Determination of Thermodynamic Properties of Select

DETERMINATION OF THERMODYNAMIC PROPERTIES OF SELECT

URANYL COMPOUNDS

A Dissertation

Submitted to the Graduate School

of the University of Notre Dame

in Partial Fulfillment of the Requirements

for the Degree of

Doctor of Philosophy

Melika Sharifironizi

Peter C. Burns, Director

Graduate Program in Civil and Environmental Engineering and Earth Sciences

Notre Dame, Indiana

November 2017

Melika Sharifironizi

DETERMINATION OF THERMODYNAMIC PROPERTIES OF SELECT

URANYL COMPOUNDS

Abstract

Melika Sharifironizi

The continued environmental, technological, and industrial importance of actinide materials sustains the need for understanding their behavior from the nanoscale to macroscale. Understanding the physical and chemical properties of actinides is a necessary step towards predicting their long- and short-term behavior in different environments. Determination of thermodynamic properties of uranium minerals and associated compounds is fundamental for understanding the alteration pathways of uranium in geological systems, predicting transport of uranium in the environment, and ultimately developing a geologic repository for spent nuclear fuel in the US. Towards this end, I focus on experimentally determining the thermodynamic properties of uranyl compounds in this research.

The work described in this dissertation explores two major topics in actinide geochemistry. The first theme is focused on determination of standard- state thermodynamic properties of the zippeite-group minerals from the uranyl sulfate family that form on uranium mine wastes and that may be important in

Melika Sharifironizi nuclear waste disposal. The result of the experimental thermochemistry of four members of the zippeite group with zippeite-type uranyl sulfate sheets and different interlayer cations is presented in this work. Pure zippeite, natrozippeite, cobaltzippeite, and zinczippeite mineral analogs were hydrothermally synthesized and thoroughly characterized using X-ray diffraction and analytical chemistry methods (ICP-OES, TGA). The standard-state enthalpy of formation of each sample was determined using high-temperature calorimetry. Solubility experiments were carried out for zippeite to measure its solubility product and hence calculate its standard state Gibbs free energy of formation. Calorimetric data revealed that there is a linear relationship between the formation enthalpies from oxides and the acidity of cation oxides, as well as the ionic radius of charge- balancing alkalis, indicating the importance of the nature and coordination environment of the interstitial complex in determining the thermodynamic properties of these minerals.

In the second component of this work, I studied the chemical thermodynamics of selected members of the family of nanoscale uranyl peroxide cage clusters. Uranyl nanoclusters, that are proposed to be energetic intermediates between dissolved aqueous uranyl species and uranyl minerals, have potential importance in an advanced nuclear fuel cycle and environmental transport of actinides following nuclear accidents. Their structures are composed of uranyl polyhedra bridged by bidentate peroxide, and other types of bridges in addition to peroxide ligands in many cases, that are charge-balanced by metal cations. Six members of this family were selected, synthesized under ambient

Melika Sharifironizi conditions, and well-characterized using single crystal X-ray diffraction, ICP-

OES, ESI-MS, and TGA. Our calorimetric data shows that the enthalpies of formation of the cluster compounds from oxides become more negative as the charge on the cluster increases. The results obtained from this work demonstrate that the energetics of uranyl peroxide cluster crystals are largely driven by the alkali cation oxide thermodynamics, rather than negatively charged uranyl cages.

The data obtained from this work proposes trends between enthalpies of formation of select uranyl phases and their chemical properties related to their crystal structure. The findings may be used to predict the energetics of other members of that family, or on a larger scale to better understand and improve the nuclear fuel cycle.

For Dad, A., and H.

CONTENTS

Figures ...... v Tables ...... viii Acknowledgments ...... ix Chapter 1: Introduction ...... 1 1.1 Uranium Discovery and its applications ...... 1 1.2 Nuclear Fuel Cycle ...... 2 1.3 U6+ Crystal Chemistry and its mineralogy ...... 6 1.3.1 The zippeite family ...... 9 1.4 Polyoxometalates ...... 12 1.4.1 Transition Metal Polyoxometalates ...... 12 1.4.2 Uranium Polyoxometalates ...... 13 1.5 Thermochemistry of uranyl compounds ...... 17 1.6 Thesis Overview ...... 22 1.7 References ...... 23 Chapter 2: Thermodynamic Studies of Zippeite, a Uranyl Sulfate Common in Mine Wastes ...... 31 2.1 Abstract ...... 31 2.2 Introduction ...... 32 2.3 Materials and Methods ...... 34 2.4 Results and Discussion ...... 38 2.5 Acknowledgments ...... 49 2.6. Supporting Information ...... 50 2.7 References ...... 55 Chapter 3: Investigation of the Structural Stability of Zippeite-Group Minerals using High-Temperature Calorimetry ...... 58

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3.1 Abstract ...... 58 3.2 Introduction ...... 59 3.3 Materials and Methods ...... 62 3.4 Results and Discussion ...... 65 3.5 Acknowledgements ...... 73 3.6 Supporting Information ...... 74 3.7 References ...... 75 Chapter 4: Energetic Studies of uranyl peroxide nanoclusters ...... 79 4.1 Abstract ...... 79 4.2 Introduction ...... 80 4.3 Experimental Methods ...... 83 4.3.1 Synthesis ...... 83 4.3.2 Characterization ...... 84 4.3.3 High-Temperature Calorimetry ...... 86 4.4 Results ...... 87 4.5 Discussion ...... 91 4.6 Acknowledgements ...... 93 4.7 Supporting Information ...... 94 4.8 References ...... 102 Chapter 5: Summary and future work ...... 106 5.1 Importance of this work ...... 106 5.2 Thermodynamic Studies of Zippeite, a Uranyl Sulfate Common in Mine Wastes ...... 107 5.3 Investigation of the Structural Stability of Zippeite-Group Minerals using High-Temperature Calorimetry ...... 109 5.4 Energetic Studies of Uranyl Peroxide Nanoclusters ...... 110 5.5 References ...... 112

FIGURES

Figure 1. 1. The open nuclear fuel cycle (Ewing, 2006)...... 3

Figure 1. 2. Schematic representation of formation of secondary uranyl minerals under oxidizing conditions (Finch and Ewing, 1992)...... 7

Figure 1.3. Coordination environments for hexavalent uranium (yellow). The uranyl ion (a) is coordinated by four, five or six equatorial ligands (red) forming square (b and e), pentagonal (c and f), and hexagonal (d and g) bipyramids (Burns, 2005; Burns et al., 1997)...... 8

Figure 1.4. The hierarchy of inorganic uranyl compounds given with frequencies of each structure type (Lussier et al., 2016)...... 9

Figure 1.5. The zippeite type uranyl-sulfate sheet ...... 11

Figure 1.6. Linkages of hexagonal bipyramids by sharing of a peroxide edge (top), U60 cluster (bottom) (Sigmon et al., 2009c)...... 16

Figure 1.7. Schematic of the AlexSYS and assembly of glassware for drop solution calorimetry in a molten oxide solvent (Navrotsky, 2014) ...... 20

Figure 1.8. Change in the baseline of the calorimeter over a reaction...... 22

Figure 2.1 Plot of experimental measurements of the solubility of zippeite in terms of dissolved U (diamonds), K (squares), and S (triangles) as a function of time from undersaturation at pH 2.8 (a), pH 3.5 (b), and pH 4 (c), and from supersaturation at pH 3.5 (d)...... 44

Figure 2.2 pH- Log a (K+) diagram for the system metaschoepite- zippeite at sulfate concentration of 11.75 × 10-3 with activity coefficients calculated using an extended Debye–Hückel equation...... 47

Figure 2.3. The sheet of uranyl and sulfate polyhedra in the zippeite structure. Uranyl pentagonal bipyramids are shown in yellow and sulfate tetrahedra are shown in blue. Green circles correspond to the hydroxyl groups (Burns et al., 2003)...... 50

Figure 2.4. Zippeite PXRD pattern. Blue trace corresponds to synthetic zippeite, and red peaks correspond to peak positions and heights calculated from the structural model (Burns et al., 2003)...... 50

Figure 2.5. DSC-TGA curve for synthetic K-zippeite...... 51

Figure 2.6. Raman spectrum of zippeite. Black represents the spectra of the starting material for solubility experiments and blue corresponds to the spectrum of the residue of solubility experiments. The Raman spectrum of the solubility experiment residue is slightly shifted for clarity...... 52

Figure 3.1. The sheet of uranyl and sulfate polyhedra in the structure of a) zippeite, b) natrozippeite, and c) cobaltzippeite and zinczippeite. Uranyl pentagonal bipyramids are shown in yellow and sulfate tetrahedra are shown in blue. Green circles correspond to the hydroxyl groups (Burns et al., 2003)...... 61

Figure 3.2. Enthalpy of formation of zippeite, natrozippeite, cobaltzippeite, and zinczippeite from oxides normalized to the number of uranium as a function of acidity of the oxides, Smith scale. (Smith, 1987) ...... 71

Figure 3.3. Enthalpy of formation of zippeite, natrozippeite, cobaltzippeite, and zinczippeite from oxides as a function of ionic radius of interlayer cation...... 72

Figure 3.4- Powder X-ray diffraction patterns for Natrozippeite, Cobaltzippeite, and Zinczippeite...... 74

Figure 4.1. Polyhedral representations of seven uranyl nanoclusters. Uranyl 4- polyhedra are shown in yellow, and blue polyhedra represent [P2O7] and 2- [HPO3] ...... 83

Figure 4.2. Enthalpy of formation from oxides as a function of the number of alkali cations per formula unit. Squares represent data from this study and triangles represent those of previous works (Armstrong et al., 2012; Kubatko et al., 2003; Tiferet et al., 2014)...... 90

Figure 4.3. Enthalpy of formation from oxides versus charge density. Squares represent data from this work and triangles represent previous works (Armstrong et al., 2012; Tiferet et al., 2014)...... 91

Figure 4.4. Polyhedral representation of U26Pp6. Yellow polyhedra are [UO2(O2)3] 4- and purple polyhedra are [P2O7] ligands...... 94

Figure 4.5. Thermogravimetric curves of Li43Na7-U24Pp12 , Na40K7- U24Pp12, K31Li4-U26Pp6, K26-U22PO3, K32U28PO3, and Li28U28...... 95

Figure 4.6. ESI-MS spectra of Li36Na12-U24Pp12 , Na43K8- U24Pp12, K31Li4-U26Pp6, K26-U22PO3, K32U28PO3, and Li28-U28...... 97

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TABLES

Table ‎2.1 Thermodynamic cycles for formation of zippeite from its elements and binary oxides ...... 40

Table ‎2.2 Formation of zippeite from uo3, compreignacite, and clarkeite ...... 48

Table ‎2.3 Chemical compositions of synthetic zippeite obtained from ICP-OES analysis ...... 51

Table ‎2.4 Aqueous complexation reactions ...... 53

Table 3.1 Calorimetric cycles for calculation of enthalpy of formation for uranyl sulfates ...... 67

Table 3.2 ICP-OES results for chemical analysis of synthetic Natrozippeite, Cobaltzippeite, and Zinczippeite ...... 75

Table ‎4.1 Sample Table ...... 82

Table ‎4.2 Comparison between theoretical mass and average observed mass of Li36Na12-U24Pp12 , Na43K8- U24Pp12, K31Li4-U26Pp6, K26-U22PO3, K32- U28PO3, and Li28-U28...... 96

Table ‎4.3 Thermochemical cycles for the calculation of the enthalpy of formation of Li36Na12-U24Pp12, Na43K8- U24Pp12, K31Li4-U26Pp6, Li28-U28, K26-U22PO3, and K32-U28PO3 from their elements and binary oxides at 25 °C ...... 98

viii

ACKNOWLEDGMENTS

First and most, I am greatly indebted to my advisor, Dr. Peter C. Burns, who always showed me great support, patience, and encouragement throughout my

Ph.D. study at the University of Notre Dame. He believed in me, supported me, and always saw me as my strengths and not my weaknesses. Without his guidance, advice, and enduring help, I could not have completed this dissertation. I am forever grateful.

I would like to thank the readers and the dissertation committee, Dr.

Jeremy Fein, Dr. Kyle Doudrick, and Dr. Amy Hixon for their time and research advice. I would also like to express my sincere thanks to Dr. Alexandra Navrotsky and Dr. Kristina Lilova for providing me with the opportunity to get trained on the

AlexSYS at the University of California, Davis, and for their valuable suggestions in designing and conducting calorimetric experiments performed in this work.

Further thanks go to Dr. Ginger Sigmon and Jennifer Szymanowski for training me on different analytical instruments, sharing their invaluable knowledge and expertise, and for their friendship.

