Cesium ion pair acidities of some N,N-dialkylacetamides and aggregation of their cesium enolates in tetrahydrofuran

Article abstract of DOI:10.1021/jo960687z

Cesium ion pair acidities are reported of some N,N-dialkylacetamides at 25°C in tetrahydrofuran. The cesium enolates of α-arylacetamides are essentially monomeric at concentrations of about 10^-4 M, but small amounts of dimers are present for the enolates of N,N-dimethyl- and N,N-diethyl(4-biphenylyl)acetamide (K_dimer is approximately 400 M^-1). The cesium enolate of N,N-dimethyldiphenylacetamide forms small amounts of dimers, but the dimerization constant of <100 M^-1 is near the limit of detection of our methods. The enolate of N,N-diethylacetamide has an average aggregation number of 3.3, consistent with a mixture of dimers and tetramers. The cesium ion pair pKs of the α-arylacetamides are about 25-26, and that of N,N-diethylacetamide is greater than 33.7. The acidity data are in good agreement with measurements that have been reported in the literature for similar systems.

Full text of DOI:10.1021/jo960687z

6
354
J . Org. Chem. 1996, 61, 6354-6359
Cesiu m Ion P a ir Acid ities of som e N,N-Dia lk yla ceta m id es a n d
Aggr ega tion of Th eir Cesiu m En ola tes in Tetr a h yd r ofu r a n
J ames A. Krom and Andrew Streitwieser*
College of Chemistry, University of California, Berkeley, California 94720-1460
Received April 16, 1996X
Cesium ion pair acidities are reported of some N,N-dialkylacetamides at 25 °C in tetrahydrofuran.
The cesium enolates of R-arylacetamides are essentially monomeric at concentrations of about 10-
M, but small amounts of dimers are present for the enolates of N,N-dimethyl- and N,N-diethyl(4-
4
-
1
biphenylyl)acetamide (Kdimer is approximately 400 M ). The cesium enolate of N,N-dimethyl-
-
1
diphenylacetamide forms small amounts of dimers, but the dimerization constant of <100 M is
near the limit of detection of our methods. The enolate of N,N-diethylacetamide has an average
aggregation number of 3.3, consistent with a mixture of dimers and tetramers. The cesium ion
pair pKs of the R-arylacetamides are about 25-26, and that of N,N-diethylacetamide is greater
than 33.7. The acidity data are in good agreement with measurements that have been reported in
the literature for similar systems.
The paucity of systematic studies of amide enolates
eration of aggregation effects is important in interpreting
reactivity, and of special interest is the possibility that
stands in contrast to the vast literature concerning
ketone enolates. The disparity is attributable to the
ubiquity of ketone enolates in synthetic chemistry.
Simple alkylation and addition reactions of amide eno-
lates have been known for some time,1 but were not
used extensively until asymmetric syntheses were
aggregates play an important role in stereoselective
reactions.1
5,27,28
We report here a study of the cesium ion
pair acidities of several N,N-dialkylamides and the
aggregation of their cesium enolates in tetrahydrofuran
(THF); the results provide a comparison to related
-5
6
-9
18
demonstrated.
Recent studies have focused on their
ketones. In the only other study of aggregation of amide
reactivities in Michael additions.1
0-14
enolates to our knowledge, the lithium salt of N,N-
1
5-19
It has become recognized that alkali metal enolates,
dimethylcycloheptatrienecarboxamide was found to be
phenolates,2
0-24
and similar compounds
25,26
generally
29
partially aggregated in melting THF.
aggregate in nonpolar solvents; dimeric, tetrameric, and
hexameric species are especially common. Papers of
Arnett and co-workers17 and of J ackman and co-workers
give comprehensive citations to this literature. Consid-
The present study includes the compounds 1-6 and
their cesium salts (Cs-1-Cs-6).
22
X
Abstract published in Advance ACS Abstracts, August 1, 1996.
(
(
(
(
(
1) Carbon Acidity. 94.
2) Needles, H. L.; Whitfield, R. E. J . Org. Chem. 1966, 31, 989.
3) Gassman, P. G.; Fox, B. L. J . Org. Chem. 1966, 31, 982.
4) Woodbury, R. P.; Rathke, M. W. J . Org. Chem. 1977, 42, 1688.
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1
947.
