Carbon monoxide

Carbon monoxide (chemical formula CO) is a colorless, odorless, tasteless, and flammable gas that is slightly less dense than air. It is the simplest oxocarbon and is isoelectronic with the cyanide ion and molecular nitrogen. Carbon monoxide is acutely toxic to haemoglobin-based oxygen-breathing life, binding to haemoglobin with an affinity approximately 240 times that of oxygen to form carboxyhaemoglobin, which impairs oxygen transport. In spite of its toxicity, CO is produced on an enormous scale industrially as a key building block for numerous commodity chemicals and plays critical roles in modern metallurgy. It is also present in trace amounts in the atmosphere, where it participates in fundamental photochemical cycles, and has been identified in interstellar space.

History

The toxic effects of breathing the fumes from smoldering charcoal were recognized in antiquity. Aristotle (384–322 BCE) observed that coal fumes caused headaches and death, and the Greek physician Galen (129–c. 216 CE) described the ‟noxious air” produced by burning charcoal. The first deliberate chemical preparation of carbon monoxide is credited to the French chemist Joseph de Lassone in 1776, who obtained the gas by heating zinc oxide with charcoal. He mistakenly identified it as a form of hydrogen because it burned with a blue flame. Its true composition as an oxide of carbon was established by the Scottish chemist William Cruickshank around 1800. The toxicological mechanism through binding to blood was elucidated by Claude Bernard in the mid‑19th century.

Physical and chemical properties

Carbon monoxide is a linear diatomic molecule with a carbon–oxygen bond length of 112.8 pm, consistent with a net‑triple bond composed of two ordinary covalent bonds and one dative π‑bond from oxygen to carbon. The molecule possesses a small electric dipole moment of 0.122 D, with the negative end located on the carbon atom; this polarity accounts for its ability to act as a strong σ‑donor and π‑acceptor ligand in metal carbonyls. The bond dissociation energy is 1072 kJ mol⁻¹, making it one of the strongest known chemical bonds.

At standard temperature and pressure, CO is a gas with a density of 1.145 kg m⁻³ (slightly less than air). It liquefies at −191.5 °C and solidifies at −205.1 °C. The gas is sparingly soluble in water (2.3 mL per 100 mL water at 20 °C), but somewhat more soluble in organic solvents such as ethanol and benzene. CO burns with a characteristic blue flame, producing carbon dioxide and releasing heat (ΔH° = −283.0 kJ mol⁻¹). It can act as a reducing agent at elevated temperatures, reducing many metal oxides to metals, and participates in the water‑gas shift reaction: CO + H₂O ⇌ CO₂ + H₂. With sulfur it forms carbonyl sulfide (COS), and with chlorine in the presence of light or a catalyst it yields phosgene (COCl₂). The gas reacts with many transition metals to give homoleptic metal carbonyls, such as Ni(CO)₄, Fe(CO)₅, and Cr(CO)₆, which are volatile liquids or solids of great importance in organometallic chemistry and catalysis.

Production

Industrial synthesis

The bulk of carbon monoxide is produced as a component of synthesis gas (a mixture of CO and H₂). The two principal routes are steam reforming of methane and partial oxidation of hydrocarbons or coal. In steam reforming, methane reacts with steam over a nickel catalyst at 700–1100 °C and 20–35 bar to yield CO and H₂:
CH₄ + H₂O → CO + 3 H₂
In the partial oxidation of coal or heavy hydrocarbons, a limited supply of oxygen yields CO:
2 C + O₂ → 2 CO
Pure CO can be separated from syngas by cryogenic distillation, pressure swing adsorption, or through chemical complexation (e.g., with copper(I) solutions). Another important source is the reduction of iron oxide in blast furnaces, where CO is generated in situ and serves as the reducing agent.

Laboratory preparation

Small quantities of CO can be prepared by dehydrating formic acid with concentrated sulfuric acid, which abstracts water and releases CO gas:
HCOOH → CO + H₂O
Alternatively, heating a mixture of oxalic acid and concentrated sulfuric acid gives a mixture of CO and CO₂, from which pure CO can be obtained by passage through a base. The decomposition of nickel tetracarbonyl or other metal carbonyls can also liberate CO of high purity.

Occurrence

Natural sources

Carbon monoxide is continuously introduced into the atmosphere by a variety of natural processes. Volcanic outgassing, forest and bushfires, and methane oxidation in the troposphere all contribute to the global CO budget. Photochemical degradation of volatile organic compounds and the biological breakdown of heme by microorganisms also release CO. Globally, natural sources account for roughly half of the total CO emissions.

Anthropogenic sources

Human activities, predominantly the incomplete combustion of carbon‑based fuels, constitute the other major source. Motor vehicle exhaust, industrial combustion, agricultural burning, and residential heating with solid fuels are significant contributors. Tobacco smoke is a concentrated source of CO, yielding inhaled concentrations that can significantly raise carboxyhaemoglobin levels in smokers.

Indoor sources

In enclosed environments, faulty gas appliances, kerosene heaters, charcoal grills used indoors, and operating internal combustion engines in attached garages can generate dangerous CO accumulations. These conditions are responsible for a substantial fraction of accidental CO poisonings.

