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Article

Liberation of Adsorbed and Co-Precipitated Arsenic from Jarosite, Schwertmannite, Ferrihydrite, and Goethite in Seawater

by
Rodrigo Alarcón
1,*,
Jenny Gaviria
2 and
Bernhard Dold
1,2,3
1
Geology Department, University of Chile, Plaza Ercilla #803, Casilla 13518, Correo 21, Santiago 8320000, Chile
2
Institute of Applied Economic Geology (GEA), University of Concepcion, Casilla 160-C, Concepción 4030000, Chile
3
Present Address: SUMIRCO (Sustainable Mining Research & Consult EIRL), Casilla 28, San Pedro de la Paz 4130000, Chile
*
Author to whom correspondence should be addressed.
Minerals 2014, 4(3), 603-620; https://doi.org/10.3390/min4030603
Submission received: 10 March 2014 / Revised: 17 June 2014 / Accepted: 18 June 2014 / Published: 8 July 2014
(This article belongs to the Special Issue Mine Waste Characterization, Management and Remediation)
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Abstract

:
Sea level rise is able to change the geochemical conditions in coastal systems. In these environments, transport of contaminants can be controlled by the stability and adsorption capacity of iron oxides. The behavior of adsorbed and co-precipitated arsenic in jarosite, schwertmannite, ferrihydrite, and goethite in sea water (common secondary minerals in coastal tailings) was investigated. The aim of the investigation was to establish As retention and transport under a marine flood scenario, which may occur due to climate change. Natural and synthetic minerals with co-precipitated and adsorbed As were contacted with seawater for 25 days. During this period As, Fe, Cl, SO4, and pH levels were constantly measured. The larger retention capability of samples with co-precipitated As, in relation with adsorbed As samples, reflects the different kinetics between diffusion, dissolution, and surface exchange processes. Ferrihydrite and schwertmannite showed good results in retaining arsenic, although schwertmannite holding capacity was enhanced due its buffering capacity, which prevented reductive dissolution throughout the experiment. Arsenic desorption from goethite could be understood in terms of ion exchange between oxides and electrolytes, due to the charge difference generated by a low point-of-zero-charge and the change in stability of surface complexes between synthesis conditions and natural media.

