In Situ Nanoscale Imaging of Struvite Formation during the Dissolution

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In situ nanoscale imaging of struvite formation during the dissolution of natural brucite: implications for phosphorus recovery from wastewaters Jörn Hövelmann, and Christine V Putnis Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.6b04623 • Publication Date (Web): 13 Nov 2016 Downloaded from http://pubs.acs.org on November 14, 2016

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Environmental Science & Technology

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In situ nanoscale imaging of struvite formation during the dissolution of

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natural brucite: implications for phosphorus recovery from wastewaters

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Jörn Hövelmann1* and ChristineV. Putnis2,3

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German Research Centre for Geosciences (GFZ), Interface Geochemistry, 14473 Potsdam, Germany 2

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Institut für Mineralogie, University of Münster, 48149 Münster, Germany

The Institute for Geoscience Research (TIGeR), Department of Chemistry, Curtin University, Perth 6845, Australia

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*Corresponding author:

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Phone: +49 331 288-28703; E-mail: [email protected] (Jörn Hövelmann)

1 ACS Paragon Plus Environment


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Abstract

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As phosphorus (P) resources are diminishing, the recovery of this essential nutrient from wastewaters

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becomes an increasingly interesting option. P-recovery through the controlled crystallization of

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struvite (MgNH4PO4•6H2O), a potential slow-release fertilizer, is highly attractive, but costly if large

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amounts of Mg have to be added. In this context, natural Mg-minerals like brucite (Mg(OH)2) could

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provide more cost effective Mg-sources compared to high-grade Mg-compounds such as MgCl2. Here

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we used in situ atomic force microscopy (AFM) to study the interactions of ammonium phosphate

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solutions with brucite (001) cleavage surfaces. Brucite dissolution was strongly enhanced in the

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presence of H2PO4- ions, most likely due to the formation of negatively charged surface complexes.

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Simultaneously with brucite dissolution, we directly observed the formation of a new phase that was

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identified as struvite by Raman spectroscopy. Our results suggest that brucite dissolution and struvite

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precipitation were coupled at the mineral-fluid interface within a thin fluid boundary layer. An

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interpretation is proposed where the heterogeneous nucleation and growth of struvite occurs via a

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particle-mediated process involving the formation of primary nanoparticles, followed by their

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continuous aggregation, fusion and possible transformation to crystalline struvite. These observations

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have implications for the feasibility of using brucite in phosphorus recovery processes.


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Introduction:

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Phosphorus (P) is a vital nutrient for plant growth, but its availability to plants is often limited due to

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its adsorption on soil mineral surfaces (mainly Fe- and Al-(hydr)oxides1) or its precipitation in the

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form of sparingly soluble phosphates (e.g., Ca-phosphate2). Thus, the use of phosphate fertilizers is

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essential in modern agriculture to ensure adequate food production for a growing global population.

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However, global resources of phosphate ores for fertilizer production are finite and may run out by the

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end of this century3. There is a substantial ambiguity regarding the actual extent of global phosphate

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reserves. However, although the timing of a ʻpeak phosphorusʼ remains uncertain, there is no dispute

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about its limited resource availability considering the growing global demand for fertilizers4. At the

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same time, high levels of fertilizers applied to land increase the potential that excess amounts of P will

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be washed out into the groundwater causing problems such as eutrophication of streams, rivers and

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coastal regions, presenting a major environmental concern. Hence, a more sustainable P management

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becomes increasingly important. In recent years, a lot of research has focused on investigating

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effective routes for the recovery of P from various industrial, farm and municipal wastewater streams.

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In this context, the mineral struvite (MgNH4PO4•6H2O) has gained strong interest5,6,7. Its controlled

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precipitation from wastewaters does not only enable the safe disposal of nutrient-laden wastes, but

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could also contribute to the conservation of natural P resources, as recovered struvite crystals may be

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reused directly as an eco-friendly slow-release fertilizer8,9. However, despite these highly attractive

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prospects, such struvite-based P recovery methods are not yet widely adopted because high costs for

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wastewater pre-treatments often limit their economic efficiency. To ensure effective struvite

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crystallization, the solution must be slightly alkaline (pH ~8-9) and contain phosphate (PO43-),

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ammonium (NH4+) and magnesium (Mg2+) ions in close to equimolar (1:1:1) concentrations10. Thus,

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for most wastewater sources pH has to be increased, which is typically done by adding an alkali source

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such as sodium hydroxide (NaOH). Moreover, the addition of magnesium is generally required since

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most nutrient-rich wastewaters are deficient in Mg2+ relative to PO43- and NH4+. Most of the

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commercially available struvite recovery technologies use pure Mg salts such as MgCl2, MgSO4 and

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MgO6. However, these high-grade compounds are expensive and may contribute up to 75% of the



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overall costs, making large-scale applications uneconomical11. Therefore, employing cheaper Mg

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sources would be an effective way for cost reduction12.

