Article pubs.acs.org/IECR
Oxidation of Ferrous Sulfate Hydrolyzed SlurryKinetic Aspects and Impact on As(V) Removal Renaud Daenzer, Thomas Feldmann, and George P. Demopoulos* McGill University, 3610 University Street, Montreal, Québec H3A 0C5, Canada S Supporting Information *
ABSTRACT: In this paper the air sparging-oxidation kinetics of hydrolyzed iron(II) sulfate slurries in connection to acidic effluent treatment by neutralization for the removal of contaminants like arsenic(V) were investigated. It was shown in the absence of As(V) and forced air sparging that Fe(II) sulfate of initial concentration 75 mmol/L hydrolyzed completely between pH 7.5 and 8.5 at 22 °C. Subsequent oxidation by forced air sparging of the ferrous hydroxide slurry was found to proceed via a series of transformations starting from ferrous hydroxide to green rust, to lepidocrocite or magnetite depending on the pH and rate of oxidation, and finally to goethite. The oxidation kinetics at pH 8 or higher were governed by oxygen mass transfer while at pH 6 they were chemically controlled with autocatalytic behavior arising from the effect of in situ formed ferric oxyhydroxide. In the presence of As(V), both Fe(II) and As(V) precipitated from solution starting at pH 4 with the latter ultimately dropping below 1 mg/L past pH 6.5 via the apparent formation of ferrous arsenate compound. Subsequent oxidation by air sparging of the Fe(II)−As(V) slurry at constant pH 8 led to destabilization of the arsenate-carrying phase resulting in partial release of As(V). The bulk control of As(V) in the latter case appears to switch from ferrous arsenate to arsenate adsorption on in situ formed iron(III) oxyhydroxide.
1. INTRODUCTION Iron is perhaps the most commonly encountered impurity in hydrometallurgical process solutions that is typically removed by hydrolytic precipitation, and it has been the subject of substantial investigation.1−3 Of the two oxidation states of iron, it is ferric that undergoes readily hydrolytic precipitation.4,5 In the case of ferrous iron, oxidation to the ferric state to induce precipitation of the latter is the typical approach.6,7 Less common is the oxidation of a ferrous hydroxide slurry. The latter system is of interest in the case of natural or forced aeration occurring during acidic sulfate effluent treatment by lime neutralization for the purpose of contaminant removal as is arsenic coprecipitation with ferric sulfate salt addition to a molar ratio Fe(III)/As(V) of at least 3.8−12 In this case part of the iron may be present as ferrous iron;9,12−14 hence, its oxidation by air sparging may be practiced given the higher affinity of ferric iron for arsenate either via surface complexation on ferrihydrite15 or via formation of ferric arsenate.16,17 Indeed in a recently reported industrial process scheme, aeration was employed during coprecipitation of arsenic with ferrous sulfate. Oxidation by air sparging was conducted in two stages at pH ∼ 6 and ∼ 8.5 resulting in effective removal of arsenic below 1 mg/L.18 The available literature on hydrolyzed ferrous slurry oxidation and its impact on the removal of arsenic(V) by coprecipitation is rather limited, hence the undertaking of the present work. Following a recent study that looked into the behavior of arsenate during hydrolysis of ferrous sulfate,14 here the focus is shifted to the oxidation kinetics of ferrous hydroxide by air/oxygen sparging at various pH set points at 22 °C and its impact on arsenic(V) removal and retention. The oxidation of ferrous iron with oxygen strongly depends on the solution pH.19 At acidic pH the rate of oxidation of © 2015 American Chemical Society
Fe(II) with oxygen is very slow as for example shown by Sato et al., who observed negligible oxidation even at 50 °C with copper(II) sulfate present as catalyst.20 Its oxidation rate though increases significantly at near neutral pH.21 In the latter case the oxidation kinetics are commonly described by the following rate equation (eq 1):21,22 −d[Fe(II)]/dt = k(OH−)2 pO2 [Fe(II)]
(1)
It is obvious that eq 1 becomes a pseudo-first order relationship, depending only on the concentration of Fe(II), at constant pH and pO2. In the meantime excess SO42− anion concentration is reported to decrease the rate constant k,22,23 but ferric hydroxide was found to have a rate increasing autocatalytic effect, via an adsorption-mediated mechanism.21,22,24 This effect was studied in detail by Davison and Seed,21 who calculated the half-lives of ferrous ions in solution at pH 6.5, 7.0, 7.5, and 8.0 to be 361, 36, 3.6, and 0.36 min, respectively, at 20 °C. However, as the oxidation rate accelerates the oxidation process may shift from chemical reaction control to oxygen mass transfer control, a rather important consideration in the design of industrial effluent treatment reactors that has not been adequately addressed in previous studies. At higher pH (>9) the oxidation rate of ferrous iron was reported to slow down because of formation of magnetite.24 These results point to a complex dependency of the oxidation rate on pH and iron (oxy)hydroxide product. At the same time while the effect of several anions (other than SO42−) on the oxidation of ferrous iron (via complex Received: Revised: Accepted: Published: 1738
