Mechanisms of ferriheme reduction by nitric oxide: Nitrite and general base catalysis

Article abstract of DOI:10.1021/ic049532x

The reductive nitrosylation (Fe^III(P) + 2NO + H₂O = Fe^II(P)(NO) + NO₂^- + 2H⁺) of the ferriheme model Fe^III(TPPS) (TPPS = tetra(4-sulfonatophenyl) porphyrinato) has been investigated in moderately acidic solution. In the absence of added or adventitious nitrite, this reaction displays general base catalysis with several buffers in aqueous solutions. It was also found that the nitrite ion, NO₂^-, is a catalyst for this reaction. Similar nitrite catalysis was demonstrated for another ferriheme model system Fe^III(TMPy) (TMPy = meso-tetrakis(N-methyl-4-pyridyl)porphyrinato), and for ferriheme proteins met-hemoglobin (metHb) and met-myoglobin (metMb) in aqueous buffer solutions. Thus, it appears that such catalysis is a general mechanistic route to the reductive nitrosylation products. Two nitrite catalysis mechanisms are proposed. In the first, NO₂^- is visualized as operating via nucleophilic addition to the Fe^III-coordinated NO in a manner similar to the reactions proposed for Fe^III reduction promoted by other nucleophiles. This would give a labile N₂O ₃ ligand that hydrolyzes to nitrous acid, regenerating the original nitrite. The other proposal is that Fe^III reduction is effected by direct outer-sphere electron transfer from NO₂^- to Fe ^III(P)(NO) to give nitrogen dioxide plus the ferrous nitrosyl complex Fe^II(P)(NO). The NO₂ thus generated would be trapped by excess NO to give N₂O³ and, subsequently, nitrite. It is found that the nitrite catalysis rates are markedly sensitive to the respective Fe^III(P)(NO) reduction potentials, which is consistent with the behavior expected for an outer-sphere electron-transfer mechanism. Nitrite is the product of NO autoxidation in aqueous solution and is a ubiquitous impurity in experiments where aqueous NO is added to an aerobic system to study biological effects. The present results demonstrate that such an impurity should not be assumed to be innocuous, especially in the context of recent reports that endogenous nitrite may play physiological roles relevant to the interactions of NO and ferriheme proteins.

Full text of DOI:10.1021/ic049532x

Inorg. Chem. 2004, 43, 5393−5402
Mechanisms of Ferriheme Reduction by Nitric Oxide: Nitrite and
General Base Catalysis1
Bernadette O. Fernandez, Ivan M. Lorkovic, and Peter C. Ford*
Department of Chemistry and Biochemistry, UniVersity of California, Santa Barbara,
Santa Barbara, California 93106-9510
Received April 9, 2004
The reductive nitrosylation (Fe (P) + 2NO + H O ) Fe (P)(NO) + NO - + 2H ) of the ferriheme model Fe (TPPS)
III
II
+
III
2
2
(TPPS ) tetra(4-sulfonatophenyl)porphyrinato) has been investigated in moderately acidic solution. In the absence
of added or adventitious nitrite, this reaction displays general base catalysis with several buffers in aqueous solutions.
It was also found that the nitrite ion, NO -, is a catalyst for this reaction. Similar nitrite catalysis was demonstrated
2
III
for another ferriheme model system Fe (TMPy) (TMPy ) meso-tetrakis(N-methyl-4-pyridyl)porphyrinato), and for
ferriheme proteins met-hemoglobin (metHb) and met-myoglobin (metMb) in aqueous buffer solutions. Thus, it appears
that such catalysis is a general mechanistic route to the reductive nitrosylation products. Two nitrite catalysis
mechanisms are proposed. In the first, NO - is visualized as operating via nucleophilic addition to the Fe -coordinated
III
2
III
NO in a manner similar to the reactions proposed for Fe reduction promoted by other nucleophiles. This would
give a labile N O ligand that hydrolyzes to nitrous acid, regenerating the original nitrite. The other proposal is that
2
3
Fe reduction is effected by direct outer-sphere electron transfer from NO - to Fe (P)(NO) to give nitrogen dioxide
III
III
2
II
plus the ferrous nitrosyl complex Fe (P)(NO). The NO thus generated would be trapped by excess NO to give
2
N O and, subsequently, nitrite. It is found that the nitrite catalysis rates are markedly sensitive to the respective
2
3
III
Fe (P)(NO) reduction potentials, which is consistent with the behavior expected for an outer-sphere electron-
transfer mechanism. Nitrite is the product of NO autoxidation in aqueous solution and is a ubiquitous impurity in
experiments where aqueous NO is added to an aerobic system to study biological effects. The present results
demonstrate that such an impurity should not be assumed to be innocuous, especially in the context of recent
reports that endogenous nitrite may play physiological roles relevant to the interactions of NO and ferriheme proteins.
Introduction
immune response, and numerous disease states involving NO
imbalances have been identified.3,4 Much of the mammalian
biology of NO can be attributed to reactions with metal
centers such as hemes; thus, it is important to understand
the fundamental mechanisms of such processes. One such
reaction is reductive nitrosylation, a process whereby the NO
reduction of the metal is concomitant with the nitrosation
of a nucleophile (eq 1). This has been demonstrated in several
ferriheme proteins including met-hemoglobin (metHb),5,6
Nitric oxide (nitrogen monoxide) is known to play key
roles in a variety of mammalian physiological processes,2,3
including blood pressure control, neurotransmission, and
* Author to whom correspondence should be addressed. E-mail: ford@
chem.ucsb.edu.
(1) Aspects of the work reported here have appeared in preliminary
communications. (a) Fernandez, B. O.; Lorkovic, I. M.; Ford, P. C.
Inorg. Chem. 2003, 42, 2-4. (b) Fernandez, B. O.; Ford, P. C. J. Am.
Chem. Soc. 2003, 125, 10510-10511.
(2) (a) Moncada, S.; Palmer, R. M. J.; Higgs, E. A. Pharmacol. ReV. 1991,
43, 109-142. (b) Feldman, P. L.; Griffith, O. W.; Stuehr, D. J. Chem.
Eng. News 1993, 71, 26. (c) Butler, A. R.; Williams, D. L. Chem.
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Grisham, M. B.; Laval, F.; Nims, R. W.; Laval, J.; Cook, J.; Pacelli,
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Cell. Regul. 1996, 34, 159-187.
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P. C.; Lorkovic, I. M. Chem. ReV. 2002, 102, 993-1017. (c)
Reichenbach, G.; Sabatini, S.; Palombari, R.; Palmerini, C. A. Nitric
Oxide 2001, 5, 395-401. (d) Cabail, M. Z.; Pacheco, A. A. Inorg.
Chem. 2003, 42, 270-272. (e) Ford, P. C.; Laverman, L. E.; Lorkovic,
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10.1021/ic049532x CCC: $27.50 © 2004 American Chemical Society
Published on Web 07/29/2004
Inorganic Chemistry, Vol. 43, No. 17, 2004 5393

Fernandez et al.
