# Nanoanalytical Identification of Siderite Dissolution-Coupled Pb Removal Mechanisms from Oxic and Anoxic Aqueous Solutions

Lisa Füllenbach, Jeffrey Paulo H. Perez, Helen Freeman, Andrew Thomas, Sathish Mayanna, Julia Parker, Jörg Göttlicher, Ralph Steininger, Jörg Radnik, Liane Benning, et al.

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Lisa Füllenbach, Jeffrey Paulo H. Perez, Helen Freeman, Andrew Thomas, Sathish Mayanna, et al. Nanoanalytical Identification of Siderite Dissolution-Coupled Pb Removal Mechanisms from Oxic and Anoxic Aqueous Solutions. ACS Earth and Space Chemistry, 2020, 4 (11), pp.1966-1977. 10.1021/ac-searthspacechem.0c00180. hal-04903548

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Füllenbach, L. C., Perez, J., Freeman, H., Thomas, A. N., Mayanna, S., Parker, J. E., Göttlicher, J., Steininger, R., Radnik, J., Benning, L. G., Oelkers, E. H. (2020): Nanoanalytical Identification of Siderite Dissolution-Coupled Pb Removal Mechanisms from Oxic and Anoxic Aqueous Solutions. - ACS Earth Space Chemistry, 4, 11, 1966-1977.

https://doi.org/10.1021/acsearthspacechem.0c00180
Nanoanalytical identification of siderite dissolution
coupled Pb removal mechanisms from oxic and
anoxic aqueous solutions

Lisa C. Füllenbach*<sup>a</sup>, Jeffrey Paulo H. Perez<sup>b</sup>, Helen M. Freeman<sup>b,c</sup>, Andrew N. Thomas<sup>d</sup>, Sathish Mayanna<sup>b,e</sup>, Julia E. Parker<sup>f</sup>, Jörg Göttlicher<sup>g</sup>, Ralph Steininger<sup>g</sup>, Jörg Radnik<sup>h</sup>, Liane G.

Benning<sup>b,i</sup>, Eric H. Oelkers*<sup>a,j</sup>

<sup>a</sup>Department of Earth Sciences, University College London, London, WC1E 6BT, United Kingdom. *l.fuellenbach@ucl.ac.uk; *e.oelkers@ucl.ac.uk

<sup>b</sup>GFZ German Research Center for Geosciences, Telegrafenberg, 14473 Potsdam, Germany

<sup>c</sup>School of Chemical and Process Engineering, University of Leeds, Leeds, LS2 9JT, United Kingdom

<sup>d</sup>Department of Applied Geosciences, Karlsruhe Institute of Technology, 76131 Karlsruhe, Germany

<sup>e</sup>Carl Zeiss Microscopy GmbH, 73447 Oberkochen, Germany

<sup>f</sup>Diamond Light Source, Harwell Science and Innovation Campus, Didcot, OX11 0DE, United Kingdom

<sup>g</sup>Karlsruhe Institute of Technology, Institute for Photon Science and Synchrotron Radiation, 76344 Eggenstein-Leopoldshafen, Germany

<sup>h</sup>Federal Institute for Material Research and Testing, 12235 Berlin, Germany

<sup>i</sup>Department of Earth Sciences, Freie Universität Berlin, 12249 Berlin, Germany

<sup>j</sup>GET, CNRS UMR 5563, 14 Avenue Edouard Belin, 3400 Toulouse, France
1 GRAPHICAL ABSTRACT. Please find the graphic at the end of this manuscript.

2 ABSTRACT.

3 Lead(II) is a toxic pollutant often found in metal-contaminated soils and wastewaters. In acidic
4 aqueous environments, Pb(II) is highly mobile. Chemical treatment strategies of such systems
5 therefore often include neutralization agents and metal sorbents. Since metal solubility and the
6 retention potential of sorbents depend on the redox state of the aqueous system, we tested the
7 efficiency of the naturally occurring redox-sensitive ferrous iron carbonate mineral siderite to
8 remove Pb(II) from acidic aqueous solutions in batch experiments under oxic and anoxic
9 conditions over a total of 1008 h. Siderite dissolution led to an increase of reactive solution pH
10 from 3 to 5.3 and 6.9, while 90% and 100% of the initial aqueous Pb(II) (0.48 x10⁻³ mol kg⁻¹) were
11 removed from the oxic and anoxic system, respectively. Scanning and transmission electron
12 microscopy (SEM, TEM), combined with X-ray absorption and photoelectron spectroscopy (XAS,
13 XPS) indicated that, under oxic conditions, Pb(II) was consumed by cerussite precipitation and
14 inner-sphere surface complexation to secondary goethite. Under anoxic conditions, Pb(II) was
15 removed by the rapid precipitation of cerussite. This efficient siderite dissolution-coupled
16 sequestration of Pb(II) into more stable solid phases demonstrates this potential method for
17 contaminated water treatment regardless of the redox environment.

18 KEYWORDS. Siderite – XANES-mapping – EXAFS – Lead carbonate – Iron oxyhydroxides –
19 Adsorption – Wastewater treatment

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# 1. INTRODUCTION

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Human use of land and natural resources often leads to surface and groundwater pollution. Divalent lead (Pb(II)) is one of the most toxic and common contaminants introduced to soils and (sub)surface waterbodies from mine tailings¹,², waste disposal and landfills³⁻⁵, and industrial activities⁶, where concentrations can reach up to several thousand mg of Pb(II) per kg of contaminated soils or sediments⁷⁻⁹. Studies of lead removal from aqueous solutions have been mainly focused on chemical sorption under oxic and controlled pH conditions to various substrates including aluminum oxides¹⁰,¹¹, clays¹², soils¹³,¹⁴, calcite¹⁵⁻¹⁷, and ferric iron (oxyhydr)oxides such as goethite¹⁸⁻²⁰. Contaminated environments such as stream sediments²¹, floodplain soils⁷,⁸ and tailing dam sites⁹, which are in consistent contact with water are highly susceptible to redox changes resulting from periodic (e.g., tidal) or seasonal (e.g., heavy rainfall) fluctuations. As the redox state of a polluted aqueous system directly affects the pH and determines the solubility of many metals and sorbents²¹⁻²³, it is necessary to understand the redox state of the system to predict accurately both metal retention and release. While acidity of wastewaters, as in acid mine drainage²⁴, is commonly overcome by the addition of alkaline, often carbonate materials²⁵,²⁶, the issue of potential redox fluctuations causing metal release in contaminated waterbodies remains a challenge.

