Water-RockInteractions, Ore Deposits, and Environmental Geochemistry:
A Tribute© The Geochemical Society, Special Publication No.7, 2002
Editors: Roland Hellmann and Scott A. Wood

to David

A Crerar

The aqueous geochemistry of the rare earth elements and yttrium. Part X.
Potentiometric determination of stability constants of acetate complexes of
La3+, Nd3+, Gd3+ and Yb3+ at 25-70 °C and 1 bar
Ru Ding AND Scott A. WOOD *Department of Geological Sciences, University of Idaho, Moscow, ID 83844-3022

USA* corresponding author. E-mail: swood@iron.mines.uidaho.edu
Abstract - The stability constants of acetate (Ac) complexes of four trivalent REE (rare earth element) ions
(La3+, Nd3+, Gd3+ and Yb3+) were determined in aqueous NaCl solutions at temperatures from 25-70 "C andionic strengths from 0.1 to 2.0 mol kg! via a potentiometric method. At the relatively low total acetate concentrations employed in this study (~0.02 m), only the first acetate complex, LnAc2+ (where Ln3+ represents
any trivalent REE ion) was required to fit the data. The data obtained are adequately described by the following equations:log Q)(La3+) = 7.081-

1326ff + 125.2·1ff - 1.892xlO-4·T·exp(l)- 6·Ay.11I2/(1+p/2)

log Q)(Nd3+) = 6.671-

1l08ff+ 141.5·1ff - 1.856x10-4·T-exp(l)- 6-A.y'11I2/Cl+p/2)

log Q)(Gd3+) = 7.092 -1258ff

2.239x10-4·T·exp(l)- 6'Ay.P/2/(1+11I2)

log Q)(Yb3+) = 8.196 - 1659ff + 137.4·1ff -1.146x10-4-T-exp(l)- 6.Ay'11I2/(1+p/2)+ 166.5·1ff -where Q, is the conditional stability constant for the association reaction Ln3+ + Ac- H LnAc2+, I is the ionicstrength in mol kg-', T is the temperature in Kelvin and Ay is the Debye- Hiickellimiting slope. Conditional
stability constants measured in this study at 25 "C and at ionic strengths ~ 1.0 mol kg-' were in good agreement (generally less than ±0.210g unit difference) with most values reported previously in the literature for
NaC104 media, suggesting that competition from chloride complexation was negligible. However, stability
constants determined in NaCl solutions of higher concentrations (2.0 mol kg+) were typically lower (bymore than 0.2 log unit for La3+) than those reported in the literature for NaCl04 solutions of the same ionicstrength, suggesting that the influence of chloride complexation is significant at this N aCl concentration.
Stability constants obtained in this study were extrapolated to 250 "C using an isocoulombic approximation, and were found to be in fair agreement (within ±O.210g unit) with the few experimental data available atelevated temperatures. However, our extrapolated values are orders of magnitude greater than previously
published theoretical estimates of these constants at 100-250 "C. Our stability constants, combined with additional thermodynamic data from the literature, should permit quantitative modeling of the influence of acetate on REE mass transfer in a variety of geologic environments and fine-tuning of models designed to
predict REE-acetate stability constants at elevated temperatures and pressures.

1. INTRODUCTION

turbance of Nd isotope and REE systematics by fluidmediated processes in sedimentary rocks (Bock et al.,
1994; McDaniel et al., 1994; Ohr et al., 1994; Kontak
and Jackson, 1995). Mobility, and particularly fractionation, of the REE in aqueous fluids can greatly
complicate the use of these elements as tracers. Assessment of the degree of mass transfer of the REE byaqueous fluids in any given geological environment
requires accurate stability constants for relevant REE
complexes and solubility products for relevant REEbearing phases, but such data are often unavailable orof uncertain quality.
The mobility of the REE in hydrothermal solutions
also is relevant to the safe disposal and effective remediation of nuclear waste materials in geologicalenvironments. The REE are produced during the fission of uranium and plutonium (Rard 1988), and are

Rare earth elements (REE) and Nd isotopes increasingly are being employed as tracers of processes
occurring during the formation and diagenesis of sedimentary rocks, and the genesis of petroleum andbase-metal deposits (e.g., Mississippi Valley-Type)that these rocks may contain (Scherer and Seitz,
1980; Graf, 1984; Banner et al., 1988a,b; Dorobek
and Filby, 1988; McLennan, 1989; Bock et al., 1994;
McDaniel et al., 1994; Ohr et al., 1994; Kontak andJackson, 1995). The REE are usually considered to berelatively immobile, and it is generally held that normalized REE patterns and Nd isotope ratios, which
are used widely as petrogenetic tracers, are comparatively immune to changes during weathering, diagenesis, metamorphismand hydrothermal alteration.
However, a number of studies have documented dis209