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Comparatively inexpensive Mg-sources include seawater, bittern (a by-product of salt manufacture),

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low-grade caustic magnesia (MgO) as well as natural Mg-rich minerals12-22. The use of seawater and

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bittern, although principally feasible17,18, appears only economical when close to the sea since

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comparatively large volumes are required. On the other hand, the efficiency of solid Mg sources

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depends to a large degree on their availability, Mg content, solubility and reactivity19. The common

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mineral magnesite (MgCO3), for example, has a high Mg content (~30 wt%), but a low solubility (Ksp

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= 10-7.8 at 25°C) and extremely slow dissolution rate in water at ambient temperature23. Therefore,

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either high doses of magnesite are required or it has to be pre-treated by acid dissolution or thermal

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decomposition to make enough Mg available for struvite formation14,21. Brucite (Mg(OH)2) is another

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natural Mg-rich mineral. It is less abundant than magnesite, but has a higher Mg content (~40 wt%).

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Brucite is widely distributed as an accessory mineral in a variety of rock types. Significant quantities

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(up to 20 wt%) are, for example, found in ultramafic rocks such as peridotites or serpentinites24,25.

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These common rock types are extensively mined for their chromite and nickel ores, hence making

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brucite also an abundant component of ultramafic mining residues26. Most brucite deposits of

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economic interest are, however, hosted by metamorphosed carbonate rocks, e.g. dolomitic marbles27.

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Using brucite as a reactant for struvite recovery could have several advantages. First of all, brucite

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dissolves relatively fast in the near-neutral pH region28,29,30, hence it could readily be added in the solid

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form without prior dissolution in acid. Furthermore, brucite dissolution not only releases Mg2+ (Mg

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source), but also OH- (alkali source), which helps to neutralize acids so as to achieve the optimal pH

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values for struvite precipitation. At the same time, brucite particles could act as nucleation sites for

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struvite, hence reducing induction times and increasing crystallization rates20. On the other hand,

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extensive precipitation of struvite onto the dissolving brucite substrate may also lead to surface

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passivation that could ultimately reduce the efficiency of the coupled dissolution-precipitation

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process19.

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Huang and co-workers15 have recently evaluated the use of natural brucite for struvite precipitation to

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remove ammonium from wastewater generated in the separation process for rare-earth elements. The 4 ACS Paragon Plus Environment

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authors performed batch-type experiments using different brucite-to-wastewater ratios and

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demonstrated that 93-95% of ammonium could be removed as struvite within 12 h. They also report

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that the precipitates collected at the end of the experiments contained a large amount of non-reacted

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brucite that could be reused in subsequent experiments for further struvite precipitation.

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While the use of brucite as a reactive medium for struvite precipitation seems principally feasible, a

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detailed understanding of how exactly brucite interacts with ammonium and phosphate bearing

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solutions is currently lacking. In particular, there is no direct molecular-scale observations of the

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mechanistic and kinetic pathways of the coupled process of brucite dissolution and subsequent struvite

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precipitation. However, such information is crucial to fully assess the potential of using brucite as an

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alternative Mg source and may provide hints of how the operational parameters of the struvite

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recovery process could be optimized.

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In this study we investigate the interactions of brucite surfaces with ammonium phosphate solutions by

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in situ, real-time imaging at the nanometer scale using atomic force microscopy (AFM) in connection

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with a fluid reaction cell. The objective was to elucidate and quantify the effect of phosphate and

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ammonium on the dissolution kinetics of brucite and to characterize the coupling between brucite

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dissolution and the following nucleation and growth of struvite at different initial pH values (6 – 9.5),

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ammonium phosphate concentrations (5 – 100 mM) and N:P molar ratios (1:1 and 2:1). Our results

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have implications for the efficiency of using natural brucite as well as synthetic Mg(OH)2 in struvite

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recovery processes.

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Methodology:

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Brucite Specimen. Natural brucite from the Tallgruvan (Norberg, Sweden) was used for the

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experiments. The initial, essentially monomineralic brucite rock sample contained minor amounts of

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dolomite (CaMg(CO3)2), magnetite (Fe3O4) and pyroaurite (Mg6Fe2(CO3)(OH)16•4H2O), which were

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avoided during AFM specimen preparation. Only optically transparent brucite crystals were used.

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Immediately before each experiment a brucite crystal was cleaved parallel to the (001) cleavage plane

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to expose a fresh surface. The final dimensions of the brucite specimens used in the AFM experiments

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were ca. 3 x 3 x 0.2 mm. 5 ACS Paragon Plus Environment


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Solutions. Aqueous solutions of phosphate and ammonium (5-100 mM) were prepared by dissolving

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reagent grade salts (Sigma Aldrich) of ammonium dihydrogen phosphate (NH4H2PO4), diammonium

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hydrogen phosphate ((NH4)2HPO4) or sodium dihydrogen phosphate (NaH2PO4) into double-

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deionized water (resistivity > 18mΩ cm-1). Adjustments of pH were made using 0.1 M NaOH or HCl.

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Ammonium- and phosphate-free solutions with pH values and ionic strengths (adjusted with NaCl)

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similar to the ammonium phosphate solutions were also used in some experiments.

