Experimental and Theoretical Insight into the Kinetics and Mechanism of the Synthesis Reaction of 2,3-Dihydro-2-phenylquinazolin-4(1H)-one Catalyzed in Formic Acid

Article abstract of DOI:10.1002/kin.21064

An efficient and simple method has been developed for the synthesis of 2,3-dihydro-2-phenylquinazolin-4(1H)-one catalyzed in formic acid. Also, the synthesis reaction between benzaldehyde and 2-aminobenzamid was monitored spectrally. On the basis of the kinetic data obtained by the UV–vis spectrophotometry, both the first and second steps of the speculative five steps mechanism were enabled to be a rate-determining step and also reaction showed second-order kinetics. Considering information obtained by the stopped-flow technique indicated that the first step is certainly a fast step. Moreover, the reaction was energetically and thermodynamically evaluated using theoretical methods and results were profoundly compared with the experimental approaches. Herein, theoretical rate constants were obtained using potential energy surfaces and the transition state theory at the B3LYP/6–311+G** level of theory. The Winger method was also applied to describe the tunneling effects. Calculations showed that the second step was the rate-determining step in accordance with the experimental data. It is also found that the oxidation step was the fastest step in the proposed mechanism. For all five steps, two possibilities were considered for generating the probable product by using the thermodynamic parameters and kinetic data. Thermodynamic parameters also showed an exothermic reaction.

Full text of DOI:10.1002/kin.21064

Experimental and
Theoretical Insight into the
Kinetics and Mechanism of
the Synthesis Reaction of
2,3-Dihydro-2-
phenylquinazolin-4(1H)-one
Catalyzed in Formic Acid
SAYYED MOSTAFA HABIBI-KHORASSANI, MEHDI SHAHRAKI, ALI EBRAHIMI,
MAHDIEH DARIJANI
Department of Chemistry, Faculty of Science, University of Sistan and Baluchestan, P. O. Box 98135-674, Zahedan, Iran
Received 12 February 2016; revised 27 October 2016; accepted 4 December 2016
DOI 10.1002/kin.21064
Published online in Wiley Online Library (wileyonlinelibrary.com).
ABSTRACT: An efficient and simple method has been developed for the synthesis of 2,3-
dihydro-2-phenylquinazolin-4(1H)-one catalyzed in formic acid. Also, the synthesis reaction
between benzaldehyde and 2-aminobenzamid was monitored spectrally. On the basis of the
kinetic data obtained by the UV–vis spectrophotometry, both the first and second steps of
the speculative five steps mechanism were enabled to be a rate-determining step and also
reaction showed second-order kinetics. Considering information obtained by the stopped-flow
technique indicated that the first step is certainly a fast step. Moreover, the reaction was
energetically and thermodynamically evaluated using theoretical methods and results were
profoundly compared with the experimental approaches. Herein, theoretical rate constants
were obtained using potential energy surfaces and the transition state theory at the B3LYP/6–
311+G** level of theory. The Winger method was also applied to describe the tunneling effects.
Calculations showed that the second step was the rate-determining step in accordance with the
experimental data. It is also found that the oxidation step was the fastest step in the proposed
Correspondence to: S. M. Habibi-Khorassani; e-mail: smhabibi
@chem.usb.ac.ir.
Contract grant sponsor: Research Council of the University of
SistꢀaCn and Baluchestan.
2017 Wiley Periodicals, Inc.

2
HABIBI-KHORASSANI ET AL.
mechanism. For all five steps, two possibilities were considered for generating the probable
product by using the thermodynamic parameters and kinetic data. Thermodynamic parameters
ꢀC
also showed an exothermic reaction.
2017
2017 Wiley Periodicals, Inc. Int J Chem Kinet 1–16,
Two mechanistic pathways can be proposed to ex-
plain the formation of the products for the reaction
between benzaldehyde and 2-aminobenzamide in the
presence of formic acid as a catalyst. Herein, all steps of
the two proposed mechanisms are appraised thermody-
namically and kinetically along with the experimental
evidences obtained by the stopped-flow technique and
UV–vis spectrophotometry to confirm or deny these
possibilities.
INTRODUCTION
2,3-Dihydroquinazolinone derivatives are an impor-
tant class of heterocyclic compounds and have a broad
range of pharmacological properties, for example, an-
ticancer, antitumor, antibiotic, antidefibrillatory, an-
tipyretic, analgesic, antihypertonic, diuretic activities
[1–4] and herbicidal agents as well as plant growth
regulators [5]. The unique biological activities have
made quinazolinone synthetic studies over the years
very attractive [6]. They are also important building
blocks in total synthesis of natural products [7] and
are the constituents of some isolated naturally occur-
ring alkaloids [8]. A wide variety of organic reac-
tions are catalyzed by different acids such as Brønsted
acid and Lewis acid. Acids are employed as cata-
lysts for hydrolysis or formation of esters, acetals, etc.
and carbon–carbon bond-forming reactions. Some chi-
ral acids are a new class of catalysts with properties
of enantioselective. The mechanism of the acid cata-
lyst is the activation of the C═X bond (X = O, NR,
CR2) and helping nucleophilic addition to the C═X
bond [9–17].
In spite of various synthesis methods and this fact
that these compounds are significant target for or-
ganic and medicinal chemists, however, there are not
comprehensive studies regarding kinetics and mech-
anism of the relevant reactions. In this work for the
first time, we describe experimental and theoretical
kinetic results together with the detailed mechanis-
tic studies for 2,3-dihydro-2 phenylquinazolin-4(1H)-
one catalyzed in formic acid based on a global ki-
netic analysis methodology. The stopped-flow tech-
nique is most commonly used when fast reactions are
involved [18,19], and in the present work there was a
significant proof for determining the fast step of the
reaction. Also, the kinetic data regarding the overall
reaction was gained using the UV–vis spectrophotom-
etry technique; this technique is generally preferred
over IR and NMR spectroscopy because (a) it is more
sensitive so more dilute solutions may be used and
(b) water and other common solvents present fewer
problems with UV–vis than with IR and NMR meth-
ods [20,21]. Moreover, theoretical methods are used
for obtaining thermodynamic and kinetic data that are
important for the accepting prophecy of the chemical
reaction mechanism and also a possible interpretation
of experimental findings [19,22].
EXPERIMENTAL
Chemicals and Instruments
The chemicals used are all analytical grade, obtained
from Fluka (Buchs, Switzerland). Extra pure formic
acid (supplied by Merck, Darmstadt, Germany) was
used throughout the study. Melting points and IR spec-
tra were measured on an Electrothermal 9100 appara-
tus and a JASCO FT-IR-460 plus spectrometer, respec-
tively. The 1H NMR spectra were recorded on Bruker
DRX-400 Advance instruments with DMSO and ace-
tone as solvents. The stopped-flow apparatus (model
SX-18MV, Surrey, UK) and the associated computer
system are obtained from Applied Photophysics (UK).
For a total of 100 μL/shot into a flow cell with 10
mm light path, the fastest time for mixing solutions
and recording the first data point is about 20 s. Ki-
netic traces are taken at different wavelengths between
400 and 300 nm, and the data are analyzed with the
SX-18MV v4.47 operating, which allows the synthe-
sis of artificial sets of time-dependent spectra as well as
spectral analysis of intermediates. In a typical single-
mixing experiment, one syringe contained 1.5 × 10−3
M 2-aminobenzamide, and the other contained 1.5 ×
10−3 M benzaldehyde in formic acid as a catalyst and
solvent. For each run, equal volumes of solutions were
mixed in the mixing chamber, and the changes in the
absorbance were monitored during a chosen time. Rate
constants are presented as an average of several kinetic
runs (at least 6–10) and are reproducible within 3%.
