Catalytic decomposition of hydrogen peroxide by ferric ion in dilute sulfuric acid solutions

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<ul><li><p>Catalytic Decomposition of Hydrogen Peroxide by Ferric Ion in Dilute Sulfuric Acid Solutions </p><p>L.E. EARY </p><p>Hydrogen peroxide decomposition in acidic solutions is catalyzed by the free ferric ion, Fe &gt;. The following rate law for this reaction is determined by the initial rate method in solutions similar to those used for acidic in situ uranium leaching: </p><p>(mH,O,) (mFe 3-) -(dmH2o:/dt)25 oc = k </p><p>(mu-) </p><p>where k = 4.3 10 -3 s -1 at 25 ~ From 25 ~ to 50 ~ the activation energy is 85.6 kJ/mol. The decomposition of hydrogen peroxide proceeds by a particular redox reaction sequence that depends on the ratio of the concentrations of hydrogen peroxide to free ferric ion. The rate law determined here is consistent with the form derived from the redox sequence for the case where the ratio of hydrogen peroxide to free ferric ion concentration is greater than 1.0. The magnitude of the rate constant indicates that the decomposition of hydrogen peroxide may cause rapid loss of this oxidant in leaching solutions containing ferric ion. </p><p>I. INTRODUCTION </p><p>HYDROGEN peroxide has been used to provide the oxidizing capacity of lixiviants for the in situ leaching of sandstone-type uranium deposits. In experimental studies that used solutions similar to those expected for in situ leaching conditions, the rate of uraninite (UO2-U3Os) dis- solution has been found to be more rapid in both basic ~ and acidic: solutions containing hydrogen peroxide than in solu- tions containing other oxidants such as dissolved oxygen, 3 ferric ion, 4 and sodium perchlorate, s Any advantages in ura- ninite dissolution rate, however, are offset by the fact that during in situ leaching, a considerable portion of the added H202 is likely to be consumed by the oxidation of iron sulfides and organic material that are commonly present in sandstone-type uranium deposits.67 Also, experimental and field observations indicate that H202 rapidly decomposes soon after injection into uranium ores, s-l~ thereby quickly reducing the oxidizing capacity of the lixiviant. </p><p>Hydrogen peroxide decomposes according to the over- all reaction: </p><p>2H202(1) = 2H20(1) + O2(g) [1] </p><p>where (1) and (g) refer to the liquid and gaseous molecular species, respectively. The rate of this reaction is catalyzed by metal ions 11 such as Fe 3+, Cu &gt;, and Co 3+. Of these, the ferric ion is undoubtedly the most abundant catalytically active metal ion present during the acidic leaching of ura- nium ores. Concentrations up to 100 ppm total dissolved iron are reported by Tweeton, et al. p- from an in situ ura- nium leaching operation that used solutions containing dilute sulfuric acid, pH = 1.5 to 2.0, and hydrogen per- oxide. However, H202 decomposition may take place in slightly acidic to strongly basic solutions, as well. It has been shown by Eligwe, et a l ) 3 that the addition of ferrous </p><p>L.E. EARY, formerly a Graduate Student with the Department of Geo- chemastry and Mineralogy, Pennsylvania State University, is now Research Scientist with Battelle Pacific Northwest Laboratory, Rlchland, WA 99352. </p><p>Manuscript submitted March 20, 1984. </p><p>ions to H202 solutions with pH between 4.0 and 6.0 caused a significant decrease in the rate of uranium extraction from a New Mexico ore. As suggested by Eligwe, et al., J3 the observed rate decrease was probably a result of the rapid consumption of the H,O2 in reactions with ferrous ions to form ferric species. Although the aqueous ferric species are limited to low concentrations above a pH of about 3.5 by the low solubility of ferric hydroxide, H20: decomposition has also been shown to be promoted by freshly precipitated ferric hydroxides ~ in solutions with pH values between 4.3 and 11.3. </p><p>In acidic solutions, dissolved