Simple soluble Bi(iii) salts as efficient catalysts for the oxidation of alkanes with H2O2

Article abstract of DOI:10.1039/c4cy01651c

A simple catalytic system based on a soluble bismuth(iii) salt, Bi(NO₃)₃/H₂O₂/HNO₃/CH₃CN + H₂O, exhibits pronounced activity towards the homogeneous oxidation of inert alkanes with the yield of oxygenate products up to 32% and TON up to 112. The experimental selectivity parameters and kinetic data together with theoretical DFT calculations indicate that the reaction occurs via a free radical mechanism involving the formation of the HO radicals which directly react with alkane molecules. The mechanism of the HO generation (which is the rate limiting step of the whole process) includes the substitution of a water ligand for H₂O₂ in the initial aqua complex [Bi(H₂O)₈]³⁺, hydrolysis of the coordinated H₂O₂, second H₂O-for-H₂O₂ substitution and the homolytic HO-OH bond cleavage in complex [Bi(H₂O)₄(H₂O₂)(OOH)]²⁺ (6). The relatively low overall activation energy for this process (ca. 20 kcal mol^-1) is accounted for by the high lability and acidity of the Bi aqua complexes and tremendous activation of coordinated H₂O₂ in 6 towards homolysis. This journal is

Full text of DOI:10.1039/c4cy01651c

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Simple soluble Bi(III) salts as efficient catalysts for the oxidation of
†
alkanes with H2O2
Bruno G. M. Rocha,a Maxim L. Kuznetsov,*a Yuriy N. Kozlov,b Armando J. L.
Pombeiroa and Georgiy B. Shul’pin*b
Abstract
Simple catalytic system based on a soluble bismuth(III) salt Bi(NO3)3/H2O2/HNO3/CH3CN+H2O
exhibits pronounced activity towards the homogeneous oxidation of inert alkanes with the yield of
oxygenate products up to 32 % and TON up to 112. The experimental selectivity parameters and
kinetic data together with theoretical DFT calculations indicate that the reaction occurs via a free
radical mechanism involving the formation of the HO• radicals which directly react with alkane
molecules. The mechanism of the HO• generation (which is the rate limiting step of the whole
process) includes the substitution of a water ligand for H2O2 in the initial aqua complex
[Bi(H2O)8]3+, hydrolysis of the coordinated H2O2, second H2OꢀforꢀH2O2 substitution and the
homolytic HO–OH bond cleavage in complex [Bi(H2O)4(H2O2)(OOH)]2+ (6). The relatively low
overall activation energy for this process (ca. 20 kcal/mol) is accounted for by the high lability and
acidity of the Bi aqua complexes and tremendous activation of coordinated H2O2 in 6 towards
homolysis.
Keywords: bismuth, hydrogen peroxide, alkanes, oxidation, DFT calculations, reaction mechanism,
homogeneous catalysis, hydroxyl radical, water exchange.
a
Centro de Química Estrutural, Instituto Superior Técnico, Universidade de Lisboa, Av. Rovisco Pais, 1049-001
Lisbon, Portugal. E-mail: max@mail.ist.utl.pt; Tel: +351 218419236
b
Semenov Institute of Chemical Physics, Russian Academy of Sciences, ulitsa Kosygina, dom 4, Moscow 119991
Russia. E-mails: Shulpin@chph.ras.ru, gbsh@mail.ru
†
Electronic supplementary information (ESI) available: Complete references 16 and 24, computational details, figures
with experimental results, tables with total energies, enthalpies, entropies, Gibbs free energies, and Cartesian atomic
coordinates. See DOI: XXX
1

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Introduction
Hydrocarbons are the most abundant carbon raw material and a search for their mild and efficient
conversion to functionalized valuable products is an important problem of modern chemistry.1 Such
a transformation of alkanes as the most chemically inert hydrocarbons requires the presence of an
appropriate oxidant and an efficient catalyst. Hydrogen peroxide is among the most promising
oxidants from the ecological point of view. Oxidation reactions of alkanes with H2O2 are usually
catalysed by various transition metal complexes (e.g. those of Fe, V, Mo, Mn, Ti, Os, Cu and
others).1 In contrast, the application of nonꢀtransition metal complexes in the homogeneous
oxidation of simple alkanes containing nonꢀactivated C–H bonds with H2O2 is very little explored.
The first work on this topic was previously published by some of us.2 It was experimentally
demonstrated that the simple catalytic system based on an aluminium salt, Al(NO3)3/H2O2,
efficiently oxidizes alkanes in aqueous acetonitrile (total yield of oxygenate products up to 31 %).
Other recently published preliminary experimental data revealed that the salts of such nonꢀtransition
metals unconventional for the functionalization of alkanes as beryllium, zinc and cadmium
[M(NO3)2], also exhibit catalytic activity in the oxidation of cyclooctane with H2O2.3
Bismuth and its species are very little explored but promising nonꢀtransition metal catalysts
for the hydrocarbon oxidation.4 Examples of the works on this topic include bismuthꢀcatalysed
allylic or benzylic oxidations with tBuOOH,5 benzylic oxidations by sodium bismuthate NaBiO3 in
acetic acid,6 the heterogeneous propylene and cyclohexane oxidations by molecular oxygen7 or
8
H2O2 catalysed by Biꢀcontaining SBAꢀ15, MCMꢀ41, ZSMꢀ5 mesoporous molecular sieves or the
MoO3–Bi2SiO5/SiO2 system, as well as the oxidation of propylene or isobutene to acrolein over the
10
11
BixMyOz (M = Mo, W),9 Bi1–x/3V1–xMoxO4 or Bi2Sn2O7 catalysts and epoxidation of stilbene by
tBuOOH with [Cp2Mo(Bi{OCH(CF3)2}2)2].12 The interaction of bismuth oxide clusters with
ethylene in the presence of O2 was also theoretically studied.13
However, to our knowledge, there are no publications on the homogeneous oxidation of
alkanes containing nonꢀactivated C–H bonds with H2O2 catalysed by Bi compounds. In this work,
2

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we discovered that the simple soluble Bi salts can efficiently catalyse the oxidation of simple inert
alkanes with hydrogen peroxide in aqueous acetonitrile solution. The detailed experimental study
demonstrated that the reaction occurs via a Fentonꢀlike mechanism including the formation of
hydroxyl radicals which attack the C‒H bonds of the alkane. The plausible mechanism of the HO•
radicals generation in this system was investigated by theoretical methods.
Experimental section
Reagents. Salt Bi(NO3)3•5H2O (98 %), hydrogen peroxide (50 % aqueous solution), cyclooctane
(99.5 %), nꢀoctane (99 %), methylcyclohexane (98 %), transꢀ/cisꢀdimethylcyclohexane (99 %) and
HNO3 (65 %) were used as received from SigmaꢀAldrich.