(
(
(
6) Evans, D. A.; Takacs, J . M. Tetrahedron Lett. 1980, 21, 4233.
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Resu lts
chemistry; Allinger, N. L., Eliel, E. L., Wilen, S. H., Eds.; Wiley: New
York, 1982; Vol. 13.
Cesiu m Ion P a ir Absor p tion Sp ectr a . The cesium
ion pairs of the compounds selected for study, with the
exception of the cesium salt of 6 (Cs-6), absorb in the
visible or near-UV. It was for this reason that the
biphenyl compounds 1 and 2 were chosen for study; the
absorbance of phenyl derivatives is too far in the UV to
be useful. The absorption spectra were found to be
slightly concentration dependent. For each ion pair, a
series of 13-15 absorption spectra at varying concentra-
(
(
9) Sauriol-Lord, F.; Grindley, T. B. J . Org. Chem. 1981, 46, 2831.
10) Stefanovsky, Y.; Gospodova, T.; Viteva, L. Tetrahedron 1986,
4
2, 5355.
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H. J . Org. Chem. 1990, 55, 132.
(
(
(
12) Viteva, L.; Stefanovsky, Y. Tetrahedron Lett. 1990, 31, 5649.
13) Viteva, L.; Gospodova, T. S.; Stefanovsky, Y. N. Tetrahedron
1
1
994, 50, 7193.
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(15) Seebach, D. Angew. Chem., Int. Ed. Engl. 1988, 21, 456.
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(17) Palmer, C. A.; Ogle, C. A.; Arnett, E. M. J . Am. Chem. Soc.
tions was acquired and processed by the singular value
decomposition (SVD) method.3
0,31
The total concentration
992, 114, 5619.
(18) Ciula, J . C.; Streitwieser, A. J . Org. Chem. 1992, 57, 431.
(19) Abbotto, A.; Streitwieser, A. J . Am. Chem. Soc. 1995, 117, 6358.
(20) J ackman, L. M.; C¸ izmeciyan, D.; Williard, P. G.; Nichols, M.
-4
ranges for the SVD experiments were (0.4-4.0) × 10
A. J . Am. Chem. Soc. 1993, 115, 6262.
(27) Seebach, D. In The Robert A. Welch Foundation Conferences
on Chemical Research XXVII. Stereospecificity in Chemistry and
Biochemistry; Houston, TX, 1983; p 93.
(
21) J ackman, L. M.; Rakiewicz, E. F.; Benesi, A. J . J . Am. Chem.
Soc. 1991, 113, 4101.
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991, 113, 3451.
23) J ackman, L. M.; Rakiewicz, E. F. J . Am. Chem. Soc. 1991, 113,
202.
(
(28) Carlier, P. R.; Lo, K. M. J . Org. Chem. 1994, 59, 4053.
(29) Bauer, W.; Seebach, D. Helv. Chim. Acta 1984, 67, 1972.
(30) Golub, G. H.; Reinsch, C. In Linear Algebra; Wilkinson, J . H.,
Reinsch, C., Eds.; Springer-Verlag: Heidelberg, 1971. Press, W. H.;
Flannery, B. P.; Teukolsky, S. A.; Vetterling, W. T. Numerical Recipes
(Fortran Version); University Press: Cambridge, 1989.
(31) Malinowsky, E. R. Factor Analysis in Chemistry; Wiley: New
York, 1991.
1
(
1
(
(
(
24) J ackman, L. M.; Chen, X. J . Am. Chem. Soc. 1992, 114, 403.
25) Arnett, E. M.; Moe, K. D. J . Am. Chem. Soc. 1991, 113, 7288.
26) Carlier, P. R.; Lucht, B. L.; Collum, D. B. J . Am. Chem. Soc.
1
994, 116, 11602.
S0022-3263(96)00687-1 CCC: $12.00 © 1996 American Chemical Society

Cesium Ion Pair Acidities of N,N-Dialkylacetamides
J . Org. Chem., Vol. 61, No. 18, 1996 6355
indicator method of acidity determination, in which the
concentrations of the ion pairs are determined by linear
least-squares best fitting of the spectra of the separate
3
3
species to the spectrum of the equilibrium mixture.
