Investigation of the redox interaction between iron(III) 5,10,15,20-tetrakis(p-sulfonatophenyl)porphyrinate and aminoiminomethanesulfinic acid in aqueous solution

Article abstract of DOI:10.1039/a804684k

Stability studies have been made on aminoiminomethanesulfinic acid [aimsa, thiourea dioxide, (NH₂)₂CSO₂] and the reactions of aimsa and its decomposition product (dithionite, S₂O₄^2-) with iron(III) 5,10,15,20-tetrakis-(p-sulfonatophenyl)porphyrinate [Fe^III(tpps)] in the presence of oxygen have been investigated. Application of NMR and stopped-flow spectrophotometry revealed direct evidence for the existence of two different forms of aimsa in aqueous solution. The slow formation of one of these forms, NH₂NHCSO₂H, is shown to determine the strong dependence of the reactivity of aimsa on the age of the stock solutions. Both aimsa and dithionite react in a similar way with Fe^III(tpps) in alkaline solutions. The SO₂^?- radical plays a key role in the redox reactions. The ratio of the oxygen and radical concentration determines which kind of reaction (oxidation, reduction or decomposition) will dominate. In general a high oxygen concentration and a low radical concentration favour the oxidation and decomposition of the metalloporphyrin, whereas a high radical concentration and a low oxygen concentration favour the reduction. In strongly basic solutions ([NaOH] = 0.5 M) a redox cycle between Fe^III(tpps) and Fe^II(tpps) is observed at low aimsa and dithionite concentration. Possible mechanisms for the decomposition of aimsa in alkaline solutions, as well as for the reactions between Fe^III(tpps) and aimsa or dithionite, are proposed.

Full text of DOI:10.1039/a804684k

View Article Online / Journal Homepage / Table of Contents for this issue
Investigation of the redox interaction between iron(III)
5,10,15,20-tetrakis(p-sulfonatophenyl)porphyrinate and
aminoiminomethanesulfinic acid in aqueous solution
Vasilios Lepentsiotis,a Rudi van Eldik,*,a Dimitri M. Stulovb and Sergei V. Makarov*,b
a Institute for Inorganic Chemistry, University of Erlangen-Nürnberg, Egerlandstrasse 1,
91058 Erlangen, Germany
b Ivanovo Academy of Chemistry and Technology, Engels str. 7, 153460 Ivanovo, Russia
Stability studies have been made on aminoiminomethanesulfinic acid [aimsa, thiourea dioxide, (NH2)2CSO2]
and the reactions of aimsa and its decomposition product (dithionite, S2O42Ϫ) with iron() 5,10,15,20-tetrakis-
(p-sulfonatophenyl)porphyrinate [FeIII(tpps)] in the presence of oxygen have been investigated. Application of
NMR and stopped-flow spectrophotometry revealed direct evidence for the existence of two different forms of
aimsa in aqueous solution. The slow formation of one of these forms, NH2NHCSO2H, is shown to determine
the strong dependence of the reactivity of aimsa on the age of the stock solutions. Both aimsa and dithionite react
in a similar way with FeIII(tpps) in alkaline solutions. The SO2 Ϫ radical plays a key role in the redox reactions.
᝽
The ratio of the oxygen and radical concentration determines which kind of reaction (oxidation, reduction or
decomposition) will dominate. In general a high oxygen concentration and a low radical concentration favour the
oxidation and decomposition of the metalloporphyrin, whereas a high radical concentration and a low oxygen
concentration favour the reduction. In strongly basic solutions ([NaOH] = 0.5 ) a redox cycle between FeIII(tpps)
and FeII(tpps) is observed at low aimsa and dithionite concentration. Possible mechanisms for the decomposition
of aimsa in alkaline solutions, as well as for the reactions between FeIII(tpps) and aimsa or dithionite, are
proposed.
Thiourea dioxide (formamidinesulfinic acid or aminoimino-
methanesulfinic acid, abbreviated as aimsa) can easily be pre-
pared via the oxidation of thiourea with aqueous solutions of
hydrogen peroxide.1 Numerous commercial applications are
related to its strong reducing properties exhibited in alkaline
solutions,2 which are attributed to its decomposition products,
mainly to sulfoxylic acid (dihydrogen dioxosulfate) H2SO2 or
conditions existed. By way of comparison, the reaction between
FeIII(tpps) and dithionite, S2O42Ϫ, a plausible decomposition
product of aimsa, in the presence of oxygen was also investi-
gated in a preliminary way.
