Phosphate-Exchanged Mg–Al Layered Double Hydroxides: A New

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Phosphate-Exchanged Mg−Al Layered Double Hydroxides: A New Slow Release Phosphate Fertilizer Maarten Everaert,*,† Ruben Warrinnier,† Stijn Baken,† Jon-Petter Gustafsson,‡ Dirk De Vos,§ and Erik Smolders† †

Department of Earth and Environmental Science, Division of Soil and Water Management, KU Leuven, Kasteelpark Arenberg 20, B-3001 Heverlee, Belgium ‡ Department of Soil and Environment, Swedish University of Agricultural Sciences (SLU), Box 7014, 750 07 Uppsala, Sweden § Department of Microbial and Molecular Systems, Centre for Surface Chemistry and Catalysis, KU Leuven, Celestijnenlaan 200F − 02461, B-3001 Heverlee, Belgium S Supporting Information *

ABSTRACT: The global phosphorus crisis provided impetus to develop fertilizers with better P use efficiency. We tested layered double hydroxides (LDHs) as slow release fertilizers with superior performance to fertilize strongly P-fixing soils. Mg−Al LDHs with varying M2+/M3+ ratios were synthesized as NO3− forms and were exchanged with HPO42−. XRD and XANES spectroscopy confirmed the identity of the phosphate-exchanged LDH. Decreasing the M2+/M3+ ratio, i.e., increasing the anion exchange capacity, increased the selectivity of P adsorption due to the increasing charge density of the LDH layers. The fertilization efficiency of the phosphate-exchanged LDH (Mg/Al ratio of 2) was compared to that of a soluble P fertilizer in two P-deficient soils, an acid weathered soil and a calcareous soil. The P use efficiency of the P-LDH in the acid soil was up to 4.5 times higher than that of soluble P. This was likely related to a liming effect of the LDH. In the calcareous soil, the P use efficiency at low doses was only 20% above that of soluble P, whereas it was lower at high doses. These overall encouraging results warrant further studies on the boundary conditions under which P-LDHs may outperform traditional fertilizers. KEYWORDS: Layered double hydroxide (LDH), Ion exchange, Slow release fertilizer, Phosphorus-use efficiency, X-ray diffraction (XRD), X-ray absorption near edge structure (XANES)

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soluble P forms.6−8 However, if slightly higher yields are obtained with struvite, this could not yet be linked to slow P release characteristics of the material.9 Also, for party-acidulated phosphate rock, a slow P release is observed due to a low solubility, which depends on the degree of acidulation,10 on the primary source of the phosphate rock11 and on the soil conditions.12 Polymer-coated soluble P fertilizers perform less in the early plant growth stages in comparison with the noncoated alternative, while the opposite is true for the later final stages, resulting in an overall higher yield and P uptake.13 In practice, mainly multinutrient formulations (e.g., NPK) are used in coated fertilizers.14,15 Several material classes that might have the desired characteristics to perform a slow release of P have already been explored, such as superabsorbent polymers, surfactantmodified zeolites, and polymeric anion exchangers.16−18 However, clear evidence for superior nutrient efficiency has not yet been found. Recently, layered double hydroxides (LDHs) have been proposed as phosphate adsorbents for the

INTRODUCTION The global phosphorus crisis in recent years has renewed interest to develop phosphate fertilizers with better phosphate use efficiency.1 Phosphorus (P) is an essential plant nutrient, but plant-available P is only present at very low concentrations in the soil solution due to strong sorption or precipitation of the PO4 anions.2 In acid-weathered soils, strong and largely irreversible P fixation on (oxy)hydroxides of Fe and Al occurs, whereas in calcareous soils P reacts with Ca2+ or Fe3+ to form poorly soluble precipitates.3,4 Current commercial phosphate fertilizers such as diammonium hydrogen phosphate (DAP) and triple superphosphate (TSP) are known to have a high water solubility. In P-deficient weathered or calcareous soils, the PO4 anions readily react after amendment, and as a result, the P use efficiency and residual values of these P fertilizers are low. A slow release of P from the fertilizer pellet may overcome this issue by gradually supplying PO4 anions in the rhizosphere, thereby limiting fixation into soils. Slow release mineral P fertilizers used in modern agriculture can be divided into poorly soluble P forms and polymer-coated soluble P forms. Struvite, a P recovery product from wastewater, is gaining interest as a poorly soluble P slow release fertilizer.5 In most cases, struvite has been found to be equally efficient for P fertilization as © 2016 American Chemical Society

Received: April 15, 2016 Revised: May 30, 2016 Published: June 9, 2016 4280

DOI: 10.1021/acssuschemeng.6b00778 ACS Sustainable Chem. Eng. 2016, 4, 4280−4287

Research Article

ACS Sustainable Chemistry & Engineering

in more acid soils, Mg2+ and Al3+ release will be harmless in terms of plant toxicity in a well-managed agricultural soil (liming, soil pH > 5). Magnesium is a macronutrient for plants,38 so this can potentially even have a beneficial effect on the plant growth if well-balanced amounts of available Mg and Ca are maintained.39 Aluminum is not a plant nutrient, but released Al3+ will result in the formation of amorphous Al(OH)3, thereby assuring a very low Al3+ availability and absence of Al toxicity.40 The risk of environmental pollution with possible adverse effects on plant growth excludes the use of Ni- or Zn-containing LDHs for a fertilizer application.41,42 In this study, different Mg−Al LDH materials with varying anion exchange capacity (AEC) are synthesized and characterized by XRD, XANES, SEM, and ICP. The uptake of P and the P release kinetics of these materials are quantified. Finally, the P fertilization efficiency of the P-LDH is compared to that of a soluble P fertilizer in a pot trial.

