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Nitrogen

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

Chemical element

A chemical element, often called simply element, is a chemical substance that canot be divided or changed into other chemical substances by any ordinary chemical technique. The smallest unit of this kind of chemical substances is an atom. An element is a class of substances that contain the same number of protons in all its atoms.

Chemistry terminology

Earlier an element or pure element was defined as a substance which "cannot be further broken down into another compound with different chemical properties" -- which should be taken to mean it consists of atoms of one element. However, due to allotropy, the isotope effect, and the confusion with the more useful term referring to the general class of atoms (irrespective of what compound it may be in), this usage is in disfavor amongst contemporary chemists, and sees restricted, mostly historical, use. This definition was motivated by the observation that these elements could not be dissociated by chemical means into other compounds. For example, water could be converted into hydrogen and oxygen, but hydrogen and oxygen could not be further decomposed, thus "elemental". There are also many counterexamples (for example "elemental oxygen" (O₂) can be decomposed by solely chemical means into oxygen ions and atoms which have drastically different chemical properties). The remainder of this article will concern itself with the first definition.

Description

The atomic number of an element, Z, is equal to the number of protons which defines the element. For example, all carbon atoms contain 6 protons in their nucleus, so for carbon Z=6. These atoms may have different amounts of neutrons, and are known as isotopes of the element. The atomic mass of an element, A, is measured in unified atomic mass units (u) is the average mass of all the atoms of the element in an environment of interest (usually the earth's crust and atmosphere). Since electrons are light, and neutrons are barely more than the mass of the proton, this usually corresponds to the sum of the protons and neutrons in the nucleus of the most abundant isotope, though this is not always the case (notably chlorine, which is about three-quarters ³⁵Cl and a quarter ³⁷Cl). Some isotopes are radioactive and decay into other elements upon radiating an alpha or beta particle. Some elements have no nonradioactive isotopes, in particular all elements with Z >= 84. The lightest elements are hydrogen and helium. Hydrogen is thought to be the first element to appear after the Big Bang. All the heavier elements, are made naturally and artificially through various methods of nucleosynthesis. As of 2005, there are 116 known elements: 93 occur naturally on earth (including technetium and plutonium), and 94 (including promethium) have been detected so far in the universe. The 23 elements not found on earth are derived artificially; the first purportedly synthesized element was technetium, in 1937, although the trace amounts of naturally occurring technetium were not known then. All artificially derived elements are radioactive with short half-lives so that any such atoms that were present at the formation of Earth are extremely likely to have already decayed. Lists of the elements by name, by symbol, by atomic number, by density, by melting point and by boiling point are available. The most convenient presentation of the elements is in the periodic table, which groups elements with similar chemical properties together.

Nomenclature

The naming of elements precedes the atomic theory of matter, although at the time it was not known which chemicals were elements and which compounds. When it was learned, existing names (e.g., gold, mercury, iron) were kept in most countries, and national differences emerged over the names of elements either for convenience, linguistic niceties, or nationalism. For example, the Germans use "Wasserstoff" for "hydrogen" and "Sauerstoff" for "oxygen," while some romance languages use "natrium" for "sodium" and "kalium" for "potassium," and the French prefer the obsolete but historic term "azote" for "nitrogen." But for international trade, the official names of the chemical elements both ancient and recent are decided by the International Union of Pure and Applied Chemistry, which has decided on a sort of international English language. That organization has recently prescribed that "aluminium" and "caesium" take the place of the US spellings "aluminum" and "cesium," while the US "sulfur" takes the place of the British "sulphur." But chemicals which are practicable to be sold in bulk within many countries, however, still have national names, and those which do not use the Latin alphabet cannot be expected to use the IUPAC name. According to IUPAC, the full name of an element is not capitalized, even if it is derived from a proper noun (unless it would be capitalized by some other rule, for instance if it begins a sentence). And in the second half of the twentieth century physics laboratories became able to produce nuclei of chemical elements that have too quick a decay rate to ever be sold in bulk. These are also named by IUPAC, which generally adopts the name chosen by the discoverer. This can lead to the controversial question of which research group actually discovered an element, a question which delayed the naming of elements with atomic number of 104 and higher for a considerable time. (See element naming controversy). Precursors of such controversies involved the nationalistic namings of elements in the late nineteenth century (e.g., as "lutetium" refers to Paris, France, the Germans were reticent about relinquishing naming rights to the French, often calling it "cassiopeium"). And notably, the British discoverer of "niobium" originally named it "columbium," after the New World, though this did not catch on in Europe. The Americans had to accept the international name just when it was becoming an economically important material late in the twentieth century.

