ISSN 1066-3622, Radiochemistry, 2013, Vol. 55, No. 2, pp. 162–167. © Pleiades Publishing, Inc., 2013.

Original Russian Text © V.P. Shilov, A.M. Fedoseev, 2013, published in Radiokhimiya, 2013, Vol. 55, No. 2, pp. 120–124.

162

Oxidation of Np(V) and Np(IV) with Fe(CN)63– Ions

in Carbonate Solutions

V. P. Shilov and A. M. Fedoseev

Frumkin Institute of Physical Chemistry and Electrochemistry, Russian Academy of Sciences, Leninskii pr. 31, block 4, Moscow, 119071 Russia; * e-mail: ShilovV@ipc.rssi.ru

Received June 13, 2012

Abstract—The formal potential of the Fe(CN)63–/Fe(CN)6

4– couple in 1 M NaHCO3 and 1–2 M Na2CO3 solu-tions was determined. It is equal to 505 and 510 mV, respectively, exceeding the potentials of the Np(VI)/(V) and Np(V)/(IV) couples in carbonate solutions. The equilibrium of the reaction Np(V) + Fe(CN)6

3– = Np(VI) + Fe(CN)6

4– was studied. Fe(CN)63– ions oxidize Np(IV) to Np(V) and then to Np(VI). The arising Np(VI) oxidizes

Np(IV). The Np(IV) oxidation accelerates in going from NaHCO3 to Na2CO3. An increase in [Na2CO3] or in the ionic strength (by adding neutral salts) decelerates the oxidation. Np(IV) introduced in an HCl solution re-acts with Fe(CN)6

3– or with Np(VI) faster than Np(IV) introduced in a Na2CO3 solution. The activation energy of the reaction of Np(IV) with Fe(CN)6

4– in the range 20–45°С is 107 kJ mol–1. The reaction mechanism in-volves formation of the activated complex from ions of Np(IV) hydroxocarbonate and oxidant.

Keywords: neptunium, carbonate solution, hexacyanoferrates, redox reactions, kinetics

The formal potential E of the Np(VI)/(V) couple is 440 mV in Na2CO3 [1] or K2СО3 [2] solution, and that of the Np(V)/(IV) couple, 100 mV in 1 M K2СО3 solution [3]. The standard potential of the Fe(CN)6

3–/Fe(CN)64– couple is 355 mV [4]. The poten-

tial of this couple increases to 457 mV in 0.5 M KBr (estimated from Fig. 2 in [4]). Cations can form ion pairs with Fe(CN)6

4–, decreasing the concentration (activity) of the free ferrocyanide ion and increasing the potential of the couple.

In <0.05 M Na2CO3 solutions, the potential of the Np(VI)/(V) couple exceeds that of the Fe(CN)6

3–/ Fe(CN)6

4– couple. Therefore, Sullivan et al. [5] studied in this medium the kinetics of the reaction

If the concentration of Na2CO3 or neutral salts is in-creased to 1 M and more, the potential of the Fe(CN)6

3–/Fe(CN)6

4– couple increases, and reaction (1) will mainly proceed from the right to the left.

In this study we measured the formal potential of the Fe(CN)6

3–/Fe(CN)64– couple in 1 M NaHCO3 and

2 M Na2CO3 solutions and examined the oxidation of Np(V) and Np(IV) with ferricyanide ions in such solu-tions.

DOI: 10.1134/S1066362213020057

EXPERIMENTAL

Np(VI) + Fe(CN)64– = Np(V) + Fe(CN)6

3–. (1)

