Aqueous Reduction Pathways Of Nitrite

At physiological pH ~7.0-7.4, nitrite anions are fairly stable in the absence of compounds like oxy-hemoglobin. In view of this stability, nitrite is often taken as an indirect measure for NO. Nitrite easily binds to many transition metal ions and decelerates their participation in redox reactions. This property is often used to cure meats or sausages by addition of nitrite. However, the presumed stability of nitrite should be approached with caution. As mentioned above, nitrite rapidly reacts with oxy-hemoglobin to methemoglobin which is unable to bind and transport oxygen. Therefore, excessive nitrite levels may contribute to methemoglobinemia, and excessive consumption of nitrite from canned spinach has been known to cause death in this way.

At acidic pH, nitrite is protonized to nitrous acid (HNO2, pK = 3.37) which is very unstable and prone to spontaneous decay via secondary reactions. In the dark, acidic nitrite solutions are in equilibrium with a very small quantity of nitrosonium cations by heterolytic fission of the bond

NO- + H+ ^ HO—NO ^ HO- ... N+ = O ^ OH- + NO+ (1)

In water, this reaction requires [47] a Gibbs energy of AG = +94 kJ/mol = -RT ln K with K = [NO+]/[NO-] ~ 4 • 10-17 at pH ~7.0.

Therefore, the release of free nitrosonium from nitrite may be considered insignificant in water at physiological pH.

Nitrite anions have a single broad absorption peak at 354 nm (£354 ~ 24 Mcm-1) which changes into a well-resolved multiplet upon protonation to nitrous acid HNO2 (cf Fig. 2 of Chapter 1). Nitrous acid is rather unstable by itself and its decomposition depends on its concentration. At low concentrations, illumination with blue light may induce homolytic fission of nitrous acid into hydroxyl radicals and nitric oxide radicals:

Atmospheric research has shown that photolytic release of OH- affects the ozone budget in the troposphere and is one of the negative effects of the nitrous smog gases in the atmosphere. In aqueous environment also, photolysis was demonstrated. Endogenous nitrite was found to be the major source of nitric oxide radicals that are released when aortic tissue was illuminated with UV light [10].

At higher concentrations, nitrous acid dismutates into nitric oxide, nitrogen dioxide and water

with NO2- subsequently dimerizing into N2O4 and reacting irreversibly with water to nitrite, nitrate and protons [for this reaction cf Chapter 1, Eq. (8)].

IN VITRO EXPERIMENTS ON eNOS IN BUFFERED SOLUTION

eNOS is the endothelial enzyme that catalyzes the synthesis of NO from L-arginine via an oxygen-consuming pathway. This arginine pathway is regulated in a complicated way by a combination of local calcium concentration, phosphorylation of serine moieties and the redox state of intrinsic protein thiols [48]. eNOS is unique among the NOS isoforms to show irreversible myristoylation of the glycine residue at its N terminus as a prerequisite for subsequent reversible palmitoylation and localization at the cell membrane [49,50]. The arginine pathway is blocked upon depletion of the substrate but also if local oxygen levels fall below a threshold value of ca 15 ^M [24]. In our experiments we studied the anoxic reduction of nitrite by buffered eNOS solutions. The anoxia was induced by removing oxygen from the solutions by bubbling with argon and conducting the experiments under an argon atmosphere.

The anoxic activity of eNOS was studied in buffered solution in an optical quartz cuvette that was sealed by a teflon stopper with three holes for the NO electrode, an inflow argon purge and an outflow exhaust, respectively. The setup allows simultaneous monitoring of the optical absorption spectrum of the enzyme and the local NO concentration inside the cuvette. Additionally, the NO content of the exhaust gas flow was determined by feeding the exhaust into a vial containing a solution of NO traps like iron-dithiocarbamate complexes. The protein was either full-length eNOS holoenzyme or only its oxygenase subdomain (eNOSoxy) and obtained by overexpression in E. coli [51]. The enzymatic reaction was started by injecting reductants into the cuvette via the inflow argon purge. The full-length protein contains the flavins of the reductase domain and was started by adding NADPH. The eNOSoxy domain does not contain any flavins and the electron flow was initiated with dithionite. In the presence of oxygen, calmodulin, BH4, L-arginine and NADPH, the holoenzyme produced copious amounts of NO as detected by the electrode. As expected, the electrode failed to detect NO release in the absence of both oxygen and nitrite.

Interestingly, the electrode showed release of significant quantities of NO under anoxic conditions if nitrite anions were present in the solution (cf Fig. 1). This formation of NO could be confirmed by two other independent methods: First, the optical absorbance showed the appearance of a shoulder near 440 nm and a resolved red-shifted band at 560 nm (cf Fig. 2). These bands indicate the formation of nitrosylated ferrous heme [52]. Second, NO radicals could also be detected in the exhaust gas flow after NO-trapping with Fe-DETC (Fe-diethyldithiocarbamate) complexes and detection of the paramagnetic NO—Fe2+-DETC adducts with EPR (cf Chapter 18). Monitoring the NADPH concentration at 340 nm showed that NO was released from nitrite under consumption of NADPH. This new nitrite reduction pathway of eNOS proceeded in Tris buffer (50 mM, pH = 7.5).

Three aspects distinguish the nitrite reduction pathway from the regular NO release via the arginine pathway: First, the nitrite reduction proceeds in the absence of oxygen. Second, the reduction is slowed but not blocked by the removal of BH4. And finally, using 15N-labeled nitrite proved that the NO was generated from nitrite anions rather than the unlabeled arginine. The 15N and 14N isotopes carry nuclear spin V2 and 1, respectively. As explained in Chapter 18, the difference in nuclear spin determines the multiplicity of the hyperfine splitting observed in the EPR spectra of the paramagnetic NO—Fe2+-DETC complexes formed upon NO trapping (Fig. 3).

Fig. 1. Left panel: Electrode traces obtained after injection of 70 ^M NADPH into a solution containing 500 ^M nitrite in the absence (bottom) or presence of 2 ^M eNOS (intermediate), and eNOS with 100 ^M nitrite (upper trace), 50 mM Tris buffer at pH = 7.6, 1 mM arginine, 12 ^M BH4, 8 ^M calmodulin. An electrode current of 195 ± 20 pA corresponds to 1 ^M NO; Right panel: Correlation between the NO concentration (black squares, electrode current) and the nitrite consumption (crossed circles, Griess reaction) from eNOS catalyzed nitrite reduction as a function of the initial nitrite concentration; the data have been normalized to 4-5 ^M eNOS; same conditions as in the left panel.

Fig. 1. Left panel: Electrode traces obtained after injection of 70 ^M NADPH into a solution containing 500 ^M nitrite in the absence (bottom) or presence of 2 ^M eNOS (intermediate), and eNOS with 100 ^M nitrite (upper trace), 50 mM Tris buffer at pH = 7.6, 1 mM arginine, 12 ^M BH4, 8 ^M calmodulin. An electrode current of 195 ± 20 pA corresponds to 1 ^M NO; Right panel: Correlation between the NO concentration (black squares, electrode current) and the nitrite consumption (crossed circles, Griess reaction) from eNOS catalyzed nitrite reduction as a function of the initial nitrite concentration; the data have been normalized to 4-5 ^M eNOS; same conditions as in the left panel.

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