Kinetic Investigation of the Reactions Connected to the System Ascorbate + O2 by Amperometric Detection of H2O2 at a Modified Platinum Electrode

Article abstract of DOI:10.1021/ac970347t

A Pt electrode modified by an electrochemically produced bilayer polymeric membrane [polypyrrole/poly(o-phenylenediamine)] entrapping the enzyme glucose oxidase proved able to detect (response time of few seconds) amperometrically (at +0.7 V vs Ag/AgCl) very low concentrations of hydrogen peroxide (micromolar range) in the presence of much higher amounts of ascorbate. The currents due to ascorbate, also electroactive at the given potential, were negligible under any conditions due to its almost complete rejection by the electrode-modifying membrane system. The very peculiar properties of the device setup were exploited to undertake a kinetic study of the reactions connected to the system ascorbate + 0₂, following the concentration of H₂O₂ produced in the reaction mixture at 27 °C, pH = 7. The reaction between ascorbate and H₂O₂ was also considered; however, different kinetic models based on the two consecutive reactions proved unable to fit the data. An investigation on the single processes by the same experimental approach was then undertaken, leading to two explanations for the inadequacy of simple kinetic models. First, the presence of metal ion traces in the reaction mixture proved to be responsible for the nonlinear dependence of the rate of both reactions on the ascorbate concentration: a mechanism involving the role of ascorbate-metal complexes as the reactants was hypothesized to explain this result. Second, the influence of the reactivity of dehydroascorbic acid, the product of ascorbate oxidation, on the kinetics was ascertained.

Full text of DOI:10.1021/ac970347t

Anal. Chem. 1997, 69, 4113-4119
Kinetic Investigation of the Reactions Connected
to the System Ascorbate + O2 by Amperometric
Detection of H2O2 at a Modified Platinum Electrode
Carlo G. Zambonin* and Ilario Losito
Dipartimento di Chimica, Universita` degli Studi, Via E. Orabona 4, I-70126 Bari, Italy
has been the object of many kinetic investigations, due to its basic
biochemical and practical significance. Most of the work done
on this reaction has been recently reviewed.1 The kinetic
determinations relevant to AA autoxidation are usually carried out
in the presence of known amounts of transition metal ions, e.g.,
Cu(II), Fe(III), Ru(III), etc., or their complexes, acting as catalysts.
This procedure masks the small catalytic effects of metal ions
impurities always present at variable and difficult-to-control
concentration levels.
A P t electrode modified by an electrochemically produced
bilayer polymeric membrane [polypyrrole/ poly(o-phen-
ylenediamine)] entrapping the enzyme glucose oxidase
proved able to detect (response time of few seconds)
amperometrically (at +0 .7 V vs Ag/ AgCl) very low con-
centrations of hydrogen peroxide (micromolar range) in
the presence of much higher amounts of ascorbate. The
currents due to ascorbate, also electroactive at the given
potential, were negligible under any conditions due to its
almost complete rejection by the electrode-modifying
membrane system. The very peculiar properties of the
device setup were exploited to undertake a kinetic study
of the reactions connected to the system ascorbate + O2 ,
following the concentration of H2 O2 produced in the
reaction mixture at 2 7 °C, pH ) 7 . The reaction between
ascorbate and H2 O2 was also considered; however, dif-
ferent kinetic models based on the two consecutive
reactions proved unable to fit the data. An investigation
on the single processes by the same experimental ap-
proach was then undertaken, leading to two explanations
for the inadequacy of simple kinetic models. First, the
presence of metal ion traces in the reaction mixture
proved to be responsible for the nonlinear dependence
of the rate of both reactions on the ascorbate concentra-
tion: a mechanism involving the role of ascorbate-metal
complexes as the reactants was hypothesized to explain
this result. Second, the influence of the reactivity of
dehydroascorbic acid, the product of ascorbate oxidation,
on the kinetics was ascertained.
Furthermore, any kinetic study of reaction 1 should also take
into account the consecutive oxidation process2 between AA and
H2O2:
CH2OH
CHOH
CH2OH
CHOH
HC
HC
OH
O
(2)
O
C
O
C
+ H2O2
+ 2H2O
C
C
C
C
OH
AA
O
O
O
DAA
In the course of the present study, an attempt was made to collect,
during the same experiment, information on processes 1 and 2
by amperometric detection of the concentration of H2O2. A Pt
electrode modified by an electrosynthesized bilayer polymeric
membrane, i.e., polypyrrole (PPy) and poly(o-phenylenediamine)
(PPD), entrapping the enzyme glucose oxidase (GOx), was
employed. In fact, this device, Pt/ PPyox/ PPD(GOx), proved3 able
to detect very small concentrations of H2O2 (micromolar), drasti-
cally rejecting the electroactive AA (see also the Experimental
Section).