Experimental
the anion radical SO2 Ϫ. It is therefore widely used as a bleach-
Chemicals of analytical reagent grade and deionized ultrapure
water were used to prepare all solutions. Nitrogen was used to
deaerate solutions when required. Aqueous HClO4 and NaOH
were used to adjust the pH, and NaClO4 to adjust the ionic
strength of the test solutions. Aminoiminomethanesulfinic acid
[(NH2)2CSO2] and sodium dithionite (Na2S2O4) were from
Aldrich; [FeIII(tpps)]3Ϫ was synthesized according to the liter-
ature.14 2-Cyclohexylaminoethanesulfonic acid (ches, Aldrich)
or NaHCO3 were used for the preparation of buffer solutions.
The UV/VIS spectra were recorded on Shimadzu UV-2102/
3102PC spectrophotometers. Rapid-scan measurements were
performed on a Bio Sequential SX-17MV Stopped-Flow Reac-
tion Analyser from Applied Photophysics equipped with a J &
M detector connected to a TIDAS 16-416 spectrophotometer.
Stopped-flow measurements were performed on a Bio Sequen-
᝽
ing agent in the textile industry,2,3 for reducing organo-sulfur,
-selenium, -tellurium and -fluorine compounds,4–7 as well as for
the determination of some nitrogen-containing compounds8
and metals.9 Aminoiminomethanesulfinic acid is also a com-
ponent of an effective initiation system for polymerization reac-
tions.10,11 Unfortunately, papers devoted to investigations of
its decomposition contain many contradictory results and con-
clusions. It is even unknown whether the decomposition process
is an ionic or a radical one.2,9
The present paper reports the results of stability studies on
aimsa in alkaline solutions and its reaction with iron()
5,10,15,20-tetrakis(p-sulfonatophenyl)porphyrinate FeIII(tpps)
in the presence of oxygen. This choice of oxidant is not only
due to the importance of porphyrin chemistry per se, but also
by way of comparison due to the availability of information on
its reactions with other sulfur-containing reductants.12,13 In a
recent study the interaction between FeIII(tpps) and sulfite in
aqueous solution was investigated in search of evidence for
redox cycling of this metalloporphyrin.13 The application of
rapid-scan spectrophotometry revealed direct evidence for such
redox cycling, which was found to strongly depend on the
selected experimental conditions. We have, therefore, in the
present study investigated the interaction between FeIII(tpps)
and aimsa in search of conditions for redox cycling. In the case
of aimsa we are dealing with a strong reducing agent, but the
presence of oxygen may significantly influence its behaviour in
aqueous solutions so that it could react in a similar way as in
the case of sulfite/O2, oxone and H2O2, where strong oxidizing
1
tial SX-18MV stopped-flow spectrophotometer. The H NMR
spectra were recorded on a 400 MHz Bruker Avance DPX-400
spectrometer at 305 K. Commercial grade deuteriated water,
D2O (Aldrich, 99.9%), was used without further purification.
Results and Discussion
Decomposition of aimsa in alkaline media
Aminoiminomethanesulfinic acid is characterized by an
absorbance maximum at 269 nm (ε = 489 Ϫ1 cmϪ1 15), such that
its decomposition can be followed spectrophotometrically. The
UV spectrum of an air-saturated alkaline solution of it was
recorded as a function of time. The absorbance at 269 nm
decreased in intensity with time in strongly alkaline solutions
J. Chem. Soc., Dalton Trans., 1998, Pages 2915–2920
2915

View Article Online
Fig. 3 Absorbance vs. time traces for the decomposition of aimsa in
a basic air-saturated solution: [NaOH] = 0.05 (1), 0.1 (2), 0.2 (3), 0.5 (4),
1 (5) and 2.58  (6); [aimsa] = 1.5 × 10Ϫ3 , 25 ЊC, λ = 315 nm
Fig. 1 Spectral changes recorded for the decomposition of aimsa in
strongly basic medium (air-saturated solutions). [aimsa] = 1.8 × 10Ϫ3 ,
[NaOH] = 0.5 , 25 ЊC, t ≈ 1200 s; ∆t = 1 min for spectra 1–4, 3 min
between spectrum 4 and 5, 1 min for spectra 5–8 and 5 min for the last
two spectra
Fig. 4 Absorbance vs. time trace for the decomposition of aimsa in a
strongly basic nitrogen-saturated solution: [aimsa] = 1.48 × 10Ϫ3 ,
[NaOH] = 1 , 25 ЊC, λ = 315 nm
partly affect the decrease in absorbance at 269 nm following the
break point.
