Effect of alkyl group size on the mechanism of acid hydrolyses of benzaldehyde acetals

Article abstract of DOI:10.1021/jo0202987

Hydrolyses of benzaldehyde acetals, PhCH(OR)2, are specific hydrogen-ion catalyzed when R = methyl, n-butyl, but with secondary and tertiary alkyl derivatives, R = i-propyl, s-butyl, t-butyl, t-amyl, hydrolyses are general-acid catalyzed. The Broonsted α values for both secondary and tertiary alkyl groups are in the range: α = 0.57-0.61. A simple iterative procedure was developed to estimate the individual rate constants for general-acid catalysis by the diacid and monoacid forms of succinic acid buffer. Plots of log k_obs (at [buffer] = 0 M) against pH are linear for the secondary and tertiary acetals, and plots of log k_H for the H³O⁺-catalyzed reaction, ¹³C and ¹H chemical shifts, and ¹J_CH coupling constants against the Charton steric parameter, v, for alkoxy groups are linear. The second-order rate constant, k_H, increases about 100-fold on going from R = Me to R = t-amyl, indicating the significant role of steric effects on reactivity. Steric effects upon ¹³C NMR chemical shifts and coupling constants indicate that increasing the bulk of the alkoxy moiety increases the electron density at the carbon reaction center, which accelerates hydrolysis. Analysis of the Jencks-More-O'Ferrall free energy diagram for the reaction provides support for concerted proton transfer and C-O bond breaking in the transition state for hydrolyses of benzaldehyde acetals with secondary and tertiary alkyl groups in contrast to specific hydrogen catalysis with R = Me and n-Bu. All our results are consistent with rate-determining acid hydrolysis of benzaldehyde dialkyl acetals to hemiacetal intermediates that breakdown rapidly to benzaldehyde.

Full text of DOI:10.1021/jo0202987

Effect of Alk yl Gr ou p Size on th e Mech a n ism of Acid Hyd r olyses
of Ben za ld eh yd e Aceta ls
Alexanders T. N. Belarmino,† Sandro Froehner,‡ and Dino Zanette*
Departmento de Quı´mica, Universidade Federal de Santa Catarina, 88040-900, Floriano´polis,
Santa Catarina, Brazil
J oa˜o P. S. Farah§
Instituto de Quı´mica, Universidade de Sa˜o Paulo, C. Postal 26077, Sa˜o Paulo-SP, 05599-970, Brazil
Clifford A. Bunton|
Department of Chemistry and Biochemistry, University of California, Santa Barbara, California 93106
Laurence S. Romsted
Department of Chemistry and Chemical Biology, Wright & Rieman Laboratories, Rutgers,
The State University of New J ersey, New Brunswick, New J ersey 08903
zanette@qmc.ufsc.br
Received April 29, 2002
Hydrolyses of benzaldehyde acetals, PhCH(OR)2, are specific hydrogen-ion catalyzed when R )
methyl, n-butyl, but with secondary and tertiary alkyl derivatives, R ) i-propyl, s-butyl, t-butyl,
t-amyl, hydrolyses are general-acid catalyzed. The Brønsted R values for both secondary and tertiary
alkyl groups are in the range: R ) 0.57-0.61. A simple iterative procedure was developed to estimate
the individual rate constants for general-acid catalysis by the diacid and monoacid forms of succinic
acid buffer. Plots of log kobs (at [buffer] ) 0 M) against pH are linear for the secondary and tertiary
1
1
acetals, and plots of log kH for the H3O+-catalyzed reaction, 13C and H chemical shifts, and J CH
coupling constants against the Charton steric parameter, ν, for alkoxy groups are linear. The second-
order rate constant, kH, increases about 100-fold on going from R ) Me to R ) t-amyl, indicating
the significant role of steric effects on reactivity. Steric effects upon 13C NMR chemical shifts and
coupling constants indicate that increasing the bulk of the alkoxy moiety increases the electron
density at the carbon reaction center, which accelerates hydrolysis. Analysis of the J encks-More-
O’Ferrall free energy diagram for the reaction provides support for concerted proton transfer and
C-O bond breaking in the transition state for hydrolyses of benzaldehyde acetals with secondary
and tertiary alkyl groups in contrast to specific hydrogen catalysis with R ) Me and n-Bu. All our
results are consistent with rate-determining acid hydrolysis of benzaldehyde dialkyl acetals to
hemiacetal intermediates that breakdown rapidly to benzaldehyde.
In tr od u ction
Hydrolyses of many acetals are catalyzed specifically
by the hydrogen ion, following the A-1 mechanism, as
illustrated for benzaldehyde dialkyl acetals in Scheme
2.1-4 For example, for acid-catalyzed hydrolyses of di-
methyl, and diethyl, benzaldehyde acetals, the first step
is considered to be equilibrium protonation followed by
rate-limiting loss of alcohol to give oxocarbenium ions
that, in water, give the final products in subsequent fast
steps.1-5 The observed first-order rate constants depend
only on hydrogen-ion concentration with no catalysis by
undissociated acids, i.e., by the acid form of a buffer.
Scheme 1 shows acid hydrolyses of benzaldehyde
dialkyl acetals with primary, secondary, and tertiary
alkyl groups in aqueous solution and the structures and
abbreviations used for each dialkyl acetal. The overall
reaction is reversible, but in dilute aqueous solutions it
is driven to the carbonyl and alcohol products.1-3
* To whom correspondence should be addressed.
† E-mail: alexanders.atnb@dpf.gov.br.
‡ E-mail: froehner@qmc.ufsc.br.
§ E-mail: jpsfarah@usp.br.
| E-mail: bunton@chem.ucsb.edu.
E-mail: romsted@rutchem.rutgers.edu.
(3) Cordes, E. H.; Bull, H. G. Chem. Rev. 1974, 74, 581-603.
(4) Fife, T. H. In Advances in Physical and Organic Chemistry; Gold,
V., Bethell, D., Eds.; Academic Press: London, 1975; Vol. 11, pp 1-122.
(5) Buckley, N.; Oppenheimer, N. J . J . Org. Chem. 1996, 61, 8048-
8062.
(1) Carey, R. A.; Sundberg, R. J . Advanced Organic Chemistry. Part
A: Structure and Mechanism, 3rd ed.; Plenum Press: New York, 1993.
(2) Lowry, T. H.; Richardson, K. S. Mechanism and Theory in
Organic Chemistry, 3rd ed.; Harper and Row: New York, 1987.
10.1021/jo0202987 CCC: $25.00 © 2003 American Chemical Society
Published on Web 01/07/2003
706
J . Org. Chem. 2003, 68, 706-717

Acid Hydrolyses of Benzaldehyde Acetals
SCHEME 1
SCHEME 2
SCHEME 3
General-acid catalysis by buffer acids, illustrated for
benzaldehyde dialkyl acetals in Scheme 3, is induced by
various structural modifications of the acetal,1-5 e.g., by
enhancing the stability of the oxocarbenium ion and/or
by releasing strain in the tetrahedral ground state.1,6-8
There are also examples of intramolecular general-acid
catalyses.9,10 Ortho esters are structurally related to
acetals and hydrolyses of some of them are general-acid
catalyzed.11 These reactions are models for hydrolyses at
glycosidic linkages in biological systems.5 With very good
leaving groups, e.g., aryloxy, hydrolyses may be sponta-
neous,5 which is mechanistically equivalent to the rapid
hydrolyses of R-haloethers following the SN1 mecha-
nism.12,13 Increasing the bulk of alkoxy groups of benz-
aldehyde acetals induces general-acid catalysis; e.g., the
hydrolysis of BTBA is general-acid catalyzed (Scheme 3).6
This mechanistic change is rationalized by assuming that
the greater bulk of the tert-butoxy group inhibits proto-
nation, but accelerates C-O cleavage, so that the two
steps become concerted. General-acid-catalyzed reactions
generally have Brønsted R values in the range 0 < R <
1.14,15 Hydrolysis of BTBA in dilute strong acid is faster
than those of the corresponding dimethyl and diethyl
derivatives, and this rate enhancement has been ascribed
to relief of steric strain in formation of the transition
state.1 An alternative mechanism, rate-determining acid-
catalyzed loss of the tert-butyl cation, was excluded by
determination of the position of bond cleavage by 18O
isotopic labeling.16
Changes in substrate conformation may also be im-
portant5,17 because steric interactions involving both the
alkoxy groups and an alkyl or aryl group at the reaction
center will affect anomeric interactions between the
unshared electron pairs of the oxygens.18 Reducing the
conformational change in going from the initial to the
transition state, e.g., by changing the bulk of the alkoxy
groups,18 should favor reaction and lower the extent of
proton transfer in the transition state. Evidence of steric
effects on initial state conformations should be indicated
1
1
by changes in 13C and H NMR chemical shifts and J CH
coupling constants at the reaction center.9 We combined
these approaches with kinetic studies on the hydrogen-
ion or buffer-mediated hydrolyses of the benzaldehyde
dialkyl acetals listed in Scheme 1 with primary, second-
ary, and tertiary alkyl groups of various sizes. These
acetals were selected because the structural changes
should have minimal electronic effects. The hydrolyses
of secondary and tertiary alkyl groups were carried out
in various buffers, including succinic acid, for which we
separated the contributions of the mono- and diacidic
forms by two different methods. The results confirm a
transition from specific to general-acid-catalyzed hydroly-
sis as the bulk of the alkoxy group increases. Because
acetal hydrolyses involve both proton transfer to oxygen
and breaking the carbon-oxygen bond, structural effects
were analyzed by using a J encks-More-O’Ferrall free
energy diagram. Consideration of the energy barriers to
the putative stepwise reactions is based on existing data
on acid-base and carbenium-ion equilibria.
