carbon monoxide: Definition from Answers.com

Carbon monoxide is a gas which is best known to us as a product of incomplete combustion. As such, mankind must have been aware of its deadly effect since the discovery and use of fire, and increasingly so as the development of the industrial revolution led to more use of combustion as a source of energy. The important producers of carbon monoxide are industrial processes, heating equipment, accidental fire, cigarettes, and the internal combustion engine. Blast furnace gas contains 25% carbon monoxide, and coal gas, which was used as a fuel in Europe up until North Sea (natural) gas became plentiful, contains 16%. Carbon monoxide poisoning is the most common cause of fatal gassing and is the cause of death in about 90% of fire victims. Domestic gas supplies still lead to carbon monoxide poisoning, but now due to leakage of products of combustion from a damaged flue or poorly maintained equipment, rather than the fuel itself, since natural gas is carbon monoxide free. In the mining industry carbon monoxide contaminates the atmosphere during and after fires or explosions. The ‘afterdamp’ occurring in such situations is a mixture of carbon dioxide and carbon monoxide.

Carbon monoxide is a colourless, odourless gas which is tasteless and non-irritant. It is somewhat less dense than air and, although it is a product of imperfect combustion, it is inflammable. The gas was first identified by Joseph Priestley in the eighteenth century, but it was Claude Bernard in 1870 who discovered the affinity between carbon monoxide and haemoglobin which accounts for its deadliness: carboxyhaemoglobin is formed and oxygen transport from the lungs to the tissues disrupted. In 1895 J. S. Haldane demonstrated that the formation of carboxyhaemoglobin is an equilibrium reaction which depends upon the relative partial pressures of carbon monoxide and oxygen in inspired gas. Haldane's interest was stimulated by the problems caused by carbon monoxide in British coal mines. By breathing carbon monoxide gas which was passed through a bottle containing a mouse, he was able to determine that man was very much more resistant to the gas. Small animals such as mice and canaries, who are more vulnerable than man due to their high metabolic rate, were used in mines to give an indication of carbon monoxide contamination. Canaries responded to the gas by falling off their perches before workers noticed any ill effects, and this normally gave ample warning. Occasionally, however, in low concentrations of the order of 0.05% carbon monoxide, the bird adapted to the gas and the workers could collapse while the bird remained well.

Carbon monoxide, like oxygen, has an affinity for iron-containing molecules, but it is about 210 times more effective in binding to iron-containing haemoglobin than oxygen is. Since air contains 21% oxygen this means that only 0.1% carbon monoxide in the air will eventually lead to 50% of the haemoglobin being combined to form carboxyhaemoglobin. Once carboxyhaemoglobin is formed, and after exposure ceases, it takes 4-5 hours for its level in the blood to fall, exponentially, by 50%. The ill effect of the gas can therefore be cumulative, and a person can be poisoned by intermittent exposure during the day.

Because carboxyhaemoglobin does not carry oxygen, a level of 50% means that the oxygen carrying capacity of the blood is reduced by 50% and there is a corresponding reduction in the ability to perform maximum exercise. The body compensates for the blood's reduction in oxygen carrying capacity by increasing cardiac output, and in the early stages of carbon monoxide poisoning the heart beats faster and more strongly. Unfortunately, haemoglobin is not the only molecule affected. Muscle myoglobin also binds carbon monoxide, 60 times more effectively than it binds oxygen. This results in a reduction of heart muscle contractility and a failure of the body's compensatory mechanisms, leading to profound tissue hypoxia, which can be fatal. The presence of carboxyhaemoglobin also diminishes the oxygen held by the normal haemoglobin, which further compounds the hypoxic effect. As tissue oxygen level falls, carbon monoxide is able to bind to other iron-containing molecules: notably cytochrome P450, an important drug-metabolizing enzyme, and cytochrome A₃, an enzyme in the terminal respiratory chain which can also be poisoned by cyanide.

The scientific history of carbon monoxide is not one of uniform gloom, however. The intense affinity of carbon monoxide for haemoglobin has allowed low concentrations to be used as a marker for measurement of the speed of blood through the lungs and the surface area of the lung available for the transfer of oxygen. This latter remains as one of the standard lung function tests. In 1951 Sjöstrand discovered that Haldane's poison gas is a normal product of the body's metabolism. The enzyme haem oxygenase breaks down the haem from senescent red blood cells, and this reaction produces carbon monoxide and bile salts. The bile salts are excreted by the liver and the carbon monoxide released gives the blood a normal carboxyhaemoglobin level of 0.2-1.0%. This endogenous carbon monoxide was thought to be just a waste product, but more recent work by Verma has demonstrated that a type of haem oxygenase is located in specific areas in the brain, and suggested that the carbon monoxide produced acts as a neurotransmitter. The carbon monoxide activates the enzyme guanylyl cyclase, as does nitric oxide, regulating the intracellular levels of the second messenger cyclic GMP, which in turn regulates cellular activity. Other workers have demonstrated the haem oxygenase enzyme system in blood vessel walls and demonstrated that the carbon monoxide released causes vasodilation, as does nitric oxide. So far, endogenous carbon monoxide release has been suggested to have a role in the sense of smell, memory, cerebellar function (and hence the body's balance and co-ordination), control of blood hormone levels from the hypothalamus, and control of smooth muscle tone and vasodilatation.

The symptoms of carbon monoxide poisoning depend on the concentration breathed. The victim may pass out without warning, but often the onset of poisoning is slow. Headache, with or without nausea, is common, and this may relate to carbon monoxide's vasodilating effect. Drowsiness and lethargy then occur, along with breathlessness on exertion. At any stage there may be chest pain; this is angina due to cardiac hypoxia. At the stage of lethargy and drowsiness, cerebral function is affected and the person may not be able to think well enough to make an escape effort. Coma follows, and death. Treatment is by removal to an uncontaminated atmosphere and the administration of 100% oxygen. Hyperbaric oxygen speeds up recovery, and there is increasing evidence that it reduces long-term neurological problems.