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Atomic Force Microscopy. Brucite (001) surfaces were imaged at room temperature (23 ± 1°C) using

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a Bruker Multimode Atomic Force Microscope (AFM) operating in contact mode. In situ experiments

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were performed within an O-ring sealed flow-through fluid cell from Digital Instruments (Bruker).

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Solutions were injected at regular time intervals between each scan (lasting ~1.5 min), giving an

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effective flow rate of 22 µL s-1. AFM images were collected using Si3N4 tips (Bruker, tip model NP-

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S20) with spring constants of 0.12 and 0.58 N m-1. Images were analyzed using the NanoScope

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Analysis software (version 1.50). Step retreat velocities or etch pit spreading rates (vs) were calculated

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measuring the length increase of etch pit step edges (s) per unit time in sequential images scanned in

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the same direction. For each experimental condition at least 5 different etch pits were analyzed in 2-10

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pairs of sequential images. Each etch pit spreading rate value thus represents an average of 10-50

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individual measurements.

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For several experiments, the outlet fluid was sampled, collecting a sequence of 2-4 aliquots of 8 ml

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(the outflow from four consecutive scans) that were later analyzed for magnesium concentration using

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ICP-OES (inductively coupled plasma – optical emission spectroscopy). The analytical uncertainties

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of the ICP-OES measurements were below 3% based on repeated measurements of aqueous standards.

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In some of the experiments scanning was stopped from time to time and the solution in the fluid cell

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was kept static for several minutes, hence, allowing the system to approach equilibrium.

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Following in situ AFM experiments, some samples were placed directly into a beaker filled with 10 ml

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of the solution used in that specific in situ experiment. After 12 – 36 hours of reaction the samples

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were recovered from the solution and immediately dried by absorbing the fluid with filter paper. The

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crystals from these ex situ experiments were then re-examined in air in the AFM.


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Scanning Electron Microscopy. Samples from the ex situ experiments were also imaged using an

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Ultra 55 Plus (Carl Zeiss SMT) scanning electron microscope (SEM) equipped with an energy

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dispersive X-ray (EDX) detector for elemental analyses of the reacted brucite surfaces and newly

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formed precipitates. Before imaging, all samples were coated with a 20-nm-thick layer of carbon.

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Raman Spectroscopy. A confocal Raman spectrometer (Horiba Jobin Yvon XploRA) operating with

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the 638 nm line of a He–Ne laser was used for the analysis of surface precipitates on brucite after

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contact with ammonium phosphate bearing solutions. Spectra were taken with a 500 µm hole, 100 µm

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slit and 1200 grooves per millimeter grating using an acquisition time of 2 x 60 s. Corrections for

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system drift were made using the 520.7 cm-1 Raman band of a silicon standard taken at the beginning

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and the end of the Raman session. Reference spectra of struvite were obtained from the RRUFF

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database (http://www.rruff.info).

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Geochemical Modelling. The hydro-geochemical software PHREEQC31 (version 3.2.0-9820) was

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used to calculate the chemical speciation of the initial solutions used in the AFM experiments as well

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as to simulate the reactions with brucite. All calculations were done using the sit.dat database (version

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9a, July 2014, www.thermochimie-tdb.com) that was supplemented with thermodynamic data for

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struvite taken from Bhuiyan et al.32.

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Results and Discussion:

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Dissolution Features. In situ AFM observations show that dissolution of brucite (001) cleavage

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surfaces is very slow when in contact with pure water or moderately saline NaCl solutions. In 100 mM

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NaCl (pH 7) dissolution only occurred by the retreat of pre-existing step edges. However, no etch pits

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were observed to form, even after a contact time of 22 min (Fig. 1A). Dissolution then rapidly

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increased upon injection of a 10 mM (NH4)2HPO4 solution (pH 7.9) with the immediate formation and

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spreading of equilateral triangular etch pits (Fig. 1B,C). Large variations in etch pit density (ranging

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from 100 in a scanned area of 5x5 µm) were observed between different surfaces, but also

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between different areas of the same surface, as well as at different reaction times, possibly indicating a

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substantial heterogeneity in the distribution of crystal defects. Most etch pits were initially shallow

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(single stepped) and randomly distributed on the surface. Lateral spreading of etch pits eventually 7 ACS Paragon Plus Environment


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resulted in their coalescence leaving behind small islands that disappeared upon further dissolution

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leading to a layer-by-layer dissolution mechanism. With time, some etch pits developed into deeper,

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concentric (multi stepped) etch pits (Fig. 1D). These etch pits most likely originate from structural

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defects that penetrate several layers, whereas the monolayer etch pits nucleated at either defect-free

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surfaces or point defects33. Occasionally, also spiral etch pits were observed, which are likely to be

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related to screw dislocations. The height profile in Fig. 1E shows that individual steps are ~0.5 nm

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deep.

