Molecular-Scale Study of Aspartate Adsorption on Goethite and

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Molecular-Scale Study of Aspartate Adsorption on Goethite and Competition with Phosphate Yanli Yang, Shengrui Wang, Yisheng Xu, Binghui Zheng, and Jingyang Liu Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.5b05450 • Publication Date (Web): 12 Feb 2016 Downloaded from http://pubs.acs.org on February 13, 2016

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Environmental Science & Technology

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Molecular-Scale Study of Aspartate Adsorption on Goethite and Competition

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with Phosphate

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Yanli Yang,†, ‡ Shengrui Wang,*, †, ‡ Yisheng Xu,† Binghui Zheng,† and Jingyang Liu§

4 †

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State Key Laboratory of Environmental Criteria and Risk Assessment, Chinese Research Academy of Environmental Sciences, Beijing 100012, China

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State Environmental Protection Key Laboratory for Lake Pollution Control,

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Research Center of Lake Eco-environment, Chinese Research Academy of

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Environmental Sciences, Beijing 100012, China

11 §

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State Key Laboratory of Environmental Protection Ecology Industry,

Chinese Research Academy of Environmental Sciences, Beijing 100012, China

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ACS Paragon Plus Environment


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ABSTRACT

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Knowledge of the interfacial interactions between aspartate and minerals, especially its

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competition with phosphate, is critical to understanding the fate and transport of amino

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acids in the environment. Adsorption reactions play important roles in the mobility,

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bioavailability, and degradation of aspartate and phosphate. Attenuated total reflectance

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Fourier-transform infrared (ATR-FTIR) measurements and density functional theory

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(DFT) calculations were used to investigate the interfacial structures and their relative

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contributions in single-adsorbate and competition systems. Our results suggest three

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dominant mechanisms for aspartate: bidentate inner-sphere coordination involving both

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α- and γ-COO−, outer-sphere complexation via electrostatic attraction and H-bonding

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between aspartate NH2 and goethite surface hydroxyls. The interfacial aspartate is

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mainly governed by pH and is less sensitive to changes of ionic strength and aspartate

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concentration. The phosphate competition significantly reduces the adsorption capacity

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of aspartate on goethite. Whereas phosphate adsorption is less affected by the presence

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of aspartate, including the relative contributions of diprotonated monodentate,

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monoprotonated bidentate, and nonprotonated bidentate structures. The adsorption

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process facilitates the removal of bioavailable aspartate and phosphate from the soil

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solution as well as from the sediment pore water and the overlying water.

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TOC Art

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INTRODUCTION

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Amino acids are an important class of chemicals in the environment, which are

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widely spread in soils, sediments, and natural waters. The net charge of amino acids

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varies with pH because of the deprotonation of NH3+ and COOH, which leads to a series

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of interesting adsorption behaviors.1 The adsorption of amino acids on minerals can

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retard their migration and alter their bioavailability in soils and sediments, and the

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interfacial structures may determine their mechanism and difficulty of degradation.

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Therefore, knowledge of amino acids’ molecular-level interactions is critical to

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understanding their fate and transport in the environment as well as the geochemical

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cycle of nitrogen. Notably, amino acids are considered to be the fundamental foundation

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for life. In-depth studies of their adsorption on minerals are needed to evaluate

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biomineralization, the viability of metal implants in the human body, and the origin and

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early evolution of life on earth.1-3

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Aspartate is the simplest acidic amino acid containing one NH3+ and two COOH

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groups, which is frequently detected at high levels in natural waters.4 In contrast to the

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extensive reports of dicarboxylate at mineral/aqueous interfaces, only few studies have

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investigated the adsorption of aspartate on TiO2,5-10 Al2O3,3 and clay minerals.11, 12

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Interestingly, wide diversities have been observed in the mechanisms and interfacial

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configurations depending on the mineral phase. Recent research on aspartate adsorption

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has mainly focused on TiO2, although it usually exists as a minor accessory phase in

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most rocks and soils.3 Parikh et al. observed that aspartate reacted with TiO2 via outer-

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sphere coordination, without inducing peak shifts indicative of covalent bonding.5 In 4


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contrast, other reports have indicated inner-sphere complexation.6-9 Jonsson et al.

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resolved a bridging bidentate structure through both aspartate COO− and an outer-

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sphere (or H-bonded) species.10 Notably, adsorption mechanisms are largely

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determined by the adsorbate structure and mineral phase. Therefore, the adsorption of

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aspartate on different minerals must still be investigated. Aluminum oxides and iron

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oxides are much more relevant than TiO2 to the fields of earth and environmental

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science. Greiner et al. observed that aspartate attached to the γ-Al2O3/D2O interface via

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bidentate and tetradentate coordination and outer-sphere complexation.3 Because of its

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high specific surface area, goethite has strong affinities for small organic acids,

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oxyanions, and heavy metals.13-16 However, very little is known about the adsorption

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behaviors of amino acids on this model mineral.17 Insight into the molecular-level

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process of aspartate on goethite motivates our experimental and theoretical research.

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Unlike the simple systems with one amino acid and one mineral phase, the natural

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environment is more complicated involving abundant oxyanions, metal ions, dissolved

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organic matter, etc. The importance of competitive adsorption (or coadsorption) with

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coexisting substance cannot be overstated.18 The adsorption of glutamate, lysine, and

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aspartate has recently been studied in the presence of Ca2+ (and Mg2+) via batch

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adsorption experiments and surface complexation modeling.1, 19 Limited knowledge is

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available for the molecular-level mechanisms of amino acids on minerals in

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competition with phosphate, particularly the in-situ adsorption process. The insufficient

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study inspires our research to reveal the interfacial structures, the adsorption capacity,

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and relative content of each surface complex of aspartate and phosphate in binary5



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adsorbate systems. The interfacial configurations of phosphate on goethite have long

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been sought but without consistent conclusions.20-23 Most reports supported the inner-

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sphere coordination in different structures, while Kubicki proposed the coexistence of

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inner- and outer-sphere complexation.23 Furthermore, there have been a few studies

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with respect to the phosphate adsorption in the presence of arsenate, U(VI), Cr(VI),

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humic acid, and organic pollutants.24-28 The existing reports provide valuable

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information for reference and comparison.

