The Chemical Interactions of Actinides in the Environment

392 Los Alamos Science Number 26 2000

The Chemical Interactions of Actinides in the Environment

From a chemist’s perspective, theenvironment is a complex, partic-ularly elaborate system. Hundreds

of chemically active compounds andminerals reside within the earth’s formations, and every patch of rock andsoil is composed of its own particularmix. The waters that flow upon andthrough these formations harbor a vari-ety of organic and inorganic ligands.Where the water meets the rock is apoorly understood interfacial region thatexhibits its own chemical processes.

Likewise, from a chemist’s point ofview, the actinides are complex elements. Interrupting the 6d transitionelements in the last row of the periodictable, the actinides result as electronsfill the 5f orbitals. Compared with the4f electrons in the lanthanide series, the5f electrons extend farther from the nucleus and are relatively exposed.Consequently, many actinides exhibitmultiple oxidation states and formdozens of behaviorally distinct molecu-lar species. To confound matters, thelight actinides are likely to undergo reduction/oxidation (redox) reactionsand thus may change their oxidationstates under even the mildest of condi-tions. Uranium, neptunium, and espe-cially plutonium often display two or

more oxidation states simultaneously inthe same solution. Accordingly, thelight actinides exhibit some of the rich-est and most involved chemistry in theperiodic table.

As a result, the chemical interactionsof actinides in the environment are in-ordinately complex. To predict how anactinide might spread through the envi-ronment and how fast that transportmight occur, we need to characterize alllocal conditions, including the nature ofsite-specific minerals, temperature andpressure profiles, and the local waters’pH, redox potential (Eh), and ligandconcentrations. We also need a quanti-tative knowledge of the competing geochemical processes that affect theactinide’s behavior, most of which areillustrated in Figure 1. Precipitation anddissolution of actinide-bearing solidslimit the upper actinide concentration insolution, while complexation and redoxreactions determine the species’ distrib-ution and stability. The interaction of adissolved species with mineral and rocksurfaces and/or colloids determines ifand how it will migrate through the environment.

Understanding this dynamic inter-play between the actinides and the en-vironment is critical for accurately as-

sessing the feasibility of storing nuclearwaste in geologic repositories. The actinides1 are radioactive, long-lived,and highly toxic. Over the last fewdecades, vast quantities of transuranicactinides (those with atomic numbersgreater than that of uranium) have beenproduced inside the fuel rods of com-mercial nuclear reactors. Currently,most spent fuel rods are stored above-ground in interim storage facilities, butthe plan in the United States and someEuropean countries is to deposit thisnuclear waste in repositories buriedhundreds of meters underground. TheDepartment of Energy (DOE) is study-ing the feasibility of building a reposi-tory for high-level waste in YuccaMountain (Nevada) and has already licensed the Waste Isolation Pilot Plant(WIPP) in New Mexico as a repositoryfor defense-related transuranic waste.Given the actinides’ long half-lives,these repositories must isolate nuclearwaste for tens of thousands of years.

1Although “actinides” refers to the fourteen elements with atomic numbers 90 through 103(thorium through lawrencium), in this article welimit the term to refer only to uranium, neptunium,plutonium, americium, and curium, unless statedotherwise. These five actinides are the only onesthat pose significant environmental concerns.

Wolfgang Runde

Actinides typically form large molecular complexes

in solution. Neptunium assumes the Np(V) species, NpO2(H2O)5

+, in many natural waters.

Complexation with different ligands can stabilize actinides in solution

and enhance their transport through the environmentRedox reactions change an actinide's

oxidation state and help to establish equilibrium between species

Microbes can facilitate actinide redox processes, while sorption or uptake by the microbes may

be a potential transport or immoblilzation mode.

At low actinide concentrations, sorption onto particulates, clays, or rock surfaces determines the environmental behavior. Actinides can also diffuse into rock, or coprecipitate with

natural ligands and become incorportated into minerals

Np(V) complex sorbed to montmorillonite

Solvated actinide species canprecipitate, forming a solid that in turndetermines the upper concentration

limit for the solvated species.

NaNpO2CO3(s)

The Np(IV) speciesNp(H2O)8

4+

NpO2(CO3)35–

Com

plex

atio

n

Dissolution

Precipitation

Actinides can migrate by water transport or by sorption

onto mobile particulates or colloids

Figure I. Overview of Actinide Behavior in the Environment

Transport

Redox

Bioavailability

Sorp

tion

Number 26 2000 Los Alamos Science 393

Chemical Interactions of Actinides in the Environment

Over the course of millennia, however,it is possible that water seeping into arepository will eventually corrode thewaste containers. The actinides willthen have access to the environment.

Research at Los Alamos has focusedon characterizing the behavior of actinides in the environments surround-ing those repository sites. (Only smallamounts of transuranic elements aregenerally found in other environments,as discussed in the box “Actinides inToday’s Environment” on page 396.)

Our solubility studies, for example,confirm that neptunium is more than athousand times more soluble in YuccaMountain waters than plutonium. Thatis because plutonium in those watersfavors the IV oxidation state while neptunium favors the far more solubleV state. Thus, for Yucca Mountain,neptunium is the actinide of primaryconcern. WIPP, however, is built withinin a geologically deep salt formation. Inthe extremely salty brines found there,the focus shifts to plutonium, since

reactions with chloride ions are knownto stabilize plutonium in the VI oxida-tion state. Plutonium(VI) species aremuch more soluble than Pu(IV) species,and plutonium mobility would be enhanced at WIPP unless a reducingenvironment can be maintained withinor around the waste containers.

Unfortunately, very few studies ofactinide geochemistry can be conductedin situ, so that we are forced tosimulate environmental conditions inthe laboratory. The concentrations ofthe actinides in natural waters are typically on the order of 10–6 molar(M) or lower. While those concentra-tions are high enough to be of environ-mental concern, they are too low toallow direct study with conventionalspectroscopies. We have thus had toadapt advanced spectroscopic tech-niques, such as x-ray absorption finestructure (XAFS) and laser-induced fluorescence spectroscopies, to studythese toxic, radioactive elements. As an example, recent XAFS investiga-tions of Pu(IV) colloids have yieldedsignificant insights into the colloids’structure, which for years has compli-cated the determination of Pu(IV) solubility. Scanning electron microscopy (SEM) has helped us eluci-date the confusing sorption behavior of U(VI) onto phosphate mineral surfaces. The box on page 412 high-lights a few examples of these spectroscopic studies.

