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

Examination of Structure and Bonding in 10-Coordinate Europium and Americium Terpyridyl Complexes Frankie D. White,† Alyssa N. Gaiser,† Evan J. Warzecha, Joseph M. Sperling, Cristian Celis-Barros, Sahan R. Salpage, Yan Zhou, Tristan Dilbeck, Andrew J. Bretton, David S. Meeker, Kenneth G. Hanson, and Thomas E. Albrecht-Schmitt* Department of Chemistry and Biochemistry, Florida State University, Tallahassee, Florida 32306, United States Downloaded via UNIV OF LOUISIANA AT LAFAYETTE on September 30, 2018 at 00:00:17 (UTC). See https://pubs.acs.org/sharingguidelines for options on how to legitimately share published articles.

S Supporting Information *

ABSTRACT: M(TpyNO2)(NO3)3(H2O)·THF (M = La, Nd, Sm, Eu, Tb, Am; TpyNO2 = 4′-nitrophenyl terpyridyl) have been prepared from the reaction of M(NO3)3·nH2O with TpyNO2 in THF. Structural analysis shows that the metal centers are 10-coordinate, providing the first example of AmIII with this coordination number. Further spectroscopic and theoretical evaluation of these complexes reveals utilization of the 5f orbitals in bonding in the AmIII complex. Comparison of Nd−L, Eu−L, and Am−L bond distances demonstrates that some caution should be taken in comparing EuIII versus AmIII in extraction experiments.

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INTRODUCTION Bonding differences between 4f and 5f elements have been of significant interest from both fundamental and applied perspectives for more than six decades.1,2 In the former case, experimental and theoretical analyses of lanthanide and actinide coordination complexes is leading to an improved understanding of periodic trends in the f-block, particularly in regard to understanding the utilization of frontier orbitals of the metal ions in bonding.3 From an applied perspective, differences in bonding between 4f and 5f elements has led to industrial-scale separation strategies for recycling used nuclear fuel, thereby enhancing the utilization of energy resources and diminishing demands on nuclear waste repositories.2,4 Numerous neutron-capture products are generated during nuclear energy production. Among these, 241Am (t1/2 = 433 y) is produced in relatively high abundance with approximately 1.3 kg being synthesized per ton of fuel under standard burn up conditions.5 When scaled to repository levels, the large quantities of this moderately long-lived isotope create a heat burden that can be mitigated by extraction prior to entombment and subsequent fissioning in fast neutron reactors.6 Thus, in a closed nuclear fuel cycle, this kind of waste product is transformed into an energy resource. The challenge in separating 241Am from other components of nuclear waste lies in its chemical similarities with lanthanides because under typical reaction conditions it shares both comparable ionic radii and a +III oxidation state with lanthanide ions. Moreover, these lanthanides are produced in high abundance during the fissioning of 235U yielding a Ln/Am ratio of ca. 15:1 complicating the separations even further © XXXX American Chemical Society

because these lanthanides can cause adverse effects if left in fuel after recycling. To experimentally assess viable complexants that possess strong selectivity for AmIII, measuring the properties of EuIII and AmIII coordination complexes is particularly informative because the radioisotopes 152,154Eu possess desirable nuclear properties, specifically hard γ emission, for use as radiotracers in binding studies that can be compared with data obtained from 241,243Am.7−9 While EuIII does have an ionic radius ∼0.03 Å smaller than AmIII,10 these ions are still ostensibly isoelectronic.7 Accordingly, when differences in the properties of isomorphous EuIII and AmIII complexes are observed, it becomes important to determine the origins of the divergence so that these differences can be augmented in improved separations strategies and provide an improved understanding of the basic differences between 4f and 5f congeners. Furthermore, of the americium crystal structures reported, the majority contain eight- or ninecoordinate metal cations with a small number of six- or sevencoordinate complexes. The most common argument made for variances in complexation between LnIII and AnIII cations with similar ionic radii is that the latter are capable of utilizing a host of frontier orbitals in bonding that include the 5f, 6p, 6d, 7s, and 7p. However, lanthanides are largely restricted to minor involvement of 4f orbitals in the early lanthanides and more significant participation of low-lying 5d orbitals, particularly early in the series.9,11,12 Notably, our understanding of bonding Received: August 6, 2018

