carbon monoxide

Production

Carbon monoxide is so fundamentally important that many methods have been developed for its production.^[1]

"Producer gas" is formed by combustion of carbon in oxygen at high temperatures 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 CO₂ 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

"Synthesis gas" is produced via the endothermic reaction of steam and carbon:

H₂O + C → H₂ + CO ΔH = 131 kJ/mol

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

MO + C → M + CO ΔH = 131 kJ/mol

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 favoured over CO₂ in high temperatures.

CO is the anhydride of formic acid. As such it is conveniently produced by the dehydration of formic acid, for example with sulfuric acid: This reaction is the basis of a convenient laboratory synthesis of CO.

Another laboratory preparation for carbon monoxide is to heat a mixture of finely divided zinc metal mixed with calcium carbonate.

Zn + CaCO₃ → ZnO + CaO + CO

Structure

The CO molecule is characterized by a bond length of 0.111 nm. This distance is consistent with a partial triple bond. The molecule has a small dipole moment and can be represented by three resonance structures:

The most important resonance form is on the left.^[2] Although one might expect a large dipole moment due to the electronegativity diferences between oxygen and carbon, the formal negative charge on the carbon cancels this difference, resulting in a small dipole moment.

Nitrogen is isoelectronic to carbon monoxide, which means it has the similar electron configuration. The physical properties of N₂ and CO are similar.

Principal chemical reactions

Industrial uses

Carbon monoxide is a major industrial gas that has many applications in bulk chemicals manufacturing.^[3]

High volume 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 to be converted to petrol.

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

Coordination chemistry

Main article: metal carbonyl

Most metals form coordination complexes containing covalently attached carbon monoxide. 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.^[4]

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 Fe(CO)₅:

Many metal-CO complexes are prepared by decarbonylation of organic solvents, not from CO. For instance, boiling iridium trichloride with triphenyl phosphine in methoxyethanol (or DMF) gives Irl(CO)(PPh₃)₂.

Organic and main group chemistry

CO reacts with carbocations and carbanions, but otherwise is unreactive toward organic compounds without the intervention of metal catalysts.^[5]

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.^[6] In the Gattermann-Koch reaction, arenes are converted to benzaldehyde derivatives in the presence of AlCl3 and HCl.^[7]

Organolithium compounds, e.g. butyl lithium react with CO, but this reaction enjoys little use.

With main group reagents, CO does undergo 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 (acetylenediolate) and Na₂C₄O₄ (sodium squarate).

Carbon monoxide in the atmosphere

Carbon monoxide, though thought of as a pollutant today, has always been present in the atmosphere, chiefly as a product of volcanic activity. 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. It also occurs naturally in bushfires. 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. Carbon monoxide is created when carbon-containing fuels are burned incompletely. 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.

Anthropogenic CO from automobile and industrial emissions may contribute to the greenhouse effect and global warming. In urban areas carbon monoxide, along with aldehydes, 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. Carbon monoxide is also a constituent of tobacco smoke.

Physiological roles

One reaction in the body produces CO. Carbon monoxide is produced naturally as a breakdown of haemoglobin, heme, is a substrate for the enzyme heme oxygenase which produces CO and biliverdin. The biliverdin is converted to bilirubin by biliverdin reductase in macrophages of the reticuloendothelial system. The lipid soluble unconjugated bilirubin is transported in the blood bound to albumin, taken up by the hepatocytes, conjugated with glucuronic acid and transported into the bile canaliculi for excretion from the body. The endogenously produced CO may have important physiological roles in the body (eg as a neurotransmitter).

CO has a significant role in microbiology being a nutrient for methanogenic bacteria,^[8] 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.^[9]

A haeme-based CO-sensor protein, CooA, is known.^[10] The scope of its biological role is still unclear, it is apparently part of a signalling pathway in bacteria and archaea, but its occurrence in mammals is not established.

History

Carbon monoxide was first prepared by the French chemist de Lassone in 1776 by heating zinc oxide with coke. He 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 English chemist William Cruikshank in the year 1800.

The toxic properties of CO were first thoroughly investigated by the French physiologist Claude Bernard around 1846. He poisoned dogs with the gas, and noticed that their blood was more rutilant in all the vessels. 'Rutilant' is a French word, but also has an entry in English dictionaries, meaning ruddy, shimmering, or golden. However, it was translated at the time as crimson, scarlet, and now is famously known as 'cherry pink'.

During World War Two, 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. The CO in this case is known as "producer gas". Carbon monoxide was also used as an extermination method during the Holocaust at some Nazi extermination camps.

Safety

Main article: Carbon monoxide poisoning

CO is toxic. Carbon monoxide detectors for homes are now readily available.

See also

• Carbon monoxide (data page)
• List of deaths from carbon monoxide poisoning

Reference

1. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
2. ^ G. L. Miessler and D. A. Tarr “Inorganic Chemistry” 3rd Ed, Pearson/Prentice Hall publisher, ISBN 0-13-035471-6.
3. ^ Elschenbroich, C.;Salzer, A. ”Organometallics : A Concise Introduction” (2nd Ed) Wiley-VCH: Weinheim, 2006. ISBN 3527281657
4. ^ Mond L, Langer K, Quincke F (1890). "Action of carbon monoxide on nickel". Journal of the Chemical Society: 749-753. DOI
5. ^ 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.
6. ^ Koch, H.; Haaf, W. "1-Adamantanecarboxylic Acid" Organic Syntheses, Collected Volume 5, p.20 (1973).
7. ^ Coleman, G. H.; Craig, D. "p-Tolualdehyde" Organic Syntheses, Collected Volume 2, p.583 (1943).
8. ^ Thauer, R. K. "Biochemistry of methanogenesis: a tribute to Marjory Stephenson" Microbiology (1998), volume 144, 2377–2406.
9. ^ Bioorganometallics: Biomolecules, Labeling, Medicine; Jaouen, G., Ed. Wiley-VCH: Weinheim, 2006.3-527-30990-X.
10. ^ Roberts, G. P.; Youn, H.; Kerby, R. L. "CO-Sensing Mechanisms" Microbiology and Molecular Biology Reviews (2004), volume 68, pages 453-473.

External links

• International Chemical Safety Card 0023
• National Pollutant Inventory - Carbon Monoxide
• NIOSH Pocket Guide to Chemical Hazards
• CID 281 from PubChem
• United States Environmental Protection Agency Carbon Monoxide page
• External MSDS data sheet
• Carbon Monoxide Kills Campaign Site

• Link page to external chemical sources.