Rate of Hydrogen Peroxide Decomposition in Nitric Acid Solutions Frend J. Miner1and Paul G. Hagan Rocky Flats Division, Dow Chemical U.S.A., Golden, Colo. 80401
Rate constants were determined for the decomposition of hydrogen peroxide in nitric acid solutions as a function of nitric acid (1.7-4.1M), hydrogen peroxide (2.3-4.9M), and an impurity mixture (Fe, Cr, Pb, Ni, and Cu in trace concentrations). The reactions were carried out at boiling temperatures. The effect of the nitric acid concentration on the rate constant is small, between 1.7 and 2.9M, but increases between 2.9 and 4.1M. The effect of the impurity concentration i s linear, increasing over the whole concentration range investigated. The hydrogen peroxide concentration has essentially no effect on the rate constant. Decomposition rate constants obtained in the absence of impurities are as follows: 1.7M HN03, 0.1 4 4 X 1 0-4 sec-l; 2.9M HN03, 0.599 X 1 0-4 sec-'; 4.1M HN03, 2.46 X 1 OV4 sec-'. The individual and combined effects of Cu(ll) and Fe(lll) on the decomposition rate constants were also investigated.
As part of a recovery process, a waste stream is produced which contains nit'ric acid, hydrogen peroxide, and trace concentrations of metallic impurit,ies. The hydrogen peroxide in t.his stream is destroyed by t.hermal decomposition using a batch distillat'ion process. This results in a continual increase in the concentration of nitric acid and metallic impurities in the still bottom. h proposed modification in this process requires recovery of the nitric acid with little or no change in its concent'ration. This could be accomplished b y thermal decomposition under reflux conditions rather than by a batch dist'illation process. For design purposes, data were needed on the rate of hydrogen peroxide decomposition as a function of the concentrations of nitric acid, hydrogen peroxide, and metallic impurities. Since there was only minimal information in the literature, a n experimental program was undertaken to obtain the required rate data. The results of this program are summarized in this paper. Experimental
Equipment and Procedure. T h e experimental decomposition work was carried out in a three-necked, lOOO-ml, round-bottom boiling flask. One neck contained a watercooled condenser open to the atmosphere, the second a thermometer, and the third a sampling device. This sampling device consisted of a dip tube extending into the bot'tom of the boiling flask. Samples were removed by connecting the sampling device to a small Erlenmeyer vacuum flask containing a 15-ml vial. Vacuum was used to transfer solution from the boiling flask into t'he vial. The vial was cooled in the Erlenmeyer flask with ice to stop the hydrogen peroxide decomposition reaction. The boiling flask was heated v d h a hemispherical heating mantle controlled b y a variable transformer. The rate of heat input was estimated b y measuring the rate of temperature rise of water in the apparatus. This rate was 21 cal/sec. Experiments were run by adding 200 ml nitric acid, 144 ml hydrogen peroxide, and 1 ml of a n impurity solution to the reaction flask. The concentrations of each of these solutions were varied according to the experimental design. The flask was then immediately placed in the heating mantle To whom correspondence should be addressed.
which was already at temperature. Samples were withdrawn periodically for hydrogen peroxide analysis. The frequency of sampling was determined by the rate of hydrogen peroxide decomposition. This rate was estimated initially by the vigorousness of bubbling (caused b y the oxygen evolution from the hydrogen peroxide decomposition) and by the rate of temperature rise (the rate of temperature rise increased with a n increase in the rate of decomposition). Six to 10 samples were taken a t recorded time intervals after the decomposition reaction began. These samples were analyzed for hydrogen peroxide concentration and the data used to calculate the decomposition rate constant. I n all runs, the total solution volume was 345 ml. The decomposjtion reactions were carried out a t the boiling temperatures of the solutions. This varied from 93' to 99OC depending on the composition of the solution. The boiling temperature was used for this investigation rather than some lower temperature because it results in a faster rate for decomposition, a situation desired for a production process. A11 chemicals, except the hydrogen peroxide, were reagent grade. The hydrogen peroxide was commercial grade, 12X (nominally 35 wt %) material. Specifications for the hydrogen peroxide require a total impurity content of less than 135 ppm. Analyses indicated the following concentrations in ppm of metallic impurities: 0.4 Cr, 0.3 Cu, 2 Fe, 0.5 Ki, 0.1 Pb, 25 Sn, 4 Ti, 0.1 V, and 2 Zn. Hydrogen peroxide concentrations were determined by a conventional spect.rophotometric method using Ti(1V) as the color-forming reagent (Egerton et al., 1954). Experimental Design. T h e region of experiment'ation was covered b y a 33 factorial design. Kine of 27 experiments required by this design were run in duplicate to obtain data on reproducibility. The parameters and levels used in the experiments are shown in Table I. The specific effects of Fe(II1) and Cu(I1) as impurities were determined in a subsequent series of experiments to be described later. The rate const.ants of the decomposition reaction were determined by plotting time vs. log of absorbance. Points lying on the straight line portion of the plot (indicating that Ind. Eng. Chem. Process Der. Develop., Vol. 1 1 , No. 4, 1972
