A Combined Density Functional Theory and Spectrophotometry Study

Apr 14, 2017 - The equilibrium constants for [NpO2·M]4+ (M = Al3+, In3+, Sc3+, Fe3+) in μ = 10 M nitric acid and [NpO2·Ga]4+ in μ = 10 M hydrochlo...
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A Combined Density Functional Theory and Spectrophotometry Study of the Bonding Interactions of [NpO2·M]4+ Cation−Cation Complexes John W. Freiderich,† Adam G. Burn,† Leigh R. Martin,‡ Kenneth L. Nash,*,† and Aurora E. Clark*,† †

Department of Chemistry, Washington State University, Pullman, Washington, United States Idaho National Laboratory, Idaho Falls, Idaho, United States

‡

S Supporting Information *

ABSTRACT: The equilibrium constants for [NpO2·M]4+ (M = Al3+, In3+, Sc3+, Fe3+) in μ = 10 M nitric acid and [NpO2·Ga]4+ in μ = 10 M hydrochloric acid media have been determined. The trend in the interaction strength follows: Fe3+ > Sc3+ ≥ In3+ > Ga3+ ≫ Al3+. These equilibrium constants are compared to those of previously reported values for NpO2+ complexes with Cr3+ and Rh3+ within the literature. Thermodynamic parameters and bonding modes are discussed, with density functional theory and natural bond orbital analysis indicating that the NpO2+ dioxocation acts as a π-donor with transition-metal cations and a σ-donor with group 13 cations. The small changes in electron-donating ability is modulated by the overlap with the coordinating metal ion’s valence atomic orbitals.

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INTRODUCTION The reprocessing of used nuclear fuel remains a topic of research for many countries. One such means of reprocessing is through exploitation of the redox chemistry of the fuel constituents. The Plutonium and Uranium Recovery by Extraction (PUREX) process, developed in the United States, has been successful in France for recovery and reuse of uranium for commercial reactors. While successful at segregation of UO22+ and Pu4+ from the remainder of the fuel, the PUREX process leaves the minor actinides (Np, Am, Cm) and other used fuel constituents together. This remaining fuel component is typically high in ionic strength and has been shown to support the formation of cation−cation complexes. Nagasaki1 and Guillaume2 have shown that such interactions dictate the redox chemistry, and thus the cation−cation complex may influence the chemistry of future nuclear fuel reprocessing efforts. The binuclear actinyl−actinyl complex reported by Sullivan demonstrates the tendency of the NpO2+ ion to bind to UO22+,3 but a pair of dioxocations is not the only configuration that may form a complex of this type. In 1962,4 the interaction of a series of trivalent cations (M3+ = Al3+, Ga3+, In3+, Sc3+, and Fe3+) with NpO2+ was examined, and the relative stability of the resulting [NpO2·M]4+ complexes was determined based on the molar absorptivity of the 980 nm f−f NpO2+ UV−vis absorption band. The strength of the complex in perchloric acid media followed the trend: Fe3+ > In3+ > Sc3+ > Ga3+ > Al3+. The neptunyl monocation has also been shown to form complexes with Cr3+ and Rh3+.5,6 In addition, several reports have discussed the bonding modes of VO2+·VO3+,7 VO2+·Ru2+,8 and NpO2+·UO22+.9 While these demonstrate the formation of cation−cation complexes involving NpO2+ and their relative © 2017 American Chemical Society

stabilities, the actual stability constants for most complexes are unknown, and the electronic structure and bonding of the dioxocation therein is far from certain. In this work, stability (or equilibrium) constants were determined for neptunyl−cation formation with Fe3+, Sc3+, Ga3+, In3+, and Al3+ using spectrophotometry as a function of the transition-metal concentration, and under similar solution chemistry conditions as prior reports.4 A detailed analysis of the bonding interactions has complemented the experimental work using density functional theory. These data indicate that the NpO2+ dioxocation acts as a π-donor with transition-metal cations and a σ-donor with group 13 cations. The small changes in electron donating ability is modulated by the overlap with the coordinating metal ion’s valence atomic orbitals.