Graphical Abstract

1. Introduction

Arsenic is one of the most toxic inorganic pollutants in aquatic systems [1]. Its source is mainly geogenic, entering the environment through volcanic emissions, hydrothermal systems [2], mineral erosion, or by reductive dissolution of iron hydroxides, such as in the case of Bangladesh [3]. Although the input from anthropogenic sources (mining, fossil fuels) is lower than natural sources, these can have a strong impact, generating local pollution episodes [2]. There are mainly three mechanisms that explain arsenic liberation into natural environments [4]: (1) oxidation of arsenic sulphides, (2) competitive desorption, and (3) reductive dissolution of iron oxides. Although each has been extensively investigated under several experimental conditions, there is still a lack of knowledge as to how these processes take place in specific environments.
Some climatic models predict that, by the end of the century, one of the consequences of global warming will be a sea level rise of up to 1.1 m above the current level [5]. This sea level change will establish a new geochemical context in the oxidation zone of mine tailings and acid sulphate soils near the shoreline, where the stability of sorbent minerals under seawater flood would be uncertain. In this paper, the effects of seawater intrusion on arsenic adsorbed to or co-precipitated with major sorbent minerals, present in acid soils or mine tailings, are simulated. Synthesized jarosite, schwertmannite, goethite, and ferrihydrite were used to represent different arsenic uptake scenarios. The aim of these experiments is to establish the stability and transport of this metalloid in these coastal environments under changing environmental conditions.
Iron oxides behave as excellent sorbents for a variety of contaminants in almost every environment. They have been widely used in the extraction of heavy metals from natural [6] and industrial sources [7], which has generated a great number of studies about mechanisms of surface interactions [8,9,10,11,12]. These minerals are common in environments with high availability of metal sulphides exposed to oxidizing conditions, as in acid sulphate soils, in the oxidation zone of mine tailings, or ore deposits [13,14]. In these systems, oxidation of the primary sulphide (pyrite) is capable of releasing large amounts of protons, sulphate, Fe(II), and trace metals (Equation (1)) [15]. Once Fe(III) is produced by oxidation of Fe(II), which may be strongly accelerated by microbial activity under low pH conditions, Fe(III) becomes the primary oxidant of pyrite [16].
FeS2 + 7/2O2 + H2O → Fe2+ + 2SO42− + 2H+
Fe2+ + 1/4O2 + H+ ↔ Fe3+ + ½ H2O
At circumneutral pH, ferrihydrite (5Fe2O3∙9(H2O)) is commonly the first mineral phase to precipitate from hydrolysis of ferric solutions [17]. Ferrihydrite is able to transform into crystalline phases of higher thermodynamic stability as goethite (α-FeOOH) or hematite (Fe2O3) [18]. When the pH of the system is acidic enough (pH 2–4) and there is a high sulphate concentration, jarosite (Na, K, H3O)∙[Fe3(SO4)2(OH)6]) (pH ~ 2) and schwertmannite (Fe16O16(OH)12∙(SO4)2) constitute the main phases [13]. Between pH 2.5 and 4, schwertmannite is, perhaps, the principal secondary mineral that forms from acid drainage [19,20].
In natural environments, redox processes commonly control the solubility of iron oxides. This occurs when interaction occurs between dissolved species, such as H+, OH− or other metal ions, with the hydroxyl groups present on the oxides surfaces [21,22]. Adsorption and formation of surface complexes with reducing species is a reaction that generates an electron transfer, reducing Fe(III) to Fe(II) [21]. Weaker Fe(II) bonds enhance reductive dissolution and release of species from oxides surfaces. For example, in Alberta (Canada) microbial activity in acid sulphate soils released significant amounts of As when the Eh was below +100 mV [23]. Mine tailings at the former Delnite gold mine in Northern Ontario showed that As(V) reduction and mobilization occurred due to reductive dissolution of goethite, influenced by a biosolid-cover [24]. Reductive dissolution also takes place in remediated marine shore tailings deposits [14]. Iron oxides under reductive conditions formed a Fe-Mn plume, which developed toward the shoreline where oxidizing and higher pH conditions triggered Fe(III) oxide-hydroxide precipitation.
When arsenic is present during hydrolysis of Fe(III) it can co-precipitate. In some minerals it can form part of the structure, for instance in jarosite at the TO4 site [25,26], or it can be adsorbed as surface complexes [27]. In relation to iron oxides, arsenate tends to be adsorbed as outer-sphere complexes [28], while arsenite can be adsorbed either as inner or outer-sphere complexes [28,29].
