Hong Ying

Xionghan Feng

Mengqiang Zhu

Bruno Lanson

Fan Liu

Wang

Xiaoming Wang

Formation and transformation of schwertmannite through direct Fe 3+ hydrolysis 2 under various geochemical conditions 3 4

Keywords: schwertmannite, Fe 3+ hydrolysis, formation, transformation, AMD 47 1. Introduction 48

Electronic supplementary information (ESI) is available free of charge via the internet at http://pubs.rsc.org, including (1) specific experimental conditions of the formation and transformation of schwertmannite; (2) FTIR spectra of schwertmannite obtained from Fe 3+ hydrolysis-dialysis pathway and of the aging products of schwertmannite formed through Fe 3+ hydrolysis by adding NaOH; (3) XRD patterns of schwertmannite with the coexistence of K + or NH 4 + and the finial aging products of schwertmannite in the presence of NH 4 + ; and (4) schwertmannite transformation rate indicated by Fe o /Fe t obtained from acidic dissolution experiments.

The Fe concentration in the solution was measured using a modified 1, 215 10-phenanthroline colorimetric method. [START_REF] Burton | Schwertmannite transformation 586 to goethite via the Fe(II) pathway: Reaction rates and implications for iron-sulfide formation, 587 Geochim[END_REF] The ratio of Fe o /Fe t indicates the 216 transformation rate of schwertmannite. enhanced peak intensity at 2.55 Å (Fig. 1a). As shown in flower-shaped and composed of small thin plates (Fig. 2a). With increasing hydrolysis 243 temperature, the thin plates transform to long strips, connecting to a "sea-urchin"-like 244 structure (Fig. 2a-d). When the hydrolysis temperature increases to 70 °C and above, 245 minor goethite impurities form, evidenced by its characteristic peaks at 4.18 Å and 1.72 246 Å (Fig. 1a), while the crystallinity of schwertmannite slightly decreases (Fig. 1a). In 247 addition, SO 4 2-content decreases and Fe/S ratio increases (Table 1), and the long strips 248 of schwertmannite become more clustered (Fig. 2e-f).

249

The FTIR spectra of these samples are similar and consistent with typical 250 schwertmannite IR fingerprint. [START_REF] Wang | Sulfate local coordination environment in 584 schwertmannite[END_REF][START_REF] Boily | FTIR spectral 592 components of schwertmannite[END_REF] These spectra consist of a broad triply degenerate 251 asymmetric stretching (ν 3 ) band at ∼1125 cm -1 with two shoulder bands at ∼1050 and 252 ∼1205 cm -1 , a ν 1 fundamental of the symmetric sulfate stretching at 980 cm -1 , a ν 4 253 bending band at ∼610 cm -1 , and a Fe-O stretching band at ~698 cm -1 (Fig. S1a).

254

Additional weak IR vibration bands at 885 cm -1 and 792 cm -1 are shown for the 255 samples hydrolyzed at 70 °C and 80 °C, indicating the formation of minor goethite (Fig.

S1a), consistent with XRD results (Fig. 1a).

257

With increasing dialysis time, diffraction peak intensities of schwertmannite 258 slightly decrease, especially for the peak at 2.55 Å (Fig. 1b), consistent with the 259 decrease of SO 4 2-content and of sulfate IR band intensities (Table 1 and Fig. S1b). In 260 addition, schwertmannite particles aggregate more closely (Figs. 2g-i), leading to a 261 "hedgehog" morphology. [START_REF] Loan | Evidence on the structure of synthetic 594 schwertmannite[END_REF] Furthermore, compared to the Na + system, the presence of appear and increase, indicative of schwertmannite formation (Fig. 3b). In contrast, the 272 pre-formed two-line ferrihydrite aged in the same sulfate concentration and solution pH 273 does not transform to schwertmannite (Fig. 3c).

274 Sulfate-bearing ferrihydrite thus appears as an intermediate phase during the 275 formation of schwertmannite (Fig. 3), similar with the observation in Zhu et al. [START_REF] Zhu | Early stage 544 formation of iron oxyhydroxides during neutralization of simulated acid mine drainage solutions, 545[END_REF] .

276

Transformation from ferrihydrite to schwertmannite is most likely because of a higher 277 thermodynamic stability than ferrihydrite. [START_REF] Loan | On the crystal growth of nanoscale 536 schwertmannite[END_REF][START_REF] Loan | Evidence on the structure of synthetic 594 schwertmannite[END_REF] However, the contrast between the 278 transformation from intermediate ferrihydrite-like mineral (Fig. 3a) and from 279 pre-formed ferrihydrite (Fig. 3c) suggests that sulfate is closely associated with 280 ferrihydrite in a way that differs from simple surface adsorption. [START_REF] Gu | X-ray absorption spectroscopic 596 quantification and speciation modeling of sulfate adsorption on ferrihydrite surfaces[END_REF] thus the formed ferrihydrite readily transforms to schwertmannite. In addition, the 285 ferrihydrite formed in the presence of sulfate may have smaller particle size and more 286 structural defects than ferrihydrite synthesized in the absence of sulfate, which further 287 enhances its transformation to schwertmannite. [START_REF] Zhang | The transformation of two-line 599 ferrihydrite into crystalline products: effect of pH and media (sulfate versus nitrate)[END_REF]

288

Schwertmannite can form at 25 °C through both Fe 3+ hydrolysis-dialysis (Fig. 1) 289 and Fe 3+ hydrolysis by adding NaOH (Fig. 3, 4a and S2). However, a theoretical 290 calculation indicates that schwertmannite formation through direct Fe 3+ hydrolysis is not 291 thermodynamically spontaneous under ambient condition (~ 25 °C), due to the reaction 292 free energy ∆G°2 98 of 6.275 kJ•mol -1 as the following equation. [START_REF] Zhou | Biomineralization: a pivotal process in developing a novel passive treatment system for 473 acid mine drainage[END_REF] 293

Fe 3+ (aq) + 7/4H 2 O(l) + 1/8SO 4 2-(aq) = 1/8Fe 8 O 8 (OH) 6 SO 4 (s) (Sch) + 11/4H + (aq) 294
This apparent discrepancy could be explained by the migration of protons (H + ) 295 through the dialysis membrane into distilled water during dialysis process. As a result, 296 the solution pH in the dialysis bag, initially very low and unfavorable to Fe 3+ hydrolysis, 297 gradually increases and then induces schwertmannite formation. On the other hand, the 298 schwertmannite readily forms from Fe 3+ hydrolysis in the presence of sulfate by adding 299 NaOH at 25 °C (Fig. 3 and4), consistent with previous reports. [START_REF] Burgos | Schwertmannite and Fe oxides formed by biological low-pH Fe(II) oxidation versus 539 abiotic neutralization: Impact on trace metal sequestration[END_REF][START_REF] Reichelt | Transformation of nanostructured schwertmannite and 542 2-line-ferrihydrite into hematite[END_REF][START_REF] Zhu | Early stage 544 formation of iron oxyhydroxides during neutralization of simulated acid mine drainage solutions, 545[END_REF] Consequently, direct 300 Fe 3+ hydrolysis should be considered as an important pathway of schwertmannite 301 formation in watershed or soils surrounded by AMD with a relative high pH. rates (above 3.33 µM Fe/min), with a lower hydrolysis rate leading to an improved 307 crystallinity, compared to the mixed directly system (Fig. 4a). In addition, 308 schwertmannite morphology varies substantially with Fe 3+ hydrolysis rate, exhibiting 309 longer strips and resulting in larger and looser aggregates at a lower Fe 3+ hydrolysis 310 rate, similar to the sea urchin-like aggregates that form from Fe 2+ oxidation. [START_REF] Zhang | Schwertmannite: occurrence, properties, synthesis 574 and application in environmental remediation[END_REF] However, 311 the morphology obtained from mixed directly system is block agglomerate (Fig. 4b). ). In addition, schwertmannite can be obtained only at low initial Fe/S molar ratio 319 (Fe/S ≤ 5) (Fig. 3 and 9b), i.e., high sulfate concentration, while ferrihydrite forms when 320 the Fe/S molar ratios exceed 8 to 10, suggesting that sulfate must be present in excess, 321 relative to the ideal mineral composition, to allow for schwertmannite formation.

