Kinetic and mechanistic studies of the reactivity of iron(IV) TAMLs toward organic sulfides in water: Resolving a fast catalysis versus slower single-turnover reactivity dilemma

Article abstract of DOI:10.1080/00958972.2015.1065974

TAML complex [Fe^III{C₆H₄-1,2-(NCOCMe₂NCO)₂CMe₂}OH₂]- (1) is oxidized by H₂O₂ or ^tBuOOH in water at pH < 10 into the corresponding iron(IV) -oxo-bridged dimer 2, which oxidizes readily ring-substituted thioanisoles p-XC6H4SMe (X=H, MeO, Me, Cl, CN) into the corresponding sulfoxides with regeneration of 1. The oxidation studied under pseudo-first-order conditions using the stopped-flow technique by monitoring the fading of the 420-nm band of 2 follows hyperbolic kinetics according to the rate law kobs = ab[p-XC6H4SMe]/(1 + b[p-XC6H4SMe]) at pH 8 and 25 °C. Parameters a, b, and ab all decrease for electron-poorer thioanisoles and the Hammett value ρ ~ 1 has been found for ab, which can be associated with the second-order rate constants for oxidation of thioanisoles by 2. The kinetics of oxidation of p-NO2C6H4SMe by H₂O₂ catalyzed by 1 has been studied under steady-state conditions. Covering the concentration of 1 in a 100-fold range has revealed that though first-order kinetics in 1 is observed at low catalyst concentrations (below 10^-6 M), there is a significant negative deviation from linearity at [1] > 10^-6 M. The latter was rationalized by the equilibrium between the monomeric and dimeric Fe^IV species 2 M M-M (Kd), both being able to oxidize p-NO2C6H4SMe with rate constants km and kd which were found to be (13 ± 1) × 10⁴ and (0.32 ± 0.01) × 10⁴ M^-1 s^-1, respectively. The difference in the rate constants is the key for resolving the dilemma of faster catalysis versus slower single-turnover reactivity of TAML activators in water.

Full text of DOI:10.1080/00958972.2015.1065974

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Kinetic and mechanistic studies of the reactivity of
iron(IV) TAMLs towards organic sulfides in water:
Resolving a fast catalysis versus slower single
turnover reactivity dilemma
Deboshri Banerjeea, Alexander D. Ryabova & Terrence J. Collinsa
a Department of Chemistry, Carnegie Mellon University, 4400 Fifth Avenue, Pittsburgh, PA
15213, USA
Accepted author version posted online: 29 Jun 2015.
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To cite this article: Deboshri Banerjee, Alexander D. Ryabov & Terrence J. Collins (2015): Kinetic and mechanistic
studies of the reactivity of iron(IV) TAMLs towards organic sulfides in water: Resolving a fast catalysis versus slower single
turnover reactivity dilemma, Journal of Coordination Chemistry, DOI: 10.1080/00958972.2015.1065974
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Publisher: Taylor & Francis
Journal: Journal of Coordination Chemistry
DOI: http://dx.doi.org/10.1080/00958972.2015.1065974
Kinetic and mechanistic studies of the reactivity of iron(IV) TAMLs towards
organic sulfides in water: Resolving a fast catalysis versus slower single
turnover reactivity dilemma
DEBOSHRI BANERJEE, ALEXANDER D. RYABOV* and TERRENCE J. COLLINS*
Department of Chemistry, Carnegie Mellon University, 4400 Fifth Avenue, Pittsburgh, PA 15213, USA
Dedicated to Prof. Rudi van Eldik on the occasion of his 70th birthday.
TAML complex
(1) is oxidized by H2O2 or
tBuOOH in water at pH < 10 into the corresponding iron(IV) μ-oxo-bridged dimer 2, which
oxidizes readily ring-substituted thioanisoles p-XC6H4SMe (X = H, MeO, Me, Cl, CN) into the
corresponding sulfoxides with regeneration of 1. The oxidation studied under pseudo-first-order
conditions by the stopped-flow technique by monitoring the fading of the 420 nm band of 2
follows hyperbolic kinetics according to the rate law kobs = ab[p-XC6H4SMe]/(1 + b[p-
XC6H4SMe]) at pH 8 and 25 °C. Parameters a, b and ab all decrease for electron-poorer
thioanisoles and the Hammett value ρ ~ 1 has been found for ab, which can be associated with
the second order rate constants for oxidation of thioanisoles by 2. The kinetics of oxidation of
p-NO2C6H4SMe by H2O2 catalyzed by 1 has been studied under steady-state conditions.
Covering the concentration of 1 in a 100-fold range has revealed that though first order kinetics
in 1 is observed at low catalyst concentrations (below 10-6 M), there is a significant negative
deviation from linearity at [1] > 10-6 M. The latter was rationalized by the equilibrium between
the monomeric and dimeric FeIV species 2 M ⇌ M-M (Kd), both being able to oxidize
p-NO2C6H4SMe with rate constants km and kd which were found to be (13 1)×104 and
(0.32 0.01)×104 M-1 s-1, respectively. The difference in the rate constants is the key for resolving
1

