Journal of Nuclear and Radiochemical Sciences, Vol. 6, No. 1, pp. 31-37, 2005

Accepted
Interaction of Actinides with Carboxylates in Solution: Complexation of U(VI), Th(IV),and Nd(III) with Acetate at Variable Temperatures
Linfeng Rao,*, a PierLuigi Zanonato,b and Plinio Di BernardobaLawrence Berkeley National Laboratory, Berkeley, CA 94720, U.S.A.
Dipartimento di Scienze Chimiche, Università di Padova, via Marzolo 1, 35131, Padova, Italyb

Received: November 15, 2004; In Final Form: November 15, 2004
The complexation of uranium (VI), thorium (IV), and neodymium (III) with acetate was studied from 10 to 70 ˚C.
The formation constants and the enthalpies of complexation were determined by titration potentiometry andcalorimetry. The complexes become stronger at higher temperatures, despite that the enthalpy of complexation
becomes more positive and unfavorable to the complexation as the temperature is increased. Extended X-ray
Absorption Fine Structure Spectroscopy (EXAFS) and optical absorption spectroscopy provide additional information to support the thermodynamic results. The effect of temperature on the thermodynamic parameters is
discussed in terms of the electrostatic model.

temperature on the complexation of U(VI), Th(IV), and Nd(III)with acetate are discussed on a comparative basis.

1. Introduction
Study of the interactions of actinides with organic materials
remains an active and challenging subject due to recent activities in the environmental management of nuclear wastes. Manyorganic materials exist in the nuclear wastes (e.g. carboxylic
acids) or occur in the natural environment (e.g. humic and fulvic
materials, microorganisms). They form fairly stable complexes
with actinides in different oxidation states and reduce someactinides, particularly neptunium and plutonium, in higher
oxidation states. Previous studies show that the redox reactions
proceed through intramolecular electron transfer in the intermediate complexes between the actinides and organic materials.1, 2
As a result, the strength of complexation affects the rate ofredox reactions. Therefore, studies of the complexation ofactinides with organic materials are critically important topredict the chemical behavior of actinides in nuclear waste
processing and in the environment.
We have studied the complexation of actinides with a seriesof carboxylates at variable temperatures because of thefollowing reasons. 1) Many carboxylic acids are known toexist in the nuclear wastes and carboxylate is one of the mostimportant functional groups in natural organic materials thatparticipate in the complexation and redox reactions withactinides. Complexation of actinides with carboxylates affects
their behavior in nuclear waste processing as well as geological disposal. 2) The temperature in the nuclear wastes instorage tanks and the geological repository is significantly
higher than the ambient temperature due to the radioactive
decay energy. Because the majority of the data in the literatureon the complexation of actinides with carboxylates are
obtained at or near 25 ˚C,3 the thermodynamic parameters thatare needed to predict the behavior of actinides in waste
processing and disposal are unavailable. This paper focuses on
the studies of the complexation of U(VI), Th(IV), and Nd(III)with acetate in solution at 10 – 70 ˚C. The formation constants,
enthalpy and entropy of complexation for U(VI) acetate andNd(III) acetate are reviewed and recalculated on the molality
scale using a conversion approach different from that previously used.4, 5 Recent data on Th(IV) acetate complexes areincluded. The thermodynamic parameters and the effect of*Corresponding author. E-mail: LRao@lbl.gov. FAX: +1-510486-5596.

2. Experimental
Chemicals. All the concentrations in molarity in this paper
are referred to 25 ˚C. All chemicals are reagent grade or higher.
Distilled water was used in preparations of all solutions.
Descriptions on the preparation and standardization of the stock
solutions of U(VI), Th(IV), and Nd(III) were given elsewhere.4−6
Solutions of sodium acetate/acetic acid were prepared by adding
calculated amounts of sodium hydroxide into solutions of acetic
acid. The ionic strength of the solutions used in potentiometry
and calorimetry was 1.0 mol dm−3 (NaClO4) in the study of
U(VI) and Th(IV), and 2.0 mol dm−3 (NaClO4) in the study ofNd(III).
Procedures
Potentiometry. Details of the potentiometric titration setup
have been provided elsewhere.4−6 Electromotive force (EMF,
in millivolts) was measured with a Metrohm pH meter (Model713) equipped with a Ross combination pH electrode (Orion
Model 8102) filled with 1.0 mol dm−3 sodium perchlorate to
avoid clogging of the electrode frit glass septum due to the low
solubility of KClO4. The EMF of the glass electrode in theacidic region can be expressed by eq 1.
E = E 0 + RT/F ln[H+] + γH[H+](1)

