by Eben van Tonder
17 Septembey 2022

Introduction

We know that nitric oxide is the curing molecule. It’s time for a review of the various reactions associated with it since meat curing is now possible by directly creating it inside the meat matrix without the need to use sodium nitrite. The various reactive nitrogen species are, however, closely linked, and very often, where you find one, you find the others.

These are my notes from the work of Peter C. Ford, B.S., M.S., PhD.. Dr Ford is the Managing Director at the University of California, Santa Barbara. His areas of specialty are catalysis, chemical biology, and chemical kinetics.

The first paper is published by Dr Ford with Katrina M. Miranda, “The solution chemistry of nitric oxide and other reactive nitrogen species.” Here I make notes to myself as to what is possibly relevant to the chemistry of meat curing.

Nitric oxide: physical and chemical properties

“The NO synthases (NOS) are themselves ferriheme proteins.”

“The interactions of NO with hemes extend across all major protein types.”

“Recent developments have suggested that nitrite may serve as a source of NO, especially under hypoxia” and a cautionary follows, that despite “the reactions of NO with metals, typically being regulatory in nature, the interactions with O₂ and free radicals often lead to the formation of harm causing reactive nitrogen species (RNS) such as NO₂ and N₂O₃ ” that induce physiological processes associated with disease or injury.

NO is more soluble in organic solvents than in water.

Nitric oxide is a colorless gas. NO₂ is brown and liquid N₂O₃ is
blue.

The simplified reaction of O₂ with NO produces NO₂, N₂O₃, and N₂O₄ in aprotic media and NO₂^– in water.

Nitrite formation acidifies aqueous solutions.

NO is neither readily oxidized nor reduced by one-electron processes, and this feature is an important factor in maintaining a lifetime sufficiently long for NO to diffuse through cells.

Reactivity and kinetics

– NO autoxidation

Autoxidation is the oxidation of a species by O2. Atmospheric NO is well known to undergo autoxidation to NO2. In aqueous media NO autoxidation leads instead to the formation of nitrite, the general stoichiometry being shown in eq.

4 NO + O₂ + 2H₂O → 4 H⁺ + 4 NO₂^–

The reaction is irreversible and provides a measurable end product that aids analysis. Two independent kinetics studies in 1993 demonstrated that, although the reaction in aqueous solution gives nitrite instead of nitrogen dioxide, the rate law remains third order. The reaction rate is relatively independent of both the solution pH and the temperature.

Several subsequent studies of this reaction under analogous conditions confirmed the form of the rate law and determined similar rate constants. However, there are, as yet, unresolved issues regarding the nature of the reactive intermediates formed during NO autoxidation.

The discovery that NO is a regulatory agent in the cardiovascular system occurred nearly concurrent with the discovery that NO elicits cytotoxicity during immune response to pathogens. These seemingly incongruent effects of NO raises the obvious question of how NO can be essential or beneficial in certain functions yet be toxic in others. The answer is in part evident from a fundamental toxicological/ pharmacological perspective in which dose response curves define the concentration ranges for deficiency, normal function and toxicity. Also crucial are the kinetics of the competing pathways depleting NO. The autoxidation reaction serves as a prime example of how both timing and concentration are critical factors in defining the effects of NO biosynthesis.

For example, if NO is consumed by autoxidation, how can it function as a signaling agent in normoxic arterial tissues? The answer lies in the third order kinetics in aqueous media, first order in [O₂] and second
order in [NO]. Consequently, the rate of the reaction is especially sensitive to the NO concentration.

NO and O₂ are more soluble in hydrophobic media such as lipids, and as a consequence, both tend to partition into such cellular regions. Thus, NO autoxidation may preferentially occur in lipid membranes rather than the cytoplasm, producing NO₂ and N₂O₃ rather than nitrite.