The approach adopted made it possible to follow continuously
the concentration of hydrogen peroxide, emphasizing the kinetic
role of H2O2 consumption, usually ignored by previous investiga-
tions (see, e.g., ref 4) based on the periodic quantification of either
AA or DAA.
In the present work, reactions 1 and 2 were studied at 27 °C
and pH ) 7 (phosphate buffer solutions), and different mathemati-
cal models were first adopted in order to fit the experimental
[H2O2] vs time curves. Their failure led to a critical kinetic
investigation on the single processes, which gave useful informa-
tion on AA oxidation. In particular, the catalyzing effect due to
The oxidation of ascorbic acid (AA) by dioxygen (also known
as autoxidation), leading to dehydroascorbic acid (DAA) and
hydrogen peroxide according1 to a reaction such as
CH2OH
CHOH
CH2OH
CHOH
HC
HC
OH
O
(1)
O
C
O
C
+ O2
+ H2O2
C
C
C
C
OH
AA
O
O
O
DAA
(1) Davies, M. B. Polyhedron 1 9 9 2 , 11, 285-321 and references cited therein.
(2) Hand, D. B.; Greisen, E. C. J. Am. Chem. Soc. 1 9 4 2 , 64, 358-361.
(3) Losito, I.; Zambonin, C. G. J. Electroanal. Chem. 1 9 9 6 , 410, 181-187.
(4) Taqui Khan, M. M.; Martell, A. E. J. Am. Chem. Soc. 1 9 6 7 , 89, 4176-4185.
* To whom correspondence should be addressed. Present address: Dipar-
timento di Chimica, Universita` della Basilicata, Via N. Sauro, I-85100 Potenza,
Italy.
S0003-2700(97)00347-8 CCC: $14.00 © 1997 American Chemical Society
Analytical Chemistry, Vol. 69, No. 20, October 15, 1997 4113

the presence of metal ions, arising from the buffer (as traces) or
added in known amounts (Cu2+), was assessed, and the influence
of DAA reactivity on the kinetics was examined.
These findings, besides permitting the rationalization of the
inadequacy of relatively simple kinetic models, can be very helpful
in drawing a reliable picture for the first steps of the quite
complicated reacting system AA-O2.
Modified electrodes showing relatively low ratios (<400)
between the current responses obtained for H2O2 and AA
respectively at the same concentration (1 mM) were not used for
kinetic experiments. This choice was made to minimize the effect
due to the AA electrooxidation products, which could alter the
electrode surface,3,5 thus influencing the sensitivity to H2O2 during
a kinetic run (see also the Results and Discussion session).
Kinetic Measurements. A thermostated electrochemical cell
was used as reaction vessel for all the kinetic experiments in this
work. Particular care was taken to prevent any further introduc-
tion of oxygen (from air) in the reacting system after a run started.
This was mandatory for the experiments involving oxygen, in
which the concentration of dissolved O2 was continuously reduced
during the kinetic run.
Some of the experiments were performed by using, as oxygen
initial concentration, the value corresponding to its partial pressure
in the atmosphere (∼0.19 atm). The concentration was evaluated
to be 0.23 mM, in good agreement with the data reported by Taqui
Khan and Martell4 for an aqueous solution with the same ionic
strength as the buffer used in this work.
EXPERIMENTAL SECTION
Materials. GOx from Aspergillus niger (EC 1.1.3.4, 168 200
units/ g) and ascorbic acid were obtained from Sigma Chemical
Co. (St. Louis, MO) and used as received. Hydrogen peroxide
was obtained as a 36% (w/ v) solution from Carlo Erba (Milano,
Italy). o-Phenylenediamine (o-PD) was purchased from Aldrich
(Steinheim, Germany) and purified just before use by vacuum
sublimation at 90 °C. Pyrrole (Aldrich) was purified by vacuum
distillation at 60 °C and stored under nitrogen at 0 °C.
The buffer solutions (pH ) 7; I ) 0.1) for the kinetic
experiments were always prepared by using monobasic and
dibasic sodium phosphates (ACS reagents) purchased from
Aldrich.