A

DOI: 10.1021/acs.inorgchem.8b02085 Inorg. Chem. XXXX, XXX, XXX−XXX

Article

Inorganic Chemistry

containing 5 mL of acetone. The acetone was vapor-diffused into the solution overnight producing Nd1 and Eu1 as amber colored crystals suitable for single crystal X-ray diffraction. Isostructural crystals of La, Sm, and Tb were prepared in a similar manner. Am(TpyNO2)(NO3)3(H2O)·THF (Am1). 243Am(NO3)3·xH2O was prepared in the following manner: 595 μL (containing 3 mg of Am3+) of a chloride solution of AmCl3 was evaporated to dryness in a 20 mL glass vial. The chloride salt was converted to the nitrate by dissolving in 500 μL of 6 M HNO3 and fuming to a residue. The fuming was repeated two more times. The residue was then dissolved in 1 mL of THF, which produced a pale-pink solution. A total of 4.6 mg of TpyNO2 (0.013 mmol) was dissolved in 4 mL of THF. The two solutions were mixed into a 20 mL glass vial producing an amber solution. The vial was placed into a 60 mL Nalgene bottle containing 10 mL of acetone. After vapor diffusion with acetone, pink crystals suitable for single crystal X-ray diffraction were obtained. Single Crystal X-ray Diffraction. Single crystals of each compound were placed on Mitogen mounts using Immersion Oil. The crystals were aligned with a Bruker D8 Quest X-ray diffractometer with an IμS X-ray source (Mo Kα, λ = 0.71073 Å) paired with a digital camera. Low temperature (28 K) data collection was performed utilizing an Oxford Cryostream N-Helix. The unit cells were determined with Quest software. Olex2 equipped with the SHELXTL program suite was used for structure determination.26,27 Spectroscopic Measurements. UV−vis−NIR measurements were performed using a Craic Technologies Microspectrophotometer. Single crystals were placed on a glass slide in Immersion Oil. The photoluminescence and excitation spectra were obtained on an Edinburgh FLS980 fluorescence spectrometer with a housed 450 W Xe lamp/single grating (1800 λ/mm, 250 nm blaze). Singe crystals were enclosed between glass slides and placed into a spectrometer housing, and the samples were excited with a Czerny−Turner monochromator. For the collection, the spectrum was obtained through a 465 nm filter. Detection was done with a Peltier-cooled Hamamatsu R928 photomultiplier tube. Raman measurements were performed with a Horiba JY LabRam HR800 Raman Spectrograph utilizing a TUI Optics DL 100 grating-stabilized diode laser (80 mW emitting power at 785 nm). The laser was focused onto the sample with an Olympus BX30 microscope with a 5× objective that simultaneously collects background scatter radiation. The scatter was then filtered through a Semrock RazorEdge Long Wave Pass edge filter before dispersing onto a 76 mm square 600 line/nm grating onto a 1024 × 256 element open electrode CCD detector with a 26 μm square pixel cooled at −70 °C. Computational Details. Given the nature of f-element coordination chemistry where correlation effects become important, multiconfigurational Complete Active Space Self Consistent Field (CASSCF) calculations were carried out using the Orca 4.0.1.2 package where a DFT wave function was used as a trial wave function.28 The DFT functional used was the hybrid PBE0 with the SARC-TZVP basis set for the metal centers, while the ligands were modeled using the def2-TZVPP set. The quasi-restricted orbitals (QRO) were calculated because they locate the f shell in the valence region, so no rotations are needed for the CASSCF calculation. Finally, the CASSCF wave function was obtained including the seven f orbitals in the active space by means of the second-order Douglas− Kroll−Hess (DKH2) Hamiltonian. State interactions via QDPT were used to take into account the spin−orbit coupling (SOC). Bader’s Quantum Theory of Atoms in Molecules (QTAIM) was used to analyze the bonding properties of the Eu1 and Am1 complexes. Parameters derived from QTAIM were calculated using the MultiWFN program based on the ab-initio wave function previously described.29 The most common QTAIM parameters reported are the electron density, ρ(r); the Laplacian of the electron density, ∇2ρ(r); the localization, λ(M); and delocalization δ(M−L) indexes. However, ∇2ρ(r) has been proven to be positive for felement compounds where purely covalent bonds are characterized by negative ∇2ρ(r) values. Instead, energy density parameters give more useful information regarding the nature of the interaction where if potential energy density, V(r), predominates over the Lagrangian

