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Geochimica et Cosmochimica Acta 108 (2013) 141–153 www.elsevier.com/locate/gca
Surface charge and growth of sulphate and carbonate green rust in aqueous media
Romain Guilbaud a,⇑,1, Maggie L. White b, Simon W. Poulton a,1
a School of Civil Engineering and Geosciences, Newcastle University, Newcastle Upon Tyne NE1 7RU, UK b School of Chemical Engineering and Advanced Materials, Newcastle University, Newcastle Upon Tyne NE1 7RU, UK
Received 16 March 2012; accepted in revised form 16 January 2013; available online 8 February 2013
Abstract
We report the first determination of the point of zero charge of sulphated and carbonated green rust particles. Green rust has been recognised as a prevalent mineral in environments such as hydromorphic soils, groundwaters and anoxic Fe(II)-rich water bodies, and the evolution of its net surface charge with pH has direct implications for the uptake of contaminants, met- als and nutrients in such settings. We find that the surface of both sulphated and carbonated green rust is positively charged at pH < 8.3, whereas it is negatively charged at pH > 8.3. Thus, alkaline settings will promote enhanced adsorption of metallic cations. However, the behaviour of ionic species surrounding green rust is more complicated than that predicted by simple pH-dependent adsorption, as our experiments suggest that green rust likely grows via dissolution–reprecipitation during Ost- wald-ripening. This implies that adsorbed species are potentially subject to repetitive steps of release into solution, re-adsorp- tion and co-precipitation during particle growth. The growth rate of green rust particles is highest within the first 50 min of aging, and appears to decrease towards an asymptote after 200 min, suggesting that particle growth controls on the uptake of dissolved species will be most important during the early steps of green rust growth. Our findings thus contribute to a better understanding of the controls that green rust may exert on dissolved ions in a variety of anoxic environments. Ó 2013 Elsevier Ltd. Open access under CC BY-NC-ND license.
1. INTRODUCTION (Loyaux-Lawniczak et al., 1999; Williams and Scherer, 2001; Bond and Fendorf, 2003; Legrand et al., 2004) Most iron oxide and oxyhydroxide minerals are highly can all be reduced by various forms of green rust. There- reactive in the environment (e.g. Poulton et al., 2004) and fore, green rust may potentially exert a key control on play a major role in toxic metal and contaminant remedia- the mobility of ions, the redox cycling of metals, and the tion (e.g. Waychunas et al., 2005). Amongst these minerals, degradation of organic and inorganic pollutants in the green rusts have recently been the focus of increasing atten- environment. tion, partly because of the presence of ferrous iron in the The structure and composition of green rust has been structure, which promotes their reducing capability. For widely documented (e.g. Refait et al., 1998; Ge´nin et al., example, chlorinated solvents (Erbs et al., 1998; Lee and 2001; Ruby et al., 2003; Simon et al., 2003), and consists Batchelor, 2002), nitrate (Hansen and Koch, 1998; Hansen of brucite-like mixed Fe(II)/Fe(III) tri-octahedral hydrox- et al., 2001; Choi and Batchelor, 2008) and heavy metals ide layers which alternate with negatively charged interlay- V 2 2 and non-metals such as As (Randall et al., 2001; Jo¨nsson ers containing planar (CO3 ,Cl) or tetrahedral (SO4 ) and Sherman, 2008), SeVI (Myneni et al., 1997) and CrIII anions. The occurrence of these different anions in the structure defines two distinct minerals, green rust 1 (con- ⇑ 2 2 Corresponding author. taining CO3 or Cl ) and green rust 2, (containing SO4 ). E-mail address: [email protected] (R. Guilbaud). Simon et al. (2003) showed that sulphated green rust 2 1 Present address: School of Earth and Environment, University (herein referred as GRSO4) is arranged in a hexagonal of Leeds, Leeds LS2 9JT, UK. geometry, differing with the rhombohedral type 1 green rust
0016-7037 Ó 2013 Elsevier Ltd. Open access under CC BY-NC-ND license. http://dx.doi.org/10.1016/j.gca.2013.01.017 142 R. Guilbaud et al. / Geochimica et Cosmochimica Acta 108 (2013) 141–153
