Ion pair properties of lithium and cesium salts of carbazole

Article abstract of DOI:10.1021/ja972839t

The lithium salt of carbazole, LiCb, is monomeric in THF in the concentration range (1-5) x 10^-3 M. It is a contact ion pair (CIP) with an effective ion pair pK of 13.48 on a scale where fluorenyllithium solvent separated ion pairs (SSIP) are 22.90 (per hydrogen). The cesium salt, CsCb, is a mixture of monomer and dimer with K₂ = 300 M^-1. The monomer forms a 1:1: complex, CsCb·CbH, with K(c) = 62 M^-1. The ion pair pK of monomeric-CsCb is 19.24.

Full text of DOI:10.1021/ja972839t

1718
J. Am. Chem. Soc. 1998, 120, 1718-1723
Ion Pair Properties of Lithium and Cesium Salts of Carbazole1
Alessandro Abbotto, Andrew Streitwieser,* Manolis Stratakis, and James A. Krom
Contribution from the Department of Chemistry, UniVersity of California, Berkeley, California 94720-1460
ReceiVed August 13, 1997
Abstract: The lithium salt of carbazole, LiCb, is monomeric in THF in the concentration range (1-5) × 10-3
M. It is a contact ion pair (CIP) with an effective ion pair pK of 13.48 on a scale where fluorenyllithium
solvent separated ion pairs (SSIP) are 22.90 (per hydrogen). The cesium salt, CsCb, is a mixture of monomer
and dimer with K2 ) 300 M-1. The monomer forms a 1:1 complex, CsCb‚CbH, with Kc ) 62 M-1. The ion
pair pK of monomeric CsCb is 19.24.
Introduction
t-BuOLi/t-BuOH in THF.8 The X-ray structures of cesium and
potassium carbazide in the presence of a donating ligand were
also obtained; in both cases a dimeric aggregate was found.9
Finally, the acidity of carbazole has been reported in several
media: DMSO (pK ) 19.9,10 19.611), DME (pK ) 14.6 (Li),
18.6 (Cs)),12 and H2O (pK ) 21.1,13 17.06,14 15.1615).
The crystal structures do not relate to the state of aggregation
of the carbazole salts in solution and, in fact, information
relevant to the possible aggregation of carbazide ion pairs in
solution is lacking. We recently showed that the cesium salt
of diphenylamine has a greater tendency to aggregate than the
lithium salt and that the monomer and dimer of the cesium salt
have measurably different absorption spectra.16 In the present
study, we compare these properties for carbazole, a closely
related but more acidic amine. The aggregation states of the
lithium and cesium salts of carbazole (LiCb and CsCb,
respectively) have been studied in THF and the equilibrium
constants have been determined. We also report the corre-
sponding ion pair acidities of carbazole (CbH) for comparison
with previous results. These data further expand the lithium16,17
and cesium5,16,17b,18 ion pair acidity scales.
The anion of carbazole has been the subject of various
experimental and theoretical investigations over the past two
decades. It is well-known that in ethereal solvents the status
of carbanions and nitranions can be generally described in terms
of an equilibrium between contact ion pairs (CIP) and solvent-
separated ion pairs (SSIP), the actual position of the equilibrium
depending on many factors such as the anion structure, the
counterion, solvent, concentration, and temperature.2 Optical
absorption and emission experiments have shown that the
potassium, sodium, and lithium salts of carbazole, indole, and
4,5-iminophenanthrene exist entirely as CIP in tetrahydrofuran
(THF) and 1,2-dimethoxyethane (DME), whereas more polar
solvents such as hexamethylphosphoric triamide (HMPT) are
needed to observe SSIP.3 By comparison, the lithium salt of
the carbon analogue of carbazole, fluorene, was reported in 1965
by Hogen-Esch and Smid to be a mixture of CIP and SSIP in
THF solution.4 In contrast, the large cesium ion always prefers
the formation of CIP.5
1H, 13C, and Li NMR spectroscopy have also been used to
7
gain information concerning the ion pairing and structure of
the carbazide salts.3c,6 In particular, 7Li NMR spectroscopy has
proven to be a useful method to determine the σ-like or π-like
interaction between the cation and the anionic substrate in the
CIP. The results showed that the nitranions of carbazole and
indole are associated with the cation through their nitrogen σ
lone pairs, whereas fluorenyl anion involves several carbon sites
in coordination with lithium, thus forming a π-like complex.
Recently, the Schleyer group found that lithium carbazide
crystallizes as a dimer from a THF solution when n-BuLi is
used as a base,7 whereas a monomeric species is obtained from
Results
UV-Visible Absorption Spectra of the Anions. The
lithium and cesium salts of carbazole were obtained in THF
(8) Lambert, C.; Hampel, F.; Schleyer, P. v. R. Angew. Chem., Int. Ed.
Engl. 1992, 31, 1209.
(9) Gregory, K.; Bremer, M.; Schleyer, P. v. R.; Klusener, P. A. A.;
Brandsma, L. Angew. Chem., Int. Ed. Engl. 1989, 28, 1224.
(10) (a) Bordwell, F. G.; Drucker, G. E.; Fried, H. E. J. Org. Chem.
1981, 46, 632. (b) Bordwell, F. F. Acc. Chem. Res. 1988, 21, 456.