Endogenous carbon monoxide function is undoubtedly disrupted during poisoning, but at our present state of knowledge it is difficult to say how this contributes to the toxic action of exogenous carbon monoxide. It may well be that our picture of the mechanisms of carbon monoxide poisoning will change as the function of endogenous carbon monoxide becomes clearer. Patients with carbon monoxide poisoning may have very poor balance and yet have good cerebral function. Short-term memory may also be severely disrupted. It is tempting to link these two features with the functions suggested for endogenous carbon monoxide. No doubt time will tell if there is a relationship.

— John A. S. Ross

Bibliography

World Health Organisation (1979). Environmental health criteria 13: carbon monoxide. WHO, Geneva.
Dawson, T. M. and Snyder, S. H. (1994). Gases as biological messengers: nitric oxide and carbon monoxide in the brain. (Review article.) Journal of Neuroscience, 14 (9), 5147-59

Carbon monoxide

Preferred IUPAC name Carbon monoxide
Other names Carbon monooxide Carbonous oxide Carbon(II) oxide Carbonyl
Identifiers
CAS number	630-08-0^Y
PubChem	281 ^Y
ChemSpider	275^Y
UNII	7U1EE4V452^Y
EC number	211-128-3
UN number	1016
KEGG	D09706^Y
MeSH	Carbon+monoxide
ChEBI	CHEBI:17245^Y
RTECS number	FG3500000
Beilstein Reference	3587264
Gmelin Reference	421
Jmol-3D images	Image 1
SMILES [C-]#[O+]
InChI InChI=1S/CO/c1-2^Y Key: UGFAIRIUMAVXCW-UHFFFAOYSA-N^Y InChI=1/CO/c1-2 Key: UGFAIRIUMAVXCW-UHFFFAOYAT
Properties
Molecular formula	CO
Molar mass	28.010 g/mol
Appearance	colourless gas
Odor	odorless
Density	0.789 g mL⁻¹, liquid 1.250 g L⁻¹ at 0 °C, 1 atm 1.145 g L⁻¹ at 25 °C, 1 atm
Melting point	−205.02 °C, 68 K, -337 °F
Boiling point	−191.5 °C, 82 K, -313 °F
Solubility in water	27.6 mg/1 L (25 °C)
Solubility	soluble in chloroform, acetic acid, ethyl acetate, ethanol, ammonium hydroxide, benzene
Refractive index (n_D)	1.0003364
Dipole moment	0.122 D
Thermochemistry
Std enthalpy of formation Δ_fH^o₂₉₈	−110.5 kJ·mol⁻¹
Standard molar entropy S^o₂₉₈	198 J·mol⁻¹·K⁻¹
Hazards
MSDS	External MSDS
EU Index	006-001-00-2
EU classification	Highly flammable (F+) Very toxic (T+)
R-phrases	R61 R12 R26 R48/23
S-phrases	S53 S45
NFPA 704	4 4 2
Flash point	−191 °C (82.2 K; −311.8 °F)
Autoignition temperature	609 °C (882 K; 1,128 °F)
Related compounds
Related carbon oxides	Carbon dioxide Carbon suboxide Oxocarbons
Supplementary data page
Structure and properties	n, ε_r, etc.
Thermodynamic data	Phase behaviour Solid, liquid, gas
Spectral data	UV, IR, NMR, MS
Y (verify) (what is: Y/N?) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Carbon monoxide (CO), also called carbonous oxide, is a colorless, odorless, and tasteless gas that is slightly lighter than air. It can be toxic to humans and animals when encountered in higher concentrations, although it is also produced in normal animal metabolism in low quantities, and is thought to have some normal biological functions. In the atmosphere it is however short lived and spatially variable, since it combines with oxygen to form carbon dioxide and ozone.

Carbon monoxide consists of one carbon atom and one oxygen atom, connected by a triple bond that consists of two covalent bonds as well as one dative covalent bond. It is the simplest oxocarbon. In coordination complexes the carbon monoxide ligand is called carbonyl.

Carbon monoxide is produced from the partial oxidation of carbon-containing compounds; it forms when there is not enough oxygen to produce carbon dioxide (CO₂), such as when operating a stove or an internal combustion engine in an enclosed space. In the presence of oxygen, carbon monoxide burns with a blue flame, producing carbon dioxide.^[1] Coal gas, which was widely used before the 1960s for domestic lighting, cooking, and heating, had carbon monoxide as a significant constituent. Some processes in modern technology, such as iron smelting, still produce carbon monoxide as a byproduct.^[2]

Worldwide, the largest source of carbon monoxide is natural in origin, due to photochemical reactions in the troposphere that generate about 5 x 10¹² kilograms per year.^[3] Other natural sources of CO include volcanoes, forest fires, and other forms of combustion.

In biology, carbon monoxide is naturally produced by the action of heme oxygenase 1 and 2 on the heme from hemoglobin breakdown. This process produces a certain amount of carboxyhemoglobin in normal persons, even if they do not breathe any carbon monoxide. Following the first report that carbon monoxide is a normal neurotransmitter in 1993,^[4] as well as one of three gases that naturally modulate inflammatory responses in the body (the other two being nitric oxide and hydrogen sulfide), carbon monoxide has received a great deal of clinical attention as a biological regulator. In many tissues, all three gases are known to act as anti-inflammatories, vasodilators, and promoters of neovascular growth.^[5] Clinical trials of small amounts of carbon monoxide as a drug are on-going.