where R is the gas constant, F is the Faraday constant and T is
the temperature in K. The last term is the electrode junction
potential for the hydrogen ion (∆Ej,H+), assumed to be proportional to the concentration of the hydrogen ion. Prior to eachtitration, an acid/base titration with standard perchloric acidand sodium hydroxide was performed to obtain the electrodeparameters of E 0 and γH. These parameters allowed the calculation of hydrogen ion concentrations from the EMF’s in thesubsequent titration.
Fifty to seventy data points were collected in each titration.
Multiple titrations were conducted at each temperature withsolutions of different concentrations of the metal cations andperchloric acid. The complexation constants on the molarity
scale were calculated with the program Hyperquad.7
Calorimetry. Details of the calorimeter and the calorimetrictitrations have been given previously.4 Usually, 50 – 70 values© 2005 The Japan Society of Nuclear and Radiochemical Sciences
Published on Web 07/31/2005

32

Rao

J. Nucl. Radiochem. Sci., Vol. 6, No. 1, 2005of experimental heat were obtained in each titration. The net
heat of complexation was determined after correction for theheat of dilution and the heat due to acid/base neutralization.
The enthalpy of metal acetate complexation was calculated by
the computer program Letagrop.8
Optical Absorption Spectroscopy. Absorption spectra ofNd(III) acetate solutions were obtained on a Varian Cary-5G
spectrophotometer equipped with a 1 × 1 peltier automatic
temperature controller. Quartz cells with 1.0 cm optical length
were used.
Extended X-ray Absorption Fine Structure (EXAFS) Spectroscopy. Uranium and thorium LIII-edge EXAFS spectra (up to
kmax ~ 15 Å−1) were collected on solutions of U(VI) and Th(IV)with different concentrations of acetate. The conditions of thesolutions were selected, based on speciation calculation withthe formation constants from this work, so that they contained
different complexes as dominant species. The average ligand– ranged from 0 to 3.0 for U(VI) acetate and from 0
number, n,
to 4.0 for Th(IV) acetate. Three scans were performed for eachsample. Energy calibration was based on assigning the firstinflection point of the absorption edge for UO2 to 17166 eV andfor ThO2 to 16300 eV. The EXAFS spectra were fit with the
R-space X-ray Absorption Package (RSXAP),9 using parameterized phase and amplitude functions generated by the program FEFF810 with the reference crystal structures including
UO2(OOCCH3)2·2H2O and Na[UO2(OOCCH3)3],11−13 ThO2 andNa6Th(CO3)5(H2O)12.14 Standard scattering paths were calculated
from the reference structures and included in the data analysis.
For the U(VI) data, the standard scattering paths include thesingle scattering U-Oax (axial oxygen), U-Oeq (equatorial oxygen)
and U-C, and the multiple scattering O=U=O (axial oxygens).
For the Th(IV) data, the standard scattering paths include thesingle scattering Th-O, Th-C, and Th-Th.
Conversion of stability constants on the molarity scale to themolality scale. To allow the comparison at different temperatures, the constants in molarity (β M ) were converted to theconstants in molality (βm) by eq 2.15
log10 βm = log10 βM + Σ rνr log10 ϑ(2)where ϑ is the ratio of the values of molality to molarity forthe ϑ specific ionic medium and is equal to 1.05 and 1.10 dm3
of solution per kg of water for 1.0 and 2.0 mol dm−3 NaClO4.
Σ rνr is the stoichiometric coefficient of the reaction (νr is positive for products and negative for reactants) and is equal to – jm+for the metal acetate complexation reactions (M + jAc− =[M(Ac)j](m−j)+ where Mm+ represents UO22+, Th4+, and Nd3+,
respectively). Therefore, the corrections for I = 1.0 and = 2.0
mol dm−3 NaClO4 are performed by eq 3a and 3b, respectively.
log10 βm, j = log10 βM, j – 0.02 × j(3a)log10 βm, j = log10 βM, j – 0.04 × j(3b)

3. Results and Discussion
U(VI) acetate
Thermodynamic parameters at 25 – 70 ˚C. The data fromboth potentiometry and calorimetry are consistent with theformation of three consecutive U(VI) acetate complexes
defined by equilibrium 4.
UO22+ + j CH3COO−