In nitrogen-saturated strongly alkaline solutions the appear-
ance of dithionite was observed immediately after mixing of an
aqueous solution of aimsa with base (Fig. 4). Interestingly, we
were not able to observe any increase in absorbance at 315 nm
during the time range mentioned in Fig. 4 if we used a fresh
nitrogen-saturated solution of aimsa (time between dissolution
and the kinetic run was less than 5 min). The dependence of
the aimsa reactivity on the age of the stock solution was
observed earlier by Mefodieva16 and Rabai et al.,17 and a pos-
sible explanation for this observation will be given later on. In
all subsequent experiments we used aged solutions of aimsa
(not less than 1 h). Longer aging (up to 24 h) of aqueous (pH
4–5) solutions of aimsa had no influence on the kinetic
characteristics.
Fig. 2 Absorbance vs. time traces for the decomposition of aimsa in a
strongly basic medium (air-saturated solution): (a) in a sealed tandem
cuvette, (b) in an open tandem cuvette and (c) after subtraction of the
dithionite absorbance at 269 nm. [aimsa] = 1.8 × 10Ϫ3 , [NaOH] = 0.5
, 25 ЊC, λ = 269 nm
([NaOH] = 0.5 ) and after some delay a band appeared at 315
nm which increased in intensity with time and reached a con-
stant value (Fig. 1). This band is the same as that observed
previously for sodium dithionite.9 The observed absorbance vs.
time traces at 269 and 315 nm are presented in Figs. 2 and 3,
respectively. The trace at 269 nm clearly shows a peculiar break
after which the absorbance decrease continues at a slower rate.
If air is bubbled through the cuvette there is no break in the
trace. Interestingly this break point was observed more clearly
if the reaction was performed in a sealed quartz cuvette. If the
cuvette was open to air only a weak break in the curve occurred.
It is important to mention that the time of the break point
corresponded to the time of the appearance of dithionite at 315
nm (Fig. 3). The latter also absorbs at 269 nm and this may
It is known that dithionite exists in equilibrium with the sul-
2Ϫ
Ϫ
fur dioxide anion radical, S2O4
2 SO2 (K = 1.4 × 10Ϫ9
᝽
,18 k1 = 1.7 sϪ1 19), which may form directly from aimsa.9 The
reaction of SO2 with O2 proceeds very rapidly (k = 2.4 × 109
Ϫ
᝽

Ϫ1 sϪ1 20). Therefore we assume that the delay in appearance of
2Ϫ
S2O4 in air-saturated strongly alkaline solutions of aimsa is
due to the reaction of SO2 Ϫ with oxygen. We were not able to
᝽
observe an increase in absorbance at 315 nm in air-saturated
solutions of aimsa ([aimsa]0 < 2 × 10Ϫ3 ) at pH р 11, but it
2916
J. Chem. Soc., Dalton Trans., 1998, Pages 2915–2920

View Article Online
changeover for d[S2O42Ϫ]/dt, depending on the pH, was previ-
ously observed in alkaline solutions.9
2Ϫ
In order to determine the concentration of S2O4 in aimsa
solutions, ε(S2O42Ϫ) at 315 nm is required. There is a relatively
large scatter in the ε(S2O42Ϫ) data in the literature (λ = 315 nm),
6900,26,27 800028 and 804329 Ϫ1 cmϪ1, which could be due to the

effect of dissolved oxygen that will oxidize S2O42Ϫ, or the purity
of dithionite. Using the most reliable value of 8043 Ϫ1 cmϪ1
(ref. 29), where analytically pure dithionite was obtained for the
2Ϫ
first time, we estimate the concentration of S2O4 produced
during aimsa decomposition in 0.5  NaOH to be 1.5 × 10Ϫ4 ,
much smaller than the initial concentration of aimsa (Fig. 3),
even taking into account the influence of oxygen. We therefore
assume that there exists another, non-radical pathway for aimsa
decomposition. The most plausible products of this reaction
2Ϫ
are sulfoxylate SO2 and urea (it is known that urea is a pre-
dominant nitrogen-containing product of aimsa decomposition
in strongly alkaline solutions30).