Alkane oxidation. Catalyst was used as a stock solution prepared by simultaneous
dissolving of Bi(NO3)3•5H2O and HNO3 (65 % aqueous) in acetonitrile. The reactions of alkanes
were carried out in air in thermostated Pyrex cylindrical vessels with vigorous stirring and using
CH3CN as solvent. Catalyst and coꢀcatalyst (nitric acid) were introduced into the reaction mixture
in the form of a stock solution in acetonitrile. Total volume of the reaction mixture was 10 mL. The
substrate was then added and the reaction started when hydrogen peroxide was introduced in one
portion. (CAUTION: the combination of air or molecular oxygen and H2O2 with organic
compounds at elevated temperatures may be explosive!). The reactions were stopped by cooling and
addition of excess of solid PPh3. Solutions were analysed after addition of nitromethane as a
standard compound by GC (instrument ‘HP 5890 – SerieꢀII’; fused silica capillary columns column
HewlettꢀPackard; the stationary phase was polyethyleneglycol: INNOWAX with parameters 25 m ×
0.2 mm × 0.4 µm; carrier gas was helium with column pressure of 15 psi) using the method
developed previously by one of us.1i,14 This method allows the estimate of concentrations of formed
alkyl hydroperoxide, alcohol and ketone. Attribution of peaks was made by comparison with
chromatograms of authentic samples.
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Computational details. The full geometry optimization of all structures and transition states
(TS) has been carried out at the DFT/HF hybrid level of theory using Becke’s threeꢀparameter hybrid
exchange functional in combination with the gradientꢀcorrected correlation functional of Lee, Yang
and Parr (B3LYP)15 with the help of the Gaussianꢀ0916 program package. The quasiꢀrelativistic
Stuttgart pseudopotential that described 78 core electrons (MWB78) and the appropriate contracted
basis set17 were employed for the Bi atoms. This basis set was augmented by addition of the fꢀ
function with exponent 0.305 optimized using the utility gauopt for the Bi atom in the ground state.
The standard basis set 6ꢀ311+G(d,p)18 was applied for all other atoms. No symmetry operations have
been applied for any of the structures calculated. Restricted approximations for the structures with
closed electron shells and unrestricted methods for the structures with open electron shells have
been employed.
Taking into account the importance of the consideration of the second coordination sphere in
the calculations of solvent effects for reactions involving highly charged species, one solvent
molecule (H2O or H2O2) was added explicitly in the second shell of the calculated structures. It was
demonstrated in the previous works3,19,20 that the combination of the B3LYP functional, the basis
sets used and the model with one explicit solvent molecule in the second coordination sphere
provide an excellent agreement between the calculated and experimental activation energies of the
water substitution reactions – two relevant steps of the calculated mechanism. Additionally, the O–
O bond dissociation energy in H2O2 calculated at this level of theory (48.7 kcal/mol, the total gasꢀ
phase energy scale) is very close to the experimental value of 48.75 ± 0.005 kcal/mol.21
The Hessian matrix was calculated analytically for the optimized structures in order to prove
the location of correct minima (no imaginary frequencies) or saddle points (only one imaginary
frequency) and to estimate the thermodynamic parameters, the latter being calculated at 25 °C. The
nature of all transition states was investigated by the analysis of vectors associated with the
imaginary frequency and, in some cases, by the calculations of the intrinsic reaction coordinates
(IRC) using the GonzalezꢀSchlegel method.22
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Total energies corrected for solvent effects (Es) were estimated at the singleꢀpoint
calculations on the basis of gasꢀphase geometries using the polarizable continuum model in the
CPCM version23 with CH3CN or, in some cases, H2O as solvents. The calculations of the solvent
effect were carried out using Gaussianꢀ0324 program package on the basis of gasꢀphase equilibrium
structures found with Gaussianꢀ09 to provide direct comparison of results obtained in this work and
previously.3,19,20 The UAKS model was applied for the molecular cavity. The entropic term for the
CH3CN solvent (Ss) was calculated according to the procedure described by Wertz25 and Cooper and
Ziegler26 using equations (C1) – (C4)
ꢀS1 = R ln(Vsm,liq/Vm,gas
ꢀS2 = R ln(V°m/Vsm,liq
)
(C1)
(C2)
(C3)
)
α = [S°,sliq – (S°,sgas + ꢀS1)]/[S°,sgas + ꢀS1]
Ss = Sg + ꢀSsol = Sg + [ꢀS1 + α(Sg + ꢀS1) + ꢀS2] =
Sg + [(–12.21 cal/mol•K) – 0.23(Sg – 12.21 cal/mol•K) + 5.87 cal/mol•K]
(C4)
where Sg = gasꢀphase entropy of solute, ꢀSsol = solvation entropy, S°,sliq, S°,s and Vsm,liq
=
gas
standard entropies and molar volume of the solvent in liquid or gas phases (149.62 and 245.48
J/mol•K and 52.16 mL/mol, respectively, for CH3CN), Vm,gas = molar volume of the ideal gas at
25 °C (24450 mL/mol), V°m = molar volume of the solution corresponding to the standard
conditions (1000 mL/mol). The Ss values for the H2O solvent were calculated using the equation
(C5).26
Ss = Sg + [(–14.3 cal/mol•K) – 0.46(Sg – 14.3 cal/mol•K) + 7.98 cal/mol•K]
(C5)
The enthalpies and Gibbs free energies in solution (Hs and Gs) were estimated using the expressions
Hs = Es + Hg – Eg and Gs = Hs – T•Ss, where Eg and Hg are the gasꢀphase total energy and enthalpy.
The relative energies discussed in the text are Gibbs free energies in solution if not stated otherwise.
In the experimental study, the mixed solvent CH3CN+H2O was used. In calculations, the
water part of this solvent was approximated by the explicit inclusion of one H2O molecule in the
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second coordination shell, whereas the acetonitrile part of the solvent was approximated by the
CPCM calculations of the bulky solvent effect for CH3CN.
Results and Discussion
Main features of alkane oxidation catalysed by bismuth salts
For this study, the catalytic system Bi(NO3)3/H2O2/HNO3/CH3CN+H2O was used. The catalyst was
introduced in the form of nitrate Bi(NO3)3•5H2O. Nitric acid was used to prevent the deep
hydrolysis of the bismuth salt and formation of insoluble mono and polynuclear hydroxoꢀ and oxoꢀ
complexes.
We have found that bismuth(III) nitrate catalyzes oxygenation of alkanes with hydrogen
peroxide in acetonitrile if nitric acid is added in a small concentration to the reaction solution.
Kinetic study of the cyclooctane oxidation was carried out by the GC method. The direct injection
of a reaction sample into the chromatograph gave comparable amounts of cyclooctanol and
cyclooctanone. An example of the experiments with direct injection is shown in Figure 1, graph A.
The reduction of the sample with PPh3 prior to the GC analysis lead to the exclusive formation of
the alcohol, while no ketone is observed (Figure 1, graph B; see also Figures S7 and S8; for this
method developed by Shul’pin, see Refs. 1i and 14). Such behaviour clearly indicates that
cyclooctyl hydroperoxide ROOH is formed as the main primary product.1i,14 The formation of
cyclooctyl hydroperoxide and its reduction to cyclooctanol with PPh3 is shown in Scheme 1.
Figure 2 represents two examples of kinetic curves of oxygenate accumulation (mainly
cyclooctyl hydroperoxide determined in the form of cyclooctanol after reduction with PPh3) under
various conditions. The maximum turnover number (TON) of 112 was attained under experimental
conditions shown in Graph A, whereas Graph B demonstrates the experiment affording maximum
yield 32 % based on initial cyclooctane. Yields of products under various conditions at different
concentration of the bismuth salt are presented in Figures S1–S4.
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0.04
0.02
A
before PPh3
-one
-ol
0
0.06
B
after PPh3
-ol
0.04
0.02
0
-one
5
10
15
20
25
30
Time (h)
Fig. 1 Accumulation of oxygenates with time in the cyclooctane (initial concentration 0.25 M)
oxidation with H2O2 (1.2 M) catalysed by Bi(NO3)3 (1.0 × 10‒3 M) in the presence of HNO3 (0.072
M) in acetonitrile (contains H2O, 2.4 M) at 60 °C. Concentrations of cyclooctanol (ꢀol) and
cyclooctanone (ꢀone) were determined before (graph A) and after (graph B) reduction of the
reaction samples with PPh3.
