A potentiometric study of Eu3+ complexation with acetate ligand from 25 to 170°C at Psat

Geochimica et Cosmochimica Acta, Vol. 66, No. 20, pp. 3599 –3613, 2002 Copyright © 2002 Elsevier Science Ltd Printed in the USA. All rights reserved 0016-7037/02 $22.00 ⫹ .00

Pergamon

PII S0016-7037(02)00967-5

A potentiometric study of Eu3ⴙ complexation with acetate ligand from 25 to 170°C at Psat ALEXANDER V. ZOTOV,*,1,4 BORIS R. TAGIROV,1,2 IGOR I. DIAKONOV,3,4 and K. VALA RAGNARSDOTTIR4 1

Institute of Geology of Ore Deposits, Petrography, Mineralogy and Geochemistry RAS, Staromonetnyi per. 35, Moscow 109017, Russia 2 Laboratoire de Geochimie, Universite Paul Sabatier, 38 rue des Trente Six Ponts, 31400 Toulouse, France 3 Schlumberger Cambridge Research, High Cross, Madingley Rd, Cambridge CB3 0EL, UK 4 Department of Earth Sciences, University of Bristol, Bristol BS8 1RJ, UK (Received September 14, 2001; accepted in revised form February 26, 2002)

Abstract—Thermodynamic properties for europium acetate (EuCH3COO2⫹ and Eu(CH3COO)2⫹) complexes were studied by potentiometry at temperatures from 25 to 170°C at the saturation pressure of water. The thermodynamic association constants (K1 and K2) for the reaction Eu3⫹ ⫹ n Ac⫺ ⫽ EuAc3⫺n . n were determined by two different experimental approaches at 25 to 75°C and 75 to 170°C. Logarithms of measured association constants increase with increasing temperature, showing that Eu-Ac complexing increases as temperature rises. Log K1 for EuCH3COO2⫹ increases from 2.91 at 25°C to 4.25 at 170°C, whereas log K2 for Eu(CH3COO)2⫹ increases from 4.83 at 25°C to 7.39 at 170°C. Species distribution calculated in this study using the experimentally determined association constants suggests that in acetatebearing solutions (ⱖ0.05 m) Eu-acetate complexes dominate over the free ion above pH of 4 to temperatures of at least 200°C. At equal total acetate and carbonate concentrations and at 25°C, the relative stability of carbonate and acetate complexes is close. However, as the temperature rises the relative stability of carbonate versus acetate complexes is model dependent. Copyright © 2002 Elsevier Science Ltd A large number of studies have been dedicated to characterising REE speciation and/or providing aqueous association constants of the REE at 25°C and 1 bar (e.g., Paul et al., 1961; Choppin and Unrein, 1963; Choppin and Strazik, 1965; Walker and Choppin, 1967; Aziz and Lyle, 1969; Usherenko and Skorik, 1972; Amaya et al., 1973; Powell, 1974; Choppin, 1976; Smith and Martell, 1976, 1989; Lundquist, 1982; Mironov et al., 1982; Davydov and Voronik, 1983; Spahiu, 1985; Cantrell and Byrne, 1987; Cossy et al., 1989; Grenthe et al., 1992; Lee and Byrne, 1992, 1993; Millero, 1992; Enderby 1995) and Diakonov et al. (1998a, 1998b) have compared thermochemical and solubility data for a wide variety of REEhydroxides. These works emphasise the importance of hydroxide, chloride, fluoride, phosphate, and carbonate complexing in these systems. Relatively few experimental studies have focused on the nature of REE bearing solutions at higher temperatures (Firsching and Mohammadzadel, 1986; Firsching and Brune, 1991; Oelkers et al., 1995; Gammons et al., 1996; Deberdt et al., 1998, 2000; Ragnarsdottir et al., 1998). To overcome the dearth of experimental data Wood (1990a, 1990b) estimated equilibrium constants of reactions forming REE fluoride, chloride, carbonate, and hydroxide complexes at temperatures from 25 to 300°C. More recently Haas et al. (1995) estimated standard partial molal thermodynamic properties of 240 aqueous REE complexes to 1000°C and 5 kbar. They concluded that REE transport may be facilitated by the formation of chloride, fluoride, and hydroxide complexes in hydrothermal fluids at acidic, neutral, and basic pH conditions, respectively. The influence of organic matter on the geochemical cycles of metals is now largely recognized. In the past two decades geochemists have shown that acetate is a common simple acid found in natural waters, basinal brines and hydrothermal solu-

1. INTRODUCTION

The trivalent rare earth element (REE) series is generally considered to behave in a similar manner in igneous to aqueous systems (e.g., Henderson 1984). The aqueous geochemistry of REE has received a growing interest in recent years in relation to radioactive waste disposal because REE are both fission products and considered as chemical analogues for trivalent actinides. A recent study of lanthanide field tracers (Nd and Eu) at Waste Area group 5-North (WAG-5N) of the Oak Ridge National Laboratory, Tennessee has shown that organic matter facilitated the almost-unretarded transport of the lanthanide tracers (McCarthy et al., 1998). It is therefore of utmost importance to measure the thermodynamic dissociation constants of lanthanide-organic complexes. While the inorganic speciation of lanthanides in natural waters has received considerable attention in recent years (e.g., Wood 1990a, 1990b; Lee and Byrne, 1992, 1993; Millero, 1992; Johannesson et al., 1994a, 1994b, 1995a, 1995b, 1997, 1999, 2000; Haas et al., 1995; Lewis et al., 1998) organic acid complexation is still poorly characterised (e.g., Byrne and Li, 1995; Johannesson et al., 1997; McCarthy et al., 1998). Modeling of the complexing of lanthanides in natural waters has shown that the relative abundance of the REE across the lanthanide series is due to the change in complex stability with increasing atomic number. For example, carbonate complexes increase in stability from lanthanum to lutetium. Using thermodynamic correlation techniques, REE-acetate complex stability have been predicted to rise from lanthanum to lutetium (Shock and Koretsky, 1993).

* Author to whom correspondence should be addressed (azotov@ igem.ru). 3599

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A. V. Zotov, et al.

Fig. 1. Examples of potentiometric cell calibration at low temperature (25 to 75°C) experiments. This E-function is shown vs. titrant volume (V, mL). E-function is calculated using e.m.f. measurements (E, mV), volumes of titrant (V, mL of 1.4M HClO4) and titrated solution (90 mL), and theoretical Nernstian (log10·RT/F) slope. As a result, Eo' (Eqn. 3) is determined by linear extrapolation. Total [H⫹] content is calculated from the E reading of the final point calibration titration.

