The Experimental Determination of the Solubility Product for Np020H in NaCl Solutions.

Kevin E. Robert$?* , Herbert B. Silberl93, Philip C. Torrettol,2, Traudel PrussinlJ, Kevin Becraftl, David E. Hobartl, and Craig F.

Keywords: Neptunium/Solubility/Waste Isolation Pilot Plant (WIPP)/Sodium chloride brines

ABSTRACT The solubility of NpOzOH(am) was measured in NaCl solutions ranging from

0.30 to 5.6 molal at room temperature (-21 * 2OC). Experiments were conducted from undersaturation and allowed to equilibrate in a COz-free environment for 37 days. The apparent solubility products varied with NaCl concentration and were between 10-’.and lo4 mo12eL-2. Using the specific ion interaction theory (SIT), the log of the solubility product of NpOzOH,,, at infinite dilution was found to be - 8.79

f 0.12. The interaction coefficient, &(NpQ+ - Cl-), was found to be (0.08 * 0.05).

INTRODUCTION The chemical behavior of actinides under environmental conditions is among

the most important phenomena in the geologic disposal of nuclear waste. Thus, efforts are in progress to characterize the solubilities and speciation of the actinides under chemical conditions that may occur in and around the repositories. The solubilities of neptunium salts have been studied under environmental conditions 11-61 and in simpler salt solutions at ionic strengths usually less than one [7-141. However, thermodynamic information in more concentrated salt media is more limited. It is necessary to build a table of solubilities of selected actinides in different media as a function of salt concentrations [E]. As part of this database, we initiated this study of the solubility of Np(V) in 0.30 to 5.6 molal NaCl solutions.

Actinide Geochemistry Group, Earth Sciences Division, Lawrence Berkeley National Laboratory, Berkeley, CA 94720 USA Lawrence Livermore National Laboratory, PO Box 808 L-231, Livermore, CA 94551, USA Address correspondence to this author at Lawrence Livermore National Laboratory. Sandia National Laboratories, MS 1320, P.O. Box 5800, Albuquerque, N M 87185-1320 USA

*

DISCLAIMER

Portions of this document may be illegible in electronic image products. Images are produced from the best available original document.

EXPERIMENTAI

Determination of Hydrogen and Hydroxide Ion Concentrations

In dilute electrolytic solutions, the “pH’ read by an electrode, or pH,, calibrated in NIST traceable buffers of known proton activity, {H+}, is by definition equal to the negative log of the proton activity in that electrolytic solution. This is true assuming the electrolytic solution is similar t o the buffers in ionic strength and that the junction potential of the electrode is the same in the buffers and in the electrolytic solution. However, in concentrated electrolytes, this electrode junction potential and the activity coefficient of the proton change with the composition and

the ionic strength of the salt solution. The pH, is shifted by a quantity, ApH, from

the negative log of the proton activity because of the change in junction potential. We use the following relation to incorporate this shift in measured pH due t o the junction potential

p{H+} = pH, + ApH. (1)

Thermodynamics relates the -log of the proton activity, p{H+}, to the -log of the proton concentration, p [H”] , with the following equation

p{H+I = p[H+l- log yH+ . (2)

We combine equations (1) and (2) to relate p[H+] t o the value read by the electrode, PHI.,

pE+l = pH, + (log yH+ + ApHh (3)

By titrating standardized solutions of HC1 with standardized solutions of NaOH, both being adjusted to the desired total ionic strength with NaC1, we relate a known pE+] to the electrode’s pHr in the acidic region of the titration. By plotting p[H+] linearly with a slope of one as function of p&, we can extrapolate to the ordinate for

an intercept value that is equal to (logyH+ + ApH). With this experimental procedure

we cannot separate the log of the proton activity coefficient from the shift in pH due to the junction potential of the electrode.

2

Beyond the equivalence point in the titration, we no longer know the p[H+],

but rather we know the p[OH-1, the -log of the hydroxide concentration. Now, we

relate p[OH-] to PHr. Using the thermodynamic relations for the dissociation of

water and substituting the appropriate expression for the pH+] into equation (3) and rearranging, we obtain

1

-p[OH-l = pH, + (log KO, - log + log h k r + ApH), (4) where KOw is the thermodynamic dissociation constant for water at ilrfinite dilution

and awater is the activity of water. We plot -p[OH-] linearly as a function of pH,,

extrapolate to the ordinate, and obtain an intercept equal to (log Kow - logyo, + log

awater + ApH). Again, we cannot separate the individual quantities in the intercept.

