Polar substituent effects in the bicyclo[1.1.1]pentane ring system: Acidities of 3-substituted bicyclo[1.1.1]pentane-1-carboxylic acids

Article abstract of DOI:10.1021/jo040236b

(Chemical Equation Presented) Experimental gas-phase acidities are reported for a series of 3-substituted (X) bicyclo [1.1.1]pent-1-yl carboxylic acids (1, Y = COOH). A comparison with available calculated data (MP2/6-311++G* *// B3LYP/6-311+G**) reveals

Full text of DOI:10.1021/jo040236b

Polar Substituent Effects in the Bicyclo[1.1.1]pentane Ring
System: Acidities of 3-Substituted
Bicyclo[1.1.1]pentane-1-carboxylic Acids
William Adcock,*,† Yakup Baran,‡ Antonello Filippi,§ Maurizio Speranza,*,§ and Neil A. Trout†
School of Chemistry, Physics, and Earth Sciences, The Flinders University of South Australia,
Adelaide, Australia 5001, the Department of Chemistry, Karaelmas University, 67100 Zonguldak, Turkey,
and the Dipartimento di Studi di Chimica e Tecnologia delle Sostanze Biologicamente Attive,
Universita` di Roma “La Sapienza”, 00185 Roma, Italy
william.adcock@flinders.edu.au
Received July 19, 2004
Experimental gas-phase acidities are reported for a series of 3-substituted (X) bicyclo [1.1.1]pent-
1-yl carboxylic acids (1, Y ) COOH). A comparison with available calculated data (MP2/6-
311++G**// B3LYP/6-311+G**) reveals good agreement. The relative substituent effects are shown
to be adequately described by a much lower level of theory (B3LYP/6-31+G*). Various correlations
are presented which clearly point to polar field effects as being the origin of the relative acidities.
Introduction
Over the years, cage polycyclic alkanes have been
deployed extensively as model substrates for the inves-
tigation of several important concepts in chemistry. In
particular, their rigid molecular frameworks coupled with
the absence of π-electrons provide ideal scaffolds for
evaluating the nature and magnitude of remote polar
substituent effects in saturated systems. Two of the most
significant systems of this kind are 1,3-disubstituted
bicyclo[1.1.1]pentanes (1) and 1,4-disubstituted bicyclo-
[2.2.2]octanes (2) in which the substituent (X) and probe
(Y) located at termini bridgehead positions are both
aligned along the major axis of the ring system. Hence,
these cage systems allow accurate measurement of
distances between substituent (X) and probe (Y) without
uncertainty being created by an angle orientation fac-
tor.
C6H5,11 CH3,14 p-FC6H4,11 and FC8H1211), and the results
of these studies have provided insightful complementary
information on the modes of transmission (“through-
bond” and “through-space”) of polar substituent effects.
Recently, Wiberg3 reported calculated acidities of 1 and
2 (Y ) COOH) at a high level of theory (MP2/6-311+ +
G**//B3LYP/6-311+G**) for a number of substituents
covering a reasonable range of polar effects. Whereas the
results of the latter system were completely validated by
comparison and correlation against appropriate available
Both chemical reactivity and NMR chemical shift
probes (energy and charge density monitors, respectively)
have been deployed (1, Y ) COOH,1-3 F,4,5 Sn(CH3)3,6 and
2, Y ) COOH,3,7-9 NH3+10, F,5,11 Sn(CH3)3,6,12 C≡CH,13
(1) Applequist, D. E.; Renken, T. L.; Wheeler, J. W. J. Org. Chem.
1982, 47, 4985.
(2) Adcock, W.; Blokhin, A. V.; Elsey, G. M.; Head, N. H.; Kristic,
A. R.; Levin, M. D.; Michl, J.; Munton, J.; Pinkhassik, E.; Robert, M.;
Saveant, J.-M.; Shtarev, A.; Stibor, I. J. Org. Chem. 1999, 64, 2618.
(3) Wiberg, K. B. J. Org. Chem. 2002, 67, 1613.
* To whom correspondence should be addressed. (W.A.) Tel: (+61
8) 8201 2134. Fax: (+61 8) 8201 3035.
† The Flinders University of South Australia.
‡ Karaelmas University.
(4) Adcock, W.; Kristic, A. R. Magn. Reson. Chem. 2000, 38, 115.
(5) Adcock, W.; Peralta, J. E.; Contreras, R. H. Magn. Reson. Chem.
2003, 41, 503.
§ Universita` di Roma “La Sapienza”.
(6) Adcock, W.; Kristic, A. R. Magn. Reson. Chem. 1997, 35, 663.