tions (e.g., Giordano, 1985; Drummond and Palmer, 1986; Helgeson et al., 1993; Giordano and Kharaka, 1994). Organic ligands are thus ubiqutous in both hydrothermal and surfical fluids and their role for transporting REE has been qualitatively demonstrated in field studies (McCarthy et al., 1998; Sholkovitz and Shen, 1995; Viers et al., 1997). However, quantitative evaluations of REE transport in organic rich fluids is hindered by the paucity of thermodynamic data at ambient temperatures and the lack of data at elevated temperatures. Shock and Koretsky (1993) have predicted the stability of REE-organic ligand complexes to fill in the gap of lacking high temperature experimental data. In the past decade experimental work has focussed on generating stability constants for metal-organic ligand complexes. Lanthanum-acetate complexes (Deberdt et al., 1998) and Gdacetate, Yb-acetate complexes (Deberdt et al., 2000) to 80°C and Nd-acetate complexes (Wood et al., 2000) have been reported to 225°C. Other studies have generatied dissociation constants for sodium-acetate complexes (Fournier et al., 1998), aluminum-acetate (Fein, 1991a; Be´ ne´ zeth et al., 1994; Palmer

and Bell, 1994), iron-acetate (Palmer and Drummond, 1988), magnesium-acetate (Semmler et al., 1990; Fein, 1991a), leadacetate (Giordano, 1989; Fournier et al., 1997), aluminum¨ hman, 1985; Fein, oxalate (Couturier et al., 1984; Sjo¨ berg and O 1991b; Thyne et al., 1992), cadmium-acetate (Be´ ne´ zeth et al., 2000), cesium-acetate and strontium-acetate (Ragnarsdottir et al., 2001). The use of these dissociation constants in species distribution calculations suggests that aqueous metal-organic

Table 1. The calculated pH values of used standard solutions. pHcalc. NN

Standard solutions

1 2 3

0.01m HCl 0.083m HAc* ⫹ 0.01m NaAc* 0.01m HCl ⫹ 0.05m NaClO4⫹ 0.01m NaCl ⫹ 0.015m NaAc 0.025m HAc ⫹ 0.025m NaAc

4

* Ac corresponds to CH3COO⫺.

75°C 110°C 140°C 170°C 2.05 3.89 4.43

2.05 4.02 4.55

2.06 4.16 4.68

2.06 4.33 4.82

4.77

4.90

5.04

5.21 Fig. 2. Calibration curves for the pH glass electrodes (see text) in 0.01m HCl and buffer acetate solutions from 75 to 170°C.

Eu(III) complexation with acetate

3601

Table 2. Measured e.m.f. (E) and calculated values of free proton concentration (h), free acetate concentration (a) and ligand number (n) based on the experimental titration results of this study. 25°C

50°C

75°C

E, mV

h·105, mol/l

a·103, mol/l

n

E, mV

h·105, mol/l

a·103, mol/l

n

139.9 133.2 127.7 118.9 112.0 106.5 101.6 97.5 92.3 87.8

7.05 5.22 4.38 3.12 2.38 1.92 1.59 1.35 1.11 0.93

3.00 3.84 4.68 6.38 8.10 9.73 11.43 13.02 15.28 17.48

0.321 0.384 0.443 0.547 0.633 0.713 0.774 0.833 0.911 0.974

156.3 148.8 142.5 132.5 124.6 118.4 111.0 104.2 98.7 94.3

8.14 6.22 4.96 3.46 2.61 2.09 1.60 1.25 1.03 0.88

2.49 3.21 3.96 5.50 7.08 8.58 10.71 13.12 15.31 17.24

140.4 133.7 128.2 119.4 112.4 106.8 102.0 97.8 92.4 88.0

7.04 5.42 4.38 3.10 2.37 1.90 1.58 1.34 1.09 0.92

3.02 3.85 4.70 6.43 8.16 9.84 11.54 13.17 15.55 17.76

0.319 0.382 0.440 0.541 0.626 0.702 0.764 0.819 0.888 0.950

159.7 152.1 145.8 135.8 127.8 121.4 113.5 106.9 101.4 96.8

8.01 6.10 4.87 3.40 2.55 2.03 1.53 1.21 0.99 0.84

I ⫽ 0.5 [Eu] ⫽ 0.02 mol/L

160.4 147.8 138.8 126.6 110.9 100.8 90.0

21.13 12.94 9.15 5.67 3.08 2.08 1.37

1.32 2.13 2.96 4.59 7.81 10.8 14.8

0.102 0.167 0.227 0.339 0.521 0.667 0.836

166.2 156.2 148.5 142.2 136.6 131.9 127.7 120.6 114.8 109.8 105.4 99.6

I ⫽ 1.0 [Eu] ⫽ 0.02 mol/L

143.2 135.4 129.2 124.1 119.8 116.1 109.7 104.6 100.1 96.3 92.9 89.9

8.52 6.30 4.94 4.05 3.42 2.97 2.31 1.89 1.59 1.37 1.20 1.07

2.83 3.75 4.68 5.60 6.50 7.34 9.05 10.7 12.2 1.37 15.1 16.4

0.181 0.238 0.290 0.339 0.385 0.431 0.509 0.582 0.645 0.704 0.756 0.806

161 152.4 145.7 140.2 135.0 130.4 123.6 117.4 112.5 107.6 103.7 100.2

Solutions I ⫽ 0.2 [Eu] ⫽ 0.012 mol/L (Electrode 1)

I ⫽ 0.2 [Eu] ⫽ 0.012 mol/L (Electrode 2)

complex formation can substantially increase mineral solubility and chemical mass transport in natural systems (e.g., Be´ ne´ zeth et al., 1994, 2000; Fein 1994; Fournier et al., 1998). To assess the effect of organics on Eu species distribution, dissociation constants of europeum-acetate complexes (EuCH3COO2⫹ and Eu(CH3COO)2⫹) were measured by potentiometry to 170°C. The purpose of this paper is to present the results of this experimental study and to apply them to assess the role of Eu-acetate complexes in natural processes. 2. MATERIAL AND METHODS 2.1. Theoretical Background Previous studies (Sonesson, 1958; Grenthe, 1962; Wood et al., 2000) suggested that in aqueous acetate solutions of less than 0.05 mol/L

