Equilibrium acidities of superacids

Article abstract of DOI:10.1021/jo101409p

In this paper, we report the most comprehensive equilibrium superacidity scale that is available to date. Contrary to most of the past works, this scale is set up in a medium of constant composition and the obtained acidity values characterize the acidities of molecules rather than acidities of media. The current scale is thus complementary to the well-known H₀ scale in the information that it provides. The solvent used is 1,2-dichloroethane (DCE). DCE has very weak basic properties (but sufficiently high polarity) and is an appropriate solvent for measuring acidities of very strong acids of diverse chemical nature. DCE acidities of well-known superacids (CF₃SO ₂OH, (CF₃SO₂)₂NH, cyanocarbon acids, etc.) as well as common mineral acids (H₂SO₄, HI, HBr, etc.) are reported. Acidities of altogether 62 acids have been determined from 176 interlinked relative acidity measurements. The scale spans 15 orders of magnitude (from picric acid to 1,1,2,3,3-pentacyanopropene) and is expected to be a useful tool in design, use, and further acidity measurements of superacidic molecules.

Full text of DOI:10.1021/jo101409p

pubs.acs.org/joc
Equilibrium Acidities of Superacids
€
€
Agnes Kutt,* Toomas Rodima, Jaan Saame, Elin Raamat, Vahur Maemets,
Ivari Kaljurand, Ilmar A. Koppel,* Romute Yu. Garlyauskayte, Yurii L. Yagupolskii,
Lev M. Yagupolskii, Eduard Bernhardt, Helge Willner, and Ivo Leito*
Institute of Chemistry, University of Tartu, Ravila 14a, Tartu, 50411, Estonia, Institute of Organic
Chemistry, National Academy of Sciences of Ukraine, Murmanskaya 5, Kiev, 02094, Ukraine, and
€
Anorganische Chemie, Bergische Universitat Wuppertal, Gaussstrasse 20, D-42097 Wuppertal, Germany
agnes.kutt@ut.ee; ilmar.koppel@ut.ee; ivo.leito@ut.ee
Received July 25, 2010
In this paper, we report the most comprehensive equilibrium superacidity scale that is available to
date. Contrary to most of the past works, this scale is set up in a medium of constant composition and
the obtained acidity values characterize the acidities of molecules rather than acidities of media. The
current scale is thus complementary to the well-known H0 scale in the information that it provides.
The solvent used is 1,2-dichloroethane (DCE). DCE has very weak basic properties (but sufficiently
high polarity) and is an appropriate solvent for measuring acidities of very strong acids of diverse
chemical nature. DCE acidities of well-known superacids (CF3SO2OH, (CF3SO2)2NH, cyanocarbon
acids, etc.) as well as common mineral acids (H2SO4, HI, HBr, etc.) are reported. Acidities of altogether 62
acids have been determined from 176 interlinked relative acidity measurements. The scale spans 15 orders
of magnitude (from picric acid to 1,1,2,3,3-pentacyanopropene) and is expected to be a useful tool in
design, use, and further acidity measurements of superacidic molecules.
Introduction
stabilizing of highly reactive species, as well as charge carriers
in electrochemical power sources.6-10
Acids of very high strength;superacids1;and their an-
ions (often behaving as weakly coordinating anions;
WCAs) have found a wide area of applications.2-5 WCAs
are of key importance as counterions in catalysis and the
Cyanocarbon acids,11-13 sulfonic acids,3 sulfonimides,5,14
and other types of strong acids have been known and used for
a long time, but equilibrium acidity data in solution on most
of these acids range from scarce to at best inconsistent15
and often rely on indirect methods.16 Although (besides their
*To whom correspondence should be addressed at the University of
Tartu.
(1) The term “superacid“ refers in this paper to a superacidic compound
not to a superacidic medium. A superacidic compound can be defined as a
compound having in a particular medium, e.g., gas phase or some solvent,
higher acidity than H2SO4.
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Chemistry; Wiley-Interscience: New York, 2009.
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Schnepp, Z. Green Chem. 2007, 9, 30–37.
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2243. (b) Posternak, A.G.;Garlyauskayte, R. Y.;Polovinko,V.V.;Yagupolskii,
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DOI: 10.1021/jo101409p
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Published on Web 12/17/2010
J. Org. Chem. 2011, 76, 391–395 391
2010 American Chemical Society

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Kutt et al.