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The aspartate adsorption on goethite was explored under the effects of pH, ionic

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strength, aspartate concentration, and phosphate competition. The objective of this

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study was to reveal the interfacial structures and relative content of each complex at the

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molecular level, which have great impact on the mobility, bioavailability, and

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degradation of the adsorbates. Besides the ATR-FTIR measurements, theoretical

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calculations were used to determine the configurations of adsorbed aspartate and

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phosphate. The two-dimensional correlation spectroscopy (2D-COS) was employed to

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identify the number of interfacial phosphate structures and peaks belonging to each.

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The relative contributions of each surface complex were analyzed because of their

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sensitive responses to spectral changes in different experimental conditions. This

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research provides new and complementary information for understanding the

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adsorption of aspartate and phosphate on metal oxides and can be used to predict and

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describe the behaviors of amino acids in the environment.

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EXPERIMENTAL SECTION

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1. Materials. L-Aspartic acid and NaH2PO4 were purchased from Sigma-Aldrich.

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All of the chemicals were of analytical or guaranteed reagent grade and were used as

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received. The samples were prepared in Milli-Q water that was boiled for 60 min and

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cooled with N2 purging to remove CO2. Goethite was prepared according to the

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procedure described elsewhere.29 The point of zero charge (PZC) and Brunauer-

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Emmett-Teller (BET) surface area of goethite were determined as pH 8.9 and 84.7 m2/g,

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respectively.18, 29 The particles are well-crystallized needles with a length of 100−200

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nm, which are mainly terminated by (110) and (100) planes.29, 30

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2. ATR-FTIR Spectroscopy Study. The ATR spectra were recorded with a Perkin-

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Elmer Spectrum 100 instrument that was equipped with an MCT detector, a constant

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flow pump, and a horizontal ATR accessory (PIKE Technologies, USA). The flow cell

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was fitted with a 45° ZnSe or Ge crystal. The usable pH for the ZnSe crystal was chosen

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within the range of 5−9 to avoid etching. The Ge crystal was used with pH levels less

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than 5 or greater than 9. Spectra acquisition, subtraction, normalization, and baseline

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correction were performed using the Spectrum software (Perkin-Elmer, Inc., USA). All

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of the spectra were collected with 256 scans at 4 cm−1 resolution to reduce noise. The

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peak numbers and positions were justified using the second derivative. Curve-fitting

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analysis of the overlapping peaks was conducted using Gaussian line shapes.31, 32

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The spectra of aspartate and phosphate solutions were obtained by subtracting the

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spectrum of background electrolyte at the same pH from the sample spectrum. The

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interfacial aspartate and phosphate were measured with a goethite film on the ATR 7



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crystal. The film was coated onto the crystal by applying 1 mL of 1 g/L goethite

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suspension and drying it in an oven at 50 °C for 1 h. Prior to use, the film was gently

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rinsed with Milli-Q water to remove loosely deposited particles. The background

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electrolyte at a predetermined pH was passed through the flow cell at a rate of 0.25

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mL/min until equilibrium was established, namely there was no further change in the

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spectra. The background spectrum was collected, and the solution was then changed to

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the sample with the same pH and ionic strength. The interfacial spectra were recorded

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as a function of time for 60 min. In single-adsorbate systems, the total concentrations

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of aspartate solutions ranged from 1 to 10 mM to obtain high-quality spectra. In

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comparison with the absorbance of interfacial aspartate, contributions from the

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dissolved species are negligible. The dynamic process at pH 6 was also investigated by

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flowing 0.1 M NaCl solution after 1 mM aspartate adsorption for ~30 min. In

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competition systems, aspartate and phosphate at the concentration of 1 mM was

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analyzed in 0.1 M NaCl solution. The slight dissolution of goethite had no detectable

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effect during the flow-cell measurements.

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3. Quantum Chemical Calculations. Geometry optimization and frequency

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calculations were performed using the Gaussian 03 program33 with the hybrid DFT

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B3LYP method. The clusters consisting of one single, two or three edge-sharing

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Fe(III)-octahedra were used to model the optimized geometries of aspartate and

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phosphate on goethite surface. The cluster size was selected to satisfy the basic

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demands for structure calculations of the interfacial configurations. For example, the

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aspartate bidentate structure involving both COO− was simulated with a three edge8


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sharing octahedra. The distance between the neighbouring surface Fe atoms is similar

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to that between the two O atoms in one aspartate COO−, but is much smaller than that

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from different COO−. These small cluster models are efficient for frequency

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calculations of small organic compounds and oxyanions on minerals.29, 31, 34-36 The

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clusters can reproduce main features of the optical response at the interface, though the

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small unit cells do not provide complete representations of the mineral surfaces.37, 38

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Moreover, this technique is computationally much more tractable than the periodic slab

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models which have not been extensively tested for the ability to calculate the vibrational

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spectra of surface complexes on minerals.23 The solvation effect was considered by

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placing explicit H2O around the dissolved and adsorbed species in a gas phase.31, 34, 36

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The number of H2O molecules for inner-sphere coordination was investigated to

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confirm the accuracy of frequency calculations and configuration determinations. The

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frequencies were calculated using the 6-31+G(d, p) basis set on C, H, O, N, and P with

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a scale factor of 0.964, coupled to a LanL2DZ basis set on Fe with a scale factor of

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0.961.39

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4. Two-Dimensional Correlation Spectroscopy. The 2D-COS can greatly enhance

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the resolution of highly overlapping peaks and facilitate the assignment of peaks

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belonging to each complex, which has successfully been applied to the adsorption

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systems of small organic compounds on goethite.14, 29, 40 The dynamic spectra in the

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region 1300−900 cm−1 from 10 to 60 min were baseline-corrected and smoothed to

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calculate the synchronous and asynchronous plots with the 2D Shige program (Shigeaki

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Morita, Kwansei-Gakuin University, 2004−2005). The averaged spectrum was set as a 9



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reference.29, 41 In the synchronous spectra, an auto peak is responsible for the changes

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of peak intensity over time, while a cross peak provides the response to the time

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perturbation at two different bands. A positive cross peak arises when the two bands

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increase or decrease simultaneously, whereas a negative cross peak demonstrates the

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opposite change in their peak intensities. The asynchronous spectra do not have auto

