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Journal of Hazardous Materials 264 (2014) 490– 497

Contents lists available at ScienceDirect

Journal of Hazardous Materials

j o ur nal homep age: www.elsev ier .com/ locate / jhazmat

ynthesis, characterization and stability of Cr(III) and Fe(III)ydroxides

. Papassiopi, K. Vaxevanidou, C. Christou, E. Karagianni, G.S.E. Antipas ∗

chool of Mining Engineering and Metallurgy, National Technical University of Athens, Zografou Campus, Athens 15780, Greece

i g h l i g h t s

Fe(III)–Cr(III) hydroxides enhance groundwater quality better than pure Cr(III) compounds.Crystalline Cr(OH)3·3H2O was unstable, with a solubility higher than 50 �g/l.Amorphous Cr(OH)3(am) was stable with a solubility lower than 50 �g/l in the range 5.7 < pH < 11.For mixed Fe0.75Cr0.25(OH)3, the stability region was extended to 4.8 < pH < 13.5.

r t i c l e i n f o

rticle history:eceived 24 July 2013eceived in revised form1 September 2013ccepted 25 September 2013vailable online 1 October 2013

eywords:r remediatione(III)–Cr(III) mixed hydroxidesixed hydroxide stability

a b s t r a c t

Chromium is a common contaminant of soils and aquifers and constitutes a major environmental prob-lem. In nature, chromium usually exists in the form of two oxidation states, trivalent, Cr(III), whichis relatively innocuous for biota and for the aquatic environment, and hexavalent, Cr(VI) which is toxic,carcinogenic and very soluble. Accordingly, the majority of wastewater and groundwater treatment tech-nologies, include a stage where Cr(VI) is reduced to Cr(III), in order to remove chromium from the aqueousphase and bind the element in the form of environmentally stable solid compounds. In the absence of ironthe final product is typically of the form Cr(OH)3·xH2O whereas in the presence of iron the precipitate isa mixed Fe(1−x)Crx(OH)3 phase. In this study, we report on the synthesis, characterisation and stability ofmixed (Fex,Cr1−x)(OH)3 hydroxides as compared to the stability of Cr(OH)3. We established that the plainCr(III) hydroxide, abiding to the approximate molecular formula Cr(OH)3·3H2O, was crystalline, highlysoluble, i.e. unstable, with a tendency to transform into the stable amorphous hydroxide Cr(OH)3(am)phase. Mixed Fe0.75Cr0.25(OH)3 hydroxides were found to be of the ferrihydrite structure, Fe(OH)3, andwe correlated their solubility to that of a solid solution formed by plain ferrihydrite and the amorphous

Cr(III) hydroxide. Both our experimental results and thermodynamic calculations indicated that mixedFe(III)–Cr(III) hydroxides are more effective enhancers of groundwater quality, in comparison to theplain amorphous or crystalline Cr(III) hydroxides, the latter found to have a solubility typically higherthan 50 �g/l (maximum EU permitted Cr level in drinking water), while the amorphous Cr(OH)3(am)phase was within the drinking water threshold in the range 5.7 < pH < 11. In comparison, the mixedFe0.75Cr0.25(OH)3 hydroxides studied were of extended stability in the 4.8 < pH < 13.5 range.

. Introduction

Chromium is one of the most common pollutants in soils andquifers and represents a major concern for the environmentaluthorities worldwide. In natural systems Cr exists in two oxidationtates, Cr(VI) and Cr(III), with very different chemical and toxico-

ogical characteristics. Hexavalent Cr typically exists as oxyanion,rO4

−2, a form which is highly soluble and mobile in the envi-onment. On the contrary, trivalent Cr is a cation characterised by

∗ Corresponding author. Tel.: +30 210 7722037.E-mail address: gantipas@metal.ntua.gr (G.S.E. Antipas).

304-3894/$ – see front matter © 2013 Elsevier B.V. All rights reserved.ttp://dx.doi.org/10.1016/j.jhazmat.2013.09.058

© 2013 Elsevier B.V. All rights reserved.

limited mobility in the environmental media, due to its low hydrox-ide solubility in neutral and alkaline pH and its tendency to formstrong complexes with common soil minerals [1–3]. Toxicologi-cally, Cr(III) is considered relatively innocuous, whereas Cr(VI) isclassified as a confirmed human carcinogen [4]. As a consequence,reduction of Cr(VI) to Cr(III) represents an important remediationstrategy.