210

R. Ding and S. A. Woodanalogues of trivalent elements of the actinide series(Moeller et a1. 1965; Choppin 1983). Moreover, amethod which is being evaluated for disposal of Pu
from dismantled nuclear warheads involves the production of a waste form containing Pu and the REE
gadolinium (Gd) which acts as a neutron moderator toprevent criticality. Thus, data are required to evaluate
the solubility of Gd and Pu under the P- T conditionsof burial of the waste.
Acetate occurs in high concentrations in some oilfield brines (up to 10,000 ppm; Willey et al., 1975;
Carothers and Kharaka, 1978) and in sediments affected by seafloor hydrothermal fluids (Martens,
1990). It is a product of thermal degradation of more
complex organic matter (Surdam et al., 1984; Kawamura et al., 1986), and can persist metastably in aqueous solution for geologically long periods of time attemperatures between 100 and 200°C (Kharaka et al.,
1983; Drummond and Palmer, 1986; Palmer andDrummond, 1986; Bell et al., 1993). Acetate also is
produced by corn roots and the roots of grasses andherbs, and occurs in soil solutions in concentrationsof 2-5 x 10-3 mol L-l (Bruckert, 1970; Sposito, 1989).
Carboxylate ligands such as acetate may also be present in mixed nuclear-organic wastes such as those occurring at Oak Ridge, Tennessee and Maxey Flats,
Kentucky (cf. Means et al., 1978; Cleveland andRees, 1981). Furthermore, acetate can be used as ananalogue of monocarboxylate groups in general, bothin individual low-molecular weight ligands and incomplex humic and fulvic acids (Wood, 1996; 2000).
The points discussed in the preceding paragraphsuggest that acetate is a potentially important ligand
for the transport of metals in geological environments, and acetate complexes of the following metals
have been studied at elevated temperatures: A13+
(Fein, 1991; Benezeth et al., 1994; Palmer and Bell,1994), Zn2+ (Hennet et al., 1988; Yang et al., 1989;
Giordano and Drummond, 1991), Fe2+ (Palmer andDrummond, 1986; 1988; Palmer and Hyde, 1993),
Mg2+ (Semmler et al., 1990; Fein, 1991), Ca2+ (Fein,
1991; Seewald and Seyfried, 1991) and Pb2+ (Hennet
et al., 1988; Giordano, 1989; Yang et al., 1989). Acetate complexes of the REE have been investigatedusing potentiometric, spectrophotometric, calorimetric, ion-exchange and solvent extraction techniques(e.g., Sones son, 1958a,b; 1959; 1960; Grenthe, 1962;
1964; Manning and Monk, 1962; Kolat and Powell,1962; Kovar and Powell, 1966; Archer and Monk,
1964; 1966; Choppin and Schneider, 1970; Antipenko et al., 1973; Gruzdev and Ermolaev, 1974;
Bukietynska et al., 1977; 1981a,b; Daniele et al.,
1985; Ilnitskaya and Ternovaya, 1988), and thesestudies are in reasonably good agreement (cf. review
by Wood, 1993). However, the above studies are re-stricted to low temperatures (20-25 °C) and, although
they collectively cover a range of ionic strength from
<0.1 to 2.0 mol L-l, individual studies have dealt generally with only a single ionic strength, making problematic the extrapolation of the conditional stability
constants to infinite dilution. Because of the lack ofexperimentally measured values at elevated temperatures and pressures, Shock and Koretsky (1993) estimated the stability
constantsof REE-acetate
complexes using the modified Helgeson-KirkhamFlowers (HKF) equation. Recently, Deberdt et al.
(1998; 2000) measured the stability constants of
La3+-, Gd3+- and Yb3+-acetate complexes at temperatures up to 80°C, and Wood et a1. (2000) determined
the stability constants of Nd3+-acetate complexes in
0.1 mol kg! NaCl from 25 to 225°C. Both Deberdt et
a1. (1998; 2000) and Wood et a1. (2000) used potentiometric methods. Also, Wruck et a1. (1997) determined stability constants of Nd3+-acetate complexes
from 20 to 70°C by means of photoacoustic spectroscopy and conventional spectrophotometry.
Because the rangesof ionic strengths
and/ortemperatures covered in these studies are somewhat
restricted, additional measurements are warranted.
This paper is the tenth in a series of contributions toimproved understanding of the aqueous geochemistryof the REE (Wood, 1990a,b; Wood, 1993; Wood andWilliams-Jones, 1994; Wood et al., 1995; Gammons
et al., 1996; van Middlesworth and Wood, 1998;
Gammons and Wood, 2000; Wood et al., 2000). Here
we report the potentiometric determination of the stability constants of the monoacetato (LnAc2+) complexes of La3+, Nd3+, Gd3+ and Yb3+ in 0.1, 0.2, 1.0
and 2.0 mol kg! NaCl from 25 to 70°C at 1 bar. Preliminary results of this investigation were reported by
Ding et a1. (1998). The potentiometric measurements
were conducted in a different manner from those ofeither Deberdt et al. (1998; 2000) or Wood et al.
(2000), and should provide an independent check on
their results.
2. METHODOLOGY
2.1. TheoreticalIn this study, cumulative stability constants (Pn) aredefined with respect to the following generalized relations:
(1)

f3n =

aL A 3-nn cnnaLn3+aAc-(2)

where Ln3+ is any trivalent REE ion, Ac- is the acetate
ion, aj is the activity of species i, and n is the ligand

The aqueous geochemistryof the rare earth elements and yttrium

number and is assumed to range from 1 to 4. The potentiometric method employed is based on the fact that,complexation of free acetate with Ln3+ according to reaction (1) above will decrease the free acetate activity,
causing the proton association reaction of acetate:(3)
to shift to the left, thus liberating protons. Therefore,
complex formation can be monitored by measuring
the pH of a solution containing REE and acetate,when neutralized by stepwise addition of a strong
base (see Baes and Mesmer, 1976 and Martell andMotekaitis, 1992).
In this study, we carried out our experiments at aconstant ionic strength maintained by the addition of
NaCl. As a result, activity coefficients are maintained
approximately constant, and we can define a conditional, cumulative stability constant according to thefollowing expression:

Qn =mLnA 3-nennmLn3+mAc-(4)

These conditional stability constants are truly constant
only at the ionic strength at which they were determined. Determination of conditional stability constants
at a series of ionic strengths permits the determinationof the stability constant at infinite dilution by extrapolation to zero ionic strength. In this study we define
pHm to be the negative logarithm of the stoichiometrichydrogen ion molality, i.e., p~ = - log mH+. The convention employed is that H+ is not complexed by thechloride ion and ion-pairing is treated implicitly by theactivity coefficient model used.
We have carried out these experiments in NaCI
medium as recommended by Millero (1992), rather
than NaCI04 or Na-trifluoromethanesulfonate
media,
so that the stability quotients determined would be directly applicable to natural hydrothermal
fluids,which are dominated by NaCl. Furthermore, REEchloride complexes are known to be relatively weakat the temperatures employed in this investigation(Wood, 1990a,b; Gammons et al., 1996 and this volume). As will be seen below, our results indicate thatchloride complexation was insignificant in our lowerionic strength experiments but not at the highest NaCI
concentration employed.
Measurements by Sonesson (1958b) indicate that
small amounts of polynuclear acetate complexes areformed by the heavy REE at high REE concentration.
However, they are not expected to be important at the
low total REE concentrations employed in our experiments. The potentiometric study of Nd-acetate complexes by Wood et al. (2000) confirmed thatpolynuclear species are not significant at REE concentrations of 0.005 mol kg ", the concentration em-

Part X

211

ployed in the present study, so such species were notconsidered in the treatment of our data.
2.2. Materials
All chemicals used in the experiments were reagent
grade, or better, and all solutions were prepared with
18-MQ·cm deionized, degassed water. The following
stock solutions were prepared: 1) 0.097 mol kg! HCI,
prepared from concentrated HCI (37 wt %, Mallinckrodt), and standardized by titration against certified
0.1 N NaOH (Fisher); 2) 0.0966 mol kg! NaOH,
which was prepared from 50 wt % NaOH (Fisher)and standardized against the 0.097 mol kg! HCI bytitration; and 3) 4.868 mol kg! NaCl, which was prepared from A.C.S. grade crystals (Fisher), standardized against silver nitrate by titration using achloride-ion selective electrode as endpoint detector.
The above three stock solutions of HCI, NaOH, and
NaCI were sparged with argon and stored under positive argon pressure in 20-L polyethylene containers to
exclude CO2, Solutions were transferred by positive
argon pressure to minimize contact with air. Stock solutions of REE chlorides, prepared from hydrated
crystalline solids (Alfa Aesar), were standardized byacidimetric titration of the eluent obtained by passing
an aliquot of the REE stock solution through a cationexchange colunm (Dowex 50WX2-200 ion-exchange
resin in H+ form, Aldrich). Excess acid was requiredin the REE stock solutions to prevent hydrolysis. In
our case, sufficient concentrated HCI was added tothe stock solution to keep the pH near 2. A stock solution of 0.196 mol kg! sodium acetate, which was prepared from A.C.S. grade crystals (Sigma ChemicalCo.), was standardized against 0.097 mol kg! HCI bytitration.
2.3. Potentiometric Measurements
The stability constants of La3+, Nd3+, Gd3+, andYb3+ complexes with acetate were determined ationic strengths of 0.1,0.2, 1.0 and 2.0 mol kg! andat temperatures of 25, 40, 50, 60, and 70 °C. The
maximum total acetate and REE concentrations employed were 0.02 and 0.005 mol kg", respectively.
Separate experiments to measure the ionization constants of acetic acid were conducted over the samerange of temperature and ionic strength. All titrations were repeated at least twice. The procedure
employed closely follows that described by Martelland Motekaitis (1992) and is described briefly
below. Additional details of the titration procedureare given by Ding (1999).
The experiments were conducted using two jacketed glass reaction vessels. One of these vessels

R. Ding and S. A. Wood

212

...-

pHelectrode----..
!Ar InIEI3.8 Fe

IIcc
HAAKE

'Ar Inl"III"I"I"~EIII;.
.......
i__,...

