Elucidation of concentrated salt effects on the solvolysis reactions of typical SN1 and SN2 substrates in sulfolane-water mixed solvent

Article abstract of DOI:10.1246/bcsj.79.751

In a 50 vol % sulfolane-H₂O mixed solvent, extensive studies have been carried out of the concentrated salt effects on the solvolysis reaction rates of aliphatic halides and related compounds (RX). In the solvent system, the "pseudo" first-order rate constants (k/s^-1) of typical S_N1 substrates, 1-adamantyl chloride and bromide, or 2-adamantyl bromide, were increased exponentially by the addition of MClO ₄ (M = Li and Na) and M(ClO₄)₂ (M = Mg and Ba); the extent of the cation effects increased as Na⁺ < Li ⁺ < Mg²⁺ ≈ Ba²⁺, which was attributed to the "chemical" interaction between the metal ions and the leaving-group anions in the "aqueous" solution modified by the mixed-organic solvents and concentrated salts. Based on the Raman spectra, we discuss the distortion of the bulk water structure and the alteration of properties of water into those of a "non-aqueous" solvent with the addition of the organic solvents and concentrated salts (LiClO₄ and Et₄NBr) as well as with increasing temperature. However, the addition of non-metallic salts, i.e., Et₄NX (X^- = ClO ₄^-, Cl^-, Br^-, and tosylate), caused decreases in the rates substantially, regardless of the leaving group (Cl or Br): the deceleration effects increasing as ClO₄^- < Br^- < Cl^- < TsO^-. The deceleration in the solvolysis rates with Et₄NX and R₄NBr (R = Pr and n-Bu) should be caused by mainly the further decrease in water activity of the solvents. In addition, we would like to report solvolyses accompanied by "ion exchange" reactions for the typical S_N2 substrates, hexyl chloride, bromide, and tosylate, in the sulfolane-H₂O solvent.

Full text of DOI:10.1246/bcsj.79.751

Ó 2006 The Chemical Society of Japan
Bull. Chem. Soc. Jpn. Vol. 79, No. 5, 751–760 (2006)
751
Elucidation of Concentrated Salt Effects on the Solvolysis Reactions
of Typical S 1 and S 2 Substrates in Sulfolane–Water Mixed Solvent
N
N
ꢀ
Masashi Hojo, Tadaharu Ueda, Eriko Ueno, Tensei Hamasaki, and Daigo Fujimura
Department of Chemistry, Faculty of Science, Kochi University, 2-5-1 Akebono-cho, Kochi 780-8520
Received May 13, 2005; E-mail: mhojo@cc.kochi-u.ac.jp
In a 50 vol % sulfolane–H2O mixed solvent, extensive studies have been carried out of the concentrated salt effects
on the solvolysis reaction rates of aliphatic halides and related compounds (RX). In the solvent system, the ‘‘pseudo’’
first-order rate constants (k/s 1) of typical SN1 substrates, 1-adamantyl chloride and bromide, or 2-adamantyl bromide,
ꢁ
were increased exponentially by the addition of MClO4 (M ¼ Li and Na) and M(ClO4)2 (M ¼ Mg and Ba); the extent of
þ
þ
2þ
2þ
the cation effects increased as Na < Li < Mg ꢂ Ba , which was attributed to the ‘‘chemical’’ interaction between
the metal ions and the leaving-group anions in the ‘‘aqueous’’ solution modified by the mixed-organic solvents and con-
centrated salts. Based on the Raman spectra, we discuss the distortion of the bulk water structure and the alteration of
properties of water into those of a ‘‘non-aqueous’’ solvent with the addition of the organic solvents and concentrated salts
(
(
(
LiClO4 and Et4NBr) as well as with increasing temperature. However, the addition of non-metallic salts, i.e., Et4NX
ꢁ
ꢁ ꢁ ꢁ
X
= ClO4 , Cl , Br , and tosylate), caused decreases in the rates substantially, regardless of the leaving group
ꢁ
ꢁ
ꢁ
ꢁ
Cl or Br): the deceleration effects increasing as ClO4 < Br < Cl < TsO . The deceleration in the solvolysis rates
with Et4NX and R4NBr (R ¼ Pr and n-Bu) should be caused by mainly the further decrease in water activity of the sol-
vents. In addition, we would like to report solvolyses accompanied by ‘‘ion exchange’’ reactions for the typical SN2 sub-
strates, hexyl chloride, bromide, and tosylate, in the sulfolane–H2O solvent.
ꢁ
ꢁ
ꢁ
ꢁ
Electrolyte effects on reaction rates are believed to play an
important role in the investigation of mechanisms. More than
TsO ꢂ NO3ꢁ ꢂ Br ꢂ Cl ꢂ no salt > Fꢁ > OH . How-
9
ever, Allen et al. have reported a great difference in the salt
1
4
0 years ago, however, Clarke and Taft described that the val-
effects of LiClO4 and n-Bu4NClO4 on the solvolysis of 4-me-
thoxybenzyl chloride in 85%(v/v) aqueous dioxane. The ef-
fects of LiClO4 and LiCl in several organic solvents and water
on the stereo-selectivity of products of the Diels–Alder reac-
tion have been discussed:10 the salt effects on Diels–Alder
idity of the theoretical treatment of salts had remained uncer-
tain; a similar situation may have lasted until nowadays. In
general, ‘‘indifferent’’ salts (or supporting electrolytes) are
never involved directly in the equilibria or kinetics of reactions
in solution. The effects by ‘‘indifferent’’ salts of lower concen-
11
reaction kinetics have been reviewed.
12–14
ꢁ3
trations (<0:1 mol dm ) in aqueous solution may be explain-
ed by the Debye–H u¨ ckel theory, which takes the electrostatic
interaction into account. A corrected Debye–H u¨ ckel theory
with size asymmetry and effective diameters has been dis-
In previous studies, we have observed that for a typical
SN1 substrate, such as 1-adamantyl halides or t-butyl chloride,
þ
þ
þ
2þ
2þ
cation effects increase as Et4N ꢃ Li , Na < Mg , Ba
12
for perchlorate salts in 80 vol % MeOH–H2O ("r ¼ 49),
2
13
cussed.
