Oxidation of Ascorbic Acid in the Presence of Nitrites

Article abstract of DOI:10.1021/jf940643w

We have shown that ascorbic acid in neutral aqueous media rather rapidly interacts with nitrite ions to form dehydroascorbic acid. If the acidity of solution is sufficiently high (pH value does not exceed much above ～7) and the nitrite concentration is of order of 0.05 M or more, the only mechanism of the reaction under aerobic conditions includes, as the first, rate-limiting reaction step, the nitrosation of ascorbic acid. A second-order reaction kinetics was observed in respect to protons, which determined a sharp acceleration of the ascorbic acid conversion when the acidity of the medium was raised. Hence, a suggestion can be drawn that the presence of nitrites in vegetables under acidic storage strongly increases the decay of vitamin C.

Full text of DOI:10.1021/jf940643w

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2948
J. Agric. Food Chem. 1996, 44, 2948−2950
ARTICLES
Oxid a tion of Ascor bic Acid in th e P r esen ce of Nitr ites
Anatoli E. Myshkin,* Vera S. Konyaeva, Klara Z. Gumargalieva, and Yuri V. Moiseev
N. N. Semenov Institute of Chemical Physics of the Russian Academy of Sciences (IKhF RAN),
Kosygin Street 4, 117977 Moscow V-334, Russia
We have shown that ascorbic acid in neutral aqueous media rather rapidly interacts with nitrite
ions to form dehydroascorbic acid. If the acidity of solution is sufficiently high (pH value does not
exceed much above ∼7) and the nitrite concentration is of order of 0.05 M or more, the only
mechanism of the reaction under aerobic conditions includes, as the first, rate-limiting reaction
step, the nitrosation of ascorbic acid. A second-order reaction kinetics was observed in respect to
protons, which determined a sharp acceleration of the ascorbic acid conversion when the acidity of
the medium was raised. Hence, a suggestion can be drawn that the presence of nitrites in vegetables
under acidic storage strongly increases the decay of vitamin C.
Keyw or d s: Ascorbic acid; nitrite; oxidation; nitrosation
INTRODUCTION
Sch em e 1
ON
OH
O•
OH
O•
O•
O
O
NO+
NO•
One of the most important problems in food ecology
is the interaction of toxicants with vitamins. A fairly
high content of nitrites in some vegetables raises a
problem of possible negative influence of the nitrite ion
on vitamins, primarily on vitamin C (ascorbic acid, AA).
A discussion of this problem is urgent, taking into
account the well-known ability of AA to undergo a rapid
nitrosation in acidic media, O-nitrosyl-AA thereby being
formed, producing unstable dehydroascorbic acid as the
result of several subsequent rapid conversion steps
(Dahn et al., 1960). The consecutive conversions start-
ing from O-nitrosyl-AA is given in Scheme 1. As
information on similar reactions in neutral media has
not yet been published, we investigated the interaction
of L-AA with nitrite ions in neutral and low-acidic media
buffered with Tris-HCl and phosphate buffers.
–NO•
of the AA destruction was calculated as a reverse value of the
reaction half-time. The final values of the above reaction rates
were further calculated as an arithmetical mean of three runs.
The experimental error did not exceed 5%.
In the next section, we shall also discuss some experiments
beyond the main subject of the paper but of interest for the
overall ecological problem.
RESULTS AND DISCUSSION
Despite complete binding of the trace heavy metal
ions with EDTA, their catalytic activity for the AA
autoxidation remained high enough. Because of this,
the effect of nitrite addition to the AA solutions might
be considered both as the result of the AA oxidation
through a phase of O-nitrosation of AA and its autoxi-
dation due to an increase in the solution ion strength
or the formation of an AA-nitrite ion complex (Gumar-
galieva et al., 1989). An answer to this question, as the
first approximation, might be obtained by means of the
pH value variation. In such a way, we have established
the period of the AA half-conversion (t1/2) in 0.05 M Tris-
HCl buffer in the absence of nitrite at pH 8.8 to be 58
min, at pH 7.3, 90 min, and at pH 6.7, 240 min. Such
regularity fully corresponds to a decrease, with reducing
pH value, in the concentration of the monoanionic form
of AA (HA-), which is more disposed to oxidation by
molecular oxygen (Khan and Shukla, 1986). However,
in the presence of 0.2 M NaNO2, the t1/2 value changes
in a reverse order and is 30, 15, and 8 min, respectively.
