Kinetics of the Deamination of Amides by Nitrous Acid

Article abstract of DOI:10.1039/a804600j

The kinetic profile of the rate constant for the nitrous acid-amide reaction in sulfuric acid as a function of acidity for a range of aliphatic and aromatic primary amides has been interpreted in terms of the HNO₂/NO⁺ and the amide/ amide · H⁺ equilibria.

Full text of DOI:10.1039/a804600j

View Article Online / Journal Homepage / Table of Contents for this issue
670 J. CHEM. RESEARCH (S), 1998
J. Chem. Research (S),
1998, 670±671$
Kinetics of the Deamination of Amides by Nitrous
Acid$
Khawla Al-Mallah and Geoffrey Stedman*
Chemistry Department, University of Wales Swansea, Singleton Park, Swansea SA2 8PP, UK
The kinetic profile of the rate constant for the nitrous acid±amide reaction in sulfuric acid as a function of acidity for
‡
a range of aliphatic and aromatic primary amides has been interpreted in terms of the HNO2/NO and the amide/
‡
amide Á H equilibria.
The mechanism of the deamination of primary amides by
nitrous acid, eqn. (1), was originally thought1±3 to involve a
rate-determining N-nitrosation reaction, as had been estab-
lished4 for primary amines. Primary aliphatic and aromatic
amides are much weaker bases than the corresponding
primary amines, and in order to get the reaction to proceed
at a reasonable rate much higher concentrations of mineral
acid are necessary. In these concentrated solutions activity
e€ects are important. Attempts to account quantitatively
for the pro®le of rate constant versus acid concentration in
terms of a single mechanism1 were unsuccessful. Subsequent
work by a number of groups5 has established that with the
presence of a strongly electron attracting group such as the
carbonyl function the initial nitrosation is reversible, and
the rate-determining stage involves proton loss. Furthermore
the site of the initial nitrosation may be the oxygen of the
carbonyl group; this is certainly the site for protonation.
The mechanism is summarised in eqns. (2)±(6).
Fig. 1 Variation of k2 with sulfuric acid concentration for
trimethylacetamide at 0 8C
the following results at the molar concentrations of sulfuric
acid speci®ed in parentheses, E/kJ mol 1: R ˆ (CH3)3C,
74.221.4 (4.5); 76.924.8 (9.3); R ˆ CNCH2, 60.124.4
(6.8), 64.022.4 (8.0). The activation energies are constant
to within one standard deviation. The results for trimethyla-
cetamide were used to correct the rate pro®le data measured
at 0 to 25 8C, in order to compare the results with those
for the other amides studied. The rate maximum is due to
the opposing e€ects of acidity on equilibria (2) and (3),
combined with the e€ect of sulfuric acid concentration on
KE and k3. Equilibrium (2) is described formally by the HR
acidity function for which DHR/D[H2SO4] is on average
ca. 1.05 dm3 mol 1. The pKNO for (2) determined spectro-
RCONH2 ‡ HNO2 4 RCO2H ‡ N2 ‡ H2O
ꢀ1†
ꢀ2†
ꢀ3†
ꢀ4†
‡
‡
NO ‡ H2O „ HNO2 ‡ H
KNO
‡
‡
RCONH2H „ RCONH2 ‡ H
KA
‡
‡
RCONH2 ‡ NO „ RCONH2 Á NO
KE
‡
‡
RCONH2 Á NO ‡ B 4 RCONHNO ‡ BH
k3 slow ꢀ5†
ꢀ6†
RCONHNO 4 RCO2H ‡ N2 fast
We summarise in this paper the results of kinetic studies
on a group of aliphatic and aromatic amides reacting with
photometrically using the HR function is
3
8.11. At
HR= 7.59, and thus at higher
3
,
[H2SO4] ˆ 8 mol dm
nitrous acid over the range [H2SO4] ˆ 2.8 to 10.5 mol dm
‡
acidities conversion of nitrite into NO is virtually com-
plete. Equilibrium (3) is described by the HA acidity func-
at 25 8C. Reactions were run with a large excess of amide
over nitrous acid, and gave good ®rst order plots of
ln[nitrite] versus time, k1/min 1. Values of k1 were directly
proportional to [amide], yielding second order rate con-
1
tion for which DHA/D[H2SO4] is only ca. 0.28 dm3 mol
.
‡
Thus at lower acidities the increase in [NO ] dominates over
the decrease in [RCONH2], while as we approach the con-
ditions where HR becomes close to pKNO the fraction of
1
stants k2/dm3 mol 1 min in terms of stoichiometric concen-
trations.
‡
NO , hR/(KNO‡hR), levels o€ and tends to 1, while the
d‰nitriteŠ=dt ˆ k2‰amideŠ‰nitriteŠ
ꢀ7†
fraction of free amide, KA/(KA‡hA), continues to decrease.
However the decrease in k2 is very much faster than the
decrease in 1/hA and the rate of proton loss is a sensitive
function of sulfuric acid concentration. This is presumably
because of variation in the activity of the various proton
All amides studied showed a similar dependence of k2 on
[H2SO4]; k2 increased with acidity and passed through a
3
sharp maximum close to 8 mol dm sulfuric acid, and then
2
decreased sharply. A typical plot is shown in Fig. 1. The
decrease in k2 at higher acidities was always markedly
steeper than the increase at lower acid concentrations. In
order to check that there was no di€erence in mechanism
for reaction at acidities on each side of the rate maximum,
Arrhenius activation energies were measured on each side
of the maximum for two of the aliphatic amides, giving
acceptors such as SO4 , H2O and possibly HSO4 . In
seeking for an empirical measure of this e€ect we used
the H0 acidity function, with the idea that as the proton
donating power measured by h0 increased so the ability
of the medium to remove a proton would decrease. The
mechanism in eqns. (2)±(6) requires the rate law (8).
‡
rate ˆ k3KE‰NO Š‰RCONH2Š
ꢀ8†
*To receive any correspondence.
To relate this to eqn. (2) we write a ˆ the fraction of
$This is a Short Paper as de®ned in the Instructions for Authors,
Section 5.0 [see J. Chem. Research (S), 1998, Issue 1]; there is there-
fore no corresponding material in J. Chem. Research (M).
‡
nitrite present as NO , and b ˆ the fraction of amide
present as RCONH2, where a and b have been de®ned in