50 vol % acetone–H2O ("r ¼ 55), and 50 vol % 1,4-dioxane–
14
Numerous studies have been performed to understand the
3
H2O ("r ¼ 44) solvent systems. The cation effects were
attributed to the interaction between the metal ions and the
leaving-group anions of substrates. For the incredibly ‘‘strong’’
interaction between the metal ions and the simple anions in
salt effects on solvolysis reactions; the Winstein’s ion-pair
4
5
scheme has had strong impact in this field. A paper appeared
entitled ‘‘Correlation of Solvolysis Rates 50 Years Later’’ after
6
12
the paper by Grunwald and Winstein of 1948. However, the
‘‘aqueous’’ solutions, we have proposed a model of water
containing highly concentrated salts. This model can be out-
lined as follows: in a highly concentrated salt solution (>5
mol dm ), too many anions and cations exist to be fully sol-
vated by H2O molecules (<55 mol dm ); under such extreme
conditions, the network of hydrogen bonding by the solvent
ꢁ3
effects by concentrated salts (>0:1 mol dm ) have not been
treated quantitatively and comprehensively. At higher ionic
strength (ꢀ), the reaction rate (k) can be expressed by
ꢁ3
ꢁ3
log k ¼ log k0 þ bꢀ;
ð1Þ
where k0 is the rate constant at infinite dilution (or ꢀ ¼ 0); b is
water is almost distorted to give isolated H O molecules.
2
7
an empirical coefficient. It is known that the rates of SN1 sol-
volysis of alkyl halides and related compounds are subject to
Isolated H O molecules are indeed very different from bulk
2
water. The donicity, for instance, of isolated H O molecules is
2
8
8
15
positive salt effects. Bunton et al. reported an exponential in-
crease in the solvolysis rate of t-butyl bromide with increasing
the concentration of LiClO4. They concluded that anion effects
very small, DN ¼ 18:0, whereas the donor number of bulk
1
6
17
water is reported to be DN
on the solution conditions. Erlich et al. have noted that
bulk
¼ 33 or 42.3, depending
16
þ
þ
23
are important but cation effects are small (for Li , Na , and
þ
Na NMR chemical shifts are concentration dependent for
NaI and NaSCN. It is supposed that the DNbulk value would
ꢁ
Et4N ), and that the order of the effects of anions is ClO4
>
Published on the web May 12, 2006; DOI: 10.1246/bcsj.79.751

7
52 Bull. Chem. Soc. Jpn. Vol. 79, No. 5 (2006)
Elucidation of Concentrated Salt Effects
approach DN ¼ 18:0 with increasing concentration of a so-
dium salt if its solubility in water had no limit. We may stress
that the donicity of ethers are very close (e.g., DN ¼ 19:2 for
solutions if they are in the ‘‘dihydrogen ether’’ ([R](H)–O–
(H)[R]) conditions. In order to support our proposal of previ-
1
4
ous solvolytic studies,1
2–14
we examined the effects of various
1
5
diethyl ether) to that of isolated H2O molecules.
types of salts, MClO4, M(ClO4)2, and R4NX, on the solvolysis
reactions of typical SN1 and SN2 substrates in a 50 vol % sulfo-
lane–H2O solvent. Coincidentally, it must be very important to
The main points proposed in our previous solvolytic stud-
ies12–14 are as follows: (1) Highly concentrated salts would re-
duce the properties of bulk water (H–O–H) into those of a non-
aqueous solvent, such as an ether (R–O–R). (2) In water-mixed
þ
2þ
distinguish the tetraalkylammonium ion (R4N ) possessing no
þ
coordination ability from metal ions (M or M ) and to distin-
guish the perchlorate ion (three-dimensionally symmetrical
ꢁ3
solvents containing concentrated metal salts (>1:0 mol dm ),
the exponential increase in solvolysis rates of SN1 substrates
are caused by the formation of carbocations through direct
ꢁ
ꢁ
ꢁ
anions, including PF6 and BF4 ) from other anions (X ), such
as halides or tosylate. During the conductometric studies in
ꢁ
20
þ
‘
‘chemical’’ interaction between metal ions and anions (X ,
aprotic solvents, we noticed that the cation of Et4N , and
not a metal ion, can form a strong electrolyte, regardless of
the kind of anions; similarly, the anion of ClO4 forms a strong
from the substrate RX) in the modified solvent. (3) The solvol-
ysis rates of SN2 substrates are decelerated by concentrated
salts because of the decrease in water activity. (4) Positive
or negative salt effects can be correlated to the m-value (by
Grunwald and Winstein)18 of the substrates.
ꢁ
electrolyte even though the counter cation is a metal ion. This
principle in aprotic solvents stands even for concentrated salts
in aqueous–organic solvent systems, at least for exploring
kinetic phenomena.
The substitution (or ‘‘ion exchange’’) of leaving groups with
added salt anions for the typical SN2 substrates, hexyl chloride,
bromide, and tosylate, were examined thoroughly in the 50
vol % sulfolane–H2O solvent system.
By means of voltammetry,19 conductometry,20,21 and UV–
visible and NMR spectroscopies,22–27 we have undeniably
demonstrated that alkali metal and alkaline earth metal ions
have ‘‘weak’’ but unexpectedly distinct chemical (i.e., covalent
bonding or coordinating) as well as Coulombic interactions
with many simple anions. This holds true not only in low per-
21
Results and Discussion
mittivity media ("r < 10) but also in higher permittivity
where
media (20 < "r < 65) with poor solvating ability,1
9,20
Effects of Metallic and Non-Metallic Salts on the Sol-
volysis Rate of 1-Adamantyl Chloride.
the simple anions include halides, carboxylates, phosphates,
sulfonates, etc. Coordination chemistry of the lithium ion with
simple anions as well as cyclic ligands has been reviewed by
Figure 1 shows
the change in the solvolysis rate of 1-adamantyl chloride with
the addition of alkali metal or alkaline earth metal perchlorates
2
8
ꢄ
Olsher et al. Stable carbocations can be produced from trityl
þ
in a 50 vol % sulfolane–water mixed solvent at 55 C. The
ꢁ1
þ
halides by the addition of the perchlorates salts of Li , Na ,
2þ
‘‘pseudo’’ first-order reaction rate constant (k/s ) increased
exponentially with increasing concentration of LiClO4 or
2
þ
2þ
2þ
22–24
Mg , (Ca , Sr ), and Ba in acetonitrile;
the carbo-
cations have been detected by UV–visible, 1H and 13C NMR
spectroscopies.