It means that, at pH 8.8, nitrite accelerates the reaction
by 2 times whereas, at pH 6.7, the acceleration is 150
times. It is clear that the mechanism of the nitrite effect
in the two cases should be essentially different. At pH
8.8, the acceleration of AA autoxidation upon the
addition of nitrite may be considered the result of a
nonspecific salt effect, while, at pH 6.7, the prevailing
nitrite effect cannot be anything other than the result
of an active specific interaction of nitrite with AA. This
EXPERIMENTAL PROCEDURES
L-Ascorbic acid was of “chemically pure” grade (Russian
classification) and was used without further purification.
Sodium nitrite, EDTA disodium salt (Trylon B), and phosphate
buffer components were of the same quality. Tris(hydroxy-
methyl)aminomethane (Tris) was of an inferior purity and was
used after the preliminary recrystallization from ethanol.
The rate of the AA conversion was measured at a temper-
ature of 25 °C spectrophotometrically at a wavelength of 267
nm, where the reaction product did not absorb. The measure-
ments were carried out on a SF-26 spectrophotometer equipped
with a thermostating device. To reduce AA autoxidation
induced by the traces of heavy metal ions (for instance, the
content of these ions in disodium phosphate is up to 0.001%,
and the content of lead in sodium nitrite up to 0.0005%), the
activity of the latter was lowered by addition of EDTA. In
each experiment, 17.5 µL of aqueous 5.2 × 10-3 M L-AA was
added to 3 mL of thermostated 0.05-0.2 M sodium nitrite
solution in 0.05 M Tris-HCl buffer, pH 6.7-7.3, or 0.067 M
sodium phosphate buffer, pH 5.7-7.1 (all the pH values
measured on a Soviet pH-340 apparatus at 25 °C with an
accuracy 0.1 pH unit) with 5 × 10-5 M EDTA disodium salt.
The concentration of dilute AA was 3 × 10-5 M. After addition
of the AA solution to a buffer system, the drop in optical
absorption at 267 nm was measured. In most cases, the rate
S0021-8561(94)00643-6 CCC: $12.00
© 1996 American Chemical Society

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Ascorbic Acid Nitrite Oxidation
J. Agric. Food Chem., Vol. 44, No. 10, 1996 2949
F igu r e 1. pH dependence of the logarithm of AA oxidation
rate (calculated as the reverse of the reaction half-time) with
0.2 M NaNO2 and 5 × 10-5 M Trylon B at 25 °C in (1) 0.05 M
Tris-HCl and (2) 0.067 M phosphate buffer. Initial AA con-
centration, 3 × 10-5 M.
F igu r e 2. Nitrosation-oxidation of AA in 0.05 M Tris-HCl
buffer, pH 6.7, at 25 °C in the presence of varied amounts of
NaNO2: [AA]0 ) 3 × 10-5 M, [Trylon B] ) 5 × 10-5 M, µ ) 0.2
M (NaNO2 + KClO4). (a) Kinetic curves: (1) 0.05, (2) 0.075,
(3) 0.10, (4) 0.15, and (5) 0.2 M NaNO2. (b) Logarithmic
dependence of the AA oxidation rate on the nitrite concentra-
tion.
effect, in turn, is most probably O-nitrosation of AA
followed by subsequent fast oxidation (Dahn et al.,
1960). In such a case, at pH 7.3, the two types of AA
oxidation would be competitive with each other.
We confirmed the above-made assumption in the
following way. After blowing off the solution with argon
to full deaeration, AA oxidation in the presence of 0.2
M NaNO2 at pH 8.8 was virtually nonexistent; pH 6.7,
blowing off O2 did not affect the oxidation rate; and at
pH 7.3, the observable rate of AA oxidation decreased
by half. A more detailed experiment showed that, in
0.05 M Tris buffer and with NaNO2 concentrations not
lower than 0.05 M, pH 7.3 was really a boundary: at
lower pH values, the AA oxidation through the nitro-
sation step was dominating whereas, at pH higher than
7.3, the contribution of the nitrosation-oxidation mech-
anism was small and became negligible with further
increase in pH.
