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Air

Air

Air is a name for the mixture of gases present in the Earth's atmosphere. Compressed air is often used in scuba diving as a shallow water breathing gas and to inflate buoyancy devices. Compressed air is also used as the means of transmission of energy to pneumatic tools.

Composition of air

By volume, air is about:
- 78.084% Nitrogen (N₂)
- 20.947% Oxygen (O₂)
- 0.934% Argon (Ar)
- 0.033% Carbon Dioxide (CO₂) With trace amounts of:
- Neon (Ne)
- Helium (He)
- Krypton (Kr)
- Sulfur dioxide (SO₂)
- Methane (CH₄)
- Hydrogen (H₂)
- Nitrous Oxide (N₂O)
- Xenon (Xe)
- Ozone (O₃)
- Nitrogen dioxide (NO₂)
- Iodine (I₂)
- Carbon monoxide (CO)
- Ammonia (NH₃) The amount of water vapor in the air varies considerably depending on weather, climate, and altitude. See Humidity. The molecular mass of air is approximately 28.96443 g/mole (molecular weight of standard air - CRC, 1983).

External link

- [http://mistupid.com/chemistry/aircomp.htm Composition of Air] Category:Atmosphere Category:Psychrometrics Category:HVAC ko:대기 ms:Udara ja:空気 simple:Air

Gas

:For other meanings see gas (disambiguation). ---- A gas is one of the four main phases of matter (after solid and liquid, and followed by plasma), that subsequently appear as a solid material is subjected to increasingly higher temperatures. Thus, as energy in the form of heat is added, a solid (e.g. ice) will first melt to become a liquid (e.g. water), which will then boil or evaporate to become a gas (e.g. water vapor). In some circumstances, a solid (e.g. "dry ice") can directly turn into a gas: this is called sublimation. If the gas is further heated, its atoms or molecules can become (wholly or partially) ionized, turning the gas into a plasma.

Properties of a gas

#All collisions are perfectly elastic #The gas fills the entire container #The molecules have negligible volume In the gas phase, the atoms or molecules constituting the matter basically move independently, with no forces keeping them together or pushing them apart. Their only interactions are rare and random collisions. The particles move in random directions, at high speeds, whose range is dependent on the temperature and defined by the Maxwell-Boltzmann distribution. Therefore, the gas phase is a completely disordered state. Following the second law of thermodynamics, gas particles will immediately diffuse to homogeneously fill any shape or volume of space that is made available to them. The thermodynamic state of a gas is characterized by its volume, its temperature, which is determined by the average velocity or kinetic energy of the molecules, and its pressure, which is determined by the average velocity and density or number of molecules. These variables are related by the fundamental gas laws, which state that the pressure in an ideal gas is proportional to its temperature and number of molecules, but inversely proportional to its volume. Like liquids and plasmas, gases are fluids: they have the ability to flow and do not tend to return to their former configuration after deformation, although they do have viscosity. Unlike liquids, however, unconstrained gases do not occupy a fixed volume, but expand to fill whatever space they occupy. The kinetic energy per molecule in a gas is the second greatest of the states of matter (after plasma). Because of this high kinetic energy, gas atoms and molecules tend to bounce off of any containing surface and off one another, the more powerfully as the kinetic energy is increased. A common misconception is that the collisions of the molecules with each other is essential to explain gas pressure, but in fact their random velocities are sufficient to define that quantity. Mutual collisions are important only for establishing the Maxwell-Boltzmann distribution. Gas particles are normally well separated, as opposed to liquid particles, which are in contact. A material particle (say a dust mote) in a gas moves in Brownian Motion. Since it is at the limit of (or beyond) current technology to observe individual gas particles (atoms or molecules), only theoretical calculations give suggestions as to how they move, but their motion is different from Brownian Motion. The reason is that Brownian Motion involves a smooth drag due to the frictional force of many gas molecules, punctuated by violent collisions of an individual (or several) gas molecule(s) with the particle. The particle (generally consisting of millions or billions of atoms) thus moves in a jagged course, yet not so jagged as we would expect to find if we could examine an individual gas molecule.

Etymology

The word "gas" was apparently coined in the early 17th century by the Belgian chemist Jan Baptist van Helmont, as a re-spelling of his pronunciation of the Greek word chaos.

Scuba diving

Scuba diving is the use of independent breathing equipment to stay underwater for long periods of time for recreational diving and professional diving. Generally the diver swims underwater, but walking and the use of diver propulsion vehicles is possible while breathing from scuba equipment. The word 'SCUBA' is an acronym for "Self Contained Underwater Breathing Apparatus", but it is grammatically acceptable to refer to 'scuba equipment' or 'scuba apparatus' in conversation. The two types of scuba equipment are the "open-circuit" Aqua-lung, developed by Jacques-Yves Cousteau and the "closed-circuit" rebreather.

History of diving

Men and women have practiced breath-hold diving for centuries. Indirect evidence comes from ancient artifacts of undersea origin found on land (e.g. mother-of-pearl ornaments), and depictions of divers in ancient drawings. In ancient Greece breath-hold divers are known to have hunted for sponges and engaged in military exploits. Of the latter, the story of Scyllis (sometimes spelled Scyllias; about 500 B.C.) is perhaps the most famous, as told by the 5th century B.C. Greek historian Herodotus (and quoted in numerous modern texts). During a naval campaign the Greek Scyllis was taken aboard ship as prisoner by the Persian King Xerxes I. When Scyllis learned that Xerxes was to attack a Greek flotilla, he seized a knife and jumped overboard. The Persians could not find him in the water and presumed he had drowned. Scyllis surfaced at night and made his way among all the ships in Xerxes's fleet, cutting each ship loose from its moorings; he used a hollow reed as snorkel to remain unobserved. Then he swam nine miles (15 kilometers) to rejoin the Greeks off Cape Artemisium. The desire to go under water has probably always existed: to hunt for food, uncover artifacts, repair ships (or sink them), and perhaps just to observe marine life. Until humans found a way to breathe underwater, however, each dive was necessarily short and frantic. One of the major hurdles of diving is to stay under water for a longer period of time. Breathing through a hollow reed allows the body to be submerged, but reeds more than two feet long do not work well; difficulty inhaling against water pressure effectively limits snorkel length. Breathing from an air-filled bag brought under water was also tried, but it failed due to rebreathing of carbon dioxide. In the 16th century people began to use diving bells supplied with air from the surface, the first effective means of staying under water for any length of time. The bell was held stationary a few feet from the surface, its bottom open to water and its top portion containing air compressed by the water pressure. A diver standing upright would have his head in the air. He could leave the bell for a minute or two to collect sponges or explore the bottom, then return for a short while until air in the bell was no longer breathable. In 16th century England and France, full diving suits made of leather were used to depths of 60 feet. Air was pumped down from the surface with the aid of manual pumps. Soon helmets were made of metal to withstand even greater water pressure and divers went deeper. By the 1830s the surface-supplied air helmet was perfected well enough to allow extensive salvage work. Starting in the 19th century, two main avenues of investigation one scientific, the other technologic - greatly accelerated underwater exploration. Scientific research was advanced by the work of Paul Bert and John Scott Haldane, from France and Scotland, respectively. Their studies helped explain effects of water pressure on the body, and also define safe limits for compressed air diving. At the same time, improvements in technology - compressed air pumps, carbon dioxide scrubbers, regulators, etc., - made it possible for people to stay underwater for long periods. Used with permission from Lawrence Martin, M.D. Web Site of Origin: [http://www.lakesidepress.com/abcindex.htm www.lakesidepress.com/abcindex.htm]. The web pages containing the scuba book are inoperative at present, and have not yet been moved to a new site. The above URL is the global or index web site. And see Timeline of underwater technology.