History

That burning coals emanated toxic fumes was first recorded by Aristotle (384–322 BC). An ancient method of execution was to shut the criminal in a bathing room with smouldering coals. What was not known was the mechanism of death. Galen (129–199 AD) speculated that there was a change in the composition of the air which caused harm when inhaled.^[6] In 1776, the French chemist de Lassone produced CO by heating zinc oxide with coke, but mistakenly concluded that the gaseous product was hydrogen, as it burned with a blue flame. The gas was identified as a compound containing carbon and oxygen by the Scottish chemist William Cumberland Cruikshank in the year 1800. Its toxic properties on dogs were thoroughly investigated by Claude Bernard around 1846.^[7]

During World War II, a gas mixture including carbon monoxide was used to keep motor vehicles running in parts of the world where gasoline and diesel fuel were scarce. External (with few exceptions) charcoal or wood gas generators were fitted, and the mixture of atmospheric nitrogen, carbon monoxide, and smaller amounts of other gases produced by gasification was piped to a gas mixer. The gas mixture produced by this process is known as wood gas. Carbon monoxide was also, it was reported, used on a small scale during the Holocaust at some Nazi extermination camps, the most notable by gas vans in Chelmno, and in the Action T4 "euthanasia" program.^[8]

Molecular properties

Carbon monoxide has a molar mass of 28.0, which makes it slightly lighter than air, whose average molar mass is 28.8. According to the ideal gas law, CO is therefore less dense than air. Neither gas is "ideal", however, so the exact densities and other comparative values depend upon temperature and pressure.

The bond length between the carbon atom and the oxygen atom is 112.8 pm.^[9]^[10] This bond length is consistent with a triple bond, as in molecular nitrogen (N₂), 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.^[11] The boiling point (82 K) and melting point (68 K) are very similar to those of N₂ (77 K and 63 K, respectively). The bond dissociation energy of 1072 kJ/mol is stronger than that of N₂ (942 kJ/mol) and represents the strongest chemical bond known.^[12]

The ground electronic state of carbon monoxide is a singlet state^[13] since there are no unpaired electrons.

Bonding and dipole moment

Carbon and oxygen together have a total of 10 valence electrons in carbon monoxide. To satisfy the octet rule for the carbon, 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 of the bonding orbitals is occupied by two electrons from oxygen, forming a dative or dipolar bond. This causes a 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 (polar) covalent bonds, and a reverse polarization is produced by the greater electronegativity of oxygen, with a small negative charge on oxygen. In the free carbon monoxide, a net negative charge δ- remains at the carbon end and the molecule has a small dipole moment of 0.122 D.^[14]

Oxygen has more electron density, but also more positive charge. Because most electron density is located between the atoms, the molecule has a net positive charge on the oxygen end. By contrast, the isoelectronic dinitrogen molecule has no dipole moment.

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.^[15] See also the section "Coordination chemistry" below.

Resonance structures and oxidation state

Resonance structures of carbon monoxide

Different (correct) Lewis structures can be drawn for carbon monoxide. In the structure with three covalent bonds, the octet rule is satisfied, but the electropositive carbon has a negative formal charge. The structure with two covalent bonds would be consistent with the very low dipole moment of the molecule if the bonds were nonpolar. The structure with one covalent bond expresses the greater electronegativity of oxygen and the calculated net atomic charges. None of them do exactly meet the real electronic structure. Calculations with natural bond orbitals show that the structure with a triple bond is the most important Lewis structure (for the free molecule); this structure is the best approximation of the real distribution of electron density, with maximal occupation of bonding orbitals and lone pair orbitals.^[16] This is in accordance with other theoretical and experimental studies that show that, despite the greater electronegativity of oxygen, the dipole moment points from the more-negative carbon end to the more-positive oxygen end.^[17]^[18] The three bonds, however, 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.^[16] 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.

Biological and physiological properties

Toxicity

Main article: Carbon monoxide poisoning

Carbon monoxide poisoning is the most common type of fatal air poisoning in many countries.^[19] Carbon monoxide is colourless, odorless, and tasteless, but highly toxic. It combines with hemoglobin to produce carboxyhemoglobin, which is ineffective for delivering oxygen to bodily tissues. Concentrations as low as 667 ppm may cause up to 50% of the body's hemoglobin to convert to carboxyhemoglobin.^[20] A level of 50% carboxyhemoglobin may result in seizure, coma, and fatality. In the United States, the OSHA limits long-term workplace exposure levels above 50 ppm.^[21] Within short time scales, carbon monoxide absorption is cumulative, since the half-life is about 5 h in fresh air (see main article).

The most common symptoms of carbon monoxide poisoning may resemble other types of poisonings and infections, including symptoms such as headache, nausea, vomiting, dizziness, fatigue, and a feeling of weakness. Affected families often believe they are victims of food poisoning. Infants may be irritable and feed poorly. Neurological signs include confusion, disorientation, visual disturbance, syncope and seizures.^[22]

Some descriptions of carbon monoxide poisoning include retinal hemorrhages, and an abnormal cherry-red blood hue.^[23] In most clinical diagnoses these signs are seldom noticed.^[22] One difficulty with the usefulness of this the effect is that carbon monoxide merely corrects what would otherwise be an unhealthy appearance, since the chief effect of removing deoxygenated hemoglobin is to make an asphyxiated person appear more normal, or a dead person appear more life-like, similar to the effect of red colorants in embalming fluid. The "false" or unphysiologic red-coloring effect in anoxic CO-poisoned tissue is related to the meat-coloring commercial use of carbon monoxide, discussed below.

Carbon monoxide binds to other molecules such as myoglobin and mitochondrial cytochrome oxidase. Exposures to carbon monoxide may cause significant damage to the heart and central nervous system, especially to the globus pallidus,^[24] often with long-term sequelae. Carbon monoxide may have severe adverse effects on the fetus of a pregnant woman.^[25]

Normal human physiology

Carbon monoxide is produced naturally by the human body as a signaling molecule. Thus, carbon monoxide may have a physiological role in the body, such as a neurotransmitter or a blood vessel relaxant.^[26] Because of carbon monoxide's role in the body, abnormalities in its metabolism have been linked to a variety of diseases, including neurodegenerations, hypertension, heart failure, and inflammation.^[26]

Microbiology

Carbon monoxide is a nutrient for methanogenic bacteria,^[27] a building-block for acetylcoenzyme A. This is the theme for the emerging field of bioorganometallic chemistry. Extremophile micro-organisms can, thus, metabolise carbon monoxide in such locations as the thermal vents of volcanoes.^[28] In bacteria, carbon monoxide is produced via the reduction of carbon dioxide by the enzyme carbon monoxide dehydrogenase, an Fe-Ni-S-containing protein.^[29]

CooA is a carbon monoxide sensor protein.^[30] The scope of its biological role is still unknown; it may be part of a signaling pathway in bacteria and archaea. Its occurrence in mammals is not established.