iron is produced in situ by the oxidation of iron sulfides, as in the oxidation of pyrite by H202: </p><p>2FeS2 + 15H20: = 2Fe~" + 4(SO4-), </p><p>+ 14820(1) + 2H + , [2] </p><p>where the subscript t refers to the total ferric or sulfate species, respectively. Ferric ions produced by such reactions will catalyze the decomposition reaction of H:O&gt; A series of experiments was conducted to determine the rate of this decomposition reaction and the rate dependencies on the concentrations of ferric ion, hydrogen peroxide, and hydro- gen ion in dilute sulfuric acid solutions with pH between 1.38 and 2.20. </p><p>II. EXPERIMENTS </p><p>The rate experiments were conducted in 800 ml open glass kettles that were immersed in a constant temperature bath. Between 25 ~ and 50 ~ the temperature in the bath could be controlled to -+ 1 ~ Magnetic stirrers and Teflon stir bars were used to keep the solutions well mixed. Rea- gent grade Fez(SOa)3 9 nH20 was used to provide the ferric species in the run solutions. The pH was adjusted initially by the dropwise addition of sulfuric acid. Distilled, deionized water was used in all experiments. </p><p>The pH was measured with a standardized combination- reference electrode. Hydrogen peroxide concentrations </p><p>METALLURGICAL TRANSACTIONS B VOLUME 16B JUNE 1985 181 </p></li><li><p>were determined colorimetrica]ly by the intensity of the yellow color produced upon the addition of H202 to a ti- tanium sulfate reagent.t4 Concentrations of ferrous, ferric, and total dissolved iron were also determined with col- orimetric methods. Total iron was determined from the color density of the ferrous-orthophenanthroline complex after reducing all dissolved iron to the ferrous state with hydroxy- lamine hydrochloride. ~ The concentrations of only the fer- rous ions in the presence of ferric were determined by the method of Tamura, et al. ~6 In this method, the reducing agent was not added, Instead, fluoride, as ammonium fluo- ride, was added to complex the ferric ions. With the ferric ions masked by the fluoride, the concentrations of only the ferrous ions were determined with the colored ferrous- orthophenanthroline complex. After the determinations of ferrous and total iron, ferric was calculated from the difference. </p><p>Rate experiments were started by adding a measured quantity of a standardized 30 pct H202 reagent to ferric sulfate solutions of known pH. The rate of H202 decom- position was measured by withdrawing 1.0 ml samples from the kettles at known time intervals to determine the remain- ing HzO2 concentration. Samples were added directly to prepared solutions of the titanium sulfate reagent to reduce the time between extraction and the determination of the H202 concentration. The color of the H202-titanium sul- fate complex was found to be stable for up to one hour in solutions containing ferric ions, but H202 determinations were completed within 5 minutes after sample extraction to minimize the effects of any continued decomposition. Ferrous determinations were also done immediately after sample withdrawal. However, ferrous concentrations were always below the detection limit of about 0.1 ppm. Total dissolved iron concentrations were determined at a Inter time. Most of the rate experiments lasted from 2 to 14 hours, but a few were extended to longer times to observe devia- tions from the initial rates of H202 decomposition. </p><p>The concentration dependencies or reaction orders of the rate of H202 decomposition were determined by the initial rate method.