Oxidation of glyoxylic acid by cerium(IV): Oxygen-induced enhancement of the primary radical concentration

Article abstract of DOI:10.1021/jp960300w

In order to help understand the role of oxygen in Ce(IV)-induced oxidation of small carbonic acids, we investigated the reaction of glyoxylic acid (HCOCOOH) and Ce(IV) in 1 M sulfuric acid. Spectrophotometric data showed that in excess of glyoxylic acid the consumption of Ce(IV) obeys pseudo-first-order kinetics, with a rate constant of 8.8 L mol^-1 s^-1 at 25°C and an activation energy of 80 kJ mol^-1. Rapid-flow EPR measurements revealed an approximately 1:2:1 triplet with a g value of 2.0071 ± 0.0005 and a hyperfine splitting of 7.1 ± 0.2 G, assignable to the primary radical formed by abstraction of a hydrogen atom from hydrated glyoxylic acid. The rate constant for the anaerobic self-decay of the radical was measured as approximately 3.7 × 10⁹ L mol^-1 s^-1. Surprisingly, oxygen had no effect on the Ce(IV) kinetics, while the radical decay was significantly inhibited under aerobic conditions (ratio of experimental rate constants ≈6.3). Amperometric measurements revealed accompanying oxygen consumption. Analyses based on numerical simulations show that the observed oxygen-induced increase in radical concentration cannot be explained in the framework of standard autooxidation mechanisms. An alternative reaction scheme is suggested which reproduces the observed aerobic radical kinetics and which thus could be relevant to similar oxidation reactions.

Full text of DOI:10.1021/jp960300w

12342
J. Phys. Chem. 1996, 100, 12342-12348
Oxidation of Glyoxylic Acid by Cerium(IV): Oxygen-Induced Enhancement of the Primary
Radical Concentration
Bettina Neumann,† Oliver Steinbock, Stefan C. Mu1ller,†,‡ and Nar S. Dalal*
Department of Chemistry, Florida State UniVersity, Tallahassee, Florida 32306-3006, and Department of
Chemistry, West Virginia UniVersity, Morgantown, West Virginia 26506
ReceiVed: January 30, 1996X
In order to help understand the role of oxygen in Ce(IV)-induced oxidation of small carbonic acids, we
investigated the reaction of glyoxylic acid (HCOCOOH) and Ce(IV) in 1 M sulfuric acid. Spectrophotometric
data showed that in excess of glyoxylic acid the consumption of Ce(IV) obeys pseudo-first-order kinetics,
with a rate constant of 8.8 L mol-1 s-1 at 25 °C and an activation energy of 80 kJ mol-1. Rapid-flow EPR
measurements revealed an approximately 1:2:1 triplet with a g value of 2.0071 ( 0.0005 and a hyperfine
splitting of 7.1 ( 0.2 G, assignable to the primary radical formed by abstraction of a hydrogen atom from
hydrated glyoxylic acid. The rate constant for the anaerobic self-decay of the radical was measured as
approximately 3.7 × 109 L mol-1 s-1. Surprisingly, oxygen had no effect on the Ce(IV) kinetics, while the
radical decay was significantly inhibited under aerobic conditions (ratio of experimental rate constants ≈6.3).
Amperometric measurements revealed accompanying oxygen consumption. Analyses based on numerical
simulations show that the observed oxygen-induced increase in radical concentration cannot be explained in
the framework of standard autooxidation mechanisms. An alternative reaction scheme is suggested which
reproduces the observed aerobic radical kinetics and which thus could be relevant to similar oxidation reactions.
Introduction
The formation of peroxy radicals is considered to be part of
the chain propagation in the generally assumed mechanism for
autooxidation reactions:15
Over the past several decades the oxidation of organic
compounds by ceric ions has been the subject of numerous
studies.1 However, most reaction mechanisms are highly
complex and thus not yet fully understood, possibly because of
the lack of direct experimental data on the involvement of radical
species. A striking example is the Ce(IV) oxidation of malonic
acid (CH2(COOH)2) which was first investigated at least as early
as 1930.2 It was generally considered that in this reaction the
first step was the abstraction of a hydrogen atom by Ce(IV) to
form the malonyl radical, which then decayed by dispropor-
tionation.3 In 1994, however, Gao et al.4 demonstrated that the
first stable product was ethanetetracarbonic acid, formed by
dimerization of the primary (malonyl) radical. Thus, their
findings were in contrast to the generally assumed radical
disproportionation model. The ceric oxidation of malonic acid
is of particular interest because of its relevance for the widely
studied Belousov-Zhabotinsky5 (BZ) reaction.6-8 It is also
known that the dynamics of the BZ reaction as well as the
kinetics of the oxidation of malonic acid are strongly influenced
by molecular oxygen.9-13 For these aerobic reaction systems,
the formation of the peroxydicarboxymethyl radical had been
long postulated as an important intermediate but only very
recently were we able to identify and characterize this species.14
In the case of the ceric oxidation of malonic acid, we also found
that the kinetics of this radical can be readily explained in terms
of a standard autooxidation mechanism. On the other hand, it
is known that peroxy radicals of small hydrocarbon species are
not easily detected. It was thus of interest to probe the
fundamental question whether the formation of peroxy radicals
is necessarily the key process of aerobic oxidation reactions of
small carboxylic acids.
