Abstract The Fukushima-Daiichi nuclear accident brought together compromised irradiated fuel and large amounts of seawater in a high radiation field.  Based on newly acquired thermochemical data for a series of uranyl peroxide compounds containing charge-balancing alkali cations, here we show that nanoscale cage clusters containing as many as 60 uranyl ions, bonded through peroxide and hydroxide bridges, are likely to form in solution or as precipitates under such conditions.  These species will enhance the corrosion of the damaged fuel and, being thermodynamically stable and kinetically persistent in the absence of peroxide, they can potentially transport uranium over long distances. 

Failed cooling systems in the reactors and spent fuel cooling ponds at the Fukushima-Daiichi nuclear power plants resulted in compromised irradiated fuel and release of radionuclides.  Copious amounts of seawater were subsequently used to cool the fuel.  The collocation of large quantities of damaged fuel, an intense radiation field, and massive amounts of seawater has created a highly heterogeneous and presumably rapidly evolving system that has the potential to release vast quantities of radionuclides to the environment.  Currently, large quantities of contaminated water remain onsite (1), and presumably some has been released to the subsurface as well as to the Pacific Ocean. 

The fuel matrix at the Fukushima-Daiichi site is mainly UO 2 , whose behavior will largely dictate release of matrix-incorporated plutonium and various other radionuclides into water used as a coolant.  The intense radiation field of the fuel will cause radiolysis of water and formation of peroxide (as well as other species) (2).  Peroxide enhances the corrosion rate of UO 2 by oxidizing U(IV) to the much more soluble U(VI) that exists as the linear dioxo uranyl cation, (UO 2 )2+.  Peroxide strongly complexes uranyl (3), which increases its aqueous solubility under alkaline conditions. 

Simple uranyl peroxide complexes contain a single uranyl ion coordinated by as many as three peroxide groups that, each being bidentate, define the equatorial edges of hexagonal bipyramidal coordination polyhedra (4) (Fig.  1).  These small complexes associate with counterions locally to balance charge in solution and readily crystallize as alkali metal salts (4–6).  When peroxide bridges uranyl ions, the configuration is bent (7–9) and nanoscale cage clusters containing as many as 60 uranyl ions self-assemble in aqueous systems (10) (Fig.  1).  These soluble clusters carry negative charges, are associated with counterions in solution, and can be crystallized.  Under acidic conditions in deionized water, the combination of uranyl and peroxide causes the precipitation of studtite, [(UO 2 )(O 2 )(H 2 O) 2 ](H 2 O) 2 (11). 

Fig.  1.  Ball-and-stick and polyhedral representations of the uranyl ion coordinated by three bidentate peroxide groups as in LiUT, NaUT, and KUT (A, B); and two bidentate peroxide groups and two hydroxides as in U60 cage clusters (C, D).  Uranium and O atoms are shown as yellow and red spheres, respectively. 

Some insight into the geochemical interactions of uranium and peroxide that may occur in the Fukushima-Daiichi systems emerge from cases where deionized water, UO 2 , and ionizing radiation have been combined.  Irradiated fuel was stored in the K-East Basins of the plutonium-production facility at the Hanford Site under 3.7 m of water maintained at 10 °C and continuously deionized by pumping through ion exchange columns.  Studtite was a major alteration phase of fuel-element claddings, and also occurred on the basin floor and in canister sludge (12).  Studtite also formed on spent fuel in deionized water under laboratory conditions (13), and on UO 2 doped with alpha emitters or irradiated by an external source in water (2, 14–17). 

Where seawater flow is relatively stagnant in the Fukushima-Daiichi systems, peroxide will accumulate, it will increase corrosion rates of exposed UO 2 , and it will complex uranyl ions in solution.  The seawater, with a pH of approximately 8, will provide abundant Na and lesser K to balance the charge of the resulting uranyl peroxide complexes that form in solution and in precipitates. 

We have measured enthalpies of formation enthalpies of formation of model uranyl peroxide compounds to gain insight into the formation and stability of simple and complex uranyl peroxide species in aqueous solution, as well as the solid phases that may form and persist (e.g., upon reduction of water volume due to evaporation or boiling, and eventual drying of the systems).  We selected model compounds that could be synthesized with high purity and yield, as detailed in Materials and Methods.  The salts M[(UO 2 )(O 2 ) 3 ](H 2 O) 9–10 [M = Li(10H 2 O), Na(9H 2 O), K(9H 2 O)] contain uranyl ions coordinated by three peroxide groups and linked through the corresponding counterions and H bonding networks (designated LiUT, NaUT, and KUT, respectively) (4–6).  The water-soluble U60 nanoscale cage cluster contains 60 uranyl ions that are bridged through bidentate peroxide as well as through hydroxyl groups (Fig.  1) (18), and crystals have composition Li 40 K 20 [UO 2 (O 2 )(OH)] 60 (H 2 O) 214 .  U60 is used here as a model for the more than 30 known cage clusters (containing between 20 and 60 uranyl ions) that are built from uranyl ions bridged through peroxide and hydroxide groups (10).  Given that the role of alkali counterions in templating uranyl peroxide cage clusters is poorly understood, it is difficult to predict which uranyl peroxide cluster(s) are most likely to form where Na is abundant; indeed we crystallized several with Na as a counterion. 

