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Direct Imaging of Nanoscale Dissolution of Dicalcium Phosphate Dihydrate by an Organic Ligand: Concentration Matters Lihong Qin,† Wenjun Zhang,† Jianwei Lu,† Andrew G. Stack,‡ and Lijun Wang*,† †

College of Resources and Environment, Huazhong Agricultural University, Wuhan 430070, China Chemical Sciences Division, Oak Ridge National Laboratory, Oak Ridge, Tennessee 37831, United States

‡

S Supporting Information *

ABSTRACT: Unraveling the kinetics and mechanisms of sparingly soluble calcium orthophosphate (Ca−P) dissolution in the presence of organic acids at microscopic levels is important for an improved understanding in determining the effectiveness of organic acids present in most rhizosphere environments. Herein, we use in situ atomic force microscopy (AFM) coupled with a fluid reaction cell to image dissolution on the (010) face of brushite, CaHPO4·2H2O, in citratebearing solutions over a broad concentration range. We directly measure the dependence of molecular step retreat rate on citrate concentration at various pH values and ionic strengths, relevant to soil solution conditions. We find that low concentrations of citrate (10−100 μM) induced a reduction in step retreat rates along both the [1̅00]Cc and [101̅]Cc directions. However, at higher concentrations (exceeding 0.1 mM), this inhibitory effect was reversed with step retreat speeds increasing rapidly. These results demonstrate that the concentration-dependent modulation of nanoscale Ca−P phase dissolution by citrate may be applied to analyze the controversial role of organic acids in enhancing Ca−P mineral dissolution in a more complex rhizosphere environment. These in situ observations may contribute to resolving the previously unrecognized interactions of root exudates (low molecular weight organic acids) and sparingly soluble Ca−P minerals.

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(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), as the thermodynamically stable phase of the final product in neutral to alkaline environments.11 Much evidence suggests that P deficiency can lead to some plants which can directly modify the root structures.12 One plant strategy to relieve this deficiency is to release large amounts of low molecular weight (LMW) organic acids by cluster roots into the rhizosphere to dissolve sparingly soluble inorganic P.13−15 Concentrations of organic acids in the bulk soil solution typically range from 0.1 μM to 0.1 mM but are estimated to exceed 50 mM in the rhizosphere of plants with cluster roots but are 90% of the total root exudates under P deficiency.17,18 The dissolution rates of phosphate minerals/precipitates can be greatly accelerated in soil in the presence of organic acids, depending on soil type and speciation and concentration of organic acids.19 Typically, however, most plants exude much lower amounts (10 points). All data with their mean values ± standard deviation (SD) of three independent sets of experiments are presented. C

dx.doi.org/10.1021/es402748t | Environ. Sci. Technol. XXXX, XXX, XXX−XXX

Environmental Science & Technology

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Figure 4. Retreat velocity of (A) the [1̅00]Cc and (B) [101̅]Cc steps for brushite crystals dissolved in low citrate concentration solutions (10−100 μM) at varying pH (4.0−8.0). The [101]Cc steps were completely inhibited (a step velocity of zero) by citrate over the above concentration range. However, introduction of relatively high concentration citrate (>0.1 mM) increases the brushite dissolution rate by increasing the etch pit density and movement velocity of the [1̅00]Cc and [101̅]Cc steps. AFM deflection images of brushite surfaces dissolving in aqueous solutions of citrate with (C) 0.1 mM or (D) 0.5 mM citrate (image taken 150 s after injection of 0.5 mM citrate, pH 5.3). Note the high etch pit density attained at 0.5 mM citrate. Images C and D, 6 × 6 μm. The blue rectangle in (A and B) indicates the retreat velocity of the [1̅00]Cc and [101]̅ Cc steps, respectively, in deionized water (pH 5.8).

