The kinetics of the oxidation of l ascor
Free Radical Biology & Medicine, Vol. 18, No. l, pp. 8 5 - 9 2 , 1995
Copyright © 1994 Elsevier Science Ltd
Printed in the USA. All rights reserved
0891-5849/95 $9.50 + .00
Pergamon
0891-5849(94)E0133-2
Original Contribution
THE KINETICS OF THE OXIDATION OF L-ASCORBIC ACID
BY PEROXYNITRITE
DELLAND BARTLETT,* DANIEL F. CHURCH,* PATRICIA L. BOUNDS,* and W. H. KOPPENOL§*
Departments of *Chemistry and *Biochemistry and ~BiodynamicsInstitute, Louisiana State University, Baton Rouge, LA,
USA; and *Department of Chemistry and Physics, Southeastern Louisiana State University, Hammond, LA, USA
(Received 7 March 1994; Revised 3 June 1994; Accepted 6 June 1994)
Abstract Peroxynitrite [O=NOO-, oxoperoxonitrate(l-)] is a strong oxidant that may be formed in vivo by the reaction of
02"- and NO'. Oxoperoxonitrate(1-) reacts with molecules in aqueous acidic solutions via pathways that involve the highly
reactive hydrogen oxoperoxonitrate either as an intermediate in a first-order reaction or as a reactive agent in a simple secondorder reaction. ESR experiments show that hydrogen oxoperoxonitrate oxidizes monohydrogen L-ascorbate by one electron:
when mixed at pH ca. 5 and passed through a flow cell within 0.1 s, the two-line ESR signal of the ascorbyl radical anion (all
= 0.18 T, g = 2.005) is observed. The overall stoichiometry of the reaction was 1 mol of ascorbate oxidized per mol of
oxoperoxonitrate(l-) added. The kinetics of the reaction were studied over the pH range 4.0-7.5 by stopped-flow spectrometry.
Hydrogen oxoperoxonitrate,observed between 300 and 350 urn, and the oxoperoxonitrate(l-) anion, at 302 nm, disappear faster
than predicted for the first-order isomerization to NO3 . The rate increases from pH 4 to 5.8, and then decreases with increasing
pH. The rate variation suggests a bimolecular reaction either between the oxoperoxonitrate(I-) anion and ascorbic acid or between
hydrogen oxoperoxonitrate and the monohydrogen ascorbate anion. Although the two pathways are kinetically indistinguishable,
the pKa values of ascorbic acid and hydrogen oxoperoxonitrate strongly suggest that the reacting species are hydrogen oxoperoxonitrate and monohydrogen ascorbate. The second-order rate constant for this reaction is 235 + 4 M ts-t at 250(7. The enthalpy
and entropy of activation are AH* = 9.3 _+ 0.5 kcal/mol and AS* = -16 _+ 2 cal/(mol.K), respectively.
Keywords---Ascorbate, Peroxynitrite, Oxoperoxonitrate(I-), Sulfhydryl, Antioxidant defense, Reaction kinetics, Stopped-flow
spectrophotometry, Electron spin resonance
duced by neurons and endothelial cells normally is of
the order of tens to hundreds nanomolar, m u c h less
than the 5 to 40 # M concentration of superoxide dismutase in cells. However, u n d e r pathologic conditions,
higher levels of nitric oxide synthase can be induced.
A concentration of 0.5 # M nitrogen m o n o x i d e has been
measured 4 near the surface of an endothelial cell. Near
activated macrophages concentrations as high as 10
/zM appear likely (J. S. Beckman, personal c o m m u n i cation). Thus, even in the presence of superoxide dismutase, oxoperoxonitrate(1-) can be formed.
O x o p e r o x o n i t r a t e ( l - ) is k n o w n to carry out reactions that are injurious to the cell: it initiates lipid
peroxidation, 5 and it reacts rapidly with sulfhydryls6
and methionine. 7 It may also play a role in reperfusion
injury, where ischaemic e n d o t h e l i u m produces superoxide and nitric oxide when blood flow is reestablished. 8 In addition, oxoperoxonitrate(1-) nitrates and
hydroxylates aromatic c o m p o u n d s , 9A° a reaction that is
catalyzed by metal complexes ~1 and C u / Z n superoxide
INTRODUCTION
There are a n u m b e r of studies that indicate that peroxynitrite [O = N O O - , oxoperoxonitrate(1-)] is formed in
vivo. Ischiropoulos et a l ) have demonstrated that
oxoperoxonitrate(1-) is formed from nitrogen m o n o x ide and superoxide near activated macrophages. The
rate constant for this r a d i c a l - r a d i c a l c o m b i n a t i o n react i o n - 6 . 7 × 109 M - I s -l as determined by flash photoly s i s 2 - - i s close to diffusion controlled. U n d e r conditions where the concentrations of nitrogen m o n o x i d e
and superoxide dismutase are comparable, this fast rate
would allow nitrogen m o n o x i d e to compete for superoxide with superoxide dismutase. The latter reacts catalytically with superoxide at a rate of 2.4 × 109
M - l s - l . 3 The concentration of nitrogen m o n o x i d e pro-
Address correspondence to: W. H. Koppenol at the Institut ftir
Anorganische Chemie, Eidgenrssische Technische Hochschule, CH
8092 ZUrich, Switzerland.
85
86
D. BARTLETTet aL
dismutase. 12 The toxicity of oxoperoxonitrate(l-) may
play a role in the familial form of amyotrophic lateral
sclerosis, ~3 where defects on chromosome 21 have
been linked to superoxide dismutase. ~4'~5
Oxoperoxonitrate(1-) is relatively stable in basic solution (pH > 12). At physiological pH, however, oxoperoxonitrate(1-) becomes hydronated to give hydrogen oxoperoxonitrate which isomerizes to nitrate.
ONOOH = H ÷ + ONOO-
ONOOH ~ NO3
pK~ = 6.8 (Ref. 16)
(1)
k2 = 1.3 s -t at 25°C (Ref. 16)
(2)
+ H÷
Compared to the anion, hydrogen oxoperoxonitrate
is more reactive, ~7 and it may be responsible for the
cytotoxic reactions attributed to oxoperoxonitrate(1-).
These harmful reactions of oxoperoxonitrate(l-) fall
into three categories: (1) reactions that are first order
in oxoperoxonitrate(1-) and zero order in substrate.
The hydroxylation and nitration of aromatic compounds as well as the oxidation of ferrocyanide (D.
Bartlett and W. H. Koppenol, unpublished) belong to
this class. (2) Oxidation during a bimolecular reaction,
such as the reaction of sulfhydryls with oxoperoxonitrate(1-); 6 (3) Metal-catalyzed nitrations of aromatic
compounds. Such reactions are first order in oxoperoxonitrate(1-) and metal complex, but zero order in
the aromatic compound (Ramezanian, Padmaja, Bartlett, and Koppenol, unpublished). Koppenol et alJ 6
have shown that formation of free hydroxyl radicals
is thermodynamically not feasible, and have suggested
that a distorted transhydrogen oxoperoxonitrate may
be the oxidizing species. Experimental data, from
which the thermodynamic and kinetic factors governing the multiple reaction pathways of oxoperoxonitrate(1-) can be discerned, are still scanty.
It is clear from the above that the cytotoxic potential
of oxoperoxonitrate(1-) is quite high. The question that
arises is: has evolution afforded a defense mechanism?
Scavenging by antioxidants such as ascorbic acid and
thiols is one possible approach toward mitigating these
potentially harmful reactions. Ascorbic acid (vitamin
C), abundant in nature, is an important antioxidant that
is present in vivo in humans at concentrations ranging
from micromolar to millimolar, depending on tissue.
Herein we describe the reaction of ascorbic acid with
oxoperoxonitrate(1-) and suggest that normal physiological concentrations of ascorbic acid are insufficient
for scavenging oxoperoxonitrate(l-).
MATERIALS AND METHODS
Oxoperoxonitrate(1-) was prepared by the autoxidation of hydroxylamine in 0.5 M NaOH solution containing 100 #M dtpaJ ~ All other solutions also contained 100 # M dtpa. Chemicals used were of reagent
grade purity. Water was purified by reverse osmosis,
deionization, and filtration (MarCor Medical). Ascorbate solutions were freshly prepared just prior to use.
Spectra were recorded on a Beckman DU-7HS recording spectrophotometer. The stoichiometry of the
reaction was determined by recording the decrease in
the absorbance of ascorbate solutions near pH 7 at
252 nm (c = 8900 M -j cm-I). Oxoperoxonitrate(1-)
concentrations are reported on the basis of an extinction coefficient of 1670 M-lcm-~. ~9 Kinetic measurements were made on an OLIS stopped-flow instrument
equipped with an OLIS RSM 1000 rapid scanning
monochromator that collects 1,000 spectra per second
from 280 nm to 430 nm. Data collected over the entire
wavelength range were analyzed with the OLIS-RSM
1000-GLOBAL FIT software to yield pseudo-first-order rate constants. The kinetics of the reaction were
studied over the range of acidities, pH 4 - 7 in acetate
and phosphate buffers at 10, 20, 25, 30, and 37°C.
Ionic strength was maintained at 0.5 M with sodium
chloride. Pseudo-first-order rate constants reported are
the average of at least six separate experiments. Temperatures were maintained to within _+ 0.1°C with a
VWR model 1160 circulating water bath.