Results and Discussion Enthalpies of formation from oxides (normalized to 1 mol of U) as a function of alkali ionic radius are shown in Fig.  2.  Fig.  2.  Comparison of enthalpies of formation from the oxides for LiUT, NaUT, KUT and the U60 nanocluster, normalized per mole of U. Note: Studtite data from Kubatko Hughes, et al. ( 19) are also included.  Insight can be gained by directly comparing formation reactions under environmentally relevant conditions, i.e., equilibria involving UO 2 (c), aqueous systems, and pertinent secondary mineral phases.  In contrast to studtite, which has limited thermodynamic stability (19), alkali uranyl peroxide formation from the oxides features highly exothermic values (Tables 1 and 2; Eq.  1): [1]Comparing the enthalpies of formation of these compounds offers insight into relative stability (Fig.  2), because each phase involves the same amount of oxygen in the reactants side of the reaction Eq.  1, differences in entropy of reaction (ΔSo) should be small.  Thus the free energy of the reactions, ΔGo (i.e., stability) parallels .  Consistent with studtite (5), the formation of alkali uranyl peroxide species in solution, and the precipitation of M 4 [(UO 2 )(O 2 ) 3 ](H 2 O) 9–10 solids, requires the corresponding alkali metal ion, water, and peroxide (Reactions A, B, and C), as reactions without peroxide are clearly endothermic (Reactions D, E, and F).  An overall increase in stability is evidenced down the alkali series, with a possible small maximum in exothermicity observed with NaUT (Fig.  2).  Analogous patterns of energetics as a function of ionic radius for alkali-bearing compounds are well established.  For example in alkali uranates (Li-Cs), the maximum stability is observed in K 2 UO 4 (20, 21).  Table 1.  Summary of thermodynamic data for uranyl peroxide compounds in kJ/mol.  ΔH ds : enthalpy of drop solution; ΔH f,ox : enthalpy of formation from the oxides; ΔH f,el : enthalpy of formation from the elements Table 2.  Enthalpies of reactions involving uranyl peroxide compounds In contrast to the M 4 [(UO 2 )(O 2 ) 3 ](H 2 O) 9–10 phases and studtite, the formation of crystals containing the U60 cluster in equilibrium with an aqueous phase and UO 2 does not require excess peroxide (Reaction G).  In addition, the formation of the U60 cluster from the pertinent M 4 [(UO 2 )(O 2 ) 3 ](H 2 O) 9–10 compounds and studtite is also favorable (Reaction H), suggesting that M 4 [(UO 2 )(O 2 ) 3 ](H 2 O) 9–10 with Li and K, as well as studtite, are metastable with respect to crystals containing U60.  At present it is not possible to rigorously determine the stability of crystals containing the U60 cluster with respect to common uranyl oxyhydrates because (i) a U60 cluster with Na as the only counterion has not yet been synthesized, and (ii) no enthalpy data exist for Li uranyl oxyhydrates.  However, because the exchange of Na for Li in the formation of U60 without peroxide (analogous to Reaction G) is also exothermic (-9,114 kJ/mol), the formation of a Na-containing U60 cluster is probably thermodynamically favorable.  The energetics of Na-bearing U60 formation and crystallization from an aqueous system in contact with UO 2 has been estimated in Reaction L. This reaction is exothermic, indicating that crystals of a Na-K bearing U60 are metastable with respect to secondary U(VI) phases (Na compreignacite and K compreignacite).  In the absence of the aqueous phase, uranyl peroxide compounds are expected to eventually decompose to a stable uranyl oxyhydrate; e.g., metaschoepite (Reaction I), clarkeite (Reaction J), Na compreignacite (Reaction K), or Na and K compreignacite (Reaction L).  In any case, due to its redox properties and its high affinity for uranyl ion complexation (3), peroxide significantly affects U mobility by acting as a catalyst in U oxidation, by increasing the solubility of the uranyl ion, and by enhancing the production of relatively stable secondary uranyl oxyhydrate phases.  Although it is not currently possible to measure the enthalpies of formation of aqueous alkali uranyl peroxide species (22), or complex cage clusters in solution, our studies of the corresponding solid phases provide useful insight pertaining to the interaction of irradiated fuel and seawater.  Peroxide production will be the highest where water contacts fuel directly (owing to the importance of alpha radiolysis), and it will achieve the highest concentration locally in the water near the surface of the fuel, where the fuel to water ratio is high, and where water is relatively stagnant.  Possible locations where peroxide buildup will be greatest at Fukushima-Daiichi include spent fuel cooling pools and areas in the reactor cores where water intimately interacts with UO 2 fuel.  We contend that simple uranyl peroxide species, as well as complex nanoscale cage clusters built from uranyl peroxide polyhedra, will form in these environments.  The thermochemical data further demonstrate the importance of uranyl peroxides in systems combining intense radioactivity, water, and UO 2 .  Precipitation of compounds such as M 4 [(UO 2 )(O 2 ) 3 ](H 2 O) 9–10 is expected in such systems, and especially the thermodynamically stable NaUT phase in the case of seawater.  Self-assembly of nanoscale uranyl peroxide clusters in such systems is also likely, and, at least in the case of U60, these are thermodynamically stable in the absence of excess peroxide.  As such, these clusters are expected to persist in water even after peroxide production stops (i.e., the water is no longer in contact with fuel).  Where clusters become sufficiently concentrated, precipitation may occur to give thermodynamically stable cluster compounds that eventually convert to alkali uranyl oxyhydrates, carbonates, silicates, and other solid phases.  These uranyl peroxide cluster compounds are thus an energetic intermediate between dissolved aqueous species and uranyl minerals.  The thermodynamic stability of such clusters in the absence of excess peroxide also indicates they may disperse in the environment through transport in water.  To verify the persistence of U60 clusters in aqueous solution in the absence of excess peroxide, crystallized U60 was dissolved in ultrapure water and the resulting solution was monitored using electrospray ionization mass spectroscopy (ESI-MS).  The spectrum attributable to U60 persisted for the solution through at least 294 d without detectable change (Fig.  S1). 