On the basis of the results mentioned above we suggest the
mechanism in equations (1)–(8) for the decomposition of aimsa
(NH2)2CSO2
NH2(NH)CSO2H
(1)
NH2(NH)CSO2H ϩ OHϪ
NH2(NH)CSO2Ϫ ϩ OHϪ
NH2(NH)CSO2Ϫ ϩ H2O (2)
2Ϫ
NH2(NH)(OH)CSO2
(3)
2Ϫ
NH2(NH)(OH)CSO2
NHCNH
(carbodiimide) ϩ SO22Ϫ ϩ H2O
(4)
Fig. 5 The NMR spectra recorded after dissolving aimsa in D2O
(nitrogen-saturated solutions). [aimsa] ≈ 0.15 , 30 ЊC
2Ϫ
NH2(NH)(OH)CSO2
Ϫ
NH2(NH)(OH)C Ϫ ϩ SO2
(5)
(6)
(7)
(8)
᝽
᝽
was observed in the nitrogen-saturated solutions.21 At pH р 11
the absorbance at 269 nm gradually decreases to close to zero
and there is no increase in absorbance at 315 nm.
NHCNH
NH2CN
2Ϫ
NH2CN ϩ H2O
(NH2)2CO
Interestingly, the rate of appearance of S2O4 does not
depend on pH in strongly alkaline solutions (Fig. 3). This result
can be explained on the assumption that aimsa exists in differ-
ent tautomeric forms in aqueous solution, or that there is a
rearrangement of it after being dissolved in water. The existence
of two tautomeric forms of aimsa in aqueous solution has been
assumed by many authors,4,21 but direct experimental evidence
is not available.
Ϫ
2Ϫ
2 SO2
S2O4
᝽
Ϫ
in alkaline solutions. In the presence of oxygen, SO2 rapidly
᝽
produces O2 Ϫ according to reaction (9).20 The equilibration in
᝽
SO2 Ϫ ϩ O2 → SO2 ϩ O2 Ϫ (k = 2.4 × 109 Ϫ1 sϪ1) (9)
᝽
᝽
reaction (1) explains the existence of ‘fast’ and ‘slow’ steps of
S2O4 appearance during aimsa decomposition. After mixing
H2N
+
O
HN
O
2Ϫ
–
C
S
C
S
aqueous solutions of aimsa with alkali the available reactive
form II rapidly produces SO2 Ϫ and S2O42Ϫ. The rest of S2O4
H2N
O
H2N
OH
2Ϫ
᝽
I
II
is formed much slower depending on the rate of formation of
NH2(NH)CSO2H. Being the rate-determining step, reaction (1)
determines the pH independence of aimsa decomposition in
strongly alkaline solutions.
It is known that aimsa exists in form I in the solid state.22 In
order to investigate possible changes in its structure after being
dissolved in water, 1H NMR spectra of nitrogen-saturated
solutions of aimsa were recorded as a function of time. Fig. 5
shows that the initial peak at δ 8.6 quickly splits into a doublet.
The changes in the NMR spectra were observed only during
the first minutes after dissolving aimsa in water (the earlier
investigators only observed the final doublet23). A reasonable
explanation for the splitting into a doublet invokes a decrease in
symmetry along the C᎐S bond, followed by intermediate
exchange between two magnetically non-equivalent nitrogen
nuclei. A similar splitting into a doublet was observed earlier in
the NMR spectra of thiourea.24 The absence of the SO2H pro-
ton may be explained by fast equilibration between deuteriated
water and the acidic hydrogen. Unfortunately, we were not able
The most challenging phenomenon is the changeover of
order in [S2O42Ϫ] with increasing pH. This changeover may be
due to attainment of steady-state conditions, where d[S2O42Ϫ]/
dt reaches a constant value. Seemingly, there is an interaction
between nitrogen-containing products of radical and ionic
aimsa decomposition at [OHϪ] у 0.2 : carbodiimide, urea and
cyanamide. The formation of these compounds was observed in
alkaline solutions of aimsa30 and some of them may be formed
from the nitrogen-containing radical, produced in reaction (5).
The equilibration of its reactions accounts for the constant
value of d[S2O42Ϫ]/dt in strongly alkaline solutions.
Reaction of aimsa with FeIII(tpps) in alkaline solutions
1
to record any H NMR spectra of aimsa in alkaline solutions.
This may be due to the base-catalysed NH-proton exchange,
The reaction between the FeIII(tpps) dimer and aimsa was
studied at pH 11 (NaHCO3 buffer) using a tandem cuvette.