10.1021/jo040236b CCC: $30.25 © 2005 American Chemical Society
Published on Web 01/06/2005
J. Org. Chem. 2005, 70, 1029-1034
1029

Adcock et al.
procedure of Meyers et al.,20 a solution of the silyloxazoline
(400 mg) in THF (10 mL) was treated with 4.5 M HCl (50 mL)
,and the resulting solution was stirred at room temperature
for ca. 20 h. The reaction mixture was then saturated with
NaCl and extracted thoroughly with CH2Cl2. After the extracts
were dried (MgSO4), the solvent was removed in vacuo to afford
the crude product which, after sublimation (80 °C/0.2 mmHg),
gave the title compound 1 (X ) SiMe3, Y ) COOH) as a white
solid (270 mg, 87%): mp 155-156 °C (lit.15 mp 165-166 °C);
1H NMR (CDCl3) δ 10.22 (bs, 1H), 1.97 (s, 6H), -0.05 (s, 9H);
13C NMR (CDCl3, relative to Me4Si) δ 175.2, 50.9, 42.8, 30.5,
-3.7.
3-Nitrobicyclo[1.1.1]pentane-1-carboxylic Acid 1 (X )
NO2, Y ) COOH). Powdered sodium azide (284 mg) was added
over 3 × 1 intervals in 1 h21 to a stirred solution of
3-phenylbicyclo[1.1.1]pentane-1-carboxylic acid 1 (X ) C6H5,
Y ) COOH; 410 mg, 2.18 mmol)17 in CHCl3 (5 mL) containing
concentrated H2SO4 (0.76 mL) maintained at 35-40 °C. The
reaction mixture was kept at this temperature for a further 1
h before being poured into water. The aqueous phase was
extracted with CHCl3 before being basified with a 20% NaOH
solution and extracted thoroughly with diethyl ether. After the
ether extract was dried (MgSO4), the solvent was removed in
vacuo to afford the free amine 1 (X ) NH2, Y ) C6H5) as a
yellow oil (300 mg, 87%): 1H NMR (CDCl3) δ 7.32-7.21 (m,
5H), 2.09 (s, 6H), 1.89 (bs, 2H); 13C NMR (CDCl3, relative to
Me4Si) δ 139.8, 128.0, 126.2, 126.17, 55.3, 48.3, 36.6.
experimental data for several substituents,8 the situation
for the former remains less defined since only one
measured gas-phase acidity result is available (1, X ) H
and Y ) COOH)3 to corroborate the calculations. Herein
we report experimental gas-phase acidities of 1 (Y )
COOH) for several substituents (X ) H, NO2, CN, CF3,
COCH3, COOCH3, F, Cl, Br, C6H5, CH3, C(CH3)3, and Si-
(CH3)3) which permit substantiation of the aforemen-
tioned calculated acidities. In addition, we report disso-
ciation constants (pKa values) of several of the carboxylic
acids (1, Y ) COOH and X ) H, NO2, CN, CF3, COOCH3,
F, and CH3) determined in water by potentiometric
titration. As we shall see below, the analysis of the new
data strengthens the view1-3 that the substituent-induced
acidities of 1 (Y ) COOH) appear to be best described in
terms of an electrostatic field model.
Experimental Section
Synthesis of Compounds. Literature procedures were
followed in the preparation of most of the carboxylic acids 1
(Y ) COOH; X ) H2, CN,15 CF3 , COCH3,15 COOCH3,16 F2, Cl,15
2
Br,15 I,15C6H5,17 CH3,18 and C(CH3)315). The preparation of the
others (1, Y ) COOH; X ) Si(CH3)3 and NO2) are described
below.
3-Trimethylsilylbicyclo[1.1.1]pentane-1-carboxylic Acid
1 (X ) Si(CH3)3, Y ) COOH). A pentane solution (3.23 mL)
of 1.7 M tert-butyllithium (5.51 mmol, 2 mol equiv) was added
dropwise to a stirred solution of 1-bromo-3-(4,4-dimethyl-2-
oxazolinyl)bicyclo[1.1.1]pentane 1 (X ) Br, Y ) Ox; 0.67 g, 2.75
mmol)19 in anhydrous diethyl ether (25 mL) maintained at -70
°C under nitrogen. The resulting solution was allowed to warm
to -50 °C for 15 min before the addition of Me3SiCl (1 mL).
The solution was maintained at this temperature with stirring
for 5 min before being allowed to warm to room temperature.
The reaction mixture was then quenched with water and
extracted with CH2Cl2. The extract was dried over MgSO4
before the solvents and volatiles were removed in vacuo.
