Catalytic Decomposition of Hydrogen Peroxide
by Ferric Ion in Dilute Sulfuric Acid Solutions
L.E. EARY
Hydrogen peroxide decomposition in acidic solutions is catalyzed by the free ferric ion, Fe >. 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:
-(dmH2o:/dt)25 oc = k
(mH,O,) (mFe 3-)
(mu-)
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.
I.
INTRODUCTION
H Y D R O G E N 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) dissolution has been found to be more rapid in both basic ~ and
acidic: solutions containing hydrogen peroxide than in solutions containing other oxidants such as dissolved oxygen,3
ferric ion,4 and sodium perchlorate, s Any advantages in uraninite 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.
Hydrogen peroxide decomposes according to the overall reaction:
2H202(1) = 2H20(1) + O2(g)
[1]
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 >, and Co 3+. Of these, the
ferric ion is undoubtedly the most abundant catalytically
active metal ion present during the acidic leaching of uranium ores. Concentrations up to 100 ppm total dissolved
iron are reported by Tweeton, et al. p- from an in situ uranium leaching operation that used solutions containing
dilute sulfuric acid, pH = 1.5 to 2.0, and hydrogen peroxide. 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
L.E. EARY, formerly a Graduate Student with the Department of Geochemastry and Mineralogy, Pennsylvania State University, is now Research
Scientist with Battelle Pacific Northwest Laboratory, Rlchland, WA 99352.
Manuscript submitted March 20, 1984.
METALLURGICAL TRANSACTIONS B
ions to H 2 0 2 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.
In acidic solutions, dissolved iron is produced in situ by
the oxidation of iron sulfides, as in the oxidation of pyrite
by H202:
2FeS2 + 15H20: = 2Fe~" + 4(SO4-),
+ 14820(1) + 2H + ,
[2]
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> 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 hydrogen ion in dilute sulfuric acid solutions with pH between
1.38 and 2.20.
II.
EXPERIMENTS
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. Reagent 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.
The pH was measured with a standardized combinationreference electrode. Hydrogen peroxide concentrations
VOLUME 16B JUNE 1985
181
were determined colorimetrica]ly by the intensity of the
yellow color produced upon the addition of H202 to a titanium sulfate reagent.t4 Concentrations of ferrous, ferric,
and total dissolved iron were also determined with colorimetric methods. Total iron was determined from the color
density of the ferrous-orthophenanthroline complex after
reducing all dissolved iron to the ferrous state with hydroxylamine hydrochloride. ~ The concentrations of only the ferrous 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 fluoride, 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 ferrousorthophenanthroline c o m p l e x . After the determinations
of ferrous and total iron, ferric was calculated from the
difference.
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 H 2 0 2 decomposition was measured by withdrawing 1.0 ml samples from
the kettles at known time intervals to determine the remaining 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 sulfate 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 deviations from the initial rates of H202 decomposition.
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 remained effectively constant. Plots of the logarithm of the
decomposition rate v s the logarithm of the concentration of
the reactant that was varied were used to determine the rate
dependence on that particular reactant.
In the dilute sulfuric acid solutions used in the experiments, 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 decomposition 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. Consequently, 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 aqueous complexes for this system, a version of the computer
code of Cathles and Breen ~9was used, Activity coefficients
were calculated by this code with the extended DebyeHuckel equation. An example of the distribution of the ferric
species at pH less than 3.5 is shown in Figure 1. The concentration units used are in terms of molality: i . e . , mFe3+
refers to the molality of the free ferric ion. The reactions and
equilibrium constants used to construct this diagram are
given in Table I.
The decomposition rate experiments were conducted
over a range of solution compositions: 1.8 x 10 . 3 tO 1.7 X
10-: m H 2 0 2 . 3 6 x 10 -; to 1,3 • 10 -2 mFe~-. 8.4 x 10 . 3
|
i
t-.
g
i
i
i
,
""./
l
i
Fe 3"
IFeSO4
~I
9 \
la.