H2O2 and O2
catalyst Bi(III)
OH
OOH
+ PPh3
− O=PPh3
Scheme 1 Oxidation of cyclooctane to cyclooctyl hydroperoxide and reduction of the latter with
triphenylphosphine to cyclooctanol.
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0.03
A
0.02
0.01
1
2
0
0.08
B
0.06
0.04
0.02
1
2
0
0
5
10
15
20
25
30
Time (h)
Fig. 2 Accumulation of oxygenates (determined as cyclooctanol, curve 1, and cyclooctanone, curve
2, after reduction with PPh3) with time in the cyclooctane (initial concentration 0.25 M) oxidation
with H2O2 (1.2 M) catalysed by Bi(NO3)3 and HNO3 (0.072 M) in acetonitrile (contains H2O, 2.4
M). Graph A: [Bi(NO3)3]0 = 2.5 × 10–4 M; at 60 °C. TON = 115; yield = 12 %. Graph B:
[Bi(NO3)3]0 = 1.0 × 10‒3 M; at 75 °C TON = 80; yield = 32 %.
In special experiments, we studied the formation of oxygenates in the absence of the
bismuth salt but in the presence of nitric acid and also in the absence of this coꢀcatalyst (Figure 3).
It can be seen that in the absence of Bi(NO3)3 the oxidation proceeds significantly less efficiently
with the yield of oxygenates of less than 3 % after 32 hours vs. 21 % in the presence of Bi(NO3)3
under the same conditions (Figures 1B and 3). In the absence of both catalyst and coꢀcatalyst (nitric
acid), the reaction affords negligible concentrations of oxygenates (Figure 3, curves 1a and 2a).
Dependence of the initial oxidation rate W0 on the concentration of added HNO3 in the “blank”
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experiment is shown in Figure 4. The initial rate W0 was determined from the slope of tangent to the
kinetic curve of total oxygenate accumulation (Figure 3, line 1b).
1
0.002
0.001
0.008
0.004
0
1b
1a
2a
2
0
0
5
10
15
20
25
30
Time (h)
Fig. 3 “Blank” experiments in the absence of Bi(NO3)3. Accumulation of oxygenates (determined
as cyclooctanol, curve 1 and cyclooctanone, curve 2 after reduction with PPh3) with time in the
cyclooctane (0.25 M) oxidation with H2O2 (1.2 M) in the presence of HNO3 (0.072 M) in
acetonitrile (contains H2O, 2.4 M) at 60 °C. Analogous experiment in the absence of both Bi(NO3)3
and HNO3 is presented by curves 1a and 2a. The tangent of kinetic curve 1 is shown by dotted
straight line 1b.
2
1
0
0
0.2 0.4 0.6 0.8 1.0
Concentration 10×c, [HNO3]0 (M)
Fig. 4 Dependence of initial reaction rate W0 on concentration of nitric acid in the cyclooctane
oxidation in the absence of Bi(NO3)3. Conditions: [cyclooctane]0 = 0.25 M, [H2O2]0 = 1.2 M (50%
aqueous), solvent MeCN, 60 °C.
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Using cyclooctane as a model substrate we studied dependences of the initial reaction rate
W0 (based on the sum of cyclohexanol and cyclohexanone concentrations measured after reduction
of the reaction sample with PPh3) on the initial concentration of each reactant at fixed
concentrations of all other components of the reaction solution. These dependences are shown in
Figures 5‒8. The initial oxidation rate W0 linearly depends on the initial concentration of Bi(NO3)3
(Figure 5), while the dependence on the concentration of added nitric acid is more complex (Figure
6). Such particular sigmoidal form of the latter trend may be accounted for by the hydrolysis
processes. In the absence of HNO3, the system is heterogeneous due to hydrolysis of the catalyst to
the insoluble oxo and hydroxo mono and polynuclear species. The addition of a low amount of
HNO3 (ca. 0.2 M) is not sufficient to suppress the hydrolysis, and the catalyst is still inactive.
Further addition of the acid results in an efficiently preventing of the hydrolysis and in
homogenization of the system providing the rapid growth of the W0 value. Finally, addition of
HNO3 at even higher concentrations leads to the saturation because the catalytic system is already
homogeneous and hydrolysis is mostly suppressed. Additionally, due to a number of equilibria
which exist in the reaction system (see below, Scheme 3) including pHꢀdependent equilibrium we
have polyfactor influence of the acid presence. Note, that the hydrolysis and/or diꢀ or even
polymerization of the catalyst are inevitable, in some extent, even in the presence of a significant
amount of HNO3 due to the highly acid nature of Bi(III), and these processes should affect the TON
values and yield of the products.
The curve of dependence of W0 vs. the initial concentration of cyclooctane approaches a
plateau at [RH]0 > 0.2 M as shown in Figure 7, curve 1. Addition of water decreases the initial
reaction rate (Figure 8, curve 3). We carried out the Biꢀcatalysed oxidation of cyclooctane at
different temperatures (Figure S5) what allowed the estimate of the effective activation energy of 20
± 2.5 kcal mol‒1 (Figure 9).
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2
1
0
0
2
3
1
5
6
4
Concentration c, 103 × [Bi(NO3)3]0 (M)
Fig. 5 Dependence of the initial cyclooctane oxidation rate W0 on the initial concentration of
catalyst Bi(NO3)3 in the oxidation of cyclooctane (0.25 M) with H2O2 (1.2 M) in the presence of
HNO3 (0.072 M). Solvent MeCN, 60 °C.
8
6
4
2
0
0
0.2 0.4 0.6 0.8 1.0
1.2
Concentration 10×c, [HNO3]0 (M)
Fig. 6 Dependence of the initial cyclooctane oxidation rate W0 on the initial concentration of HNO3
in the oxidation of cyclooctane (0.25 M) with H2O2 (1.2 M) catalysed by Bi(NO3)3 (1.0 × 10‒3 M) in
acetonitrile at 60 °C.
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2
1.6
1.2
0.8
0.4
0
6
4
2
0
1
0.4
Concentration c, [cyclooctane]0 (M)
0.2
0.8 1.0
0
0.6
Fig. 7 Curve 1: dependence of oxidation rate W0 on initial concentration of cyclooctane (RH) in the
oxidation of cyclooctane with H2O2 (1.2 M) in acetonitrile (contains H2O, 2.4 M). Conditions:
[Bi(NO3)3]0 = 1.0 × 10‒3 M; [HNO3]0 = 0.072 M; 60 °C. Concentrations of the products
(cyclooctanol and cyclooctanone) were measured after reduction with PPh3. Line 2: linearization of
dependence presented by curve 1 using coordinates [RH]0 – [RH]0/W0.