(6) Anderson, E.; Fife, T. H. J . Am. Chem. Soc. 1971, 93, 1701-
1704.
(7) Craze, G.-A.; Kirby, A. J . J . Chem. Soc., Perkin Trans. 2 1978,
354-356.
(8) Anderson, E.; Fife, T. H. J . Am. Chem. Soc. 1969, 91, 7163-
7166.
(9) Fife, T. H.; Bembi, R.; Natarajan, R. J . Am. Chem. Soc. 1996,
118, 12956-12953.
(10) Brown, J . B.; Kirby, A. J . J . Chem. Soc., Perkin Trans. 2 1997,
1081-1093.
(11) DeWolfe, R. H. Carboxylic Ortho Acid Derivatives: Preparation
and Synthetic Applications; Academic Press: New York, 1970; Vol. 14.
(12) Buckley, N.; Maltby, D.; Burlingame, A. L.; Oppenheimer, N.
J . J . Org. Chem. 1996, 61, 2753-2762.
(13) Knier, B. L.; J encks, W. P. J . Am. Chem. Soc. 1980, 102, 6789-
6798.
(14) Zuman, P.; Patel, R. Techniques in Organic Reaction Kinetics;
Wiley: New York, 1984.
(15) Bell, R. P. The Proton in Chemistry; Cornell University Press:
Ithaca, NY, 1959.
Finally, hydrolyses of primary benzaldehyde acetals
show induction periods by stopped-flow spectroscopy for
primary alkyl acetals that are too fast to be observed by
(16) Cawley, J . J .; Westheimer, F. H. Chem. Ind. (London) 1960,
656-656.
(17) Li, S.; Kirby, A. J .; Deslongchamps, P. Tetrahedron Lett. 1993,
34, 7757-7758.
(18) Anderson, J . E.; Heki, K.; Hirota, M.; J orgensen, F. S. J . Chem.
Soc., Chem. Commun. 1987, 554-555.
J . Org. Chem, Vol. 68, No. 3, 2003 707

Belarmino et al.
TABLE 1. Secon d -Or d er Ra te Con sta n ts, kH a n d kHA, for th e Acid -Ca ta lyzed Hyd r olyses of Dia lk yl Ben za ld eh yd e
Aceta ls a t 25 °C in Aqu eou s Bu ffer s w ith 0.5 M KCl
buffer (pH),b no. of data pts,
r ) linear correlation coefficient
kH,
kHA,
acetal
M-1 s-1
M-1 s-1 a
BMA (Me)c
BBA (n-Bu)c
BIPA (i-Pr)
succinate (4.50, 4.90), 8 pts
succinate (4.50, 4.90), 8 pts
21.3( 2.3%
78.3 ( 0.38%
formate (3.60, 4.20, 4.50), 15 pts, r ) 9971
acetate (4.49, 4.70, 5.15), 15 pts, r ) 9920
formate (3,55, 4.10), 8 pts, r ) 0.9957
acetate (4.50, 5.00), 8 pts, r ) 0.9978
formate (3.60, 4.20, 4.50), 15 pts, r ) 9977
acetate (4.49, 4.70, 5.15), 15 pts, r ) 0.9968
phosphate (6.60, 7.10), 10 pts, r ) 0.9996
acetate (4.60, 5.00), 10 pts, r ) 0.9996
phosphate (6.60, 7.10) 10 pts, r ) 0.9999
176
244
161
0.129
0.0266
0.170
BSBA (s-Bu)
BTBA (t-Bu)c
245
0.0323
3.04
2080
2450
2560
2110
2530
0.358c
0.0308c
0.556
BTAA (t-Am)
0.0610
a
b
Values for kHA obtained by using eq 2 (see text). Buffer pKa values: formate, 3.85; acetate, 4.75; phosphate, 6.75; succinate (H2A),
4.21; succinate (HA-), 5.64 (ref 25). c Literature values (25 °C, M-1 s-1): BMA, 60.5 (ref 19); BBA, 73 (ref 24); BTBA, 2000 (ref 24); BTBA,
kH ) 2950, kHA ) 0.325; acetate, 0.029; phosphate (ref 6).
TABLE 2. Secon d -Or d er Ra te Con sta n ts, kH, kH A, a n d kHA, for th e Gen er a l-Acid -Ca ta lyzed Hyd r olyses of Secon d a r y
a n d Ter tia r y Dia lk yl Ben za ld eh yd e Aceta ls a t 252 °C in Su ccin a te Bu ffer s w ith 0.5 M KCla
succinate form (pH), no. of data pts,
r ) linear correlation coefficient
kH,
kH A or kHA
,
2
acetal
M-1 s-1
M-1 s-1
BIPA (i-Pr)
H2A, I (4.21), 10 pts, r ) 0.9753
200
207b
221
0.109
0.00421
0.116
0.0144
0.106
0.00608
0.119
0.0144
1.62
HA-, I(5.10, 5.40), 10 pts, r ) 0.7625
H2A, II (4.21, 4.50), 10 pts., r ) 0.9994
HA-, II (5.40, 80 & 100 mM buffer), av 2 pts
H2A, I (4.01, 4.56), 10 pts, r ) 0.9767
HA-, I (5.10, 5.40) ,10 pts, r ) 0.9711
H2A, II (4.01, 4.56), 10 pts, r ) 0.9710
HA-, II (5.40, 80 & 100 mM buffer), av 2 pts
H2A, I(4.67), 5 pts, r ) 0.9996
BSBA (s-Bu)
BTBA (t-Bu)
BTAA (t-Am)
209c
1920
2260d
2080
2240e
HA-, I (5.50, 5.73), 10 pts, r ) 0.9332
H2A, II (4,67), 5 pts, r) 0.9997
0.0785f
1.71
HA-, II (5.50, 5.73), 10 pts, r ) 0.8927
H2A, I (4.67, 5.01), 10 pts, r ) 0.9873
HA-, I (5.69, 5.91), 10 pts, r ) 0.9759
H2A, II (4.67, 5.01), 8 pts, r ) 0.9858
HA-, II (5.69, 5.91), 10 pts, r ) 0.9803
0.0584f
2.51
0.164
2.55
0.158
a
Values of kH A and kHA Calculated by methods I and II. Method I: kH, kH A, and kHA obtained by the iteration method. Method II: kHA
2
2b
obtained by difference; kobs - kH(av)[H3O+] - kH A[H2A]. See text for details. kH ) (244 + 176 + 200)/3 ) 207 M-1 s-1
.
c kH ) (161 + 245
f Literature
2
+ 220)/3 ) 209 M-1 s-1
.
kH ) (1920 + 2080 + 2450 + 2600)/4 ) 2260 M-1 s-1
.
e kH ) (2110 + 2530 + 2080)/3 ) 2240 M-1 s-1
.
d
value of kHA (25 °C M-1 s-1, succinate buffer) for BTBA 0.237 (ref 6). Buffer pKa values: see Table 1.
g
conventional methods.19,20 These short induction periods
are attributed to formation of transient hemiacetal
intermediates within the first few seconds of the reaction
and do not perturb the relatively slow overall hydrolyses.