Reduction of Cr(VI) can be achieved by conventional chem-ical methods, using agents such as Na2S2O4 [5], CaSx [6], Fe0

[7,8] or FeSO4·7H2O [9–11]. Bioremediation is another alterna-tive, driven by the discovery of a wide variety of microorganismspossessing chromate reductive activity. Several sulfate and ironreducing bacteria have been found to reduce Cr(VI) under anaerobic

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N. Papassiopi et al. / Journal of Ha

onditions, either directly, by using Cr(VI) as a final electroncceptor [12–14], or indirectly, through biologically producedeductants, such as S2− or Fe2+ [15–18]. A crucial issue for theuccess of remediation is the environmental stability of the finalolid products. When reduction occurs in the absence of iron, e.g.s is the case with sulfur based reductants, the resulting com-ound is the Cr(III) hydroxide, Cr(OH)3. When the Cr(VI) reducinggent is an iron based compound, e.g. Fe(II), Fe0 or FeS2, the solidroduct is often a mixed Fe(III)–Cr(III) hydroxide, with a generalolecular formula of the type Fe(1−x)Crx(OH)3. Thermodynamic

tability studies of Cr(OH)3 over a wide pH range [1–3,19] indi-ate that Cr(OH)3 has good stability at neutral and alkaline pH,ut that there is also an abrupt increase of solubility at slightlycidic pH, typically below 5. Also, the solubility of Fe(1−x)Crx(OH)3ixed hydroxides was found to be lower to that of pure Cr(OH)3,

epending on the molar ratio between Fe and Cr [20]; a recent studystablished that such a mixed hydroxide with x = 0.25 had a solu-ility one order of magnitude lower compared to simple Cr(OH)3ith stability extended towards the acidic pH region [21], e.g. atH 4.5 the mixed hydroxides had a solubility lower than 50 �g/lEU drinking water limit), whereas pure Cr(OH)3 had a solubilityf the order of 500 �g/l. The two forms of Cr(III) highlighted inelevant literature are (a) a crystalline hydrous phase of the formr(OH)3·3H2O [22–24] and (b) an amorphous phase, often referredo as Cr(OH)3, although a more appropriate formula unit would ber(OH)3·xH2O, where x is lower than 3 [3,19,25–27]. In a sufficientlylkaline environment, Cr(III) salt solutions lead to precipitation ofydrogen-bonded Cr(OH)3(OH2)3 crystalline networks; the latterre metastable, eventually turning into an amorphous OH-bridgedetwork via dehydration (heat treatment above 50 ◦C accelerateshis transition) [24]. Trivalent Cr is normally found in the form ofr(OH)3 in neutral or slightly alkaline environments; it is also ofelatively low solubility (ksp = 6.7 × 10−3) and it is generally consid-red amorphous [8,18,21,28,29]. In the case of underground watereservoirs, the presence of iron may lead to precipitation of mixedxides of the sort Cr(III)–Fe(III) by a simultaneous entrapment ofr(VI) ions via reduction by Fe(II) cations [20,30].

In the current study, we discuss the synthesis, characterisationnd stability (by means of solubility) of mixed (Fex,Cr1−x)(OH)2ydroxides as compared to the stability of pure Cr(OH)3. Morepecifically, we investigate the structure of the plain Cr(OH)3 andixed chromium hydroxides and proceed to identify how struc-

ural characteristics determine their environmental stability over wide pH range. Our experimental work has involved synthesisf four solids: (1) the Fe hydroxide, Fe(OH)3, (2) the Cr hydroxider(OH)3, (3) a mixed Fe–Cr hydroxide, Fe0.75Cr0.25(OH)3 and (4) aixed Fe–Cr hydroxide, of the same stoichiometry as (3) whichas produced via redox mixing of Fe(II) and Cr(VI) solutions. We

onclude by qualifying our experimental stability results by com-arison against theoretically calculated solubilities. It is noted thatixed Fe–Cr hydroxides corresponding to the molecular formula

e0.75Cr0.25(OH)3 are produced when reduction of Cr(VI) is carriedut with ferrous iron compounds [9–11], due to the stoichiometryf the specific redox reaction, where three ferrous ions are requiredor the reduction of one chromate to the trivalent state.

. Materials and methods

.1. Synthesis routes

The compounds synthesized were Fe(OH)3, Cr(OH)3 and two

ixed Fe(III)–Cr(III) hydroxides with the same molar ratio

e(III)/Cr(III) = 3/1 and different synthesis procedures. The simpleydroxides and one of the mixed compounds were produced viaydrolysis reactions; the second mixed hydroxide was produced

s Materials 264 (2014) 490– 497 491

through a redox reaction between divalent iron, Fe(II), and hexa-valent chromium, Cr(VI). Following, we outline the synthesis routesfor each of the systems discussed.