The scope of actinide interactions inthe environment is too broad to coverin a single article. We will concentratetherefore on the solubility and specia-tion of actinides in the presence of hydroxide and carbonate ions, whichare the two most environmentally rele-vant ligands. We will then briefly dis-cuss some aspects of actinide sorptionand actinide interactions with microor-ganisms. The solubility and sorption ofactinides pose two key natural barriersto actinide transport away from a geologic repository, while microorgan-isms pose a less-studied but potentialthird barrier. While we are fully awareof additional geochemical reactions in



1.0

0.5

–0.5

0

Red

ox p

oten

tial,

Eh (

volts

)

0 2 4 6pH

8 10 12 14

PuO2+

PuF22+

Pu3+

PuO2CO3(aq)

PuO2(OH2)2(aq)

PuO2OH(aq)

Pu(OH)4(aq)

PuOH2+

PuO2(CO3)22–

PuO2(CO3)34–

Groundwater

Rain/streams

Pu(III)

Pu(IV)

Pu(V)

Pu(VI)

natural waters

PuO2F+

PuO2CO3–Ocean

Figure 2. Pourbaix Diagram for PlutoniumThis Eh-vs-pH diagram is calculated for plutonium in water containing hydroxide, carbonate, and fluoride ions. (The ligand concentrations are comparable to those foundin water from well J-13 at Yucca Mountain, Nevada, while the plutonium concentrationis fixed at 10–5 M). Specific complexes form within defined Eh/pH regions, while the stable oxidation states (colors) follow broader trends. For example, more-oxidizing conditions (higher Eh values) stabilize redox-sensitive actinides like plutonium in thehigher oxidation states V and VI. The red dots are triple points, where plutonium canexist in three different oxidation states. The range of Eh/pH values found in natural waters is bounded by the solid black outline. In ocean water or in groundwater, plutonium is likely to be found as Pu(IV), while in rainwater or streams, plutonium can assume the V state. Other natural environments will favor Pu(III) or Pu(VI) complexes.The dashed lines define the area of water stability. Above the upper line, water is thermodynamically unstable and is oxidized to oxygen; below the lower line, water is reduced to hydrogen.

the environment, such as coprecipita-tion, mineralization, and diffusionprocesses (and have thus included themin Figure 1), we will not discuss theirroles here.

Oxidation States and RedoxBehavior

Water is the dominant transportmedium for most elements in the envi-ronment. 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 actinide oxidation states predominate and whichactinide species are stable.

For example, Figure 2 shows thePourbaix diagram (Eh vs pH) for pluto-nium in water containing the two mostenvironmentally relevant ligands, thehydroxide (OH–) and carbonate (CO3

2–)ions. Even in this simple aqueous system, plutonium can exist in four oxidation states: III, IV, V, and VI.(Plutonium is unique in this regard. The Pourbaix diagrams for the other actinides 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 oxida-tion states are in equilibrium.

Because of intrinsic differences inredox potentials, each actinide will exhibit a different set of oxidationstates for a given set of solution condi-tions. In contrast to plutonium, U(III) isunstable under most conditions and oxidizes 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 stability region 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 domi-nated by 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 groundwater environments. But in other aqueous en-vironments, including streams, brines,or bogs, U(IV), Np(IV), andPu(III,V,VI) are also common and like-ly to be stable. Table 1 summarizes theoxidation states of the actinides andhighlights the environmentally most relevant 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 in natural systems.

Additional chemical processes occur-ring in 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 magni-tude greater than those of Pu(IV) solids,so the enhanced stability of Pu(V)would serve to increase the total pluto-nium concentration 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 (H2O2) 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 discus-sions, therefore, we often refer to“generic” actinides—e.g., An(IV) complexes—where An is shorthand for actinide.

Effective Charge. It has beenknown for many years that An(IV)forms the strongest, most stable complexes and An(V) the weakest. Thisbehavior follows directly from the effective charges of the ion.

In the III and IV states, the actinidesform hydrated An3+ and An4+ 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, AnO2

+ and



Table I. Oxidation States of Light Actinidesa

Th Pa U Np Pu Am Cm

III III III III III III IIIIV IV IV IV IV IV IV

V V V V VVI VI VI VI

VII VII

aThe environmentally most important oxidation states are bolded.

AnO22+, with overall charges of +1 and

+2, respectively. These cations areoften referred to as the actinyl ions.The covalent bonding between the actinide and the two oxygen atoms inthe actinyl ion (O=An=O)n+, where n = 1 or 2, enhances the effectivecharge of the central actinide ion to2.3±0.2 for AnO2

+ and 3.3±0.1 forAnO2

2+ (Choppin 1983). A ligand approaching 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 com-plexes generally follows their effectiveion charges, as shown below:

The second and third lines show theion’s effective charge and formal oxida-tion states.

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 in-creases from 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 highest effective charge, it forms the most stable 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 there-fore the 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 most

soluble 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 strong

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 numbers

An4+ > AnO22+ ≥ An 3+ > AnO2

+

+4 +3.3 +3 +2.3IV VI III V



Actinides in Today’s Environment

In 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 (now Sellafield) 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 (account-ing only for announced nuclear tests). Atmospheric and underground tests have also injected 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 tests propelled 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 and oceans and poses little health risk to the general public.

Underground nuclear tests injected transuranics into highly localized regions surround-ing the 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.

is allowed. Coordination numbers be-tween 6 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 mole-cules in the first coordination spheresurrounding the ion, exhibit coordina-tion numbers that vary with oxidationstate: 8–10 for An3+, 9–12 for An4+,4–5 for AnO2

+ and 5–6 for AnO22+.

These coordination numbers have beendetermined from nuclear magentic reso-nance (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. This increase in stability follows the trend ofthe actinide contraction, wherein the

ionic radii decrease with atomic number. The ionic radii for An(IV) ionswith coordination numbers of 8 are reported to be 1.00, 0.98, 0.96, 0.95,and 0.95 Å for uranium, neptunium,plutonium, americium, and curium, re-spectively (Shannon 1976).

