Role of Alcoholic Hydroxyls of Dicarboxylic Acids ... - ACS Publications

Research Article pubs.acs.org/journal/ascecg

Role of Alcoholic Hydroxyls of Dicarboxylic Acids in Regulating Nanoscale Dissolution Kinetics of Dicalcium Phosphate Dihydrate Lihong Qin,† Lijun Wang,*,† and Baoshan Wang*,‡ †

College of Resources and Environment, Huazhong Agricultural University, 1 Shizishan Street, Wuhan 430070, Hubei, China College of Chemistry and Molecular Sciences, Wuhan University, Luojiashan Road, Wuhan 430072, Hubei, China

‡

S Supporting Information *

ABSTRACT: Due to the potential shortage of phosphate (P) rock resources and a faster growth in demand for phosphate fertilizers, unraveling the kinetics of calcium phosphate (Ca−P) crystallization and dissolution is important for understanding the P mobility and bioavailability. Plants have developed different strategies, such as carboxylic acid exudation into the rhizosphere, to cope with low P bioavailability through dissolution of sparingly soluble Ca−P minerals. However, the dissolution kinetics may be more complicated in the presence of both carboxylate and hydroxyl groups in organic acids. Here in situ atomic force microscopy (AFM) is used to directly observe the kinetics of nanoscale dissolution on the (010) surface of dicalcium phosphate dihydrate (brushite, CaHPO4·2H2O) in the presence of succinic acid (SA, 0 alcoholic hydroxyl (−OH)), malic acid (MA, 1 −OH), and tartaric acid (TA, 2 −OH), respectively, over a broad concentration range. We demonstrate that the role of dicarboxylic acids varies with the number of alcoholic hydroxyls and that fully deprotonated hydroxy-dicarboxylic acids play a critical role in controlling the dissolution rate of steps and morphology modification of etch pits. Direct AFM imaging shows that only TA can adsorb along specific directions of the [1̅01̅]Cc steps on the brushite (010) surface at pH ≥ 6 to induce the formation of trapezium-shaped etch pits. This depends on specific molecular recognition and stereochemical conformity between hydroxylcarboxyl of TA and atomic [1̅01̅]Cc steps by molecular modeling using density functional theory. The effectiveness of alcoholic hydroxyls can be enhanced by deprotonated brushite interfaces with the increase of the solution pH. This combined AFM and molecular modeling study may provide microscopic insights into understanding P mobilization by dissolution in soils. KEYWORDS: Hydroxy-dicarboxylic acids, Dicalcium phosphate dihydrate, P mobilization, Nanoscale dissolution, AFM, Phosphate fertilizer

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of organic acids (OA) by roots into the rhizosphere.7−10 Secreted citrate, malate (MA), and succinate (SA) derive from intermediates of tricarboxylic acid cycle occurring inside plant cells,11,12 and tartarate (TA) comes from the metabolic product of ascorbic acid (vitamin C) in cells.13,14 In addition to OA species, the pH and concentrations in the rhizosphere are also critical in controlling the dissolution of sparingly soluble P minerals/precipitates. Generally, soluble OA concentrations in bulk soils range from 0.1 μM to 0.1 mM,15 but in the rhizosphere the concentrations are likely to be much higher (about 1−10 mM).16 Moreover, the pH range in the

INTRODUCTION Phosphorus (P) is an essential macronutrient of plants, and its bioavailability is often limited as a result of the low solubility of P precipitates in soils.1,2 After P fertilizer is applied, dissolved orthophosphates can rapidly adsorb and precipitate at various mineral surfaces such as calcite3 and (hydro)oxides of iron and aluminum.4 During precipitation metastable intermediate phases form in calcium (Ca)-rich soils including amorphous Ca−P (ACP) and dicalcium phosphate dihydrate (brushite, CaHPO4·2H2O, DCPD) followed by the transformation to octacalcium phosphate (Ca8(HPO4)2(PO4)4·5H2O, OCP), tricalcium phosphate (α/β-Ca3(PO4)2, α/β-TCP), and the least soluble hydroxyapatite (Ca10(PO4)6(OH)2, HAP) in neutral to alkaline environments,5,6 lowering P bioavailability. The most effective plant strategy to mobilize sparingly soluble P precipitates is the release of relatively large quantities © 2017 American Chemical Society

Received: December 19, 2016 Revised: March 21, 2017 Published: March 23, 2017 3920

DOI: 10.1021/acssuschemeng.6b03105 ACS Sustainable Chem. Eng. 2017, 5, 3920−3928

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ACS Sustainable Chemistry & Engineering

Figure 1. AFM (A) deflection and (B) height images of a brushite (010) surface illustrating the triangular etch pits along the [10̅ 0]Cc, [101]̅ Cc, and [101]Cc directions with the degrees 29−55−96 in water at pH ≥ 4.0. Images A and B, 5.5 × 5.5 μm. (C) Depth profile along line 1 → 2 in (B) showing the monomolecular step height of about 7.6 Å. (D) The interlayer distance between (010) surfaces (step height) is 7.6 Å (half the size of the lattice constant of the b axis) as marked in the side view of the brushite (010) face (Ca, blue; P, gray; O, red; H, white). Projections of the unit cell are shown in a dashed blue rectangle, and the (010) cleavage plane is shown as the black dashed line.

directly explore the role of alcoholic −OH groups in an OA molecule in P release due to dissolution at brushite−water interfaces under various OA concentrations.

