On the Hydrophobicity of Nitrogen Dioxide: Could there be a “lens” effect for NO₂ reaction kinetics?

Abstract

Solvent “lens” effects for the reaction kinetics of NO₂ can be evaluated on the basis of published Henry’s law constants for nitrogen dioxide in various solvents. Water-to-organic solvent partition coefficients were derived from Henry’s law constants and used to assess the tendencies of NO₂ toward fleeing the aqueous environments and concentrating in biological hydrophobic media. It is concluded, based only on the estimated aqueous medium-to-cell membrane partition coefficient for NO₂, that such tendencies will be relatively small, and that they may account for an acceleration of chemical reactions in biological hydrophobic media with reaction kinetics that are first order on NO₂ by a factor of approximately 3 ± 1. Thus, kinetic effects due to mass action will be relatively small but it is also important to recognize that because NO₂ will tend to dissolve in cell membranes, reactions with cell membrane components will not be hindered by lack of physical solubility at these loci. In comparison to other gases, nitrogen dioxide is less hydrophobic than NO, O₂ and N₂.

Introduction

Exposures to nitrogen dioxide (NO₂) can occur in industrial and occupational settings, and during war and acts of terror involving the use of explosives. Moreover, NO₂ is present in polluted air [1–4] and also can be formed endogenously [5–11]. The reaction of peroxynitrite with carbon dioxide, reactions catalyzed by myeloperoxidase using nitrite as the electron donor, and acidification of nitrite followed by oxidation of the resulting nitric oxide (NO) are among the biological processes that can result in the endogenous formation of NO₂ [5;7–9;12]. Nitrogen dioxide is involved in biological oxidations, including the nitration of protein tyrosine residues and possibly also the nitration of membrane lipids, reactions with low molecular weight antioxidants, and the oxidation of thiol residues. Cell signaling and covalent modifications mediated by NO₂ will depend in part on its physicochemical properties relevant to passive transport, such as its ability to diffuse and penetrate into membranes, its relative solubility in membranes and aqueous solutions, and its ability to chemically react within the membranous or aqueous milieu. Despite the biological importance of NO₂, it has been recently pointed out that the relative solubility of NO₂ in membranes and aqueous medium is unknown [13–15], revealing the lack of even the most essential parameter necessary to interpret the reactivity of NO₂ across different cellular and biological compartments. In contrast, the solubility of NO in several solvents has been studied in detail and this has led to the proposal for the existence of an acceleration or “lens” effect [13;14] whereby the hydrophobicity of NO helps focus some of the chemical reactivity of NO, such as its autoxidation, in cell membranes and lipid particles. This brief exercise presented here represents an attempt to estimate the relative solubility of NO₂ in biological membranes and any possible ensuing acceleration or “lens” effects in NO₂ reaction kinetics.

Methods

Solvent “lens” effects [13;14] for the reaction kinetics of NO₂ are evaluated herein on the basis water-to-organic solvent partition coefficients ( $K_w^s$) as these coefficients may result in high, compartmentalized concentrations of NO₂, that thereby increase the rates of reaction of NO₂ in these compartments. Values of $K_w^s$ were derived from Henry’s law constants for NO₂ ( $K_H^s$) calculated from published solubility studies of gaseous NO₂ in various solvents [16;17] (see equations 5, 6 and 8 below, for operations applied in this process). When necessary, data were corrected for temperature to 37 °C using equation 7 and published enthalpies of solution (Δ_sol H) for NO₂ in various solvents [17;18].

Results

On Solvent “Lens” Effects in Reaction Kinetics

Moller et al. studied the acceleration of the autoxidation of NO when NO is allowed to partition between water and organic media [13]. Since under a wide range of experimental conditions the rate law for the autoxidation of NO was found to be

$$-\frac{d[NO]}{dt}=4k{[NO]}^2[O_2]$$ (1)

and diffusion across compartments is not rate-limiting, these authors interpreted the acceleration they observed as arising from the larger concentration of NO in the organic media resulting from the hydrophobicities of NO and O₂ [13]. It is important to recognize that solvent effects on the rate constant k could either amplify or dampen effects due to reactants’ hydrophobicities since the rate constant is subject to solvent effects. However, it is probable that hydropobicity is the dominant factor in the “lens” effect since the rate constant in equation (1) in the solvent carbon tetrachloride has been found to be close to that in aqueous solution [14;19].

