Redox and Coordination Chemistry Differences of the 4F and 5F Elements

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2018

Redox and Coordination Chemistry DFrainfkfiee Dr. Wehniteces of the 4f and 5f Elements

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COLLEGE OF ARTS AND SCIENCES

REDOX AND COORDINATION CHEMISTRY DIFFERENCES OF THE 4f AND 5f

ELEMENTS

FRANKIE D. WHITE

A Dissertation submitted to the Department of Chemistry and Biochemistry in partial fulfillment of the requirements for the degree of Doctor of Philosophy

2018 Frankie D. White defended this dissertation on September 28, 2018. The members of the supervisory committee were:

Thomas E. Albrecht-Schmitt Professor Directing Dissertation

Samuel L. Tabor University Representative

Kenneth G. Hanson Committee Member

Yan-Yan Hu Committee Member

The Graduate School has verified and approved the above-named committee members and certifies that the dissertation has been approved in accordance with university requirements.

This work is dedicated to God, my parents, Frankie D. White Sr. and Pamella Potter White, my family, and supporting friends. I am extremely grateful for all the love and support you have given me and would not be here without you.

iii ACKNOWLEDGMENTS

First, I would like to thank God, who provided me with the abilities and gift to pursue my goals in life. Secondly, I would like to thank my parents. You two have provided me with everything that I ever needed. You continuously have placed me before yourselves in every aspect to make sure I was able to strive to be the best I could be. You constantly reminded me to keep my priorities in order. Thank you. To my sisters and brothers, Kia, Ty, Trey, and Calvin, thank you for all the love and support you have given me throughout all my academic years. I would also like to thank my friends Amanda, Elijah, and Jeff. You always supported my decisions to obtain a higher education. I tremendously thank my two academic mentors Dr. Richard Sykora and Dr. Thomas Albrecht-Schmitt. Dr. Sykora, I spent the first part half of my research years under you at the University of South Alabama. You were a large part of my decision to obtain a doctorate in chemistry and I am forever thankful for your presence in my life. Additionally, you introduced me to Tom, my current mentor who I also want to thank. Tom, you have had a substantial impact on my life and have always been supportive of me. You challenged me tremendously, but also had the same amount of faith in me. Thank you both. I would also like to thank all of the members of the Albrecht-Schmitt group that I had the pleasure of working. Samantha Cary, you got me started in the lab and had a major impact on my beginnings. I would also like to thank all of the following: Teresa Eaton, Shane Galley, Ali Arico, Kevin Seidler, Matt Marsh, Sasha Chemey, Wes Potter, David Dan, Alyssa Gaiser, Evan Warzecha, Ryan Greer, Joe Sperling, Carla McKinley Bonnie Klamm, David Meeker, Brian Long, Joanna Campbell, Vanessa Proust, Renaud Jeanine, Cory Windorff, and Ashini Jayasinghe. In addition to these research members, I would also like to thank the Alabugin group for constantly letting me borrow materials and the Hanson group for performing all my spectroscopic studies that I needed, especially Yan Zhou and Tristan Dilbeck. Lastly, I would like to thank Florida State University and my committee for extending my educational knowledge and for allowing me to be a graduate student and researcher in their facilities. I also would like to thank the Department of Energy Grant under the award number of DE-FG02-13ER16414 for providing the funding for my research.

iv TABLE OF CONTENTS

List of Tables ...... viii List of Figures ...... x Abstract ...... xiii

1. INTRODUCTION ...... 1

1.1 History of f-elements...... 1 1.2 Problems in f-elements ...... 2 1.3 Chemistry of f-elements ...... 3 1.3.1 The Characteristics of 4f and 5f-elements ...... 4 1.3.2 Trends in the f-elements ...... 5 1.3.3 Dissertation Objectives ...... 7 1.4 Figures...... 7

2. SYNTHESIS AND CHARACTERIZATION METHODS ...... 10

2.1 Past and Current Methods for f-element Crystal Growth ...... 10 2.2 Characterization Techniques ...... 10 2.2.1 Single Crystal X-ray Diffraction ...... 10 2.2.2 Spectroscopic Measurements ...... 11 2.2.3 Computational Details ...... 12 2.2.4 Electrochemical Studies ...... 12 2.2.5 Syntheses of Reactions ...... 13 2.2.6 Recycling of 243Am and 249Cf ...... 13 2.3 Figures...... 14

3. EXAMINATION OF STRUCTURE AND BONDING IN 10-COORDINATE EUROPIUM AND AMERICIUM TERPYRIDYL COMPLEXES ...... 15

3.1 Introduction ...... 15 3.2 Results and Discussion ...... 17 3.2.1 Structural Characterization ...... 17 3.2.2 Ionic Radii Calculations ...... 19 3.2.3 Spectroscopic Measurements ...... 20 3.2.4 Computational Analysis ...... 21 3.3 Conclusion ...... 21 3.4 Methods and Characterizations ...... 22 3.4.1 Syntheses ...... 22 3.4.2 Single Crystal X-ray Diffraction ...... 23 3.4.3 Computational Details ...... 23 3.5 Figures...... 24

v 4. STRUCTURAL, PHOTOPHYSICAL, AND ELECTROCHEMICAL INVESTIGATIONS OF F-ELEMENTS IN NON-AQUEOUS CONDITIONS ...... 26

4.1 Introduction ...... 26 4.2 Results and Discussion ...... 27 4.2.1 Structural Characterization ...... 27 4.2.2 Spectroscopic Studies ...... 28 4.2.3 Electrochemical Studies ...... 29 4.3 Conclusion ...... 30 4.4 Methods and Characterization ...... 31 4.4.1 Syntheses...... 31 4.4.2 Single Crystal Crystallography ...... 32 4.4.3 Spectroscopic Studies ...... 32 4.4.4 Electrochemical Studies ...... 32 4.5 Figures ...... 33

5. TUNING THE ENERGETICS OF SAMARIUM(II): MOLECULAR AND ELECTRONIC STRUCTURE, AND HYDROLYTIC REACTIVITY OF A SAMARIUM(II) CROWN ETHER COMPLEX ...... 40

5.1 Introduction ...... 40 5.2 Results and Discussion ...... 41 5.2.1 Syntheses...... 41 5.2.2 Structural Features ...... 42 5.2.3 Spectroscopic Properties ...... 43 5.2.4 Broken-Symmetry (DFT) Calculations ...... 43 5.2.5 Ab-initio Calculations ...... 44 5.2.6 Time-Dependent DFT ...... 46 5.3 Conclusion ...... 46 5.4 Methods and Characterizations ...... 47 5.4.1 Experimental Syntheses ...... 47 5.4.2 Crystallographic Studies ...... 48 5.4.3 UV-Vis-NIR, Excitation, and Photoluminescence Spectroscopy...... 48 5.4.4 Computational Details ...... 49 5.5 Figures...... 50

6. EXAMINING THE SECOND BREAK IN THE 5F SERIES: ADVANCEMENTS TOWARDS MOLECULAR CF(II) ...... 56

6.1 Introduction ...... 56 6.2 Results and Discussion ...... 57 6.2.1 Selection of Starting Materials ...... 57 II * 6.2.2 Facile Synthesis of [Sm (DB30C10)][BPh4]2 and SmTp2 ...... 58 6.2.3 Reaction of 249Cf with DB30C10 and Tp* ...... 59 6.2.4 Spectroscopic Measurements ...... 60 6.3 Conclusion ...... 61

vi 6.4 Methods and Characterizations ...... 61 6.4.1 Syntheses ...... 61 6.4.2 Single Crystal X-ray Crystallography ...... 63 6.4.3 UV-Vis and Luminescence Measurements ...... 63 6.5 Figures...... 63

7. CONCLUSION ...... 66

APPENDICES ...... 69

A. TABLES FROM CHAPTER 1 ...... 69

B. TABLES FROM CHAPTER 3 ...... 70

C. TABLES FROM CHAPTER 4 ...... 72

D. TABLES FROM CHAPTER 5 ...... 73

E. CRYSTALLOGRAPHIC DATA ...... 78

References ...... 89

Biographical Sketch ...... 101

vii LIST OF TABLES

Table A.1. Breakdown of heavy metal single crystal structures curently in the CSD as of 2018 69

Table B.1. Values for ligands used in ionic radii calculations. Values shown are in angstroms . 70

Table B.2. Ionic radii of 10-coordinate f-elements by different calculation methods. Values are in units of angstroms...... 70

Table B.3. Bonding parameters derived from the Bader’s theory of AIM. The electron density, the Lagrangian kinetic energy, potential energy, energy density, were evaluated in the bond critical point. All parameters are a.u...... 70

Table B.4. Bonding parameters derived from the Bader’s theory of AIM for the metal-ligand BCPs. The electron density, the Langragian kinetic energy, potential energy, and energy density were evaluated in the bond critical point. All parameters are a.u...... 71

Table C.1. Values from the cyclic voltammetry experiments involving Ln/AnCrypt...... 72

Table D.1. Calculated excitation spectrum for both Sm(III) centers in Sm2(DB30C10)(OH)2I4 .73

Table D.2. Ab-initio computed crystal-field parameters that show the axial character of the III dinuclear Sm complex, Sm2(DB30C10)(OH)2I4...... 74

Table D.3. Calculated excitation spectrum for [Sm(DB30C10)][BPh4]2...... 75

Table E.1. Crystal data and structure refinement for Nd1 ...... 78

Table E.2. Bond lengths for Nd1 ...... 78

Table E.3. Crystallographic details of Eu1...... 79

Table E.4. Bond lengths for Eu1 ...... 80

Table E.5. Crystal data and structure refinement for Am1 ...... 81

Table E.6. Bond lengths for Am1...... 82

Table E.7. Crystal data and structure refinement for SmDBC ...... 82

Table E.8. Bond lengths for SmDBC ...... 83

Table E.9. Crystal data and structure refinement for Sm2(DB30C10)I4(OH)2...... 85

Table E.10. Selected bond lengths for Sm2(DB30C10)I4(OH)2...... 86

viii II Table E.11. Crystallographic details of [Sm (DB30C10)][BPh4]2...... 87

II Table E.12. Selected bond lengths for [Sm (DB30C10)][BPh4]2 ...... 88

ix LIST OF FIGURES

Figure 1.1 Schematic of the PUREX process used to extract uranium and plutonium from spent nuclear fuel12 ...... 7

Figure 1.2 Breakdown of crystal structures in the Cambridge Structural Database by Metal Group. This demonstrates how little the f-elements research is performed in contrast to the transition metals ...... 8

Figure 1.3 Radial extension of Sm3+ compared to Pu3+. The radial extension of plutonium extends beyond the core allowing the 5f orbitals to have more participation in bonding. Adapted from Clark.23 ...... 8

Figure 1.4 Reaction scheme utilized to obtain all the lanthanides in the divalent oxidation states.24 ...... 9

Figure 1.5 The III-II redox potentials of the lanthanides and actinides. Other than the noted lanthanides, Sm, Eu, and Yb, the +2-oxidation state is relatively difficult to obtain. In the actinides, the +2-oxidation state becomes increasingly stable...... 9

249 Figure 2.1 Purification scheme of Cf. The precipitation of CfF3 allows for a greater confidence of purification as many other metal fluorides are soluble ...... 14

Figure 3.1 Crystal structure of [Am(TpyNO2)(NO3)3(H2O)]∙THF (Am1) shown with the 50% ellipsoid probability. Hydrogen atoms and the co-crystallized THF molecule have been omitted for clarity...... 24

Figure 3.2 Excitation and photoluminescence spectra of [Eu(TpyNO2)(NO3)3(H2O)]∙THF (Eu1). The sharp f-f transitions of Eu3+ are displayed...... 25

Figure 3.3 Absorption spectra of [Eu(TpyNO2)(NO3)3(H2O)]∙THF (Eu1) and [Am(TpyNO2)(NO3)3(H2O)∙THF (Am1) ...... 25

Figure 4.1 The crystal structure of SmDBC shown with the 50% thermal ellipsoids. The structure contains three samarium sites. Two sites show the ligand around the metal, while the third samarium site is surrounded by five nitrates in order to charge balance the structure. Hydrogen bonding is displayed to show how further stabilization is achieved through intramolecular interactions...... 33

Figure 4.2 The encapsulation of Sm with 2.2.2-cryptand is presented. The metal center in the LnCrypt structures have the same geometry resulting in a 9-coordinate metal site. The structure is shown with the 50% thermal ellipsoid probability (hydrogens omitted). The metal is shown in green, oxygen in red, nitrogen in blue, and carbon in grey ...... 34

x Figure 4.3 The luminescence of SmDBC shows the f-f transitions commonly associated with trivalent lanthanides. The molecule does not have efficient energy transfer. When excited at 320 nm, the Sm3+ peaks are not observed as in compounds that exhibit antenna effects……………………...34

Figure 4.4 Emission spectra of samarium before and after reduction. Before reducing to Sm2+, the f-f transitions associated with Sm3+ are visible. After reduction, the broad band emission is characteristic of f-d transitions which are allowed ...... 35

Figure 4.5 The absorbance of [Sm(2.2.2-cryptand)(THF)][BPh4]2 over the course of one hour. As the compound begins to decompose, it becomes amorphous and loses its properties...... 35

Figure 4.6 Absorbance spectrum of Yb(BPh4)2 as a solid crystalline sample. Like the LnCrypt structures, the compound possesses the divalent broad spectroscopic properties ...... 36

Figure 4.7 The absorbance and emission properties of [Eu(2.2.2-cryptand)(THF)][BPh4]2. Unlike the other divalent lanthanides studied in this work, EuCrypt has fluorescent properties in the blue region...... 36

Figure 4.8 The cyclic voltammagram of Sm2+ in DMF with and without 2.2.2-cryptand. This demonstrates the effects ligands can play on the redox potential of f-elements. When 2.2.2- cryptand is present in solution of divalent lanthanides, the redox potential can shift drastically to more favorable conditions...... 37

Figure 4.9 CV of various crown ethers with Sm3+ in acetonitrile. The CV’s show that the ligand has a small effect on samarium. However, it was observed that 2.2.2-cryptand gave the best response for studying the lanthanides and actinides ...... 37

Figure 4.10 The CV’s of Sm, Eu, Yb, and Cf in the presence of 2.2.2-cryptand. The spectra show that europium has the most favorable redox potential due to the stabilization of the half- filled f-orbital. Californium, however, was shown to have an easier path to stabilize the divalent oxidation state than samarium...... 38

Figure 4.11 The color change exhibited by divalent Sm, Eu, and Yb (left to right) when complexed by 2.2.2-cryptand in solution ...... 38

Figure 4.12 Precipitation of AgI from the reaction of CfI3 with silver triflate. The mother liquor was removed from the solid and was used to obtain the first cyclic voltammogram of californium in non-aqueous solutions...... 39

Figure 5.1 A view of the structure of the [Sm(DB30C10)]2+ cation in II II [Sm (DB30C10)][BPh4]2∙2CH3CN. The Sm cation is bound by all ten etheric oxygen atoms from the dibenzon-30-crown-10 ligand creating a sphenocorona geometry around the metal center. 50% probability ellipsoids are depicted ...... 50

Figure 5.2 An illustration of the structure of the dinuclear, SmIII dibenzo-30-crown-10 complex, III Sm2(DB30C10)I4(OH)2. The Sm cations are bridging by μ-OH anions that form hydrogen

xi bonds with the crown ether. Two trans iodide anions complete the coordination sphere. 50% probability ellipsoids are depicted...... 51

Figure 5.3 Depictions of the conformations of dibenzo-30-crown-10 in solid state. The top figure (a) shows the form adopted by the crown ether without a metal ion. The second conformation (b) is adopted with alkali metal cations such as Na+ and K+. The final Pac-Man II conformation (c) is found with Sm in [Sm(DB30C10)][BPh4]2∙2CH3CN ...... 51

Figure 5.4 Absorption (left at 298 K), and excitation and photoluminescence (right at 77 K) II spectra of [Sm (DB30C10)][BPh4]2∙2CH3CN. The absorption spectrum displays broad-band peaks indicative of 4f →5d transitions; whereas the excitation and photoluminescence spectra obtained at 77 K reveal characteristic fine 4f → 4f transitions. The photoluminescence spectrum 5 7 shows the D0 → FJ (J = 0, 1, 2, 3) transitions ...... 52

Figure 5.5 Spin Density representation for the High-Spin (HS) and Broken-Symmetry (BS) states ...... 52

Figure 5.6 Energy diagram showing the lanthanide states (black horizontal lines) and the interval of energy in which the f → d transitions appears (green square) ...... 53

II Figure 5.7 Calculated TDDFT absorption spectrum of [Sm (DB30C10)][BPh4]2...... 54

II Figure 5.8 Types of orbitals involved in excitation of [Sm (DB30C10)][BPh4]2 ...... 54

Figure 5.9 Reaction scheme for forming the Sm(II) and Sm(III) products. The presence of water causes the oxidation of the Sm(II) site to form the dinuclear product ...... 55

Figure 6.1 SmII structures deemed suitable for Cf. The structures have interesting crystallographic properties that would be ideal to study with Cf. The syntheses of the two structures are relatively simple to synthesize compared to other divalent f-element compounds .63

Figure 6.2 The UV-Vis and emission spectra (inlet) of the californium product obtained from the reaction with DB30C10 ...... 64

Figure 6.3 Emission from the amorphous Cf product at 420 nm (left) and 365 nm (right) ...... 64

2+ * 2+ Figure 6.4 (a) Emission of Cf :CaTp2 . The small amount of dopant of Cf can give a red shift. The luminescence of Cf2+ could possibly be tunable like Eu2+...... 65

Figure 6.5 Calculated TDDTF absorption spectrum of divalent californium within DB30C10. .65

xii ABSTRACT

This dissertation seeks to determine the differences in the lanthanides and later actinides in non-aqueous media. Research in the f-elements is significantly understudied compared to the other metals of the periodic table. Even more so are the later actinides which were largely unstudied for an extended period as it was believed later actinides were identical to lanthanides. A review by Neidig et al, “The Covalency in f-element Complexes” has ignited significant interest in the bonding of the actinides.1 A tremendous amount of research in the f-elements, particularly the actinides, has been performed in aqueous conditions at high temperatures and pressures. Chemistry under these conditions limit the research possible for lower oxidation states. Additionally, non-aqueous techniques allow for the investigation of these elements in more organic environments. The goal of this work is to pave a greater understanding of knowledge for lanthanides and actinides by examining their redox and coordination chemistries in these environments that could lead to applications other than nuclear energy and weapons. The first portion of this dissertation examines the chemistry that is already heavily acknowledged about f-elements: coordination chemistry. When modeling later actinides, a common notion is to utilize the isoelectronic lanthanide as the surrogate. Although for electronic comparisons this is useful, it is often not the case for examining isostructural compounds. The isoelectronic lanthanide is often smaller in ionic radius, which is a factor that dominates the chemistry of the lanthanides. Despite this, isolation of isostructural coordination compounds was obtained for the isoelectronic and size analogs of americium; europium and neodymium. This seemingly mundane study showed that americium portrays a small amount of covalency in its bonds which is not observed in the lanthanides. These small differences lead to profound changes in chemical properties as observed later in this work. The second portion of this work focuses on analyzing the divalent oxidation state of f- elements with crown ethers. The divalent oxidation state has been obtained for all lanthanides using potassium and 2.2.2-cryptand. The next step was to determine the extent to which crown ethers and solvents have on the redox properties of f-elements. Because all the lanthanides had been obtained in the divalent oxidation state in a similar matter, it was expected that the redox chemistries would behave identically. To surprise, ytterbium behaves differently and shows greater reversibility than the most stable divalent lanthanide, europium. Additionally, it was

xiii found that californium also behaves like ytterbium electrochemically, even though it would be expected to behave like samarium. It was proposed that this may be attributed to the 5f orbitals. The last of this work involves obtaining californium in the divalent oxidation state as a molecular system. This was done by modeling with samarium which is the most similar to californium in its redox and coordination properties. Quick and simple routes to synthesizing divalent samarium structures were obtained in ordinary glovebox conditions for attempts with californium. Under identical reaction conditions, isolation of Cf(II) crystals in the solid state were unsuccessful. However, interesting spectroscopic properties where observed that portrayed divalent californium as having tunable luminescence similar to divalent europium compounds. To our surprise, even though samarium resembles californium, the chemistry between the two elements are very different, further broadening the gap between the 4f and 5f elements.

xiv CHAPTER 1

INTRODUCTION

1.1 History of f elements

The chemistry of the 4f and 5f elements have been severely understudied in comparison to the rest of the periodic table. These reasons vary from the lack of reactivity, to the cost of materials to the supposed identical properties between the two series. However, these elements have found their way to becoming essential in everyday matters. The amount of applications for the 4f elements are numerous and inlcude cell phones, batteries, LEDs, and permanent magnets.2,3 Although the 5f elements have far fewer implementations in society, there have been vital applications discovered with these radioactive elements. These include, but are not limited to, energy sources, space exploration, and cancer treatment.4–6 The 4f block of ther periodic table consists of the elements lanthanum through lutetium. The first of these elements was discovered in the town of Ytterby, Sweden in 1787.7 All of these elements are naturally occurring, except for Promethium, which is the only lathanide which exists as a radioactive isotope. In contrast, the radioactive 5f block, which spans from actinium to lawrencium, has quite a difference in how it was discovered. Uranium and thorium have been known the longest since their discovery in the late 1700’s and early 1800’s. They were found in Pitchblende, a mineral found in the mines of the Czech Replubic.7 The actinides beyond uranium, are known as the transuranics (Np-Lr) and these elements are all man-made and weren’t discovered until 1939 thru 1961. The discovery of these compounds was accomplished by bombarding a uranium-238 target with neutrons. Once uranium-239 forms from neutron capture, it decays via β-emission to form neptunium-239.7 In a similar manner, the rest of the actinides were obtained. The impact of these radioactive materials were felt shortly after their discovery. Scientist quickly found use of the energy coming from some of the actinides isotopes. In particular uranium-235 and plutonium-239, became of significant interest during World War II. These isotopes were known for the capability of fissioning, ultimately leading to the discovery of the nuclear bomb. After the devastation of the nuclear bombs dropped on Hiroshima and Nagasaki, scientist searched for more positive ways to use the elements. One such way included using the

energy from these isotopes to provide power to cities. This lead to the development of nuclear power plants. The first of these plants to be incoporated into a power grid was the USSR’s Obninsk Nuclear Power Plant in 1954.8 Following this, other ways of incorporating these actinides into society involved using californium-252 as a medical isotope for treating women with terminal stages of cervical cancer and utilizing amercium-241 in smoke detectors.9,10 However, with the creation of these radiolytic applications came the problem of waste management.