322

XRD patterns indicate that the presence of Cl -does not affect schwertmannite 323 formation (Fig. 9d), suggesting that although Cl -can promote akaganéite 324 (FeO(OH) 1-x Cl x ) formation under acidic conditions, [START_REF] Zhu | Early stage 544 formation of iron oxyhydroxides during neutralization of simulated acid mine drainage solutions, 545[END_REF] the presence of numerous sulfate 325 probably inhibits the formation of akaganéite, similar to the reports in Fe 2+ 326 bio-oxidation system. [START_REF] Xiong | Influence of chloride and sulfate on formation of akaganéite and 579 schwertmannite through ferrous biooxidation by acidithiobacillus ferrooxidans cells[END_REF][START_REF] Ishikawa | Influence of anions on the formation 602 of β-FeOOH rusts[END_REF] Pure schwertmannite could thus be synthesized through FeCl 3 327 hydrolysis-dialysis if sufficient sulfate is present. [START_REF] Bigham | A poorly crystallized 461 oxyhydroxysulfate of iron formed by bacterial oxidation of Fe(II) in acid-mine waters[END_REF] Additionally, the presence of K + or With increasing aging time to 24 h, goethite is present in all samples, with more 338 goethite occurring at higher Fe 3+ hydrolysis rate (Fig. 5b), suggesting that the 339 schwertmannite obtained at a higher Fe 3+ hydrolysis rate is less stable, possibly due to 340 its lower crystallinity. FTIR spectra of the samples aged for 3 h (Fig. 5c) and 24 h (Fig. 341 5d) show that the characteristic Fe-O stretching band of schwertmannite at ~698 cm -1 342 decreases with aging time, whereas the characteristic OH bending vibration bands of 343 goethite at 885 cm -1 and 792 cm -1 increase with increasing Fe 3+ hydrolysis rate and 344 aging time, confirming XRD results (Fig. 5a and5b). Thus, lower Fe 3+ hydrolysis rate 345 leads to larger crystallite size and higher structural stability of schwertmannite, thus 346 disfavoring its transformation to other phases. This observation was also made for 347 schwertmannite formed from Fe 2+ oxidation. At 80 ºC and Fe/S = 2, the concentration of dissolved Fe 3+ decreases with 366 increasing pH from 2.0 to 3.5 (Fig. 7a), while that of dissolved SO 4 2-increases (Fig. 8a), 367 indicating that a higher pH enhances SO 4 2-release and thus the transformation from 368 schwertmannite to goethite. This transformation is supported by the XRD patterns (Fig. bending vibration bands of goethite at 885 cm -1 and 792 cm -1 increase with increasing 373 pH (Fig. S3a), further confirming the above analyses. The increased transformation rate 374 of schwertmannite to goethite can be partially ascribed to the easier sulfate release at a 375 higher pH. [START_REF] Jonsson | Schwertmannite precipitated from acid mine 550 drainage: phase transformation, sulphate release and surface properties[END_REF]

376

For a given pH and Fe/S molar ratio, the concentration of dissolved Fe 3+ 377 substantially decreases with increasing aging temperature (from 25 to 80 o C) and aging 378 time (Fig. 7c). As to dissolved SO 4 2-, it increases slightly at 25 ºC and more 379 significantly at higher temperatures with increasing aging time (Fig. 8b), suggesting that 380 SO 4 2-release is enhanced at a higher temperature during schwertmannite 381 transformation. Additionally, XRD patterns (Fig. 9c andS3c) and acidic dissolution 382 results (Fig. S4a) indicate that goethite formation is favored at higher aging 383 temperatures and longer aging time, consistent with lower concentration of dissolved 384 Fe 3+ (Fig. 7c) and higher concentration of dissolved SO 4 2-(Fig. 8b). In addition, the 385 mineral phases formed at different Fe/S molar ratios exhibit different stabilities at 60 °C 386 and pH 3 (Fig. 9b andS3b), with lower Fe/S molar ratio resulting in lower 387 concentration of dissolved Fe 3+ (Fig. 8b). As a consequence, schwertmannite obtained at 388 a lower Fe/S molar ratio (Fe/S ≤ 5) transforms more readily to goethite than ferrihydrite 389 which is obtained at higher Fe/S ratios (Fe/S > 5) (Fig. 9b andS3b). Cl -(Fe/Cl -= 0.2) inhibits SO 4 2-release and the transformation of schwertmannite to 395 goethite (Fig. 8c and9d), possibly because massive Cl -adsorption on schwertmannite 396 surface impedes SO 4 2-release and thus the dissolution-recrystallization process. As 397 reported previously, adsorption of ions on schwertmannite surface could stabilize its 398 structure. [START_REF] Burton | Arsenic 560 effects and behavior in association with the Fe(II)-catalyzed transformation of schwertmannite, 561[END_REF][START_REF] Schoepfer | Contrasting effects of phosphate on the rapid 620 transformation of schwertmannite to Fe(III) (oxy)hydroxides at near-neutral pH[END_REF] 399

The presence of K + promotes hydrolysis and precipitation of dissolved Fe 3+ (Fig. 400 7d) and inhibits SO 4 2-release (Fig. 8d) through the formation of K + -bearing jarosite, 401 exhibiting aggregated spherical particle morphology (Fig. 9d and10d). K + -bearing 402 jarosite formation is also confirmed by FTIR spectrum with bands at around 1080 ~ 403 1200 cm -1 , 1008 cm -1 , and 628 cm -1 (Fig. S3d), respectively, assigned to the vibrational 404 modes of v 3 (SO 4 2-), v 1 (SO 4 2-), and v 4 (SO 4 2-). [START_REF] Sasaki | Distinction of jarosie-group compounds by raman 623 spectroscopy[END_REF] Similarly, a high NH

Conclusion and implications

[Type here] 416

In the present study, schwertmannite formation through direct Fe 3+ hydrolysis and 417 its subsequent transformation have been systematically investigated under various 418 geochemical conditions. Schwertmannite can be obtained over a Fe 3+ hydrolysis 419 temperature range of 25 -60 °C, a pH range of 2.0 -3.5, molar ratios of Fe/S lower than 420 5, Fe 3+ hydrolysis rates higher than 3.33 µM/min, and in the presence of Cl -, K + or NH 4 + 421 (Fig. 11). Such an extended range of conditions suggests that direct Fe 3+ [Type here]

  of 0.2 M NaOH was added dropwise to a 74 mL mixed solution containing 173 Fe(NO 3 ) 3 and Na 2 SO 4 at addition rates of 0.5, 0.1, and 0.05 mL/min, respectively, hydrolysis rates of 33.33, 6.67, and 3.33 µM/min. The direct mix of 175 NaOH, Fe(NO 3 ) 3 and Na 2 SO 4 solutions was recognized as control system (named as 176 mixed directly), approximately representing an infinite Fe 3+ hydrolysis rate. Solution pH 177 was then adjusted to 3 for all experiments with a Fe/S molar ratio of 2 [c(Fe 3+ ) = 48.55 178 mM, c(SO 4 2-) = 24.275 mM]. In addition, to determine the effects of Fe 2+ on 179 schwertmannite transformation, a Fe 2+ solution [c(Fe 3+ )/c(Fe 2+ ) = 10] was added to the 180 fresh schwertmannite suspension formed form the condition of a NaOH addition rate of 181 0.1 mL/min. Subsequently, all samples were sealed and aged at 60 °C for 96 h, during 182 which, the solution pH were maintained at pH 3 by adding 0.1 M NaOH or HNO 3 at 183 regular time intervals. The containers were open to air during pH adjustment and 184 sampling, allowing entrance of atmospheric O 2 to the reaction systems. At pre-set time 185 intervals, a 15 mL suspension was sampled and passed through a 0.22 μm membrane 186 filter mounted on a vacuum apparatus. The dissolved SO 4 2-and Fe concentrations in the 187 filtrates were measured as described above, whereas the wet solids on the membrane 188 were rinsed with 30 mL deionized water, air dried, ground, and analyzed by XRD, 189 FTIR, SEM, and acidic dissolution.