the dilemma of faster catalysis versus slower single turnover reactivity of TAML activators in
water.
Keywords: Iron complex; TAML; Kinetics; Mechanism; Thioanisoles; Stopped-flow
1. Introduction
Iron complexes in higher oxidation states have attracted much attention from academic and
industrial researchers due to their enormous oxidizing potential in numerous instances [1-7]. Our
interests focus on the iron TAML activators of peroxides (TAML = tetraamido macrocyclic
ligand or their iron complexes) [8, 9] which have a broad spectrum of environmentally oriented
applications [10] and bio-inspired and -competitive mechanistic chemistry [5, 11]. In 2008 we
reported that TAML complex 1 (chart 1) is transformed by tBuOOH in water into iron(IV)
derivatives 2 and/or 3 depending on reaction conditions[12, 13]. Later, the kinetics and
mechanism of this transformation were explored in detail in collaboration with Professor Rudi
van Eldik, to whom this issue is devoted [14]. Our investigations of the reactivity of FeV and FeIV
complexes derived from TAMLs, which were performed at –40 °C in MeCN, revealed that both
forms oxidize thioanisoles, though FeV is by several orders of magnitude more reactive than FeIV
[15]. More recently we explored the reactivity of iron(IV) species in aqueous solution towards
hexacyanoferrate(II) as an electron donor [16]. Here we report the results of kinetic and
mechanistic studies of the oxidation of ring-substituted thioanisoles into the corresponding
sulfoxides under catalytic and single-turnover conditions using conventional and stopped-flow
spectroscopy, respectively. The data revealed that FeIV TAML species behave as electrophiles
toward p-XC6H4SMe and allowed us to rationalize a pseudo-controversial discrepancy in
reactivity measured under the catalytic conditions (below micromolar concentrations) and the
single-turnover conditions, which require significantly higher concentrations of TAML
activators. This effect is ascribed to the reversible formation of less reactive dimeric iron(IV)
species, which start to play a noticeable role at TAML concentrations above 10-6 M.
2

Chart 1. TAML activators used in this study.
2. Experimental
2.1. Materials and methods
Thioanisole and p-substituted thioanisoles purchased from Sigma and Aldrich were used as
received. Orange II (Aldrich) was purified by passing through a column filled with SMT Bulk-
C18 (Separation Methods Technologies Inc., Newark) with water-methanol mixture (9:1) as an
eluent. H2O2 (30% w/w) and tBuOOH (70% w/w in water) were purchased from Fluka. Other
reagents, components of buffer solutions, and solvents (at least ACS reagent grade) were
obtained from Aldrich, Fisher, or Acros and used as received or purified in an appropriate
manner [17]. TAML 1 was synthesized as described elsewhere [18]. Uv-vis spectrophotometric
measurements were performed using a Hewlett Packard Diode Array spectrophotometer (model
8453) equipped with a thermostated cell holder and an automatic 8-cell positioner. Temperature
was controlled by Thermo digital temperature controller RTE17 with an accuracy of 1 °C. Both
quartz and appropriate plastic 1.0 cm cuvettes were used. Stopped-flow kinetic measurements
were carried out using a two-syringe BioKine32 SFM-20 model (BioLogic Science Instruments)
at 25 °C. Technical details are summarized elsewhere [14]. Products of the oxidation of
thioanisoles were analyzed by high performance liquid chromatography (HPLC) using a Varian
C18 Omnispher column with acetonitrile and water as eluting solvents, as well as by gas
chromatography – mass spectrometer (GC-MS) using a ThermoDSQ instrument. Kinetic
measurements were performed at pH 8 in 0.01 M phosphate buffer prepared using HPLC grade
water. Stock solutions of thioanisoles and other substrates used in this study were prepared in
acetonitrile (HPLC grade) or buffer/acetonitrile mixture. Stock solution of Orange II was
prepared in water (HPLC grade). Stock solutions of H2O2 were prepared from 30% H2O2 and
3