in the f-block has evolved past historic descriptions of lanthanide bonding as being purely ionic and actinides as being partially covalent because of advances in both experiment and theory. In particular, the use of X-ray absorption spectroscopy and the development of fully-relativistic ab initio wave function calculations have provided both a more detailed and a more quantitative understanding of dissimilarities between 4f and 5f elements. Examples of this include some of the first observations of non-negligible involvement of 4f orbitals in bonding in benchmark compounds such as the sesquioxides and hexachlorides, Ln2O3 and [LnCl6]3−,9,11 and utilization of all possible frontier orbitals in compounds like Cf[B6O8(OH)5].12 Ironically, increased covalency does not always manifest in larger bond dissociation energies or more favorable complexation enthalpies when 4f and 5f metal ions are compared.13−15 This necessitates a more holistic understanding of the dynamics of complexation and solvation of the molecules when they are in solution or the gas phase.16,17 One of the recent focuses in f-block separation chemistry has been on utilizing nitrogen-rich ligands that show discrimination between lanthanides and actinides. The most simplistic argument that has been put forward is that N-donor ligands are softer on a Pearson scale than O donors,18 and therefore show a preference for binding actinides versus lanthanides19 However, this softness effect is amplified by the energetic degeneracy of the ligand 2p orbitals and metal-based orbitals, typically the 5f, that leads to energy-degeneracy-driven covalency with a small, but non-negligible, orbital overlap bonding component.20 Many of these N-donor ligands designed for lanthanide/actinide separations contain one or more pyridine moieties. Among the simplest N-based chelators is terpyridine (terpy), which should be considered as an archetypal ligand for forming stable coordination complexes with d- and f-block metal ions. Moreover, lanthanide terpy complexes exhibit the antenna effect and enhance photoluminescence from well-known luminescent ions, especially Eu3+ and Tb3+, leading to a variety of applications.21−23 Terpy is also an ideal ligand for benchmarking and understanding bonding in actinide complexes with N-donor ligands. However, well-characterized examples of such complexes, especially with elements beyond uranium, are scarce.8,16 In order to provide insights into similarities and differences in bonding between lanthanides and actinides, specifically Eu3+ and Am3+, complexes containing a terpy derivative, 4′-nitrophenyl terpyridyl (TpyNO2), have been prepared and characterized using a variety of experimental and theoretical methods.

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EXPERIMENTAL SECTION

Synthesis. Caution! 243Am (t1/2 = 7.38 × 103 years) has potential health risks due to its α and γ emission, along with the emission of its daughter 239Np (t1/2= 2.35 days). 239Np undergoes β and γ emission. This element was handled in a Category II nuclear hazard facility. All manipulations were performed in a radiologic fume hood without exclusion of air and water. M(TpyNO2)(NO3)3(H2O)·THF (Nd1, Eu1). Nd(NO3)3·6H2O and Eu(NO3)3·6H2O were prepared by dissolving Nd2O3 or Eu2O3 (Sigma, 99.999%) in concentrated HNO3 with gentle heating and slow evaporation to dryness. 4′-Nitrophenyl terpyridyl (TpyNO2) was prepared by literature methods.24,25 A total of 12.9 mg (0.029 mmol) of Nd(NO3)3·6H2O or 13.1 mg (0.029 mmol) of Eu(NO3)3·6H2O was dissolved in 2 mL of tetrahydrofuran (THF). TpyNO2 (10.3 mg, 0.029 mmol) was dissolved in 8 mL of THF. The two solutions were mixed together in a 20 mL glass vial with Nd1 and Eu1, each resulting in an amber solution. The vial was placed into a 60 mL Nalgene bottle B

DOI: 10.1021/acs.inorgchem.8b02085 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry kinetic energy, G(r), the total energy density will be negative and a certain degree of covalency can be attributed to that bond.

