Enthalpy of formation of sulfate green rusts

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Geochimica et Cosmochimica Acta 72 (2008) 1143–1153 www.elsevier.com/locate/gca

Enthalpy of formation of sulfate green rusts Lena Mazeina a

a,1

, Alexandra Navrotsky

a,*

, Darby Dyar

b

Peter A. Rock Thermochemistry Laboratory and NEAT ORU, University of California at Davis, One Shields Avenue, Davis, CA 95616, USA b Department of Earth and Environment and Department of Astronomy, Mount Holyoke College, South Hadley, MA 01075, USA Received 22 May 2007; accepted in revised form 28 November 2007; available online 8 December 2007

Abstract As a contribution to the systematic study of iron oxide thermodynamics, this work reports enthalpies of formation of green rust, a double layered (FeII, FeIII) hydroxide with the ideal stoichiometry FeII 1x FeIII x ðOHÞ2 ½Am y=m  nH2 O, with sulfate as the anion in the interlayer. Samples were characterized by X-ray powder diffraction, thermogravimetric analysis, infrared spectroscopy, and Mo¨ssbauer spectroscopy. Full chemical analysis was performed. Contents of FeII, FeIII, water, and sulfate were obtained. We report standard enthalpies of formation for green rust with different FeII/FeIII ratios. Enthalpies of formation from single cation compounds, namely, Fe(OH)2, Fe(OH)3, FeSO4 and H2O show reasonable agreement with Gibbs free energies of formation from single cation compounds recalculated from the reported literature values. These values show that green rust has little stabilization over a mechanical mixture of these single cation compounds and there is no thermodynamic preference for any particular FeII/FeIII ratio. Ó 2007 Elsevier Ltd. All rights reserved.

1. INTRODUCTION Green rust is a layered double (FeII, FeIII) hydroxide (LDH) with the ideal stoichiometry FeII 1x FeIII x ðOHÞ2 ½Am y=m  nH2 O, where anion Am is commonly SO4 2 ; Cl ; CO3 2 ; m is charge of the anion, x is content of ferric iron and n is content of water. Green rusts occur in soils (Koch and Moerup, 1991; Randall et al., 2001; Erbs, 2004), in corrosion products (Rodriguez and Gonzalez, 2006) and as a by-product of bacterial activity (Rodriguez and Gonzalez, 2006). Green rusts also influence soil redox properties and can control iron redox cycling in many environments (Roden, 2006). They are often fine-grained and/or poorly crystalline (Delnavaz and Allmann, 1988; Gehin et al., 2002), and have high surface area (Cuttler et al., 1984). Because of their large active surface, green rusts sorb a wide variety of contaminants including heavy metals (Bond and Fendorf, 2003; Herbert, 2003), toxic inorganic and organic com-

*

Corresponding author. Fax: +1 530 752 9307. E-mail address: [email protected] (A. Navrotsky). 1 Present address: Naval Research Laboratory, 4555 Overlook Avenue, Washington, DC 20375, USA. 0016-7037/$ - see front matter Ó 2007 Elsevier Ltd. All rights reserved. doi:10.1016/j.gca.2007.11.032

pounds (Johnson and Bullen, 2003; O’Loughlin et al., 2003a) and radioactive elements (Dodge et al., 2002; O’Loughlin et al., 2003b), thus preventing them from moving into the environment or spreading further into the corrosion layer. Quantitative description of all these processes at ambient and elevated temperatures requires a full set of thermodynamic data, namely, Gibbs free energy, enthalpy and entropy, for the green rust phases. Green rusts, as all LDHs, are solids with rather variable compositions (different FeII/FeIII ratios, various interlayer anions, and variable amount of interlayer water). The most studied green rust compounds are those with FeII–FeIII ratios equal to 2 and 3 and the structure of relatively wellcrystalline green rust with sulfate in the interlayer was determined by Rietveld analysis (Simon et al., 2003). The Gibbs free energies of formation of green rust with FeII/ FeIII ratios of 2–2.5 have been determined using potentiometric titration (Detournay et al., 1975; Hansen et al., 1994; Refait et al., 1999) and/or solubility measurements (Detournay et al., 1975; Hansen, 2001). Nevertheless, green rust has wider range of FeII/FeIII ratios—0.6–4.0 (Cuttler et al., 1990; Cornell and Schwertmann, 1996). These green rusts as well as the poorly crystalline ones are less investigated both structurally and

