Carbon monoxide

Carbon monoxide is a poisonous, flammable gas that is colorless, odorless, tasteless, and slightly less dense than air. Carbon monoxide consists of one carbon atom and one oxygen atom connected by a triple bond. It is the simplest carbon oxide. In coordination complexes, the carbon monoxide ligand is called carbonyl. It is a key ingredient in many processes in industrial chemistry.
The most common source of carbon monoxide is the partial combustion of carbon-containing compounds. Numerous environmental and biological sources generate carbon monoxide. In industry, carbon monoxide is important in the production of many compounds, including drugs, fragrances, and fuels.
Indoors CO is one of the most acutely toxic contaminants affecting indoor air quality. CO may be emitted from tobacco smoke and generated from malfunctioning fuel-burning stoves and fuel-burning heating systems and from blocked flues connected to these appliances. Carbon monoxide poisoning is the most common type of fatal air poisoning in many countries.
Carbon monoxide has important biological roles across phylogenetic kingdoms. It is produced by many organisms, including humans. In mammalian physiology, carbon monoxide is a classical example of hormesis where low concentrations serve as an endogenous neurotransmitter and high concentrations are toxic, resulting in carbon monoxide poisoning. It is isoelectronic with both cyanide anion and molecular nitrogen.

Physical and chemical properties

Carbon monoxide is the simplest oxocarbon and is isoelectronic with other triply bonded diatomic species possessing 10 valence electrons, including the cyanide anion, the nitrosonium cation, boron monofluoride and molecular nitrogen. It has a molar mass of 28.0, which, according to the ideal gas law, makes it slightly less dense than air, whose average molar mass is 28.8.
The carbon and oxygen are connected by a triple bond that consists of a net two pi bonds and one sigma bond. The bond length between the carbon atom and the oxygen atom is 112.8 pm. This bond length is consistent with a triple bond, as in molecular nitrogen, which has a similar bond length and nearly the same molecular mass. Carbon–oxygen double bonds are significantly longer, 120.8 pm in formaldehyde, for example. The boiling point and melting point are very similar to those of . The bond-dissociation energy of 1072 kJ/mol is stronger than that of and represents the strongest chemical bond known.
The ground electronic state of carbon monoxide is a singlet state since there are no unpaired electrons.

Temperature	Temperature	Density	Specific heat	Dynamic viscosity	Kinematic viscosity	Thermal conductivity	Thermal diffusivity	Prandtl number
−73.15	200	1.6888	1.045	1.27	0.0752	1.7	0.0963	0.781
−53.15	220	1.5341	1.044	1.37	0.0893	1.9	0.119	0.753
−33.15	240	1.4055	1.043	1.47	0.105	2.06	0.141	0.744
−13.15	260	1.2967	1.043	1.57	0.121	2.21	0.163	0.741
6.85	280	1.2038	1.042	1.66	0.138	2.36	0.188	0.733
26.85	300	1.1233	1.043	1.75	0.156	2.5	0.213	0.73
46.85	320	1.0529	1.043	1.84	0.175	2.63	0.239	0.73
66.85	340	0.9909	1.044	1.93	0.195	2.78	0.269	0.725
86.85	360	0.9357	1.045	2.02	0.216	2.91	0.298	0.725
106.85	380	0.8864	1.047	2.1	0.237	3.05	0.329	0.729
126.85	400	0.8421	1.049	2.18	0.259	3.18	0.36	0.719
176.85	450	0.7483	1.055	2.37	0.317	3.5	0.443	0.714
226.85	500	0.67352	1.065	2.54	0.377	3.81	0.531	0.71
276.85	550	0.61226	1.076	2.71	0.443	4.11	0.624	0.71
326.85	600	0.56126	1.088	2.86	0.51	4.4	0.721	0.707
376.85	650	0.51806	1.101	3.01	0.581	4.7	0.824	0.705
426.85	700	0.48102	1.114	3.15	0.655	5	0.933	0.702
476.85	750	0.44899	1.127	3.29	0.733	5.28	1.04	0.702
526.85	800	0.42095	1.14	3.43	0.815	5.55	1.16	0.705