Kugelrohr distillation (50 °C/0.5 mmHg) of the residue gave
the silyloxazoline 1 (X ) SiMe3, Y ) Ox) as a colorless oil (400
mg, 61%): 1H NMR (CDCl3) δ 3.57 (s, 2H), 1.90 (s, 6H), 1.26
(s, 6H), -0.043 (s, 9H); 13C NMR (CDCl3, relative to Me4Si) δ
168.4, 78.5, 67.7, 51.1, 42.6, 30.7, 22.7, -3.6. By use of the
A solution of the amine (300 mg) in acetone (5 mL) was
added to a freshly prepared solution of dimethyldioxirane22
in acetone (150 mL), and the mixture was stirred overnight
at room temperature. Removal of the solvent in vacuo afforded
the crude nitrophenyl derivative 1 (X ) NO2, Y ) C6H5) as a
1
solid (220 mg, 64%): mp 74-75 °C (lit.1 mp 81.5-82 °C); H
NMR (CDCl3) δ 7.40-7.24 (m, 5H), 2.64 (s, 6H); 13C NMR
(CDCl3, relative to Me4Si) δ 135.7, 128.5, 127.6, 126.4, 68.5,
55.2, 35.3. The crude nitrophenyl compound 1 (X ) NO2, Y )
C6H5; 200 mg, 1.05 mmol) was oxidatively cleaved to the nitro
acid 1 (X ) NO2, Y ) COOH) by use of ruthenium tetraoxide23
as recently described for the preparation of the fluoro acid 1
(X ) F; Y ) COOH) from the corresponding fluorophenyl
compound 1 (X ) F, Y ) C6H5).2 The reddish colored oxidation
product was sublimed (80 °C/0.5 mmHg) to afford the desired
acid 1 (X ) NO2, Y ) COOH) as a light brown powder. A
further sublimation followed by recrystallization from benzene
gave the nitro acid as white colored rods (100 mg, 55%): mp
1
156-157 °C (lit.1 mp 165-165.5 °C); H NMR (CDCl3) δ 7.87
(7) (a) Roberts, J. D.; Moreland, W. T., Jr. J. Am. Chem. Soc. 1953,
75, 2167. (b) Holtz, H. D.; Stock, L. M. J. Am. Chem. Soc. 1964, 86,
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(8) Koppel, I. A.; Mishima, M.; Stock, L. M.; Taft, R. W.; Topsom,
R. D. J. Phys. Org, Chem. 1993, 6, 685.
(9) Exner, O.; Bo¨hm, S. Chem. Eur. J. 2002, 8, 5147 and references
therein.
(10) Adcock, W.; Anvia, R.; Butt, G.; Cook, A.; Duggan, P.; Grob, C.
A.; Marriott, S.; Rowe, J.; Taegepera, M.; Taft, R. W.; Topsom, R. D.
J. Phys. Org. Chem. 1991, 4, 353.
(11) Adcock, W.; Trout, N. A. Chem. Rev. 1999, 99, 1415 and
references therein.
(12) (a) Adcock, W.; Kok, G. B.; Abeywickrema, A. N.; Kitching, W.;
Drew, G. M.; Olszowy, H. A.; Schott, I. J. Am. Chem. Soc. 1983, 105,
290. (b) Adcock, W.; Gangodawila, H.; Kok, G. B.; Iyer, V. S.; Kitching,
W.; Drew, G. M.; Young, D. Organometallics 1987, 6, 156.
(13) Adcock, W.; Butt, G.; Kok, G. B.; Marriott, S.; Topsom, R. D. J.
Org. Chem. 1986, 50, 2551.
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Reson. Chem. 1986, 24, 213.
(15) (a) Della, E. W.; Taylor, D. K. J. Org. Chem. 1994, 59, 2986.
(b) Taylor, D. K. Ph.D. Dissertation, Flinders University of South
Australia, Adelaide, 1993.
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114, 601.
(18) Lochert, I. D. Ph.D. Dissertation, Flinders University of South
Australia, Adelaide, 1996.
(19) Adcock, W.; Binmore, G. T.; Krstic, A. R. Walton, J. C.; Wilkie,
J. J. Am. Chem. Soc. 1995, 117, 2758.
(bs, 1H), 2.68 (s, 6H); 13C NMR (CDCl3, relative to Me4Si) δ
173.2, 64.2, 55.0, 31.4.
Measurement of Gas-Phase Acidities. (i) Methodology.