22
NaClO4, i.e., linearity between logðk=s Þ and the concentra-
ꢁ1
2
3,24
For a fluoran-based black color former in
tion of MClO4 was observed. With Mg(ClO4)2 or Ba(ClO4)2,
although linearity was observed at lower salt concentrations,
acetonitrile, the color has been developed by the addition of
the alkali metal or alkaline earth metal perchlorates2 as well
as phenol which is used as a normal color developer.
4
the slope appeared to become lower somewhat at higher salt
þ
concentrations. The effects of cations increased as Na
<
We have explained the large salt effects on the indicator
acidity (or Hammett acidity function) using p-(dimethylami-
29
no)azobenzene and on the proton transfer from tropolone26
2
5
2
7
or nitrophenols to amine bases in acetonitrile solution in
terms of the chemical interaction between the anions and
þ
2þ
M or M , not merely as ion-exchange reactions or the alter-
nation between the contact ion pair (CIP) and the solvent-
27
30
separated ion pair (SSIP). Binder and Kreevoy examined
þ
the interaction of Li with a betaine dye, 2,6-diphenyl-4-
(2,4,6-triphenyl-1-pyridino)phenolate, and p-nitrophenolate in
acetonitrile. Very recently, we obtained decisive H NMR evi-
1
26
dence of the formation of a ‘‘reversed-coordinated’’ species,
þ
ꢁ
ꢁ
(
Li )2L , for the anions (L ) from tropolone and hinokitiol
4-isopropyltropolone) in acetonitrile-d3 but not in DMF-
(
d6. The extraordinary chemical reactions with alkali metal
and alkaline earth metal ions in higher permittivity media have
3
1
3
2
33
been reviewed. Fuoss and Kraus have proposed the concept
of ‘‘symmetrical’’ triple ion formation, which is based on elec-
trostatic interaction between (spherical) ions in an electric field
of low permittivity.
Fig. 1. Changes in the solvolysis rate constants of 1-ada-
mantyl chloride with the addition of various salts in 50
Further, we may postulate that a very small but direct
þ
2þ
ꢄ
‘‘chemical’’ interaction between M or M and simple anions,
such as halide, occurs even in aqueous or organic-aqueous
vol % sulfolane–H O at 55 C: ( ) NaClO ; ( ) LiClO4;
2
4
(
) Mg(ClO4)2; ( ) Ba(ClO4)2.

M. Hojo et al.
Bull. Chem. Soc. Jpn. Vol. 79, No. 5 (2006)
753
Table 1. ‘‘Pseudo’’ First-Order Reaction Rate Constants of the Solvolyses for Typical SN1 Substrates in the Presence of
Various Salts in the 50 vol % Sulfolane–H2O Solvent
ꢄ
Salta)
ꢁ1
logðk=sꢁ1
ꢁ1 b)
Substrate
Temp/ C
k/s
6:69 ꢅ 10ꢁ
Þ
ꢀ logðk=s
—
Þ
5
4
5
4
4
4
5
4
5
5
5
1
-Adamantyl chloride
55
No salt
LiClO4
NaClO4
Mg(ClO4)2
Ba(ClO4)2
Ni(ClO4)2
Co(ClO4)2
No salt
Et4NClO4
Et4NCl
Et4NBr
Et4NOTs
Pr4NBr
ꢁ4:17
ꢁ3:94
ꢁ4:02
ꢁ3:60
ꢁ3:56
ꢁ3:97
ꢁ4:08
ꢁ3:97
ꢁ4:19
ꢁ4:37
ꢁ4:23
ꢁ4:67
ꢁ4:57
ꢁ5:12
ꢁ4:89
ꢁ
55
55
55
55
55
55
60
60
60
60
60
60
60
60
1:14 ꢅ 10
þ0:23
þ0:15
þ0:57
þ0:61
þ0:20
þ0:09
—
ꢁ0:22
ꢁ0:40
ꢁ0:26
ꢁ0:70
ꢁ0:60
ꢁ1:15
ꢁ0:92
ꢁ
ꢁ
ꢁ
ꢁ
ꢁ
ꢁ
ꢁ
ꢁ
ꢁ
ꢁ5
ꢁ
ꢁ
9:49 ꢅ 10
2:52 ꢅ 10
2:76 ꢅ 10
1:08 ꢅ 10
8:24 ꢅ 10
1:08 ꢅ 10
6:49 ꢅ 10
4:26 ꢅ 10
5:92 ꢅ 10
c)
2:14 ꢅ 10
2:71 ꢅ 10
7:55 ꢅ 10
5
6
5
n-Bu4NCl
n-Bu4NBr
1:29 ꢅ 10ꢁ
8:10 ꢅ 10ꢁ
6
5
5
5
5
6
6
6
2
-Adamantyl bromide
80
No salt
LiClO4
ꢁ5:09
ꢁ4:91
ꢁ4:92
ꢁ4:65
ꢁ4:69
ꢁ5:19
ꢁ5:40
ꢁ5:26
ꢁ5:48
ꢁ3
—
ꢁ
ꢁ
ꢁ
ꢁ
ꢁ
ꢁ
ꢁ
ꢁ6
80
80
80
80
80
80
80
80
1:22 ꢅ 10
þ0:18
þ0:17
þ0:44
þ0:40
ꢁ0:10
ꢁ0:31
ꢁ0:17
ꢁ0:39
NaClO4
Mg(ClO4)2
Ba(ClO4)2
Et4NClO4
Et4NCl
1:20 ꢅ 10
2:24 ꢅ 10
2:03 ꢅ 10
6:53 ꢅ 10
4:00 ꢅ 10
Et4NBr
Et4NOTs
5:45 ꢅ 10
c)
3:29 ꢅ 10
a) The salt concentration of 1.0 mol dm , except for transition-metal salts of 2:0 ꢅ 10 mol dmꢁ3. b) ꢀ logðk=sꢁ1Þ =
ꢁ
3
logðk1=sꢁ1Þ[with salt] ꢁ logðk0=sꢁ1Þ[no salt] for each substrate. c) p-Toluenesulfonate (tosylate) salt.
Scheme 1. The solvolysis of 1-adamantyl halide in the pres-
ence of alkali metal and alkaline earth metal ions in the
modified ‘‘aqueous’’ solution.