Chemical symbols

Specific chemical elements

Before chemistry became a science, alchemists had designed arcane symbols for both metals and common compounds. These were however used as abbreviations in diagrams or procedures; there was no concept of one atoms combining to form molecules. With his advances in the atomic theory of matter, John Dalton devised his own simpler symbols, based on circles, which were to be used to depict molecules. These were superseded by the current typographical system in which chemical symbols are not used as mere abbreviations though each consists letters of the Latin alphabet - they are symbols intended to be used by peoples of all languages and alphabets. The first of these symbols were intended to be fully international, for they were based on the Latin abbreviations of the names of metals: Fe comes from Ferrum; Ag from Argentum. The symbols were not followed by a period (full stop) as abbreviations were. Besides a name, later chemical elements are also given a unique chemical symbol, based on the name of the element, not necessarily derived from the colloquial English name. (e.g., sodium has chemical symbol 'Na' after the Latin natrium). The same applies to "W" (wolframium) for Tungsten , "Hg" (Hydrargyrum) for mercury and "K" for potassium. Stricly taken, a symbol like Tu for tungsten or M or Me for mercury seems to be more logical. Chemical symbols are understood internationally when element names might need to be translated. There are sometimes differences; for example, the Germans have used "J" instead of "I" for iodine, so the character would not be confused with a roman numeral. The first letter of a chemical symbol is always capitalized, as in the preceding examples, and the subsequent letters, if any, are always minuscule (small letters).

General chemical symbols

There are also symbols for series of chemical elements, for comparative formulas. These are one capital letter in length, and the letters are reserved so they are not permitted to be given for the names of specific elements. For example, an "X" is used to indicate a variable group amongst a class of compounds (though usually a halogen), while "R" is used for a radical (not to be confused with radical_(chemistry), meaning a compound structure such as a hydrocarbon chain. The letter "Q" is reserved for "heat" in a chemical reaction. "Y" is also often used as a general chemical symbol, although it is also the symbol of Yttrium. "Z" is also frequently used as a general variable group. "L" is used to represent a general ligand in inorganic and organometallic chemistry. "M" is also often used in place of a general metal.

Nonelement symbols

Nonelements, especially in organic and organometallic chemistry, often acquire symbols which are inspired by the elemental symbols. A few examples: Cy - cyclohexyl; Ph - phenyl; Bz - benzoyl; Bn - benzyl; Cp - Cyclopentadiene; Pr - propyl; Me - methyl; Et - ethyl; Tf - triflate; Ts - tosyl.

External links

- [http://www.vanderkrogt.net/elements/ Elementymology & Elements Multidict] word history and language dictionary

Chemical information

- [http://www.webelements.com/ WebElements]
- [http://www.vcs.ethz.ch/chemglobe/ptoe/ ChemGlobe]
- [http://pearl1.lanl.gov/periodic/default.htm Los Alamos National Laboratory]
- [http://www.chemicalelements.com/ ChemicalElements] ko:화학 원소 ms:Unsur kimia ja:元素 simple:Element th:ธาตุเคมี

Inert

In English, to be inert is to be in a state of doing little or nothing. In chemistry, the term inert is used to describe something that is not chemically active. The noble gases were described as being inert because they did not react with the other elements or themselves. Now scientists know that the reason that inert gases are completely inert to basic chemical reactions (such as combustion), for example, is because their outer valence shell is completely filled with electrons. With a filled outer valence shell, an inert atom is not able to acquire or lose an electron, and is therefore not able to participate in any chemical reactions. Molecules such as elemental nitrogen, is inert under room conditions and exists as a diatomic molecule, N₂. The inertness of nitrogen is due to the presence of the very strong triple covalent bond in the N₂ molecule. For inert atoms or molecules, a lot of energy is involved before it can combine with other elements to form compounds. A high temperature and pressure is necessary, and sometimes requires the presence of a catalyst.

Earth's atmosphere

Earth's atmosphere is a layer of gases surrounding the planet Earth and retained by the Earth's gravity. It contains roughly 78% nitrogen and 21% oxygen, with trace amounts of other gases. The atmosphere protects life on Earth by absorbing ultraviolet solar radiation and reducing temperature extremes between day and night. The atmosphere has no abrupt cut-off. It slowly becomes thinner and fades away into space. There is no definite boundary between the atmosphere and outer space. Three-quarters of the atmosphere's mass is within 11 km of the planetary surface. In the United States, persons who travel above an altitude of 50.0 miles (80.5 km) are designated as astronauts. An altitude of 120 km (75 mi or 400,000 ft) marks the boundary where atmospheric effects become noticeable during re-entry. The Karman line, at 100 km (62 mi), is also frequently used as the boundary between atmosphere and space.