rhizosphere can be broad from alkalinity (about 8−9) to acidity (as low as pH 3.6).17 The dissolution of these sparingly soluble Ca−P precipitates by OAs exhibits the H+- and ligand-promoted mechanisms. The ligand-promoted dissolution occurs by various interactions through chelation18 and hydrogen bonding,19 direction specific adsorption20,21 of OA molecules at the mineral surfaces.22,23 Moreover, the presence of alcoholic hydroxyls (−OH group) in OAs may have significant contributions to dissolution after synergistic binding of −OH and −COOH groups to mineral surfaces. Recently, Lindegren et al.24 have shown that the alcoholic −OH group in citrate was critical for the formation of inner sphere surface complexes on goethite surfaces. Chung et al.25 also demonstrated that hydroxycitrate binding to calcium oxalate crystal surfaces is energetically more favorable than citrate. Additionally, succinic acid would prefer to transform to the bisuccinate through deprotonation upon the initial adsorption on Cu (110) surfaces, compared with tartaric acid having two additional alcoholic −OH groups.26 The aim of the present study is to investigate the role of alcoholic −OH groups of dicarboxylic acids in Ca−P dissolution at the nanoscale. A series of in situ AFM experiments combined with molecular modeling using density functional theory were performed, in which the brushite (010) surfaces interacted with succinic acid (SA, (CH2)2(CO2H)2, 0 alcoholic −OH group), malic acid (MA, CH2CH OH(CO2H)2, 1 alcoholic −OH group), and tartaric acid (TA, (CHOH)2(CO2H)2, 2 alcoholic −OH groups), respectively, under chemical conditions that mimic rhizosphere environments. The mineral brushite was chosen because (1) it is an important product in the direct application of fertilizers to soils;6 (2) as a model mineral it is large enough to be employed as substrates for in situ atomic force microscopy (AFM) observations to monitor nanoscale dissolution of Ca−P crystals.27 To our knowledge, there has been no in situ experimental effort combined with molecular modeling to

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EXPERIMENTAL SECTION

Brushite Crystal Synthesis. Brushite single crystals were synthesized by a gel method28 and were characterized as single phase by Bruker D8 X-ray diffraction (Billerica, Massachusetts, USA).27 In Situ AFM Imaging of Brushite Dissolution. All in situ dissolution experiments were performed using an Agilent 5500 AFM (Agilent 5500, Phoenix, USA) operating in contact mode. An optically clear brushite crystal (about 1.0 × 1.0 × 0.1 mm in size) was cleaved to expose a fresh (010) surface. Water at pH of 3.7−9.0 was adjusted by the addition of 0.01 M NaOH or HNO3. We used water as a reference solution to counter the effect that the true dissolution rates can be complicated by the stress applied from AFM tips to the crystal surface.29 At a specific pH of 5.9, water was passed over the brushite (010) cleavage surface in the AFM fluid cell which dominates the macroscopic habit at a constant flow rate of 1 mL/min to ensure surface-controlled reaction rather than diffusion control,27 and dissolution immediately occurred with the formation of typical triangular etch pits on the exposed surfaces (Figure 1A, B). These were used to establish the crystallographic orientation of the substrates. Then, the solutions of tartaric acid (L-TA), malic acid (LMA), or succinic acid (SA) at different concentrations (0.01 mM to 10.0 mM) and pH values (4.0, 6.0, and 8.0) were passed over the cleaved brushite crystals. All reaction solutions were prepared from high-purity solids (Sigma-Aldrich, St. Louis, Missouri) using purity water (resistivity >18 MΩ-cm at 25 °C, pH 5.8−6.0) from a two-step purification treatment including triple distillation (YaR, SZ-93, Shanghai, China) and deionization (Milli-Q, Billerica, MA, USA). The AFM images were collected using Si3N4 tips (nanosensor point probe plus contact with a force constant of 0.2 N/m) with scan rates of 3−5 Hz and an average scan time of 75 s for scan area of 7 × 7 μm2. Five different locations of three different crystals per solution composition were imaged to ensure reproducibility of the results. The details on the calculation of the retreat rates during dissolution can be found in ref 27. 3921


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Figure 2. Retreat velocity of the [1̅00]Cc, [101̅]Cc, and [101]Cc steps of a brushite (010) surface dissolved in the presence of SA, MA, or TA over a broad concentration range 0.01−10.0 mM at (A) pH 8.0, (B) 6.0, and (C) 4.0. The black horizontal line and two dashed lines in A−C indicate the retreat rate range in water at pH 4.0−8.0.

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Molecular Modeling Using Density Functional Theory. The adsorption of SA, L-MA, or L-TA on the neutral or partially deprotonated [1̅01̅]Cc steps on the brushite (010) surface has been calculated using the slab models with density functional theory (DFT). Gradient-corrected functional Perdew−Burke−Ernzerhof (PBE)30 was employed with the all-electron double numerical plus polarization (DNP) basis set.31 The long-range dispersion was considered on the basis of Grimme’s correction (PBE-D).32 The conductor-like screening model (COSMO)33 was used to simulate the aqueous environment for the calculations. The surfaces were cleaved from the fully optimized 2 × 2 × 3 brushite primitive unit cell with the size of u = 16.175 Å, v = 11.884 Å, and θ = 79.1° and a total of 312 atoms, that is, 24(CaHPO4· 2H2O). The vacuum slabs were set to the thickness of 20 Å. Sodium ions (Na+) were added to the surfaces to neutralize the cell. The optimizations were carried out using the Dmol3 program.34 The convergence tolerances of energy, maximum force, and maximum displacement were set to 10−5 au, 0.002 au/Å, and 0.005 Å, respectively. The binding energy was calculated as ΔE = Ead − (Emol + Eslab), where Ead is the total energy of the complex formed by the adsorption of SA, MA, or LA molecules on the [1̅01̅]Cc steps. Emol and Eslab are the total energies of the isolated molecules and the surface, respectively. The PBE-D/DNP optimized unit cell parameters for brushite, a = 5.942 Å, b = 14.83 Å, c = 6.379 Å, β = 117.0°, are in good agreement with the experimental data (a = 5.812 Å, b = 15.18 Å, c = 6.239 Å, β = 116.4°).35 The average deviation between theory and experiment is only 1.8%, indicative of reliability of the current theoretical modeling.