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EXPERIMENTAL SECTION

General Procedure and Materials. All solutions were prepared from reagent grade materials with distilled 18 MΩ deionized water. Standard reagent grade HCl (36.5%, J.T. Baker) and HNO3 (69%, J.T. Baker) solutions were standardized through potentiometric titrations. Sodium nitrate (NaNO3, Reagent ACS) solid was dissolved in degassed deionized water, filtered through a fine glass frit filter, and recrystallized from hot water. The NaNO3 solution was then standardized using ion-exchange chromatography (Dowex 50 × 4 beads, H+ form) and potentiometric titrations. The Np working solution was created at the Idaho National Laboratory from a 237Np (in HF/3 M HNO3) mother solution. Neptunium-237 is a radioactive element with a half-life of 2.2 × 106 y and emits both α and γ radiation. Special engineering controls including shielding and a fume hood were used along with personal Received: October 5, 2016 Published: April 14, 2017 4788

DOI: 10.1021/acs.inorgchem.6b02369 Inorg. Chem. 2017, 56, 4788−4795

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Inorganic Chemistry

Figure 1. Absorbance spectra of NpO2+ cation−cation complexes in μ = 10 M with HNO3 or HCl. [NpO2+] = 1 mM, [M3+] = 0.62−2 M. Spectra shown for metals: Fe3+ (A), Sc3+ (B), In3+ (C), Al3+ (D), Ga3+ (E). Spectra of 1 mM NpO2+ is identical to (D). The transition metals used include Fe3+, Sc3+, Al3+, Ga3+, and In3+. Iron(III) nitrate nonahydrate (Fe(NO3)3·9H2O, 99.99%, SigmaAldrich) and aluminum nitrate nonahydrate (Al(NO3)3·9H2O, ≥ 98%, Sigma-Aldrich) were used as received. Scandium nitrate hydrate (99.99%, Sigma-Aldrich) was analyzed for water content by Karl Fischer (KF) coulometric analysis and used as received. Anhydrous gallium chloride (GaCl3, ∼10 mesh, 99.99%, Sigma-Aldrich) was used as received. Indium nitrate hydrate was prepared from indium perchlorate hydrate. The hydrolysis products of indium have been reported by Baes and Mesmer,13 and based on these data, it was possible to precipitate the In hydroxide(s) at pH 8 and suppress most of the water-soluble hydrolysis products. After precipitation the solid was washed with deionized water and tested for perchlorate with KNO3. The solid was then dissolved in 1 M HNO3. This solution was dried in an oven at 70 °C and analyzed for water content (KF analysis), H+, and In3+ using ion-exchange chromatography (Dowex 50 × 8, H+ form), and potentiometric titrations. Instrumentation. Experiments involving NpO2+ were monitored using a Varian Cary 6000i spectrophotometer and were conducted in the Materials and Fuels Complex (MFC) at the Idaho National Laboratory. A temperature-controlled 12-compartment cassette was used to hold 2 mL capacity with 1.0 cm path length quartz cuvettes

protection equipment to safely handle this material. The dominant neptunium oxidation states in this matrix were Np(IV, VI). On the basis of elution profiles by Horwitz et al.10 for transuranic (TRU) resin columns (Eichrom Technologies Inc.), a 5 mL aliquot of 237Np in HF/ HNO3 media was injected onto the TRU resin column (part number: TR-R50-S) and eluted with 2 M HNO3. The Np4+ and NpO22+ were retained within the column as evidenced by observation of the green/ brown color of the resin. Spectrophotometry confirmed the presence and concentration of NpO2+ as ∼8% by an intense absorption band at 980 nm in the absorption spectrum. The resin-bound Np4+ and NpO22+ were eluted with 0.01 M HCl. This Np4+/NpO22+ solution was titrated with 30% hydrogen peroxide (H2O2) until the solution turned a teal-green color, indicating that a majority of the NpO2+ had been produced from Np4+ and NpO22+.11 The redox titration proceeded according to the following reactions:

