Formation of a promazine radical and promazine 5-oxide in the reaction of promazine with hydrogen peroxide: Mechanistic insight from kinetic and epr measurements

Article abstract of DOI:10.1002/kin.20454

The kinetics of the oxidation of promazine (PMZ) by hydrogen peroxide was studied in the presence of a large excess of H₂O₂ in acidic chloride media using UV-vis spectroscopy. The reaction proceeds via two consecutive steps. In the first step, oxidation leads to formation of a promazine radical. In the second step, the promazine radical is oxidized to promazine 5- oxide. Electron paramagnetic resonance spectroscopy (EPR) results provide clear evidence for the formation of an intermediate promazine radical. Linear dependences of the pseudo-first- order rate constants (k₁ and k₂)on[H₂O₂] with a nonzero intercept were established for the first and the second process, respectively. The rate of the first stage of the reaction increased slightly with increasing concentration of O₂, indicating the role of the OH? radicals on the redox process, which are transformed into the Cl₂-? radicals. The mechanism of the overall reaction is discussed on the basis of all these kinetic measurements.

Full text of DOI:10.1002/kin.20454

Formation of a Promazine
Radical and Promazine
5-Oxide in the Reaction of
Promazine with Hydrogen
Peroxide: Mechanistic
Insight from Kinetic and
EPR Measurements
´
JOANNA WISNIEWSKA, GRZEGORZ WRZESZCZ, STANISLꢀ AW KOTER
Department of Chemistry, Nicolaus Copernicus University, 87-100 Torun´, Poland
Received 4 August 2008; revised 21 January 2009, 28 May 2009; accepted 15 July 2009
DOI 10.1002/kin.20454
Published online in Wiley InterScience (www.interscience.wiley.com).
ABSTRACT: The kinetics of the oxidation of promazine (PMZ) by hydrogen peroxide was studied
in the presence of a large excess of H2O2 in acidic chloride media using UV–vis spectroscopy.
The reaction proceeds via two consecutive steps. In the first step, oxidation leads to formation
of a promazine radical. In the second step, the promazine radical is oxidized to promazine 5-
oxide. Electron paramagnetic resonance spectroscopy (EPR) results provide clear evidence for
the formation of an intermediate promazine radical. Linear dependences of the pseudo-first-
order rate constants (k1 and k2) on [H2O2] with a nonzero intercept were established for the
first and the second process, respectively. The rate of the first stage of the reaction increased
slightly with increasing concentration of O2, indicating the role of the OH• radicals on the redox
process, which are transformed into the Cl2−• radicals. The mechanism of the overall reaction
is discussed on the basis of all these kinetic measurements. ꢀ 2009 Wiley Periodicals, Inc. Int
C
J Chem Kinet 42: 1–9, 2010
INTRODUCTION
Although considerable research has already been car-
ried out in an attempt to understand the role of free
Correspondence to: Joanna Wis´niewska; e-mail: wisnia@
chem.uni.torun.pl.
Contract grant sponsor: Nicolaus Copernicus University.
Contract grant number: 311-Ch.
2009 Wiley Periodicals, Inc.
radicals in the pathogenesis of many diseases and in
normal healthy cell metabolism, the Þeld still remains
active in most disciplines in the life sciences [1,2].
Among others, reactive oxygen species are involved
c
ꢀ

´
WISNIEWSKA, WRZESZCZ, AND KOTER
2
in a wide variety of physiological and pathophysio-
logical processes [3,4]. These active oxygen species,
such as the superoxide anion (O−2 •), hydrogen peroxide
(H2O2), and hydroxyl radical (OH•) can interact with
cellular membrane structures and induce lipid perox-
idation, which has been indicated in oxidative stress
in cells [2]. However, only OH• radical can directly
initiates a lipoperoxidation, but H2O2 and O−2 • induce
this process indirectly. Hydrogen peroxide is one of
the products in the dismutation of the superoxide an-
ion radicals catalyzed by superoxide dismutases and
appears as an intermediate in many reactions of or-
ganic material with dioxygen. Dioxygen is unreactive
toward many organic species but can be activated by
reduction to HO2/O2−•, H2O2, OH•, or by coordination
to transition metal ions [5]. It is reasonable that hy-
drogen peroxide may oxidize organic substances more
efÞciently than oxygen. Under physiological condition
in cells, hydrogen peroxide oxidizes thiols, indoles,
and the imidazole, phenolic, thioester, and methionine
groups in other biologically active species [6].