1.2 Problems in f-elements The 4f lanthanides have not posed as much of a waste problem as their 5f counterparts. In general, this is due to the lack of radioactivity in lanthanides. Therefore, much of the problems associated with wasted involves the actinides. This waste problem has become more of an issue due to the maintenance and care of nuclear weapons and power plants. Over time, and after multiple incidents involving radioactive materials being released into the environement, certain measures have been taken to improve the fate of nuclear waste. Despite some of these problems, nuclear energy still serves as one of the significant and reliable sources of energy. In fact, some countries rely on nuclear energy as the primary source of energy for delivering power to the grids of its cities.11 The handling of nuclear waste continues to be an ongoing problem for radioactive materials. The United States reported that 45,000 metric tons of nuclear waste was being stored within the country and that estimate would increase by approximately 2,000 metric tons each year.12 A large problem associated with this waste is the amount of time it takes before some of these elements decay into isotopes that are deemed safe. Therefore, significant studies are needed to determine the chemistry of these elements while they are in their radioactive form to better understand how these elements behave in certain conditions. The large area of research in 5f elements is based on sepeartions of long-lived isotopes from short-lived isotopes and recyclable isotopes from non-recyclable isotopes. There have been many established methods for trying to achieve these seperations. The first, and one of the more notable processes is the Plutonium Uranium Recovery by Extraction, better known as the “PUREX” process. This process involves employing an organic solvent that is 30% tributyl phosphate to extract uranium and plutonium as solid complexes (Figure 1.1).13 This process

takes care of the activity from plutonium and uranium, however, the remaining activity stems from the transuranics and short lived lanthanide isotopes. The process largely known for this secondary seperation is the Trans Uranic Extraction (TRUEX).14 TRUEX seperations employ octyl(phenyl)-N,N-diiso-butylcarbamoylemetliylphosphine oxide in the PUREX solvent. Due to its implementation into the PUREX process, this sepeartion has become ideal for minimizing transuranic waste from spent nuclear fuel. Recently, the difficulty in seperations comes from the similarities between amercium and curium. This is attributed to their +3 oxidation state and near identical coordination chemistry. A few experiments have shown the seperation of these two elements, but only on small scales.15,16 The larger problem that lies with actinides is the amount of research performed on these elements. This is perhaps due to the idea that actinides behave identically to the lanthanides. However, even if lanthanide and actinide research were to be combined, it is magnitudes less than research performed on transition metals. By examining the Cambridge Structural Database (CSD), one can readily observed the extent of wich lanthanide and actinide chemistry is studied. Figure 2.1 and Table A.1 show the breakdown of crystal structures by orbital group and elements respectively. The combination of both the 4f and 5f series takes up approximately 8.5% of metal crystal structures known. In comparison, platinum, with a price per gram higher than every lanthanide, takes up approximately 4% of all metal crystal structures in the CSD. It is this lack of research that has hindered the growth of chemistry in the f-elements.

1.3 Chemistry of f-elements The depth of research performed on transition metals have paved the way for tremendous applications and vital uses. In a similar manner, this kind of depth of knowledge has had similar impacts on the main group metals. Therefrore, the greater understanding we have about an element, the more impactful applications can be obtained. For example, iron is a well- characterized metal that continues to receive a monumental amount of research, and because of this, there is a very deep understanding of iron. However, it was not until the combination of iron

with neodymnium, that one of the strongest permanent magnets came into existance, Nd2Fe14B. Once the chemistry of lanthanides and actinides are as well-characterized as many of the transition metals, the applications of these metals could be endless.

1.3.1 The Characteristics of 4f and 5f-elements

There are numerous kinds of bonds present in chemistry. However, lanthanides tend to be dominated by one particular type of bonding: cooridination bonding. In general terms, this is due to the ionic character observed in the 4f orbitals. The f-orbitals of lanthanides are very shielded by the xenon core, 5s, and 5p orbitals, and therfore have very little participation in stronger types of bonding. Given this ionic-like character, lanthanides have relatively large ionic radii and can take on large coordination numbers. Interestingly, the large variety in coordiantion numbers can have an extenisve effect on the ionic radii of the metal ions.17 This ionic radii effect is the main factor in the coordination chemistry of lanthanides. The bonding of actinides is more diverse than lanthanides. In the 5f series, the most notable difference from that of the 4f series is the actinyl group. This is most commonly seen in x+ uranium, neptunium, and plutonium, where the metal takes the form AnO2 . Unlike the bonds observed in lanthanides, this bond is very strong and has covalent character. It was often debated whether the Ac—O bond was better viewed as a double bond or triple bond, and the latter thought of these bonds is becoming more accepted.18 This bonding character makes these actinides unique from the lanthanides. Another notable property that makes some actinides differ from lanthanides is the array of oxidation states observed, especially in the beginning of the series. The +4 through +6 oxidation state is seen from thorium to americium. Uranium, neptunium, and plutonium can be obtained in just as many oxidation states as some transition metals by occupying oxidation states from +2 through +7.19–22 What is even more unique about actinide bonding is their radial extension from the core. And Figure 3.1 clearly shows this difference. In the radial extension plots of Sm3+ and Pu3+, the orbital behavior looks very similar. However, at the 1 Å mark, the Pu3+ 5f orbital is extended beyond the core unlike Sm3+. In the supposed analogs of uranium and neptunium, neodymium and promethium respectively, there is also a ground state configuration difference. Neodymium and promethium take on the expected ground state configurations of 4f4 and 4f5, while uranium and neptunium have ground state configurations of 5f36d1 and 5f46d1.23 These differences have made these few actinide elements to be of significant interest. The other actinides, notably actinium and americium through californium, have been largely unstudied. This is perhaps due to Glenn T. Seaborg’s Actinide Concept. The lanthanides core-like orbitals make the +3 oxidation state dominant across the entire series, with a few 4

having +2 and +4 oxidations states reasonably achievable. In the actinides mentioned, this +3 oxidation state is also predominant. This lead to the idea that these later actinides begin to behave identically to lanthanides and therefore, were not necessary to study given their radioactivity, “uninteresting” chemistry, and cost.

1.3.2 Trends in the f-elements The entirety of the lanthanides behave similarly in their chemistry. As previously mentioned, the +3 oxidation state dominates the series with a few lanthanides having other achievable oxidation states. These are cerium, praeseodymium, samarium, europium, terbium, and ytterbium. Of these, cerium, praeseodymium, and terbium have reasonably acheiveable +4 oxidation states. The loss of an electron from these elements results in the stability of an empty orbital, nearly empty orbital, and a half-filled orbital respectively. Samarium, europium, and ytterbium have reasonably achieveable +2 oxidation states. The gain of an electron in these elements results in a nearly half-filled orbital, half-filled orbtial, and full orbital respectively. Outside of these elements, other oxidation states were thought to be non-existant. However, with recent advances in lanthanide chemistry, the +2 oxidation state has been obtained with every lanthanide.24 This was obtained by the scheme observed in Figure 4.1. The earlier actinides were described previously as having numerous oxidation states. To focus on the later actinides, the chemistry was descibed as being lanthanide like. Although this is true to some extent, later actinides tend to also be susceptible to multiple oxidation states. Americium has been obtained in the +3 through +6 state. Berkelium has a readily accessible +3 and +4 oxidation states, while californium has been observed in the +2, +3, and +4 oxidation states. Interestingly, the +2 oxidation state becomes increasingly stable as the series continues after californium until nobelium’s dominant oxidation state is +2. In the lanthanides, gradual trends are often observed, and the lanthanide conctraction is a major one in particular. This contraction refers to the greater than expected decrease in ionic radii due to the poor nuclear shielding of the 4f orbitals, and this poor shielding makes it so the 6s electrons are attracted towards the nucleus. Additionally, this contraction is enhanced by aproximately 10% due to relitavistic effects. Another gradual feature examined in lanthanides is the dominant coordination number, which gradually decreases with ionic radius.25 Generally speaking, lanthanide chemistry can be explained as having linear trends.

While the lanthanides generally have linear trends when tranversing the series, actinide trends lend to be more stepwise. By examining the III→ II redox potentials of both f-series, another difference can be observed. Figure 5.1 shows the linear behavior of the increasing reduction potentials required to reach the +2 oxidation state in the lanthanides with the exception of samarium, europium, and ytterbium as explained earlier. In comparison, the actinides seem to jump in stability at plutonium and again at californium. This could lead to the identification of two transition points in the actinide series. The transition at plutonium can be explained by the diverse chemistry plutonium is capable of partacipating in. Plutonium’s chemistry has been better characterized than californium, in part due to the latter having substantially fewer studies. However, three recent studies have been done that obtained interesting phenomena in amercium and californium. In 2014, a californium borate structure was obtained that exhibited bizarre properties not normally observed in lanthanides.26 Such properties were the disparity of the observed magnetism from the theoretcial value and vibronic coupling. In californium, disparity of the magnetic properties has been observed in other compounds, however, not to the extent of this borate structure.27 The vibronic progression observed in this compound also portrayed the compound as unique because this property is an indicator of the prescence of stronger types of bonding, which is not present in lanthanides. Shortly after this, a californium pyridine structure was prepared that had a similar reduction in the magnetic moment to the californium borate structure.28 The properties observed in the californium compounds implied that there was a second transition in chemistry in the actinide series at californium. Since then, other studies on berkelium and americium have shown that the later actinides are not as identical as the lanthanides as previously thought. One example of this was displayed in berkelium(IV) iodate, where it was expected that this compound would be isostructural to cerium(IV) iodate. However, the berkelium compound is isostructural to the zirconium structure.29 Another EXAFS and XANES experiment with an americium complex determined that the bonding exhibited some covalent character that was not observed in the lanthanide analogs.30 These properties have allowed the actinides to further distinguish themselves from their 4f counterparts.

1.3.3 Dissertation Objectives This work focuses on gaining a deeper understanding of lanthanides and actinides in non- aqueous solutions and the solid-state, by examing the redox and coordination properties of select elements. In particular, studies utilize nitrogen and oxygen donor systems within terpyridines and crown ethers. The chemistry of the transuranics has been largely understudied. Therefore these studies focus on samarium and europium as these lanthanides are often considered analogs of the later actinides. This work consists of crystallography, electrochemisty, and computational studies, that aim to provide a better understanding of the redox chemistry in lanthanides and actinides as well as their coordination chemistry in non-aqueous media. A large portion of the work in transuranic chemistry has been done in aqueous conditions, and this could be due to the radiolytic effects transuranics could have on organic solvents. However, under the right conditions, this work shows that this chemistry is feasible. By examining this type of chemistry in the the f-elements, the extent of knowledge in these elements can be increased. All extensions of knowledge in the actinides are even more pronounced due to the scarcity of studies performed on these elements. Utilizing non-aqueous media hopes to provide new and critical information in two different aspects: determining how different lanthanides are from the transuranics, and uncovering the reasons for the second transition in the actinide series.

1.4 Figures

Figure 1.1. Schematic of the PUREX process used to extract uranium and plutonium from spent nuclear fuel.12

Figure 1.2. Breakdown of crystal structures in the Cambridge Structural Database by Metal Group. This demonstrates how little the f-elements research is performed in contrast to the transition metals

Figure 1.3. Radial extension of Sm3+ compared to Pu3+. The radial extension of plutonium extends beyond the core allowing the 5f orbitals to have more participation in bonding. Adapted from Clark.23

Figure 1.4. Reaction scheme utilized to obtain all the lanthanides in the divalent oxidation states.24

Figure 1.5. The III-II redox potentials of the lanthanides and actinides. Other than the noted lanthanides, Sm, Eu, and Yb, the +2-oxidation state is relatively difficult to obtain. In the actinides, the +2-oxidation state becomes increasingly stable.

CHAPTER 2

SYNTHESIS AND CHARACTERIZATION METHODS

2.1 Past and Current Methods for f-element Crystal Growth Crystal growth has been a large part of understanding f-elements chemistry. This is because crystal structures give insight to structural traits such as bond lengths, radii, and crystal packing that contribute to certain properties of compounds. Methods that are often employed are hydrothermal or solvothermal syntheses to produce crystallization of rather insoluble compounds. These types of syntheses have been around since the mid 1800’s.31 Hydrothermal or solvothermal synthesis typically involve heating a reaction in a sealed teflon liner surrounded by a steel jacket at high temperatures and pressures, where the formation of crystals is usually from the least soluble product. This type of reaction was often utillized in the crystallization of many transuranic compounds.26,28,29,32–34 However, if rapid crystallization can be accomplished, evaporation and solution crystallization can be utilized before radiolytic damage begins. There is less concern involved when using this type of technique with lanthanides as radiolysis is often not a problem. Using this type of method is advantageous over hydrothermal syntheses for a few reasons. First, the variety of solvents is greater as many organic solvents can decompose quickly at higher temperatures and pressures to form peroxide or CO-systems. Furthermore, the use of more bulky, reactive ligands can be utilized as these complexes can also break down under hydrothermal syntheses conditions. Lastly, non-aqueous media can be used, and this is important when considering redox reactions that cannot employ solvents such as water and alcohols. These are often used for hydrothermal syntheses, but they can oxidize metals in a reduced state. Additionally, these redox reactions often use reductants such as potassium and sodium that react violently with water and alcohols.

2.2 Characterization Techniques 2.2.1 Single Crystal X-ray Diffraction Single crystal X-ray diffraction is one of the most powerful tools in characterizing materials in the solid state. As previously mentioned, crystal structures give structural information leading to particular properties. Crystals that are not air or water-sensitive were

placed in immersion oil and adhered to a MiTeGen mount on a magnetic base. Crystals that were air or water-senstive were mounted in the same fashion, but they were either prepared in a glovebox or in Parabar 10312 oil. The crystals were then aligned on a Bruker D8 Quest diffractometer equipped with a digital camera. The intensity measurements were performed with a ΙμS X-ray source with a 30 watt microfused sealed tube (Mo K Kα, λ = 0.71073 Å) and multilayer focusing optics. Unit cell determinations utlized APEX III, and collections were measured by either a standard hemisphere or calculated strategy to give the highest possible resolution. SAINT software was used for data integration. XPREP was used for space group determination and correct symmetry. OLEX2 equipped with the SHELXT suite was used for crystal structure determination.35,36 Crystal structure images were prepared in OLEX2 or Mercury. 2.2.2 Spectroscopic Measurements Another excellent method for characterizing f-elements is by spectroscopy. This is especially true for lanthanides, which exhibit very specific forbidden f-f transitions. These characteristic traits have been well studied.37 Absorbance measurments were performed on a CRAIC Technologies microspectraphotometer. These measurements were accomplished by placing single crystals on a glass slide in either Immersion or Parabar 10312 oil which are then aligned using a built-in microscope. In rare cases, this instrument was used for photoluminescence measurements with a 365 nm and 420 nm excitation source. The majority of photoluminescence measurments displayed in this work was performed on an Edinburgh FLS980 fluorescence spectrometer equipped with a nitrogen cryostat cold finger. A housed 450 W Xe lamp/single grating (1800 λ/mm, 250 nm blaze). To excite samples, a Czernty-Turner monochromator was utilized. The data was then passed through a Czerny- Turner monochromator (once passed through the necessary wavelength filter) and collected on a Peltier-cooled Hamamatsu R928 photomultiplier tube. Raman measurements were performed on a Horiba JY LabRam HR800 Raman Spectrograph utilizing a TUI Optics DL 100 grating-stabilized diode laser with 80 mW emitting power at 785 nm. The laser source was then focused on the sample with an Olympus BX30 microscope with a 5x objective that simultaneously collected on the background scatter radiation. The scatter was then filtered through a Semrock RazorEdge Long Wave Pass edge filter before being dispersed onto a 76 mm square 600 line/nm grating onto a 1024 x 256 element

open electrode. This was then collected by a CCD detector with a 26 μm square pixel cooled at - 70 ℃. 2.2.3 Computational Details The arrangement of this work is based on a compilation of independent studies, and therefore, many computational details are specific to certain molecules. For this reason, this section is devoted to describing the computational details that transverse the entirety of the work studied. Other computational details that are specific to a chapter are discussed in further detail. The software used to run wavefunction-based calculations were the ORCA4.02 package developed by Neese and coworkers and MOLCAS 8.2.38,39 The CASSCF approach was used to construct all possible configurations within a chosen active orbital space because it provides an accurate multiconfigurational character present in open-shell actinide systems. All these calculations account for static and dynamical correlation, which is the most common problem associated with single-configuration calculations such as Hartree-Fock (HF) and Density Functional Theory (DFT) along with the spin-orbit coupling (SOC) correction to the states. All CASSCF wave functions were obtained using DFT/HF orbitals. For DFT calculations, the PBE040 hybrid functional was used due to its proven effectiveness in predicting geometrical parameters for heavy elements.41 Finally, the need to understand bonding differences between some lanthanide/actinide systems made it necessary to analyze bonding parameters, which were calculated using the MultiWFN program relying on the Bader’s theory of Atoms in Molecules.42 2.2.4 Electrochemical Studies Electrochemical experiments were performed to assess the stability of f-element cations in non-aqueous conditions. Electrochemical experiments utilized a CH Instruments electrochemical analyzer CHI620E. The electrodes were passed into an argon-filled glovebox with a traditional three electrode configuration. This consisted of a platinum working electrode, platinum wire auxiliary electrode, and a Ag/Ag+ reference electrode for many of the non-aqueous solutions. In the case of DMF, a double junction Ag+/Ag(2.2.2-cryptate)+ electrode was used. The use of this double junction was necessary to protect the silver ions from DMF.43 For the

electrolyte, tetrapropylammonium tetrakis[3,5-bis(trifluoromethyl)phenyl]borate (NPr4BArF) was utilized.

