Nitroxyl
Nitroxyl (HNO), also known as azanone, is a highly reactive nitrogen oxide species and the one-electron reduced and protonated form of nitric oxide (NO), featuring a bent triatomic structure with the formula H–N=O, where the nitrogen-oxygen bond length is approximately 1.212 Å and the H–N–O angle is 108.6° in its singlet ground state.[1] As the smallest nitroso compound, HNO is unstable in aqueous solutions, rapidly dimerizing (rate constant k ≈ 8 × 10⁶ M⁻¹ s⁻¹) to form hyponitrous acid (H₂N₂O₂), which decomposes to nitrous oxide (N₂O), and it exhibits a pKₐ of 11.4, existing in equilibrium with its deprotonated form, the nitroxyl anion (NO⁻).[2] This reactivity distinguishes HNO from NO, enabling it to act as both a nucleophile and reducing agent in biological contexts.[3] Chemically, HNO undergoes fast reactions with soft nucleophiles such as thiols (e.g., glutathione, k ≈ 3.1 × 10⁶ M⁻¹ s⁻¹), forming sulfinamides or disulfides, and with oxygen (k ≈ 1.8 × 10⁴ M⁻¹ s⁻¹) to produce peroxynitrite (ONOO⁻), while also binding to metalloproteins like heme-containing enzymes (e.g., myoglobin, k ≈ 2.75 × 10⁵ M⁻¹ s⁻¹).[1] These properties have been studied through techniques like pulse radiolysis and photolysis, highlighting HNO's short half-life (≈0.6 ms at 100 µM concentration) and its generation from donors such as Angeli's salt (Na₂N₂O₃).[1] In biological systems, HNO modulates cardiovascular function by enhancing cardiac contractility and inducing vasodilation without developing tolerance, unlike NO-based therapies, and it shows potential cardioprotective effects against ischemia-reperfusion injury through thiol modifications and interactions with signaling pathways.[3][2] Ongoing research emphasizes HNO's therapeutic promise for conditions like heart failure, with advances in selective detection methods (e.g., fluorescent probes) and donor compounds to overcome its instability and enable precise delivery.[2] Despite its enigmatic nature and historical underappreciation compared to NO, HNO's unique redox chemistry and biological activity position it as a key player in nitrogen oxide signaling, warranting further exploration in pharmacology and physiology.[3]Properties
Physical Properties
Nitroxyl (HNO) is a simple triatomic molecule with the formula HNO and C_s symmetry, featuring a bent geometry characteristic of its singlet ground state. The H-N-O bond angle measures approximately 108°, reflecting the V-shaped structure similar to water but with distinct bonding due to the nitrogen-oxygen double bond. Experimental and computational studies have determined key bond lengths as N-O ≈ 1.21 Å and N-H ≈ 1.06 Å, consistent with a partial double bond character in the N-O linkage and a standard N-H single bond. These structural parameters contribute to HNO's reactivity, distinguishing it from linear nitric oxide (NO). HNO has a significant dipole moment of approximately 2.1 D, contributing to its polarity and interactions.[4][5][1] Spectroscopic signatures provide essential tools for identifying and characterizing nitroxyl. In the infrared spectrum, the N=O stretching vibration appears at approximately 1565 cm⁻¹, a frequency lower than that of NO (around 1876 cm⁻¹) due to the weakened double bond from protonation and reduction. This band is prominent in gas-phase measurements and shifts predictably with isotopic substitution, aiding confirmation in matrix isolation experiments. Ultraviolet-visible spectroscopy reveals an absorption maximum near 350 nm, associated with π→π* transitions in the N=O moiety, though the exact position can vary slightly in solution due to solvent interactions. These spectral features are crucial for detecting transient HNO in chemical and biological systems without interference from NO.[6][7][8] The thermodynamic profile of nitroxyl underscores its instability and fleeting existence. Its high reactivity manifests in a short half-life of about 1 ms in aqueous solution at pH 7, primarily driven by rapid dimerization to hyponitrous acid (second-order rate constant ≈ 8 × 10⁶ M⁻¹ s⁻¹), though this lifetime extends at ultralow concentrations. Solubility in water is limited to around 100 μM, limited by self-reaction before saturation, while it is a polar gas under standard conditions akin to other small reactive nitrogen species. Compared to nitric oxide, HNO is the one-electron reduced and protonated form, exhibiting higher proton affinity (pK_a ≈ 11.4 for HNO/NO⁻) that favors the neutral form at physiological pH and enhances its electrophilic character.[1][9][10]Chemical Properties
