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1730, 1770

Free molecule [4,5]

Free molecule [8]

Free molecule [9]

1708 cm-1 for W-methyl-d, [16] l-glucamine dithiocarbamate (MGD)

Two distinct modes with [20] Av ~ 40 cm-1

Nitrosyl ligands are inequivalent [12] with Av ~ 40 cm-1

The UV spectrum of nitric oxide shows absorption below 200 nm due to weak electronic transitions to unoccupied orbitals, and the onset of photoionization to nitrosonium NO+ by ejection of the unpaired electron from the n* orbital. As expected for an antibonding electron, the ionization threshold is fairly low with 9.26 eV. This photoionization produces a prominent photoabsorption peak near 14 eV. For still higher photon energies >20 eV, fragmentation of the molecule occurs [21].

Nitric oxide is a highly corrosive gas with a boiling point of — 151.7°C at 1 atm and is prone to oxidation to nitrogen dioxide radicals when in contact with dioxygen. The reaction chemistry of NO^ is very complex due to several reasons: first and foremost, it is thermody-namically unstable. If kinetically allowed by conditions of high pressure and temperature, it may dismutate via the pathways [22]

The prime example for the first pathway is when pure NO^ is compressed to high density for storage in pressure containers. Interestingly, this reaction appears as a dispropor-tionation as it generates products with higher as well as lower oxidation state than the original NOV A prominent example for the second pathway is the automotive catalytic converter where rhodium catalyzes the reductive decomposition of NO^ into dinitrogen and dioxygen.

The second reason is the fact that many higher oxides of nitrogen have a rather unstable electronic structure: N2O, N2O4 and N2O3 all appear as resonance hybrids resonating between several isomeric forms with different atomic and electronic structures [23]. In addition, NO2 is a radical with tendency to dimerize and HNO dimerizes to metastable hyponitrous acid HONNOH which decomposes into N2O and water [see below, Eq. (6)].

In addition, the analysis of the reaction chemistry and identification of the reaction products is technically challenging since many nitrogen oxides have strong broad overlapping absorptions in the ultraviolet at wavelength shorter than 280 nm. The visible bands are usually better resolved but have very small extinctions (N2O3 with £620 ~ 20 Mcm-1; nitrate NO— with £300 ~ 7 Mcm-1; nitrite NO— with £354 ~ 24 Mcm-1; nitrous acid HNO2 with a characteristic tetrad of peaks at 347, 358, 371 and 386 nm and £371 ~ 54 Mcm-1, pKa = 3.37; NO2 with £385 ~ 30 Mcm-1). This last radical, nitrogen dioxide, is a common industrial and automotive pollutant and is responsible for the orange-brown hue of smog and polluted air. Peroxynitrite anions have a stronger UV absorption with £302 ~ 1704 Mcm-1. The lifetime of peroxynitrite at physiological pH is below a second because of its reactivity towards proteins and propensity to protonate to cis or trans forms of pernitrous acid [24,25]. Far stronger absorptions in the visible region are known for the hyponitrite radical anion N2O-'" (£290 ~ 6 x 103 Mcm-1); its protonized form, the radical HN2O2 (£290 ~ 3 x 103 Mcm-1 pKa = 5.5) [32] and the unstable N3O- anion (£380 ~ 3.8 x 103 Mcm-1) with decay rate of 300 s-1 in water at room temperature [26].

The optical spectra of nitrite and nitrate are shown in Fig. 2.

Nitric oxide itself can readily participate in a wide variety of redox reactions [28]. It may be oxidized to the nitrosonium cation NO+, which is isoelectronic to carbon monoxide CO

Fig. 2. Room temperature spectra of nitrite, nitrous acid and nitrate. The curves are: (1) 22.6 mM NaNO2; (2) 92.8 mM NaNO3; (3) 22.6 mM NaNO2 plus 92.8 mM HNO3; (4) 22.6 mM NaNO2 plus 418 mM HNO3. (From Ref. [27].)

Wavelength (nm)

Fig. 2. Room temperature spectra of nitrite, nitrous acid and nitrate. The curves are: (1) 22.6 mM NaNO2; (2) 92.8 mM NaNO3; (3) 22.6 mM NaNO2 plus 92.8 mM HNO3; (4) 22.6 mM NaNO2 plus 418 mM HNO3. (From Ref. [27].)

and N2. It has a prominent UV absorption (£220 ~ 3850 Mcm 1 [29]). The oxidation reaction is

The reduction potential of 1.21 V is quite high and exceeds range of the typical reduction potentials found in vivo. The latter range from -0.3 V (NAD+/NADH) to +0.82 V (O2/H2O). Therefore, nitrosonium is not generated by simple oxidation of nitric oxide radicals. Acidic reduction of nitrite anions provides an alternative pathway [Eq. (3b)]

However, in the normal physiological range pH ~ 7.0 -7.4, this equilibrium is shifted to the far left side [28] with [NO+]/[NO-] ~ 10-17 approaching infinitesimally small values. In biological systems, the formation of nitrosonium is thought to be dominated by heterolytic fission of N2O4 [see below Eq. (8)]. However, even when formed, nitrosonium is short lived in the presence of water. Reaction (3b) shows that the nitrosonium cation reacts rapidly with water to nitrite. It may also react with other nucleophiles, and can nitrosate proteins in biological systems. Other nitrosating pathways will be discussed below [cf Eq. (10)].

Alternatively, nitric oxide radicals may be reduced to the nitroxyl anion NO- which is isoelectronic to molecular oxygen. Having two half-occupied orbitals, this biradical may exist [30,31] as a spin triplet (S = 1) or as a spin singlet (S = 0)

Just as with the dioxygen molecule, the spin triplet 3NO- is ground state, with the spin singlet 17-20 kcal/mol (0.74-0.87 eV) higher in energy.

The two spin states of the nitroxyl anion have markedly different reaction rates with nitric oxide to form the hyponitrite radical N2O22-, or even the bluish trinitrogen trioxide anion N3O3- [32]. The rates are strongly affected by the constraints imposed by spin conservation. The protonation of the ground state triplet nitroxyl is kinetically slow as it requires spin conversion between singlet HNO and triplet NO- as well as a nuclear reorganization:

Whereas older literature often quotes a low equilibrium constant of pKa ~ 4.7, newer estimates [26,31] favor a sharply upward revision to a value of pKa ~ 11 for the protonation of 3NO-. This value makes the protonized HNO the predominant species at physiological pH. At higher concentrations, the chemistry of HNO is complicated by its tendency to dimerize and decompose into dinitrogen oxide N2O and water. The irreversible decomposition has a high rate of k = 8 x 106 (Ms)-1:

The nitroxyl anion NO- may act directly on a range of biological molecules [33,34]. Angeli's salt Na2N2O3 is a well-known water-soluble nitroxyl donor.

In water, NO^ has a solubility [35] of 1.9 mM at 25°C and against a Pno of 1 atm (Fig. 3 and Table 2). The solubility is somewhat higher than that of dioxygen.

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