Direct determination of the Gibbs' energy of formation of peroxynitrous acid

Article abstract of DOI:10.1021/ic025698r

The kinetics of decomposition of peroxynitrous acid (ONOOH) was investigated in the presence of 0.1-0.75 M HClO₄ and at a constant ionic strength. The decay rate of ONOOH decreased in the presence of H₂O₂, approaching a limiting value well below 75 mM H₂O₂. It also decreased in the presence of relatively low [HNO₂] but did not approach a lower limiting value, since ONOOH reacts directly with HNO₂. The latter reaction corresponds to a HNO₂- and H⁺-catalyzed isomerization of ONOOH to nitrate, and its third-order rate constant was determined to be 520 ± 30 M^-2 s^-1. The mechanism of formation of O₂NOOH from ONOOH in the presence of H₂0₂ was also scrutinized. The results demonstrated that in the presence of 0.1-0.75 M HClO₄ and 75 mM H₂O₂ the formation of O₂NOOH is insignificant. The most important finding in this work is the reversibility of the reaction ONOOH + H₂O ? HNO₂ + H₂O₂ and its equilibrium constant was determined to be (7.5 ± 0.4) × 10^-4 M. Using this value, the Gibbs' energy of formation of ONOOH was calculated to be 7.1 + ± 0.2 kcal/mol. This figure is in good agreement with the value determined previously from kinetic data using parameters for radicals formed during homolysis of peroxynitrite.

Full text of DOI:10.1021/ic025698r

Inorg. Chem. 2003, 42, 3796−3800
Direct Determination of the Gibbs’ Energy of Formation of
Peroxynitrous Acid
,†
Ga´bor Mere´nyi,* Johan Lind,† Gideon Czapski,‡ and Sara Goldstein‡
Department of Chemistry, Nuclear Chemistry, The Royal Institute of Technology,
S-10044 Stockholm 70, Sweden, and Department of Physical Chemistry,
The Hebrew UniVersity of Jerusalem, Jerusalem 91904, Israel
Received May 7, 2002
The kinetics of decomposition of peroxynitrous acid (ONOOH) was investigated in the presence of 0.1−0.75 M
HClO and at a constant ionic strength. The decay rate of ONOOH decreased in the presence of H O , approaching
4
2 2
a limiting value well below 75 mM H O . It also decreased in the presence of relatively low [HNO ] but did not
2
2
2
approach a lower limiting value, since ONOOH reacts directly with HNO . The latter reaction corresponds to a
2
+
HNO - and H -catalyzed isomerization of ONOOH to nitrate, and its third-order rate constant was determined to be
2
520 ± 30 M-2 s-1. The mechanism of formation of O NOOH from ONOOH in the presence of H O was also
2
2 2
scrutinized. The results demonstrated that in the presence of 0.1−0.75 M HClO and 75 mM H O the formation
4
2 2
of O NOOH is insignificant. The most important finding in this work is the reversibility of the reaction ONOOH +
2
H O h HNO + H O , and its equilibrium constant was determined to be (7.5 ± 0.4) × 10-4 M. Using this value,
2
2
2 2
the Gibbs’ energy of formation of ONOOH was calculated to be 7.1 ± 0.2 kcal/mol. This figure is in good agreement
with the value determined previously from kinetic data using parameters for radicals formed during homolysis of
peroxynitrite.
Introduction
for free radicals formed via homolysis of peroxynitrite, and
such parameters may not always be as accurately known as
those of nonradical species. Therefore, a direct determination
of the Gibbs’ energy of ONOOH avoiding free radicals
should be of value.
It has been known for some time that, in acidic media,
the decomposition of ONOOH is proton-catalyzed, and it
has been assumed that this reaction, just like the uncatalyzed
consumption of ONOOH, yields nitrate as the sole product.9-12
In the present work, we shall scrutinize the acidic decom-
position of ONOOH, which will lead to an accurate value
for the Gibbs’ energy of formation of ONOOH.
Peroxynitrite (ONOOH/ONOO-) is a powerful oxidant
that can form in biological systems in the reaction of nitric
oxide with superoxide radicals.1,2 Peroxynitrite is well-known
for its deleterious effects,3-5 which oftentimes derive from
the formation of free radicals during its decomposition.6
Accurate knowledge of thermodynamical data for peroxyni-
trite is necessary for quantitative assessment of its chemistry.