While it is clear that the reactive intermediates formed during the autoxidation of aqueous NO have cytotoxic and mutagenic activities during immune response, these have not been well defined. The
autoxidation mechanism is commonly assumed to parallel the gas-phase reaction, where the rate-limiting step produces NO₂, which is trapped by NO to give N₂O₃. Hydration of the latter would give nitrous acid. Wink et al. attempted to identify the key autoxidation intermediates by examining relative reactivities with trapping agents such as ferrocyanide, 2-2′ -azinobis(3-ethylbenzthiazoline-6-sulfonic acid) (ABTS), azide and GSH. Comparing the results of competition studies with reported reactivities, led to the conclusion that the relevant oxidant generated during aqueous NO oxidation is not NO₂. Kinetics simulations attempted to model these results using the published rate constants for NO₂ reactions with such trapping agents came to a similar conclusion, although an alternative mechanism was not proposed.

– The reactivity of other nitrogen oxides

Since NO is neither readily reduced nor oxidized by simple one-electron processes, both the production of NO and the conversion of NO into other nitrogen oxides is typically driven by metalloenzymes or by reactions with other free radicals or O₂.

Nitrogen-containing species that are more reduced than NO are often precursors of NO. The most common process for biosynthesis of NO begins with two-electron oxidation by nitric oxide synthase (NOS) of L-arginine to N^ω-hydroxy-L-arginine (NOHA), converting a guanidino nitrogen atom from the oxidation state of an amine (- 3) to that of a hydroxylamine (- 1). this is followed by a three-electron oxidation resulting in release of NO (+2). The primary end-products of mammalian NO biology are nitrite (+3) and nitrate (+5), so the eventual fate of NO in vivo is further oxidation.

– Oxidation state +1: HNO, NO^– and N₂O

The slow rate of proton transfer dictates that NO^–and HNO should effectively function independently, although of the two, HNO is expected to be the major species under physiological conditions. While there has been considerable speculation about endogenous production of nitroxyl, there are known and emerging pharmacological applications of exogenously applied HNO in treatments of alcoholism, sickle cell disease, heart failure, cancer, and pain. Such studies have established that HNO can be produced by metabolic pathways.

Of potential biosynthetic pathways for HNO production, perhaps the most intriguing is by NOS under hypoxic or low cofactor conditions or by oxidative degradation of the intermediate NOHA. Rather than a three-electron oxidation needed to generate NO from NOHA, HNO production would involve a two-electron oxidation of that intermediate.

Hypoxia reduces the intensity of the HNO response, potentially offering a protective mechanism toward ischemia/reperfusion injury in the brain.

Multiple comparisons of NO or HNO have shown that in nonvasoactive assays the responses to these two nitrogen oxides are generally discrete, due to distinct chemical modifications. HNO is in many cases more reactive than NO. For example, it undergoes rapid dimerization, forming hyponitrous acid, which then dehydrates irreversibly to nitrous oxide (eq. (31)). This is a major sink for HNO, and the high second order rate constant of self-consumption (k₂ = 8 × 10⁶ M^-1s^-1 in ambient aqueous solution) precludes the storage of HNO. The N₂O produced has often been cited as an indirect marker of HNO formation.

2 HNO → [HONNOH] → N₂O + H₂O

In addition to dimerization, HNO can be consumed in a cellular environment by a number of molecular species. In particular, reaction with a oxidizing metal center will convert HNO into NO (eq. (32)). HNO and NO can also react with each other at a near diffusion-limited rate to give the radical HN₂O₂, which may provide a rapid mechanism to switch from HNO to NO signaling.

– Oxidation state +3: NO⁺, NO₂^– , N₂O₃ and N-Nitrosoamines

The chemistry of this oxidation state is dominated by nitrosative modifications in which NO⁺ is formally donated to a nucleophile. Under physiological conditions, NO⁺ itself is at best a transient species since the one-electron reduction to NO is so highly favorable and NO⁺ reacts rapidly with water to form nitrous acid.

NO⁺ + H₂O⇄HNO₂ +H⁺

The equilibrium constant for the above eq. has been estimated at 0.7 × 10⁸ M, such that NO⁺ is a prominent species only in highly concentrated acid (e.g., 60% sulfuric acid).

The identification of NO⁺ in concentrated sulfuric acid solutions is possible only because the activity of the acid is so high and that of H₂O is so low under these conditions.