Different studies using Extended X-ray Absorption Fine Structure (EXAFS) and Fourier transform infrared spectroscopy (FTIR) have attempted to establish the possible bond between arsenic and iron oxide surfaces without reaching consensus to date. Waychunas et al. [30] stated that, due to thermodynamic constraints, there is only a low chance for the formation of mononuclear monodentate and bidentate complexes. Some authors established that bidentate binuclear complexes are most likely to form due to their higher thermodynamic stability [31,32,33]. MICRO-EXAFS spectra of individual Fe(III) oxy-hydroxide grains point to inner-sphere bidentate-binuclear forms as the predominant As(V) complex and the existence of a second sphere corresponding mainly to bidentate-mononuclear [34]. Furthermore, other authors postulate that As complexes at the surface of goethite would be exclusively monodentate complexes [35] and when the As load increases significantly or pH level is greater than 6, then, only bidentate binuclear complexes are formed [30,33,36].
Some species can compete with As for available vacancies restricting the metalloid adsorption. Studies on ferrihydrite indicate that the extent of adsorption can be affected by ionic competition, mainly by PO4 >> organic ligands > SO4 > Cl [37,38]. When As is part of the structure of iron oxides, the long term release is controlled by media dissolution, pH-dependent sorption/desorption, ion exchange, and transformation processes [27,39].
Tailings disposal in bays and beaches has been a widely used practice in the past and is still performed in places like Papua New Guinea and Indonesia [40]. Lihir gold mine on Niolam Island in Papua New Guinea annually produces over 35 Mt of waste rock which are dumped into nearshore deepwater valleys, and about 100,000 ML of post-processing tailings slurry deposited at depth from a sub-surface pipeline [41]. By representing a potential risk to marine ecosystems, one of the main challenges in this field is to determine the geochemical stability and pollution potential of tailings in flooded environments [42].
In Chañaral (Chile) about 220 Mt of tailings fill the homonymous bay covering an area of about 4 km2. Dold [43] indicated the presence of an oxidation zone (>1 m) with significant amounts of arsenic associated with secondary Fe(III) oxide-hydroxides and jarosite. The instability of the sorbent phases and the intrusion of seawater, a product of variations within the coastal cycle, promoted the release of dissolved As, Mo, and colloidal adsorbed Cu and Zn into the ocean, generating an impact over the meiofaunal assemblages through increasing the bioavailability of heavy metals [44] and there has been no sign of recovery 35 years after ceasing disposal [45]. During the summer, high rates of evaporation promote capillary transport of Cu and Zn, which precipitate as efflorescent salts on the tailings surface, increasing the risk of metal exposure for the local community [43,46]. Remediation experiences of similar systems have been successfully implemented in environments where the hydrological characteristics allowed it (e.g., Bahia de Ite, Peru [14]).
Concentrations of arsenic in open ocean seawater usually do not exceed 2 µg/L [47]. It is found mainly as arsenate (As (V)), although, concentration of arsenite (As(III)) in coastal waters affected by anthropogenic activity can be as high as 19% of AsTotal [48]. Arsenic in oxidizing environments can be mainly found as protonated species of AsO43−. Biological activity plays an important role in marine arsenic speciation, reducing arsenate to arsenite, given the low thermodynamic stability of arsenite in oxidizing environments (As(III)/As(V) ≈ 10−26) [49]. It also plays a role in the formation of monomethylarsenic acid and dimethylarsenic acid; however, the concentration of this arsenic species is much smaller than the inorganic forms.
Low concentrations of organic matter present in seawater (1–3 mg/L CH3O [50]) do not allow reductive dissolution to be an effective mechanism for As release, though in coastal environments organic matter is usually more abundant than that found in the open ocean, which could be an important factor in speciation and solubility of iron [51].
Although As(III) species are mainly limited by biological processes, As(III) can be found in large quantities in polluted mining areas, especially in reducing groundwater environments [48]; however, it was not considered for the purpose of this study.

