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Article Cite This: Inorg. Chem. XXXX, XXX, XXX−XXX

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Complexation of Uranium(VI) with N‑(2Hydroxyethyl)ethylenediamine‑N,N′,N′‑triacetic Acid in Aqueous Solution: Thermodynamic Studies and Coordination Analyses Xingliang Li,†,‡ Zhicheng Zhang,*,† Leigh R. Martin,§,∥ Shunzhong Luo,‡ and Linfeng Rao*,† †

Chemical Sciences Division, Lawrence Berkeley National Laboratory, Berkeley, California 94720, United States Institute of Nuclear Physics and Chemistry, China Academy of Engineering Physics, Mianyang, Sichuan 621999, China § Aqueous Separations and Radiochemistry Department, Idaho National Laboratory, PO Box 1625, Idaho Falls, Idaho 83415, United States

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‡

S Supporting Information *

ABSTRACT: N-(2-Hydroxyethyl)ethylenediamine-N,N′,N′-triacetic acid (HEDTA, denoted as H3L in this work, and the three dissociable protons represent those of the three carboxylic groups) is a strong chelating ligand and plays an important role in the treatment and disposal of nuclear wastes as well as separation sciences of f-elements. In this work, the complexation of HEDTA with U(VI) was studied thermodynamically and structurally in aqueous solutions. Potentiometry and microcalorimetry were used to measure the complexation constants (298−343 K) and enthalpies (298 K), respectively, at I = 1.0 mol·L−1 NaClO4. Thermodynamic studies identified three 1:1 U(VI)/HEDTA complexes with different degrees of deprotonation, namely, UO2(HL)(aq), UO2L−, and UO2(H−1L)2−, where H−1 represents the deprotonation of the hydroxyl group. The results indicated that all three complexation reactions are endothermic and driven by entropy only. Coordination modes of the three complexes were investigated by NMR and extended X-ray absorption fine structure spectroscopies. In the UO2(HL)(aq) complex, HEDTA holds a tridentate mode, and the coordination occurs to the end of the ethylenediamine backbone. Two oxygens of the two carboxylic groups and one nitrogen of the amine group participate in the coordination. In both UO2L− and UO2(H−1L)2−, HEDTA holds a tetradentate mode and coordinates to U(VI) along the side of the ethylenediamine backbone. The difference is that in the UO2(H−1L)2− complex, the alkoxide form of the HEDTA hydroxyl group directly binds to the U(VI) atom, forming a highly strong chelation.

■

INTRODUCTION N-(2-Hydroxyethyl)ethylenediamine-N,N′,N′-triacetic acid, denoted as HEDTA (Figure 1), is an amino-poly(carboxylic

knowledge of the coordination chemistry of U(VI) with amino-polycarboxylate ligands. A massive campaign has been and will continue to be performed in the United States to treat and finally dispose of the legacy nuclear wastes generated from the nuclear weapon production program. An understanding of the chemical behavior of actinides in the waste treatment and disposal is essential for safely and efficiently managing the wastes. The complexation of actinides with inorganic and organic ligands, which exist in the wastes, chemical treatment processes, and the disposal environments, directly affects the chemical behavior of actinides in the waste streams. And thus, studies of the actinide complexation with those ligands have received increasingly more attention in the past few decades. HEDTA was intentionally added to the chemical separation processes during the nuclear weapon production campaign. As

Figure 1. Structure of HEDTA.

acid) and holds strong chelation ability to metal cations due to its multiple coordination sites. The motivation of this work is twofold: (1) gaining thermodynamic data of U(VI) complexation with HEDTA at varied temperatures to help understand the chemical behavior of U(VI) in the nuclear waste treatment and disposal, and (2) investigation into the coordination modes of the U(VI)/HEDTA complexes to enhance our © XXXX American Chemical Society

Received: March 14, 2018

A

DOI: 10.1021/acs.inorgchem.8b00655 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry

+ H2O = (UO2)2(OH)L+ + H+, respectively. In the other work by Lai et al.,6 the U(VI)/HEDTA complexes were studied with a polarographic method at t = 30 °C and I = 0.3 mol·L−1 NaClO4. At the pH range of 1.0−8.0, proposed were the presence of three complexes UO2(H2L)2(aq), UO2(HL)22−, and UO2L24−. The formation reactions and the estimated constants were reported as follows: UO22+ + 2(H2L)− = UO2(H2L)2(aq), log K = 5.57; UO22+ + 2(HL)2− = UO2(HL)22−, log K = (5.73−7.13); UO22+ + 2L3− = UO2L24−, log K = (8.61−9.81). Clearly, discrepancy exists in the complexation models between the two studies. The inconsistency not only raises the questions for the proposed complexation models but also leads to the suggestion that an understanding of coordination modes is critical to correctly define the complexation reactions for such a ligand with many donor atoms. In this work, we integrated thermodynamic measurements and structural analyses to investigate the complexation of U(VI) with HEDTA in aqueous solutions at varied temperatures. In the thermodynamic work, potentiometry was used to determine the stability constants of U(VI)/HEDTA complexes in 1.0 mol·L−1 NaClO4 at 298−343 K, and calorimetry used to measure the enthalpy of the complexation at 298 K. The coordination modes of the complexes were analyzed by two advanced structural analysis techniques, namely, NMR and extended X-ray absorption fine structure (EXAFS) spectroscopies.

