Runde 2000 - Waste Isolation Pilot Plant

Runde 2000 - Waste Isolation Pilot Plant

Chemical Interactions of Actinides in the EnvironmentTable I. Oxidation States of Light Actinides aTh Pa U Np Pu Am CmIII III III III III III IIIIV IV IV IV IV IV IVV V V V VVI VI VI VIVII VIIa The environmentally most important oxidation states are bolded.the environment, such as coprecipitation,mineralization, and diffusionprocesses (and have thus included themin Figure 1), we will not discuss theirroles here.Oxidation States and RedoxBehaviorWater is the dominant transportmedium for most elements in the environment.Compared with the pH valuesand ionic strengths that can be obtainedin the laboratory, most natural watersare relatively mild. Typically, they arenearly neutral (pH 5 to 9), with a widerange of redox potentials (from –300 to+500 millivolts) and low salinities(ionic strengths ≤ 1 molal). The waterconditions determine which actinideoxidation states predominate and whichactinide species are stable.For example, Figure 2 shows thePourbaix diagram (Eh vs pH) for plutoniumin water containing the two mostenvironmentally relevant ligands, thehydroxide (OH – ) and carbonate (CO 3 2– )ions. Even in this simple aqueoussystem, plutonium can exist in fouroxidation states: III, IV, V, and VI.(Plutonium is unique in this regard.The Pourbaix diagrams for the otheractinides are simpler.) The normalrange of natural waters is outlined inthe figure and overlaps with the stabilityfields of plutonium in the III, IV, and Voxidation states. Within this range,plutonium exhibits two triple points,where species in three different oxidationstates are in equilibrium.Because of intrinsic differences inredox potentials, each actinide willexhibit a different set of oxidationstates for a given set of solution conditions.In contrast to plutonium, U(III) isunstable under most conditions andoxidizes easily to U(IV), while U(V)disproportionates easily to U(IV) andU(VI). Neptunium(III) and Np(VI) areon the edges of the water stabilityregion and can exist only under stronglyreducing or oxidizing conditions,respectively. Americium and curiumwill only be found in the III oxidationstate under most conditions. Similarly,all actinides beyond curium are dominatedby the lanthanide-like trivalentoxidation state.As a rule of thumb, we expect tofind U(VI), Np(V), Pu(IV), Am(III),and Cm(III) as the prevalent oxidationstates in most ocean or groundwaterenvironments. But in other aqueous environments,including streams, brines,or bogs, U(IV), Np(IV), andPu(III,V,VI) are also common and likelyto be stable. Table 1 summarizes theoxidation states of the actinides andhighlights the environmentally mostrelevant ones. Some of the states listedin the table, such as Pa(III) or Pu(VII),can be synthesized only under extremeconditions, far from those found innatural systems.Additional chemical processes occurringin solution are likely to influencethe actinide’s oxidation state stability.The stability of Pu(V) in natural waterscontaining carbonate is an example.The plutonium will complex with thecarbonate ligands, and if the plutoniumconcentration is low (less than about10 –6 M), radiolytically induced redoxreactions are minimized. Consequently,the stability of Pu(V) is enhanced andits disproportionation to Pu(IV) andPu(VI) is reduced. As discussed later,the solubilities of solids formed fromPu(V) complexes are orders of magnitudegreater than those of Pu(IV) solids,so the enhanced stability of Pu(V)would serve to increase the total plutoniumconcentration in solution.Another example for the stabilizationof actinides in solution is the oxidationof Am(III) and Pu(IV) through radiolyticformation of oxidizing species, such asperoxide (H 2 O 2 ) or hypochlorite(ClO – ). At high actinide concentrationsand hence under the influence of (itsown) alpha radiation, those normallystable oxidation states are oxidized toform Am(V) and Pu(VI), respectively,especially in concentrated chloridebrines.Despite the complexity of actinideredox behavior, however, we shouldemphasize that within a given oxidationstate, actinides tend to behave similarly.For example, we can study a U(IV)complex and from it infer the behaviorof the analogous Np(IV) and Pu(IV)complexes. In the following discussions,therefore, we often refer to“generic” actinides—e.g., An(IV)complexes—where An is shorthandfor actinide.Effective Charge. It has beenknown for many years that An(IV)forms the strongest, most stablecomplexes and An(V) the weakest. Thisbehavior follows directly from theeffective charges of the ion.In the III and IV states, the actinidesform hydrated An 3+ and An 4+ ions insolution, respectively. In contrast, thehighly charged ions in the V and VIstates are unstable in aqueous solutionand hydrolyze instantly to form lineartrans-dioxo cations, AnO 2 + andNumber 26 2000 Los Alamos Science 395

Chemical Interactions of Actinides in the EnvironmentAnO 2 2+ , with overall charges of +1 and+2, respectively. These cations areoften referred to as the actinyl ions.The covalent bonding between theactinide and the two oxygen atoms inthe actinyl ion (O=An=O) n+ , wheren = 1 or 2, enhances the effectivecharge of the central actinide ion to2.3±0.2 for AnO 2 + and 3.3±0.1 forAnO 22+ (Choppin 1983). A ligandapproaching the actinyl ion sees thisenhanced effective charge and bondsdirectly to the actinide in the equatorialplane of the linear actinyl ion. Thus, thepreference for actinides to form complexesgenerally follows their effectiveion charges, as shown below:An 4+ > AnO 2+ 2 ≥An 3+ > AnO +2+4 +3.3 +3 +2.3IV VI III VThe second and third lines show theion’s effective charge and formal oxidationstates.Recent measurements of bondlengths in the actinyl ions of V and VIcomplexes reflect the above sequence.A higher effective charge correlateswith a stronger, hence shorter, bond tothe coordinated ligands. Accordingly,the bond length of the actinyl bonds increasesfrom about 1.75 angstroms (Å)in the VI oxidation state (U=O is 1.76 Å;Pu=O is 1.74 Å) to about 1.83 Å in theV oxidation state (Np=O is 1.85 Å;Pu=O is 1.81 Å).Because An(IV) has the highesteffective charge, it forms the moststable complexes in solution and alsoforms the most stable precipitates withthe lowest solubility. Conversely,An(V) complexes are the weakest, andits solubility-controlling solids are themost soluble. An(V) species are thereforethe most likely to migrate throughthe environment. Table 1 shows that theonly transuranics that favor the V stateare neptunium and plutonium. But inthe environments surrounding repositorysites, as well as in most groundwatersand ocean environments, plutonium willmost likely assume the IV state. Thusneptunium is predicted to be the mostsoluble and easily transported actinideand is the actinide of major concern inassessing the performance of YuccaMountain.Coordination Number and IonicRadii. Another fundamental aspectabout the actinide cations is that theytend to act as hard Pearson “acids,”which means that they form strongActinides in Today’s EnvironmentIn 1942, Enrico Fermi built mankind’s first fission reactor, and nuclear reactions withinthat nuclear pile created the transuranic elements neptunium, plutonium, americium,and curium. With the subsequent development of nuclear power, the inventory oftransuranics has increased dramatically. As of 1990, the estimated accumulations inspent nuclear fuel were approximately 472 tons of plutonium, 28 of tons of neptunium,6 tons americium, and 1.6 tons of curium.Fortunately, the amount of radioactive material accidentally released from nuclearpower plants has been relatively small. Environmental contamination is severe in only afew locations. Major releases were due to a fire in a reactor at Windscale (nowSellafield) in the United Kingdom in 1957, an explosion at the Kyshtym nuclear wastestorage facility in the former USSR in 1957, and an explosion and fire at the Chernobylreactor in Russia in 1986. In addition, a number of DOE sites in the United States,including the Hanford Site in Washington, Rocky Flats Environmental Technology Sitein Colorado, and Idaho National Engineering and Environmental Laboratory in Idaho,contain high local concentrations of actinides.The main source of transuranics in the environment has been nuclear weapons testing.Since the first detonation of a nuclear device in 1945 at Trinity Site in Alamogordo,New Mexico, more than 4 tonnes of plutonium have been released worldwide (accountingonly for announced nuclear tests). Atmospheric and underground tests have alsoinjected about 95 kilograms of americium into the environment, along with tiny amountsof the still heavier element curium.Until they were banned by international treaty in 1963, aboveground nuclear testspropelled the actinides directly into the atmosphere. There they dispersed and migratedaround the globe before settling back to earth. Little can be done to isolate or retrievethis material. It is a permanent, albeit negligible, component of the earth’s soils andoceans and poses little health risk to the general public.Underground nuclear tests injected transuranics into highly localized regions surroundingthe detonation sites. These concentrated actinide sources became mineralized andintegrated into the rock matrix once their areas cooled and solidified. For the most part,the nuclear material is fixed in place. However, plutonium from a test conducted at theNevada Test Site in 1968 has been detected in a well more than a kilometer away. Itsremarkably fast migration rate (at least 40 meters per year) is likely due to colloidaltransport through the highly fractured, water-saturated rock surrounding the detonation site.complexes with highly anionic ligandsby electrostatic interactions. Becausethe actinide-ligand bond is substantiallyionic, the number of bound ligands (thecoordination number) and the relativepositions of the ligands around thecation are determined primarily bysteric and electrostatic principles.Consequently, for a given oxidationstate, a range of coordination numbers396 Los Alamos Science Number 26 2000