Scheme 1. Proposed Ferrihemeprotein Reductive Nitrosylation at pH
> 7 6
metHb(NO) reaction with water that is activated by general
base catalysis. It was in this context that the research
described here was initiated to determine whether analogous
general base catalysis pathways are apparent in the NO
reductions of water-soluble ferriheme models such as
FeIII(TPPS)(H2O)2 (1, TPPS ) tetra(4-sulfonato-phenyl)-
porphyrinato) in aqueous buffered solutions (eq 3). General
base catalysis is indeed demonstrated; however, perhaps an
-
even more interesting discovery is that the nitrite ion (NO2 )
strongly catalyzes this reaction. This catalytic pathway also
extends to another ferriheme model, FeIII(TMPy)(H2O)2 (2,
TMPy ) meso-tetrakis(N-methyl-4-pyridyl)porphyrinato),
and to the ferriheme proteins metMb and metHb.
cupric centers in cytochrome c oxidase and laccase, and
Cu(II) and Fe(III) model systems.7,8 Reductive nitrosylation
has also received attention as a possible pathway for the
formation of biological S-nitrosothiols and N-nitrosoamines.9
In addition, it has been suggested that such a pathway may
be responsible for the nitrosation of the â-cys-93 of
hemoglobin to form S-nitrosohemoglobin (SNO-Hb),10 the
subject of a controversial proposal as a nitric oxide carrier
in the cardiovascular system.11
FeIII(TPPS) + 2 NO + H2O h
FeII(TPPS)(NO) + NO2- + 2H+ (3)
As noted above, the transfer of NO+ from the heme iron
to a â-cys-93 in the hemoglobin analogous to the reaction
shown in eq 1 was proposed as a possible mechanism for
the formation of SNO-Hb.10 This proposal is similar to the
report that the oxidized myoglobin (metMb) reacts with NO
plus glutathione (GSH) to give S-nitrosoglutathione (GSNO).5c
However, it has also been argued that the formation of SNO-
Hb may be the result of mixing artifacts when NO is added
to ferriheme protein or red blood cell solutions, especially
when dioxygen is present.13 A related issue is the recent
attention directed toward the possible role of the nitrite ion
LxMn+ + NO + Nuc h LxM(n - 1) + Nuc-NO+ (1)
In an earlier investigation, Hoshino and co-workers6 re-
ported the kinetics for the NO reductions of ferriheme pro-
teins, ferric cytochrome-c (CytIII), met-myoglobin (metMb),
and met-hemoglobin in buffered aqueous media. At pH
values >6.5, the rates of these reactions displayed a first-
order dependence on [OH-] and a complex dependence on
[NO], consistent with the equilibrium formation of the
nitrosyl complex FeIII(NO) followed by the nucleophilic
attack of OH- at the coordinated nitrosyl (Scheme 1). Such
coordination leads to charge transfer from the NO to the
Fe(III) center,12 rendering the nitrosyl more electrophilic than
the free ligand.
-
(NO2 ) in the cardiovascular system.14 Because nitrite is
reported to be a major vascular storage pool of NOx species,14
the chemical interplay of ferriheme models and proteins with
NO and nitrite continues to be an issue of major importance
to cardiovascular biology.
Experimental Section
The rate law predicted by Scheme 1 is described by eq 2,
and the observed rates for the NO reduction of CytIII, metMb,
and metHb proved to be functions of both [NO] and [OH-]
in accord with this analysis.
Materials. All solutions were prepared with doubly distilled
water from a Nanopure distillation system. The iron(III) porphyrin
salts (Na4[FeIII(TPPS)] and [FeIII(TMPy)(Cl)]Cl4) were obtained
from Midcentury (Posen, IL), human hemoglobin (Hb) was obtained
from Sigma, and horse skeletal myoglobin (Mb) was obtained from
Calbiochem. Reagent-grade sodium nitrite (Fisher Scientific) was
used without further purification. Buffer solutions were prepared
from sodium bisulfate (NaHSO4‚H2O (Allied Chemical) (pH 2
K
OH[OH-]
d[FeIII]
KNO[NO]
) -kd ×
×
[FeIII] (2)
1 + KOH[OH-]
dt
1 + KNO[NO]
However, metHb also demonstrated a pH-independent
(11) (a) Jia, L.; Bonaventura, C.; Bonaventura, J.; Stamler, J. S. Nature
1996, 380, 221-226. (b) Gow, A. J.; Luchsinger, B. P.; Pawloski, J.
R.; Singel, D. J.; Stamler, J. S. Proc. Natl. Acad. Sci. U.S.A. 1999,
96, 9027-9032. (c) See also Gladwin, M. T.; Lancaster, J. R., Jr.;
Freeman, B. A.; Schechter, A. N. Nat. Med. (N.Y.) 2003, 9, 496-
500.
(12) (a) Scheidt, W. R.; Lee, Y. J.; Hatano, K. J. Am. Chem. Soc. 1984,
106, 3191-3198. (b) Wylie, G. R. A.; Scheidt, W. R. Chem. ReV.
2002, 102, 106-1089.
(13) (a) Han, T. H.; Hyduke, D. R.; Vaughn, M. W.; Fukuto, J. M.; Liao,
J. C. Proc. Natl. Acad. Sci. U.S.A. 2002, 99, 7763-7768. (b) Zhang,
Y.; Hogg, N. Free Radical Biol. Med. 2002, 32, 1212-1219. (c)
Herold, S.; Rock, G. J. Biol. Chem. 2003, 278, 6623-6634.
(14) (a) Cosby, K.; Partovi, K. S.; Patel, R. P.; Reiter, C. D.; Martyr, S.;
Yang, B. K.; Waclawiw, M. A.; Zaloa, G.; Xu, X.; Huang, K. T.;
Shields, H.; Kim-Shapiro, D. B.; Schechter, A. N.; Cannon, R. O.,
III; Gladwin, M. T. Nat. Med. (N.Y.) 2003, 9, 1498-1505. We thank
Dr. M. T. Gladwin for an advance copy of this manuscript. (b)
Rodriquez, J.; Maloney, R. E.; Rassaf, T.; Bryan, N. S.; Feelisch, M.
Proc. Natl. Acad. Sci. U.S.A. 2003, 100, 336-341.
pathway for reductive nitrosylation at lower pH values (<6)
(kH O ) 1.1 × 10-3 s-1),6 possibly the result of the
2
(7) (a) Torres, J.; Cooper, C. E.; Wilson, M. T. J. Biol. Chem. 1998, 273,
8756-8766. Torres, J.; Svistunenko, D.; Karlsson, B.; Cooper, C. E.;
Wilson, M. T. J. Am. Chem. Soc. 2002, 124, 963-967. (c) Tran, D.;
Skelton, B. W.; White, A. H.; Laverman, L. E.; Ford, P. C. Inorg.
Chem. 1998, 37, 2505-2511.
(8) For example, (a)Wayland, B. B.; Olson, L. W. J. Am. Chem. Soc.
1974, 96, 6037-6041. (b) Trofimova, N. S.; Safronov, A. Y.; Ikeda,
O. Inorg. Chem. 2003, 42, 1945-1951.
(9) (a) Feelisch, M. S.; Rassaf, T.; Mnaimneh, S.; Singh, N.; Byran, N.
S.; Jourd’Heuil, D.; Kelm, M. FASEB J. 2002, 16, 1775-1785 (b)
Bryan, N. S.; Rassaf, T.; Maloney, R. E.; Rodriguez, C. M.; Saijo, F.;
Rodriguez, J. R.; Feelisch, M. Proc. Nat. Acad. Sci. U.S.A. 2004, 101,
4308-4313.