a result, large quantities of HEDTA are present in the nuclear wastes. For example, the initial concentration of HEDTA in some high-level nuclear waste tanks at the Hanford site is as high as 0.2 mol·L−1, and the total amount of HEDTA from the Hanford mixed waste alone is ∼1500 tons.1−4 Therefore, an investigation of the complexation of U(VI) with HEDTA is necessary to help understand the effect of HEDTA on the chemical behavior of uranium in the waste treatment and disposal, as U(VI) is the most abundant actinide in the waste. To predict the chemical behavior of actinides and assess the safety of waste treatment and disposal, thermodynamic properties of the complexation are required. More importantly, the data at elevated temperatures are particularly needed, because the temperatures of the waste storage, chemical treatment processes, and the surrounding environment of the nuclear waste in the repository are substantially higher than the ambient temperature. Currently, the majority of thermodynamic data on the complexation of actinides are determined at or near 298 K. And thus, a substantial effort has been and will continue to be made to determine the thermodynamic data of complexation at the elevated temperatures, including stability constants, enthalpy, and entropy. Amino-poly(carboxylic acid)s are Lewis bases, containing both nitrogen and oxygen donor atoms, and generally form strong complexes with metal cations through a multidentate coordination structure. As is shown in Figure 1, the HEDTA molecule holds nine donor atoms, including two nitrogen atoms located on its amine backbone and seven oxygen atoms on its three acetic pendant arms and one ethanol arm. All of them potentially participate in coordination, forming various coordination modes. Uranyl, a linear dioxo U(VI) cation, only allows coordinating atoms to approach the U(VI) atom through its equatorial plane in the complexes. This hindrance could somehow place the limitation on the access of some HEDTA donor atoms to the U(VI) atom in the U(VI)/ HEDTA complexes and thus render the HEDTA ligand into a specific configuration to accommodate a coordination geometry. From the coordination chemistry point of view, an understanding of those coordination structures could enhance our knowledge of the actinyl coordination behavior with multidentate ligands. For the thermodynamic study purpose, those coordination details could also provide substantial help in verifying the complexation patterns and explaining the complex strength. Studies of U(VI) complexation with HEDTA are surprisingly scarce. Only two publications in the 1960s could be found from the literature.5,6 Interestingly, the observations presented in the two studies differ significantly, even though the studied pH ranges were much overlapped between them.5,6 In the work by Bhat et al.,5 the U(VI)/HEDTA complexes were investigated with combined potentiometry and spectrophotometry at t = 25 °C and I = 0.2 mol·L−1 NaClO4. At low pH (1.5−3.0), two complexes, namely, UO2(HL)(aq) and (UO2)2L+, where H3L represents the HEDTA ligand and the three protons represent those of three carboxylic groups, were identified. The formation constants were determined to be log K = 6.27−6.37 and 16.44−16.87 for the reactions UO22+ + HL2− = UO2(HL)(aq) and 2UO22+ + L3− = (UO2)2L+, respectively. At relatively high pH (4.0−6.0), Bhat et al. proposed that these two complexes underwent hydrolysis, and the hydrolysis constants were determined to be log KH = −(5.23−5.43) and −(9.81−10.09) for the hydrolysis reactions UO2(HL)(aq) + H2O = UO2(OH)(HL)− + H+ and (UO2)2L

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EXPERIMENTAL SECTION

Chemicals. All chemicals were reagent grade or higher. Boiled Milli-Q water (18.2 MΩ·cm at 298 K) was used for preparing solutions. The uranyl stock solution in perchloric acid was prepared as follows. Solid U3O8 was dissolved in 2 mol·L−1 HNO3 under low heating. The solution was filtered to remove any undissolved solid. U(VI) was precipitated as hydroxide with 2.0 mol·L−1 NH4OH. The precipitate was washed with water at pH 7−8 and then dissolved in 0.2 mol·L−1 HClO4. The concentrations of U(VI) and H+ in the stock solution were determined, respectively, by fluorimetry7 using standard solutions of U(VI) in 1.0 mol·L−1 H3PO4 and by Gran titration.8 The stock solution of HEDTA was prepared by dissolving a weighed amount of HEDTA (Sigma-Aldrich, >98%) in water and neutralizing it to the desired extent with 1.0 mol·L−1 NaOH solution (Metrohm). For the NMR experiments, the solution of HEDTA in D2O was prepared by dissolving a weighed amount of HEDTA in D2O (CIL, D: 99.9%) and neutralizing it to the desired extent with NaOD (CIL, NaOD: 40%, D: >99%). The ionic strength of all working solutions except those for the NMR measurements was maintained at 1.0 mol·L−1 NaClO4 (298 K). In this paper, all concentrations in the molarity unit are referred to 298 K. Potentiometry. Potentiometry was used to determine the stability constants of U(VI)/HEDTA complexes in a 1.0 mol·L−1 NaClO4 solution at the temperatures of 298, 313, 328, and 343 K. The titrations were conducted in a computer-controlled titration unit. Details of the unit have been described elsewhere.9,10 The titration unit is composed of a double-jacket glass titration vessel, a Metrohm dosimat (907 Titrando), and a pH meter equipped with a pH electrode. The temperature of the titration vessel was maintained by circulating water from a thermostat water bath. An inert atmosphere was maintained in the titration vessel by purging Ar gas into the vessel. The original electrode filling solution (3.0 mol·L−1 KCl) was replaced with 1.0 mol·L−1 NaCl to prevent the clogging of the electrode junction due to the low solubility of KClO4. The hydrogen ion concentration in a titration solution is potentiometrically obtained with the following equations: B

DOI: 10.1021/acs.inorgchem.8b00655 Inorg. Chem. XXXX, XXX, XXX−XXX

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Figure 2. Potentiometric titrations of UO22+/HEDTA complexation at varied temperatures. I = 1.0 mol·L−1 NaClO4. T = 298 K (a), 313 K (b), 328 K (c), 343 K (d). Cup solution: V0 = 17.44 mL (a), 14.79 mL (b), 14.52 mL (c), 14.78 mL (d); mL = 50.68 μmol (a), 50.82 μmol (b), 37.90 μmol (c), 50.64 μmol (d); mH = 177.14 μmol (a), 177.63 μmol (b), 138.79 μmol (c), 177.01 μmol (d). mU = 22.10 μmol, titrant = 100.3 mmol·L−1 NaOH for (a−d). Symbols (Δ)−experimental data; dashed red line−fit (related to right axis, −log[H+]). Solid lines: U(VI) speciation (related to left axis, species %), black−UO22+, green−UO2HL, blue−UO2L−, cyan−UO2(H−1L)2−. E = E0 +

RT ln[H+] + γH[H+] F

(1)

E = E0 +

KW RT ln + γOH[OH−] F [OH−]

