Kinetics and mechanism of the cyclization of ω-(p-nitrophenyl)-hydantoic acid amides: Steric hindrance to proton transfer causes a 104-fold change in rate

Article abstract of DOI:10.1039/b211040g

The pH-rate profiles for the cyclization of primary 2,3-dimethyl and 2,2,3-trimethyl-hydantoinamides (2-UAm and 3-UAm respectively) differ strikingly from those for the cyclizations of the corresponding N-methylated amides 2-MUAm and 3-MUAm; which are dominated by the water reaction, spanning some 6 pH units. For the cyclization of UAm the plateau extends over no more than two pH units. The difference is due to the slower base-catalyzed cyclization of the N-methylamides. The solvent kinetic isotope effect for this hydroxide-catalyzed reaction is close to 1.2, consistent with a slow protonation by water of the amino-group of the negatively charged tetrahedral intermediate. General base catalysis was observed with bases of pK_BH up to 8. The Bronsted β are compatible with a hydrogen bonding mechanism for the GBC. In the gem-dimethyl compounds 3 the leaving group is flanked by substituents on both sides. The N-methyl group in 3-MUAm hinders frontal access of the proton, causing a 14000 fold decrease in rate. This is only 3800 fold in the compound with one methyl group at position 2.

Full text of DOI:10.1039/b211040g

View Article Online / Journal Homepage / Table of Contents for this issue
Kinetics and mechanism of the cyclization of ꢀ-(p-nitrophenyl)-
hydantoic acid amides: steric hindrance to proton transfer causes
a 104-fold change in rate†
Violina T. Angelova,a Anthony J. Kirby,b Asen H. Koedjikova and Ivan G. Pojarlieff*a
a Institute of Organic Chemistry, Bulgarian Academy of Sciences, ul. Acad. G. Bonchev block 9,
Sofia 1113, Bulgaria. E-mail: ipojarli@orgchm.bas.bg
b University Chemical Laboratory, Cambridge, UK CB2 1EW
Received 11th November 2002, Accepted 27th January 2003
First published as an Advance Article on the web 12th February 2003
The pH-rate profiles for the cyclization of primary 2,3-dimethyl and 2,2,3-trimethyl-hydantoinamides (2-UAm and
3-UAm respectively) differ strikingly from those for the cyclizations of the corresponding N-methylated amides
2-MUAm and 3-MUAm; which are dominated by the water reaction, spanning some 6 pH units. For the cyclization
of UAm the plateau extends over no more than two pH units. The difference is due to the slower base-catalyzed
cyclization of the N-methylamides. The solvent kinetic isotope effect for this hydroxide-catalyzed reaction is close
to 1.2, consistent with a slow protonation by water of the amino-group of the negatively charged tetrahedral
ϩ
intermediate. General base catalysis was observed with bases of pKBH up to 8. The Brønsted β are compatible
with a hydrogen bonding mechanism for the GBC. In the gem-dimethyl compounds 3 the leaving group is
flanked by substituents on both sides. The N-methyl group in 3-MUAm hinders frontal access of the proton,
causing a 14000 fold decrease in rate. This is only 3800 fold in the compound with one methyl group at position 2.
One of the main problems in studying chemical reactions to
mimic in vivo processes is that they are very slow under
biologically relevant conditions. Attaching the reactants to the
same molecule converts bond forming reactions into cycliz-
ations; the consequent entropic gain, reinforced by thermo-
dynamically favourable ring formation, can bring about
accelerations of up to 108-fold.1 Reaction can be made faster
still by introducing steric strain in the substrate, and a con-
venient way to do this is by introducing substituents into the
interconnecting chain: the resulting increase in rate defines the
gem-dimethyl effect (GDME).2
In a series of papers3–5 on the cyclization of hydantoate‡
esters, UE, however, unexpected behaviour was observed in the
OH- catalyzed reaction.§ When all hydrogens in the chain were
replaced by methyl groups, structure 3, the GDME either dis-
appeared or was reversed, as a result of a change of mechan-
ism. These observations were explained by steric hindrance
slowing down proton transfer to the leaving ethoxy group, and
thus making k2 in Scheme 1 the rate determining step (rds) as
opposed to k1 in esters 1-UE and 2-UE. This does not happen
upon acid catalysis going through T؉ ¶ because only T؊ is so
unstable that protonation competes with heavy atom reorganiz-
ation. When the ester OEt group was replaced by OH, in the
Scheme 1
permethylated acid 3-UAc, a normal GDME and no change of
mechanism (i.e. rate determining formation of T؊) was
observed.6 Evidently the replacement opened sufficient access
to the leaving group that protonation did not become the
slow step.
acids and unsubstituted amides on the other, the presence of
an N-methyl group in structure 3 should similarly hinder the
approach of the proton. In this case the expected result is
simply a decrease in the reaction rate since proton donation to
the leaving group already is rate limiting. The present paper
identifies this effect and shows it to be remarkably strong: the
ratio of the corresponding rate constants for 3-UAm and for
3-MUAm amounts to four powers of ten.