2.2.5 Syntheses of Reactions The reactions presented in this work were performed by crystallizations in non-aqueous solvents. Redox reactions often required reduction potentials greater than -1.5 V. Therefore, strong reductants were often employed, such as sodium and potassium metal or potassium graphite. These reductants react violently with water, and caution should be taken when working with these materials. All redox reactions were performed under Schlenk like or glovebox techniques with rigorous exclusion of air and water. Handling of the transuranic materials presented were carefully handled in a facility equipped with HEPA filters and continuously monitored alpha, beta, and gamma detectors dedicated to this type of work. 2.2.6 Recycling of 243Am and 249Cf Transuranic metals are highly radioactive, scarce, and costly, especially in terms of academic research. Therefore, recycling of these metals is vital as it keeps radioactive waste from accumulating as rapidly. There have been procedures for reprocessing 243Am and 244,248Cm after chemical reactions.44 This process involves digesting the unknown residue in acid and running the solution through a column. However, in the case a californium, this process had not been tested. Additionally, extra steps were necessary due to the contamination of transition metals that can form anionic complexes in the mixture. Figure 2.1 shows the basic scheme of this

process. To obtain a pure CfCl3∙xH2O solution the following procedure was performed. An 249 unknown residue containing Cf was carefully taken up in a mixture of 8M HNO3 and 6M HCl and carefully transferred to a 23 mL Teflon liner. The mixture was sealed and heated at 200 ℃ for 24 hours. The solution was evaporated to a soft residue and transferred to a platinum crucible with DI water. The solution was heated to 900 ℃ over the course of 24 hours to ash any organics left in the mixture. After burning, the solution was dissolved in concentrated HCl and dried to a soft residue. This dissolution in HCl followed by subsequent evaporation was repeated two additional times. After the last evaporation, the solid residue was taken up in 5 mL of 6M HCl, transferred to a 15 ml Falcon tube, and diluted with 5 mL of DI water. Then 3 mL of concentrated HF was added to the solution to precipitate CfF3 as a bright greenish yellow solid. This solution was centrifuged, and the pellet washed with DI water. The pellet was then dissolved in a 50/50 mixture of concentrated HCl and saturated boric solution with gentle heat

(~50-70 ℃). Upon dissolution of the CfF3, the solution was dried to a residue.

The solid matrix was then dissolved in 1M HCl and loaded onto a 5 mL cation column (AG-50 X8 resin) in a 20 mL BioRad column. The cation column was prepped by the method proposed for americium. The column was washed with 1M HCl (100 mL), then the californium was eluted from the column with concentrated HCl (~40 mL). To remove transition metal ionic species, an anion column was run immediately after the cation column. The 10 mL anion column (AG-MP1 resin) was conditioned by running 30 mL of ultrapure water and 30 mL of concentrated HCl through the column. The californium solution was then run through the column and collected in the smallest possible volume. Elution was stopped once the activity of the column was less than 10% of the original activity of the solution. This solution was then carefully dried to a residue and taken up in 5 mL of 1M HCl. The concentration of the solution was calculated by UV-Vis.

2.3 Figures

249 Figure 2.1. Purification scheme of Cf. The precipitation of CfF3 allows for a greater confidence of purification as many other metal fluorides are soluble.

CHAPTER 3 EXAMINATION OF STRUCTURE AND BONDING IN 10-COORDINATE EUROPIUM AND AMERICIUM TERPYRIDYL COMPLEXES

Adapted with permission from 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, and Thomas Albrecht-Schmitt. Submitted to Inorganic Chemistry.

3.1 Introduction Bonding differences between 4f and 5f elements have been of significant interest from both fundamental and applied perspectives for more than six decades.45,46 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.47 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.14,46 Numerous neutron-capture 241 products are generated during nuclear energy production. Among these, Am (t½ = 433 y) is produced in relatively high abundance with approximately 1.3 kg being synthesized per ton of fuel under standard burn up conditions.48 When scaled to repository levels, the large quantities of this moderately long-lived isotope creates a heat burden that can be mitigated by extraction prior to entombment and subsequent fissioning in fast neutron reactors.49 Thus, in a closed nuclear fuel cycle of this kind a 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 an Ln:Am ratio of ca. 15:1, complicating the separations even further. 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 15

ionic radius ~0.03 Å smaller than AmIII,52 these ions are still ostensibly isoelectronic.27 Accordingly, when differences in the properties of isomorphous EuIII and AmIII complexes are observed, it is 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. 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. Whereas 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 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 3- 51,53 compounds such as the sesquioxides and hexachlorides, Ln2O3 and [LnCl6] , and utilization 54 of all possible frontier orbitals in compounds like Cf[B6O8(OH)5]. 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.55–57 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.58,59 One of the recent focuses in f-block 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,60 and therefore show a preference for binding actinides versus lanthanides61 However, this softness effect is amplified by 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.62 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) that should be considered as an archetypal ligand for forming stable 16

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.63–65 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.44,58 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.

3.2 Results and Discussion 3.2.1 Structural Characterization

M(TpyNO2)(NO3)3(H2O)·THF (An1/Ln1) (M = La, Nd, Sm, Eu, Tb, Am) for the sake of brevity only Eu1, Nd1, and Am1 are discussed and details on the rest can be found in the ESI.

All compounds are isomorphous and crystallize in the monoclinic space group P21/c. The 3+ structure features M cations bound by one, tridentate TpyNO2 ligand, three nitrate anions, and one water molecule creating a ten-coordinate environment around the metal centers as depicted in Figure 1. The co-crystallized THF solvent molecule does not 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 ten-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, but on the most fundamental level the belief that lanthanide and actinide contractions are easily demonstrated periodic trends is usually false. This is attributable to two features of bond distances measured from X-ray diffraction experiments. First, 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.66,67 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 17

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+, but this subsumes that said compounds can be prepared. In fact, there are often thermodynamic, synthetic, or practical barriers that render such studies impractical or even implausible in a number of chemical systems. 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 are also 0.048(2) Å with the Am‒

OH2 bond distances being longer again. Again, the ionic radius of EuIII is ca. 0.03 Å smaller than AmIII with the latter’s ionic radius comparing well with NdIII.68 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 re-emerge 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),44,51 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

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 152,154Eu 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. The ionic radius of 10-coordinate 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 methods 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.68 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 B.1). Recalculating the average ionic radii for the nitrate oxygen atoms, terpy nitrogen atoms, and water oxygen atom based on the lanthanum complex gives the expected ionic radius of 1.35 Å for the ionically bound nitrate oxygens. Conversely, the terpyridine nitrogens give an ionic radius of 1.39 Å. While the water oxygen atom yields 1.25 Å, reflecting the degree of covalency observed. 3.2.2 Ionic Radii Calculations The nitrate bond-based calculations are the most accurate because the nitrate anion interaction with the cationic centers 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 the 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 atoms bonds of H/ρ provided in Table B.2. 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.51 Therefore, their calculated ionic radii are expected to be similar (Table B.3). 19

Unfortunately, due to the solvent molecule in the structure, the volume to ionic radius 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.69 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 the oxygen instead of 1.35 Å. Using 1.35 Å for the nitrate oxygen atoms and 1.36 Å for the crown-ether oxygens gives an Eu3+ ionic radius of 1.176(4) Å. This number is within certainty of the calculated ionic radius for the 10- coordinate europium complexes presented in this work. One example of a previously reported lanthanum complex containing a 10-coordinate III 70 La center is La(terpy)(NO3)2(acac). 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 acac oxygens similar to the nitrate oxygens), 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 1. 3.2.3 Spectroscopic Measurements 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 intra-ligand transitions of the TpyNO2 ligand. When the complex is excited at 365 nm, sharp photoluminescence bands are observed as shown in Figure 3 that are assigned 7 7 III to the F4 – F0 transitions of Eu . 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 nm 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 non-radiative 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 non-radiative.71 20

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 7 5 nm corresponding to the F0→ L6 transition as shown in Figure 4.

3.2.4 Computational Analysis A useful theoretical tool for addressing bonding is the Quantum Theory of Atoms in Molecules (QTAIM), and its parameters derived from the Bader’s theory.72 These parameters have been used to understand the nature of interactions between metal ions and ligands in the f- block. Table B.2 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 III III observed experimentally in the Ln and Am 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 bond displaying some degree of covalency in only one of the ten 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.

3.3 Conclusion In summary, a series of lanthanide and americium terpyridine coordination complexes have been prepared and characterized. These compounds reveal ten-coordinate metal sites, and

the first example of this coordination number with AmIII. The ionic radii have been determined for ten-coordinate EuIII and AmIII to be 1.17 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 X-ray diffraction datasets 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.

3.4 Methods and Characterizations 3.4.1 Syntheses

243 M(TpyNO2)(NO3)3(H2O)∙THF (M = La, Nd, Sm, Eu, Tb). Caution! Am (t1/2 = 7.38 x 103 years) has potential health risks due to its α and γ emission, along with the emission of its 239 239 daughter Np (t1/2= 2.35 days). Np 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.

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. 73,74 4’-nitrophenyl terpyridyl (TpyNO2) was prepared by literature methods. 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 containing 5 mL of acetone. The acetone was vapor-diffused into the solution overnight producing Nd1 and Eu1 as amber colored crystals, respectively, suitable for single crystal X-ray diffraction. Isostructural crystals of La, Sm, and Tb were prepared in a similar manner.

243 Am(TpyNO2)(NO3)3(H2O)∙THF. Am(NO3)3∙xH2O was prepared in the following 3+ manner: 595 µL (containing 3 mg of Am ) 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. 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.

3.4.2 Single Crystal X-ray Diffraction Single crystals of each compound were placed on a Mitogen mounts using Immersion oil. The crystals were aligned with a Bruker D8 Quest X-ray diffractometer with an IμS X-ray source (MoKα, λ = 0.71073 Å) paired with a digital camera. Low temperature (28 K) data collection 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.35,36 3.4.3 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 wavefunction was used as a trial wavefunction.39 The DFT functional used was the hybrid PBE0 with the SARC- TZVP basis set for the metal centers while the ligands were modelled 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 wavefunction 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 was used to take into account the spin-orbit coupling (SOC), The 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 wavefunction previously described.75 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) 2 indexes. However, ρ(r) has been∇ proven to be positive for f-element compounds where purely 2 covalent bonds are ∇characterized by negative ρ(r) values. Instead, energy density parameters give more useful information regarding the nature∇ of the interaction where if potential energy

23 density, V(r), predominates over the Lagrangian kinetic energy, G(r), the total energy density will be negative and a certain degree of covalency can be attributed to that bond.

3.5 Figures

Figure 3.1. Crystal structure of [Am(TpyNO2)(NO3)3(H2O)]·THF (Am1) shown with the 50% ellipsoid probability. Hydrogen atoms and the co-crystallized THF molecule have been omitted for clarity.

Figure 3.2. Excitation and photoluminescence spectra of [Eu(TpyNO2)(NO3)3(H2O)]·THF (Eu1). The sharp f-f transitions of Eu3+ are displayed.

Figure 3.3. Absorption spectra of [Eu(TpyNO2)(NO3)3(H2O)]·THF (Eu1) and [Am(TpyNO2)(NO3)3(H2O)]·THF (Am1).

CHAPTER 4

STRUCTURAL, PHOTOPHYSICAL, AND ELECTROCHEMICAL INVESTIGATIONS OF F-ELEMENTS IN NON-AQUEOUS CONDITIONS

4.1 Introduction Electrochemistry has been an essential part of fundamental chemistry since its recognition in the mid 1800’s.76 In particular, this research has been fundamental to f-elements as many of the oxidation states observed today were not previously recognized.24 Based on the calculated Ln3+→Ln2+ redox potentials of some lanthanides, achieving the +2 oxidation state was not considered possible even with the strongest reducing agents. Considering gadolinium(III) has a very stable f7 configuration and calculated redox potential of -3.9 V, it has been obtained in the divalent oxidation state.77 This was obtained using potassium as the reductant which has a reduction potential of potassium (-2.9 V). Because of this, there is two ways of thinking that this can be possible. The first is that the calculated potential of many lanthanides are incorrect, and that the actual reductional potential must be lower than that of potassium’s. The second is that the conditions in which the reaction was performed lowered the Ln3+→Ln2+ redox potential to a potential lower than -2.9 V. Although lanthanides quite often have periodic trends, there is one aspect of these elements that do not portray a trend: the filling of the f or d-orbital in the +2-

oxidation state. This seems to be sporadic. The magnetic moments of [K(2.2.2-cryptand)][Cp’3 Ln] were measured and it was found that Sm, Eu, and Tm take on 4fn+1 configuration while Y, La, Gd. Tb, Dy, Ho, and Er have 4f5d1 configuration.78 Additionally, it was found that for some lanthanides, the electron could be tuned to go into either the f or d-orbital.79–82 Therefore, one can assume that the ligand plays a role in adjusting the redox potential of lanthanides. This type of redox chemistry is very scarce in the actinides. However, a few examples of divalent 5f compounds have been structurally characterized. These metals are such that the ground state configuration is 5f6d1.83–85 Recently, this formal +2-oxidation state has been achieved where the addition of an election begins to fill the essentially non-bonding 5f orbitals.22,23 Along with the scarcity and radioactivity of the later actinides (Am, Cm, Bk), these elements also begin to exhibit high III→II redox potentials (-2.3, -2.8, -2.5 respectively). As previously mentioned, californium seems to be a break in this trend. The estimated and closest measurement of californium’s reduction potential by is -1.6 V.86 This is very similar to

samarium III→II redox potential of -1.55 V. Therefore, by examining the lanthanides redox and electrochemistry, a model can be constructed for the later actinides that could be suitable. The divalent oxidation state is becoming more intriguing to examine because it gives a different look at the chemistry in which f-elements can participate. Obtaining these oxidation states were made attainable by changing the environment that f-metal was in. This was often achieved using a crown ether of some sort. Therefore, this work investigates structural, spectroscopic, and electrochemical properties of lanthanides and actinides with various crown ethers.

4.2 Results and Discussion 4.2.1 Structural Characterization The structure of each crown ether compound was able to be determined by single crystal

X-ray crystallography. [Sm(DBC)(NO3)2(H2O)]2[Sm(NO3)5] (DBC = dibenzo-30-crown-10) (SmDBC) is structurally similar to the heavier lanthanide compounds synthesized by Lu et al with the exception of the solvent molecule.87 The structure has three independent Sm(III) sites. Two of these sites are 9-cooridinate with a DBC ligand, two nitrate anions, and two water molecules coordinating. The other site is a 10-cooridinate site with the Sm(III) center being bound by five nitrate anions as illustrated in Figure 4.1. The DBC ligand is tridentate with average Sm—O distances of 2.500(3) Å. The nitrates bind at bond distances of 2.554(4) Å, and the water bond distances average 2.490(3) Å. The geometry of these two metal sites can be best described as tricapped trigonal prism. The penta-nitrate Sm(III) site has 10 oxygens bound by an average bond distance of 2.487(3) Å forming a bicapped square antiprism. The two ligand sites are also stabilized by intramolecular hydrogen bonding between the coordinating waters and crown ether.

The structures of the [Ln(2.2.2-cryptand)(THF)][BPh4]2 (Ln = Sm, Eu) show that the lanthanide sites are 9-cooridinate. The crown ether encapsulates the metal site and it capped by a THF solvate to give a mono-capped square antiprism geometry of the Ln2+ center (Figure 4.2). The crystallographic bond distances for these compounds were found to be very close between the two compounds. These bond distances are similar to other in-crypt lanthanide compounds.88 However, these compounds presented a challenging, interesting crystallographic problem. Although the crystal symmetry is triclinic, the compounds cannot be readily placed into one of

the two triclinic space groups (P1 and P-1) over the other. The space group P1 appears to be the more correct space group. In P-1, there is a fully occupied cryptand site along with a half occupied cryptand site on the inversion center. In P1, there are three cryptand sites, with one site displaying whole molecule disorder. In this model, the site is not half occupied but approximately 70/30 for each orientation. For the SmCrypt structure, when cooled to 100 K, the structure undergoes additional ordering that gives a slight expansion in the unit cell.

Ytterbium crystallized without the 2.2.2-cryptand in the structure as Yb(Ph4)2. This is perhaps due to the size difference of ytterbium from the larger samarium and europium lanthanides. Ytterbium might not be bound as tightly to 2.2.2-cryptand as the crown ether is often used for the chelation of large cations such as alkali metals. 4.2.2 Spectroscopic Studies The excitation and emission of SmDBC was obtained (Figure 4.3). The excitation spectrum shows a broad absorbance at around 300 nm followed sharp peaks related to f-f transitions that are forbidden. When the sample was excited at 320 nm, or the DBC ligand, only a broad band in observed in the emission. If excited at 403 nm, the feature of a broad band along 3+ 4 6 4 6 with the Sm ion f-f transition peaks are observed to correlate to the G5/2- H5/2, G5/2- H7/2, and 4 6 G5/2- H9/2 transitions. It is interesting to note that DBC phosphoresces on its own but does not work well as an antenna for intramolecular charge transfer to Sm3+.

When a solution of Sm(trif)3 was reduced over potassium in DMF, the spectroscopy was measured before and afterwards (Figure 4.4). This showed the transition from sharp f-f peaks to broad features. This indicates the reduction to the divalent oxidation state, which is explained in further detail below.

The divalent SmCrypt and Yb(BPh4)2 compounds do not exhibit luminescence properties like SmDBC does. Therefore, another such method of verifying these compounds exist in the divalent oxidation state is by measuring the absorbance properties of the compounds. When in the trivalent oxidation state, lanthanides exhibit f-f transitions although in some divalent oxidation states exhibit this property as well.89 Excluding EuCrypt, these compounds were reactive with oxygen and water. Even when coated in Parabar 10312 oil, these compounds begin to degrade after approximately 30 minutes. The samarium complex was most reactive as it has the highest III→II redox potential of the three compounds. As observed in Figure 4.5, the compound was completely degraded after one hour. This is either due to oxidation or solvation

of the THF from the structure. The structure of SmCrypt was attempted after oxidation, but no diffraction peaks were observed indicating an amorphous product. Divalent lanthanides often display broad absorption spectra which are attributed to the Laporte-allowed 4f-5d transitions.90 This transits to the deep color often observed in divalent lanthanides. The atomic spectra of these compounds show transitions 20,000 – 30,000 cm-1 91,92 above the 4f levels. Figure 4.6 shows the absorbance spectra of the Yb(BPh4)2 compound. The compound displays broad band features often observed in the divalent lanthanides confirming the oxidation state of these compound to be +2. Interestingly, of the three compounds, EuCrypt was the only one to have fluorescence which is shown in Figure 4.7. However, Eu(III) and Eu(II) have been shown to be capable of being white light emitters.93,94 4.2.3 Electrochemical Studies Electrochemical studies on these elements revealed interesting characteristics of f- elements in non-aqueous conditions. The first experiment observed the effects of the solvent on

samarium. Figure 4.8 shows the cyclic voltammetry of Sm(trif)2 in dimethylformamide (DMF) and THF. The potential becomes much more favorable in THF. Both spectra show samarium acting as quasi-reversible, but the THF spectrum shows more reversibility. This is likely due to THF not being as strong of a coordinating solvent which can withdraw electron density from the metal center making the it more positively charged. The effect of the ligand field on samarium’s redox potential was then considered. The ligands used were 18-crown-6, 2.2.2-cryptand, and DBC. These ligands provide different coordination spheres around the metal center by changing the coordination number, geometry, and N- or O-donors. Each one of the ligands produced an almost indistinguishable change in potential from each other. However, as observed in Figure 4.9, the 2.2.2-cryptand system gave the best response. For this reason, the electrochemistry of the cryptand system was studied in more depth with europium, ytterbium, and californium. As mentioned previously, these elements have reasonably high reduction potentials. Therefore, care was taken to handle these experiments in a glovebox to exclude air and water. The metals all underwent relatively large shifts in reduction potential with each exceeding a shift in potential of at least 0.5 V (Table C.1). In Figure 4.10, europium shows the most favorable

reduction potential, even though ytterbium has the most reversibility as indicated by Ipc/Ipa in

Table C.1. The greater this value, the less reversible the analyte of examination becomes. A value of 1 generally indicates reversibility, while values lower than 1 show greater reversibility. Interestingly, the redox potentials of samarium and californium are similar but give very different shifts in potential. Samarium has a shift in potential of +520 mV while californium displays a shift in potential of +715 mV. This shift in californium is similar to the shift observed in ytterbium. However, this is believed to be for two different reasons. By examining the crystal structure of the three lanthanides, ytterbium forms a different complex in the same conditions. This means the coordination chemistry of ytterbium is giving rise to different environments in solution that could produce a different electrochemical behavior. In californium, we believe this larger shift in redox potential is attributed to the difference in 5f and 6d orbitals. In lanthanides, certain configurations can lead to the expression of d-orbital involvement the divalent oxidation state. This is further examined in Chapter 5. In californium, due to the radial extension of the 5f orbital versus the 4f orbitals, it is considered that the divalent oxidation state involves the addition of the electron to the f-orbital to give a 5f10 configuration. In this case, obtaining the divalent oxidation state of californium should be easier to obtain than samarium. This belief is further explained by the slightly lower reduction potential measured than what previously considered for californium of -1.6 V. Samarium’s potential of -1.55 V was slightly lower than californium’s, but this experiment shows the opposite. Against the ferrocene couple, samarium’s redox potential was measured to be -1.66 V and californium’s to be -1.53 V. Although this is only a difference of 13 mV, is adds to the notion of the divalent oxidation state becoming easier to obtain as the 5f series transverses to the later actinides.

