M. D. SILVERMAN, G. M. WATSON, and H. F. McDUFFlE Chemistry Division, O a k Ridge National Laboratory, O a k Ridge, Tenn.
Peroxide Decomposition in Aqueous Homogeneous Reactor Fuels
THE
power level of an aqueous homogeneous reactor appears to be limited a t low temperatures by the concentration of peroxide produced by fission fragments acting on the solution. Uranyl peroxide, formed by the reaction of hydrogen peroxide with uranyl ion, is a slightly soluble compound; its precipitation in a circulating homogeneous reactor system would undoubtedly pose serious problems. The rate of formation of peroxide in the reactor is dependent on the power density. However, the rate of decomposition of hydrogen peroxide in aqueous solution is known to depend largely on a number of other factors-e.g., temperature, concentration, and trace impurities. A kinetic study of the decomposition of hydrogen peroxide in uranyl sulfate solutions was, therefore, undertaken. Small amounts of corrosion products (iron, nickel, and chromium), formed by the action of uranyl sulfate solution on 347 stainless steel a t elevated temperatures, and measurable quantities of fission products \vi11 undoubtedly be present in the reactor system. I t was deemed advisable to expand this kinetic study in order to include an investigation of the effect of some of these species on the rate of decomposition of peroxide. In the radiation-induced decomposition of water, two reactions appear to occur simultaneously ( I ) :
I n the study of water decomposition by fission recoils ( 5 ) ,it was concluded that the decomposition of water in aqueous homogeneous-reactor solutions occurred principally as shown by Equation 1. Thus the G value for hydrogen production (molecules produced per 100 e.v. of energy absorbed in solution) may also be used as a measure of the rate of peroxide production. Although considerable work has been reported on the kinetics of systems involving hydrogen peroxide ( Z ) , there is no mention of studies in aqueousuranium sulfate systems. When hydrogen peroxide is introduced into a uranyl sulfate solution, there is an immediate reaction UOnf+
+
H202-+
+ 2H+
(3)
That the hydrogen peroxide reacts essentially quantitatively can be demonstrated in dilute solutions by the fact that the change in hydrogen ion concentration is consistent with the amount of peroxide added according to the stoichiometry of Equation 3, whether or not the solubility of uranyl peroxide is exceeded. Experiments were planned so that precipitation was avoided. The decomposition of peroxide can be represented by two alternate mechanisms : A:
UOc
+ 2H'
+
H?O?--t Hn0
B:
uo.4 UOs
1238
UO,(aq.)
INDUSTRIAL AND ENGINEERING CHEMISTRY
+
UOa
+ 21Jf
UO;*
4-H202
4- '/2 On
+ '/z +
UO,'
(4) (5) (6)
0 2
+
H20
If the rate-determining step is either Equation 5 or 6 and whether or not the mechanism is A or B, a conductance method for peroxide decomposition should approximate the chemical or titration method, provided a linear relation exists between conductance and hydrogen ion concentration. The fact that the equilibrium in Equation 3 is far to the right increases the sensitivity of the conductance method. Chemical titration, which determines both hydrogen peroxide and uranyl peroxide additively, also does not distinguish between alternative mechanisms A and B. Dilute solutions (about 8 grams of uranium per liter) were used for the conductance work in order to avoid buffering and nonlinear effects. Standard chemical titration methods were first employed for determining rates of decomposition of peroxide in this system. Chemical methods, however, proved inadequate for the determination of progressively faster rates brought about by the addition of various catalytic species. Therefore, it was necessary to develop a method for following these fast rates in situ. -4conductance method was developed which was found to be reproducible and consistent with that of chemical titration. The details of the method may be found elsewhere ( 7 7).
(7)
Apparatus
Chemical Method. The reaction vessel for all experiments was a modified three-necked flask. A motor-driven
Table 1. Decomposition of Hydrogen Peroxide in Aqueous Uranyl Sulfate Solutions at 78.0" C.