initiation:
formation of R• (e.g. by reaction with
a metal catalyst)
propagation:
R• + O2 f ROO•
ROO• + RH f ROOH + R•
termination: ROO• + ROO• f nonradical products + O2
where RH represents the organic substrate, R• the primary
radical, ROO• the peroxy radical, and ROOH the hydroperoxide
of RH that might further react with present metal ions.
Here we present results on the ceric oxidation of glyoxylic
acid (GOA; HCOCOOH) under aerobic and anaerobic condi-
tions as obtained from spectrophotometric, rapid-flow electron
paramagnetic resonance (EPR), amperometric measurements,
and numerical simulations. Glyoxylic acid was selected as a
test case for the standard model of autooxidation for several
reasons. First, it is one of the smallest carbonic acids that can
be oxidized by Ce(IV). Second, GOA is possibly an intermedi-
ate in the ceric oxidation of malonic acid under oxygen:
Recombination of the peroxymalonyl radical could form tar-
tronic acid (HCOH(COOH)2) and either GOA or mesoxalic acid
(CO(COOH)2) as the immediate decay products. GOA has also
been postulated to be a stable intermediate in the Ce(IV)
oxidation of tartronic acid and can substitute malonic acid as
the organic substrate for oscillating BZ reactions.16-18 Hence,
kinetic data on the oxidation of GOA by Ce(IV) should provide
important information for a more complete understanding of
the aerobic oxidation of malonic acid and other related reactions
(as for example the BZ reaction).
† Max-Planck-Institut fu¨r molekulare Physiologie, Rheinlanddamm 201,
D-44139 Dortmund, Germany.
In one of the earlier studies of the reaction between Ce(IV)
and glyoxylic acid Jwo and Noyes3 measured the stoichiometric
ratio ∆[Ce(IV)]/∆[GOA] as 1.9. Furthermore, formic acid
(HCOOH, FA) and CO2 were identified as the reaction products.
‡ Otto-von-Guericke-Universita¨t Magdeburg, Institut fu¨r Experimentelle
Physik, Universita¨tsplatz 2, 39106 Magdeburg, Germany.
X Abstract published in AdVance ACS Abstracts, June 15, 1996.
S0022-3654(96)00300-0 CCC: $12.00 © 1996 American Chemical Society

Oxidation of Glyoxylic Acid by Cerium(IV)
J. Phys. Chem., Vol. 100, No. 30, 1996 12343
These results were obtained under unspecified atmospheric
conditions. Nevertheless, these findings suggest the following
simple reaction mechanism involving the primary radical •CO-
(COOH):
HCOCOOH + Ce(IV) f •COCOOH + Ce(III) + H+ (R1)
•COCOOH + •COCOOH + H2O f
HCOCOOH + HCOOH + CO2 (R2)
To our knowledge the radical in this reaction (•COCOOH) has
neither been identified nor kinetically studied so far. However,
glyoxylic acid related radicals produced by flash photolysis18,19
have been detected earlier: Using EPR spectroscopy Steenken
et al.18 reported the detection of radicals formed from photolysis
of GOA at pH 3 and higher. At lower pH, however, no radicals
were observed after irradiation unless the sensitizer acetone was
added to the solution. We were intrigued by the fact that in
contrast to the case of malonic acid, the possibility of an
intermediate formed by the dimerization of GOA• and any role
of oxygen were not discussed. The present study suggests that
these are indeed key processes in the oxidation of GOA by Ce-
(IV) and thus likely in other related systems.
Experimental Section
Kinetic and stoichiometric measurements of Ce(IV) were
performed utilizing a Hewlett-Packard HP8452 diode array
spectrophotometer. The time dependence of the Ce(IV) con-
centration was followed at a wavelength of 320 nm. Kinetic
experiments with initially added Ce(III) were carried out with
anhydrous cerous (Ce(III)) sulfate (Fluka, purum). The stoi-
chiometry was determined by photometric analysis of the
residual Ce(IV) concentration approximately one day after
initiation of the reaction. These experiments were performed
at a wave length of 398 nm due to higher Ce(IV) concentrations.
