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carbon monoxide

carbon monoxide is a gas which is best known to us as a product of incomplete combustion. As such, mankind must have been aware of its deadly effect since the discovery and use of fire, and increasingly so as the development of the industrial revolution led to more use of combustion as a source of energy. The important producers of carbon monoxide are industrial processes, heating equipment, accidental fire, cigarettes, and the internal combustion engine. Blast furnace gas contains 25% carbon monoxide, and coal gas, which was used as a fuel in Europe up until North Sea (natural) gas became plentiful, contains 16%. Carbon monoxide poisoning is the most common cause of fatal gassing and is the cause of death in about 90% of fire victims. Domestic gas supplies still lead to carbon monoxide poisoning, but now due to leakage of products of combustion from a damaged flue or poorly maintained equipment, rather than the fuel itself, since natural gas is carbon monoxide free. In the mining industry carbon monoxide contaminates the atmosphere during and after fires or explosions. The ‘afterdamp’ occurring in such situations is a mixture of carbon dioxide and carbon monoxide.

Carbon monoxide is a colourless, odourless gas which is tasteless and non-irritant. It is somewhat less dense than air and, although it is a product of imperfect combustion, it is inflammable. The gas was first identified by Joseph Priestley in the eighteenth century, but it was Claude Bernard in 1870 who discovered the affinity between carbon monoxide and haemoglobin which accounts for its deadliness: carboxyhaemoglobin is formed and oxygen transport from the lungs to the tissues disrupted. In 1895 J. S. Haldane demonstrated that the formation of carboxyhaemoglobin is an equilibrium reaction which depends upon the relative partial pressures of carbon monoxide and oxygen in inspired gas. Haldane's interest was stimulated by the problems caused by carbon monoxide in British coal mines. By breathing carbon monoxide gas which was passed through a bottle containing a mouse, he was able to determine that man was very much more resistant to the gas. Small animals such as mice and canaries, who are more vulnerable than man due to their high metabolic rate, were used in mines to give an indication of carbon monoxide contamination. Canaries responded to the gas by falling off their perches before workers noticed any ill effects, and this normally gave ample warning. Occasionally, however, in low concentrations of the order of 0.05% carbon monoxide, the bird adapted to the gas and the workers could collapse while the bird remained well.

Carbon monoxide, like oxygen, has an affinity for iron-containing molecules, but it is about 210 times more effective in binding to iron-containing haemoglobin than oxygen is. Since air contains 21% oxygen this means that only 0.1% carbon monoxide in the air will eventually lead to 50% of the haemoglobin being combined to form carboxyhaemoglobin. Once carboxyhaemoglobin is formed, and after exposure ceases, it takes 4–5 hours for its level in the blood to fall, exponentially, by 50%. The ill effect of the gas can therefore be cumulative, and a person can be poisoned by intermittent exposure during the day.

Because carboxyhaemoglobin does not carry oxygen, a level of 50% means that the oxygen carrying capacity of the blood is reduced by 50% and there is a corresponding reduction in the ability to perform maximum exercise. The body compensates for the blood's reduction in oxygen carrying capacity by increasing cardiac output, and in the early stages of carbon monoxide poisoning the heart beats faster and more strongly. Unfortunately, haemoglobin is not the only molecule affected. Muscle myoglobin also binds carbon monoxide, 60 times more effectively than it binds oxygen. This results in a reduction of heart muscle contractility and a failure of the body's compensatory mechanisms, leading to profound tissue hypoxia, which can be fatal. The presence of carboxyhaemoglobin also diminishes the oxygen held by the normal haemoglobin, which further compounds the hypoxic effect. As tissue oxygen level falls, carbon monoxide is able to bind to other iron-containing molecules: notably cytochrome P450, an important drug-metabolizing enzyme, and cytochrome A3, an enzyme in the terminal respiratory chain which can also be poisoned by cyanide.

