The Kinetics and Mechanism of the Ferrate(VI) Oxidation of Hydroxylamines

Article abstract of DOI:10.1021/ic020705x

Aqueous solutions of potassium ferrate(VI) cleanly and rapidly oxidize hydroxylamine to nitrous oxide, N-methylhydroxylamine to nitrosomethane, N-phenylhydroxylamine to nitrosobenzene, and O-methylhydroxylamine to methanol and nitrogen. The kinetics show first-order behavior with respect to each reactant and a two term component representing acid dependent and independent pathways. A general mechanism involving intermediate formation coupled with a two-electron oxidation is proposed.

Full text of DOI:10.1021/ic020705x

Inorg. Chem. 2003, 42, 6923−6928
The Kinetics and Mechanism of the Ferrate(VI) Oxidation of
Hydroxylamines
Michael D. Johnson* and Brooks J. Hornstein
Department of Chemistry and Biochemistry, New Mexico State UniVersity,
Las Cruces, New Mexico 88003
Received December 4, 2002
Aqueous solutions of potassium ferrate(VI) cleanly and rapidly oxidize hydroxylamine to nitrous oxide,
N-methylhydroxylamine to nitrosomethane, N-phenylhydroxylamine to nitrosobenzene, and O-methylhydroxylamine
to methanol and nitrogen. The kinetics show first-order behavior with respect to each reactant and a two term
component representing acid dependent and independent pathways. A general mechanism involving intermediate
formation coupled with a two-electron oxidation is proposed.
Introduction
Rapid coupling of these radicals to produce molecular
nitrogen was observed in reactions of hydroxylamine with
the aforementioned oxidants.
Hydroxylamines play a significant role in modern indus-
trial chemistry due to their wide range of chemical properties.
These include changes in reactivity with pH, formation of
N and O bonded compounds, and solubility in both aqueous
and nonaqueous solvents.1 This diversity of properties and
uses makes the study of hydroxylamine reactions an impor-
tant area of investigation.
The oxidation of hydroxylamine, NH2OH, is complex
since it can function as a one or two-electron donor and form
a variety of products. A one-electron pathway has been
suggested for the reductions of Pt(IV),2 Ce(IV),3 and V(V)4,
and the mechanism is thought to involve initial complex
formation followed by rate determining dissociation to
product NH2O‚.
2NH2O‚ f N2 + 2H2O
When the coupling rate is slower than the oxidation rate of
the radical, alternate products are observed. For example,
the radical is quantitatively oxidized by an excess of Fe(III)
in acid to NOH which then rapidly dimerizes to form N2O.5
NH2O‚ + Fe3+ f NOH + Fe2+ + H+
2NOH f N2O + H2O
Mononitrogen species may also be produced from initial one-
electron oxidations of hydroxylamine. For example, the
oxidation of hydroxylamine by Mn(III)6 or Ag(II)7 in acidic
solution produces only nitrate. In this case, the rapid
oxidation of H2NO‚ surpasses NOH coupling thereby sup-
pressing the dinitrogen pathway as shown here.
NH2OH + Mn+ f M-NH2OHn+
M-NH2OHn+ f NH2O‚ + M(n-1)+ + H+
This radical is subject to different chemical fates depending
on the oxidant and reaction conditions.
M
n+ + NH3OH+ f M(n-1)+ + NH2O‚ + 2H+
k
* To whom correspondence should be addressed. E-mail: johnson@
nmsu.edu. Phone: (505) 646-3627. Fax: (505) 646-2394.
(1) Encyclopedia of Industrial Chemical Analysis; Wiley: New York,
1971; Vol. 14. ComprehensiVe Inorganic Chemistry; Pergamon: New
York, 1973; Vol. 2. Cuisia, K. G.; Hwa, C. M.; Jacob, J. T.; Salutsky,
M. L. U.S. Patent 4 067 690; Chem. Abstr. 1978, 88, 141457h.
Takatomi, H.; Yamauchi, S.; Ogawa, K. Japanese Patent 02 233 503;
Chem. Abstr. 1991, 114, 65261c. Wilson, W. L. Br. Patent 739 144;
Chem. Abstr. 1971, 75, 65754s.
(2) Gupta, K.; Sen, P. K. J. Inorg. Nucl. Chem. 1977, 39, 1651.
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3456.