The acidity of 6 was determined against diphenyl-
methane (pKCs ) 33.25),32 but the lack of a suitable
enolate absorption necessitated the use of the less precise
single indicator method.3
4,35
The observed acidities at 25
°
C are presented in Figure 2.
Discu ssion
We first consider the possibility that the cesium amide
enolates investigated here are significantly aggregated
to ion pair dimers or higher aggregates. It has been
shown that aggregation of the ion pair of interest causes
the apparent acidity of its conjugate acid to increase as
3
5
the formal concentration of the ion pair is increased.
This behavior arises because the ion pair acidity is
defined relative to monomeric species, and aggregation
is essentially a side reaction that consumes the monomer.
Thus, the apparent acidity, which we determine by
measuring the formal (total) concentration, increases
because aggregation is favored by increasing concentra-
tion. Concentration-dependent aggregation equilibria
sometimes can also be detected by UV-visible absorption
spectroscopy. Since higher order aggregates generally
absorb at shorter wavelengths than lower order ag-
gregates, the absorption bands of the solutions exhibit a
blue shift as the concentration is increased. Usually the
shift is rather small, so we analyze the spectroscopic data
by singular value decomposition (SVD), which is sensitive
to systematic changes in the shapes of the absorption
F igu r e 1. Concentration dependence of the UV-visible
absorption spectra of the cesium enolates of 1 (A) and 2 (B).
The isosbestic points are at 412 nm (A) and 416 nm (B).
2
9,30
bands. Detailed discussions of SVD
and its applica-
31
tion to spectroscopy can be found elsewhere.
The SVD results show that the UV-visible absorption
spectra of each of the ion pairs investigated here (except
Cs-6, which was not examined spectroscopically because
of the lack of a suitable chromophore) consist of two
components, which suggests the presence of two distinct
chemical species in equilibrium. The actual spectra of
these species are not provided by the SVD treatment.
Nevertheless, we have shown that such spectra can be
constructed from the SVD output with the assumption
of a particular aggregation model.33 For example, if we
assume that an ion pair system consists of an equilibrium
between monomer and dimer, then the hypothetical
monomer and dimer spectra can be constructed, and the
equilibrium constant can be calculated with the aid of
the experimentally measured extinction coefficient. Our
experience has shown that many aggregation models will
be consistent with the spectroscopic data, but in general
Ta ble 1. Sp ectr oscop ic Da ta for th e Cesiu m En ola tes of
th e Ar yl-Su bstitu ted Am id es
compounda
λmax (nm)
ꢀ (M-1 cm-1
)
S1b
S2b
c
c
1
2
3
4
5
, BiphCH2CONMe2
, BiphCH2CONEt2
, Ph2CHCONMe2
, Ph2CHCONEt2
, Ph2CHCONC4H8
412
416
392
398
388
35,600
36,600
19,400
18,700
19,000
352.7
351.9
344.7
396.9
356.8
5.5
4.4
2.2
1.6
2.0
c
c
a
Spectroscopic data for the cesium enolate of the indicated
amide. b S1 and S2 are the first and second singular values from
the SVD procedure. Data are for the isosbestic point (see text).
c
-
4
M for the biphenylylacetamides and (0.3-7.0) × 10
M
for the diphenylacetamides. The SVD results show that
the spectra of each ion pair are the sums of two
components having different concentration dependence.
Furthermore, extinction coefficient measurements of Cs-1
and Cs-2 at varying concentrations revealed isosbestic
points at 412 and 416 nm, respectively (Figure 1). We
could not locate clear isosbestic points in the absorption
spectra of the diphenylacetamide enolates, but the con-
centration dependencies of these spectra are probably
smaller than the experimental error in our extinction
coefficient measurements. (See the Experimental Section
for details.) The spectroscopic data are collected in Table
3
3
only one will also be consistent with the acidity data;
this combination is an example of coupled equilibria.
The spectra of Cs-1 and Cs-2 clearly show concentra-
tion-dependent behavior. When extinction coefficient
measurements are made, the normalized spectra show
isosbestic points (Figure 1) as well as blue shifts with
increasing concentration. These observations are evi-
dence of simple equilibria between two aggregated spe-
1
.
(
32) Streitwieser, A.; Ciula, J . C.; Krom, J . A.; Thiele, G. J . Org.
Chem. 1991, 56, 1074.
33) Krom, J . A.; Petty, J . T.; Streitwieser, A. J . Am. Chem. Soc.
