NO — Nitric Oxide

Nitric oxide (NO) is a simple diatomic molecule consisting of one nitrogen and one oxygen atom. Despite its small size, NO is tremendously influential in industrial chemistry, atmospheric processes, and human physiology. In the atmosphere, NO is a key participant in smog chemistry and ozone cycling; in living systems, it functions as a short-lived signaling molecule that regulates blood vessel dilation, neurotransmission, and immune responses. Because NO possesses a single unpaired electron, it behaves as a neutral radical, giving it distinctive reactivity and paramagnetism.

Under standard conditions, nitric oxide is a colorless gas. When exposed to air, it rapidly reacts with oxygen to form brown nitrogen dioxide (NO\(_2\)), a visible change that often demonstrates its presence in the laboratory. NO is sparingly soluble in water but dissolves better in organic solvents. Its role as an intermediate in large-scale processes (e.g., the Ostwald process for nitric acid) underscores its industrial importance.

From a bonding perspective, NO is unusual. Molecular orbital (MO) theory describes NO as having a bond order between a double and triple bond (approximately 2.5), accounting for its relatively strong N–O bond and radical character. The presence of one unpaired electron explains the molecule’s paramagnetism.

Structurally, NO is linear (diatomic) and exhibits a small but nonzero dipole moment, making it slightly polar. Thermally, NO is fairly stable; however, in oxidizing environments it is readily converted to higher oxides of nitrogen like NO\(_2\) and N\(_2\)O\(_3\). Since NO is a radical, many reactions proceed via radical chain mechanisms or through nitrosyl intermediates when coordinated to transition metals.

Laboratory methods commonly generate NO by reduction of nitric acid under carefully controlled conditions. For example, copper metal reacts with dilute nitric acid to give NO (if the acid is sufficiently dilute and cold), along with copper(II) nitrate.

In other setups, nitrite salts can be reduced in acidic media to liberate NO. Because multiple nitrogen oxides may form depending on conditions, gas scrubbing and analysis are often required to isolate pure NO.

Industrial production of NO is largely in situ within the Ostwald process, where ammonia is catalytically oxidized to NO, which is then further oxidized to NO\(_2\) and absorbed to produce nitric acid.

This step is exothermic and highly optimized in modern plants, reflecting NO’s central role as an intermediate in nitrogen-oxygen chemistry at scale.

Physical properties: NO is a colorless, nearly odorless gas with low boiling and melting points, indicating weak intermolecular forces. It is only slightly soluble in water but dissolves more readily in organic solvents. Liquid NO can be condensed at cryogenic temperatures and exhibits paramagnetism similar to the gas.

Chemical reactivity: The radical nature of NO makes it reactive toward oxygen, halogens, ozone, and various radicals. A hallmark reaction is its rapid oxidation by oxygen to NO\(_2\):

NO also forms nitrosyl halides with halogens, such as nitrosyl chloride (NOCl), and reacts with ozone to form nitrogen dioxide while regenerating oxygen. In aqueous oxygenated environments, interconversions among NO, NO\(_2\), nitrite (NO\(_2^{-}\)), and nitrate (NO\(_3^{-}\)) are common, underpinning atmospheric and environmental nitrogen cycles.

In mammalian biology, NO is a critical signaling molecule produced enzymatically by nitric oxide synthases (NOS) from L-arginine. It diffuses readily across membranes and activates soluble guanylate cyclase in target cells, increasing cyclic GMP (cGMP) and triggering vasodilation.

Physiological functions include regulation of blood pressure, inhibition of platelet aggregation, neural signaling, and host defense mechanisms in macrophages. Clinically, inhaled NO is used in neonatal intensive care to manage pulmonary hypertension, improving oxygenation by selectively dilating pulmonary vasculature. Pharmacologically, drugs that enhance NO signaling (e.g., through cGMP preservation) are key in cardiovascular therapy. Because NO is short-lived, its therapeutic use is tightly monitored to avoid methemoglobinemia and oxidative stress.

Atmospherically, NO is emitted from high-temperature combustion (vehicles, power plants), lightning, and soil microbial processes. Together with NO\(_2\) (collectively NO\(_x\)), it participates in photochemical smog formation and influences tropospheric ozone levels. In the stratosphere, nitrogen oxides can catalyze ozone depletion cycles.

Safety considerations: NO is toxic at elevated concentrations. It can bind to hemoglobin (via oxidation to NO\(_2\) or formation of nitrosyl species), impairing oxygen transport. In air, its immediate oxidation to NO\(_2\) produces an irritant brown gas harmful to lungs. Handling requires adequate ventilation, gas monitoring, and appropriate materials (stainless steel, certain polymers) to avoid corrosion from NO\(_x\) and acids formed in moist air.

NO is monitored using chemiluminescence analyzers where NO reacts with ozone to form excited NO\(_2\) that emits light proportional to concentration:

Aqueous analysis often involves Griess chemistry after converting NO to nitrite, enabling colorimetric quantification. In coordination chemistry, NO acts as a ligand (nitrosyl, NO\(^+\)/NO\(^-\) resonance forms) binding to transition metals, significantly affecting electronic spectra and reactivity. Classic complexes like nitroprusside exhibit characteristic NO-derived signals used in spectroscopy and reactivity studies.