Solubility and Speciation

The solubility of an actinide is lim-ited primarily by two properties: thestability of the actinide-bearing solid(the solubility-controlling solid) and the stability of the complexed species insolution. An actinide species will precipitate when its dissolved concen-tration is above the thermodynamic(equilibrium) solubility of its solubility-controlling solid, or if the kinetics

favors the formation of a solid of high-er solubility. Formation of the solid setsan upper concentration limit for the actinide 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 concen-tration of An(III) in solution. At lowpH with increasing carbonate concen-tration, the hydrated carbonated solidAn2(CO3)3•2–3H2O(s), where (s) indi-cates the solid phase, readily precipi-tates and hence lowers the concentrationof An3+ in solution. But complexationof An3+ with carbonate stabilize the solution species and increases the over-all An(III) solubility. As the carbonateconcentration increases and the differentcarbonate species form, the total actinide concentration in solution first



Figure 3. Total An(III) ConcentrationVersus Carbonate Concentration(a) The partial pressure of CO2 in naturalwaters spans about three orders of magnitude (gray area). The stability diagram 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 increased stability 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 An2(CO3)3•2–3H2O(s). Calculatedconcentrations for individual solutionspecies are indicated by the gray linesthat run tangent to the solubility curve, or by the gray dashed lines. At low carbonate concentrations, the An3+ iondominates the solution species. The totalAn(III) concentration drops rapidly withhigher carbonate concentrations as moreof the solid phase precipitates. But carbonate complexation begins to stabi-lize the amount of actinide in solution. By the time the anionic bis- and triscar-bonate complexes dominate, the solutionconcentration increases.

10–7 10–6 10–5 10–4 10–3 10–2

An2(CO3)3 •2–3 H2O(s)

Carbonate concentration (M)

Tota

l sol

vate

d A

n(III

) con

cent

ratio

n (M

)

10–7

10–6

10–5

10–4

244Cm241Am

An(CO3)33–

An(CO3)2–

An(CO3)+

An(OH)2+

An3+

10–2

10–4

10–6

10–8

10–10

10–5 10–3 10–210–4 10–1 10

An3+(aq)An(

III) c

once

ntra

tion

(M)

AnOHCO3(s)

CO2 Partial Pressure (atm)

An(OH)3(s)

An2(CO3)3(s)

(a)

(b)

reaches 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 by performing experiments in the laboratoryunder well-controlled conditions inwhich the actinide concentration ismeasured while varying the ligand concentration. Such experiments enableus to measure the solubility product Kspand the stability constant β. These thermodynamic parameters are the basisfor modeling the solubility boundariesfor actinides in natural waters, as discussed in the box on “Solubility andSpeciation Parameters.”

We should note, however, thatmeaningful interpretation of solubilitydata requires a detailed knowledge ofthe composition, crystallinity, and solu-bility of the solubility-controlling solid,along with the steady-state concentra-tion and composition of the solutionspecies. Steady state is assumed to beestablished when the actinide concen-tration remains stable for several weeksor longer. But the solubility-controllingsolids of the actinides generally precipi-tate as an amorphous phase because radiolysis 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 not represent the most thermodynamicallystable solids with the lowest solubility.Laboratory-based solubility studiestherefore provide an upper limit to actinide concentrations in a potential release scenario from a nuclear wasterepository.

Because of their omnipresence, hydroxide ions (OH–) and carbonateions (CO3

2–) are the most importantligands that complex to actinides in theenvironment. While other ligands, suchas phosphate, sulfate, and fluoride, canlower the actinide concentration



Solubility and Speciation Parameters

Solubility products (Ksp) 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. For example, the solubility product of AmOHCO3(s), where (s) refers to the solid phase, isgiven by the product of the concentration of the dissolved ions Am3+, OH–, and CO3

2–

according to the following reaction:

(1)

The solubility product is then given by

(2)

where the brackets indicate ion concentration. The formation of an actinide complex insolution, for example,

(3)

is described by the complex’s formation constant, βn. For the forward reaction in Equation (3), the formation constant is

(4)

where n refers to the number of ligands. For the reverse reaction, βn 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:

(5)

Accordingly, the total actinide concentration in solution is given by the sum of the con-centrations of species present in solution as follows:

(6)

Complexation of the dissolved actinide ion with ligands in solution (large βn values)generally increases the total actinide concentration in solution.

Am(III) = Am + +

= Ksp 1 + + .

3+

nn

nm

m

[ ] [ ] ⎡⎣⎢

⎤⎦⎥ [ ]

[ ][ ]× [ ] [⎛

⎝⎜⎞⎠⎟∑ ∑β β

Am(CO3)n3–2n

CO32–

CO32–

OH–OH–

Am(OH)m3–m

m]

= Am CO = K CO

OH .n

3+32– n

n sp32– n–1

–⎡⎣⎢

⎤⎦⎥ [ ] [ ] [ ]

[ ]β βAm(CO3)n3–2n

βn n=

Am3+ ,

⎡⎣⎢

⎤⎦⎥

[ ][ ]Am(CO3)n

3–2n

CO32–

Am + 3+ nCO32– Am(CO3)n

3–2n

K = [Am ][OH ][CO32– ] , sp

3+ –

AmOHCO3(s) Am + OH + CO32– . 3+ –

(because of the low solubility of thecorresponding solid phase), their concentrations in natural waters aregenerally low. Consequently, they cannot compete successfully with hydroxide or carbonate complexation.However, organic biodegradation prod-ucts, such as humates or fulvates, aregenerally present in natural aquifers andcan potentially play a role in actinidesolubility and migration.

Hydrolysis. The interaction of actinide 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), trivalent actinides form the positively chargedor neutral hydroxo complexes of general formula An(OH)m

3–m, where m = 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 hydroxide

An(OH)n(s), where n = 3 for An(III)and 4 for An(IV) or the hydratedoxide AnO2•nH2O for An(IV) are expected to control the solubility ofactinides in both oxidation states.

Understanding the An(IV) solutionchemistry is still a formidable chal-lenge. The high charge allows the An4+

ions to easily hydrolyze and form mul-tiple species simultaneously, even atlow pH. Plutonium(IV) is known to hydrolyze even in dilute acidic solutions and, at concentrations greaterthan 10–6 M, will undergo oligomeriza-tion and 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-micron-scale colloid particles are nominallycomposed of Pu(IV) hydroxo polymersand contain structural characteristics of

plutonium oxide, PuO2. 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 amor-phous oxyhydroxide compound. Thefunctional groups originate from thepolymerization of monomeric speciesformed through hydrolysis, and thenumber of functional groups dependson the history of the solution prepara-tion and aging process. While thechemical “structure” of every Pu(IV)colloid may be similar, each colloid canhave a unique distribution of functionalgroups depending on solution prepara-tion, 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 pluto-nium atoms together in the polymer areconverted into chemically resistant oxy-gen bridges.