Occurrence

Carbon monoxide occurs in various natural and artificial environments. Typical concentrations in parts per million are as follows:

**Composition of dry atmosphere, by volume**^[31]
Concentration	Source
ppmv: parts per million by volume (note: volume fraction is equal to mole fraction for ideal gas only, see volume (thermodynamics))
0.1 ppmv	Natural atmosphere level (MOPITT)^[32]
0.5 to 5 ppmv	Average level in homes^[33]
5 to 15 ppmv	Near-properly adjusted gas stoves in homes, modern vehicle exhaust emissions^[34]
17 ppmv	Atmosphere of Venus
100 to 200 ppmv	Exhaust from automobiles in the Mexico City central area^[35]
700 ppmv	Atmosphere of Mars
5,000 ppmv	Exhaust from a home wood fire^[36]
7,000 ppmv	Undiluted warm car exhaust without a catalytic converter^[34]

Atmospheric presence

The streak of red, orange, and yellow across South America, Africa, and the Atlantic Ocean in this animation points to high levels of carbon monoxide on September 30, 2005.

MOPITT 2000 global carbon monoxide.

Carbon monoxide is present in small amounts in the atmosphere, chiefly as a product of volcanic activity but also from natural and man-made fires (such as forest and bushfires, burning of crop residues, and sugarcane fire-cleaning). The burning of fossil fuels also contributes to carbon monoxide production. Carbon monoxide occurs dissolved in molten volcanic rock at high pressures in the Earth's mantle.^[37] Because natural sources of carbon monoxide are so variable from year to year, it is extremely difficult to accurately measure natural emissions of the gas.

Carbon monoxide has an indirect radiative forcing effect by elevating concentrations of methane and tropospheric ozone through chemical reactions with other atmospheric constituents (e.g., the hydroxyl radical, OH^.) that would otherwise destroy them.^[38] Through natural processes in the atmosphere, it is eventually oxidized to carbon dioxide. Carbon monoxide concentrations are both short-lived in the atmosphere and spatially variable.

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.

Urban pollution

Carbon monoxide is a temporary atmospheric pollutant in some urban areas, chiefly from the exhaust of internal combustion engines (including vehicles, portable and back-up generators, lawn mowers, power washers, etc.), but also from incomplete combustion of various other fuels (including wood, coal, charcoal, oil, paraffin, propane, natural gas, and trash).

Role in ground level ozone formation

Main article: Tropospheric Ozone

Carbon monoxide is part of the series of cycles of chemical reactions that form Photochemical smog. Along with aldehydes, it reacts photochemically to produce peroxy radicals. Peroxy radicals subsequently oxidize nitrogen oxide (NO) to nitrogen dioxide(NO₂).^[39] Although this creation of NO₂ is the critical step leading to low level ozone formation, it also increases this ozone in another, somewhat mutually exclusive way, by reducing the quantity of NO that is available to react with ozone.^[39]

Simplified, the net effect of the ozone cycle is:

CO + 2O₂ → CO₂ + O₃

Indoor pollution

In closed environments, the concentration of carbon monoxide can easily rise to lethal levels. On average, 170 people in the United States die every year from carbon monoxide produced by non-automotive consumer products.^[40] However, according to the Florida Department of Health, "every year more than 500 Americans die from accidental exposure to carbon monoxide and thousands more across the U.S. require emergency medical care for non-fatal carbon monoxide poisoning" ^[41] These products include malfunctioning fuel-burning appliances such as furnaces, ranges, water heaters, and gas and kerosene room heaters; engine-powered equipment such as portable generators; fireplaces; and charcoal that is burned in homes and other enclosed areas. The American Association of Poison Control Centers (AAPCC) reported 15,769 cases of carbon monoxide poisoning resulting in 39 deaths in 2007.^[42] In 2005, the CPSC reported 94 generator-related carbon monoxide poisoning deaths.^[40] Forty-seven of these deaths were known to have occurred during power outages due to severe weather, including Hurricane Katrina.^[40] Still others die from carbon monoxide produced by non-consumer products, such as cars left running in attached garages. The Centers for Disease Control and Prevention estimates that several thousand people go to hospital emergency rooms every year to be treated for carbon monoxide poisoning.^[43]

Carbon monoxide is also a minor constituent of tobacco smoke.

Blood presence

Carbon monoxide is absorbed through breathing and enters the blood stream through gas exchange in the lungs. Normal circulating levels in the blood are 0% to 3%, and are higher in smokers. Carbon monoxide levels cannot be assessed through a physical exam. Laboratory testing requires a blood sample (arterial or venous) and laboratory analysis on a CO-Oximeter. Additionally, a noninvasive carboxyhemoglobin (SpCO) test method from Pulse CO-Oximetry exists and has been validated compared to invasive methods.^[44]

Astrophysics

Outside of Earth, carbon monoxide is the second-most common molecule in the interstellar medium, after molecular hydrogen. Because of its asymmetry, the carbon monoxide 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 our information about the molecular clouds in which most stars form.^[45]

Production

Many methods have been developed for carbon monoxide's production.^[46]

Industrial production

A major industrial source of CO is producer gas, a mixture containing mostly carbon monoxide and nitrogen, formed by combustion of carbon in air at high temperature when there is an excess of carbon. In an oven, air is passed through a bed of coke. The initially produced CO₂ equilibrates with the remaining hot carbon to give CO. The reaction of O₂ with carbon to give CO is described as the Boudouard equilibrium. Above 800 °C, CO is the predominant product:

O₂ + 2 C → 2 CO (ΔH = −221 kJ/mol)

Another source is "water gas", a mixture of hydrogen and carbon monoxide produced via the endothermic reaction of steam and carbon:

H₂O + C → H₂ + CO (ΔH = +131 kJ/mol)

Other similar "synthesis gases" can be obtained from natural gas and other fuels.

