Reduced Amino Acid Schiff Base-Iron(III) Complexes Catalyzing Oxidation of Cyclohexane with Hydrogen Peroxide

Article abstract of DOI:10.1002/ejic.202100356

The reduced amino acid Schiff base ligands have been prepared and were coordinated with ferric chloride to generate the iron(III) complexes. The ligands and complexes have been characterized using FT-IR, UV-vis, elemental analysis, ICP-AES analysis, mass spectra etc. After the structural characterization, these complexes were applied for the oxidation of cyclohexane using hydrogen peroxide as the oxidant under mild conditions. The activity tests showed that the L-phenylalanine-derived reduced Schiff base iron(III) complex(Ph?FeCl) afforded the highest yield of cyclohexanol and cyclohexanone(total yield up to 23.2 %). Notably, the Ph?FeCl complex catalyzes the reaction via a heterogeneous approach, allowing the complex to be separated and recycled conveniently after the oxidation reaction. Besides, the Ph?FeCl catalyst can also be extended for the selective oxidation of other alkanes and aromatics into alcohols, ketones and phenols etc. Finally, the reaction mechanism of cyclohexane oxidation on the iron(III) complex was proposed as well by the free radical inhibitors and EPR study of active intermediates.

Full text of DOI:10.1002/ejic.202100356

European Journal of Inorganic Chemistry
Accepted Article
Title: Reduced Amino Acid Schiff Base-Iron(III) Complexes Catalyzing
Oxidation of Cyclohexane with Hydrogen Peroxide
Authors: Anna Zheng, Qingqing Zhou, Bingjie Ding, Difan Li, Tong
Zhang, and Zhenshan Hou
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To be cited as: Eur. J. Inorg. Chem. 10.1002/ejic.202100356
Link to VoR: https://doi.org/10.1002/ejic.202100356
01/2020

10.1002/ejic.202100356
European Journal of Inorganic Chemistry
Reduced Amino Acid Schiff Base-Iron(III) Complexes Catalyzing Oxidation of
Cyclohexane with Hydrogen Peroxide
Anna Zheng,[a] Dr. Qingqing Zhou,[a] Bingjie Ding,[a] Dr. Difan Li,[a]
Tong Zhang,[a] Prof. Dr. Zhenshan Hou*[a],[b]
aKey Laboratory for Advanced Materials, Research Institute of Industrial Catalysis, School of
Chemistry and Molecular Engineering, East China University of Science and Technology,
Shanghai 200237, China.
bShanghai Key Laboratory of Green Chemistry and Chemical Processes, East China Normal
University, School of Chemistry and Molecular Engineering, Shanghai 200062, China.
*Corresponding author (Prof. Dr. Zhenshan Hou). houzhenshan@ecust.edu.cn;
https://chem.ecust.edu.cn/2014/1113/c6655a50161/page.htm
Abstract The reduced amino acid Schiff base ligands have been prepared and were
coordinated with ferric chloride to generate the iron(III) complexes. The ligands and
complexes have been characterized using FT-IR, UV-vis, elemental analysis, ICP-AES
analysis, mass spectra etc. After the structural characterization, these complexes were
applied for the oxidation of cyclohexane using hydrogen peroxide as the oxidant under
mild conditions. The activity tests showed that the L-phenylalanine-derived reduced
Schiff base iron(III) complex(Ph-FeCl) afforded the highest yield of cyclohexanol and
cyclohexanone(total yield up to 23.2%). Notably, the Ph-FeCl complex catalyzes the
reaction via a heterogeneous approach, allowing the complex to be separated and
recycled conveniently after the oxidation reaction. Besides, the Ph-FeCl catalyst can
also be extended for the selective oxidation of other alkanes and aromatics into alcohols,
ketones and phenols etc. Finally, the reaction mechanism of cyclohexane oxidation on
the iron(III) complex was proposed as well by the free radical inhibitors and EPR study
of active intermediates.
1
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European Journal of Inorganic Chemistry
Introduction
Petroleum and natural gas derived hydrocarbons play a huge role in modern
civilization. Petroleum includes a significant amount of saturated hydrocarbons
(alkanes), aromatic and other unsaturated hydrocarbons.[1] Nowadays, the chemical
process for the conversion of the feed hydrocarbons is normally carried out in most
cases with the participation of heterogeneous catalysts, and requires the elevated
temperatures (>300 ºC).[2] In particular, a process of partial oxidation plays a crucial
role in the production of many feedstocks and chemical intermediates from
hydrocarbons.[3]
The catalytic oxidation of C-H bonds of hydrocarbons to the oxygen-containing
products under mild conditions is an appealing challenge. It has been found that the
earth abundant transition metal complexes, especially Fe, Cu, Co and Mn salts
coordinated by a large variety of N-, O- and P-containing ligands such as Salen, heme-
and non-heme N4-, and phosphine ligands are highly stable, catalytically active, and
have attracted much attention in catalytic oxidation.[4] Cyclohexane has been used as a
typical model compound for evaluating catalyst activity for oxidation of hydrocarbons
owing to the following reasons. First, it has a sufficiently high C-H bond dissociation
energy of 99.3 kcal/mol, therefore its oxidation via non-catalytic processes is
challenging;[5] Second, as a cyclic alkane with all equal positions, the products from
cyclohexane oxidation are non-isomers; Third, cyclohexane oxidation can produce
cyclohexanol(A) and cyclohexanone(K), both of which are called KA oil and they are
important intermediates for the synthesis of Nylon-6 and Nylon-6,6 in industry.[6]
The oxidation of cyclohexane is carried out using cobalt based homogeneous
catalysts such as cobalt isooctanoate and cobalt porphyrins with high temperatures
(between 170 and 230 ºC) and under 10-15 bar of pressure. The conversion of
cyclohexane needs to be around 5-12% in order to inhibit the over-oxidation of KA oil
in the chemical process.[6c,7] In general, KA oil production involves two consecutive
steps: the non-catalytic auto-oxidation of cyclohexane to cyclohexyl hydroperoxide
(CHHP) and the decomposition of CHHP to KA oil. However, a small amount of
2
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10.1002/ejic.202100356
European Journal of Inorganic Chemistry
carboxylic acids and esters could be produced as by-products under the harsh
conditions.[8] Additionally, the high-energy cost caused by the high reaction temperature
has become a big disadvantage in the industrial production process.
