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Structure and Acidity of Organic Compounds.

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ANGEWANDTE CHEMIE
V O L U M E 3 * N U M B E R 10
OCTOBER 1964
P A G E S 661-712
Structure and Acidity of Organic Compounds [l]
BY PR1V.-DOZ. DR. W. SIMON
LABORATORIUM FUR ORGANISCHE CHEMIE
DER EIDGENOSSISCHEN TECHNISCHEN HOCHSCHULE, ZURICH (SWITZERLAND)
Elucidation of the structures of organic acids and bases with the aid of their acidity constants
is greatly hindered by their low solubilities in water and by the small quantities of material
normally available. Apparent acidity constants K* can be determined potentiometrically
on small samples (about 0.5 rng) for many classes of compounds of medium acidity. The
apparent crcidity constants K&,, determined in the solvent system methyl cellosolve/
water can be correlated, in a similar manner to thermodynamic quantities, with structural
parameters (e.g. Hammett’s LT and TaJi’sD* values). K&-, values can yield information
on the axial or equatorial position of carboxyl groups in cyclohexanecarboxylic acids.
I. Introduction
Some excellent reviews have been written----on the
reIationships between structure and acidity of-a wide
range of compounds 12-51,
In this connection acids and bases are defined according
to Brmsted [6] and Lowry [7] as proton donors and
proton acceptors, respectively, so that acid-base equilibria in an amphiprotic [8] solvent LH of high dielectric
constant may be formulated as follows:
-
- -
AH + LH =-A@ + LH2@
.cLH~~.~LH~~
- c A-Q . ~ A B ~K ’ = aA@’aLH2@
(2)
~
~~
~ A H ’ ~ L H
CAH
AH . C L H . ~ L H
(a = activity; c = concentration; f = activity coefficient)
obtained by applying the law of mass action to Equilibrium (l), the following relationship is valid:
A F = -2.303 RT log K’ = 2.303 RT pK’
(1)
In Equation (l), AH is the acid to be investigated and
A Q the base formed from AH by the removal of a
proton. No assumptions will be made regarding the
nature of AH and A Q ; it should merely be noted that
A9 contains one hydrogen atom less, and carries one
[I] Lecture delivered a t a meeting of the Chemical Society of
Bern (Switzerland) on January 18th, 1963.
[2] J. F. King in K . W. Bentley: Elucidation of Structures by
Physicat and Chemical Methods. Part I, Interscience Publishers,
New York-London 1963, pp. 317-403.
[3] E. M. Arneft in S. G . Cohen, A . Streitwieser, Jr., and R . W.
Tuft: Progress in Physical Organic Chemistry. Interscience Publishers, New York-London 1963, Vol. I, pp. 223-403.
[4] H. C.Brown, D. H . McDaniel, and 0. Hufliger in E. A . Braude
and F. C.Nuchod: Determination of Organic Structures by Physical Methods. Academic Press, New York 1955, pp. 567-662.
151 R. P. Bell: The Proton in Chemistry. Methuen & Co., London 1959.
[6] J. N . Brmsted, Recueil Trav. shim. Pays-Bas 42, 718 (1923).
[7] 7’. M. Lowry, Chem. and Ind. 42, 43 (1923).
181 L. F. Audrieth and J. Kleinberg: Non-Aqueous Solvents.
Wiley New York and Chapman & Hall, London 1953.
Angew. Chern. internat. Edit. 1 Vol. 3 (1964) / No. 10
unit charge more, than AH. The pairs AH/Aa-and
LHzQ/LH are referred to as conjugate acid-base
pairs. For the equilibrium constant
(3)
According to Equation (3), the pK’ value of the acid
AH is directly proportional to the change in free
energy A F of the acid-base reaction of Equation (1):
AF
=
AH - TAS
(4)
For a change in AF of 1.O kcal/mole at 25 OC,Equation
(3) gives a change in pK’ of 0.73 units. In principle,
both AF (or pK’) and AH can be correlated with the
structure of the compound studied. However, opinions
are divided as to which of the two quantities is the more
suitable one [3, 9-14].
191 L . P. Hurnrnetf: Physical Organic Chemistry. McGraw-Hill
Book Co., New York 1940.
[lo] R. W,Tafr, Jr. in M . S. Newmun: Steric Effects in Organic
Chemistry. Wiley, New York 1956.
[ l l ] H . C . Longuet-Higgins and C . A . Coulson, J. chem. SOC.
(London) 1949, 971.
[I21 H. S. Frank and M. W. Evans, J. chem. Physics 13, 507
(1945).
1131 J. E. Leffler, J. org. Chemistry 20, 1202 (1955).
1141 W . B. Person, J. Amer. chem. SOC.84, 536 (1962).
66 1
Quantitative investigations on acid-base equilibria are
generally carried out on systems in which relatively
small quantities of AH are allowed to react with a large
quantity of solvent LH. The activity aLH of LH at
constant temperature and pressure may therefore be
regarded as constant, and the dissociation constant is
obtained from Equation (2) :
Instead of t h e p K value ( p K = -log K ) of a n acid-base pair,
the so-called basicity constant ( ~ K B is
) sometimes used. I n
contrast t o t h e p K value, this constant applies t o a n equilibrium of t h e type:
B + L H Z B H @ + L3
(6)
To avoid confusion, this terminology will be avoided here a s
far as possible. Consequently, in the remainder of this
article t h e basicity of t h e base B will be defined by t h e acidity
of the conjugate acid BHa. Thus, for t h e acid BH" there is
a n acid-base equilibrium similar t o Equation (1):
BH@+ LH
B + LH2@
(7)
11. Determination of Acidity Constants
Routine determinations of dissociation constants are
generally carried out by potentiometric or spectrophotometric methods [15]. In the potentiometric method the
activity of the hydrogen ions is measured potentiometrically (e.g. with a glass electrode/calomel electrode
combination [16]) for a given ratio cA'/cA~. Assuming
that f A B / f A H = 1, the constant K can then be calculated
from Equation (5). Errors due to neglect of the activity
coefficients are about 0.05 pK units [17] for 1:1 electrolytes in aqueous systems with sample concentrations of
0.01 M. The ratio cA@/cAH
can be adjusted by mixing
a given amount of the acid with varying amounts
of its salt. At the half-neutralization point in the potentiometric titration [18-221 cAG = cAH and c ~ ~
M 1, so that at this point the pH value of the solution
is equal to the pK of AH [cf.