diffraction experiments. The difference in bond distances between neighboring f-block elements is on the order of 0.01 Å until one reaches the curium/berkelium boundary when a discontinuity occurs.30,31 Thus, M−L bond distance esd’s of ca. 0.001(4) that would be typical of a standard resolution crystal structure (0.77 Å resolution is common) are insufficient to detect these periodic trends in a rigorous manner (i.e., within 3Δ limits). While one could argue that detecting these aforementioned trends could be easily accomplished by simply choosing metal ions that are farther apart in the series such as comparing La3+ to Lu3+ or Pu3+ to Cm3+, this subsumes that said compounds can be prepared. In fact, there are often thermodynamic, synthetic, or practical barriers that render such studies unworkable or even implausible in a number of chemical systems in which isostructural compounds cannot be obtained throughout the entire series. A number of comparisons can be made between the Eu1 and Am1 bond distances. If one considers the bonding metrics derived from diffraction data obtained at 28 K, the average M− N bond distances from the metal centers to the TpyNO2 ligand are 2.5577(8) and 2.585(2) Å for Eu1 and Am1, respectively, providing a Δ of 0.027(2) Å. Similarly, the average M−O bond distances to the nitrate anions are 2.5234(8) and 2.571(2) Å for Eu1 and Am1, respectively, yielding a Δ of 0.048(2) Å. The Δ for the M−OH2 interactions is also 0.048(2) Å with the Am−OH2 bond distances also being longer (Table 1).

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RESULTS AND DISCUSSION Structural Characterization. M(TpyNO2)(NO3)3(H2O)· THF (M = La, Nd, Sm, Eu, Tb, Am) and for the sake of brevity only Eu1, Nd1, and Am1 are discussed, and details on the rest can be found in the Supporting Information. TpyNO2 was selected for this study as it is the precursor to obtaining the −NH2 derivative, which could be utilized in photoreduction separations of Eu from Am and Cm. However, TpyNH2 has poor solubility, and single crystals were not able to be obtained. All compounds are isomorphous and crystallize in the monoclinic space group P21/c. The structure features M3+ cations bound by one tridentate TpyNO2 ligand, three bidentate nitrate anions, and one water molecule creating a 10coordinate environment around the metal centers, as depicted in Figure 1. The cocrystallized THF solvent molecule does not

Table 1. Metal Bond Lengths of Nd1, Eu1, and Am1 TpyN TpyN TpyN nitrateO nitrateO nitrateO nitrateO nitrateO nitrateO H2O average N bond length average nitrate O bond length

Nd1 (Å)

Eu1 (Å)

Am1 (Å)

2.6038(16) 2.5959(16) 2.6033(15) 2.5837(14) 2.5590(14) 2.5437(14) 2.5950(14) 2.5759(14) 2.5394(14) 2.4619(14) 2.6010(16) 2.5661(14)

2.5644(8) 2.5420(8) 2.5666(8) 2.5439(8) 2.5407(7) 2.4886(8) 2.5722(8) 2.5081(7) 2.4869(8) 2.4192(8) 2.5577(8) 2.5234(8)

2.589(2) 2.578(2) 2.589(2) 2.583(2) 2.561(2) 2.553(2) 2.604(2) 2.589(2) 2.536(2) 2.467(2) 2.585(2) 2.571(2)

Again, the ionic radius of EuIII is ca. 0.03 Å smaller than AmIII with the latter’s ionic radius comparing well with NdIII.32 This leads to the expectation that if one factors in slight contractions induced by the addition of a small covalent contribution in americium bonding that one would expect that Am−L bond lengths would be approximately equal in length to EuIII bonds. Yet, they are notably longer. However, if one compares the Am−L and Nd−L distances, expected values reemerge with an average Nd−N(TpyNO2) distance of 2.6010(16) Å, yielding a Δ of 0.016(2) Å with the Am−N bonds being shorter. Again, based on previous observations concerning An−N (An = actinide) bonds, it is the Am− N(TpyNO2) bonds that would be expected to show covalent contributions (vide infra),8,9 and if the Nd−O(nitrate) and Am−O distances are compared, the former are 0.005(2) Å shorter than those measured with americium. The Nd−O and Am−O distances to the bound water molecule differ by the same amount and in the same direction. Thus, one can conclude that there is a statistically significant shortening of the