(11) Terekhova, M. I.; Petrov, E. S.; Rokhlina, E. M.; Kravtsov, D. N.;
Shatenshtein, A. I. Khim. Geterotsikl. Soedin. 1979, 1104.
(12) Petrov, E. S.; Terekhova, M. I.; Basmanova, V. M.; Shatenshtein,
A. I. Zh. Org. Khim. 1980, 16, 2457. The reported acidities are normalized
to the DMSO value of fluorene (22.9).
(1) Carbon Acidity. 98. For part 97 see: Streitwieser, A.; Schriver, G.
W. Heteroatom Chem. 1997, 8, 533-7.
(2) For a review, see: Ions and Ion Pairs in Organic Reactions; Szwarc,
M., Ed.; Wiley: New York, 1972.
(13) Schulman, S. G.; Capomacchia, A. C. Anal. Chim. Acta 1972, 59,
471.
(3) (a) Vos, H. W.; Blom, H. H.; Velthorst, N. H.; MacLean, C. J. Chem.
Soc., Perkin Trans. 2 1972, 635. (b) Vos, H. W.; MacLean, C.; Velthorst,
N. H. J. Chem. Soc., Faraday Trans. 2 1976, 72, 63. (c) Velthorst, N. H.
Pure Appl. Chem. 1979, 51, 85.
(4) (a) Hogen-Esch, T. E.; Smid, J. J. Am. Chem. Soc. 1965, 87, 669.
(b) Hogen-Esch, T. E.; Smid, J. J. Am. Chem. Soc. 1966, 88, 307. (c) Hogen-
Esch, T. E.; Smid, J. J. Am. Chem. Soc. 1966, 88, 318.
(5) Bors, D. A.; Kaufman, M. J.; Streitwieser, A. J. Am. Chem. Soc.
1985, 107, 6975.
(14) Gaboriaud, R.; Halle, J.-C.; Letellier, P. Bull. Soc. Chim. Fr. 1976,
1093.
(15) Balo´n, M.; Carmona, M. C.; Mun˜oz, M. A.; Hidalgo, J. Tetrahedron
1989, 45, 7501.
(16) Krom, J.; Petty, J. T.; Streitwieser, A. J. Am. Chem. Soc. 1993,
115, 8024.
(17) (a) Gronert, S.; Streitwieser, A. J. Am. Chem. Soc. 1986, 108, 7016.
(b) Kaufman, M. J.; Gronert, S.; Streitwieser, A. J. Am. Chem. Soc. 1988,
110, 2829.
(6) Cox, R. H. Can. J. Chem. 1971, 49, 1377.
(7) Hacker, R.; Kaufmann, E.; Schleyer, P. v. R.; Mahdi, W.; Dietrich,
H. Chem. Ber. 1987, 120, 1533.
(18) (a) Streitwieser, A.; Bors, D. A.; Kaufman, M. J. J. Chem. Soc.,
Chem. Commun. 1983, 1394. (b) Streitwieser, A.; Ciula, J. C.; Krom, J.
A.; Thiele, G. J. Org. Chem. 1991, 56, 1074.
S0002-7863(97)02839-4 CCC: $15.00 © 1998 American Chemical Society
Published on Web 03/04/1998

Ion Pair Properties of Li and Cs Salts of Carbazole
J. Am. Chem. Soc., Vol. 120, No. 8, 1998 1719
Table 1. Aggregation Equilibria of Cesium Carbazide CsCb
K2
2(CsCb)1 y\z (CsCb)2
a
[(CsCb)1]
(10-4 M)
[(CsCb)2]
(10-4 M)
K2
(102 M-1
)
25.2
21.1
17.5
15.4
14.2
12.9
12.2
11.0
9.8
19.1
13.3
9.14
7.13
6.08
4.99
4.45
3.61
2.85
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
3.0
a K2 ) [(CsCb)2]/[(CsCb)1]2.
present in solution. These spectra are obtained from a linear
transformation of the basis spectra as described previously.16
The resulting spectra and extinction coefficients of the two
species in solution provide their concentrations at each deter-
mination. The assumption that the two species are the monomer
(CsCb)1 and dimer (CsCb)2 is consistent with the results of
independent ion pair acidity studies described below.
Figure 1. Concentration dependence of the UV-vis spectrum of CsCb
in THF at 25 °C showing the equilibrium monomer/dimer. Concentra-
tions range from 6.365 × 10-3 (top spectrum) (λmax ) 401.5 nm) to
1.565 × 10-3 M (bottom spectrum) (λmax ) 407.0 nm) in a 1-mm cell.
solution by using fluorenyllithium and diphenylmethylcesium
as the base, respectively. All of the UV-visible experiments
were carried out in THF solution at 25.0 ( 0.1 °C. The window
of study of the anions was limited at low wavelengths to 390
nm, since neutral carbazole starts absorbing significantly below
this value.
