Kinetics and Mechanism of the Ferrate Oxidation of Thiosulfate and Other Sulfur-Containing Species

Article abstract of DOI:10.1021/ic960480o

The kinetics of the reaction of ferrate, FeO₄^2-, with several sulfur-containing species in aqueous media have been investigated, and the results are reported. It was found that, when the reductant is in excess, ferrate rapidly oxidizes thiosulfate to sulfite, benzenesulfinate to benzenesulfonate, methionine to its corresponding sulfoxide, and dimethyl sulfoxide to dimethyl sulfone. The rate law for each reaction is first order with respect to each reactant and first order with respect to the hydrogen ion concentration. A mechanism for each oxidation reaction is discussed.

Full text of DOI:10.1021/ic960480o

Inorg. Chem. 1996, 35, 6795-6799
6795
Kinetics and Mechanism of the Ferrate Oxidation of Thiosulfate and Other
Sulfur-Containing Species
Michael D. Johnson* and John F. Read†
Department of Chemistry and Biochemistry, New Mexico State University,
Las Cruces, New Mexico 88003
ReceiVed May 1, 1996X
The kinetics of the reaction of ferrate, FeO42-, with several sulfur-containing species in aqueous media have been
investigated, and the results are reported. It was found that, when the reductant is in excess, ferrate rapidly
oxidizes thiosulfate to sulfite, benzenesulfinate to benzenesulfonate, methionine to its corresponding sulfoxide,
and dimethyl sulfoxide to dimethyl sulfone. The rate law for each reaction is first order with respect to each
reactant and first order with respect to the hydrogen ion concentration. A mechanism for each oxidation reaction
is discussed.
Introduction
has been recently reinvestigated by Johnson,8 and his results
Although the ferrate ion, FeO42-, has been known for over a
century, its chemistry remains relatively unexplored. In view
of its potential utility in environmental and synthetic applica-
tions, one goal of this laboratory is to develop the chemistry of
this unusual oxidation state of iron. We have begun a series of
studies on the redox kinetics of ferrate with a variety of
substrates.1
Ferrate is a tetrahedral ion, is isostructural with chromate or
manganate,2 and is a strong oxidant as well (0.9 and 1.9 V (vs
NHE) in basic and acidic media, respectively).3 The oxidations
of a few organic substrates have been reported, all from a
synthetic perspective. Although from its redox potentials and
structure one might anticipate its chemistry to be essentially
identical with that of permanganate or chromate, there are
significant differences in its reactivity. Aqueous potassium
ferrate oxidizes alcohols to aldehydes or ketones,4 thiols to
disulfides or sulfonic acids,5 aromatic amines to azo com-
pounds,6 and 1° aliphatic amines deaminate to form aldehydes.4
The products and yields of all these reactions depend upon
reaction times and conditions. Surprisingly, however, ferrate
is unreactive toward addition across double or triple bonds. Such
distinct reaction preferences should make ferrate an important
and selective oxidant for synthetic organic studies.
were different from those originally reported. In addition,
Johnson reported the oxidation of selenite and found evidence
for a quasi-stable ferrate/selenite-bridged intermediate. Bielski
and co-workers have reported the ferrate oxidation of amino
acids11 and a pulse-radiolytic study10 of the ferrate reduction
by the hydrated electron to form FeVaq. Most recently Johnson
and Hornstein have reported the oxidation of hydrazine and
monomethylhydrazine to molecular nitrogen:9,12
FeO42- + N2H4 f N2 + FeII
(1)
(2)
FeO42- + CH3N2H3 f N2 + CH3OH + FeII
For these reactions, diazene intermediates were postulated on
the basis of the reduction of an unsaturated carboxylic acid using
ferrate and hydrazine.
This paper reports the kinetics for the oxidation of thiosulfate,
benzenesulfinate, dimethyl sulfoxide, and methionine with
ferrate in basic media. Mechanisms for these oxidation reactions
are also presented.
Experimental Section
Materials. Sodium thiosulfate, sodium benzenesulfinate, dimethyl
sulfoxide, and methionine were purchased from Aldrich Chemicals and
used without further purification. The thiosulfate was standardized by
iodometric titration. All other chemicals were of reagent grade.
Potassium ferrate was prepared by the method reported by Thomp-
son, Ockermann, and Schreyer.13 The crude product was recrystallized
and the purity checked by spectrophotometric analysis at 505 nm (ꢀ )
1100 M-1 cm-1).14 Purities up to 95% were obtained, and the measured
rates were found to be independent of the sample purity.
