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22 Sep.,2023

 

Carbon monoxide

History

Sources

Physiological effects

Uses

Resources

Carbon monoxide is a compound of carbon and oxygen with the chemical formula CO. It is a colorless, odorless, tasteless, toxic gas. Carbon monoxide is poisonous to all warm-blooded animals (when it is inhaled and combined with hemoglobin in the blood, which prevents the absorption of oxygen) and to many other life forms. It has a density of 1.250 g/L at 32°F (0°C) and 760 mm Hg pressure. Carbon dioxide can be converted into a liquid at its boiling point of -312.7°F (-191.5°C) and then to a solid at its freezing point of -337°F (-205°C). It is about 3% lighter than air.

The discovery of carbon monoxide is often credited to the work of the English chemist and theologian Joseph Priestley (1733–1804). In the period between 1772 and 1799, Priestley gradually recognized the nature of this compound and showed how it was different from carbon dioxide, with which it often appeared. Nonetheless, carbon monoxide had been well known and extensively studied in the centuries prior to Priestley’s work. As early as the late 1200s, Spanish alchemist Arnold of Villanova (c.1238–c.1310) described a poisonous gas produced by the incomplete combustion of wood that was almost certainly carbon monoxide.

In the five centuries between the work of Arnold and that of Priestley, carbon monoxide was studied and described by a number of prominent alchemists and chemists. Many made special mention of the toxicity of the gas. French scientist Johann (Jan) Baptista van Helmont (1580–1644) in 1644 wrote that he nearly died from inhaling gas carbonum, apparently a mixture of carbon monoxide and carbon dioxide.

An important milestone in the history of carbon monoxide came in 1877 when French physicist Louis Paul Cailletet (1832–1913) found a method for liquefying the gas. Two decades later, a particularly interesting group of compounds made from carbon monoxide, the carbonyls, were discovered by the French chemist Paul Sabatier (1854–1941).

Carbon monoxide is the twelfth most abundant gas in the atmosphere. It makes up about 1.2→× 10-5% of a sample of dry air in the lower atmosphere. The major natural source of carbon monoxide is the combustion of wood, coal, and other naturally occurring substances on the Earth’s surface. Huge quantities of carbon monoxide are produced, for example, during a forest fire or a volcanic eruption. The amount of carbon monoxide produced in such reactions depends on the availability of oxygen and the combustion temperature. High levels of oxygen and high temperatures tend to produce complete oxidation of carbon, with carbon dioxide as the final product. Lower levels of oxygen and lower temperatures result in the formation of higher percentages of carbon monoxide in the combustion mixture.

Commercial methods for producing carbon monoxide often depend on the direct oxidation of carbon under controlled conditions. For example, producer gas is made by blowing air across very hot coke (nearly pure carbon). The final product consists of three gases, carbon monoxide, carbon dioxide, and nitrogen in the ratio of 6 to 1 to 18. Water gas is made by a similar process, by passing steam over hot coke. The products in this case are hydrogen (50%), carbon monoxide (40%), carbon dioxide (5%) and other gases (5%). Other methods of preparation are also available. One of the most commonly used involves the partial oxidation of hydrocarbons obtained from natural gas.

The toxic character of carbon monoxide has been well known for many centuries. At low concentrations, carbon monoxide may cause nausea, vomiting, restlessness, and euphoria. As exposure increases, a person may lose consciousness and go into convulsions. Death is a common result. The U.S. Occupational Safety and Health Administration (OSHA) has established a limit of 35 ppm (parts per million) of carbon monoxide in workplaces where a person may be continually exposed to the gas (Figure 1).

The earliest explanation for the toxic effects of carbon monoxide was offered by the French physiologist Claude Bernard in the late 1850s. Bernard pointed out that carbon monoxide has a strong tendency to replace oxygen in the respiratory system. Someone exposed to high concentrations of carbon monoxide may actually begin to suffocate as his or her body is deprived of oxygen.

Today a fairly sophisticated understanding is known of the mechanism by which carbon monoxide poisoning occurs. Normally, oxygen is transported from the lungs to cells in red blood cells. This process occurs when oxygen atoms bond to an iron atom at the center of a complex protein molecule known as oxyhemoglobin. It is an unstable molecule that decomposes in the intercellular spaces to release free oxygen and hemoglobin. The oxygen is then available to carry out metabolic reactions in cells, reactions from which the body obtains energy.