Nitrous acid is a relatively weak acid (eq. below). The most recent reevaluation gave the pK_a of HNO₂ as 3.16 at 25 ^◦C (3.11 at 37 ^◦C). Therefore, at physiological pH, the conjugate base nitrite ion is ~99.99% of the total distribution. The equilibrium constants noted above can be used to estimate the concentration of NO⁺ in a 1 mM NO₂^– solution at pH 6 to be ~10^-22 M. Thus, despite occasional suggestions in the literature that free NO⁺ is participating in nitrosation reactions, the direct involvement of NO⁺ seems unlikely to have mechanistic validity in an aqueous environment, except in concentrated acids.

HNO₂ ⇄ H⁺ + NO₂^–

Solutions of nitrite in dilute aqueous acid are well known to promote nitrosation. Currently, there are several scenarios under consideration for this chemistry:

The first involves a protonated nitrous acid, H₂NO₂⁺, while the second involves dehydration of HNO₂ to produce the powerful nitrosating reagent N₂O₃.

2 HNO₂⇄ H₂O + N₂O₃

The N₂O₃ anhydride of nitrous acid can dissociate to NO and NO₂. The net result of the following two eq. is the acid promoted disproportionation of nitrite to nitric oxide and nitrogen dioxide.

2 HNO₂⇄ H₂O + N₂O₃

and

Acid promoted nitrite disproportionation has been suggested to be a viable pathway toward NO formation from NO2^–, especially in the acidic fluids of the stomach.

There has long been a concern that nitrate and nitrite in the diet promotes colorectal and other cancers owing to amine nitrosation to give N-nitroso compounds (NOCs). The type of amine impacts the outcome in that secondary amines are often stably nitrosated, while N-nitrosated primary amines undergo facile deamination.

The rate of NO production and the NO concentration in equilibrium with dissolved NO₂^– is a complex function of pH and nitrite concentration.

For example, at equilibrium, a 10 μM NO₂^– solution at pH 6 would contain ~5 nM NO and an equal concentration of NO₂, if there are no other pathways depleting one or the other of these species. At pH 7.4 this number drops to ~200 pM. However, these are dynamic equilibria, so if for example NO₂ were removed by a trapping agent, higher concentrations of NO would result (Le Chatelier’s principle). Thus, in evaluating the potential biological impact of HNO₂ disproportionation on steady state concentrations of NO, one must consider not only these equilibria, but also the holistic dynamics of all individual processes that deplete or enhance the concentrations of the key species involved.

N₂O₃ not only is a potential source of NO and NO₂ from nitrite disproportionation, but also is a powerful nitrosating reagent, hence it is a RNS in its own right. Again, the word “nitrosation” represents a chemical reaction where (formally) NO+ is added to a nucleophile (X^– ) or replaces an H⁺ of a compound XH. Given the extremely short lifetime of NO⁺ under biological conditions, NO⁺ must be directly transferred in nitrosation reactions, rather than released and captured in separate steps. A particularly valuable resource on this topic is a monograph by Williams.

Biologically, the two most important targets for N₂O₃ would be thiols and amines (eq below) and it is likely that nitrosation by N₂O₃ occurs by direct transfer of NO⁺ to a nucleophile from N₂O₃. However, it should be emphasized that, in addition to such metal-free nitrosation processes, NO in the presence of a redox active metal center will also nitrosate such nucleophiles

N₂O₃ + RSH → RSNO + H⁺ + NO²

N₂O₃ + RR’NH → RR’N(NO) + H⁺ + NO₂

When initiated by nitrite, the kinetics for nitrosation of nucleophiles such as amines and thiols follow the rate law described in eq.

rate = k’ [HNO₂]²[X^–]

The second order dependence on nitrous acid concentration is interpreted as reflecting the formation of N₂O₃ according to eq.

2 HNO₂⇄ H₂O + N₂O₃

According to this interpretation, k’ = k_x × KN₂O₃, with k_x being the rate constant for the direct reaction of N₂O₃ with X^– and K_N2O3 being the equilibrium constant for eq. above (3 × 10^-3 M^-1). The value of k_x depends on the identity of the nucleophile. For amines, k_x is in the 10⁷-10⁹ M^-1s^-1 range while values ranging from 3 × 10⁵ to 7 × 10⁷ M^-1 s^-1 have been reported for GSH.