E, mV

h·105, mol/l

a·103, mol/l

n

0.363 0.436 0.501 0.618 0.716 0.807 0.930 1.004 1.073 1.14

162.8 155.9 145.0 135.9 128.7 119.5 112.3 106.2 101.1 96.4

6.60 5.25 3.65 2.69 2.12 1.56 1.23 1.00 0.84 0.72

2.88 3.57 4.98 6.54 8.07 10.49 12.77 15.03 17.12 19.28

0.463 0.534 0.661 0.761 0.850 0.948 1.032 1.098 1.156 1.192

2.54 3.29 4.06 5.64 7.29 8.90 11.30 13.72 16.05 18.20

0.359 0.429 0.494 0.607 0.699 0.781 0.881 0.954 1.013 1.067

170.0 162.9 151.6 142.5 135.0 125.8 118.4 112.1 106.7 102.0

6.42 5.07 3.48 2.57 2.00 1.47 1.15 0.93 0.78 0.67

2.96 3.68 5.21 6.84 8.52 11.08 13.58 16.08 18.51 20.84

0.459 0.527 0.644 0.738 0.814 0.901 0.967 1.011 1.042 1.063

18.45 12.88 9.77 7.78 6.36 5.38 4.63 3.59 2.91 2.43 2.08 1.69

1.57 2.21 2.87 3.53 4.24 4.91 5.60 6.96 8.27 9.57 10.8 12.6

0.175 0.245 0.310 0.372 0.428 0.482 0.531 0.623 0.705 0.777 0.843 0.928

162.2 153.1 145.4 139.2 133.8 129.2 124.35 117.0 111.1 104.85 100.2 94.15

15.88 11.73 9.07 7.38 6.16 5.29 4.50 3.52 2.89 2.35 2.01 1.64

1.45 1.93 2.45 2.96 3.47 3.97 4.57 5.63 6.60 7.85 8.85 10.3

0.253 0.331 0.403 0.472 0.537 0.599 0.653 0.758 0.856 0.930 1.01 1.11

12.78 9.38 7.38 6.07 5.02 4.26 3.33 2.67 2.24 1.88 1.63 1.44

2.02 2.70 3.37 4.01 4.75 5.50 6.76 8.12 9.34 10.7 11.9 13.1

0.209 0.277 0.343 0.405 0.459 0.510 0.611 0.696 0.777 0.838 0.901 0.959

171.7 162.3 154.8 148.1 143.0 138.4 130.6 123.8 118.15 113.35 109.4 105.55

9.58 7.00 5.45 4.36 3.68 3.16 2.43 1.94 1.61 1.37 1.20 1.06

2.06 2.77 3.49 4.28 4.97 5.68 7.09 8.56 9.97 11.3 12.4 13.7

0.284 0.350 0.411 0.465 0.520 0.571 0.663 0.739 0.808 0.871 0.934 0.985

metal concentration, Eu forms mononuclear complexes with acetate in accordance to a generalised reaction: Eu3⫹ ⫹ n Ac⫺ ⫽ EuAc3⫺n n . The stability quotients and constants for this reaction are 关Eu(Ac)3⫺n n 兴 关Eu3⫹兴 䡠 关Ac⫺兴 n

(1)

关Eu(Ac)3⫺n ␥ EuAcn3⫺n n 兴 3⫹ ⫺ n 䡠 3⫹ n 关Eu 兴 䡠 关Ac 兴 ␥ E␷ 䡠 ␥ Ac ⫺

(2)

Qn ⫽ and Kn ⫽

where Qn and Kn are the stability quotient and constant, respectively, n is a number of ligand (Ac-ion) in complex, [ ] and ␥ denote the concentration and activity coefficient of species, respectively.

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A. V. Zotov, et al.

Fig. 3. Examples of the typical dependence of the average number of acetate ions (n) connected to Eu3⫹ on the logarithm of free acetate ion concentration (a) at 25 and 75°C. Symbols represent the experimental data given in Table 2. Electrode 1 (closed circles) and Electrode 2 (open circles). Lines correspond to fitting curves derived from log Q1 and log Q2 values reported in Table 3.

In this study we used only mildly acidic solutions (pH up to 5) and did not take into account the hydrolysis of Eu3⫹. At 25°C the hydrolysis can occur only in alkaline solutions (Baes and Mesmer, 1976; Rard, 1985; Wood et al., 1990a) but hydrolysis constants are expected to rise with increasing temperature (Wood et al., 1990b; Shock and Koretsky, 1993). However, the resent potentiometric study of Wood et al. (2000) provided no evidence of hydrolysis in acid acetate solutions until pH ⫽ 5 to 5.5 at 150 to 175°C. Furthermore, Deberdt et al. (1998) showed no significant hydrolysis of La3⫹ and Gd3⫹ until pH ⬎ 9 up to 150°C. In this study the standard state for water corresponds to unit activity of pure liquid at given temperature and pressure. For aqueous species the standard state corresponds to a 1 m (molal) hypothetical solution which behaves ideally. 2.2. Materials All chemicals used in this study were of reagent grade quality. All water used was deionised (MilliQ) or distilled. HClO4 concentration in stock solutions (2.021 mol/L and 1.40 mol/L) were determined by acid/base titrations. NaClO4 stock solutions (8.331 and 5.00 mol/L) were prepared by dissolving solid (hydrated or unhydrated) sodium perchlorate, in freshly boiled distilled water with subsequent filtration through a 0.2 ␮ filter. The concentration of the aqueous salt was determined by weighting dry salt after evaporation of several portions of solution at 140°C. 0.5 and 0.3099 mol/L Eu(ClO4)3 stock solutions were prepared by dissolving Eu2O3 powder in a small excess of HClO4. Eu-oxide was preheated at 800°C for 2 h to remove traces of water and carbonate and subsequently stored in a dessicator. The [H⫹] concentration of the Eu stock solutions was slightly less than 10⫺2 M. Solutions of sodium acetate were prepared using freshly boiled deionised (distilled) water and dry salt that had been heated at 200oC for 2 h and stored in a dessicator. 2.3. Methods of the Potentiometric Study at 25 to 75°C 2.3.1. Experimental methods Measurements were performed at 25, 50, and 75°C in solutions of ionic strength (I) of 0.2, 0.5, and 1.0 (NaClO4) using Shott H62 combination electrode (runs at I ⫽ 0.5 and 1.0), Metrohm 6.0219.100

combination electrode, Metrohm 6.0733.100 single pH electrode and Metrohm 6.0726.100 reference electrode at room temperature, connected with salt bridge (runs at I ⫽ 0.2). The e.m.f. measurements were performed using I-130 pH-meter and Metrohm 713 pH Meter, respectively (the sensitivity is 0.1 mV). To avoid precipitation of KClO4 the combination electrodes and salt bridge were filled with a 3M NaCl solution. The electrode stability was checked before every series of measurements against a NBS pH 4.001 buffer. A double walled glass cell of approximately 200 mL volume was used for the experiments. Water circulation assured constant temperature (⫾0.1°C). Electrode, thermometer, inlet and outlet tubings for Ar gas and burettes for titrant solutions were fitted in a silicon rubber lid. Ar was first bubbled through a NaClO4 solution of the same ionic strength and temperature as the test solution. Experimental solutions were purged of CO2 by bubbling argon through the solution for ⬃1 h before every series of measurements. Ar purging continued during measurements. A two-step titration procedure was used in this study to investigate Eu-Ac complexation in solutions with a constant ionic strength. First, to calibrate the electrode system, the initial solution of Eu(ClO4) with a constant ionic strength (NaClO4) and molar (M) concentration of hydrogen ion ([H⫹]) of approximately 0.003 mol/L was titrated by a HClO4 solution of the same ionic strength to a proton concentration of approximately 0.01 mol/L. Two examples of calibration are shown in Figure 1. At all temperatures, the calibration data (E-function vs. HClO4 concentration) define a straight line with a strictly theoretical Nernstian slope. As a rule, the discrepancy between the experimental and fitted values of E is less than 0.08 mV and never exceeds 0.18 mV. After calibration, the acid solution was titrated in the opposite (basic) pH direction with a Na-Ac solution of the same ionic strength to determine Eu-Ac stability constants using the method of competitive reactions (Rossotti and Rossotti, 1961). To keep the Eu concentration constant, the same quantity of a second solution with Eu content twice as that in the experimental solution was added simultaneously. This titration was continued until the H⫹ concentration reached approximately 10⫺5 M. No correction for the change in ionic strength resulting from complexation was not introduced. The greatest possible decrease of ionic strength at the end of titration did not exceed 20, 10 and 5% at ionic strength of 0.2, 0.5 and 1, respectively.

Eu(III) complexation with acetate

3603

Table 3. Eu-Ac stability quotients (concentration constants), determined in this study as a function of ionic strength (I) and temperature.