Temperature and the atmospheric layers

The temperature of the Earth's atmosphere varies with altitude; the mathematical relationship between temperature and altitude varies between the different atmospheric layers:
- troposphere: From the Greek word tropos meaning to turn or mix. The troposphere is the lowest layer of the atmosphere starting at the surface going up to between 7 km at the poles and 17 km at the equator with some variation due to weather factors. The troposphere has a great deal of vertical mixing due to solar heating at the surface. This heating warms air masses, which then rise to release latent heat as sensible heat that further buoys the air mass. This process continues until all water vapor is removed. In the troposphere, on average, temperature decreases with height due to expansive cooling.
- stratosphere: from that 7–17 km range to about 50 km, temperature increasing with height.
- mesosphere: from about 50 km to the range of 80 km to 85 km, temperature decreasing with height.
- thermosphere: from 80–85 km to 640+ km, temperature increasing with height. The boundaries between these regions are named the tropopause, stratopause, and mesopause. The average temperature of the atmosphere at the surface of earth is 14 °C.

Various atmospheric regions

Atmospheric regions are also named in other ways:
- ionosphere — the region containing ions: approximately the mesosphere and thermosphere up to 550 km.
- exosphere — above the ionosphere, where the atmosphere thins out into space.
- magnetosphere — the region where the Earth's magnetic field interacts with the solar wind from the Sun. It extends for tens of thousands of kilometers, with a long tail away from the Sun.
- ozone layer — or ozonosphere, approximately 10 - 50 km, where stratospheric ozone is found. Note that even within this region, ozone is a minor constituent by volume.
- upper atmosphere — the region of the atmosphere above the mesopause.
- Van Allen radiation belts — regions where particles from the Sun become concentrated.

Pressure

:Barometric Formula: (used for airplane flight) barometric formula :Main article: Atmospheric pressure :Nasa mathematical model: NRLMSISE-00 Atmospheric pressure is a direct result of the weight of the air. This means that air pressure varies with location and time because the amount (and weight) of air above the earth varies with location and time. Atmospheric pressure drops by ~50% at an altitude of about 5 km (equivalently, about 50% of the total atmospheric mass is within the lowest 5 km). The average atmospheric pressure, at sea level, is about 101.3 kilopascals (about 14.7 pounds per square inch).

Thickness of the atmosphere

The atmosphere is present to heights of 1000 km. or more. But at this height it is so thin and at such low pressure that it's almost like it isn't there.
- 57.8% of the atmosphere is below the summit of Mount Everest.
- 72% of the atmosphere is below the common flight height of airplanes, (about 10000 m or 32800 ft).
- 99.99999% of the atmosphere is below the highest X-15 plane flight on August 22, 1963 which reached an altitude of 354,300 ft or 108 km. Therefore, most of the atmosphere is below 100 km (99.9999%) although in the rarified region above this there are auroras, and other atmospheric effects.

Composition

aurora
Source for figures above: [http://nssdc.gsfc.nasa.gov/planetary/factsheet/earthfact.html NASA]
Carbon dioxide and methane updated (to 1998) by IPCC TAR table 6.1 [http://www.grida.no/climate/ipcc_tar/wg1/221.htm]

Minor components of air not listed above include:
- The mean molecular mass of air is 28.97 g/mol.

Heterosphere

Below an altitude of about 100 km, the Earth's atmosphere has a more-or-less uniform composition (apart from water vapor) as described above. However, above about 100 km, the Earth's atmosphere begins to have a composition which varies with altitude. This is essentially because, in the absence of mixing, the density of a gas falls off exponentially with increasing altitude, but at a rate which depends on the molecular mass. Thus higher mass constituents, such as oxygen and nitrogen, fall off more quickly than lighter constituents such as helium, molecular hydrogen, and atomic hydrogen. Thus there is a layer, called the heterosphere, in which the earth's atmosphere has varying composition. As the altitude increases, the atmosphere is dominated successively by helium, molecular hydrogen, and atomic hydrogen. The precise altitude of the heterosphere and the layers it contains varies significantly with temperature.[http://www.oma.be/BIRA-IASB/Public/Research/Thermo/Thermotxt.en.html]

Density and mass

The density of air at sea level is about 1.2 kg/m³. Natural variations of the barometric pressure occur at any one altitude as a consequence of weather. This variation is relatively small for inhabited altitudes but much more pronounced in the outer atmosphere and space due to variable solar radiation The atmospheric density decreases as the altitude increases. This variation can be approximately modeled using the barometric formula. More sophisticated models are used by meteorologists and space agencies to predict weather and orbital decay of satellites. The total mass of the atmosphere is about 5.1 × 10¹⁸ kg, or about 0.9 ppm of the Earth's total mass. The above composition percentages are done by volume. Assuming that the gases act like ideal gases, we can add the percentages p multiplied by their molar masses m, to get a total t = sum (p·m). Any element's percent by mass is then p·m/t. When we do this to the above percentages, we get that, by mass, the composition of the atmosphere is 75.523% N₂, 23.133% O₂, 1.288% Ar, 0.053% CO₂, 0.001267% Ne, 0.00029% CH₄, 0.00033% Kr, 0.000724% He, and 0.0000038 % H₂. ppm This graph is from the NRLMSISE-00 atmosphere model, which has as inputs: latitude, longitude, date, time of day, altitude, solar flux, and the earth's magnetic field daily index.

