Electrochemical studies on the reduction of uranyl ions in nitric acid-hydrazine media

Journal of Electroanalytical Chemistry 776 (2016) 127–133

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Electrochemical studies on the reduction of uranyl ions in nitric acid-hydrazine media Satyabrata Mishra a, Sini K. a, Ch. Jagadeeswara Rao b, C. Mallika a,⁎, U. Kamachi Mudali a a b

Reprocessing Group, Indira Gandhi Centre for Atomic Research, Kalpakkam 603 102, India Corrosion Science and Technology Group, Indira Gandhi Centre for Atomic Research, Kalpakkam 603 102, India

a r t i c l e

i n f o

Article history: Received 1 March 2016 Received in revised form 25 June 2016 Accepted 1 July 2016 Available online 04 July 2016 Keywords: Uranyl reduction Holding reductant Diffusion coefficient Cyclic voltammetry Chronopotentiometry Current efficiency

a b s t r a c t The redox behaviour of uranyl ions (UO2+ 2 ) in nitric acid and nitric acid-hydrazine media were investigated by the transient electrochemical techniques cyclic voltammetry (CV) and chronopotentiometry (CP), at the working electrodes platinum, gold, titanium and glassy carbon (GC) at 298 K. At a very low concentration of 0.05 M nitric acid and in the absence of hydrazine, the reduction of uranyl ions was found to be under charge transfer and diffusion control. Nevertheless, as the acidity of the supporting electrolyte increased, the reduction was purely under kinetic control. The diffusion coefficient (Do) values for the reduction of U (VI) at the working electrodes platinum, gold GC as well as Ti were determined. The current efficiency and conversion efficiency for the reduction of uranyl ions were found to be better under constant potential conditions than that at constant current conditions. © 2016 Elsevier B.V. All rights reserved.

1. Introduction Partitioning of uranium from plutonium is a major step in the aqueous reprocessing of spent nuclear fuels using PUREX process [1]. Partitioning in the organic phase can be achieved by adopting either non-reductive or reductive stripping method. In the former method, separation is accomplished by forming aqueous favouring complexes of Pu (IV) with sulphate, oxalate, or by saturating the organic solvent with uranium [2]. However, this method is not a favoured process for commercial application. In the latter method, separation is achieved by reducing Pu(IV) to aqueous favouring Pu(III) using suitable reducing agents and process conditions. The reducing agents which have been investigated for the separation of plutonium from uranium include hydrazine [3,4], hydroxylamine [5,6], ferrous ion (introduced as sulphamate, acetate, nitrate, etc.) [7–9], uranous ion (introduced as nitrate) [10–15], ascorbic acid [16] and recently hydrogen in the presence of a catalyst [17,18]. Among all the reductants available for reducing Pu(IV) to Pu(III), uranous nitrate and ferrous sulphamate are widely employed. Ferrous sulphamate reduces Pu (IV) with good efficacy. However, it introduces non-volatile and corrosive ions into the process streams and complicates nuclear waste management operation, in addition to forming a strong complex with plutonium. Further, it is necessary to maintain high ferrous to the ferric ratio (20 to 40 ⁎ Corresponding author. E-mail addresses: [email protected] (S. Mishra), [email protected] (S. K.), [email protected] (C.J. Rao), [email protected] (C. Mallika), [email protected] (U.K. Mudali).

http://dx.doi.org/10.1016/j.jelechem.2016.07.002 1572-6657/© 2016 Elsevier B.V. All rights reserved.

fold) to ensure complete reduction of Pu(IV). Hence, considerable attention has been paid to the use of uranous nitrate as the reductant for Pu(IV). The data available in the literature for the preparation and reducing action of uranous nitrate in stripping Pu are exhaustive [11,12]. Uranous nitrate can be produced from uranyl nitrate by chemical [4,18], photochemical [14,15] and electrochemical methods [19–22]. The former two approaches are not employed for the commercial production. Electro-reduction method is relatively simple and suitable for large scale production. Thus, production of uranous nitrate with excellent conversion efficiency is an important process step in nuclear fuel reprocessing. Conversion of U(VI) to U(IV) with high current efficiency can be achieved in continuous as well as in batch mode electro-reduction process if it is voltage-controlled [22]. Reduction of U(VI) at Ti cathode in nitric acid-hydrazine media was studied by Kim et al. [23]. These authors reported that a high activation potential was required for the reduction of uranyl ions at Ti cathode. The electrochemical behaviour of U (VI) in nitric acid solutions of concentration 0.1 to 3 M on platinum and titanium electrodes was investigated [10,24]. Polarographic studies in strong acid solutions, conducted by Kolthaff and Harris [25] revealed that at a higher acid concentration, the primary reduction product was UO+ 2 , which undergoes disproportionation in the presence of H+ to yield U (VI), U(IV) and H2O. Wei et al. [20] studied the effect of acidity on the reduction of U(VI) by cyclic voltammetric technique and had reported that high acidity is electrochemically favourable for the reduction of U(VI) to U(IV). Their observation is consistent with the results reported by Ghandour et al. [10], based on polarographic studies. Though several researchers had studied the electrochemical behavior of U(VI) at different electrodes