I would especially like to thank my father and my brother for their love and support in all these years. Special thanks go to Houtan Jebelli who always makes me feel so special. I am also grateful to many of the past and present members

of‎Dr.‎Burns’s‎lab‎for‎being‎good‎friends‎and‎helping‎me‎in‎different‎ways‎from‎ the first day I joined the group.

The research presented here is supported by the Office of Basic Energy

Sciences of the U.S. Department of Energy as part of the Materials Science of

Actinides Energy Frontiers Research Center (DESC0001089). Powder X-Ray

Diffraction, Raman spectroscopy and calorimetric analyses were done at the

Materials Characterization Facility (MCF) as part of the Center for Sustainable

Energy at Notre Dame. ICP-OES analysis was conducted at the Center for

Environmental Science and Technology (CEST) at Notre Dame. I would also like to acknowledge the CEST-Bayer Predoctoral fellowship for providing financial aid during my graduate studies.

CHAPTER 1:

INTRODUCTION

1.1 Uranium Discovery and its applications

Uranium is a member of the actinide elements series and has atomic number‎92.‎Uranium‎is‎a‎fairly‎abundant‎element‎in‎Earth’s‎crust‎(2.7‎ppm‎in‎the‎ upper crust and 1 to 1.7 ppm is the global crustal average) and occurs in economic concentrations in a variety of geological environments (Cuney, 2010).

Naturally‎occurring‎uranium‎found‎in‎Earth’s‎crust‎contains‎two‎main‎isotopes‎of‎ uranium: 99.3% 238U and 0.7% 235U. Uranium was discovered in pitchblende ore from Johanngeorgenstadt in the Erzgebirge by the Berlin pharmacist Klaproth in

1789, and only gained wider interest 150 years later, when Hahn and

Strassmann, again in Berlin, discovered nuclear fission of 235U in 1938. For several years uranium was used as a coloring agent for glasses and ceramic glazes. With the discovery of radium by Marie Curie in 1898 and its implications for medicine, uranium was mined for the recovery of radium and uranium was treated as a waste. Uranium production at this time came from the Colorado

Plateau, the ex Belgian Congo and Bohemia.

In 1938, the world entered the atomic age with the discovery of nuclear fission. Nuclear fission is a process in which the nucleus of a heavy atom splits into lighter nuclei, resulting in the release of a large amount of energy and the

production of neutrons and gamma rays. If the fissionable atom is capable of sustaining a nuclear fission chain reaction with neutrons of any energy (i.e., fast neutrons and slow neutrons), then it is called fissile. 235U is the only naturally occurring isotope with fissile properties:

235 1 U+ n0  fission fragments + 2-3 neutrons (1--‐2MeV) + energy

The energy released by a fission reaction is large enough to be used in different ways from producing electricity to making nuclear weapons depending on the degree of control on the fission reaction rate and the percentage of 235U in the fuel (Ewing, 1999). Based on the degree of enrichment of uranium in 235U, uranium is classified into five different grades: depleted uranium (less than 0.71%

235U), natural uranium (~0.71% 235U), low-enriched uranium (0.71-20% 235U), highly-enriched uranium (20-90% 235U), and weapons-grade uranium (>90%

235U) (Bruno and Ewing, 2006). For most commercial nuclear power reactors, the uranium is enriched to 3-5% 235U, whereas weapons use higher enrichment

(>90% 235U) (Bruno and Ewing, 2006). More than 30 isotopes of uranium have been discovered, but only three of them are of significant interest because of their long half-lives, 234U, 235U, and 238U (Ewing, 1999).

1.2 Nuclear Fuel Cycle

A series of industrial processes with the goal of production of nuclear energy makes up the nuclear fuel cycle. The front-end steps in the cycle are for production of nuclear fuel for nuclear reactors and back-end steps are activities to safely manage and dispose of spent nuclear fuel. The commercial fuel cycle used in the United States is an open cycle (figure 1.1). In an open cycle, there is

no treatment of the spent nuclear fuel (SNF) to recover the remaining 235U or the newly created 239Pu. In this cycle, the SNF is treated as waste and disposed of in a geological repository without any further reprocessing or recycling activities

(Ewing, 2005; Ewing, 2006).

Figure 1. 1. The open nuclear fuel cycle (Ewing, 2006).

The‎fuel‎cycle‎begins‎with‎mining‎natural‎uranium‎ore‎from‎Earth’s‎crust.‎

The ore is then milled and chemically treated. If it is high in silica, sulfuric acid is

2+ used for leaching, which leads to the production of aqueous (UO2) ; if it is

4- carbonate-rich, aqueous (UO2)(CO3)3 is produced by ammonium carbonate leaching. These steps are followed by solvent extraction or ion exchange processes to recover the uranium and produce a fine powder called yellowcake.

Then, a purification process using nitric acid and extractors is used to convert the yellow cake to uranium trioxide (UO3). This compound is subsequently converted to UF6 using a fluidized bed reactor. At this point, uranium is compatible with enrichment processes, either by gas centrifuge or

235 gaseous diffusion. After being enriched in U, pellets of UO2 cladded by zirconium alloy are fabricated and ready for use in a reactor. Upon irradiation in a reactor, the 235U concentration decreases while radioactive fission products

(137Cs, 99Tc, 90Sr, etc.) and transuranium elements (Pu, Np, Am and Cm) increase. The production of these elements and isotopes increases the radioactivity of the fuel and makes the used fuel matrix more complicated (Bruno and Ewing, 2006). After being irradiated in the reactor, spent nuclear fuel is transferred to spent fuel pools for cooling. The water cools the fuel that is still hot due to the heat generated by decay of short-lived isotopes and provides shielding from radiation. After being cooled in spent fuel pools for several years, the SNF

(Spent Nuclear Fuel) is ready to be placed in dry casks for long-term storage or to undergo reprocessing in a chemical facility. As the United States employs an open cycle, there is no reprocessing and the SNF is stored at the site of the reactor or eventually hypothetically in a geological repository. At this point in time there is no geological repository or reprocessing plant for spent nuclear fuel in the United States, and the SNF is stored at spent fuel pools onsite at reactors.

When the capacity of the pool is reached, the facility may use above-ground dry storage casks to store the SNF at independent spent fuel storage facilities

(ISFSIs) either at the reactor site or away from the reactor (Ewing, 2015).

Searching for a suitable site for a deep geologic repository for nuclear waste in the US started in 1982 after the Congress passed the Nuclear Waste

Policy Act. The US Department of Energy (DOE) was appointed to locate, select, and evaluate repository sites. In 1987, the Congress stopped the repository site selection,‎and‎Nevada’s‎Yucca‎Mountain‎was‎designated‎as‎the‎only‎site‎to‎be‎ considered for further evaluation. After several years of study on different properties of the proposed site, the project was approved in 2002 by the

Congress, and the DOE started working on its license application to the Nuclear

Regulatory Commission (NRC). However, the federal funding for the project was cut to $390 million in 2008, and the Secretary of Energy in 2009 announced plans to terminate the project. The federal funding for the Yucca Mountain repository project ended in 2011.

Up to 2012, there were about 65000 metric tons of spent nuclear fuel stored at the site of commercial reactors in the USA, and it was estimated that this number increases by 2400 metric tons per year (DOMENICI et al., 2010).

Hence thorough studies of the spent nuclear fuel matrix and its characteristics are needed to develop a new and improved method of separations/reprocessing or for choosing a site for a geological repository. Not only does the USA have high-level waste from fuel irradiation in commercial reactors, but it also has waste from uranium mining, milling, enrichment facilities, and fuel fabrication steps that are low-level waste. These wastes are not enriched in 235U but are sometimes dumped at their own facility sites. Although they are low-level waste and have almost the same activity as natural uranium, they are still environmentally

important and they should be comprehensively studied as well. Finally, the USA accumulated large quantities of weapons-related nuclear waste during the Cold

War (Ewing, 2005), some of which has been released to the environment. Much of this waste remains in temporary storage.

1.3 U6+ Crystal Chemistry and its mineralogy

In the environment, uranium occurs mainly as the tetravalent and hexavalent oxidation states, depending on the redox conditions. Uranium in the tetravalent oxidation state exists in reducing conditions as uraninite, UO2+X

(Maher et al., 2012), as well as several less common minerals. Uraninite, a major ore of uranium, is black and quite insoluble, and is isostructural with fluorite CaF2

(Finch and Ewing, 1992). Under oxidizing conditions, UO2 is first converted to

U3O8 and then to UO3 to reach its most stable oxidation state, U(VI) (Finch and

Ewing, 1992). While tetravalent uranium is almost insoluble in water, hexavalent uranium is soluble under ambient conditions and can migrate more easily in the environment (Maher et al., 2012). Formation of secondary minerals from

2+ dissolved UO2 in the oxidation zone is highly affected by the pH of the aqueous solution as well as the availability of cations and anions in the environment

(Figure 1.2). Out of nearly 200 minerals that contain uranium as a fundamental structural constituent, only six of them- uraninite (UO2+X), coffinite (USiO4), brannerite ((U,Ca,Fe)(Ti,Fe)2O6), uranmicrolite ((U,Ca,Ce)2(Nb,Ta)2O6(OH,F), uran-pyrochlore (U,Ca,Ce)2(Ta,Nb)2O6(OH,F), and carnotite

(K2(UO2)2(VO4)2•3H2O)- comprise the bulk material of most of the known uranium deposits (Hore-lacy, 2016).

Figure 1. 2. Schematic representation of formation of secondary uranyl minerals under oxidizing conditions (Finch and Ewing, 1992).

Hexavalent uranium in the presence of water is almost always present as

+2 +6 the nearly linear uranyl ion (UO2) , with an average (U) - O bond length of 1.79

+2 Å (Burns et al., 1997). The linear uranyl ion (UO2) that is observed in both aqueous and solid-state phases is typically coordinated by 4, 5 or 6 ligands in the equatorial planes of square, pentagonal, and hexagonal bipyramids (Figure 1.3).

In the case of coordination by oxygen, the average bond distance of the uranyl ion and equatorial oxygens are 2.264 Å, 2.368 Å and 2.460 Å for the square, pentagonal, and hexagonal bipyramids, respectively (Burns et al., 1997). The average bond strength for U-O bonds to the equatorial oxygens for the square,

pentagonal, and hexagonal bipyramids are 0.71, 0.53, and 0.44 vu (valence units), respectively (Burns, 2005). In the uranyl ion, the U6+ cation is strongly linked to two oxygen atoms with bond strengths ranging from 1.59 to 1.67 vu

(Burns, 2005). The strong U-O bond within the uranyl ion satisfies most of the bonding requirements of the oxygen atoms, leading to limited further bonding by the two oxygen atoms. Hexavalent uranium polyhedra tend to form sheet structures in which equatorial oxygens/ligands are linked to other high valence polyhedra to form structural units (Burns, 2005). These sheets can be connected through some low valence cations and/or water and hydrogen bonds, resulting in extended structures (Krivovichev et al., 2003; Li and Burns, 2001; Tabachenko et al., 1983).

Based upon the linkage of polyhedra of higher valence cations, U6+ crystal structures can be arranged into a structural hierarchy. This hierarchy, developed in 1996 initially, updated in 2005, and further expanded in 2016 consists of 727 known, well-defined inorganic structures that contain stoichiometric quantities of

U6+. As shown in Figure 1.4. hexavalent uranium structures in this hierarchy are classified into 5 different groups: a) isolated polyhedra (24 compounds/ 0 minerals), b) finite clusters (70 compounds/10 minerals), c) infinite chains (94 compounds/15 minerals), c) infinite sheets (353 compounds/79 minerals), and d) frameworks (186 compounds/13 minerals) (Burns, 2005; Burns et al., 1996;

Lussier et al., 2016). Dominance of sheet structures (~50%) is the result of edge or vertex sharing of uranyl polyhedra and the almost unreactive nature of oxygen atoms in the uranyl ion due to their bond valence satisfaction by U.

Figure 1.4. The hierarchy of inorganic uranyl compounds given with frequencies of each structure type (Lussier et al., 2016).

1.3.1 The zippeite family

Based‎on‎Hawthorne’s‎theory‎a‎mineral‎structure‎may‎be‎divided‎into‎two‎ parts, a structural unit and an interstitial complex (Hawthorne, 1985). The structural unit is an array of strongly bonded polyhedra that is usually anionic, and is characterized by its Lewis basicity. The interstitial complex (usually

cationic) is an array of larger low-valence cations and (H2O) groups, and is characterized by its Lewis acidity. Interstitial complexes link the structural units with relatively weak cation-anion and hydrogen bonds into a continuous structure, and the breakdown of a structure is usually controlled by the strengths of the weak interactions between the structural unit and the interstitial complex

(Schindler and Hawthorne, 2001; Schindler and Hawthorne, 2008). The interaction between these two units can be examined using the valence-matching principle (Brown, 1981; Brown, 2002); stable structure will form when the Lewis acidity and basicity of the constituent parts match. The uranyl mineral family includes the zippeite group, which are uranyl sulfate minerals with a specific structural sheet consisting of uranyl pentagonal bipyramids and sulfate tetrahedra, and a variety of interstitial complexes.