"
k,
pH,
x
104
1.6
4.8 4.70 5.07 4.78 4.60 4.90 6.00a 6.39a 5.97a 4.94 5.05
3.84 4.37 4.32 1.02 4.95
2.60
0.0486 0.0497 0.0497 0.0526 0.0526 0.0176 0.0176 0.0175 0.0178 0.0178 0.0045 0.004.5 0 0
4.08 3.80 4.02 1.24 1.27 0.976 1.14 1.37 0.496 0.420 0.380 0.463 5.19 4.72
3.05
Corrosion products
Sec.-[
c.
2.50
0.162 0.162 0.162 0.176 0.162
a
.. 3.3
Fircinn n m d i i r t s
Others
Promoter combination
5.29 5.30 5.26 4.09 4.07 3.35 4.26 0.0685 0.0748
3.5 2.5
I-
t \ A
2;O
S'eo
$-
2bO
3bO
3iO
Ah, Sec
kl, Liters Mole-' Set.-'
11 16 24
0.081 0.003 0.005
2240 44
Fe+2
Ni +l Rii
3?0
40'
Figure 1. Peroxide decomposition as a function of temperature
.
cu
Ti
..
2020 76
+4
+(
Fe+2 Cu i-2
Reagents
glass stirrer was inserted through the center opening. Hydrogen peroxide was pipetted through a side opening of the flask into the reacting solution after the latter was brought u p to temperature in a thermostat. The other side opening was kept closed during the chemical experiments. Conductance Method. A simplified version of the Jones conductivity bridge, applicable for rapid null-point measurements ( 7 7), was employed in conjunction with small platinum bead electrodes.
2'60
Concn.,
P.P.M. Cr +3
Precipitation observed.
"k50
Catalysts for Peroxide Decomposition in Uranyl Sulfate Solutions
Species Added
Initial Concn. Concn. U02S04, Hz02, M M x 103 0.65 5.2
i
Table II.
Unstabilized 30% hydrogen peroxide, diluted with double-distilled water, was employed for all experiments. Three different preparations of uranyl sulfate were used. Preparation A (glass) was prepared in the laboratory; preparation B (glass and steel) was prepared in plant production, and preparation C was a sample of inactive fuel solution which had been circulated in the homogeneous reactor experiment loop system for a week a t 250' C. This solution doubtless contained impurities which could be expected to give a higher rate constant for the decomposition of peroxide. All other reagents were ACS grade.
850 3
0.018 0.0010
6 793
0.34
20 8370
..
..
ing the peroxide. Whenever fast rates of decomposition were encounterede.g., those caused by catalytic effects of iron and ruthenium-approximate settings were made on the Helipot before adding hydrogen peroxide to the reaction solution to take full advantage of the anticipated shift in conductance. Some difficulty was encountered in obtaining representative samples of the heterogeneous catalysts because of the size of the solid particles suspended in solution. Results
The data obtained by the chemical method are given in Table I and Figure 1. The data given in Table I1 and
Procedure
For the chemical method, standard kinetics techniques were employed. Portions of hydrogen peroxide of '/z to 1 ml. were added to 50 to 100 ml. of uranyl sulfate solution held at constant temperature. At various times after addition, samples of the reaction mixture were withdrawn and pipetted into a n excess of ceric sulfate, which was then back-titrated with ferrous ammonium sulfate solution using o-phenanthroline as indicator. A sharp end point for the titration was obtained even in the presence of 1M of uranyl sulfate. Catalytic studies were performed as follows: Definite concentrations of the various species were added to the reaction flask about 30 seconds before the hydrogen peroxide was added, and conductance readings were taken after allowing enough time for mixing (approximately 30 seconds). Additional catalyst was usually inserted for successive runs until further additions had no effect on the rate of decomposition. Whenever it was suspected that hydrolysis and/or precipitation of the catalyst might occur rapidly--e.g., Ru(1V)only enough time was allowed for its dispersion (5 to 10 seconds) before add-
0 021 0
20
,
40
I
60
80
r00
T I M E , SEC.