The scientific history of carbon monoxide is not one of uniform gloom, however. The intense affinity of carbon monoxide for haemoglobin has allowed low concentrations to be used as a marker for measurement of the speed of blood through the lungs and the surface area of the lung available for the transfer of oxygen. This latter remains as one of the standard lung function tests. In 1951 Sjöstrand discovered that Haldane's poison gas is a normal product of the body's metabolism. The enzyme haem oxygenase breaks down the haem from senescent red blood cells, and this reaction produces carbon monoxide and bile salts. The bile salts are excreted by the liver and the carbon monoxide released gives the blood a normal carboxyhaemoglobin level of 0.2–1.0%. This endogenous carbon monoxide was thought to be just a waste product, but more recent work by Verma has demonstrated that a type of haem oxygenase is located in specific areas in the brain, and suggested that the carbon monoxide produced acts as a neurotransmitter. The carbon monoxide activates the enzyme guanylyl cyclase, as does nitric oxide, regulating the intracellular levels of the second messenger cyclic GMP, which in turn regulates cellular activity. Other workers have demonstrated the haem oxygenase enzyme system in blood vessel walls and demonstrated that the carbon monoxide released causes vasodilation, as does nitric oxide. So far, endogenous carbon monoxide release has been suggested to have a role in the sense of smell, memory, cerebellar function (and hence the body's balance and co-ordination), control of blood hormone levels from the hypothalamus, and control of smooth muscle tone and vasodilatation.

The symptoms of carbon monoxide poisoning depend on the concentration breathed. The victim may pass out without warning, but often the onset of poisoning is slow. Headache, with or without nausea, is common, and this may relate to carbon monoxide's vasodilating effect. Drowsiness and lethargy then occur, along with breathlessness on exertion. At any stage there may be chest pain; this is angina due to cardiac hypoxia. At the stage of lethargy and drowsiness, cerebral function is affected and the person may not be able to think well enough to make an escape effort. Coma follows, and death. Treatment is by removal to an uncontaminated atmosphere and the administration of 100% oxygen. Hyperbaric oxygen speeds up recovery, and there is increasing evidence that it reduces long-term neurological problems.

Endogenous carbon monoxide function is undoubtedly disrupted during poisoning, but at our present state of knowledge it is difficult to say how this contributes to the toxic action of exogenous carbon monoxide. It may well be that our picture of the mechanisms of carbon monoxide poisoning will change as the function of endogenous carbon monoxide becomes clearer. Patients with carbon monoxide poisoning may have very poor balance and yet have good cerebral function. Short-term memory may also be severely disrupted. It is tempting to link these two features with the functions suggested for endogenous carbon monoxide. No doubt time will tell if there is a relationship.

John A. S. Ross

Bibliography

World Health Organisation (1979). Environmental health criteria 13: carbon monoxide. WHO, Geneva.
Dawson, T. M. and and Snyder, S. H. (1994). Gases as biological messengers: nitric oxide and carbon monoxide in the brain. (Review article.) Journal of Neuroscience, 14(9), 5147–59.

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Carbon Monoxide

Carbon monoxide

Carbon monoxide is a compound of carbon and oxygen in which the ratio of the two elements is one atom of carbon to one atom of oxygen. Its formula is CO. Carbon monoxide is a colorless, odorless, tasteless, poisonous gas. Most people have heard about carbon monoxide because of its toxic effects. People who live or work in crowded urban areas may become ill with headaches and nausea because of exposure to carbon monoxide in polluted air. In higher concentrations, the gas can even cause death.

History

The early history of gases such as carbon monoxide is sometimes difficult to trace. Until the early 1600s, scientists did not realize that the material we call air is actually a mixture of gases. As early as the late thirteenth century, Spanish alchemist Arnold of Villanova (c. 12351311) described a poisonous gas formed by the burning of wood; this gas was almost certainly carbon monoxide.

Flemish scientist Jan Baptista van Helmont (c. 15801644; some sources give death date as 1635) nearly died as a result of inhaling gas carbonum, apparently a mixture of carbon monoxide and carbon dioxide. Credit for the discovery of carbon monoxide, however, is usually given to English chemist and theologian Joseph Priestley (17331804). During the period between 1772 and 1799, Priestley gradually recognized the difference between carbon dioxide and carbon monoxide and correctly stated the properties of the latter gas.

Sources

Like carbon dioxide, carbon monoxide is formed naturally during the combustion (burning) of wood, coal, and other naturally occurring substances. Huge quantities of carbon monoxide are produced, for example, during a forest fire or a volcanic eruption.

Words to Know

Combustion: Oxidation that occurs so rapidly that noticeable heat and light are produced; burning.

Hemoglobin: An complex iron-containing molecule that transports oxygen through the circulatory system.

Incomplete combustion: Combustion that occurs in such a way that fuel is not completely oxidized ("burned up"). The incomplete combustion of carbon-containing fuels (such as coal and oil) always results in the formation of some carbon monoxide.