The gas-phase acidity (GPA) of an acid AH is defined as the
Gibbs free energy change (∆G) of the AH f A- + H+ reaction
usually defined at 298 K. The gas-phase acidities of the [1.1.1]
acids (AH; 1, Y ) COOH) were measured by means of the
kinetic method introduced by Cooks and co-workers.24 This
method is based upon the generation of cluster ions [A‚H‚B]-,
where BH is a suitable reference compound, in an electrospray
(20) (a) Meyers, I. A.; Pansegrau, P. D.; Ricker, W. F. J. Am. Chem.
Soc. 1988, 110, 7178. (b) Meyers, I. A. Gabel, R.; Mihelich, E. D. J.
Org. Chem. 1978, 43, 1372.
(21) Vogel, A. I. Practical Organic Chemistry, 4th ed.; Longman:
London, 1978.
(22) (a) Murray, R. W.; Jeyaraman, R.; Mohan, L. Tetrahedron Lett.
1986, 27, 2335. (b) Murray, R. W. Rajadhyaksha, S. N. Mohan, L. J.
Org. Chem. 1989, 54, 5783.
(23) Carlsen, P. H. J.; Katsuki, T.; Martin, V. S.; Sharpless, K. B.
J. Org. Chem. 1981, 46, 3936.
(24) (a) Cooks, R. G.; Kruger, T. L. J. Am. Chem. Soc. 1977, 99, 1279.
(b) McLuckey, S. A.; Cameron, D.; Cooks, R. G. J. Am. Chem. Soc. 1981,
103, 1313. (c) Cooks, R. G.; Patrick, J. S.; Kotiaho, T.; McLuckey, S. A.
Mass Spectrom. Rev. 1994, 13, 287. (d) Cooks, R. G.; Koskinen, J. T.;
Thomas, P. D. J. Mass Spectrom. 1999, 34, 85. (e) Armentrout, P. B.
J. Am. Soc. Mass Spectrom. 2000, 11, 371.
1030 J. Org. Chem., Vol. 70, No. 3, 2005

Polar Substituent Effects in the Bicyclo[1.1.1]pentane Ring System
(ESI) source of a FT-ICR spectrometer.25 The mass-selected
[A‚H‚B]- ion is activated by collision with a suitable gas (e.g.,
N2), and the fragments ([A]- and [B]-) arising from the
collision-induced fragmentation (CID; eg 1) are mass analyzed.
each [1.1.1]-acid. In addition, the pKa value of benzoic acid was
similarly measured for comparison with literature values. The
new value (4.11 ( 0.01) is about 0.10 units less than the best
literature values.27
k1
9
k2
[A]- 7
[A‚H‚B]- 98[B]-
(1)
Computational Methods
All calculations reported below were carried out at the
B3LYP/6-31+G* level of theory utilizing the GAUSSIAN 94
and 98 program packages.28 Frequency calculations were
performed on most of the density functional theory (DFT)
optimized geometries to determine zero-point vibrational
energies (ZPVE) and, as well, to ensure that there were no
imaginary frequencies at the stationary point.
The standard form of the kinetic method allows quantitative
determination of the gas-phase acidity of AH provided that
certain requirements such as the absence of reverse activation
barriers for the two dissociation channels of eq 1 are met.
Within this framework the ∆GPA difference (GPA(BH) - GPA-
(AH)) is calculated from the relative abundances of the
individual deprotonated fragments ([A]-/[B]-) through eq 2.
[A]-
k
Results and Discussion
∆GPA
RTeff
ln 1 ) ln
)
(2)
[B]-
k2
Gas-Phase Measurements. The gas-phase acidities
(∆G300; kcal mol-1) of the [1.1.1]-acids (1, Y ) COOH) are
displayed (in red) in Figure 1 together with the values
for the various reference acids (in black). The relevant
∆∆G300 values are given in blue. It should be noted that
no value is reported for the iodo acid (1, Y ) COOH; X )
I) because the corresponding carboxylate ion proved to
be fragile under the conditions of measurement (see
above). A noteworthy feature is that the methodology of
measurement does not provide a reliable determination
of the ∆∆S values associated with the competitive colli-
sion-induced fragmentation of the adduct (eq 1) between
the [1.1.1]-carboxylates (1, Y ) COO-) and the reference
carboxylate ions. However, since both competitors are
carboxylic acids, it is assumed that the ∆∆S values are
close to zero. It follows that the ∆∆G values coincide with
the corresponding ∆∆H parameters. Anchoring the mea-
sured ∆∆G values to the reference gas-phase acidities
produces the absolute acidity scale for 1 (Y ) COOH)
shown in red in Figure 1. The relevant absolute ∆H scale
(Table 1) is obtained from this by adding to the ∆G300
values the T∆S term calculated for a free proton with
the Sackur-Tetrode equation (7.75 kcal/mol). It can be
seen (Table 1) that there is fairly good agreement
between the observed and available calculated ∆H values
given the respective uncertainties in the numbers. This
is exemplified by the very good linear correlation between
them (Table 2, entry 1). Thus, the calculated acidities
reported by Wiberg3 at the MP2/6-311++G** theoretical
level adequately account for the experimental absolute
Accordingly, the slope of the linear plot of ln([A]-/[B]-)
versus the GPA of a series of reference acids (BH) provides
an estimate of the effective temperature term (Teff ) 300 K),
while the y-intercept gives the GPA(AH) value.The constraints
for the measurement of GPA values by the standard kinetic
method are likely met in the case of the [1.1.1] acids (AH; 1,
Y ) COOH) since the loss of a proton from carboxylic acids
hardly involves reverse activation barriers.
