Novel Oxidizing Reagent Based on K2FeO4
J . Org. Chem., Vol. 61, No. 18, 1996 6361
combination with an aluminosilicate solid. Association
of a ferric salt with a clay support had been pioneered in
this laboratory in the 1980s, leading to the design of the
“clayfen” supported reagent,10 with numerous uses in
organic synthesis,11 including the mild and selective
oxidation of alcohols to carbonyl derivatives.12 We rea-
soned that replacing iron(III) with iron(VI) should boost
the oxidizing ability of the reagent, while the continued
presence of a microporous adsorbent of the clay or zeolite
type would be conducive to the high selectivities our
supported reagents and catalysts have repeatedly shown.
By contrast with other transition metals, iron is singled
out as being considered nontoxic.13 In addition, replace-
ment of both Brønsted (H2SO4, HNO3, etc.) and Lewis
acids (AlCl3, TiCl4, etc.) with aluminosilicate solids also
contributes to decreasing the amount of toxic and cor-
rosive wastes.
conversely, it is possible to form the Na4FeO5 salt by
heating to 370 °C a mixture of Fe2O3 and Na2O2 under
an atmosphere of dioxygen. Rigorous control of the
experimental conditions is required in order to minimize
the amount of iron(III) and iron(IV) derivatives formed
as byproducts contaminating the desired Fe(VI) salt.22,23
Fortunately, K2FeO4 provides a better entry into the
group of iron(VI) derivatives. Many preparations of this
salt in aqueous solution take advantage of both the
stability of the ferrate dianion in basic medium and the
insolubility of K2FeO4 in a saturated solution of potas-
sium hydroxide. This ensures first formation and second
isolation through precipitation of potassium ferrate.
Such a procedure can be transposed to preparation of
both rubidium and cesium ferrates (Rb2FeO4 and Cs2-
FeO4).24 The cost of the RbOH and CsOH hydroxides
makes impractical any but small-scale preparation of
these ferrates.
Wh y P ota ssiu m F er r a te(VI)?
P r ep a r a tion of P ota ssiu m F er r a te(VI)
Potassium ferrate (K2FeO4) is the best known member
among the family of iron(VI) derivatives. It is made and
purified more easily, and it is also used in making other
ferrates. A quick survey shows that it is both more stable
and more readily made. The ferrates of alkaline earths
SrFeO4 or BaFeO4 and their hydrates can be obtained
indirectly from reaction between a solution of the acetate
or chloride of the metal and potassium ferrate.14,15
Barium ferrate (BaFeO4), as the monohydrate, can be
prepared in like manner by precipitation from an alkaline
solution of sodium ferrate upon addition of a saturated
solution of barium nitrate.16 To form the mixed potas-
sium-strontium ferrate (K2Sr(FeO4)2), a more complex
system consisting of Sr(OAc)2, Sr(OH)2, KCl, KOH, and
K2FeO4 was used.17 Such methods lack generality how-
ever. Losana described in 1925 the preparation of
numerous metallic ferrates by exchange from BaFeO4 or
K2FeO4.18 But the truth is that subsequent attempts at
The history thread gives meaning to the present
procedures. As early as 1702, the German chemist and
physician Georg Stahl mentioned the appearance of an
unstable purplish red color when the molten mass
resulting from detonation of a mixture of saltpeter and
iron filings was dissolved in water.25 When Eckeberg26
and Becquerel27 in 1834 heated to red mixtures of potash
and various iron ores, they also observed similar colors,
which we now know to be diagnostic of the FeO42- ferrate
dianion. As for attribution of this color to a high-valent
iron species indeed, credit goes to Fre´my, who in the
1840s suggested a formula of the FeO3 type.28 Even
though this FeO3 oxide was never isolated, the presence
of hexavalent iron in potassium ferrate and barium