The generation of Pu(IV) colloids isone of the main processes that compli-cate determining thermodynamic constants and accurately predicting soluble plutonium concentrations in natural waters. Both the colloids andother amorphous Pu(IV) hydroxidesolids span a wide range of structuralfeatures, crystallinities, and stabilities,and hence exhibit a wide range of solubilities. The presence of Pu(IV)colloids will therefore complicate anycalculation of plutonium solubility, aswill be evident when we review plutonium solubility and speciation inYucca Mountain waters.

Solid hydroxides and oxides are alsoexpected to limit the solubility ofAn(V) and An(VI) complexes. For example, Np2O5(s) has been determinedto be the solubility-controlling solid inwaters from Yucca Mountain, whileschoepite, UO2(OH)2•nH2O(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 forma-

tion of anionic hydroxo complexes:AnO2(OH)2

– for An(V) andAnO2(OH)n

2–n (n = 3 and 4) forAn(VI). The hydrolysis of U(VI) iscomplicated even more by the formation of polymeric hydroxidespecies at high U(VI) concentration(>10–5 M), i.e., (UO2)3(OH)5

+ or(UO2)4(OH)7

+. We are currently conducting experiments to see if Pu(VI)undergoes a similar polymerization.

Carbonate Complexation. Carbon-ate is ubiquitous in nature. In surfacewaters, such as oceans, lakes, orstreams, its concentration ranges gener-ally between 10–5 and 10–3 M. Ingroundwaters, its concentration isenhanced (up to 10–2 M) because of theincreased CO2 partial pressure, which isabout 10–2 atm deep underground compared 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 to actinides 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 in solution and in the solid state.

We have investigated the structureof carbonate complexes involving theactinyl ions AnO2

+ and AnO22+. (These

complexes are more soluble and thusmore easily studied.) As shown in Figure 4, three monomeric solutioncomplexes form. Depending on thesolution conditions, these complexesmay coexist with each other and withhydroxo complexes. It is therefore diffi-cult to obtain clean, unambiguous struc-tural data about a single species.

We can make inroads into that prob-lem by investigating the stability andstructure of the triscarbonato species,AnO2(CO3)3

m-6, shown in Figure 4(c).This is the limiting (fully coordinated)An(V,VI) carbonate complex in solu-tion. At carbonated concentrationsabove 10–3 M, this complex forms atthe exclusion of the others. Using a



battery of spectroscopic techniques, including multinuclear NMR spec-troscopy, x-ray diffraction, and XAFS,we have obtained formation constants,bond lengths, and other structural parameters for the An(V) and (VI) car-bonate complexes (Clark et al. 1999).This information has helped us under-stand various aspects of the An(V,VI)carbonate system. (The article by Clarkthat begins on page 310 describes struc-tural investigations 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 formu-las An(OH)(CO3)(s) for An(III) andAn(OH)2m(CO3)n(s) for An(IV), whereAn stands for uranium, neptunium, plu-tonium, 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 different compositions by increasing the carbon-ate concentration. This increase reducesthe extent of hydrolysis and allows the formation, for example, of pureAn(III) carbonate solids, such asAm2(CO3)3•nH2O(s) orNaAm(CO3)2•nH2O(s), which are the solubility-controlling solids in solu-tions containing high concentrations of sodium and carbonate. These solidsare analogous to the ones formed by the trivalent lanthanides Nd(III) and Eu(III).

Similarly, we have observed the formation of solid An(VI) and An(V)compounds in solutions containing highconcentrations of alkali metals and carbonate. 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 AnO2CO3(s) and are madefrom highly ordered actinyl carbonatelayers. However, An(V) solids need to

incorporate a monovalent cation, suchas Na+, K+, or NH4

+, into their struc-ture (otherwise they would have anoverall charge). Thus, we observe onlyternary An(V) solids with general com-positions M(2n–1)AnO2(CO3)n•mH2O(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 in-creases (stoichiometrically from 0 to 1to 3), the structures become increasing-ly open, and the solids generally become more soluble under conditionsin which the uncomplexed actinide iondominates solution speciation. Replace-ment of the cations with larger ones

enhances the solubility of the actinideby several orders of magnitude. We exploited that fact when we studied theneptunyl “double carbonate” solidsM3NpO2(CO3)2. These solids exhibitsuch low solubilities in near neutralsolutions (10–4 to 10–7 M) that the

solution species cannot be studied bymany spectroscopic methods. To modi-fy the solubility and obtain more-concentrated neptunyl solutions, we examined the use of large, bulkycations in an attempt to generate an unfavorable fit of the complex into thestable lattices shown in Figures 5(b)and 5(c). We found that tetrabutyl-ammonium provided stable solutions ofNpO2CO3

– and NpO2(CO3)23– in



Figure 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 monomer-ic complexes shown in (a)–(c), m = 1 for An(V) and m = 2 for An(VI). Actinides in the V state have the lowest effective charge and form the weakest complexes. In these complexes, 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 carbonate ligands: 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 concentra-tions to form the polynuclear actinyl carbonate complex (UO2)3(CO3)66–. 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, AnO2(CO3)22–.

(c) AnO2(CO3)3m–6

(d) (UO2)3(CO3)66–

Terminal carbonate

Bridgingcarbonate

(b) AnO2(CO3)2m–4(a) AnO2(CO3)m–2

Oxygen

Carbon

Actinide

millimolar concentrations over a widepH range. The enhanced solubility allowed us to study the temperature dependence of the carbonate complexa-tion reactions with near-infrared (NIR)absorption spectroscopy and to obtainadditional structural information aboutthe Np(V) solution carbonate complex-es with XAFS spectroscopy.

The compositions and stabilities ofthese actinyl carbonate solids dependon the ionic strength and carbonateconcentration and are controlled by theactinide, carbonate, and sodium concen-trations in solution. For example, thestability of the Np(V) solidsNa(2n–1)NpO2(CO3)n•mH2O(s) is givenby the solubility product Ksp:

Ksp = [Na+]2n-1 [NpO2+] [CO3

2–]n .