Reducing agent: A substance that removes oxygen from some other material.

Toxic: Poisonous.

The relative amounts of carbon monoxide or carbon dioxide that form during combustion depend on two factors: the amount of oxygen present and the combustion temperature. When a large supply of oxygen is present and when the combustion temperature is high, carbon dioxide is more likely to be formed. With limited supplies of oxygen and at lower temperatures, carbon monoxide is produced.

Carbon monoxide is not extracted from the air very easily but is produced commercially by the controlled oxidation of carbon. For example, producer gas is a product made by blowing air across very hot coke (nearly pure carbon). Producer gas 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 processpassing steam over hot coke. The products in this case are hydrogen, carbon monoxide, carbon dioxide, and other gases in the ration of 10 to 8 to 1 to 1.

Physiological effects

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

Scientists now know how carbon monoxide poisoning occurs. Normally, oxygen is transported from the lungs to cells by means of red blood cells. This process occurs when oxygen atoms bond to an iron atom in the middle of a complex molecule known as oxyhemoglobin. Oxyhemoglobin is a fairly unstable molecule that breaks down to release free oxygen and hemoglobin for use by the body's cells. The oxygen is then available to carry out reactions in cells from which the body gets energy.

If carbon monoxide is present in the lungs, this sequence of reactions is disrupted. Carbon monoxide bonds with iron in hemoglobin to form carbonmonoxyhemoglobin, a complex somewhat similar to oxyhemoglobin. Carbonmonoxyhemoglobin, however, is a more stable compound than oxyhemoglobin. When it reaches cells, it has little tendency to break apart; instead, it continues to circulate in the bloodstream in its bound form. As a result, cells are unable to obtain the oxygen they need for energy production, and the symptoms of carbon monoxide poisoning begin to appear.

Carbon monoxide poisoningat least at moderate levelsis so common in everyday life that carbon monoxide detectors, similar to smoke alarms, are found in many businesses and homes. Poorly ventilated charcoal fires, improperly installed gas appliances, and exhaust from automobiles and trucks are the most common sources of the gas. In fact, levels of carbon monoxide in the air can become dangerously high in busy urban areas that have large numbers of cars and trucks. Cigarette smokers may also be exposed to harmful levels of the gas. Studies have shown that the one-to-two pack-a-day smoker may have up to 7 percent of the hemoglobin in her or his body tied up in the form of carbonmonoxyhemoglobin.

Uses

Carbon monoxide is used in industry primarily as a source of energy and as a reducing agent. Both producer and water gas are burned as fuels for a variety of industrial operations. As a reducing agent, carbon monoxide is used to convert the naturally occurring oxide of a metal to the pure metal. When carbon monoxide is passed over hot iron oxides, for example, the oxides are converted to metallic iron.

[See also Carbon dioxide; Carbon family ]

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Carbon Monoxide

CARBON MONOXIDE

Carbon monoxide (CO) is a clear, colorless, odorless, and insidious poison that is responsible for hundreds of inadvertent and preventable deaths in the United States each year. The major environmental source of CO is incomplete combustion of carbonaceous fossil fuels. The reason for its toxicity is that it combines with the oxygen-carrying site of hemoglobin, the red protein within red blood cells that is responsible for delivering oxygen from the lung to body tissues. CO has a more than two-hundredfold greater affinity for this oxygen-carrying site than does oxygen. This means that, at sea level, exposure to 1,000 parts per million (ppm) CO in 20 percent oxygen (200,000 ppm) would lead, at equilibrium, to about 50 percent of hemoglobin sites being combined with CO rather than oxygen. Fortunately, it requires eight to twelve hours for maximum blood levels to be achieved when the body encounters a new CO concentration, otherwise mainstream cigarette smoke, which contains even higher levels of CO, might be instantaneously lethal. When CO combines with hemoglobin, the resulting chemical is called carboxy hemoglobin (COHb).

The negative effect of CO on the delivery of oxygen to the tissues extends beyond just the simple blockage of oxygen-combining sites. Each hemoglobin molecule contains four oxygen-carrying sites. Once the first oxygen molecule is released at the tissue level the second, third, and fourth come off even more rapidly. Oxygen release is delayed by CO so that there is even less oxygen delivered than would be expected purely on the basis of the amount of oxygen not being carried by hemoglobin. For this reason, overt symptoms due to lack of oxygen can be observed at COHb levels of approximately 15 to 20 percent, or even less, in healthy people. Levels of COHb over 40 percent can be lethal.