---........ r

L

Heating
Bath/Circulator
08

Figure 1: Schematic diagram of the apparatus employed in this study for potentiometric titrations. The
reference cell is used for calibration of the pH electrode and contained solutions with precisely known
concentrations of N aCl and strong acid or base. The test cell was used for the stability constant measurements and contained solutions with precisely known concentrations of LnCI3, NaCl, HCl and NaAc. Aftercalibration in the reference cell (as described in the text), the electrode was transferred immediately into
the test cell and the measurement of stability constants was initiated.

served as the reference cell in which the pH electrode
was calibrated, and the other served as the test cell inwhich the stability constants were determined. The
apparatus used to carry out the potentiometric titrations is illustrated schematically in Figure 1. The
caps for each reaction vessel have five openings toallow a combination glass pH electrode (Orion,
ROSS), PEEK (polyetheretherketone)tubing (thetitrant inlet), a mercury thermometer, and gas inlet
and outlet tubes to be introduced into the cell. Ultrapure (99.99%) argon was continually
bubbledthrough the cells to exclude CO2, Before being introduced into the reaction cells, the argon was passed
through a polyethylene drying tube containing ASCARITE II (to absorb any traces of CO2) and bubbled through an aqueous solution at the same ionicstrength and a temperature near that of the reaction
cell (to inhibit evaporation of water from the cell). A
0.8-mm thick Teflon gasket was used to effect a seal
between each cell and its lid. The temperature wascontrolled by circulation of thermos tatted water
through the jacket of the cell using a circulatingwater bath (Haake Model D8). The desired temperature could be maintained throughout the titration experiment to within ±0.05 °C. The solutions weremixed thoroughly during titration by means of a
Teflon-coated magnetic stirring bar. When not in use,

the glass electrode was stored in a N aCI solution containing 0.001 mol kg! H+ with a similar ionicstrength to that of the next experiment.
At the start of each experiment, weighed amountsof reference solution (containing a known amount of
HCI) and test solution (containing known amounts of
LnC13, HCI, NaCI and acetate) were placed into thecalibration and test cells, respectively, using polyethylene transfer pipettes. After sealing the cells, air was
removed by sparging with water-saturated Ar and thesolutions were allowed to attain thermal equilibrium.
The potential of the glass-electrode in the calibration
cell was monitored with a pH meter (Radiometer
Copenhagen Model PHM84). Once a stable cell potential was attained « 0.01 mVImin), the potential
was recorded. Then, a precise amount of standard basewas titrated into the cell to change its pH a known
amount, and a second potential reading was taken
once a stable value was attained. This procedure wasrepeated until a calibration curve of mV vs. calculated
pHm' containing 5-10 points, was obtained. Then theelectrode was transferred immediately to the test celland measurement of the stability constant was begun.
The measured potential was converted to pHm of thetest solutions using the calibration curve describedabove. This calibration procedure is based on the
Nernst Equation:

The aqueous geochemistry of the rare earth elements and yttrium Part X

o

E=E

+Eu+

RT
H
2.3025FPm(5)where E is the measured potential, EOis a constant depending on the glass and reference electrodes, I\J isthe liquid junction potential, R is the gas constant andF is Faraday's constant. If EO and ELJ together are
constant in both the calibration cell and the test cell,then a calibration plot of potential vs. pHm can beused to determine the pHm of the test cell during stability constant measurements.
Calibration was conducted under the same conditions of ionic strength and temperature as the test cell.
The major ion compositions in the calibration and test
cells were made as similar as possible to minimize
liquid junction and activity coefficient effects. The
ionic strength of the solution and the concentration of
the titrant-base were also matched to avoid changes inionic strength during titration. Thus, we believe the
ELJ term in equation (5) to be small and constant inour experiments. Therefore, no explicit corrections
were made for liquid-junction effects.
During both calibration and stability constant measurements, the titrant solution (NaOH) was transferred
by volume using a syringe pump (KD Scientific,
Model 210) which was connected via PEEK tubing to
the cell. This syringe pump is capable of delivering
volumes down to 0.001 mL, and was calibrated byweighing amounts of deionized water delivered. In
initial experiments it was found that addition of thebase titrant to the test solution caused local supersaturation with respect to solid REE hydroxide, whichthen precipitated and could not be redissolved. This
problem was eliminated by employing PEEK tubing
with a very fine inner diameter, relatively slow titrant
addition rates, a relatively low concentration of NaOH
in the titrant and vigorous mixing in the cell.
During titrations of REE-bearing solutions with
NaOH, eventually a pH was reached where precipitation of Ln(OH)3 was observed. This precipitation
usually occurred where the rate of change of pH withtitrant concentration suddenly increased. Once precipitation occurred, the titration was stopped, andseveral data points taken prior to precipitation weredeleted from the data set to insure that hydrolysis did
not affect the results.
2.4. Determination of Conditional Stability Constants
Conditional stability constants of REE-acetate
complexes and dissociation constants of acetic acid
were obtained from the titration data in this study
using the program "BEST" provided by Martell andMotekaitis (1992). This program obtains the best-fit
stability constants for a system by minimizing the213

Table I. Comparison of dissociation quotients of acetic
acid in aqueous NaCl media measured in this study
with those of Mesmer et al. (1989).
Ionic

T(OC)

log Qllog Q2

DiIT.