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thermodynamically. Systematic thermodynamic data for the wide compositional range and different crystallinity degree are necessary for calculations involving green rust equilibrium with other phases, during redox, corrosion and other processes. Additionally, measurements of Gibbs free energy often cover too small a temperature range for accurate calculation of DH f and DS f from the temperature derivative of DGf . Values of enthalpy and entropy, that normally require direct measurements, are lacking, in part because of extreme instability of green rust in air. Rapid FeII oxidation makes characterization and analysis difficult. The main goal of this study is to measure the enthalpy of formation of green rusts with different FeII/FeIII ratios and crystallinity degree and to further examine the thermochemical systematics of green rusts as members of the larger group of LDHs. We report enthalpies of formation of fully characterized green rusts with sulfate in the interlayer and FeII/FeIII ratios ranging from 0.50 to 1.34. Acid solution calorimetry, used in this study, allows preserving FeII in the solution and provides a reproducible final state of both green rust and reference phases. The enthalpy from this work and free energy from literature sources allow an estimation of entropy of formation. 2. EXPERIMENTAL 2.1. Materials All chemicals, NaOH (EM Science, Merck), FeSO4Æ7H2O and Fe2(SO4)3ÆxH2O (both P99%, Alfa Aesar) were used without preliminary treatment. Water content x in Fe2(SO4)3ÆxH2O was determined by thermogravimetric analysis (TGA) just before the synthesis and is 8.7 ± 0.1. Deionized water used for solution preparations was purged of dissolved O2 and CO2 by bubbling with Ar for at least 1 h before use. We used purchased FeSO4Æ7H2O (99.999% metal basis, Alfa Aesar) as a reference phase for calorimetry. It was characterized prior to calorimetry by powder X-ray diffraction (XRD) and by TGA. The XRD patterns of FeSO4Æ 7H2O corresponded to that of melanterite, and the number of waters of hydration was confirmed to be 7 by TGA. 2.2. Synthesis A common method of green rust synthesis is controlled oxidation of FeII solutions. Oxygen is introduced at a very slow rate and the solution is titrated with base to increase pH and cause co-precipitation of FeII and FeIII hydroxides. Most of the samples were synthesized starting with FeII solutions. Only one sample (described last in Section 2.2) was synthesized from a solution containing both FeII and FeIII at the beginning. Sample nomenclature is given according to the FeII/FeIII ratio determined by chemical analysis. Samples GR-1.34 and GR-1.21 were synthesized by titration of FeII-solution (prepared by dissolving 30 and 25 g of FeSO4Æ7H2O, respectively, in 250 ml distilled water) by 1 M NaOH up to pH = 7 (Schwertmann and Cornell, 1991). At this pH a green precipitate was formed and bubbling by O2 was stopped.

A solution of 0.01 M Fe2+ prepared from FeSO4Æ7H2O was used to synthesize sample GR-0.89 (method modified from Drissi et al., 1994). The solution was first titrated under an Ar atmosphere by freshly prepared NaOH/Na2CO3 (0.3 M/0.18 M) up to pH 6.8. Then, O2 was bubble through the solution with a flow rate 0.5 to 0.8 ml/min till the solution became green. Sample GR-0.50 was synthesized by the co-precipitation method of Gehin et al. (2002). A quantity of 11.63 g of FeSO4Æ7H2O and 3.48 g of Fe2(SO4)3Æ8.7H2O were mixed in distilled water in a three-neck flask under an Ar atmosphere. The mixture was then titrated by 0.3 M NaOH with constant stirring till the pH reached 6.8 and the color became bluish-green. The precipitates were held in the mother solution for 24 h for aging with constant Ar bubbling through the solution. Then samples were freeze dried, transferred under N2 to a glove box and stored there under an Ar atmosphere with <1 ppm O2. No color change was observed during storage of samples. 2.3. Characterization After freeze-drying, samples were analyzed without further treatment. To prevent oxidation, all sample preparation was performed under Ar in a glove box. XRD patterns were obtained using a Scintag PAD V diffractometer operating at 45 kV and 40 mA using Cu Ka (k = 0.154056 nm) radiation. Patterns were collected from 10° to 65–70° 2h with a step size of 0.02° 2h angle and dwell time of 2 s. The length of scan and dwell time were chosen to obtain distinguishable peaks before the sample oxidizes. The sample was covered by glycerol to retard FeII oxidation as recommended in the literature (Hansen, 1989; Schwertmann and Cornell, 1991). Samples covered with glycerol remained unoxidized for 1.5 h, a typical duration of an XRD run. For some runs over two hours, a slight reddish color to the surface was observed, but the bulk of the sample still remained green. The peak positions and particle size were determined using JADE 6.1 software (Materials Data Inc., Livermore, 2001). Lattice parameters were calculated from a formula relating d-spacings and hkl-indexes for the hexagonal lattice of the brucite-like structure (Cornell and Schwertmann, 1996). Lattice parameters c were calculated for each d-spacing (0 0 1), (0 0 2), (0 0 3), (0 0 4) and (0 0 5) where possible. The average value of c was then used to calculate the a lattice parameter using peaks (1 0 1) and (1 0 2) when possible. Error in the calculation of the c and a lattice parameter is one standard deviation from the average values. Particle size was estimated from peak broadening using formula (Sherrer equation) implemented in the software. Fourier transform infrared (FTIR) spectra were measured with a Bruker Equinox 55 spectrometer using the KBr pellet technique. Pellets (13 mm diameter, 150 mg) were pressed until the mixture of KBr and sample became translucent. The infrared (IR) spectra were recorded immediately after pellet preparation. The spectrometer was flushed continuously with nitrogen to minimize contamination by atmospheric water and CO2 during analysis. Spectra were collected in the 400–4000 cm1 range with a

Enthalpy of formation of green rusts

resolution of 4 cm1. A baseline correction was applied before interpretation. Both sulfate analysis, performed by ion chromatography, and water content, determined by Karl-Fisher titration, were done at Galbraith Laboratories, Inc. Additionally to this water analysis, weight loss experiments were performed. The samples were heated to 1100 °C for at least 12 h. The water content was recalculated from the measured weight loss according to the following reaction: FeII 1x FeIII x ðOHÞ2þx2y ðSO4 Þy  nH2 O þ

1x 1100  C O2 ƒƒƒ ƒ! 2

x Fe2 O3 þ ðn þ 1 þ 0:5x  yÞH2 O " þySO3 " 2

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areas of all FeIII components relative to the total area are accurate to within ±3–5% absolute. The very small sextet labeled FeII in the 295 K data is poorly defined (and its parameters are in fact difficult to interpret), but its presence was needed to obtain satisfactory fits to the data. We did not correct for differential recoil-free fraction because the appropriate f values are not known for this material; however, based on analogies to melanterite and other hydrous sulfates, such a correction is likely to be small and within error bars (Rothstein et al., 2005). 2.4. Calorimetry