The catalytic oxidation of cyclohexane under mild conditions (temperature < the
boiling point of cyclohexane) would relieve the hazard and also allow for reaction
pathway and selectivity control.[9] Thus, many efforts have been devoted to developing
highly efficient catalysts that are active under low reaction temperature in order to
impede over-oxidation reaction. The coordination compounds indeed show some
catalytic activity for alkane oxidation, which has attracted much attention in the recent
years.[10] Especially, Schiff bases derived transition metal complexes are highly
attractive catalysts for the oxidation of alkanes considering that they are cheap, easy
available and highly stable.[10d,11]
Among transition metal, iron is inexpensive, abundant, and non-toxic.[12]
Specialized enzymes in nature such as porphyrins, a model of heme enzymes, or non-
heme enzymes including methane monooxygenase, rieske dioxygenase, lipoxygenase
and catechol dioxygenase usually contain iron in their active sites.[13] These enzymes
are widely applied in the oxidation of biochemical fields. Taking inspiration from nature,
biomimetic model iron complexes have been designed and used in C-H compounds’
functionalization.[14] More recently, simple Schiff base iron complexes have attracted
considerable attention as bioinspired oxidation catalysts mimicking non-heme iron
oxygenases by using H2O2 as the terminal oxidant.[15]
Notably, Schiff base ligands, derived from different aldehydes and various natural
amino acids which have the advantages of wide sources, low prices and easy
availability,[16] could be composited to several corresponding transition metal
complexes with interesting structural features. Additionally, owing to the good
biocompatibility of amino acid Schiff bases and their corresponding iron complexes,
the study of their design and application has become a hot topic.[17] Some Schiff base
iron complexes have been reported in catalytic cyclohexane oxidation over the past few
years. For instance, [FeIII(HL)Cl2(DMF)]Cl·DMF[6d] afforded 37% yield of KA oil
3
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10.1002/ejic.202100356
European Journal of Inorganic Chemistry
under mild conditions by using H2O2 and HNO3. Dithiocarbazate Schiff base iron
complexes (FeSBdiAP)[6b] and FeSMdiAP[6c] can catalyze cyclohexane oxidation and
the yield of KA oil could reach up to ca.30% and 16.5%, respectively.
Additionally, the reduced Schiff bases are more stable, easily available and highly
bio-compatible, and are expected to generate much more interesting coordination
chemistry due to their conformationally flexible backbone. Inspired by these
particularities and in pursuit of our interest in the development of novel iron catalysts
for the oxidation of cyclohexane using H2O2 under mild conditions, in this work, we
attempted to synthesize the amino acid-derived reduced Schiff base iron(III) complexes
Ala-FeCl, Ph-FeCl and Tyr-FeCl(whereAla, Ph and Tyr = N, N, O-donor reduced Schiff
base ligands) and determined their structure by various methods. Furthermore, these
iron(III) complexes were employed for cyclohexane oxidation with H2O2 as an oxidant.
It was found that the L-phenylalanine-derived reduced Schiff base iron(III) complex
(Ph-FeCl) showed the high catalytic activity under mild conditions without additives,
and the KA oil yield was comparable to the reported iron complexes. Particularly, Ph-
FeCl is insoluble in most of organic solvents, which provides a convenient way for
catalyst separation and recovery. Finally, the reaction mechanism was also proposed on
the basis of EPR characterization.
Results and Discussion
Preparation and characterization of iron(III) complexes
In this work, three Schiff bases were synthesized from the reaction of L-alanine, L-
phenylalanine and L-tyrosine with 2-pyridinecarboxaldehyde, respectively, which were
subsequently reacted with NaBH4 to generate the corresponding reduced Schiff base
ligands Ala, Ph and Tyr, respectively. The complexes: Ala-FeCl, Ph-FeCl and Tyr-FeCl
were obtained by the reaction of the ligands with FeCl3·6H2O (Scheme 1 and Scheme
S1).The reduced Schiff bases are in general more stable and easily separable than Schiff
bases.[18] Besides, the reduced Schiff bases show another advantage that they can
provide a choice of coordination geometry around the metal ions.[19] Afterwards, all
ligands and iron(III) complexes were firstly subjected to elemental analysis and ICP-
4
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10.1002/ejic.202100356
European Journal of Inorganic Chemistry
AES analysis (Table S1). The structural formulas of the iron(III) complexes were then
obtained and shown in Scheme 1.
Scheme 1. Synthesis of the amino acid Schiff base ligands and iron(III) complexes.
As the iron(III) complexes are not dissolved in most organic solvents but are
sparsely soluble in DMSO, the UV-vis spectra of ligands and the iron(III) complexes
were carried out in DMSO by the following steps: UV-vis spectra of a sole DMSO was
recorded as background and then the UV-vis absorption of the ligands and complexes
solution in DMSO were recorded by a subtraction of background. As shown in Figure
S1, the spectra displayed two main absorption bands in the 190-800 nm wavelength
range. The absorption bands at 260-270 nm can be assigned to the π–π* transitions of
the pyridine group and phenyl moieties in the ligands and such absorption bands did
not show any changes when the ligands were complexed with FeCl3·6H2O.
Comparatively, the absorption bands appearing around 321-324 nm can be assigned to
the ligand to metal charge transfer (LMCT) transitions, which indicated that the
coordination interaction occurred between iron(III) sites and the ligands.[20] All
absorption bands and the extinction coefficients of ligands and the corresponding iron
complexes have been listed in Table S1.
Next, the ligands Ala, Ph and Tyr and their iron(III) complexes have been
5
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10.1002/ejic.202100356
European Journal of Inorganic Chemistry
characterized by FT-IR and their characteristic bands were shown in Figure S2 and
Table S2. It can be seen that the absorption peaks of three ligands were very similar.
The absorption bands (3422-3450 cm-1) were attributed to the free N-H bond or O-H in
water as the overlap in the vibration absorption of O-H and N-H,[21] while the absorption
bands of C=N bond on the pyridine ring were around 1575-1591 cm-1. The
antisymmetric stretching vibration absorption of COO- in the terminal carboxyl group
of three ligands appeared at 1630 cm-1, 1656 cm-1 and 1610 cm-1, respectively.[22] After
the ligands were bounded with iron(III) center, the absorption of the N-H bond red-
shifted to 3244-3281 cm-1 arising from the dispersion of the charges.[23] Besides, the
peaks corresponding to the ring stretching frequencies C=N at 1577, 1575 and 1591 cm-
1 of free pyridine ring also blue-shifted to 1610, 1607 and 1612 cm-1, respectively.[24,25]
Meantime, the antisymmetric stretching vibration absorption of COO- tended to give a
blue shift, from 1630 cm-1 to 1649 cm-1, 1656 cm-1 to 1672 cm-1 and 1610 cm-1 to 1644
cm-1 owing to the coordination effect, respectively. Notably, the free COOH group of
each ligand gave an absorption band, appearing at 3059, 3064 and 3056 cm−1,
respectively. However, these bands disappeared once iron(III) salts were complexed
with ligands, confirming clearly that the free COOH group in ligands have participated
in the coordination.[24] In addition, several new absorption peaks appeared, which can
be assigned to the formation of new forming Fe-N(py) bond (1003, 1002 and 1004 cm-
1), Fe-O bond (459, 450 and 454 cm-1) and Fe-N bond (423, 425 and 424 cm-1).[26] All
these results supported strongly that the iron(III) sites were coordinated with carboxyl
group, pyridine and amine group of the reduced Schiff base ligands as shown in Scheme
1.