We used 237Np purified by anion exchange. To a Np(IV) solution eluted from the ion-exchange column we added HClO4 and evaporated until thick white fume appeared. The dry residue was dissolved in a mixture of HNO3 and HClO4, and the evaporation was repeated. This operation resulted in conversion of Np(IV) to Np(VI). Several crystals of NpO2(ClO4)2· nH2O were dissolved in 0.01 M HClO4. The remaining amount of dry Np(VI) perchlorate was dissolved in water, NaNO2 was added, and the resulting Np(V) was precipitated with ammonia. Np(V) hydroxide was separated, washed with ice-cold double-distilled water, and dissolved in the minimal amount of HClO4. A part of the Np(V) solution was mixed with 4 M НСl, NH2OH·HCl was added, and the mixture was heated for 0.5 h to convert Np(V) to Np(IV). A saturated CsCl solution was added for precipitating Cs2NpCl6. The precipitate was separated on a glass frit, washed with concentrated HCl, and dried in air. Prior to oxidation experiments, a weighed portion of Cs2NpCl6 was dis-solved in 2 M HCl or 2 M Na2CO3. An aliquot of the stock solution of neptunium was mixed with 2.6 ml of 1.1 M HCl and placed in a quartz cell (l = 1 cm). The valence forms of Np were monitored spectrophotomet-

OXIDATION OF Np(V) AND Np(IV) WITH Fe(CN)63– IONS 163

RADIOCHEMISTRY Vol. 55 No. 2 2013

RESULTS AND DISCUSSION rically. The Np concentration in stock solutions was determined by titration with a standard Na2EDTA so-lution using Xylenol Orange indicator, with the pre-liminary reduction of neptunium to Np(IV) [6]. We used analytically pure grade NaHCO3 and recrystal-lized chemically pure grade Na2CO3. The Na2CO3 con-centration in the stock solution was determined by acid titration. The NaHCO3 solution was prepared from an analytically weighed portion and used within the preparation day. K3Fe(CN)6 and K4Fe(CN)6·3H2O were of chemically pure grade. All the solutions were prepared in double-distilled water.

The redox potential of the Fe(CN)63–/Fe(CN)6

4– cou-ple was found by measuring emf (Еmeas) of the circuit

Pt|Fe(CN)63–/Fe(CN)6

4–, NaHCO3 (Na2CO3)||KCl saturated,

AgCl|Ag, Pt.

The formal potential, i.e., the potential of the cou-ple at [Fe(CN)6

3–] = [Fe(CN)64–], was calculated by the

formula Еf = Еmeas + E(Ag/AgCl) – Ed, where E(Ag/AgCl) is the potential of the reference electrode in saturated KCl solution, equal to 198.9 mV at 25°С [7], and Еd is the diffusion potential.

Electrochemical measurements were performed with an ОР-211/1 digital pH-meter (Radelkis, Hun-gary) operating in the millivoltmeter mode with an accuracy of ±1 mV.

Experiments were performed as follows. A cell was charged with 2.94 ml of the carbonate solution, the working (1 × 1 cm Pt plate) and reference electrodes were submerged, and Ar was bubbled. Then, 0.030-ml portions of 0.1 M K3Fe(CN)6 and 0.1 M K4Fe(CN)6 were added, the pH-meter and stopwatch were switched on, and the pH-meter readings were taken at regular intervals. When studying reaction (1) going in one or another direction and the Np(IV) oxidation, the cell (l = 1, 2, or 5 cm) was charged with a carbonate solution and a Np solution, and the absorption spec-trum was recorded in the range 350–800 nm with a Shimadzu PC 3100 (Japan) or SF-46 (Leningrad Opti-cal and Mechanical Association, Russia) spectropho-tometer. An aliquot of a K3Fe(CN)6 or K4Fe(CN)6 so-lution was added, and changes in the absorption spec-trum were monitored. In some cases, the Np solution was added last. Because the Np(V) solubility at low СО3

2– concentrations or at pH < 9 is low [8], we used Na2CO3 or NaHCO3 + Na2CO3 mixture.