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peaks. An asynchronous cross peak indicates the uncorrelated response of two bands,

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which originate from different surface complexes or moieties of the same complex in

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different molecular environments.29

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RESULTS AND DISCUSSION

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1. ATR Spectra and DFT Calculations of Dissolved Aspartate. Aspartic acid

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contains one NH3+ group with pKa3 = 9.8 and two COOH groups with pKa1 = 2.1 and

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pKa2 = 3.9.5,

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monoanionic, and dianionic species, respectively (Figure 1 and Figure S1 in the

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Supporting Information, SI). On the basis of the second derivative, more frequencies

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were observed within the 1700−1500 cm−1 range than in previous studies. Curve-fitting

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results of the overlapping peaks are of high quality with r2 more than 0.99. The peak

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assignments are on the basis of DFT calculations (Table S1). Our correlation analysis

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shows a reasonable agreement between the theoretical frequencies and the experimental

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data (SI Tables S2−S3 and Figure S2).

6

The solution spectra at pH 3, 6, and 11 represent the zwitterionic,

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Figure 1. (A) Optimized structures of the dissolved aspartate. Eleven explicit H2O molecules are not

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At pH 3, aspartate is in the zwitterionic form with protonated γ-COOH and NH3+.

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The γ-COOH vibrations contribute to a band at 1724 cm−1 associated with ν(C=O),5, 42

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a band at 1243 cm−1 associated with coupled ν(C−O) and δ(O−H),42 and a band at 1351

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cm−1 associated with νs(γ-COO−). The δas(NH3+), δs(NH3+),42 and ρr(NH3+) modes are

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observed at 1637, 1511, and 1272 cm−1, respectively. The two bands at 16015, 42 and

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1396 cm−1 are attributed to the asymmetric and symmetric ν(α-COO−), respectively.

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The δ(CH2) scissor band appears at 1416 cm−1, and the ρ(CH2) wag appears near 1213

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cm−1. The 1305 cm−1 peak is attributed to δ(CH).42

shown for clarity. (B) Curve-fitting analysis of the solution spectra for 0.1 M aspartate in 0.1 M NaCl. The spectra were normalized to the most intense peak.

The solution spectra exhibit significant changes with increasing pH. Deprotonation

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11



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of γ-COOH at pH 6 increases the symmetry in the molecular structure, which results in

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downward shifts of νas(α-COO−) from 1601 to 1572 cm−1 and δs(NH3+) from 1511 to

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1476 cm−1. Furthermore, the νas(γ-COO−) mode begins to appear at 1596 cm−1.

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Interestingly, the δas(NH3+), νs(α-COO−), and νs(γ-COO−) vibrations are affected less

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by COOH deprotonation. The two bands at 1417 and 1308 cm−1 are assigned to CH2

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scissoring and wagging, respectively. The band at 1327 cm−1 is due to δ(CH), and the

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1233 cm−1 peak is due to NH3+ rocking.

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For the aspartate dianion at pH 11, the NH3+ deprotonation further increases the

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similarity between α- and γ-COO−. Accordingly, their asymmetric stretches shift

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downward to 1550 and 1568 cm−1, respectively. NH2 scissoring occurs at 1598 cm−1,

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and the rocking mode is centered at 1325 cm−1. The vibrations associated with νs(α-

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COO−) at 1395 cm−1, νs(γ-COO−) at 1359 cm−1, δ(CH2) scissoring at 1419 cm−1, and

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ρω(CH2) at 1307 cm−1 are not sensitive to NH3+ deprotonation.

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2. ATR Spectra of Adsorbed Aspartate at the Goethite/Aqueous Interface. The

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interfacial spectra with 1 mM aspartate in 0.1 M NaCl exhibit different characteristics

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in peak numbers, locations, and profiles at pH 3, 6, and 11, which suggests disparate

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complexation mechanisms at different pH (Figure 2). Notably, the interfacial aspartate

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at a certain pH remains unchanged during adsorption, as indicated by the same peak

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positions in the dynamic spectra (SI Figure S3).

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Figure 2. Curve-fitting analysis of the interfacial spectra collected at 20 min for 1 mM aspartate in 0.1

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Compared with the spectrum of aspartate zwitterion, the interfacial spectrum at pH

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3 exhibits significant changes in peak locations and shapes, which indicates inner-

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sphere coordination.31 The absence of γ-COOH vibrations at 1724 and 1243 cm−1

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suggests the involvement of distal carboxyl upon adsorption. The surface complex with

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deprotonated COO− has also been proposed for glutamate and carboxylic acids on metal

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oxides.3, 5, 6, 43, 44 The dianionic-type configurations could be excluded because of the

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persistent existence of δs(NH3+) at 1518 cm−1 after adsorption. The nonparticipation of

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the NH3+ group in surface complexation is supported by most previous studies of amino

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acid adsorption.3,

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mostly positive sites of goethite at pH 3 (below pHpzc 8.9). The strong adsorption of

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aspartate on goethite is primarily attributed to the γ-COOH group because little

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adsorption on metal oxides is observed for amino acids with neutrally charged side

M NaCl. The spectra were normalized to the most intense peak.

5, 6

The NH3+ group experiences electrostatic repulsion from the

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chains or for carboxylic acids with a single COOH and no other participating groups.3,

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6, 17

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could be ambiguous when characterized on the sole basis of ATR-FTIR analysis,

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thereby necessitating in-depth DFT calculations for theoretical confirmation. Notably,

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no apparent outer-sphere adsorption is detected with 1 mM aspartate in 0.1 M NaCl at

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pH 3. One reason for this lack of adsorption is that the protonated γ-COOH is not

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electrostatically attracted to the positively charged goethite surface. Another reason is

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that the deprotonated α-COO− and adjacent NH3+ can preclude electrostatic binding to

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the mineral surface.6

However, further differentiation between the bidentate and monodentate structures

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The peak numbers and positions of the interfacial spectrum at pH 6 resemble those

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of the solution spectrum of aspartate monoanion, except for the two additional bands at

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1517 and 1285 cm−1. The slight shifts of main peaks are due to the smaller distortions

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of aspartate upon adsorption. The nearly unchanged COO− stretches (1600, 1564, 1389,

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and 1356 cm−1) suggest the weak outer-sphere coordination5, 31, 43 via electrostatic

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interactions that occur between the monoanionic aspartate and the positive sites of

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goethite. The goethite surface herein still has a net amount of positive charge (below

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pHpzc), though the elevated pH reduces the relative concentration of positive to neutral

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to negative sites. The significant adsorption is closely associated with the electrostatic

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attraction via aspartate γ-COO−, which facilitates the further attachment of α-COO−, as

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evidenced by distinct variations in the relative intensities of νs(α-COO−) at 1389 cm−1.