Diving Issues

See Diving hazards and precautions.

Equipment to allow underwater breathing

The two most common types of equipment are:
- surface supplied diving, where the diver's breathing gas (usually air) is pumped down from the surface. Standard diving dress is a historically interesting type of surface supplied diving equipment.
- scuba, where the breathing gas supply is carried by the diver.

Need to see underwater

Diving masks and diving helmets solve this problem. Occasionally commando frogmen use special contact lenses instead.

Avoiding losing body heat

Water conducts heat from the diver 25 times better than air, which can lead to hypothermia. Except in very warm water, the diver needs the thermal insulation provided by wetsuits and drysuits. See the main article: Diving suit.

Avoiding skin cuts and grazes

Diving suits also help prevent the diver's skin being damaged by rough or sharp underwater objects and marine animals and coral.

Diving longer and deeper safely

There are a number of techniques to increase the diver's ability dive deeper and longer:
- technical diving - Diving Deeper than 130 feet and/or using mixed gases.
- surface supplied diving - use of umbilical gas supply and diving helmets
- saturation diving - long-term use of underwater habitats under pressure and a gradual release of pressure, over several days, in a decompression chamber at the end of a dive

Being mobile underwater

- The diver needs to be mobile underwater. Personal mobility is enhanced by fins worn on the feet and Diver Propulsion Vehicles. Other equipment to improve mobility includes diving bells and diving shots.

Sources

The Diving Manual, BSAC, ISBN 095389192534 Dive Leading, BSAC, ISBN 0953891941 The Club 1953-2003, BSAC, ISBN 095389195X

External links

-
- [http://www.sportdiver.com Sport Diver] - Dive magazine covering dive travel adventures, scuba diving gear reviews, and PADI Diving Society events & news. Official magazine of the PADI Diving Society.
- [http://www.fishid.com/ marine life learning center] - Marine Life identification and behavior books.
- [http://www.diversalertnetwork.org/ Divers Alert Network]
- [http://www.bsac.com British Sub-Aqua Club] - BSAC Welcome
- [http://scuba.rinkes.nl/ Brief history of diving] - From antiquity to the present.
- [http://www.scuba-guide.com Scuba Diving Guide] - Information for scuba divers.
- [http://www.thescubaguide.com The Scuba Guide] - Massive directory of gear and articles for beginning scuba divers.
- [http://www.scubamonster.com/Uwe/ForumList.aspx Scuba Monster] - Scuba Usenet discussions and archive.
- [http://www.nauiww.org/ NAUI Worldwide] is the world's oldest not-for-profit membership training agency organized solely to support and promote dive safety through education.
- [http://www.padi.com PADI International] - Professional Association of Diving Instructors.
- [http://www.ukdiving.co.uk UK Diving - Diving and Scuba resource] - Large Internet resource for divers. Category:Acronyms Category:Diving ms:Scuba ja:スキューバ・ダイビング

Breathing gas

Air is the most common and only natural breathing gas. Other artificial gases, either pure gases or mixtures of gases, are used in enclosed breathing environments such as SCUBA equipment, recompression chambers, submarines, space suits and anaesthetic machines. A safe breathing gas has three essential features:
- it must contain sufficient oxygen to support the life, consciousness and work rate of the breather.
- it must not contain harmful gases. Carbon monoxide and carbon dioxide are common poisons in breathing gases. There are many others.
- it must not become toxic when being breathed at high pressure such as when underwater. Oxygen and nitrogen are examples of gases that become toxic under pressure. Most breathing gases are a mixture of oxygen and one or more inert gases. The techniques used to fill diving cylinders with gases other than air are called gas blending.

Common diving breathing gases

Common diving breathing gases are:
- Air is a mixture of 21% oxygen, 78% nitrogen, and approximately 1% other trace gases; to simplify calculations this last 1% is usually treated as if it were nitrogen. Being cheap and simple to use, it is the most common diving gas. As its nitrogen component causes nitrogen narcosis it has a safe depth limit of 40 metres (130 feet) for most divers.
- Pure oxygen is mainly used to speed the shallow decompression stops at the end of a technical dive. It was much used in frogmen's rebreathers.
- Nitrox is a mixture of oxygen and air, and generally refers to mixtures which are more than 21% oxygen. It is mainly used instead of air to accelerate decompression or to decrease the risk of Decompression sickness.
- Trimix is a mixture of oxygen, nitrogen and helium and is often used during the deep phase of a technical dive.
- Heliox is a mixture of oxygen and helium and is often used in the deep phase of a commercial deep dive.
- Heliair is a mixture of oxygen, nitrogen, and helium. It is suitable for in the deep phase of a technical dive. Blended from helium and air, it always has a 21:79 ratio of oxygen to nitrogen; the balance of the mix is helium.
- Neox is a mixture of oxygen and neon sometimes employed for technical dives. It is rarely used due to its cost.

Individual component gases

Oxygen

Oxygen (O₂) must be present in every breathing gas. This is because it is essential to the human body's metabolic process, which sustains life. The human body cannot store oxygen for later use as it does with food. If the body is deprived of oxygen for more than a few minutes, unconsciousness results. The tissues and organs within the body (notably the heart and brain) are damaged if deprived of oxygen for much longer than four minutes. The proportion of oxygen in a breathing gas determines the depth at which the mixture gas can safely be used:
- hypoxic mixes have lower proportion of oxygen than air, 21%, or more strictly less than 16% oxygen and are designed only to be breathed at depth as a "bottom gas". Trimix, Heliox and Heliair are used to create typical hypoxic mixes and are used in technical diving as deep breathing gases.
- normoxic mixes have the same proportion of oxygen as air, 21%. The maximum operating depth of a normoxic mix could be as shallow as 47 mtres/155 feet. Trimix is often described as normoxic even the level of oxygen is less than 21% but high enough to be safe to breath on the surface.
- hyperoxic mixes have a more oxygen than 21%. Nitrox is a typical hyperoxic breathing gas. Breathing Nitrox, as opposed to air, can result in less nitrogen narcosis. Hyperoxic mixtures, when compared to air, cause oxygen toxicity at shallower depths but can be used to shorten decompression stops by drawing the dissolved nitrogen come out of the body more quickly. The minimum safe partial pressure of oxygen in a breathing gas is 16 kPa (0.16 bar). Below this partial pressure the diver risks unconsciousness and death due to hypoxia. When a hypoxic mix is breathed is shallow water it may not have a high enough ppO2 to keep the diver conscious. For this reason normoxic or hyperoxic "travel gases" are used at medium depth between the "bottom" and "decompression" phases of the dive. The maximum safe partial pressure of oxygen in a breathing gas depends on exposure time, but for dives of less than 3 hours is commonly considered to be 140 kPa (1.4 bar), although the U.S. Navy has been known to authorize dives with a partial oxygen pressure of as much as 180 kPa (1.8 bar). At high partial pressures or longer exposures, the diver risks oxygen toxicity including a seizure similar to an epileptic fit. Each breathing gas has a maximum operating depth that is determined by its oxygen content. Oxygen analysers are used to measure the concentration of oxygen in the gas mix. Filling a diving cylinder with pure oxygen costs around five times more than filling it with compressed air. As oxygen supports combustion and causes rust in diving cylinders, it should be handled with respect when gas blending.