(pH = -log aLHz@)
Equation (5)].
In spectrophotometric determinations of pK values
[15,23 -261, the sample is generally dissolved in a
[I51 A . AIbert and E. P. Serjeant: Ionization Constants of Acids
and Bases, A Laboratory Manual. Methuen & Co., London
and Wiley, New York 1962.
[I61 R. G. Bates: Electrometric p H Determinations. Interscience
Publishers, New York-London 1954.
[I71 P. Debye and E. Hiickel, Physik 2. 24, 185 (1923).
[I81 M. T .Kelley, D. J. Fisher, and E. B. Wagner, Analytic. Chem.
32, 61 (1960).
[I91J . B. Neilandsand M. D.Cannon,Analytic. Chem. 27,29(1955).
[20] W . Simon, E. Kovats, L. H. Chopard-dit-Jean, and E. Heilbronner, Helv. chim. Acta 37, 1872 (1954).
[21] J.-P. Phillips: Automatic Titrators. Academic Press, New
York 1959.
[22] T. V. Parke and W. W. Davis, Analytic. Chem. 26, 642 (1954).
1231 0. Redlich, E. K. Holt, and J . Bigeleisen, J. Amer. chem. SOC.
66, 1 3 (1944).
[24] G. C. Hood, 0. Redlich, and C. A . Reilly, J. chem. Physics 22,
2067 (1954).
[25] G. C. Hood, 0. Redlich, and C. A . Reilly, J. chem. Physics 23,
2229 (1955).
[26] G. C. Hoodand C. A. Reilly, J. chem. Physics 27, 1126 (1957);
28, 329 (1958).
662
system of constant hydrogen ion activity (buffer
solution), and the ratio CA&AH is obtained by
spectrophotometry. If the pH of the solution is known,
K can be calculated from Equation ( 5 ) , ignoring activity coefficients. Constants which have been obtained
applying this approximation are often called apparent
dissociation constants.
In the systematic determination of dissociation constants for qualitative analysis of organic natural
products, two main restrictions must be considered :
a) It should be possible to carry out the measurements
on small quantities of sample (0.3-1.5 mg).
b) Owing to the low solubility of most organic compounds in water, the determinations must be carried
out in non-aqueous solvents.
Under these conditions, determinations of thermodynamic dissociation constants K must be abandoned
in favor of those of apparent acidity or dissociation
constants K*. Potentiometric microtitration [20,22,27]
is the most suitable method for rapid routine determinations.
Solvents used so far include: alcohol a n d alcohol/water [28 to
321, ethylene glycol mixed with other organic solvents [33-35],
acetonitrile [36- 381, dimethylformamide [39 -411, cyclohexanone [42], dioxane,'water [43-451, benzene [46], nitrobenzene [36,48], acetone [47,49], pyridine [50,51], glacial acetic
[27] W. Ingold, Helv. chim. Acta 29, 1929 (1946); Mikrochem.
verein. Mikrochim. Acta 36, 276 (1951).
[28] J . I. Bodin, Dissertation Abstr. 19, 235 (1958).
[29] G. Kortiinz and N . Buck, Z. Elektrochem., Ber. Bunsenges.
physikal. Chem. 62, 1083 (1958).
[30] L. Michaelis and M. Mizutani, Z. physik. Chem. 116, 135
(1925).
[31] S . B. Knight, R . H . Wallick, and C . Balch, J . Amer. chem.
SOC.77, 2577 (1955).
[32] L. A. Wooten and L. P . Hanzmett, J. Amer. chem. SOC.57,
2289 (1935).
1331 C. H. Kalidas and M. N . Das, J. Indian chem. SOC.36, 231
(1959).
/ [34]c M. ~N . Das
~ and S . R. Palit, J. Indian chem. SOC.31, 34 (1954).
[35] S . R . Palit and U. N . Singh, J. Indian chem. SOC. 33, 507
(1956).
[36] H. K . Hall, Jr., J. physic. Chem. 60, 63 (1956).
[37] J. F. Coetzee and I. M. Kolthofi, J. Amer. chem. SOC.79, 61 10
(1 957).
[38] E. J. Forrnan and D. N. Hume, J. physic. Chem. 63, 1949
(1959).
[39] P. H. Given, M. E. Peover, and J. Schoen, J. chern. SOC.(London) 1958, 2674.
[40] 1.A. Dean and C. Cain, Jr., Analytic. Chem. 27, 212 (1955).
1411 S. M. Kaye, Analytic. Chern. 27, 292 (1955).
[42] V. I. Dulova and J. N . Kim, Khim. Nauka i Prom. 4, 134
(1959).
[43] J. C. James and J . G. Knox, Trans. Faraday SOC.46, 254
(1950).
1441 H. P . Marshall and E. Grunwald, J. Amer. chem. SOC.76,
2000 (1954).
1451 S . Kertes, J. chem. SOC.(London) 1955, 1386.
[46] M. M. Davis and H. B. Hetrer, J. Res. nat. Bur. Standards
60, 569 (1958).
[47] H. V. Malmstadt and D. A. Vassallo, Analytic. Chem. 31,
206 (1959).
[48] S. Aronoff, J. physic. Chem. 62, 428 (1958).
[49] J . S. Fritz and S . S . Yamamura, Analytic. Chem. 29, 1079
(1957).
[SO] C. A. Streuli and R. R. Miron, Analytic. Chem. 30, 1978
(1958).
[51] A. Anastasi and E. Mecarelli, Mikrochem. Acta 1956, 252.
Angew. Chem. internat. Edit.
Vol. 3 (1964)
No.10
acid [52,53], and methyl cellosolve [27,54]. In these solvents
the results of potentiometric determinations of acidity constants are reproducible only if all operations of the determination procedure are highly standardized [55,56].