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Table I. Experimental Parameters and levels Investigated Level 0
Parameter
Level 1
Level 2
Nitric acid, M 1.7 2.9 4.1 Hydrogen peroxide, M 2.3 3.4 4.9 Impurities, ml" 0 0.5 1 a Metal ions which were investigated as impurities were combined in a single solution. The concentration of these ions in the final HNOS-HZOZsolutions resulting from the addition of 1 ml of this impurity solution is as follows: Fe, 19.1 mg/l.; Cu, 1.7 mg/l.; Cr, 2.7 mg/l.; Ni, 3.5 mg/l.; Pb, 1.3 mg/l. Table II. Effect of HN03Concentration on the HzOz Decomposition Rate Constant HNOa concn, k x 104 M
1.7 2.9 4.1
Sec-l
SD, sec-I
0.144 0.599 2.46
3=0.012 +O. 005 ~0.006
the decomposition reaction had begun) were used to calculate the slope with a least-squares computer program. This slope was the rate constant of the reaction. The relative standard deviation of the analytical method used for hydrogen peroxide was 0.5%. The precision of the decomposition rate constants varied with the magnitude of the rates. The pooled standard deviation for the constants in various ranges are as follows (all values x sec-1): k < 1, 0.020; 1 < k < 10, 0.36; 11 < k < 20, 0.9; 21 < k < 70, 2.1; k > 70,21. Results and Discussion
An early reference indicates that the decomposition of hydrogen peroxide deviates from a firsborder reaction with respect to the hydrogen peroxide concentration in the presence of even low concentrations of nitric acid and Fe(II1) (Peterson, 1951). However, a more recent reference indicates the reaction is still first order in nitric acid concentrations as high as 8 M and in both the presence and absence of Fe(II1) (Elson, 1961). Our experimental results agree with this latter reference.
'0 0
---- HN03 0
-.-.-"
2 2
Figure 1 . Average effects of parameters on rate of HzOz decomposition 548 Ind. Eng. Chem. Process Des. Develop., Vol. 1 1 , No. 4, 1 9 7 2
n
3 0.2
0
l
I 5000
I 10000
Figure 2. Effect of " 0 3 composition rate constant
I 15000 T I M E , sac.
l 20000
, 25000
J 30000
concentration on the HzOz de-
Factorial Experiments. T h e average effects of t h e parameters on the decomposition rate constants are shown in Figure 1 in the form of average effect curves. D a t a for these curves were obtained from the experimental values of the 27 points in the factorial experiment b y averaging results from all experiments that contained a given parameter at a given level. For example, to calculate the decomposition rate constant a t 4.131 Hxoa, rate constants were averaged from all experiments that contained 4.1M "01, irrespective of what the other parameters and levels were. The average effect curves show that the rate of constants for the decomposition reach a maximum a t the highest acid and impurity concentrations investigated. There is initially only a small change in the rate constant as the nitric acid concentration is increased from 1.7 to 2.9M, but then there is a considerably larger change as the acid concentration is increased from 2.9 to 4.1-11. The effect of the impurity concentration is linear, increasing over the whole concentration range investigated. The hydrogen peroxide concentration has essentially no effect on the rate constants. This relative effect of the three variables is also indicated by the statistical evaluation of the factorial experiment. The changes in both the impurity concentration and the nitric acid concentration have statistically significant effects a t the 957, confidence level on the rate constants. However, the change in hydrogen peroxide concentration does not, indicating that the rate constants are independent of the hydrogen peroxide concentration. This is the requirement for a first-order reaction. The rate constants for hydrogen peroxide decomposition are summarized in Table I1 as a function of the nitric acid concentration. No impurities were present. In Figure 2 the effect of the nitric acid concentration on the rate of decomposition is shown in the form of the fraction of hydrogen peroxide remaining vs. time. Effect of Individual M e t a l I o n s o n t h e Decomposition Rate Constants. M a n y metal ions have been found to catalyze the decomposition of hydrogen peroxide in solution (Schumb et al., 1955, p 467). Most of the investigations, however, were carried out in neutral or near-neutral hydrogen peroxide solutions : for nitric acid-containing solutions, the nitric acid concentrations were no higher than 0.4,TI (Peterson, 1951; Garten, 1962). Information was needed in HxO3 solutions up to approximately 451. Exploratory experiments were run, therefore, in 4.1M HN03-2.3JI H202 solutions using each of the metal ions listed in Table I individually and a t the concentrations listed in the table. The results showed that only Fe(II1) and Cu(I1) increased the decomposition rate constant above that of 4.1M "03-2.3JI HZOZ
Table Ill. Effect of Fe(lll) Concentration on HzO, Decomposition Rate Constant
Solution: 1.7X "03, k
Fe(lll) concn
m/l.
