðsÞ þ xHgðIIÞ/ Fe1x ; Hgx SðsÞ þ xFeðIIÞ FeSðsÞ þ HgðIIÞ/HgSðsÞ þ FeðIIÞ
(2)
hFeS þ HgðIIÞ/hFeS HgðIIÞ
(3)
In Eq. (3), FeS can be hydrolyzed over pH, resulting in formation of surface charge as similarly observed in metal oxides such as iron or aluminum oxide (Amirbahman et al., 2013; Butler and Hayes, 1998; Wolthers et al., 2005). In this study, the symbol of “hFeS” encompasses the charged FeS and also “Hg(II)” includes all types of divalent mercury complexed with inorganic anions such as Cl- or sulfur-containing oxyanions, if
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w a t e r r e s e a r c h 5 3 ( 2 0 1 4 ) 3 1 0 e3 2 1
Since Hg(II) was shown to be successfully removed from solution by FeS(s) in batch kinetic tests, experiments were conducted to evaluate ultrafiltration as a means of removing the Hg-loaded FeS(s) particles from the solution to produce clean water. In the kinetic test, a wide range of molar ratio of Hg/ FeS(s) was applied by varying initial Hg concentration at given loading of FeS(s) in order to differentiate sorption behaviors. However, the DE/UF system chose the initial concentration of mercury with concentration level (