In order to avoid spontaneous decomposition of aimsa
before mixing with FeIII(tpps), only the metalloporphyrin solu-
tion was buffered. After mixing the two solutions the pH was
since the same behaviour was observed for thiourea.25
Fig. 3 shows also that the absorbance increase at 315 nm
exhibits good zero-order behaviour at [OHϪ] у 0.2 . At lower
pH we observed deviations from zero-order behaviour. Such a
J. Chem. Soc., Dalton Trans., 1998, Pages 2915–2920
2917

View Article Online
Fig. 6 Spectral changes recorded for the reaction between the FeIII
(tpps) dimer and aimsa in a nitrogen-saturated basic solution. [FeIII
-
-
(tpps)] = 6 × 10Ϫ6 , [aimsa] = 7.5 × 10Ϫ4 , pH 11; I = 0.1 ; 25 ЊC;
∆t = 1 min between the first two spectra and 5 min between the others
Fig. 8 (a) Spectral changes recorded for the reaction between the FeIII
-
-
(tpps) dimer and aimsa in an air-saturated basic solution. [FeIII
(tpps)] = 6 × 10Ϫ6 , [aimsa] = 5 × 10Ϫ5 , pH 11, I 0.1 , 25 ЊC,
∆t = 900 s. (b) Absorbance vs. time dependence for the reaction between
the FeIII(tpps) dimer and aimsa in an air-saturated basic solution. [FeIII
-
(tpps)] = 6 × 10Ϫ6 , [aimsa] = 5 × 10Ϫ5 , pH 11, I 0.1 , 25 ЊC, λ = 408
Fig. 7 Spectral changes recorded for the back reaction between the
FeIII(tpps) dimer and aimsa in a nitrogen-saturated basic solution.
Details as in Fig. 6 except ∆t = 150 min (see Discussion)
nm
decay of the Soret band of the FeIII(tpps) is larger and the
intensity of the FeII(tpps) band is much lower than in the
case of nitrogen-saturated solutions, indicating that oxygen
favours the oxidation and decomposition of the complex. At
high oxygen concentration the oxidation and decomposition
of FeIII(tpps) is favoured, whereas the reduction is inhibited.
This can be explained by the fact that in the presence of
approximately 11. The aimsa and oxygen concentrations were
varied, whereas the FeIII(tpps) concentration (6 × 10Ϫ6 ) and
the ionic strength (0.1 ) were kept constant. Fig. 6 shows the
spectral changes observed for the reaction of FeIII(tpps) with a
large excess of aimsa (7.5 × 10Ϫ4 ) in nitrogen-saturated solu-
tions. The cycle time between the first two spectra is 1 min and
5 min between the others. After 1 min a clear decrease in the
Soret band of the FeIII(tpps) dimer and the appearance of a
new band at λ = 426 nm is observed (Fig. 6). Maximum absorb-
ance of the new band is reached after approximately half an
hour. These spectral changes can be assigned to the reduction
of the FeIII(tpps) dimer to the FeII(tpps) dimer and are in full
agreement with the literature.12,13,18,31 Following the formation
of the FeII(tpps) band, which remains stable for about 1 h, a
slow back reaction takes place and the spectral changes show
that the FeII(tpps) band slowly decreases and after about 15 h
the FeIII(tpps) band is formed again (Fig. 7). The intensity of
the FeIII(tpps) band is low in comparison to its intensity at the
beginning of the reaction, which demonstrates that, beside the
back oxidation to FeIII(tpps), decomposition of the complex
also occurs. Thus three different reactions takes place: reduc-
tion of FeIII(tpps) to FeII(tpps), back oxidation and overall
decomposition.
oxygen SO2 Ϫ, which is a product of the aimsa decomposition
᝽
Ϫ
in alkaline solutions, rapidly produces O2 according to reac-
᝽
tion (9). We assume that SO2 Ϫ, being a radical, reacts with
᝽
biradical O2 more rapidly than sulfoxylate; the same situation
was observed for dithionite.21 The radical O2 Ϫ is known in the
᝽
literature32 to oxidize iron() complexes to FeIII and to produce
hydrogen peroxide, which is also able to oxidize FeII to FeIII,
equations (10)–(12). Thus a whole range of species can be
Fe2ϩ ϩ O2 Ϫ ϩ 2 H2O → Fe3ϩ ϩ H2O2 ϩ 2 OHϪ (10)
᝽
Ϫ
Fe2ϩ ϩ H2O2 → Fe3ϩ ϩ OH ϩ OH
(11)
(12)
ؒ
Fe2ϩ ϩ OH → Fe3ϩ ϩ OHϪ
ؒ
formed, which can all oxidize and decompose FeIII(tpps). The
more oxygen present in solution, the more of these species can
be formed.