RESULTS AND DISCUSSION

Dissolution on (010) Brushite Surfaces in Water. Following injection of water (pH ≥ 4.0), slow dissolution occurred on the exposed (010) surfaces with the formation of shallow, triangular etch pits with the degrees 29−55−96 (Figure 1A). The measured depth of shallow etch pits (7.6 Å, Figure 1B, C) exactly matches the interlayer distance35 (Figure 1D). The interlayer between (010) surfaces covers a corrugated Ca2+ with HPO42− bilayer structure between the two water layers along the b axis (b = 15.18 Å) according to a unit cell with four CaHPO4·2H2O molecules (Figure 1D).36 The weak H-bonding of water molecules creates a cleavage (010) plane perpendicular to the b-axis.36,37 At pH 8.0, 6.0, or 4.0, step retreat rates along the [1̅00]Cc direction were 5.1 ± 0.4 (n = 5, the number of crystals which were imaged), 4.8 ± 0.5 (n = 10), and 2.9 ± 0.4 (n = 5) nm/s, respectively (Figure S1); spreading velocities along the [101̅]Cc direction were 2.7 ± 0.2 (n = 5), 2.6 ± 0.3 (n = 10), and 2.3 ± 0.2 (n = 5) nm/s, respectively; these results are consistent with the rate minimum of the brushite dissolution, especially at lower pH.38 The [101]Cc steps have the lowest dissolution rates, 0.5 ± 0.4 (n = 5), 0.5 ± 0.4 (n = 10), and 1.2 ± 0.3 (n = 5) nm/s, respectively, at pH 8.0, 6.0, and 4.0 (Figure S1). Effects of SA, MA, and TA on the Brushite (010) Surface Dissolution. Dissolution Kinetics. In the presence of 0.01−10.0 mM SA (pH 8.0), retreat rates of pits along the 3922


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ACS Sustainable Chemistry & Engineering [10̅ 0]Cc, [101]̅ Cc, and [101]Cc directions were hardly changed, compared to that in water (pH 8.0) (Figure 2A). However, the presence of MA or TA (≥1.0 mM at pH 8.0) significantly promoted dissolution along both [1̅00]Cc and [101̅]Cc directions with the concentration increase (Figure 2A), and little changes in the retreat rates along the [101]Cc steps were observed (Figure 2A). Surprisingly, as the pH was lowered to 6.0, the presence of MA or TA in the concentration range of 0.01−10.0 mM had no obvious effects on pit retreat velocities of three steps compared to that in pure water (Figure 2B). In contrast, the pit retreat velocities along the [1̅00]Cc, [101̅]Cc, and [101]Cc directions declined with increasing the SA concentration in the range of 0.01 to 0.5 mM and reached a minimum of 2.3 ± 0.1 (n = 3), 2.0 ± 0.1 (n = 3), and 0 (n = 3) nm/s, respectively, at 0.5 mM (Figure 2B). When the SA concentrations were further increased to 10.0 mM, the rate of pit retreat suddenly increased to 8.5 ± 0.2 (n = 3), 4.1 ± 0.1 (n = 3), and 5.6 ± 0.2 (n = 3) nm/s, respectively, along the [1̅00]Cc, [101̅]Cc, and [101]Cc directions (Figure 2B). When the pH was further lowered to 4.0, the enhancement in the pit retreat speeds along three step directions following the order TA > SA > MA at all concentrations tested (Figure 2C), especially for the [101]Cc direction of steps reaching 11.6 ± 0.2 (n = 3), 9.4 ± 0.2 (n = 3), and 6.1 ± 0.3 (n = 3) nm/s in 0.5 mM TA, SA, and MA solutions, respectively (Figure 2C). Modification of Etch Pit Morphology. When both pH (6.0−9.0) and TA concentrations were raised, a new step [1̅01̅]Cc appeared (shown by arrows in Figure 3B′ and C′), that

is, the characteristic triangular shape of etch pits was changed to a four-sided trapezium. This became dominant in 10.0 mM TA solutions at pH 9.0 by a faster transformation (within 75 s) (Figure 3C and C′). We noted that there exhibited a systematic trend that the simultaneous increase of the pH values (6.0−9.0) and TA concentrations (1.0−10.0 mM) can induce the express of the [1̅01̅]Cc steps (Table 1), whereas this change in the pit shape was absent in SA or MA solutions even at higher concentrations (100 mM) and the same pH range 6.0−9.0 (Table 1). Moreover, the pH values in 10.0 mM TA solutions were increased to 9.0 from 6.0; spreading velocity along the newly expressed [1̅01̅]Cc direction decreased to 10.1 ± 0.5 (n = 3) nm/s from 18.9 ± 0.8 (n = 3) nm/s (Figure S2). The emergence of another new direction of the [102]Cc steps was observed (shown by arrows in Figure 4), intersecting at the [1̅00]Cc and [101]Cc steps with 84 ± 3° and 151 ± 3°, respectively, following about 34 min of exposure to 0.1 mM SA solutions at pH 4.0 (Figure 4A). These new steps rapidly appeared after only 3 min of exposure to 0.5 mM SA solutions at pH 4.0 (Figure 4B) and completely replaced the [101]Cc steps at 10 min (Figure 4B). This results in distorted triangular pits with the degrees 29−67−84 (Figure 4B). Moreover, when the SA concentrations were increased to 10.0 mM from 0.5 mM at pH 4.0, the pit retreat rates along the [102]Cc direction increased to 23.0 ± 0.3 nm/s (n = 3) from 9.0 ± 0.5 nm/s (n = 3) (Figure S3A), and the retreat rate of the [102]Cc steps relative to the [101̅]Cc and [1̅00]Cc steps was about 2.6 and 1.9, respectively, at 10.0 mM (Figure S3B and C). For TA, only a narrow concentration range 0.1−1.0 mM (pH 4.0) produced the [102]Cc steps (Table 2, Figure S4), this new step disappeared, and the normal triangular etch pit shape recovered when the concentrations are higher than 1.0 mM. Moreover, the retreat rates of the [102]Cc steps also increased to 16.7 ± 0.7 nm/s (n = 3) at 1.0 mM from 10 ± 0.2 nm/s (n = 3) at the TA concentration of 0.5 mM (Figure S3A). The retreat rate of the [102]Cc steps (at 1.0 mM) relative to the [101̅]Cc and [1̅00]Cc steps was 2.0 and 1.5, respectively (Figure S3B and C). In addition, MA (pH 4.0) at all concentrations tested (0.5−10 mM) cannot change the pit morphology (Table 2, Figure S4). H+-Promoted Dissolution on the Brushite (010) Surface. In order to assess which one of two mechanisms of H+- and ligand-promoted dissolution is dominant at pH 4.0, water was adjusted by 0.01 M HNO3 to generate pH 3.7, the newly formed [102]Cc steps gradually replaced the [101]Cc steps after 11 min of H+-promoted dissolution (Figure 5A). Distorted triangular pits with the degrees 29−67−84 formed in the presence of only H+ ions (Table 2). The use of HCl at pH 3.7 led to the same result as HNO3 that ruled out the anion effects. After 4, 8, and 11 min of dissolution, the retreat rates of the [101]Cc steps gradually increased to 11.1 ± 0.5 (n = 3), 12 ± 0.5 (n = 3), and 13.3 ± 0.5 nm/s (n = 3) from 1.0 ± 0.3 nm/ s (n = 5) (Figures 5C and S5). Finally, the [101]Cc steps disappeared, and the normal triangular pits completely converted to distorted ones (Figure 5A-C). The newly triangular pits dissolved in water (pH 3.7), having anisotropic spreading velocities along the [1̅00]Cc, [101̅]Cc, and [102]Cc directions, 6.1 ± 0.4 (n = 5), 5.0 ± 0.3 (n = 5), and 8.4 ± 0.3 (n = 5) nm/s, respectively (Figures S5 and 5C). Moreover, the [101]Cc step exhibits the higher H+ charge density, compared to that of the [102]Cc step (Figure 5B and C). Molecular Modeling on the Adsorption of TA onto the [1̅01̅]Cc Steps. To understand TA molecules specifically

