[2π + 2π]-Cycloaddition of biadamantylidene to 4-phenyl-3H-1,2,4-triazole-3,5(4H)-dione. Effects of temperature, high pressure, and solvent

Article abstract of DOI:10.1134/S1070428017120144

The effects of temperature, solvent nature, and high hydrostatic pressure on the rate of the reaction of biadamantylidene with 4-phenyl-3H-1,2,4-triazole-3,5(4H)-dione have been estimated. Significant shielding of the C=C double bond in biadamantylidene is responsible for the high entropy and volume of activation. Quantitative yield of the reaction in the temperature range 25?45°C is related to its exothermicity. The rate of the [2π + 2π]-cycloaddition unexpectedly weakly depends on the solvent polarity, which makes it radically different from the [2π + 2π]-reaction with tetracyanoethylene.

Full text of DOI:10.1134/S1070428017120144

ISSN 1070-4280, Russian Journal of Organic Chemistry, 2017, Vol. 53, No. 12, pp. 1864–1869. © Pleiades Publishing, Ltd., 2017.
Original Russian Text © V.D. Kiselev, D.A. Kornilov, O.V. Anikin, I.A. Sedov, A.I. Konovalov, 2017, published in Zhurnal Organicheskoi Khimii, 2017,
Vol. 53, No. 12, pp. 1828–1833.
[2π+2π]-Cycloaddition of Biadamantylidene
to 4-Phenyl-3H-1,2,4-triazole-3,5(4H)-dione. Effects
of Temperature, High Pressure, and Solvent
V. D. Kiseleva,* D. A. Kornilova, O. V. Anikina, I. A. Sedova, and A. I. Konovalovb
a Butlerov Institute of Chemistry, Kazan Federal University, ul. Kremlevskaya 18, Kazan, 420008 Tatarstan, Russia
*e-mail: vkiselev.ksu@gmail.com
b Arbuzov Institute of Organic and Physical Chemistry, Kazan Scientific Center, Russian Academy of Sciences,
ul. Arbuzova 8, Kazan, 420088 Tatarstan, Russia
Received April 7, 2017
Abstract—The effects of temperature, solvent nature, and high hydrostatic pressure on the rate of the reaction
of biadamantylidene with 4-phenyl-3H-1,2,4-triazole-3,5(4H)-dione have been estimated. Significant shielding
of the C=C double bond in biadamantylidene is responsible for the high entropy and volume of activation.
Quantitative yield of the reaction in the temperature range 25‒45°C is related to its exothermicity. The rate of
the [2π+2π]-cycloaddition unexpectedly weakly depends on the solvent polarity, which makes it radically
different from the [2π+2π]-reaction with tetracyanoethylene.
DOI: 10.1134/S1070428017120144
A large number of studies have been performed on
[2+2]-cycloadditions with various pairs of substituted
ethylenes, where one component exhibits pronounced
electron donor properties (vinyl ethers, tetramethoxy-
ethylene, cyclopropyl-substituted ethylenes) and the
other is an electron acceptor (di-, tri-, and tetracyano-
ethylenes) [1–5]. In most publications, an appreciable
effect of molecular asymmetry of the donor and ac-
ceptor reagents on the reaction rate was noted. Reac-
tions of vinyl ethers with tetracyanoethylene were
found to accelerate in polar medium [1, 5, 6]. The
stereospecificity of the addition was conserved by
about 90%. The reduced stereospecificity in [2+2]-
cycloadditions gave grounds to presume that thermal
reactions forbidden by the Woodward–Hoffmann
orbital symmetry conservation rules involve inter-
mediate formation of zwitterionic species. Huisgen [1]
concluded that the rate of zwitterion transformation
into cyclobutane derivative is approximately 5 times
higher than the rate of its decomposition into initial
tetracyanoethylene and vinyl ether. For the reaction of
tetracyanoethylene with anethole the rate constant ratio
k(acetonitrile)/k(cyclohexane) is 29000, for the reaction
with ethoxyisobutene, 10800, and for the reaction with
butyl vinyl ether, 2600. The ratio k(MeCN)/k(CCl4) for
the reaction of tetracyanoethylene with 3,4-dihydro-
2H-pyran is 17000 [6]. The observed differences in
the rates of [2 +2]-cycloadditions with tetracyano-
ethylene in polar and nonpolar solvents is consistent
with the greater stabilization of polar transition state
by 5‒6 kcal/mol relative to low-polar initial mole-
cules [1, 6].