Chemical Interactions of Actinides in the EnvironmentAn(III) concentration (M)Total solvated An(III) concentration (M)(a)10 –2An(OH) 3 (s)10 –4AnOHCO 3 (s)An 2 (CO 3 ) 3 (s)10 –610 –8An 3+ (aq)10 –10 10 –5 10 –4 10 –3 10 –2 10 –1 10CO 2 Partial Pressure (atm)(b)10 –4 244 CmAn 2 (CO 3 ) 3 •2–3 H 2 O(s)241 Am10 –5An 3+An(CO 3 ) 3–An(CO 3 ) +310 –6An(CO 3 ) –2An(OH) 2+10 –710 –7 10 –6 10 –5 10 –4 10 –3 10 –2Carbonate concentration (M)Figure 3. Total An(III) ConcentrationVersus Carbonate Concentration(a) The partial pressure of CO 2 in naturalwaters spans about three orders ofmagnitude (gray area). The stabilitydiagram for An(III) solid phases (at pH 7)shows that the An(III) hydroxocarbonateis favored in most natural waters. Lowerpartial pressures or soluble carbonateconcentration translate to increasedstability of the An(III) hydroxide; higherpressures mean increased stability forthe normal carbonate solids. (b) The redcurve is the total concentration of An(III)in solution when the solubility-controllingsolid is An 2 (CO 3 ) 3 •2–3H 2 O(s). Calculatedconcentrations for individual solutionspecies are indicated by the gray linesthat run tangent to the solubility curve,or by the gray dashed lines. At lowcarbonate concentrations, the An 3+ iondominates the solution species. The totalAn(III) concentration drops rapidly withhigher carbonate concentrations as moreof the solid phase precipitates. Butcarbonate complexation begins to stabilizethe amount of actinide in solution.By the time the anionic bis- and triscarbonatecomplexes dominate, the solutionconcentration allowed. Coordination numbers between6 and 12 have been reported forcomplexes of An(III) and An(IV) andbetween 2 and 8 for the actinyl ions.For example, the aquo ions of theactinides, which have only water moleculesin the first coordination spheresurrounding the ion, exhibit coordinationnumbers that vary with oxidationstate: 8–10 for An 3+ , 9–12 for An 4+ ,4–5 for AnO 2 + and 5–6 for AnO 2 2+ .These coordination numbers have beendetermined from nuclear magentic resonance(NMR) and XAFS spectroscopystudies (see the article by Conradson etal. on page 422).Generally, the stabilities of actinidecomplexes at a given oxidation stateincrease with atomic number. Thisincrease in stability follows the trend ofthe actinide contraction, wherein theionic radii decrease with atomicnumber. The ionic radii for An(IV) ionswith coordination numbers of 8 arereported to be 1.00, 0.98, 0.96, 0.95,and 0.95 Å for uranium, neptunium,plutonium, americium, and curium, respectively(Shannon 1976).Solubility and SpeciationThe solubility of an actinide is limitedprimarily by two properties: thestability of the actinide-bearing solid(the solubility-controlling solid) and thestability of the complexed species insolution. An actinide species willprecipitate when its dissolved concentrationis above the thermodynamic(equilibrium) solubility of its solubilitycontrollingsolid, or if the kineticsfavors the formation of a solid of highersolubility. Formation of the solid setsan upper concentration limit for theactinide in solution.Figure 3 illustrates these concepts forAn(III), where An is either americium orcurium. The data points in the lowergraph show the measured total concentrationof An(III) in solution. At lowpH with increasing carbonate concentration,the hydrated carbonated solidAn 2 (CO 3 ) 3 •2–3H 2 O(s), where (s) indicatesthe solid phase, readily precipitatesand hence lowers the concentrationof An 3+ in solution. But complexationof An 3+ with carbonate stabilize thesolution species and increases the overallAn(III) solubility. As the carbonateconcentration increases and the differentcarbonate species form, the totalactinide concentration in solution firstNumber 26 2000 Los Alamos Science 397

Chemical Interactions of Actinides in the Environmentreaches a minimum and then increasesdue to the formation of anionic An(III)complexes. The solubility would changeif conditions favored the formation of anew solubility-controlling solid.We obtain solubility data byperforming experiments in the laboratoryunder well-controlled conditions inwhich the actinide concentration ismeasured while varying the ligandconcentration. Such experiments enableus to measure the solubility product K spand the stability constant β. Thesethermodynamic parameters are the basisfor modeling the solubility boundariesfor actinides in natural waters, asdiscussed in the box on “Solubility andSpeciation Parameters.”We should note, however, thatmeaningful interpretation of solubilitydata requires a detailed knowledge ofthe composition, crystallinity, and solubilityof the solubility-controlling solid,along with the steady-state concentrationand composition of the solutionspecies. Steady state is assumed to beestablished when the actinide concentrationremains stable for several weeksor longer. But the solubility-controllingsolids of the actinides generally precipitateas an amorphous phase becauseradiolysis affects their crystallization.Actinide solids may become less solublewith time as they transform from theirinitially formed, disordered structuresinto ordered crystalline solids, therebylowering their free energy. Such achange, however, may not appear forseveral years. Thus, the solids formedin laboratory experiments may notrepresent the most thermodynamicallystable solids with the lowest solubility.Laboratory-based solubility studiestherefore provide an upper limit toactinide concentrations in a potentialrelease scenario from a nuclear wasterepository.Because of their omnipresence,hydroxide ions (OH – ) and carbonateions (CO 3 2– ) are the most importantligands that complex to actinides in theenvironment. While other ligands, suchas phosphate, sulfate, and fluoride, canlower the actinide concentrationSolubility and Speciation ParametersSolubility products (K sp ) of actinide-bearing solid phases are indispensable primary datafor predicting the concentration limits of actinides in the environment. The solubilityproduct describes the dissolution reaction of the solid into its ionic components. Forexample, the solubility product of AmOHCO 3 (s), where (s) refers to the solid phase, isgiven by the product of the concentration of the dissolved ions Am 3+ , OH – , and CO 2–3according to the following reaction:3+ –AmOHCO 3 (s) Am + OH + CO 2–3 .The solubility product is then given by3+ –K sp = [Am ][OH ][CO 2– 3 ] ,where the brackets indicate ion concentration. The formation of an actinide complex insolution, for example,3+Am + nCO 2–3 Am(CO 3 ) 3–2nnis described by the complex’s formation constant, β n . For the forward reaction inEquation (3), the formation constant isβ n=⎡⎤⎣⎢Am(CO 3 ) 3–2nn ⎦⎥[ ][ ],nAm 3+ CO 2–3where n refers to the number of ligands. For the reverse reaction, β n changes sign andis known as the stability constant. The concentration of an actinide solution complexcan be calculated by using the known solubility product of the solubility-limiting solidand the formation constant of the species of interest as follows:⎡Am(CO ⎤⎣⎢ 3 ) 3–2nnAccordingly, the total actinide concentration in solution is given by the sum of the concentrationsof species present in solution as follows:[ ] [ ]⎡3+= βnAm CO 3 2– n = βn K sp⎦⎥ [ ][ ]3+Am(III) = Am + Am(CO ⎤ +⎣⎢ 3 ) 3–2nn Am(OH) 3–mmK sp ⎛n⎞=× ⎜1+ n[ ] + m[⎟ .⎝∑ β CO 2–3 ∑ β OH – m]CO 2–nm⎠3 OH –[ ][ ][ ][ ]⎦⎥ [ ]2–n–1CO3 .–OHComplexation of the dissolved actinide ion with ligands in solution (large β n values)generally increases the total actinide concentration in solution.(1)(2)(3)(4)(5)(6)398 Los Alamos Science Number 26 2000