The cell was partially filled with phosphate buffer solution (I
) 0.1, pH ) 7), kept in a thermostatic bath at 27 °C, allowing O2
to reach, under stirring, its equilibrium concentration. The three
electrodes (working, counter, and reference) were then inserted,
and O2 saturated buffer was added to fill the reaction vessel. After
the operating potential (0.7 V vs Ag/ AgCl) was applied and the
system allowed to achieve a steady value for the background
current, the cell was protected from light and the kinetic
experiment started by adding a known amount of AA to the
solution (final concentration ranging from 0.1 to 1 mM). The AA
stock solutions (50 mM) used for the injection were prepared just
before each kinetic run and deaerated by flowing a UHP inert
gas, in order to avoid any preliminary reaction with oxygen.
Other experiments were performed adopting lower oxygen
concentrations. In these cases, a different preparation stage was
necessary before starting the kinetic run. After the cell was filled,
an inert gas was bubbled in the solution through the overflow
pipe until the desired O2 concentration was reached. During this
stage, the O2 concentration was followed (and finally calculated)
by amperometric detection at an additional Pt electrode (operating
potential, -0.4 V vs Ag/ AgCl). The chosen AA aliquot was then
injected and the kinetic study started.
In any case, particular care was used to eliminate small air
bubbles inside the cell before each run.
Kinetic Models. Kinetic models based on two pseudo-first-
order (1/ 1) or second-order (2/ 2) consecutive reactions were
adopted, as a first approach, to study the oxidation of AA by
oxygen and by H2O2 in the same system.
In particular, the 2/ 2 model was chosen on the basis of
previous kinetic investigations1,3,4,6 on processes 1 and 2, indicating
a second-order dependence for their rate.
The mathematical treatment already available7 for this model
(with a ratio of 2 between the initial concentrations of the
reactants) was extended to any ratio between the initial concentra-
tions of the reactants. Briefly, as a final result, the following
equation was obtained:
Apparatus. Modified electrode preparation was carried out
by a PAR 273 potentiostat-galvanostat (EG&G, Princeton Applied
Research, Princeton, NJ) coupled to a conventional three-electrode
system with a Pt wire as counter electrode and a Ag/ AgCl/ KCl
(saturated) electrode as reference. The amperometric response
during kinetic experiments was measured by a Model 400 EC
Detector (EG&G) connected to a Y-t strip chart recorder (LCI
100, Perkin-Elmer, Norwalk, CT).
The UV spectra were recorded in the range 210-400 nm using
a Perkin-Elmer Lambda 2 UV/ visible spectrometer.
Modified Electrode P reparation. The preparation of the
modified electrode has been described elsewhere3. Briefly, a
platinum disk (diameter, 3 mm), embedded in a Teflon cylinder,
was used as the working electrode: its surface was polished with
0.3 µm alumina, washed, and pretreated by potential cycling
between -0.21 and +1.19 V vs Ag/ AgCl in 0.5 M H2SO4 until a
steady state voltammogram was obtained. A PPy film was
potentiostatically grown (at +0.7 V vs Ag/ AgCl) from a 10 mM
KCl solution containing pyrrole (0.4 M). The deposition charge
employed was typically 300 mC/ cm2. The Pt/ PPy-modified
electrode was washed and transferred to another cell containing
5 mM o-PD (and 500 units/ mL of GOx, when indicated) in a 0.1
M phosphate buffer (pH ) 7.0) solution. A potential of +0.7 V
vs Ag/ AgCl was applied in order to potentiostatically deposit a
PPD film; the process was stopped when no appreciable current
flowed in the circuit. After preparation and thorough washing,
the modified electrode was potentiostated at +0.7 V overnight in
a 0.1 M phosphate buffer (pH ) 7.0) in order to complete the
PPy overoxidation process.
A calibration test was performed before each kinetic experi-
ment in order to assess the sensitivity of the Pt/ PPyox/ PPD(GOx)
electrodes to H2O2 in the range 2-100 µM, including the
maximum H2O2 concentrations reached during the reactions. A
response time (t0.9) of about 5 s was estimated. The current values
were controlled to be constant with time, which indicated that no
decomposition of H2O2 was occurring under the present experi-
mental conditions. The rejection of ascorbate from the electrode
surface was also checked by measuring the amperometric
response after addition of AA (final concentration, 1 mM) to
phosphate buffer.
(5) Palmisano, F.; Zambonin, P. G. Anal. Chem. 1 9 9 3 , 65, 2690-2692.