In the dienophile series 1,1-dicyanoethylene (elec-
tron affinity Ea = 1.53 eV), tricyanoethylene (Ea =
2.10 eV), and tetracyanoethylene (Ea = 2.88 eV) [7, 8],
the rate of the Diels–Alder reactions with dienes in-
creases in parallel with the electron affinities of dieno-
philes [8, 9], whereas [2 +2]-cycloadditions of the
same dienophiles with vinyl ether display the opposite
variation of the reaction rate [1]. A strong effect of
the alkene asymmetry on the reaction rate was also
observed in [2 + 2]-cycloadditions of tetracyano-
ethylene to cyclopropyl-substituted ethylene. 1,1-Di-
cyclopropylethylene turned out to be more reactive by
almost 6 orders of magnitude than cis- and trans-1,2-
dicyclopropylethylenes [3, 4]. The strong influence of
solvent on the rate of reactions involving tetracyano-
ethylene may be rationalized by significant change of
solvation and reactivity of tetracyanoethylene. The rate
of the nonpolar Diels–Alder reaction of 9,10-dimethyl-
1864

[2π+2π]-CYCLOADDITION OF BIADAMANTYLIDENE
1865
Scheme 1.
Ph
N
Ph
O
O
N
O
O
(2)
N
N
NC)2C=C(CN)2 (4)
N N
1
3
anthracene with tetracyanoethylene increases almost
400-fold in going from weakly polar π-donor alkyl-
benzenes to chloroform [10].
(1) to 2 with formation of adduct 3 (Scheme 1) and
determined the enthalpies of these reactions and
equilibrium constants.
It seemed reasonable to study the effect of solvent
polarity on [2+2]-cycloadditions to 4-phenyl-3H-
1,2,4-triazole-3,5(4H)-dione (2) which exhibits mod-
erate acceptor properties comparable with those of
maleic anhydride [11]. If their is an accessible allylic
proton, compound 2 readily reacts with substituted
alkenes according to the ene reaction pathway, in con-
trast to the symmetry-forbidden [2+2]-cycloaddition.
The Alder-ene reaction of biadamantylidene (1) and
triazole 2 is forbidden by Bredt’s rule, since double
bond migration to the bridgehead carbon atoms in 1
would lead to a high strain energy.
According to [13], the reaction 1 + 2 → 3 in ben-
zene is characterized by an enthalpy of activation of
44.4 kJ/mol and an entropy of activation of –138 J×
mol–1 K–1. The rate constant at 25°C reported in [13]
coincided with our data (Table 1), but the enthalpy and
entropy of activation were strongly different. As shown
by X-ray analysis [13, 14], the four-membered ring in
adduct 3 is appreciably bent due to increased steric
hindrances to the approach of molecules 1 and 2 to
each other. This is consistent with increased negative
entropies of activation in all the examined solvents;
furthermore, the free energy of activation is con-
tributed mainly by the entropy term T∆S≠ (Table 1).