(herein referred as GRCO3), although both display posi- sphere and the pH was stable at the beginning of each II III 2þ tively charged ½Fe4 Fe2 ðOHÞ12 hydroxide layers. experiment. Mixed Fe(II)–Fe(III) solutions were made by II III In the environment, the formation of green rust is not dissolution of Fe SO4 7H2O and Fe2 (SO4)3 5H2Oata uncommon and the mineral has been found in association ratio of x = 0.33 (where x = nFe(III)/[nFe(III) + nFe(II)]) as with microbial communities and various mineral assem- suggested by Ruby et al. (2003) in order to solely form II III II III blages. Since its first identification as a corrosion product the stoichiometric Fe4 Fe2 (OH)12SO4 mH2O and Fe4 Fe2 2 of steel water pipes (Stampfl, 1969), authigenic occurrences (OH)12CO3 mH2O species. Na2CO3 was used as the CO3 of green rust have been reported in soils (e.g. Trolard bearing reagent. et al., 1997; Abdelmoula et al., 1998), groundwaters (Chris- tiansen et al., 2009) and an anoxic Fe-rich lacustrine water 2.2. Precipitation and ageing of GRSO4 and GRCO3 column (Zegeye et al., 2012). Because green rusts are mixed Fe(II)–Fe(III) species, the roles of abiotic and biotic redox A first set of experiments consisted of the progressive reactions in their formation has been investigated. Shewanel- formation of GRSO4 and GRCO3 following the protocol la putrefaciens cultures can use Fe(III) mineral species as elec- described by Ruby et al. (2003). Fe(II) and Fe(III) ions were tron acceptors to form green rusts (e.g. Ona-Nguema et al., co-precipitated in a borosilicate reaction vessel by dropwise 2001; Zegeye et al., 2007a). Green rusts have also been ob- addition of a 0.1 M NaOH solution to a 0.05 M Fe solution 1 served during Fe(II) oxidation by nitrate reducing bacteria (100 mL) at a rate of 0.625 mL min to form GRSO4(See (Pantke et al., 2011) and during an oxygenic photoferrotro- Supplementary Information 1 for further discussion regard- phy (Kappler and Newman, 2004), while GRSO4 can be used ing the potential effect of dissolved Si on GR formation). as a sulphate source and electron acceptor during bacterial Similarly, GRCO3 was formed by dropwise addition of a sulphate reduction (Zegeye et al., 2007b). Green rust can also 0.1 M NaOH + 0.2 M Na2CO3 mixture (Bocher et al., form abiotically by partial oxidation of Fe(OH)2 (e.g. Schw- 2004) or a 0.2 M Na2CO3 solution to the 0.05 M Fe solu- ertmann and Fechter, 1994; Refait et al., 1997, 1998; Randall tion at a rate of 0.625 mLmin 1. Throughout the base addi- et al., 2001) or steel (e.g. Stampfl, 1969; S´wietlik et al., 2012), tion, pH was monitored, aliquots of the reaction were by partial reduction of ferrihydrite or lepidocrocite (Hansen, sampled, and the solids were vacuum filtered pending 1989), by reaction of Fe(II)aq with ferrihydrite, goethite or XRD and TEM analyses in order to check for potential hematite (Usman et al., in press), and by co-precipitation changes in mineralogy. of Fe(II)aq and Fe(III)aq under anoxic conditions (e.g. Ge´hin A second set of experiments consisted of ageing GRSO4 et al., 2002; Ruby et al., 2003; Bocher et al., 2004; Ruby et al., and GRCO3 in situ after rapid precipitation. 5 mL of 2006a,b; Ahmed et al., 2010; this study). NaOH or Na2CO3 was injected into 45 mL of Fe(II)/ Despite the increasing knowledge of biotic and abiotic Fe(III) solution (0.05 M, x = 0.33, OH:Fe = 2:1) in serum green rust formation mechanisms, and of the evident inter- bottles in order to quantitatively remove Fe from solution actions between green rust and metals or contaminants, lit- as GRSO4 or GRCO3. The solids were collected by vacuum tle is known about the surface chemistry and evolution of filtration after ageing on various timescales, from 5 min to green rust particles in aqueous media. This is due to the 69 h (in the case of GRSO4), and end-member minerals high reactivity towards oxidation by molecular oxygen, were characterised by XRD. and due to green rust metastability, which makes it difficult to experimentally assess the properties of green rust. In this 2.3. Solubility contribution, we report experimental insight into funda- mental aspects of the surface properties of green rust parti- Our aim here was not to specifically investigate the