An important parameter is the Gibbs’ energy of ONOOH,
which has previously been determined to be 7.5 ( 0.4 kcal/
mol.7,8 However, this determination has involved parameters
* To whom all correspondence should be directed. E-mail: gm@
nuchem.kth.se.
Experimental Section
† The Royal Institute of Technology.
‡ The Hebrew University of Jerusalem.
Materials. All chemicals were of analytical grade and were used
as received. Solutions were prepared with deionized water that was
distilled and purified using a Milli-Q water purification system.
(1) Beckman, J. S.; Beckman, T. W.; Chen, J.; Marshall, P. A.; Freeman,
B. A. Proc. Natl. Acad. Sci. U.S.A. 1990, 87, 1620-1624.
(2) Pryor, W. A.; Squadrito, G. L. Am. J. Physiol. 1995, 268, L699-
L722.
(3) Beckman, J. S. Chem. Res. Toxicol. 1996, 9, 836-844.
(4) Hurst, J. K.; Lymar, S. V. Chem. Res. Toxicol. 1997, 10, 802-810.
(5) Ischiropoulos, H. Arch. Biochem. Biophys. 1998, 356, 1-11 and
references therein.
(6) Goldstein, S.; Czapski, G.; Lind J.; Mere´nyi, G. J. Biol. Chem. 2000,
275, 3031-3036.
(7) Mere´nyi, G.; Lind, J. Chem. Res. Toxicol. 1998, 11, 243-246.
(8) Mere´nyi, G.; Lind, J.; Goldstein, S.; Czapski, G. J. Phys. Chem. A
1999, 103, 5685-5691.
(9) Halfpenny, E.; Robinson, P. L. J. Chem. Soc. 1952, 928-938.
(10) Anbar, M.; Taube, H. J. Am. Chem. Soc. 1954, 76, 6243-6247.
(11) Benton, D. J.; Moore, P. J. Chem. Soc. A 1970, 3179-3182.
(12) Saha, A.; Goldstein, S.; Cabelli, D.; Czapski, G. Free Radical Biol.
Med. 1988, 23, 653-659.
3796 Inorganic Chemistry, Vol. 42, No. 12, 2003
10.1021/ic025698r CCC: $25.00 © 2003 American Chemical Society
Published on Web 05/08/2003

Gibbs’ Energy of Formation of Peroxynitrous Acid
Figure 2. Observed decay rate constant of 0.21 mM ONOOH in the
presence of residual H2O2 and HNO2 (A) and 50 mM added H2O2 (B) at I
) 0.75 M and 25 °C.
Figure 1. Decay of 0.22 mM ONOOH in the presence of residual H2O2
and HNO2 (ca. 15 µM each) (a) and added 50 mM H2O2 (b). All solutions
contained 0.5 M HClO4 and 0.25 M NaClO4 at 25 °C. The trace in the
presence of added H2O2 was corrected against a blank without ONOOH.
at 0.5 M H+ without and with 50 mM added H2O2 are given
in Figure 1. The decay of ONOOH in the presence of residual
HNO2 and H2O2 could not be fitted to a single exponential
decay, but to two consecutive first-order reactions, where
the contribution of the second reaction was less than 8%.
These observations will be discussed in following paragraphs.
In the presence of 50 mM H2O2 added, a small fast build-up
of the absorption was observed, which subsequently decayed
via a first-order reaction. The small build-up of the absorption
is attributed to the formation of ONOOH via the reaction of
H2O2 with the residual HNO2,9-12 which is present in all
preparations. Figure 2 displays the observed first-order decay
rate constant of ONOOH as a function of [H+] in the
presence of residual H2O2 and HNO2 as well as with 50 mM
H2O2 added. These findings can be rationalized in terms of
reactions 1-3.
Peroxynitrite and peroxynitrate were synthesized through the
reaction of nitrite with acidified H2O2 using a quenched-flow with
a computerized syringe pump (World Precision Instruments model
SP 230IW) as described elsewhere.12,13 Briefly, peroxynitrite was
synthesized by mixing 0.60 M H2O2 in 0.7 M HClO4 with 0.63 or
0.60 M nitrite, and the mixture was quenched with 3 M NaOH at
room temperature. The stock solutions contained ca. 0.12 M
peroxynitrite. Starting with 0.63 M nitrite, the stock solution ended
up containing about 3% residual H2O2 and 12% residual nitrite.