On Estimating the Solvent “Lens” effect for NO₂

Studies on the solubility of NO₂ are experimentally challenging and it is not surprising that related literature is rather scarce. The solubility studies for NO₂ are complicated by the reactivity of NO₂ and N₂O₄ with many solvents. [Caution: The authors are aware of one explosion with fire that occurred in a research laboratory while using NO₂ in tetrahydrofuran.] It is also important to consider that NO₂ dimerizes to form N₂O₄ and the latter can react with traces of water present in solvents (Equations 2–4).

$$2{\text{NO}}_2\rightleftharpoons {\text{N}}_2{\text{O}}_4$$ (2)

$${\text{NO}}_2/{\text{N}}_2{\text{O}}_4+\text{Solvents}\to \text{Products}$$ (3)

$${\text{N}}_2{\text{O}}_4+{\text{H}}_2\text{O}\to {\text{NO}}_2^{-}+{\text{NO}}_3^{-}+2{\text{H}}^{+}$$ (4)

Despite the complications arising from the hydrolysis reaction (Equation 4), water is perhaps the most studied solvent, owing to the environmental importance of NO₂, and a number of reactive uptake studies have been conducted in water from which Henry’s law constants for NO₂ in water can be derived [18;20]. In contrast, few gaseous NO₂-organic solvent partitions [16;17] had been measured and apparently, no organic solvent-to-water partition coefficients for NO₂ had been reported to date.

The solubility of a gas in a solvent can be described in different ways, but usually, the Henry’s law constant $K_H^s$ is used to this effect. Thus, for NO₂, we can write equation (5):

$$K_H^s=\frac{{[{NO}_2]}_s}{p_{{NO}_2}}$$ (5)

where [NO₂]_s is the molar concentration of NO₂ in organic solvent s and p_NO2 is the partial pressure of NO₂ in the gas phase in units of atmospheres. When the solvent s is water (w), we write $K_H^w$. Water-to-organic solvent partition coefficients ( $K_w^s$) can then be calculated by dividing the Henry’s law constant for organic solvent ( $K_H^s$) by the equivalent for water ( $K_H^w$) using equation (6):

$$K_w^s=K_H^s/K_H^w=\frac{{[{NO}_2]}_s}{p_{{NO}_2}}/\frac{{[{NO}_2]}_w}{p_{{NO}_2}}={[{NO}_2]}_s/{[{NO}_2]}_w$$ (6)

Cheung et al.[18] compared literature values for $K_H^w$ for NO₂ with values they obtained and recommend $K_H^w=1.4\pm 0.2\times {10}^{-2}\text{M}{\text{atm}}^{-1}$ for 20 °C and $K_H^w=2.3(+0.3-0.9)\times {10}^{-2}\text{M}{\text{atm}}^{-1}$ for 3 °C. From these two values, using equation (7), one can calculate an approximate enthalpy of solution (Δ_sol H) for NO₂ in water of −20 kJ mol⁻¹ and extrapolate $K_H^w=9.1\times {10}^{-3}\text{M}{\text{atm}}^{-1}$ for 37 °C.

$$K_H^w(T_2)=K_H^w(T_1)\cdot exp\left(\frac{-{\mathrm{\Delta }}_{\mathit{sol}}H}{R}\left(\frac{1}{T_2}-\frac{1}{T_1}\right)\right)$$ (7)

We identified two published solubility studies for gaseous NO₂ in organic solvents[16;17] and the data published in these articles are summarized in Table 1. Mendiara et al. published Henry’s law constants in terms of molar fractions and in units of atm ( $k_{{NO}_2}^H$) [16] and these were converted to $K_H^s$ (units of M atm⁻¹) using equation 8, which is valid for dilute solutions, where MW^s and ρ^s represent the molecular weight and the density of solvent s, respectively.

Table 1.

