Interaction of uranyl with acetate in aqueous solutions at variable temperatures

J. Chem. Thermodynamics 71 (2014) 148–154

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Interaction of uranyl with acetate in aqueous solutions at variable temperatures Vladimir Sladkov ⇑ CNRS, Institut de Physique Nucléaire (IPN), UMR 8608, Orsay F-91406, France Univ Paris-Sud, Orsay F-91405, France

a r t i c l e

i n f o

Article history: Received 26 July 2013 Received in revised form 9 December 2013 Accepted 11 December 2013 Available online 17 December 2013 Keywords: Uranyl Acetate Complexation Stability constants Temperature effect Thermodynamic parameters Affinity capillary electrophoresis

a b s t r a c t Increasing activities in the environmental management of nuclear wastes incite significant interest in the study of the interaction of actinides with organic matter in aqueous solution, especially at elevated temperature. The system U(VI)-acetic acid is studied in aqueous solutions at pH 2 in the temperature range from 15 °C to 55 °C by affinity capillary electrophoresis (ACE). The formation of two complex species UO2CH3COO+ and UO2(CH3COO)2 is observed. Thermodynamic parameters (the molar Gibbs energy of reaction (DrGm), the molar enthalpy of reaction (DrHm) and the molar entropy of reaction (DrSm)) are determined at fixed ionic strength of 0.05 mol  L1 (NaClO4–HClO4) and calculated at 0 ionic strength with extended Debye–Hückel equation for the activity coefficients. Obtained results are compared with literature data. Ó 2013 Elsevier Ltd. All rights reserved.

1. Introduction The interaction of U(VI) with natural organic matter is of great interest. It deals not only with a fundamental knowledge, but also with the migration of U(VI) species in the environment and their distribution in the geosphere and biosphere [1–3]. Knowledge about the impact of natural organic matter on U(VI) migration is required for the long term risk assessment for nuclear waste repositories, for facilities of the former uranium mining and milling sites [4–7]. Humic acids are the principal components of humic substances, which are the major constituents of organic matter. Humic acids are characterized by heterogeneous structural and functional properties. They are composed of aliphatic and aromatic structural elements and a variety of different, mainly oxygen containing functional groups [4]. The most simple model substances for humic acids represent low-molecular short chain organic acids which can occur as humic acid building blocks and simulate their structure and functionality. The acetic acid is the smallest organic acid with a carbon chain. Among the different geochemical processes, aqueous complexation can also control the transport behaviour of radionuclides. Moreover, the most probable transport media is aqueous phase ⇑ Address: CNRS, Institut de Physique Nucléaire (IPN), UMR 8608, Orsay F-91406, France. Tel.: +33 169156406; fax: +33 169156470. E-mail address: [email protected] 0021-9614/$ - see front matter Ó 2013 Elsevier Ltd. All rights reserved. http://dx.doi.org/10.1016/j.jct.2013.12.011

[6,8–11]. In this project we are interested in the complexation of U(VI) with acetic acid in aqueous solutions at various temperatures. The interaction of U(VI) with acetic acid at room temperature has been studied in several projects. In paper [12] the data from these projects are reviewed. The three complex species of uranyl with acetate are evidenced (UO2Ac+, UO2Ac2 and UO2(Ac)3). The logarithms of the thermodynamic stability constants (b°) are log b°1 = 2.94 ± 0.08 [12], log b°2 = 5.50 ± 0.15 [12] and log b°3 = 7.25 [13]. Very few papers are devoted to study of uranyl interaction with acetic acid at various temperature [14–18]. However, actually the increasing activities in the environmental management of nuclear wastes incite significant interest in the study of the chemistry of actinides in solution, especially at elevated temperature. Usually, potentiometric and calorimetric methods are used to determine the thermodynamic parameters. The most studies of U(VI)-acetate interactions are performed at high ionic strength (mostly at 1 M) [14,16–18]. To the best of our knowledge there is no the standard thermodynamic parameters communicated (the molar enthalpy (DrH0m), and the molar entropy (DrS0m) for complex formation reactions) at zero ionic strength. In this project we study of uranyl interaction with acetic acid at various temperatures (from 15 °C to 55 °C) by means of affinity capillary electrophoresis (ACE). Affinity capillary electrophoresis methods are successfully used for the determination of stability constant values for different systems in the last years [12,19–40].