(6) Taqui Khan, M. M.; Shukla, R. S. J. Mol. Catal. 1 9 8 7 , 39, 139-146.
(7) Pannetier, G.; Souchay, P. Chemical Kinetics; Elsevier: Amsterdam, 1967;
pp 196-199.
4114 Analytical Chemistry, Vol. 69, No. 20, October 15, 1997

Figure 2. Kinetic analysis of the hydrogen peroxide concentration
profile relevant to the reaction mixture AA (1 mM)-O2 (0.23 mM) in
phosphate buffer (0.1 M, pH ) 7) solution at 27 °C. A model based
on two consecutive second-order reactions was used. The H2O2
concentration profile was obtained from the current response cor-
rected for the small contribution due to ascorbate.
Figure 1. Calibration curves for hydrogen peroxide obtained with
a Pt/PPYox/PPD(GOx) electrode in phosphate buffer (0.1 M, pH ) 7)
solution in the absence (a) and in the presence (b) of ascorbate (1
mM). Each step corresponds to a 2.5 µM increase in H2O2 concentra-
tion. Temperature, 27 °C; applied potential, 0.7 V vs Ag/AgCl
(saturated KCl).
1
1 - r
k [O ] t )
r+1 dâ (3)
∫
1
2 0
â (1 - 2r)â2 + (2 - R)(r - 1)â + â
where k1 is the second-order kinetic constant for reaction 1, r the
ratio between the kinetic constants of reaction 2, k2 and k1, â )
[O2]/ [O2]0, and R ) [AA]0/ [O2]0.
A program was developed in Quick Basic (version 4.0)
environment to perform the simulation of the concentration-time
profiles and fit the experimental curves by means of the variable
step-size simplex algorithm,8 k1 and k2 being the dependent
variables.
Figure 3. Current-time response obtained with a Pt/PPYox
/
PPD(GOx) electrode after addition of (A) ascorbate, final concentra-
tion 30 µM, and (B) H2O2, final concentration 2 µM, to a stirred
phosphate buffer (0.1 M, pH ) 7) solution containing Cu2+ (10 µM).
Temperature, 27 °C; applied potential, 0.7 V vs Ag/AgCl (saturated
KCl). The ratio between the sensitivities to H2O2 and AA at the given
electrode configuration is around 500 (see text). The curve corre-
sponds to the fourth experiment in Table 1 (vide infra).
As far as the 1/ 1 model is concerned, commercial software
was employed to perform the optimization of the constants values.
The best curve calculated by means of the 2/ 2 model (dotted
line) shows that the latter cannot fit the experimental data. The
inadequacy of the model could be due to a difference between
the real kinetic orders of the involved reactions (1 and 3) and
those reported in the literature. Another possible explanation
could be the occurrence of other reactions involving H2O2 and/
or O2 (e.g., with DAA, the first oxidation product of AA).
An investigation on the kinetic orders for the single reactions
was then undertaken; the results are reported in the following
paragraphs.
Investigation on the Kinetic Order: Reaction 2 . A typical
experimental curve for a reaction between AA and H2O2 is shown
in Figure 3. In this case, oxygen is first removed from the
phosphate buffer, and then AA and H2O2 are sequentially added.
As was already reported, a first-order dependence of the reaction
rate on [H2O2] is found3 under the same conditions used in the
present work.
RESULTS AND DISCUSSION
A comparison between the calibration curves for H2O2 obtained
by the Pt/ PPyox/ PPD(GOx) electrode both in the buffer solution
as such and in the presence of ascorbate at the highest concentra-
tion level adopted in this work (1 mM) is shown in Figure 1.
The same increase of the current signal is observed for equal
additions of hydrogen peroxide in both cases, which indicates that
ascorbate does not influence the electrode sensitivity. However,
the presence of AA results in a current decrease after each
addition due to the occurrence of reaction 2; of course, its slope
increases with H2O2 concentration.
Figure 2 shows the variation of [H2O2] during a kinetic run
with initial [AA] ) 1 mM: the curve clearly suggests that
hydrogen peroxide is a stable intermediate between consecutive
processes, easily detectable by the electroanalytical method used,
although its concentration was quite low under our experimental
conditions (maximum recorded value was around 7 µM).
The dependence on the [AA] was then evaluated by performing
kinetic runs in deareated solutions with different initial concentra-
tions of AA. The pseudo-first-order kinetic constants were
(8) Nelder, J. A.; Mead, R. Comput. J. 1 9 6 5 , 7, 308-313.