(2)

the dilution heat of the titrant (Qdil,j). The value of dilution heat was determined in separate runs with 900 μL of 1.0 mol·L−1 NaClO4 in the cell (in the absence of U(VI) and HEDTA) titrated with the same NaOH solution. The net reaction heat at the jth point (Qr,j) was obtained from the difference: Qr,j = Qex,j − Qdil,j. The program HypDeltaH13 was used to analyze the data and calculate the enthalpy of complexation. Multiple titrations with different concentrations of UO22+ and HEDTA and varied initial acidity were performed to reduce the uncertainty of the results. NMR Spectroscopy. To investigate the coordination mode(s) in the U(VI)/HEDTA complexes, NMR experiments were performed on a series of U(VI)/HEDTA samples. The sample conditions, including ligand/metal ratios and pH, were carefully selected through simulation by using the program Hyss200914 so that a desired complex was dominated in the individual samples. With the selected conditions, an NMR solution was prepared by the following procedures. A weighed amount of solid UO2(NO3)2·6H2O and the desired volume of the HEDTA stock in D2O were added to deuterated water. Then, the pD (−log[D+]) of the solution was adjusted to a desired value by solutions of DClO4 or NaOD in D2O. The pD measurements were conducted with a combined pH electrode previously calibrated by three pH buffers (4.0, 7.0, and 10.0). For each sample, 1H and 13C NMR spectra were collected on a Bruker Avance DPX 300 spectrometer, operated at 300 MHz for the measurement of 1H signals. The 1H NMR spectra were recorded with a standard pulse program ZG30, averaging 32 scans for each spectrum. The obtained 1H NMR spectra were referenced to the solvent peak (HOD, 4.79 ppm). The 13C NMR spectra were collected with a decoupler on pulse program ZGDC, each averaging over 10 000 scans. Extended X-ray Absorption Fine Structure Spectroscopy. EXAFS was used to gain additional structural information on the U(VI)/HEDTA complexes. U LIII-edge data were collected at Stanford Synchrotron Radiation Lightsource (SSRL) for four solutions: a U(VI) solution in the absence of HEDTA and three U(VI)/HEDTA solutions with different ligand/metal ratios and pH values. Each U(VI)/HEDTA solution was prepared under the conditions where a desired complex is dominant. The solutions (each 2.0 mL) were contained in plastic vials and then sealed in plastic bags. The samples were mounted on a sample positioner with Scotch tape and measured on Beamline 11-2 with Y

where R is the gas constant, F is the Faraday constant, and T is the temperature in kelvin. E is the electromotive force measured through the pH electrode during the titration. KW is the water dissociation constant (KW = [H+] [OH−]). Equations (1) and (2) are used to address the acidic and basic regions, respectively, for the calibration fit. The last term is the electrode junction potential for the hydrogen ion (ΔEj,H+) or the hydroxide anion (ΔEj,OH‑) and assumed to be proportional to the concentration of the hydrogen ion or the hydroxide anion. The values of E°, γH, and γOH depend on the specific electrode, the temperature, and the ionic strength, and can be determined through an acid/base calibration titration prior to the measurement titration by using standard perchloric acid and sodium hydroxide. A hydrogen ion concentration is thus calculated from the electromotive force (EMF) measured in the ligand/metal titration. The ligand/metal titration was performed by adding sodium hydroxide solution into the acidic solution containing U(VI) and HEDTA. For each addition, the EMF was collected with the following criterion: the drift of EMF (ΔE) was less than 0.1 mv for 60 s. Multiple titrations were conducted at each temperature with the vessel solutions containing different concentrations of U(VI) (CM), HEDTA(CL), and total proton (CH). The stability constants of U(VI)/HEDTA complexes on the molarity scale were calculated with the program Hyperquad.11 Microcalorimetry. Microcalorimetry was used to determine the complexation enthalpy of U(VI) with HEDTA in a 1.00 mol·L−1 NaClO4 solution at 298 K. The calorimetric titrations were conducted on an isothermal microcalorimeter (ITC 4200, Calorimetry Sciences Corp). The performance of the calorimeter was tested by measuring the enthalpy of protonation of tris(hydroxymethyl)-aminomethane (THAM). The value (−47.6 ± 0.3) kJ·mol−1 was obtained at 298 K and is in agreement with the literature.12 A titration was initially set up with 900 μL of U(VI)/HEDTA solution in the calorimeter cell. An initial solution was controlled to be acidic. After thermal equilibrium was achieved, a solution of NaOH was added stepwise through a 250 μL syringe to the cell (5.0 μL per addition). In each titration, n additions were made (usually n = 40− 50), resulting in n experimental values of the heat generated during the titration (Qex,j, where j = 1 to n). These values were corrected for C

DOI: 10.1021/acs.inorgchem.8b00655 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry Table 1. Thermodynamic Parameters for the UO22+/HEDTA Complexes reaction 2+

2−

UO2 +HL

= UO2(HL)(aq)

UO22+ + L3− = UO2L−

UO22++ L3−=UO2(H−1L)2−+ H+

T, K 298 298 298 313 328 343 298 298 313 328 343 298 298 313 328 343

I, mol L−1

log β (±3σ)

ΔH (±3σ), kJ·mol−1

ΔS (±3σ), J·mol−1·K−1

ref.

1.0 1.0 0.2 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0 1.0

5.98 ± 0.03

a

24 ± 3 25 ± 3c

195 ± 10 198 ± 10

15 ± 3a 20 ± 3c

239 ± 10 256 ± 10

36 ± 1a 40 ± 3c

185 ± 3 198 ± 10

pwb pw 5 pw pw pw pw pw pw pw pw pw pw pw pw pw

6.23−6.37 6.14 ± 0.03 6.39 ± 0.03 6.48 ± 0.06 9.88 ± 0.06 10.07 ± 0.06 10.11 ± 0.06 10.27 ± 0.09 3.35 ± 0.09 3.69 ± 0.09 3.93 ± 0.09 4.20 ± 0.09

a

Values were calculated from the van’t Hoff plots. bpw stands for present work. cValues were obtained from calorimetry at 298 K.