(4) Fattah, A.; Mazer, J.; Wells, C. F. J. Chem. Soc., Dalton Trans. 1980,
1532.
5Mn+ + 2H2O + NH2O‚ f 5M(n-1)+ + NO3- + 6H+ rapid
Alternatively, the oxidation of hydroxylamine can take
place via a two-electron pathway. Margerum and co-workers
have shown hydroxylamine oxidation by halogens to occur
(5) Butler, J. H.; Gordon, L. I. Inorg. Chem. 1986, 25, 4573.
(6) Davies, G.; Kustin, K. Inorg. Chem. 1963, 8, 484.
(7) Honig, D. S.; Kustin, K.; Martin, J. F. Inorg. Chem. 1978, 17, 1895.
10.1021/ic020705x CCC: $25.00 © 2003 American Chemical Society
Published on Web 09/19/2003
Inorganic Chemistry, Vol. 42, No. 21, 2003 6923

Johnson and Hornstein
through two-electron, halogen transfer steps (X ) Br, Cl, I).8
X2 + NH2OH a X2NH2OH
Ferrate is a tetrahedral ion, isostructural with chromate or
manganate.13 Its reduction potentials are 0.9 and 1.9 V
(versus NHE) in base and acid, respectively.14 Although one
might expect its reactivity to resemble that of chromate or
manganate, significant differences are emerging. Potassium
ferrate oxidizes alcohols to aldehydes or ketones,15 thiols to
disulfides or sulfonic acids, and arylamines to azo or nitro
compounds16-18 and deaminates primary alkylamines to form
aldehydes.18 Surprisingly, under no conditions does ferrate
add to double or triple carbon-carbon bonds. An advantage
ferrate has over many other oxidants is that the final iron
product is rust, which is easily separated from the desired
products and disposed of safely. With such advantages over
other transition metal oxidants, ferrate has the potential to
become an important “green” reagent in organic oxidation
studies.
To date, relatively few kinetic studies of ferrate oxidations
have appeared in the literature. In 1974, Goff and Murmann
published the first kinetic study for the ferrate oxidation of
hydrogen peroxide and sulfite along with an oxygen ex-
change study.19 The oxidation was reinvestigated by Johnson
and found to be different from the original report. Read and
Sharma have examined the oxidation of several sulfur centers
by ferrate.20 Bielski has reported the oxidation of amino acids
by ferrate21 occurs via one-electron radical pathways. He also
proposed that the oxidation of phenol by ferrate occurs by a
one-electron pathway to produce Fe(V) and phenoxyl radi-
cal.22 In this system, Fe(V) rapidly undergoes a two-electron
transfer to form an inner-sphere Fe(III) complex. The exact
mechanism by which this occurs is not known; however, it
is thought to involve either inner-sphere substitution or an
electrophilic addition.
X2NH2OH f X- + XNHOH + H+
XNHOH f X- + NOH + H+
2NOH f N2O + H2O
The proposed mechanism involves rate determining forma-
tion of a nitrogen-halogen complex followed by electron
transfer and halogen loss. In excess hydroxylamine, NOH
is produced and undergoes rapid dimerization to form N2O
and H2O, k ) (4.5 ( 2.7) × 109 M-1s-1.9 With excess
bromine, HONO was observed which led Margerum to
propose a mechanism that includes two two-electron oxida-
tions of N(-I) to eventually form N(III).8
Br2 + NH2OH h BrNHOH + HBr
BrNHOH + Br2 f Br2NOH + Br- + H+
Br2NOH f BrNO + H+ + Br-
BrNO + H2O f Br- + HONO + H+
Haight proposed that the Cr(VI) oxidation of hydroxy-
lamine also proceeds by two-electron steps.10 Initial formation
of a H2NOCrO3H ester is followed by an intramolecular two-
electron redox reaction to produce Cr(IV) and NOH. This
mechanism dominates when the reaction is carried out with
excess reductant. In contrast, excess Cr(VI) rapidly oxidizes
NOH to HNO2 before dimerization can take place.
High oxidation state complexes, in particular those that
involve Fe(IV) and Fe(V), are proposed to play a significant
role in biological systems.11 To further our understanding
of high oxidation state iron chemistry, the oxidation of
hydroxylamine and substituted hydroxylamines with potas-
sium ferrate, K2FeO4, was examined.