Electron spin resonance (ESR) spectra were obtained with a Varian E-109 instrument interfaced to
a computer for data acquisition and postrun analysis,
operating at X-band frequencies, and equipped with a
T M l l 0 microwave cavity. The instrument settings
used in these experiments were: gain = 4 × 103; modulation amplitude = 0. l0 mT; center field = 335.1 mT;
sweep width = 4.0 mT; time constant = 1 s; sweep
time = 240 s; microwave power = 20 mW. These
experiments were carried out with a quartz flow-mixing cell mounted in the TM1 l0 cavity. The temperature
of the solutions and the cell were constant at approximately 2 I°C throughout the experiment. A reciprocating pump with a stainless steel head attached to the
outlet of the flow-mixing cell was used to draw approximately equal amounts of the reagents through the
flow-mixing cell at a total flow rate of approximately
7.0 ml per minute.
RESULTS
Stoichiometry
At pH 7 in the presence of 0.25 mM ascorbate the
absorbance changes at 252 nm resulting from the addi-
L-Ascorbic acid oxidation
10
Or)
0
0102
0104
0106
0108
0.10
([ Ascorbate ]T) / M
87
shown in Fig. 2. From pH 4.0, the rate increases to a
maximum near pH 5.8 and then decreases steadily as
the pH is increased, clearly indicating that two protonation equilibria are important. The low rates of reaction
below pH 4 and above pH 7 indicate that the interaction
between the species AH2 and ONOOH, and that between the charged species AH- and ONOO-, are insignificant. For the present, we will assume that the reaction takes place between the most abundant species
near pH 5.8, the pH of the maximum in Fig. 2, namely
AH- and ONOOH (see the Discussion section).
Fig. 1. Variation of pseudo-first-order rate constants for the reaction
of ascorbic acid with oxoperoxonitrate(l-) at temperature, T = 25°C
and pH = 4.0, [ascorbate]r = 0.01 - 0.07 M, I = 0.50 M (NaCI).
AH- + ONOOH -'-' Products
(4)
A rate law to express the observed rate data for the
competitive kinetics is given in Eq. 5:
tion of various concentrations (25-200/~M) oxoperoxonitrate(1-) corresponded to a stoichiometry of 1.03
mol of ascorbate per mol of oxoperoxonitrate(1-)
added.
-d[ONOO-]
dt
kE[H+]
[ONOO-]
KI + [H +1
k4Ka[H +]
[AH2][ONOO-]
(K, + [H÷])(K3 + [H÷])
+
(5)
Total ascorbate variation
At pH 4.0 and 25°C, the decay of oxoperoxonitrate(l) in the presence of excess ascorbic acid is
clearly a bimolecular process, first order in both oxidant and reductant, A plot of pseudo-first-order rate
constant (l%b,) as a function of the total concentration
of all ascorbic acid species, [ascorbate]T, is linear, as
shown in Fig. 1. The intercept of the extrapolated line
with the Y-axis, 1.65 s -t, corresponds to the spontaneous decomposition of hydrogen oxoperoxonitrate in
acidic solution and compares well with the previously
obtained value of 1.3 s -1 at 2 5 ° C ) 6
By accounting for the decomposition reaction under
the given conditions, a corrected pseudo-first-order rate
constant, k'obs Can be derived (Eq. 6):
kaK3[H+][AH2]
(Kl + [H+])(K3 + [H+])
k~,s =
(6)
A graphical fit of the proposed mechanistic scheme
400
v
"7
300
•e
200
pH variation and mechanistic fit
Ascorbic acid, AH2, dissociates in aqueous solution
to give an acid solution containing the monohydrogenascorbate anion, AH-:
O
tO
.~
_4'
AH2 = AH- + H ÷
p K , = 4 . 0 4 ( R e f . 20)
(3)
100
0
3.0
The second acid dissociation constant of ascorbic
acid, pK = 11.3, need not be considered over the range
of hydrogen ion concentrations studied. The variation
of the apparent rate of reaction with pH at 25°C is
i
i
4.0
I
i
t
5.0
i
6.0
t
i
7.0
L
8.0
pH
Fig. 2. Variation of apparent second-order rate constant with pH
for the reaction of oxoperoxonitrate(1-) with ascorbic acid at 25°(2,
[ONOO-] = 1.0-8.0 mM, [ascorbate]r = 0.010-0.15 M, I = 0.50
M (NaCI).
88
D. BARTLETTet al.
20
3000
16
)
• •
25O0
"
~2000
12
~ 1500
.0
--~ 8
1000
4
500
0
0.00
0.02
0.04
0.06
K3[H"][AH2]
(K1 +[H*])(K3+[H+])
0.08
Temperature dependence
The reaction between ascorbic acid and oxoperoxonitrate(l-) was studied at 10, 20, 25, 30, and 37°C;
second-order rate constants for the reaction at these
temperatures are 100, 192, 235, 324, and 569 M-~s -t,
respectively. A n enthalpy o f activation, AH* = 9.3 _+
0.5 kcal/mol and entropy o f activation, AS* = - 1 6 _+
2 cal/(K-mol) for the reaction between ascorbic acid
and oxoperoxonitrate(1-) is derived from the Eyring
plot 2t shown in Fig. 4.
-185
-190
R -195
-200
c
rr
-205
-210
i
0.00320
i
i
i
0.00330
i
0.00340
i
335.1
336.1
Field (mT)
is shown in Fig. 3 - - a plot of k'ob~ as a function o f
K3[H+][AH2]/(KI +[H+])(K3+[H+]) at 25°C. From this
plot, a second-order rate constant of 235 __+ 4 M - t s -t
for k4 is derived.
.lz
0
334.1
/M
Fig. 3. Plot of corrected pseudo-first-order rate constant, k'ob~against
[H+I/(K~ +[H+]) × K3[AH2]/(K3+ [H+]) for the reaction of AH
+ ONOOH (data in Fig. 2). From the slope a second-order rate
constant of 235 +_ 4 M ts ~ is calculated.
V
E
Y
(B)
I
0.00350
r
J
0.00360
1/T (K q)
Fig. 4. Eyring plot for the reaction of ascorbic acid with oxoperoxonitrate(l-) at pH 4-7.5, [ONOO-] = 0.08-8.0 mM and [ascorbate]a= 0.010-0.15 M. Temperature variation rate data are (TPC, k/
M Js J): 10, 100; 20, 192; 25, 235; 30, 324; 37, 569.
Fig. 5. Detection of the ascorbyl radical at pH 5. (A) Two-line signal
obtained when ascorbic acid and oxoperoxonitrate(l-) are mixed at
pH 5 to give solution 1 mM in each reactant. (B) Background,
solution same as (A) without the addition of oxoperoxonitrate(l-).
ESR studies
A sharp two-line E S R spectrum is observed within
100 ms of mixing solutions o f oxoperoxonitrate(1-)
and ascorbic acid at final concentrations of 1.0 m M
each, at pH 5 (Fig. 5). This signal is 35-fold more
intense than that obtained from ascorbic acid alone.
The two lines are separated by 0.18 mT, with a g-value
o f 2.005, clearly indicating formation o f the ascorbyl
radical. 22
DISCUSSION
Due to the instability o f the hydronated form, there
have only been limited investigations L6'tL23 into the
reactivity o f oxoperoxonitrate(1-). The second-order
rate constants reported for reactions o f oxoperoxonitrate(1-) with cysteine and glutathione are on the order
o f ( 2 - 5 ) × 103 M - t s -j at ca. pH 7 and 37°C 6 with
activation energies of ca. 10 kcal/mol, t6 The rate constant for the reaction presented here is smaller by an
order o f magnitude, but the enthalpy of activation is
similar, which suggests that a c o m m o n mechanism
m a y be operative.
We have assumed that hydrogen oxoperoxonitrate
oxidizes the m o n o h y d r o g e n ascorbate anion. On the
basis of the kinetic evidence, we cannot rule out that
the reaction occurs between O N O O - and AH2. This
reaction is kinetically indistinguishable from the reaction between O N O O H and A H , and both reactions
may be occurring simultaneously. Considering the
pKas o f the reactant species (4.04 and 6.8 for AH2 and
O N O O H , respectively) both A H - and O N O O H are
present in much higher concentrations than are AH2
L-Ascorbic acid oxidation
and ONOO- at or near the pH optimum, and are, therefore, more likely to be the reactive species in this reaction. Furthermore, ONOOH is likely to be a better
electron acceptor than the negatively charged ONOO .
The ESR experiment shows the presence of the ascorbyl radical, which suggests that hydrogen oxoperoxonitrate oxidizes monohydrogen ascorbate via a
one-electron transfer mechanism. In principle, the
oxidation can take place by a one- or two-electron
mechanism, because both reduction potentials, E°'(O NOOH, 2H+/NO2 ", H20) and E°(ONOOH, 2H+/NO2-,
H20), 1.4 V and 0.99 V, respectively, ~6are higher than
that of the ascorbyl/monohydrogenascorbate couple,
0.28 V, and the dehydroascorbate/monohydrogen
ascorbate couple, 0.054 V. 24 The one-electron oxidation with the ascorbyl radical and nitrogen dioxide as
products is the simpler reaction, and the most exergonic per electron. A more complicated mechanism
would involve two-electron oxidation of monohydrogen ascorbate to dehydroascorbate, followed by comproportionation with unreacted monohydrogen ascorbate to yield two ascorbyl radicals. Considering that
the comproportionation reaction is unfavorable by 10.5
kcal (for 2 mol of ascorbyl radical formed), the more
complicated mechanism does not seem likely. One
could also argue that hydrogen oxoperoxonitrate undergoes homolysis during the isomerization to nitrate,
and that the hydroxyl radical formed in this reaction
oxidizes the monohydrogen ascorbate. There are three
arguments against that pathway. The first two are based
on thermodynamic considerations? 6 The hydroxyl radical reacts with nitrogen dioxide to form hydrogen
oxoperoxonitrate, not nitric acid, 25-27 Reaction 7. The
reason for this is that the unpaired electron of nitrogen
dioxide is delocalized over its two oxygens.