foil as the reference (four scans for each sample). The data were collected up to kmax ≈ 15.0 Å−1 in both transmission and fluorescence modes. The EXAFS data reduction was performed with the program Athena.15 A spectrum was energy-calibrated to the reference (Y) by assigning the first inflection point of the K edge of yttrium to be 17 038 eV. The data reduction included pre-edge background subtraction followed by spline fitting and normalization. The EXAFS data were extracted above the threshold energy E0, defined as 17 166 eV. Fitting of the EXAFS data was conducted with the program Artemis.15 The fit utilized the theoretical phases and amplitudes calculated by the program FEFF716 with two model structures: piperazinium di-μ2-hydroxo-bis[iminodiacetato-dioxouranate(VI)] octahydrate17 and (m2-ethylene-1,2-diamine-N,N′-diacetato)-tetraoxo-bis(ethylene-1,2-diamine-N,N′-diacetato)-diuranium(VI) dihydrate.18 For each sample, selected single scattering (SS) and multiple scattering (MS) paths from the FEFF calculation were used in the fit based on the proposed coordination mode. In all the fits, an amplitude factor (S02) and a threshold energy shift (ΔE0) were considered to be global parameters. Hanning windows with a k range (3.0−14.0 Å−1) and Fourier transform with an R range (0.95−3.5 Å) were used. The fit R factor (r) is provided as an indication of the fit quality.

hydroxylic group upon the complexation. The resultant ligand is denoted by (H −1 L) 4− , where H −1 represents the deprotonation of the hydroxyl group. However, note that potentiometric titrations are unable to distinguish the deprotonation of the HEDTA hydroxyl group (as shown by Reaction 5) from that of water through the hydrolysis of uranyl (as shown by the reaction UO22+ + L3− + H2 O = UO2(OH)L2− + H+). Integration of the thermodynamic data and structrual anlaysis through NMR and EXAFS helped to clarify the deprotonation process and was discussed below. In the thermodynamic data evaluation, if assuming that the deprotonation is caused by the hydrolysis of uranyl, the hydrolyzed U(VI)/HEDTA complex would be UO2(OH)L2−. In this case, reaction 5 (log K(5) = 3.35 at 298 K, Table 1) is the sum of the following two reactions: hydrolysis of uranyl (Khyd: UO22+ + H2O = UO2(OH)+ + H+) and complexation of HEDTA with the hydrolyzed uranyl (Kcom: UO2(OH)+ + L3− = UO2(OH)L2−). The Khyd was obtanied from the literature19 to be log Khyd = −5.98 at 298 K, and the Kcom could thus be calculated to be log Kcom = log K(5) − log Khyd = 9.33 at 298 K, which is very close to the value for the complexation of L3− with UO22+ to form UO2L− shown by reaction 4 (log K(4) = 9.88 at 298 K, Table 1). Considering the diffrence in electric charges between the hydrolyzed uranyl (UO2(OH)+) and uranyl (UO22+), the stability constant for UO2(OH)L2− should be significantly smaller than that for UO2L−. Therefore, it is more reasonable to assume that the deprotonation occurred on the hydroxyl group of HEDTA in the hydrolyzed complex. More convincing evidence to support this assumption is presented in the structural analysis through the NMR and EXAFS work in subsequent sections. The enthalpies of complexation were evaluated by two approaches: The van’t Hoff equation using the equilibrium constants at different temperatures and the direct measurements by calorimetric titrations. The van’t Hoff plots of log β versus 1/T were depicted in Figure 3. The sound linearity of log β as a function of 1/T for each reaction suggests that the enthalpy of complexation could be considered constant, and the heat capacity of complexation is zero over this temperature range. From the slope of each line, the enthalpy values were calculated to be 24 ± 3, 15 ± 3, and 36 ± 1 kJ·mol−1 for reactions 3, (4), and (5), respectively.

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RESULTS AND DISCUSSION Stability Constants, Enthalpy, and Entropy. The stability constants of U(VI)/HEDTA complexes at varied temperatures (298−343 K) were determined by potentiometry. Details of titration conditions are listed in Table S1 in Supporting Information. Figure 2 shows representative potentiometric titrations of U(VI)/HEDTA at individual temperatures. The titrations were conducted at the pH range of 2.0−7.0 with varied ligand/metal ratios from 1.0 to 2.5. In the fit, various complexation models were tried, and the best fit was achieved with the following complexation reactions: UO22 + + HL2 − = UO2 (HL)(aq)

(3)

UO22 + + L3 − = UO2 L−

(4)

UO22 + + L3 − = UO2 (H−1L)2 − + H+

(5)

The obtained stability constants are sumarized in Table 1. Reaction 5 indicates that the L3− (three carboxylic groups fully deprotonated) experiences further deprotonation of its D

DOI: 10.1021/acs.inorgchem.8b00655 Inorg. Chem. XXXX, XXX, XXX−XXX

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determined stability constants of U(VI)/HEDTA complexes were used as known values. And the protonation constants and enthalpy values of HEDTA (cf. Table S3 in Supporitng Information) were also fixed with the values determined previously.10 The enthalpies of reactions 3, (4), and (5) were calculated to be 25 ± 3, 20 ± 3, and 40 ± 3 kJ·mol−1, respectively. Taking into account the uncertainties, the enthalpy results from calorimetry are in agreement with those from the van’t Hoff method (Table 1). It is noteworthy that an excellent fit was achieved with the same complexation model for both calorimetric and potentiometric data collected under the similar experimental conditions, such as the ligand to metal ratios (∼1.8/1.0) and the pH range (2.5−8.0). The mutual agreement between calorimetry and potentiometry further supports the proposed three complexation reactions. The thermodynamic data of U(VI)/HEDTA complexation (Table 1) indicate that all three complexation reactions are endothermic and exclusivley driven by large positive entropies. The large entropy changes suggest that the complexes are innersphere, and both the uranyl cation and the HEDTA ligand are substantially dehydrated upon complexation, thereby forming strong chelate complexes. These observations are well-supported by the structural analysis discussed below. Coordination Modes by NMR. Before analyzing the NMR data to obtain information on the coordination modes in the three U(VI)/HEDTA complexes, it is helpful to convert the equilibrium constant for reactions 5 into the “stability constants” for reaction 6, so that the binding strengths in the three complexes can be directly compared.