UO2(OOCCH3)j(2−j)+ j = 1, 2, 3

βj = [UO2(OOCCH3)j(2−j)+]/([UO22+][CH3COO−] j )(4)

The equilibrium constants and the corresponding Gibbs freeenergy, enthalpy and entropy of complexation are given in
Table 1. The values of log βM and ∆H were obtained by potentiometry and calorimetry, respectively. The values of log βmwere recalculated from the experimental log βΜ using eq 3a.
As shown by the data in Table 1, the overall stability constantsof all three U(VI) acetate complexes increase as the temperature is elevated.
Structures of U(VI) acetate complexes in solution. The Fourier
Transforms of the background-subtracted uranium L III-edge
EXAFS spectra for three U(VI) acetate solutions are shown in
Figure 1. Based on the speciation calculation using the stability– is equal toconstants in Table 1, the average ligand number, n,
1, 2, and 3 for Solutions I, II and III, respectively. The 1:1 and
1:3 complexes, UO2(OOCCH3)+ and UO2(OOCCH3)3−, are thedominant species in Solutions I and III, respectively. However,
Solution II contains a mixture of three complexes (1:1, 1:2,
and 1:3) in comparable amounts without a dominant species.
Systematic changes in the Fourier Transforms of the
EXAFS spectra are observed in the region around 2 Å (before
the correction for the phase shift). From the 1:1 uranyl-acetate
complex (Solution I), through the mixture (Solution II), to the
1:3 complex (Solution III), the broad peak at ~ 2 Å splits into
two at 1.8 and 2.15 Å, respectively. The feature at 2.15 Å
becomes more significant for Solution III, suggesting that

TABLE 1: Formation Constants and Corresponding Thermodynamic Parameters for Uranyl Acetate Complexation, I = 1.05
mol kg–1 (NaClO4)t˚C

log βM

UO22+ + Ac− = UO2Ac+

25
35
45
55
70

2.58 ± 0.03
2.67 ± 0.02
2.74 ± 0.03
2.85 ± 0.02
2.98 ± 0.05

UO22+ + 2Ac− = UO2Ac2

25
35
45
55
70

UO22+ + 3Ac− = UO2Ac3−

25
35
45
55
70

The error limits represent 3σ

log βm∆G
kJ mol−1∆H
kJ mol−1∆S
J K mol−1

2.56 ± 0.03
2.65 ± 0.02
2.72 ± 0.03
2.83 ± 0.02
2.96 ± 0.05

-14.61-15.63-16.57-17.78-19.4410.6 ± 0.8
11.8 ± 0.5
13.0 ± 0.6
14.3 ± 0.5
15.4 ± 0.7

85 ± 389 ± 293 ± 298 ± 2102 ± 34.37 ± 0.14
4.60 ± 0.08
4.76 ± 0.10
4.94 ± 0.06
5.27 ± 0.09

4.33 ± 0.14
4.56 ± 0.08
4.72 ± 0.10
4.90 ± 0.06
5.23 ± 0.09

-24.71-26.90-28.75-30.78-34.3620 ± 321 ± 222 ± 224 ± 127 ± 1150 ± 13155 ± 8160 ± 8167 ± 4179 ± 5

6.86 ± 0.04
7.11 ± 0.03
7.23 ± 0.04
7.38 ± 0.03
7.62 ± 0.06

6.80 ± 0.04
7.05 ± 0.03
7.17 ± 0.04
7.32 ± 0.03
7.56 ± 0.06

-38.81-41.59-43.67-45.98-49.6617.5 ± 0.6
18.8 ± 0.3
20.8 ± 0.3
22.8 ± 0.3
24.6 ± 0.4

189 ± 3196 ± 2203 ± 2210 ± 2216 ± 2

−1

J. Nucl. Radiochem. Sci., Vol. 6, No. 1, 2005

Interaction of Actinides with Carboxylates in Solution

33

Figure 2. Proposed structures of uranyl acetate complexes in solution.
(a) UO2(OOCCH3)+. (b) UO2(OOCCH3)2(aq). (c) UO2(OOCCH3)3−.