selective recovery of P from waste streams, and these materials may also act as slow release P fertilizers. LDHs are inorganic anion exchangers, often consisting of layered hydroxides of divalent (M2+) and trivalent (M3+) cations. They are commonly represented by the general formula [M2+1−x M3+x (OH)2]x+ [Am−x/m]·nH2O, with x equal to M3+/(M2+ + M3+), and Am− the intercalated anion. Until now, the use of LDHs mainly focused on P recovery from waste streams, followed by a desorption of P from the materials rather than on its use as fertilizer. Promising results are obtained for P uptake with Mg− Al,19−22 Mg−Fe,23,24 and Zn−Al25,26 LDHs from both synthetic solutions and wastewater. Also, Ca-containing materials with good P removal properties are reported, but their precipitation-based P interaction is a major disadvantage for P recovery.27,28 Overall, calcination of as synthesized LDHs,25,29 as well as the introduction of quaternary cations (e.g., Zr4+) in the material,30 improve the P uptake properties of LDHs. The feasibility of state-of-the-art P desorption and LDH regeneration is, however, arguable. Highly saline or alkaline desorption solutions (up to 2 M NaCl30 or 1 M NaOH31) are required to replace the intercalated HPO42− anions by Cl− or OH− anions, which have a lower affinity for the LDH than HPO42−.20 This would lead to large saline or caustic waste streams. Alternatively, the HPO42− anions can be effectively exchanged in a CO32− solution,31 given that CO32− has an even higher affinity for the LDH than HPO42−, but this methodology would require material calcination to regenerate the LDH, which is energy intensive. A successful attempt has been made to obtain a P desorption at lower NaCl−NaOH concentrations using anionic surfactants, although a large scale application remains doubtful.32 Clearly, the current practices to recycle PLDHs have major drawbacks, and there is a need for new ways to efficiently use P-LDHs. P-LDHs might be used directly as slow and controlled release P fertilizers in soils. LDHs are able to slowly release intercalated anions, as proven in many other applications.33,34 However, the fertilization efficiency of P-LDHs has not yet been investigated. In pH neutral soils, a slow release of P from LDHs may be obtained by the ion exchange reaction between intercalated HPO42− and HCO3− or CO32− anions from the soil solution. In an acid-weathered soil, this ion exchange could be accompanied by P release caused by LDH dissolution.35,36 In addition to these slow release mechanisms, LDHs may release P in response to specific signals from the plant roots, which we term controlled release. This additional P release from LDHs may be triggered by rhizosphere conditions and by plant strategies to overcome P deficiency, such as acidification, excretion of organic anions, and a higher carbonate concentration around the respiring roots. Therefore, P release from LDHs may be spatiotemporally controlled: P release is not only slow, as is the case for common slow release fertilizers, but it may also be localized near the plant roots. The development of an LDH fertilizer material may allow us to bridge the gap between the recycling of P from waste streams and P fertilization in problematic soils. Against this background, this study was set up to develop a synthesis and exchange protocol for P-exchanged LDHs and test their fertilizer value. The Mg2+ and Al3+ were selected as metal cations for the LDH fertilizer since it is not harmful to soil organisms, including plants. After P desorption, the residual phase is hydrotalcite (Mg6Al2(OH)16CO3.4H2O), a naturally occurring mineral in certain alkaline soils.37 If long-term LDH addition would result in partial dissolution of the material, e.g.,

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EXPERIMENTAL SECTION

LDH Synthesis and Characterization. Nitrate forms of Mg−Al LDHs were synthesized by coprecipitation using mixed solutions of Mg(NO3)2·6H2O and Al(NO3)3·9H2O with 3 M total metal concentration.43 Solutions with Mg/Al ratios of 2, 3, and 4 were used to obtain three LDH materials with a different AEC: Every Al3+ cation in the mineral lattice gives rise to a positive charge which contributes to the AEC. Here, 50 mL of the metal solutions were added at a rate of 10 mL/h to 400 mL of a 0.01 M NaOH solution at pH 12 under vigorous stirring using a B. Braun Perfusor pump. The pH was maintained at 12 ± 0.2 by adding 3 M NaOH using a Metrohm Titrino 702. The synthesis was performed at room temperature and in a nitrogen atmosphere to prevent the formation of carbonate in solution. Afterward, the material was ripened in the synthesis mixture for 0.5 h and recovered by centrifugation (20 min, 15,000g). The materials were washed three times with milliQ water and subsequently dried at 80 °C. The chemical composition of as-synthesized LDHs was determined by inductively coupled plasma optical emission spectrometry (ICPOES; PerkinElmer Optima 3300 DV) after dissolution in HNO3. Crystallographic and morphological characteristics of as-synthesized and P-exchanged LDHs were investigated by scanning electron microscopy (SEM) and X-ray diffraction (XRD). The SEM micrographs were obtained using a PHILIPS XL 30 FG microscope on Aucoated samples. Powder XRD analysis was performed with a STOE STADI P Combi instrument in the Debye−Scherrer geometry (Cu Kα1) using an image plate detector (2θ = 0−62°; Δ2θ = 0.03°). The AEC of the materials was calculated from the Mg/Al ratio, assuming dehydrated materials; the charge density of the hydroxide layers (Cd, in e/nm2) using the following formula:44