For the production of Fe(OH)3, 40.4 g of Fe(NO3)3·9H2O weredissolved in 500 ml of deionized water followed by the addition of330 ml KOH 1 M, the last 20 ml of which were added drop wise whilesolution alkalinity was constantly monitored in order to achieve apH value within a range of 7 to 8. The solution was then vigor-ously stirred to avoid possible regional pH concentrations higherthan 8, known to lead in the formation of goethite [31]. The nextstage involved solution filtration and separation of the initial solu-tion ions from the solid by dialysis (passive transport through semipermeable membranes) [32]. The solid was then freeze-dried andthe resulting powder was sealed in a N2-atmosphere container.

Synthesis of the Cr(OH)3 compound involved dissolution of 40 gCr(NO3)3·9H2O in 500 ml of deionized water followed by the sameprocedure used in the synthesis of Fe(OH)3.

The mixed hydroxide 3Fe(OH)3·Cr(OH)3 was synthesized viatwo different routes. In the first, 30.3 g of Fe(NO3)3·9H2O were dis-solved in 500 ml of water followed by the sequential addition of10 g of Cr(NO3)3·9H2O and 330 ml of KOH. Secondarily, for the pro-duction of 3Fe(OH)3·Cr(OH)3 via a redox route, 750 ml of a FeSO4solution (0.2 M) and 200 ml of K2Cr2O7 solution (0.125 M) weremixed in a spherical reactor under a N2 flux and continuous stirring.The 3Fe(OH)3·Cr(OH)3 compound precipitated immediately with asynchronous decrease of the pH to a value of approximately 2. Inthe next stage, NaOH 5 M was added gradually to the solution, upto a pH of 7 to 8. The suspension was then continuously stirredfor 24 h after which it was filtered and the solid was subjected todialysis and freeze-drying.

2.2. Characterization

Chemical analysis of the compounds was carried out by HCldigestion (0.2 g of the compound in 20 ml 6 N acid) while deter-mination of the constituent elements in the solution was carriedout via a Perkin Elmer 2100 atomic absorption spectrophotometer(AAS). X-ray Diffraction (XRD) Cu K� spectra were recorded on aBruker D8 Focus diffractometer. Differential thermal analysis (DTA)and thermogravimetry (TG) were performed on a Setaram thermalanalyser in an Ar atmosphere and at a temperature increase rate of10 ◦C/min. Fourier-transformed infrared spectrometry (FTIR) mea-surements were recorded on a Perkin Elmer spectrometer usingKBr pellets with a 0.1 wt% content in solid sample.

2.3. Solubility tests

Stability of the synthesized hydroxides was tested within apH range of 2.5 to 6.5. The solubility tests were carried out in50 ml centrifugal tubes by mixing 1 g of solid with 30 ml of 0.1 MNaNO3 solution and adjusting the pH to the desired initial valuewith HNO3 or NaOH solutions. Prior to the solubility tests, thesolids were subjected to a washing procedure for the removal ofpotentially residual adsorbed ions, notable Cr(VI) or Cr(III). The pro-cedure involved addition of 30 ml of the background solution (0.1 MNaNO3) to the tube containing the 1 g soil, agitation for approx-imately 2 min, centrifugation of the suspension, and removal ofthe supernatant solution. These steps were repeated three times.Finally, a fourth dose of 30 ml background solution was added toeach tube and the pH was adjusted with HNO3 or NaOH, to achievediscrete pH values within a range of 2.5 to 6.5. The suspensions

were then kept under ambient temperature conditions (22 ± 1 ◦C)and constantly agitated for 24 h or a week. After this period, theresulting pH was recorded and the supernatant liquor was sepa-rated by centrifugation, filtered through a 0.45 �m membrane and

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92 N. Papassiopi et al. / Journal of Ha

nalyzed by atomic absorption spectroscopy (AAS), flame emissionr furnace graphite, for the determination of final Cr and Fe content.