Figure 1. Experimental (solid lines) and fitted (dashed lines) uranium
LIII-edge EXAFS data. (I) Solution I, n– = 1, UO2(OOCCH3)+ dominant; (II) Solution II, n– = 2, no dominant species; (III) Solution III, n– =3, UO2(OOCCH3)3− dominant.

more oxygens and/or carbons are found in a distant shell in the
1:3 complex than in the 1:1 complex. The best-fit parameters
are summarized in Table 2.
As Denecke et al. showed,16 the bidentate and unidentate
acetate in the U(VI) complex can be differentiated by two
features in EXAFS: (1) the RU-O is longer in the former (~ 2.45 −2.48 Å) than in the latter (~ 2.35 Å); (2) the RU-C in the former
is ~ 2.85 − 2.88 Å, easily identified by EXAFS. However,detecting the carbon in a unidentate acetate may be difficult
due to its longer distance (RU-C 3.5 Å) and the overlap of oscillations resulting from the U-C single scattering and the linear
O=U=O multiple scattering (1.78 × 2 Å). Therefore, data in
Table 2 suggest the following: 1) The acetate ligand in the 1: 1
uranyl acetate complex (Solution I) is bidentate to uranium(two oxygens at 2.50 Å and one carbon at 2.91 Å). 2) In the
TABLE 2: Best Fit Parameters for Uranium LIII-Edge EXAFS
Shell

R,a Å

Na

σ, bÅ

∆E0, eV

Solution I
U-Oax
1:1 Uranyl/acetate U-Oeq1
U-Oeq2
U-C

1.78
2.38
2.5
2.91

2.0
4.0
2.0
1.3

0.04110.07030.09200.0500-14.48-14.48-14.48-14.48

Solution II
Mixture

U-Oax
U-Oeq
U-C

1.78
2.42
2.9

2.0
5.9
2.2

0.03700.08880.0500-12.45-12.45-12.45

Solution III
U-Oax
1:3 Uranyl/acetate U-Oeq1
U-Oeq2
U-C

1.78
2.34
2.48
2.87

2.0
1.9
4.1
2.1

0.03440.05330.04820.0500-12.37-12.37-12.37-12.37

Samplesa

The 95% confidence limits for the bond lengths (R) and coordination numbers (N) for each shell are: U-Oax, 0.01 Å and ± 15%;
U-Oeq, 0.02 Å and ± 25%; U-C, 0.02 Å and ± 25%, respectively.
b
σ is the EXAFS Debye-Waller term, which accounts for theeffects of thermal and static disorder through damping of the
EXAFS oscillations by the factor exp(−2k 2 σ 2 ).

1:3 complex (Solution III), two acetates are bidentate (fouroxygens at 2.48 Å and two carbons at 2.87 Å) and the third
acetate is unidentate. One of the two oxygens at 2.34 Å isfrom the unidentate acetate while the other is from water. 3)
Because no single species is dominant in Solution II, wecannot unambiguously assign the structure of the 1:2 complex
based on the EXAFS data. However, the two carbons at 2.90
Å implies that there are two bidentate acetates per uranium by
average in Solution II. Based on the EXAFS data, the structures of U(VI) acetate complexes in solution are proposed in
Figure 2.
Thermodynamic trends and the structure of complexes insolution. The stepwise enthalpy and entropy of complexation
can be calculated from the data in Table 1. The trends are: (1)The first two steps are endothermic but the third is exothermic
and the enthalpy of complexation decreases in the order: ∆H1(10 to 15 kJ mol−1) > ∆H2 (9 to 12 kJ mol−1) >> ∆H3 (about −2
kJ mol −1 ). (2) The entropy change becomes smaller in theorder: ∆S1 (85 to 102 J K−1 mol−1) > ∆S2 (65 to 77 J K−1 mol−1)> ∆S3 (37 to 39 J K−1 mo1−1) and ∆S3 is less than 50% of ∆S1.
These trends can be rationalized in terms of the perturbation in
the primary hydration sphere and the bulk structure of water
upon complex formation. Since the degree of charge neutralization decreases as the stepwise complexation progresses, theenergy required for restructuring the bulk solvent is expected
to become less in successive steps. Thus the stepwise enthalpy
change decreases in the order: ∆H1 > ∆H2 > ∆H3. Furthermore,
the unidentate acetate in the third complex (Figure 2c) replaces
only one water molecule in the primary hydration sphere sothat it requires less desolvation energy than a bidentate coordination. Consequently, the enthalpy change in the third step ismuch less than the first two and, in this particular case, thestepwise formation of the third complex becomes exothermic.
Similarly, the trend in entropy can be interpreted based on
the structural information of the uranyl acetate complexes insolution. In general, the entropy change upon the complexation consists of ∆St (translational), ∆Sr (rotational), and ∆Sc(conformational). The change in vibrational entropy (∆Sv) can
usually be neglected. Quantitative evaluation of the entropy
effect in complexation is difficult because the quantities ∆St,
∆Sr , and ∆Sc are not easy to assess and the information on thesolvation of the ligand is rarely available. However, integration of the thermodynamic data and structural information
obtained in this work allows qualitative discussions. A larger
gain in the translational entropy is certainly expected for abidentate complex (steps 1 and 2) than a unidentate complex(step 3), because the former replaces more water moleculesfrom the primary hydration sphere. Besides, higher degree ofcharge neutralization in steps 1 and 2 causes larger net increase