In all experiments Ce(IV)(SO4)2 × 4H2O (Merck) and GOA
(Fluka) were used. Arrhenius parameters were obtained from
variable temperature measurements using a thermostated water
bath.
EPR measurements were made with a Bruker ER200D
X-band (9.5 GHz) EPR spectrometer and a rapid flow attach-
ment.14,20 The two reactant solutions ([GOA] ) 1.0 M and [Ce-
(SO4)2] ) 14.0 mM, both in 1 M H2SO4) were pumped to a
quartz flat cell equipped with a T-shaped mixing device (pump
used: Gilson Miniplus3). The quartz flow cell (Wilmad WG-
804-Q) was mounted in the EPR cavity and had an active and
a dead volume of 125 µL and <50 µL, respectively. The stable
radicals 4-hydroxy-2,2,6,6-tetramethylpiperidine-N-oxyl (TEMPO)
and 1,1-diphenyl-2-picrylhydrazyl (DPPH) (Aldrich) were used
as standards for the determination of radical concentrations and
g values. The microwave frequency was measured with a
Hewlett-Packard (HP5420) digital frequency counter and the
magnetic field was calibrated with a gaussmeter for measuring
the g values and hyperfine couplings.
Figure 1. Spectrophotometric data for the Ce(IV) oxidation of
glyoxylic acid. (a) Dependence of the amount of reduced Ce(IV) on
the initial concentration of glyoxylic acid. Open circles: aerobic
conditions; solid circles: anaerobic conditions. (b) Semilogarithmic
plot of the Ce(IV) concentration evolution during the reaction at 25.0
°C. Data were obtained under anaerobic (solid lines) and aerobic ([O2]0
≈ 1.0 mM; dashed lines) conditions. Initial concentrations: [Ce(IV)]
) 0.1 mM, [GOA] ) 4.0, 6.0, 10.0 mM, [H2SO4] ) 1 M.
revealed three peaks with the chemical shifts: 174.1 ppm (1),
166.8 ppm (2), and 87.2 ppm (3). Peak (2) was identified as
formic acid, while signals (1) and (3) correspond to residual
glyoxylic acid. From comparison with literature spectra and
numerical predictions, peak (1) and (3) were assigned to the
hydrated form (HCOCOOH × H2O) of glyoxylic acid.21
Results
Stoichiometric Analysis. With the view of understanding
the overall Ce(IV)-GOA reactivity we first carried out mea-
surements on the reaction stoichiometry under well-defined
aerobic and anaerobic conditions. Figure 1a shows the con-
sumption of Ce(IV) for various initial GOA concentrations. The
consumption was measured spectrophotometrically as the dif-
ference between the initial (here 1.0 mM) and the residual Ce-
(IV) concentration (wavelength 398 nm). The applicability of
Beer’s law was checked in additional experiments which
revealed no significant deviations from a linear dependence
between absorbance and [Ce(IV)]. The curves in Figure 1a
consist of two linear segments. The horizontal segments arise
from experiments in which the initial GOA concentration was
too high for a complete reaction with 1 mM Ce(IV), while the
points on the tilted segment represent experiments with no
residual GOA. Based on these data we measured the stoichio-
metric ratio ∆[Ce(IV)]/∆[GOA] under aerobic and anaerobic
conditions as 1.98 ( 0.1 and 2.02 ( 0.1, respectively. The
results obtained indicate no significant dependence of the
stoichiometry on the atmospheric conditions.
Oxygen consumption was measured amperometrically using
a Clark-type electrode. Prior to the experiments the electrode
was calibrated utilizing solutions of known oxygen concentration
(solutions were equilibrated by the means of a gas mixing
pump).
A 270 MHz FT NMR system (JNM-EX270, JEOL) was used
for the identification of the reaction products. Samples of 0.6
M Ce(IV)(NH4)2(NO3)6 (Fluka) and 0.6 M GOA in 1 M sulfuric
acid were allowed to react at room temperature for more than
12 h. Proton-decoupled 13C NMR spectra were then taken with
dioxane (67.4 ppm) as an internal standard. The spectra
This result is important because it implies a basic constraint
on the possible mechanisms of the reaction. Note that the

12344 J. Phys. Chem., Vol. 100, No. 30, 1996
Neumann et al.
dependent line broadening and microwave power saturation
effects, as deduced from several detailed measurements. The
g values was measured as 2.0071 ( 0.0005. These features
imply that the radical’s unpaired electron is delocalized (at least
partially) over a heavy-atom framework involving two magneti-
cally equivalent protons. Further clues to radical identification
were derived from the fact that the radical is not likely to be of
alkoxy or peroxy type, because in that case the g value difference
from g ) 2.0027 would be an order of magnitude larger while
the proton hyperfine coupling would be significantly (about
50%) smaller.22 Another helpful clue was that Ce(III) com-
plexation leads to an increase in g values. For example, the g
•
value of the HO2 radical increases from 2.0140 to 2.0180 on
Figure 2. EPR spectra observed from the reaction between Ce(IV)
and glyoxylic acid under anaerobic (a) and aerobic (b) conditions. Both
spectra were obtained under identical experimental conditions. Initial
concentrations: [Ce(IV)] ) 7.0 mM, [GOA] ) 0.5 M, [H2SO4] ) 1
M.