Bonding and dipole moment

The strength of the bond in carbon monoxide is indicated by the high frequency of its vibration, 2143 cm⁻¹. For comparison, organic carbonyls such as ketones and esters absorb at around 1700 cm⁻¹.
Carbon and oxygen together have a total of 10 electrons in the valence shell. Following the octet rule for both carbon and oxygen, the two atoms form a triple bond, with six shared electrons in three bonding molecular orbitals, rather than the usual double bond found in organic carbonyl compounds. Since four of the shared electrons come from the oxygen atom and only two from carbon, one bonding orbital is occupied by two electrons from oxygen, forming a dative or dipolar bond. This causes a C←O polarization of the molecule, with a small negative charge on carbon and a small positive charge on oxygen. The other two bonding orbitals are each occupied by one electron from carbon and one from oxygen, forming covalent bonds with a reverse C→O polarization since oxygen is more electronegative than carbon. In the free carbon monoxide molecule, a net negative charge δ⁻ remains at the carbon end and the molecule has a small dipole moment of 0.122 D.
The molecule is therefore asymmetric: oxygen is more electron dense than carbon and is also slightly positively charged compared to carbon being negative.
Carbon monoxide has a computed fractional bond order of 2.6, indicating that the "third" bond is important but constitutes somewhat less than a full bond. Thus, in valence bond terms, is the most important structure, while :C=O is non-octet, but has a neutral formal charge on each atom and represents the second most important resonance contributor. Because of the lone pair and divalence of carbon in this resonance structure, carbon monoxide is often considered to be an extraordinarily stabilized carbene. Isocyanides are compounds in which the O is replaced by an NR group and have a similar bonding scheme.
If carbon monoxide acts as a ligand, the polarity of the dipole may reverse with a net negative charge on the oxygen end, depending on the structure of the coordination complex.
See also the section "Coordination chemistry" below.

Bond polarity and oxidation state

Theoretical and experimental studies show that, despite the greater electronegativity of oxygen, the dipole moment points from the more-negative carbon end to the more-positive oxygen end. The three bonds are in fact polar covalent bonds that are strongly polarized. The calculated polarization toward the oxygen atom is 71% for the σ-bond and 77% for both π-bonds.
The oxidation state of carbon in carbon monoxide is +2 in each of these structures. It is calculated by counting all the bonding electrons as belonging to the more electronegative oxygen. Only the two non-bonding electrons on carbon are assigned to carbon. In this count, carbon then has only two valence electrons in the molecule compared to four in the free atom.

Occurrence

Carbon monoxide occurs in many environments, usually in trace levels. Photochemical degradation of plant matter, for example, generates an estimated 60 million tons/year. Typical concentrations in parts per million are as follows:

Atmospheric presence

Carbon monoxide is present in small amounts in the Earth's atmosphere. Most comes from chemical reactions with organic compounds emitted by human activities and natural origins due to photochemical reactions in the troposphere that generate about 5 × 10¹² kilograms per year. Other natural sources of CO include volcanoes, forest and bushfires, and other miscellaneous forms of combustion such as fossil fuels. Small amounts are also emitted from the ocean, and from geological activity because carbon monoxide occurs dissolved in molten volcanic rock at high pressures in the Earth's mantle. Because natural sources of carbon monoxide vary from year to year, it is difficult to accurately measure natural emissions of the gas.
Carbon monoxide has an indirect effect on radiative forcing by elevating concentrations of direct greenhouse gases, including methane and tropospheric ozone. CO can react chemically with other atmospheric constituents that would otherwise destroy methane. Through natural processes in the atmosphere, it is oxidized to carbon dioxide and ozone. Carbon monoxide is short-lived in the atmosphere, and spatially variable in concentration.
Due to its long lifetime in the mid-troposphere, carbon monoxide is also used as a tracer for pollutant plumes.

Astronomy

Beyond Earth, carbon monoxide is the second-most common diatomic molecule in the interstellar medium, after molecular hydrogen. Because of its asymmetry, this polar molecule produces far brighter spectral lines than the hydrogen molecule, making CO much easier to detect. Interstellar CO was first detected with radio telescopes in 1970. It is now the most commonly used tracer of molecular gas in general in the interstellar medium of galaxies, as molecular hydrogen can only be detected using ultraviolet light, which requires space telescopes. Carbon monoxide observations provide much of the information about the molecular clouds in which most stars form.
Beta Pictoris, the second brightest star in the constellation Pictor, shows an excess of infrared emission compared to normal stars of its type, which is caused by large quantities of dust and gas near the star.
In the atmosphere of Venus carbon monoxide occurs as a result of the photodissociation of carbon dioxide by electromagnetic radiation of wavelengths shorter than 169 nm. It has also been identified spectroscopically on the surface of Neptune's moon Triton.
Solid carbon monoxide is a component of comets. The volatile or "ice" component of Halley's Comet is about 15% CO. At room temperature and at atmospheric pressure, carbon monoxide is actually only metastable and the same is true at low temperatures where CO and are solid, but nevertheless it can exist for billions of years in comets. There is very little CO in the atmosphere of Pluto, which seems to have been formed from comets. This may be because there is liquid water inside Pluto.
Carbon monoxide can react with water to form carbon dioxide and hydrogen:
This is called the water-gas shift reaction when occurring in the gas phase, but it can also take place in an aqueous solution.
If the hydrogen partial pressure is high enough, formic acid will be formed:
These reactions can take place in a few million years even at temperatures such as found on Pluto.