Divox

"Divox" is oxygen. In the Netherlands, pure oxygen is regarded as medicinal (as opposed to industrial oxygen, such as that used in welding) and is only available on medical prescription. The diving industry "created" Divox and registered it as a trademark to circumvent the strict rules concerning medicinal oxygen thus making it easier for (recreational) scubadivers to obtain oxygen for blending their breathing gas.

Nitrogen

Nitrogen (N₂) is an inert gas and the main component of air, the cheapest and most common breathing gas used for diving. It causes nitrogen narcosis in the diver, so its use is limited to shallower dives. Nitrogen can cause decompression sickness. Equivalent air depth is used often used to help design a breathing gas mix by determining the maximum nitrogen content for a particular depth of dive. Many divers find that the level of narcosis caused by a 30-metre (100-foot) dive, whilst breathing air, is a comfortable maximum. The partial pressure of nitrogen at this depth on air is 316 kPa (3.16 bar) (Fraction of nitrogen x absolute pressure = 0.79 x 400 kPa). So, what fraction of nitrogen would cause the same narcosis at 60 metres? The answer is 45% nitrogen. (316 kPa/700 kPa)

Helium

Helium (He) is an inert gas that is less narcotic than nitrogen at diving pressures, so it is more suitable for deeper dives than nitrogen. But helium can still cause decompression sickness. It also causes High Pressure Nervous Syndrome. It is not very suitable for dry suit inflation due to its poor thermal insulation properties — helium is a very good conductor of heat, but air is a rather poor conductor of heat. Helium fills typically cost ten times more than an equivalent air fill. Helium also distorts the diver's voice, which may impede communication.

Neon

Neon (Ne) is an inert gas sometimes used in deep commercial diving but is very expensive. Like helium, it is less narcotic than nitrogen, but unlike helium, it does not distort the diver's voice.

Hydrogen

Hydrogen (H₂) has been used in deep diving gas mixes but is very explosive when mixed with more than about 4 to 5% oxygen (such as the oxygen found in breathing gas). This limits use of hydrogen to deep dives and complicated protocols to insure that oxygen is cleared from the lungs, the blood stream and the breathing equipment before breathing hydrogen starts. Like helium, it distorts the diver's voice. See

Unwelcome components of breathing gases

Many gases are not suitable for use in diving breathing gases. Here is an incomplete list.

Argon

Argon (Ar) is an inert gas that is more narcotic than nitrogen, so is not suitable as a diving breathing gas. It is used for dry suit inflation because of its good thermal insulation properties. Argon is much more expensive than air.

Carbon dioxide

Carbon dioxide (CO₂) is produced by the metabolism of the human body and causes carbon dioxide poisoning.

Carbon monoxide

Carbon monoxide (CO) is produced by incomplete combustion. Two common sources are:- : Internal combustion engine exhaust gas in the air being drawn into a diving air compressor. : Lubricants of the compressor firing under compression like happens in a diesel engine. It causes carbon monoxide poisoning.

Hydrocarbons

Hydrocarbons (C_xH_y) can occur due to compressor lubricants leaking, or due to incomplete combustion near the air intake, as for carbon monoxide. : They cause explosions, especially in high-oxygen mixtures. : Oil mist in breathed air can slowly damage the lungs and finally cause the lungs to degenerate into severe emphysema.

Moisture content

The process of compressing gas into a diving cylinder removes moisture from the gas. This is good for corrosion prevention in the cylinder but means that the diver inhales very dry gas. The dry gas extracts moisture from the divers lungs while underwater contributing to dehydration, which is also thought to be a predisposing risk factor of decompression sickness. It is also uncomfortable, causing a dry throat and making the diver thirsty. This problem is reduced with rebreathers because the soda lime reaction to remove carbon dioxide puts moisture back into the breathing gas. In a hot tropical climate, open circuit diving can accelerate heat exhaustion because of dehydration.

Gas detection and measurement

Divers find it difficult to detect most gases that are likely to be present in diving cylinders because they are colourless, odourless and tasteless. Electronic sensors exist for some gases, such as oxygen analysers, helium analyser, carbon monoxide detectors and carbon dioxide detectors. Oxygen analysers are commonly found underwater in rebreathers. Oxygen and helium analysers are often used on the surface during gas blending. Chemical and other types of gas detection methods are not often used in diving.

External links

- [http://www.onderwatersport.org/upload/documents/Facsheet2.pdf Dutch language fact sheet on Divox] Category:Diving

Buoyancy

In physics, buoyancy is an upward force on an object immersed in a fluid (i.e. a liquid or a gas), enabling it to float or at least to appear to become lighter. If the buoyancy exceeds the weight, then the object floats; if the weight exceeds the buoyancy, the object sinks. If the buoyancy equals the weight, the body has neutral buoyancy and may remain at its level. If its compressibility is less than that of the surrounding fluid, it is in stable equilibrium and will, indeed, remain at rest, but if its compressibility is greater, its equilibrium is unstable, and it will rise and expand on the slightest upward perturbation, but fall and compress on the slightest downward perturbation. It was the ancient Greek, Archimedes of Syracuse, who first discovered the law of buoyancy, sometimes called Archimedes' principle: :The buoyant force is equal to the weight of the displaced fluid. Typically, the weight of the displaced fluid is directly proportional to the volume of their displaced fluid (Specifically if the surrounding fluid is of uniform density.) Thus, among objects with equal masses, the one with greater volume has greater buoyancy. Suppose a rock's weight is measured at 10 Newtons when suspended by a string in a vacuum. Suppose that when the rock is lowered by the string into water, it displaces water whose weight is 3 Newtons. The force it then exerts on the string from which it hangs will be 10 Newtons minus the 3 Newtons of buoyant force: 10 - 3 = 7 Newtons. Buoyancy is the underlying principle of many vehicles such as boats, ships, balloons, and airships.

Density

If the weight of an object is less than the weight of the fluid that the object would displace if it was fully submerged, then the object is less dense than the fluid and it floats at a level so it displaces the same weight of fluid as the weight of the object. An object made of a material of higher density than the fluid, e.g. a metal object in water, can still float if it has a suitable shape (e.g. a hollow which is open upwards or downwards) that keeps a large enough volume of air below the surface level of the fluid. In that case, for the average density mentioned above, the air is included also, which may reduce this density to less than that of the fluid.