V. Sladkov / J. Chem. Thermodynamics 71 (2014) 148–154

This method attract attention due to a number of advantages, including high resolution, speed of analysis, low running costs, minimal sample consumption and aptitude for automation, among others [41]. The recent achievements in this field are reviewed [40,42,43]. Special review is devoted to the actinide speciation by capillary electrophoresis [41]. Due to the stability constants obtained from electrophoretic mobility measurements at different temperatures, the thermodynamic parameters can be determined. The values of stability constants, thermodynamic parameters, such as the molar Gibbs energy (DrGm), the molar enthalpy (DrHm), and the molar entropy (DrSm) for complex formation reactions are calculated at ionic strengths of 0.05 mol  L1 and extrapolated to zero ionic strength with aid of extended Debye–Hückel equation for the activity coefficients [44]. 2. Experimental section 2.1. Chemicals and solutions All chemicals used are of analytical reagent grade. The stock UO2(ClO4)2 solution (0.1 M in 0.63 M HClO4) is obtained by dissolving UO2(NO3)2  6H2O (>99% FLUKA puriss) in 12 M HClO4 (MERCK Suprapur) and evaporating the resulting solution to almost dryness on a sand bath. The residue is dissolved in concentrated HClO4 and evaporated again. This last operation is repeated three times. The uranium concentration is checked by PERALS spectrometric measurements described elsewhere [45]. Glacial acetic acid (P99.99%) is supplied by Sigma–Aldrich. Concentrated perchloric acid (60% solution from Sigma–Aldrich) is diluted in water to the requisite concentration. The exact concentration is determined by acid–base titration with certified NaOH solution. Sodium perchlorate (99%) is provided by Merck. The impurity content in this chemical is low enough (Sulphate and chloride contents are at ppm level.) and does not have any impact on results obtained. Moreover the work is done with diluted solutions. The exact concentration of diluted solution is determined by gravimetric method. All solutions are prepared with deionised water (Millipore direct Q, R = 18.2 MX). 2.2. Apparatus and software Capillary Electrophoresis measurements. P/ACE system MDQ capillary electrophoresis instrument (Beckman Coulter, France) is used. The system is comprised of 0–30 kV high-voltage built in power supply, equipped with a UV–vis spectrophotometric diode array detector. UV direct detection at 200 nm is used in this work. A capillary (50 lm I.D., 363 lm O.D.) made from fused silica is obtained from Beckman Instruments and had a total length (Lt) of 31.2 cm and an effective separation length (Ld) of 21 cm. The capillary is housed in an interchangeable cartridge with circulating liquid coolant. The length of non-thermostated capillary inlet (Linlet) is about 4.2 cm (about 20% of the effective separation length). The ambient temperature is about 20 °C. The data are obtained at the standard state pressure (0.1 MPa). Every measurement is repeated at least three times. Data acquisition and processing are carried out with Karat 32 software (Beckman Coulter, France). All concentration and constant calculations are done with the EXCELÒ and ORIGINÒ software programs. The solver module program is used for fitting experimental points by least squares curve fitting. pH measurements. A pH-meter GLP-21 (Crison, France) and a combination electrode are used for pH measurements after calibration against NIST standards (4.01 and 7.00). An aliquot of solution is used for each measurement. The pH values of solutions is 2.00 ± 0.05.

149

2.3. Procedure The capillary is conditioned prior to use by successive washes with 0.1 M sodium hydroxide, deionised water and the buffer solution under study. It is rinsed for 2 min (at a pressure of 103.4 kPa) with the buffer between two runs and kept filled with deionised water overnight. In order to avoid hydrolysis and/or polymerisation of the uranyl ion, that could potentially lead to the formation of additional species and/or to the modification of uranyl mobility values, we use the perchloric acid–sodium perchlorate aqueous solutions at pH 2.0 as background electrolyte (BGE). At such values and uranyl concentration of 1  104 M, the contribution of hydrolysed and polymeric species to U(VI) is negligible [46]. The ionic strength (I) of BGE is 0.05 mol  L1. The relative standard deviation estimated is less than 5%. The pH is controlled with pH-meter (see Section 2.2). The normal polarity mode is applied (injection is performed at the positive end). The sample, containing only uranyl ions and the neutral marker, is injected in the capillary. The capillary contains BGE (perchloric aqueous solution at pH 2.0) with fixed concentrations of acetic acid. Sets of runs are then made for different concentrations of acetic acid (from 0 to 0.8 mol  L1) in BGE. The separation voltage applied is 5 kV and the injection time (by a pressure of 3.45 kPa) is 4 s. This value of applied voltage is chosen to respect the Ohm’s law. The current is about 40 lA and the input power is about 0.2 W. In these conditions we can estimate that the effect of Joule heating is insignificant in the thermostated part of capillary and the deviation from the desired temperature is 61 °C [19,47]. The separations are performed at constant forward pressure of 1.4 kPa. 2.4. Data treatments 2.4.1. Electrophoretic mobility determination The electrophoretic mobility l (cm2  s1  V1) can be calculated by using the following expression:

l ¼ Lt Ld

1=t  1=teof ; U

ð1Þ

where Lt (m) is the total capillary length, Ld (m) is the length between capillary inlet and the detection window, U is applied voltage in V, t (s) is the migration time of the studied species, teof (s) is the migration time of dimethylsulfoxide (DMSO), used as a neutral marker for electroosmotic flow mobility determination. In this work the hydrodynamic transfer of the sample through the non-thermostated inlet into the thermostated region of capillary is used to avoid the influence of non efficiently thermostated short inlet [23]. To calculate the ion mobility for this mode, expression (1) is modified. In this case Ld is no longer the length between the capillary inlet and the detection window, but is now the length between the position of the sample in the capillary thermostated zone and the detection window (L0d ). To find L0d we measure the elution time of marker (DMSO), using only forward pressure. In this case we can write:

tp ¼

pd2 L0d 4tm

;

ð2Þ

where tp is volumetric flow rate (generated by forward pressure, m3/s), d is the capillary diameter (m), tm is the elution time of marker after the hydrodynamic transfer through the non-thermostated inlet (s). The L0d can be calculated with aid of Poiseuille equation for laminar viscous and incompressible flow and the expression (2): 2

L0d ¼

DPd t m ; 32gLt

ð3Þ

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V. Sladkov / J. Chem. Thermodynamics 71 (2014) 148–154

where DP is the applied pressure (Pa), g is the dynamic viscosity (Pa  s). Knowing the water viscosity at the given temperature [48], we can calculate Ld’ with expression (3). It should be noted that here we use the aqueous BGE, consisting mainly of sodium perchlorate electrolyte with an ionic strength of 0.05 mol  L1. Such concentrations of sodium perchlorate do not modify the viscosity of water [48,49] and we can suppose that the viscosity of the BGE is the same as that of the water. L0d is calculated as the average value from the 10 measurements for each temperature. The time chosen for the hydrodynamic transfer of the sample through the non-thermostated inlet is 30 s (under the pressure of 6.9 kPa). With these parameters, as it has been found, we can be sure that the sample is present in the thermostated part of capillary at all temperatures studied (i.e., L0d < Ld  Linlet). The viscosity of aqueous acetic acid solutions is not constant with increasing of acetic acid concentration from 0.1 M to 0.8 M. This fact should be taken into account, as the ion mobility depends on the media viscosity (exp. 4):

l¼

q ; 6pgr

ð4Þ

where q is the charge of the ionized solute and r is the solute radius. The mobilities calculated with modified expression (1) needs to be corrected [12]. The viscosity is changed from 0.89 mPa  s for aqueous solution [48] to 0.99 mPa  s (calculated with equation given in [50]) for 0.8 M acetic acid aqueous solution (25 °C). Thus the mobility values calculated with modified expression (1) will be referred to the media with different viscosities. To take it into account, the mobility values (l) obtained with expression (1) are corrected with coefficient k = g0 /g, where g0 refers to the acetic acid aqueous solution and g refers to the aqueous solution.

l ¼ kl0 : The viscosity values for aqueous solutions with different concentrations of acetic acid at different temperatures are calculated with equation given in [50]. The relative standard deviation (Sr) for measured mobility values is in the interval 0.5–5% (n = 3, P = 0.95). 2.4.2. Complex species mobility estimation at different temperatures In order to estimate the complex species mobility values we used the expression proposed in the paper [51], based on the relationship between the limiting conductivity k0 of an ion and the Stokes radius and a postulate that the volume of a hydrated complex ion is, to a first approximation, equal to the sum of the hydrated volumes of the constituent simple ions.