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in various media including room temperature ionic liquids [26,27] as well as the stability of U(V) in basic carbonate and in aqua complexes [28–30] by cyclic voltammetric technique, kinetic studies on the reduction of U(VI) in nitric acid medium is not reported reliably. Nevertheless, chronopotentiometric oxidation and reduction of Pu in mineral acid solutions using Pt wire micro-electrodes were investigated by Peters and Shults [31] and cyclic voltammetric studies of Pu and Np in nitric acid media were reported by Cassadio et al. [32]. To understand the mechanism of reduction of uranyl ions and to determine the kinetic parameters for the reduction reaction, the redox behavior of uranyl ions in nitric acid medium at different working electrodes was investigated in the present study using the transient electrochemical techniques, cyclic voltammetry (CV) and chronopotentiometry (CP). The reduction behavior of U (VI) in the presence of hydrazine was also investigated to compare the results with those data generated in the nitric acid medium. The value of heterogeneous electron transfer rate constant (ks) estimated for uranyl reduction at a low concentration of nitric acid (0.05 M) using Klingler and Kochi equation [33] based on peak separation method was validated using Nicholson equation [34]. The current, as well as conversion efficiencies in the reduction of uranyl to uranous ions at Ti working electrode, were compared on constant potential/current conditions. 2. Experimental 2.1. Chemicals The chemicals and reagents used were nuclear grade U3O8 (supplied by NFC, Hyderabad), AR grade nitric acid (Hi-Pure Chemicals, Chennai), potassium hydrogen phthalate (GR, 99.9%, Merck, Mumbai), sodium hydroxide and sulphuric acid (GR, 98%, Merck, Mumbai), hydrazine hydrate (AR, 99%, S d fine Chem Ltd., Mumbai), potassium dichromate (GR, 99.5%, Merck, Mumbai), potassium oxalate (AR, 99.5%, Ranbaxy Fine Chemicals Ltd., New Delhi) and Ferroin indicator (Merck, Mumbai). Solutions of uranyl nitrate in different concentrations of nitric acid (0.05, 1.0 and 1.5 M) and hydrazine (0.25 M) were prepared by dissolving accurately weighed quantities of U3O8 in hot nitric acid. 0.25 M hydrazine nitrate was added as the holding reductant to scavenge nitrous acid (produced as an intermediate in the reduction of nitric acid), which would otherwise catalyse the oxidation of U(IV) ions. 2.2. Estimation of uranyl [U(VI)] and uranous [U(IV)] ions Uranyl ions in the stock solutions were quantitatively analyzed by Davis and Gray method [35] after reducing U(VI) to U(IV) in a strong phosphoric acid medium using Fe(II) and titrating the resultant U(IV) ions against standard potassium dichromate solution potentiometrically, using Metrohm auto titroprocessor-670. The amount of U(IV) generated during electro-reduction was estimated periodically in the sulphuric acid medium by titrating against standard potassium dichromate in the presence of ferroin indicator. The end point was the disappearance of reddish orange and the appearance of blue colour. Uranous ions were also characterized by UV‐Visible spectrophotometry.

hydrazinium nitrate occurs at pH 5.6 (for the concentrations of nitric acid and hydrazine encountered in the test solutions). Thus, for the estimation of the free acid, 0.1 ml of the aliquot (test solution) was added to 30 ml of 0.1 M (pH adjusted) potassium oxalate solution. Owing to the presence of free acid in the sample, the pH of potassium oxalate solution decreased, and it was brought back to the initial pH of 5.6 by titrating against standardized NaOH solution and using a combined glass electrode in an auto titroprocessor. From the titre value, the concentration of free acid in the sample was calculated. To determine the concentration of hydrazine, 3 ml of 30% formaldehyde was added to the same solution, the pH of which was 5.6. Formaldehyde releases HNO3 from N2H4.HNO3 and binds hydrazine, the quantity of which is equivalent to the amount of acid liberated from N2H4.HNO3. As free acid was released from hydrazine nitrate, the pH of the solution was found to decrease below pH 5.6, which was again adjusted back to 5.6 with standardized sodium hydroxide. The concentration of acid determined from the titre value is equal to the concentration of hydrazine. 2.4. Electrochemical setup Electrochemical behaviour in the reduction of U(VI) was investigated by performing CV and CP experiments at 298 K in a standard threeelectrode cell comprising Pt foil (Surface area, SA: 1.2 cm2) as a counter electrode and Ag-AgCl as a reference electrode. Platinum (SA: 0.416 cm2), glassy carbon (SA: 0.448 cm2), Titanium (SA: 0.328 cm2) and Gold (SA: 0.228 cm2) were employed as working electrodes for the electrochemical studies using Autolab Model PGSTAT 30 (EcoChemie, the Netherlands) electrochemical system, equipped with General Purpose Electrochemical Software. 25 ml of 50 mM uranyl nitrate in different concentrations of nitric acid (0.05, 1.0 and 1.5 M), without or with hydrazine (0.25 M) served as the electrolyte. The voltammograms were recorded at different scan rates from 0.01 to 0.1 Vs-1.. Similarly, chronopotentiograms were recorded at various applied constant currents. All the experiments were performed after de-oxygenation of the test solutions by purging with argon gas. 3. Results and discussion 3.1. Cyclic voltammograms of U(VI) in HNO3 medium The cyclic voltammograms of 50 mM U(VI) in 0.05 M nitric acid, recorded at a platinum working electrode over the potential range 0.3 to ‐0.65 V against Ag/AgCl reference electrode at the scan rates 0.01 and 0.1 V s−1 are shown in Fig. 1. The CV recorded for the supporting