However, J ensen21,22 and co-workers and Capon23 report
significant induction periods (1-2 min) for some primary
acetals and significant deviations from first-order kinetics
in the pH range 5-7 for the hydrolysis of BTBA, a break
in the log kobs-pH profile for BTBA, and a marked
increase, by a factor of approximately 10, in the observed
second-order rate constant for BTBA hydrolysis at pH >
7. They conclude that at a low pH, ca. e5, the rate-
determining step is a breakdown of the hemiacetal
intermediate (second step in Scheme 3), followed by a
transition to rate-determining acetal hydrolysis above pH
7. We see no break in the log kobs-pH profiles for any
hydrolyses, in particular those of the tertiary acetals,
BTBA and BTAA. In sum, all our results are consistent
with rate-determining acid hydrolysis of primary, sec-
ondary, and tertiary benzaldehyde dialkyl acetals, fol-
lowed by rapid breakdown of the hemiacetal intermediate
to benzaldehyde.
Resu lts
Kin etics. Tables 1 and 2 list all the second-order rate
constants (M-1 s-1) for hydronium ion, kH, and buffer-
catalyzed hydrolyses, kHA and kH A, of benzaldehyde
2
acetals at 25 °C in 0.5 M KCl. Table 1 lists second-order
rate constants, kH, for hydrolyses of BMA and BBA in
formate, acetate, and phosphate buffers. Table 2 lists
estimates of second-order rate constants for buffer-
catalyzed hydrolyses by succinic acid, H2A, and its
monoanion, HA-, obtained by methods I and II (see
Appendix). All values of the observed first-order rate
constants, kobs, obtained at each buffer concentration and
pH that were used to calculate the second-order rate
constants are compiled in the Supporting Information.
Values of kobs are averages of three kinetic runs and were
determined in 3-4 different buffers across a concentra-
(19) Engell, K. M.; McClelland, R. A.; Sorensen, P. E. Can. J . Chem.
1999, 77, 978-989.
(20) Findley, R. L.; Kubler, D. G.; McClelland, R. A. J . Org. Chem.
1980, 45, 644-648.
(21) J ensen, J . L.; Martinez, A. B.; Shimazu, C. L. J . Org. Chem.
1983, 48, 4175-4179.
(22) J ensen, J . L.; Lenz, P. A. J . Am. Chem. Soc. 1978, 100, 1291-
1293.
(23) Capon, B. Pure Appl. Chem. 1977, 49, 1001-1007.
708 J . Org. Chem., Vol. 68, No. 3, 2003

Acid Hydrolyses of Benzaldehyde Acetals
tion range of 10-250 mM buffer at several different pH
values. Where comparable, second-order rate constants
are in reasonable agreement with literature values6,24
(see Tables 1 and 2). Our value of kHA for BTBA in
hydrogen succinate buffer is lower than earlier values
(see Table 2), but these values probably included minor
contributions from catalysis by succinic acid, because
catalysis by the two acid forms had not been separated
(see below).
Hydrolyses of BMA and BBA are not buffer catalyzed
up to 0.25 M succinate. Reactions are first-order in [H+]
and values of kH in Table 1 are averages of eight sets of
rate data. Note: square brackets indicate concentration
in molarity here and throughout the text.
Hydrolyses of the other acetals are catalyzed by the
acid form of the buffer, Tables 1 and 2, and by H3O+. For
reactions in formate, acetate and phosphate buffers, the
relationship of kobs to acid concentration is given by
F IGURE 1. Plot of kobs/[H+] versus [HA]/[H+] based on eq 2
for BTBA in formate buffer at pH 3.60, 4.20, and 4.50. Values
of kH and kHA were estimated from the slopes and intercepts,
respectively, of the least-squares fit, correlation coefficient: cc
) 0.9977 (see inserted equation).
kobs ) kH[H+] + kHA[HA]
(1)
Values of the two rate constants were obtained from the
slope, kHA, and intercepts, kH, of data plots based on
kobs
[H+]
k
HA[HA]
[H+]
) kH
+
(2)
Figure 1 shows a typical plot of this type for hydrolysis
of BTBA in formate buffer at pH 3.6, 4.2, and 4.5. For
all the substrates, correlation coefficients for such plots
are g0.992 (Table 1).
For reactions in succinate buffers (Table 2), kobs is given
by
kobs ) kH[H+] + kH A[H2A] + kHA[HA-]
(3)
2
Contributions by H2A and HA- to kobs were separated by
working over a significant pH range and using literature
values of the first and second dissociation constants of
pKa(1) ) 4.21 and pKa(2) ) 5.6425 and by assuming that
these values are insensitive to ionic strength up to 0.5
M KCl. To check this assumption, we measured pKa of
formic and acetic acid under our kinetic conditions, i.e.,
0.5 M KCl (see Experimental Section), and obtained 3.89
and 4.74, which are similar to the literature values,
respectively, of 3.85 and 4.7525 and support our assump-
tion that pKa values of all the buffers are insensitive to
ionic strength.
F IGURE 2. Plots of Y values (see text and Appendix) for the
acid-catalyzed hydrolysis of BTBA in succinic acid buffer as a
function of buffer-hydronium concentration ratios at 25 °C.
Plots are from the fourth (last iteration) used to estimate final
values of kHA (A) and kH (B). The insets indicate the pH of
A
2
the buffers and the fitting equation and correlation coefficient,
cc, used to estimate kH from the intercepts (plots A and B)
and kHA (plot A) and kH (plot B).
A
2
and concentrations of both forms of the buffer, H2A and
HA-. The concentrations of H2A and HA- at each pH are
obtained by solving quadratic equations that contain the
acidity constants Ka(1) and Ka(2). An iterative procedure
was used to solve for the two rate constants until their
values converged, which usually took four iterations.
Figure 2 shows the results for BTBA in succinic acid. The
value of Y for the ordinate contains several terms
Values of kH and kHA for catalysis by succinic acid
2A
and succinate monoanion were estimated by methods I
and II, which are fully described in the Appendix. Method
I is an iterative process in which initial estimates of kH
2A
and kHA are made by using equations similar to eq 2 for
H2A and HA- at a pH at which H2A or HA- is the
and is different for plots A and B in Figure
including kobs
dominant form of the buffer. Initial estimates of kH A and
2 (see caption). The definitions of Y are in the Appendix.
2
kHA are then substituted into equations that contain
terms that include both rate constants, the acidity
constants Ka(1) and Ka(2) for ionization of succinic acid
These results show that values of kH are generally
better defined, i.e., less scatter in the data, than those of
2A
k
HA, which is true for BTBA (Figure 2) and the other
acetals (see correlation coefficients in Table 2). The
primary reason for the greater scatter in the estimates
and uncertainty of the values of kHA is that they are an
(24) Froehner, S. J .; Nome, F.; Zanette, D.; Bunton, C. A. J . Chem.
Soc., Perkin Trans. 2 1996, 673-676.
(25) Kortum, G.; Vogel, W.; Andrussow, K. Dissociation Constants
of Organic Acids in Aqueous Solution; Butterworth: London, 1961.
order of magnitude smaller than those of kH
.
2A
J . Org. Chem, Vol. 68, No. 3, 2003 709

Belarmino et al.
F IGURE 4. log kobs ([buffer] ) 0 M)-pH profiles for the
secondary and tertiary acetals. Values of kobs were obtained
by extrapolation from intercepts of kobs-[buffer] profiles, e.g.,
Figure 3. Solid lines are least-squares fits with the slopes and
correlation coefficients listed in this figure. The dashed line
has a slope of 1 positioned to intersect as many secondary
acetal data points as possible.