ð1Þ

Values of x, y, and n were calculated from elemental analysis. We used slightly different than ideal chemical formulas of green rusts for reasons described in Section 4. The analysis of FeII and FeIII was performed at the Department of Geology at the University of New Mexico (Albuquerque, NM) using slightly modified procedures as outlined by Kolthoff and Sandell (1952) and Furman (1962). The samples were dissolved with hydrofluoric and sulfuric acid. The ferrous iron was titrated with standard potassium dichromate solution using sodium diphenylamine sulfonate as an indicator. The total iron content was determined using ICP-MS, and FeIII content was the difference between FeII and total iron concentration. In addition, FeII and FeIII concentrations were determined by Mo¨ssbauer spectroscopy at Mount Holyoke College on sample GR-0.50, in order to corroborate results of elemental analysis. To avoid the possibility of oxidation of the sample during shipping, mounting and handling special techniques were employed. The sample was mounted in a holder and sealed in a glove box under Ar and then shipped in a triple container for analysis, where it was stored in a freezer at 20 K. For analysis, the sample was attached to the brass rod sample holder in the Mo¨ssbauer apparatus and then inserted into a cryostat pre-cooled to 12 K, followed by evacuation and then purging with He gas. Total exposure time of the sealed mount to ambient conditions was less than 1 min. After a run at 12 K, the sample temperature was increased to 293 K without opening the sample chamber. Although the color of the sample did not change, oxidation (both during shipment and storing prior to analysis) cannot be fully excluded due to the very high reactivity of green rust. The Mo¨ssbauer spectra were acquired using a source of 60 mCi 57Co in Rh on a WEB Research Co. Model 302 spectrometer equipped with a Janus 4 K cryostat. Spectra were corrected to remove the fraction of the baseline due to Compton scattering. Run times were 24 h, and baseline counts were 19 and 22 million for the 15 and 293 K spectra, respectively, after the Compton correction. Spectra were collected in 2048 channels and modeled using an in-house program Dist3e (University of Ghent, Belgium). Isomer shift (IS or d) and quadrupole splitting (QS or D) of the doublets were allowed to vary, and widths of all peaks were coupled to vary in unison. Errors on isomer shifts are estimated at ±0.04 mm/s because of high peak overlap and low signal-to-noise ratios. Quadrupole splitting values have uncertainties of ±0.05 mm/s. The summed

For the determination of enthalpies of formation of green rust, the enthalpy of solution (DHsol) was measured in a Hart Scientific IMC-4400 isothermal calorimeter by dissolving approximately 4–7 mg of a pelletized sample in 25.0 ml of 5.0 N HCl (Alfa Aesar), maintained at 25.0 ± 0.1 °C. Pellet preparation was carried out in an Ar atmosphere. The measured heat effect of each drop was 1–3 J. In a separate set of experiments, we determined that the calorimeter in this configuration is capable of detecting thermal effects of 0.02 J. Calorimeter calibration was performed by dissolving KCl (NIST standard reference material 1655) in deionized water at the same temperature. To have a complete calorimetric cycle, one needs to know final states of all components after dissolution. The initial state in every calorimetric experiment is a crystalline green rust and 5 N HCl. The final state for FeIII after the complete dissolution of the sample is a dilute solution of FeIII in HCl, always at nearly the same concentration, namely 30–50 mg/L. The final concentration of FeII in 5 N HCl was determined by iodometric titration of FeII present after dissolution. Due to the insufficient amount of other samples, only sample GR-0.50 was titrated to determine if any FeII oxidizes to FeIII during dissolution. A pelletized sample (4 to 5 mg) was dissolved in HCl at the same conditions as during calorimetry. Immediately after sample dissolution (which took less than 5 min), the beaker was placed under Ar to prevent iodine oxidation by air. The solution was then titrated with standardized thiosulfate solution according to the method described in Vogel (1961). The molar amount of Na2S2O3 spent to titrate the solution was equal to the molar amount of FeII in solution. The results (Table 1) show that concentration of FeII after sample dissolution corresponds to the initial concentration of FeII in the sample. Thus, no FeII oxidizes to FeIII, and all the FeII remains in this oxidation state. This is supported by the Pourbaix diagram of iron in aqueous solutions at low pH (Pourbaix, 1963). A solution of FeSO4Æ7H2O in 5 N HCl did not show any yellow color that would indicate FeIII, further confirming the stability of FeII in 5 N HCl. At high concentrations of HCl, the predominant FeII and FeIII species are FeCl+ (Heinrich and Seward, 1990) and FeCl2 þ (Tagirov et al., 2000) complexes. Because all such FeCl+ and FeCl2 þ complexes cancel out in the thermochemical cycles (as long as overall concentrations are similar), these aqueous species will not be considered fur-

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Table 1 Iodometric titration data for sample GR-0.50 Pellet weight

lmol of Na2S2O3 used

5.68 4.67 4.60 4.31 Average

31.3 23.9 20.2 24.5 25.0 ± 4.6

a

Fe2+ present after dissolution lmol

mg

wt%

31.3 23.9 20.2 24.5 25.0 ± 4.6

1.75 1.34 1.13 1.37 1.4 ± 0.3

30.8 28.6 24.6 31.7 28.9 ± 3.2a

Compare with elemental analysis, 28.2 wt% (Table 3).