JOCArticle
-
gas-phase acidities17,18) the pKa values of some of them have
been published in acetic acid,19 most of the acidity data have
been obtained in solvent mixtures of varying composition,
such as aqueous H2SO4 at different concentrations2 or aceto-
nitrile acidified to different extents,15 i.e. not in constant-
composition liquid media. The reported infrared spectro-
scopic scale of the ion-pairing ability of superacid anions
with trioctylammonium counterions;the νNH scale16;is
useful for characterizing anions of superacids, but is not an
equilibrium acidity scale and characterizes the hydrogen
bond acceptor strength of anions rather than the Brønsted
acidity of their parent acids.
and protonated base (HBþA1 and HBþA2-) during the
relative direct acidity measurements:
ΔKip
HA1 þ HBþA2- a HA2 þ HBþA1-
ð3Þ
ð4Þ
aðHBþA1- Þ aðHA2Þ
3
3
ΔpKip ¼ log
aðHBþA2- Þ aðHA1Þ
Assuming that the ratio of activity coefficients of the ion-
paired and neutral species of both acids are similar at any
acidity of the solution;f(HBþA1-)/f(HA1) = f(HBþA2-)/
f(HA2);concentrations of the species can be used instead of
their activities. The counterion HBþ of the acid anions is the
protonated phosphazene base t-BuP1(pyrr) (in a few cases
Alk4Nþ). It is important to emphasize that the directly
measured acidities are so-called ion-pair acidities. Never-
theless, their differences are expected to reflect the differences
of the free ion acidities of the acids: ΔpKa ≈ ΔpKip
The choice of the medium for measurements of the acid
strength of strong acids is not simple. Typical solvents are
either not convenient for acid-base measurements because
of low polarity, poor solvating ability of ions, and polar solutes
(e.g., heptane,20 SO2) or are too basic (e.g., DMSO,21 MeCN22)
for studying strong acids.
In this work, the equilibrium acidities of a large range of
strong acids, many of them superacids, have been measured
in 1,2-dichloroethane (DCE) and we present the most com-
prehensive equilibrium superacidity scale that is available to
date in a medium of constant composition.
pKaðHA2Þ - pKaðHA1Þ
½HBþA1- ꢁ ½HA2ꢁ
3
3
¼ ΔpKa ꢀ ΔpKip ¼ log
ð5Þ
½HBþA2- ꢁ ½HA1ꢁ
The necessary condition for this is that the ion-pairing
affects both anions in a similar way and, in particular, there
must be no hydrogen bonding involved. Some of the smaller
anions (such as Cl- and BF4-) may be more strongly influenced
by ion pairing and this may lead to a shift of their pKa values in
the scale relative to the acids with larger anions. Never-
theless, this is not expected to influence the span of the scale
itself because its backbone is built by using acids that on
deprotonation give anions with efficiently delocalized charge.
Such acids and their anions are also not significantly influ-
enced by traces of water in the solvent.25 Since correcting for
this ion-pairing would be somewhat artificial and introduce
additional assumptions (see ref 26) we decided not to correct
and leave the data as they are.
The formation of hydrogen bonding between the acid
anions and the protonated base is hindered by the choice
of the base;t-BuP1(pyrr). The proton in protonated t-BuP1-
(pyrr) is buried between the bulky substituents and the charge
of the bulky cation is highly delocalized. There is strong
evidence against hydrogen bonding between the acid anions
and the protonated phosphazene base:
(1) Low hydrogen bond donicity of protonated phospha-
zenes was among the goals for which phosphazene bases
were initially developed by the Schwesinger group.27 The
proton in the protonated forms of the bulky phosphazene
bases is buried between substituents. This is especially true if
the substituent on the imino nitrogen is bulky (t-Bu in our case)
and if the substituents on the phosphorus are large, such as
pyrrolidinyl (our case), tetramethyl guanidinyl, etc. The
hindrance of the proton in protonated phosphazenes has
DCE is an appropriate solvent for studying superacids
because it combines negligible basic properties and inertness
with the ability to dissolve many polar and ionic compounds.