Analytical Chemistry, Vol. 69, No. 20, October 15, 1997 4115

concentrations of the main catalyzing ions (Cu2+ and Fe3+) were
measured, by means of ETAAS, in the buffer solution, prepared
with analytical grade chemicals used for all the kinetic runs. The
corresponding values (0.16 µM for Cu2+ and 0.15 µM for Fe3+),
though much smaller than those commonly used for studying
catalyzed processes (often4 larger than 10-5 M), do not exclude
the influence of ion-involving mechanisms on the kinetics.
As a result, an expression for the overall reaction rate under
“uncatalyzed” conditions could be the following:
Table 1. Kinetic Data Relevant to the Reaction
between Ascorbate and H2O2, Performed in the
Absence of Oxygena
initial concentrations (mM)
b
[AA]
[H2O2]
[Cu2+
]
t1/ 2 (s)
kc (s-1
)
0.030
0.100
1.000
0.030
0.100
1.000
0.002
0.002
0.002
0.002
0.002
0.002
1728
852
672
396
240
198
4.01 × 10-4
8.14 × 10-4
1.03 × 10-3
1.75 × 10-3
2.89 × 10-3
3.50 × 10-3
0.010
0.010
0.010
d[H2O2]
a
v ) -
) knc[AA][H2O2] +
Conditions: phosphate buffer solution (pH ) 7, I ) 0.1); temper-
ature, 27 °C. b Final concentration due to Cu2+ added to the buffer
solution. c Pseudo-first-order constants.
dt
kc,MAA[MAA][H2O2] + kc,MAA [MAA2][H2O2] (6)
∑
∑
2
M
M
calculated from the half-life time (t1/ 2) of H2O2, whose initial
concentration was always 2 µM; the relevant data are summarized
in Table 1.
It is apparent that, although an increase of the kinetic constant
with the initial AA concentration is observed, there is not a linear
correlation between them. This finding can be explained by
considering the simultaneous presence of two different mecha-
nisms for the AA oxidation by H2O2, similar to those proposed4
for its reaction with O2.
The simplest model can be based on a first reaction path
involving, as rate-determining step, a direct electron transfer
between the ascorbate ion and H2O2 (eq 4), while a second one
would involve H2O2 and one or more complexes between ascor-
bate and catalyzing metal ions, even present as traces in the
reaction mixture, e.g., Cu2+, Fe3+, etc.
k
nc being the kinetic constant for the uncatalyzed process, whereas
kc,MAA and kc,MAA are those relevant to the catalyzed ones. In eq
2
6, the effect on the kinetics due to the 1:1 (MAA) and 1:2 (MAA2)
complexes, which are those commonly formed by ascorbate with
different metal ions,10 has been taken into account.
According to the values of the stability constants for MAA and
MAA2 complexes reported in the literature,10 it can be predicted
that a large ascorbate excess may result in a complete displace-
ment of the complexation equilibria toward the MAA2 species,
thus making the [MAA] values negligible and the [MAA2] ones
no more correlated to [AA].
As for the catalytic effect, the MAA2 complexes are expected
to be less efficient than the MAA ones, in analogy with the
behavior10 of the mixed chelates involving ascorbate and another
ligand, like EDTA, in the AA oxidation by oxygen. The steric
hindrance due to the second ligand, making the electron transfer
between the metal ion and oxygen more difficult, has been
hypothesized to explain the lower reactivity observed in this case.
It is likely that some of the metal ions initially present in the
buffer solution used for the kinetic experiments of Table 1 are
almost completely in the MAA2 form, even when low AA
concentrations are adopted, as in the first run. As a consequence,
the contribution due to the MAA2 species to the reaction rate (see
eq 6) is not dependent on AA concentration, and a nonlinear
increase of the kinetic constant is observed when [AA] goes from
0.03 to 0.1 mM.
CH2OH
CH2OH
CHOH
CHOH
HC
HC
O
O
O
C
(4)
O
C
+ H2O2
H
+ 2H2O
C
C
–
O
C
C
O
O
O
ascorbate
DAA
The case relevant to the mononuclear complex CuAA+ is
indicated in eq 5; it is worth noting that DAA formed by the
reaction, though reported as the lactone, is probably present as
2,3-diketogulonate at the pH used for the kinetic study.
A further increase of AA concentration could have the same
effect also on the ions present at higher concentrations, thus
reducing their catalytic effect, mainly due to the MAA complexes.