Figure 3. Van’t Hoff plots of log β vs 1/T. Graphs (a−c) correspond to reactions 3−5, respectively. Symbol (■) - experimental data; red solid line - linear fitting; dash line -95% confidence band.

Figure 4 shows a representative calorimetric titration of U(VI)/HEDTA at 298 K. Table S2 in Supporting Information lists the detailed conditions of all titrations. The total reaction heat (Q) was fitted with the complexation model proposed in the potentiometry work. In the fit, the potentiometrically

UO22 + + (H−1L)4 − = UO2 (H−1L)2 −

(6)

2−

The formation constant of UO2(H−1L) by reaction 6 is not directly available but can be roughly estimated by assuming the deprotonation constant for the hydroxyl group in HEDTA (L3− = H−1L4− + H+) is similar to that of water (log KH ≈ −14 at 298 K). Then the formation constant of the UO2(H−1L)2− by reaction 6 was calculated to be log K6 = log K5 − log KH = 17.4 at 298 K. The three formation constants, log K3 for UO2(HL)(aq) (5.98, Table 1), log K4 for UO2L− (9.88, Table 1), and log K6 for UO2(H−1L)2− (17.4), clearly indicate that HEDTA holds different binding strength in the three complexes, forming intermediate (UO2(HL)(aq)), strong (UO2L−), and very strong (UO2(H−1L)2−) complexes. The increasingly higher binding strength of HEDTA in the three complexes implies that the electronegativity as well as the denticity of the HEDTA ligand in the three complexes increases from (UO2(HL)(aq)) to (UO2L−), and further to (UO2(H−1L)2−). Accordingly, we proposed different coordination modes of HEDTA in the three complexes as shown in Figure 5. These coordination modes not only represent the implication of the binding strength but also are welldemonstrated by the NMR work, which is discussed below. The 1H and 13C NMR spectra for a series of five HEDTA solutions, in the absence and presence of U(VI), are presented in Figure 6 with solution conditions and calculated speciation in the caption. For each solution, the speciation calculation was conducted by the program Hyss200914 with stability constants of the HEDTA protonation10 and complexation with U(VI) (this work). Solution (1), in the absence of U(VI), contains only free ligand with different protonated forms. From solution (2) to (5), the dominant HEDTA complexes vary with the

Figure 4. Calorimetric titration of UO22+/HEDTA complexation. T = 298 K and I = 1.0 mol·L−3 NaClO4. Cup solution: V0 = 900 μL, mU = 0.834 μmol, mL = 1.415 μmol, mH = 5.192 μmol. Titrant: 25.15 mmol· L−1 NaOH. (upper) Thermogram (dilution not corrected). (lower) Fit of total reaction heat (Q) and specition of U(VI); measured (○)and calculated (red dash line) total reaction heat (right y-axis) as a function of titrant volume; percentage of U(VI) species (solid line, left y-axis) as a function of titrant volume, green−UO22+, blue− UO2(HL), cyan−UO2L−, magenta−UO2(H−1L)2−. E

DOI: 10.1021/acs.inorgchem.8b00655 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry

Figure 5. Proposed coordination modes of U(VI)/HEDTA complexes. (a) UO2(HL)(aq), (b) UO2L−, (c) UO2(H−1L)2−.

Figure 6. NMR spectra of U(VI)/HEDTA solutions. (left) 1H. (right) 13C. Solution conditions: CL = 103 (1), 36.5 (2), 35.5 (3), 35.0 (4), 34.5 mmol·L−1 (5); CU = 0.0 (1), 18.4 (2), 17.9 (3), 17.6 (4), 17.4 (5) mmol·L−1; CD = 229 (1), 57.6 (2), 31.8 (3), 19.4 (4), 7.37 (5) mmol·L−1; pD = 6.50 (1), 3.80 (2), 5.90 (3), 6.44 (4), 7.20 (5). Speciation (relevant to total ligand concentration): Solution (1): 20% D2L−, 78% DL2−; solution (2): 43% D2L−, 49% UO2(DL)(aq); solution (3): 25% D2L−, 25% DL2−, 16% UO2(DL)(aq), 30% UO2L−; solution (4): 11% D2L−; 38% DL2−, 34% UO2L−, 11% UO2(D−1L)2−; solution (5): 47% DL2−, 18% UO2L−, 32% UO2(D−1L)2−.

following: UO2(DL)(aq) → UO2L− → UO2(D−1L)2−. Inspection of Figure 6 reveals that the NMR features of both 1 H and 13C synchronically vary with the complexation. Therefore, an analysis of the variations of NMR features with the formation of different complexes could well yield the coordination details of individual complexes, thereby demonstrating the proposed coordination modes in Figure 5. For the convenience of interpretation of the NMR data, the alkyl groups of HEDTA are labeled by a, a′, b, b′, c, d, and e as shown in Figure 1. The spectra of solution (1) show clean, sharp peaks of chemical shifts in both 1H and 13C NMR (Figure 6(1)). These peaks were assigned referring to the early work.10 In the presence of U(VI), solution (2) gave different 1H and 13C spectra (cf. Figure 6(2)), presenting appreciable variations such as peak broadening and disappearing. Taking into account the calculated speciation of solution (2) (43% D2L−, 49% UO2(DL)(aq)), it is readily recognized that this newly formed complex, UO2(DL)(aq), was exclusively responsible

for such variations. To understand the coordination structure, three unique changes in its 13C spectrum are particularly useful: (1) one of the two carboxylic carbon peaks (a/a′) disappeared; (2) the peak of the a/a′ methyl carbons was broadened; (3) the peak of the b methyl carbon was also broadened. While the collected one-dimensional (1D) NMR data could not provide a direct way or evidence to confirm which carboxylic carbon peak disappeared, the broadening of the a/a′ methyl carbon peak suggests that it is very likely that the disappearance occurred to the a/a′ carboxylic carbons based on the following rationalizations. Assuming that the a/a′ carboxylic groups participate in the coordination to U(VI), the large electric charges of U(VI) could change the electromagnetic environments of its proximate atoms, thereby smearing out the a/a′ carboxylic carbon peak and broadening their neighboring carbon peak (a/a′ methyl carbons). The same argument could also be made to the b amine group, as its neighboring carbon, the b carbon, shows similar broadening behavior. Therefore, we could conclude that in UO2(DL)(aq), F