The λmax of the monomer (CsCb)1 and of the dimer (CsCb)2
were found to be 408.5 and 401.5 nm, respectively. At 404.5
nm the extinction coefficient for the monomeric species is ꢀ )
ꢀob ) 2110 M-1 cm-1, while the dimeric species has ꢀ ) 2ꢀob
) 4220 M-1 cm-1. By using the actual spectra of the monomer
and the dimer, the nine experimental spectra were deconvoluted
and the concentrations of each species were determined at each
step. The data are summarized in Table 1. From these data
the equilibrium constant K2 for reaction 1 was determined to
The absorption spectrum of LiCb does not depend on the
concentration of the anion in the range of concentrations used
((1 × 10-3) ÷ (5 × 10-3) M). The λmax was found to be 392.0
nm (ꢀ ) 2860 ( 20 M-1 cm-1). By contrast, the spectrum of
CsCb was found to be dependent on the concentration of the
anion CsCb as well as the presence of neutral CbH. Because
of the latter dependence the absorption of CsCb was studied
first in the absence of neutral CbH. The λmax gradually shifted
from 401.5 to 407 nm as the concentration of CsCb was changed
from 6.365 × 10-3 to 1.565 × 10-3 M. At 404.5 nm the value
of the extinction coefficient was found to be independent of
be 3.0 × 102 M-1
.
2(CsCb)1 h (CsCb)2
(1)
Interaction of CsCb with the Neutral CbH: Formation
of the Adduct CsCb‚CbH. The UV-vis spectrum of CsCb
was found to be dependent on the presence of the neutral as
well. This result shows that an interaction between the nitranion
Cb- and neutral CbH occurs in THF solution with the formation
of a complex, C. Our goal was to investigate and determine
the stoichiometry of the adduct between the anion and the neutral
and the corresponding complexation constant.
the formal concentration of CsCb, ꢀ ) 2110 ( 4 M-1 cm-1
These spectra are shown in Figure 1.
.
The dependence of the shape of the absorption spectrum on
the concentration of CsCb is indicative that more than one
species (aggregates), each with a different absorption spectrum,
exist in THF solution. To investigate the number and nature
of the aggregates we submitted the data for the absorption of
CsCb at nine different concentrations to “singular-value de-
composition” (SVD) analysis in the same way as applied
previously to the cesium salt of diphenylamine.16 The experi-
mental data (values of absorbance for each determination at 0.5
nm intervals over the range 500.0-390.0 nm) were decomposed
leading to nine spectral bases, each associated with a “singular
value”, S. The singular values constitute the “weighting factors”
by which each basis describes the experimental set of data. Two
main singular values were found, S1 ) 12.15 and S2 ) 0.43.
The remaining seven values were found to be between 0.01
and 0.004 and only describe noise. That is, the linear combina-
tion of two basis spectra are able to reproduce the nine original
experimental spectra, demonstrating that only two distinguish-
able species are present in solution whose relative concentrations
differ in each determination.
To analyze the experimental spectra of the CsCb in the
presence of the neutral CbH we start by assuming that the
complex is an adduct between one molecule of the salt CsCb
and one molecule of the neutral CbH. This model is obviously
the simplest and most likely case. The model was tested for
consistency with results from spectroscopic and ion pair acidity
studies. By addition of increasing amounts of neutral carbazole
to a solution of CsCb in THF the absorption spectrum was found
to shift toward the blue. A 5 × 10-3 M solution of CsCb in
THF treated with a large excess of the neutral (final concentra-
tion was 0.7-1 M) gave rise to a spectrum that was almost
insensitive to further additions of CbH. At this point the
spectrum is mostly that of the complex. A set of different
experiments in which the starting concentration of CsCb was
varied from 1.51 × 10-3 to 8.59 × 10-3 M and then treated
with excess neutral carbazole to the order of the concentrations
above showed that each of the final spectra cannot be considered
due only to the absorption of the complex, but that an unknown
amount of free CsCb is still present; that is, the conversion to
the complex is not complete. Unfortunately, further addition
of CbH is limited by its solubility in THF. A procedure of
The two basis spectra have only a linear-algebraic meaning
and are not the actual UV-visible spectra of the two species

1720 J. Am. Chem. Soc., Vol. 120, No. 8, 1998
Abbotto et al.
successive interations was therefore used to obtain the spectrum
of the complex and the equilibrium constant, Kc, for reaction
2.
(CsCb)1 + CbH a C
(2)
When both the salt CsCb and the neutral CbH are present in
solution, from eqs 1 and 2 the concentrations of all the species
are as follows:
[(CsCb)1] ) {CsCb}o - [C] - 2[(CsCb)2]
(3a)
[(CsCb)2] ) K2[(CsCb)1]2
[CbH] ) [CbH]o - [C]
(3b)
(3c)
where [C] ) concentration of the complex, [CbH] ) concentra-
tion of the uncomplexed neutral carbazole, {CsCb}o ) initial
total formal concentration of cesium carbazide, [CbH]o ) initial
concentration of carbazole.
If we know the concentrations {CsCb}o and [CbH]o and the
values of K2 and Kc, we are therefore able to determine [(CsCb)1]
and [(CsCb)2] and, by using the actual spectra of the monomer
and of the dimer, to obtain the spectrum of the complex by
subtraction of these components. We call experiment A that
in which the complexation equilibrium is studied and where
the concentrations {CsCb}o and [CbH]o are known quantities.