~6 According to this method, the concentrations of each of the reactants, HzO2, ferric ion, and hydrogen ion were varied in turn while the other two concentrations re- mained effectively constant. Plots of the logarithm of the decomposition rate vs the logarithm of the concentration of the reactant that was varied were used to determine the rate dependence on that particular reactant. </p><p>In the dilute sulfuric acid solutions used in the experi- ments, ferric ion was present both as the uncomplexed free ion and as ferric sulfate and hydroxide complexes. The effects of ferric speciation on the rate of H202 decom- position were not examined, but previous studies H'18 have suggested that it is only the free ferric ion, Fe 3+, that is the active catalytic agent promoting H202 breakdown. Con- sequently, the rate dependencies are given here in terms of the concentrations of free ferric ion and hydrogen ion. In order to calculate the distributions of the free ions and aque- ous complexes for this system, a version of the computer code of Cathles and Breen ~9 was used, Activity coefficients were calculated by this code with the extended Debye- Huckel equation. An example of the distribution of the ferric species at pH less than 3.5 is shown in Figure 1. The con- centration units used are in terms of molality: i .e . , mFe 3+ refers to the molality of the free ferric ion. The reactions and equilibrium constants used to construct this diagram are given in Table I. </p><p>The decomposition rate experiments were conducted over a range of solution compositions: 1.8 x 10 .3 tO 1.7 X 10-: mH202.36 x 10 -; to 1,3 10 -2 mFe~-. 8.4 x 10 .3 </p><p>g </p><p>la. </p><p>| i i i i i </p><p>t - . Fe 3" , l ""./ IFeSO 4 </p><p>~ I 9 \ / ',, , i ' . . -. </p><p>),' ",, o X 7( "",, '* / \ .,' ", _. .~+" </p><p>/ \ / ",, F~OH. L.-- / \,'" '-, e(SO, L </p><p>o / \ -,. ,z_..~:- -=. / ---~-._~ N , , \ _ .D . . , ,~" ' - . . . . . - ". . . . . . ',,,'- </p><p>~ . -"F,(OH); , , ,~ </p><p>0.0 I.O 2.0 3.0 pH </p><p>Fig. 1--Distribution of the femc species for the Fe~--SO.]--H:O system at 25 ~ w~th 10 -" 7~ m Fe~- and tv,,o concentratmns of sulfate The solid and dashed lines correspond to total sulfate concentratmns of 10 138 and 10-2 '~ m, respectlvel~ </p><p>Table I. Reactions, Equilibrium Constants, and References for the Fe,3+-SO]--H20 System. In Column 3, ~ Is an Adjustable Parameter Corresponding to the Size of the Ion </p><p>That Is Required in the Extended Debye-Huckel Equation for Calculating Activity Coefficients. </p><p>Reaction log K (25 ~ d X 10 l~ (Angstroms) Reference </p><p>H + + SO]- = HSO2 Fe 3+ + SO] = FeSO2 </p><p>Fe 3t- 4- 2804 = Fe(SO4) 2 Fe 3+ + HSO4 = FeHSO] + Fe 3+ + H20 = FeOH 2+ + H + Fe 3+ + 2H20 = Fe(OH) + + 2H + 2Fe 3. + 2H20 = Fez(OH) 4+ + 2}t + </p><p>1.99 9 (H+) ~ Smith and Martell 2~ 4.04 9 (Fe3-) a Smith and Martell 2~ </p><p>5 (FeSO2)" 5,38 7 (Fe(SO,)i) r Wagman, et al, 2~ 0.6 7 (FeHSOj+) c Sapieszko, et a l ) 2 </p><p>-2.19 5 (FeOH:*) ~ Smith and Martell 2~ -5.67 5.4 (Fe(OH)~) ~ Smith and Martell 2~ -2.95 11 (Fe:(OH)~+) ~ Smith and Martell 2~ </p><p>~Truesdel[ and Jones -~ bStumm and Morgan 2~ Cestlmated </p><p>182--VOLUME 16B, JUNE 1985 METALLL RGICAL TRANSACTIONS B </p></li><li><p>to 5.0 x 10 -2 m(SO ] ),, and pH = 1.38 to 2.20. The con- centrations of the free ferric and hydrogen ions were calcu- lated by the iron distribution program for each experiment. For the above ranges of solution composition, the concen- trations of free ferric ion were generally equal to about 5 pct of the total ferric concentration. Hydrogen ion concentra- tions were close to values that would be approximated from the pH of the solutions. </p><p>I I I . RESULTS AND DISCUSSION </p><p>In general, the rates of H202 decomposition were ob- served to be fairly rapid. Typical results of the H202 concen- tration change as a function of time are shown in Figure 2. The measured change in H202 concentration was a linear </p><p>O </p><p>~o "- O C~ </p><p>Fig </p><p>I I I L I l I I I I </p><p>i i ~ 004 0 0.0013 </p><p>0 -- 00.0030 </p><p>O </p><p>~176 S 10 l's 20 25 ao T IME, HOURS </p><p>2- -Typ ica l experimental results showing the decrease in H:O2 con- centratlon with time for three molal concentratmns of Fe 3~ at pH = I 8 and 25 ~ </p><p>function of time for an initial period, before the decom- position rate slowed significantly. Initial rates were deter- mined