Henry’s law constants for nitrogen dioxide in various solvents and water-to-organic solvent partition coefficients at 37 °C. From left to right, second column: Henry’s law constants $K_{\text{H}}^{\text{s}}$ (M atm⁻¹) at temperatures specified in the corresponding reference; third column: Henry’s law constants $K_{\text{H}}^{\text{s}}$ for 37 °C calculated using Eq. (7) as indicated in the text; fourth column, water-to-organic solvent partition coefficients $K_{\text{w}}^{\text{s}}$ calculated using Eq. (6) with values for $K_{\text{H}}^{\text{s}}$ and $K_{\text{H}}^{\text{w}}$ for 37 °C.

Solvent	${10}^2\times K_{\text{H}}^{\text{s}}$	${10}^3\times K_{\text{H}}^{\text{s}}$;37 °C	$K_{\text{w}}^{\text{s}}$ 37 °C	Ref.
CCl4	1.69; 20 °C	11	1.2	[16]
CHCl3	2.15; 20 °C	17	1.9	[16]
n-Decane	5.96; 25 °C	50	5.5	[17]
1,1,2,2-Tetrachloroethane	7.34; 25 °C	61	6.7	[17]
Nitrate esters	4.33; 25 °C	38	4.2	[17]
Water	1.4; 20 °C	9.1 ( $K_{\text{H}}^{\text{w}}$)	1.0	[18]

$$K_H^s={(k_{NO2}^H\cdot M{W}^s/{\rho }^s)}^{-1}$$ (8)

Values of $K_H^s$ were extrapolated to 37 °C using Δ_sol H = −12 kJ mol⁻¹ for 1,1,2,2-tetrachloroethane [17], chloroform and carbon tetrachloride. Δ_sol H = −11 kJ mol⁻¹ was used for n-decane [17]; Δ_sol H = −9 kJ mol⁻¹ was used for nitrate esters[17] and Δ_sol H = −20 kJ mol⁻¹, which was derived from data recommended by Cheung et al. [18], was used for water. Nitrate esters refer to 1,4- and 2-3-butylene glycol dinitrate, 1,2-propylene glycol dinitrate, and glyceryl trinitrate, all of which exhibited similar solubility properties [17]. Although not specified by Lur’e et al. [17], it is assumed here that these authors used pure chemicals (and not mixtures of isomers) in their study and that they refer to the linear isomer of decane (i.e. n-decane) and to the 1,1,2,2-tetrachloroethane isomer of tetrachloroethane that is produced at an industrial scale and is unreactive towards NO₂.

Henry’s law constants for NO₂, $K_H^s$, shown in Table 1, are 1.2 to 6.7 times larger for organic solvents compared to water. The relative insensitivity of $K_w^s$ to changes in solvent suggests the NO₂-solvent interactions are of a very general type, as has been previously interpreted for NO [21]. This slight preference that NO₂ has for dissolving in organic solvents relative to water has implications, for example, on reactive uptake of gas phase NO₂ by the lung epithelial lining fluid layer (ELF) that is covered by a thin film composed primarily of phospholipids and surfactant proteins and must be taken into account when modeling the local dose of NO₂ in the respiratory tract. Thus, for example, it becomes important to compare Henry’s law constants for NO₂ in water to those in organic solvents since a higher solubility of NO₂ in the lipids near the air/lung ELF interface will tend to enhance NO₂ reactive uptake [18] by providing a subphase where NO₂ can dissolve that is vicinal to biological reactive targets (if all other experimental parameters remained unchanged).

Henry’s law constants and water-to-organic solvent partition coefficients for N₂, O₂, NO, and NO₂ are shown on Table 2 [22–24]. In order to simplify comparisons, only the same or similar solvents for which these constants are known for NO₂ are shown. It is interesting to note that NO₂ is more soluble in all solvents, including water, than the other gases. However, since the most marked difference occurs when the solvent is water, this results in a lower hydrophobicity for NO₂.

Table 2.

Henry’s law constants $K_{\text{H}}^{\text{s}}$ (M atm⁻¹) and water-to-organic solvent partition coefficients $K_{\text{w}}^{\text{s}}$ for N₂, O₂, NO, and NO₂ at 25 °C [16,17,21–24].