In contrast to the one-electron mechanisms suggested by
Bielski, Johnson and Lee have proposed two-electron reduc-
tions of ferrate. Johnson favored a quasistable ferrate/
substrate bridged intermediate for the reaction with selenite
and sulfite24 as well as for thiosulfate.19 The proposed bridged
species contains an ester linked, Fe-O-S moiety (S )
substrate) accompanied by consecutive two-electron reduc-
Although the ferrate ion, FeO42-, has been known for over
a century, its chemistry remains relatively unexplored. With
recent developments in the synthesis of potassium ferrate,12
its applications in the areas of environmental and organic
chemistry are likely to increase.
(13) Hoppe, M. L.; Schlember, E. O.; Murmann, R. K. Acta Crystallogr.
1982, B38, 2237.
(14) Wood, R. H. J. Am. Chem. Soc. 1958, 80, 2038.
(15) Audette, R. J.; Quail, J. W.; Smith, P. J. Tetrahedron Lett. 1971, 270.
(16) Bartzatt, R. L.; Carr, J. Transition Met. Chem. 1986, 11, 116.
(17) Johnson, M. D.; Hornstein, B. J. J. Chem. Soc., Chem. Commun. 1996,
965. Johnson, M. D.; Hornstein, B. J. Submitted for publication.
(18) Fibrouzabadi, H.; Ghaderi, E. Tetrahedron Lett. 1978, 839.
(19) Goff, H.; Murmann, R. K. J. Am. Chem. Soc. 1971, 93, 6586. Johnson,
M. D.; Bernard, J. Inorg. Chem. 1992, 31, 5140.
(8) (a) Liu, R. M.; McDonald, M. R.; Margerum, D. W. Inorg. Chem.
1995, 34, 6093. (b) Beckwith, R. C.; Cooper, J. N.; Margerum, D. W.
Inorg. Chem. 1994, 33, 5144.
(9) Bazylinski, D. A.; Hollcher, T. C. Inorg. Chem. 1985, 24, 4285.
(10) Haight, G. P.; Scott, R. A.; Cooper, J. N. J. Am. Chem. 1974, 96, 4126.
(11) Oxidases and Related Redox Systems; King, T. E., Mason, H. S.,
Morrison, M., Eds.; Liss: New York, 1988. Fox, B. G.; Froland, W.
A.; Dege, J. E.; Lipscomb, J. D. J. Biol. Chem. 1989, 264, 1023. Fox,
B. G.; Borneman, J. G.; Wackett, L. P.; Lipscomb, J. D. Biochemistry
1990, 29, 6419. Elgren, T. E.; Lynch, J. B.; Juarez-Garcia, C.; Mu¨nck,
E.; Sjo¨berg, B. M.; Que, L. J. Biol. Chem. 1991, 266, 19265. The
Enzymes; Sigman, D. S., Ed.; Academic Press: San Diego, CA, 1992;
Vol. 20. Cytochrome P-450. Structure, Mechanism and Biochemistry;
Ortiz de Montellano, P. R., Ed.; Plenum Press: New York, 1995.
(12) Caddick, S.; Murtagh, L.; Weaving, R. Tetrahedron 2000, 56, 9367.
Lapicque, F.; Valentin, G. Electr. Commun. 2002, 109, 67. Licht, S.;
Tel-Vered, R.; Halperin, L. Electr. Commun. 2002, 4, 933. Johnson,
M. D. May 5, 1988, U.S. Patent 5,746,9984.
(20) Read, J. F.; Boucher, K. D.; Mehlman, S. A.; Watson, K. J. Inorg.
Chim. Acta 1998, 267, 159. Read, J. F.; Adams, E. K.; Gass, H. J.;
Shea, S. E.; Theirault, A. Inorg. Chem. Acta 1998, 281, 43. Read, J.
F.; John, J.; MacPherson, J.; Schaubel, C.; Theriault, A. Inorg. Chem.
Acta 2001, 315, 96. Sharma, V. K.; Rendon, R. A.; Millero, F. J.;
Vazquez, F. G. Mar. Chem. 2000, 70, 235. Sharma, V. K.; Smith, J.
O.; Millero, F. J. EnViron. Sci. Technol. 1997, 31, 2486. Sharma, V.
K.; Burnett, C. R.; O’Connor, D. B.; Cabelli, D. EnViron. Sci. Technol.
2002, 36, 4182.
(21) Bielski, B. H. J.; Sharma, V. K. J. Am. Chem. Soc. 1991, 30, 4306.