The uptake of CO increases as respiratory rates increase. This puts children at greater risk since they breathe more rapidly, in proportion to their body weight, than adults. This explains the unfortunate situation of a family in an automobile stuck in a snowstorm with the motor running being found with the adults unconscious and the children dead. The fetus is also at higher risk due to the greater affinity of CO for fetal, as compared to adult, hemoglobin.

All cases of fatal CO poisoning are readily preventable. In addition to automobile exhaust, other lethal sources of CO are often related to home heating systems. Blockage of flues, or inappropriate repair work on the home heating source or on ducts, is often responsible for CO toxicity. Symptoms of CO toxicity, such as headache, weakness, and listlessness, tend to be worse in the morning and to go away during the day if people leave the home. Many fatal cases are preceded by visits to physicians or emergency departments with only symptomatic treatment. Home CO alarms are relatively cheap and are an effective means of prevention. CO poisoning occurs more rapidly at high altitude due to the relative lack of oxygen to compete for the oxygen-combining site of hemoglobin. Conversely, symptomatic CO poisoning is treated with oxygen.

CO is also made in the human body through the normal catabolism of heme (oxygen-carrying hemoglobin), which leads to a background concentration in the blood of approximately 0.5 percent COHb. Concentrations of 2 to 3 percent COHb have been associated with an increased risk of angina attacks in susceptible individuals with preexisting arteriosclerotic heart disease. Preventing this adverse consequence is the major basis for the current U.S. ambient standard for CO. There has been a significant decline in outdoor CO levels in the United States as a result of decreased automotive emissions of carbon monoxide.

Bernard D. Goldstein

(see also: Ambient Air Quality [Air Pollution] )

Bibliography

Ernst, A., and Zibrak, J. D. (1998). "Carbon Monoxide Poisoning." New England Journal Medicine 339:16031608.

Raub, J. A.; Mathieu-Nof, M.; Hampson, M. B.; and Thom, S. R. (2000). "Carbon Monoxide Poisoning A Public Health Perspective." Toxicology 145(1):114.

Tomaszewski, C. (1999). "Carbon Monoxide Poisoning Early Awareness and Intervention Can Save Lives." Postgrad Medicine 105(1):3940.

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Carbon Monoxide

Carbon Monoxide


Carbon monoxide is an invisible, odorless, and poisonous gas with the chemical formula CO. Because of its toxicity, the U.S. Environmental Protection Agency (EPA) regulates CO. The gas is a by-product of incomplete combustion (burning with insufficient oxygen). Its major source is vehicle exhaust (60 percent). Other sources include water heaters and furnaces, gas-powered engines (boats and lawn mowers), charcoal and wood fires, agricultural burning, and tobacco smoke.

CO is classified as an indirect greenhouse gas. It does not contribute to global warming directly, but leads to the formation of ozone. Ozone is the major air pollutant formed in photochemical smog and a potent greenhouse gas.

Human exposure to elevated CO impairs oxygen uptake in the bloodstream. Under CO-free conditions, oxygen is transported from the lungs to tissues by hemoglobin. When CO is present, it mimics the shape of oxygen and binds instead to the hemoglobin. The molecule is not easily released, blocking further oxygen uptake, and ultimately depriving organs and tissues of life-sustaining oxygen. The symptoms of CO poisoning range from dizziness, mild headaches, and nausea at lower levels to severe headaches, seizures, and death at higher levels.

The EPA national outdoor air quality standard for CO is nine parts per million or ppm (0.0009 percent) averaged over an eight-hour period. The gas is life-threatening after three hours at 400 ppm (0.04 percent) and within minutes at 1.28 percent. In 1996, 525 deaths in the United States were attributed to unintentional and 1,988 deaths to intentional CO poisoning.

Exposure to CO can be reduced by assuring adequate ventilation when near any combustion source. Indoor cooking with charcoal and running gaspowered engines inside a garage are both dangerous and should be avoided. Fuel-burning appliances and fireplaces ought to be routinely inspected.