. Results and discussion

.1. Compound synthesis and characterization

Chemical analysis results for the pure and mixed Cr(III)–Fe(III)

ompounds are presented in Table 2. The weight loss, �G, corre-ponding to adsorbed and structural water molecules is at a highestf 50.1% in the case of the pure Cr hydroxide, ranging between 25.3nd 32.0 wt% for the rest of the compounds. Fe content in the pure

ig. 1. Fe(III) hydroxide spectra: (a) XRD, (b) TG–DTA (vertical axis notation: for TG:weight loss, for DTA: %weight loss/min), (c) FTIR (vertical axis is % transmittance).otation is followed throughout.

s Materials 264 (2014) 490– 497

hydroxide is 52.2 wt%, while for the mixed hydroxides it varies from31.9 to 34.7 wt%. Respective Cr content is 36.1 and 9.9–10.8 wt%.Following, we discuss each of the synthesized hydroxides sepa-rately.

3.1.1. XRD, DTA–TG and FTIR analyses of the pure Fe(III) andCr(III) compounds

The XRD spectrum of the pure Fe(III) oxide is shown in Fig. 1a and

is typical of 2-line ferrihydrite amorphicity or nano-crystallinity.From the DTA–TG diagram of the pure Fe(III) hydroxide, shownin Fig. 1b, it is inferred that the weight loss owing to crystalline

Fig. 2. Cr(III) hydroxide spectra: (a) XRD, (b) TG–DTA, (c) FTIR.

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N. Papassiopi et al. / Journal of Ha

ater is 25.33%. Ferrihydrites synthesized in this manner have aormula unit of the type Fe2O3·xH2O; based on our own obser-ations of weight loss during thermal analysis the factor x wasefined as x = 3.02. As mentioned earlier, the Fe content was

ig. 3. Spectra of the two mixed Fe(III)–Cr(III) hydroxides: (a) XRD, (b) DTA–TG of the hyd) DTA–TG of the redox-synthesised hydroxide(e) FTIR of the redox-synthesised hydrox

s Materials 264 (2014) 490– 497 493

established to be 52.2 wt%, which in turn yields a value for the fac-tor x equal to 3.14, in agreement with our thermal analysis-derivedvalue; we, thus, conclude that this particular ferrihydrite may beadequately described by the formula unit Fe(OH)3. The compound’s

drolysis-synthesised hydroxide, (c) FTIR of the hydrolysis-synthesised hydroxide,ide.

494 N. Papassiopi et al. / Journal of Hazardous Materials 264 (2014) 490– 497

Table 1Chemical analyses and specific gravity measurements: (a) Total weight loss from heating at 25 to 800 ◦C during DTA–TG analysis, (b) based on AAS analysis results, (c)empirical molecular type derived from chemical analyses, (d) determined with the use of a water displacement pycnometer.

Solid �G, (%) (800 ◦C) Fe (%) Cr (%) Molecular type � (g/cm3)

(a) (b) (b) (c) (d)

Fe 25.3 52.2 Fe2O3·3.0 H2O 2.772311

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FeCr (hydrolysis) 27.8 36.7FeCrA (redox) 33.5 37.0

TA diagram (Fig. 2) also reveals a pronounced endothermic peakt approximately 130 ◦C assigned to the dissociation of crystallineater. The latter is followed by an exothermic peak at 380 ◦C mostrobably due to recrystallisation. No weight change was observedbove 400 ◦C. The system’s FTIR spectrum, depicted in Fig. 3, showswo characteristic peaks at 3500 and 1550–1600 cm−1 correspond-ng to stretching and bending vibrations in the H–O bond of freeurface or bulk (structural) OH groups [33]. The sharp peak at350 cm−1was assigned to N–O stretching, owing to the presence ofesidual NO3

− traces, despite the prolonged dialysis against water,n effect which is also reported in [31]. The broad peak in the band00–600 cm−1 consists actually of two peaks, approximately at 650nd 450 cm−1, which are due to deformation vibrations of the bulkH group, typical in the IR spectrum of ferrihydrite [28]. Absorptiont the 1000–1100 cm−1 range can be attributed to hydroxo bridgesonds of the type Fe–(OH)2–Fe according to Kopylovich et al. [34].

Pure Cr(III) hydroxide was produced using an experimental pro-edure similar to that followed for ferrihydrite. The system’s XRDpectrum is shown in Fig. 2a; the material appears to be quasirystalline, much in contrast to crystallographic studies of Cr(III)ydroxides synthesized under comparable conditions, yieldingmorphous systems [3,19,21]. In fact, the spectrum in Fig. 2a resem-les that of hydrated Cr(III) hydroxide of the form Cr(OH)3·3H2O22]; our own TG and chemical analysis data also support this for-

ula unit. The system’s DTA–TG diagram is presented in Fig. 2b. Aseen from the TG curve, heating of the material between room tem-erature and 550 ◦C, results in a total weight loss of approximately0%. On the assumption that the thermal decomposition reaction is