2. Materials and Methods

Sorbent minerals used for this experiment were synthesized according to several procedures [13,17,52,53,54,55,56,57]. No sample of schwertmannite with co-precipitated arsenic was synthesized, instead, a natural sample form the acid mine drainage of Monte Romero mine (Iberian Pyrite Belt, Spain) was used [58]. Incorporation of arsenic was considered to be by surface adsorption and co-precipitation. Every mineral was characterized using a Rigaku RADII-C X-ray diffractometer (XRD) (35 kV, 15 mA) from 20° to 100° 2θ using a range of 0.05° 2θ and a counting time of 10 s per step. X-ray diffraction patterns confirmed the correct synthesis for the procedures used and are presented in Figure 1 and Figure 2. Further information about the experimental procedure can be found in the supplementary information section. In order to simulate As release from Fe(III) oxide-hydroxides and jarosite during seawater intrusion, the synthesized minerals were contacted with seawater during a 25-day period.
Minerals 04 00603 g001 1024
Figure 1. X-ray diffraction pattern for jarosite (jt), goethite (gt), schwertmanntite (sh), and ferrihydrite (fh).
Figure 1. X-ray diffraction pattern for jarosite (jt), goethite (gt), schwertmanntite (sh), and ferrihydrite (fh).
Minerals 04 00603 g001
Minerals 04 00603 g002 1024
Figure 2. X-ray diffraction pattern for goethite (gt) and schwertmannite (sh-gt) with adsorbed As.
Figure 2. X-ray diffraction pattern for goethite (gt) and schwertmannite (sh-gt) with adsorbed As.
Minerals 04 00603 g002
Twenty-five liters of seawater were taken 10 km offshore Concepción, Chile, by a research vessel. The chemical analysis of the seawater was carried out using titration methods with AgNO3 for Cl− and turbidimetry for SO42−. As and Fe were both measured using an atomic absorption spectrometer (AAS), model Hitachi Z-8100 Polarized Zeeman (Hitachi, Tokyo, Japan). For the Fe analysis, Chelex-100 resin was used to pre-concentrate the sample with HNO3, which was then flushed to 10 mL. Arsenic hydrides were generated in the presence of H2SO4, HNO3, and HClO4 which were later used for readings (Table 1 and Table 2). The respective detection limits were 5 µg/L for Fe and 1 µg/L for As.
Adsorbed and co-precipitated arsenic, associated with mineral samples, was measured using the same instrumental techniques as for sea water but with different acids for liberating the arsenic. Arsenic was detected in all samples as concentrations were above the detection limit. For every mineral used in this experiment, 1.5 g of the synthesized material (natural schwertmannite for co-precipitated As) was brought into contact with 400 mL of seawater in sealed vessels for 596 h (~25 days). During this period, continuous stirring was applied (Figure 3). Over the duration of the experiments, 14 extractions (15 mL each) were taken, of which about 40% took place within the first 24 h. This was decided in order to track the period with the greatest release kinetics [59]. Every extraction was filtered using a 0.45 µm membrane filter, acidified with 50 µL of HNO3 and stored at 4 °C until analysis. Concentrations of the released species are available in Table 3, Table 4, Table 5 and Table 6. One of the limitations of this study is that pH and extractions were made at different times so there is some uncertainty regarding the correlation of these variables. However, long term trends, indicative of their relative behavior, can be identified in the figures and Table 7. All analyses were conducted in GEA Facilities (Instituto de Geología Económica Aplicada, Universidad de Concepción, Chile). In this work, computations involving arsenic and surface speciation were performed using PhreePlot [60]. Thermodynamic constants were taken from the wateq4f database. For the modeling of surface speciation, a charge distribution multisite ion complexation (CDMUSIC) model was chosen.
Table
Table 1. Composition of sea water used in this study and comparison with literature values.
Table 1. Composition of sea water used in this study and comparison with literature values.
Ions/MetalsThis workNordstrom [61]Turekian [62]
Cl (g/L)19.519.3519.4
SO4 (g/L)2.62.712.58
As (µg/L)1-2.6
Fe (mg/L)<0.0080.0020.0034
Table
Table 2. Arsenic load in minerals used for this study.
Table 2. Arsenic load in minerals used for this study.
MineralAs (wt %)As(mol/g)As (mg/kg)
SchwertmanniteCo-precipitated0.567.47 × 10−521
Adsorbed0.811.08 × 10−430.375
FerrihydriteCo-precipitated2.973.96 × 10−4111.375
Adsorbed2.963.95 × 10−4111
JarositeCo-precipitated0.162.14 × 10−56
Adsorbed0.385.07 × 10−514.25
GoethiteCo-precipitated0.719.48 × 10−526.625
Adsorbed3.344.46 × 10−4125.25
Minerals 04 00603 g003 1024
Figure 3. Magnetic stirrer with vessels containing schwertmannite samples.
Figure 3. Magnetic stirrer with vessels containing schwertmannite samples.
Minerals 04 00603 g003
Table
Table 3. As release from ferrihydrite.