CO detectors are available to detect less obvious sources, such as a malfunctioning furnace. The sensors operate in one of three ways: They mimic the body's response to CO (biomimetic detectors), they allow a heated metal oxide to react with the gas (metal oxide detectors), or they facilitate a reaction using platinum electrodes immersed in an electrolyte solution (electrochemical detectors). The lowest level that a CO alarm can detect is 70 ppm.

see also Air Pollution; Global Warming; Greenhouse Gases; Indoor Air Pollution; Ozone; Vehicular Pollution.

Bibliography

Phillips, William G. (1998). "Carbon Monoxide Detectors, What You Need to Know." Popular Science 252(1):7681.

Turco, Richard P. (1997). Earth under Siege: From Air Pollution to Global Change. New York: Oxford University Press.


Internet Resource

U.S. Environmental Protection Agency Web site. Available from http://epa.gov/IAQ.

Marin Sands Robinson

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carbon monoxide

carbon monoxide, chemical compound, CO, a colorless, odorless, tasteless, extremely poisonous gas that is less dense than air under ordinary conditions. It is very slightly soluble in water and burns in air with a characteristic blue flame, producing carbon dioxide; it is a component of producer gas and water gas, which are widely used artificial fuels. At high pressures and elevated temperatures it reacts with hydrogen in the presence of a catalyst to form methanol, and under similar conditions reacts with methanol to produce acetic acid. It is also used in the production of polycarbonate and polyurethane as well as detergents. As a reducing agent, removing oxygen from many compounds, it is used in the reduction of metals, e.g., iron (see blast furnace), from their ores.

Carbon monoxide is formed by combustion of carbon in oxygen at high temperatures when there is an excess of carbon. It is also formed (with oxygen) by decomposition of carbon dioxide at very high temperatures (above 2,000°C). It is present in the exhaust of internal-combustion engines (e.g., in automobiles) and is generated in coal stoves, furnaces, and gas appliances that do not get enough air (because of a faulty draft or for other reasons).

Carbon monoxide is an extremely poisonous gas. Breathing air that contains as little as 0.1% carbon monoxide by volume can be fatal; a concentration of about 1% can cause death within a few minutes. The gas is especially dangerous because it is not easily detected by human senses. Early symptoms of carbon monoxide poisoning include drowsiness and headache, followed by unconsciousness, respiratory failure, and death. First aid for a victim of carbon monoxide poisoning requires access to fresh air; administration of artificial respiration and, if available, oxygen; and, as soon as possible, expert medical attention. When carbon monoxide is inhaled, it reacts with hemoglobin, the red blood pigment that normally carries oxygen to all parts of the body. Because carbon monoxide is attracted to the hemoglobin about 210 times as strongly as is oxygen, it takes the place of oxygen in the blood, causing oxygen starvation throughout the body. Carbon monoxide detectors for homes are now readily available.

Carbon monoxide from automobile and industrial emissions is a dangerous pollutant that may contribute to the greenhouse effect and global warming. In urban areas carbon monoxide, along with aldehydes, react 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 (see smog). Carbon monoxide is also a constituent of tobacco smoke.

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carbon monoxide

carbon monoxide Colourless, odourless poisonous gas (CO) formed during the incomplete combustion of fossil fuels, occurring for example in coal gas and the exhaust fumes of cars. Carbon monoxide poisons by combining with the haemoglobin in red blood cells and thus preventing them from carrying oxygen around the body. (This happens if the inhaled air contains only 0.1% of carbon monoxide by volume.) It is used as a reducing agent in metallurgy. Properties: density 0.968 (air = 1); m.p. −205°C (−337°F); b.p. −191.5°C (−312.7°F).

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carbon monoxide

carbon monoxide A colourless odourless gas, CO. It is formed by the incomplete combustion of carbon and is present in car-exhaust gases. Carbon monoxide is able to bond with metals; this accounts for its toxicity, which is due to the binding of the CO to the iron in haemoglobin, thereby blocking the uptake of oxygen. See carboxyhaemoglobin.

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carbon monoxide

carbon monoxide (mŏn-ok-syd) n. a colourless almost odourless gas that is very poisonous. When breathed in it combines with haemoglobin in the red blood cells (see carboxyhaemoglobin). Carbon monoxide is present in coal gas and motor exhaust fumes. Formula: CO.

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carbon monoxide

car·bon mon·ox·ide • n. a colorless, odorless, toxic flammable gas, CO, formed by incomplete combustion of carbon.

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Carbon monoxide

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.