When 0.60 M nitrite was used, the amounts of H2O2 and nitrite
remaining were about 7% each.12 The yield of peroxynitrite was
determined from its absorption at 302 nm using ꢀ ) 1670 M-1
cm-1 14
. Peroxynitric acid was prepared by mixing nitrite with excess
H2O2 in HClO4 at room temperature, and its concentration was
determined through its reaction with iodide at pH 4.4.13 The
concentration of H2O2 was determined using the Fricke method.
Methods. Stopped-flow kinetic measurements were carried out
using the Bio SX-17MV sequential stopped-flow from Applied
Photophysics with a 1-cm optical path. Peroxynitrite in 0.01 M
NaOH was mixed with HClO4 solutions at a 1:1 ratio. In some
experiments, the acidic solution also contained H2O2, while in others
the alkaline peroxynitrite solutions contained nitrite. In another set
of experiments, equal volumes of acidified H2O2 and nitrite were
mixed. A constant ionic strength of 0.75 M was maintained in all
experiments using NaClO4 in addition to HClO4. The formation
and decomposition of ONOOH was followed at 290 nm, where all
other species absorb negligibly. All experiments were carried out
at 25 °C.
ONOOH + H2O + H+ yk1z
k-1
HNO2 + H2O2 + H+ K1 ) k1/k-1 (1)
k2
9
ONOOH
8 NO3- + H+
(2)
(3)
k3
ONOOH + H+ 98 NO3- + 2H+
When the concentration of ONOOH is initially well below
1 mM and only residual H2O2 and HNO2 are present, reaction
-1 is negligible, and the rate law for the decay of ONOOH
simplifies to:
Simulations were carried using the GLINT global analysis
package (Applied Photophysics), and a noncommercial program
INTKIN written by Dr. H. A. Schwarz.
-d[ONOOH]/dt ) (k1[H+] + k2 + k3[H+])[ONOOH] (4)
At high [H2O2], reaction 1 is quenched; i.e., equilibrium
1 is completely shifted to the left, and therefore eq 4 reduces
to eq 5:
Results and Discussions
Effect of H2O2 on the Decay of ONOOH. The decay
rate of ONOOH in the presence of 0.1-0.75 M H+ was
found to decrease with increasing [H2O2], and to approach
a limiting lower value. For instance, at 0.5 M H+, the limiting
rate was reached above 20 mM H2O2. Typical kinetic traces
-d[ONOOH]/dt ) (k2 + k3[H+])[ONOOH]
(5)
We obtained k2 ) 1.3 ( 0.3 s-1 and (k1 + k3) ) 10.6 ( 0.6
M-1 s-1 from the intercept and slope of line A in Figure 2,
and from line B in Figure 2, k2 ) 1.1 ( 0.1 s-1 and k3 ) 4.3
( 0.1 M-1 s-1. Both intercepts agree with the well-known
(13) Goldstein, S.; Czapski, G.; Lind J.; Mere´nyi, G. Inorg. Chem. 1998,
37, 3943-3947.
(14) Hughes, M. N.; Nicklin, H. G. J. Chem. Soc. A 1968, 450-452.
Inorganic Chemistry, Vol. 42, No. 12, 2003 3797

Mere´nyi et al.
Table 1. Summary of All the Rate Constants Derived from the Present
Study at I ) 0.75 M'
no.