A standard method has been developed in our laboratory for the determination of apparent dissociation
constants in the system methyl cellosolve/water (80:20 %
by weight), using samples of 0.1 mg and more [*I. The
compound to be investigated is titrated according to a
predetermined program using a glass-electrode/calomelelectrode combination which has been conditioned in
water and calibrated with aqueous buffer solutions. The
apparent pH of the solution [**I at half-neutralization
of the functional group to be determined is assumed to
be equal to pKhCs, i.e. equal to the negative decadic
logarithm of the apparent acidity constant KGCs(MCS:
methyl cellosolve, ethylene glycol monomethyl ether
Figure 1 shows the glass-electrode/calomel-elecrrodemeasuring cell with a n interchangeable electrode membrane (cuptype electrode) [20,60,62], which serves as titration vessel,
and a piston burette with a digital readout; the titrant is
displaced from the burette by means of a tantalum piston
(normal rate of titration: 1.4 ylimin). A control unit starts
and stops the titration and controls the flow of nitrogen for
agitation of the solution. The signal of the p H meter [Type
7666, Leeds & Northrup Co., Philadelphia, Pa. (U.S.A.)] is
registered by a strip chart recorder [Speedomax Type G ,
Leeds & Northrup Co.].
CH30CH2CH20H).
Details of the standard method can be obtained from earlier
publications [20,55-601. The standard deviations found for
a single determination were 0.07 [ 5 8 ] and 0.04 p K b c s units
[SS] for semi-automatic [20] and fully automatic [56) titrations, respectively. With the semi-automatic titration equipment, 1 ml of a 3x 10-3 to 4 x 10-3 M solution is normally
titrated at a constant speed, and the neutralization curve is
plotted simultaneously (20, 611. The p K G c s values of well
over a thousand organic acids and bases have been published
so far [63-651.
~~
~
.-
1521 J . A . Riddick, Analytic. Chem. 32, 172 R (1960).
[53] S. Bruckenstein and I . M . Kolthof, J. Amer. chem. SOC.78,
2974 (1956).
[54] A . E . Ruehle, Ind. Engng. Chem., analyt. Edit. 10, 130 (1938).
[55] W. Simon and E, Heilbronner, Chimia 10, 256 (1956).
[56] W.Simon and E. Heilbronner, Helv. chim. Acta 40,210(1957).
[*I Methyl cellosolve is a good solvent for many compounds.
Admixture of20 yoby weight of water increases the stability of the
potential of the glass electrode, without greatly reducing its
universal solvent power.
[**I Measured in accordance with the definition of pH values
on thc so-called p r a c t i c a l pH scale [16,57]. According to this
scale, the pH values of aqueous solutions are determined on the
basis of pH standards. By analogy, the apparent pH of the
sample (PHGCS) is defined with reference to a standard buffer
solution as follows:
where
pHs= 4.005 (pH of 0.05 M potassium hydrogen phthalate at
25 "C);
E = electromotive force [abs. volts] obtained at 2 5 "C with
the electrode system in the solution being investigated;
E, = electromotive force [abs. volts] obtained at 25 "C with the
electrode system in 0.05 M potassium hydrogen phthalate in
water.
Es and E are given the sign of the conductor system of the glass
electrode.
The denominator 0.05916 corresponds to the theoretical electrode
function; it is adjusted to match the electrode function determined in aqueous systems.
[57] K . Schwabe: Fortschritte der pH-MeBtechnik. Verlag Technik, Berlin 1953.
[58] W. Simon, Helv. chim. Acta 41, 1835 (1958).
[59] W. Simon, Helv. chim. Acta 39, 883 (1956).
[60] W. Simon, Chimia 10, 286 (1956).
[61] R. Dohner and W. Simon, unpublished work.
[62] P. M . T.Kerridge, Biochem. J. 19, 61 1 (1925); J. sci. Instruments 3, 404 (1926).
[63] W. Simon, G . H . Lyssy, A . Morikofer, and E. Heilbronner: Zusammenstellungvon scheinbaren Dissoziationskonstanten im Losungsmittelsystem Methylcellosolve/Wasser. Juris-Verlag, Zurich
1959, Vol. I.
[64] P. F. Sommer and W. Simon: 1631, Vol. 11.
[65] W. Simon and P. F. Sommer: [63], Vol. 111.
Angew. Chem. internat. Edit.
1 Vol. 3 (1964) { No.
10
Fig. 1. Glass-electrode/calomel-electrodemeasuring cell with piston
burette.
I . Nitrogen inlet for agitation of solution; 2. cover (chromium-plated
brass); 3. burette tip; 4. glass envelope; 5. insert with electrode membrane; 6. glass-electrode buffer solution; 7. Araldit@block; 8,lO. inlet
and outlet for water (temperature regulation); 9. glass-electrode lead;
1 I . Ag/AgCl inner reference electrode; 12. sample; 13. porous sintered
oxide plug; 14. tip of reference electrode (filled with saturated KCI);
15. reference-electrode lead; 16. reference electrode; 17. tantalum
piston; 18. Teflon" stopper; 19. mount (aluminum); 20. adapter;
21. handle for raising burette; 22. digital readout; 23. drive; 24. motor;
25. burette support.
III. Interpretation of Apparent Acidity
Constants
At constant temperature and pressure the apparent
acidity constants K* determined by the method outlined
above are functions of the structure of the compound
under investigation, as well as of the solvent, concentration, and determination procedure. These constants
K * no longer have the well-defined thermodynamic
significance of an equilibrium constant in the sense
of Equations (2)-(5). Standardization (solvent, concentration, determination procedure, temperature, and
pressure are kept constant) leads however to pK&,
values which depend only on the structures of the
compounds under investigation. Since conclusions are
to be drawn from these apparent acidity constants
KGCs regarding the structure of organic compounds,
the question arises as to how far it is justifiable to
interpret these pKGCs values in a manner analogous
663
to thermodynamic pK values. It has been shown
that such an analogous interpretation is justified [66,67].