Cesiu m Ion P a ir Acid ity. The acidities of the aryl-
substituted acetamide derivatives were measured against
(
the indicator 9-tert-butylfluorene (tBuFl), which has a pK
1993, 115, 8024.
of 24.39 on the cesium ion pair acidity scale in THF.32
(34) Streitwieser, A., J r.; Reuben, D. M. E. J . Am. Chem. Soc. 1971,
9
3, 1794.
35) Kaufman, M. J .; Streitwieser, A., J r. J . Am. Chem. Soc. 1987,
109, 6092.
The existence of suitable absorption spectra for the ion
pairs of these compounds allows the use of the double
(

6
356 J . Org. Chem., Vol. 61, No. 18, 1996
Krom and Streitwieser
F igu r e 2. Concentration dependence of the observed cesium ion pair acidities of the amides.

Cesium Ion Pair Acidities of N,N-Dialkylacetamides
J . Org. Chem., Vol. 61, No. 18, 1996 6357
monomeric at concentrations of about 10- M. The
acidities of 1 and 2 increase slightly with increasing ion
pair concentration (Figures 2A and 2B), consistent with
enolate aggregation. The acidity of 3 also increases, but
the acidity change over the entire concentration range
is about the same magnitude as the experimental error
(Figure 2C). Hypothetical aggregation equilibrium con-
stants were calculated from the acidity data according
4
Ta ble 2. Hyp oth etica l Aggr ega tion Con sta n ts
Ca lcu la ted fr om Sp ectr oscop ic a n d Acid ity Da ta
calcd aggregation constant
assumed
from
spectroscopy
from
acidity
enolate pKCsa equilibrium
b
470 M-1
470 M-1
Cs-1 24.87
Cs-2 25.26
Cs-3 25.48
Cs-4 26.09
Cs-5 25.85
M/Di
M/Tr
M/Te
M/Di
M/Tr
M/Te
M/Di
M/Tr
M/Te
M/Di
M/Tr
M/Te
M/Di
M/Tr
M/Te
6
-2
5
-2
-3
3.0 × 10 M
5.5 × 10 M
1
0
-3
8
1.4 × 10
M
9.4 × 10 M
-
1
-1
33
440 M
350 M
to a published procedure; the results are collected in
6
-2
5
-2
-3
3.2 × 10 M
4.9 × 10 M
Table 2. Comparison with the aggregation constants
calculated from the spectroscopic data shows best agree-
ment for the monomer/dimer model in each case, but the
discrepancies between the two analyses are rather large
for Cs-2 and Cs-3. We conclude that the cesium enolates
of the biphenylyl-substituted acetamides are slightly
aggregated, probably to ion pair dimers; the data for the
enolate of 3 are consistent with aggregation, but the
putative dimerization constant is barely significant when
compared to experimental error. The higher tendency
of the enolates from the monosubstituted acetamide 1 to
aggregate when compared to the disubstituted isomer 3
is probably associated with steric effects.
1
0
-3
8
1.6 × 10
M
9.6 × 10 M
-
1
-1
120 M
87 M
5
-2
-3
5
-2
-3
6.8 × 10 M
1.1 × 10 M
9
8
2.0 × 10 M
2.0 × 10 M
-
1
83 M
c
5
-2
-3
4.4 × 10 M
9
1.0 × 10 M
-
1
69 M
c
5
-2
-3
6.6 × 10 M
9
2.1 × 10 M
a
True cesium ion pair acidity of the conjugate acid with respect
to the monomeric ion pair. The pKs are statistically corrected per
b
acidic hydrogen. M/Di, monomer-dimer; M/Tr, monomer-tri-
c
mer; M/Te, monomer-tetramer. The acidity data are not con-
sistent with aggregation hypotheses. See text for discussion.
The apparent decreases in the acidities of 4 and 5
(
Figures 2D and 2E) are not consistent with aggregation
cies. The shifts in the spectra of the 2,2-diphenylacet-
amide enolates are not obvious, and the SVD output
indicates that they are only slightly greater than the
noise of the instrument. Although the spectroscopic
shifts of these ion pairs are consistent with aggregation,
the experimental results are not compelling. Addition-
ally, we shall see below that the spectra of the structur-
ally related 4 and 5 enolates change under the conditions
of the acidity experiments.