Although the values for k_x are large, one must keep in mind that the equilibrium concentration of HNO₂ ([HNO₂] = K_a^-1[NO₂^– ][H⁺]) is very low at physiological pH. For example, in a 1 mM nitrite solution at pH 7, the concentration of HNO₂ would be < 10^-7 M, so significant nitrosation promoted by equilibrium concentrations of N₂O₃ formed by nitrite disproportionation would not be expected near neutral pH. On the other hand, N₂O₃ generated transiently in a chemical or biological process could contribute to nucleophile nitrosation pathways occurring in competition with hydrolysis to nitrous acid (see below).

rate = k’ [HNO₂]²[X^–]

Addition of NO to aerated solutions containing such nucleophiles can also lead to nitrosation and this can be attributed to the generation of N₂O₃ intermediates formed during NO autoxidation. For example, alkyl nitrites dissolved in aerated organic solvents are formed when alcohols are treated with NO. Such reactions of N₂O₃ generated by NO autoxidation are likely to be a major source of amine and thiol nitrosations and subsequent transformations during immune response. Furthermore, since biological experiments with endogenously generated or exogenously added NO are typically carried out under normoxic conditions, it is important to consider nitrosation as a possible outcome.

Hydrolysis of N₂O₃ to nitrous acid:

N₂O₃ + H₂O → 2H⁺ + 2 NO₂^–

the reverse of eq:

2 HNO₂⇄ H₂O + N₂O₃

occurs with a rate constant of 2 × 10⁴ s^-1 so a lifetime of ~50 μs might be expected for N₂O₃ that is generated in an aqueous solution. The time frame is likely to be much longer in a hydrophobic cellular environment where the activity of water would be much lower. Thus, nitrosation of thiols and amines by NO autoxidation may have biological relevance, particularly under conditions where the concentration of NO is high enough that autoxidation is a significant sink for NO.

An interesting feature of N₂O₃ is that it has several isomers that may be chemically important. Computational studies indicate that the asymmetric isomer ON-NO₂ is the most stable, but the other two species depicted are only slightly less so, and the three can interchange readily. Each isomer has been argued to be capable of nitrosation reactions, although it is quite possible that the selectivities
might differ, as has been suggested.

Amine nitrosations have long been of interest owing to concerns about the potentially carcinogenic character of N-nitrosoamines when meats are preserved by adding nitrite or nitrate salts. Such N-nitrosation can also be mutagenic, for example upon deamination of nitrosated exocyclic primary amines in DNA. Again, N-nitrosoamines are most likely formed endogenously under conditions of immune response or in the acid conditions of the stomach. A recent development was the findings in 2018 and 2019 that certain commonly used medications contained N-nitrosodimethylamine (NDMA) and related nitrosoamines as impurities at levels sometimes significantly exceeding the Federal Drug Administration limits. The likely sources of this impurity lay in changes to the manufacturing protocols.

– Oxidation state +4: NO₂ and N₂O₄

Although, acidic nitrite solutions are a means to NO formation via the disproportionation reactions depicted in eqs.

2 HNO₂⇄ H₂O + N₂O₃

and

the coproduct of this reaction is nitrogen dioxide. This free radical is also produced from NO autoxidation, at least in hydrophobic media

2 NO + O₂ → 2 NO₂

from decomposition of peroxynitrite and by metal catalyzed oxidation of nitrite. The chemical and physiological roles of NO₂ have not received the attention given to those of nitric oxide, but may be quite significant.

NO₂ is a very strong one-electron oxidant with a one electron reduction potential of +1.04 V in aqueous media

NO₂ + e^– → NO₂^– E^◦ = +0.895 V (in aq. solution)

As a consequence, it readily undergoes redox reactions with various biological reductants such as ascorbate (k₂ = 1.8 × 10⁷ M^-1 s^-1 at pH 6.5, ferrocytochrome c (6.6 × 10⁷ M^-1 s^-1) and thiols and
thiolates (5 × 10⁷ M^-1 s^-1 for cysteine at pH 7.4, the reaction being largely attributed to the thiolate form). The reversible one-electron oxidation of tyrosine by NO₂ is of particular interest since elevated levels of 3-nitrotyrosine have been detected in a variety of inflammatory disease states.

