Carbon dioxide

General
Name	Carbon dioxide
Chemical Formula	CO2
Appearance	Colourless gas
Physical
Formula weight	44.0 amu
Melting point	Liquifies under high pressure at 216 K (−57 °C)
Boiling point	sublimes at 195 K (−78 °C)
Density	1.6 Mg/m3 (solid) 1.98 kg/m3 (gas at 298 K)
Solubility	0.145 g in 100 g water
Thermochemistry
ΔfH0gas	−393.52 kJ/mol
ΔfH0solid	? kJ/mol
S0gas, 100 kPa	213.79 J/(mol·K)
S0solid	? J/(mol·K)
Safety
Ingestion	May cause nausea, vomiting, GI hemorrhage.
Inhalation	Asphyxiant (suffocating), causes hyperventilation. Repeated exposure dangerous.
Skin	Dry ice may cause frostbite.
Eyes	Can be dangerous.
More info	Hazardous Chemical Database
SI units were used where possible. Unless otherwise stated, standard conditions were used. Disclaimer and references

Carbon dioxide is an atmospheric gas composed of one carbon and two oxygen atoms. One of the best known of chemical compounds, it is frequently called by its formula:

CO₂ (pronounced "see oh two")

Carbon dioxide results from the combustion of organic matter if sufficient amounts of oxygen are present. It is also produced by various microorganisms from fermentation and cellular respiration. Plants utilize carbon dioxide during photosynthesis, using both the carbon and the oxygen to construct carbohydrates. In addition, plants also release oxygen to the atmosphere which is subsequently used for respiration by heterotrophic organisms, forming a cycle. It is present in the Earth's atmosphere at a low concentration and acts as a greenhouse gas. It is a major component of the carbon cycle.

Chemical and physical properties

Carbon dioxide is a colorless gas which, when inhaled at high concentrations (a dangerous activity due to the associated asphyxiation risk), produces a sour taste in the mouth and stinging sensation in the nose and throat. These effects result from the gas dissolving in the mucous membranes and saliva, forming a weak solution of carbonic acid.

Its density at 298 K is 1.98 kg m⁻³, about 1.5 times that of air. The carbon dioxide molecule (O=C=O) contains two double bonds and has a linear shape. It has no electrical dipole. As it is fully oxidized, it is not very reactive and in particular not flammable.

At temperatures below –78°C, carbon dioxide condenses into a white solid called dry ice. Liquid carbon dioxide forms only at pressures above 5.1 atm; at atmospheric pressure, it passes directly between the gaseous and solid phases at in a process called sublimation.

Water will absorb its own volume of carbon dioxide, and more than this under pressure. About 1% of the dissolved carbon dioxide turns into carbonic acid. The carbonic acid in turn dissociates partly to form bicarbonate and carbonate ions.

Uses

Liquid and solid carbon dioxide are important refrigerants, especially in the food industry, where it is employed during the transportation and storage of ice cream and other frozen foods.

Carbon dioxide is used to produce carbonated beverages such as soft drinks and soda water. Traditionally, the carbonation in beer and sparkling wine comes about through natural fermentation, but some manufacturers carbonate these beverages artificially.

The leavening agents used in baking produce carbon dioxide to cause dough to rise. Baker's yeast produces carbon dioxide by fermentation within the dough, while chemical leaveners such as baking powder and baking soda release carbon dioxide when heated or exposed to acids.

Carbon dioxide is often used as an inexpensive, non-flammable pressurized gas. Life jackets often contain canisters of pressured carbon dioxide for quick inflation. Steel capsules are also sold as supplies of compressed gas for airguns and paintball markers. Rapid vaporization of liquid CO₂ is used for blasting in coal mines.

Carbon dioxide extinguishes flames, and some fire extinguishers, especially those designed for electrical fires, contain liquid carbon dioxide under pressure. Carbon dioxide also finds use as an atmosphere for welding, although in the welding arc, it reacts to oxidize most metals. Use in the automotive industry is common despite significant evidence that welds made in carbon dioxide are brittler than those made in more inert atmospheres, and that such weld joints deteriorate over time due to the formation of carbonic acid. It is used as a welding gas primarily because it is much less expensive than more inert gases such as argon or helium.

Liquid carbon dioxide is a good solvent for many organic compounds. It has begun to attract attention in the pharmaceutical and other chemical processing industries as a less toxic alternative to more traditional solvents such as organochlorides. (See green chemistry.)

Plants require carbon dioxide to conduct photosynthesis, and greenhouses may enrich their atmospheres with additional CO₂ to boost plant growth. It has been proposed that carbon dioxide from power generation be bubbled into ponds to grow algae that could then be converted into biodiesel fuel. High level of carbon dioxide in the atmosphere effectively exterminate many pests. Greenhouses will raise the level of CO₂ to 10,000 ppm (1%) for several hours to eliminate pests such as whitefly, spider mites, and others.