2NpO2 2 + + H 2O2 ⇆ NpO2+ + O2(g) + 2H+ 2Np4 + + 2H+ + O2(g) ⇆ H 2O2 + 2NpO2+ + 8H+ This solution was then brought to volume in a volumetric flask and standardized using Beer’s Law and a literature value for the molar absorptivity (ε = 395 M−1 cm−1).12 4789

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Inorganic Chemistry

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with screw caps. The instrument was set into a dual-beam mode so that each sample solution could be run with its respective background solvent simultaneously. All solutions were prepared by mass or by calibrated pipettes prior to titration additions in controlled radiological areas. NpO2+ in aqueous media has a strong near-infrared (NIR) absorbance at 980 nm, attributed to f−f transitions that are highly sensitive to changes in the coordination environment. Complexation of NpO2+ leads to new, typically red-shifted, absorption peaks near this f−f transition.14 Previously reported neptunyl cation−cation complexes have absorption peaks between 992 and 1004 nm.4 Here, the vis−NIR spectra was measured between 900 and 1050 nm, using 0.2 nm resolution. The fitting of the spectrophotometric data was done using Hyperquad v1.1.1. developed by Gans.15 Note that the reported equilibrium constants are conditional. In all other literature accounts involving cation−cation complexes, the equilibrium constants are fit assuming a simple solvation environment whether the aqueous system is chloride-, nitrate-, sulfate-, or perchlorate-based. The data presented here maintain this literature standard for purposes of consistency and comparison. Computational Methods. Density functional theory (DFT) was employed using the B3LYP hybrid functional to optimize the gas- and solution-phase geometries of NpO2+ and [NpO2·M]4+ species.16−19 The triplet spin states of the [NpO2·Al]4+, [NpO2·Ga]4+, [NpO2·In]4+, and [NpO2·Sc]4+ were studied, while the octet spin state of the [NpO2·Fe]4+ structure was pursued. The atomic orbitals of the trivalent metals (Al, Sc, Fe, and Ga), hydrogen, and oxygen were described using the all-electron aug-cc-pVTZ basis set.20−23 The basis set for indium utilized the pseudopotential-based aug-cc-pVTZ-pp, while the Stuttgart RSC 1997 relativistically corrected effective core potentials (RECP) and associated basis set was used for neptunium.24,25 This computational protocol (method and basis set) has been shown to yield reasonable geometric parameters across a wide range of actinide and lanthanide solvated clusters. The optimized structures were confirmed to be local minima on the potential energy surface using normal-mode analysis. Similarly, the explicitly solvated NpO2(H2O)4+, M(H2O)63+, and [NpO2(H2O)4·M(H2O)5]4+ that have full solvation within the primary coordination sphere of each ion were also optimized in the gas and solution phases. These hydration numbers were chosen based upon prior study of the fully solvated metal ions that have indicated their preference in the dilute aqueous limit.26−28 The solution phase (beyond the first solvation shell) was described by a polarizable continuum model with the integral equation formalism variant (PCM) and the UFF cavity.29 Single-point energy calculations of the low-spin analogues of these systems revealed that they were higher in energy than the high-spin cases. Interestingly, the explicitly solvated complex [NpO2(H2O)4·Fe(H2O)5]4+aq was not able to be fully optimized, and as such its data is not discussed herein. After analysis of the electronic structure, charge donation from the solvating waters indicates that extended solvation shells may be necessary to optimize this system, which is beyond the scope of the current work and does not change the conclusions derived herein within the series studied. Prior work has also illustrated that extended solvation shells (perhaps up to the third shell) would also likely be needed to obtain accurate thermochemistry for the free energy of cation−cation formation30 (within 5 kcal/mol) and that the very small experimentally obtained free energies ( Sc3+ > In3+ ≥ Ga3+ ≫ Al3+. Though this trend generally agrees with prior work,4 the stability constants themselves provide evidence that Sc3+ forms a slightly more stable complex than In3+. In terms of the free energy of complex formation, all ΔGcc are under 1 kcal/mol, and no complex was observed for Al3+ in nitric acid media (μ = 10 M) in the current work. Trends in Geometric Features of Neptunyl−Cation Complexes. Figure 2 presents the B3LYP implicit solvation geometries of NpO2+aq and [NpO2·M]4+aq and the explicitly solvated NpO2(H2O)4+aq, M(H2O)63+aq, and [NpO2(H2O)4· M(H2O)5]4+aq clusters. While gas-phase geometries for NpO2+ and [NpO2·M]4+ were obtained, only the solution-phase data are discussed herein, with the gas-phase data presented in Supporting Information. The transition and main-group metal−water distances of the explicitly solvated clusters are well-reproduced with respect to the reported experimental literature data, having an average deviation of 0.04 Å relative to values obtained by a variety of X-ray diffraction and spectroscopic methods.34−38 The Np−O distance in the NpO2+ cation has been measured as 1.80 ± 0.02 Å in aqueous 1 M HClO439 using in situ extended X-ray absorption fine 4790