25
30
40
50
60
70
25
40
50
60
70
25
40
50
60
70
25
40
50
60
70

-4.52
-4.54
-4.54
-4.54
-4.58
-4.58
-4.50
-4.49
-4.52
-4.53
-4.55
-4.49
-4.47
-4.48
-4.49
-4.49
-4.61
-4.58
-4.57
-4.56
-4.56

-4.54
-4.54
-4.55
-4.56
-4.58
-4.60
-4.50
-4.50
-4.51
-4.53
-4.54
-4.49
-4.48
-4.47
-4.48
-4.48
-4.61
-4.57
-4.56
-4.55
-4.55

-0.02
-0.00
-0.01
-0.02
0.00
-0.02
-0.00
-0.01
0.01
0.00
0.01
-0.00
-0.01
0.01
0.01
0.01
0.00
0.01
0.01
0.01
0.01

strength(mol kg-I)

0.1

0.2

1.0

2.0

)This study - average uncertainty is ±O.03 log units.
2Mesmer et al. (1989) - average uneertainty is ±O.02log units.

sum of the squares of the differences between calculated and measured pH for all the points in the titration, i.e., the standard error (see Martell andMotekaitis, 1992 for details). Raw titration data are
input into "BEST", and calculated stability constants
are output, in volume and molarity (mol L-l) units.
All concentrations and conditional stability constants
reported in this paper are given in molality units (molkg"). Conversions between weight and volume, andmolarity and molality were accomplished using thedensity of aqueous sodium chloride solutions (calculated using the expression of Fabuss et al., 1966), assuming that, because NaCI is the dominantcomponent in solution, the density of the solution isequal to the density of an NaCI solution of the same
concentration. Additional data required by "BEST"are the dissociation products of water, which havebeen reported by Busey and Mesmer (1978).
In order to determine the number and identity of
species present in the titration experiments, a numberof models containing various combinations of thespecies LnAc2+, LrrAc.,", LnAc30 were fit to the titration data. The best fit was deemed to be the one thatminimized the overall standard error of the fit using the
fewest species. Moreover, for a species to be includedin the model, it had to account for a significant percent-R. Ding and S. A. Wood

214

t(°C)

a)

70

60

25

40

50

3.0

La3+
2.5

I

if
en 2.0.Q

1.5

1.0+-------------r---------~~--------~------------~------------,_--------~
0.00310.00320.00330.00280.00290.00300.00341fT(K)

b)

3.0

t(°C)

70

60

40

50

25

Nd3+

2.5

if
en

.Q

2.0

1.5 +----,.----r----,.----r-----r-----1
0.00320.00330.00280.00290.00300.00310.00341fT(K)

Figure 2: Plots of log Q, vs. 1ff(K) for Ln3+-acetate complexes in NaCl solutions at various ionicstrengths (in mol kgl). The error bars represent three standard deviations about the mean of at least two replicates. The solid curves are the predicted values of log Q, calculated from equations (7-10) in the text, which
represent global fits to all the stability constant data for each REE: a) La3+; b) Nd3+;

age of the total REE present over at least some portionof the titration and the standard error of the conditional
stability constant for the species had to be relatively
small. The details on the mechanism by which the best
fit stability constants are calculated by "Best" are given
in Martell and Motekaitis (1992).

3. RESULTS
3.1. Conditional Dissociation Constants of Acetic Acid
The conditional dissociation constants of acetic
acid in NaCI media were measured in this study as a

The aqueous geochemistryof the rare earth elements and yttrium

60
t(°C)
70
3.0 T""__;_....;__I..-_.......JL...- __c)

50

I..- __Part X

215

25

40

--l.----,

...I-

2.5

cf
OJ

2.0

..Q

1.5

1.0

+----r-----,r-----r----r-----,r------1
0.00290.00280.00300.00310.00320.00340.0033

1!T(K)

t(°C)

d)

70

60

25

40

50

3.0

Yb3+

0.1

2.5

r:

cf
OJ

2.0

..Q

1.5

+---.....,..----r---.....,..---..-----r----f

1.0
0.00280.00290.00300.00310.00320.00330.0034

1!T(K)

Figure 2 (continued):c) Gd3+; d) Yb3+.

check on our experimental methodology and data reduction procedures. The results of these measurements are given in Table 1, where they are comparedto the results of Mesmer et al. (1989). It can be seenthat the two sets of measured dissociation quotientsare in excellent agreement at all temperatures. Thisagreement suggests that our experimental apparatus
and procedures are capable of yielding high-quality
conditional equilibrium constants.

3.2. Conditional Stability
Constants of REE-Acetate Complexes
Under all experimental conditions and for all REE
studied, only the monoacetato complex, LnAc2+, wasrequired to fit the titration data. Other complexes suchas LnAc2 + were too low in concentration to measureaccurately at the acetate concentrations studied. Thiswas particularly evident because, on fitting a two-R. Ding and S. A. Wood

216

Table 2. Unsmoothed conditional stability constants for
LnAc2+ in NaCl solutions obtained in this study at
various temperatures and ionic strengths.
t
I
La
Nd
Gd
Yb
(0C) (mol kg-I)

25

40

50

60

70

0.1
0.2
1.0
2.0
0.1
0.2
1.0
2.0
0.1
0.2
1.0
2.0
0.1
0.2
1.0
2.0
0.1
0.2
1.0
2.0

1.8S±O.OS 2.17±0.09 2.07±O.OS 1.88±O.09
1.78±O.18 2.14±0.09 2.03±O.18 1.79±O.09
1.34±0.OS 1.73±O.09 1.71±0.14 1.4S±0.OS
l.38±O.09 1.73±O.OS 1.79±0.02 1.S8±0.OS
2.12±O.09 2.33±0.22 2.29±O.14 2.16±O.OS
1.80±0.OS 2.17±0.18 2.09±O.22 1.88±O.09
l.39±O.02 1.80±0.14 1.77±0.02 1.60±0.14
l.39±O.09 1.76±O.OS l.78±O.OS 1.69±0.OS
2.16±0.142.37±0.27
2.34±O.27 2.22±O.14
1.92±O.OS 2.30±O.02 2.26±0.14 2.19±O.09
1.48±O.02 1.90±0.l8 1.89±0.09 1.73±0.02
1.36±O.OS 1.79±0.OS l.79±O.02 l.7S±O.OS
2.34±O.14 2.S2±O.14 2.49±0.22 2.44±O.02
2.l4±O.32 2.38±O.OS 2.37±0.22 2.33±0.OS
1.81±0.27 2.00±0.22 1.98±O.18 1.96±O.OS
1.48±O.14 1.84±0.02 1.82±O.09 1.80±0.04
2.37±O.27 2.S8±O.09 2.56±0.14 2.S2±0.OS
2.17±0.14 2.48±0.18 2.47±0.09 2.40±0.09
l.79±O.05 2.01±0.18 2.04±O.4S 1.97±O.09
1.4S±O.OS 1.81±O.27 1.86±0.14 1.90±0.14

Errors given represent three standard deviations from the mean of
two or more replicate measurements.

species model (LnAc2+ and LnAc2 +) to the data andcalculating the distribution of species based on thestability quotients obtained, it was found that the
maximum percentage of total REE for the second
complex was always very low (::::0.18%) and the standard error for the estimate of Q2 was often larger thanthe value of Q2 itself. Clearly, at the acetate concentrations employed in this study, the diacetato andhigher complexes were present in insufficient quantity to obtain their conditional stability constants.
Thus, we report conditional stability constants only
for the monoacetato LnAc2+ complex.
The unsmoothed, conditional stability constants for
the LnAc2+ complexes (i.e., Ql) determined in thisstudy are given as a function of temperature and ionicstrength in Table 2. In each case the quoted error represents three standard deviations from the mean ofreplicate measurements. The temperature and ionicstrength dependencies of the stability constants arealso shown in Figure 2. For the most part, the measured conditional stability constants increase steadily
with increasing temperature. However, for La, Nd and
Gd, log Q1 values are nearly independent of temperature at I = 2.0 mol kg", It is also evident that, in general, the conditional stability constants increase withdecreasing ionic strength. However, the difference inconditional stability constants at I = 1.0 and I = 2.0
mol kg! is relatively small, especially for Gd and Yb.