Then, the molar conductivity (m) values of iron(III) complexes were obtained in
DMSO solution at 25 ºC and were in the range 4.34-50.43 Ω-1·cm2·mol-1. The
conductivity(0) of pure DMSO solvent is 0.51 Ω-1·cm-1, which means that the
conductivity of solvent is negligible. The molar conductivity can be deduced from
Figure S3 and the values were shown in Table 1. Since the value above 40 Ω-1·cm2·mol-
1 was often expected for a 1:1 electrolyte,[27] Ala-FeCl was a non-electrolyte in DMSO
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10.1002/ejic.202100356
European Journal of Inorganic Chemistry
due to the low molar conductivity (4.34 Ω-1·cm2·mol-1). Conversely, Ph-FeCl and Tyr-
FeCl gave the molar conductivity values as high as 43.59 and 50.43 Ω-1·cm2·mol-1,
revealing that both of complexes were 1:1 electrolytes and had two counter ions in
DMSO solution.[10a] The results based on molar conductivity were in well agreement
with those of the molecular formulas (Scheme 1) and the MALDI-TOF mass spectra
(Figure S4).
Table 1. Molar conductivities (m) of iron(III) complexes at 25 oC in DMSO solution.
m (Ω-1 cm2 mol-1)
Complexes
Ala-FeCl
Ph-FeCl
Tyr-FeCl
4.34
43.59
50.43
Catalytic performance
The catalytic activity of a series of iron(III) complex catalysts were investigated
for cyclohexane oxidation using 30 wt% H2O2 as an oxidant at 40 °C. For comparison,
Ala, Ph, Tyr, Bpy-FeCl, Phen-FeCl, Fe2O3 and Fe3O4 have also been employed as
catalysts and their catalytic activity have been shown in Table 2.
Table 2. Cyclohexane oxidation with H2O2 on different catalysts
Yield of
KA oil
(%)a
A/K
molar
ratio
Entries
1
Catalysts
Blank
0
0
2
3
4
5
6
7
Fe2O3
Fe3O4
0
0
0
0
Bpy-FeCl
Phen-FeCl
FeCl3·6H2O
Fe(OTf)3
11.6
20.6
10.1
12.3
2.6
1.5
3.1
1.6
7
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European Journal of Inorganic Chemistry
8
Fe(OAc)2(OH)
Fe(acac)3
Ala
0.8
0
2.4
0
9
10
11
12
13
14
15
0
0
Ph
0
0
Tyr
0
0
Ala-FeCl
Tyr-FeCl
Ph-FeCl
15.9
17.1
23.2
1.2
1.1
1.0
16b
15.4
1.2
FeCl3·6H2O
Reaction conditions: 4 mmol cyclohexane, 8 mmol (30 wt%)H2O2, 0.017 mmol catalyst,
3 mL CH3CN, 40 oC, 4 h. aGC yield with toluene as an internal standard. The yield was
based on cyclohexane. bThe ligand Ph was an additive.
The catalytic reaction could barely occur without catalysts (Table 2, entry 1),
indicating that H2O2 could not oxidize cyclohexane without catalysts. As for Fe2O3 and
Fe3O4, neither of them showed catalytic activity (Table 2, entries 2 and 3). Bpy-FeCl
and Phen-FeCl catalysts gave 11.6% and 20.6% yield of KA oil, respectively (Table 2,
entries 4 and 5). Unfortunately, the reaction system was a homogeneous phase under
this circumstance and thus the catalysts could not be separated from reaction mixture
after reaction. Besides, the simple iron salts, including FeCl3·6H2O, Fe(OTf)3,
Fe(OAc)2(OH) and Fe(acac)3 only afforded low catalytic activity for the reaction (Table
2, entries 6-9). Three ligands (Ala, Ph and Tyr) were inactive (Table 2, entries 10-12).
Conversely, once the iron(III) complexes were formed and used as catalysts, they
offered much higher yield of the KA oil, as compared with iron salt itself (Table 2,
entries 6 vs 13-15). It can be seen that if the ligand Ph was added to FeCl3 in CH3CN
just prior to reaction, a yield of KA oil approximately 15.4% was still achieved (Table
2, entry 16), which was higher than a sole iron salt. Although the activity was slightly
lower than that of Ph-FeCl, the coordination interaction of iron salt with the amino acid-
derived reduced Schiff base ligands indeed exerted a positive effect on the cyclohexane
8
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10.1002/ejic.202100356
European Journal of Inorganic Chemistry
oxidation.
To identify the reason why the three iron(III) complexes provided the different
catalytic performance for cyclohexane oxidation (Table 2, entries 13-15), the rate of
H2O2 decomposition was determined on three complexes. The concentration of residual
H2O2 in the liquid phase were determined by potential difference titration of Ce3+/Ce4+.
As shown in Figure S5, it can be seen that the rate of H2O2 decomposition decreased
according to the following order: Ala-FeCl > Tyr-FeCl > Ph-FeCl. As a result, the
catalytic activity of the three complexes was highly close to the rate of H2O2
decomposition. A decreased rate of H2O2 decomposition would increase the utilization
efficiency of H2O2 and thus was beneficial to the cyclohexane oxidation. On the other
hand, Ala-FeCl showed the lowest catalytic activity among the three complexes likely
due to the strong nucleophilic chlorine ions. Ala-FeCl has four Cl- coordinated with Fe
sites. Nevertheless, the dissociation of Cl- is favorable for generating vacant sites for
H2O2 activation on Fe(III) sites. There have been some reports regarding the number of
iron-bound chloride ions that might be related to the catalytic activity of alkane
oxidation. The iron complexes containing one and two chloride ions in the iron
coordination sphere exhibit the better activity. When the number of iron-bound chloride
ions grows (to three and four), the catalyst activity decreases.[28] Consistently, Ala-FeCl
contains four chloride ions showed the lowest activity in cyclohexane oxidation.
The KAoil yield/time profiles for cyclohexane oxidation over the different iron(III)
complexes with H2O2 as an oxidant were examined and shown in Figure 1. For the three
iron(III) complexes, the yield of KA oil increased with the reaction time, and the KA
oil yield levelled off after 4 h. It was obvious that Ph-FeCl afforded the highest yield of
KA oil among those catalysts (Figure 1a), which was consistent with the rate of the
H2O2 decomposition by the iron(III) complexes (Figure S5). In addition, it was
observed that the A/K molar ratio increased with reaction time within 1 h (Figure 1b),
which might be due to the high initial concentration of H2O2 at the beginning of reaction,
promoting preferentially the oxidation of cyclohexanol(A) to cyclohexanone(K). After
reaction for longer time (> 1 h), the A/K molar ratio remained almost constant on Ala-
9
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10.1002/ejic.202100356
European Journal of Inorganic Chemistry
FeCl but slight decreasing on Ph-FeCl and Tyr-FeCl catalysts (Figure 1b). The rapid
rate of H2O2 decomposition at initial stage of reaction (ca. 1 h) on Ala-FeCl catalyst
can account for the constant A/K molar ratio, when the reaction time was prolonged
further (> 2 h), the residual H2O2 became less and thus was unable to promote further
oxidation of cyclohexanol(A) to cyclohexanone(K). In contrast, for Ph-FeCl and Tyr-
FeCl catalysts, even after 2 h more residual H2O2 still played a role in further oxidizing
cyclohexanol(A) into cyclohexanone(K), and thus the A/K molar ratio slightly
decreased with reaction time (Figure 1b).