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Inorganic Chemistry V (V)O2+ + V (V)O3 + ⇆ [VO2 ·VO]4 +

In this case, two pentavalent vanadium oxocations (VO2+ and VO3+) form a cation−cation complex in highly acidic solutions (11.8 M HClO4 and 10 M H2SO4). The favorability of the complex was highly dependent on μ (Keq(μ = 5M) ≈ 0.8 M−1 and Keq(μ = 8M) = 8.1 ± 0.1 M−1); however, it was fundamentally proposed to be based upon a relatively small separation in energy levels between the VO2+ 2pπ highest occupied molecular orbital (HOMO) of the oxygen and the VO3+ lowest unoccupied molecular orbital (LUMO) t2g orbital having dxy symmetry (Figure 3A), enabling on some level of a d−p π bond.

Figure 2. Optimized geometries of implicitly solvated NpO2+, [NpO2· M]4+, and their explicitly solvated analogues (M = Al3+, Ga3+, In3+, Fe3+, and Sc3+). Bond distances in angstroms. The presentation of the transition metals is in order of increasing ionic radius (note that [NpO2(H2O)4·Fe(H2O)5]4+ is not discussed, because it was not able to be fully optimized) .

structure (EXAFS) spectroscopy, which is within 0.01 Å of both the implicit and explicitly solvated DFT calculated distances in Figure 2. All neptunyl−cation complexes adopt a linear geometry that may facilitate spatial overlap of the valence s, p, and d atomic orbitals of the transition metals with the molecular orbitals of the NpO2+ moiety. Both the implicitly and explicitly solvated structures are predicted to have perturbations to the two neptunyl bond lengths as a result of the cation interaction with the bridging diyl oxo-atom, referred to hereafter as the μ-oxo (or μO). This comes in the form of lengthening of the Np−μO bond, and slight contraction of the neptunyl bond trans to the interaction. Use of explicit solvation leads to smaller perturbations in the neptunyl bond length than that observed using implicit solvation, less sensitivity of the neptunium-μ-oxo bond distance to the particular transition metal involved, and longer bond distances between the transition metal and μ-oxo atom. Considering the exact trends in the bond distances as a function of transition metal, it is observed that M-μO distance generally increases as the ionic radius of the trivalent metal increases.40 At the same time the predicted bond distance of the Np−μO generally decreases with the ionic radius of the transition metal in the implicit solvation complex, while in the explicitly solvated clusters the Np−μO lengths are all within 0.03 Å of each other. This is demonstrative of a change in the electronic structure, most likely electrostatic in nature, imparted by the first solvation shell of H2O (vide infra). Electronic Structure of Neptunyl-Cation Complexes. While electrostatic contributions no doubt contribute to the favorable cation−cation interaction, as a consequence of the octahedral geometry of the trivalent metals in aqueous solution,40,41 they may also participate in σ and/or π bonding interactions due to overlap of the metal d orbitals and the s, px, py, and pz orbitals of the neptunyl oxygen. To consider this, an interesting analogue to the neptunyl−cation complexes can be found in the work of Madic and co-workers,7 who examined the complexation reaction of two pentavalent vanadium oxocations:

Figure 3. Proposed bonding mode for cation−cation complexes involving oxocation species. Arrows indicate direction of electron migration for (A) [VO 2 ·VO] 4+ 7 and (B) [(H 2 O) 5 V (IV) O· Ru(II)(NH3)5]4+.8

De Smedt also considered vanadium cation−cation complex formation; however, it was in the context of the oxocation [(H2O)5V(IV)O·Ru(II)(NH3)5]4+,8 instead of the dioxo species. This kinetically stable8 complex containing VO2+ was compared to that of octahedral Ru(II) complexes containing N2, CH3CN, and CO, which are weak σ-donors, but good π-acceptors. In this case, the orbitals of VO2+ are spatially positioned to overlap with the dxz and dyz orbitals of Ru. The strength of this complex is attributable to Ru(II) back-bonding with the dxz and dyz orbitals, which have the proper energy and symmetry to interact with the lower energy vacant orbitals of π-acceptor ligands such as N2 and VO2+ as illustrated in Figure 3B. One may infer from these data that if the oxocation VO2+ has similar bonding properties to dioxocations, then all cations of these type may also be poor σ-donors and that NpO2+ likely behaves as a π-acceptor. While no detailed bonding analysis has been provided within the experimental studies of the cation−cation complexes of neptunyl, previous DFT studies9 of [NpO2· UO2]3+ and [NpO2·NpO2]2+ gas-phase complexes with examination of the molecular orbitals have suggested that the [NpO2·NpO2]2+ dimer interacts in a donation/backbonding type of interaction between the fz2x of the Np metal center and π* of the NpO bond, though this was not discussed in detail. Herein we systematically examine the electronic structure and bonding properties of each neptunyl−cation complex, so as to determine what, if any, bonding interactions may facilitate the experimentally determined small negative ΔGcc in this work. The electronic structure of the implicitly and explicitly solvated neptunyl transition-metal complexes has been considered from several different perspectives. It is useful to 4791

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Inorganic Chemistry

Table 2. Natural Electron Configurationa of the Atoms in the Implicitly and Explicitly Solvated NpO2+ Reactant and cation− Cation Products, Where M = Al3+, Sc3+, Fe3+, Ga3+, and In3+ reactants NpO2(H2O)4+aq Sc(H2O)63+aq Al(H2O)63+aq Ga(H2O)63+aq In(H2O)63+aq 3d metal products [NpO2·Fe] 4+aq [NpO2·Sc] 4+aq [NpO2(H2O)4·Sc(H2O)5]4+aq group 13 metal products [NpO2·Al] 4+aq [NpO2(H2O)4·Al(H2O)5]4+aq [NpO2·Ga] 4+aq [NpO2(H2O)4·Ga(H2O)5]4+aq [NpO2·In]4+aq [NpO2(H2O)4·In(H2O)5]4+aq a

μO

Np

O

[He]2s1.902p4.89 [He]2s1.802p4.80

[Rn]5f3.846d0.70 [Rn]7s0.125f4.106d1.067p0.20

[He]2s1.902p4.89 [He]2s1.802p4.80

M3+

NpO2+aq [Ar]4s0.163d0.804p0.32 [Ne]3s0.35 3p0.61 [Ar]4s0.473d104p0.59 [Kr]5s0.444d105p0.49 M3+ [Ar] 4s03d5.50 [Ar]4s03d0.29 [Ar]4s0.153d0.844p0.314d0.05 M3+ [Ne]3s0.263p0.11 [Ne]3s0.343p0.58 [Ar]4s1.043d104p0.16 [Ar]4s0.473d104p0.57 [Kr]5s1.104d105p0.04 [Kr]5s0.454d105p0.48