Promazine (Scheme 1), as have many other N-
substituted phenothiazines, has been found to modify
various neurotransmission processes [7]. In medicine,
phenothiazines are also applied as antiarrhythmic
drugs or photodynamic virus inactivators [8]. Their
efÞciency in affecting the antibiotic resistance of bac-
teria and tumor cells is the subject of continuous and
intense investigations [9]. A quenching action of the
phenothiazines on radical formation has been found to
be responsible for inhibitory effects on lipid peroxi-
dation [10,11] and characterized redox activity of this
class of tricyclic compounds.
cation radical and copper(I) was reoxidized by dioxy-
gen to copper(II). Dioxygen dissolved in the solution
was the direct oxidant in the next process leading to
sulfoxidation of the cation radical. Direct interaction
of O2 with paramagnetic reductants results in an ap-
pearance of H2O2 as intermediate [12]. The particular
emphasis of the present work was to understand the
kinetics and mechanism of the nonenzymatic and non-
catalyzed oxidation of promazine by hydrogen per-
oxide. The redox activity of hydrogen peroxide re-
vealed the redox potential, which is equal to 0.38 V for
the H2O2/OH• couple and 1.35 V for the H2O2/H2O
species (pH 7) [13]. The hydroxyl radical generated in
the reduction of hydrogen peroxide is known to be the
most reactive oxygen species with a redox potential
for the OH•/H2O couple of 2.31 V (pH 7) [13] higher
than for the H2O2/H2O compounds. The hydroxyl rad-
ical reacts usually at rates close to diffusion-controlled
with practically all biological molecules, its half-life
in cells has been estimated to be 10−9 s [14]. In the
presence of H2O2 in aqueous solution, the OH• radical
can be transformed into the O−2 • (HO2) radical anion,
which is less reactive toward most biological materials
than the OH• radical. In the present work, the OH•
production in the reaction was studied using the spin-
trapping method of electron paramagnetic resonance
spectroscopy (EPR) to establish further experimental
evidence its inßuence on the mechanism of the re-
action, although this inßuence is not direct. In acidic
chloride solutions, the OH• radical reacts with chloride
ion and oxidizes to a dichloro radical anion, Cl−2 •. This
radical is a very strong oxidant; its redox potentials is
equal to 2.30 V (Cl−2 •/2Cl−) [15]. The rate constants of
the one-electron oxidation reactions of the Cl−2 • radical
with numerous organic and inorganic compounds were
determined using pulse radiolysis and ßash photolysis
experiments [16].
In an earlier article, we reported on the kinetics
of the slow oxidation of promazine by dioxygen in
solutions of aqueous copper(II) halides [12]. In this re-
action, the copper(II) halides oxidized promazine to a
O
S
S
S
+
N
N
N
N
N
N
.
PMZ
PMZ+
PMZO
Scheme 1
International Journal of Chemical Kinetics DOI 10.1002/kin

FORMATION OF A PROMAZINE RADICAL AND PROMAZINE 5-OXIDE
3
EXPERIMENTAL
Materials
one scan only was taken for each measurement. The
spectra were scanned every 10 min during reaction
time. The microwave frequency was monitored with
a frequency divider. The magnetic Þeld was measured
with an automatic MNR-type magnetometer.
Hydrogen peroxide (30%; Sigma-Aldrich, St.
Louis, MO, USA), DMPO (5,5-dimethyl-1-pyrroline
N-oxide; Aldrich, Milwaukee, WI, USA) used in spin-
trapping experiments, and all other chemicals were an-
alytical grade reagents. Water puriÞed in a Millipore
Milli-Q system was used to prepare all the solutions.
The concentration of a stock solution of hydrogen per-
oxide was determined manganometrically. The solu-
tion was stable for 1 month when stored at 5◦C. Hy-
drogen peroxide is unstable at higher temperatures and
undergoes a dismutation reaction, which can be facili-
tated by Na+, K+ ions, as well as by traces of transition
metal ions [17]. The stability of hydrogen peroxide has
been analyzed in different media using manganomet-
ric titrations. The loss of H2O2 in the HCl solution was
equal 0.4% within the standard errors (t = 500 min,
[H2O2] = 0.05 mol dm−3, [H+] = 0.1 mol dm−3, I =
0.1 mol dm−3 (H+, Cl−), T = 298 K), in water it was
1%, while in an aqueous solution of NaCl this percent-
Kinetic Measurements
The kinetic measurements were performed on a
Lambda 20 Perkin–Elmer spectrophotometer equipped
with an external Julabo F25 ultrathermostat, which en-
abled all the solutions to be thermostatted to the de-
sired temperature (298
0.1 K). Kinetic data were