ꢁ
3
NaClO4, respectively. The Arrhenius plots for 2.0 mol dm
of the salts gave similar values, suggesting that the SN1 mech-
anism without salt remains even in the presence of concentrat-
z
z
z
ꢄ
ed salts. The ꢀH , ꢀS , and ꢀG values at 55 C were eval-
uated to be 97.4, ꢁ29:5, and 107.0 (R ¼ 0:998) for no salt,
ꢁ3
Fig. 2. Arrhenius plots of the solvolysis reaction of 1-ada-
mantyl chloride in 50 vol % sulfolane–H2O: ( ) no salt;
97.0, ꢁ24:4, and 105.0 (R ¼ 0:990) for 1.0 mol dm LiClO4,
ꢁ1
ꢁ1
and 101.0 kJ mol , ꢁ14:2 eu, 105.7 kJ mol (R ¼ 0:997) for
ꢁ
3
ꢁ3
(
(
) 1.0 mol dm
LiClO4; ( ) 2.0 mol dm
LiClO4;
ꢁ3
1
.0 mol dm NaClO4, respectively.
As shown above, the solvolysis rate of 1-adamantyl chloride
ꢁ
3
ꢁ3
) 1.0 mol dm NaClO4; ( ) 2.0 mol dm NaClO4.
increased remarkably in the presence of concentrated salts in
spite of the (expected) decrease in the activity of solvent water.
The observed results can be accounted for by the reaction
mechanism in Scheme 1: First of all, the properties of the sul-
folane–water or bulk water should be altered to be those of a
‘‘non-aqueous’’ solvent by the concentrated salts; the ‘‘chemi-
þ
2þ
2þ
Li < Mg , Ba (cf. Table 1).
The Arrhenius plots, as shown in Fig. 2, of the solvolysis re-
ꢄ
action (45–65 C) of 1-adamantyl chloride in the 50 vol % sul-
folane–H2O solvent gave 100.1, 99.8, and 103.8 kJ mol acti-
vation energies (Ea) for no salt, 1.0 mol dm LiClO4, and
ꢁ1
ꢁ3

7
54 Bull. Chem. Soc. Jpn. Vol. 79, No. 5 (2006)
Elucidation of Concentrated Salt Effects
ꢁ
that someone is suspicious that the interaction between X and
þ
2þ
M or M in the 50% aqueous solution causes cleavage of the
34
ꢁþ
ꢁꢁ
C
–X covalent bond. Rudakov and Kozhevnikov have re-
ported exponential increases in the solvolysis rates of t-butyl
halides in EtOH–H2O mixed solvents with the complex forma-
þ
ꢁ
tion constants (e.g., K = [MX]/[M ][X ]) of transition-metal
2þ
þ
3þ
2þ
2þ
þ
2þ
35
halides, such as of Ag , Bi , Cd , Pb , Tl , Co , Mn ,
2
þ
3þ
2þ
Cu , Ga , or Zn . On the other hand, Cox and Maskill
have discovered the solvolyses of t-butyl iodide catalyzed by
I2 in aqueous ethanol; the same effect, but to a smaller extent,
is also found for t-butyl bromide in the solvent. We would like
ꢁ
ꢁ
to note that polyhalogen ion formation, such as I3 and I2Br
(
ꢁ
ꢁ
e.g., I2 + Br ꢀ I2Br ), in the solution can account for the
catalytic effects by I2. The stability of the polyhalogen ions
ꢁ
ꢁ
ꢁ
36
has been reported to decrease as I3 > I2Br > Br3 .
As a reference study, the effects of transition-metal salts
were examined; with the assistance of the complex formation
ꢁ
ꢁþ
between a transition-metal ion and halide ions (X ), the C
ꢁꢁ
–
bond may be cleaved to produce the reaction intermediate
Fig. 3. Changes in the solvolysis rate constants of 1-ada-
mantyl chloride with the addition of various salts in 50
vol % sulfolane–H2O at 60 C: ( ) Et4NBr; ( ) Et4NCl;
X
species (the carbocation). Upon the addition of a small amount
of AgClO4 to 1-adamantyl chloride in 50 vol % sulfolane–H2O
at 55 C, precipitation of AgCl occurred to disturb further titra-
tion with NaOH.
ꢄ
(
) Et4NClO4; ( ) Et4NOTs.
ꢄ
þ
2þ
2þ
cal’’ interaction may operate between M or M and the chlo-
ride of (or from) the substrate to produce the 1-adamantyl cat-
ion in the deformed solvent; the 1-adamantyl cation can react
with H2O promptly, even if the activity of water is weakened.
The effects of non-metallic salts on the SN1 solvolysis rate
were examined in the same solvent. Figure 3 shows the change
in the solvolysis rate of 1-adamantyl chloride with increasing
The addition of Ni at very low concentrations [ð2:0{
4:0Þ ꢅ 10ꢁ mol dm ], however, caused an increase in the sol-
3
ꢁ3
ꢁ
1
volysis rate of the substrate. The slope, logðk=s Þ vs concn of
the metal ion, was extremely larger than that for LiClO4 or
Mg(ClO4)2: the slope of 96.5 for Ni , cf., 0.28 and 0.56 for
2þ
þ
2þ
Li and Mg , respectively. Cobalt(II) perchlorate also caused
a linear increase in the logðk=s Þ of 1-adamantyl chloride
ꢁ1
þ
ꢁ
ꢁ
ꢁ
ꢁ
ꢁ
2þ
the concentration of Et4N X (X = ClO4 , Cl , Br , and
in the same solvent system; however, the slope for Co
2þ
ꢁ
ꢄ
TsO ) at 60 C. All the tetraethylammonium salts caused de-
celeration in the solvolysis rate; the tosylate with the most
bulky among the four anions caused the most remarkable de-
(ꢀ log k=ꢀ c = ca. 40) was smaller than that of Ni . The dif-
2þ
2þ
ference in the acceleration effects between Ni and Co can
be explained by the complex formation (chemical interaction)
ꢁ1
2þ
2þ
ꢁ
crease in the logðk=s Þ value of ꢁ0:70, i.e., the solvolysis rate
abilities of Ni and Co with Cl in aqueous solution:
log ꢂ ¼ 0:34 and ꢁ0:2 for NiCl2 and CoCl2, respectively, at
ꢁ
3
decreased to be one-fifth at 1.0 mol dm of the salt, compared
ꢁ
2
1
with that in the absence of the salt [logðk=s Þ ¼ ꢁ3:97] at
room temperature, where ꢂ is the overall formation constant
of M (Cl )2. Halide salts (Cl and Br ) of Ni
2
ꢄ
2þ
ꢁ
37
ꢁ
ꢁ
2þ
and
6
0 C. The tetraethylammonium salts caused the deceleration
ꢁ
ꢁ
ꢁ
ꢁ
2þ
effect in the order of ClO4 < Br < Cl < TsO for the
anions. The effects must be brought by the decrease in the
activity of solvent water; that is, the solvolysis reaction in
the present system should be initiated by the attack of water
molecules towards the substrate in order to produce the carbo-
cation. The solvation abilities of the anions may result in the
difference of ‘‘effective’’ water activities.
