Density Functional Theory Calculations on the Complexation of p

Article pubs.acs.org/JPCA

Density Functional Theory Calculations on the Complexation of p‑Arsanilic Acid with Hydrated Iron Oxide Clusters: Structures, Reaction Energies, and Transition States Adrian Adamescu,† Ian P. Hamilton,‡ and Hind A. Al-Abadleh‡,* †

Department of Chemistry, University of Waterloo, Waterloo, ON Canada N2L 3G1 Department of Chemistry and Biochemistry, Wilfrid Laurier University, Waterloo, ON Canada N2L 3C5

‡

S Supporting Information *

ABSTRACT: Aromatic organoarsenicals, such as p-arsanilic acid (pAsA), are still used today as feed additives in the poultry and swine industries in developing countries. Through the application of contaminated litter as a fertilizer, these compounds enter the environment and interact with reactive soil components such as iron and aluminum oxides. Little is known about these surface interactions at the molecular level. We report density functional theory (DFT) calculations on the energies, optimal geometries, and vibrational frequencies for hydrated pAsA/ iron oxide complexes, as well as changes in Gibbs free energy, enthalpy, and entropy for various types of ligand exchange reactions leading to both inner- and outer-sphere complexes. Similar calculations using arsenate are also shown for comparison, along with activation barriers and transition state geometries between inner-sphere complexes. Minimum energy calculations show that the formation of inner- and outer-sphere pAsA/iron oxide complexes is thermodynamically favorable, with the monodentate mononuclear complexes being the most favorable. Interatomic As−Fe distances are calculated to be between 3.3 and 3.5 Å for inner-sphere complexes and between 5.2 and 5.6 Å for outer-sphere complexes. In addition, transition state calculations show that activation energies greater than 23 kJ/mol are required to form the bidentate binuclear pAsA/iron oxide complexes, and that formation of arsenate bidentate binuclear complexes is thermodynamically -rather than kinetically- driven. Desorption thermodynamics using phosphate ions show that reactions are most favorable using HPO42− species. The significance of our results for the overall surface complexation mechanism of pAsA and arsenate is discussed.

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INTRODUCTION In developing countries like China with intense poultry operations,1,2 and prior to their ban in Europe3 and voluntary suspension of sale in North America,4,5 roxarsone (4-hydroxy3-nitrobenzenearsenic acid, ROX) and p-arsanilic acid (4aminophenyl arsonic acid, pAsA) are, or have been, used as feed additives to control disease, stimulate growth, and improve both feed efficiency and feed conversion.6−8 The majority of these compounds are not metabolized in the poultry and are excreted chemically unchanged, and the manure is turned into fertilizer pellets for commercial use.9 After application of the fertilizer to soil, microbial activity and ultraviolet light10−14 lead to the formation of other organic arsenic species such as dimethylarsinic acid (DMA) and inorganic arsenic (iAs).11,15,16 This in turn poses a number of health and environmental concerns since arsenic, in its various forms, is a known carcinogen17−21 and has been correlated with hypertension as well as other cardiometabolic diseases.22 It is well established that interaction of pollutants with soil particles controls both their transport and bioavailability.23 Hence the fate of arsenic contaminants, whether introduced to the environment naturally or anthropogenically, depends upon their interaction with reactive soil components and their competition with organic matter and phosphate for sites on soil © 2014 American Chemical Society

particles. Since these interactions occur at the water/solid interface, it is necessary to use surface-sensitive techniques to quantify the thermodynamics and kinetics of binding, as well as obtaining structural data on the adsorbed surface complexes. Commonly used techniques, which have proven to be effective when studying surface complexes under environmentally relevant conditions, are attenuated total internal reflectance Fourier transform infrared (ATR-FTIR)24 spectroscopy and extended X-ray absorption fine structure (EXAFS) spectroscopy.25 For spectroscopic results to be helpful in building surface complexation models that can predict binding thermodynamics, the experimental results need to be accompanied by theoretical calculations for surface processes of environmental relevance. The application of computational chemistry to studies in geochemistry is increasingly becoming invaluable in providing molecular-level details about interfacial processes that either aid in the interpretation of experimental data or provide information that is difficult to obtain experimentally.26 Examples of systems investigated include the interaction of oxyanions and organic molecules with iron and aluminum Received: May 13, 2014 Revised: July 8, 2014 Published: July 9, 2014 5667

dx.doi.org/10.1021/jp504710b | J. Phys. Chem. A 2014, 118, 5667−5679

The Journal of Physical Chemistry A

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oxyhydroxides.27−31 Hybrid molecular orbital/density functional theory (MO/DFT) proved to be adequate for calculating electronic energies, optimized structural parameters, vibrational frequencies, and binding thermodynamics.29,32 Solvent effects were accounted for in these calculations both implicitly through the use of a solvation model such as the integral equation formalism polarizable continuum model (IEFPCM)33 and explicitly by adding water molecules around the clusters.29 Conclusions about the nature of surface complexes, whether monodentate mononuclear (MM) or bidentate binuclear (BB), and their degree of protonation can only be drawn with confidence after correlating experimental vibrational frequencies with calculated ones. Because there are fewer theoretical studies on the surface chemistry of organic oxyanions in general, and organoarsenicals in particular, we previously reported DFT calculations of geometries and vibrational frequencies for hydrated arsenate and organoarsenicals to show the effect of protonation and increasing organic substitution on the stretching frequency of As−O bonds [v(As−O)].34 These calculations were helpful in explaining trends in v(As−O) observed in ATR-FTIR spectra of these compounds as a function of pH.35 We also reported DFT calculations on the thermodynamic favorability for the formation of hydrated DMA-iron oxide clusters in the outersphere (OS), MM, and BB configurations.36,37 After correlating the calculated v(As−O) with experimental frequencies from equilibrium studies, results indicated simultaneous formation of both inner-sphere, MM and BB, as well as OS complexes for DMA on hematite and goethite particles. We recently reviewed field and bulk studies that examine the transformation of arsenic in soils treated with contaminated poultry litter,38−42 which clearly showed that mechanistic details of the adsorption of organoarsenicals and their degradation products are needed to better understand their environmental fate. These studies also highlighted the need for surface-sensitive measurements on the interaction of aromatic arsenical compounds relevant to those released from poultry litter with reactive components in soil. Structural characterization and thermodynamics of pAsA adsorption and desorption to/from hematite and goethite were investigated in our lab using ATR-FTIR. Results from these studies indicated that pAsA forms at least two types of inner-sphere complexes that are most likely MM over a wide pH range, that desorption of pAsA is more efficient by phosphate in solution than chloride, and that the binding constant of pAsA is higher than that of DMA.43,44 We recently reported in situ kinetic studies for the interaction of pAsA with hematite nanoparticles under neutral conditions using ATR-FTIR.45 The initial adsorption rate constant for pAsA was lower by a factor of 1.6 than that of iAs(V), suggesting an average behavior for the formation of quantitatively more weakly bonded MM and/or hydrogen-bonded complexes for the former relative to strongly bonded BB surface complexes for the latter under our experimental conditions. These conclusions were confirmed from trends observed in the adsorption kinetics of phosphate on hematite surfaces with either adsorbed pAsA or iAs(V). In this paper, we investigate, computationally, the formation of neutral and charged inner- and outer-sphere pAsA/iron oxide clusters, their minimum-energy geometries, vibrational frequencies and the thermodynamic favorability of their formation. Frequency calculations are used for correlation with the spectral features observed experimentally, while thermodynamic state functions of adsorption are calculated

for various hypothetical ligand exchange reactions to gain insight into the types of reactions that are most thermodynamically favorable. Intercomplex reactions are also analyzed using multiple single-point energy calculations along the reaction pathway to quantify relative energy barriers between MM and BB complexes of pAsA. Similar calculations were performed on iAs(V)−iron oxide complexes for comparison in an effort to explain the dominant formation of BB complexes. In addition, various hypothetical desorption reactions due to phosphate are considered for which the Gibbs free energies of desorption were calculated.

