Carbon monoxide

Carbon monoxide (CO) is a colourless, odourless, tasteless and highly toxic gas composed of one carbon atom and one oxygen atom linked by a triple bond. As the simplest oxocarbon, it plays significant roles in atmospheric chemistry, industrial processes and biological systems, despite its well-known toxicity. Its physical properties, reactivity and environmental behaviour make it an important subject in both inorganic and organic chemistry.

Physical and Chemical Characteristics

Carbon monoxide is slightly less dense than air, with a molar mass of 28.0 g mol⁻¹ compared to air’s average of 28.8 g mol⁻¹. Owing to its triple bond, the CO molecule is exceptionally stable, exhibiting a short bond length of 112.8 pm and a high bond dissociation energy of around 1072 kJ mol⁻¹—one of the strongest known chemical bonds. Its melting point (−205 °C) and boiling point (−191 °C) are similar to those of nitrogen, reflecting the comparable mass and bonding structure of the two gases.
Carbon monoxide is isoelectronic with several other diatomic species, including nitrogen (N₂), the cyanide anion (CN⁻) and the nitrosonium cation (NO⁺). Like these species, it contains ten valence electrons and forms a triple bond made up of two π bonds and one σ bond. The molecule exhibits a small dipole moment, with electron density shifted slightly towards the carbon atom despite oxygen’s higher electronegativity. This unusual polarity arises from the presence of a dative component in one of the bonding orbitals, with two electrons donated by oxygen.
Spectroscopically, CO shows a characteristic stretching frequency of approximately 2143 cm⁻¹, considerably higher than typical carbonyl compounds in organic chemistry. Its electronic ground state is a singlet, with all electrons paired, contributing to its stability in both terrestrial and extraterrestrial environments.

Bonding, Polarity and Oxidation State

The bond in carbon monoxide is often represented by multiple resonance forms, the most significant being a triple-bonded structure with a negative charge localised on carbon. Alternative non-octet resonance structures contribute to its reactivity, particularly in coordination chemistry. The oxidation state of carbon in CO is conventionally assigned as +2, assuming all bonding electrons are attributed to oxygen.
CO’s polarity influences its ability to act as a donor and acceptor ligand in metal complexes. Depending on the metal centre, the effective dipole may invert, with carbon or oxygen bearing a partial negative charge in different contexts.

Occurrence and Environmental Behaviour

Carbon monoxide occurs naturally in trace quantities in Earth’s atmosphere, with typical background levels of around 80 parts per billion. It is generated by:

• partial combustion of organic matter,
• photochemical oxidation of hydrocarbons,
• volcanic and geothermal activity,
• wildfires and biomass burning,
• limited release from oceans and geological substrates.

In the atmosphere, CO has a short lifetime of about one to two months. It reacts with hydroxyl radicals (OH), reducing the availability of these radicals for methane oxidation and thereby influencing atmospheric composition indirectly. Its presence contributes to the formation of tropospheric ozone.
Due to its relatively long mid-tropospheric lifetime, CO is used as a tracer gas in atmospheric and climate studies. It enables tracking of pollutant transport and large-scale atmospheric circulation patterns.

Extraterrestrial Occurrence

Beyond Earth, CO is exceptionally abundant in the interstellar medium—second only to hydrogen—as it readily forms under low-temperature astrophysical conditions. Its asymmetric structure makes it bright in radio and millimetre wavelengths, allowing astronomers to map molecular clouds and star-forming regions with high sensitivity.
Carbon monoxide has been detected:

• in interstellar molecular clouds,
• in the atmosphere of Venus (from photodissociation of CO₂),
• on Triton, a moon of Neptune,
• in cometary nuclei such as Halley’s Comet, in which CO constitutes a major volatile component.

Although metastable at room temperature, CO ice can persist for billions of years in the cold environments of cometary bodies.

Sources and Industrial Relevance

Most anthropogenic carbon monoxide comes from incomplete combustion of carbon-based fuels. Key sources include:

• vehicle exhausts,
• industrial furnaces,
• heating systems and stoves,
• biomass burning and tobacco smoke.

CO is a major industrial reagent used in processes such as:

• synthesis of methanol,
• hydroformylation of alkenes,
• Fischer–Tropsch synthesis of hydrocarbons,
• production of various pharmaceuticals and fine chemicals,
• formation of metal carbonyl complexes.

These applications depend on CO’s ability to bind strongly to transition metals and to participate in controlled reduction and carbonylation reactions.

Biological Significance and Toxicology

Although toxic at high concentrations, carbon monoxide is produced endogenously in mammals through the action of haem oxygenase, which breaks down haem into biliverdin, iron and CO. At low levels it acts as a gasotransmitter, participating in cellular signalling pathways. This dual nature makes CO a clear example of hormesis, where small doses have biological regulatory functions while higher concentrations are lethal.
Inhalation of elevated CO levels leads to carbon monoxide poisoning, a major cause of accidental death due to its strong affinity for haemoglobin. Binding to haemoglobin forms carboxyhaemoglobin, impairing oxygen transport and causing tissue hypoxia.
Common indoor sources of hazardous CO exposure include malfunctioning heaters, gas appliances, charcoal grills and inadequately ventilated combustion systems.