Acceleration and energy

Although Archimedes' principle gives the force on a buoyant object, this does not determine the related acceleration of the object in the usual way over Newton's first law. This is for two reasons: Not only has the mass of the object to be accelerated but also the mass of the displaced fluid. One can compare the situation to a scale, where the weight on one side is given by the object, and the weight on the other side by the displaced fluid element. Depending on which of the two is heavier, one side of the scale will drop and the other rise, but since both sides are rigidly connected, both masses have to be accelerated together at the same rate (albeit in opposite directions). The second reason is that viscosity dissipates energy, so that, even taking into account the kinetic and potential energies of the object and the fluid (but ignoring heat energy), energy is lost to viscosity, a form of friction. It is obvious that without taking the displaced fluid element into account, energy would not be conserved during the buoyant motion of an object as it would gain both potential and kinetic energy when rising in the fluid.

Nitrogen

Nitrogen is the chemical element in the periodic table that has the symbol N and atomic number 7. Commonly a colorless, odorless, tasteless and mostly inert diatomic non-metal gas, nitrogen constitutes 78 percent of Earth's atmosphere and is a constituent of all living tissues. Nitrogen forms many important compounds such as amino acids, ammonia, nitric acid, and cyanides.

Notable characteristics

Nitrogen is a non-metal, with an electronegativity of 3.0. It has five electrons in its outer shell, so is trivalent in most compounds. Pure nitrogen is an unreactive colorless diatomic gas at room temperature, and comprises about 78.08% of the Earth's atmosphere. It condenses at 77 K at atmospheric pressure and freezes at 63 K. Liquid nitrogen is a common cryogen.

Applications

Nitrogen Compounds

Molecular nitrogen in the atmosphere is relatively non-reactive, but in nature it is slowly converted into biologically (and industrially) useful compounds by some living organisms, notably certain bacteria (see Biological role below). The ability to combine or fix nitrogen is a key feature of modern industrial chemistry, where nitrogen (along with natural gas) is converted into ammonia (via the Haber process). Ammonia, in turn, can be used directly (primarily as a fertilizer), or as a precursor of many other important materials including explosives, largely via the production of nitric acid by the Ostwald process. The salts of nitric acid include important compounds like potassium nitrate (or saltpeter, important historically for its use in gunpowder) and ammonium nitrate, an important fertilizer. Various other nitrated organic compounds, such as nitroglycerin and trinitrotoluene, are used as explosives. Nitric acid is used as an oxidizer in liquid fueled rockets. Hydrazine and hydrazine derivatives find use as rocket fuels.

Molecular nitrogen (gas and liquid)

Nitrogen gas is readily produced by allowing liquid nitrogen (see below) to warm up and evaporate. It has a wide variety of applications, including serving as a more inert replacement for air where oxidation is undesirable;
- to preserve the freshness of packaged or bulk foods (by delaying rancidity and other forms of oxidative damage)
- on top of liquid explosives for safety It is also used in:
- the production of electronic parts such as transistors, diodes, and integrated circuits
- the manufacture of stainless steel
- filling automotive tires due to its inertness and lack of moisture or oxidative qualities, as opposed to air. A further example of its versitility is its use (as a preferred alternative to carbon dioxide) to pressurize kegs of some beers, particularly thicker stouts and Scottish and English ales, due to the smaller bubbles it produces, which make the dispensed beer smoother and headier. A modern application of a pressure sensitive nitrogen capsule known commonly as a "widget" now allows nitrogen charged beers to be packaged in cans and bottles. A very popular example of this is Guinness Draught. Liquid nitrogen is produced industrially in large quantities by distillation from liquid air and is often referred to by the quasi-formula LN₂. It is a cryogenic (extremely cold) fluid which can cause instant frostbite on direct contact with living tissue. When appropriately insulated from ambient heat it serves as a compact and readily transported source of nitrogen gas without pressurization. Further, its ability to maintain an unearthly temperature as it evaporates (77 K, -196 °C or -320 °F) makes it extremely useful in a wide range of applications as an open-cycle refrigerant, including;
- the immersion freezing and transportation of food products
- the preservation of bodies, reproductive cells (sperm and egg), and biological samples and materials
- in the study of cryogenics
- for demonstrations in science education
- in dermatology for removing unsightly or potentially malignant skin lesions,e.g., warts, actinic keratosis, etc.

History

Nitrogen (Latin nitrum, Greek Nitron meaning "native soda", "genes", "forming") is formally considered to have been discovered by Daniel Rutherford in 1772, who called it noxious air or fixed air. That there was a fraction of air that did not support combustion was well known to the late 18th century chemist. Nitrogen was also studied at about the same time by Carl Wilhelm Scheele, Henry Cavendish, and Joseph Priestley, who referred to it as burnt air or phlogisticated air. Nitrogen gas was inert enough that Antoine Lavoisier referred to it as azote, which stands for without life; this term has become the French word for "nitrogen" and later spread out to many other languages. Compounds of nitrogen were known in the Middle Ages. The alchemists knew nitric acid as aqua fortis. The mixture of nitric and hydrochloric acids was known as aqua regia, celebrated for its ability to dissolve gold. The earliest industrial and agricultural applications of nitrogen compounds used it in the form of saltpeter (sodium- or potassium nitrate), notably in gunpowder, and much later, as fertilizer, and later still, as a chemical feedstock.

Occurrence

Nitrogen is the largest single component of the Earth's atmosphere (78.084% by volume, 75.5% by weight) and is acquired for industrial purposes by the fractional distillation of liquid air or by mechanical means of gaseous air (i.e. pressurised reverse osmosis membrane or PSA (Pressure Swing Adsorption). Compounds that contain this element have been observed in outer space. Nitrogen-14 is created as part of the fusion processes in stars. Nitrogen is a large component of animal waste (for example, guano), usually in the form of urea, uric acid, and compounds of these nitrogenous products. Molecular nitrogen has been known to occur in Titan's atmosphere for some time, and has now been detected in interstellar space by David Knauth and coworkers using the Far Ultraviolet Spectroscopic Explorer.

Compounds

The main hydride of nitrogen is ammonia (NH₃) although hydrazine (N₂H₄) is also well known. Ammonia is somewhat more basic than water, and in solution forms ammonium ions (NH₄⁺). Liquid ammonia is in fact slightly amphiprotic and forms ammonium and amide ions (NH₂^-); both amides and nitride (N^3-) salts are known, but decompose in water. Singly and doubly substituted compounds of ammonia are called amines. Larger chains, rings and structures of nitrogen hydrides are also known but virtually unstable. Other classes of nitrogen anions are azides (N₃^-), which are linear and isoelectronic to carbon dioxide. Another molecule of the same structure is dinitrogen monoxide (N₂O), or laughing gas. This is one of a variety of oxides, the most prominent of which are nitrogen monoxide (NO) and nitrogen dioxide (NO₂), which both contain an unpaired electron. The latter shows some tendency to dimerize and is an important component of smog. The more standard oxides, dinitrogen trioxide (N₂O₃) and dinitrogen pentoxide (N₂O₅), are actually fairly unstable and explosive. The corresponding acids are nitrous (HNO₂) and nitric acid (HNO₃), with the corresponding salts called nitrites and nitrates. Nitric acid is one of the few acids stronger than hydronium.