0.73 1.45 7.3 38 95 190 381
mM
0.013 0.026 0.131 0.680 1.70 3.40 6.82
Table V. Combined Effect of Fe(lll) and Cu(ll) on H202 Decomposition Rate Constant
Solution: 1 . 7 N "03, 3.4M HZOZ
3.4M HZOZ
x
104, sec-'
k/mM Fe(lll), sec-l mM-'
0.192 0.201 1.26 12.2 33.6 69.6 133
14.8 7.7 12.2 18.0 20.0 20.5 19.6
Table IV. Effect of Cu(ll) Concentration on H202 Decomposition Rate Constant
k X
Fe(lll) concn,
ms/l. 0 38 95 190 381
lo4,sec-l
No Cu(ll)
0.83 mg/l. Cu(ll)
0.144 12.2 33.6 69.3 133
2.16 21.2 38.8 66.2 147
Table VI. Effect of HNO, Concentration on the Cu(ll) and Fe(ll1)-Cu(ll) Catalyzed Decomposition Rate Constant of H202
Solution: 3.4M H202
Solution: 1.7M HN03, 3.4M Hz02
~
Cu(ll) _ concn _
_
mdl.
mM
0.83 1.66 8.28 43.2 108 216 433
0.013 0.026 0.130 0.681 1.70 3.40 6.82
k X lo4, _ sec
-'
2.16 2.71 4.58 8.04 9.29 11.7 13.9
k/mM Cu(ll), rec-' mM-'
166 104 35.3 11.8 5.47 3.44 2.03
without impurities present. These two ions were investigated, therefore, in great,er detail. Effect of Fe(II1) and Cu(I1) on the Decomposition Rate Constants. Iron and copper were investigated individually in 1.751 HX03-3.4M HzOz solutions. The concentration of each met'al ion was varied individually, and the c0nstant.s for the rates of decomposition measured. The result's are summarized in Tables I11 and IV. Iron does not have the same effect as copper on the rate constants. For Fe(III), there is essentially a stmight line relationship between the rate const'ant and the concentrat'ion of Fe(II1). This is indicated by the relatively constant value for the ratio k / m X Fe(II1). Such is not the case with Cu(I1). The value of its ratio is init,ially quite large-11 times that of Fe(II1). But as the concentration of Cu(1I) increases, the value of the ratio k / m M Cu(I1) decreases rapidly so that a t 6.82 m d l Cu(I1) it is only about one-t,enth that of Fe(II1). Early work showed that Cu(I1) has a promoter (synergist'ic) action on the catalytic effect of Fe(II1) (Bohnson and Robertson, 1923; Schumb et al., 1955, p 534). This was indicated by t'he fact t'hat, the rate constant for decomposition was greater when the t'wo were present together than was t'he sum of the individual Fe(II1) and Cu(I1) rates. This promotion effect was investigated in this work and the experimental data are summarized in Table V. They show t.hat a t the lower Fe(II1) concentrations, this effect exists. But a t the higher Fe(II1) concentrations, the influence of Cu(I1) decreases. I n fact, based on the precision of t'he replicates at the two highest Fe(II1) concent,rations, the presence of Cu(I1) has no significant, effect on the decomposition rate constant's. Effect of Acidity on the Cu(I1) a n d Fe(II1)-Cu(I1) Catalyzed Decomposition Rate Constants. T h e d a t a in Table I1 show that' the rate constant of hydrogen peroxide decomposition increases with a n increase in the nitric acid
0.05 1.7 2.9 4.1
0.139 2.71 6.89 16.2
131 21.1 21.8 29.5
composition. T o determine if the acid concentration would also affect the rate of decomposition catalyzed b y Cu(I1) and a n Fe(II1)-Cu(I1) mixture, a series of experiments were run in which the nitric acid concentration was varied from 0.05 to 4.111.The results are summarized in Table VI. Increasing the nitric acid concent,ration increases the rate const,ants when the decomposition is catalyzed by Cu(I1). But when Fe(II1) is added, the effect is different. I n 0.05~16 acid, the decomposition rate constant is quite large. But in 1.7.11 acid the rate drops about, sixfold and increases only a little as the acid concentration is subsequently raised. X possible explanation for this behavior is based on the formation, in the low acid concentration, of a colloidal Fe(II1) hydrous oxide (or possibly a slightly soluble basis salt) (Schumb et' al., 1955, p 528). This m-ould have a large surface area and could be more active, catalytically, than Fe(II1) ions in solution. Acknowledgment
The authors acknowledge with thanks the assistance of Yvonne AI. Ferris and John Lynch in bhe design of t h e factorial experiment and the interpretation of the results, and Reuben Sironen in obtaining some of the experiment'al dat,a. literature Cited
Bohnson, V. L., Robert,son, A. C., J . Amer. Chem. SOC.,45, 2.512 (1923). Egert,on, A. C., Everett, A. J., Minkoff, G. J., Rudrakanchana, S., Salooja, K. C., Anal. Chim. Acta, 10, 422 (1954). Elson, R . E., U.S. At. Energy Comm. Rept,. UCRL-6536, (September 1961). Garten, V. A., Aust. J . Chem., 15, 719 (1962). Peterson, S.,J . Amer. Chem. Soc., 73,3521 (1951). Schumb, W. C., Sat,terfield, C. N., Wentworth, R. L., "Hydrogen Peroxide," Reinhold, New York, N.Y., 1953. RECEIVED for review February 7, 1972 ACCEPTED June 14, 1972 Work performed under U.S. Atomic Energy Commission Contract AT(29-1)-1106.
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