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COMPUTATIONAL METHODS Geometry Optimization. Calculations were performed using the Gaussian 09 program running on Sharcnet,46 typically on 8 or 16 processors. All structures were energy minimized, without any symmetry or geometry constraints, using DFT with the B3LYP functional47,48 and the 6-311+G(d,p) basis set. Frequencies were computed and used to correlate theoretical v(As−O) to experimental values. We confirmed that all transition state structures had one imaginary frequency and that the corresponding vibrational mode shows movement that goes from reactant to product. The calculations were performed on clusters with a net charge of −2, −1, 0, or +1, at 298.15 K (25 °C) and 1 atmospheric pressure. Solvation was simulated by explicit inclusion of water molecules around each structure and also by using the IEFPCM.33 Clusters containing Fe3+ atoms were calculated at high spin. It was observed that for small iron oxide clusters (containing 1 to 5 Fe atoms) the lowest electronic energy was obtained when considering five unpaired electrons for each Fe atom.49 Since two Fe atoms are used to represent the iron oxide surface in this paper, calculations are performed with 10 unpaired electrons (hence a multiplicity of 11) for each of the complexes. To minimize computational cost, all calculations were initially performed with a smaller 6-31G(d) basis set with explicit water molecules added one or two at a time. Once the fully hydrated complexes had been optimized at this level of theory, the calculations were run again at the B3LYP/6-311+G(d,p) level of theory. B3LYP/ 6-311+G(d,p) with IEFPCM was selected as the method of choice for these calculations because it yielded the best results compared to experiment when measured against three other methods for similar type clusters.49 The M062X functional with the same basis set and solvation model was also used to calculate the structures needed for a MM pAsA reaction to compare with the B3LYP method since this functional is parametrized for use with transition metals.50 The MM complexes proved to be more challenging to compute because of their rotational freedom and the tendency of the AsOx moiety to rotate away from the iron core. To deal with this we fixed the starting geometry of the MM with the pAsA at the center of the Fe-oxide and added a few explicit waters to form hydrogen bonds (FeOH2-−OH2--OAs) to keep the pAsA from rotating away. This was in contrast with the other geometries for OS and BB of pAsA and arsenate, where the starting geometry was optimized first with no explicit waters and then waters were added one or two at a time. Transition States. To calculate activation barriers for the intercomplex transitions from MM to BB, multiple constraint distance optimizations were performed along the reaction pathway. This involved decreasing the distance between the unbound oxygen of the AsO moiety and the Fe center atom by 0.1 Å and optimizing the complex at each step (note that while 5668



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the AsO-Fe distance was held fixed, the rest of the cluster was allowed to relax). A similar procedure was performed by Watts et al. for arsenate complexes.51 This was repeated iteratively until a second covalent bond (i.e., BB complex) was formed and a local minimum was reached. This is illustrated in Figure 1 for pAsA/iron oxide.

each negatively charged O atom hydrogen bonded to either 2 or 3 explicit waters. Structures c, f, and i in Figure 2 have 12 explicit waters with 4 waters added around the NH2 group. Table 1 shows that increasing the number of explicit water molecules from 4 to 8 to 12 only alters the bond distances by 0.01−0.02 Å. The indented pAsA in Table 1 shows that choosing a different functional, namely M06-2X with the same basis set and solvation model, does not affect the bond distances by more than 0.02 Å. Thus, for remainder of the paper, only the negatively charged species (pAsA−) is considered since it is the dominant species in a pH range of 4.1(pKa2) to 9.2 (pKa3) where most environmental processes occur. The point of zero charge for many iron (oxyhydr)oxides (such as goethite) is around 953 and thus pAsA− would most likely be attracted by the positively charged iron oxide surface to form a neutral pAsA/iron oxide complex. Also, data in Table 1 shows that protonation of pAsA affects d(As−C), which increases as deprotonation increases. This observation could be explained in light of our earlier calculation34 of atomic charges as a function of organic substitution and deprotonation. The formation of As−O− groups increases electronic density around the O atoms relative to As−OH and that causes the As−C bonds to weaken slightly. The effect of rotating the aromatic ring of pAsA− and the resulting electronic energy for various levels of solvation was also explored. Four constraint energy minimization calculations were performed for pAsA− in the gas phase, with the IEFPCM solvation model, as well as with IEFPCM and both 4 and 8 explicit water molecules. The energies of the pAsA structures were computed after rotating the C−C−As−O dihedral angle every 10° from −180 to 180° while the rest of the bonds and angles were allowed to relax. In Figure 3(a) the relative energy of the pAsA− molecule is shown as a function of the dihedral angle and it may be seen that the relative energy minima for all four solvation cases were at about −164° and 20° for the C− C−As−O dihedral angle. These minimum relative energies also correspond with the position of the −OH group on pAsA− being approximately perpendicular to the aromatic ring as shown in Figures 3(b). The structure of that minimum energy pAsA− is shown in Figure 4(a). The relative energy maxima were at about −69° and 111° for the C−C−As−O dihedral angle for the two cases with no explicit water molecules, and −77° and 103° for the cases with 4 and 8 explicit water molecules. These angles correspond with the position of the −OH group on pAsA− being in the same plane as the aromatic ring as seen in Figure 4(b). One reason for this might be that there is hyperconjugation between the As−O σ bond when lined up with the π bonds of the aromatic ring. Hyperconjugation is known to stabilize molecules and that affects the length of the sigma bonds with carbon leading to shorter bond distances.54 Figure 3(c) shows the variation of the As−C bond distances with the dihedral angles. Near the angle of the minimum relative energy structure of pAsA− shown in Figure 4(a), gas phase calculations yielded d(As−C) = 1.967 Å, which increased to 1.972 Å near the angles of the maximum relative energy structure shown in Figure 4(a). Figure 3(c) also shows that using the IEFPCM solvation model with no explicit waters, a slight change in the As−C bond by about 1.4% is observed in our calculations. Adding explicit waters further decreased d(As−C), and values approached the experimentally determined distance reported for the crystalline structure of pAsA−.55

Figure 1. Reaction pathway for the transition of a pAsA MM complex to a BB complex. The AsO-Fe distance was decreased iteratively by 0.1 Å as the rest of the cluster was allowed to relax. The BB complex is higher in energy than the MM (but lower than the OS complex) and there is an activation barrier of 23 kJ/mol going from MM to BB.

The geometry with the highest energy along the reaction pathway was used as the starting point for the transition state calculation [using the command OPT=(TS, CalcFC)] to find the transition state structure, complex G, which exhibited one imaginary frequency. The corresponding vibrational mode shows movement that goes from MM to BB. As discussed in the next section, the explicit inclusion of water molecules often plays a critical role in lowering the energy of the transition state.

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RESULTS AND DISCUSSION Geometries of Hydrated pAsA Clusters and pAsA/Iron Oxide Complexes. The optimized geometries of hydrated pAsA in simulated bulk acidic, neutral, and basic conditions are shown in Figure 2. These clusters were considered given the pKa values of pAsA(aq), 1.9, 4.1, and 9.2 corresponding to the deprotonation of the −NH3+ and the two As−OH groups, respectively.52 Table 1 shows the bond distances of As−O and As−C [d(As−O) and d(As−C)], for the fully protonated, the singly protonated, and the fully deprotonated pAsA clusters with four explicit water molecules surrounding each molecule. In the case of the fully protonated pAsA, the As−O bond ranges from 1.65−1.67 Å and d(As−OH) ranges from 1.76− 1.79 Å depending on the number of explicit water molecules used in the simulation. In the case of the singly protonated species, the electrons are delocalized over two As−O bonds to form a slightly longer As−O bond distance ranging from 1.68− 1.70 Å, and d(As−OH) is 1.81 Å in all cases where explicit water molecules were added. In the case of fully deprotonated pAsA, the electrons are delocalized over all three As−O bonds resulting in slightly longer As−O bonds ranging from 1.71− 1.72 Å. When increasing the number of explicit waters from 4 to 8, we also see waters starting to populate the second solvation shell. Structures b, e and h in Figure 2, showing pAsA, pAsA−, and pAsA2− with 8 explicit waters, all have 6 waters in the first solvation shell and 2 waters in the second solvation shell, with 5669



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Figure 2. Optimized structures of pAsA species surrounded with explicit water molecules. Structures a, b, and c are dominant species in acidic conditions (pH < 4.1), structures d, e, and f are dominant in neutral conditions (4.1 < pH < 9.2), and structures g, h, and i are the dominant species in basic conditions (pH > 9.2) based on the pKa values of pAsA. Calculated bond distances are listed in Table 1.

The adsorption of pAsA− onto iron (oxyhydr)oxides can produce different types of complexes via ligand exchange. Figure 5 shows minimum energy geometries of hydrated pAsAiron oxide clusters in the OS (A), MM (B), and BB (C) configurations. In OS complex A, pAsA− is adsorbed via hydrogen bonds and the interatomic As−Fe distance ranges

from 5.21 to 5.58 Å between the two iron centers. MM complex B is formed when ligand exchange takes place and one covalent bond formed between the oxygen of the arsenical and the Fe atom of the iron oxide surface. In addition, BB complex C has two covalent bonds between two oxygen atoms of the arsenical and two Fe atoms of the iron oxide surface. Structural 5670



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Table 1. Calculated As−O and As−C Bond Distances (Å) for pAsA Clusters with Various Explicit Water Moleculesa Geometries

As−O1

As−O2

As−O3

As−C

pAsA·(H2O)4 pAsA·(H2O)8 pAsA·(H2O)12 pAsA−·(H2O)4 pAsA−·(H2O)4b pAsA−·(H2O)8 pAsA−·(H2O)12 pAsA2−·(H2O)4 pAsA2−·(H2O)8 pAsA2−·(H2O)12

1.65 1.66 1.67 1.69 1.68 1.68 1.68 1.71 1.71 1.71

1.77 1.76 1.76 1.69 1.67 1.69 1.70 1.71 1.72 1.72

1.79 1.78 1.79 1.81 1.79 1.81 1.81 1.71 1.72 1.72

1.90 1.89 1.88 1.92 1.90 1.91 1.91 1.95 1.94 1.94

a

Structures are shown in Figure 1, which were calculated using B3LYP/6-311+G(d,p) with the IEFPCM solvation model. bCalculated with M06-2X/6-311+G(d,p) and IEFPCM for comparison.