Table 3. As release from ferrihydrite.
Experiment typeWithout AsCo-precipitated AsAdsorbed As
HoursClSO4FeAsClSO4FeAsClSO4FeAs
(g/L)(g/L)(mg/L)(µg/L)(g/L)(g/L)(mg/L)(µg/L)(g/L)(g/L)(mg/L)(µg/L)
1019.782.46<0.05<119.922.37<0.0525319.621.53<0.051227
1019.452.33<0.05<119.432.49<0.0526019.362.48<0.051556
1119.732.25<0.05<119.552.36<0.0526019.722.29<0.051533
1319.552.35<0.05<119.942.36<0.0525819.732.30<0.051366
1719.412.45<0.05<119.652.24<0.0526319.802.25<0.051254
2321.742.39<0.05519.372.47<0.0522419.632.47<0.051088
3620.652.35<0.05219.082.31<0.0522619.442.26<0.05950
5619.562.22<0.05219.092.06<0.0521719.582.28<0.05849
8519.452.63<0.05819.632.57<0.0518918.932.51<0.05836
11919.462.32<0.05-19.122.17<0.0518819.392.46<0.05692
18019.502.24<0.05<119.502.63<0.0517319.502.44<0.05693
26019.502.39<0.05<119.502.34<0.0517819.502.55<0.05682
41619.502.14<0.05<119.502.75<0.0511819.502.11<0.05308
59619.502.43<0.05<119.502.72<0.057619.502.34<0.05273
Table
Table 4. As release from schwertmannite.
Table 4. As release from schwertmannite.
Experiment typeWithout AsCo-precipitated AsAdsorbed As
HoursClSO4FeAsClSO4FeAsClSO4FeAs
(g/L)(g/L)(mg/L)(µg/L)(g/L)(g/L)(mg/L)(µg/L)(g/L)(g/L)(mg/L)(µg/L)
1019.732.31<0.05<121.292.50<0.05<119.642.65<0.05<1
1019.512.76<0.05<120.002.40<0.05<120.282.42<0.05<1
1119.732.62<0.05<119.682.27<0.05219.942.52<0.05<1
1319.922.38<0.05<118.932.43<0.05<119.122.48<0.05<1
1719.352.52<0.05<119.212.38<0.05<118.962.40<0.053
2319.092.70<0.05<119.052.640.42919.162.43<0.052
3619.152.56<0.05<118.962.450.14220.302.500.967
5619.232.550.16<119.132.490.17<119.102.77<0.051
8519.332.540.24<119.152.690.18119.222.21<0.052
11919.992.240.20119.362.510.27<119.032.69<0.052
18019.502.640.20<119.502.930.26119.502.76<0.053
26019.502.680.26719.502.750.29219.502.70<0.052
41619.502.490.20219.502.290.33319.502.20<0.052
59619.502.350.10219.502.390.42219.502.32<0.051
Table
Table 5. As release from goethite.
Table 5. As release from goethite.
Experiment typeWithout AsCo-precipitated AsAdsorbed As
HoursClSO4FeAsClSO4FeAsClSO4FeAs
(g/L)(g/L)(mg/L)(µg/L)(g/L)(g/L)(mg/L)(µg/L)(g/L)(g/L)(mg/L)(µg/L)
1019.542.38<0.05<118.972.97<0.055419.632.20<0.056,572
1019.442.37<0.05<119.592.47<0.055419.872.41<0.059,378
1119.442.14<0.05<119.542.35<0.055319.602.19<0.0515,404
1319.302.35<0.05<119.442.62<0.055119.252.39<0.0519,721
1720.282.20<0.05<119.442.48<0.054719.282.44<0.0520,620
2319.272.20<0.05<119.393.19<0.054819.252.43<0.0525,776
3619.352.14<0.05<119.972.48<0.054819.532.03<0.0524,873
5619.332.19<0.05<119.712.18<0.055019.992.42<0.0522,994
8519.332.29<0.05<119.652.37<0.054819.192.20<0.0524,000
11919.712.51<0.05119.092.27<0.055019.402.52<0.0521,834
18019.502.69<0.05319.502.530.145019.502.770.0822,776
26019.502.62<0.05619.502.600.144319.502.57<0.0519,578
41619.502.070.08119.502.260.083419.502.26<0.0515,627
59619.502.340.20219.502.220.092719.502.561.1913,043
Table
Table 6. As release from jarosite.
Table 6. As release from jarosite.
Experiment typeWithout AsCo-precipitated AsAdsorbed As
HoursClSO4FeAsClSO4FeAsClSO4FeAs
(g/L)(g/L)(mg/L)(µg/L)(g/L)(g/L)(mg/L)(µg/L)(g/L)(g/L)(mg/L)(µg/L)
1019.632.68<0.05<119.801.84<0.05419.932.18<0.05261
1019.362.44<0.05<119.772.60<0.05619.642.31<0.05362
1119.342.19<0.05<119.662.70<0.05519.232.39<0.05441
1319.432.27<0.05<119.682.52<0.05719.422.46<0.05532
1719.692.40<0.05<119.452.46<0.05719.682.43<0.05744
2319.672.56<0.05<119.342.40<0.05719.792.48<0.051000
3619.352.39<0.05<119.242.61<0.05519.452.43<0.051037
5619.352.37<0.05519.292.54<0.05519.342.45<0.051607
8519.582.33<0.05<119.262.45<0.05519.212.52<0.051799
11919.682.72<0.05<119.362.55<0.05519.282.31<0.051654
18019.502.56<0.05219.502.80<0.05919.502.46<0.052101
26019.502.70<0.05<119.502.78<0.05719.502.740.492968
41619.502.17<0.05<119.502.38<0.05519.502.273.026452
59619.502.32<0.05119.502.55<0.05219.502.52<0.054758
Table
Table 7. pH during seawater saturation.
Table 7. pH during seawater saturation.
Ferrihydrite
Without AsHours0244872120168192336504648
pH7.26.47.78.07.77.77.37.97.37.4
Co-precipitated AsHours0244872120168192336504648
pH7.26.57.78.07.77.87.67.97.57.3
Adsorbed AsHours0244896144168312480648
pH7.27.78.07.87.97.77.97.57.5
Schwertmannite
Without AsHours0247296120192216240288336360504672
pH7.25.23.63.33.63.63.63.63.63.63.53.53.6
Co-precipitated AsHours0487296168192216264312336480648
pH7.24.63.63.93.73.93.83.83.83.73.73.5
Adsorbed AsHours0244896144168312480648
pH7.24.04.34.24.24.24.24.14.1
Jarosite
Without AsHours0247296120192216240288336360504672
pH7.27.96.98.08.06.87.67.97.87.67.45.45.9
Co-precipitated AsHours0487296168192216264312336480648
pH7.28.08.38.26.97.37.97.97.77.77.57.0
Adsorbed AsHours0244896144168312480648
pH7.27.68.08.07.97.87.77.67.5
Goethite
Without AsHours0247296120192216240288336360504672
pH7.26.85.78.28.17.26.88.07.77.77.77.87.3
Co-precipitated AsHours0487296168192216264312336480648
pH7.28.08.28.17.16.78.07.87.87.77.97.7
Adsorbed AsHours0244896144168312480648
pH7.27.38.18.08.07.98.27.87.7