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8
3
4
1
6
4
2
0
4
8
0
[H2O]total (M)
0.4
0.2
0
1
2
0
0.5
1
1.5
2
Concentration c, [H2O2]0 (M)
Fig. 8 Curve 1: dependence of the initial cyclooctane oxidation rate W0 on initial concentration of
H2O2 in the oxidation of cyclooctane (0.25 M) catalysed by Bi(NO3)3 (1.0 × 10‒3 M) in the presence
of HNO3 (0.007 M) in acetonitrile at 60 °C. Concentration of water [H2O]total was constant = 3.53
M. Curve 2: linearization of dependence presented by curve 1 using coordinates [H2O2]0 –
[H2O2]0/W0. Curve 3: dependence of W0 on total concentration of water present in the reaction
solution in the oxidation of cyclooctane (0.25 M) with H2O2 (1.2 M) catalysed by Bi(NO3)3 (1.0 ×
10‒3 M) in the presence of HNO3 (0.007 M) in acetonitrile at 60 °C.
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−16.0
−15.0
−14.0
−13.0
0.0029
0.0031
0.0030
1/T (K−1)
Fig. 9 The Arrhenius plot for the cyclooctane oxidation with H2O2 catalysed by complex Bi(NO3)3.
Conditions: [Bi(NO3)3]0 = 1.0 × 10‒3 M; [cyclooctane]0 = 0.25 M; [H2O2]0 = 1.2 M; [HNO3]0 =
0.072 M) in acetonitrile (contains H2O, 2.4 M). This dependence corresponds to the effective
activation energy Ea = 20 ± 2.5 kcal mol˗1. The original kinetic curves at different temperatures are
presented in Figure S5.
A comparison of kinetic curves for the cyclooctane oxidations catalysed by bismuth and
aluminum salts in the presence and absence of nitric acid is demonstrated in Figure S6. It can be
seen that effect of addition of nitric acid is weaker in the case of aluminum in comparison with
bismuth.
Selectivity of alkane oxidation
We measured the selectivity parameters of oxidations of certain linear and branched alkanes
with H2O2 catalysed by the bismuth salt. These values coincide with the parameters determined
previously for other systems based on vanadium, iron, osmium, nickel, rhenium and aluminium.2,27
These systems have been shown to oxidize alkanes with the formation of free hydroxyl radicals.
Thus, the selectivity parameter in the nꢀoctane oxidation (see Figure S7) is C(1) : C(2) : C(3) : C(4)
= 1.0 : 4.4 : 4.3 : 4.0 (reactivities of the H atoms at carbons 1, 2, 3 and 4 of the nꢀoctane chain
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which are relative normalized taking into account the number of hydrogen atoms at each carbon).
The selectivity parameter in the oxidation of methylcyclohexane (see Figure S8) is 1º : 2º : 3º = 1.0 :
4.2 : 16 (the relative normalized reactivities of hydrogen atoms at primary, secondary and tertiary
carbons of methylcyclohexane). The oxidation of cisꢀ1,2ꢀdimethylcyclohexane and transꢀ1,2ꢀ
dimethylcyclohexane proceeds nonꢀstereoselectively (see Table 1 and Figure S9) with the trans : cis
ratio being 0.7 : 0.8 (the ratio of isomers of tertꢀalcohols with mutual transꢀ and cisꢀorientation of
the methyl groups formed in the oxidation of cisꢀ and transꢀ1,2ꢀdimethylcyclohexane).
Table 1 Oxidation of cisꢀ and transꢀisomers of 1,2ꢀdimethylcyclohexanea
Content of products (%)
Entry Time (h)
transꢀ1,2ꢀDimethylꢀ1ꢀhydroxycyclohexane cisꢀ1,2ꢀDimethylꢀ1ꢀhydroxycyclohexane
Oxidation of cisꢀ1,2ꢀdimethylcyclohexane
1
2
3
4
1
2
3
8
42
41
39
35
58
59
61
65
Oxidation of transꢀ1,2ꢀdimethylcyclohexane
5
6
7
8
1
2
3
8
51
43
42
38
49
57
58
62
a Reaction conditions: [1,2ꢀdimethylcyclohexane]0 = 0.25 M, [Bi(NO3)3]0 = 1.0 × 10‒3 M; [H2O2]0 =
1.2 M; [HNO3]0 = 0.072 M in acetonitrile (contains H2O, 2.4 M), 45 °C.
Kinetic analysis of cyclooctane oxidation
Selectivity parameters measured for the oxidation of linear and branched alkanes by the
Bi(NO3)3/H2O2/HNO3/CH3CN+H2O system C(1) : C(2) : C(3) : C(4), 1° : 2° : 3° and trans : cis
indicate that the oxidation species operating in this system is hydroxyl radical (see previous
section). Our kinetic data also support an assumption that hydroxyl radicals attack C‒H bonds in
saturated hydrocarbons. In the kinetic analysis described below we operate with the initial rate of
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the cyclooctyl hydroperoxide (ROOH) formation, W0 = (d[ROOH]/dt)0, which is equal to the initial
rate of the oxygenate formation. This initial rate W0 was determined from the slope of a dotted
straight line which is tangent to the kinetic curve.
The mode of dependence of the alkane oxidation rate on its initial concentration (see Figure
7, curve 1) are in agreement with our second assumption about a competition between substrate
(alkane) and solvent (acetonitrile) in the interaction with hydroxyl radicals generated by the system
under consideration. It is reasonable to propose the simplest kinetic scheme which includes a stage
of the radical generation (1) and two parallel stages of their consumption (2) and (3):
H2O2 → → HO•
(1)
(2)
(3)
HO• + RH → → ROOH
HO• + CH3CN → → products
Let us designate the hydroxyl radical generation rate in stage (1) as Wi. A detailed
discussion of this parameter is given below. In the sequence of transformations of RH into ROOH,
the rateꢀlimiting step (with constant k2) is the reaction between hydroxyl radical and the alkane
leading to the formation of alkyl radical R•. Reaction (3) with constant k3 is a rateꢀlimiting stage in
the sequence of transformations (3) which affords products of acetonitrile oxidation. If we assume
that concentration of hydroxyl radicals is quasiꢀstationary, analysis of the kinetic scheme (1)–(3),
gives the following expression for the dependence of the ROOH accumulation rate on the
concentration of cyclooctane, RH:
k2[RH]0Wi
d[ROOH]
=
W0 =
dt
k2[RH]0 + k3[CH3CN]0
0
(4)
(5)
The following is the same dependence in more convenient form for the analysis form:
k3
k2
[RH]0
1
[CH3CN]0
=
[RH]0 +
W0
Wi
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Straight line 2 in Figure 7 clearly testifies that the experimental dependence presented by curve 1 in
this Figure satisfies eq. (5) because term [RH]0/W0 linearly depends on [RH]0. For the conditions of
experiments reflected by Figure 7, the following parameters of eq. (5) were calculated:
k3[CH3CN]0/k2 = 0.14 M and Wi = 6 × 10‒7 M s‒1. The first ratio characterizes the rates of
interactions between hydroxyl radicals and acetonitrile and cyclooctane. The values of parameters
k3[CH3CN]/k2 = 0.14 M and k3/k2 = 0.008 (taking into account that [CH3CN]0 = 17 M) measured
using the kinetic data presented in Figure 7 agree with the parameters expected for free hydroxyl
radical. Indeed, these values for the catalytic systems which generate hydroxyl radicals have been
reported to vary in the intervals k3[CH3CN]/k2 = 0.10–0.20 M and k2/k1 = 0.006–0.012 (Table 2).
Table 2 Kinetic parameters for the oxidation of cyclohexane and acetonitrile with various systems
a
based on H2O2
Entry
System
k3[CH3CN]/k2 (M)
k3/k2
Ref.