Ce(III) complexation.23,24 Generally, it is expected that Ce-
(III) bonds to the radical via the CO oxygens. Based on these
arguments we tentatiVely assign the observed triplet to the
following radical complex, formed by hydrogen abstraction from
the hydrated form of glyoxylic acid:
identical stoichiometry obtained from anaerobic and aerobic
measurements can be due to the same major reaction pathway
but might as well be caused by different pathways.
Reaction Kinetics of Ce(IV). The kinetics of the oxidation
of GOA by Ce(IV) was studied in excess of GOA. Figure 1b
shows the photometrically detected decay of Ce(IV) for three
different GOA concentrations. The curves reveal only small
deviations from the expected linear dependence. Furthermore,
we found no systematic differences between the cerium kinetics
under aerobic and anaerobic conditions. The pseudo-first-order
value of k1 at 25.0 °C was calculated from various experimental
b. Oxygen Effect on Radical Concentration. A significant
observation from the results presented in Figure 2, a and b, is
that admission of oxygen to the GOA/Ce(IV) reaction system
leads to a significant increase in the radical signal (about a factor
of 3 under the given conditions) without any new EPR peak.
This is in contrast to the case of malonic acid where oxygen
exposure caused a rapid decrease in the malonyl radical
concentration until oxygen was entirely consumed, with con-
comitant formation of the peroxy radical.14 As will be shown
in the following, this unexpected difference in radical concentra-
tions can be attributed to an additional oxygen-induced pathway
in the reaction mechanism.
c. Radical Kinetics. The radical decay kinetics were
monitored under anaerobic as well as aerobic conditions. These
measurements were performed using the EPR stopped-flow
technique with the initial concentrations and apparatus as
described in the Experimental Section. Figure 3a shows the
decay curve for an anaerobic (i.e., nitrogen equilibrated) reaction
mixture. Based on the reactions R1 and R2 we expect a second-
order self-decay of the radical that is described by
rate constants as 8.8 ( 0.2 L mol-1 s-1
.
For probing the reaction further we measured the temperature
dependence of k1 as summarized below. The Arrhenius plot
of ln k1 vs 1/T was analyzed over the temperature range of 14-
40 °C. The initial concentration of Ce(IV) in all experiments
was 0.1 mM, while the initial GOA concentration was varied
between 4.0 and 20.0 mM. The Arrhenius plot revealed a linear
dependence, allowing the determination of the activation energy
Ea and the preexponential factor A of reaction (R1) to be Ea ≈
80 kJ mol-1 and A ≈ 1.4 × 1015 L mol-1 s-1. The measured
activation energy of 80 kJ mol-1 is in the typical range for ceric
oxidation reactions of organic acids; corresponding values for
malonic (MA) and tartronic acid (TA) are 46 and 74 kJ mol-1
,
respectively.3 The trend Ea(MA) < Ea(TA) < Ea(GOA) reflects
the expected significance of steric factors.
Additional measurements were carried out to study the
possible back reaction between Ce(III) and reaction products
of GOA. In these experiments the initial concentrations of Ce-
(IV) and GOA were 0.1 and 6.0 mM, respectively. It was found
that addition of Ce(III) in concentration of <1 mM caused no
detectable effects on the kinetics. Higher concentrations of Ce-
(III) (10 and 20 mM), however, slightly decelerated the reaction,
resulting in small deviations from the simple exponential decay.
Since the overall effect of Ce(III) on the reaction was considered
to be rather small, no further quantitative analysis was carried
out.
EPR Measurements. a. Radical Identification. Since an
important goal of this study was to directly detect and identify
any free-radical involvement and the radical species were
expected to be short-lived, we utilized the fast-flow EPR
technique as described in the Experimental Section. Figure 2
shows typical EPR spectra obtained from reaction mixtures of
GOA and Ce(IV) that were purged with nitrogen (Figure 2a)
and oxygen (Figure 2b), respectively. In both cases the major
feature is the approximately 1:2:1 triplet, with a hyperfine
splitting of 7.1 ( 0.2 G. The observed deviation from the 1:2:1
line shape under oxygen (Figure 2b) can be ascribed to the mI-
d[GOA•]/dt ) k1[GOA]0[Ce(IV)] - 2k2[GOA•]2 (1)
if GOA is in excess (i.e., its initial concentration, [GOA]0 .