1
reaction
k
ONOOH + H+ f H2O2 + HNO2 + H+ 6.3 ( 0.7 M-1 s-1
-1 H2O2 + HNO2 + H+ f ONOOH + H+ (9.6 ( 1.0) × 103 M-2 s-1
2
3
9
ONOOH f NO3- + H+
1.2 ( 0.2 s-1
ONOOH + H+ f NO3- + 2H+
4.3 ( 0.1 M-1 s-1
ONOOH + HNO2 + H+
f
520 ( 30 M-2 s-1
HNO2 + NO3- + 2H+
Table 2. Observed Formation (kf) and Decay (kd) Rate Constants of
ONOOH under Limiting Concentrations of HNO2 or H2O2 in the
Presence of 0.5 M HClO4 and 0.25 M NaClO4
[H2O2], mM
[HNO2], mM
kf ,a s-1
kd, s-1
75
50
50
20
10
5
1
1
1
1
1
1
2.25
1
1
1
310 ( 20
210 ( 10
210 ( 10
104 ( 2
58 ( 1
b
3.5 ( 0.1
3.5 ( 0.1
3.5 ( 0.1
3.4 ( 0.1
3.3 ( 0.1
3.1 ( 0.1
4.4 ( 0.1
6.3 ( 0.1
9.5 ( 0.1
16.7 ( 0.2
Figure 3. Double reciprocal plot of the dependence of the observed rate
constant of the decay of ONOOH on [H2O2]. The data were taken from
Table 1.
5
b
10
20
50
56 ( 1
104 ( 3
210 ( 16
a The decay of the absorption followed first-order kinetics, and kd was
determined. The formation of the absorption could not be well separated
from its decay, and therefore, kf was determined by fitting the kinetic trace
to two consecutive first-order reactions, where the rate constant for the
second reaction was fixed to the observed kd. b Non-pseudo-first-order
conditions.
rate constant of neutral decomposition of ONOOH at 25
°C.2,8 The difference between the two limiting rate constants,
calculated from the slopes at high and residual H2O2, yields
k1 ) 6.3 ( 0.7 M-1 s-1. All the rate constants obtained in
the present study are compiled in Table 1.
Reaction of HNO2 with H2O2. Upon mixing HNO2 at
constant [H+] with varying amounts of excess H2O2, or vice
versa, the rapid build-up of the absorption at 290 nm is
followed by a slower decay. The data are presented in Table
2. The rate of formation of ONOOH was first order with
respect to [HNO2] and [H2O2], as long as the latter was kept
well below ca. 0.25 M. This behavior is in agreement with
earlier observations.9-12 In the presence of a sufficiently large
excess of H2O2 over HNO2 or vice versa, the rate of reaction
-1 becomes so much faster than that of reaction 1 that only
the former can be observed. Under these conditions, the rate
of formation of ONOOH is given by eq 6:
Figure 4. Observed decay rate constant of ONOOH vs [HNO2] under
limiting concentrations of H2O2 in the presence of 0.1 M HClO4 (b) and
0.5 M HClO4 (9). The ionic strength was 0.75 M.
(k1 + k-1[H2O2])[H+] . (k2 + k3[H+]), and hence, eq 7 is
obtained:
-d([HNO2] + [ONOOH])/dt ) (k2 + k3[H+])[ONOOH]
(7)
These are equilibrium conditions, which imply that [HNO2]
) K1[ONOOH]/[H2O2]. Consequently, eq 8 is obtained:
-d[ONOOH]/dt )
(k2 + k3[H+])([H2O2]/(K1 + [H2O2]))[ONOOH] (8)
d[ONOOH]/dt ) k-1[H+][H2O2][HNO2]
(6)
An inverse plot of kobs versus [H2O2] yields 1/intercept )
(k2 + k3[H+]) ) 3.5 ( 0.1 s-1 and slope/intercept ) K1 )
(7.5 ( 0.4) × 10-4 M (Figure 3). The former value is
virtually identical to the one obtained from Figure 2B. Most
importantly, the quantitative agreement between the directly
and indirectly determined K1 firmly settles its value.
While the decay rate of ONOOH approaches a limiting
value as [H2O2] increases, it keeps on increasing with
increasing [HNO2], when the latter is in excess. This is
clearly seen in Table 2 and Figure 4. The intercepts in this
figure yield (k2 + k3[H+]) ) 1.5 ( 0.2 and 3.6 ( 0.4 s-1 at
0.1 and 0.5 M H+, as expected, while the slopes correspond
to the rate constants for the direct reaction between ONOOH
and HNO2 at 0.1 and 0.5 M H+, their values being 42 ( 8
From the data in Table 2, one obtains k-1[H+] ) (4.8 (
0.5) × 103 M-1 s-1, and hence, k-1 ) (9.6 ( 1.0) × 103
M-2 s-1. We determined above k1 ) 6.3 ( 0.7 M-1 s-1 in
an independent kinetic measurement, and hence, k1/k-1 yields
the equilibrium constant K1 ) (6.6 ( 1.5) × 10-4 M.