The overall rate of reaction is followed by monitor-
ing absorbance changes of the products with time on a
Varian (model Cary Bio-300, CA, USA) UV–vis spec-
trophotometer with a 10-mm light-path cell. In two
types of instruments, the temperature of the reaction
was maintained to within 0.1°C at various tempera-
tures by circulating water bath.
International Journal of Chemical Kinetics DOI 10.1002/kin.21064

EXPERIMENTAL AND THEORETICAL INSIGHT INTO THE KINETICS
3
Figure 1 Reaction between benzaldehyde 1 and 2-aminobenzamide 2 in the presence of formic acid as a catalyst and solvent
for generation of 2,3-dihydro-2-phenylquinazolin-4(1H)-one 3.
General Procedure for the Synthesize of 2,
3-Dihydro-2-phenylquiazolin-4(1H)-One in
Formic Acid
The mixture of benzaldehyde 1 (1 mmol), 2-
aminobenzamide 2 (1 mmol), and formic acid (1 mL)
were stirred at various temperatures (Fig. 1). The
progress of the reaction was monitored by TLC (ethyl
acetate: petroleum ether 3:1). The reaction mixture
was poured into 10 mL ice water; on solidification,
it was filtered, washed with ice water, and recrys-
tallized from ethanol to give the pure 2,3-dihydro-2-
phenylquinazolin-4(1H)-one. At 30˚C, the product was
obtained in 98% yield.
Figure 2 The UV spectra of the reaction between benzalde-
hyde 1 (1.5 × 10−3) and 2-aminobenzamide 2 (1.5 × 10−3
)
in the presence of formic acid over the wavelength range
of 295–450 nm at 25°C; the upward direction of the arrow
indicates the progress of product versus time. [Color figure
can be viewed at wileyonlinelibrary.com]
EXPERIMENTAL RESULTS AND
DISCUSSION
Kinetics
For a kinetic study, it is necessary to find a suit-
able wavelength that the reactants have relatively no
absorbance value. For this reason in the first exper-
iment, 1.5 × 10−3 M solution of each reactant 1
and 2 and product 3 were prepared in formic acid.
The relevant spectrum of each reactant and product
were recorded over the wavelength range 200–800 nm.
In the second experiment, the reaction mixture was
started into a 10-mm black quartz spectrophotome-
ter cell along with a 1.5 × 10−3 M solution of each
reactant according to the stoichiometry of each reac-
tant in the overall reaction (Fig. 1). The reaction was
monitored by conducting scans of the entire spectrum
with 3 min intervals during the whole reaction time
at ambient temperature (Fig. 2). From these experi-
ments, the appropriate wavelengths were found out
to be in the range of 350–390 nm. In the third ex-
periment under same concentration of each compound
(1.5 × 10−3 M), the experimental absorbance curve
was recorded versus time at 25°C temperature and
wavelength 360 nm. This is shown in Fig. 3 (dotted
line), which exactly fits to the second-order curve (solid
line).
Figure 3 Second-order fit curve (solid line) accompanied
by the original experimental curve (dotted line) for the reac-
tion between 1(1.5 × 10−3) and 2 (1.5 × 10−3) in formic
acid at 360 nm and 25°C. [Color figure can be viewed at
wileyonlinelibrary.com]
In this case, the overall order of the rate law can be
written as
Rate = kove[2]β, Rate = kove[1][2], α + β = 2
(1)
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4
HABIBI-KHORASSANI ET AL.
Table I Values of the Overall Rate Constants for the
Reaction between 1 and 2 in the Presence of Different
Solvents at All Temperatures Investigated
kaove(M–1 min–1
)
Wavelength
(nm)
ε (D)
20°C
25°C
30°C
35°C
Solvent: mix formic acid/acetic acid (1:1)
360
Solvent: formic acid
360 58
32.1
35.3
50.4
84.3
130.5
330.3
87.8
162.3
228.4
Figure 4 The dependence of the second-order rate constant
(ln k1) on reciprocal temperature for the reaction between
compounds 1 and 2 measured in formic acid at wavelength
of 360 nm conforming to the Arrhenius equation. [Color
figure can be viewed at wileyonlinelibrary.com]
aOverall rate constant.
Herein, in all the experiments, the UV–vis spectrum of
product 3 was measured over the concentration range
(1.5 × 10−3 M ࣘ [3] M ࣘ1.5 × 10−2 M) to check
a linear relationship between absorbance values and
concentrations.
Effect of Solvent and Temperature
To determine the effect of solvent and temperature
environment on the rate of reaction, it was needed
to perform various experiments at different solvent
polarities and temperatured but otherwise under the
same concentration as the third experiment. For this
cause, formic acid with a dielectric constant (58 D)
was chosen as a suitable solvent since it not only
dissolved all compounds but also did not react with
them.
The effects of solvent and temperature on the rate
constant are given in Table I. The results show that
the rate of reaction in each solvent was increased
at higher temperature; also, the rate of reaction be-
tween 1 and 2 speeds up in a higher dielectric con-
stant environment (formic acid) at all temperatures
investigated.
Figure 5 Eyring plots according to Eqs. (2) and (3), for
the reaction between 1 and 2 in formic acid as a solvent and
catalyst.
In the temperature range studied, the dependence
of the second-order rate constant (ln kove = ln k1)
of the reaction on reciprocal temperature was con-
sistent with the Arrhenius equation, giving activa-
tion energy (64.43 k J mol−1) from the slope of
Fig. 4.
used):
ꢀ
ꢁ
ꢀ
ꢁ
k1
T
kB
h
ꢀS‡
R
ꢀH‡
RT
Ln
=
Ln
+
−
ꢂ
ꢀH‡ = 50.4277 0.0015
ꢀS‡ = 4.2817 0.0048
The activation parameters that involve ꢀG‡, ꢀS‡,
and ꢀH‡ can be now calculated for the first step (rate-
determining step (RDS), k1), as an elementary reac-
tion, on the basis of the Eyring equation (2) (Fig. 5A,
ln(k1)/T vs. 1/T) and also a different linearized form of
the Eyring equation (3) (Fig. 5B, T × ln(k1)/T against
(T)) [23]. The results are collected in Figs. 5A and
5B. Statistical analysis of the Eyring equation clearly
confirms that the standard errors of ꢀS‡ and ꢀH‡ cor-
relate (where Tav is the center of the temperature range
k J
mol
ꢂ
J
(2)
(3)
mol
ꢀ
ꢁ
ꢀ
ꢁ
k1
T
kB
h
ꢀS‡
ꢀH‡
R
T ×
= T × Ln
+
−
R
ꢂ
ꢀH‡ = 50.2839 0.0015
ꢀS‡ = 3.5168 0.0048
k J
mol
ꢂ
J
mol
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EXPERIMENTAL AND THEORETICAL INSIGHT INTO THE KINETICS
5
Table II Effect of Various Catalysts on a Reaction
between 1 (1.5 × 10−3 M) and 2 (1.5 × 10−3 M)
Compounds in the Presence of Formic Acid (as a
Solvent and Catalyst) at Wavelength of 360 nm
kove (M−1min−1
)
SD
Catalyst
T = 25°C
162.3
Formic acid
0.0076
0.0035
Trichloroacetic acid
257.3
Figure 6 Pseudo–first-order fit curve (solid line) accom-
panied by the original experimental curve (dotted line) with
respect 2-aminobenzamide 2 (1.5 × 10−3), for the reaction
between 1 (1.5 × 10−2) and 2 (1.5 × 10−3) which pro-
ceeded in formic acid at 25°C and 370 nm. [Color figure can
be viewed at wileyonlinelibrary.com]
σ(ꢀS‡) = 1/Tavσ(ꢀH ‡)
(4)
It follows that in most solution-phase studies σ(ꢀS‡)