View Article Online
J. CHEM. RESEARCH (S), 1998 671
Table 1 Parameters used for plots of eqn. (9)
[H2SO4]/
and there are deviations from eqn. (9); values of log (k2/ab)
are greater than predicted from the equations in Table 1 for
the aliphatic amides. The measurements for the aromatic
amides were not extended to such high acidities. It is poss-
ible that an extra pathway comes into play at high acidities.
3
R
Slope
Intercept
mol dm
pKA
1.40
(CH3)3C*
ClCH2
CNCH2
1.2120.02 7.2220.08 2.8 to 9.3
1.0120.03 5.8420.12 2.8 to 9.3
0.9220.04 5.2120.16 3.4 to 9.1
1.0320.06 5.1220.23 2.8 to 9.3
1.2220.04 7.4020.16 5.5 to 10.4
1.2920.07 7.4720.27 5.5 to 10.5
1.3220.04 7.5420.13 3.2 to 10.5
2.74
3.73
4.18
1.67
2.02
1.97
1.54
2.70
Finally we turn to
a consideration of the di€ering
reactivities of the amides studied. Since there are variations
in the slopes of the plots of eqn. (9) it is simplest to make
comparisons at a given sulfuric acid concentration by substi-
tuting the appropriate value of H0. The same reactivity
sequence is obtained over our range of acidities. For the
aliphatic amides it is (CH3)3C>ClCH2>CNCH2>Cl2CH,
which is the same as the sequence of pKA values. For the
aromatic amides the sequence is p-CH3O>p-CH3>p-Br>
p-Cl>p-NO2, again the same sequence as the pKA values,
except for the interchange of positions of chlorine and
bromine which are very close together anyway. Clearly the
greater the electron releasing power of R in RCONH2
Cl2CH
p-MeC6H4
p-BrC6H4
p-ClC6H4
p-MeOC6H4 1.2820.02 7.9020.07 3.1 to 9.1
p-NO2C6H4 1.2520.05 6.9920.19 5.5 to 10.2
*Corrected from 0 to 25 8C.
terms of KNO, hR, KA and hA in the previous paragraph.
To allow for the variation of k3KE with 1/h0 we write
k3KE=k4/h0. With these substitutions, and the assumption
that there is only a small fraction of the amide present as
‡
RCONH2NO , we obtain eqn. (9).
‡
the higher is the reactivity. The addition of NO to the free
amide may be expected to vary with R in the same sense as
logꢀk2=ab† ˆ log k4 ‡ H0
ꢀ9†
‡
the addition of H , and this is undoubtedly favoured by
Plots of this type yielded good straight lines with slopes
between 0.92 and 1.27 over the range of acidities speci®ed in
Table 1. Values of pKA for aliphatic amides are due to
Liler6 and for aromatic amides come from the review7 by
Boyd. The value for (CH3)3CONH2 comes from our own
work8 based upon NMR measurements, and is very close
to Liler's value. Although our model gives a reasonably
satisfactory description of the data the choice of H0 was
arbitrary. However most acidity functions vary in an
approximately linear manner with [H2SO4] in the higher
range of acid concentrations, and so the choice of other
acidity functions to describe the variation of KEk3 would
still have yielded linear plots, though with di€erent slopes.
Since of the use of 1/h0 is a purely empirical measure of the
ability of the medium to accept a proton we do not o€er
any comment on the variation of the slopes. There must be
considerable di€erences in the solvation of the aliphatic
amides in view of the di€erences in R; for the aromatic
amides which all have a phenyl ring the slopes are quite
similar. An alternative explanation of the dependence of
k2/ab upon 1/h0 would be an acid±base equilibrium for (5)
followed by a rate-determining rearrangement as in eqn. (6).
However as Williams5 points out the evidence for a rate-
determining loss of a proton is very strong, so we reject this
alternative.