Figure 4. (a) Energy minimum for pAsA− is when the dihedral angle C−C−As−O ∼ 20°. This dihedral angle also corresponds with the position of the OH group on pAsA− being perpendicular to the aromatic ring as shown. (b) The energy maximum for pAsA− is when the dihedral angle C−C−As−O ∼ 117°. This also corresponds with the position of the OH group on pAsA− being parallel to the aromatic ring.

transitioning between the MM and BB configurations. As a result of complexation with the iron oxide, the As−O bonds for OS complex A are 1.68 and 1.69 Å, and the As−OH bond is 1.80 Å. Similarly, the AsO bond for complex B (MM) is 1.66 Å, showing more double bond character while As−OH is 1.82 Å. The As−O(Fe) bond is 1.70 Å. For complex C (BB) the As−O(Fe) bonds are 1.68 and 1.69 Å and the As−OH bond is 1.80 Å. The interatomic As−Fe distances for the OS complex A are 5.21 and 5.58 Å, while for the inner-sphere complexes they are 3.54 and 4.92 Å for complex B, and 3.29 and 3.43 Å for complex C. Moreover, Table 2 shows the geometries of the energy minimized MM arsenate structures D and E, as well as BB structure F for comparison. There are two possible MM

Figure 3. Relative energy of pAsA− as a function of (a) dihedral angle (C−C−As−O) and (b) OH angle relative to the aromatic ring, and (c) the As−C bond distance as a function of the dihedral angle (C− C−As−O), at various levels of solvation using B3LYP/6-311+G(d,p) method.

parameters of these energy-minimized clusters are listed in Table 2, as well as those of the transition state complex G, 5671



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Figure 5. Energy-minimized structures of hydrated pAsA/iron oxide clusters (A−C), arsenate/iron oxide clusters (D−F), transition state (TS) complexes for the conversion of the MM to BB structures of pAsA (G) and arsenate (H), and hydrated PO4/iron oxide clusters (J−M) calculated using B3LYP/6-311+G(d,p) and the IEFPCM solvation model. Calculated bond distances are listed in Table 2.

seems to be slightly pulling the molecule away from iron centers relative to iAs(V). To date, there are no reports on d(As−Fe) for pAsA adsorbed on iron (oxyhydr)oxides from X-ray absorption studies. However, d(As-M), where M = Fe, Al, Ti were reported for surface complexes of monomethyarsonic acid (MMA), which is a monosubstituted arsenical like pAsA. Shimizu et al.60 reported structural parameters of MMA adsorbed on goethite from fits to EXAFS data, where a coordination number (CN) of 1.8 ± 1.1 yielded the best fit with d(As−Fe) = 3.31 ± 0.03 Å.60 For a fixed CN of 3, the average value of the As−O bond lengths is reported to be 1.70 ± 0.008 Å. Also, the same group studied MMA coordination to aluminum oxide and reported best fits with CN = 2.1 ± 0.94 and d(As−Al) =3.16 ± 0.022 Å.61 Similarly, for a fixed CN of 3, the average value of the As−O bond lengths is reported to be 1.69 ± 0.007 Å.61 Moreover, Jing et al.62 reported structural data for MMA surface complexes on TiO2 nanoparticles (CN= 1.9 ± 0.4 and d(As−Ti) =3.32 ± 0.01 Å, and CN= 3.0 ± 0.1 and As−O =1.69 ± 0.01 Å). The above experiments were conducted at pH 5 where the singly protonated form of MMA, CH3AsO2H−, is the most dominant species based on its pKa1 (3.6).63 These experimental results indicate that the formation of BB complexes is similar to that of arsenate complexes, with reported As-M interatomic distances in the range 3.23−3.37

configurations as shown in Figure 5 for complexes D and E for which the interatomic As−Fe distances are 4.32, 3.41 and 3.97, 3.50 Å, respectively. Complex E is an intermediate complex between the MM complex D and BB complex F. Watts et al.51 calculated BB complexes for arsenate on Fe(III) (oxyhydr)oxide. Our equivalent structure, complex F from Table 2, shows a similar geometry to theirs with a d(As−Fe) of 3.20 and 3.28 Å, d(As−O) between 1.65−1.72 Å and a d(As−OH) of 1.79 Å. Table 2 also lists comparisons with selected experimental distances using X-ray absorption spectroscopy for arsenate surface complexes on hematite and goethite. In addition, we previously36 calculated structural parameters for iAs(V)-iron oxide BB complexes, neutral and positively charged, which yielded As−O distances of 1.66 Å (uncomplexed), As−OH distances between 1.74 and 1.81 Å and As−O(Fe) distances between 1.68 and 1.72 Å. The slight differences in these distances could also be due to small differences in the Hbonding occurring among the explicit water molecules, the Fe cluster, and the As species in our models. Also, the relatively small differences in the calculated bond distances between pAsA and iAs(V) complexes can be explained by the presence of the aromatic group, which promotes resonance between the two uncomplexed As−O bonds in the case of MM complexes. Also, the aromatic group 5672



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Table 2. Predicted As−O, Fe−O Bond Distances (Å) and As−Fe Inter-Atomic Distances (Å) of pAsA/Iron Oxide Complexes A, B, C, Arsenate/Iron Oxide Complexes D, E, F and Transition State Complexes G, H shown in Figure 5 as Calculated Using B3LYP/6-311+G(d,p) and the IEFPCM Solvation Model OS, MM and BB pAsA complexes A,B,C

As−Fe1

A. pAsA · (H2O)4 Fe2(OH)4(OH2)4 5.21 B. pAsA-Fe2(OH)5(OH2)3 · (H2O)5 4.92 B. pAsA-Fe2(OH)5(OH2)3 · (H2O)5a 4.66 C. pAsA-Fe2(OH)5(OH2)3 · (H2O)5 3.43 MM arsenate complexes D and E and BB complex F D. HAsO4−Fe2(OH)4(OH2)5 · (H2O)3 E. HAsO4−Fe2(OH)4(OH2)4 · (H2O)4 F. HAsO4−Fe2(OH)4(OH2)4 · (H2O)4 Transition state for pAsA G. pAsA-Fe2(OH)5(OH2)3 · (H2O)5 Transition state for arsenate

As−Fe1

As−Fe2 5.58 3.54 3.48 3.29 As−Fe1

arsenate arsenate arsenate arsenate arsenate

on on on on on

hematite, BB (ref 56) goethite, BB (ref 56) goethite, BB (ref 57) goethite, BB (ref 58) goethite, BB (ref 59)

1.68 1.66 1.65 1.68

1.69 1.70 1.69 1.69

3.41 3.50 3.40 As−O1

3.34 As−Fe2

H. HAsO4−Fe2(OH)4(OH2)4 · (H2O)4 3.71 3.40 Comparison with computational results (ref 51) As−Fe Fe2(OH)4(OH2)4HAsO4 (BB) Fe2(OH)4(OH2)4HAsO4·4H2O (BB) Fe2(OH)4(OH2)4HAsO4·8H2O (BB) [Fe2(OH)2(OH2)6H2AsO4]3+ (BB) [Fe2(OH)2(OH2)6H2AsO4]3+ (BB) Comparison with XAS datab

As−O2

As−Fe2

4.32 3.97 3.27 As−Fe2

3.94 As−Fe1

As−O1

As−O1

As−O2

1.68 1.68 1.70

1.71 1.70 1.70 As−O2

1.66 As−O1 1.68 3.25 3.28 3.30 3.24 3.29

1.71 1.69 1.70 1.70 1.71

As−Fe

As−Fe

As−OFe

3.24 3.30

3.35 3.30

1.70 1.70

3.30 3.28 ± 0.01 3.30 (±0.008)

1.69 1.69 1.67 As−C 1.91 As−O3

1.72 As-OFe

Fe1−O1

As−OH

1.91 1.90 1.88 1.90 As−O3

1.71 As−O2

As−Fe

3.13 3.20 3.30 3.24 3.29

As−C

1.67 As-OFe

1.80 1.82 1.79 1.80 As−OH

1.98 Fe1−O1

1.98 1.98 2.08 Fe2−O2

1.81 As−OH

Fe1−O1

2.08 Fe2−O2

As−OH

2.05 AsO

1.83 1.76 1.76 1.72 1.73 As−OH

1.70 1.70

2.19 Fe1−O1

2.06 2.03 2.11 Fe2−O2

1.80 1.82 1.79 As−OH

1.82 As−OH

1.73 1.72 1.7 1.70 1.71 As−OFe

Fe2−O2

1.63 1.67 1.67 1.72 1.73 As−OH

AsO

1.70 1.70

1.62 1.63

1.69 1.69 (±0.004)

a

Calculated with M06-2X/6-311+G(d,p) and IEFPCM for comparison. bXAS = X-ray absorption spectroscopy. Selected references. For data on other iron (oxyhydr)oxides, see Table 4 in ref 51.