sol- cles. In particular, we demonstrate the mechanism of green ubility of green rust, which would require a completely dif- rust particle growth in solution, and provide the first deter- ferent study. Instead, our aim was to determine the balance mination of the particle surface charge as a function of pH. between adsorbed and desorbed protons at the surface of GRSO4 and GRCO3 as a function of pH, and therefore 2. EXPERIMENTAL METHODS the point of zero charge (PZC) for each mineral in aqueous solutions. However, during proton titrations, a portion of 2.1. Reagents the green rust may dissolve, depending on the pH. This mineral dissolution consumes protons and thus it is impor- Experiments were conducted under CO2-free, oxygen- tant that this consumption is not included in our determina- free conditions (<1 ppmv O2) in a re-circulating anoxic cab- tion of “sorbed protons” during the PZC measurement. Ò inet (Coy Laboratory Products Inc. ). A 98% N2 +2%H2 Therefore, it is critical to estimate the concentration of atmosphere was maintained through time by automatized H+ consumed during such dissolution. injection of the gas mixture, while minute amounts of oxy- The solubility equations for “anhydrous” GRCO3 and gen were scavenged by palladium catalysts. All reagents GRSO4 can be expressed as follows (Eqs. (1) and (2)): and acids were of analytical grade and solutions were pre- pared under oxygen-free conditions using 18.2 MX cm þ 2þ 3þ Fe6ðOHÞ CO3 þ 13H ¼ 4Fe þ 2Fe þ HCO þ 12H2O deionised water sparged with O -free grade N for 20– 12 3 2 2 ð1Þ 40 min (Butler et al., 1994). Rigorous exclusion of oxygen þ 2þ 3þ 2 is essential because Fe(II)aq and green rusts are prone to Fe6ðOHÞ12SO4 þ 12H ¼ 4Fe þ 2Fe þ SO4 þ 12H2O oxidation. Solutions were equilibrated with the inert atmo- ð2Þ R. Guilbaud et al. / Geochimica et Cosmochimica Acta 108 (2013) 141–153 143
The associated solubility constants have been docu- 0.1, 0.5 or 0.7 M KNO3) of known pH, and any subsequent mented by Rickard and Luther (2007) to be, log K = 39.1 pH change was tracked over 6 min (as suggested by Tewari and log K = 3.9, respectively, based on the potentiometric and McLean, 1972). NaCl was avoided as it may have pro- determination of the free energy of formation (Drissi moted the formation of chlorinated green rust. The rates of et al., 1995; Refait et al., 1999). In fact, poorly soluble ferric nitrate reduction by GRSO4 and GRCO3 are slower than iron may precipitate as secondary solid phases such as lep- those observed using chlorinated and fluorinated green idocrocite or goethite, and the overall reactions will reflect rusts (e.g. Hansen et al., 2001; Choi et al., 2012), and there an incongruent solubility where solids coexist, as described was no evidence for ammonium production during our by Eqs. (3) and (4) for the formation of ferric oxyhydroxide experiment (See Supplementary Information 2 for further (FeOOH) discussion regarding the potential nitrate reduction in our
þ 2þ experiment).The procedure was repeated using starting Fe6ðOHÞ12CO3 þ 7H ¼ 4Fe þ 2FeOOH þ HCO3 þ 8H2O electrolytes of various pH. An apparent pHPZC value is ob- ð3Þ tained when pH remains constant after mineral addition. þ 2þ 2 Although this approach can point towards the PZC va- Fe6ðOHÞ12SO4 þ 6H ¼ 4Fe þ 2FeOOH þ SO4 þ 8H2O ð4Þ lue, it does not take into account any matrix or dissolution- consuming H+ effects. Therefore, two titrations were per- Depending on which reaction governs green rust solubil- formed in 0.01, 0.04, 0.1, 0.5 or 0.7 M KNO3within situ pre- ity during the titrations, total [Fe(aq)T] in the supernatant cipitated GRSO4 and GRCO3 or with solid GRSO4 and will vary, whereby 1 mol of green rust is equivalent to GRCO3 added to the electrolyte. In both cases green rust 1 6 mol of Fe(aq)T for Eqs. (1) and (2). For Eqs. (3) and concentrations were 5gL . The solution pH was al- (4), however, 1 mol of green rust is equivalent to 4 mol of lowed to stabilise before HCl (0.02 M) addition (usually dissolved Fe(aq)T and 2 mol of FeOOH. Therefore, we <15 min), and the titration was performed from pH 9 to measured total [Feaq] in the supernatant throughout the 7.5 to (i) prevent back precipitation when increasing the experiment. Aliquots of the solution were