Cd = xe/(a 2 sin 60°) with a = 2d011 = 0.309 nm, x = M3+/(M2+ + M3+), and e the elementary charge. Phosphate Uptake by LDHs. The uptake of phosphate by synthesized LDHs was determined by equilibrating LDHs in duplicate with solutions containing different concentrations of KH2PO4. The amount of P added to the solution was calculated as a percentage of the theoretical AEC of the materials assuming adsorption of the divalent HPO42− and ranges between 0% and 250% of the AEC. The pH of the KH2PO4 solutions was adjusted to 7.2 with a dilute NaOH solution before the LDH addition. The suspensions were shaken endover-end for 24 h and centrifuged (3000g, 20 min), and the concentrations of P, Mg, and Al in the supernatants were determined by ICP-OES. The recovered LDH material was characterized using XRD. Mg/Al (2/1) material treated in 10%, 70%, and 250% P solutions were analyzed using P K-edge XANES spectroscopy on beamline BL-8 of the Synchrotron Light Research Institute, Thailand.45 The XANES data processing was performed by means of the 4281

DOI: 10.1021/acssuschemeng.6b00778 ACS Sustainable Chem. Eng. 2016, 4, 4280−4287

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ACS Sustainable Chemistry & Engineering Table 1. Selected Properties of Soils

oxalate extractione soil location

soil type

sand (%)

silt (%)

clay (%)

pH

eCEC (cmolc kg )

Al (mg kg )

Fe (mg kg−1)

P (mg kg−1)

Kuinet, Kenia Guadalajara, Spain

ferrosol calcic cambisol

16 44

6 31

78 25

4.2 7.0

8.9 14.1

700 440

2100 160

126 120

a

b

b

b

c

−1

d

−1

a

Major FAO soil grouping.50 bParticle size analysis by pipet method.51 cpH (1:5) determined in 0.01 M CaCl2. deCEC (effective CEC) determined at soil pH by silver thiourea method.52 eAmmonium oxalate-extractable Al, Fe, and P.53 content at field capacity (pF 2). After the last rewetting phase, the soils were incubated for 5 days at 20 °C in darkness. The soils were transferred into pots, and five pregerminated barley seedlings were planted in every pot, 0.5−1 cm beneath the soil surface. The pots were installed in a growth chamber with a day/night regime of 16 h light (20 °C) and 8 h darkness (16 °C), a light intensity of 650 mol photons m−2 s−1 during the day, and a relative humidity of 75%. Water losses were corrected daily using deionized water. The soil surface was covered by polyethylene beads to reduce evaporation losses. After 4 days, the number of plants per pot was reduced to three by removing the smallest plants in each pot. After 17 days, the shoots were harvested, dried at 60 °C for 48 h, and their weights determined. Their elemental composition was determined by ICP-OES after digestion in boiling HNO3 in a hot block. Statistical analysis of the results was performed in JMP Pro 11 (SAS).

Athena program in the Demeter Software Package (v 0.9.24).46 By means of a linear combination fitting (LCF) approach, a set of spectra of known standards (variscite, amorphous aluminum phosphate, and struvite; more are listed in the SI) were combined and fitted to the sample spectra.47 Phosphate Release Kinetics from P-LDHs. The release kinetics of PO4 were determined in a zero sink system at pH 8.3 in 2 mM NaHCO3. These conditions mimic those in the rhizosphere of a calcareous soil: a total dissolved HCO3− concentration around 2 mM and a constant withdrawal of P from soil solution by plant roots. First, P-LDHs were prepared at pH 7.2 by equilibrating the synthesized LDH materials for 24 h with a solution containing an amount of phosphate (added as KH2PO4) equal to 200% of the theoretical AEC of the LDH. Afterward, the suspensions were centrifuged, and obtained P-LDHs were washed with milliQ water and dried at 60 °C. The P release kinetics from P-LDHs and from two soluble P fertilizers (KH2PO4 and granular triple super phosphate) were determined by equilibrating them with anion exchange membranes (AEMs) as infinite P sink. These AEMs (BDH Laboratory Supplies, Poole, England; anion exchange capacity: 0.037 mequiv cm−2) have previously been used to determine plant available P.4 Before the experiment, the membranes were preconditioned in 0.5 M NaHCO3 in order to saturate the binding sites with carbonate. During equilibration, dialysis membranes (Spectra/Por 4; regenerated cellulose; MWCO, 12−14 kDa; diameter, 16 mm) were used to physically separate the fertilizer material from the infinite sink. The internal solution contained either 10 mg of the P-LDH, 3 mg of KH2PO4, or 3 mg of industrial triple super phosphate (TSP, 18.3 wt % P) in 2 mM NaHCO3 at pH 8.3. The external solution contained two 6 cm2 AEM strips suspended in 35 mL of 2 mM NaHCO3 at pH 8.3. The ratio of the AEM binding capacity to the HPO42−, initially present in the inner compartment, was always larger than 30. For every P material, three replicates were used. These suspensions were shaken end-over-end for 1600 h, and the P release from the fertilizers was monitored by recovering the AEMs from the external compartment at selected moments and by replacing them with fresh AEMs. The recovered AEMs were shaken in 0.5 M HCl for 24 h to desorb the membrane bound P, and the P in these solutions was measured by ICP-OES. Pot Trial. The P fertilization efficiency of the phosphate-exchanged Mg−Al (2/1) LDH was determined and compared with that of KH2PO4 by growing barley (Hordeum vulgare L.) in two P-deficient soil samples collected from the surface horizon of a Ferrasol (Kuinet, Kenia; pH 4.3) and a calcic cambisol (Guadelajara, Spain; pH 7.0) (soil characteristics are listed in Table 1). Apart from these two soils, an additional treatment was prepared by mixing the Kuinet soil with 1 wt % of powdery gypsum (CaSO4·2H2O) to identify the effect of neutral cations (Mg2+ or Ca2+) on the minimization of Al3+ toxicity. Gypsum is commonly used in weathered soils to control subsurface acidity and improve soil structure.48,49 The soils were amended with P doses of 0, 10, 30, 100, and 300 mg P/kg soil, either as powdered PLDH (2/1) material (mortar grinded) or as dissolved KH2PO4. The K concentrations in all treatments were adjusted to 100 mg K (kg soil)−1 by adding dissolved KCl. The soils were also amended with 100 mg N (kg soil)−1 and 10 mg Mg (kg soil)−1 using Ca(NO3)2 and Mg(NO3)2· 6H2O. For every treatment, three replicate pots were prepared. After addition of the nutrients, the soils were subjected to four drying/ rewetting cycles to increase P fixation and simulate conditions in the field. One cycle consisted of air drying (24 h) at 25 °C, followed by rewetting the soils with deionized water to 80% of their moisture