acquired and processed using the UV KinLab (Perkin–
Elmer) software. Kinetic runs were initiated by mixing
a promazine solution with a solution containing hydro-
gen peroxide and hydrochloric acid. For a few runs, at
pH > 1, NaCl was included to maintain constant ionic
strength, I = 0.1 mol dm−3 (H+, Na+, Cl−). The ini-
tial promazine concentration in reaction solutions was
usually 5 × 10−5 mol dm−3. The concentrations of
other reagents were as follows: [H2O2] = (5–35) ×
10−3 mol dm−3, [H+] = 0.02–0.1 mol dm−3. In a few
experiments, kinetics was studied in the presence of
transition metal ions scavenger, disodium ethylenedi-
aminetetraacetate dihydrate ([EDTA] = 1 × 10−3 mol
dm−3). Because the rate of reaction was not affected by
iron and copper chelator, the measurements were done
without this species. Progress of the reaction was fol-
lowed spectrophotometrically at 514 nm (a maximum
for the promazine radical) and at 345 nm (a maximum
for promazine 5-oxide in 0.1 mol dm−3 HCl where pro-
mazine radical and H2O2 absorbed negligibly). Usu-
ally, the kinetic measurements were done in the air
atmosphere, but in a few series of experiments, the
reaction was studied under either an oxygen or argon
atmosphere. The spectra were scanned every 100 s up
to 2t1/2. An excess of H2O2 over promazine was used
to simplify the reactions into the pseudo-Þrst-order
processes. The pseudo-Þrst-order rate constants were
calculated from a multiexponential dependence for
two consecutive one-electron processes A → B → C
with Þxed value of k2 because of the long reaction
time [18]. The kinetic traces at 345 nm represented
changes in the concentration of the Þnal product, and
no induction time was observed. On account of this
fact, a minimalization of the sum of the squares of the
deviations for the absorbance dependence versus time,
calculated for C, was performed from the Þrst-order
dependence. The kinetic runs were performed in a se-
ries for various reactant concentrations. The relative
standard errors of the pseudo-Þrst-order rate constants
for a single kinetic trace were ca. 0.2%.
age was 10% (t = 500 min, [H2O2] = 0.05 mol dm−3
,
I = 0.1 mol dm−3 (Na+, Cl−), T = 298 K). Promazine
(Sigma-Aldrich) and its aqueous solution were stored
in the dark because of their photosensitivity. Argon and
oxygen gases of high purity were applied in the exper-
iments. EDTA (disodium ethylenediaminetetraacetate
dihydrate) was purchased from Sigma and was used to
mask transition metal ions in the solutions.
HPLC Measurements
The Þnal product, promazine 5-oxide, was estimated
on a Shimadzu HPLC chromatograph equipped with a
C18 column (Supelco; 0.46-cm i.d., 15-cm i.l., 5-μm
particles), Rheodyne 7125 sample injector, Shimadzu
SPD-10A (VP) UV–vis detector (250 nm), and Shi-
madzu LC-10AD (VP) pump. A mobile phase compo-
sition (90 vol% of acetonitrile and 10 vol% of 50 mmol
dm−3 K3PO4, pH 7) was performed for chromato-
graphic separations.
EPR Measurements
EPR spectra were recorded with Radiopan SE/X 2547
spectrometer in the X band (ca. 9.25 GHz) using a ßat
quartz cell. Typical parameters for acquisition of EPR
spectra were as follows: central Þeld, 330 mT; sweep
width, 10 mT; microwave power, 16 mW; modulation
frequency, 100 kHz; modulation amplitude, 0.1 mT;
time constant, 1 s; sweep time, 120 s; and receiver gain,
(1.6–20) × 104. Because of long sweep time (120 s),
International Journal of Chemical Kinetics DOI 10.1002/kin

´
4
WISNIEWSKA, WRZESZCZ, AND KOTER
1.0
0.9
0.8
0.7
0.6
0.5
0.4
0.3
0.2
0.1
0.0
300
350
400
450
500
550
600
Wavelength, nm
Figure 1 Spectral changes during the oxidation of promazine by hydrogen peroxide. Experimental conditions: [H2O2] =
0.035 mol dm−3, [PMZ] = 5 × 10−5 mol dm−3, [H+] = 0.1 mol dm−3, I = 0.1 mol dm−3 (H+, Cl−), 298 K, ꢀt = 30 min,
ttotal = 495 min, 17 scans.
ical, PMZ+•, which absorbs intensively in the visible
range (Fig. 1) at 514 nm (λmax, nm (εmax, dm3 mol−1
cm−1): 513(9200)) [19] and to promazine 5-oxide de-
tected within the UV range after the HPLC separation.
The spectra of promazine 5-oxide revealed two low-
Þeld electronic transition bands at 335 and 297 nm
(Fig. 1) in the ultra-violet spectral range (λmax, nm
(εmax, dm3 mol−1 cm−1): 228(24,300), 269(10,300),
297(6500), 335(4700)) [20]. The reaction was moni-
tored in the acidic chloride media, in which promazine
exists in its protolytic form with a protonated nitro-
gen atom in the alkyloamine group in a substituent. A
positive charge of this substituent is not included in
Scheme 1.