This would balance the contribution to the reaction rate due to
the nonmediated mechanism (corresponding to the first term in
eq 6), which remarkably increases. As a result, only a slight
variation of the kinetic constant is observed.
CH2OH
CH2OH
CHOH
HC
CHOH
OH2
H
HC
OH2
O
C
O
(5)
O
+ 2H2O + Cu2+
O
C
+ H2O2
Cu
C
C
Kinetic runs were also performed by using the same concen-
trations of AA and H2O2 but with the addition of Cu2+ ions as
catalysts, at a final concentration of 10 µM. The relevant data,
reported in Table 1, confirm the small influence of the [AA] on
the rate of reaction 2, though the latter is always greater than the
one observed in uncatalyzed conditions. This was expected, since
the concentration of Cu2+ complexes is higher in this case.
O
OH2
C
C
O
O
OH2
O
CuAA+
DAA
As recently pointed out,9 the contribution due to metal ions
should always be considered, even if they are in trace amounts,
when the oxidation of AA is studied. In the present case, the
(10) Martell, A. E. Ascorbic Acid: Chemistry, Metabolism and Uses; ACS Advances
in Chemistry Series 200; American Chemical Society: Washington, DC, 1982;
Chapter 7, pp 153-178.
(9) Schaich, K. M. In Methods in Enzimology; Packer, L., Glazer, A. L., Eds.;
Academic Press Inc.: New York, 1990; Vol. 186, pp 121-125.
4116 Analytical Chemistry, Vol. 69, No. 20, October 15, 1997

µM) than those in experiments 4-6 was adopted. As a result,
though the correlation between v0 and [AA]0 is still not linear,
the influence of [AA]0 on v0 is higher (an apparent order 0.67 for
AA can be estimated in this case, compared with 0.40 obtained
when 5 µM Cu2+ was added). This is in agreement with eq 7,
since the contribution due to the complexes is lowered on going
Table 2. Kinetic Data Relevant to the Reaction
a
between Ascorbate and O2
initial concentrations (mM)
c
v0
(mM‚s-1
)
b
expt
[AA]
[O2]
[Cu2+
]
1
2
3
4
5
6
7
8
0.3
0.3
0.1
0.3
0.3
0.1
0.1
0.3
0.23
0.11
0.23
0.23
0.055
0.23
0.23
0.23
-
-
-
3.21 × 10-5
1.45 × 10-5
1.06 × 10-5
5.98 × 10-4
1.45 × 10-4
3.86 × 10-4
2.02 × 10-4
4.22 × 10-4
from 5 to 2 µM Cu2+
.
0.005
0.005
0.005
0.002
0.002
It is worth noting that the reaction between AA and O2 shows
the “expected” second-order kinetics (first order in respect to each
reagent) in uncatalyzed conditions. Thus, the contribution due
to the last two terms in eq 7 seems to be negligible in this case,
whereas the corresponding ones for the process AA + H2O2 have
already an influence on the kinetics (vide ante).
a
Conditions: phosphate buffer solution (pH ) 7, I ) 0.1), temper-
ature, 27 °C. b Final concentration due to Cu2+ added to the buffer
solution. c Initial rates, v0 ) d[H2O2]/ dt ≈ -d[O2]/ dt.
The method based on the initial rate has also enabled us to
evaluate directly the influence of metal ion traces on the AA
reactivity when virtually identical systems are considered.
In effect, kinetic experiments performed by using buffer
solutions prepared with salts (monobasic and dibasic sodium
phosphate) from different sources showed variations, even re-
markable, in the v0 value. This can be attributed only to
differences in the trace amounts of metal ions in the buffer
solutions, since the same concentrations of AA (1 mM) and O2
(0.23 mM) were always used to make the comparison.
The results suggest that an equation like v ) [AA]n[H2O2]
cannot be generally applied, even considering an apparent order
n < 1, to describe the kinetics of the reaction between AA and
H2O2 in uncatalyzed conditions. As a consequence, the 2/ 2 model
is not suitable for studying the overall process due to reactions 1
and 2.
On the other hand, even when the 1/ 1 model was applied,
considering the ascorbate concentration as a constant throughout
the reaction, it was not possible to obtain a good fit of the
experimental curve.
In any case, the experimental [H2O2] is lower than the
predicted one for long reaction times; this finding could indicate
the occurrence of other H2O2-consuming processes along with
reactions 1 and 2 (vide infra).