Stock solutions of the reductants were prepared using deionized
water. Ionic strength was maintained with sodium perchlorate, and
sodium phosphate buffers were used to control pH. The phosphate
ion also prevented precipitation of iron hydroxides, which are produced
To date, only a handful of kinetic studies of ferrate oxidations
have appeared in the literature.7-11 Goff and Murmann7
reported the first kinetic investigation in 1974 for the ferrate
oxidation of hydrogen peroxide and sulfite. Their sulfite work
* To whom correspondence should be addressed.
† On sabbatical leave from Mount Allison University, Sackville, New
Brunswick, Canada.
X Abstract published in AdVance ACS Abstracts, October 1, 1996.
(1) Johnson, M. D.; Hornstein, B. J. J. Chem. Soc., Chem. Commun. 1996,
965.
(2) Hoppe, M. L.; Schlemper, E. O.; Murmann, R. K. Acta Crystallogr.
1982, B38, 2237.
(3) Wood, R. H. J. Am. Chem. Soc. 1958, 80, 2038.
(4) Audette, R. J.; Quail, J. W.; Smith, P. J. Tetrahedron Lett. 1971, 279.
(5) Bartzatt, R. L.; Carr, J. Transition Met. Chem. 1986, 11, 116.
(6) Firouzabadi, H.; Ghaderi, E. Tetrahedron Lett. 1978, 839. Firouzabadi,
H.; Mohajer, D.; Entezari-Moghadam, M. Bull. Chem. Soc. Jpn. 1988,
61, 2185.
(7) Goff, H.; Murmann, R. K. J. Am. Chem. Soc. 1971, 93, 6058.
(8) Johnson, M. D.; Bernard, J. Inorg. Chem. 1992, 31, 5140.
(9) Johnson, M. D.; Hornstein, B. Inorg. Chim. Acta 1994, 225, 145.
(10) (a) Bielski, B. H. J.; Thomas, M. J. Inorg. Chem. 1989, 28, 3947.
Bielski, B. H. J.; Thomas, M. J. J. Am. Chem. Soc. 1987, 109, 7761.
(b) Bielski, B. H. J.; Sharma, V. K. Inorg. Chem. 1991, 30, 4306.
(11) Carr, J. D.; Kelter, P. B.; Tabatabai, A.; Spichal, D.; Erickeson, J.;
McLaughlin, C. W. Proc. Conf. Water Chlorination Chem. EnViron.
Impact Health Eff. 1985, 1285.
(12) Johnson, M. D.; Hornstein, B.; Wingo, R. Proceedings of the 5th
Annual Conference on Waste Technology DeVelopment; WERC
Administrative Office: Las Cruces, NM, 1995; p 358.
(13) Thompson, G. W.; Ockerman, L. T.; Schreyer, J. M. J. Am. Chem.
Soc. 1951, 73, 1379.
(14) Carrington, A.; Schonland, D.; Ingram, D. J. E.; Symons, M. C. R. J.
Chem. Soc. A 1956, 4710.
S0020-1669(96)00480-6 CCC: $12.00 © 1996 American Chemical Society

6796 Inorganic Chemistry, Vol. 35, No. 23, 1996
Johnson and Read
during the reduction process. Solutions were bubbled with argon gas
for 15 min to remove dissolved oxygen from solution prior to kinetic
measurements.
Kinetic Measurements. The reaction rates were measured spec-
trophotometrically using a D110 Durrum spectrophotometer with an
OLIS computer interface. The reactions were monitored at 505 nm
under pseudo-first-order conditions with the reductant in at least 10-
fold excess. Typical reductant ranges were between 0.05 and 0.002
M. Absorbance changes were fitted using an OLIS data fitting routine.
Each rate constant is an average of three to five trials. The temperature
was controlled to within 0.1 °C. For determination of the activation
parameters, five to seven temperatures ranging from 12 to 34.0 °C were
used. All studies were conducted using 0.05 M phosphate buffers (I
) 1.0 M NaClO4) to control pH and to prevent any precipitation of
iron hydroxide/oxides by complexation of the iron final product. No
effect on the observed rate constants was found between the phosphate
concentrations of 0.01 and 0.05 M. The ferrate concentrations were
between 1 × 10-4 and 5 × 10-4 M.