For the An(VI) solid AnO2CO3(s),n is equal to 1 and its solubility is nominally independant from the sodiumconcentration in solution. At the lowionic strengths of most natural waters,an unrealistically high NpO2

+ concen-tration would be needed to precipitate aternary Np(V) carbonate solid. Thus,solid An(V) or An(VI) carbonates arein general not observed in natural wa-ters. However, because the ternary compounds of Np(V) have been observed to form in concentrated electrolyte solutions, these solids areimportant for solubility predictions inbrines found at WIPP.

Solubility and Speciation inYucca Mountain Waters

As a way of tying together many ofthe concepts presented in the previoussections, we can consider the solubilityand speciation of neptunium and pluto-nium in 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 effec-tive charge of the AnO2

+ ion, formsweak complexes. Thus, neptunium hasthe distinction of being the most solubleand transportable actinide and is the



UO2CO3 KPuO2CO3 K3NpO2(CO3)2•nH2O

Figure 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 UO2CO3(s) or PuO2CO3(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 NH4

+. 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 MAnO2CO3(s). This opens the structure, in comparison with the An(VI) binary structure, and increases solubility. (c) At higher metal ion concentrations, metal ions will be incorporated into each layer to form very open structures of general formula M3AnO2(CO3)2(s).

one of greatest concern when assessingthe environmental safety of the poten-tial underground storage site at YuccaMountain. Plutonium is considered themost toxic actinide and is always of environmental 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).

However, groundwater conditions with-in and around Yucca Mountain vary.The carbonate concentration changes(because the CO2 partial pressurechanges with depth), as does the pH,Eh, and temperature. We mimicked thisvariation in the natural waters by con-ducting our studies at various pH val-ues while maintaining a constant ionicstrength. We also conducted our studiesat different temperatures (25˚C, 60˚C,and 90˚C) to approximate the range of

temperatures 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 the actinides in J-13 water—in all oxidationstates—were mainly oxides and hydrox-ides. We observed only the formationof greenish brown Np(V) crystallinesolids of general formula Np2O5•nH2O(s),green Pu(IV) solids of general formulaPuO2•nH2O 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 NpO2

+.The remaining soluble Np(V) was com-plexed with either carbonate or hydrox-ide, with about 58 percent NpO2CO3

–

and 11 percent NpO2OH(aq). At pH 6,however, Np(V) hydrolysis and carbon-ate complexation reactions were mini-mized, and all of the neptunium waspresent as NpO2

+. By comparison, at all pH values

examined 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 geochemical models, 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 amorphous oxides/hydroxides or colloidal forms ofPu(IV). (Note that if crystalline PuO2(s)were the solubility-controlling solid,Pu(IV) solubility would be about 10–17

M, 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 dramatically affect neptunium solubility, as illustrated



Figure 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 differ-ent solubility-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 Np2O5(s), rather than the previously postulated sodium carbonate solidNaNpO2CO3(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 Ksp data of Pu(IV) hydroxide or hydrousoxide reported in the open literature (gray area).

0.01 0.1 1.0

Spe

cies

con

cent

ratio

n (m

ol k

g−1)

Ionic strength (mol kg−1)J-13 water

NaNpO2CO3 •nH2O(s)

Np2O5 •nH2O(s)

Np undersaturated

Pu (Nitsche et al. 1993)

Np (Nitsche et al. 1993)

Pu oversaturated

Pu(OH)4(s) (Lemire and Tremaine 1980)

10−2

10−4

10−6

10−8

10−10

10−12

Np oversaturated

in Figure 7 (Kaszuba and Runde 1999).Under strong reducing conditions(below about –120 millivolts), the totalsoluble neptunium concentration is limited by a Np(IV) solid phase,Np(OH)

4(s), and is conservatively

assumed to be about 10–8 M. Similar toplutonium, the presence of a more ordered solid oxyhydroxide or oxide,such as NpO2(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 oxidiz-

ing and Np(V) becomes more stable insolution. Eventually, the Np(V) concen-tration is great enough to allow aNp(V) solid phase, Np2O5(s), to precip-itate. 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 the

far-field environment. Spent nuclearfuel contains mainly UO2 and actinidesin low oxidation states (III and IV). Ifconditions at the disposal site are reducing—for example, because of microbial activities or the reduction capacity of corroding iron from storagecontainers—the actinides will be main-tained in their lower oxidation statesand soluble actinide concentrations maybe low. If local reducing conditionschange because oxidizing waters infil-trate from the surroundings, actinideswill enter their more-soluble oxidationstates and begin to migrate into the environment. Once far from the reposi-tory, conditions may change drastically,



pH

pH 7

4 6 8 1210 −400 −200 2000 600400

400 mV

300 mV

200 mV

100 mV

Np(OH)4(s)

Np(OH)4(s)

Np2O5(s)

Np2O5(s)

−100 mV0 mV10−8

10−6

10−4

10−2

1

10−10

Np

conc

entra

tion

(mol

kg−1

)

Np

conc

entra

tion

(mol

kg−1

)

10−6

10−4

10−8

Eh (mV)

(b)(a)

−200 mV

Figure 7. Neptunium Solubility in Yucca Mountain J-13 Water Neptunium solubility varies tremendously depending on the Eh and pH of the water. (a) The graph shows the neptunium concentra-tion as 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,Np2O5(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 six orders 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 condi-tions (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 Np2O5(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 Eh.

and the transport behavior of redox-sensitive elements such as plutonium orneptunium may change. Clearly, evalu-ation of actinide mobility in the geo-chemical environment around nuclearwaste repositories must account for soluble actinide concentrations that arecontrolled by the redox potential of thetransport medium and as well by thesolubility of the actinides’ solid phases.

Complexation with Chlorideand Implications for WIPP

In most natural waters, actinides areusually associated with hydroxide orcarbonate ligands. However, WIPP,which is designed to hold defense-related transuranic waste, is carved outof a salt formation. Water at the WIPPsite thus consists of concentrated brines

that are saturated with chloride saltsand exhibit very high ionic strengths(between 4 and 8 molal). With increas-ing distance from the repository, thebrines’ ionic strength decreases andmay ultimately reach the lower valuesof most natural surface and ground-waters. 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 of actinide compounds significantly. Asmentioned earlier, in most natural waters plutonium favors the IV and Voxidation states, but radiolytichypochlorite formation in concentratedchloride solution may result in the formation of Pu(VI). Furthermore,complexation by chloride may enhance stabilization 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 asPuO2Cl+ and PuO2Cl2, formed as afunction of chloride concentrations (seeFigure 8). Interestingly, the tris- andtetrachloro complexes are not formed insignificant concentrations in NaCl solutions (up to the point of NaCl saturation). However, we have observedspectroscopically UO2Cl4

2– andPuO2Cl4

2– at very high LiCl concentrations and obtained structuraldata from single crystals synthesizedfrom concentrated chloride solutions.