History

The discovery of carbon monoxide is often credited to the work of the English chemist and theologian Joseph Priestley (17331804). 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 Priestleys work. As early as the late 1200s, Spanish alchemist Arnold of Villanova (c.1238c.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 (15801644) 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 (18321913) 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 (18541941).

Sources

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 Earths 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.

Physiological effects

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 poisoningat least at moderate levelsis 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.

Uses

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 122176°F (5080°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.

Resources

BOOKS

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

Emsley, John. Natures 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

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Carbon Monoxide

Carbon monoxide

A colorless, odorless, tasteless gas that is produced in only very small amounts by natural processes. By far the most important source of the gas is the incomplete combustion of coal , oil, and natural gas . In terms of volume, carbon monoxide is the most important single component of air pollution . Environmental scientists rank it behind sulfur oxides, particulate matter, nitrogen oxides , and volatile organic compounds, however, in terms of its relative hazard to human health. In low doses, carbon monoxide causes headaches, nausea, fatigue, and impairment of judgment. In larger amounts, it causes unconsciousness and death.

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Carbon Monoxide

Carbon monoxide

Carbon monoxide is a compound of carbon and oxygen with the chemical formula CO. It is a colorless, odorless, tasteless, toxic gas. 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).


History

The discovery of carbon monoxide is often credited to the work of the English chemist and theologian Joseph Priestley. 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. None the less carbon monoxide had been well known and extensively studied in the centuries prior to Priestley's work. As early as the late 1200s, the Spanish alchemist Arnold of Villanova 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. Johann (or Jan) Baptista van Helmont 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 the French physicist Louis Paul Cailletet 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.


Sources

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 .

Physiological effects

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 final result. The U.S. Occupational Safety and Health Administration 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.

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 we have a fairly sophisticated understanding 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. Oxyhemoglobin is a fairly 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.


Uses

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 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.

See also Metallurgy.

David E. Newton

Resources

books

Boikess, Robert S., and Edward Edelson. Chemical Principles. 2nd edition. New York: Harper & Row Publishers, 1981, pp. 672 - 673.

Brown, Theodore L., and H. Eugene LeMay Jr. Chemistry: The Central Science. 3rd edition. Englewood Cliffs, NJ: Prentice-Hall, 1985, pp. 390-392, 66 -669.

Budavari, Susan, ed. The Merck Index. 11th edition. Rahway, NJ: Merck and Company, 1989, pp. 1821.

Greenwood, N. N., and A. Earnshaw. Chemistry of the Elements. Oxford: Butterworth-Heinneman Press, 1997.

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 edition. London: Macmillan & Company, 1957, pp. 49, 116, 142, 151.


other

The United Nations. "The Conference and Kyoto Protocol." (March 2003). <http://unfccc.int/resource/convkp.html>.

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 which reduces materials by donating electrons to them.

Toxicity

—The extent to which a substance is poisonous.

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Carbon Monoxide

Carbon Monoxide


Carbon monoxide is an odorless, tasteless, colorless, and poisonous gas. Although most of the carbon monoxide in the atmosphere comes from natural sources, a great deal is also added by the burning of fossil fuels by automobiles and industry. It is an extremely poisonous gas if inhaled, since it kills by preventing oxygen from reaching the cells.

Carbon monoxide (CO) is a compound consisting of one carbon atom (the smallest unit of an element) and one oxygen atom. It is normally in the air people breathe, although in extremely small amounts. It is produced naturally when a substance that contains carbon decays or breaks down in the absence of oxygen. This happens in swamps where there is little oxygen. Carbon monoxide is also produced in greater quantity when carbon-containing substances are burned without enough oxygen being present. This often happens when gas furnaces, wood stoves, and space heaters malfunction or are not properly vented. Automobile engines and other gasoline or diesel-powered motors also generate carbon monoxide. The exhaust from these motors can be deadly if they are operated in enclosed areas or attached garages.

Carbon monoxide can kill a person who breathes it. It does this by preventing the blood from being able to carry oxygen. Without oxygen, people and animals soon die. Once inhaled, carbon monoxide combines with the hemoglobin (the oxygen-carrying substance in the blood) to the exclusion of oxygen. In fact, carbon monoxide combines with hemoglobin two hundred times faster than oxygen does. Additionally, the hemoglobin does not release the carbon monoxide as it does the oxygen. Thus, as more and more red blood cells pick up carbon monoxide, the total number available to deliver oxygen to the cells keeps decreasing, and soon the person slowly falls into a sleeplike state. Deprived of oxygen, the brain begins to slow its functions. Eventually, bodily functions stop, and the person dies.