μO

Np

O

[He]2s1.922p5.07 [He]2s1.902p5.23 [He]2s1.732p5.07 μO

[Rn]5f3.466d0.44 [Rn]5f3.546d0.49 [Rn]7s0.125f3.856d0.967p0.19 Np

[He]2s1.922p4.69 [He]2s1.912p4.72 [He]2s1.802p4.66 O

[He]2s1.912p5.46 [He]2s1.752p5.17 [He]2s1.912p5.19 [He]2s1.752p5.15 [He]2s1.882p4.93 [He]2s1.772p5.11

[Rn]5f3.376d0.39 [Rn]7s0.115f3.846d0.957p0.19 [Rn]5f3.046d0.42 [Rn]7s0.125f3.866d0.967p0.19 [Rn]5f3.266d0.62 [Rn]7s0.115f3.856d0.967p0.19

[He]2s1.932p4.60 [He]2s1.812p4.66 [He]2s1.932p4.39 [He]2s1.802p4.66 [He]2s1.902p4.46 [He]2s1.802p4.66

Electron occupancies of less than 0.05 are not included, and the coordinating metal ion reactant has an ideal electron configuration in all systems.

Table 3. Change in the Natural Atomic Chargesa in the Implicitly Solvated [NpO2·M]4+ and Explicitly Solvated [NpO2(H2O)4· M(H2O)5]4+, Where M = Al3+, Sc3+, Fe3+, Ga3+, and In3+ 3d metal

rion28

M

μO

[NpO2·Fe] aq [NpO2·Sc]4+aq [NpO2(H2O)4·Sc(H2O)5]4+aq group 13 metal

0.645 0.745 rion28

−0.55 −0.30 −0.02 M

−0.21 −0.34 −0.19 μO

[NpO2·Al]4+aq

0.535

4+

[NpO2(H2O)4·Al(H2O)5]4+aq [NpO2·Ga]4+aq [NpO2(H2O)4·Ga(H2O)5]4+aq [NpO2·In]4+aq [NpO2(H2O)4·In(H2O)5]4+aq

0.620 0.800

−0.38 (0.02) −1.19 (0.01) −1.14 −0.01

−0.59 (−0.32) −0.32 (−0.29) −0.03 −0.27

∑M‑μO

Np

O

−0.76 −0.64 −0.21 ∑M‑μO

0.58 0.48 0.29 Np

0.17 0.15 0.14 O

−0.97 (−0.30) −1.51 (−0.28) −1.17 −0.28

0.70 (0.31) 1.04 (0.28) 0.74 0.28

0.26 (0.14) 0.46 (0.14) 0.42 0.14

a

The change in charge is determined relative to the implicitly solvated NpO2+ and M3+ and explicitly solvated NpO2(H2O)4+ and M(H2O)63+. Comparisons to the ionic radius of M, rion (in Å), presume a coordination number of 6.28

metal and μO gain significant electron density upon cation− cation formation, where the majority (60−75%) of those electrons have migrated from the Np center. In accordance with the trends in bond distances discussed above, the charge of the coordinating metal in the cation−cation complex generally becomes more negative with increasing ionic radius of the group 13 cations, while the opposite trend is observed for the 3d transition metals studied (Table 3). The charge of the bridging “yl” oxygen increases with the group 13 cation ionic radius while decreasing as a function of the 3d cation ionic radius. The bonding within the cation−cation complexes can be discerned via study of both the Natural Bond Orbitals (which are orbitals that localize the electron density between pairs of atoms) as well as the molecular orbitals that represent the standard delocalized description of chemical bonding. Let us first consider these data for the cation−cation complexes composed of the 3d coordinating metals. Trivalent scandium is a d0 cation, and thus the available bonding (beyond purely ionic) is to accept electrons from the filled μO 2p to form a d− p π bond. The stability constant for NpO2+ with trivalent scandium is determined to be Keq = 2.0 ± 0.1 M−1 (Table 1). Both implicitly and explicitly solvated trivalent scandium acts