2.3. Simultaneous determination of free acidity and hydrazine As uranium (heavy metal ion) present in the electrolyte solution interferes in the estimation of free acidity by forming hydroxide with the titrant (NaOH) at very low concentrations of acid, it was essential to mask the uranyl ions by complexing with pre-neutralized potassium oxalate, before the determination of free acidity. Further, to eliminate the interference of hydrazine present in the solution, the pH of the potassium oxalate solution was adjusted to 5.6, because during free acidity determination, hydrazine (the weak base) would produce the conjugate acid, whose pKa value is 7.96. As per Henderson‐Hasselbalch equation, the equivalence point for the neutralisation reaction to form

Fig. 1. Cyclic voltammograms of U (VI) in 0.05 M HNO3 recorded with Pt electrode at 298 K.

S. Mishra et al. / Journal of Electroanalytical Chemistry 776 (2016) 127–133

electrolyte (HNO3; without uranyl ions) under identical experimental conditions is reproduced in the same figure as an inset. Two distinct well resolved reduction peaks (Ic and IIc) and two oxidation peaks (Ia and IIa) observed in presence of uranium in the same potential range 0.3 to − 0.65 V at the scan rate of 0.1 V s − 1 (Fig. 1) indicates that reduction of U(VI) in very low concentration of nitric acid (0.05 M) is a two-step single electron transfer process. The reduction peaks Ic and IIc correspond to the reduction of U(VI) to U(V), and U(V) to U(IV) respectively, at Pt electrode. The oxidation peaks Ia and IIa are attributed to the oxidation of U (IV) to U (V), and U (V) to U (VI) respectively. However, the peak pertaining to the reduction of U (V) to U (IV) and the corresponding oxidation peak were not well resolved in the CV recorded at the lower scan rate of 0.01 V s − 1 . In the two-step reduction of U (VI) to U (IV), the unstable U (V) ion produced in the first reduction step undergoes a bimolecular disproportionation reaction to give U (VI) and U (IV) as reported earlier [36] and hence, no peak corresponding to the formation of U (V) appeared at low scan rates. Accordingly, the oxidation peak of U(V) to U(VI) did not appear when the scan rate was 0.01 V s− 1 in Fig. 1. Controlled potential electrolysis of a solution containing 50 mM U(VI) in 0.05 M HNO3 was carried out for two hour at −0.12 V, which was marginally more negative than the onset potential of −0.05 V for uranyl reduction. Estimation of uranium in the solutions before and after electrolysis by wet chemical as well as UV–Visible spectral analyses revealed that the solutions contained only U(VI) ions, implying that in the absence of a stabilizing agent like a holding reductant, U(IV) generated in nitric acid medium would be oxidized to U(VI) instantaneously. 3.2. Electrochemical behaviour of U(VI)–HNO3-N2H4 system at Pt electrode The electro-reduction of U(VI) to U(IV) in nitric acid solutions proceeds along with the simultaneous reduction of nitric acid to oxides of nitrogen and water via the intermediate, HNO2 [37]. This nitrous acid is always in equilibrium with nitric acid, and since it is a powerful oxidizing agent, it chemically reoxidizes U(IV) to U(VI) spontaneously. Thus, it is essential to add a holding reductant to scavenge nitrite ions during electrolysis. As hydrazine does not introduce any corrosive metal ion into the waste stream, it is used for stabilizing U(IV) in nitrate solutions. Therefore, it will be worthwhile to investigate the redox behavior of uranyl ions in nitric acid medium in the presence of hydrazine. The cyclic voltammograms recorded for a solution of 50 mM U(VI) in 1 M nitric acid in the presence of 0.25 M N2H4 at two different scan rates with platinum working electrode at 298 K are presented in Fig. 2. The CVs recorded for the same solution without hydrazine, under identical conditions are also presented in the same figure as an inset. In Fig. 2 the narrow range of 0 to − 0.3 V was selected for the potential to visualize the peak clearly since the peak will not be discernable in wide scan. Since hydrazine stabilizes U(IV) in nitric acid medium, one would expect a prominent oxidation peak in the reverse scan for the electrolyte solution in the absence of hydrazine. However, the CVs of the two solutions (with and without hydrazine) could not be distinguished; the reason for which is not known at present. Only one reduction wave was observed even over a wide scan of potential up to − 1 V and no oxidation wave appeared (Fig. 2), both in the presence or absence of hydrazine. Ikeda-Ohno et al. [38] reported that uranium exists in various complex species in aqueous nitric acid medium, and the nature of coordination complexes varies with the concentration of nitric acid. Further, U(IV) forms spherical coordinating U 4 + cation, [U IV(H2O) x(NO 3) 5 ] − and U(VI) forms trans dioxouranyl cation [UVIO2(OH)5]3 − in aqueous nitric acid solutions. Oxidation of U(IV) to U(VI) as per the following reaction involves the formation of oxygen bonds with uranium (oxygen bridges in the formation of U(VI) complex in aqueous solutions).