Table 3. Calculated v(As−O) Frequencies (cm−1) for pAsA/Iron Oxide Complexesa v(As−O)

OS complex A A. pAsA · (H2O)4 Fe2(OH)4(OH2)4

823(s)b,c,d, 828(s)b,c,d, 840(s)b,d, b,c,d b,d b,c,d b,d 859

MM complex B

a

, 861 , 868 v(As-OFe)

v(As−OH) 853(s)

,

629(s)

b,c,d

, 635b,d, 644(s)c

, 909

B. pAsA-Fe2(OH)5(OH2)3 · (H2O)5 BB complex C

778b,c,d, 789b,c,d, 830(s)b,c,d, 840b,c,d, 861c, 865c, 918b,d v(As-OFe)

C. pAsA-Fe2(OH)5(OH2)3 · (H2O)5

832(s)b,c,d, 845(s)b,c,d, 862b,c,d, 867c, 903b,d, 930b,d b

b,c,d

v(As−O)

v(As−OH)

892(s)b,c,d

603(s)b,c,d v(As−OH) 631(s)b,c,d 648(s)b,c,d

c

Frequencies adjusted for anharmonic behavior using F = 1.03. Coupled with waters rocking/wagging. Coupled with C−C−C bending in the aromatic ring. dCoupled with OH bending. (s) = strong.

Å.56,60,64−66 This comparison shows that reported As-M interatomic distances in MMA BB complexes are on the higher end of the range reported for arsenate. In light of the calculated As-M interatomic distances for MM pAsA being higher than those for BB, and given the uncertainly in CN from EXAFS fits, the existence of MM complexes cannot be ruled out in samples prepared for EXAFS measurements. Also, because distances between As and Fe greater than 5 Å are usually undetectable using EXAFS, conclusions about OS complexes cannot be inferred from this data. In situ resonant surface scattering utilized by Catalano et al.65,67 for studying arsenate adsorption on hematite and corundum single crystal surfaces can measure OS arsenate complexes with As-metal distances >5 Å. Hence, studying surface complexes of organoarsenicals using these techniques will provide invaluable insight into their geometries. Calculated v(As−O) of Hydrated pAsA/Iron Oxide Complexes. In order to aid in the interpretation and assignment of spectral features observed in earlier ATR-FTIR

measurements for pAsA complexation to iron (oxyhdyr)oxides,43,44 Table 3 shows the calculated As−O stretching frequencies, v(As−O), after adjustment for anharmonicity by a scaling factor of 1.03, which is a rounded up figure as suggested by Irikura et al.68 Although it is common practice to use scaling factors from tables calculated by others in the gas phase, we have used our own experimental and theoretical data to create scaling factors that are precise for the molecules and the level of theory we use. Specifically, when using iron and arsenic with B3LYP/6-311+G(d,p) IEFPCM, it was shown that a factor of 1.03 will give the best correlation with experimental frequencies.49 When the spectrum of each complex was analyzed, medium and strong vibrations with a bond stretch greater than 0.2 Å were considered meaningful and labeled with (s) (Table 3). The weak v(As−O) vibrations below this limit are also listed. Experimental v(As−O) results obtained from pAsA adsorption on goethite (FeOOH) show spectral features at 769, 789, 837, and 870 cm−1, and those on hematite (Fe2O3) 5673



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are at 795, 837, and 877 cm−1.43 Figure 6 shows the correlation plots between these experimental v(As−O) values and the

° = ΔHads

∑ (E0 + Hcorr)products − ∑ (E0 + Hcorr)reactants (1)

° = ΔHads

∑ (E0 + Gcorr)products − ∑ (E0 + Gcorr)reactants (2)

° = (ΔHads ° − ΔGads ° )/T ΔSads

(3)

Table 4 shows the calculated ΔH°ads, ΔG°ads, and ΔS°ads values for various adsorption reactions of pAsA and arsenate from electronic energy and thermal corrections listed in Table S1 in the Supporting Information for all reactants and products. Reactions 7−9 show that the type of water leaving group affects the sign and values of these state functions. Reactions 8 and 9 are the same as reaction 7 forming complex F, but with different water clusters as leaving groups. These are shown to illustrate the importance of the chosen water cluster leaving group when considering reactions of hydrated pAsA− with Fe2(OH)5(OH2)5+. While reactions 8 or 9 yielded more exergonic values of ΔG°ads (i.e., more negative), they do so with an increase in entropy, as seen in the fourth column in Table 4,

Figure 6. Correlation between experimental and calculated v(As−O) frequencies with pAsA on goethite (empty markers) and hematite (solid markers). Assignment is listed in Table 3.

Table 4. ΔH°ads, ΔG°ads (kJ/mol), and ΔS°ads (kJ/mol K) for Ligand Exchange Reactions of pAsA and Arsenate with Iron Oxide Clusters

calculated ones listed in Table 3. The lower frequencies correlate best with inner-sphere MM complexes, while the high experimental frequencies have components from both innerand outer-sphere complexes. This result further confirms our earlier interpretation of spectral data, where we suggested the formation of at least two types of surface complexes for pAsA.43 Despite being a monosubstituted organoaresnical, this finding is similar to our earlier reports on the simultaneous formation of inner- and outer-sphere complexes of DMA, which is a disubstituted organoarsenical under experimental equilibrium conditions with continuous flow of the aqueous phase over the iron (oxyhdyr)oxide film. 37 When these results were complemented with initial adsorption and desorption kinetic measurements,45,69,70 it became clear that with increasing organic substitution on the AsO4 moiety, the number of weakly bonded complexes increases. This results in relatively smaller rates of adsorption for organoarsenicals and faster desorption due to flowing phosphate accompanied by faster ligand exchange of phosphate for surfaces with adsorbed DMA and pAsA relative to adsorbed iAs(V). To obtain further insight into the mechanistic aspects of the ligand exchange process, the sections below investigate in detail the thermodynamic favorability of the formation of pAsA complexes shown in Figure 5, and interconversion between MM and BB complexes for pAsA and iAs(V). Calculated ΔH°ads, ΔG°ads, and ΔS°ads for pAsA Reaction with Iron Oxide Clusters under Simulated Environmental Conditions. Calculating the enthalpy (ΔH°ads) and Gibbs free energy (ΔG°ads) of adsorption for hypothetical reactions can reveal the types of complexes that are thermodynamically favorable. In a neutral environment of 4.1 < pH < 9.2, the pAsA− species would react with a positively charged iron (oxyhydr)oxide surface. Reactions will occur via ligand exchange to form the neutral outer- and inner-sphere complexes A, B, and C. The ΔH°ads and ΔG°ads values of the adsorption reactions were calculated as shown in eqs 1 and 2, where E0 is the electronic energy and Hcorr and Gcorr are the thermal corrections to enthalpy and Gibbs free energies, respectively.71 The entropy, ΔS°ads, was extracted from eqs 1 and 2 using eq 3 where T = 298.15 K.

pAsA adsorption

ΔH°ads ΔG°ads

ΔS°ads

pAsA− · (H2O)4 + [Fe2(OH)5(OH2)5]+ · (H2O)4 → Forming OS, MM and BB complexes A, B, and C respectively 1. (OS) pAsA · (H2O)4 Fe2(OH)5(OH2)4 + −59.3 −48.2 −0.0374 (H2O)5 −103 0.00646 2. (MM) pAsA-Fe2(OH)5(OH2)3 · (H2O)5 + −101 (H2O)5 0.0318 3. (MMa) pAsA-Fe2(OH)5(OH2)3 · (H2O)5 + −77.1 −86.6 (H2O)5 4. (BB) pAsA-Fe2(OH)5(OH2)3 · (H2O)5 + −73.3 −83.8 0.0352 (H2O)5 Arsenate adsorption ΔH°ads ΔG°ads ΔS°ads [H2AsO4]− · (H2O)4 + [Fe2(OH)5(OH2)5]+ · (H2O)4 → Forming MM complexes D and E and BB complex F respectively 5. HAsO4−Fe2(OH)4(OH2)5 ·(H2O)3 + −13.1 −14.1 +0.00310 (H2O)6 −47.0 −41.5 −0.0183 6. HAsO4−Fe2(OH)4(OH2)4 ·(H2O)4 + (H2O)6 7. HAsO4−Fe2(OH)4(OH2)4 ·(H2O)4 + −57.1 −46.0 −0.0372 (H2O)6 Forming BB complex F with different leaving groups (shown in bold) 8. HAsO4−Fe2(OH)4(OH2)4 ·(H2O)4 + −14.2 −71.6 +0.192 2(H2O)3 +36.1 −138 +0.585 9. HAsO4−Fe2(OH)4(OH2)4 ·(H2O)4 + 6(H2O) a

Calculated with M06-2X/6-311+G(d,p) and IEFPCM for comparison.

where ΔS°ads for each reaction is displayed. Using larger water clusters as leaving groups approximates bulk water better than multiplying one water molecule several times. The choice of five-water cluster as the leaving group for reactions 1−4 and a six-water cluster for reactions 5−7 was made to minimize the entropy of each reaction. The energy minimization also correlates with having the same number of clusters on the products side as on the reactants side. We advise using the same number of reactants and products whenever calculating ΔG°ads since that minimizes the entropy change and more accurately depicts bulk water interactions. The minimum energy structures of the water clusters used in these reactions are 5674