Carbon monoxide is also a byproduct of the reduction of metal oxide ores with carbon, shown in a simplified form as follows:

MO + C → M + CO

Since CO is a gas, the reduction process can be driven by heating, exploiting the positive (favorable) entropy of reaction. The Ellingham diagram shows that CO formation is favored over CO₂ in high temperatures.

Laboratory preparation

Carbon monoxide is conveniently produced in the laboratory by the dehydration of formic acid, for example with sulfuric acid.^[47]^[48] Another method is heating an intimate mixture of powdered zinc metal and calcium carbonate, which releases CO and leaves behind zinc oxide and calcium oxide:

Zn + CaCO₃ → ZnO + CaO + CO

Coordination chemistry

Main article: metal carbonyl

The HOMO of CO is a σ MO.

The LUMO of CO is a π*antibonding MO.

Most metals form coordination complexes containing covalently attached carbon monoxide. Only metals in lower oxidation states will complex with carbon monoxide ligands. This is because there must be sufficient electron density to facilitate back-donation from the metal d_xz-orbital, to the π*molecular orbital from CO. The lone pair on the carbon atom in CO, also donates electron density to the d_x²−y² on the metal to form a sigma bond. In nickel carbonyl, Ni(CO)₄ forms by the direct combination of carbon monoxide and nickel metal at room temperature. For this reason, nickel in any tubing or part must not come into prolonged contact with carbon monoxide (corrosion). Nickel carbonyl decomposes readily back to Ni and CO upon contact with hot surfaces, and this method is used for the industrial purification of nickel in the Mond process.^[49]

In nickel carbonyl and other carbonyls, the electron pair on the carbon interacts with the metal; the carbon monoxide donates the electron pair to the metal. In these situations, carbon monoxide is called the carbonyl ligand. One of the most important metal carbonyls is iron pentacarbonyl, Fe(CO)₅:

Many metal-CO complexes are prepared by decarbonylation of organic solvents, not from CO. For instance, iridium trichloride and triphenylphosphine react in boiling 2-methoxyethanol or DMF) to afford IrCl(CO)(PPh₃)₂.

Organic and main group chemistry

In the presence of strong acids and water, carbon monoxide reacts with alkenes to form carboxylic acids in a process known as the Koch–Haaf reaction.^[47] In the Gattermann-Koch reaction, arenes are converted to benzaldehyde derivatives in the presence of AlCl₃ and HCl.^[48] Organolithium compounds (e.g. butyl lithium) react with carbon monoxide, but these reactions have little scientific use.

Although CO reacts with carbocations and carbanions, it is relatively nonreactive toward organic compounds without the intervention of metal catalysts.^[50]

With main group reagents, CO undergoes several noteworthy reactions. Chlorination of CO is the industrial route to the important compound phosgene. With borane CO forms an adduct, H₃BCO, which is isoelectronic with the acylium cation [H₃CCO]⁺. CO reacts with sodium to give products resulting from C-C coupling such as sodium acetylenediolate 2Na⁺·C₂O2−
2. It reacts with molten potassium to give a mixture of an organometallic compound, potassium acetylenediolate 2K⁺·C₂O2−
2, potassium benzenehexolate 6K⁺ C₆O6−
6,^[51] and potassium rhodizonate 2K⁺·C₆O2−
6.^[52]

The compounds cyclohexanehexone or triquinoyl (C₆O₆) and cyclopentanepentone or leuconic acid (C₅O₅), which so far have been obtained only in trace amounts, can be regarded as polymers of carbon monoxide.

At pressures of over 5 gigapascals, carbon monoxide disproportionates into carbon dioxide (CO₂) and a solid polymer of carbon and oxygen, in 3:2 atomic ratio.^[53]^[54]

Uses

Chemical industry

Carbon monoxide is an industrial gas that has many applications in bulk chemicals manufacturing.^[55]

Large quantities of aldehydes are produced by the hydroformylation reaction of alkenes, carbon monoxide, and H₂. Hydroformylation is coupled to the Shell Higher Olefin Process to give precursors to detergents. Methanol is produced by the hydrogenation of carbon monoxide. In a related reaction, the hydrogenation of carbon monoxide is coupled to C-C bond formation, as in the Fischer-Tropsch process where carbon monoxide is hydrogenated to liquid hydrocarbon fuels. This technology allows coal or biomass to be converted to diesel.

In the Monsanto process, carbon monoxide and methanol react in the presence of a homogeneous rhodium catalyst and hydroiodic acid to give acetic acid. This process is responsible for most of the industrial production of acetic acid.

An industrial scale use for pure carbon monoxide is purifying nickel in the Mond process.

Meat coloring

Carbon monoxide is used in modified atmosphere packaging systems in the US, mainly with fresh meat products such as beef, pork, and fish to keep them looking fresh. The carbon monoxide combines with myoglobin to form carboxymyoglobin, a bright-cherry-red pigment. Carboxymyoglobin is more stable than the oxygenated form of myoglobin, oxymyoglobin, which can become oxidized to the brown pigment metmyoglobin. This stable red color can persist much longer than in normally packaged meat.^[56] Typical levels of carbon monoxide used in the facilities that use this process are between 0.4% to 0.5%.

The technology was first given "generally recognized as safe" (GRAS) status by the U.S. Food and Drug Administration (FDA) in 2002 for use as a secondary packaging system, and does not require labeling. In 2004 the FDA approved CO as primary packaging method, declaring that CO does not mask spoilage odor.^[57] Despite this ruling, the process remains controversial for fears that it masks spoilage.^[58] In 2007 a bill^[59] was introduced to the United States House of Representatives to label modified atmosphere carbon monoxide packaging as a color additive, but the bill died in subcommittee. The process is banned in many other countries, including Canada, Japan, Singapore, and the European Union.^[60]^[61]^[62]

Medicine

Following the first report that carbon monoxide is a normal neurotransmitter in 1993,^[4] as well as one of three gases that naturally modulate inflammatory responses in the body (the other two being nitric oxide and hydrogen sulfide), carbon monoxide has received a great deal of clinical attention as a biological regulator. In many tissues, all three gases are known to act as anti-inflammatories, vasodilators, and encouragers of neovascular growth.^[5] However, the issues are complex, as neovascular growth is not always beneficial, since it plays a role in tumor growth, and also the damage from wet macular degeneration, a disease for which smoking (a major source of carbon monoxide in the blood, several times more than natural production) increases the risk from 4 to 6 times.

