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Model for predicting hydroxylamine nitrate stability in plutonium process solutions
Journal of Loss Prevention in the Process Industries 24 (2011) 76e84
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Model for predicting hydroxylamine nitrate stability in plutonium process solutions G. Scott Barney a, *, Paul B. Duval b a b
105903 N. Harrington Rd., West Richland, WA, USA Shaw AREVA MOX Services, LLC, Aiken, SC, USA
a r t i c l e i n f o
a b s t r a c t
Article history: Received 29 July 2010 Received in revised form 27 October 2010 Accepted 29 October 2010
A mathematical model is described that may be used to determine the safety of hydroxylamine nitrate (HAN) solutions used in solvent extraction purification of plutonium. The most significant hazard associated with hydroxylamine use in processing plutonium is its rapid, autocatalytic reaction with nitric acid which can result in an explosion or pressurization of process vessels with radiological consequences to humans. In addition, heat is produced by the reaction that could potentially ignite process solvents. The HAN decomposition reaction can occur only under specific process conditions (temperature; HAN, plutonium and nitric acid concentrations) and the model is used to identify these conditions so that they can be avoided. A kinetics model has been developed using all of the known significant reactions that could occur in process solutions containing HAN and nitric acid as well as plutonium and iron. The reaction kinetics data (rate laws, rate constants, activation energies) used in the model were obtained from chemical literature sources. The model shows that the autocatalytic HAN reaction with nitric acid is very rapid and is catalyzed by Pu(III) and Fe(II) in process solutions. High temperatures and nitric acid concentrations also promote the reaction. Ó 2010 Elsevier Ltd. All rights reserved.
Keywords: Hydroxylamine stability Plutonium kinetics Nitric acid Model
1. Introduction Solvent extraction processing in the nuclear industry often utilizes hydroxylamine nitrate (NH3OHNO3 or HAN) to separate plutonium from uranium and other impurities by stripping plutonium from an organic solvent containing tributyl phosphate. Aqueous HAN is used to reduce Pu(IV) in the organic phase to Pu(III) which is not very soluble in the organic phase. Industrial experience with HAN has been acquired in France, Japan, and the UK for more than 20 years without any reported incidents (DOE, 1998). However, the use of HAN in the plutonium purification process with nitric acid solutions introduces the possibility of an autocatalytic reaction that may produce unacceptable temperature and pressure increases to the system and consequently result in overpressurization and/or explosions within the process vessels. Several accidents have been reported in the United States (DOE, 1998) involving explosions or pressurization of process vessels by HAN reactions in the nuclear process industry. It is noted that all these accidents occurred in storage, chemical make-up vessels, or evaporation vessels and not in the extraction process. However it is
* Corresponding author. Tel.: þ1 509 967 2549. E-mail address: [email protected] (G.S. Barney). 0950-4230/$ e see front matter Ó 2010 Elsevier Ltd. All rights reserved. doi:10.1016/j.jlp.2010.10.004
prudent to determine the conditions for safe operations to prevent any incident in the extraction process. Recent research on the stability of HAN in aqueous solutions includes kinetic modeling of the oxidation of HAN by nitric acid using ab initio molecular orbital calculations (Ashcraft, Raman, & Green, 2008; Raman, Ashcraft, Vial, & Klasky, 2005), calorimetric and spectroscopic measurements (Kumasaki, 2004; Kumasaki, Fujimoto, & Ando, 2003; Liu, Papadaki, et al., 2009; Liu, Wei, Guo, Rodgers, & Mannan, 2009), and effects of metal ions on HAN stability (Sugikawa, Umeda, Sekino, Matsuda, & Kodama, 2002). A significant amount of work has also been performed on the stability of hydroxylamineewater solutions (Cisneros, Rodgers, & Mannan, 2002, 2004; Cisneros et al., 2003) prompted by two tragic industrial accidents involving hydroxylamine solutions (Long, 2004; Yusaku, Masahide, Takashi, & Hiroshi, 2005). None of these studies include the effects of dissolved plutonium on HAN stability. In addition, the nitric acid concentrations and temperatures were not in the range used in plutonium purification. A mathematical model is described in this paper that calculates the rates of the reactions involved in the oxidation of HAN under conditions that exist in a plutonium purification solvent extraction process. The process conditions required for the rapid, autocatalytic HAN reaction are identified by the model so that these conditions can be avoided. The model predicts the rate of the reactions
G.S. Barney, P.B. Duval / Journal of Loss Prevention in the Process Industries 24 (2011) 76e84
involved and resulting temperature changes and gas generation rates in the process solutions. 2. Chemical reactions 2.1. Autocatalytic reaction The chemical reactions involved in the autocatalytic oxidation of hydroxylamine nitrate (NH3OHNO3) with nitric acid have been proposed by several laboratories (Gowland & Stedman, 1981; Bennett, Brown, Maya, & Posey, 1982). Hydroxylammonium ion is the major hydroxylamine species in solutions with a pH lower than about 5.8 and will be the major species in the highly acidic process solutions. The overall autocatalytic reaction can be written as follows: NH3OHþ þ 0.5(x þ 1)HNO3 / xHNO2 þ 0.25(3 x)N2O þ Hþ þ (1.75 0.25x)H2O
(Reaction 1)
Nitrous acid production (Reaction 2)
Nitrous acid scavenging
NH3OHþ þ HNO2 / N2O þ Hþ þ 2H2O
Gowland and Stedman (1981) and Pembridge and Stedman (1979) have proposed a mechanism for the autocatalytic reaction that includes Reactions (4), (5), (8), and (9). They state that the autocatalytic reaction is limited by Reaction (5), which is relatively slow. Nitric acid solutions of HAN are stable as long as the nitrous acid scavenging reaction (3) is dominant over the autocatalytic reaction, which produces nitrous acid. Reaction (7) accounts for the alternate decomposition of nitrous acid which yields nitric oxide. The reverse reaction can occur to produce nitrous acid and nitric dioxide [Reaction (6)]. Reactions of hydrazinium ion (N2Hþ 5 ) and hydrazoic acid (HN3) with nitrous acid [Reactions (10) and (11)] are important because hydrazine (N2H4) is used as a nitrous acid suppression agent in solvent extraction processes using HAN. 2.3. Reactions of plutonium species in solution
The reaction is strongly exothermic and produces a gaseous reaction product (N2O) that can pressurize process equipment. The key to initiation of the autocatalytic reaction of HAN is the production of nitrous acid (HNO2). Nitrous acid is produced and consumed by NH3OHþ by the following overall reactions:
NH3OHþ þ 2HNO3 / 3HNO2 þ Hþ þ H2O
77
(Reaction 3)
If the nitrous acid concentration reaches a critical threshold, Reaction (2) dominates and nitrous acid is rapidly produced. Any reaction that can increase the nitrous acid concentration can lower the decomposition temperature of the solutions. Dissolved iron has been shown to lower the decomposition temperatures of HAN solutions (Sugikawa et al., 2002) and there are reactions of iron in nitric acid solution that will produce nitrous acid. The detailed reactions that were considered in this paper are listed below. 2.2. Reactions of nitrogen species in solution NH3OHþ þ 2HNO3 / 3HNO2 þ Hþ þ H2O
(Reaction 2)
HNO2 þ NH3OHþ / N2O þ 2H2O þ Hþ
(Reaction 3)
þ 2NO2 þ H2O / HNO2 þ NO 3 þH
(Reaction 4)
þ HNO2 þ NO 3 þ H / 2NO2 þ H2O
(Reaction 5)
þ NO þ NO 3 þ H / HNO2 þ NO2
(Reaction 6)
þ NO2 þ HNO2 / NO þ NO 3 þH
(Reaction 7)
2NO2 þ NH3OHþ / HNO þ 2HNO2 þ Hþ
(Reaction 8)
2NO2 þ HNO þ H2O / 3HNO2
(Reaction 9)
þ N2Hþ 5 þ HNO2 / HN3 þ 2H2O þ H
(Reaction 10)
HN3 þ HNO2 / N2 þ N2O þ H2O
(Reaction 11)
Dissolved plutonium and iron can participate in reactions with HAN. For most nuclear materials that are processed, the concentrations of plutonium in the solutions downstream from the plutonium stripping column will be relatively high compared to iron. Plutonium will likely catalyze the autocatalytic HAN reaction with nitric acid by reactions analogous to the iron reactions that destabilize HAN solutions. Plutonium in these solutions will be reduced species of Pu(III) which can react with nitric acid to produce nitrous acid. Any reaction mechanism that can produce nitrous acid will potentially destabilize the solutions with respect to the autocatalytic HAN reaction. The reactions likely to be important are Reactions (12) and (13) below. Nitrate complex formation as given by Reactions (14) and (15) are important in the model calculations, but are very rapid and reversible so they are considered equilibrium reactions. 2Pu4þ þ 2NH3OHþ / 2Pu3þ þ N2 þ 4Hþ þ 2H2O
(Reaction 12)
4þ þ HNO2 þ H2O 2Pu3þ þ 3Hþ þ NO 3 / 2Pu
(Reaction 13)
3þ Pu4þ þ NO 3 / PuNO3
(Reaction 14)
2þ Pu3þ þ NO 3 / PuNO3
(Reaction 15)
The kinetics of these reactions has been studied and rate laws and rate constants have been proposed and measured. These rate data have been included in the HAN stability model. 2.4. Reactions of iron species in solution Iron(III) can be reduced by HAN to Fe(II) (Bengtsson, Fronaeus, & Bengtsson-Kloo, 2002) which can then react with nitric acid to produce nitrous acid (Epstein, Kustin, & Warshaw, 1980). Several additional reactions of Fe3þ and Fe2þ can occur with NO2, HNO2, and NO that can exist in the process solutions. They are listed below. Fe3þ þ NH3OHþ / Fe2þ þ ½N2 þ 2Hþ þ H2O
(Reaction 16)
Fe2þ þ NO2 þ Hþ / Fe3þ þ HNO2
(Reaction 17)
Fe2þ þ HNO2 þ Hþ / Fe3þ þ NO þ H2O
(Reaction 18)
Fe3þ þ NO þ H2O / Fe2þ þ HNO2 þ Hþ
(Reaction 19)
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3. Rate equations used in model
4.1. Rate constants with limited information
The model calculates concentrations of each significant species involved in the autocatalytic reaction over time for a given set of initial conditions (temperature, ionic strength, initial concentrations of HAN, HNO3, HNO2, Fe3þ, Pu3þ, Pu4þ, etc.) The significant species, based on the proposed reaction mechanism, are Fe3þ/Fe2þ, Pu4þ/ þ Pu3þ, NH3OHþ, NO2, HNO2, NO, HNO, Hþ, NO 3 , N2, N2O, N2H5 , and HN3. The differential equations describing the changes in concentrations over time are given in Tables 1 and 2. The second order equation, d[NH3OHþ]/dt ¼ k3[HNO2][NH3OHþ], is used for Reaction (3) in these combined equations and the acid dependency of k3 (the second order rate constant) is calculated using published data. The model receives the input data for initial conditions of the HANeHNO3eH2N4eFeePu solutions and then simultaneously solves this set of differential equations for concentrations of the significant species and solution temperature. IGOR PRO version 6.0.6 (Wavemetrics, Inc.) programming language was utilized to solve the equations.
Some of the rate constants used in the model are limited to a single value at a specific temperature and solution ionic strength. For others, activation energies are available. Table 3 lists values for the rate constants where limited information is available.