Figure 5. Retreat velocity of (A) the [101]Cc, (B) [1̅00]Cc, and (C) [101̅]Cc steps for brushite crystals dissolved in the presence of citrate at higher concentrations ranging from 1.0 to 30.0 mM and varying pH (4.0−8.0). The blue rectangle in (A) indicates the retreat velocity of the [101]Cc steps in deionized water (pH 5.8).

axis.31 According to its atomic structure, brushite is a noncentrosymmetric monoclinic crystal with four CaHPO4· 2H2O molecules per unit cell.31 The structure is composed of corrugated rows of calcium cations with HPO42− anions such that adjacent calcium (or phosphate) clusters are at alternate heights,30 forming a corrugated bilayer structure (Figure 1). Water molecules bound to the calcium cations point outward at the top and bottom edges of these bilayers, resulting in layers of water between the Ca2+ and HPO42− containing sheets (Figure 1).31 For this reason, the fully hydrated {010} faces have a relatively low interfacial energy26 of 4.5 mJ m−2 compared with other Ca−P phases such as 8 mJ m−2 for apatite.32 It is also clear that dissolution involves the removal of surface steps and then the step orientation plays a role in dissolution due to their anisotropic nature. The [101]Cc step is the

other forms of Ca−P. The crystal surface composition/ structure on different samples was constant and unchanged using XRD, SEM-EDX, and XPS. Dissolution Features on Brushite (010) Cleavage Surfaces. Dissolution of brushite mainly proceeded by the formation and spreading of etch pits, which were observed on the smooth (010) crystal surfaces by SEM (Figure 2C) and AFM (Figure 2D). In situ AFM experiments show that brushite crystals in deionized water or citrate solutions (Figure 3A) inside the AFM fluid cell dissolve on atomic steps to generate the same etch pits with the degrees 29−55−96 triangular shapes. Owing to the chiral nature of this crystal, the (010) and (01̅0) faces have crystallographically distinct steps along the [101]Cc, [101̅]Cc, and [1̅00]Cc directions (Figure 2D), Cc is recommended to describe brushite crystals with a unique bD



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calcium-terminated polar step,31 whereas the [1̅00]Cc step is OH terminated with the slowest growth kinetics (Figure 1).33 Shallow etch pits with depths of about 0.77 nm, that is, close to 0. 76 nm (Figure 1), half the size of the lattice constant of the b axis, formed randomly on the surface (Figure 3B and C). No deep etch pits were observed. It is interesting to note that the triangular half a unit cell etch pits do not reverse in successive monolayers during dissolution. This seems to confirm a Cc space group rather than the C2/c space group. Brushite crystals dissolve in deionized water at pH 5.8 (Figure 2D) along the [1̅00]Cc, [101̅]Cc, and [101]Cc directions, having anisotropic spreading velocities, 4.97 ± 0.22 nm/s (n = 3, the number of crystals which were imaged), 2.48 ± 0.15 nm/s (n = 3) (Figure 4A and B) and 0.50 ± 0.08 (n = 3) (Figure 5A), respectively. The [101]Cc steps have the lowest dissolution rate in water regardless of pH. However, in low concentration citrate solutions (10−100 μM, pH 4.0−7.7), step retreat rates along the [1̅00]Cc direction are the pH- and concentration-dependent (Figure 4A). When both pH and citrate level were elevated, the obvious enhancement in dissolution inhibition along the [1̅00]Cc direction was observed (Figure 4A), whereas dissolution for the [101̅]Cc steps is only concentration-dependent, that is., step retreat rates decreased with increasing citrate concentration (Figure 4B). However, dissolution was promoted when the concentration of citrate was further increased: the step velocity of the [1̅00]Cc and [101̅]Cc abruptly increased from about 3.10 and 1.9 nm/s at 0.1 mM (100 μM) citrate (pH 5.3), to about 5.01 and 2.7 nm/s at 0.5 mM citrate, an increase of 62 and 42%, respectively. In addition to the increase of the movement velocity of these two steps, introduction of 0.5 mM citrate concentration also increased the etch pit density (Figure 4C and D). The dissolution of the [101]Cc steps was completely inhibited (no observable velocity) by citrate ranging from 10 to 100 μM at pH 4.0−7.7. However, dissolution was initiated when solution pH values were greater than 7.0 and citrate concentrations were further increased (>5 mM) (Figure 5A). The presence of 10 or 30 mM citrate (pH 7.7), increased the [101]Cc step velocities to 1.29 ± 0.02 (n = 4) and 2.37 ± 0.05 nm/s (n = 4), respectively (Figure 5A). Similar results were obtained for the [1̅00]Cc steps at pH 7.7: the dissolution rates rapidly increased to 9.64 ± 0.14 nm/s (n = 3) at 30 mM citrate, from 6.29 ± 0.09 (n = 4), 6.8 ± 0.10 (n = 3), or 7.47 ± 0.11 nm/s (n = 3) at citrate concentrations of 1.0, 5.0, or 10.0 mM, respectively (Figure 5B). At pH values below 7.0, the velocities of the [10̅ 0]Cc steps remained almost constant (about 6.30 nm/s) in the presence of all tested citrate (1.0−30.0 mM) (Figure 5B). For the [101̅]Cc steps, spreading velocity gradually increased with decreasing pH and increasing citrate concentrations (Figure 5C). Modification of Etch Pit Morphology by Citrate. Although the number of brushite specimens was limited, the data showed systematic changes in etch pit shapes during dissolution. AFM deflection images illustrated the etch pit transformation from the normal triangular shape in water to a four-sided trapezium following exposure of the brushite (010) surface to the citrate solution (Figure 6). The length of a newly expressed step (shown by an arrow in Figure 6A and B), that is, the ratio of the lengths of the step [10̅ 1̅]Cc to the [101]Cc increased with increasing citrate concentration from 1.0 to 30.0 mM at pH 4.0−6.0. At pH values exceeding 7.0, the emergence of another new step direction shown by arrows in Figure 7A−C during brushite dissolution after a few minutes of exposure to