There is a theory that, in some nerve cell synapses, when long-term memories are being laid down, the receiving cell makes carbon monoxide, which back-transmits to the transmitting cell, telling it to transmit more readily in future. Some such nerve cells have been shown to contain guanylate cyclase, an enzyme that is activated by carbon monoxide.^[4]

Studies involving carbon monoxide have been conducted in many laboratories throughout the world for its anti-inflammatory and cytoprotective properties. These properties have potential to be used to prevent the development of a series of pathological conditions including ischemia reperfusion injury, transplant rejection, atherosclerosis, severe sepsis, severe malaria, or autoimmunity. Clinical tests involving humans have been performed, however the results have not yet been released.^[63]

References

^ Carbon Monoxide – Molecule of the Month, Dr Mike Thompson, Winchester College, UK.
^ Robert U. Ayres, Edward H. Ayres (2009). Crossing the Energy Divide: Moving from Fossil Fuel Dependence to a Clean-Energy Future. Wharton School Publishing. p. 36. ISBN 0137015445. http://books.google.com/?id=KQqLicqcTM8C&pg=PA36.
^ Weinstock, B.; Niki, H. (1972). "Carbon Monoxide Balance in Nature". Science 176 (4032): 290–2. Bibcode 1972Sci...176..290W. doi:10.1126/science.176.4032.290. PMID 5019781.
^ ^a ^b ^c New York Times article. Accessed May 2, 2010
^ ^a ^b Li, L; Hsu, A; Moore, PK (2009). "Actions and interactions of nitric oxide, carbon monoxide and hydrogen sulphide in the cardiovascular system and in inflammation--a tale of three gases!". Pharmacology & therapeutics 123 (3): 386–400. doi:10.1016/j.pharmthera.2009.05.005. PMID 19486912.
^ David G. Penney, Carbon Monoxide Toxicity, p.5, CRC Press, 2000 ISBN 0849320658.
^ Rosemary H. Waring, Glyn B. Steventon, Steve C. Mitchell (2007). Molecules of death. Imperial College Press. p. 38. ISBN 1860948146. http://books.google.com/?id=FvgNKPxb43IC&pg=PA38.
^ Martin Kitchen (2006). A history of modern Germany, 1800-2000. Wiley-Blackwell. p. 323. ISBN 1405100419. http://books.google.com/?id=_04R6DoPxLoC&pg=PT323.
^ O. R. Gilliam, C. M. Johnson and W. Gordy (1950). "Microwave Spectroscopy in the Region from Two to Three Millimeters". Physical Review 78 (2): 140. Bibcode 1950PhRv...78..140G. doi:10.1103/PhysRev.78.140.
^ Haynes, William M. (2010). Handbook of Chemistry and Physics (91 ed.). Boca Raton, Florida: CRC Press. p. 9-33. ISBN 978-1439820773.
^ Haynes, William M. (2010). Handbook of Chemistry and Physics (91 ed.). Boca Raton, Florida: CRC Press. p. 9-39. ISBN 978-1439820773.
^ Common Bond Energies (D) and Bond Lengths (r)
^ http://www.mpe.mpg.de/lab/CO/co.html Last accessed June 22, 2010.
^ Scuseria, Gustavo E.; Miller, Michael D.; Jensen, Frank; Geertsen, Jan (1991). "The dipole moment of carbon monoxide". J. Chem. Phys. 94 (10): 6660. Bibcode 1991JChPh..94.6660S. doi:10.1063/1.460293. http://jcp.aip.org/resource/1/jcpsa6/v94/i10/p6660_s1.
^ Lupinetti, Anthony J.; Fau, Stefan; Frenking, Gernot; Strauss, Steven H. (1997). "Theoretical Analysis of the Bonding between CO and Positively Charged Atoms". J. Phys. Chem. A 101 (49): 9551–9559. doi:10.1021/jp972657l.
^ ^a ^b Stefan, Thorsten; Janoschek, Rudolf (February 2000). "How relevant are S=O and P=O Double Bonds for the Description of the Acid Molecules H2SO3, H2SO4, and H3PO4, respectively?". Journal of Molecular Modeling 6 (2): 282–288. doi:10.1007/PL00010730.
^ Blanco, Fernando; Alkorta, Ibon; Solimannejad, Mohammad; Elguero, Jose (2009). "Theoretical Study of the 1:1 Complexes between Carbon Monoxide and Hypohalous Acids". J. Phys. Chem. A 113 (13): 3237–3244. doi:10.1021/jp810462h. PMID 19275137.
^ Meerts, W (1 June 1977). "Electric and magnetic properties of carbon monoxide by molecular-beam electric-resonance spectroscopy". Chemical Physics 22 (2): 319–324. Bibcode 1977CP.....22..319M. doi:10.1016/0301-0104(77)87016-X.
^ Omaye ST. (2002). "Metabolic modulation of carbon monoxide toxicity". Toxicology 180 (2): 139–150. doi:10.1016/S0300-483X(02)00387-6. PMID 12324190.
^ Tikuisis, P; Kane, DM; McLellan, TM; Buick, F; Fairburn, SM (1992). "Rate of formation of carboxyhemoglobin in exercising humans exposed to carbon monoxide.". Journal of Applied Physiology 72 (4): 1311–9. PMID 1592720.
^ "OSHA CO guidlines". OSHA. http://www.osha.gov/SLTC/healthguidelines/carbonmonoxide/recognition.html. Retrieved May 2009.
^ ^a ^b Blumenthal, Ivan (1 June 2001). "Carbon monoxide poisoning". J R Soc Med (The Royal Society of Medicine) 94 (6): 270–272. PMC 1281520. PMID 11387414. http://www.pubmedcentral.nih.gov/articlerender.fcgi?tool=pmcentrez&artid=1281520.
^ Ganong, William F (2005). "37". Review of medical physiology (22 ed.). McGraw-Hill. p. 684. ISBN 0071440402. http://books.google.com/?id=OLa8vDBXDD4C&dq=Ganong+WF.+Review+of+Medical+Physiology.+Norwalk+Ct:+Appleton+%26+Lange,+1995&printsec=frontcover. Retrieved May 2009.
^ Prockop LD, Chichkova RI (2007). "Carbon monoxide intoxication: an updated review". J Neurol Sci 262 (1–2): 122–130. doi:10.1016/j.jns.2007.06.037. PMID 17720201.
^ Susan Tucker Blackburn (2007). Maternal, fetal, & neonatal physiology: a clinical perspective. Elsevier Health Sciences. p. 325. ISBN 1416029443. http://books.google.com/?id=2y6zOSQcn14C&pg=PA325.
^ ^a ^b Wu, L; Wang, R (December 2005). "Carbon Monoxide: Endogenous Production, Physiological Functions, and Pharmacological Applications". Pharmacol Rev 57 (4): 585–630. doi:10.1124/pr.57.4.3. PMID 16382109. http://pharmrev.aspetjournals.org/cgi/content/full/57/4/585#XI._Conclusions_and_Perspectives. Retrieved May 26, 2009.
^ R. K. Thauer (1998). "Biochemistry of methanogenesis: a tribute to Marjory Stephenson. 1998 Marjory Stephenson Prize Lecture" (Free). Microbiology 144 (9): 2377–2406. doi:10.1099/00221287-144-9-2377. PMID 9782487. http://mic.sgmjournals.org/cgi/reprint/144/9/2377.
^ C.Michael Hogan. 2010. Extremophile. eds. E.Monosson and C.Cleveland. Encyclopedia of Earth. National Council for Science and the Environment, washington DC