4. Rate constants Rate constant data were found for each of the 19 reactions listed above except for Reactions (8) and (9). Rate constants for Reactions (8) and (9) were estimated using methods described below. Most rate constant data were obtained at a single temperature or at nitric acid solution concentrations lower than those expected in the solvent extraction process. Where necessary, rate constant values were extrapolated to process-relevant conditions using standard techniques for each of the significant reactions. A relatively small number of reactions were found to significantly affect the autocatalytic HAN reaction. In the absence of iron and plutonium, the significant reactions are (3), (4), (5), and (8)e (11). In the presence of plutonium, Reactions (12) and (13) were significant. In the presence of iron, Reactions (16)e(18) are important. Each of these reactions and the rate constants for them are discussed in the following sections.
4.2. Rate constants with known dependencies on solution composition and temperature 4.2.1. Rate constant k3 Reaction (3) is key in determining the stability of HANeHNO3 solutions as discussed above. Values for the rate constant, k3, have been published by a number of workers (Barney, 1971; Bennett et al., 1982; Pembridge & Stedman, 1979). Measurements have been made in both nitric acid and perchloric acid solutions at ionic strengths of 0.1e6.0 M and temperatures of 0e50 C. The mechanism for this reaction appears to change above approximately 2.0 M ionic strength. At lower ionic strengths, the second order rate constant increases with increasing acid concentration as seen in Fig. 1. At higher ionic strength a maximum is reached and then the rate constant decreases with increasing acid concentration. This strongly affects the model results, because nitrous acid scavenging is not as effective at higher nitric acid concentrations and causes the solutions to become more unstable. The data from the higher ionic strengths were used to model effects of nitric acid concentration, ionic strength, and temperature on k3. Activation energies for Reaction (3) have been published in Gowland and Stedman (1981), Bennett et al. (1982) and Barney (1971). A value of 41.5 kJ mol1 is used in the model to predict the dependency of k3 on temperature. The fitted equation used to predict k3 in the model is thus,
h i k3 ¼ 106 eð41;500=ð8:314*TÞÞ 234:08 þ 63:067m 85:702 Hþ h i h i2 ð12Þ 9:9333m2 þ 3:8233m Hþ þ 5:5347 Hþ
Table 1 Rate equations used in the model for calculation of HAN, HNO2, NO2, NO, HNO, N2Hþ 5 , and HN3 concentrations.
i h d NH3 OHþ dt
h ¼ k16
i2 h i2 h i2 h i h i Fe3þ NH3 OHþ ½PuðIVÞ2 NH3 OHþ 2 þ þ þ k ½HNO NH OH OH þk þ k NH ½NO h i2 h i4 h i4 2 3 2 3 8 3 2 12 Fe2þ Hþ Kd þ NO ½PuðIIIÞ2 Hþ 3
! i h ½NO ½HNO h 3þ i þ k3 ½HNO2 NH3 OHþ Fe k19 h i þ k’19 ½NO þ k}19 h i2 þ þ H H ih i h i h i2 h i h h i 2 þ þ þ þ þk7 ½HNO2 ½NO2 H 2k8 NH3 OH ½NO2 2 3k9 ½NO2 2 ½HNO 0:5k4 ½NO2 þk5 ½HNO2 NO3 H k6 NO3 ½NO H h i h i0:5 h i0:5 ih i h h i k13 þ k’13 Hþ ½PuðIIIÞ½HNO2 0:5 Hþ NO 3 h i Hþ þ k11 ½HNO2 ½HN3 Hþ 0:5 þ k10 ½HNO2 N2 Hþ 5 1 þ b NO3
h h i h i i d½HNO2 ½HNO2 ¼ k17 Fe2þ ½NO2 þ k18 þ k’18 Hþ þ k}18 Fe2þ ½HNO2 dt ½NO
h i ih i h i h i2 h i h h i d½NO2 þ ¼ k4 Fe2þ ½NO2 þ 2k4 ½NO2 2 2k5 ½HNO2 NO Hþ k6 NO þk7 ½HNO2 ½NO2 Hþ þ 2k8 NH3 OHþ ½NO2 2 þ2k9 ½NO2 2 ½HNO 3 3 ½NO H dt ! h h i i h i h i2 h i d½NO ½HNO2 ½NO ½HNO h 3þ i þ ¼ k18 þ k’18 Hþ þ k}18 Fe2þ ½HNO2 þ k19 h i þ k’19 ½NO þ k}19 h i2 þ k6 NO Fe k7 ½HNO2 ½NO2 Hþ 3 ½NO H þ þ dt ½NO H H
h i d½HNO ¼ k8 NH3 OHþ ½NO2 2 þk9 ½NO2 2 ½HNO dt h i ih i h d N2 Hþ 5 Hþ ¼ k10 ½HNO2 N2 Hþ 5 dt
(1)
ih i h h i d½HN3 ¼ k10 ½HNO2 N2 Hþ Hþ þ k11 ½HNO2 ½HN3 Hþ 5 dt
(2)
(3)
(4)
(5)
(6) (7)
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79
Table 2 Rate equations used in the model to calculate concentrations of Pu(IV), Fe(III), Hþ, and NO3.
d½PuðIVÞ ¼ k12 dt
i h d Fe3þ dt h i d Hþ dt
h ¼ k16
(8)
i2 h i2 h h i h i i Fe3þ NH3 OHþ ½HNO2 Fe2þ ½HNO2 þ k17 Fe2þ ½NO2 k18 þ k’18 Hþ þ k}18 h i2 h i4 ½NO 2þ þ Fe H
! ½NO ½HNO h 3þ i Fe k19 h i þ k’19 ½NO þ k}19 h i2 Hþ Hþ
i2 h i2 ! h h i h i i Fe3þ NH3 OHþ ½HNO2 ½NO ½HNO h 3þ i Fe2þ ½HNO2 k19 h i þ k’19 ½NO þ k}19 h i2 Fe þ k17 Fe2þ ½NO2 þ k18 þ k’18 Hþ þ k}18 ¼ 2k16 h i2 h i4 ½NO Hþ Hþ Fe2þ Hþ h h h i i ih i h i h i2 h i þ k7 ½HNO2 ½NO2 Hþ k8 NH3 OHþ k3 ½HNO2 NH3 OHþ 0:5k4 ½NO2 2 þk5 ½HNO2 NO Hþ þ k6 NO 3 3 ½NO H h i2 h i h i0:5 h i0:5 ih i h ½PuðIVÞ2 NH3 OHþ k13 þ k’13 Hþ ½PuðIIIÞ½HNO2 0:5 Hþ NO 3 2 h i Hþ k10 ½HNO2 N2 Hþ ½NO2 2k12 h i4 2 þ 1:5 5 2 þ b 1 þ NO Kd þ NO3 ½PuðIIIÞ H 3
h i d NO 3 dt
h i2 h i0:5 h i0:5 k13 þ k’13 ½Hþ ½PuðIIIÞ½HNO2 0:5 Hþ NO ½PuðIVÞ2 NH3 OHþ 3 h i h i4 2 1 þ b NO Kd þ NO ½PuðIIIÞ2 Hþ 3 3
(9)
h
2
¼ 0:5k4 ½NO2 þk5 ½HNO2
h
NO 3
ih
þ
H
i
þ
h
k6 NO 3
i
h i h i2 ½NO Hþ k7 ½HNO2 ½NO2 Hþ þ 0:5
where m is the molar ionic strength and T is the temperature in Kelvin. 4.2.2. Rate constant k5 Reaction (5) is the rate-limiting step for the autocatalytic reaction of HAN with HNO3 according to Gowland and Stedman (1981). It is very important to know the value of k5 accurately to predict the conditions required for the autocatalytic reaction. The k5 values over a range of solution compositions and temperatures must be calculated. The solution concentrations of interest are 0e0.5 M for HAN and 0.5e2 M for HNO3. The temperature range is from 25 C to 60 C. The reaction rate can be expressed in term of solution activities as follows:
d½HNO2 ¼ k5 aHNO2 aNO3 aHþ dt ih i h Hþ ¼ k5 fHNO2 fNO3 fHþ ½HNO2 NO 3
(13)
where a is the activity and f is the molar activity coefficient for a given component of the solution. Molar activity coefficients were
h i0:5 h i0:5 k13 þ k’13 ½Hþ ½PuðIIIÞ½HNO2 0:5 Hþ NO 3 h i 1 þ b NO 3
(10)
(11)
calculated for the range of solution compositions and temperatures required. Schmid and Bahr (1964) have measured k5 values over HNO3 concentrations of 0.3e5.0 M and temperatures of 0e20 C. Their data are used to calculate k5 values for different temperatures at zero ionic strength and to determine the activation energy for this reaction (66.8 kJ mol1). Molar activity coefficients for Hþ and NO 3 are calculated in the model for each time step using the Pitzer (1991, chap. 3) method. 4.2.3. Rate constant k8 Reaction (8) is very important in determining the rate of the autocatalytic reaction of HAN with nitric acid. The applicable rate law is
h i d½NO2 ¼ k8 NH3 OHþ ½NO2 2 dt
(14)
There are no published values of the rate constant for Reaction (8). It can be estimated from an equation derived by Gowland and Stedman (1981) that describes the rate of nitrous acid formation
Table 3 Rate constants used in the model for which only temperature and ionic strength effects are known. Rate constant
Value
Temperature ( C)
Ionic strength (M)
Activation energy (kJ mol1)
Reference
k4 k6 k7 k10 k11 k12 k13 k130 k16 k17 k18 k180 k1800 k19 k190 k1900
7.0E7 M1 s1 5E3 M1 s1 9.3 M2 s1 2.339E3 M2 s1 1.09E3 M2 s1 2.9E2 M5 s1 1.49E2 M1.5 s1 1.32E3 M2.5 s1 1.85E9 M3 s1 3.1E4 M1 s1 7.8E3 M1 s1 2.3E1 M2 s1 7.6E3 M1 s1 5.6E4 s1 1.6E2 M1 s1 5.4E4 s1
25 25 25 25 25 30 16 16 25 25 25 25 25 25 25 25
0.014e0.02 1.5 1.5 2.0 2.0 2.3e2.7 2.0 2.0 1.0 1.5 1.5 1.5 1.5 1.5 1.5 1.5
3.77 N/Aa N/A 56.5 54.5 130 55.8 55.8 133 N/A N/A N/A N/A N/A N/A N/A
Schwarz & White, 1983 Orban & Epstein, 1982 Orban & Epstein, 1982 Perrott & Stedman, 1976 Perrott & Stedman, 1976 Barney, 1976 Newton, 2002 Newton, 2002 Bengtsson et al., 2002 Orban & Epstein, 1982 Orban & Epstein, 1982 Orban & Epstein, 1982 Orban & Epstein, 1982 Orban & Epstein, 1982 Orban & Epstein, 1982 Orban & Epstein, 1982
a
Not available in the reference.
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G.S. Barney, P.B. Duval / Journal of Loss Prevention in the Process Industries 24 (2011) 76e84
Fig. 1. Second order rate constants for Reaction (3) versus acid concentration for different ionic strengths, temperatures, and acidesalt solutions [Bennett et al. (1982): 25 C in HNO3eLiNO3 mixtures and Barney (1971): 30 C in HNO3eNaNO3 mixtures].