The evolution of the Earth's atmosphere

model The history of the Earth's atmosphere prior to one billion years ago is poorly understood, but the following presents a plausible sequence of events. This remains an active area of research. The modern atmosphere is sometimes referred to as Earth's "third atmosphere", in order to distinguish the current chemical composition from two notably different previous compositions. The original atmosphere was primarily helium and hydrogen. Heat (from the still-molten crust, and the sun) dissipated this atmosphere. About 3.5 billion years ago, the surface had cooled enough to form a crust, still heavily populated with volcanoes which released steam, carbon dioxide, and ammonia. This led to the "second atmosphere", which was primarily carbon dioxide and water vapor, with some nitrogen but virtually no oxygen (though very recent simulations run at the University of Waterloo and University of Colorado in 2005 suggested that it may have had up to 40% hydrogen [http://newsrelease.uwaterloo.ca/news.php?id=4348]). This second atmosphere had approximately 100 times as much gas as the current atmosphere. It is generally believed that the greenhouse effect, caused by high levels of carbon dioxide, kept the Earth from freezing. During the next few billion years, water vapor condensed to form rain and oceans, which began to dissolve carbon dioxide. Approximately 50% of the carbon dioxide would be absorbed into the oceans. One of the earliest types of bacteria were the cyanobacteria. Fossil evidence indicates that these bacteria existed approximately 3.3 billion years ago and were the first oxygen-producing evolving phototropic organisms. They were responsible for the initial conversion of the earth’s atmosphere from an anoxic state to an oxic state (that is, from a state without oxygen to a state with oxygen). Being the first to carry out oxygenic photosynthesis, they were able to convert carbon dioxide into oxygen, playing a major role in oxygenating the atmosphere. Photosynthesizing plants would later evolve and convert more carbon dioxide into oxygen. Over time, excess carbon became locked in fossil fuels, sedimentary rocks (notably limestone), and animal shells. As oxygen was released, it reacted with ammonia to create nitrogen; in addition, bacteria would also convert ammonia into nitrogen. As more plants appeared, the levels of oxygen increased significantly, while carbon dioxide levels dropped. At first the oxygen combined with various elements (such as iron), but eventually oxygen accumulated in the atmosphere, resulting in mass extinctions and further evolution. With the appearance of an ozone layer (ozone is an allotrope of oxygen) lifeforms were better protected from ultraviolet radiation. This oxygen-nitrogen atmosphere is the "third atmosphere".

References

- [http://www.oma.be/BIRA-IASB/Public/Research/Thermo/Thermotxt.en.html The thermosphere: a part of the heterosphere], by J. Vercheval (viewed 1 Apr 2005)

External links

- [http://nssdc.gsfc.nasa.gov/space/model/models_home.html#atmo NASA atmosphere models]
- [http://nssdc.gsfc.nasa.gov/planetary/factsheet/earthfact.html NASA's Earth Fact Sheet]
- [http://atmospheres.agu.org/ American Geophysical Union: Atmospheric Sciences]
- [http://www.srh.noaa.gov/srh/jetstream/atmos/layers.htm Layers of the Atmosphere] Category:Atmospheric sciences Category:Atmosphere Category:Environments ko:대기권 ms:Atmosfera ja:大気

Amino acids

In chemistry, an amino acid is any molecule that contains both amino and carboxylic acid functional groups. In biochemistry, this shorter and more general term is frequently used to refer to alpha amino acids: those amino acids in which the amino and carboxylate functionalities are attached to the same carbon, the so-called α–carbon. An amino acid residue is what is left of an amino acid once a molecule of water has been lost (an H⁺ from the nitrogenous side and an OH^- from the carboxylic side) in the formation of a peptide bond.