Biological role

Nitrogen is an essential part of amino and nucleic acids which makes nitrogen vital to all life. Legumes like the soybean plant, can recover nitrogen directly from the atmosphere because their roots have nodules harboring microbes that do the actual conversion to ammonia in a process known as nitrogen fixation. The legume subsequently converts ammonia to nitrogen oxides and amino acids to form proteins.

Isotopes

There are two stable isotopes: N-14 and N-15. By far the most common is N-14 (99.634%), which is produced in the CNO cycle in stars. The rest is N-15. Of the ten isotopes produced synthetically, one has a half life of nine minutes and the remaining isotopes have half lives on the order of seconds or less. Biologically-mediated reactions (e.g., assimilation, nitrification, and denitrification) strongly control nitrogen dynamics in the soil. These reactions almost always result in N-15 enrichment of the substrate and depletion of the product. Although precipitation often contains subequal quantities of ammonium and nitrate, because ammonium is preferentially retained by the canopy relative to atmospheric nitrate, most of the atmospheric nitrogen that reaches the soil surface is in the form of nitrate. Soil nitrate is preferentially assimilated by tree roots relative to soil ammonium.

Precautions

Nitrate fertilizer washoff is a major source of ground water and river pollution. Cyano (-CN) containing compounds form extremely poisonous salts and are deadly to many animals and all mammals.

References

- [http://periodic.lanl.gov/elements/7.html Los Alamos National Laboratory – Nitrogen]

External links

- [http://www.webelements.com/webelements/elements/text/N/index.html WebElements.com – Nitrogen]
- [http://education.jlab.org/itselemental/ele007.html It's Elemental – Nitrogen]
- [http://www.sunysccc.edu/academic/mst/ptable/n.html Schenectady County Community College – Nitrogen]
- [http://www.uigi.com/nitrogen.html Nitrogen N2 Properties, Uses, Applications]
- [http://box27.bluehost.com/~edsanvil/wiki/index.php?title=Nitrogen_gas Computational Chemistry Wiki] Category:Nonmetals Category:Pnictogens Category:Nitrogen metabolism ko:질소 ja:窒素 simple:Nitrogen th:ไนโตรเจน

Oxygen

Oxygen is a chemical element in the periodic table. It has the symbol O and atomic number 8. The element is very common, found not only on Earth but throughout the universe, usually covalently bonded with other elements. Unbound oxygen (usually called molecular oxygen, O₂, a diatomic molecule) first appeared on Earth during the Paleoproterozoic era (between 2500 million years ago and 1600 million years ago) and as a product of the metabolic action of early anaerobes (archaea and bacteria). The presence of free oxygen drove most of the organisms then living to extinction. The atmospheric abundance of free oxygen in later geological epochs and up to the present has been largely driven by photosynthetic organisms, roughly three quarters by phytoplankton and algae in the oceans and one quarter from terrestrial plants.

Characteristics

At standard temperature and pressure, oxygen is mostly found as a gas consisting of a diatomic molecule with the chemical formula O₂. O₂ has two energetic forms:
- The low-energy predominant single-bonded diradical triplet oxygen. This native diradical quality of oxygen contributes to its destructive chemical nature. This form is stabilized by the degeneracy effect.
- The high-energy double-bonded molecule singlet oxygen. Oxygen is a major component of air, produced by plants during photosynthesis, and is necessary for aerobic respiration in animals. The word oxygen derives from two words in Greek, οξυς (oxys) (acid, sharp) and γεινομαι (geinomai) (engender). The name "oxygen" was chosen because, at the time it was discovered in the late 18th century, it was believed that all acids contained oxygen. The definition of acid has since been revised to not require oxygen in the molecular structure. Liquid O₂ and solid O₂ have a light blue color and both are highly paramagnetic. Liquid O₂ is usually obtained by the fractional distillation of liquid air. Liquid and solid O₃ (ozone) have a deeper color of blue. A recently discovered allotrope of oxygen, tetraoxygen (O₄), is a deep red solid that is created by pressurizing O₂ to the order of 20 GPa. Its properties are being studied for use in rocket fuels and similar applications, as it is a much more powerful oxidizer than either O₂ or O₃.

Applications

Liquid oxygen finds use as an oxidizer in rocket propulsion. Oxygen is essential to respiration, so oxygen supplementation has found use in medicine (as oxygen therapy). People who climb mountains or fly in airplanes sometimes have supplemental oxygen supplies (as air). Oxygen is used in welding (such as the oxyacetylene torch), and in the making of steel and methanol. Oxygen presents two absorption bands centered in the wavelengths 687 and 760 nanometers. Some scientists have proposed to use the measurement of the radiance coming from vegetation canopies in those oxygen bands to characterize plant health status from a satellite platform. This is because in those bands, it is possible to discriminate the vegetation's reflectance from the vegetation's fluorescence, which is much weaker. The measurement presents several technical difficulties due to the low signal to noise ratio and due to the vegetation's architecture, but it has been proposed as possibility to monitor the carbon cycle from satellite, thus in a global scale. Oxygen, as a mild euphoric, has a history of recreational use that extends into modern times. Oxygen bars can be seen at parties to this day. In the 19th century, oxygen was often mixed with nitrous oxide to promote an analgesic effect; indeed, such a mixture (Entonox) is commonly used in medicine today.

History

Oxygen was first discovered by Michał Sędziwój, Polish alchemist and philosopher in late 16th century. Sędziwój assumed the existence of oxygen by warming nitre (saltpeter). He thought of the gas given off as "the elixir of life". Oxygen was again discovered by the Swedish pharmacist Carl Wilhelm Scheele sometime before 1773, but the discovery was not published until after the independent discovery by Joseph Priestley on August 1, 1774, who called the gas dephlogisticated air (see phlogiston theory). Priestley published his discoveries in 1775 and Scheele in 1777; consequently Priestley is usually given the credit. It was named by Antoine Laurent Lavoisier after Priestley's publication in 1775.

Occurrence

Oxygen is the second most common component of the earth's atmosphere (20.947% by volume).

Compounds

Due to its electronegativity, oxygen forms chemical bonds with almost all other elements (which is the origin of the original definition of oxidation). The only elements to escape the possibility of oxidation are a few of the noble gases. The most famous of these oxides is dihydrogen monoxide, or water (H₂O). Other well known examples include compounds of carbon and oxygen, such as carbon dioxide (CO₂), alcohols (R-OH), aldehydes, (R-CHO), and carboxylic acids (R-COOH). Oxygenated radicals such as chlorates (ClO₃⁻), perchlorates (ClO₄⁻), chromates (CrO₄²⁻), dichromates (Cr₂O₇²⁻), permanganates (MnO₄⁻), and nitrates (NO₃⁻) are strong oxidizing agents in and of themselves. Many metals such as iron bond with oxygen atoms, iron (III) oxide (Fe₂O₃). Ozone (O₃) is formed by electrostatic discharge in the presence of molecular oxygen. A double oxygen molecule (O₂)₂ is known and is found as a minor component of liquid oxygen. Epoxides are ethers in which the oxygen atom is part of a ring of three atoms.