Fig. 2. Comparison of calculated values of k8 from fitting the model to experimental decomposition data and from the Pembridge and Stedman (1979) measurements of HNO2 generation.
and disappearance during the autocatalytic reaction based on rate laws for Reactions (3)e(5) and (8). The equation is
solution was considered unstable if the HNO2 concentration increased over a 1000 s time period. An example of how this criterion was used is shown in Fig. 3 where HNO2 concentrations are plotted over time for different values of k8. The solution stability is quite sensitive to k8 values and these curves show that there is a definite boundary for stability at 25 C for a solution containing 0.05 M HAN and 3.0 M HNO3. The curves show that a k8 value of 4.12E8 results in a stable solution while a value of 4.13E8 gives an unstable solution. A maximum error of several degrees for the decomposition temperature can be expected when using this criterion to adjust the k8 values. Model results show that once the concentration of HNO2 begins to increase, the solution will become unstable. The experimental HAN decomposition temperature data and results of these calculations are shown in Table 4. This approach gave k8 values that are quite close to the values in Fig. 2. Fig. 2 shows a comparison of k8 values calculated from a fit to the experimental decomposition data with those of Pembridge and Stedman over a range of nitric acid concentrations. The two different sources of k8 values are in reasonably good agreement, although the slopes of the lines are slightly different. The decomposition data yield a fitted equation that predicts k8 values over a wide range of acid and HAN
i h ih i h ½HNO2 NH3 OHþ 3k5 Hþ NO 3 d½HNO2 i h ¼ k4 dt þ 2 NH3 OHþ k8 i h i h k3 ½HNO2 NH3 OHþ ¼ kn NH3 OHþ ½HNO2 ð15Þ where kn is a measured rate constant for the formation of HNO2 and steady state approximations, d[NO2]/dt ¼ 0 and d[HNO]/dt ¼ 0 are assumed. This equation can be solved for k4/k8 (and thus k8) by using the published rate data (kn values) for HNO2 generation from HAN oxidation with HNO3 given by Pembridge and Stedman (1979). Values of k3, k4, and k5 that are required for this calculation are estimated by methods described in this paper. The equation solved for k4/k8 is
h ih i i h 3k5 Hþ NO 3 k4 ¼ 2 NH3 OHþ kn þ k3 k8
(16)
The results for calculation of k8 for 2.68e7.86 M HNO3 and 0.002e0.01 M HAN are shown in Fig. 2 for 25 C. The activation energy for this reaction is not known so the effect of temperature cannot be estimated using the Pembridge and Stedman data. No estimated k8 values are available for solutions with low acid concentrations. The HAN model can be applied at acid concentrations lower than 2.86 M by extrapolating the k8 data calculated from the 25 C Pembridge and Stedman published data on HNO2 generation rates given in Fig. 2. Values of k8 calculated in this manner are plotted in Fig. 2 against HNO3 concentration. The concentration of HAN in these experiments seems to have little effect on the k8 values. Low HAN concentrations are used because the reaction is exothermic and significant heating of the solutions must be avoided to obtain accurate data. Fig. 2 shows that the fitted data for a logarithmic plot of k8 are linear and are likely to be reasonably accurate in the extrapolation. Since no temperature data other than 25 C are available for this reaction, k8 values at other temperatures were estimated by fitting published experimental data (Gowland & Stedman, 1981) with the model where HAN decomposition temperatures are measured over a range of acid and HAN concentrations. Calculated values of k3, k4, and k5 as given above are used in these calculations. The HAN
Fig. 3. Example of a stability boundary calculation for the HANeHNO3 autocatalytic reaction ([HAN]0 ¼ 0.05 M, [HNO3] ¼ 3.0 M, temperature ¼ 25 C).
G.S. Barney, P.B. Duval / Journal of Loss Prevention in the Process Industries 24 (2011) 76e84 Table 4 Results of k8 calculation using measured HAN decomposition temperatures of Gowland and Stedman (1981). [HAN] (M)
[HNO3] (M)
Temperature ( C)
Estimated k8 (M2 s1)
0.02 0.02 0.02 0.05 0.05 0.05 0.01 0.005 0.02 0.01 0.005 0.01 0.005
3 3.5 4 3 3.5 4 3 3 2.5 2.5 2.5 2 2
17 11 2 25 13 5 10 4 30 19 12 30 21
3.48eþ08 1.77e þ 08 1.25eþ08 4.13eþ08 2.66eþ08 1.60eþ08 3.56eþ08 3.46eþ08 5.56eþ08 6.31eþ08 5.96eþ08 1.28eþ09 1.19eþ09
concentrations and temperatures. A predictor equation for k8 as a function of nitric acid concentration, HAN concentration and temperature was obtained by fitting the data in Table 4 to a multivariate function and the resulting function is used in the HAN safety model. The equation used in the model is: þ k8 ¼ eð27:456þ10:635½HAN1:3686½H 0:013308T Þ
(17)
where T is the temperature in Kelvin. 4.2.4. Rate constant k9 The rate constant for Reaction (9) has not been measured. The reaction has been described in Gowland and Stedman (1981) and Pembridge and Stedman (1979) as being relatively fast compared to other reaction steps in the autocatalytic mechanism. The reaction occurs in two steps. The two elementary reactions are N2O4 þ HNO / HNO2 þ N2O3
(Reaction 20)
and N2O3 þ H2O / 2HNO2
(Reaction 21)
Both reactions are very fast but, according to Gowland and Stedman (1981), Reaction (9) is controlled by elementary Reaction (20). The activation energy calculated from transition state theory (TST) (Ashcraft et al., 2008) for the first reaction is 0.7 kcal mol1 (2.93 kJ mol1). A value for k9 was estimated at 25 C, 0.05 M HAN, and 3.0 M HNO3 using the model and setting the k8 value (k8 ¼ 8.64E8 M2 s1) for these same conditions by interpolating the values from the Pembridge and Stedman data (see previous section). The k9 value that leads to an unstable solution (see previous section for the criteria used for determining stability) for these conditions was k9 ¼ 1.5E9 M2 s1. This value, along with the TST activation energy was used to derive an expression for k9 as a function of temperature (in K). Solution stability is not nearly as sensitive to k9 as it is to k8. The Arrhenius expression used in the model is:
k9 ¼ 4:891 109 eð2930=8:314TÞ
81
published data were used for this comparison. Pembridge and Stedman (1979) give a curve showing the concentration of nitrous acid over time for a solution initially containing 0.005 M HAN and 5.0 M HNO3 at 25 C. Bennett et al. (1982) published a similar curve for a solution initially containing 0.003 M HAN and 3.0 M HNO3 at 25 C. The initial nitrous acid concentration for the Bennett, et al. curve was specified as <4E5 M. This is important because the initial nitrous acid concentration determines the induction period for the autocatalytic reaction (the lower the initial nitrous acid concentration, the longer will be the induction period). These curves are compared with the model output curves in Fig. 4. The induction period for both measured curves is about 100 s before the autocatalytic reaction accelerates. The model output curves reasonably reproduce this induction period. The nitrous acid yield predicted by the model is slightly lower than that measured by Pembridge and Stedman and slightly higher than that measured by Bennett et al. However, the model appears to reproduce measured HAN stability data reasonably well. The model is used to calculate boundary conditions for stability of solutions containing a range of HAN and nitric acid concentrations. Again, the criterion for assessing solution stability is that the HAN solution is considered unstable if the HNO2 concentration increases over a 1000 s time period (see the example in Fig. 3). Plots of nitric acid concentrations versus initiation temperatures for the autocatalytic reaction are linear for a given initial HAN concentration, as shown by Gowland and Stedman (1981). Their data are reproduced along with the model predictions for HAN solution initiation temperatures in Fig. 5. This figure shows both measured values and model results for initiation temperatures for nitric acid concentrations of 2e4 M and HAN concentrations of 0.005e0.05 M. There is good agreement between the measured and modeled values. For a given initial HAN concentration, the solutions are unstable above the line and stable below the line. For example, a 0.05 M HAN solution becomes unstable above about 25 C at 3.0 M HNO3 and above about 5 C at 4.0 M HNO3. Because of the relatively high concentrations of reduced plutonium [Pu(III)] in the aqueous stripping solution, the potential exists for reoxidation of the plutonium with nitric and nitrous acid. These reactions yield additional nitrous acid that can promote the autocatalytic decomposition of HAN and make the process solutions more unstable. Unfortunately, there are no direct experimental data available showing the effects of Pu(III) concentrations on HAN solution stability. This lack of direct experimental data requires that
(18)
where T is the temperature in Kelvin. 5. Results and discussion The validity of the model was confirmed by comparing the output of the model to published curves showing the concentration of nitrous acid over time for solutions of HAN in nitric acid solutions that undergo the autocatalytic HAN reaction. Several sources of
Fig. 4. Comparison of published data (Bennett et al., 1982; Pembridge & Stedman, 1979) on HNO2 concentration over time with model output curves during the HANenitric acid autocatalytic reaction.
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respectively at a constant ionic strength of 2.0 M. These rate constants are significantly lower than those measured by Dukes (1960). Newton (2002) calculated an activation energy of 55.6 kJ mol1 for the rate constants and this value was used to extend the rate constants to higher temperatures using the Arrhenius equation. The effect of ionic strength on the rate constants was not determined. The predictor equations used in the 0 model to calculate values of k13 and k13 are given below:
k13 ¼ 1:648 108 eð55;600=8:314TÞ
(19)
and
k013 ¼ 1:465 107 eð55;600=8:314TÞ
Fig. 5. Boundary conditions [experimental data from Gowland and Stedman (1981) with open symbols and calculated values using the model (closed symbols)] for the HAN autocatalytic reaction over a range of initial HAN concentrations [(A) ¼ 0.05 M, (B) ¼ 0.02 M, (C) ¼ 0.01 M, and (D) ¼ 0.005 M].
assumptions be made concerning the reaction mechanisms involved and that literature data on reaction rate constants be used. The two overall reactions involving plutonium species assumed to be significant in determining the stability of process solutions are Reactions (12) and (13). The kinetics of Reaction (12) has been determined by Barney (1976). The kinetics of Reaction (13) has been determined by Newton (2002), Koltunov and Marchenko (1973), Andreychuk, Frolov, Rotmanov, and Vasiliev (1990), and by Dukes (1960). If these reaction mechanisms are correct, the HANeHNO3ePu(III) solutions are inherently unstable because of the continuous reoxidation of Pu(III) with HNO3 and subsequent reduction of Pu(IV) with HAN. These reactions will gradually deplete the HAN by reaction with HNO2 and Pu(IV). At some point in the depletion of HAN, the autocatalytic HAN decomposition will be rapid. Thus, the stability of these solutions is not as well-defined as for HANeHNO3 and HANeHNO3eFe3þ solutions at expected concentrations of plutonium and iron in process solutions ([Pu]/ [Fe] > 150). If iron concentrations are higher, a gradual depletion of HAN will occur that will be similar to the effect of Pu(III). The most recent rate data on the reaction of Pu(III) with nitric acid and nitrous acid have been reported by Newton (2002). He justifies the empirical rate law found by Koltunov and Marchenko (1973) with 0.5 dependencies on [HNO2], [Hþ], and [NO 3 ] (see Equation (8) above) by proposing a mechanism having the rate determining step as a reaction of Pu(III) with NO2. Since NO2 is in rapid equilibrium with HNO2, Hþ, and NO 3 according to Reaction (5), the rate law is compatible with this mechanism. Using the Koltunov data, he calculated a value for the formation constant of 1 the PuNO2þ 3 complex (b) to be 0.123 M . Andreychuk, et al. (1990) have also measured rates for Pu(III) oxidation in nitric acid. They found that the rate of oxidation is controlled by chemical mechanisms rather than radiolysis at nitric acid concentrations greater than 0.55 M. The Pu(IV) concentration versus time curves cover a nitric acid concentration range of 0.12e2.90 M. A comparison of rates for reoxidation from the papers of Andreychuk et al. and Koltunov and Marchenko is very good, providing strong support for the Pu(III) reoxidation part of the model. Model calculations show that HAN solution stability is significantly affected by the rate constants k13 and k130 and by the formation constant, b. The values given for the rate constants at 16 C are 0.0149 M1.5 s1 and 0.00132 M2.5 s1 for k13 and k130 ,
(20)
where T is the temperature in Kelvin. Effects of temperature and ionic strength on b have not been accurately measured. A literature review of available data on formation constants for the PuNO2þ 3 complex did not find reliable measurements of this constant. By analogy, the formation constant complex ranges from about 1.8 to 1.6 at for the Am(III)NO2þ 3 25e30 C at an ionic strength of 1.0e2.01 M (Schulz, 1976). The b values for Pu(III) are likely near these values. Values of b for Pu(III) that are appropriate for the current model were obtained by fitting data obtained by Andreychuk et al. (1990) using the model. Their published rate curves show [Pu(IV)] versus time for a range of HNO3 concentrations, using an initial Pu(III) concentration of 0.010 M. They did not specify a temperature for the measurements, but previous papers indicated a temperature of 22 2 C. These rate curves are reproduced in Fig. 6 along with the model output for the same conditions (no HAN present). Values of b were adjusted to fit the measured data. No other parameters were adjusted. The fit shows that the model is reasonably accurate and gives confidence in the mechanism chosen and the rate laws used in the model. The b values are within the expected range and decrease linearly as the acid concentrations increase. The three data points were fitted to the equation,
h
b ¼ 6:044 1:4913 Hþ
i
(21)
Using the rate laws and constants described above for reactions of plutonium in the HANeHNO3ePu solution system, the model shows that plutonium has a significant effect on the stability of these solutions. A significant amount of nitrous acid is generated by plutonium(III) reoxidation. This nitrous acid lowers the stability of the HAN solutions by promoting the nitrous acid scavenging reaction and the autocatalytic HAN decomposition. An example of the
Fig. 6. Comparison of measured Pu(III) reoxidation curves from Andreychuk et al. (1990) with curves calculated using the model and fitted beta values. The initial [Pu (III)] was 0.10 M for the model calculation.