This work
28a
1
2
3
4
5
6
7
H2O2/O2/Bi(NO3)3/HNO3
H2O2/O2/(nꢀBu4N)VO3/PCA
H2O2/O2/“Cu4”/CF3COOH
H2O2/O2/“Cu4”/HCl
0.14
0.008
0.008
0.012
0.006
0.008
0.14
0.20
28b
0.10
28b
H2O2/O2/[Co4Fe2OSae8]/HNO3
H2O2/O2/Cp*2Os/py
0.14
28c
0.09–0.19
0.19
0.0055–0.011
0.011
28d
H2O2/O2/ Cp2Fe/Py/PCA
H2O2/O2/“Fe2(TACN)”/PCA
28e
8
a
0.19
0.011
27a
Abbreviations: PCA = pyrazineꢀ2ꢀcarboxylic acid, “Cu4” = tetracopper(II) triethanolaminate
[O⊂Cu4{N(CH2CH2O)3}4(BOH)4][BF4]2, [Co4Fe2OSae8] complex
[Co4Fe2OSae8]·4DMF·H2O, where H2Sae salicylideneꢀ2ꢀethanolamine, Cp*2Os
complex
=
=
=
decamethylosmocene, Cp2Fe = ferrocene, “Fe2(TACN)” = an ironIII) complex with 1,4,7ꢀ
triazacyclononane. Concentration [CH3CN] was assumed to be 17 M.
Experimentally found dependence of the cyclooctane oxidation rate W0 on [H2O2]0 (Figure
8) indicates the intermediate formation of an adduct A of the Bi species with hydrogen peroxide.
Due to this, it is reasonable to add a stage of this complex formation (6) to the previously
considered kinetic scheme (1)‒(3):
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K6
“Bi” + H2O2
A
(6)
Here “Bi” is the bismuth species introduced into the reaction mixture. Now we can say that the
transformation (1) which affords hydroxyl radical is a monomolecular decay of the hydrogen
peroxide complex A, and this decomposition is a rateꢀlimiting stage:
A → → HO•
(1a)
Let us assume that concentration [A] is quasiꢀstationary, that is k‒6 >> k1, where k‒6 is the
rate constant of the reversed reaction in stage (6) and k1 is the rate constant of the monomolecular
transformation (1a) of A into hydroxyl radical with the rate Wi. Analysis of the sequence of
transformations (6) and (1a) gives the following expression for the rate Wi taking into account that
[“Bi”]0 << [H2O2]0:
k1K6["Bi"]0[H2O2]
Wi =
1 + K6[H2O2]
(7)
Here [“Bi”]0 is total concentration of the bismuth salt introduced into the solution and K6 is
equilibrium constant of transformation (6). Analysis of this equation demonstrates that at higher
concentrations of H2O2 the rate Wi should not depend on the concentration of H2O2. Such a
saturation was indeed observed experimentally (Figure 8).
Using eqs. (4) and (7) we obtain the expression for the alkane oxidation rate:
k2[RH]0
d[ROOH]
×
=
0
W0 =
dt
k2[RH]0 + k3[CH3CN]0
k1["Bi"]0[H2O2]
1 + K6[H2O2]
[H2O2]
1 + K6[H2O2]
×
×
Φ
=
(8)
where
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k1k2K6[RH]0["Bi"]0
=
Φ
k2[RH]0 + k3[CH3CN]0
(9)
Let us now transform eq. (8) into the form which is convenient for analysis of the experimental data:
[H2O2]
1
(1 + K6[H2O2])
=
Φ
W0
(10)
Figure 8, line 2 testifies that the experimental dependence of W0 on [H2O2] shown by Figure 8,
line 1 is described by eq. (10). In accordance with this equation we can calculate K6 = 1.3 M‒1 and
7.7 × 10‒7 s‒1. Inserting value
M we can calculate k1 = 9 × 10‒4 s‒1 for the conditions of experiments shown in Figure 8.
Φ
=
Φ
into eq. (9) using parameters [RH]0, [“Bi”]0 and k3[CH3CN]/k2 = 0.14
Theoretical analysis of the catalytic system under discussion (see below) demonstrates that
the most probable route of HO• radical generation is a monomolecular decay of a bismuth derivative
containing two peroxide fragments. Due to this, we need to assume that the species denoted as “Bi”
includes one peroxide molecule, and the following equilibrium exists:
K11
Bi + Н2О2
“Bi”
(11)
The constant of equilibrium (11) is so high that under the conditions of our experiment
shown in Figure 8 practically all bismuth is included into species “Bi”. We can estimate this
constant K11 if we accept the condition K11[H2O2] > 1. Taking into account that in the experiment
the minimum concentration of hydrogen peroxide is 0.1 M the value K11 should be higher than 10
M‒1 and the second equilibrium constant of the diperoxo derivative formation K6 = 1.3 M‒1 is
noticeably lower that the first one K11. This result obtained in the kinetic analysis of the
experimental data is in good qualitative agreement with the result of the theoretical consideration of
the system (see below).
Measuring the experimental effective activation energy for the catalytic system which
includes bismuth salt we took into account that the cyclooctane oxidation proceeds in some extent
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also in the absence of the metal. The lower the temperature, the larger the contribution of the
“blank” reaction Wblank to the total rate of the Biꢀcatalysed process Wtotal. Due to this, we determined
the rate of the real Biꢀcatalysed reaction Wcat as the difference between the reaction rates measured
in the presence and in the absence of bismuth salt: Wcat
°C the rate of nonꢀcatalysed reaction is comparable with a half of the total observed oxidation rate,
we analysed data obtained for temperature > 55 °C. The dependence of ln(Wcat) = ln(Wtotal Wblank
= Wtotal – Wblank. Since at temperature < 50
–
)
on 1/T is presented in Figure 9. Analysis of this dependence leads to a value of the activation
energy for the really Biꢀcatalysed reaction which is equal to 20 ± 2.5 kcal mol‒1.
Theoretical study of the reaction mechanism
The experimental kinetic data and reaction selectivity parameters discussed above indicate
that the Biꢀcatalysed oxidation of alkanes with H2O2 occurs via formation of highly reactive HO•
radicals. The HO• radical then abstracts hydrogen from the alkane molecule to give alkyl radical R•.
The latter reacts with molecular oxygen producing ROO• and then alkylhydroperoxide ROOH
(Scheme 2). The mechanism of the HO• radical formation from H2O2 (the rate limiting step of the
whole process) catalysed by transition metal complexes was initially proposed by Haber and
Weiss29 for the Fenton reagent, Fe(II) + H2O2. Since that time, it was revised in a number of
works30 and adapted later to the case of catalysts based on metals in the highest oxidation state. All
these mechanisms include oneꢀelectron oxidation/reduction of the metal atom. However, their
realization is more than questionable for the Bi(III) catalysts because the Bi(II) or Bi(IV) species
are unstable and their formation is highly unfavourable.
O2
O
R H
Cat
ROO•
C R''
HO•
R•
R'
+
ROOH
ROH
H2O2
H2O
Scheme 2 Radical mechanism of the Biꢀcatalysed alkanes oxidation with H2O2 (the rate limiting
step of the whole process is boxed).