Solvent	N2		O2		NO		NO2
	${10}^2\times K_{\text{H}}^{\text{s}}$	$K_{\text{w}}^{\text{s}}$	${10}^2\times K_{\text{H}}^{\text{s}}$	$K_{\text{w}}^{\text{s}}$	${10}^2\times K_{\text{H}}^{\text{s}}$	$K_{\text{w}}^{\text{s}}$	${10}^2\times K_{\text{H}}^{\text{s}}$	$K_{\text{w}}^{\text{s}}$
CCl4	0.663	10.1	1.24	9.90	1.43	7.35	1.49	1.21
CHCl3	0.552	8.40	1.16	9.26	1.62a	8.3a	1.99	1.62
1,1,2,2-Tetrachloroethane	—	—	—	—	—	—	7.34	5.95
n-Decane	0.623	9.52	1.13	9.01	1.27a	6.49a	5.96	4.85
n-Hexane	1.06	16.2	1.53	12.2	1.80	9.26	—	—
Nitrate esters	—	—	—	—	—	—	4.33	3.52
1-Propyl nitrate ester	0.589	9.00b	0.832b	6.62	—	—	—	—
Nitrobenzene	—	—	—	—	0.786	4.03	—	—
Waterc	0.0656	1	0.125	1	0.195	1	1.23	1

^aCalculated using Abraham’s solvation equations.

^bAt 20 °C.

^c $K_{\text{H}}^{\text{s}}=K_{\text{H}}^{\text{w}}$.

In order to compare NO₂ to NO (since biological “lens” effects have been reported for the latter [13;14]) we begin by selecting a $K_H^w$ value for NO at 37 °C. The most recent $K_H^w$ values for NO compiled by Sander [20] as well as a recent determination by Zacharia et al. [25] average 1.9×10⁻³ M atm⁻¹ at 25 °C. Then, to correct this value to 37 °C, we employ the enthalpy of solution Δ_sol H for NOin water, taken as −13 kJ mol⁻¹, which is the average of values compiled by Sander [20] and the heat of solution derived from solubility data at various temperatures given in Perry and Green [26]. Then, using equation (6), one arrives at $K_H^w=1.6\times {10}^{-3}\text{M}{\text{atm}}^{-1}$ for NO at 37 °C.

Henry’s law constants and water-to-organic solvent partition coefficients $K_w^s$ for NO and NO₂ at 37 °Care shown in Table 3. Henry’s law constants for NO at 25 °C and enthalpies of solvation in several solvents, which were used here for temperature corrections to 37 °C, were reported by Shaw et al. [27]. In addition, the solvation equation of Abraham [21;28] was used as predictor for the partition coefficients for NO in chloroform and n-decane since experimental data for NO in these solvents are not available but are available for NO₂. Since the enthalpies of solvation are unknown for chloroform and n-decane, the values for carbon tetrachloride and n-hexane, respectively, were used as surrogate values to calculate $K_w^s$. We are now in a position to compare $K_w^s$ values for NO with NO₂ at 37 °C (Table 3). In order to facilitate comparisons, when experimental data are not available for the same solvent for NO and NO₂, experimental values for related solvents are shown in Table 3.

Table 3.

Henry’s law constants $K_{\text{H}}^{\text{s}}$ (M atm⁻¹) for NO and water-to-organic solvent partition coefficients $K_{\text{w}}^{\text{s}}$ for NO and NO₂ in various solvents at 37 °C.

Solvent	NO		NO2
	${10}^2\times K_{\text{H}}^{\text{s}}$; 37 °C	$K_{\text{w}}^{\text{s}}$	$K_{\text{w}}^{\text{s}}$
CCl4	1.38	8.3	1.2
CHCl3	1.58	10	1.9
1,1,2,2-Tetrachloroethane	—	—	6.7
Cyclohexane	1.77	11	—
n-Hexane	1.74	11	—
n-Decane	1.23	7.7	5.5
Nitrate esters	—	—	4.2
Nitrobenzene	0.76	4.8	—
Water	0.16	1.0	1.0