Table 5. Eu-Ac concentration stability quotients from the literature.

I* 0.2 0.5 1.0

T, °C

log Q1

log Q2

T, °C

Ia

log Q1

25 50 75 25 50 75 25 50 75

2.11 ⫾ 0.03 2.28 ⫾ 0.02 2.40 ⫾ 0.03 1.91 ⫾ 0.02 2.08 ⫾ 0.02 2.27 ⫾ 0.02 1.83 ⫾ 0.02 2.04 ⫾ 0.02 2.18 ⫾ 0.02

3.42 ⫾ 0.04 3.69 ⫾ 0.08 3.76 ⫾ 0.10 3.40 ⫾ 0.02 3.68 ⫾ 0.02 4.12 ⫾ 0.02 3.25 ⫾ 0.02 3.70 ⫾ 0.02 3.66 ⫾ 0.02

0 10 25 40 55 20 20 25

2 2 2 2 2 0.5 0.1 0.07

1.839 1.886 1.903 1.908 2.064 1.942 2.305 1.903

* Molar (NaClO4).

The e.m.f. reading attained steady state usually within 2 to 5 min at all temperatures. Consequently they were recorded 5 to 15 min after addition of a small (0.5 to 4 mL) portions of titrant solution to the test solution (initially 70 or 90 mL). From time to time, the electrode stability was checked for longer time (up to 60 min). Sometimes, slight non-systematic variations were observed at 50 and 75°C. Consequently, the overall uncertainty of the e.m.f. readings was accepted as 0.1 mV at 25°C and 0.2 mV at 50° and 75°C, respectively. 2.3.2. Calculation methods Measured electrode potential of an experimental solution with a constant ionic strength may be written as: E ⫽ E o' ⫹ ␯ log关H⫹兴,

(3)

where [H⫹] corresponds to a molar concentration (M) of hydrogen ion, ␯ ⫽ 2.303·R·T/F, R and F are gas and Faraday constants, respectively, and T is temperature in K. Eo' corresponds to a sum of the standard electrode (Eo) and liquid junction (EJ) potentials and a molar activity coefficient term (␯ log ␥H⫹). Calculation of the stability quotients were performed in the present study according to classical methods of potentiometric data treatment summarized by Rossoti and Rossoti (1962), and first applied to REEacetate complexation by Sonesson (1958). First, the Eo' and total [H⫹] content are calculated from calibration data. Eo' is determined by a linear extrapolation and total [H⫹] content is generated from the E reading of the final point calibration titration. Then, the nH function is calculated, which corresponds to the number of protons related to acetate, not linked in complexes with Eu3⫹: nH ⫽

关HAc兴 H⫺h ⫽ 关Ac⫺兴 ⫹ 关HAc兴 共H ⫺ h兲Q HAc ⫹ 共H ⫺ h兲 h

(4)

where H corresponds to the total [H⫹] content in the system corrected o' for dilution and is determined from basic titration data (h ⫽ 10(E⫺E )/␯). QHAc corresponds to the concentration ionization constant for acetic acid. This constant was experimentally determined in the present study

log Q2

3.188 3.904

log Q3

3.785

Ref. b b b b b c d e

a–ionic strength, molar (NaClO4); b–Choppin and Schneider (1970), solvent extraction, ⫾0.02 uncertainty in log Q except for 50°C (⫾0.04); c–Grenthe (1962), potentiometry, Eu 0.02– 0.04 M, Ac 0.001– 0.14 M, n 0.09 –1.8, ⬍⫾0.02; d–Kolat and Powell (1962), potentiometry, Eu 0.004 M, Ac 0 – 0.08 M, n ⬍ 2, ⫾0.04; e–Manning and Monk (1962), solvent extraction, ⫾0.01.

by a calibration titration procedure performed at 25°C on a 0.5 mol/L NaClO4 solution without Eu. The average of two values (log QHAc ⫽ ⫺4.49 ⫾ 0.01) is close to the log QHAc precisely determined by Mesmer et al. (1989) in NaCl media and recalculated to molar scale (⫺4.46). Consequently the HAc equilibrium quotients from Mesmer et al. (1989) were used throughout this study, after correction for molar concentrations, using densities of NaCl solutions as a function of concentration and temperature from Potter and Brown (1977). Knowledge of nH allows calculation of the average number of acetate ions (n) connected to Eu3⫹: A⫺ n⫽

H⫺h nH B

(5)

where A and B correspond to total acetate and Eu concentration, respectively. The concentration of free acetate ions (a) can be calculated as: a ⫽ A⫺(H⫺h) ⫺ nB.

(6)

The number of Eu-Ac complexes in experimental solutions may be deduced from the dependence of n on log a. The n value allows the estimation of the number of complexes. If the n is not much greater than 1 as in our study (see section 3.1.), only two complexes (EuAc2⫹ and Eu(Ac)2⫹) dominate in solutions. In this case, their stability quotients are related by this expression (Rossoti and Rossoti, 1961): P 1Q 1 ⫹ P 2Q 2 ⫽ 1,

(7)

where P1 and P2 are functions of n and a. Three different forms of Eqn.

Table 4. Measured pH of tested solutions (the average values from measurements by two pH electrodes). pHmeasured NN 1 2 2a 3 4

Tested solutions

75°C

110°C

140°C

170°C

Acid Eu-solution* ⫹ 0.012m NaAc** Acid Eu-solution* ⫹ 0.015m NaAc 0.01m HCl ⫹ 0.015 m NaAc ⫹ 0.06m NaClO4 Acid Eu-solution* ⫹ 0.03m NaAc Acid Eu-solution* ⫹ 0.04m NaAc

3.44 ⫾ 0.01 3.88 ⫾ 0.01 4.44 4.73 4.88

3.39 3.79 4.57 4.75 4.94

3.33 ⫾ 0.01 3.71 ⫾ 0.01 4.68 ⫾ 0.01 4.76 5.07 ⫾ 0.01

3.25 ⫾ 0.02 3.62 ⫾ 0.02 4.86 ⫾ 0.02 4.80 ⫾ 0.01 5.20 ⫾ 0.01

* 0.01m Eu(ClO4)3 ⫹ 0.00977 m HCl ⫹ 0.00023 m HClO4. ** Ac corresponds to CH3COO⫺.

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Fig. 4. Dependence of the calculated function log Q1 ⫺ A·⌬Z2·I0.5/(1 ⫹ å·B·I0.5) on ionic strength and temperature from 25 to 75°C. Symbols and solid lines correspond to experimental values and fitting curves, respectively.