sampled during pH, and (ii) maintain the experiment at a pH below the do- proton titrations (ranging from pH 9 to 7.5), filtered main of magnetite formation (see Sections 3.1 and 3.2). through a 0.45 lm membrane Milliporee filter, and fil- Blank titrations were performed from pH 9 to 7.5 on the trates were acidified with a few drops of concentrated supernatant (pH 9) collected after filtration using a H2SO4 for subsequent spectrophotometric [Fe(aq)T] deter- 0.45 lm membrane Milliporee filter. In this way matrix ef- mination. When smaller sized filters were used, particles fects were taken into account. However, this blank titration tend to clog off the filter and we preferred using 0.45 lm does not account for effects due to mineral dissolution and membranes as both filter sizes gave identical results. We solubility equilibration during the pH 9–7.5 titration, as the suspect that it is the propensity of GR particles to flocculate ionic background may change when pH decreases. Hence, that produces individual flocks which are large enough to solubility data were used to correct for these effects. For be trapped on a 0.45 lm membrane filter. each measured pH, the charge balance at the mineral sur- 1 2.4. Potentiometric determination of the PZC face, Q (in mol g ), was calculated using Eq. (5): C C ðC C ½H þ þ½OH Þ ½Hþ Q ¼ A B Ablank Bblank c PZC determination of mineral suspensions is usually de- m rived from electrokinetic studies or potentiometric titra- ð5Þ tions. Wolthers et al. (2005) compared the PZC of iron sulphides obtained by proton titrations (Widler and Sew- where CA and CB are the concentrations of acid and base ard, 2002; Wolthers et al., 2005) to the isoelectric point ob- added to the reaction or to the blank, [H+] and [OH ] tained using electrophoresis (e.g. Bebie´ et al., 1998) and are calculated from the measured pH using the Debye–Hu¨c- þ concluded that electrokinetic studies may result inoxidation kel activity coefficients, ½H c is the concentration of pro- of reduced mineral species. This suggests that the surface of tons consumed as GR solubility increases (as explained in oxygen sensitive minerals may be “refreshed” by proton Section 2.3) and m is the mass of the mineral. dissolution during anoxic potentiometric titrations, and We performed back-titrations from pH 7.5 to 9 to check thus needs to be considered with respect to green rust. for reversibility. This showed that depending on the exper- The principal difficulty in using potentiometric titrations imental duration, some titration and back-titration curves for green rust is its rapid conversion to other mineral were poorly matched, suggesting a significant degree of phases, because OH addition leads to the formation of green rust dissolution and precipitation of new Fe phases magnetite and H+ dissolution may form ferrihydrite or goe- when the pH increased during the back-titration. This thite (see below for the mechanisms of green rust forma- was confirmed by XRD analysis of the solids at the end tion). It is thus essential to (i) perform the titration within of the titration, which showed a magnetite component in the narrowest pH range possible; (ii) account for protolytic the XRD spectra. Kapetas et al. (2011) point out the ad- compounds in the solution; and (iii) account for green rust verse effect of allowing a long equilibration time between solubility and H+ consumption during mineral dissolution. each pH step on the reversibility of titrations. Although We constrained the pH domain of the PZC using the exper- their inspection of the optimal stability criterion before imental approach described by Ahmed and Maksimov the next acid/base addition accounted specifically for bacte- (1969) and Tewari and McLean (1972). A weighed amount riological systems, it is possible that similar effects occur of solid was added to a nitrate electrolyte (0.006, 0.01, 0.04, during the titration of metastable phases such as green rust. 144 R. Guilbaud et al. / Geochimica et Cosmochimica Acta 108 (2013) 141–153