Isotopes

Oxygen has fifteen known isotopes with atomic masses ranging from 12 to 26. Three of them are stable and twelve are radioactive. The radioisotopes all have half lives of less than three minutes. The stable isotopes have mass numbers of 16, 17 and 18, of which oxygen-16 is the most common (over 99%).

Precautions

Oxygen can be toxic at elevated partial pressures (i.e. high relative concentrations). This is important in some forms of scuba diving, such as with a rebreather. Certain derivatives of oxygen, such as ozone (O₃), singlet oxygen, hydrogen peroxide, hydroxyl radicals and superoxide, are also highly toxic. The body has developed mechanisms to protect against these toxic species. For instance, the naturally-occurring glutathione can act as an antioxidant, as can bilirubin which is normally a breakdown product of hemoglobin. Highly concentrated sources of oxygen promote rapid combustion and therefore are fire and explosion hazards in the presence of fuels. This is true as well of compounds of oxygen such as chlorates, perchlorates, dichromates, etc. Compounds with a high oxidative potential can often cause chemical burns. The fire that killed the Apollo 1 crew on a test launchpad spread so rapidly because the pure oxygen atmosphere was at normal atmospheric pressure instead of the one third pressure that would be used during an actual launch. (See partial pressure.) Oxygen derivatives are prone to form free radicals, especially in metabolic processes. Because they can cause severe damage to cells and their DNA, they are thought to be related to cancer and aging.

References

- [http://periodic.lanl.gov/elements/8.html Los Alamos National Laboratory – Oxygen]
- [http://physics.nist.gov/cgi-bin/AtData/main_asd Nist atomic spectra database]
- [http://chartofthenuclides.com/default.html Nuclides and Isotopes Fourteenth Edition]: Chart of the Nuclides, General Electric Company, 1989

External links

- [http://www.priestleysociety.net Priestley Society, Dedicated to Joseph Priestley the man who discovered oxygen]
- [http://www.best-home-remedies.com/minerals/oxygen.htm Oxygen - Benefits, Deficiency Symptoms And Food Sources]
- [http://www.josephpriestley.info Joseph Priestley Information Website, about the man who discovered oxygen]
- [http://periodic.lanl.gov/elements/8.html Los Alamos National Laboratory – Oxygen]
- [http://www.webelements.com/webelements/elements/text/O/index.html WebElements.com – Oxygen]
- [http://education.jlab.org/itselemental/ele008.html It's Elemental – Oxygen]
- [http://members.tripod.com/tjaartdb0/html/oxygen_toxicity.html Oxygen Toxicity]
- [http://www.uigi.com/oxygen.html Oxygen (O2) Properties, Uses, Applications]
- [http://www.compchemwiki.org/index.php?title=Oxygen Computational Chemistry Wiki]
- [http://koti.mbnet.fi/antitz/dime/en Tests with liquid oxygen :-)] Category:Nonmetals Category:Chalcogens als:Sauerstoff ko:산소 ms:Oksigen ja:酸素 simple:Oxygen th:ออกซิเจน

Carbon dioxide

Carbon dioxide is an atmospheric gas comprised of one carbon and two oxygen atoms. A very widely known chemical compound, it is frequently called by its formula CO₂. In its solid state, it is commonly known as dry ice. Carbon dioxide derives from multiple sources including volcanic outgassing, the combustion of organic matter and respiration processes of living aerobic organisms. 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 because of the associated asphyxiation risk), produces a sour taste in the mouth and a 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 25 °C 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 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. Test For Carbon Dioxide. When a lighted splint is inserted into a test tube containing carbon dioxide, the flame is immediately extinguished, as carbon dioxide does not support combustion. (Certain fire extinguishers contain carbon dioxide to extinguish the flame). To further confirm that the gas is carbon dioxide, the gas may be bubbled into calcium hydroxide solution. The calcium hydroxide turns milky because of the formation of calcium carbonate.

Uses

Liquid and solid carbon dioxide are important refrigerants, especially in the food industry, where they are employed during the transportation and storage of ice cream and other frozen foods. Carbon dioxide is used to produce carbonated 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, nonflammable 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, paintball markers, for inflating bicycle tires, and for making seltzer. 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 because of 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, and is used to remove caffeine from coffee. 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 levels 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 medicine, up to 5% carbon dioxide is added to pure oxygen for stimulation of breathing after apnea and to stabilize the O₂/CO₂ balance in blood. A common type of industrial gas laser, the carbon dioxide laser, uses carbon dioxide as a medium. Carbon dioxide is commonly injected into or adjacent to producing oil wells. It will act as both a pressurizing agent and, when dissolved into the underground crude oil, will significantly reduce its viscosity, enabling the oil to flow more rapidly through the earth to the removal well. In mature oil fields, extensive pipe networks are used to carry the carbon dioxide to the injection points.

Dry Ice

Dry ice is a genericized trademark for solid ("frozen") carbon dioxide. The term was coined in 1925 by Prest Air Devices, a company formed in Long Island City, New York in 1923. Dry ice at normal pressures does not melt into liquid carbon dioxide but rather sublimates directly into carbon dioxide gas at −78.5 °C (−109.3 °F). Hence it is called "dry ice" as opposed to normal "wet" ice (frozen water). Dry ice is produced by compressing carbon dioxide gas to a liquid form, removing the heat produced by the compression (see Charles' law), and then letting the liquid carbon dioxide expand quickly. This expansion causes a drop in temperature so that some of the CO₂ freezes into "snow", which is then compressed into pellets or blocks.

Uses

temperature, New York, USA)]]
- Cooling foodstuffs, biological samples, and other perishable items.
- Producing "dry ice fog" for special effects. When dry ice is put into contact with water, the frozen carbon dioxide sublimates into a mixture of cold carbon dioxide gas and cold humid air. This causes condensation and the formation of fog; see fog machine. The effect of fog by the mixture of dry ice with water, is best formed when the water is warm, rather than cold.
- Tiny pellets of dry ice (instead of sand) are shot at a surface to be cleaned. Dry ice is not as hard as sand, but it speeds processing by sublimating to nothing and does not produce nearly as much lung-damaging dust.
- Increasing precipitation from existing clouds or decreasing cloud thickness by cloud seeding.
- Producing carbon dioxide gas as needed in such systems as the fuel tank inerting system in the B-47 aircraft.
- Brass or other metallic bushings are buried in dry ice to shrink their size so they will fit inside a machined hole. When the bushing warms back up, it expands and makes an extremely tight fit.

Handling

Because of its particular characteristics, dry ice requires special precautions when handling. It is extremely cold and there should be no direct contact with skin (i.e., wear proper insulating gloves). It is constantly sublimating to carbon dioxide gas, so it cannot be stored in a sealed container as the pressure buildup will quickly cause the container to explode. The sublimated gas must be ventilated; otherwise, it may fill the enclosed space and create a suffocation hazard. Special care for ventilating vehicles is needed as well because of the small space. People who handle dry ice should also be aware that carbon dioxide is heavier than air and will sink to the floor. Some markets require those purchasing dry ice to be of 18 years of age or older.