The rapid-scan spectrophotometer used was a Harrick rapid-scan
monochromator interfaced with a Durrum stopped-flow spectrometer.
The data collection and manipulation system was designed by the OLIS
Corporation.
IR measurements were made using a Perkin-Elmer 1720X FT-IR
spectrometer. UV-vis measurements were made on an HP diode array
spectrophotometer.
Figure 1. Typical plot of kobs versus [thiosulfate]. Conditions: pH
8.5, T ) 25.0 °C, I ) 1.0 M (NaClO4), 0.05 M (Na phosphate) buffer.
reaction were undertaken. Spectrophotometric quantification
of the malachite green reaction with the sulfite produced also
supports complete conversion of thiosulfate to sulfite. The
stoichiometry is pH independent between 7.5 and 11.0. On the
basis of the above information, the following reaction may be
written:
Results and Discussion
Oxidation of Thiosulfate. Stoichiometry. Thiosulfate
rapidly reduces ferrate in aqueous media, and when the reduction
is carried out under phosphate-buffered (basic) aerobic condi-
tions, no ferric oxy hydroxides precipitate from solution. The
presence of Fe(III) was demonstrated by the addition of
potassium thiocyanate to the spent solutions to form the red
ferric thiocyanate complex. When oxygen was excluded, Fe-
(III) was again the only iron species found.
14H+ + 4FeO42- + 3S2O32- f
4Fe3+ + 6SO32- + 7H2O (3)
Tetrathionate and dithionate are also logical oxidation products
on the basis of other reactions of thiosulfate with strong oxidants.
If tetrathionate were produced, however, the stoichiometry
would be 3:1, and if dithionate were formed, a stoichiometry
of 1:2 would be observed, both of which are well outside the
experimental error of this study. In fact, the 1.3:1 stoichiometry
indicates that these are not produced to any significant extent
in side reactions. Qualitative tests show that tetrathionate is
only slowly oxidized by ferrate (minutes) and dithionate is not
oxidized at all. As far as we are aware, this represents the only
example of direct thiosulfate oxidation to sulfite.
Kinetics. The reaction kinetics were studied anaerobically,
under pseudo-first-order conditions. Absorbance decreases for
the loss of ferrate fit nicely to a single-exponential decay. Figure
1 shows a typical plot of the observed rate constants versus the
thiosulfate concentration at pH 8.5. The linear dependence
indicates a first-order term in thiosulfate for the rate law, and
the slope is the observed second-order rate constant. Unlike
other ferrate oxidations, no positive intercept (approximately
0.1-0.2 s-1 due to ferrate decomposition in the absence of
reductant at this pH) was observed. This is because the thio-
sulfate reaction is very rapid compared with the rate of ferrate
decomposition. The plots of kobs versus [reductant], however,
do show significant y intercepts at pHs approaching the pKa of
ferrate. This is due to the competitive pathway for the oxidation
of water and makes measurement of reaction rates through the
pKa of ferrate difficult, particularly for slow redox reactions.7,8
Hydrogen ion concentrations were kept constant with the use
of phosphate buffers. No significant differences in rate were
observed upon changing phosphate buffer concentrations.
Figure 2 shows a plot of the observed second-order rate
constants versus the hydrogen ion concentration. The linear
dependence on H+ with a zero y intercept indicates a first-order
As a check for the production of sulfate as the final product,
barium chloride was added to the spent, filtered solutions where
thiosulfate was kept in excess (no phosphate buffer used).
Under the experimental concentrations where barium sulfite
would remain in solution but barium sulfate would precipitate,
no precipitate of barium sulfate was observed upon oxidation
of thiosulfate with ferrate. On the basis of earlier work,8 the
reaction with sulfite will be slower than the thiosulfate reaction
under the conditions of this study; therefore, conversion of the
sulfite to sulfate would not be observed. When ferrate is used
in slight excess, however, a white precipitate is immediately
formed upon treatment of the spent solution with aqueous
barium chloride. This suggests the subsequent oxidation of
sulfite to sulfate.
To confirm the presence of sulfite, malachite green was added
to the reaction mixtures. Sulfite preferentially decolorizes
malachite green by formation of an addition compound.15 Tests
of spent ferrate/thiosulfate reaction mixtures immediately gave
a positive spot test for the presence of sulfite. As a check for
possible interferences, ferric ion, thiosulfate, sulfate, nor
tetrathionate decolorized malachite green.