However, the difference between these two intercepts, (logH+ + ApH) - Oog Kow -

10gyOH- + log + ApH), is equal to pM+] - (-p[OH-]) which is also the -log of the

ion product of water, or pK',, in that sodium chloride solution. We combine both

acidic and basic data to evaluate for an intercept, (logy,, + ApH), that defines the

shift from p[H+] to the value read by the electrode, p y , due to the activity coefficient and the liquid junction potential. "his method is similar to that of others working in concentrated salt solutions [11,12,16].

Using this procedure, we calibrated two Ross combination pH electrodes (Orion Research Co.) filled with 3 M NaCl for p[H+] (mo1.L-l) and determined the pK', (mo12.L-2 and then molal') of water in 0.01, 0.03, 0.1, 0.3, 0.6, 1.0, 1.8, 2.3, 3.0, 4.1, and 5.6 molal sodium chloride solutions using single titrations. The [OH-] concentrations were then determined from the evaluated p[H'] values and the experimentally determined pK', values.

Solution Preparation

Green, amorphous neptunium(V) hydroxide solid was precipitated fkom ionic Np02+ with the addition of CO,-fkee NaOH [ll]. Absorption spectrophotometry showed that the neptunium in the initial stock solution was in the +5 oxidation state. To approximately 100 mg of neptunium, we added enough 1 M NaOH to form

3

a green solid suspension. We immediately started six solubility experiments from undersaturation by placing -8 mg of freshly precipitated neptunium0 hydroxide into approximately 80 mL of 0.3, 0.6, 1.0, 1.8, 3.0, and 5.6 molal sodium chloride solutions. The solubility experiments were contained in polyetheretherketone (PEEK) cells (materials obtained from Cadillac Plastics, San Francisco, California and machined to specifications at Lawrence Berkeley National Laboratory). The six cells were sealed, placed on a shaker table, and shaken continuously.

Solubility Measurements

The aqueous phase was separated from any solids or suspended particles using Centricon-30 centrifugal filters (Amicon Corporation, Danvers, Massachusetts) containing a YM-type membrane with a calculated pore size of 4.1 nm. Routine separations were performed after presaturating the filters with 500 pL,

of solution. Concentration measurements of the aqueous portions were made by y-

spectroscopy with a low-energy germanium counting system of Lawrence Berkeley National Laboratory design, calibrated bi-weekly with neptunium standards. The average error in determining dissolved neptunium concentrations is 6 to 7%, including errors in counting, dilutions, and the standard calibration samples. The solution pH, was measured immediately after sampling the test solutions for neptunium.

RESULTS AND DISCUSSION

Figure 1 is a comparison of our results for the ion product of water with literature values near room temperature [17]. Our pKL results are within 2 0.02 units of the 20°C values of Harned and Mannweiler up to 3 molal NaCl [17], indicating that our pH measurement method is reasonable.

N p O hydrolysis is unimportant below pH 10 1113 and the solubility experiments cited here never exceeded p[H+] 10.0. Spectroscopic measurements confirmed the absence of CO, in these samples through day 37. The apparent

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solubility product, mo12.L-2, of the amorphous Np0,OH solid in these solutions is given by

K, (1,) = [NpO,+I [OH-] (5 )

for the dissolution reaction

NpO,OH(am) u NpO; +OH-. (6) The measured apparent solubility products, mof.L-2, as a fhc t ion of time for each NaCl concentration are plotted in Figure 2. Over the 37 days of the experiments, the solubility products remained fairly constant but did begin to show signs of decreasing.

In similar experiments in NaC10, media, Neck et al. [ l l l observed amorphous Np020H solid in 0.1 M NaClO,, both amorphous Np0,OH and an “aged” white Np0,OH solid in 1.0 M NaClO,, and only the aged, white Np020H solid in 3.0 M NaClO,. In a comparison between a thermodynamic model derived in part from the data of Neck et al. [ll] and our experiments, good agreement between our data and the model was achieved by assuming the solid phase as NpO,Oq,, [El . Although this suggests that the aging behavior Np0,OH solid is slower in NaCl media than in NaC10, media, the solid phases at the end of the experiment were not characterized because the amorphous solid is unrecognizable by x-ray powder =action.