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RESULTS AND DISCUSSION Physical and Chemical Properties of LDHs. Elemental analysis of LDHs shows that their M2+/M3+ ratios slightly deviate from those in the synthesis solution, with a general tendency of lower Mg incorporation as the Mg/Al ratio in the synthesis increases (Table 2). The XRD patterns of asTable 2. Material Characteristics of Three LDHs

LDH material

molar Mg/Al ratio in synthesis solution (−)

molar Mg/Al ratio in LDH materials (−)

calculated AEC (meq g−1)

charge density (e nm−2)

Mg−Al (2/1) Mg−Al (3/1) Mg−Al (4/1)

2/1 3/1 4/1

1.95 2.74 3.64

4.2 3.5 3.0

4.1 3.3 2.7

synthesized LDH materials correspond to the expected structural characteristics of layered double hydroxides (Figure 1): sharp and symmetric reflections for the (003) and (006) crystal planes and broader and asymmetric reflections for the (015) and (018) planes. The observed (003) d-spacings of assynthesized LDHs are almost equal, suggesting a similar orientation of interlayer nitrate anions in all materials. Given the constant layer thickness of 4.8 Å,54 the comparison of these (003) d-spacings to the geometry of the nitrate anion (approximated as a cylinder with radius 2.77 Å and height 2.76 Å; details in Section 4, SI) shows that nitrate does not have to be oriented parallel to the layer but can be tilted up to approximately 58°−60° in the interlayer space in all three LDH materials. This orientation would allow a more efficient compensation of the charges of the layer. The small decrease of about 0.1 Å in (003) d-spacing with decreasing Mg/Al ratio (increasing AEC) of the as-synthesized materials is attributed to the shortening of the hydrogen bonds.55 This contrasts with earlier work: It has been reported that the basal spacing of nitrate-exchanged LDHs increases with decreasing Mg/Al ratio, which was explained by a shift to a more vertical orientation of the interlayer nitrate anions in materials with higher AEC.36 4282

DOI: 10.1021/acssuschemeng.6b00778 ACS Sustainable Chem. Eng. 2016, 4, 4280−4287

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ACS Sustainable Chemistry & Engineering

Figure 1. Right: XRD patterns of the LDH materials obtained after treatment in exchange solutions containing 250% P of the AEC of the specific LDH, compared to the patterns of the as-synthesized materials. Left: Graphic presentation of the intercalated anions and their effect on the basal spacings (given in Å).