Overview

Amino acids are the basic structural building units of proteins. They form short polymer chains called peptides or polypeptides which in turn form structures called proteins. protein Twenty amino acids are encoded by the standard genetic code and are called proteinogenic or standard amino acids. At least two others are also coded by DNA in a non-standard manner as follows:
- Selenocysteine is incorporated into some proteins at a UGA codon, which is normally a stop codon.
- Pyrrolysine is used by some methanogens in enzymes that they use to produce methane. It is coded for similarity to selenocysteine but with the codon UAG instead. Other amino acids contained in proteins are usually formed by post-translational modification, that is modification after translation (protein synthesis). These modifications are often essential for the function of the protein. Proline is the only proteinogenic amino acid whose side group is cyclic and links to the a-amino group, forming a secondary amino group. Formerly, proline was misleadingly called an imino acid. Over one hundred amino acids have been found in nature. Some of these have been detected in meteorites, especially in a type known as carbonaceous chondrites. Microorganisms and plants often produce very uncommon amino acids, which can be found in peptidic antibiotics (e.g., nisin or alamethicin). Lanthionine is a sulfide-bridged alanine dimer which is found together with unsaturated amino acids in lantibiotics (antibiotic peptides of microbial origin). 1-Aminocyclopropane-1-carboxylic acid (ACC) is a small disubstituted cyclic amino acid and a key intermediate in the production of the plant hormone ethylene. In addition to protein synthesis, amino acids have other biologically-important roles. Glycine and glutamate are neurotransmitters as well as standard amino acids in proteins. Many amino acids are used to synthesize other molecules, for example:
- tryptophan is a precursor of the neurotransmitter serotonin);
- glycine is one of the reactants in the synthesis of porphyrins such as heme. Numerous non-standard amino acids are also biologically-important: GABA (another neurotransmitter), carnitine (used in lipid transport within a cell), ornithine, citrulline, homocysteine, hydroxyproline, hydroxylysine, and sarcosine. Some of the 20 standard amino acids are called essential amino acids because they cannot be synthesized by the body from other compounds through chemical reactions, but instead must be taken in with food. In humans, the essential amino acids are lysine, leucine, isoleucine, methionine, phenylalanine, threonine, tryptophan, valine, and (in children) histidine and arginine. The phrase "branched-chain amino acids" is sometimes used to refer to the aliphatic amino acids: leucine, isoleucine and valine.

General structure

The general structure of proteinogenic alpha amino acids is: COOH | H-C-R | NH₂ Where "R" represents a side chain specific to each amino acid. Amino acids are usually classified by properties of the side chain into four groups: acidic, basic, hydrophilic (polar), and hydrophobic (nonpolar).

Isomerism

Except for glycine, where R = H, amino acids occur in two possible optical isomers, called D and L. The L amino acids represent the vast majority of amino acids found in proteins. D amino acids are found in some proteins produced by exotic sea-dwelling organisms, such as cone snails. They are also abundant components of the cell walls of bacteria.

Reactions

Proteins are created by polymerization of amino acids by peptide bonds in a process called translation. translation Peptide bond formation
1. Amino acid; 2, zwitterion structure; 3, two amino acids forming a peptide bond. (See also bond.)

List of standard amino acids

Structures

Structures and symbols of the 20 amino acids present in genetic code. Image:amino_acids_2.png

Chemical properties

Following is a table listing the one letter symbols, the three-letter symbols, and the chemical properties of the side chains of the standard amino acids. The mass listed is the weighted average of all common isotopes, and includes the mass of H₂O. The one-letter symbol for an undetermined amino acid is X. The three-letter symbol Asx or one-letter symbol B means the amino acid is either asparagine or aspartic acid, whereas Glx or Z means either glutamic acid or glutamine.

Hydrophilic and hydrophobic amino acids

Depending on how polar the side chain, aminoacids can be hydrophilic or hydrophobic to various degree. This influences their interaction with other structures, both within the protein itself and within other proteins. The distribution of hydrophilic and hydrophobic aminoacids determines the tertiary structure of the protein, and their physical location on the outside structure of the proteins influences their quaternary structure. For example, soluble proteins have surfaces rich with polar aminoacids like serine and threonine, while integral membrane proteins tend to have outer ring of hydrophobic aminoacids that anchors them to the lipid bilayer, and proteins anchored to the membrane have a hydrophobic end that locks into the membrane. Similarly, proteins that have to bind to positive-charged molecules have surfaces rich with negatively charged aminoacids like glutamate and aspartate, while proteins binding to negative-charged molecules have surfaces rich with positively charged chains like lysine and arginine. Hydrophilic and hydrophobic interactions of the proteins do not have to rely only on aminoacids themselves. By various posttranslational modifications other chains can be attached to the proteins, forming hydrophobic lipoproteins or hydrophylic glycoproteins.

Nonstandard amino acids

Aside from the twenty standard amino acids, there is a vast number of nonstandard amino acids not used in the body's regular manufacturing of proteins. Examples of nonstandard amino acids include the sulfur-containing taurine and the neurotransmitters GABA and dopamine. Other examples are lanthionine, 1-amino isobutyric acid, dehydroalanine, dehydro-amino-butyric acid, Nonstandard amino acids are usually formed through modifications to standard amino acids. For example, taurine can be formed by the decarboxylation of cysteine, while dopamine is synthesized from tyrosine and hydroxyproline is made by a posttranslational modification from proline.