4.3 Conclusion The structure, spectroscopic, and electrochemical investigations of divalent lanthanides and actinides were performed to gain a better understanding of the non-aqueous redox chemistry of these elements. Different crown ethers were explored for synthetic and electrochemical purposes. Although DBC proved to coordinate to lanthanides, 2.2.2-cryptand gave structural insight to divalent lanthanides as well as providing the best response to electrochemistry experiments. After performing electrochemistry experiments on samarium, europium, ytterbium, and californium, a better understanding of the redox chemistry was gained. Ytterbium, which did

not give isostructural information like samarium and europium, showed to have the greatest reversibility of all the elements examined due to the solution confirmation. Californium showed to have a similar behavior to ytterbium in its shift of redox potential. Albeit similar, these reasons why were very different as we believe californium’s shift in potential is due to the stabilization of an electron in the f-orbital to give a 5f10 configuration. This electrochemical experiment has now provided physical evidence for the stabilization of the divalent oxidation state in the latter actinides.

4.4 Methods and Characterization 4.4.1 Syntheses

[Sm(DBC)(NO3)2(H2O)]2[Sm(NO3)5]. SmDBC. In a 6 mL glass vial Sm(NO3)3∙6H2O (23.1 mg, .052 mmol) was dissolved in 1 mL of acetonitrile. In a separate glass vial, 26.8 mg of DBC (0.050 mmol) was dissolved in 1:1 mixture of acetonitrile/chloroform. The two solutions were mixed and layered with 1 mL of a 1:1 mixture of toluene/hexane. The reaction was capped and placed in a desiccator for two weeks in which large colorless, plate-like crystals formed.

[Ln(2.2.2-cryptand)(THF)][BPh4]. SmCrypt. In an argon filled glovebox, 10.3 mg of

SmI2 (0.0255 mmol) added to a 20 mL glass vial in 3 mL of acetonitrile dried over molecular

sieves producing a dark green solution. Then 2 mL of THF containing 23.1 mg of Bu4NBPh4 (0.0411 mmol) was added. The solution was stirred. Then a 1 mL THF solution of 2.2.2-cryptand (10.3 mg, 0.0274 mmol) was added dropwise. The solution turned a lighter green. The solution was placed in a – 34 ℃ freezer overnight producing SmCrypt as green, irregular crystals.

EuCrypt. In an argon-filled glovebox was added 10.6 mg of EuI2 (0.0261 mmol) to a 20 mL glass vial in 3 mL of acetonitrile. The solution turned light yellow. Then 2 mL of THF

containing 23.1 mg of Bu4NBPh4 was added to the reaction in which the solution turned faint yellow. Upon addition of 1 mL of THF containing 10.3 mg of 2.2.2-cryptand (0.0274 mmol) the solution turned colorless. The solution was placed in a freezer at -34 ℃ overnight producing EuCrypt as irregular, colorless crystals.

Yb(BPh4)2. In an argon-filled glovebox was added 9.9 mg of YbI2 (0.0232 mmol) to a 20 mL glass vial in 3 mL of acetonitrile. The solution turned dark yellow. Then 2 mL of THF

containing 23.1 mg of Bu4NBPh4 was added to the reaction in which the solution turned faint yellow. Upon addition of 1 mL of THF containing 10.3 mg of 2.2.2-cryptand (0.0274 mmol) the

solution remained pale yellow. The solution was placed in a freezer at -34 ℃ overnight producing YbCrypt as irregular, pale-yellow crystals. The color of these solutions before and after addition of THF are displayed in Figure 4.11. 249 Electrochemistry Solutions. Caution! Cf is a highly radioactive material (t(1/2) = 351 years) with a specific activity of 4.1 Ci g-. It has a very serious health risk potential (γ emission = 249 245 0.388 MeV). Cf decays to Cm (t(1/2) = 8,500 years) from is 6.194 MeV alpha decay energy. These studies were performed in an argon glovebox, fitted to a HEPA filter system, in a lab that is continuously monitored with β and γ monitors. Lanthanide triflates were prepared dissolving the anhydrous lanthanide triflate in THF dried over sodium metal in an argon filled glovebox. An

equimolar amount of 2.2.2-cryptand was added to the reaction vial. Cf(trif)3 was prepared in the 249 following manner. 1.0 mg of CfCl3∙xH2O from Oak Ridge National Lab was converted to

CfI3∙xH2O. This was done by dissolving the chloride salt in concentrated hydroiodic acid with gentle heat. The solution was then evaporated to dryness and the process repeated two additional times. The sample was then pumped into a glovebox overnight. Once in the glovebox, the iodide salt was dissolved in 0.5 mL of THF to produce an amber colored solution. Then 3.2 mg of silver triflate in 0.6 mL THF was added in which a white precipitate formed (AgI) (Figure 4.12). The AgI was allowed to settle and the amber colored solution was filtered for use in electrochemistry experiments. 4.4.2 Single Crystal Crystallography Crystals of SmDBC were mounted in immersion oil on a MiTeGen mount. Single crystals of LnCrypt were mounted in Parabar 10312 oil. Further examination was proceeded as mentioned in Chapter 2. 4.4.3 Spectroscopic Studies The UV-Vis of LnCrypt was done on a Craic Technologies Microspectrophotometer by placing single crystals on a glass slide in Parabar 10312 oil to minimize exposure to air and water. The luminescence of SmDBC was measured on the setup explained in Chapter 2. The sample was prepared by placing single crystals of SmDBC on a glass slide in immersion oil. A secondary glass slide was placed on top, and the sample was sealed with epoxy. 4.4.4 Electrochemical Studies The electrochemical investigations were prepared in the ways discussed in Chapter 2. However, in the case of LnCrypt, the solid samples had to dissolved in acetonitrile as they were

32 insoluble in many of glovebox solvents capable of the solvent range required for these electrochemistry experiments. To get electrochemical studies, BuN4BPh4 was not added to avoid precipitation. The conditions in which the experiment was performed were as follows: THF as f solvent, 0.1 M NPr4Bar 4 as electrolyte, 25 ℃, with a 50 mV/sec scan rate.

4.5 Figures

Figure 4.1. The crystal structure of SmDBC shown with the 50% thermal ellipsoids. The structure contains three samarium sites. Two sites show the ligand around the metal, while the third samarium site is surrounded by five nitrates in order to charge balance the structure. Hydrogen bonding is displayed to show how further stabilization is achieved through intramolecular interactions.

Figure 4.2. The encapsulation of Sm with 2.2.2-cryptand is presented. The metal center in the LnCrypt structures have the same geometry resulting in a 9-coordinate metal site. The structure is shown with the 50% thermal ellipsoid probability (hydrogens omitted). The metal is shown in green, oxygen in red, nitrogen in blue, and carbon in grey.

Figure 4.3. The luminescence of SmDBC shows the f-f transitions commonly associated with trivalent lanthanides. The molecule does not have efficient energy transfer. When excited at 320 nm, the Sm3+ peaks are not observed as in compounds that exhibit antenna effects.

Figure 4.4. Emission spectra of samarium before and after reduction. Before reducing to Sm2+ , the f-f transitions associated with Sm3+ are visible. After reduction, the broad band emission is characteristic of f-d transitions which are allowed.

Figure 4.5. The absorbance of [Sm(2.2.2-crytpand)(THF)][BPh4]2 over the course of one hour. As the compound begins to decompose, it becomes amorphous and loses its properties.

Figure 4.6. Absorbance spectrum of Yb(BPh4)2 as a solid crystalline sample. Like the LnCrypt structures, the compound possesses the divalent broad spectroscopic properties.

Figure 4.7. The absorbance and emission properties of [Eu(2.2.2-cryptand)(THF)][BPh4]2. Unlike the other divalent lanthanides studied in this work, EuCrypt has fluorescent properties in the blue region.

Figure 4.8. The cyclic voltammagram of Sm2+ in DMF with and without 2.2.2-cryptand. This demonstrates the effects ligands can play on the redox potential of f-elements. When 2.2.2- cryptand is present in solution of divalent lanthanides, the redox potential can shift drastically to more favorable conditions.

Figure 4.9. CV of various crown ethers with Sm3+ in acetonitrile. The CV’s show that the ligand has a small effect on samarium. However, it was observed that 2.2.2-cryptand gave the best response for studying the lanthanides and actinides.

Figure 4.10. The CV's of Sm, Eu, Yb, and Cf in the prescence of 2.2.2-cryptand. The spectra show that europium has the most favorable redox potential due to the stabilization of the half- filled f-orbital. Californium, however, was shown to have an easier path to stabilizing the divalent oxidation state than samarium.

Figure 4.11. The color change exhibited by divalent Sm, Eu, and Yb (left to right) when complexed by 2.2.2-cryptand in solution.

Figure 4.12. Precipitation of AgI from the reaction of CfI3 with silver triflate. The mother liquor was removed from the solid and was used to obtain the first cyclic voltammogram of californium in non-aqueous solutions.

CHAPTER 5 TUNING THE ENERGETICS OF SAMARIUM(II): MOLECULAR AND ELECTRONIC STRUCTURE, AND HYDROLYTIC REACTIVITY OF A SAMARIUM(II) CROWN ETHER COMPLEX

Adapted with permission from Frankie D. White, Cristian Celis-Barros, Jillian Rankin, Eduardo Solís-Céspedes, David Dan, Alyssa N. Gaiser, Yan Zhou, Jasmine N. Colangelo, Dayán Páez- Hernández, Ramiro Arratia-Pérez, and Thomas E. Albrecht-Schmitt. Submitted to Chemical Science.

5.1 Introduction Cryptands and crown ethers possess the ability to strongly bind lanthanide cations concomitant with substantial changes in the relative stabilities of the III+ versus II+ oxidation states.24,88,95 The resultant complexes can be used as sensors through selective and reversible quenching of the photoluminescence from several lanthanides, most notably EuIII and TbIII.96–98 Moreover, these shifts in the reduction potentials can be utilized to stabilize and perhaps isolate rare examples of divalent lanthanide complexes that can be used to understand coordination chemistry,99,100 electronic structure,24 and reactivity.88 The electronic structure of classical LnII compounds that are typically thought of as adopting 4f n+1 configurations is readily explained through the stabilization imparted by self-exchange that maximizes in half- and fully-filled orbitals as exemplified by EuII (4f 7) and YbII (4f 14). In contrast, other divalent lanthanides such as NdII and DyII have been observed in both 4f n+1 and 4f n5d1 configurations where calculated differences in energies between the configurations are small.101 Thus, the particular configuration adopted is currently challenging to predict either experimentally or theoretically. The consequences of the placement of the additional electron profoundly affects the resultant physical and chemical properties of the complexes. Hence, an expansion of this chemistry beyond the rare examples that are currently known may provide the key to unlocking our understanding of what dictates configurational preferences in the f-block. Despite the unusual and potentially useful properties of encapsulated divalent lanthanide complexes, structurally characterized examples of cryptand and crown-ether complexes other than those containing EuII remain scarce owing to their substantially increased reactivity with water and oxygen.102 Oddly, SmII compounds have been known since the early 20th century, and are in fact historic enough to have once been denoted by samarous (or samaric 3+), and yet few 40

structurally characterized examples exist outside of binary halides.101 Among the few well- characterized examples with SmII are [Sm(2.2.2-cryptand)]2+ that was recently obtained via an + + 88 103 unanticipated displacement of K from [K(2.2.2-cryptand)] , Sm(18-crown-6)(ClO4)2, where SmII is bound by both the crown ether and chelating perchlorate anions, and the sandwich 103 complex, [Sm(15-crown-5)2][ClO4]2, where the perchlorate anions are outer sphere. The cavity size of these sequestering agents can also be expanded to accommodate multiple metal sites within a single ligand as is made possible with dibenzo-30-crown-10. However, examples of such complexes are largely restricted to alkali metals where they have been utilized in the selective extraction of radioactive 135,137Cs+ from nuclear waste using calixarene-crowns.104 The purpose of this paper is to provide the synthesis, structure, and electronic properties of a SmII dibenzo-30-crown-10 complex, the product of its reaction with water, and a comparison of the electronic structure of both complexes.

5.2 Results and Discussion 5.2.1 Syntheses

II The synthesis of [Sm (DB30C10)][BPh4]2·2CH3CN requires rigorous exclusion of air and water from the reaction mixture. Although much like observed for [Sm(15-crown- 103 II 5)2][ClO4]2, crystals of [Sm (DB30C10)][BPh4]2·2CH3CN are surprisingly stable and can be stored under immersion oil for days in moist air without decomposition. Shockingly, the crystals can be placed directly in deoxygenated water for several days without decomposition. In contrast to the low-reactivity of the solid, it was observed that if the solvents used at any point in this synthesis contained even trace amounts of water that the initial red colour gradually faded, even in an argon-filled glovebox, and large, colorless

crystals would form prior to the addition of [Bu4N][BPh4]. X-ray analysis of these crystals III revealed that they consist of the dinuclear Sm complex, Sm2(DB30C10)(OH)2I4. One could speculate that this product forms via the reduction of water by SmII. However, SmII might first be oxidized by O2 followed by hydrolysis to create the Sm2O2 core (vide infra). Surprisingly, this compound decomposes much more rapidly than II [Sm (DB30C10)][BPh4]2·2CH3CN upon exposure to air.

5.2.2 Structural Features II II The crystal structure of [Sm (DB30C10)][BPh4]2·2CH3CN contains a Sm cation bound by all ten etheric oxygen atoms from the DB30C10 as shown in Figure 5.1. The geometry around the samarium center is best approximated as a sphenocorona, which is the most isotropic geometry for a ten-coordinate metal ion and likely results from the conformational flexibility of this large crown ether. The Sm‒O bond distances range from 2.645(4) to 2.791(3) Å, with the longest interactions occurring to the etheric oxygen atoms that are bound to the phenyl rings. 103 These distances are quite similar to those measured in [Sm(15-crown-5)2][ClO4]2, but are much longer than found in SmIII crown-ether complexes where distances of 2.4 to 2.5 Å have been measured.105,106 Of further interest is that the crown ether is folded into a so-called “Pac-Man” conformation that fully encapsulates the SmII cation. This conformation is reinforced by -

stacking of the two phenyl groups as depicted in Figure 5.1 with a centroid to centroid 𝝅𝝅 𝝅𝝅distance of 3.594(7) Å. The form adopted by the crown ether in II + + [Sm (DB30C10)][BPh4]2·2CH3CN differs substantially from that observed with Na and K cations whose ionic radius SmII lies between.107,108 When the crown ether hosts alkali metal cations the phenyl groups are staggered and not eclipsed, and a Pac-Man conformation is not adopted. Figure 2 shows a comparison of conformations of dibenzo-30-crown-10 as pure substance, with alkali metal cations, and with divalent samarium.

III The structure of the Sm hydrolysis product, Sm2(DB30C10)(OH)2I4, is comprised of a Sm2(OH)2 diamond core within a DB30C10 ligand with two, trans, iodide anions bound to each SmIII cation as illustrated in Figure 3. The core bears some similarities to those observed with other LnIII cations encapsulated within crown ethers where nitrate most often serves to complete the coordination sphere and balance charge.87 Each SmIII center is also ligated by three etheric oxygen atoms from the crown ether. Thus, the SmIII cations are bound by five oxygen atoms that are also found to be approximately co-planar. The SmIII‒O etheric oxygen atom distances are significantly shorter than those observed in [Sm(DB30C10)]2+ and average 2.464(7) Å. The SmIII‒O distance to the bridging hydroxyl group is 2.249(9) Å. The coordination sphere is completed by the two, trans iodide anions with expected long distances of 3.115(11) Å. These distances agree with those observed in other lanthanide compounds with peripheral iodide ligands.109,110 The accommodation of the 42

[Sm2(OH)2] core within the crown ether causes it to unfold from that observed in [Sm(DB30C10)]2+ into a twisted conformation as shown in Figure 3. This conformation is buttressed by hydrogen bonds between the bridging hydroxyl groups and the two etheric oxygen atoms that are not bound to the SmIII centres on each side of the crown ether with hydrogen-bonding distances of 3.096(7) Å. One could speculate that the weak SmIII‒I and SmIII‒O (etheric) interactions, the low coordination number of samarium, and the rather open pentagonal faces create opportunities for degradation of this complex in moist air. 5.2.3 Spectroscopic Properties

Divalent samarium displays rich photophysics that are quite distinct from that exhibited by the trivalent state. Excitation and photoluminescence spectra were obtained from crystals of

[Sm(DB30C10)][BPh4]2·2CH3CN at 77 K and are depicted in Figure 4. The excitation spectrum shows both broad and sharp features with the former being assigned to transitions between the 4f 6 7 5 1 ground state ( F0) and a 4f 5d configuration (e.g. an 4f → 5d transition), with the latter fine 5 features owing their origin to intra-f transitions. The f→f transitions occurring between the D0 7 → FJ (J = 0, 1, 2, 3) states produce the fine features that are characteristic of lanthanides. The

luminescence from [Sm(DB30C10)][BPh4]2·2CH3CN agrees well with that obtained from 103 II [Sm(15-crown-5)2][ClO4]2. A detailed discussion and review of Sm doped into a variety of hosts was used in the evaluation of [Sm(15-crown-5)2][ClO4]2 where it was shown that 5 7 the energy of the D0 → F0 transition in a variety of host lattices is typically found in the range -1 of ~14475 to 14700 cm . In [Sm(15-crown-5)2][ClO4]2 the energy of this transition is at 14636 ‒1 -1 cm , whereas in [Sm(DB30C10)][BPh4]2·2CH3CN it is found at 14577 cm placing both complexes intermediate in the range of energies. 5.2.4 Broken-Symmetry (DFT) Calculations

The first method used to understand the nature of the interaction between SmIII centers in