In order to investigate the influence of aimsa on the redox
reactions, the oxygen concentration was kept constant (air-
saturated solutions) and the reaction between FeIII(tpps) and
aimsa was performed using different aimsa concentrations. In
In order to investigate the influence of oxygen on these reac-
tions the same experiment was performed with air-saturated
solutions. The spectral changes clearly showed that in the pres-
ence of oxygen the formation of FeII(tpps) is inhibited, whereas
decomposition and the back reaction are accelerated. The
2918
J. Chem. Soc., Dalton Trans., 1998, Pages 2915–2920

View Article Online
Fig. 9 (a) Rapid-scan spectra recorded for the reaction between the
FeIII(tpps) dimer and aimsa in an air-saturated strongly basic solution.
[FeIII(tpps)] = [aimsa] = 6 × 10Ϫ6 , [NaOH] = 0.5 , 25 ЊC, ∆t = 20 s.
(b) Absorbance vs. time trace for the reaction between the FeIII(tpps)
Fig. 10 (a) Rapid-scan spectra recorded for the reaction between the
FeIII(tpps) dimer and dithionite in a basic solution in the presence of
air. [FeIII(tpps)] = 6 × 10Ϫ6 , [S2O42Ϫ] = 6 × 10Ϫ5 , pH 11; I = 0.1 ,
25 ЊC, ∆t = 4 s. (b) Absorbance vs. time trace for the reaction between
the FeIII(tpps) dimer and dithionite in basic solution in the presence of
air. [FeIII(tpps)] = 6 × 10Ϫ6 , [S2O42Ϫ] = 6 × 10Ϫ5 , pH 11, I = 0.1 ,
25 ЊC, λ = 421 nm
dimer and aimsa in an air-saturated strongly basic solution. [FeIII
(tpps)] = [aimsa] = 6 × 10Ϫ6 , [NaOH] = 0.5 , 25 ЊC, λ = 428 nm
-
the case of a small increase or a small decrease in the aimsa
concentration no significant change could be observed in terms
of the participating redox reactions. If the aimsa concentration
is however decreased significantly (from 7.5 × 10Ϫ4 to 5 × 10Ϫ5
) the spectral changes show a different behaviour [Fig. 8(a)
and 8(b)]. Only a decay in the Soret band of the FeIII(tpps)
dimer due to decomposition (cycle time 15 min) and no form-
ation of the FeII(tpps) band are observed. Obviously the ratio
between oxygen and aimsa plays an important role in the redox
reactions. In this case the oxygen concentration is higher than
the aimsa concentration, and the oxidation as well as decom-
position of the complex is favoured. If the aimsa concentration
is higher than the oxygen concentration, partial reduction of
the complex can be observed along with its decomposition and
back oxidation. In order to increase the concentration of rad-
icals the same experiment was repeated in a strongly alkaline
medium ([NaOH] = 0.5  after mixing) where the decom-
position of aimsa is more efficient than at pH 11. The spectral
changes and the absorbance vs. time trace at 428 nm clearly
show that, under the strongly basic conditions, redox cycling
can be observed [Figs. 9(a) and 9(b)]. The formation of the
FeII(tpps) band is fast (about 30 s) compared to the slow back
(pH ≈ 11) and a large excess of dithionite (6 × 10Ϫ4 ), reduc-
tion of the FeIII(tpps) dimer to the FeII(tpps) dimer occurs. This
observation is in agreement with the literature.12 If oxygen is
added to the FeIII(tpps)–dithionite system decomposition and
oxidation of the metalloporphyrin plays a significant role. Fig.