k0complex ¼

jzcomplex j : P 3 1=3 ½ ðzi =k0i Þ 

ð5Þ

k0complex and zcomplex refers to the limiting conductivity and the charge of the complex, k0i and zi refers to the limiting conductivities and the charges of the species i, constitutive of the complex. By applying the relation between the limiting conductivity and the electrophoretic mobility, given by

l0 ¼

k0 ; F

ð6Þ

where F is the Faraday constant, estimation of complex species mobility is possible [20]. The values of individual ion mobilities used for the calculation of the complex species mobilities with the expression (5) and (6) were taken from papers [52–55]. The values of limiting conductivities for uranyl at various temperatures

have not been found in the literature. They are estimated with aide of Walden’s rule [56]. For extrapolation to required ionic strength the equation based on Debye–Huckel–Onsager (DHO) limiting law and Pitts’ equation were used [57,58]. 2.4.3. Thermodynamic data treatment For small temperature intervals (between 0 and 50 °C) changes of equilibrium constants can be calculated with the constant heatcapacity method [59]:

  Dr HðT 0 Þ 1 1 log bðTÞ ¼ log bðT 0 Þ þ  R ln 10 T 0 T    Dr C p T 0 T ; þ  1 þ ln T0 R ln 10 T

ð7Þ

where R is the molar gas constant, T is the experimental temperature and T0 is the temperature at the standard state (25 °C), DrCp – heat capacity of complexation. It is assumed, that DrCp is constant. The results can also be treated with Van’t Hoff plot. A plot of the natural logarithm of the equilibrium constant versus the reciprocal temperature gives a straight line. In this case DrCp is near zero and it can be neglected.

ln b ¼ 

Dr Hm Dr Sm þ : RT R

ð8Þ

2.4.4. Ionic strength corrections Extended Debye–Hückel expression. As a weak ionic strength is used (0.05 mol  L1), the extended Debye–Hückel expression (9) is applied for activity coefficient (c) calculations.

log ci ¼ z2i A

I1=2 m 1=2 1 þ Bai Im

:

ð9Þ

A and B are constants which are temperature dependant, and ai is the effective diameter of hydrated ion i. The values of A and B as a function of temperature are given in [44] In our paper the value Bai is kept constant at 1.5 kg1/2  mol1/2. The error of this simplification is insignificant in our experimental conditions. Im is the ionic strength in molality units (mol  kg1). Specific ion interaction theory (SIT). By applying the specific ion interaction theory (SIT) [46] the values of stability constants obtained at different ionic strengths can be extrapolated to I = 0 (b°). Especially this approach is useful, when high ionic strengths are used (more than 0.1 mol  kg1). The following formula is used for the extrapolation to I = 0:

log b þ DZ 2 D ¼ log b  DeIm ;

ð10Þ

MZM þ nLZL ¡MLZnM nZL ;

ð11Þ

where b is the stability constant at the given ionic strength, Dz2 = (zM  nzL)2  zM2  nzL2, where (zM  nzL), zM, and zL are the charges of MLn, M and L species, respectively (equation (11)), p p D = A Im/(1 + Ba Im) is the Debye–Huckel term, A = 0.5091 kg1/2  1/2 mol at 25 °C, Ba = 1.5, De is the difference of interaction solute coefficients. For ionic strength corrections at the temperature other than 298 K, values the Debye–Huckel parameter A are taken from the table given in [11]. 3. Results and discussion 3.1. Determination of thermodynamic parameters (DrGm, DrHm, DrSm) The electrophoretic mobility of U(VI) increases with increasing of temperature [12], as the mobility value is inversely proportional of media viscosity (exp. (4)). With the temperature increasing the

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V. Sladkov / J. Chem. Thermodynamics 71 (2014) 148–154

A

0.8 M

1 mAU

0.6 M 0.5 M 0.4 M 0.3 M

 UO2þ 2 þ 2CH3 COO ¡UO2 ðCH3 COOÞ2 :

The complexation stability constants are

b1 ¼

0.2 M 0.1 M 0.05 M 0.01 M

b2 ¼

0.8 M 0.6 M 0.5 M 0.4 M

1 mAU

0.3 M 0.2 M 0.1 M 0.05 M 0.01 M 0 0.5

1.0

1.5

2.0

Time. min FIGURE 1. Electropherograms of 1  104 M U(VI) at different concentration of acetic acid in BGE at 15 °C (A) and at 55 °C (B). pH = 2.0. I = 0.05 mol  L1. For other experimental conditions see Section 2.2.