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Recently, we applied theoretical methods for the elucidation of a plausible mechanism of the
HO• generation from H2O2 catalysed by complexes of metals bearing a single stable nonꢀzero
oxidation state (Al, Ga, In, Sc, Y, La, Be, Zn, Cd).3,19,20 This mechanism includes (
i) substitution of
a water molecule for H2O2 in the aqua complexes [M(H2O)n]k+ formed upon dissolution of
corresponding nitrate M(NO3)k in aqueous CH3CN, (ii) hydrolysis of the ligated hydrogen peroxide,
(iii) substitution of other H2O ligand for the second H2O2 molecule, (iv) elimination of a water
ligand (in the case of Zn and Cd) and (v) monomolecular homolytic HO–OH bond cleavage in the
key intermediate [M(H2O)(n–m)(H2O2)(OOH)](k–1)+ (Scheme 3).
H2O2
–H2O
H2O
H2O2
OH2
HOO
H2O
OH2
[M(H2O)(n–2)
H2O2
OH2
–H3O+
k+
(k–1)+
k+
[M(H2O)(n–2)
]
]
[M(H2O)(n–2)
]
H3O+
–(m–1)H2O
(m–2)H2O
•
HOO O2H2
[M(H2O)(n–m)
HOO
OH
OH
k
( –1)+
(k–1)+
]
(k–1)+
[M(H2O)(n–m)
]
–HOO•
[M(H2O)(n–m)
]
–HO•
M = Al, Ga, In, Sc, n = 6, m = 2, k = 3
M = Y, n = 8, m = 2, k = 3
M = La, n = 9, m = 2, k = 3
M = Be, n = 4, m = 2, k = 2
M = Zn, Cd, n = 6, m = 3, k = 2
Scheme 3 Mechanism of the HO• radical generation proposed for the systems [M(H2O)n]k+/H2O2
(M = Al, Ga, In, Sc, Y, La, Be, Zn, Cd).
This mechanism does not require a change of the metal oxidation state. Instead, the OOH–
coꢀligand in intermediate [M(H2O)n–m(H2O2)(OOH)](k–1)+ plays the same role as the transition metal
does in the classical Fenton or Fentonꢀlike processes, that is, the OOH– ligand is easily oxidized by
one electron upon the homolytic HO–OH bond cleavage leading to the stable complex [M(H2O)n–
m(OH)(•OOH)](k–1)+. Here, we used theoretical DFT methods to investigate if the same mechanism
may also be applied for the Bi catalyst, and results are discussed in the following sections.
(i) Initial catalytic form. Dissolution of bismuth nitrate in water in the presence of a strong
acid results in the formation of aquaꢀcomplexes [Bi(H2O)n]3+. The coordination number of Bi3+ in
such solutions is highly variable, the formation of octa (
n n
= 8)31 or/and ennea ( = 9)32 hydrates is
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usually proposed, while the structure [Bi(H2O)9](SO3CF3)3 was determined by Xꢀray analysis for
the triflate complex in solid state.33
In order to estimate the relative stability of the octaꢀ and enneahydrate forms of [Bi(H2O)n]3+
within the theoretical model used in this work, the calculations of complexes [Bi(H2O)9]3+ 19),
(
3+
[Bi(H2O)9](H2O)3+ (19-H2O), [Bi(H2O)8](H2O)3+ 18-H2O) and [Bi(H2O)8](H2O)2 18-2H2O) were
( (
carried out. Both octahydrates were found to be slightly more stable than corresponding
enneahydrates with the same number of H2O molecules (by 1.4–2.0 kcal/mol in terms of ꢀGs, Table
3) what is in qualitative agreement with previous ab initio theoretical data.34
Table 3 Calculated Gibbs free energies of activation (
ꢀ
Gs≠) and reaction (
ꢀ
Gs) in solution (in
kcal/mol)
Reaction
ꢀ
Gs≠
ꢀGs
[Bi(H2O)9]3+ (19) → [Bi(H2O)8](H2O)3+ (18-H2O)
[Bi(H2O)9](H2O)3+ (19-H2O) → [Bi(H2O)8](H2O)23+ (18-2H2O)
[Bi(H2O)8](H2O2)3+ (18-H2O2) → [Bi(H2O)7](H2O)(H2O2)3+ (27-H2O-H2O2)
[Bi(H2O)7](H2O)(H2O2)3+ (27-H2O-H2O2) → [Bi(H2O)7(H2O2)](H2O)3+ (38-H2O)
[Bi(H2O)7(H2O2)](H2O)3+ (38-H2O) → [Bi(H2O)5(OOH)](H2O)2+ (46-H2O) + H3O+ + H2O
[Bi(H2O)5(OOH)](H2O)2+ (46-H2O) + H2O2 →
–1.4
–2.0
1.2
5.8
4.3
via TS1
via TS2
–0.4
2.6
[Bi(H2O)4(OOH)](H2O)(H2O2)2+ (55-H2O-H2O2) + H2O
via TS3
via TS4
9.8
3.0
7.9
[Bi(H2O)4(OOH)](H2O)(H2O2)2+ (55-H2O-H2O2) →
[Bi(H2O)4(H2O2)(OOH)](H2O)2+ (66-H2O)
–1.5
[Bi(H2O)4(H2O2)(OOH)](H2O)2+ (66-H2O) → [Bi(H2O)4(OH)(OOH)](H2O)2+• (76-H2Oa) + HO•
[Bi(H2O)4(OH)(OOH)](H2O)2+• (76-H2Oa) → [Bi(H2O)4(OH)(OOH)](H2O)2+• (76-H2Ob)
[BiIII(H2O)7(H2O2)](H2O)3+ (38-H2O) → [BiIV(H2O)7(OH)](H2O)3+ + HO•
[BiIII(H2O)5(OOH)](H2O)2+ (46-H2O) → [BiII(H2O)5](H2O)2+ + HOO•
[BiIII(H2O)4(H2O2)(OOH)](H2O)2+ (66-H2O) → [BiII(H2O)4(H2O2)](H2O)2+ + HOO•
[Bi(H2O)5(OOH)](H2O)2+ (46-H2O) → [Bi(H2O)5(O)](H2O)2+ + HO•
[Bi(H2O)4(H2O2)(OOH)](H2O)2+ (66-H2O) → [Bi(H2O)4(H2O2)(O)](H2O)2+ + HO•
H2O2 → 2HO•
5.5
–1.8
39.8
37.5
35.9
38.5
37.7
39.4
24.1
[Bi(H2O)7(H2O2)](H2O)3+ (38-H2O) → [Bi(H2O)3(OH)3](H2O)2+ (9) + H3O+ + 2H2O
The calculated structures of complexes 19 and 18 correspond to the tricapped trigonal
prismatic and square antiprismatic polyhedra without noticeable distortions demonstrating no
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stereochemical influence of the Bi lone electron pair (Figure 10
i
and ii). The calculated mean Bi–O
bond length in 18-2H2O (2.48 Å) is in reasonable agreement with experimental EXAFS and LAXS
solution data for the octaꢀcoordinated Bi(III) ion (2.41(1) Å).31
iii
i
ii
v
iv
Fig. 10 Idealized polyhedra of some calculated structures.
(ii) First H2O-for-H2O2 substitution. At the first step of the investigated mechanism, one
of the water ligands in [Bi(H2O)n]3+ is substituted for the H2O2 molecule. The water exchange
process is known for aquaꢀcomplexes of the most of metals.35 Apparently, in the presence of a
significant amount of H2O2, the similar substitution of coordinated H2O for H2O2 should occur.