October 8, 2014 January 20, 2015 January 23, 2015 January 23, 2015 DOI: 10.1021/ie503976k Ind. Eng. Chem. Res. 2015, 54, 1738−1747
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Industrial & Engineering Chemistry Research formation) by oxygen has been studied23 no similar study for the effect of arsenate species has been reported. In addition to pH the rate of oxidation can impact the type of iron (II,III) (oxy)hydroxide phase that would form by reaction of air/ oxygen with ferrous species, and this can also have an effect on arsenic removal/retention.25 Thus, lepidocrocite (γ-FeOOH) is produced via rapid oxidation in neutral to slightly basic solutions, while the oxidation product is magnetite (Fe3O4) if oxidation is done at slow rate at room temperature.26 Furthermore, the formation of γ-FeOOH or Fe3O4 by oxidation of ferrous hydroxide is typically preceded by other intermediate iron precipitates such as green rust (GR).26 Green rust can vary in composition but essentially consists of mixed Fe(II)/Fe(III) hydroxide layers which alternate with negatively charged interlayers in which anions like carbonate or sulfate are located.27 The sulfated green rust that is relevant in the present work is symbolized as GRII and has the following composition: [FeII4FeIII2(OH)12SO4].26−29 A detailed account of the pH and oxidation−reduction potential dependent pathway by which ferrous hydroxide undergoes oxidation/transformation to GRII and finally FeOOH was provided by Ruby et al.28,29 It is known that arsenic present in solution adsorbs on ferrihydrite, this being more extensive in the case of arsenate, and in particular as pH increases, molar ratio of Fe(III)/As(V) increases and ferrihydrite forms via in situ hydrolysis (coprecipitation mode).10,15,30−32 However, as discussed above if a ferrous solution is neutralized under oxidizing conditions, oxidation products such as green rust, magnetite, or other oxyhydroxides like lepidocrocite may form that can interact with arsenic contributing to its retention. Thus, Jönsson and Sherman33 found As(V) to adsorb on different types of green rust as well as magnetite at pH < 8. But the system apparently is very dynamic as the relative amount of As(V) present during oxidation of ferrous can have an effect on the type of Fe(II,III) oxy-hydrolyzed phase formed as illustrated by Refait et al.’s work.34 These authors observed green rust to form at low arsenic(V) presence (Fe/As ratios of 12 and 24), but no green rust formed at higher arsenic level (Fe/As = 4.8). In the latter case the final product was a mixture of ferrihydrite and other poorly crystalline Fe(III)oxyhydroxides.34 Randall et al.25 conducted a detailed study on the mechanism of As(V) adsorption on GRII (sulfate green rust) at constant pH = 7. It was found that As(V) adsorbs (via surface complexation) onto GRII, during its formation and subsequent oxidation to γ-FeOOH. However, if both solids were present simultaneously As(V) would be preferably adsorbed on GRII and only on γ-FeOOH once all GRII had been oxidized. In contrast to these results it seems that in other cases (carbonated green rust) arsenate was reported to be incorporated rather than simply adsorbed.34 The preceding discussion illustrates that the rate of oxidation of ferrous as a function of pH in ambient temperature effluent treatment systems can lead to different Fe(II,III) oxidation products with variable capacity of arsenate retention via different mechanisms (adsorption, incorporation, compound formation). Further, the rate of oxidation of ferrous iron may be subjected to the limitations of gas−liquid mass transfer kinetics,35,36 and the impact this can have on iron oxyhydroxide formation and arsenic removal has not been studied. Previous research at McGill has focused on the effect of process parameters on the efficiency of the Fe(III)/As(V) coprecipitation process and the stability of the generated coprecipitates.8,10,37 In a recent paper, the role of ferrous iron during
coprecipitation of As(V) with mixed Fe(II,III) media was investigated and shown to contribute to arsenate removal via the formation of ferrous arsenate.14 It is the purpose of the present paper to continue the study of this system by focusing on the oxidation kinetics of ferrous iron in solution or as ferrous hydroxide slurry by oxygen/air sparging at various pH set points and its impact on the uptake of As(V) under conditions of acidic effluent treatment by neutralization.