(10) Luchsinger, B. P.; Rich, E. N.; Gow, A. J.; Williams, E. M.; Stamler,
J. S.; Singel, D. J. Proc. Natl. Acad. Sci. U.S.A. 2003, 100, 461-466.
5394 Inorganic Chemistry, Vol. 43, No. 17, 2004

Mechanisms of Ferriheme Reduction by Nitric Oxide
solutions)), chloroacetic acid (Mallinckrodt) (pH 3), sodium acetate,
(Mallinckrodt) mixed with triflic acid (CF3SO3H) (Aldrich) (pH
4-5), N,N′-diethyl-N,N′-bis(sulfopropyl)ethylenediamine, DESPEN
(HO3SC3H6)(C2H5)NC2H4N(C2H5)-(C3H6-SO3H) (GFS Chemi-
cals) (pH 4.5-6), and NaH2PO4/Na2HPO4 (Fisher Scientific) (pH
6-7.4).
Chromatographic-grade argon (Praxair) was passed through an
Oxiclear oxygen trap (Alltech) before use to provide an inert
atmosphere. Nitric oxide (99%, Aire Liquide) was purified by
passage through a stainless steel column filled with Ascarite II
(Thomas Scientific) to remove higher nitrogen oxides. All NO
transfers were carried out on a greaseless vacuum line.
General Methods. The purchased hemoglobin and myoglobin
samples were mixtures of oxidized hemo/myoglobin and the oxy
form of reduced hemo/myoglobin.15 To obtain purified metHb, we
subjected the solutions to gentle degassing on a vacuum line because
removing dioxygen from the solutions stabilizes the oxidized form
of the protein.15b Another method to obtain metHb is to react the
protein solution with excess of K3[Fe(CN)6] (Mallinckrodt) to
oxidize the remaining Hb in the sample.15c The metHb solution
was then purified via molecular exclusion chromatography by
passing through a column of CM Sephadex C-25 resin (Pharmacia).
Purified metMb was obtained by analogous procedures.
Buffer solutions were prepared according to standard procedures.
Unless otherwise noted, solutions were also adjusted to a common
ionic strength (µ), generally 0.10 M, by adding sodium triflate,
Na[CF3SO3] (Aldrich), as the supporting electrolyte. Dilute solutions
of triflic acid were added to the reaction solutions to adjust the
pH.
syringes via argon-purged airtight syringes (Hamilton Gastight no.
1010 or no. 1005) equipped with removable needles, Teflon
plungers, and exit valves. All solutions were protected from room
light with aluminum foil or a black cloth and were used within the
same day of preparation.
Instrumentation. The solution pH was monitored using a deep-
vessel, Corning combination electrode with a silver ion barrier
(model 476306). The pH meter (Orion Research) and electrode were
calibrated over the pH 4-10 range with standard buffers (EM
Science and Fisher Scientific). Electronic absorption spectra and
kinetics measurements (Agilent Chemstation) were recorded on a
Hewlett-Packard 8452A diode array spectrophotometer using
custom-made reaction vessels equipped with 1.0-cm quartz cells
(NSG Precision Cells and Hellma). All kinetics analyses and
measurements were made using a thermostated cell holder attached
to a cell stirring module that was controlled by a temperature bath
to ensure proper equilibration and mixing of solutions while data
was being accumulated. Other kinetics measurements were per-
formed with an Applied Photophysics model SX-18 MV stopped-
flow spectrometer using standard mixing techniques. Drive syringes
were thoroughly purged with the appropriate solutions after
degassing, and samples were equilibrated for 15-30 min at the
appropriate temperature using the attached temperature control unit.
Data Analysis. Absorbance intensity versus time traces obtained
by single-wavelength detection from all instruments were converted
to ∆Abs versus time plots and were fit to appropriate kinetics
models using Igor 4.0 (Wavemetrics) software.
Results and Discussion
Aqueous solutions for kinetics measurements were deaerated.
Iron porphyrin solutions were subject to at least three freeze-
pump-thaw cycles (f-p-t) using specialized cells described below.
Because ferriheme proteins denature under the temperature extremes
of the f-p-t cycles, we degassed them via boiling under reduced
pressure (evacuating the samples on the vacuum manifold for
several cycles at room temperature).
NO solutions for kinetics studies were prepared by vacuum
transfer techniques using flasks of known volumes with greaseless
fittings, high-vacuum Teflon stopcocks, a four-sided quartz cuvette,
a septum-sealed electrode adaptor (as injection ports) for anaerobic
use, and an O-ring adaptor for the connection to the gas/vacuum
manifold. All manipulations of NO gas and NO-containing solutions
were carried out on vacuum lines using stainless steel or glass with
greaseless fittings. When necessary, NO samples were passed
through a secondary flask containing additional Ascarite II or 6 M
NaOH to remove other NOx species. The NO was frozen at 77 K
into this secondary scrubber, and then the cleaned NO was
quantitatively distilled into the desired reaction flask after the
secondary flask was thawed. After NO transfer, the flask contents
were thawed and shaken for equilibration. NO concentrations were
calculated from PNO and the assumption that the solubility is the
same as it is in water.16
Proton-Coupled Redox Potentials. The redox reaction
depicted in eq 3 can be separated into the equilibrium
formation of the ferriheme nitrosyl complexes (eq 4) and
the subsequent redox reaction (eq 5). The potential for the
latter is drawn from the half cells shown in eqs 6 and 7.17
A
-
key feature is the marked acid dependence of the NO2 /NO
half cell (E° ) 1.20 V).18 Under standard conditions (1 N
H+), the NO reductions of neither FeIII(TPPS)(NO) nor
FeIII(TMPy)(NO) are favorable (∆E° ) -0.62 and -0.42
V, respectively, for eq 5); however, in both cases, the reaction
becomes favorable at higher pH values (∆E ) 0.21 and 0.41
V, respectively, at pH 7) and at lower nitrite concentrations.
In this context, kinetics experiments with these two ferriheme
model systems were carried out at pH > 3 but at pH < 6 to
suppress the hydrolysis of FeIII-coordinated H2O and the
formation of oxo-bridged dimers FeIIIOFeIII.17,19,20
KNO
FeIII(P) + NO y z FeIII(P)(NO)
(4)
FeIII(P)(NO) + H2O + NO h
FeII(P)(NO) + NO2- + 2H+(5)
Solutions for stopped-flow experiments were transferred from
the equilibrated reaction flasks and loaded onto the evacuated drive
FeIII(P)(NO) + e- h FeII(P)(NO)
NO2- + 2H+ + e- h H2O + NO
(6)
(7)
(15) (a) Information from the Sigma-Aldrich website (www.
sigma-aldrich.com) and from Sigma-Aldrich customer service/
technical support. (b) Budavari, S., Ed. Merck Index: An Encycopedia
of Chemicals, Drugs, and Biologicals, 11th ed.; Merck & Co: Rahway,
NJ, 1989; p 734. (c) Joshi, M. S.; Ferguson, T. B., Jr.; Han. T. H.;
Hyduke, D. R.; Liao, J. C.; Rassaf, T.; Bryan, N.; Feelisch, M.;
Lancaster, J. R., Jr. Proc. Natl. Acad. Sci. U.S.A. 2002, 99, 10341-
10346.