In the case of hydantoic amides the amino group is a poorer
leaving group and its protonation in T؊ is usually rate determin-
ing under base catalysis.7,8 Because of the similarities in steric
requirements of esters and methylamides on the one hand and
† Electronic supplementary information (ESI) available: Observed first-
order rate coefficients, constants for solvent and buffer catalysis for the
cyclization reactions. See http://www.rsc.org/suppdata/ob/b2/b211040g/
‡ The IUPAC name for hydantoic acid is ureidoacetic acid or N-
carbamoylglycine.
§ The study encompassed also 5-methyl and 5-phenyl hydantoic esters,
refs. 3 and 4
Experimental
Uncorrected melting points were measured in capillaries, IR
spectra on a Specord IR 75 or Bruker IFS 113v instrument, UV
spectra on a Specord UV Vis or a UNICAM SP 800 spectro-
photometer, and NMR spectra on a Bruker DRX 250 instru-
¶ According to ref. 6 most likely the amino group is positively charged.
T h i s j o u r n a l i s © T h e R o y a l S o c i e t y o f C h e m i s t r y 2 0 0 3
O r g . B i o m o l . C h e m . , 2 0 0 3 , 1, 8 5 9 – 8 6 5
859

View Article Online
ment. Chemical shifts are quoted in ppm as δ values against
TMS and couplings in Hz.
exclude participation of this second pathway, which has been
observed by Machácek et al.9 with some simple hydantoin
amides when cyclization is slow. However, such participation
can be confidently eliminated upon consideration of the rates
involved. The second order rate constant for alkaline hydrolysis
of hydantoinamide itself can be estimated as 2 × 10Ϫ4 dm3
molϪ1 sϪ1 from the data of ref. 9 while the second order rate
constant for OH--catalyzed cyclization of 2-UAm is about 1 ×
104 dm3 molϪ1 sϪ1. The sterically hindered amides studied in this
paper should be hydrolyzed much more slowly than unsubsti-
tuted hydantoinamide as is well documented with carboxylic
acid derivatives. Furthermore, the reaction of 3-UAm com-
pared to 2-UAm (vide infra) is faster due to the GDME specific
to cyclization reactions.
Materials
Inorganic reagents and buffer components were of analytical
grade and used without further purification. Potassium
hydroxide and buffer solutions were prepared with CO2-free
distilled water. D2O, 99 atom%, was from Aldrich.The prep-
arations of 1,5-dimethyl-3-(4-nitrophenyl)hydantoin|| and
1,5,5-trimethyl-3-(4-nitrophenyl)hydantoin have been described
previously.7
3-Methyl-5-(4-nitrophenyl)hydantoinamides
At pH-values above 7 the change of absorbance of 2-UAm
with time takes the form depicted in Fig. 1. This is due to the
consecutive reactions:7
The parent amides of 2-methylaminopropionic acid and 2-
methylaminoisobutyric acid were prepared from the respective
esters as described in ref. 7 and papers quoted therein. Several
days in a saturated ammonia solution in methanol at room
temperature were needed to complete conversion conveniently
followed by IR of the dry residues. Literature melting points
were obtained after crystallization from CHCl3. The methyl-
amino amides were treated with p-nitrophenylisocyanate in
dry benzene as described in ref. 7 and used without further
recrystallization because of their ready cyclization.