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Rao

J. Nucl. Radiochem. Sci., Vol. 6, No. 1, 2005

TABLE 3: Equilibrium Constants and Corresponding Thermodynamic Parameters for the Formation of Th(IV) Acetate
Complexest˚C

log βMlog βm∆G
kJ mol−1∆H
kJ mol−1∆S
J K−1 mol−1

Th4+ + Ac− = ThAc3+

10
25
40
55
70

3.71 ± 0.03
3.81 ± 0.02
3.96 ± 0.02
4.11 ± 0.05
4.30 ± 0.02

3.69 ± 0.03
3.79 ± 0.02
3.94 ± 0.02
4.09 ± 0.05
4.28 ± 0.02

-20.0
-21.6
-23.6
-25.7
-28.1

10.2± 0.2312.2± 0.18
15.2± 0.75
17.3 ± 0.8
17.0 ± 0.9

107 ± 1114 ± 1124 ± 2131 ± 3131 ± 3Th4+ + 2Ac- = Th(Ac)22+

10
25
40
55
70

6.76 ± 0.02
6.83± 0.02
6.94 ± 0.04
7.31 ± 0.06
7.67 ± 0.02

6.72 ± 0.02
6.79 ± 0.02
6.90 ± 0.04
7.27 ± 0.06
7.63 ± 0.02

-36.4
-38.8
-41.4
-45.7
-50.1

12.7± 0.33
16.9± 0.27
20.7 ± 1.4
25.5 ± 1.3
28.3 ± 1.4

174 ± 1187 ± 1
198 ± 5217 ± 4229 ± 4Th4+ + 3Ac− = Th(Ac)3+

10
25
40
55
70

8.58 ± 0.06
8.77± 0.13
9.05 ± 0.07
9.41± 0.15
9.75 ± 0.13

8.52 ± 0.06
8.71 ± 0.13
8.99 ± 0.07
9.35 ± 0.15
9.69 ± 0.17

-46.2
-49.7
-53.9
-58.7
-63.7

17.5 ± 1.1
30.8 ± 0.7
35.4 ± 2.6
39.2 ± 3.7
38.9 ± 4.7

225 ± 4270 ± 3285 ± 8298 ± 12299 ± 10

Th4+ + 4Ac− = Th(Ac)4

10
25
40
55
70

10.12 ± 0.12
10.25 ± 0.19
10.51 ± 0.12
11.35 ± 0.11
11.54 ± 0.20

10.04 ± 0.12
10.17 ± 0.25
10.43 ± 0.18
11.27 ± 0.11
11.46 ± 0.2

-54.4
-58.0
-62.5
-70.8
-75.3

29.1 ± 1.7
30.7 ± 1.4
44.8 ± 4.5
65.9 ± 3.1
99.5 ± 6.1

295 ± 6298 ± 7343 ± 6417 ± 10509 ± 18

Th4+ + 5Ac− = Th(Ac)5−

10
25
40
55
70

11.37 ± 0.15
11.51 ± 0.22
11.69 ± 0.21
12.48 ± 0.28
12.94 ± 0.24

11.27 ± 0.15
11.41 ± 0.22
11.59 ± 0.21
12.38 ± 0.28
12.84 ± 0.24

-61.1
-65.1
-69.5
-77.8
-84.3

34.2 ± 1.2
38.9 ± 0.8
39.9 ± 3.5
34.2 ± 3.2
40.1 ± 4.2

337 ± 5349 ± 5349 ± 12341 ± 11363 ± 13

I = 1.05 mol kg–1 NaClO4. The error limits represent 3σ

in the disorder of the bulk water. In addition, it is reasonable
to assume that the structuring effect of a unidentate acetate on
the bulk water could be stronger than a bidentate acetate
because the former is more capable of forming hydrogen bonds
with bulk water due to its “free” oxygen (Figure 2c). The
combination of all these effects results in a much smaller ∆S3
and the decreasing order of the stepwise entropy of complexation: ∆S1 > ∆S2 >> ∆S3.
Th(IV) acetate
Thermodynamic parameters at 10 – 70 ˚C. The data fromboth potentiometry and calorimetry are consistent with the formation of five consecutive Th(IV) acetate complexes. The
equilibrium constants and the corresponding Gibbs freeenergy, enthalpy and entropy of complexation are given in
Table 3.
EXAFS. The EXAFS spectra and associated Fourier Transforms of six Th(IV) solutions are shown in Figure 3. The valuesof n– range from 0 to 4. The Fourier Transforms show prominent features in the region around 2 Å that represent the oxygen shell of coordinating water molecules or carboxylategroups of the acetate. The peak around 2 Å broadens from