(ii) Procedure. The CID experiments were performed by
using a commercial APEX 47e FT-ICR mass spectrometer
equipped with an ESI source and a syringe pump. Operating
conditions for the ESI source were as follows: spray voltage,
3.8 kV; capillary temperature, 298 K; sheath gas (N2) flow rate,
30 units (roughly 0.75 L/min). The selected gas-phase com-
plexes were generated by electrospraying at a flow rate of 10
mL/min an acetonitrile solution containing equimolar amounts
(10 µM each) of the [1.1.1] acid (AH) and the reference acid
(BH). The FT-ICR-CID experiments were carried out in the
negative ion mode. The so-formed [A‚H‚B]- complexes were
injected from the external source of the instrument into the
analyzer, quenched by multiple collision with N2, admitted into
the analyzer via a pulsed valve assembly (the nominal peak
pressure was ca. 5 × 10-6 mbar), and eventually isolated by
broad-band ejection of the unwanted fragments. Sustained off-
resonance irradiation CID experiments were carried out by
accelerating the quenched ions with an activation frequency
0.30 kHz higher than the cyclotron frequency of the ion and
by allowing them to collide for 1.5 s with N2 present in the
analyzer at the constant pressure of 2 × 10-8 mbar.
pKa Measurements. Potentiometric titrations were carried
out in a water-jacketed vessel maintained at 298.2 K under a
stream of argon. Data were obtained from 10 mL aliquots of
solution containing 5.0 × 10-3 M HCl, 10-1 M KCl, and 1.0 ×
10-3 M [1.1.1]-acid titrated with standardized 0.1 M KOH. The
latter carbonate-free solution was titrated against oven-dried
potassium hydrogen phthalate which was used as the primary
standard. An autoburet equipped with a 5 mL buret was used
to deliver the titrant, and the potential was measured by a
Sure Flow 8165 BN combination electrode connected to a pH
meter. The autoburet and pH meter were interfaced to an IBM-
compatible personal computer which controlled the addition
of titrant using a program AUTOTIT (written by Dr. A. P.
Arnold and Dr. P. A. Duckworth) so that successive addition
of titrant caused a decrease of ∼4 mV in potential. The
electrode was calibrated by a titration in the absence of organic
acid and fitting the resulting data from this strong acid-base
titration to the Nernst equation to find the correct values for
E° and pKw. The pKa values were determined using the
program SUPERQUAD.26 Three titrations were performed for
(27) Shorter, J. Pure Appl. Chem. 1994, 66, 2451.
(28) (a) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Gill, P. M. W.;
Johnson, B. G.; Robb, M. A.; Cheeseman, J. R.; Keith, T.; Petersson,
G. A.; Montgomery, J. A.; Raghavachari, K.; Al-Laham, M. A.;
Zakrzewski, V. G.; Ortiz, J. V.; Foresman, J. B.; Cioslowski, J.;
Stefanov, B. B.; Nanayakkara, A.; Challacombe, M.; Peng, C. Y.; Ayala,
P. Y.; Chen, W.; Wong, M. W.; Andres, J. L.; Replogle, E. S.; Gomperts,
R.; Martin, R. L.; Fox, D. J.; Binkley, J. S.; Defrees, D. J.; Baker, J.;
Stewart, J. P.; Head-Gordon, M.; Gonzalez, C.; Pople, J. A. Gaussian
94, Revision D.3; Gaussian: Pittsburgh, PA, 1995. (b) Frisch, M. J.;
Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman,
J. R.; Zakrzewski, V. G.; Montgomery, J. A., Jr.; Stratmann, R. E.;
Burant, J. C.; Dapprich, S.; Millan, J. M.; Daniels, A. D.; Kudin, K.