2-(1Ј-Methyl -3Ј-(4-nitrophenyl)ureido)propionamide. Yield
25%, mp 90 ЊC, λmax(H2O)/nm 330 nm, νmax/cmϪ1 3315 and 3175
(NH), 1690 (COurea), 1640 (COamide); δH (DMSO-d6) 1.273 (3H,
d, J 7.1, 2-Me), 2.903 (3H, s, 1Ј-Me), 4.710 (1H, q, J 7.2, 2-H),
7.763 (2H, d, J 9.2, o-H), 8.1458 (2H, d, J 9.2, m-H), 7.356 and
7.0557 (1H, s, NHamide, hindered rotation, exchangeable with
D2O), 9.0260 (1H, s, NHurea); MS electrospray Mϩ ϩ Na
289.0921
2-Methyl-2-(1Ј-methyl -3Ј-(4-nitrophenyl)ureido)propionam-
ide. Yield 20%, mp 146–148 ЊC, λmax(H2O)/nm 330 nm, νmax
/
cmϪ1 3440 and 3270 (NH), 1680 (COurea), 1660 (COamide);
δH (DMSO-d6)** 1.347 (3H, s, 2-Me), 2.991 (3H, s, 1Ј-Me),
7.717 (2H, d, J 9.5, o-H), 8.130 (2H, d, J 9.5, m-H), 6.9 br s (1H)
6.6 br s (1H); MS CI Mϩ ϩ 1, 281.3.
Fig. 1 The change of absorbance of 2-UAm with time in a 0.008 M
Tris buffer at 0.5 fraction base pH 8.30 (I = 1 M KCl) at 25.0 ЊC.
The line is calculated using eqn. (2) and the rate constants reported in
Table 1 and ref. 6.
Kinetic measurements
(1)
These were carried out as described previously.7 Due to the
ready conversion of 3-UAm into hydantoin in DMSO, stock
solutions were freshly prepared before each experiment.
Multiple scan spectra taken during the course of the cyclization
Since the extinction coefficients of amide and acid are prac-
tically the same, the rate constants of the integrated kinetic
equation of the system can be readily determined from curves
as that of Fig. 1, using the following eqn. (2):
showed good isosbestic points and the infinity spectra (10τ1/2
)
were found to be identical with the spectra of the respective
3-(4-nitrophenyl)hydantoins at the same concentration. A
“clean” reaction was also supported by well-behaved first order
kinetics. In the case of 2-UAm the kinetics at pH > 7 were
complicated by partial hydrolysis of the product hydantoin: this
process reached an equilibrium with the respective hydantoic
acid (see Results). Solvent kinetic isotope effects were
determined as described previously.6
(2)
Here At and Aeq are the absorbances at times t and infinity, ∆A =
(εH Ϫ εA)C0 where εH and εA are the extinction coefficients of
Hyd and UAm or UAc (3600 dm3 molϪ1 cmϪ1 and 13600 dm3
molϪ1 cmϪ1 respectively at 330 nm), C0 the initial amide concen-
tration, a = k12, b = k23 ϩ k32 and c = k32, the pseudo-first-order
rates at the specific pH. In order to reduce the number of
parameters fitted k23 and k32 were calculated from the rate con-
stants and equations given in ref. 6 using Ke = [Hyd][aOH]/[UAc]
= 4.44 × 10Ϫ6. Thus k12 and Aeq were treated as adjustable
parameters in nonlinear curve fitting by means of the GRAFIT
program. For 0.9 fraction base (FB) phosphate and 0.1, 0.2 and
0.3 FB Tris only the first part of the curve shown on Fig. 1 was
monitored, and the data treated by means of eqn. (2), to correct
for deviations caused by participation of the second reaction.
Results
The data described in the Experimental Section show that the
hydantoinamides are coverted quantitatively into hydantoins.
This could occur directly (Scheme 1) or in two consecutive
steps: hydrolysis of the hydantoinamides to hydantoic acids fol-
lowed by cyclization of the latter. Our recent results with acids
2-UAc and 3-UAc show that they cyclize faster than the amides
in most of the pH-range studied. Thus the kinetic data can not
|| The IUPAC name for hydantoin is imidazolidine-2,4-dione.
** Due to the relatively rapid cyclization in DMSO the spectra show also
the signals of product hydantoin.
The pH-rate profiles
The rate profiles for the cyclization of the hydantoic acid amide
O r g . B i o m o l . C h e m . , 2 0 0 3 , 1, 8 5 9 – 8 6 5
860

View Article Online
Table 1 Rate constants for Hϩ, OHϪ and H2O catalyzed ring closure of amides of 5-(4-nitrophenyl)hydantoic acids at 25.0 ЊC and ionic strength
1.0 M
Compound
kH/dm3 molϪ1sϪ1
kw/sϪ1
kOH/dm3 molϪ1sϪ1
2-UAma
3-UAm
(2.24 0.25) × 10Ϫ4
(1.60 0.12) × 10Ϫ2
(6.33 0.46) × 10Ϫ6
(6.47 0.73) × 10Ϫ5
(1.17 0.08) × 104
(1.12 0.08) × 106
a Parameters calculated by means of eqn. (3) omitting data in phosphate at FB 0.8 and 0.9 and Tris buffers.