190 2 . 3 . 3 2 194M 2 . 3 . 4 NH 4 + 2 . 3 . 5

 23322344235 Effects of pH, Fe/S molar ratio, and temperature on the long-term aging of 191 schwertmannite 192 Solutions containing Fe(NO 3 ) 3 and Na 2 SO 4 with Fe/S molar ratios of 1.5, 2.0, 2.5, 193 5.0, 8.0 or 10.0 [c(Fe 3+ ) = 48.55 mM] were directly mixed with different amounts of 0.NaOH solution, leading to final pH values of pH 2.0, 2.5, 3.0 and 3.5 with total 195 suspension volume of 200 mL. The suspensions with different pHs (Fe/S molar ratio = [Type here] 196 2.0) were aged at 80 °C for 96 h, while the suspensions with different Fe/S molar ratios 197 (pH = 3.0) were aged at 60 °C for 96 h. Subsequent experimental details are the same as 198 in Section 2.3.2. 199 Effects of co-existing Cl -, K + , or NH 4 + on the long-term aging of 200 schwertmannite 201 Solutions containing Fe(NO 3 ) 3 , Na 2 SO 4 , and NaCl [c(Fe 3+ ) = 48.55 mM, Fe/S = 2, 202 Fe/Cl = 2, 0.2] were directly mixed with 0.2 M NaOH solution to obtain 203 schwertmannite suspensions with a total volume of 200 mL. The schwertmannite 204 suspensions were then aged at 60 °C and pH 3 for 96 h. In addition, solutions containing 205 Fe(NO 3 ) 3 , K 2 SO 4 or (NH 4 ) 2 SO 4 [c(Fe 3+ ) = 48.55 mM, Fe/S = 2 and 0.2, Fe/K + or 206 =1 and 0.1] were directly mixed with 0.2 M NaOH solution to obtain 207 schwertmannite suspensions. The suspensions were subsequently aged at 80 °C and pH 208 2 for 96 h. Subsequent experimental details are the same as in Section 2.3.2. 209 Acidic dissolution experiments 210 To explore the transformation rate of schwertmannite under different conditions, 211 10 mg of dried intermediate samples were dissolved in 10 mL of 0.2 M acidic 212 ammonium oxalate solution (pH 3, Fe o ) or 4 M HCl solution (Fe t ) at RT for 2 h under 213 stirring. Resulting solution was then immediately filtered through a 0.22 µm membrane.214

3 . Results and discussion 234 3 . 1

 331 for phase identification were recorded from 10 to 70° 219 at a scan speed of 1°•min -1 on a Bruker D8 Advance X-ray diffractometer equipped with 220 a LynxEye detector using Ni-filtered Cu Kα radiation (λ = 0.15418 nm). To identify 221 mineral phases with poor crystallinity, synchrotron-based XRD data of part samples 222 were collected at beamline BL14B1 (λ = 0.6895 Å, scanning from 0.5 to 30° and 223 exposure time of 30 s) at the Shanghai Synchrotron Radiation Facility (SSRF) 52 or at 224 beamline 11-ID-B at the Advanced Photon Source (APS), Argonne National Laboratory 225 (X-ray energy of 58.86 keV, λ = 0.2112 Å). 50 226 The morphology and particle size of the samples were observed by SEM (SU8000) 227 at an accelerating voltage of 10 or 20 kV. Specifically, a small amount of sample was 228 pasted on conductive glue, and then plated using a sputtering apparatus. In addition, 229 FTIR of the samples was recorded on a Bruker VERTEX 70 spectrophotometer. The 230 samples were mixed gently with KBr (~1% sample weight) in an agate mortar and 231 pelletized. Each sample was collected 128 scans with a resolution of 4 cm -1 over the 232 4000-400 cm -1 against the air background. 233 Formation of schwertmannite through Fe 3+ hydrolysis-dialysis 235 XRD patterns of the products obtained from different Fe 3+ hydrolysis temperatures 236 are shown in Figure 1a. With increasing hydrolysis temperature from 25 -60 °C, pure 237 schwertmannite is obtained, and its crystallinity gradually increases evidenced by the 238

302 3 . 3 3 . 3 . 1

 33331 Effects of geochemical conditions on schwertmannite formation through Fe 3+ 303 hydrolysis by adding NaOH 304 Effect of Fe 3+ hydrolysis rate 305 XRD patterns indicate that pure schwertmannite forms at different Fe 3+ hydrolysis Page 33 of 61 Environmental Science: Nano [Type here] 306

312315 3 . 3 . 2 317 ( 4 318- 3 .2 46

 3323174346 These schwertmannite samples, obtained from different Fe 3+ hydrolysis rates, should 313 have distinct surface reactivity and geochemical behaviors, which need to be further 314 studied. Effects of pH, Fe/S molar ratio and co-existing ions 316 Schwertmannite can form through Fe 3+ hydrolysis over the pH range of 2.0 -3.5 Fig 4a, 9a and S2) that is slightly wider compared to that of Fe 2+ oxidation (i.e., pH 2.

328NH 4 + 3 . 4 3 . 4 . 1

 434341 slightly increases schwertmannite crystallinity (Fig.S2), similar to the Fe 3+ 329 hydrolysis-dialysis pathway (Fig.1c), probably because K + or NH 4 + ions enter the 330 schwertmannite tunnel structure during Fe 3+ hydrolysis, thus enhancing schwertmannite 331 crystal growth. 332 Effects of geochemical conditions on the transformation of schwertmannite 333 formed through Fe 3+ hydrolysis by adding NaOH 334 Effect of Fe 3+ hydrolysis rate 335During the transformation of schwertmannite formed at different Fe 3+ hydrolysis 336 rates, minor goethite form only in the mixed directly sample after a 3 h aging (Fig.5a).

  337

3699a)

  and the lower Fe o /Fe t (ratio of dissolved Fe in oxalate to total Fe) measured at 370 higher pH values (Fig.S4b). Additionally, FTIR spectra indicate that the characteristic 371 Fe-O stretching band of schwertmannite at ~701 cm -1 gradually decreases whereas OH Page 36 of 61 Environmental Science: Nano [Type here]

390 3 . 4 . 4

 344 Effects of co-existing Cl -, K + , or NH 4 + 391 Compared to the system without Cl -, the presence of Cl -(Fe/Cl -= 2) accelerates the 392 release of SO 4 2-(Fig. 8c), possibly because Cl -promotes SO 4 2-release by ligand 393 exchange, enhancing goethite formation (Fig. 9d). In contrast, a high concentration of Page 37 of 61 Environmental Science: Nano [Type here] 394

4 + concentration 405 (

 4405 Fe/NH 4 + = 0.1) favors jarosite formation with very low concentration of dissolved 406 SO 4 2-(Fig.8d), leading to the coexistence of goethite and NH 4 + -bearing jarosite (Fig.407S5), coincident with the results observed in Fe 2+ bio-oxidation system.[START_REF] Jones | Synthesis and properties of ternary 625 (K, NH 4 , H 3 O)-jarosites precipitated from Acidithiobacillus ferrooxidans cultures in simulated 626 bioleaching solutions[END_REF][START_REF] Jones | Formation and characterization of 628 ternary (Na, NH 4 , H 3 O)-jarosites produced from Acidithiobacillus ferrooxidans cultures[END_REF] More 408 lath-like particles and looser aggregates are observed at a higher concentration of NH 4 + 409 (Fig.10 b). However, the presence of low NH 4 + (Fe/NH 4 + = 1) increases SO 4 2-release 410 (Fig.8d), thus favoring the formation of goethite (Fig.S5). Compared to NH 4 + , the 411 presence of K + is easier to induce jarosite formation, consistent with a previous study 412 indicating that the ability of K + to promote jarosite formation is ~ 75× greater than that 413 of NH 4 + ,[START_REF] Bai | Both initial concentrations of 631 Fe(II) and monovalent cations jointly determine the formation of biogenic iron hydroxysulfate 632 precipitates in acidic sulfate-rich environments[END_REF] probably because that the ionic radius of hydrated K + (1.32 Å) fits better in 414 jarosite structure than that of NH 4 + (1.44 Å).

  415

Figure Captions 635 Fig. 1 .Fig. 2 .Fig. 3 .Fig. 4 .Fig. 5 .

 63512345 Figure Captions

Fig. 6 .Fig. 7 .Fig. 8 .Fig. 9 .Fig. 10 .Fig. 11 .[

 67891011 Fig. 6. XRD patterns of the transformation products of schwertmannite, obtained from

Fig. 1 .Fig. 2 .

 12 Fig. 1. XRD patterns of the products obtained from Fe 3+ hydrolysis-dialysis at different

Fig. 3 .Fig. 4 .

 34 Fig. 3. Synchrotron-based XRD patterns (a) and pair distribution functions [G(r)s] (b)

Fig. 5 .Fig. 6 .Fig. 7 .Fig. 10 .Fig. 11 .

 5671011 Fig. 5. Synchrotron-based XRD patterns and FTIR spectra of the products obtained

25 Fig. S1 . 26 Fe

 25S126 Fig. S1. FTIR spectra of the products obtained from Fe 3+ hydrolysis-dialysis at different

.