Uses of substances derived from amino acids

- Aspartame (aspartyl-phenylalanine-1-methyl ester) is an artificial sweetener.
- 5-HTP (5-hydroxytryptophan) has been used to treat neurological problems associated with PKU (phenylketonuria), as well as depression (as an alternative to L-Tryptophan).
- L-DOPA (L-dihydroxyphenylalanine) is a drug used to treat Parkinsonism.
- Monosodium glutamate is a food additive to enhance flavor.

References

- Doolittle, R.F. (1989) Redundancies in protein sequences. In Predictions of Protein Structure and the Principles of Protein Conformation (Fasman, G.D. ed) Plenum Press, New York, pp. 599-623
- David L. Nelson and Michael M. Cox, Lehninger Principles of Biochemistry, 3rd edition, 2000, Worth Publishers, ISBN 1572591536
- [http://www.sciencedirect.com/science?_ob=ArticleURL&_udi=B6T36-3XB0N6H-H&_coverDate=09%2F10%2F1999&_alid=241945989&_rdoc=1&_fmt=&_orig=search&_qd=1&_cdi=4938&_sort=d&view=c&_acct=C000050221&_version=1&_urlVersion=0&_userid=10&md5=3cb10a335716303532fc517906a12b3a On the hydrophobic nature of cysteine.]

External links

- [http://micro.magnet.fsu.edu/aminoacids/index.html Molecular Expressions: The Amino Acid Collection] - Has detailed information and microscopy photographs of each amino acid.
- [http://researchnews.osu.edu/archive/aminoacd.htm 22nd amino acid] - Press release from Ohio State claiming discovery of a 22nd amino acid. Category:Amino acids Category:Nitrogen metabolism ko:아미노산 ja:アミノ酸

Ammonia

Ammonia is a compound of nitrogen and hydrogen with the formula NH₃. At standard temperature and pressure ammonia is a gas. It is toxic and corrosive to some materials, and has a characteristic pungent odor. An ammonia molecule is not flat, but has the shape of a compressed tetrahedron known as a trigonal pyramid, as would be expected from VSEPR theory. This shape gives the molecule an overall dipole moment and makes it polar so that ammonia very readily dissolves in water. The nitrogen atom in the molecule has a lone electron pair, and ammonia acts as a base. In acidic or even neutral aqueous solutions, it can bond to a hydronium ion (H₃O⁺), releasing a water molecule (H₂O) to form the positively charged ammonium ion (NH₄⁺), which has the shape of a regular tetrahedron. The degree to which ammonia forms the ammonium ion depends on the pH of the solution. The main uses of ammonia are in the production of fertilizers, explosives and polymers. It is also an ingredient in certain household glass cleaners. Ammonia is found in small quantities in the atmosphere, being produced from the putrefaction of nitrogenous animal and vegetable matter. Ammonia and ammonium salts are also found in small quantities in rainwater, while ammonium chloride (sal-ammoniac) and ammonium sulfate are found in volcanic districts; crystals of ammonium bicarbonate have been found in Patagonian guano. Ammonium salts also are found distributed through all fertile soil and in seawater. Substances containing ammonia or that are similar to it are called ammoniacal.

History

Salts of ammonia have been known from very early times; thus the term Hammoniacus sal appears in the writings of Pliny, although it is not known whether the term is identical with the more modern sal-ammoniac. In the form of sal-ammoniac, ammonia was known to the alchemists as early as the 13th century, being mentioned by Albertus Magnus. It was also used by dyers in the Middle Ages in the form of fermented urine to alter the colour of vegetable dyes. In the 15th century, Basilius Valentinus showed that ammonia could be obtained by the action of alkalis on sal-ammoniac. At a later period, when sal-ammoniac was obtained by distilling the hoofs and horns of oxen and neutralizing the resulting carbonate with hydrochloric acid, the name Spirit of hartshorn was applied to ammonia. Gaseous ammonia was first isolated by Joseph Priestley in 1774 and was termed by him alkaline air. In 1777 Karl Wilhelm Scheele showed that it contained nitrogen, and Claude Louis Berthollet, in about 1785, ascertained its composition. The Haber process to produce ammonia from the nitrogen contained in the air was developed by Fritz Haber and Carl Bosch in 1909 and patented in 1910. It was first used on an industrial scale by the Germans during World War I. The ammonia was used to produce explosives to sustain their war effort.