Table 2 Buffer catalysis data for the cyclization of the amide 2,3-dimethyl-5-(4-nitrophenyl)hydantoic acid, 2-UAm, at 25.0 ЊC and ionic strength
1.0 M
a
Buffer acid
pKa
Conc. rangeb/mol dmϪ3
FBc
kB/dm3molϪ1sϪ1 b
d
H3NϩCH2CO2H
2.45
0.01–1
0.01–1
0.01–1
0.01–1
0.01–1
0.01–1
0.01–1
0.01–1
0.3
0.5
0.7
0.9
0.3
0.7
0.3
0.7
0.9
0.3
0.5
0.7
0.1
0.3
0.5
0.7
0.8
0.84e
0.9
0.1
0.2
0.3
0.5
0.7
HCO2H
3.57
4.62
(6.67 0.82) × 10Ϫ5
(3.47 0.50) × 10Ϫ4
CH3CO2H
0.01–1
(CH3)AsO2H
6.19
6.48
0.01–0.15
0.01–0.15
0.01–0.15
0.016–0.2
0.032–0.6
0.032–0.5
0.0032–0.4
0.0016–0.008
0.0016–0.008
0.0016–0.008
0.0016–0.008
0.0016–0.008
0.0016–0.008
0.0016–0.008
0.0016–0.008
(7.02 0.45) × 10Ϫ3
Ϫ
H2PO4
0.0196 0.010f
kBOH = (2.22 0.32) × 105
Tris
8.42
0.121 0.012g
Ϫ2
HPO4
12.32
508h
a pKa-values at ionic strength 1 M from ref. 13. b Four or more runs carried out within each concentration range. c The fractions base of glycine and
formate buffers were corrected for the change in pH with dilution. The initial FB values are quoted in the Table. d See text. e Determined from pH.
f Calculated by means of eqn. (6), k2OH = (7,49 0.82) × 106. g Calculated from a linear fit of kbuff against FB. h Calculated from kBOH
.
Table 3 Buffer catalysis data for the cyclization of the amide of 2,2,3-trimethyl-5-(4-nitrophenyl)hydantoic acid, 3-UAm, at 25.0 ЊC and ionic
strength 1.0 M (KCl)
a
Buffer acid
pKa
Conc. rangeb mol dm3
f.b.c
kB/dm3molϪ1sϪ1
H3NϩCH2CO2H
2.45
0.01–1
0.01–1
0.01–1
0.01–1
0.01–1
0.01–1
0.01–1
0.01–1
0.01–1
0.01–1
0.01–1
0.01–1
0.016-0.2
0.016-0.2
0.3
0.5
0.7
0.9
0.3
0.5
0.7
0.9
0.3
0.5
0.7
0.9
0.1
0.3
(5.53 0.65) × 10Ϫ4 d
HCO2H
3.57
4.62
6.48
(2.24 0.14) × 10Ϫ3
0.0203 0.0011
0.927 0.035
CH3CO2H
Ϫ
H2PO4
a pKa-values from ref. 13. b Four or more runs carried out within each concentration range. c The fractions base of glycine and formate buffers were
corrected for the change in pH with dilution. The initial FB values are quoted in the Table. d See text.
2-UAm and 3-UAm are k0-values extrapolated to zero buffer
concentration from the buffer experiments listed in Tables 2 and
3††, above. In the case of 2-UAm above pH 7 these are the k12-
values discussed above. For the HCl solutions, activities, instead
of the usual convention of concentrations, were used to bring
these in line with the buffer data. An activity coefficient of
0.851 at I = 1 M KCl was used7.
In the pH range studied the two hydantoinamides 2-UAm
and 3-UAm gave rise to simple rate profiles (Fig. 2) with slopes
of Ϫ1, 0 and 1 for H3Oϩ, water and OHϪ-catalyzed reactions
respectively:
†† All observed pseudo-first-order rate constants are available as elec-
tronic supplementary information (ESI)†.
