Entropy and enthalpy effects of 4-(phenylthio)-substituted phenols

Article abstract of DOI:10.1039/a707856k

Electron paramagnetic monitoring of free phenoxyl radicals in benzene-toluene (1:1) solution has been used to study the temperature dependence of the equilibrium (a), for which the equilibrium constant, K₁ is (formula presented) R¹ Bu^t; R² = S-(3-methyl-4-hydroxy-5-tert-butylphenyl), R³ = CH₃ given by eqn. (b). (formula presented) The fit of experimental data to eqn. (c) yields ΔH₁ = 28.80 ± 3.39 kJ mol^-1 and ΔS₁ = 116.9 ± 12.7 J K^-1 In K₁ = (-ΔH/RT) + (ΔS/R) (c) mol^-1. The entropy change value is about 30 J KT^-1 mol^-1 in excess of the absolute entropy of the molecular rotor evolving about each of the C-S bonds in the thiobis(phenol) 1, thus denoting the enhancement by solute-solvent interactions. As well as using the enthalpy change to account for the stability of the generated phenoxyls, and/or the ability of hydrogen cession by phenols, the entropy change must also be measured when evaluating the antioxidant behaviour of this type of molecule, as its influence on the Gibbs free energy change is shown to be relevant.

Full text of DOI:10.1039/a707856k

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Entropy and enthalpy effects of 4-(phenylthio)-substituted phenols
Daniela E. Paulon and Marta E. J. Coronel*
Departamento de Quimica, Facultad de Ciencias Exactas y Naturales, Universidad Nacional de
Mar del Plata, Funes 3350, 7600 Mar del Plata, Argentina
Electron paramagnetic monitoring of free phenoxyl radicals in benzene–toluene (1:1) solution has been
used to study the temperature dependence of the equilibrium (a), for which the equilibrium constant, K1 is
OH
O•
O•
OH
R1
R3
R1
R1
R1
R3
R1
R1
(a)
+
+
R2
R3
R2
R3
1
2
3
4
R
1 = But; R2 = S-(3-methyl-4-hydroxy-5-tert-butylphenyl), R3 = CH3
given by eqn. (b).
[3][4]
K1 ؍
(b)
[1][2]
The fit of experimental data to eqn. (c) yields ÄH1 ؍ 28.80 3.39 kJ mol؊1 and ÄS1 ؍ 116.9 12.7 J K؊1
ln K1 ؍ (؊ÄH/RT) ؉ (ÄS/R) (c)
mol؊1. The entropy change value is about 30 J K؊1 mol؊1 in excess of the absolute entropy of the molecular
rotor evolving about each of the C᎐S bonds in the thiobis(phenol) 1, thus denoting the enhancement by
solute–solvent interactions. As well as using the enthalpy change to account for the stability of the
generated phenoxyls, and/or the ability of hydrogen cession by phenols, the entropy change must also be
measured when evaluating the antioxidant behaviour of this type of molecule, as its influence on the
Gibbs free energy change is shown to be relevant.
O•
OH
Introduction
R1
R1
R1
R3
The antioxidant strength of phenols has been largely correlated
with phenoxyl radical stability and evaluated by the phenolic
O᎐H bond dissociation enthalpy [D(O᎐H)].1 The resonance stabil-
ization of the unpaired electron is mainly determined by the
type and location of the substituents on the aromatic ring. The
presence of an oxygen atom connecting a para-alkyl group
brings maximal antioxidant character when the whole molecule
is a chromane ring system, as in the tocopherols.1–8 However, in
non-rigid phenolic molecules having internal rotation about the
C4᎐O bond, an entropy effect should be expected. This feature
has been the subject of a previous work from our laboratory,
where the cost of the entropy loss was estimated for a sterically
hindered 4-methoxyphenoxyl radical.7 Those results indicated
that the origin of the entropic effect can be, not only the loss of
degrees of freedom for that motion, but also solvent ordering
surrounding the radical.