Cr(OH)3·xH2O → Cr2O3 + (2x + 3)H2O (1)

he factor x was estimated as 2.7, in agreement with [34] citing value of 2.5. An exothermic peak in the region 420–500 ◦Cbserved in our own DTA diagram (Fig. 2b), was also found in theTA spectra in [23,26]. However, the hydroxide system in [26] washaracterized as amorphous while the one in [23] was reportedo be crystalline. In the compound’s FTIR spectrum, shown inig. 2c, the absorption in the wave number ranges of 3500 and550–1600 cm−1 was assigned to the stretching and bendingibrations of adsorbed and structural water molecules. Addition-lly, the peaks at 450 and 1050 cm−1 were assigned to Cr–OHibrations, also supported by findings of Amonette and Rai [35].

.1.2. XRD, DTA–TG and FTIR analyses of the mixed Fe(III) andr(III) compounds

The XRD spectra of the mixed hydroxides, presented in Fig. 3a,ppear largely similar; the spectrum of hydrolysis hydroxide, FeCr,ears a strong resemblance to that of Fe(III) 2-line ferrihydriteindicative broad peaks at 35◦ and 63◦ 2� for Cu K�, Fig. 1). Theedox hydroxide, FeCrA, appears to be overall amorphous, whilee note that peaks related to crystalline Cr(OH)3·3H2O were con-

lusively absent in both spectra(see Fig. 2a for comparison). Hence

e deduce that the structure of the mixed hydroxides could be

dequately described by that of ferrihydrite and it is largely unaf-ected by the presence of Cr, the increased participation of whichromotes amorphicity [36].

6.1 Cr2O3·8.5 H2O 2.1851.0 (Fe2O3)(0.76) ·(Cr2O3)0.24·3.4 H2O 2.69760.5 (Fe2O3)(0.77)·(Cr2O3)0.23·3.7 H2O 2.8018

DTA–TG profiles for both hydroxides are shown in Fig. 3b andd. In the case of the hydrolysis-synthesized FeCr system, a double-stepped weight reduction of 16.5% and 11.3% was observed (Fig. 3b),corresponding to the removal of 2.0 and 1.4 mol of water per molof mixed M2O3 (also, see Table 1). The two endothermic peaks at160 and 210 ◦C coincide with the two-step loss of water, whilethe exothermic peak at 500 ◦C can be attributed to alterations inthe crystal structure of mixed oxides after dehydration. A simi-lar exothermic peak at 500 ◦C was observed by Kopylovich et al.[34] for a mixed hydroxide, containing Fe/Cr at 1/l molar ratio, andwas attributed to the formation of a crystalline solid solution ofhexagonal Fe2O3 and Cr2O3, which was confirmed by the X-raydiffractogram of the thermolysis product. The FeCr thermal profileresembles also those of mixed Fe(III)–Cr(III) hydroxides (at a Fe/Crratio of 7:3) studied in [37] that were reported to behave as solidsolutions of crystalline �-Fe2O3 and �-Cr2O3. In contrast, the redox-synthesized FeCrA system (Fig. 2d) did not exhibit any exothermicactivity in the region 400–500 ◦C. The DTA–TG profile revealed twoendothermic peaks 160 ��� 660 ◦C, corresponding to 25.5% ���7.5% weight loss steps, respectively. The latter can be attributedto desorption of SO3. A similar profile is observed during the ther-mal analysis of schwertmannite, which is a sulfate bearing Fe(III)hydroxide with the approximate formula Fe16O16(OH)(16−2x)(SO4)x

[33]. The presence of sulfates, which is due to the use of FeSO4salt for the production of mixed hydroxides via the redox route, isalso evident in the FTIR spectrum presented in Fig. 3e. The complexpeaks around wave number 1100 cm−1 are characteristic of sulfateanions’ stretching vibrations of [33]. Other features of interest werethe H–OH vibrations at 1600 and 3500 cm−1, Cr–OH stretching at450 and 1000 cm−1, and the Fe–OH peaks at 500 and 600 cm−1.Similarly, the principal FTIR features of the FeCr system, shownin Fig. 3c, are the H–OH peaks at 1550–1600 and 3500 cm−1, N–Ostretching vibrations at 1350 cm−1, Cr–OH at 450 and 1050 cm−1

and Fe–OH at 600 and 500 cm−1.