Figure 2. Catalytic effect of iron on peroxide decomposition A. Uncatalyzed 6. Fs(ll), 0.22 p.p.m. C. 0.45 P.p.m. D. 0.67 P.p.m. VOL. 48, NO. 8
E.
0.90 P.p.m.
F. G.
2.0 P.p.m. 4.3 P.p.m. 6.5 P.p.m.
H.
AUGUST 1956
1239
I
I20
25,600, and 23,700 calories for preparations A , B, and C, respectively, were calculated from the slopes. No reference could be found regarding the activation energy for the decomposition of hydrogen peroxide in dilute aqueous solutions in the absence of catalytic materials. Data for correlating the chemical and conductance methods, which agree within 57G over a wide range of concentration and rates, are given elsewhere ( 7 7 ) . The results of catalytic studies employing the conductimetric technique are listed in Table 11. The data have been analyzed by the method of Guggenheim (7). The different catalysts studied are grouped in Table I1 based on thc probable source of origin in a homogeneous reactor fuel. Catalytic action
-
0
I
I
'
2
4
6
'
8
I
'
I
I
I
'
k = Xu
40 42 14 46 18 20 22
CONCENTRATION OF CATALYST, PPM
Figures 2 to 5 were obtained by the conductance method. Examination of the time dependence of the peroxide concentration indicates that the rate is first order with respect to peroxide concentration and apparently independent of the uranyl sulfate concentration in this range. The firstorder rate constants, k, in column 4 of Table I were calculated by
The concentration of peroxide determined in these experiments by the chemical method must include peroxide in the form of both hydrogen peroxide and dissolved uranyl peroxide. There is a n apparent decrease in the rate constant for the decomposition of peroxide in 0.0045-mole uranyl sulfate experiments. Because of experimental difficulties, however, it is not possible to ascertain whether or not the difference is real a t this concentration. The large difference between these values in the neighborhood of 4 X and that for .'pure'' water, 7 X 10-6, are indeed real and should be noted. Apparently the rate constant is also independent of the p H of the solution. Uranyl sulfate solutions of the concentration range listed cover a p H range between 1.5 to 3.5. I t seems extremely doubtful that the effects caused by changing the p H would so closely cancel any inverse effects due to variation of the uranyl sulfate concentration. The temperature dependence of the decomposition reaction of peroxide in all three preparations is shown in Figure 1, Activation energies of 25,300,
+ kijc]
which is valid for this region. Promoter action was observed, as expected, with combinations of catalysts reported in the literature (2). An illustration, iron-copper promotion, is shown in Figure 5. Data obtained in separate experiments indicate that similar catalytic effects are observed if iron is added to uranyl sulfate in the form of ferrous or ferric sulfate. I n either case, it was not possible to determine whether or not the ferrous-ferric ion concentrations had adjusted to similar ratios before undertaking the rate measurements. I t was very apparent, however, that the iron was undergoing relatively rapid hydrolysis a t 100' C. and a p H of 3. At low concentrations of iron (approximately 0.2 p,p,m.), reproducible reaction rates were obtained rept-atedly with addi-
Figure 3. Rate of peroxide decomposition as a function of catalyst concentra tion
1 240
ments become proportionately less with increasing concentration and a "plateau" is finally reached where the rate is independent of concentration. The height of the plateau in these experiments appears to depend on the nature of the catalyst used. The relative heights of the plateaus can be obtained from the magnitudes of Ak listed in Table 11. The corresponding concentrations of catalyst are also listed. Reproducible values of 4 k were obtained, regardless of whether uranyl sulfate solution A or B was employed. The catalytic activity of some species has been expressed alternatively in terms of the catalytic coefficient, kl, as given in the last column. These were obtained from the linear portions of the experimental curves and the relation
"4
40
20
30 40 TIME, SEC.
50
60
Figure 4. Loss o f ruthenium catalytic activity with time A.