Investigation on the Kinetic Order: Reaction 1 . Kinetic
information on process 1 was also obtained through H2O2
monitoring, using the slope of the [H2O2] vs t curves for t f 0 as
a reliable estimate for the initial rate v0 of reaction 1.
Results relevant to kinetic runs performed in the presence of
oxygen are reported in Table 2.
Kinetic Runs with Low Oxygen or Ascorbate Concentra-
tions: Effects of DAA Reactivity. As cited before, the dehy-
droascorbic acid (DAA) produced by reactions 1 and 2 could be
responsible for the decrease of [H2O2] after long reaction times
not explained by kinetic models based only on two consecutive
processes. The possibility of a reaction between H2O2 and DAA
has been reported in a recent work11 based on a GC/ MS study,
suggesting that DAA, supposed12 to be rapidly hydrolyzed to 2,3-
diketogulonate in aqueous solution at pH ) 7, can be oxidized by
H2O2 to 2,3-diketo-4,5,5,6-tetrahydroxyhexanoic acid.
On the other hand, the possibility of a reaction between oxygen
and DAA has been also previously hypothesized,13 though no
kinetic data are available on this process.
Data relevant to experiments 1-3 suggest that the reaction is
unimolecular both toward AA and O2 in uncatalyzed conditions.
In fact, the initial rate is directly proportional to the different
concentrations of one of the reactants when the other is constant.
A different behavior is observed when Cu2+ is added as a
catalyst, at final concentration 5 µM: while v0 is still proportional
to [O2]0 (experiments 4 and 5), the linear correlation between v0
and [AA]0 is no longer observed (experiments 4-6). This finding
appears to be in agreement with the observations made on the
A careful analysis of the [H2O2] vs t profiles for experiments
performed under particular conditions provided an indirect
confirmation of the DAA reactivity. As an example, the response
obtained during experiment 5 in Table 2 is shown in Figure 4. In
this case, the global process is limited by the oxygen concentra-
tion, and the almost complete disappearance of H2O2 is observed
within a relatively short time (about 20-25 min).
Contrary to the case of the uncatalyzed reaction between AA
(1 mM) and O2 (0.23 mM), shown in Figure 2, a very good fit
was obtained by using the 1/ 1 model, whose application was
possible because AA was in excess in comparison with O2 and
H2O2. The fitting results suggest that the possible influence of
the reaction between H2O2 and DAA is quite small, at least under
these conditions, also due to the fact that H2O2 has been already
consumed when DAA reaches appreciable concentrations.
A very different behavior is observed in experiments 4 and
6-8 of Table 2, in which the initial concentration of AA is
data obtained for reaction 2 under comparable conditions:
a
prevailing mechanism based on parallel processes, like the one
summarised in eq 6, can be invoked for reaction 1 in the presence
of small concentrations of catalysts.
An expression for v0 in the case of reaction 1 could be then
d[O2]
v ) -
) knc[AA]0[O2]0 +
(
)
dt t)0
kc,MAA[MAA]0[O2]0 + kc,MAA [MAA2]0[O2]0 (7)
∑
∑
2
M
M
(11) Deutsch, J. C.; Santosh-Kuwar, C. R.; Kolhouse, J. F. Anal. Chem. 1 9 9 4 ,
66, 345-350.
(12) Tolbert, B. M.; Ward, J. B. Ascorbic Acid: Chemistry, Metabolism and Uses;
Advances in Chemistry Xeries 200; American Chemical Society: Washington,
DC, 1982; Chapter 5, pp 101-123.
where the symbols have the same meaning as in eq 6.
Data relevant to experiments 7 and 8 in Table 2 further support
this hypothesis: in this case, a concentration of Cu2+ lower (2
(13) Dunn, J. A.; Ahmed, M. U.; Murtiashaw, M. H.; Richardson, J. M.; Walla,
M. D.; Thorpe, S. R.; Baynes, J. W. Biochemistry 1 9 9 0 , 29, 10964-10970.
Analytical Chemistry, Vol. 69, No. 20, October 15, 1997 4117

the total consumption of AA. The concentration of the latter was
then monitored by UV spectroscopy (absorption at λ ≈ 260 nm)
on small aliquots of the reaction mixture with Cu2+ ) 5 µM at
different times. The corresponding values are also reported in
Figure 5a; in the inset, the detail of the UV spectra (with the
exception of point 3, for which no absorption at 260 nm is
detectable) is shown. It is apparent that a decrease of [H2O2]
occurs even after ascorbate has been completely consumed. An
experiment performed by adding 50 µM H2O2 to a phosphate
buffer solution containing 5 µM Cu2+ showed that the presence
of Cu2+ ions has a negligible effect on H2O2 decomposition. These
findings seem to confirm the already observed11 reactivity between
H2O2 and DAA. At the same time, reaction with oxygen cannot
be excluded.