ther in the thermochemical cycle in this paper. The reference phases were lepidocrocite, c-FeOOH (data from Majzlan et al., 2004), a-MgSO4 and MgO (data from Majzlan et al., 2005), FeSO4Æ7H2O and liquid water. 3. RESULTS 3.1. Characterization All obtained materials were bluish-green to dark green in color. Although some samples showed poor crystallinity and broad peaks, XRD patterns corresponded to the ones known for well-crystalline green rust II (Hansen et al., 1994; Hansen, 2001) with sulfate in the interlayer. Sample crystallinity decreases in the order GR-0.89 (particle size 30 nm), GR-0.50 (particle size 20 nm), GR-1.34 (particle size 15 nm), GR-1.21 (particle size 5–10 nm). We also observed possible admixtures of X-ray amorphous material indicated by a broad hump in the X-ray pattern at 2h = 20– 25°, but additional XRD tests on sample with known X-ray pattern covered with glycerol showed that glycerol is responsible for the hump in the X-ray patterns. Due to the low intensity of some peaks, their breadth, iron fluorescence and consequently high background, it was not possible to refine the lattice parameters using Rietveld analysis; therefore lattice parameters (shown in Table 2) were manually calculated from the interplanar distances dhkl. For the same reasons, the errors for some samples are relatively large. Both lattice parameters a and c correspond to those reported for green rusts with sulfate in the interlayer (Hansen et al., 1994; Hansen, 2001). FTIR confirmed that all samples have only sulfate in the interlayer, which is especially important for the sample GR-0.89 because a solution containing carbonate was used during its synthesis. All samples showed bands at 1111– 1126 cm1, at 668 cm1 and at 612–617 cm1 corresponding to the SO4 2 frequencies (Peulon et al., 2003). No peaks

Table 2 Lattice parameters of green rusts ˚) Sample ID d(hkl = 001) (A

˚) c (A

GR-1.34 GR-1.21 GR-0.89 GR-0.50

10.95 10.85 10.99 10.94

11.0250 10.8560 10.9523 10.9089

˚) a (A (8) (4) (4) (3)

3.156 (1) 3.16 (3) 3.171 (3) 3.18 (1)

Errors (in brackets, referring to the last digit) were calculated as standard deviation from the mean.

at 1300–1360 cm1 corresponding to carbonate vibrational modes were detected. Elemental analyses are given in Table 3. OH-content is calculated from FeII, FeIII, and SO4 2 content based on charge balance. Results of the loss on ignition experiments together with calculated weight loss from chemical analysis are shown in Table 4. The errors for calculated results are estimated from analytical error of 1 wt% for H2O and S, and 5 wt% for FeII, FeIII and therefore for OH-content. Experimental values are a little smaller than calculated, but both in agreement within the experimental error. Mo¨ssbauer data (Fig. 1) are presented in Table 5, with magnetic FeIII and FeII components listed first for 15 and 295 K spectra. The 15 K spectrum is dominated by a magnetic sextet outlined with a dashed line in Fig. 1 (44% of the total area), with d = 0.48 mm/s, D = 0.26 mm/s, and BHf = 49.4 T. At first approximation, this feature could be attributed to goethite or another FeOOH polymorph (Murad and Johnston, 1987). Such an assignment would be confirmed if a sextet with appropriate BHf (38.1 T) and quadrupole splitting (0.26 mm/s), or paramagnetic doublet with QS = 0.48 mm/s arising from nanophase goethite is present at 295 K (which it is not). According to Koch (1998), these features (both sextets) might represent disorder in cation distribution in green rust. Random cation distributions in green rusts can give rise to distributions in the electric field gradient, and smear out the electric field distribution. This can result in additional peaks in the Mo¨ssbauer spectrum. Furthermore, substitution of by O2 and H2O for OH in the green rust structure might cause additional interactions to produce asymmetry in line intensity and peak distortion, though peak areas will be unaffected. Cation disorder might also explain the relatively broad peaks in our Mo¨ssbauer spectra relative to the more narrow ones reported in the literature for well crystalline phases (e.g., Gehin et al., 2002). The next largest component in the 15 K spectrum is a FeII doublet with d = 1.30 mm/s and D = 2.79 mm/s, which has been unambiguously identified as green rust (e.g., Gehin et al., 2002). Small components corresponding to Fe3+ in the rust and a small broad Fe2+ peak that resembles a partially-ordered oxide are also found. The 295 K spectrum is equally complex. It is dominated by paramagnetic species, particularly two strong doublets (d = 0.30 and 0.39 mm/s and D = 0.60–0.90 mm/s, respectively) assigned to FeIII in green rust (e.g., Koch, 1998; Gehin et al., 2002). The magnetic sextet that was dominant in the 15 K spectrum (with d = 0.48 mm/s, D = 0.26 mm/s, and BHf = 49.4 T) is not present, as well as a paramagnetic doublet with QS = 0.48 mm/s, as would be expected to occur from bulk or nanophase goethite, respectively (e.g., Cornell and Schwertmann, 1996). This supports the idea the sextet in the 15 K spectrum arises from magnetic ordering of some phase that is paramagnetic at 295 K. At 295 K, the two sextets have BHf = 34.1 and 33.1 T. Possible assignment of these sextets might include goethite (d = 0.37 mm/s, D = 0.26 mm/s, and BHf = 38.2 T), magnetite (d = 0.26 and 0.67 mm/s, D = 0.02 and 0.00 mm/s, and BHf = 49.0 and 46.0 T), and maghemite (d = 0.32 mm/s, D = 0.02 mm/s, and BHf = 49.9 T). However,

Enthalpy of formation of green rusts

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Table 3 Chemical composition Sample ID

GR-1.34 GR-1.21 GR-0.89 GR-0.50 a

FeII

FeIII

S

H2O

Recalculated

mol

wt%

mol

wt%

mol

wt%

mol

wt%

OH

0.57 0.55 0.47 0.33

23.3 23.5 20.1 14.1

0.43 0.45 0.53 0.67

19.2 17.5 22.5 28.2

0.36 0.38 0.34 0.33

9.2 8.3 8.3 8.0

0.91 0.90 0.82 0.59

12.0 12.36 11.26 8.09

1.72 1.70 1.85 2.01

Sample ID

Experimental

Calculated from elemental analysis and Eq. (1)