Materials and Methods Synthesis of Li 4 UO 2 (O 2 ) 3 ·10H 2 O (LiUT).  Aqueous lithium hydroxide solution (4 M, 4 mL) and 30% H 2 O 2 (3 mL) are combined in a 40 mL beaker with a stir bar and placed in an ice bath.  Uranyl nitrate hexahydrate (0.5 g, ∼1 millimole) was dissolved in 6 mL deionized H 2 O, that was previously cooled in the ice bath.  While stirring, the uranyl nitrate solution was added to the LiOH/H 2 O 2 mixture via pipette, and a clear, orange, bubbly solution formed.  The beaker was then placed inside of a large jar containing a layer of ethanol on the bottom, and the jar was capped with a lid.  The ethanol diffuses slowly into the aqueous-uranyl solution.  After 1 d, small orange-yellow dodecahedral crystal have appeared; these are the prior reported [UO 2 (O 2 ) 3 ] 12 [(UO 2 (OH) 4 Li 16 (H 2 O) 28 ] 3 ·Li 6 [H 2 O] 26 (6).  After 5 d, long yellow blade-like crystals grew out from the original crop of crystals.  Crystals of LiUT were harvested by breaking of the dodecahedral crystals at the tips.  Yield ∼0.32 g, 55% based on U. The blade-like crystals were characterized by single-crystal X-ray diffraction, and the bulk was characterized by powder X-ray diffraction (5–60° 2-theta, Cu-Kα radiation) to ensure that the single-crystals are representative of the bulk, and the bulk is pure-phase (Fig.  S2).  A representative figure of the structure is shown in Fig.  S3.  Synthesis of Na 4 UO 2 (O 2 ) 3 ·9H 2 O (NaUT).  Aqueous sodium hydroxide solution (4 M, 4 mL) and 30% H 2 O 2 (3 mL) were combined in an 80 mL beaker with a stir bar and placed in an ice bath.  Uranyl nitrate hexahydrate (0.5 g, ∼1 millimole) was dissolved in 6 mL deionized H 2 O, that was previously cooled in the ice bath.  While stirring, the uranyl nitrate solution was added to the NaOH/H 2 O 2 mixture via pipette, and a clear, orange, bubbly solution formed.  Upon rapid addition of ethanol, a heavy orange precipitate formed, and the liquid became colorless.  The crystalline precipitate was isolated via vacuum filtration and washing with additional ethanol, and left to dry in air.  Yield = 0.55 grams, 88%.  This product is characterized by powder X-ray diffraction (5–60° 2-theta, Cu-Kα radiation) and thermogravimetric analysis.  Both characterization techniques confirm the sodium uranyl peroxide salt reported by Alcock was formed (4, 6).  Synthesis of K 4 UO 2 (O 2 ) 3 ·9H 2 O (KUT).  The synthesis of a mixed Cs/K or Rb/K salt similar to KUT was described previously (22).  The pure K-salt was synthesized in this study.  Aqueous potassium hydroxide solution (4 M, 4 mL) and 30% H 2 O 2 (3 mL) was combined in a 40 mL beaker with a stir bar and placed in an ice bath.  Uranyl nitrate hexahydrate (0.5 g, ∼1 millimole) was dissolved in 6 mL deionized H 2 O, that was previously cooled in the ice bath.  While stirring, the uranyl nitrate solution was added to the KOH/H 2 O 2 mixture via pipette.  Initially an orange solution formed followed by the rapid precipitation of a bright yellow crystalline salt.  The salt was isolated by vacuum filtration and washing with ethanol.  X-ray powder diffraction, an SEM image, and TG-DTA (thermogravimetric-differential thermal analysis) analysis of KUT are shown in Figs.  S4, S5, and S6, respectively.  Due to the high solubility of this salt, and the fact that LiUT and NaUT dominate under identical conditions in a freshly prepared solution, we reason that this salt is KUT.  