[Ce(IV)]). Moreover, spectrophotometric measurements showed
that [Ce(IV)] obeys a simple first-order equation
[Ce(IV)] ) [Ce(IV)]0 exp(-kexpt)
(2)
with the initial concentration [Ce(IV)]0 and the experimental
rate constant kexp ) [GOA]0k1. Further, assuming a steady-
state radical concentration, i.e., d[GOA•]/dt ) 0, eqs 1 and 2
yield
[GOA•] ) {([Ce(IV)]0kexp)/(2k2)}1/2 exp(-0.5kexpt) (3)
Hence, the radical concentration is expected to decrease with
the rate 0.5kexp. In order to test this prediction we first utilized
the spectrophotometric data to calculate kexp for the conditions
of the EPR stopped-flow experiment (Figure 3a). This value
was found to be kexp ) 4.4 s-1. We then used this kexp and eq
3 to predict the time dependence of the radical decay. The

Oxidation of Glyoxylic Acid by Cerium(IV)
J. Phys. Chem., Vol. 100, No. 30, 1996 12345
TABLE 2: Experimental Values of the Oxygen
Consumption for Several Initial Concentrations of GOA and
a
Ce(IV)a
[GOA]0/M [Ce(IV)]0/mM consumed [O2]/mM ∆[Ce(IV)]/∆[O2]
0.2
0.2
0.4
0.4
0.4
0.4
0.4
0.70
1.40
0.42
0.70
0.98
1.40
1.96
0.03
0.05
0.03
0.05
0.07
0.09
0.11
23
28
13
15
14
16
18
a The initial concentration of oxygen was kept constant at [O2]0 ≈
0.15 mM.
b
Figure 4. Two examples of oxygen consumption obtained from
amperometric measurements. Initial concentrations: [GOA] ) 0.2 M,
[Ce(IV)] ) 1.4 mM (top curve, solid line), [GOA] ) 0.4 M, [Ce(IV)]
) 1.96 mM (bottom curve, dashed line).
Figure 3. Experimental EPR signal decay and calculated curves for
the GOA• kinetics under anaerobic (a) and aerobic (b, [O2]0 ≈ 1.0 mM)
conditions. Initial concentrations: [Ce(IV)] ) 7.0 mM, [GOA] ) 0.5
M. Calculated curves represent an exponential dependence with
experimental rate constants: kexp ) 4.4 s-1 (a, cf. eq 3), kexp ) 0.35
s-1 (b).
the EPR signal integration and that in the residence time tr we
estimate the relative error of k2 as (30%.
TABLE 1: Experimental Values of the Rate Constant k2
for Different Initial Concentrations of Glyoxylic Acida
Figure 3b shows a characteristic example for the radical decay
under aerobic conditions. In these experiments both of the stock
solutions were equilibrated with oxygen for at least 30 min prior
to the start of the measurement. The oxygen concentration was
estimated as 1.0 × 10-3 M. Concentrations, sample tempera-
ture, and EPR parameters were the same as in the previous
experiments (Figure 3a). The decreasing radical concentration
obeys a simple exponential function with an experimental rate
constant kexp ) 0.35 s-1, in contrast to the value of 0.5 × 4.4
s-1 under anaerobic conditions. In the presence of oxygen,
therefore, the radical decay is significantly inhibited: The
average ratio of the experimental rate constants under anaerobic
and aerobic conditions was found to be 6.3. This result implies
that the radical decay inhibition is essentially the source of the
observed oxygen influence on the intensity of the EPR signal
(Figure 2a,b), i.e., increased radical accumulation under aerobic
conditions.
[GOA]0/M [Ce(IV)]fl/mM kexp/s-1 [GOA•]fl/µM k2/L mol-1 s-1
0.1
0.2
0.3
0.4
0.5
4.92
3.46
2.43
1.71
1.20
0.88
1.77
2.65
3.53
4.41
0.75
0.82
1.02
0.98
0.83
3.9 × 109
4.5 × 109
3.1 × 109
3.1 × 109
3.8 × 109
a The table summarizes data that was used to derive k2 on the basis
of eq 4. Differences in k2 mainly reflect the experimental error in the
determination of [GOA•]fl from the EPR signals.
dashed curve in Figure 3a shows this calculated behavior, which
can be seen to be in good agreement with the observed EPR
signal decay. These results thus strongly support our initial
hypothesis that the radical consumption is caused by a second-
order self-decay.