NO~ + HO" = ONOOH
(7)
The rate for the forward reaction, kT, has recently
been redetermined to be 4.5 × 10 9 M - i s - t , 27 and the
Gibbs energy change is - 2 1 _+ 2 kcal/mol) 6 From
these data, one calculates via A G ° = -RTln(kT/k -7)
that the rate of homolysis, k - 7 lies between 10-6S -I tO
10 SS ~, which is at least 10 6 times slower than the
rate of isomerization, k2! The second thermodynamic
argument is based on the experimentally determined
entropy of activation of the isomerization reaction.
This entropy of activation is rather small, inconsistent
with homolysis. Instead, the small value is suggestive
of a "tight" transition state, presumably trans-hydrogen oxoperoxonitrate with a reduced N - O - O bond
89
angle and a lengthened O - O bond. ~6 The third argument is a kinetic one. A reaction that involves the
hydroxyl radical would be first order in hydrogen
oxoperoxonitrate, and zero order in ascorbate. The reactions of ascorbic acid and its anion with the hydroxyl
radical to yield the ascorbyl radical are known to be
very fast, 1.0 x 10 m M ~s-~, and are not pH dependent. 28 Because the decomposition of peroxynitrite is
accelerated in the presence of and, indeed, dependent
on the concentration of ascorbate, it is not possible to
account for the observed second-order behavior and the
pH dependence if the hydroxyl radical were tormed.
As to the mechanism of the reaction, we propose
that the reaction involves the monohydrogen ascorbic
acid species, AH- transferring an electron to a vacant
orbital on the O - O bond in hydrogen oxoperoxonitrate, ONOOH. This mechanism can be represented by
the following reaction:
AH- + ONOOH ~ NO2 + A" + H 2 0
(4)
The reaction stoichiometry indicates that ascorbate
is also oxidized in a nonrate-determining manner by
the decaying trans hydrogen oxoperoxonitrate species,
probably giving the same products as in Reaction 4.
In the presence of excess ascorbate, the stoichiometry
of the reaction is consistent with the follt)w-up reactions:
NO~ + AH- --. NO~ + A'2A"
+ H +--+A + AH
(8)
(9)
One might ask: What was the concentration of the
ascorbyl radical during the ESR experiment? An estimate can be made by equating the rate of formation
of the of the ascorbyl radical in Reactions 4 and 8,
and in the first-order decay of peroxynitrite, with the
rate of diproportionation of the ascorbyl radical in Reaction 9. After elimination of [NO2"] we obtain Eq. 10:
dIA" ]
dt
-
=
0
=
k2[ONOOH]
+ 2k4[AH-][ONOOH] - 2 k d A ' - ] 2
(10)
A second-order rate constant of 1.5 x l 0 7 M-~s ~
for the decay of the ascorbyl radical (kg) at pH 5 can
be calculated from the literature. 29 The rate of formation for the ascorbyl radical under the conditions of
the ESR experiment described herein [1.0 mM each
ascorbic acid and oxoperoxonitrate(l-) mixed at pH 5]
is 1.7 x 10 -3 M-~s -~. Thus, the steady-state approxi-
90
_
D. BARTLETFet
CH20H
CH20H
I
I
_
, .,
I
,
,~o
÷ H,o÷
o
,~
.o;
.o
H. O,,~HI
~,,/
CH=OH
I
O ~CH
o
0
\
Scheme 1.
mation that resulted in Eq. 10 leads to a maximum
concentration of the ascorbyl radical of ca. 7 #M.
The mechanism of the reaction of monohydrogen
ascorbate with hydrogen oxoperoxonitrate may involve
the formation of a hydrogen-bonded complex shown
in Scheme 1. It is possible to visualize a transition
state in which the rate determining electron transfer
reaction takes place in this intermediate followed by
the break up of the complex to give A'-, NO2 °,
and H20.
Some support for equilibrium formation of this
loosely bonded intermediate is given by the observation of a moderately large negative value for the entropy of activation. This implies an activated complex
that is more ordered than its precursors, typically signifying a bonded complex. Clearly, I~q for the formation
of the hydrogen-bonded complex is small because saturation is not observed.
An alternative mechanism for the reaction involves
nucleophilic attack of A H - and ONOOH to form a
short-lived covalently bonded ascorbyl nitrosyl peroxide intermediate with displacement of O H - (see
Scheme 2).
This intermediate decomposes to give AH °, which
has a pica of -0.4522 and would dissociate to give A'-,
H ÷, and NO2". The mechanism in Scheme 1 gives rise
to an ascorbyl radical where the unpaired electron is
located on the C-2 oxygen, whereas in the mechanism
in Scheme 2, the unpaired electron would reside on C3 oxygen. Due to resonance, the radical may be located
on those oxygen atoms as well as on the oxygen atom
attached to C-1. ESR evidence indicates that the ascorbyl radical is O-centered. It is somewhat curious that
the reaction is so slow, and it is conceivable that the
neatly hydrogen-bonded complex shown in Scheme
1 is actually nonproductive: a small displacement of
aL
hydrogen oxoperoxonitrate may be necessary for the
reaction to proceed, as shown in Scheme 2. The small
enthalpy of activation obtained for this reaction also
indicates that the bimolecular reaction involves the c i s hydrogen oxoperoxonitrate isomer.
Can ascorbate be a scavenger of oxoperoxonitrate(1-) under physiological conditions? The small
rate constant for the reaction of hydrogen oxoperoxonitrate with monohydrogen ascorbate indicates that the
former cannot be scavenged by the latter under physiological conditions. We will consider three reactions:
(1) the isomerization to nitrate, (2) the reaction with
monohydrogen ascorbate, and (3) the reaction with
sulfhydryls. As follows from Eq. 11, the question is
whether the product of the bimolecular rate constant,
kb and the concentration of reductant is larger than k~,
the rate constant for the isomerization to nitrate.
Rate of decomposition(s- ~)
= k~ + kb[Reductant]
(11)
Near pH 7.4 only 20% of oxoperoxonitrate(1-) is
hydronated. Ascorbic acid is fully in the monohydrogen
ascorbate, AH -t form. Thus, the effective bimolecular
rate constant, kb, equals 235/5 M-~s -~ or 47 M-~s -~.
Monohydrogen ascorbate concentrations between 4 0 110 #M are found in human serum, 3° while estimated 3j
concentrations in tissues are higher, between 190/zM
in heart and 19 mM in the adrenal gland, with 1.6 mM
in liver and brain. At a median concentration of ca. 0.5
mM, one calculates a rate of disappearance of oxoper-
CH2OH
I
o/CH
.0
CH2OH
I
0
..
0
O--N
C~
0
\
H
\0
o
o
o
I
I
\
HI + O H -
H
CH2OH
I
o/CH
O
o
0
//
O
o®
+ NO~ + H20
Scheme 2.
N=O
L-Ascorbic acid oxidation
oxonitrate(1-) via the b i m o l e c u l a r m e c h a n i s m (kb[AH I)
o f 0.025 s - ~. T h i s reaction c o m p e t e s with the first-order
i s o m e r i z a t i o n to nitrate, lq = 0.26 s -~ a n d with the
reaction with sulfhydryl c o m p o u n d s . If w e a s s u m e an
overall rate c o n s t a n t o f 4 × l03 M - I s -j for that reaction, 6
and a cytosolic c o n c e n t r a t i o n o f 5 m M o f glutathione
a n d cysteine then oxoperoxonitrate(1-) disappears at a
rate o f 10 s ~. A l t h o u g h these n u m b e r s are estimates,
a n d apply to 25°C, the c o n c l u s i o n is clear: ascorbate
c a n n o t play a direct role in the d e f e n s e against oxoperoxonitrate(1-). Instead, it is likely that sulfur c o m p o u n d s
do protect. I n d e p e n d e n t l y , Radi also c o n c l u d e d that the
b i m o l e c u l a r reaction o f oxoperoxonitrate(1-), with sulfh y d r y l s at p h y s i o l o g i c a l concentrations, results in a
m u c h faster d i s a p p e a r a n c e o f oxoperoxonitrate(1-) than
the i s o m e r i z a t i o n reaction. 32 S o m e peroxidases react
with rate constants o n the order o f l06 to 107 M - I s -I
with h y d r o g e n oxoperoxonitrate. 33 H o w e v e r , not every
cell c o n t a i n s such e n z y m e s . T h e reaction o f oxoperoxon i t r a t e ( l - ) with glutathione results in the f o r m a t i o n of
a thiy134 radical that reacts with another glutathione to
f o r m a G S S G ' - radical; this strongly r e d u c i n g radical
reacts with o x y g e n to form superoxide, w h i c h is subseq u e n t l y s c a v e n g e d b y superoxide dismutase, 34'35 or reacts with n i t r o g e n m o n o x i d e to f o r m o x o p e r o x o n i trate(1-). Alternatively, GSSG" m a y react with n i t r o g e n
m o n o x i d e to yield oxonitrate(1-), N O - , w h i c h can also
react with d i o x y g e n to yield o x o p e r o x o n i t r a t e ( 1 - ) ) 6
Acknowledgements - - This work was supported by the National
Institutes of Health (GM48829) and the Council for Tobacco Research, Inc. We are grateful to Dr. J. S. Beckman for helpful discussions.