/
o'* /X \
/
o
N
/
~
0.0
',, , i
'..
),'
\
/
.,'
7(",
\,'"
/\
,,
\
-.
"",,
",,
'-, -,. e(SO,
L
,z_..~:--=.
.
",,
_..~+"
F~OH.L.--
/ ---~-._~
_.D..,,~"'-.....". .....
',,,'-"F,(OH); , , , ~
I.O
2.0
3.0
pH
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~
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
That Is Required in the Extended Debye-Huckel Equation for Calculating Activity Coefficients.
Reaction
H + + SO]- = HSO2
Fe 3+ + SO] = FeSO2
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. + 2 H 2 0 = Fez(OH)4+ + 2}t +
~Truesdel[ and Jones -~
bStumm and Morgan2~
Cestlmated
182--VOLUME 16B, JUNE 1985
log K (25 ~
1.99
4.04
5,38
0.6
-2.19
-5.67
-2.95
d X 10 l~ (Angstroms)
Reference
9 (H+) ~
9 (Fe3-)a
5 (FeSO2)"
7 (Fe(SO,)i) r
7 (FeHSOj+) c
5 (FeOH:*) ~
5.4 (Fe(OH)~) ~
11 (Fe:(OH)~+)~
Smith and Martell 2~
Smith and Martell 2~
Wagman, et al, 2~
Sapieszko, et a l ) 2
Smith and Martell 2~
Smith and Martell 2~
Smith and Martell2~
METALLL RGICALTRANSACTIONS B
to 5.0 x 10-2 m(SO ] ),, and pH = 1.38 to 2.20. The concentrations of the free ferric and hydrogen ions were calculated by the iron distribution program for each experiment.
For the above ranges of solution composition, the concentrations of free ferric ion were generally equal to about 5 pct
of the total ferric concentration. Hydrogen ion concentrations were close to values that would be approximated from
the pH of the solutions.
III.
RESULTS AND DISCUSSION
In general, the rates of H202 decomposition were observed to be fairly rapid. Typical results of the H202 concentration change as a function of time are shown in Figure 2.
The measured change in H202 concentration was a linear
O
I
I
I
L
I
i
i ~
~o
0
"- O
• C~
0
l
I
I
I
I
O
004
m~
A
-00.0030
0.0013
tO
oOd~
-r
"O
O
~176
function of time for an initial period, before the decomposition rate slowed significantly. Initial rates were determined from the slopes of plots of H202 concentration v s
time for the initial period when the linear relationship was
clearly observed. Slopes were calculated by linear regression of this rate data. Depending on the reactant concentrations, the initial period of linear change in H202 lasted
from 2 to about 12 hours before deviations to slower decomposition rates were obvious.
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 complete 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
On
O
t%l
r6
i
S
10
l's
TIME,
20
25
3.0
ao
5.7 X IO- 2
3'.4
318
-log(m~e3,)
HOURS
Fig 2 - - T y p i c a l experimental results showing the decrease in H:O2 concentratlon with time for three molal c o n c e n t r a t m n s o f Fe 3~ at p H = I 8
and 25 ~
~
Fig. 3 - - D e p e n d e n c e of the H , O - d e c o m p o s m o n rate on m~e3- with m ~
HzO2 = I 8 x 1 0 - 3 a t 2 5 ~
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.
Figure
m~162
*
3
4
5
6
7
2.30
2.30
1.20
1.20
3.30
3.30
3.20
2.30
1.20
x
x
•
x
•
x
•
•
x
10-5
10-5
10-4
10-4
10-4
10-'~
10 -4
10-5
10.4
METALLURGICAL TRANSACTIONS B
m~ +
1.80
3.70
3.70
1.80
9.00
3.70
1.80
3.70
1.80
3.70
1.80
x
•
x
x
•
x
x
x
x
•
x
10 2
10 2
10 2
10 2
10 3
10 -2
10-2
10 -2
10 -2
10-2
10 2
o
/~H202
1.80
1.80
5.80
5.80
5.80
x 10-3
• 10-3
x 10 3
X 10 -3
X 10 -3
~Fe 3.
