Carbon monoxide


Carbon monoxide
Structure of the carbon monoxide molecule
Space-filling model of the carbon monoxide molecule
IUPAC name
Other names Carbonic oxide
Identifiers
CAS number [630-08-0]
PubChem 281
EC number 211-128-3
UN number 1016
ChEBI 17245
RTECS number FG3500000
ChemSpider ID 275
Properties
Molecular formula CO
Molar mass 28.010 g/mol
Appearance colourless, odorless gas
Density 0.789 g/mL, liquid
1.250 g/L at 0 °C, 1 atm
1.145 g/L at 25 °C, 1 atm
Melting point

-205 °C (68 K)
Boiling point

-191.5 °C (81 K)
Solubility in water 0.0026 g/100 mL (20 °C)
Solubility soluble in chloroform, acetic acid, ethyl acetate, ethanol, ammonium hydroxide
Dipole moment 0.112 D
Hazards
MSDS External MSDS
EU Index 006-001-00-2
EU classification Highly flammable (F+)
Repr. Cat. 1
Toxic (T)
R-phrases R61, R12, R23, R48/23
S-phrases S53, S45
NFPA 704
2
4
0
 
Flash point -191 °C
Autoignition
temperature		 609 °C
Related compounds
Related carbon oxides Carbon dioxide
Carbon suboxide
Oxocarbons
Supplementary data page
Structure and
properties		 n, εr, etc.
Thermodynamic
data		 Phase behaviour
Solid, liquid, gas
Spectral data UV, IR, NMR, MS
Except where noted otherwise, data are given for
materials in their standard state
(at 25 °C, 100 kPa)
Infobox references

Carbon monoxide, with the chemical formula CO, is a colorless, odorless and tasteless, yet highly toxic gas. Its molecules consist of one carbon atom and one oxygen atom, connected by a covalent double bond and a dative covalent bond. It is the simplest oxocarbon, and can be viewed as the anhydride of formic acid (CH2O2).

Carbon monoxide is produced from the partial oxidation of carbon-containing compounds; it forms in preference to the more usual carbon dioxide (CO2) when there is a reduced availability of oxygen, such as when operating a stove or an internal combustion engine in an enclosed space. Carbon monoxide has significant fuel value, burning in air with a characteristic blue flame, producing carbon dioxide. Despite its serious toxicity, it was once widely used (as the main component of coal gas) for domestic lighting, cooking and heating, and in the production of nickel. Carbon monoxide still plays a major role in modern technology, in industrial processes such as iron smelting and as a precursor to myriad products.

Contents

History

Carbon monoxide has been unknowingly used by humans since prehistoric times, for the smelting of iron and other metallic ores.[citation needed] The gas was used for executions by the Greek and Romans in Classical Antiquity,[1] and was described by the Spanish doctor Arnaldus de Villa Nova in the 11th century[citation needed]. 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.[citation needed] The gas was identified as a compound containing carbon and oxygen by the English chemist William Cumberland Cruikshank in the year 1800. Its toxic properties on dogs were thoroughly investigated by Claude Bernard around 1846.

During World War II, carbon monoxide was used to keep motor vehicles running in parts of the world where gasoline was scarce. External charcoal or wood burners were fitted, and the carbon monoxide produced by gasification was piped to the carburetor.[citation needed] The CO in this case is known as "wood gas". Carbon monoxide was also reportedly used on a small scale during the Holocaust at some Nazi extermination camps (most notably by gas vans in Chelmno), and in the Action T4 "euthanasia" program.[citation needed]

Molecular properties

The carbon monoxide molecule consists of one atom of carbon and one atom of oxygen, covalently bonded by a double bond and a dative covalent bond. Its bond length is 0.1128 nm.[2] The effects of atomic formal charge and electronegativity result in a small bond dipole moment with its negative end on the carbon atom[3]. The reason for this, despite oxygen's greater electronegativity, is that the highest occupied molecular orbital has an energy much closer to that of carbon's p orbitals, meaning that greater electron density is found near the carbon. In addition, carbon's lower electronegativity creates a much more diffuse electron cloud, enhancing the polarizability. This is also the reason that almost all chemistry involving carbon monoxide occurs through the carbon atom, and not the oxygen.

The bond length of CO is consistent with a partial triple bond, and the molecule can be represented by three resonance structures:

In this classical model, the leftmost structure contributes the most. As such, carbon monoxide resembles molecular nitrogen, and in addition, it has nearly the same molecular mass. Indeed, their physical properties (boiling point, melting point, etc.) are very similar.