ࣈ σ(ꢀH‡) × 0.0033 K−1. This correlation has been
mentioned elsewhere [23,24]. The standard errors for
activation parameters have been calculated according
to above instructions [23–25], and they have been re-
ported along with this parameters in Figs. 5A and
5B. With respect to the values of ꢀS‡ and ꢀH‡,
ꢀG‡ = 60.88 kJ mol−1 is reported for the reac-
tions between 1 and 2 in formic acid at 360 nm and
289.2 K.
Effect of Catalyst
The rate of reaction was increased in the presence of
trichloroacetic acid as a second catalyst in comparison
with the first catalyst (formic acid) (see Table II).
Figure 7 Pseudo–first-order fit curve (solid line) accom-
panied by the original experimental curve (dotted line) with
respect to benzaldehyde 1 (1.5 × 10−3) for the reaction be-
tween 1 (1.5 × 10−3) and 2 (1.5 × 10−2) which proceeded in
formic acid at 25°C and 370 nm. [Color figure can be viewed
at wileyonlinelibrary.com]
Effect of Concentration
Also to find the partial order of reaction in relation
to compounds 1 and 2 under the pseudo-order con-
dition, two separate experiments (fourth and fifth ex-
periments) were performed at 25°C and wavelength of
370 nm. In the fourth experiment, the reaction mixture
was started into a 10-mm quartz spectrophotometer
cell with concentration [(1.5 × 10−2 M compound 1)]
and [(1.5 × 10−3 M compound 2)]. The experimental
absorbance curve versus time along with a first-order
fitting was recorded (Fig. 6). For this case, the rate law
can be expressed as Eq. (5):
Then in the fifth experiment [(1.5 × 10−3 M com-
pound 1)] and [(1.5 × 10−2M compound 2), the ex-
perimental absorbance curve versus times along with
a first-order fit was recorded at 25°C and wavelength
of 370 nm (Fig. 7). For this case, the rate law can be
expressed as Eq. (7):
Rate = kove[1]α[2]β Rate = kobs[1]α
(7)
Rate = kove[1]α[2]β Rate = kobs[2]β
(5)
kobs
[2]β
0.2737
kobs = kove[2]β kove
=
kove
=
= 18.25
1.5×10−2
kobs
[1]α
0.2770
kobs = kove[1]α kove
=
kove
=
= 18.46
1.5×10−2
(8)
(6)
In fact, in the fifth experiment, the obtained kove (18.25)
from Eq. (8) is nearly equal to kove (18.46) of the fourth
experiment (Eq. (6)). Herein, in the fifth experiment the
So the order of reaction in relation to 2-
aminobenzamide 2 is one (β = 1).
International Journal of Chemical Kinetics DOI 10.1002/kin.21064

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HABIBI-KHORASSANI ET AL.
Figure 8 The proposed mechanism involves, step1, step2, step3, step4, and step5; the reaction between benzaldehyde 1
and 2-aminobenzamide 2 in the presence of formic acid for the generation of product 3. [Color figure can be viewed at
wileyonlinelibrary.com]
absorbance versus time provided a pseudo–first-order
curve. This is possible when α is one in Eq. (7). As
a result, since β = 1 (as determined previously), it is
reasonable to accept that the reaction is first order with
respect to compounds 2 (β = 1) and 1 (α = 1). This is
compatible with the overall order of reaction (α + β
= 2) in the third experiment.
Utilizing the above results, a simplified scheme of
the proposed reaction mechanism (Fig. 8) is shown in
Fig. 9.
International Journal of Chemical Kinetics DOI 10.1002/kin.21064

EXPERIMENTAL AND THEORETICAL INSIGHT INTO THE KINETICS
7
For obtaining the concentration of intermediate [I2],
the following equations yielded by applying the steady-
state assumption:
d[I2]
= k2[I1] − k3[I2] = 0
(16)
dt
k2[I1] − k3[I2]
(17)
And with the placement of Eq. (17) in (15), the follow-
ing equation is obtained:
Rate = k2[I1]
(18)
Figure 9 The simplified scheme for the proposed reaction
between benzaldehyde 1 and 2-aminobenzamide 2 in the
presence of formic acid as a solvent and catalyst.
[I1] is generated in accordance with the steady-state
assumption:
d[I1]
= k1[1][2] − k−1[I1] − k2[I1]
(19)
(20)
d t
To investigate which steps of the proposed mecha-
nism could be the RDS, the rate law is written using
the final step for the product:
k1[1][2]
[I1] =
k
−1 + k2
Rate = k5[I4][o]
(9)
The value of Eq. (20) can be put into Eq. (19):
The steady-state assumption can be employed for ob-
taining the concentration of [I4], which is generated
from the following equations:
k1[1][2]
Rate =
(21)
k
−1 + k2
Equation (21) is not involved k3, k4, and k5; hence steps
3, 4, and 5 not enable to be a RDS. Nevertheless, the
rate law contains k1 and k2; therefore, there is two pos-
sibilities for the RDS. If k2 is a RDS, this speculation
k–1 ꢀ k2 is logical, then the rate law can be stated as
d[I4]
= k4[I3] − k5[I4][o] = 0
(10)
dt
k4[I3] = k5[I4][o]
(11)
k1k2
k1k2
The value of Eq. (11) can be replaced in Eq. (9), so the
rate equation becomes
Rate =
[1][2], kove
=
⇒ Rate = kove[1][2]
k
k
−1
−1
(22)
Rate = k4[I3]
(12)
This equation is compatible with the second-order ex-
For obtaining the concentration of intermediate [I3],
the following equations yielded by applying the steady-
state assumption:
perimental rate law (Eq. (1)).
Other possibility can be employed for k1 (step1)
as a RDS; in this case, it is reasonable to accept this
assumption k–1 ꢁ k2. Under this condition, the rate
law can be written as
d[I3]
= k3[I2] − k4[I3]
(13)
dt
Rate = k1[1][2][cat], kove = k1[cat], Rate = kove[1][2]
(23)
k3[I2] = k4[I3]
(14)
Both Eqs. (22) and (23) are consistent with the second-
order rate experimental rate law (Eq. (1)}. In fact, there
is two possibilities (k1 or k2) for the RDS. It seems that
further study is necessary for a right decision. Next
sections discuss regarding this difficulty.
The value of Eq. (14) can be replaced in Eq. (12), so
the rate equation becomes
Rate = k3[I2]
(15)
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HABIBI-KHORASSANI ET AL.