O r g . B i o m o l . C h e m . , 2 0 0 3 , 1, 8 5 9 – 8 6 5
861

View Article Online
cyclization of amide 3-UAm was studied in glycine, formate,
acetate and phosphate buffers. For 2-UAm the buffer range was
extended to cacodylate and Tris. With the more acidic glycine
and formate buffers pH increased upon dilution and the
changes in pH and % free base were calculated as described
previously.3 The kinetic behaviour of the substrates in buffers
was examined as follows. The rates determined at a certain
buffer ratio were plotted against total buffer concentration.
These plots were linear allowing the intercept k0 and the slope
kbuff to be obtained from a least squares fit. The kbuff vs. FB plots
were also reasonably linear as demonstrated by the example of
Fig. 3. The intercepts at fraction base 0 and 1 equal the rate
constants for general acid and general base catalysis (GBC)
respectively. No general acid catalysis, GAC, was observed and
in this case the slope of a linear plot kbuff against FB equals kB.
The kinetics are thus described by eqn. (4):
kobs = kHaH ϩ kw ϩ kOH
a
OH ϩ kB[B]
(4)
Two exceptions to this straightforward behaviour were
observed. In the case of 2-UAm in phosphate buffer catalysis
increased much more rapidly than the linear dependence on
percentage of free base expected for simple GBC (Fig. 4a).
Fig. 2 pH-rate profiles for: 2-UAm (triangles: also ϩ: data in 0.8–0.9M
phosphate and *: in Tris buffers) 3-UAm (closed circles) 2-MUAm
(diamonds) and 3-MUAm (open circles).
kobs = kHaH ϩ kw ϩ kOHaOH
(3)
In the case of 2-UAm, however, the data obtained in phos-
phate buffers with >80% free base suggested a leveling of the
rate increase with aOH. Attempts to expand the pH-rate profile
by studying the reaction in Tris buffers proved unsuccessful
because the k0-values in the latter were found to be about three
times smaller than those obtained at the same pH in phosphate
buffers (Fig. 2). This can be attributed to inhibition by the
amine component of the buffer as has been observed in similar
reactions.10 In some cases11 inhibition affects only k0. This
seems to be the present case as judged from the altogether
“normal” behaviour of the kbuff versus fraction free base plot
shown on Fig. 3.
Fig. 3 Plot of kbuff in Tris buffers against fraction base.
The complication in Tris buffers compromised attempts for
kinetic analysis of the deviation from linearity of the data for k0
in phosphate buffers at FB greater than 0.8. For this reason the
rate constants for 2-UAm were derived by means of the simple
eqn. (3) from the data collected in buffers more acidic than 0.8
FB phosphate. These, together with the results for 3-UAm, are
listed in Table 1.
Fig. 4 (a) Plot of kbuff determined in phosphate buffers against
fraction free base (FB). (b) Plot of kbuff/(fraction base) against
hydroxide ion activity, aOH
.
The rising curve in Fig. 4a can be accommodated by adding a
Solvent kinetic isotope effects, kH/kD, SKIE, were measured
for the alkaline reaction in 0.002 M phosphate buffers at 0.3 FB
for 2-UAm and 0.2 FB for 3-UAm are shown in Table 4. Under
these conditions buffer catalysis is less than 10% of the total
reaction.
cross term kBOH[B]aOH
.
As [B] = [buffer]FB
kbuff[buffer] = kB[buff]FB ϩ kBOHaOH[buffer]FB
then
General base catalysis
Buffer catalysis was clearly apparent in all cases studied. The
kbuff/FB = kB ϩ kBOHaOH
(5)
O r g . B i o m o l . C h e m . , 2 0 0 3 , 1, 8 5 9 – 8 6 5
862

View Article Online
Table 4 Brønsted β-values for GBC of hydantoinamides and solvent
kinetic isotope effects in phosphate buffers
The validity of eqn. (5) is illustrated by Fig. 4b. The most likely
cause for the form of this equation is catalysis3 by the phos-
Ϫ3
phate trianion, PO4 its concentration being proportional to
b
Compound
β-values
pHa
kOH/kOD
[HPO4Ϫ2] and [OHϪ] because of the equilibrium:
2-UAm
3-UAm
0.69c
0.79d
5.95
5.88
1.28
1.21
a Average value in H2O of the SKIE experiments. b See Experimental.
c From linear regression of data for catalysis by formate, acetate,
cacodylate, monohydrogenphosphate and Tris (r = 0.996). d From linear
regression of data for catalysis by formate, acetate and monohydrogen-
phosphate (r = 0.998). Inclusion of the point for glycine yields β = 0.75.