Solution I to VI as n– is increased from 0 to 4, suggesting a single oxygen shell for Solution I ( n– = 0) but multiple oxygen
shells for Solution VI (n– = 4). Best fit of the EXAFS data indicate that the spectrum for Solution I is represented by a single
Th-O shell with ~ 10 oxygen atoms at 2.45 Å, while the spectrum for Solution VI is represented by two distinct Th-O
shells: 2 – 3 oxygen atoms at 2.37 Å and 8 oxygen atoms at
2.52 Å. The spectra for Solutions I I – V are represented by
gradual transition from a single Th-O shell (Solution I) to two
Th-O shells (Solution VI).
Conclusive assignments of the two oxygen shells for Solution
VI are difficult since there are no data on the crystal structure

Figure 3. Experimental (solid line) and fitted (dotted line) thorium
LIII-edge EXAFS spectra (right) and associated Fourier Transforms(left). (I-VI): Solutions I – VI, n– = 0, 1.0, 1.5, 2.0, 3.0, 4.0, respectively.

of Th(IV) acetate complexes. However, the most reasonableassignments can be made in comparison with the structural
information on U(VI) acetate complexes in solution and crystalline Na6Th(CO3)5·12H2O.14 As discussed in previous sections,acetates coordinate to the linear UO22+ cation in the equatorial
plane with either bidentate or unidentate modes. The RU-Oeq forthe bidentate acetate is 2.45 – 2.48 Å, longer than the RU-Oeq for
the unidentate acetate or the solvating water (~ 2.35 Å).4, 16
Since the effective ionic radius of Th(IV) is larger than that ofU(VI),17 the RTh-O for bidentate or unidentate acetate should beslightly longer than corresponding RU-O. This appears to be inagreement with the structural data of Na6Th(CO3)5·12H2O(c)where the five carbonate ligands are bidentate with the average

J. Nucl. Radiochem. Sci., Vol. 6, No. 1, 2005

Interaction of Actinides with Carboxylates in Solution

35

TABLE 4: Formation Constants and Corresponding Thermodynamic Parameters for Nd(III) Acetate Complexationt
˚C

log βMlog βm∆G
kJ mol−1∆H
kJ mol−1∆S
J K−1 mol−1

Nd3+ + Ac− = NdAc2+

25
45
70

1.92 ± 0.04
2.05 ± 0.03
2.24 ± 0.01

1.88 ± 0.04
2.01 ± 0.03
2.20 ± 0.01

-10.73-12.24-14.457.1 ± 0.03
9.1 ± 0.3
11.7 ± 0.2

59.8 ± 0.4
67 ± 276 ± 1

Nd3+ + 2Ac− = NdAc2+

25
45
70

3.02 ± 0.05
3.40 ± 0.05
3.65 ± 0.03

2.94 ± 0.05
3.32 ± 0.05
3.57 ± 0.03

-16.78-20.2-23.4

14.6 ± 0.16
17.5 ± 0.4
20.1 ± 0.4

105 ± 2119 ± 3127 ± 2

Nd3+ + 3Ac− = NdAc3

25
45
70

3.58 ± 0.08
4.36 ± 0.09
4.71 ± 0.10

3.46 ± 0.08
4.24 ± 0.09
4.59 ± 0.10

-19.75-26.9-30.2

18.4 ± 0.46
23.6 ± 0.5
29.2 ± 0.5

128 ± 4159 ± 4173 ± 3

I = 2.20 mol kg−1 NaClO4. The error limits represent 3σ

Figure 5. Normalized peak area for the 4I9/2 → 4G5/2, 2G7/2 transition ofNd(III) vs. Cacetate/CNd (a) and n– (b) at 25 and 70 ˚C.