Co caused no obvious differences in the solvolysis rate con-
stants, compared with the perchlorate salts. Strong acids, such
as HClO4 and HCl, of ð2:0{4:0Þ ꢅ 10ꢁ mol dm also caused
exponential increases in the rate constants with ꢀ log k=ꢀc =
ca. 25.
3
ꢁ3
The above results indicate that complex formation with
ꢁ
transition-metal ions can extract the Cl ion from 1-adamantyl
Similar but more remarkable deceleration in the solvolysis
rate was observed with other tetraalkylammonium bromides
chloride, which may induce the exponential increase in the sol-
volysis rate constant with the metal ion concentration. We
imagine that the ‘‘special salt effect’’ proposed by Winstein
(
R4NBr: R ¼ Pr and n-Bu). The longer the chain in tetraalkyl-
4
ammonium bromide, the more remarkable the deceleration of
ꢁ1
et al. might be brought by the ‘‘direct’’ chemical interaction
þ
þ
þ
þ
ꢁ
the solvolysis rate as Et4N < Pr4N < n-Bu4N : logðk=s
Þ
between Li (ClO4 ) and benzenesulfonate anions from sub-
strates in lower permittivity media, such as acetic acid. Raman
spectral evidences will be given in the final section to show the
distortion of hydrogen-bonding structure of aqueous solution
with the addition of organic solvents and concentrated salts
or with increasing temperature.
ꢁ3
values with 1.0 mol dm bromide salts were ꢁ4:23, ꢁ4:57,
ꢄ
and ꢁ4:89, respectively, at 60 C. The chloride salt of n-
þ
ꢁ1
Bu4N caused slightly larger deceleration (logðk=s Þ ¼
ꢁ
5:12) than the bromide salt (cf. Table 1).
Chemical Interaction between Anions and Metal Ions.
No interaction, except for the Coulombic force, should operate
þ
On the Solvolysis Rate of 2-Adamantyl Bromide. Anoth-
er typical SN1 substrate, 2-adamantyl bromide, in the same sol-
vent system was subject to similar acceleration effects by the
metal perchlorates as shown in Fig. 4. The solvolysis rate con-
ꢁ
between the halide ions (X ) and alkali metal (M ) or alkaline
earth metal ions (M2 ) of lower concentrations in aqueous so-
þ
lution at room temperature. Therefore, it must be quite natural

M. Hojo et al.
Bull. Chem. Soc. Jpn. Vol. 79, No. 5 (2006)
755
Fig. 4. Changes in the solvolysis rate constants of 2-ada-
mantyl bromide with the addition of various salts in 50
Fig. 5. Changes in the solvolysis rate constants of SN2 sub-
strates with the addition of LiClO4 or Et4NClO4 in
ꢄ
vol % sulfolane–H2O at 80 C: ( ) NaClO4; ( ) LiClO4;
) Mg(ClO4)2; ( ) Ba(ClO4)2.
ꢄ
50 vol % sulfolane–H2O at 80 C: (circles) hexyl tosylates;
triangles) hexyl bromide; (squares) hexyl chloride. Open
(
(
and solid symbols are for LiClO4 and Et4NClO4, respec-
tively.
ꢁ
1
ꢁ6
stant of the substrate was observed to be k=s ¼ 8:10 ꢅ 10
ꢄ
ꢁ1
at 80 C. The logðk=s Þ value increased linearly with increas-
ing concentration of LiClO4. However, the salts of the other
þ
Solvolyses of SN2 Substrates, Hexyl Chloride, Bromide,
and Tosylate. As typical SN2 substrates, hexyl halides and
tosylate were selected in the present study; aliphatic halides
of shorter aliphatic chains, such as methyl or ethyl chloride,
are apt to escape from the reaction vessel because of their
low boiling points: ꢁ23:7, 12.3, 46–47, 78.5, 107.8, and 134
2þ
2þ
metal ions, Na , Mg , and Ba , gave convex curves at high-
er salt concentrations. The acceleration effects by the metal
þ
þ
2þ
2þ
ions increased as Na < Li < Mg , Ba . The metal cat-
ions should assist the formation of a stable carbocation from
2
-adamantyl bromide in a manner similar to 1-adamantyl chlo-
ꢄ
ride. However, the carbocation of a secondary carbon type
may be less stable than of a tertiary carbon type. Therefore,
the degree of acceleration caused by the metal salts in the
solvolysis rate for 2-adamantyl bromide should be less than
that for 1-adamantyl chloride, regardless of the difference in
the leaving group. For instance, the solvolysis rates of the
C for methyl, ethyl, propyl, butyl, pentyl, and hexyl chlorides,
respectively.43 In Fig. 5 are shown the changes in the solvoly-
ꢁ1
sis rates, logðk=s Þ of hexyl chloride, bromide, and tosylate
ꢁ3
with 0–2.0 mol dm LiClO4 in 50 vol % sulfolane–H2O at
ꢄ
80 C. As contrasted with the cases of the typical SN1 sub-
strates, the solvolysis rates of the SN2 substrates were not ac-
1
- and 2-derivatives increased with the logarithm factors
celerated but rather slightly decelerated by the addition of the
ꢁ1
ꢀ logðk=sꢁ Þ] of þ0:85 and þ0:62 at 3.0 mol dm LiClO4,
1
ꢁ3
metal perchlorates. The logðk=s Þ value of hexyl tosylate de-
[
respectively. Ionization potentials and electron affinities of
adamantane (Ad) have been evaluated by Yan et al.; 1-
Ad is predicted to lie 47 kJ mol below 2-Ad . The influ-
ence of alkyl substitution on the solvent effects in solvolysis
of 1-adamantyl bromide has been examined by Takeuchi et
al. In the above discussion, only the stability of the carbo-
cation was taken into account; the role of the leaving group is
also important. However, the problem of solvation for anions
prevented us from discussing further.