Figure 6. AFM deflection images showing evolution of the etch pit morphology of a brushite (010) surface after 2 min of dissolution in (A) 1.0 mM, (B) 5.0 mM, (C) 10.0 mM, or (D) 30.0 mM citrate within a range of pH 4.0−6.0. Arrows in A and B indicate a new step, mirror step [10̅ 1]̅ Cc to the [101]Cc, and insets in A−D depict schematically a continuous evolution of etch pit morphology, where the dashed triangle represents the case without additive and the red solid line reflects the case after citrate addition. Images A and B, 6 × 6 μm; C and D, 6.5 × 6.5 μm.

0.5−5.0 mM citrate solutions. Surprisingly, however, this new step disappeared and the triangular etch pit morphology recovered when the citrate concentration was further increased to 10.0 mM (Figure 7D). At a low concentration range (10− 100 μM citrate) and varying pH (4.0−8.0), no change of etch pit morphology was observed.

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DISCUSSION Dissolution Kinetics at Brushite−Water Interfaces. Figure 3B and C show that shallow etch pits with depths are all 0.75 ± 0.02 nm that corresponds to half the size of the lattice constant of the b axis, suggesting that the etch pit is formed by two-dimensional (2D) nucleation on dislocations or point defects;34,35 or by spontaneous 2D-nucleation. The respective rates of dissolution in [101]Cc, [101̅]Cc, and [1̅00]Cc directions in the presence of citrate at two different concentration ranges were measured (Figures 4 and 5). All three step movement rates may be changed, but the relative behavior remained constant allowing comparison of differences in step retreat rate between samples and experiments.36,37 In the present study, the primary concentration-dependent effect at the low concentration range (10−100 μM) at varying solution pH values was the steady slowing of both the [101̅]Cc and [1̅00]Cc steps, as the citrate concentration increased (Figure 4): the retreat velocity of the [101̅]Cc and [1̅00]Cc steps had decreased by a factor of approximately 1.6 in a 100 μM citrate solution (pH 6.9). The [101]Cc step velocity was completely inhibited (a step velocity of zero) in the presence of 10−100 μM citrate, suggesting that citrate binds strongly to the [101]Cc step. It seems likely that the negative carboxyl groups E