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The redox sensitive ferrous carbonate siderite (FeCO₃) is a potential candidate material for redox-dependent metal remediation. Several studies have investigated siderite as a reductive sorbent for redox sensitive metal(loid)s such as Cr²⁷⁻²⁹, As³⁰,³¹, and Hg³². Under controlled pH and redox conditions these metal(loid)s were reduced and adsorbed to either the siderite surface or, due to coupled redox reactions, to secondary ferric (oxyhydr)oxides. Little attention has been drawn to illuminate the complete reaction pathway between siderite and contaminant over time,
where siderite can act as both dissolution-induced neutralizer and sorbent under either redox condition. Siderite is highly soluble in acidic aqueous solutions³³, and its dissolution leads to aqueous phase neutralization through bicarbonate production. Under oxic conditions, the release and oxidation of Fe(II) from dissolving siderite results in the precipitation of Fe(III) (oxyhydr)oxides³⁴. During dissolution under reduced conditions, iron oxidation is suppressed, and aqueous metals can interact with the siderite surfaces and its dissolution products.

To test whether siderite reactivity in acidic aqueous systems is also sufficient to remove less redox-sensitive metals such as Pb in either redox environment over time, and to better understand the mechanisms involved, we have monitored natural siderite dissolution, pH and metal concentrations in initially acidic Pb(II)-bearing aqueous solutions under oxic and anoxic conditions. We used scanning and transmission electron microscopy to investigate the siderite-precipitate interface and the reaction products. Complimentary X-ray absorption and photoelectron spectroscopic analyses, including high-resolution elemental X-ray fluorescence and absorption spectroscopic mapping, allowed localization and determination of Fe(II/III) and Pb(II) speciation and bonding environment. Data acquired from this set of analyses contributes to the evaluation of the metal retention efficiency of siderite as a potential low-cost in situ remediation material.

## 2. MATERIAL AND METHODS

### 2.1. Materials

Natural macrocrystalline siderite from Peyrebrune Quarry, France³³, was ground in an agate mortar and sieved to a < 63 µm size fraction. The siderite powder was cleaned by repeated rinsing with ethanol after ultrasonication; its surface area was 9.95 m² g⁻¹, as determined by BET
multipoint Krypton adsorption (Quantachrome Autosorb-1MP). Powder X-ray diffraction (XRD; PANalytical X'Pert) confirmed the purity of the cleaned siderite powder (> 98%) with minor amounts of fluorite, quartz and traces of calcite and hematite impurities. Inductively coupled plasma optical emission spectroscopy (ICP-OES; Varian 720ES) of the siderite powder revealed contents of 2.43 wt% Mn, 1.27 wt% Ca, and 0.34 wt% Mg, which are likely structural impurities in siderite and minor calcite. Minor and trace contents of 0.74 wt% S, 0.13 wt% K, 0.02 wt% Na, 0.02 wt% Si, 37.3 ppm Cu, and 24.5 ppm Cr also detected in the ground material are likely contributions from mineral impurities (e.g., quartz, trace sulfides).

## 2.2. Batch Experiments

All chemicals use in the experiments were reagent or analytical grade from Honeywell Fluka and VWR and used as received. All aqueous stock solutions were prepared using ultrapure deionized water (Milli-Q, resistivity ~18 MΩ·cm). The aqueous solutions used in the oxic experiments were prepared under ambient laboratory conditions, whereas those used in the anoxic experiments were prepared inside an CO₂-free anaerobic chamber (97% N₂, 3% H₂; COY Laboratory Products Inc.) using de-oxygenated deionized water that was sparged with argon at ~90 °C for at least 4 h prior to the experiments. For both the oxic and anoxic experiments, the pH of all starting solutions was adjusted to 3.0 ±0.1 by adding 15.5 M HNO₃ and 2.5 M NaOH, and the ionic strength was adjusted to 2-5 x10⁻³ mol kg⁻¹ by adding 0.1 M NaNO₃. For all lead-bearing solutions, 0.48 x10⁻³ mol kg⁻¹ lead(II) nitrate (Pb(NO₃)₂) was dissolved into deionized water prior to ionic strength and pH adjustment. The concentrations of lead used in the experiments are based on average concentrations reported in contaminated environments⁶⁻⁹.
All oxic and anoxic experiments were performed in duplicate in batch reactors that were continuously mixed on a rotary mixer. A list of all experimental details is provided in Table S1. For each experiment, based on preliminary experiments (not shown) and studies using siderite as metal sorbent<sup>27,29,30</sup>, 6 g L<sup>-1</sup> siderite was reacted with the oxic and anoxic acidic Pb(II)-bearing aqueous solutions. Control experiments in the absence of Pb were also conducted following the same procedures (Figure S1). To allow the analysis of the fluid and the solid phases at various stages of the siderite dissolution reaction, a series of 15 distinct single short-term (ST) batch experiments were performed and ended at different times. The experiments in these series ran for 0.5, 1, 2, 3, 6, 12, 20, 28, 36, 48, 60, 72, 96, 144, and 240 h. Additional long-term (LT) experiments were run for a total of 1008 h; during these experiments, fluid aliquots were collected every 168 h. Sampling of liquid aliquots from the oxic experiment was conducted within minutes to avoid equilibration of the reactive fluid with atmospheric CO<sub>2</sub>. Final solids and fluids from all experiments were separated for analysis by vacuum filtration using 0.2-µm nylon filters. To avoid potential iron oxidation and to ensure all iron remained dissolved in the filtrates, aliquots were immediately acidified after. The collected solid residues were kept dry or in ethanol at ambient conditions, and in case of the anoxic samples in the anaerobic chamber, to prevent further phase transformation.

The elemental concentrations of all 214 collected fluid samples were measured using ICP-OES (Varian 720ES / Agilent Technologies 5110); the average relative standard deviation (RSD) of the analyses ranged from 0.3 to 1.2% (Table S2). Final liquid samples with Pb concentrations below the detection limit of the ICP-OES were remeasured using the more sensitive inductively coupled plasma mass spectrometer (ICP-MS; Bruker 90); analyses were within an accuracy of 0.57% RSD. To determine the oxidation of dissolved Fe<sup>2+</sup> to Fe<sup>3+</sup> in the reacted aqueous solutions, Fe speciation
110 was measured by the colorimetric Ferrozine method³⁵ using a segmented flow analysis device
111 (SFA; SEAL AutoAnalyzer 3HR) equipped with a digital colorimeter. The average RSD of these
112 measurements ranged between 0.34 and 0.35%. Comparison of the total iron concentrations
113 determined using SFA ($Fe_{SFA}^{total}$) to those obtained by ICP-OES ($Fe_{OES}^{total}$) revealed that $Fe_{SFA}^{total} \approx Fe_{OES}^{total}$,
114 whereupon $Fe_{OES}^{total}$ is a close approximation of the $Fe^{2+}$ released during siderite dissolution (Figure
115 S2).

116 **2.3. Solid phase characterization**

117 Initial and reacted siderite and selected secondary solid reaction products were imaged and
118 analyzed using a ZEISS Ultra Plus Gemini field-emission scanning electron microscope (FE-
119 SEM) equipped with an In-lens secondary electron and UltraDry SDD energy dispersive X-ray
120 (EDX) detectors. Solids collected from anoxic experiments were mounted onto SEM stubs inside
121 the anaerobic chamber and transported to the SEM facility using an anaerobic jar to avoid sample
122 oxidation.