standardized by measuring the absorbance at 230 nm (ε = 72.8 M-1 cm-1) [19]. Fresh stock
solutions of tBuOOH were prepared from 70% tBuOOH in water (ca. 7.2 M) without calibration.
Stock solutions of all TAML catalysts used in this study were made in water (HPLC grade) or
acetonitrile (HPLC grade).
2.2. Reactions of 2
Solutions of 2 were prepared at 25 °C by adding 0.5 equivalents of H2O2 or tBuOOH to solutions
of 1 in 0.01 M phosphate buffer at pH 8.0. Reactions with thioanisoles were followed by
monitoring uv-vis spectra of 2 at 420 and 780 nm after addition of variable concentrations of the
substrates. Products of the reaction between 2 and thioanisole were analyzed by HPLC and
identified by comparing retention times to that of the authentic compound.
2.3. Kinetic studies
The kinetics of oxidation of 4-nitrothioanisole was monitored by following its absorbance
decrease at 350 nm, while that for Orange II was monitored at 485 nm (pH 8). The temperature
was always maintained at 25 °C. A typical kinetic run was performed as follows. The phosphate
buffer (2 mL) was first added to a thermostated cuvette to which appropriate amounts of the
stock solutions of a substrate and 1 were added. The reaction was initiated by the addition of an
aliquot of the stock solution of tBuOOH or H2O2. The rates were calculated from the measured
changes in absorbance with time using the estimated extinction coefficients for the substrates.
3. Results and discussion
3.1. Introductory remarks
In the solid state, iron(III) TAML compounds are five-coordinate square pyramidal species with
an axial aqua, methanolic or acido ligand [5, 11]. In water, they are six-coordinate diaqua species
1 with axial water ligands as shown in chart 1. A first aqua ligand of 1 deprotonates with pKa of
10.1 [20]. A half equivalent of H2O2 (faster) [14] or tBuOOH (slower) [12] converts millimolar
solutions of 1 in water into an iron(IV) TAML species, which is either the μ-oxo dimer 2 or
oxoiron monomer 3 as confirmed by the results of Mössbauer spectroscopy [12]. Species 2,
which is structurally similar to the complex generated from 1 by dioxygen in organic solvents
4