129

Fig. 2. Cyclic voltammograms of 50 mM U (VI)–1 M HNO3 with and without hydrazine recorded at Pt electrode at 298 K.

Hence, reaction (1) would be highly irreversible, and the reduction of U (VI) to U (IV) alone is possible. þ − U 4þ þ 2H2 O→UO2þ 2 þ 4H þ 2e

ð1Þ

This indicates that the reduction of U (VI) to U (IV) is an irreversible process in the presence of hydrazine, as the oxidation of U (IV) involves the formation of oxygen bridges [22]. In nuclear reprocessing plants, reduction of 0.21–0.42 M (50–100 g/l) of uranyl ions is carried out at a higher acidity of 1.5 to 2 M HNO3. During electro-reduction of U (VI), the concentration of nitric acid decreases owing to its reduction simultaneously. If the final acidity drops below 0.5 M, precipitation of uranyl ions occurs, which is undesirable and should be avoided. Hence, CV studies were also performed at a higher acidity of 1.5 M HNO3 with 0.25 hydrazine and the voltammograms are presented in Fig. 3. In Fig. 3, at the low scan rate of 0.01 V.s−1 a reduction wave was clearly found nevertheless, it became broad and started at the onset potential of −0.07 V for the high scan rate of 0.1 V.s−1and ended at −0.22 V. To

Fig. 3. Cyclic voltammograms of 50 mM U(VI)–1.5 M HNO3 with hydrazine recorded at Pt electrode at 298 K.

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confirm it to be a peak, the CV was recorded for the same system without uranyl ions at 0.01 V.s−1scan rate included in the same figure as an inset, which clearly indicates there is no current maxima in the measured potential range compared to the peaks obtained in presence of uranyl ions. This indicates the peak obtained corresponds to reduction of uranyl ions. The cathodic peak current (icp) in Fig. 3 was observed to increase and the cathodic peak potential (Ecp) shifted cathodically with an increase in scan rate from 0.01 to 0.10 V s−1. The difference between the values of cathodic peak and half-peak potentials ðjEcp −Ecp =2 jÞ was found to be marginally higher (69 mV at 298 K) than the value required for a reversible process (55 mV at 298 K). This indicates that the reduction of uranyl ions in nitric acid-hydrazine medium was not a reversible process [39]. The peak current for the reduction of U (VI) increased, and the reduction wave shifted to more positive direction with an increase in nitric acid concentration. Theoretically, the reduction of uranyl ions should occur at 0.123 V vs. Ag/AgCl; however, the apparent potential for the reduction of U (VI) was found to be about −0.16 V vs. Ag/AgCl in the presence of acid and hydrazine, indicating that the reduction of U (VI) at Pt working electrode requires a higher over potential of 0.037 V. This observation upholds the fact that the reduction is not a reversible process. Controlled potential electrolysis of a solution containing 100 mM U(VI) in 1 M HNO3 and 0.25 M N2H4 was carried out for three hour at −0.12 V (slightly more negative than the onset potential for the reduction wave) vs. Ag/AgCl reference electrode. The surface areas of working and counter electrodes were 35 and 40 cm2 respectively. Samples were collected from the cell at intervals of every one hour during electrolysis for chemical and spectral analyses. The results of the analyses are given in Table 1 and Fig. 4 respectively. For the generation of 33.5 mM uranous ions at the applied potential of −0.12 V, the electricity consumed experimentally was 918.4 coulomb, which value is close to the theoretical value of 911.0 coulomb if the reduction process had followed a 2 electron transfer or two stage single electron transfer. The UV–Visible spectra shown in Fig. 4 indicates that the reduction could be a 2 electron transfer process. The intensities of absorption bands at 478, 545 and 657 nm due to tetravalent uranium were found to increase with time. The absorption bands corresponding to U (V) species could not be observed in the UV–Visible spectra. Probably, the U(V) species produced from the reduction of U(VI) might have disproportionated immediately to U(VI) and U(IV). The controlled potential electrolysis of a solution containing a higher concentration of 199 mM U(VI) (47.4 g/l) in 1 M HNO3 and 0.25 M N2H4 (similar to the concentration that is used in the actual plant) was carried out for 5 h at −0.12 V against Ag/AgCl reference electrode. Based on the results of controlled potential electrolysis experiments, it is concluded that the reduction of uranyl ions is a two electron transfer process even with 1 M acidity. One of the objectives of the present study was to compare the current and conversion efficiencies between constant potential and constant current in the electro-reduction process for the generation of U (IV). Hence, constant potential electrolysis was performed at an applied potential of −0.12 V (close to the onset potential of −0.05 V obtained from the CV run for the reduction wave). The results of the experiment listed in Table 1 for the initial concentrations of U(VI) as 100 mM, revealed that the conversion efficiency was poor when the reduction was carried out at this potential, though the current efficiency was nearly 100%. As Ti is