Figure 10 The UV absorbance changes with time that oc-
curred upon mixing reagents in the chamber of the stopped-
flow instrument at various wavelengths. [Color figure can be
viewed at wileyonlinelibrary.com]
Figure 11 Second-order fit curve for step1 (k1) together
with the original experimental curve at 370 and 25°C in
formic acid by the stopped-flow technique.
nates (IRC) [31] were performed to verify that the tran-
sition states connected to the expected reactants and
products for each step. For proton transfer reactions,
the tunneling effect is expected to play a significant
role when the imaginary frequency is large [32]. So we
use simple methods, including the Wigner method for
calculation of tunneling correction [33]. In this work,
we carried out a detailed theoretical study of the reac-
tion shown in Figure 8 to gain insight into the reaction
mechanism. All calculations were performed with the
GAMESS program [34].
Stopped-Flow Study
The experimental data obtained from the previous sec-
tion by the UV-spectrophotometry technique indicated
that the both step1 (k1) and step2 (k2) have a possibil-
ity to be a RDS. Therefore, using of the stopped-flow
technique can be useful for determining the fast step.
Figure 10 shows the UV spectral changes of the re-
action versus time recorded upon mixing of three
reactants in formic acid and in the chamber of the
stopped-flow instrument. As can be seen in Fig. 10, the
absorbance changes appeared instantly within 20 s;
then their intensities increased slowly more than 200 s
at the various wavelengths. So, it can be deduced that
the reaction includes an initial fast step (step1) and con-
sequently a slow step (step2) as a RDS. Step1, which
is associated with the increase in the absorbance, was
analyzed on a fast timescale with a double exponential
function at 370 nm (Fig. 11).
Theoretical Results and Discussion
The reaction between benzaldehyde 1 and 2-
aminobenzamide 2 in the presence of formic acid for
generating product 3 is shown in Figure 8. The pro-
posed mechanism consists of five steps that are shown
in Fig. 8. The role of formic acid is as a catalyst and
solvent together in the reaction. At the beginning of
reaction, carbonyl of benzaldehyde is activated with
formic acid (role of catalyst, in the first step). Each
step is discussed in the following sections.
COMPUTATIONAL SECTION
Methods
Step1. In the first step of the reaction, N6 and C2 atoms
draw close and lead to the C2─N6 bond formation at
intermediate I1. The C2─N6 bond length of the first
The hybrid density functional B3LYP method, which
combines the three parameter Becke exchange func-
tional, B3, with the Lee–Yang–Parr nonlocal correla-
tion functional (LYP) is used for mechanical calcu-
lation [26,27]. The geometries of reactants, products,
intermediates, and transition states (TS) involved in the
reaction were optimized at the B3LYP/6–311+G(d,p)
level [28]. Frequency calculations were performed with
the same level of theory to determine the nature of the
stationary points. Reactants and intermediates were de-
termined with no imaginary frequency and the transi-
tion state with one imaginary frequency. A scaling fac-
tor of 0.9668 for zero-point vibrational energies was
used for correction [30]. The intrinsic reaction coordi-
˚
transition state (TS1, Fig. 12) is 2.02 A. At TS1, the
formic acid transfers its proton to aldehyde oxygen
for generation of O3–H4 bond, and amine transfers
its proton to the anion format (see step1, Fig. 8). The
C2─N6, H4─O3, and N6─H13 bond lengths are 2.02,
˚
1.001, and 1.52 A in the TS1. The transition state of
step1 (TS1) showed one imaginary frequency at 259.82i
cm−1. The most important geometrical features, which
are undergoing changes in the first step, are illustrated
in Table III (see step1, Fig. 8).
At the TS1, C2─O3 and C17─O16 bond lengths are
˚
1.33 and 1.31 A and the corresponding bonds at I1
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EXPERIMENTAL AND THEORETICAL INSIGHT INTO THE KINETICS
9
To prove the role of formic acid as a catalyst, scan-
ning of C2─N6 bond formation between reactants 1 and
2 was carried out in the absence formic acid; herein,
there was no TS1 and I1. As a result, the role of formic
acid as a catalyst is recognized as a necessary tool for
beginning of the reaction.
Step2. In step2, H2O is eliminated in a concerted re-
action that is initiated with the O3─H14 bond forma-
tion and C2─O3 and N6─H14 bonds cleavage while
passing through a transition state (TS2) on the basis
of a four-membered ring geometry for the generation
of intermediate I2. To obtain the profile energy, we
scanned the distance between O3 and H14 atoms. The
Figure 12 The potential energy profile for step1 of the pro-
posed mechanism at the B3LYP/6–311+G(d,p) level. [Color
figure can be viewed at wileyonlinelibrary.com]
are 1.40 and 1.34 Ǻ, respectively. The C2─O3 and
C17─O16 bond lengths in the I1 are longer than in the
reactants (1 + 2) that showed double bonds (C2═O3
and C17═O16) in the reactant changed to single bonds
in the I1. The C17 atom in formic acid has two bonds
with O15 and O16 atoms (Fig. 8, step1), containing a
single bond (C17─O15) and double bond (C17═O16). In
the proton transfer of step1, these bonds are exchanged
to each other. So in the I1 (C17─O16) bond is a single
bond and (C17═O15) bond is a double bond. The struc-
ture and potential profile energy of TS1 is shown in
Fig. 12. There is hydrogen bonding between hydrogen
and heteroatom. In the literature, distance and angle
˚
TS2 has a forming of 1.15 A O3─H14 bond length.
The N6─H14 and C2─O3 bond lengths are 1.64 and
˚
˚
1.43 A in the TS2, whereas these are 3.23 and 4.95 A
in the intermediate 2 (I2). Step2 proceeded via a four-
member ring formation of TS2 because of the change
in H14─N6─C2─O3 dihedral angle. The hydrogen (H3)
atom and hydroxyl group (OH) should be located at
the same side during the progress of this event. After
this, N6─H14 bond starts to bond cleavage. A change
in dihedral angle (H14─N6─O3─C2) and N6─H14 bond
length during the scanning of O3─H14 bond is shown in
Figs. 13a and 13b, respectively.
˚
of hydrogen bonding are definite 1.5–2.2 A and 90°–
Transition state(TS2) is located on at 41st step,
whereas the dihedral angle (O3─C2─N6-H14) is exactly
0° (see Fig. 13a); and N6─H14 bond distance sharply
increased (see Fig. 13b). The TS2 showed one imagi-
nary frequency at 1572.45i cm−1. The most important
geometrical features, which are undergoing changes in
the second step, are illustrated in Table IV (see step2
180°, respectively [35]. In the reactant (1 + 2), two
H-bonding interaction were observed between the O3
and H4 also O16 and H13 atoms with hydrogen bond
˚
length of 1.74 A and 2.01 that exchange to O3─H4
˚
and O16─H13 bond with bond length 0.97 and 1 A,
respectively.
Table III Some of Bond Lengths, Bond Angles, and Dihedral Angles for step1 at the B3LYP/6–311+G(d,p) Level of the
Proposed Mechanism
Step1
Geometrical Parameter
Reactant (1 + 2)
TS1
I1
˚
Bond lengths (A)
C1−C2
C2−O3
1.47
1.23
1.74
0.99
1.22
1.34
2.01
2.80
1.39
1.47
1.33
1.002
1.67
1.31
1.24
1.07
2.02
1.40
1.52
1.40
0.97
1.93
1.34
1.21
1.00
1.49
1.42
163.25
71.58
−86.44
24.95
−88.59
149.178
O3−H4
H4−O15
C17−O16
C17−O15
O16−H13
C2−N6
N6−C7
Bond angle (deg)
O15−H4−O3
N6−C2−O3−H4
C7−N6−C1−C2
H14−N6−C2−O3
H14−N32−C2−H5
H14−N6−C2−C1
175.91
64.74
158.36
83.77
−34.86
−148.913
175.91
65.79
132.31
60.61
−52.54
−170.44
Dihedral angle (deg)
International Journal of Chemical Kinetics DOI 10.1002/kin.21064

10
HABIBI-KHORASSANI ET AL.