N.; Strain, M. C.; Farkas, O.; Tomasi, J.; Barone, V.; Cossi, M.; Cammi,
R.; Mennucci, B.; Pomelli, C.; Adamo, C.; Clifford, S.; Ochterski, J.;
Petersson, G. A.; Ayala, P. Y.; Cui, Q.; Morokuma, K.; Rega, N.;
Salvador, P.; Dannenberg, J. J.; Malick, D. K.; Rabuck, A. D.;
Raghavachari, K.; Foresman, J. B.; Cioslowski, J.; Ortiz, J. V.; Baboul,
A. G.; Stefanov, B. B.; Liu, G.; Liashenko, A.; Piskorz, P.; Komaromi,
I.; Gomperts, R.; Martin, R. L.; Fox, D. J.; Keith, T.; Al-Laham, M. A.;
Peng, C. Y.; Nanayakkara, A.; Challacombe, M.; Gill, P. M. W.;
Johnson, B.; Chen, W.; Wong, M. W.; Andres, J. L.; Gonzales, C.; Head-
Gordon, M.; Replogle, E. S.; Pople, J. A. Gaussian 98, Revision A.11.3;
Gaussian: Pittsburgh, PA, 2002.
(25) Fenn, J. B.; Mann, N.; Meng, C. K.; Wong, S. F. Mass Spectrom.
Rev. 1990, 9, 37.
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J. Org. Chem, Vol. 70, No. 3, 2005 1031

Adcock et al.
FIGURE 1. Interlocking ladder scale (kcal/mol) of gas-phase acidities of [1.1.1]-acids (1, Y ) COOH) (in red) and reference acids
(in black). ∆∆G300 values are given in blue.
TABLE 1. Enthalpies (∆H, kcal/mol) of Dissociation of
[1.1.1]-acids (1, Y ) COOH)
Note also (Table 2, entry 3) that the observed ∆H
values correlate very well against polar field constants
(σF) determined from the 19F substituent chemical shifts
(SCS) in c-C6H12 of 1-substituted (X) 4-(p-fluorophenyl)-
bicyclo[2.2.2]octanes (2, Y ) p-FC6H4).11 The fact that
these σF values in nonpolar solvents mirror the gas-phase
environment is clearly underlined by a correlation of very
high precision (r ) 0.997, n ) 12) between them and
recently published available energies (∆E4, kcal/mol)
calculated at the B3LYP/6-311+G** level of theory of the
isodesmic reaction (eq 4) for system 2 (Y ) COOH). This
result is particularly significant since the solution acidity
data of 2 (Y ) COOH)
a
b
X
∆Hexp
∆Hcalcd
CH3
345.15
343.5
C(CH3)3
Si(CH3)3
H
343.55
343.55
343.45 (342.9)b
342.55
343.0
C6H5
COOCH3
COCH3
F
340.25
339.85
338.85
336.9
336.1
334.8
Cl
337.35
CF3
336.95
Br
336.35
CN
333.85
332.6
330.9
NO2
332.65
X-[2.2.2]-COOH + H-[2.2.2]-COO- f
H-[2.2.2]-COOH + X-[2.2.2]-COO- (4)
a Calculated from the ∆G300 values (Figure 1) using the T∆S
contribution of 7.75 kcal/mol (calculated for a free proton by the
Sackur-Tetrode equation). b Taken from ref 3.
together with those of 4-substituted (X) quinuclidinium
ions, underpin the empirical parameters (σI ≡ σF) which
are believed to characterize polar field effects of substit-
uents.30 Thus, an obvious corollary from the very good
correlation (Table 2, entry 3) between ∆H of 1 (Y )
COOH) and σF is that the gas-phase acidities of 1 (Y )
COOH) and 2 (Y ) COOH) should parallel one another.
This is clearly the case as revealed by the good correlation
between them (Table 2, entry 4). The regression coef-
ficient (1.295), which indicates an enhancement of polar
substituent effects in 1 (Y ) COOH), is in good agreement
with the value (1.33) obtained by Wiberg3 from a plot
between the respective calculated ∆H parameters. The
values. However, it is worth noting that a much lower
level of theory (B3LYP/6-31+G*) is equally adequate to
describe the relative substituent effects on the acidities.
Indeed, as shown in entry 2 of Table 2, an excellent
correlation exists between the measured ∆H values
(Table 1) and the B3LYP/6-31+G*-calculated energy
changes (∆E3, kcal/mol) for the isodesmic reaction (eq 3;
Table 3).29
X-[1.1.1]-COOH + H-[1.1.1]-COO- f
H-[1.1.1]-COOH + X-[1.1.1]-COO- (3)
(29) The isodesmic energies listed in Table 3 are uncorrected (0.98
ZVPE corrections effect only minor perturbations).