Therefore:
β-values were calculated using only the rate data for the
“general” bases listed in Tables 2 and 3. The results of these
excellent linear correlations are summarized in Table 4.
3Ϫ
where kphos is the rate constant for catalysis by PO4
.
A similar upward deviation from linearity was found
for 2-UAm in glycine buffers. The data, however, could not
be accommodated by any simple equation derived for
similar situations observed with hydantoic esters5,12 and
acids.5 The pH of glycine buffers coincides with a region
where the pH-rate profile changes its slope from 0 to 1
indicating a change in mechanism as the likely cause for
non-linearity of the plot. With 3-UAm the extent of the plateau
is much smaller and eqn. (4) describes the data in glycine
satisfactorily.
The rate constants for GBC listed in Tables 2 and 3, were
obtained by a multivariable fit to eqn. (4) of the entire data set
for a given buffer. In this way any variation of pH at constant
FB was accounted for. The constants for specific catalysis,
determined from the rate profile, were inserted as known
parameters. For phosphate catalysis in the case of 2-UAm eqn.
(6) was used which includes the above cross term and the appar-
ent leveling off in catalysis by the solvent species is accounted
for by a denominator term:
Discussion
Mechanism
In our previous report7 on the cyclization of hydantoinmethyl-
amides detailed evidence indicated that the hump in the MUAm
rate profiles between pH 6 and 8 (Fig. 2) is due to a change in
the rate determining step from k1KNH/Kw to k2 (Scheme 1),
resulting in two parallel reactions catalyzed by water and two
reactions catalyzed by OHϪ. In the present case of UAm only a
single OHϪ reaction is observed in the pH-region studied.‡‡
The similarity in Brønsted plots for GBC and SKIE repre-
sented in Table 4 leaves little doubt that the mechanism of
cyclization of the two compounds is the same. For the reac-
tion catalyzed by bases in general and that catalyzed by OHϪ
the mechanistic criteria: β-values and SKIE are best met by a
hydrogen bonding mechanism for the former and trapping of
an unstable intermediate for the second.14 When diffusion con-
trolled proton transfer is rate limiting in the so-called trapping
of unstable intermediates, their further instability enforces
preassociation15 of the substrate and the catalyzing acid or
base because proton transfer within the complex can proceed
faster than diffusion. This rapid donation or abstraction of a
proton can prevent the intermediate from returning to react-
ants and direct its transformation to products. Hydrogen
bonding can stabilize the preassociation complex and so
further lower the energy of the transition state. The dipolar
tetrahedral intermediate T is such a highly unstable inter-
mediate:
(6)
Values of kw and kOH were taken from the rate profile and the
remaining three constants determined as adjustable parameters.
The rate constant for catalysis by Tris is the slope of the linear
dependence kbuff/FB shown in Fig. 3.
Brønsted plots of the statistically corrected rate constants
for GBC are presented in Fig. 5. In order to avoid the strong
influence of the widely spaced points for water (kw/55.5) and
hydroxide anion the slopes of the linear fits, the Brønsted
The α- or β-values for hydrogen bonding equilibria are
around 0.2 and similar values are expected for GBC by hydro-
gen bonding. Catalysis by the general base of the overall reac-
tion can be formulated in the following manner (charges of base
species are omitted):
(7)
Fig. 5 Brønsted plots for GBC of the cyclization of hydantoinamides:
open symbols 2-UAm, closed symbols 3-UAm. Squares represent
catalytic constants listed in Table 2 and Table 3, circle the deviant point
for PO43Ϫ, triangles the data for water and OHϪ catalysis. Least squares
lines are drawn to fit the squares (see Table 4).
‡‡ The apparent leveling in the pH-rate profile upon increasing pH
cannot be interpreted unambiguously.
O r g . B i o m o l . C h e m . , 2 0 0 3 , 1, 8 5 9 – 8 6 5
863

View Article Online
When V = kBH[T؊][BH] it can be readily shown from eqn. (7)
that for the experimentally observed GBC, eqn. (8):
Table 5 Ratio of the rate constants of the primary (UAm) and
secondary (MUAm) 5-(p-nitrophenyl)hydantoinamidesa
Reaction
k2-UAm/k2-MUAm
k3-UAm/k3-MUAm
(8)
kHb
kw
kOH
kH2N
kHCO2
kCH3CO2
kHPO4
kTris
15
3
3800
10
2
14000
15
37
164
260
V = kB[UAm][B]
c
ϩCH2CO2Ϫ
Ϫ
34
58
127
49
Ϫ
Ϫ
β = 1 Ϫ α because β for an ionization equilibrium, KBH, is unit-
y.Thus with α = 0.2 a β-value of 0.8 is expected for the overall
GBC eqn (8). The β-values listed in Table 4 are a little smaller
than 0.8, but still compatible with hydrogen bonding.