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The charge distribution of the monoanionic aspartate could be rearranged when it is

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attracted to the goethite surface as a result of the different molecular environment from 14


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the solution phase. Notably, the bands at ~1517 and 1285 cm−1 were also detected in

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the interfacial aspartate at pH 3, which indicates the same origin of inner-sphere

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coordination. The dominant outer-sphere complexation is further confirmed by the

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dramatically decreased peak intensities along with the flowing of back ground

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electrolyte (SI Figure S4).

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The weak absorbance in the poor spectrum at pH 11 indicates unfavorable adsorption

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capacity, which has also been observed for aspartate on TiO2 and Al2O3 under strong

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basic conditions.3, 6, 10 The bands at 1600−1300 cm−1 remain nearly unchanged upon

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adsorption, which rules out inner-sphere complexation. The electrostatic outer-sphere

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coordination is also impossible due to the repulsion between the COO− groups in

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aspartate dianion and goethite surface with net negative charge (above pHpzc 8.9). The

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slight adsorption can be attributed to the H-bonding between aspartate NH2 and goethite

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surface hydroxyl groups. Furthermore, these H-bonds are weakened by the electrostatic

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repulsion from aspartate COO−. The weak adsorption and strong interference from

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H−O−H bending of water result in the poor-quality spectra, especially in the range of

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1650−1500 cm−1 (SI Figure S3). Consequently, we focused on investigating aspartate

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adsorption at pH 3 and 6 in this study.

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3. Effects of Ionic Strength and Aspartate Concentration on Adsorption. The

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interfacial spectra with 1 mM aspartate in Figure 3 exhibit sensitive responses to the

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decreased ionic strength (0.01 M NaCl) at pH 3 and 6 (SI Figure S5). The spectra at pH

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3 in 0.01 and 0.2 M NaCl both resolve the same inner-sphere complex as that in 0.1 M

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NaCl, as demonstrated by the main peak positions in both spectra being identical to 15



301

those of the spectrum in 0.1 M NaCl solution. The increased NaCl concentration (0.2

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M) has no detectable effect on the peak locations and profiles of the adsorbed spectrum.

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In contrast, the overlapping peaks from 1700 to 1550 cm−1 become broader at the lower

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ionic strength (0.01 M NaCl) because of the appearance of two bands at 1598 and 1567

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cm−1. Notably, these two peaks are similar to those of the monoanionic aspartate, which

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suggests electrostatic outer-sphere coordination between the positively charged

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goethite and the small amount of aspartate monoanion (SI Figure S1).

308 309 310

Figure 3. Curve-fitting analysis of the interfacial spectra collected at 20 min with 1 mM aspartate in 0.01

311

At pH 6, aspartate adsorption is governed by the dominant electrostatic outer-sphere

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complexation and slight inner-sphere coordination. The peak locations remain nearly

and 0.2 M NaCl. The spectra were normalized to the most intense peak.

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the same at different NaCl concentrations. Interestingly, the contribution of the ~1518

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cm−1 peak to the interfacial spectrum increases along with elevated ionic strength,

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whereas the other frequencies are much more insensitive. The peak-area ratio can be

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used to reflect the relative adsorption capacity between outer-sphere and inner-sphere

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coordination.29 The ratios in the peak area of ~1484 cm−1 (outer-sphere complex) to

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~1518 cm−1 (inner-sphere complex) are 4.23, 2.22, and 1.54 for 0.01, 0.1, and 0.2 M

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NaCl, respectively. At higher ionic strengths, the reduced contribution of outer-sphere

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complexation to the total aspartate adsorption is mainly due to the weakened

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electrostatic attraction between the dissolved aspartate and the goethite surface. The

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diffuse layer thickness decreases with increasing electrolyte ion density.45

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In this study, we also investigated the effect of aspartate concentration on the

324

interfacial structures and their relative contributions (SI Figure S6). The peak locations

325

and profiles of the spectra for 5 and 10 mM aspartate at pH 3 resemble those for 1 mM

326

aspartate, which indicates the same inner-sphere complexation (Figures 2 and 4).

327

Furthermore, no evidence of an additional surface complex (appearance of new bands)

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is detected. At pH 6, the dominant outer-sphere complexation and unfavorable inner-

329

sphere coordination are also resolved in the spectra for 5 and 10 mM aspartate because

330

the frequencies are identical to those for 1 mM aspartate. Interestingly, the elevated

331

aspartate concentration slightly enhances the inner-sphere contribution to the overall

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adsorption. The peak-area ratios between ~1518 and ~1484 cm−1 are 0.45, 0.65, and

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0.72 for 1, 5, and 10 mM aspartate, respectively.

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334 335 336

Figure 4. Curve-fitting analysis of the interfacial spectra collected at 20 min with 5 and 10 mM aspartate

337

4. DFT Calculations of Adsorbed Aspartate on Goethite. Seven possible

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interfacial configurations, including two mononuclear monodentate (M-M), two

339

mononuclear bidentate (M-B), and three binuclear bidentate (B-B) structures, were

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calculated to confirm the proposed inner-sphere complex (Figure 5). The M-M, M-B,

341

B-B I, and B-B II structures coordinate to the surface iron via single aspartate α- or γ-

342

COO−. The B-B III complex binds to two Fe atoms while involving both α- and γ-COO−.

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As suggested by Ha and Hwang et al., the optimized configuration can be determined

344

by the agreement between the experimental and theoretical vibrations.31, 44 The results

345

indicate that the B-B III structure fits much better with the observed inner-sphere

346

complex than the other configurations (SI Tables S4−S6 and Figures S7−S8). Peak

in 0.1 M NaCl. The spectra were normalized to the most intense peak.