In continuation of our studies of the kinetics and
thermochemistry of Diels–Alder and Alder-ene reac-
tions of 4-phenyl-3H-1,2,4-triazole-3,5(4H)-dione (2)
[12], in the present work we estimated the effects of
the solvent nature, temperature, and high pressure on
the rate of [2+2]-cycloaddition of biadamantylidene
Unlike [2+2]-cycloaddition reactions with tetra-
cyanoethylene (4), the rate of which increases by 3–
4 orders of magnitude in going from nonpolar solvents
to acetonitrile, variation of the rate of the reaction
1 + 2 → 3 in the examined solvents by almost 3 orders
Table 1. Rate constants k2 (L mol–1 s–1), enthalpies ∆H≠ (kJ/mol), entropies ∆S≠ (J mol–1 K–1), and Gibbs energies of activa-
tion ∆G≠ (kJ/mol, 25°C) for the reaction 1 + 2 → 3 in different solvents
Solvent
εa
k2 (25°C)
k2 (35°C)/krel
k2 (45°C)
∆H≠
–∆S≠
∆G≠
Ethyl acetate
1,4-Dioxane
Benzene
06.0
02.2
02.3
0.00118
0.00138
0.00641
b0.00640b
–
0.00191/1.0
0.00232/1.2
0.00909/4.7
–
0.00929c/4.9
0.0170d/8.90
0.0236/12.3
0.111/58.1
0.577/302
1.78/932
0.00309
0.00353
0.13400
–
35.4
35.0
26.6
b44.4b
–
182
182
198
b138b
–
89.6
89.2
85.6
b85.5b
–
Acetone
20.7
36.7
02.3
37.5
10.3
04.6
–
Dimethylformamide
Toluene
–
–
–
–
–
0.01750
0.0716
0.4070
1.4600
0.03430
0.156
0.785
2.270
24.0
28.3
23.4
14.9
198
172
175
192
83.0
79.5
75.5
72.1
Acetonitrile
1,2-Dichloroethane
Chloroform
a
b
c
d
Dielectric permittivity [15].
Data of [13].
Calculated with correction for the reaction of 2 with acetone, which contributes ~10% to the overall rate.
Calculated with correction for the reaction of 2 with DMF, which contributes ~20% to the overall rate.
RUSSIAN JOURNAL OF ORGANIC CHEMISTRY Vol. 53 No. 12 2017

KISELEV et al.
1866
Table 2. Experimental (∆V≠exp, cm3/mol) and corrected (∆V≠corr, cm3/mol) volumes of activation for the reaction 1 + 2 → 3 in
toluene and ethyl acetate at 25°C
Solvent
Toluene
Ethyl acetate
k(1000 bar)/k(1 bar)
dlnkp/dr, 1 bar
‒∆V≠exp
βT RTa
‒∆V≠corr
6.45
7.90
0.00214
0.00238
53.1±1.4
58.8±1.5
2.3
3.0
50.8±1.4
55.8±1.5
a
Compressibility coefficients: toluene, 91.9×10–6 bar–1; ethyl acetate, 120×10–6 bar–1 [17].
d–1 = –(0.04064±0.00028) c3 + (1.1136663±0.0000013);
R2 = 0.9991; ∆Vr = –36.5 cm3/mol.
(2) c02 = 5.35×10–3 M; c01 = 6.37×10–2 M;
of magnitude is not determined by solvent polarity.
The rate constants in benzene, acetone, dimethylform-
amide, and acetonitrile are related as 1:1:1.9:12.2.
Considerable increase of the reaction rate in 1,2-di-
chloroethane and chloroform is likely to result from
hydrogen bonding between compound 2 and solvent
molecules, as noted for a number of Diels–Alder and
Alder-ene reactions with 2.
d–1 = –(0.04277±0.00047) c3 + (1.114418±0.0000008);
R2 = 0.9976; ∆Vr = –38.4 cm3/mol;
∆Vr (av.) = ‒37.5±0.8 cm3/mol;
∆V≠corr/∆Vr = –55.8/–37.5 = 1.48.
The enthalpy of the reaction 1 + 2 → 3 was deter-
mined in 1,2-dichloroethane at 25°C by seven double
measurements by adding a known volume (V, μL) of
a solution of 2, c02 = 0.1455 M, to 25 mL of a solution
of 1, c01 0.0126 M: (1) V2 = 60, ∆H = ‒85.2; (2) V2 =
60, ∆H = –86.0; (3) V2 = 60, ∆H = –85.5; (4) V2 = 60,
∆H = –87.4; (5) V2 = 80, ∆H = –87.0; (6) V2 = 80,
∆H = –86.7; (7) V2 = 80 μL, ∆H = ‒86.0 kJ/mol;
∆H(av.) = –86.3±0.8 kJ/mol. Thus, the heat effect of
the [2π+2π]-cycloaddition reaction under study is much
smaller than those of the [4π+2π]-, [2π+2π+2π]-, and
[2π+2σ+2σ]-additions and ene reaction [12].