Biology

Carbon dioxide is an end 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 from the body's tissues to the lungs where it is exhaled. Carbon dioxide content in fresh air is approximately 0.04%, and in exhaled air approximately 4.5%. When inhaled in high concentrations (about 5% by volume), it is toxic to humans and other animals. Hemoglobin, the main oxygen-carrying molecule in red blood cells, can carry both oxygen and carbon dioxide, although in quite different ways. The decreased binding to oxygen in the blood due to increased carbon dioxide levels is known as the Haldane Effect, and is important in the transport of carbon dioxide from the tissues to the lungs. Conversely, a rise in the partial pressure of CO₂ or a lower pH will cause offloading of oxygen from hemoglobin. This is known as the Bohr Effect. According to a study by the USDA [http://itest.slu.edu/articles/90s/hannan.html], an average person's respiration generates approximately 450 liters (roughly 900 grams) of carbon dioxide per day. CO₂ is carried in blood in three different ways. Most of it (about 80%–90%) is converted to bicarbonate ions HCO₃⁻ by the enzyme carbonic anhydrase in the red blood cells. 5%–10% is dissolved in the plasma and 5%–10% is bound to hemoglobin as carbamino compounds. The exact percentages vary depending whether it is arterial or venous blood. The CO₂ bound to hemoglobin does not bind to the same site as oxygen; rather it combines with the N-terminal groups on the four globin chains. However, because of allosteric effects on the hemoglobin molecule, the binding of CO₂ does decrease the amount of oxygen that is bound for a given partial pressure of oxygen. Carbon dioxide may be one of the mediators of local autoregulation of blood supply. If it is high, the capillaries expand to allow a greater blood flow to that tissue. Bicarbonate ions are crucial for regulating blood pH. As breathing rate influences the level of CO₂ in blood, too slow or shallow breathing causes respiratory acidosis, while too rapid breathing, hyperventilation, leads to respiratory alkalosis. It is interesting to note that although it is oxygen that the body requires for metabolism, it is not low oxygen levels that stimulate breathing, but is instead higher carbon dioxide levels. As a result, breathing low-pressure air or a gas mixture with no oxygen at all (e.g., pure nitrogen) leads to loss of consciousness without subjective breathing problems. This is especially perilous for high-altitude fighter pilots, and is also the reason why the instructions in commercial airplanes for case of loss of cabin pressure stress that one should apply the oxygen mask to oneself before helping others—otherwise one risks going unconscious without being aware of the imminent peril. 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. 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. Carbon dioxide, either as a gas or as dry ice, should be handled only in well-ventilated areas. See also: Arterial blood gas.

Atmosphere

Arterial blood gas As of 2004, the earth's atmosphere is about 0.038% by volume (380 µL/L or ppmv) or 0.053% by weight CO₂. This represents about 2.7 × 10¹² tonnes of CO₂. Because of 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% of the amount which is released by human activities. short tons–2000.]] Since the start of the Industrial Revolution, the atmospheric CO₂ concentration has increased by approximately 110 µL/L or about 40%, most of it released since 1945. Monthly measurements taken at Mauna Loa [http://cdiac.esd.ornl.gov/trends/co2/sio-mlo.htm] since 1958 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 is the second major cause. In 1997, Indonesian peat fires may have released 13%–40% as much carbon as fossil fuel burning does [http://en.wikipedia.org/wiki/Peat#Peat_fires]. Various techniques have been proposed for removing excess carbon dioxide from the atmosphere in carbon dioxide sinks. Not all the emitted CO₂ remains in the atmosphere; some is absorbed in the oceans or biosphere. The ratio of the emitted CO₂ to the increase is atmospheric CO₂ is known as the airborne fraction (Keeling et al., 1995); this varies for short-term averages but is typically 57% over longer (5 year) periods.

carbon dioxide sink 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. The effect of combustion-produced carbon dioxide on climate is called the Callendar effect.

Variation in the past

Callendar effect The most direct method for measuring atmospheric carbon dioxide concentrations for periods before direct sampling is to measure bubbles of air (fluid or gas inclusions) trapped in the Antarctic or Greenland ice caps. The most widely accepted of such studies come from a variety of Antarctic cores and indicate that atmospheric CO₂ levels were about 260–280µL/L immediately before industrial emissions began and did not vary much from this level during the preceding 10,000 years. The longest ice core record comes from East Antarctica, where ice has been sampled to an age of 650,000 years before the present. [http://pubs.acs.org/cen/news/83/i48/8348notw1.html] 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. Some studies have disputed the claim of stable CO₂ levels during the present interglacial (the last 10 kyr). Based on an analysis of fossil leaves, Wagner et al. argued that CO₂ levels during the period 7–10 kyr ago were significantly higher (~300 µL/L) and contained substantial variations that may be correlated to climate variations. Others have disputed such claims, suggesting they are more likely to reflect calibration problems than actual changes in CO₂. Relevant to this dispute is the observation that Greenland ice cores often report higher and more variable CO₂ values than similar measurements in Antarctica. However, the groups responsible for such measurements (e.g., Smith et al.) believe the variations in Greenland cores result from in situ decomposition of calcium carbonate dust found in the ice. When dust levels in Greenland cores are low, as they nearly always are in Antarctic cores, the researchers report good agreement between Antarctic and Greenland CO₂ measurements. calcium carbonate On longer timescales, 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.[http://www.grida.no/climate/ipcc_tar/wg1/fig3-2.htm] On long timescales, 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 favored 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, present carbon dioxide levels are likely higher now than at any time during the past 20 million years [http://www.grida.no/climate/ipcc_tar/wg1/107.htm#331] and at the same time lower than at any time in history if we look at time scales longer than 50 million years.

Oceans

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 illustrate that carbon dioxide is produced by animal respiration and microbial fermentation. In 1772, Joseph Priestley used carbon dioxide produced from the action of sulfuric acid on limestone to prepare soda water, the first known instance of an artificially carbonated beverage. Carbon dioxide was first liquefied (at elevated pressures) in 1823 by Humphrey Davy and Michael Faraday. The earliest description of solid carbon dioxide was given by Charles Thilorier, who in 1834 opened a pressurized container of liquid carbon dioxide, only to find that the cooling produced by the rapid evaporation of the liquid yielded a "snow" of solid CO₂.