Sm2(DB30C10)(OH)2I4 was the Broken Symmetry (BS) approach. The molecule was optimized using two different spin polarizations, Δρ = 10 and Δρ = 0, to represent both the high (HS) and low spin (LS) states, respectively. In both cases, the geometrical parameters have practically the same values with some small deviations in the distances of the peripheral ligands. The Sm–Sm and Sm–I distances were 3.673 and 3.125 Å, respectively, while the distance from the Sm to the bridging oxygens were ca. 2.256 Å. The HS state shows a strong localization of the

unpaired electrons on each SmIII center (Δρ = 5.187) which is a common feature in lanthanide compounds due to the shielded nature of the 4f shell. The LS state is near the HS state, but lower in energy. The stabilization of this state could be explained at the DFT level using the broken- symmetry (BS) formalism to avoid the use of the extensive post–Hartree-Fock corrections. The BS wavefunction is formed by one Slater determinant constructed using molecular orbitals localized onto the two paramagnetic centers with opposite spins (Figure 5). This determinant is 2 not an S eigenstate but corresponds to a total Ms = 0 state that occasionally can be related to the

real Ms = 0 state in the so-called strong delocalization limit. However, in this case, a more proper approach would be to use the strong localization limit owing to the small overlap between the orbitals containing unpaired electrons.111 Using this approach, the calculated coupling constant between both states is 16.2 cm-1, which is consistent with antiferromagnetic coupling between both SmIII ions. However, considering the inherent overestimation of GGA functionals of these values, it is possible to conclude that in this molecule the metal centers are essentially uncorrelated. The energy of the BS state contains the most important contributions to the HS-LS separation, including the ligand-bridge effects often called ligand-spin polarization observed in Figure 5. 5.2.5 Ab-initio Calculations The electronic structures of both SmIII and SmII compounds were investigated by means of the CASSCF multiconfigurational approach. In the dinuclear compound the minimal active space containing the seven 4f orbitals for each metal center was sufficient to properly reproduce the electronic structure of the ground and low-lying excited states. In the case of SmII the 5d subshell was necessary to be included in the active space because its energy is modifiable by the local environment allowing for potential changes in the nature of the ground state. III III Sm2(DB30C10)(OH)2I4. As expected, in the dinuclear Sm compound, each Sm site shows an electronic structure very close to the free ion that is typical for trivalent lanthanides (Table D.1). As shown in Table D.1, the first five multiplets have pure contribution from the 6H spin-free ground term and are ordered in increasing energy. Additionally, the energy difference between multiplets is sufficient to confirm that all of them are well isolated as it generally occurs in trivalent lanthanide compounds. The analysis of the crystal-field parameters shows that both metal centers experience almost the same effects from the surrounded ligands (Table D.2). It is noticeable that the parameter is 0 2 44 𝐵𝐵

negative and corresponds to the largest value, indicating some axial character of the crystal-field. This corresponds, according to the form of the related Stevens operator = 3 ( + 1), to 0 2 a stabilization of the MJ =1/2 state as ground state. At the same time, the𝑂𝑂 non2 -axial𝐽𝐽𝑧𝑧 − parameters𝐽𝐽 𝐽𝐽 (mainly ) are also large which reduces the axial character and produces a more isotropic −1 crystal-field.𝐵𝐵2 This can be explained in terms of the charge distribution of the free lanthanide ion at the lowest J state, where SmIII is associated with a prolate distribution. This means that strong- field ligands in axial positions introduce high locale anisotropy in the electron density. However, in our case, the locale anisotropy is reduced because the weak-field iodide anions in axial positions in combination with the equatorial oxo ligands. This combination produces a more intense deformation of the charge distribution that can be diminished by replacing the axial weak-field ligands with strong-field ligands such as chlorine and fluorine anions. In that case, anisotropy is sought. II II [Sm (DB30C10)][BPh4]2. As previously mentioned, the electronic structure of Ln ions is richer than LnIII ions because of the well-known stabilization of the 5d subshell in the LnII ion that allows for crossover between 4f n+1 and 4f n5d1 configurations in the ground state.112 Thus, the ligand field plays a key role as it affects the energy of the unoccupied 5d orbitals more strongly than the 4f core orbitals. Consequently, the spectroscopic properties, particularly the 4f→5d transitions, can be tuned by modifying the ligand field. Based on these criteria, our calculations aimed to determine the ground state configuration and the optical properties of SmII in a deca-coordinated environment. The analysis of the multiconfigurational wave function 6 7 reveals that the 4f configuration is the ground state, which corresponds to the F0 ground term 7 -1 followed by the other FJ (J=1-6) terms distributed by about 7500 cm (Table D.3). The excited states with configurations of 4f 55d appear at ~14,500 cm-1 and form a broad band that extends to 25,000 cm-1 (starts at 26,200 cm-1 in the free ion112). The arrangement of these states agrees with the experimental absorption spectrum determined for divalent samarium. However, the quintuplet states start at ~15,000 cm-1 and extends beyond 25,000 cm-1 (Table D.3). A more 5 5 -1 detailed analysis on the emissive DJ states shows that the D0 state appears at 15,234 cm which corresponds to an overestimation of only 657 cm-1 (less than 0.1 eV) with respect to the 5 7 experiment ( D0→ F0 luminescent transition). Accordingly, a simplified photophysical mechanism underlying the observed emission might be based on the energy transfer from an 5 5 excited septuplet with 4f 5d configuration to a quintuplet DJ followed by a vibrational

relaxation and subsequent emission. This mechanism is also supported by the observed spectral overlap between the absorption and emission spectra (Figures 5.6). 5.2.6 Time-Dependent DFT

TD-DFT calculations were carried out to understand in depth the nature of the transitions in the [Sm(DB30C10)]2+ ion. Figure 7 shows the calculated absorption spectrum that can be associated with the room temperature absorption spectrum (Figure 5.4). Table 4 summarizes the primary transitions involved in the experimental measured region. It is remarkable that all of the transitions are combinations of 4f → 5d and 4f → π* metal—ligand charge transfer (MLCT) transitions in nature, differing only in the main contributor. The first three bands are dominated by the promotion of one electron from the 4f subshell to the 5d subshell, whereas the last three are dominated by CT transitions. The nature of the orbitals involved in these transitions can be largely classified into pure 4f, 5d, and unoccupied ligand orbitals (π*) (Figure 5.8). It is interesting that the Pac-Man conformation (Figure 5.3c) is stabilized by π–π interactions that are also responsible for the broadening in the absorption spectrum of divalent samarium by MLCT transitions from 4f electrons to low-lying unoccupied orbitals with bonding character. This stabilizing interaction can be observed in the Non-Covalent Interactions (NCI) surface shown in Figure S4 featuring a typical π–π stacking.

5.3 Conclusion Samarium, europium, and ytterbium are considered the classical examples of lanthanide elements that are readily reduced to the divalent state, and examples of compounds containing these Ln2+ ions have been known for well over a century. All three ions have been described as adopting a 4f n+1 configuration, i.e. 4f 6, 4f 7, 4f 14 for SmII, EuII, and YbII, respectively. The synthetic accessibility of the latter two examples is readily explained by the energetic stabilization imparted by self-exchange in spherically- symmetric ground states, and is in turn reflexed in the small, negative, standard reduction potentials. While SmII borders a half-filled 4f shell, this does not explain the rather moderate reduction potential for reducing SmIII to SmII in a rigorous manner. Furthermore, the classical description of SmII as adopting a 4f 6 ground state also appears to be overly simplistic. Our examination of the electronic structure of SmII in 46

II [Sm(DB30C10)][BPh4]2 and the electronic spectra of previous reported Sm complexes is consistent with a 4f 6 with low-lying 4f 55d1 states. The strong crystal-field effect produced by the 10-coordinate crown ether allows for the stabilization these 4f 55d1 states to be lowered by approximately 12,000 cm-1 with respect to the free ion. The changes from broad features at room temperature to fine features at 77 K conform to what would be expected for a state with an isolated 4f 6 configuration in the ground state. This description provides a much more satisfying explanation for the changes in the electronic absorption of SmII.

5.4 Methods and Characterizations

5.4.1 Experimental Syntheses

All manipulations were performed under Schlenk-type or glovebox

conditions with exclusion of air. In the case of [Sm(DB30C10)][BPh4]2, tetrahydrofuran (THF) and acetonitrile (MeCN) were obtained from pressurized silica columns and further dried (NaK for THF, 3 Å molecular sieves for MeCN) before use in reactions. For

Sm2(DB30C10)(OH)2I4, acetonitrile (Sigma) was used as obtained from the manufacturer without further purification or drying. SmI2 (99.9 %, Sigma), dibenzo-30-crown-10 (98%, Sigma), and tetrabutylammonium tetraphenylborate (99%, Sigma) were used as received without further purification. The general reaction scheme is observed in Figure 5.9.

[Sm(DB30C10)][BPh4]2. The synthesis of [Sm(DB30C10)][BPh4]2 was obtained

by mixing SmI2 (10.1 mg, 0.025 mmol) with [Bu4N][BPh4] (27.7 mg, 0.049 mmol) in 0.5 mL of THF and 2.0 mL of CH3CN until all solids were completely dissolved. In a separate vial DB30C10 (13.9 mg, 0.026 mmol) was dissolved in 2 mL of CH3CN. The DB30C10 solution was carefully added dropwise to avoid rapid precipitation, and the solution turned from dark green to bright red. After standing in glovebox freezer a -34 ℃ for 12 hours, bright red crystals suitable for single crystal crystallography formed in 78%

yield (Figure ES1). Anal. calcd. CHN for SmO10N2B2H43C80 (%): C, 68.27; H, 6.16; N, 1.99. Found (%). C, 68.12; H, 6.27; N, 1.31. High H and low N content due to loss of solvent and slight oxidation while under vacuum. Calculated CHN is more consistent with a 1 single acetonitrile molecule being retained. H NMR of [Sm(DB30C10)]I2 (600 MHz,

CD3CN): δ 8.4 (s, XH); 7.7 (m, XH), 7.2 (s, XH); 7.14 (s, XH); 6.9 (m, 8H); 6.1 (s, 1H); 4.2 47

(s, 8H); 4.12 (m, 8H); 4.07 (m, 8H); 3.8 (s, 8H); 3.7 (m, 8H); 3.6 (m, 8H); 2.8 (s, 8H). DB30C10 ligand: 6.9 (d, 8H); 4.1 (m, 8H); 3.8 (m, 8H); 3.7 (m, 8H); 3.6 (m, 8H); 1.97 (s, acetonitrile).

Sm2(DB30C10)(OH)2I4. A 5 mL Schlenk flask was charged with SmI2 (10.2 mg,

0.025 mmol) and DB30C10 (14.7 mg, 0.027 mmol). CH3CN (3 mL) was added to the Schlenk flask without further drying and stoppered. After approximately three to five days, colorless crystals suitable for single crystal X-ray diffraction formed. Rapid degradation of the crystals takes place shortly after forming, and the crystals change color 1 and decompose into a pale-yellow powder. H NMR (600 MHz, CD3CN): δ 6.9 (s, 8H); 4.0 (s, 8H); 3.7 (s, 16 H); 3.3 (s, 8H).

5.4.2 Crystallographic Studies

Single crystals of each compound were adhered to a Mitogen mount while contained within Parabar 10312 oil. The crystals were then aligned utilizing a Bruker D8 Quest X-ray diffractometer with an IμS X-ray source (MoKα, λ = 0.71073 Å) with a digital camera. Unit cells were determined with QUEST software and standard hemisphere collections were measured. The structure was solved in OLEX2 equipped with the SHELXTL program suite by the structure expansion solution method.35,36 These structural CIF files can be found in the CSD under the deposition numbers 1843892 and 1843893.

5.4.3 UV-Vis-NIR, Excitation, and Photoluminescence Spectroscopy UV-Vis-NIR measurements were performed using a Craic Technologies Microspectrophotometer and were performed by placing single crystals of II [Sm (DB30C10)][BPh4]2 onto a glass slide under Parabar 10312 oil. The II photoluminescence and excitation spectra of [Sm (DBC)][BPh4]2 were measured on an Edinburgh FLS980 fluorescence spectrometer equipped with a nitrogen cryostat cold finger dewar at 77 K. A housed 450 W Xe lamp/single grating (1800 λ/mm, 250 nm blaze). A Czerny-Turner monochromator was used to excite the solid sample. For the data collection, the spectrum was collected through a 610 nm long-pass color filter (emission) or a 475 nm long-pass color filter (excitation), single grating (1800 I/mm, 500 nm blaze) Czerny-Turner monochromator and then detected with a Peltier-cooled Hamamatsu R928 photomultiplier tube.

5.4.4 Computational Details Geometry optimization and electronic properties were obtained by using the Amsterdam Density Functional (ADF) code, where the scalar relativistic and spin–orbit effects were incorporated using the zero-order regular approximation (ZORA).38 The structure was fully optimized via the GGA (Generalized Gradient Approximation) BP86 functional including the standard Slater-type-orbital (STO) basis sets with triple-zeta quality plus double polarization functions (TZ2P) for all atoms.113,114 The same level of theory was used to analyze the interactions between the two SmIII centers by means of the broken-symmetry approach (BS) developed by Noodleman and co-workers.115–120 In this approach, the real multideterminant electronic state with unpaired α and β-electrons involving both metal centers is described by a single determinant Kohn-Sham wavefunction with electrons partially localized at different metal sites. Even though the energies and spin densities of the broken-symmetry state differ from those of the multideterminant antiferromagnetic state, the interaction parameter can be obtained from “mapping procedures” using spin-Hamiltonians and spin projection methods. 2+ For both Sm2(DB30C10)I4(OH)2 and [Sm(DB30C10)] , multiconfigurational Complete Active Space Self Consistent Field (CASSCF) calculations were performed in MOLCAS 8.2121 for a better understanding of the nature of the ground state and low-lying excited states. The dinuclear SmIII local properties were investigated making a diamagnetic substitution by replacing one SmIII site with LaIII. First, CASSCF calculation was performed employing an active space consistent in five electrons in seven 4f orbitals CAS(5,7). Spin-orbit coupling (SOC) was introduced in the second step by diagonalizing the SOC operator based on the optimized CASSCF wavefunctions by the RASSI (Restricted Active Space States Interaction) method. Scalar relativistic effects are considered by means of the Douglas-Kroll-Hess transformation and the spin-orbit integrals are calculated using the AMFI (Atomic Mean Field Integrals) method. In each case the all-electron ANO-RCC basis set with TZP quality was employed for every atom. The same method was also applied to [SmII(DB30C10)]2+. However, expanded active space was utilized. Divalent lanthanides are characterized by a 5d subshell closer to the 4f in energy making possible to have a crossover between the 4f n+1 and 4f n5d1 states. Therefore, the expanded active space consisted in six electrons in twelve orbitals CAS(6,12), although only one of the 5d orbitals was considered as truly active because the occupation numbers of the remaining orbitals is negligible. Using this active space allows to reproduce the proper ground

49 state and understand the nature of the low-lying states. Also, since dynamical correlation could play a relevant role in this system, second order perturbation theory corrections (CASPT2) were included in the final results. Additionally, time-dependent density functional theory (TD-DFT) was used as complement to the ab-initio electronic structure by calculating the absorption spectrum of [Sm(DB30C10)]2+. The ORCA 4.0.1 package122 was used to perform these calculations. Two functionals were tested for this aim, the parametrized CAMB3LYP and PBE0, which are both hybrid functionals. The former was discarded since it was unable to properly reproduce the ground state even though it is the most common functional used for these purposes. The PBE0 ground state provided an agreement with the CASSCF results.

5.5 Figures

Figure 5.1 A view of the structure of the [SmII(DB30C10)]2+ cation in II II [Sm (DB30C10)][BPh4]2·2CH3CN. The Sm cation is bound by all ten etheric oxygen atoms from the dibenzo-30-crown-10 ligand creating a sphenocorona geometry around the metal center. 50% probability ellipsoids are depicted.

Figure 5.2. An illustration of the structure of the dinuclear, SmIII dibenzo-30-crown-10 complex, III Sm2(DB30C10)I4(OH)2. The Sm cations are bridging by µ-OH anions that form hydrogen bonds with the crown ether. Two trans iodide anions complete the coordination sphere. 50% probability ellipsoids are depicted.

Figure 5.3. Depictions of the conformations of dibenzo-30-crown-10 in the solid state. The top figure (a) shows the form adopted by the crown ether without a metal ion. The second conformation (b) is adopted with alkali metal cations such as Na+ and K+. The final Pac-Man II conformation (c) is found with Sm in [Sm(DB30C10)][BPh4]2∙2CH3CN.

Figure 5.4. Absorption (left at 298 K), and excitation and photoluminescence (right at 77 K) II spectra of [Sm (DB30C10)][BPh4]2·2CH3CN. The absorption spectrum displays broad-band peaks indicative of 4f → 5d transitions; whereas the excitation and photoluminescence spectra obtained at 77 K reveal characteristic fine 4f → 4f transitions. The photoluminescence spectrum 5 7 shows the D0 → FJ (J = 0, 1, 2, 3) transitions.

Figure 5.5. Spin Density representation for the High-Spin (HS) and Broken-Symmetry (BS) states.

Figure 5.6. Energy diagram showing the lanthanide states (black horizontal lines) and the interval of energy in which the f→d transitions appears (green square).

II Figure 5.7. Calculated TDDFT absorption spectrum of [Sm (DB30C10][BPh4]2.

II Figure 5.8. Types of orbitals involved in excitation of [Sm (DB30C10)][BPh4]2.

Figure 5.9. Reaction scheme for forming the Sm(II) and Sm(III) products. The presence of water causes the oxidation of the Sm(II) site to form the dinuclear product.

CHAPTER 6

EXAMINING THE SECOND BREAK IN THE 5F SERIES: ADVANCEMENTS TOWARDS MOLECULAR CF(II)

6.1 Introduction Californium is the last element on the periodic table that milligram amounts are capable of being obtained. In addition to this, the chemistry of this element has been of significant interest because of the unexpected properties in its magnetism, spectroscopy, and structure.26,28 The element also represents a second break in the 5f series in which the divalent oxidation state becomes more readily accessible (the first to have a III→II redox potential below -2.0 V). This has led to the question of what makes the divalent oxidation state more accessible towards the end of the 5f series. Of the actinides, nobelium is the only one to have a stable divalent oxidation state that is preferential over the trivalent oxidation state. The reason the divalent oxidation state becomes preferred has not been readily evaluated. One reason for this is because it has not been experimentally determined if the electron is added to the d-orbital or f-orbital to give a 5f 9d1 or 5f10 configuration. The only divalent californium 123–126 compounds presented in the solid state thus far are CfCl2, CfBr2, CfI2, and SrB4O7:Cf(II). Of these, the latter of these divalent structures gave the most insight on this possible electron configuration. Peterson et al. explain that f-f transitions were observed at lower energies, but there were peaks observed at higher energies that could be from possible f-d transitions. The lack of experimental evidence has made the conclusion of this state difficult to assign. The assignment and study of these states can now be readily studied applied with modern computational details. However, these calculations are better suited for studying the bonds in molecular systems. Divalent actinide compounds of this type had not been known until the 83 discovery of the organometallic compound [K(2.2.2-cryptand)][C5H4SiMe3)3U] in 2013. Since, more divalent organometallic compounds have been prepared with thorium, uranium, neptunium, and plutonium.22,127–129 This work has proven vital by providing knowledge on the whereabouts of the added electron in the early actinides. Even so, the configuration of divalent compounds still seems to be interchangeable depending on the compound as a d- or f-orbital addition. This is more evident at neptunium.

Ideally, to better obtain knowledge of the addition of the electron orbital placement in the divalent actinides, americium would be the next ideal element to study. However, there are many difficulties with doing so with americium. The divalent organometallic compounds mentioned previously had a synthetic advantage over the later actinides: the availability of their pure metals. This is vital to air- and water-sensitive chemistry as the metal is often used as the precursor to synthesizing anhydrous starting materials in the f-elements.130,131 In addition to this, larger quantities (> 20 mg of metal content) are able to be obtained and handled for many of the earlier actinides up to plutonium. Obtaining pure metals of this context with later actinides from americium to californium are very expensive, difficult, and hazardous to obtain and handle. In the case of californium, it is the last element in which milligram quantities are able to be obtained. The energies emitted from americium and californium are also significant which limit reactions to reactions on scales less than 10 mg. Americium has a much higher III→II redox potential than californium (-2.3 V compared to -1.6 V) that would make obtaining its divalent oxidation state extremely challenging. Therefore, californium is the transactinide of choice for pursuing this feat. There have been few studies performed on californium, and even far less concerning its divalent oxidation state. However, as discussed in Chapter 3, we were able to successfully obtain electrochemical measurements on californium. Additionally, it was conferred that this divalent oxidation state was more readily accessible than samarium’s, and the addition of crown ethers to the solution further stabilized this oxidation state. Here we discuss methods for obtaining anhydrous californium starting materials as well as the structure, electronic, and spectroscopic properties of the first molecular divalent californium compound in the solid state.

6.2 Results and Discussion 6.2.1 Selection of Starting Materials

The synthesis of [Sm(DB30C10)][BPh4]2 was discussed in Chapter 5. This compound exhibited intriguing properties that would be of significant interest for characterization of californium (II). Another compound of interest with divalent californium is bis[hydrotris(3,5- * dimethylpyrazolyl)borato]samarium(II) (SmTp2 ) because of its sandwich-like structure as observed in Figure 6.1. This compound was first synthesized in 1993 by Takats et al.132

However, like most divalent samarium compounds, these reactions utilize SmI2 as the starting

material. As mentioned earlier, CfI2 has been prepared in the solid state on a microgram scale, but these conditions are not available in the laboratory currently equipped to hand 249Cf. Because of this, routes to synthesize these divalent molecules from the trivalent starting material had to be determined. When seeking starting materials for air-sensitive californium manipulations, ease of preparation, solubility, and purity are of significant importance due to the small amounts that are handled. The common trans uranium material to obtain is the chloride salt. Similar to

lanthanides, the anhydrous An/LnCl3 compounds are insoluble in most non-aqueous solvents. The An/Ln bromide and iodide salts have greater solubility. The preparation of the anhydrous materials can be obtained in multiple ways. These ways involve an in-situ reaction of the f-metal with the thionyl halide or by sublimation with the appropriate ammonium halide salt.133,134 However, these reactions would be very challenging and hazardous for milligram amounts of material with radioactive elements such as 243/241Am and 249Cf. Another such method of

preparing anhydrous iodides is by synthesizing the An/LnI2 compounds by the reaction with diiodoethane. It has been shown with the lanthanides that if excess metal is stirred in THF with 1,2-diiodoethane at room temperature in an inert atmosphere.135 This proved particularly effective with samarium, even in the presence of trace amounts of peroxide and water. It was then found that the metal quality was the most important factor. However, obtaining pure metals of the trans uranium is extremely difficult beyond plutonium. Due to the concerns mentioned above, the triflate salt was selected as the best available starting material. The lanthanide triflates are able to be made anhydrous in a relatively simple procedure in comparison to the other lanthanide salts. The triflate can be dried by heating under vacuum at 200 ℃ for 48 hours which does not from the oxy-salt, LnOX (X = Cl, Br, I).136 II * 6.2.2 Facile Synthesis of [Sm (DB30C10)][BPh4]2 and SmTp2 Due to the radioactivity of 249Cf, the reaction must be performed as quickly as possible in the fewest number of steps possible to avoid a great loss of material and radiolytic degradation. II The preparation of single crystals of [Sm (DB30C10)][BPh4]2 from Sm(otf)3 can be obtained in under 4 hours. This is a significant decrease in the crystallization time discussed in Chapter 5

that involves freezing overnight at -35 ℃. Anhydrous Sm(otf)3 was dissolved in THF and stirred over sodium metal in which the solution turned from colorless to dark purple after 30 minutes.