(22) Bielski, B. H. J.; Rush, J. D. Free Radical Res. 1995, 22, 571.
(23) Hornstein, B. J. Ph.D. Dissertation, New Mexico State University,
1999.
(24) Organic Syntheses; Wiley & Sons: New York, 1941; Collect. Vol. I,
p 445.
6924 Inorganic Chemistry, Vol. 42, No. 21, 2003

Ferrate(VI) Oxidation of Hydroxylamines
were carried out using a Nuclide RMS-16 isotope ratio mass
spectrometer at the University of MissourisColumbia with the aid
of Professor R. Kent Murmann.
tions of Fe(VI) that results in Fe(II). Direct oxygen transfer
was observed by oxygen tracer studies thereby supporting
this mechanism. The reduction of ferrate in two-electron steps
is also postulated by Lee in the oxidation of secondary
alcohols.29 As opposed to an ester bridge between ferrate
and substrate, Lee suggests a direct inner-sphere attack of
the alcohol on the iron center prior to electron transfer to
form a five-coordinate intermediate.
In recent investigations, Johnson and Hornstein studied
the oxidation of aniline with ferrate. New kinetic and spectral
evidence suggests a key step in the reaction mechanism is
the formation of an iron-imido complex that either yields
nitrosobenzene or reacts with another aniline to produce cis-
azobenzene.23 To shed light on these complex mechanisms,
we have studied the reaction of the ferrate ion with
hydroxylamines.
Results and Discussion
Stoichiometry. Hydroxylamine, N-methylhydroxylamine,
O-methylhydroxylamine, and N-phenylhydroxylamine all
rapidly (<20 s) bleach aqueous solutions of ferrate over the
pH range 8-11. When the reaction is carried out in the
presence of oxygen and no buffer, the precipitation of iron-
(III) hydroxide is observed. In contrast, when oxygen is
excluded, iron(II) is formed as shown by the formation of
an intense red color upon acidification immediately after the
reaction is over and addition of 1,10-phenanthroline. Further
evidence for the sole production of iron(II) in the absence
of oxygen was obtained by adding potassium thiocyanate to
the spent reaction mixture. Under these conditions, the blood
red [Fe(SCN)4]- ion is not formed. However, upon exposure
to oxygen, the addition of SCN- turned the solutions red,
indicating the formation of iron(III).28 These observations
are consistent with the initial production of iron(II) followed
by its rapid oxidation to iron(III) by molecular oxygen under
alkaline conditions.
Gas was rapidly evolved in the ferrate/hydroxylamine
reaction. To identify this gas, solutions of each reactant were
first flushed with argon to remove dissolved oxygen or
carbon dioxide and then mixed anaerobically in a sealed vial.
After the reaction was complete, a headspace sample was
removed with a gastight syringe and analyzed using mass
spectroscopy. A single peak at m/e 44 indicated the sole
production of N2O. The gas phase IR shows absorptions at
2230 and 1280 cm-1 which are also consistent with N2O.27
The presence of nitrate or nitrite was excluded by the absence
of any adsorption in the UV after iron(III) was removed by
ion exchange chromatography.
The N-methylhydroxylamine and O-methylhydroxylamine
oxidation products were determined with a combination of
GC and GC/MS. Characterization of the gaseous product was
the same as for NH2OH. The oxidation of O-methylhydroxy-
lamine produced N2O whereas analysis of the N-methylhy-
droxylamine reaction headspace showed only a peak for
argon, the gas used for flushing, indicating the absence of
gaseous products.
The nongaseous reaction products were identified by gas
chromatography. The spent aqueous reaction mixture was
extracted into methylene chloride and analyzed by GC. For
the oxidation of O-methylhydroxylamine, a single peak was
obtained with a retention time identical to that of a methanol
standard. The oxidation of N-methylhydroxylamine with an
excess of ferrate was allowed to proceed for 1 h, and only
traces of nitromethane were detected. Similar results were
observed using phenylhydroxylamine where only nitrosoben-
zene was observed as the final product.
Experimental Section
Materials. Hydroxylamine, N-methylhydroxylamine, and O-
methylhydroxylamine were purchased from Aldrich Chemicals as
their hydrochloride salts and dried on a vacuum line before use.