TABLE 3. F ittin g P a r a m eter s, Slop es, In ter cep ts, a n d
Cor r ela tion Coefficien ts for kobs-[Bu ffer ] P lots fr om th e
F its for th e Rep r esen ta tive Da ta in F igu r e 3, Mea n
Cor r ela tion Coefficien ts a n d Sta n d a r d Devia tion s for All
kobs-[Bu ffer ] P lots for BIP A, BSBA, BTAA, a n d BTBA
Used To Cr ea te F igu r e 4a
F IGURE 3. Effect of increasing [buffer] on kobs for the buffer-
catalyzed hydrolyses of secondary and tertiary acetals at
constant pH, 0.5 M KCl at 25 °C. The acetal and buffer types
and pH for each data set are indicated in the figure. Lines are
least squares correlations and were used to obtain kobs at
[buffer] ) 0 M; see Figure 4.
acetal (pH)b
slope
intercept
corr coef
BTBA (5.01
BSBA (4.01)
BTBA (5.15)
BTAA (5.91)
BIPA (5.40)
BSBA (5.15)
0.120
0.051
0.098
0.073
0.0067
0.011
0.021
0.022
0.018
0.029
0.0013
0.0025
0.977
0.975
0.996
0.981
0.988
1.000
Method II was used to check the results obtained by
method I. In this approach, values of kH A were estimated
2
first at the lower pH by assuming that kHA[HA-] ≈ 0,
that kH was best defined by the average of kH values from
plots of eq 1 for the other buffers, and then estimates of
acetalc
no. of plots
mean corr coef
std dev
BIBA
BSBA
BTAA
BTBA
10
8
9
0.996
0.994
0.997
0.995
0.004
0.008
0.003
0.006
kH were obtained from slopes of plots of kobs - kH(av)-
2A
[H3O+] against [H2A]. This value of kH was used to
2A
12
obtain estimates of kHA from similar plots that include
terms for [HA-] at higher pH values. Table 2 summarizes
the results from method II, pH values of the buffers used
for each calculation, and the correlation coefficients. As
with method I, correlation coefficients in method II are
a
b
Complete data sets in the Supporting Information. From
least-squares fits of data in Figure 3. Corr coef ) correlation
coefficient, Std. Dev. ) standard deviation. c Least squares data
on kobs-[buffer] plots used to calculate log kobs values at each pH.
usually better for kH A than for kHA. The value of kHA for
2
all kobs-[buffer] plots obtained for each acetal. Note that
the mean correlation coefficients are between 0.99 and
0.999.
BSBA by method II is the least well-defined because only
two data points at the highest buffer concentration gave
reasonable values of kHA due to errors caused by small
differences between two large numbers in the subtrac-
Figure 4 shows that for all substrates and buffers, plots
of log k ([buffer] ) 0 M) against pH are linear with slopes
ca. 0.96 for the tertiary and about 0.83 for the secondary
acetals. Note that the slopes for the two tertiary and the
two secondary acetals are essentially the same. The
dashed line shows an imposed slope of -1.0 on the data
for the secondary acetals. Because the buffers involve
different anions, deviations from unit slopes are probably
due to specific salt effects, even though we used 0.5 M
KCl as a “swamping” electrolyte. Changing ion type may
affect estimates of hydrogen-ion activity, or concentra-
tion, from use of a glass electrode, and also reaction rates.
Specific salt effects have been observed in hydrolyses of
acetals and ortho esters catalyzed by dilute strong acids,
tion. For the remaining buffers, however, values of kH
2A
and kHA by methods I and II are in reasonable agreement.
Correlations of log kobs (at [buffer] ) 0 M) against pH
(Figure 4) were obtained by extrapolating plots of kobs
against [buffer], which gives values of the first-order rate
constants of the hydrogen-ion-catalyzed reactions at the
appropriate pH. Data are in the Supporting Information.
Figure 3 shows six example plots, including four with the
poorest correlation coefficients, and two with much better
correlation coefficients. Table 3 summarizes the slopes,
intercepts, and correlation coefficients, for each least-
squares fit in Figure 3. Table 3 also summarizes the
mean correlation coefficients and standard deviations for
710 J . Org. Chem., Vol. 68, No. 3, 2003

Acid Hydrolyses of Benzaldehyde Acetals
TABLE 4. Br øn sted r Va lu es for Secon d a r y a n d
Ter tia r y Dia lk yl Ben za ld eh yd e Aceta ls a t 25 °C in
Aqu eou s Bu ffer s w ith 0.5 M KCl Obta in ed by Meth od s I
a n d IIa
acetal
R (H3O+)
ccd
Secondary
0.61
BIPA, Ib
BIPA, IIc
BSBA, Ib
BSBA, IIc
0.9944
0.9969
0.9968
0.9981
0.58
0.60
0.57
Tertiary
0.60
BTBA, Ib
BTBA, IIc
BTAA, Ib
BTAA, IIc
0.9926
0.9910
0.9994
0.9994
0.61
0.57
0.57
a
Values of R including kH for the H3O+-catalyzed reaction and
-
b
a statistical correction of a -0.3 for H2A and H2PO4
.
Method I
(see text). c Method II (see text). Linear correlation coefficient.
d
F IGURE 5. Typical Brønsted plot used to estimate R, showing
results for BTBA from both method I and method II including
the correlation coefficient, cc.
where there is no uncertainty in hydrogen-ion concentra-
tion.26,27 These medium effects inevitably complicate
comparisons of rate data from experiments with different
buffers and “swamping” electrolytes. In this context we
note that J ensen et al.21 reported a break in plots of
extrapolated values of log k against pH for hydrolysis of
BTBA in acetate, cacodylate, and methylphosphonate
buffers, which they ascribe to a change of mechanism
from to rate-determining acetal hydrolysis to rate-
determining breakdown of the hemiacetal.21 J ensen et
al. also report nonlinearity in the double reciprocal plots
(as in eq 2 and Figures 1 and 2 here) used to estimate
the rate constant for buffer catalysis of BTBA hydrolysis
in cacodylate buffer between pH 6-7.21 All our double
reciprocal plots for secondary and tertiary acetals are
linear at all pH values, although some contain scatter,
e.g., Figure 2. Also, Anderson and Fife6 obtained linear
kobs-[buffer] plots, as we do in Figure 3. On the basis of
the linearity of our log k-pH plots for BTBA and BTAA
in Figure 4, the overall consistency of our results, and
the evidence of specific electrolyte effects on acetal
hydrolyses,26,27 we suspect that the reported break in the
log k-pH profile and the nonlinearity of their buffer plots
in cacodylate buffer reported by J ensen et al. as evidence
for the buildup of hemiacetal intermediate in the hy-
drolysis of BTBA may be caused by a specific electrolyte
effect.21 There is also the possibility of decomposition of
BTBA, which generates impurities and may complicate
determination of rate constants (see Experimental Sec-
tion).
TABLE 5. 13C a n d 1H NMR Ch em ica l Sh ifts a n d C-H
Cou p lin g Con sta n ts in CDCl3 a t 25 °C, th e Ster ic
P a r a m eter , Aver a ge Secon d -Or d er Hyd r on iu m Ion Ra te
Con sta n ts, a n d Loga r ith m s of th e Secon d -Or d er Ra te
Con sta n ts Ra tios Rela tive to BMA a t 25 °C in 0.5 M KCl
13Ca
1Ha
1J CH
kH(av),c
Me
acetal δ (ppm) δ (ppm) (Hz)
νb
M-1 s-1 log(kHR/kH
)
BMA
BBA
BIPA
BSBA
BTBA
BTAA
102.7
101.8
99.8
99.7
94.2
93.0
5.40
5.49
5.58
5.56
5.71
5.73
166.8 0.36
166.7 0.58
162.7 0.75
162.4 0.86
158.7 1.22 2260
156.5 1.35 2240
21.3
78.3
207
209
0
0.565
0.988
0.992
2.03
2.02
a
b
Complete 13C and 1H data in Supporting Information. Spe-
cifically for alkoxy groups from refs 28 and 29. c From Tables 1
and 2.
buffers (see footnotes to Table 2). The R values calculated
by these four separate methods are summarized in the
Supporting Information. Table 4 includes only the R
values that include both average kH values and the
statistical correction. Figure 5 shows the Brønsted plots
for BTBA obtained from Methods I and II. When average
values of kH for the H3O+-catalyzed reactions are excluded
from the Brønsted plots, numerical values of R are higher
than when kH is included and spreads in R are greater.