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on the degree of polymerization of such additives. However, in our system, it may be impossible to form polymerization of citrate. Across the continuum of driving force, mineral dissolution can be, in fact, understood through the same mechanistic theory of nucleation developed for mineral growth.43 Therefore, it seems more reasonable that citrate alters the water layer near the (010) surface or acts as a bridge between the solution and the crystal, effectively lowering the magnitude of the desolvation barrier, EK, by perturbations that displace water molecules or of the kinetic barrier, Eb, to removing atoms from the surface to initiate a pit.44 Moreover, the introduction of a high concentration of citrate could not be related to a higher driving force for dissolution because our experimental system is far away from equilibrium (i.e., the uncontrolled saturation in our experiments). At citrate concentrations far greater than 100 μM, there are a number of proposed mechanisms, including enhancing chelation and complexation by inner/outer-sphere adsorbed species etc.45 For example, Stumm and co-workers46,47 have demonstrated that LMW anions that bind in an inner-sphere, mononuclear, bidentate manner can dramatically enhance the dissolution kinetics of a variety of (oxyhydr)oxide minerals by polarizing bridging metal−oxygen bonds in the mineral substrate. In contrast, LMW anions that adsorb to form inner-sphere, binuclear, bidentate complexes tend to inhibit mineral dissolution.48 Similarly, outer-sphere adsorbed LMW anions can inhibit mineral dissolution through steric blocking of access to dissolution-active surface sites for dissolutionpromoting species such as protons.49 For citric acid used in our study, it is a tricarboxylic acid, with pKa1 = 3.13, pKa2 = 4.76, pKa3 = 6.40 at 298 K, that can form several species (H3L, H2L−, HL2−, and L3−) depending on solution pH.50 The major fraction of citrate species in the solution is neutral/ monoanionic, mono/dianionic, and trianionic in these three pH ranges of 6.0, respectively. As the solution pH increases, there is an obvious increase in the [101]Cc and [1̅00]Cc step retreat rates in the presence of higher concentration citrate (1.0−30.0 mM) (Figure 5A and B), suggesting that the dissociated citrate species is more likely to complex with calcium. This may imply that complexation with Ca then promotes the dissolution process. On the other hand, between pH 6.0 and 8.0, the dominant surface species of phosphate is HPO42−,11 whose presence inhibits complexation of citrate due to an increase of electrostatic repulsive forces between HPO42−/H2PO4− surface species of the mixed charge [101̅]Cc steps and the negatively charged citrate anions. This results in an observed decrease in the [101̅]Cc step retreat rates with increasing pH at a constant citrate concentration (Figure 5C). This observed trend is in good agreement with previous studies on citric acid adsorption by goethite.51 Net increase in the overall dissolution rates depends on retreat and deepening rates as well as the nucleation density of etch pits.52 In our experiments, no significant deepening of the brushite surface exposed to the experimental solution was observed. Therefore, differences in dissolution rates in the presence of different concentration citrate result from changes in etch pit density and/or retreat rates of all three steps. The interfacial dissolution process includes mass transfer, adsorption equilibria, and the kinetics of acid dissolution and surface complexation. All possible combinations of above processes at different citrate concentrations and pH values may explain the trends of the dissolution rates. In general, one or a combination

Figure 7. AFM deflection images showing evolution of the etch pit morphology of a brushite (010) surface after 5 min of dissolution in (A) 0.5 mM, (B) 1.0 mM, (C) 5.0 mM, and (D) 10.0 mM citrate at a constant pH 7.3. Arrows in A−C indicate a new step perpendicular to the [1̅00]Cc, and insets in A−C demonstrate a schematic drawing of the evolution of etch pit morphology with increasing citrate concentration (solid line), and inset in D shows its recovery to triangular shape (dashed line) as citrate concentrations exceeding 10.0 mM. Image A, 5.5 × 5.5 μm; B−D, 6.5 × 6.5 μm.