123 Bulk mineral structure analyses were performed with a high resolution PANalytical X'pert PRO
124 MPD diffractometer equipped with a Co anode X-ray tube ($\lambda$ = 1.789 Å), a Ge incident beam
125 monochromator and a X'celerator detector. Repeated measurements of anoxic samples stored
126 under dry ambient laboratory conditions for 19 h and 5 weeks verified the long-term stability of
127 siderite and secondary phases.

128 High resolution morphological and structural analyses were performed with a FEI Tecnai G2
129 F20 X-Twin transmission electron microscope (TEM) operated at 200 kV with a Schottky field
130 emitter as an electron source. The TEM is equipped with a Gatan Imaging Filter (Tridiem™) and
131 a Fishione high-angle angular dark-field detector (HAADF) and an EDAX X-ray analyzer for
simultaneous elemental analysis. Combined fast Fourier transform (FFT) was used to determine
local structure and morphology of the secondary particles.

X-ray absorption near edge spectroscopy (XANES) combined with X-ray fluorescence (XRF) mapping at beamline I14 of the Diamond Light Source (UK) was used on the oxic sample to determine the elemental distribution of Fe and Pb and to identify the speciation and oxidation state of Fe and Pb in the metal-binding phases. The Python-based software ATHENA³⁶ and MANTIS were used for data processing. Complimentary extended X-ray absorption fine structure (EXAFS) spectra were collected at the ANKA-KIT synchrotron (Germany) SUL-X beamline to investigate the local bonding environment of Pb in both oxic and anoxic samples. Data reduction was performed using ARTEMIS software³⁶. Details on the fitting procedures of XANES and EXAFS data are given in the Supplementary Information. Pure mineral model compounds were used to fit the XAS data; Pb(II)-adsorbed ferrihydrite was used as representative for Pb(II) adsorption to ferric oxyhydroxide surfaces. The high carbonate solubility and reactivity of siderite, impeded the preparation of Pb(II)-adsorbed siderite as model.

The surface chemistry of the anoxic samples was investigated by X-ray photoelectron spectroscopy (XPS) to determine Pb speciation on the siderite surfaces using a KRATOS Axis Ultra DLD. Data reduction was performed using UNIFIT software³⁷.

## 2.4. Geochemical Modelling

Saturation indices of the phases involved in this study were calculated based on measured aqueous solution chemistry data using the software PHREEQC³⁸ together with its 2017 Lawrence Livermore National Library database.
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## 3. RESULTS

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### 3.1. Pb(II) uptake during siderite dissolution

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In Pb(II)-bearing oxic aqueous solutions, the pH reached 4.5 within 2-3 h of reaction and continued to increase steadily to a final pH of 5.3 (Figure 1A). Simultaneously, Fe<sup>total</sup> concentrations increased to 0.14 x10<sup>-3</sup> mol kg<sup>-1</sup> during the first 2-4 h of reaction. A subsequent distinct drop in aqueous Fe<sup>total</sup> concentrations and a visible change in the color of the suspension to reddish ochre indicated the precipitation of ferric (oxyhydr)oxides; the onset of ferrous iron oxidation and ferric oxyhydroxide precipitation may have initiated as soon as pH 4 was reached<sup>39</sup>.

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Within 2 h after the onset of siderite dissolution, 9 ±3.6% of the initial aqueous Pb(II) was removed from the aqueous solution; after 1008 h 90% of the Pb(II) was removed (Figure 1C). Under anoxic conditions, the pH of the Pb(II)-bearing aqueous solution increased to 5.1 within 0.5 h and to 6.5 after 96 h of reaction; from this time onwards, the pH increased steadily to 6.9 at the end of the experiment (Figure 1B). The Fe<sup>total</sup> concentrations peaked at 0.36 x10<sup>-3</sup> mol kg<sup>-1</sup> after 168 h of reaction. Notably, Pb(II) was completely removed from the aqueous solution within 96-144 h of reaction (Figure 1D). The final aqueous Pb(II) concentrations were < 10 µg L<sup>-1</sup>, the upper limit of Pb(II) allowed in drinking water by the EU ECD<sup>40</sup>, which demonstrates the Pb(II) removal efficiency of this method.

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<figure><img src="image_2.png" /><figcaption>**Figure 1.** Total metal concentrations (black symbols) and pH (grey symbols) shown as a function of time. Fe<sup>total</sup> and pH during siderite dissolution in the presence of Pb(II) under oxic (A; diamonds) and anoxic conditions (B; triangles); corresponding Pb(II) concentrations are given in C and D, respectively (circles = control solutions). Error bars are smaller than the symbols. Standard error of the pH probe is ±0.01, RSD% (ICP-OES) is 0.29 to 1.23%. Shown are data from all Pb(II)-bearing experiments (Table S1). For clarity, Fe<sup>total</sup> concentrations and pH measured during the first 24 h of the oxic experiment are shown in the inset in A.</figcaption></figure>

**3.2. Characterization of the reaction products**

The solids recovered from the oxic and anoxic Pb(II)-bearing aqueous solutions after 1008 h were characterized using SEM imaging combined with EDX analyses, TEM and XRD. An overview of the XRD results is given in Figure S3 and Table S3. Measured *d*-spacings from TEM images and their corresponding FFT patterns are compiled in Table S4.
<figure><img src="image_3.png" /><figcaption>**Figure 2.** Scanning and transmission electron micrographs of siderite reacted in Pb(II)-bearing solutions for 1008 h under oxic conditions (PbO LT). **A** Rod-shaped and intercalated pseudo-hexagonal tabular precipitates covering a siderite surface. **B** Ultra-fine precipitates collected from the reactor suspension: a nearly completely reacted siderite fragment overgrown and surrounded by nanoparticulate precipitates forming clumps (**C**) and rods (**D**) via (nearly) oriented attachment; respective FFT patterns of the nanoparticles are indicative of goethite (2.25-2.26 Å, 2.56 Å), traces of hematite (3.87 Å) and residual siderite (2.40 Å, 2.79 Å). See Table S4 for measured *d*-spacings and reference data.</figcaption></figure>

After 1008 h of reaction under oxic conditions, the recovered siderite grains were heterogeneously covered with secondary phases identified as nanoparticulate goethite (α-FeOOH) (Figure 2). Minor amounts of hematite (α-Fe₂O₃) were also present in the unreacted siderite starting material and are thus not considered a reaction product. The precipitates formed a discontinuous layer of ≤ 100-200 nm rods and needles randomly oriented to the siderite surface, partially radially intergrown and occasionally intercalated with tabular crystallites (Figure 2A).
197 Ultra-fine precipitates collected from the reacted suspension consisted of ca. 5 nm long
198 nanoparticles forming aggregates of ≤ 100 nm clumps with 100-200 nm long rods growing out of
199 the clumps (Figure 2B). The lattice spacings in the clumped nanoparticles measured 2.25 Å (Figure
200 2C), and 2.56 Å in the rods (Figure 2D), identifying both precipitates as goethite. Faint spots in
201 the FFT patterns (insets Figure 2C and 2D) from the particle aggregates, exhibited lattice spacings
202 of 2.40 and 2.79 Å, which suggest the presence of residual siderite, and of 3.87 Å which suggests
203 traces of hematite.