2. EXPERIMENTAL SECTION 2.1. Chemicals and Procedure. All chemicals used were of reagent grade quality supplied by Sigma-Aldrich or Fisher Scientific. All solutions were prepared with deionized water. All tests were batch experiments conducted in a 2 L baffled reactor. These experiments involved the use of different types of impellers, gas (air/oxygen) flow rates, agitation speeds, and pH set points. Solutions were prepared by dissolving calculated amounts of ferrous sulfate in 1650 mL of DI water to yield an iron concentration of 74.74 mmol/L (4174 mg/L) and, in later experiments, the addition of dissolved As2O5 to yield an Fe/As molar ratio of 4 and 1.65 (experiments OE8As-1 and -2), respectively. Air or pure oxygen was supplied to the reactor via a gas flow controller (Matheson Model 3270). The gas outlet was directly underneath the impeller. The starting pH of the solutions was approximately 1.5 unless otherwise noted. Base either as 0.5 M Ca(OH)2 or 0.5 M NaOH (OE8-9, OE8As-1) and in one case 1 M NaOH (OE8As-2) was added automatically using a pH controller to maintain pH at the set point. In the case of experiments OE8As-1 and -2, NaOH was added at a constant flow rate up to pH 8 before being manually added to maintain that pH for a duration of 1 h as well as during oxidation at constant pH 8. Typically the pH of the solution was ramped up to pH 8 within 20 min before the onset of oxidation by gas sparging. After the targeted pH was achieved the gas flow controller was switched on such that oxidation could commence. During that time, automatic base addition continued to maintain the select pH set point. For pH drift oxidation experiments OEDrift-1 and -2, 10 min after addition of ferrous sulfate, a calculated amount of NaOH base ([Fe(II)]/[OH−] = 0.58) was added to solution at once thereby immediately increasing the pH to ∼8.5. The pH adjusted solution was left to equilibrate for an additional 10 min prior to the onset of air sparging; thereafter, no further pH adjustments were made. 2.2. Analysis and Characterization Methods. The solid samples produced by the addition of calcium hydroxide base (Ca(OH)2) to iron sulfate media formed gypsum (CaSO4· 2H2O) as a byproduct. In order to characterize the solids the gypsum was dissolved in pH ≈ 1.5 (adjusted with HCl or H2SO4) water after which the solids were rinsed, filtered, and dried at 40 °C in an oven. Slurry samples were drawn out of the agitated reactors and centrifuged and the supernatant was filtered using 0.02 μm Whatman Anatop or 0.2 μm MCE Fisherbrand syringe filters. The resulting filtrate solution was acidified using three drops of trace metal grade nitric acid for analysis of the solution composition. Arsenic analysis was carried out using a flow injection metal hydride spectrometer (PerkinElmer Analyst 100 AAS equipped with a FIAS-400 Hydride generator) for samples containing less than 0.5 mg/L As(V). Samples with As(V) concentration greater than 0.5 mg/L and total iron measurements were analyzed using a Thermo Jarrell Ash Trace Scan ICP-OES. 1739
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Table 1. Summary of Experiments Performed at Different Oxidation Conditions and the Empirical Linear Equation Giving Fe(II) Concentration Decrease with Time during Gas Sparging gas flow rate experiment OEpHox−x [mL/min]
impeller type 5BF/3BPa (agitation speed [rpm])
OE4-1 OE4-2 OE4-3 OE6-1
air, 500 oxygen, 500 ″ air, 300
5BF (1000−2400) ″ 5BF (750) ″
OE8-1 OE8-2 OE8-3 OE8-4 OE8-5 OE8-6 OE8-7 OE8-8 OE8-9 OEDrift-1 OEDrift-2 OE8As-1; Fe/As = 3.52 OE8As-2; Fe/As = 1.68
air, 500 ″ air, 300 oxygen, 300 none air, 300 ″ ″ ″ ″ ″ ″
5BF 5BF 5BF ″ ″ 5BF 5BF 3BP 5BF 5BF 5BF 5BF
″
5BF (500, 750)
a
(1000) (2000) (750)
(1000) (500) (1000) (750) (700−740) (721) (500, 750)
pH initial
oxidation pH set point/actual
base used for pH adjustment
Oxidation Experiments at pH = 4, 6 ∼3.7 4 0.5 M Ca(OH)2 ∼3.7 4/4.06 ″ 2.55 4/3.78 ″ 3.77 6/5.17 ″ Oxidation Experiments at pH = 8 3.70 8/7.67 ″ 3.70 8/7.76 ″ 1.58 8/8.18 ″ 1.68 8/6.59 ″ 1.32 8/7.95 ″ 1.32 8/8.05 ″ 1.47 8/8.19 ″ 1.50 8/8.13 ″ 1.58 8/7.77 0.5 M NaOH 8.60 drift pHf = 4.37 50 w/w NaOH 8.4 drift pHf = 4.48 ″ 1.51 8/7.88 0.56 M NaOH (3.39 mL/min) 1.48 8/7.94 1.00 M NaOH (1.88 mL/min)
linear equation y = [Fe(II)] in mg/L, x = time in minutes
oxidation efficiency [%]
n.a. n.a. y = −0.90x + 4285.2 y = −54.19x + 5523.6
n.a. n.a. n.a. n.a.