Figure 14 The potential energy profile for step2 of the pro-
posed mechanism at the B3LYP/6–311+G(d,p) level. [Color
figure can be viewed at wileyonlinelibrary.com]
which indicate that the single bond C2─N6 changed to
the double bond. The enhancement of C7─N6─C2─C1
dihedral angle implies that the two benzene rings are
the farthest from each other. The distance between H4
and N10 atoms changed from 5.57 in I1 to 1.97 in I2 that
show hydrogen bonding formed in I2 (see I2 in step2,
Fig. 8). The TS2 energy is 47.58 and 58.49 kcal mol−1
higher than the energy of reactant (I1) and intermediate
I2, respectively (see Fig. 14.). Therefore, this process is
exothermic (࢞r = −10.91 kcal mol−1) in comparison
with the reactants.
Figure 13 (a) Change in the dihedral angle of
H14─N6─C2─C3 during the scanning of N6─H14 bond;
(b) change in the N6─H14 bond during the scanning of the
O3─H14 bond. [Color figure can be viewed at wileyonlineli-
brary.com]
Fig. 8). The structure and potential energy profile of
TS2 are indicated in Fig. 14.
Step3. The third step is accompanied with a six-
membered ring formation during internal nucleophilic
attack of N10 atom on C2 atom of I2. The concerted
reaction initiated with a N10─C2 bond formation while
passing through a third transition state (TS3). The N10
and C2 atoms draw close and led to N10─C2 bond at in-
termediate I3. The TS3 has a forming bond distance of
At the TS2, C1─C2, C2─N6 and C2─O3 bond
lengths, C7─N6─C2 and C1─C2─N6 angles, and also
H14─N6─C2─H5 and C7─N6─C2─C1 dihedral angles
˚
are 1.50, 1.39, and 1.64 A; 119.85° and 119.63°; and
also 105.72° and 127.23°, respectively. The compari-
son between molecular geometries of I1, TS2, and I3
species show that the bond length of C2─N6 decreased
˚
1.79 A (N10─C2). The transition state of step3 showed
one imaginary frequency at 461.02i cm−1. At the TS3
˚
˚
˚
from 1.46 A in I1 to 1.39 A in TS2 and 1.28 A in I2,
C1─C2, C2─C6, C6─C7, C8─C9, and C9─N10 bond
Table IV Some of Bond Lengths, Bond Angles, and Dihedral Angles for Step2 at the B3LYP/6–311+G(d,p) Level of the
Proposed Mechanism
Step2
Geometrical Parameter
I1
TS1
1.50
1.39
1.64
1.89
1.43
I2
˚
Bond length (A)
C1−C2
C2−N6
1.53
1.46
1.41
2.02
1.01
1.46
1.28
3.23
1.99
4.95
C2−O3
H13−O15
N6−H14
H14−O3
2.28
1.15
0.96
H14−N6−C2−H5
C7−N6−C2−C1
H14−N6−O3−C2
H13−O15−H4
C7−N6−C2
C1−C2−N6
95.13
82.46
161.38
105.015
121.74
113.64
105.72
127.23
−178.011
115.66
119.851
119.63
46.44
175.40
143.96
129.46
120.045
122.612
Dihedral angle (deg)
Bond angle (deg)
International Journal of Chemical Kinetics DOI 10.1002/kin.21064

EXPERIMENTAL AND THEORETICAL INSIGHT INTO THE KINETICS
11
Table V Some of Bond Lengths, Bond Angles, and Dihedral Angles for Step3 at the B3LYP/6–311+G(d,p) Level of the
Proposed Mechanism
Step3
Geometrical Parameter
I2
TS3
1.49
1.35
1.34
1.43
1.49
2.69
2.07
1.79
1.79
I3
˚
Bond length(A)
C1−C2
C2−N6
1.46
1.28
1.41
1.50
1.34
1.94
2.25
1.69
3.76
175.80
120.77
123.49
118.46
1.52
1.40
1.32
1.40
1.54
2.64
1.96
1.79
1.55
N6−C7
C8−C9
C9−N10
H11−N6
H12−O15
H4−O18
N10−C2
Dihedral angle (deg)
Bond angle (deg)
C1−C2−N6−C7
C7−N6−C2
C1−C2−N6
C8−C9−N10
−167.56
117.01
117.97
114.429
−166.835
115.58
114.72
112.20
As can be seen, the major changes can be expressed as follows:
(1) C1−C2, C2−N6, C9−N10 bonds decreased in this order: I2 > TS3 > I3.
(2) N6−C7, C8−C9 bonds increased in this order: I2 < TS3 < I3.
(3) C7−N6−C2, C1−C2−N6, C1−C2−N6 and C8−C9−C10 angles are planar in the I2, but decreased in the I3.
lengths; C7─N6─C2, C1─C2─N6, and C8─C9─C10 an-
Fig. 15). Therefore, this process is endothermic (࢞r =
36.1 kcal mol−1) in comparison with the reactants.
gles; and C1─C2─N6─C7 dihedral angles are 1.49,
˚
1.35, 1.34, 1.43, and 1.49 A; 117.01°, 117.97°, and
114.429°; and –167.56°, respectively.
Step4. Step4 is accompanied with intermolecular hy-
drogen shift in I3, for generation of the H12─N6bond
in I4. The forth transition state TS4 has a distance of
The most important geometrical features, which are
undergoing changes in the third step, are illustrated in
Table V (see step3, Fig. 8). The structure and potential
energy profile of TS3 is shown in Fig. 15.
Herein, there are three hydrogen bondings
(H11─N6, H12─O15, and H4─O18), the distance of
H11─N6 hydrogen bond increased, but in H12─O15 de-
creased and the distance of H4─O18 hydrogen bond-
˚
˚
N6-H12 1.48 A whereas it is 2.64 A in reactant (I3). In
the TS4, the C2─C6, N6─C7, C7─C8, C8─C9, C9─C10,
N10─C2, and N10─H12 bond lengths; and C1─C2─N10,
C1─C2─N6, and N6─C2─N10 angles are 1.45, 1.38,
˚
1.43, 1.44, 1.45, 1.53, 1.26 A; and 110.81°, 114.76°,
and 96.21°, respectively.
ing is approximately constant (see steps3 and
,
4
The transition state of step4 showed one imaginary
frequency at 1762.43i cm−1. The most important ge-
ometrical features, which are undergoing changes in
the fourth step, are illustrated in Table VI (see step4,
Fig. 8). The structure and potential energy profile of
TS4 is shown in Fig. 16.
Fig. 8). These cases represent the relationship between
the amid group and formic acid. The TS3 energy is
38.03 and 2.13 kcal mol−1 higher than the energy
of reactant (I2) and intermediate I3, respectively (see
The distance of hydrogen bonding of H12─O15 is
˚
˚
1.97 A in the I3 that changed to 2.22 A in the I4.
The strength of hydrogen bonding has been during
the hydrogen shift process reduced. The TS4 energy
is 24.04 and 63.34 kcal mol−1 higher than the energy
of reactant I3 and intermediate I4, respectively (see
Fig. 16). Therefore, this process is exothermic (࢞r =
−39.3 kcal mol−1) in comparison with the reactants.