^ Jaouen, G., Ed. (2006). Bioorganometallics: Biomolecules, Labeling, Medicine. Weinheim: Wiley-VCH. ISBN 3-527-30990-X.
^ Roberts, G. P.; Youn, H.; Kerby, R. L. (2004). "CO-Sensing Mechanisms". Microbiology and Molecular Biology Reviews 68 (3): 453–473. doi:10.1128/MMBR.68.3.453-473.2004. PMC 515253. PMID 15353565. http://www.pubmedcentral.nih.gov/articlerender.fcgi?tool=pmcentrez&artid=515253.
^ Source for figures: Carbon dioxide, NOAA Earth System Research Laboratory, (updated 2010.06). Methane, IPCC TAR table 6.1, (updated to 1998). The NASA total was 17 ppmv over 100%, and CO₂ was increased here by 15 ppmv. To normalize, N₂ should be reduced by about 25 ppmv and O₂ by about 7 ppmv.
^ Committee on Medical and Biological Effects of Environmental Pollutants (1977). Carbon Monoxide. Washington, D.C.: National Academy of Sciences. pp. 29. ISBN 0-309-02631-8.
^ Green W. "An Introduction to Indoor Air Quality: Carbon Monoxide (CO)". United States Environmental Protection Agency. http://www.epa.gov/iaq/co.html. Retrieved 2008-12-16.
^ ^a ^b Gosink, Tom (1983-01-28). "What Do Carbon Monoxide Levels Mean?". Alaska Science Forum. Geophysical Institute, University of Alaska Fairbanks. http://www.gi.alaska.edu/ScienceForum/ASF5/588.html. Retrieved 2007-12-01.
^ Singer, Siegfried Fred. The Changing Global Environment. Dordrecht: D. Reidel Publishing Company. p. 90.
^ Gosink T (January 28, 1983). "What Do Carbon Monoxide Levels Mean?". Alaska Science Forum. Geophysical Institute, University of Alaska Fairbanks. http://www.gi.alaska.edu/ScienceForum/ASF5/588.html. Retrieved December 16, 2008.
^ Astrid Sigel, Roland K. O. Sigel (2009). Metal-Carbon Bonds in Enzymes and Cofactors. Royal Society of Chemistry. p. 243. ISBN 1847559158. http://books.google.com/?id=kcllhZcy53cC&pg=PA243.
^ James Carrick White et al. (1989). Global climate change linkages: acid rain, air quality, and stratospheric ozone. Springer. p. 106. ISBN 0444015159. http://books.google.com/?id=gnaNHeU4rgQC&pg=PA106.
^ ^a ^b Ozone and other photochemical oxidants. National Academies. 1977. p. 23. ISBN 0309025311. http://books.google.com/?id=QkErAAAAYAAJ&pg=PA23.
^ ^a ^b ^c U.S Consumer Product Safety Commission, Carbon Monoxide Questions and Answers, accessed 2009-12-04
^ (http://www.doh.state.fl.us/Environment/programs/Environmental_Public_Health_Tracking/HTML_pages/Data_Carbon_Monoxide.html)
^ American Association of Poison Control Centers 2007 Annual Report
^ Centers for Disease Control and Prevention, National Environmental Public Health Tracking Network, Carbon Monoxide Poisoning, accessed 2009-12-04
^ Roth D., Herkner H., Schreiber W., Hubmann N., Gamper G., Laggner A.N., Havel C. Ann Emerg Med. 2011 Jul;58(1):74-9.
^ Combes, Françoise (1991). "Distribution of CO in the Milky Way". Annual Review of Astronomy & Astrophysics 29: 195. doi:10.1146/annurev.aa.29.090191.001211.
^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 200. ISBN 0-12-352651-5.
^ ^a ^b Koch, H.; Haaf, W. (1973), "1-Adamantanecarboxylic Acid", Org. Synth., http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=cv5p0020 ; Coll. Vol. 5: 20
^ ^a ^b G. H. Coleman, David Craig (1943), "p-Tolualdehyde", Org. Synth., http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=cv2p0583 ; Coll. Vol. 2: 583
^ Mond L, Langer K, Quincke F (1890). "Action of carbon monoxide on nickel". Journal of the Chemical Society 57: 749–753. doi:10.1039/CT8905700749.
^ Chatani, N.; Murai, S. "Carbon Monoxide" in Encyclopedia of Reagents for Organic Synthesis (Ed: L. Paquette) 2004, J. Wiley & Sons, New York. doi:10.1002/047084289
^ Büchner, W.; Weiss, E. (1964). "Zur Kenntnis der sogenannten "Alkalicarbonyle" IV[1] Über die Reaktion von geschmolzenem Kalium mit Kohlenmonoxid". Helvetica Chimica Acta 47 (6): 1415–1423. doi:10.1002/hlca.19640470604.
^ George Fownes (1869). A Manual of elementary chemistry. H.C. Lea. p. 678. http://books.google.com/?id=WM8HAAAAIAAJ&pg=PA678.
^ Katz, Allen I.; Schiferl, David; Mills, Robert L. (1984). "New phases and chemical reactions in solid carbon monoxide under pressure". The Journal of Physical Chemistry 88 (15): 3176. doi:10.1021/j150659a007.
^ Evans, W. J.; Lipp, M. J.; Yoo, C.-S.; Cynn, H.; Herberg, J. L.; Maxwell, R. S.; Nicol, M. F. (2006). "Pressure-Induced Polymerization of Carbon Monoxide: Disproportionation and Synthesis of an Energetic Lactonic Polymer". Chemistry of Materials 18 (10): 2520. doi:10.1021/cm0524446.
^ Elschenbroich, C.;Salzer, A. ”Organometallics : A Concise Introduction” (2nd Ed) Wiley-VCH: Weinheim, 2006. ISBN 3-527-28165-7
^ Sorheim, S, Nissena, H, Nesbakken, T (1999). "The storage life of beef and pork packaged in an atmosphere with low carbon monoxide and high carbon dioxide". Journal of Meat Science 52 (2): 157–164. doi:10.1016/S0309-1740(98)00163-6.
^ Eilert EJ (2005). "New packaging technologies for the 21st century". Journal of Meat Science 71 (1): 122–127. doi:10.1016/j.meatsci.2005.04.003.
^ "Low-Oxygen Packaging with CO: A Study in Food Politics That Warrants Peer Review". http://www.foodsafetymagazine.com/article.asp?id=644&sub=sub1. Retrieved April 18, 2007.
^ "Carbon Monoxide Treated Meat, Poultry, and Seafood Safe Handling, Labeling, and Consumer Protection Act (Introduced in House)". http://www.thomas.gov/cgi-bin/query/z?c110:H.R.3115:.
^ "Proof in the Pink? Meat Treated to Give It Fresh Look". ABC News. November 14, 2007. http://abcnews.go.com/GMA/Consumer/Story?id=3863064&page=1. Retrieved May 27, 2009.
^ Carbon Monoxide in Meat Packaging: Myths and Facts. American Meat Institute. 2008. http://www.meatami.com/ht/a/GetDocumentAction/i/40141. Retrieved May 2009.
^ "CO in packaged meat". Carbon Monoxide Kills Campaign. http://www.carbonmonoxidekills.com/packages_meat.htm. Retrieved May 2009.
^ Johnson, Carolyn Y. (October 16, 2009). "Poison gas may carry a medical benefit". The Boston Globe. http://www.boston.com/news/local/massachusetts/articles/2009/10/16/poison_gas_may_carry_a_medical_benefit/?page=full. Retrieved October 16, 2009.