The difficulty is that Kc is an unknown quantity and cannot be
determined unless the spectrum of the complex is available. The
spectrum obtained by treating a THF solution of CsCb with a
large excess of CbH was taken as a first approximation to the
spectrum of the complex. This spectrum plus the separate
spectra of the cesium carbazide ion pairs and of the indicator
ion pair were used to fit a total of 20 acidity equilibria from
two independent experiments (experiments B and C) (see the
Ion Pair Acidity section). The fitting of the 20 acidity equilibria
allows the determination of the concentrations of all of the
species in solution. In particular, we now know the first
approximations to the concentrations of [(CsCb)1], [CbH], and
[C] in 20 different equilibria. By plotting the values of [C] vs
the values of [(CsCb)1]‚[CbH] a linear correlation was found
(slope ) 76.6, intercept ) 5.6 × 10-6, r2 ) 0.996). The
linearity of the correlation is evidence that the hypothesis of
the stoichiometry of the complex is correct. The slope of the
line provides us with a first approximation to Kc and, by
applying eqs 3a-3c, a first set of values for [C], [(CsCb)1],
and [(CsCb)2] is now available for the complexation equilibrium
(experiment A). In particular, the concentrations of the
monomer and the dimer allow determination of their contribu-
tions to the total spectra of their equilibria with the complex.
These contributions are relatively small and by subtraction
provide the next approximation to the spectrum of the complex
itself. This next approximation to the spectrum of C was used
again to fit the acidity equilibria and a new set of values for
Kc, [C], [(CsCb)1], and [(CsCb)2] was obtained. The procedure
was repeated until self-consistency was reachedsthat is, the
value of Kc coming from two consecutive sets differed by less
than 0.1%. Self-consistency was reached after three iterations.
The results are summarized in Table S2 (Supporting Informa-
tion). The plot of [C] vs [(CsCb)1]‚[CbH] is shown in Figure
S2 (Supporting Information). The separated spectra of the
monomer, dimer, and the complex of cesium carbazide are
shown in Figure 3. The extinction coefficient of the complex
was found to be 845 M-1 cm-1 at 404.5 nm. The complexation
Figure 3. Separated spectra of the monomer (CsCb)1 (spectrum A,
λmax 408.5 nm), the dimer (CsCb)2 (spectrum B, λmax 401.5 nm), and
the complex (CsCb)1‚CbH (spectrum C) in THF at 25 °C ([(CsCb)1]
) [(CsCb)1‚CbH] ) [(CsCb)2]/2 ) 2.522 × 10-3 M in a 1-mm cell).
Table 3. pKa and Extinction Coefficient Measurements of Lithium
Salt Carbazole and of CH Acid Indicators in THFa,b
d
compdc
pKLi/THF
λmax (ꢀ)
CbH
Ph-3,4-BF
13.48
14.88
392(2860)
394(23900)
418(26000)
373(14300)
498(33400)
526(25600)
411(25300)
9-BpFl
16.98
9-PhFle
17.60
a On a per hydrogen basis. b For a lithium ion pair acidity scale of
other CH acids see ref 17b. c Abbreviations: CbH, carbazole; Ph-3,4-
BF, 7-phenyl-3,4-benzofluorene; 9-BpFl, 9-biphenylylfluorene; 9-PhFl,
9-phenylfluorene. d pKa values referenced to the SSIP of fluorene at
22.90. e Reference 17b.
Ion Pair Acidity. The lithium ion pair acidity of CbH was
measured against the indicator 7-phenyl-3,4-benzofluorene (Ph-
3,4-BF, pKLi ) 14.88;17 Table 3), equilibrium 4:
CbH + LiPh-3,4-BF a LiCb + Ph-3,4-BF
(4)
The UV-vis spectra of the lithium salt of the indicator, LiPh-
3,4-BF, and of the lithium salt of carbazole, LiCb, were used
to fit a set of 10 acidity equilibria spectra in the range 650-
390 nm. The concentrations of the anions [LiPh-3,4-BF] and
[LiCb] were determined directly, and those of the neutrals, [Ph-
3,4-BF] and [CbH], by difference from the initial concentrations.
The data are collected in Table S4 (Supporting Information).
A small deviation of the observed constant Kobs for reaction 4
from constancy was attributed to the slight but significant
dissociation of the lithium solvent-separated ion pair of the
indicator to free ions in THF:
LiPh-3,4-BF a Li+ + Ph-3,4BF-
(5)
The dissociation constant Kd for equilibrium 5 has not been
determined in THF but can be reasonably assumed to be
comparable to that of similar carbon acids 9-phenylfluorene and
3,4-benzofluorene in THF (Kd ) 1.0 × 10-5 M at 25 °C). The
spectra of LiPh-3,4-BF SSIP and the free carbanion are
sufficiently similar that a straightforward correction using Kd
constant for reaction 2 is Kc ) 62 M-1
.

Ion Pair Properties of Li and Cs Salts of Carbazole
J. Am. Chem. Soc., Vol. 120, No. 8, 1998 1721
then provides the value of [LiPh-3,4-BF] to be used in eq 4.