N-Phenylhydroxylamine was prepared by reduction of nitrobenzene
with zinc powder as described in the literature24 and recrystallized
from warm benzene. Water enriched in 18O or 17O was a gift from
Dr. R. Kent Murmann. All other chemicals were of reagent grade.
Potassium ferrate was prepared by the method reported by
Thompson, Ockermann, and Schreyer.25 The crude product was
recrystallized from saturated potassium hydroxide and the purity
checked by spectrophotometric analysis at 505 nm (ꢀ ) 1170 M-1
cm-1). Purities up to 95% were obtained, and the measured rates
were found to be independent of the sample purity. H217O (27%
atom) or H218O (99.9% atom) was used to label ferrate oxygens.
Samples of ferrate were dissolved in a minimum of water, allowed
to stand for about 10 min, and dried under reduced pressure.19
Kinetic Measurements. Stock solutions of the reductants were
prepared using Nanopure water, and ionic strength was maintained
with sodium perchlorate. Sodium phosphate was used to control
pH, and it also prevents precipitation of iron hydroxides formed
during the reaction. Argon bubbling was used to remove dissolved
oxygen and maintain anaerobic conditions.
The reaction rates were measured spectrophotometrically using
a D110 Durrum stopped-flow spectrophotometer with an OLIS
computer interface.26 All reactions were carried out under pseudo-
first-order conditions with the reductant in at least 10-fold excess.
Reductant concentrations ranged from 0.005 to 0.050 M. Changes
in absorbance versus time were fit to appropriate kinetic equations
with OLIS data fitting software. Each rate constant determined from
stopped-flow experiments was an average of 3-5 trials. Rapid
scanning experiments were carried out using an OLIS RSM 1000
rapid scanning spectrophotometer.
Product Analysis. Spent reaction mixtures were extracted with
methylene chloride, dried over anhydrous sodium sulfate, and
analyzed with a Hewlett-Packard 5880 gas chromatograph equipped
with a Superox II capillary column. Spectral characterization used
an HP 8452A diode array spectrophotometer. Isotopic tracer studies
(25) Thompson, G. W.; Ockerman, L. T.; Schreyer, J. M. J. Am. Chem.
Soc. 1951, 73, 1379.
A precise direct titration of a ferrate solution with excess
hydroxylamine to determine the reaction stoichiometry was
not possible due to the production of reduced iron species
that catalyze the decomposition of ferrate before the end point
was reached. Crude determinations, made by rapidly mixing
(26) Standard interface package from OLIS Corporation, GA.
(27) Laane, J.; Ohlsen, J. R. Prog. Inorg. Chem. 1980, 27, 466.
(28) Vogel, A. Vogel’s QualitatiVe Inorganic Analysis; Wiley: New York,
1987; pp 101-103.
(29) Norcross, B. E.; Lewis, W. C.; Huifa, G.; Noureldin, N. A.; Lee, D.
G. Can. J. Chem. 1997, 75, 129.
Inorganic Chemistry, Vol. 42, No. 21, 2003 6925

Johnson and Hornstein
Figure 2. Representative hydrogen ion dependence for the hydroxylamine
and N-methylhydroxylamine oxidation by ferrate. Conditions: T ) 25 °C,
I ) 1.0 M (NaClO4), [Na phosphate] ) 50 mM.
Figure 1. Typical plot of kobs versus [hydroxylamine]. Conditions: T )
25 °C, I ) 1.0 M (NaClO4), [Na phosphate] ) 50 mM; NH2OH, pH )
9.0; CH3NHOH, pH 8.8; CH3ONH2, pH 8.8.
Table 1. Proton Independent and Proton Dependent Rate Constants for
Hydroxylamine/Ferrate Reactions at 25 °C, 0.05 M Phosphate Buffer,
1.0 M Ionic Strength
ferrate and the hydroxylamine in known ratios until the
purple ferrate color persists, indicate 2:1 reductant to ferrate
stoichiometries for hydroxylamine and N-methylhydroxy-
lamine and N-phenylhydroxylamine. In an attempt to better
define the stoichiometry of these reactions, the addition of
excess hydroxylamine to a known quantity of aqueous ferrate
followed by back-titration of the remaining hydroxylamine
with dichromate gave 2.01 ( 0.35:1 and 2.1 ( 0.8:1
reductant/ferrate ratios for hydroxylamine and N-methylhy-
droxylamine, respectively, over the entire pH range of the
studies (pH 8.0-11.0). Similar experiments were attempted
with the O-methyl derivative; however, inconsistent results
were obtained presumably due to its relatively slow reaction
with ferrate or reaction of Cr(VI) with the methanol produced
from the Fe(VI) reaction. On the basis of product analysis
and the available stoichiometric data, the following reactions
can be written.