Figure 3. AFM deflection images showing the etch pit morphology of the brushite (010) surfaces dissolving in 10.0 mM TA at (A and A′) pH 6.0, (B and B′) pH 8.0, or (C and C′) pH 9.0 after 10 min, 3 min, and 75 s of dissolution, respectively. Arrows in B′ and C′ indicate the newly expressed [1̅01̅]Cc steps. 3923


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ACS Sustainable Chemistry & Engineering Table 1. Summary of Observed Changes in the Pit Morphology at pH 6.0−9.0 organic acid

SA

MA

pH concn (mM) pit shape

6.0−9.0 0.01−100

6.0−9.0 0.01−100

TA 6.0 ≥10 trapezium

8.0 ≥5.0 trapezium

9.0 ≥1.0 trapezium

Figure 4. Time sequence AFM (deflection mode) images showing that the [101]Cc steps gradually were replaced by the [102]Cc steps (demonstrated by arrows) in the presence of (A) 0.1 mM and (B) 0.5 mM SA at pH 4.0. Note the [102]Cc steps intersecting with [1̅00]Cc and [101]Cc steps at 84 ± 3° and 151 ± 3°, respectively, were observed after 34 min of dissolution in (A) 0.1 mM SA, and the [102]Cc steps completely replaced the [101]Cc steps after 10 min of dissolution in (B) 0.5 mM SA. This results in the formation of a distorted triangular pit with the degrees 29−67−84.

pH 6.0 or 9.0 can use one −COO− group (with two O atoms) bound to a Ca cation [d1(COO−···Ca) = 2.54 and 2.60 Å; d2(COO−···Ca) = 2.51 and 2.64 Å, respectively] (Figure 6A and B). One alcoholic −OH group and another elongated −COO− group bind adjacent Ca cation [d(HO···Ca) = 2.47 and 2.44 Å; d(COO−···Ca) = 2.41 and 2.47 Å, respectively]. Another alcoholic −OH group forms H-bonding to OPO3H [d(HO···OPO3H) = 1.70 and 1.66 Å at pH 6.0 and 9.0, respectively] (Figure 6A and B). As a result, the average binding energy of a single TA molecule on neutral and deprotonated [1̅01̅]Cc steps can be calculated to be −210.9 and −300.0 kJ/mol, respectively (Figure 7), whereas binding of SA and MA to the [1̅01̅]Cc steps is less favorable with average binding energies of −202.3 and −201.5 kJ/mol on neutral [1̅01̅]Cc steps, as well as 256.6 and 284.8 kJ/mol on deprotonated [1̅01̅]Cc steps, respectively (Figure 7). Roles of Alcoholic −OH Groups of TA/MA at Higher pH: Negatively Charged Brushite−Water Interfaces Is Important. At pH 8.0, the major species in the SA (pKa1 = 4.2, pKa2 = 5.6), MA (pKa1 = 3.4, pKa2 = 5.2), and TA (pKa1 = 2.9, pKa2 = 4.4) solutions are all dianionic,40,42 and the presence of MA or TA at concentrations greater than 1.0 mM strongly enhanced retreat rates of the [1̅00]Cc and [101̅]Cc steps (Figure 2A); whereas SA at 0.01−10.0 mM (pH 8.0) had no effects on retreat velocities of three steps (Figure 2A). The most likely explanation is the lack of an alcoholic −OH group in a SA molecule. When the pH of the MA and TA solutions was decreased to 6.0, no changes in retreat velocities of three steps were observed compared to that in water (Figure 2B). This may suggest that the brushite mineral interface is another factor in influencing the role of alcoholic −OH groups. As the pH increases to 8.0 from 6.0, the brushite interface presumably

Table 2. Summary of Observed Changes in the Pit Morphology at pH 4.0 organic acid