The rate constant k2 (0.43±0.02 L mol–1 s–1) for the
reaction of 1 (c01 = 0.0126 M) with 2 (c02 = 4.64×
10–4 M) in 1,2-dichloroethane was calculated from
the curve of heat evolution which is proportional to
the concentration of adduct 3. A similar value
(0.407±0.010 L mol–1 s–1) was obtained from the rate
of variation of the concentration of 2 (Table 1). This
means that the rates of the reaction 1 + 2 → 3,
calculated from the change of the concentration of 3
(determined by calorimetry) and of the concentration
of 2 (determined by spectrophotometry), are similar.
Therefore, the rate of the transformation of possible
intermediate to adduct 3 should be higher than the rate
of its formation.
We previously found [16] that the change in en-
tropy of chemical reactions is often proportional to the
change in reaction volume. Therefore, increased
volume of activation may be expected. The volumes
of activation were determined for the reactions in
toluene and ethyl acetate from the reaction rates at
1 and 1000 bar (Table 2).
We obtained very high negative values of the
volumes of activation for the reaction 1 + 2 → 3 in
toluene (–50.8) and ethyl acetate (–55.8 cm3/mol),
which were in agreement with the large negative
entropies of activation of this reaction. It seemed
reasonable to compare the volumes of activation with
the volumes of reaction in these solvents. For this
purpose, we determined changes in the densities of
solutions and calculated concentrations of adduct 3 in
the course of the reaction 1 + 2 → 3, which were sub-
stituted into Eq. (12) (see Experimental).
By two measurements for the reaction in toluene
we obtained:
(1) c02 = 5.00×10–3 M, c01 = 2.46×10–2 M:
d–1 = –(0.04240±0.00020) c3 + (1.1574369±0.0000004);
R2 = 0.9994; ∆Vr = ‒36.6 cm3/mol;
(2) c02 = 2.02×10–2 M, c01 = 5.03×10–2 M;
d–1 = –(0.04393±0.00025) c3 + (1.1544040±0.0000032);
The equilibrium constants for the reaction 1 + 2 → 3
were determined in toluene. The following data were
obtained: at 70°C: c01 = c02 = 0.01317 M, D∞2 = 0.111,
R2 = 0.9990; ∆Vr = –38.0 cm3/mol;
∆Vr(av.) = –37.3±0.7 cm3/mol;
∆V≠corr/∆Vr = –50.8/–37.3 = 1.36.
c∞2 = 4.53×10–4 M, Keq = 6.2×104; at 80°C: c01
=
c02 = 0.01317 M, D∞2 = 0.165, c∞2 = 6.73×10–4 M,
Keq = 2.7×104. Then, the enthalpy of the reaction is
‒81.5 kJ/mol, and the entropy, –145 J mol–1 K–1. Fairly
close values of the reaction enthalpies in toluene
By two measurements for the reaction in ethyl
acetate we obtained:
(1) c02 = 1.50×10–2 M, c01 = 6.37×10–2 M;
RUSSIAN JOURNAL OF ORGANIC CHEMISTRY Vol. 53 No. 12 2017

[2π+2π]-CYCLOADDITION OF BIADAMANTYLIDENE
1867
Table 3. Rate constants k2 (L mol–1 s–1) for the [2π+2π]-cycloaddition of compound 2 with biadamantylidene (1) at 35°C,
[2π+2σ+2σ]-reaction of 2 with quadricyclane (5) at 25°C [18], [2π+2π]-reaction of 2 with 2-chloroethyl vinyl ether (6) at
25°C [19]; ene reaction of 2 with 2-methylbut-2-ene (7) at 10°C [20]; [4π+2π]-reaction of 2 with anthracene (8) at 25°C [21],
ene reaction of 2 with norbornene 9 at 40°C accompanied by rearrangement [22], and ene reaction of 2 with trans-hex-3-ene
(10) at 20°C [23] and enthalpies of dissolution of 2 at 25°C (ΔsolnH2, kJ/mol) [11] in different solvents