from the slopes of plots of H202 concentration vs time for the initial period when the linear relationship was clearly observed. Slopes were calculated by linear re- gression of this rate data. Depending on the reactant concen- trations, the initial period of linear change in H202 lasted from 2 to about 12 hours before deviations to slower decom- position rates were obvious. </p><p>The initial rates as a function of each of the reactant concentrations were used to determine the form of the rate law describing H202 decomposition. The log-log plots from which the rate dependencies on reactant concentrations were determined are shown in Figures 3 through 7. A com- plete summary of these rate dependencies and the reactant concentrations in the experimental solutions is given in Table II. The slopes and intercepts of the lines shown in Figures 3 through 7 were calculated by linear regression of the logarithm of the H202 decomposition rate as a function of the logarithm of the reactant concentrations. The error </p><p>O </p><p>A </p><p>tO o~ Od -r </p><p>"O </p><p>On O </p><p>t%l i r6 </p><p>m~ </p><p>~ 5.7 X IO -2 3.0 3'.4 318 </p><p>-log(m~e3,) Fig. 3 - -Dependence of the H,O- decomposmon rate on m~e3- with m ~ HzO2 = I 8 x 10-3at25~ </p><p>Table II. Summary of the Rate Dependencies, rife3+, nnzoz, nn*, and Reactant Concentrations. The Errors Associated with the Rate Dependencies Indicate Two Standard Deviations about the Slopes of the Lines Shown in Figures 3 through 7. </p><p>Figure m~162 * m~ + o nH~ /~H202 ~Fe 3. ~'2H202 3 1.80 x 10 2 1.80 x 10 -3 </p><p>3.70 10 2 1.80 10 -3 </p><p>4 3.70 x 10 2 5.80 x 10 3 1.80 x 10 2 5.80 X 10 -3 9.00 10 3 5.80 X 10 -3 </p><p>5 2.30 x 10 -5 3.70 x 10 -2 2.30 x 10 -5 1.80 x 10 -2 1.20 10 -4 3.70 x 10 -2 1.20 x 10 -4 1.80 x 10 -2 </p><p>6 3.30 10 -4 3.70 10 -2 3.30 x 10 -'~ 1.80 x 10 2 </p><p>7 3.20 10 -4 2.30 10 -5 1.20 x 10 .4 </p><p>1.80 x 10 3 5.80 x 10 3 5.80 x 10 .3 </p><p>1.16 (_+0.10) 0.95 (_+0.12) </p><p>1.11 (+_0.10) 0.91 (_+0.08) 0.92 (+0.06) </p><p>0.99 (_+0.14) 1.19 (_+0.18) 1.19 (_+0.12) 1.02 (+-0.06) </p><p>1.01 (_+0.08) 1.17 (_+0.06) </p><p>-0.89 (+__0.14) -1.21 (+0.18) -1.16 (+0.16) </p><p>METALLURGICAL TRANSACTIONS B VOLUME 16B, JUNE 1985-- 183 </p></li><li><p>q:- </p><p>o </p><p>0,,I 1,5 </p><p>I I I I I I I </p><p>~.8 '~X i 0 -2 </p><p>/ . . . . - , / 3'.8 ,i z . ;6 s:o </p><p>0 -log (~"/Fe3+) Fig. 4 - -Dependence of the H20_, decomposition rate on m~ ~ with m ~ H202 = 5.8 10 -3 at 25 ~ </p><p>O </p><p>A </p><p>O </p><p>v </p><p>O </p><p>,6 </p><p>l I I I </p><p>m~ -2 </p><p>10-2 </p><p>J - m~.:,.sx,o 2 d8 s 2.6 </p><p>-log (m~2o2) Fig 5- -Dependence of the H202 decomposition rate at 25 ~ on m ~ H202 with m%&gt; = 2.3 I0 -s (lines with open square and star symbols) and m~3* = 1 2 I0 -4 (hnes wlth open mrcle and filled square symbols) </p><p>bars associated with these lines represent two standard devi- ations about the slopes calculated by the regression analysis and give a measure of the experimental uncertainty in the determinations of the orders with respect to a given reactant. Figures 3 and 4 are plots of -log(dmH~O2/dt) against -log(m%3-), where -(dmHeO2/dt) is the rate of H202 de- composition in mol/kg H20-sec, and the superscript zero refers to the initial molal concentration. The slopes of the lines shown in these figures give the dependence, nF~3-, of the H202 decomposition rate on the initial concentration of </p><p>I I l I I I </p><p>"O cu T//H+=3.7XIO-2 : .5 \ ~- </p><p>-r </p><p>! </p><p>s.o z:o s s </p><p>-log (m~ Fig 6 - -Dependence of the U~O2 decomposatlon rate on m ~ H,O2 with m~ - = 33 x 10 4at25oc </p><p>I f i I I I </p><p>/7,/0 3~3. 2 X I0 -4 \T </p><p>9 .- tM Fe3, = 2.3X I0 5 "0 ,4" - .~ rg/ </p><p>o: I </p><p>2 ~ _o </p><p>! </p><p>"5 /7/0 =I .2X lO TM ~T Fe3+ </p><p>o 1:2 1:6 "2".0 0 -log (mH+) </p><p>Ftg 7- -Dependence of the H_,O_~ decomposition rate at...</p></li></ul>


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