204 Siderite reacted with aqueous Pb(II) under anoxic conditions resulted in the precipitation of ≤ 5
205 µm rounded particles and micron-scale aggregates on the smooth but etched siderite surfaces
206 (Figure 3A). EDX analyses of the reacted siderite grains, their surfaces and the precipitates showed
207 enrichment of Pb (Figure 3B; Figure S4). XRD analyses identified nanocrystalline cerussite
208 (PbCO3) as the secondary phase (Figure S3). This was confirmed by TEM imaging and FFT
209 analyses of the polycrystalline material attached to the siderite surfaces exhibiting d-spacings of
210 2.17 Å, 2.63 Å and 2.92 Å, which are characteristic for cerussite (Figure 3C), while lattice spacings
211 of 1.53 Å likely resulted from the presence of nearby siderite. No residual hematite from the
212 unreacted material was found in the reacted solids, suggesting Fe(II)-induced reductive dissolution
213 of this ferric component41-43. Notably, within the top 20 nm of a reacted siderite, a reaction zone
214 was observed exhibiting a poorly (poly)crystalline phase oriented in a seemingly random angle to
215 the underlying siderite. The lattice spacings of this phase matched that of the cerussite (Figure 3D-
216 E).

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<figure><img src="image_4.png" /><figcaption>**Figure 3.** Scanning and transmission electron micrographs of siderite reacted in Pb(II)-bearing aqueous solutions for 218 1008 h under anoxic conditions (Pb A LT). **A** Reacted siderite grains showing dissolution-features, micron-scale precipitates and Pb-enrichment on the surface (**B**). **C** TEM and the respective FFT pattern identify the precipitates as cerussite (2.63 Å, 2.92 Å) attached to siderite (1.53 Å, 2.17 Å). **D** Structurally independent cerussite (2.63-2.65 Å) growth on the reacted siderite surface (2.44 Å). **E** Widening of the siderite crystal lattice (2.83 Å) near the reacted surface to *d*-spacings indicative for cerussite (2.90 Å). See Table S4 for measured *d*-spacings and reference data.</figcaption></figure>

**3.3. Speciation and local bonding environment of Pb**

A set of complimentary XAS and XPS analyses were performed to determine the bonding environment of Pb associated with the solids recovered from the oxic and anoxic experiments. Results are presented in Figures 4, 5, and S5; respective fitting parameters are given in Tables 1, 2, and S5.
230 **3.3.1. Pb speciation and bonding in the oxic system**

231 Linear combination fitting (LCF) results of the Fe K-edge XANES spectra obtained from the
232 bulk solid sample of the oxic experiment indicated that ca. 5% of the reacted solid-bound iron
233 Fe(II) oxidized to Fe(III) during siderite dissolution in the presence of Pb(II) (Table 1). XRF
234 elemental mapping of individual reacted siderite grains showed that Pb was present on all analyzed
235 siderite surfaces, however, relatively higher concentrations were occasionally observed along the
236 edges, while lower Pb concentrations were detected on (fracture) planes (Figure 4D-G). High-
237 resolution XANES mapping of selected areas on grain surfaces showed that in areas of increased
238 Pb concentrations, 83-85% of Fe was present as Fe(III). In areas with lower Pb concentrations,
239 79% of solid-bound Fe had remained as Fe(II) (Figure 4C). This Fe(II)-Fe(III) distribution on the
240 reacted siderite surfaces indicates heterogeneous surface oxidation and/or covering with secondary
241 ferric precipitates. Notably, LCF results of the Pb L<sub>III</sub>-edge spectra of the same mapping area did
242 not indicate a spatial correlation with or preference of Pb for either iron oxidation state. In contrast,
243 39% of the solid-associated Pb was oxide-bound, and 61% was present as Pb carbonate (Figure
244 4F, area 1). In a mapped area of another grain of reacted siderite, the XRF element map indicated
245 an equal distribution of Pb and a predominance of Fe(III) in the XANES map (Figure 4G, area 2).
246 The LCF of the Pb L<sub>III</sub>-near edge spectrum of this area, however, indicated a higher contribution
247 of oxide-bound Pb over Pb carbonate (Figure 4B, area 2).
<figure><img src="image_5.png" /><figcaption>**Figure 4.** XANES spectra of model compounds and reacted siderite in the presence of aqueous Pb(II) under oxic and anoxic conditions after 1008 h of reaction. **A** Fe K-edge of samples and standards. **B** Pb L<sub>III</sub>-edge of samples and standards. Pb A LT = anoxic long-term experiment, Pb O LT = oxic long-term experiment; sid = siderite, goe = goethite, hem = hematite, PbO = massicot (orthorhombic); Pb-Hfo = Pb-adsorbed hydrous ferric oxide (ferrihydrite), cer = cerussite, hyc = hydrocerussite. Black lines represent bulk sample analyses. Dotted lines represent linear combination fits; colored lines correspond to high-resolution XANES mapping areas (C) in **F** (area 1; 6.5 x 5.5 µm) and **G** (area 2; 1.5 x 1.5 µm) mapped at a 100 x 100 nm pixel size. Note that in all XANES spectra are mixtures of Fe and Pb bonding environments, respectively (see Table 1); the dominant species are indicated in **C**. **D-G** High-resolution X-ray fluorescence maps of reacted siderite grains reacted in Pb(II)-bearing aqueous solution under oxic conditions (Pb O LT); white outlines indicate magnified areas shown in **E-G**. Red = Pb, blue = Fe, magenta = Fe + Pb. Note that the Pb distribution varies only slightly and locally; strongly blue (Fe-rich) areas possibly indicate precipitate-free surfaces or fracture planes (white arrows). Image dimensions: **D** = 165 x 145 µm, **E** = 45 x 38 µm, **F** = 18 x 13 µm, **G** = 2.5 x 2.5 µm.</figcaption></figure>
<table><caption>**Table 1.** XANES Fe K-edge and Pb L<sub>III</sub>-edge linear combination fitting parameters and speciation of solids collected from both oxic<sup>a</sup> and anoxic<sup>b</sup> long-term experiments measured in bulk and at high resolution mapping. Area 1 measured 6.5 x 5.5 µm (Figure 4F), area 2 covered 1.5 x 1.5 µm (see Figure 4G).</caption><thead><tr><th rowspan="2">Edge</th><th rowspan="2">Sample</th><th rowspan="2">χ<sup>2</sup></th><th rowspan="2">χ<sup>2</sup><sub>reduced</sub></th><th colspan="3">Standard material*</th><th rowspan="2">Error<br/>±(%)</th></tr><tr><th>Pb-adsorbed</th><th>Oxide</th><th>Carbonate</th></tr></thead><tbody><tr><td rowspan="4"><b>Fe K-edge</b><br/>Pb O LTa</td><td>bulk</td><td>0.04</td><td>0.0006</td><td></td><td><b>Fe<sub>2</sub>O<sub>3</sub> (Fe<sup>3+</sup>)</b></td><td><b>FeCO<sub>3</sub> (Fe<sup>2+</sup>)</b></td><td>3.1</td></tr><tr><td>area 1<sup>†</sup></td><td>0.93</td><td>0.0095</td><td></td><td>4.9</td><td>95.1</td><td>9.0</td></tr><tr><td>area 1<sup>†</sup></td><td>0.07</td><td>0.0013</td><td></td><td>83.1</td><td>16.9</td><td>4.5</td></tr><tr><td>area 2<sup>‡</sup></td><td>0.04</td><td>0.0007</td><td></td><td>20.8</td><td>79.2</td><td>3.4</td></tr><tr><td>Pb A LT<sup>b</sup></td><td>bulk</td><td>0.27</td><td>0.0040</td><td></td><td>–</td><td>100.00</td><td>0.0</td></tr><tr><td rowspan="4"><b>Pb L<sub>III</sub>-edge</b><br/>Pb O LT</td><td>bulk</td><td>0.01</td><td>0.0001</td><td><b>Pb-Hfo<sup>c</sup></b></td><td><b>PbO</b></td><td><b>PbCO<sub>3</sub></b></td><td>2.0</td></tr><tr><td>area 1<sup>†</sup></td><td>0.03</td><td>0.0008</td><td>–<sup>d</sup></td><td>25.4<sup>d</sup></td><td>18.7</td><td>5.2</td></tr><tr><td>area 2<sup>‡</sup></td><td>0.06</td><td>0.0012</td><td>19.0</td><td>38.6</td><td>61.4</td><td>1.8</td></tr><tr><td>Pb A LT</td><td>bulk</td><td>0.02</td><td>0.0001</td><td>–</td><td>–</td><td>100.0</td><td>3.6</td></tr></tbody></table>