Solvent
2+1
2+5
2+6
2+7
2+8
2+9
2+10
∆solnH2
Tetrahydrofuran
Ethyl acetate
1,4-Dioxane
–
0.0382
0.0680
0.0910
0.2820
0.9160
0.4170
0.9900
3.6130
6.0800
–
–
0.060
0.109
0.150
0.640
0.570
1.260
2.560
4.900
14.3000
8.100
0.077
0.106
0.028
0.057
0.094
0.330
0.320
0.520
1.010
1.550
5.090
–
0.327
0.312
–
0.0032
06.7
09.0
08.8
18.3
14.2
20.6
21.8
21.9
24.4
–
0.00191
0.00232
0.02360
0.11100
0.00909
–
0.014
–
0.0046
–
Toluene
–
0.775
0.623
–
–
Acetonitrile
0.043
0.270
–
–
Benzene
0.0354
Chlorobenzene
1,2-Dichloroethane
Chloroform
–
–
0.57700
1.78000
–
0.120
0.180
0.130
–
–
0.1440
3.050
–
–
–
–
–
Methylene chloride
Dimethylformamide
Acetone
0.01700
0.00929
0.1620
–
–
0.453
–
0.014
–
–
(–81.5 kJ/mol) and 1,2-dichloroethane (–86.3±0.8 kJ×
mol–1) should be noted. The ratio of the entropies of
activation and reaction in toluene (–198/–145 = 1.36)
is consistent with the ratio of the activation and reac-
tion volumes (–50.8/‒37.3 = 1.36±0.06). The calculat-
ed equilibrium constant (4.63×106) and rate constant
for the reverse reaction (k1 = 3.78×10–9 s–1) at 25°C
correspond to almost complete transformation of 1
and 2 into 3.
reaction of 2 with 2-chloroethyl vinyl ether 6 (R =
0.98), [4π+2π]- and [2π+2σ+2σ]-cycloadditions, and
ene reaction. As follows from the slopes of depen-
dences (1)‒(6), the change of the lnk(2+1) value for
the reaction 1 + 2 → 3 in the examined solvent series
is almost twice as large as those for the other reactions,
which indicates stronger solvent effect on that reaction.
However, this variation is not related to the solvent
polarity. The rates of the reaction 1 + 2 → 3 in di-
methylformamide and acetone are comparable with the
rates in benzene and toluene. Change of the enthalpy
of solvation of 2 is not the only factor responsible for
the variation of the rate of the reaction 1 + 2 → 3 (7).
We failed to obtain [2+2]-cycloaddition product
from biadamantylidene (1) and tetracyanoethylene (4).
The absorbance of the complex of 4 with toluene in
a mixture with 1 in 1,2-dichloroethane with addition of
toluene did not change over 2 days. However, it is
known that 9,10-dimethoxyanthracene reacts with
compound 4 at a fairly high rate, but the reaction stops
when equilibrium establishes in the system [10]. The
available data on the enthalpies of reactions of various
dienes with compounds 2 and 4 indicate that the
reactions with 4 are less exothermic (by 20‒30 kJ/mol)
than those with 2 [24]. Assuming ∆Hr = –56 kJ/mol
and ∆Sr = –145 J mol–1 K–1 for the reaction of 1 with 4,
the conversion should exceed 80‒90%. Thus, invari-
ance of the absorbance of a solution of 1+4 suggests
very low reactivity of 4 toward 1.
lnk(2+5) = (0.6464±0.0642)lnk(2+1) + (1.4277±0.2628);
(1)
lnk(2+6) = (0.4071±0.0540)lnk(2+1) ‒ (2.0364±0.1968);
(1) (2)
R2 = 0.9500, N = 5;
lnk(2+7) = (0.6665±0.1527)lnk(2+1) + (1.742±0.6358);
(1) (3)
R2 = 0.7313, N = 9;
lnk(2+8) = (0.5386±0.0982)lnk(2+1) + (0.8989±0.4024);
(1) (4)
R2 = 0.8575, N = 7;
lnk(2+9) = (0.3204±0.0709)lnk(2+1) + (0.6923±0.2747);
(1) (5)
R2 = 0.8718, N = 5;
lnk(2+10) = (0.5485±0.2081)lnk(2+1) ‒ (1.4489±0.9430);
(1) (6)
R2 = 0.8742, N = 3;
lnk(2+1) = (0.3442±0.1133)∆solnH2 ‒ (9.0405±2.0060);
(1) (7)
R2 = 0.6486, N = 7.