GR-1.34 GR-1.21 GR-0.89 GR-0.50

38.2a ± 3.0b (6)c 36.5 ± 0.9 (2) 32.6 ± 2.7 (2) 28.1 ± 2.6 (3)

41.0 ± 6.0 41.8 ± 5.9 40.7 ± 6.1 39.6 ± 6.2

a

c

II

Formula III

III

Fe /Fe

Fe /S

1.34 1.21 0.89 0.50

1.20 1.20 1.55 2.02

FeII 0:57 FeIII 0:43 ðOHÞ1:72 ðSO4 Þ0:36  0:91H2 O FeII 0:55 FeIII 0:45 ðOHÞ1:70 ðSO4 Þ0:38  0:90H2 O FeII 0:47 FeIII 0:53 ðOHÞ1:85 ðSO4 Þ0:34  0:82H2 O FeII 0:33 FeIII 0:67 ðOHÞ2:01 ðSO4 Þ0:33  0:59H2 O

Recalculated from the charge balance.

Table 4 Weight loss, wt%

b

a

Average. Two standard deviations of the mean. Number of measurements.

the sextets are very poorly resolved, with d = 0.1 and 0.8 mm/s and D = 0.46 and 0.06 mm/s, respectively. We cannot definitely assign the sextets here to any specific mineral phase, though phases such as magnetite and maghemite can probably be ruled out on the basis of their high BHf. If goethite is present, its abundance would be 21% of the total Fe in the sample, it means that the volumetric mixture of phases present would be roughly 10% goethite and 90% green rust. Such a large amount of goethite would have been seen by other methods, e.g., FTIR or/and XRD. Since these features are present in both 15 and 295 K spectra, we exclude that they appeared due to the oxidation of green rust during warming up from 15 to 295 K. Thus, we think that, these features indeed represent disorder in cation distribution in green rust as mentioned by Koch (1998). An FeII doublet with d = 1.18 mm/s and D = 3.24 mm/s is present in the 295 K spectra, although it is half the size of the analogous doublet in the 15 K spectrum; this implies that some peak area is being shared with the other heavily-overlapping FeII doublet at d = 1.25 mm/s and D = 1.95 mm/s. The minor components in these fits might well be caused by the fitting algorithm attempting to compensate for and model the asymmetry in line intensity that results from substitution and cation disorder. For both spectra, the FeIII contents, determined from the peak areas, are 60 mol% at 15 K and 67 mol% at 295 K (both with the maximum error ±5 mol%), agreeing very well with the results of elemental analysis of that sample (67 ± 3 mol%). We conclude that there is no compelling evidence for the presence of additional phases although they cannot be fully ruled out. 3.2. Calorimetric data Results of acid solution calorimetry are given in Table 6. For each sample, the reported enthalpy of solution, DHsol,

represents the mean of several measurements with an associated error of two standard deviations of the mean. The enthalpies of formation from the elements, DH f , were calculated from the thermodynamic cycle shown in Table 7. For better comparison of our data on enthalpy and literature data on Gibbs free energy, we normalized literature data for 1 mol of iron. Additionally, we calculated Gibbs free energy based on literature data and enthalpy of formation relative to single cation components (DfGscc and DfHscc, respectively) as done by Allada et al. (2006) for analogous hydrotalcite structures: xFeðOHÞ3 þ ð1  x  yÞFeðOHÞ2 þ yFeSO4 þ nH2 O ! FeII 1x FeIII x ðOHÞ2þx2y ðSO4 Þy  nH2 O

ð2Þ

The reason we write the chemical formula of the green rust this way is discussed below. According to reaction (2) enthalpy of formation from single cation compounds (DfHscc) would represent the difference between enthalpies of formation of green rust (DfHGR) and sum of enthalpies of formation of single cation compounds (DfHi) multiplied by the coefficient (ni) according to the reaction (2): Df H scc ¼ Df H GR  Rni  Df H i

ð3Þ

Gibbs free energies of formation from single cation compounds (DfGscc) were calculated analogously: Df Gscc ¼ Df GGR  Rni  Df Gi

ð4Þ

where DfGscc is the Gibbs free energy of green rust from single cation compounds; DfGGR is Gibbs free energy taken from literature (Detournay et al., 1975; Refait et al., 1999; Hansen, 2001); and DfGi is Gibbs free energy of single cation compounds from reaction (2) taken from different sources (see Table 8) multiplied by the chemical coefficient ni according to reaction (2). Because the most stable ferrous sulfates are hydrated, we also calculated the enthalpy of formation relative to single cation components using ferrous sulfate heptahydrate: xFeðOHÞ3 þ ð1  x  yÞFeðOHÞ2 þ yFeSO4  7H2 O þ ðn  7yÞ  H2 O ! FeII 1x FeIII x ðOHÞ2þx2y  ðSO4 Þy  nH2 O

ð5Þ

Enthalpies of formation of green rusts from elements and from single cation compounds are given in Tables 6 and 8, respectively. Enthalpies of formation from single cation compounds and Gibbs free energies from single cation compounds are plotted as a function of FeII/FeIII ratio in Fig. 2.

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Fig. 1. Mo¨ssbauer spectra at 15 and 295 K of sample GR-0.50. Sextets are labeled by their hyperfine field, and are shown in gray or as dashed lines. Solid black lines represent doublets. The smallest doublet (o) in the 15 K spectra probably represents a magnetic phase that is ordering at this temperature.