The acidity data, discussed below, clearly indicate that
the aryl-substituted ion pairs are mainly monomeric at
the concentrations in these experiments (about 10-4 M).
In the analyses of the spectra, therefore, we only con-
sidered aggregation models that include the monomer.
Hypothetical spectra of the aggregates were constructed
from the SVD output according to monomer/dimer,
monomer/trimer, and monomer/tetramer aggregation
models, as mentioned above; data for the hypothetical
spectra and equilibrium constants are presented in Table
and are probably artifacts caused by anomalies in the
spectra of their enolates; the shapes of the absorption
bands of the enolates alone or in the presence of their
conjugate acids are slightly different than in the acidity
experiments, i.e., in the presence of the t-BuFl cesium
ion pair. Consequently, small errors in the concentration
determinations are caused by the inability to precisely
model the equilibrium spectra from the spectra of the
separate components. The errors are relatively larger
at the lower concentrations, and we assume that the
measurements at higher concentrations are more ac-
curate. We considered the possibility that the spectro-
scopic anomalies in the equilibrium measurements are
caused by mixed aggregate formation between the eno-
late and the indicator ion pair, but the spectrum of the
indicator ion pair is unchanged in the acidity experiment.
Although we have no explanation for the anomalies in
their spectra, these enolates seem to be essentially
monomeric.
As expected from previous studies of similar systems,38
an aryl substituent on the acetamide enolate stabilizes
it with respect to aggregation by delocalizing the charge.
The slightly greater tendency of the biphenylylacetamide
enolates to aggregate when compared with their diphenyl
isomers is perhaps due to subtle steric effects. Another
possible explanation is that a small population of the
biphenylylacetamide enolates adopts the E configuration,
which would be expected from steric considerations to
dimerize more readily; of course, a similar configuration
is not available to the diphenyl systems. NMR spectro-
scopic studies suggest that the enolate of N,N-dimeth-
2
. The hypothetical aggregation constants for the di-
phenylacetamide enolates are prone to error because we
were not able to find isosbestic points in their absorption
spectra; thus, we have simply used the extinction coef-
ficient at the λmax in the calculations.
We now turn to the discussion of the acidity experi-
ments. Inspection of Figure 2F shows that the acidity
of 6 depends on the formal ion pair concentration; the
quantitative treatment by the usual method35 gives an
average aggregation number of 3.3 for the cesium enolate
of 6 (Cs-6). Note that this method gives only average
values, so the value of 3.3 does not necessarily indicate
a trimer. Comparison with literature data for similar
compounds suggests that Cs-6 is probably a mixture of
dimers and tetramers. Trimeric species are known,
however, for some alkali amides in the absence of
donating ligands, both in solution and in crystal struc-
tures.3
ylphenylacetamide assumes only the Z configuration in
DMSO solution,3
9,40
but this free enolate ion is not
aggregated under these conditions. Under our condi-
tions, the E isomer could be stabilized by aggregation.
We note that small amounts of E configured amide
6,37
4
1
enolates have been observed in certain cases.
In a
The acidities of the aryl-substituted acetamides are
nearly independent of concentration, showing that the
cesium enolates of these compounds are essentially
(
38) Kaufman, M. J .; Gronert, S.; Bors, D. A.; Streitwieser, A., J r.
J . Am. Chem. Soc. 1987, 109, 602.
39) Bobylev, V. A.; Petrov, M. L.; Komarov, V. Y.; Vasil’ev, V. V.;
(
(
36) Mootz, D.; Zinnius, A.; B o¨ ttcher, B. Angew. Chem., Int. Ed. Engl.
969, 8, 378.
37) Lucht, B. L.; Collum, D. B. J . Am. Chem. Soc. 1994, 116, 7949.
Ionin, B. I. Zh. Org. Khim. 1983, 19, 20.
(40) Bradamante, S.; Pagani, G. A. J . Chem. Soc., Perkin Trans. 2
1986, 1035.
1
(

6
358 J . Org. Chem., Vol. 61, No. 18, 1996
Krom and Streitwieser
reasonable structure for the dimer, such as 7, the E
that an extra aryl substituent is actually destabilizing.
phenyl group is far from the coordination center.