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When taking into account the precision of the DFT method for calculating thermodynamic functions (±10 kJ/mol),51 we find that ΔG°ads from B3LYP would range from −93 to −113 kJ/ mol, which slightly overlaps with values from M06-2X, −76.6 to −96.6 kJ/mol. Hence, both of these functionals reproduced experimental observations on pAsA adsorption. On the other hand, Watts et al.51 used the latter functional to calculate ΔG°ads for the formation of arsenate BB complexes on iron clusters: H 2 AsO 4 − ·8H 2 O + Fe 2 (OH)6 (OH 2 ) 4 ·4H 2 O → Fe2(OH)4(OH2)4HAsO4·4H2O + OH−·4H2O + 5H2O. They found that ΔG°ads values would range −13 to +7 kJ/mol, which does not accurately reflect that arsenate adsorption is favorable and was observed experimentally. The arsenate reactions 5−7 in Table 4 show different thermodynamic favorability, with formation of the BB complex being the most favorable followed by the MM complexes. These theoretical results are in agreement with recent experimental findings which used ATR-FTIR to show that the majority of pAsA adsorbs on hematite nanoparticles in a monodentate fashion45 and other experimental and computational work on arsenate that reports BB complexation being the most favorable.29,56,70,73 The activation barriers of pAsA and arsenate for the formation of BB complexes from MM complexes were quantified to explain the differences in complexation for the two arsenicals. The transition state complexes for pAsA and arsenate are shown as complexes G and H respectively in Tables 2 and S1, and their structures are shown in Figure 5. Figure 7 shows transition state (TS) complex G with an activation barrier of 23 kJ/mol for the formation of BB complex C from MM complex B. This is a significant barrier considering that it is almost 10× higher than the thermal energy at room temperature, which is 8.315 J K−1 mol−1 × 298 K = 2.5 kJ/mol. A similar activation barrier of 26 kJ/mol was found for transition state complex H for the formation of BB complex F from MM complex E. The difference between the two activation barriers is not significant enough to explain why arsenate preferentially forms BB complexes while pAsA preferentially forms MM complexes. Hence these theoretical results suggest that the formation of mostly MM complexes for pAsA and mostly BB complexes for arsenate is thermodynamically- rather than kinetically driven due to the high exothermicity of the reactions listed in Table 4. Recently, Farrell and Chaudhary studied the adsorption of arsenate on ferric hydroxides in an aqueous environment using DFT calculations.74 A ΔG°ads of −55 kJ/mol and an activation barrier of 112 kJ/mol were reported for the formation of the neutral BB complex from the MM complex. Also, He et al. studied the adsorption of arsenate on titanium dioxide in an aqueous environment using DFT calculations.75 They calculated a ΔG°ads of −31 kJ/mol and an activation barrier of 94 kJ/ mol for the formation of the neutral BB complex from the MM complex. Although our conclusions regarding arsenate adsorption are generally in agreement with the ones in these studies and our ΔG°ads of −32 kJ/mol is similar, our activation barrier of 23 kJ/mol for the formation BB complex F from MM complex E is much lower. We believe that the most significant difference between these studies and our study is the inclusion of explicit water molecules in our calculations. As noted in the previous section, the explicit inclusion of water molecules often plays a critical role in lowering the energy of the transition state. While the energy-minimized structures of arsenate shown in Figure 5 are stabilized via bonding to explicit water molecules, the inherently unstable

shown in Figure S1 in the Supporting Information. It may be seen that each pair of hydrogen bonded water molecules has a nonplanar structure.72 For comparison, Watts et al.51 recently investigated a number of factors that influence the values of calculated ΔG°ads for arsenate adsorption, namely charge of the iron clusters, and degree of explicit hydration of the iron clusters and arsenate species, H2AsO4− and HAsO42−. It was found that adsorption of either arsenate species onto the more highly charged iron cluster is more energetically favorable, with ΔG°ads values ranging from −263 to −338 kJ/mol. For neutral clusters, adsorption of H2AsO4− is found to be endergonic, whereas adsorption of HAsO42− is exergonic. Also, the authors found that increasing the number of explicit water around the initial neutral iron cluster in the adsorption of H2AsO4− results in more negative ΔG°ads. Values of ΔG°ads for arsenate-Fe product clusters with eight explicit waters were within the ±10 kJ/mol relative to that with four waters. Hence, it is more practical to simulate explicit hydration with four waters to save on computational time, which we did in this paper. Moreover, it was found that hydrating the initial H2AsO4− with eight explicit waters produced more realistic ΔG°ads values with hydrated neutral iron cluster. Hence, the main conclusion was that reactants and products should be hydrated in order to obtain results that are meaningful for comparison with experimental data. While the reactions listed in Table 4 are not exactly the same as those studied by Watts et al.,51 they take into account the above findings on hydration of reactants and products. In this paper, we limited arsenate adsorption reactions to those with H2AsO4− and hydrated positively charged iron clusters as compared with other reactions we calculated earlier.36 For this arsenate species, the formation of neutral MM and BB complexes with water as the main leaving group is more thermodynamically favorable than negatively charged ones.36 Hence, we believe reactions 5−7 are relevant to experimental conditions near neutral pH. Moreover, the pAsA reactions 1−4 in Table 4 are thermodynamically favorable with a negative ΔG°ads, but reactions leading to the formation of MM pAsA are the most favorable, followed by BB and OS as illustrated in Figure 7. Reaction 3 in Table 4 was calculated using the M06-2X/6311+G(d,p) method for comparison with results from the B3LYP functional. Both methods predict exothermic and exergonic reactions for the formation of MM pAsA complexes.

Figure 7. Reaction pathway for pAsA adsorption onto an iron oxide surface and calculated ΔG°ads for complexes A, B, and C and transition state complex G. 5675



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Table 5. Calculated ΔH°des, ΔG°des (kJ/mol), and ΔS°des for pAsA and Arsenate from Iron Oxide using H2PO4− and HPO42− as the Desorbing Agents

transition state complexes shown in Figure 5 are stabilized to a greater extent. This effect can be so pronounced that the terms water-assisted reaction76 and water-catalyzed reaction77 have been used. Although this effect is well-established, understanding the role of explicit water molecules in lowering the activation energy remains an area of current research.78,79 In experimental studies, ΔG°ads is typically calculated from binding affinity constants, Keq extracted from applying the Langmuir adsorption model, or a surface complexation model to room-temperature isotherm data at a given pH, or pH edge data at a given bulk phase concentration (see ref 69 for literature review). While these studies showed that ligand exchange reactions of arsenic compounds with metal (oxyhydr)oxides is spontaneous, they cannot be used to estimate ΔH°ads or ΔS°ads. Few batch studies have been conducted for arsenate adsorption as a function of temperature, typically from 293 to 338 K.80−82 Tuna et al. reported an exothermic reaction (−52 kJ/mol) accompanied by a decrease in entropy (−0.084 kJ/mol K) using a hybrid Fe(III)-activated carbon adsorbent at pH 3.80 Partey et al. reported an endothermic adsorption reaction for arsenate (17.8 kJ/mol) on laterite iron concretions, which contain a mixture of iron, manganese, titanium and aluminum (oxyhydr)oxides, with a net positive entropy change (+0.033 kJ/mol K) at pH 7.81 Similar inconsistencies were found in batch isotherm temperaturedependent studies of phosphate adsorption on soils as critically analyzed by Harvey and Rhue.82 Direct measurements of enthalpy changes for surface reactions have been made through flow adsorption calorimetry.83 For example, Kabengi et al. reported exothermic ligand exchange reactions for aqueous nitrate replacing adsorbed chloride (−4.5 kJ/mol) and potassium replacing adsorbed sodium (−1.2 kJ/mol) on amorphous aluminum hydroxide.84 The same group reported a range of values, from −3 to −66 kJ/ mol, for arsenate adsorption on this material at pH 5.7.85 To our knowledge, there have been no direct measurements of entropy changes upon adsorption of arsenical compounds. As detailed above, reactions in Table 4 with equal numbers of reactants and products were found to have the smallest ΔS°ads values. Yet, from a statistical thermodynamics standpoint,86 it is very likely that the increase in the degrees of freedom in surface arsenic species that form upon adsorption due to their size relative to coordinated water would result in a net increase in the entropy of adsorption. Clearly, more temperature-dependent and calorimetric experiments need to be conducted to further explore this area. Calculated ΔG°des of pAsA(ads) and iAs(ads) Due to Phosphate. Manure contaminated with arsenic also contains phosphate and organic matter, which compete for reactive sites on soil. To simulate the desorption of pAsA(ads) due to phosphate, Gibbs free energies of desorption (ΔG°des) were calculated using eq 1 for a number of reactions that result in phosphate adsorption, as shown Table 5. Phosphate has a second pKa around 7 resulting in two dominant phosphate species, H2PO4− and HPO42−. When pAsA− leaves due to H2PO4− acting as a desorbing agent, the neutral PO4/iron oxide complex L forms. Likewise, when HPO42− acts as a desorbing agent, the negatively charged PO4/iron oxide complex M forms. Similarly, when the phosphate species H2PO4− and HPO42− act as desorbing agents for arsenate, complexes J and K will form. Complexes J and L are essentially the same except for an extra explicit water that is added to complex L to make the reaction with pAsA balance. This is also true for complexes K and M,

Desorption of pAsA using H2PO4‑

ΔH°des ΔG°des

ΔS°des

Desorbing OS, MM and BB Complexes A, 1. Complex A + H2PO4− · (H2O)4→ Complex L + pAsA− · (H2O)4 2. Complex B + H2PO4− · (H2O)4→ Complex L + pAsA− · (H2O)4 3. Complex C + H2PO4− · (H2O)4→ Complex L + pAsA− · (H2O)4 Desorption of pAsA using HPO42−