3.2. Solubility tests

3.2.1. Cr(III) compoundThe solubility results of the crystalline Cr(III) hydroxide are pre-

sented in Fig. 4. Two sets of data corresponding to the solubilityof amorphous Cr(III) hydroxides are also shown in Fig. 4. The firstset was published by Rai et al. [19] and represents the solubilityof Cr(III) hydroxides, which were left in contact with the aqueoussolutions for a long equilibration period of 502–507 days. The sec-ond set corresponds to Cr(III) hydroxides, which were subjected to athermal treatment immediately after their precipitation by hydrol-ysis, consisting of keeping the hydrolysis pulp at 70 ◦C for 60 h [21].It is noted that the hydrolysis procedure followed in the currentstudy is similar to that applied by Papassiopi et al. [21], but specificattention was paid to avoid any exposition of the resulting precip-itate to high temperature conditions; for this reason the removal

of retained moisture was carried out by freeze drying. This modifi-cation in the experimental procedure resulted in the production ofsolids with different chemical and crystallographic characteristics,i.e. the solid produced in [21] was amorphous and corresponded

N. Papassiopi et al. / Journal of Hazardous Materials 264 (2014) 490– 497 495

-7.0

-6.0

-5.0

-4.0

-3.0

-2.0

-1.0

0.0

2 3 4 5 6 7

log[

Cr],

M

pH

Amorphous (R ai et al. , 20 04)

Amorphous (Papassiopi et al., 2012)

Crystall ine, 7 d ( This stud y)

Crystalline, 24 h (This study)

Calculated with logKsp=4.09

Calculated with logKsp=7.28

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ig. 4. Aqueous chromium concentrations at different pH values as measured in suuration, 24 h and 7 days, and (ii) aged amorphous Cr(III) hydroxides, Cr(OH)3(amMinteq software and thermodynamic data presented in Table 2.

o the formula Cr(OH)3, whereas the hydroxide of current study isrystalline with 3 additional water molecules in the crystal struc-ure, Cr(OH)3·3H2O.

The solubility data presented in Fig. 4 suggest that the crystallineydroxide (shown as curve (2)) is much more soluble within thecidic region (pH 2.5 to 5.0), compared to its amorphous counter-arts (shown as curve (1) in Fig. 4). The high stability (low solubility)f the amorphous phase has been reported in a number of previ-us studies [3,19], however solubility data for the crystalline Cr(III)ydroxide are not available in the relevant scientific literature.r(OH)3·3H2O is a metastable solid phase in the Cr(III)–H2O sys-em which is gradually transformed to the more stable Cr(OH)3morphous phase.

We note here that curve (1) in Fig. 4 was calculated via the spe-iation software VMinteq [38], based on equilibrium constants forhe hydrolysis of Cr(III) ions and the precipitation of amorphousr(OH)3 as estimated in [19] which, in turn, incorporated experi-ental data from work in [1,19]; the latter study considered the

r(III) hydroxide solubility within a wide pH range (pH = 2.8–14)s well as extended equilibration times, up to 698 days, in ordero achieve steady state conditions in the equilibrium between ther(III) hydroxide and the solvent. The Cr(III) hydrolysis and precip-

tation reactions and the corresponding constants are presented inable 2. Curve (2) was also calculated considering the same Cr(III)ydrolysis reactions for the aqueous species (reactions (1)–(5) inable 2) and changing the value of the solubility product describ-ng the “pseudo equilibrium” of the crystalline solid phase (seeeaction (6a)). The trend of experimental solubility data suggests

hat crystalline compound’s Ksp is at least 3 orders of magnitudeigher compared to that of amorphous, namely log Ksp > 7.28 forhe crystalline compound against log Ksp = 4.09 for the amorphousne.

able 2r(III) hydrolysis and precipitation reactions and corresponding reaction constantsI = 0 M) used for thermodynamic calculations with VMinteq.

Reaction log K

1 Cr3+ + H2O � CrOH2+ + H+ −3.57 [3]2 CrOH2+ + H2O � Cr(OH)2

+ + H+ −6.27 [3]3 CrOH2+ + 2H2O � Cr(OH)3(aq) + 2H+ −10.93 [1,19]4 Cr(OH)3(aq) + H2O � Cr(OH)4

− + H+ −11.52 [1,19]5 2Cr(OH)4

− � Cr2O2(OH)42− + 2H2O 3.48 [1,19]

6 Cr(OH)3(am) + 2H+ � CrOH2+ + 2H2O 4.09 [1,19]6a Cr(OH)3·3H2O(cr) + 2H+ � CrOH2+ + 4H2O 7.28 [Current study]

sions of (i) crystalline Cr(III) hydroxide, Cr(OH)3·3H2O, with contact times of shortm available published data [19,21]. Solid and dashed lines were calculated using