G.S. Barney, P.B. Duval / Journal of Loss Prevention in the Process Industries 24 (2011) 76e84
83
decomposition is accelerated by the presence of iron. Fig. 8 shows the model output for HAN concentration curves over time with different initial concentrations of iron. The initial conditions for these model runs were temperature ¼ 40 C, [NH3OHþ] ¼ 0.05 M, [Pu(IV)] ¼ 0.000001 M, [HNO3] ¼ 2.0 M, and [HNO2] ¼ 0.001 M. Iron definitely lowers the stability of HANenitric acid solutions at high iron concentrations. However, at the maximum expected iron concentration of about 0.001 M, HAN stability is only slightly lowered. It can be concluded that plutonium concentrations will have a much greater effect on lowering HANenitric acid solution stabilities than will iron in these solutions. 6. Conclusions
Fig. 7. Model output curves showing the effects of several initial Pu(III) concentrations on HAN concentration over time and showing the time required to reach the autocatalytic HAN reaction (vertical curves) (temperature ¼ 25 C and [HAN]0 ¼ 0.05 M, [HNO3] ¼ 2.0 M).
effects of a range of Pu(III) concentrations on HAN concentrations at 25 C over a 1E6 s time period is shown in Fig. 7. The initial conditions for this calculation are [NH3OHþ] ¼ 0.05 M, [Pu(IV)] ¼ 0.000001 M, [HNO3] ¼ 2.0 M, [HNO2] ¼ 0.001 M, and [Fe] ¼ 0 M. The curves show that increasing the plutonium concentration significantly lowers the stability of the solutions, resulting in the autocatalytic HAN reaction. The HAN concentration is gradually depleted while the concentrations of Pu(IV) and HNO2 increase over time. Plutonium concentrations normally encountered in solvent extraction purification processes can range up to about 0.2 M which will strongly affect the stability of HANenitric acid solutions. Dissolved iron will be present in nitric acid process solutions used in solvent extraction due to corrosion of the stainless steel equipment and piping in addition to iron impurities in the feed solutions. The highest concentration of dissolved iron in these solutions is estimated to be about 1E3 M. Iron can affect the autocatalytic HAN decomposition by generating or consuming NO2 or HNO2 in solution [see Reactions (17)e(19)]. Experimental measurements (Sugikawa et al., 2002) show that the
Fig. 8. Model output curves showing effects of several initial Fe(III) concentrations on HAN concentration over time and showing the time required to reach the autocatalytic HAN reaction (vertical curves) (temperature ¼ 40 C and [HAN]0 ¼ 0.05 M, [HNO3] ¼ 2.0 M).
A mathematical model has been developed to predict the safety of hydroxylamine use in solvent extraction purification of plutonium. The model allows prediction of the stability of solutions containing HAN in processing vessels under a range of conditions. Conditions for which the hazardous autocatalytic reaction of HAN with nitric acid can occur are predicted so that they can be avoided. The model calculates concentrations of important chemical species in solution over time for a given set of input parameters. It consists of a chemical kinetics model in which the reaction rates for 19 nitrogen, plutonium, and iron species are estimated. The stability of HAN in nitric acid process solutions is controlled by complex relationships between temperature and concentrations of HAN, nitric acid, plutonium and iron. Process parameters (flow rates, volumes, initial reagent concentrations and temperatures) can affect the values of each of these parameters. It is found that plutonium(III) in nitric acid solutions containing HAN significantly lowers the stability of the solutions toward the autocatalytic HAN reaction with nitric acid. Dissolved iron also catalyzes this reaction, but is not as significant as plutonium because of its expected low concentrations. High acid concentrations and temperatures also promote the autocatalytic reaction. For normal extraction process conditions, the model confirms the absence of significant risk, in accordance with extensive industrial experience gained in reprocessing plants in Europe and Japan. References Andreychuk, N. N., Frolov, A. A., Rotmanov, K. V., & Vasiliev, VYa (1990). Plutonium (III) oxidation under a-irradiation in nitric acid solutions. Journal of Radioanalytical and Nuclear Chemistry, 142, 427e432. Ashcraft, R. W., Raman, S., & Green, W. H. (2008). Predicted reaction rates of HxNyOz intermediates in the oxidation of hydroxylamine by aqueous nitric acid. Journal of Physical Chemistry A, 112, 7577e7593. Barney, G. S. (1976). A kinetic study of the reaction of plutonium(IV) with hydroxylamine. Journal of Inorganic and Nuclear Chemistry, 38, 1677e1681. Barney, G. S. (1971). The reaction of hydroxylamine with nitrous acid. ARH-SA-97. Richland, Washington: Atlantic Richfield Hanford Company. Bengtsson, G., Fronaeus, S., & Bengtsson-Kloo, L. (2002). The kinetics and mechanism of oxidation of hydroxylamine by iron(III). Journal of the Chemical Society Dalton Transactions, 2002, 2548e2552. Bennett, M. R., Brown, G. M., Maya, L., & Posey, F. A. (1982). Oxidation of hydroxylamine by nitrous and nitric acids. Inorganic Chemistry, 21, 2461e2468. Cisneros, L. O., Rodgers, W. J., & Mannan, M. S. (2002). Effect of air in the thermal decomposition of 50 mass% hydroxylamine/water. Journal of Hazardous Materials, A95, 13e25. Cisneros, L. O., Rodgers, W. J., & Mannan, M. S. (2004). Comparison of the thermal decomposition behavior for members of the hydroxylamine family. Thermochimica Acta, 414, 177e183. Cisneros, L. O., Wu, X., Rodgers, W. J., Mannan, M. S., Park, J., & North, S. W. (2003). Decomposition products of 50 mass% hydroxylamine/water under runaway reaction conditions. Transactions of the IChemE, 81, 121e124. DOE/EH-0555. (1998). Technical report on hydroxylamine nitrate. Washington, DC: U. S. Department of Energy. Dukes, E. K. (1960). Kinetics and mechanisms for the oxidation of trivalent plutonium by nitrous acid. Journal of the American Chemical Society, 82, 9. Epstein, I. R., Kustin, K., & Warshaw, L. J. (1980). A kinetics study of the oxidation of iron(II) by nitric acid. Journal of the American Chemical Society, 102(11), 3751e3758.
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Gowland, R. J., & Stedman, G. (1981). Kinetic and product studies on the decomposition of hydroxylamine in nitric acid. Journal of Inorganic and Nuclear Chemistry, 43(11), 2859e2862. Koltunov, V. S., & Marchenko, V. I. (1973). Reaction between plutonium(III) and nitrous acid. Radiokhimiya, 15, 777e781. Kumasaki, M. (2004). Calorimetric study on the decomposition of hydroxylamine in the presence of transition metals. Journal of Hazardous Materials, 115, 57e62. Kumasaki, M., Fujimoto, Y., & Ando, T. (2003). Calorimetric behaviors of hydroxylamine and its salts caused by Fe(III). Journal of Loss Prevention in the Process Industries, 16, 507e512. Liu, L., Papadaki, M., Pontiki, E., Stathi, P., Rodgers, W. J., & Mannan, M. S. (2009). Isothermal decomposition of hydroxylamine and hydroxylamine nitrate in aqueous solutions in the temperature range 80e160 C. Journal of Hazardous Materials, 165, 573e578. Liu, L., Wei, C., Guo, Y., Rodgers, W. J., & Mannan, M. S. (2009). Hydroxylamine nitrate self-catalytic study with adiabatic calorimetry. Journal of Hazardous Materials, 162, 1217e1222. Long, L. A. (2004). The explosion at concept sciences: hazards of hydroxylamine. Process Safety Progress, 23, 114e120. Newton, T. W. (2002). Redox reactions of plutonium ions in aqueous solutions. In D. Hoffman (Ed.), Advances in plutonium chemistry 1967e2000 (pp. 24e57). La Grange Park, Illinois: American Nuclear Society. Orban, M., & Epstein, I. R. (1982). Bistability in the oxidation of iron(II) by nitric acid. Journal of the American Chemical Society, 104, 5918e5922.
Pembridge, J. R., & Stedman, G. (1979). Kinetics, mechanism, and stoichiometry of the oxidation of hydroxylamine by nitric acid. Journal of the Chemical Society Dalton Transactions, 1979, 1657. Perrott, J. R., & Stedman, G. (1976). The kinetics of nitrite scavenging by hydrazine and hydrazoic acid at high acidities. Journal of Inorganic and Nuclear Chemistry, 39, 325e327. Pitzer, K. S. (1991). Activity coefficients in electrolyte solutions (2nd ed.). Boca Raton, Florida: CRC Press. Raman, S., Ashcraft, R. W., Vial, M., & Klasky, M. L. (2005). Oxidation of hydroxylamine by nitrous and nitric acids. Model development from first principle SCRF calculations. Journal of Physical Chemistry A, 109, 8526e8536. Schmid, G., & Bahr, G. (1964). Formation of HNO2 from NO and HNO3 at high acid strength. Zeitschrift für Physikalische Chemie, 41, 8e25. Schulz, W. W. (1976). The chemistry of americium. ERDA critical review series, report no. TID-26971. Schwartz, S. E., & White, W. H. (1983). Kinetics of reactive dissolution of nitrogen oxides into aqueous solutions. In J. O. Nriagu (Ed.), Advances in environmental science and technology, Vol. 12 (pp. 3e109). New York: John Wiley & Sons. Sugikawa, S., Umeda, M., Sekino, J., Matsuda, T., & Kodama, T. (2002). Evaluation of safety limit for handling hydroxylamine nitrate/nitric acid solution as a plutonium reductant. In JAERI conf 2002e004 (pp. 475e480). Tokai-mura, Naka-ken 319-1195, Japan: Japan Atomic Energy Research Institute. Yusaku, I., Masahide, W., Takashi, U., & Hiroshi, K. (2005). Cause investigation of explosive accident of hydroxylamine and risk evaluation of its water solution. Report of National Research Institute of Fire and Disaster, 99, 114e119.