4. DISCUSSION Table 5 Results of Mo¨ssbauer spectroscopy for sample GR-0.50 d (mm/s)

D (mm/s)

BHf (T)

Area (%)

4.1. Chemical composition of green rust

Attributed to

2

15 K, v = 1.47

0.85 0.48 0.42 1.30 0.15 1.04

0.20 0.26 0.07 2.79 0.62 1.76

28.5 49.9 41.6

8 44 8 20 8 2

Fe FeIII FeIII FeII FeIII FeII

295 K, v2 = 0.89

0.10 0.80 0.30 0.39 1.25 1.18

0.46 0.06 0.60 0.90 1.95 3.24

34.1 33.1

11 10 33 21 15 10

FeIII FeII FeIII FeIII FeII FeII

Values of d and D are referenced to Fe foil.

II

Chemical composition of green rust depends on many factors, including final pH of synthesis, initial FeII/FeIII ratio, oxidation rate and duration of the aging in mother solution. The ideal FeII/FeIII ratio for LDHs should be between 2 and 3. As noted by Hansen (2001), the complete chemical composition was seldom described. Any amorphous material present could alter the FeII/FeIII ratio of the green rust. Often the FeII/FeIII ratio was determined only by Mo¨ssbauer spectroscopy. The majority of works report FeII/FeIII ratios of 2:1 or 3:1. However, the total range of FeII/FeIII ratio is reported to be 0.6–4.0 (Cuttler et al., 1990; Cornell and Schwertmann, 1996). Lewis (1997) noted that FeIII content determined by direct chemical analysis does not always agree with values from Mo¨ssbauer spectroscopy. In our work the

Enthalpy of formation of green rusts

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Table 6 Enthalpies of solution and enthalpies of formation from the elementsd Sample ID

Molecular weight

GR-1.34 GR-1.21 GR-0.89 GR-0.50

136.1 137.5 134.7 132.4

a b c d

Enthalpies of solution, DHsol J/g

kJ/mol

334.8a ± 7.7b 349.8 ± 7.5 348.7 ± 13.9 311.4 ± 6.6

45.6 ± 1.1 48.1 ± 1.0 47.0 ± 1.9 41.2 ± 0.8

Enthalpy of formation from elements, DH f , kJ/mol (8)c (5) (5) (6)

1079.6 ± 1.6 1084.4 ± 1.6 1068.6 ± 2.2 1036.8 ± 1.5

Average. Two standard deviations of the mean. Number of measurements. Reported enthalpies of formation from elements are calculated using FeSO4Æ7H2O as a reference phase.

Table 7 Thermodynamic cycle and reference calorimetric data for calculation of enthalpy of formation, DH f , of green rusts Reaction

Enthalpy of reaction, kJ/mol

½FeII 1x FeIII x ðOHÞð2þx2yÞ ðSO4 Þy nH2 Ocr þ½ð2 þ x  2yÞ Hþ aq ¼ ½ð1  xÞ  Fe2þ þ x  Fe3þ þy  SO4 2 þ ðn þ 2 þ x  2yÞ H2 Oaq FeSO4  7H2 Ocr ¼ ½Fe2þ þSO4 2 þ 7H2 Oaq c-FeOOHcr + [3H+]aq = [Fe3+ + 2H2O]aq a-MgSO4cr ¼ ½Mg2þ þ SO4 2 aq MgOcr + [2H+]aq = [Mg2+ + H2O]aq H2Ol,25°C = H2Oaq,25°C Fecr þ O2g þ 12H2g ¼ c-FeOOHÞcr Fecr + 5.5O2g + 7H2g = FeSO4Æ7H2Ocr Mgcr + 2O2g + Scr = a-MgSO4cr Mgcr þ 12O2g ¼ MgOcr 1  2O2g þ H2g ¼ H2 Ol;25 C Fecr þ yScr þ ð12n þ 12x þ y þ 1Þ  O2g þðn þ 12x  y þ 1Þ  H2g ¼ ½FeII 1x FeIII x ðOHÞð2þx2yÞ ðSO4 Þy  nH2 Ocr DH fðGRÞ ¼ DH solðGRÞ þ ð1  xÞ ½DH solðFeSO4 7H2 OÞ þ DH f ðFeSO4  7H2 OÞ þx  ½DH solðFeOOHÞ þ DH fðFeOOHÞ  þ ð1  x  yÞ ½DH solðMgOÞ  DH solðMgSO4 Þ þ DH fðMgOÞ  DH fðMgSO4 Þ  þðn þ 7x  y  6Þ  ½DH dilðH2 OÞ þ DH fðH2 OÞ 

DHsol(GR) DH solðFeSO4 7H2 OÞ ¼ 47:1  0:3y ð4Þz DHsol(cFeOOH) = 46.5* ± 0.2a DHsol(a-MgSO4) = 53.5 ± 0.5b DHsol(MgO) = 149.7 ± 0.6b DHdilution = 0.4a DH fðc-FeOOHÞ ¼ 549:4  1:4c DH fðFeSO4 7H2 OÞ ¼ 3014:0  0:6d DH fðMgSO4 Þ ¼ 1289:0  0:5e DH fðMgOÞ ¼ 601:6  0:3d DH fðH2 OÞ ¼ 285:8  0:1d DH fðGRÞ

Temperature is 25 °C. *Average, two standard deviations of the mean, number of measurements. a Majzlan et al. (2004). b Majzlan et al. (2005). c Majzlan et al. (2003). d Robie and Hemingway (1995). e DeKock (1986).