This solution was then mixed with a solution of Bu4NBPh4 and DB30C10 in acetonitrile. The

ratio of THF to acetonitrile needs to be a maximum of 1:4, because the product will precipitate as a powder. Once the mixture is obtained, crystals can be formed with four hours by vapor diffusion with diethyl ether. * SmTp2 is very insoluble in many solvents. If the reaction is done in pure THF or acetonitrile, an immediate purple precipitate is formed. Takas et al were able to obtain single * crystals of this compound by the diffusion of NaTp2 in THF to a solution of SmI2 in THF over

the course of a few weeks. However, the same product can be obtained by reducing Sm(otf)3 in THF over sodium metal then filtering the resultant purple solution into an equal volume * acetonitrile solution of KTp2 . From this method, crystals can be formed overnight (~12 hours) or a greater yield can be obtained by placing in the freezer at -35 ℃ overnight. The ability to produce single crystals of these compounds in a short period of time is ideal for studying divalent californium in a molecular molecule. 6.2.3 Reactions of 249Cf with DB30C10 and Tp*

The preparation of Cf(otf)3 is described below. The green solution was dissolved in THF and stirred over sodium metal. The reduction with sodium is a bit milder in comparison to potassium metal or KC8. Previous attempts to synthesize the [Cf(Crypt)][BPh4]2 that is isostructural to the LnCrypts discussed in Chapter 4 utilized the latter reductants. In the case of

KC8, much of the californium was lost in the graphite byproduct. And the reduction attempt with potassium for reaction with DB30C10 produced a brown precipitate after stirring for 1 hour. In the first reduction reaction with sodium, the californium solution in THF turned from pale green

to a golden amber color. Upon filtering this solution into a solution of Bu4NBPh4 and DB30C10 in acetonitrile, an insoluble pale green/yellow precipitate formed immediately. This solution was removed in vacuo and the precipitate partially dissolved in pyridine. Overnight, amorphous green-yellow crystals formed. Because it was thought that the californium reduction had not been to completion in the reaction with DB30C10, the same reduction process was repeated with recycled 249Cf. The californium solution was reduced over sodium metal for 5 hours in which the solution turned from pale green to pale yellow. Upon addition of this solution to acetonitrile, a similar suspension as describe in the DB30C10 reaction occurred. However, it appeared that some californium appeared to be in solution and was left to crystallized overnight. Colorless crystals of formed which is contributed to the contamination of Ca in the recycled 249Cf material. Upon

observing the spectroscopic properties of this material, it appeared that divalent Cf could be doped in the product as discussed below. 6.2.4 Spectroscopic Measurements Although crystal structures of the compounds could not be obtained, spectroscopic measurements were made to discern the oxidation state of californium. The luminescence of californium (III) is known, although some of the emission peaks were reassigned recently.28 Very little is known on californium (II) spectroscopic properties, except for the study on doped

SrB4O7. These spectra that were difficult to perceive as the data seems to be low resolution. Additionally, those measurements were performed on Cf(II) in a salt lattice rather than a molecular system. The UV-vis on the amorphous material had some ambiguity in the spectra. This is due to what seems to be the presence of Cf3+ and Cf2+ in the spectrum in Figure 6.2. This would support the indication that the reduction of californium did not go to completion. One interesting thing to observe is the luminescence of the product at 420 nm and 365 nm. Californium is known for luminescing green when excited at 420 nm, however, there is very little emission from this wavelength. The little bit of emission at 420 nm excitation could be from trace amounts of Cf3+ in the material. To our surprise, the compound produced bright white-blue emission when excited at 365 nm (Figure 6.3). It was originally thought that this observance was due to DB30C10. However, it was confirmed that DB30C10 does emit when exited at 365 nm. This blue emission could then possibly be attributed to Cf2+ as the emission slowly faded with time and was not able to be excited again because of oxidation. In the reaction with Tp*, colorless crystals formed. The product was determined to be * from an impurity of calcium to give the structure CaTp2 . The structure of these compounds has been previously prepared and studied.137 When the luminescence of the compound was measured, it was found to produce a red emission with a broad peak at 610 nm. This is interesting as many compounds with calcium have emissions at much higher energies.138–140 This emission wavelength is consistent with the emission wavelengths observed by Peterson et al in 2+ Cf :SrB4O7 where emission is observed between 500 and 700 nm (Figure 6.4). Additionally, the spectra measured in the two compounds do not match extremely well with calculated spectra 2+ * peaks a in Figure 6.5. This emission could be attributed to the doping of Cf in CaTp2 which single crystals exhibited a significant amount gamma energy. Additionally, the transition metal doping can be eliminated from this as well for multiple reasons. When, calcium is doped with

transition metals, emissions are not observed at wavelengths longer than 575 nm which was observed with manganese.139 The second reason for eliminating transition metals that could be of possible contamination (Cr, Mn, Fe, Ni) is due to the redox potentials of these metals. Of these metals, chromium has the most difficult redox potential from the Mx+ oxidation state to the M0 state at -0.91 V. Over the course of the reduction reaction which utilized sodium (-2.71 V), these transition metals would have been reduced to the elemental metal.

6.3 Conclusion Californium is not a heavily characterized element and many of these studies consist in the trivalent oxidation state. Although divalent californium has been prepared as the dihalides, little is known about this particular oxidation state, especially in molecular systems. Two attempts to synthesized divalent californium in a molecular system were made that did not successfully form single crystals to characterized in the solid state. However, some information has been obtained that could be of great insight to Cf2+ spectroscopy properties. It could be of possibility that divalent californium has a tunable luminescence that is similar to europium. With electrochemical studies141, and possible spectroscopic determinations, the path to the divalent californium in the molecular state may be more readily obtained in the near future.

6.4 Methods and Characterizations 6.4.1 Syntheses All manipulations were performed in an inert atmosphere with rigorous exclusion of air

and water (excluding Cf(otf)3 preparation). All reagents were used without any further purification. THF was further dried over NaK, and acetonitrile was further dried with Al2O3. 249 249 Preparation of Anhydrous Cf(otf)3. Caution! Cf is a Caution! Cf is a highly - radioactive material (t(1/2) = 351 years) with a specific activity of 4.1 Ci g . It has a very serious 249 245 health risk potential (γ emission = 0.388 MeV). Cf decays to Cm (t(1/2) = 8,500 years) from is 6.194 MeV alpha decay energy. These studies were performed in an argon glovebox, fitted to a HEPA filter system, in a lab that is continuously monitored with β and γ monitors. 5.0 mg Cf3+ as the chloride (0.02 mmol 249Cf) from Oak Ridge National Lab was dissolved in 2 mL of 6M HCl. The resulting green solution was dried down to a residue and dissolved in 1 mL of ultrapure

water. A 2 mL solution of recycled CfCl3∙xH2O was dissolved in 1 mL of ultrapure water. Then a

1 mL solution containing excess (15.6 mg, 0.162 mmol) of (NH4)2CO3 was added. The solutions were mixed in a 15 mL centrifuge tube in which a green precipitation formed. The precipitate was centrifuged, then the pellet washed twice with 1 mL of ultrapure water. Then 44 μL of a 50% solution of trifluoromethanesulfonic acid was added in which a gas evolved (presumably

CO2) and immediate dissolution of the solid. The resulting green solution was transferred to a Schlenk flask and carefully heated to a dry residue with a nitrogen stream to aid in evaporation of the solution. Once dried, the Schlenk flask was pumped into an argon atmosphere glovebox for 24 hours. The Schlenk flask was then placed under vacuum at 50 ℃ for 15 minutes, 100 ℃ for 15 minutes, and 130 ℃ for 4 hours to form the anhydrous Cf(otf)3 salt. II Synthesis of [Sm (DB30C10)][BPh4]2. Sm(otf)3 (3.1 mg, 0.0052 mmol) was dissolved in 1 mL of THF over a sodium mirror. The solution turned from colorless to purple after stirring for approximately 30 minutes. Then 4.1 mg of DB30C10 (0.0076 mmol) and 6.8 mg of

Bu4NBPh4 (0.0121 mmol) were dissolved in 1.5 mL of acetonitrile. The solutions were mixed in a 6 mL shell vial and vapor diffused with 3 mL of diethyl ether. Bright red crystals formed after a few hours at glovebox temperature. * Synthesis of SmTp2 . 4.6 mg (0.0077 mmol) of Sm(otf)3 was added to a sodium mirror with 600 μL of THF and stirred for six hours. The solution was then filtered into a solution containing 5.3 mg (0.0158 mmol) of KTp* in 495 μL of acetonitrile. The solution was capped and crystals formed after 12 hours at glovebox temperature.

Reaction with Dibenzo-30-crown-10. The Cf(otf)3 salt was dissolved in 1 mL of THF. The solution was then transferred over a sodium mirror and stirred for 30 minutes in which the solution turned from green to golden yellow. The solution was removed from the mirror and filtered slowly into a solution containing 23.1 mg of Bu4NBPh4 (0.0411 mmol) and 11.3 mg of DB30C10 (0.0211 mmol). The solution remained golden yellow. It was then vapor diffused with diethyl ether in which no crystals were produced over night. The solvents were vacuumed off and the green residue was partially dissolved in pyridine which produced green amorphous crystals. Reaction with KTp*. Cf(otf) (8.7 mg, 0.0125 mmol) with dissolved in 600 μL of THF producing a pale green solution. The solution was then transferred to a vial containing a sodium mirror and stirred for five hours. The solution turned from pale green to pale yellow. Upon addition to a 595 µL solution of KTp* (9.1 mg, 0.026 mmol), a pale yellow-green suspension

formed. The reaction was capped, and colorless crystals formed overnight which were 2+ * determined to be Cf :CaTp2 .

6.4.2 Single Crystal X-ray Crystallography Single crystal X-ray diffraction measurements were performed by selecting single crystals of compounds and adhering them to a Mitogen mount while contained within Parabar 10312 oil. The crystals were then aligned utilizing a Bruker D8 Quest X-ray diffractometer with an IμS X-ray source (MoKα, λ = 0.71073 Å) with a digital camera. Unit cells were determined with QUEST software and standard hemisphere collections were measured. The structure was solved in OLEX2 equipped with the SHELXTL program suite by the structure expansion solution method.35,36 6.4.3 UV-VIS and Luminescence Measurements UV-Vis and fluorescence measurements were measured by a CRAIC Technologies microspectrophotometer. Crystals of each sample were prepared on a glass slide in an inert atmosphere in Parabar 10312 oil.

6.5 Figures

Figure 6.1. SmII structures deemed suitable for Cf. The structures have interesting crystallographic properties that would be ideal to study with Cf. The syntheses of the two structures are relatively simple to synthesize compared to other divalent f-element compounds

Figure 6.2. The UV-Vis and emission spectra (inlet) of the californium product obtained from the reaction with DB30C10.

Figure 6.3. Emission from the amorphous Cf product at 420 nm (left) and 365 nm (right).

2+ * 2+ Figure 6.4. Emission of Cf :CaTp2 . The small amount dopant of Cf can give a red shift. The luminescence of Cf2+ could possibly be tunable like Eu2+.

Figure 6.5. Calculated TDDTF absorption spectrum of divalent californium within DB30C10

CHAPTER 7

CONCLUSION

A common mistake that is often made in sciences is the over grouping of similar items with the assumption that the entire group performs identically. In chemistry, this often done by assuming all elements in a group behave the same. This misconception had plagued f-elements for many decades from the conception that 4f and the latter 5f have identical chemistry and therefore were not worth studying due to radioactivity and cost. However, recent discoveries have led to error in this way of thinking. This began with the discovery of the unusual bonding observed in californium borate, followed shortly by the intriguing spectroscopic properties in californium dipicolinate. After this, small differences in other late actinides were noticed that 3- separated them from their lanthanide counterparts. The covalency detected in AmCl6 and the radial extension of the 5f-orbitals in comparison to the 4f-orbitals have furthered the gap in differences in the lanthanides and later actinides. The earlier actinides, in particular uranium, neptunium, and plutonium have been well characterized compared to elements from americium and beyond. This is attributed mostly to their applications in nuclear energy and weapons. An additional aspect to these elements in the easily accessible numerous oxidation states. Though the +2-oxidation state has become known for all the lanthanides recently, the chemistry of these ions behaves very differently from the +3 state as the divalent oxidation state can only be obtained in non-aqueous conditions with the exclusion of samarium, europium, and ytterbium. Even so, with the exception of europium, these oxidation states only last for short periods of time in aqueous conditions. When in the divalent oxidation state, lanthanides can partake in chemistry that is not so predictive as the trivalent oxidation state. The addition of an electron can go into d- or f-orbitals. In the later actinides, the divalent oxidation state is so scarcely studied that it is uncertain what the properties these elements may display. For these reasons, this work attempts to further delve in the chemistry of f-elements in non-aqueous conditions. For elements in which characteristics are dominated by coordination chemistry, it is imperative to gain a complete understanding of their behavior beyond aqueous conditions. Furthermore, the examination of the divalent oxidation state could provide even more fundamental knowledge and disparity between the 4f and 5f elements.

This work begins by determining a difference in the primary chemistry of the elements. In Chapter 3, a difference was observed in the coordination chemistry of lanthanides and actinides in non-aqueous media. This was shown by the small difference in covalency observed

in [Am(TpyNO2)(NO3)3(H2O)]∙THF compared to the [Ln(TpyNO2)(NO3)3(H2O)]∙THF systems. Additionally, such little research has been performed on later actinides that this structure proved to be the first characterization of a 10-cooridinate americium compound in which the ionic radius was calculated. It is these small differences that lead to applications that have great importance like nuclear waste separations. The next step was to examine the redox chemistry of the f-elements in greater detail in non-aqueous conditions. This was done by examining the better studied lanthanides, samarium, europium, and ytterbium with different solvents and crown ethers. Vast information was learned from these experiments. All of the lanthanides have been obtained in the divalent oxidation state with reduction by potassium which has a reduction potential of -2.93 V even though some lanthanides have been calculated to have reduction potentials greater than -3.0 V. This is made possible by either the calculated reduction potentials being incorrect, or by changing the environment the reduction potential can be shifted. The latter seems to be the case as the study showed that in THF with a crown ether, reduction potentials can shift by -0.5 V. This would allow potassium to have enough reduction potential to reduce every lanthanide to the divalent oxidation state. Additionally, another difference was observed between the chemistry of lanthanides and later actinides. In the lanthanides, where sized and coordination number dominate the properties, an unexpected outcome was observed when ytterbium showed to have a more reversible III→II reduction than europium. This could be explained by the size difference and the different structure obtained by ytterbium versus samarium and europium in the 2.2.2-cryptand system. However, in one of the first electrochemistry experiments on californium, it produced an intriguing observance as well. Californium has a calculated reduction potential of -1.6 V, while samarium has a reduction potential of -1.55 V. It was found that the reduction potential of californium was measured to be slightly lower than samarium’s while behaving electrochemically like ytterbium. The identicalness could not be for the same reason as californium (III) is similar in size to samarium. It was therefore determined that the 5f orbitals had to be a cause for this different outcome.

Chapter 5 observed the chemistry of samarium, the analog to californium, in the divalent oxidation state in greater detail. The work showed the interactions of samarium with dibenzo-30- crown-10 in the trivalent and divalent oxidation state. Two interesting things were obtained from the study. The first being that divalent samarium can out-compete alkali metals for encapsulation in the crown ether due the great stability achieved in the ligand metal charge transfer and ᴨ- stacking of the benzyl groups. Secondly, the configuration of Sm2+ was confirmed to be 4f6, but had an enormous shift in the energy levels of the 5d-orbtials of 12,000 cm-1. This study lead to the importance of examining the system with Cf2+. The ending of this work, attempted to isolate divalent californium in the solid state in a molecular system. Since only the dihalide salts are known with the later actinides, isolation of a molecular Cf(II) system would allow for a substantial amount of knowledge to be obtained as many perceive californium as the second transition in the actinide series. However, attempts were unsuccessful. The failed attempts lead to a possible conclusion on the spectroscopic properties of divalent californium. The luminescence of divalent californium in molecular environments may be tunable as observed in some divalent europium compounds. In addition to this, it was noted that under the same reaction conditions, the redox chemistry of californium does not behave like its analog samarium. This work provides a greater understanding to the differences in the redox and coordination chemistry of the lanthanides and actinides. By exploring simple systems that were deemed unnecessary at one-point, small fundamental behaviors were obtained. This work concludes that although the lanthanides and actinides are both in the f-block, each element should be treated as unique, especially in the 5f series.

APPENDIX A

TABLES FROM CHAPTER 1

Table A.1. Breakdown of heavy metal single crystal structures currently in the CSD as of 2018. Element Crystal Structures Pt 20881 U 6332 Pu 163 Am 24 Cm 5 Bk 2 Cf 3

APPENDIX B

TABLES FROM CHAPTER 3

Table B.1. Values for ligands used in ionic radii calculations. Values shown are in angstroms.

Method Terpyridine N Nitrate O Water O Shannon 1.46 1.35 1.36 Nitrate - 1.35 - La-Based 1.39 1.35 1.25 Ce-Based 1.38 1.35 1.24

Table B.2. Ionic radii of 10-cooridinate f-elements by different calculation methods. Values are in units of angstroms.

Vol. Unit Nitrate Known Shannon La-based Ce-Based Cell Based La 2859.4(14) 1.27 1.239(2) 1.272(2) 1.27 1.275(2) Ce 2835.1(8) 1.25 1.212(2) 1.241(2) 1.244(2) 1.25 Nd 2829.6(7) 1.1832(16) 1.2161(16) 1.2142(16) 1.2182(16) Sm 2809.9(8) 1.155(3) 1.190(3) 1.187(3) 1.191(3) Eu 2789.7(7) 1.1393(8) 1.1734(8) 1.1713(8) 1.1753(8) Tb 2791.8(9) 1.123(3) 1.162(3) 1.155(3) 1.159(3) Am 2825.3(12) 1.181(2) 1.221(2) 1.213(2) 1.217(2)

Bond ρ(r) G(r) V(r) H(r) |V/G| H/ρ * Eu-Nterpy 0.0421 0.0413 -0.0415 -0.0003 1.0067 -0.0066 * Eu-ONO2 0.0420 0.0468 -0.0459 0.0009 0.9801 0.0222

Eu-OH2 0.0447 0.0545 -0.0523 0.0022 0.9595 0.0492 * Am-Nterpy 0.0445 0.0446 -0.0465 -0.0020 1.0443 -0.0444

Table B.3. - Continued Bond ρ(r) G(r) V(r) H(r) |V/G| H/ρ * Am-ONO2 0.0431 0.0486 -0.0492 -0.0006 1.0132 -0.0148

Am-OH2 0.0474 0.0573 -0.0577 -0.0003 1.0058 -0.0070 *Remaining bonds are shown in Table B.4 Table B.4. Bonding parameters derived from the Bader’s theory of AIM for the metal-ligand BCPs. The electron density, the Lagrangian kinetic energy, potential energy, and energy density were evaluated in the bond critical point. All parameters are in a.u.