As an example, Figure 2 shows the pKhCs values for
m- and p-substituted benzoic acids, phenols, and salicylic acids plotted as a function of the corresponding
Hammett o values [66] (which are independent of the
I
(l)]. With suitable solvents it is therefore possible to
bring compounds of extremely high or low acidities
into an easily measurable acid-base equilibrium with
the solvent. For example, phenols can be studied in
pyridine [70], and aromatic amines in glacial acetic
acid [71].
In the interpretation of the effect of the dielectric constant, two fundamental types of acid-base equilibria
must be distinguished:
‘0.
\
AH
+ LH +----t
LH2@-k A ”
(9)
+A
(lo)
(DC small)
AH%+ LH
I
301
,
,
-10
-95
0
-1
,
05
\
,
13
’5
Fig. 2. p K & c s values for benzoic acids, phenols, and salicylic acids as
a function o f the corresponding Harnmett a values.
Curve I : Salicylic acids: 4 - N H 2 ; 4-OH; 4-OCH3; 4-CH3; 5-OH; 3-OH;
4-CsHs-CONH; 5-CI; 5-61.
Curve 11: Benzoic acids: p-NH2; p - O H ; p-OCH,; rn-NH2; m-OH; H ;
p - F ; m-OCH3; p-Cl; p-Br; p - I ; rn-F; m-I; m - C l ; m-Br; 91r-NO2;p-NOz.
Curve 111: Phenols: p - C I ; p-Br; p - I ; m-CI; m-Br; m-NOz; p - N O 2 .
(The order o f the acids corresponds to the order of the plotted points,
from left t o right).
nature of the reaction, but depend on the type of
substituent) in order to test the validity of Hammett’s
po rule [91
~K&cs=pK;;lcs(O)- P
(8)
In Equation (8), the pKGCs(O) value corresponds to
the pK& value of the unsubstituted compound. It
can be seen from Figure 2 that the apparent acidity
constants of the m- and p-substituted phenols, benzoic
acids, and salicylic acids are affected in the same way
by the substituents on the aromatic nucleus as the
equilibrium and rate constants of other reactions of
side-chain derivatives of benzene. The accuracy of the
linear regression of the pK&cs-values on cr is not
significantly different from that of the corresponding
K values [66].
LH?
According to the Born relationship[72], the energy required to
form an ion is inversely proportional to the dielectric constant of the solvent and to the ionic radius. When a neutral
acid AH dissociates, two ions are formed. This process
requires more energy in organic media, as their dielectric
constants are normally low compared with that of water
( e . g . for methyl cellosolve: water (80: 20 by weight) at
25”C, DC = 32.0 [67]), so that in changing from water
to organic solvents, the extent of dissociation should be
expected to decrease (increase in the pK* values). For
positively charged conjugate acids, the effect of the dielectric
constant should be very small, since for each ion formed,
another simultaneously disappears. However, some individuality should result from the different ionic radii, so
that a simple correlation between the pK values determined
in different media (cf. [32,68,731) is not to be expected. As
a result of the superposition of the effect of the dielectric
constant and that of the protolytic character of the solvent,
the pK* values of acids of the type shown in Equation (9)
generally increase on changing from water to organic solvents, such as alcohols, glycols, and glycol ethers, while those
of the type shown in Equation (10) decrease [30,74] as a
result of the increased basicity of the solvent [ 3 ] .
In Figure 3, the pKGCs values of about 200 functional
groups are compared with the pK values determined in
water (pK,,,)
[75]. Two regression lines can be recognized, one of which (IIj can be assigned to equilibria
of the type shown in Equation (9), the other (I) to that
of Equation (10). As expected, the pKGcS values for
the first type are higher, and those for the second type
lower, than the corresponding pKHZovalues.
If a single regression line [76] is calculated for all the compounds considered in Figure 3, the standard deviation of the
p K k c s values from this line .is found to be 3.15 pKkcs
units (95 % confidence limit). Calculation of the pKGcs
values from P K H ~ O- and vice versa
is thus in general
completely useless. The p K k c s values can be predicted
with much greater accuracy if we restrict ourselves to a single
class of compounds (cf. Figure 3 ) . Thus, for acids with a n equilibrium of the type of Equation (9), the standard deviation of
the ~ K ~ values
C S from the regression line is 1.15 p K k c s units
(95 ,?(confidence limit). The standard deviation for acids of
the equilibrium type of Equation (10) is of the same order
~
IV. Influence of Solvent on the Acidity Constant
The influence of the solvent used on the acidity constant
is essentially due, on the one hand, to its dielectric
constant DC [68], and on the other hand, to its protolytic
character [69]. A basic solvent accepts a proton easier
than an acidic one, so that acidity constants are higher
(lower pK value) in basic solvents, e.g. pyridine, than
in acidic solvents, e.g. glacial acetic acid [cf. Equation
[66] W. Simon, A. Morikofer, and E. Heilbronner, Helv. chim.
Acta 40, 1918 (1957).
[67] W. Simon, Doris Meuche, and E. Heilbronner, Helv. chim.
Acta 39, 290 (1956).
[681 W. F. K. Wynne-Jones, Proc. Roy. SOC.(London) A 140,
440 (1933).
[69] N. Bjerrum and E. Larsson, Z . physik. Chem. 127,358 (1927).
664
[70] C. A. Streuli, Analytic. Chem. 32, 407 (1960).
1711 M. F. Hall and J. B. Conant, J. Amer. chem. SOC.49, 3047,
3062 (1927).
[72] G. Kortiim and J. O’M. Bockris: Textbook of Electrochemistry. Elsevier Publishing Company, New York-Amsterdam-London-Brussels 1951, Vol. I.
[73] N. A. Izmailov,Zhur. Fiz. Khim. 24, 321 (1950).
[74] M. Mizutani, Z. physik. Chem. 116, 350 (1925).
[75] P. F. Sommer, Ph. D. Thesis, Eidgenossische Technische
Hochschule Zurich, 1961.