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Journal of Loss Prevention in the Process Industries journal homepage: www.elsevier.com/locate/jlp
Model for predicting hydroxylamine nitrate stability in plutonium process solutions G. Scott Barney a, *, Paul B. Duval b a b
105903 N. Harrington Rd., West Richland, WA, USA Shaw AREVA MOX Services, LLC, Aiken, SC, USA
a r t i c l e i n f o
a b s t r a c t
Article history: Received 29 July 2010 Received in revised form 27 October 2010 Accepted 29 October 2010
A mathematical model is described that may be used to determine the safety of hydroxylamine nitrate (HAN) solutions used in solvent extraction purification of plutonium. The most significant hazard associated with hydroxylamine use in processing plutonium is its rapid, autocatalytic reaction with nitric acid which can result in an explosion or pressurization of process vessels with radiological consequences to humans. In addition, heat is produced by the reaction that could potentially ignite process solvents. The HAN decomposition reaction can occur only under specific process conditions (temperature; HAN, plutonium and nitric acid concentrations) and the model is used to identify these conditions so that they can be avoided. A kinetics model has been developed using all of the known significant reactions that could occur in process solutions containing HAN and nitric acid as well as plutonium and iron. The reaction kinetics data (rate laws, rate constants, activation energies) used in the model were obtained from chemical literature sources. The model shows that the autocatalytic HAN reaction with nitric acid is very rapid and is catalyzed by Pu(III) and Fe(II) in process solutions. High temperatures and nitric acid concentrations also promote the reaction. Ó 2010 Elsevier Ltd. All rights reserved.
Keywords: Hydroxylamine stability Plutonium kinetics Nitric acid Model
1. Introduction Solvent extraction processing in the nuclear industry often utilizes hydroxylamine nitrate (NH3OHNO3 or HAN) to separate plutonium from uranium and other impurities by stripping plutonium from an organic solvent containing tributyl phosphate. Aqueous HAN is used to reduce Pu(IV) in the organic phase to Pu(III) which is not very soluble in the organic phase. Industrial experience with HAN has been acquired in France, Japan, and the UK for more than 20 years without any reported incidents (DOE, 1998). However, the use of HAN in the plutonium purification process with nitric acid solutions introduces the possibility of an autocatalytic reaction that may produce unacceptable temperature and pressure increases to the system and consequently result in overpressurization and/or explosions within the process vessels. Several accidents have been reported in the United States (DOE, 1998) involving explosions or pressurization of process vessels by HAN reactions in the nuclear process industry. It is noted that all these accidents occurred in storage, chemical make-up vessels, or evaporation vessels and not in the extraction process. However it is
* Corresponding author. Tel.: þ1 509 967 2549. E-mail address: [email protected] (G.S. Barney). 0950-4230/$ e see front matter Ó 2010 Elsevier Ltd. All rights reserved. doi:10.1016/j.jlp.2010.10.004
prudent to determine the conditions for safe operations to prevent any incident in the extraction process. Recent research on the stability of HAN in aqueous solutions includes kinetic modeling of the oxidation of HAN by nitric acid using ab initio molecular orbital calculations (Ashcraft, Raman, & Green, 2008; Raman, Ashcraft, Vial, & Klasky, 2005), calorimetric and spectroscopic measurements (Kumasaki, 2004; Kumasaki, Fujimoto, & Ando, 2003; Liu, Papadaki, et al., 2009; Liu, Wei, Guo, Rodgers, & Mannan, 2009), and effects of metal ions on HAN stability (Sugikawa, Umeda, Sekino, Matsuda, & Kodama, 2002). A significant amount of work has also been performed on the stability of hydroxylamineewater solutions (Cisneros, Rodgers, & Mannan, 2002, 2004; Cisneros et al., 2003) prompted by two tragic industrial accidents involving hydroxylamine solutions (Long, 2004; Yusaku, Masahide, Takashi, & Hiroshi, 2005). None of these studies include the effects of dissolved plutonium on HAN stability. In addition, the nitric acid concentrations and temperatures were not in the range used in plutonium purification. A mathematical model is described in this paper that calculates the rates of the reactions involved in the oxidation of HAN under conditions that exist in a plutonium purification solvent extraction process. The process conditions required for the rapid, autocatalytic HAN reaction are identified by the model so that these conditions can be avoided. The model predicts the rate of the reactions
G.S. Barney, P.B. Duval / Journal of Loss Prevention in the Process Industries 24 (2011) 76e84
involved and resulting temperature changes and gas generation rates in the process solutions. 2. Chemical reactions 2.1. Autocatalytic reaction The chemical reactions involved in the autocatalytic oxidation of hydroxylamine nitrate (NH3OHNO3) with nitric acid have been proposed by several laboratories (Gowland & Stedman, 1981; Bennett, Brown, Maya, & Posey, 1982). Hydroxylammonium ion is the major hydroxylamine species in solutions with a pH lower than about 5.8 and will be the major species in the highly acidic process solutions. The overall autocatalytic reaction can be written as follows: NH3OHþ þ 0.5(x þ 1)HNO3 / xHNO2 þ 0.25(3 x)N2O þ Hþ þ (1.75 0.25x)H2O
(Reaction 1)
Nitrous acid production (Reaction 2)
Nitrous acid scavenging
NH3OHþ þ HNO2 / N2O þ Hþ þ 2H2O
Gowland and Stedman (1981) and Pembridge and Stedman (1979) have proposed a mechanism for the autocatalytic reaction that includes Reactions (4), (5), (8), and (9). They state that the autocatalytic reaction is limited by Reaction (5), which is relatively slow. Nitric acid solutions of HAN are stable as long as the nitrous acid scavenging reaction (3) is dominant over the autocatalytic reaction, which produces nitrous acid. Reaction (7) accounts for the alternate decomposition of nitrous acid which yields nitric oxide. The reverse reaction can occur to produce nitrous acid and nitric dioxide [Reaction (6)]. Reactions of hydrazinium ion (N2Hþ 5 ) and hydrazoic acid (HN3) with nitrous acid [Reactions (10) and (11)] are important because hydrazine (N2H4) is used as a nitrous acid suppression agent in solvent extraction processes using HAN. 2.3. Reactions of plutonium species in solution
The reaction is strongly exothermic and produces a gaseous reaction product (N2O) that can pressurize process equipment. The key to initiation of the autocatalytic reaction of HAN is the production of nitrous acid (HNO2). Nitrous acid is produced and consumed by NH3OHþ by the following overall reactions:
NH3OHþ þ 2HNO3 / 3HNO2 þ Hþ þ H2O
77
(Reaction 3)
If the nitrous acid concentration reaches a critical threshold, Reaction (2) dominates and nitrous acid is rapidly produced. Any reaction that can increase the nitrous acid concentration can lower the decomposition temperature of the solutions. Dissolved iron has been shown to lower the decomposition temperatures of HAN solutions (Sugikawa et al., 2002) and there are reactions of iron in nitric acid solution that will produce nitrous acid. The detailed reactions that were considered in this paper are listed below. 2.2. Reactions of nitrogen species in solution NH3OHþ þ 2HNO3 / 3HNO2 þ Hþ þ H2O
(Reaction 2)
HNO2 þ NH3OHþ / N2O þ 2H2O þ Hþ
(Reaction 3)
þ 2NO2 þ H2O / HNO2 þ NO 3 þH
(Reaction 4)
þ HNO2 þ NO 3 þ H / 2NO2 þ H2O
(Reaction 5)
þ NO þ NO 3 þ H / HNO2 þ NO2
(Reaction 6)
þ NO2 þ HNO2 / NO þ NO 3 þH
(Reaction 7)
2NO2 þ NH3OHþ / HNO þ 2HNO2 þ Hþ
(Reaction 8)
2NO2 þ HNO þ H2O / 3HNO2
(Reaction 9)
þ N2Hþ 5 þ HNO2 / HN3 þ 2H2O þ H
(Reaction 10)
HN3 þ HNO2 / N2 þ N2O þ H2O
(Reaction 11)
Dissolved plutonium and iron can participate in reactions with HAN. For most nuclear materials that are processed, the concentrations of plutonium in the solutions downstream from the plutonium stripping column will be relatively high compared to iron. Plutonium will likely catalyze the autocatalytic HAN reaction with nitric acid by reactions analogous to the iron reactions that destabilize HAN solutions. Plutonium in these solutions will be reduced species of Pu(III) which can react with nitric acid to produce nitrous acid. Any reaction mechanism that can produce nitrous acid will potentially destabilize the solutions with respect to the autocatalytic HAN reaction. The reactions likely to be important are Reactions (12) and (13) below. Nitrate complex formation as given by Reactions (14) and (15) are important in the model calculations, but are very rapid and reversible so they are considered equilibrium reactions. 2Pu4þ þ 2NH3OHþ / 2Pu3þ þ N2 þ 4Hþ þ 2H2O
(Reaction 12)
4þ þ HNO2 þ H2O 2Pu3þ þ 3Hþ þ NO 3 / 2Pu
(Reaction 13)
3þ Pu4þ þ NO 3 / PuNO3
(Reaction 14)
2þ Pu3þ þ NO 3 / PuNO3
(Reaction 15)
The kinetics of these reactions has been studied and rate laws and rate constants have been proposed and measured. These rate data have been included in the HAN stability model. 2.4. Reactions of iron species in solution Iron(III) can be reduced by HAN to Fe(II) (Bengtsson, Fronaeus, & Bengtsson-Kloo, 2002) which can then react with nitric acid to produce nitrous acid (Epstein, Kustin, & Warshaw, 1980). Several additional reactions of Fe3þ and Fe2þ can occur with NO2, HNO2, and NO that can exist in the process solutions. They are listed below. Fe3þ þ NH3OHþ / Fe2þ þ ½N2 þ 2Hþ þ H2O
(Reaction 16)
Fe2þ þ NO2 þ Hþ / Fe3þ þ HNO2
(Reaction 17)
Fe2þ þ HNO2 þ Hþ / Fe3þ þ NO þ H2O
(Reaction 18)
Fe3þ þ NO þ H2O / Fe2þ þ HNO2 þ Hþ
(Reaction 19)
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3. Rate equations used in model
4.1. Rate constants with limited information
The model calculates concentrations of each significant species involved in the autocatalytic reaction over time for a given set of initial conditions (temperature, ionic strength, initial concentrations of HAN, HNO3, HNO2, Fe3þ, Pu3þ, Pu4þ, etc.) The significant species, based on the proposed reaction mechanism, are Fe3þ/Fe2þ, Pu4þ/ þ Pu3þ, NH3OHþ, NO2, HNO2, NO, HNO, Hþ, NO 3 , N2, N2O, N2H5 , and HN3. The differential equations describing the changes in concentrations over time are given in Tables 1 and 2. The second order equation, d[NH3OHþ]/dt ¼ k3[HNO2][NH3OHþ], is used for Reaction (3) in these combined equations and the acid dependency of k3 (the second order rate constant) is calculated using published data. The model receives the input data for initial conditions of the HANeHNO3eH2N4eFeePu solutions and then simultaneously solves this set of differential equations for concentrations of the significant species and solution temperature. IGOR PRO version 6.0.6 (Wavemetrics, Inc.) programming language was utilized to solve the equations.
Some of the rate constants used in the model are limited to a single value at a specific temperature and solution ionic strength. For others, activation energies are available. Table 3 lists values for the rate constants where limited information is available.