A [1.3((0.2)]:1 stoichiometry was determined for ferrate to
thiosulfate by rapid mixing of known ratios of thiosulfate and
ferrate until the purple color of ferrate no longer rapidly
disappears and sulfate does not appear, since sulfite is also
oxidized by ferrate. Such strictures provide difficulties in
obtaining a good end point in the “titration” method; therefore,
attempts to quantify the amount of sulfite produced by this
(15) Vogel, A. QualitatiVe Inorganic Analysis; Wiley: New York, 1987
(revised by G. Svehla).
(16) (a) Howelett, K. E.; Wedzicha, B. L. Inorg. Chim. Acta 1976, 18, 33.
(b) Baldea, I.; Niac, G. Inorg. Chem. 1970, 9, 110. (c) Kirschenbaum,
L. J.; Rush, J. D. Inorg. Chem. 1983, 22, 3304.

Ferrate Oxidation of Thiosulfate
Inorganic Chemistry, Vol. 35, No. 23, 1996 6797
From the above mechanism the following expression may be
written for k:
k ) kH[H+]/([H+] + Ka)
(5)
Under the conditions that the rates were measured (between
pH 8 and 11), [H+] is less than the ferrate Ka (pKa ) 7.3)11)
and eq 5 reduces to
k ) kHKa-1[H+]
(6)
The slope, k0, of the plot of k versus [H+] is a composite of kH
and Ka
k0 ) kHKa-1
(7)
and the bimolecular rate constant kH at 25 °C may be calculated
as 36 000 M-1 s-1
.
This two-electron mechanism is in contrast with the ferrate
oxidation of cysteine, where Bielski proposed that the oxidation
occurs via a one-electron pathway to produce Fe(V) and cysteine
radicals.10b The latter subsequently dimerize to produce cystine.
If thiosulfate were oxidized initially via a one-electron process
to produce S2O3-, dimerization to produce tetrathionate might
be anticipated. Since our studies indicate that tetrathionate is
a minor product at best, a one-electron pathway appears unlikely.
The proposed two-electron oxidation does conform with other
ferrate oxidations that we have studied. Apparently, when
provided a choice, ferrate will take the two-electron pathway
preferentially over a one-electron pathway.
Figure 2. Hydrogen ion dependence on the rate of the thiosulfate/
ferrate reaction. Conditions: T ) 25.0 °C, I ) 1.0 M (NaClO4), 0.05
M (Na phosphate) buffer.
term in H+ and no acid-independent term. An observed rate
law may be written as
2-
-d[Fe(VI)]/dt ) k0[H+][FeO42-][S2O3
]
(4)
The value for k0 is (7.1 ( 0.2) × 1011 M-2 s-1 at 25 °C and I
) 1.0 M (NaClO4). This value is similar to that for the ferrate
oxidation of sulfite, 1.8 × 1012 M-2 s-1 8
. A study of the tem-
Using pulse radiolysis, Bielski has established that Fe(IV) is
a strong oxidant and appears to react at least 1 order of magni-
tude faster than ferrate with reductants. This supports the as-
sumption that the steps in the mechanism involving the oxi-
dations by Fe(IV) are not rate-determining.10 The steps involv-
ing oxidation of OSSO32- by Fe(IV) (eqs I.5 and I.6 in Scheme
1) may be of minimal importance since they represent two
highly reactive intermediates. They are included, however, for
completeness.
perature dependence provides the activation parameters 18.0 (
3.5 kJ mol-1 and 42 ( 20 J mol-1 K-1 for ∆H* and ∆S*,
respectively.
Mechanism. The rate law shows a single proton-dependent
pathway for the oxidation of thiosulfate by ferrate ion. Since
the protonation of thiosulfate does not occur in the pH range
studied, it is safe to assume the dependence is due to the pro-
tonation of ferrate itself. Since ferrate is a stronger oxidant upon
protonation, the reaction rate would be expected to increase.