Information about AA autoxidation in neutral media
is widely available (Khan and Martell, 1968; Hsieh and
Harris, 1987; Yatsimirskii and Labuda, 1984). Due to
this fact, the main subject of the discussion in the
current paper will be that of data on AA oxidation in
the presence of NaNO2 at a pH lower than 7.3, that is,
in the pH value region where the nitrosation-oxidation
is a dominating pathway of the AA degradation.
Variation of the concentrations of AA, sodium nitrite,
and protons within the pH range 6.7-7.3 made it
possible to derive an approximate equation for the AA
nitrosation-oxidation rate in Tris buffer under low-
acidic pH values:
function of [H+]2 convincingly explains a rather sharp
transition from the AA autoxidation catalyzed by nitrite
to the nitrosation-oxidation of AA when passing from
pH ∼7.3 to lower pH values. Indeed, in the pH region
near pH 7.3, the rate of AA autoxidation undergoes a
rather considerable drop with lowering pH value (by
∼2.5 times under the pH transition from 7.3 to 6.7),
whereas the AA nitrosation-oxidation reaction is sharply
accelerated.
A nitrite order of 1.25 is also of interest. In contrast
to the proton order, whose magnitude steadily keeps
close to 2, the nitrite order changes quite noticeably,
depending on the medium composition. Particularly, at
different pH values in the phosphate buffer , the nitrite
order is close to 1.4-1.6. However, in both phosphate
and Tris-HCl buffers, the nitrite order does not reach a
magnitude of 2, which has been found in more acidic
media (Dahn, 1960). Figure 2a represents a number of
kinetic curves for the nitrosation-oxidation of AA in
Tris buffer, pH 6.7 with varying NaNO2 concentrations.
The logarithmic dependence of the v1/2 on the nitrite
concentration (Figure 1b) allowed us to calculate the
above-mentioned reaction order (precisely 1.23 ( 0.009)
for nitrite ions. The most probable nitrosating agent-
nitrous anhydride-however, formed according to reac-
tions 2 and 3 (Ridd, 1961), requires the nitrite order to
v ) d[AA]0.3[NO2
]
- 1.25[H+]2
(1)
HNO2 + HNO2 h N2O3 + H2O
H2NO2+ + NO2- h N2O3 + H2O
(2)
(3)
The most important practical conclusion of the above
equation is that the rate of AA decomposition in the
presence of nitrite rises as the square of the proton
concentration. It is well demonstrated by the depen-
dence of log v1/2 ) 103/t1/2 (min-1) upon the pH value as
shown in Figure 1 (curve 1). The slope of this depen-
dence is 2.16 ( 0.031, which approximates well enough
to 2. A similar quantity obtained in phosphate buffer
within the pH range 5.1-7.1 (Figure 1, curve 2) is 1.79
( 0.04, which to some extent corroborates the previous
result. Dependence of the rate of AA decomposition as
be 2. The nitrite first order might imply the nitrosating
agent to be H2NO2+, whose formation rate depends on
[H+] and [NO2-]. Calculations were made, however,
that showed the participation of H2NO2+ in the limiting
reaction step to be impossible for reasons both of very
low equilibrium constant of the HNO2 protonation and
of very high rate constant of the same reaction. Besides,

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2950 J. Agric. Food Chem., Vol. 44, No. 10, 1996
Myshkin et al.
the proton second order suggests the nitrosating agent
to be nitrous anhydride formed in reaction 2. So we
cannot explain a low order for nitrite by simple kinetic
reasons.
A combination of factors of stability and reactivity of
NO+B is capable of considerably changing the rate of
AA nitrosation.
The mechanism of AA nitrosation is more com-
prehensively discussed in our previous publications
(Myshkin et al., 1990, 1991; Konyaeva et al., 1992).
A rather low AA reaction order indicates that forma-
tion of a nitrosating agent is mostly the limiting step.
Clearly, there should also be a concurrent reaction
pathway in which AA is involved in the limiting step.
The kinetic data obtained make it possible to draw
some practical conclusions as to a possible influence of
nitrites on vegetables under storage conditions. Per-
form a rough calculation. According to our data, 0.83
mg/L AA in the presence of 9.2 g/L NO2- ions at pH 6.7
(Tris-HCl buffer) and a temperature of 25 °C decom-
poses by half in 1 min. Decrease now the concentration
LITERATURE CITED
Dahn, H.; Loewe, L.; Bunton, C. A. Uber die Oxidation von
Askorbinsaure durch Salpetrigsaure. Teil VI: Ubersicht und
Diskussion der Ergebnisse. Helv. Chim. Acta 1960, 43, 320-
333.