The resulting values of Kobs are now constant. The dissociation
of LiCb to free ions is negligible by analogy to that of lithium
diphenylamide (Kdiss ) 10-12-10-11 M),16 and as shown by its
relatively low pK (vide infra). The N-Li association in the
CIP is substantially stronger than the ion pair association in
the indicator SSIP. Moreover, the constancy of the corrected
Kobs confirms that LiCb is present as one species in THF
solutionssthat is, aggregation is negligible at these concentra-
tions. Thus, Kobs is taken as Ko for equilibrium 4; the resulting
mean value of ∆pK (-1.40 ( 0.01) gives the lithium ion-pair
pK of carbazole in THF as 13.48 on a scale in which the
reference is that of fluorenyllithium SSIP whose pK is taken as
22.90 (per H).17
[(CsCb)1][3,4-BF]
Ka )
(10)
[CbH][Cs3,4-BF]
Algebraic manipulation relates the observed concentration of
{CsCb} to the observed quantity {CsCb}/Kob through the
following equation:16
{CsCb} ) Ka({CsCb}/Kob) + 2K2Ka2({CsCb}/Kob)2 (11)
The plot of {CsCb} vs {CsCb}/Kob is shown in Figure S4
(Supporting Information). Statistics of the correlations of the
data reported in Table S5 are shown in Table S6 (Supporting
Information). The dimerization constants from acidity experi-
ments B and C were found to be K2 ) 2.9 × 102 and 2.5 × 102
M-1, respectively, in excellent agreement with the results from
the spectroscopic data, where a monomer/dimer equilibrium was
assumed. The ion pair acidity constant with respect to the
indicator 3,4-BF was found to be Ka ) 0.86, from which the
ion pair pKCs/THF of carbazole was calculated to be, after
statistical correction, 19.24.
The cesium ion pair acidity of carbazole was measured against
the indicator 3,4-benzofluorene (3,4BF, pKCs ) 19.47),5,17b,18
equilibrium 6:
CbH + Cs3,4BF a CsCb + 3,4BF
(6)
The separated spectra of the cesium salt of the indicator, Cs3,4-
BF, and of the monomer (CsCb)1, dimer (CsCb)2, and complex
of the cesium salt of carbazole were used to fit a total of 26
acidity equilibria recorded in two independent experiments. The
concentrations of the ion pair species were determined directly.
The concentration of the neutral indicator [3,4-BF] was obtained
by difference of the concentration of the corresponding anion,
while the value of the neutral uncomplexed carbazole [CbH]
was calculated from eq 7:
Discussion and Conclusions
This UV-vis study has shown some important and unex-
pected differences between the lithium and the cesium salts of
carbazole in THF solution at 25 °C. LiCb was found to be
entirely monomeric whereas CsCb exists as a mixture of three
distinguishable ionic species, each with distinct absorption shape,
the monomer (CsCb)1, dimer (CsCb)2, and an adduct (or
complex) CsCb‚CbH. All of the equilibrium constants among
these species were determined. The dimerization constant (eq
1) is K2 ) 3.0 × 102 M-1 and the complexation constant for
reaction 2 is Kc ) 62 M-1. The comparison of acidity equilibria
with SVD analysis of spectra as a function of concentration
shows that the two distinctive species identified by SVD are
the ion pair monomer and dimer; this result is a further
example16 that demonstrates the power of this combination in
determining both the extent of aggregation and the stoichiometry
of the aggregates even in highly dilute solution. A comparison
with the analogous salts of diphenylamine DPA, LiDPA and
CsDPA,16 is instructive. In both cases the lithium salt is
monomeric whereas the cesium salt consists of monomers and
dimers in the concentration region 10-3-10-4 M. DPA is less
acidic than CbH but its cesium salt is less prone to dimerize
(K2 ) 1.6 × 102 compared to 3.0 × 102 for CsCb); usually, the
more basic anions tend to form the more aggregated ion pairs.
In the present case, the enforced planarity of the carbazole ring
system undoubtedly results in lower steric congestion in the
dimer and leads to a higher dimerization constant although an
alternative explanation is implied below.
One striking difference between the two cesium salts, CsCb
and CsDPA, is the ability of the former to form an adduct with
the neutral species, which was shown to be formed from one
molecule of CsCb and one molecule of CbH. A structural
hypothesis can be forwarded to account for this different
behavior. Ion pairs of nitranions usually coordinate to the
nitrogen lone pair. For CsCb, however, we can postulate that
π-coordination between the pyrrole anion moiety and the large
cesium cation is competitive; in such an ion pair, the nitrogen
lone pairs are now available for hydrogen bonding to a free
carbazole molecule. For the smaller lithium cation, the π-co-
ordinated structure presumably cannot compete with σ-coordi-
nation particularly since the π-coordinated structure has fewer
coordination sites for solvation of lithium cation; in the favored
σ-coordination structure such hydrogen bonding is now blocked.
[CbH] ) [CbH]o - {CsCb} - 2[C]
(7)
where
{CsCb} ) [(CsCb)1] + 2[(CsCb)2]
(8)
is the total concentration of the uncomplexed ion-pair cesium
carbazide. The acidity data are summarized in Table S5
(Supporting Information).