REFERENCES
1. Ischiropoulos, H.; Zhu, L.; Beckman, J. S. Peroxynitrite formation from macrophage-derived nitric oxide. Arch. Biochem. Biophys. 298:446-451; 1992.
2, Huie, R. E.; Padmaja, S. Reaction of "NO with 02"-. Free Radic.
Res. Commun. 18:195-199; 1993.
3, Fielden, E. M.; Roberts, P. B.; Bray, R. C.; Lowe, D. J.; Mautner,
G. N.; Rotilio, G.; Calabrese, L. The mechanism of action of
superoxide dismutase from pulse radiolysis and electron paramagnetic resonance. Biochem. J. 139:49-60; 1974.
4. Malinski, T.; Taha, Z. Nitric oxide release from a single cell
measured in situ by a porphyrinic-based microsensor. Nature
358:676-678; 1992.
5. Radi, R.; Beckman, J. S.; Bush, K. M.; Freeman, B, A. Peroxynitrite-induced membrane lipid peroxidation: The cytotoxic potential of superoxide and nitric oxide. Arch. Biochem. Biophys.
288:481-487; 1991.
6. Radi, R.; Beckman, J. S.; Bush, K. M.; Freeman, B. A. Peroxynitrite oxidation of sulfhydryls. J. Biol. Chem. 266:4244-4250;
1991.
7. Moreno, J. J.; Pryor, W. A. Inactivation ofa-I-proteinase inhibitor by peroxynitrite. Chem. Res. ToxicoL 5:425-431; 1992.
8. Beckman, J, S. lschaemic injury mediator. Nature 345:27-28;
1990.
91
9. Trifonow, I. Anwendung der Persalpetersaure for Analytische
Zwecke. Z. Anorg. Chem. 124:136-139; 1922.
10. Halfpenny, E.; Robinson, P. L. The nitration and hydroxylation
of aromatic compounds by pernitrous acid. J. Chem. Soc. 939946; 1952.
I 1. Beckman, J. S.; Ischiropoulos, H.; Zhu, L.; van der Woerd, M.;
Smith, C. D.; Chen, J.; Harrison, J.; Martin, J. C.; Tsai, M.
Kinetics of superoxide dismutase- and iron-catalyzed nitration
of phenolics by peroxynitrite. Arch. Biochem. Biophys.
298:438-445; 1992.
12. lschiropoulos, H.; Zhu, L.; Chen, J.; Tsai, M.; Martin, J. C.;
Smith, C. D.; Beckman, J. S. Peroxynitrite-mediated tyrosine
nitration catalyzed by superoxide dismutase. Arch. Biochem.
Biophys. 298:431-437; 1992.
13. Beckman, J. S.; Carson, M.; Smith, C. D.; Koppenol, W. H.
ALS, SOD and peroxynitrite. Nature 364:584; 1993.
14. Rosen, D. R.; Siddique, T.; Patterson, D.; Figlewicz, D. A.;
Sapp, P.; Hentati, A.; Donaldson, D.; Goto, J.; O'Regan, J. P.;
Deng, H.-X.; Rahmani, Z.; Krizus, A.; McKenna-Yasek, D.;
Cayabyab, A.; Gaston, S. M.; Berger, R.; Tanzi, R. E.; Halperin,
J. J.; Herzfeldt, B.; van den Bergh, R.; Hung, W.-Y.; Bird, T.;
Deng, G.; Mulder, D. W.; Smyth, C.; Laing, N. G.; Soriano, E.;
Pericak-Vance, M. A.; Haines, J.; Rouleau, G. A.; Gusella, J. S.;
Horvitz, H. R.; Brown, R. H., Jr. Mutations in Cu/Zn superoxide
dismutase gene are associated with familial amyotrophic lateral
sclerosis. Nature 362:59-62; 1993.
15. Deng, H.-X.; Hentati, A.; Tainer, J. A.; Iqbal, Z.; Cayabyab,
A.; Hung, W,-Y.; Getzoff, E. D.; Hu, P.; Herzfeldt, B.; Roos,
R. P.; Warner, C.; Deng, G.; Soriano, E.; Smyth, C.; Parge,
H. E.; Ahmed, A.; Roses, A. D.; Halleweli, R. A.; PericakVance, M. A.; Siddique, T. Amyotrophic lateral sclerosis and
structural defects in Cu,Zn superoxide dismutase. Science
261:1047 - 1051 ; 1993,
16. Koppenol, W. H.; Moreno, J. J.; Pryor, W. A.; Ischiropoulos,
H.; Beckman, J. S. Peroxynitrite, a cloaked oxidant formed by
nitric oxide and superoxide. Chem. Res. Toxicol. 5:834-842;
1992.
17. Edwards, J. O.; Plumb, R. C. The chemistry of peroxonitrites.
Prog. lnorg. Chem. 41:599-635; 1993.
18. Hughes, M. N.; Nicklin, H G. Autoxidation of hydroxylamine
in alkaline solution. 3". Chem. Soc. (A) 164-168; 1971.
19. Benton, D. J.; Moore, P. Kinetics and mechanism of the formation and decay of peroxynitrous acid in perchloric acid solutions.
J. Chem. Soc. (A) 3179-3182; 1970.
20. Taqui Khan, M. M.; Martell, A, E. The kinetics of the reaction
of iron (III) chelates of aminopolycarboxylic acids with ascorbic
acid. J. Am. Chem. Soc. 90:3386-3389; 1968.
21. Atkins, P. Physical chemistD'. Oxford, UK: Oxford University
Press; 1990:858.
22. Laroff, G. P.; Fessenden, R. W.; Schuler, R. H. The electron
spin resonance spectra of radical intermediates in the oxidation
of ascorbic acid and related substances. J. Am. Chem. Soc.
94:9062-9073; 1972.
23. Crow, J. P.; Spruell, C.; Chen, J.; Gunn, C.; lschiropoulos, H.;
Tsai, M.; Smith, C. D.; Radi, R.; Koppenol, W. H.; Beckman,
J. S. On the pH-dependent yield of hydroxyl radical products
from peroxynitrite. Free Radic. Biol. Med. 16:331-338; 1994.
24. Williams, M. H.; Yandell, J. K. Outer-sphere electron transfer
reactions of ascorbate anions. Aust. J. Chem. 35:1133-1144;
1982.
25, Gr~itzel, M.; Henglein, A.; Taniguchi, S. Pulsradiolytische Beobachtungen fiber die Reduktion des NO3 -Ions und tiber die
Bildung und Zerfall der persalpetrigen S~iure in wassriger LOsung. Ber. Bunsenges. Physik. Chem. 94:292-298; 1970.
26. Barat, F.; Gilles, L.; Hickel, B.; Sutton, J. Flash photolysis of
the nitrate ion in aqueous solution: Excitation at 200 rim. J.
Chem. Soc. (A) 982-1986; 1970.
27. Logager, T.; Sehested, K. Formation and decay of peroxynitrous
acid: A pulse radiolysis study. J. Phys. Chem. 97:6664-6669;
1993.
92
D. BARTLETTet aL
28. Ross, A. B.; Mallard, W. G.; Heiman, W. P.; Bielski, B. H. J.;
Buxton, G. V.; Cabelli, D. E.; Greenstock, C. L.; Huie,
R. E.; Neta, P. NDRL-N1ST solution kinetics database: Ver. I.
Gaithersburg, MD: National Institute of Standards and Technology; 1992.
29. Cabelli, D. E.; Bielski, B. H. J. Kinetics and mechanism for the
oxidation of ascorbic acid/ascorbate by hydroperoxyl/superoxide radicals. A pulse radiolysis and stopped flow photolysis
study. J. Phys. Chem. 87:1809-1812; 1983.
30. Margolis, S. A., Davis, T. Stabilization of ascorbic acid in human plasma, and its liquid-chromatographic measurement. Clin.
Chem. 34:2217-2223; 1988.
31. Behrens, W. A.; Mad~re, R. A highly sensitive high-performance liquid chromatography method for the estimation of
ascorbic and dehydroascorbic acid in tissues, biological fluids
and foods. Anal. Biochem. 165:102-107; 1987.
32. Radi, R. Oxidative reactions of peroxynitrite in biological systems: Direct attack vs. the hydroxyl radical-like pathway. In:
Davies, K. J. A., ed. The oxygen paradox. Padova: CLEUP
Press; (in press).
33. Floris, R.; Piersma, S. R.; Yang, G.; Jones, P.; Wever, R. Interac-
tion of myeloperoxidase with peroxynitrite--A comparison with
lactoperoxidase, horseradish peroxidase and catalase. Eur. Z
Biochem. 215:767-775; 1993.
34. Ohara, A.; Gatti, R. M.; Radi, R. Spin-trapping studies of peroxynitrite decomposition and of 3-morpholinosydnonimine NEthylcarbamide autoxidation: Direct evidence for metal-independent formation of free radical intermediates. Arch. Biochem.
Biophys. 310:118-125; 1994.
35. Winterbourn, C. C. Superoxide as an intracellular radical sink.
Free Radic. Biol. Med. 14:85-90; 1992.
36. Koppenol, W. H. A thermodynamic appraisal of the radical sink
theory. Free Radic. BioL Med. 14:91-94; 1993.