Discussion

One can notice that all $K_w^s$ values are higher than one, indicating that both NO and NO₂ are more soluble in organic solvents than they are in water, more so NO than is NO₂. However, even for NO, the water-to-organic solvent relative solubilities are not very pronounced. Nevertheless, there will be some thermodynamic tendency for NO and NO₂ to flee the aqueous environment where they will most likely be produced. Thus, NO is 7.7–11 times more soluble in hydrocarbons, 8.3 times more soluble in carbon tetrachloride, 10 times more soluble in chloroform, and 4.8 times more soluble in nitrobenzene than it is in water. Indeed, the “lens” effect or acceleration of the autoxidation of NO in liposomes by a factor of ~ 30 relative to aqueous solution has been explained using water-to-liposome partition coefficients of 3.6 and 3.2 for NO and O₂, respectively [13;14;29]. Thus, when one considers the form of the rate expression in equation (1), a theoretical acceleration factor of (3.6)²(3.2) = 41, which is not very different from the observed value of ~30, can be calculated. It is important to recognize that it is because the rate expression is second order in NO and first order in O₂ and that because both NO and O₂ are relatively hydrophobic it is that the overall effect results in a large acceleration factor.

In comparison to NO, the relative solubilities of NO₂ for organic solvents are less pronounced. NO₂ is 5.5 times more soluble in the hydrocarbon n-decane, 1.2 times more soluble in carbon tetrachloride, 1.9 times more in chloroform, and 6.7 times more soluble in 1,1,2,2-tetrachloroethane than it is in water. Going towards more polar solvents, NO₂ is 4.2 times more soluble in the nitrate esters, 1,4- and 2-3-butylene glycol dinitrate, 1,2-propylene glycol dinitrate, and glyceryl trinitrate than it is in water. Thus, for NO₂ in these solvents, the water-to-organic solvent partition coefficients vary from 1.2 to 6.7, suggesting that NO₂ is slightly hydrophobic but less so than is NO.

Averaging across the hydrocarbons, the $K_w^s$ for NO is 9.9 while the $K_w^s$ for NO₂ in n-decane is 5.5; across the chlorocarbons, the average $K_w^s$ for NO is 9.2 while the average $K_w^s$ for NO₂ is 3.3. The $K_w^s$ for NO in nitrobenzene (4.8) is 14% larger than the $K_w^s$ for NO₂ in nitrate esters (4.2).

Water-to-solvent partition coefficients, in particular those for 1-octanol and chloroform, are often used in predicting solubilities and derived properties for various chemical species in and across cell membranes [30]. The paucity of solubility data for NO₂ does not allow for the estimation of $K_w^{1-\mathit{oc}tan\mathit{ol}}$ or $K_w^{CH3Cl}$ using, for example, the solvation equations of Abraham as has been done for NO (5.5 and 8.3, respectively[21]; however, since, on the average, water-to-solvent partition coefficients for NO₂ are about 40% of those for NO and are bracketed between 1.2 and 6.7, and the solvent the NO₂-solvent interactions are of a very general type and do not appear to change much across solvents of various types, it is reasonable to expect the partition coefficient for water to the fatty acyl chains region of the cell membrane to be 3 ± 1.

Salting Out Effects

The biological aqueous medium ionic strength will result in a small salting out effect for the uncharged NO₂ radical. Salting out effects for uncharged species like NO₂ usually are small and obey a simple model for ionic strengths up to 5 M (equation 9) [31].

$$ln\gamma =bI$$ (9)

where γ is the activity coefficient for the uncharged species, b is a constant that is typically between 0 and 0.3 for small dissolved gases, and I is the ionic strength, usually less than 0.25 M for biological fluids [32]. Thus, 0 ≤ γ ≤ 1.19 for NO₂ and since this effect will occur only in the aqueous phase, the net effect will be an increase in $K_w^s$ for NO₂ of as much as 19 %.