We therefore developed a new titration technique in order 1980), k is the wavelength of the X-rays (0.154 nm for Cu to minimise the titration duration and to account for all Ka), and b is the full width of the peak at half maximum matrix effects in the blank. Here, the titration was per- height (FWHM). Although the Scherrer equation has been formed from pH 9 to 7.8 using increments of 0.2 pH units. used in nanomaterial studies to determine the size of the GRSO4 and GRCO3 were precipitated in duplicates, and crystal domain (e.g. Huang et al., 2003; Wolthers et al., after pH equilibration one replicate was filtered and the 2003; Jeong et al., 2008; Guilbaud et al., 2010), our aim here supernatant kept for the blank titration. The proton bal- was to track any difference in L over time. It is not known ance was calculated by the difference in acid or base addi- whether the use of glycerol may have a deteriorating effect tion to bring the sample and the blank to the desired pH on crystallinity and the FWHM. Nevertheless, we did not using Eq. (6): attempt to provide absolute crystal size but rather aimed to track variations over time. Precision on L is given by C C ðC C ½H þ þ½OH Þ Q ¼ A B Ablank Bblank ð6Þ the reproducibility between replicates and is ±1.4 nm (1r). m Using this method, the blank solution accounted for sol- 2.5.3. TEM analysis ubility effects and dissolution-consumed H+ at any pH, Aliquots of the solid phase were re-suspended in etha- since the supernatant was titrated after equilibration be- nol, and a drop of the solution was loaded and dried on tween green rust and the solution. We used this time con- a copper grid. Ethanol was used in order to ensure rapid suming method as a validity check for the titration curves evaporation and to minimise the air exposure time during obtained from the previous method. the drying process. TEM analysis was performed on a Phi- lips CM100 electron microscope operated at 100 kV. 2.5. Analyses 3. RESULTS Total dissolved Fe was measured by UV–Vis spectro- photometry. Solid phases were analysed by X-ray diffrac- 3.1. Formation of GRSO4 tion (XRD) for mineral characterisation, crystallinity and particle average size determination, and by transmission We performed the precipitation experiments as de- electron microscopy (TEM) for crystallinity and particle scribed and reviewed by Ruby et al. (2003, 2006a,b) in order size verification. to guarantee that we could reproduce the same mechanism of green rust formation, and also obtain the same products. 2.5.1. Spectrophotometry Fig. 2 plots R as a function of pH for the sulphated species Filtrates were acidified with a few drops of concentrated and shows very good agreement between our experiments H2SO4 and treated with potassium permanganate in order and those of previous studies. Ruby et al. (2003, 2006a,b) to quantitatively oxidise Fe(II)aq into Fe(III)aq. Total [Feaq] proposed a mass balance diagram where observed mineral (i.e. Fe(III)aq) was determined spectrophotometrically using assemblages are plotted as a function of x and R (with the thiocyanate method at 480 nm (e.g. Vogel, 1951)ona R = nNaOH/nFe(aq)T), and showed that pure green rust forms Genesis 6 spectrophotometer (Thermo Spectronice). for x = 0.33 and R = 2. They showed that for x = 0.33, the first mineral phase to form during base addition is a ferric 2.5.2. XRD analysis hydroxysulphate salt (i.e. a ferric oxyhydroxide where some - 2 Directly after filtration, aliquots of the solid phase were of the OH ions are substituted by SO4 ) and when R =1, coated with a drop of glycerol to prevent oxidation (e.g. the solid fraction of the system is only comprised of 2 Hansen, 1989; Williams and Scherer, 2001; Bocher et al., FeOOH. OH substitution by SO4 in the surface layer 2004) and loaded onto a glass sample holder. Fig. 1 illus- of the ferric oxyhydroxideis then followed by cluster sepa- trates typical XRD scans obtained for GRSO4 (Fig. 1A) ration from the ferric oxyhydroxide substrate to react with II and GRCO3 (Fig. 1B). The analysis of a green rust sample Fe (OH)2 and form GRSO4,when R = 2. Green rust con- was repeated after 8 and 24 h and displayed the same XRD version to magnetite and Fe(OH)2 occurs when R >2. spectrum, inferring that the glycerol coating efficiently pre- The whole mechanism is reflected in the titration curves by vented oxidation. Samples were analysed on a PANalyti- the twopH plateaux, and is confirmed by the TEM images that calX’Pert Pro MPD diffractometer using Cu–Ka radiation show the transition from [FeOOH + GRSO4]to[GRSO4 + ˚ (characteristic wavelength, k = 1.5418 A) generated at an Fe3O4 (magnetite) + possibly Fe(OH)2] assemblages. accelerating voltage of 40 kV and a current of 40 mA in a