n.a. y = −242.48x + 762.7 y = −88.12x + 2952.9 y = −914.33x + 3354.9 y = −12.49x + 3202.6 y = −126.12x + 2484.1 y = −29.91x + 3100 y = −38.13x + 3242.2 y = −159.88x + 2740.5 y = −177.24x + 7319.6 y = −198.65x + 7930.8 y500 = −51.38x + 2989.6; y750 = −104.99x + 3487 y500 = −75.95x + 3507.1; y750 = −166.42x + 4354
n.a. n.a. 15.9 30.1 n.a. 20.6 5.1 7.0 23.9 30.8 33.4 18.2 29.0
5BF = five blade flat impeller; 3BP = three blade pitched impeller.
Figure 1. Oxidation of ferrous hydrolyzed slurry (test OE8-7) at pH = 8 with air sparging rate (Qg) of 300 mL/min; vertical lines denote the start (1) and end (2) of oxidation. (a) pH and ESHE profile. (b) Linear conversion plot representing the oxidation of Fe2+ in solution by gas sparging.
first neutralization of the acidic solution (initial pH ∼ 1.5) containing ferrous iron (4,174 mg/L) either in the absence or presence of arsenic(V) to the pH value followed by oxidation. In most tests the pH was kept constant during oxidation, but some tests were also run at drifting pH during oxidation. All experiments and their specific parameters investigated are summarized in Table 1 (additional details can be found in Table SI 1 in the Supporting Information). As a result of the adopted procedure, the oxidation involved in most cases a heterogeneous reaction with ferrous hydroxide or ferrous arsenate/hydroxide slurry. 3.1. Oxidation of Hydrolyzed Ferrous at Constant pH. Before the effects of the various process parameters on ferrous hydroxide/arsenate oxidation are presented, a typical oxidation experiment is described in detail. The test chosen for description (OE8-7) involved oxidation of ferrous hydroxide
To determine the Fe(II) concentration a modified potassium dichromate titration method was used, and details on this can be found in a previous publication.14 Solids characterization was carried out via X-ray diffraction using a Philips PW 1710 diffractometer equipped with a copper target (Cu Kα1 radiation, λ = 1.5406 Å).
3. RESULTS AND DISCUSSION As discussed in the introduction it is of interest to study the oxidation chemistry of ferrous iron in the context of an iron− arsenate coprecipitation process. To this end multiple oxidation experiments (OE) were conducted initially involving only hydrolyzed ferrous solutions. Subsequently, oxidation experiments containing variable molar fractions of As(V) were carried out to study the coprecipitation behavior of arsenate during oxidation of Fe(II) to Fe(III). The procedure adopted involved 1740
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Figure 2. Iron concentration and ESHE and pH evolution profiles for experiment OE6-1 conducted at pH = 6 with 300 mL/min air flow rate; (1) base addition, (2) start of air sparging and automatic pH control, (3) end of oxidation. (a) Fe2+ and total iron concentration evolution over time in solid and filtrate. (b) pH and ESHE profile.