The interaction of actinides withchloride and other weak anionic ligandssignificantly affects solubility predic-tions and must be considered in risk assessments of WIPP and other nuclearwaste repositories in salt formations.With regard to neptunium, chloride enhances the Np(V) solubility becauseof a strong ion interaction between theNp(V) compounds and chloride ions



870860850840830820Wavelength (nm)

PuO22+

PuO2Cl+

PuO2Cl2PuO2Cl3

–

PuO2Cl42–

–10

–5

5

10(a)

(b)

0 1 2 3 4

Inte

nsity

(arb

itrar

y un

its)

R – φ (Å)

1-M HClO4

Pu=O Pu–Oeq

Pu–Cl

15-M LiCl

4-M NaCl

1-M NaCl

[Cl](M)

1 4 15

Pu=O(Å)

1.741.751.75

(Å)

2.402.432.49

Pu–Cl(Å)

—2.752.70

Pu–Oeq

Four

ier T

rans

form

mag

nitu

de

0

Figure 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 non-complexing 1-M HClO4, 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 water molecules (indicated by the Pu–Oeq 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 PuO2

2+ ion, indicate the formation of inner-sphere Pu(VI) chlo-ride complexes. These bands allowed us to determine the formation constant, β, of themost important Pu(VI) chloride complexes in NaCl solutions, PuO2Cl+ and PuO2Cl2(aq).

(see Figure 9). Runde et al. (1996)simulated this interaction using thePitzer ion-interaction model and devel-oped a Pitzer parameter set to predictthe Np(V) solubility at different NaClconcentrations and in artificial brine solutions (Runde et al. 1999). Neglect-ing the 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 chemical potential 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) solubil-ities reported 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, respec-tively. 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 for electrolyte ions such as K+, Mg2+,Ca2+, and SO4

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.

Sorption

The solubility of an actinide speciesprovides an upper limit to the actinideconcentration in solution and may beconsidered as the first barrier to actinide transport. Once an actinidehas dissolved in water, its sorptiononto the surrounding rock or mineralsurfaces is expected to act as a secondary barrier. When the actinideconcentrations are far below the solubility limits, as is likely in mostsurface and groundwater systems,water/rock interfacial processes are

expected to dominate actinide transportthrough the environment. (Simple dilu-tion will lower the actinide concentra-tion, typically to trace concentrations,once the actinide migrates away fromthe contaminant or waste source.)

Actinide sorption onto mineral surfaces depends on many factors, suchas the actinide’s speciation and concen-

tration, the flow rate and chemistry ofthe water, and the composition of thesurrounding rocks. In addition, differentmechanisms, such as physical adsorp-tion, ion exchange, chemisorption, oreven surface precipitation, can fix theactinide to a mineral/rock surface or ingeneral inhibit its transport. Further-more, redox processes complicate the



10−7

10−6

10−5

10−4

10−3

10−7 10−6 10−5 10−4 10−3 10−2 10−1

An(V

) (m

ol k

g−1)

Spat

ial d

istri

butio

n (%

)

Carbonate concentration (mol kg−1)

NaNpO2CO3•3.5H2O(s)

Na3NpO2(CO3)2•nH2O(s)

100

80

60

40

20

NpO2+

NpO2(CO3)35−

NpO2(CO3)23−

NpO2(CO3)−

5-M NaClO4

5-M NaCl

Figure 9. Solubility of Np(V) in Concentrated Electrolyte Solutions The 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 NaClO4solution over the same range. Similar data are obtained for Am(V). The neptunium solubility is about an order of magnitude higher in NaCl because of the stabilizing effect of actinide-chloride complexation reactions. The solubility-controlling solid forboth curves at low carbonate is the hydrated ternary Np(V) carbonate,NaNpO2CO3•3.5H2O(s). Interestingly, this solid initially forms at high carbonate concentrations. While it is kinetically favored to precipitate, it is thermodynamically unstable. Over time, it transforms to the thermodynamically stable equilibrium phaseNa3NpO2(CO3)2•nH2O(s) (dashed curve).

sorption behavior of neptunium andplutonium. Unfortunately, we are onlybeginning to understand such actinidesorption and transport mechanisms.

Generally, batch-sorption experi-ments are performed to investigate howactinides will be retained in the envi-ronment. An actinide-containing solu-tion is allowed to come in contact withthe geomaterial of interest, and wemeasure the amount of actinide that

remains in solution. The difference be-tween the remaining and initial actinideconcentrations equals the amount thatsorbed to the sample. This differenceallows us to calculate the batch-sorptiondistribution coefficient Kd, defined asthe moles of radionuclide per gram ofsolid phase divided by the moles of ra-dionuclide per milliliter of solution andreported as milliliters per gram (mL/g).Low Kd values mean low sorption, and

high values imply effective immobiliza-tion due 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 extrapo-late to 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 varying but known compositions. Sorption coefficients from these studies canserve as input data into a transportmodel that will estimate the actinide retention along water flow paths ofknown mineral content. The model predictions can then be compared withexperimental results obtained from dynamic 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 exam-ple, K

dvalues for neptunium sorption

onto hematite range from 100 to



Figure 10. Uptake of Pu(V) by ColloidsLu et al. (1998) obtained data for Pu(V) adsorption onto colloids, including hematite(Fe2O3), silica (SiO2), and montmorillonite (Mont), in natural J-13 water and in a synthentic (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 of the sorption onto montmorillonite suggests a more complicated interaction than occurs with oxide minerals.

Table II. Desorption of Plutonium from Colloids in J-13 Water after 150 Days

Colloid Material Amount DesorbedPu(IV) Pu(V)

(%) (%)

Hematite (α-Fe2O3) 0.0 0.01Goethite (α-FeOOH) 0.57 0.77Montmorillonite 8.23 0.48Silica (SiO2) 19.93 1.04

Ads

orbe

d P

u(V

) (%

)

120

Fe2O3 (syn)

SiO2 (syn)

SiO2

Fe2O3

100

80

60

40

20

00 40 80 120

Time (h)160 200 240

Mont (syn)

Mont

2Flow-through, or dynamic, column experiments,wherein the actinide solution is allowed to migrate through a short column of packed geomaterial, provide information about dynamictransport properties. Dynamic column experi-ments are difficult to perform and interpret, andmost of the sorption studies on the actinides arestill performed as batch experiments.