SA

MA

pH concn (mM) pit shape

4.0 0.1−10 distorted triangle

4.0 0.1−10

TA 4.0 0.1−1.0 distorted triangle

4.0 >1.0

interacting with the [10̅ 1]̅ Cc steps to form four-sided trapeziumlike etch pits with increasing the pH (Figure 3, Table 1), we used molecular modeling with energy minimization to calculate binding energies of TA, MA, and SA docking to [1̅01̅]Cc steps with increasing pH from 6.0 to 9.0 (Figure 6). Ideally, the adsorption of acids on the [1̅01̅]C steps will also involve the interactions with an underlying terrace on the [010] direction due to the [1̅01̅]C steps having only a height of 7.6 Å. However, for adsorption of dicarboxylic acids on the fully water-covered (010) surface, it occurs only via relatively weak hydrogen bonding (Figure S6), thus we chose the [101] steps of brushite as the dominant adsorption sites. As pH increased to 9.0 from 6.0, brushite can be partially deprotonated due to approaching the isoelectric point of brushite (6.1).39 The dianionic species of two −COOH groups in a TA molecule is present in solutions, and alcoholic −OH groups do not dissociate.40,41 In addition, no dissociation of interfacial water has been observed during the geometrical optimizations; the DFT calculations in this work only concern the adsorption of dicarboxylic acid on the preferable [1̅01̅]Cc steps. The predicted deprotonated sites on the [10̅ 1]̅ Cc steps can be from H2O molecules, because H2O molecules are on the outermost surfaces and the bulk protons of HPO4− ions are along the [1̅01̅]Cc directions (Figure 6B). The stereochemistry of the TA binding reveals that TA at either 3924


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Figure 5. Evolution of a distorted triangular etch pit in acidic water (pH 3.7). (A) Time sequence AFM (deflection mode) images showing that [101]Cc steps were replaced progressively by the [102]Cc steps after 11 min reaction of dissolution in water (pH 3.7). Distorted triangular pits along the [10̅ 0]Cc, [101]̅ Cc, and [102]Cc directions with the degrees 29−67−84 were formed. (B) The pit atomic structure of the [10̅ 0]Cc, [101]̅ Cc, [101]Cc, and [102]Cc steps (Ca, blue; O, red; P, gray; H, white). (C) Schematic of etch pit evolution. Yellow lines indicate the [101]Cc steps, black arrows indicate the relative step retreat velocities, and red circles represent the density of Ca2+ cations, the coordinated oxygen anions, or protonated oxygen anions along four crystallographically distinct steps in acidic water at pH 3.7.

Figure 6. Structural conformation of a single SA, MA, and TA molecule adsorbed to (A) neutral and (B) deprotonated [1̅01̅]Cc steps on a brushite (010) surface (Ca, blue; O, red; P, gray; H, white). An OA molecule is highlighted as ball−stick (C, gray; O, red; H, white; Na, brown green). The number is distance in angstroms.

further increasing the pH to 9.0, TA caused intense dissolution on the brushite (010) surface with the increase of the pit density (Figure 3C and C′). Because the alcoholic −OH groups of TA and MA did not dissociate even at pH 9.0,41 dissolution was enhanced only by deprotonated mineral interfaces. However, so far, we have not obtained the sight of the overall dissolution process of the deprotonated brushite (010) face in the presence of hydroxy-dicarboxylic acids by parametrized kinetic Monte Carlo simulation or ab initio molecular dynamics simulation because the piece of work is very computationally demanding. Generally, there are three main factors that affect the interfacial structure: adsorbate-mineral interactions, intermolecular interactions in the interfacial fluid layer, and entropy.43−46 First, although the more negatively charged mineral surface can be repulsive to the −COO− groups at high pH,47,48 the energy for TA molecules binding to the

Figure 7. DFT-calculated binding energy of SA (0 alcohol−OH), MA (1 alcohol−OH,) and TA (2 alcohol−OH) molecules with calcium of a brushite (010) surface at pH 6.0 or 9.0.

deprotonates because of brushite having an isoelectric point of about 6.2.39 The negatively charged brushite (010) surfaces due to deprotonation may be favorable to the ligand-promoted dissolution by the presence of alcoholic −OH groups. After 3925


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solution pH because of the presence of the stronger interaction between hydroxy-dicarboxylic acids molecules and the deprotonated (010) brushite surface. At acidic pH values, the deprotonation degree of OAs increases in the order of MA < TA < SA, and this correlates with the alcoholic hydroxyldependent adsorption of OAs on the solid surface. This study shows that mineral dissolution kinetics is complicated in the presence of organic ligands, especially for hydroxy-carboxylic acids using multiple binding sites that allow for specific molecular recognition and/or stereochemical interacting with crystal surfaces to initiate H+- and ligand-promoted dissolution. These direct AFM observations combined with molecular modeling may aid in understanding mobilization of sparingly soluble phosphate minerals/fertilizers in a more complex rhizosphere−soil environment.