This oxygen deprivation does not always kill immediately. Low-level exposure can cause flu-like symptoms including shortness of breath, mild headaches, fatigue, and nausea. Higher levels may cause dizziness, severe headaches, mental confusion, and fainting. Prolonged exposure can cause death. According to the U. S. Consumer Product Safety Commission, more than 2,500 people will die and 100,000 will be seriously injured by carbon monoxide poisoning over the next 10 years.

People who smoke also run the risk of harming themselves with this toxic gas. The carbon monoxide present in cigarette smoke not only excludes oxygen from binding to hemoglobin, but it also prevents it from picking up carbon dioxide (which is the waste product of breathing). This means that the person's heart has to pump harder to try and rid the body of carbon dioxide wastes.

Large cities had an especially bad problem in the past as unhealthy amounts of carbon monoxide would build up during rush hour because of automobile exhaust. Newer cars are now equipped with catalytic converters that chemically change carbon monoxide into harmless carbon dioxide. Also, in the last decade or so, carbon monoxide detectors for the

home have become practical, and many cities are now requiring that at least one carbon monoxide detector be installed in every home, apartment, and hotel. Since carbon monoxide is odorless, tasteless, and colorless, these monitors are invaluable in alerting people to possible exposure.

[See alsoCarbon Family ]

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Carbon Monoxide

Carbon Monoxide

OVERVIEW

Carbon monoxide is a colorless, odorless, tasteless, toxic gas. It is one of the most common poisons in the environment and is responsible for thousands of deaths and hospital emergency room visits each year in the United States. Carbon monoxide is produced from fuel-burning appliances, such as space heaters, furnaces, stoves, and vehicles. It is also a component of cigarette smoke. Carbon monoxide is flammable and capable of forming an explosive mixture with air.

KEY FACTS

OTHER NAMES:

None

FORMULA:

CO

ELEMENTS:

Carbon, oxygen

COMPOUND TYPE:

Nonmetallic oxide

STATE:

Gas

MOLECULAR WEIGHT:

28.01 g/mol

MELTING POINT:

−205.02°C (−337.04°F)

BOILING POINT:

−191.5°C (−312.7°F)

SOLUBILITY:

Slightly soluble in water; soluble in alcohol and chloroform

Under ideal circumstances, a carbon-containing fuel such as charcoal, natural gas, or wood burns in air to form carbon dioxide and water. Lacking an adequate supply of oxygen, at low temperatures, or under certain other conditions, incomplete combustion occurs. Incomplete combustion is a process in which a fuel is not completed oxidized. In such instances, carbon monoxide is produced along with carbon dioxide and water. An inadequate supply of oxygen can occur in enclosed, poorly vented spaces, such as the interior of houses, garages, and cars.

The discovery of carbon monoxide is usually credited to the English chemist Joseph Priestley (1733–1804). Between 1772 and 1799, Priestley investigated the properties of carbon monoxide and recognized the difference between carbon monoxide and carbon dioxide. The gas was first prepared synthetically by the French chemist Joseph Marie François de Lassone (1717–1788) in 1776, although he mistakenly identified it as hydrogen. The correct chemical formula for carbon monoxide was first identified by the English chemist William Cruikshank (1745–1800) in 1800.

HOW IT IS MADE

A number of methods are available for the commercial production of carbon monoxide. In one procedure, air is passed over hot coke, graphite, or anthracite coal to make producer gas, a mixture of carbon dioxide, carbon monoxide, hydrogen, and water vapor. In a similar procedure, steam is passed over hot coke or graphite to make water gas, a mixture of carbon monoxide, carbon dioxide, hydrogen, and nitrogen. In a third procedure, steam is mixed with natural gas to form synthesis gas, consisting of carbon dioxide, carbon monoxide, and hydrogen. In all three procedures, the carbon monoxide component of the gas mixture produced in the reaction can be separated out from the other gases.

Yet a fourth method for making carbon monoxide involves the partial oxidation of hydrocarbon gases obtained from natural gas or petroleum. These gases consist of carbon-hydrogen compounds which, when oxidized, are converted to carbon monoxide, carbon dioxide, and water. By controlling the amount of reactants used and the conditions of the reaction (temperature and pressure), the portion of carbon monoxide produced can be increased.