Below are the values of Еmeas and formal potentials Еf of the Fe(CN)6

3–/Fe(CN)64– couple (plus Еd). The

quantity Ed can be estimated by the Henderson equa-tion, but the reference values of the ion mobilities used in this equation are given for infinite dilution and are hardly applicable to concentrated solutions and to esti-mation of the diffusion potential.

The potentials that we obtained for the Fe(CN)63–/

Fe(CN)64– couple exceed the potential of the

Np(VI)/(V) couple starting from the NaHCO3 concen-tration of 0.1 M. Therefore, small excess of ferricya-nide in 0.2–1 M NaHCO3 or Na2CO3 solutions should convert Np(V) to Np(VI). It follows from comparison of the potentials of these couples that, in 1 M NaHCO3 or Na2CO3, the equilibrium constant of reaction (1) going from the right to the left is

[NaHCO3], M 0.1 0.18 0.29 0.44 0.53 0.74 1.0

[Na2CO3], M 1.0 2.0

Emeas, mV 261 270 279 288 294 299 306 309 312

Ef + Ed, mV 460 469 478 487 493 498 505 508 511

K = [Np(VI)][Fe(CN)64–][Np(V)]–1[Fe(CN)6

3–]–1 = 10. (2)

Indeed, Np(VI) is formed on mixing the carbonate solution of Np(V) with a K3Fe(CN)6 solution. By the moment of recording the spectrum, i.e., within 25– 30 s, the reaction is complete. Hence, the bimolecular rate constant k2 is no less than ~103 l mol–1 s–1. The equilibrium concentrations and values of K are given in Table 1. The Np(VI) concentration was estimated using the molar extinction coefficient of Np(VI), ε350 = 1700 l mol–1 cm–1.

The experimental values of K increase in going from 1 M NaHCO3 to 0.1–1.96 M Na2CO3 and remain below the expected level. Probably, in the Np(V) car-bonate complex, the СО3

2– group is replaced by the Fe(CN)6

4– ion, which leads to a decrease in the concen-tration of the Np(V) carbonate complex, to an increase in E[Np(VI)/(V)], and to a decrease in K.

In a solution containing 0.062 M Na2CO3, 4.75 M NaClO4, 1 mM Np(V), and 0.7–1.9 mM K3Fe(CN)6, the equilibrium constant is 1.7 ± 0.1, i.e., in a solution with a high concentration of cation but a low concen-tration of СО3

2–, the equilibrium constant is lower be-cause of formation of a mixed Np(V) complex.

The mechanism of Np(V) oxidation includes the

SHILOV, FEDOSEEV 164

RADIOCHEMISTRY Vol. 55 No. 2 2013

Table 1. Influence of conditions on the equilibrium of the reaction Np(V) + Fe(CN)63– = Np(VI) + Fe(CN)6

4– at 23°Cа

[NaHCO3] [Na2CO3] [Fe(CN)63–]0 [Np(V)]0 [Np(VI)]eq K

M mM 0.46 0.02 1.36 0.51 0.27 0.28 0.93 0.04 1.41 0.54 0.42 1.48

0.1 1.10 1.04 0.66 2.60 0.5 0.97 1.04 0.68 4.43 1.0 1.12 1.04 0.71 3.72 1.96 1.13 1.04 0.73 4.30

а Subscripts 0 and eq refer to the initial and equilibrium concentrations; the same in Table 2.

step in which the СО32– group of the Np(V) carbonate

complex is replaced by the Fe(CN)63– ion:

Table 2. Influence of conditions on the equilibrium of the reaction Np(VI) + Fe(CN)64– = Np(V) + Fe(CN)6

3– at 23°C

[NaHCO3] [Na2CO3] [Fe(CN)64–]0 [Np(VI)]0 [Np(VI)]eq K

M mM 0.95 0.04 0.77 1.02 0.58 0.99

0.95 0.04 1.54 1.02 0.41 1.02

0.95 0.04 2.29 1.01 0.32 1.08

0.91 0.04 7.45 0.99 0.2 2.1

1.96 0.77 1.02 0.68 2.53

1.96 1.54 1.02 0.49 1.76

1.96 3.05 1.01 0.34 1.80

1.96 4.54 1.00 0.27 1.93

1.96 6.00 0.99 0.25 2.40

[NaClO4]