The basis spectra of the monomer (CsCb)1 and of the dimer
(CsCb)2 were obtained as linear combinations of the chemically
meaningless basis spectra derived directly from the SVD output
(U1 and U2).16 The assumption that the two distinct species
that describe the behavior of the uncomplexed cesium carbazide
in solution are actually the monomer and the dimer implies a
stoichiometry. The advantage of using the two derived spectra
is that the concentrations of the monomer and the dimer can be
calculated and therefore the extinction coefficient of each
aggregate and the equilibrium constants of dimerization (K2)
and complexation (Kc) can be determined. Alternatively, the
deconvolution of the experimental UV-vis spectra using the
“direct” components U1 and U2 do not make use of any further
assumptions but have the limitation that only the total formal
concentration of the cesium carbazide ion pairs {CsCb} can be
determined. That is, both the derived spectra assuming the
stoichiometry of (CsCb)1 and (CsCb)2 and the direct use of U1
and U2 lead to the same values of {CsCb}. Next, we use only
the values of {CsCb}, as reported in Table 5, to determine the
relative acidity Ka and the dimerization constant K2.
The observed constant Kobs of equilibrium 6 is given by
{CsCb}[3,4-BF]
Kob
)
(9)
[CbH][Cs3,4-BF]
while the true ion pair acidity equilibrium constant Ka is

1722 J. Am. Chem. Soc., Vol. 120, No. 8, 1998
Abbotto et al.
Table 7. Comparison of pKa of Carbazole (CbH) and Diphenylamine (DPA) in Various Solvents
compd
Li/THFa
Cs/THFa
Li/DMEa,b
Cs/DMEa,b
Li,crypt/THFa,c
DMSO
H2O
CbH
DPA
13.48d
19.05j
19.24d
24.20j
14.6
19.9
18.6
23.9
20.0
19.9,e 19.6f
24.95k
15.16,g 17.06,h 21.1i
22.44, 23.39l
a The reported acidities are referred to the DMSO value of fluorene (22.9). b Reference 12. c THF with lithium cation as the [2.1.1]cryptate: ref
24. d This work. e Reference 10. f Reference 11. g Reference 15. h Reference 14. i Reference 13. j Reference 16. k Bordwell, F. G.; Algrim, D. J. J.
Am. Chem. Soc. 1988, 110, 2964. l Cox, R. A.; Stewart, R. J. Am. Chem. Soc. 1976, 98, 488.
This hypothesis is supported by theoretical studies19 and by some
X-ray structures. The central 6-π-electron ring of the carbazide
anion is understandably better at π-coordination than the
extended π-system of diphenylamide anion. Moreover, the fact
that the λmax of LiCb is at a shorter wavelength than that of
monomeric (CsCb)1 shows that both LiCb and CsCb are CIP
in THF solution.4b,20
The X-ray structure of a cesium carbazide crystal has been
reported.9 From a 0.2 M solution of carbazole in ether in the
presence of t-BuOCs and PMDTA (N,N,N′,N′′,N′′-penta-
methyldiethylenetriamine), crystals of the cesium nitranion were
obtained having a dimeric Ci structure in which the cation is in
an apical position with respect to the π-system of the heterocycle
and simultaneously facing the σ-nitrogen lone pair of the second
molecule. Thus, a π-like and a σ-like interaction occurs
simultaneously between cesium and the carbazide in this
structure. The formation of the dimer (CsCb)2 in THF solution
can be seen as an alternative to complexation but using similar
structural motifs. Furthermore, neutral carbazole itself has been
reported to form strongly hydrogen-bonded dimeric structures
in which the two molecules assume an antiparallel stack
geometry, both theoretically (AM1) and experimentally (fluo-
rescence excitation and emission spectra) ascertained.21 In fact,
one can expect that the deprotonation of one carbazole can lead
to an even more stable hydrogen-bonded structure, with the
atoms of nitrogen of the two carbazole moieties facing one
another in the same plane and sharing an atom of hydrogen.
Lithium carbazide crystallizes from THF as a dimer.7 Although
even in this case simultaneous σ and π interactions between
the cations and the nitranions occurs, unlike other lithium
amides, the lithium cation does not occupy the same apical
position as was found for cesium, but instead a planar rhomboid
N-Li-N-Li configuration is preferred, where each lithium is
at the same position with respect to the two molecules of
nitranions. Furthermore, when t-BuOLi is used as base, a
monomeric crystal was obtained from THF, in which the lithium
coordinates the nitrogen σ lone pair.8 In this case, the additional
coordination to t-BuOH and t-BuO- competes effectively with
another LiCb ion pair. These results support the hypothesis
that at low concentrations in THF the lithium is bonded to the
σ-nitrogen lone pair, and blocks the formation of a hydrogen
bond complex with a molecule of CbH.
Figure 5. Effect of HMPA on the spectrum of LiCb in THF: (a) in
pure THF; (b) in 2 mol % HMPA; (c) in 17 mol % HMPA.
ion are clearly decisive in promoting such higher coordination.
Nevertheless, the lithium coordination within these aggregates
is that of a tight CIP. In LiCb, the lithium is tightly coordinated
to the nitrogen and the second nitrogen lone pair is conjugated
within the carbazole π-system and less available for aggregation.
The strong N-Li interaction in the LiCb CIP is also indicated
by its lower basicity relative to CsCb and to the ionic acidity
in DMSO (see Table 7). When the lithium cation is cryptated,
the ion pair is now effectively an SSIP and the corresponding
acidity is lower by more than 6 pK units24 (see Table 7).