Figure 4. Absorption spectra of Nd/acetate solutions (the hypersensitive 4I9/2 → 4G5/2, 2G7/2 transition) at 25 and 70 ˚C. Optical length: 1.0
cm, pH = 4.7, [Nd(III)]total = 0.025 mol dm−3, the ratios of Cacetate/CNd
for spectra 1 – 6 are 0, 4, 8, 16, 28, and 40, respectively.

RTh-O = 2.505 Å.14 Based on these results, the eight oxygens at
2.52 Å for Solution VI can be most reasonably assigned to theoxygens from bidentate acetate ligands, corresponding to four
acetates by average for every Th 4+ . Then the 2 − 3 oxygen
atoms at 2.37 Å can be reasonably assigned to the solvatingwater molecules, making a total coordination number of 10 – 11
for Th4+. It should be pointed out that the EXAFS data alone
do not provide compelling evidence on the structures of thecomplexes in solution, but they are consistent with the thermodynamic results that suggest a value of n– = 4 for Solution VI.
Nd(III) acetate
Thermodynamic parameters at 25 – 70 ˚C. The equilibrium
constants and the corresponding Gibbs free energy, enthalpy
and entropy of complexation are given in Table 4. The valuesof log βm were recalculated from the experimental log βM using
eq 3b.
Absorption Spectra at 25 and 70 ˚C. The absorption spectraof Nd(III) acetate solutions (CAcetate/CNd = 0, 4, 8, 16, 28, and40) at 25 and 70 ˚C are shown in Figure 4. The data indicatethat the 4I9/2 → 4G5/2, 2G7/2 transition was significantly intensified as CAcetate/CNd was increased from 0 to 40, in a similar

fashion at both temperatures. Attempts to deconvolute thespectra and calculate the oscillator strengths for individual
species were not successful. However, when the peak area
(normalized against the peak area in the absence of acetate)
was plotted against CAcetate/CNd, the data did indicate that theintensity of the 4I9/2 → 4G5/2, 2G7/2 transition was enhanced by thecomplexation more significantly at 70 ˚C than at 25 ˚C (Figure5a). Based on the generalization by Henrie et al.,18 this difference
implies that there are a greater number of coordinated acetate
ligands around Nd3+ at 70 ˚C than at 25 ˚C when the ratio ofCAcetate/CNd is constant. In other words, the spectrometric datasuggest that Nd(III)/acetate complexation is stronger at 70 ˚Cthan at 25 ˚C. This is consistent with the trend in the thermodynamic constants obtained by potentiometry (Table 4). Figure 5b
indicates that the peak area is proportional to the average numberof acetate ligands around Nd3+ at both temperatures.
The effect of temperature on the complexation of U(VI),
Th(IV), and Nd(III)The stability of complexes. Data in Tables 1, 3, and 4 showthat the U(VI), Th(IV), and Nd(III) acetate complexes become
stronger as the temperature is increased. As discussed inprevious publications, the effect of temperature on the formation constants of the complexes between a hard acid (e.g.
actinide and lanthanide cations) and a hard base (e.g. oxygen
donor ligands such as acetate) could be interpreted with anelectrostatic model. Combining a modified Born equation 19
and the empirical expression for the dielectric constant ofwater (ε = ε0exp(−T/ϑ), where ε0 = 305.7; ϑ = 219 K),20, 21 thetemperature coefficient of the complexation constant (β) is
expressed as∂ (log β)/∂ T = Ne2Z1Z2/(0.2303Rd12)(1/T − 1/ϑ)/(εT)(5)

where the symbols are explained in Reference 19. Since ϑ isfar below the freezing point of water, T is always higher than ϑin the whole accessible temperature range of an aqueous solu-

36

Rao

J. Nucl. Radiochem. Sci., Vol. 6, No. 1, 2005

tion. As a result, (1/T – 1/ϑ) is always negative. Thus, theelectrostatic model predicts that the complexation between
species of opposite charges is enhanced by the increase intemperature and that the magnitude of enhancement is proportional to Z1Z2. These predictions appear to agree with thedata on the formation of 1:1 complexes of U(VI), Th(IV), andNd(III) with acetate, as shown in Table 5. The agreement
implies that the complexation between actinides and acetate isdominantly electrostatic interactions.
TABLE 5: Temperature Effect on the Formation Constants of
Acetate Complexes
Reaction