1.4
a
b
25
20
15
10
5
1.2
1.0
0.8
0.6
Ph-FeCl
Tyr- FeCl
Ala- FeCl
Ph-FeCl
Tyr- FeCl
Ala- FeCl
0
0
2
4
6
8
10
12
0
2
4
6
8
10
12
Reaction time /h
Reaction time /h
Figure 1. a) The KA oil yield/time profiles and b) the A/K molar ratio/time profiles of
cyclohexane oxidation over the different catalysts. Reaction conditions: 4 mmol
cyclohexane, 8 mmol (30 wt%)H2O2, 0.017 mmol catalyst, 3 mL CH3CN, 40 oC.
Additionally, Ph-FeCl catalyst offered the lowest A/K molar ratio among three
catalysts, implying that Ph-FeCl enabled more cyclohexanol(A) to be oxidized into
cyclohexanone(K). Actually, if cyclohexanol(A) was employed as a substrate,
cyclohexanone(K) was produced with 62.7% yield on Ph-FeCl catalyst. However,
cyclohexanone(K) cannot be oxidized further into other products under the present
condition. As a result, the oxidation of a partial cyclohexanol(A) into cyclohexanone(K)
led to an obvious decrease of A/K molar ratio on Ph-FeCl catalyst as reaction time was
prolonged (Figure 1b).
Afterwards, the effect of reaction temperature on the cyclohexane oxidation was
investigated on Ph-FeCl catalyst. As shown in Figure 2, the yield of KA oil increased
10
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10.1002/ejic.202100356
European Journal of Inorganic Chemistry
with time at different reaction temperatures (Figure 2a). The optimal yield of KA oil
can be achieved at 313 K and higher reaction temperature would result in lower yield
of KA oil due to invalid decomposition of H2O2, although the yield of KA oil increased
sharply at the early stage of reaction at the higher reaction temperature. In addition, the
A/K molar ratio tended to decrease with reaction time at different temperature (Figure
2b), resulting from the further oxidization of cyclohexanol(A) into cyclohexanone(K).
Moreover, the A/K molar ratio values were highly dependent on the reaction
temperatures. At low reaction temperature (293 K and 313 K), the A/K molar ratio
reached nearly 1.0. However, the A/K molar ratio increased up to 1.4 under the higher
reaction temperature (333 K and 353 K).
a
b
25
20
15
10
5
1.5
1.2
0.9
0.6
0.3
293K
313K
333K
353K
293K
313K
333K
353K
0
0
2
4
6
8
10
12
0
2
4
6
8
10
12
Reaction time /h
Reaction time /h
Figure 2. a) The KA oil yield/time profiles and b) the A/K molar ratio/time profiles of
cyclohexane oxidation at different temperatures. Reaction conditions: 4 mmol
cyclohexane, 8 mmol (30 wt%)H2O2, 0.017 mmol Ph-FeCl, 3 mL CH3CN.
Next, the effect of the molar ratio of H2O2/Cyclohexane on the cyclohexane
oxidation was investigated and shown in Figure 3a. It was observed that two equivalents
amount of H2O2 were enough for the oxidation of cyclohexane on the Ph-FeCl catalyst,
and more H2O2 actually had a negative impact on catalytic efficiency. This was
probably due to the accumulation of water along with the consumption of more H2O2,
and those accumulated water would reduce the oxidation rate in the reaction.[29]
Obviously, the oxidation reaction occurred effectively at the appropriate molar ratio of
H2O2/Cyclohexane. As the molar ratio of H2O2/Cyclohexane increased from 1:2 to 4:1,
11
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10.1002/ejic.202100356
European Journal of Inorganic Chemistry
the A/K molar ratio slightly decreased from 1.0 to 0.9, which might be caused by the
oxidation of cyclohexanol(A) under excessive H2O2.[30]
Yield of KA oil
A/K molar ratio
Yield of KA oil
A/K molar ratio
a
b
25
20
15
10
5
25
20
15
10
5
1.5
1.0
0.5
0.0
3
2
1
0
0
0
1:2
1:1
2:1
3:1
3.5:1
4:1
5
H
2
O
2
Cl
2
CN
2
H
The molar ratio of H2O2: Cyclohexane
DMSO COOH
3
CH
CH
3
CH
COOC
3
CH
Figure 3. a) The effect of the molar ratio of H2O2/Cyclohexane. Reaction conditions: 4
mmol cyclohexane, 0.017 mmol Ph-FeCl, 30 wt% H2O2, 40 oC, 4 h, 3 mL CH3CN; b)
The effect of different solvents on cyclohexane oxidation. Reaction conditions: 4 mmol
cyclohexane, 0.017 mmol Ph-FeCl, 8 mmol (30 wt%)H2O2, 40 oC, 4 h, 3 mL solvent.
Subsequently, we investigated the solvent effect on the cyclohexane oxidation
using Ph-FeCl as a catalyst. In a typical experiment, the reaction was carried out in 3
mLof solvent at 40 ºC for 4 h. As shown in Figure 3b, cyclohexane was hardly oxidized
using H2O, DMSO or dichloromethane as a solvent, and the yield of KA oil declined in
solvents by the following order: acetonitrile > acetic acid > ethyl acetate. This behavior
indicated that the reaction media played a vital role to the oxidation of cyclohexane
with H2O2 as an oxidant. The strong polar solvents seemed to favor this reaction.
However, water with strong polarity afforded no activity due to complete immiscibility
of cyclohexane with H2O2. Additionally, although DMSO could dissolve trace amount
of Ph-FeCl catalyst, no catalytic activity was observed as the cyclohexane. The strong
polar DMSO molecules could be competitors for the active sites of metal center on the
catalyst. Besides, dichloromethane and ethyl acetate are weak polar solvents, H2O2 and
organic phase existed in a two-phase state under reaction conditions, which was
unfavorable for the oxidation of cyclohexane. Notably, the oxidation reaction could
proceed smoothly in acetic acid and acetonitrile, and cyclohexane can give higher yield
12
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European Journal of Inorganic Chemistry
of KA oil in acetonitrile. One possible reason was that it was observed visually that Ph-
FeCl catalyst could be much better dispersed in acetonitrile than in acetic acid, leading
to more easy accessibility of substrate molecules with catalytic iron sites. In addition,
acetonitrile acted as a solvent which can donor the electron to iron sites and likely
facilitated to reduce the iron sites.[31] It could be seen that A/K molar ratio was the
lowest in acetonitrile, which revealed that cyclohexanol(A) was more prone to be
oxidized into cyclohexanone(K) in acetonitrile in comparison with other solvents
(Figure 3b).
As a result, it was found that the cyclohexane oxidation can offer an optimal KA
oil yield in CH3CN. Since Ph-FeCl catalyst was well dispersed in the reaction media, a
hot filtration experiment was performed in order to identify whether some iron(III)
active species were leached and the leached part participated in the catalytic process.