Table 1 lists steady-state values for -log KJI,), and -logK&J, for each of the solubility experiments. The reported values are the averages of the seven determinations made in each experiment over the 37 days of reaction time with an uncertainty of two times the standard deviation.

We evaluated for the solubility product at infinite dilution, K$,=O), using the specific ion interaction theory (SIT) [18,19,20], where

log Ksp(Im = 0 ) = log Ksp(Im)+ log(YNpo;)+log(Y 1. (7) OH-

According to SIT, the activity coefficient of ionj , %, with charge zj in a solution of

ionic strength I , is given by

is the Debye-Huckel term at 20°C and E( j,k) is the 0.5050& 1+ 1 . 5 G

where D =

1 interaction coefficient of ion j with ion k. The summation extends over all ions k present in the solution. The interaction coefficients for ions of the same charge sign

I are zero. Introducing equation 8 into equation 7 for the activity coefficients and

(9)

where A& = &(Np02+ - C1-) + E(Na+ - OH-). From the plot of ‘log Ksp(IJ - 20” as a

function of ionic strength (molal), we obtain the solubility product at infinite

dilution, K,(I,=O), as the intercept and -A& as the slope. As outlined by Grenthe et

al. [211, we used a l/(20? linear regression, of the ‘log KJIJ - 20” values for all six

ionic strengths. We obtained values for Ksp(Im=O) and A& of (-8.79 f 0.12) and (0.12

f 0.051, respectively. Figure 3 shows our values of “log KJIJ - 20’’ as a function of ionic strength and the resulting linear regression of those data. Our value of (-8.79 k 0.12) for Ksp(Im=O) is in agreement with the result from Neck et al. of (-8.76 0.05) [ill within the uncertainties and our preliminary Pitzer analysis of Neck‘s data (-

8.72) [ E l . Grenthe et al. [21] list a value for E(Na+ - OH-) of (0.04 * 0.01) that when

rearranging gives log Ksp(Im) - 2 0 = log Ksp(Im = O ) - A d m

subtracted from our AE gives the resulting interaction coefficient, ~(Np0,’ - Cl-1, a

value of (0.08 -c 0.05).

ACKNOWLEDGMENTS

This work was performed as part of the Waste Isolation Pilot Plant (WIPP) Actinide Source Term Program at the Lawrence Berkeley National Laboratory for Sandia National Laboratories under Contract No. 40-2516 and AH-5592. Lawrence Berkeley National Laboratory is operated by the University of California for the U.S. Department of Energy under Contract DE-AC-03-76SF00098. Partial support for Herbert B. Silber came from the NIH-Minority Biomedical Research Support (MBRS) Grant and PRF Grant 26067-B3 a t San Jose State University. This work was performed under an NQA-1 equivalent Quality Assurance Program.

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References

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Novak, C.F.: An Evaluation of Radionuclide Batch Sorption Data on Culebra Dolomite for Aqueous Compositions Relevant to the Human Intrusion Scenario for the Waste Isolation Pilot Plant WIPP). SAND91-1299. Albuquerque, New Mexico: Sandia National Laboratories (1992).

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Nitsche, H.: “Solubility Studies of Transuranium Elements for Nuclear Waste Disposal: Principles and Overview.” Radiochimica Acta, 52/53,3-8 (1991).

Mathur, J.N., Choppin, G.R.: “Phosphate Complexing of NPO~~+”, Radiochimica Acta, 64,175-177 (1994).

Silber, H.B., Nitsche, H., Gatti, R.C., Gehmeker, H., Feige, G., Bucher, J., Edelstein, N.: “The Effects of Radiolysis Upon Speciation and Solubility of Neptunium in Brine Solutions”, Radiochimica Acta, 66/67, 15-18 (1994).

Khalili, F.1, V. Symeopoulos, J.-F. Ch-en, and G.R. Choppin: “Solubility of Nd in Brine”, Radiochimica Acta, 66/67: 5 1-54 (1994).

Moskvin, A. I.: “Hydrolytic Behavior of Neptunium(W, V, VI)”, Radiokhimiya, 13,681-687 (197 1). (English translation) Soviet Radiochem., 13,700-705 (1971).

Sevost’yanova, E.P., Khalturin, G.V.: “Hydrolytic Behavior of Neptunium(V)”, Radiokhimiya, 18,8704376 ( 1976). (English translation) Soviet Radiochem., 18,738-743 (1976).