Chemical Interactions of Actinides in the Environment(because of the low solubility of thecorresponding solid phase), theirconcentrations in natural waters aregenerally low. Consequently, theycannot compete successfully withhydroxide or carbonate complexation.However, organic biodegradation products,such as humates or fulvates, aregenerally present in natural aquifers andcan potentially play a role in actinidesolubility and migration.Hydrolysis. The interaction ofactinide and hydroxide ions producesmonomeric and—particularly forPu(IV)—polymeric solution species.The solid-phase structures are oxides,oxyhydroxides, or hydroxides of lowsolubility.In the absence of carbonate orother strong ligands (or with very lowconcentrations of them), trivalentactinides form the positively chargedor neutral hydroxo complexes ofgeneral formula An(OH) m 3–m , wherem = 1, 2, or 3, that is An(OH) 2+ ,An(OH) 2 + , and An(OH) 3 (aq). Similarly,An(IV) forms complexes in solutionof general formula An(OH) n 4–n , wheren = 1, 2, 3, or 4. The solid hydroxideAn(OH) n (s), where n = 3 for An(III)and 4 for An(IV) or the hydratedoxide AnO 2 •nH 2 O for An(IV) areexpected to control the solubility ofactinides in both oxidation states.Understanding the An(IV) solutionchemistry is still a formidable challenge.The high charge allows the An 4+ions to easily hydrolyze and form multiplespecies simultaneously, even atlow pH. Plutonium(IV) is known tohydrolyze even in dilute acidicsolutions and, at concentrations greaterthan 10 –6 M, will undergo oligomerizationand polymerization to form verystable intrinsic colloids.Recent x-ray absorption fine-structure(XAFS) studies, discussed on page 432in the article by Conradson, have givenmore insight into the structural detailsof Pu(IV) colloids. The sub-micronscalecolloid particles are nominallycomposed of Pu(IV) hydroxo polymersand contain structural characteristics ofplutonium oxide, PuO 2 . However, theXAFS data reveal that these particlespossess multiple functional groups,such as terminal hydroxo and bridgingoxo and hydroxo groups. Each colloidparticle can be considered an amorphousoxyhydroxide compound. Thefunctional groups originate from thepolymerization of monomeric speciesformed through hydrolysis, and thenumber of functional groups dependson the history of the solution preparationand aging process. While thechemical “structure” of every Pu(IV)colloid may be similar, each colloid canhave a unique distribution of functionalgroups depending on solution preparation,particle size, and water content.Once they form, however, the colloidsare inordinately stable in near-neutralsolution, presumably because over time,the hydroxo groups that join the plutoniumatoms together in the polymer areconverted into chemically resistant oxygenbridges.The generation of Pu(IV) colloids isone of the main processes that complicatedetermining thermodynamicconstants and accurately predictingsoluble plutonium concentrations innatural waters. Both the colloids andother amorphous Pu(IV) hydroxidesolids span a wide range of structuralfeatures, crystallinities, and stabilities,and hence exhibit a wide range ofsolubilities. The presence of Pu(IV)colloids will therefore complicate anycalculation of plutonium solubility, aswill be evident when we reviewplutonium solubility and speciation inYucca Mountain waters.Solid hydroxides and oxides are alsoexpected to limit the solubility ofAn(V) and An(VI) complexes. Forexample, Np 2 O 5 (s) has been determinedto be the solubility-controlling solid inwaters from Yucca Mountain, whileschoepite, UO 2 (OH) 2 •nH 2 O(s) is moststable for U(VI). Interestingly, thesesolids exhibit both acidic and basiccharacteristics. With increasing pH,their solubilities tend to decrease inacidic media and increase in alkaline(basic) solutions because of the formationof anionic hydroxo complexes:AnO 2 (OH) 2 – for An(V) andAnO 2 (OH) n 2–n (n = 3 and 4) forAn(VI). The hydrolysis of U(VI) iscomplicated even more by theformation of polymeric hydroxidespecies at high U(VI) concentration(>10 –5 M), i.e., (UO 2 ) 3 (OH) 5 + or(UO 2 ) 4 (OH) 7 + . We are currentlyconducting experiments to see if Pu(VI)undergoes a similar polymerization.Carbonate Complexation. Carbonateis ubiquitous in nature. In surfacewaters, such as oceans, lakes, orstreams, its concentration ranges generallybetween 10 –5 and 10 –3 M. Ingroundwaters, its concentration isenhanced (up to 10 –2 M) because of theincreased CO 2 partial pressure, which isabout 10 –2 atm deep undergroundcompared with the atmospheric partialpressure of about 3 × 10 –4 atm. Theserelatively high carbonate concentrationsinfluence the environmental chemistryof actinides in all oxidation states.Carbonate generally bonds toactinides in a bidentate fashion, that is,two of the three oxygen atoms in thecomplex bond to the actinide. Thus,actinide carbonate complexes are verystrong and highly stable both insolution and in the solid state.We have investigated the structureof carbonate complexes involving theactinyl ions AnO 2 + and AnO 2 2+ . (Thesecomplexes are more soluble and thusmore easily studied.) As shown inFigure 4, three monomeric solutioncomplexes form. Depending on thesolution conditions, these complexesmay coexist with each other and withhydroxo complexes. It is therefore difficultto obtain clean, unambiguous structuraldata about a single species.We can make inroads into that problemby investigating the stability andstructure of the triscarbonato species,AnO 2 (CO 3 ) 3 m-6 , shown in Figure 4(c).This is the limiting (fully coordinated)An(V,VI) carbonate complex in solution.At carbonated concentrationsabove 10 –3 M, this complex forms atthe exclusion of the others. Using aNumber 26 2000 Los Alamos Science 399

Chemical Interactions of Actinides in the Environmentbattery of spectroscopic techniques,including multinuclear NMR spectroscopy,x-ray diffraction, and XAFS,we have obtained formation constants,bond lengths, and other structuralparameters for the An(V) and (VI) carbonatecomplexes (Clark et al. 1999).This information has helped us understandvarious aspects of the An(V,VI)carbonate system. (The article by Clarkthat begins on page 310 describes structuralinvestigations of the Pu(VI)tris-carbonate.)Carbonate Solids. In many naturalwaters, carbonate effectively competeswith hydrolysis reactions, and mixedhydroxocarbonate solids are likely toform. These solids have general formulasAn(OH)(CO 3 )(s) for An(III) andAn(OH) 2m (CO 3 ) n (s) for An(IV), whereAn stands for uranium, neptunium, plutonium,or americium. (The equilibriain these systems are quite complicated,and neither the structure nor the valuesfor n has been determined accuratelyfor An(IV) compounds.)In the laboratory, however, we canstudy actinide solids of differentcompositions by increasing the carbonateconcentration. This increase reducesthe extent of hydrolysis and allowsthe formation, for example, of pureAn(III) carbonate solids, such asAm 2 (CO 3 ) 3 •nH 2 O(s) orNaAm(CO 3 ) 2 •nH 2 O(s), which arethe solubility-controlling solids in solutionscontaining high concentrationsof sodium and carbonate. These solidsare analogous to the ones formedby the trivalent lanthanides Nd(III)and Eu(III).Similarly, we have observed theformation of solid An(VI) and An(V)compounds in solutions containing highconcentrations of alkali metals andcarbonate. These An(V,VI) carbonatesolids are inherently interesting.Together, they can be viewed as astructural series, as evident in Figure 5.Binary An(VI) solids have a generalcomposition AnO 2 CO 3 (s) and are madefrom highly ordered actinyl carbonatelayers. However, An(V) solids need to(a) AnO 2 (CO 3 ) m–2(d) (UO 2 ) 3 (CO 3 ) 66–TerminalcarbonateBridgingcarbonateFigure 4. Carbonate Complexes of the Actinyl IonsActinides in the V and VI state form similar carbonate complexes in solution, althoughthe charge of the species depends on the oxidation state. Thus, in the three monomericcomplexes shown in (a)–(c), m = 1 for An(V) and m = 2 for An(VI). Actinides in theV state have the lowest effective charge and form the weakest complexes. In thesecomplexes, Np(V) has the longest actinyl bond (An=O) length (1.86 Å). We have alsodetermined only slight changes in the equatorial An–O distances to the carbonateligands: they vary between 2.42 and 2.44 Å for An(VI) and between 2.48 and 2.53 Å forNp(V). (d) Uranium(VI) may also polymerize at high carbonate and uranium concentrationsto form the polynuclear actinyl carbonate complex (UO 2 ) 3 (CO 3 ) 6– 6 . If present,this polynuclear complex stabilizes U(VI) in solution and increases the overall U(VI)solubility in comparison with the solubility of its monomeric counterpart, AnO 2 (CO 3 ) 2– 2 .incorporate a monovalent cation, suchas Na + , K + , or NH 4 + , into their structure(otherwise they would have anoverall charge). Thus, we observe onlyternary An(V) solids with general compositionsM (2n–1) AnO 2 (CO 3 ) n •mH 2 O(s),where n = 1 or 2, M is the monovalentcation, and An stands for neptunium,plutonium, or americium.As the number of cations incorporatedinto the An(V,VI) carbonate solids increases(stoichiometrically from 0 to 1to 3), the structures become increasinglyopen, and the solids generallybecome more soluble under conditionsin which the uncomplexed actinide iondominates solution speciation. Replacementof the cations with larger ones(b) AnO 2 (CO 3 ) 2m–4(c) AnO 2 (CO 3 ) 3m–6OxygenCarbonActinideenhances the solubility of the actinideby several orders of magnitude. Weexploited that fact when we studied theneptunyl “double carbonate” solidsM 3 NpO 2 (CO 3 ) 2 . These solids exhibitsuch low solubilities in near neutralsolutions (10 –4 to 10 –7 M) that thesolution species cannot be studied bymany spectroscopic methods. To modifythe solubility and obtain moreconcentratedneptunyl solutions, weexamined the use of large, bulkycations in an attempt to generate anunfavorable fit of the complex into thestable lattices shown in Figures 5(b)and 5(c). We found that tetrabutylammoniumprovided stable solutions ofNpO 2 CO 3 – and NpO 2 (CO 3 ) 2 3– in400 Los Alamos Science Number 26 2000