In the theater, dry ice is used to produce fog as a special effect: when dry ice added to water, the evaporating mixture of CO₂ and cold humid air condenses as a fog.

Dry ice is also used in cleaning: shooting tiny dry ice pellets at a surface cools the dirt and causes it to pop off.

A common type of industrial laser uses carbon dioxide as a medium.

Biology

Carbon dioxide is a waste product in organisms that obtain energy from breaking down sugars or fats with oxygen as part of their metabolism, in a process known as cellular respiration. This includes all plants, animals, many fungi and some bacteria. In higher animals, the carbon dioxide travels in the blood (where most of it is held in solution) from the body's tissues to the lungs where it is exhaled.

Carbon dioxide content in fresh air is less than 1%, in exhaled air ca. 4.5%. When inhaled in high concentrations (about 5% by volume), it is toxic to humans and other animals. Hemoglobin, the main molecule in red blood cells, can bind both to oxygen and to carbon dioxide. If the CO₂ concentration is too high, then all hemoglobin is saturated with carbon dioxide and no oxygen transport takes place (even if plenty of oxygen is in the air). As a result, people in a poorly ventilated room will experience difficulty breathing due to accumulated carbon dioxide, even before lack of oxygen becomes a problem. Carbon dioxide, either as a gas or as dry ice, should be handled only in well ventilated areas.

OSHA limits carbon dioxide concentration in the workplace to 0.5% for prolonged periods, or to 3% for brief exposures (up to ten minutes). OSHA considers concentrations exceeding 4% as "immediately dangerous to life and health." People who breathe 5% carbon dioxide for more than half an hour show signs of acute hypercapnia, while breathing 7–10% carbon dioxide can produce unconsciousness in only a few minutes.

According to a study by the USDA [1], an average person's repiration generates approximately 450 liters (roughly 900 grams) of carbon dioxide per day.

The CO₂ that is carried in blood can be found in different areas. 8% of CO₂ is in the plasma as a gas. 20% of it is bound to hemoglobin. The CO₂ bounded to hemoglobin is not competing with oxygen binding since it binds to amino acids rather than heme molecules. The remaining 72% of it is carried as bicarbonate HCO₃⁻ which is a buffer important in our pH regulation. The level of bicarbonate is regulated and if it is high then we breath more rapidly to get rid of the excess carbon dioxide. The level of carbon dioxide/bicarbonate in the blood affects the thickness of the blood capillaries. If it is high, the capillaries expand and more blood rushes in and carries the excess bicarbonate to the lungs. To help avoid the loss of carbon dioxide to a deadly low level, the body has developed certain defensive mechanisms. These include contractions of the air pipes and blood pipes, and the increased production of mucus.

Plants remove carbon dioxide from the atmosphere by photosynthesis, which uses light energy to produce organic plant materials by combining carbon dioxide and water. This releases free oxygen gas. Sometimes carbon dioxide gas is pumped into greenhouses to promote plant growth. Plants also emit CO₂ during respiration; but on balance they are net sinks of CO₂.

Atmosphere

Atmospheric CO₂ concentrations, measured at Mauna Loa.

As of 2004, the earth's atmosphere is about 0.038% by volume (380 µL/L or ppmv) CO₂. Due to the greater land area, and therefore greater plant life, in the northern hemisphere as compared to the southern hemisphere, there is an annual fluctuation of about 5 µL/L, peaking in May and reaching a minimum in October at the end of the northern hemisphere growing season, when the quantity of biomass on the planet is greatest.

Despite its small concentration, CO₂ is a very important component of Earth's atmosphere, because it absorbs infrared radiation and enhances the greenhouse effect.

The initial carbon dioxide in the atmosphere of the young Earth was produced by volcanic activity; this was essential for a warm and stable climate conducive to life. Volcanic activity now releases about 130 to 230 teragrams (145 million to 255 million short tons) of carbon dioxide each year. Volcanic releases are about 1% the amount which is released by human activities.

Since the start of the Industrial Revolution, the atmospheric CO₂ concentration has increased by approximately 110 µL/L or about 40%.

The actual increase in overall atmospheric CO₂ is 110 ppvm - which, by itself, represents 0.01% of the total volume of Earth's atmosphere.

Some environmentalists object to attention being drawn to the proportion of the increase because they feel it it an attempt to marginalize the increase. They prefer to point to the 40% figure, possibly because it is larger and seems much more dramatic than 0.01%.

Monthly measurements taken at Mauna Loa [2] since 1959 show an increase from 316 µL/L in that year to 376 µL/L in 2003, an overall increase of 60 µL/L during the 44-year history of the measurements. Burning fossil fuels such as coal and petroleum is the leading cause of increased man-made CO₂; deforestation the second major cause. Various techniques have been proposed for removing excess carbon dioxide from the atmosphere in carbon dioxide sinks.