Formation of the PMZ+• radical during oxidation of
promazine by hydrogen peroxide was observed using
EPR spectroscopy. Figure 2 shows the EPR spectrum of
the reaction system after ca. 15 min from the initiation,
and clearly indicates the presence of the PMZ+• rad-
ical. The spectrum has unresolved hyperÞne lines but
its g value and overall shape is consistent with the EPR
spectrum of PMZ+•, published previously [12,21]. The
EPR signal of promazine radical with g = 2.005 could
be explained as arising mainly from coupling of the un-
paired electron with heterocyclic ring nitrogen nucleus
(IN = 1) and two equivalent β-protons (IH = 1/2) of
the side chain (aN ∼ 0.7 mT, aH(β) ∼ 0.35 mT), and
also from weaker coupling with the four proton pairs
of the lateral rings (aH(β) ꢁ aH(ring)). Owing to the
large number of lines, the relatively low spin density
RESULTS AND DISCUSSION
The reaction between promazine (PMZ) and hydrogen
peroxide was followed in the aqueous acidic media
under an argon, air, or oxygen atmosphere ([PMZ] =
5 × 10−5 mol dm−3, [H2O2] = (5–35) × 10−3 mol
dm−3, [H+] = 0.02–0.1 mol dm−3, I = 0.1 mol dm−3
(H+, Na+, Cl−), T = 298 K) (Table I). The oxidation
of promazine leads to the appearance of the cation rad-
Table I Observed Pseudo-First-Order Rate Constants
for the Reaction between Promazine (k1) or Promazine
Radical (k2) and Hydrogen Peroxide
Atmosphere [H2O2] (mol dm−3) 105 k1 (s−1) 105 k2 (s−1
)
Argon
0.0049
0.0148
0.0247
0.0346
0.0049
0.0148
0.0247
0.0346
0.0049
0.0148
0.0247
0.0346
5.60
6.88
7.76
8.49
5.93
7.03
8.34
9.34
6.51
8.18
9.97
1.82
3.11
3.98
5.38
1.59
2.62
3.79
5.06
1.62
2.69
4.25
5.16
Air
Oxygen
11.5
[PMZ] = 5 × 10−5 mol dm−3, T = 298 K, pH 1, I = 0.1 mol
dm−3 (H+, Cl−).
International Journal of Chemical Kinetics DOI 10.1002/kin

FORMATION OF A PROMAZINE RADICAL AND PROMAZINE 5-OXIDE
5
2.025
2.015
2.005
1.995
1.985
g
Figure 2 Simulated (top) and experimental (bottom) EPR spectra of the promazine cation radical generated via oxidation of
the acidic solution of 6 × 10−3 M promazine by 5 × 10−2 M H2O2 under air atmosphere. Experimental conditions: [HCl] =
1 × 10−1 M, I = 0.1 (H+, Cl−), T = 295 K, signal after ca. 15 min, modulation amplitude, 0.02 mT, microwave frequency,
9.257 GHz; other parameters as described in the Experimental section.
of the unpaired electron at the lateral ring protons, and
also incomplete averaging of the g factor and/or hyper-
Þne anisotropy, the EPR spectrum of promazine radical
cation remains poorly resolved.
dent of the acid concentration within the 0.02–0.1 mol
dm−3 range (data not shown).
Plots of k2 versus [H2O2] for the second stage—
the oxidation of the cation radical (PMZ+•) to pro-
mazine 5-oxide—are presented in Fig. 4. The pseudo-
Þrst-order rate constants were found to be linearly de-
pendent on the H2O2 concentration. The rate constants
did not differ substantially with increasing [O2] and the
differences in the k2 values did not affect the intercepts
and the slopes under the oxygen atmosphere.
These kinetic data can be rationalized by the fol-
lowing reaction scheme:
The preliminary experiments showed the following
interesting features of the reaction: (i) the appearance
of promazine 5-oxide follows in two consecutive re-
actions: formation of the promazine radical and its
disappearance leading to promazine 5-oxide; (ii) the
highest stability of hydrogen peroxide was observed in
acidic aqueous media, i.e. loss of 0.05 mol dm−3 H2O2
in acidic solution (0.1 mol dm−3 HCl) was less than
0.4%, in H2O was equal to 1%, while in neutral solu-
tion of salt (0.1 mol dm−3 NaCl) decomposition was
up to 10% during 500 min at 298 K; (iii) the oxida-
tion of promazine in the solution at the ionic strength,
I = 1.0 mol dm−3 (H+, Na+, Cl−), T = 298 K follows
two-fold faster than at I = 0.1 mol dm−3 (H+, Cl−)
under the same conditions and temperature.