Although its relative permittivity (εr =10.60) is among the
highest known for chloroalkanes, it is still too low to efficiently
support the dissociation of ion pairs into free ions.23 For this
reason and because of its low ion-solvating ability, relative
ion-pair acidities (ΔpKip) are directly measured. The low
polarity also causes several side processes in acidity measure-
ments: ion-pair formation and homoconjugation are more
extensive in DCE than, for example, in MeCN.
The pKa value is used to describe the equilibrium acidity of
an acid HA in a solvent S:
Ka
-
HA þ S a A þ SHþ
ð1Þ
ð2Þ
-
aðSHþÞ aðA Þ
3
pKa ¼ - log
aðHAÞ
Using the approach22,24 based on relative acidity measure-
ments of acids HA1 and HA2 (see the SI for a description of
the measurement method) there is no need to determine the
activity of the solvated proton in the medium, a(HSþ), or to
take into account the formation of ion pairs of the acid anion
(17) Koppel, A.; Taft, R. W.; Anvia, F.; Zhu, S.-Z.; Hu, L.-Q.; Sung,
K.-S.; DesMarteau, D. D.; Yagupolskii, L. M.; Yagupolskii, Y. L.; Ignatev,
N. V.; Kondratenko, N. V.; Volkonskii, A. Y.; Vlasov, V. M.; Notario, R.;
Maria, P.-C. J. Am. Chem. Soc. 1994, 116, 3047–3057.
€
(18) Leito, I.; Raamat, E.; Kutt, A.; Saame, J.; Kipper, K.; Koppel, I. A.;
Koppel, I.; Zhang, M.; Mishima, M.; Yagupolskii, L. M.; Garlyauskayte,
R. Y.; Filatov, A. A. J. Phys. Chem. A 2009, 113, 8421–8424.
(19) (a) Engelbrecht, A.; Rode, B. M. Monatsh. Chem. 1972, 103, 1315–
1319. (b) Foropoulos, J.; Desmarteau, D. D. Inorg. Chem. 1984, 23, 3720–
3723.
(25) Kaupmees, K.; Kaljurand, I.; Leito, I. J. Phys. Chem. A 2010, 114,
11788–11793.
~~
€
(20) Room, E.-I.; Kaljurand, I.; Leito, I.; Rodima, T.; Koppel, I. A.;
(26) Kaljurand, I.; Rodima, T.; Pihl, A.; Maemets, V.; Leito, I.; Koppel,
I. A.; Mishima, M. J. Org. Chem. 2003, 68, 9988–9993.
Vlasov, V. M. J. Org. Chem. 2003, 68, 7795–7799.
(21) Bordwell, G. Acc. Chem. Res. 1988, 21, 456–463.
(27) Schwesinger, R.; Schlemper, H.; Hasenfratz, C.; Willaredt, J.;
Dambacher, T.; Breuer, T.; Ottaway, C.; Fletschinger, M.; Boele, J.; Fritz,
H.; Putzas, D.; Rotter, H. W.; Bordwell, F. G.; Satish, A. V.; Ji, G.-Z.; Peters,
E.-M.; Peters, K.; Schnering, H. G. V.; Walz, L. Liebigs Ann. 1996, 1055.
(28) Kolomeitsev, A. A.; Koppel, I. A.; Rodima, T.; Barten, J.; Lork, E.;
€
€
(22) Kutt, A.; Leito, I.; Kaljurand, I.; Soovali, L.; Vlasov, V. M.;
Yagupolskii, L. M.; Koppel, I. A. J. Org. Chem. 2006, 71, 2829–2838.
(23) Reichardt, C. Solvents and Solvent Effects in Organic Chemistry, 3rd
ed.; Wiley-VCH: Weinheim, Germany, 2003.
€
Roschenthaler, G.-V.; Kaljurand, I.; Kutt, A.; Koppel, I.; Maemets, V.;
Leito, I. J. Am. Chem. Soc. 2005, 127, 17656–17666.
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(24) Leito, I.; Rodima, T.; Koppel, I. A.; Schwesinger, R.; Vlasov, V. M.
J. Org. Chem. 1997, 62, 8479–8483.
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been demonstrated by X-ray diffractograms.27,28 A particularly
noteworthy feature of phosphazenium cations in the context
of this paper is their ready applicability in the preparation of
the so-called “naked fluorides”.29 Cations behaving as inert
toward the highly HB-acceptory fluoride anion can be safely
considered the same with respect to all anions of this study.