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The observed etch pit and step edge morphologies are in agreement with previous studies28,29 and can

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be understood by considering the crystal structure of brucite (Fig. 1F). Brucite is composed of sheets

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of edge-sharing Mg(OH)6 octahedra. Each hydroxyl (OH) group is coordinated by three Mg atoms

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with the O-H-vector perpendicular to the (001) plane. The H-atom forms hydrogen bonds to three

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oxygen atoms on the adjacent brucite sheet. The thickness of a single brucite sheet is 0.47 nm, which

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corresponds to the depth of individual steps observed in AFM. The equilateral triangular shape of the

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etch pits results from the 3-fold rotation axis normal to the (001) plane. The edges of the triangles are

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parallel to the symmetrically equivalent [100], [010] and [110] directions, which are characterized by

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strong periodic bond chains (PBCs) consisting of edge-sharing Mg(OH)6 octahedra.

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Effects of phosphate and ammonium on brucite dissolution. In general, the presence of ammonium

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and phosphate strongly increased the dissolution rate of brucite as shown by the much faster spreading

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of etch pits. As well, increasing the ammonium phosphate concentration resulted in a continuous

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acceleration of etch pit spreading (Fig. S1).

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For equilateral triangular etch pits, the spreading rate (vs) can be defined as =

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denotes the distance from the center of a triangle to its side and s is the side length of the triangle (Fig.

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1F). In our experiments, the reproducibility of absolute spreading rate values was somewhat limited as

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shown by the fact that rate measurements on different surfaces, but under the same conditions (i.e.,

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same concentrations and pH) gave variable results. For example, spreading rates in 20 mM NH4H2PO4

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(pH 8.5) measured on three different surfaces were 0.24 (±0.04), 0.33 (±0.07) and 0.49 (±0.04) nm s-1

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(Fig. 2A). The reason for this surface dependent variability of spreading rates remains elusive, but may

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be related to inherent variations in surface energy, e.g., due to differences in the number, distribution 8 ACS Paragon Plus Environment

=

, √

where x

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and nature of reactive sites34, possibly also implying variable degrees of residual stress/strain in the

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naturally grown brucite crystals. Nevertheless, an adequate comparison of spreading rates was possible

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for values measured in the same area of the same surface.

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On a given surface, measured etch pit spreading rates generally showed an approximately linear

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increase with increasing ammonium phosphate concentration at a fixed pH, e.g., from 0.07 (±0.02) nm

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s-1 in 5 mM NH4H2PO4 to 0.51 (±0.06) nm s-1 in 50 mM NH4H2PO4 at pH 8.5 (Fig. 2A). For

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comparison, measured values in NaCl solutions with the same pH and similar ionic strength only

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ranged between 0.01 and 0.03 nm s-1, thus any increase in etch pit spreading rates must be due to the

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presence of phosphate and ammonium. On the other hand, increasing the pH from 8 to 9.5 at a fixed

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concentration resulted in a nearly exponential decrease in measured etch pit spreading rates, e.g., from

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0.76 (± 0.11) nm s-1 at pH 8 to 0.04 (± 0.02) nm s-1 at pH 9.5 in 20 mM NH4H2PO4 (Fig. 2B). This

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trend may be explained by an increased activity of OH- at higher pH values resulting in a lower

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thermodynamic driving force for brucite dissolution. Another likely explanation for our observations

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is that, depending on pH, the adsorption of phosphate ions on the brucite substrate either enhances or

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inhibits its dissolution. It has been shown that the dihydrogen phosphate (H2PO4-) ion accelerates the

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dissolution of metal (oxy)hydroxides such as brucite35 and goethite36,37 by increasing their surface

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protonation through the formation of mononuclear negatively charged surface complexes. Conversely,

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the fully deprotonated phosphate ion (PO43-) may inhibit brucite dissolution by forming binuclear

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surface complexes that bridge two Mg centers, thereby increasing the energy needed for the

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detachment of Mg atoms from the brucite crystal lattice35. Thus, it seems reasonable to assume that in

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our system brucite dissolution kinetics are mostly controlled by the presence of H2PO4- and PO43- ions

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in solution and their respective adsorption on the brucite surface. This is further corroborated by

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speciation calculations (Table S1 and Fig. S2) showing that an increase in solution pH from 8 to 9.5

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decreases the activity of the dissolution enhancing H2PO4- species while increasing the activity of the

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inhibiting PO43- species, i.e., consistent with the observed decrease in etch pit spreading rates at higher

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pH values. On the other hand, the activity of the (dominant) singly protonated phosphate (HPO42-) ion

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remains virtually constant in the considered pH range, suggesting that this species most likely does not

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affect brucite dissolution. From Figure 2C it can be seen that there is generally a good positive linear 9 ACS Paragon Plus Environment


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correlation (Rsq = 0.945 - 0.981) between the measured spreading rate values of Fig.2A and B and the

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calculated H2PO4- activities in the corresponding solution.

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It should be noted that the enhancement of brucite dissolution in the presence of phosphate may also

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be related, to some degree, to the formation of Mg-phosphate bearing complexes or clusters in solution

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(Table S2), that lower the saturation state with respect to brucite, i.e., maintaining far-from

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equilibrium conditions and hence promoting further dissolution. A similar effect will arise from the

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precipitation of a phosphate-bearing phase such as struvite.