Nitroxyl (HNO) features a singlet ground state in which all electrons are paired in molecular orbitals or exist as lone pairs, including a non-bonding lone pair on the nitrogen atom that influences its reactivity profile. This electronic configuration distinguishes HNO from nitric oxide (NO), which possesses a triplet ground state due to an unpaired electron in a π* orbital.[11] HNO exhibits weak acidity, with a pK_a value of approximately 11.4 for the equilibrium HNO ⇌ NO⁻ + H⁺, ensuring that the neutral HNO species predominates under physiological conditions (pH ~7.4).[9] The redox chemistry of HNO is characterized by defined potentials that govern its interconversions. The standard reduction potential for the one-electron couple NO + H⁺ + e⁻ → HNO is approximately -0.14 V versus the normal hydrogen electrode (NHE), facilitating the formation of HNO from NO under reducing conditions. Further two-electron reduction of HNO to hydroxylamine (NH₂OH) occurs with E° ≈ +0.7 V for HNO + 2H⁺ + 2e⁻ → NH₂OH. In contrast, the one-electron reduction of NO to its anion NO⁻ has a more negative potential of about -0.8 V versus NHE, highlighting HNO's position as an intermediate in the nitrogen oxide redox ladder.[9] The lone pair on the nitrogen atom enables HNO to behave as a nucleophile, particularly in interactions with electrophilic centers such as transition metals, where it donates electrons to form coordination complexes. This nucleophilic character stands in opposition to the electrophilicity of NO, which readily accepts electrons due to its radical nature and empty π* orbital.[12] HNO's stability is limited in aqueous environments, primarily due to rapid dimerization to hyponitrous acid (H₂N₂O₂), which proceeds via a second-order process with a rate constant of approximately 8 × 10⁶ M⁻¹ s⁻¹ at neutral pH. The dimer is transient and often dehydrates to nitrous oxide (N₂O) and water.[13][14]Synthesis and Generation
Chemical Synthesis
HNO was first detected spectroscopically in 1958 by F. W. Dalby through flash photolysis of mixtures including nitric oxide and ammonia, providing initial insights into its transient nature.[15][1] One of the most established laboratory routes for generating HNO involves the acid-catalyzed decomposition of Angeli's salt (Na₂N₂O₃), a diazeniumdiolate that undergoes protonation at the nitroso group, followed by tautomerization and N–N bond cleavage to yield HNO and nitrite (NO₂⁻). This process occurs efficiently at pH 4–8, with first-order kinetics (rate constant ~6.8 × 10⁻⁴ s⁻¹ at 25 °C), though below pH 4, the product shifts to nitric oxide (NO) due to further oxidation. The reaction stoichiometry is Na₂N₂O₃ + H⁺ → HNO + NaNO₂ + Na⁺, but effective HNO yields are approximately 50% owing to rapid dimerization of HNO to hyponitrous acid (k = 8 × 10⁶ M⁻¹ s⁻¹), which decomposes to N₂O and H₂O.[16][17] HNO can also be produced via two-electron oxidation of hydroxylamine (NH₂OH), a direct and clean method that avoids certain by-products associated with donor decomposition. Mild chemical oxidants, such as chloramine-T (N-chlorobenzenesulfonamide) or periodate (IO₄⁻), facilitate this transformation: NH₂OH + oxidant → HNO + reduced oxidant products. For example, periodate oxidation in acidic media proceeds through intermediate species like NH₂O• radicals, leading to HNO as a key product before further oxidation to nitrite or nitrate depending on conditions. These reactions are typically conducted in aqueous solution at neutral to mildly acidic pH, with yields ranging from 50–80% under optimized conditions, though excess oxidant can promote over-oxidation. Recent studies have shown catalytic generation of HNO from NH₂OH oxidation by hydrogen peroxide using heme proteins like myoglobin (as of 2025).[17][18][19] Photolysis represents another controlled synthetic approach, particularly through UV irradiation of hydroxylamine derivatives or certain nitroso compounds. For instance, UV photolysis (λ ~254 nm) of hydroxylamine (NH₂OH) in aqueous solution generates HNO alongside hydrogen radicals (NH₂OH → HNO + H•), though competing pathways produce ammonia, nitrogen, and water. More selective methods involve photo-uncaging of N-alkoxysulfonamides or Piloty's acid (N-hydroxybenzenesulfonamide) under UV or visible light, yielding HNO with efficiencies up to 70% in deoxygenated media.[20][14] Recent advances include solid-gas reactions for generating HNO in the gas phase by contacting solid base-catalyzed HNO donors with gaseous bases (as of 2022). Despite these advances, achieving high-purity HNO remains challenging due to its inherent instability and propensity for dimerization, which limits overall yields to typically less than 70%. Contamination by NO is a common issue, particularly from pH-dependent side reactions in decompositions like that of Angeli's salt, necessitating inert atmospheres, low temperatures, and rapid trapping to minimize unwanted oxidation products.[21][16][14]Biological Generation