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assignments of the B-B III complex are on the basis of DFT calculations. The δas(NH3+)

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and δs(NH3+) deformations contribute to the bands at 1649 and 1518 cm−1, respectively.

349

The νas(α-COO−) vibrations occur at 1608 cm−1, and the νas(γ-COO−) vibrations occur

350

near 1578 cm−1. The two peaks at 1421 and 1394 cm−1 are both ascribed to the mixed

351

symmetric stretches of α- and γ-COO−. The bands at 1448 and 1310 cm−1 correspond

352

to CH2 scissoring and wagging, respectively. The band at 1349 cm−1 is caused by δ(CH),

353

and the 1285 cm−1 peak is caused by ρr(NH3+). The DFT calculations of outer-sphere

354

coordination were also considered in this study. Our results accurately simulate the

355

observed frequencies of electrostatic and H-bonded outer-sphere complexes (SI Tables

356

S5−S6 and Figure S8). Because of the slight distortions upon adsorption, peak

357

assignments of the two complexes are consistent with those of the corresponding

358

dissolved species.

359 360 361

Figure 5. Optimized configurations of interfacial aspartate and phosphate in mononuclear monodentate (M-M), mononuclear bidentate (M-B), and binuclear bidentate (B-B) structures. Explicit H2O molecules 19



362

are not shown for clarity.

363

5. Phosphate Only Adsorption on Goethite. The interfacial spectra with 1 mM

364

phosphate are of the same peak numbers and locations at pH 3 and 6, which

365

demonstrates the identical surface complexes (Figure 6 and SI Figure S9). The dramatic

366

shape differences at these two pH indicate the uneven content of each configuration,

367

which merits further analysis on the overlapping peaks with 2D-COS. As detailedly

368

discussed in SI (Figures S10−S11 and Table S7), three frequency groups are resolved

369

at both pH 3 and 6, namely (A) ~1253, 1222, 1128, and 1008 cm−1; (B) ~1172, 1073,

370

and 969 cm−1; and (C) ~1096, 1046, and 939 cm−1. In comparison with those of the

371

dissolved H2PO4− and HPO42− (SI Figures S12−S13), the interfacial spectra experience

372

distinct peak shifts upon adsorption, which indicates the inner-sphere coordination.31

373

Our DFT results suggest that the combination of diprotonated M-M (group A),

374

monoprotonated B-B (group B), and nonprotonated B-B (group C) structures can well

375

describe the interfacial phosphate on goethite (SI Tables S8−S10 and Figures S14−S15).

376

The optimized configurations are different from those reported in previous studies,

377

which provide complementary insight into the inconsistent mechanisms of phosphate

378

adsorption (SI Table S11).20-23 For example, Luengo proposed that the adsorbed

379

phosphate on goethite were in the monoprotonated and nonprotonated bidentate forms,

380

whereas Kubicki supported the monodentate and bidentate HPO42− complexes as well

381

as the outer-sphere coordination.21, 23

20


Page 20 of 32

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382 383 384 385 386


Figure 6. (A–D) Curve-fitting analysis of the interfacial spectra collected at 20 min with 1 mM phosphate in single- and binary-adsorbate systems (I = 0.1 M NaCl). (E) The interfacial spectra collected at 10 min with 1 mM aspartate in competition with 1 mM phosphate in 0.1 M NaCl. The spectra were normalized to the most intense peak.

387

The significant spectral changes motivate our further research on the effect of pH on

388

the relative contents of each complex. The peak-area ratios of the bands at about 1008,

389

1073, and 1046 cm–1 are 5.69:1:1.06 and 4.13:1:4.24 at pH 3 and 6, respectively. The

390

nonprotonated B-B complexation (1046 cm–1) becomes more important at pH 6. The

391

P–OH deprotonation is easier at high pH, which facilitates the binding of O atom in P–

392

OH to the surface iron. The relative contribution of diprotonated M-M structure (1008

393

cm–1) decreases markedly with the increasing pH. However, it is still considerable at

394

pH 6 where H2PO4− is not the dominant species in solution phase (90% at pH 3 and 10%

395

at pH 6, SI Figure S13). In one, more phosphate can be adsorbed on the limited goethite 21



396

surface sites via M-M coordination. In another, the diprotonated complex is stabilized

397

by the monoprotonated and nonprotonated B-B structures via H-bonding.

398

6. Competitive Adsorption between Aspartate and Phosphate. When competition

399

with each other, the structures of adsorbed aspartate and phosphate are identical to those

400

in their single-adsorbate systems (Figure 6). Within the range of 1300−900 cm−1, the

401

interfacial spectra with 1 mM aspartate and phosphate are of exactly similar peak

402

numbers and locations to those of the individual phosphate adsorption on goethite.

403

Interestingly, poor spectral quality is observed in the 1800−1300 cm−1 region, which in

404

general resembles the features of aspartate only adsorption at the same pH.

405

As for aspartate, the rather low absorbance indicates the unfavorable adsorption as a

406

result of phosphate competition. Besides the dramatically reduced capacity in the initial

407

adsorption phase, the weaker affinity of aspartate on goethite is demonstrated by the

408

decreasing peak intensities along with the reaction time (SI Figure S16). Notably, the

409

influencing mechanisms of phosphate competition are different at pH 3 and 6. At pH 3,

410

most of the goethite surface sites are occupied by phosphate, and the adsorbed aspartate

411

B-B III complex is partly removed by the phosphate inner-sphere coordination. The

412

slight adsorption at pH 6 is attributed to the weakened electrostatic attractions between

413

aspartate monoanion and goethite surface. The net positive charge of goethite surface

414

decreases significantly with the attachment of phosphate (H2PO4− and HPO42−). The

415

more phosphate adsorption, the less aspartate outer-sphere complexation.

416

In contrast to the partial removal of adsorbed aspartate in the initial phase, phosphate

417

experiences fast and considerable adsorption within 30 min (SI Figure S17). Because 22


Page 22 of 32

Page 23 of 32


418

of the much stronger adsorption affinity, the interfacial phosphate on goethite is hardly

419

affected by the presence of aspartate (Figure 6). Moreover, no significant variations are

420

observed from the contributions of each complex to the overall phosphate adsorption

421

when competing with aspartate. The peak-area ratios of the ~1008, 1073, and 1046 cm–

422

1

peaks herein are 5.41:1:1.11 and 4.18:1:4.39 at pH 3 and 6, respectively.