gradual formation and transformation of ferrous hydroxide and green rust to ultimately some form of poorly crystalline iron(III) oxy-hydroxide, FeOOH,26 which can also be inferred from the ESHE variations during the experiment. Thus, in terms of ESHE it is interesting to note that precipitation of ferrous during the ramp-up period resulted in a dramatic drop of ESHE from ∼500 mV at pH ∼1.5 to −500 mV at pH ∼8. This was followed upon air sparging by an increase to a plateau at −200 mV until oxidation was complete thereafter increasing gradually to near 400 mV. These ESHE changes are consistent with the oxidation pathways determined by Ruby et al.28 When comparing the experimental pH and ESHE values with their pH−ESHE diagram it is seen that just before the oxidation by air sparging the conditions represent the equilibria between Fe2+/ GRII, Fe2+/Fe(OH)2, and Fe(OH)2/GRII, respectively. During oxidation between times 20−50 min it is seen that Fe(OH)2 oxidizes to form GRII, which coincides with a drift of the ESHE and pH values across the GRII equilibrium region. During the pH = 8 controlled oxidation stage the corresponding ESHE was ∼−170 mV, which corresponds to the equilibrium couple of GRII/γ-FeOOH. This suggests that during the air sparging oxidation period GRII and lepidocrocite are likely present with the former slowly oxidizing into the latter final product, which was in part confirmed by XRD analysis as shown later in section 3.2.2. The predominant character of mass transfer-limited oxidation kinetics for the test done at pH 8 (linear plot in Figure 1b) was further validated via an extensive investigation of the effects of parameters such as agitation speed, impeller type (radial > axial), and gas type (O2 > air) as described in the Supporting Information section of this article (see particular plots in Figure SI1a−c). The linear rate equations corresponding to different conditions are summarized in Table 1. The oxidation rate was strongly dependent on pH having as a result a shift upon decrease of pH from mass transfer control to chemical reaction control. At pH 6 a complex chemical reaction rate-limiting regime becomes evident as can be deduced from the sigmoidal curve (Figure 2a). It is suggested that in the case of pH 6 the initial slow oxidation step corresponds to oxidation of soluble ferrous ions to ferric hydroxide, which forms in situ as result of spontaneous hydrolysis. Upon significant formation of ferric hydroxide (or ferric oxyhydroxide) owing to the sorption capacity of the
at constant pH 8. As can be seen in Figure 1a, the automated controlled constant base (0.5 M Ca(OH)2) addition to the starting solution to raise the pH from 1.47 to 8 took 20 min. This ramp-up to pH 8 is equivalent to a hydrolysis test as described in a previous publication,14 which results in the bulk of ferrous iron to precipitate as ferrous hydroxide. Upon reaching pH 8, oxidation was started via constant flow rate air sparging. The overall reaction therefore can be written as in eq 2: 6Fe(OH)2 + 1.5O2 → 6FeOOH + 3H 2O
(2)
The average pH between the onset of air sparging and the end of oxidation (later determined by titration) for this test was 8.19. The start of oxidation is marked by a slight drop in pH and a simultaneous increase in ESHE (from ∼−100 mV to ∼+380 mV at the end of oxidation). The end of the total (involving both soluble and precipitated forms) Fe(II) oxidation is marked by an increase in ESHE and the stabilization of pH, signaling no further reaction. The corresponding oxidation curve for the conversion of ferrous to ferric upon air sparging at pH 8 is shown in Figure 1b. The ferrous concentration data was obtained by titrating slurry samples for total (soluble and precipitated) Fe(II) at different time intervals. The starting Fe(II) concentration was 4233 mg/L and dropped to approximately 3100 mg/L (∼14% drop) over the course of the 20 min ramp-up to pH 8. This loss in ferrous iron (typically in the range between 14 and 35%) represents the portion of Fe(II) that was oxidized by air infiltration into the reactor, which was closed with a lid but subjected to otherwise atmospheric conditions. A green, mixed valence iron precipitate having the appearance of sulfate green rust26,38 formed. Therefore, the bulk of the remaining ferrous just before air sparging commenced and can be thought to be in the form of precipitated ferrous hydroxide. The oxidation of the slurry (monitored by measuring the ferrous concentration) clearly followed a linear trend with time: [Fe(II)] = −29.91t + 3100 (R2 = 0.995). The linear plot suggests that the oxidation of ferrous is controlled by oxygen mass transfer at the gas/liquid interface. During the oxidation experiments, color changes were noticed that are consistent with those observed by Lin et al.26 going from green to dark green, blue, greenish-yellow, and finally orange-yellow. These color changes most likely reflect 1741
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Figure 3. Oxidation of ferrous sulfate hydrolyzed slurry at drifting pH (OEDrift-1 and -2) with [Fe(II)]/[OH−] = 0.583; (1) start of NaOH addition, (2) start of air sparging at 300 mL/min. (a) pH and ESHE profiles. (b) Fe2+ concentration profiles for solid and filtrate.