[21], dominates at pH below 10, while species 3 dominates at pH above 12. Variation of pH
converts reversibly 2 into 3 and vice versa suggesting the following equilibrium [12]:
[LFe−O−FeL]2- (2) + 2 OH-
2[LFeIV=O]2- (3) + H2O
(1)
Major spectral differences between 2 and 3 manifest in the visible region at wavelengths above
700 nm where 3 does not absorb light. In contrast, 2 has two broad metal-to-ligand charge
transfer bands around 800 and 1100 nm (figure 1A). Equilibrium 1 controlled by hydroxide ions
suggests that a similar process such as shown by eq. 2 may occur at lower pH in dilute solutions
of 2.
[LFe−O−FeL]2- (2) + H2O
2[LFeIV−OH]- (4)
(2)
Dilution is a driving force for reaction 2. Species 4 as shown is analogous to the postulated
compound II of haloperoxidases [22, 23], although the recently suggested actual speciation of
iron(IV) TAMLs in aqueous solution is significantly more complex [16].
Control experiments with 2 generated from a sodium salt of 1 in the presence of 0.5
equivalents of H2O2 or tBuOOH have shown that 2 is rapidly reduced by electron donors to
afford 1. When an orange solution of 1 (1×10-4 M) at pH 8 (0.01 M phosphate buffer) is treated
with tBuOOH, the solution turns dark and new bands appear at 800 and 1100 nm with a shoulder
at 540 nm (figure 1A). It takes ca. 10 min to achieve this conversion at 25 °C. TAML-catalyzed
degradation of Orange II by peroxides was studied in detail previously [24] and therefore it was
used in this work. Orange II has a strong band around 480 nm and does not absorb light at 800
nm. Therefore both the dye degradation and the reduction of 2 into 1 could be confirmed in one
experiment (figure 1B). Both processes occur at similar rates. Similar bleaching of Orange II by
2 was observed when H2O2 (0.5 eq) was used instead of tBuOOH.
3.2. Stopped-flow kinetic studies of reactions of 2 with p-XC6H4SMe
Thianisole reacts with 2 instantaneously at room temperature. Dimer 2 is reduced back into 1
producing methyl phenyl sulfoxide as confirmed by HPLC. Kinetics of the reduction of 2 into 1
by p-XC6H4SMe was studied using stopped-flow spectroscopy in excess of sulfides to ensure
pseudo-first-order conditions. A decrease in the absorbance of 2 (ca. 10-5 M) at 420 nm was
5

registered. Exponential kinetic traces such as shown in figure 2 were generated in all cases
studied. Satisfactory first-order kinetics was observed in a matter of 3-4 half-lives. The pseudo-
first-order rate constants kobs were measured at different concentrations of thioanisoles and
-1
revealed hyperbolic dependencies such as in figure 3 for all p-XC6H4SMe studied. Plots kobs
versus [XC6H4SMe]-1 were straight lines with non-zero intercepts consistent with rate law 3,
which agrees with a mechanism involving the formation of an intermediate produced from 2 and
thioanisole followed by the rate-limiting step. Interestingly, the same rate law was established for
the reaction between the same partners in acetonitrile at -40 °C [15].
ꢁꢂ[ArSMe]
ꢀobs
=
(3)
ꢃꢄꢂ[ArSMe]
The a and b parameters of eq. 3 were calculated for all thioanisoles used in this study by fitting
the data such as in figure 3 to eq. 3 and the corresponding values are summarized in table 1. Rate
law 3 and the dimeric structure of 2 combined suggest several mechanistic possibilities, some of
which are included in table 2. Mechanism M1 assumes that the intact dimer 2 coordinates
reversibly an organic sulfide followed by rate-limiting electron transfer to afford a radical cation
MeArS+· which is then rapidly transformed into a sulfoxide through a sequence of events
proposed elsewhere [25].
Mechanism M2 (table 2) involves a reversible conversion of dimer 2 into a monomeric
reactive FeIV-sulfide adduct as a result of a bridge cleavage by the nucleophilic organic sulfide
followed by the electron transfer step. Mechanism M3 assumes that sulfide attack at dimer 2
causes its disproportionation that leads to the iron(V)oxo species coordinated by sulfide and an
iron(III) derivative. The iron(V)oxo species may be involved in a concerted oxygen atom transfer
affording sulfoxide and iron(III). Mechanism M4 is based on the assumption that dimer 2 is
noticeably less reactive than a monomeric form such as 4 in eq. 2. The 4 monomer forms a
reactive intermediate with sulfide followed by an electron transfer step. There is one essential
feature that distinguishes M4 from mechanisms M1−M3. If the equilibrium driven by K4 is
shifted strongly to the left (dimer 2 is a dominating species in solution), the reaction should have
a kinetic order of one-half in the oxidant. Kinetic results obtained under the steady-state
conditions, which are described below, will highlight the importance of additional mechanism
M5, which is actually a combination of M1 and M4.
6