and −6.5 e nm−2, depending on its orientation (details in Section 5, SI). In LDHs with high charge density (2/1 and 3/1 Mg−Al LDHs), the higher charge density of phosphate compared to nitrate favors the intercalation of phosphate by minimizing the energy of the system: The positive charge in the layers is better compensated, and the entropy of the system is maximized by releasing two nitrates into solution for every phosphate sorbed. In contrast, for the Mg−Al (4/1) LDH with the lowest charge density, nitrate is energetically favored since a too high anion charge density would cause electrostatic repulsion due to a local excess of negative charge. This explains the relatively low selectivity for phosphate of the Mg−Al (4/1) LDH. The (003) reflections in the XRD patterns of LDHs provide information about the nature and orientation of the anions in the interlayer gallery. To this end, the XRD patterns of assynthesized LDH materials are compared with those of LDHs equilibrated with exchange solutions containing 250% of their AEC as phosphate (Figure 1). Given the constant layer thickness of 4.8 Å, the effect of phosphate binding on the interlayer spacing can be determined from the basal spacing. In the XRD pattern of the P-exchanged Mg−Al (2/1) LDH, a shift of the (003) reflection to lower 2θ values is observed upon exchange. This is attributed to intercalation of phosphate, which replaces the smaller nitrate anion and therefore leads to a larger basal and interlayer spacing. The theoretical interlayer spacing of a phosphate-intercalated LDH can be determined from the size of the phosphate tetrahedron. The phosphate tetrahedron has two different stable orientations in the interlayer of an LDH: either face sharing with one hydroxide layer and corner sharing with the other or edge sharing with both hydroxide layers (Figure S2). Assuming a P−O bond length of 1.53 Å, an ionic radius of the O2− ion of 1.38 Å, and an O−P−O bond angle of 109°,58 geometrical calculations show that the interlayer spacing of the P-LDH may range between 3.38 and 3.86 Å. The observed interlayer spacing of 3.59 Å (8.39 Å basal spacing minus 4.8 Å hydroxide layer spacing) falls perfectly in this range. The (003) reflection of the nitrate-intercalated LDH, observed in the as-synthesized Mg− Al (2/1) LDH, disappears completely after equilibration of this material with excess phosphate, showing that nitrate is completely replaced by phosphate in the interlayer gallery.

The SEM micrographs show a crystal size of around 20 nm in all as-synthesized LDHs (Figure S1, left), which is similar to a previously reported value.56 Due to the small crystal size, the typical hexagonal crystal morphology of LDHs is not readily recognized with SEM. With decreasing Mg/Al ratio of the materials, the XRD reflections are less intense and broader, which hints at lower crystallinity or smaller sizes of the crystalline domains (Figure 1). Wang et al. postulate two effects influencing the crystallinity of LDHs.57 On one hand, stability increases with decreasing Mg/Al ratio because the increased positive charge leads to stronger bonding between the hydroxide layers and interlayer anions. On the other hand, increased Al3+ substitution for Mg2+ introduces additional strain in the lattice. Since crystallinity decreases here with decreasing Mg/Al ratio of the material, the latter cation size effect outweighs the former effect in the net effect on the crystallinity. Phosphate Uptake by LDHs. The exchange isotherms reveal higher uptake of P, expressed relative to the AEC, for LDHs with increasing AEC, i.e., with lower Mg/Al ratio. This means that a larger AEC increases the selectivity for HPO42− over NO3− (Figure 2). The difference in P uptake can be attributed to both the charge density of the LDH layers (Table 2) and that of the nitrate and phosphate anions. The cross section charge density of the phosphate anion is approximately −7.5 e nm−2, whereas that of the nitrate anion is between −4.1

Figure 2. P-exchanged isotherms for the Mg−Al (2/1) (diamonds), Mg−Al (3/1) (squares), and the Mg−Al (4/1) (triangles) materials. 4283

DOI: 10.1021/acssuschemeng.6b00778 ACS Sustainable Chem. Eng. 2016, 4, 4280−4287

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ACS Sustainable Chemistry & Engineering

The feature at 2158 eV is very probably due to an Mg−PO4 interaction since it is also, but more clearly, present in the spectrum of the Mg phosphate mineral struvite. Such a weak feature can also be seen in the spectrum of newberyite, MgHPO4·3H2O.60 Whether this feature results from a minor amount of Mg3(PO4)2 or from electrostatically bound intercalated P is inconclusive. From comparison with the spectra of amorphous AlPO4 and variscite, it can be concluded that these phases are not major components in P-LDH samples, as K-edge energies of the former are ∼0.5 eV higher and as they have different post-edge features at 2161 and 2171.5 eV. Slow Release of Phosphate from P-LDHs. A clear difference in P release is observed between the soluble and the slow release P fertilizers during equilibration with a zero sink for P (AEMs) in 2 mM HCO3− (Figure 3). The solid TSP and