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We also observed a tendency of higher etch pit spreading rates in the presence of (NH4)2HPO4 relative

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to those measured in the presence of NH4H2PO4 at the same pH and total phosphate concentration

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(Fig. 2A). This effect is further confirmed when comparing etch pit spreading rates in NaH2PO4,

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NH4H2PO4 and (NH4)2HPO4 measured in the same 5x5 µm surface area during a continuous

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experiment (Fig. 2D). At a total phosphate concentration of 10 or 20 mM and a pH of 8.5, rates in

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NaH2PO4 (no NH4+ and NH3 present) were consistently ~20% lower compared to rates in NH4H2PO4

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and ~50% lower compared to rates in (NH4)2HPO4. Moreover, measured Mg concentrations in the

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outflow solutions showed a slightly increasing trend from NaH2PO4 to (NH4)2HPO4 (Table S3, Fig.

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S3), thus corroborating the in situ AFM observations. As in the case of phosphate already described,

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these observations may be explained in several ways. Firstly, an increase in the NH4+ activity will

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increase the ion activity product of struvite, thus favoring struvite precipitation. The formation of

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struvite will consume dissolved Mg2+ lowering the ion activity product of brucite and leading to a

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larger thermodynamic driving force for brucite dissolution. The consumption of dissolved Mg2+ by

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aqueous Mg-NH3 bearing complexes, on the other hand, likely has a negligible effect on the saturation

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with respect to brucite as such complexes are predicted to be only present in very low concentration

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(Table S2). Another possibility is that NH4+ promotes brucite dissolution via the formation of surface

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complexes or modifications in the water structure and surface hydration dynamics38,39. However, this

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would have to be tested separately from phosphate, i.e., using ammonium chloride or nitrate solutions.

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Struvite nucleation and growth onto brucite (001) cleavage surfaces. Simultaneously with brucite

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dissolution a new phase was observed to form in the presence of all ammonium phosphate solutions.

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In the earliest nucleation stages, small, isolated particles ( 6 struvite should

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always be the first phase to become stable, thus being the most likely phase to precipitate. On the other

struvite)

are

newberyite

(Mg(PO3OH)·3H2O),

bobbierite


(Mg3(PO4)2·8H2O)

and

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hand, for pH values ≤ 6, newberyite may become stable first. Depending on the initial fluid

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composition and pH, the amount of Mg that needs to be released to reach struvite saturation ranges

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between 0.06 – 57.5 mg/L (= 0.0025 – 2.4 mmol/L) (Table S4). However, Mg concentrations in the

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outflow solutions were in most cases below the detection limit ( 8.5) brucite

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dissolution slows down quickly due to the inhibitory effect of PO43- and/or OH- ions. On the other

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hand, in the pH range 6 – 8.5 the dissolution promoting effect of H2PO4- ions dominates, thus ensuring

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high Mg release rates, while the potential of precipitating undesired phases is minimized.

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While brucite is a relatively cheap Mg source, its solubility in water (Ksp = 10-10.9 at 25°C) is much

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lower compared to more expensive artificial compounds like MgCl2. This fact may have negative

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consequences for the efficiency of the struvite formation and harvesting process. However, several

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aspects also have to be taken into account when assessing the usefulness of brucite as compared to the

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highly soluble salts. Firstly, brucite dissolution is strongly enhanced in the presence of phosphate.

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Thus, sufficient amounts of brucite may readily be dissolved without the use of strong acids. In

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addition, the release of OH- ions during brucite dissolution would lower the amount of NaOH needed

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to achieve the optimal, slightly alkaline pH conditions for struvite precipitation. Secondly, if brucite

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dissolution and struvite precipitation are coupled at the reaction interface (as observed in our AFM

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experiments), absolute solubilities are not necessarily the most relevant factor, but rather the

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solubilities of the phases in the solution at the mineral interface. The total amount of dissolved Mg

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may be very low at any time, but the reaction could still be efficient because even the dissolution of

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just a few monolayers of brucite may highly supersaturate the solution at the reaction interface with

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respect to struvite43,44. Moreover, as pointed out by Adnan et al.45, high magnesium concentrations in

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the process fluid may also be of concern because of an increased risk of unintentional struvite 14 ACS Paragon Plus Environment

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precipitation elsewhere in the system, especially if the effluent of the reactor is continuously pumped

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back to the inlet of the treatment plant. This risk could be effectively lowered if Mg is not dosed via

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the external addition of dissolved MgCl2, but via the dissolution of solid brucite particles inside the

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reactor. A tight interfacial coupling between brucite dissolution and struvite precipitation, would allow

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to better control the location of struvite precipitation because supersaturated conditions will only be

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reached within the small solution volume at the particle-fluid interface, but not within the bulk

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solution.

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In our experiments, the elevated Mg concentrations at high dissolution sites most likely played the

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most important role in controlling the struvite nucleation rates. However, it seems also reasonable to

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assume that the underlying brucite substrate provided energetically favorable sites for struvite

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nucleation considering the clear alignment of the formed crystals. Thus, we may infer that the presence

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of steps on the brucite surfaces triggers heterogeneous nucleation leading to faster precipitation as

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compared to homogeneous nucleation. This would be consistent with results by Stolzenburg et al.20

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who observed faster struvite precipitation and higher phosphorous recovery rates when using

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MgO/Mg(OH)2 suspensions instead of MgCl2 solutions.