The corresponding kinetic equations for all mechanisms in table 2 show that the kinetic
parameter a is equal or proportional to the rate constant k2, which corresponds to the electron (or
atom in M3) transfer from sulfide at FeIV (or FeV in M3). The parameter b, which is equal or
proportional to K1, characterizes an ability of sulfides to coordinate iron. The product ab, which
is proportional to k2K1, links the reaction conditions at low concentrations of p-XC6H4SMe and
therefore corresponds to effective second order rate constants for the interaction between
iron(IV) species and thioanisoles. Accordingly, figure 4 compares the substituent effects of X in
p-XC6H4SMe on a, b and ab in terms of the Hammett formalism. As shown, all kinetic
parameters decrease from electron-rich to electron-poorer sulfides, i.e. electron-withdrawing
substitutions disfavor the reaction via all possible channels. Both electron-transfer rates and
binding of sulfides to iron(IV) decrease as the basicity of sulfides drops. The best Hammett
correlation is observed for log(a/aH), which is in fact ~log(k2/k2H). The ρ slope of –0.18 is
however very low pointing to an insignificant sensitivity of the intramolecular electron-transfer
step to electronic effects in sulfides when the reaction is run in water. The ρ of –0.18 indicates
considering an additional mechanistic option (M6) similar to M3. The rate-limiting step of M6 is
the disproportion during the oxo-bridge cleavage of 2 to afford iron(III) and iron(V) species, the
latter rapidly oxidizing sulfide (table 2).
Parameter b, which is associated with binding of sulfides to iron(IV) in multiple ways, is
noticeably more sensitive to the electronic effects (ρ = –0.79). As anticipated, the binding
weakens for electron-poor sulfides. The ab product, which reflects contributions from both the
binding and the rate-limiting electron transfer step, has the most negative ρ of ca. –1. Its absolute
value is however still appreciably lower than that of –2.1 observed for the oxidation of the same
sulfides to sulfoxides by the TAML iron(V)oxo species in wet MeCN as solvent [15]. The
analytical expression 4 for the line was used to estimate the ab value of (1.2 0.8)×104 M-1 s-1 for
p-NO2C6H4SMe using the Hammett σ constant for NO2 of 0.78 [26].
log(ab/abH) = −(0.16 0.07) − (0.98 0.21)×σ
(4)
3.3. Kinetics of 1-catalyzed oxidation of 4-nitrothioanisole by hydrogen peroxide under the
steady-state conditions
The previous section summarizes the results of single turnover kinetic studies of oxidation of
organic sulfides by iron(IV) species 2 in aqueous solution. Complex 2 does not absorb visible
7

light strongly and therefore such studies in excess of electron donors were possible for colorless
thioanisoles p-XC6H4SMe with X = MeO, Me, H, Cl and CN. The nitro derivative
p-NO2C6H4SMe could not be investigated similarly because it has a band at 350 nm. In contrast,
kinetic studies of 1-catalyzed oxidation of p-NO2C6H4SMe by H2O2 are feasible as opposed to
the other colorless thioanisoles. The reaction between 1 and H2O2 affords 2 in a more than 80%
yield at pH below 10 [12] and therefore we decided to compare the rate constants obtained under
the steady-state and single turnover conditions, the latter being available for p-NO2C6H4SMe
from extrapolation using eq. 4. Under the catalytic conditions, p-NO2C6H4SMe is oxidized by
H2O2 into the corresponding sulfoxide confirmed by HPLC and gas chromatography−mass
spectrometry (GC-MS) after the reaction was completed. The oxidation of p-NO2C6H4SMe into
p-NO2C6H4SOMe is a very clean process. As seen in figure 5, the sulfide is gradually converted
into the sulfoxide, which has a maximal absorbance at 280 nm. Two isosbestic points hold at 239
and 305 nm suggesting no accumulation of other absorbing intermediates. The kinetic data
collected at 350 and 280 nm match perfectly. No further oxidation of the sulfoxide into sulfone
was observed under the reaction conditions used (pH 8, 25 °C).
We have performed several kinetic studies of TAML-catalyzed oxidations of different
targets (D) by peroxides [16, 24, 27-31]. The oxidations were first-order in the catalyst at
concentrations below 10-6 M. Mixed orders in H2O2 and D were established as in eq. 5.
[ ]
ꢅ D
[
][ ]
ꢀIꢀII HꢇOꢇ D
III
[
Fe
]
( )
5
−
=
[
]
[ ]
ꢅꢆ
ꢀ-I + ꢀI HꢇOꢇ + ꢀII D
Equation 5 is consistent with a minimalistic reaction mechanism presented in scheme 1. If the
oxidation of D occurs ‘in the kII kinetic regime’ when kI[H2O2] >> k-I + kII[D] at high
concentrations of H2O2, eq. 4 simplifies to –d[D]/dt = kII[D][FeIII] and the second order rate
constant kII should be similar to the product ab (table 1). The fact that the oxidation of iron(III)
TAML by H2O2 may afford an iron(V) intermediate [14, 27] is kinetically unimportant because
the reactivity of FeV to sulfides is by four orders of magnitude higher than that of FeIV [15] and
under the steady-state conditions the rate-limiting step should be the oxidation by FeIV. The
reader may find more detailed justification elsewhere [16].
kI
k-I
8