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Inorganic Chemistry

Figure 7. EXAFS of U(VI)/HEDTA samples. Solution conditions: I = 1.0 mol·L−1 NaClO4 or HClO4, V0 = 1.80 mL, CU = 24.5 (solution (I)), 24.6 (solution (II)), 24.6 (solution (III)), 1.11 mmol L−1 (solution (IV)); CL = 0.0 (I), 120.0 (II), 140.0 (III), 3.24 mmol L−1 (IV); CH = 1000 (I), 250 (II), 139 (III), 0.13 mmol L−1 (IV); pH = 2.92 (II), 6.10 (III), 7.20 (IV).

two oxygens of the a/a′ carboxylic group and a nitrogen of the b amine group participate in the coordination, forming a tridentate mode as depicted in Figure 5a. From solution (2) to solution (3), the NMR features in both 1 H and 13C continue to change. Regarding the speciation of solution (3), a new complex (UO2L−) formed by 30%, whereas the UO2(DL)(aq) remained in an appreciable amount (16%). It is thus believed that those NMR feature changes are attributed to the formation of this new complex. Compared to that of solution (2), the 13C spectrum of solution (3) varied dramatically, and only the b/b′ carbon peaks could be recognized. The dramatic variations imply that the coordination structure of UO2L− must be quite different from that of UO2(DL)(aq). A significant feature of the 13C spectrum of solution (3) is that the c carboxylic carbon peak also disappeared along with the a/a′ carboxylic carbon peak, which suggests that the c carboxylic group participates in the coordination of the new complex (UO2L−). As discussed earlier, UO2L− is much stronger than the UO2(HL)(aq) (the formation constant differs by ∼4 log units). Such enhancement in the binding strength suggests that HEDTA has a higher degree of chelation in the UO2L− complex than that in the UO2(HL)(aq) complex. It is very likely that HEDTA is tetradentate in the UO2L− complex, as depicted in Figure 5b. Two oxygens of the a and c carboxylic groups and two nitrogens of the amine groups bind to the U(VI) atom. An outstanding variation in the NMR features from solution (3) to (4) and to (5) is described below. In both 1H and 13C NMR, a set of peaks started to show up in solution (4), and became appreciable in solution (5), which are marked by asterisk (*) in Figure 6. Importantly, the advent of this new set of peaks was synchronized with the formation of UO2(D−1L)2− in solutions (4) and (5). It is thus reasonable to assign the new peaks to the UO2(D−1L)2− complex. The observation of the resonances of HEDTA in UO2(D−1L)2− as isolated peaks implies that the HEDTA ligand in the complex experienced a slow exchange,20 the rate of which is much below the NMR

time scale. A similar NMR behavior has been observed on the complexation of U(VI) with other chelating ligands like gluconate in aqueous solution.21 A slow ligand exchange rate indicates the strong binding strength of HEDTA in the UO2(D−1L)2− complex, which is in agreement with the thermodynamic constant calculated in a previous section (log K6 = 17.4, highest among the three complexes). This agreement is, in turn, provides support for the proposed reaction process for the formation of UO2(H−1L)2− (reaction 5), in which deprotonation of the hydroxyl group occurs upon the complexation and the strong chelation forms. Figure 5c depicts its coordination structure, where the alkoxide of HEDTA directly binds to the U(VI) atom. The new NMR features for solutions (4) and (5) also support the proposed coordination mode. Before further discussion, note that the majority of the new 13C peaks were assigned with our best knowledge of the reasonable ranges and possible displacements of resonances. While the assignment of three carboxyl carbon peaks holds more confidence, the assignment of others may carry less certainty. As shown in the 13 C spectrum of solution (5), the three carboxyl carbon peaks are separated. This is very much in agreement with the proposed coordination structure depicted in Figure 5c, in which three carboxylic groups are in different chemical environments. The most downfield peak was assigned to the a carboxyl carbon, as the a carboxylic group directly binds to U(VI) in the complex and receives the largest influence from it. Another feature is a large displacement of the e methyl carbon peak (cf. spectra of solutions (4) and (5)). This strongly supports the hydroxyl group deprotonation and participation in the complexation. In general, the alkoxide form holds strong complexation ability and could make a short U−O bond in the complex. As a result, the large influence of U(VI) on the electromagnetic environment of its neighboring carbon would be generated, causing the large downfield displacement of the e carbon resonance as we observed in its NMR spectra. G

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Inorganic Chemistry Table 2. EXAFS Fitting Results of U(IV) Species in Solutions solutiona I speciation: 100% UO22+ II speciation: 99% UO2(HL)(aq) (H3L = HEDTA) III speciation: 10% UO2(HL)(aq) 67% UO2L− 23% UO2(H−1L)2− IV speciation: 10% UO2L− 90% UO2(H−1L)2−

shell

Nb

Rb (Å)

σ2

notice

U−Oax U−Oeq U−Oax U−(O/N)eq U−C U−Oax U−(O/N)eq U−C U−Oax U−Oeq U−Neq U−C

2.0 5.1 2.0 4.2 4.0 2.0 4.2 5.9 2.0 2.1 2.0 5.8

1.77 2.41 1.78 2.36 3.36 1.78 2.35 3.36 1.79 2.31 2.46 3.35

0.0020 0.0053 0.0019 0.0066 0.0107 0.0020 0.0084 0.0120 0.0018 0.0059 0.0025 0.0101

S02 = 1.00, ΔE0 = 9.5 eV r = 0.009 S02 = 1.0, ΔE0 = 7.4 eV r = 0.010

S02 = 1.0, ΔE0 = 8.3 eV r = 0.010

S02 = 0.97, ΔE0 = 9.5 eV r = 0.016

a

The U(VI) speciation (relevant to total [U]) was calculated with the simulation program HySS200914 by using the complexation constants, which were determined in this work. bSolution I: U−Oax, R ± 0.01, N = 2.0 held constant; U−Oeq, R ± 0.01, N ± 0.6. Solution II: U−Oax, R ± 0.01, N = 2.0 held constant; U−(O/N)eq, R ± 0.01, N ± 0.8; U−C, R ± 0.05, N ± 1.0. Solution III: U−Oax, R ± 0.01, N = 2.0 held constant; U−(O/N)eq, R ± 0.01, N ± 1.0; U−C, R ± 0.04, N ± 1.1. Solution IV: U−Oax, R ± 0.01, N = 2.0 held constant; U−Oeq, R ± 0.02, N ± 0.4; U−Neq, R ± 0.03, N = 2.0 held constant; U−C, R ± 0.05, N ± 1.2.