The present study has been to investigate the influence of a
sulfur containing substituent, because of the availability of a
class of such antioxidants, of which phenol 1 (synonyms Thio-
alkofen BM, Antioxidant 736) is a good example. The presence
of a 4-(phenylthio) substituent (4-PTS) allows us additional
insight into the molecular factors that can control the thermo-
chemistry of the hydrogen transfer to phenoxyl radicals.
For that purpose, equilibrium (1) was monitored by means of
electron paramagnetic resonance (EPR) in its equilibrium state
and K1 values were determined at varying temperature.
+
R3
R2
2
1
(1)
O•
OH
R1
R3
R1
R1
+
R2
R3
3
4
R
1 = But; R2 = S-(3-methyl-4-hydroxy-5-tert-butylphenyl),
R3 = CH3
Experimental
Reaction (1) was carried out in benzene–toluene (1:1) solvent
mixtures in order to obtain stable solutions along the temper-
ature range of K1 measurements. Radicals were generated in the
EPR sampling tubes by adding a known concentration of
galvinoxyl radical [G; 2,6-di-tert-butyl-α-(3,5-di-tert-butyl-4-
oxohexa-2,5-dien-1-ylidene)-p-tolyloxy free radical: Aldrich;
J. Chem. Soc., Perkin Trans. 2, 1998
885

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Table 1 Variation in equilibrium constant and Gibbs’ free energy
values with change in temperature for equilibrium (1)
a
T/K
ln K1
σ2 b
∆G/kJ molϪ1 c
243.5
253.5
263.5
273.5
283.5
293.5
0.075(6)
0.175(6)
0.864(8)
1.39(13)
1.62(6)
0.04
0.14
0.26
0.28
0.06
0.01
Ϫ0.15 0.12
Ϫ0.37 0.13
Ϫ1.89 0.14
Ϫ3.16 0.09
Ϫ3.82 0.10
Ϫ6.16 0.04
2.52(5)
a Number of measurements given in parentheses. b Variance of ln K1.
c Mean standard error.
Fig. 2 1H NMR spectrum of: (a) 1, δH 1.3675 (s, 9 H, o-But), 2.1845
(s, 3 H, o-Me), 4.7460 (s, 1 H, aryl-OH), 6.9886 and 6.9975 (poorly
resolved d, 1 H, m-aryl-H), 7.1621 and 7.1720 (poorly resolved d,
1 H, m-aryl-H), 7.2466 (s, Cl3CH impurity in Cl3CD); (b) MOPH,
δ 1.4323 (s, 18 H, 2,6-di-But), 3.7637 (s, 3 H, p-OMe), 4.7488 (s, 1 H,
aryl-OH), 6.7592 (s, 2 H, two m-aryl-H), 7.2467 (s, Cl3CH impurity in
Cl3CD)
Fig. 1 X-band first derivative EPR spectrum of: (a) 2 in a benzene–
toluene (1:1) solution of 4 (0.137 mol dmϪ3) and galvinoxyl (0.2 mmol
dmϪ3); (b) similar condition EPR spectrum for 3; (c) EPR spectrum of
the solution resulting from the addition of 1 (0.85 mol dmϪ3) to the
mixture described in (a)
them by EPR. Their slow decay in the equilibria follows the
kinetics reported for 28 which means that the hydrogen transfer
is very fast with respect to the lifetime of the radicals. Hence,
[2] ϩ [3] = [G0], the concentration of galvinoxyl radical added
to the parent phenols in generating the reaction system. [2] was
measured in the presence and in the absence of 1 to find its
concentration in the equilibria, and this value was subtracted
from [G0] to give [3]. [G0] was obtained by direct weighing of
highly pure G. This device avoids absolute calibrations with
external standards.
Fig. 1(a), (b) and (c) shows the EPR records of 2, 3 and their
combined radical equilibrium, respectively. The values of ln K1
and ∆G1 at different temperatures are gathered in Table 1.