Production

Because of its many uses, ammonia is one of the most highly-produced inorganic chemicals. Before the start of WWI most ammonia was obtained by the dry distillation of nitrogenous vegetable and animal products; by the reduction of nitrous acid and nitrites with nascent hydrogen; and also by the decomposition of ammonium salts by alkaline hydroxides or by quicklime, the salt most generally used being the chloride (sal-ammoniac) thus ::2NH₄Cl + 2CaO → CaCl₂ + Ca(OH)₂ + 2NH₃ It has also been obtained by decomposing magnesium nitride (Mg₃N₂) with water, ::Mg₃N₂ + 6H₂O → 3Mg(OH)₂ + 2NH₃ Today the Haber process is the most important method for production of ammonia. In this process, nitrogen and hydrogen gases combine directly on an iron catalyst at a pressure of 200 bar (20 MPa, 3000 lbf/in²) and a temperature of 500 °C to produce ammonia. ::N₂ + 3H₂ → 2 NH₃ Compared to older methods, the feedstocks of the Haber process are relatively inexpensive—nitrogen makes up 78% of the atmosphere, while hydrogen can be readily produced from natural gas.

Properties

Ammonia is a colourless gas with a characteristic pungent smell; it is lighter than air, its density being 0.589 times that of air. It is easily liquefied and the liquid boils at -33.7 °C, and solidifies at -75 °C to a mass of white crystals. Liquid ammonia possesses strong ionizing powers (ε = 22), and solutions of salts in liquid ammonia have been much studied. Liquid ammonia has a very high standard enthalpy change of vaporization (23.35 kJ/mol, c.f. water 40.65 kJ/mol, methane 8.19 kJ/mol, phosphine 14.6 kJ/mol) and can therefore be used in laboratories in non-insulated vessels at room temperature, even though it is well above its boiling point. It is miscible with water. All the ammonia contained in an aqueous solution of the gas may be expelled by boiling. The aqueous solution of ammonia is basic. The maximum concentration of ammonia in water (a saturated solution) has a density of 0.880 g cm^-3 and is often known as '.880 Ammonia'. It does not sustain combustion, and it does not burn readily unless mixed with oxygen, when it burns with a pale yellowish-green flame. At high temperature and in the presence of a suitable catalyst, ammonia is decomposed into its constituent elements. Chlorine catches fire when passed into ammonia, forming nitrogen and hydrochloric acid; unless the ammonia is present in excess, the highly explosive nitrogen trichloride (NCl₃) is also formed. The ammonia molecule readily undergoes nitrogen inversion at normal pressures, that is to say that the nitrogen atom passes through the plane of the three hydrogen atoms as if it were an umbrella turning inside out in a strong wind. The energy barrier to this inversion is 24.7 kJ/mol in ammonia, and the resonance frequency is 23.79 GHz, corresponding to microwave radiation of a wavelength of 1.260 cm. The absorption at this frequency was the first microwave spectrum to be observed (C. E. Cleeton & N. H. Williams, 1934).

Formation of salts

One of the most characteristic properties of ammonia is its power of combining directly with acids to form salts; thus with hydrochloric acid it forms ammonium chloride (sal-ammoniac); with nitric acid, ammonium nitrate, etc. However perfectly dry ammonia will not combine with perfectly dry hydrogen chloride, moisture being necessary to bring about the reaction. ::NH₃ + HCl → NH₄Cl The salts produced by the action of ammonia on acids are known as the ammonium salts and all contain the ammonium ion (NH₄⁺).

Acidity

Although ammonia is well-known as a base, it can also act as an extremely weak acid. It is a protic substance, and is capable of dissociation into the amide (NH₂⁻) ion, for example when solid lithium nitride is added to liquid ammonia, forming a lithium amide solution: Li₃N(s)+ 2NH₃(l) → 3Li⁺(am) + 3NH₂⁻(am). This is a Bronsted-Lowry acid-base reaction in which ammonia is acting as an acid.

Formation of other compounds

Ammonia can act as a nucleophile in substitution reactions. Amines can be formed by the reaction of ammonia with alkyl halides, although the resulting –NH₂ group is also nucleophilic and secondary and tertiary amines are often formed as by-products. Using an excess of ammonia helps minimise multiple substitution, and neutralises the hydrogen halide formed. Methylamine is prepared commercially by the reaction of ammonia with chloromethane, and the reaction of ammonia with 2-bromopropanoic acid has been used to prepare racemic alanine in 70% yield. Ethanolamine is prepared by a a ring-opening reaction with ethylene oxide: the reaction is sometimes allowed to go further to produce diethanolamine and triethanolamine. Amides can be prepared by the reaction of ammonia with a number of carboxylic acid derivatives. Acyl chlorides are the most reactive, but the ammonia must be present in at least a two-fold excess to neutralise the hydrogen chloride formed. Esters and anhydrides also react with ammonia to form amides. Ammonium salts of carboxylic acids can be dehydrated to amides so long as there are no thermally sensitive groups present: temperatures of 150–200 °C are required. The hydrogen in ammonia is capable of replacement by metals, thus magnesium burns in the gas with the formation of magnesium nitride Mg₃N₂, and when the gas is passed over heated sodium or potassium, sodamide, NaNH₂, and potassamide, KNH₂, are formed. Where necessary in substitutive nomenclature, IUPAC recommendations prefer the name azane to ammonia: hence chloramine would be named chloroazane in substitutive nomenclature, not chloroammonia.