Zippeite-group minerals form from water at low pH (<4) during and subsequent to the alteration of uraninite, and sulfide minerals under oxidizing conditions. The zippeite group consists of twelve uranyl sulfate minerals: zippeite, natrozippeite, magnesozippeite, cobaltzippeite, zinczippeite, nickelzippeite, rabejacite, bluelizardite, marecottite, meisserite, sejkoraite, and pseudojohannite (Brugger et al., 2003; Brugger et al., 2006; Frondel et al., 1976;

Plášil et al.,‎2014a;‎Plášil et al.,‎2011a;‎Plášil et al.,‎2013;‎Plášil et al., 2014c).

The zippeite structure contains the linear uranyl ion in which U6+ is linked to two

O atoms by covalent bonds and each uranyl ion is coordinated by five equatorial

O atoms at the vertices of a uranyl pentagonal bipyramid. Each bipyramid shares two of its edges with two other bipyramids, forming a zig-zag chain that is two

bipyramids wide (Figure 1.5). These chains are connected to each other through sulfate tetrahedra to form the sheet structure. Burns et al. (2003) designated the sheet anion topology based on six minerals of the zippeite group (potassium, sodium, magnesium, cobalt, zinc, and nickel) that is represented by a sheet of triangles, squares and pentagons. The uranyl sulfate sheets are linked by interlayer mono and divalent cations as well as water groups (Brugger et al.,

2003).

Figure 1.5. The zippeite type uranyl-sulfate sheet

The solubility of uranyl minerals plays an important role in the distribution and transport of uranium under oxidizing conditions. Uranyl sulfates significantly impact the mobility of uranium in acid sulfate-rich mine drainage waters. In remediation activities of abandoned uranium mining sites where sulfuric acid leaching was used to exploit uranium, the interaction of uranium with sulfate and

uranium sulfate minerals with the environment is important (Meinrath et al., 2006;

Plášil et al.,‎2014a;‎Plášil et al.,‎2011a;‎Plášil et al.,‎2013;‎Plášil et al., 2014c;

Plášil et al., 2011c). Moreover, uranyl sulfates may form in stainless steel canisters that are expected to house spent nuclear fuel at a nuclear repository as several studies have shown that manganese sulfide inclusions in stainless steels are susceptible to corrosion (Ryan et al., 2002; Williams and Zhu, 2000). Under oxidizing conditions, such as in the formerly proposed repository within Yucca

Mountain, sulfide inclusions that are initiation sites for pitting may convert to sulfate. Hence, zippeite minerals may form within a repository due to the presence of sulfate groups and these may affect the overall concentration and transportation of uranium (Merkel et al., 2002).

1.4 Polyoxometalates

1.4.1 Transition Metal Polyoxometalates

Polyoxometalates (POMs), especially transition-metal polyoxometalates, have been of interest for decades due to their unusual chemical and physical properties and potential applications in different fields, such as catalysis, medicine, materials science, nanotechnology, molecular magnetism, and photochemistry (Hill, 2007; Long et al., 2007; Long et al., 2010; Mal et al., 2008;

Micoine et al., 2009; Mitchell et al., 2010; Pradeep et al., 2010; Zhao et al.,

2016). A transition-metal polyoxometalate is a discrete polyatomic ion, usually anionic, in which transition metals, Group V and Group VI - V5+, Nb5+, Ta5+, Mo6+, and W6+- are linked together via oxygen atoms making a closed structure. These

self-assembling water-soluble structures form due to the prevalence of a double- bonded oxygen ligand, known as an –yl oxygen, in their structure. The –yl oxygen ligand passivates the surface of the metal-cluster due to its resistance to hydrolysis and condensation reactions. Thus, if the pH is appropriate, the POM may remain intact in solution without ancillary protecting ligands (Nyman and

Burns, 2012).

Countercations (e.g., alkali, ammonium and phosphorous cations) are known to play an important role in synthesis, assembly, stability, solubility, and even phase transitions into non-aqueous solvents of POMs structures (Gao et al., 2017; Nyman and Burns, 2012; Yin et al., 2011). Polyoxometalates in solution may transform into single-layered hollow, spherical blackberry superstructures due to the large difference between the size of charge-balancing cations and the

POM macroions and their hydrophilic natures (Liu, 2002; Liu, 2010; Soltis et al.,

2015; Yin et al., 2011). Blackberries are self-assembled structures that adjust their size upon a change in the solution pH, charge density on the macroions, and change of solvent content (Yin et al., 2011).

1.4.2 Uranium Polyoxometalates

Uranium polyoxometalates have also yl oxygens in their structure, but two in a trans configuration unlike to the d0 transition metals. They possess two yl ligands in their structure forming U(VI)O2, U(V)O2, and the formal bond order is three; i.e., O‎≡‎U‎≡‎O‎(Burns et al., 2005; Nyman and Burns, 2012). In contrast to the d0 transition metal polyoxometalates that have been known for decades, the assembly of uranyl ions into POMs has only recently (2005) been discovered

(Burns et al., 2005). Burns et al. discovered the family of uranyl peroxide nanoclusters that spontaneously form upon adding peroxide to uranyl nitrate and adjusting the pH of the solution using base (Burns et al., 2005). They have different shapes, with the majority being closed-cage clusters, and the rest being hollow structures with ring or bowl shapes. They are composed of uranyl polyhedra bridged through bidentate peroxide ligands making structures with different chemical and physical properties. In addition to peroxide ligands, uranyl ions have sometimes other types of bridges such as hydroxyl groups, pyrophosphate, phosphite, and oxalate present in their structure (Ling et al.,

2010a; Ling et al., 2010b; Qiu and Burns, 2012; Qiu et al., 2012). Uranyl nanoclusters having peroxide bridges with or without hydroxyl groups form spontaneously at room temperature in alkaline solutions (pH 7 to 13), while incorporation of other ligands decreases their assembly pH to as low as 4. These negatively charged uranyl species are charge balanced by alkali cations that are both encapsulated and located exterior to the cluster.

2+ The pair of yl oxygen pointing in two directions in the structure of (UO2) gives stability to the inner surface of the uranyl shell and the outer surface of the cluster by its unreactive nature. In addition, the inherently bent and pliable dihedral angle of uranyl-peroxide-uranyl bridges (U-O2-U), which is due to the combinations of electronic orbitals, counters the two-dimensional polymerization, and provides curvature for the formation of clusters (Qiu et al., 2015; Sigmon et al., 2009a). DFT calculations have shown that the dihedral angle of U-O2-U is about 140°, and this is consistent with the bending in U-O2-U moiety of reported

cage clusters in general (Miró et al., 2010; Sigmon et al., 2009a; Vlaisavljevich et al., 2010). The presence of hydroxyl bridges in nanocluster structures generally lead to the formation of bigger clusters due to the fact that U–(OH)2–U bridges are flatter, as shown by DFT simulations (Nyman and Burns, 2012). The pliability of the dihedral angel is also strongly dependent on the presence of nearby alkali cations (Gil et al.,‎2012;‎Miró et al., 2012; Nyman and Alam, 2012; Vlaisavljevich et al., 2010). Thus, counter cations have been considered to be one of the factors contributing to size, curvature and assembly of nanoclusters from their building blocks that are uranyl monomers. The combination of the small covalent components of the dihedral angel of the uranyl-peroxide-uranyl bridges as well as the intrinsic unreactive nature of the yl oxygens in the uranyl ion lead to the self-assembly of hollow spherical-shapes of uranyl peroxide nanoclusters in solution.

Since the initial discovery of uranyl peroxide nanoclusters, this family has expanded to more than 60 published structures of high-symmetry topologies with different chemical composition, charge, size, and properties (Qiu et al., 2014;

Sigmon et al., 2009c; Sigmon et al., 2009f; Unruh et al., 2011). The smallest is

U16 (16 uranyl polyhedra) and the largest is U124 (124 uranyl polyhedra) with average sizes of 1.5 and 5 nm, respectively. Of the over 60 published uranyl nanoclusters, the water-soluble U60 (Figure 1.6), with a topology identical to C60 buckminsterfullerene and with the formula Li48+mK12(OH)m[UO2(O2)(OH)]60(H2O)n

(where‎m≈20‎and‎n≈310)‎is‎the‎most‎studied‎with‎a‎diameter‎of‎2.7‎nm‎(Sigmon et al., 2009c).

Figure 1.6. Linkages of hexagonal bipyramids by sharing of a peroxide edge (top), U60 cluster (bottom) (Sigmon et al., 2009c).

Uranyl peroxide nanoclusters exist in the crystalline form, and as dissolved species in an aqueous phase. The crystalline phase is composed of an assemblage of nanoclusters as well as counter cations and water, and in the aqueous phase nanoclusters behave as isolated structures dissolved in solution.

The crystals of uranyl nanoclusters are generally soluble in water. They dissolve readily in undersaturated solutions liberating isolated nanoclusters to the solution

(Qiu et al., 2012). A qualitative study has shown that isolated uranyl nanoclusters are stable in solution for months (Armstrong et al., 2012). In a study done on the

60− aqueous solubility of U60 ([(UO2)(O2)(OH)]60 ) and U24Pp12

48− ([(UO2)24(O2)24(P2O7)12] ), it has been shown that uranyl nanocluster solubility depends heavily on the identity of their charge-balancing alkali cations and the thickness of their electrical double layer (Peruski et al., 2017). The amount of uranium released upon dissolution of U60 and U24Pp12 crystals in water was in the

range of 38,000-300,000 ppm, which is about 3 or more orders of magnitude higher than would be expected for simple inorganic species under similar conditions (Peruski et al., 2017).

Due to their physical and chemical properties, uranyl peroxide nanoclusters have potential applications in the nuclear industry and nuclear fuel cycles (Burns, 2011; Wylie et al., 2013). For example, they can be used to extract and recover uranium from solution using ultrafiltration or SBA15 due to the size and mass difference between uranyl nano-clusters and other smaller solutes in solution (Liu et al., 2015; Wylie et al., 2016; Wylie et al., 2013). The uranyl peroxide nanocluster family may form at contaminated sites such as

Fukushima and Hanford (Armstrong et al., 2012; Burns et al., 2012).

1.5 Thermochemistry of uranyl compounds

In order to predict the solubility and stability of a phase of interest, the standard state Gibbs free energy must be determined. The Gibbs free energy is the energy associated with a reaction that can be used to do work and is a function of temperature, standard state enthalpy and entropy of formation. Using the Gibbs free energy, one can calculate if a reaction is spontaneous and if the products are stable relative to the reactants. The free energy G is a quantity that becomes more negative during the course of natural process. It is described as a function of enthalpy (ΔH), entropy (ΔS), and temperature, (T), in the following way:

ΔGreaction = ΔHreaction –TΔSreaction (1)

There‎are‎three‎ways‎to‎calculate‎∆G:‎

1- Having the standard state Gibbs free energy of reactants and products: ΔGreaction= ΔGproducts –ΔGreactants (2)

2- Based on solubility analysis of the phase of interest: ΔGreaction° = -RTlnKsp° (3)

Where Ksp is the equilibrium constant of a reaction, T is the temperature, and R is 0.08206 L atm mol-1 K-1. The equilibrium constant of a chemical reaction is the ratio of the activity of products over the activity of reactants at equilibrium. The Ksp of a reaction is independent of the concentrations of reactants and products, but it does depend on the temperature and pressure of the system.

3- Based on measuring or calculating the standard state enthalpy and entropy of formation, equation 1. Thermodynamic studies of uranium and higher actinide materials are crucial in understanding their behavior in the nuclear fuel cycle. Studies of the thermodynamic properties of uranium minerals and related compounds are motivated by a need to understand their behavior in the genesis of uranium deposits, the exploration and extraction of U ore, the transport of radionuclides in the environment and nuclear waste repositories, and the operation of nuclear reactors (Burns and Sigmon, 2013; Peeters et al., 2008). Several studies have shown that using a high-temperature twin Calvet oxide melt drop-solution calorimeter, the AlexSYS 1000 made by Setaram, one can obtain reliable data for the enthalpy of formation of uranyl compounds. To date, accurate thermodynamic data for minerals including uranyl oxide hydrates, peroxides, uranyl carbonates, uranyl phosphates, uranyl silicates, and uranyl vanadates have been obtained using this calorimetric method as have solubilities for most of them (Dzik et al., 2017; Gorman-Lewis et al., 2007; Gorman-Lewis et al.,

2009; Guo et al., 2015; Guo et al., 2014; Kubatko et al., 2006; Kubatko et al.,

2005; Kubatko et al., 2003; Shvareva et al., 2011; Spano et al., 2017).