  Fig. S2. Synchrotron based XRD patterns of the products obtained from quick

37 Fig. S3 .

 37S3 Fig. S3. FTIR spectra of the mineral phases obtained from quick hydrolysis of Fe 3+ and

Table

.2 Mineral evolution during Fe 3+ hydrolysis by adding NaOH

	Page 31 of 61	Environmental Science: Nano
	262	K + or NH 4 + (Fig. 1c) slightly increases the crystallinity of schwertmannite but decreases
	263	the SO 4 2-content in schwertmannite (Table 1).
	265	During Fe 3+ hydrolysis in the presence of SO 4 2-, XRD patterns of the intermediate
	266	products indicate that ferrihydrite is the initial product and schwertmannite is observed
	267	later but within 45 min (Fig. 3a). The corresponding PDF data [G(r)s] of intermediate
	268	products show that the PDF data of initial sample (~3.2 min) is similar to that of a
	269	ferrihydrite reference. With increasing aging time, the intensity of Fe-Fe peak at 5.42 Å
	270	increases, and the Fe-Fe peaks at 7.40 Å and 11.06 Å and Fe-O peak at 6.08 Å gradually
	271
		[Type here]

264 3

.2 Effect of the presence of Fe 2+

	15, 20, 58 Environmental Science: Nano stronger if Fe 2+ is present (Fig. 6a), compared to the Fe 2+ -free system (Fig 5b), suggesting that Fe 2+ significantly accelerates schwertmannite transformation to goethite. In addition, the concentration of dissolved SO 4 2-is higher in the presence of Fe 2+ than that in the Fe 2+ -free system (Fig. 6b), indicating that Fe 2+ promotes the release of sulfate during schwertmannite dissolution and transformation. During the aging process, the concentration of dissolved Fe 2+ decreases with increasing aging time, while that of dissolved SO 4 2-increases (Fig. 6b), suggesting that schwertmannite transformation to goethite occurs through a dissolution-recrystallization mechanism by consuming Fe 2+ 3.4Page 35 of 61 348 350 351 352 353 354 355 356 357 358 and releasing SO 4 2-. Enhancement of schwertmannite transformation in the presence of
	[Type here]

349

After aging for 24 h, XRD patterns indicate that goethite diffraction peaks are 359 Fe 2+ possibly results from two aspects. On one hand, aqueous Fe(II) species can 360 exchange structural Fe(III), thus enhancing the mineral phase reorganization, as 361 evidenced by stable Fe isotope tracers.

[START_REF] Frierdich | Iron atom 607 exchange between hematite and aqueous Fe(II)[END_REF][START_REF] Liu | Fe(II)-induced phase 609 transformation of ferrihydrite: The inhibition effects and stabilization of divalent metal cations[END_REF][START_REF] Neumann | Atom exchange between aqueous Fe(II) and structural Fe in clay minerals[END_REF]

On the other hand, electron transfer between 362 adsorbed Fe(II) and structural Fe(III) promotes the reductive dissolution of 363 schwertmannite and its subsequent recrystallization.

[START_REF] Liu | Fe(II)-induced phase 609 transformation of ferrihydrite: The inhibition effects and stabilization of divalent metal cations[END_REF][START_REF] Larese-Casanova | Heterogeneous oxidation of Fe(II) on iron 615 oxides in aqueous systems: Identification and controls of Fe(III) product formation[END_REF][START_REF] Jones | Redox characterization of the Fe(II)-catalyzed 618 transformation of ferrihydrite to goethite[END_REF]

364 3.

4.3 Effects of pH, temperature, and Fe/S molar ratio 365

Table 1 .

 1 The Fe and S contents in the schwertmannite samples formed through Fe 3+ 679 hydrolysis-dialysis at different hydrolysis temperatures, dialysis time, and in the

	680	presence of K + or NH 4 + .
			Sample		Fe	(mmol/g)	SO 4	molar ratio of Fe/S
			25 o C -7d	10.69	1.71	6.25
			40 o C -7d	10.27	1.78	5.77
			50 o C -7d	10.65	1.86	5.73
			60 o C -7d	10.52	2.03	5.18
			70 o C -7d	10.28	1.73	5.94
			25 o C -15d	11.55	1.67	6.92
			60 o C -15d	11.69	1.79	6.53
		60 o C -7d -K +	9.83	1.66	5.92
		60 o C -7d -NH 4 +	11.63	2.13	5.46
	681
	682	7 6 5 4	3	d spacing (Å)	2

Table S1 .

 S1 Experimental conditions on formation and transformation of schwertmannite

	22	through direct Fe 3+ hydrolysis
		Sample	Fe 3+ (mM)	SO 4 2-(mM)	pH	Temperature	Co-existing ions (mM)	Dialysis time
		(a) Schwertmannite formation though Fe 3+ hydrolysis-dialysis pathway
		Hydrolysis temperature	19.98	10.56	-	25, 40, 50, 60, 70, 80 o C	-	7d
		Dialysis time 19.98	10.56	-	25, 60 o C	-	1, 3, 7, 15 d
		Coexistence of K + or NH 4 + 19.98	10.56	-	60 o C	K + , NH 4 + (21.12 mM)	7d
		(b)

Mineral evolution during Fe 3+ hydrolysis by adding NaOH

	OH -/Fe 3+ = 1	400	600	2.5	25 o C	-	-
	Fhy + SO 4 2-	48.55	72.825	2.5	25 o C	-	-
	(c)

Fe 3+ hydrolysis rate on the formation and long-term aging of schwertmannite

	Aging temperature	48.55		24.275	3.0	25, 60, 80 °C	-	-
	Co-existing Fe 2+	48.55		24.275	3.0	Aging: 60	°C	Fe 2+ (4.855 mM)	-
	Co-existing Cl -	48.55		24.274	3.0	Aging: 60	°C	Cl -(24.275, 242.75 mM)	-
	Co-existing K + or NH 4 +	48.55		24.275, 247.25	2.0	Aging: 80	°C	K + , NH 4 + mM) (48.55, 485.5	-
	23
				Fe-O	v 4
	1200 1100 1000	900	800	700	600
	24
	33.33 µM/min	48.55		24.275	3.0	Formed:25 Aging: 60	°C °C	-	-
	6.67 µM/min 48.55		24.275	3.0	Formed:25 Aging: 60	°C °C	-	-
	3.33 µM/min 48.55		24.275	3.0	Formed:25 Aging: 60	°C °C	-	-
	(d) Schwertmannite transformation under various geochemical conditions
						2.0,
	pH effects	48.55		24.275	2.5, 3.0,	Aging: 80 °C	-	-
						3.5
					32.37,
	Fe/S molar ratios	48.55		24.275, 9.71, 6.07, 19.42,	3.0	Aging: 60 °C	-	-
					4.855

[Type here]

2-(mM)

Acknowledgements

448

This study was supported by the National Natural Science Foundation of China

Electronic supplementary information


Chaoyun Ying

Bruno Lanson

Cheng Wang

Xiaoming Wang

Hui Yin

Yupeng Yan

Wenfeng Tan

Fan Liu

Xionghan Feng

Highly enhanced oxidation of arsenite at the

Keywords: birnessite, arsenite oxidation, pyrophosphate, Mn(III) chelating, depassivation

Introduction

In natural waters, arsenite, As(III), mainly occurs as the neutral species H 3 AsO 3 , while arsenate, As(V), is present as H 2 AsO 4 -or HAsO 4 2-oxyanions under a variety of pH conditions [START_REF] Cerkez | Coupled redox transformation of chromate and arsenite on ferrihydrite[END_REF]. Both forms are highly toxic and carcinogenic, and the maximum contaminant level for arsenic in drinking water defined by the World Health Organization (WHO) is thus extremely low at 10 µg/L. Arsenite is even more toxic and mobile than arsenate in aqueous environments [START_REF] Cerkez | Coupled redox transformation of chromate and arsenite on ferrihydrite[END_REF][START_REF] Fischel | The influence of environmental conditions on kinetics of arsenite oxidation by manganese-oxides[END_REF][START_REF] Neppolian | Sonochemical oxidation of arsenic(III) to arsenic(V) using potassium peroxydisulfate as an oxidizing agent[END_REF] and more difficult to remove owing to its low affinity for sorbents [START_REF] Luong | Iron-based subsurface arsenic removal technologies by aeration: a review of the current state and future prospects[END_REF]. Oxidizing As(III) into the more easily extractable As(V) thus appears appropriate to achieve efficient arsenic immobilization and removal [START_REF] Zhang | Simultaneous oxidation and adsorption of As (III) from water by cerium modified chitosan ultrafine nanobiosorbent[END_REF].