(2) The absence of hydrogen bond between the protonated
phosphazene base and acid anions is also seen from the UV-vis
spectra that are recorded during the measurement. When
phosphazene base t-BuP1(pyrr) is used as titrant no distor-
tion of the spectra of acid-anion mixtures is observed. At the
same time, titration of some of the acids (1, 13, 17, 21, 24, 28,
and 29) with triethylamine (or being originally triethylam-
monium or pyridinium salts) caused distortion of the spec-
tral plot (a hypsochromic shift of the maximum of spectrum)
and strongly inconsistent results. This is a clear indication of
formation of a hydrogen-bonded complex. The same effect
was never experienced with t-BuP1(pyrr). For double-checking,
some of the compounds were also titrated with an even
bulkier base t-BuP4(dma) and the spectra matched with the
spectra obtained from the titration with base t-BuP1(pyrr).
The t-BuP1(pyrr) base was preferred because it is a sufficiently
strong base30 and is commercially available with reasonable
purity and price.
well-defined protonation centers in the anion.7 It is fair to
say that designing a superacid depends largely on the design
of the corresponding anion.32
From Table 1 two different design concepts of superacid
anions emerge as the most potent in enhancing acidity:33
cyanocarbon anions34 (denoted with green color) and anions
of poly-SO2X-imides and methanes (blue color). Both of these
anion groups obtain their stability from the excellent charge
delocalization but with different mechanisms. Cyanocarbon
anions are a beautiful example of delocalization of the negative
charge by efficient polar resonance in a planar (or nearly
planar) C3 (45, 49, 52, 54, 62, etc.) or C5 (19, 61) moiety.35,36
In the sulfo-anions the efficient combination of a polarizable
sulfur atom with the high electron-acceptor power of the dO
atoms (or the even much higher electron-acceptor power of
the dN-Tf37 fragments) is engaged. Both of these concepts
lead to the formation of anions that have a number of
protonation centers with similar (and very low) basicity.
Of the acids not belonging to either of these groups the
strongest is HB(CN)4 58. Although composed according to a
different principle, its anion displays charge delocalization
similar to anions of the cyanocarbon acids by forming 4 equally
weakly basic protonation centers. Anion B(CN)(CF3)3- 31 is a
good example of the loss of this symmetry: the CF3 groups
with their weaker electron-acceptor power leave the single -CN
significantly more basic than in B(CN)4- and as a result 31
is by 4 orders of magnitude weaker acid than 58. HB(CN)4
58 is also significantly stronger than HBF4 33: the small size
Thus, we might have relatively loosely bound ion pairs
between a sterically hindered huge cation (with diameter
˚
over 9 A) and different anions of (mostly) CH acids with well
delocalized negative charge, but not hydrogen-bonded com-
plexes.
-
of the BF4 anion creates high partial charges on the -F
substituents so that the very low intrinsic basicity of the -F
center increases considerably.
Results
The results of acidity measurements are presented in Table 1.
pKa values of 62 acids are interconnected with 176 relative
acidity measurements to obtain a self-consistent acidity scale
ranging in total for 15 pKip units. The exact positions of
individual acids on the scale were found as in previous
works22,24 by minimizing the sum of squares of differences
between the measured ΔpKip values and the differences of the
assigned pKa values. The agreement between the overlapping
measurements (i.e., the ΔpKip of A vs B plus ΔpKip of B vs C
must be equal to ΔpKip of A vs C) is very good as indicated by
the consistency standard deviation of the scale 0.032 units.22,24
Our method allows for the measurement of relative acid-
ities expressed as ΔpKip values. To obtain absolute pKa values,
the scale has to be anchored to an acid with a known pKa
value. There are currently no reliable absolute experimental
pKa values available in DCE (see discussion in the SI31). We thus
keep the scale relative at this time, and anchor the scale to the pKa
value of picric acid 1, which is arbitrarily assigned to 0.0 units.
Except for HClO4 55, all classical strong mineral acids are
in the upper half of the scale. Their relative weakness com-
pared to the acids discussed above is especially noticeable in
the case of the hydrogen halides: HCl 2(pKa ca. -8inwater21) is
just slightly stronger than picric acid 1 (pKa = 0.3 in water22),
HI 22 (pKa ca. -10 in water38) is more than 5 orders of
magnitude weaker than HClO4 (pKa ca. -10 in water38). The
reason is obvious: most of the classical mineral acids have
small anions and their strength in water comes largely from
the efficient solvation of the anions.