If carbon monoxide is present in the lungs, this sequence is disrupted. Carbon monoxide bonds with iron in hemoglobin to form carbonmonoxyhemoglobin, a complex somewhat similar to oxyhemoglobin. Carbonmonoxyhemoglobin is, however, a more stable

compound than is oxyhemoglobin. When it reaches cells, it has much less tendency to break down, but continues to circulate in the bloodstream in its bound form. As a result, cells are unable to obtain the oxygen they need for metabolism and energy production dramatically decreases. The clinical symptoms of carbon monoxide poisoning described above are manifestations of these changes.

Carbon monoxide poisoning—at least at moderate levels—is common in everyday life. Poorly vented charcoal fires, improperly installed gas appliances, and the exhaust from internal combustion vehicles are among the most common sources of the gas. In fact, levels of carbon monoxide in the air can become dangerously high in busy urban areas where automotive transportation is extensive. Cigarette smokers may also be exposed to dangerous levels of the gas. Studies have shown that the one to two pack-a-day smoker may have up to 7% of the hemoglobin in her or his body tied up in the form of carbonmonoxyhemoglobin.

Carbon monoxide is a very important industrial compound. In the form of producer gas or water gas, it is widely used as a fuel in industrial operations. The gas is also an effective reducing agent. For example, when carbon monoxide is passed over hot iron oxides, the oxides are reduced to metallic iron, while the carbon monoxide is oxidized to carbon dioxide.

In another application a mixture of metallic ores is heated to 122–176°F (50–80°C) in the presence of producer gas. All oxides except those of nickel are reduced to their metallic state. This process, known as the Mond process, is a way of separating nickel from other metals with which it commonly occurs.

Yet another use of the gas is in the Fischer-Tropsch process for the manufacture of hydrocarbons and their oxygen derivatives from a combination of hydrogen and carbon monoxide. Carbon monoxide

KEY TERMS

Combustion— A form of oxidation that occurs so rapidly that noticeable heat and light are produced.

Hemoglobin— An iron-containing, complex molecule carried in red blood cells that binds oxygen for transport to other areas of the body.

Incomplete combustion— Combustion that occurs in such a way that fuel is not completely oxidized. The incomplete combustion of carbon-containing fuel, for example, always results in the formation of some carbon monoxide.

Intercellular spaces— The spaces between cells in tissue.

Reductant (reducing agent)— A chemical substance that reduces materials by donating electrons to them.

Toxicity— The extent to which a substance is poisonous.

also reacts with certain metals, especially iron, cobalt, and nickel, to form compounds known as carbonyls. Some of the carbonyls have unusual physical and chemical properties that make them useful in industry. The highly toxic nickel tetracarbonyl, for example, is used to produce very pure nickel coatings and powders.

Catalytic converters are used in automobiles to reduce carbon monoxide emissions. Recent nanotechnology advances (those technologies involving microscopic devices) have developed a nanoparticle catalyst made of nonreactive metals, helping to reduce more efficiently such poisonous gases as nitrogen oxides and carbon monoxide.

See also Metallurgy.

BOOKS

Ede, Andrew. The Chemical Element: A Historical Perspective. Westport, CT: Greenwood Press, 2006.

Emsley, John. Nature’s Building Blocks: An A-Z Guide to the Elements. Oxford, UK: Oxford University Press, 2001.

Merck. The Merck Index. Whitehouse Station, NJ: Merck; London: Harcourt, 2001.

Lide, D.R., ed. CRC Handbook of Chemistry and Physics Boca Raton: CRC Press, 2001.

Matthews, John A., E.M. Bridges, and Christopher J. Caseldine The Encyclopaedic Dictionary of Environmental Change. New York: Edward Arnold, 2001.

Partington, J.R. A Short History of Chemistry. 3rd ed. London: Macmillan & Company, 1957.

Stwertka, Albert. A Guide to the Elements. New York: Oxford University Press, 2002.

David E. Newton