Fig. 4. UV–Visible spectra of 100 mM U(VI) in 1 M nitric acid and 0.25 M N2H4 during constant potential electrolysis.

the material of construction for the cathode vessel in the plant, controlled potential electrolysis of a solution containing 199 mM (47.4 g/l) U(VI) in 1 M HNO3 and 0.25 M N2H4 was carried out for the complete reduction of U(VI) at three different potentials, viz. -0.5 V (onset potential), −0.75 V (first reduction peak) and −1.0 V (second reduction peak) using Ti as the working electrode (instead of Pt) against Ag/AgCl reference, based on the CV results. The results are discussed in Section 3.6. 3.3. Estimation of kinetic parameters for the electro-reduction of uranyl ions Fig. 5 represents the CVs of 50 mM U(VI) ions in 0.05 M HNO3 recorded at Pt electrode with various scan rates from 0.01 to 0.10 V s−1 at 298 K. The cathodic peak at −0.245 V corresponds to the reduction of U(VI) to U(V) and the corresponding oxidation peak at − 0.03 V was obtained at 0.01 V s−1 scan rate. For a reversible charge transfer process, the peak potential (Ep) is independent of scan rate (ν) and

Ecp −Eap ¼ 2:29

RT nF

ð2Þ

Table 1 Concentration of U(IV), free acidity and N2H4 during electro-reduction of 100 mM U(VI) in 1 M HNO3 and 0.25 M N2H4; Initial volume: 145 ml; final volume: 141 ml. Time/h

[U(IV)]/mM

Free acidity/M

[N2H4]/M

0 1 2 3

0 8.9 18.4 33.5

1.006 0.996 0.989 0.988

0.250 0.246 0.244 0.245

*CE: Current efficiency: 98.4%.

Fig. 5. Cyclic voltammograms of 50 mM UVI) in 0.05 M HNO3 recorded at Pt electrode with different scan rates at 298 K.

S. Mishra et al. / Journal of Electroanalytical Chemistry 776 (2016) 127–133

It is evident from Fig. 5 that the reduction peak potential shifted significantly from −0.245 to −0.288 V (against Ag/AgCl reference) as the scan rate increased from 0.01 to 0.1 V s−1. The difference in the value of (Ecp − Eap) was higher (172 mV at 298 K) than the value required for a reversible process (58 mV at 298 K) and it indicates that the process is not a reversible process [39]. The measurable parameters from the cyclic voltammograms in Fig. 5 are tabulated in Table 2. The important criterion for the irreversible charge transfer kinetics is the shift in the peak potential with scan rate. The relation between the cathodic peak current, icp and the scan rate is  1 1 = ðαnα ÞFυ =2 icp ¼ 0:496nFC 0 AD0 2 RT

ð3Þ

where n is the number of electrons involved in the charge transfer reaction, F is the Faraday constant, A is the area of the electrode (in cm2), Co is the bulk concentration of electro-active species (mol cm−3), R is the gas constant, T is the absolute temperature (K), ν is the scan rate (V s−1), Do is the diffusion coefficient (cm2 s−1), α is the charge transfer coefficient and nα is the number of electrons involved in the rate determining step. The value of αnα can be obtained using the equation 1:857 RT c Ep −Ecp =2 ¼ αnα F

ð4Þ

Where Ecp =2 is the half-peak potential. Making use of the slope of the plot, icp vs. ðυÞ =2 and substituting the value of 0.58 for αnα and the values of other parameters in Eq. (3), the diffusion coefficient was determined to be 2.19 × 10-6 cm2s-1 at 298 K for the scan rate 0.01 V.s-1 . The shift in peak potential and broadening of peak shape with scan rate were observed in the cyclic voltammograms (Fig. 5) for the reduction of uranyl ions in a very low concentration of nitric acid (0.05 M). The difference between the cathodic and anodic peak potentials (ΔEp)  increased with scan rate, and the average of the peak potentials 1 2 ðEcp þ 1