3. Results and Discussion

3.1. Arsenic Release from Ferrihydrite

The highest rate of liberation from ferrihydrite with co-precipitated arsenic occurred within the first 20 h, with a maximum of 260 µg/L. This corresponds to approximately 0.2% of the total load capacity (TLC) for this synthesis. After that, a decrease in dissolved arsenic showed two main steps until reaching 76 µg/L after 25 days from the beginning of the experiment (Figure 4; Table 3). The pH strongly decreased during the first 24 h; however, this did not show any relationship to Fe concentrations, which always remained below detection limit. It is possible that, during arsenic co-precipitation, a small fraction of the metalloid could be adsorbed as surface complexes, in which case the initial release could be related to ion exchange rather than dissolution.
Ferrihydrite with adsorbed As instantly reacted with seawater, releasing about 1.4% TLC (1555 µg/L) within 10 h, followed by a reduction to 273 µg/L towards the end of the experiment (Figure 4). The pH remained stable around 7.5 and no Fe release was measured. This was expected at neutral pH, with concentrations of OH− and H+ being inadequate to promote dissolution [18]. The high load capacity for synthetic ferrihydrite (2.96 wt %) could be explained in terms of surface area, which for this mineral can reach up to 600 m2/g (Table S1 in supplementary information) and enables ferrihydrite to better react with the surrounding media. Under this scenario, stability of inner-sphere bidentate complexes [30] at the given conditions would be responsible for high arsenic retention. Nevertheless, Jain [63] determined that adsorbed As(V) in ferrihydrite can cause a reduction of the point of zero charge (PZC), by as much as 2.4 pH units.
Minerals 04 00603 g004 1024
Figure 4. Arsenic release from ferrihydrite. (A) Arsenic release from co-precipitated ferrihydrite and (B) arsenic release from adsorbed ferrihydrite.
Figure 4. Arsenic release from ferrihydrite. (A) Arsenic release from co-precipitated ferrihydrite and (B) arsenic release from adsorbed ferrihydrite.
Minerals 04 00603 g004
This could be what triggers limited desorption by ion exchange in this case, due to the difference in electric charge between the electrolyte and ferrihydrite PZC (Table S2). Ferrihydrite transformation to goethite should not affect the results of this study (~1% transformed after 500 h [64]).