The hot filtration procedure was presented as follows. After the reaction was carried out
for 30 minutes, Ph-FeCl was centrifuged and removed from the reaction mixture. The
residual reaction mixture continued to react for another 12 h. It could be seen from
Figure 4a that the yield of KA oil only had a slight increase in the next 2 h after hot
filtration, and then remained almost unchanged. ICP-AES analysis confirmed the
leaching of the iron species was neglectable (ca. 9 ppm) after reaction, which revealed
that the Ph-FeCl catalyst was highly stable and leaching-resistant. As a result, a slight
increase yield of KA oil after hot filtration might be caused by some free radicals which
have generated in the reaction before hot filtration and continued to involve in the
oxidation reaction. After the free radicals were consumed completely, the yield did not
increase any more after 4 h (Figure 4a), and the yield of KA oil after hot filtration was
much lower than that without hot filtration. The hot filtration test strongly suggested
that the present reaction would proceed in a heterogeneous manner.
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European Journal of Inorganic Chemistry
35
30
25
20
15
10
5
none
1 equiv. of t-butanol
1 equiv. of BHT
No filteration
30
Hot filteration
25
0.2 equiv. of 1,4-benzoquinone
20
15
10
5
0
0
0
2
4
6
8
10
12
0
1
2
3
4
Reaction time /h
Reaction time /h
Figure 4. a) Hot filtration curve of Ph-FeCl catalyst and b) Inhibition effect of free
radical inhibitors on cyclohexane oxidation. Reaction conditions: 4 mmol cyclohexane,
8 mmol (30 wt%)H2O2, 0.017 mmol Ph-FeCl, 3 mL CH3CN, 40 oC. The equivalent
amount of radical inhibitors was based on cyclohexane.
It has been reported that cyclohexane oxidation using H2O2 as oxidant seems to
proceed via a radical mechanism initiating from transition metal catalyzed
decomposition of H2O2. In order to have a deeper understand of the reaction mechanism,
it was important to understand whether the reaction was a free radical-initiated reaction.
As shown in Figure 4b, different radicals annihilators were added to the acetonitrile
solution. The inhibitory effect of 2,6-di-tertbutyl-4-methylphenol(BHT) was very
significant because BHT reacted with oxygen-containing free radicals to form stable
compounds, thus interrupting the oxidative formation of KA oil.[32] Besides, t-butanol
was added to the reaction mixture as an inhibitor for HO· radicals, leading to a slight
loss of the yield of KA oil. This result indicated that HO· radicals could contribute
partially to the cyclohexane oxidation.[33] Moreover, when only 0.2 equivalent amount
of 1,4-benzoquinone(based on cyclohexane) was added as an annihilator for superoxide
radicals(O2·-) or peroxy radicals(HOO·), it could be seen that the reaction was almost
suppressed. This implied the crucial role of superoxide radicals (O2·-)/peroxy radicals
(HOO·) in the oxidation reaction. Although the addition of t-butanol could annihilate
HO· radicals, HO· radicals could be still regenerated through the reaction of peroxy
(HOO·) radicals with hydrogen peroxide,[34] resulting in the slight decrease activity over
cyclohexane oxidation. On the contrary, the use of 1,4-benzoquinone can annihilate O2·-
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10.1002/ejic.202100356
European Journal of Inorganic Chemistry
/HOO· radicals so that HO· radicals could not be generated continuously in the reaction
system. As such, it could be concluded that the cyclohexane oxidation is driven mainly
by the participation of O2·-/ HOO· radical species, and partially by the contribution of
HO· radicals.
The spin trap method is widely used for the fixation of radicals that arise in various
chemical and biochemical reactions.[35] Direct detection of O2·-/ HOO· and HO· active
species is rather difficult, due to their extremely high reactivity and short lifetime (HO·:
10-10 s, [36] O2·-: 51 s [37]). In this work, 5,5-dimethyl-1-pyrroline N-oxide(DMPO) was
.-
used as the spin trap of O2 /HOO· radicals.[38] The possible presence of radicals and
their participation in cyclohexane oxidation on Ph-FeCl catalyst could be identified
since the interaction of N-oxide with a peroxy radical gave rise to the formation of the
radical adduct. As shown in Figure 5, the signals of DMPO kept silent in the absence
of catalyst and oxidant (Spectrum 1). In contrast, once H2O2 was added to the reaction
mixture, the four characteristic peaks with an intensity ratio of 1:1:1:1 appeared
(Spectrum 2), corresponding to the EPR signals of DMPO-OOH adducts (G =3506,
3520, 3532; g = 2.0062).[39]
Spectrum 2
.
.
.
.
Spectrum 1
3490 3500 3510 3520 3530 3540 3550
Magnetic Field /G
Figure 5. EPR spectra of spin adducts observed using DMPO as a spin trap reagent in
the cyclohexane oxidation. Spectrum 1 was observed under the conditions: 0.03 M of
DMPO in 3 mL CH3CN. Spectrum 2 was observed under the conditions: 1.33 M of
cyclohexane, 0.017 mmol of Ph-FeCl, 2.66 M of H2O2, 0.03 M of DMPO, 3 mLCH3CN.
As such, in the light of the characterization by the free radical annihilator and EPR,
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10.1002/ejic.202100356
European Journal of Inorganic Chemistry
the conclusion can be drawn that Ph-FeCl operated likely via a mechanism involving
HOO· radical species. The reaction of H2O2 with the ferric(III) ions can produce FeIV=O
species along with HO·radicals (Eq. 1),[40] while H2O2 afforded HOO· species by the
reaction with FeIV=O species, and meantime ferric(III) ions regenerated in the cycle
under the reaction conditions (Eq. 2). At the same time, O2·-/ HOO· radical species
could be converted into HO· radicals in the presence of H2O2 (Eq. 3).[34b] Notably, the
A/K molar ratio was close to 1.0 on three complexes. This result suggested strongly
that a Fenton-like reaction denominated, in which the hydroxyl radicals HO· are active
species for the sequential oxidation reaction.[15a,15c,41]
Then, the HO· radicals reacted with the cyclohexane generating cyclohexyl
radicals (Cy.) (Eq. 4), which soon reacted with molecular oxygen to form the
cyclohexylperoxyl radical CyOO· (Eq. 5). Integrating cyclohexane with CyOO· radical
would reform the cyclohexyl radicals(Cy.) and generate the cyclohexyl-hydroperoxide
(CyOOH) (Eq. 6). Some CyOO· intermediate species would directly react with HOO·-
radicals to form cyclohexanone(K), dioxygen and water (Eq. 7), and cyclohexanone(K)
could also be generated by homolytic decompositions of CyOOH (Eq. 8). Meantime,
cycloxexanol(A) was generated via the heterolytic (Eq. 9) decompositions of CyOOH.
Notably, cycloxexanol(A) could be oxidized into cyclohexanone(K) by H2O2 likely via
a radical mechanism (Eqs.10 and 11),[42] which was well consistent with the present
obseravtion on Ph-FeCl catalyst.