FeII/FeIII ratio determined by direct chemical analysis is 0.50–1.3. The FeII/FeIII ratio equal to 0.50 was confirmed by Mo¨ssbauer spectroscopy for sample GR-0.50. Inadequate and lacking analyses of interlayer anion and water contents reflect difficulties due to the very reactive nature of green rusts. As a results, content of the anion was often set to maintain charge balance according to idealized composition where FeIII/SO4 = 2 (Lewis, 1997; Refait et al., 1999). Water analysis in Simon et al. (2003) was determined from differential scanning calorimetry, which brings some uncertainty as well because one does not know the energetics and the release temperature of the bound water. Thus, the present work emphasizes complete characterization (sulfate, water content both by Karl-Fisher titration and thermogravimetric

analysis, and ferric and ferrous iron content both by wet chemical methods and Mo¨ssbauer spectroscopy) of samples used for calorimetry. As can be seen from Table 3, FeIII/S ratios for samples GR-1.34, GR-1.21, and GR-0.89 differ from the ideal ratio 2. Since no other sulfate containing phases were identified either by XRD or by FTIR, we assign all found sulfate to the green rust itself. Although, the deviation from the ideal ratios could be affected by the error in the chemical analysis, fluctuations within ±5 wt% for both FeIII and S will not give a substantial deviation from the observed FeIII/S ratios. In the literature, FeIII/S ratio was often reported to be different than 2 (e.g., Refait et al., 1999) or were not analyzed at all. Additionally, water and hydroxyl groups can be

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Table 8 Enthalpies and free energies of formation from single cation compound, DfHscc and DfGscc, given together with the literature data which were used to calculate these values Formula of green rust

FeII 0:57 FeIII 0:43 ðOHÞ1:72 ðSO4 Þ0:36  0:91H2 O FeII 0:55 FeIII 0:45 ðOHÞ1:70 ðSO4 Þ0:38  0:90H2 O FeII 0:47 FeIII 0:53 ðOHÞ1:85 ðSO4 Þ0:34  0:82H2 O FeII 0:33 FeIII 0:67 ðOHÞ2:01 ðSO4 Þ0:33  0:59H2 O

DfHscc, kJ/mol

DfGscc, kJ/mol

Formula of green rust

Relative to FeSO4

Relative to FeSO4Æ7H2O

6.6 ± 1.9 2.3 ± 1.9 2.5 ± 2.4 3.8 ± 1.8

24.0 ± 2.4 29.9 ± 2.4 26.3 ± 2.8 24.2 ± 2.2

FeII 0:67 FeIII 0:33 ðOHÞ2 ðSO4 Þ0:17 a FeII 0:67 FeIII 0:33 ðOHÞ2 ðSO4 Þ0:17 b FeII 0:71 FeIII 0:29 ðOHÞ2 ðSO4 Þ0:14 c

Relative to FeSO4

Relative to FeSO4Æ7H2O

13.9 ± 2.0 14.1 ± 2.3 10.1 ± 3.1

9.5 ± 1.9 9.6 ± 2.2 6.5 ± 3.0

References phase used

DH f ; kJ=mol

DGf ; kJ=mol

H2O Fe(OH)2 Fe(OH)3 FeSO4 FeSO4Æ7H2O

285.8 ± 0.1d 574.1 ± 1.0e 832.6 ± 1.0e 928.8 ± 0.5e 6014.0 ± 0.6d

237.1 ± 0.1d 486 ± 1.0e 711 ± 1.0e 824.1 ± 0.5e 2509.5 ± 1.3e

a b c d e

Refait et al. (1999). Hansen (2001). Detournay et al. (1975). Robie and Hemingway (1995). McBride et al. (2002).

ing, a mixture of fine-grained minerals including green rusts can form (Christiansen and Stipp, 2003). Full oxidation of green rust leads to formation of ferrihydrite, Fe(OH)3[ÆxH2O] (Refait et al., 2003), which in turn can coarsen and transform to iron oxyhydroxides (FeOOHÆxH2O). Thus a general reaction of green rust formation and further transformation can be written as O2

O2

FeðOHÞ2 ! Green Rust ! FeðOHÞ3

Fig. 2. Enthalpies (DfHscc, kJ/mol) and free energies (DfGscc, kJ/mol) of formation from single cation compound as function of FeIII/Fetotal and FeII/FeIII ratios. Note: Values of DfHscc kJ/mol and DfGscc refer to samples with different compositions. [1] Detournay et al. (1975); [2] Refait et al. (1999); [3] Hansen (2001).

substituted by O2 2 (Koch, 1998). Thus, since green rusts represent a complicated group of phases, we believe that the general green rust formula should be written as FeII 1x FeIII x ðOHÞ2þx2y ½Am y=m  nH2 O, where x is FeIII content, y is the content of anion A with the charge m and n is the amount of water. We used these chemical formulas to calculate enthalpies of formation from single cation compounds (reactions (4) and (5)). 4.2. Stability of green rust When the environment in aqueous solutions and soils, containing dissolved FeII, changes from reducing to oxidiz-

ð6Þ

and Fe(OH)2 and Fe(OH)3 can be considered as end-members. Therefore, we considered stability of green rust relative to these compounds and FeSO4. We also compared stability of green rust relative to hydrated iron (II) sulfate, FeSO4Æ7H2O, because it is the most stable ferrous iron sulfate. Enthalpies of formation from single cation compounds according to reaction (2) (see Table 8) show that green rust of a given composition is more stable in enthalpy than a stoichiometric mechanical mixtures of Fe(OH)3, Fe(OH)2, FeSO4 and liquid water. The stabilization by 2 to 7 kJ/ mol is of the same magnitude as that reported for the analogous hydrotalcites (layered double hydroxides with Al as trivalent species) (Allada et al., 2006). Green rust is less stable by 24–30 kJ/mol than mechanical mixture of single cation compounds and hydrated ferrous iron sulfate according to reaction (3). Nevertheless, because green rust is never made from solid iron sulfates, this metastability is not important for synthesis in the laboratory or natural environment. Free energy of formation of green rust from elements was reported in earlier work. The values, recalculated per mole of iron for a better comparison of earlier and current data, are 631.7 ± 1.7 kJ/mol (Refait et al., 1999) and 632 kJ/mol (Hansen, 2001) for the composition FeII 0:67 FeIII 0:33 ðOHÞ2 ðSO4 Þ0:17 , and 608.7 ± 2.8 kJ/mol for the composition FeII 0:71 FeIII 0:29 ðOHÞ2 ðSO4 Þ0:14 .