1.16
0.95
1.11
0.91
0.92
(_+0.10)
(_+0.12)
(+_0.10)
(_+0.08)
(+0.06)
0.99
1.19
1.19
1.02
1.01
1.17
1.80 x 10 3
5.80 x 10 3
5.80 x 10.3
nH~
~'2H202
(_+0.14)
(_+0.18)
(_+0.12)
(+-0.06)
(_+0.08)
(_+0.06)
- 0 . 8 9 (+__0.14)
-1.21 (+0.18)
- 1 . 1 6 (+0.16)
VOLUME 16B, JUNE 1985-- 183
I
I
I
I
I
I
I
I
I
l
I
I
I
q:-
o
~.8'~X
i 0 -2
"O
cu
:.5
T//H+=3.7XIO-2
\
~-
-r
!
0,,I
/
/
1,5
....
-,
z:o
s.o
3'.8
,i z
0
-log (~"/Fe3+)
Fig 6 - - D e p e n d e n c e of the U~O2 decomposatlon rate on m ~ H,O2 with
m~ - = 3 3 x 10 4 a t 2 5 o c
Fig. 4 - - D e p e n d e n c e of the H20_, decomposition rate on m ~
H202 = 5.8 • 10 -3 at 25 ~
~
with m ~
I
l
I
I
m~
s
-log(m~
s:o
.;6
s
I
f
i
I
I
I
\T
/7,/0 3~3. 2 X I 0 - 4
-2
O
"9.0
A
tM
,4"
rg/ Fe3, = 2.3X I0 5
-.~
o:
I
O
10-2
v
2 ~
_o
!
O
"5
,6
J
-
m~.:,.sx,o 2
o
s
d8
2.6
-log (m~2o2)
Fig 5 - - D e p e n d e n c e of the H202 decomposition rate at 25 ~ on m ~ H202
with m%> = 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)
bars associated with these lines represent two standard deviations 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 decomposition 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
t84
VOLUME 16B, JUNE 1985
/7/0
Fe3+
1:2
=I.2XlO
TM
~T
"2".0
1:6
0
-log (mH+)
Ftg 7 - - D e p e n d e n c e of the H_,O_~ decomposition rate at 25 ~ on m ~
with m v H:O: = 1.8 x 10 -3 (hne with open square symbols) and m ~
H202 = 5 8 • l0 -3 (hnes ~]th star and open circle symbols).
free ferric ion. The slopes shown in Figures 3 and 4 indicate
nve> is equal to 1.0 (see Table II); hence, the rate of the
decomposition reaction is directly proportional to the initial
concentration of free ferric ion. Similarly, the slopes shown
in Figures 5 and 6 give the rate dependence on the H202
concentration, nil,o,. These slopes are also nearly equal to
1.0 (see Table II),-indicating the rate of H202 decomposition
is directly proportional to the concentration of H202. The
last rate dependence is determined from Figure 7. The
slopes in Figure 7 give the rate dependence on hydrogen ion
concentration, nil-. From Table II, it can be seen that n.- is
METALLURGICAL TRANSACTIONS B
approximately equal to -1.0; hence, it can be concluded
that the rate of H202 decomposition is inversely proportional to the concentration of hydrogen ion. The following
rate expression is determined from the rate data shown in
Figures 3 through 7:
0
- (dmn2o:/dt)25 ~ =
0 3
k (mH2o2)(mFe ")
(mO~)
[3]
where k = 4.3 (+ 1.3) x 10-3 (s-l). The magnitude of this
rate constant, k, is an average value calculated from all the
measured decomposition rates and the rate dependencies
given in [3].