Step5. The final step is the oxidation process of I4 by
the oxygen molecule (O20═O19) in the air. In the litera-
ture, singlet oxygen is used for oxidation of secondary
amine to imines [35,36]. So, the singlet oxygen is used
in our calculations. In step5, there are three hydrogens
Figure 15 The potential energy profile for step3 of the pro-
posed mechanism at the B3LYp/6–311+G(d,p) level. [Color
figure can be viewed at wileyonlinelibrary.com]
International Journal of Chemical Kinetics DOI 10.1002/kin.21064

12
HABIBI-KHORASSANI ET AL.
Table VI Some of Bond Lengths and Bond Angles for
step4 at the B3LYP/6–31G(d,p) Level of the Proposed
Mechanism
Geometrical
Step4
Bond length (A)
Parameter
I3
TS4
I4
˚
C2−N6
N6−C7
C7−C8
1.40
1.32
1.46
1.40
1.54
1.55
1.03
1.79
2.65
1.97
1.45
1.38
1.43
1.44
1.45
1.53
1.26
1.77
1.48
2.13
1.46
1.39
1.41
1.48
1.36
1.46
3.26
1.71
1.01
2.22
C8−C9
C9−C10
N10−C2
N10−H12
O18−H4
H11−N6
H12−O
C1−C2−1N510 107.85 110.81 111.22
Bond angle (deg) C1−C2−N6 114.723 114.76 111.30
N6−C2−N10 110.13
96.215 107.07
As can be seen, the major changes can be expressed as follows:
Figure 17 (a) The potential energy profile for step5 of the
proposed mechanism at the B3LYP/6–311+G(d,p) level
when the oxygen molecule is closed to the hydrogen atom
(H5) from the C2─N6 side; (b) the potential energy profile for
step5 of the proposed mechanism at the B3LYP/6–311+G(d,
p) level when the oxygen molecule is close to the hydro-
gen atom (H5) from the C2─N10 side. [Color figure can be
viewed at wileyonlinelibrary.com]
(1) C7−C8, C9−C10, N10−C2 bonds decreased in this order: I3 >
TS4 > I4.
(2) C2−N6, N6−C7, C8−C9 bonds increased in this order: I4 <
TS4 < I3.
(3) C1−C2−N10, C1−C2−N6, N6−C2−N10 angles changed
about 4°.
In the second possibility, the oxygen molecule should
be close to H5 from the location near to the C2─N10
bond.
So the first path is the concerted reaction between
oxygen number 19 in the O2 molecule (O20═O19) and
H5 upon the C2─N6 bond, and then reaction with oxy-
gen number 20 in the O2 molecule (O20═O19) while
passing through a TS5a for generation of product 3.
The second path is concerted reaction between O19
atom and H5 upon the C2─N10 bond, and then reaction
with O20 atom while passing through a TS5b for gen-
eration of product 4 (see step5, Fig. 8). The distance
between O19 and H5 atoms was scanned for obtaining
the profile of energy. The structures and potential en-
ergy profiles of TS5a and TS5b for the two pathways
are shown in Fig. 17a and 17b, respectively.
Figure 16 The potential energy profile for step4 of the pro-
posed mechanism at the B3LYp/6–311+G(d,p) level. [Color
figure can be viewed at wileyonlinelibrary.com]
(H12, H5, and H11, step5 Fig. 8) that could participate
in the oxidation process of I4. Oxygen molecule can
attack to hydrogen numbers 11, 12, or 5. When an
oxygen molecule is close to hydrogen number five, the
product 3 or 4 is generated while passing through a
transition state (TS5). Now, the question is, which hy-
drogen atoms (12 or 11) can be reacted to other oxygen
in O2 (O20═O19) molecule for generation of products.
This is related to the orientation of oxygen molecule
(O20═O19) for capturing hydrogen number five. There
are two basic features, in the first possibility, oxygen
molecules could be close to hydrogen number five from
the side near to the C2─N6 bond (see step5, Fig. 8).
When the scan process was carried out from the
C2─N6 side, the TS5a energy is 2.47 and 66.79 kcal
mol−1 higher than the energy of reactant (I4) and
product 3, respectively (see Fig. 17a). While the scan
process performed from the C2─N10 side, the energy
TS5b is 2.67 and 63.28 kcal mol−1 higher than the en-
ergy of (I4) reactant and product 4, respectively (see
Fig. 17b). The barrier from the C2─N6 side is slightly
less than C2─N10 side and product 3 (66.79 kcal mol−1
)
International Journal of Chemical Kinetics DOI 10.1002/kin.21064

EXPERIMENTAL AND THEORETICAL INSIGHT INTO THE KINETICS
13
Table VII Some of Bond Angles and Dihedral Angles for step5 at the B3LYP/6–311+G(d,p) Level of the Proposed
Mechanism
Step5
Geometrical Parameter
I4
TS5
1.27
1.40
1.01
1.42
1.37
1.28
2.20
1.30
113.06
114.88
114.71
103.67
−68.86
72.52
Product3
C2−H5
C2−N6
1.22
1.44
1.008
1.45
1.38
1.84
2.83
1.30
4.06
1.37
1.40
0.98
0.96
1.45
N6−H12
C2−N10
N10−C9
O20−H5
O19−H12
2.35
1.25
O20−O19
N6−C2−N10
N6−C2−C1
C1−C2−N10
H5−C2−C1
H12−N6−C2−H5
H5−C2−N10−H11
110.60
111.31
112.43
106.72
−90.133
65.58
121.97
121.19
121.97
89.35
−25.88
−175.91
Bond angle (deg)
Dihedral angle (deg)
is more stable than the product 4 (63.28 kcal mol−1),
so the product 3 kinetically and thermodynamically
generated easily than product 4. The TS5a has a O20-
˚
˚
H5 distance of 1.28 A, whereas it is 1.84 and 0.98 A
in reactant (I4) and product 3, respectively. At 1.135
˚
A of H5─O20 bond length, energy decreases sharply
Figure 18 Two possible products during the oxidation pro-
cess in step5 of the reaction.
because of the H5─C2 and N6-H12 bonds cleavage
and the two H5─O19 and H12─O19 bonds formation
(Table VII).
In the TS5a C2─H5, C2─N6, N6─H14, C2─N10,
O20─H5, O19─H12, and O20─O19 bond length;
N6─C2─C1, C1─C2─N10, and N6─C2─N10 angles;
and H12─N6─C2─H5 dihedral angle are 1.27, 1.40,
thermodynamically more stable than product 4, and the
tautomer formed and kinetically generated easier than
product 4 (see Table VIII).
˚
1.01, 1.42 1.28, 2.20, and 1.25 A; 114.88°, 114.7°, and
113.06°; and −68.865°, respectively (Table VII).
As expected, the C2─N6 and C2─N10 bond length
decreased along with the change from I4 to product
3 and N6─C2─C1 and C1─C2─N10 and N6─C2─N10
angle increased.
The two different products (3 and 4) formed from
the two possibilities. These are shown in Fig. 18. Both
of them optimized at the B3LYP/6–311+G (d,p) level.
Also, the tautomer of products 3 and 4 is shown in
Fig. 18 that was optimized at the B3LYP/6–311+G
(d,p) level. The results indicated that the product 3 is
Rate Constants
The other goal of theoretical study is to predict the reac-
tion rate constants. The kinetic parameter for various
steps calculated with the transition state theory [36].