deprotonated (010) surface of brushite may increase due to the stronger interaction between the alcoholic −OH groups and the more negatively charged surface.49,50 Second, the stronger hydrogen-bonding interactions among interfacial water molecules tend to pull alcohol−OH groups to the brushite solid surface.45,51 This can be weakened at higher pH due to a stronger disruption on the hydrogen-bonding structure of water at a more negatively charged mineral surface.52,53 The role of interfacial water may be relatively passive for the alcoholic −OH groups that can bind to the mineral surface more strongly than water.54−57 Finally, the stereochemical conformity will lead to the increase of binding energy between TA and more negatively charged surfaces (Figures 6B and 7). Roles of Alcoholic −OH Group of OA at Acidic pH. H+Promoted Dissolution. Without OA, the [101]Cc steps exhibit higher retreat rates in the presence of H+ (pH 3.7) (Figure S5) and finally disappear (Figure 5A). This can be related to the protonation of oxygen sites along this step direction. In addition to higher solubility of dihydrogen phosphate in water compared to monohydrogen phosphate, protonation of O atoms on the brushite (010) surface presumably causes the increase of charge density that improves the penetration of water into the mineral.37 The [101]Cc steps of etch pits is the most attractive for water due to the presence of the highest H+ charge density, whereas the [102]Cc steps exhibit the least affinity to water (Figure 5 B). This suggests that the [102]Cc steps are more stable than the [101]Cc steps to explain why the [102]Cc steps can replace the [101]Cc steps in acidic water. H+ from Dissociation of OA on the Brushite (010) Surface Participating in Interfacial Dissolution. The [102]Cc steps can also replace the [101]Cc steps in the presence of 0.1−10 mM SA (pH 4.0) (Figure 4 and Table 2). The [102]Cc steps occurred at lower pH (Figure 5A), and interfacial pH values should be more acidic than that of the bulk SA solutions (0.1−10.0 mM, pH 4.0) due to the deprotonation of −COOH groups of SA molecules adsorbed at the interface.58 Furthermore, the deformed triangular pits were also induced by TA in a narrow concentration range 0.1−1.0 mM (pH 4.0) (Table 2), suggesting that the interfacial deprotonation of the high concentration TA (>1.0 mM) is not favorable. Humblot et al.26 attributed this behavior to an increased propensity for intermolecular interactions in the TA system, that is, hydrogen bonding interactions between adjacent monotartrate molecules.59 This promotes the formation of locally high coverage of TA molecules at the mineral interface,26,59 thereby lowering the possibility for the formation of thermodynamically stable bitartrate by fully deprotonation. The preferential deprotonation of TA (pH 4.0) compared to MA at the interface can be attributed to stereochemical conformity due to the presence of an additional alcoholic −OH group in TA molecules.60 Thus, an additional alcoholic −OH group can make the interface more acidic due to deprotonation. It is also because a stronger intermolecular interaction in the MA system effectively interferes with the deprotonation of MA at pH 4.0 at the brushite interface.61,62 The degree of deprotonation of OAs at interfaces closely correlates with the nanoscale dissolution kinetics of the brushite (010) face (Figure 2C).

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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.6b03105. Retreat velocities, Figures S1−S3 and S5; AFM images of etch pits, Figure S4; and calculated adsorption of SA, Figure S6 (PDF)

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AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected] (L.W.). *E-mail: [email protected] (B.W.). ORCID

Lijun Wang: 0000-0001-7125-9480 Baoshan Wang: 0000-0003-3417-9283 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the National Natural Science Foundation of China (41471245 and 41071208 to L.J.W. and 21573165 to B.S.W.), a Specialized Research Fund for the Doctoral Program of Higher Education (20130146110030), and the Fundamental Research Funds for the Central Universities (2662015PY206, 2662015PY116).

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REFERENCES

(1) López-Arredondo, D. L.; Leyva-González, M. A.; GonzálezMorales, S. I.; López-Bucio, J.; Herrera-Estrella, L. Phosphate nutrition: Improving low-phosphate tolerance in crops. Annu. Rev. Plant Biol. 2014, 65, 95−123. (2) Shen, J. B.; Yuan, L. X.; Zhang, J. L.; Li, H. G.; Bai, Z. H.; Chen, X. P.; Zhang, W. F.; Zhang, F. S. Phosphorus dynamics: From soil to plant. Plant Physiol. 2011, 156, 997−1005. (3) Wang, L. J.; Ruiz-Agudo, E.; Putnis, C. V.; Menneken, M.; Putnis, A. Kinetics of calcium phosphate nucleation and growth on calcite: Implications for predicting the fate of dissolved phosphate species in alkaline soils. Environ. Sci. Technol. 2012, 46, 834−842. (4) Wang, L. J.; Putnis, C. V.; Ruiz-Agudo, E.; Hövelmann, J.; Putnis, A. In situ imaging of interfacial precipitation of phosphate on goethite. Environ. Sci. Technol. 2015, 49, 4184−4192. (5) Wang, L. J.; Nancollas, G. H. Calcium orthophosphates: Crystallization and dissolution. Chem. Rev. 2008, 108, 4628−4669. (6) Wang, L. J.; Lu, J. W.; Xu, F. S.; Zhang, F. S. Dynamics of crystallization and dissolution of calcium orthophosphates at the nearmoleculear level. Chin. Sci. Bull. 2011, 56, 713−721.

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CONCLUSIONS At alkaline pH values, hydroxy-dicarboxylic acids significantly increase the dissolution rates of brushite crystals. The effectiveness of alcoholic hydroxyls can be enhanced by deprotonated brushite interfaces with the increase of the 3926