5–45° 2h range, with a 0.0334° 2h step size. As suggested 3.2. Formation of GRCO3 by Randall et al. (2001), the green rust patterns used for data interpretation where those described by Vinsˇ et al. We constructed a similar titration graph for GRCO3 as (1987). for GRSO4 (Fig. 3A), and compared the results obtained We used the Scherrer equation (Eq. (7)) on the first dif- using the 0.1 M NaOH + 0.2 M Na2CO3 or the 0.2 M fraction peak to determine the average crystallite size: Na2CO3-only titrant solutions. Here, the solid assemblage is only comprised of green rust for R = 4 (as opposed to L ¼ kkðbcoshÞ 1 ð7Þ R = 2 for GRSO4; Fig. 2), since Na2CO3 +H2O= where L is the average crystallite dimension perpendicular 2NaOH + H2CO3 (and therefore RNa2CO3 =2RNaOH). to the reflecting planes (Warren, 1990), k is 0.91 (Brindley, Unsurprisingly, the 0.1 M NaOH + 0.2 M Na2CO3 titrant R. Guilbaud et al. / Geochimica et Cosmochimica Acta 108 (2013) 141–153 145
30000 A 25000 B GRSO4 GRCO3 25000 20000 20000 freshly precipitated 15000 freshly precipitated 15000 after 22 hours after 22 hours 10000
Relative intensity 10000
5000 5000
0 0 0 5 10 15 20 25 30 35 40 45 50 0 5 10 15 20 25 30 35 40 45 50 2oθ 2oθ
Fig. 1. Typical XRD spectra of freshly precipitated and glycerol-coated GRSO4 (Fig. 1A) and GRCO3 (Fig. 1B). Open circles represent the scan of the same glycerol-coated sample re-analysed after 22 h outside the anoxic cabinet. The similarity of the scans highlights the efficiency of glycerol to prevent oxidation.
1 2 3
100 nm 100 nm 500 nm
12
10 x = 0.33 3 x = 0.35 2 8 Ruby et al. (2003) (x = 0.33) 1 6 pH
GRSO4 4 => Mag
FeOOH => GRSO4 2 FeOOH 0 0 0.5 1 1.5 2 2.5 R = n(base)/n(Fe tot)
Fig. 2. Evolution of pH versus the relative amount of titrant added during the formation of sulphated green rust. Light grey circles represent the experimental data obtained starting with x = 0.33. Darker circles were obtained starting with x = 0.35. Our data are in good agreements with previous work (light grey curve). The two red crosses indicate the points where the reaction was sub-sampled and the solid products analysed by TEM. At pH 6.5, solid products consist of a mixture of green rust hexagons and lepidocrocite/goethite spikes. At pH 10.2, green rust hexagons have partially transformed to darker magnetite crystals. brings the system towards higher alkalinity than for 0.2 M fact, XRD analysis of the solid phase throughout the titra- Na2CO3 alone, and when GRCO3 is formed, there is a dif- tion revealed that although Na2CO3 was used and not ference of more than 2 pH units between the two experi- NaOH, GRSO4was still forming (Fig. 3B). This is due to ments (Fig. 3A). We found the higher pH system much the high proportion of sulphate (the starting Fe(II) and less stable than when using Na2CO3 only, as green rust rap- Fe(III) source are both iron sulphate salts) in the solution 2 idly transformed into magnetite, and we thus preferred to compared with CO3 at the start of the progressive base 2 2 use the Na2CO3-only solution for subsequent experiments. addition. The additional plateau reflects SO4 CO3 sub- 2 The difference between the GRCO3 (Fig. 3) and the stitution when CO3 becomes the dominant anion. GRSO4 (Fig. 2) formation graphs lies in the presence of a Using Mo¨ssbauer spectroscopy, Ruby et al. (2003, supplementary transformation for GRCO3asR increases. In 2006a,b) showed that when RNaOH > 2 (similarly, when 146 R. Guilbaud et al. / Geochimica et Cosmochimica Acta 108 (2013) 141–153
14 pH
NaOH + Na2CO3 A 67891011 12 -2 Na2CO3 -2.5 10 y = -1.8x + 11.3 -3 8 R² = 0.91 ]
T -3.5 pH 6 GRCO3 -4 GRSO4 => GRCO3 => Mag 4 FeOOH -4.5 => GRSO 4 GRSO 2 -5 4
FeOOH log [Fe(aq) GRCO3 0 -5.5 012345 -6 R=n(base)/n(Fetot) -6.5 -7
GRSO 4 B GRCO 3 Fig. 4. Total dissolved [Fe(aq)T] plotted against pH for both green GRCO GRSO 3 4 rust species. The dependence of log Fe(aq) upon pH is supported GRSO4 T GRSO 2 4 by the coefficient of determination (R = 0.91) on the regression GRCO3 line.