In the context of the above discussion, it is notable that
FeIII(TPPS) solutions prepared at pH 2.04 (100 mM bisulfate)
(17) (a) Barley, M. H.; Takeuchi, K. J.; Meyer, T. J. J. Am. Chem. Soc.
1986, 108, 5876-5885. (b) Barley, M. H.; Rhodes, M. R.; Meyer, T.
J. Inorg. Chem. 1987, 26, 1746-1750.
(18) Bard, A. J., Parsons, R., Jordan, J., Eds. Standard Potentials in Aqueous
Solutions; Marcel Dekker: New York, 1985; pp 127-139.
(16) Young, C. L., Ed. Oxides of Nitrogen; IUAPC Solubility Data Series;
Pergamon Press: Oxford, U.K., 1983; Vol. 8.
Inorganic Chemistry, Vol. 43, No. 17, 2004 5395

Fernandez et al.
demonstrated no reduction by NO (1.9 mM), whereas a pH
3.06 solution (100 mM monochloroacetic acid) exhibited
minimal spectral changes.
The equilibrium constant for eq 3 was determined for
FeIII(TPPS) in 298 K, pH 4.5 NaOAc/HOAc buffer (50 mM)
with [NO] ) 1.9 mM by recording spectral changes as the
nitrite concentration was varied from 0 to 40 mM (Supporting
Information, Figure S1). From these spectral changes and a
Lineweaver-Burke-type analysis, the equilibrium constant
was determined to be (3.35 ( 0.89) × 10-2 M. From this
value, the reduction potential for eq 6 (P ) TPPS) was found
to be (0.582 ( 0.007) V, a value that compares well with
that (0.586 V) determined by Meyer and co-workers using
cyclic voltammetry.17a
Figure 1. Spectral changes upon adding different [NO] to an aqueous
solution of FeIII(TMPy) with 100 mM triflic acid (µtot ) 0.10 M) at 298 K.
The inset is a linear plot of ∆Abs-1 (at 394 nm) versus [NO]-1 from which
the equilibrium constant KNO ) (4.6 ( 1.7) × 102 M-1 was calculated.
Equilibrium Formation of FeIII(P)(NO). Upon mixing
a solution of FeIII(TPPS) (1) with NO, the equilibrium
formation of FeIII(TPPS)(NO) (3) is evident in the shift of
the Soret band λmax at 394 nm (ꢀ ) 9.7 × 104 M-1 cm-1)
characteristic of 1 to the λmax (422 nm, ꢀ ) 12.0 × 104 M-1
cm-1) characteristic of 3. At low pH (<3), no significant
further reaction is observed. At higher pH values, there is a
progressive change in the spectrum both in the Soret and
Q-band regions consistent with the reduction of 3 to
FeII(TPPS)(NO) (4) (λmax(Soret) ) 412 nm, ꢀ ) 10.0 × 104
M-1 cm-1) as indicated by eq 5 (Supporting Information,
Figure S2). Because the second reaction under these condi-
tions is slow, it is possible to evaluate the equilibrium
represented by eq 4 from the spectral changes observed upon
changing the NO partial pressure for pH 5 buffered solutions
(µ ) 0.10 M) at 298 K. The value of (1.32 ( 0.09) × 103
M-1 determined for K is in good agreement with the
previously reported values of (1.2 ( 0.2) × 103 M-1 and
1.1 × 103 M-1 in acidic solution.21
techniques were used to measure the kobs for the reaction of
NO with FeIII(TMPy) for four different [NO] ((0.20-1.0
mM) at pH 1.0, µ ) 0.10) and T ) 298 K. The kobs values
determined by monitoring the temporal spectral changes at
542 nm were plotted versus [NO] to give linear plots (data
not shown) with slopes equal to kon (2.9 × 104 M-1 s-1) and
intercepts equal to koff (68 s-1). Interestingly, both kon and
k
off in this case are significantly smaller than the analogous
rate constants measured for 1 (5 × 105 M-1 s-1 and 460
s-1, respectively) by laser flash photolysis techniques under
comparable conditions.21a The equilibrium constant KNO
obtained from the ratio of these values was 4.2 × 102 M-1,
in agreement with that determined spectroscopically.
kon
FeIII(P) + NO y z FeIII(P)(NO)
(8)
koff
The values of KNO for ferriheme proteins metHb and
metMb have been previously reported as 1.4 × 104 and 1.3
× 104 M-1, respectively, in buffered solutions at pH 7.21,23
NO Reduction of FeIII(TPPS). General Base Catalysis.
The kinetics for the NO reduction and nitrosylation of 1 to
4 were investigated over pH range 3-6 using various buffers
at 298 K. Under these conditions, µ-oxo dimerization of
FeIII(TPPS) is minimal. Equilibrating a buffered aqueous
solution of 1 with NO (760 Torr, [NO] ) 1.9 mM) led to
the rapid formation of an equilibrium mixture of 1 and the
ferric nitrosyl complex 3 (eq 4), followed by the slower
reduction to the ferrous analogue 4 (eq 5). The solution
spectra exhibited a corresponding shift in the Soret band λmax
from 422 to 412 nm. The Q-band region also exhibited
temporal changes that are characteristic of this reaction: a
decrease at 536 nm and the growth of a band at 542 nm.
Under such a large excess of NO, the system cleanly forms
the reduced ferrous nitrosyl complex (Figure 2), and the
spectral changes seen were independent of the buffers used.
Temporal plots of the disappearance of 1 monitored at 422
nm or the appearance of 4 monitored at 412 nm could be fit
to exponential functions under these conditions and were in
agreement with each other. The kobs values were determined
Because the rate constants for the NO reduction of
FeIII(TMPy) (2) proved to be much larger than those for
FeIII(TPPS) (10-3 vs 10-4 s-1), the equilibrium constant for
NO association was determined at lower pH, where the ferric
nitrosyl species would be stable. In pH 1.0 (µtot ) 0.10 M,
using 100 mM triflic acid) aqueous solution, the equilibrium
constant for eq 4 at 298 K was determined for FeIII(TMPy)
from the changes in the absorption spectrum as NO was
added. Spectrophotometric determination of KNO in this
manner gave (4.6 ( 1.7) × 102 M-1 (Figure 1), roughly a
factor of 3 smaller than the KNO determined for 1.
An alternative way to measure KNO is to determine the
rate constants, kobs, for the approach to equilibrium concen-
trations of the FeIII(P)(NO) and FeIII(P) complexes after
mixing solutions of FeIII(P) and of NO. The observed rate
constant is equal to the sum of kon[NO] + koff, where kon
and koff are the respective rate constants for the formation
and dissociation of Fe(P)(NO) (eq 8).22 Stopped-flow kinetics
(19) Forshey, P. A.; Kuwana, T. Inorg. Chem. 1981, 20, 693-700.
(20) (a) El-Awady, A. A.; Wilkins, P. C.; Wilkins, R. G. Inorg. Chem.
1985, 24, 2053-2057. (b) Harris, F. L.; Toppen, D. L. Inorg. Chem.
1978, 17, 71-73.
(21) (a) Laverman, L. E.; Ford, P. C. J. Am. Chem. Soc. 2001, 123, 11614-
11622. (b) Hoshino, M.; Ozawa, K.; Seki, H.; Ford, P. C. J. Am. Chem.