4.7 0.07 0.77 1.02 0.56 0.82

4.7 0.07 1.54 1.02 0.37 0.78

4.7 0.07 3.05 1.01 0.25 0.99

(3) NpO2

V(CO3)23– + Fe(CN)6

3– = NpO2VCO3Fe(CN)6

4– + CO32–,

NpO2

VCO3Fe(CN)64– → Np(VI) + Fe(CN)6

4–. (4)

We performed several experiments on the Np(VI) reduction with Fe(CN)6

4– ions in concentrated solu-tions. However, we calculated the equilibrium constant of this reaction by Eq. (2). Addition of K4Fe(CN)6 to the carbonate solution of Np(VI) with a high concen-tration of sodium ions leads to a decrease in the Np(VI) concentration. The reaction is complete in 3– 5 min. The rate constant k1 is close to 102 l mol–1 s–1. The initial and equilibrium concentrations of Np(VI) and the equilibrium constants are given in Table 2.

The equilibrium constants in Tables 1 and 2 are close and are lower than the expected values, which is attributed to the formation of a mixed Np(V) complex.

The Np(VI) reduction passes through a step of com-plexation of Np(VI) with Fe(CN)6

4– [5].

Beitz et al. [9] studied the reaction

(5) Pu(VI) + Fe(CN)64– = Pu(V) + Fe(CN)6

3–

in 0.05 M Na2CO3. At 25°С, the rate constants of for-ward (kf) and reverse (kr) reactions are 1214 and 3.64 × 104 l mol–1 s–1, respectively, and K5 = kf /kr = 0.0334. As can be seen, the Pu(VI) reduction with Fe(CN)6

4– is far from being complete even with excess Fe(CN)6

4–. In going to more concentrated Na2CO3 solutions, the Pu(VI) reduction becomes complicated. The formal potential of the Pu(VI)/(V) couple in 0.1 and 1.0 M K2СО3 is 300 and 320 mV, respectively [2]; hence, K5 = 10–3.

In carbonate solutions, excess ferricyanide oxidizes Np(IV) at moderate rate to Np(V) and then rapidly to Np(VI). The figure shows the kinetic curves of Np(IV)

OXIDATION OF Np(V) AND Np(IV) WITH Fe(CN)63– IONS 165

RADIOCHEMISTRY Vol. 55 No. 2 2013

oxidation. They are not linearized in semilog coordi-nates; the process deviates from the first order with its progress, i.e., it is autocatalytic. This is due to the oc-currence of reactions (6), (1) (from the right to the left), and (7):

(6) Np(IV) + Fe(CN)63– = Np(V) + Fe(CN)6

4–,

Np(IV) + Np(VI) = 2Np(V). (7)

Reaction (7) was studied previously [10] in 1.71–4.0 M K2СО3 solutions at 52–76.5°С. Np(IV) was in-troduced as an aliquot of a solution in 1 M HClO4. It was found that the rate constant of reaction (7) de-creased by a factor of 4 with an increase in [K2CO3] from 1.71 to 4 M, but at constant K2СО3 concentration addition of 0.21 M SO4

2– also decreased the rate con-stant by a factor of 4, which seems unexpected.