References

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External links

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- [http://www.dryiceinfo.com/science.htm Dry Ice information]
- Bassam Z. Shakhashiri: [http://scifun.chem.wisc.edu/chemweek/CO2/CO2.html Chemical of the Week: Carbon Dioxide]
- Keeling, C.D. and T.P. Whorf: [http://cdiac.esd.ornl.gov/trends/co2/sio-mlo.htm Atmospheric carbon dioxide record from Mauna Loa], 2002
- [http://www.usatoday.com/weather/news/2004-03-21-co2-buildup_x.htm Mauna Loa 2004 update]
- [http://www.uigi.com/carbondioxide.html CO2 Carbon Dioxide Properties, Uses, Applications]
- [http://www.compchemwiki.org/index.php?title=Carbon_dioxide Computational Chemistry Wiki]
- [http://scifun.chem.wisc.edu/chemweek/CO2/CO2_phase_diagram.gif Pressure-Temperature phase diagram for carbon dioxide] Category:Inorganic carbon compounds Category:Oxides Category:Greenhouse gases Category:Propellants Category:Household chemicals Category:Solvents Category:Refrigerants Category:Fire suppression agents ko:이산화 탄소 ms:Karbon dioksida ja:二酸化炭素 simple:Carbon dioxide th:คาร์บอนไดออกไซด์

Helium

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- Atmospheric value, abundance may differ elsewhere. :This page is about the chemical element helium. For the American indie rock band Helium see Helium (band). Helium (He) is a colorless, odorless, tasteless, non-toxic, nearly inert monatomic chemical element that heads the noble gas series in the periodic table and whose atomic number is 2. Its boiling and melting points are the lowest among the elements and it exists only as a gas except in extreme conditions. Extreme conditions are also needed to create the small handful of helium compounds, which are all unstable at standard temperature and pressure. Its most abundant stable isotope is helium-4 and its rare stable isotope is helium-3. The behavior of liquid helium-4's two varieties—helium I and helium II—is important to researchers studying quantum mechanics (in particular the phenomenon of superfluidity) and those looking at the effects that near absolute zero temperatures have on matter (such as superconductivity). Helium is the second most abundant and second lightest element in the periodic table. In the modern Universe almost all new helium is created as a result of the nuclear fusion of hydrogen in stars. On Earth it is created by the radioactive decay of much heavier elements (alpha particles are helium-4 nuclei produced by alpha-decay). After its creation, part of it is trapped with natural gas in concentrations up to 7% by volume. It is extracted from the natural gas by a low temperature separation process called fractional distillation. In 1868 the French astronomer Pierre Janssen first detected helium as an unknown yellow spectral line signature in light from a solar eclipse. Since then large reserves of helium have been found in the natural gas fields of the United States, which is by far the largest supplier of the gas. Helium is used in cryogenics, in deep-sea breathing systems, to cool superconducting magnets, in helium dating, for inflating balloons, for providing lift in airships and as a protective gas for many industrial uses (such as arc welding and growing silicon wafers). Inhaling a small volume of the gas temporarily changes the quality of one's voice.

Electron energy levels

Depending on the spin orientation of the two electrons in the Helium atom, one speaks of parahelium for two anti-parallel spins (S=0) and of orthohelium for two parallel spins (S=1). For the orthohelium one of the electrons does not sit in the ground orbital (1s). [http://hyperphysics.phy-astr.gsu.edu/hbase/quantum/helium.html]

Applications

orthohelium Pressurized helium is commercially available. Helium is used for many purposes that require one or more of its unique properties; low boiling point, low density, low solubility, high thermal conductivity, or its inertness. Airships and balloons (toy, weather, and research) are inflated with helium because it is lighter than air (1 m³ of helium will lift 1 kg). Helium is currently preferred to hydrogen in airships because, while it is more expensive, it is not flammable and has 92.64% of the lifting power of hydrogen. Trimix, a mixture of helium, oxygen, and nitrogen, is used in deep-sea breathing gas systems to reduce the risk of nitrogen narcosis (high pressure nitrogen having a narcotic effect on the brain) and oxygen toxicity at high pressures. Higher pressures require a greater proportion of helium and reduced amounts of nitrogen and oxygen (every ten-meter increase in depth yields a one atmosphere increase of pressure). Heliox, a mixture of helium and oxygen, and heliair, a mixture of air and helium, is also used in this way. Below 130 meters (430 ft) a mixture of hydrogen, helium, and oxygen called hydreliox is used to help prevent high pressure nervous syndrome. All these uses rely on helium's very low solubility in water (the major component of blood). The extremely low boiling point makes helium useful as a coolant in magnetic resonance imaging, superconducting magnets, cryogenics, and to remove thermal noise from detectors used in astronomy. The extreme coldness of liquid helium is also used to produce superconductivity in some ordinary metals such as lead (lead becomes superconductive at 7.3 K), allowing for a completely free flow of electrons in the metal. Other uses:
- Because of its high thermal conductivity and inertness, helium is used as a coolant in some nuclear reactors (for example, pebble-bed reactors) and in arc welding air-sensitive metals that require heavy welds.
- The high thermal conductivity and sound velocity of helium is also desirable in thermoacoustic refrigeration. The inertness of helium adds to the environmental advantage of this technology over conventional refrigeration systems which may contribute to ozone depleting and global warming effects.
- Its inertness makes it useful as a protective gas in growing silicon and germanium crystals, in titanium and zirconium production, protecting important historical documents, and in gas chromatography. This property also makes it useful in pressurizing liquid fuel rockets (see below) and in supersonic wind tunnels.
- The gain medium of the helium-neon laser (the first gas laser) most commonly used to scan bar codes is a mixture of helium and neon.
- This gas' rate of diffusion through solids is three times that of normal air, making it an excellent component in leak detection in high-vacuum equipment and high pressure containers.
- In rocketry helium is used as an ullage medium to displace fuel and oxidizers in storage tanks and to condense hydrogen and oxygen to make rocket fuel. It is also used to purge fuel and oxidizer from ground support equipment prior to launch and to precool liquid hydrogen in space vehicles. For example, the Saturn 5 booster used in the Apollo program needed about 13 million ft³ (370,000 m³) of helium to launch.
- Physics researchers use alpha particles (helium nuclei) in particle accelerators and nuclear reaction experiments.
- Helium gas is used to fill the space between lenses in some solar telescopes because its extremely low index of refraction reduces the distorting effect of temperature variations in the gas filling the telescope (some telescopes are filled with vacuum instead).
- Radioactive decay of uranium and thorium produces alpha particles that quickly become helium. This happens at a known constant rate so if the containing rock or mineral can retain its helium then the ratio of helium to its radioactive parent atoms indicates its age. Alternatively, if the helium is not well-retained, the ratio of helium-3 to helium-4 contains some of the same information, since only helium-4 is produced by radioactive decay. Use of helium in this way is called helium dating.

History

Discoveries

Helium was first detected on August 18, 1868 as a bright yellow line with a wavelength of 587.49 nm in the spectrum of the chromosphere of the Sun, by French astronomer Pierre Janssen during a total solar eclipse in India. Janssen was at first ridiculed since no element had ever been detected in space before being found on Earth. October 20th the same year, English astronomer Norman Lockyer also observed the same yellow line in the solar spectrum and concluded that it was caused by an unknown element after unsuccessfully testing to see if it were some new type of hydrogen. Since it was near the Fraunhofer D line he later named the new line D₃, distinguishing it from the nearby D₁ and D₂ double lines of sodium. He and English chemist Edward Frankland named the element after the Greek word for the Sun god, Helios, and, assuming it was a metal, gave it an -ium ending (a mistake that was never corrected). British chemist William Ramsay isolated helium on March 26, 1895 by treating cleveite (now known to be uraninite) with mineral acids. Ramsay