[76] 0. L. Davies: Statistical Methods in Research and Production. 2nd Edit., Oliver and Boyd, Edinburgh 1954.
Angew. Chem. infernat. Edit. / Vol. 3 (1964)
/ No. I 0
of magnitude. If we consider only classes in which the compounds are structurally closely related, e . g . m- and psubstituted benzoic acids (cf. Figure 4), m- and p-substituted
phenols, benzenesulfanilides, anilines, etc., the p K h c s
values can be predicted with much greater accuracy on the
basis of the p K ~ , ovalues. The p K h c s values for m- and
p-substituted phenols, for example, can be predicted from the
~ K H ~values
o
with a n accuracy of 0.18-0.22 units ( 9 5 %
confidence limit).
A
2 1 - 1
1
L391,3!
2
3
I
I
I
I
'
L
5
6
7
9
-.
9
10
11
12
PG$J-
Fig. 3. Correlation of p K % c s values with pKII,o values [751.
I : Benzoic acid
Benzoic ucids:
2: m-Hydroxy; 3: m-Fluoro; 4: m-Chloro; 5 : m-Bromo; 6 : m-lodo;
7: m-Nitro; 8: m-Amino; 9: m-Methoxy; 10: p-Hydroxy; 1 1 : p Fluoro; 12: p-Chloro; 13: p-Bromo; 14: p-Nitro; 15: p-Amino;
16: p-Methoxy; 17: o-Fluoro; 18: o-Chloro; 19: o-Bromo; 20: oIodo; 21: o-Nitro; 22: o-Amino; 23: o-Methoxy; 24: o-Hydroxy.
25: Formic acid; 26: Acetic acid; 27: Isovaleric acid; 28: Cyclohexanecarboxylic acid; 29: Pivalic acid; 30-31: Oxalic acid pK1-pKz;
32-33: Malonic acid pKl-pK2;
34-35: Succinic acid pKl-pK2;
36-37: Glutaric acid pKI-pK2; 38-39: Adipic acid pK1-pK1; 40-41 :
Pimelic acid pKt-pKz; 42-43: Suberic acid pKl-pK2;
44: Azelaic
acid p K 1 ; 45-46: n-Propylmalonic acid pK1-pKz; 47-48: Isopropylmalonic acid pKl-pK2; 49-50: a,=-Dimethylsuccinic acid pKl-pKn;
51-52: a-Methylglutaric acid pKl-pKz;
53: p-Methylglutaric acid
p K i ; 54-55: a,a-Dimethylglutaric acid pKl-pKz;
56-57: P,P-Dimethylglutaric acid pKl-pK2;
58-59: Citraconic acid pK1-pK2
60-61 : Itaconic acid p K l - p K ~ ; 62-63: Mesaconic acid pK1-pK2
64-65: Maleic acid pKl-pK2;
66-67: Fumaric acid pK1-pK2,
68-69: Phthalic acid pKl-pK2;
70-71 : Tsophthalic acid pK1-pKz;
72-73: Terephthalic acid PKI-PKZ.
Phenol :
74: in-Chloro; 75: m-Bromo; 76: m-Nitro; 77: p-Chloro; 78: p Bromo; 79: p-Iodo; 80: p-Nitro; 81: o-Chloro; 82: o-Nitro.
83: H;
Benzenesulfanilides of the fype X-GjH4-SOz-NH-CaHs:
84: m-Nitro; 85: p-Chloro; 86: p-Nitro; 87: p-Amino; 88: p-Methyl;
89: p-Methoxy.
:
H;
Benzencslrifanilides of rhe type C ~ H S - - S O ~ - N H - C ~ H ~ X90:
91: m-Nitro; 92: p-Chloro; 93: ?-Nitro; 94: ?-Methyl; 95: p-Methoxy; 96: p-Acetoxy; 97: p-Carboxy.
98: Ammonia; 99: Methylamine; 100: Ethylamine; 101: Propylaniine; 102: Butylamine; 103: Heptylamine; 104: Octylamine; 105:
Cyclohexylamine; 106: Isobutylamine; 107: Isoamylamine; 108: Dimethylamine; 109: Diethylamine; 110: Dipropylamine; l l 1 : Dibutylamine; 112: Diamylamine; 113: Diisopropylamine; 114: Diisobutylamine; 115: Trimethylamine; 116: Triethylamine; 117: Tripropylamine; ' 118: Tributylamine; 119: o-Phenylene diamine; 120: mPhenylene diamine; 121: p-Phenylenediamine; 122: Aniline; 123: 1Naphthylamine: 124: 2-Naphthylamine; 125: o-Toluidine; 126: mToluidine: 127: p-Toluidine; 128: p-Anisidine; 129: N,N-Dimethylaniline; 130: N.N-Dimethyl-p-toluidine; 131: N-Methylaniline; 132:
N-Ethylaniline; 133: Piperidine; 134: N-Ethylpiperidine; 135: Piperazine; 136-137:
2,s-Dimethylpiperazine pKl-pK2;
138: Morpholine: 139: N-Methylmorpholine; 140: N-Ethylmorpholine; 141 :
Pyrrolidine;
142: p-Hydroxyethylamin;
143: Bis-p-hydroxyethylamine; 144: Tris-3-hydroxyethylamine; 145: Hydrazine; 146-147:
Ethylene diamine pK1-pK2;
148-149: Trimethylene diamine pK1pK2; 150--151: Tetramethylene diamine pKl-pK2;
152-153: Pentamethylene diamine pK~-pKz.
154: Glycolic acid; 155: Benzilic acid; 156: 3-Indolylacetic acid;
157: 2.6-Dichloro-p-cresol; 158: I-Naphthoic acid; 159: 2-Naphthoic
acid; 160: m-Nitrobenzenesulfonamide; 161 : p-Nitrobenzenesulfonamide; 162: CaH5-SOz-NH-C6H3(0H)-COOH
163-166: Pyroinellitic acid pKI-pK4; 167-168: D-Tartaric acid p K 1 - p K ~ ; 169-170:
meso-Tartaric acid pKl-pK2.