When the statistical corrections are excluded, there is a
small change in R of about 0.02-0.05 log units. However,
in all four cases there is no significant difference between
R values for the secondary and tertiary acetals. In view
of complications by specific salt effects of the buffer ions
(see above), we believe that this is the significant conclu-
sion regarding effects of secondary and tertiary alkyl
groups on general-acid-catalyzed reactions of dialkyl
benzaldehyde acetals despite differences in steric bulks.
To estimate R for the general-acid-catalyzed reactions,
four different Brønsted plots were prepared for each
acetal from the rate constants listed in Tables 1 and 2,
i.e., plots with and without the value of kH for the H3O+-
catalyzed reaction and plots with and without the
statistical correction for succinic acid, H2A, and phos-
phate monoanion, H2PO4-, which as dibasic acids can
donate two protons. Literature pKa values were used in
the calculations and, as is standard practice, we assume
that the pKa of H3O+ is -1.74 and that at constant ionic
strength, -log [H+] ) pH. The value of kH(av) used in
the Brønsted plots is the average obtained in different
1
1
Table 5 lists 13C and H NMR chemical shifts and J CH
coupling constants at the reaction center, Charton steric
parameters for alkoxy groups, ν,28,29 average values of
rate constants for acid hydrolysis, kH(av), and the log
ratio of the rate constants for each substrate relative to
BMA, log(kHR/kHMe). Figure 6 shows plots of 13C and H
1
(26) Bunton, C. A.; Reinheimer, J . D. J . Phys. Chem. 1970, 74,
(28) Charton, M. J . Org. Chem. 1978, 43, 3995-4001.
4457-4464.
(29) Charton, M. In Similarity Models in Organic Chemistry,
Biochemistry and Related Fields; Zalewski, R. I., Kryowski, T. M.,
Shorter, J ., Eds.; Elsevier: Amsterdam, 1991; Vol. 42, pp 629-688.
(27) McIntyre, D. M.; Long, F. A. J . Am. Chem. Soc. 1954, 76, 3243-
7.
J . Org. Chem, Vol. 68, No. 3, 2003 711

Belarmino et al.
F IGURE 7. J encks-More-O’Ferrall diagram for the acid
hydrolysis of a benzaldehyde dialkyl acetal in aqueous solution
at 25 °C. Corner 1 represents the initial state, corner 2 the
intermediate from preequilibrium protonation, corner 3 the
oxocarbenium ion intermediate, and corner 4 the oxocarbenium
ion/alkoxide intermediates from initial cleavage of the C-O
bond. The vertical axes represent C-O bond cleavage and the
horizontal axes represent O-H bond formation from the initial
state. Corners 2-4 also include estimates of the free energy
difference ∆∆G° relative to that of the initial state, corner 1,
∆G° ) 0.
F IGURE 6. Correlation of the Charton steric parameter, ν,
with the log of relative rate constants (to methoxy), A, with
the C-H coupling constant, B, and proton, C, and carbon, D,
chemical shifts at the central acetal carbon at 25 °C. The
equations are linear least-squares fits of the data showing
slopes, intercepts, and correlation coefficients.
ably, hydrolyses of all our acetals follow a common SE2
mechanism in which protonation on oxygen is concerted
with C-O cleavage, but for hydrolyses of BMA and BBA,
oxygen protonation is essentially complete in their tran-
sition states. This ambiguity may be common for many
specific hydrogen-ion-catalyzed reactions, except where
substrate protonation is observed independently.31 How-
ever, it is generally assumed that acetals of simple
aliphatic aldehydes hydrolyze by A-1 mechanisms.1-4
Structural changes caused in acetals by adding a
phenyl ring, i.e., by going from aliphatic aldehyde to
benzaldehyde acetals, may induce significant changes in
mechanism and the role of catalyzing acids. The oxygens
on benzaldehyde acetals should be less basic that those
of aliphatic aldehyde acetals because inductive electron
withdrawal by the phenyl group at saturated carbon
inhibits preequilibrium proton transfers. However, the
phenyl group should stabilize both oxocarbenium ion
intermediates3-5 and oxocarbenium ion-like transition
states.3,4 Sufficient stabilization should lead to slow
proton transfer probably concerted with C-O bond
cleavage and observable general-acid catalysis. Increas-
ing the bulk of the alkoxy groups should also inhibit
protonation on oxygen by reducing hydration of the
developing oxonium ion. Thus, there is concerted proton
transfer with some combination of a good leaving group,
generation of a very stable oxocarbenium ion, and
electronic and steric effects inhibiting proton transfer to
oxygen.5
Me
chemical shifts, 1J CH coupling constants, and log(kHR/kH
)
values versus ν. The four data sets were also plotted
against a variety of other steric parameters, Es, Es′, Es*,
c
and Es , and Charton’s v for alkyl groups (not shown).30
Correlation coefficients for all the experimentally mea-
sured values of steric bulk are generally g0.91 for each
parameter, showing that all four experimental properties
depend similarly on steric bulk. Figure 6 shows only the
results for ν for alkoxy groups, cc ≈ 0.99, because this
correlation is mostly closely related to substrate struc-
tures, i.e., to calculated molecular volumes of alkoxy
groups based on van der Waals radii.28,29 Note that steric
effects are substantial because kH(av) increases by 2
orders of magnitude in going from methyl (BBA) to t-Am
(BTAA) derivatives.
Discu ssion
Failure to observe general-acid catalysis, as in hydroly-
ses of BMA and BBA, is necessary, but not sufficient,
evidence that reaction involves preequilibrium protona-
tion. As R f 1, the contribution of the kH[H+] term to
k
obs, eq 1, becomes so great that catalysis by HA is almost
undetectable. Going to high buffer concentrations is not
feasible because they may introduce adventitious, and
specific, electrolyte effects.26,27 In practice, demonstrating
general-acid catalysis at R > 0.8 is difficult.15 Conceiv-
Analysis of substituent effects in terms of a J encks-
More-O’Ferrall free energy diagram, is illustrated in
Figure 7.2,32,33 The diagram shows only the step(s)
(31) Cox, R. A. Acc. Chem. Res. 1987, 20, 27-31.
(32) More O’Ferrall, R. A. J . Chem. Soc. B 1970, 274-277.
(33) J encks, W. P. Chem. Rev. 1972, 72, 705-718.
(30) Isaacs, N. S. Physical Organic Chemistry; Longman Group:
Essex, U.K., 1987.
712 J . Org. Chem., Vol. 68, No. 3, 2003

Acid Hydrolyses of Benzaldehyde Acetals
SCHEME 4
generating the oxocarbenium ion (Schemes 2 and 3),
because in water with dilute substrate, formation of the
oxocarbenium ion is rate limiting.2 The two stepwise
pathways along the axes from the reactant corner, 1, are
(a) the classical A-1 mechanism of initial protonation
followed by C-O cleavage or (b) initial C-O cleavage
followed by rapid protonation of the alkoxide ion, i.e.,
spontaneous (uncatalyzed) hydrolysis.5
The similarity of ∆∆G° values at corners 2 and 3 for
BMA is probably fortuitous, but it is consistent with the
idea that oxocarbenium formation from BMA and BBA
involves extensive, although not necessarily complete,
proton transfer, consistent with there being no indication
of buffer catalysis of these reactions, and that relatively
modest structural changes could promote the general-
acid pathway. These comparisons indicate a “late” transi-
tion state. The free energy of the protonated acetal,
corner 2, is much higher than that of the reactants,
corner 1, thus the reaction coordinate for hydrolyses of
the primary acetals should be near the right-hand side
of the diagram and the transition state should be between
species at corners 2 and 3 (Figure 7). The absence of
observed general-acid catalysis in hydrolyses of the
primary acetals BMA and BBA is consistent with rate-
limiting cleavage of the protonated substrate,3-5,15 but a
concerted pathway with almost complete proton transfer
in the transition state, i.e., R between 0.8 and 1,15 cannot
be distinguished kinetically. The free energies of activa-
tion, calculated by use of the Eyring equation, are 15.8
and 10.9 kcal‚mol-1 for hydrogen-ion-catalyzed hydroly-
ses of BMA and BTBA, respectively. The value for
hydrolysis of BMA is higher than the estimated free
energies of protonated MBA and the oxocarbenium ion,
as expected for a stepwise reaction. If we assume that
basicities of BMA and BTBA are similar, the free energy
of activation for hydrolysis of BTBA is lower than that
of protonated substrate, consistent with a concerted
reaction, but higher than that of the putative oxocarbe-
nium ion product; i.e., the free energy of activation is
higher than the predicted free energy of the intermediate.