Table 3. Values of log ~, for REE-acetate complexes at zeroionic strength calculated in this study from equations (7-10).
we)
La
Nd
Gd
Yb
25
2.58
2.90
2.81
2.60
40
2.79
3.08
3.00
2.86
50
2.92
3.18
3.13
3.03
60
3.04
3.28
3.24
3.18
70
3.15
3.38
3.35
3.32
Average uncertainty

is ±O.llog

units.

For each REE studied, an equation of the following
form was fit to the entire set of conditional stability
constants:
log Q1 = A + BIT + C- lIT + DoT-exp(I) +LQ;/-Ay.II/2/(1 +11/2)(6)where ~ is the Debye-Hiickellimiting
slope (takenfrom Ananthaswamy and Atkinson, 1984), I is theionic strength in mol kg", T is the temperature in
Kelvin, LQ;12 is the change in the charges squared forthe complexation reaction (here LQ;12 = -6), and A, B,
C and D are fit parameters. The first two terms inequation (6) are of the van't Hoff type. The last term
represents the Debye-Hiickel contribution to activity
coefficients. The remaining two terms were found,after an exhaustive search, to provide the overall best
fit to the stability constant data as judged by the standard error of fit and examination of the residuals. The
best fit A, B, C and D parameters were obtained using
the Marquardt-Levenberg least squares algorithm in
Sigmaplot® (Jandel Scientific). The resulting equations for the conditional stability constants for each of
the REE studied are:log Ql (La3+) = 7.081- 13261T + 125.2·1IT1.892xlO-4·T·exp(I) - 6.Ay.II/2/(I+II/2)(7)
log QI(Nd3+) = 6.671-11081T + 141.5·1IT1.856xlQ-4·T·exp(I) - 6.~.I1I2/(1 +1112)(8)
log Ql(Gd3+) = 7.092 -12581T + 166.5-1IT2.239xlQ-4·T·exp(I) - 6.~.I1I2/(I+III2)(9)log Ql(Yb3+) = 8.196 - 16591T + 137.4·1IT1.146xl0-4·T·exp(I) - 6.~.I1I2/(1 +11/2) (10)The solid curves in Figure 2 represent the fit to theconditional stability constant data provided by theabove equations. It can be seen that equations (7-10)
fit the majority of the data within their respective uncertainties, and almost all the data within ±0.1 log
units. Smoothed values of the stability constants at
25-70 °C and infinite dilution, calculated from equations (7-10), are given in Table 3.

The aqueous geochemistryof the rare earth elements and yttrium

217

Part X

Table 4. Thermodynamic data obtained in this study for the formation of
LnAc2+ complexes at zero ionic strength and 25°C.
Gd
Nd
La

Property

Yb~GOr,298.'5 (kJ mol")-14.7±O.6-16.6±O.6-16.0±O.6-14.8±0.6

MIOr,298.'5(kJ mol")

2S.1±2.4

20.9±1.4

23.7±1.S

31.6±1.8~SOr,298.'5 (J mol! K-')

133±1OI 26±7

133±7

IS3±4

21.0

-

17.2

23.9

~HOr,298.'5 (kJ mol")!
lDeberdt et al. (1998, 2000)

Table S. Values of log Q, reported in the literature for 20-2ScC.
Yb
Gd
Nd
La
!(Be)
MediumSource
1.84
ND
1.56
1.90
2.0 mol L-' NaClO4
20
Sonesson (19S8a)
1.64
ND
ND
ND
2.0 mol L-' NaClO4
20
Sonesson (19S8b)
ND
1.82
ND
ND
1.0 mol L-l NaClO4
20
Sonesson (1959)
2.03
2.22
2.16
2.02
0.1 mol L-' NaCI04
20
Kolat & Powell (1962)
2.56
ND
2.67
2.55
infinite dilution
25
Archer & Monk (1964)
2.45
ND
ND
ND
infinite dilution
25
Archer & Monk (1966)
2.02
1.84
1.82
2.11
0.1 mol L-' NaClO4
25
Kovar & Powell (1966)
ND
ND
1.54
ND
2.0 mol L-' NaClO4
25
MigaI' & Chebotar (1967)
ND
ND
ND
1.90
2.0 mol L-' NaClO4?
Migal' & Chebotar (1970)
ND
ND
1.35
1.93
0.5 mol L-' NaClO4?
Migal' et al. (1971)
ND
ND
ND
1.80
2.0 mol L-' NaClO4
20
Gruzdev & Ermolaev (1974)
ND
ND
ND
1.83
2.0 mol L-' NaClO4
21
Bukietynska et al. (1977)
ND
ND
1.81
ND
2.0 mol L-' NaClO4
21
Bukietynska et al. (1981a)
ND
ND
ND
1.93
2.0 mol L-I NaClO4
21
Bukietynska et al. (1981b)
ND
1.96
1.66
1.93
2.0 mol L-' NaClO4
Bukietynska & Mondry (1984) 21
ND
ND
ND
1.74
0.2SmoIL-'
25
Daniele et al. (1985)
1.20
1.70
1.4
1.76
1.0 mol L-' KCl
20
Ternovaya (1985)
2.24
2.43
2.52
2.24
infinite dilution
25
Wood (1993)
ND
ND
1.90
ND
2.2 mol L-' NaClO4
20
Wruck et al. (1997)
ND
ND
2.20
ND
infinite dilution
25
Deberdt et al. (1998)
2.32
2.52
ND
ND
infinite dilution
25
Deberdt et al. (2000)
ND
ND
ND
2.1
0.1 mol L-' NaCl
25
Wood et al. (2000)
2.81
2.60
2.90
2.58
infinite dilution
25
This study
1.93
2.14
2.23
1.90
0.09 mol L-' NaCl
25
This study'
1.76
1.98
1.73
2.06
0.19moIL-'
NaCl
25
This study'
1.48
1.76
1.73
0.98 mol L-I NaCl
1.38
25This study!
1.50
1.70
1.26
1.70
2.03 mol L-' NaCl
25
This study'
IThe conditional stability constants obtained in this study have been converted to molar units for the purpose of comparison with values
obtained in the literature. ND= not determined.