creased linearly with increasing LiClO4 concentration: the de-
ꢁ1
38
crease in logðk=s Þ, ꢀ log k ¼ ꢁ0:13 and ꢁ0:20 at 1.0 and
þ
ꢁ1
þ
ꢁ3
2.0 mol dm LiClO4. As for the halide substrates, more dis-
tinct deceleration appeared to be caused by the addition of
LiClO4 of higher concentrations: the decrease for the chloride
ꢁ3
3
9
substrate; ꢀ log k ¼ ꢁ0:46 and ꢁ0:66 at 1.0 and 2.0 mol dm
LiClO4, respectively (cf. Table 2). The activation energy (Ea)
ꢁ1
was found to be 94.6 kJ mol for the solvolysis of hexyl tosy-
ꢄ
z
z
late without salts in the solvent (50–80 C). The ꢀH , ꢀS ,
z
ꢄ
ꢁ1
The deceleration effects on 2-adamantyl bromide in the pres-
ꢁ
and ꢀG values at 55 C were evaluated to be 91.9 kJ mol
ꢁ1
,
ence of non-metallic salts, Et4NX, were observed as ClO4
<
ꢁ58:2 eu, and 110.0 kJ mol , respectively, with R ¼ 0:9999.
ꢁ
ꢁ
ꢁ
ꢁ1
Br < Cl < TsO : the logðk=s Þ values of ꢁ5:19, ꢁ5:26,
It goes without saying that in the SN2 solvolysis (hydroysis)
ꢁ
ꢁ
3
ꢁ
5:40, and ꢁ5:48, respectively, for 1.0 mol dm salts. The
system, the metal cations may have no way to assist to give X
ꢁ
þ
ꢁꢁ
fact that the solvolysis rates of 2-adamantyl bromide and 1-ada-
mantyl chloride were decelerated by the non-metallic salts in
the same order of the salt anion indicated that the common
ion effect was not much concerned in the solvent system,
although the common ion salt effects have been reported in
on the SN2 transition state (H2OꢆꢆꢆC H2RꢆꢆꢆX ). The deceler-
ation in the solvolysis in the presence of LiClO4 can be attrib-
uted to the decrease in the activity of water, which attacks the
substrate in order to initiate the solvolysis reaction. A non-
metallic salt, Et4NClO4, also caused the deceleration in the
solvolysis rates of the SN2 substrates to a similar extent as
LiClO4 (cf. Fig. 5).
40
41
SN1 solvolyses in acetic acid, phenol, and aqueous ace-
2
4
tone.

7
56 Bull. Chem. Soc. Jpn. Vol. 79, No. 5 (2006)
Elucidation of Concentrated Salt Effects
Table 2. ‘‘Pseudo’’ First-Order Reaction Rate Constants of the Solvolyses for Typical SN2 Substrates in the
ꢄ
Presence of Various Salts in 50 vol % Sulfolane–H2O Solvent at 80 C
ꢄ
Salta)
No salt
ꢁ1
logðk=sꢁ1
ꢁ1 b)
Substrate
Temp/ C
k/s
7:67 ꢅ 10ꢁ
Þ
ꢀ logðk=s
—
Þ
7
7
7
7
Hexyl chloride
80
ꢁ6:12
ꢁ6:58
ꢁ6:19
ꢁ6:16
ꢁ4:94
ꢁ5:86
ꢁ
80
80
80
80
80
LiClO4
2:62 ꢅ 10
ꢁ0:46
ꢁ0:07
ꢁ0:04
þ1:18
þ0:26
ꢁ
ꢁ
ꢁ5
ꢁ6
Et4NClO4
Et4NCl
6:52 ꢅ 10
6:99 ꢅ 10
d)
Et4NBr
1:15 ꢅ 10
1:37 ꢅ 10
c),d)
Et4NOTs
ꢁ
ꢁ
ꢁ
ꢁ7
ꢁ
5
6
6
Hexyl bromide
Hexyl tosylate
80
No salt
LiClO4
Et4NClO4
Et4NCld)
Et4NBr
1:41 ꢅ 10
9:49 ꢅ 10
8:24 ꢅ 10
4:95 ꢅ 10
2:03 ꢅ 10
2:58 ꢅ 10
ꢁ4:85
ꢁ5:02
ꢁ5:08
ꢁ6:31
ꢁ4:69
ꢁ4:59
ꢁ3:77
ꢁ3:90
ꢁ3:89
ꢁ6:28
ꢁ4:72
ꢁ3:82
—
80
80
80
80
80
ꢁ0:17
ꢁ0:23
ꢁ1:46
þ0:16
þ0:26
5
c),d)
ꢁ5
Et4NOTs
1:70 ꢅ 10ꢁ
4
—
80
No salt
LiClO4
Et4NClO4
Et4NCl
Et4NBr
ꢁ
ꢁ
ꢁ
ꢁ
4
4
7
5
80
80
80
80
80
1:27 ꢅ 10
ꢁ0:13
ꢁ0:12
ꢁ2:51
ꢁ0:95
ꢁ0:05
1:30 ꢅ 10
5:27 ꢅ 10
1:89 ꢅ 10
c)
ꢁ4
Et4NOTs
1:50 ꢅ 10
ꢁ
3
ꢁ1
ꢁ1
ꢁ1
a) The salt concentration of 1.0 mol dm . b) ꢀ logðk=s Þ = logðk1=s Þ[with salt] ꢁ logðk0=s Þ[no salt]
for each substrate. c) p-Toluenesulfonate (tosylate) salt. d) Only the linear parts in the ln½Sꢇ vs t plot were
evaluated because there is a leaving group exchange reaction.
Fig. 6. The (apparent) substrate concentration changes for
hexyl bromide with time during solvolysis in the presence
Fig. 7. The (apparent) substrate concentration changes for
hexyl chloride with time during solvolysis in the presence
ꢄ
ꢄ
of Et4NCl in 50 vol % sulfolane–H2O at 80 C: ( ) 0; ( )
of Et NBr in 50 vol % sulfolane–H O at 80 C: ( ) 0; ( )
4
2
ꢁ3
ꢁ3
0
.1; ( ) 0.2; ( ) 0.5; ( ) 1.0 mol dm Et4NCl.
0.1; ( ) 0.2; ( ) 0.5; ( ) 1.0 mol dm Et NBr.