water viscosity decreases (from 1.139 cP at 15 °C to 0.5040 cP at 55 °C [48]). In figure 1 the electropherograms of 1  104 M U(VI) in the presence of various acetic acid concentrations in BGE are presented at 15 °C (figure 1A) and 55 °C (figure 1B). Only one peak is observed, whose mobility is diminished with the increase of ligand concentration in BGE. We deal here with kinetically labile complex system [12,40]. The rate of ligand exchange between uranyl species is sufficiently fast compared to the separation time. The mobility observed corresponds to the sum of mobilities of coexisting different charged species of U(VI). The broadening of peak, observed with increasing acetate concentration (figure 1), reflects equilibration between U(VI) species having different electrophoretic mobilities. U(VI) is present as uranyl (aquacomplex UO2(H2O)5)2+), whose mobility is positive, in the absence of acetic acid in BGE. In the presence of acetic acid the U(VI) mobility is diminished, as the mobility of the first complex species formed (UO2CH3COO+) is smaller and the mobility of the second complex species formed UO2(CH3COO)2 is zero. The mobility of U(VI) can be calculated from the weighted average of the mobilities of the respective species [60]:

lUðVIÞ ¼

X

li ai ;

ð12Þ

where li is the mobility of the respective U(VI) species and ai is its molar fraction. We can assume the following complexation equilibriums in the studied range of acetate concentration:  þ UO2þ 2 þ CH3 COO ¡UO2 CH3 COO ;

½UO2 ðCH3 COOÞ2 

ð16Þ

:

ð13Þ

lUðVIÞ ¼ a0 lUO2þ þ a1 lUO2 Acþ ; 2

ð17Þ

where a0 = [UO22+]/CU(VI) = 1/(1 + b1[CH3COO] + b2[CH3COO]2) and a1 = [UO2CH3COO+]/CU(VI) = b1[CH3COO]a0. (CU(VI) is the total concentration of U(VI) species). Experimental values of mobility, obtained at different ligand concentrations, are fitted with the expression (17). The obtained experimental points and fitting curves are presented in figure 2. The values of stability constants b1 and b2 are varied to obtain the best fit. The mobility of complex species UO2CH3COO+ is estimated with expressions (5) and (6). The set of experimental values of U(VI) mobilities are obtained at different temperatures. The acetate concentrations are calculated taking into account the dissociation constants of acetic acid at given ionic strength and at given temperature [44]. The values of dissociation constants for acetic acid at different temperatures are taken from paper [61]. The obtained stability constants values at different temperatures are presented in table 1. From this table we can see that the complexes become stronger with the temperature increasing. The Van’t Hoff plots (ln b vs 1/T) (exp. (8)) are constructed for the values of the first and the second stability constants (figure 3). The fairly-good linear fits in this plot suggest that the enthalpy of complexation in this temperature range (15 to 55 °C) can be assumed constant. In other words, the heat capacity value of complexation (DrCp) in this temperature range is near zero and can be neglected. Indeed, if we use expression (7) for determination of DrHm and DrSm and we neglect the contribution of second order term, we obtain the same values of thermodynamic parameters. The thermodynamic parameters obtained at ionic strength of 0.05 mol  L1 are presented in table 2.

7 6 5 -1 -1

6

Time. min

B

ð15Þ

;

 2 ½UO2þ 2 ½CH3 COO 

55 °C

45 °C

35 °C

4

2

5

8

4

 ½UO2þ 2 ½CH3 COO 

µ×10 , m V s

3

½UO2 CH3 COOþ 

The expression for mobility values observed (derived from (12)) is

0 2

ð14Þ

25 °C

3 15 °C

2 1 0

1E-5

1E-4

1E-3 -

[CH3COO ], M FIGURE 2. Experimental points (solid symbols) for U(VI) mobilities and fitting curves (dashed lines) as a function of acetate concentration at different temperatures. pH = 2.0. I = 0.05 mol  L1. CU(VI) = 1  104 M. For other experimental conditions see Section 2.2.