To determine the contribution of reaction (7) in NaHCO3 and Na2CO3 solutions at 20°С, we performed separate experiments. In NaHCO3 or Na2CO3 solutions at [Np(IV)] = [Np(VI)] = 1 mM, the kinetic curves are linearized in the coordinates 1/D350–t (time), i.e., the reaction follows the second-order rate law. In a NaHCO3 by the end of the experiment, Np(V) carbon-ate precipitates. From the relationship

Kinetic curves of Np(IV) oxidation with ferricyanide ions in 1 M Na2CO3. l = 1 cm, 20°С, [Fe(CN)6

3–] = 40 mM; [Np(IV)], mM: (1) 1.33 and (2) 2.66.

k = εlΔ(1/D)/Δt,

where D = Dt – D∞, Dt and D∞ are the values of the optical density at the given time moment and by the end of the experiment, respectively, and l is the cell length, using the ε350 value for Np(VI), we calculated the rate constants. The results are given in Table 3.

The rate constant k7 increases in going from NaHCO3 to Na2CO3 in the case of using stock solution of Np(IV) both in HCl and in Na2CO3. An increase in [Na2CO3] by a factor of 2 leads to a decrease in k7 by a factor of 4. When the same ionic strength is made with NaCl, k7 decreases by a factor of 3.6. Comparison of the k7 values in Na2CO3 + Na2SO4 and Na2CO3 + NaCl so-lutions shows that the SO4

2– ions decelerate the reaction owing exclusively to an increase in the ionic strength.

A decrease in the Np(IV) and Np(VI) concentrations to 0.15 mM leads to an increase in the rate constant to 0.261 and 0.615 l mol–1 s–1 in 0.994 M NaHCO3 + 0.008 M Na2CO3 and in 1 M Na2CO3, respectively [20°С, stock solution of Np(IV) in 2 M Na2CO3].

The specific features of reaction (7) are primarily associated with the Np(IV) speciation. In the aqueous system K+–НСО3

––СО32––ОН– at low and high concen-

trations of carbonate and high concentration of bicar-bonate, Np(CO3)5

6– prevails [11]. In solutions with low bicarbonate concentration, Np(OH)2(CO3)2

2– is formed. However, according to [12], in NaHCO3–Na2CO3–NaClO4 solutions at ionic strengths of 0.5, 1.0, and 2.0 M, Np(OH)2(CO3)2

2– prevails in the pH range 8.5–10.5, and Np(OH)4(CO3)2

4– prevails at pH > 12. In per-chloric acid solutions, the Np(IV) + Np(VI) reaction occurs by the mechanism of the hydrogen atom trans-fer from the hydrate shell of the Np(IV) ion to the Np(VI) ion [13]. If such a mechanism takes place in carbonate solutions, the Np(IV) ion should contain a labile Н atom. Therefore, the activated complex is formed with the participation of a hydroxycarbonate ion, e.g., Np(OH)2(CO3)2

2–. The OH groups are sources of “yl” oxygen atoms for NpO2

+. On dilution of the stock solution of Np(IV) in 2 M Na2CO3, the yield of this complex in 1 M Na2CO3 is higher than in 1 M NaHCO3, which is reflected in k7. High yield of hydro-lyzed ions is observed in the case of dilution of acid stock solution of Np(IV). An increase in [Na2CO3] over 1 M decreases the yield of the hydrolyzed ion and

Table 3. Influence of conditions on the kinetics of the reac-tion Np(IV) + Np(VI) = 2Np(V) at 20°C

[NaHCO3] [Na2CO3] [NaCl] Stock solution

of Np(IV) I, M

k7, l mol–1 s–1 M

1.0 HCl 1 0.44

0.03 0.97 HCl 3 0.77

0.98 0.03 Na2CO3 1 0.112

1.0 Na2CO3 3 0.300

0.485 Na2SO4 1.0 Na2CO3 4.45 0.126

1.0 1.46 Na2CO3 4.46 0.124

1.973 Na2CO3 5.92 0.067

1.0 2.98 Na2CO3 5.98 0.088

3 +

log

SHILOV, FEDOSEEV 166

RADIOCHEMISTRY Vol. 55 No. 2 2013

impedes the process. An increase in the ionic strength leads to a similar result. The hydrolyzed ions are capa-ble of forming dimers in the initial steps of polymeri-zation. In a solution with [Np(IV)] = [Np(VI)] = 0.15 mM, the dimerization is slow, and k' is larger than in the solution in which [Np(IV)] = [Np(VI)] = 1 mM.