Decomposition rates were also measured as a function of
temperature from 25 ~ to 50 ~ for two initial solution compositions: (1) H202 = 6.3 x 10-3m, Fe~+ = 2.2 x 10-3m,
pH = 1.5, and (2) H202 = 2.3 x 10-3m, Fe~* = 7.2 x
10-3m, pH = 1.5. From the decomposition rates measured
at five degree intervals between 25 ~ and 50 ~ apparent
activation energies of 82.4 kJ/mol and 88.7 kJ/mol were
calculated using the Arrhenius equation ~7 for the two respective solution compositions given above.
The rate equation given in [3] is similar in form to the
expression reported by Barb, et al. n for the H202 decomposition rate in dilute perchloric acid solutions with concentration ratios of H202 to free ferric ion greater than 1.0. In
dilute perchloric acid solutions, ferric ion is present predominantly as the free ion, and Barb, et a l Y describe the
H202 decomposition reaction path as a sequence of redox
reactions that are catalyzed by the free ferric ion. According
to Barb, et al.," and later substantiated by Walling and
Weil, 25 the redox sequence is:
k4
Fe 3+ + H202(1) ~ Fe e+ + HO2" + H +
[4]
k5 Fe3+ + HO. + OH
Fe z+ + H202(1)~
[5]
k6
HO. + H202(1)--~ H20(1) + HO2"
[6]
[7]
HO2" + Fe z+ k8 HO_~ + Fe3+
[8]
HO" + Fe z+ k~ Fe3+ + OH
[9]
According to Walling and Weil,25 steps [4] and [7] also
involve equilibria with hydrogen ion prior to going to completion. As such, both steps [4] and [7] can each be considered as the summation of two reactions. For step [4], the
first of these two reactions is the equilibria with hydrogen
ion which may be written as:
[10]
This equilibrium reaction is then followed by the completion
of the step:
FeOOH 2+ ~ Fe > + HO2".
[11]
Similarly, step [7] can be considered as the sum of two
reactions) s First, is the equilibrium with hydrogen ion,
HO2- K~2=H* + O~- ,
[12]
followed by step completion,
O2 + Fe 3+ ~2~ O2(g) + Fe 2+ .
METALLURGICAL T R A N S A C T I O N S B
-(dmn2o,/dt) = 2(Kl~
[13]
v2 (m~,o,) (mFe3.)
-
~
/
[14]
(m.-)
The form of expression [14] is in agreement with the expression determined from the initial rate experiments carried out
in dilute sulfuric acid solutions, Eq. [3].
In contrast, for solutions with concentration ratios of
H202 less than 1.0. Barb, et al." expect the termination
step [9] to replace [8]. Considering [9] and once again assuming steady state concentrations for the same chain
carriers, the rate expression that can be derived from [4]
to [7], [9], and [10] to [13] is:
- ( d m m o f f d t ) = (2Kl~
- -
\
1'2 (mH2~
k9
/
(gt/Fe3")l'2 [15]
( m . - ) la
Experimental results from Barb, et al., n however, yielded
the following rate expression for solutions with concentration ratios of H202 to Fe 3. less than 1.0:
-(dmmoffdt)
-
HO2" + Fe > --~
k7 Fe > + H+ + O2(g)
Fe 3+ + H202(1)x,0=H+ + FeOOH2 +.
As pointed out by Walling and Weil, 25the equilibrium reactions, [10] and [12], should lead to an inverse order dependency on the concentration of hydrogen ion for the overall
steps [4] and [7], respectively.