4. Rate constants Rate constant data were found for each of the 19 reactions listed above except for Reactions (8) and (9). Rate constants for Reactions (8) and (9) were estimated using methods described below. Most rate constant data were obtained at a single temperature or at nitric acid solution concentrations lower than those expected in the solvent extraction process. Where necessary, rate constant values were extrapolated to process-relevant conditions using standard techniques for each of the significant reactions. A relatively small number of reactions were found to significantly affect the autocatalytic HAN reaction. In the absence of iron and plutonium, the significant reactions are (3), (4), (5), and (8)e (11). In the presence of plutonium, Reactions (12) and (13) were significant. In the presence of iron, Reactions (16)e(18) are important. Each of these reactions and the rate constants for them are discussed in the following sections.
4.2. Rate constants with known dependencies on solution composition and temperature 4.2.1. Rate constant k3 Reaction (3) is key in determining the stability of HANeHNO3 solutions as discussed above. Values for the rate constant, k3, have been published by a number of workers (Barney, 1971; Bennett et al., 1982; Pembridge & Stedman, 1979). Measurements have been made in both nitric acid and perchloric acid solutions at ionic strengths of 0.1e6.0 M and temperatures of 0e50 C. The mechanism for this reaction appears to change above approximately 2.0 M ionic strength. At lower ionic strengths, the second order rate constant increases with increasing acid concentration as seen in Fig. 1. At higher ionic strength a maximum is reached and then the rate constant decreases with increasing acid concentration. This strongly affects the model results, because nitrous acid scavenging is not as effective at higher nitric acid concentrations and causes the solutions to become more unstable. The data from the higher ionic strengths were used to model effects of nitric acid concentration, ionic strength, and temperature on k3. Activation energies for Reaction (3) have been published in Gowland and Stedman (1981), Bennett et al. (1982) and Barney (1971). A value of 41.5 kJ mol1 is used in the model to predict the dependency of k3 on temperature. The fitted equation used to predict k3 in the model is thus,
h i k3 ¼ 106 eð41;500=ð8:314*TÞÞ 234:08 þ 63:067m 85:702 Hþ h i h i2 ð12Þ 9:9333m2 þ 3:8233m Hþ þ 5:5347 Hþ
Table 1 Rate equations used in the model for calculation of HAN, HNO2, NO2, NO, HNO, N2Hþ 5 , and HN3 concentrations.
i h d NH3 OHþ dt
h ¼ k16
i2 h i2 h i2 h i h i Fe3þ NH3 OHþ ½PuðIVÞ2 NH3 OHþ 2 þ þ þ k ½HNO NH OH OH þk þ k NH ½NO h i2 h i4 h i4 2 3 2 3 8 3 2 12 Fe2þ Hþ Kd þ NO ½PuðIIIÞ2 Hþ 3
! i h ½NO ½HNO h 3þ i þ k3 ½HNO2 NH3 OHþ Fe k19 h i þ k’19 ½NO þ k}19 h i2 þ þ H H ih i h i h i2 h i h h i 2 þ þ þ þ þk7 ½HNO2 ½NO2 H 2k8 NH3 OH ½NO2 2 3k9 ½NO2 2 ½HNO 0:5k4 ½NO2 þk5 ½HNO2 NO3 H k6 NO3 ½NO H h i h i0:5 h i0:5 ih i h h i k13 þ k’13 Hþ ½PuðIIIÞ½HNO2 0:5 Hþ NO 3 h i Hþ þ k11 ½HNO2 ½HN3 Hþ 0:5 þ k10 ½HNO2 N2 Hþ 5 1 þ b NO3
h h i h i i d½HNO2 ½HNO2 ¼ k17 Fe2þ ½NO2 þ k18 þ k’18 Hþ þ k}18 Fe2þ ½HNO2 dt ½NO
h i ih i h i h i2 h i h h i d½NO2 þ ¼ k4 Fe2þ ½NO2 þ 2k4 ½NO2 2 2k5 ½HNO2 NO Hþ k6 NO þk7 ½HNO2 ½NO2 Hþ þ 2k8 NH3 OHþ ½NO2 2 þ2k9 ½NO2 2 ½HNO 3 3 ½NO H dt ! h h i i h i h i2 h i d½NO ½HNO2 ½NO ½HNO h 3þ i þ ¼ k18 þ k’18 Hþ þ k}18 Fe2þ ½HNO2 þ k19 h i þ k’19 ½NO þ k}19 h i2 þ k6 NO Fe k7 ½HNO2 ½NO2 Hþ 3 ½NO H þ þ dt ½NO H H
h i d½HNO ¼ k8 NH3 OHþ ½NO2 2 þk9 ½NO2 2 ½HNO dt h i ih i h d N2 Hþ 5 Hþ ¼ k10 ½HNO2 N2 Hþ 5 dt
(1)
ih i h h i d½HN3 ¼ k10 ½HNO2 N2 Hþ Hþ þ k11 ½HNO2 ½HN3 Hþ 5 dt
(2)
(3)
(4)
(5)
(6) (7)
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79
Table 2 Rate equations used in the model to calculate concentrations of Pu(IV), Fe(III), Hþ, and NO3.
d½PuðIVÞ ¼ k12 dt
i h d Fe3þ dt h i d Hþ dt
h ¼ k16
(8)
i2 h i2 h h i h i i Fe3þ NH3 OHþ ½HNO2 Fe2þ ½HNO2 þ k17 Fe2þ ½NO2 k18 þ k’18 Hþ þ k}18 h i2 h i4 ½NO 2þ þ Fe H
! ½NO ½HNO h 3þ i Fe k19 h i þ k’19 ½NO þ k}19 h i2 Hþ Hþ
i2 h i2 ! h h i h i i Fe3þ NH3 OHþ ½HNO2 ½NO ½HNO h 3þ i Fe2þ ½HNO2 k19 h i þ k’19 ½NO þ k}19 h i2 Fe þ k17 Fe2þ ½NO2 þ k18 þ k’18 Hþ þ k}18 ¼ 2k16 h i2 h i4 ½NO Hþ Hþ Fe2þ Hþ h h h i i ih i h i h i2 h i þ k7 ½HNO2 ½NO2 Hþ k8 NH3 OHþ k3 ½HNO2 NH3 OHþ 0:5k4 ½NO2 2 þk5 ½HNO2 NO Hþ þ k6 NO 3 3 ½NO H h i2 h i h i0:5 h i0:5 ih i h ½PuðIVÞ2 NH3 OHþ k13 þ k’13 Hþ ½PuðIIIÞ½HNO2 0:5 Hþ NO 3 2 h i Hþ k10 ½HNO2 N2 Hþ ½NO2 2k12 h i4 2 þ 1:5 5 2 þ b 1 þ NO Kd þ NO3 ½PuðIIIÞ H 3
h i d NO 3 dt
h i2 h i0:5 h i0:5 k13 þ k’13 ½Hþ ½PuðIIIÞ½HNO2 0:5 Hþ NO ½PuðIVÞ2 NH3 OHþ 3 h i h i4 2 1 þ b NO Kd þ NO ½PuðIIIÞ2 Hþ 3 3
(9)
h
2
¼ 0:5k4 ½NO2 þk5 ½HNO2
h
NO 3
ih
þ
H
i
þ
h
k6 NO 3
i
h i h i2 ½NO Hþ k7 ½HNO2 ½NO2 Hþ þ 0:5
where m is the molar ionic strength and T is the temperature in Kelvin. 4.2.2. Rate constant k5 Reaction (5) is the rate-limiting step for the autocatalytic reaction of HAN with HNO3 according to Gowland and Stedman (1981). It is very important to know the value of k5 accurately to predict the conditions required for the autocatalytic reaction. The k5 values over a range of solution compositions and temperatures must be calculated. The solution concentrations of interest are 0e0.5 M for HAN and 0.5e2 M for HNO3. The temperature range is from 25 C to 60 C. The reaction rate can be expressed in term of solution activities as follows:
d½HNO2 ¼ k5 aHNO2 aNO3 aHþ dt ih i h Hþ ¼ k5 fHNO2 fNO3 fHþ ½HNO2 NO 3
(13)
where a is the activity and f is the molar activity coefficient for a given component of the solution. Molar activity coefficients were
h i0:5 h i0:5 k13 þ k’13 ½Hþ ½PuðIIIÞ½HNO2 0:5 Hþ NO 3 h i 1 þ b NO 3
(10)
(11)
calculated for the range of solution compositions and temperatures required. Schmid and Bahr (1964) have measured k5 values over HNO3 concentrations of 0.3e5.0 M and temperatures of 0e20 C. Their data are used to calculate k5 values for different temperatures at zero ionic strength and to determine the activation energy for this reaction (66.8 kJ mol1). Molar activity coefficients for Hþ and NO 3 are calculated in the model for each time step using the Pitzer (1991, chap. 3) method. 4.2.3. Rate constant k8 Reaction (8) is very important in determining the rate of the autocatalytic reaction of HAN with nitric acid. The applicable rate law is
h i d½NO2 ¼ k8 NH3 OHþ ½NO2 2 dt
(14)
There are no published values of the rate constant for Reaction (8). It can be estimated from an equation derived by Gowland and Stedman (1981) that describes the rate of nitrous acid formation
Table 3 Rate constants used in the model for which only temperature and ionic strength effects are known. Rate constant
Value
Temperature ( C)
Ionic strength (M)
Activation energy (kJ mol1)
Reference
k4 k6 k7 k10 k11 k12 k13 k130 k16 k17 k18 k180 k1800 k19 k190 k1900
7.0E7 M1 s1 5E3 M1 s1 9.3 M2 s1 2.339E3 M2 s1 1.09E3 M2 s1 2.9E2 M5 s1 1.49E2 M1.5 s1 1.32E3 M2.5 s1 1.85E9 M3 s1 3.1E4 M1 s1 7.8E3 M1 s1 2.3E1 M2 s1 7.6E3 M1 s1 5.6E4 s1 1.6E2 M1 s1 5.4E4 s1
25 25 25 25 25 30 16 16 25 25 25 25 25 25 25 25
0.014e0.02 1.5 1.5 2.0 2.0 2.3e2.7 2.0 2.0 1.0 1.5 1.5 1.5 1.5 1.5 1.5 1.5
3.77 N/Aa N/A 56.5 54.5 130 55.8 55.8 133 N/A N/A N/A N/A N/A N/A N/A
Schwarz & White, 1983 Orban & Epstein, 1982 Orban & Epstein, 1982 Perrott & Stedman, 1976 Perrott & Stedman, 1976 Barney, 1976 Newton, 2002 Newton, 2002 Bengtsson et al., 2002 Orban & Epstein, 1982 Orban & Epstein, 1982 Orban & Epstein, 1982 Orban & Epstein, 1982 Orban & Epstein, 1982 Orban & Epstein, 1982 Orban & Epstein, 1982
a
Not available in the reference.
80
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Fig. 1. Second order rate constants for Reaction (3) versus acid concentration for different ionic strengths, temperatures, and acidesalt solutions [Bennett et al. (1982): 25 C in HNO3eLiNO3 mixtures and Barney (1971): 30 C in HNO3eNaNO3 mixtures].