Thermodynamic(L1Go r,298' L1Ho r,298'

quantities

and L1Sor298) for reaction(1) at 25 °C and infinite

lution4. The values

ar'e given

were obtainedin Table-log

constant

was obtained(7-10) according

was set(7-10) and the resultingwith respectfrom the expression:

AH~,298 - L1G~298
T

(11)

ex-

to tempera-(13)(8.314 J mol-l K-l)

by differentiating

4. DISCUSSION

The enthalpy
equations4.1. Comparison to Previous Workto the relation:
The

dlOg,8I)=((12), the ionic strength

were differentiatedgiven in Table 3 for(1) at 25 °C and infinite dilution.

of reaction

pressions,81,298

Where R is the ideal gas constantand PI 298 is the stability

equationture. Finally, L1S0r.298was obtained

from the formula:

L1G~,298 = 2.3025R(298.15)

reactio~di-of L1Gor,298

To applyequal to zero in equations

st

AH~2982.3025RT2

conditional

stability

constants

this study at 25 °C are compared(12)

in the literature

at 20-25

obtainedinwith those reported°C in Figure

3. DetailsofR. Ding and S. A. Wood

218

3.0

a)

La3+

•

f

...0

0

~

Ii .w

\l

1.5 -l

~

W'

I

1.0

E1

0

2.5.

a
~
2.0

•

0.0

0.2

0.4

8

•

•

Archer and Monk (1964)
Bukietynska & Mondry (1984)
Wood (1993)
This study
Kolat & Powell (1962)
Kovar & Powell (1966)
Migal' et al. (1971)
Ternovaya (1985)
Sonesson (1958a)

Migal' & Chebotar (1967)<>
8 Deberd! et al. (1998)

W

Daniele et al. (1985)

0

Wruck et al. (1997)

E1

Bukietynska & Mondry (1984)
Archer & Monk (1964)

0

Wood (1993)
This study

0

Kola! & Powell (1962)

fj.

Wood et al. (in press)
Migal' et al. (1971)

0.6

11/2/(1 +1112)

b)

3.0.,

Nd3+

a
Ol

I[j

_Q

l!l
2.0

•

~
~
0

&

•

1.5
0.0

•

•...

2.5

0.2

0.4

\l

Kovar & Powell (1966)•

Ternovaya (1985)
Sonesson (1958a)0

Bukie!ynska

Migal' & Chebotar (1970)<> Gruzdev& Ermolaev (1974)•

et al. (1977)

0.6

11/2/(1 +1112)

3.0

C)

J
a
Ol

Gd3+
10 Wood (1993)

•
...

... •

LJ
2.01

_Q

~
15

1.0

Kolat & Powell (1962)\l

Kovar & Powell (1966)
Ternovaya (1985)

E1

Sonesson (1958a)
Bukietynska & Mondry (1984)

•• •
E1

0 Sonesson (1959)
0

1
0.0

0.2

0.4

11/2/(1 +1112)

This study

0

0.6

Deberdt et al. (in press)

The aqueous geochemistryof the rare earth elements and yttrium Part X

219

3.0

d)

•

Archer & Monk (1964)•

This study...

Kolat & Powell (1962)
Kovar & Powell (1966)\1

Ternovaya (1985)

o Wood (1993)
25~

o
Ol

o

2.0

.Q

•

o

• •

1.5

o Sonesson (1958b)oOeberdt et al. (in press)

IB

Archer & Monk (1966)

1.0
0.0

0.2

0.4

0.6

11/2/(1 +11/2)
Figure 3: Comparison of log Q, values for Ln3+-acetate complexes determined in this study with thosereported previously in the literature at 20-25 "C, plotted as a function of ionic strength (11/2/(1 +1"2): a)La3+; b) Nd3+; c) Gd3+; d) Yb3+. Values of conditional stability constants and ionic strengths from thisstudy have been converted from mol kg! to mol L-' units for the purpose of comparison. The average uncertainty for most of the data depicted is less than or equal to ±O.1log units.

the values from the literature are given in Table 5. In
both Figure 3 and Table 5, we have converted ourconcentrations and stability constants to molar (molL-l) units for comparison with the literature values
which are given almost exclusively in these units.
Note that we did not include previously publisheddata in the fitting process described above, primarily
because our experiments were conducted in NaCI
solutions, whereas most experiments reported in theliterature to date were conducted in NaCI04 solutions. Independent comparison of our stability constants with those determined in non-complexing
perchlorate media should permit assessment of thedegree to which chloride complexation of the REE
may have occurred.
Figure 3 shows that, in general, our results at finite
ionic strengths agree very well with those in the literature when comparison is made at the same ionicstrength. In particular, our data for all four REE at 0.1
mol L-l ionic strength agree to within 0.1 log unit withthose of Kolat and Powell (1962) and Kovar and Powell (1966), with our conditional stability constants
falling between those reported by the latter two authors. Moreover, our data at 1.0 mol L-l ionic strength
are nearly identical to those of Ternovaya (1985) at thesame ionic strength, with the exception of the datum
for Yb. The agreement among our results and those reported in the literature is particularly notable because
the majority of the results reported previously havebeen determined in non-complexing NaCI04 media,

whereas ours were determined in NaCI media. The
agreement among the data obtained in these two
media suggest that chloride complexation is quite
weak at room temperature and the lower ionicstrengths (NaCl concentrations) considered, in concordance with the known weak complexation of thetrivalent REE with chloride under these conditions (cf.
Wood, 1990a,b; Gammons et al., 1996 and this volume). However, at 2 mol kg! ionic strength, there is adiscrepancy between our results and those obtained in
NaCI04 media. This discrepancy is greatest for La3+,
but decreases with atomic number. For Yb3+, our stability quotient at 2 mol kg! is in good agreement withthat determined in the landmark study by Sones son(1958b). There are several reasons why we believethat the observed discrepancy is a result of the competing effect of chloride complexation at the highest ionicstrength: 1) the discrepancy is only apparent at thehighest concentrations of chloride investigated; 2) the
direction of the discrepancy is compatible with the effect of chloride complexation. As shown by Wood etal. (2000), if chloride complexation occurs, stability
quotients for REE-acetate complexes calculated neglecting chloride complexation will be underestimated; and 3) the discrepancy decreases on goingfrom La3+ to Yb3+, consistent with the apparent decrease in the stability of chloride complexes along the
REE series (Wood, 1990a). Furthermore, the relatively weak temperature dependence of the stability
quotients for La- and Nd-acetate complexes at I = 2.0