For the oxidation of promazine to the cation radical,
the pseudo-Þrst-order rate constants, k1, were found to
be linearly dependent on [H2O2] and were higher in
the presence of oxygen than under the argon atmo-
sphere. The change of atmosphere affected the slopes,
but the statistically insigniÞcant, nonzero intercepts in
the plots of k1 versus [H2O2] are illustrated in Fig. 3.
The observed rate constants were found to be indepen-
H+
PMZ + H2O2 −→ PMZ+• + OH•
+ H2O (rate determining)
PMZ+• + H2O2 −→ PMZ2+ + OH•
(1)
H+
+ H2O (rate determining)
(2)
(3)
PMZ2+ + H2O −→ PMZO + 2H+
−•
− −→
OH• + Cl
ClOH
←−
4.3 × 109mol−1dm3s−1 [22]
ClOH−• −→ Cl• + H2O
(4)
(5)
H+
2.1 × 1010mol−1dm3s−1 [22]
International Journal of Chemical Kinetics DOI 10.1002/kin

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WISNIEWSKA, WRZESZCZ, AND KOTER
0.00012
0.00010
0.00008
k
0.00006
0.005
0.010
0.015
0.020
0.025
0.030
0.035
[H2O2], mol dm
Figure 3 Plots of k1 versus [H2O2] for the oxidation of promazine with hydrogen peroxide. Experimental conditions:
[PMZ] = 5 × 10−5 mol dm−3, T = 298 K, pH 1, I = 0.1 mol dm−3 (H+, Cl−), ( ) argon, ( ) air, (ꢁ) oxygen atmosphere.
ꢀ
•
PMZ+• + O2 + H+ −→ PMZ2+
+ HO•2 (very slow process)
2PMZ+• −→ PMZ
−•
− −→
←−
Cl• + Cl
Cl
2
8.0 × 109mol−1dm3s−1 [23]
PMZ + Cl−2 • −→ PMZ+• + 2Cl−
5 × 109mol−1dm3s−1 [24]
(6)
(7)
(8)
(9)
+ PMZ2+ (very slow process)
0.000055
0.000050
0.000045
0.000040
0.000035
0.000030
0.000025
0.000020
0.000015
0.005
0.010
0.015
0.020
0.025
0.030
0.035
[H2O2], mol dm
Figure 4 Plots of k2 versus [H2O2] for the oxidation of promazine radical with hydrogen peroxide. Experimental conditions:
[PMZ] = 5 × 10−5 mol dm−3, T = 298 K, pH 1, I = 0.1 mol dm−3 (H+, Cl−), ( ) argon, ( ) air, (ꢁ) oxygen atmosphere.
ꢀ
•
International Journal of Chemical Kinetics DOI 10.1002/kin

FORMATION OF A PROMAZINE RADICAL AND PROMAZINE 5-OXIDE
7
Promazine undergoes the oxidation reaction with
hydrogen peroxide according to reaction (1), in which
the hydroxyl and promazine radicals are produced.
The promazine radical as an intermediate product is
transformed Þnally into the promazine dication (2),
which hydrolyzes with the diffusion-controlled rate to
promazine 5-oxide (3). The PMZO is a Þnal product
of the overall reaction with H2O2. It is known that
the hydroxyl radical reacts unselectively with organic
species, and hydroxo derivatives or other character-
istic products are produced in the reaction with this
radical. As is seen in the reaction scheme, the hy-
droxyl radical is scavenged by chloride ion (4, 5),
which is transformed into the dichloro radical anion,
Cl−2 • (6). Under the conditions used in the kinetic ex-
periments, when [H+, Cl−] = 0.1 mol dm−3, OH•
almost never reacts with promazine. The second-order
rate constant for the reaction of OH• with chlorpro-
mazine is equal to 1.5 × 1010 mol−1 dm3 s−1 [25]
and is similar to the rate constant for the reaction of
Cl−2 • with this phenothiazine derivative. Thus, the re-
actions of the hydroxyl radical with PMZ and PMZ+•
are suppressed and the dichloro radical anion reacts
with PMZ in reaction (7). A very slow oxidation of
PMZ+• by dioxygen produces the superoxo radical (8).