Determinations of enthalpies of formation using the AlexSYS have also been done for non-mineral uranyl peroxide cluster compounds. Unlike uranyl minerals, calorimetric studies of uranyl clusters have been reported for only two examples, U60 containing 60 uranyl polyhedra and U28 containing 28 uranyl polyhedra (Armstrong et al., 2012; Tiferet et al., 2014). Despite the considerable amount of work on calorimetric and solubility studies of uranyl compounds, their thermodynamic properties remain largely unknown and more research is needed in this area.

The AlexSYS calorimeter provides the drop solution enthalpy or the sum of the heat content of the sample from room temperature, its heat of dissolution in the solvent and its heat of reaction (if a reaction occurs). It has thermocouples that measure the temperature difference between the sample and its constant surroundings (Figure 1.7), and it has been shown that this method can give reliable data for the enthalpy of dissolution (Navrotsky, 1977; Navrotsky, 1997).

Based on the drop solution enthalpy, the enthalpy of formation of the sample from elements, oxides, or other compounds is calculated using proper thermochemical cycles.

Figure 1.7. Schematic of the AlexSYS and assembly of glassware for drop solution calorimetry in a molten oxide solvent (Navrotsky, 2014)

The AlexSYS is‎calibrated‎using‎high‎purity‎α-Al2O3 at 973K. Sample pellets of alumina are dropped into the instrument and the well-known heat capacity of alumina is used to calculate the calibration factor. Prior to each experiment, the complete dissolution of the phase in 3Na2O3.4MoO3 is confirmed by experiments in a furnace at 973 K followed by visual inspection. O2 or an inert gas (e.g. Ar, N2) is used as the flushing gas over the solvent at a constant rate to evolve any gas or water vapors associated with dissolution of the samples. The same gas is also bubbled at a constant rate through the

solvent to provide stirring as well as an oxidizing environment and to prevent local saturation in the solvent.

The AlexSYS is a very stable yet sensitive instrument. When it is placed at relatively constant temperature and humidity, it gives reliable, accurate, and reproducible data. However, every imperfection in the setup, experiment, or any disturbance in the environment can cause errors in the results. The AlexSYS measures the temperature difference between the sample and its constant surroundings, using thermocouples, wired in series – i.e. thermopile. The baseline of the calorimeter represents the difference in the signal between the two thermopiles. In the ideal case, it should be zero. In practice, the most important is for the baseline to be constant (Figure 1.8). Having a stable baseline before starting an experiment is necessary in order to ensure that the final oxidation state of the sample dropped has been reached and dissolution and reactions are done. To check the stability of the baseline, the baseline is monitored for about 10 minutes and the fluctuation of the baseline must be less than‎0.04‎μV.‎In‎general,‎the‎fluctuation‎of‎the‎baseline‎is‎expected‎to‎be‎ between 0.02-0.1‎μV‎within‎an‎hour‎depending‎on the environmental conditions to be designated stable. Although a stable baseline is an important factor in the functionality of the calorimeter, it is not enough for acquiring accurate and reproducible data. Having pure, single-phase, and homogenized samples that are thoroughly characterized is also necessary in order to obtain reliable data.

High-temperature drop-solution calorimetry has several advantages over other techniques used for determination of enthalpy of formation. This technique

requires only a small quantity of sample (around 50 mg is enough for calorimetric analysis), has high sensitivity and hence more accurate results, introduces the non-mechanical stirring technique, i.e. bubbling, and ensures complete dissolution of the sample at high temperature using different types of solvents.

Figure 1.8. Change in the baseline of the calorimeter over a reaction.

1.6 Thesis Overview

Important gaps remain in understanding the behavior of uranium minerals and compounds under different environmental conditions despite their importance in environmental, material, and nuclear science fields. This dissertation is focused on synthesis, characterization, solubility experiments, and calorimetric studies of selected uranyl minerals and synthetic compounds.

Chapter 1 provides some background on the importance of actinide materials and nuclear energy, hexavalent uranium mineralogy, and uranyl

polyoxometalates. Specifically, the environmental importance of the zippeite family (uranyl sulfate minerals) and uranyl peroxide nanoclusters (recently- discovered uranium polyoxometalates) as well as their novel properties that make them important in the nuclear industry are discussed in Chapter 1.

Chapter 2 presents the first complete set of thermodynamic properties (i.e. standard state enthalpy, entropy, and Gibbs free energy of formation) of zippeite derived from solubility and calorimetric studies. Chapter 3 describes thermodynamic studies of synthetic analogs of four minerals from the zippeite family to investigate the relationship between their chemical properties and enthalpy of formation acquired from calorimetric studies. Chapter 4 is focused on synthesis, characterization, and measurement of enthalpy of formation for six members of the family of uranyl peroxide nanoclusters to better understanding of their solid-state behavior. Chapter 5 concludes research presented in this dissertation and comments on possible future work.

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Zhao, J.-W., Li, Y.-Z., Chen, L.-J., Yang, G.-Y., 2016. Research progress on polyoxometalate-based transition-metal–rare-earth heterometallic derived materials: synthetic strategies, structural overview and functional applications. Chemical Communications, 52(24): 4418-4445.

CHAPTER 2:

THERMODYNAMIC STUDIES OF ZIPPEITE, A URANYL SULFATE COMMON

IN MINE WASTES

2.1 Abstract

Zippeite is a potassium uranyl sulfate mineral that forms on uranium mine wastes and that may be important in nuclear waste disposal. Little is known about its thermodynamic properties. In this study, we synthesized zippeite,

K3(H2O)3.3[(UO2)4(SO4)2O3(OH)], and measured its thermodynamic properties.

The standard enthalpy of formation of‎zippeite‎from‎the‎elements,‎−‎

8655.97±12.55 kJ/mol, was calculated from the measured enthalpy of dissolution in a high-temperature oxide melt solution calorimeter. Solubility measurements were conducted from both undersaturation and supersaturation in order to constrain the equilibrium state. Using the solubility product (Ksp) obtained from the solubility data, the standard Gibbs free energy of formation of zippeite at 298

K‎was‎calculated‎to‎be‎−‎7783.44±6.87‎kJ/mol.‎This‎value,‎coupled‎with‎the‎ standard enthalpy of formation value determined from the calorimetric data, was then‎used‎to‎calculate‎the‎standard‎entropy‎of‎formation,‎−‎2926.49±45.64‎J/mol.‎

Our studies yield the first complete set of thermodynamic properties of zippeite, and the results allow us to predict the conditions under which its formation is favorable relative to other uranyl minerals under a wide range of conditions of geologic and engineering interest.

2.2 Introduction

Zippeite, K3(H2O)3[(UO2)4(SO4)2O3(OH)], is a relatively soluble mineral that generally forms under low pH conditions during and subsequent to the alteration of uraninite, UO2+x and sulfide minerals, such as pyrite and galena, under oxidizing conditions (Brugger et al., 2003; Frondel et al.,‎1976;‎Plášil et al., 2014c). The aqueous solubility of uranium minerals is important in determining the distribution and transport of uranium under oxidizing conditions.

Uranyl sulfates significantly impact the mobility of uranium in water, especially in former uranium mines and related mine tailings (Finch and Murakami, 1999).

Zippeite is particularly important in abandoned uranium mines that discharge acid drainage due to sulfide oxidation. It is common in abandoned uranium mining sites where sulfuric acid leaching was used to recover uranium (Meinrath et al.,‎2006;‎Plášil et al.,‎2014a;‎Plášil et al.,‎2011a;‎Plášil et al.,‎2013;‎Plášil et al.,‎2014b;‎Plášil et al., 2011b). In addition to these occurrences in the environment, uranyl sulfates like zippeite may form in the stainless steel canisters intended to house spent nuclear fuel in a nuclear repository, as several studies have shown that manganese sulfide inclusions in stainless steel are susceptible to corrosion (Ryan et al., 2002; Williams and Zhu, 2000). Under oxidizing conditions, such as in the previously proposed repository at Yucca

Mountain NV, sulfide inclusions that are initiation sites for pitting may convert to sulfate, and uranyl sulfates may affect the overall concentration and transport of uranium in a repository setting (Merkel et al., 2002). The presence of sulfate in

Hanford tank wastes arising from the bismuth phosphate process used during

nuclear weapons production adds to the need to understand the thermodynamic properties of zippeite (Tingey et al., 2004).

There currently is a lack of reliable thermodynamic data for zippeite. To date, there has been only one report concerning the Gibbs free energy of zippeite (O'Brien and Williams, 1981), with a value derived from dissolution studies‎and‎metal‎aqueous‎complexation‎calculations.‎O’Brien‎and‎Williams‎ reported the Gibbs free energy of formation of zippeite from the elements to be

−12603‎±10‎‎kJ/mol‎(O'Brien and Williams, 1981). In that study, it was not clear that equilibrium was attained during the experiment, the enthalpy and entropy of formation were not determined, the composition of the zippeite studied was uncertain (see below), and the possibility of non-stoichiometric (incongruent) dissolution was not evaluated by analyzing the residual solid phase after the experiments. Prediction of the Gibbs free energy and enthalpy of formation of various U6+ phases, including K-zippeite, has been conducted based on models involving the polyhedral geometries and coordination numbers of the cations involved (Chen et al., 1999).

Subsequent to the earlier studies, Burns et al. characterized crystals of synthetic material and defined the structure and formula of zippeite,

K3(H2O)3[(UO2)4(SO4)2O3(OH)] (Burns et al., 2003), both of which differ from the results from the earlier studies. The zippeite structure contains sheets of uranyl pentagonal bypiramids and sulfate tetrahedra (Burns et al., 2003) (Figure 2.3 in

Supporting Information). Each bipyramid shares two of its equatorial edges with two other uranyl bipyramids, forming a zig-zag chain that is two bipyramids wide.

These chains are connected to each other through four-connected sulfate tetrahedra to form sheets. The uranyl sulfate sheets are linked through bonds to potassium cations, as well as through hydrogen bonds associated with water groups. Using a bond valence approach and the electro neutrally requirement,

Burns et al. showed that previous formulas and structures were based on natural zippeite with impurities or were not accurate (Burns et al., 2003).

In this work, we have synthesized zippeite and confirmed the purity of the phase using powder X-ray diffraction (PXRD), inductively coupled plasma optical emission spectroscopy (ICP-OES), and thermogravimetric analysis (TGA). We conducted solubility studies both from undersaturation and oversaturation, as well as high-temperature oxide melt solution calorimetry in order to determine the thermodynamic properties of zippeite, and we use these results to illustrate the calculation of the thermodynamic stability of zippeite relative to other uranyl- bearing phases under a range of conditions.

2.3 Materials and Methods

Caution: Although the uranium used in the experiments described here was isotopically depleted, precautions for handling radioactive materials should be followed. In our experiments, all of the reagents, unless otherwise stated, were analytical grade and used as received without further purification. Zippeite was synthesized hydrothermally by combining 0.254 g uranyl acetate, 0.052 g potassium sulfate,‎and‎4‎mL‎ultrapure‎18‎MΩ‎‎H2O in a 23 mL Teflon-lined Parr reaction vessel. The pH of the solution was adjusted to 4.3 using 1M KOH.

Several batches of material were made and mixed together in order to

accumulate enough material for characterization, calorimetric studies and solubility measurements. The reaction vessel was placed inside a mechanical convection oven and heated to 150 ºC. After 3 days, the reaction vessel was removed from the oven and allowed to cool to ambient temperature. The yellow micro-crystalline zippeite was recovered by air-filtration and washed with ultrapure water. To verify the purity, powder X-ray diffraction (PXRD) patterns were collected using a Bruker D8 Davinci diffractometer equipped with monochromatic‎Cu‎Kα-radiation over the‎2θ‎range‎from‎5‎to‎55‎with‎a‎step‎size‎ and exposure time of 0.01º and 1 s, respectively.

The elemental composition of the synthesized zippeite was determined by complete dissolution of the synthesized material followed by elemental analysis using a Perkin Elmer Optima 8000 DV-ICP-OES instrument. Approximately 0.12 mg of the sample was dissolved in 10 mL diluted (5%) HNO3 and analyzed in triplicate for the concentrations of U, S, and K in solution. Six external standards having the elements of interest with concentration ranging from 0.1 to 20 ppm were used for calibration. Also, yttrium was added to all samples and standards at a concentration of 0.6 ppm to monitor instrument drift. The experimental ratios between the elements of interest, U, K, and S were in good agreement with the chemical formula associated with the crystalline phase of zippeite (Burns et al.,

2003).

Thermogravimetric and differential scanning calorimetric analyses were carried out using a Mettler Toledo TGA/DSC with a vertical balance calibrated by alumina pellets coupled with a mass spectrometer (Thermo Star, Pfeiffer) for

evolved gas analysis. The number of molecules of water associated with the synthetic phase was calculated based on mass change upon temperature increase. A 25 mg pellet of the sample was heated up to 950 ºC at a 5 ºC/min heating rate under flowing Ar gas at 50 mL/min.