Although As(III) is stable in homogeneous systems [START_REF] Katsoyiannis | Application of biological processes for the removal of arsenic from groundwaters[END_REF][START_REF] Tsang | Determination of phosphate/arsenate by a modified molybdenum blue method and reduction of arsenate by S(2)O(4)(2-)[END_REF], a variety of metal oxides such as iron (hydr)oxides [START_REF] Amirbahman | Kinetics of sorption and abiotic oxidation of arsenic(III) by aquifer materials[END_REF][START_REF] Wang | Effects of pH, dissolved oxygen, and aqueous ferrous iron on the adsorption of arsenic to lepidocrocite[END_REF][START_REF] Yang | Adsorption of As (III) on porous hematite synthesized from goethite concentrate[END_REF][START_REF] Zhao | Adsorption and heterogeneous oxidation of As (III) on ferrihydrite[END_REF] and titanium oxides [START_REF] Dutta | Adsorption of arsenate and arsenite on titanium dioxide suspensions[END_REF][START_REF] Ferguson | TiO2-photocatalyzed As (III) oxidation in aqueous suspensions: reaction kinetics and effects of adsorption[END_REF][START_REF] Guan | Application of titanium dioxide in arsenic removal from water: a review[END_REF] can induce its oxidation directly or indirectly. As a consequence, As concentration in natural aqueous environments is essentially controlled by interactions with mineral surfaces. Composite oxides have also been developed to enhance their efficiency as oxidants and adsorbents for both As(III) and As(V) [START_REF] Chakravarty | Removal of arsenic from groundwater using low cost ferruginous manganese ore[END_REF]Feng et al., 2006a;[START_REF] Ma | Arsenite simultaneous sorption and oxidation by natural ferruginous manganese ores with various ratios of Mn/Fe[END_REF][START_REF] Mccann | In situ arsenic oxidation and sorption by a Fe-Mn binary oxide waste in soil[END_REF][START_REF] Wu | Iron and arsenic speciation during As(III) oxidation by manganese oxides in the presence of Fe(II): Molecular-level characterization using XAFS, Mössbauer, and TEM analysis[END_REF][START_REF] Ying | Oxidation and competitive retention of arsenic between iron-and manganese oxides[END_REF][START_REF] Zhang | Removal mechanism of As(III) by a novel Fe-Mn binary oxide adsorbent: Oxidation and sorption[END_REF][START_REF] Zhang | Enhanced removal of arsenite and arsenate by a multifunctional Fe-Ti-Mn composite oxide: photooxidation, oxidation and adsorption[END_REF][START_REF] Zheng | As(III) adsorption on Fe-Mn binary oxides: Are Fe and Mn oxides synergistic or antagonistic for arsenic removal?[END_REF]. Manganese (Mn) oxides can also oxidize As(III) [START_REF] Manning | Arsenic (III) oxidation and arsenic (V) adsorption reactions on synthetic birnessite[END_REF][START_REF] Scott | Reactions at oxide surfaces. 1. Oxidation of As(lll) by synthetic birnessite[END_REF], as reported in natural lacustrine environments [START_REF] Oscarson | The oxidation of arsenite by aquatic sediments[END_REF](Oscarson et al., , 1981a(Oscarson et al., , 1981b)), and Mn oxides have a key role in As geochemical cycling [START_REF] Driehaus | Oxidation of arsenate(III) with manganese oxides in water treatment[END_REF][START_REF] Nesbitt | XPS study of reductive dissolution of 7Å-birnessite by H 3 AsO 3 , with constraints on reaction mechanism[END_REF][START_REF] Scott | Reactions at oxide surfaces. 1. Oxidation of As(lll) by synthetic birnessite[END_REF][START_REF] Tournassat | Arsenic(III) oxidation by birnessite and precipitation of manganese(II) arsenate[END_REF]. Produced Mn(II) adsorbs to the particle edges, thus potentially blocking reactive sites for further oxidation [START_REF] Villalobos | The influence of particle size and structure on the sorption and oxidation behavior of birnessite: I. Adsorption of As(V) and oxidation of As(III)[END_REF]. Lafferty and coworkers (Lafferty et al., 2010a(Lafferty et al., , 2010b[START_REF] Lafferty | Arsenite oxidation by a poorly-crystalline manganese oxide. 3. Arsenic and manganese desorption[END_REF] identified Mn(II) as the unique reduced product during the initial phase of the reaction.

They also concluded that Mn(III) was subsequently formed as the result of the comproportionation of Mn(II) adsorbed onto birnessite surface rather than from Mn(IV) reduction, as previously hypothesized [START_REF] Nesbitt | XPS study of reductive dissolution of 7Å-birnessite by H 3 AsO 3 , with constraints on reaction mechanism[END_REF]. Manganite (-Mn(III)OOH) then accumulates at the birnessite surface leading to its passivation, as Mn(III) sites are less reactive than Mn(IV) ones with respect to As(III) oxidation, as shown by DFT calculations [START_REF] Zhu | Quantum chemical study of arsenic (III, V) adsorption on Mn-oxides, implications for arsenic(III) oxidation[END_REF]. Finally, Lafferty and coworkers showed that As(V) was the sole As species adsorbed to manganese oxides (Lafferty et al., 2010a(Lafferty et al., , 2010b[START_REF] Lafferty | Arsenite oxidation by a poorly-crystalline manganese oxide. 3. Arsenic and manganese desorption[END_REF].

In contrast to the optimization of the oxidizing solid, little attention has been paid

to the influence of additional chemicals present in the reactive medium. For example, the presence of phosphate in solution hampers As(III) oxidation by Mn oxides [START_REF] Chiu | Arsenic adsorption and oxidation at manganite surfaces. 1. Method for simultaneous determination of adsorbed and dissolved arsenic species[END_REF][START_REF] Parikh | Evaluating environmental influences on As III oxidation kinetics by a poorly crystalline Mn-oxide[END_REF], whereas that of tartaric acid promotes it (Feng et al., 2006b). More generally, the presence in the reactive medium of compounds chelating Mn, and more especially Mn(III), is expected to enhance the oxidative process [START_REF] Ma | Arsenite simultaneous sorption and oxidation by natural ferruginous manganese ores with various ratios of Mn/Fe[END_REF]. With this respect, pyrophosphate (PP) appears especially relevant owing to its strong affinity for Mn(III) [START_REF] Liu | Influence of pyrophosphate on the generation of soluble Mn(III) from reactions involving Mn oxides and Mn(VII)[END_REF][START_REF] Marafatto | Crystal growth and aggregation in suspensions of δ-MnO2 nanoparticles: implications for surface reactivity[END_REF][START_REF] Parker | Manganese (III) binding to a pyoverdine siderophore produced by a manganese (II)-oxidizing bacterium[END_REF][START_REF] Soldatova | Mn(II) oxidation by the multicopper oxidase complex Mnx: A coordinated two-stage Mn(II)/(III) and Mn(III)/(IV) mechanism[END_REF]. The relevance of PP is further increased by its common formation in natural environments, as the simplest polymer of orthophosphate resulting from the breakdown of ATP and ADP (Klewicki and Morgan, 1999b;[START_REF] Trouwborst | Soluble Mn(III) in suboxic zones[END_REF].

The present experimental work thus investigated in detail the influence of PP on As(III) oxidation by the Mn oxide birnessite at circumneutral pH. The effect of PP addition on the completeness of the reaction and the involved reaction mechanisms were specifically studied from a combination of wet chemical analyses and of solid phase characterization using electron microscopy, X-ray diffraction, X-ray absorption and photoelectron spectroscopies, and infrared spectroscopy.

Materials and Methods

Chemicals

All chemical used in the present study were purchased from Sinopharm Chemical Reagent, except for the manganese(III) acetate dihydrate which was purchased from Sigma Chemical. All chemicals were of analytical grade. Atomic absorption spectroscopy Mn standard was prepared by dilution of a 1000 mg/L standard. Deionized water was used throughout the experiments.

Synthetic Acid Birnessite Preparation

Synthetic acid birnessite was prepared according to a method modified from that of [START_REF] Mckenzie | The synthesis of birnessite, cryptomelane, and some other oxides and hydroxides of manganese[END_REF]. In brief, 45 mL of a 6 M HCl solution were added dropwise to 300 mL of a boiling 0.667 M solution of KMnO 4 . The obtained suspension was stored at 60 °C for 12 h to increase birnessite crystallinity. Resulting solid was then washed with deionized water and centrifuged (10×) to remove K + and Cl -in excess.

Batch As(III) Oxidation by Birnessite and Wet Chemical Analyses

Batch experiments were conducted using 50 mL of a 0.2 g/L suspension of the obtained synthetic birnessite at two As(III) concentrations (0.5 and 1 mM) in the presence/absence of sodium pyrophosphate (PP). When present, PP concentrations were 1, 2, 3, 4, 5, 8, and 10 mM, with a 2.5 mM NaCl background electrolyte. Mixed suspensions, whose pH was pre-adjusted to 7.2 using 0.2-1 mM HCl, were then stirred for 24 h. The solids were subsequently filtered and washed thoroughly with deionized water to remove residual dissolved ions.