B, and C, respectively −28.3 −27.6 −0.002 54 13.7 27.5 −0.0464

Desorbing OS, MM and BB Complexes A, 4. Complex A + HPO42− · (H2O)4→ Complex M + pAsA− · (H2O)4 5. Complex B + HPO42− · (H2O)4→ Complex M + pAsA− · (H2O)4 6. Complex C + HPO42− · (H2O)4→ Complex M + pAsA− · (H2O)4 Desorption of arsenate using H2PO4‑

B, and C, respectively −53.0 −47.0 −0.0203

Desorbing MM Complexes D and E and BB 7. Complex D + H2PO4− · (H2O)4→ Complex J + H2AsO4− ·(H2O)4 8. Complex E + H2PO4− · (H2O)4→ Complex J + H2AsO4− ·(H2O)4 9. Complex F + H2PO4− · (H2O)4→ Complex J + H2AsO4− ·(H2O)4 Desorption of arsenate using HPO42−

Complex F respectively −66.3 −55.6 −0.0359

Desorbing MM Complexes D and E and BB 10. Complex D + HPO42− · (H2O)4→ Complex K + H2AsO4− ·(H2O)4 11. Complex E + HPO42− · (H2O)4→ Complex K + H2AsO4− ·(H2O)4 12. Complex F + HPO42− · (H2O)4→ Complex K + H2AsO4− ·(H2O)4

Complex F respectively −81.6 −65.1 −0.0551

−14.3

8.08

ΔH°des ΔG°des

−0.0752 ΔS°des

−11.0

8.13

−0.0642

−39.0

−11.3

−0.0929

ΔH°des ΔG°des

ΔS °des

−32.5

−28.2

−0.0145

−10.7

−6.94

−0.0125

ΔH°des ΔG°des

ΔS°des

−47.7

−37.7

−0.0338

−25.9

−16.4

−0.0318

where complex M has an extra explicit water to make the reaction with pAsA balance. The optimized geometries of complexes J − M are shown in Figure 5, while the energies for Complexes J - M are shown in Table S1 of the Supporting Information. Our complexes J−M with 4 and 5 explicit waters were compared to those calculated by Kubicki et al.30 with 8 explicit waters and only very minor differences in geometries were observed. The BB phosphate complexes calculated by Kubicki et al. had P−O distances of 1.57 Å, while our values ranged from 1.51−1.55 Å for complexes J through M. Similarly, for the interatomic distance of P−Fe, the calculated values of Kubicki et al. ranged from 3.16−3.22 Å while ours ranged from 3.19−3.29 Å for complexes J through M. For reactions 1−3 and 7−9 in Table 5, H2PO4− is used as the desorbing species for pAsA complexes A−C, and arsenate complexes D−F respectively to form the phosphate complexes L and J. Similarly, for reactions 4−6 and 10−12, HPO42− is used as the desorbing agent to form phosphate complexes K and M. Table 5 shows that while arsenate removal may be thermodynamically favorable with either of the two phosphate species, pAsA might require the more deprotonated HPO42− species for better removal, especially for inner-sphere complexes (reactions 5 and 6). Desorption of the OS pAsA complex is highly favorable with both desorbing species. Thus, desorption favorability of pAsA decreases in this order: OS > BB > MM, where desorption of inner-sphere complexes may depend greatly on the acidity of the environment. 5676



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(3) European Union. Council Directive 1999/29/Ec of 22 April 1999 on the Undesirable Substances and Products in Animal Nutrition. http://ec.europa.eu/food/fs/afs/substances/substances01_en.pdf. (4) Health Canada. Pfizer and Dominion Veterinary Laboratories Suspending Sales of Veterinary Drugs Containing Roxarsone (Information Note on June 22, 2011). http://www.hc-sc.gc.ca/dhpmps/vet/issues-enjeux/roxarsone-eng.php. (5) U.S. Food and Drug Administration, FDA: Pfizer Will Voluntarily Suspend Sale of Animal Drug 3-Nitro (FDA Press Release June 8, 2011). http://www.fda.gov/NewsEvents/Newsroom/ PressAnnouncements/ucm258342.htm. (6) Jones, F. T. A Broad View of Arsenic. Poult. Sci. 2007, 86, 2−14 and references therein. (7) Cullen, W. R.Is Arsenic an Aphrodisiac? The Sociochemistry of an Element; RSC Publishing: Cambridge, 2008. (8) Canadian Food Inspection Agency-Arsanilic Acid (Date Revised 2006−09). http://www.inspection.gc.ca/animals/feeds/medicatingingredients/mib/mib-4/eng/1330714521085/1330716893318. (9) Nachman, K. E.; Graham, J. P.; Price, L. B.; Silbergeld, E. K. Arsenic: A Roadblock to Potential Animal Waste Management Solutions. Environ. Health Perspect. 2005, 113, 1123−1124. (10) Garbarino, J. R.; Bednar, A. J.; Rutherford, D. W.; Beyer, R. S.; Wershaw, R. L. Environmental Fate of Roxarsone in Poultry Litter. I. Degradation of Roxarsone During Composting. Environ. Sci. Technol. 2003, 37, 1509−1514. (11) Cortinas, I.; Filed, J. A.; Kopplin, M.; Garbarino, J. R.; Gandolfi, A. J.; Sierra-Alvarez, R. Anaerobic Biotransformation of Roxarsone and Related N-Substituted Phenylarsonic Acids. Environ. Sci. Technol. 2006, 40, 2951−2957. (12) Arroyo-Abad, U.; Mattusch, J.; Moder, M.; Elizalde-Gonzalez, M. P.; Wennrich, R.; Matysik, F. M. Identification of Roxarsone Metabolites Produced in the System: Soil-Chlorinated Water-Light by Using HPLC-ICP-MS/ESI-MS, HPLC-ESI-MS/MS and High Resolution Mass Spectrometry (ESI-TOF-MS). J. Anal. At. Spectrom. 2011, 26, 171−177. (13) Bednar, A. J.; Garbarino, J. R.; Ferrer, I.; Rutherford, D. W.; Wershaw, R. L.; Ranville, J. F.; Wildeman, T. R. Photodegradation of Roxarsone in Poultry Litter Leachates. Sci. Total Environ. 2003, 302, 237−245. (14) Stolz, J. F.; Perera, E.; Kilonzo, B.; Kail, B.; Crable, B.; Fisher, E.; Ranganthan, M.; Wormer, L.; Basu, P. Biotransformation of 3-Nitro-4Hydroxybenzene Arsonic Acid (Roxarsone) and Release of Inorganic Arsenic by Clostridium Species. Environ. Sci. Technol. 2007, 41, 818− 823. (15) Makris, K. C.; Quazi, M.; Punamiya, P.; Sarkar, D.; Datta, R. Fate of Arsenic in Swine Waste from Concentrated Animal Feeding Operations. J. Environ. Qual. 2008, 37, 1626−1633. (16) Jackson, B. P.; Bertsch, P. M. Determination of Arsenic Speciation in Poultry Wastes by IC-ICP-MS. Environ. Sci. Technol. 2001, 35, 4868−4873. (17) Bissena, M.; Frimmel, F. H. Arsenic - a Review. Part I: Occurrence, Toxicity, Speciation, Mobility. Acta Hydrochim. Hydrobiol. 2003, 31, 9−18. (18) Dopp, E.; Hartmann, L. M.; Florea, A.-M.; Rettenmeier, A. W.; Hirner, A. V. Enviornmental Distribution, Analysis, and Toxicity of Organometal(Loid) Compounds. Critical Rev. Toxic. 2004, 34, 301− 333. (19) Kenyon, E. M.; Hughes, M. F. A Concise Review of the Toxicity and Carcinogenicity of Dimethylarsinic Acid. Toxicology 2001, 160, 227−236. (20) Tchounwou, P. B.; Centeno, J. A.; Patlolla, A. K. Arsenic Toxicity, Mutagenesis, and Carcinogenesisa Health Risk Assessment and Management Approach. Mol. Cell. Biochem. 2004, 255, 47−55. (21) Hirano, S.; Kobayashi, Y.; Hayakawa, T.; Cui, X.; Yamamoto, M.; Kanno, S.; Shraim, A. Accumulation and Toxicity of Monophenyl Arsenicals in Rat Endothelial Cells. Arch. Toxicol. 2005, 79, 54−61. (22) Abhyankar, L. N.; Jones, M. R.; Guallar, E.; Navas-Acien, A. Arsenic Exposure and Hypertension: A Systematic Review. Environ. Health Perspect. 2012, 120, 494−500.