TGA analysis gave a formula consistent with nine water molecules (calculated, 31% wt loss; observed, 32% wt loss).  Energy Dispersive Spectroscopy gave a K∶U ratio of 4∶1.  Synthesis of U60 Nanoclusters.  U60 crystals were synthesized using methods described by Sigmon, et al. ( 18).  In summary, using standardized materials in a 20 mL scintillation vial: 1 mL (0.436 M uranyl nitrate hexahydrate), 0.25 mL (0.285 M potassium chloride), and 1 mL of 30% H 2 O 2 were combined and the pH was adjusted to 9.0 using 2.203 M lithium hydroxide aqueous solution.  After 7 d, large, yellow crystals formed in solution.  Identity of U60 was confirmed by single-crystal X-ray diffraction.  Single-Crystal X-Ray Diffraction.  Single-crystal X-ray diffraction of LiUT was performed at -85 °C on a Bruker AXS SMART-CCD diffractometer with graphite monochromated Mo-Kα (0.71073 Å) radiation.  Data collection and reduction were carried out with SMART 5.054 (23) and SAINT 6.02 (24) software, respectively.  A numerical absorption correction from face indexing was applied.  The structure was solved by Direct Methods [program SIR97 (25)] and refined by full matrix least-squares on the basis of F2 using SHELX97 (26).  Formula: H 20 O 18 Li 4 U; FW = 575.97; Monoclinic, P2 1 /c(#14), a = 7.4761(8), b = 13.7915(14), c = 14.9976(16) Å, β = 98.659(2)°, V = 1,528.7(3) Å3, Z = 4, ρ calcd = 2.503 Mg·m-3; μ(MoKα) = 10.704 mm-1; 2.02 ≤ θ ≤ 26.40°; R 1 [I > 2σ(I)] = 0.0168; wR 2 [I > 2σ(I)] = 0.0391, GOF = 1.03.  Electrospray Ionization Mass Spectra (ESI-MS).  ESI-MS were recorded on a Bruker microOTOF-Q II high resolution quadrapole time of flight instrument (Q-TOF) (3,200 V capillary voltage, 0.4 Bar nebulizer gas, 4 L/ min dry gas, 180 °C dry gas temperature, negative-ion mode).  The samples were introduced by direct infusion at 7 μL/ min and scanned over 500–5,000 m/z with data averaged over 2–5 min.  Data was deconvoluted using the MaxEnt algorithm software.  Calorimetry.  High temperature oxide melt solution calorimetry was conducted using a Tian-Calvet twin microcalorimeter.  The details of this instrument and experimental conditions for study of uranyl-bearing materials are provided elsewhere (19, 27, 28).  Small pressed pellets (∼5 mg) of powdered sample were dropped from room temperature into a melt of 20 g of sodium molybdate (3Na 2 O·4MoO 3 ) at 976 K inside platinum crucibles.  The calorimeter was calibrated prior to analysis with ∼5 mg pellets of a standard reference of α-Al 2 O 3 .  To ensure an oxidizing environment, O 2 was continuously flushed over the melt head space.  Gas flushing also assisted in sweeping evolved H 2 O out of the calorimeter.  Upon rapid and complete dissolution of the sample, the enthalpy of drop solution ΔH ds , was measured, and, using appropriate thermochemical cycles (see Tables S1, S2, and S3), enthalpies of formation from the oxides, ΔH f,ox , and from the elements, ΔH f,el , were calculated. 

Acknowledgments Jennifer Szymanowski conducted TG and ESI-MS analysis.  This work was supported as part of the Materials Science of Actinides, an Energy Frontier Research Center funded by the Department of Energy, Office of Science, Office of Basic Energy Sciences under award number DE-SC0001089.  Sandia National Laboratories is a multiprogram laboratory operated by Sandia Corporation, a wholly owned subsidiary of Lockheed Martin company, for the Department of Energy’s National Nuclear Security Administration under contract DE-AC04-94AL85000.