*Fit results normalized to 100%; †see Figure 4C and F; ‡see Figure 4C and G;

<sup>a</sup> Oxic, <sup>b</sup> anoxic long-term experiment; see Table S1 for experimental specifications.

<sup>c</sup> Hfo = hydrous ferric oxide (ferrihydrite). <sup>d</sup> Fits using only Pb-Hfo and no PbO resulted in poorer fit statistics.

Shell-by-shell fits of the Pb L<sub>III</sub>-edge EXAFS region of the solids indicate a first shell Pb–O interatomic distance *R* of *R*<sub>Pb-O</sub> = 2.31 Å; this distance is indicative of Pb(II) adsorbed to iron (oxyhydr)oxides (Table 2). Previous studies reported radial distances of *R*<sub>Pb-O</sub> ≤ 2.35 Å for a distorted trigonal pyramidal coordination of Pb with the surface oxygen of goethite<sup>44-47</sup>. A second single scattering path of *R*<sub>Pb-Fe</sub> = 3.31 Å falls within the reported range of *R*<sub>Pb-Fe</sub> distances of 2.91-3.49 Å, indicative of the bidentate mononuclear inner-sphere surface complexation with FeO<sub>6</sub>-octahedra in the iron (oxyhydr)oxide surface<sup>44</sup>. The coordination number of this pathway is similar (0.55) to values observed in previous EXAFS studies of Pb adsorbed on Fe(III) oxyhydroxide surfaces<sup>20</sup> and lies within the range of values expected for a bidentate mononuclear surface complex (≤ 1.0). In contrast to the clear indication of Pb-carbonate environments in the XANES
spectra, the EXAFS data lack unambiguous evidence of carbon scattering. Reasons for this may be a combination of methodologic limitations and sample specifics. Since carbon is a weak backscattering atom and lead is a strong absorber, the contribution of the Pb–C scattering with R<sub>Pb–C</sub> ≥ 3.5 Å to the EXAFS spectrum<sup>20</sup> may be insignificant in poorly crystalline materials such as the reacted siderite surface and Pb surface complexes<sup>48</sup>. The wide but discontinuous coverage of siderite grain surfaces with precipitates observed in SEM imaging, and the continuous presence of Pb-carbonate species indicated in XANES mapping (Figure 4), suggest that Pb is also associated with carbonate at the siderite surface. It seems likely, therefore, that the observed contributions of Pb–C scatterings to the EXAFS spectrum of the bulk sample were overshadowed by the more prominent Pb–Fe scattering of Pb.

<table><caption>**Table 2.** EXAFS Pb L<sub>III</sub>-edge *k*<sup>3</sup>-weighed shell fitting parameters and interatomic path lengths measured on bulk solid samples from both oxic<sup>a</sup> and anoxic<sup>b</sup> long-term experiments. Uncertainties are given in brackets.</caption><thead><tr><th></th><th>ΔE<sub>0</sub> (eV)</th><th>R-range (Å)</th><th>Shell</th><th>N</th><th>R (Å)</th><th>σ<sup>2</sup> (Å<sup>2</sup>)</th></tr></thead><tbody><tr><td rowspan="2">Pb O LT<sup>a</sup></td><td rowspan="2">-6.7 (2.1)</td><td rowspan="2">1.2 – 4.0</td><td>Pb–O</td><td>2.24 (0.17)</td><td>2.31 (0.03)</td><td>0.01*</td></tr><tr><td>Pb–Fe</td><td>0.55 (0.25)</td><td>3.31 (0.02)</td><td>0.01*</td></tr><tr><td rowspan="3">Pb A LT<sup>b</sup></td><td rowspan="3">-0.05 (3.5)</td><td rowspan="3">1.5 – 5.0</td><td>Pb–O</td><td>8.4 (2.91)</td><td>2.65 (-0.02)</td><td>0.027 (0.006)</td></tr><tr><td>Pb–C</td><td>3*</td><td>3.15 (0.07)</td><td>0.027 (0.006)</td></tr><tr><td>Pb–Pb</td><td>6*</td><td>4.15 (-0.07)</td><td>0.014 (0.004)</td></tr></tbody></table>

* Parameter was set during fitting.

ΔE<sub>0</sub> = energy shift. N = coordination number (degeneracy). R = interatomic path lengths. σ<sup>2</sup> = Debye-Waller disorder term. The amplitude reduction factor (S<sub>0</sub><sup>2</sup>) was set to 0.9 in all fits.