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Recyclability of the catalyst and substrate scope
The reusability of the Ph-FeCl catalyst was determined because catalyst
recyclability exhibits an indispensable part in the catalytic performance. After each run,
the catalyst was separated from the reaction mixture by filtration, followed by washing,
and then it was dried under vacuum for the next catalytic reaction. As shown in Figure
6a, after five cycles, the catalytic activity of Ph-FeCl almost kept unchanged. The ICP-
AES analysis indicated the leaching of the iron species was neglectable (ca. 9 ppm).The
FT-IR spectra of the spent catalyst Ph-FeCl after the fifth run was almost the same as
that of the fresh one (Figure 6b), confirming the catalyst was quite stable in the course
of reaction recycles.
30
a
b
25
20
15
10
5
Ph-FeCl
The spent Ph-FeCl
0
1
5
4000 3500 3000 2500 2000 1500 1000 500
2Recycl3e times4
Wavenumber /cm-1
Figure 6. a) Catalytic recyclability over Ph-FeCl for cyclohexane oxidation and b) The
FT-IR spectra of Ph-FeCl and the spent Ph-FeCl. Reaction conditions: 4 mmol
cyclohexane, 8 mmol (30 wt%)H2O2, 0.017 mmol Ph-FeCl, 3 mL CH3CN. 40 oC, 4 h.
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Next, it was demonstrated that the present Ph-FeCl catalyst was not only utilized
for the oxidation of sp3 hybrid C-H bonds, but also extended to the oxidation of aromatic
C-H bonds without any other additives. As shown in Table 3, the Ph-FeCl catalyst was
active against both cycloalkanes and bridged alkanes (Table 3, entries 1-4). In addition
to the catalytic activity on sp3 hybrid C-H bonds, Ph-FeCl was also active for the
catalytic oxidation of sp2 hybrid C-H bonds of aromatic compounds (Table 3, entry 5).
Notably, due to the σ-π hyper conjugation between the α-position C-H bond of the
benzene ring and the π system of benzene, the bond energy of the α-position C-H bond
decreased and the activity of α-H increased, which could lead to the product arising
from not only the hydroxylation of the aromatic ring, but also the oxidation of C-H
bonds on the substituted alkyl group (Table 3, entries 6-8). Moreover, the conversion of
benzene homologs with an electron-donating group afforded better results (Table 3,
entries 6-8). In particular, the conversion of cumene could reach up to 22.7% due to the
electron-rich isopropyl group. Because the electron donating abilities of methyl and
ethyl groups are less than that of the isopropyl group, the reactivity of toluene and
ethylbenzene was less than that of cumene.
Table 3. Oxidation of different substrates with Ph-FeCl complex
H2O2
T
(C)
t
Con.
Product
Sel. (%)
Entries Substrates (equiv.)a
(h) (%)
2-methylcyclohexanone (43)
2-methylcyclohexanol (57)
Cyclooctanol (55)
1
2
3
4
5
2
2
2
2
3
r.t
8
4
8
15.6
16.5
7.8
50
60
40
60
Cyclooctanone (45)
2-norbornyl alcohol (49)
2-norbornone (51)
1-adamantanol (55)
2-adamantanone (45)
Phenol(99)
12 14.9
12
4.7
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Benzaldehyde (60)
6
2
2
40
40
12 15.5
Benzoic acid (13)
2-methylphenol (7)
3-methylphenol (20)
Benzaldehyde (5)
7
8
12
20
Acetophenone (72)
1-phenylethanol (17)
2-ethylphenol (3)
3-ethylphenol (3)
Acetophenone (34)
2-phenyl-2-propanol (57)
2-isopropylphenol (3)
2
40
12 22.7
4-isopropylphenol (6)
a
Reaction conditions: 4 mmol substrate, 0.017 mmol Ph-FeCl, 3 mL CH3CN. The
equivalent amount of H2O2 was based on substrates.
Finally, the catalytic performance of the catalytic system was compared with those
reported systems in the oxidation of cyclohexane to KA oil. As shown in Table S3, the
catalytic performance of the current Ph-FeCl is comparable to those of some other
reported catalytic systems using hydrogen peroxide as an oxidant. However, the current
catalytic system exhibits the advantages of cheaper, more available and environmental
friendly in catalyst synthesis. Particularly, the present catalytic system could be
recycled easily due to heterogeneous catalysis and exhibits C-H bond oxidation with a
broad substrate scope, including cycloalkanes, bridged alkanes and aromatic
hydrocarbons. For building a more efficient C-H bond oxidation catalysis system, the
studies to further examine the utility of a platform derived from amino acid Schiff bases
for other reactions are in progress in our group.
Conclusion
This work has developed three reduced amino acid based Schiff base ligands using
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European Journal of Inorganic Chemistry
L-alanine, L-phenylalanine and L-tyrosine as precursors, and then these derived ligands
were coordinated with FeCl3·6H2O to form the corresponding iron(III) complexes Ala-
FeCl, Ph-FeCl and Tyr-FeCl, respectively. After thorough characterization of the
iron(III) complexes, the as-synthesized iron(III) complexes have been employed as
highly efficient catalysts for the oxidation of cyclohexane with H2O2 as an oxidant. In
particular, Ph-FeCl catalyst afforded an approximately 23.2% total yield of KA oil
without adding any co-catalyst at 40 °C in 4 h. Notably, the present Ph-FeCl complex
followed a heterogeneous mechanism for the catalytic oxidation, which allowed the
catalyst to be recycled easily after reaction. Besides, the catalyst showed almost no loss
of activity even after five cycles, arising from high stability of the complex catalyst and
negligible leaching of iron species. Moreover, Ph-FeCl catalyst can be extended to the
activation of the other sp3 hybridized C-H and aromatic sp2 hybridized C-H bonds. EPR
characterization demonstrated that Phen-FeCl operated via a radical mechanism.
Undoubtedly, this work represents a vital research subject in the field of green catalysis
that can be extended definitely for the efficient catalytic organic transformations via an
easily available and sustainable iron-based catalytic system.
Experimental section
Materials
The commercial ferric chloride (FeCl3·6H2O), L-phenylalanine and L-tyrosine
were obtained from Taitan(Shanghai) Chemical Technology Co., Ltd. L-alanine was
obtained from Macklin(Shanghai) Biochemical Technology Co., Ltd. 2-
pyridinecarboxaldehyde and 2,2’-bipyridine(bpy) were obtained from Bide(Shanghai)
Medicine Chemical Reagent Co., Ltd. The spin trap 5,5-dimethyl-1-pyrroline-N-oxide
(DMPO) was obtained from Energy Chemical(Shanghai) Co., Ltd. 1,10-phenanthroline
(phen) was obtained from Aladdin Co., Ltd. Fe2O3 and Fe3O4 were obtained from
Yonghua Chemical Co., Ltd., and their XRD patterns were shown in Figure S6.