COMMON USES AND POTENTIAL HAZARDS

Carbon monoxide has three major industrial uses. The first is in the synthesis of a large variety of organic compounds. For example, it takes part in a group of reactions known as the Fischer-Tropsch reactions in which carbon monoxide is first reduced with hydrogen gas and then converted to any number of organic compounds that contain oxygen. The gas is also used to make acetic acid, a major industrial chemical used in the synthesis of polymers and other organic products.

Interesting Facts

  • Some authorities believe that the American writer Edgar Allan Poe (1809–1849) owed his wild imagination and early death to chronic carbon monoxide poisoning, caused by the gas lighting used in his home.
  • Public health authorities estimate that the number of suicides resulting from the inhalation of carbon monoxide fumes from automobiles has decreased by about 80 percent since the introduction of catalytic converters in cars. Catalytic converters reduce the amount of carbon monoxide released in a car's exhaust.
  • Carbon monoxide bonds to hemoglobin in the blood stream 200 times as efficiently as does oxygen. This ability to exclude oxygen from the blood is responsible for the toxic effects caused by carbon monoxide in the body.

A second application of carbon monoxide, either by itself or in conjunction with other gases, is as an industrial fuel. The gas burns very efficiently with the release of large amounts of heat and relatively few undesirable by-products (the most important being carbon dioxide).

A third industrial use for carbon monoxide is in the refining of metals. Most metal ores exist in the form of oxides or sulfides when extracted from the earth. For example, the two most important ores of iron are magnetite (Fe3O4) and hematite (Fe2O3). After an ore has been mined, it is treated to remove the oxygen or sulfur in the ore to obtain a pure metal. Carbon monoxide is often used for this purpose with oxide ores because it combines with oxygen from the ore to form carbon dioxide, leaving the metal behind: Metal oxide + CO → Metal + CO2.

Carbon monoxide is a highly toxic gas that has noticeable health effects even in relatively small concentrations. When the concentration of carbon monoxide reaches levels of about 100 ppm (parts per million), an individual is likely to experience mild headaches, fatigue, shortness of breath, and errors in judgment. As the concentration of carbon monoxide increases, these symptoms become more pronounced. Exposure to a concentration of more than 400 ppm for more than three hours is likely to put a person at serious health risk. He or she may begin to lose consciousness and experience serious disorientation. At concentrations of more than 1,500 ppm, death is likely in less than an hour. These symptoms vary somewhat depending on a person's age and overall health.

Words to Know

POLYMER
A compound consisting of very large molecules made of one or two small repeated units called monomers.
REDUCTION
Chemical reaction in which oxygen is removed from a substance or electrons are added to a substance.
SYNTHESIS
A chemical reaction in which some desired chemical product is made from simple beginning chemicals, or reactants.
SYNTHESIS GAS
A mixture of several gases, such as carbon dioxide and hydrogen, used to produce compounds such as methanol and ammonia that can be separated out from the synthesis gas.

Most governmental agencies have set a recommended limit of 35 ppm for periods of up to eight hours. For purposes of comparison, the normal concentration of carbon monoxide in the atmosphere in an open area tends to be less than 1 ppm. But in urban areas or other locations with many vehicles, gas heaters, wood burning stoves, or other sources of carbon monoxide, carbon monoxide levels can be much higher. Someone traveling inside a car on a busy freeway, for example, may be exposed to carbon monoxide levels of up to 25 ppm. Concentrations of up to 100 ppm have been measured in the center of busy downtown urban areas.

FOR FURTHER INFORMATION

"Carbon Monoxide." Chemical Fact Sheet. Plastics and Chemicals Industries Association. http://www.pacia.org.au/_uploaditems/docs/3.carbon_monoxide.pdf (accessed on October 2, 2005).

"Carbon Monoxide Questions and Answers." Consumer Product Safety Commission. CPSC Document #466. http://www.cpsc.gov/cpscpub/pubs/466.html (accessed on October 3, 2005).

Dwyer, Bob, et al. Carbon Monoxide: A Clear and Present Danger. Mount Prospect, Ill.: ESCO Press, 2004.

Patnaik, Pradyot. Handbook of Inorganic Chemicals. New York: McGraw-Hill, 2003, 1887–191.

See AlsoAcetic Acid; Carbon Dioxide

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