Biological and physiological properties

Toxicity

Carbon monoxide poisoning is the most common type of fatal poisoning in many countries.[4] Carbon monoxide is colorless and odorless, but extremely toxic: it combines with hemoglobin in the blood to produce carboxyhemoglobin (HbCO), which is ineffective for delivering oxygen to the body tissues (a condition known as anoxemia). Concentrations as low as 667 ppm can cause up to 50% of the body's hemoglobin to convert to HbCO. In the United States, OSHA limits long-term workplace exposure levels to 50 ppm.[5]

The most common symptoms of CO poisoning can resemble the flu, including headache, nausea and vomiting, dizziness, lethargy and a feeling of weakness. Infants may be irritable and feed poorly. Neurological signs include confusion, disorientation, visual disturbance, syncope and seizures.[1].

In his pioneering 1846 study, Claude Bernard observed that that the blood of poisoned dogs was more rutilant ("gleaming"[6] or "glowing"[7]) in all the vessels, a fact now known to be due to the formation of HbCO. Some classic descriptions of CO poisoning cite also retinal hemorrhages, bright reddish skin, and an abnormal "cherry-red" blood hue;[8] but in most clinical diagnoses these signs are seldom seen.[1]

Carbon monoxide is believed to compromise other important molecules such as myoglobin, and mitochondrial cytochrome oxidase. Exposures can lead to significant damage to the heart and central nervous system, especially to the globus pallidus,[9] often with long-term sequelae. Carbon monoxide can also have severe effects on the fetus of a pregnant woman.

Human Physiology

Carbon monoxide is produced naturally in the human body as part of normal metabolism, such as the breakdown of heme (a part of the hemoglobin molecule) by the enzyme heme oxygenase to CO, biliverdin and a Fe3+ cation. The endogenously produced CO may have important physiological roles in the body, such as a neurotransmitter or a blood vessels relaxant.[10] In the neuronal system it has been shown to be involved in learning and memory and odor response, among others.[citation needed] It provides cardiac protection in the circulatory system. It also has roles in the immune, respiratory, reproductive, and gastrointestinal systems, as well as in the kidneys and liver. Because of its expansive role, abnormalities in CO metabolism have been linked to a variety of disease processes, including neurodegenerations, hypertension, heart failure, and inflammation.[10] In addition CO regulates inflammatory reactions in a manner that prevents the development of several diseases such as atherosclerosis or severe malaria.[citation needed]

Microbiology

CO is a nutrient for methanogenic bacteria,[11] a building block for acetylcoenzyme A. This theme is the subject for the emerging field of bioorganometallic chemistry. In bacteria, CO is produced via the reduction of carbon dioxide via the enzyme carbon monoxide dehydrogenase, an Fe-Ni-S-containing protein.[12]

A heme-based CO-sensor protein, CooA, is known.[13] The scope of its biological role is still unclear, it is apparently part of a signaling pathway in bacteria and archaea, but its occurrence in mammals is not established.

Occurrence

Carbon monoxide commonly occurs in various natural and artificial environments. Here are some typical concentrations:

  • 0.1 ppm - natural background atmosphere level (MOPITT)
  • 0.5 to 5 ppm - average background level in homes[14]
  • 5 to 15 ppm - levels near properly adjusted gas stoves in homes[14]
  • 100-200 ppm - Mexico City central area from autos etc.[15]
  • 5,000 ppm - chimney of a home wood fire [16]
  • 7,000 ppm - undiluted warm car exhaust - without catalytic converter[16]

Atmospheric presence

MOPITT 2000 global carbon monoxide

Carbon monoxide has always been present as a minor constituent of 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) and the burning of fossil fuels. It occurs dissolved in molten volcanic rock at high pressures in the earth's mantle. Carbon monoxide contents of volcanic gases vary from less than 0.01% to as much as 2% depending on the volcano.[citation needed] 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. 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.

Urban pollution

Carbon monoxide is a major atmospheric pollutant in urban areas, chiefly from exhaust of internal combustion engines (including vehicles, portable and back-up generators, lawn mowers, power washers, etc.), but also from improper burning of various other fuels (including wood, coal, charcoal, oil, kerosene, propane, natural gas, and trash). Along with aldehydes, it reacts photochemically to produce peroxy radicals. Peroxy radicals react with nitrogen oxide to increase the ratio of NO2 to NO, which reduces the quantity of NO that is available to react with ozone.

Indoor pollution

In closed environments, the concentration of carbon monoxide can easily rise to lethal levels. On average, about 170 people in the United States die every year from CO produced by non-automotive consumer products. These products include malfunctioning fuel-burning appliances such as furnaces, ranges, water heaters and room heaters; engine-powered equipment such as portable generators; fireplaces; and charcoal that is burned in homes and other enclosed areas. In 2005 alone, CPSC staff is aware of at least 94 generator-related CO poisoning deaths. Forty-seven of these deaths were known to have occurred during power outages due to severe weather, including Hurricane Katrina. Still others die from CO 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 CO poisoning.