In the absence of PP, As(V) was determined using the ammonium molybdenum method [START_REF] Oscarson | The oxidation of arsenite by aquatic sediments[END_REF]. In the presence of PP, As(III) and As(V) were determined using coupled LC-ICP-MS, As(III) and As(V) being separated by liquid chromatography (LC -Dionex IonPac AS19 Columns) using a four-component mobile phase (pH 10.22) at 1 mL/min flow rate. The outflow from the LC was directly connected to the inductively coupled plasma mass spectrometer (ICP-MS) for

As quantification. In specific experiments, PP was replaced by other chelating agents, such as sodium oxalate. In this case, As(V) and total As concentrations were measured using the molybdenum blue method [START_REF] Feng | A quantitative model for the coupled kinetics of arsenic adsorption/desorption and oxidation on manganese oxides[END_REF], the concentration of As(III)

being calculated from that of total As by subtracting As(V).

Mn concentration in solution was determined using atomic absorption spectrometry (AAS -Agilent Technologies 200 series). In addition, concentration of Mn(III)-PP complexes were determined with a UV-Vis spectrophotometer (Agilent Technologies Cary 8454), using 10 mm path length cuvettes and the absorption peaks at 258 nm and 480 nm [START_REF] Webb | Evidence for the presence of Mn(III) intermediates in the bacterial oxidation of Mn(II)[END_REF].

Characterization of the Solid Products.

Mn K-edge XANES and Mn speciation analysis

X-ray absorption near edge structure (XANES) spectroscopy data were collected at room temperature on the 1W1B beamline at the Beijing Synchrotron Radiation Facility (BSRF) over the 6.4-7.0 keV range (Mn K-edge). Energy calibration was systematically performed using a Mn metal foil before data collection. Mn K-edge spectra were averaged and background-subtracted using the following parameters: E 0 = 6538 eV, Rbkg = 1 Å and k-weight = 2.

Field emission scanning electron microscopy

Gold-coated samples were observed using a field-emission scanning electron microscope (FESEM -SU8010, Hitachi) with a maximum resolution of 1 nm. For high-resolution images, the microscope was operated at 10 kV using a working distance of 10-15 mm, and an in-lens secondary electron detector.

Synchrotron radiation X-ray powder diffraction

Synchrotron radiation X-ray powder diffraction (SR-XRD) was performed on the BL14B1 beamline of the Shanghai Synchrotron Radiation Facility (SSRF). Data were collected over the 2-45 ° 2θ range (λ = 0.6895 Å), with 30 s exposure times [START_REF] Yang | Introduction of the X-ray diffraction beamline[END_REF].

Fourier transform infrared (FTIR) spectroscopy

FTIR spectra were collected using a Bruker Vertex 70 spectrometer equipped with a deuterated triglycine sulfate detector (Bruker Optics Inc., Ettlingen, Germany).

Spectra were measured over the 4000-400 cm -1 range with a 4 cm -1 resolution in transmission mode. Thirty-two scans were collected and averaged for each sample, the data being collected, processed, and analyzed with the OPUS program.

X-ray photoelectron spectroscopy

X-ray photoelectron spectra (XPS) were collected using a VG Multilab2000 X-ray photoelectron spectrometer with an Al K X-ray source (1486 eV) and a base pressure of 3×10 -9 Torr in the analytical chamber. The scans were recorded using the large area mode. The survey scans were collected using a fixed pass energy of 100 eV and an energy step size of 1.0 eV, whereas the narrow scans have a pass energy of 25 eV and an energy step size of 0.1 eV. Spectra were charge-corrected to C1s with a binding energy of 284.80 eV collected from the surface adventitious carbon [START_REF] Wang | Study on the shape control and photocatalytic activity of high-energy anatase titania[END_REF] and analyzed with the Avantage software. The parameters proposed by Ilton et al. (2016) and Bang et al. (2005) for the spectral fitting of Mn3p and As3d multiple peaks were used.

Results and Discussion

The Oxidation of As(III) by Birnessite with/without Chelating Agents

Based on their relative redox potentials, As(III) could be oxidized by dissolved oxygen under circum-neutral conditions (△E for O 2 /H 2 O and As(V)/As(III) couples is 0.70 V at neutral pH). However, As(III) is mostly hydrophilic and present as neutral species below pH 9, and the kinetics of its oxidation by O 2 is very slow at neutral pH [START_REF] Lan | Efficient catalytic As(III) oxidation on the surface of ferrihydrite in the presence of aqueous Mn(II)[END_REF]. Consistently, the control experiments performed in the absence of manganese oxides showed that As(III) was essentially stable over a 24 h period (Fig. S1).

In the absence of PP and under similar circum-neutral conditions, birnessite was able to oxidize ~ 35 % of the As(III) initially present (1 mM) to As(V) after 24 h (Fig. 1a), whereas ~10 % of As adsorbed/precipitated at the birnessite surface. The As(III) concentration dropped exponentially with time over the first 2 h. The application of a pseudo-first order model yielded an apparent As(III) oxidation rate constant (k) of 0.1804 ± 0.0349 h -1 (k 1 ) over this 2 h interval (Fig. 1b). The reaction slowed down over the subsequent 10 h with an apparent rate constant decreased by about one order of magnitude (k 2 = 0.0187 ± 0.0015 h -1 ), most likely as the result of the formation of a Mn precipitate passivating birnessite surface (Lafferty et al., 2010a). After 12 h, As(III)

was not oxidized by birnessite anymore, consistent with the following equations proposed for As(III) oxidation (Lafferty et al., 2010a(Lafferty et al., , 2010b)):

>Mn(IV)-OH + H 3 As(III)O 3 + H + → Mn(II) + H 2 As(V)O 4 -+ H 2 O (1) >Mn(IV)-OH + Mn(II)+ H 2 O + H + → 2 (>Mn(III)-OH 2 ) (2)
According to these equations, As(III) oxidation occurs as a two-electron transfer process, the formation of Mn(III) resulting from the comproportionation of surface-adsorbed Mn(II) and structural Mn(IV) (Eq. 2). The proportion of As(III) oxidized by birnessite over 24 h was doubled in the presence of 5 mM PP in solution compared to the PP-free system (Fig. 1a). In the presence of PP, As(III) oxidation was not significantly promoted over the first 2 hours, the apparent rate constant being 0.1993 ± 0.0125 h -1 (k 3 , compared to k 1 = 0.1804 ± 0.0349 h -1 -Fig. 1b). Over the next 10 hours, As(III) oxidation was highly enhanced in the presence of 5 mM PP, however, with an apparent rate constant of 0.0438 ± 0.0101 h -1 (k 4 ). Moreover, addition of PP allowed releasing to solution As species initially adsorbed/precipitated at the birnessite surface (Fig. 1a). Enhancement of As(III) oxidation results from the chelation by PP of Mn(III) formed as the result of Eq. 2 and from the release of these complexes to solution (Eq. 3 -Klewicki and Morgan, 1999b):

Mn(III)OOH + HP 2 O 7 3- + 3 H + → (Mn(III)HP 2 O 7 ) aq + 2 H 2 O (3)
Previous works showed that the excess of PP relative to Mn(III) should be at least fourfold to prevent the complexes from disproportionating [START_REF] Qian | Geochemical stability of dissolved Mn(III) in the presence of pyrophosphate as a model ligand: complexation and disproportionation[END_REF].

Thus, the difference between k 4 and the oxidation rate determined over the subsequent 12 hours (k 5 = 0.0091 ± 0.0025 h -1 ) is likely related to the formation of Mn(III) that could not be fully chelated by PP (Fig. 1b).

Mn concentration in solution increased steadily in the presence of PP (Fig. 2a), nevertheless, as shown by the increase in intensity of the 258 nm UV-Vis absorption band (Fig. 2b) that is characteristic of the Mn(III)-PP complex [START_REF] Soldatova | Mn(II) oxidation by the multicopper oxidase complex Mnx: A coordinated two-stage Mn(II)/(III) and Mn(III)/(IV) mechanism[END_REF].

Consistently, the filtered solution appeared brownish red (Fig. S2), a color typical for Mn(III)-PP [START_REF] Qian | Geochemical stability of dissolved Mn(III) in the presence of pyrophosphate as a model ligand: complexation and disproportionation[END_REF]. High-valent Mn in solution, which hereafter refers to Mn(III)-PP, was quantified using Leucoberbelin Blue I [START_REF] Zhu | Rapid determination of the Mn average oxidation state of Mn oxides with a novel two-step colorimetric method[END_REF]. The similar increases of aqueous Mn and of Mn(III)-PP indicated that Mn in solution was overwhelmingly trivalent and present as Mn(III)-PP (Fig. 2a), consistent with XANES results (Fig. S3). In the presence of both As(III) and As(V), Mn(III)-PP complexes remained essentially stable over a 24 h interval in circum-neutral conditions (Table S1 and Fig. S4), consistent with the slow disproportionation reported for Mn(III)-PP at circum-neutral pH [START_REF] Qian | Geochemical stability of dissolved Mn(III) in the presence of pyrophosphate as a model ligand: complexation and disproportionation[END_REF]. The minor decrease of Mn(III)-PP after 12 h is possibly related to Mn oxides catalyzing PP hydrolysis [START_REF] Wan | Manganese Oxide Catalyzed Hydrolysis of Polyphosphates[END_REF].