Figure 1. Crystal structure of [Am(TpyNO2)(NO3)3(H2O)]·THF (Am1) shown with 50% ellipsoid probability. Hydrogen atoms and the cocrystallized THF molecule have been omitted for clarity.

participate in bonding to the metal center. The geometry is best approximated by a distorted bicapped square antiprism and provides the first example of a 10-coordinate AmIII complex. The ability of these crystals to diffract at higher than normal angles motivated X-ray diffraction studies at temperatures as low as 28 K in an effort to obtain high-resolution crystal structures for both complexes and to decrease the thermal contributions to the bond distances. There are a number of important results that could be derived from such studies. On the most fundamental level, the belief that monotonic decreases in lanthanides and actinide radii is an easily demonstrated periodic trend is often false. This is attributable to two features of bond distances measured from X-ray C

DOI: 10.1021/acs.inorgchem.8b02085 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry Table 2. Ionic Radii Values of Each Compound with Four Different Methods, Values Presented in Ångstroms La Ce Nd Sm Eu Tb Am

vol. unit cell (Å3)

known

shannon

nitrate-based

La-based

Ce-Based

2859.4(14) 2835.1(8) 2829.6(7) 2809.9(8) 2789.7(7) 2791.8(9) 2825.3(12)

1.27 1.25

1.239(2) 1.212(2) 1.1832(16) 1.155(3) 1.1393(8) 1.123(3) 1.181(2)

1.272(2) 1.241(2) 1.2161(16) 1.190(3) 1.1734(8) 1.162(3) 1.221(2)

1.27 1.244(2) 1.2142(16) 1.187(3) 1.1713(8) 1.155(3) 1.213(2)

1.275(2) 1.25 1.2182(16) 1.191(3) 1.1753(8) 1.159(3) 1.217(2)

Table 3. Bonding Parameters Derived from the Bader’s Theory of AIMa bond

ρ(r)

G(r)

V(r)

H(r)

|V/G|

H/ρ

Eu−Nterpy* Eu−ONO2* Eu−OH2 Am−Nterpy* Am−ONO2* Am−OH2

0.0421 0.0420 0.0447 0.0445 0.0431 0.0474

0.0413 0.0468 0.0545 0.0446 0.0486 0.0573

−0.0415 −0.0459 −0.0523 −0.0465 −0.0492 −0.0577

−0.0003 0.0009 0.0022 −0.0020 −0.0006 −0.0003

1.0067 0.9801 0.9595 1.0443 1.0132 1.0058

−0.0066 0.0222 0.0492 −0.0444 −0.0148 −0.0070

a

The electron density, the Lagrangian kinetic energy, potential energy, and energy density were evaluated at the bond critical point. All parameters are in a.u. Remaining bonds are shown in Table S2.

is largely ionic. Because of the presence of nitrate anions in the americium complex, this calculation provides the most reliable value for calculating both the unknown lanthanides and americium’s ionic radii. Additionally, the lanthanum and cerium-based calculations for Nd, Sm, Eu, Tb, and Am are all within the nitrate-based values (1.22, 1.19, 1.17, 1.16, and 1.22 Å, respectively) given their uncertainties. The slight difference between the ionic radii for Shannon’s and lanthanum/cerium-based TpyNO2 nitrogen atoms is due to the covalent character of the bonding. This covalent character significantly shortens the predicted ionic radius of nitrogen. This covalency is validated by the negative values of the TpyNO2 nitrogen atom bonds of H/ρ provided in Table 3. The neodymium and americium complexes have unit cell volumes within 4 Å3 of each other. This is expected as neodymium is often utilized as an analog for americium.9 Unfortunately, due to the solvent molecule in the structure, the volume versus i.r.3 plot is not a good verification tool for determining the ionic radii of the unknown lanthanide and actinide 10-coordinate centers. The volume of the europium compound is smaller than that of the terbium compound; however, the ionic radius of europium is calculated to be larger than that of terbium. Placherel et al. previously calculated the ionic radius of 10-coordinate europium to be 1.21 Å.33 When recalculating with the current Shannon ionic radii, the ionic radius for this reported structure is 1.176(4) Å. The reported calculation used a currently accurate value of 1.46 Å for the nitrogen atoms; however, 1.31 Å was used for oxygen instead of 1.35 Å. Using 1.35 Å for the nitrate oxygen atoms and 1.36 Å for the crown-ether oxygen atoms gives a Eu3+ ionic radius of 1.176(4) Å. This number is within the uncertainties of the calculated ionic radius for the 10-coordinate europium complex presented in this work. One example of a previously reported lanthanum complex containing a 10-coordinate La III center is La(terpy)(NO3)2(acac).34 When calculating the ionic radius of lanthanum using Shannon’s method, the ionic radius is determined to be 1.24 Å. Conversely, when calculating the ionic radius of lanthanum using the previously discussed lanthanum values for terpy, nitrate, and water (treating the