FeIII-TAML + ROOH
(Oxidized TAML)
k
III
II
(Oxidized TAML) + Substrate ⎯⎯→ Fe -TAML + Product
fast
Primary Product ⎯⎯⎯→ → → Final Product(s)
Scheme 1. General mechanism of catalysis by TAML activators in aqueous media.
However, the rate of oxidation deviated from linearity at higher concentrations of TAML
catalysts in certain cases and this issue is addressed here because the iron(IV) dimeric species
dominate at millimolar concentrations in aqueous solution at pH below 10. The catalysis is
however performed using micromolar and lower concentrations of TAMLs and therefore dimer 2
could undergo monomerization to form 4 as in eq. 2.
It has been confirmed that the oxidation of p-NO2C6H4SMe follows ‘normal’ kinetics
under the conditions typically used for catalysis by TAMLs. Figure 6 shows that the reaction rate
is directly proportional to the concentration of 1 below 10-6 M. Figure 7 illustrates a hyperbolic
dependence of the rate on [H2O2]. The inset to figure 7 shows that the inverse rate is proportional
to [H2O2]-1 with a positive intercept as predicted by eq. 4. Assuming that k-I is negligible [11], the
rate constants kI and kII of 167 20 and (1.6 0.2)×105 M-1 s-1, respectively, have been calculated.
The kI value agrees with that of 260 M-1 s-1 obtained previously using a cyclometalated
ruthenium(II) dye as an electron donor under the same conditions [27]. The kII value appeared by
one order of magnitude higher than ab estimated from the Hammett plot.
To resolve the rate controversy, we have measured initial rates of oxidation of
p-NO2C6H4SMe by H2O2 in the broadest range of concentrations of 1 at a high concentration of
H2O2 to ensure that the limiting rate expression –d[D]/dt = kII[D][FeIII] holds. The data presented
in figure 8, which covers a 100-fold concentration range, indicate that a non-linear dependence is
observed. The reaction order in 1 changes from one (figure 6) to one half (inset to figure 8) at
lower and higher concentrations of the catalyst, respectively. This observation is consistent with
the occurrence of the postulated equation 2. The fraction of monomeric species such as 4
increases on dilution and, if the reactivity of the monomeric species is appreciably higher than
that of the dimeric counterparts, the dependence shown in figure 8 can be rationalized
quantitatively. The corresponding mechanism assumes the coexistence of both active monomeric
9

and dimeric species (scheme 2) and results in eq. 6 for the dependence of the reaction rate on the
total catalyst concentration Mt = [M] + 2[M2].
Kd
2 M
M
M
kd
km
+ S
+ S
Products
Scheme 2. A reaction scheme involving monomeric M and dimeric M−M oxidized intermediates
in catalysis by TAMLs (see text for details).
ꢅ[D]
ꢅꢆ
[D]
(
)
−
=
ꢉ 2ꢀ − ꢀ × ꢊ 8ꢈ M + 1 − 1ꢌ + 4ꢀ ꢈ M ꢍ
(6)
ꢋ
m
d
d
t
d
d
t
8ꢈd
Equation 6 includes terms proportional both to M and M , the latter being negligible at lower
ꢋ
t
t
catalyst concentrations. Equation 6 is in qualitative agreement with the data presented in figure 8
and therefore it was used for fitting the data. The value of km ((13 1)×104 M-1 s-1) was
determined directly from the data collected at the lowest concentrations of the catalyst (figure 6)
and therefore kd and Kd, i.e. the rate constant for the dimer and the dimerization constant,
respectively, were found in the fitting routine. The values of kd and Kd equal
(0.32 0.01)×104 M-1 s-1 and (4.5 0.5)×104 M-1, respectively. The value of kd for p-NO2C6H4SMe
obtained in the fit should be compared to that of (1.2 0.8)×104 M-1 s-1 estimated from the
Hammett plot in figure 4. Though the latter is slightly higher than kd obtained from the data in
figure 8, they are rather close taking into account experimental uncertainties. The observed
difference may be due to the fact that km and kd differ by almost two orders of magnitude and the
monomerization via eq. 2 could not be taken into consideration in the single turnover
experiments.
The value of equilibrium constant Kd has been used to calculate the concentration profile
for a relative amount of the monomeric form in solution at pH 8 (figure 9). The relative amount
of the monomeric form is higher than 90% at concentrations below 1.4×10-6 M and is 50% at
10