Figure 4. XRD patterns documenting the evolution of iron oxy-hydroxide phases as a function of oxidation time for experiment OEDrift-1 (M = Magnetite, G = Goethite, L = Lepidocrocite).
latter, the remaining soluble Fe2+ species ions adsorb on it resulting in an autocatalytic effect (second fast step) on the oxidation of ferrous as proposed by Tamura et al.24 and represented by eq 3. 2+ Fe2 + + Fe(OH)3 · H+ → Feads · Fe(OH)3 + H+
3.2.1. Oxidation at Initial Ferrous to Hydroxide Ratio of 0.583 at pH ∼ 8.5. The ratio [Fe(II)]/[OH−] = 0.583 was selected based on the investigations by Olowe and Génin39 and Lin et al.,26 who used a similar ratio to study the oxidation of ferrous hydroxide in slightly basic solutions (pH 8.2). Similar to the work of these authors, the pH and ESHE were monitored during the course of the oxidation of ferrous sulfate and correlated to changes in the composition of the precipitated phases. The pH and ESHE profiles for two such drift experiments are shown in Figure 3. Prior to initiation of oxidation (OEDrift-1 and 2) a calculated amount of NaOH ([Fe(II)]/[OH−] = 0.583) was added at once to the solution thereby immediately increasing the pH to 8.6 and 8.4, respectively. A few minutes (3−6 min) after the initiation of air sparging an apparent rise in pH (see arrow at around 25 min) occurred. Moreover, it is noticeable that throughout the oxidation period (from the onset of air sparging) there are other inflection points and pH plateaus, namely, at pH ∼ 7 and ∼ 4.5 (final). At the same time the ESHE is seen to drop to −550 mV upon rise of pH to 8.5 thereafter increasing in a stepwise form during oxidation and finally reaching the terminal value 400−450 mV, in other words
(3)
Finally, the oxidation of dissolved ferrous iron at pH 4 has been experimentally shown to occur very slowly. Thus, even when pure O2 was used instead of air at an agitation speed of 750 rpm, there was less than 5% of the initial [Fe(II)] oxidized after over 3.5 h of sparging. At this or lower pH the oxidation of ferrous iron requires elevated temperature and high oxygen partial pressures.6,35 3.2. Oxidation of Hydrolyzed Ferrous at Drifting pH. In addition to oxidation experiments carried out at constant pH, tests were conducted whereby the pH of the solution was allowed to drift during oxidation giving insight into the reaction mechanism. Two types of tests were run at a [Fe(II)]/[OH−] ratio of 0.583. The first one at initial pH ∼ 8.5 (OEDrift experiments) and the second one at initial pH 11.5 (HEFe experiments described in the Supporting Information). 1742
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Industrial & Engineering Chemistry Research exhibiting a similar behavior as observed with the constant pH oxidation tests. The associated precipitation−oxidation behavior of hydrolyzed ferrous slurry is illustrated in Figure 3b. As it can be seen the bulk of ferrous (>90%) had precipitated (see solid symbols)apparently as ferrous hydroxide prior to the onset of oxidationthereafter converting to ferric iron at a constant rate. This is consistent with the mass transfer control mechanism discussed above. The jump in pH (see arrow in Figure 3a) is attributed to the enrichment of the solution with OH− by the intermediate formation of GRII (Fe(II)4Fe(III)2(OH)12SO4) from the oxidation of Fe(OH)2 at 0.5 ≤ Fe(II)/OH ≤ 0.58 (eq 4) as discussed and reported by Olowe and Génin39
et al.40 that fast aerial oxidation of an intermediate green complex resulted in lepidocrocite and that slower oxidation by air resulted in magnetite. Given that Fe(OH)2 is more sensitive than GRII to air oxidation, it is reasonable to assume that samples at t = 20−23 min contained mostly Fe(OH)2 and therefore were reported as lepidocrocite, whereas samples at t = 26−35 min contained decreasing amounts of GRII, which oxidized as crystalline magnetite. 3.3. Oxidation of Hydrolyzed Ferrous in the Presence of Arsenic(V). In this section the oxidation at constant pH = 8 of hydrolyzed ferrous sulfate slurries containing arsenic(V) at two Fe(II)/As(V) molar ratios, namely, 4 and 1.65 (targeted), is investigated (Table 2). The initial pH of the solution was
6Fe(OH)2 + SO4 2 − + 0.5O2 + H 2O → GRII + 2OH−
Table 2. Pre-Oxidation Hydrolysis Conditions and Data for Experiments OE8As-1 and -2
(4)
Going back to the pH curve of Figure 3a, after the initial pH rise there is a pH decrease (going through an initial plateau at pH 7 and a final one at pH ∼ 4.5) due to acid generation arising apparently from the oxidation of the intermediate GRII phase into the final ferric oxyhydroxide phase (identified as γFeOOH as discussed in section 3.2.2) according to eq 5: Fe(II)4 Fe(III)2 (OH)12 SO4 + O2 → 6FeOOH + H 2SO4 + 2H 2O
exp.