Chemical Interactions of Actinides in the EnvironmentUO 2 CO 3 KPuO 2 CO 3 K 3 NpO 2 (CO 3 ) 2 •nH 2 OFigure 5. Actinide(V,VI) Alkali Metal/Carbonate SolidsActinide(V,VI) carbonate solids form a “structural series” that progresses from dense, highly ordered layered structures to moreopen structures. A single layer is shown above a perspective drawing of each solid. (a) An(VI) will combine with carbonate to formbinary solids, such as UO 2 CO 3 (s) or PuO 2 CO 3 (s). Within a layer, each actinide is bonded to two carbonates in a bidentate fashionand to two others in a monodentate fashion. (This is in contrast to solution species, in which carbonate always bonds to an An(VI)in a bidentate fashion.) (b) Actinide(V) ions cannot form binary carbonate solids, since the structural unit would have an overallcharge. Solid An(V) carbonates therefore form ternary complexes in conjunction with monovalent metal cations (M), such as Na + ,K + , or NH + 4 . Within a layer, each actinyl ion binds to three carbonate ligands in a bidentate fashion. The metal ions sit between thelayers and form the ternary complex MAnO 2 CO 3 (s). This opens the structure, in comparison with the An(VI) binary structure, andincreases solubility. (c) At higher metal ion concentrations, metal ions will be incorporated into each layer to form very openstructures of general formula M 3 AnO 2 (CO 3 ) 2 (s).millimolar concentrations over a widepH range. The enhanced solubilityallowed us to study the temperaturedependence of the carbonate complexationreactions with near-infrared (NIR)absorption spectroscopy and to obtainadditional structural information aboutthe Np(V) solution carbonate complexeswith XAFS spectroscopy.The compositions and stabilities ofthese actinyl carbonate solids dependon the ionic strength and carbonateconcentration and are controlled by theactinide, carbonate, and sodium concentrationsin solution. For example, thestability of the Np(V) solidsNa (2n–1) NpO 2 (CO 3 ) n •mH 2 O(s) is givenby the solubility product K sp :K sp = [Na + ] 2n-1 [NpO 2 + ] [CO 3 2– ] n .For the An(VI) solid AnO 2 CO 3 (s),n is equal to 1 and its solubility isnominally independant from the sodiumconcentration in solution. At the lowionic strengths of most natural waters,an unrealistically high NpO 2 + concentrationwould be needed to precipitate aternary Np(V) carbonate solid. Thus,solid An(V) or An(VI) carbonates arein general not observed in natural waters.However, because the ternarycompounds of Np(V) have beenobserved to form in concentratedelectrolyte solutions, these solids areimportant for solubility predictions inbrines found at WIPP.Solubility and Speciation inYucca Mountain WatersAs a way of tying together many ofthe concepts presented in the previoussections, we can consider the solubilityand speciation of neptunium and plutoniumin Yucca Mountain waters. (Thearticle by Eckhardt, starting on page410, discusses more of the researchconducted for Yucca Mountain.) Asmentioned earlier, neptunium tends toexist in the V oxidation state in naturalwaters and because of the lower effectivecharge of the AnO 2 + ion, formsweak complexes. Thus, neptunium hasthe distinction of being the most solubleand transportable actinide and is theNumber 26 2000 Los Alamos Science 401

Chemical Interactions of Actinides in the EnvironmentSpecies concentration (mol kg −1 )10 −210 −4J-13 waterNp undersaturatedFigure 6. Solubility of Neptunium(V) and Pu(IV) J-13 WaterThe figure shows our measured total soluble concentrations of neptunium (red circles)and plutonium (red square) in J-13 water at a fixed carbonate concentration of 2.8 mM(indicated by the dashed line), as well as earlier measurements conducted at a higherionic strength (black circle and square). The lines are theoretical predictions for differentsolubility-controlling solids. The Np(V) concentration is about 100 times greaterthan that of Pu(IV). Our measurements confirm that the controlling Np(V) solid is theoxide Np 2 O 5 (s), rather than the previously postulated sodium carbonate solidNaNpO 2 CO 3 (s). The solubility of plutonium in J-13 water presumably is controlled byamorphous oxides/hydroxides or colloidal forms of Pu(IV), since the measured Pu(IV)solubilities are nearly an order of magnitude higher than predicted by Lemire andTremaine for the solubility product of Pu(OH) 4 (s). Our data are consistent with therange of Pu(IV) solubilities as calculated from K sp data of Pu(IV) hydroxide or hydrousoxide reported in the open literature (gray area).one of greatest concern when assessingthe environmental safety of the potentialunderground storage site at YuccaMountain. Plutonium is considered themost toxic actinide and is always ofenvironmental concern.Our studies were conducted in waterfrom a particular borehole in YuccaMountain (J-13 water). The sodium andcarbonate concentrations are low (botharound 2 to 3 mM), as is the ionicstrength of J-13 water (0.003 molal).NaNpO 2 CO 3 •nH 2 O(s)0.01 0.1 1.0Ionic strength (mol kg −1 )Np (Nitsche et al. 1993)Np 2 O 5 •nH 2 O(s)Np oversaturated10 −6Pu (Nitsche et al. 1993)Pu oversaturated10 −8Pu(OH) 4 (s) (Lemire and Tremaine 1980)10 −1010 −12However, groundwater conditions withinand around Yucca Mountain vary.The carbonate concentration changes(because the CO 2 partial pressurechanges with depth), as does the pH,Eh, and temperature. We mimicked thisvariation in the natural waters by conductingour studies at various pH valueswhile maintaining a constant ionicstrength. We also conducted our studiesat different temperatures (25˚C, 60˚C,and 90˚C) to approximate the range oftemperatures expected at the repository.Not surprisingly, changing these basicparameters greatly affected the actinidebehavior (Efurd et al. 1998). Figure 6summarizes the results of our studies.The controlling solids for theactinides in J-13 water—in all oxidationstates—were mainly oxides and hydroxides.We observed only the formationof greenish brown Np(V) crystallinesolids of general formula Np 2 O 5 •nH 2 O(s),green Pu(IV) solids of general formulaPuO 2 •nH 2 O and/or amorphous Pu(IV)hydroxide/colloids. The solubility ofneptunium was about 10 –3 M at pH 6,but dropped to about 5 × 10 –6 M at pH8.5. The neptunium species distributionsimilarly changed with conditions. AtpH 8.5, about 31 percent of Np(V)was present as the neptunyl ion NpO 2 + .The remaining soluble Np(V) was complexedwith either carbonate or hydroxide,with about 58 percent NpO 2 CO 3–and 11 percent NpO 2 OH(aq). At pH 6,however, Np(V) hydrolysis and carbonatecomplexation reactions were minimized,and all of the neptunium waspresent as NpO 2 + .By comparison, at all pH valuesexamined in our study, 99 percent ofthe soluble plutonium was calculated tobe in the IV state as Pu(OH) 4 (aq).When using Pu(OH) 4 (s) as the solubility-limiting solid in our geochemicalmodels, the plutonium concentration insolution was predicted to be about10 –9 M (using the thermodynamic dataavailable in the open literature). Asseen in Figure 6, experimental data areabout one to two orders of magnitudehigher. The discrepancy presumablyarises because plutonium solubility isinstead controlled by amorphousoxides/hydroxides or colloidal forms ofPu(IV). (Note that if crystalline PuO 2 (s)were the solubility-controlling solid,Pu(IV) solubility would be about 10 –17M, well below the Pu(IV) hydroxidesolubility range.)The J-13 water is almost neutral(pH ~7) and has a redox potential ofabout +430 millivolts. But differentwater conditions would dramaticallyaffect neptunium solubility, as illustrated402 Los Alamos Science Number 26 2000