The values of the solvent kinetic isotope effect in Table 4
apply for the OHϪ-catalyzed reaction where BH is a water
molecule. The amino group in T؊ should have a pK of about
8.4§§ so the proton donation by H2O is an endothermic reaction
for which a preassociation or hydrogen bonding mechanism is
no longer possible. Simple proton transfer to T؊ is the most
likely mechanism. Such transfers are characterized by a SKIE
of ca.1.2¶¶ when the participating acid and bases differ con-
siderably in pK and this is consistent with the values found for
UAm.||||
a Data for MUAm from ref. 7. b Comparison with kw of ref. 7. c Compa-
b
rison with kOH of ref. 7.
gem-dimethyl derivatives 3. Structure TS-3-MUAm shows how
the amino group is flanked on one side by the geminal methyl
groupsandontheotherbythearylgroup,****withtheN-methyl
group barring the water molecule from the most accessible
route from the front side of the nitrogen atom.
The chosen mechanism for the observed OHϪ catalyzed
cyclization of UAm coincides with the mechanism assigned to
the OHϪ catalyzed reaction of the N-methyl derivatives MUAm
at higher pH. The remaining two reactions, acid and water
catalyzed cyclization of UAm were not our main focus of
attention. Our previous study on MUAm left the question of
the rds of the acid reaction without a definite answer. Two
mechanisms involving slow proton transfers were assigned to
the two water-catalyzed reactions: deprotonation of the
p-nitrophenylureido group concerted with formation of T؊ and
secondly, a rate determining water-mediated switch converting
T0 into T .
When the N-methyl group is removed in gem-dimethyl
derivative 3-MUAm the acceleration is 14000-fold, a rate
increase that is spectacular even though it contains a contribu-
tion from steric strain in T؊ (a factor of 10–15-fold is suggested
by the results from kH). Interestingly the effect, 3600-fold, of
removing the N-methyl group in the derivative with the
monomethylated side chain 2-MUAm is still very large. This
indicates that even a single methyl group at the α-carbon atoms
can exert considerable hindrance to proton transfer to the
methylamino group.
The lower kUAm/kMUAm ratios observed for the catalytic
constants for GBC are not surprising because in a hydrogen-
bonding mechanism the rate determining step does not involve
the proton transfer itself, which is very fast, but formation and
breakdown of T؊ within the hydrogen bonded complex.
Steric hindrance to proton transfer
The major difference in the rate profiles of UAm and the
N-methyl derivatives MUAm demonstrated in Fig. 2 is due
formally to the much larger decrease in kOH for alkaline cycliz-
ation for the latter, compared with the remaining constants. A
decrease in the rate upon N-methylation is often observed in
similar cyclizations and is due to an increase of steric strain in
the tetrahedral intermediates. The effect on the rate is typically
5–100-fold: examples involving simple hydantoinamides are
9
ˇ
provided by the work of Sterba and coworkers. Methylation
**** Accordingtothecrystalstructureofahydantoinderivative18 thearyl
group is twisted about 70Њ out of plane. Molecular mechanics on the
tetrahedral intermediate in a related system yield a much smaller value
of ca. 25Њ (unpublished results).
of the amide group slows down alkaline cyclization five-fold
for 3-methylhydantoinamide and 15-fold for 3,5-dimethyl-
hydantoinamide. These are very similar to the ratios kUAm/kMUAm
listed in Table 5 for the acid and water catalyzed cyclizations.
Bearing in mind the well-established assumption in physical
organic chemistry17 that steric effects are closely similar in the
acid and base catalyzed hydrolysis of carboxylic acid deriv-
atives, the drastically greater effect for kOH cannot be explained
in terms of additional steric strain within T؊. We consider,
as suggested previously for the corresponding reactions of
hydantoic acid esters, that steric hindrance to proton transfer to
the leaving group is the basis of the unusual behaviour of the
Acknowledgements
We thank the Bulgarian Academy of Sciences for support and
the Royal Society for travel funds and Dr I. B. Blagoeva for
helpful discussions.