The fit of experimental data to eqn. (2) yielded a = Ϫ3464
2 × 10Ϫ4 ] to solutions of the parent phenols. Samples, 120 µl
in volume, were frozen and degassed by vacuum pumping up
to 5 kPa to ensure constant pressure in the experiments. EPR
spectra were recorded at fixed periods of time after thawing in
the thermostatted cavity of a Bruker ER-200 X-band spec-
trometer equipped with a B-VT 1000 variable temperature unit.
Peak-to-peak amplitudes of the first derivative spectral lines in
the lowest field triplet of 2 [Fig. 1(a)] were measured in the
absence of 1 (blank samples) and in the presence of 1 (reaction
mixtures) to evaluate every radical concentration.
Proton NMR (200 MHz, in deuterochloroform) and FTIR
(in nujol dispersion) were recorded on a Bruker AC-200 NMR
and on a Nicolet Magna FTIR spectrometer, respectively, at the
spectroscopy service unit UMYMFOR (CONICET-FCEN,
University of Buenos Aires).
Phenols 1 (Antioxidant Kit, Chem Service, Illinois, USA)
and 4 (Aldrich) were crystallized from hexane and their purity
checked by TLC. Benzene and toluene (Merck, Analytical
Grade) were dried on potassium carbonate and distilled before
use.
ln K1 = aTϪ1 ϩ b
(2)
4.08 and b = 14.06 1.53 (mean standard error) from which
we can derive ∆H1 = 28.80 3.39 kJ molϪ1 and ∆S1 = 116.9
12.7 J KϪ1 molϪ1. By knowing ∆H = 334.1 kJ molϪ1 for the
reaction 4 = 2 ϩ H, previously obtained against the absolute
calorimetric value D(O᎐H) = 340.1 kJ molϪ1 in 2,4,6-tri-tert-
butylphenol,6,9 we obtain D(O᎐H) = 362.9 3.2 kJ molϪ1 for
phenol 1.
The present results indicate that an enthalpic stabilization
has been introduced into the O᎐H of phenol 1 by the 4-PTS
with respect to the 4-methyl analog, although the entropy
change has made reaction (1) spontaneous.
Results and discussion
The persistence of radicals 2 and 3 makes it possible to monitor
886
J. Chem. Soc., Perkin Trans. 2, 1998

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Fig. 4 X-band first derivative EPR spectrum of MOP: an overlapped
quartet of triplets (see the text for hyperfine interactions)
This explanation agrees well with the difference in pKa values
of para-anisidine (5.34) and 4-methylthioaniline (4.35);12 those
values mean an easier heterolytic cleavage of the phenolic O᎐H
bond, as opposed to a more difficult homolytic dissociation,
with sulfur, rather than oxygen, at C4. In other words, this is a
good picture of the change in stability of the anions compared
with the corresponding radicals. Those differences can be
ascribed to the electron-acceptor character of the sulfur-
containing substituent when facing a stronger electron donor
such as the phenolic O᎐H group.
Hence, the high entropy gain in reaction (1) becomes under-
standable on the basis of the molecular polarization borne by 1
that is lost in its daughter radical. As a matter of fact, radical 3
can account for as much as 84 J KϪ1 molϪ1 (the absolute
entropy of the free rotation about each of the C᎐S bonds13) plus
a little more than 30 J KϪ1 molϪ1 in excess. As long as we are
dealing with reactions in liquids that excess of entropy gain
involves the loss or the decrease of the solvent ordering built up
by the parent phenol.
It is likely that the dipole strength of MOPH is not compar-
able to compound 1, mainly because of the better balanced
electron affinity of the groups located on the opposite sides of
the aromatic ring. Actually we had previously found the
enthalpy and entropy change values for reaction (3)7 as follows:
Fig. 3 Infrared (FTIR) spectrum of (a): 1; (b): MOPH
MOPH ϩ 2
MOP ϩ 4
(3)
Proton NMR [Fig. 2(a) and (b)] and IR [Fig. 3(a) and (b)]
spectra were recorded for phenol 1 and 2,6-di-tert-butyl-4-
methoxyphenol (MOPH) for comparative purposes.