Although it has been reported previously that the rate of
decay of ONOOH is independent of [H2O2],12 closer inspec-
tion of the data in Table 2 reveals that it increases slightly
but systematically as [H2O2] varies from 5 to 50 mM. From
this variation, the equilibrium constant K1 can be determined
directly as follows. At high [H2O2], the establishment of
equilibrium 1 is much more rapid than the decay of the
equilibrated mixture of HNO2 and ONOOH. This means that
3798 Inorganic Chemistry, Vol. 42, No. 12, 2003

Gibbs’ Energy of Formation of Peroxynitrous Acid
Finally, we note that, due to the reversibility of reaction
1, residual amounts of nitrite and H2O2 in stock solutions of
peroxynitrite cannot be avoided, not even when ONOOH is
synthesized from equal amounts of the reactants (see the
Experimental Section).
Mechanism and Yield of Formation of O2NOOH from
ONOOH in the Presence of H2O2. The most important
observation in the present study was that the decay rate of
ONOOH at low pH decreased with increasing [H2O2], and
we interpreted this finding in terms of the reversibility of
reaction 1, where both forward and reverse rates are being
catalyzed by [H+]. However, it is also known that O2NOOH
13,15,16
is formed from ONOOH in the presence of H2O2
and
is relatively stable in acidic solutions. Hence, one might argue
that its formation in our system could affect the kinetic
interpretation. To address this point, we synthesized peroxy-
nitrate through the reaction of nitrite with acidified H2O2.
The yield of O2NOOH increased with increasing [H2O2] and
[H+]. Figure 6 presents a double reciprocal plot of the yield
of O2NOOH versus [H2O2] obtained when 0.18 M HNO2
reacted with excess H2O2 in the presence of 1.65 M HClO4.
The figure demonstrates that a yield of 100% O2NOOH can
be obtained at infinite concentrations of H2O2. The present
results as well as the literature data13,15,16 suggest that the
formation of O2NOOH from ONOOH in the presence of
H2O2 takes place via the following mechanism.
Figure 5. Decay of 0.21 mM ONOOH in the presence of residual HNO2
and H2O2 (a), 0.75 mM added HNO2 (b), and 4.4 mM added HNO2 (c).
All solutions contained 0.5 M HClO4 and 0.25 M NaClO4 and were at 25
°C.
and 267 ( 16 M-1 s-1, respectively. This reaction is expected
to be similar to reaction 1, as ONOOH, like H2O2, is a
hydroperoxide. Hence, the reaction of ONOOH with HNO2
should form ONOONO, which decomposes fast to yield 2
•NO2, and the latter hydrolyses to yield nitrate and HNO2
(reactions 9-11).
k9
HNO2 + ONOOH + H+ 98 ONOONO + H2O + H+ (9)
(i) uncatalyzed pathway, where k12 + k13 ) k2 )
1.2 s-1 and k13/k12 ) 2.5:16
k10
9
ONOONO
8 2•NO2
(10)
ONOOH yk z •NO2 + •OH
(12)
12
k11
k-12
2•NO2 + H2O
9
8 HNO2 + NO3- + H+
(11)
k13
9
ONOOH
8 NO3- + H+
(13)
(14)
(15)
Thus, k9 ) 520 ( 30 M-2 s-1, and eqs 9-11 correspond
to a simultaneous HNO2- and H+-catalyzed isomerization
of ONOOH to nitrate.