An interesting case is HBF4. This acid actually is not
-
supposed to exist. Nevertheless titration of Bu4Nþ BF4 with
acid 64 and afterward with base t-BuP1(pyrr) demonstrated the
reversibility of the protonation-deprotonation process. The
formed complex HF BF3 seems to be stable in solution
3 3 3
because of its very low concentration and the extremely low
Lewis basicity of DCE. The acid also behaves very consistently:
three measurements were carried out with it (see Table 1)
and their results are consistent with the measurements not
Discussion
(32) (a) Lipping, L.; Leito, I.; Koppel, I.; Koppel, I. A. J. Phys. Chem. A
2009, 113, 12972–12978. (b) Koppel, I. A.; Burk, P.; Koppel, I.; Leito, I.;
Sonoda, T.; Mishima, M. J. Am. Chem. Soc. 2000, 122, 5114–5124.
(33) It is expected that the carborane-based superacids are still much
stronger.8,16,32 Our attempts to measure the equilibrium acidity of the
unsubstituted HCB11H12 failed. See the SI and ref 32a.
The set of 62 acids directly measured to the present scale
(plus acids 63 and 64 with estimated acidities) includes all
common strong mineral acids and a number of superacids
from different families. The critical factor for the acidity of
a superacid is the stability of its anion and the lack of
ꢀ
(34) Vianello, R.; Maksic, Z. B. New J. Chem. 2009, 33, 739–784. and
references cited therein.
(35) Vianello, R.; Maksic, Z. B. New J. Chem. 2008, 32, 413–427.
ꢀ
(36) Vianello, R.; Liebman, J. F.; Maksic, Z. B. Chem.;Eur. J. 2004, 10,
ꢀ
(29) Schwesinger, R.; Link, R.; Thiele, G.; Rotter, H.; Honert, D.; Limbach,
€
H.-H.; Mannle, F. Angew. Chem., Int. Ed. Engl. 1991, 30, 1372–1375.
5751–5760.
€
~~
(30) Kaljurand, I.; Koppel, I. A.; Kutt, A.; Room, E.-I.; Rodima, T.;
Koppel, I.; Mishima, M.; Leito, I. J. Phys. Chem. A 2007, 111, 1245–1250.
(31) Bos, M.; Dahmen, E. A. M. F. Anal. Chim. Acta 1973, 63, 185–196.
(37) Koppel, I. A.; Burk, P.; Koppel, I.; Leito, I. J. Am. Chem. Soc. 2002,
124, 5594–5600.
(38) Brownstein, S.; Stillman, A. E. J. Phys. Chem. 1959, 63, 2061–2062.
J. Org. Chem. Vol. 76, No. 2, 2011 393

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TABLE 1. Equilibrium Acidity Scale Valid in 1,2-Dichloroethane As Solvent (Acidity Increases Downwards)h
aDirectly measured relative acidity values in DCE. bPredicted pKa values of MeCN (see the SI for details). cpKa value of picric acid is arbitrarily set
to 0. dTos represents the 4-Me-C6H4SO2- group. eTf represents the CF3SO2- group. fX-TCNP represents 2-X-1,1,3,3-tetracyanopropene. gEstimated
DCE pKa values, see text. hScheme 1 shows compound and group structures.
394 J. Org. Chem. Vol. 76, No. 2, 2011

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SCHEME 1. Abbreviations for Compounds and Groups Used
values measured in this work with the respective MeCN pKa
values shows quite good correlation. The most deviating points
in the first case out of 21 points are acids whose anions have
rather localized negative charge (HBr 12, HI 22) but also
some cyanocarbon acids (19 and 27) and one sulfonic acid
(29). In general, it is possible to predict MeCN pKa values
knowing DCE pKa values at least within (1.5 pKa units. This
opens the possibility to estimate the MeCN pKa values of
very strong acids that cannot be measured experimentally in
MeCN. From the correlation of the gas-phase acidities with
DCE pKa values it seems that acid anions with small size and
localized charge (2, 5, 12, 22, 41) behave differently from
anions with large size and delocalized charge indicating
stabilization of the anion by solvent molecules in the solution
phase. The correlations indicate that DCE has ca. 10% better
differentiating ability (on a logarithmic scale) than MeCN,
but ca. 1.5-2 times lower differentiating ability than the gas
phase (see the SI, pages S4 and S5).
involving HBF4. During this experiment, it is impossible that
the acidity of pure HF was measured by mistake (based on
the moles of titrants consumed). HF is also such a weak
acid (considerably weaker than HCl) that it could not be
included into the present scale. We cannot, however, rule out
that the acidic species involved was actually a hydrate, such
as H2O HBF4. The reported pKa value should therefore be
3
treated with caution.