Rate Laws for Reactions of NO₂ with Biological Target Molecules

Rate laws for chemical reactions of NO₂ (and its dimer N₂O₄ with which it trends to exist in equilibrium) with substrate molecules generally have the form qiven in equation (10)

$$-\frac{d[S]}{dt}=k_1[{NO}_2][S]+k_2{[{NO}_2]}^2[S]=k_1[{NO}_2][S]+k_2{K}_{11}[N_2{O}_4][S]$$ (10)

where S represents the biological target molecule, and k₁ and k₂ are observed rate constants for the first and second order terms, respectively[33]. The second order term occurs not because the reaction has a termolecular component but instead because the reaction involves N₂O₄. This relationship occurs because [NO₂]² and [N₂O₄] are linearly related through the equilibrium constant for reaction 2 (equation 11).

$$K_{10}=\frac{{[{NO}_2]}^2}{[N_2{O}_4]}$$ (11)

On The Formation of N₂O₄

Because of equation (11), the [NO₂] is one of the parameters that determine whether or not N₂O₄ will participate in these reactions. The equilibrium constant K₁₁ for the dissociation of N₂O₄ is relatively insensitive to solvent effects and values obtained in various solvents range from 0.25×10⁻⁴ to 4.6×10⁻⁴ M at 37 °C [16;34–36]. The solution equilibria favor the formation of the dimer relative to the gas phase where K₁₁ is 1.46×10⁻² M [37] so it is important to evaluate whether or not N₂O₄ may play a role in the biological reactions of NO₂. If one assumes [NO₂] does not surpass 0.1 μM (a rather high concentration for biological systems, even for the less reactive NO [9]), then, one can calculate [NO₂]/[N₂O₄]= K₁₁/[NO₂] will be between 250 and 4600. Thus, the equilibrium mixture will contain between 0.02 and 0.4% N₂O₄ on a molar basis. There is no doubt that there will be little N₂O₄ formed at equilibrium. However, it is unlikely that equilibrium will be reached, and consequently, that there will be even less N₂O₄ formed compared to equilibrium conditions. This is because the relatively high reactivity of NO₂ for biological molecules will outcompete dimerization to form N₂O₄ since biological molecules can react with NO₂ with rate constants [38] that although they can be smaller than the rate constant for the dimerization of NO₂ (k_N2O4 = 9×10⁸ M⁻¹s⁻¹ in water [39]), the biological molecules can be present in concentrations that are much larger than NO₂ can achieve in tissues and more than compensate for the difference in rate constants. The rates of disappearance of NO₂ due to its dimerization (R₁₂) and due to reaction with a biological substrate S (R₁₃) can be expressed as depicted in equations 12 and 13, respectively.

$${\text{R}}_{12}=-\frac{d[{NO}_2]}{dt}=2k_{N2O4}{[{NO}_2]}^2$$ (12)

$${\text{R}}_{13}=-\frac{d[{NO}_2]}{dt}=k_S[{NO}_2][S]$$ (13)

As an example, one can evaluate R₁₃/R₁₂ = 0.5(k_S/k_N2O4)([S]/[NO₂]) when [NO₂] = 0.1×10⁻⁶ M, a high [NO₂] that represents a favorable biological scenario for dimerization to N₂O₄. Now, let us consider that the substrate that will react with NO₂ is intracellular glutathione (GSH) ([GSH]_{intracellular} ≈ 2 mM; k_GSH = 2×10⁷ M⁻¹s⁻¹ [38]). Evaluation of R₁₃/R₁₂ for these values of rate constants and concentrations yields 200 revealing that it is very unlikely that any significant N₂O₄ will be formed even under conditions where high [NO₂] is produced that favor formation of N₂O₄. Hence, it would be safe to assume that formation of N₂O₄ in biological systems will be very inefficient, that reaction products resulting from reactions of biological molecules with N₂O₄ will be negligible, and that N₂O₄ will not function as a substantial reservoir for NO₂.

On The Formation of N₂O₃

NO₂ can react with NO to form dinitrogen trioxide (N₂O₃) in a reversible reaction as depicted in reaction 14

$$\begin{array}{l}{\text{N}}_2{\text{O}}_3\rightleftharpoons \text{NO}+{\text{NO}}_2\\ K_{14}=\frac{[NO][{NO}_2]}{[N_2{O}_3]}\end{array}$$ (14)