2.2×10-5 M. Thus, dilute solutions of TAMLs should preferably be used in kinetic studies to
avoid complicating kinetic analysis association phenomena.
4. Conclusion
Reactive iron(IV) species produced from TAML activators in the presence of peroxides may
undergo spontaneous dimerization when the catalyst concentration exceeds 10-6 M in aqueous
solution at pH 8. The dimerization is negligible in water when more dilute solutions of TAML
activators are used. The dimerization may cause a switch of the reaction order in the catalysts
from 1 to ½, though mixed reaction order may also be observed. The results presented here
support our previous assumption that reactive monomeric iron(IV) TAMLs are not affected by
aggregation to any significant extent when used at concentrations below 10-6 M in water [16].
This study would not be possible without using the stopped-flow spectroscopy, the technique
which one of us (ADR) learned from Rudi (van Eldik) when we all were much younger [32].
Acknowledgements
Operating support is acknowledged from the Heinz Endowments (T.J.C.), the Institute for Green
Science (T.J.C.), and the Environmental Protection Agency (grant RD 83 to T.J.C.). ADR thanks
Genoa R. Warner for critical reading of the manuscript and helpful comments.
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13

Table 1. Parameters a, b and ab of eq. 3 experimentally obtained for TAML
compound 2 at 25 °C, pH 8 in 0.01 M phosphate. Compound 2 (~1×10-5 M)
was obtained from 1 and H2O2 prior to addition of methyl aryl sulfides.
X in p-XC6H4SMe
10-3×b / M-1
a / s-1
10-4×ab / M-1 s-1
OCH3
CH3
H
12
25
16
24
14
4
7
3
4
6
4.8 0.3
4.6 0.2
4.3 0.2
4.1 0.5
3.2 0.5
7.4
10.3
5.7
Cl
CN
NO2
1.3
1.2a)
1
a) Value extrapolated from the Hammett plot in figure 4.
14

Table 2. Mechanistic options for oxidation of methyl aryl sulfides by 2; all but M4 are consistent
with rate law 3.
Me
Ar
Me
Ar
Me
Ar
k2
M1
K1
O
fast
O
O
MIV
MIV
S
Products
MIV
MIV
MIV
MIII
S
S
ꢎ ꢏ [S] ; a = k2, b = K1
2
1
ꢀobs
=
ꢃꢄ ꢏ [S]
1
Me
Ar
Me
Ar
Me
k2
M2
K1
O
fast
MIV
MIV
S
Products
MIV
S
MIV
S
MIV
MIII
O
Ar
ꢔꢕ
IV
ꢐ.ꢑꢎ ꢏ ꢒM
ꢓ
ꢀobs
=
[S]; a = 0.5k2, b = K1[MIV]-1
2
1
IV ꢔꢕ
[
]
ꢃꢄ ꢏ
M
[S]
1
Me
Ar
Me
Ar
Me
S
M3
M4
k2
K1
O
MIV
MIV
S
MIII
S
2MIII
O
MV
O
Ar
ꢔꢕ
III
ꢐ.ꢑꢎ ꢏ ꢒM
ꢓ
[S]; a = 0.5k2, b = K1[MIII]-1
2
1
ꢀobs
=
III ꢔꢕ
[
]
ꢃꢄ ꢏ
M
[S]
1
K4
+H2O
Me
Ar
Me
Ar
Me
k2
K1
O
fast
1
MIV
MIV
OH
MIV
S
2
S
Products
HO
MIV
MIII
S
-H2O
Ar
ꢎ [S] ꢏ M
ꢎ ꢏ [S] ꢏ M
ꢖ[2]
t ; − ꢖ[2]
=
ꢋ
ꢋ
ꢘ
2
ꢘ
2
1
t
−
=
ꢙ
ꢚꢙ
2
[ ]
ꢖꢗ
ꢖꢗ
ꢃꢄꢏ
S
-1
1
[ ]
ꢄ S
ꢙ
1
M5
M6
M1 + M4
Me
Me
Me
Ar
k2
K1
O
fast
O
MIV
MIV
S
Products
MIV
MIV
O
MV
S
MIII
S
Ar
Ar
15