ferrate was demonstrated by various methods during the
ensuing period, allowing Moeser to write a detailed
review of ferrates and their chemistry in 1897.29
Moeser described three types of preparation for potas-
sium ferrate: viz., a dry way, heating to red various
potassium- and iron-containing minerals; an electro-
chemical way, electrolyzing a potash solution with an iron
anode; and a wet way, oxidizing a basic solution of a Fe-
(III) salt by an hypochlorite or hypobromite. Among
these three approaches, the 20th century has privileged
the third and last. Dry reactions imply detonation and
elevated temperatures. They are considered dangerous
(we agree) and too difficult to implement, and they have
become obsolete. The only relatively recent contribution
of that nature, by Russian authors, describes preparation
of potassium ferrate by calcination of a mixture of ferric
oxide and potassium peroxide at 350-370 °C under a
dioxygen flow.22 As for anodic oxidation by an iron
electrode dipped in a concentrated solution of an alkaline
preparing calcium or transition metal (Co2+, Fe3+, Hg2+
,
Pb2+, Zn2+) ferrates from K2FeO4 have failed.19,20 The
expected products did not precipitate, or, when they did,
their instability led them to immediate decomposition
and evolution of dioxygen. Only silver ferrate (Ag2FeO4)
could be obtained from K2FeO4 and AgNO3. It is not very
stable at ambient temperature, though, and has to be
stored in a refrigerator.21
Sodium ferrate (Na2FeO4) has a behavior different from
those of other ferrates and remains soluble in an aqueous
solution saturated in sodium hydroxide. Its preparation
from an aqueous medium is thus made difficult and leads
to rather impure samples. In the absence of solvent,
(10) Fieser, M. Fieser and Fieser’s Reagents for Organic Synthesis;
Wiley: New York, 1984; Vol. 11, p 237; 1986; Vol. 12, p 231.
(11) Corne´lis, A.; Laszlo, P. Synthesis 1985, 909-918.
(12) Corne´lis, A.; Laszlo, P. Synthesis 1980, 849-850.
(13) Knepper, W. A. In Kirk-Othmer Encyclopedia of Chemical
Technology, 3rd ed.; Wiley: New York, 1981; Vol. 13, pp 735-753.
(14) Gump, J . R.; Wagner, W. F.; Schreyer, J . M. Anal. Chem. 1954,
26, 1957.
(15) Scholder, R.; Bunsen, H. v.; Kindervater, F.; Zeiss, W. Z. Anorg.
Allg. Chem. 1955, 282, 268-279.
(16) Firouzabadi, H.; Mohajer, D.; Entezari-Moghaddam, M. Bull.
Chem. Soc. J pn. 1988, 61, 2185-2189.
(17) Ogasarawa, S.; Tanako, M.; Bando, Y. Bull. Inst. Chem. Res.
Kyoto Univ. 1988, 66, 64-67; Chem. Abstr. 1989, 110, 17644a.
(18) Losana, L. Gazz. Chim. Ital. 1925, 55, 468-497; Chem. Abstr.
1926, 20, 156.
(19) Herber, R. H.; J ohnson, D. Inorg. Chem. 1979, 18, 2786-2790.
(20) Firouzabadi, H.; Mohajer, D.; Entezari-Moghaddam, M. Synth.
Commun. 1986, 16, 723-731.
(21) Firouzabadi, H.; Mohajer, D.; Entezari-Moghaddam, M. Synth.
Commun. 1986, 16, 211-223.
(22) Kiselev, Y. M.; Kopelev, N. S.; Zavyalova, N. A.; Perfiliev, Y.
D.; Kazin, P. E. Russ. J . Inorg. Chem. (Transl. of Zh. Neorg. Khim.)
1989, 34, 1250-1253.
(23) Kopelev, N. S.; Perfiliev, Y. D.; Kiselev, Y. M. J . Radioanal.
Nucl. Chem. 1992, 162, 239-251.
(24) Audette, R. J .; Quail, J . W. Inorg. Chem. 1972, 11, 1904-1908.
(25) Stahl, G. E. Opusculum Chymico-Physico-Medicum; Halae-
Magdeburgiae, 1715, p 742.
(26) Referred to in the following: Moeser, L. J . Prakt. Chem. 1897,
56, 425-437.
(27) Becquerel, A. Ann. Chim. Phys. 1834, 51, 105. If the name
sounds familiar, it should: the Becquerels were a scientific dynasty
and the Nobel Prize winner for the discovery of radioactivity was the
grandson of this scientist of the 1830s.
(28) (a) Fre´my, E. C. R. Acad. Sci. Paris 1841, 12, 23; (b) 1842, 14,
442; (c) Ann. Chim. Phys. 1844, 12 (Se´r. 3), 361-382.
(29) Moeser, L. J . Prakt. Chem. 1897, 56, 425-437.