Fig. 2. Comparison of calculated values of k8 from fitting the model to experimental decomposition data and from the Pembridge and Stedman (1979) measurements of HNO2 generation.
and disappearance during the autocatalytic reaction based on rate laws for Reactions (3)e(5) and (8). The equation is
solution was considered unstable if the HNO2 concentration increased over a 1000 s time period. An example of how this criterion was used is shown in Fig. 3 where HNO2 concentrations are plotted over time for different values of k8. The solution stability is quite sensitive to k8 values and these curves show that there is a definite boundary for stability at 25 C for a solution containing 0.05 M HAN and 3.0 M HNO3. The curves show that a k8 value of 4.12E8 results in a stable solution while a value of 4.13E8 gives an unstable solution. A maximum error of several degrees for the decomposition temperature can be expected when using this criterion to adjust the k8 values. Model results show that once the concentration of HNO2 begins to increase, the solution will become unstable. The experimental HAN decomposition temperature data and results of these calculations are shown in Table 4. This approach gave k8 values that are quite close to the values in Fig. 2. Fig. 2 shows a comparison of k8 values calculated from a fit to the experimental decomposition data with those of Pembridge and Stedman over a range of nitric acid concentrations. The two different sources of k8 values are in reasonably good agreement, although the slopes of the lines are slightly different. The decomposition data yield a fitted equation that predicts k8 values over a wide range of acid and HAN
i h ih i h ½HNO2 NH3 OHþ 3k5 Hþ NO 3 d½HNO2 i h ¼ k4 dt þ 2 NH3 OHþ k8 i h i h k3 ½HNO2 NH3 OHþ ¼ kn NH3 OHþ ½HNO2 ð15Þ where kn is a measured rate constant for the formation of HNO2 and steady state approximations, d[NO2]/dt ¼ 0 and d[HNO]/dt ¼ 0 are assumed. This equation can be solved for k4/k8 (and thus k8) by using the published rate data (kn values) for HNO2 generation from HAN oxidation with HNO3 given by Pembridge and Stedman (1979). Values of k3, k4, and k5 that are required for this calculation are estimated by methods described in this paper. The equation solved for k4/k8 is
h ih i i h 3k5 Hþ NO 3 k4 ¼ 2 NH3 OHþ kn þ k3 k8
(16)
The results for calculation of k8 for 2.68e7.86 M HNO3 and 0.002e0.01 M HAN are shown in Fig. 2 for 25 C. The activation energy for this reaction is not known so the effect of temperature cannot be estimated using the Pembridge and Stedman data. No estimated k8 values are available for solutions with low acid concentrations. The HAN model can be applied at acid concentrations lower than 2.86 M by extrapolating the k8 data calculated from the 25 C Pembridge and Stedman published data on HNO2 generation rates given in Fig. 2. Values of k8 calculated in this manner are plotted in Fig. 2 against HNO3 concentration. The concentration of HAN in these experiments seems to have little effect on the k8 values. Low HAN concentrations are used because the reaction is exothermic and significant heating of the solutions must be avoided to obtain accurate data. Fig. 2 shows that the fitted data for a logarithmic plot of k8 are linear and are likely to be reasonably accurate in the extrapolation. Since no temperature data other than 25 C are available for this reaction, k8 values at other temperatures were estimated by fitting published experimental data (Gowland & Stedman, 1981) with the model where HAN decomposition temperatures are measured over a range of acid and HAN concentrations. Calculated values of k3, k4, and k5 as given above are used in these calculations. The HAN
Fig. 3. Example of a stability boundary calculation for the HANeHNO3 autocatalytic reaction ([HAN]0 ¼ 0.05 M, [HNO3] ¼ 3.0 M, temperature ¼ 25 C).
G.S. Barney, P.B. Duval / Journal of Loss Prevention in the Process Industries 24 (2011) 76e84 Table 4 Results of k8 calculation using measured HAN decomposition temperatures of Gowland and Stedman (1981). [HAN] (M)
[HNO3] (M)
Temperature ( C)
Estimated k8 (M2 s1)
0.02 0.02 0.02 0.05 0.05 0.05 0.01 0.005 0.02 0.01 0.005 0.01 0.005
3 3.5 4 3 3.5 4 3 3 2.5 2.5 2.5 2 2
17 11 2 25 13 5 10 4 30 19 12 30 21
3.48eþ08 1.77e þ 08 1.25eþ08 4.13eþ08 2.66eþ08 1.60eþ08 3.56eþ08 3.46eþ08 5.56eþ08 6.31eþ08 5.96eþ08 1.28eþ09 1.19eþ09
concentrations and temperatures. A predictor equation for k8 as a function of nitric acid concentration, HAN concentration and temperature was obtained by fitting the data in Table 4 to a multivariate function and the resulting function is used in the HAN safety model. The equation used in the model is: þ k8 ¼ eð27:456þ10:635½HAN1:3686½H 0:013308T Þ
(17)
where T is the temperature in Kelvin. 4.2.4. Rate constant k9 The rate constant for Reaction (9) has not been measured. The reaction has been described in Gowland and Stedman (1981) and Pembridge and Stedman (1979) as being relatively fast compared to other reaction steps in the autocatalytic mechanism. The reaction occurs in two steps. The two elementary reactions are N2O4 þ HNO / HNO2 þ N2O3
(Reaction 20)
and N2O3 þ H2O / 2HNO2
(Reaction 21)
Both reactions are very fast but, according to Gowland and Stedman (1981), Reaction (9) is controlled by elementary Reaction (20). The activation energy calculated from transition state theory (TST) (Ashcraft et al., 2008) for the first reaction is 0.7 kcal mol1 (2.93 kJ mol1). A value for k9 was estimated at 25 C, 0.05 M HAN, and 3.0 M HNO3 using the model and setting the k8 value (k8 ¼ 8.64E8 M2 s1) for these same conditions by interpolating the values from the Pembridge and Stedman data (see previous section). The k9 value that leads to an unstable solution (see previous section for the criteria used for determining stability) for these conditions was k9 ¼ 1.5E9 M2 s1. This value, along with the TST activation energy was used to derive an expression for k9 as a function of temperature (in K). Solution stability is not nearly as sensitive to k9 as it is to k8. The Arrhenius expression used in the model is:
k9 ¼ 4:891 109 eð2930=8:314TÞ
81
published data were used for this comparison. Pembridge and Stedman (1979) give a curve showing the concentration of nitrous acid over time for a solution initially containing 0.005 M HAN and 5.0 M HNO3 at 25 C. Bennett et al. (1982) published a similar curve for a solution initially containing 0.003 M HAN and 3.0 M HNO3 at 25 C. The initial nitrous acid concentration for the Bennett, et al. curve was specified as <4E5 M. This is important because the initial nitrous acid concentration determines the induction period for the autocatalytic reaction (the lower the initial nitrous acid concentration, the longer will be the induction period). These curves are compared with the model output curves in Fig. 4. The induction period for both measured curves is about 100 s before the autocatalytic reaction accelerates. The model output curves reasonably reproduce this induction period. The nitrous acid yield predicted by the model is slightly lower than that measured by Pembridge and Stedman and slightly higher than that measured by Bennett et al. However, the model appears to reproduce measured HAN stability data reasonably well. The model is used to calculate boundary conditions for stability of solutions containing a range of HAN and nitric acid concentrations. Again, the criterion for assessing solution stability is that the HAN solution is considered unstable if the HNO2 concentration increases over a 1000 s time period (see the example in Fig. 3). Plots of nitric acid concentrations versus initiation temperatures for the autocatalytic reaction are linear for a given initial HAN concentration, as shown by Gowland and Stedman (1981). Their data are reproduced along with the model predictions for HAN solution initiation temperatures in Fig. 5. This figure shows both measured values and model results for initiation temperatures for nitric acid concentrations of 2e4 M and HAN concentrations of 0.005e0.05 M. There is good agreement between the measured and modeled values. For a given initial HAN concentration, the solutions are unstable above the line and stable below the line. For example, a 0.05 M HAN solution becomes unstable above about 25 C at 3.0 M HNO3 and above about 5 C at 4.0 M HNO3. Because of the relatively high concentrations of reduced plutonium [Pu(III)] in the aqueous stripping solution, the potential exists for reoxidation of the plutonium with nitric and nitrous acid. These reactions yield additional nitrous acid that can promote the autocatalytic decomposition of HAN and make the process solutions more unstable. Unfortunately, there are no direct experimental data available showing the effects of Pu(III) concentrations on HAN solution stability. This lack of direct experimental data requires that
(18)
where T is the temperature in Kelvin. 5. Results and discussion The validity of the model was confirmed by comparing the output of the model to published curves showing the concentration of nitrous acid over time for solutions of HAN in nitric acid solutions that undergo the autocatalytic HAN reaction. Several sources of
Fig. 4. Comparison of published data (Bennett et al., 1982; Pembridge & Stedman, 1979) on HNO2 concentration over time with model output curves during the HANenitric acid autocatalytic reaction.
82
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respectively at a constant ionic strength of 2.0 M. These rate constants are significantly lower than those measured by Dukes (1960). Newton (2002) calculated an activation energy of 55.6 kJ mol1 for the rate constants and this value was used to extend the rate constants to higher temperatures using the Arrhenius equation. The effect of ionic strength on the rate constants was not determined. The predictor equations used in the 0 model to calculate values of k13 and k13 are given below:
k13 ¼ 1:648 108 eð55;600=8:314TÞ
(19)
and
k013 ¼ 1:465 107 eð55;600=8:314TÞ
Fig. 5. Boundary conditions [experimental data from Gowland and Stedman (1981) with open symbols and calculated values using the model (closed symbols)] for the HAN autocatalytic reaction over a range of initial HAN concentrations [(A) ¼ 0.05 M, (B) ¼ 0.02 M, (C) ¼ 0.01 M, and (D) ¼ 0.005 M].
assumptions be made concerning the reaction mechanisms involved and that literature data on reaction rate constants be used. The two overall reactions involving plutonium species assumed to be significant in determining the stability of process solutions are Reactions (12) and (13). The kinetics of Reaction (12) has been determined by Barney (1976). The kinetics of Reaction (13) has been determined by Newton (2002), Koltunov and Marchenko (1973), Andreychuk, Frolov, Rotmanov, and Vasiliev (1990), and by Dukes (1960). If these reaction mechanisms are correct, the HANeHNO3ePu(III) solutions are inherently unstable because of the continuous reoxidation of Pu(III) with HNO3 and subsequent reduction of Pu(IV) with HAN. These reactions will gradually deplete the HAN by reaction with HNO2 and Pu(IV). At some point in the depletion of HAN, the autocatalytic HAN decomposition will be rapid. Thus, the stability of these solutions is not as well-defined as for HANeHNO3 and HANeHNO3eFe3þ solutions at expected concentrations of plutonium and iron in process solutions ([Pu]/ [Fe] > 150). If iron concentrations are higher, a gradual depletion of HAN will occur that will be similar to the effect of Pu(III). The most recent rate data on the reaction of Pu(III) with nitric acid and nitrous acid have been reported by Newton (2002). He justifies the empirical rate law found by Koltunov and Marchenko (1973) with 0.5 dependencies on [HNO2], [Hþ], and [NO 3 ] (see Equation (8) above) by proposing a mechanism having the rate determining step as a reaction of Pu(III) with NO2. Since NO2 is in rapid equilibrium with HNO2, Hþ, and NO 3 according to Reaction (5), the rate law is compatible with this mechanism. Using the Koltunov data, he calculated a value for the formation constant of 1 the PuNO2þ 3 complex (b) to be 0.123 M . Andreychuk, et al. (1990) have also measured rates for Pu(III) oxidation in nitric acid. They found that the rate of oxidation is controlled by chemical mechanisms rather than radiolysis at nitric acid concentrations greater than 0.55 M. The Pu(IV) concentration versus time curves cover a nitric acid concentration range of 0.12e2.90 M. A comparison of rates for reoxidation from the papers of Andreychuk et al. and Koltunov and Marchenko is very good, providing strong support for the Pu(III) reoxidation part of the model. Model calculations show that HAN solution stability is significantly affected by the rate constants k13 and k130 and by the formation constant, b. The values given for the rate constants at 16 C are 0.0149 M1.5 s1 and 0.00132 M2.5 s1 for k13 and k130 ,
(20)
where T is the temperature in Kelvin. Effects of temperature and ionic strength on b have not been accurately measured. A literature review of available data on formation constants for the PuNO2þ 3 complex did not find reliable measurements of this constant. By analogy, the formation constant complex ranges from about 1.8 to 1.6 at for the Am(III)NO2þ 3 25e30 C at an ionic strength of 1.0e2.01 M (Schulz, 1976). The b values for Pu(III) are likely near these values. Values of b for Pu(III) that are appropriate for the current model were obtained by fitting data obtained by Andreychuk et al. (1990) using the model. Their published rate curves show [Pu(IV)] versus time for a range of HNO3 concentrations, using an initial Pu(III) concentration of 0.010 M. They did not specify a temperature for the measurements, but previous papers indicated a temperature of 22 2 C. These rate curves are reproduced in Fig. 6 along with the model output for the same conditions (no HAN present). Values of b were adjusted to fit the measured data. No other parameters were adjusted. The fit shows that the model is reasonably accurate and gives confidence in the mechanism chosen and the rate laws used in the model. The b values are within the expected range and decrease linearly as the acid concentrations increase. The three data points were fitted to the equation,
h
b ¼ 6:044 1:4913 Hþ
i
(21)
Using the rate laws and constants described above for reactions of plutonium in the HANeHNO3ePu solution system, the model shows that plutonium has a significant effect on the stability of these solutions. A significant amount of nitrous acid is generated by plutonium(III) reoxidation. This nitrous acid lowers the stability of the HAN solutions by promoting the nitrous acid scavenging reaction and the autocatalytic HAN decomposition. An example of the
Fig. 6. Comparison of measured Pu(III) reoxidation curves from Andreychuk et al. (1990) with curves calculated using the model and fitted beta values. The initial [Pu (III)] was 0.10 M for the model calculation.