Preliminary experiments showed that the HO•2 radical
in acidic media did not oxidize PMZ under the con-
centration conditions used in these experiments, when
the HO•2 radical is present in a very low steady-state
concentration. In the acidic media, the HO•2 radical
326
328
330
332
334
B (mT)
Figure 5 Experimental EPR spectrum of radicals gener-
ated via oxidation of the acidic solution of 0.1 M DMPO
by 5 × 10−2 M H2O2 under air atmosphere. Experimen-
tal conditions: [HCl] = 1 × 10−1 M, I = 0.1 (H+, Cl−),
T = 295 K, signal after ca. 23 min, modulation amplitude,
0.1 mT, microwave frequency, 9.256 GHz; other parameters
as described in the Experimental section.
stant of the OH• radical reaction with Cl−, 4.3 × 109
mol−1 dm3 s−1 [22]. For that reason, the concentration
of DMPO was high, ca. 0.1 mol dm−3, and their reac-
tion with OH• can proceed. A signal of a DMPO·OH•
adduct was also detected in the same solution of H2O2
([DMPO] = 0.1 mol dm−3, [H2O2] = 0.05 mol dm−3
,
pH 1, I = 0.1 mol dm−3 (H+, Cl−), T = 295 K, sig-
nal after ca. 23 min, modulation amplitude, 0.1 mT,
microwave frequency, 9.256 GHz). The adduct inten-
sity increased during experiments (at the same time
a signal of DMPO·OOH• adduct completely disap-
peared). EPR spectrum of DMPO-H2O2 acidic mixture
is shown in Fig. 5, where 1:2:2:1 quartet characteristic
of DMPO-OH• adduct is clearly seen [32]. Another
weak radical 1:1:1 triplet signal with aN ∼ 1.4 mT is
also seen. It comes probably from the further oxida-
tion or/and degradation of DMPO-OH• adduct. In the
mixture of promazine, H2O2, and DMPO, however, a
very intensive signal characteristic for the PMZ+• rad-
ical was observed, and the signal of DMPO·OH• dis-
(pKa HO = 4.88) [26] disproportionates and recom-
2
bines with other radicals in the solution (E◦ = 0.152 V
(O2/HO2), pH 0, pO = 1 atm) [26]. This suggestion
2
was supported by spin-trapping experiments in EPR
([DMPO] = 0.1 mol dm−3, [H2O2] = 0.05 mol dm−3
,
pH 1, I = 0.1 mol dm−3 (H+, Cl−), T = 295 K, sig-
nal after ca. 23 min, modulation amplitude, 0.1 mT,
microwave frequency, 9.256 GHz), in which a signal
of a DMPO·OOH• adduct was detected directly only
in the solution of H2O2. The rate of the DMPO·OOH•
adduct formation is slow at neutral pH, however at
acidic pH increases and then the second-order rate con-
stant is equal to 27 mol−1 dm3 s−1 [27,28] at pH 6.2
and 103 mol−1 dm3 s−1 at pH 5.0 [27,29]. The high
reactivity of O−2 • to DMPO in acidic pH was proposed
to be due to the protonation of O−2 • to form the hy-
droperoxyl radical, HO•2, and the fact that HO2• is a
stronger oxidizer than O−2 • (E◦ = 1.06 and 0.94 V
at pH 0, respectively) [30]. Although the rate of the
DMPO·OOH• adduct formation is not fast, this species
was observed in the H2O2 solution in the presence of
chloride ions. The rate constant of the OH• radical ad-
dition to DMPO, (2.7 – 4.3) × 109 mol−1 dm3 s−1
[31], is on the same level of magnitude as the rate con-
appeared on the timescale of the experiments (t1/2
=
23 min) [33]. It was mentioned earlier that the loss of
H2O2 was less than 0.4% in 0.1 mol dm−3 HCl and
did not affect the kinetics of the oxidation reaction by
H2O2, but the reactive oxygen species may inßuence
the mechanism of reaction. The rate of the H2O2 de-
composition is 1 × 10−6 mol−1 dm3 s−1 ([H2O2] =
0.1 mol dm−3, pH 10, T = 308 K) [17] and is negligi-
ble on the timescale of this experiment.
The cation radical of promazine is very stable in
the solution [34] and very slowly disproportionates
(9). The second-order rate constant for their dispro-
portionation in 0.1 mol dm−3 H2SO4 calculated from
the dependence of log k on [H2SO4] [34] is equal to
4.55 mol−1 dm3 s−1. This means that the rate of this re-
action, 1 × 10−10 mol dm−3 s−1, is lower than the rate
International Journal of Chemical Kinetics DOI 10.1002/kin

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WISNIEWSKA, WRZESZCZ, AND KOTER
8
Cl−2 • radicals, according to reaction scheme (1)–(9),
Table II Linear Regression Data for the k1
Dependences on [H2O2] (Fig. 3) for the Reaction
between Promazine and Hydrogen Peroxide
II
and the Þrst-order (k ) and second-order rate con-
−2
stants (k2II) for the second stage are almost unaffected
by the reactive oxygen species and oxygen atmosphere.