At high supersaturation levels the least stable phase precipi-tates first [39,40] in accordance with Stranski’s rule or the OstwaldStep rule; such behaviour is often driven by sheer stereochemistry.Reaction products with simple structures form more rapidly thancomplicated, although thermodynamically stable, compounds. Thekinetics of transformation to the thermodynamically stable phasemay be variable, from hours to months and years. The results ofRai et al. [3,19] indicate that for freshly prepared Cr(III) hydrox-ides, if there is not any thermal aging procedure as in the casein [21], the establishment of equilibrium conditions in the acidicpH region between 3 and 5 requires a long equilibration time ofapproximately 500 days. This time frame is shorter in the neutraland alkaline pHs, where contact times in the order of 8 to 20 dayswere found to be sufficient [3,19]. Our current work suggests thatabove a pH value of 5.5, chromium(III), which is released from theunstable crystalline hydroxide, re-precipitates within 24 h approx-imating the solubility of amorphous compound (see Fig. 4).

3.2.2. Mixed Fe(III) and Cr(III) compoundsStability test data for the mixed Fe, Cr hydroxides are presented

in Fig. 5. Both systems are similar in their solubility behaviour. Onthe assumption that both phases behave as solid solutions of theform [Fe(OH)3](1−x)·[Cr(OH)3]x, the Cr(III) content in solution canbe estimated on the basis of the pure chromium hydroxide reactionconstant, KCr, and the Cr(OH)3 activity in the solid solution [41] as

Cr(OH)2+/(H+)2 = KCrxCr�Cr = KCr

′ (2)

where xCr is the molar fraction and �Cr is the activation coefficientof the chromium hydroxide in solid solution. According to [20], therelation between activity coefficient and molar fraction is as follows

log �Cr = 0.28(1 − xCr)2 − 1.79(1 − xCr)3 (3)

If it is assumed that the molar fraction is equal to 0.25, then

log �Cr = −0.60 (4)

Calculation of the apparent solubility constant, KCr′, was based

on the amorphous hydroxide, Cr(OH)3(am), for which log KCr = 4.09(Table 2) and from Eqs. (2) and (3), we derive

log KCr′ = 3.49 (5)

The continuous line in Fig. 5 was calculated by VMinteq on thebasis of an equilibrium constant for which log KCr

′= 3.49 and as

496 N. Papassiopi et al. / Journal of Hazardous Materials 264 (2014) 490– 497

-7.0

-6.5

-6.0

-5.5

-5.0

-4.5

-4.0

-3.5

-3.0

-2.5

-2.0

2 3 4 5 6 7

log[

Cr],

M

pH

FeCr (hydrolysis)

FeCr A (redox)

Calcula ted

Cr, detec�on limit

Cr, drin king wat er limit

Fig. 5. Chromium concentration in the aqueous phase at different pH values dur-ing the solubility tests of mixed hydroxides FeCr (hydrolysis) and FeCrA (redox). Thesolid line depicts the calculated Cr concentrations, assuming that the mixed hydrox-ifC

cth

sFr

FdTc(ss

-9

-8

-7

-6

-5

-4

-3

-2

2 4 6 8 10 12 14

log[

Cr],

M

pH

Cr, drinking water limit

Cr(OH)3.3H2O

Cr(OH)3 (am)

Fe0.75Cr0.25(OH)3

des are solid solutions of ferrihydrite and amorphous Cr(OH)3, corresponding to theormula [Fe(OH)3](1–x)[Cr(OH)3]x , where the mol fraction and activity coefficient ofr hydroxide are xCr = 0.25 and log �Cr = −0.60, respectively.

an be seen, it correlates satisfactorily with experimental data. Inurn, the fit indicates that Eq. (3) is a fair approximation of the Crydroxide activity in solid solution.

Fig. 6 represents the Fe concentration in solution for the variousolubility tests; in the figure, curve (1) corresponds to the calculatede(III) content on the basis of congruent dissolution according to theeaction (6)

+

[Fe(OH)3]0.75[Cr(OH)3]0.25 + 3H → 0.75Fe(III)aq

+ 0.25Cr(III)aq + 3H2O (6)

ig. 6. Iron concentration in aqueous solution at different pH values, as determineduring the solubility tests of mixed hydroxides FeCr (hydrolysis) and FeCrA (redox).he solid line (1) corresponds to the expected iron concentration assuming stoi-hiometric dissolution of the solid solution [Fe(OH)3](1−x)[Cr(OH)3]x , x = 0.25 (Eq.6)). The long-dashed line (2) depicts Fe concentration at equilibrium with theolid solution and activity coefficient of Fe hydroxide log �Fe = −0.12 (Eq. (11)). Thehort-dashed line shows the solubility of pure ferrihydrite.