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ACS Sustainable Chemistry & Engineering (7) Lambers, H.; Raven, J. A.; Shaver, G. R.; Smith, S. E. Plant nutrient-acquisition strategies change with soil age. Trends Ecol. Evol. 2008, 23, 95−103. (8) Stutter, M. I.; Shand, C. A.; George, T. S.; Blackwell, M. S. A.; MacKay, R. L.; Richardson, A. E.; Condron, L. M.; Turner, B. L.; Haygarth, P. M. Recovering phosphorus from soil: A root solution? Environ. Sci. Technol. 2012, 46, 1977−1978. (9) Lambers, H.; Martinoia, E.; Renton, M. Plant adaptations to severely phosphorus-impoverished soils. Curr. Opin. Plant Biol. 2015, 25, 23−31. (10) Lambers, H.; Finnegan, P. M.; Jost, R.; Plaxton, W. C.; Shane, M. W.; Stitt, M. Phosphorus nutrition in Proteaceae and beyond. Nature Plants 2015, 1, 15109. (11) Raghothama, K. G. Phosphate acquisition. Annu. Rev. Plant Physiol. Plant Mol. Biol. 1999, 50, 665−93. (12) Plaxton, W. C.; Tran, H. T. Metabolic adaptations of phosphatestarved plants. Plant Physiol. 2011, 156, 1006−1015. (13) Loewus, F. A.; Stafford, H. A. Observations on the incorporation of C14 into tartaric acid and the labeling pattern of D-glucose from an excised grape leaf administered L-ascorbic acid-6-C. Plant Physiol. 1958, 33, 155−156. (14) DeBolt, S.; Cook, D. R.; Ford, C. M. L-Tartaric acid synthesis from vitamin C in higher plants. Proc. Natl. Acad. Sci. U. S. A. 2006, 103, 5608−5613. (15) Jones, D. L.; Darah, P. R.; Kochian, L. V. Critical evaluation of organic acid mediated iron dissolution in the rhizosphere and its potential role in root iron uptake. Plant Soil 1996, 180, 57−66. (16) Ryan, P. R.; James, R. A.; Weligama, C.; Delhaize, E.; Rattey, A.; Lewis, D. C.; Bovill, W. D.; McDonald, G.; Rathjen, T. M.; Wang, E.; Fettell, N. A.; Richardson, A. E. Can citrate efflux from roots improve phosphorus uptake by plants? Testing the hypothesis with nearisogenic lines of wheat. Physiol. Plant. 2014, 151, 230−242. (17) Neumann, G.; Römheld, V. Root excretion of carboxylic acids and protons in phosphorus-deficient plants. Plant Soil 1999, 211, 121− 130. (18) Perry, T. D.; Duckworth, O. W.; Kendall, T. A.; Martin, S. T.; Mitchell, R. Chelating ligand alters the microscopic mechanism of mineral dissolution. J. Am. Chem. Soc. 2005, 127, 5744−5745. (19) Norén, K.; Persson, P. Adsorption of monocarboxylates at the water/goethite interface: The importance of hydrogen bonding. Geochim. Cosmochim. Acta 2007, 71, 5717−5730. (20) Teng, H. H.; Chen, Y.; Pauli, E. Direction specific interactions of 1, 4-dicarboxylic acid with calcite surfaces. J. Am. Chem. Soc. 2006, 128, 14482−14484. (21) Wu, C.; Wang, X.; Zhao, K.; Cao, M.; Xu, H.; Xia, D.; Lu, J. R. Molecular modulation of calcite dissolution by organic acids. Cryst. Growth Des. 2011, 11, 3153−3162. (22) Weng, Y.-X.; Li, L.; Liu, Y.; Wang, L.; Yang, G. Z. Surfacebinding forms of carboxylic groups on nanoparticulate TiO2 surface studied by the interface-sensitive transient triplet-state molecular probe. J. Phys. Chem. B 2003, 107, 4356−4363. (23) Mudunkotuwa, I. A.; Grassian, V. H. Citric acid adsorption on TiO2 nanoparticles in aqueous suspensions at acidic and circumneutral pH: Surface coverage, surface speciation, and its impact on nanoparticle-nanoparticle interactions. J. Am. Chem. Soc. 2010, 132, 14986−14994. (24) Lindegren, M.; Loring, J. S.; Persson, P. Molecular structures of citrate and tricarballylate adsorbed on α-FeOOH particles in aqueous suspensions. Langmuir 2009, 25, 10639−10647. (25) Chung, J.; Granja, I.; Taylor, M. G.; Mpourmpakis, G.; Asplin, J. R.; Rimer, J. D. Molecular modifiers reveal a mechanism of pathological crystal growth inhibition. Nature 2016, 536, 446−450. (26) Humblot, V.; Lorenzo, M. O.; Baddeley, C. J.; Haq, S.; Raval, R. Local and global chirality at surfaces: Succinic acid versus tartaric acid on Cu (110). J. Am. Chem. Soc. 2004, 126, 6460−6469. (27) Qin, L. H.; Zhang, W. J.; Lu, J. W.; Stack, A. G.; Wang, L. J. Direct imaging of nanoscale dissolution of dicalcium phosphate dihydrate by an organic ligand: Concentration matters. Environ. Sci. Technol. 2013, 47, 13365−13374.