Am−N(tpyNO2) bonds relative to those observed with neodymium, but that the other bonds are essentially equivalent. There are additional conclusions that one can draw from these studies. First, that the decrease in bond lengths between NdIII and EuIII is not the expected range with a Δ of 0.0433(15) Å instead of ∼0.03 Å. Second, the Eu−L distances are shorter than expected. Finally, by extension of this line of reasoning, comparisons of solvent extraction data between 152,154 Eu and 241,243Am need to be treated with a great deal of caution because the differences in effective ionic radii may be larger (or smaller) than one anticipates based on simple extrapolations. Ionic Radii Determination. The ionic radius of 10coordinate AmIII is unknown. To determine this, X-ray diffraction data from the isomorphous M(TpyNO2 )(NO3)3(H2O)·THF (M = La, Nd, Sm, Eu, Am) series were utilized. Determining ionic radii using Shannon’s method involves using a known cationic or anionic radius and bond lengths to calculate the unknown radii based on the assumption that a bond length should be equal to the sum of a cationic radius and an anionic radius. For the complexes presented in this paper, covalency must be accounted for.32 Using Shannon’s ionic radii to calculate the metal centers was not the best method due to the degree of covalency in the terpy nitrogen bonds and water bonds. Four different calculations were used: Shannon’s ionic radii, radii calculated using known the lanthanum 10-coordinate radius, (1.27 Å), cerium 10-coordinate radius (1.25 Å), and radii calculated using only the nitrate bonds (Table 2). Back-calculating the average ionic radii for the nitrate oxygen atoms, terpy nitrogen atoms, and water oxygen atom based on the known lanthanum ionic radius, 1.27 Å, gives the expected ionic radius of 1.35 Å for the ionically bound nitrate oxygen atoms. Conversely, the terpyridine nitrogen atoms give an ionic radius of 1.39 Å that is quite contracted from the 1.46 Å value given in the Shannon tables. The water oxygen atom yields 1.25 Å, reflecting the degree of covalency observed. The nitrate bond-based calculations are the most accurate because the nitrate anion interaction with the cationic centers D

DOI: 10.1021/acs.inorgchem.8b02085 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry

Theoretical Examination of Bonding. A useful theoretical tool for addressing bonding is the Quantum Theory of Atoms in Molecules (QTAIM), and its parameters derived from Bader’s theory.36 These parameters have been used to understand the nature of interactions between metal ions and ligands in the f-block. Table 3 summarizes the main QTAIM parameters derived from the electron density, ρ(r), such as the Lagrangian kinetic energy, G(r), the potential energy density, V(r), and the energy density, H(r). Two useful ratios are also shown in the table, the balance of the energy density (V/G) and the normalized total energy density (H/ρ). Similarities between electron densities are in agreement with the similar bond lengths observed experimentally in the LnIII and AmIII TpyNO2 complexes. However, energies corresponding to those densities denote striking differences between the two complexes. When |V/G| > 1, the electron density is dominated by potential energy making H(r) < 0 and characterizing the interaction as covalent. The degree of covalency is analyzed in terms of normalized H(r) where more negative values denote more covalent electrons. In this sense, bonds in Am1 are characterized by weak covalent bonds according to the increasing degree of covalency of Am−OH2 < Am−ONO2 < Am−Nterpy. In contrast, Eu1 interactions are dominated mainly by purely ionic bonds displaying some degree of covalency in only one of the 10 bonds. This can be better explained in terms of the role of the f shell over the total energy density value. The difference observed between Eu1 and Am1 bonds is clear; the 4f contribution to the total energy is always positive while the 5f contribution is always negative. This means that the observed partial covalency in Am1 is due to the stabilizing energy produced by the 5f electrons to the total energy density.