For example, Bordwell found that R,R-diphenylacetophe-
4
6
none is less acidic than R-phenylacetophenone. Since
the steric environment in the acetophenones is similar
to the acetamides studied here, we assume that a similar
mechanism operates. Further evidence is provided by
altering the substitution at nitrogen. For the biphenyl-
ylacetamides, replacement of dimethylamino by diethyl-
amino is deacidifying by 0.4 pK unit in the biphenylyl-
acetamides and 0.6 pK unit in the diphenylacetamides.
Although the difference of 0.2 pK unit is small, it is
probably genuine since locking the ethyl groups into a
pyrrolidino ring (5) is acidifying by just that amount.
We were surprised to find that the amide enolates
examined in this study seem to be less aggregated than
ketone enolates of similar structure.18 For example, the
cesium enolates of 1,3-di(4-biphenylyl)acetone and 1,1,3,3-
tetraphenylacetone have dimerization constants of 595
-1
and 570 M , respectively. On the basis of electronic and
steric considerations, the cesium enolates of R,R-diphen-
ylacetophenone and of 3 offer the best comparison; the
Con clu sion
Cesium ion pair acidity and UV-visible spectroscopic
studies have established that the cesium enolates of N,N-
dialkyl(4-biphenylyl)acetamides aggregate in THF solu-
tion, forming small amounts of dimeric species at con-
-
1
ketone enolate has a dimerization constant of 2310 M
,
but the dimerization constant of the amide enolate is
-1
<
100 M . We have no satisfactory explanation for these
observations.
-
4
centrations of about 10 M. Similar studies of cesium
enolates of N,N-dialkyldiphenylacetamides show that
their dimerization constants must be less than about 100
Ion pair aggregation causes the apparent pK to de-
crease from the true value defined relative to monomeric
species. Thus, the true cesium ion pair pK of N,N-
diethylacetamide is greater than 33.7, consistent with
Bordwell’s estimate of 34-35 for the pK of N,N-dimeth-
ylacetamide in DMSO.42 For further comparison, Fraser
and co-workers determined the pK of N-methylpyrroli-
done in THF with lithium counterion to be less than 33.8
on their acidity scale.43 After adjusting their scale so that
-
1
M
. The greater tendencies of the biphenylacetamide
enolates to aggregate is probably due to lower steric
demand when compared with their diphenyl isomers,
possibly because they have a less sterically demanding
E-configuration available. The cesium enolate of N,N-
diethylacetamide is rather highly aggregated, probably
forming a mixture of dimers and tetramers. For reasons
that are not clear, the aryl amide enolates are less
aggregated than structurally related ketone enolates.
The cesium ion pair acidities of the arylacetamides are
similar, with pKCs values of about 25-26 at 25 °C. N,N-
Diethylacetamide has a pKCs of greater than 33.7. The
acidities determined here are in good agreement with
Bordwell’s measurements in DMSO42 and Fraser’s mea-
surement of N-methylpyrrolidone with lithium counter-
ion in THF.43 The data suggest that the effective C-H
acidities of simple amides are about 35 on the cesium
ion pair scale and about 30 on the lithium ion pair scale.
the pK of triphenylmethane coincides with its value on
our lithium ion-pair acidity scale,4
4,45
we find that N-
methylpyrrolidone would have a pK of less than 30.2 on
our lithium scale. This is in rough accord with our
measurement of the cesium ion pair acidity of N,N-
diethylacetamide because, for contact ion paired species,
lithium pK values are generally lower by 3-7 pK units
when compared with their cesium pK values.3
5,38
Thus,
it seems safe to conclude that simple N,N-dialkylamides
have effective pKs of about 35 and 30 in THF on the
cesium and lithium scales, respectively, but we reem-
phasize that the acidities will be concentration-dependent
because of ion pair aggregation.
The cesium ion pair pK of 1 (24.87) compares favorably
with the ionic pK of N,N-dimethylphenylacetamide in
DMSO42 (pK ) 26.9 when statistically corrected per
hydrogen), with some of the difference being due to the
extra acidifying effect of biphenylyl compared with phen-
yl.