electron release by R. Step (5) however involving proton
loss will be reduced in rate by increase in electron release by
R. We conclude therefore that the reactivity in the reaction
of nitrous acid with primary amides is controlled by the
basicity of the amide, by favouring the formation of an
‡
amide Á NO species, and that this is more important than
the e€ect on the rate of proton loss.
Experimental
Materials.ÐThe amides used were either commercially available
materials (Aldrich, BDH, Koch-Light) or were prepared from the
acid chlorides by addition to stirred, ice-cold 0.88 ammonia. In
some cases the acid chlorides were prepared by reaction of thionyl
chloride with the carboxylic acid. The amides were recrystallised
from water to constant melting point.
Kinetic Methods.ÐSome reactions were followed by colorimetric
analysis for nitrous acid involving diazotisation of sulfanilic acid
and coupling with alkaline 2-hydroxynaphthol-3,6-disulfonic acid.
Other reactions were followed by direct UV spectrophotometry
at wavelengths where nitrite absorbed. For slow runs a blank exper-
iment was carried out to correct for the self-decomposition of
nitrous acid.
Thanks are due to the University of Mosul for study
leave (to K. A.-M.)
Although the present model gives a reasonable description
of the data in terms of the generally accepted mechanism,
it is important to realise its limitations. One concerns
Received, 17th June 1998; Accepted, 10th July 1998
Paper E/8/04600J
the value used for pKNO
. Various values have been
suggested,9,10 and in his interpretation of the kinetics of
diazotisation in dilute acid Ridd used11 a value of ca. 6.5,
which yields limiting rates in good agreement with the calcu-
References
1 M. N. Hughes and G. Stedman, J. Chem. Soc., 1964, 5840.
2 M. L. Bender and H. Ladenheim, J. Am. Chem. Soc., 1960, 82,
1895.
3 J. Jaz and A. Bruylants, Bull. Soc. Chim. Belges, 1961, 70, 99.
4 J. H. Ridd, Q. Rev. Chem. Soc., 1961, 15, 418.
5 D. L. H. Williams, in Nitrosation, Cambridge University Press,
Cambridge, 1988, p. 101.
‡
lated value of the encounter rate between NO and ArNH2.
As we are concerned with data over a wide range of sulfuric
acid concentration, using the HR data to describe (2) we
prefer to use the more negative ®gure of 8.11 deduced by
Deno et al.12 when ®tting spectrophotometric estimates of
6 M. Liler, J. Chem. Soc. B, 1969, 385.
‡
[NO ]/[HNO2] to his HR data obtained by the ionisation of
7 J. H. Boyd, in Solute±Solvent Interactions, ed. J. F. Coetzee and
C. D. Ritchie, Marcel Dekker, New York, 1969, vol. 1, ch. 3,
p. 97.
8 K. Al-Mallah, Ph.D. Thesis, University of Wales, 1974.
9 Stability Constants, Special Publication No. 17, The Chemical
Society, London, 1964, p. 163.
10 Stability Constants, Special Publication No. 25, The Chemical
Society, London, 1971, p. 91.
11 J. H. Ridd, Adv. Phys. Org. Chem., 1978, 16, 1.
12 N. C. Deno, H. E. Berkheimer, W. L. Evans and H. J. Peterson,
J. Am. Chem. Soc., 1959, 81, 2344.
arylcarbinol indicators. At lower [H2SO4] changing the
value of pKNO merely displaces the line of log (k2/ab) versus
H0 without signi®cantly changing the slope. As HR
approaches pKNO the choice is more important. We tried
using pKNO
concentrations greater than 6.5 mol dm
=
6.5, but found deviations at sulfuric acid
3
,
whereas with
Deno's value linearity was observed up to the speci®ed con-
centrations shown in Table 1. We have also extended a few
3
measurements to higher acidities, [H2SO4] ˆ 13.1 mol dm
,