For solutions with concentration ratios of H202 to Fe 3+
greater than 1.0, Barb, et al." assumed that the redox sequence involved only steps [4] through [8] by reasoning that
in these solutions the hydroxyl radical would react predominately with H202, reaction [6], rather than with ferrous
ion, reaction [9]. By assuming steady state concentrations
for the chain carriers Fe 2+, HO., and HO2" for steps [4]
through [8] and considering the expanded relations in [10]
to [13] for steps [4] and [7], the following overall rate expression can be derived:
-
=
k (mH2~
(tnFe'-)
[16]
(m.-)
The discrepancies between [15] and [16] are explained by
Walling and Well 25 as being caused by poor resolution in
determining the ratio of H202 to Fe > where the change in
mechanisms takes place. Barb, et al. u assumed the change
occurred cleanly at ratios equal to 1.0. However, as later
pointed out by Walling and Weil, 25 the mechanism change
probably does not take place at a single concentration ratio
for all solutions.
The lowest concentration ratio of H202 to Fe 3. at which
the H2OR decomposition rate could be measured in this study
was about 3.0. In the dilute sulfuric acid solutions used
in the experiments, approximately 5 pct of the total ferric
species are calculated to be free ferric ions at 25 ~
Ferric sulfate complexes are the dominant ferric species
(Figure 1). In order to lower the ratio of H202 to Fe 3+ to
below 1.0 in the experiments, either very low H202 concentrations or very high total ferric ion concentrations
would have to be used. Analytical difficulties in determining
changes in low H202 concentration made rate data unreliable under such conditions. Increasing total ferric ion
concentrations to values higher than those used in the experinaents would have exceeded the concentrations applicable to in situ uranium leaching. Also, in solutions with
high total ferric ion concentrations it was expected that the
H20,. decomposition rate would be too rapid to allow initial
V O L U M E 16B, JUNE 1 9 8 5 - - 185
rate methods to be used. However, it is expected that the rate
expression and rate constant determined here, Eq. [3],
should be applicable for most of the conditions that are
likely to be encountered during in situ uranium leaching
with acidic, H202 solutions. This experimentally determined rate law and rate constant should provide an improved means to assess quantitatively the effects of the
decomposition of H202 that would occur if dissolved iron is
present in leaching solutions, and may aid in the design of
more efficient lixiviant compositions. In addition, this rate
expression is in good agreement with the expression that can
be derived from the redox reaction sequence of Barb, et al. H
and Walling and Well, 25 Eq. [14]. This sequence was developed by these workers to describe the reaction path leading
to H202 decomposition in dilute perchloric acid solutions in
which the dominant ferric species is the free ferric ion. The
experimental work described here demonstrates that a similar reaction path may describe the decomposition of H202 in
dilute sulfuric acid solutions in which the ferric species are
expected to be present largely as ferric sulfate complexes
(Figure 1). It should be noted, though, that it was assumed
here that the active catalyzing species was the free ferric
ion, as suggested by Barb, e t a l . , H and the rate law is
written accordingly. It is possible that future work will
determine that the rate law should be written in terms of
the dominant ferric species, such as FeSO2. If this is the
case, then the rate constant measured here may be in error
by some constant factor that reflects the effects of ferric
speciation. However, the orders of the reaction rate with
respect to the reactant concentrations, as given in [3], should
still be correct.
manuscript. Thanks are also due to A. C. Lasaga and L. M.
Cathles for their help and comments. This research has been
supported in most part by the United States Bureau of Mines
Contract No. J0100065 to L. M. Cathles and H. L. Barnes.
Support was also received from the National Science Foundation, Geochemistry Program Grant No. EAR-8206777 to
H.L. Barnes, and from the Mineral Conservation Section,
Department of Geosciences, Pennsylvania State University.
REFERENCES
1.
2.
3.
4
5
6.
7
8
9
10
11.
12.
13.
14
15
IV.
SUMMARY
The rate expression describing the ferric ion catalyzed
decomposition of HzO 2 in dilute sulfuric acid solutions is
found to agree well with the proposed redox reaction sequence of Barb, et al. H and Walling and Well 25 for solutions
with concentration ratios of H202 to Fe 3+ greater than 1.0.
The measured rates of the ferric ion catalyzed decomposition of H202 indicate that this process can cause a significant
reduction in the H202 concentration in a relatively short
period of time (Figure 2). This reaction may be at least
partly responsible for the experimental and field observations of rapid loss of H202 during in situ uranium leaching
and should be taken into account when determining optimum lixiviant compositions.