Fe/As molar ratio
% Fe(II) oxidized during hydrolysis to pH = 8
time to reach pH = 8 [min]
Fe(II)in/[OH−] at pH = 8
OE8As-1 OE8As-2
3.52 1.68
27 20
98 147
0.588 0.379
approximately 1.5 in order to prevent any premature precipitation. The pH was increased by constant NaOH base addition up to pH = 8, during which precipitation occurred. The pH was subsequently manually maintained at that pH for 1 h after which the slurry solution was sparged with air at constant pH = 8. The concentration profiles of Fe(II)slurry, Fe(II)filtrate, Fetot, and As(V)filtrate over time and pH for the two Fe/As starting solutions (equal to 3.52 and 1.68; OE8As-1 and -2) are shown in Figure 5a,b. The ramping up of the oxidation experiments containing arsenic before air sparging followed the same procedure as described in a previous publication.14 3.3.1. Preoxidation Hydrolysis of the Fe(II)/As Solution. During preoxidation hydrolysis, induced by constant base addition, 20−27% of the ferrous iron was found to oxidize due to air infiltration (data in Table 2). The pH (and ESHE) evolution during solution titration (hydrolysis) was monitored and is shown in Figure 6 for the OE8As-1 test. The pH evolution curves provide some interesting features (inflection points summarized in Table 3 and indicated by arrows at ∼70 min and ∼95 min in Figure 6) that can be linked to the precipitation behavior of ferrous and arsenate depicted graphically in Figure 5a,b. As it can be seen from the plots in Figures 5 and 6, As(V) was slowly removed from solution at the start of base addition between pH = 1.5 and pH ≈ 3.5−4. During that pH range, the As(V) concentration dropped together with Fe(II) at an Fe/As molar ratio of 1.94 and 1.58 for OE8As-1 and -2 respectively. It should be emphasized that, up to the first pH inflection point, the arsenic that was removed from solution ranges between 7.5 and 30.5% of the total. Subsequently, near the inflection point noticeable in the pH and ORP profile (Figure 6), As(V) was removed from solution at a much faster rate. By calculating the molar ratio of Fe(II)/As(V) that precipitated between the first and second inflection points (Table 3) the values of 1.40 and 1.55 for OE8As-1 and -2, respectively, were obtained. As discussed in more detail in a previous publication,14 these molar ratios may represent the phase that is formed between the two inflection points, which is close to the stoichiometric Fe(II)/ As(V) ratio of 1.5 for ferrous arsenate (symplesite).41
(5)
The above transformation reaction may involve other intermediate steps as evident by the intermediate plateau of pH 7 and the detection of magnetite in an intermediate product as reported in the next section (Figure 4). 3.2.2. Characterization of Oxidation Products. The oxidation products for OEDrift-1 are shown in Figure 4 as a function of oxidation time. At the onset of oxidation (t = 20) it was expected to find a mixture of Fe(OH)2 and GRII (some GRII was suspected to have had already formed prior to air sparging due to the decrease (∼10−15% in total [Fe(II)] between t = 10−20 min). However, instead, lepidocrocite (γFeOOH) was found, which is suspected to be the result of the fast oxidation of the ferrous phase solids upon exposure to air after separation from the solution for XRD characterization, and hence is not considered a true representative. Between t = 26 and t = 35 min the formation of magnetite (Fe3O4) appears. During sampling, it was observed that the oxidation slurries responded to a permanent magnet confirming the presence of magnetite. As oxidation continues past t = 38 min, goethite is the only solid phase left in the slurry as shown by the final XRD pattern. The final product of oxidation for OEDrift-2 was also goethite (not shown). The final goethite products of pH drift oxidation experiments are consistent with the results by Olowe and Génin39 for an initial FeSO4/NaOH ratio of 0.58. In this characterization work by XRD no green rust could be detected despite the indirect evidence for its formation discussed earlier in connection to the pH and ESHE profile characteristics. Thus, according to the discussion about the different reaction stages above, the sample at t = 26 min is at the height of the transformation of Fe(OH)2 into GRII (eq 4), yet the XRD scan indicates the presence of crystalline magnetite. This discrepancy is attributed to an artifact of the problem of preserving the sample for XRD (failure to prevent its oxidation during characterization). This was also the case in previous investigations as for example commented by Misawa 1743
DOI: 10.1021/ie503976k Ind. Eng. Chem. Res. 2015, 54, 1738−1747
Article
Industrial & Engineering Chemistry Research
Figure 5. Iron and arsenic concentration evolution profiles for tests (a) OE8As-1 (Fe(II)/As(V) = 3.52 and (b) OE8As-2 (Fe(II)/As(V) = 1.68). (a) Fe(II)/As(V) = 3.52 solution; (1) start of hydrolysis by constant base addition (3.33 mL/min, 0.56 mol/L NaOH), (2) manual base addition maintaining pH = 8 for 1 h (0.56 mol/L NaOH), (3) start of oxidation by air sparging and manual base addition maintaining pH = 8 during oxidation (0.5 mol/L NaOH). (b) Fe(II)/As(V) = 1.68 solution; (1) start of hydrolysis by constant base addition (1.88 mL/min, 1 mol/L NaOH), (2) manual base addition maintaining pH = 8 for 1 h (1 mol/L NaOH), (3) start of oxidation by air sparging and manual base addition maintaining pH = 8 during oxidation (1 mol/L NaOH).