A high-temperature twin Calvet calorimeter, model AlexSYS 1000 produced by Setaram, was used to measure the enthalpy of dissolution of zippeite in the melt (3Na2O-4MoO3) at 700 ºC. It has been shown that this method gives reliable data for the enthalpy of dissolution of uranyl compounds

(Navrotsky, 1977; Navrotsky, 1997). Based on the drop solution enthalpy, the enthalpies of formation of the sample from its elements and binary oxides were calculated using thermochemical cycles. The instrument was calibrated using the heat‎content‎of‎high‎purity‎α-Al2O3. The complete dissolution of the zippeite in

3Na2O3-4MoO3 was confirmed by experiments in a furnace at 700 ºC and visual inspection, where ∼ 5 mg of zippeite was dropped into the molten solvent and a clear solution was observed after about 25 min. O2 gas was used as the flushing gas over the solvent at a constant rate, 35 mL/min, to evolve any gas or water vapors associated with dissolution. O2 was also the bubbling gas and was bubbled at 5 mL/min through the solvent to provide stirring as well as an oxidizing environment and to prevent local saturation in the solvent.

Raman spectroscopy was conducted using a Bruker Sentinel system linked via fiber optics to a video-assisted Raman probe equipped with a 785 nm,

400 mW light source, TE-cooled, 1024*255 CCD array. The spectra were collected over a frequency range of ∼ 300-2500 cm-1 for 15 seconds with 3

integrations and a 45-second background. Raman spectra of the starting material, zippeite, were compared with the Raman spectra of the residue of the solubility experiments.

Solubility experiments were conducted approaching equilibrium from both under- and super-saturation. To approach equilibrium from undersaturation, approximately 500 mg of the synthesized zippeite sample was added to 8 mL of

0.1M NaClO4 in a reaction vessel. The pH of the solution was kept below 4 using diluted HCl to simplify the aqueous speciation of uranium and to ensure that aqueous uranyl carbonate and uranyl hydroxide complexes are not dominant species in the experimental solution (Choppin and Mathur, 1991; Tripathi, 1983).

Solubility measurements from under-saturation were conducted at three different pH values, 2.8, 3.5 and 4. A solubility experiment that approached equilibrium from supersaturation was conducted at pH 3.5 by addition of ~ 500 mg zippeite to

8 mL of 0.1 M NaClO4 solution that already contained dissolved U, K, and S from the addition of UO2(NO3)2 and K2SO4. The starting concentrations of U, K, and S were 0.044, 0.05, and 0.026 mol/L, respectively, and were identical to those measured for the undersaturation experiment at pH 3.5. All experiments were conducted in Teflon reaction vessels, which were sealed and the pH of each solution was measured and adjusted each day for four weeks using a pH meter calibrated by 3 NIST standards. At various intervals, aliquots of solution were taken from‎each‎tube‎and‎filtered‎with‎0.1‎μL‎nylon‎filters,‎to‎remove‎any‎residual‎ solid. Subsequently, the solution samples were diluted in about 10 mL of diluted

(5%) HNO3 acid for ICP-OES analyses to measure the concentration of the

elements of interest. At the end of each experiment, about 20 mg of the remaining powder was used for PXRD and ICP-OES analysis in order to test whether the solid phase was stable over the duration of the experiment and that the composition of the remaining solid did not change.

2.4 Results and Discussion

The PXRD pattern of finely ground synthetic zippeite deposited onto a zero-background quartz plate is shown in Figure 2.4 in Supporting Information.

The observed pattern is in good agreement with that calculated from the previously published zippeite structure model by Burns et al. (Burns et al., 2003) and there are no additional peaks. Chemical analysis results are in good agreement with the accepted formula (Table 2.2 in Supporting Information).

Thermogravimetric analysis (Figure 2.5 in supporting Information) yielded a somewhat higher water content than expected (3.3 moles of water instead of 3), likely due to surface water that was lost at 40-80 ºC. Loss of structural water occurred at 80-150 ºC. During heating to 650-900 ºC, SO2 was liberated from the sample, as confirmed by mass spectrometry of the evolved gas.

High-temperature oxide melt solution calorimetry using molten sodium molybdate at 700 ºC was used to determine the standard enthalpy of formation of zippeite using a well-established methodology (Navrotsky, 1997; Navrotsky,

2014). The sample readily dissolved in the melt and the calorimetric reaction was complete after about 33 minutes, as indicated by the calorimeter signal returning to the original baseline. Using thermochemical cycles (Table 2.1) the enthalpy of formation of zippeite from its binary oxides and elements was calculated‎to‎be‎−‎

1338.53‎±‎11.74‎kJ/mol‎and‎−‎8655.97‎±‎12.55‎kJ/mol,‎respectively.‎Reactions‎1‎ and 2 represent zippeite formation from its oxides and elements:

1.5K2O(xl, 298 K) + 4UO3(xl, 298 K) + 2SO3(g, 298 K) + 3.8H2O (l, 298K)

= K3(H2O)3.3[(UO2)4(SO4)2O3(OH)](xl, 298K) (1)

3K(xl, 298K) + 4U(xl, 298 K) + 2S (s, 298K) + 3.8H2(g, 298K) + 11.65O2

(g, 298K) = K3(H2O)3.3[(UO2)4(SO4)2O3(OH)](xl, 298K) (2)

TABLE 2.1

THERMODYNAMIC CYCLES FOR FORMATION OF ZIPPEITE FROM ITS ELEMENTS AND BINARY OXIDES

Reactions ∆H

∆H1: ΔHds (K-zippeite) K3(H2O)3.3[(UO2)4(SO4)2O3(OH)](xl, 298K) 754.289 ± 3.20 kJ//mol = 1.5 K2O(soln, 976 K) + 4UO3(soln, 976K) + 2SO3(soln, 976K) +

3.8H2O(gas, 976K) 40 (2)‎‎‎‎‎‎‎‎‎‎‎‎∆H2:‎ΔHds (UO3) (6)‎‎‎‎‎‎‎‎‎‎‎‎∆H6:‎ΔHf°(UO3) − 1223.8 ± 0.8 9.49 ± 1.53 kJ/mol UO = UO U + 1.5O = kJ/mol (Robie 3 (xl, 298 K) 3(soln, 976 K) (Helean et al., 2002) (xl,298K) 2 (g, 298K) UO 3 (xl, 298K) et al. , 1978) (3)‎‎‎‎‎‎‎‎‎‎‎‎‎∆H3:‎ΔHds (SO3) − 203.7 ± 4.091(7)‎‎‎‎‎‎‎‎‎‎‎‎∆H7:‎ ΔHf° (H2O) − 285.8 ± 0.1 SO3 (g, 298 K) = SO3 (soln, 976 kJ/mol (Navrotsky, H2 (g, 298K) + 0.5O2 (g, 298K) = kJ/mol (Robie K) 2014) H 2 O (l, 298K) et al. , 1978) (4) ∆H4:‎ΔHds (K2O) − 318.0 ± 3.1 (8)‎‎‎‎‎‎‎‎‎‎‎∆H8:‎ΔHf° (SO3) − 395.7 ± 0.7 K2O (xl, 298 K) = K2O (soln, 976K) kJ/mol (Molodetsky et S (S, 298K) + 1.5O2 (g, 298K) = kJ/mol (Robie al. , 2000) SO 3 (g, 298K) et al. , 1978) (5)‎‎‎‎‎‎‎‎‎‎‎‎‎∆H5:‎ΔH (H O) hc 2 (9)‎‎‎‎‎‎‎‎‎‎‎‎∆H9:‎ΔH ° (K O) − 363.2 ± 2.1 H O = H O 69.0 kJ/mol (Robie f 2 2 (l, 298 K) 2 (g 976 K) 2K + 0.5O = kJ/mol (Robie et al., 1978) (xl, 298 K) 2 (g, 298K) K O et al. , 1978) 2 (xl, 298K)

TABLE 2.1 (CONTINUED)

Reactions ∆H

(12) K3(H2O)3.3[(UO2)4(SO4)2O3(OH)] o ΔHf-ox (K-zippeite) = -ΔH1 + 1.5ΔH4 + 4ΔH2 + 2ΔH3 + 3.8ΔH5 1.5K2O(xl, 298 K) + 4UO3(xl, 298 K) + 2SO3(g, 298 K) + − 1338.53 ± 11.61 kJ/mol 3.8H2O (l, 298K) = K3(H2O)3.3[(UO2)4(SO4)2O3(OH)](xl, 298K)

(13) K3(H2O)3.3[(UO2)4(SO4)2O3(OH)]

ΔHf° (K-zippeite) = -ΔH1 + 1.5ΔH4 + 4ΔH2+ 2ΔH3 + 3.8ΔH5 + 1.5ΔH9 + 4ΔH6+ 2ΔH8 +3.8ΔH7 − 8655.97 ± 12.54 kJ/mol 3K(xl, 298K) + 4U(xl, 298 K) + 2S (s, 298K) + 3.8H2(g, 298K) +

11.65O2 (g, 298K) = K3(H2O)3.3[(UO2)4(SO4)2O3(OH)](xl,

298K ) xl, solid material; g, gaseous; soln ,solution; and l, liquid. Error is represented as two standard deviations of the mean.

Solubility experiments were conducted both from undersaturation and supersaturation to constrain the equilibrium state. Measurements of dissolved U,

S, and K in the experimental samples were used to calculate the equilibrium constant and Gibbs free energy of the dissolution reaction (3) as well as the standard-state Gibbs free energy of formation of zippeite. The dissolution reaction of zippeite can be expressed as:

+ + 2+ 2- 7H + K3(H2O)3.3[(UO2)4(SO4)2O3(OH)] = 3K + 4UO2 + 2SO4 + 7.3H2O [3]

The standard state for zippeite and water are assumed to be the pure phases at the pressure and temperature of interest; the standard state for aqueous species is assumed to be a hypothetical one molal solution at the pressure and temperature of interest that behaves as if infinitely dilute (Gorman-Lewis et al.,

2008). Using these standard states, the solubility product of zippeite can be expressed as:

3 4 2 푎퐾+ ∗푎 2+ ∗푎 2− 푈푂2 푆푂4 Ksp = 7 [4] 푎퐻+ where ' 푎 ' represents the aqueous activity of the subscripted species. All experiments reached steady-state within two weeks with the highest concentration of dissolved U, S, and K observed in the pH 2.8 experiment, Figure

1. The average concentration of dissolved U, S, and K for data points after 15 days of the start of each experiment were used in the calculations of the solubility product. All experiments yielded almost identical extents of non-stoichiometric dissolution, with K slightly in excess of U relative to stoichiometric dissolution, and S slightly lower in concentration relative to U. The pH 3.5 experiments from under- and super-saturation yielded steady-state concentrations of the dissolved

elements that were not significantly different from each other, suggesting that equilibrium was reached within the experimental time period. Powder X-ray diffraction patterns of the solid phase residues from the experiment did not reveal any additional crystalline phases and Raman spectra of the residues matched those of the starting material (Figure 2.6 in Supporting Information). These observations indicate that a secondary precipitating phase, if present, was minor

(less than 5%) or had the same vibrational modes as the crystalline starting material.

The elemental analyses of the experimental samples yielded total elemental concentrations of each solution, but in order to calculate the value of the solubility product, the activities of individual U, S, and K species in solution must be determined. We used the aqueous complexation reactions listed in

Table 2.4 in supporting information in conjunction with FITEQL 2.0 (Herbelin and

Westall, 1999) to calculate the activity of each dissolved species in each

experimental sample. The Davies equation was used within FITEQL 2.0 to calculate activity coefficients for the ionic species. The calculated value of the Ksp did not vary consistently with pH. The calculated logarithm of the solubility product averaged for the four different‎solubility‎experiments‎with‎its‎1δ‎error‎is‎

4.14±0.11. This average solubility product value was used to determine the standard state Gibbs free energy of the dissolution reaction using the following equation:

ΔG°reaction =‎−‎RTlnKsp [5] where‎T‎is‎the‎absolute‎temperature‎and‎R‎is‎the‎gas‎constant.‎The‎ΔG°reaction value‎with‎its‎1δ‎error‎was‎calculated‎to‎be‎-23.68±0.65 kJ/mol. Using literature values for standard state Gibbs free energies of formation‎(ΔG°f) for each component in the dissolution reaction(Cox et al., 1989; Robie et al., 1978), and the‎calculated‎value‎of‎ΔG°reaction, we calculated the standard-state Gibbs free energy‎of‎formation‎of‎zippeite‎to‎be,‎−‎7783.44‎±‎6.87‎kJ/mol:

ΔG°f (Zippeite)‎=‎−‎ΔG°reaction − 7ΔG°f H+ +3ΔG°f K+ +‎4ΔG°f

UO22+ +‎2ΔG°f SO42- +‎3.3‎ΔG°f H2O [6]

The standard-state Gibbs free energy of formation of zippeite, in conjunction with the standard enthalpy of formation of zippeite from the calorimetry experiments, was used to obtain the standard entropy of formation of zippeite from the following relationship:

ΔG°f = ΔH°f –TΔS°f [7]

We‎calculated‎the‎value‎of‎ΔS°f for‎zippeite‎with‎propagated‎1δ‎error‎to‎be‎

−2926.49‎±‎45.64‎J/mol.