7 (Eqn. 5–15 to 5–17 in Rossoti and Rossoti, 1961) were used to derive Q values by a least-square regression method. Values of the stability quotients (Q1 and Q2) for different ionic strengths from 0.1– 0.2 to 1–2 were extrapolated to zero ionic strength using an extend Debye-Hu¨ ckel equation: logQ ⫺

A⌬Z 2 冑I

1 ⫹ a˙ B 冑I

⫽ logK ⫹ bI

(8)

where K corresponds the thermodynamic stability constant, A and B are the Debye-Hu¨ ckel electrostatic coefficients, I is ionic strength, Z denotes a species charge and b is a fitting parameter (Vasil’ev, 1962). The ion size parameter å was taken as 4.5 which is close to the average value for 3⫹ charged cations and their complexes (Vasil’ev, 1962; Borisov and Shvarov, 1992). 2.4. Methods of Potentiometric Study at 75 to 170°C 2.4.1. Experimental methods Measurements were performed at 75, 110, 140 and 175°C in solutions of low ionic strength, varied from 0.05 to 0.07, using a high temperature nonisothermal cell. The cell’s construction and electrode specifications is similar with those described by Krukov et al. (1966),

Table 6. Stability constants for Eu acetate complexes (K1 and K2), and constant of reaction (11) determined in this study. T, °C

log K1

log K(11)

log K2

Method

25 50 75 75 110 140 170

2.91 ⫾ 0.05 3.11 ⫾ 0.05 3.30 ⫾ 0.05 3.30 ⫾ 0.07 3.61 ⫾ 0.07 3.89 ⫾ 0.07 4.25 ⫾ 0.06

1.92 ⫾ 0.14 1.98 ⫾ 0.14 1.95 ⫾ 0.14 1.88 ⫾ 0.19 2.34 ⫾ 0.17 2.84 ⫾ 0.16 3.14 ⫾ 0.13

4.83 ⫾ 0.15 5.09 ⫾ 0.15 5.25 ⫾ 0.15 5.18 ⫾ 0.17 5.95 ⫾ 0.15 6.73 ⫾ 0.14 7.39 ⫾ 0.11

a a a b b b b

a–low temperature measurements at constant ionic strength and extrapolation to zero ionic strength; b– high temperature measurements in low concentration solutions and direct calculation using activity coefficients.

Fig. 5. Dependence of the calculated function log Q(11) ⫺ A·⌬Z2·I0.5/(1⫹å·B·I0.5) on ionic strength and temperature from 25 to 75°C. Symbols and lines correspond to experimental values and fitting curves, respectively: Solid (25°C), long dashed (50°C) and short dashed line (75°C).

Pokrovski et al. (1993, 1995) and Ragnarsdottir et al. (2001). High temperature solid contact pH-selective glass electrodes produced by A. Sergeev (“Izmeritel’naya tekhnika,” Moscow) were used. They allow the undertaking of pH measurements from 60 to 70° to 200°C. Electromotive force (e.m.f.) of the cell was measured using a Metrohm 713 pH Meter and two pH-selective electrodes and an external Ag/AgCl reference electrode that was thermostated at 19 to 22°C. Both a reference electrode and a salt bridge with the ceramic junction were filled with 3M KCl solution. The temperature of the cell is regulated and measured with thermocouples of K-type using CAL 3200 Autotune temperature controller with uncertainty of ⫾1°C. The electrodes were calibrated by measuring the potentials in 0.01 m HCl and acetate buffer solutions at the same temperature and pressure as the solutions of interest. Composition of standard solutions is listed in Table 1. Calibration and tested solutions (approximately 210 to 220 cm3) were placed in a Ti-cell with a volume of 380 cm3. before pouring solution into the cell the cell was filled with argon. All solutions were prepared using freshly boiled deionised water. The e.m.f. measurements for all solutions (both calibration and tested) were performed at slowly increasing temperature. Each solution is heated from room temperature to 170°C. The rate of temperature increase was approximately equal to 0.4 to 0.5°C per min. and decreased to 0.1 to 0.2 °C per min. at recording temperatures (75, 110, 140 and 170°C). As it has been shown at low temperature (see section 2.3.1.) these conditions are suitable to approach equilibrium in the system. 2.4.2. Calculation methods The pH of solutions studied with nonisothermal potentiometrical cell is related to the measured e.m.f. (E) as: pH ⫽

E0 ⫺ E ⫺ Ej S

(9)

where Eo represents the sum of the standard glass electrode potential at temperature T, the total reference electrode potential at its temperature (Tref.), the thermodiffusion potential in the salt bridge, and the difference in hydrogen electrode potential at temperatures T and Tref.; Ej and

Eu(III) complexation with acetate

3605

Fig. 6. The pH values measured (symbols) and calculated (lines) as a function of NaAc total concentration at temperatures from 75 to 170°C. Closed symbols are pH measured in Eu-acetate solutions (Table 4, NN 1, 2, 3 and 4), open symbols represent pH, measured in acetate solution without Eu (Table 4, N 2a). Lines were calculated taking into account: two Eu-Ac complexes (EuAc2⫹ and Eu(Ac)2⫹, Table 7) - solid lines; three complexes (EuAc2⫹, Eu(Ac)2⫹ and Eu(Ac)3°(aq), Table 7) - short dashed lines; and without Eu acetate comlexing - long dashed lines. S correspond to the liquid junction potential and the electrode slope (RTln10/F), respectively. Assuming that difference in Ej for calibration solutions was insignificant relative to the uncertainty in other measurements of 2 to 3 mV, the Eo and S were deduced from calibration data as a function of temperature. The experimental solutions were close to the used buffer acetate calibration solutions with respect to both the pH and ionic strength. That gives the basis of defining the pH of investigated solutions using pH ⫽

E0 ⫺ E S

where Eo and S are calculated from calibration data (Fig. 2).

(10)

The pH of calibration solutions was calculated (Table 1) using the GIBBS computer code for the HCh software package (Shvarov, 1999; Shvarov and Bastrakov, 1999) and thermodynamic properties for all species given in the Slop98 database. Calibration curves (Fig. 2) for both pH electrodes give a linear dependency of E on the pH at all temperatures with the slope close to the theoretical Nernstian one. The deviation of experimental points from the fitted line with the theoretical slope dose not exceed 2.5 mV at 75°C and 1.5 mV at other temperatures. There is no systematic deviation in the whole pH range from 2 to 5. The variation of pH in a particular solution calculated from measurements by two pH electrodes does not exceed 0.04 pH units at 170°C and 0.02 to 0.03 pH units at lower temperature. In general, the uncertainty in measured pH values is estimated to be ⫾0.04 pH.

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Fig. 7. Logarithm of the EuAc2⫹ (a) and Eu(Ac)2⫹ (b) stability constants as a function of reciprocal temperature comparing the prediction on the base of Slop98 and the experimental values from this study. Symbols present experimental data obtained by two different experimental approaches: at 25 to 75°C (closed) and 75 to 170°C (open symbols). Solid line denotes the fitting using the data from Table 7. For comparison also shown are the stability constants for La-acetate (Deberdt et al., 1998, Nd-acetate (Wood et al., 2000), Yb and Gd-acetate complexes (Deberdt et al., 2000). (c) - Logarithm of the Eu(Ac)3° stability constant as a function of reciprocal temperature comparing the predictions on the base of Slop98 (dashed line) and this study (solid line). The symbol presents the data of Grenthe (1962) for Q3 (Table 5), extrapolated to infinite dilution in this study.