Z1Z2

β70/β25a

Nd3+ + Ac− → Nd(Ac)2+

-3

2

2+

−

+

4+

−

3+

UO2 + Ac → UO2(Ac)Th + Ac → Th(Ac)
a
b

b

-3.2

2.5

-4

2.8

Taken as a measure of the magnitude of temperature effect.
Based on the effective charge on UO22+ in the literature.19, 22

Enthalpy and entropy of complexation. Data in Tables 1, 3,
and 4 indicate that the overall complexation of U(VI), Th(IV),and Nd(III) with acetate is endothermic and is driven by large
and positive entropies, consistent with the general observation
for the interactions between “hard acid” cations and “hardbase” anions. As shown in Figure 6, though the enthalpies
become more positive and unfavorable to the complexation at
higher temperatures, the complexes are more stable because ofincreasingly more positive entropy of complexation. These
trends can be related to the perturbation of the solvent structureby thermal motions that results in a more disordered bulk
water structure at higher temperatures. In the process ofcomplexation, the highly ordered solvating water molecules
are released to a more disordered bulk solvent. As a result, thegain in the entropy of complexation is larger at higher temperatures. This effect has been discussed in more detail elsewhere.5, 23, 24
Figure 6 shows that the overall enthalpies of formation of all
the acetate complexes, except Th(Ac)4(aq), can be fitted withlinear functions of the temperature, suggesting that the heat
capacities of formation of these complexes are constant from25 to 70 ˚C. From the slopes of the linear fits, the overall heat
capacities of complexation (J K−1 mol−1) are calculated to be:(110 ± 19) for UO2Ac+, (172 ± 48) for UO2(Ac)2(aq), and (170
± 12) for UO2(Ac)3−; (125 ± 22) for ThAc3+, (265 ± 11) forTh(Ac)22+, (341 ± 94) for Th(Ac)3+, and (46 ± 68) for Th(Ac)5−;and (102 ± 13) for NdAc2+, (122 ± 19) for Nd(Ac)2+, and (239± 27) for Nd(Ac)3(aq). Unlike these complexes, the enthalpy
of formation of Th(Ac)4(aq) drastically increases in a non-linear
manner when the temperature is increased from 10 to 70 ˚C. Thissuggests that heat capacity of formation of Th(Ac)4(aq) is large
and temperature-dependent. This may not be surprising since
the stepwise formation of Th(Ac)4(aq) is a reaction of completecharge neutralization:
Th(Ac)3+ + Ac− → Th(Ac)4(aq)(6)

The heat capacity of a reaction consists of electrostatic andnon-electrostatic contributions and the former is expected to be
dominant in the reaction between a “hard acid” and a “hardbase”. For an isocolumbic reaction where the sign and magnitude of the electrical charges on the reactants and the products
are identical, the electrostatic contributions from the reactants
and products may balance out so that the heat capacity is smallor near zero, and independent of temperature.25 However, forthe reactions where the electrical charges of the reactants andproducts are drastically different, large heat capacity changes
are often associated. This could be particularly true for thereactions of complete charge neutralization such as Reaction 6.
In fact, the heat capacities of the formation of Nd(Ac)3(aq) andUO2(Ac)2(aq), also reactions of complete charge neutralization,
are found to be large.
4. Summary
The complexation of U(VI), Th(IV), and Nd(III) with acetate
becomes stronger as the temperature is increased from 10 to 70
˚C, despite that the enthalpy of complexation becomes more
endothermic and unfavorable to the complexation at higher
temperatures. The enhancement of the complexation is mainly
due to a larger entropy effect at higher temperatures and can be
explained by the effect of temperature on the solvent structure
and a simple electrostatic model.

Figure 6. Overall thermodynamic parameters of U(VI), Th(IV) and Nd(III) acetate complexation at different temperatures, Symbols: (1) ML; (2)ML2; (3) ML3; (4) ML4; (5) ML5.

Interaction of Actinides with Carboxylates in Solution

Acknowledgments. This work was supported by the Director,
Office of Science, Office of Basic Energy Sciences and Office of
Biological and Environmental Research under U.S. Departmentof Energy Contract No. DE-AC03-76SF0098 at Lawrence
Berkeley National Laboratory, and by the Ministero dell'Università e della Ricerca Scientifica e Tecnologica (MURST,
Roma) within the program COFIN02. The EXAFS experiments were conducted at SSRL, which is operated by the
Department of Energy, Division of Chemical Sciences.
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