Coordination Geometry by EXAFS. EXAFS analysis was conducted for four U(VI) solutions. Each solution was carefully prepared under selected conditions, so that a specific U(VI) species was dominant. As a result, structural information for each solution could be used to address the coordination mode of the specific U(VI) species. The raw and fitted data of EXAFS oscillations as well as the Fourier transformations (FT) are depicted in Figure 7. The fitting results and the U(VI) speciation, which was calculated in terms of solution conditions, are listed in Table 2. An excellent fit was achieved for each solution with proposed scattering paths. Solution (I) contains only free uranyl in the 1.0 mol·L−1 HClO4 solution in the absence of HEDTA. Under this condition, U(VI) is expected to be hydrated by five water molecules in its equatorial plane.22,23 The EXAFS fitting result agrees excellently with the above prediction. Two dioxo oxygens are located at the distance of 1.77 Å and 5.1 equatorial oxygens at the distance of 2.41 Å (Table 2), all of which agree with those observed in the early work.22−26 Solutions (II), (III), and (IV), in the presence of HEDTA, contain the U(VI)/HEDTA complexes: 99% UO2(HL)(aq) in solution (II), 67% UO2L− in solution (III) and 90% UO2(H−1L)2− in solution (IV). By using the obtained coordination number of the U−O/Neq shell (NU−O/Neq) and the calculated speciation results of the U(VI) species for individual solutions (Table 2), in combination of the proposed coordination number of HEDTA in the complexes (Figure 5), we estimated or calculated the remaining hydration number of the UO2(HL)(aq), UO2L−, and UO2(H−1L)2− complexes to be 1.2, 0.2, and 0.1, respectively. In view of the uncertainties of NU−O/Neq, it is concluded that only one hydration water remains in UO2(HL)(aq), and no hydration water exists in both UO2L− and UO2(H−1L)2−. This means that the formation of UO2(HL)(aq), UO2L−, and UO2(H−1L)2− accompanies the dehydration of four, five, and five water molecules, respectively, which well explains the observed large entropy changes of those complexation reactions in the thermodynamic study (see Table 1). As indicated in the FT spectra of solutions (II), (III), and (IV), existence of the U−C scattering shell demonstrates the occurrence of the HEDTA coordination to U(VI). An increase of the NU−C from 4.0 (solution II) to 5.9 (solution III) and 5.8

(solution IV) very well agrees with the proposed coordination structures of the dominant complexes in those solutions. As depicted in Figure 5, the number of carbon atom in proximity of the U(VI) atom is 4.0 in the UO2(HL)(aq), and 6.0 in both UO2L− and UO2(H−1L)2−. The variations in fitting results of the U−(O/N)eq shell from solution (I) to (IV) also support the proposed coordination modes of the complexes. For example, the RU‑(O/N)eq for solutions (II) and (III) is appreciably smaller than the RU−O for solution (I) and accompanied by a relatively large Debye− Waller factor. This tells that the carboxylic oxygen binds to the U(VI) atom in the complexes (UO2(HL)(aq) and UO2L−) more closely than the oxygen of water in the hydrated U(VI) species. Also, multiple subshells may exist in the scattering shell, including U−Ocarboxyl, U−N, and U−Owater, with the large range of distance distributions. For solution (IV), a shoulder occurred in this scattering shell (cf. the FT graph of solution (IV)). By the fit, the shell was resolved into two subshells: U− Oeq and U−Neq. A quite small RU‑Oeq value suggests the existence of a short U−O bond in the complex. This bond is very likely formed from the binding of the alkoxide form to the U(VI) in the complex because of its strong binding ability.

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CONCLUSION This work demonstrates that HEDTA holds unique complexation behavior with U(VI) in aqueous solutions. In the pH range of 2.0−7.0, it forms 1:1 U(VI)/HEDTA complexes with different degrees of deprotonation, that is, UO2(HL)(aq), UO2L−, and UO2(H−1L)2−. Thermodynamic parameters of the complexation were determined by potentiometry and calorimetry, including stability constants, enthalpy, and entropy. The results indicate that formation of the three complexes are endothermic and entropy-driven only. The complexation patterns as well as the thermodynamic parameters were well-supported by structural analysis. The structural analysis of three complexes was conducted by NMR and EXAFS. The analysis demonstrates the presence of different coordination modes in the complexes. In UO2(HL)(aq), HEDTA holds a tridentate binding mode to U(VI) through two oxygens of the two carboxylic groups and one nitrogen of the amine group at the end of HEDTA molecule. H

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Inorganic Chemistry In UO2L−, HEDTA holds a tetradentate mode through two oxygens of the two carboxylic groups and two nitrogens of the two amine groups along the side of HEDTA molecule. Particularly interesting is a coordination structure of the UO2(H−1L)2−, in which the U(VI) atom was bound by two nitrogens of the two amine groups, one oxygen of the carboxylic group and one oxygen of the alkoxide form of the HEDTA hydroxyl group. By integrating thermodynamic and structural data, the formation of this strong chelate complex through the deprotonation of the hydroxyl group was evidently demonstrated for the first time in the present study.