consider both implicitly and explicitly solvated systems, as they provide insight into the impact of H2O in the primary coordination shell upon the cation−cation interactions beyond the dielectric continuum of the bulk solvent. Natural Population Analysis defines localized 1-center orbitals that, among other things, are optimized to represent the effective atomic charge in a molecular environment. Analysis of the natural electron configuration of each atom in a molecule can help understand the nature of migration of electron density between the isolated NpO2+ and coordinating metal ions upon formation of the cation−cation complex. In the implicitly solvated NpO2+, the Np center (which is formally of oxidation state V with a 5f2 electron configuration) receives significant electron density from the diyl−oxo atoms such that its initial natural electron configuration is [Rn]5f3.846d0.70 (Table 2). Significant migration of electron density from the neptunium and diyl-oxo trans to the μO−M interaction occurs upon formation of the cation−cation complex. Table 3 presents the change in natural atomic charges of the coordinating metal cation, the μO, Np, and trans-oxo atoms upon formation of the cation−cation complex relative to the implicitly and explicitly solvated reactants (charge of products minus charge of reactants). In the implicitly solvated systems the coordinating 4792

DOI: 10.1021/acs.inorgchem.6b02369 Inorg. Chem. 2017, 56, 4788−4795

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Inorganic Chemistry primarily to polarize electron density from the Np to the μO, and any molecular or natural bonding orbitals between the Sc and μO contains minimal transition metal (TM) character (Tables 3 and 4). The implicit model has only 4% TM

the 3d period. In the case of the group 13 metals, the octahedral d10 metals, Ga3+ and In3+, exist with completely filled eg and t2g orbitals,40 and thus the electron density migrates from the neptunyl π orbital into predominantly the empty valence s AO (Table 2) in the implicitly solvated systems. The larger the s AO (by going down the group), the more charge transfer is observed and the larger the participation of the group 13 atomic orbitals in both the NBOs and MOs of the M−μO interaction. The nature of the bonding interaction is decidedly σ in its characteristics (Figure 4). No [NpO2·Al]4+ complex forms, as shown from spectral measurements of NpO2+ in 2 M Al3+. This is somewhat in contrast to Sullivans work in perchlorate media,4 where he observed a decrease in ε(Np(V), λ = 980 nm) from 400 to 360 cm−1 M−1 and attributed this to a weak [NpO2·Al]4+ complex. No spectral changes were observed in this work. The extent of charge donation from Np spans 0.5 to 1 e− with most of the electrons leaving the 5f natural atomic orbital (Table 3), and thus in the most extreme example, [NpO2·Ga]4+, the neptunium has gone from a [Rn]5f3.846d0.70 electron configuration in implicitly solvated NpO2+ to a [Rn]5f3.046d0.42 configuration in the cation−cation complex. These observations agree with the increase in stability constant going down group 13. In combination, the changes in electronic structure upon formation of the cation−cation complex within the implicitly solvated system indicates that the NpO2+ acts as an electron donor, wherein the filled μO donates electrons to the empty valence orbitals of the coordinating cation and forms a partial d−p π bond with the transition-metal cations and donates into the empty s AOs of the group 13 cations in a σ-bonding fashion. A slightly more nuanced description emerges when studying the changes in electronic structure of the explicitly solvated species. In general, within these systems, electron density does not migrate to the coordinating metal ion but rather to the μO from the Np center. This is because explicit solvation of the coordinating metal ion enables electron density migration from the solvent to the ion and thus dampens the electrostatic attraction for the coordinating metal and the NpO2+ unit. When the μO becomes part of the first solvation shell of the coordinating metal ion, it donates essentially the same number of electrons to the coordinating ion as H2O. However, this in turn leads to the μO pulling electron density from the Np center to it. Thus, in general, explicit solvation leads to more Natural Bonding Orbitals between the coordinating ion and the neptunyl unit but less charge transfer. Indeed, this may be one reason why the fully solvated [NpO2(H2O)4·Fe(H2O)5]4+ experienced difficulty being optimizedas migration of electron density from the H2O in the primary solvation shell may necessitate a second solvation shell to maintain stability.