171: Pyridine; 172: 2-Aminopyridine; 173: 4-Aminopyridine; 174:
Lepidine; 175: Quinoline; 176: Acridine; 177: Strychnine; 178:
Neostrychnine; 179: a-Dihydrolysergic acid; 180: y-Dihydrolyserpic
acid; 181 : Allylamine; 182: Diphenylguanidine; 183: Triphenylpuanidine.
Angew. Chem. inter'nat. Edit. 1 Vol. 3 (1964)
1 No. 10
Fig. 4. Correlation of pK+M,-s values with P K H ~ Ovalues for benzoic
acids 1751.
Curve I: m- and p-substituted benzoic acids; p-NOz; m-NOz; m-Br;
m-CI; m - I ; m-F; p - B r ; p-CI; m-OH; m-OCH3; p-F; benzoic acid;
p-OCH3; p-OH; m-NH2; p-NHz.
Curve 11: o-substituted benzoic acids; o-NOz; o-Br; 0-1; 0-CI; o-OH;
0-F; o-OCH3; o-NH2.
Salicylic acid forms a n exception owing t o the pronounced hydrogen
bonding between its o-hydroxyl group a n d its carboxylate group; it was
therefore not considered in the calculation of the regression lines. The
order of the acids corresponds to the order of the plotted points, from
left t o right.
The change in the acidity constant with the composition
of the solvent can be used for the determination of the
nature of the acid functions. Thus, for example, it has
been shown [77] that in dimethylformamide/HzO (2: 1)
protonated l-alkyl-l-azacyclononan-5-ol-6-one
with
&&)
Ho OH
(11
6
121
R = CH3 or CzH5 exists as a bicyclic structure (2) and
that it is the tertiary hydroxyl group which dissociates.
The isopropyl and t-butyl compounds [R = CH(CH3)z
or C(CH3)3], on the other hand, are present in the
monocyclic form ( I ) , and dissociation takes place at
the nitrogen function [77].
V. Acidity of Cyclohexanecarboxylic Acids
Valuable conclusions regarding the structure of cyclohexanecarboxylic acids can be drawn from their pK&
values [78-801, since equatorial and axial carboxyl
-
[ 7 7 ] N . J.Leonnrdand M . o k i , J . Amer. chem. SOC.76,3463 (1954).
[78] P.F. Somwwr, V .P. Arya, and W. Simon, Tetrahedron Letters
20, 18 (1960).
665
Solvent
A A F = AFaxial- AFequat.
[kcal/molel
APK = PKaxiaI-PKequat.
C ~ H I I C O O H C6HllC00-
from A A F
I
Methyl cellosolvelwater
(8O:ZO by weight)
EthanoI/water
(50:50 by volume)
I
1.6 -1- 0.3
1.5
1
2.2 i 0.3
2. I
1
I I*]
-
0.4
0.48
0.50
0.4
Dimetbylformamidef
water (66: 34 by volume)
0.50
0.44
1'1 The values in this column are the differences in the pK values of
cis- and trans-4-t-butylcyclohexane1-carboxylic acids, and were deter-
mined potentiometrically. The values in the other columns were obtained from thermodynamic data in combination with the dissociation
constants of unsubstituted or alkyl-substituted acids.
groups differ appreciably in acidity [81-841. As can be
seen from Table 1, the difference in the free energies of
cyclohexanecarboxylic acids with axial and those with
equatorial carboxyl groups is about 1.6 kcal/mole; the
difference for the corresponding anions is about 2.2
kcal/mole [81,82,84]. The fact that the acidity of the
axial isomer is 0.5 pKGCs units lower than that of the
equatorial isomer can be attributed to hindrance of
solvation of its anion by about 0.6 kcal/mole relative to
the anion of the equatorial isomer [Sl,S2,SS] (cf. also
[86]). This additional hindrance of solvation is essentially due to interactions with axial substituents in the
y-positions with respect to the axial carboxyl group
(1,3-interactions) [cf. Formula (3)J.
-1
rn
[79] P. F. Sommer, C.Pascual, V.P. Arya, andW. Simon, Helv.
chim. Act. 46, 1734 (1963).
[80] C. Pascualand W. Simon, Helv. chim. Acta 47, 683 (1964).
[81] R. D . Stolow, J. Amer. chem. SOC.81, 5806 (1959).
[82] M . Ti&, J. JonriS, and J . Sicher, Collect. czechoslov. chem.
Commun. 24, 3434 (1959).
[83] H . F. J. Dippy, S. R. C. Hughes, and J. W. Laxton, J. chem.
SOC.(London) 1954,4102.
[84] H . Van Bekkum, P . E. Verkade, and B. M . Wepster, Kon.
Akad. Wetensch. Amsterdam, Proc. B 64, 161 (1961).
1851 D . H . R. Burton, Experientia 6, 316 (1950).
1861 G. S. Hammondand D . H. Hogle, J. Amer. chem. SOC.77,
338 (1955).
666
1
2
3
6'-
Fig I. pKt,-s values of acids R-COOH as a function of the polar
substituent constants o f .
1 : Pivalic acid; 2: Diethylacetic acid; 3: Cyclohexanecarboxylic acid;
4: a-Methylbutyric acid; 5 : Cyclohexylacetic acid; 6: Isovaleric acid;
7: y-Phenylbutyric acid; 8 : P-Phenylpropionic acid; 9: Acetic acid;
10: Diphenylacetic acid: [ I : Phenylacetic acid; 12: Glycolic acid;
13: Methoxyacetic acid; 14: Formic acid; 15: Iodoacetic acid; 16:
Mandelic acid; 17: Phenoxyacetic acid; 18: Chloroacetic acid; 19:
Fluoroacetic acid; 20: Cyanoacetic acid; 21 : Dichloroacetic acid.
From Helv. Chim. Acta 46, 1736 (1963).
in Figure 5, in which pKhcs values are plotted as a
function of Taft's polar substituent constants o* [lo]
[cf. Equation (S)]. If a hydrogen atom on the carbon
atom a to the carboxyl group is replaced by a methyl
group, or if ring fusion takes place in the a-position,
c* changes by -0.10 to -0.12 units [lo], corresponding
to a mean increment in pKkcs of 0.22 units [cf.