of citrate interact with the calcium-terminated [101]Cc polar step because the geometry and stereochemistry may be matched between citrate and this polar step.31 Moreover, citrate also appears to bind to the mixed charge steps based on the evidence that the velocity of the [1̅00]Cc and [101̅]Cc steps decreased when citrate was added (Figure 4A and B). All three steps maintain straight and well-defined, suggesting that adsorption rather than step pinning dominates during dissolution.26 When citrate concentration is greater than 100 μM, the dissolution kinetics under the present experimental conditions was significantly altered by citrate, causing brushite to rapidly dissolve along the [101̅]Cc and [1̅00]Cc directions (Figure 5A). In addition to increasing the step retreat velocity, increasing citrate concentration from 0.1 to 0.5 mM resulted in an obvious increase in etch pit density (Figure 4C and D). Values of etch pit density higher than defect density have been related to nucleation of pits on defect-free surfaces under far from equilibrium conditions (as is the case for our solutions). Similar to the role of background electrolytes,38 citrate may disrupt the hydration environment of the brushite surface and hence to reduce the energy barrier needed for unassisted etch pit nucleation. Whatever the mechanism, the energy barrier for etch pit formation must be reduced to result in a higher etch pit density (Figure 4D). This sudden change from inhibition to promotion in step kinetics by an additive has been observed in a number of crystal growth systems.39−42 In the case of the promotion/inhibition of crystal growth it has been often observed that the effect of the additives depends on the chemical speciation in the aqueous solution and, particularly, F



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mentioned before, in the presence of small organic anions, the mechanism underlying the brushite dissolution can be the same as in the case of a simple inorganic electrolyte (Note the relatively high etch pit density attained at 0.5 mM citrate in Figure 4D). Dissociated citrate can then promote dissolution of brushite as a result of the stabilizing effect of its charge on water of solvation of brushite crystal building units (Ca2+ and HPO42−).62 We suggest that water molecules actually mobilize ions from the brushite crystal structure and then dissolved ions can be sequestered by ligand species in solutions. At higher concentration of citrate (1.0−30.0 mM), the chelating capacity of the ligand-containing solution toward dissolved calcium ions increases. Evolution of Etch Pit Morphology on Brushite (010) Faces in the Presence of Citrate. The characteristic triangular shape changed to a four-sided trapezium in acidic conditions, and shape recovered to the normal in alkaline conditions with increasing citrate concentration. In acidic conditions, citrate causes the step [1̅01̅]Cc to appear, changing the normal triangular step pattern into a four-sided trapezium (Figure 6A−D). It should be noted that the second polar step ([1̅01̅]Cc) likely always has the same dissolution kinetics as the observed polar step because the two steps have the same surface chemistry.30 But the second polar step is not normally expressed owing to the geometry of the two bounding steps ([101̅]Cc and [1̅00]Cc). It is uncertain whether the pit shape change by citrate is due to thermodynamics (lowering step energies) or kinetics (surface poison or impurity inhibition), or both. This is largely because the system is far away from equilibrium (i.e., the uncontrolled saturation in our experiments). We propose that the observed changes in pit morphology may result from a selective interaction/binding of citrate along certain specific directions based on the well-defined anisotropy in the etch pits (Figure 6). Three mechanisms may exist for citrate and brushite steps at a microscopic level, including electrostatic interaction, geometric matching, and stereochemical conformity based on previous results.63,67 Both the [101]Cc, a calcium-terminated polar step, and its mirror step [1̅01̅]Cc satisfy the requirement that a strong electrostatic interaction with a chelating ligand is favored at steps where Ca atoms are contiguous, this will slow the step motion (no observable rates for the [101]Cc steps at lower citrate concentrations, Figures 4 and 5) either by enhancing the step stability or by inhibiting the dissolution kinetics. This specificity of the citrate interaction with [101]Cc steps may be due to its tridentate structure containing three carboxylate groups, serving as multiple binding sites that allow for specific molecular recognition63,64 or stereochemical interacting with crystal surfaces.65−67 Moreover, the step geometry may allow multiple accommodations of carboxylate groups with the least amount of distortion.65 In general, the interaction of carboxyl groups of citrate with brushite surfaces may have three distinct effects: alteration of step free energies; specific interactions that stabilize facets; and the chelation of calcium. Similar etch pit and two-dimensional island or spiral morphological modifications are frequently observed in calcite dissolution and growth system in the presence of different inorganic and organic species.34,52,68,69 Morphological changes in growth and dissolution features can also be understood in terms of the changes in the hydration of the mineral surface induced by the background ions at certain concentrations.70 In terms of this recognition, the evolution in etch pit morphology

of the following reactions dominates brushite dissolution in citrate solutions. At low pH (OPO3H20 or OPO3H− (the notation “>” means surface sites) at pH < 5,57 whereas at pH > 7, rates are controlled by the hydrolysis of calcium centers >CaOH2+.54,57 Therefore, overall brushite dissolution rates in the presence of an organic ligand should be given by eq 4 when considering the CaL1−m surface complex concentration at pH > 7,58 that is, r+ = [k Ca{>CaOH+2 } + kL{>CaL1 − m}]