Chemical Interactions of Actinides in the Environment(a)10 −2 1400 mVNp 2O 5(s)(b)10 −4pH 7Np 2O 5(s)Np concentration (mol kg −1 )10 −410 −610 −8300 mV200 mV100 mV0 mV−100 mVNp concentration (mol kg −1 )10 −610 −8 Eh (mV)Np(OH) 4(s)Np(OH) 4(s)−200 mV10 −104 6 8 10 12−400 −200 0 200 400 600pHFigure 7. Neptunium Solubility in Yucca Mountain J-13 WaterNeptunium solubility varies tremendously depending on the Eh and pH of the water. (a) The graph shows the neptunium concentrationas a function of pH. Calculated solubility boundaries at various Eh values are superimposed on the graph. The colored regioncontains mainly Np(V) in solution, while the upper and lower bold lines are the solubility boundaries of the Np(V) solid phase,Np 2 O 5 (s), and the Np(IV) solid phase, Np(OH) 4 (s), respectively. Solutions with neptunium concentrations above these boundaries areoversaturated and will precipitate the respective solid. The total neptunium concentration in solution may vary by more than sixorders of magnitude over the entire pH range under oxidizing conditions. (b) This graph shows the solubility as a function of Eh ata fixed pH of 7, and corresponds to the Np solubility as one moves along the red line in the previous graph. Under reducing conditions(below –120 mV), Np(IV) dominates the solution speciation. Its concentration in solution is limited to less than 10 –8 molal bythe solubility of Np(OH) 4 (s). As conditions become more oxidizing, Np(V) becomes the redox-stable oxidation state and begins todominate the speciation. [The Eh determines the Np(IV)/Np(V) ratio in solution, while the overall Np(IV) concentration is maintainedby the solubility of Np(OH) 4 (s).] The neptunium concentration rises with redox potential until it reaches approximately 10 –4 molal atan Eh of ~250 mV. Solutions containing a higher neptunium concentration are oversaturated, and the Np(V) solid phase Np 2 O 5 (s)precipitates. The soluble neptunium concentration remains constant from then on. The shaded area spans the region where Eh andsolubility of the Np(IV) solid control the neptunium concentration in solution. Outside this area, the solubilities of the Np(IV) andNp(V) solids govern neptunium solution concentration, independent of Figure 7 (Kaszuba and Runde 1999).Under strong reducing conditions(below about –120 millivolts), the totalsoluble neptunium concentration islimited by a Np(IV) solid phase,Np(OH) 4(s), and is conservativelyassumed to be about 10 –8 M. Similar toplutonium, the presence of a moreordered solid oxyhydroxide or oxide,such as NpO 2 (s), would reduce thatconservative boundary by several ordersof magnitude. Under these reducingconditions, Np(IV) also dominates thesolution speciation in the form of theneutral complex Np(OH) 4 (aq). The totalneptunium concentration will rise as thewater conditions become more oxidizingand Np(V) becomes more stable insolution. Eventually, the Np(V) concentrationis great enough to allow aNp(V) solid phase, Np 2 O 5 (s), to precipitate.Formation of this solid sets anupper limit to the soluble neptuniumconcentration. Within the pH and Ehranges of most natural aquifers(between 4 and 9 and between –200and +700 millivolts, respectively), thatconcentration may vary by more thanfour orders of magnitude.When modeling actinide releasefrom a radioactive waste repository, wemust anticipate the conditions at thesource (e.g., close to a spent nuclearfuel container) as well as those in thefar-field environment. Spent nuclearfuel contains mainly UO 2 and actinidesin low oxidation states (III and IV). Ifconditions at the disposal site arereducing—for example, because ofmicrobial activities or the reductioncapacity of corroding iron from storagecontainers—the actinides will be maintainedin their lower oxidation statesand soluble actinide concentrations maybe low. If local reducing conditionschange because oxidizing waters infiltratefrom the surroundings, actinideswill enter their more-soluble oxidationstates and begin to migrate into theenvironment. Once far from the repository,conditions may change drastically,Number 26 2000 Los Alamos Science 403

Chemical Interactions of Actinides in the Environment(a)10Fourier Transform magnitude(b)Intensity (arbitrary units)50–5–100 1 2 3 4R–φ (Å)8201-M HClO 4PuO 22+Pu=O Pu–O eqPu–Cl830and the transport behavior of redoxsensitiveelements such as plutonium orneptunium may change. Clearly, evaluationof actinide mobility in the geochemicalenvironment around nuclearwaste repositories must account forsoluble actinide concentrations that arecontrolled by the redox potential of thetransport medium and as well by thesolubility of the actinides’ solid phases.1-M NaClPuO 2 Cl + PuO 2 Cl 2PuO 2 Cl 3–840850Wavelength (nm)4-M NaCl15-M LiCl[Cl](M)1415Pu=O(Å)1.741.751.75PuO 2 Cl 42–860Pu–O eq(Å)2.402.432.49Pu–Cl(Å)—2.752.70Figure 8. XAFS Data and NIR Spectra of Pu(VI) in NaCl Solutions(a) By comparing XAFS data for Pu(VI) in 1-M NaCl, 4-M NaCl, and 15-M LiCl, and noncomplexing1-M HClO 4 , we have obtained condition-specific structural informationabout the solution species. Chloride appears to bond above 1-M chloride in solution.Bonding is clearly evident by 4-M chloride. As seen in the table (inset), the plutonylbond length (Pu=O) increases slightly after the chlorine bonds and replaces watermolecules (indicated by the Pu–O eq bond) in the equatorial plane of the O=Pu=O moiety.(b) The absorption bands at 838, 843, 849, and about 855 nm, plus the characteristicband at 830 nm for the PuO 2+ 2 ion, indicate the formation of inner-sphere Pu(VI) chloridecomplexes. These bands allowed us to determine the formation constant, β, of themost important Pu(VI) chloride complexes in NaCl solutions, PuO 2 Cl + and PuO 2 Cl 2 (aq).870Complexation with Chlorideand Implications for WIPPIn most natural waters, actinides areusually associated with hydroxide orcarbonate ligands. However, WIPP,which is designed to hold defenserelatedtransuranic waste, is carved outof a salt formation. Water at the WIPPsite thus consists of concentrated brinesthat are saturated with chloride saltsand exhibit very high ionic strengths(between 4 and 8 molal). With increasingdistance from the repository, thebrines’ ionic strength decreases andmay ultimately reach the lower valuesof most natural surface and groundwaters.This variation complicates ourtransport models, since the species thatform within the WIPP site will not bethose that control actinide behaviorfar from the site.Chloride has been shown to affectthe solubility and speciation ofactinide compounds significantly. Asmentioned earlier, in most naturalwaters plutonium favors the IV and Voxidation states, but radiolytichypochlorite formation in concentratedchloride solution may result in theformation of Pu(VI). Furthermore,complexation by chloride may enhancestabilization of Pu(VI) with respect toreduction to Pu(V) and Pu(IV).We have studied the complexationreactions of U(VI) and Pu(VI) withchloride using a variety of spectroscopictechniques, including ultraviolet,visible, and NIR absorption and XAFSspectroscopies (Runde et al. 1999). Wecould determine when U(VI) andPu(VI) chloride species, such asPuO 2 Cl + and PuO 2 Cl 2 , formed as afunction of chloride concentrations (seeFigure 8). Interestingly, the tris- andtetrachloro complexes are not formed insignificant concentrations in NaClsolutions (up to the point of NaClsaturation). However, we have observedspectroscopically UO 2 Cl 4 2– andPuO 2 Cl 4 2– at very high LiClconcentrations and obtained structuraldata from single crystals synthesizedfrom concentrated chloride solutions.The interaction of actinides withchloride and other weak anionic ligandssignificantly affects solubility predictionsand must be considered in riskassessments of WIPP and other nuclearwaste repositories in salt formations.With regard to neptunium, chlorideenhances the Np(V) solubility becauseof a strong ion interaction between theNp(V) compounds and chloride ions404 Los Alamos Science Number 26 2000

Chemical Interactions of Actinides in the Environment(see Figure 9). Runde et al. (1996)simulated this interaction using thePitzer ion-interaction model and developeda Pitzer parameter set to predictthe Np(V) solubility at different NaClconcentrations and in artificial brinesolutions (Runde et al. 1999). Neglectingthe chloride interaction leads to apredicted Np(V) solubility about anorder of magnitude too low at a neutralpH. The change of water activity withionic strength affects the chemicalpotential of the solid phase because ofhydration water molecules in the solid.Np(V) solubilities were measuredfrom oversaturation in two artificialWIPP brines, AISinR and H-17. Thesolutions reached a steady state after241 days. The measured Np(V) solubilitiesreported in the literature are2 × 10 –6 molal and 1 × 10 –6 molal inAISinR and H-17, respectively, whileour model predicts Np(V) solubilities of1 × 10 –5 and 3 × 10 –6 molal, respectively.The prediction is acceptable forthe H-17 brine but lacks accuracy forthe AISinR brine. This is because themodel accounts only for interactionsbetween Np(V) and Na + and Cl – ions,while the Pitzer parameters forelectrolyte ions such as K + , Mg 2+ ,Ca 2+ , and SO 4 2– (other constituents inthe brines) are unknown. Additionaldata in various electrolyte solutions arerequired to develop a more accuratemodel for the Np(V) solution chemistryin complex brines.SorptionThe solubility of an actinide speciesprovides an upper limit to the actinideconcentration in solution and may beconsidered as the first barrier toactinide transport. Once an actinidehas dissolved in water, its sorptiononto the surrounding rock or mineralsurfaces is expected to act as asecondary barrier. When the actinideconcentrations are far below thesolubility limits, as is likely in mostsurface and groundwater systems,water/rock interfacial processes areAn(V) (mol kg −1 ) Spatial distribution (%)10080604020NpO 2+10 −3 5-M NaCl5-M NaClO 410 −4NaNpO 2 CO 3 •3.5H 2 O(s)10 −510 −6Na 3 NpO 2 (CO 3 ) 2 •nH 2 O(s)10 −710 −7 10 −6 10 −5 10 −4 10 −3 10 −2 10 −1NpO 2 (CO 3 ) −NpO 2 (CO 3 ) 23−Carbonate concentration (mol kg −1 )NpO 2 (CO 3 ) 35−Figure 9. Solubility of Np(V) in Concentrated Electrolyte SolutionsThe upper graph shows the distribution of predominant Np(V) solution species as afunction of carbonate concentrations, with environmentally relevant concentrationsshown in gray. The lower graph shows the soluble Np(V) concentrations in 5-M NaClsolution—conditions that are relevant for WIPP—and noncomplexing 5-M NaClO 4solution over the same range. Similar data are obtained for Am(V). The neptuniumsolubility is about an order of magnitude higher in NaCl because of the stabilizingeffect of actinide-chloride complexation reactions. The solubility-controlling solid forboth curves at low carbonate is the hydrated ternary Np(V) carbonate,NaNpO 2 CO 3 •3.5H 2 O(s). Interestingly, this solid initially forms at high carbonateconcentrations. While it is kinetically favored to precipitate, it is thermodynamicallyunstable. Over time, it transforms to the thermodynamically stable equilibrium phaseNa 3 NpO 2 (CO 3 ) 2 •nH 2 O(s) (dashed curve).expected to dominate actinide transportthrough the environment. (Simple dilutionwill lower the actinide concentration,typically to trace concentrations,once the actinide migrates away fromthe contaminant or waste source.)Actinide sorption onto mineralsurfaces depends on many factors, suchas the actinide’s speciation and concentration,the flow rate and chemistry ofthe water, and the composition of thesurrounding rocks. In addition, differentmechanisms, such as physical adsorption,ion exchange, chemisorption, oreven surface precipitation, can fix theactinide to a mineral/rock surface or ingeneral inhibit its transport. Furthermore,redox processes complicate theNumber 26 2000 Los Alamos Science 405