The J encks-More-O’Ferrall diagram is for a methoxy
derivative, and increasing the bulk of the alkoxy groups
should affect free energies of the structures at corners
1-4 to different extents. Relative to the products, corner
3, the free energy of the initial state, corner 1, should
increase because of steric interactions between the alkoxy
and phenyl groups. Increasing the bulk of the alkoxy
groups will also destabilize the species at corner 2 by
steric crowding and steric hindrance to hydration of the
protonated acetal, and also destabilize the oxocarbenium
ion, corner 3, although to a lesser extent because hydra-
tion of the protonated substrate, corner 2, should be
stronger than that of the charge-dispersed oxocarbenium
ion, corner 3. The free energy of the species at corner 4
will be increased to some extent by steric hindrance to
hydration of the alkoxide ion but there should be release
of steric strain in going from an sp3 to an sp2 carbon at
the reaction center. In any event, ∆∆G° at this corner is
much higher than at corners 2 and 3 and uncertainties
in structural effects on the energy at corner 4 should not
affect interpretation of the results.
We estimated the free energies differences, ∆∆G°
values, of hypothetical intermediates at corners 2-4,
relative to the initial state, corner 1, ∆G° ) 0 at the
standard state of 1 M BMA and 1 M hydronium ions at
25 °C from the equation ∆∆G° ) -RT ln K, with R )
0.001 987 kcal K-1 M-1 and T ) 298.15 K (25 °C). ∆∆G°
for protonation of the acetal, corner 1, BMA, to its
conjugate acid, corner 2, is given by the pKa of the
conjugate acid, ca. -7,3,34 and ∆∆G° ≈ 9.5 kcal/mol. The
free energy difference between corner 4 and the oxocar-
benium ion product, corner 3, is given by the pKa of
MeOH, which is 15.5,35 and ∆∆G° (corner 3 to 4) ≈ 9.2
kcal/mol. The free energy difference between the product
corner, 3, and the initial state, corner 1, is more difficult
to estimate because it includes proton transfer and C-O
bond cleavage of an aryl alkyl ether. We made an
estimate from pKR+ for conversions of arylmethyl car-
bocations into ROH and H3O+ (Scheme 4). This hydration
reaction is similar to the transformation of the oxocar-
benium ion and HOR, corner 3, into the acetal and H3O+,
corner 1. Some values of pKR+ are Ph3C+, -6.6; 4-Me-
OC6H4CPh2+, -3.4, (4-MeOC6H4)2CH+, -5.7.36 We as-
sume that for formation of an oxocarbenium ion pKR+ is
ca. -3, i.e., that the electronic release from the two alkoxy
groups is similar to that from two phenyl and one
2-methoxyphenyl groups, and that pKR+ is similar for loss
of HOH (carbocation acidity) and MeOR ()Me) (corners
1 and 3). We include the molarity of water because water
is lost from the protonated alcohol for pKR+, but in acetal
hydrolyses 1 M alcohol is lost from protonated substrate.
Values of pKR+ are estimated in strongly acidic solvents
from 1 to 10 M acid, and the water molarity will be on
the order of 20-55.5 M, which means that pKR+ is more
negative by 3 (ln 20) to 4 (ln 55.5) units. We therefore
assume that pKR+ for hypothetical formation of the acetal
from the oxocarbenium ion and MeOH ca. -7, and the
free energy for formation of oxocarbenium ion from the
acetal is ∆∆G° (corner 1 to 3) ≈ 9.5 kcal/mol at 25 °C.
Finally, ∆∆G° for formation of the oxocarbenium ion and
MeO-, corner 4, is the sum of ∆∆G° values for formation
of intermediates at corner 3 from corner 1 and corner 4
from corner 3, i.e., 9.5 + 9.2, and ∆∆G° (corner 1 to 4) ≈
18.7 kcal/mol. In sum, the approximate ∆∆G° values of
species at corners 2-4 relative to corner 1 are 9.5, 9.5,
and 18.7 kcal/mol, respectively. These ∆∆G° values are
also shown in Figure 7. They indicate, qualitatively, the
relative free energies of the hypothetical intermediates
and are consistent with the transition state not having
extensive oxocarbenium character, i.e., contributions
from corners 3 and 4.
In general, increasing the bulk of the alkoxy groups
increases the energy of the species at corners 1 and 2
relative to 3 and 4 and the transition state moves toward
the initial state along the reaction coordinate (toward
corner 1 from 3) and orthogonal to the diagonal (toward
corner 4 from 2).2 This movement of the transition state
is consistent with a decrease in R as observed for
(34) Bunton, C. A.; DeWolfe, R. H. J . Org. Chem. 1965, 30, 1371-
1375.
(35) Ballinger, P.; Long, F. A. J . Am. Chem. Soc. 1960, 82, 795-
798.
(36) Coetzee, J . F.; Ritchie, C. D. Solute-Solvent Interactions; Marcel
Dekker: New York, 1969.
J . Org. Chem, Vol. 68, No. 3, 2003 713

Belarmino et al.
hydrolyses of secondary and tertiary acetals and in-
creases in the rate constants corresponding to “earlier”
transition states along the curved pathway in Figure 7.
There should be limited oxocarbenium ion character in
these transition states, consistent with a relatively low
sensitivity to electronic donation from para-substituents
on the phenyl group, which follows σ instead of σ+
Hammett substituent parameters.3,6 In addition, F values
are significantly less negative than for SN1 reactions of
nonionic substrates.2 Thus, proton transfer rather than
C-O cleavage, is a significant source of the energy
barrier in hydrolyses of primary, secondary, and tertiary
acetals that can be regarded as SE2 reactions in which
an oxocarbenium ion is displaced by attack on oxygen
with varying extents of bond making and breaking, Our
conclusions differ from those of J ensen et al. who
examined the effect of substituents in the primary alkoxy
groups in terms of Taft’s electronic and steric parme-
ters.37 They concluded that the transition states had
extensive oxocarbenium ion character, but Taft’s param-
eters contain a mixture of electronic and steric contribu-
tions that are difficult to separate unambiguously in
terms of their contributions to C-O bond cleavage and
proton transfer to the alkoxy oxygen in the transition
state.
inhibits proton transfer by an inductive effect and assists
C-O bond breaking by an electronic effect. Protonation
and formation of oxocarbenium ions have been treated
theoretically for hydrolyses of simple acetals,39-41 but the
phenyl group introduces complicating steric interactions,
and it may also stabilize a forming oxocarbenium ion to
an extent that depends on the torsional angle between
it and the breaking bond. However, the resonance (me-
someric) effect is not all important because electronic
effects follow σ rather than σ+ parameters in Hammett
plots of acetal hydrolyses.3 In sum, increasing the alkyl
group bulk accelerates these reactions and alters extents
of proton transfer and C-O bond breaking in the transi-
tion state. However, ascribing these effects to simple
relief of initial state steric strain, as in SN1 reactions of
congested alkyl derivatives,2 is probably too simplistic in
view of steric effects upon the 13C NMR chemical shifts
and C-H coupling constants discussed below.
NMR Sp ectr a . Both 13C and 1H NMR spectra are
consistent with structures of the acetals as regards
chemical shifts and multiplicities of the aliphatic groups
(see Supporting Information). In addition, variations in
the NMR parameters of the central (R-) C-H atoms are
informative as regards initial state conformations and
1
reactivities. Plots of 13C chemical shifts, δ, and J CH
Values of R for the secondary alkyl derivatives are
essentially the same as those of the tertiary derivatives,
although the latter are significantly more reactive (Tables
1 and 2). Apparently, increasing the bulk of these alkyl
groups has little effect on the extent of proton transfer
but significantly increases that of C-O bond cleavage;
i.e., the transition state shifts toward the bottom side of
the diagram, consistent with an increase in the energies
at corners 1 and 2 relative to 3 and 4 in Figure 7.