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The formation of a dense surface layer of struvite could eventually armour the brucite from further

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reaction as has been previously observed in other systems46,47,48. In our case, however, the struvite

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precipitates did not completely cover the brucite surface. Indeed, the preferential precipitation along

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deep step edges may eventually block these highly reactive sites and consequently slow down the

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reaction. However, a complete inhibition is not expected because the flatter areas of the brucite surface

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remained uncovered during our experiments and continued dissolution (i.e., formation and spreading

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of new etch pits) in these areas was observed. .

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Overall, our experimental results suggests that natural brucite is a highly suitable reactive medium for

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struvite crystallization. It should be noted, however, that our experiments were performed with

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synthetic ammonium phosphate solutions. The reactions occurring in real wastewaters, which contain

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many more dissolved species, remain to be studied in more detail to verify whether using brucite is

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indeed feasible for phosphorus recovery applications. For wastewaters from the rare-earth industry,

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Huang et al.15 already demonstrated that high struvite recovery rates can indeed be achieved with 15 ACS Paragon Plus Environment


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brucite as a Mg source. Another aspect that needs to be considered in feasibility studies is the

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availability of brucite. The current worldwide market for brucite is comparatively small and there are

391

only a few large, high-grade brucite deposits in production27. However, ample occurrences of brucite

392

in nature24,25 and a potentially large number of still unexplored and undeveloped deposits27 make the

393

availability of Mg(OH)2 an unlikely limiting factor.

394

Despite the fact that the present study focused on natural brucite, it should be emphasized that our

395

findings are also relevant for the use of synthetic Mg(OH)2 and MgO. Low grade MgO, a by-product

396

of the magnesite industry, has already been tested as a cheap alternative Mg source in a range of

397

laboratory and pilot-scale studies12,19-22. Results by Stolzenburg et al.20 and Castro et al.22, for example,

398

have shown that MgO particles quickly convert to Mg(OH)2 when suspended in water. The subsequent

399

precipitation of struvite in the presence of ammonium and phosphate was found to be largely

400

controlled by the dissolution of the newly formed Mg(OH)2 demonstrating that the new knowledge

401

gained from our study is readily transferable to such MgO-based struvite crystallization approaches.

402 403

Supporting Information. Additional AFM image sequences of brucite dissolution, thermodynamic

404

calculations (solution speciation, saturation indices), chemical analyses of fluid composition, XRD

405

analyses of reacted brucite powders and further experimental details. This material is available free of

406

charge via Internet at http://pubs.acs.org.

407 408

Acknowledgements. The authors thank V. Rapelius for help with ICP-OES analyses and A. M.

409

Schleicher for help with XRD analyses. The financial support from the Helmholtz Recruiting initiative

410

provided to Liane G. Benning and JH is kindly acknowledged. CVP acknowledges funding received

411

through the European Union Marie Curie initial training networks, CO2React and Flowtrans.

412 413

References.

414

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Wang, L.; Ruiz-Agudo, E.; Putnis, C. V.; Menneken, M.; Putnis, A. Kinetics of calcium 16 ACS Paragon Plus Environment

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phosphate nucleation and growth on calcite: implications for predicting the fate of dissolved

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phosphate species in alkaline soils. Environ. Sci. Technol. 2012, 46 (2), 834–842.

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for thought. Glob. Environ. Chang. 2009, 19 (2), 292–305.

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Heffer, P.; Prud’homme, M. Fertilizer Outlook 2015-2019. In 83rd IFA Annual Conference; International Fertilizer Industry Association, Istanbul (Trukey), 2015.

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Cordell, D.; Drangert, J.-O.; White, S. The story of phosphorus: global food security and food

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Kataki, S.; West, H.; Clarke, M.; Baruah, D. C. Phosphorus recovery as struvite from farm,

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municipal and industrial waste: feedstock suitability, methods and pre-treatments. Waste

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Kataki, S.; West, H.; Clarke, M.; Baruah, D. C. Phosphorus recovery as struvite: recent

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nutrient-rich wastewater: a review. Env. Sci Pollut Res Int 2015, 22 (22), 17453–17464.

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El Diwani, G.; El Rafie, S.; El Ibiari, N. N.; El-Aila, H. I. Recovery of ammonia nitrogen from

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Etter, B.; Tilley, E.; Khadka, R.; Udert, K. M. Low-cost struvite production using source-

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Huang, H. M.; Xiao, X. M.; Yang, L. P.; Yan, B. Removal of ammonium from rare-earth

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Lahav, O.; Telzhensky, M.; Zewuhn, A.; Gendel, Y.; Gerth, J.; Calmano, W.; Birnhack, L.

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CO2 sequestration during brucite [Mg(OH)2] dissolution. Environ. Sci. Technol. 2012, 46 (9),

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in situ atomic force microscopy observations. Clays Clay Miner. 2006, 54 (5), 598–604

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calculations. U.S. Geological Survey, Water Resources: Denver, CO 1999, p 99−4259.