The aryl-substituted acetamides have similar acidities,
as might have been anticipated from acidity measure-
ments of the aryl-substituted acetones.18 In that study,
however, it was found that diphenyl substitution is more
acidifying than biphenylyl (1,3-di(4-biphenylyl)acetone,
pKCs ) 17.10; 1,1,3,3-tetraphenylacetone, pKCs ) 16.57),
just the opposite of the findings for the acetamides. In
fact, diphenyl is usually more acidifying than either
monophenyl or biphenylyl, but exceptions occur when the
steric congestion in the anion is great enough that the
aryl substituents are forced out of planarity to the extent
Exp er im en ta l Section
With the exception of 1, all of the amides are known
compounds. They were prepared by the reaction of the
corresponding acyl chloride with an excess of the amine. The
solid amides were purified by recrystallization from ethanol
or diethyl ether, and N,N-diethylacetamide was purified by
distillation through a spinning-band fractionating column.
N,N-Diethyl-2-(4-biphenylyl)acetamide is an oil and was puri-
fied by column chomatography over silica gel. The purification
of THF was accomplished by described procedures, and our
methods of acidity determination, which make use of an inert
atmosphere glovebox and fiber-optic UV-visible spectropho-
tometer, have been published.33
N,N-Dim eth yl-2-(4-bip h en ylyl)a ceta m id e (1). The com-
mercial reagents were used without purification. A solution
of 3.5 g (16 mmol) of biphenylylacetic acid (Aldrich) in 25 mL
of thionyl chloride (Aldrich) was gently warmed under a drying
tube until the effervescence of HCl was evident (about 60 °C).
The reaction mixture was stirred for an additional 2 h, and a
solution of dimethylamine in diethyl ether was prepared by
extracting a 40% aqueous solution of the amine with diethyl
ether. The reaction mixture was allowed to cool, and the
thionyl chloride was removed by distillation under reduced
pressure. The crude acyl chloride was dissolved in diethyl
(41) Heathcock, C. H.; Buse, C. T.; Kleschick, W. A.; Pirrung, M.
C.; Sohn, J . E.; Lampe, J . J . Org. Chem. 1980, 45, 1066.
(
(
42) Bordwell, F. G.; Fried, H. E. J . Org. Chem. 1981, 46, 4327.
43) Fraser, R. R.; Bresse, M.; Mansour, T. S. J . Chem. Soc., Chem.
Commun. 1983, 620.
44) Gronert, S.; Streitwieser, A., J r. J . Am. Chem. Soc. 1986, 108,
016.
45) Kaufman, M. J .; Gronert, S.; Streitwieser, A., J r. J . Am. Chem.
Soc. 1988, 110, 2829.
(
7
(46) Bordwell, F. G.; Bares, J . E.; Bartmess, J . E.; McCollum, G. J .;
Van Der Puy, M.; Vanier, N. R.; Matthews, W. S. J . Org. Chem. 1977,
42, 321.
(

Cesium Ion Pair Acidities of N,N-Dialkylacetamides
J . Org. Chem., Vol. 61, No. 18, 1996 6359
ether and carefully added to the vigorously agitated ethereal
dimethylamine solution. The mixture was filtered, and the
solids were washed with two portions of ether. The combined
ethereal extracts were dried over MgSO , and the solvent was
4
removed by rotary evaporation. The crude solid (3.4 g) was
ylmethyl)cesium was added to the cell until the red-brown color
of cesiated PSX persisted, and the spectrum was obtained. The
solution was successively diluted with known amounts of THF,
and the spectrum was obtained after each dilution. The color
of the cesiated PSX persisted throughout the experiment. The
spectrum of the end point indicator was subtracted from each
spectrum, and the extinction coefficient at each concentration
was calculated in a straightforward manner.
recrystallized twice from diethyl ether to give pure amide: mp
1
8
(
8-89 °C; H NMR (CDCl
3
) δ 3.01 (s, 3 H), 3.10 (s, 3 H), 3.76
s, 2 H), 7.3-7.7 (m, 9 H). Anal. Calcd for C16
H, 7.16; N, 5.85. Found: C, 80.34; H, 7.37; N, 6.09.
H
17NO: C, 80.30;
Extin ction Coefficien t Mea su r em en ts. A solution of
known concentration of the amide in THF was prepared in a
UV cuvette, and a tiny amount (about 50 µg) of 9-phenylthio-
xanthene (PSX) was added as an end point indicator. (Diphen-
Ack n ow led gm en t. This research was supported in
part by NSF Grants 92-21277 and 95-28273.
J O960687Z