Ster ic a n d Electr on ic Effects on Rea ctivity. Steric,
rather than electronic, effects appear to be primarily
responsible for the increase in reactivity in going from
methyl to tert-alkyl derivatives (Tables 1, 2, and 4 and
Figure 6). The steric parameter, ν, for alkoxy groups,
developed by Charton, is based on van der Waals radii
of the constituent atoms and therefore should be free of
complications caused by mixing of electronic and steric
effects that may contribute to other steric parameters
obtained from structure-reactivity correlations.29,30 The
against ν for the alkoxy groups are linear,18,42 as are log
k values (Figure 6). There is extensive evidence that, for
acetals of aliphatic aldehydes, 13C chemical shifts and
1J CH coupling constants at the central (R-) C-H decrease
with the increasing steric bulk of the alkoxy groups. This
decrease in chemical shift is associated with an increase
in electron density at the reaction center, which favors
development of a cationic center in transition state
formation, consistent with observed increases in reac-
tivities. The changes in coupling constants are related
to changes in the dihedral angles between C-H and the
oxygen unshared lone pairs, induced by the anomeric
effect, and steric repulsions between alkyl groups on
oxygen and the substituent on the central carbon. In
contrast to the 13C NMR spectra, 1H NMR chemical shifts
increase with the increasing bulk of the alkoxy groups
(Figure 6), possibly due to changing orientations of the
phenyl group with respect to the axis of the C-H bond
induced by steric interactions with an alkoxy group, as
discussed above. Interference between a tert-alkoxy group
and phenyl group will orient the ring in line with the
C-H bond generating downfield shifts, whereas orienta-
tion toward the face of the phenyl group gives upfield
shifts because of interactions with the ring current of the
phenyl group. This question will be discussed in a
separate paper.
Me
plot of log kHR/kH against v (Figure 6) is linear (cc )
0.989) with a slope of ca. 2, which is clear evidence for
the role of steric effects in controlling reactivity and a
gradual change in extents of bond making and breaking
in the transition state, but we have to consider how they
may mediate electronic effects of the alkoxy groups.
In a simple acetal, the anomeric effect controls initial
state conformation because of stabilizing interactions
between the oxygen lone pairs, but introduction of a
substituent at the reaction center, such as the phenyl
group in these experiments, may generate steric interac-
tions that decrease the oxygen-oxygen electronic inter-
action.18 Chang and Su recently analyzed the factors
contributing to the anomeric effect.38 The situation is
particularly complex with a phenyl substituent, which
not only affects the initial state conformation but also
Con clu sion s
Steric bulk in the alkoxy groups accelerates acid
hydrolyses. For primary alkyl groups the proton is
extensively or completely transferred in the transition
state, but general-acid catalysis of hydrolyses of second-
(39) Woods, R. J .; Andrews, C. W.; Bowen, J . P. J . Am. Chem. Soc.
1992, 114, 859-864.
(40) Webster, A. C.; Fraser-Reid, B.; Bowen, J . P. J . Am. Chem. Soc.
1991, 113, 8293-8298.
(37) J ensen, J . L.; Herold, L. R.; Lenz, P. A.; Trusty, S.; Sergi, V.;
Bell, A. T.; Rogers, P. J . Am. Chem. Soc. 1979, 101, 4672-4677.
(38) Chang, Y.-P.; Su, T.-M. J . Phys. Chem. A 1999, 103, 8706-
8715.
(41) Liang, G. Y.; Sorensen, J . B.; Whitmire, D. J . Compt. Chem.
2000, 21, 329-339.
(42) Pihlaja, K.; Nurmi, T. Isr. J . Chem. 1980, 20, 160-167.
714 J . Org. Chem., Vol. 68, No. 3, 2003

Acid Hydrolyses of Benzaldehyde Acetals
from data obtained over 4-5 half-lives by Guggenheim’s
method with HP-89532-K kinetic software. Rate constants
were determined in triplicate with <5% deviations.
p Ka Deter m in a tion s. The pKa values of formic and acetic
acids were determined by potentiometric titration. A 0.1 M
solution of the acid containing 0.5 M KCl was titrated with
standardized 1.02 M NaOH in 0.5 M KCl. The pKa was
obtained from the inflection point of the pH-volume NaOH
solution curves.
ary and tertiary derivatives indicates concerted proton
transfer and C-O bond breaking. The totality of our
results is consistent with rate-determining acetal hy-
drolysis followed by rapid breakdown of the hemiacetal
intermediate from pH 3.6-7 and at all buffer concentra-
tions. Consideration of probable reaction paths in terms
of a J encks-More-O’Ferrall free energy diagram indi-
cates that there is significant proton transfer in the
transition state and limited buildup of positive charge
at the reaction center, consistent with evidence on
electronic substituent effects.3 Steric effects upon 13C
chemical shifts at the central carbon show that induced
changes in electron densities should affect reactivities
and that accelerations are not due solely to relief of steric
strain involving the alkoxy groups and changes from
tetrahedral to trigonal geometries at the reaction center.
Ack n ow led gm en t. D.Z. is grateful for financial
support from PRONEX and CNPq, 910039/97-6, which
provided support the Brasil-USA collaboration. C.A.B.
and L.S.R. are grateful for support from the National
Science Foundation U. S.-Brasil Cooperative Science
Program (INT-9722458). L.S.R. thanks Lata Racha-
konda for patient word processing, Eric Romsted for
assistance in the analysis of kinetic data, and NSF
(CHE-9985774) for financial support.
Exp er im en ta l Section
Su p p or tin g In for m a tion Ava ila ble: Seventeen tables
containing (a) the complete set of buffer catalysis data used
to estimate the second-order rate constants for the hydronium
ion and buffer-catalyzed reactions for the hydrolyses of all
primary, secondary, and tertiary dialkyl benzaldehyde acetals,
(b) Brønsted R values for secondary and tertiary dialkyl
benzaldehyde acetals comparing three methods of calculating
R, and (c) the 13C and 1H NMR and GC-MS analytical data
on the acetals. This material is available free of charge via
the Internet at http://pubs.acs.org.
Ma ter ia ls. Analytical grade reagents: glacial acetic acid,
40% aqueous formic acid, succinic acid, 85% phosphoric acid,
potassium hydroxide, potassium chloride, anhydrous Et2O and
MeCN, CDCl3 and C2D6OD were used as received. Aqueous
solutions were prepared with either doubly distilled or Milli-
pore treated water.
Acetals were prepared from R,R-dichlorotoluene and the
corresponding alcohol by the method of Cawley and West-
heimer.16 One-half equivalent of R,R-dichlorotoluene was added
dropwise to a solution of the potassium alkoxide in its alcohol,
which was refluxed for 5-7 h. Excess alcohol was distilled off,
the acetal was extracted into Et2O, the mixture was filtered,
and Et2O was evaporated, and the residual oil was vacuum
distilled. The 1H and 13C spectra of the acetals were consistent
with their structures with no stray peaks (Supporting Infor-
mation). The GLC chromatograms for each acetal were single
peaks and their MS spectra are consistent with their struc-
tures (Supporting Information). We did not see molecular ions
in the mass spectra. The highest m/e peaks correspond to
formation of oxocarbenium ions by loss of RO, and the base
peaks correspond to subsequent loss of CnH2n where n is the
number of carbons in the alkyl groups. The 200 1H NMR
spectrum of the sample BTBA that had been used earlier in
the experiments reported here showed that it had decomposed
significantly with time, as indicated by the presence of signals
for both tert-butyl alcohol, benzaldehyde, and unidentified
signals that were absent in the freshly prepared BTBA sample.
BTBA was resynthesized recently, and values of kobs measured
under a few different buffer conditions reproduced those
obtained earlier. It appears that it is desirable to use freshly
prepared samples of the tertiary alkoxy acetals.
Ap p en d ix
This appendix summarizes the processes and math-
ematical derivations for estimating the rate constants for
general-acid catalysis of acetal hydrolysis for the di- and
monoacid forms of succinic acid by Methods I and II. The
original data for these calculations is in the Supporting
Information. The final values of the rate constants for
both methods are listed in Table 2 in the text.