Reaction (6) was studied in solutions containing 15–30-fold excess of Fe(CN)6

3– to decrease the effect of reaction (7). The reaction progress was monitored by variation of D480 or D705, corresponding to the absorp-tion of Np(VI) or Np(IV), respectively. From the initial portion of the kinetic curves in the logD–time coordi-nates, we found the pseudo-first-order rate constant k' by the relationship

Table 4. Influence of conditions on the rate constant of reaction (6). [Np(IV)] = 1.33, [Fe(CN)63–] = 40 mM

T, оC [NaHCO3] [Na2CO3] Stock solution of Np(IV) I, M k6 × 103, l mol–1 s–1

M 20 0.92 HCl 1.16 4.57 20 0.04 0.96 HCl 3.16 2.3–4.9 20 0.90 0.04 Na2CO3 1.26 2.3 20 0.5 0.5 Na2CO3 2.24 3.9 20 1 Na2CO3 3.24 4.7 20 1.84 Na2CO3 5.76 1.20 20 2.52 NaCl 1 Na2CO3 5.76 1.40 34 1 Na2CO3 3.24 46.7 45 1 Na2CO3 3.24 147

2.3logD = k't + const,

where D = D∞ – Dt, D∞ is the maximal optical density at 480 nm, or D = Dt – D∞, D∞ is the minimal optical density at 705 nm.

Below we give the values of k' and the calculated values of k6 = k'/[Fe(CN)6

3–] in 1 M Na2CO3 + 1.33 mM Np(IV) solution at 20°С.

With an increase in the ferricyanide concentration, k' increases almost proportionally, i.e., the reaction follows the first-order rate law with respect to the oxidant. A decrease in the bimolecular rate constant is associated with an increase in the ionic strength, as seen from Table 4.

The use of a stock solution of Np(IV) in 2 M HCl leads to the scatter of the results. In a 0.04 M NaHCO3 +

[Fe(CN)63–], mM 20 40 80 160

I, M 3.12 3.24 3.48 3.96

k' × 105, s–1 8.68 18.72 26.9 49.3

k6 × 103, l mol–1 s–1 4.34 4.68 3.75 3.08

0.96 M Na2CO3 solution, the extreme values differ by a factor of 2. With the carbonate stock solution of Np(IV), the scatter of the results obtained under the same conditions did not exceed 20%. The majority of data given in Table 4 were obtained with carbonate stock solution of Np(IV). In going from 1 M NaHCO3 + Na2CO3 to 1 M Na2CO3, the pH value, [CO3

2–], ionic strength, and rate constant increase. An increase in [Na2CO3] over 1 M or addition of NaCl leads to an increase in the ionic strength without significant change in pH and to a decrease in the rate constant. The activation energy is close to 107 kJ mol–1, which is comparable with the activation energy of reac- tion (7), equal to 103 kJ mol–1 [10].

In a 1 M Na2CO3 solution containing 2.66 mM Np(IV) introduced in the carbonate stock solution, k6 is 5.3 × 10–3 and 4 × 10–3 l mol–1 s–1 in the presence of 40 and 80 mM Fe(CN)6

3–, respectively (20°C). In the first case, reaction (7) plays noticeable role, whereas in the second case higher ionic strength inhibits the process.

The course of reaction (6) depends on the content of hydrolyzed Np(IV) ions. Their yield decreases with an increase in [CO3

2–] from 1 to 1.84 M, because k6 decreases by a factor of almost 4; therefore, k6 is proportional to [CO3

2–]–n. Essentially similar effect is exerted by an increase in the ionic strength. In going from NaHCO3 + Na2CO3 to 1 M Na2CO3, pH increases from <9 to 12, i.e., [OH–] increases. Hence follows that k6 depends on [OH–]m, and an increase in [CO3

2–] does not prevent the hydrolysis. Thus, k6 = k0[OH–]m[CO3

2–]–n, with m > n. The activated complex is formed from the Np(IV) hydroxocarbonate ion and Fe(CN)6

3–. The charge transfer in the activated complex gives rise to Np(V).