Chemical Interactions of Actinides in the EnvironmentAdsorbed Pu(V) (%)1201008060402000 40 80 120Time (h)Figure 10. Uptake of Pu(V) by ColloidsLu et al. (1998) obtained data for Pu(V) adsorption onto colloids, including hematite(Fe 2 O 3 ), silica (SiO 2 ), and montmorillonite (Mont), in natural J-13 water and in asynthentic (syn) J-13 water. The colloid concentration was maintained at 200 mg/L,with the colloids being about 100 nm is size. The sorption clearly depends on the typeof mineral, but the sorption mechanism remains unknown. The time-dependence ofthe sorption onto montmorillonite suggests a more complicated interaction thanoccurs with oxide minerals.Table II. Desorption of Plutonium from Colloids in J-13Water after 150 DaysColloid Materialsorption behavior of neptunium andplutonium. Unfortunately, we are onlybeginning to understand such actinidesorption and transport mechanisms.Generally, batch-sorption experimentsare performed to investigate howactinides will be retained in the environment.An actinide-containing solutionis allowed to come in contact withthe geomaterial of interest, and wemeasure the amount of actinide thatFe 2 O 3 (syn)Fe 2 O 3SiO 2MontSiO 2 (syn)Mont (syn)160 200 240Amount DesorbedPu(IV)Pu(V)(%) (%)Hematite (α-Fe 2 O 3 ) 0.0 0.01Goethite (α-FeOOH) 0.57 0.77Montmorillonite 8.23 0.48Silica (SiO 2 ) 19.93 1.04remains in solution. The difference betweenthe remaining and initial actinideconcentrations equals the amount thatsorbed to the sample. This differenceallows us to calculate the batch-sorptiondistribution coefficient K d , defined asthe moles of radionuclide per gram ofsolid phase divided by the moles of radionuclideper milliliter of solution andreported as milliliters per gram (mL/g).Low K d values mean low sorption, andhigh values imply effective immobilizationdue to sorption onto rock surfaces.Batch-sorption experiments providehelpful information about the sorptioncapacity of a mineral or soil. But whilethey are easy to perform, they havesome drawbacks. The experiments arefundamentally static and provide onlypartial information about the dynamic,nonequilibrium processes that occur innature. 2 Additionally, the results ofbatch-sorption experiments are verysample specific and difficult to extrapolateto the different conditions foundthroughout a repository.To apply sorption data accurately,we must obtain sorption coefficientsfrom both individual, well-characterizedminerals, such as zeolites, manganeseoxides, or iron oxyhydroxides, andfrom mineral mixtures with varyingbut known compositions. Sorptioncoefficients from these studies canserve as input data into a transportmodel that will estimate the actinideretention along water flow paths ofknown mineral content. The modelpredictions can then be compared withexperimental results obtained fromdynamic experiments that use complexsoils and rocks.Studies using minerals (hematite,calcite, gibbsite, albite, and quartz),clays (montmorillonite), and zeolites(clinoptilolite) have shown that Am(III)and Pu(IV) sorb very rapidly to thesesolids in the following descendingorder: hematite, montmorillonite,clinoptilolite, and calcite, followed to amuch lesser extent by gibbsite, albite,and quartz. But other factors, such asthe concentration, solution speciation,and oxidation-state distribution of theactinide species, significantly affect theextent of surface sorption. As an example,K dvalues for neptunium sorptiononto hematite range from 100 to2 Flow-through, or dynamic, column experiments,wherein the actinide solution is allowed tomigrate through a short column of packedgeomaterial, provide information about dynamictransport properties. Dynamic column experimentsare difficult to perform and interpret, andmost of the sorption studies on the actinides arestill performed as batch experiments.406 Los Alamos Science Number 26 2000

Chemical Interactions of Actinides in the EnvironmentCaPCaPUCaCaOO0 1 2 3 4 5keV0 1 2 3 4 5keVFigure 11. SEM Micrograph of Uranium Precipitate on ApatiteA solution containing U(VI) at 10 –4 M was kept in contact with apatite for 48 hours. The central SEM micrograph (1100 × magnification)shows crystalline plates on the smooth apatite mineral surface. Analysis of these plates by energy-dispersive spectroscopy(right graph) indicated the presence of calcium, oxygen, uranium, and phosphorus, consistent with crystalline uranium phosphate.Similar analysis of the bare regions of apatite (left graph) showed signals corresponding to only calcium, oxygen, and phosphorus.The precipitates were observed at uranium concentrations above 10 –5 M. Studies show that at lower concentrations, uranium sorbs,rather than precipitates, to the mineral surface.2000 mL/g, while plutonium sorptiononto hematite results in K dvaluesabove 10,000 mL/g. This difference ismainly caused by the different stabilitiesof the actinide oxidation states andtheir complexation strengths. The lowcharge of the NpO 2 + ion results in alow sorption of neptunium, whilePu(IV) dominates the complexation andsorption processes of plutonium. Plutonium(V)is expected to follow the lowsorption affinity of Np(V).Lu et al. (1998) determined thesorption of Pu(V) onto oxide and claycolloids (Figure 10). Sorption onto ironoxides, such as hematite, is strong andirreversible, but only 50 percent of thePu(V) is sorbed on the montmorillonite(clay). However, the sorption mechanismsare largely unknown. The timedependenceof the sorption onto theclay suggests a more complicated interactionthan occurs with oxide minerals.Interestingly, the desorption ofPu(V) does not follow this general pattern.Lu et. al. (1998) also discussed thedesorption of Pu(IV) and Pu(V) fromvarious colloids. Table II summarizestheir findings: Surprisingly, Pu(IV)exhibits a greater tendency to desorbthan Pu(V). Particularly striking is therelatively high desorption of Pu(IV)from silica, which indicates a differentsorption mechanism or differentbinding to the silica functional groupsthan to the Fe(III) minerals, includingpossible changes in plutoniumoxidation states.These investigations indicate theimportance of understanding the oxidation-statestability of the actinide sorbedon a mineral surface. Redox-activeminerals, such as manganese or ironoxide/hydroxides, are present in subsurfaceenvironments and may alter theredox states of uranium, neptunium,and plutonium. At the Rocky FlatsEnvironmental Technology Site,seasonal variations of plutonium concentrationsin settling ponds correlateastonishingly well with solublemanganese concentrations, indicating apotential release of plutonium by wayof seasonal redox cycling.It has been generally hypothesizedthat the environmentally mobile Pu(V)is reduced to Pu(IV) through theformation of strong mineral-surfacecomplexes. Quite surprisingly, Neuet al. (2000) showed that Pu(V) is notreduced when sorbed on MnO 2 . X-rayabsorption near-edge structure indicatedthe presence of Pu(V) on MnO 2 , eventhough the starting material of thebatch-sorption experiment was colloidalPu(IV) hydroxide. Their data alsoindicated that the plutonium wasoxidized to Pu(VI). The only time theydid not observe the higher-oxidationstatesurface species was when Pu(IV)was complexed with a strong ligand,such as nitrilo triacetic acid (NTA).These observations strongly support thehypothesis that plutonium undergoes aredox cycling that is enhanced bysurface interactions. We need more dataabout the surface species, however, tomodel such interactions.Obviously, batch experiments alonecannot elucidate sorption mechanisms.Specifically, we have to distinguish betweensurface precipitation and interfacialmolecular reactions in order tomodel sorption behavior. In the eventNumber 26 2000 Los Alamos Science 407