CF3SO(dNTf)NHTf 64;the acidic titrant used in this
work to protonate the anions of the strongest acids;is a very
strong neutral acid. The Yagupolskii substitution,37 i.e.
replacement of dO by dNTf in aromatic or aliphatic
sulfonimides, leads to an acidity increase by 6.4 (6 and 28)
to 5.3 (23 and 57) units. By introducing the Yagupolskii sub-
stituent to Tf2NH 47, compound 64 is obtained. Taking into
account the average acidity increase (5.9 units), the pKa value
of 64 can be estimated as -18 units.
The best-known way to describe acidity of a superacidic
medium is the H0 value of Hammett acidity function. The H0
value has been determined for most liquid superacidic media,
also for those that have no constant composition, such as
mixtures of Brønsted and Lewis acids and mixtures of different
Brønsted acids (HSO3F and SbF5; CF3SO3H and SbF5; HF
and HSO3F; etc.). The H0 values of the liquid media depend
strongly on their actual composition. For example, while HF
as a molecule has low acidity, the acidity of liquid HF in
terms of H0 is very high: -15.1.2 This is caused by formation
of the highly stable FHF- ion on dissociation of HF.2
Different from the H0 function, the acidity scale of this
work characterizes acidity of molecules (similar to the intrinsic
GA values in the gas phase) rather than media, and therefore
offers information that is complementary to the H0 values.
The latter for the different neat superacidic media are deter-
mined, in addition to their inherent acidity, also by the variable
nonspecific and specific solute-solvent interaction charac-
teristics23 of the media. Because of this, no good correlation
between the pKa values in DCE and H0 values for individual
acids could be expected (R2=0.37, S=1.30, with four acids:
7, 34, 41, and 55).
Experimental Section
The UV-vis spectrophotometric titration method used in
this work is essentially the same method as described earlier.22,24
It is based on the UV-vis spectrophotometric titration of a
solution of two acids with the solution of a nonabsorbing base
t-BuP1(pyrr)39 to obtain the relative acidity of the two com-
pounds. The reversibility of the protonation-deprotonation pro-
cess was verified with all compounds. Compounds that did not
behave reversibly (HPF6, HAl[OCH(CF3)2]4) were not included
in the scale. For further details see the SI.
Acknowledgment. This work was supported by the Grant
Nos. 7374 and 8162 from the Estonian Science Foundation
and by the target financing projects SF0180061s08 and
SF0180089s08 from the Ministry of Education and Science
of Estonia. We are indebted to Prof. Takaaki Sonoda and
JEMCO Inc. for compounds 47, 50, 51, 53, and 56. We are
indebted to Prof. O. W. Webster for compounds 19 and 61.
Supporting Information Available: Choosing an anchor
compound for the acidity scale, comparison and correlations
of DCE pKa values with IR frequencies (νNH), MeCN pKa values,
and GA values, homoconjugation, estimate of DCE pKa values
of compounds 63 and 64, description of failed measurements of
some compounds, description of the ΔpKa measurement proce-
dure, MeCN pKa values of hydrogen halides and a revision of
the MeCN pKa values, description of the origin, purification,
and synthesis of used chemicals and solvents, NMR spectra of
newly synthesized compounds, and UV-vis spectra of all com-
pounds on the scale that have useful UV-vis spectra. This material
is available free of charge via the Internet at http://pubs.acs.org.
Correlation of the DCE pKa values with the gas-phase
acidities is also poor (see the SI for details of the correlations).
At the same time the comparison of some of the DCE pKa
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(39) Kaljurand, I.; Kutt, A.; Soovali, L.; Rodima, T.; Maemets, V.; Leito,
I.; Koppel, I. A. J. Org. Chem. 2005, 70, 1019–1028.
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