NO2 + e^– → NO₂^– E^◦ = +0.895 V (in aq. solution)

The second order pH-dependent rate constants for the reaction of NO₂ with tyrosine have been reported (3.2 × 10⁵ M^-1 s^-1 at pH 7.5 and 2.0 × 10⁷ M^-1 s^-1 at pH 11.3). Tyrosine reacts even more rapidly with the carbonate radical at physiological pH (k₂ = 4.5 × 10⁷ M^-1 s^-1). In this context, subsequent reaction of the resulting tyrosine radical with NO₂ (k₂ = 1.3 × 10⁹ M^-1 s^-1) would be a potential pathway for the formation of 3-nitrotyrosine.

Unlike NO, the free radical NO₂ is prone to dimerization

2 NO₂⇄ N₂O₄

The second order rate constant for dimerization is 4.5 × 10⁸ M^-1 s^-1, and the equilibrium constant is 7.7 × 10⁴ M^-1 in water. The NO₂/N₂O₄ mixture has been used extensively to modify organic compounds. The dimer can serve as a nitrating (NO₂⁺) or nitrosating (NO⁺) agent in relatively polar, non-aqueous media or in highly concentrated acid solutions

N₂O₄ ⇄ NO₂⁺ + NO₂^–

N₂O₄ ⇄ NO⁺ + NO₃^–

The equilibrium between NO₂, N₂O₄ and the related ions is dependent on solvent, with heterolytic cleavage being significant only in polar solvents. In nonpolar media, the modifications are assumed to arise from direct reaction with N₂O₄ in a similar fashion to that for N₂O₃. Exposure to NO₂/N₂O₄ often leads to a mixture of nitrated products, with nitrosated and oxidized species also possible.

In aqueous solution, the dimer undergoes fairly rapid disproportionative hydrolysis to nitric and nitrous acid, (k = 1 × 10³ s^-1), such that its inherent lifetime in water is short, ~1 ms. Given the second order dependence of the concentration of N₂O₄ on NO₂, which is highly reactive, and the sink represented by the hydrolysis pathway

N₂O₄+ H₂O ⇄ 2 H⁺+ NO₃^– + NO₂^–

the role of the dimer in physiological nitration pathways may be limited. However, N-nitrosation of both primary and secondary amines in neutral or alkaline conditions has been observed to outcompete hydrolysis, although the yields are typically quite poor compared to organic media.

-> Oxidation state +5: NO₃^– , NO₂⁺ and ONOO^–

Nitrate (NO₃^– ) is a planar ion that is isoelectronic and structurally analogous to carbonate (CO²/₃^– ). HNO₃ is a very strong acid (pK_a = – 1.3), so, unlike nitrite, nitrate is never protonated under physiologically relevant conditions. Nitrate is often viewed as a physiological endproduct of nitrogen metabolism and is excreted in the urine. Notably, the higher nitrate concentrations in the urine of patients with bacterial infections were an important clue to Tannenbaum, Hibbs and others that elevated NO concentrations are generated during the immune response.

Although certain mammalian enzymes such as xanthine oxidoreductase have been shown to exhibit nitrate reductase activity such reactions do not appear to play a major role in mammalian systems. As noted above, excessive dietary nitrate was generally considered to be a harmful component of food that caused infantile methemoglobinemia and carcinogenesis.

However, more recent discoveries that certain high nitrate foods such as beets and green leafy vegetables have a positive effect in lowering blood pressure and that nitrated fatty acids are endogenous anti-inflammatory signaling mediators (NO₂-FA) have generated reevaluations of those concerns. An explanation of the former effect has been that nitrate in circulating fluids is transported to the mouth where it concentrates in the saliva. Oral bacteria have nitrate reductase activity that converts nitrate to nitrite, which is then swallowed. The acidity of the stomach fluids then results in disproportionation of the resulting nitrous acid, and the NO equivalents generated may be transferred to the cardiovascular system leading to vasodilatory effects. The mechanism of such transfer is uncertain, but it has been proposed that antihypertensive effects of orally administered nitrite or nitrate result from formation of gastric S-nitrosothiols.