### 3.3.2. Pb speciation and bonding in the anoxic system

XANES Fe K-edge spectra of the bulk solids collected from the anoxic experiments confirmed that Fe(II) oxidation did not occur and all solid-bound Fe(II) remained as carbonate (Figure 4A). LCF of the respective Pb L<sub>III</sub>-edge spectra indicate that 100 ±3.6% of the solid-bound Pb(II) was
present as carbonate (Table 1). In the EXAFS Pb L<sub>III</sub>-edge shell fitting, the spectrum could be closely fit to cerussite with physically reasonable deviations from the ideal structure indicating that Pb is present in a 9-fold coordination<sup>49</sup> (Table 2).

Additional XPS analyses of the anoxic sample (Table S5) showed Pb 4f<sub>7/2</sub> and 4f<sub>5/2</sub> peaks at binding energies of 138.6 and 143.5 eV, respectively. These agree well with the binding energies reported for Pb(II) in Pb-carbonates<sup>50,51</sup>, Pb-calcite<sup>15</sup> and Pb-aragonite<sup>52</sup> (Figure 5). Shifts towards higher binding energies by 0.3 and 0.4 eV observed in the O 1s and C 1s spectra of the reacted sample indicative for MeCO<sub>3</sub><sup>53,54</sup> suggest a change from the calcite-type into the aragonite-type carbonate structure<sup>52</sup>.

<figure><img src="image_6.png" /></figure>

**Figure 5.** XPS spectra of the (A) O 1s, (B) C 1s and (C) Pb 4f photoelectron lines obtained from the siderite sample reacted with aqueous Pb(II) for 1008 h under anoxic conditions (Pb A LT) shown as solid black curves; spectra obtained from the unreacted siderite sample are shown for comparison in dashed curves. Grey curves show the best fits of the spectra of the reacted sample. Adv. = adventitious carbon. Shifts in the O 1s and C 1s spectra of the reacted sample relative to the unreacted siderite samples toward higher energies by 0.3 and 0.4 eV, respectively, indicate a change in the carbonate structure, possibly due to recrystallization from siderite to cerussite on the reacted siderite surface<sup>52</sup>. The ethanol detected in the C 1s spectra originates from sample storage in ethanol.

308
310 **4.1. Reaction mechanisms of Pb(II) removal during siderite dissolution**

311 Siderite dissolves readily in acidic aqueous solution, with its solubility decreasing with 312 increasing pH³⁴,⁵⁵. Its reaction products differ strongly under different redox conditions due to the 313 redox sensitivity of the ferrous iron released into the reactive solution, which influences the 314 mechanisms of Pb(II) removal from the reactive solution.

315 **4.1.1. Siderite dissolution coupled iron oxidation and goethite precipitation**

316 During siderite dissolution under oxic conditions, the released ferrous iron oxidizes and 317 precipitates as goethite nanoparticles after > 2 h of reaction at a pH of ≥ 4.5 consistent with the 318 reaction

$$
319 \qquad \mathrm{FeCO_3(s)} + \frac{1}{2} \mathrm{O_2} + \mathrm{H_2O} \rightarrow \alpha\text{-FeOOH(s)} + \mathrm{HCO_3^-} \qquad (\text{eq. 1}).
$$

320 Note that the aqueous solution pH continuously increased during the experiment. This suggests 321 that proton-consuming siderite dissolution dominates over proton-releasing reactions of iron 322 oxidation, precipitation and adsorption during the oxic experiments. TEM and SEM imaging 323 showed heterogeneously distributed precipitates particularly growing along crystal edges, and 324 nanorods growing from the siderite surface at random orientations. Similar observations were 325 reported by Renard et al.⁵⁶, showing interface-coupled siderite dissolution and structurally 326 independent secondary ferric iron phase precipitation. Structurally independent discontinuous 327 precipitation suggests heterogeneous nucleation and growth on the siderite surface, arising from 328 the structural dissimilarity between the dissolving trigonal parent material siderite and the 329 orthorhombic goethite precipitate⁵⁷. Such coatings contain pores allowing the reactive fluid to
330 continuously dissolve the substrate without affecting the dissolution rate (Figure S6 and Table S6);
331 a similar process has been suggested for otavite (Cd-carbonate) precipitation on aragonite⁵⁸. There
332 was no evidence for potentially dissolved atmospheric CO₂(g) enhancing siderite dissolution upon
333 exposure to air during sampling. As the aqueous solution pH was less than 6, such effects would
334 be insignificant⁵⁹.

335
4.1.2. *Coupled siderite dissolution and cerussite precipitation*

336 Under both oxic and anoxic conditions, initial aqueous Pb(II) concentrations decreased with the
337 onset of siderite dissolution. Under oxic conditions, ca. 9% of Pb(II) was removed from solution
338 during the first 2 h of reaction, before the visible precipitation of goethite. Within the first 0.5 h of
339 reaction under anoxic conditions, ca. 13% Pb(II) was removed from solution. Complementary
340 PHREEQC calculations based on the measured aqueous solution composition indicate that both
341 oxic and anoxic Pb(II)-bearing aqueous solutions were saturated with respect to cerussite within
342 less than 0.5 h of reaction, and thus shortly after the onset of siderite dissolution (Table S7). Since
343 cerussite is markedly less soluble (log Ksp = -13.2)⁶⁰,⁶¹ than siderite (log Ksp = -10.9)³³ under the
344 investigated experimental conditions, Pb(II) preferentially precipitated as cerussite as soon as
345 sufficient aqueous carbonate was available from siderite dissolution regardless the redox condition
346 of the reactive solution.
347

347 Under oxic conditions, early siderite dissolution-induced cerussite precipitation may explain the
348 Pb-carbonate scattering contributions of up to ca. 19% in bulk and locally up to ca. 60% of the
349 solid-bound Pb(II) in the XANES spectra. The tabular crystallites observed in SEM images
350 growing from the siderite surface of the solids collected from the oxic experiments after 2-72 h of
351 reaction, and the pseudo-hexagonal tabular crystallites observed among the surface layer of
goethite formed by the end of the experiment (Figure S7) resemble cerussite precipitates formed on aragonite surface reported by Godelitsas et al.⁵². Similar crystallites were also observed in the solids recovered from the anoxic experiments, but were absent, however, in the solids collected from the Pb(II)-free experiments (Figure S8). These observations confirm cerussite formation under both redox conditions despite the lack of evidence for cerussite in the XRD analysis of the oxic sample, in which the minor quantities of cerussite formed were most likely overshadowed by the secondary goethite in this sample.