Fig. 7. Comparison of calculated solubilities for a wide range of pH of hydroxidesCr(OH)3·3H2O, Cr(OH)3(am) and Fe0.75Cr0.25(OH)3.

Based on this, Fe(III) concentration in solution is expected to bethree-fold that of Cr(III). Curve (2) in Fig. 6 corresponds to the Fe(III)concentration at thermodynamic equilibrium with the solid solu-tion [Fe(OH)3]0.75·[Cr(OH)3]0.25, while curve (3) indicates Fe(III)equilibrium with pure ferrihydrite. Both curves (2) and (3) werecalculated via use of VMinteq software. Calculation of curve (2)involved the estimation of the apparent reaction constant KFe

′ ofthe solid solution according to reaction

[Fe3+]/[H+]3 = KFexFe�Fe = KFe′ (7)

where KFe is the solubility constant of pure Fe(OH)3 (for whichlog KFe = 3.5, [33]), while xFe and �Fe are the molar fraction andactivity coefficient of Fe(OH)3 in solid solution. Coefficient �Fe wascalculated based on �Cr by implementation of the Gibbs–Duhemequation [42]

dln gB

dxA=

[ −xA

(1 − xA)

](dln gA

dxA

)(8)

By combining Eqs. (3) and (8), we derive �Fe in solid solution asa function of Cr hydroxide molar ratio

log �Fe = −2.41xCr2 + 1.79xCr

3 (9)

If we let xCr = 0.25, we calculate

log �Fe = −0.12 (10)

and from Eq. (7), for xFe = 0.75 and log KFe = 3.5, we get

log KFe′ = 3.25 (11)

As seen in Fig. 6, the experimentally determined concentrationsof Fe are located between the theoretically calculated curves (1) and(2) or (3), suggesting that probably iron is initially released accord-ing to reaction (6) and tends to re-precipitate, since the aqueousphase is oversaturated with respect to either pure Fe(III) hydroxideor mixed Fe(III)–Cr(III) hydroxides.

Thermodynamically predicted solubilities for the three Crhydroxides in this study are compared in Fig. 7 for a pH range of2 to 14. As can be seen, the crystalline hydroxide Cr(OH)3·3H2Oexhibits a solubility in excess of 50 �g/l (drinking water EU limit)for all pH values. The amorphous hydroxide, Cr(OH)3(am), yieldssolubility values below the 50 �g/l threshold in the pH range of

5.7 to 11, while the mixed Fe0.75Cr0.25(OH)3 hydroxide expressesstability within the pH region of 4.8 to 13.5, where we note thatfurther stability testing of the mixed hydroxide needs to be carriedout towards the alkaline pH range.

zardou

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wlbtMtsi

ietthwbIr

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[

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199–214.

N. Papassiopi et al. / Journal of Ha

. Conclusions

The plain Cr(III) hydroxide, produced in the absence of iron,as found to be crystalline, abiding to the approximate molecu-

ar formula Cr(OH)3·3H2O. Based on stability tests, it was found toe highly soluble, therefore unstable, with a tendency to graduallyransform into the more stable amorphous hydroxide Cr(OH)3(am).

ixed Fe0.75Cr0.25(OH)3 hydroxides were found to possess a struc-ure similar to plain ferrihydrite, Fe(OH)3. Their solubility could beatisfactorily predicted as that of a solid solution of two hydroxides,.e. ferrihydrite and the amorphous Cr(III) hydroxide.

Both our experimental results and thermodynamic calculationsndicated that mixed Fe(III)–Cr(III) hydroxides can lead to a moreffective remediation of groundwater quality, in comparison withhe plain amorphous or crystalline Cr(III) hydroxides. The crys-alline hydroxide Cr(OH)3·3H2O was found to have a solubilityigher than 50 �g/l (maximum EU permitted Cr level in drinkingater), in all pH values studied. For amorphous Cr(OH)3(am), solu-

ilities lower than 50 �g/l can be achieved in the range 5.7 < pH < 11.n the case of mixed hydroxides Fe0.75Cr0.25(OH)3 the stabilityegion was extended, covering the range 4.8 < pH < 13.5.

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