Eap Þ was more or less constant for all scan rates, indicating that the process could be quasi-reversible. The heterogeneous charge transfer rate constant (ks) can be calculated using Eq. (5), which was proposed by Klingler and Kochi [33] based on peak separation.  2     vF 1=2 α nF c exp Ep −Eap ks ¼ 2:18 D0 ðαna Þ RT RT

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The heterogeneous charge transfer rate constant (ks) can also be computed by Nicholson method [34] using the following equation: ψ¼

ks 1= 2

½πD0 ðnF=RT Þν 

ð6Þ

Eq. (6) is applicable if (Ecp − Eap) is within 200 mV for a particular scan. The value of ψ was obtained from Nicholson theoretical plot. The value of ks calculated by this method at the scan rate of 0.01 V s−1 was 2.78 × 10−4 cm s−1. This is in good agreement with the value obtained using Klingler and Kochi method. 3.4. Chronopotentiometric studies on the reduction of uranyl ions The electrochemical behavior of U(VI) ions in 0.05 M nitric acid was also investigated using chronopotentiometry. Fig. 6 shows the chronopotentiograms recorded at various applied currents at a platinum electrode. The transient time (τ) observed in chronopotentiograms is a measure of time elapsed between the commencement of constant potential and the time at which the concentration of electro-active species resulting from diffusion reaches zero at the electrode. A relation between the applied current and transient time is given by Sand's equation [39], which enables the determination of diffusion coefficient. 1

1 nFAðD0 πÞ iτ =2 ¼ 2

=2

C0

ð7Þ

The Do value of 1.90 × 10−6 cm2 s−1 calculated using Eq. (7) at room temperature is in reasonably good agreement with the values of diffusion coefficient determined from cyclic voltammetric results. 3.5. Reduction of uranyl ions at glassy carbon/gold working electrodes Cyclic voltammograms were recorded using the working electrodes glassy carbon (GC) and gold for a solution of 50 mM U(VI) in 0.05 M nitric acid. Two well-distinguished reduction peaks were observed with both working electrodes (GC and Au) at the higher scan rate of 0.1 V s−1, as could be found with Pt electrode, but at lower scan rates the reduction peak at higher negative potential disappeared. The reduction waves for uranyl ions appeared at the peak potentials −0.217 and − 0.214 V (vs. Ag/AgCl) with glassy carbon and gold electrodes

ð5Þ

According to Eq. (5), if [ks/(ν) ½] N 0.11, the process is reversible. If 0.11 N [ks/ (ν) ½] N 3.7 × 10−6, it is a quasi-reversible process. If 3.7 × 10−6 N [ks/ (ν) ½], the process is irreversible. The value of ks at the scan rate of 0.01 V s−1 was determined to be 2.52 × 10− 4 cm s− 1, using Do = 2.19 × 10− 6 cm2 s−1. The value of 2.52 × 10−3 cm s(− 1/2) V(− 1/2) calculated for [ks/(ν) ½] for the same scan rate revealed that the reduction of uranyl ions at a very low concentration of nitric acid, without hydrazine is a quasi–reversible process. Table 2 Peak parameters obtained from the CVs recorded with Pt for 50 mM U(VI) in 0.05 M HNO3. 

ðEcp þ Eap Þ=V

Scan rate/(V s−1)

Ecp/V

icp/A

Eap/V

(Ecp − Eap)/V

½1

0.01 0.02 0.04 0.06 0.08 0.10

−0.245 −0.253 −0.265 −0.276 −0.283 −0.288

−7.15E-04 −9.60E-04 −1.36E-04 −1.69E-04 −1.89E-04 −2.07E-04

−0.073 −0.057 −0.045 −0.037 −0.034 −0.026

0.172 0.196 0.220 0.239 0.249 0.262

0.159 0.155 0.155 0.156 0.158 0.157

2

Fig. 6. Chronopotentiograms of 50 mM U(VI) in 0.05 M HNO3 recorded at Pt electrode at different applied currents.

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respectively at the scan rate of 0.01 V s−1 compared to the value of −0.245 V (vs. Ag/AgCl) at a platinum working electrode. Peak potential shift and the broadening of peak shape with scan rate were observed in the CVs with both the working electrodes. The increasing trend in the difference between the cathodic and anodic peak potentials (ΔEp) with scan rate and the more or less constant value for the average of  the peak potentials, 1 2 ðEcp þ Eap Þ at different scan rates revealed that the reduction process was quasi-reversible for the various working electrodes employed in the present study. The values of diffusion coefficient for the reduction of uranyl ions at glassy carbon and gold electrodes calculated based on Eq. (3) were 3.46 × 10−6 and 3.13 × 10−6 cm2 s− 1 respectively. These values are in close agreement with the value of 2.19 × 10− 6 cm2 s− 1 obtained for Do, using Pt electrode. The Do values derived were also consistent with those values obtained from chronopotentiometric studies at glassy carbon (3.02 × 10−6 cm2 s−1) and gold (2.73 × 10−6 cm2 s−1) working electrodes respectively. The heterogeneous electron transfer rate constant, ks determined using glassy carbon (1.6 × 10−4 cm s− 1) and gold electrode (2 × 10−4 cm s−1) also confirmed the reduction of uranyl ions at very low concentrations of nitric acid in the absence of hydrazine to be a quasi–reversible reaction. The values of heterogeneous electron transfer rate constant (ks) for the reduction of uranyl ions at a low concentration of nitric acid (0.05 M), derived using Klingler and Kochi equation [33] as well as by Nicholson equation [34] at different working electrodes are compared in Table 3. The diffusion coefficient values calculated using the results of cyclic voltammetric and chronopotentiometric experiments are also compared in this Table.