ABBREVIATIONS
A--dehydroascorbate
A ' - - - a s c o r b y l radical a n i o n
dtpa--diethylenetriamine-N,N' ,N'"-pentaacetate
H2A--ascorbic acid
HA---monohydrogen
ascorbate anion
Copyright © 1994 Elsevier Science Ltd
Printed in the USA. All rights reserved
0891-5849/95 $9.50 + .00
Pergamon
0891-5849(94)E0133-2
Original Contribution
THE KINETICS OF THE OXIDATION OF L-ASCORBIC ACID
BY PEROXYNITRITE
DELLAND BARTLETT,* DANIEL F. CHURCH,* PATRICIA L. BOUNDS,* and W. H. KOPPENOL§*
Departments of *Chemistry and *Biochemistry and ~BiodynamicsInstitute, Louisiana State University, Baton Rouge, LA,
USA; and *Department of Chemistry and Physics, Southeastern Louisiana State University, Hammond, LA, USA
(Received 7 March 1994; Revised 3 June 1994; Accepted 6 June 1994)
Abstract Peroxynitrite [O=NOO-, oxoperoxonitrate(l-)] is a strong oxidant that may be formed in vivo by the reaction of
02"- and NO'. Oxoperoxonitrate(1-) reacts with molecules in aqueous acidic solutions via pathways that involve the highly
reactive hydrogen oxoperoxonitrate either as an intermediate in a first-order reaction or as a reactive agent in a simple secondorder reaction. ESR experiments show that hydrogen oxoperoxonitrate oxidizes monohydrogen L-ascorbate by one electron:
when mixed at pH ca. 5 and passed through a flow cell within 0.1 s, the two-line ESR signal of the ascorbyl radical anion (all
= 0.18 T, g = 2.005) is observed. The overall stoichiometry of the reaction was 1 mol of ascorbate oxidized per mol of
oxoperoxonitrate(l-) added. The kinetics of the reaction were studied over the pH range 4.0-7.5 by stopped-flow spectrometry.
Hydrogen oxoperoxonitrate,observed between 300 and 350 urn, and the oxoperoxonitrate(l-) anion, at 302 nm, disappear faster
than predicted for the first-order isomerization to NO3 . The rate increases from pH 4 to 5.8, and then decreases with increasing
pH. The rate variation suggests a bimolecular reaction either between the oxoperoxonitrate(I-) anion and ascorbic acid or between
hydrogen oxoperoxonitrate and the monohydrogen ascorbate anion. Although the two pathways are kinetically indistinguishable,
the pKa values of ascorbic acid and hydrogen oxoperoxonitrate strongly suggest that the reacting species are hydrogen oxoperoxonitrate and monohydrogen ascorbate. The second-order rate constant for this reaction is 235 + 4 M ts-t at 250(7. The enthalpy
and entropy of activation are AH* = 9.3 _+ 0.5 kcal/mol and AS* = -16 _+ 2 cal/(mol.K), respectively.
Keywords---Ascorbate, Peroxynitrite, Oxoperoxonitrate(I-), Sulfhydryl, Antioxidant defense, Reaction kinetics, Stopped-flow
spectrophotometry, Electron spin resonance
duced by neurons and endothelial cells normally is of
the order of tens to hundreds nanomolar, m u c h less
than the 5 to 40 # M concentration of superoxide dismutase in cells. However, u n d e r pathologic conditions,
higher levels of nitric oxide synthase can be induced.
A concentration of 0.5 # M nitrogen m o n o x i d e has been
measured 4 near the surface of an endothelial cell. Near
activated macrophages concentrations as high as 10
/zM appear likely (J. S. Beckman, personal c o m m u n i cation). Thus, even in the presence of superoxide dismutase, oxoperoxonitrate(1-) can be formed.
O x o p e r o x o n i t r a t e ( l - ) is k n o w n to carry out reactions that are injurious to the cell: it initiates lipid
peroxidation, 5 and it reacts rapidly with sulfhydryls6
and methionine. 7 It may also play a role in reperfusion
injury, where ischaemic e n d o t h e l i u m produces superoxide and nitric oxide when blood flow is reestablished. 8 In addition, oxoperoxonitrate(1-) nitrates and
hydroxylates aromatic c o m p o u n d s , 9A° a reaction that is
catalyzed by metal complexes ~1 and C u / Z n superoxide
INTRODUCTION
There are a n u m b e r of studies that indicate that peroxynitrite [O = N O O - , oxoperoxonitrate(1-)] is formed in
vivo. Ischiropoulos et a l ) have demonstrated that
oxoperoxonitrate(1-) is formed from nitrogen m o n o x ide and superoxide near activated macrophages. The
rate constant for this r a d i c a l - r a d i c a l c o m b i n a t i o n react i o n - 6 . 7 × 109 M - I s -l as determined by flash photoly s i s 2 - - i s close to diffusion controlled. U n d e r conditions where the concentrations of nitrogen m o n o x i d e
and superoxide dismutase are comparable, this fast rate
would allow nitrogen m o n o x i d e to compete for superoxide with superoxide dismutase. The latter reacts catalytically with superoxide at a rate of 2.4 × 109
M - l s - l . 3 The concentration of nitrogen m o n o x i d e pro-
Address correspondence to: W. H. Koppenol at the Institut ftir
Anorganische Chemie, Eidgenrssische Technische Hochschule, CH
8092 ZUrich, Switzerland.
85
86
D. BARTLETTet aL
dismutase. 12 The toxicity of oxoperoxonitrate(l-) may
play a role in the familial form of amyotrophic lateral
sclerosis, ~3 where defects on chromosome 21 have
been linked to superoxide dismutase. ~4'~5
Oxoperoxonitrate(1-) is relatively stable in basic solution (pH > 12). At physiological pH, however, oxoperoxonitrate(1-) becomes hydronated to give hydrogen oxoperoxonitrate which isomerizes to nitrate.
ONOOH = H ÷ + ONOO-
ONOOH ~ NO3
pK~ = 6.8 (Ref. 16)
(1)
k2 = 1.3 s -t at 25°C (Ref. 16)
(2)
+ H÷
Compared to the anion, hydrogen oxoperoxonitrate
is more reactive, ~7 and it may be responsible for the
cytotoxic reactions attributed to oxoperoxonitrate(1-).
These harmful reactions of oxoperoxonitrate(l-) fall
into three categories: (1) reactions that are first order
in oxoperoxonitrate(1-) and zero order in substrate.
The hydroxylation and nitration of aromatic compounds as well as the oxidation of ferrocyanide (D.
Bartlett and W. H. Koppenol, unpublished) belong to
this class. (2) Oxidation during a bimolecular reaction,
such as the reaction of sulfhydryls with oxoperoxonitrate(1-); 6 (3) Metal-catalyzed nitrations of aromatic
compounds. Such reactions are first order in oxoperoxonitrate(1-) and metal complex, but zero order in
the aromatic compound (Ramezanian, Padmaja, Bartlett, and Koppenol, unpublished). Koppenol et alJ 6
have shown that formation of free hydroxyl radicals
is thermodynamically not feasible, and have suggested
that a distorted transhydrogen oxoperoxonitrate may
be the oxidizing species. Experimental data, from
which the thermodynamic and kinetic factors governing the multiple reaction pathways of oxoperoxonitrate(1-) can be discerned, are still scanty.
It is clear from the above that the cytotoxic potential
of oxoperoxonitrate(1-) is quite high. The question that
arises is: has evolution afforded a defense mechanism?
Scavenging by antioxidants such as ascorbic acid and
thiols is one possible approach toward mitigating these
potentially harmful reactions. Ascorbic acid (vitamin
C), abundant in nature, is an important antioxidant that
is present in vivo in humans at concentrations ranging
from micromolar to millimolar, depending on tissue.
Herein we describe the reaction of ascorbic acid with
oxoperoxonitrate(1-) and suggest that normal physiological concentrations of ascorbic acid are insufficient
for scavenging oxoperoxonitrate(l-).
MATERIALS AND METHODS
Oxoperoxonitrate(1-) was prepared by the autoxidation of hydroxylamine in 0.5 M NaOH solution containing 100 #M dtpaJ ~ All other solutions also contained 100 # M dtpa. Chemicals used were of reagent
grade purity. Water was purified by reverse osmosis,
deionization, and filtration (MarCor Medical). Ascorbate solutions were freshly prepared just prior to use.
Spectra were recorded on a Beckman DU-7HS recording spectrophotometer. The stoichiometry of the
reaction was determined by recording the decrease in
the absorbance of ascorbate solutions near pH 7 at
252 nm (c = 8900 M -j cm-I). Oxoperoxonitrate(1-)
concentrations are reported on the basis of an extinction coefficient of 1670 M-lcm-~. ~9 Kinetic measurements were made on an OLIS stopped-flow instrument
equipped with an OLIS RSM 1000 rapid scanning
monochromator that collects 1,000 spectra per second
from 280 nm to 430 nm. Data collected over the entire
wavelength range were analyzed with the OLIS-RSM
1000-GLOBAL FIT software to yield pseudo-first-order rate constants. The kinetics of the reaction were
studied over the range of acidities, pH 4 - 7 in acetate
and phosphate buffers at 10, 20, 25, 30, and 37°C.
Ionic strength was maintained at 0.5 M with sodium
chloride. Pseudo-first-order rate constants reported are
the average of at least six separate experiments. Temperatures were maintained to within _+ 0.1°C with a
VWR model 1160 circulating water bath.