Gumargalieva, K. Z.; Davydov, R.; Kalinina, I. G.; Moiseev,
Yu. V. Nitrite and decomposition of ascorbic acid. Zh. Fiz.
Khim. 1989, 63, 2539-2542.
Hsieh, I. P.; Harris, N. D. Oxidation of ascorbic acid by nitrite
ion in copper-catalyzed sucrose solution. J . Food Sci. 1987,
52, 1384-1388.
Konyaeva, V. S.; Myshkin, A. E.; Gumargalieva, K. Z.; Moiseev,
Yu. V. Nitrosation of ascorbic acid by nitrite ion ion in non-
buffered neutral and moderately acid aqueous media. Izv.
Akad. Nauk SSSR, Ser. Khim. 1992, (7), 1540-1544.
Myshkin, A. E.; Konyaeva, V. S.; Gumargalieva, K. Z.; Moiseev,
Yu. V. Principal mechanisms of the oxidation of ascorbic acid
in presence of nitrite in non-acid aerobic aqueous media.
Dokl. Akad. Nauk SSSR 1990, 315, 383-387.
Myshkin, A. E.; Konyaeva, V. S.; Gumargalieva, K. Z.; Moiseev,
Yu. V. Mechanism of nitrosation of ascorbic acid by nitrite
in neutral aqueous media. Izv. Akad. Nauk SSSR, Ser.
Khim. 1991, (10), 2242-2247.
Read, J . H. Nitrosation, diazotisation, and deamination. Q.
Rev. 1961, 15, 418-441.
Taqui Khan, M. M.; Martell, A. E. Kinetics of metal ion and
metal chelate catalyzed oxidation of ascorbic acid. III.
Vanadyl ion catalyzed oxidation. J . Am. Chem. Soc. 1968,
90, 6011-6017.
-
of NO2 ions to a reasonable practical value, for in-
stance, to 10 mg/L. In this case, if nitrite order is equal
to 1.25, the half-decay time of AA rises by (9.2 × 103/
10)1.25 times and reaches ∼5000 min, that is, close to
85 h. However, at rather slight acidification of the
medium, even to pH 6.0, the AA half-decay time
decreases by (10-6/2 × 10-7 2
) ) 25 times and now
amounts to ∼3.5 h (at 5 °C ∼25 h). At a higher acidity
inside a vegetable, which is either specially created for
storage or is endogenously developed during the storage,
the time of AA half-decay at 25 °C can be reduced by
tens and hundreds times: 100 times per each pH unit.
Then it is possible to affirm that an elevated content of
nitrate ions in vegetables and therefore correspondingly
elevated content of nitrite ions produce a highly unfa-
vorable medium for vitamin C under low-acid condi-
tions. Naturally, our calculations are rather conven-
tional and do not take into account many other chemical
factors, which are determined by both the initial state
of vegetable tissue and vegetable storage conditions. So
our conclusions should be regarded only on a qualitative
level. And still, as various chemical factors may not
only inhibit the nitrite-induced AA destruction but, in
some cases, also increase it (salt effects etc.), a danger
of elevated AA destruction in the presence of nitrite ions
must be taken into consideration. The following obser-
vation, in particular, could support the last statement:
we have found that different nucleophiles (neutral bases
and anions) may efficiently modify the nitrosating agent
in accordance with eq 4
Taqui Khan, M. M.; Shukla, R. S. Thermodynamics of the
homogeneous oxidation of L-ascorbic acid by molecular
oxygen catalyzed by ruthenium ion and ruthenium chelates.
J . Mol. Catal. 1986, 37, 269-285.
Yatsimirskii, K. B.; Labuda, Ya. Oxidation of ascorbic acid by
molecular oxygen catalyzed by tetrabenzo-tetraazamacro-
cyclic complex of copper. Dokl. Akad. Nauk SSSR 1984, 276,
880-883.
Received for review November 15, 1994. Revised manuscript
received May 30, 1995. Accepted J uly 10, 1996.X This work
was supported from the State budget of the Russian Federation.
H2NO2+ + B h NO+B + H2O
(4)
J F940643W
Reagent NO+B being more stable as compared to H2-
NO2+ may be both less or more reactive than the latter.
X Abstract published in Advance ACS Abstracts, Sep-
tember 1, 1996.