R2 = 0.9442, N = 8;
(1)
The data in Table 3 show the effect of solvents on
the rate of reactions (1)‒(6) involving compound 2.
There is proportionality in the solvent effects on the
rates of the [2π+2π]-cycloaddition 1 + 2 → 3, [2π+2π]-
RUSSIAN JOURNAL OF ORGANIC CHEMISTRY Vol. 53 No. 12 2017

KISELEV et al.
1868
V≠corr = ∆V≠exp + βTRT,
(9)
EXPERIMENTAL
where βT is the isothermal compressibility coefficient
of the solvent.
Biadamantylidene 1 was synthesized from adaman-
tan-2-one according to [25]; it was recrystallized from
ethanol and additionally purified by alumina column
chromatography using hexane as eluent; mp 181‒
183°C [25]. 4-Phenyl-3H-1,2,4-triazole-3,5(4H)-dione
(2, 97%; from Aldrich, Germany) was purified by
sublimation at 100°C (100 Pa), mp 165–170°C
(decomp.) [26]). The purity of 2 was checked by
spectrophotometry [12]. Adduct 3 was isolated in
quantitative yield, mp 228‒230°C (decomp.) [13, 14]);
The ratio of the activation and reaction volumes
∆V≠corr/∆Vr was determined by the kinetic method, ac-
cording to which the volume of reaction was calculated
from the dependence of the density of solution on the
concentration of adduct 3. For the reaction 1 + 2 → 3,
the overall volume of a solution containing initial
reactants 1 and 2 and adduct 3 may be expressed by
Eqs. (10) and (11):
1
its structure was confirmed X-ray analysis and H and
13C NMR [13, 14]. All solvents were purified by
known methods [15].
Vτ = Vs + (c01 – c3,τ)V1 + (c02 – c3,τ)V2 + c3,τ V3;
Vτ = [Vs + (c01V1 + c02 V2)] + c3,τ (V3 – V1 – V2)
= Vτ=0 + c3,τ ∆Vr;
(10)
(11)
(12)
Kinetic measurements. The progress of the reac-
tion was monitored following variation of the absor-
bance of 2 at λ 530‒550 nm (Hitachi-2900 spectro-
photometer, Japan). A quartz cell containing 3 mL of
a solution of 1 was kept in the cell compartment until
temperature equilibration (the temperature was main-
tained with an accuracy of ±0.1°C), and 100‒150 μL
of a solution of 2 with a required concentration was
injected. The working concentration of 2 was (4–6)×
10–3 M, which was 3‒12 times lower than the concen-
tration of 1, to minimize temperature fluctuations of
the reaction mixture.
1/dτ = 1/dτ=0 + c3,τ ∆Vr–n/1000dτ=0
.
Equation (12) derived from (11) by division by the
solution weight is more convenient for measuring the
density of the reaction mixture during the reaction.
Here, Vτ=0 and Vτ are the volumes of the mixture at the
initial moment and time τ; Vs is the solvent volume; V1,
V2, and V3 are partial molar volumes of compounds 1,
2, and 3, respectively; c01, c02, and c3,τ are the initial
molar concentrations of 1 and 2 and the current
concentration of 3; and ∆Vr is the volume of reaction.
The current concentration of 3 was calculated from
the kinetic data. The current densities of the reaction
mixture were measured with an Anton Paar DSA
5000M precision densimeter (Austria) at 25±0.002°C
with an accuracy of ±2×10–6 g/cm3.
The stability of compound 2 in all solvents used
was checked by measuring variation of its absorbance
duering the reaction time. The concentration of 2 in
DMF and acetone during the reaction time decreased
by 20 and 10%, respectively, which was taken into
account while calculating the rate constants. The
relative standard deviations for the rate constants were
within ±3%, for the enthalpy of activation, ±2 kJ/mol,
and for the entropy of activation, ±6 J mol–1 K–1.