∆H3 = Ϫ12.40 kJ molϪ1 and ∆S3 = Ϫ18.47 J KϪ1 molϪ1.
As a comment, we dare say that the quantities of the
enthalpy and entropy change for reaction (1) fall into the
correlation found for Frank and Evans for the vaporization
of inert gases from water: ∆S = 14 ϩ 0.0019 ∆H,14 a fact that
told us about entropy driven equilibria. Since antioxidants are
sometimes added to food or pharmaceuticals to preserve them
from oxidation, it should be emphasized that only the toco-
pherols (in their natural sources and amounts) could be safe
enough for avoiding undesirable effects on hydrogen transfer
equilibria.
At this point, it seems appropriate to mention that the
presence of steric hindrance due to the ortho-substituents
in phenols of reaction (1) cancels only partially the repulsion
cooperating with the thermochemical stabilization of the
phenoxyl radicals.15
The NMR spectrum of compound 1 shows two signals for
the meta-aryl protons as is expected for the different alkyl sub-
stituents on the aromatic ring, both protons having chemical
shifts above δ 7, while the singlet due to the same protons in
MOPH is located at δ 6.759 [see captions to Fig. 2(a) and (b)].
Those data clearly indicate that the para-methoxy exerts a
greater diamagnetic screening over the neighbouring sites of the
molecule, a fact that correlates well with a stronger electron
release ability towards the aromatic ring.
On the other hand, the infrared spectrum of MOPH exhibits
a band at 3589.4 cmϪ1 for the O᎐H stretching, whereas that of
1 shows two distinct absorptions which can be assigned to
the same vibration, 3550.8 and 3410.0 cmϪ1; the latter is fairly
broad and indicates hydrogen bonding to some extent, but the
former shows a shift to a lower frequency, which suggests elec-
tron release towards the 4-PTS group,10 a feature attributable
to the low lying 3d orbitals of the sulfur atom.11 As is often
suggested, the sulfur atom can use its low energy unoccupied 3d
orbitals to partially form double bonds, while maintaining
a low bond angle value, and contribute to delocalizing the
negative charge. This fact, combined with the electron donor
character of the phenolic moiety, enhances the contribution of
canonical structures such as 5.
The EPR spectra of the phenoxyl radicals 2 [Fig. 1(a)], 3
[Fig. 1(b)] and MOP (Fig. 4) have the basis structure of a
quartet of triplets, although this is only evident for 2 and 3 due
to the relationship between the values of the meta-ring proton
hyperfine interactions and those of the hydrogen nuclei in the
para- or ortho-methyl substituents on the phenoxyl: a(m-H)
=
0.169, a(p-Me) = 1.125 mT for 2;16 a(m-H) = 0.124, a(o-Me) = 0.571
17
mT for 3;
a
= 0.086, a(p-OMe) = 0.162 mT for MOP.18 All the
(m-H)
EPR spectra show equivalence for both the methyl and the two
meta-ring protons. As we can see in the NMR spectrum of 1, its
meta-ring hydrogen atoms have different chemical shifts. It may
be that such a difference is lost in 3, its daughter phenoxyl, but it
is also possible that this small gap in magnetic moments is not
enough to be displayed as distinct hyperfine interactions and
only contributes to the line width in its radical EPR spectrum.