4
Solvolyses Accompanied by ‘‘Ion Exchange’’ Reactions
of hexyl bromide by the addition of Et4NCl. Oppositely, the
plots in Fig. 7 demonstrate the drastic increases of the (appa-
rent) solvolysis rates of hexyl chloride by the addition of
Et4NBr. The relatively slow solvolysis reaction, compared
of SN2 Substrates. It must be noticed that the effects by tetra-
ethylammonium salts other than the perchlorate on the solvol-
ysis rates of the SN2 substrates were completely different from
those of Et4NClO4.
Figure 6 shows the plots of ln½Sꢇ vs t during the solvolysis
of hexyl bromide in the presence of various concentrations of
Et4NCl, where the amount of substrate, [S], is expressed by the
titration volume (mL) with NaOH (½V1 ꢁ Vtꢇ). The plots can
demonstrate the drastic decreases in (apparent) solvolysis rates
ꢁ
ꢁ
with the exchange between Cl and Br , may cause the curved
lines for the whole observed period. With a careful examina-
tion of the plots of ln½Sꢇ vs t in Fig. 6, we found that, in the
very early stage of the reaction, the ln½Sꢇ value decreased fol-
lowing that for hexyl bromide itself even in the presence of
ꢁ3
Et4NCl (e.g., 0.2 mol dm ); however, the plots formed linear

M. Hojo et al.
Bull. Chem. Soc. Jpn. Vol. 79, No. 5 (2006)
757
Scheme 2. The ‘‘anion exchange’’ and solvolysis (hydroly-
sis) reactions in the presence of sufficient amounts of dif-
ferent anions.
on methyl esters in water as follows: the log k value for
CH3OSO2C6H5 was increased by factors of 2.0 and 2.6 with
ꢁ
ꢁ
ꢄ
Cl and Br , respectively, compared with H2O at 25 C.
Judging from the symmetrical feature (cf. Fig. 8) of the
changes in the solvolysis rates of hexyl bromide and chloride
with increasing concentrations of Et4NCl and Et4NBr, respec-
ꢁ
tively, we may suppose that the nucleophilicities of Cl and
ꢁ
Fig. 8. Changes in the solvolysis rate constants of SN2
substrates with the addition of various salts in 50 vol %
Br are at the same level in a 50 vol % sulfolane–H2O solution
ꢄ
at 80 C. The proposed mechanisms of the ‘‘anion exchange’’
and hydrolysis reactions are illustrated in Scheme 2.
ꢄ
sulfolane–H2O at 80 C: ( ) Et4NBr; ( ) Et4NCl; (
)
Et4NOTs for hexyl tosylate; ( ) Et4NBr; ( ) Et4NCl;
In the series of studies, we have realized that, in the co-
existence of Et4NCl, some substrates possessing much higher
boiling points than the reaction temperatures, such as ethyl
(
(
) Et4NOTs for hexyl bromide; ( ) Et4NBr; ( ) Et4NCl;
) Et4NOTs for hexyl chloride.
ꢄ
43
tosylate (bp ¼ 173 C) or butyl bromide (bp760 ¼ 101:3
1
5
ꢄ
43
curves parallel to the plot curve for hexyl chloride afterward
C), have been lost in reaction vessels easily and no sub-
strates remain to be solvolized; this suggests some parts of
the substrates have been already converted into more volatile
compounds (eg., EtCl, the boiling point of 12.3 C) while sol-
volysis reactions are occurring.
(
cf. Figs. 6 and 7). The linear parts in the plots of ln½Sꢇ vs t
were utilized for evaluating the k values.
Figure 8 shows the changes in the solvolysis rates of hexyl
chloride, bromide, and tosylate in the presence of 0–1.0
ꢄ
ꢁ
3
þ
ꢁ
ꢁ
ꢁ
ꢁ
ꢁ
mol dm Et4N X (X = Cl , Br , and TsO ). The solvol-
ysis rate of hexyl bromide was remarkably decelerated with in-
creasing concentration of Et4NCl and approached that of the
The technique of adding a strong nucleophile (e.g., lyate
ions or the azide ion) has been often used so as to examine
the tendency of a substrate to react through either the SN1 or
ꢁ3
46,47
chloro-substrate (the intrinsic value) at 0.2 mol dm Et4NCl;
further addition of the salt caused slightly further deceleration
of the solvolysis rate. However, Et4NBr or Et4NOTs caused no
or minor acceleration in the hexyl bromide solvolysis. On the
other hand, the solvolysis rate of hexyl chloride was remark-
ably accelerated with increasing the concentration of Et4NBr;
SN2 pathway.
In 75% aqueous dioxane, the addition of
ꢁ3
0.1 mol dm NaN3 to 1-adamantyl bromide (a typical SN1
ꢄ
substrate, at 100 C) and 2-propyl tosylate (an SN1–SN2 bor-
ꢄ
derline substrate, at 50 C) caused the formation of RN3 in 6
and 87%, respectively.46
Raman Spectral Changes of Water OD Stretching with
Adding Organic Solvents and Salts, or with Increasing
Temperature. Figure 9 shows the Raman spectra of water
(D2O) containing sulfolane at room temperature. The pure
D2O gives an OD stretching band with two peaks at 2380
ꢁ
3
at 1.0 mol dm Et4NBr, the rate approached that of hexyl bro-
ꢁ
ꢁ
mide (the intrinsic value). The salts of Cl and TsO caused
no and a minor acceleration effect, respectively, on the solvol-
ysis rate of the chloro-substrate. The acceleration and deceler-
ation effects with the addition of different anions may suggest
that ‘‘ion exchange’’ reactions occur prior to or in parallel to
the solvolysis reaction.
The addition of 0.5 mol dm 3 Et4NBr and Et4NCl to hexyl
tosylate caused deceleration in the solvolysis rates towards
those of (as if) hexyl bromide and chloride, respectively.
The addition of Et4NOTs caused no change in the solvolysis
rate of hexyl tosylate as well as of the halide substrates. The
above results may verify that the tosylate ion in the added salt
has a weak ability to substitute (or exchange) the anions from
the substrates, i.e., hexyl bromide and chloride. However, the
bromide and chloride ions can substitute anions easily from all
ꢁ1
and 2520 cm ; these peaks are assigned to strongly hydro-
gen-bonded D2O and weakly (or non-) hydrogen-bonded
D2O, respectively.48 Upon the addition of sulfolane to D2O,
the peak at 2520 increased at the expense of the intensity at
ꢁ
ꢁ1
2380 cm . The increasing content of sulfolane caused further
decreases in the intensity at 2380 cm and an increase in the
ꢁ1
ꢁ
1
intensity of the peak at 2520 cm or at higher wavenumbers.