On the basis of its reactions with sulfite, where ferrate acts
as a two-electron oxidant and as an oxygen transfer agent, the
reaction mechanism given in Scheme 1 for the oxidation of
thiosulfate is proposed. The charges are not balanced, since
the nature of the Fe(IV) moiety has yet to be characterized,
although its existence in aqueous media has been established
by Bielski and co-workers.10a
The sulfur oxidation intermediate, OSSO32-, proposed in this
mechanism has been suggested in the monoperoxysulfate
oxidation of coordinated thiosulfate:17
-
-
HSO5
HSO5
[Co(NH3)5S2O3]+
8 [Co(NH3)5SOSO3]+
[Co(NH3)5SO2SO3]+ (8)
8
Scheme 1
This “isodithionite” species is rapidly oxidized to metabi-
sulfite by monoperoxysulfate. Qualitative experiments with
ferrate show a similar reaction pattern with coordinated thio-
sulfate.
Ka
\
HFeO4- y
z FeO42- + H+
(I.1)
kH
HFeO4- + S2O32- 98 FeIV + OSSO32- + H+ (I.2)
HFeO4- + OSSO32- f FeIV + S2O52- + H+ (I.3)
During the oxidation of sulfite to sulfate, Murmann used 18
O
labeling to demonstrate that an oxygen in the first coordination
sphere of ferrate is incorporated in the sulfate reaction product.7
He proposed that an Fe-O-S bond was formed during this
process. This observation was also used by Johnson to explain
the kinetics for the oxidation of sulfite and selenite.8 Since an
FeO42- + OSSO32- f FeIV + S2O5
FeIV + S2O32- f FeII + OSSO3
FeIV + OSSO32- f FeII + S2O5
(I.4)
(I.5)
(I.6)
(I.7)
2-
2-
2-
-
outer-sphere mechanism would likely produce S2O3 radicals,
we suggest that a ferrate/thiosulfate ester is a probable inter-
mediate in this reaction:
FeIV + FeII f 2FeIII
Kd
\
H2O + S2O52- y
z 2SO32- + 2H+
(I.8)
(17) Johnson, M. D.; Balahura, R. J. Inorg. Chem. 1988, 27, 3104.

6798 Inorganic Chemistry, Vol. 35, No. 23, 1996
Johnson and Read
Such an intermediate would be helpful to understand the
oxidation of thiosulfate (eq I.2 in Scheme 1) by hydrogen ferrate.
Adducts have been proposed for the two-electron oxidation of
thiosulfate by Ag(III) and CrO42-, except that these species have
an M-S bond.16b,c Although the thiosulfate complexes were
spectrally observable with these oxidants, attempts to observe
a ferrate/thiosulfate adduct using rapid-scanning spectroscopy
(at pH 9 between 300 and 650 nm within 11 ms of mixing)
were unsuccessful. This may be due to a combination of small
equilibrium constant coupled with rapid redox reactions.
Oxidation of Other Sulfur Compounds. In order to
examine ferrate/sulfur redox reactions further, where dimeriza-
tion of sulfur products should not be important, and to provide
fundamental information regarding the relative rates of sulfur
oxidation, the oxidation of benzenesulfinate, methionine, and
dimethyl sulfoxide with ferrate was studied. Ferrate solutions
are rapidly decolorized upon addition of excess reductant under
conditions similar to those for thiosulfate. The reaction stoi-
chiometry for benzenesulfinate, methionine, and dimethyl
sulfoxide with ferrate are (1.5 ( 0.1):1, (1.3 ( 0.1):1, and (1.5
( 0.1):1, respectively, at pH 9.0. Using the same approach as
described for thiosulfate, the final iron product was shown to
be Fe(III).
The iron(III) in the spent solutions was removed using ion
exchange chromatography (SP-Sephadex). For benzenesulfi-
nate, the UV spectrum of the final organic product was recorded
in the chromatographed aqueous solution. For methionine and
DMSO, the solutions were extracted with methylene chloride,
the extracts evaporated to dryness, and IR spectra obtained of
the final product. These techniques showed that the final
organic oxidation products correspond to addition of an oxygen
to the sulfur atom, which agrees with a preliminary study of
ferrate oxidation of other sulfur compounds by Carr and Barzatt.5
On the basis of these two observations, an overall reaction for
each reductant may be written as
Figure 3. Hydrogen ion dependence on the rate of the S/ferrate
reactions. Conditions: T ) 25.0 °C, I ) 1.0 M (NaClO4), 0.05 M (Na
phosphate) buffer.