The standard Gibbs free energy of europium-acetate complexes was computed by iteration using the HCh software package (Shvarov, 1999; Shvarov and Bastrakov, 1999) from measured solution pH by manually adjusting these ⌬fGo values until the computed pH matched its measured counterpart. Species considered in these calculations were CH3COOHo, CH3COO⫺, NaCH3COOo, Na(CH3COO)2⫺, Eu3⫹, Eu(CH3COO)2⫹, Eu(CH3COO)2⫹, Cl⫺, HClo, Na⫹, H⫹, OH⫺, and H2O. Thermodynamic properties for all species except Eu-Ac complexes were taken from the Slop98 database which is an extension of the SUPCRT92 database (Johnson et al., 1992; Shock and Koretsky, 1993; Haas et al., 1995; Shock, 1995; Shock et al., 1997 and references therein). Throughout these calculations Cl⫺ was used instead of ClO4⫺. Activity coefficients were calculated using an extended Debye-Hu¨ ckel

equation (Helgeson et al., 1981). Ion size parameter å was taken as 4.5 for all charged species. Activity coefficients of neutral species are assumed to be unity. Note that a variation of the Gibbs free energy for Na-Ac species does not significantly influence the final results because these complexes in tested solutions are presented in negligible quantities. 3. EXPERIMENTAL RESULTS AND DISCUSSION

3.1. Potentiometric Measurements and Stability Quotient for Eu Acetate Complexes From low temperature experimental data (25 to 75oC) the values of the free acetate ion molar concentration (a) and ligand

Eu(III) complexation with acetate

3607

From high temperature measurements (75 to 150oC) the pH values of Eu-Ac solutions were determined (Table 4). The uncertainty given in this Table demonstrates only the difference in measurements by two electrodes. Total uncertainty of pH is ⫾0.04 at all temperatures. 3.2. Stability Constants and Thermodynamic Properties for Eu Acetate Complexes Values of log Q1 and log Q2 obtained in this study (Table 3) and taken from literature data (Table 5) were extrapolated to zero ionic strength. Figure 4 shows the log Q1 extrapolation at 25, 50 and 75oC. Data of Grenthe (1962) and Kolat and Powell (1962) at 20oC and Choppin and Schneider (1970) at 55oC were recalculated to 25 and 50oC, respectively, using the temperature dependence from Choppin and Schneider (1970). This Figure demonstrates a good consistency of data from the literature and those obtained in the present study, except log Q1 value of Manning and Monk (1962) at 25oC. The latter was thus excluded from our fit. The fitting parameter b was accepted to be one and the same at all temperatures (b ⫽ 0.19). Obtained values of the stability constant at infinite dilution are listed in Table 6. The overall uncertainty in log K1 was estimated to be of ⫾0.05. As for Q2, the direct extrapolation of log Q2 to zero ionic strength is not reliable due to a very high value of the ⌬Z2 for reaction (2) that is equal to ⫺10. It is more reliable to recalculate the experimental values of log Q1 and log Q2 using the logarithm of the stability quotient of the reaction: Eu(Ac)2⫹ ⫹ Ac⫺ ⫽ Eu(Ac)2⫹,

Fig. 8. Logarithm of the first (a) and second (b) REE-acetate stability constants as a function of REE atomic number at 25, 75, 110, and 170°C. Symbols present the experimental data of Deberdt et al. (1998) for La, Deberdt et al. (2000) for Yb and Gd, Wood et al. (2000) for Nd and this study for Eu. Closed and open symbols denote the measured (or interpolated) and extrapolated values for REE-acetate stability constants, respectively.

number (n) were calculated (Table 2). Figure 3 demonstrates a typical dependence of the ligand number on a log (a). As n ⬍ 1.2 for all experiments, data from Table 2 were converted to log Q1 and log Q2, assuming that only two Eu-Ac complexes (EuAc2⫹ and Eu(Ac)2⫹) dominate in the experimental solutions. The average values for stability quotients calculated using three equations and two electrodes (at I ⫽ 0.2) as described above, are given in Table 3.

(11)

for which ⌬Z2(11) ⫽ ⫺4, then extrapolate the log Q(11) to infinite dilution and finally convert the obtained values of log K(11) to log K2. In this procedure the data of Grenthe (1962) and Kolat and Powell (1962) at 20°C were recalculated to 25°C using the temperature dependence from this study. Uncertainty for this extrapolation is remarkably higher than for log K1. The fitting parameter b is determined only at 25 and 50°C (0.31 and 0.50, respectively); at 75°C it was adopted as the mean value of these ones (0.40). Also one experimental point (at 75°C and I ⫽ 0.5) was excluded from extrapolation because it was in disagreement with other data. The results of extrapolation to infinite dilution is presented in Figure 5. The generated thermodynamic constants (log K(11) and log K2) are given in Table 6. Their uncertainties are estimated as ⫾0.15 log unit. Log K1 and log K2 values were also generated from high temperature pH measurements at 75 to 150°C (Table 6). Figure 6 demonstrates a difference in pH of real Eu-Ac solutions (experimental points and fitted dashed line) and those calculated for hypothetical analogous solutions without Eu-Ac complexation (fitted solid line). At experimental conditions this difference is approximately 0.2 to 1 pH unit and shows the tendency to increase with increasing temperature. Note the good agreement of the calculated hypothetical curves with the pH value measured in solution of N 2a (Table 4) with the same ionic strength and Ac concentration as solution N 2 but without Eu. In general, the fitted pH values excellently describe the experimental data with deviations less than 0.02 to 0.03 of pH

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Table 7. Standard partial molal thermodynamic properties at 25°C and 1 bar and HKF equation of state parameters for Eu acetate complexes adopted in this study. ° ° ° ° ⌬fG298.15 S298.15 Cp298.15 V298.15 a1·10 cal· cal·mol⫺1· cal·mol⫺1· cm3· cal·mol⫺1· K⫺1 K⫺1 mol⫺1 bar⫺1 mol⫺1 3⫹

Eu EuAc2⫹ Eu(Ac)⫹ 2 ° Eu(Ac)3, aq

⫺137300 ⫺229579 ⫺320303 ⫺409817

⫺53.0 ⫺13.97 18.26 50.60

⫺36.4 50.1 151.1 235.5

⫺42.0 3.6 54.3 108.7

a2·10⫺2 a3 cal· cal·K· mol⫺1 mol⫺1·bar⫺1

⫺3.1037 ⫺15.3599 2.7500 ⫺1.0666 9.3029 14.9307 16.6413 32.8512

11.7871 6.1690 ⫺0.1123 ⫺7.1605

a4·10⫺4 c1 c2·10⫺4 ␻·10⫺5 cal·K· cal·mol⫺1· cal·K· cal· K⫺1 mol⫺1 mol⫺1 mol⫺1 ⫺2.1440 ⫺2.7348 ⫺3.3961 ⫺4.1370

References

6.0548 ⫺10.4900 2.3161 Slop98 47.3933 7.1730 1.2998 This study (a) 97.1602 27.7439 0.2676 This study (a) 144.1841 44.9368 0.0 This study (b)

(a) Calculated from experimental data (Table 6) and thermodynamic properties of Eu3⫹; parameters a1–a4 are taken from the Slop98 database (Shock et al., 1997; Shock and Koretsky, 1993). ° ° (b) Estimated on the base of logQ3 (Table 5) at 20°C (Grenthe, 1962) and linear correlation of the nonsolvation contribution into S298.15 and Cp298.15 with the number of acetate ligands in the complexes.