These results may be compared with a limited study of the
effect of hexamethylphosphoramide (HMPA) on the LiCb in
THF as shown in Figure 5. Addition of 2 mol % HMPA results
in a modest change in spectrum; the λmax increases to 397 nm
and may reflect peripheral solvation of lithium in the CIP by
HMPA. Addition of more HMPA to a mole fraction of 0.17
causes appearance of a new peak at λmax 424 nm that could
now be the solvent separated ion pair. At this point 3,4-
benzfluorene (Li pK 19.3)17 is completely deprotonated; thus,
LiCb with HMPA is effectively much more basic than in THF
alone. A more precise comparison is inappropriate since the
solvent has been changed significantly from that of pure THF.
The pK value for CsCb compares well with the free ion value
in DMSO, confirming a trend already observed for CsDPA16
and for cesium salts generally of carbon acids.5,16,17b,18 The
difference in acidity between diphenylamine and carbazole in
DMSO (5 pK units) is also found for the cesium and lithium
ion pair acidities in THF and DME, showing that the relative
stabilization of carbazide anion is about the same in all of these
The absence of aggregation recorded for LiCb and LiDPA,
at least at the typical concentrations used in our studies, points
up a dramatic difference from the situation occurring in many
enolate lithium salts of comparable steric bulk for which much
higher aggregation numbers are found.22,23 The additional
coordination site provided by the three lone pairs of an oxide
(22) (a) Jackman, L. M.; Lange, B. C. Tetrahedron 1977, 33, 2737-69.
(b) Bauer, W.; Seebach, D. HelV. Chim. Acta 1984, 67, 1972-88. (c) Arnett,
E. M.; Moe, K. D. J. Am. Chem. Soc. 1991, 113, 7288-93. (d) Abbotto,
A.; Streitwieser, A. J. Am. Chem. Soc. 1995, 117, 6358-9. (e) Abu-
Hasanayn, F.; Stratakis, M.; Streitwieser, A. J. Org. Chem. 1995, 60, 4688-
9.
(23) For reviews see: Seebach, D. Angew. Chem., Int. Ed. Engl. 1988,
27, 1624. Jackman, L M.; Bortiatynski, J. AdV. Carbanion Chem. 1992, 1,
45-87.
(24) Gareyev, R. F.; Antipin, I. S.; Konovalov, A. I. Results to be
submitted for publication.
(19) Neuhaus, A.; Streitwieser, A. Paper in preparation.
(20) (a) Carter, H. V.; McClelland, B. J.; Warhurst, E. Trans. Faraday
Soc. 1960, 56, 455. (b) Hoijtink, G. J. Ind. Chim. Belge 1963, 12, 1371. (c)
Buschow, K. H. J.; Hoijtink, G. J. J. Chem. Phys. 1964, 40, 2501. (d)
Buschow, K. H. J.; Dieleman, J.; Hoijtink, G. J. J. Chem. Phys. 1965, 42,
1993. (e) Velthorst, N. H.; Hoijtink, G. J. J. Am. Chem. Soc. 1965, 87,
4529.
(21) Taylor, A. G.; Jones, A. C.; Phillips, D. Chem. Phys., 1989, 138,
413.

Ion Pair Properties of Li and Cs Salts of Carbazole
J. Am. Chem. Soc., Vol. 120, No. 8, 1998 1723
the initial carbazole had been converted completely to CsCb. Known
amounts of THF were added (about 0.3 g each time) and the absorption
spectrum was recorded after each addition. CsCb does not absorb in
the region 455-700 nm. This region of the spectrum was used to
evaluate the component due to the absorption of Cs9-t-BuFl in the nine
original spectra. After subtraction of the absorbance band of Cs9-t-
BuFl, nine spectra of CsCb at different concentrations (from 1.565 ×
10-3 M to 6.365 × 10-3 M) were obtained and processed by SVD.
The same set of spectra was used to calculate the extinction coefficient
at 404.5 nm according to Beer’s law (at this wavelength the extinction
coefficient was found to be concentration-independent). SVD analysis
and the extinction coefficient for the lithium salt LiCb were obtained
analogously using fluorenyllithium as a base and its spectrum as end-
point indicator of complete deprotonation of CbH.
cases. The slightly higher difference in acidity observed in the
case of LiCb, with respect to LiDPA, can be interpreted in terms
of a somewhat stronger interaction of the lithium cation with
the nitrogen atom of the anionic substrate in the former case,
perhaps for steric reasons.
Experimental Section
General. Ion pair acidity and spectral studies were carried out in a
Vacuum Atmospheres glovebox under an argon atmosphere. The UV-
vis spectra were recorded at 25.0 ( 0.1 °C on a computer-driven
Shimadzu UV-2101PC spectrometer unit. Details of the apparatus have
been previously described.16 Parameters for UV-vis measurements
were as follows: scan speed 200 nm min-1, slit width 2.0 nm, sampling
interval 0.5 nm, and cell path length 0.1 cm.