Bond ρ(r) G(r) V(r) H(r) |V/G| H/ρ

Eu-Nterpy 0.0405 0.0388 -0.0389 -0.0002 -1.0045 -0.0043

Eu-Nterpy 0.0407 0.0391 -0.0393 -0.0002 -1.0051 -0.0049

Eu-Nterpy 0.0421 0.0413 -0.0415 -0.0003 -1.0067 -0.0066

Eu-ONO2 0.0420 0.0468 -0.0459 0.0009 -0.9801 0.0222

Eu-ONO2 0.0362 0.0401 -0.0378 0.0023 -0.9425 0.0638

Eu-ONO2 0.0400 0.0443 -0.0430 0.0013 -0.9700 0.0333

Eu-ONO2 0.0367 0.0404 -0.0384 0.0021 -0.9484 0.0568

Eu-ONO2 0.0337 0.0371 -0.0343 0.0028 -0.9252 0.0824

Eu-ONO2 0.0420 0.0469 -0.0459 0.0009 -0.9799 0.0224

Eu-OH2 0.0447 0.0545 -0.0523 0.0022 -0.9595 0.0492

Am-Nterpy 0.0442 0.0436 -0.0457 -0.0021 -1.0479 -0.0473

Am-Nterpy 0.0441 0.0433 -0.0454 -0.0020 -1.0473 -0.0465

Am-Nterpy 0.0445 0.0446 -0.0465 -0.0020 -1.0443 -0.0444

Am-ONO2 0.0368 0.0406 -0.0399 0.0007 -0.9824 0.0195

Am-ONO2 0.0416 0.0468 -0.0472 -0.0003 -1.0069 -0.0078

Am-ONO2 0.0431 0.0486 -0.0492 -0.0006 -1.0132 -0.0148

Am-ONO2 0.0381 0.0423 -0.0418 0.0005 -0.9891 0.0120

Am-ONO2 0.0408 0.0456 -0.0458 -0.0002 -1.0038 -0.0043

Am-ONO2 0.0388 0.0431 -0.0428 0.0003 -0.9939 0.0068

Am-OH2 0.0474 0.0573 -0.0577 -0.0003 -1.0058 -0.0070

APPENDIX C

TABLES FROM CHAPTER 4

Table C.1. Values from cyclic voltammetry experiments involving Ln/AnCrypt.

Ipc/Ipa (100 |Ecath-Eanod| + 0 E1/2 (vs. Fc/Fc ) (V) ΔE (vs. E ) (mV) mV/sec) (mV)

Eu -0.48 +524 1.18 340

Yb -0.99 +705 0.75 90

Sm -1.66 +520 2.64 696

Cf -1.53 +715 2.10 448

APPENDIX D

TABLES FROM CHAPTER 5

Table D.1. Calculated excitation spectrum for both Sm(III) centers in Sm2(DB30C10)(OH)2I4.

Center I Center II Term J E(cm-1) E(cm-1) 0.0 0.0 6H 5/2 199.5 202.8 399.6 401.6

1098.9 1108.3 6H 7/2 1273.4 1278.9 1353.7 1340.6 1440.4 1446.2

2385.9 2392.4 2524.1 2528.3 6H 9/2 2604.4 2602.4 2639.2 2633.0 2735.9 2740.5

3800.1 3804.8 3921.1 3929.8 6H 11/2 4002.2 4001.1 4053.7 4044.8 4092.3 4093.3 4196.2 4200.3

Table D.1. -Continued. Center I Center II Term J E(cm-1) E(cm-1) 5391.6 5401.2

5467.3 5473.9 5526.8 5521.1 5590.5 5579.9 5661.4 5664.4

5743.4 5752.1 6731.8 6725.6 6851.7 6855.7 6H, 6G, 6F 15/2 6904.8 6919.7 6982.7 6978.8 7081.8 7076.9 7163.7 7167.9 7281.3 7278.7

Table D.2. Ab-initio computed crystal-field parameters that show the axial character of the III dinuclear Sm complex, Sm2(DB30C10)(OH)2I4.

Center I Center II k k k Q B q B q -2 -0.62 10.10 -1 10.68 5.77 2 0 -16.87 -16.97 1 3.79 2.20 2 -4.47 -4.66 -4 -0.59 0.38

Table D.2.-Continued Center I Center II k k k Q B q B q -3 -2.23 -3.63 -2 1.79 -1.77 -1 -2.33 -2.30 4 0 -0.35 -0.30 1 -1.39 1.51 2 0.73 0.32 3 -1.78 0.60 4 -0.25 -8.79

Table D.3. Calculated excitation spectrum for [Sm(DB30C10)][BPh4]2. State Energy (cm-1) State Energy (cm-1)

7 5 F0 0.0 D0 15234.4

116.5 17793.1 7 5 F1 162.4 D1 17794.9 561.4 17897.1 571.4 20484.2 1136.6 20511.2 7 5 F2 1254.6 D2 20550.9 1356.8 20606.3 1443.6 20884.7 1513.6 24657.3 1752.0 24668.5 1781.6 24686.1 7 5 F3 1963.9 D3 24686.9 2084.5 24697.6 2334.5 24699.1

Table D.3.-Continued State Energy (cm-1) State Energy (cm-1) 2481. 24701.9

2781.8 2851.1 2895.8 3045.2 7 F4 3057.2 3929.5 3980.5 4490.9 4496.2

4786.7 4787.7 4883.8 4934.6 7 F5 5113.8 5256.3 5261.3 5264.6 5486.1 5497.6 5659.9

5676.7 5737.1 5741.1 6046.9 6117.1

Table D.3.-Continued. State Energy (cm-1) State Energy (cm-1) 7 F6 6458.6 6475.2 6821.0 6822.2 7144.7 7144.7

APPENDIX E

CRYSTALLOGRAPHIC DATA

Table E.1. Crystal data and structure refinement for Nd1. Identification code 1857533 Empirical formula C25H24N7NdO13 Formula weight 774.75 Temperature/K 28 Crystal system monoclinic Space group P21/c a/Å 11.5357(16) b/Å 17.691(2) c/Å 14.3924(19) α/° 90 β/° 105.555(3) γ/° 90 Volume/Å3 2829.6(7) Z 4 3 ρcalcg/cm 1.819 μ/mm-1 1.917 F(000) 1548.0 Crystal size/mm3 0.358 × 0.28 × 0.197 Radiation MoKα (λ = 0.71073) 2Θ range for data collection/° 4.328 to 66.588 Index ranges -17 ≤ h ≤ 17, -27 ≤ k ≤ 27, -21 ≤ l ≤ 22 Reflections collected 77048 Independent reflections 10868 [Rint = 0.0568, Rsigma = 0.0337] Data/restraints/parameters 10868/0/420 Goodness-of-fit on F2 1.059 Final R indexes [I>=2σ (I)] R1 = 0.0276, wR2 = 0.0593 Final R indexes [all data] R1 = 0.0401, wR2 = 0.0635 Largest diff. peak/hole / e Å-3 1.00/-1.48

Table E.2. Bond lengths for Nd1. Atom Atom Length/Å Atom Atom Length/Å Nd1 O2 2.5437(14) C16 C28 1.394(3) Nd1 O3 2.5837(14) C16 C30 1.384(3) Nd1 O5 2.5590(14) C16 N34 1.470(2) Nd1 O10 2.5759(14) C17 C19 1.490(2) Nd1 O11 2.5394(14) C17 N25 1.356(2) 78

Table E.2. – Continued. Atom Atoms Length/Å Atom Atom Length/Å Nd1 N13 2.6033(15) C17 C32 1.399(2) Nd1 N22 2.9794(16) C19 N31 1.348(2) Nd1 N25 2.6038(16) C19 C37 1.400(2) Nd1 N26 2.9838(15) C20 C39 1.391(3) Nd1 N31 2.5959(16) C20 C43 1.390(3) Nd1 O35 2.5950(14) C23 C27 1.406(3) Nd1 O0AA 2.4619(14) C23 C36 1.484(2) O2 N12 1.2812(19) C23 C44 1.405(3) O3 N22 1.2716(19) C24 N25 1.345(2) O4 N22 1.236(2) C24 C40 1.391(3) O5 N22 1.2767(19) C27 C30 1.389(3) O8 N34 1.235(2) C28 C44 1.388(3) C9 C29 1.487(2) C29 C39 1.395(2) C9 N31 1.355(2) C32 C42 1.389(3) C9 C33 1.398(2) C33 C36 1.399(2) O10 N26 1.276(2) C36 C37 1.396(2) O11 N26 1.284(2) C40 C42 1.394(3) N12 O18 1.227(2) C41 C43 1.395(3) N12 O35 1.278(2) O1 C2 1.451(2) N13 C29 1.356(2) O1 C3 1.448(2) N13 C41 1.350(2) C2 C1AA 1.541(3) O14 N26 1.225(2) C0AA C1AA 1.537(3) O15 N34 1.235(2) C0AA C3 1.517(3)

Table E.3. Crystallographic details of Eu1.

Identification code 1854046-1854048 Empirical formula C25H24EuN7O13 Formula weight 782.47 Temperature/K 28.0(1) Crystal system monoclinic Space group P21/c a/Å 11.5374(15) b/Å 17.564(2) c/Å 14.321(2) α/° 90 β/° 106.004(5) γ/° 90 Volume/Å3 2789.7(7)

Table E.3. – Continued. Z 4 3 ρcalcg/cm 1.863 μ/mm-1 2.332 F(000) 1560.0 Crystal size/mm3 0.233 × 0.201 × 0.170 Radiation MoKα (λ = 0.71073) 2Θ range for data collection/° 3.76 to 99.07 Index ranges -24 ≤ h ≤ 24, -36 ≤ k ≤ 37, -30 ≤ l ≤ 30 Reflections collected 195028 Independent reflections 28524 [Rint = 0.0606, Rsigma = 0.0343] Data/restraints/parameters 28524/0/511 Goodness-of-fit on F2 1.037 Final R indexes [I>=2σ (I)] R1 = 0.0274, wR2 = 0.0559 Final R indexes [all data] R1 = 0.0399, wR2 = 0.0602 Largest diff. peak/hole / e Å-3 1.62/-1.45

Table E.4. Bond lengths for Eu1. Atom Atom Length/Å Atom Atom Length/Å Eu1 O2 2.4886(8) C15 C20 1.4875(12) Eu1 O3 2.5407(7) C15 C37 1.3973(12) Eu1 N4 2.9304(8) N17 O36 1.2205(11) Eu1 O5 2.5439(8) C18 C24 1.4804(12) Eu1 N6 2.5666(8) C18 C29 1.3999(12) Eu1 O7 2.5722(8) C18 C32 1.3979(12) Eu1 N8 2.5644(8) C20 C33 1.3969(12) Eu1 O10 2.5081(7) N21 O23 1.2294(11) Eu1 N11 2.5420(8) N21 C28 1.4664(11) Eu1 O14 2.4869(8) N21 O31 1.2272(12) Eu1 N17 2.9379(8) C22 C46 1.3886(12) Eu1 O1 2.4192(8) C24 C26 1.3932(12) O2 N12 1.2797(10) C24 C33 1.3963(12) O3 N4 1.2614(10) C25 C44 1.3945(12) N4 O10 1.2710(10) C27 C28 1.3832(13) N4 O19 1.2345(10) C27 C29 1.3891(12) O5 N17 1.2673(10) C28 C40 1.3904(13) N6 C25 1.3536(11) C30 C37 1.3915(12) N6 C38 1.3411(11) C30 C46 1.3898(13) O7 N12 1.2707(11) C32 C40 1.3916(12) N8 C15 1.3505(11) C34 C38 1.3944(13)

Table E.4. – Continued. Atom Atom Length/Å Atom Atom Length/Å N8 C22 1.3442(12) C34 C39 1.3878(13) N11 C13 1.3490(11) C39 C44 1.3875(12) N11 C20 1.3458(11) O16 C43 1.4417(12) N12 O35 1.2224(10) O16 C45 1.4461(12) C13 C25 1.4826(11) C41 C43 1.5146(14) C13 C26 1.3968(12) C41 C48 1.5368(13) O14 N17 1.2812(11) C45 C48 1.5320(14)

Table E.5. Crystal data and structure refinement for Am1. Identification code 1854050 Empirical formula C25H24AmN7O13 Formula weight 873.51 Temperature/K 28.0(1) Crystal system monoclinic Space group P21/c a/Å 11.531(3) b/Å 17.690(5) c/Å 14.383(4) α/° 90 β/° 105.630(5) γ/° 90 Volume/Å3 2825.3(12) Z 4 3 ρcalcg/cm 2.054 μ/mm-1 2.798 F(000) 1688.0 Crystal size/mm3 0.326 × 0.215 × 0.136 Radiation MoKα (λ = 0.71073) 2Θ range for data collection/° 4.33 to 75.69 Index ranges -19 ≤ h ≤ 19, -30 ≤ k ≤ 30, -24 ≤ l ≤ 23 Reflections collected 95448 Independent reflections 14727 [Rint = 0.0824, Rsigma = 0.0627] Data/restraints/parameters 14727/0/420 Goodness-of-fit on F2 1.042 Final R indexes [I>=2σ (I)] R1 = 0.0373, wR2 = 0.0666 Final R indexes [all data] R1 = 0.0662, wR2 = 0.0731 Largest diff. peak/hole / e Å-3 1.17/-1.80

Table E.6. Bond lengths for Am1. Atom Atom Length/Å Atom Atom Length/Å Am1 O2 2.553(2) O14 N39 1.275(3) Am1 O3 2.583(2) C18 C25 1.400(4) Am1 O4 2.561(2) C18 C42 1.400(4) Am1 N5 2.589(2) C19 C32 1.385(4) Am1 N6 2.578(2) C20 C24 1.402(4) Am1 N7 2.589(2) C20 C36 1.390(4) Am1 O11 2.589(2) O21 N40 1.229(3) Am1 O14 2.536(2) C23 N40 1.468(3) Am1 O15 2.604(2) C23 C41 1.391(4) Am1 N17 2.982(2) C23 C46 1.389(4) Am1 N39 2.987(2) C24 C42 1.490(4) Am1 O1 2.467(2) C25 C26 1.398(4) O2 N10 1.274(3) C25 C30 1.482(4) O3 N17 1.270(3) C27 C29 1.393(4) O4 N17 1.270(3) C27 C43 1.391(4) N5 C19 1.349(4) C28 C29 1.394(4) N5 C24 1.355(3) C30 C31 1.407(4) N6 C8 1.352(3) C30 C38 1.403(4) N6 C42 1.352(3) C31 C41 1.383(4) N7 C28 1.357(3) C32 C36 1.391(4) N7 C45 1.345(3) O33 N39 1.226(3) C8 C26 1.392(4) C38 C46 1.391(4) C8 C28 1.494(4) C43 C45 1.393(4) N10 O15 1.278(3) O16 C35 1.455(4) N10 O37 1.226(3) O16 C47 1.445(3) O11 N39 1.278(3) C34 C44 1.530(4) O12 N40 1.235(3) C34 C47 1.515(4) O13 N17 1.242(3) C35 C44 1.538(4)

Table E.7. Crystal data and structure refinement for SmDBC. Identification code DisserationFile Empirical formula C48H84N9O20Sm3 Formula weight 1558.29 Temperature/K 137.52 Crystal system triclinic Space group P-1 a/Å 9.8629(8) b/Å 20.8521(16)

Table E.7. – Continued. c/Å 21.0553(16) α/° 90.0358(17) β/° 98.8686(17) γ/° 101.3616(17) Volume/Å3 4192.5(6) Z 4 3 ρcalcg/cm 2.469 μ/mm-1 4.258 F(000) 3124.0 Crystal size/mm3 ? × ? × ? Radiation MoKα (λ = 0.71073) 2Θ range for data collection/° 4.266 to 69.926 Index ranges -15 ≤ h ≤ 15, -32 ≤ k ≤ 32, -33 ≤ l ≤ 33 Reflections collected 97614 Independent reflections 35064 [Rint = 0.0970, Rsigma = 0.1290] Data/restraints/parameters 35064/0/1086 Goodness-of-fit on F2 1.124 Final R indexes [I>=2σ (I)] R1 = 0.0741, wR2 = 0.1867 Final R indexes [all data] R1 = 0.1265, wR2 = 0.2089 Largest diff. peak/hole / e Å-3 5.70/-5.16

Table E.8. Bond lengths for SmDBC. Atom Atom Length/Å Atom Atom Length/Å Sm1 O3 2.441(3) C32 C41 1.500(6) Sm1 O4 2.432(3) O33 C93 1.397(8) Sm1 O6 2.509(3) O33 C112 1.460(6) Sm1 O7 2.434(3) N34 O71 1.217(7) Sm1 O8 2.560(3) O35 C79 1.423(7) Sm1 O9 2.473(4) O35 C131 1.451(7) Sm1 O23 2.547(4) O37 C76 1.447(7) Sm1 O31 2.512(4) O37 C106 1.429(6) Sm1 N34 2.957(5) O38 N44 1.272(6) Sm1 O36 2.552(4) C39 C46 1.387(8) Sm1 O43 2.689(5) C39 C77 1.387(7) Sm1 N58 3.002(6) O40 C118 1.420(8) Sm2 O27 2.487(4) O40 C120 1.436(9) Sm2 O28 2.502(3) O42 N63 1.275(7) Sm2 O42 2.473(4) O43 N58 1.244(6) Sm2 O45 2.486(4) N44 O49 1.269(6) Sm2 O47 2.503(4) N44 O80 1.200(6) 83

Table E.8. – Continued. Atom Atom Length/Å Atom Atom Length/Å Sm2 O53 2.454(4) O45 N60 1.263(6) Sm2 O57 2.498(4) C46 C84 1.400(7) Sm2 O59 2.472(5) O47 N95 1.259(7) Sm2 N60 2.904(5) O48 C98 1.437(7) Sm2 O69 2.509(4) O48 C108 1.447(7) Sm2 O72 2.507(5) C51 C86 1.398(7) Sm2 N95 2.891(5) C51 C87 1.412(8) Sm3 O5 2.373(4) O53 N95 1.293(7) Sm3 O10 2.419(4) C54 C75 1.496(8) Sm3 O15 2.468(4) O56 C116 1.429(8) Sm3 O18 2.515(4) O56 C144 1.415(8) Sm3 O35 2.520(4) O57 N60 1.270(7) Sm3 O38 2.495(4) N58 O67 1.207(7) Sm3 N44 2.904(5) N60 O85 1.228(6) Sm3 O48 2.472(4) O61 C97 1.405(7) Sm3 O49 2.454(4) O61 C119 1.439(7) Sm3 N90 2.898(5) N63 O72 1.274(7) Sm3 O102 2.481(5) N63 O82 1.223(7) O4 C21 1.419(6) C65 C88 1.405(8) O4 C54 1.433(5) C65 C110 1.370(7) O6 C32 1.443(6) O68 C113 1.410(7) O6 C75 1.445(6) O68 C121 1.425(9) O8 C22 1.443(6) C70 C76 1.527(8) O8 C30 1.428(6) C73 C92 1.379(9) O9 N58 1.283(7) C73 C127 1.385(8) O11 C20 1.433(6) C74 C78 1.411(9) O11 C39 1.374(5) C74 C84 1.391(8) O13 C46 1.363(6) C77 C78 1.368(7) O13 C70 1.428(7) C79 C94 1.485(8) O14 C88 1.362(6) C83 C101 1.510(8) O14 C104 1.435(6) C86 C99 1.395(9) O15 N90 1.260(7) C87 C96 1.380(7) O16 C51 1.369(6) C88 C105 1.368(7) O16 C100 1.444(6) N90 O102 1.277(7) O17 C41 1.435(6) N90 O117 1.209(7) O17 C87 1.365(6) C91 C119 1.465(8) O18 C83 1.426(7) C92 C114 1.374(8) O18 C109 1.468(7) C93 C104 1.510(8) O19 C73 1.376(6) N95 O111 1.207(6) O19 C91 1.433(7) C96 C122 1.386(7) 84

Table E. 8. – Continued. Atom Atom Length/Å Atom Atom Length/Å C20 C30 1.499(6) C97 C116 1.485(9) C21 C22 1.508(6) C98 C131 1.473(9) O23 N34 1.273(5) C99 C122 1.362(9) N24 O27 1.253(6) C100 C121 1.482(8) N24 O52 1.219(6) C103 C127 1.396(9) N24 O69 1.274(6) C103 C129 1.346(11) N25 O28 1.251(7) C105 C123 1.378(9) N25 O55 1.217(6) C106 C120 1.472(10) N25 O59 1.251(6) C108 C109 1.502(10) O26 C65 1.374(6) C110 C115 1.407(9) O26 C101 1.429(7) C112 C144 1.428(11) O29 C92 1.400(7) C113 C118 1.453(11) O29 C94 1.451(8) C114 C129 1.389(10) O31 N34 1.260(6) C115 C123 1.355(11)