Equation 3 also enabled us to measure the rate constant k2.
By defining [GOA•]fl and [Ce(IV)]fl as the actual concentrations
of GOA• and Ce(IV) established during the EPR flow experi-
ment, we find that
Oxygen Consumption Measurements. Since the oxygen
enhancement of radical concentration points toward a direct
consumption of oxygen during the reaction, we performed
systematic measurements of the oxygen concentration using a
Clark electrode. Table 2 summarizes the total oxygen con-
sumption for various initial Ce(IV) and GOA concentrations.
The data reveal that an increase in [Ce(IV)] as well as in [GOA]
causes enhanced consumption of oxygen. Equally significantly,
the data show that the oxygen consumption is approximately
10-30 times smaller than the total amount of reduced Ce(IV).
The precise stoichiometric ratio [O2]consumed:[Ce(IV)]consumed
depends on the initial values of [GOA] and [Ce(IV)]. Two
characteristic examples of the temporal evolution of the oxygen
concentration during the reaction are shown in Figure 4. From
these data it is clear that the oxygen concentration decreases
2
k2 ) ([Ce(IV)]flkexp)/(2[GOA•]fl )
(4)
[Ce(IV)]fl can be calculated from eq 2 by considering the
residence time of the flow system (tr ) 0.4 s) as the reaction
time t. The radical concentration [GOA•]fl was measured in
numerous EPR experiments and calibrated by comparison with
the EPR signal of 4-hydroxy-2,2,6,6-tetramethylpiperidine-N-
oxyl. Table 1 summarizes the results for k2 obtained under five
different experimental conditions at 25 °C. The entries for k2
in Table 1 show reasonable consistency with the average value
of k2 ) 3.7 × 109 L mol-1 s-1. Considering the error due to

12346 J. Phys. Chem., Vol. 100, No. 30, 1996
Neumann et al.
Figure 5. Typical example of the simulated time evolution of radical
concentrations obtained on the basis of the standard autooxidation
mechanism (cf. Introduction) including the anaerobic reactions R1 and
R2. The dashed line represents the [GOA•] under anaerobic conditions.
Under aerobic conditions the model predicts the [GOAO2•] and [GOA•]
kinetics as shown by the solid and dotted line, respectively. The circles
represent experimental data for [GOA•] under aerobic conditions.
slowly compared to the reduction of Ce(IV) and is therefore
not directly influenced by this catalyst.
Theoretical Modeling. In order to obtain further clues to
the mechanism for the observed, rather unusual, effect of radical
decay inhibition by oxygen, we carried out theoretical modeling
of the basic reaction kinetics. As a first step, we wanted to
assess the extent to which standard autooxidation mechanisms
(summarized in the Introduction) could explain the observed
oxygen effect. All numerical simulations were performed by
numerical integration of the underlying rate equations using the
FACSIMILE program.
Figure 6. (A) Scheme of the modified standard autooxidation model
that leads to an increase in [GOA•] under oxygen atmosphere. The
key process is the direct reaction of the primary radical with oxygen.
The loop between GOA•, oxygen, and X represents an unknown reaction
subset which could involve the peroxy radical and other intermediates.
(B) Alternative reaction scheme which is capable of modeling the
reported radical dynamics. The first stable intermediate of the primary
radical decay, Y, reacts with oxygen to efficiently produce the
experimentally observed radical concentration. Similar to the oxygen
involving reaction steps (R3, R4) in (A) the oxygen-catalyzed back
reaction (R7) represents a possibly complex mechanism.
As outlined in the Introduction, it is known that in most
autooxidation reactions a peroxy radical and its corresponding
peroxide are formed. The peroxy radicals form from a fast
reaction between the primary radical (here GOA•) with molec-
ular oxygen. However, our EPR spectra revealed no evidence
increase of [GOA•]: Typically, the aerobic decay of GOA• was
nearly identical with the anaerobic behavior. In particular, the
results showed no influence of oxygen on the rate of the GOA•
decay when simulations where performed with the expected,
nearly diffusion-controlled, rate for the reaction between oxygen
and GOA•. Additional simulations of the observed oxygen
consumption usually revealed very fast depletion that is not in
agreement with the experimental data. We therefore conclude
that the standard autooxidation mechanism is not even quali-
tatively capable of reproducing the experimental observations.