Fe(bpy)Cl3(Bpy-FeCl) and Fe(phen)Cl3(Phen-FeCl) were prepared according to the
previous reports (see supporting information),[43] and their FT-IR spectra were shown
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10.1002/ejic.202100356
European Journal of Inorganic Chemistry
in Figure S7. Hydrogen peroxide (30 wt%), free radical inhibitors (2,6-di-tertbutyl-4-
methylphenol(BHT), t-butanol and 1,4-benzoquinone) and other substrates were
obtained from Sinopharm Chemical Reagent Co., Ltd. (Shanghai, China).
Catalyst characterization
The FT-IR spectra were recorded at room temperature on a Nicolet Magna 550 FT-
IR spectrometer. UV-Vis measurements were carried out on a Varian Cary 500
spectrophotometer in optically transparent cells, the path length was 1 cm. The
elemental analysis of C, H and N was performed using an Elementar vario EL III C H
N O S elemental analyzer and the ICP-AES analysis of iron on a Varian 710 instrument,
respectively. The ESI-MS of ligands were recorded on a XEVO G2 TOF mass
spectrometer in CH3OH. The mass spectra of the Fe (III) complex catalysts were
investigated by MALDI-TOF, with α-cyano-4-hydroxycinnamic acid (CHCA) as a
matrix in DMSO. The X-ray diffraction (XRD) analysis was performed with a D/MAX
2550 VB/PC using CuKα radiation (λ = 1.5406 Å) operated at 40 kV and 100 mA (step:
0.02°; scan area: 10-80°). The electron-spin resonance (EPR) spectrum was recorded
with a Bruker EMX-8/2.7 at 298 K, using 5,5-dimethylpyrroline N-oxide (DMPO) as a
trap reagent. Instrument settings: Microwave frequency, 9.880 GHz; Microwave power,
6.412 mW; Modulation frequency, 100 kHz; Modulation amplitude, 1.00 G; Center
field, 3520 G. 1H NMR and 13C NMR spectra were recorded on a Bruker AVANCE 600
MHz instrument using D2O as a solvent. Conductivity measurements were carried out
at 298 K in DMSO with a DDS-307 conductivity meter using a DJS-0.1C electrode.
The procedure used for conductivity measurements was as the following: 4 mL of dry
DMSO was introduced into the cell and the conductivity value of the blank was
measured (0). Subsequently, a trace of iron(III) complexes were dissolved in DMSO
with different concentration (CFe), and the relative conductivity of the catalyst in DMSO
was determined (1), giving the conductivity values of the iron(III) complex solution
by subtraction (/CFe,  = 1-0). The plot of the concentration(CFe) vs conductivity()
gave the slope as the molar conductivity(m). The reaction products were analyzed by
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European Journal of Inorganic Chemistry
gas chromatography (GC) [a Shimadzu GC-2014 gas chromatograph equipped with an
RTX-5 capillary column (30 m × 0.25 mm × 0.25m)]. The products were also
identified by GC-MS [Agilent 6890/5973 GC-MS equipped with an HP-5MS column
(30 m × 0.25 mm × 0.25 m)].
Synthesis of ligands
Preparation of N-(2-Pyridylmethyl) alanine (Ala)
The ligand was prepared following a modified literature procedure.[24] Firstly,
Na2CO3 (1.06 g, 10 mmol) and L-alanine (0.89 g, 10 mmol) were mixed in 10 mL of
water, then 2-pyridinecarboxaldehyde (1.08 g, 10 mmol) in 5 mL ethanol was slowly
added to it. The mixed solution was stirred at room temperature for 2 h and the color of
the solution changed from colorless to yellow. Next, the solution was cooled to 0 ºC,
NaBH4 (0.46 g, 12 mmol) was added slowly, and bubbles were generated. After 2 h, the
pH of solution was adjusted to 4.0-5.0 with HCl. The mixture system was stirred further
for 2 h. Finally, the white solid was filtered, washed with water and ethanol. The as-
obtained solid was dried under vacuum at 60 ºC for 12 h. Yield: 1.61 g, 89.1%. ESI-MS
(CH3OH, m/z): calcd for C9H12N2O2: 180.20, observed: 179.08[M - H]-(Figure S8). 1H
NMR (600 MHz, D2O, ppm): 1.43 (d, JH-H = 7.2 Hz, 3H, CH3), 3.65 (q, JH-H = 7.2 Hz,
1H, NHCH), 4.25 (s, 2H, pyridine-CH2), 7.35 (dd, J1 = 5.0 Hz, J2 = 7.4 Hz, 1H, 5-H,
pyridine), 7.41 (d, JH-H = 8.0 Hz,1H, 3-H, pyridine), 7.79 (dd, J1 = 7.8 Hz, J2 = 7.8
Hz,1H, 4-H, pyridine), 8.47 (d, JH-H = 4.9 Hz, 1H, 6-H, pyridine). 13C NMR (150 MHz,
D2O, ppm): δ 174.74, 150.41, 149.35, 138.43, 124.40, 124.07, 57.68, 49.81, 15.08
(Figure S9). UV-vis (DMSO) λmax, nm (ε, M−1 cm−1): 266(25351) (Table S1, Figure S1).
FT-IR (KBr, cm-1) 3429(N-H), 3059(COOH), 1630(COO-), 1577(C=N pyridine) (Table
S2, Figure S2).
Preparation of ligand N-(2-Pyridylmethyl) phenylalanine (Ph)
The synthesis of Ph was similar to that ofAla, but L-phenylalanine was used instead
of L-alanine. Yield: 2.33 g, 84.9%. ESI-MS (CH3OH, m/z): calcd for C15H16N2O2:
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European Journal of Inorganic Chemistry
256.30, observed: 255.11[M - H]-(Figure S8). 1H NMR (600 MHz, D2O, ppm): 3.25 (m,
2H, phenyl-CH2), 3.93 (t, JH-H = 6.8 Hz, 1H, NHCH), 4.30-4.35 (d, JH-H = 14 Hz, 2H,
pyridine-CH2), 7.30-7.45 (m, 7H, 3-H, 5-H pyridine, 5H, phenyl), 7.88 (t, JH-H = 7.8
Hz, 1H, 4-H, pyridine), 8.52 (d, JH-H = 5.0 Hz, 1H, 6-H, pyridine). 13C NMR (150 MHz,
D2O, ppm): δ 171.46, 147.04, 145.42, 145.14, 143.38, 143.10, 134.02, 129.05, 126.66,
126.39, 62.32, 48.09, 35.58 (Figure S10). UV-vis (DMSO) λmax, nm (ε, M−1 cm−1):
263(21074) (Table S1, Figure S1). FT-IR (KBr, cm-1) 3450(N-H), 3064(COOH),
1656(COO-), 1575(C=N pyridine) (Table S2, Figure S2).