For this step, only a mechanism of the dissociative type D was found. It includes the
elimination of one water ligand from [Bi(H2O)8](H2O2)3+ 18-H2O2 to the second coordination
sphere and the formation of intermediate [Bi(H2O)7](H2O)(H2O2)3+ 27-H2O-H2O2) with a reduced
(
coordination number of Bi via TS1 (Scheme 4). This intermediate is then transformed into the
hydrogen peroxide adduct [Bi(H2O)7(H2O2)](H2O)3+ 38-H2O) via TS2 upon coordination of H2O2
(
to the metal. The coordination polyhedron of 27-H2O-H2O2 is capped trigonal prism (Figure 10iii),
whereas that of 38-H2O is distorted square antiprism (Figure 10ii). Both these complexes are only
slightly endoergonic relative to the starting 18-H2O2 (by 1.2 and 0.9 kcal/mol, respectively)
reflecting very flexible coordination sphere of Bi. The Bi–O2H2 bond in 38-H2O (2.606 Å) is longer
than the Bi–OH2 bonds by 0.11–0.15 Å demonstrating a weaker binding of hydrogen peroxide to
the metal in comparison with water. The energies of TS1 and TS2 are very similar, the former
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being less stable only by 0.3 kcal/mol (Figure 11). The activation barrier of the water substitution is
5.8 kcal/mol (Table 3). Such a low value indicates a high lability of the bismuth aquaꢀcomplexes
that correlates with the waterꢀexchange rate for Bi(III) of ~5
×
106 s–1 (Ref. 36) comparable with
In3+ (Ref. 37). Thus, the formation of the hydrogen peroxide adduct 38-H2O in solution is
favourable in terms of both kinetic and thermodynamic arguments.
H2O2
H2O
H2O2
[Bi(H2O)7]3+
[Bi(H2O)7]3+
[Bi(H2O)7]3+
=
=
27
–H2O
TS2
TS1
H2O2
H2O
[Bi(H2O)7]3+
[Bi(H2O)7]3+
18
38
Scheme 4 Mechanism of the H2OꢀforꢀH2O2 substitution in [Bi(H2O)8]3+ (secondꢀsphere solvent
molecules are omitted).
All attempts to located a nonaꢀcoordinated intermediate [Bi(H2O)8(H2O2)]3+ of the
associative mechanism A or transition state of the concerted (interchange) mechanisms I, Ia or Id
failed and, therefore, the realization of these mechanisms may be ruled out.
(iii) Hydrolysis of the hydrogen peroxide adduct. It is well known that Bi(III) in aqueous
solution is a strong acid with pKa value of 1.1 (Ref. 38) and it undergoes an intensive hydrolysis. At
the first step of the hydrolysis, the H+ abstraction from a protic ligand by an outerꢀsphere water
molecule occurs. Taking into account the more acidic nature of H2O2 compared to H2O, the former
ligand should preferentially undergo such a deprotonation in [Bi(H2O)7(H2O2)](H2O)3+ 38-H2O
leading to the formation of the hydroperoxo species [Bi(H2O)m(OOH)]2+ (Scheme 5). In neutral or
weakly acidic media, the hydrolysis goes further and leads to the formation of various mono and
polynuclear hydroxoꢀ and oxoꢀcomplexes many of them being insoluble.39 However, in strongly
acidic media the hydrolysis is expected to be mostly restricted by its first step, i.e. deprotonation of
H2O2 in 38-H2O
.
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H2O2
[Bi(H2O)7]3+
HOO
H2O
HOO
–H3O+
–2H2O
[Bi(H2O)7]2+
[Bi(H2O)5]2+
38
46
No minimum
Scheme 5 Hydrolysis of coordinated hydrogen peroxide in [Bi(H2O)7(H2O2)]3+ (secondꢀsphere
solvent molecules are omitted).
ꢀG , kcal/mol
s
[Bi(H2O)4(OH)(OOH)]2+
20
76-H2Oa
+ HO•
16
12
8
TS4
14.4
15.2
TS3
13.4
13.2
[Bi(H2O)4(OH)(OOH)]2+
11.3
76-H2Ob
9.8
[Bi(H2O)4(OOH)]2+
55-H2O-H2O2
[Bi(H2O)4(H2O2)(OOH)]2+
66-H2O
TS1
TS2
5.8
5.5
3.4
4
[Bi(H2O)5(OOH)]2+
0.9
1.2
46-H2O
[Bi(H2O)7]3+
0
[Bi(H2O)7(H2O2)]3+
[Bi(H2O)8]3+
27-H2O-H2O2
38-H2O
18-H2O2
•
Water substitution
Hydrolysis
Water substitution
HO radical
generation
Fig. 11 General energy profile for the proposed mechanism of the HO• generation (only metalꢀ
containing species and the HO• radical are indicated, secondꢀsphere solvent molecules are omitted,
numbers indicate the relative energies).
The
geometry
optimization
of
the
octaꢀcoordinated
hydrolytic
product
[Bi(H2O)7(OOH)](H2O)2+ was undertaken. However, all attempts resulted in an extrusion of two
water molecules from the first coordination sphere. As a result, the equilibrium structure of the
hexaꢀcoordinated complex [Bi(H2O)5(OOH)](H2O)2+ 46-H2O) was obtained. The coordination
(
polyhedron of 46-H2O is a highly distorted pentagonal pyramid with the OOH– ligand being in the
axial position (Figure 10iv). Such a structure reflects a wellꢀknown effect of the stereochemically
active lone electron pair of Bi, and the similar structural motif was found in solid state for various
Bi complexes.40
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The CPCM method and the model used in this work do not allow accurate estimates of the
ꢀ
Gs value of hydrolysis of 38.19 Hence, this value was calculated using the experimental pKa value
of Bi(III) (1.1 [Ref. 38]) and the corresponding thermodynamic cycle (see ESI for details). Thus
Gs of hydrolysis of 38 is only 2.6 kcal/mol demonstrating that the H2O2 ligand in 38-
calculated
ꢀ
H2O is highly acidic. At the same time, it is curious that the ꢀGs values of H2O2 hydrolysis in the
aquaꢀcomplexes [M(H2O)5(H2O2)]3+ (M = Al, Ga, In) are even lower than that for Bi (0.6–2.2
kcal/mol),19,20 that is, the H2O2 ligand in the complexes of Al, Ga and In is even more acidic.
(iv) Second H2O-for-H2O2 substitution. At the next step of the proposed mechanism, the
substitution of another H2O ligand for H2O2 occurs. For this step, the dissociative mechanism D is
also realized and it includes dissociation of a water ligand from [Bi(H2O)5(OOH)](H2O2)2+ 46-
H2O2, via TS3, to give complex [Bi(H2O)4(OOH)](H2O)(H2O2)2+
coordination of H2O2 to Bi and formation of [Bi(H2O)4(H2O2)(OOH)](H2O)2+
(Scheme 6). The coordination polyhedron of 66-H2O is also a distorted pentagonal pyramid, while
that of 55-H2O-H2O2 is a highly distorted square pyramid (Figure 10 ). In both cases, the OOH–
(
55-H2O-H2O2) followed by the
(
66-H2O) via TS4
v
ligand occupies the axial position. The square pyramidal structural motif is also known for Bi(III)
species.40a,41 Both intermediate 55-H2O-H2O2 and complex 66-H2O are noticeably endoergonic
relative to 46-H2O (by 7.9 and 6.4 kcal/mol, respectively, Table 3). The activation energies of TS3
and TS4 are similar, the latter having a slightly higher energy. The activation barrier of the second
water substitution step is 10.9 kcal/mol and this is noticeably higher than that of the first
substitution (5.8 kcal/mol, Table 3).