a
k1 ) ko (M-1 s-1
)
k2 (M-2 s-1
)
kh (M-1 s-1
)
NH2OH
(4.8 ( 0.3) × 103 (6.5 ( 0.8) × 1010
(3.3 ( 0.4) × 104
CH3NHOH (3.5 ( 5) × 103
(3.2 ( 0.4) × 1011 (1.60 ( 0.2) × 105
CH3ONH2
1.9 ( 0.9
(2.2 ( 0.4) × 109
110 ( 20
a kh ) Kak2.
Table 2. Phosphate Buffer Concentration on the Hydroxylamine/
Ferrate Reaction at 25 °C, 1.0 M Ionic Strength (NaClO4) and pH 10.3
[phosphate]T, M
k′ (M-1 s-1
)
0.010
0.050
0.10
850
770
930
0.15
960
0.25
1100
was too fast to study under these conditions. On the basis of
these observations, a general rate law for the disappearance
of ferrate may be written as follows.
2NH2OH + FeO42- + 4H+ f N2O + Fe2+ + 5H2O
2CH3NHOH + FeO42- + 4H+ f
2CH3NO + Fe2+ + 4H2O
-d[Fe(VI)]/dt ) k′[Fe(VI)][reductant]
Figure 2 shows a representative plot of k′ versus the hydrogen
ion concentration for hydroxylamine and N-methylhydroxy-
lamine. A similar linear proton dependence was observed
for O-methylhydroxylamine, and each is best described by
the following equation.
2CH3ONH2 + FeO42- + 4H+ f
2CH3OH + Fe2+ + N2O + 3H2O
2PhNHOH + FeO42- + 4H+ f 2PhNO + Fe2+ + 4H2O
k′ ) k1 + k2[H+]
Kinetics. The reaction rates were measured anaerobically
under pseudo-first-order conditions with at least a 10-fold
excess of reductant. Absorbance decreases were monitored
at 505 nm, λmax for ferrate, and observed rate constants, kobs’s,
were calculated using single-exponential fit routines, indicat-
ing a first-order dependence on ferrate(VI). Linear plots of
kobs versus the concentration of hydroxylamine show first-
order behavior for all the hydroxylamines, see Figure 1,
where the slope of these plots equals the second-order rate
constant, k′. The ferrate oxidation of phenylhydroxylamine
Values of k1 and k2 for each of the hydroxylamines were
determined, see Table 1. Since the pH was maintained with
phosphate buffer, a buffer concentration dependence study
was carried out. Table 2 shows that an increase in buffer
concentration, keeping the ionic strength constant with
sodium perchlorate, resulted in only a slight increase in rate,
30% increase over a 25-fold increase in phosphate concen-
tration. This effect was considered minimal, probably due
to medium effects and not to general acid catalysis. The lack
6926 Inorganic Chemistry, Vol. 42, No. 21, 2003

Ferrate(VI) Oxidation of Hydroxylamines
of sensitivity to phosphate also probably indicates that no
significant amounts of phosphatoferrate(VI) complexes were
formed.
explain the transfer of a ferrate oxygen to the final products
in the oxidation of sulfite.19
Fe18O42- + SO32- f 18O3Fe-18O-SO3
4-
Mechanism. The simplest mechanism for the oxidation
of hydroxylamine by ferrate consistent with this above
information is given in Scheme 1. Since the pH range over
which the study was carried out was higher than the pKa of
the hydroxylamines,30 protonation was assumed to occur at
the ferrate ion, pKa 7.3.30 In addition, as will be discussed
later, if the protonation occurs on hydroxylamine, a one-
electron or outer-sphere process is expected and is incon-
sistent with the experimental data.