acac oxygen atoms similar to the nitrate oxygen atoms), the ionic radius is the known value, 1.27 Å. Therefore, one concludes again the ionic radius of 10-coordinate AmIII using the value obtained from nitrate complexation is the most reliable, and these values are presented in Table 2. Optical and Vibrational Properties. The photoluminescence, absorption, and Raman spectra of Eu1 and Am1 were measured at room temperature. The excitation spectrum of Eu1 exhibits a broadband feature from ∼300 to 420 nm attributed to intraligand transitions of the TpyNO2 ligand. When the complex is excited at 365 nm, sharp photoluminescence bands are observed as shown in Figure 2 that are

Figure 2. Excitation (592 nm) and emission (365 nm) spectra of Eu1.

assigned to the 7F4−7F0 transitions of EuIII. In short, this is a classic case of the so-called antenna effect where the TpyNO2 ligand is absorbing short wavelengths and transferring energy to intra-f excitations. Excitation of the ligand overlaps with EuIII’s excitation bands 395 and 465 nm. The optical properties of Am1 are consistent with other AmIII compounds that possess bound water molecules in that it decays in a nonradiative manner. While AmIII compounds can exhibit photoluminescence at near 650 nm, this energy is close enough to that of the overtones of water that energy loss becomes entirely nonradiative.35 The absorbance of each compound was also measured on single crystals. The absorbance spectrum further confirms the obscuring of the EuIII excitation peaks, and the normal sharp absorption bands are not observed. In the Am1 spectrum, the sharp features are observed at 502 nm corresponding to the 7F0 → 5L6 transition as shown in Figure 3.

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CONCLUSIONS In summary, a series of lanthanide and americium terpyridine coordination complexes have been prepared and characterized. These compounds reveal 10-coordinate metal sites, and the first example of this coordination number with AmIII. The ionic radii have been determined for 10-coordinate EuIII, NdIII, and AmIII to be 1.17 Å, 1.22 Å, and 1.22 Å, respectively. The closest analog of AmIII in the lanthanide series is NdIII, and the Am−N bonds have been shown to be contracted relative to the Nd−N bonds. Great care has been taken to collect high-resolution Xray diffraction data sets to reduce the esd’s on the bond distances to provide more statistically significant comparisons between bond lengths. QTAIM analyses of the bonding demonstrates a small degree of covalency in the Am−N bonds that is substantially diminished in the EuIII complex.

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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.8b02085. Experimental details for synthesis, luminescence, absorption, and Raman measurements (PDF) Accession Codes

CCDC 1854046−1854051, 1857533−1857535, and 1857537 contain the supplementary crystallographic data for this paper. These data can be obtained free of charge via www.ccdc.cam.ac.uk/data_request/cif, or by emailing data_request@ccdc. cam.ac.uk, or by contacting The Cambridge Crystallographic

Figure 3. Absorption spectra of Eu1 and Am1. E

DOI: 10.1021/acs.inorgchem.8b02085 Inorg. Chem. XXXX, XXX, XXX−XXX

Article

Inorganic Chemistry

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

Frankie D. White: 0000-0002-8644-1231 Evan J. Warzecha: 0000-0001-7007-9250 Cristian Celis-Barros: 0000-0002-4685-5229 Yan Zhou: 0000-0002-7290-1401 Kenneth G. Hanson: 0000-0001-7219-7808 Thomas E. Albrecht-Schmitt: 0000-0002-2989-3311 Author Contributions

The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript. Author Contributions †

These authors contributed equally.

Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the Department of Energy, Heavy Elements Chemistry Program under Award Number DEFG02-13ER16414. We would also like to thank Jason Johnson and Ashley Gray at Florida State University for their role in radiation work.

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REFERENCES

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DOI: 10.1021/acs.inorgchem.8b02085 Inorg. Chem. XXXX, XXX, XXX−XXX

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DOI: 10.1021/acs.inorgchem.8b02085 Inorg. Chem. XXXX, XXX, XXX−XXX