GRSO4 different electrolyte concentrations. However, the given pre-
Relative Intensity (counts) pH 9 cision does not permit us to strictly differentiate between the pH 8 pH 7 GRCO3 and the GRSO4 results. A broader examination of the plot gives us some indication on the mechanism that 10 15 20 25 30 35 40 governs GR solubility in our experiment. Reaction con- θ 2° stants K for Eqs. (1) and (2) can be written as (Eqs. (10) Fig. 3. Evolution of pH versus the relative amount of titrant added and (11)): during the formation of carbonated green rust (Fig. 3A) using ½HCO ½FeðaqÞ 6 different titrant solutions. During the titration, GRSO is the first K 3 T 10 4 ¼ þ 13 ð Þ green rust species to form as indicated by the red spectrum ½H 2 6 (Fig. 3B, pH 7). Anion substitution, i.e. GRSO4 to GRCO3 ½SO4 ½FeðaqÞT conversion, starts when [CO 2 ] > [SO 2 ]. K ¼ ð11Þ 3 4 ½Hþ 12
Isolating [Fe(aq)T] in these equations gives a pH dependent 1 K RNa2CO3 > 4), green rusts rapidly convert to magnetite and linear equation where log½FeðaqÞT 6 log ½anion 2 pH, and Fe(OH)2. The equation can be written as an acid–base reac- therefore the slope would be 2 in a log [Fe(aq)T] versus tion (Eq. (8)): pH plot. Similarly, for Eqs.3 and 4, the slope would be 1.6. Note that the intervention of Fe(III) complexes II III aq Fe Fe ðOHÞ CO3 þ OH + 4 2 12 such as Fe(OH)2 would display even lower slopes, around II III 1.5, and the formation of stoichiometric ferrihydrite ¼ Fe1 Fe2 O4 þ 3FeðOHÞ2 þ HCO3 þ 3H2O ð8Þ (Fe5HO8 4H2O) would either be pH independent or would Ferrous hydroxide is poorly crystalline in aqueous med- display a slope of 1. Our data suggest that during the pH ia, and we could not detect its presence in the XRD spec- 9–7.5 titrations, a small fraction of FeOOH solids may con- trum of the end products. Alternatively, the common tribute to the reaction, but ferrihydrite is unlikely to be anoxic Fe(OH)2 to magnetite conversion (the so-called present. The PZC for freshly precipitated goethite has been “Schikorr reaction”; Schikorr, 1933) may take place and determined to be 9.3 ± 0.1 (Antelo et al., 2005), implying the overall reaction is (Eq. (9)): that throughout our titrations, the surface of potential FeOOH solids is positively charged, minimising any H+ Fe6ðOHÞ12CO3 þ OH ¼ 2Fe3O4 þ H2 þ HCO3 þ 5H2O adsorption effect. Instead, in their model using phosphate ð9Þ as an adsorbed species, Antelo et al. (2005) suggest that for a pH range of 7.8–9, non-protonated phosphate com- 3.3. Solubility plexes largely dominate the inner-sphere species. Therefore, the positively charged goethite particles sites are fully com- Total dissolved Fe measured during the course of the pensated by non-protonating anions and no further H+ titrations performed on in situ precipitated GR shows some adsorption is likely to occur. One could argue that the lin- dependency on pH for both GRCO3 (using 0.01 M KNO3) earity discussed above is in fact a combination of an -2 and GRSO4 (using 0.04 M KNO3). Fig. 4 shows that slope correlation line for pH > 8 and a lower slope correla- although the duplicate analysis gave a poor precision (max- tion line for pH < 8. This may indicate that incongruent imum ± 0.5log Fe(aq)T),log [Fe(aq)T] tends to plot linearly dissolution of green rust only takes place at lower pH, as a function of pH, with a slope of 1.8 for both mineral implying that green rust is the only solid phase from pH species. It is possible that the apparent shift between the 9 to 8. The very high green rust concentration used in our two plots may be due to different solubility products using experiments ( 5gL 1) suggests that the green rust surface R. Guilbaud et al. / Geochimica et Cosmochimica Acta 108 (2013) 141–153 147
A
B