Meth od I. The process is divided into two steps: (a)
obtaining initial estimates of the rate constants by using
simplified forms of the full rate expression and (b) using
the initial estimates in an iterative process to obtain more
accurate values of each rate constant. In general, for each
secondary and tertiary acetal value of kHA were obtained
from kobs values measured at the one or two highest pH
values and values of kH A were obtained from the one or
2
two lowest pH values measured. The decision of what
values to use was based on data quality, i.e., the correla-
tion coefficients of the plots.
1. Estim atin g kHA. The full rate expression for general-
acid catalysis by succinic acid is
Buffer solutions were prepared by dissolving weighed
amounts of the acid and KCl in water (ca. 20 mL), titrating
the solution with ca. 0.9 M KOH to the desired pH and then
diluting with water to 25 mL. The pH was remeasured before
each experiment.
Meth od s. 1H (200 MHz) and 13C (50.27 MHz) NMR spectra
were obtained at 25 °C in CDCl3 with TMS as the internal
standard. 13C spectra of BTBA, BIPA, and BBA are almost
identical in C2D6OD, showing that there are no specific solvent
effects on chemical shifts. Mass spectral data were obtained
on a GC-MS instrument by injecting acetals in CH2Cl2 onto
a CBP1 (nonpolar) column (25 cm length, i.d. 0.22 mm, 0.25
mm particle size) using a temperature gradient of 40-310 °C
at 10 °C/min followed by 5 min at 310 °C and a 70 V EI source
at 25 ° C. Hydrolyses were followed on a diode array spectro-
photometer, and pH was measured on a potentiometer.
Kin etics. A stock solution (10 µL of acetal, 8.0 × 10-3 M in
dry MeCN) was injected into 2 mL of aqueous buffer (see
preparation above) at 25.0 °C. Reactions were followed at 250
nm for at least 10 half-lives. Rate constants were calculated
kobs ) kH[H3O+] + kHA[HA-] + kH A[H2A] (A1)
2
To obtain an initial estimate of kHA, we assume that at
the highest pH used kH A[H2A] ≈ 0 to give
2
kobs ) kH[H3O+] + kHA[HA-]
(A2)
(A3)
Equation A2 is rearranged to give
kobs
[H3O+]
k
HA[HA-]
[H3O+]
) kH
+
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Belarmino et al.
The value of kobs is measured and the value of [H3O+] is
obtained from the measured pH. The value of [HA-] at a
particular pH is obtained by solution of a quadratic
expression (see below).
To obtain a refined value of kHA by iteration, eq A1 is
combined with the expression for the acidity constant of
succinic acid, eq A4, and rearranged to give eq A5:
[AT] ) [H2A] + [HA-] + [A2-
]
(A9)
to give
[H2A]Ka(1) [HA-]Ka(2)
[AT] ) [H2A] +
+
(A10)
[H3O+]
[H3O+]
Substitution for [HA-] gives
[HA-][H3O+]
Ka(1) )
(A4)
(A5)
[H2A]
[H2A]Ka(1) [H2A]Ka(1) Ka(2)
[H3O+]
[AT] ) [H2A] +
+
(A11)
[H3O+]2
kobs
k
HA[HA-]
[H3O+]
k
H A[HA]
2
Y(kHA) )
-
) kH +
[H3O+]
Solving for [H2A] gives
Ka(1)
[AT][H3O+]
[H3O+]2 + Ka(1)[H3O+] + Ka(1) Ka(2)
and Y(kHA), which is equal to the next equality, is plotted
against [HA-]/[H3O+], with the slope ) kHA and the
intercept ) kH.
[H2A] )
(A12)
(b) Der iva tion of th e Con cen tr a tion of HA-. The
expressions for the acidity constants are substituted into
eq A9 to give
2. Estim a tin g k A. To obtain an initial estimate of
H2
k
A, we assume that at the lowest pH used kHA[HA-] ≈
H2
0 and eq A1 becomes
[H3O+][HA-]
Ka(1)
Ka(2)[HA-]
[H3O+]
kobs ) kH[H3O+] + kH A[H2A]
(A6)
+ [HA-] +
(A13)
2
[AT] )
Equation A6 is rearranged to
Equation A13 is solved for [HA-]:
k
H A[H2A]
kobs
[H3O+]
2
[AT][H3O+]Ka(1)
[H3O+]2 + [H3O+]Ka(1) + Ka(1)Ka(2)
) kH
+
(A7)
[H3O+]
[HA-] )
(A14)
The value kobs is measured and the value of [H3O+] is
obtained from the measured pH. The value of [H2A] at a
particular pH is obtained by solution of a quadratic
expression (see below).
Meth od II. To check on the quality of estimates of kHA
and kH obtained by method I, we also estimated these
2A
constants by a second method based on difference.
To estimate kH A, the definition of the first acidity
2
To obtain a refined value of kH by iteration, eq A1
2A
constant Ka(1) for succinic acid was combined with the
mass balance equation
was rearranged to give
k
[H2A]
kobs
[H3O+]
kHA[HA]
[H3O+]
[HA-][H3O+]
H2A
Y(kH A) )
-
) kH
+
(A8)
[H3O+]
2
Ka(1) )
(A15)
(A16)
[H2A]
[AT] ) [H2A] + [HA-]
and Y(kH A), which is equal to the next equality, is plotted
2
against [H2A]/[H3O+], with the slope ) kH and the
2A
intercept ) kH.
to give eq A17, which was used to solve for [H2A] at the
lowest or two lowest pH values used, assuming that
[HA-] was negligible because kHA is about 10 times
3. Iter a tive P r oced u r e. The first iteration of eq A5,
in which the initial estimate of kH is obtained from eq
2A
A7, gives a value of kHA. The concentrations of HA-, H2A,
and A2- are obtained from quadratic equations for their
concentrations by using the first and second pKa values
for succinic acid (see below). The value of kHA obtained
from eq A5 is then plugged into eq A8 to give the first
smaller than kH A (see Table 2 in the text). The values of
2
[AT][H3O+]
Ka(1) + [H3O+]
[H2A] )
(A17)
iterative estimate of kH A. This value of kH is then
plugged back into eq A5 to obtain the second iterative
estimate of kHA. This process is repeated until the
2A
2
[H2A] at each buffer concentration were substituted into
eq A18 and values of kH were obtained from the slope
2A
of kobs - kH(av)[H3O+] vs [H2A], where kH(av) is the
average value of kH obtained from values of kH from the
intercepts of plots using the other buffers, e.g., formate,
acetate, and phosphate (see Table 2 in text).
separate values of kHA and kH A converge, which usually
2
takes four iterations.
4. Der iva tion of Qu a d r a tic Equ a tion s for th e
Dia cid , H2A, a n d Mon oa cid , HA-, F or m s of Su ccin ic
Acid in Ter m s of th e Tota l Stoich iom etr ic Su ccin ic
Acid Con cen tr a tion , AT. (a ) Der iva tion of th e Con -
cen tr a tion of H2A. The expressions for the acidity
constants are substituted into the mass balance equation
for succinic acid
kobs - kH(av)[H3O+] ) kH A[H2A]
(A18)
2
To estimate kH A, the definition of the second acidity
2
constant Ka(2) for succinic acid, eq A19, was combined
716 J . Org. Chem., Vol. 68, No. 3, 2003

Acid Hydrolyses of Benzaldehyde Acetals
with the mass balance equation (eq A20), to give eq A21,
[A2-][H3O+]
[AT][H3O+]
Ka(2) + [H3O+]
[HA-] )
(A21)
Ka(2) )
(A19)
values of [HA-] at each buffer concentration were sub-
[HA-]
[AT] ) [HA-] + [A2-
stituted into eq A22 and values of kHA were obtained from
the slope of kobs - kH(av)[H3O+] - kH A[H2A] vs [HA-].
2
]
(A20)
kobs - kH(av)[H3O+] - kH A[H2A] ) kHA[HA-] (A22)
2
which was used to solve for [HA-] at the two lowest pH
values used, assuming that [H2A] was negligible. The
J O0202987
J . Org. Chem, Vol. 68, No. 3, 2003 717