2000 mL/g, while plutonium sorptiononto hematite results in K

dvalues

above 10,000 mL/g. This difference ismainly caused by the different stabili-ties of the actinide oxidation states andtheir complexation strengths. The lowcharge of the NpO2

+ ion results in alow sorption of neptunium, whilePu(IV) dominates the complexation andsorption processes of plutonium. Pluto-nium(V) is expected to follow the lowsorption affinity of Np(V).

Lu et al. (1998) determined the sorption 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 mecha-nisms are largely unknown. The time-dependence of the sorption onto theclay suggests a more complicated inter-action than occurs with oxide minerals.

Interestingly, the desorption ofPu(V) does not follow this general pat-tern. Lu et. al. (1998) also discussed thedesorption of Pu(IV) and Pu(V) fromvarious colloids. Table II summarizes

their 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 different binding to the silica functional groupsthan to the Fe(III) minerals, includingpossible changes in plutonium oxidation states.

These investigations indicate the importance of understanding the oxida-tion-state stability of the actinide sorbedon a mineral surface. Redox-activeminerals, such as manganese or ironoxide/hydroxides, are present in subsur-face environments and may alter theredox states of uranium, neptunium,and plutonium. At the Rocky Flats Environmental Technology Site, seasonal variations of plutonium con-centrations in settling ponds correlateastonishingly well with soluble manganese 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 the formation of strong mineral-surfacecomplexes. Quite surprisingly, Neu et al. (2000) showed that Pu(V) is not reduced when sorbed on MnO2. X-rayabsorption near-edge structure indicatedthe presence of Pu(V) on MnO2, eventhough the starting material of thebatch-sorption experiment was colloidalPu(IV) hydroxide. Their data also indicated that the plutonium was oxidized to Pu(VI). The only time theydid not observe the higher-oxidation-state surface 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 by surface interactions. We need more dataabout the surface species, however, tomodel such interactions.

Obviously, batch experiments alonecannot elucidate sorption mechanisms.Specifically, we have to distinguish be-tween surface precipitation and interfa-cial molecular reactions in order tomodel sorption behavior. In the event



keV

0 1 2 3 4 5

Ca

P

Ca

O

Ca

Ca

O

P

U

keV

0 1 2 3 4 5

Figure 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 × magnifica-tion) 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.

of a surface precipitation, the Kd valuemay represent the solubility product ofthe surface precipitate but will not reflect the sorption behavior of an actinide species.

For example, phosphate-containingminerals, such as apatite, are known toreduce thorium and uranium concentra-tions in contaminated solutions. Sorp-tion distribution 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 microcrys-talline precipitates on the apatite whenthe initial U(VI) concentrations wereabove 10–5 M (see Figure 11). At lowerU(VI) concentrations, no surface precipitation was observed. Thus, at thehigher concentrations, the Kd values actually 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 influ-

ence the 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 An3+ ion is stable anddoes not participate in bonding with hydroxide 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 lig-ands, such as SO4

2–, CrO42– or AsO4

3–.In addition to An(III) species, the sametrend of enhanced stability and low surface retention has been observed forthe An(VI) complex UO2(CO3)3

4–.Enhanced modeling efforts have

begun to address the fundamentalmechanism of actinide adsorption ontomineral surfaces. Janecky and Sattel-berger (unpublished results) modeledthe adsorption behavior of Np(V) oncalcite by minimizing the free energyof the coordination. The structural parameters of NpO2CO3

–, which is themain solution complex in neutral andlow-carbonate waters (such as J-13water), were used to mimic the adsorp-tion. 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 under-stand the sorption/desorption reactionsof actinides with colloids and the



pH

Spe

cies

dis

tribu

tion

(%)

20

40

60

80

An(CO3)33–An(CO3)2

–An(CO3)+

An(OH)2+

An3+

4 5 7 96 8 10

100

80

(a)

(b)

60

40

20

0

0.1-M NaCl1.0-M NaCl

Sor

ptio

n (%

)

Figure 12. Correlation between An(III)Solution Speciation and Sorption onto Montmorillonite(a) Even with an order-of-magnitude difference in NaCl concentration, sorptionof the trivalent actinides Am(III) andCm(III) peaks in near-neutral waters. Atlower pH, the An3+ 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 sorption because they are repelled from the negatively charged functional groups onthe mineral surface. (b) Calculations for a weak sodium chloride solution (0.1-M NaCl and CO2 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.

actinides’ resulting transport character-istics. This area of environmental migration 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 239Pu/240Puisotope 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 this remarkably fast movement of plutoniumthrough the water-saturated rock. It isnot clear, however, whether the transport was facilitated by intrinsicplutonium colloids or natural (clay orzeolite) colloids. What is clear is thattransport models to date have underesti-mated the extent of colloidal transporton plutonium mobility.

Microbial Interactions

We have discussed the most promi-nent chemical processes involving actinides in the environment: dissolu-tion and precipitation of solids, com-plexation in solution, and sorption ontosurfaces. But there is growing attentionto the importance of actinide interac-tions with microorganisms. Great vari-eties of 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 the hypersalinity 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 be associated with microbes, eitherthrough surface binding or throughmetabolic uptake, and be transported bythem through the environment. However,the microorganisms’ role in environ-mental transport behavior of radioactive

contaminants 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 metabolic activity, microbes can also establish reducing 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 of uranium, neptunium, and plutoniuminto less-soluble forms. Several bacteriaspecies have been reported in the litera-ture to 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 from geologic repositories and could alsohelp remediate actinide-contaminatedsoil, groundwater, and waste streams.The ultimate goal is to use microorgan-isms as “living” backfill to immobilizeactinides through processes such asbioreduction, biosorption, or bioprecip-itation and 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 the bacteria’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 to be negligible.