(30) Hansch, C.; Leo, A.; Taft, R. W. Chem. Rev. 1991, 91, 165 and
references therein.
1032 J. Org. Chem., Vol. 70, No. 3, 2005

Polar Substituent Effects in the Bicyclo[1.1.1]pentane Ring System
TABLE 2. Regression Statistics of Correlative Analysesa
entry
yb
xc
md
ce
rf
Fg
ph
ni
1
2
3.
4.
5.
6
∆Hobs[1.1.1]
∆Hobs[1.1.1]
∆Hobs[1.1.1]
∆Gobs[1.1.1]
∆Ecalcd[1.1.1][1.1.1]
∆pKa
∆Hcalcd[1.1.1]
∆Ecalcd[1.1.1]
0.950
0.857
-16.977
1.295
0.495
0.081
-1.623
18.322
344.197
344.319
-108.480
0.095
-23.316
4.023
0.994
0.992
0.983
0.974
0.996
0.945
0.974
398.896
586.726
315.759
109.528
1129.848
41.739
<0.0001
<0.0001
<0.0001
<0.0001
<0.0001
0.0013
7
12
13
8
10
7
k
σF
∆Gobs[2.2.2]j
∆Ecalcd[1.1.1]
∆Gobs[1.1.1]
l
7.
∆pKa
σF
91.690
0.0002
7
a Simple linear regression, y ) mx + c. b Dependent, variable. c Independent variable. d Regression coefficient. e Intercept. f Correlation
coefficient. g F-test of variance for overall correlation. h Probability value. i Number of data points in correlation. j Taken from ref 8.
k c-C6H12, taken from ref 11. l CH3OH, taken from ref 11.
TABLE 3. Energies (∆E, kcal/mol) of Isodesmic
TABLE 4. Acidities of [1.1.1]-Acids (1, Y ) COOH) in
Reactions (eqs 3 and 5, Respectively)a
H2O at 298.2 K
a
X
∆E3(eq 3)
∆E5(eq 5)
X
pKa
CH3
4.12 ( 0.03
3.95 ( 0.00
3.51 ( 0.10
3.29 ( 0.05
3.25 ( 0.05
3.21 ( 0.01
3.08 ( 0.03
CH3
0.57
-0.04
-0.79
0.00
-0.34
-2.29
-2.70
-4.63
-5.67
-6.55
-8.92
-8.53
-11.09
-13.16
0.32
H
C(CH3)3
Si(CH3)3
H
NH2
C6H5
OH
COOCH3
COCH3
F
Cl
CF3
COOCH3
F
CF3
CN
NO2
0.00
-0.12
-1.18
-1.96
a Determined by potentiometric titration. Given as the mean
of 3 × 1 titrations.
-3.09
-3.94
-4.50
-5.48
-6.47
three reasons: (i) the distance between the two systems
being compared
CN
NO2
X-[1.1.1]-[1.1.1]-COOH + H-[1.1.1]-[1.1.1]-COO- f
a See the Supporting Information for B3LYP/6-31+G*-calcu-
lated energies and zero-point energies (ZPE) of the various
structures.
H-[1.1.1]-[1.1.1]-COOH + X-[1.1.1]-[1.1.1]-COO-
(5)
origin of the enhancement was ascribed to an increase
in the field effect as the distance decreases and, moreover,
is in agreement with the ratio (1.32) based on a distance
dependency of 1/r2 from simple electrostatic theory (Kirk-
wood-Westheimer (K-W) equation).31 However, it is
highly probable that the latter accord is fortuitous given
that one of the key approximations in the derivation of
the K-W equation, namely, that the length of the
substituent dipole is very small compared to the distance
between the substituent and reaction center, is question-
able on the molecular scale. Recently, a modified field
effect model has been proposed with a “falloff” with
distance considerably less than that expected from simple
electrostatics.32 Interestingly, a comparison of the gas-
phase acidities of 4-substituted (X) quinuclidinium and
bicyclo[2.2.2]octylammonium ions appears to support this
proposition.10 In view of the uncertainty in the actual
dependence of the field effect on distance we decided to
perform a stringent test by comparing the energetics of
reaction 3 with those of reaction 5 in which the corre-
sponding bis-bicyclopentyl compounds ([2]staffanes)33 are
deployed. These rod-shaped cage-structure compounds
appear to be ideal molecular spacers for this purpose for
is now markedly different; (ii) no problem is created by
an angle orientation factor; (iii) the B3LYP/6-31+G* level
of theory, which was shown above to be adequate for
describing relative substituent effects on acidities, can
be economically and conveniently applied to these rela-
tively large systems.