Zippeite can form as an alteration product of uraninite, and in some cases, it has been found associated with (meta)schoepite in uranium mines (De Vivo et al., 1984). Reaction 8 describes the formation of zippeite from metaschoepite:

+ 2- + 4UO3•2H2O + 3K + 2SO4 + H = K3(H2O)3.3[(UO2)4(SO4)2O3(OH)] + 4.7 H2O [8]

The calculated value of Gibbs free energy of formation for zippeite from this study‎was‎used‎in‎order‎to‎determine‎the‎ΔG°rxn for‎reaction‎8‎(−32.47‎±‎16.33‎ kJ/mol). As described by Equation 9, alteration of metaschoepite is dependent on the concentration of the cations in the system:

1 ΔGrxn =‎ΔG°rxn + RT ln ( 3 2 ) [9] 푎퐾+∗ 푎 2−∗ 푎퐻+ 푆푂4

For example, for an aqueous sulfate concentration of 11.75 × 10-3 mol/L, as reported for the post-mining groundwater around the Crow Butte uranium mine (Davis and Curtis, 2007), an activity-activity diagram can be calculated,

Figure 2. As shown in Figure 2, zippeite is stable at lower pH and in the presence of high concentrations of K+. For instance, at pH 3, metaschoepite is unstable relative to zippeite for any concentration of dissolved potassium approximately above 1 mol/L.

Figure 2.2 pH- Log a (K+) diagram for the system metaschoepite- zippeite at sulfate concentration of 11.75 × 10-3 with activity coefficients calculated using an extended Davies equation.

In another example, we can use the results from this study to calculate whether zippeite or other uranyl phases in Hanford tanks are the more stable phase within the tanks and in groundwaters at the site. Uranyl oxides and oxyhydrates such as compreignacite and clarkeite have been found in contaminated sediments (Um et al., 2009) around the Hanford site and have been reported as potential uranium-bearing phases present in Hanford tank residual waste. The thermodynamics of the formation of zippeite from UO3, Na- compreignacite (Na2(UO2)6O4(OH)6(H2O)7), and clarkeite (Na(UO2)O(OH)) are shown in Table 2.2 Zippeite is stable only with respect to UO3 for pH below 3.

TABLE 2.2

FORMATION OF ZIPPEITE FROM UO3, COMPREIGNACITE, AND CLARKEITE

ΔG ° Reactions reaction ΔG at Hanford (kJ/mol) rxn <0 4UO + 3K+ + 2SO 2- + H+ + 3.3 H O = K (H O)−‎78.30‎±‎7.17[(UO ) (SO ) O (OH)] 3 4 2 3 2 3.3 2 4 4 2 3 For pH below 3 + + 2(Na2(UO2)6O4(OH)6(H2O)7) + 7H + 9K + 6SO 2- = >0 4 −‎165.93‎±‎32.06 3(K3(H2O)3.3[(UO2)4(SO4)2O3(OH)]) + For all pH +

48 4Na + 12.1 H2O + 2- +

4(Na(UO )O(OH)) + 3K + 2SO + 5H 2 4 +33522.095 ± >0 = K3(H2O)3.3[(UO2)4(SO4)2O3(OH)] + + 49.48 For all pH 4Na + 0.7 H2O

Zippeite can form from ions in solution when there is significant

2+ 2- + evaporation to increase the concentration of UO2 , SO4 , and K (Finch and

Murakami, 1999). Using our calculated solubility product for zippeite, we calculated that Hanford tank wastes are supersaturated with respect to zippeite

+2 + 2- for any pH above 5 based on the activities of UO2 , K , and SO4 at the Hanford site. Since the pH of the tanks is above 4 (Um et al., 2009), our calculations suggest that zippeite can form directly from precipitation from Hanford tank waste solutions. The reason not been found outside of the tanks is likely that uranyl oxyhydrates and other less soluble minerals such as phosphates have formed

2+ + (Um et al., 2009) and that the concentrations of dissolved UO2 , sulfate, and K , are lower in the groundwater beneath the tanks than in the tank solutions and hence are below the saturation state for zippeite.

2.5 Acknowledgments

This work was funded by the Office of Basic Energy Sciences of the U.S.

Department of Energy as part of the Materials Science of Actinides Energy

Frontiers Research Center (DESC0001089). Powder X-Ray Diffraction, Raman spectroscopy and calorimetric analyses were done at the Materials

Characterization Facility (MCF) as part of the Center for Sustainable Energy at

Notre Dame. ICP-OES analysis was conducted at the Center for Environmental

Science and Technology (CEST) at Notre Dame.

2.6. Supporting Information

Figure 2.4. Zippeite PXRD pattern. Blue trace corresponds to synthetic zippeite, and red peaks correspond to peak positions and heights calculated from the structural model (Burns et al., 2003).

Figure 2.5. DSC-TGA curve for synthetic K-zippeite.

TABLE 2.3

CHEMICAL COMPOSITIONS OF SYNTHETIC ZIPPEITE OBTAINED FROM

ICP-OES ANALYSIS

Sample Ratio Ideal Actual mol K/mol U 0.75 0.72±0.3 K-zippite mol S/mol U 0.5 0.49±0.02 The error is stated as two standard deviations of the mean.

TABLE 2.4

AQUEOUS COMPLEXATION REACTIONS

Reactions Log K (Reference) 2+ + + UO2 + H2O = UO2OH + H -5.25 (Guillaumont et al., 2003) 2+ + UO2 + 2H2O = UO2(OH)2° + 2H -12.15 (Guillaumont et al., 2003) 2+ - + UO2 + 3H2O = UO2(OH)3 + 3H -20.25 (Guillaumont et al., 2003) 2+ 2- + UO2 + 4H2O = UO2(OH)4 + 4H -32.4 (Guillaumont et al., 2003) 2+ 3+ + 2UO2 + H2O = (UO2)2OH + H -2.70 (Guillaumont et al., 2003) 2+ 2+ + 53 2UO2 + 2H2O = (UO2)2(OH)2 + 2H -5.62 (Guillaumont et al., 2003)

2+ + + 3UO2 + 5H2O = (UO2)3(OH)5 + 5H -15.55 (Guillaumont et al., 2003) 2+ - + 3UO2 + 7H2O = (UO2)3(OH)7 + 7H -32.20 (Guillaumont et al., 2003) 2+ + + 4UO2 + 7H2O = (UO2)4(OH)7 + 7H -21.90 (Guillaumont et al., 2003) 2+ 2- UO2 + CO3 = UO2CO3 9.94 (Guillaumont et al., 2003) 2+ 2- 2- UO2 + 2CO3 = UO2(CO3)2 16.61 (Guillaumont et al., 2003) 2+ 2- 4- UO2 + 3CO3 = UO2(CO3)3 21.84 (Guillaumont et al., 2003) 2+ 2- 6- 3UO2 + 6CO3 = (UO2)3(CO3)6 54.00 (Guillaumont et al., 2003) 2+ 2- - + 2UO2 + CO3 + 3H2O = (UO2)2CO3(OH)3 + 3H -0.86 (Guillaumont et al., 2003) 2+ 2- + + 3UO2 + CO3 + 3H2O = (UO2)3O(OH)2(HCO3) + 3H 0.65 (Guillaumont et al., 2003) 11UO 2+ + 6CO 2- + 12H O = (UO ) (CO ) (OH) 2- + 2 3 2 2 11 3 6 12 36.40 (Guillaumont et al., 2003) 12H+ 2+ 2- UO2 + SO4 = UO2SO4 36.40 3.150 (Guillaumont et al., 2003) 2+ 2- 2- UO2 + 2SO4 = UO2 (SO4)2 4.140 (Guillaumont et al., 2003)

TABLE 2.4 (CONTINUED)

Reactions Log K (Reference) 2+ 2- 4- UO2 + 3SO4 = UO2 (SO4)3 3.020 (Guillaumont et al., 2003) 2+ - + UO2 + HSO4 = UO2 (HSO4) 2.95 (Martell et al., 2004) + 2- - H + SO4 = HSO4 1.980 (Guillaumont et al., 2003) + - K + OH = KOH(aq) -14.46 (Martell et al., 2004) + - K + HSO4 = KHSO4 0.85 (Martell et al., 2004) + - K + ClO4 = KClO4 -0.01 (Martell et al., 2004) + 2- - Na + CO3 = NaCO3 -1.27 (Martell et al., 2004) + 2- + Na + CO3 + H = NaHCO3(aq) -10.03 (Martell et al., 2004)

54 + - Na + OH = NaOH(aq) -14.18 (Martell et al., 2004) + 2- + Na + CO3 + H = NaHCO3(aq) -10.03 (Martell et al., 2004)

2.7 References

Brugger, J., Burns, P.C., Meisser, N., 2003. Contribution to the mineralogy of acid drainage of Uranium minerals: Marecottite and the zippeite-group. American Mineralogist, 88(4): 676-685.

Burns, P.C., Deely, K.M., Hayden, L.A., 2003. The crystal chemistry of the zippeite group. The Canadian Mineralogist, 41(3): 687-706.

Chen, F., Ewing, R.C., Clark, S.B., 1999. The Gibbs free energies and enthalpies of formation of U6+ phases: An empirical method of prediction. American Mineralogist, 84(4): 650-664.

Choppin, G., Mathur, J., 1991. Hydrolysis of actinyl (VI) cations. Radiochimca Acta, 52(1): 25-28.

Cox, J., Wagman, D.D., Medvedev, V.A., 1989. CODATA key values for thermodynamics. Chem/Mats-Sci/E.

Davis, J., Curtis, G.P., 2007. Consideration of geochemical issues in groundwater restoration at uranium in-situ leach mining facilities. Division of Fuel, Engineering, and Radiological Research, Office of Nuclear Regulatory Research, US Nuclear Regulatory Commission.

De Vivo, B., Ippolito, F., Capaldi, G., Simpson, P., 1984. Uranium geochemistry, mineralogy, geology, exploration and resources. Springer.

Finch, R., Murakami, T., 1999. Systematics and paragenesis of uranium minerals. Reviews in Mineralogy, 38: 91-180.

Frondel, C., Ito, J., Honea, R.M., Weeks, A.M., 1976. Mineralogy of the zippeite group. Canadian Mineralogist, 14(4): 429-436.

Gorman-Lewis, D., Burns, P.C., Fein, J.B., 2008. Review of uranyl mineral solubility measurements. The Journal of Chemical Thermodynamics, 40(3): 335-352.

Guillaumont, R. et al., 2003. Update on the Chemical Thermodynamics of Uranium, Neptunium, and Plutonium. Elsevier, Amsterdam, 919 pp.

Helean, K. et al., 2002. Enthalpies of formation of Ce-pyrochlore, 4+ 6+ Ca0.93Ce1.00Ti2.035O7.00, U-pyrochlore, Ca1.46U 0.23U 0.46Ti1.85O7.00 and Gd- pyrochlore, Gd2Ti2O7: three materials relevant to the proposed waste form for excess weapons plutonium. Journal of nuclear materials, 303(2): 226- 239.

Herbelin, A.L., Westall, J.C., 1999. FITEQL: A computer program for determination of chemical equilibrium constants from experimental data. Version, 4: 99-01.

Martell, A.E., Smith, R.M., Motekaitis, R., 2004. NIST critically selected stability constants of metal complexes. NIST standard reference database, 46(6.0).

Merkel, B., Planer-Friedrich, B., Wolkersdorfer, C., 2002. Uranium in the Aquatic Environment: proceedings of the International Conference [on] Uranium Mining and Hydrogeology III and the International Mine Water Association Symposium, Freiberg, Germany, 15-21 September 2002: with 453 figures, 151 tables and a CD-ROM. Springer Science & Business Media.

Molodetsky, I., Navrotsky, A., DiSalvo, F., Lerch, M., 2000. Energetics of oxidation of oxynitrides: Zr–N–O, Y–Zr–N–O, Ca–Zr–N–O, and Mg–Zr–N– O. Journal of Materials Research, 15(11): 2558-2570.

Navrotsky, A., 1977. Progress and new directions in high temperature calorimetry. Physics and Chemistry of Minerals, 2(1-2): 89-104.

Navrotsky, A., 1997. Progress and new directions in high temperature calorimetry revisited. Physics and Chemistry of Minerals, 24(3): 222-241.