The premise from Eq. 3 was supported by the experimental ratio between As(V), produced from Eq. 1, aqueous Mn(III), produced from Eq. 2 and complexed by PP, that steadily approached a 1:2 ratio with increasing PP in the presence of 0.5 mM As(III) (Fig. S5), consistent with the stoichiometric ratio from Eqs. 1, 2. At low PP concentration, the experimental ratio was significantly larger than 1:2 due to the limited formation of Mn(III)-PP complexes. Increasing PP concentrations increased Mn(III)-PP concentration in solution, birnessite being essentially dissolved over 24 h in the presence of 10 mM PP and 1 mM As(III) (Fig. 3). The stoichiometric 1:2 ratio between As(V) and Mn indicates that the oxidation of 1 mM As(III) required the consumption of 2 mM Mn(IV), which exceeded the total Mn content in the suspension. As a result, As(V) in solution never reached 1 mM even in the presence of 10 mM PP.

In addition, time-resolved As K-edge XANES spectroscopy was used to determine the overall relative proportions of As(III) and As(V) in the reaction system (Fig. S6). The position of As K absorption edge steadily shifted to higher energies with increasing reaction time, showing the formation of As(V) at the expense of As(III). Linear combination fitting (LCF) of the spectra unraveled the in-situ evolution kinetics of As(III) and As(V) species. The kinetic of As(III) oxidation using pseudo-first order model was linear over the first 180 min in the presence of PP (apparent rate constant k = 0.0049 ± 0.0001 h -1 -Fig. S7) and of the induced alleviation of surface passivation.

Identification of the Solid Products

Birnessite synthesized in the present study exhibited a uniform nano-flower morphology resulting from the assemblage of nano-flakes as revealed by FESEM (Fig. 4). When reacted with As(III) in the absence of PP, these nano-flowers appeared "filled" after 24 h of reaction, nano-flake edges becoming blurred, owing to the presence of precipitates at the particle surface. Increasing PP concentration in solution appeared to prevent the presence of these precipitates, leaving clean birnessite surfaces and thus significantly alleviating the passivation of birnessite surface.

SR-XRD patterns of the unreacted birnessite exhibited reflections at 0.723 (not shown), 0.361, 0.244, and 0.142 nm typical for birnessite (Fig. 5a). After reacting with 1 mM As(III) for 24 h in the absence of PP, additional reflections typical for manganite (ICDD#41-1379) were visible (Fig. 5a). Manganite precipitation was consistent with previous reports of birnessite interactions with aqueous Mn(II) at circum-neutral pH [START_REF] Elzinga | Reductive transformation of birnessite by aqueous Mn(II)[END_REF][START_REF] Lefkowitz | Influence of pH on the reductive transformation of birnessite by aqueous Mn(II)[END_REF][START_REF] Tournassat | Arsenic(III) oxidation by birnessite and precipitation of manganese(II) arsenate[END_REF].

Manganite peaks were not visible when PP was present in the reacting suspension, however (Fig. 5a). None of the various manganese arsenate compounds with contrasting As: Mn molar ratios reported in the literature as secondary products of As(III) oxidation by birnessite [START_REF] Moore | Reaction scheme for the oxidation of As(III) to As(V) by birnessite[END_REF][START_REF] Tournassat | Arsenic(III) oxidation by birnessite and precipitation of manganese(II) arsenate[END_REF] were detected either in the present study; pyrochroite (Mn(OH) 2 -ICDD#01-73-1133) was not observed either.

The 920 cm -1 band in the FT-IR spectrum of unreacted birnessite corresponded to the bending vibration of -OH located at vacancy sites (Fig. 5b - [START_REF] Yin | Distinct effects of Al 3+ doping on the structure and properties of hexagonal turbostratic birnessite: A comparison with Fe3+ doping[END_REF][START_REF] Zhao | Fourier transform infrared spectroscopy study of acid birnessites before and after Pb 2+ adsorption[END_REF]. After reaction with As(III) for 24 h in the absence of PP, new bands were visible at 1150, 1116, and 1087 cm -1 that were related to in-plane and out-of-plane bending modes of structural OH in manganite [START_REF] Kohler | Hydrogen bonding and Jahn-Teller distortion in groutite, α-MnOOH, and manganite, γMnOOH, and their relations to the manganese dioxides ramsdellite and pyrolusite[END_REF][START_REF] Lefkowitz | Influence of pH on the reductive transformation of birnessite by aqueous Mn(II)[END_REF], consistent with the presence of this phase shown by XRD data. Consistent with XRD data, these bands, typical for manganite, were not visible when PP was added to the reactive system.

In addition, XPS allowed identifying the chemical state of elements in the few uppermost atomic layers of solids and monitoring the evolution of Mn oxidation state at the birnessite surface. Mn3p XPS spectra were fitted using scale factors for the three multiplets used (Table S2) (Ilton et al., 2016), and the relative contributions of the different Mn oxidation states were quantified from the overall multiplet area (Fig. 6)

(Table 1). The best fit to Mn3p spectrum of unreacted birnessite yielded 67% Mn(IV), 29% Mn(III), and 4% Mn(II), and the relative proportion of Mn(IV) at birnessite surface decreasing to 41% after reaction with As(III) in the absence of PP.

Simultaneously, Mn(III) and Mn(II) increased to 50% and 9%, respectively, as the result of As(III) oxidation. The significant increase of Mn(III) was consistent with the key role played by Mn(IV) in the oxidation of As(III) and with the observed formation of manganite at the particle surface (Figs. 4,5). If PP was added to the reactive system, the relative content of Mn(III) at the birnessite particle surface decreased (Table 1), consistent with increased proportion of Mn(III)-PP in solution and the absence of manganite formation. Enhancement of As(III) oxidation thus appeared to be induced by the absence of birnessite surface passivation, resulting from manganite precipitation, which was caused by Mn(III) complexation by PP. This enhancement was consistent with theoretical calculations showing the higher affinity of both As(III) and As(V) for Mn(IV) sites compared to Mn(III) sites and with the increased electron transfer rate between As(III) and Mn(IV), compared to Mn(III) [START_REF] Zhu | Quantum chemical study of arsenic (III, V) adsorption on Mn-oxides, implications for arsenic(III) oxidation[END_REF]. Finally, XPS As3d spectra were characteristic of coexisting As(V) and As(III) species (Fig. S8), whose binding energies were 45.5 ± 0.1 eV and 44.2 ± 0.1 eV, respectively (Table S3 - Bang et al., 2005). In the absence of PP, adsorbed/precipitated As was essentially present as As(V), consistent with previous reports (Fig. S8 - Lafferty et al., 2010b).

The Effects of Different Chelating Agents on As(III) Oxidation

Naturally occurring organic ligands other than PP, such as oxalate and citrate, are also known to chelate Mn [START_REF] Duckworth | Siderophore-manganese(III) interactions. I. Air-oxidation of manganese(ll) promoted by desferrioxamine B[END_REF][START_REF] Jiang | The reduction of Cr(VI) to Cr(III) mediated by environmentally relevant carboxylic acids: State-of-the-art and perspectives[END_REF]Klewicki and Morgan, 1999a;[START_REF] Mu | Mn 2+ promoted Cr(VI) reduction with oxalic acid: The indispensable role of In-situ generated Mn 3+[END_REF][START_REF] Taube | Catalysis of the reaction of chlorine and oxalic acid. Complexes of trivalent manganese in solutions containing oxalic acid[END_REF][START_REF] Zhang | Enhanced degradation of methyl parathion in the ligand stabilized soluble Mn(III)-sulfite system[END_REF]. As for PP, the presence of oxalate in the reactive system promoted As(III) oxidation (by ~15% after leucoberbelin blue dye [START_REF] Zhu | Rapid determination of the Mn average oxidation state of Mn oxides with a novel two-step colorimetric method[END_REF], however. The limited enhancement of As(III) oxidation by oxalate was most likely linked to the reduced release of Mn(III) in solution, and thus to the lower ability of oxalate, compared to PP, to prevent surface passivation by chelating Mn(III) formed during the reaction (Eqs. 1, 2; Fig. 7b). The lower chelating activity of oxalate, compared to PP, was consistent with the respective stability constants of these complexing agents with Mn(III) (Log K are 9.98 and 11.7