In order to help answer the query whether similar models
can reproduce our EPR findings, we generalized the chain
propagation reactions of the standard model. Since a fast
depletion of oxygen obviously diminishes a possible inhibition
process, we excluded oxygen-consuming reactions by modifying
the second step of the propagation loop. However, the reaction
between oxygen and the primary radical remains the key process
of the oxygen effect on GOA•:
•
for the existence of GOAO2 . Possible explanations for our
•
failure to detect the EPR signal from GOAO2 include the
•
following: (1) GOAO2 is so short-lived that it reaches EPR-
detectable concentrations only within the dead time of our EPR
flow-experiment; (2) concentrations are below the detection limit
during the whole course of the reaction, implying that although
•
GOAO2 is formed, it undergoes a fast conversion to a
nonradical species; (3) the formation of a peroxy radical is just
not favored or is, at least, not a major pathway here. Notice,
of course, that the explanations 1 and 2 explicitly require a
reaction between the primary radical GOA• and molecular
oxygen, i.e., assume that the reaction follows the standard
autooxidation pathway.
Numerical analyses were performed on the basis of the
standard autooxidation mechanism including the anaerobic
reaction steps (R1) and (R2). A typical example is shown in
Figure 5, presenting the anaerobic and aerobic evolution of the
radical concentration, dashed and dotted line, respectively. The
formation of a peroxy radical (solid line) would inhibit the
production of GOA• within the first 100 ms. This time interval,
however, was not resolved in our flow experiments. The crucial
test for the standard model was to reproduce the observed
temporal evolution of [GOA•] under aerobic conditions, shown
by open circles in Figure 5. The depicted example reveals an
irreconcilable discrepancy between the observed and calculated
decay behaviors. On the basis of the results obtained by varying
the rate constants systematically within broad intervals and their
comparison with the experimental data, we conclude that this
discrepancy is characteristic for the standard model. The
involved reaction steps cannot reproduce an oxygen-induced
GOA + Ce(IV) f GOA• + Ce(III) + H+
GOA• + O2 f X
(R1)
(R3)
X f GOA• + O2
(R4)
(R5)
GOA• + GOA• f P
A simple sketch of this mechanism is shown in Figure 6A.
The loop between GOA• + O2 and X represents an unknown
reaction subset which includes the peroxy radical and possibly

Oxidation of Glyoxylic Acid by Cerium(IV)
J. Phys. Chem., Vol. 100, No. 30, 1996 12347
experiments. Similar to the oxygen-involving reactions in (R3)
and (R4), the oxygen-catalyzed back reaction (R7) represents a
possibly complex mechanism. The solid line in Figure 7 shows
the radical decay calculated on the basis of the above model
using the rate constants k6 ) k2 ) 3.7 × 109 L mol-1 s-1, k7 )
6.6 × 103 L mol-1 s-1, k8 ) 0.75 s-1 and an oxygen
concentration of 1 mM. This curve is in excellent agreement
with the experimental data obtained under aerobic conditions.
We note that the model is also consistent with the described
anaerobic radical decay, and it violates neither the measured
stoichiometry nor the observed Ce(IV) kinetics. However, it
does not consider the consumption of oxygen. Since the
consumption is small and fairly slow, it seems likely that an
intermediate of the reaction loop between Y and GOA• involves
an additional pathway, controlling the oxygen depletion process.
Figure 7. Comparison of the experimental data (dotted lines: (a)
anaerobic, (b) aerobic conditions) of the radical kinetics to the simulated
aerobic radical decay curves (solid line, Model B; dashed line, Model
A) in a semilogarithmic presentation.
Discussion
This study establishes that the anaerobic oxidation of gly-
oxylic acid by Ce(IV) in 1 M sulfuric acid can indeed be
described by the two reactions (R1) and (R2). While this
conclusion is in agreement with the earlier studies,3 the present
investigation identifies for the first time the involved radical
and furthermore reports the rate constants k1 and k2. The
analysis of k2 was based on the result that the anaerobic kinetics
of GOA• follow a second-order self-decay that obeys a steady-
state approximation. It is reasonable, however, to assume that
the reaction step (R2) consists of at least two irreversible reaction
steps in which a dimer of the primary radical occurs as an
intermediate. In this context k2 has to be considered as the rate
constant of the dimer formation.
various other intermediates. Its rate constants correspond to
the rate-limiting steps of this subset. Numerical simulations
showed that additional reactions involving Ce(IV) are of minor
relevance due to the low concentrations of the possible reaction
species. A significant influence of Ce(III) (i.e., back reactions)
was also ruled out, since addition of Ce(III) to the reaction
mixture showed no significant effect in our spectrophotometric
data.