G.S. Barney, P.B. Duval / Journal of Loss Prevention in the Process Industries 24 (2011) 76e84
83
decomposition is accelerated by the presence of iron. Fig. 8 shows the model output for HAN concentration curves over time with different initial concentrations of iron. The initial conditions for these model runs were temperature ¼ 40 C, [NH3OHþ] ¼ 0.05 M, [Pu(IV)] ¼ 0.000001 M, [HNO3] ¼ 2.0 M, and [HNO2] ¼ 0.001 M. Iron definitely lowers the stability of HANenitric acid solutions at high iron concentrations. However, at the maximum expected iron concentration of about 0.001 M, HAN stability is only slightly lowered. It can be concluded that plutonium concentrations will have a much greater effect on lowering HANenitric acid solution stabilities than will iron in these solutions. 6. Conclusions
Fig. 7. Model output curves showing the effects of several initial Pu(III) concentrations on HAN concentration over time and showing the time required to reach the autocatalytic HAN reaction (vertical curves) (temperature ¼ 25 C and [HAN]0 ¼ 0.05 M, [HNO3] ¼ 2.0 M).
effects of a range of Pu(III) concentrations on HAN concentrations at 25 C over a 1E6 s time period is shown in Fig. 7. The initial conditions for this calculation are [NH3OHþ] ¼ 0.05 M, [Pu(IV)] ¼ 0.000001 M, [HNO3] ¼ 2.0 M, [HNO2] ¼ 0.001 M, and [Fe] ¼ 0 M. The curves show that increasing the plutonium concentration significantly lowers the stability of the solutions, resulting in the autocatalytic HAN reaction. The HAN concentration is gradually depleted while the concentrations of Pu(IV) and HNO2 increase over time. Plutonium concentrations normally encountered in solvent extraction purification processes can range up to about 0.2 M which will strongly affect the stability of HANenitric acid solutions. Dissolved iron will be present in nitric acid process solutions used in solvent extraction due to corrosion of the stainless steel equipment and piping in addition to iron impurities in the feed solutions. The highest concentration of dissolved iron in these solutions is estimated to be about 1E3 M. Iron can affect the autocatalytic HAN decomposition by generating or consuming NO2 or HNO2 in solution [see Reactions (17)e(19)]. Experimental measurements (Sugikawa et al., 2002) show that the
Fig. 8. Model output curves showing effects of several initial Fe(III) concentrations on HAN concentration over time and showing the time required to reach the autocatalytic HAN reaction (vertical curves) (temperature ¼ 40 C and [HAN]0 ¼ 0.05 M, [HNO3] ¼ 2.0 M).
A mathematical model has been developed to predict the safety of hydroxylamine use in solvent extraction purification of plutonium. The model allows prediction of the stability of solutions containing HAN in processing vessels under a range of conditions. Conditions for which the hazardous autocatalytic reaction of HAN with nitric acid can occur are predicted so that they can be avoided. The model calculates concentrations of important chemical species in solution over time for a given set of input parameters. It consists of a chemical kinetics model in which the reaction rates for 19 nitrogen, plutonium, and iron species are estimated. The stability of HAN in nitric acid process solutions is controlled by complex relationships between temperature and concentrations of HAN, nitric acid, plutonium and iron. Process parameters (flow rates, volumes, initial reagent concentrations and temperatures) can affect the values of each of these parameters. It is found that plutonium(III) in nitric acid solutions containing HAN significantly lowers the stability of the solutions toward the autocatalytic HAN reaction with nitric acid. Dissolved iron also catalyzes this reaction, but is not as significant as plutonium because of its expected low concentrations. High acid concentrations and temperatures also promote the autocatalytic reaction. For normal extraction process conditions, the model confirms the absence of significant risk, in accordance with extensive industrial experience gained in reprocessing plants in Europe and Japan. References Andreychuk, N. N., Frolov, A. A., Rotmanov, K. V., & Vasiliev, VYa (1990). Plutonium (III) oxidation under a-irradiation in nitric acid solutions. Journal of Radioanalytical and Nuclear Chemistry, 142, 427e432. Ashcraft, R. W., Raman, S., & Green, W. H. (2008). Predicted reaction rates of HxNyOz intermediates in the oxidation of hydroxylamine by aqueous nitric acid. Journal of Physical Chemistry A, 112, 7577e7593. Barney, G. S. (1976). A kinetic study of the reaction of plutonium(IV) with hydroxylamine. Journal of Inorganic and Nuclear Chemistry, 38, 1677e1681. Barney, G. S. (1971). The reaction of hydroxylamine with nitrous acid. ARH-SA-97. Richland, Washington: Atlantic Richfield Hanford Company. Bengtsson, G., Fronaeus, S., & Bengtsson-Kloo, L. (2002). The kinetics and mechanism of oxidation of hydroxylamine by iron(III). Journal of the Chemical Society Dalton Transactions, 2002, 2548e2552. Bennett, M. R., Brown, G. M., Maya, L., & Posey, F. A. (1982). Oxidation of hydroxylamine by nitrous and nitric acids. Inorganic Chemistry, 21, 2461e2468. Cisneros, L. O., Rodgers, W. J., & Mannan, M. S. (2002). Effect of air in the thermal decomposition of 50 mass% hydroxylamine/water. Journal of Hazardous Materials, A95, 13e25. Cisneros, L. O., Rodgers, W. J., & Mannan, M. S. (2004). Comparison of the thermal decomposition behavior for members of the hydroxylamine family. Thermochimica Acta, 414, 177e183. Cisneros, L. O., Wu, X., Rodgers, W. J., Mannan, M. S., Park, J., & North, S. W. (2003). Decomposition products of 50 mass% hydroxylamine/water under runaway reaction conditions. Transactions of the IChemE, 81, 121e124. DOE/EH-0555. (1998). Technical report on hydroxylamine nitrate. Washington, DC: U. S. Department of Energy. Dukes, E. K. (1960). Kinetics and mechanisms for the oxidation of trivalent plutonium by nitrous acid. Journal of the American Chemical Society, 82, 9. Epstein, I. R., Kustin, K., & Warshaw, L. J. (1980). A kinetics study of the oxidation of iron(II) by nitric acid. Journal of the American Chemical Society, 102(11), 3751e3758.
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G.S. Barney, P.B. Duval / Journal of Loss Prevention in the Process Industries 24 (2011) 76e84
Gowland, R. J., & Stedman, G. (1981). Kinetic and product studies on the decomposition of hydroxylamine in nitric acid. Journal of Inorganic and Nuclear Chemistry, 43(11), 2859e2862. Koltunov, V. S., & Marchenko, V. I. (1973). Reaction between plutonium(III) and nitrous acid. Radiokhimiya, 15, 777e781. Kumasaki, M. (2004). Calorimetric study on the decomposition of hydroxylamine in the presence of transition metals. Journal of Hazardous Materials, 115, 57e62. Kumasaki, M., Fujimoto, Y., & Ando, T. (2003). Calorimetric behaviors of hydroxylamine and its salts caused by Fe(III). Journal of Loss Prevention in the Process Industries, 16, 507e512. Liu, L., Papadaki, M., Pontiki, E., Stathi, P., Rodgers, W. J., & Mannan, M. S. (2009). Isothermal decomposition of hydroxylamine and hydroxylamine nitrate in aqueous solutions in the temperature range 80e160 C. Journal of Hazardous Materials, 165, 573e578. Liu, L., Wei, C., Guo, Y., Rodgers, W. J., & Mannan, M. S. (2009). Hydroxylamine nitrate self-catalytic study with adiabatic calorimetry. Journal of Hazardous Materials, 162, 1217e1222. Long, L. A. (2004). The explosion at concept sciences: hazards of hydroxylamine. Process Safety Progress, 23, 114e120. Newton, T. W. (2002). Redox reactions of plutonium ions in aqueous solutions. In D. Hoffman (Ed.), Advances in plutonium chemistry 1967e2000 (pp. 24e57). La Grange Park, Illinois: American Nuclear Society. Orban, M., & Epstein, I. R. (1982). Bistability in the oxidation of iron(II) by nitric acid. Journal of the American Chemical Society, 104, 5918e5922.
Pembridge, J. R., & Stedman, G. (1979). Kinetics, mechanism, and stoichiometry of the oxidation of hydroxylamine by nitric acid. Journal of the Chemical Society Dalton Transactions, 1979, 1657. Perrott, J. R., & Stedman, G. (1976). The kinetics of nitrite scavenging by hydrazine and hydrazoic acid at high acidities. Journal of Inorganic and Nuclear Chemistry, 39, 325e327. Pitzer, K. S. (1991). Activity coefficients in electrolyte solutions (2nd ed.). Boca Raton, Florida: CRC Press. Raman, S., Ashcraft, R. W., Vial, M., & Klasky, M. L. (2005). Oxidation of hydroxylamine by nitrous and nitric acids. Model development from first principle SCRF calculations. Journal of Physical Chemistry A, 109, 8526e8536. Schmid, G., & Bahr, G. (1964). Formation of HNO2 from NO and HNO3 at high acid strength. Zeitschrift für Physikalische Chemie, 41, 8e25. Schulz, W. W. (1976). The chemistry of americium. ERDA critical review series, report no. TID-26971. Schwartz, S. E., & White, W. H. (1983). Kinetics of reactive dissolution of nitrogen oxides into aqueous solutions. In J. O. Nriagu (Ed.), Advances in environmental science and technology, Vol. 12 (pp. 3e109). New York: John Wiley & Sons. Sugikawa, S., Umeda, M., Sekino, J., Matsuda, T., & Kodama, T. (2002). Evaluation of safety limit for handling hydroxylamine nitrate/nitric acid solution as a plutonium reductant. In JAERI conf 2002e004 (pp. 475e480). Tokai-mura, Naka-ken 319-1195, Japan: Japan Atomic Energy Research Institute. Yusaku, I., Masahide, W., Takashi, U., & Hiroshi, K. (2005). Cause investigation of explosive accident of hydroxylamine and risk evaluation of its water solution. Report of National Research Institute of Fire and Disaster, 99, 114e119.