A
B
Figure 1. A. Formation of 2 from 1 and 0.5 equiv tBuOOH. Inset to A shows kinetics of the
formation of 2. B. Dashed line is the spectrum of 1 before adding tBuOOH. The ssolution of
Orange II (5×10-5 M) was added after the formation of 2 was complete and the soolid spectra
visualize the dye bleaching. Inset B shows a rapid decrease in absorbance at 800 nnm due to
reduction of 2 restoring the spectrum of 1 which is affected by the products of deggradation of
Orange II. Conditions: [1] 1×10-4 M, [tBuOOH] 5×10-5 M, pH 8 (0.01 M phosphate), 25 °C.
16

Figure 2. Kinetic trace corresponding to reduction of
monitored at 420 nm by stopped-flow spectroscopy.
2
(1
×
10-5 M) by thioanisole (5×10-5 M) as
and H2O2 prior to the
2
was made from
1
addition of sulfide. Conditions: pH 8 (0.01 M phosphate), 25 °C.
17

Figure 3. Pseudo-first order rate constants obs for the reduction of
k
2 into 1 against different
concentrations of -XC6H4SMe ( , X = H; ●, X = CN). Solid and dashed lines are calculated
p
○
curves using the best-fit parameters of eq. 3 summarized in table 1. Conditions: pH 8 (0.01 M
phosphate), 25 °C.
18

Figure 4. Hammet plots of parameters of eq. 3 in a log (p/pH) form (p = a, b, or ab) against σp of
p-substituted thianisoles for their oxidation by 2 at pH 8 and 25 °C. Data are from table 1.
19

Figure 5. Spectral changes during
1
-catalyzed oxidation of
-NO2C6H4SMe (7.1×10-5 M) by
p
H2O2 (0.04 M). Other conditions: [
1
] = 2.3×10-7, pH 8, 25 °C.
20

Figure 6. Linear dependence of initial rates of oxidation of
p
-NO2C6H4SMe on concentration of
1
below 10-6 M at [sulfide] 2.4×10-5 M, [H2O2] 0.4 M, pH 8 (0.01 M phosphate) and 25 °C.
21

Figure 7. Initial rate of oxidation of 4-nitrothioanisole by H2O2 catalyzed by 1 as a function of
peroxide concentrations. Inset is the corresponding double reciprocal plot. Conditions: [1]
2.3×10-7 M, [4-nitrothioanisole] 7.1×10-5 M, pH 8 (0.01 M phosphate), 25 °C.
22

A
B
Figure 8. A. Deviation from linearity of the initial rate for oxidation of
(2.4×10-5 M) by H2O2 (0.4 M) as a function of concentrations of 1 covering a 100-fold range.
p
-NO2C6H4SMe
The solid line was calculated using the best-fit values of d and d (see text for details). B. The
k
K
rate as a linear function of a square root of [1] at [1] > 1×10-6 M. Conditions: pH 8 (0.01 M
phosphate), 25 °C.
23

Figure 9. Fraction of the monomeric form M of iron(IV) as a function of total iron concentration
calculated using Kd = 4.5×105 M-1. Inset is the same dependence at low Mt.
24

Graphical Abstract
25