Carbon monoxide is also a constituent of tobacco smoke.

Production

Carbon monoxide is so fundamentally important that many methods have been developed for its production.[17]

Laboratory preparation

Carbon monoxide is conveniently produced in the laboratory by the dehydration of formic acid, for example with sulfuric acid. 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

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 important 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 is 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.

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 those 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. 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 was once used for the industrial purification of nickel in the Mond process.[18]

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.

Organic and main group chemistry

In the presence of strong acids and water, carbon monoxide reacts with olefins to form carboxylic acids in a process known as the Koch-Haaf reaction.[19] In the Gattermann-Koch reaction, arenes are converted to benzaldehyde derivatives in the presence of AlCl3 and HCl.[20] Organolithium compounds, e.g. butyl lithium react with CO, but this reaction enjoys little use.

Although CO reacts with carbocations and carbanions, it is relatively unreactive toward organic compounds without the intervention of metal catalysts.[21]

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 Na2C2O2 (sodium acetylenediolate), and potassium to give K2C2O2 (potassium acetylenediolate) and K2C6O6 (potassium rhodizonate).

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 high pressure (over 5 gigapascals), carbon monoxide disproportionates into carbon dioxide CO2 and a solid polymer of carbon and oxygen (in 3:2 atomic ratio).[22][23]

Uses

Chemical industry

Carbon monoxide is a major industrial gas that has many applications in bulk chemicals manufacturing.[24]

Large quantities of aldehydes are produced by the hydroformylation reaction of alkenes, CO, and H2. In one of many applications of this technology, hydroformylation is coupled to the Shell Higher Olefin Process to give precursors to detergents. Methanol is produced by the hydrogenation of CO. In a related reaction, the hydrogenation of CO is coupled to C-C bond formation, as in the Fischer-Tropsch process where CO 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 red and fresh. The CO 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 colour can persist much longer than in normally packaged meat.[25] Typical levels of CO used are 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 odour.[26] Despite this ruling, the technology remains controversial in the US for fears that it is deceptive and masks spoilage.[27] In 2007 a bill[28] was introduced to the United States House of Representatives to label modified atmosphere carbon monoxide packaging as a "color additive;" however, the bill died in subcommittee. The practice is banned in many other countries, including Canada, Japan, Singapore and the European Union.[29] [30] [31]

Medicine

CO is also currently being studied in several research laboratories throughout the world for its anti-inflammatory and cytoprotective properties that can be used therapeutically to prevent the development of a series of pathologic conditions such as ischemia reperfusion injury, transplant rejection, atherosclerosis, sepsis, severe malaria or autoimmunity. However, there are yet no clinical applications of CO in humans[citation needed].