Table 4. α and β Natural Bond Orbitals between the Trivalent Metal and Bridging Neptunyl Oxygen for the Implicitly Solvated [NpO2·M]4+ and Explicitly Solvated [NpO2(H2O)4·M(H2O)5]4+ Complexes No. (α, β)

%M

%O

e− occupancy

[NpO2·Sc]4+aq

α (2)

[NpO2(H2O)4·Sc(H2O)5]4+aq

α (1) β (3)

[NpO2·Fe]4+aq

β (3)

No. (α, β)

4 4 9 10 5 4 19 19 16 %M

96 96 91 90 95 96 81 81 84 %O

α (0.92) α (0.92) α (0.97) β (0.98) β (0.91) β (0.91) β (0.96) β (0.96) β (0.94) e− occupancy

α (1) β (1) β (1) α (1) β (1)

9 9 16 12 12

91 91 84 88 88

α (0.98) β (0.98) β (0.94) α (0.97) β (0.98)

α (1) β (1)

11 11

89 89

α (0.97) β (0.97)

3d metal

group 13 metal [NpO2·Al]4+aq [NpO2(H2O)4·Al(H2O)5]4+aq [NpO2·Ga]4+aq [NpO2(H2O)4·Ga(H2O)5]4+aq [NpO2·In]4+aq [NpO2(H2O)4·In(H2O)5]4+aq

participation in two predicted Sc-μO NBOs, while the explicitly solvated cation−cation complex has three NBOs with 4−10% participation of the TM atomic orbitals. Accordingly there is very little electron migration to the Sc center, and the electron configuration is very similar to that of the trivalent Sc reactant. Moving across the TM series to Fe changes the bonding description significantly. As illustrated in Figure 4 and

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CONCLUSIONS From experimental spectrophotometric measurements of NpO2+ species in μ = 10 M acidic aqueous media, with trivalent metals, and computational studies, the following trend in complex stability was determined: Fe3+ > Sc3+ > In3+ > Ga3+ ≫ Al3+. These different metal centers provide insight into the factors influencing the actinyl−cation complex. Both free energies of complexation and available bonding interactions for these metals were analyzed. The computational results from this study confirm a prior computational study9 suggesting that NpO2+ can behave as a π-donor with a cation that has empty p

Figure 4. (A) The α and β d−p π molecular orbitals in implicitly solvated [NpO2·Fe]aq4+, and (B) the α and β s−p σ molecular orbitals in implicitly solvated [NpO2·Ga]aq4+.

quantified in Tables 3 and 4, the cation−cation complexation with Fe3+aq has more overlap of the partially filled 3d atomic orbitals (AOs) with the μO 2pπ to form at least a partial d−p π bond, leading to a change in the natural electron configuration of the Fe-center from d5 to d5.5 in the implicitly solvated [NpO2·Fe]3+ complex. This observation agrees well with the experimentally determined trend in stability constants across 4793

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AOs as in transition metals; however, when empty s AOs are present, as in the group 13 metals, it can act as a σ-donor.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.6b02369. A brief thermodynamic discussion, optimized structures, and electronic configurations (PDF)

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AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. (K.L.N.) *E-mail: [email protected]. (A.E.C.) ORCID

John W. Freiderich: 0000-0002-5792-8892 Leigh R. Martin: 0000-0001-7241-7110 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS K.N. acknowledges support from the Department of Energy (DOE) Office of Nuclear Energy’s Nuclear Energy University Programs for funding of this research (Contract No. 14002 02/ Project No. 10-881). A.C. acknowledges support from the U.S. DOE, Office of Science, Heavy Elements program, DE-SC-0001815, for oversight of the computational studies. This research also used resources of the Oak Ridge Leadership Computing Facility at the Oak Ridge National Laboratory, which is supported by the Office of Science of the U.S. DOE under Contract No. DE-AC05-00OR22725.

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REFERENCES

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