220

R. Ding and S. A. Woodmol kg! may be explained by the fact that the temperature dependencies of REE-acetate and -chloride
complexes are quite similar (Wood et al., 2000), suchthat the two offset one another.
The agreement among the results at infinite dilution is more variable. For La3+ at infinite dilution(Fig. 3a), the stability constant of this study agrees
well with the value determined by Archer and Monk
(1964). On the other hand, the result of Deberdt et al.
(1998) agrees more closely with the value extracted
by Wood (1993) via a Pitzer treatment of all the available literature data. However, the value determined inthis study is more than 0.3 log units greater than thatof Deberdt et ai. (1998). For Nd3+ (Fig. 3b), the value
determined in this study is considerably larger thanany of the values determined by Archer and Monk
(1964), Wood (1993), or Wood et a1. (2000), all ofwhich are relatively close. For Gd3+ (Fig. 3c) the values from Wood (1993) and Deberdt et al. (2000) are
relatively close, but our value is more than 0.2 log
units greater. Finally, for Yb3+ (Fig. 3d) the values
from this study and Archer and Monk (1964, 1966)
are similar, falling within 0.20 log units of one another, whereas the values of Wood (1993) and Deberdt et a1. (2000) are lower.
The discrepancies noted among values of log PI
are most likely due to uncertainties in the activity coefficients of highly charged species such as the trivalent REE. Figure 3 demonstrates that the agreement
among log Ql values determined at finite ionicstrengths is generally much better than that amonglog PI values at infinite dilution. Moreover, quite different methods of extrapolation to infinite dilution
were employed in each of these studies. For example,
Archer and Monk (1964; 1966) employed the Davies
equation to correct their constants to infinite dilution.
On the other hand, Wood (1993) used the data ofKovar and Powell (1966) at 0.1 mol kg! ionicstrength and those of Sonesson (1958a,b) at 2.0 mol
kg! ionic strength combined with a modification of
the Pitzer equations proposed for REE complexes byMillero (1992). Deberdt et al. (1998; 2000) obtainedlog PI values at infinite dilution by applying the extended Debye-Hiickel equation to calculate activity
coefficients for their low, but variable ionic strength
solutions. Wood et al. (2000) estimated the value oflog PI for Nd using the extended Debye-Hiickel equation recommended by Helgeson (1969). Finally, ourstudy is the only one in which the constants were determined over a range of ionic strengths and extrapolated semi-empirically to infinite dilution.
As mentioned in the Introduction, there are relatively few data available in the literature on the stability constants of REE-acetate complexes at elevatedtemperature. However, in the next few paragraphs wecompare our data to those from the relatively few studies reported in the literature conducted at temperatures
greater than 25 °C. Our data for stability constants atelevated temperatures are compared to those determined by Deberdt et a1.(1998, 2000) in Figures 4a,b,c.
In general, the data exhibit similar trends with respect
to temperature, but the stability constants of Deberdt
et a1. (1998, 2000) are consistently lower than ours by
0.2-0.4 log units. As discussed in the preceding paragraph, the discrepancies shown in Figure 4 most likely
arise from differences in the manner in which the results were extrapolated to zero ionic strength.
The conditional stability constants measured by
Wruck et a1. (1997) for Nd3+ as a function of temperature (20-70 "C) in 2.2 mol kg! NaCI04 are morethan 0.2 log units higher than those measured by us in
2.0 mol kg! NaCl. Similarly, the results of Bukietynska et al. (1981a) for 2.0 mol kg! NaCI04 at temperatures from 0 to 70 °C are 0.3-0.4 log units higher
than those of this study for 2.0 mol kg! NaCl. These
differences between our results in 2.0 mol kg! NaCI
solution and previous measurements in -2.0 mol kg!
NaCI04 solutions are most likely attributable to theeffect of chloride complexation.
The recently reported conditional stability constants for Nd3+ in 0.1 mol kg! NaCI solutions from
25-225 °C of Wood et a1. (2000) are compared withconditional stability constants derived in this study
from 25-70 °C in Figure 5. It is apparent that the stability constant determined in this study based on a
one-species fit of the experimental data at 25 °C is ingood agreement with the corresponding value measured by Wood et al. (2000), but only fair agreementwith the data at the higher temperatures. There is a discrepancy of 0.2-0.3 log unit between these two data
sets at temperatures greater than 25 °C. The uncertainty in the log Ql values of Wood et al. (2000) below
100 °C is on the order of ±O.l log unit, and therefore
the error limits of the two sets of data just overlap oralmost overlap. It is interesting to note that, although
there was no statistical justification for the introduction of a second Nd-acetate complex to fit our data,
Figure 5 shows that the value of log Ql is sensitive towhether or not a second complex is included in thespeciation model. Log Q1 values calculated with a
two-species model are lower than those calculatedusing a one-species model, and these values are closer
to those of Wood et a1. (2000). Nevertheless, because
there is no statistical justification for including a second complex in the fit to the titration data, and because
the conditional stability constants derived for the second complex have unreasonably large uncertainties
and exhibit highly irregular trends with respect to temperature and ionic strength, we report only log Ql and
log PI values obtained from the one-species model inThe aqueous geochemistryof the rare earth elements and yttrium Part X

a)
80

70

60

50

40

25

La3+

4

1=0
3
C!l.

OJ

.2

2

0.00320.00300.00280.00341fT(K)

b)

t(°C)
70

80

50

60

40

25

4

Gd3+

3

1=0

I

y

~

~

co,

OJ

.2
2

0.00320.00300.00280.00341fT(K)

C)

t(°C)
80

70

60

50

40

25

4

Yb3+
1=0
3
C!l.

OJ

.2
2

0.00280.00320.00300_00341fT(K)
Figure 4: Comparison of log ~, values obtained in this study with those determined by Deberdt et al.
(1998,2000) at various temperatures for: a) La3+; b) Gd3+; and c) Yb3+ The open squares are the results ofthis study and the open circles are the results of Deberdt et aI. (1998, 2000). The solid lines are linear leastsquares fits to the data and are provided as visual aids only.

221

R. Ding and S. A. Wood

222

200

50

25

5.0

Nd3+
1=0.1 molal

4.5
4.0
3.5

a

3.0

en
0

2.5
2.0
1.5
1.0
0.00200.00240.00280.0032

1!T(K)
Figure 5: Comparison oflog Q, values determined in this study for Nd3+ in 0.1 mol kg! NaCl solution
with the smoothed values obtained by Wood et al. (2000) in the same medium at various temperatures up
to 250 DC.

7~-----------------------------------------------------------------~
Nd3+
6 infinite dilution
5

C!:l..
0')

o

4

o

3

8¢

2+--~---,..---~-----,,.------r-l

o

50

100

150

200

250

Figure 6: Comparison of stability constants (log 13,) at infinite dilution estimated in this study (filled
circles) for the first Nd3+ - acetate complex at elevated temperatures with those determined by Wood etal. (2000).

this paper. Results of the two-species model fits to thetitration data can be found in Ding (1999).
As shown in the preceding paragraphs, our values
oflog Ql and log
agree to within the combined experimental errors (±0.1-0.3 log units) with the data ofPI

Deberdt et a1. (1998, 2000) and Wood et al. (2000).
The agreement among these studies is encouraging
because, although all three studies employed a potentiometric method, there are significant differences inthe experimental methodology. In this study, we em-

The aqueous geochemistry

6

~

~A

of the rare earth elements and yttrium

rA I

La3+
5J

en..
Ol

0

~~

:1
0

6

~

~

o ......
0·
I

50

I

100

I

I

150

200

250

co,

Ol

0

4

i

3 ~

I

0I

5J

..0

~I

.>

3

I

j ~d3+

Nd3+

5

41

..0

0"
.......
0······-

6~

223

Part X

..0
...0"

2
0

50

6~

Yb3+

5

100

150

200

4

..0

O·
3J

...0"

-?

250

/
.....
0··
0 ...
...0.........
0 ........