The Global Warming Theory (GWT) predicts that increased amounts of CO₂ in the atmosphere tend to enhance the greenhouse effect and thus contribute to global warming.

Variation in the past

Graph of Vostok CO₂ (black)/T (blue)/Dust (red) over the last 400 kyr. The current value is 380 µL/L

The most direct method for measuring atmospheric carbon dioxide concentrations in the distant past is to measure its concentration in bubbles of air (fluid or gas inclusions) trapped in the Antarctic or Greenland ice caps. The most extensive such study has been undertaken at Vostok, Antarctica, where ice has been sampled to a depth of 3,600 meters, corresponding to an age of 420,000 years before the present. [3] During this time, the atmospheric carbon dioxide concentration has varied between 180–210 µL/L during ice ages, increasing to 280–300 µL/L during warmer interglacials. In 2004, a European team announced the recovery of a 740,000-year ice core, but carbon dioxide levels from this core have not yet been published.

Various proxy measurements have been used to attempt to determine atmospheric carbon dioxide levels millions of years in the past. These include boron and carbon isotope ratios in certain types of marine sediments, and the number of stomata observed on fossil plant leaves. While these measurements give much less precise estimates of carbon dioxide concentration than ice cores, there is evidence for very high CO₂ concentrations (>3,000 µL/L) between 600 and 400 Myr BP and between 200 and 150 Myr BP.[4] On long time-scales, atmospheric CO₂ content is determined by the balance among geochemical processes including organic carbon burial in sediments, silicate rock weathering, and vulcanism. The net effect of slight imbalances in the carbon cycle over tens to hundreds of millions of years has been to reduce atmospheric CO₂. The rates of these processes are extremely slow, hence they are of limited relevance to the atmospheric CO₂ response to emissions over the next hundred years. In more recent times, atmospheric CO₂ concentration continued to fall after about 60 Myr BP and there is geochemical evidence that concentrations were <300 µL/L by about 20 Myr BP. Low CO₂ concentrations may have been the stimulus that favoured the evolution of C4 plants, which increased greatly in abundance between 7 and 5 Myr BP. Although contemporary CO₂ concentrations were exceeded during earlier geological epochs, they are likely higher now than at any time during the past 20 million years [5]. A comparison of the world's average temperature and and the associated CO₂ levels throughout the past 600 million years is presented graphically here.

Oceans

File:Liquid carbon dioxide on the bottom of the ocean.jpg

An image of liquid carbon dioxide spilling from a beaker on the ocean floor. The image was taken from video of a carbon sequestration experiment done by the Monterey Bay Aquarium Research Institute.

The Earth's oceans contain a huge amount of carbon dioxide in the form of bicarbonate and carbonate ions--much more than the amount in the atmosphere. The bicarbonate is produced in reactions between rock, water, and carbon dioxide. One example is the dissolution of calcium carbonate:

CaCO₃ + CO₂ + H₂O <--> Ca²⁺ + 2 HCO₃^-

Reactions like this tend to buffer changes in atmospheric CO₂. Reactions between carbon dioxide and non-carbonate rocks also add bicarbonate to the seas, which can later undergo the reverse of the above reaction to form carbonate rocks, releasing half of the bicarbonate as CO₂. Over hundreds of millions of years this has produced huge quantities of carbonate rocks. If all the carbonate rocks in the earth's crust were to be converted back into carbon dioxide, the resulting carbon dioxide would weigh 40 times as much as the rest of the atmosphere.

The vast majority of CO₂ added to the atmosphere will eventually be absorbed by the oceans and become bicarbonate ion, but the process takes on the order of a hundred years because most seawater rarely comes near the surface.

History

Carbon dioxide was one of the first gases to be described as a substance distinct from air. In the 17th century, the Flemish chemist Jan Baptist van Helmont observed that when he burned charcoal in a closed vessel, the mass of the resulting ash was much less than that of the original charcoal. His interpretation was that the rest of the charcoal had been transmuted into an invisible substance he termed a "gas" or "wild spirit" (spiritus sylvestre).

Carbon dioxide's properties were studied more thoroughly in the 1750s by the Scottish physician Joseph Black. He found that limestone (calcium carbonate) could be heated or treated with acids to yield a gas he termed "fixed air." He observed that the fixed air was denser than air and did not support either flame or animal life. He also found that it would, when bubbled through an aqueous solution of lime (calcium hydroxide), precipitate calcium carbonate, and used this phenomenon to demonstrate that carbon dioxide is produced by animal respiration and microbial fermentation.

External links

• Dry Ice information
• Bassam Z. Shakhashiri: Chemical of the Week: Carbon Dioxide
• Keeling, C.D. and T.P. Whorf: Atmospheric carbon dioxide record from Mauna Loa, 2002
• Mauna Loa 2004 update