Uncatalyred
D. 7 Min. after addi-
B. C.
RdIV) 1.8 p.p.m. 3 Min. after addi-
E.
tian
tion 12 Min. after addition
is shown by two different types of graphs. One type is given in Figure 2 where the fraction of peroxide present is plotted against time on semilog scale for varying amounts of species added. The other type is shown in Figure 3 where Ak = k - ko is plotted as ordinate against concentration of the catalyst. Here ko is the first-order specific reaction-rate constant for the decomposition of peroxide in uranyl sulfate solution in the absence of any added catalyst, and k is the value obtained in the presence of catalyst. I n all cases studied, the catalytic action follows a similar pattern. At extremely dilute concentrations of catalyst the decomposition rates increase essentially linearly. The rate incre-
INDUSTRIAL AND ENGINEERING CHEMISTRY
Figure 5. Promoter effect o f copper in iron catalyzed peroxide decomposition
,
35
30t
F e ( l I l , C u ( I I i 7 9 3 RRM.
1
4
i
I
m ---/
Fe (El
I ~
CONCENTRATION OF F e ( I I ) , FeM.
tional portions of fresh peroxide without further introduction of iron. Other investigators (3, 70) have also shown that the catalytic decomposition of peroxide by iron is due to ionic iron, but there is still considerable uncertainty as to whether ferrous or ferric ion is the stronger catalyst. There is general agreement, however, that an equilibrium exists between the two ions so that catalytic action occurs when both are present (2). Catalytic data for ruthenium(1V) as shown in Figures 3 and 4 indicate that this species hydrolyzes very rapidly in this system at 100’ C. The chemical behavior of ruthenium is known (6) to agree with this observation. I t appears that Ru(IV) is an even better catalyst than iron, but it hydrolyzes a t a more rapid rate. A half life for the rate of hydrolysis of Ru(1V) at 100’ C. in this system, estimated from Figure 4, is approximately 6 minutes. Application t Q Homogeneous Reactor Operation
The experimental results of this investigation suggest that the rate of change of peroxide concentration with time at any fixed temperature in an aqueous system may be represented by
factor of 2 over the temperature range from 25’ to 100’ C. It is evident from the data presented that catalysis by ionic species can greatly accelerate the rate of decomposition of peroxide. Since some of the species investigated would undoubtedly be present in a homogeneous reactor system (especially iron, nickel, and chromium in 347 stainless steel), the effect these accelerated rates will have on reactor power operating levels is of interest. Uncertainty as to the possible valence states in which some of the fission products occur or are formed precludes any attempt to use the rate data obtained for fission products in this work for this purpose. I t seems reasonable, however, to make such an estimate by using Equation 8 and the rate constants obtained for aged iron-containing solutions. This estimate would therefore tend to be on the conservative side in predicting what power levels could be attained in a n aqueous-uranyl sulfate homogeneous reactor. For the steady state in a reactor system, the term dC/dt (in equation 8) becomes equal to zero and K = kc,,. The production of peroxide in a reactor is considered to be a function of the G value (molecules of hydrogen produced per 100 e.v. of energy absorbed in solution) and the energy input (power density) as given by the formula
time, it is conceivable that uranium peroxide precipitation could occur following shutdown if the solution were cooled below a certain temperature. The amount of peroxide formed by delayed neutrons and radiation due to fission products (gamma and beta activity) is only a small percentage of that formed during operation. Following shutdown, a “safe” temperature (100’ C.) at which the solution could be held may be estimated by using Figure 1, where k is plotted against the reciprocal of the absolute temperature. The quantitative aspects of this work were experimentally verified during the operation of the Homogeneous Reactor Experiment in the fall of 1953. The usefulness of the foregoing data was also aptly demonstrated in a recent interpretation of the reactivity loss encountered during the operation of the aqueous homogeneous research reactor at North Carolina State College (4). Acknowledgment
The authors wish to acknowledge their gratitude to R . A. Dandl of the Instrument Division, Oak Ridge National Laboratory, for the design and construction of the modified Jones conductivity bridge used in this investigation and to C. H. Secoy for his kind encouragement.