A realistic model has to explain the observed inhibition of
GOA• decay under aerobic conditions. The dashed line in
Figure 7 shows a typical logarithmic plot of the primary radical
concentration as a function of time, obtained on the basis of
the above model (rate constants k3 ) 1 × 109 L mol-1 s-1, k4
) 500 s-1, k5 ) k2 ) 3.7 × 109 L mol-1 s-1). During the first
two seconds the concentration exhibits a fast decay governed
by the anaerobic reaction dynamics. It is followed by a slower,
nearly exponential, decay. Although the rate of this slower
decay is consistent with the observed aerobic kinetics, the
realized concentrations are too low and clearly not in agreement
with the experimental data. Systematic simulations, testing the
parameter space (k3, k4) within the limits 10-2 and 1010, reveal
that this class of reaction mechanisms (Figure 5a) is simply not
capable of generating a fast transition to the observed slow
exponential decay and thus cannot reproduce the measured
radical concentration of >2 × 10-6 M. Again, these results
emphasize the failure of standard autoxidation mechanisms to
reproduce the measured oxygen-induced inhibition of the
primary radical reactivity.
According to standard autoxidation mechanisms, oxidation
reactions that are initiated by the formation of carbon-centered
radical species and exposed to molecular oxygen are generally
assumed to form peroxy radicals at an almost diffusion-
controlled rate. So far, however, only a few peroxy radicals of
small organic molecules have been directly identified. This fact
is generally explained in terms of low concentration and fast
decay of this species. Our experiments show no evidence for
•
the formation of the peroxy radical GOAO2 . Furthermore,
exposure of the reaction mixture to oxygen does not lead to
quenching of the primary radical but rather to the opposite
effect: Oxygen causes a significant decrease of the experimental
decay constant of the primary radical disappearance (Figure 3).
On the basis of our theoretical analysis as presented in the
previous section, we can conclude that the general model of
autoxidation processes fails to describe these experimental
results. Accordingly, the formation of a possible peroxy radical
cannot lead to the observed fast increase of the GOA•
concentration, under aerobic conditions. This result strongly
supports our suggestion that certain organic acids like GOA
must follow a fundamentally different reaction mechanism. The
scheme presented in Figure 6b can explain our experimental
results. The main feature of this mechanism is the oxygen-
catalyzed reaction of the first intermediate that is formed by
the radical decay process. While the exact identity of this
intermediate (species labeled Y in Figure 6b) remains to be
elucidated, the proposed scheme readily reproduces the observed
increase in [GOA•] and the observed decrease in experimental
rate constants.
In the following, we suggest an alternative reaction scheme
that readily models the reported radical dynamics (general
structure shown in Figure 6B):
GOA + Ce(IV) f GOA• + Ce(III) + H+
GOA• + GOA• f Y
(R1)
(R6)
(R7)
(R8)
Y + O2 f GOA• + O2
Y f P
This mechanism postulates a reaction between oxygen and an
unstable product of the GOA• dimerization. This product, the
species Y in the above reaction scheme, could be a precursor
or unstable conformation of the expected recombination product
tetrahydroxysuccinic acid (THSA). Our preliminary results
show that the disodium salt of THSA, a substance that was first
reported in 1885 by Miller, decays under the given experimental
conditions (1 M H2SO4) on a time scale of many hours.25-28
Its self-decay is therefore not relevant for our fast EPR
The investigated reaction system shows, for example, remark-
able contrast to the recently reported free-radical chemistry of
the ceric oxidation of malonic acid.14 In the malonic acid
system, molecular oxygen was considered as a fast kinetical
switch, influencing mainly the kinetics of Ce(IV) and the
peroxymalonyl radical. The slow decay of the primary malonyl
radical revealed therefore no significant dependence on the initial

12348 J. Phys. Chem., Vol. 100, No. 30, 1996
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oxygen concentration. The elucidation of the detailed structure
of aerobic reaction pathways for small organic acids is therefore
an important challenge for future studies. The results could
have the potential for revealing a different aerobic oxidation
mechanism. Also the role of glyoxylic acid in the complex
oxidation of malonic acid and related cross reactions are of
interest for future studies. A thorough knowledge of these
reaction mechanisms should furthermore provide a better
understanding of oxygen effects8-11 and aging29,30 in the self-
organizing Belousov-Zhabotinsky reaction.
The comparison of the aerobic radical mechanisms during
the Ce(IV) oxidation of glyoxylic and malonic acid gives rise
to another intriguing question of general importance: What are
the structural differences between these two organic molecules
that cause the observed, significant differences in their reaction
pathways and radical kinetics? This question provides, in itself,
ample reason for further systematic studies on the aerobic
oxidation of organic acids and the necessary accumulation of
additional data.
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Acknowledgment. This research was supported by the
Deutsche Forschungsgemeinschaft, the U.S. Department of
Energy, and the Florida State University. O.S. thanks the Fonds
der Chemischen Industrie for a Liebig Fellowship.
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