Detailed investigation on the rate constants was per-
formed to provide insight into the competition between
various reaction steps. The potential energy surface and
transition state parameter were utilized to predict the
rate constant as shown in Eq. (24), where kB, h, and
ꢁ are the Boltzmann, Plank’s constant, and tunneling
Table VIII Energy, Enthalpy, and Gibbs Free Energy of Products 3 and 4 Calculated at the B3LYP/6–311+G(d,p) Level
at 298 K in kcal mol−1
4
3
Tautomer Form
(E°) Energy
(H°) Enthalpy
(G°) Gibbs free energy
−454,463.41
−454,462.82
−454,496.29
−454,472.55
−454,471.96
−454,505.14
−454,466.9049
−454,466.3125
−454,499.5091
International Journal of Chemical Kinetics DOI 10.1002/kin.21064

14
HABIBI-KHORASSANI ET AL.
Table IX Rate Constant and Activation Energies for All Steps of the Speculative Mechanism at the B3LYP/6–311+G
(d,p) Level
Step
Rate Constant
Preexponential Factor (Frequency Factor)
Ea Forward (kcal/mol)
1
k1 = 1.34 × 10−8a
k–1 = 8.97
A1 = 9.56 × 107a
A–1 = 1.90 × 1011b
A2 = 3.79 × 1013b
A3 = 5.12 × 1011b
A4 = 6.88 × 1012b
A5 = 2.07 × 1012b
28.19
−1
2
k2 = 1.77 × 10−19b
k3 = 1.03 × 10−16b
K4 = 9.49E × 104
K5 = 1.40 × 10+12
44.69
38.37
22.22
0.82
3
4
5
aL mol−1s−1
bs−1
.
.
Table X Comparison of the Tunneling Corrections for
Each Step of Mechanism Using the Wigner’s Approach
correction, respectively. The tunneling correction fac-
tor was estimated using Wingers formula as shown in
Eq. (27). The activation energy and Arrhenius factor
(A) could also be obtained from Eqs. (25) and (26):
Tunneling Correction
Step
Using Wigner’s Method
1
2
3
4
5
1.01
3.40
1.193
4.02
1.42
kBT
h
=/RT
(
)
k = ꢁ
e− ꢀG
(24)
(25)
(26)
Ea = ꢀH= + nRT
enkBT
Table XI Activation Parameters Involving for Each
Step of the Speculative Mechanisms at 298.15 K at the
B3LYP/6–311+G(d,p) Level
‡
A =
eꢀS /R
h
Step1 ࢞G‡a 25.02 ࢞H‡a
21.18 ࢞S‡b −0.012
1
1
1
ꢀ
ꢁ
Step-1 ࢞G‡
16.05 ࢞H‡
13.39 ࢞S‡
−0.008
2
hv‡
–1
–1
–1
1
Step2 ࢞G‡
43.03 ࢞H‡
43.51 ࢞S‡
0.001
ꢁ = 1 +
(27)
Step3 ࢞G‡2 39.26 ࢞H‡2 37.19 ࢞S‡2 −0.006
24 kBT
Step4 ࢞G‡3 21.57 ࢞H‡3 21.04 ࢞S‡3 −0.001
4
0.88 ࢞G‡4 −0.361 ࢞S‡4 −0.004
Table IX illustrates the comparison between rate con-
stants and activation energies in all steps. Step1 is a bi-
molecular reaction, whereas other steps are unimolec-
ular. However, the higher and smaller rate constant val-
ues can be attributed to the RDS and the fastest step,
respectively. Among unimolecular steps, step two is
the RDS, which has a Ea,forward = 44.69 and k2 = 1.77
× 10−19 s−1 As expected, the reveres step1 rate con-
stant is 8.97 s−1 which is larger than step2 (1.77 ×
Step5 ࢞G‡
–
5
5
5
tropy, and free energy for each step of the speculative
mechanism at 298.15 K are listed in Table XI. Activa-
tion free energy like the rate constant shows whether a
step is fast or slow. Thus, steps2 and have the max-
5
10−19
s
−1) .Thus, step1 has an equilibrium state.
imum and minimum values of activation free energy,
respectively. Therefore, step2 and step5 are the slow
and fast steps, respectively.
The oxidation step (step5) has the largest rate con-
stant and smallest activation energy, so step five is the
fastest step. So this step is virtually without barrier
energy (0.82 kcal mol−1).
The tunneling correction using the Wigner’s ap-
proach is compared and tabulated in Table X.
Step2 and step4 have the maximum of tunneling cor-
rection since proton transfer reactions have the largest
imaginary frequency.
In addition, Table XII shows thermodynamic pa-
rameters containing enthalpy, free energy, and entropy
values obtained at the B3LYP/6–311+G(d, p) level for
each step involved in the mechanism. The comparison
between the free energy value show that steps2, 4, and
are exothermic, and steps1 and are endothermic.
5
3
The ࢞rG0 and ࢞rH0 values of step5 are more nega-
tive than all steps, and step3 is a more positive step
with respect to others. Step2 (k2) is the RDS and ki-
netically is a hard step; nevertheless it is a favorable
Mechanical calculations for obtaining activation pa-
rameters performed at the B3LYP/6–311+G (d,p) level
for each step of the reaction. Activation enthalpy, en-
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EXPERIMENTAL AND THEORETICAL INSIGHT INTO THE KINETICS
15
Table XII Reaction Parameters Involving ࢞rH0,
In addition, full kinetic and a mechanistic investi-
gation of the reaction was undertaken using UV and
stopped-flow spectrophotometry techniques. The re-
sults can be summarized as follows:
࢞rG01, ࢞rS0 for Each Step of the Speculative
1
Mechanisms at 298.15 K
Step1 ࢞rH0a
7.53 ࢞rG0a
12.03 ࢞rS0b −0.015
1
1
1
Step2 ࢞rH0 −11.92 ࢞rG0 −12.21 ࢞rS0
0.0009
2
3
2
3
2
Step3 ࢞rH0
36.26 ࢞rG0
38.366 ࢞rS0 −0.007
1. The overall reaction order followed second-
order kinetics and the partial reaction orders with
respect to each of reactants, benzaldehyde 1, 2-
aminobenzamide 2, are 1 and 1, respectively.
2. The values of the overall rate constants of the
reaction were calculated successfully at all sol-
vents and temperatures investigated.
3. The rates of all reactions decreased in solvents
with lower dielectric constant, and this can be
related to differences in stabilization of the re-
actants and the activated complex by the solvent
in the transition state.
3
Step4 ࢞rH0 −38.60 ࢞rG0 −38.94 ࢞rS0
0.001
0.005
4
4
4
5
Step5 ࢞rH0 −63.52 ࢞rG0 −65.16 ࢞rS0
5
5
Table XIII Thermodynamic Parameters Involving
࢞rH0total, ࢞rG0 and ࢞rS0
for Overall Reaction
total,
total
of the Speculative Mechanisms at 298.15 K at the
B3LYP/6–311+G(d,p) Level
Mechanism Involving
Steps1, 2, 3, 4, 5
࢞rH0a
࢞rG0a
࢞rS0b
total
total
total
4. With respect to the experimental data, the second
step of proposed mechanism was recognized as
a RDS(k2) and this was confirmed based upon
the steady-state approximation and stopped-flow
method.