Under anoxic conditions, XRD, XPS and XANES analyses confirmed crystalline cerussite formation and that 100% of the Pb is bound as a carbonate at the end of the long-term experiment. Our TEM data of the reacted siderite suggest local changes of its near-surface structure. A 20 nm thick surface reaction zone, exhibiting a widening of the siderite lattice and overgrowths of poorly (poly)crystalline material were observed, with both features exhibiting lattice distances indicative of cerussite (Figure 3D-E). Pb(II) has an ionic radius of 1.19 Å in octahedral coordination, and an ionic radius of 1.35 Å in the 9-fold coordination as in cerussite⁴⁹,⁶². These Pb(II) radii are 53% and 75% larger than the radius of octahedrally coordinated Fe(II) in siderite (0.780 Å)⁶²,⁶³. Incorporation of Pb(II) into the dissolution-induced vacated FeO₆ site in siderite would thus lead to distortion and a marked widening of the siderite lattice, despite possible relaxation of the siderite surface structure due to dissolution. Although the relatively large Pb(II) and Ba(II) ions have been reported to incorporate into the [⁶]Ca(II) site in calcite⁶⁴,⁶⁵, dilation around potentially incorporated Pb(II) ions into the much smaller [⁶]Fe(II) site in siderite is expected to be much more significant than that observed for the incorporation of Pb(II) and Ba(II) into the larger metal site of Ca(II) (1.0 Å)⁶² in calcite⁶⁴. EXAFS shell fitting, however, clearly indicates the 9-fold coordination of Pb(II) in cerussite. The shift towards higher binding energies in the O and C 1s spectra further imply non-
isotype mineral replacement. Although studies of Pb(II) sorption to calcite at pH 8.3 reported no evidence of Pb incorporation or solid solution formation¹⁷, slight variations in the initial Pb(II) concentrations and pH conditions can lead to the (co)precipitation of Pb-carbonates⁶⁶,⁶⁷. Similar studies of CaCO₃ interaction with Pb and transition metals at pH 2.8 to 7 showed dissolution-induced (co)precipitation of divalent metal carbonates⁵²,⁵⁸,⁶⁸⁻⁷². While the replacement of orthorhombic aragonite by isotypic cerussite is fast and results in an epitaxial layer⁵²,⁷², the replacement of trigonal siderite by orthorhombic cerussite requires a structural change. Solvent-mediated transformation of non-isostructural carbonates show random orientation of the precipitate relative to its host substrate⁶⁹,⁷¹,⁷², similar to what was observed in TEM analyses (Figure 3D). Therefore, a coupled process of dissolution and re-precipitation-controlled transformation as shown by Yuan et al.⁷¹ and Di Lorenzo et al.⁷² for the structurally independent replacement of trigonal calcite by orthorhombic cerussite likely explains the observed changes in surface structure and mineral phase. This was likely facilitated by micro-pore generation due to the differences in structure, solubility and molar volume between siderite and cerussite⁵⁸,⁷¹,⁷².

### 4.1.3. Pb(II) adsorption to goethite and siderite

Under oxic conditions, the continued precipitation of goethite over time led to a wide, discontinuous coverage of the siderite surfaces and to the growth of these precipitates in the reactive fluid suspension through aggregation (Figure 2A-B; Figure S7). Shell-by-shell fitting of the Pb L<sub>III</sub>-edge EXAFS spectra of the bulk solid reaction products indicated interatomic distances of R<sub>Pb-O</sub> = 2.31 Å and R<sub>Pb-Fe</sub> = 3.31 Å in the first and second coordination shell, respectively (Table 2). Both interatomic distances are in perfect agreement with distances reported for Pb-adsorbed to goethite⁴⁷ and characteristic for inner-sphere bidentate edge-sharing Pb(II)-adsorption to iron (oxyhydr)oxides⁴⁴⁻⁴⁷. This conclusion is supported by the XANES spectra of the bulk sample and
area 2 (Figure 4). LC fits indicate a distorted trigonal pyramidal coordination of Pb(II) with surface oxygens characteristic for Pb(II) adsorbed to ferric (oxyhydr)oxides⁴⁴. Some of the scattering contributions also indicate a distorted tetragonal pyramidal coordination of Pb(II) as in PbO (orthorhombic)⁷³,⁷⁴, which implies the presence of an additional oxygen atom in the distorted adsorption coordination sphere. This may be supported by the first shell Pb–O distance, which is slightly longer than for a trigonal sorption geometry ($R_{Pb-O} \sim 2.25–2.28 \ Å$)⁴⁴,⁷⁴. Contributions of these scatterings are particularly prominent in the locally resolved XANES mapping spectra and potentially result from dehydration and enclosure of Pb(II) ions during advanced aggregation and growth of the iron oxyhydroxide surface precipitates on the reacted siderite grains. The absence of second-neighbor oxygens and the low coordination number suggest that Pb(II) is adsorbed rather than incorporated⁴⁷,⁷⁴. Also, the absence of third shell scatterings exclude ternary absorption of carbonate²⁰; in contrast, the first-shell sensitive XANES spectrum suggests the Pb-carbonate coordination of cerussite.

Since at least ca. 9% of the initial Pb(II) concentrations in the reactive aqueous solution was likely consumed by cerussite formation preceding goethite precipitation, the majority of remaining carbonate-bound Pb(II) was likely consumed by the growth of the cerussite precipitates to ≥ 100 nm large crystallites (Figure 2A, Figure S7) until the reactive aqueous solution reached equilibrium with respect to cerussite after approximately 500-600 h of reaction and a pH of 5.0 (Figure 1; Table S7). However, the early sorption of Pb(II) to siderite, as reported to occur within 30 min under acidic pH conditions⁷⁵, cannot be ruled out. Bivalent metal carbonates $MeCO_3$ (with Me = Ca, Fe, Mn) generally exhibit a high affinity to sorb heavy metals such as Zn(II), Cd(II) and Pb(II) to their surfaces¹⁵⁻¹⁷,⁵³,⁶⁶⁻⁶⁸,⁷⁵,⁷⁶. For siderite, the pH of neutral surface charge ($pH_{PZC}$; point of zero charge) is 5.3 ±0.1⁷⁷. At pH < $pH_{PZC}$ carbonate mineral surfaces exposed to aqueous solution adsorb
protons<sup>77</sup>; in the presence of dissolved divalent metals, these can compete with protons for the carbonate surface sites, contributing to the development of a more positive surface charge below the pH<sub>PZC</sub><sup>78</sup>. This suggests that Pb(II) could also have sorbed to siderite early in our experiments, forming >CO<sub>3</sub>Pb<sup>+</sup> surface complexes<sup>77,78</sup>, preceding precipitation and growth of cerussite on the siderite surface. Therefore, Pb(II) uptake from the oxic system was likely three-fold: (I) potential sorption to siderite, (II) precipitation as cerussite, and (III) adsorption to goethite. As siderite surface coverage by goethite precipitates advanced over time, and the siderite dissolution rate significantly decreases at a pH ≥ 5,<sup>34</sup> the adsorption to goethite likely dominated the drawdown of aqueous Pb(II) toward the end of these experiments.