Maya, L.: “Hydrolysis and Carbonate Complexation of DioxoneptuniumW) in 1.0 M HC104 at 25’C.” Inorganic Chemistry vol. 22 #14: 2093-2095 (1983)

Nitsche, H., Standifer, E.M., Silva, R.J.: “Neptunium Complexation with Carbonate”, Lanthanide Actinide Res., 3,203-211 (1990).

Neck, V., J.I. Kim, Kanellakopolos, B.: “Solubility and Hydrolysis Behaviour of Neptunium(V).” Radiochimica Acta, 56,2530 (1992).

Neck, V., R u d e , W., Kim, J.I., Kanellakopulos, B.: “Solid-Liquid Equilibrium Reaction of Neptunium(V) in Carbonate Solution at Different Ionic Strength”, Radiochimica Acta, 65,29-37 (1994).

13. Lierse, Ch., Treiber, W. and . Kim, J.I.: “Hydrolysis Reactions of Neptunium(V).” Radiochimica Acta, 38, 27-28 (1985).

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Moriyama, H., Pratopo, M.I., Higashi, K.: “Systematics of Hydrolysis and Carbonate Complexation Constants of Ions of Transuranium Elements”, Radiochimica Acta, 66/67,73-79 (1994).

Novak, C.F., and Roberts, ICE.: “Thermodynamic Modeling of Neptunium(V) Solubility in Na-C03-HC03-Cl-C104-H-OH-H20 Electrolytes.” Materials Research Society Symposium Proceedings. Volume 353. Scientific Basis for Nuclear Waste Management XVIII. T. Murakami and R.C. Ewing, editors. Pittsburgh, Pennsylvania: Materials Research Society. Vol. 353 Part 2, 1119-1128 (1995).

Rai, D., Felmy, A.R., Juracich, S.P., and Rao, L.: Estimating the Hydrogen Ion Concentration in Concentrated NaCl and Na2SO4 Electrolytes. SAND94 1949. Albuquerque, New Mexico: Sandia National Laboratories (1995).

Harned, H.S., Mannweiler, G.E.: “The Thermodynamics of Ionized Water in Sodium Chloride Solutions”, J. Am, Chem. SOC., 57,1873-76 (1935).

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DISCLAIMER

.

Table 1. The solubility products for NpO,OH(am) in NaCl solutions.

~ ~~ ~

NaCl Concentration,

rno1.L-l

11.7

log KSp&J * 2a* NaCl log Ksp(T,.J * 2a* (rnoP*L-') Concentration, (molal2>

molal

-8.61 * 0.28 0.30 -8.61 * 0.28 -8.56 0.28 0.60 -8.55 * 0.28 -8.5 1 0.14 1.0 -8.49 * 0.14 -8.56 * 0.11 1.8 -8.53 * 0.11

I I -8.73 f 0.09 3.0 -8.68 f: 0.09 I I

I I

-9.02 +. 0.24 5.6 -8.92 f 0.24 I I I I I

* Mean values over days 1 through 37.

9

-1 3.6

-1 3.8

-14.0 n ?L ca -

-14.2 W

3

0 0

k -14.4 -

-14.6

-1 4.8

-1 5.0

A 25"C[17]

0 2O"C[17]

15"C[17]

. This work, (21 f 2)" C

U

0 1 2 3 4 5 6 Ionic Strength (molal)

Figure 1. The ion product of water (molal2> in sodium chloride solutions. Open symbols are from Harned and Mannweiler [17]; filled symbols are from this work.

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1 X I 0-8

1 XI o - ~

1 X I 0-’0 0

L

1

++ 0.3 molal NaCl + 1.8 molal NaCl

-€+ 0.6 molal NaCl -i- 3.0 molal NaCl

-6- 1.0 molal NaCl + 5.6 molal NaCl I I I I I I I

5 10 15 20 25 Time (Days)

30 35 40

Figure 2. The solubility products (mo12.L-2) of NpO,OH(am) solid in solutions of sodium chloride as a function of time.

11

-8.50

-8.70

-8.90 Q cy ' -9.10 2. = k' -9.30

-9.70

-9.90 0 1 2 3 4

Ionic Strength (molal) 5 6

Figure 3. Plot of log Ksp(IJ - 20 as a function of ionic strength for the determination of the solubility product of NpO,OH(am) at infinite dilution, log Ksp(Im=O).

12