It can be seen (Table 2, entry 5) that the linearity
between ∆E5 and ∆E3 (listed in Table 3) is excellent with
a regression coefficient of 0.495. This measure of the
“falloff” with distance is clearly not in agreement with
either the square of the ratio of the distances between
the bridgehead hydrogens and the carboxylate oxygens
in the respective systems ((5.184/8.507)2 ) 0.371) or
simply the ratio of the distances (5.184/8.507 ) 0.609).
In fact, it lies somewhere between 1/r and 1/r2 and, hence,
gives credence to the proposal that the field effect model
requires modification.32
pKa Measurements. The pKa values for some of the
[1.1.1]-acids (1, Y ) COOH) were determined in water
and are given in Table 4. Values for the chloro and bromo
acids could not be determined due to their rapid solvoly-
ses. In several cases (X ) C(CH3)3, Si(CH3)3, C6H5, and
COCH3), solubility was an intractable problem. The
correlation between these pKa values (Table 4) and the
corresponding gas-phase acidities (Figure 1) is not of high
precision (Table 2, entry 6), but a strong linear trend is
indicated. Hence, the solution acidity trends can clearly
be predominantly attributed to intrinsic structural fac-
tors. The significantly improved fit of the correlation
between the pKa values and the σF values for methanol
(31) (a) Westheimer, F. H.; Kirkwood, J. G. Trans. Faraday Soc.
1947, 43, 77 and references therein. (b) Bowden, K.; Grubbs, E. J.
Chem. Soc. Rev. 1996, 25, 171 and references therein.
(32) (a) Charton, M. J. Phys. Org. Chem. 1999, 12, 275. (b) Charton,
M.; Charton, B. I. J. Chem. Soc., Perkin Trans. 2 1999, 2203.
(33) Kaszynski, P.; McMurdie, N. D.; Michl, J. J. Am. Chem. Soc.
1991, 56, 307.
J. Org. Chem, Vol. 70, No. 3, 2005 1033

Adcock et al.
(Table 2, entry 7) attests to specific substituent-solvent
interactions having a minor bearing on the trends.
Y ) F and Sn(CH3)3) reveals the importance of electron
delocalization effects rather than polar field effects.4-6
Conclusions
Acknowledgment. We are grateful to the Austral-
ian Research Council for partial financial support of this
work and, as well, to the Italian Ministero dell’Istruzione,
della Universita` e della Ricerca (MIUR) and the Con-
siglio Nazionale delle Ricerche (CNR) of Italy. We also
thank Associate Professor K. P. Wainwright for his
assistance in the measurement of the solution acidities
of the [1.1.1] acids.
The experimental gas-phase acidity data of the [1.1.1]-
acids (1, Y ) COOH) presented in this paper corroborate
calculated values determined at a high level of theory
(MP2/6-311++G**//B3LYP/6-311+G**) recently reported
by Wiberg.3 The relative substituent effects on the
acidities, however, are accurately described by a much
lower level of theory (B3LYP/6-31+G*). The origin of
these effects can be ascribed exclusively to an electro-
static effect of the substituent dipole acting directly
through space. Cross-cage interactions due to the prox-
imity of the bridgehead positions (ca. 1.85 Å) appear to
be negligible. This picture corroborates an earlier deduc-
tion of Applequist et al.1 that some solution acidities of
such acids could be described satisfactorily in terms of
an electrostatic field model. Further, it is in agreement
with recent evidence presented by Wiberg3 for the opera-
tion of polar field effects rather than inductive transmis-
sion through bonds. By contrast, it is worth noting that
the use of charge density probes in the [1.1.1]-system (1,
Supporting Information Available: Electronic energies
(E) and zero-point vibrational energies (ZPVE) obtained from
DFT calculations (B3LYP/6-31+G*) for 3-substituted (X) bicyclo-
[1.1.1]pentane-1-carboxylic acids, 3′-substituted (X) 1′,3-
bibicyclo[1.1.1]pentane-1-carboxylic acids, and their corre-
sponding anions (Tables S1 and S2); Cartesian coordinates of
bicyclo[1.1.1]pentane-1-carboxylic acids and their correspond-
ing anions; 1H and 13C NMR spectra of 3-nitro- and 3-trimethyl-
silylbicyclo[1.1.1]pentane-1-carboxylic acids 1 (Y ) COOH; X
) NO2 and Si(CH3)3, respectively). This material is available
free of charge via the Internet at http://pubs.acs.org.
JO040236B
1034 J. Org. Chem., Vol. 70, No. 3, 2005