ACKNOWLEDGMENTS
The author would like to thank primarily H.L. Barnes
and K.M. Krupka for their constructive reviews of this
186
VOLUME 16B, JUNE 1985
16
17
18
19,
20.
21
22
23
24
25.
J. B. Hlskey. Trans. lnstn Mm Metall. (Sect C Mineral Process
Extr Metall.). 1980, vol 89, pp 145-52.
L.E. Eary and L M Cathles: Metall. Trans B, 1983, vol. 14B,
pp. 325-34.
D. E Grandstaff. Econ Geol , 1976, vol 71, pp. 1493-506
A. R Amell and D. Langmulr' U.S. Bur. Mines Final Rep , Contract
No. H0272019, Dept. of the Interior, Washington. DC, 1978
P Gatichon, R S Schecter, A Cowle~. and M Breland. In Situ.
1977. vol. 2, pp. 125-46.
H C Granger and C G Warren Econ Geol 1969, vol 64,
pp 160-71.
R . L Reynolds and M B Goldhaber' Econ Geol 1983, vol. 78,
pp 160-71
B C Lawes In &tu, 1978, vol. 2. pp 75-92
D C Grant. U S Bur Mines Open File Rep. 52-79, Dept. of the
Interior. Washington. DC, 1979
L . M . Lttz' Mm. Eng., 1982, vol. 34, pp. 52-56
W G Barb, J. H Baxendale, P. George, and K R. Hargrave Trans
Faradm Soe , 1951, vol 47. pp 591-616.
D R Tweeton, G. R. Anderson, J K. Ahlness, O. M. Peterson, and
W. H Engelman. Soc. Mm Eng AIME, 1978, Preprmt No 70-43
C A Ehgwe, A E Torma, and F W DeVnes H)drometall., 1982,
vol 9, pp. 83-95
G M Elsenburg. lnd Eng Chem, Anal E d . 1943, vol 15,
pp 327-28
G. D Chnstlan Anah'twal Chemtstr). Xerox College Pubhshlng,
Waltham. MA, 1971, pp 351-52
H Tamura, K Goto. T Yotsoyanagl, and M Nagagama. Talanta,
1874. vol 21, pp. 314-18
A . C . Lasaga: Kmettcs of Geochemwal Processes, Revzews t n
Mineralogy, A C. Lasaga and R J Klrkpamck. eds,, Mineralogical
Society of Amenca, Washington, DC. 1981, ,~ol. 8, pp 1-67.
W. G Barb, J H. Baxendale, R George. and K. R Hargrave' Nature,
1949. vol 163, pp 692-94
L.M. Cathles and K J. Breen. U.S Bur Mines Final Rep , Grant
Nos. G5105007, G5115007, and Gl115427, Dept. of the Interior,
Washington, DC, 1983
R.M. Smith and A.E. Martell Crmcal Sta&ho Constants, Vol. 4.
lnorgamc Complexes, Plenum Press. New York, NY, 1976, 257 pp.
D D Wagman, W. H. Evans, V B. Parker, I Harlow, S M Bailey,
and R H Schumm Nat Bur. Stds Tech. Note 270-4, 1969
R S Sapleszko, R. C. Patel, and E. Matljevlc. J Phvs Chem., 1977,
vol 81, pp 1061-68.
A H Truesdell and B E Jones. J. Res U S Geol. Survey, 1974,
vol. 2, pp. 233-48
W. Stumm and J J Morgan: Aquauc Chemtstr~" An Introduction
Emphastzmg Chemical Equihbrta m Natural Waters, 2nd ed., John
Wdey and Sons, New York, NY, 1981, p. 135
C Walhng and T Well' Int J. Chem K m , 1974, vol VI,
pp. 507-16.
METALLURGICAL TRANSACTIONS B