Extracts from the work of Dr. Peter C. Ford and its possible aplication to meat science

by Eben van Tonder
17 Septembey 2022


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., Ph.D.. dr Pord is Managing Director at the University of California, Santa Barbara. His areas of speciality is 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.

“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 N2O3that 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

The simplefied 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 Nω-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 NOand 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 pKa 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


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’ = kx × KN2O3, with kx being the rate constant for the direct reaction of N2O3 with X and KN2O3 being the equilibrium constant for eq. above (3 × 10-3 M-1). The value of kx depends on the identity of the nucleophile. For amines, kx 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 kx are large, one must keep in mind that the equilibrium concentration of HNO2 ([HNO2] = Ka-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