I

2

2
0

50

100

150

200

250

t(°C)

0

50

100

150

200

250

t(°C)

Figure 7: Comparison of stability constants (log ~,) at infinite dilution estimated in this study using theisocoulombic approximation (solid circles, solid line) with those estimated by Shock and Koretsky (1993)(open circles, dotted line) for the first acetate complex with: a) La3+; b) Nd3+; c) Gd3+; d) Yb3+

ployed glass electrodes, and titrated NaOH into a solution containing LnC13, NaCl, HCl and acetate, holding ionic strength constant throughout the titration.
Deberdt et al. (1998, 2000) also used glass electrodes,
but titrated a LnCl3 solution into a sodium
acetate/sodium hydroxide buffer solution and ionicstrength was low but not constant. Finally, Wood et al.
(2000) employed a hydrogen-concentration
cell, andtitrated a sodium acetate solution into a solution containing LnCI3, HCI and NaCl, maintaining ionicstrength constant throughout.
In Table 4 we compare the enthalpies of complexation calculated by us with those determined by Deberdt et ai. (1998, 2000). For La the agreement
between the two studies is quite reasonable (i.e., thetwo results agree within their combined uncertainty).
However, for Gd and Yb there is a significant discrepancy, with our values being 5-8 kJ mol-l higher
than those of Deberdt et al. (2000). Although thetemperature trends of the stability constants of Gdand Yb-acetate complexes determined in this study
and by Deberdt et al. (2000) are subparallel (Fig. 4),
our results exhibit a temperature trend with a slightly

more positive slope than those of Deberdt et al.
(2000). It is well known that enthalpies derived from
van't Hoff type plots are very sensitive to uncertainties in the log K values, so these slight differences inslope probably account for the observed differencesin enthalpies of complexation.
4.2. Extrapolation of Stability Constants
to Higher Temperature
Although we have not derived heat capacitychanges for the REE-acetate complexation reactions
studied here, examination of Figure 4 suggests that, at
infinite dilution over the temperature range investigated, LlCo p,r,298 "" O. This would suggest that a simple
van't Hoff treatment might be used to extrapolate themeasured stability constants to higher temperatures.
However, there is no reason to believe that LlCo p,r,298
for reaction (1), a non-isocoulombic reaction, would
remain near zero over a temperature range greaterthan the approximately 50 °C range covered in thisstudy. Therefore, in order to extrapolate the stability
constants for REE-acetate complexes to higher tem-

224

R. Ding and S. A. Wood

Table 6. Stability constants (log [3,) for REE-acetate complexes at elevated temperatures calculated in this study
using the isocoulombic method.
WC)
La
Nd
Gd
Yb
100
3.6±O.2
3.7±O.2
3.7±0.2
3.8±O.2150
4.2±O.3
4.3±O.3
4.3±O.3
4.S±O.3200
4.8±O.5
4.9±0.S
S.O±O.S
S.3±0.5250
S.5±O.S
S.S±O.S
S.7±O.S
6.0±O.S

perature, the isocoulombic approximation (Lindsay,
1980; Cobble et al., 1982) was used. The complex
formation reaction:sky (1993) are lower than ours by several orders ofmagnitude at temperatures above 50 "C. Deberdt et
aI. (1998; 2000) and Wood et a1. (2000) also foundthat their experimentally determined stability constants for REE-acetate complexes were higher thanthe predictions by Shock and Koretsky (1993). Furthermore, our measured and estimated stability constants are in much closer agreement with theexperimentally determined values of Deberdt et al.
(1998; 2000) and Wood et al. (2000), than are the estimates of Shock and Koretsky (1993). Thus, we believe that the theoretical estimates of Shock and
Koretsky (1993) require revision in the light of these
new experimental data.

(14)
is not isocoulombic, but it can be converted to an isocoulombic reaction by addition of other appropriate
reactions. In our case, we added reactions (15) and(16) to reaction (14) to obtain reaction (17), which is
isocoulombic.

Ln3+ + AI(OH)2+ + Ac- H
A13+ + LnAc2+ + OH-(17)

The change of heat capacity of reaction (17) is more
likely than that of reaction (14) to remain approximately zero over a wide temperature range. Log K ofreaction (17) then can be extrapolated to elevatedtemperatures using a plot of log K vs. lIT. Values oflog K required for reactions (15) and (16) at elevatedtemperatures
were takenfrom Palmer
andWesolowski (1993) and Busey and Mesmer (1978),respectively. The uncertainty in the extrapolated values of the stability constants is considerably greaterthan in the raw experimental data determined here.
This is a result of the accumulation of errors resulting
from the combination of reactions (14), (15) and (16),as well as potential errors introduced by the isocoulombic approximation. The results of the extrapolation, with estimated uncertainties, are given in Table
6 and compared with the values reported by Wood et
a1. (2000) in Figure 6. The extrapolated stability constants (this study) are in reasonable agreement withthe experimentally derived data (Wood et al., 2000)
considering their combined uncertainties. It is particularly interesting that the stability constants from thetwo studies are nearly identical at 250 "C.
Our extrapolated stability constants are comparedwith those predicted by Shock and Koretsky (1993) inFigure 7. The values estimated by Shock and Koret-

5. CONCLUSIONSThe stability constants of La3+, Nd3+, Gd3+, andYb3+ complexes with acetate have been determined
potentiometrically from 25 to 70 °C at ionic strengths
of 0.1, 0.2,1.0 and 2.0 mol kg'".
Under the conditions of our experiments, complexes higher than the 1:1 complex LnAc2+ were notrequired to fit the data.
The measured conditional stability constants show
a mild, positive temperature dependence over therange investigated. The constants decrease with increasing ionic strength.
The measured conditional stability constants foreach REE can be represented adequately by an equation of the type: log Q1 = A + BIT + C·I/T +D·T·exp(I) + ~?Ay.II/2/(1+1112).
Our conditional stability constants at 25 °C andionic strengths between 0.1 and 1.0 mol kg! are in excellent agreement with the conditional stability constants reported in the literature for NaCI04 media. The
similarity of REE-acetate stability constants in chloride and perchlorate media suggests that REE-chloride
complexes are very weak under the conditions investigated. However, at 2 mol kg! ionic strength, the difference between the data obtained here and previous worksuggest that chloride complexation has become more
important at higher NaCl concentrations.
The experimentally derived stability constants wereextrapolated to 250 °C using the isocoulombic approximation. These extrapolated constants are in fair agreement with previously measured values (Wood et aI.,
2000), but suggest that previous theoretical predictions
by Shock and Koretsky (1993) are in need of revision.

Acknowledgments-This
paper is dedicated to the memory of
Dr. David A. Crerar, who was one of the first geochemists torecognize the importance of dissolved organic matter in thecomplexation of metals in natural aqueous environments andThe aqueous geochemistry of the rare earth elements and yttrium

was a pioneer in the study of the thermodynamics of metal- organic complexation at elevated temperatures. SAW is indebted
to Professor Crerar for the positive influence on his career. Acknowledgment is made to the Donors of the Petroleum Research Fund, administered by the American Chemical Society,and the State Board of Education of Idaho for financial supportof this research. Some of the data reduction and preparation ofthis manuscript were conducted while SAW was on sabbaticalleave at Oak Ridge National Laboratory. This sabbatical leave
was financially supported by the Geosciences Program of the
Office of Basic Energy Research, U.S. Department of Energy
under contract DE-ACOS-960R22464 with Lockheed Martin
Energy Research Corporation. Reviews by Moira Ridley, Dick
Kettler and Tom Giordano resulted in significant improvementof the manuscript and are gratefully acknowledged.

Editorial handling: T.H. Giordano

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