## 4.2. Environmental implications

In nature, the redox conditions of metal-contaminated sites may not be constant over time, potentially causing sorbed metals to be released<sup>14,21</sup>. Siderite, however, can effectively remove and sustainably retain toxic Pb(II) by precipitation in cerussite from acidic aqueous solutions at varying redox conditions, with no additional harm to the environment. Under either oxygenated or reduced conditions, Pb(II) uptake is governed by siderite dissolution, which is accompanied by rapid increase in pH to less acidic conditions. Initially, some Pb(II) rapidly precipitated as cerussite under either redox condition, followed by its simultaneous sorption onto oxidation-induced goethite nanoprecipitates under oxic conditions, or further precipitation as cerussite under anoxic conditions.

Both mechanisms generate products that are stable under the investigated experimental and comparable natural conditions. Potentially, these dissolution-coupled mechanisms complementarily ensure Pb(II) retention in systems undergoing redox fluctuations, such as
443 contaminated soils, floodplains, or stream sediments. For example, goethite as the reaction product
444 of the oxic experiment is very insoluble under oxic conditions (log Ksp = -39.80)⁷⁹. If, however,
445 redox conditions change from oxic to anoxic, reductive dissolution of goethite may occur²³,⁴¹⁻⁴³,⁸⁰.
446 As a result, the strong inner-sphere complexes of Pb(II) with the goethite surface may become
447 unstable and susceptible to desorption¹¹,¹⁸. Under reduced and acidic conditions, however, Pb(II)
448 readily precipitates as sparingly soluble cerussite. Since the solubility product of cerussite is
449 substantially lower than that of siderite, Pb(II) is much less susceptible to be liberated from
450 cerussite. Redox sensitivity of carbonate-bound Pb(II) is also negligible compared to Fe(III)
451 (oxyhydr)oxide adsorbed Pb(II). Notably, when Pb(II) was completely consumed by PbCO3
452 formation within 96-144 h of the anoxic experiment, no release of Pb(II) was observed over the
453 remaining experiment duration of > 850 h. This is because of the low solubility and high stability
454 of cerussite over a wide range of pH conditions⁶⁰,⁸¹, inhibiting Pb remobilization. In turn, the
455 oxygenation of the FeCO3–Pb/PbCO3–H2O system would provoke goethite precipitation and
456 Pb(II) adsorption to goethite. The observed aggregation of the nanoparticulate goethite may
457 additionally promote Pb(II) retention in the oxic system⁸²,⁸³. Elevated pCO2(g) at above neutral
458 pH, as often present in soil fluids or carbonate-rich surface waters, may enhance Pb(II) retention
459 by enhanced siderite dissolution⁵⁹, cerussite precipitation⁶⁰,⁸¹ and/or ternary surface complexation
460 on ferric (oxyhydr)oxide precipitates²⁰.

461

462 5. CONCLUSIONS

463 Siderite dissolution effectively removes 90% and 100% of initially 0.48 x10⁻³ mol kg⁻¹ Pb(II)
464 from acidic oxic and anoxic aqueous solutions via three main pathways: neutralization, Pb-
465 carbonate formation and adsorption. Under either redox condition, the liberation of carbonate
466 during siderite dissolution leads to rapid increase in reactive fluid pH. Simultaneously, Pb(II)
467 rapidly precipitates as PbCO₃, which was identified by XRD, TEM and EXAFS as nanoparticulate
468 cerussite. Additional XANES-mapping and XPS analyses confirmed Pb-carbonate formation on
469 the siderite surface, which particularly controls Pb(II) consumption under anoxic conditions as
470 well as during early stages of the reaction under oxic conditions. EXAFS shell-by-shell fitting
471 showed that Pb(II) also readily forms inner-sphere Pb(II) adsorption complexes with the oxidation-
472 induced goethite nanoprecipitates under oxic conditions, which dominates Pb(II) uptake from the
473 oxic aqueous solution at a later stage of the reaction.

474 Overall, the nanoparticulate and scarcely soluble reaction products remained stable and no Pb
475 was re-released over the entire experiment duration of 1008 h. Reduced conditions, however, allow
476 faster and more stable capture of Pb(II) than oxic conditions due to cerussite formation. We further
477 showed that the presence of Pb(II) does not affect the dissolution rate of the micro-granular natural
478 siderite material, which remained governed by the reactive solution pH and its saturation state in
479 the aqueous system. As a common iron ore (current market price $103.30/t)⁸⁴, the material costs
480 for siderite used as remediation reagent can be estimated assuming a reaction efficiency based on
481 this study of 13.5 mg Pb sequestered per g siderite. Treatment of water contaminated with 100 mg
482 L⁻¹ aqueous Pb(II) using natural siderite would cost approximately $0.69 per ton contaminated
483 water, or $2.00/t if using synthetic siderite⁸⁵. This study thus demonstrates that natural siderite
484 dissolution is an effective, low-cost and sustainable Pb(II) sequestration method for potential
485 acidic wastewater remediation applications under either redox condition.
486 ASSOCIATED CONTENT

487 **Supporting Information.** The following files are available free of charge: 1) Supporting information on methodology, microscopic and spectroscopic results, and geochemical calculations (PDF). 2) Chemical solution data (MS Excel).

490 AUTHOR INFORMATION

491 **Corresponding Authors**

492 * Lisa C. Füllenbach. E-Mail: l.fuellenbach@ucl.ac.uk

493 * Eric H. Oelkers. E-Mail: e.oelkers@ucl.ac.uk

494 **Author Contributions**

495 The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript.

497 **Notes**

498 The authors declare no competing financial interest.

499 ACKNOWLEDGMENTS

500 This work was funded by the European Union's Horizon 2020 Marie Skłodowska-Curie Innovative Training Network Grant No. 675219. L.G. Benning acknowledges the financial support from the German Helmholtz Recruiting Initiative funding (award number I-044-16-01). Synchrotron mapping work was carried out with the support of the Diamond Light Source, instrument I14 (proposal MG21719-1). EXAFS work was conducted with the support of the ANKA-KIT, instrument SUL-X (proposal A2017-028-012253). Many thanks go to I. Wood for the XRD
analyses, to G. Saldi for providing BET measurements, to A. Schreiber for preparing FIB sections
for TEM, and to J. Stockmann (BAM) for performing the XPS measurements.

## 508 ABBREVIATIONS

Pb, lead; Pb O (LT), siderite dissolution experiment in Pb(II)-bearing solution under oxic conditions (long-term); Pb A (LT), siderite dissolution experiment in Pb(II)-bearing solution under anoxic conditions (long-term); ICP-OES, inductively coupled plasma optical emission spectroscopy; SFA, segmented flow analysis; XRD, X-ray (powder) diffraction; (FE-)SEM, (field emission) scanning electron microscopy; EDX, energy dispersive X-ray spectroscopy; FIB, focused ion beam; (HR)TEM, (high-resolution) transmission electron microscopy; XAS, X-ray absorption spectroscopy; XANES, X-ray absorption near edge structure; EXAFS, extended X-ray absorption fine structure; XRF, X-ray fluorescence; XPS, X-ray photoelectron spectroscopy.

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