Soc. 1993, 115, 9568-9575. (c) Nakagawa, S.; Yashiro, T.; Munakata,
H.; Imai, H.; Uemori, Y. Inorg. Chim. Acta 2003, 349, 17-22.
(22) Lorkovic, I. M.; Miranda, K. M.; Lee, B.; Bernhard, S.; Schoonover,
J. R.; Ford, P. C. J. Am. Chem. Soc. 1998, 120, 11674-11683.
(23) Laverman, L. E.; Wanat, A.; Oszajca, J.; Stochel, G.; Ford, P. C.; van
Eldik, R. J. Am. Chem. Soc. 2001, 123, 285-293.
5396 Inorganic Chemistry, Vol. 43, No. 17, 2004

Mechanisms of Ferriheme Reduction by Nitric Oxide
For example, a majority of the FeIII(TPPS) experiments were
carried out with [NO] ) 1.9 mM at 298 K where kred would
be 1.4 kobs
.
kred
8
FeIII(P)(NO) + H2O + NO
FeII(P)(NO) + NO2- + 2H+ (10)
In the pH range of 4-6, kred did prove to be dependent on
the identity and concentration of the buffer used. For the
linear plot of kobs versus the concentration of the NaOAc/
HOAc buffer at 0.1 M ionic strength and PNO ) 1 atm, a
slope of 1.73 × 10-3 M-1 s-1 and an intercept of 1.95 ×
10-4 s-1 were found (Supporting Information, Figure S3).
This implies the relationship kred ) k0 + kbuffer[buffer], where
k0 is a buffer-independent term as indicated by the nonzero
y intercept. For NaOAc/HOAc at pH 5.0, the buffer-
dependent term, kacetate was determined to be (2.4 ( 0.1) ×
10-3 M-1 s-1 (calculated by multiplying the slope by 1.4 as
discussed above), and k0 was determined to be 2.8 × 10-4
s-1. At the same pH, the buffer DESPEN exhibited similar
behavior but with a 5-fold greater slope (kDESPEN ) (1.1 (
0.1) × 10-2 M-1 s-1). Such behavior was previously seen
with the NO reduction of the copper(II) complex Cu(dmp)22+
(dmp ) 2,9-dimethyl-1,10-phenanthroline) and was attributed
to general base catalysis.7c In this context, one may propose
a mechanism for the NO reduction of 1 under conditions of
general base catalysis as illustrated by Scheme 2.
Similar behavior was seen for phosphate buffer solutions
at pH 6.0 (Supporting Information, Figure S4) where a slope
of 3.0 × 10-3 M-1 s-1 (kphosphate ) 4.2 × 10-3 M-1 s-1) was
obtained; however, constant ionic strength was not main-
tained. In this case, the extrapolated value of kobs at [buffer]
) 0 (k0 ≈ 6 × 10-4 s-1) is several times larger than the k0
found in the pH 5.0 acetate solutions, suggesting that specific
base catalysis by OH- may be contributing at the higher pH.
Modest specific base catalysis was demonstrated for the
ferriheme protein reductive nitrosylations in pH 7-8 solu-
tions.6
Figure 2. Temporal spectral changes after the introduction of NO (1.9
mM) to FeIII(TPPS) (1.2 × 10-5 M) in pH 4.98 NaOAc/HOAc buffered
(80 mM) aqueous solution at 298 K with µtot ) 0.10 M. The first seven
spectra were taken at 600 s intervals; subsequent scans were recorded every
1200 s. (Inset) Temporal absorbance trace taken at 412 nm, representing
the appearance of FeII(TPPS)(NO). The data and the fit (kobs ) 3.32 ×
10-4 s-1) are indistinguishable.
in this manner for the NO reduction under a variety of
conditions including a range of [NO] (0-1.9 mM) and
several buffers at different pH values. These proved to be
modestly sensitive to ionic strength, so most kinetics
experiments were carried out at a low µ, generally 0.1 M
but sometimes 0.2 M. In all experiments, spectral charac-
teristics were similar to those observed with the phosphate
buffer. Within the NaOAc/HOAc buffer system, kobs dem-
onstrated little or no dependence on the pH. For example,
the kobs for the reaction of 1 with 1.9 mM NO (PNO ) 1.0
atm) in pH 5.0 aqueous solution buffered with 20 mM
NaOAc/HOAc (µ ) 0.1 M) was determined to be (2.14 (
0.37) × 10-4 s-1. At pH 4.0, with the same NaOAc/HOAc
concentration and ionic strength, kobs ) (2.36 ( 0.30) × 10-4
s-1. Thus, at these relatively low solution pH values, there
appears to be no specific base catalysis (by OH-) of the
reductive nitrosylation of FeIII(TPPS). However, the rates
were somewhat faster when the reaction was carried out at
pH 6.0. For example, kobs was measured to be (6.8 ( 0.5) ×
10-4 s-1 in 50 mM phosphate buffer, about 3-fold higher
than that measured at pH 5.0. Although a small fraction of
this increase can be attributed to the higher buffer concentra-
tion (see below), this result suggests that OH- may contribute
to the overall reductive nitrosylation kinetics at this somewhat
higher pH as was seen with the ferriheme proteins.6
Reduction of FeIII(TMPy). The NO reduction of
FeIII(TMPy) (2) also occurs readily.8b Under conditions
analogous to those used for the reactions of 1, the kobs values
for the reduction of FeIII(TMPy)/FeIII(TMPy)(NO) solutions
to FeII(TMPy)(NO) proved to be much faster at all pH values
examined. For example, the value of kobs recorded for the
reductive nitrosylation of 2 in pH 4.5 acetate buffer solution
(20 mM) under 1.0 atm of NO was 7 × 10-3 s-1. This is
more than an order of magnitude larger than that determined
for 1 under analogous conditions.
Consistent with the observations of Hoshino et al., the kobs
values determined for the reduction to FeII(TPPS)(NO) under
specific conditions increase nearly linearly with NO at low
concentrations but approach a limiting value at high [NO].
According to the model described in eqs 4 and 5, the kinetics
should follow the behavior indicated by eq 9
Catalysis by Nitrite Ion. Our initial kinetics studies
suffered reproducibility problems despite considerable efforts
to ensure the consistency of the reaction conditions. These
problems, however, led us to suspect a possible kinetic role
for nitrite, which is not only a reaction product (eq 6) but
also a common impurity in aqueous NO solutions.24 Prepar-
kredKNO[NO][FeIII]
1 + KNO[NO]
d[FeIII]
)
) kobs[FeIII]
(9)
dt
where kred is the rate constant for the reduction of 3 and may
represent the sum of several terms dependent on the medium
and [FeIII] is the sum of [1] and [3]. In general, the kred for
the nitrosyl complex FeIII(P)(NO) (eq 10) can be calculated
(24) (a) Wolak, M.; Stochel, G.; Hamza, M.; van Eldik, R. Inorg. Chem.
2000, 39, 2018-2019. (b) Nemes, A.; Pestovsky, O.; Bakac, A. J.
Am. Chem. Soc. 2002, 124, 421-427.
from the relationship kred ) {(1+KNO[NO])/KNO[NO]}kobs
.
Inorganic Chemistry, Vol. 43, No. 17, 2004 5397

Fernandez et al.