Figure 4. Kinetic analysis of the concentration profile for hydrogen
peroxide relevant to the reaction mixture AA (0.3 mM) + O2 (0.05
mM) in phosphate buffer (0.1 M, pH ) 7) solution containing 5 µM
Cu2+ (temperature, 27 °C). The calculated concentrations were
obtained by using a model based on two consecutive pseudo-first-
order reactions. The reported concentrations of H2O2 were calculated
on the basis of the current response corrected for the contribution
due to ascorbate.
It is also interesting to point out that the slopes of the [H2O2]
vs t curves after AA disappearance are not significantly influenced
by Cu2+ concentration, thus suggesting that metal ions have a
slight catalytic effect in this case. This finding is consistent with
the weak affinity10 of DAA for metal ions, which minimizes the
possibility of an ion-mediated mechanism for the DAA-H2O2
process.
The simultaneous presence of ascorbic and dehydroascorbic
acids in the reaction mixture makes the system too complex to
draw direct kinetic information on the oxidation processes
involving DAA. However, the peculiar experimental device setup
for the present study seems also suitable for investigating systems
in which only DAA (or other oxidation products of AA) is present
as an oxidizable species. This investigation could contribute to
clarify, from a more general point of view, the network of reactions
involving ascorbate oxidation.
CONCLUSIONS
In the present work, a novel approach has been adopted to
investigate, from a kinetic point of view, the processes connected
to ascorbate oxidation by oxygen.
In particular, hydrogen peroxide, present as a stable intermedi-
ate in the reaction mixture, was monitored amperometrically by
a specially modified electrode able to drastically prevent the
simultaneous electrooxidation of ascorbate.
This peculiar experimental technique allowed us to obtain an
interesting picture of the first steps of the AA-O2 reacting system,
likely the most complete presented up to now by a single
experimental approach (comparison of homogeneous data), show-
ing that other processes probably involving the byproduct of AA,
dehydroascorbic acid, have to be considered, together with the
already known reactions AA + O2 and AA + H2O2. Moreover, it
emphasized the role played by catalyzing metal ions, even if
present in very low amounts.
Figure 5. Concentration profiles for hydrogen peroxide relevant to
the reaction mixtures AA (0.1 mM) + O2 (0.23 mM) in phosphate
buffer (0.1 M, pH ) 7) solution containing (a) Cu2+ (5 µM) and (b)
Cu2+ (2 µM). In (a), the concentration of ascorbate in the reaction
mixture, determined by UV spectroscopy at different times (see the
inset for the relevant spectra, with the exception of point 3, for which
no appreciable absorption at 260 nm is detectable), is also reported.
H2O2 concentrations were obtained from the corresponding current
values corrected for the contribution due to ascorbate.
In most cases, these effects could be quantitatively determined,
though, as a matter of fact, they prevented the successful
application of kinetic models based on simple consecutive reac-
tions.
Because of its complexity, the problem of ascorbate “autoxi-
dation” is far from being solved; however, the experimental device
presented in the paper offers new, original possibilities for further
studies on this system.
stoichiometrically lower than that of oxygen, the “correct” value
being equal to 2[O2]0, according to eqs 1 and 2. As an example,
the [H2O2] vs t profiles obtained in two cases (with different Cu2+
concentrations) are reported in Figure 5.
According to the reaction scheme based on processes 1 and
2, a stabilization of the H2O2 concentration should be reached after
Furthermore, the experimental approach adopted seems to be
a very promising method to get useful information also on similar
reaction systems involving hydrogen peroxide.
4118 Analytical Chemistry, Vol. 69, No. 20, October 15, 1997

ACKNOWLEDGMENT
Received for review April 1, 1997. Accepted July 24,
1997.X
The authors thank P. G. Zambonin for his stimulating interest
in this work. Financial support from Italian Research Ministery
(M.U.R.S.T.), National Research Council (CNR), and A.I.R.C. (a
fellowship to C.G.Z.) is gratefully acknowledged. Part of this work
is matter of the Ph.D. thesis by I.L.
AC970347T
X Abstract published in Advance ACS Abstracts, September 1, 1997.
Analytical Chemistry, Vol. 69, No. 20, October 15, 1997 4119