At 90 vol % sulfolane, the wavenumber of weakly (or non-)
hydrogen-bonded D2O seemed to approach that of the D2O
monomer (ꢃ1 ¼ 2671 cmꢁ ) in the gas phase. It has been
suggested that the bulk water, in which the H2O molecules
may highly associate through hydrogen bonding, would be
altered into rather isolated water molecules by the addition
1
49
the substrates. It is well known that the nucleophilicity order of
ꢁ
14
halide ions in aprotic solvents, such as acetone, is Cl
ꢁ
>
Br > I , whereas it is Cl < Br < I in aqueous solu-
of aprotic solvents.
Further distortion of the water structure was brought by
the addition of salts into water–organic solvent mixtures.
ꢁ
ꢁ
ꢁ
ꢁ
4
4
45
tion. Harris et al. have noted the attack of nucleophiles

7
58 Bull. Chem. Soc. Jpn. Vol. 79, No. 5 (2006)
Elucidation of Concentrated Salt Effects
Temperature effects on the Raman spectra of 50 vol % sul-
ꢄ
folane–D2O were also examined over the range of 5–75 C.
At higher temperatures, the intensity around 2400 nm decreas-
ꢁ1
ed, and the peak intensity at ca. 2520 cm slightly increased:
The distortion of hydrogen-bonding structure at higher temper-
ature was demonstrated. It can be concluded that the hydrogen-
bonding structure of bulk water is distorted not only by the ad-
dition of good amounts of organic solvents and concentrated
salts but also by increasing temperature, admitting that D2O
is a more structural liquid than H2O.50 Recently, Raman spec-
tra of water monomers were observed in highly supersaturated
droplets of NaClO4, LiClO4, and Mg(ClO4)2; where the hydro-
gen-bonding network of water molecules were broken.51 The
water structure has been widely discussed for water–organic
solvent mixtures.52
Now, we may conclude that the properties of the sulfolane–
H2O solution containing concentrated salts are approaching
those of ‘‘dihydrogen ether.’’ In ‘‘dihydrogen ether’’ or a modi-
fied ‘‘aqueous’’ solution, direct chemical interaction should
Fig. 9. The Raman spectra of D2O containing sulfolane at
room temperature: (——) 0; (- - - -) 50; (– – –) 75; (– - –)
þ
operate between alkali metal or alkaline earth metal ions (M
or M ) and anions from substrates (RX).
90 vol % of sulfolane.
2þ
Experimental
Sulfolane as the solvent, purchased from Nacalai Tesque and
Wako were used as received. The substrates from Tokyo Kasei
(
TCI) were used without further purification: 1-adamantyl chlo-
ride (>98%) and bromide (>97%), 2-adamantyl bromide
>97%), t-butyl chloride (>98%), hexyl chloride (>95%), bro-
(
mide (>98%), and p-toluenesulfonate (tosylate) (>98%). Salts
of NaClO4, Et4NBr, Et4NOTs, and Pr4NBr from Aldrich as well
as LiClO4 and Et4NCl from Wako, and n-Bu4NCl from TCI were
used. Tetraethylammonium perchlorate (Et4NClO4) was prepared
by adding an equivalent amount of HClO4 to Et4NBr in water.
The precipitates were filtered, washed with cold water and were
recrystallized from water several times, followed by drying at
ꢄ
7
0 C in vacuo. Tetrabutylammonium bromide was prepared as
follows: tributylamine and butyl bromide in acetonitrile solvent
were reflaxed for 36 h, followed by evaporation of the solvent to
dryness; the n-Bu4NBr precipitates were recrystallized from ethyl
ꢄ
acetate twice and dried at 70 C in vacuo. Commercially obtained
transition-metal salts, Ni(ClO4)2 6H2O (Aldrich), Co(ClO4)2
Fig. 10. The Raman spectra of 50 vol % sulfolane–D2O
mixed solvent in the presence of concentrated LiClO4
at room temperature: (——) R = 1 (no salt); (– – –)
R = 20; (- - - -) R = 10; where R = [sulfolane + D2O]/
ꢆ
ꢆ
6
(
H2O (Alfa), NiCl2 6H2O, NiBr2, CoCl2 6H2O, and CoBr2
Wako), were used as received.
The initial concentrations of the substrates were all kept at
0.010 mol dm in the reaction vessels. In principle, 0–90% con-
sumption of a substrate was followed for evaluating the rate con-
ꢆ
ꢆ
[
LiClO4].
ꢁ3
Figure 10 shows the changes in Raman spectra of a 50 vol %
sulfolane–D2O mixed solvent containing concentrated LiClO4.
The concentration of LiClO4 is represented by the ratio of the
sum of moles of sulfolane and D2O ([sulfolane] + [D2O]) to
the moles of LiClO4 ([LiClO4]). With increasing the concen-
tration of LiClO4, i.e., decreasing the R value, the intensity
stant (relative error of <ꢈ1:0%), except for some cases; in S 2
N
substrates containing Et4NX salts, only the linear parts in the
ln½Sꢇ vs t curve were utilized. Details for the other reagents and
the procedure for the kinetic experiments have been described pre-
viously.12–14 Raman spectra of D2O (99.9 atom % D from Aldrich)
were recorded with a Horiba-Jobin Yvon LabRam Spectrograph
HR-800. The Raman intensity was standardized with the CD
stretching band of 5 vol % CD3OD–D2O as an internal intensity
standard; CD3OD, 99.8 atom % D, from Aldrich was used. The
temperature of the sample was regulated by circulating water of
ꢁ1
around 2400 cm decreased, indicating that hydrogen-bonded
water molecules were altered into non-hydrogen bonded, i.e.,
isolated water molecules. The addition of Et4NBr caused ef-
fects similar to those by LiClO4; however, the wavenumber
ꢄ
ꢁ
1
constant temperature (5–75 C).
of the ca. 2520 cm peak remained almost constant with in-
creasing concentration of Et4NBr. The changes of Raman
spectra of D2O containing Et4NBr or LiBr have been dis-
We are grateful to Dr. Ryoji Fujiyama of Kochi University
for his valuable suggestions and support to the present study.
1
4
cussed.

M. Hojo et al.
Bull. Chem. Soc. Jpn. Vol. 79, No. 5 (2006)
759
2
996, 69, 971.
5
M. Hojo, H. Hasegawa, H. Yoneda, Bull. Chem. Soc. Jpn.
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