value of 0.65 mg/L at pH = 8. In the meantime the excess ferrous (not associated with As(V)) following the precipitation of ferrous arsenate can be seen to continue precipitating, this time as ferrous hydroxide, but at a slower rate than ferrous arsenate (compare slopes of filtrate [Fe(II)] vs time plots before and after point (2) in Figure 5a). 3.3.2. Constant pH 8 Posthydrolysis Oxidation. After the initial hydrolysis of the Fe(II)/As solution up to pH = 8, the solution was equilibrated at that pH for 1 h before the onset of air sparging. During that time the concentration of As(V) did not change for OE8As-1 and increased only slightly to 0.94 mg/L for OE8As-2. Figure 7 presents the change in arsenic concentration (on a logarithmic scale) during the combined hydrolysis−oxidation treatment of experiments OE8As-1 and -2. In contrast to Figure 5a,b, which did not clearly show any change in [As] during oxidation due to scale, Figure 7 reveals that oxidation of the ferrous arsenate/ferrous hydroxide slurry at pH 8 leads to arsenic release. For OE8As-1 (Fe(II)/As(V) = 3.5), a rise in [As] is not noticed until 30 min after air sparging is initiated. Within this period all remaining Fe(II) in solution (346 mg/L in OE8As-1 and 871 mg/L in OE8As-2) precipitated out via its oxidation to Fe(III) and the in situ hydrolysis of the latter to ferric oxyhydroxide. At the same time the bulk of ferrous iron in the precipitated solids was oxidized as well. There appears however that a fraction of ferrous resists further oxidation. In the case of the OE8As-1 test (Figure 5a) this fraction corresponds to Fe(II) concentration equal to 391 mg/L at t = 210 min or ∼9% of the total iron. It is assumed that the remaining fraction of ferrous iron that is in the form of ferrous arsenate (symplesite) as ferrous hydroxide judging from the single-iron tests (see Figure 1b) is fully oxidized under the same conditions (pH 8). The amount of arsenic associated with this residual (oxidation-resisting) ferrous arsenate represents 25% of the original amount of arsenic. (In the case of OE8As-2 test (Figure 5b) this amount was estimated at ∼28% of the total arsenic). The remaining arsenic in the solids is most likely adsorbed on the in situ formed iron(III) oxyhydroxide as further discussed next.
Figure 6. pH and ESHE evolution during hydrolysis and oxidation of ferrous−arsenate containing solution (OE8As-1 test); (1) start of hydrolysis by constant base addition at 3.33 mL/min of 0.56 mol/L NaOH, (2) manual base addition for maintaining pH = 8 with 0.56 mol/L NaOH, (3) start of oxidation by air sparging and manual base addition of 0.5 mol/L NaOH.
Table 3. pH Inflection Points for Pre-Oxidation Hydrolysis of Ferrous−Arsenic(V) Containing Solutions Fe/As = 4
Fe/As = 1.65
experiment
HEAs-1
OE8As-1
HEAs-2
OE8As-2
pH of 1st inflection point pH of 2nd inflection point
4.05 7.08
4.05 7.38
3.41 7.64
3.64 8a
a The second inflection point is assumed to have occurred exactly where the hydrolysis stopped.
The precipitation of ferrous arsenate in this context has led to effective removal of arsenic(V). Thus, in the case of Fe(II)/ As(V) molar ratio of 3.52, the residual arsenic concentration past pH = 7.7 was found to be below the detection limit of ICPOES instrument (