3.6. Reduction of uranyl ions with titanium as the working electrode The use of platinum as an electrode material is limited due to its high cost in bulk form even though Pt offers high conductivity, good electrocatalytic activity and excellent resistance to oxidation and corrosion. Hence, titanium is extensively used as electrodes in electrochemical industries. The reduction behavior of U(VI) ions in 0.05 M nitric acid was investigated by cyclic voltammetry employing Ti as the working electrode (SA: 0.385 cm2). The CVs recorded for different scan rates at 298 K over the potential range 1.0 to - 1.5 V (versus Ag/AgCl reference electrode) are reproduced in Fig. 7. Two distinct reduction peaks at − 0.764 and − 1.05 V (against Ag/ AgCl) and one oxidation peak at 0.182 V were observed in the potential range 1.0 to–1.5 V, for the scan rate of 0.01 V s−1 in 0.05 M HNO3. With increase in the scan rate, the peak current increased and the two reduction waves merged, resulting in the disappearance of the peak at − 0.764 V. This behavior, with respect to the disappearance of the peak corresponding to the reduction of U(VI) to U(V) at higher scan rates at Ti electrode was contrary to the behavior observed at Pt electrode, wherein the reduction wave for U(VI) to U(V) merged at the very low scan rate of 0.01 V s−1. Similar to Pt working electrode, a shift in the peak potential and broadening of peak shape with scan rate were observed in the cyclic voltammograms recorded using Ti electrode. However, it was not possible to determine the charge transfer rate constant from the CVs, due to the absence of prominent oxidation peak. Based on the CV results, controlled potential electrolysis of a Table 3 Kinetic parameters determined using different working electrodes. Kinetic parameters

50 mM U(VI) in 0.05 M HNO3 Platinum

Glassy carbon

Gold

Do (cm2 s−1)/CV expt. Do (cm2 s−1)/CP expt. ks (cm s−1)/Klingler and Kochi [33] ks (cm s−1)/Nicholson [34]

2.19 × 10−6 1.90 × 10−6 2.52 × 10−4 2.78 × 10−4

3.46 × 10−6 3.02 × 10−6 2.61 × 10−4 3.69 × 10−4

3.13 × 10−6 2.73 × 10−6 2.20 × 10−4 3.85 × 10−4

Fig. 7. Cyclic voltammograms of 50 mM U(VI) in 0.05 M HNO3 recorded with Ti electrode for different scan rates at 298 K.

solution containing 199 mM (47.4 g/l) U(VI) in 1 M HNO3 and 0.25 M N2H4 was carried out for the reduction of U(VI) at three different potentials, viz. −0.5 V (onset potential), −0.75 V (first reduction peak) and −1.0 V (second reduction peak) against Ag/AgCl reference. The surface area of working (Ti) and counter (Pt) electrodes were 120 and 140 cm2 respectively. The results of the three controlled potential electrolysis experiments with titanium working electrode are presented in Table 4. This Table reveals that current efficiency of N 90% can be achieved if the reduction is carried out at the peak potentials; nevertheless, better conversion efficiency could be achieved when the electrolysis was carried out at the second reduction potential. To compare the current efficiency of the process in the electro-reduction of U(VI) at constant potential conditions with that of the reduction reaction carried out under constant current condition, an electroreduction experiment was performed using a Ti mesh cylindrical cathode of surface area, 120 cm2 in a divided cell assembly (a porous ceramic diaphragm separated the cathode and anode compartments) using the stock solution of 199 mM U(VI) with 1 M acidity and 0.25 M hydrazine at a fixed current density of 6 mA/cm2. The current density was decided based on our earlier studies [40]. The concentration of U(IV), nitric acid and hydrazine analyzed as a function of time during electrolysis in the divided cell are tabulated in Table 5. It is apparent from Tables 4 and 5 that the constant potential reduction process is an energy efficient process compared to constant current reduction process since a part of the energy is utilized for the reduction Table 4 Results of the electrolysis experiments for the electro-reduction of 199 mM U(VI) in 1 M HNO3 and 0.25 M N2H4 with Ti working electrode at the onset, first and second reduction potentials. Time (h)