We still can derive some additional ideas from our previous6,7
and present studies with these reaction systems. Hydrogen
+OH
R1
R3
–S
C
14H21O
5
J. Chem. Soc., Perkin Trans. 2, 1998
887

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transfer reactions have been extensively studied as a model for
the correlation of reactivity, selectivity and structure in free
radicals.1–5,19 Though absolute rate constants are actual estim-
ations of phenol reactivity, small changes in selectivity between
substrates may be made more apparent by evaluating the com-
petitive hydrogen abstraction equilibria with a couple of radicals
offering rather similar attachment and environment to the
exchangeable hydrogen atom. The hindered phenoxyls used
as substrates in our studies come from simple molecules
and, besides allowing easy recognition and measurement, have
their main structure differences confined to the para-
substituent, whose influence on thermochemical parameters
becomes straightforward. In addition, if the transition state
may be regarded as a hydrogen bonded complex, then the O᎐H
stretching vibration can be taken as the reaction coordinate.19
Several studies about the deuterium effect have been performed,
among which we should quote those by K. U. Ingold and co-
workers, who found a good deal of isotope effect for the hydro-
gen transfer from α-tocopherol.5
References
1 J. A. Howard, in Free Radicals, ed. J. K. Kochi, Wiley Interscience,
New York, 1973, vol. II, ch. 12.
2 J. A. Howard and K. U. Ingold, Can. J. Chem., 1963, 41, 1744.
3 J. A. Howard and K. U. Ingold, Can. J. Chem., 1964, 42, 1044.
4 J. A. Howard and E. Furimsky, Can. J. Chem., 1973, 51, 3378.
5 G. W. Burton, T. Doba, L. Hughes, F. L. Lee, L. Prasad and
K. U. Ingold, J. Am. Chem. Soc., 1985, 107, 7053.
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7 M. E. J. Coronel and A. J. Colussi, J. Chem. Soc., Perkin Trans. 2,
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8 K. U. Ingold, in Free Radicals, ed. J. K. Kochi, Wiley Interscience,
New York, 1973, vol. I, ch. 2, p. 3.
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Soc., 1989, 91, 3883.
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Infrared and Raman Spectroscopy, Academic Press, New York, 1990,
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11 J. E. Bennet, B. Mile, A. Thomas and B. Ward, in Advances in
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1970, vol. 8, pp. 38–39.
12 CRC Handbook of Chemistry and Physics, 75th edn., 1994–1995;
section 8–52; Chief Editor, D. H. Lide, Associate Editor, H. P. R.
Frederiske, CRC Press, Boca Raton, 1994.
In such a case fast, diffusion controlled, hydrogen transfers
are expected20 between phenolic substrates whose D(O᎐H) values
range within 42.5 kJ molϪ1, the energy gap from the ground to
the first excited vibrational level to a typical wavenumber about
3500 cmϪ1. For those systems hydrogen transfer rate constants
should be inherently independent of the equilibrium constant
values.
13 S. W. Benson, Thermochemical Kinetics, Wiley, New York, 1976,
2nd edn., ch. 2.
14 H. S. Frank and M. W. Evans, J. Chem. Phys., 1945, 13, 507.
15 D. D. M. Wayner, E. Lusztyk, K. U. Ingold and P. Mulder, J. Org.
Chem., 1996, 61, 6430.
We may conclude that although potential antioxidant reac-
tivities of phenols can be well correlated with the stability of the
phenoxyls generated, and both entities judged by means of the
enthalpy change of the hydrogen transfer, entropy effects
should not be neglected. They are indicative of polar influences
which, together with resonance stabilization, are the origin of
unexpected changes in selectivity, for reactions taking place
in liquids. Moreover, the amount of the entropy change might
be correlated with the average polarity of solvent mixtures,
an aspect that we think deserves further study.
16 The a(m-H) value reported in ref. 7 was mistaken.
17 Data from our laboratory.
18 J. Pannell, Chem. Ind., 1962, 1797.
19 G. A. Russell, in Free Radicals, ed. J. K. Kochi, Wiley Interscience,
New York, 1973, vol. I, ch. 7, pp. 275–331.
20 L. Valgimigll, J. T. Banks, K. U. Ingold and J. Lusztyk, J. Am.
Chem. Soc., 1995, 117, 9966.
Acknowledgements
The authors thank Dr Norma S. Nudelman for reviewing the
manuscript and Drs Keith U. Ingold and Carlos G. Diaz for
helpful comments.
Paper 7/07856K
Received 31st October 1997
Accepted 16th January 1998
888
J. Chem. Soc., Perkin Trans. 2, 1998