Nitroxyl (HNO) was first proposed as an endogenous signaling molecule in the 1990s through cardiovascular studies demonstrating its role in vasorelaxation distinct from nitric oxide (NO), particularly via effects of HNO donors like cyanamide in rabbit aortic tissue.[22] Early experiments linked HNO release from such precursors to thiol-sensitive relaxation mechanisms, differentiating it from NO pathways.[22] Enzymatic production of HNO occurs primarily through nitric oxide synthase (NOS) isoforms, which can generate HNO as a byproduct during the oxidation of the intermediate N-hydroxy-L-arginine (NOHA) to NO, especially under conditions of cofactor depletion such as low tetrahydrobiopterin levels.[23] This pathway has been observed in neuronal NOS (nNOS) and inducible NOS (iNOS), where uncoupling leads to partial reduction products including HNO rather than full NO formation.[24] Non-enzymatic routes contribute to HNO generation, notably the reaction of NO with thiolates (RS⁻), such as glutathione, under hypoxic conditions, yielding HNO and S-nitrosothiols (R-S-NO): \text{NO} + \text{RS}^- \rightarrow \text{HNO} + \text{R-S-NO} This mechanism is supported by kinetic studies showing efficient HNO formation in anaerobic environments mimicking tissue hypoxia. Mitochondrial HNO production arises from partial reduction of NO during electron transport chain leaks, particularly involving complex I or III, leading to HNO formation in cardiac mitochondria and contributing to redox signaling.[25] HNO generation is regulated by physiological factors, including pH-dependent release from prodrugs mimicking Angeli's salt, which accelerates HNO liberation at neutral pH in cellular environments. Estimated steady-state concentrations of HNO in tissues are in the low nanomolar range (~10^{-9} M), limited by rapid dimerization and scavenging by thiols. Recent reviews highlight ongoing exploration of additional enzymatic and non-enzymatic pathways in mammals and plants (as of 2024).[26][27]Reactions
Reactions with Inorganic Species
Nitroxyl (HNO) readily coordinates to metal centers, particularly heme iron in proteins such as myoglobin. In deoxymyoglobin, featuring ferrous iron (Fe(II)), HNO binds rapidly to form a stable adduct described as {Fe(II)–NO}⁸ in Enemark–Feltham notation, where the nitrogen of HNO coordinates to the iron, and the complex exhibits dual hydrogen bonding involving the distal histidine and a water molecule for enhanced stability. The second-order rate constant for this binding is approximately 10⁷ M⁻¹ s⁻¹, indicating efficient trapping of HNO under physiological conditions.[28] This interaction contrasts with slower binding to ferric heme (Fe(III)), where rates are around 10⁵–10⁶ M⁻¹ s⁻¹, highlighting the preference for reduced iron states.[29] A prominent reaction of HNO involves its dimerization, which proceeds via second-order kinetics to form hyponitrous acid (H₂N₂O₂), subsequently decomposing to nitrous oxide (N₂O) and water:$2 \mathrm{HNO} \rightarrow \mathrm{H_2N_2O_2} \rightarrow \mathrm{N_2O + H_2O}
The rate constant for dimerization is approximately 8 × 10⁶ M⁻¹ s⁻¹ at neutral pH, making it a dominant decay pathway in aqueous solutions without trapping agents.[14] This process is pH-dependent; at acidic pH, protonated HNO favors rapid dimerization, while at higher pH (>7), deprotonation to NO⁻ slows the reaction due to electrostatic repulsion, though the overall kinetics remain influenced by acid-base equilibria of the dimer.[1] The cis isomer of hyponitrous acid predominates and decomposes more readily than the trans form. HNO also undergoes oxidation by molecular oxygen (O₂) in aerated solutions to form peroxynitrite (ONOO⁻):
\mathrm{HNO + O_2 \rightarrow HOONO}
(leading to ONOO⁻ + H⁺). This reaction exhibits a second-order rate constant of 1.8 × 10⁴ M⁻¹ s⁻¹ at physiological pH (7.4), underscoring its relevance in oxygenated aqueous environments.[30] This pathway predominates in neutral solutions, with peroxynitrite as the major product, though under certain conditions HNO may contribute to NO production via secondary reactions. In reductive environments, such as anaerobic conditions, HNO can be further reduced to hydroxylamine (NH₂OH):
\mathrm{HNO + 2e^- + 2H^+ \rightarrow NH_2OH}
This two-electron, two-proton process is thermodynamically favorable and occurs via intermediates like the aminoxyl radical, relevant in biological systems lacking oxygen where HNO serves as an intermediate in nitrogen metabolism.[31] The reaction highlights HNO's role in redox cascades, though specific rate constants vary with the reducing agent and pH.[1]