This explains why phosphate can occupy most of the AEC of this LDH (Figure 2). In contrast to the observations with the Mg−Al (2/1) LDH, the (003) reflection of the as-synthesized Mg−Al (4/1) LDH hardly shifts after equilibration with phosphate. A basal spacing below 8 Å is maintained, resulting in an interlayer spacing of 3.18 Å. This is too small for the phosphate tetrahedron (Figure S2). Therefore, phosphate cannot be present in the interlayer of this LDH, and so the interlayer gallery will still be occupied by nitrate. The exchange isotherm has a very steep (selective) part followed by a very small increase in sorbed P (Figure 2). The phosphate removed from the exchange solutions (up to 46% of the AEC) must, therefore, be bound to the outer surfaces of the LDH particles. The positive charge density on the crystal edges of an LDH is much higher compared to that of the basal plane,59 and therefore, an edge sorption site can be treated as a point charge instead of a charged surface. Under such circumstances, size and orientation of competing anions are less important, and Coulomb’s law dictates the interaction of anions with the sorption sites. This results in a high selectivity for divalent HPO42− over monovalent NO3− for sorption at the outer surfaces of LDHs, which explains the relative high P uptake by the Mg−Al (4/1) LDH in absence of P intercalation. The Mg−Al (3/1) LDH combines the P uptake mechanisms from both Mg−Al (2/1) and Mg−Al (4/1) materials. After equilibration with phosphate, the XRD pattern of the Mg−Al (3/1) material shows a double (003) and a double (006) reflection, highlighting the presence of two different anion species in the LDH interlayer, i.e., nitrate and phosphate. Deconvolution of this peak is difficult due to a large overlap. The maximum value of this (003) reflection at a 2θ value of 11.14° results in an interlayer spacing comparable to that of the as-synthesized material and can be attributed to nitrate intercalation. To more accurately calculate the basal spacing corresponding to phosphate intercalation in this material, the smallest 2θ value of the splitted 006 reflection is used (21.61°). This approach yields a basal spacing of 3.38 Å, which is perfectly in agreement with the calculated basal spacing corresponding to one of two stable phosphate orientations in the LDH interlayer spacing (Figure S2). Material stability during the anion exchange reaction is crucial: Dissolved metal species may interact with the free P to form insoluble precipitates such as AlPO4 and Mg3(PO4)2. These P forms are much less available for plant uptake compared to P bound through ion exchange, which would adversely affect the P fertilization efficiency of the material. The stability of the LDH material during equilibration with phosphate is supported by the preservation of the material crystallinity as evidenced by the XRD patterns, the absence of new XRD reflections on the diffraction pattern of LDHs, and absence of visual changes of the material on the SEM micrographs. Also, from free Mg and Al concentrations in the exchange solutions after a 24 h reaction (Figure S3) and the solution pH at that moment (Figure S4), a good LDH stability can be concluded; e.g., maximally 3% of LDH-derived Mg is dissolved, and this fraction is highest for the Mg−Al (4/1) LDH (details in Section 2, SI). Finally, the XANES spectra of the P-exchanged Mg/Al (2/1) LDH material at different P loadings are mutually very similar and display a clearly unique P bonding character that cannot be considered as mixtures of the reference materials (details in Section 3, SI). The most special feature of P-LDH XANES spectra is a post-edge shoulder at 2158 eV, and a broad post-edge peak at 2170 eV (Figure S5).

Figure 3. Semilog plot of P release kinetics from soluble P fertilizers (KH2PO4 and TSP) and from slow release P fertilizers (P-exchanged LDHs) upon reaction with an anion exchange membrane acting as a zero sink. The vertical axis expresses the desorbed P as percentage of the total initial P.

KH2PO4 particles dissolve quickly. Within 24 h, nearly all added P is recovered by the AEMs, and no more P is detected in solution. The release of P by LDHs is clearly slower compared to soluble P forms, confirming that the P-LDH can act as a slow release compound. A kinetic desorption model (details in Section 4, SI) shows a comparable P release for different LDHs, with the half-life of the P release ranging from 30 h for the Mg−Al (2/1) to 41 h for the Mg−Al LDH (3/1). After about 650 h, no detectable desorption of P from LDHs is observed. Unlike for the soluble P fertilizers, not all P could be released from the LDH phase in this experimental setup; up to 49% of P for the (2/1), 53% for the (3/1), and 75% for the (4/ 1) LDH is not desorbed from the initial materials after 1600 h. This can be explained by the relatively low amount of divalent carbonate anion at pH 8.3 and the possible formation of insoluble P precipitates such as Mg3(PO4)2 as mentioned above. The %P desorbed after 1600 h was distinctly larger for the (2/1) and (3/1) than for the (4/1) Mg−Al LDH. This is explained by the selectivity of LDHs for carbonate over phosphate, which is, in turn, explained by the charge density of LDHs and anions. The high cross section charge density of carbonate (−8.37 to −13.52 e nm−2; details in Section 3, SI) explains why more phosphate is released from P-LDHs with higher charge density. Despite the incomplete P release, the observed slow release is encouraging and suggests that P-LDHs may have potential as slow release fertilizers. The phosphateexchanged Mg−Al 2/1 LDH is selected for use in the pot trial: 4284

DOI: 10.1021/acssuschemeng.6b00778 ACS Sustainable Chem. Eng. 2016, 4, 4280−4287

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ACS Sustainable Chemistry & Engineering

Figure 4. Total shoot P (mg P/pot, three seedlings/pot) plotted against the P fertilizer dose. The fertilizer was either dissolved KH2PO4 (diamonds) or a powdered P-LDH slow release fertilizer (spheres). The ∗ indicates significant differences between both fertilizer forms (p < 0.05). Data of dry matter yields are in Figure S7.