unit. There is only one exception at the highest Ac concentrations (0.04) at lower temperatures (75 and 110°C) when the deviation reaches 0.07 and 0.08 of pH units, respectively. Both used potentiometric experimental approaches give a very good agreement at temperature of 75°C. At this temperature the generated values of log K1 and log K2 (Table 6) coincide in an uncertainty range. Figure 7a and 7b shows the log K1 and log K2 as a function of temperature. In the whole temperature range under investigation both thermodynamic constants are higher than those predicted by Shock and Koretsky (1993). Moreover, the generated temperature dependence for both log K1 and log K2 are close to those recently determined for acetate complexes of other rare earth elements: La (Deberdt et al., 1998), Nd (Wood et al., 2000), Gd and Yb (Deberdt et al., 2000). The variation of stability constants (log K1 and log K2) as a function of REE atomic number at 25, 75, 110 and 170°C are shown in Figures 8a and 8b. At all investigated temperatures the stability constants increase from La to Eu and then decrease to Yb as it was noted by Deberdt et al. (2000). 3.3. Prediction of the Thermodynamic Properties of Eu Acetate Complexes at Elevated Temperature and Pressure The stability constants for the formation of Eu(Ac)2⫹ and Eu(Ac)2⫹ (Table 6) were used to determine the standard thermodynamic properties and HKF parameters of these species within the framework of the revised Helgeson-Kirkham-Flowers model (Tanger and Helgeson, 1988). Using this procedure the thermodynamic properties for Eu3⫹, HAc°, and Ac⫺ and volumetric properties of Eu-acetate complexes were adopted from the Slop98 database (Shock and Koretsky, 1993; Shock et al., 1995; Shock et al., 1997). Optimization of the parameters (⌬fGo, So, Cpo) was performed using the UT-HEL computer code (Shvarov, 1999). The values of c1, c2 and ␻ were estimated with the correlations of Shock and Helgeson (1988). Results of these calculations for EuAc2⫹ and Eu(Ac)2⫹ are given in Table 7. Finally, we would like to estimate, highly approximately, the stability of a third Eu acetate complex (Eu(Ac)3o). The good fit to experimental data using EuAc2⫹ and Eu (Ac)2⫹ only indicates that this complex did not appear under our experimental conditions because of the low acetate concentrations (ⱕ0.04 mol·kg⫺1 H2O). The neutral complex has only been observed by Grenthe (1962) at 20°C, at high

acetate concentrations (up to 0.14 mol/L) and ionic strength, I ⫽ 0.5. According to Grenthe (1962), log Q3 ⫽ 3.785 ⫾ 0.02 and log Q2 ⫽ 3.188 ⫾ 0.02 at I ⫽ 0.5 (Table 5) at 20°C. Thus, for the reaction: Eu(Ac)2⫹ ⫹ Ac⫺ ⫽ Eu(Ac)3o

(12)

Log Q(12) ⫽ 0.597 ⫾ 0.03 at the same ionic strength (I ⫽ 0.5). Our extrapolation of this constant to infinite dilution using an extended Debye-Hu¨ ckel Eqn. 8 gives log K(12) ⫽ 0.8 ⫾ 0.1, resulting in log K3 ⫽ log K2 ⫹ log K(12) ⫽ 5.6 ⫾ 0.2 at 20°C. The standard partial molal entropy and heat capacity for Eu(Ac)3° were generated using the linear correlation proposed by Hovey (1988) between the nonsolvation contribution of the thermodynamic properties of complexes and their respective ligand number. Calculated values of the nonsolvation contribution in entropy and heat capacity for Eu3⫹, EuAc2⫹ and Eu(Ac)2⫹ are given in Table 8. The linear extrapolation of these values to ligand number n⫽3 (Fig. 9) permits the approximation of the corresponding data for Eu(Ac)3° and subsequently to estimate So298.15 and Cpo298.15 of this neutral complex (Table 8), assigning the Born coefficient (␻) to be equal to zero. On the basis of ⌬fGo293.15 calculated from log K3, values of o S 298.15 and Cpo298.15 given above and the coefficients a1–a4 from the Slop98 database, the standard thermodynamic properties and HKF parameters for Eu(Ac)3° were generated (Table 7). Predicted values of log K3 as a function of temperature are shown in Fig. 7c. To control the predicted thermodynamic data for Eu(Ac)3° (Table 7) we calculated the pH values of experimental solutions in high temperature experiments with regard to three Eu acetate complexes. They do not agree with measured pH at higher temperature (140 to 170°C) and higher acetate concentration (0.03 to 0.04 mol·kg⫺1 H2O). This can be seen graphically in Fig. 6. The temperature dependency of ⌬fG° for Eu(Ac)3° is defined rather accurately but the value of ⌬fGo298 is based on the data of the only experimental study (Grenthe, 1962) at a single value of ionic strength (0.5M NaClO4). Therefore, the thermodynamic properties for Eu(Ac)3° presented in Table 7 should be considered as a preliminary approximation. Thus, to solve this problem, future experimental studies need to be designed for this system at higher acetate concentrations.

Eu(III) complexation with acetate

3609

° Table 8. Calculated values of solvation (⌬s) and nonsolvation (⌬n) contributions into S298.15 and Cp° plexes.

Entropy, cal·mol⫺1·K⫺1

Eu3⫹ EuAc2⫹ Eu(Ac)⫹ 2 Eu(Ac)3,aq

298.15

for Eu⫹3 and Eu acetate com-

Heat capacity, cal·mol⫺1·K⫺1

° S298.15

° ⌬sS298.15 (a)

° ⌬nS298.15 (b)

° Cp298.15

° ⌬sCp298.15 (a)

° ⌬nCp298.15 (b)

␻·10⫺5 cal·mol⫺1

⫺53.0(c) ⫺13.97(d) 18.26(d) 50.6

⫺13.45 ⫺7.55 ⫺1.55 0.0

⫺39.55 ⫺6.42 19.81 50.6(e)

⫺36.4(c) 50.1(d) 151.1(d) 235.5

⫺21.33 ⫺11.97 ⫺2.47 0.0

⫺15.07 62.07 153.57 235.5(e)

2.3161(c) 1.2998(d) 0.2676(d) 0.0

° (a) ⌬sS298.15 ⫽ ␻·Y ⫽ ⫺5.81·10⫺5·␻; ⌬sCp° 298.15 ⫽ ␻·T·X⫽ ⫺3.09·10⫺7·298.15·␻ (Tanger and Helgeson, 1988); ° ° ° (b)⌬nS298.15 ⫽ S298.15 ⫺ ⌬sS298.15 and ⌬nCp° 298.15 ⫽ Cp° 298.15 ⫺ ⌬sCp° 298.15; (c)Slop98 database; (d)This study (Table 7); (e)Generated with linear extrapolation to the ligand number n ⫽ 3 (Figure 9).