References
1 M. I. Page and W. P. Jencks, Proc. Natl. Acad. Sci. USA, 1971, 68,
1678; M. I. Page, Chem. Soc. Rev., 1973, 2, 295.
§§ For the N-methyl analogues we estimated previously7 a pK of
8.7, using the procedure of Fox and Jencks.16a The reduction of 0.3
is based on Taylor’s16b compilation of pK-values of amines with
geminal substituents. The difference between NH2 and NHMe averages
to 0.31.
2 C. K. Ingold, J. Chem. Soc., 1921, 119, 305; C. K. Ingold, S. Sako
and J. F. Thorpe, J. Chem. Soc., 1922, 1117; N. L. Allinger and
V. Zalkow, J. Org. Chem., 1960, 25, 701; I. B. Blagoeva, B. J. Kurtev
and I. G. Pojarlieff, J. Chem. Soc. Perkin Trans. 2, 1979, 1115;
A. J. Kirby, Adv. Phys. Org. Chem., 1980, 17, 183; R. E. Valter, Usp.
Khim., 1982, 51, 1374; L. Mandolini, Adv. Phys. Org. Chem., 1986,
22, 17; S. P. Verevkin, M. Kümmerlin, H.-D. Beckhaus, C. Galli and
C. Rüchardt, Eur. J. Org. Chem., 1998, 579.
¶¶ Due to the change of viscosity with temperature.
|||| Depending on the life-time of the intermediate another possibility is
a concerted mechanism but this particular proton transfer step goes
however against Jencks’s libido rule.
O r g . B i o m o l . C h e m . , 2 0 0 3 , 1, 8 5 9 – 8 6 5
864

View Article Online
3 E. Atay, I. B. Blagoeva, A. J. Kirby and Ivan G. Pojarlieff, J. Chem.
Soc., Perkin Trans. 2, 1998, 2289.
4 E. Atay, I. B. Blagoeva and I. G. Pojarlieff, Dokl. Bulg. Akad. Nauk.,
2000, 53(2), 61.
10 I. B. Blagoeva, I. G. Pojarlieff and A. J. Kirby, J. Chem. Soc. Perkin
Trans. 2, 1984, 745.
11 A. Thibbin and W. P. Jencks, J. Am. Chem. Soc., 1979, 101,
4963.
5 I. B. Blagoeva, A. J. Kirby, A. H. Koedjikov and I. G. Pojarlieff, Can.
J. Chem., 1999, 77, 849.
6 I. B. Blagoeva, A. J. Kirby, I. I. Kochiashki, A. H. Koedjikov,
I. G. Pojarlieff and M. M. Toteva, J. Chem. Soc., Perkin Trans. 2,
2000, 1953.
7 A. H. Koedjikov, I. B. Blagoeva, I. G. Pojarlieff and A. J. Kirby,
J. Chem. Soc., Perkin Trans. 2, 1996, 2479, and references quoted
therein.
12 A. H. Koedjikov, P. M. Ivanov and I. G. Pojarlieff, http://
www.arkat.org/arkat/journal/Tee/content.htm.
13 I. B. Blagoeva, J. Chem. Soc. Perkin Trans. 2, 1987, 127.
14 M. Page and A. Williams, Organic & Bioorganic Mechanisms,
Longman, Singapore, 1997, pp. 177–178 and p. 92.
15 W. P. Jencks, Acc. Chem. Res., 1976, 9, 425.
16 (a) J. P. Fox and W. P. Jencks, J. Am. Chem. Soc., 1974, 96, 1436; (b)
P. J. Taylor, J. Chem. Soc., Perkin Trans. 2, 1993, 1423.
17 R. W. Taft, Steric Effects in Organic Chemistry., ed. M. S. Newman,
Wiley, New York, 1956, p. 556.
8 M. Page and A. Williams, Organic & Bioorganic Mechanisms,
Longman, Singapore, 1997, p. 175.
ˇ
9 V. Machácek, G. Svobodová and V. Sterba, Collect. Czech. Chem.
18 E. Schaumann, S. Grabley and G. Adiwidjaja, Liebigs Ann. Chem.,
1981, 264.
Commun., 1987, 52, 140.
O r g . B i o m o l . C h e m . , 2 0 0 3 , 1, 8 5 9 – 8 6 5
865