Of all LDHs tested, this material has the highest phosphate binding capacity and shows the highest quantity of P released per mass of LDH in the desorption test. Pot Trial. The biomass yield and the total shoot P uptake increase with increasing P fertilizer dose in both soils (Figure S7 and Figure 4). With increasing KH2PO4 fertilizer dose, the shoot P concentrations increase from deficient values (830 mg P kg−1 for acid soil and 1000 mg P kg−1 for calcareous soil) to adequate values (>1600 mg P kg−1) illustrating that both soils are P deficient.61 In the acid soil, high doses of the P-LDH significantly increase both shoot P concentrations and plant yield compared to the corresponding doses of dissolved KH2PO4 fertilizer. The shoot P uptake (mg P/pot) was 4.5fold larger for the P-LDH than for the KH2PO4 at the intermediate dose of 100 mg P/kg soil. Surprisingly, if this soil is additionally amended with 1 wt % CaSO4, no significant differences in shoot yield or shoot P between the LDH and KH2PO4 are observed. This indicates that the high yield in PLDH fertilized pots, compared to those amended with KH2PO4, is probably not largely due to a slow release effect of P. The soil pH of the acid Kuinet soil is low (Figure S8), indicating the possibility of manganese or aluminum toxicity in the plants.62 Plant analysis suggests Mn toxicity in the unamended acid soil, which is partly alleviated by KH2PO4 addition but strongly reduced by the P-LDH addition. The Mn concentration of plants grown in the acid soil amended with low doses of KH2PO4 (0−30 mg P/kg soil) ranges between 1000 and 1500 mg Mn kg−1. This largely exceeds toxic thresholds for Mn in barley.63 In contrast, the plants grown in the soil amended with high doses of P-LDH only contain 45 mg Mn kg−1. In plants grown in the gypsum treated Kuinet soil with KH2PO4, the shoot Mn content is reduced to 420 mg Mn kg−1. Most likely, a fraction of the alkaline LDH material added to the acid soil dissolves.35 Soil pH indeed increases with an increasing P-LDH dose from 4.3 to 6.0 (Figure S8), thereby reducing Mn solubility and toxicity. Therefore, alleviation of toxicity of Mn (and likely Al) explains most of the added value of the LDH over KH2PO4 in soils without gypsum application. In the gypsum treated soils, Mn toxicity is likely reduced by the dissolution of gypsum and the Ca2+−Mn2+ competition for root uptake64 and by the small but important increase in soil pH from 4.3 to 4.5 compared to the treatment without gypsum (Figure S8). The promising effects of P-LDHs in acid soils can be compared with that of reactive rock phosphates where dissolution is accompanied by alkalinization.65 In the calcareous Guadalajara soil, the shoot P concentrations and biomass yields increase with an increasing dose of KH2PO4

(Figure 4 and Figure S7). At low fertilizer doses, the P-LDH performs equally or better than KH2PO4. At a dose of 30 mg P/ kg soil, the plant internal P is significantly higher for the P-LDH fertilizer than for KH2PO4, and the shoot P/pot is 1.2-fold larger than in the corresponding KH2PO4 treatment. This may be due to a slow release of P from the P-LDH. However, this trend is reversed at high P doses, where the shoot P concentration and yield of the KH2PO4 treatments are above P-LDH treatments. This trend is unrelated to pH changes. Given the relatively high pH of this soil, the LDH dissolution is improbable, and the changes in soil pH are minor (Figure S8). Elemental analysis suggests S deficiency in the soils; none of the soils were amended with S (except gypsum treatment). In the calcareous soil, the shoot S concentration decreases from 1700 mg S kg−1 (control) to 970 mg S kg−1 (highest LDH treatment), whereas adequate concentrations are about 1500− 4000 mg S kg−1.61 We speculate that the low S availability for plants in this LDH treatment may be related to SO42− sorption on the LDH phase. Given a soil pH of 7.0, PO4 in solution is only partly present as divalent HPO42−, thereby favoring the adsorption of divalent SO42− and limiting S availability.66 We have replanted the same soils three months after the first trial and amended all soils with 100 mg SO4−S kg−1 (NaHSO4· H2O). Plant yield in the calcareous soil increased with increasing P application, and there was no more decline in yield at the highest LDH-treated soils, with shoot S concentrations of about 2800 mg S kg−1. The shoot P uptake was not significantly different between the LDH or KH2PO4 treatments; but all values were lower in the replanted test than in the first test, likely due to the P fixation upon aging of the soil (details not shown). This suggests that the adverse effects of high LDH doses in the first test with the calcareous soil are due to S fixation. In the acid soils, LDHs probably dissolve and do not sequester significant amounts of S. This study highlights the potential for direct use of phosphate-exchanged LDH materials as fertilizers. The synthesized P-LDHs bind phosphate in the interlayer gallery and at sorption sites on the outer surface. The slow release properties of these materials are confirmed in an environment that mimics the conditions in soil solutions. The fertilization efficiency of the Mg−Al (2/1) P-LDH, containing 4.0% P (by weight), was compared to that of dissolved KH2PO4 (20.2% P by weight) in a pot trial with two contrasting P-fixing soils. In most treatments, the P-LDH performs equally or better than the soluble P fertilizer, although the better performance of PLDHs could not directly be related to its slow release properties. More research is needed to develop LDH materials 4285