Electron spin resonance (ESR) spectra were obtained with a Varian E-109 instrument interfaced to
a computer for data acquisition and postrun analysis,
operating at X-band frequencies, and equipped with a
T M l l 0 microwave cavity. The instrument settings
used in these experiments were: gain = 4 × 103; modulation amplitude = 0. l0 mT; center field = 335.1 mT;
sweep width = 4.0 mT; time constant = 1 s; sweep
time = 240 s; microwave power = 20 mW. These
experiments were carried out with a quartz flow-mixing cell mounted in the TM1 l0 cavity. The temperature
of the solutions and the cell were constant at approximately 2 I°C throughout the experiment. A reciprocating pump with a stainless steel head attached to the
outlet of the flow-mixing cell was used to draw approximately equal amounts of the reagents through the
flow-mixing cell at a total flow rate of approximately
7.0 ml per minute.
RESULTS
Stoichiometry
At pH 7 in the presence of 0.25 mM ascorbate the
absorbance changes at 252 nm resulting from the addi-
L-Ascorbic acid oxidation
10
Or)
0
0102
0104
0106
0108
0.10
([ Ascorbate ]T) / M
87
shown in Fig. 2. From pH 4.0, the rate increases to a
maximum near pH 5.8 and then decreases steadily as
the pH is increased, clearly indicating that two protonation equilibria are important. The low rates of reaction
below pH 4 and above pH 7 indicate that the interaction
between the species AH2 and ONOOH, and that between the charged species AH- and ONOO-, are insignificant. For the present, we will assume that the reaction takes place between the most abundant species
near pH 5.8, the pH of the maximum in Fig. 2, namely
AH- and ONOOH (see the Discussion section).
Fig. 1. Variation of pseudo-first-order rate constants for the reaction
of ascorbic acid with oxoperoxonitrate(l-) at temperature, T = 25°C
and pH = 4.0, [ascorbate]r = 0.01 - 0.07 M, I = 0.50 M (NaCI).
AH- + ONOOH -'-' Products
(4)
A rate law to express the observed rate data for the
competitive kinetics is given in Eq. 5:
tion of various concentrations (25-200/~M) oxoperoxonitrate(1-) corresponded to a stoichiometry of 1.03
mol of ascorbate per mol of oxoperoxonitrate(1-)
added.
-d[ONOO-]
dt
kE[H+]
[ONOO-]
KI + [H +1
k4Ka[H +]
[AH2][ONOO-]
(K, + [H÷])(K3 + [H÷])
+
(5)
Total ascorbate variation
At pH 4.0 and 25°C, the decay of oxoperoxonitrate(l) in the presence of excess ascorbic acid is
clearly a bimolecular process, first order in both oxidant and reductant, A plot of pseudo-first-order rate
constant (l%b,) as a function of the total concentration
of all ascorbic acid species, [ascorbate]T, is linear, as
shown in Fig. 1. The intercept of the extrapolated line
with the Y-axis, 1.65 s -t, corresponds to the spontaneous decomposition of hydrogen oxoperoxonitrate in
acidic solution and compares well with the previously
obtained value of 1.3 s -1 at 2 5 ° C ) 6
By accounting for the decomposition reaction under
the given conditions, a corrected pseudo-first-order rate
constant, k'obs Can be derived (Eq. 6):
kaK3[H+][AH2]
(Kl + [H+])(K3 + [H+])
k~,s =
(6)
A graphical fit of the proposed mechanistic scheme
400
v
"7
300
•e
200
pH variation and mechanistic fit
Ascorbic acid, AH2, dissociates in aqueous solution
to give an acid solution containing the monohydrogenascorbate anion, AH-:
O
tO
.~
_4'
AH2 = AH- + H ÷
p K , = 4 . 0 4 ( R e f . 20)
(3)
100
0
3.0
The second acid dissociation constant of ascorbic
acid, pK = 11.3, need not be considered over the range
of hydrogen ion concentrations studied. The variation
of the apparent rate of reaction with pH at 25°C is
i
i
4.0
I
i
t
5.0
i
6.0
t
i
7.0
L
8.0
pH
Fig. 2. Variation of apparent second-order rate constant with pH
for the reaction of oxoperoxonitrate(1-) with ascorbic acid at 25°(2,
[ONOO-] = 1.0-8.0 mM, [ascorbate]r = 0.010-0.15 M, I = 0.50
M (NaCI).
88
D. BARTLETTet al.
20
3000
16
)
• •
25O0
"
~2000
12
~ 1500
.0
--~ 8
1000
4
500
0
0.00
0.02
0.04
0.06
K3[H"][AH2]
(K1 +[H*])(K3+[H+])
0.08
Temperature dependence
The reaction between ascorbic acid and oxoperoxonitrate(l-) was studied at 10, 20, 25, 30, and 37°C;
second-order rate constants for the reaction at these
temperatures are 100, 192, 235, 324, and 569 M-~s -t,
respectively. A n enthalpy o f activation, AH* = 9.3 _+
0.5 kcal/mol and entropy o f activation, AS* = - 1 6 _+
2 cal/(K-mol) for the reaction between ascorbic acid
and oxoperoxonitrate(1-) is derived from the Eyring
plot 2t shown in Fig. 4.
-185
-190
R -195
-200
c
rr
-205
-210
i
0.00320
i
i
i
0.00330
i
0.00340
i
335.1
336.1
Field (mT)
is shown in Fig. 3 - - a plot of k'ob~ as a function o f
K3[H+][AH2]/(KI +[H+])(K3+[H+]) at 25°C. From this
plot, a second-order rate constant of 235 __+ 4 M - t s -t
for k4 is derived.
.lz
0
334.1
/M
Fig. 3. Plot of corrected pseudo-first-order rate constant, k'ob~against
[H+I/(K~ +[H+]) × K3[AH2]/(K3+ [H+]) for the reaction of AH
+ ONOOH (data in Fig. 2). From the slope a second-order rate
constant of 235 +_ 4 M ts ~ is calculated.
V
E
Y
(B)
I
0.00350
r
J
0.00360
1/T (K q)
Fig. 4. Eyring plot for the reaction of ascorbic acid with oxoperoxonitrate(l-) at pH 4-7.5, [ONOO-] = 0.08-8.0 mM and [ascorbate]a= 0.010-0.15 M. Temperature variation rate data are (TPC, k/
M Js J): 10, 100; 20, 192; 25, 235; 30, 324; 37, 569.
Fig. 5. Detection of the ascorbyl radical at pH 5. (A) Two-line signal
obtained when ascorbic acid and oxoperoxonitrate(l-) are mixed at
pH 5 to give solution 1 mM in each reactant. (B) Background,
solution same as (A) without the addition of oxoperoxonitrate(l-).
ESR studies
A sharp two-line E S R spectrum is observed within
100 ms of mixing solutions o f oxoperoxonitrate(1-)
and ascorbic acid at final concentrations of 1.0 m M
each, at pH 5 (Fig. 5). This signal is 35-fold more
intense than that obtained from ascorbic acid alone.
The two lines are separated by 0.18 mT, with a g-value
o f 2.005, clearly indicating formation o f the ascorbyl
radical. 22
DISCUSSION
Due to the instability o f the hydronated form, there
have only been limited investigations L6'tL23 into the
reactivity o f oxoperoxonitrate(1-). The second-order
rate constants reported for reactions o f oxoperoxonitrate(1-) with cysteine and glutathione are on the order
o f ( 2 - 5 ) × 103 M - t s -j at ca. pH 7 and 37°C 6 with
activation energies of ca. 10 kcal/mol, t6 The rate constant for the reaction presented here is smaller by an
order o f magnitude, but the enthalpy of activation is
similar, which suggests that a c o m m o n mechanism
m a y be operative.
We have assumed that hydrogen oxoperoxonitrate
oxidizes the m o n o h y d r o g e n ascorbate anion. On the
basis of the kinetic evidence, we cannot rule out that
the reaction occurs between O N O O - and AH2. This
reaction is kinetically indistinguishable from the reaction between O N O O H and A H , and both reactions
may be occurring simultaneously. Considering the
pKas o f the reactant species (4.04 and 6.8 for AH2 and
O N O O H , respectively) both A H - and O N O O H are
present in much higher concentrations than are AH2
L-Ascorbic acid oxidation
and ONOO- at or near the pH optimum, and are, therefore, more likely to be the reactive species in this reaction. Furthermore, ONOOH is likely to be a better
electron acceptor than the negatively charged ONOO .
The ESR experiment shows the presence of the ascorbyl radical, which suggests that hydrogen oxoperoxonitrate oxidizes monohydrogen ascorbate via a
one-electron transfer mechanism. In principle, the
oxidation can take place by a one- or two-electron
mechanism, because both reduction potentials, E°'(O NOOH, 2H+/NO2 ", H20) and E°(ONOOH, 2H+/NO2-,
H20), 1.4 V and 0.99 V, respectively, ~6are higher than
that of the ascorbyl/monohydrogenascorbate couple,
0.28 V, and the dehydroascorbate/monohydrogen
ascorbate couple, 0.054 V. 24 The one-electron oxidation with the ascorbyl radical and nitrogen dioxide as
products is the simpler reaction, and the most exergonic per electron. A more complicated mechanism
would involve two-electron oxidation of monohydrogen ascorbate to dehydroascorbate, followed by comproportionation with unreacted monohydrogen ascorbate to yield two ascorbyl radicals. Considering that
the comproportionation reaction is unfavorable by 10.5
kcal (for 2 mol of ascorbyl radical formed), the more
complicated mechanism does not seem likely. One
could also argue that hydrogen oxoperoxonitrate undergoes homolysis during the isomerization to nitrate,
and that the hydroxyl radical formed in this reaction
oxidizes the monohydrogen ascorbate. There are three
arguments against that pathway. The first two are based
on thermodynamic considerations? 6 The hydroxyl radical reacts with nitrogen dioxide to form hydrogen
oxoperoxonitrate, not nitric acid, 25-27 Reaction 7. The
reason for this is that the unpaired electron of nitrogen
dioxide is delocalized over its two oxygens.