k14
•OH + H2O2 f HO2 + H2O
•
k15
Decay of ONOOH When [HNO2] or [H2O2] Is Below
a Few Millimolar. When ONOOH decayed in the presence
of less than 10 mM H2O2 or HNO2, the decay obeyed two
consecutive first-order reactions, where the contribution of
the second reaction increased with increasing [H2O2] or
[HNO2]. Such a dual decay was also observed when no extra
H2O2 (or HNO2) was added (see Figure 1a). These observa-
tions can now be explained. The fast decay involves the
approach to equilibrium 1. Subsequently, the equilibrated
mixture of ONOOH, HNO2, and H2O2 decays to nitrate, and
this decay is monitored by way of the absorbance of the
ONOOH fraction. Although one can obtain analytical solu-
tions in closed form, these are not very useful, and thus,
simulation of the traces is to be preferred. Some typical
kinetic traces with excess [HNO2] are presented in Figure
5. These traces, as well as those with excess [H2O2], were
quantitatively simulated by means of rate constants k1, k-1,
k2, k3, and k9, taking into account the residual HNO2 and
H2O2 present in the stock solution of peroxynitrite. For
example, the modeling results in the presence of 0.75 mM
added HNO2 (Figure 5b) are presented in the Supporting
Information (Figure 5sup).
HO2 + •NO2
98 O2NOOH
•
(ii) acid-catalyzed pathway, where k16 ) k3 ) 4.3 M-1 s-1:
k16
ONOOH + H+ 98 NO2+ + H2O
(16)
k17
9
NO2+ + H2O
8 NO3- + 2H+ fast
(17)
(18)
k18
9
NO2+ + H2O2
8 O2NOOH + H+
From the slope of the line in Figure 6, we determined k17/
k18 ) 2.8 ( 0.2 M. Thus, at high pH and moderate
concentrations of H2O2, the maximum yield of O2NOOH
does not exceed ca. 28% of the initial ONOOH,16 but it can
approach 100% at low pH and very high [H2O2]. Neverthe-
less, under our experimental conditions, namely at 0.5 M
H+ and at the relatively low H2O2 concentrations of 50 -
75 mM, the acid-catalyzed pathway produces no more than
(15) Appelman, E. H.; Gosztola, D. J. Inorg. Chem. 1995, 34, 787.
(16) Hodges, G. R.; Ingold, K. U. J. J. Am. Chem. Soc. 1999, 121, 10695-
10701.
Inorganic Chemistry, Vol. 42, No. 12, 2003 3799

Mere´nyi et al.
these values and K1, we obtain ∆Gf°(ONOOH) ) (7.1 (
0.2) kcal/mol. Using pKa(ONOOH) ) 6.6 ( 0.1,21 we
calculate ∆Gf°(ONOO-) ) (16.1 ( 0.3) kcal/mol. This is
in good agreement with (16.6 ( 0.4) kcal/mol,7 the latter
having been determined using tetranitromethane as an
efficient O2•- scavenger, and employing the relevant radical
parameters.
Utilizing data extracted from a number of publica-
tions,2,8,21-24 we can calculate a very accurate equilibrium
constant for reaction 12, with K12 ) k12/k-12 ) (0.34 ( 0.08)/
(5 ( 1) × 109 ) (7 ( 3) × 10-11 M. By means of the well-
known value of ∆Gf°(•NO2) ) 15.1 kcal/mol,25 we obtain
∆G°f(•OH) ) (5.9 ( 0.3) kcal/mol. This figure is in excellent
agreement with ∆G°f(•OH) ) (6.0 ( 0.5) kcal/mol as
obtained by Schwarz and Dodson26 and is only slightly lower
than 6.4 ( 0.3 kcal/mol, as reported by Klaening et al.27
Figure 6. Dependence of the yield of O2NOOH on [H2O2] at low pH.
Double reciprocal plot of the yield of O2NOOH vs [H2O2] obtained when
0.18 M nitrite reacted with excess of H2O2 in the presence of 1.65 M HClO4
at room temperature.
Conclusions
2% O2NOOH. Here, the uncatalyzed pathway is the main
source of O2NOOH. However, not even this source can
produce more than {1.2/(4.3[H+] + 1.2)} × 28% ) ca. 10%
of ONOOH. Thus, 12% is the maximum yield that we can
get. Furthermore, it should be considered that at 290 nm
ONOOH (ꢀ ≈ 190 M-1 cm-1) has by far the strongest
absorbance of all occurring species, including O2NOOH (ꢀ
< 20 M-1 cm-1). Most importantly, below pH 2 O2NOOH
is a stable species on our experimental time scale, its lifetime
being thousands of seconds.17 Therefore, the accumulation
of no more than ca. 12% O2NOOH during the decomposition
of ONOOH in a first-order reaction has no effect whatsoever
on the rate constant extracted from the signal.