CONCLUSIONS AND SIGNIFICANCE Results reported herein from DFT calculations show that the formation of inner- and outer-sphere pAsA/iron oxide complexes is thermodynamically favorable, with the MM being the most favorable. Transition state calculations between MM and BB complexes showed that activation energies greater than 23 kJ/mol are required to form the BB pAsA/iron oxide complexes, and that formation of arsenate BB complexes is thermodynamically, rather than kinetically, driven. The explicit inclusion of water molecules plays a critical role in lowering the energies of the transition states for these reactions and is essential in order to obtain accurate activation energies. Desorption energies of pAsA using phosphate ions showed higher favorability using HPO42− species suggesting that desorption efficiency is pH-dependent. While the above studies were conducted on a simple model of the iron oxide surface, the results are significant as they constitute systematic theoretical investigations of pAsA surface interactions with iron oxides and have provided mechanistic details that are challenging to obtain experimentally. For example, from the frequency calculations, it is now confirmed that the 840 cm−1 spectral component we used previously to construct adsorption isotherms, pH envelopes, and kinetic curves43,45,87 arises from inner- and outer-sphere surface complexes of pAsA. However, when combined with thermodynamic favorability and transition state calculations, it is very likely that this spectral component has major contributions from MM surface complexes. Moreover, the results on enthalpy and entropy of adsorption are reported for the first time for these complexes, which highlight three key areas of further investigation: effect of leaving group types and number (e.g., multiples of water monomers vs water clusters), need for temperature-dependent calculations and “wet” experiments in addition to calorimetric measurements of adsorption heats, and need for similar calculations on extended model surfaces of iron oxides.

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ASSOCIATED CONTENT

S Supporting Information *

Figures and tables showing geometries and listing structural parameters. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*Phone: (519)884-0710, ext. 2873. Fax: (519)746-0677. Email: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors acknowledge Marshall Linder and partial funding from Laurier, NSERC and Early Researcher Award from Ontario’s Ministry of Research and Innovation.

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REFERENCES

(1) Zhang, F. S.; Li, Y. X.; Yang, M.; Li, W. Content of Heavy Metals in Animal Feeds and Manures from Farms of Different Scales in Northeast China. Int. J. Environ. Res. Pub. Health 2012, 9, 2658−2668. (2) Hygienical Standard for Feeds [HSF (GB 13078−2001)] http:// www.chinafeed.org.cn/mem/jiangsu/center/center3-1.htm. 5677



Article

(23) Sparks, D. L. Environmental Soil Chemistry, 1st ed.; Academic Press: San Diego, 1995. (24) Hind, A. R.; Bhargava, S. K.; McKinnon, A. At the Solid/Liquid Interface: FTIR/ATR- the Tool of Choice. Adv. Colloid Interface Sci. 2001, 93, 91−114. (25) Brown Jr., G. E.; Parks, G. A.; O’Day, P. A. Sorption at MineralWater Interfaces: Macroscopic and Microscopic Perspectives. In Mineral Surfaces; Vaughan, D. J., Pattrick, R. A. D., Eds.; Chapman & Hall: New York, 1995; pp 129−183. (26) Cygan, R. T. Molecular Modeling in Mineralogy and Geochemistry. In Reviews in Mineralogy and Geochemistry; Mineralogical Society of America: Washington, DC, 2001; Vol. 42, pp 1−35. (27) Kubicki, J. D. Comparison of As(III) and As(V) Complexation onto Al- and Fe-Hydroxides. In Advances in Arsenic Research: Integration of Experimental and Observational Studies and Implications for Mitigation; O’Day, P. A., et al., Eds.; ACS Symposium Series: Washington, DC, 2005; Vol. 915, pp 104−117. (28) Kubicki, J. D. Computational Chemistry Applied to Studies of Organic Contaminants in the Environment: Examples Based on Benzo[a]Pyrene. Am. J. Sci. 2005, 305, 621−644. (29) Kubicki, J. D.; Kwon, K. D.; Paul, K. W.; Sparks, D. L. Surface Complex Structures Modelled with Quantum Chemical Calculations: Carbonate, Phosphate, Sulphate, Arsenate and Arsenite. Eur. J. Soil Sci. 2007, 58, 932−944. (30) Kubicki, J. D.; Paul, K. W.; Kabalan, L.; Zhu, Q.; Mrozik, M. K.; Aryanpour, M.; Pierre-Louis, A.-M.; Strongin, D. R. ATR−FTIR and Density Functional Theory Study of the Structures, Energetics, and Vibrational Spectra of Phosphate Adsorbed onto Goethite. Langmuir 2012, 28, 14573−14587. (31) Kubicki, J. D.; Schroeter, L. M.; Itoh, M. J.; Nguyen, B. N.; Apitz, S. E. Attenuated Total Reflectance Fourier-Transform Infrared Spectroscopy of Carboxylic Acids Adsorbed onto Mineral Surfaces. Geochim. Cosmochim. Acta 1999, 63, 2709−2725. (32) Kubicki, J. D.; Blake, G. A.; Apitz, S. E. Molecular Orbital Calculations for Modeling Acetate-Aluminosilicate Adsorption and Dissolution Reactions. Geochim. Cosmochim. Acta 1997, 61, 1031− 1046. (33) Tomasi, J.; Mennucci, B.; Cammi, R. Quantum Mechanical Continuum Solvation Models. Chem. Rev. 2005, 105, 2999−3093. (34) Adamescu, A.; Gray, H.; Stewart, K.; Hamilton, I. P.; AlAbadleh, H. A. Trends in the Frequencies of v(AsOxHx−1) [x = 2−4] in Selected As(V)-Containing Compounds Investigated Using Quantum Chemical Calculations. Can. J. Chem. 2010, 88, 65−77. (35) Cowen, S.; Duggal, M.; Hoang, T. N.; Al-Abadleh, H. A. Vibrational Spectroscopic Characterization of Some EnvironmentallyImportant Organoarsenicals: A Guide for Understanding the Nature of Their Surface Complexes. Can. J. Chem. 2008, 86, 942−950. (36) Adamescu, A.; Hamilton, I. P.; Al-Abadleh, H. A. Thermodynamics of Dimethylarsinic Acid and Arsenate Interactions with Hydrated Iron-(Oxyhydr)Oxide Clusters: DFT Calculations. Environ. Sci. Technol. 2011, 45, 10438−10444. (37) Adamescu, A.; Mitchell, W.; Hamilton, I. P.; Al-Abadleh, H. A. Insights into the Surface Complexation of Dimethylarsinic Acid on Iron (Oxyhydr)Oxides from ATR-FTIR Studies and Quantum Chemical Calculations. Environ. Sci. Technol. 2010, 44, 7802−7807. (38) Brown, B. L.; Slaughter, A. D.; Schreiber, M. E. Controls on Roxarsone Transport in Agricultural Watersheds. Appl. Geochem. 2005, 20, 123−133. (39) Jackson, B. P.; Seaman, J. C.; Bertsch, P. M. Fate of Arsenic Compounds in Poultry Litter Upon Land Application. Chemosphere 2006, 65, 2028−2034. (40) Rutherford, D. W.; Bednar, A. J.; Garbarino, J. R.; Needham, R.; Staver, K. W.; Wershaw, R. L. Environmental Fate of Roxarsone in Poultry Litter: II. Mobility of Arsenic in Soils Amended with Poultry Litter. Environ. Sci. Technol. 2003, 37, 1515−1520. (41) Oyewuni, O.; Schreiber, M. E. Release of Arsenic and Other Trace Elements from Poultry Litter: Insights from a Field Experiment on the Delmarva Peninsula, Delaware. Appl. Geochem. 2012, 27, 1979−1990.

