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Carbon Monoxide

Carbon monoxide (CO)

Carbon monoxide, with the chemical formula CO, is a colorless , odorless and tasteless, yet highly toxic gas. Its molecules consist of one carbon atom covalently bonded to one oxygen atom. It is the simplest oxocarbon, and can be viewed as the anhydride of formic acid. There are two covalent bonds and a coordinate covalent bond between the oxygen and carbon atoms.

Carbon monoxide is produced from the partial oxidation of carbon-containing compounds, notably in internal-combustion engines. Carbon monoxide forms in preference to the more usual carbon dioxide when there is a reduced availability of oxygen present during the combustion process. Carbon monoxide has significant fuel value, burning in air with a characteristic blue flame, producing carbon dioxide. Despite its serious toxicity, CO plays a highly useful role in modern technology as a precursor to myriad products.

Uses

Chemical industry

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

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 and pork. 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, giving the appearance of freshness. Typical levels of CO used are 0.4% to 0.5%. The technology was first given generally recognized as safe status by the U.S. Food and Drug Administration (FDA) in 2002 for use as a secondary packaging system. In 2004 the FDA approved CO as primary packaging method, declaring that CO does not mask spoilage odour. Despite this ruling, the technology remains controversial in the US for fears that it is deceptive and masks spoilage. Elsewhere, coloring meat to make it appear fresh is banned in many other countries.

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.

Biological and physiological properties

Toxicity

Carbon monoxide is colorless and odorless, but very toxic. Carbon monoxide poisoning is the most common type of fatal poisoning in many countries. Exposures can lead to significant toxicity of the central nervous system and heart, often with long-term sequelae. Carbon monoxide can also have severe effects on the fetus of a pregnant woman.

Carbon monoxide combines with haemoglobin in the blood converting it to carboxy-haemoglobin (HbCO), which is ineffective for delivering oxygen (a condition known as anoxemia). Myoglobin, and mitochondrial cytochrome oxidase are thought to be compromised too. Concentrations as low as 667 ppm can cause up to 50% of the body's haemoglobin to be. Carboxy-haemoglobin can revert to haemoglobin, but the recovery takes time bcause HbCO is fairly stable.

Symptoms of mild poisoning include headaches and dizziness at concentrations less than 100 ppm. As a result, exposures of this level can be life-threatening. In the United States, OSHA limits long-term workplace exposure levels to 50 ppm.

reatment largely consists of administering 100% oxygen or hyperbaric oxygen therapy, although the optimum treatment remains controversial. Domestic carbon monoxide poisoning can be prevented by the use of household carbon monoxide detectors.

Natural processes

Carbon monoxide is produced naturally in the human body as a breakdown of heme (which is one of hemoglobin moieties), a substrate for the enzyme heme oxygenase. The enzymatic reaction results in breakdown of heme to CO, biliverdin and Fe3+ radical. The endogenously produced CO may have important physiological roles in the body (eg as a neurotransmitter or a blood vessels relaxant). In addition CO regulates inflammatory reactions in a manner that prevents the development of several diseases such as atherosclerosis or severe malaria.

CO is a nutrient for methanogenic bacteria, 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.

A haeme-based CO-sensor protein, CooA, is known. 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
  • 5 to 15 ppm - levels near properly adjusted gas stoves in homes
  • 100-200 ppm - Mexico City central area from autos etc.
  • 5,000 ppm - chimney of a home wood fire
  • 7,000 ppm - undiluted warm car exhaust - without catalytic converter

As atmospheric component

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.

As urban pollutant

Carbon monoxide is a major atmospheric pollutant in urban areas, chiefly from exhaust of internal combution 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.

As an indoor pollutant

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.

In industry

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:

    + 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 → H + 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
    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 favored over CO2 in high temperatures.

In the laboratory

Carbon monoxide is conveniently produced in the lab 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 by the reaction

    Zn + CaCO3 → ZnO + CaO

Chemistry

Molecular structure

The CO molecule possesses a bond length of 0.1128 nm. Formal charge and electronegativity difference cancel each other out. The result is a small dipole moment with its negative end on the carbon atom. 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 dipole moment. This is also the reason that almost all chemistry involving carbon monoxide occurs through the carbon atom, and not the oxygen.

Coordination chemistry

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 x−y 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.

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. In the Gattermann-Koch reaction, arenes are converted to benzaldehyde derivatives in the presence of AlCl3 and HCl. 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. 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).

Effects of Various CO Levels
Carbon Monoxide Level in PPM Resulting Conditions on Humans
50 Permissible Exposure Level for 8 hours (OSHA).
200 Possible mild frontal headache in 2 to 3 hours.
400 Frontal headache and nausea after 1 to 2 hours. Occipital after 2-1/2 to 3-1/2 hours.
800 Headache, dizziness, and nausea in 45 minutes. Collapse and possible death in 2 hours
1,600 Headache, dizziness, and nausea in 20 minutes. Collapse and death in 1 hour.
3,200 Headache and dizziness in 5 to 10 minutes. Unconsciousness and danger of death in 30 minutes
6,400 Headache and dizziness in 1 to 2 minutes. Unconsciousness and danger of death in 10 to 15 minutes.
12,800 Immediate effects-unconsciousness. Danger of death in 1 to 3 minutes.