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(3) Okemgbo, A. A.; Hill, H. H.; Metcalf, S. G.; Bachelor, M. Metal ion interferences in reverse polarity capillary zone electrophoretic analysis of Hanford Defense Waste for ethylenediaminetetraacetic acid (EDTA) and n-hydroxyethylethylenediaminetriacetic acid (HEDTA). Anal. Chim. Acta 1999, 396 (2−3), 105−116. (4) Okemgbo, A. A.; Hill, H. H.; Metcalf, S. G.; Bachelor, M. Reverse polarity capillary zone electrophoretic determination of ethylenediaminetetraacetic acid and N-hydroxyethyl ethylenediaminetriacetic acid in Hanford defense waste. J. Microcolumn Sep. 2000, 12 (1), 48−56. (5) Bhat, T. R.; Venkateswara Rao, T. Complexes of uranium(VI) with N-hydroxyethyl ethylene-diaminetriacetic acid and ethyleneglycol bis (2-aminoethyl) ether tetraacetic acid. Z. Anorg. Allg. Chem. 1967, 354, 201−207. (6) Lai, T.-T.; Lee, S.-T. Polarography of uranium(VI) complexes with N-hydroxyethyl-ethylenediamine-triacetic acid. Anal. Chem. 1969, 41, 1316−1319. (7) Sill, C. W.; Peterson, H. E. Fluorescence Test for Uranium in Aqueous Solution. Anal. Chem. 1947, 19, 646−651. (8) Gran, G. Determination of the equivalence point in potentiometric titrations. Part II. Analyst 1952, 77 (920), 661−671. (9) Zanonato, P.; Di Bernardo, P.; Bismondo, A.; Liu, G.; Chen, X.; Rao, L. Hydrolysis of Uranium(VI) at Variable Temperatures (10−85 °C). J. Am. Chem. Soc. 2004, 126, 5515−5522. (10) Li, X.; Zhang, Z.; Endrizzi, F.; Martin, L. R.; Luo, S.; Rao, L. Effect of temperature on the protonation of N-(2-hydroxyethyl)ethylenediamine-N,N′,N′-triacetic acid in aqueous solutions: Potentiometric and calorimetric studies. J. Chem. Thermodyn. 2015, 85, 35− 41. (11) Gans, P.; Sabatini, A.; Vacca, A. Investigation of equilibrium in solution. Determination of equilibrium constants with the Hyperquad suite of programs. Talanta 1996, 43, 1739−1753. (12) Christensen, J. J.; Hansen, L. D.; Izatt, R. M. Handbook of Proton Ionization Heats and Related Thermodynamic Quantities; John Wiley and Sons: New York, 1976. (13) Gans, P.; Sabatini, A.; Vacca, A. Simultaneous Calculation of Equilibrium Constants and Standard Formation Enthalpies from Calorimetric Data for Systems with Multiple Equilibria in Solution. J. Solution Chem. 2008, 37 (4), 467−476. (14) Alderighi, L.; Gans, P. r.; Ienco, A.; Peters, D.; Sabatini, A.; Vacca, A. Hyperquad simulation and speciation (HySS): a utility program for the investigation of equilibria involving soluble and partially soluble species. Coord. Chem. Rev. 1999, 184, 311−318. (15) Ravel, B.; Newville, M. ATHENA, ARTEMIS, HEPHAESTUS: data analysis for X-ray absorption spectroscopy using IFEFFIT. J. Synchrotron Radiat. 2005, 12, 537−541. (16) Zabinsky, S. I.; Rehr, J. J.; Ankudinov, A.; Albers, R. C.; Eller, M. J. Multiple-scattering calculations of x-ray-absorption spectra. Phys. Rev. B: Condens. Matter Mater. Phys. 1995, 52, 2995. (17) Jiang, J.; Sarsfield, M. J.; Renshaw, J. C.; Livens, F. R.; Collison, D.; Charnock, J. M.; Helliwell, M.; Eccles, H. Synthesis and Characterization of Uranyl Compounds with Iminodiacetate and Oxydiacetate Displaying Variable Denticity. Inorg. Chem. 2002, 41, 2799. (18) Palladino, G.; Szabo, Z.; Fischer, A.; Grenthe, I. Structure, equilibrium and ligand exchange dynamics in the binary and ternary dioxouranium(VI)-ethylenediamine-N,N′-diacetic acid-fluoride system: a potentiometric, NMR and X-ray crystallographic study. Dalton Trans. 2006, 5176. (19) Martell, A. E.; Smith, R. M. NIST Critically Selected Stability Constants of Metal Complexes. In NIST Standard Reference Database 46, Version 6.0; Motekaitis, R. J., Ed.; NIST, 2001. (20) Drago, R. S. Physical Methods in Chemistry; Saunders: Philadelphia, 1977. (21) Zhang, Z.; Helms, G.; Clark, S.; Tian, G.; Zanonato, P.; Rao, L. Complexation of Uranium(VI) by Gluconate in Acidic Solutions: a Thermodynamic Study with Structural Analysis. Inorg. Chem. 2009, 48, 3814−3824.

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.8b00655. Tabulated experimental conditions for potentiometric and calorimetric titrations and thermodynamic parameters for the protonation of HEDTA (PDF)

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AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. Phone: +1-510-486-5142. (Z.Z.) *E-mail: [email protected]. Phone: +1-510-486-5427. (L.R.) ORCID

Xingliang Li: 0000-0002-4182-1572 Zhicheng Zhang: 0000-0002-2192-3846 Leigh R. Martin: 0000-0001-7241-7110 Linfeng Rao: 0000-0002-1873-0066 Present Address ∥

Nuclear Security & Isotope Technology Division, Oak Ridge National Laboratory, Oak Ridge, Tennessee 37831, United States. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The experimental work was supported by the Fuel Cycle Research and Development (FCR&D) Thermodynamics and Kinetics Program, Office of Nuclear Energy, the U.S. Department of Energy (DOE) under Contract No. DEAC02-05CH11231 at Lawrence Berkeley National Laboratory (LBNL). Preparation of the manuscript was supported by the Director, Office of Science, Office of Basic Energy Sciences under U.S. DOE Contract No. DE-AC02-05CH11231 at LBNL. L.R.M. acknowledges the support from DOE NE FCR&D Thermodynamics and Kinetics program, under DOE Idaho Operations Office Contract No. DE-AC07-05ID14517. The EXAFS data were collected at Stanford Synchrotron Radiation Laboratory, which is a user facility operated for the U.S. DOE by Stanford Univ.

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REFERENCES

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