Scheme 2. Proposed Scheme for General Base Catalysis
ingsolutions with NO2- levels below 50 µM proved difficult
to achieve even with strict precautions. The likely source of
such contamination is from NO autoxidation25 by trace
amounts of air. The problem is markedly accentuated by the
tendency of NOx impurities to concentrate in the solution
phase. The reexamination of the spectra from the early
with eq 8 was demonstrated for FeIII(TPPS) in pH 5.0
aqueous NaOAc/HOAc buffer (20 mM) solution by measur-
ing the kobs values for the NO reduction of 1 at a fixed NO2
-
concentration (2.0 mM) but varing [NO]. The rates increased
with [NO] but approached a limiting value. According to eq
9, a double-reciprocal plot (kobs-1 vs [NO]-1) should be linear
with a nonzero intercept equal to kred-1 and a slope equal to
(KNOkred)-1. This plot is shown in Figure 3, from which the
values kred ) (7.2 ( 0.4) × 10-3 s-1 and KNO ) 1.39 × 103
M-1 were obtained. The latter is within the experimental
uncertainty of the KNO values measured for FeIII(TPPS)
from the spectral changes under varied PNO at pH 5 (see
above) and by other workers.21 The kred term equals k0 +
-
kinetics studies suggested that NO2 (at λmax ) 354 nm, ꢀ
) 22.6 M-1 cm-1) levels in the aqueous NO solutions were
in the range of 40-180 µM (with most samples falling in
the range of <100 µM). Although such low levels are rarely
attained with aqueous NO solutions, the problems with
reproducibility led us to add NaNO2 deliberately to the
reaction solution to evaluate its possible catalytic effects on
NO reductions of ferrihemes. As will be seen below, the
effects of nitrite on the kinetics are profound.
-
kbuffer[buffer] + knitrite[NO2 ]; however, under these condi-
-
tions, knitrite[NO2 ] is more than an order of magnitude larger
When NaNO2 (0-2.5 mM) was deliberately added to
buffered solutions of 1 and NO (NaOAc/HOAc buffer at pH
4.96 with µtot ) 0.1 M and [NO] ) 1.9 mM), the rates of
FeIII reduction to 4 were dramatically enhanced. The kobs
than the sum of the other two.
The catalytic rate constant knitrite proved to be independent
of the buffer solution in which the experiment was conducted.
A plot similar to that shown in Figure S5 of the Supporting
-
values proved to be a linear function of [NO2 ] (Sup-
porting Information, Figure S5); that is, kobs ) (kred
k
+
nitrite[NO2-])f(NO), where f(NO) ) (KNO[NO]/(1 + KNO[NO])
and is the function described in eq 9 to account for the role
of NO under a specified set of conditions. The slope of this
plot is 2.2 ( 0.1 M-1 s-1. Once the equilibrium formation
of FeIII(TPPS)(NO) is taken into account (i.e., f(NO)), the
value of the catalytic rate constant knitrite can be calculated
from this slope as 3.1 M-1 s-1, more than 3 orders of
magnitude larger than the effect of the acetate buffer alone.
(kacetate ) 2.4 × 10-3 M-1 s-1.)
In contrast, when millimolar amounts of nitrite (0-100
mM) were added to FeIII(TPPS) in the absence of added NO,
no reduction of 1 to 4 could be seen spectroscopically.
Thus, nitrite is indeed a catalyst for the NO reduction of
-
FeIII(TPPS). Because NO2 is also a reaction product, the
system is, in principle, autocatalytic, although the very small
concentrations of ferric porphyrin present would lead to very
-
low concentrations of NO2 formed via the reaction stoi-
chiometry.
The functional dependence of the nitrite-catalyzed reduc-
tive nitrosylation rates on NO concentration that is consistent
Figure 3. (Upper) Plot of kobs versus [NO] from the NO reduction of
FeIII(TPPS) with varying [NO] (0-1.92 mM) using 50 mM NaOAc/HOAc,
µtot ) 0.35 M, pH 4.99 at 298 K with added NaNO2 (2.01 mM). (Lower)
Plot of kobs-1 versus [NO]-1 (same data). From this linear plot, a KNO value
of (1.39 ( 0.27) × 103 M-1 was determined.
(25) Ford, P. C.; Wink, D. A.; Stanbury, D. M. FEBS Lett. 1993, 326,
1-3.
5398 Inorganic Chemistry, Vol. 43, No. 17, 2004

Mechanisms of Ferriheme Reduction by Nitric Oxide
Information was determined for the nitrite catalysis in the
DESPEN buffer at pH 5.0 with µ ) 0.1 M at 298 K. The
slope in this medium was determined to be 2.4 ( 0.1 M-1
s-1. Taking into account the equilibrium constant for NO
association and the proposed model, we obtained an adjusted
nitrite-catalysis rate constant of knitrite ) 3.4 M-1 s-1 in the
DESPEN buffer solution, which was experimentally equiva-
lent to that measured in acetate buffer at the same pH.
Nitrite catalysis was also seen for the NO reduction of
FeIII(TMPy). This system is sufficiently more reactive,
requiring stopped-flow kinetics techniques to measure rel-
-
evant kobs values over a range of [NO2 ]. By using the
stopped-flow spectrophotometer, solutions of FeIII(TMPy)
(8.5 × 10-5 M) in pH 5.0 NaOAc/HOAc buffer solution
(µtot ) 0.10 M) were mixed with similarly buffered NO
solutions (2.0 mM) containing NaNO2 at various concentra-
tions (10-50 mM) at 298 K. Again, linear plots of kobs versus
-
[NO2 ] with near-zero intercepts were obtained (Supporting
Information, Figure S6). The slopes of these plots reproduc-
ibly fell in the range of 24.6 ( 0.9 M-1 s-1, and from these
slopes and the relationship between kobs and kred described
above, the knitrite value of 83 ( 5 M-1 s-1 was obtained for
the NO reduction of FeIII(TMPy). Notably, this value is
nearly 30-fold larger than that obtained for FeIII(TPPS).
Nitrite Catalysis of Ferriheme Protein Reduction.
The observation of nitrite catalysis with FeIII(TPPS) and
FeIII(TMPy) led us to reexamine the NO reductions of
ferriheme proteins metHb and metMb. These kinetics experi-
ments were carried out in pH 7.0 aqueous phosphate buffer
solutions at 298 K at constant ionic strength (µ ) 0.20 M)
with various NaNO2 concentrations under an atmosphere of
Figure 4. (Upper) Difference spectra recorded upon the reaction of NO
(1.8 mM) with metHb (7 × 10-6 M) in pH 6.99 phosphate buffer solution
(41.3 mM) with [NaNO2] ) 20 mM, µtot ) 0.20 M, 298 K. (Lower)
Difference spectra for the reaction of NO (1.8 mM) with metMb (2.71 ×
10-5 M) in pH 7.00 phosphate buffer solution (50.3 mM) with [NaNO2] )
20 mM, µtot ) 0.20 M, T ) 298 K. For both metHb and metMb, the first
spectrum is taken after 60 s. The next trace is taken after 150 s. The three
subsequent spectra are taken at 150-s intervals. The sixth and seventh spectra
are taken after 900 and 1500 s, respectively. The remaining spectra were
recorded at 600 s intervals.
0.1) × 10-2 M-1 s-1 for metMb according to eq 9. Similar
to the heme models described above, in the absence of NO,
there was no reduction of metHb or metMb by the added
nitrite alone. However, the possible formation of nitrite
complexation to the ferric centers needs to be considered.