In previously published data [12], copper(II) in the ox-
idation reaction of promazine by dioxygen accelerated
the Þrst stage of the redox process into the submi-
crosecond scale and the very slow oxidation of PMZ+•
by dioxygen was the rate-limiting step. The product of
the second-order rate constant for the copper(II) halide
mediated oxidation of PMZ+• by dioxygen, calculated
from the activation parameters and the O2 concentra-
k1I (slope)
(mol−1 dm3 s−1
k−I 1 (intercept)
Atmosphere
)
(s−1
)
Argon
Air
Oxygen
(0.96 0.08) × 10−3 (5.27 0.20) × 10−5
(1.17 0.04) × 10−3 (5.35 0.09) × 10−5
(1.71 0.06) × 10−3 (5.68 0.16) × 10−5
T = 298 K, pH 1, I = 0.1 mol dm−3 (H+, Cl−), standard errors.
of the oxidation of PMZ to PMZ+• (1), 5.75 × 10−10
tion, k [O2] = 3 × 10−5 s−1 (I = 1.0 mol dm−3, [O2]
mol dm−3
s
−1, and the rate of oxidation of PMZ+• to
O2
= 1 × 10−3 mol dm−3, T = 298 K) [12] is higher than
PMZO (2), 2.5 × 10−10 mol dm−3 s−1, for 10% of con-
version degree ([H2O2] = 0.0346 mol dm−3, [PMZ] =
5 × 10−5 mol dm−3, [H+] = 0.1 mol dm−3, I = 0.1
(H+, Cl−). The concentration of PMZ+• at the time of
overall reaction is not higher than 30% of total pro-
mazine concentration. Under such circumstances, the
rate of the oxidation reaction of PMZ to PMZ+• (1)
and the PMZ+• disproportionation are similar. How-
ever, including the disproportionation reaction in the
scheme of two one-electron consecutive processes has
not improved kinetic results.
II
−2
the Þrst-order rate constant, k , for the second stage
of the reaction with H2O2; moreover, this last value re-
mains unchanged in different atmospheres. This means
that the very slow reaction of PMZ+• with O2 is inef-
fective as a reaction path in the spontaneous oxidation
using H2O2.
CONCLUSIONS
The results of the present study indicate that the degra-
dation of promazine using hydrogen peroxide proceeds
via two consecutive reactions, which occur with very
similar rates and appear to be inßuenced by reactive
oxygen species, i.e. the OH• radicals, which are trans-
formed into the Cl−2 • radical anions. However, the di-
rect proof of its participation in the reaction was not
obtained. Under physiological condition, the redox po-
tential for H2O2/OH• species, 0.38 V (pH 7), protects
many organic compounds against oxidation by H2O2
[13]. In acidic solution, this redox potential is equal
to 0.8 V (pH 0) [13] and enables production of the
OH• radical in the reaction. The very high stability of
the promazine radical and N-substituted phenothiazine
radicals evoked the very characteristic reaction step in-
volving the oxidative degradation of the promazine,
PMZ → PMZ+• → PMZO, as well as the other N-
substituted phenothiazine drugs and the kinetics of this
process.
The differential equations are too complicated to be
solved. No analytical solutions have been found for
the proposed reaction scheme. Moreover, if all the re-
actions were included, too many variables would have
to be determined simultaneously in the numeric calcu-
lations.
The linear regression data for the dependences of
the observed rate constants k1 and k2 on [H2O2] are
presented in Tables II and III. Thus, those for the Þrst
and second stages of the reaction can be expressed by
the following equations:
k1 = k−I 1 + k1I [H2O2]
k2 = k−II2 + k2II[H2O2]
(10)
(11)
in which the Þrst-order (k−I 1) and second-order rate
constants (k1I ) are dependent on the concentrations of
reactive oxygen forms, which are transformed into
Table III Linear Regression Data for the k2
Dependences on [H2O2] (Fig. 4) for the Reaction
between Promazine Radical and Hydrogen Peroxide
The authors wish to express their gratitude to Prof. Grzegorz
Schroeder and Mr. Jarogniew Bartoszewicz (Faculty of
Chemistry, Adam Mickiewicz University, Poznan´) for their
kind help in carrying out the EPR measurements.
II
−2
k2II (slope)
(mol−1 dm3 s−1
k
(intercept)
(s−1
)
Atmosphere
)
BIBLIOGRAPHY
Argon
Air
Oxygen
(1.17 0.07) × 10−3 (1.27 0.16) × 10−5
(1.17 0.04) × 10−3 (0.96 0.09) × 10−5
(1.23 0.08) × 10−3 (1.00 0.19) × 10−5
1. Mathews, C. K.; van Holde, K. E.; Ahern, K. G.
Biochemistry, 3rd ed.; Addison-Wesley Longman: San
Francisco, CA, 1999; pp. 553–555.
T = 298 K, pH 1, I = 0.1 mol dm−3 (H+, Cl−), standard errors.