Preparation of ligand N-(2-Pyridylmethyl) tyrosine (Tyr)
The synthesis of Tyr was similar to that of Ala, but L-alanine was substituted by L-
tyrosine. Yield: 2.14 g, 78.5%. ESI-MS (CH3OH, m/z): calcd for C15H16N2O3: 272.30,
1
observed: 271.11[M - H]-(Figure S8). H NMR (600 MHz, D2O, ppm): 3.01 (m, 2H,
phenyl-CH2), 3.70 (t, JH-H = 6.7 Hz, 1H, NHCH), 4.15 (s, 2H, pyridine-CH2), 6.69-7.00
(m, 4H, phenol 4C-H), 7.25 (dd, J1 = 4.9 Hz, J2 = 7.8 Hz, 1H, 5-H, pyridine), 7.29 (d,
JH-H = 8.0 Hz, 1H, 3-H, pyridine), 7.73 (t, JH-H = 7.8 Hz, 1H, 4-H, pyridine), 8.35 (d,
JH-H = 5.0 Hz, 1H, 6-H, pyridine). 13C NMR (150 MHz, D2O, ppm): δ 170.55, 155.16,
146.34, 144.97, 143.71, 130.73, 127.89, 127.39, 125.07, 115.45, 62.03, 46.96, 34.76
(Figure S11). UV-vis (DMSO) λmax, nm (ε, M−1 cm−1): 266(27065) (Table S1, Figure
S1). FT-IR (KBr, cm-1) 3422(N-H), 3056 (COOH), 1610(COO-), 1591(C=N pyridine)
(Table S2, Figure S2).
Synthesis of iron(III) complexes
Ala-FeCl. The aqueous solution (10 mL) of Ala (1 mmol, 0.18 g) was added to a
methanol solution (10 mL) of FeCl3·6H2O (0.5 mmol, 0.14 g). The reaction mixture
was stirred for 2 h and the precipitate was formed. The resultant solid was then dried
under vacuum at 60 ºC to give Ala-FeCl as a yellow powder. Yield 0.26 g (85%).
C18H22Cl4Fe2N4O4 (611.89): Calc. C, 35.33, H 3.62, N 9.16, Fe 18.25; Found: C 35.93,
H 3.98, N 9.08, Fe 18.23. UV-vis (DMSO) λmax, nm (ε, M−1 cm−1): 265(28315),
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10.1002/ejic.202100356
European Journal of Inorganic Chemistry
321(16965) (Table S1, Figure S1). FT-IR (KBr, cm-1) 3281(N-H), 1649(COO-),
1610(C=N pyridine), 1003(Fe-N pyridine), 459(Fe-O), 423(Fe-N) (Table S2, Figure
S2). MALDI-TOF (DMSO, CHCA as matrix, m/z) Calc. for [Fe2(Ala)2Cl4] 611.89.
Found [Fe2(Ala)2Cl4 + 2H - Cl]+ 579.00; [Fe(Ala)Cl2 + CHCA]+ 495.05 (Figure S4).
The complexes Ph-FeCl and Tyr-FeCl were also prepared as yellow powder in the
similar method, except the ligand Ala was replaced by Ph or Tyr, respectively.
Ph-FeCl. Yield 0.21 g (70%). C30H30ClFeN4O4 (601.88): Calc. C 59.87, H 5.02, N
9.31, Fe 9.28; Found: C 60.49, H 5.46, N 9.18, Fe 9.33. UV-vis (DMSO) λmax, nm (ε,
M−1 cm−1): 261(18595), 320(7952) (Table S1, Figure S1). FT-IR (KBr, cm-1) 3244(N-
H), 1672(COO-), 1607(C=N pyridine), 1002(Fe-N pyridine), 450(Fe-O), 425(Fe-N)
(Table S2, Figure S2). MALDI - TOF (DMSO, CHCA as matrix, m/z) Calc. for
[Fe(Ph)2Cl] 601.88. Found [Fe(Ph)2Cl - Cl]+ 566.38; [Fe(Ph)Cl + 2Cl - 2H]+ 415.24;
[Ph + H]+ 257.14 (Figure S4).
Tyr-FeCl. Yield 0.26 g (69%). C30H30ClFeN4O6 (633.89): Calc. C 56.84, H 4.77;
N 8.84, Fe 8.81; Found: C 56.72, H 4.92, N 8.10, Fe 8.96. UV-vis (DMSO) λmax, nm (ε,
M−1 cm−1): 260(24530), 290(19910), 324(10229) (Table S1, Figure S1). FT-IR (KBr,
cm-1) 3249 (N-H), 1664(COO-), 1612(C=N pyridine), 1004(Fe-N pyridine), 454(Fe-O),
424(Fe-N) (Table S2, Figure S2). MALDI-TOF (DMSO, CHCA as matrix, m/z) Calc.
for [Fe(Tyr)2Cl] 633.89. Found [Fe(Tyr)2Cl - Cl - H - CO2]+ 553.23; [Fe(Tyr)Cl + Cl +
2H]+ 401.04; [Tyr + H]+ 273.09 (Figure S4).
Typical procedure for catalytic oxidation of cyclohexane
For the cyclohexane oxidation reaction, cyclohexane (0.336 g, 4 mmol) and the
iron catalyst (0.017 mmol) were added into CH3CN (3 mL) in a 25 mL reaction tube.
H2O2 (30 wt%, 0.9 g, 8 mmol) was gradually added dropwise as the oxidant, and the
reaction mixture was stirred at 40 ºC for 4 h. After the reaction, 2 mL ethyl acetate was
added to the mixture, and the catalyst was centrifuged, and the clear organic phase was
reduced by PPh3 according to the previous method reported by Shul’pin.[44] The yield
of products was determined by GC analysis using toluene as the internal standard.
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European Journal of Inorganic Chemistry
Quantitative determination of cyclohexane, cyclohexanol(A) and cyclohexanone(K)
were carried out by GC using the individual calibration curve method.
The rate of H2O2 decomposition on the iron(III) complexes
The rate of H2O2 decomposition on the iron(III) complexes (Ala-FeCl, Ph-FeCl
and Tyr-FeCl) were determined. The concentration of residual H2O2 in the liquid phase
was quantified by potential difference titration of Ce3+/Ce4+. The typical procedure was
described as following: after the reaction, the reaction liquid was separated by
centrifugation, and then one drop of the mixed aqueous solution of 1,10-phenanthroline
(0.075 M) and FeSO4 (0.025 M) (indicator) and two drop of the concentrated HCl (37%)
were added into the reaction mixture. The end point of titration was determined until
the solution turned blue and did not change within 30 s. The blank sample was also
tested in the absence of cyclohexane and catalyst.
Acknowledgements
The authors thank for support from the National Natural Science Foundation of China
(21773061, 21978095) and the Open Research Fund of Shanghai Key Laboratory of Green
Chemistry and Chemical Processes.
Conflict of interest
The authors declare no conflict of interest.
Keywords: Amino acid • Cyclohexane • Iron(III) complex • Oxidation• Schiff base
References
[1] G. B. Shul’pin, Y. N. Kozlov, L. S. Shul’pina, Catalysts 2019, 9, 1046.
[2] a) J. M. Thomas, Nature 1985, 314, 669-670; b) T. Ito, J. H. Lunsford, Nature 1985,
314, 721-722.
[3] C. C. Liu, T. S. Lin, S. I. Chan, C. Y. Mou, J. Catal. 2015, 322, 139-151.
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