Navrotsky, A., 2014. Progress and New Directions in Calorimetry: A 2014 Perspective. Journal of the American Ceramic Society, 97(11): 3349-3359.

O'Brien, T., Williams, P., 1981. The aqueous chemistry of uranium minerals. Part 3. Monovalent cation zippeites. Inorganic and Nuclear Chemistry Letters, 17(3): 105-107.

Plášil,‎J. et al., 2013. Meisserite, Na5(UO2)(SO4)3(SO3OH)(H2O), a new uranyl sulfate mineral from the Blue Lizard mine, San Juan County, Utah, USA. Mineralogical Magazine, 77(7): 2975-2988.

Plášil,‎J.,‎Kampf,‎A.R.,‎Kasatkin,‎A.V.,‎Marty, J., 2014b. Bluelizardite, Na7(UO2)(SO4)4Cl(H2O)2, a new uranyl sulfate mineral from the Blue Lizard mine, San Juan County, Utah, USA. Journal of Geosciences, 59(2): 145-158.

Plášil,‎J.‎et‎al.,‎2011b. The crystal structure of natural zippeite, K1. + 85H 0.15[(UO2)4O2(SO4)2(OH)2](H2O)4, from Jáchymov, Czech Republic. The Canadian Mineralogist, 49(4): 1089-1103.

Plášil,‎J.,‎Sejkora,‎J.,‎Škoda,‎R.,‎Škácha,‎P.,‎2014c.‎The‎recent‎weathering‎of‎ uraninite from‎the‎Červená‎vein,‎Jáchymov‎(Czech‎Republic):‎a‎fingerprint‎ of the primary mineralization geochemistry onto the alteration association. Journal of Geosciences, 59(3).

Robie, R.A., Hemingway, B.S., Fisher, J.R., 1978. Thermodynamic properties of minerals and related substances at 298. 15 K and 1 bar (10/sup 5/pascals) pressure and at higher temperatures, Geological Survey, Washington, DC (USA).

Ryan, M.P., Williams, D.E., Chater, R.J., Hutton, B.M., McPhail, D.S., 2002. Why stainless steel corrodes. Nature, 415(6873): 770-774.

Tingey, J.M., Bryan, G.H., DesChane, J.R., 2004. Dangerous waste characteristics of waste from Hanford Tank 241-S-109. Pacific Northwest National Laboratory.

Tripathi, V., 1983. Uranium transport modeling: geochemical data and sub- models Ph. D, thesis, Stanford University, Stanford, CA, USA.

Um, W. et al., 2009. Uranium phases in contaminated sediments below Hanford’s‎U‎tank‎farm.‎Environmental‎science‎&‎technology,‎43(12):‎4280- 4286.

Williams, D.E., Zhu, Y.Y., 2000. Explanation for initiation of pitting corrosion of stainless steels at sulfide inclusions. Journal of the Electrochemical Society, 147(5): 1763-1766.

CHAPTER 3:

INVESTIGATION OF THE STRUCTURAL STABILITY OF ZIPPEITE-GROUP

MINERALS USING HIGH-TEMPERATURE CALORIMETRY

3.1 Abstract

Accurate thermodynamic data for uranyl minerals are needed to understand their role in uranium deposit formation and subsequent alteration, transport of uranium in contaminated subsurface environments, and for predicting the performance of a geological repository for nuclear waste. Three members of the uranyl sulfate zippeite group, natrozippeite, cobaltzippeite, and zinczippeite have been synthesized and characterized prior to determination of their standard state enthalpies of formation from binary oxides and from elements using high-temperature drop solution calorimetry. Calculated standard state enthalpies of formation from the elements, ΔHf, at 298 K, were found to be -

18015.84 ± 28.76, -4362.09 ± 10.44, and -4698.22 ± 13.68 kJ/mol for natrozippeite (Na5[(UO2)8(SO4)4O5(OH)3]•12H2O), cobaltzippeite

(Co[(UO2)2(SO4)O2]•3.7H2O), and zinczippeite (Zn[(UO2)2(SO4)O2]•4.1H2O), respectively. Current results, together with the standard state enthalpy of formation value for zippeite (containing K) reported earlier, allows investigation of the role of acid-base interactions between the structural unit and interstitial complexes relative to the enthalpy of formation of zippeite group minerals. There

is a positive linear relationship between the formation enthalpies from oxides and the acidity of cation oxides. There is also a linear correlation between the ionic radius of charge-balancing alkalis and the enthalpy of formation of the studied compounds from their binary oxides, indicating the importance of the nature and coordination environment of the interstitial complex in determining the thermodynamic properties of these minerals.

3.2 Introduction

2+ Uranyl minerals, which contain the (UO2) uranyl ion, are chemically and structurally diverse, are important for understanding the history of uranium ore deposits (Finch and Murakami, 1999) impact the mobility of uranium in contaminated sites (Finch and Ewing, 1992) and form due to the alteration of irradiated nuclear fuel in laboratory simulations intended to model performance of a geological repository for nuclear waste (Finch and Ewing, 1992; Wronkiewicz et al., 1996; Zhang and Navrotsky, 2004). Uranyl sulfate minerals are amongst the more soluble of the uranyl minerals (Finch and Murakami, 1999). They often form at low pH in oxidizing conditions due to alteration of uraninite, UO2+x, and sulfides such as pyrite, marcasite, and chalcopyrite (Brugger et al., 2003; Frondel et al.,

1976;‎Plášil et al., 2014e). Where uraninite is undergoing oxidation and dissolution, uranyl sulfate aqueous complexes are important in the transport of U in some low pH groundwaters (Finch and Murakami, 1999). Uranyl sulfate minerals are particularly common in abandoned uranium mines where acid drainage high in sulfate ions flows over or through mine tailings or mine workings, or where uranium extraction was done by a sulfuric acid leaching

process (Meinrath et al.,‎2006;‎Plášil et al.,‎2014a;‎Plášil et al.,‎2011a;‎Plášil et al.,‎2013;‎Plášil et al.,‎2014c;‎Plášil et al., 2011c). Studies have also shown that uranyl sulfates may be present in tanks storing nuclear waste from weapons production at the Hanford, Washington site (Sharifironizi et al., 2016).

Zippeite group minerals are uranyl sulfates containing anionic sheets of uranyl pentagonal bipyramids and sulfate tetrahedra (Figure 3.1). The interlayer regions of the different species contain H2O and various cations that balance the charge of the uranyl sulfate sheet (Burns et al., 2003). Chemical analyses of synthetic materials and natural samples by Frondel at al. provided a general formula for zippeite-group minerals: Ax(UO2)6(SO4)3(OH)10•yH2O (with A = K, Na,

2+ 2+ NH4, Co, Ni, Fe , Mn , Mg, Zn; x=4 and y=4 for monovalent cations; and x=2 and y=16 for divalent cations) (Frondel et al., 1976). Later, a single-crystal X-ray study of several members of the zippeite group by Burns et al. led to revision of the formulae for K-, Na-, Co-, Zn-, Mg-, and NH4-bearing zippeites (Burns et al.,

2003). To date, different monovalent, divalent, and trivalent cations and REEs have been reported as interlayer cations in zippeite species: zippeite, natrozippeite, magnesozippeite, cobaltzippeite, zinczippeite, nickelzippeite, rabejacite, bluelizardite, marecottite, meisserite, sejkoraite, and pseudojohannite

(Brugger et al., 2003; Brugger et al., 2006; Frondel et al.,‎1976;‎Plášil et al.,

2014a;‎Plášil et al.,‎2011a;‎Plášil et al.,‎2013;‎Plášil et al., 2014c; Vochten et al.,

1995).

It is important to measure thermodynamic properties of uranyl minerals to underpin modeling efforts. High-temperature calorimetry in which well- characterized materials are dissolved in flux has emerged as a preferred method for measurement of reproducible and accurate thermodynamic data, and the enthalpies of formation have been reported for various uranyl oxide hydrates, silicates, carbonates, phosphates, vanadates, and sulfates using this approach

(Dzik et al., 2017; Gorman-Lewis et al., 2007; Gorman-Lewis et al., 2009; Guo et al., 2015; Guo et al., 2014; Kubatko et al., 2006; Kubatko et al., 2005; Kubatko et

al., 2003; Shvareva et al., 2011; Spano et al., 2017). Recently, the standard state thermochemical‎properties‎(∆Hfº,‎∆Sfº,‎∆Gfº) of zippeite,

K3(H2O)3.3[(UO2)4(SO4)2O3(OH)], were reported (Sharifironizi et al., 2016). Here we expand our study of the thermodynamic properties of zippeite-group minerals to synthetic analogs of natrozippeite, cobaltzippeite, and zinczippeite and probe the importance of interlayer cations and interlayer configuration on the thermodynamic properties of synthetic zippeites using high-temperature calorimetry.

3.3 Materials and Methods

Zippeite-structure compounds containing monovalent and divalent cations were synthesized by mild hydrothermal reactions of uranyl acetate or UO3 and

Na2SO4, CoSO4, or ZnSO4. To synthesize natrozippeite

(Na5[(UO2)8(SO4)4O5(OH)3]•12H2O), 0.244 g uranyl acetate was combined with

0.074 g sodium sulfate in a Teflon-lined 23 mL Parr reaction vessel with 4 mL of

18‎MΩ‎H2O. The pH of the solution, measured using an Orion Combination pH electrode calibrated against BDH Buffer Reference Standards, was 4.1.

Synthesis for cobaltzippeite (Co[(UO2)2(SO4)O2]•3.7H2O) was done by combining

0.228 g UO3 and 0.310 g CoSO4 in‎a‎reaction‎vessel‎with‎4‎mL‎of‎18‎MΩ‎H2O.

The initial pH of the mixture was 5 and was adjusted to 3 by adding ~1M HCl

(Burns et al., 2003). Zinczippeite (Zn[(UO2)2(SO4)O2]•4.1H2O) was synthesized by combining 0.228 g UO3 with 0.576 g ZnSO4 and‎4‎mL‎of‎18‎MΩ‎H2O in a reaction vessel (Burns et al., 2003). The initial pH was 5.5, and was adjusted to 5 using dilute HCl.

Reaction vessels were sealed and placed in a Fisher Scientific Isotemp oven at 170°C. Reactions for cobaltzippeite and zinczippeite were removed from the oven after 1 day, and that for natrozippeite after one week. Vessels were allowed to cool to room temperature after removal from the oven, and were then opened and the products were recovered by air-filtration followed by washing several‎times‎using‎18‎MΩ‎H2O. In order to accumulate enough material for characterization and calorimetric studies, the synthesis procedures were repeated several times and the products were combined for each synthetic phase.

Powder X-ray diffraction was used to characterize the purity and bulk phase composition of each sample. A 15 mg aliquot of each sample was ground and placed on a zero-background quartz plate for data collection using a Bruker

D8 Advance Davinci powder diffractometer equipped with CuKα‎radiation.‎A‎

LynxEye solid-state detector was used to collect data over an angular range of 5-

55°‎2θ‎with‎a‎step‎size‎of‎0.01°‎and‎an‎exposure‎time‎of‎1s.‎

Elemental analyses of the samples were done using a Perkin Elmer

Optima 8000 inductively coupled plasma-optical emission spectrometer (ICP-

OES). Samples were prepared by dissolving 0.1 - 0.15 mg of sample in 10 mL of aqueous HNO3 (5%). Concentrations of U, S, Na, Co, and Zn were calculated using external calibration and the internal standard addition method. To prepare external standards, standard solutions containing 1000 ppm of U, S, Na, Co, or

Zn were diluted in 5% HNO3 to concentration ranging from 0.1 to 15 ppm. Yttrium

was added to all samples and standards as an internal standard at a concentration of 0.6 ppm to monitor instrument drift.

Thermogravimetric and differential scanning calorimetry were done using a Mettler Toledo TGA-DSC1 instrument. Ten to 15 mg of sample was placed in an alumina crucible and heated at 5°C/min from room temperature to 700°C with an argon purge gas flow rate of 50 mL/min. The change in the mass of each sample upon heating in the range of 25-280°C was used to calculate the corresponding number of molecules of water per formula unit.

Calorimetry was done using a Setaram AlexSYS high-temperature oxide- melt drop-solution calorimeter. This instrument is appropriate for measuring the heat of dissolution of uranium minerals and compounds, and gives reliable and reproducible data (Navrotsky, 1977; Navrotsky, 2014). Four to six mg pressed pellets of a finely ground sample are dropped from room temperature to the working temperature of the calorimeter, 700 °C in this case, into a solvent.