Figure 51.
Since the pKEfcs value of cyclohexanecarboxylic acid
is 7.44 [87], the pKGcs values of alicyclic monocarboxylic acids can be estimated, under certain assumptions, using the following rule [78-801:
(31
On the assumption that the effects of these axial substituents are a d d i t i v e , the pKGCs increment for one
axial hydrogen atom in y position to an axial carboxyl
group is 0.25 units, relative to the acid without such an
interaction. Within the range of validity of this assumption, it should be possible to estimate the pK>&,,, values
by counting the 1,3-interactions [78-801. The prediction
of the pKhCs values of cyclohexanecarboxylic acids
can be extended by taking into account the effects of
substituents in the a-position relative to the carboxyl
group.
Some indication of the effect of substituents R on the
acidity of carboxylic acids of the type R-COOH is given
0
pKSc,
=
7.44
+ 0.25 a + 0.22 b
(1 1 )
where a: number of 1,3-interactions; and
b=l for an a-CH, or for ring-closure in the a-position;
b=O for a-H.
Since the equilibrium contribution of the epimeric
cyclohexanecarboxylic acid with an axial carboxyl
group is only about 5 %, it may be neglected within the
limits of accuracy of the pK&
determinations
(standard deviation s = 0.04 or 0.07 pKGCs units [SS]).
The rule given in Equation (11) for the evaluation of
pKGCs values has so far been found to be valid
for 60 compounds, some of which were very complex,
e.g. vinhaticoic acid (4) [*I. No exceptions are yet
0-7
&F13
H,C
COOH
COOH
H
PKLcs
calcdlated: 7.91
measured: 8.01
1871 W. Simon, unpublished work.
[*] The interaction of the equatorial hydrogen with the equatorial
carboxyl group is geometrically equivalent to that between the
corresponding axial substituents.
Angew. Chem. internot. Edit. / Val. 3 (1964) / N o . I 0
known. The standard deviation of the calculated values
from the measured ones is 0.11 pKhCs units; the
largest deviation observed was 0.24 pKGCs units.
For example, the rule expressed by Equation (11) was
used for the determination of the structures of two
isomeric decalincarboxylic acids (5a) and (5b) [m.p.
5 1 and 142-143 OC,respectively]prepared by Brown [89].
The problem was that of assigning the cis and trans ring
'
Maximum p K h c range
~
CHQ
The pKGcS ranges for various classes of compounds
are shown in Figure 6 [75]. The limiting pK%cs values
for each class of compounds were in part determined
experimentally ; in some cases the range was calculated
COOH
calc. 8.66
found 8.86
P K + M C ~ calc. 8.16
found 8.30
I
.
I
I
fusion. Titration of 0.7 mg of each compound yielded the
answer within an hour. On the basis of a comparison
of the measured pKGCs values with the 'calculated
ones, the compound which melts at 51 "C must be
assigned the trans-ring fusion.&Definite assignment
would hardly have been possible by other methods, and
would certainly have required considerably more effort
[891 (cf. also C90-971).
-
I
I
I
I
I
I
I
-
7
I
I
I
The maximum usable pK& range is determined by
titration of a strong acid and titration of a strong base
(hydrochloric acid and tetramethylammonium hydroxide) [*I. According to Auevbach and Smolczyk [98], two
inflections are recognizable in the titration curves of
dicarboxylic acids when the difference in p K of the two
carboxyl groups is at least 1.2. Accordingly, an inflec~~
[88] W. Simon, Chimia 16, 312 (1962).
[89] We are indebted to R . F. C . Brown, Australian National
University, Canberra, for samples of this acid.
[90] J. W. ApSimon, 0 . E. Edwards, and R . Howe, Canad. J.
Chem. 40,630 (1962).
[91] B. E. Cross, J. R . Hanson, L. H . Briggs, R . C . Cambie, and P .
S . Rutledge, Proc. chem. SOC.(London) 1963, 17.
[92] V. P . Arya and B. G . Engel, Helv. chem. Acta 44, 1650 (1961).
[93] K . Nakanishi, Yong- Yeng Lin, H. Kakisawa, Hong- Yen Hsii,
and H . C . Hsiu, Tetrahedron Letters 22, 1451 (1963).
[94] C . A. Henrick and P . R . Jeferies, Chem. and Ind. 1963, 1801.
[95] A . I. Scott, S . A. Suthertand, D . W. Young, L. Guglielmetti,
D . Arigoni, and G . A . Sim, Proc. chem. SOC. (London) 1964, 19.
[96] R. C . Cambie, W. R . J . Simpson, and L. D . Colebrook, Tetrahedron 19, 209 (1963).
[97] H. Vorbrueggen and C . Djerassi, J. ..Amer. chem. SOC.84,
2990 (1962).
[*I The p K k s range is governed by the levelling effect of the
solvent. The alkali error of the glass electrode should be negligible
under the experimental conditions.
[98] F..R. Auerbach and-E. Smolczyk, Z . physik. Chem. 110, 65
(1924).
Atigeis.
Chem. internat. Edit.
/
Vol. 3 (1964) / No. 10
3
I
I
I
-
I
I
I
I
I
.
I
I
Range
11.1-3.7
.
.
I
I
I
1 I
I
I
M. p. 142-143 "C
M. p. 51 "C
VI. The pK&,
12.3-2.5
Extreme p K h c s values at which inflections in the
titration curve are still recognizable
C OOH
PKhcs
Table 2. p K k c s range for acid-base titrations in methyl cellosolve/HzO
(80:20 by weight) t751.
COOII
P
113'2
tion should be discernible in the titration curve of an
acid or base, when the pKhcS values are 1.2 units or
more within the pKkcs range. The limiting values
shown in Table 2 agree well with the observed values.
I
4
I
5
I
6
I
7
I
8
I
9
I
I
I
I
I
I
I
1 0 1 1 1 2 1 3 1 4
PK.,sFig. 6. pKh,-s ranges for various classes of compounds [751. Limiting
values according to Table 2 : (---)
and (- - -1.