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V. Sladkov / J. Chem. Thermodynamics 71 (2014) 148–154

TABLE 1 Stability constant values obtained at different temperatures. I = 0.05 mol  L1 (NaClO4-HClO4, pH 2.0). Temperature (°C)

log b1

log b2

15 25 35 45 55

2.44 ± 0.10 2.57 ± 0.08 2.63 ± 0.10 2.71 ± 0.10 2.73 ± 0.15

4.88 ± 0.10 4.97 ± 0.15 5.14 ± 0.15 5.24 ± 0.15 5.41 ± 0.15

14

I=0

UO2(CH3COO)2

12

I=0,05

lnβ

10

8

I=0

6

+

UO2CH3COO

I=0,05

4 3.0

3.1

3.2

3.3

1/T×1000, K

3.4

3.2. Comparison with literature data

3.5

-1

FIGURE 3. Van’t Hoff plot (ln b vs. 1/T) for the complexation of uranyl with acetate. Solid symbols – experimental points (pH = 2.0), solid line – linear fit and dashed lines – upper and lower limits of the confidence band at the 95% level.

In order to determine the thermodynamic parameters at I = 0, the values of stability constants obtained at ionic strength of 0.05 mol  L1 are calculated at zero ionic strength [44] with the aid of expressions ((9), (18), and (19)) at each temperature.

b1 ¼

b2 ¼

b1

cUO2þ 2

ð18Þ

;

b2

cUO2þ c2Ac 2

These data show that the enthalpy and entropy of complexation are positive in this temperature range. The complexation is entropy driven, characteristic of ‘‘hard acid’’ and ‘‘hard base’’ interactions and inner-sphere complexation [62]. Dehydratation of uranyl and acetate plays important roles in the complexation reaction. The energy required for dehydratation contributes to the positive enthalpy change and the number of water molecules released from hydration sphere contributes to the positive entropy change. The complexation is enhanced at higher temperature because the increase in the entropy term (TDS) exceeds the increase of enthalpy. The magnitude of enthalpy and entropy is informative of the denticity in complexes, especially the entropy because it is directly related to the number of water molecules that the ligand replaces [63]. For example, the entropy values for UO2Ac+ and UO2Ac+ are comparable to those of known bidentate inner-sphere complexes such as the uranyl-sulfate (126 ± 8 J  K1  mol1 and 222 ± 28 J  K1  mol1) [46] and uranyl-malonate (130 ± 2 J  K1  mol1 and 256 ± 14 J  K1  mol1) [64]. The entropy of complexation for a monodentate U(VI) complex would be much smaller . When the ligand is monodentate, as the third ligand in the uranyl-acetate complex, the stepwise entropy of complexation is only about 40 J  K1  mol1. [64]. The thermodynamic data support that the acetate ion is bidentate. This was demonstrated by IR spectroscopy, DFT calculations and EXAFS measurements [65].

ð19Þ

:

The Van’t Hoff plots are constructed for zero ionic strength (figure 3). The obtained values of DrG°m, DrH°m and DrS°m are given in table 2.

The obtained thermodynamic values are compared with the literature data (table 2). The values of thermodynamic parameters obtained in the present paper are higher, than the ones obtained in the literature [14,16–18]. To compare thermodynamic stability constants obtained in this paper with the ones taken from literature data, the graphs log b1 + 4D for the first constant and log b2 + 6D for the second constant against Im (equation (10)) are constructed (figure 4). It follows to keep in mind, that these graphs are constructed with data obtained with perchloric or chloric media. The fairy good linear correlations are observed, although the aqueous media is not the same for the points used. The correlations coefficients are 0.972 for the first constant and 0.9196 for the second constant. From these graphs the log b°1 is found to equal 2.99 ± 0.04 and log b°2 is found to equal 5.23 ± 0.15. DrG°m value for the first equilibrium (13) is (17.0 ± 0.3) kJ  mol1 and for the second equilibrium (14) is (29.8 ± 0.9) kJ  mol1. These values are very close to the ones obtained only with ACE data and extrapolated with Debye–Hückel equation (table 2). The

TABLE 2 Thermodynamic parameters for the equilibriums UO22+ + CH3COO ¡ UO2CH3COO+ and UO22+ + 2CH3COO ¡ UO2(CH3COO)2. T = 298 K. Ionic strength, M (electrolyte used)

i

log bi

DrGm(298 K) (kJ  mol1)

DrHm (kJ  mol1)

DrSm,(298 K) (J K1 mol1)

Method

Reference

0

1 2

2.94 ± 0.08 5.50 ± 0.15

16.8 ± 0.4 31.4 ± 0.8

14.5 ± 1.5 25 ± 2

104 ± 6 187 ± 5

ACE

p.w.