Sodium molybdate, 3Na2O-4MoO3, was used as the solvent and the measured drop solution heats were used to calculate the enthalpies of formation of each sample using thermodynamic cycles (Table 3.1). The heat content of high-purity

α-Al2O3 was used to calibrate the AlexSYS. To ensure complete dissolution of samples in the solvent, ∼ 5 mg of zippeite samples were dropped into molten solvent at 700 ºC in a furnace, followed by visual inspection that confirmed a clear solution. Calorimetry analyses were conducted using a stream of O2 gas to flush the calorimeter and to mix the solvent by bubbling, thereby sweeping away any evolving gas/water vapors and maintaining an oxidizing environment.

3.4 Results and Discussion

The X-ray powder diffraction patterns for natrozippeite, cobaltzippeite, and zinczippeite were compared to patterns calculated using their structures from

Burns et al. (Burns et al., 2003) (Figure 3.4 in Supporting Information), and confirmed the purity of the material. Chemical analyses yielded compositions consistent with the previously published formula for each compound (Table 3.2 in

Supporting Information). Thermal analyses indicated 12.0, 3.7, and 4.1 moles of water per formula unit for natrozippeite, cobaltzippeite, and zinczippeite, respectively. The somewhat higher water contents, as compared to the ideal formulas, for cobaltzippeite and zinczippeite, are likely more reliable than those derived earlier from crystal-structure analysis that would not have been sensitive to minor departures from fully occupancy of H2O sites.

The standard state enthalpy of formation of each phase from its binary

o oxides‎(ΔHf-ox )‎and‎its‎elements‎(ΔHf°) was calculated from its respective measured drop solution enthalpy and the heat content of binary oxides as shown in Table 3.1. The standard-state enthalpies of formation from the elements are -

18015.84 ± 28.76, -4362.09 ± 10.44, and -4698.22 ± 13.68 kJ/mol for natrozippeite, cobaltzippeite, and zinczippeite, respectively. The calculated enthalpy of formation from the oxides for natrozippeite (-2110.19 ± 27.85 kJ/mol), cobaltzippeite (-322.89 ± 10.21 kJ/mol), and zinczippeite (-442.98 ± 13.56 kJ/mol) indicates they are stable in enthalpy relative to their binary oxides at standard state. Combining the results of this study with the previously reported standard state enthalpy of formation of zippeite, K3(H2O)3.3[(UO2)4(SO4)2O3(OH)],

(-1338.53 kJ/mol) (Sharifironizi et al., 2016) indicates that the presence of Na cations in the interlayer space leads to the release of more energy of formation when compared to mechanical mixtures of corresponding binary oxides. Given that these compounds contain very similar uranyl sulfate sheets, differences in entropy of their formation reactions should be relatively small, and thus the free energy of the reactions should follow a similar trend as seen here for enthalpy.

TABLE 3.1

CALORIMETRIC CYCLES FOR CALCULATION OF ENTHALPY OF FORMATION FOR URANYL SULFATES

Reactions ∆H ∆H (kJ/mol) ∆H (kJ/mol)

Na5[(UO2)8(SO4)4O5(OH)3]•12H2O (xl, 298K) = 1 ∆Hds (natrozippeite) 2.5Na2O (soln, 973K) + 8UO3 (soln, 973K) + 4SO3 (soln, 1758.91 ± 15.67 973K) + 13.5H2O (g,973K)

Co[(UO2)2(SO4)O2]•3.7H2O (xl, 298K) = CoO (soln, 973K) 412.87 ± 8.82 2 ∆Hds (cobaltzippeite) + 2UO3 (soln, 973K) + SO3 (soln, 973K) + 3.7H2O (g,973K) Zn[(UO ) (SO )O ]•4.1H O = ZnO 2 2 4 2 2 (xl, 298K) (soln, 560.56 ± 12.54 3 ∆Hds (zinczippeite) 973K) + 2UO3 (soln, 973K) + SO3 (soln, 973K) + 4.1H2O

(g,973K)

4 ∆Hds (UO3) UO3 (xl, 298K) = UO3 (soln, 976K) 9.49 ± 1.53

5 ∆Hds (SO3) SO3 (g, 298K) = SO3 (soln, 976K) -203.7 ± 4.09

6 ∆Hhc (H2O) H2O (l, 298K) = H2O (g, 976K) 69

7 ∆Hds (Na2O) Na2O (xl, 298K) = Na2O (soln, 976K) -217.56 ± 4.25

8 ∆Hds (CoO) CoO (xl, 298K) = CoO (soln, 976K) 15.66 ± 0.59

9 ∆Hds (ZnO) ZnO (xl, 298K) = ZnO (soln, 976K) 19.4 ± 0.7 0 10 ∆Hf (UO3) U (xl, 298K) + 3/2 O2 (g, 298K) = UO3 (xl, 298K) -1224 ± 0.80 0 11 ∆Hf (SO3) S (xl, 298K) + 3/2 O2 (g, 298K) = SO3 (xl, 298K) -395.7 ± 0.7 0 12 ∆Hf (H2O) H2 (g, 298k) + 1/2 O2 (g, 298K) = H2O (l, 298K) -285.8 ± 0.10 3

TABLE 3.1 (CONTINUED)

Reactions ∆H ∆H (kJ/mol) ∆H (kJ/mol) 0 13 ∆Hf (Na2O) 2Na (xl, 298K) + 1/2 O2 (g, 298K) = Na2O (xl, 298K) -414.82 ± 0.28 0 14 ∆Hf (CoO) Co (xl, 298K) + 1/2O2 (g, 298K) = CoO (xl, 298K) -237.94 ± 1.25 0 15 ∆Hf (ZnO) Zn (xl, 298K) + 1/2O2 (g, 298K) = ZnO (xl, 298K) -350.46 ± 0.27

2.5Na2O (s, 298K) + 8UO3 (s,298K) + 4SO3 (s, 298K) + 13.5H2O (l, 16 ∆H 0 (natrozippeite) 298K) -2110.19 ± 27.85 f-ox -∆H1‎+‎2.5∆H7‎+‎8∆H4‎+‎4∆H5‎+‎13.5∆H6

CoO (s, 298K) + 2UO3 (s, 298K) + SO3 (s, 298K) + 3.7H2O (l, 298K) 0 17 ∆Hf-ox (cobaltzippeite) -∆H2‎+‎∆H8‎+‎2∆H4‎+‎∆H5‎+‎3.7∆H6 -322.89 ± 10.21 68 ZnO + 2UO + SO + 4.1H O ∆H 0 (zinczippeite) (s, 298K) 3 (s, 298K) 3 (s, 298K) 2 (l, 298K) 18 f-ox -∆H3‎+‎∆H9‎+‎2∆H4 +‎∆H5‎+‎4.1∆H6 -442.98 ± 13.56

5Na (s, 298K) + 8U (s,298K) + 4S (s, 298K) + 13.5H2 (g, 298K) + 26 O2 (g, 298K) 0 19 ∆Hf (natrozippeite) -∆H1‎+‎2.5∆H7‎+‎8∆H4‎+‎4∆H5‎+‎13.5∆H6‎+‎2.5∆H13‎+‎ -18015.84 ± 28.76 8∆H10‎+‎4∆H11+‎13.5∆H12

Co (s, 298K) + 2U (s, 298K) + S (s, 298K) + 3.7H2 (g, 298K) + 6.85O2 (g, 298K) 0 20 ∆Hf (cobaltzippeite) -∆H2‎+‎∆H8‎+‎2∆H4‎+‎∆H5‎+‎3.7∆H6‎+‎∆H14‎+‎2∆H10‎+‎ -4362.09 ± 10.44 ∆H11‎+‎3.7‎∆H12

TABLE 3.2 (CONTINUED)

∆H Reactions ∆H ∆H (kJ/mol) (kJ/mol) Zn (s, 298K) + 2U (s, 298K) + S (s, 298K) + 4.1H2 (g, 298K) + 7.05O2 (g,

298K) -4698.22 ± 13.68 21 ∆H 0 (zinczippeite) -∆H3‎+‎∆H9‎+‎2∆H4‎+‎∆H5‎+‎4.1∆H6‎+‎∆H15‎+‎2∆H10‎+‎ f ∆H11‎+‎4.1∆H12

xl, solid material; g, gaseous; soln ,solution; and l, liquid. Error is represented as two standard deviations of the mean.

The compatibility of the Lewis-acidity of the structural unit (usually anionic overall, containing high cation-valence polyhedra) and the Lewis-basicity of the interstitial complex (interlayer cations and water) is important in determining the stability of the overall structure (Brown, 2002; Brown, 2009; Hawthorne, 2012;

Hawthorne, 2015; Hawthorne and Schindler, 2008). Given that zippeite, natrozippeite, cobaltzippeite and zinczippeite contain similar sheets of uranyl pentagonal bipyramids and sulfate tetrahedra, it is possible to evaluate the impact of cations on structural stability. Oxide acidity is used to probe the interactions between charge-balancing cations and the uranyl sulfate sheets of the zippeite group (Spano et al., 2017).

The contribution of acid-base interactions to structure stability was probed by plotting the formation enthalpies of the zippeite compounds as a function of cation acidity using the Smith scale (Figure 3.2) (Smith, 1987). To facilitate comparison, all values were normalized to the number of moles of uranium in the formula unit. The data show that the normalized enthalpy of formation from constituent oxides of the zippeite compounds becomes more exothermic with the decreasing acidity of the cation oxides, K2O, Na2O, CoO, and ZnO. The relationship in Figure 3.2 confirms the positive correlation between the acidity of the oxides and energetically favorable structures in enthalpy when compared to a mixture of binary oxides. Similar trends have been reported for uranyl silicates, uranyl hydroxide oxides, and uranyl vanadates (Navrotsky et al., 2013; Shvareva et al., 2012; Spano et al., 2017).

According to Hawthorne, chemical breakdown of a mineral is usually controlled by the strength of the relatively weak interactions between the structural unit and the interstitial complex (Schindler and Hawthorne, 2001;

Schindler and Hawthorne, 2008). Figure 3.3 shows that there is a linear relationship between the ionic radius of the interstitial cation in the zippeites studied and their normalized enthalpy of formation from their corresponding binary oxides. The most negative normalized formation enthalpy corresponds to zippeite, which contains the largest interstitial cation. The enthalpy of formation from binary oxides of the members of the zippeite group considered here is strongly impacted by the interstitial complexes (i.e. interlayer charge balancing cations), more so than the uranyl sulfate sheet. This is consistent with previous studies of uranyl minerals that highlighted the importance of alkali cations in the

energetics of their formation (Armstrong et al., 2012; Navrotsky et al., 2013;

Shvareva et al., 2012; Spano et al., 2017).

Figure 3.3. Enthalpy of formation of zippeite, natrozippeite, cobaltzippeite, and zinczippeite from oxides as a function of ionic radius of interlayer cation.

In summary, experimentally-determined enthalpies of formation for natrozippeite, cobaltzippeite, and zinczippeite indicate that they are energetically stable in enthalpy relative to their elements and corresponding binary oxides at standard state at 298 K. The relationship between the standard state enthalpy of the compounds and ionic radius and acidity of charge-balancing cations illustrates the importance of acid-base interactions between the zippeite-type sheet and interlayer cations in understanding the thermodynamic properties of this family of minerals.

3.5 Acknowledgements

This research was funded by the Office of Basic Energy Sciences of the

U.S. Department of Energy as part of the Materials Science of Actinides Energy

Frontier Research Center (DESC0001089). Partial support of this project was from the University of Notre Dame Center for Environmental Science and

Technology (CEST)/Bayer Predoctoral Fellowship. Powder X-Ray Diffraction and calorimetric data were collected in the Materials Characterization Facility (MCF) as part of the Center for Sustainable Energy at Notre Dame. ICP-OES analysis was conducted at the Center for Environmental Science and Technology (CEST) at Notre Dame.

3.6 Supporting Information

Figure 3.4- Powder X-ray diffraction patterns for Natrozippeite, Cobaltzippeite, and Zinczippeite.

TABLE 3.2

ICP-OES RESULTS FOR CHEMICAL ANALYSIS OF SYNTHETIC

NATROZIPPEITE, COBALTZIPPEITE, AND ZINCZIPPEITE

Sample Ratio Ideal Actual mol Na/mol U 0.63 0.63±0.01 Natrozippeite mol S/mol U 0.5 0.48±0.03 mol Co/mol U 0.5 0.5±0.02 Cobaltzippeite mol S/mol U 0.5 0.48±0.03 mol Zn/mol U 0.5 0.53±0.02 Zinczippeite mol S/mol U 0.5 0.47±0.03

3.7 References

Armstrong, C.R. et al., 2012. Uranyl peroxide enhanced nuclear fuel corrosion in seawater. Proceedings of the National Academy of Sciences, 109(6): 1874-1877.

Brown, I.D., 2002. The chemical bond in inorganic chemistry: the bond valence model, 12. Oxford University Press on Demand.

Brown, I.D., 2009. Recent developments in the methods and applications of the bond valence model. Chemical Reviews, 109(12): 6858-6919.