The enthalpy of the reaction 1 + 2 → 3 in 1,2-di-
chloroethane was determined at 25°C using a high-
precision TAM III calorimeter with a 25-mL glass
isoperibolic calorimetric cell equipped with a stirrer,
heater, and thermistor. The cell was charged with a so-
lution of 1 (0.0126 M); after temperature equilibration,
it was calibrated by supplying a known amount of heat
through the heater. A solution of 2 (0.1445 M) was
then added in 60‒80-μL portions using an automatic
dosing unit. Each next portion was added after a time
exceeding the time necessary for the reaction com-
pletion by more than 99.9%. The heat effect of the
addition of a 0.145 M solution of 2 into pure 1,2-di-
chloroethane was measured in a separate experiment; it
proved to be negligible. Successive measurements
gave the following enthalpies of the reaction 1 + 2 → 3
in 1,2-dichloroethane: ‒85.2, ‒86.0, ‒85.5, ‒87.4,
‒87.0, ‒86.0, and –86.7 kJ/mol; average value
–86.3±0.8 kJ/mol.
The reaction 1 + 2 → 3 at elevated pressure was
carried out in toluene and ethyl acetate at 25°C using
an HP-500 pressure regulator (Japan), a PCI-500
variable-volume quartz cell (Japan), and a SCINCO
program-controlled spectrophotometer (Korea). The
observed volume of activation (∆V≠exp) for the reaction
1 + 2 → 3 in both solvents was calculated from the
rate constants at 1 and 1000 bar using Eq. (8) [27]:
[lnkp/p](1 bar) = (1.15±0.03)ln[k(1000 bar)/k(1 bar)]×10–3.
(8)
The volume of activation was corrected for solvent
compressibility [19]:
RUSSIAN JOURNAL OF ORGANIC CHEMISTRY Vol. 53 No. 12 2017

[2π+2π]-CYCLOADDITION OF BIADAMANTYLIDENE
1869
12. Kiselev, V.D., Kornilov, D.A., and Konovalov, A.I., Int.
J. Chem. Kinet., 2017, vol. 49, no. 8, p. 562. doi
10.1002/kin.21094
13. Cheng, C.C., Seymour, C.A., Petti, M.A., Greene, F.D.,
and Blount, J.F., J. Org. Chem., 1984, vol. 49, p. 2910.
14. Seymour, C.A. and Greene, F.D., J. Am. Chem. Soc.,
Equilibrium constants. A low-melting tube made
of molybdenum glass was soldered to a quartz cell.
After loading a solution of compounds 1 and 2 (c01 =
c02 = 0.01317 M) in toluene, the tube was sealed. The
equilibrium constants were calculated from the absor-
bance of 2 after heating the cell for 3‒4 h in the dark at
70 and 80°C (molar absorption coefficient of 2 at
λ 540 nm ε = 245 L mol–1 cm–1).
1980, vol. 102, p. 6384.
15. Riddick, J.A., Bunger, W.B., and Sakano, T.K., Organic
Solvents: Physical Properties and Methods of Purifica-
tion, New York: Wiley, 1986, 4th ed.
This study was performed under financial support
by the Russian Foundation for Basic Research (project
no. 16-03-00071) and by the Program of the Russian
Federation for the competitive development of the
Kazan Federal University.
16. Kornilov, D.A. and Kiselev, V.D., J. Chem. Eng. Data,
2015, vol. 60, p. 3571.
17. Kiselev, V.D., Bolotov, A.V., Satonin, A.P., Shakiro-
va, I.I., Kashaeva, H.A., and Konovalov, A.I., J. Phys.
Chem. B, 2008, vol. 112, p. 6674.
18. Kiselev, V.D., Kornilov, D.A., Anikin, O.V., Latypo-
va, L.I., Bermeshev, M.V., Chapala, P.P., and Konova-
lov, A.I., Russ. J. Org. Chem., 2016, vol. 52, p. 777.
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