Chemical Interactions of Actinides in the EnvironmentFigure 12. Correlation between An(III)Solution Speciation andSorption onto Montmorillonite(a) Even with an order-of-magnitudedifference in NaCl concentration, sorptionof the trivalent actinides Am(III) andCm(III) peaks in near-neutral waters. Atlower pH, the An 3+ ion is stable and doesnot participate in complexation reactionswith the most common ligands, hydroxideor carbonate, no matter if the ligands arein solution or on a surface. At higher pH,negatively charged An(III) species areformed that increase the species’solubility and reduce their sorptionbecause they are repelled from thenegatively charged functional groups onthe mineral surface. (b) Calculations fora weak sodium chloride solution(0.1-M NaCl and CO 2 partial pressure of0.01 atm) suggest that sorption is mainlycorrelated with the presence of cationicsolution species. The gray area denotesthe pH range of natural waters.(a)100800.1-M NaCl1.0-M NaClSorption (%)6040200(b)80604020Species distribution (%)An 3+An(OH) 2+An(CO 3 ) +An(CO 3 ) 2–An(CO 3 ) 33–4 5 6 7 8 9 10pHof a surface precipitation, the K d valuemay represent the solubility product ofthe surface precipitate but will notreflect the sorption behavior of anactinide species.For example, phosphate-containingminerals, such as apatite, are known toreduce thorium and uranium concentrationsin contaminated solutions. Sorptiondistribution coefficients for U(VI)and Th(IV) have been reported in theliterature, but there are alsoreports ofU(VI) and Th(IV) precipitates. OurSEM studies have identified microcrystallineprecipitates on the apatite whenthe initial U(VI) concentrations wereabove 10 –5 M (see Figure 11). At lowerU(VI) concentrations, no surfaceprecipitation was observed. Thus, at thehigher concentrations, the K d valuesactually reflect a solubility product,while at the lower concentrations, thedecrease of U(VI) in solution resultsfrom sorption of the actinide onto themineral surface.Complexation reactions can influencethe actinide sorption processes thatalso depend strongly on pH (Figure 12).For the trivalent actinides Am(III) andCm(III), the sorption reaches a peakunder conditions of near-neutral waters.There, the cationic An(III) solutionspecies predominate solution speciation.At lower pH, the An 3+ ion is stable anddoes not participate in bonding withhydroxide or carbonate ligands. Athigher pH, negatively charged An(III)species are formed that exhibit reducedinteractions with chemically activegroups at the mineral surface. Mineralor rock surfaces at high pH are generallynegatively charged, and apparently thenegatively charged actinide species arerepelled from the anionic surfacegroups. As such, anionic actinidespecies behave as simple anionic ligands,such as SO 4 2– , CrO 4 2– or AsO 4 3– .In addition to An(III) species, the sametrend of enhanced stability and lowsurface retention has been observed forthe An(VI) complex UO 2 (CO 3 ) 3 4– .Enhanced modeling efforts havebegun to address the fundamentalmechanism of actinide adsorption ontomineral surfaces. Janecky and Sattelberger(unpublished results) modeledthe adsorption behavior of Np(V) oncalcite by minimizing the free energyof the coordination. The structuralparameters of NpO 2 CO 3 – , which is themain solution complex in neutral andlow-carbonate waters (such as J-13water), were used to mimic the adsorption.The resulting surface-complexmodeling indicates a distance of about4 Å between the neptunium center andthe nearest calcium ions in the calcitestructure. Data from preliminary XAFSstudies support this distance. Althoughthis agreement between experiment andtheory is encouraging, more detailedstudies have to be performed to developmuch more accurate transport models,including this surface complexationmodel.We are also trying to better understandthe sorption/desorption reactionsof actinides with colloids and the408 Los Alamos Science Number 26 2000

Chemical Interactions of Actinides in the Environmentactinides’ resulting transport characteristics.This area of environmentalmigration received heightened attentionwith the discovery of plutonium in aborehole at the Nevada Test Site(Kersting et al. 1999). The plutoniumhad evidently migrated 1.3 kilometersin only 30 years. (Each nuclear test ischaracterized by a unique 239 Pu/ 240 Puisotope ratio, and that ratio can be usedto identify the plutonium source.) Asdiscussed in the article by MaureenMcGraw, we now believe that colloidaltransport was responsible for thisremarkably fast movement of plutoniumthrough the water-saturated rock. It isnot clear, however, whether thetransport was facilitated by intrinsicplutonium colloids or natural (clay orzeolite) colloids. What is clear is thattransport models to date have underestimatedthe extent of colloidal transporton plutonium mobility.Microbial InteractionsWe have discussed the most prominentchemical processes involvingactinides in the environment: dissolutionand precipitation of solids, complexationin solution, and sorption ontosurfaces. But there is growing attentionto the importance of actinide interactionswith microorganisms. Great varietiesof microorganisms exist in allaquifers and soils, and many surviveunder the most extreme conditions,such as around hydrothermal vents atthe bottom of the ocean, within deepsubterranean aquifers, or within thehypersalinity of geologic salt beds.Microorganisms can accumulate onsolid surfaces often via formation ofbiofilms, or they may be suspended inthe aquifer. As such they act as mobileor even self-propelled colloids. Theconcern is that the actinides can beassociated with microbes, eitherthrough surface binding or throughmetabolic uptake, and be transported bythem through the environment. However,the microorganisms’ role in environmentaltransport behavior of radioactivecontaminants is still virtually unknown.Microorganisms, especially bacteria,are also known to mediate redoxprocesses. They may reduce elementsfrom higher oxidation states and use theenthalpy of the redox reaction to growand reproduce. Through metabolicactivity, microbes can also establishreducing conditions by changing the pHand Eh of the local waters. While directexperimental evidence is still lacking, itis almost certain that some microbescan catalyze the transformation ofuranium, neptunium, and plutoniuminto less-soluble forms. Several bacteriaspecies have been reported in the literatureto effectively reduce the highlymobile U(VI) into the far less solubleand mobile U(IV).Given these characteristics,microbes could pose a third naturalbarrier to actinide transport fromgeologic repositories and could alsohelp remediate actinide-contaminatedsoil, groundwater, and waste streams.The ultimate goal is to use microorganismsas “living” backfill to immobilizeactinides through processes such asbioreduction, biosorption, or bioprecipitationand mineralization. At this time,however, actual applications of bacteriain remediation strategies and fielddemonstrations are rare.Recent WIPP-related studies at LosAlamos have focused on the interactionof urnaium and plutonium with thebacteria’s cell walls or byproducts, suchas extracellular biopolymers, and thetoxicity and viability of microorganismsin the presence of the actinides. Franciset al. (1998) have shown that actinidesare associated only at low levels withthe halophilic bacteria isolated frommuck pile salts around the WIPP site.The extent of association varied withbacterial culture, actinide species, pH,and the presence of organic or inorganiccomplexants. The radiolytic effects onthe microbial viability were found tobe negligible.However, other microorganisms canconcentrate metal ions to a great extent.Hersman et al. (1993) observed a stronguptake of plutonium by pseudomonassp. Within the first week of the experiment,the bacteria concentrated plutoniumto levels that were nearly 10,000times greater than a sterile control organism.Lower uptake was observed forcontact times of two weeks or longer.Another strain of this bacteria,pseudomonas mendocina, is known toremove lead from Fe(III) hydroxide/oxides and was investigated for itspotential to remove radionuclides fromthose environmentally abundant minerals.In fact, Kraemer et al. (2000) foundthat 20 percent of the sorbed plutoniumwas removed by this aerobically growingbacteria after about four days. Theexperimental conditions are presentlyoptimized to remove the majority of thesorbed plutonium and to maximize itspotential application in soil remediationstrategies.Leonard et al. (1999) also studiedthe inhibitory and toxic effects of magnesiumoxide, MgO, on these WIPP-derivedbacterial cultures. Magnesiumoxide is of specific interest because itwas chosen as backfill material at theWIPP repository to maintain the pHbetween 9 and 10, to reduce the amountof infiltrating water, and to limit theamount of CO 2 that can accumulatethrough degradation of cellulose andother organic compounds in the wastes.Initial studies indicated an inhibited cellgrowth as the amount of MgOincreased, as well as a decrease inmicrobially generated gas (N 2 O) byactive halophilic bacteria. Ultimately,the toxic effect of MgO inhibits microbialgrowth, and consequently microorganismsmay be unable to create areducing environment at the repositorythat would stabilize actinides.A more detailed study on the microbialassociation mechanism includedanalysis of the uptake of uranium andplutonium by sequestering agents andextracellular polymers (exopolymers).As an example, some exopolymers producedby bacillus licheniformis consistof repeating γ-polyglutamic acid (γ-PGA). Neu et al. (manuscript in preparation)observed actinide hydroxide precipitationat metal-to-glutamate ratiosNumber 26 2000 Los Alamos Science 409