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derived struvite fertilizers as phosphorus sources for plants. Chemosphere 2012, 89 (10), 1202−1210. (10) McLay, C. D. A.; Rajan, S. S. S.; Liu, Q. Agronomic effectiveness of partially acidulated phosphate rock fertilizers in an allophanic soil at near-neutral pH. Commun. Soil Sci. Plant Anal. 2000, 31 (3−4), 423− 435. (11) Use of Phosphate Rocks for Sustainable Agriculture; FAO Fertilizer and Plant Nutrition Bulletin; Volume 13; Food and Agriculture Organization of the United Nations: Rome, 2004. (12) Kumar, V.; Gilkes, R. J.; Bolland, M. D. A. The agronomic effectiveness of reactive rock phosphate, partially acidulated rock phosphate and monocalcium phosphate in soils of different pH. Fert. Res. 1993, 34 (2), 161−171. (13) Pauly, D. G.; Malhi, S. S.; Nyborg, M. Controlled-release P fertilizer concept evaluation using growth and P uptake of barley from three soils in greenhouse. Can. J. Soil Sci. 2002, 82 (2), 201−210. (14) Tomaszewska, M.; Jarosiewicz, A. Use of polysulfone in controlled-release NPK fertilizer formulations. J. Agric. Food Chem. 2002, 50 (16), 4634−4639. (15) Jarosiewicz, A.; Tomaszewska, M. Controlled-Release NPK Fertilizer Encapsulated by Polymeric Membranes. J. Agric. Food Chem. 2003, 51, 413−417. (16) Zhan, F.; Liu, M.; Guo, M.; Wu, L. Preparation of superabsorbent polymer with slow-release phosphate fertilizer. J. Appl. Polym. Sci. 2004, 92, 3417−3421. (17) Pickering, H. W.; Menzies, N. W.; Hunter, M. N. Zeolite/rock phosphate - A novel slow release phosphorus fertiliser for potted plant production. Sci. Hortic. (Amsterdam, Neth.) 2002, 94, 333−343. (18) Sengupta, S.; Pandit, A. Selective removal of phosphorus from wastewater combined with its recovery as a solid-phase fertilizer. Water Res. 2011, 45 (11), 3318−3330. (19) Das, J.; Patra, B. S.; Baliarsingh, N.; Parida, K. M. Adsorption of phosphate by layered double hydroxides in aqueous solutions. Appl. Clay Sci. 2006, 32, 252−260. (20) Park, K. Y.; Song, J. H.; Lee, S. H.; Kim, H. S. Utilization of a Selective Adsorbent for Phosphorus Removal from Wastewaters. Environ. Eng. Sci. 2010, 27 (9), 805−810. (21) Novillo, C.; Guaya, D.; Allen-Perkins Avendaño, a.; Armijos, C.; Cortina, J. L.; Cota, I. Evaluation of phosphate removal capacity of Mg/Al layered double hydroxides from aqueous solutions. Fuel 2014, 138, 72−79. (22) Kuzawa, K.; Jung, Y. J.; Kiso, Y.; Yamada, T.; Nagai, M.; Lee, T. G. Phosphate removal and recovery with a synthetic hydrotalcite as an adsorbent. Chemosphere 2006, 62 (1), 45−52. (23) Drenkova-Tuhtan, A.; Mandel, K.; Paulus, A.; Meyer, C.; Hutter, F.; Gellermann, C.; Sextl, G.; Franzreb, M.; Steinmetz, H. Phosphate recovery from wastewater using engineered superparamagnetic particles modified with layered double hydroxide ion exchangers. Water Res. 2013, 47, 5670−5677. (24) Seida, Y.; Nakano, Y. Removal of phosphate by layered double hydroxides containing iron. Water Res. 2002, 36, 1306−1312. (25) Cheng, X.; Huang, X.; Wang, X.; Sun, D. Influence of calcination on the adsorptive removal of phosphate by Zn-Al layered double hydroxides from excess sludge liquor. J. Hazard. Mater. 2010, 177 (1−3), 516−523. (26) Yang, K.; Yan, L. G.; Yang, Y. M.; Yu, S. J.; Shan, R. R.; Yu, H. Q.; Zhu, B. C.; Du, B. Adsorptive removal of phosphate by Mg-Al and Zn-Al layered double hydroxides: Kinetics, isotherms and mechanisms. Sep. Purif. Technol. 2014, 124, 36−42. (27) Xu, Y.; Dai, Y.; Zhou, J.; Xu, Z. P.; Qian, G.; Lu, G. Q. M. Removal efficiency of arsenate and phosphate from aqueous solution using layered double hydroxide materials: intercalation vs. precipitation. J. Mater. Chem. 2010, 20, 4684. (28) Ashekuzzaman, S. M.; Jiang, J. Q. Study on the sorptiondesorption-regeneration performance of Ca-, Mg- and CaMg-based layered double hydroxides for removing phosphate from water. Chem. Eng. J. 2014, 246, 97−105.

that meet the requirements for both selective P adsorption from wastewater and efficient P fertilization. Also, a comparison of P release and P use efficiency between P-LDH fertilizers and different commercial P fertilizers, including slow release fertilizers, would allow the boundary conditions of the PLDH use to be better determined.

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ASSOCIATED CONTENT

* Supporting Information S

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acssuschemeng.6b00778. Additional details of the material stability, XANES measurement, anion geometry, desorption kinetics, and pot trial are provided. (PDF)

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS M.E. thanks IWT for a Ph.D. fellowship, and S.B. thanks the Onderzoeksfonds KU Leuven for the postdoctoral mandate. We are grateful to KU Leuven for support in the interdepartmental PB3 project. Also acknowledged are Ann Kristin Eriksson (SLU, Sweden), Dean Hesterberg (North Carolina State University, U.S.A.), and Steve Hillier (James Hutton Institute, Aberdeen, U,K,) for compiling the XANES reference spectra. Wantana Klysubun and the staff at BL-8 at SLRI, Thailand, are thanked for support and organization of the beamline.

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REFERENCES

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