18O3Fe-18O-SO34- f 18OSO32- + FeIV
The formation of an Fe-O-S species is followed by oxygen
transfer to the sulfate products. Subsequent studies by
Johnson and Bernard on the oxidation of selenite22 suggested
a similar oxidation mechanism, which is supported by recent
17O NMR studies that show oxygen transfer from ferrate to
form selenate.31
2-
In contrast, the oxidation of hydroxylamine by Fe18O4
Scheme 1
shows no incorporation of labeled oxygen into the final
product. The same result is found when the reaction is carried
out in H218O with unlabeled ferrate. These studies demon-
strate that the oxygen in the N2O final product comes from
the hydroxylamine itself; i.e., no oxygen transfer occurs
during oxidation. Similar studies were carried out with the
methylated hydroxylamines and 17O labeled ferrate. In these
experiments, the absence of 17O signals in the final products
again suggests a lack of oxygen transfer from the ferrate
ion. On the basis of these observations, the formation of
bridged intermediates via the ferrate oxygen seems unlikely.
Stanbury et al. reported large reorganizational energies in
the oxidation of hydroxylamine by hexachloroiridate(IV).32
After tying up any trace metals with oxalate, they were able
to calculate an unprecedently small self-exchange rate
constant for the NH2OH/NH2OH+ couple of 5 × 10-13 M-1
s-1 at 25 °C. On the basis of their observation, they
concluded that outer-sphere redox reactions with hydroxy-
lamine would be rare. Ferrate does not oxidize oxalate under
these experimental conditions, and studies carried out in 1
mM sodium oxalate showed no diminution of reaction rates
of ferrate with hydroxylamine. This suggests that catalysis
by trace metals is negligible, and the rapid oxidation rate
implies that an outer-sphere redox reaction is unlikely
between ferrate and hydroxylamine. A one-electron oxidation
is unlikely on the basis of negative results for radical
formation using acrylonitrile as a radical trap. These trapping
experiments have been shown to be successful for the
detection of radicals for the oxidation of some substituted
anilines and are expected to be reliable in the present studies
with hydroxylamines.
Two-electron transfers are implied in Scheme 1 for each
of the rate determining steps. Since an outer-sphere process
is energetically unfavorable for two-electron transfers, we are
forced to propose an inner-sphere mechanism. As discussed
in the Introduction, rapid dimerization of NOH and produc-
tion of N2O in the case of hydroxylamine oxidation is
indicative, although not proof, of a two-electron mechanism,
whereas one-electron oxidations of hydroxylamine typically
produce N2.2-4 Two-electron reductions of ferrate have been
proposed in other studies,24,29,31 and while the Fe(IV) species
has yet to be characterized, its existence in aqueous media
has been established by Bielski and co-workers.33 Their
studies show Fe(IV) reacts at least an order of magnitude
(R ) H or CH3)
HFeO42- a H+ + FeO4
Ka
2-
NH2OR + FeO42- f NOR + FeIV ko
NH2OR + HFeO4- f NOR + FeIV kh
NH2OR + FeIV f NOR + Fe2+ fast
H2O + 2NOR f N2O + 2ROH rapid
Similar oxidation mechanisms may be written for the
N-substituted hydroxylamines:
RNHOH + FeO42- f RNO + FeIV ko
RNHOH + HFeO4- f FeIV + RNO kh
FeIV + RNHOH f RNO + Fe2+ rapid
Because the precise nature of the FeIV species is unknown,
we have not balanced these reactions with respect to oxygens
or protons. On the basis of this mechanism, k1 ) k0 and k2
) Ka-1kh. Values for k0 and kh are shown in Table 1.
Since nitrate was not observed in the oxidation of
hydroxylamine, the coupling reaction of NOH with itself
must be significantly faster than further oxidation by either
ferrate(VI) or ferrate(IV). Under our conditions, the coupling
constant may be calculated; it is 5 × 108 M-1 s-1 at pH 9.9
At this pH, the ferrate/hydroxylamine rate constant is 5.38
× 103 M-1 s-1, see Figure 2, well below the coupling rate.
Assuming even a 103 faster rate for Fe(IV),33 the coupling
remains much faster than further oxidation to nitrate.
Ferrate intermediates previously proposed involve a sub-
strate bound to the iron through a ferrate oxygen. For
example, Murmann first suggested such an intermediate to
(30) Ka ) 7.24 × 10-7, 1.1 × 10-6, 2.51 × 10-5 for NH2OH, CH3NHOH,
and CH3ONH2, respectively. Marque, H. M. J. Chem. Soc., Dalton
Trans. 1991, 339. pKa for HFeO4- is 7.3. Sharma, V. K.; Burnett, C.
R.; Millero, F. Chem. Phys. 2001, 3, 2059.