(4)

where kCa and kL are rate constants. Actually, organic ligands may affect rates of reaction by affecting the stability and concentration of any intermediate species.59 Specifically, the effect of citrate on step velocity could depend on whether the citrate complexes calcium ions comprising the step themselves or an intermediate, inner-sphere adsorbed species on top of the step. Whether citrate inhibits or facilitates dissolution maybe affected by number of bonds citrate is able to make to a given species, for example, through steric hindrances to binding, and thus its capacity to liberate the ion from the mineral surface. In the presence of salt, NaCl (0.5 M), velocities of etch pit retreat (Supporting Information Figure S1) in 1.0 mM citrate solutions were significantly higher than in the absence of salt. The salt effect may arise through a transition from dissolution controlled by the population of dislocation defects to 2D nucleation of vacancy islands, possibly leading to the enhancement of step retreat rates.43 At high ionic strengths though free Cl−(aq) is limited due to ion pair formation, the activity is reduced relative to lower ionic strengths because of the change in the activity coefficient. In addition, specifically for anions of background electrolytes, Cl− ions determine the solvent structure around Ca2+ and calcium removal from the surface structure is the rate limiting step for dissolution.60,61 As G



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observed in this study, that is, from trapezium to triangular shape, may be explained as the result of progressive stabilization of polar faces/step edges (those with a net dipole moment) by citrate at a certain concentration range (0.5−5.0 mM) (Figure 7). Collectively, we have presented in situ AFM observations on the dissolution of brushite by citrate, and the results indicate that citrate has a bimodal effect (promoter/inhibitor) on the dissolution of brushite (010) surfaces, which mainly depends on the concentration. We have systematically shown a more complete view about how an organic ligand influence the dissolution kinetics, as different effects dominate in different concentrations of organic acids and solution pH values. Our results provide a detailed analysis of surface dissolution of a Ca−P mineral in the presence of an organic acid over a significantly broader range of concentrations at various pH values and ionic strengths, which would aid in understanding phosphate mineral dissolution in soils and the bioavailability of P. The obtained results exhibit possible implications for analyzing the controversial role of organic acids in a more complex soil−rhizosphere environment since the studied compositions are relevant to rhizospheric solutions, and mimic natural environmental conditions. Because of differences in the solution conditions used in our AFM system and the soil environment, it is difficult to directly relate our results to actual rhizosphere interactions. However, concerning the pH and organic concentrations relevant to soils, as well as how the microscopic work with single crystals relates to bulk processes in soils, this study has taken an important initial step toward a further understanding of an isolated factor, namely how an organic ligand promotes or inhibits Ca−P dissolution at the microscopic level. This study may provide insights for furthering the understanding of the release of P by mineral dissolution in more complex soil systems.

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ASSOCIATED CONTENT

S Supporting Information *

Results of brushite dissolution in the absence and presence citrate at 0.1 mM or 1.0 mM with different levels of NaCl (Figure S1). This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*(L.J.W.) Phone/Fax: +86-27-87288095; e-mail: ljwang@mail. hzau.edu.cn. 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 (Grant No. 41071208) and a startup grant (2010BQ063) and two Independent Innovation Foundation (2012MBDX014, 2011JQ008) from the Huazhong Agricultural University to Lijun Wang. It is also supported in part by the Fundamental Research Funds for the Central Universities (2011PY150). Research is sponsored by the Chemical Sciences, Geosciences, and Biosciences, Office of Basic Energy Sciences, U.S. Department of Energy to AGS. H



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