International Journal of Chemical Kinetics DOI 10.1002/kin

FORMATION OF A PROMAZINE RADICAL AND PROMAZINE 5-OXIDE
9
2. Suematsu, M.; Schmid-Scho¨nbein, G. W.; Ishimura, Y.;
Tsuchiya, M. In Free Radicals; Punchard, N. A.; Kelly,
J. (Eds.); Oxford University Press: Oxford, UK, 1996;
p. 83.
16. Neta, P.; Huie, R. E.; Ross, A. B. J Phys Chem Ref Data
1988, 17, 1027–1284.
17. Galba´cs Z. M.; Csa´nyi L. J. J Chem Soc, Dalton Trans
1983, 2353–2357.
3. Kraatz, H.-B.; Metzler-Nolte, N. (Eds). Concepts
and Models in Bioinorganic Chemistry; Wiley-VCH:
Weinheim, Germany, 2006.
18. Wilkins, R.G. Kinetics and Mechanism of Reactions of
Transition Metal Complexes, 2nd ed.; VCH: Weinheim,
Germany, 1991; pp. 18–21.
4. Da Silva, J. J. R.; Williams, R. J. P. The Biological
Chemistry of the Elements, 2nd ed.; Oxford University
Press: Oxford, UK, 2001.
5. Fabian, I.; Csordas V. In Advances in Inorganic Chem-
istry, van Eldik R.; Hubbard C. D., (Eds.); Vol. 54,
Academic Press: Amsterdam, 2003; pp. 395–461.
6. Brot, W.; Weissbach, H. Arch Biochem Biophys 1983,
223, 271–281.
19. Pelizzetti, E.; Mentasti, E. Inorg Chem 1979, 18, 583–
588.
20. Wis´niewska, J. Trans Met Chem 2007, 32, 811–815.
21. Ruperez, F. L.; Conesa, J. C.; Soria, J. J Mol Struct 1983,
98, 165.
22. Jayson, G. G.; Parsons, B. J.; Swallow A. J. J Chem Soc,
Faraday Trans 1 1973, 69, 1597–1607.
23. Nagarajan, V.; Fessenden, R. W. J Phys Chem 1985, 89,
2330–2335.
24. Willson, R. L. Biochem Soc Trans 1973, 1, 929–
931.
7. Valzelli, L.; Garattini, S. In Principles of Psychopharma-
cology, Clark, W. G.; del Giudice, J. (Eds.); Academic
Press: New York, 1970; p. 255.
8. Molnar, J.; Tarodi, B.; GalÞ, M.; Matkovics, B.; Moto-
hashi, N. Anticancer Res 1992, 12, 273–280.
9. Kidd, S. E.; Hambley, T. W.; Heve´r, A.; Nelson, M. J.;
Molna´r, J. J Inorg Biochem 1996, 62, 171–181.
10. Yonaha, M.; Tampo, Y.; Clarke, W.; Davis, H. W.;
Schenkman J. B. Arch Biochem Biophys 1992, 292,
62–69.
11. Hara, S.; Endo, T.; Kuriiwa, S; Iwata, N. Res Commun
Chem Pathol Pharmacol 1992, 63, 163–173.
12. Wis´niewska, J.; Kita P.; Wrzeszcz G. Trans Met Chem
2007, 32, 857–863.
13. Sawyer, D. T. In Encyclopedia of Inorganic Chemistry;
King E., (Ed.); Wiley: Chichester, UK, 1994; p. 2951.
14. Kaur, H.; Halliwell, B. In Free Radicals; Punchard, N.
A.; Kelly, J., (Eds.); Oxford University Press: Oxford,
UK, 1996; p. 102.
25. Joshi, R.; Ghanty, T. K.; Mukherjee, T. J Mol Struct
2008, 888, 401–408.
26. Petlicki, J.; van de Ven T. G. M. J Chem Soc, Faraday
Trans 1998, 94, 2763–2767.
27. Burgett, R. A.; Bao, X; Villamena, F. A. J Phys Chem A
2008, 112, 2447–2455.
28. Allouch, A.; Lauricella, R. P.; Tuccio, B. N. Mol Phys
2007, 105, 2017–2024.
29. Finkelstein, E.; Rosen, G. M.; Rauckman, E. J. J Am
Chem Soc 1980, 102, 4994–4999.
30. Buettner, G. R. Arch Biochem Biophys 1993, 300,
535.
31. Kochany, J.; Bolton, J. R. J Phys Chem. 1991, 95, 5116–
5120.
32. Buettner, G. R., Free Rad Biol Med 1987, 3, 259–303.
33. Carmichael A. J.; Makino K.; Riesz P. Radiat Res 1984,
100, 222–234.
15. Wardman P. J. Phys Chem Ref Data 1989, 18, 1711–
1755.
34. Levy L. J Med Chem 1972, 15, 898–905.
International Journal of Chemical Kinetics DOI 10.1002/kin