423

7. Environmental Significance. The molecular-level adsorption mechanisms of

424

aspartate and phosphate on goethite were investigated both in their individual and

425

competition systems. As for aspartate only adsorption, we resolve one inner-sphere and

426

two outer-sphere complexes that differ from those on Al2O3 and TiO2, which further

427

demonstrates the critical influence of mineral phase on aspartate adsorption.3, 5 The B-

428

B III inner-sphere coordination and electrostatic attraction result in favorable

429

adsorption under acidic and near-neutral conditions. The markedly decreased

430

adsorption at pH 11 is mainly due to the weakened H-bonding (by electrostatic

431

repulsion) and the small number of available adsorption sites. The adsorption of amino

432

acids on minerals can reduce their biological degradation rates in the environment,

433

which has important impact on the geochemical cycle of nitrogen.46 Moreover, the

434

adsorption modes significantly affect the stability of interfacial aspartate. In inner-

435

sphere and electrostatic outer-sphere coordination, the aspartate NH3+ remains

436

relatively free of surface complexation, which leaves it potentially subject to

437

degradation.47, 48

438

The phosphate competition dramatically changes the adsorption behaviors of

439

aspartate on goethite. The unfavorable aspartate adsorption and partial removal of the 23



440

adsorbed complex are attributed to the much weaker adsorption affinity than that of

441

phosphate. The abundant phosphate in natural waters can inhibit the aspartate

442

adsorption and accelerate its migration in the environment. On the contrary, the

443

phosphate adsorption is less affected by the presence of aspartate, including the

444

interfacial configurations (diprotonated M-M, monoprotonated B-B, and nonprotonated

445

B-B) and the relative contributions of each complex. The favorable adsorption

446

promotes the removal of bioavailable phosphate from natural solutions, i.e. the soil

447

solution, the sediment pore water, and the overlying water. These results have important

448

implications for understanding the fate and transport of amino acids and phosphate.

449

ASSOCIATED CONTENT

450

Supporting Information

451

Speciation distribution of dissolved aspartate and phosphate as a function of pH.

452

Dynamic spectra of aspartate and phosphate on goethite in single-adsorbate systems.

453

DFT calculations of dissolved aspartate in solution and interfacial configurations on

454

goethite. 2D correlation analysis of phosphate only adsorption. ATR spectra of

455

dissolved phosphate and DFT calculations of interfacial phosphate on goethite.

456

Interfacial phosphate on goethite reported by previous studies. Dynamic spectra of

457

competitive adsorption between aspartate and phosphate. This material is available free

458

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

459

AUTHOR INFORMATION

460

Corresponding Author 24


Page 24 of 32

Page 25 of 32


461

Shengrui Wang*

462

Address: Chinese Research Academy of Environmental Sciences, 8 Dayangfang,

463

Beiyuan, Beijing 100012, China

464

Tel.: +86-10-84915277. Fax: +86-10-84915190. E-mail: [email protected].

465

Notes

466

The authors declare no competing financial interest.

467

ACKNOWLEDGEMENTS

468

This research was supported by the National Natural Science Foundation of China

469

(U1202235), the National High-Level Talents Special Support Plan (for Science and

470

Technology Innovation Talents to Special Support Plan), the National Key Science and

471

Technology Special Program “Water Pollution Control and Treatment” (2012ZX07102-

472

004), and the China Postdoctoral Science Foundation (2014M561024).

473

REFERENCES

474

(1) Lee, N.; Sverjensky, D. A.; Hazen, R. M. Cooperative and competitive adsorption

475

of amino acids with Ca2+ on rutile (α-TiO2). Environ. Sci. Technol. 2014, 48 (16),

476

9358−9365.

477

(2) Sverjensky, D. A.; Jonsson, C. M.; Jonsson, C. L.; Cleaves, H. J.; Hazen, R. M.

478

Glutamate surface speciation on amorphous titanium dioxide and hydrous ferric oxide.

479

Environ. Sci. Technol. 2008, 42 (16), 6034−6039.

480

(3) Greiner, E.; Kumar, K.; Sumit, M.; Giuffre, A.; Zhao, W.; Pedersen, J.; Sahai, N.

481

Adsorption of L-glutamic acid and L-aspartic acid to γ-Al2O3. Geochim. Cosmochim. 25



482

Acta 2014, 133, 142−155.

483

(4) Chu, W.-H.; Gao, N.-Y.; Deng, Y.; Dong, B.-Z. Formation of chloroform during

484

chlorination of alanine in drinking water. Chemosphere 2009, 77 (10), 1346−1351.

485

(5) Parikh, S. J.; Kubicki, J. D.; Jonsson, C. M.; Jonsson, C. L.; Hazen, R. M.;

486

Sverjensky, D. A.; Sparks, D. L. Evaluating glutamate and aspartate binding

487

mechanisms to rutile (α-TiO2) via ATR-FTIR spectroscopy and quantum chemical

488

calculations. Langmuir 2011, 27 (5), 1778−1787.

489

(6) Roddick-Lanzilotta, A. D.; McQuillan, A. J. An in situ infrared spectroscopic study

490

of glutamic acid and of aspartic acid adsorbed on TiO2: Implications for the

491

biocompatibility of titanium. J. Colloid Interface Sci. 2000, 227 (1), 48−54.

492

(7) Pászti, Z.; Guczi, L. Amino acid adsorption on hydrophilic TiO2: A sum frequency

493

generation vibrational spectroscopy study. Vib. Spectrosc 2009, 50 (1), 48−56.

494

(8) Giacomelli, C. E.; Avena, M. J.; De Pauli, C. P. Aspartic acid adsorption onto TiO2

495

particles surface. Experimental data and model calculations. Langmuir 1995, 11 (9),

496

3483−3490.

497

(9) Guo, Y.-n.; Lu, X.; Zhang, H.-p.; Weng, J.; Watari, F.; Leng, Y. DFT Study of the

498

adsorption of aspartic acid on pure, N-doped, and Ca-doped rutile (110) surfaces. J.

499

Phys. Chem. C 2011, 115 (38), 18572−18581.