0.05 (NaClO4)

1 2

2.57 ± 0.10 4.97 ± 0.15

14.7 ± 0.5 28.3 ± 0.9

13.7 ± 1.5 23 ± 2

95 ± 5 174 ± 5

ACE

p.w.

60.1

1

3.00 ± 0.06

17.2

20.1

125.4

Pot.

[15]

1 (NaClO4)

1 2

2.46 ± 0.01 4.38 ± 0.02

14.00 ± 0.04 25.0 ± 0.1

11.8 ± 0.1 17.9 ± 0.3

86.5 ± 0.8 144 ± 1

Pot., cal.

[16]

1.05 (NaClO4)

1 2

2.58 ± 0.03 4.37 ± 0.14

14.9 25.3

10.6 ± 0.8 20 ± 3

86 ± 3 152 ± 13

Pot., cal.

[17]

1 (NaClO4)

1 2

2.42 ± 0.02 4.42 ± 0.03

13.81 ± 0.09 25.2 ± 0.3

10.54 ± 0.10 20.2 ± 0.3

81.7 ± 0.5 152.3 ± 1.5

Cal.

[14]

1 (NaClO4)

1 2

2.56 ± 0.02 4.08 ± 0.06

14.6 ± 0.1 23.3 ± 0.3

10.5 ± 0.8 18.6 ± 4.0

84.2 ± 2.7 140.5 ± 13.5

Pot., cal.

[18]

Abbreviations: pot. – potentiometry, cal. – calorimetry, p.w. – present work. in italic – molality units (mol  kg1) are used.

V. Sladkov / J. Chem. Thermodynamics 71 (2014) 148–154

A

eventual dispersion of hazardous materials. To predict this dispersion reliable thermodynamic data are indispensable. To understand the mechanism of interaction of complicated natural organic matter with U(VI), study of complexation of this radionuclide with acetic acid is performed at different temperatures. Although acetic acid is very important organic species, there is a lack of data on thermodynamic properties of its complex species.

4.5

4.0

logβ1+4D

153

3.5 3

Acknowledgment

4 5,6

I am grateful to Dr. Claire Le Naour for valuable remarks and fruitful discussion.

3.0 1,2

References 2.5

0

1

2

3

4

5

Im

B 7.5 7.0

logβ2+6D

6.5

6.0

3,5,6

1,2

5.5

4

5.0

4.5

0

1

2

3

4

Im FIGURE 4. Extrapolation to I = 0 of experimental data for the formation of UO2CH3COO+ (A) and UO2(CH3COO)2 (B) using the specific ion interaction theory. The points refer to NaClO4 media – 1–6, the others points refer to NaCl media. The experimental points are taken: 1, 2 – from this work and [12], 3 – [17], 4 – [18], 5 – [16], 6 – [14] and other points – from [67].

confidence intervals overlap for the values obtained by different manner. The difference between these values may be insignificant, especially for the first constant. The De (kg  mol1) obtained from slops of lines are 0.26 ± 0.02 (figure 4A) and 0.48 ± 0.08 (figure 4B). If we estimate De for the reactions (13) and (14), we obtain 0.14 and 0.37, respectively for chloric media, and 0.27 and 0.62 for perchloric media. We can note that the values obtained from the slops of lines in figure 4 have intermediate magnitudes between the values calculated for chloric and perchloric media. The values eUO2þ ;ClO , eUO2þ ;Cl and eAc ;Naþ are taken from [46]. The 4 2 2 values of eUO2 Acþ ;ClO4 and eUO2 Acþ ;Cl are estimated as the average of eUO2þ and eAc ;Naþ , eUO2þ ;Cl and eAc ;Naþ [66]. The comparison by ;ClO 4 2 2 the same manner DrH°m and DrS°m, using the dependence of this parameters from ionic strength, is not performed due to the lack of experimental data at different ionic strength. Only data obtained with ionic strength of 0.05 and 1.05 mol  kg1 are available (table 2). 4. Conclusion The modelling of the behaviour of hazardous materials under environmental condition is important task for the environment protection. Among the geochemical processes, the complexation of radionuclide with natural organic matter can determine the

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JCT 13-447