OH2
O2H2
HOO
[Bi(H2O)4]2+
HOO
HOO
H O
2
2
=
[Bi(H2O)4]2+
=
[Bi(H2O)4]2+
5
–H O
2
5
TS3
TS4
HOO
OH2
HOO
O2H2
[Bi(H2O)4]2+
[Bi(H2O)4]2+
6
6
6
4
Scheme 6 Mechanism of the second H2OꢀforꢀH2O2 substitution (secondꢀsphere solvent molecules
are omitted).
26

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DOI: 10.1039/C4CY01651C
(v) Formation of HO• radicals. Theoretical studies of the discussed mechanism for the
catalytic systems based on the aquaꢀcomplexes of the Group III metals [M(H2O)n]3+ (M = Al, Ga,
In, Sc, Y, La) indicated that the hydrogen peroxide ligand in intermediates [M(H2O)n–
2(H2O2)(OOH)]2+ is highly activated towards the homolytic HO–OH bond cleavage.19,20 Careful
analysis and scan of the potential energy surface revealed that the HO• generation upon homolytic
HO–OH bond rupture in [Bi(H2O)4(H2O2)(OOH)](H2O)2+ 66-H2O results in the formation of
complex [Bi(H2O)4(OH)(OOH)](H2O)2+• 76-H2Oa) (Scheme 7) and this process requires only 5.5
(
kcal/mol (vs. 39.4 kcal/mol in free H2O2, Table 3). This value is quite comparable with the
calculated adiabatic HO–OH bond energies in [M(H2O)4(H2O2)(OOH)](H2O)2+ for M = Al, Ga, In
and Sc (4.4–7.5 kcal/mol).19,20 Thus, the ability of Bi(III) to activate hydrogen peroxide towards the
HO• formation is similar to that of the nonꢀtransition metals of the Group III. The formed complex
[Bi(H2O)4(OH)(OOH)](H2O)2+•
(
76-H2Oa) is then transformed to another isomer 76-H2Ob with the
energy lower by 1.8 kcal/mol than 76-H2Oa
.
The principal reasons of such tremendous activation of H2O2 in [M(H2O)4(H2O2)(OOH)]2+
i
) ability of the OOH– ligand in the product [M(H2O)4(OH)(OOH)]2+• to be easily oxidized to
are (
HOO• (Scheme 7) and (ii) delocalization of the spin electron density among two oxygen atoms of
the OOH ligand in the product [M(H2O)4(OH)(OOH)]2+• (Figure 12).19,20 Note that the oxidation
state of Bi(III) is not changed upon the reaction 66-H2O
→
76-H2O. Instead, the oxidation of the
OOH– coꢀligand stabilizes the product of the HO–OH bond cleavage 76-H2O and, hence, activates
H2O2 towards its decomposition.
27

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Page 28 of 38
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DOI: 10.1039/C4CY01651C
H
–I
•
0
••
H
•
•
–II
•
•
•
•
•
•
•
• •O• •O
•
•O
•
•
•
[Bi(H2O)4]2+
H
H –I
H
•
•
O
•
•
–I
••
–I
•
–I
•
•
•
•
•
•
•
•
•
•
• •
•
• •
•O
O•
•O
•
•
•
H
–I
•
[Bi(H2O)4]2+
+
•
•
•
O
•
•
66
•
H
H
0
–I
••
•
•
–II
•
•
•
•
•
•
•
• •O•• •O
•
O
•
•
•
•
[Bi(H2O)4]2+
76
Scheme 7 HO–OH bond cleavage in [Bi(H2O)4(H2O2)(OOH)]2+ (valence electrons and oxidation
states of the oxygen atoms are shown).
Fig. 12 Spin density distribution in 76-H2O
.
It is important that the
ꢀ
Gs energies required for the formation of the HO• or HOO• radicals
accompanied by the oneꢀelectron oxidation or reduction of Bi(III) are very high being 39.8, 37.5
and 35.9 kcal/mol for the reactions 12, 13 and 14, respectively (Table 3), and, therefore, these
routes may be ruled out.
[BiIII(H2O)7(H2O2)](H2O)3+ (38-H2O
[BiIII(H2O)5(OOH)](H2O)2+ (46-H2O
)
→
→
[BiIV(H2O)7(OH)](H2O)3+ + HO•
[BiII(H2O)5](H2O)2+ + HOO•
(12)
(13)
(14)
)
[BiIII(H2O)4(H2O2)(OOH)](H2O)2+ (66-H2O [BiII(H2O)4(H2O2)](H2O)2+ + HOO•
) →
Additionally, the O–O bond cleavage in the OOH– ligand is also not favourable; the
values of reactions 15 and 16 being 38.5 and 37.7 kcal/mol, respectively.
ꢀGs
[Bi(H2O)5(OOH)](H2O)2+ (46-H2O
)
→
[Bi(H2O)5(O)](H2O)2+ + HO•
[Bi(H2O)4(H2O2)(O)](H2O)2+ + HO•
(15)
(16)
[Bi(H2O)4(H2O2)(OOH)](H2O)2+ (66-H2O
)
→
28

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Catalysis Science & Technology
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DOI: 10.1039/C4CY01651C
(vi) Closure of the catalytic cycle. At the final steps of the catalytic cycle, the HOO• radical
may be formed from complex 76-H2Ob to give [Bi(H2O)4(OH)](H2O)2+ 85-H2O) (Scheme 8). The
(
initial catalytic form is then restored as a result of the protonation of 85-H2O and subsequent water
addition.
H3O+
2H2O
•
OH
[Bi(H2O)4]2+
H2O
[Bi(H2O)7]3+
HOO
OH
[Bi(H2O)4]2+
–HOO•
76
18
85
Scheme 8 Closure of the catalytic cycle.
(vii) Energy profile and activation energy. The energy profile for the whole process of the
HO• radical formation is shown in Figure 11. Its inspection demonstrates that the rate limiting step
of the HO• generation in the system [Bi(H2O)8]3+/H2O2 is the monomolecular HO–OH bond
cleavage in the key intermediate [Bi(H2O)4(H2O2)(OOH)](H2O)2+ 66-H2O. The overall Gibbs free
energy of activation relative to the initial level 18-H2O2 is 15.2 kcal/mol. The estimate of the overall
activation enthalpy
of one of the reaction steps, i.e. hydrolysis, was calculated on the basis of the experimental pKa
value which is related to G rather than to H. However, assuming that the system entropy is not
ꢀ ꢀ
Hs≠ is not that straightforward as the calculations of Gs≠. Indeed, the energy
ꢀ
ꢀ
changed during the hydrolysis step, the overall activation enthalpy of the HO• formation should be
22.0 kcal/mol. This value correlates well with the experimentally determined effective activation
energy of 20
± 2.5 kcal/mol (see above).
(viii) Non-radical pathway. Additionally, the possibility of a nonꢀradical mechanism
involving the formation of Bi(V) species was also considered (the pathway similar to that proposed
by Bray and Gorin for the Fenton system)42 (Scheme 9). However, all attempts to calculate the
Bi(V) oxoꢀaquaꢀcomplexes [Bi(H2O)n(=O)]3+
(n = 7, 5) failed. The geometry optimization always
led to the hydroperoxo species indicating that the Bi(V) oxoꢀaquaꢀcomplexes are not stable and the
coordinated water ligand should be spontaneously oxidized by Bi(V) to the peroxo derivatives. The
equilibrium structure of the trihydroxo Bi(V) complex [Bi(H2O)3(OH)3](H2O)2+
29
(9) was