(42) Arai, Y.; Lanzirotti, A.; Sutton, S.; Davis, J. A.; Sparks, D. L. Arsenic Speciation and Reactivity in Poultry Litter. Environ. Sci. Technol. 2003, 37, 4083−4090 and references therein. (43) Depalma, S.; Cowen, S.; Hoang, T. N.; Al-Abadleh, H. A. Adsorption Thermodynamics of p-Arsanilic Acid on Iron (Oxyhydr)Oxides: In-situ ATR-FTIR Studies. Environ. Sci. Technol. 2008, 42, 1922−1927. (44) Chabot, M.; Hoang, T. N.; Al-Abadleh, H. A. ATR-FTIR Studies on the Nature of Surface Complexes and Desorption Efficiency of p-Arsanilic Acid on Iron-(Oxyhydr)oxides. Environ. Sci. Technol. 2009, 43, 3142−3147. (45) Arts, D.; Sabur, M. A.; Al-Abadleh, H. A. Surface Interactions of Aromatic Organoarsenical Compounds with Hematite Nanoparticles Using ATR-FTIR: Kinetic Studies. J. Phys. Chem. A 2013, 117, 2195− 2204. (46) Shared Hierarchical Academic Research Computing Network (SHARCNET). http://www.sharcnet.ca. (47) Becke, A. D. A New Mixing of Hartree−Fock and Local Density-Functional Theories. J. Chem. Phys. 1993, 98, 1372−1377. (48) Lee, C. T.; Yang, W. T.; Parr, R. G. Development of the ColleSalvetti Correlation-Energy Formula into a Functional of the ElectronDensity. Phys. Rev. B 1988, 37, 785−789. (49) Adamescu, A. Density Functional Theory Calculations on Hydrated Dimethylarsinic Acid and Iron Oxide Clusters. MSc, Wilfrid Laurier University, 2012. http://scholars.wlu.ca/etd/1125. (50) Zhao, Y.; Truhlar, D. G. The M06 Suite of Density Functionals for Main Group Thermochemistry, Thermochemical Kinetics, Noncovalent Interactions, Excited States, and Transition Elements. Theor. Chem. Acc. 2008, 120, 215−241. (51) Watts, H. D.; Tribe, L.; Kubicki, J. D. Arsenic Adsorption onto Minerals: Connecting Experimental Observations with Density Functional Theory Calculations. Minerals 2014, 4, 208−240. (52) Nuallain, C. O.; Cinneide, S. O. Thermodynamic Ionization Constants of Aromatic Arsonic Acids. J. Inorg. Nucl. Chem. 1973, 35, 2871−2881. (53) Kosmulski, M. Ph-Dependent Surface Charging and Points of Zero Charge III. Update. J. Colloid Interface Sci. 2006, 298, 730−741. (54) Dewar, M. J. S.; Schmeising, H. N. A Re-Evaluation of Conjugation and Hyperconjugation: The Effects of Changes in Hybridisation on Carbon Bonds. Tetrahedron 1959, 5, 166−178. (55) Nuttall, R. H.; Hunter, W. N. P-Arsanilic Acid, a Redetermination. Acta Cryst. C: Cryst. Struc. Comm. 1996, C52, 1681−1683. (56) Sherman, D. M.; Randall, S. R. Surface Complexation of Arsenic(V) to Iron(III) (Hydr)Oxides: Structural Mechanism from Ab Initio Molecular Geometries and EXAFS Spectroscopy. Geochem. Cosmochim. Acta 2003, 67, 4223−4230. (57) O’Reilly, S. E.; Strawn, D. G.; Sparks, D. L. Residence Time Effects on Arsenate Adsorption/Desorption Mechanisms on Goethite. Soil Sci. Soc. Am. J. 2001, 65, 67−77. (58) Waychunas, G. A.; Davis, J. A.; Fuller, C. C. Geometry of Sorbed Arsenate on Ferrihydrite and Crystalline FeOOH: ReEvaluation of EXAFS Results and Topological Factors in Predicting Sorbate Geometry, and Evidence for Monodentate Complexes. Geochem. Cosmochim. Acta 1995, 59, 3655−3661. (59) Farquhar, M. L.; Charnock, J. M.; Livens, F. R.; Vaughan, D. J. Mechanisms of Arsenic Uptake from Aqueous Solution by Interaction with Goethite, Lepidocrocite, Mackinawite, and Pyrite: An X-Ray Absorption Spectroscopy Study. Environ Sci Technol. 2002, 36, 1757− 1762. (60) Shimizu, M.; Arai, Y.; Sparks, D. L. Multiscale Assessment of Methylarsenic Reactivity in Soil. 1. Sorption and Desorption on Soils. Environ. Sci. Technol. 2011, 45, 4293−4299. (61) Shimizu, M.; Ginder-Vogel, M.; Parikh, S. J.; Sparks, D. L. Molecular Scale Assessment of Methylarsenic Sorption on Aluminum Oxide. Environ. Sci. Technol. 2010, 44, 612−617. (62) Jing, C.; Xiaogunag, M.; Liu, S.; Baidas, S.; Patraju, R.; Christodoulatos, C.; Korfiatis, G. P. Surface Complexation of Organic 5678



Article

Arsenic on Nanocrystalline Titanium Oxide. J. Colloid Interface Sci. 2005, 290, 14−21. (63) Wu, J.-Z.; Ho, P. C. Speciation of Inorganic and Methylated Arsenic Compounds by Capillary Zone Electrophoresis with Indirect UV Detection. Application to the Analysis of Alkali Extracts of As2S2(Realgar) and As2S3(Orpiment). J. Chromatogr. A 2004, 1026, 261−270. (64) Loring, J. S.; SAndstrom, M. H.; Noren, K.; Persson, P. Rethinking Arsenate Coordination at the Surface of Goethite. Chem. Eur. J. 2009, 15, 5063−5072. (65) Catalano, J. G.; Park, C.; Fenter, P.; Zhang, Z. Simultaneous Inner- and Outer-Sphere Arsenate Adsorption on Corundum and Hematite. Geochem. Cosmochim. Acta 2008, 72, 1986−2004. (66) Jing, C.; Meng, X.; Calvache, E.; Jiang, G. Remediation of Organic and Inorganic Arsenic Contaminated Groundwater Using a Nanocrystalline TiO2-Based Adsorbent. Environ. Pollut. 2009, 157, 2514−2519. (67) Catalano, J. G.; Zhang, Z.; Park, C.; Fenter, P.; Bedzyk, M. J. Bridging Arsenate Surface Complexes on the Hematite (012) Surface. Geochem. Cosmochim. Acta 2007, 71, 1883−1897. (68) Irikura, K. K.; Johnson, R. D.; Kacker, R. N. Uncertainties in Scaling Factors for Ab Initio Vibrational Frequencies. J. Phys. Chem. A 2005, 109, 8430−8437. (69) Tofan-Lazar, J.; Al-Abadleh, H. A. ATR-FTIR Studies on the Adsorption/Desorption Kinetics of Dimethylarsinic Acid on Iron(Oxyhydr)Oxides. J. Phys. Chem. A 2012, 116, 1596−1604. (70) Tofan-Lazar, J.; Al-Abadleh, H. A. Kinetic ATR-FTIR Studies on Phosphate Adsorption on Iron-(Oxyhydr)oxides in the Absence and Presence of Surface Arsenic: Molecular-Level Insights into the Ligand Exchange Mechanism. J. Phys. Chem. A 2012, 116, 10143−10149. (71) Thermochemistry in Gaussian [Online]. http://www.gaussian. com/g_whitepap/thermo.htm. (72) Xu, X.; Goddard, W. A. Bonding Properties of the Water Dimer: A Comparative Study of Density Functional Theories. J. Phys. Chem. A 2004, 108, 2305−2313. (73) Fukushi, K.; Sverjensky, D. A. A Predictive Model (ETLM) for Arsenate Adsorption and Surface Speciation on Oxides Consistent with Spectroscopic and Theoretical Molecular Evidence. Geochem. Cosmochim. Acta 2007, 71, 3717−3745. (74) Farrell, J.; Chaudhary, B. K. Understanding Arsenate Reaction Kinetics with Ferric Hydroxides. Environ. Sci. Technol. 2013, 47, 8342− 8347. (75) He, G.; Pan, G.; Zhang, M. Studies on the Reaction Pathway of Arsenate Adsorption at Water−TiO2 Interfaces Using Density Functional Theory. J. Colloid Interface Sci. 2011, 364, 476−481. (76) Antonczak, S.; Ruiz-Mpez, M. F.; Rivail, J. L. J. Am. Chem. Soc. 1994, 116, 3912−3921. (77) Vö hringer-Martinez, E.; Hansmann, B.; Hernandez, H.; Francisco, J. S.; Troe, J.; Abel, B. Water Catalysis of a RadicalMolecule Gas-Phase Reaction. Science 2007, 315, 497−501. (78) Du, B.; Zhang, W. Theoretical Study on the Water-Assisted Reaction of NCO with HCHO. J. Phys. Chem. A 2013, 117, 6883− 6892. (79) Hatanaka, M.; Morokuma, K. Role of Water in Mukaiyama− Aldol Reaction Catalyzed by Lanthanide Lewis Acid: A Computational Study. J. Am. Chem. Soc. 2013, 135, 13972−13979. (80) Tuna, A. Ö .; Ö zdemir, E.; Simsek, E. B.; Beker, U. Removal of As(V) from Aqueous Solution by Activated Carbon-Based Hybrid Adsorbents: Impact of Experimental Conditions. Chem. Eng. J. 2013, 223, 116−128. (81) Partey, F.; Norman, D.; Ndur, S.; Nartey, R. Arsenic Sorption onto Laterite Iron Concretions: Temperature Effect. J. Colloid Interface Sci. 2008, 321, 493−500. (82) Harvey, O. R.; Rhue, R. D. Kinetics and Energetics of Phosphate Sorption in a Multi-Component Al(III)−Fe(III) Hydr(Oxide) Sorbent System. J. Colloid Interface Sci. 2008, 322, 384−393. (83) Groszek, A. J. Flow Adsorption Microcalorimetry. Thermochim. Acta 1998, 312, 133−143.

(84) Kabengi, N. J.; Rhue, R. D.; Daroub, S. H. Using Flow Calorimetry to Determine the Molar Heats of Cation and Anion Exchange and the Point of Zero Net Charge on Amorphous Aluminum Hydroxides. Soil Sci. 2006, 171, 13−20. (85) Kabengi, N. J.; Daroub, S. H.; Rhue, R. D. Energetics of Arsenate Sorption on Amorphous Aluminum Hydroxides Studied Using Flow Adsorption Calorimetry. J. Colloid Interface Sci. 2006, 297, 89−94. (86) McQuarrie, D. A. Statistical Thermodynamics; Harper & Row Publishers: New York, 1973. (87) Mitchell, W.; Goldberg, S.; Al-Abadleh, H. A. In-situ ATR-FTIR and Surface Complexation Modeling Studies on the Adsorption Thermodynamics of Mono- and Di-Substituted Organoarsenicals on Iron-(Oxyhydr)Oxides. J. Colloid Interface Sci. 2011, 358, 534−540.

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