500

(10) Jonsson, C. M.; Jonsson, C. L.; Estrada, C.; Sverjensky, D. A.; Cleaves Ii, H. J.;

501

Hazen, R. M. Adsorption of L-aspartate to rutile (α-TiO2): Experimental and theoretical

502

surface complexation studies. Geochim. Cosmochim. Acta 2010, 74 (8), 2356−2367.

503

(11) Ikhsan, J.; Johnson, B. B.; Wells, J. D.; Angove, M. J. Adsorption of aspartic acid 26


Page 26 of 32

Page 27 of 32


504

on kaolinite. J. Colloid Interface Sci. 2004, 273 (1), 1−5.

505

(12) Naidja, A.; Huang, P. M. Aspartic acid interaction with Ca-montmorillonite:

506

adsorption, desorption and thermal stability. Appl. Clay Sci. 1994, 9 (4), 265−281.

507

(13) Hanna, K.; Martin, S.; Quilès, F.; Boily, J. F. Sorption of phthalic acid at goethite

508

surfaces under flow-through conditions. Langmuir 2014, 30 (23), 6800−6807.

509

(14) Lindegren, M.; Loring, J. S.; Persson, P. Molecular structures of citrate and

510

tricarballylate adsorbed on α-FeOOH particles in aqueous suspensions. Langmuir 2009,

511

25 (18), 10639−10647.

512

(15) Hongshao, Z.; Stanforth, R. Competitive adsorption of phosphate and arsenate on

513

goethite. Environ. Sci. Technol. 2001, 35 (24), 4753−4757.

514

(16) Villalobos, M.; Trotz, M. A.; Leckie, J. O. Surface complexation modeling of

515

carbonate effects on the adsorption of Cr(VI), Pb(II), and U(VI) on goethite. Environ.

516

Sci. Technol. 2001, 35 (19), 3849−3856.

517

(17) Norén, K.; Loring, J. S.; Persson, P. Adsorption of alpha amino acids at the

518

water/goethite interface. J. Colloid Interface Sci. 2008, 319 (2), 416−428.

519

(18) Yang, Y.; Duan, J.; Jing, C. Molecular-scale study of salicylate adsorption and

520

competition with catechol at goethite/aqueous solution interface. J. Phys. Chem. C 2013,

521

117 (20), 10597−10606.

522

(19) Estrada, C. F.; Sverjensky, D. A.; Pelletier, M.; Razafitianamaharavo, A.; Hazen,

523

R. M. Interaction between L-aspartate and the brucite [Mg(OH)2]–water interface.

524

Geochim. Cosmochim. Acta 2015, 155, 172−186.

525

(20) Tejedor-Tejedor, M. I.; Anderson, M. A. The protonation of phosphate on the 27



526

surface of goethite as studied by CIR-FTIR and electrophoretic mobility. Langmuir

527

1990, 6 (3), 602−611.

528

(21) Luengo, C.; Brigante, M.; Antelo, J.; Avena, M. Kinetics of phosphate adsorption

529

on goethite: Comparing batch adsorption and ATR-IR measurements. J. Colloid

530

Interface Sci. 2006, 300 (2), 511−518.

531

(22) Rahnemaie, R.; Hiemstra, T.; van Riemsdijk, W. H. Geometry, charge distribution,

532

and surface speciation of phosphate on goethite. Langmuir 2007, 23 (7), 3680−3689.

533

(23) Kubicki, J. D.; Paul, K. W.; Kabalan, L.; Zhu, Q.; Mrozik, M. K.; Aryanpour, M.;

534

Pierre-Louis, A.-M.; Strongin, D. R. ATR-FTIR and density functional theory study of

535

the structures, energetics, and vibrational spectra of phosphate adsorbed onto goethite.

536

Langmuir 2012, 28 (41), 14573−14587.

537

(24) Zeng, H.; Fisher, B.; Giammar, D. E. Individual and competitive adsorption of

538

arsenate and phosphate to a high-surface-area iron oxide-based sorbent. Environ. Sci.

539

Technol. 2008, 42 (1), 147−152.

540

(25) Cheng, T.; Barnett, M. O.; Roden, E. E.; Zhuang, J. Effects of phosphate on

541

uranium(VI) adsorption to goethite-coated sand. Environ. Sci. Technol. 2004, 38 (22),

542

6059−6065.

543

(26) Wang, X.-H.; Liu, F.-F.; Lu, L.; Yang, S.; Zhao, Y.; Sun, L.-B.; Wang, S.-G.

544

Individual and competitive adsorption of Cr(VI) and phosphate onto synthetic Fe–Al

545

hydroxides. Colloids Surf., A 2013, 423, 42−49.

546

(27) Hiemstra, T.; Mia, S.; Duhaut, P.-B.; Molleman, B. Natural and pyrogenic humic

547

acids at goethite and natural oxide surfaces interacting with phosphate. Environ. Sci. 28


Page 28 of 32

Page 29 of 32


548

Technol. 2013, 47 (16), 9182−9189.

549

(28) Qin, X.; Liu, F.; Wang, G.; Li, L.; Wang, Y.; Weng, L. Modeling of levofloxacin

550

adsorption to goethite and the competition with phosphate. Chemosphere 2014, 111,

551

283−290.

552

(29) Yang, Y.; Yan, W.; Jing, C. Dynamic adsorption of catechol at the goethite/aqueous

553

solution interface: A molecular-scale study. Langmuir 2012, 28 (41), 14588−14597.

554

(30) Ding, X.; Song, X.; Boily, J.-F. Identification of fluoride and phosphate binding

555

sites at FeOOH surfaces. J. Phys. Chem. C 2012, 116 (41), 21939−21947.

556

(31) Ha, J.; Hyun Yoon, T.; Wang, Y.; Musgrave, C. B.; Brown, J. G. E. Adsorption of

557

organic matter at mineral/water interfaces: 7. ATR-FTIR and quantum chemical study

558

of lactate interactions with hematite nanoparticles. Langmuir 2008, 24 (13), 6683−6692.

559

(32) Foerstendorf, H.; Jordan, N.; Heim, K. Probing the surface speciation of uranium

560

(VI) on iron (hydr)oxides by in situ ATR FT-IR spectroscopy. J. Colloid Interface Sci.