(See below.)
NO (1.8 mM). Protein concentrations were 7 × 10-6
M
(metHb) and 2.7 × 10-5 M (metMb). Consistent with the
earlier studies by Hoshino et al.,6 reactions carried out in
pH 7.0 solution in the absence of added nitrite (and in a
-
Nitrite Binding to FeIII(Por) Models. NO2- has recently
been shown to bind to metHb and metMb with equilibrium
constants of 223 and 42.4 M-1, respectively, at 298 K.26 With
this in mind, we instituted studies to evaluate the possible
nitrite ion binding to heme model systems FeIII(TPPS) and
FeIII(TMPy) (eq 10). These studies were carried out at pH 5
to reduce possible complications arising from hydrolysis or
dimerization of the FeIII(P) aquo complexes at higher pH
values and nitrite protonation at lower pH values (pKa
(HNO2) 3.37).27 The concentrations of the ferriheme com-
plexes were 1 × 10-4 M in each case, and spectral changes
in these solutions as a function of added NO2- were evaluated
over the ranges of 0-800 and 0-600 mM for 1 and 2,
respectively. For both complexes, only very small changes
manner to minimize adventitious NO2 ) showed that metHb
and metMb underwent slow reductive nitrosylation but that
the reaction was considerably slower for the latter. Upon
mixing, the protein solutions react rapidly with NO to
generate an equilibrium mixture of the ferriheme protein and
its nitrosyl complex. Under these conditions, this fast process
was followed by the slow formation of the ferrous analogues
Hb(NO) and Mb(NO), with lifetimes of 103 and 104 s,
respectively. However, in the presence of added NaNO2 (0-
20 mM for metHb; 0-80 mM for metMb), the rates
dramatically increased, consistent with the catalytic role of
-
NO2 in these systems as well. Figure 4 displays typical
spectral changes, and Figure S7 in the Supporting Informa-
tion illustrates the linear dependence of the kobs values for
the NO reductions of metHb and metMb proteins under these
conditions.
The slopes of these plots are 0.13 ( 0.01 and (1.09 (
0.10) × 10-2 M-1 s-1 for metHb and metMb, respectively.
Taking into account the equilibrium constants for the forma-
tion of the nitrosyl ferriheme protein complexes metHb(NO)
(KNO ) 1.3 × 104 M-1)6 and metMb(NO) (1.4 × 104
M-1),21b,23 the catalytic rate constants knitrite would be
calculated as 0.14 ( 0.01 M-1 s-1 for metHb and (1.1 (
-
in the UV-vis spectra were seen for [NO2 ] e 50 mM, a
range where the catalysis of NO reduction was clearly
-
evident. However, at higher [NO2 ], there were progressive
spectra changes indicating the formation of complexes,
presumably via the displacement of an axial water of
(26) Wanat, A.; Gdula-Argasinska, J.; Rutkowska-Zbik, D.; Witko, M.;
Stochel, G.; van Eldik, R. J. Biol. Inorg. Chem. 2002, 7, 165-176.
(27) Weast, R. C., Ed. Handbook of Chemistry and Physics, 67th ed.; CRC
Press: Boca Raton, FL, 1986; p D-163.
Inorganic Chemistry, Vol. 43, No. 17, 2004 5399

Fernandez et al.
Scheme 3. Dead-End Equilibrium
Scheme 4. Inner-Sphere Mechanism for Nitrite Catalysis
FeIII(P)(H2O)2. The treatment of these data by standard
methods gave small KNO values of ∼3 and <10 M-1 at
-
2
298 K for 1 and 2, respectively, values that are much smaller
than those reported for the ferriheme proteins.26 (The
-
approximate values of KNO described above are the result
2
of the high nitrite concentrations necessary to effect complex
formation.)
KNO
-
2
-
FeIII(P) + NO2- y
z FeIII(P)(NO2 )
(11)
One model for the potential kinetics effect of eq 11 would
be to treat this as a nitrite “dead-end” equilibrium that inhibits
reaction with NO as illustrated in Scheme 3. In that model,
the formation of ferric nitrosyl complexes is inhibited by
the competition with nitrite for the coordination site.
However, whereas such competition might be expected for
metHb and metMb, which have a proximal histidine ligand
so there is only one potential NO coordination site, it is not
clear what the effect might be for 1 and 2 because these
each have two axial sites available for NO coordination. In
the case of Scheme 3, the rate law would still have the form
Scheme 5. Outer-Sphere Electron-Transfer Mechanism for Nitrite
Catalysis
-
kobs ) (kred + knitrite[NO2 ]) f(NO), but in this case, f(NO) )
-
-
(KNO[NO]/(1 + KNO[NO] + KNO [NO2 ]). The significance
2
of such inhibition (if any) would therefore be based on the
-
-
relative magnitudes of KNO [NO2 ] versus the rest of the
2
denominator, 1 + KNO[NO]. Given the large values of KNO
for the ferriheme systems described here, any such inhibition
must be small in each case because the rate studies that
probed the catalytic role of nitrite were carried out at
relatively high [NO] (1.9 mM). For example, among the
systems studied metMb (42.4 M-1) and metHb (223 M-1)
-
display the largest values of KNO , and the catalytic role of
2
nitrite was examined over the respective concentration ranges
-
of 0-80 and 0-20 mM. At the highest [NO2 ] for which
nitrite catalysis was probed for these two systems, the
-
-
respective values of KNO [NO2 ] were 3.4 and 4.5; by
comparison, values of 1 +2KNO[NO] would be 24.4 and 26.2
under the conditions studied. Thus, we conclude that the
formation of nitrite complexes according to a dead-end
equilibrium of the type indicated by Scheme 3 would have,
at most, a minor effect on the overall rates of the ferriheme
reductions by NO described here.
nism follows the general theme discussed above with regard
to specific and general base catalysis; however, there is no
obvious reason why aqueous NO2- should have the unusual
nucleophilicity necessary for such a catalytic pathway.
A markedly different mechanism is proposed in Scheme
5. The key step in this scheme would be the direct reduction
of the ferriheme nitrosyl to the ferroheme analogue by outer-
Possible Mechanisms of Nitrite Catalysis. There are two
limiting mechanisms by which we can envision nitrite
catalysis to occur. The first (Scheme 4) would be inner sphere
-
-
sphere electron transfer from NO2 . The other product would
in nature where NO2 serves as the primary nucleophile
•
toward the FeIII-coordinated nitrosyl to give an intermediate
FeII-N2O3 complex. The dissociation of the intermediate
releases the ferroheme complex or protein plus N2O3, which
should undergo hydrolysis in aqueous solution to nitrous
acid.28 The ferroheme intermediate is then trapped by excess
NO to produce the FeII(P)(NO) final product. This mecha-
be nitrogen dioxide, NO2 , and this would be rapidly
scavenged by the reaction with excess NO (k ) 1.1 × 109
M-1 s-1)29 to give N2O3, an intermediate with the same
(28) (a) Graetzel, M.; Henglein, A.; Lilie, J.; Beck, G. Ber. Bunsen-Ges.
Phys. Chem. 1969, 73, 646-653. (b) Graetzel, M.; Taniguchi S.;
Henglein, A. Ber. Bunsen-Ges. Phys. Chem. 1970, 74, 488-493.