I and 2: Aliphatic dicarboxylic acids HOOC-(CH2)n-COOH,
pKl
pK1 and
and pK2; 3 and 4: Phthalic acids HOOC-C6H4-COOH,
pK2; 5: Saturated aliphatic monocarboxylic acids H-(CHZ)~-COOH;
6: Aliphatic monocarboxylic acids with a substituent in the a-position
and a conjugation barrier between the substituent and carboxyl group;
7 : Cyclohexanecarboxylic acids without an additional heteroatom;
8: Benzenesulfanilides C ~ H S - S O ~ - N H - C ~ H ~ X
(X in the m- or pposition): 9: Benzenesulfanilides X-C6H4-S02-NH-C6Hs
(X in the
m- or p-position); 10: m- and p-substituted phenols; 11: m- and psubstituted benroic acids: 12: 3-, 4-, or 5-monosubstituted salicylic
acids: 13: Primary aliphatic amines n-alkyl-NH2; 14: Secondary
aliphatic amines (n-alkyl)2NH: 15: Tertiaryaliphatic amines(n-alkyl)AN.
using the Hammett equation [9], e.g. for m- and p substituted benzoic acids, phenols, disubstituted benzoic acids, and benzenesulfanilides, o r the Tngold-Taft
equation [lo], e.g. for a-monosubstituted alkanoic acids.
VII. Acidity Measurements in Acetic Acid
It can be seen from Figure 6 and Table 2 that differentiation between the acidities of very strong and very
weak acids and bases is not possible in the solvent
system methyl cellosolve/water (80: 20 by weight). This
667
requires the use of solvents with other levelling properties [81, i.e. other proton-donor and proton-acceptor
properties.However, potentiometric pK* determinations
according to the standard procedure described are possible only if electrode systems can be found which are
suitable to measure the activity or concentration of the
protonated solvent [LH @z in Equation (l)] in these
solvents. It has been shown that certain glass electrodes
satisfy this requirement for the solvents acetic acid,
acetic acid/0.5 % water [99-1021, and methyl isobutyl
ketone [l02]. Thus, in acetic acid with a water content
of 0-0.5 % by weight, the glass electrode responds
quantitatively over a wide acidity range to the concentration of protonated acetic acid (symbolized by
[ A c O H ~ ~ ]so
) , that the quantity:
can be determined in the same manner as the pH of
aqueous systems.
Acid-base equilibria in solvents of low dielectric constants are not easy to interprete, owing to the presence
of ion-pairs and other associates. For example, in a
solvent such as acetic acid (DC = 6.2 at 25 ‘C [107]), an
acid A H forms an ion-pair :
AH
+ AcOH + AcOHzOAe
value of the half-neutralized sample [99,101]. If the
pressure, temperature, sample and water concentrations,
parameters of the apparatus, and determination procedure, are kept constant or are standardized, pK;,,H
values are obtained which under certain conditions are
essentially a function of the “overall dissociation constant” KAH of the acid AH (or base) under investigation [99,101]:
The strength of an acid AH in solvents of low dielectric
constant is expected to be determined by the ionization
constant KfH [lo51 assigned to the equilibrium of
Equation (13). The constant K,,
is in turn related to
this ionization constant K f H and to the dissociation
constant K t H of the equilibrium (14) by Equation
(15). Since KfH is of the same order of magnitude
for various acids (about 10-7) [104], PK&,H values
are mainly determined by the values of KFH, and can
therefore be correlated with the structure of organic
compounds, analogously to other equilibrium constants,
as shown, for example, in Figure 7. Similar considerations naturally also apply to bases B.
(13)
with the ionization constant
9
AH
Ki
- [AcOH?A’]
I
[AH]
The ion-pair dissociates according to
(14)
with the dissociation constant [53,103-1061:
In order to avoid misunderstandingin the characterization of
acid-base equilibria in non-aqueous media, the expressions
“ionization constant” and “dissociation constant” have been
avoided in favor of the non-commital term “acidity constant”
[cf. Equation ( S ) ] .
By analogy to the standard method for the determination of pKGCs, it is possible to determine PK&H
values which correspond essentially to the pHAcOH
[99] DorothPe Wegmann, Ph.D. Thesis, Eidgenossische Technische Hochschule Zurich, 1962.
[I001 DorothPe Wegmann, J. P. Escarfail, and W . Simon, Helv.
chim. Acta 45, 826 (1962).
[ l o l l Dorothie Wegmann and W . Simon, Helv. chim. Acta 45,
962 (1962).
[I021 Jean-Pierre Escarfuil, Ph. D. Thesis, Eidgenossische Technische Hochschule Zurich, 1963.
[I031 S. Bruckenstein and I. M . Kolthoff, J. Arner. chem. Soc. 79,
5915 (1957).
[lo41 I. M . Kolthoffand S. Bruckenstein, J. Arner. chern. Soc. 79,
1 (1957).
[I051I. M . Kolthoffand S. Bruckenstein, J. Amer. chem. SOC.78,
1 (1956).
[I061 S. Bruckenstein and I . M . Kolthoff, J. Amer. chem. SOC.78,
10 (1956).
[I071 J. W.Tomecko and W . H . Hafcher,Trans. Roy. SOC.
Canada,
Sect. I11 45, 39 (1951).
668
Fig. 7.
(0)
p K L C o values
~
of pyridines (0)and glucose phenylosazones (6)
as a function of the Hammett substituent constant 0.
The lines were obtained by calculating linear regressions, assuming
equal standard deviations for the individual experimental values 1761.
From Helv. chim. Acta 45, 962 (1962).
It is thus now possible, using new standard procedures,
to carry out routine measurements of acid-base equilibria
on small amounts of compounds having very high or low
acidity [99,lOl].
Received. March 19th, 1964
[A 394/181 IE]
German version: Angew. Chem. 76, 112 (1964)
Translated by Express Translation Service, London
Angew. Chenl. internot. Edit. 1 Vol. 3 (1964)
1 No. I 0
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