K = 0.0062 X G X P D where K represents the rate of addition of peroxide (or rate of formation in the reactor), k is the specific reaction-rate constant for the decomposition of peroxide, and C is the concentration of peroxide in the solution. Rigorously Equation 8 should include a negative third term to account for the decomposition of peroxide by radiation only. The rate due to this effect depends on the inverse square root of radiation intensity and the square root of peroxide concentration in the range above 0.01 M . Since the steady-state concentration of concern in reactor solutions is estimated at 0.0039M, Equation 8 affords a n excellent approximation; moreover, it is a conservative estimate in that it gives a lower limit for the allowable peroxide concentration. I n a reactor, K is dependent on the fission density and G value and is independent of temperature; k is strongly dependent on temperature and other solution factors; and the upper limit for C depends on the solubility of uranyl peroxide in uranyl sulfate solutions a t that temperature. The solubility of uranyl peroxide as a function of temperature is unknown, although exploratory experiments (8) at ORNL have given a n order of magnitude, 4 X 10-3M, for the solubility at room temperature and 100’ C. These preliminary data indicate that the solubility does not vary by more than a
References
where PD is expressed in kilowatts per liter. For a dilute uranium sulfate solution (approximately 0.17M), the G value is approximately 1.5 and the peroxide formation is K = 0.0093 PD, in moles of hydrogen peroxide produced/ minute-liter of solution. The power density which may be tolerated (without fear of uranyl peroxide precipitation) may then be expressed as
Allen. A. 0..Davis. T. W.. Hochanadel, C. J., Ghormlev, ’J. A., J . Phys. Chem. 5 6 , 575 (1952).
Baxendale, J. H., Decomposition of Hydrogen Peroxide by Catalysts in Homogeneous Aqueous Solution in “Advances in Catalysis,” vol. IV, pp. 31-86, Academic Press, New York, 1952. (3) Ibid., p. 58. (4) Beck, C. K., Nucleonics 13, No. 7, 58 (1955 - -)., \
PD
=
k-- C,, 0,0093
Employing a value of 0.0039 mole for C,,, the allowable steady state peroxide concentration obtained from other experiments, and a value of 1.2 min.-1 for k, the reaction-rate constant for an aged iron-containing uranyl sulfate solution, an allowable power density of 0.5 kw./ liter a t 100’ C. is estimated. The accumulation or addition of catalytic materials, such as corrosion products, fission products, and possible recombination catalysts, and the conservative value for the allowable peroxide concentration in solution will all tend to permit actual reactor operation a t higher power levels. This statement has been confirmed where power levels as high as tenfold that given have been attained under actual operating conditions. After operation of the reactor a t appreciable power densities for some
( 5 ) Boyle, J. W., Ghormley, J. A., Hochanadel, C. J., Kieffer, W. F., Sworski, T. J., Geneva Conference, on Peacetime Uses of the Atom, Paper No. 524,1955. ( 6 ) Gmelin’s Handbuch der anorqischen Chemie, 8 Auflage, Verlag Chemie, Berlin. 1938. (7) Gugqenheim, E. A., Phil. Mag. 2, 538 (1926). (8) Silverman, M. D., Quarterly Rept. of
Homogeneous Reactor Project, period ending Oct. l , 1952, ORNL1424, p. 161. (9) Silverman, M. D., Watson, G. M., McDuffie. H. F.. Classified Journal, Reactor Sii. Technol. 3, No. 3, 43
(1953). (10) Simon, A., Haufe, W., Reetz, T.,
Preissler, R., Z . anorq. u. allgem.
Chem. 230, 129-47 (1936). (11) Watson, G. M., Silverman, M. D., McDuffie, H. F., Anal. Chem. 28, 1107 (1956).
RECEIVEDfor review December 27, 1955 ACCEPTED March 9, 1956 Division of Industrial and Engineering Chemistry, 129th Meeting, ACS, Dallas, Tex., April 1956 VOL. 48, NO. 8
AUGUST 1956
124 1