10(a) shows the spectral changes of the reaction between FeIII-
(tpps) and dithionite (6 × 10Ϫ5 ) in the presence of air
(approximately 50% air-saturated solution). There is no form-
ation of the FeII(tpps) band at λ = 426 nm, but a new band at
λ = 421 nm appears for a short time and decomposes [Fig. 10(a)
and 10(b)]. According to earlier results13 this band can be
assigned to the oxidation of FeIII(tpps) to FeIVO(tppsϩ) or
FeVO(tpps). During the decomposition of the band at 421 nm a
shift back towards the wavelength of the FeII(tpps) dimer band
occurs [Fig. 10(a) and 10(b)]. At very low dithionite concen-
tration (6 × 10Ϫ6 ) and in a strong alkaline medium
([NaOH] = 0.5 ) a redox cycle between FeIII(tpps) and FeII-
(tpps) does occur [Fig. 11(a) and 11(b)] as was observed in the
case of aimsa under these conditions. In agreement with the
redox cycle observed for aimsa, the formation of FeII(tpps) is
fast (approximately 30 s) whereas the back oxidation is slow
[about 1000 s; Fig. 11(b)]. The intensity of the FeII(tpps) band
observed during the redox cycling with dithionite is much lower
than the FeII(tpps) band observed during the redox cycling with
aimsa [Figs. 9(a), 9(b) and 11(a), 11(b)] indicating that more
FeII(tpps) is formed when aimsa is used.
oxidation [more than 1000 s; Fig. 9(b)]. This experiment
Ϫ
demonstrates that the formation of the SO2 radical during
᝽
the decomposition of aimsa, as well as the ratio between the
radical and oxygen concentration, determine the redox
processes.
Reaction of dithionite with FeIII(tpps) in alkaline solutions
Conclusion
Similar experiments as in the case of aimsa were also performed
with dithionite. In the case of nitrogen-saturated solutions
The experiments show that both aimsa and dithionite react
in a very similar way with FeIII(tpps) in alkaline solutions.
J. Chem. Soc., Dalton Trans., 1998, Pages 2915–2920
2919

View Article Online
Fe2ϩ ϩ O2 Ϫ ϩ 2 H2O → Fe3ϩ ϩ H2O2 ϩ 2 OHϪ (10)
᝽
Ϫ
Fe2ϩ ϩ H2O2 → Fe3ϩ ϩ OH ϩ OH
(11)
(12)
ؒ
Fe2ϩ ϩ OH → Fe3ϩ ϩ OHϪ
ؒ
rate determining step will be the decomposition of aimsa to
Ϫ
form the SO2 radicals. The last four reactions are important
᝽
when oxygen is present (especially in excess) and can account
for the oxidation and decomposition of the metalloporphyrin.
Acknowledgements
The award of a fellowship by the Deutsche Akademische Aus-
tauschdienst to S. V. Makarov is greatly appreciated.
References
1 E. B. Barnett, J. Chem. Soc., 1910, 97, 63.
2 T. K. Das, A. K. Mandavawalla, A. Lahiri and S. K. Datta,
Colourage, 1984, 31, 15.
3 M. Arifoglu, W. N. Marmer and R. R. Dudley, Text. Res. J., 1992,
62, 94.
4 J. Drabowicz and M. Mikolaiczyk, Synthesis, 1978, 8, 542.
5 E. S. Lang and J. V. Comasseto, Synth. Commun., 1988, 18, 301.
6 J. T. B. Ferreira, A. R. M. De Oliveira and J. B. Comasseto, Synth.
Commun., 1989, 19, 239.
7 W. Y. Huang and J. L. Zhuang, Chin. Chem. Lett., 1990, 1, 191.
8 W. B. Koniecky and A. L. Linch, Anal. Chem., 1958, 30, 1134.
9 J. E. McGill and F. Lindstrom, Anal. Chem., 1977, 49, 26.
10 M. H. El-Rafie, M. K. Zahran, K. F. El Tahlawy and A. Hebeish,
Polym. Degrad. Stab., 1995, 47, 73.
11 S. A. Abdel-Hafiz, J. Appl. Polym. Sci., 1995, 58, 2005.
12 A. El Awady, P. C. Wilkins and R. G. Wilkins, Inorg. Chem., 1985,
24, 2053.
13 V. Lepentsiotis, R. van Eldik, F. F. Prinsloo and J. J. Pienaar,
unpublished work; V. Lepentsiotis and R. van Eldik, J. Chem. Soc.,
Dalton Trans., 1998, 999.
14 E. B. Fleischer, J. M. Palmer, T. S. Srivastava and A. Chatterjee,
J. Am. Chem. Soc., 1971, 93, 3162.
15 C. R. Chinake, R. H. Simoyi and S. B. Jonnalagadda, J. Phys.
Chem., 1994, 98, 545.
16 M. P. Mefodieva, Radiokhimya, 1974, 16, 193 (in Russian).
17 G. Rabai, R. T. Wang and K. Kustin, Int. J. Chem. Kinet., 1993, 25,
53.
