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'''Carbon monoxide''' ([[carbon|C]][[oxygen|O]]) is a colorless, odorless, and tasteless gas that is slightly lighter than air. It is 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 spatially variable, short lived, having a role in the formation of ground-level ozone.
'''Carbon monoxide''' ([[carbon|C]][[oxygen|O]]) is a really good gas to sniff and smell. It is is slightly lighter than air. It is 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 spatially variable, short lived, having a role in the formation of ground-level ozone.


Carbon monoxide consists of one [[carbon]] atom and one [[oxygen]] atom, connected by a triple bond that consists of two [[covalent bond]]s as well as one [[Dipolar bond|dative]] covalent bond. It is the simplest [[oxocarbon]], and [[isoelectronic]] with the [[cyanide]] ion and molecular [[nitrogen]]. In [[coordination complex]]es the carbon monoxide [[ligand]] is called [[carbonyl]].
Carbon monoxide consists of one [[carbon]] atom and one [[oxygen]] atom, connected by a triple bond that consists of two [[covalent bond]]s as well as one [[Dipolar bond|dative]] covalent bond. It is the simplest [[oxocarbon]], and [[isoelectronic]] with the [[cyanide]] ion and molecular [[nitrogen]]. In [[coordination complex]]es the carbon monoxide [[ligand]] is called [[carbonyl]].

Revision as of 00:47, 15 October 2012

Carbon monoxide
Wireframe model of carbon monoxide
Wireframe model of carbon monoxide
Spacefill model of carbon monoxide
Spacefill model of carbon monoxide
Names
Preferred IUPAC name
Carbon monoxide
Other names
Carbon monooxide
Carbonous oxide
Carbon(II) oxide
Carbonyl
Identifiers
3D model (JSmol)
3587264
ChEBI
ChemSpider
ECHA InfoCard 100.010.118 Edit this at Wikidata
EC Number
  • 211-128-3
421
KEGG
MeSH Carbon+monoxide
RTECS number
  • FG3500000
UNII
UN number 1016
  • InChI=1S/CO/c1-2 checkY
    Key: UGFAIRIUMAVXCW-UHFFFAOYSA-N checkY
  • InChI=1/CO/c1-2
    Key: UGFAIRIUMAVXCW-UHFFFAOYAT
  • [C-]#[O+]
Properties
CO
Molar mass 28.010 g/mol
Appearance colourless gas
Odor odorless
Density 789 kg/m3, liquid
1.250 kg/m3 at 0 °C, 1 atm
1.145 kg/m3 at 25 °C, 1 atm
Melting point −205.02 °C (−337.04 °F; 68.13 K)
Boiling point −191.5 °C (−312.7 °F; 81.6 K)
27.6 mg/1 L (25 °C)
Solubility soluble in chloroform, acetic acid, ethyl acetate, ethanol, ammonium hydroxide, benzene
1.0003364
0.122 D
Thermochemistry
198 J·mol−1·K−1
−110.5 kJ·mol−1
Hazards
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 4: Very short exposure could cause death or major residual injury. E.g. VX gasFlammability 4: Will rapidly or completely vaporize at normal atmospheric pressure and temperature, or is readily dispersed in air and will burn readily. Flash point below 23 °C (73 °F). E.g. propaneInstability 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g. white phosphorusSpecial hazards (white): no code
4
4
2
Flash point −191 °C (82.1 K; −311.8 °F)
Related compounds
Supplementary data page
Carbon monoxide (data page)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
checkY verify (what is checkY☒N ?)

Carbon monoxide (CO) is a really good gas to sniff and smell. It is is slightly lighter than air. It is 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 spatially variable, short lived, having a role in the formation of ground-level 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, and isoelectronic with the cyanide ion and molecular nitrogen. 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 ongoing.

History

Aristotle (384–322 BC) first recorded that burning coals emanated toxic fumes. 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 a few exceptions) charcoal or wood gas generators were fitted, and the mixture of atmospheric nitrogen, carbon monoxide, and small 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 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 neither exactly has the densities predicted by the ideal gas law.