Enthalpy of formation of green rusts

The enthalpies of formation and Gibbs free energies of formation from single cation compounds according to reactions (4) and (5) for several compounds along the FeII 1x FeIII x ðOHÞ2þx2y ðSO4 Þy  nH2 O join show small differences as a function of FeII/FeIII ratio (Fig. 2 and Table 8). A similar minor thermodynamic dependence as a function of Al content was observed for different hydrotalcites (Allada et al., 2005). It was argued that these differences were not substantial enough to promote significant thermodynamic preference for any given metal ratio (Allada et al., 2006). The general agreement in enthalpies of formations from single cation compounds, DfHscc with free energies of formation from single cation compounds, DfGscc, obtained from different studies, different techniques and at different times strongly support this statement. The associated uncertainties with these free energy values are probably somewhat larger than reported for different reasons. Even though the FeIII/SO4 ratio deviates from ideal and is equal to 4.66 (Refait et al., 1999), the ideal 2:1 ratio was used in calculations of the chemical potential and free energy from constituent reactions (Refait et al., 1999). Although the green rusts formed definitely contain interlayer water, many works calculate the free energy of formation from a series of equilibrium reactions for anhydrous green rust, which also can introduce some uncertainty. Indeed, Allada et al. (2005) argued that the confinement of water in the interlayer with more negative enthalpy and entropy than liquid water is a thermodynamic driving force for hydrotalcite formation. Values of Gibbs free energy were either obtained using measured redox potentials (Refait et al., 1999) or solubility measurements (Detournay et al., 1975; Hansen, 2001), but without performing reversals needed to show equilibrium. Detournay et al. (1975) made an attempt to perform both reactions of formation from Fe(OH)2 and transformation to goethite for the same green rust, but the reported DfG values differ by 40 kJ/mol. This gives a rough idea of the possible uncertainty in determination of free energy of formation using either solubility or redox methods. The errors associated with these values of Gibbs free energy are thus probably larger than reported. Nevertheless, the recalculated free energies of formation from single cation compounds agree with our values of enthalpies of formation from singe cation compounds within a few kJ/mol depending on FeII/FeIII ratio and literature source (Fig. 2 and Table 8). Because the entropy of formation from condensed phases according to reaction (2) is expected to be small, such agreement is encouraging. 4.3. Entropy estimations As discussed above, green rust shows a small stabilization in both enthalpy and Gibbs free energy relative to the mixture of single cation components. Such small values reflect the structural similarities of the mixture of single cation compounds and green rust as argued for hydrotalcites (Allada et al., 2005, 2006). These similarities suggest that the magnetic contributions to entropy of the mechanical mixture and the green rust may also be similar. Thus, we might expect the entropy of formation from single cation

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compounds, DfSscc, would be similar to that found for hydrotalcites (about 5 J/mol K; Allada et al., 2005). Instead, values of DfSscc calculated from our values of DfHscc and literature values of Gibbs free energies of formation, are positive and in the range of 10–25 (±10) J/ mol K. The calculation did not take into account slight compositional differences between samples (FeII/FeIII ratio, water and sulfate contents as well as crystallinity and disorder) from this work and from the literature that would increase the uncertainty of these estimations. Positive values of entropies may also reflect the differences between magnetic transitions between green rusts and the single cation compounds. They might also suggest that interlayer water is less tightly bound in green rust than in hydrotalcites. 5. CONCLUSIONS Green rust, belonging to the group of hydrotalcite-like compounds or layered double hydroxides, with sulfate anion in the interlayer was investigated by direct calorimetric measurements of the enthalpy of formation of samples with various FeII/FeIII ratios. Prior to calorimetry, each sample was fully characterized by XRD, FTIR, TGA, chemical analysis, and Mo¨ssbauer spectroscopy. Iron content determined by chemical analysis and by Mo¨ssbauer spectroscopy agreed. Standard enthalpies of formation are 1036 to 1079 kJ/mol depending on the composition. Comparison with literature standard Gibbs free energies of formation are not possible because Gibbs free energies are reported for the anhydrous phases. Nevertheless, when one compare enthalpies and Gibbs free energies of formation from single cation compounds, namely, Fe(OH)2, Fe(OH)3, FeSO4 and H2O they agree within a few kJ/mol. Small differences in the enthalpy of formation from single cation compounds as a function of FeII/FeIII ratio indicate that there is no significant thermodynamic preference for any given ratio. Estimated entropy of formation from single cation compounds is 10–25 (±10) J/mol K. The trends in enthalpies of formation from single cation compounds are similar for green rusts and aluminum-bearing hydrotalcites but the positive entropies of formation of green rusts appear in contrast to the slightly negative values for hydrotalcites. ACKNOWLEDGMENTS Financial support from DOE Grant DEFG03-97ER14749 is gratefully acknowledged. John Husler from the Department of Geology at the University of New Mexico is thanked for the iron analysis, and Elizabeth Sklute for assistance with Mo¨ssbauer data processing. J. Majzlan and J.M. Neil are gratefully acknowledged for useful discussions. Three anonymous reviewers and the editor are thanked for their useful comments and critique.

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