3.4. Species Distribution Calculations To assess the effect of aqueous europium acetate complexing in natural systems, species distribution calculations were performed for EuCl3 ⫹ NaCl ⫹ NaAc solutions as a function of pH at 25, 100 and 200°C. The total Eu concentration was set at 10⫺3 m, the ionic strength was set in a range of 0.12⫾0.03, and total acetate concentration was set 0.005, 0.05 and 0.15 m. The upper limit of total acetate concentration is close to maximum concentration measured in sedimentary basinal brines (Wood et al., 2000). These calculations were performed with regard to three Eu acetate complexes for which thermodynamic properties are given in Table 7. Note that the Eu(Ac)3° properties were

estimated approximately and the other Eu species (hydroxy or chloride complexes) were not taken into consideration. Results of these calculations are shown as a function of pH at different temperatures and total acetate concentrations in Figure 10. It can be seen that at the total acetate concentration ⱖ0.05 mol·kg⫺1 H2O and pH ⬎4 the free ion (Eu3⫹) is practically absent and Eu is present as Eu acetate complexes through the temperature range (25 to 200oC). Note that with increasing temperature and acetate concentration a portion of Eu(Ac)3° significantly increases up to 50 to 80% at pH ⬎5. However, the thermodynamic properties of this species, were predicted in this study only approximately and

Fig. 9. Correlation of the standard partial molal non-solvation entropies and heat capacities for Eu acetate complexes with the number of acetate ions present in the complexes. The symbols correspond to the calculated values using data in Tables 7 and 8, the solid lines are obtained by a linear regression. Intersections of lines with the vertical axes at n ⫽ 3 give the predicted values of the standard partial molal non-solvation entropies and heat capacity for Eu(Ac)3°.

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Figure 10. Aqueous Eu acetate speciation as a function of pH in EuCl3 ⫹ NaCl ⫹ NaAc solution (ionic strength is ⬃0.12 ⫾ 0.03) at 25, 100 and 200°C, total Eu concentration is 0.001 m and total acetate concentration in a range from 0.005 to 0.15 m. The diagrams were calculated with the GIBBS computer code (Shvarov and Bastrakov, 1999), using thermodynamic properties for Eu acetate complexes taken from Table 7 and those for other species reported in the Slop98 database (see text).

the stability of Eu(Ac) 3° could be overestimated. In any case, at total acetate concentration of ⱖ0.005 m Eu acetate complexes dominate over Eu3⫹ in natural solutions. Relative stability of acetate and carbonate complexes of Eu are compared in Figure 11. Thermodynamic properties of Eu carbonate species (EuHCO3⫹2, EuCO3⫹ and Eu(CO3)2⫺) at room temperature are well known (Slope98 database; Wood, 1990a, 1990b). However, there is no a common agreement about the extrapolation of stability constants to high tempera-

ture. For this reason, two different plots (Figs. 11B and 11C) are constructed for 150°C using the thermodynamic properties of Eu carbonate complexes given by the Slop98 (Haas et al., 1995) or Wood (1990) (Figs. 11B and 11C, respectively). Approximate range of Ac⫺ and HCO3⫺ concentrations in sedimentary basinal brines are shown by rectangle. It can be seen that under these conditions at 25°C (Fig. 11A) the relative stability of acetate and carbonate complexes of Eu is close. However, relative stability of these species with increasing

Eu(III) complexation with acetate

3611

Fig. 11. Predominance fields of Eu⫹3, Eu-acetate and Eu-carbonate complexes as a function of logarithm of the Ac⫺ and HCO3⫺ activity at 25 (A) and 150°C (B and C). Stability constants for Eu-acetate complexes are generated using Table 7. Stability constants for Eu-carbonate complexes at 25°C (A) are taken from Slop98 and at 150°C from also Slop98 (B) or from the alternative study by Wood (1990b) (C). The rectangle denotes the approximate range of conditions found in sedimentary basinal brines.

temperature is highly questionable. The relative stability of Eu-acetate complexes will increase or, on the contrary, slightly decrease according to which thermodynamic properties for Eu-carbonate species are taken (Fig. 11B and 11C, respectively). 4. CONCLUSIONS

Our study shows that aqueous europium is represented almost entirely as Eu-Ac-complexes at acetate concentrations commonly found in natural brines (0.02 to 0.15 m) and at moderate temperatures. Moreover, this example gives us the basis for proposing an important significance of neutral acetate species for the natural transport of REE in general, especially at higher temperatures (above 100°C). Unfortunately, the stability of these neutral species neither for Eu nor other REE has yet been studied experimentally. Future experimental studies need to define the thermodynamic properties of neutral acetate complexes of REE with confidence. Acknowledgments—This study was initiated at the Institute of Ore Deposits while Igor Diakonov was a postdoctoral Fellow in Bristol, funded by the Leverhulme Trust. The study was concluded during a Leverhulme Trust funded visit of Alexander Zotov to Bristol, through the Institute of Advanced Studies. We are grateful for insightful discussions with Eric Oelkers, Silvie Castet, Jacques Schott and Gleb Prokovski. We thank three anonymous reviewers for their careful comments and suggestions resulting in improvement of the manuscript. Funding from the Leverhulme Trust and NERC grant no RE1919 is gratefully acknowledged. Associate editor: W. H. Casey

REFERENCES Amaya T., Kakihana H., and Maeda M. (1973) The hydrolysis of Y3⫹, La3⫹, Gd3⫹ and Er3⫹ ions in an aqueous solution containing 3 M (Li)ClO4 as an ionic medium. Bull. Chem. Soc. Jpn. 46, 1720 –1723. Aziz A. and Lyle S. J. (1969) Equilibrium constants for aqueous fluoro complexes of scandium, yttrium, americium(III) and curium(III) by extraction into di-2-ethylhexyl phosphoric acid. J. Inorg. Nucl. Chem. 31, 3471–3480. Baes C. F. Jr. and Mesmer R. E. (1976) The hydrolysis of cations. Wiley-Interscience. Be´ ne´ zeth P., Castet S., Dandurand J. L., and Schott J. (1994) Experimental study of aluminum acetate complexing between 60 and 200° C. Geochim. Cosmochim. Acta. 58, 4561– 4571. Be´ ne´ zeth P. and Palmer D. A. (2000) Potentiometric determination of cadmium-acetate complexation in aqueous solutions to 250oC. Chem. Geol. 167, 11–24. Borisov M. V. and Shvarov Yu. V. (1992) Thermodynamics of Geochemical Processes. Moscow State University Publishing Office (in Russian). Byrne R. H. and Li B. (1995) Comparative complexation behaviour of the rare earths. Geochim. Cosmochim. Acta 59, 4575– 4590. Cantrell K. J. and Byrne R. H. (1987) Rare earth element complexation by carbonate and oxalate ions. Geochim. Cosmochim. Acta 51, 597– 605. Choppin G. R. and Unrein P. J. (1963) Halide complexes of the lanthanide elements. J. Inorg. Nucl. Chem. 25, 387–393. Choppin G. R. and Strazik W. F. (1965) Complexes of trivalent lanthanide and actinide ions, I. Outer-sphere ion pairs. Inorg. Chem. 4, 1250 –1254. Choppin G. R. and Schneider J. K. (1970) The acetate complexing by trivalent actinide ions. J. Inorg. Nucl. Chem. 32, 3283–3288. Choppin G. R. (1976) Factors in the complexation of lanthanides. Proc. 12th Rare Earth Res. Conf. (ed. C. E. Lundin) 1, 130 –139. Cossy C., Barnes A. C., and Enderby J. E. (1989) The hydration of Dy3⫹ and Yb3⫹ in aqueous solution: A neutron scattering first order difference study. J. Chem. Phys. 90, 3254 –3260.

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