Table 1. Kinetic and Thermodynamic Data for Ferrate-Sulfur
Oxidations
k0
kH
∆H*
∆S*
substrate
thiosulfate
(M-1 s-1
)
(M-1 s-1
)
(kJ/mol) (J/(mol K))
(7.1 ( 0.2) × 1011 3.6 × 104 18.0 ( 3.5 42 ( 20
benzenesulfinate (1.3 ( 0.2) × 1011 7.0 × 103 20.6 ( 2.6 33 ( 15
methionine
(1.4 ( 0.1) ( 1011 6.5 × 103 34.0 ( 2.7 84 ( 9
dimethyl sulfoxide (2.2 ( 0.1) × 109 1.1 × 102 41.0 ( 1.8 55 ( 10
pH range studied, the acid dependence must be due to the
protonation of ferrate itself. The general reaction mechanism
given by Scheme 2 is proposed. Step II.2 in Scheme 2 is
probably similar to that for thiosulfate and involves an oxygen
transfer. Using the same arguments as for thiosulfate, the
bimolecular rate constant kH for each reductant may be
calculated (see Table 1).
10H+ + 2FeO42- + 3PhSO2- f
2Fe3+ + 3PhSO3- + 5H2O (9)
Scheme 2
Ka
\
HFeO4- y
z FeO42- + H+
(II.1)
(II.2)
10H+ + 2FeO42- + 3CH3SMet f
2Fe3+ + 3CH3S(O)Met + 5H2O (10)
kH
HFeO4- + S
9
8 FeIV + S(O) + H+
FeIV + S f FeII + S(O)
FeII + FeIV f 2FeIII
(II.3)
(II.4)
10H+ + 2FeO42- + 3CH3S(O)CH3 f
2Fe3+ + 3CH3S(O)2CH3 + 5H2O (11)
Conclusions
The kinetics were studied under conditions identical with
those for the thiosulfate and show a simple first-order depen-
dence on H+ with no acid-independent term (Figure 3). The
following general rate law may be written, where S represents
the sulfur reductant.
This is the first complete study of the oxidation of sulfur-
containing compounds by ferrate. When the reductant is present
in excess, thiosulfate is oxidized to sulfite, benzenesulfinate is
oxidized to benzenesulfonate, methionine is converted to its
corresponding sulfoxide, and dimethyl sulfoxide is oxidized to
dimethyl sulfone. Each oxidation is presumed to occur via a
two-electron step and involve oxygen transfer.7 This is
consistent with other ferrate oxidations1,9,11,18 and supports an
emerging theme for ferrate; that is, it preferentially acts as a
-d[Fe(VI)]/dt ) k0[H+][FeO42-][S]
(12)
Table 1 shows k0 values for the three reductants at 25 °C
and I ) 1.0 M (NaClO4). The temperature dependence of this
reaction was studied to obtain ∆H* and ∆S* for k0, and these
values are also presented in Table 1. Since benzenesulfinate,
methionine and dimethyl sulfoxide are not protonated in the
(18) We have studied the oxidation of hydroxylamine and substituted hy-
droxylamines: Johnson, M. D.; Hornstein, B. J. To be submitted for
publication.

Ferrate Oxidation of Thiosulfate
Inorganic Chemistry, Vol. 35, No. 23, 1996 6799
two-electron oxidant when either a one- or two-electron oxida-
tion pathway is available.
fashion, the SO3 moiety attached to the S(-II) atom in
thiosulfate also decreases the reaction rate compared to HS-.
The order of oxidation is thiosulfate > benzenesulfinate ≈
methionine > dimethyl sulfoxide. For thiosulfate, the terminal
S(-II) is presumably the site for oxidation, since it would be
expected to be most readily oxidized. This is supported by
preliminary ferrate oxidation studies of HS-, where the rates
are too fast to measure with stopped-flow techniques.19 The
slower oxidation of dimethyl sulfoxide compared to methionine
is most likely due to the presence of an oxygen on the sulfur
atom and its electron-withdrawing effects.20 In an analogous
Unfortunately, there appears to be little consistency in the
activation parameters for the oxidations reported herein. Due
to the paucity of thermodynamic data for ferrate oxidations, little
can be concluded about these parameters at this time. This study
includes them for completeness and to build up the available
information regarding ferrate oxidations.
Acknowledgment. We thank the Petroleum Research Fund
for their support of this work. We also thank Dr. Robert
Balahura and Dr. David Stanbury for their helpful comments
during the preparation of this paper.
(19) Johnson, M. D.; Read, J. Work in progress.
(20) Deutsch, E.; Root, M.; Nosco, D. L. AdV. Inorg. Bioinorg. Mech. 1982,
1, 269.
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