5400 Inorganic Chemistry, Vol. 43, No. 17, 2004

Mechanisms of Ferriheme Reduction by Nitric Oxide
Scheme 6. Born-Haber-Type Cycle to Estimate the Reduction
Potential for Nitrosyl Complexes of Ferriheme Proteins
demonstrated. The four points on this plot include the two
model compounds for which the ∆E1/2(FeIII/FeII) are well
characterized. To our knowledge, direct measurements of the
metHb(NO)/Hb(NO) and metMb(NO)/Mb(NO) half-cell
reduction potentials have not been reported, but estimates
of 0.49-0.57 and 0.47 V (vs NHE), respectively, were
generated from known reduction potentials and equi-
librium constants6,27,34 using Born-Haber-type cycles (Scheme
6). The values estimated correlate with the knitrite trend:
FeIII(TMPy) > FeIII(TPPS) > metHbT > metMb. Quantita-
tively, Figure 5 is less convincing because the slope of that
plot is 27.8 V-1, whereas the predicted slope would be 38.9
V-1. Nonetheless, such a difference can easily be attributed
to the assumptions inherent in our use of the Marcus cross-
relation treatment as well as the uncertainties in the ferriheme
protein reduction potentials. Thus, the qualitative agreement
is consistent with the outer-sphere electron-transfer mecha-
nism (Scheme 5).
Summary and Commentary. Nitrite ion catalysis of
reductive nitrosylation has been demonstrated for heme
model complexes FeIII(TPPS) and FeIII(TMPy) as well as for
ferriheme proteins metHb and metMb. The increased sen-
sitivity of the kinetics to the FeIII(P)(NO) reduction potential
is consistent with the behavior expected for an outer-sphere
electron-transfer mechanism; however, an inner-sphere pro-
cess may also be sensitive to the overall driving force of the
reaction, as reflected by this parameter. Nitrite is one of the
products of NO autoxidation in aqueous solution24 and is a
ubiquitous component of experiments when aqueous NO is
added to an aerobic system to study biological effects and
should not be assumed to be innocuous. For example,
reactions of NO and red blood cells or metHb have been
reported to give SNO-Hb.10,11 It has been argued that this
product may be the result of the direct transfer of a NO+
moiety from the iron center to the â-cysteine-93 thiol.
However, according to the crystal structure data,35 the
distance between the two sites is quite large, so a direct
reaction of the cysteine-SH with the coordinated NO does
not appear to be likely. However, both the inner- and outer-
Figure 5. Plot of ln(knitrite) vs E1/2 (V), where knitrite is the catalytic rate
constant for the nitrite catalysis of reductive nitrosylation and E1/2 represents
the known or estimated (see text) reduction potentials for the nitrosyl
ferriheme complexes: FeIII(TPPS)(NO), FeIII(TMPy)(NO), metHb(NO), and
metMb(NO).
composition as that proposed for the inner-sphere mecha-
nism. Again, the latter would be subject to hydrolysis to give
the nitrite product. The key step here is the electron transfer,
and although this is not thermodynamically favored (∆E )
-0.31 V), the fast trapping of one of the initial products,
•
NO2 , and subsequent N2O3 hydrolysis may make it viable.
The potential role of an outer-sphere electron transfer as
the rate-limiting step may be evaluated by using the Marcus
cross relation30 to estimate the rate constant for the second-
-
order reaction between FeIII(P)(NO) and NO2 . From the
Marcus cross relation, kOS ≈ (k11k22KOS)1/2, where k11 is the
rate constant for FeII(P)(NO)/FeIII(P)(NO) self-exchange, k22
-
is that for NO2 /NO2 self-exchange via outer-sphere electron
transfer (0.3 M-1 s-1),31 and KOS is the equilibrium constant
of the electron-transfer step (6 × 10-6 for P ) TPPS).32
Accordingly, for kOS ) knitrite, the calculated value for k11
for P ) TPPS would need to be ∼6 × 106 M-1 s-1. This
estimated value falls within the range of the previously
observed self-exchange rates of low-spin Fe porphyrin
complexes;33 however, it is important to recognize that the
FeIII/II(TPPS)(NO) self-exchange rate constant is unknown.
Furthermore, unlike other low-spin ferriheme/ferroheme
couples, the nitrosyl complexes may be subject to re-
organization energy contributions resulting from the different
configurations (linear and bent, respectively) of the Fe-NO
bond in the two oxidation states.12
If an outer-sphere-type mechanism is viable for the nitrite
catalysis of FeIII(P) reduction, then the rate constant knitrite
should be a function of the FeIII(P)(NO)/FeII(P)(NO) reduc-
tion potentials ∆E1/2(FeIII/FeII). This is derived from the
1/2
Marcus cross relation becasue kOS is proportional to KOS
if the rather large assumption is made that k11 is ap-
proximately constant for different hemes. In this context, a
plot of ln(knitrite) versus ∆E1/2(FeIII/FeII) should be linear.
Figure 5 is such a plot, and the predicted pattern is roughly
(34) (a) The equilibrium constant for the formation of FeII(P)(NO) is 8.7
× 109 M-1 for metHbT, 1.7 × 1011 M-1 for metHbR, and 1.4 × 1011
M-1 for metMb.5e The reduction potential for metHb is 0.144 V (vs
NHE) and for metMb is 0.06 V (vs NHE).34b Electrochemical studies
in this laboratory agree with the estimated reduction potential for
metHb(NO) (∼0.530 V vs NHE).34c (b) Fasman, G. D., Ed. Handbook
of Biochemistry and Molecular Biology; CRC Press: Cleveland, OH,
1976; pp 131-150. (c) Kilnic, M. E.; Fernandez, B. O., unpublished
work.
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Inorganic Chemistry, Vol. 43, No. 17, 2004 5401

Fernandez et al.
sphere mechanisms described above for nitrite catalysis call
for the intermediacy of N2O3, a known nitrosating agent.
Although N2O3 is known to undergo hydrolysis, this reaction
has a roughly millisecond lifetime in aqueous solution that
is likely to be longer in the hydrophobic pocket of a heme
protein. Thus, if the reaction were catalyzed by nitrite, then
N2O3 formation at the heme site could easily lead to protein
modification such as â-cys-93 nitrosation. If one considers
the (very low) NO concentrations present under normal
physiological conditions in the cardiovascular system, then
the generation of SNO-Hb via such pathways may not be
relevant under such circumstances. However, in experiments
where a bolus of NO is added to Hb or red blood cell
solutions, the higher localized concentrations of NO might
well lead to products resulting from the intermediacy of
N2O3.
intermediates in biological transformations often attributed
to NO alone.
Acknowledgment. This work was supported by the U.S.
National Science Foundation (CHE0095144). We thank
Chosu Khin, Mark Lim, undergraduate Joy Jackson (URCA
Grant), and Kelly McCord (CNSI Apprentice Researcher)
for experimental measurements relevant to the results
reported. B.O.F. thanks the UCSB Graduate Division for a
Graduate Opportunity Fellowship and a Dissertation Fel-
lowship.
Supporting Information Available: Figures showing spectral
changes, an equilibrium plot, the effects of buffers on rate constants,
and the effects of nitrite catalysis. This material is available free
of charge via the Internet at http://pubs.acs.org.
In summary, the unexpected catalysis pathway described
further emphasizes the potentially important roles of NOx
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