−70.26
−65.91
−0.0145
5. Activation energy (64.43 kJ mol−1) was ob-
tained for the overall reaction by UV–vis experi-
ment. In addition, activation parameters (࢞S‡ =
50.4277 0.0015 J (mol K)–1, ࢞H‡ = 4.2817
0.0048, and ࢞G‡ = 60.88 kJ mol–1) were
obtained for the first step of reaction as an ele-
mentary step.
thermodynamical process. Thermodynamic values for
the overall reaction of the proposed mechanisms at
298.15 K are given in Table XIII. The free energy and
enthalpy for the overall reaction shows that reaction is
exothermic and thermodynamically and kinetically is
favorable and executable.
Moreover, theoretical calculations at the B3LYP/6–
31+G(d,p) level have been used for determination of
mechanism and RDS of reaction between benzalde-
hyde 1 and 2-aminobenzamide 2 within in the presence
of formic acid. The proposed mechanism involved five
steps. The results show that step2 is RDS and they are
in agreement with the results obtained by the UV and
stopped-flow spectrophotometry techniques.
CONCLUSION
In summary, an efficient and simple method has
been developed for the formation of highly func-
tionalized 2,3-dihydro-quinazolin-4(1H)-ones via two-
component under mild conditions using formic acid
as a solvent and catalyst. This reaction can be em-
ployed as an efficient approach for the preparation
of synthetically and pharmaceutically important 2,3-
dihydroquinazolin-4(1H)-ones systems. This method-
ology offered several advantages such as mild re-
action conditions, easy work-up procedure, simple
and readily available precursors, high atom efficiency,
clean reaction profiles, nontoxic and inexpensive cata-
lyst, no need to column chromatography, environmen-
tally friendly catalyst, and good to high yields. 2,3-
Dihydroquinazolin-one derivatives have been reported
to possess a wide range of pharmacological activities
such as anti inflammatory and analgesic activities and
antitumor, anticancer, antibacterial, and diuretic activ-
ities. In addition, these compounds as chemical inter-
mediates can easily be oxidized to their corresponding
quinazolin-4(3H)-ones, which are also important bio-
logically active heterocyclic compounds.
BIBLIOGRAPHY
1. Yale, H. L.; Kalkstein, M. J Med Chem 1967, 10, 334–
336.
2. Neil, G. L.; Li, L.; Buskirk, H. H.; Moxley, T. E. Cancer
Chemother Rep, Part 1 1972, 56, 163–173.
3. Bonola, G.; Da Re, P.; Magistretti, M.; Massarani, E.;
Setnikar, I. J Med Chem 1968, 11, 1136–1139.
4. Bolger, J. W.; Google Pat. 1966.
5. Sohn, B. Chem Abstr 1964, 16075h.
6. Alaimo, R. J.; Russell, H. E. J Med Chem 1972, 15,
335–336.
7. Cohen, E.; Klarberg, B.; Vaughan Jr, J. R. J Am Chem
Soc 1959, 81, 5508–5509.
8. Okumura, K.; Oine, T.; Yamada, Y.; Hayashi, G.;
Nakama, M. J Med Chem 1968, 11, 348–352.
International Journal of Chemical Kinetics DOI 10.1002/kin.21064

16
HABIBI-KHORASSANI ET AL.
9. Khoobi, M.; Ma’mani, L.; Rezazadeh, F.; Zareie, Z.;
Foroumadi, A.; Ramazani, A.; Shafiee, A. J Mol Catal
A: Chem 2012, 359, 74–80.
10. Ahankar, H.; Ramazani, A.; Joo, S. W. Res Chem In-
termed 2016, 42, 2487–2500.
11. Mirhadi, E.; Ramazani, A.; Rouhani, M.; Joo, S. W.
Chemija 2013, 24.
12. Ramazani, A.; Reza Kazemizadeh, A. Curr Org Chem
2011, 15, 3986–4020.
22. Darijani, M.; Habibi-Khorassani, S. M.; Shahraki, M.
Int J Chem Kinet 2015, 47, 751–763.
23. Lente, G.; Fa´bia´n, I.; Poe¨, A. J. New J Chem 2005, 29,
759–760.
24. Espenson, J. Chemical Kinetics and Reaction Mecha-
nism, McGraw-Hill: New York, 1995.
25. Poe, A.; Twigg, M. Mechanisms of Inorganic and
Organometallic Reactions. Plenum: New York, 1994,
vol. 8, p. 218.
13. Reza Kazemizadeh, A.; Ramazani, A. Curr Org Chem
2012, 16, 418–450.
14. Ramazani, A.; Rouhani, M.; Joo, S. W. Ultrason
Sonochem 2016, 28, 393–399.
26. Becke, A. J Chem Phys 1993, 104, 1040.
27. Lee, C.; Yang, W.; Parr, R. Phys Rev B 1988, 37, 785.
28. Montgomery Jr, J.; Ochterski, J.; Petersson, G. J Chem
Phys 1994, 101, 5900–5909.
15. Khoobi, M.; Ramazani, A.; Hojjati, Z.; Shakeri, R.;
Khoshneviszadeh, M.; Ardestani, S. K.; Shafiee, A.;
Foroumadi, A.; Joo, S. W. Phosphorus Sulfur Silicon
Relat Elem 2014, 189, 1586–1595.
29. Merrick, J. P.; Moran, D.; Radom, L. J Phys Chem A
2007, 111, 11683–11700.
30. Gonzalez, C.; Schlegel, H. B. J Chem Phys 1989, 90,
2154–2161.
31. Som, S.; Liu, W.; Zhou, D. D.; Magnotti, G. M.; Sivara-
makrishnan, R.; Longman, D. E.; Skodje, R. T.; Davis,
M. J. J Phys Chem Lett 2013, 4, 2021–2025.
32. Hirschfelder, J.; Wigner, E. J Chem Phys 1939, 7, 616–
628.
´
16. Ahankar, H.; Ramazani, A.; Slepokura, K.; Lis, T.; Joo,
S. W. Green Chem 2016, 18, 3582–3593.
17. Ramazani, A.; Rouhani, M.; Mirhadi, E.; Sheikhi, M.;
´
Slepokura, K.; Lis, T. Nanochem Res 2016, 1, 87–
107.
18. Shahraki, M.; Habibi-Khorassani, S. M. J Phys Org
Chem 2015, 28, 396–402.
19. Habibi-Khorassani, S. M.; Ebrahimi, A.; Maghsoodlou,
M. T.; Shahraki, M.; Price, D. Analyst 2011, 136(8),
1713–1721.
20. Shahraki, M.; Habibi-Khorassani, S. M.; Dehdab, M.
RSC Adv 2015, 5, 52508–52515.
21. Pourpanah S. S.; Habibi-Khorassani, S. M.; Shahraki,
M. Chin J Catal 2015, 36(5), 757–763.
33. Schmidt, M. W.; Baldridge, K. K.; Boatz, J. A.; Elbert, S.
T.; Gordon, M. S.; Jensen, J. H.; Koseki, S.; Matsunaga,
N.; Nguyen, K. A.; Su, S. J Comput Chem 1993, 14,
1347–1363.
34. Jiang, G.; Chen, J.; Huang, J.-S.; Che, C.-M. Org Lett
2009, 11, 4568–4571.
35. Lin, S.; Yang, L. Tetrahedron Lett 2005, 46, 4315–4319.
36. Glasstone, S.; Eyring, H.; Laidler, K. J. The Theory of
Rate Processes; McGraw-Hill: New York, 1941.
International Journal of Chemical Kinetics DOI 10.1002/kin.21064