(31) Johnson, M. D.; Hornstein, B. J.; Vogels, C. M. Unpublished results.
(32) Hung, M.; McKee, M. L.; Stanbury, D. M. Inorg. Chem. 1994, 33,
5108.
(33) Bielski, B. J. H.; Thomas, M. J. Inorg. Chem. 1989, 28, 3947.
Inorganic Chemistry, Vol. 42, No. 21, 2003 6927

Johnson and Hornstein
faster than ferrate with reductant. This is consistent with the
assumption that mechanistic steps involving Fe(IV) are not
rate determining. In light of the experimental evidence, an
intermediate in which the substrate is bound directly to the
iron in the ferrate ion is likely. Since the ferrate/hydroxyl-
amine reaction occurs much faster than oxygen exchange
on ferrate, an expanded first coordination sphere intermediate
seems likely. Three proposed structures are shown here.
Scheme 2
HFeO4- a H+ + FeO4
Ka
2-
FeO42- + HNROR′ a [O4Fes(NHROR′)]2- Khf
HFeO4- + HNROR′ a [HO4Fes(NHROR′)]- Kf
when R ) H and R′ ) H or CH3
[HO4Fes(NH2OR)]- f NOR + FeIV kh′
[O4Fes(NH2OR)]2- f NOR + FeIV ko′
NH2OH + FeIV f NOR + Fe2+ rapid
H2O + 2NOR f N2O + 2ROH rapid
when R ) CH3 and R′ ) H
Unfortunately, it is impossible to distinguish between these
possibilities on the basis of the present kinetic studies, and all
attempts to trap or spectrally observe an intermediate were
unsuccessful. Although there are no examples of similar
iron-hydroxylamine complexes, Weighardt and others have
established analogous structures for the hydroxylamine
adducts of tungstate, molybdate and vanadate.34,35 On the
basis of these analogues, the side-on binding of hydroxyl-
amine to ferrate would seem most likely.
[HO4FesNHCH3(OH)]- f CH3NO + FeIV kh′
[O4FesNHCH3(OH)]2- f CH3NO + FeIV ko′
CH3NHOH + FeIV f CH3NO + Fe2+ rapid
is presumed to occur via a two-electron mechanism. This is
consistent with recent observations of ferrate oxida-
tions17,19,23,29 and supports an emerging theme for ferrate in
that it preferentially acts as a two-electron oxidant.
Although no intermediates are observed spectrally for the
oxidation of hydroxylamines by ferrate, their presence is
suggested by the oxidation products. An intermediate formed
between the hydroxylamine oxygen and the iron(VI) center
is similar to the one proposed by Lee in the oxidation of
alcohols by ferrate.29 Oxygen bridged (through the iron
oxygen) structures are disfavored on the basis of the lack of
oxygen transfer from ferrate to the hydroxylamine final
product. Further kinetic studies involving the oxidation of
nitrogen based compounds are needed to elucidate these
suggested intermediates and the propensity of ferrate to
oxidize by two electrons.
A general mechanism that incorporates both the proposed
intermediate and tracer studies is shown in Scheme 2.
From this mechanism, ko and kh reported in Table 1 are
actually composites of ko′Kf and kh′Khf. Unfortunately, the
formation constants are unknown so values for the intramo-
lecular steps cannot be determined. Rapid scanning spec-
trophotometry does not show any change in the shape or
position of the absorption spectrum of ferrate during the
redox reaction. This implies that the majority of the iron-
(VI) species present exists as “free” ferrate and that the values
of Khf and Kf are small.
Conclusions
A thorough study of the oxidation of hydroxylamine,
N-methylhydroxylamine, and O-methylhydroxylamine is
reported. When the reductant is present in excess, NH2OH
is oxidized to N2O, CH3NHOH is oxidized to CH3NO, and
CH3ONH2 is converted to CH3OH and N2O. Each oxidation
Acknowledgment. The authors with to thank Professor
R. Kent Murmann for his kind gift of H217O and help with
the isotope tracer studies. This work was funded in part by
an ACS/PRF grant (PRF #35481-AC3).
(34) Weighart, K. AdV. Inorg. Bioinorg. Mech. 1984, 3, 213.
(35) Keramidas, A.; Miller, S.; Anderson, O.; Crans, D. C. J. Am. Chem.
Soc. 1997, 119, 8901.
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