3.2. Arsenic Release from Schwertmannite

Arsenic release from schwertmannite with co-precipitated As remained at values close to normal seawater throughout the experiment (~2 µg/L), with the exception of up to 9 µg/L dissolved arsenic after 23 h, after this peak the value reduced again to normal seawater values within 13 h (Figure 5). An increase in iron and sulphate concentrations (Fe > 0.4 mg/L and SO4 > 2.6 g/L) showed consistency with arsenic liberation, which also decreased during the next hours, but at a slower rate than arsenic. During the first 72 h, pH dropped from seawater (pH = 8.1) to acidic values (pH = 3.6) confirming the high potential of schwertmannite to buffer the pH. Arsenic release after 23 h could indicate that schwertmannite was able to dissolve or partially transform to a stable mineral while the circumneutral pH allowed it, however, due to the great charge difference between the environment and schwertmannite’s PZC, surface re-adsorption as well as co-precipitation processes were able to control ion release after it reached acid pH values.
Schwertmannite with adsorbed As in contact with seawater only released 7 µg/L of arsenic, within the first 36 h. The concentration then decreased to normal seawater values toward the end of the experiment, similarly to the previous case. At pH ~ 3.7, a slight increase in arsenic (7 µg/L), iron (0.96 mg/L), and sulphate (2.5 g/L) concentrations can be interpreted as a minor dissolution that ends when the system normalizes around pH 4. In order to understand the low TLC (~0.81 wt %) compared with values given in the literature (~10 wt %) [56,65], it should be considered that under the Regenspurg method a quick precipitation process could affect the morphology of the grains and subsequently the development of surface area [54].
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Figure 5. Arsenic release test for schwertmannite. (A) Arsenic release from co-precipitated schwertmannite and (B) arsenic release from adsorbed schwertmannite.
Figure 5. Arsenic release test for schwertmannite. (A) Arsenic release from co-precipitated schwertmannite and (B) arsenic release from adsorbed schwertmannite.
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3.3. Arsenic Release from Goethite

The synthesis with co-precipitated As started with a release of 55 µg/L, which decreased to 27 µg/L after 586 h. From hour 50 to 180, a sustained decrease in pH from 8.2 to 6.7 set the start of Fe release, with a maximum of 0.14 mg/L that declined as pH stabilized at 7.8 by the end of the experiment (Figure 6). After 180 h, despite its high stability range, goethite released a significant amount of Fe that could be understood in terms of dissolution. When synthetic goethite presents short aging time, it also can present a similar solubility product (Ksp) to ferrihydrite [66]. It is possible that given our synthesis conditions, a short aging time and a high Ksp, protonization over the goethite surface could be the responsible mechanism behind iron and arsenic release.
The synthesis with adsorbed As showed the greatest arsenic desorption by releasing 25.8 mg/L of arsenic at a rate of 1.43 × 10−5 mol/L/h during the first day in contact with seawater (Figure 6). The pH underwent multiple changes, going from normal seawater to pH 7.2 to 8.1 within the first 48 h. These trends also correlated with arsenic liberation. Iron continued below detection limit in almost every extraction and sulphate presented the lowest average levels for this mineral. PhreePlot computations for goethite surface complexes showed that above pH 8, soluble arsenic species predominate in a seawater environment, so natural arsenic release is expected at high pH. This occurs due to a lower stability in surface complexes as arsenic ions exchange with seawater species, mainly sulphate. At lower pH, re-adsorption would occur as (OAsO2OH−1.5) complexes that can result in a reduction of up to 50% of the initial release of dissolved arsenic by the end of the experiment.
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Figure 6. Arsenic release test for goethite. (A) Arsenic release from co-precipitated goethite and (B) arsenic release from adsorbed goethite.
Figure 6. Arsenic release test for goethite. (A) Arsenic release from co-precipitated goethite and (B) arsenic release from adsorbed goethite.
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3.4. Arsenic Release from Jarosite

Among the studied minerals, jarosite with co-precipitated arsenic showed the lowest levels of uptake in adsorbed and co-precipitated synthesis. Arsenic release during contact with seawater was less than 9 µg/L, which represents about 0.15% of the TLC. Iron remained below the detection limit even though the stability field for jarosite is restricted to acid environments. Sulphate kept stable during the experiment. Considering theories that could explain the limited arsenic release, it is possible that As co-precipitation caused a decrease in solubility [67] or release of sulphate caused precipitation of a Fe-OOH coating, protecting jarosite from the surrounding media [68].
Similarly to goethite, jarosite with adsorbed As showed high desorption by releasing up to 6452 µg/L, which corresponds to 45% TLC (Figure 7). The pH maintained around 7.6 but with variations within the first hours. After contact with seawater, the pH first decreased to 7.2 to thereafter increase to 8 at the same time that a large quantity of arsenic was released. A relationship between pH and desorption range can be observed. When the system reached pH 8, arsenic gradually increased in concentration, until pH dropped again. At this stage, the arsenic release began to occur at a higher rate. Given that there was no release of iron but sulphate remained at nominal values for seawater, it is possible that sulphate exchange could be taking place on the jarosite surface, promoting arsenic release. In this case, re-adsorption would not take place due to higher pH values [69].
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Figure 7. Arsenic release test for jarosite. (A) Arsenic release from co-precipitated jarosite and (B) arsenic release from adsorbed jarosite.
Figure 7. Arsenic release test for jarosite. (A) Arsenic release from co-precipitated jarosite and (B) arsenic release from adsorbed jarosite.
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4. Conclusions