Table E.9. Crystal data and structure refinement for Sm2(DB30C10)I4(OH)2. Identification code 1843892 Empirical formula C48H69N3O19I6Sm3 Formula weight 2204.51 Temperature/K 230.0 Crystal system monoclinic Space group C2/c a/Å 37.219(2) b/Å 19.2589(12) c/Å 21.1740(13) α/° 90 β/° 114.2980(10) γ/° 90 Volume/Å3 13833.1(15) Z 8 3 ρcalcg/cm 2.117 μ/mm-1 5.254 F(000) 8272.0 Crystal size/mm3 0.55 × 0.27 × 0.22 Radiation MoKα (λ = 0.71073) 2Θ range for data collection/° 4.222 to 55.184 Index ranges -48 ≤ h ≤ 48, -25 ≤ k ≤ 25, -27 ≤ l ≤ 27 Reflections collected 105747 Independent reflections 16015 [Rint = 0.0410, Rsigma = 0.0260]

Table E.9. – Continued. Data/restraints/parameters 16015/0/717 Goodness-of-fit on F2 0.974 Final R indexes [I>=2σ (I)] R1 = 0.0218, wR2 = 0.0525 Final R indexes [all data] R1 = 0.0790, wR2 = 0.0711 Largest diff. peak/hole / e Å-3 1.09/-0.57

Table E.10. Selected bond lengths for Sm2(DB30C10)I4(OH)2. Atom Atom Length/Å Atom Atom Length/Å Sm1 Sm2 3.6635(7) O13 C32 1.424(11) Sm1 I1 3.1279(11) O9 C25 1.413(12) Sm1 I2 3.1027(11) O9 C24 1.369(12) Sm1 O11 2.240(9) O3 C5 1.411(12) Sm1 O12 2.249(8) O3 C4 1.418(11) Sm1 O6 2.484(7) O4 C10 1.376(12) Sm1 O7 2.448(7) O4 C11 1.435(10) Sm1 O5 2.448(7) C8 C7 1.366(11) Sm2 I3 3.1253(10) C8 C9 1.379(14) Sm2 I4 3.1043(11) C2 C1 1.480(12) Sm2 O11 2.266(9) C10 C5 1.404(10) Sm2 O12 2.240(8) C10 C9 1.359(14) Sm2 O2 2.463(7) C39 C38 1.501(12) Sm2 O1 2.473(7) C17 C18 1.485(13) Sm2 O10 2.465(7) C19 C24 1.410(11) 1 Sm3 Sm3 3.6625(12) C19 C20 1.413(14) Sm3 I6 3.1274(11) C26 C25 1.509(13) 1 Sm3 I5 3.1028(11) C31 C31 1.411(18) Sm3 O18 2.244(6) C31 C30 1.412(12) Sm3 O19 2.276(6) C35 C34 1.507(12) Sm3 O16 2.450(7) C30 C29 1.403(13) Sm3 O14 2.463(7) C5 C6 1.362(13) Sm3 O15 2.492(7) C15 C16 1.499(12) O6 C15 1.452(10) C36 C37 1.492(13) O6 C14 1.429(11) C28 C27 1.491(13) O7 C17 1.455(10) C14 C13 1.516(13) O7 C16 1.436(11) C12 C11 1.467(12) O16 C37 1.442(11) C7 C6 1.388(13) O16 C38 1.445(11) C3 C4 1.526(12) O2 C2 1.457(11) C33 C32 1.485(13) 1 O2 C3 1.441(11) C29 C29 1.35(2) O14 C34 1.439(11) C45 N3 1.118(15) 86

Table E.10. – Continued. Atom Atom Length/Å Atom Atom Length/Å O14 C33 1.434(11) C45 C46 1.428(14) O1 C1 1.459(11) C41 C40 1.369(13) O1 C28 1.432(11) C41 C42 1.381(14) 1 O10 C26 1.430(11) C40 C40 1.38(2) O10 C27 1.448(10) C24 C23 1.355(13) O15 C35 1.424(10) C20 C21 1.399(15) O15 C36 1.440(10) C43 C44 1.453(15) O17 C39 1.438(12) C23 C22 1.382(15) O17 C40 1.386(11) C44 N2 1.104(14) O5 C12 1.435(12) C47 C48 1.436(16) O5 C13 1.441(11) C47 N1 1.109(16) O8 C19 1.357(12) C21 C22 1.361(13) 1 O8 C18 1.413(11) C42 C42 1.38(2) O13 C31 1.333(10)

II Table E.11. Crystallographic details for [Sm (DB30C10)][BPh4]2. Identification code 1843893 Empirical formula C40H43BNO5Sm0.5 Formula weight 703.74 Temperature/K 230.0 Crystal system monoclinic Space group C2/c a/Å 25.160(3) b/Å 12.4775(17) c/Å 24.342(3) α/° 90 β/° 112.582(3) γ/° 90 Volume/Å3 7055.7(15) Z 8 3 ρcalcg/cm 1.325 μ/mm-1 0.893 F(000) 2928.0 Crystal size/mm3 0.626 × 0.224 × 0.094 Radiation MoKα (λ = 0.71073) 2Θ range for data collection/° 4.41 to 53.884 Index ranges -31 ≤ h ≤ 32, -15 ≤ k ≤ 15, -30 ≤ l ≤ 30 Reflections collected 42263 Independent reflections 7394 [Rint = 0.2936, Rsigma = 0.1601]

Table E.11. – Continued. Data/restraints/parameters 7394/0/430 Goodness-of-fit on F2 0.979 Final R indexes [I>=2σ (I)] R1 = 0.0592, wR2 = 0.0938 Final R indexes [all data] R1 = 0.1370, wR2 = 0.1187 Largest diff. peak/hole / e Å-3 0.92/-1.19

II Table E.12. Selected bond lengths for [Sm (DB30C10)][BPh4]2. Atom Atom Length/Å Atom Atom Length/Å Sm1 O2 2.688(3) C26 C25 1.387(7) 1 Sm1 O2 2.688(3) C33 C34 1.396(7) Sm1 O3 2.791(3) C33 B1 1.634(8) 1 Sm1 O3 2.791(3) C33 C38 1.386(7) 1 Sm1 O1 2.709(3) C34 C35 1.389(8) Sm1 O1 2.709(3) C18 C17 1.379(8) 1 Sm1 O4 2.708(3) C6 C5 1.380(8) Sm1 O4 2.708(3) C6 C7 1.361(8) Sm1 O5 2.645(4) C23 C24 1.386(8) 1 Sm1 O5 2.645(4) C27 C32 1.399(7) O2 C4 1.383(6) C27 C28 1.399(7) O2 C3 1.452(6) C27 B1 1.650(7) O3 C9 1.387(6) C32 C31 1.392(7) O3 C10 1.453(6) C28 C29 1.393(7) O1 C2 1.424(6) C8 C7 1.383(8) O1 C1 1.431(6) C10 C11 1.491(7) O4 C11 1.441(6) C25 C24 1.363(8) O4 C12 1.446(7) C12 C13 1.477(8) O5 C14 1.438(6) C35 C36 1.378(9) 1 O5 C13 1.432(6) C14 C1 1.487(7) C4 C9 1.395(7) C17 C16 1.357(9) C4 C5 1.379(7) C36 C37 1.368(9) C21 C22 1.392(7) C29 C30 1.369(8) C21 C26 1.396(6) C31 C30 1.378(9) C21 B1 1.650(7) C38 C37 1.375(8) C22 C23 1.387(7) C16 C15 1.366(9) C9 C8 1.373(7) C20 C15 1.376(9) C19 C18 1.381(7) C39 C0AA 1.414(8) C19 B1 1.642(8) C0AA N2B 1.075(14) C19 C20 1.391(7) C0AA N2A 1.244(19) C3 C2 1.478(7)

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Evaluation of the Effect of Water, Oxygen, and Peroxide Content, Preparative Methods, and the Activation of Samarium Metal. J. Org. Chem. 2012, 77 (7), 3049–3059.

(136) Kobayashi, S. (Shu); Anwander, R. (Reiner). Lanthanides : Chemistry and Use in Organic Synthesis; Springer, 1999.

(137) Michel, O.; Martin Dietrich, H.; Litlabø, R.; To, K. W.; Maichle-Mo, cilia; Anwander, R. Tris(Pyrazole)Borate Complexes of the Alkaline-Earth Metals: Alkylaluminate Precursors and Schlenk-Type Rearrangements. Organometallics 2012, 31, 23.

(138) Zhang, X.; Pan, Q.; Kim, S. Il; Yu, Y. M.; Seo, H. J. Temperature Dependence of the Luminescence of Calcium-Magnesium Phosphate Ca 3 Mg 3 (PO 4 ) 4 :Eu 2+ , a Blue- Emitting Material for White Light-Emitting Diodes. 2014.

(139) Brightwell, J. W.; Ray, B.; Buckley, C. N. Preparation, Crystal Growth and Luminescence in Calcium Sulphide. J. Cryst. Growth 1982, 59 (1–2), 210–216.

(140) Groenink, J. A.; Hakfoort, C.; Blasse, G. The Luminescence of Calcium Molybdate The Luminescence of Calcium Molybdate; 1979; Vol. 54.

(141) Marsh, M. L. Non-Aqueous Electrochemical Studies of Lanthanide and Actinide Complexes, Florida State University, 2018.

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BIOGRAPHICAL SKETCH

Frankie D. White Personal Information Birthdate: September 13, 1989 Birthplace: Jackson, Mississippi Citizenship: United States of America Current Career Graduate Research Assistant Thomas Albrecht-Schmitt, Principle Investigator Florida State University Department of Chemistry and Biochemistry 95 Chieftan Way Tallahassee, FL 32304 Professional Experience and Education Doctor of Philosophy in Chemistry, Florida State University 2015 – Present Advisor: Dr. Thomas Albrecht-Schmitt Thesis: Redox-Coordination Chemistry of Lanthanides and Transuranics in Non- aqueous Media Master of Science in Environmental Toxicology, University of South Alabama 2013 – 2015 Advisor: Dr. Richard Sykora Thesis: Tetracyanoplatinate and Dicyanoaurate Compounds as Potential Chemical Sensors Bachelor of Science in Chemistry and Biochemistry 2008 – 2013 Advisor: Dr. Richard Sykora Research: Lanthanide Cyanometallate Coordination Chemistry

Honors and Awards Graduate Assistantship, University of South Alabama, ($42,000) 2013 – 2015 National American Chemical Society Undergraduate Research Assistant, 2013 Inorganic Chemistry, Spring Chemistry Departmental Award, University of South Alabama ($1,500) 2013 Presidential Scholarship, University of South Alabama, ($12,500) 2008 – 2012

Research Interests To date my research interests have focused on advancing understanding of how f-element coordination chemistry is unique in comparison to the other elements in the periodic table. To achieve this, I am inserted in comparatively characterizing the redox properties, stability, and reactivity of the transuranic elements (Np – Cf) vs lanthanides and d-block transition elements.

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Another research interest is centered on identifying structure to function relationships for plutonium and the other actinide elements using single crystal X-ray crystallography.

Publications Submitted and in Press Citations determined as of June 2, 2018 using SciFinder = 14 16. 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, Dave E. Hobart, and Thomas Albrecht-Schmitt. “Examination of Structure and Bonding in 10-Coordinate Europium and Americium Terpyridyl Complexes. Accepted to Inorganic Chemistry with revisions, July 2018.

15. Frankie D. White, Cristian Celis-Barros, Jillian Rankin, Eduardo Solis-Cespedes, David Dan, Jasmine N. Colangelo, Dayan Paez-Hernandez, Ramiro-Arratia-Perez, Alyssa Gaiser, Yan Zhou, and Thomas E. Albrecht-Schmitt. “Encapsulation of Samarium(II) within a Crown Ether, Electronic Structure, and Hydrolytic Reactivity”. Submitted to Chemical Sciences, July 2018.

14. Samantha K. Cary, Jing Su, Shane S. Galley, Thomas E. Albrecht-Schmitt, Enrique R. Batista, Maryline G. Ferrier, Stosh A. Kozimor, Veronika Mocko, Brian L. Scott, Cayla E. Van Alstine, Frankie D. White, Ping Yang. “A series of Dithiocarbamates for Americium, Curium, and Californium”. Accepted to Dalton Transactions July 2018

13. Todd N. Poe, Frankie D. White, Vanessa Proust, Eric M. Villa, Matthew J. Polinski. IV IV “[Ag2RE(Te2O5)2]SO4 (RE = Ce or Th ): A New Purely Inorganic d/f-Heterometallic Cationic Material”. Inorganic Chemistry. 2018, 57, 4816-4819. Citations(0) DOI: 10.1021/acs.inorgchem.8b00504

12. Matthew L. Marsh, Frankie D. White, Shane Galley, and Thomas Albrecht-Schmitt. “Comparison of the Electronic Properties of f7, f8, and f9 Lanthanides with Formally Isoelectronic Actinides”. Handbook on the Physics and Chemistry of Rare Earths. 2018. Citations(0) DOI: 10.1016/bs.hpcre.2018.01.001.

11. Nikolay P. Tsvetkov, Edgar Gonzalez-Rodriguez, Audrey Hughes, Gabriel dos Passos Gomes, Frankie D. White, Febin Kuriakose, and Igor V. Alabugin. “Radical Alkyne peri-Annulation Reactions for the Synthesis of Functionalized Phenalenes, Benzanthrenes, and Olympicene”. Angewandte Chemie. 2018, 130, 1-6. Citations(1) DOI: 10.1002/anie.201712783 102

10. Matthew L. Marsh, Frankie D. White, Wes M. Potter, and Thomas E. Albrecht-Schmitt. “Synthesis and Structural Diversity in Lanthanide and Actinide Complexes”. Solid-State Chemistry Textbook Chapter. Submitted June 2017.

9. C. J. Evoniuk, G. P. Gomes, M. Ly, F. D. White, and I. V. Alabugin. “Coupling Radical Homoallylic Expansions with C-C Fragmentations for the Synthesis of Heteroaromatics: Quinolines from Reactions of o-Alkenylarylisonitriles with Aryl, Alkyl, and Perfluoroalkyl Radicals”. Journal of Organic Chemistry. 2017, 85, 4265-4278. Citations(5) DOI: 10.1021/acs.joc.7b00262

8. J. C. Gaitor, M. S. Zayas, D. J. Myrthil, F. White, J. M. Hendrich, R. E. Sykora, R. A. O’Brien, J. T. Reilly, A. Mirjafari. “Crystal structure of a methimazole-based ionic liquid”. Acta Crystallographica Section E: Crystallographic Communications 2015, 71, o1008-o1009. Citations(1) DOI: 10.11007/S2056989015022136

7. F. White, L. N. Pham, K. R. Xiang, R. Thomas, P. Vogel, Z. Assefa, and R. E. Sykora. “Synthesis, structures, and photoluminescence properties of lanthanide dicyanoaurates containing dimeric aurophilic interactions,” Inorganica Chimica Acta, 2014, 414, 240- 249. Citations(4) DOI: 10.1016/j.ica.2014.02.012

6. F. White and R. E. Sykora. “Crystal structure of catena-poly[[aqua-(2,2’:6’,2’’- terpyridine-κ3N,N’,N’’)-cobalt(II)]-μ-cyanido-κ2N:C-[dicyanidoplatinum(II)]-μ-cyanido- 2 κ C:N][Co(C15H11N3)(H2O){Pt(CN)4}]” Acta Crystallographica Section E: Structure Reports 2014, 70(9), m322-m323. Citations(0) DOI: 10.1107/S1600536814017425

5. F. White and R. E. Sykora. “Crystal structure of di-μ-benzato-κ4 O:O'-bis[aqua(benzato- κO)(benzato-κ2 O,O')(2,2':6',2''-terpyridine-κ3 N,N',N'')europium(III)]-benzoic acid (1/2)” Acta Crystallographica Section E: Structure Reports 2014, 70(9), m328-m329. Citations(0) DOI: 10.1107/S1600536814018182

4. F. White and R. E. Sykora. “Crystal structure of bis(2,2’:6’,2”-terpyridine-κ3 N,N’,N”)nickel(II) dicyanoaurate (I)” Acta Crystallographica Section E: Research Communications 2014, 70, 519-521. Citations(0) DOI: 10.1107/S1600536814024672

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3. J. L. Bradfield, R. A. Braun, F. White, J. M. Hendrich, R. E. Sykora, and A. Mirjafari. “Crystal Structure of triphenyl(vinyl)phosphonium tetraphenylborate” Acta Crystallographica Section E: Structure Reports 2014, E70, o1143. Citations(0) DOI: 10.1107/S1600536814021357

2. A. T. Thames, F. D. White, L. N. Pham, K. R. Xiang, and R. E. Sykora. “Tris(thiocyanato-κN)tris(triphenylphosphine oxide-κO)europium(III)−(nitrato-κ2O,O') bis(thiocyanato-κN)tris(triphenylphosphine oxide-κO) europium(III) (1/1)” Acta Crystallographica Section E: Structure Reports 2012, E68, m1530. Citations(1) DOI: 10.1107/S1600536812047472

1. L. N. Pham, A. T. Thames, F. D. White, K. R. Xiang, and R. E. Sykora. “Tris(thiocyanato-κN)tris(triphenylphosphine oxide-κO)terbium(III),” Acta Crystallographica Section E: Structure Reports 2012, E68, m1531. Citations(2) DOI: 10.1107/S1600536812047289

Oral Presentations 5. F. D. White, M. L. Marsh, D. E. Hobart, and T. E. Albrecht-Schmitt. Florida State Candidacy Seminar. “Exploration of redox coordination chemistry in f-elements in aqueous and non-aqueous solutions”. Fall 2017.

4. F. D. White, M. L. Marsh, D. E. Hobart, and T. E. Albrecht-Schmitt. National American Chemical Society Meeting, Washington, D. C. “Exploring Redox Coordination Chemistry in Transuranic Elements with Various Crown Ethers and Cryptands through Lanthanides”. August 2017.

3. F. D. White. Florida State University Inorganic Seminar. “Low Valent f-block Coordination Chemistry”. Spring 2017.

2. F. D. White. University of South Alabama Thesis Seminar. “Tetracyanoplatinate and Dicyanoaurate Compounds as Potential Chemical Sensors”. Summer 2015

1. F. D. White, M. Pischek, L. Pham, and R. E. Sykora. University of South Alabama Department of Chemistry Student Seminar. “Structural and Spectroscopic Characterization of Lanthanide Dicyanoaurates”. Spring 2012.

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Poster Presentations 1. National American Chemical Society Meeting, San Francisco, CA. “Modeling Redox Coordination Chemistry for Transuranium Elements”. Spring 2017. 2. National American Chemical Society Meeting, Dallas, TX. “Structural, Photoluminescence, and Energy Transfer Properties of Lanthanide Dicyanoaurates containing Terpyridine Ligands”. Spring 2014. 3. National American Chemical Society Meeting, San Diego, CA. “Effect of Energy Transfer on Various Lanthanide Ions by Au2Pt4 Clusters”. Spring 2017.

Expertise Single Crystal X-ray Diffraction Spectroscopic Measurements (Absorbance, Luminescence) Air sensitive manipulations and syntheses (Glovebox, Schlenk Line Techniques) Handling Radioactive Materials (Certified for a Category II Nuclear Facility) Recycling of Transuranic Elements (237Np, 239Np, 239Pu, 243Am, 249Cf)

Teaching Graduate Teaching Assistant for General Chemistry Laboratory, Department of Chemistry, University of South Alabama 2014 – 2015. Graduate Teaching Assistant for Honor’s General Chemistry Laboratory, Department of Chemistry and Biochemistry, Florida State University 2015 – 2016. Graduate Teaching Assistant for General Chemistry, Department of Chemistry and Biochemistry, Florida State University 2016.

Mentorship and Outreach Mentor to Undergraduate Research Assistants, Florida State University, 2016 – Present Invited 2nd Grade Black History Science Talk, Tallahassee School of Math and Science, February 2017. Science Olympiad, Springhill University, 2014.

Favorite Quote at Florida State University “Do you know what now means?” – Dr. Thomas Albrecht-Schmitt

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