NO~ + HO" = ONOOH
(7)
The rate for the forward reaction, kT, has recently
been redetermined to be 4.5 × 10 9 M - i s - t , 27 and the
Gibbs energy change is - 2 1 _+ 2 kcal/mol) 6 From
these data, one calculates via A G ° = -RTln(kT/k -7)
that the rate of homolysis, k - 7 lies between 10-6S -I tO
10 SS ~, which is at least 10 6 times slower than the
rate of isomerization, k2! The second thermodynamic
argument is based on the experimentally determined
entropy of activation of the isomerization reaction.
This entropy of activation is rather small, inconsistent
with homolysis. Instead, the small value is suggestive
of a "tight" transition state, presumably trans-hydrogen oxoperoxonitrate with a reduced N - O - O bond
89
angle and a lengthened O - O bond. ~6 The third argument is a kinetic one. A reaction that involves the
hydroxyl radical would be first order in hydrogen
oxoperoxonitrate, and zero order in ascorbate. The reactions of ascorbic acid and its anion with the hydroxyl
radical to yield the ascorbyl radical are known to be
very fast, 1.0 x 10 m M ~s-~, and are not pH dependent. 28 Because the decomposition of peroxynitrite is
accelerated in the presence of and, indeed, dependent
on the concentration of ascorbate, it is not possible to
account for the observed second-order behavior and the
pH dependence if the hydroxyl radical were tormed.
As to the mechanism of the reaction, we propose
that the reaction involves the monohydrogen ascorbic
acid species, AH- transferring an electron to a vacant
orbital on the O - O bond in hydrogen oxoperoxonitrate, ONOOH. This mechanism can be represented by
the following reaction:
AH- + ONOOH ~ NO2 + A" + H 2 0
(4)
The reaction stoichiometry indicates that ascorbate
is also oxidized in a nonrate-determining manner by
the decaying trans hydrogen oxoperoxonitrate species,
probably giving the same products as in Reaction 4.
In the presence of excess ascorbate, the stoichiometry
of the reaction is consistent with the follt)w-up reactions:
NO~ + AH- --. NO~ + A'2A"
+ H +--+A + AH
(8)
(9)
One might ask: What was the concentration of the
ascorbyl radical during the ESR experiment? An estimate can be made by equating the rate of formation
of the of the ascorbyl radical in Reactions 4 and 8,
and in the first-order decay of peroxynitrite, with the
rate of diproportionation of the ascorbyl radical in Reaction 9. After elimination of [NO2"] we obtain Eq. 10:
dIA" ]
dt
-
=
0
=
k2[ONOOH]
+ 2k4[AH-][ONOOH] - 2 k d A ' - ] 2
(10)
A second-order rate constant of 1.5 x l 0 7 M-~s ~
for the decay of the ascorbyl radical (kg) at pH 5 can
be calculated from the literature. 29 The rate of formation for the ascorbyl radical under the conditions of
the ESR experiment described herein [1.0 mM each
ascorbic acid and oxoperoxonitrate(l-) mixed at pH 5]
is 1.7 x 10 -3 M-~s -~. Thus, the steady-state approxi-
90
_
D. BARTLETFet
CH20H
CH20H
I
I
_
, .,
I
,
,~o
÷ H,o÷
o
,~
.o;
.o
H. O,,~HI
~,,/
CH=OH
I
O ~CH
o
0
\
Scheme 1.
mation that resulted in Eq. 10 leads to a maximum
concentration of the ascorbyl radical of ca. 7 #M.
The mechanism of the reaction of monohydrogen
ascorbate with hydrogen oxoperoxonitrate may involve
the formation of a hydrogen-bonded complex shown
in Scheme 1. It is possible to visualize a transition
state in which the rate determining electron transfer
reaction takes place in this intermediate followed by
the break up of the complex to give A'-, NO2 °,
and H20.
Some support for equilibrium formation of this
loosely bonded intermediate is given by the observation of a moderately large negative value for the entropy of activation. This implies an activated complex
that is more ordered than its precursors, typically signifying a bonded complex. Clearly, I~q for the formation
of the hydrogen-bonded complex is small because saturation is not observed.
An alternative mechanism for the reaction involves
nucleophilic attack of A H - and ONOOH to form a
short-lived covalently bonded ascorbyl nitrosyl peroxide intermediate with displacement of O H - (see
Scheme 2).
This intermediate decomposes to give AH °, which
has a pica of -0.4522 and would dissociate to give A'-,
H ÷, and NO2". The mechanism in Scheme 1 gives rise
to an ascorbyl radical where the unpaired electron is
located on the C-2 oxygen, whereas in the mechanism
in Scheme 2, the unpaired electron would reside on C3 oxygen. Due to resonance, the radical may be located
on those oxygen atoms as well as on the oxygen atom
attached to C-1. ESR evidence indicates that the ascorbyl radical is O-centered. It is somewhat curious that
the reaction is so slow, and it is conceivable that the
neatly hydrogen-bonded complex shown in Scheme
1 is actually nonproductive: a small displacement of
aL
hydrogen oxoperoxonitrate may be necessary for the
reaction to proceed, as shown in Scheme 2. The small
enthalpy of activation obtained for this reaction also
indicates that the bimolecular reaction involves the c i s hydrogen oxoperoxonitrate isomer.
Can ascorbate be a scavenger of oxoperoxonitrate(1-) under physiological conditions? The small
rate constant for the reaction of hydrogen oxoperoxonitrate with monohydrogen ascorbate indicates that the
former cannot be scavenged by the latter under physiological conditions. We will consider three reactions:
(1) the isomerization to nitrate, (2) the reaction with
monohydrogen ascorbate, and (3) the reaction with
sulfhydryls. As follows from Eq. 11, the question is
whether the product of the bimolecular rate constant,
kb and the concentration of reductant is larger than k~,
the rate constant for the isomerization to nitrate.
Rate of decomposition(s- ~)
= k~ + kb[Reductant]
(11)
Near pH 7.4 only 20% of oxoperoxonitrate(1-) is
hydronated. Ascorbic acid is fully in the monohydrogen
ascorbate, AH -t form. Thus, the effective bimolecular
rate constant, kb, equals 235/5 M-~s -~ or 47 M-~s -~.
Monohydrogen ascorbate concentrations between 4 0 110 #M are found in human serum, 3° while estimated 3j
concentrations in tissues are higher, between 190/zM
in heart and 19 mM in the adrenal gland, with 1.6 mM
in liver and brain. At a median concentration of ca. 0.5
mM, one calculates a rate of disappearance of oxoper-
CH2OH
I
o/CH
.0
CH2OH
I
0
..
0
O--N
C~
0
\
H
\0
o
o
o
I
I
\
HI + O H -
H
CH2OH
I
o/CH
O
o
0
//
O
o®
+ NO~ + H20
Scheme 2.
N=O
L-Ascorbic acid oxidation
oxonitrate(1-) via the b i m o l e c u l a r m e c h a n i s m (kb[AH I)
o f 0.025 s - ~. T h i s reaction c o m p e t e s with the first-order
i s o m e r i z a t i o n to nitrate, lq = 0.26 s -~ a n d with the
reaction with sulfhydryl c o m p o u n d s . If w e a s s u m e an
overall rate c o n s t a n t o f 4 × l03 M - I s -j for that reaction, 6
and a cytosolic c o n c e n t r a t i o n o f 5 m M o f glutathione
a n d cysteine then oxoperoxonitrate(1-) disappears at a
rate o f 10 s ~. A l t h o u g h these n u m b e r s are estimates,
a n d apply to 25°C, the c o n c l u s i o n is clear: ascorbate
c a n n o t play a direct role in the d e f e n s e against oxoperoxonitrate(1-). Instead, it is likely that sulfur c o m p o u n d s
do protect. I n d e p e n d e n t l y , Radi also c o n c l u d e d that the
b i m o l e c u l a r reaction o f oxoperoxonitrate(1-), with sulfh y d r y l s at p h y s i o l o g i c a l concentrations, results in a
m u c h faster d i s a p p e a r a n c e o f oxoperoxonitrate(1-) than
the i s o m e r i z a t i o n reaction. 32 S o m e peroxidases react
with rate constants o n the order o f l06 to 107 M - I s -I
with h y d r o g e n oxoperoxonitrate. 33 H o w e v e r , not every
cell c o n t a i n s such e n z y m e s . T h e reaction o f oxoperoxon i t r a t e ( l - ) with glutathione results in the f o r m a t i o n of
a thiy134 radical that reacts with another glutathione to
f o r m a G S S G ' - radical; this strongly r e d u c i n g radical
reacts with o x y g e n to form superoxide, w h i c h is subseq u e n t l y s c a v e n g e d b y superoxide dismutase, 34'35 or reacts with n i t r o g e n m o n o x i d e to f o r m o x o p e r o x o n i trate(1-). Alternatively, GSSG" m a y react with n i t r o g e n
m o n o x i d e to yield oxonitrate(1-), N O - , w h i c h can also
react with d i o x y g e n to yield o x o p e r o x o n i t r a t e ( 1 - ) ) 6
Acknowledgements - - This work was supported by the National
Institutes of Health (GM48829) and the Council for Tobacco Research, Inc. We are grateful to Dr. J. S. Beckman for helpful discussions.