During this laboratory study, some common Fe(III) hydroxides and oxyhydroxy-sulphates, including ferrihydrite, goethite, schwertmannite, and jarosite, present in the oxidation zone of mine tailings and in acid soils, showed efficacy as arsenic sorbents under several experimental conditions. Goethite and ferrihydrite were able to adsorb large loads of arsenic at their surfaces (~3 wt %) while schwertmannite and jarosite were able to incorporate less than 0.8 and 0.3 wt % of arsenic, respectively. Once in contact with seawater, each mineral showed different sorbent capacities, depending on the type of arsenic load and mineral stability.
Our results demonstrated that during seawater intrusion in coastal tailings, arsenic release can be attributed mainly to ion exchange and dissolution processes. Ferrihydrite and schwertmannite, two meta-stable minerals with large reactive surface areas, were less likely to release arsenic due to a higher ZPC and schwertmannite’s buffering capacity that acidified the pH of contact seawater. Moreover, highly stable goethite and jarosite showed the greatest release among minerals with adsorbed arsenic.
In general terms, synthesized minerals with co-precipitated arsenic were less inclined to liberation in comparison with syntheses with adsorbed arsenic, where ion exchange was a key parameter for liberation of the metalloid.
Schwertmannite and ferrihydrite presented the highest retention; however in the first case it mostly depended on the pH buffering capacity. It can be hypothesized that, under a real scale flooding scenario, schwertmannite would not be able to acidify the seawater in the same way it did in this experiment and that sea water geochemistry would dominate. In this case, a lower retention would be expected. In the medium term, transformation processes can release significant amounts of As that would not be fully retained by new sorbents. In the case of schwertmannite, exchange between sulphate and several species can affect stability and increase transformation rates to goethite [70]. The increase in dissolved iron and sulphate from schwertmannite towards the end of the experiment may indicate low stability or the beginning of transformation/dissolution processes. Ferrihydrite demonstrated its importance in coastal environments by representing a sink for arsenic at alkaline pH. However, the meta-stable nature for this mineral implies that in the long term dissolution and transformation processes can release arsenic regardless.
Goethite, another common stable mineral in the coastal environment, was not able to retain adsorbed arsenic after contact with seawater. This was due to the lower stability that surfaces complexes presented under seawater conditions. Although some authors have proposed bidentate complexes as the main linkage for arsenic in goethite’s surface, in this case the release of 20% TLC would be more in accordance with less stable monodentate complexes as proposed by Loring [35].
In coastal environments, stability of sorbent minerals should be considered as a whole, taking into account interaction between sorbents, fate and transport of toxic elements, surface complexation, interaction with seawater species, transformation into stable minerals, and dissolution processes.

Acknowledgments

We acknowledge the Staff of the laboratory of the Institute for Applied Economic Geology (GEA) University of Concepcion, Chile for their support during experiment and analysis. We also gratefully acknowledges to Pierre Rousseau who helped us by reviewing this work.

Author Contributions

Experimental design, data collection, and analysis was conducted by Jenny Gaviria under the supervision of Bernhard Dold. Rodrigo Alarcón was in charge of data interpretation, analysis and scientific writing under the supervision of Bernhard Dold.

Conflicts of Interest

The authors declare no conflict of interest.

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MDPI and ACS Style

Alarcón, R.; Gaviria, J.; Dold, B. Liberation of Adsorbed and Co-Precipitated Arsenic from Jarosite, Schwertmannite, Ferrihydrite, and Goethite in Seawater. Minerals 2014, 4, 603-620. https://doi.org/10.3390/min4030603

AMA Style

Alarcón R, Gaviria J, Dold B. Liberation of Adsorbed and Co-Precipitated Arsenic from Jarosite, Schwertmannite, Ferrihydrite, and Goethite in Seawater. Minerals. 2014; 4(3):603-620. https://doi.org/10.3390/min4030603

Chicago/Turabian Style

Alarcón, Rodrigo, Jenny Gaviria, and Bernhard Dold. 2014. "Liberation of Adsorbed and Co-Precipitated Arsenic from Jarosite, Schwertmannite, Ferrihydrite, and Goethite in Seawater" Minerals 4, no. 3: 603-620. https://doi.org/10.3390/min4030603

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