The reaction ONOOH + H2O h HNO2 + H2O2 is
reversible, and its equilibrium constant, K1, was determined
to be (7.5 ( 0.4) × 10-4 M. From K1, the Gibbs’ energy of
formation of ONOOH was calculated to be 7.1 ( 0.2 kcal/
mol. This value is in good agreement with the one determined
using parameters for radicals formed during homolysis of
peroxynitrite.
Acknowledgment. S.G. and G.C. thank The Israel
Science Foundation.
Supporting Information Available: Additional figure. This
material is available free of charge via the Internet at
http://pubs.acs.org.
Thermodynamics of ONOOH. In the present work, a
very accurate value has been determined directly for the
equilibrium constant of reaction 1, K1 ) (7.5 ( 0.4) × 10-4
M. This allows us to determine both the standard reduction
potential of ONOOH as well as the Gibbs’ energy of
formation of ONOOH. The standard reduction potential of
H2O2, E°19, is 1.76 V.18 By combining eqs 1 and 19, we
obtain E°20 ) 1.67 V.
IC025698R
(19) Wagman, D. D.; Evans, W. H.; Parker, V. B.; Hallow, I.; Bailey, S.
M.; Schumm, R. H. Selected Values of Chemical Thermodynamic
Properties; National Bureau of Standards and Technology Note 270-
3; U.S. Government Printing Office: Washington, DC, 1968.
(20) Schmid, G.; Neumann, U. Z. Phys. Chem. N. F. 1967, 54, 152-159.
(21) Logager, T.; Sehested, K. J. Phys. Chem. 1993, 97, 6664-6669.
(22) Mahoney, L. R. J. Am. Chem. Soc. 1970, 92, 5262-5263.
(23) Gerasimov, O. V.; Lymar, S. V. Inorg. Chem. 1999, 38, 4317-4321.
(24) Hodges, G. R.; Ingold, K. U. J. Am. Chem. Soc. 1999, 121, 10695-
10701.
(25) Stanbury, D. M. AdV. Inorg. Chem. 1989, 33, 69-138.
(26) Schwarz, H. A.; Dodson, R. W. J. Phys. Chem. 1984, 88, 3643-
3647.
(27) Klaening, U. K.; Sehested, K.; Holcman, J. J. Phys. Chem. 1985, 89,
760-763.
H2O2 + 2H+ + 2e- h 2H2O E°19 ) 1.76 V (19)
ONOOH + 2H+ + 2e- h HNO2 + H2O E°20 ) 1.67 V
(28) In the following reference: Wagman, D. D., et al. J. Phys. Chem.
Ref. Data, Suppl. 1982, 11 (2), another value is given. However, this
value was shown to be erroneous in the following: Ram, M. S.;
Stanbury, D. M. Inorg. Chem. 1985, 24, 2954. Park. J.-Y.; Lee, Y.-
N. J. Phys. Chem. 1988, 92, 6294.
(29) O’Sullivan, D. W.; Lee, M.; Noone, B. C.; Heikes, B. G. J. Phys.
Chem. 1996, 100, 3241. The combination of Henry’s law constant
for H2O2 taken from this work with the accurately known gas-phase
value of ∆fG° (H2O2, g) yields the same value for ∆fG° (H2O2, aq) as
tabulated in ref 19. This value is also in excellent agreement with the
one derived from electrochemical measurements in acidic media, as
described in ref 18.
(20)
The Gibbs’ energies of formation of HNO2 and H2O2 in
water and that of H2O are accurately known, with ∆Gf°-
(HNO2) ) -13.3 kcal/mol,19,20,28 ∆Gf°(H2O2) ) - 32.0 kcal/
mol,18,19,29 and ∆Gf°(H2O) ) - 56.7 kcal/mol. By means of
(17) Re´gimbal, J.-M.; Mozurkewich, M. J. Phys. Chem. A 1997, 101, 8822.
(18) Standard Potentials in Aqueous Solutions; Bard, A. J., Parsons, R.,
Jordan, J., Eds.; Marcel Dekker, Inc.: New York, 1985; p 57.
3800 Inorganic Chemistry, Vol. 42, No. 12, 2003