External links

Wikimedia Commons has media related to: Carbon monoxide

Explanation of the structure
Carbon Monoxide Safety Association
International Chemical Safety Card 0023
National Pollutant Inventory – Carbon Monoxide
NIOSH Pocket Guide to Chemical Hazards
CID 281 from PubChem
External MSDS data sheet
Carbon Monoxide Detector Placement
Carbon Monoxide Kills Awareness Campaign Site
Carbon Monoxide Purification Process
Carbon Monoxide Hazards with Backpacking Stoves
USFDA IMPORT BULLETIN 16B-95, May 1999
FDA Agency Response Letter GRAS Notice No. GRN 000083
Microscale Gas Chemistry Experiments with Carbon Monoxide
Instant insight outlining the physiology of carbon monoxide from the Royal Society of Chemistry
Pictures of CO Poisoning Radiology and Pathology Images from MedPix.

v d e Oxocarbons

Common oxides	CO₂ CO

Exotic oxides	CO₃ CO₄ C₂O C₂O₂ C₂O₃ C₂O₄ C₃O₂ C₄O₂ C₄O₄ C₄O₆ C₅O₂ C₅O₅ C₆O₆ C₁₂O₆ C₁₂O₉

Compounds derived from oxides	Metal carbonyls Carbonic acid Bicarbonates Carbonates Dicarbonates Tricarbonates

Neurotransmitters

Amino acids

Alanine · Aspartate · Cycloserine · DMG · GABA · Glutamate · Glycine · Hypotaurine · Kynurenic acid (Transtorine) · NAAG (Spaglumic acid) · NMG (Sarcosine) · Serine · Taurine · TMG (Betaine)

Endocannabinoids

2-AG · 2-AGE (Noladin ether) · AEA (Anandamide) · NADA · OAE (Virodhamine) · Oleamide · PEA (Palmitoylethanolamide) · RVD-Hpα · Hp (Hemopressin)

Gasotransmitters

Carbon monoxide · Hydrogen sulfide · Nitric oxide · Nitrous oxide

Monoamines

Dopamine · Epinephrine (Adrenaline) · Melatonin · NAS (Normelatonin) · Norepinephrine (Noradrenaline) · Serotonin (5-HT)

Purines

Adenosine · ADP · AMP · ATP

Trace amines

3-ITA · 5-MeO-DMT · Bufotenin · DMT · NMT · Octopamine · Phenethylamine · Synephrine · Thyronamine · Tryptamine · Tyramine

Others

1,4-BD · Acetylcholine · GBL · GHB · Histamine

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