To reduce the effect of reaction (7), we performed

REFERENCES

OXIDATION OF Np(V) AND Np(IV) WITH Fe(CN)63– IONS 167

RADIOCHEMISTRY Vol. 55 No. 2 2013

experiments with solutions containing 0.15 mM Np(IV) and 40 mM Fe(CN)6

3– in cells with the optical path length of 5 cm. We used a stock solution of Np(IV) in 2 M Na2CO3, and it was added last. At 20°С in 0.916 M NaHCO3 + 0.007 M Na2CO3 solution, k6 is 18 × 10–3 in the period from 0 to 5 min and 2.7 × 10–3 l mol–1 s–1 in the period from 5 to 120 min. In 1 M Na2CO3, k6 is 57 × 10–3 in the period from 0 to 3 min, 9.6 × 10–3 in the period from 3 to 10 min, and 5.4 × 10–3 l mol–1 s–1 in the later period. The course of the process is apparently associated with dimerization of Np(IV) hydroxycarbonate complexes.

It should be noted in conclusion that k6 is lower than k1 and k2 by 5–6 orders of magnitude. Reac- tion (6) is accompanied by a change in the structure of the central ion, whereas reaction (1) going in any direction does not involve changes in the structure of the participating ions.

1. Wester, D.W. and Sullivan, J.C., J. Inorg. Nucl. Chem., 1981, vol. 43, no. 11, pp. 2919–2923.

2. Simakin, G.A., Volkov, Yu.F., Visyashcheva, G.I., et al., Radiokhimiya, 1974, vol. 16, no. 6, pp. 859–863.

3. Fedoseev, A.M., Peretrukhin, V.F., and Krot, N.N., Dokl. Akad. Nauk SSSR, 1979, vol. 244, no. 5, pp. 1187–1190.

4. Hanania, G.I.H., Irvine, D.H., Eaton, W.A., and George, P., J. Phys. Chem., 1967, vol. 71, no. 7, pp. 2022–2030.

5. Sullivan, J.C., Brown, W., Nash, K., and Thomp- son, M., Radiochim. Acta, 1982, vol. 30, no. 1, pp. 45–47.

6. Smirnov-Averin, A.P., Kovalenko, G.S., Ermola- ev, N.P., and Krot, N.N., Zh. Anal. Khim., 1966, vol. 21, no. 1, pp. 76–78.

7. Midgley, D. and Torrance, K., Potentiometric Water Analysis, Chichester: Wiley, 1978.

8. Neck, V., Runde, W., Kim, J.I., and Kanellakopulos, B., Radiochim. Acta, 1994, vol. 65, no. 1, pp. 29–37.

9. Beitz, J., Jonah, C., Sullivan, J.C., and Woods, M., Ra-diochim. Acta, 1986, vol. 40, no. 1, pp. 7–9.

10. Frolova, L.M., Chistyakov, V.M., Ermakov, V.A., and Rykov, A.G., Radiokhimiya, 1976, vol. 18, no. 2, pp. 281–283.

11. Rao, D., Hess, N.J., Felmy, A.R., and Moore, D.A., Radiochim. Acta, 1999, vol. 84, no. 3, pp. 159–169.

12. Kitamura, A. and Kohara, Y., J. Nucl. Sci. Technol. (Jpn.), 2002, suppl. 3, pp. 294–297.

13. Hindman, J.C., Sullivan, J.C., and Cohen, D., J. Am. Chem. Soc., 1959, vol. 81, no. 10, pp. 2316–2319.