(28) Roop Kumar, R.; Wang, M. Biomimetic deposition of hydroxyapatite on brushite single crystals grown by the gel technique. Mater. Lett. 2001, 49, 15−19. (29) Colombani, J. Pitfalls in the measurement of the true dissolution kinetics of soft minerals. Procedia Earth Planet. Sci. 2013, 7, 179−182. (30) Perdew, J. P.; Burke, K.; Ernzerhof, M. Generalized gradient approximation made simple. Phys. Rev. Lett. 1996, 77, 3865. (31) Delley, B. An all-electron numerical method for solving the local density functional for polyatomic molecules. J. Chem. Phys. 1990, 92, 508. (32) Grimme, S. Semiempirical GGA-type density functional constructed with a long-range dispersion correction. J. Comput. Chem. 2006, 27, 1787−1799. (33) Delley, B. The conductor-like screening model for polymers and surfaces. Mol. Simul. 2006, 32, 117−123. (34) Delley, B. From molecules to solids with the DMol3 approach. J. Chem. Phys. 2000, 113, 7756. (35) Curry, N. A.; Jones, D. W. Crystal structure of brushite, calcium hydrogen orthophosphate dihydrate: A neutron-diffraction investigation. J. Chem. Soc. A 1971, 3725−3729. (36) Qiu, S. R.; Orme, C. A. Dynamics of biomineral formation at the near-molecular level. Chem. Rev. 2008, 108, 4784−4822. (37) Lin, T. J.; Heinz, H. Accurate force field parameters and pH resolved surface models for hydroxyapatite to understand structure, mechanics, hydration, and biological interfaces. J. Phys. Chem. C 2016, 120, 4975−4992. (38) Zhang, J. W.; Nancollas, G. H. Unexpected pH dependence of dissolution kinetics of dicalcium phosphate dihydrate. J. Phys. Chem. 1994, 98, 1689−1694. (39) Herschke, L.; Lieberwirth, I.; Wegner, G. Zinc phosphate as versatile material for potential biomedical applications Part II. J. Mater. Sci.: Mater. Med. 2006, 17, 95−104. (40) Weast, R. C. Handbook of Chemistry and Physics, 57th ed.; CRC Press: Cleveland, OH, 1976. (41) Beck, M. T.; Csiszar, B.; Szarvas, P. Acidic dissociation constants of alcoholic hydroxyls of hydroxy-carboxylic acids: Tartaric Acid. Nature 1960, 188, 846−847. (42) Tung, L. A.; King, C. J. Sorption and extraction of lactic and succinic acids at pH > pKal. 1. Factors governing equilibria. Ind. Eng. Chem. Res. 1994, 33, 3217−3223. (43) Fang, F.; Szleifer, I. Effect of molecular structure on the adsorption of protein on surfaces with grafted polymers. Langmuir 2002, 18, 5497−5510. (44) Fang, F.; Szleifer, I. Competitive adsorption in model charged protein mixtures: Equilibrium isotherms and kinetics behavior. J. Chem. Phys. 2003, 119, 1053. (45) Zhang, L.; Liu, W. T.; Shen, Y. R. Competitive molecular adsorption at liquid/solid interfaces: A study by sum-frequency vibrational spectroscopy. J. Phys. Chem. C 2007, 111, 2069−2076. (46) Calle-Vallejo, F.; Sautet, P.; Loffreda, D. Understanding adsorption-induced effects on platinum nanoparticles: An energydecomposition analysis. J. Phys. Chem. Lett. 2014, 5, 3120−3124. (47) Kang, S.; Xing, B. Adsorption of dicarboxylic acids by clay minerals as examined by in situ ATR-FTIR and ex situ DRIFT. Langmuir 2007, 23, 7024−7031. (48) Hwang, Y. S.; Lenhart, J. J. Adsorption of C4-dicarboxylic acids at the hematite/water interface. Langmuir 2008, 24, 13934−13943. (49) Yang, Y.; Yan, W.; Jing, C. Dynamic adsorption of catechol at the goethite/aqueous solution interface: A molecular-scale study. Langmuir 2012, 28, 14588−14597. (50) Gulley-Stahl, H.; Hogan, P. A.; Schmidt, W. L.; Wall, S. J.; Buhrlage, A.; Bullen, H. A. Surface complexation of catechol to metal oxides: An ATR-FTIR, adsorption, and dissolution study. Environ. Sci. Technol. 2010, 44, 4116−4121. (51) Israelachvili, J. N. Intermolecular and Surface Forces; Academic Press: London, UK, 1985. (52) Lis, D.; Backus, E. H. G.; Hunger, J.; Parekh, S. H.; Bonn, M. Liquid flow along a solid surface reversibly alters interfacial chemistry. Science 2014, 344, 1138−1142. 3927


Research Article

ACS Sustainable Chemistry & Engineering (53) Nihonyanagi, S.; Yamaguchi, S.; Tahara, T. Counterion effect on interfacial water at charged interfaces and its relevance to the Hofmeister series. J. Am. Chem. Soc. 2014, 136, 6155−6158. (54) Mian, S. A.; Yang, L. M.; Saha, L. C.; Ahmed, E.; Ajmal, M.; Ganz, E. A fundamental understanding of catechol and water adsorption on a hydrophilic silica surface: Exploring the underwater adhesion mechanism of mussels on an atomic scale. Langmuir 2014, 30, 6906−6914. (55) Okhrimenko, D. V.; Nissenbaum, J.; Andersson, M. P.; Olsson, M. H. M.; Stipp, S. L. S. Energies of the adsorption of functional groups to calcium carbonate polymorphs: The importance of −OH and −COOH groups. Langmuir 2013, 29, 11062−11073. (56) Cooke, D. J.; Gray, R. J.; Sand, K. K.; Stipp, S. L. S.; Elliott, J. A. Interaction of ethanol and water with the {1014} surface of calcite. Langmuir 2010, 26, 14520−14529. (57) Sand, K. K.; Yang, M.; Makovicky, E.; Cooke, D. J.; Hassenkam, T.; Bechgaard, K.; Stipp, S. L. S. Binding of ethanol on calcite: The role of the OH bond and its relevance to biomineralization. Langmuir 2010, 26, 15239−15247. (58) van den Brand, J.; Blajiev, O.; Beentjes, P. C. J.; Terryn, H.; de Wit, J. H. W. Interaction of anhydride and carboxylic acid compounds with aluminum oxide surfaces studied using infrared reflection absorption spectroscopy. Langmuir 2004, 20, 6308−6317. (59) Lawton, T. J.; Pushkarev, V.; Wei, D.; Lucci, F. R.; Sholl, D. S.; Gellman, A. J.; Sykes, E. C. H. Long range chiral imprinting of Cu (110) by tartaric acid. J. Phys. Chem. C 2013, 117, 22290−22297. (60) Kunkel, D. A.; Hooper, J.; Simpson, S.; Beniwal, S.; Morrow, K. L.; Smith, D. C.; Cousins, K.; Ducharme, S.; Zurek, E.; Enders, A. Rhodizonic acid on noble metals: Surface reactivity and coordination chemistry. J. Phys. Chem. Lett. 2013, 4, 3413−3419. (61) Walch, H.; Dienstmaier, J.; Eder, G.; Gutzler, R.; Schlögl, S.; Sirtl, T.; Das, K.; Schmittel, M.; Lackinger, M. Extended twodimensional metal-organic frameworks based on thiolate-copper coordination bonds. J. Am. Chem. Soc. 2011, 133, 7909−7915. (62) Roth, C.; Passerone, D.; Merz, L.; Parschau, M.; Ernst, K. H. Two-dimensional self-assembly of chiral malic acid on Cu (110). J. Phys. Chem. C 2011, 115, 1240−1247.

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Role of Alcoholic Hydroxyls of Dicarboxylic Acids ... - ACS Publications

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