561

2014, 416, 133−138.

562

(33) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.;

563

Cheeseman, J. R.; Montgomery, Jr., J. A.; Vreven, T.; Kudin, K. N.; Burant, J. C.;

564

Millam, J. M.; Iyengar, S. S.; Tomasi, J.; Barone, V.; Mennucci, B.; Cossi, M.; Scalmani,

565

G.; Rega, N.; Petersson, G. A.; Nakatsuji, H.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda,

566

R.; Hasegawa, J.; Ishida, M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Klene, M.;

567

Li, X.; Knox, J. E.; Hratchian, H. P.; Cross, J. B.; Bakken, V.; Adamo, C.; Jaramillo, J.;

568

Gomperts, R.; Stratmann, R. E.; Yazyev, O.; Austin, A. J.; Cammi, R.; Pomelli, C.;

569

Ochterski, J. W.; Ayala, P. Y.; Morokuma, K.; Voth, G. A.; Salvador, P.; Dannenberg, J. 29



570

J.; Zakrzewski, V. G.; Dapprich, S.; Daniels, A. D.; Strain, M. C.; Farkas, O.; Malick,

571

D. K.; Rabuck, A. D.; Raghavachari, K.; Foresman, J. B.; Ortiz, J. V.; Cui, Q.; Baboul,

572

A. G.; Clifford, S.; Cioslowski, J.; Stefanov, B. B.; Liu, G.; Liashenko, A.; Piskorz, P.;

573

Komaromi, I.; Martin, R. L.; Fox, D. J.; Keith, T.; Al-Laham, M. A.; Peng, C. Y.;

574

Nanayakkara, A.; Challacombe, M.; Gill, P. M. W.; Johnson, B.; Chen, W.; Wong, M.

575

W.; Gonzalez, C.; and Pople, J. A. Gaussian 03, Revision E.01, Gaussian, Inc.,

576

Wallingford CT, 2004..

577

(34) Bhandari, N.; Hausner, D. B.; Kubicki, J. D.; Strongin, D. R. Photodissolution of

578

ferrihydrite in the presence of oxalic acid: An in situ ATR-FTIR/DFT study. Langmuir

579

2010, 26 (21), 16246−16253.

580

(35) Johnston, C. P.; Chrysochoou, M. Investigation of chromate coordination on

581

ferrihydrite by in situ ATR-FTIR spectroscopy and theoretical frequency calculations.

582

Environ. Sci. Technol. 2012, 46 (11), 5851−5858.

583

(36) Li, W.; Pierre-Louis, A.-M.; Kwon, K. D.; Kubicki, J. D.; Strongin, D. R.; Phillips,

584

B. L. Molecular level investigations of phosphate sorption on corundum (α-Al2O3) by

585

31

586

Cosmochim. Acta 2013, 107, 252−266.

587

(37) Pierre-Louis, A.-M.; Hausner, D. B.; Bhandari, N.; Li, W.; Kim, J.; Kubicki, J. D.;

588

Strongin, D. Adsorption of carbon dioxide on Al/Fe oxyhydroxide. J. Colloid Interface

589

Sci. 2013, 400, 1−10.

590

(38) Sanchez-de-Armas, R.; San-Miguel, M. A.; Oviedo, J.; Marquez, A.; Sanz, J. F.

591

Electronic structure and optical spectra of catechol on TiO2 nanoparticles from real time

P solid state NMR, ATR-FTIR and quantum chemical calculation. Geochim.

30


Page 30 of 32

Page 31 of 32


592

TD-DFT simulations. Phys. Chem. Chem. Phys. 2011, 13 (4), 1506−1514.

593

(39) NIST Computational Chemistry Comparison and Benchmark Database; NIST

594

Standard Reference Database Number 101; Johnson, R. D., III, Ed.; Release 16a,

595

August 2013, http://cccbdb.nist.gov/.

596

(40) Norén, K.; Persson, P. Adsorption of monocarboxylates at the water/goethite

597

interface: The importance of hydrogen bonding. Geochim. Cosmochim. Acta 2007, 71

598

(23), 5717−5730.

599

(41) Noda, I.; Ozaki, Y. Two-Dimensional Correlation Spectroscopy: Applications in

600

Vibrational and Optical Spectroscopy; John Wiley & Sons: Chichester, England, 2004.

601

(42) Pearson, J. F.; Slifkin, M. A. The infrared spectra of amino acids and dipeptides.

602

Spectrochim. Acta 1972, 28A (12), 2403−2417.

603

(43) Johnson, S. B.; Yoon, T. H.; Kocar, B. D.; Brown, G. E. Adsorption of organic

604

matter at mineral/water interfaces. 2. Outer-sphere adsorption of maleate and

605

implications for dissolution processes. Langmuir 2004, 20 (12), 4996−5006.

606

(44) Hwang, Y. S.; Liu, J.; Lenhart, J. J.; Hadad, C. M. Surface complexes of phthalic

607

acid at the hematite/water interface. J. Colloid Interface Sci. 2007, 307 (1), 124−134.

608

(45) Gulley-Stahl, H.; Hogan, P. A.; Schmidt, W. L.; Wall, S. J.; Buhrlage, A.; Bullen,

609

H. A. Surface complexation of catechol to metal oxides: An ATR-FTIR, adsorption, and

610

dissolution study. Environ. Sci. Technol. 2010, 44 (11), 4116−4121.

611

(46) Kitadai, N.; Yokoyama, T.; Nakashima, S. In situ ATR-IR investigation of L-lysine

612

adsorption on montmorillonite. J. Colloid Interface Sci. 2009, 338 (2), 395−401.

613

(47) Sheals, J.; Sjöberg, S.; Persson, P. Adsorption of glyphosate on goethite: 31



614

Molecular characterization of surface complexes. Environ. Sci. Technol. 2002, 36 (14),

615

3090−3095.

616

(48) Jonsson, C. M.; Persson, P.; Sjöberg, S.; Loring, J. S. Adsorption of glyphosate on

617

goethite (α-FeOOH): Surface complexation modeling combining spectroscopic and

618

adsorption data. Environ. Sci. Technol. 2008, 42 (7), 2464−2469.

32


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Molecular-Scale Study of Aspartate Adsorption on Goethite and

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