Potential: −0.5 V 4+

0 0.5 1.0 1.5 2.0 3.0 4.0 C.E. (%)

Potential: −0.75 V

Potential: −1.0 V

[U ] Free (M) acidity (M)

[N2H4] [U ] Free (M) (M) acidity (M)

[N2H4] [U4+] Free (M) (M) acidity (M)

[N2H4] (M)

0 – 0.047 – 0.126 0.160 0.198 67

0.247 – 0.212 – 0.187 0.175 0.168

0.247 0.223 0.191 0.170 0.125 – –

0.247 0.226 0.145 – – – –

1.070 – 0.936 – 0.841 0.822 0.791

4+

0 0.059 0.128 0.161 0.199 – – 91.3

1.070 0.966 0.914 0.886 0.879 – –

0 0.114 0.199 – – – – 90.2

1.070 0.988 0.897 – – – –

S. Mishra et al. / Journal of Electroanalytical Chemistry 776 (2016) 127–133 Table 5 Concentration of U(IV), free acidity and N2H4 during the electro-reduction of 199 mM of U(VI) in 1 M nitric acid and 0.25 M N2H4 using Ti cathode at 6 mA/cm2 in a divided cell. Initial volume of catholyte: 140 ml; Final volume: 136 ml. Sl. No.

Time (h)

[U(IV)] (M)

Free acidity (M)

[N2H4] (M)

1 2 3 4

0 1 2 2.5

0 0.096 0.166 0.198

1.001 0.828 0.654 0.552

0.498 0.470 0.451 0.435

*CE: Current efficiency: 81%.

of nitric acid in the latter case. In the constant current reduction process, almost 0.5 mol of nitric acid got reduced, whereas only 0.1 to 0.12 mol of nitric acid was reduced in the constant potential reduction process. 4. Conclusions At very low concentrations of nitric acid, reduction of uranyl ions proceeded through the two steps: U (VI) to U (V) and subsequently U (V) to U (IV). As the concentration of the supporting electrolyte (nitric acid) increased, the reduction reaction followed a two electron single step. Further, at very low concentration of nitric acid, in the absence of hydrazine, the reduction of uranyl ions was found to be quasi–reversible and hence, under mixed control. As the acidity of the electrolyte was increased, the reduction was found to be irreversible and was purely under kinetic control. The diffusion coefficient values at different working electrodes were comparable. Chronopotentiometric results validated the diffusion coefficient values determined from cyclic voltammetric experiments. The value of heterogeneous electron transfer rate constant for uranyl reduction at Pt working electrode, estimated at low concentration of nitric acid (0.05 M) using Klingler and Kochi equation was in good agreement with the value obtained using Nicholson equation. The energy efficiency of constant potential reduction process was found to be better than that of the constant current reduction process. References [1] D.D. Sood, S.K. Patil, Chemical aspects of the back end of nuclear fuel cycle, Frontiers in Nuclear Chemistry, Perfect Prints, Thane, 1996 104. [2] S.K. Sali, D.M. Noronha, H.R. Mhatre, M.A. Mahajan, K. Chander, S.K. Agarwal, V. Venugopal, A novel methodology for processing of plutonium bearing waste as ammonium plutonium(III)-oxalate, Nucl. Technol. 151 (2005) 289. [3] P. Regnaut, P. Faugeras, A. Braut, R. Helou, A. Redon, The Processing Of Irradiated Uranium In The Fontenary-aux-ROSES PILOT PLANT, Proc. 2nd U.N. Internatl. Conf. on Peaceful Uses of Atomic Energy, Geneva, Article 17, 73, 1958. [4] S.K. Rao, R. Shyamlal, C.V. Narayan, U. Jambunathan, A. Ramanujam, V.P. Kansra, Uranous nitrate production for PUREX process applications using PtO2 catalyst and hydrazine nitrate as a reductant, Report BARC/2003/E/009, 2003. [5] J.M. McKibben, J.E. Bercaw, Hydroxylamine nitrate as a plutonium reductant in the Purex solvent extraction process, Report DP-1248, 1971. [6] D.J. Reif, Applicability of hydroxylamine nitrate reductant in Pulse-column contactors, Report DP-1649, 1983. [7] A. Naylor, Pu–U separation techniques in TBP systems: use of ferrous sulphamate, U(IV) and aqueous soluble complexing agents, Advanced course in reprocessing of fuel from present and future power reactors; organized by Netherlands –Norwagian Reactor School at Institute for Atomenergi, Norway, August 21–Sept. 01 1967, pp. 172–200. [8] S. Kumar, R. Kumar, S.B. Koganti, Contributions of non-extracting Pu reductants (ferrous sulphamate and hydroxylamine nitrate) and holding reductant (hydrazine nitrate) to the aqueous density in U–Pu partitioning system, Report IGC-273, 2005.

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