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nanoscale imaging reveals the growth of calcite crystals via amorphous nanoparticles. Cryst.

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Putnis, A.; Putnis, C. V. The mechanism of reequilibration of solids in the presence of a fluid phase. J. Solid State Chem. 2007, 180 (5), 1783–1786.

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struvite crystallization — II: Applying in-reactor supersaturation ratio as a process control

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Béarat, H.; McKelvy, M. J.; Chizmeshya, A. V. G.; Gormley, D.; Nunez, R.; Carpenter, R. W.;

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water content on mineral reactions in unsaturated porous media: implications for brucite

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carbonation and CO2 sequestration. Geochim. Cosmochim. Acta 2015, 148, 477–495.

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Harrison, A. L.; Dipple, G. M.; Power, I. M.; Mayer, K. U. The impact of evolving mineral-

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water-gas interfacial areas on mineral-fluid reaction rates in unsaturated porous media. Chem.

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Geol. 2016, 421, 65–80.

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Figure captions.

538

Figure 1: (A-D) Time-lapse in situ AFM deflection images of a brucite surface during dissolution in a

539

flow-through cell. (A) After 22 min in 100 mM NaCl (pH 7). No etch pits are observed. Dissolution

540

occurred via slow retreat of pre-existing step edges. (B - D) Injection of 10 mM (NH4)2HPO4 (pH 8)

541

resulted in the immediate formation and spreading of equilateral triangular etch pits. White arrow in

542

(A-C) serves as reference point. (D) Development of concentric etch pits after 34 min in contact with

543

10 mM (NH4)2HPO4 (pH 8). (E) Depth profile along section a → b as indicated by the dashed line in

544

(D). The height of single steps is ~0.5 nm. (F) Left: the brucite structure projected along the b-axis.

545

The thickness of one Mg(OH)2 layer is 0.47 nm. Right: the brucite structure projected along the c-axis.

546

The morphology and crystallographic orientation of the edges of the triangular etch pits is outlined.

547 548

Figure 2: Measured etch pit spreading rates as a function of (A) total phosphate concentration, (B) pH

549

and (C) H2PO4- activity. (D) Comparison of etch pit spreading rates in NaH2PO4, NH4H2PO4 and

550

(NH4)2HPO4. Values measured on the same surface during a continuous experiment are indicated by

551

same colors. Symbols indicate different salt solutions (squares: NH4H2PO4, diamonds: (NH4)2HPO4,

552

triangles: NaH2PO4). Error bars are standard deviations of the measured values.

553 554

Figure 3: Time-lapse in situ AFM deflection images showing the dissolution of a brucite surface

555

(indicated by the continuous growth of etch pits) and the simultaneous nucleation of a new phase in

556

100 mM (NH4)2HPO4 (pH 8). Fresh solution (2 ml) was injected before each image. Slow growth of 21 ACS Paragon Plus Environment


557

nucleated particles is observed (e.g., particle indicated by blue arrows). After a short growth period

558

some particles were detached by the scanning tip (e.g., particle indicated by red arrows) or started to

559

dissolve again (e.g., particle indicated by white arrows).

560 561

Figure 4: (A) AFM deflection image showing globular shaped nanoparticles nucleated on a brucite

562

surface during reaction with 100 mM NH4H2PO4 (pH 8.5). (B) Same area as in (A) after the solution

563

had been kept stagnant for 20 min. Growth of particles is observed, e.g., in the area outlined by the

564

dashed rectangle. (C, D) Height profiles along sections a → b (dashed lines in (A) and (B)) showing

565

the height increase of a nanoparticle from ~8 to ~16 nm. (E) AFM deflection image showing larger

566

particle aggregates formed after 40 min in contact with 100 mM (NH4)2HPO4 (pH 8). (F) Higher

567

magnification AFM deflection image of the area outlined by the dashed square in (E) showing the

568

cluster nature of the larger aggregates. (G) Height profile along section a → b as indicated by the

569

dashed line in (F). The total height of the particle aggregate is ~30 nm.

570 571

Figure 5: AFM deflection images showing precipitates formed on brucite surfaces after longer periods

572

of reaction during ex situ experiments. (A) After ~38 h in 100 mM (NH4)2HPO4 (pH 8). Large particle

573

clusters have formed, mainly along deep steps. (B) After ~14 h in 100 mM NH4H2PO4 (pH 8.5). Some

574

elongated precipitates composed of fused-together nanoglobular particles have formed in areas of high

575

step density. (C) After ~14 h in 100 mM (NH4)2HPO4 (pH 8.5). Some thick precipitates with more

576

well-defined straight edges have formed on the brucite surface. (D-F) Height profiles along sections a

577

→ b as indicated by the dashed lines in (A-C), respectively. Some precipitates have reached

578

thicknesses of more than 200 nm. (G) Higher magnification AFM deflection image of the area

579

outlined by the dashed square in (C) revealing the nanoglobular texture of the precipitated phase.