Absorption Spectra and Extinction Coefficient of the Complex
(CsCb)1‚CbH (C). The spectra were obtained over the wavelength
range of 396-600 nm. A solution of CbH (0.957 mg, 5.723 × 10-3
mmol) and of 9-t-BuFl (ca. 0.5 mg) in THF (1.115 g, 1.267 mL) was
prepared in a UV cell, and the baseline spectrum was taken. A solution
of CsDPM in THF was added via microsyringe until the absorption of
Cs9-t-BuFl persisted (complete conversion of CbH to CsCb). From
this spectrum the concentration of Cs9-t-BuFl was determined as
described above (9.752 × 10-4 M). Then excess CbH (0.159 g, 0.95
mmol) was added in order to shift the equilibrium toward the formation
of the complex, and the spectrum recorded (this spectrum is constituted
by the absorption bands of the complex (CsCb)1‚CbH and of the
unreacted CsCb). From the concentration of Cs9-t-BuFl, the fraction
of the latest added CbH which was converted to its cesium salt by
reaction with Cs9-t-BuFl was calculated, and the initial concentrations
of CbH and CsCb were exactly determined ({CsCb}o ) 5.49 × 10-3
M, [CbH]o ) 0.750 M). Using the value of Kc as obtained from the
analysis of the ion-pair acidity data (see the Results section), the
concentrations of the monomer (CsCb)1, of the dimer (CsCb)2, and of
the complex were obtained. This treatment allowed to subtract the
absorbance band of CsCb to be subtracted from the original spectrum,
leaving only the absorbance band of the complex, which was used to
determine its extinction coefficient and to deconvolute the spectra of
the ion pair acidity studies.
Materials. Anhydrous THF16 and (diphenylmethyl)cesium (Cs-
DPM)16 were prepared as described previously. Anhydrous hexane
was obtained by distillation over CaH2. The hydrocarbon indicator
acids used in this work either were available from our previous studies
or were synthesized by published procedures. All of the indicators
were carefully purified prior to use by repeated recrystallization
followed by vacuum sublimation.
Carbazole (CbH). Commercial material (99%, Aldrich) was
recrystallized twice from absolute EtOH and sublimed three times under
vacuum (2 × 10-2 Torr) at 70 °C, mp 245-247 °C (lit.25 mp 240-
243 °C). Further evidence of purity was provided by satisfactory
elemental analysis.
Lithium Diisopropylamide (LDA). A 2.16 M solution of n-BuLi
in hexane (2.3 mL, 5.0 mmol) was added to a solution of freshly distilled
diisopropylamine (0.570 g, 5.6 mmol) in anhydrous hexane (5 mL) at
5 °C under nitrogen. After the mixture was stirred for 1 h at room
temperature, the solvent and the excess diisopropylamine were carefully
removed under vacuum to leave a yellow residue, which was then taken
into the glovebox and transferred to a sublimation apparatus. After
two sublimations (90-95 °C at 2 × 10-2 mmHg for 48 h, and 75-80
°C at 10-3-10-4 mmHg for 24 h) a white crystalline solid (0.10 g)
was obtained.
Fluorenyllithium. Commercial fluorene (Aldrich) was recrystallized
twice from absolute EtOH, sublimed twice under vacuum, and taken
into the glovebox. In the glovebox, freshly sublimed LDA (25 mg,
0.23 mmol) was added to a solution of fluorene (55 mg, 0.33 mmol)
in THF (2 mL). The resulting orange mixture was allowed to stand
overnight. The solvent and the formed diisopropylamine were removed
under vacuum, and the residue was taken into the glovebox. THF (2
mL) was finally added to the yellow solid, giving rise to a bright orange
solution (ca. 0.1 M solution of fluorenyllithium in THF). The solution,
which was kept in the glovebox at 4 °C, was stable for at least 1 week.
After this period, it started to turn slightly darker and was discarded.
Absorption Spectra and Extinction Coefficient of Cesium Carb-
azide (CsCb) and Lithium Carbazide (LiCb). The spectra were
obtained over the wavelength range of 390-700 nm. The neutral
carbazole (CbH) does not absorb in this region. A solution of CbH
(0.909 mg, 5.436 × 10-3 mmol) and of 9-tert-butylfluorene (ca. 0.5
mg) (9-t-BuFl, pKCs/THF ) 24.39) in THF (0.752 g, 0.854 mL) was
prepared in a UV cell, and the baseline spectrum was recorded.
Aliquots of a stock solution of CsDPM in THF were added via
microsyringe until the absorption of Cs9-t-BuFl persisted. At this point
Ion Pair Acidity Studies. The procedures used in the acidity
determinations have been previously described in detail.16 The
lithium17b and cesium18b ion pair indicator pK’s used are those of the
revised scales.
Acknowledgment. This research was supported in part by
grants from the National Science Foundation and the National
Institutes of Health. We are also indebted to Dr. Christoph
Lambert for preliminary studies and to NATO for travel funds.
Supporting Information Available: Additional figures (S2,
plot of C as [(CsCb)1][CbH]; S4, determination of Ka and K2
from cesium ion pair acidity studies) and tables (S2, determi-
nation of KC; S4, lithium ion pair acidity of CbH; S5, cesium
ion pair acidity of CbH; S6, equilibrium constants from Cs ion
pair acidity experiments) (5 pages). See any current masthead
page for ordering and Internet access instructions.
(25) Perrin, D. D.; Armarego, W. L. F. Purification of Laboratory
Chemicals; Pergamon Press: Oxford, 1988.
JA972839T