See also

References

  1. ^ a b c Blumenthal, Ivan (June 2001). "Carbon monoxide poisoning" (in English). J R Soc Med (The Royal Society of Medicine) 94 (6): 270–272. PMID 11387414. PMC: PMC1281520. http://www.pubmedcentral.nih.gov/articlerender.fcgi?tool=pubmed&pubmedid=11387414. Retrieved on May, 2009. 
  2. ^ 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. doi:10.1103/PhysRev.78.140. 
  3. ^ W. Kutzelnigg (2002). Einführung in die Theoretische Chemie. Wiley-VCH. ISBN 3-527-30609-9. 
  4. ^ Omaye ST. (2002). "Metabolic modulation of carbon monoxide toxicity". Toxicology 180 (2): 139–50. doi:10.1016/S0300-483X(02)00387-6. 
  5. ^ "OSHA CO guidlines" (in English). OSHA. http://www.osha.gov/SLTC/healthguidelines/carbonmonoxide/recognition.html. Retrieved on May, 2009. 
  6. ^ Reverso Dictionary. Accessed April 14, 2009
  7. ^ Webster's Online
  8. ^ Ganong, William F (2005). "37" (in English). Review of medical physiology (22 ed.). Published by McGraw-Hill Professional. p. 684. ISBN 0071440402. http://books.google.com/books?id=OLa8vDBXDD4C&dq=Ganong+WF.+Review+of+Medical+Physiology.+Norwalk+Ct:+Appleton+%26+Lange,+1995&printsec=frontcover&source=bn&hl=en&ei=QU8dSvnhG5OeMoHU0J8P&sa=X&oi=book_result&ct=result&resnum=4#PPA684,M1. Retrieved on May, 2009. 
  9. ^ Prockop LD, Chichkova RI (2007). "Carbon monoxide intoxication: an updated review". J Neurol Sci 262 (1-2): 122-30. PMID 17720201. 
  10. ^ a b Wu, L; Wang, R. (2005 Dec). "Carbon Monoxide: Endogenous Production, Physiological Functions, and Pharmacological Applications". Pharmacol Rev 57 (4): 585-630. PMID 16382109. http://pharmrev.aspetjournals.org/cgi/content/full/57/4/585#XI._Conclusions_and_Perspectives. Retrieved on May 26, 2009. 
  11. ^ 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. http://mic.sgmjournals.org/cgi/reprint/144/9/2377. 
  12. ^ Jaouen, G., Ed. (2006). Bioorganometallics: Biomolecules, Labeling, Medicine. Weinheim: Wiley-VCH. ISBN 3-527-30990-X. 
  13. ^ Roberts, G. P.; Youn, H.; Kerby, R. L. (2004). "CO-Sensing Mechanisms". Microbiology and Molecular Biology Reviews 68: 453–473. doi:10.1128/MMBR.68.3.453-473.2004. PMID 15353565. 
  14. ^ a b "Basic Information : Carbon Monoxide". http://www.epa.gov/iaq/co.html. Retrieved on 2007-12-01. 
  15. ^ Singer, Siegfried Fred (1975). The Changing Global Environment. Dordrecht: D. Reidel Pub. Co.. pp. 90. ISBN 9789027704023. http://books.google.com/books?id=Ww3DnCF_KZcC&pg=PA90&lpg=PA90&dq=%22carbon+monoxide%22+ppm+concentration+traffic+%22mexico+city%22&source=web&ots=SzIyYgUxWh&sig=ZM4p8whF1mtj1kE3XSx1YUx70zw#PPA90,M1. 
  16. ^ a b Gosink, Tom (1983-01-28). "What Do Carbon Monoxide Levels Mean?" (HTML). Alaska Science Forum. Geophysical Institute, University of Alaska Fairbanks. http://www.gi.alaska.edu/ScienceForum/ASF5/588.html. Retrieved on 2007-12-01. 
  17. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 200. ISBN 0-12-352651-5.
  18. ^ 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. 
  19. ^ Koch, H.; Haaf, W. (1973). "1-Adamantanecarboxylic Acid". Org. Synth.; Coll. Vol. 5: 20. 
  20. ^ G. H. Coleman, David Craig (1943). "p-Tolualdehyde". Org. Synth.; Coll. Vol. 2: 583. 
  21. ^ 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
  22. ^ A. I. Katz, D. Schiferl, and R. L. Mills (1984): J. Phys. Chem., 88 (15), 3176–3179
  23. ^ W. J. Evans, M. J. Lipp, C.-S. Yoo, H. Cynn, J. L. Herberg, R. S. Maxwell, and M. F. Nicol (2006). "Pressure-Induced Polymerization of Carbon Monoxide: Disproportionation and Synthesis of an Energetic Lactonic Polymer". Chemistry of Materials, vol. 18, 2520–2531. doi: 10.1021/cm0524446
  24. ^ Elschenbroich, C.;Salzer, A. ”Organometallics : A Concise Introduction” (2nd Ed) Wiley-VCH: Weinheim, 2006. ISBN 3-527-28165-7
  25. ^ 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): 157–64. doi:10.1016/S0309-1740(98)00163-6. 
  26. ^ Eilert EJ (2005). "New packaging technologies for the 21st century". Journal of Meat Science 71 (1): 122–27. doi:10.1016/j.meatsci.2005.04.003. 
  27. ^ "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 on 2007-04-18. 
  28. ^ "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:. 
  29. ^ "Proof in the Pink? Meat Treated to Give It Fresh Look" (in English). ABC News. November 14, 2007. http://abcnews.go.com/GMA/Consumer/Story?id=3863064&page=1. Retrieved on May 27, 2009. 
  30. ^ Carbon Monoxide in Meat Packaging: Myths and Facts, American Meat Institute, 2008, http://www.meatami.com/ht/a/GetDocumentAction/i/40141, retrieved on May 2009 
  31. ^ "CO in packaged meat" (in English). Carbon Monoxide Kills Campaign. http://www.carbonmonoxidekills.com/packages_meat.htm. Retrieved on May 2009. 

External links

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Inorganic compounds of carbon
Common oxides
CO2 · CO
Exotic oxides
C2O2 · C3O2 · C5O2 · C2O · CO3
Compounds derived from oxides
Ionic compounds






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