580 581

Figure 6: (A) SEM image of a brucite surface after ~14 h of reaction in 100 mM (NH4)2HPO4 (pH

582

8.5). Numerous elongated crystals have formed along the step edges. (B) Higher magnification SEM

583

image of the same surface revealing a clear alignment of the lath-shaped crystals on the brucite

584

substrate. Insets display EDX analyses taken from the brucite surface and the newly formed crystals, 22 ACS Paragon Plus Environment

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585

demonstrating that high amounts of P were incorporated in the new phase. (C) Representative Raman

586

spectrum of the new phase (red curve), showing perfect agreement with a reference spectrum for

587

struvite (black curve).



(A-D) Time-lapse in situ AFM deflection images of a brucite surface during dissolution in a flow-through cell. (A) After 22 min in 100 mM NaCl (pH 7). No etch pits are observed. Dissolution occurred via slow retreat of pre-existing step edges. (B - D) Injection of 10 mM (NH4)2HPO4 (pH 8) resulted in the immediate formation and spreading of equilateral triangular etch pits. White arrow in (A-C) serves as reference point. (D) Development of concentric etch pits after 34 min in contact with 10 mM (NH4)2HPO4 (pH 8). (E) Depth profile along section a → b as indicated by the dashed line in (D). The height of single steps is ~0.5 nm. (F) Left: the brucite structure projected along the b-axis. The thickness of one Mg(OH)2 layer is 0.47 nm. Right: the brucite structure projected along the c-axis. The morphology and crystallographic orientation of the edges of the triangular etch pits is outlined. 109x85mm (300 x 300 DPI)

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Measured etch pit spreading rates as a function of (A) total phosphate concentration, (B) pH and (C) H2PO4activity. (D) Comparison of etch pit spreading rates in NaH2PO4, NH4H2PO4 and (NH4)2HPO4. Values measured on the same surface during a continuous experiment are indicated by same colors. Symbols indicate different salt solutions (squares: NH4H2PO4, diamonds: (NH4)2HPO4, triangles: NaH2PO4). Error bars are standard deviations of the measured values. 95x81mm (300 x 300 DPI)



Time-lapse in situ AFM deflection images showing the dissolution of a brucite surface (indicated by the continuous growth of etch pits) and the simultaneous nucleation of a new phase in 100 mM (NH4)2HPO4 (pH 8). Fresh solution (2 ml) was injected before each image. Slow growth of nucleated particles is observed (e.g., particle indicated by blue arrows). After a short growth period some particles were detached by the scanning tip (e.g., particle indicated by red arrows) or started to dissolve again (e.g., particle indicated by white arrows). 61x27mm (300 x 300 DPI)


Page 26 of 30

Page 27 of 30


(A) AFM deflection image showing globular shaped nanoparticles nucleated on a brucite surface during reaction with 100 mM NH4H2PO4 (pH 8.5). (B) Same area as in (A) after the solution had been kept stagnant for 20 min. Growth of particles is observed, e.g., in the area outlined by the dashed rectangle. (C, D) Height profiles along sections a → b (dashed lines in (A) and (B)) showing the height increase of a nanoparticle from ~8 to ~16 nm. (E) AFM deflection image showing larger particle aggregates formed after 40 min in contact with 100 mM (NH4)2HPO4 (pH 8). (F) Higher magnification AFM deflection image of the area outlined by the dashed square in (E) showing the cluster nature of the larger aggregates. (G) Height profile along section a → b as indicated by the dashed line in (F). The total height of the particle aggregate is ~30 nm. 127x237mm (300 x 300 DPI)



Figure 5: AFM deflection images showing precipitates formed on brucite surfaces after longer periods of reaction during ex situ experiments. (A) After ~38 h in 100 mM (NH4)2HPO4 (pH 8). Large particle clusters have formed, mainly along deep steps. (B) After ~14 h in 100 mM NH4H2PO4 (pH 8.5). Some elongated precipitates composed of fused-together nanoglobular particles have formed in areas of high step density. (C) After ~14 h in 100 mM (NH4)2HPO4 (pH 8.5). Some thick precipitates with more well-defined straight edges have formed on the brucite surface. (D-F) Height profiles along sections a → b as indicated by the dashed lines in (A-C), respectively. Some precipitates have reached thicknesses of more than 200 nm. (G) Higher magnification AFM deflection image of the area outlined by the dashed square in (C) revealing the nanoglobular texture of the precipitated phase. 82x52mm (300 x 300 DPI)


Page 28 of 30

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(A) SEM image of a brucite surface after ~14 h of reaction in 100 mM (NH4)2HPO4 (pH 8.5). Numerous elongated crystals have formed along the step edges. (B) Higher magnification SEM image of the same surface revealing a clear alignment of the lath-shaped crystals on the brucite substrate. Insets display EDX analyses taken from the brucite surface and the newly formed crystals, demonstrating that high amounts of P were incorporated in the new phase. (C) Representative Raman spectrum of the new phase (red curve), showing perfect agreement with a reference spectrum for struvite (black curve). 123x246mm (300 x 300 DPI)



61x44mm (300 x 300 DPI)


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In Situ Nanoscale Imaging of Struvite Formation during the Dissolution

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