This equilibrium has been studied by Shaw et. al. [27] and using their data one can calculate K₁₄ = 0.34×10⁻⁴ M in carbon tetrachloride, and 2.2×10⁻⁴ in acetonitrile, at 37 °C. In water, K₁₄ = 7.3×10⁻⁵ M at 25 °C [40] and assuming Δ₁₄H = 62 kJmol⁻¹ (as in acetonitrile [24]), one can calculate K₁₄ = 1.9×10⁻⁴ M at 37 °C. As we observed with K₁₁, K₁₄ is also relatively insensitive to solvent effects. Moreover, K₁₁ ≈ K₁₄. The relative amount of NO₂ that will react with NO to form N₂O₃ at equilibrium will depend on the concentrations of both of these reactants. For [NO] = 0.1 μM, one can calculate [NO₂]/[N₂O₃]= K₁₄/[NO] will be between 340 and 2200. Thus, only 0.05 to 0.3% of NO₂ will form N₂O₃ at equilibrium when [NO] = 0.1 μM. The rate constant for the reaction of NO₂ with NO (k_N2O3 = 1.1×10⁹ M⁻¹s⁻¹ in water [40]) to form N₂O₃ is also similar to the rate constant for the dimerization of NO₂ (k_N2O4 = 9×10⁸ M⁻¹s⁻¹ [39]) to form N₂O₄. The rate laws for these processes are given in equations 15 and 16, respectively.

$$\frac{d[N_2{O}_3]}{dt}=k_{N2O3}[NO][{NO}_2]$$ (15)

$$\frac{d[N_2{O}_4]}{dt}=k_{N2O4}{[{NO}_2]}^2$$ (16)

The overall case for N₂O₃ formation is thus similar to the case for N₂O₄, however, it is likely that [NO₂] < [NO] and that R₁₅ > R₁₆, because of the higher reactivity of NO₂ for biological molecules as compared to NO. Consequently, although conditions will not favor formation of either N₂O₃ or N₂O₄, N₂O₃ will generally be more likely to form than will be N₂O₄.

On the Reactions of NO₂ with Target Molecules in Cell Membranes

As seen above, the reactions of NO₂ with biological molecules follow rate laws that are first order in NO₂ and first order in the biological molecule for environmentally and biologically relevant concentrations of NO₂. Any gain in mass action due to gains in [NO₂] in the hydrophobic region of the cell membrane will be modest because we estimate a water to the hydrophobic region of the cell membrane partition coefficient of 3 ± 1 (Figure 2). Reactions of NO₂ in cell membranes will proceed with rates that are consistent with rate constants, concentrations for biological molecules in the cell membrane, and an estimated three-fold increase in [NO₂].

Figure 2.

Conclusions

NO₂ has a slight preference to dissolve in organic solvents compared to water and this may have some implications, for example, during reactive uptake of NO₂ by the lung epithelial lining fluid. NO₂ is 1.2 times more soluble in CCl₄, 1.9 times in CHCl₃, 4.2 times more soluble in the nitrate esters, 1,4- and 2-3-butylene glycol dinitrate, 1,2-propylene glycol dinitrate, and glyceryl trinitrate, 5.5 times in n-decane, and 6.7 times more soluble in 1,1,2,2-tetrachloroethane than it is in water, and the preference for solvent increases by approximately 19% when the ionic strength of the aqueous medium is taken into account. Henry’s law constants indicate NO₂ is more soluble than NO in all solvents studied here with the most marked difference in water, resulting in a lower hydrophobicity for NO₂, as compared to NO. The available data on the solubility of NO₂ leads us to conclude that “lens” effects due to the slight hydrophobicity of NO₂ will be small, probably around a factor of approximately 3 ± 1. Thus, only a small acceleration of NO₂ reactions is expected to occur in hydrophobic cellular compartments (if all other experimental parameters remained unchanged). Nitrogen dioxide is less hydrophobic than NO, O₂ and N₂. Due to the low concentrations of NO₂ that can be achieved in biological systems and in relation to the dissociation constant of N₂O₄, it is unlikely that N₂O_4, at any time, represents any significant reservoir of NO₂, or significantly contributes to the reactivity of NO₂. Finally, although biological conditions do not favor formation of either N₂O₃ or N₂O₄, it is estimated that N₂O₃ will generally be more likely to form than will be N₂O₄.

Figure 1.