REFERENCES
1. Ischiropoulos, H.; Zhu, L.; Beckman, J. S. Peroxynitrite formation from macrophage-derived nitric oxide. Arch. Biochem. Biophys. 298:446-451; 1992.
2, Huie, R. E.; Padmaja, S. Reaction of "NO with 02"-. Free Radic.
Res. Commun. 18:195-199; 1993.
3, Fielden, E. M.; Roberts, P. B.; Bray, R. C.; Lowe, D. J.; Mautner,
G. N.; Rotilio, G.; Calabrese, L. The mechanism of action of
superoxide dismutase from pulse radiolysis and electron paramagnetic resonance. Biochem. J. 139:49-60; 1974.
4. Malinski, T.; Taha, Z. Nitric oxide release from a single cell
measured in situ by a porphyrinic-based microsensor. Nature
358:676-678; 1992.
5. Radi, R.; Beckman, J. S.; Bush, K. M.; Freeman, B, A. Peroxynitrite-induced membrane lipid peroxidation: The cytotoxic potential of superoxide and nitric oxide. Arch. Biochem. Biophys.
288:481-487; 1991.
6. Radi, R.; Beckman, J. S.; Bush, K. M.; Freeman, B. A. Peroxynitrite oxidation of sulfhydryls. J. Biol. Chem. 266:4244-4250;
1991.
7. Moreno, J. J.; Pryor, W. A. Inactivation ofa-I-proteinase inhibitor by peroxynitrite. Chem. Res. ToxicoL 5:425-431; 1992.
8. Beckman, J, S. lschaemic injury mediator. Nature 345:27-28;
1990.
91
9. Trifonow, I. Anwendung der Persalpetersaure for Analytische
Zwecke. Z. Anorg. Chem. 124:136-139; 1922.
10. Halfpenny, E.; Robinson, P. L. The nitration and hydroxylation
of aromatic compounds by pernitrous acid. J. Chem. Soc. 939946; 1952.
I 1. Beckman, J. S.; Ischiropoulos, H.; Zhu, L.; van der Woerd, M.;
Smith, C. D.; Chen, J.; Harrison, J.; Martin, J. C.; Tsai, M.
Kinetics of superoxide dismutase- and iron-catalyzed nitration
of phenolics by peroxynitrite. Arch. Biochem. Biophys.
298:438-445; 1992.
12. lschiropoulos, H.; Zhu, L.; Chen, J.; Tsai, M.; Martin, J. C.;
Smith, C. D.; Beckman, J. S. Peroxynitrite-mediated tyrosine
nitration catalyzed by superoxide dismutase. Arch. Biochem.
Biophys. 298:431-437; 1992.
13. Beckman, J. S.; Carson, M.; Smith, C. D.; Koppenol, W. H.
ALS, SOD and peroxynitrite. Nature 364:584; 1993.
14. Rosen, D. R.; Siddique, T.; Patterson, D.; Figlewicz, D. A.;
Sapp, P.; Hentati, A.; Donaldson, D.; Goto, J.; O'Regan, J. P.;
Deng, H.-X.; Rahmani, Z.; Krizus, A.; McKenna-Yasek, D.;
Cayabyab, A.; Gaston, S. M.; Berger, R.; Tanzi, R. E.; Halperin,
J. J.; Herzfeldt, B.; van den Bergh, R.; Hung, W.-Y.; Bird, T.;
Deng, G.; Mulder, D. W.; Smyth, C.; Laing, N. G.; Soriano, E.;
Pericak-Vance, M. A.; Haines, J.; Rouleau, G. A.; Gusella, J. S.;
Horvitz, H. R.; Brown, R. H., Jr. Mutations in Cu/Zn superoxide
dismutase gene are associated with familial amyotrophic lateral
sclerosis. Nature 362:59-62; 1993.
15. Deng, H.-X.; Hentati, A.; Tainer, J. A.; Iqbal, Z.; Cayabyab,
A.; Hung, W,-Y.; Getzoff, E. D.; Hu, P.; Herzfeldt, B.; Roos,
R. P.; Warner, C.; Deng, G.; Soriano, E.; Smyth, C.; Parge,
H. E.; Ahmed, A.; Roses, A. D.; Halleweli, R. A.; PericakVance, M. A.; Siddique, T. Amyotrophic lateral sclerosis and
structural defects in Cu,Zn superoxide dismutase. Science
261:1047 - 1051 ; 1993,
16. Koppenol, W. H.; Moreno, J. J.; Pryor, W. A.; Ischiropoulos,
H.; Beckman, J. S. Peroxynitrite, a cloaked oxidant formed by
nitric oxide and superoxide. Chem. Res. Toxicol. 5:834-842;
1992.
17. Edwards, J. O.; Plumb, R. C. The chemistry of peroxonitrites.
Prog. lnorg. Chem. 41:599-635; 1993.
18. Hughes, M. N.; Nicklin, H G. Autoxidation of hydroxylamine
in alkaline solution. 3". Chem. Soc. (A) 164-168; 1971.
19. Benton, D. J.; Moore, P. Kinetics and mechanism of the formation and decay of peroxynitrous acid in perchloric acid solutions.
J. Chem. Soc. (A) 3179-3182; 1970.
20. Taqui Khan, M. M.; Martell, A, E. The kinetics of the reaction
of iron (III) chelates of aminopolycarboxylic acids with ascorbic
acid. J. Am. Chem. Soc. 90:3386-3389; 1968.
21. Atkins, P. Physical chemistD'. Oxford, UK: Oxford University
Press; 1990:858.
22. Laroff, G. P.; Fessenden, R. W.; Schuler, R. H. The electron
spin resonance spectra of radical intermediates in the oxidation
of ascorbic acid and related substances. J. Am. Chem. Soc.
94:9062-9073; 1972.
23. Crow, J. P.; Spruell, C.; Chen, J.; Gunn, C.; lschiropoulos, H.;
Tsai, M.; Smith, C. D.; Radi, R.; Koppenol, W. H.; Beckman,
J. S. On the pH-dependent yield of hydroxyl radical products
from peroxynitrite. Free Radic. Biol. Med. 16:331-338; 1994.
24. Williams, M. H.; Yandell, J. K. Outer-sphere electron transfer
reactions of ascorbate anions. Aust. J. Chem. 35:1133-1144;
1982.
25, Gr~itzel, M.; Henglein, A.; Taniguchi, S. Pulsradiolytische Beobachtungen fiber die Reduktion des NO3 -Ions und tiber die
Bildung und Zerfall der persalpetrigen S~iure in wassriger LOsung. Ber. Bunsenges. Physik. Chem. 94:292-298; 1970.
26. Barat, F.; Gilles, L.; Hickel, B.; Sutton, J. Flash photolysis of
the nitrate ion in aqueous solution: Excitation at 200 rim. J.
Chem. Soc. (A) 982-1986; 1970.
27. Logager, T.; Sehested, K. Formation and decay of peroxynitrous
acid: A pulse radiolysis study. J. Phys. Chem. 97:6664-6669;
1993.
92
D. BARTLETTet aL
28. Ross, A. B.; Mallard, W. G.; Heiman, W. P.; Bielski, B. H. J.;
Buxton, G. V.; Cabelli, D. E.; Greenstock, C. L.; Huie,
R. E.; Neta, P. NDRL-N1ST solution kinetics database: Ver. I.
Gaithersburg, MD: National Institute of Standards and Technology; 1992.
29. Cabelli, D. E.; Bielski, B. H. J. Kinetics and mechanism for the
oxidation of ascorbic acid/ascorbate by hydroperoxyl/superoxide radicals. A pulse radiolysis and stopped flow photolysis
study. J. Phys. Chem. 87:1809-1812; 1983.
30. Margolis, S. A., Davis, T. Stabilization of ascorbic acid in human plasma, and its liquid-chromatographic measurement. Clin.
Chem. 34:2217-2223; 1988.
31. Behrens, W. A.; Mad~re, R. A highly sensitive high-performance liquid chromatography method for the estimation of
ascorbic and dehydroascorbic acid in tissues, biological fluids
and foods. Anal. Biochem. 165:102-107; 1987.
32. Radi, R. Oxidative reactions of peroxynitrite in biological systems: Direct attack vs. the hydroxyl radical-like pathway. In:
Davies, K. J. A., ed. The oxygen paradox. Padova: CLEUP
Press; (in press).
33. Floris, R.; Piersma, S. R.; Yang, G.; Jones, P.; Wever, R. Interac-
tion of myeloperoxidase with peroxynitrite--A comparison with
lactoperoxidase, horseradish peroxidase and catalase. Eur. Z
Biochem. 215:767-775; 1993.
34. Ohara, A.; Gatti, R. M.; Radi, R. Spin-trapping studies of peroxynitrite decomposition and of 3-morpholinosydnonimine NEthylcarbamide autoxidation: Direct evidence for metal-independent formation of free radical intermediates. Arch. Biochem.
Biophys. 310:118-125; 1994.
35. Winterbourn, C. C. Superoxide as an intracellular radical sink.
Free Radic. Biol. Med. 14:85-90; 1992.
36. Koppenol, W. H. A thermodynamic appraisal of the radical sink
theory. Free Radic. BioL Med. 14:91-94; 1993.
ABBREVIATIONS
A--dehydroascorbate
A ' - - - a s c o r b y l radical a n i o n
dtpa--diethylenetriamine-N,N' ,N'"-pentaacetate
H2A--ascorbic acid
HA---monohydrogen
ascorbate anion