Chemical Interactions of Actinides in the Environmenthigher than 1:20, but both U(VI) andPu(IV) ions formed soluble complexesat ratios 1:100 and 1:200, respectively.Presumably at the higher PGA concentrations,the actinide ion is encapsulatedand isolated from hydroxo-colloid formation.The binding strength for U(VI)and Pu(IV) is lower than for otherstrong ligands and chelators, such asTiron or NTA, but on the same order asfor other exopolymers or humic acids.An upper binding boundary for thepolyglutamate exopolymer was determinedto be about 0.5 micromolesPu(IV) per milligram of γ-PGA. Activecells of bacillus licheniformis bonded tomore than 90 percent of the introducedPu(IV)—up to 8.4 micromoles ofPu(IV) were added. As expected, otherfunctional groups on the exterior cellsurface (in addition to the PGA exoploymer)increase the number of bindingsites and enhance metal binding.This research clearly shows thatactinides bind with microorganisms andtheir metabolic byproducts. Dependingon the static or dynamic nature of themicroorganism, the interaction of plutoniumand other actinides with extracellularligands may affect theactinides’ mobility both away fromstorage containers and within theenvironment. Future studies will focuson understanding the mechanisms ofactinide/microbial interactions andexploiting their potential for remediationand immobilization efforts. Uptakestudies of plutonium by means ofsequestering agents are highlighted inthe box on “Siderophore-MediatedMicrobial Uptake of Plutonium” onpage 416.Concluding RemarksFrom this overview of actinide interactionsin the environment, it shouldbe clear that such interactions are complex.Each local environment is unique,and the site-specific conditions determinewhich actinide species predominateas well as each species’ overalltransport and migration characteristics.In general, actinide solubilities arelow in most natural waters—belowmicromolar concentrations. Actinidesin the V oxidation state will have thehighest solubilities; those in the IVoxidation state, the lowest. Conversely,An(V) species show low sorption ontomineral and rock surfaces, and An(IV)species tend to sorb strongly. Typically,neptunium is the only actinide thatfavors the V oxidation state, and assuch it is the most soluble and mosteasily transportable actinide under environmentalconditions. Plutonium favorsthe IV oxidation state in many naturalwaters, and normally its solubility isquite low. Once plutonium is in solution,its migration is dominated by avariety of processes, including redoxreactions in solution or on mineralsurfaces, sorption, colloidal transport,simple diffusion into rock interfaces,coprecipitation, and mineralization.Furthermore, increasing researchactivites indicated that microbes canaffect actinide migration, either throughdirect association or by helping tomaintain a reducing environment.To understand this complex set ofinteracting processes requires that wehave a strong, multidisciplinary foundationin environmental chemistry, molecularscience, and interfacial science andthe requisite spectroscopies, subsurfacetesting, and laboratory experiments. Butwe also need to interpret experimentalresults within the context of the multicomponentnatural system.A case in point is the recent identificationof a higher oxide of plutoniumof general formula PuO 2+x(s) , which isformed by water-catalyzed oxidation ofPuO 2 (s). As discussed by Haschke inthe article beginning on page 204, thisPu(IV) “superoxide” may also containPu(VI), which would increase plutoniumsolubility. In a recently publishedpaper (Haschke et al. 2000), the authorsspeculated that the Pu(VI) coulddissolve and be transported more easilyfrom a nuclear waste storage site thanpreviously estimated. Further, theyconjectured that Pu(VI) might havecaused the rapid migration of plutoniumat the Nevada Test Site. In these speculations,however, they failed to includeinfluential environmental factors.While we acknowledge that thepotential presence of Pu(VI) in aPuO 2 (s) matrix may increase thedissolution kinetics, the chemical conditionsof the local environment (waterand rock compositions) determine thestability of plutonium oxidation statesin the aqueous phase and the overallplutonium mobility. Under ambientconditions and in most aqueous environments,Pu(VI) is unlikely to bestable. Once in solution, it is typicallyreduced to lower oxidation states and,depending on the water redox potential,may eventually precipitate asPu(OH) 4 (s) or PuO 2 (s). Consequently,the fact that the superoxide PuO 2+x (s)may include Pu(VI) does not automaticallyimply enhanced plutonium solubilityor mobility in the natural environment.Furthermore, given theresults on colloidal transport, it isimprobable that Pu(VI) is responsiblefor the plutonium migration at theNevada Test Site.As the inventory of actinides ininterim storage facilities grows andstorage containers age, increasing therisk of accidental releases, and as thesafety of permanently storing nuclearwaste in geologic repositories must beevaluated, it becomes increasinglyimportant to understand how actinideswill interact with the environment.More sophisticated models are neededto account for all the potentialmigration paths away from an actinidesource. Theoretical and experimentalscientists will be challenged for yearsby the demands of developing thosemodels. 410 Los Alamos Science Number 26 2000

Chemical Interactions of Actinides in the EnvironmentFurther ReadingBrown, G. E., V. E. Heinrich, W. H. Casey,D. L. Clark, C. Eggleston, A. Felmy, D. W.Goodman, et al. 1999. Chem. Rev. 99:77–174.Clark, D. L., S. D. Conradson, D. W. Keogh,M. P. Neu, P. D. Palmer, W. Runde, et al.1999. X-Ray Absorption and DiffractionStudies of Monomeric Actinide Tetra-, Penta-,and Hexavalent Carbonato Complexes. InSpeciation, Techniques and Facilities forRadioactive Materials at Synchrotron LightSources, OECD Nucl. Energy Agency. 121.Choppin, G. R., and E. N. Rizkalla. 1994.Solution Chemistry of Actinides andLanthanides. In Handbook on the Physics andChemistry of Rare Earths. Edited by K. A.Gschneidner Jr. and L. Eyring. New York:North-Holland Publishing Co.Choppin, G. R., J.-O. Liljenzin, and J. Rydberg.1995. Behavior of Radionuclides in theEnvironment. In Radiochemistry and NuclearChemistry. Oxford: Butterworth-Heinemann.Choppin, G. R., 1983. Radiochim. Acta 32: 43.Dozol, M., R. Hagemann, D. C. Hoffman,J. P. Adloff, H. R. Vongunten, J. Foos, et al.1993. Pure and Appl. Chem. 65: 1081.Efurd, D. W., W. Runde, J. C. Banar,D. R. Janecky, J. P. Kaszuba, P. D. Palmer,et al. 1998. Environ. Sci. Tech. 32: 3893.Francis, A. J., J. B. Gillow, C. J. Dodge,M. Dunn, K. Mantione, B. A. Strietelmeier,et al. 1998. Radiochim. Acta 82: 347.Hascke, J., T. Allen, and L. Morales. 2000.Science 287: 285.Hersman, L.E., P. D. Palmer, and D. E. Hobart.1993. Mater. Res. Soc. Proc. 294: 765.Kraemer, S. M., S. F. Cheah, R. Zapf, J. Xu,K. N. Raymond, and G. Sposito. 2000.Geochim. Cosmochim. Acta (accepted).Kazuba, J. P., and Runde, W. 1999. Environ. Sci.Tech. 33: 4427.Kersting, A. B., D. W. Efurd, D. L. Finnegan,D. J. Rokop, D. K. Smith, J. L. Thompsom.1999 Nature 397: 56.Kim, J. I. 1986. Chemical Behavior ofTransuranic Elements in Natural AquaticSystems. In Handbook on the Physics andChemistry of the Actinides. Edited by A. J.Freeman and G. H. Lander. New York:North-HollandPublishing Co.Langmuir, D. 1997. Chap. 13. In AqueousEnvironmental Geochemistry. Upper SaddleRiver, NJ: Prentice Hall.Leonard, P. A., B. A. Strietelmeier,L. Pansoy-Hjelvik, R. Villarreal. 1999.Microbial Characterization for the Source-Term Waste Test Program (STTP) at LosAlamos. Conference on Waste Management99, Tucson, AZ.Lemire, R. J. and P. R. Tremaine. 1980.J. Chem. Eng. Data. 25: 361.Lu, N., I. R. Triay, C. R. Cotter, H. D. Kitten,and J. Bentley. 1998. "Reversibility ofsorption of plutonium-239 onto colloids ofiron oxides, clay and silica." Abstract ofAgronomy, p.31 ASA, CAA, SSSA AnnualMeetings, Baltimore, MD.Neu, M. P., W. Runde, D. M. Smith,S. D. Conradson, M. R. Liu, D. R. Janecky,and D. W. Efurd. 2000. EMSP Workshop,Atlanta, GA.Nitsche, H., A. Muller, E. M. Standifer, R. S.Deinhammer, K. Becraft, T. Prussin, R. C.Gatti. 1993. Radiochima. Acta. 62: 105.Runde, W., S. D. Reilly, M. P. Neu. 1999.Geochim. Cosmochim. Acta. : 3443.Runde, W., M. P. Neu, D. L. Clark. 1996.Geochim. Cosmochim. Acta 60: 2065.Shannon, R. D. 1976. Acta Cryst. 32A: 751.Wolfgang Runde earned B.S. and Ph.D.degrees in chemistry from the Technical Universityof Munich, Germany, in 1990 and 1993,under the direction of Professor Dr. J. I. Kim.He then worked for two years as a scientific staffmember at theInstitute forRadiochemistry,TU Munich.During this time,he was invited towork as a guestscientist at SandiaNational Laboratorieson technicalaspects ofthe actinidesource termmodel for WIPP. Wolfgang then spent a year aspostdoctoral fellow at the G. T. Seaborg Institutefor Transactinium Science at the LawrenceLivermore National Laboratory before joiningthe Chemical Science and Technology Divisionat Los Alamos National Laboratory in 1996.He is presently a staff member and leader of theActinide Environmental Chemistry team in theEnvironmental Science and Waste TechnologyDivision with a joint appointment in the Nuclearand Radiochemistry group in the ChemicalScience and Technology Division. His researchinterests are in the areas of inorganic, environmental,and radiochemistry of f elements,including the actinides, with the focus onnuclear waste isolation and the environment. Hisresearch activities have been closely affiliatedwith the national nuclear waste programs in theUS (WIPP and Yucca Mountain) and Germany(Gorleben).Number 26 2000 Los Alamos Science 411

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