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 cherry-red effect is that it corrects, or masks, 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 lifelike, 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]
ppmv: parts per million by volume (note: volume fraction is equal to mole fraction for ideal gas only, see volume (thermodynamics))
Concentration Source
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, along with aldehydes, part of the series of cycles of chemical reactions that form photochemical smog. It reacts with hydroxyl radical (•OH) to produce a radical intermediate •HOCO, which transfers rapidly its radical hydrogen to O2 to form peroxy radical (HO2•) and carbon dioxide (CO2).[39] Peroxy radical subsequently reacts with nitrogen oxide (NO) to form nitrogen dioxide (NO2) and hydroxyl radical. NO2 gives O(3P) via photolysis, thereby forming O3 following reaction with O2. Since hydroxyl radical is formed during the formation of NO2, the balance of the sequence of chemical reactions starting with carbon monoxide and leading to the formation of ozone is:

CO + 2O2 + hν → CO2 + O3

(where hν refers to the photon of light absorbed by the NO2 molecule in the sequence)

Although the 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.[40]

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.[41] 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"[42] 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.[43] In 2005, the CPSC reported 94 generator-related carbon monoxide poisoning deaths.[41] Forty-seven of these deaths were known to have occurred during power outages due to severe weather, including Hurricane Katrina.[41] 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.[44]

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.[45]

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.[46]

Production

Many methods have been developed for carbon monoxide's production.[47]

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.[48][49] 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 dxz-orbital, to the π*molecular orbital from CO. The lone pair on the carbon atom in CO, also donates electron density to the dx²−y² on the metal to form a sigma bond. This electron donation is also exhibited with the cis effect, or the labilization of CO ligands in the cis position. 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.[50]

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:

Structure of iron pentacarbonyl. Iron pentacarbonyl.

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.

Metal carbonyls in coordination chemistry are usually studied using infrared spectroscopy.

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.[48] In the Gattermann-Koch reaction, arenes are converted to benzaldehyde derivatives in the presence of AlCl3 and HCl.[49] 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.[51]

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+
·C
2O2−
2. It reacts with molten potassium to give a mixture of an organometallic compound, potassium acetylenediolate 2K+
·C
2O2−
2, potassium benzenehexolate 6K+
 C
6O6−
6,[52] and potassium rhodizonate 2K+
·C
6O2−
6.[53]

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.[54][55]

Uses

Chemical industry

Carbon monoxide is an industrial gas that has many applications in bulk chemicals manufacturing.[56]

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.[57] 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.[58] Despite this ruling, the process remains controversial for fears that it masks spoilage.[59] In 2007 a bill[60] 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.[61][62][63]

Medicine

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 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.[64]

See also

References

  1. ^ Carbon Monoxide – Molecule of the Month, Dr Mike Thompson, Winchester College, UK.
  2. ^ 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 0-13-701544-5.
  3. ^ 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.
  4. ^ a b c New York Times article. Accessed May 2, 2010
  5. ^ 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.
  6. ^ David G. Penney, Carbon Monoxide Toxicity, p.5, CRC Press, 2000 ISBN 0-8493-2065-8.
  7. ^ Rosemary H. Waring, Glyn B. Steventon, Steve C. Mitchell (2007). Molecules of death. Imperial College Press. p. 38. ISBN 1-86094-814-6.{{cite book}}: CS1 maint: multiple names: authors list (link)
  8. ^ Martin Kitchen (2006). A history of modern Germany, 1800-2000. Wiley-Blackwell. p. 323. ISBN 1-4051-0041-9.
  9. ^ 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.
  10. ^ Haynes, William M. (2010). Handbook of Chemistry and Physics (91 ed.). Boca Raton, Florida, USA: CRC Press. p. 9-33. ISBN 978-1-43982077-3.
  11. ^ Haynes, William M. (2010). Handbook of Chemistry and Physics (91 ed.). Boca Raton, Florida, USA: CRC Press. p. 9-39. ISBN 978-1-43982077-3.
  12. ^ Common Bond Energies (D) and Bond Lengths (r)
  13. ^ Vidal, C. R. (28 June 1997). "Highly Excited Triplet States of Carbon Monoxide". Archived from the original on 2006-08-28. {{cite web}}: Unknown parameter |accessedate= ignored (help)
  14. ^ 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.
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  16. ^ a b Stefan, Thorsten; Janoschek, Rudolf (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. {{cite journal}}: Unknown parameter |month= ignored (help)
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