Biological and toxicological effects

Mechanism of toxicity

Carbon monoxide exerts its primary toxic effect by competing with oxygen for the ferrous heme site of haemoglobin. The affinity of haemoglobin for CO is approximately 240 times its affinity for O₂, so even low ambient concentrations (as little as 0.04% by volume) can be lethal over extended periods. The binding locks haemoglobin in the high‑affinity R‑state, shifting the oxygen–haemoglobin dissociation curve leftward and inhibiting the release of oxygen to tissues. Moreover, CO interferes with cellular respiration by inhibiting cytochrome c oxidase, a terminal enzyme of the mitochondrial electron transport chain, further exacerbating tissue hypoxia.

Symptoms and treatment

Early symptoms of acute CO poisoning include headache, dizziness, nausea, confusion, and a flu‑like malaise, which can easily be misdiagnosed. Progression leads to loss of consciousness, convulsions, cardiorespiratory failure, and death. The classic (but often post‑mortem) sign of cherry‑red skin is a consequence of the bright‑red colour of carboxyhaemoglobin. Chronic low‑level exposure can cause persistent neurological deficits. Treatment involves immediate removal from the contaminated environment, administration of high‑flow 100% oxygen via a non‑rebreather mask, and, in severe cases, hyperbaric oxygen therapy to accelerate the dissociation of CO from haemoglobin and cytochrome c oxidase.

Endogenous production and signalling roles

Carbon monoxide is not solely an exogenous toxin; it is produced endogenously in mammals as a by‑product of heme catabolism by heme oxygenase isozymes (HO‑1 and HO‑2). Under normal physiological conditions, endogenous CO maintains basal carboxyhaemoglobin levels of about 0.5–0.8%. CO has been established as a gaseous signalling molecule (gasotransmitter) with anti‑inflammatory, anti‑apoptotic, and vasodilatory properties. It modulates several intracellular pathways, including soluble guanylyl cyclase, and is under investigation as a potential therapeutic agent in conditions such as pulmonary hypertension and organ transplantation.

Uses

Chemical industry

Carbon monoxide is a cornerstone of large‑scale chemical synthesis. In the Fischer–Tropsch process, syngas is converted to liquid hydrocarbons over iron or cobalt catalysts. CO is the direct precursor to methanol, produced by hydrogenation (CO + 2 H₂ → CH₃OH), and to acetic acid via the Monsanto or Cativa processes (methanol carbonylation: CH₃OH + CO → CH₃COOH). Hydroformylation (oxo reaction) combines CO, H₂, and an alkene over a cobalt or rhodium catalyst to yield aldehydes, central intermediates for plasticizers and detergents. Other important products include dimethyl carbonate and phosgene, the latter used in the manufacture of polycarbonates and polyurethanes.

Metallurgy

In the Mond process, impure nickel is reacted with CO at 50–60 °C to form volatile nickel tetracarbonyl (Ni(CO)₄), which is then decomposed at 200–250 °C to deposit ultra‑pure nickel. Carbon monoxide also serves as a reducing agent in the extraction of iron from its ores within the blast furnace, and is used as a fuel gas (producer gas and water gas) in some industrial heating applications.

Food technology

CO is employed in modified atmosphere packaging of fresh meat and fish to preserve a bright‑red colour by stabilizing myoglobin in the form of carboxymyoglobin, which does not oxidize to brown metmyoglobin as readily. This practice is approved in several countries but prohibited in others due to concerns that it may mask spoilage.

Medicine and research

In preclinical and clinical research, low‑dose inhaled carbon monoxide is being explored as an anti‑inflammatory therapeutic for conditions ranging from acute lung injury to ischemia‑reperfusion injury. Organometallic CO‑releasing molecules (CORMs) have been developed to deliver controlled amounts of CO to specific tissues.

Environmental and atmospheric roles

Carbon monoxide is a trace atmospheric gas with an average tropospheric mixing ratio of about 50–150 ppb in the unpolluted Northern Hemisphere. Its major atmospheric sink is oxidation by the hydroxyl radical (OH·): CO + OH· → CO₂ + H·, a reaction that generates hydroperoxyl radicals and influences the oxidative capacity of the atmosphere. By consuming OH·, CO indirectly prolongs the atmospheric lifetime of methane (a potent greenhouse gas) and contributes to the formation of tropospheric ozone, making it a secondary greenhouse gas. CO has a relatively short atmospheric lifetime of around one to two months, leading to highly variable concentrations. It is used as a tracer of pollution transport and combustion sources. In interstellar clouds, CO is the second‑most abundant molecule after H₂ and is routinely used to map molecular gas in galaxies.

Safety and detection

Occupational exposure limits for CO are set by organizations such as OSHA (permissible exposure limit: 50 ppm, 8‑hour time‑weighted average) and NIOSH (recommended exposure limit: 35 ppm, ceiling 200 ppm). Carbon monoxide sensors based on electrochemical cells, semiconductor oxide films, or biomimetic opto‑chemical detectors are widely available in domestic alarm systems. In the event of a leak, calm ventilation and evacuation are essential. Firefighters and emergency personnel utilize portable multi‑gas monitors to assess potentially hazardous environments. Because CO is both invisible and imperceptible to human senses, it is frequently dubbed a ‟silent killer”.