However, other microorganisms canconcentrate metal ions to a great extent.Hersman et al. (1993) observed a stronguptake of plutonium by pseudomonas

sp. Within the first week of the experi-ment, the bacteria concentrated plutoni-um to levels that were nearly 10,000times greater than a sterile control or-ganism. 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 its potential to remove radionuclides fromthose environmentally abundant miner-als. In fact, Kraemer et al. (2000) foundthat 20 percent of the sorbed plutoniumwas removed by this aerobically grow-ing bacteria after about four days. The experimental 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 mag-nesium oxide, MgO, on these WIPP-de-rived bacterial cultures. Magnesiumoxide is of specific interest because itwas chosen as backfill material at theWIPP repository to maintain the pH between 9 and 10, to reduce the amountof infiltrating water, and to limit theamount of CO2 that can accumulatethrough degradation of cellulose andother organic compounds in the wastes.Initial studies indicated an inhibited cellgrowth as the amount of MgO increased, as well as a decrease in microbially generated gas (N2O) by active halophilic bacteria. Ultimately,the toxic effect of MgO inhibits micro-bial growth, and consequently microor-ganisms may be unable to create a reducing environment at the repositorythat would stabilize actinides.

A more detailed study on the micro-bial association mechanism includedanalysis of the uptake of uranium andplutonium by sequestering agents andextracellular polymers (exopolymers).As an example, some exopolymers pro-duced by bacillus licheniformis consistof repeating γ-polyglutamic acid (γ-PGA). Neu et al. (manuscript in prepa-ration) observed actinide hydroxide pre-cipitation at metal-to-glutamate ratios



higher 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 concen-trations, the actinide ion is encapsulatedand isolated from hydroxo-colloid for-mation. 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 deter-mined to 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 exo-ploymer) increase the number of bind-ing sites and enhance metal binding.

This research clearly shows that actinides bind with microorganisms andtheir metabolic byproducts. Dependingon the static or dynamic nature of themicroorganism, the interaction of pluto-nium and other actinides with extra-cellular ligands may affect the actinides’ mobility both away fromstorage containers and within the environment. Future studies will focuson understanding the mechanisms ofactinide/microbial interactions and exploiting their potential for remedia-tion and immobilization efforts. Uptakestudies of plutonium by means of sequestering agents are highlighted inthe box on “Siderophore-Mediated Microbial Uptake of Plutonium” onpage 416.

Concluding Remarks

From this overview of actinide inter-actions in the environment, it should be clear that such interactions are com-plex. Each local environment is unique,and the site-specific conditions deter-mine which actinide species predomi-nate as well as each species’ overalltransport and migration characteristics.

In general, actinide solubilities arelow in most natural waters—below micromolar concentrations. Actinides in the V oxidation state will have thehighest solubilities; those in the IV oxidation 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 that favors the V oxidation state, and assuch it is the most soluble and mosteasily transportable actinide under envi-ronmental conditions. Plutonium favorsthe IV oxidation state in many naturalwaters, and normally its solubility isquite low. Once plutonium is in solu-tion, its migration is dominated by avariety of processes, including redoxreactions in solution or on mineral surfaces, sorption, colloidal transport,simple diffusion into rock interfaces, coprecipitation, and mineralization. Furthermore, increasing research activites indicated that microbes can affect 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 founda-tion in environmental chemistry, molec-ular science, and interfacial science andthe requisite spectroscopies, subsurfacetesting, and laboratory experiments. Butwe also need to interpret experimentalresults within the context of the multi-component natural system.

A case in point is the recent identifi-cation of a higher oxide of plutoniumof general formula PuO2+x(s), which isformed by water-catalyzed oxidation ofPuO2(s). As discussed by Haschke inthe article beginning on page 204, thisPu(IV) “superoxide” may also containPu(VI), which would increase plutoni-um solubility. In a recently publishedpaper (Haschke et al. 2000), the authorsspeculated that the Pu(VI) could dissolve and be transported more easilyfrom a nuclear waste storage site thanpreviously estimated. Further, they conjectured that Pu(VI) might havecaused the rapid migration of plutonium

at the Nevada Test Site. In these specu-lations, however, they failed to includeinfluential environmental factors.

While we acknowledge that the potential presence of Pu(VI) in aPuO2(s) matrix may increase the dissolution kinetics, the chemical condi-tions of 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 envi-ronments, Pu(VI) is unlikely to be stable. 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 PuO2(s). Consequently,the fact that the superoxide PuO2+x(s)may include Pu(VI) does not automati-cally imply enhanced plutonium solubil-ity or mobility in the natural environ-ment. Furthermore, given the results on colloidal transport, it is improbable that Pu(VI) is responsiblefor the plutonium migration at theNevada Test Site.

As the inventory of actinides in interim 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 increasingly important to understand how actinideswill interact with the environment.More sophisticated models are neededto account for all the potential migration paths away from an actinidesource. Theoretical and experimentalscientists will be challenged for yearsby the demands of developing thosemodels. �



Further Reading

Brown, 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 Diffraction Studies of Monomeric Actinide Tetra-, Penta-,and Hexavalent Carbonato Complexes. In Speciation, Techniques and Facilities for Radioactive Materials at Synchrotron Light Sources, OECD Nucl. Energy Agency. 121.

Choppin, G. R., and E. N. Rizkalla. 1994. Solution Chemistry of Actinides and Lanthanides. In Handbook on the Physics and Chemistry 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 the Environment. In Radiochemistry and Nuclear Chemistry. 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 of Transuranic Elements in Natural Aquatic Systems. In Handbook on the Physics and Chemistry of the Actinides. Edited by A. J. Freeman and G. H. Lander. New York: North-HollandPublishing Co.

Langmuir, D. 1997. Chap. 13. In Aqueous Environmental Geochemistry. Upper Saddle River, 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 Los Alamos. Conference on Waste Management 99, 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 of sorption of plutonium-239 onto colloids of iron oxides, clay and silica." Abstract of Agronomy, p.31 ASA, CAA, SSSA Annual Meetings, 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 Univer-sity of Munich, Germany, in 1990 and 1993,under the direction of Professor Dr. J. I. Kim.He then worked for two years as a scientific staff

member at theInstitute for Radiochemistry,TU Munich.During this time,he was invited towork as a guestscientist at San-dia National Lab-oratories on tech-nical aspects ofthe actinidesource term

model for WIPP. Wolfgang then spent a year aspostdoctoral fellow at the G. T. Seaborg Institutefor Transactinium Science at the Lawrence Livermore National Laboratory before joining the 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 Chemical Science and Technology Division. His researchinterests are in the areas of inorganic, environ-mental, and radiochemistry of f elements, including the actinides, with the focus on nuclear 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).

The Chemical Interactions of Actinides in the Environment

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