Fig. 11 (a) Rapid-scan spectra recorded for the reaction between the
FeIII(tpps) dimer and dithionite in an air-saturated strongly basic solu-
tion. [FeIII(tpps)] = [S2O42Ϫ] = 6 × 10Ϫ6 , [NaOH] = 0.5 , 25 ЊC,
∆t ≈ 20 s. (b) Absorbance vs. time trace for the reaction between the
FeIII(tpps) dimer and dithionite in an air-saturated strongly basic solu-
tion. [FeIII(tpps)] = [S2O42Ϫ] = 6 × 10Ϫ6 , [NaOH] = 0.5 , 25 ЊC,
λ = 406 nm
Ϫ
Obviously the SO2 radical plays a key role during the redox
᝽
reactions. The ratio of the oxygen and radical concentrations
determines which reaction (oxidation, reduction or decom-
position) will dominate. In general a high oxygen concentration
and a low radical concentration favour the oxidation and
decomposition of the metalloporphyrin, whereas high concen-
trations of radicals and low oxygen concentration favour the
reduction. In strong basic solutions ([NaOH] = 0.5 ) a redox
cycle between FeIII(tpps) and FeII(tpps) was observed at low
aimsa and dithionite concentration.
18 K. Hatano and Y. Ishida, Bull. Chem. Soc. Jpn., 1982, 55, 3333.
19 D. O. Lambeth and G. Palmer, J. Biol. Chem., 1973, 248, 6095.
20 R. E. Huie, C. L. Clifton and N. Altstein, Radiat. Phys. Chem., 1989,
33, 361.
21 V. V. Budanov and S. V. Makarov, Chemistry of Sulfur-containing
Reducing Agents, Khimia (Chemistry), Moscow, 1994, p. 140 (in
Russian).
22 J. S. Song, E. H. Kim, S. K. Kang, S. S. Yun, I. H. Suh, S. S. Choi,
S. Lee and W. P. Jensen, Bull. Korean Chem. Soc., 1996, 17, 201.
23 G. Gattow and W. Manz, Z. Anorg. Allg. Chem., 1988, 561, 66.
24 D. R. Eaton and K. Zaw, Can. J. Chem., 1971, 49, 3315.
25 Y. Chun, I. Yavari and J. D. Roberts, Org. Magn. Reson., 1982, 18,
74.
The proposed mechanism for the reaction between FeIII(tpps)
and aimsa consists of reactions (1), (2), (5) and (8)–(14). The
26 C. W. J. Scaife and R. G. Wilkins, Inorg. Chem., 1980, 19, 3244.
27 H. S. Kim, B. Claude and C. Tondre, J. Phys. Chem., 1990, 94, 7711.
28 S. G. Mayhew, Eur. J. Biochem., 1978, 85, 535.
29 C. E. McKenna, W. G. Gutheil and Wei Song, Biochim. Biophys.
Acta, 1991, 1075, 109.
30 A. E. Miller, J. J. Bischoff and K. Pae, Chem. Res. Toxicol., 1988, 1,
169.
31 S. M. Chen, J. Electroanal. Chem. Interfacial Electrochem., 1996,
407, 123.
32 C. Brandt and R. van Eldik, Chem. Rev., 1995, 95, 119 and refs.
therein.
(NH2)2CSO2
NH2(NH)CSO2H
(1)
NH2(NH)CSO2H ϩ OHϪ
NH2(NH)CSO2Ϫ ϩ H2O (2)
2Ϫ
NH2(NH)(OH)CSO2
NH2(NH)(OH)C Ϫ ϩ SO2
(5)
(8)
Ϫ
᝽
᝽
Ϫ
2Ϫ
2 SO2
S2O4
᝽
SO2 Ϫ ϩ HO2Ϫ → products (k = 5.9 Ϫ1 sϪ1 33
)
(13)
(14)
33 C. Creutz and N. Sutin, Inorg. Chem., 1974, 13, 2041.
᝽
Fe3ϩ ϩ SO2 Ϫ → Fe2ϩ ϩ SO2
᝽
SO2 Ϫ ϩ O2 → SO2 ϩ O2 Ϫ (k = 2.4 × 109 Ϫ1 sϪ1) (9)
Received 22nd June 1998; Paper 8/04684K
᝽
᝽
2920
J. Chem. Soc., Dalton Trans., 1998, Pages 2915–2920