for oxalate and PP, respectively) [START_REF] Nico | Mn(III) center availability as a rate controlling factor in the oxidation of phenol and sulfide on δ-MnO 2[END_REF]. Consistently, the first-order rate constant determined for the breakdown of Mn(III)-oxalate complexes (Mn(III)(C 2 O 4 ) + , k = 1.97 × 10 -1 s -1 ) was six orders of magnitude higher than that of Mn(III)-PP (Mn(III)HP 2 O 5 , k = 3 × 10 -7 s -1 ) under neutral conditions [START_REF] Nicholson | Geochemistry of manganese oxides: metal adsorption in freshwater and marine environments[END_REF][START_REF] Taube | Catalysis of the reaction of chlorine and oxalic acid. Complexes of trivalent manganese in solutions containing oxalic acid[END_REF][START_REF] Zhang | Enhanced degradation of methyl parathion in the ligand stabilized soluble Mn(III)-sulfite system[END_REF]. Moreover, organic ligands, such as oxalate, are considered as reducing agents interacting with Mn (hydr)oxides [START_REF] Flynn | Reductive transformations of layered manganese oxides by small organic acids and the fate of trace metals[END_REF][START_REF] Wang | Structural transformation of birnessite by fulvic acid under anoxic conditions[END_REF][START_REF] Wang | Reaction of Mn III,IV (hydr)oxides with oxalic acid, glyoxylic acid, phosphonoformic acid, and structurally-related organic compounds[END_REF]. Therefore, the ligands with both low reducing ability (or no redox activity) and high complexation ability with Mn(III), such as PP [START_REF] Klewicki | Kinetic Behavior of Mn(III) Complexes of Pyrophosphate, EDTA, and Citrate[END_REF], appear as ideal inhibitors of Mn(III) (hydr)oxide formation during reaction with As(III).

Environmental significance

Birnessite is ubiquitous in aquatic and terrestrial environments [START_REF] Butterfield | Mn(II,III) oxidation and MnO 2 mineralization by an expressed bacterial multicopper oxidase[END_REF], where As(III) potentially coexists with pyrophosphate (PP). The major source of PP in the environment appears to be ATP released from cells that can be hydrolyzed quickly to produce PP [START_REF] Orriss | Pyrophosphate: a key inhibitor of mineralisation[END_REF]. The above results provide a mechanistic understanding for the enhancement of As(III) oxidation by birnessite in the presence of PP through alleviation of surface passivation when the three species coexist.

As As(III) oxidation to As(V) is an effective pathway for arsenic detoxification in arsenic-containing drinking and waste water, birnessite is a potentially attractive oxidant to be used in large-scale treatment. Previous studies have shown that the presence of aqueous Fe(II) and Mn(II) may inhibit As(III) oxidation by birnessite [START_REF] Gude | As(III) oxidation by MnO 2 during groundwater treatment[END_REF], however, as these species are oxidized at the birnessite surface leading to its passivation and favoring As(III) [START_REF] Ehlert | Impact of birnessite on arsenic and iron speciation during microbial reduction of arsenic-bearing ferrihydrite[END_REF]. In the presence of PP, precipitation of Mn(III) on the surface of birnessite will likely be removed, thus alleviating surface passivation and enhancing As(III) oxidation. Hydrous ferric oxides, such as ferrihydrite [START_REF] Ehlert | Impact of birnessite on arsenic and iron speciation during microbial reduction of arsenic-bearing ferrihydrite[END_REF], formed at the birnessite surface are also likely to be dissolved in the presence of PP [START_REF] Kassim | Ferrihydrite in pyrophosphate extracts of podzol B horizons[END_REF], thus releasing adsorbed As(III) to solution and promoting its oxidative interaction with birnessite.

Conclusions

As previously reported, oxidation of As(III) in presence of birnessite and at circum-neutral pH conditions proceeds as a two-electron transfer. The subsequent formation of Mn(III) results from the comproportionation of Mn(II) and Mn(IV) and leads to the precipitation of manganite at the surface of the initial birnessite particles.

Manganite precipitation passivates birnessite surface and restrains As(III) oxidation completeness. In the presence of PP, As(III) oxidation is significantly enhanced at circum-neutral pH. The major influence of PP on As(III) oxidation results from its strong chelating affinity for Mn(III), causing the relative content of Mn(III) at the birnessite surface to decrease. PP chelates newly formed Mn(III) and thus inhibits the formation of solid Mn(III) phases, such as manganite, that may passivate birnessite surface. The addition of PP to the oxidative system has major implications both for the efficiency of As(III) removal during water treatment and for the durability of birnessite as oxidizing catalyzer. Compared to other Mn(III) chelating agents, such as oxalate, pyrophosphate appears especially efficient in impeding surface passivation and limiting possible reductive activity.

Declaration of interest statement

Declarations of interest: none. Table 1 Relative proportions of the different Mn oxidation states at the birnessite surface determined from the fits of Mn3p XPS spectra shown in Fig. 6.

Legends of figures and tables

Tables Table 1 Relative proportions of the different Mn oxidation states determined from the fits of Mn3p XPS spectra shown in Fig. 6.

Fig. 1

 1 Fig. 1 (a) Evolution of As(III), As(V), and total As concentrations in solution during

Fig. 3

 3 Fig. 3 As(III), As(V), and Mn concentrations in solution and the ratio between Mn and

Fig. 4

 4 Fig. 4 FESEM images of unreacted synthetic birnessite (Bir) and of reaction products

Fig. 6 Fig. 7

 67 Fig. 6 XPS spectra (Mn3p) of unreacted birnessite (Bir) and of solid reaction products

Fig. 1

 1 Fig. 1 (a) Evolution of As(III), As(V), and total As concentrations in solution during

FigureFig. 2

 2 Figure

FigureFig. 3

 3 Figure

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 4 Figure

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 5 Figure

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Fig. S2

 S2 Fig. S2 Evolution of solution color from As(III) oxidation experiments performed in

Fig. S3

 S3 Fig. S3 Mn K-edge XANES spectrum of the solution resulting from an As(III)

Fig. S4

 S4 Fig. S4 Evolution as a function of time of UV-Vis absorption spectra of a solution

Fig. S5

 S5 Fig. S5 As(III), As(V), and Mn concentrations in solution as a function of the

Fig

  Fig. S6 (a) Selected As K-edge XAS spectra used to determine the relative proportions

Fig. S7

 S7 Fig. S7 Kinetic analysis of As(III) removal deduced from the linear combination fitting

Fig. S8

 S8 Fig. S8 As3d XPS spectrum of reaction products from an As(III) oxidation experiment

Table Click here to download Table: Table1.docx

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Acknowledgments

This work was supported by National Key R&D Program of China (No. 2017YFD0200201), and the National Natural Science Foundation of China (NSFC Grant Nos. 41471194).

Appendix. Supporting Information

Supplementary Information includes Tables S1-S3, Figs. S1-S8.

Supporting information to

Highly enhanced oxidation of arsenite at the surface of birnessite in the presence of pyrophosphate and the underlying reaction mechanisms

Materials and Methods

In-situ quick As K-edge XANES spectroscopy.

In-situ quick As K-edge X-ray Absorption Near Edge Structure (XANES) spectroscopy data were collected at ~25 ℃ at the 1W2B beamline of Beijing Synchrotron Radiation Facility in fluorescence mode, to track the evolution of As(III) and As(V) relative proportions during the oxidation of 2.0 mM As(III) in the presence of 1.6 g/L birnessite and 5 mM PP at pH ~7.5. Experimental conditions were similar to those reported by [START_REF] Ginder-Vogel | Quantification of rapid environmental redox processes with quick-scanning X-ray absorption spectroscopy (Q-XAS)[END_REF] and could be described as follows. The 50 mL suspensions were stirred in 100 mL plastic beakers with a small hole (Ø = 1 cm) close to the bottom. The spectroscopy data were continuously collected as a function of reaction time between 11.67 and 12.46 keV. Elementary spectra were collected within 40 s each. The Athena program was used for background removal and linear combination fittings (LCF) for quantifying the speciation variations of As using As(III) and As(V) references.

Table S1

Evolution as a function of time of the solution UV-Vis absorbance at 480 nm from As(III) oxidation experiments in the presence of Mn(III)-PP (0.5 mM initial As(III) concentration, initial pH 7.2). Blank experiments without As(III) were also conducted. Samples 0.5 mM Mn(III)-PP (Abs480 nm)

1 mM Mn(III)-PP (Abs480