Peroxynitrite: The peroxynitrite ion is an isomer of nitrate, but formation of ONOO􀀀 has been related to the pathogenesis of multiple diseases

The principal source of ONOO^– is the very fast reaction
of NO with O₂^– according to the following reaction

O₂^– + NO → ONOO^–

Both NO and O₂^– are produced during the inflammatory immune response. For this reason, the chemistry, biochemistry and biology of ONOO^– have drawn considerable attention and has been reviewed extensively elsewhere by

J. S. Beckman, W.H. Koppenol, Nitric oxide, superoxide, and peroxynitrite: the good, the bad, and the ugly, Am. J. Physiol. 271 (5) (1996) C1424–C1437. Pt. 1.
C. Szabo, H. Ischiropoulos, R. Radi, Peroxynitrite: biochemistry, pathophysiology and development of therapeutics, Nat. Rev. Drug Discov. 6 (2007) 662–680;
S. Bartesaghi, R. Radi, Fundamentals on the biochemistry of peroxynitrite and protein tyrosine nitration, Redox Biology 14 (2018) 618–625

G. Ferrer-Sueta, N. Campolo, M. Trujillo, S. Bartesaghi, S. Carballa, N. Romero, B. Alvarez, R. Radi, Biochemistry of peroxynitrite and protein tyrosine nitration, Chem. Rev. 118 (2018) 462;
O. Augusto, S. Goldstein, J.K. Hurst, J. Lind, S.V. Lymar, G. Merenyi, R. Radi, Carbon dioxide-catalyzed peroxynitrite reactivity – the resilience of the radical mechanism after two decades of research, Free Radic. Biol. Med. 135 (2019) 210–215;
W.H. Koppenol, P.L. Bounds, T. Nauser, R. Kissner, H. Rüegger, Peroxynitrous acid: Controversy and consensus surrounding an enigmatic oxidant, Dalton Trans. 41 (2012) 13779–13787;
W.H. Koppenol, S. Serrano-Luginbuehl, T. Nauser, R. Kissner, Thinking outside the cage: A new hypothesis that accounts for variable yields of radicals from the reaction of CO2 with ONOO-, Chem.Res.Toxicol. 33 (2020) 1516–1527.

The brief discussion here will focus on reactions of ONOO^– in aqueous media.

The pK_a of ONOO^– is reported to be ~6.8, so it is significantly protonated at normal physiological pH values. Peroxynitrous acid (ONOOH) is very reactive, and one pathway for decomposition is via
unimolecular isomerization to nitrate. A secondary process involves decay through homolytic cleavage of the peroxide bond to give NO₂ and the hydroxyl radical. Cage recombination of these two radicals would provide another mechanism for isomerization to nitrate, but escape from the solvent cage of these species could lead to significant damage. The extent of this process has been reviewed elsewhere.

Another pathway to ONOO􀀀 decomposition is the reaction with carbon dioxide (k2 = 3 × 104 M􀀀 1 s􀀀 1) to form the nitrosoperoxycarbonate ion ONOOCO2􀀀 (Scheme 5). Given the millimolar concentrations of CO2 in the cytosol, this reaction rate is sufficient to significantly consume ONOO􀀀 formed during inflammation. One decay pathway is to undergo rearrangement to the nitrocarbonate ion O2NOCO2􀀀 , followed by hydrolysis to nitrate and bicarbonate. The CO2 adduct is also inherently unstable toward homolytic cleavage of the peroxo O–O bond to NO2 and the carbonate radical anion, CO3•–, and this pathway accounts for much of the overall decay of ONOO􀀀. Both NO2 and CO3•– are strong oxidants and while they are readily trapped by GSH (k2 = 2 × 107 and 5 × 106 M􀀀 1 s􀀀 1 at pH 7.4, respectively), they react at comparable rates with tyrosine to give Tyr• radicals. The very fast reaction of Tyr• with NO2 (k = 1.3 × 109 M􀀀 1 s􀀀 1) is a likely predecessor to the formation of nitrotyrosine residues (Scheme 6). Protein nitration is a marker for several diseases, such as amyotrophic lateral sclerosis (ALS or Lou Gerig’s disease). Lim et al. have published an extensive kinetics analysis of the reactions of ONOO􀀀 under physiologically relevant conditions and concluded that the radical scavenging ability of ascorbate and GSH left 3-nitrotyrosine as the only tyrosine derivative that is likely to be formed at a significant rate.

The solution chemistry of nitric oxide and other reactive nitrogen species