Ammonia as a ligand

Ammonia can act as a ligand in transition metal complexes. It is a pure σ-donor, in the middle of the spectrochemical series, and shows intermediate hard-soft behaviour. For historical reasons, ammonia is named ammine in the nomenclature of coordination compounds. Some notable ammine complexes include:
- Hexamminecopper(II), [Cu(NH₃)₆]²⁺, a characteristic dark blue complex formed by adding ammonia to solution of copper(II) salts.
- Diamminesilver(I), [Ag(NH₃)₂]⁺, the active species in Tollens' reagent. Formation of this complex can also help to distinguish between precipitates of the different silver halides: AgCl is soluble in dilute (2 M) ammonia solution, AgBr is only soluble in concentrated ammonia solution while AgI is insoluble in aqueous solution of ammonia. Ammine complexes of chromium(III) were known in the late 19th century, and formed the basis of Alfred Werner's theory of coordination compounds. Werner noted that only two isomers (fac- and mer-) of the complex [CrCl₃(NH₃)₃] could be formed, and concluded that the ligands must be arranged around the metal ion at the vertices of an octahedron. This has since been confirmed by X-ray crystallography. An ammine ligand bound to a metal ion is markedly more acidic than a free ammonia molecule, although deprotonation in aqueous solution is still rare. One example is the Calomel reaction, where the resulting amidomercury(II) compound is highly insoluble. ::Hg₂Cl₂ + 2NH₃ → Hg + HgCl(NH₂) + NH₄⁺ + Cl⁻

Uses

The most important single use of ammonia is in the production of nitric acid. A mixture of one part ammonia to nine parts air is passed over a platinum gauze catalyst at 850 °C, whereupon the ammonia is oxidized to nitric oxide. ::4NH₃ + 5O₂ → 4NO + 6H₂O The catalyst is essential, as the normal oxidation (or combustion) of ammonia gives dinitrogen and water: the production of nitric oxide is an example of kinetic control. As the gas mixture cools to 200–250 °C, the nitric oxide is in turn oxidized by the excess of oxygen present in the mixture, to give nitrogen dioxide. This is reacted with water to give nitric acid for use in the production of fertilizers and explosives. In addition to serving as a fertilizer ingredient, ammonia can also be used directly as a fertilizer by forming a solution with irrigation water, without additional chemical processing. This later use allows the continuous growing of nitrogen dependent crops such as maize (corn) without crop rotation but this type of use leads to poor soil health. Ammonia has thermodynamic properties that make it very well suited as a refrigerant, since it liquefies readily under pressure, and was used in virtually all refrigeration units prior to the advent of haloalkanes such as Freon. However, ammonia is a toxic irritant and its corrosiveness to any copper alloys increases the risk that an undesirable leak may develop and cause a noxious hazard. Its use in small refrigeration units has been largely replaced by haloalkanes, which are not toxic irritants and are practically not flammable. (Note: Butane and isobutane, which have very suitable thermodynamic properties for refrigerants, are extremely flammable.) Ammonia continues to be used as a refrigerant in large industrial processes such as bulk icemaking and industrial food processing. Ammonia is also useful as a component in absorption-type refrigerators, which do not use compression and expansion cycles but can exploit heat differences. Since the implication of haloalkane being major contributors to ozone depletion, ammonia is again seeing increasing use as a refrigerant. Ammonia is a primary ingredient in old-style household cleaners. It is also sometimes added to drinking water along with chlorine to form chloramine, a disinfectant. Unlike chlorine on its own, chloramine does not combine with organic (carbon containing) materials to form carcinogenic halomethanes such as chloroform.

Liquid ammonia as a solvent

:See also: Inorganic nonaqueous solvent Liquid ammonia is the best-known and most widely studied non-aqueous ionizing solvent. Its most conspicuous property is its ability to dissolve alkali metals to form highly coloured, electrically conducting solutions containing solvated electrons. Apart from these remarkable solutions, much of the chemistry in liquid ammonia can be classified by analogy with related reactions in aqueous solutions. Comparison of the physical properties of NH₃ with those of water shows that NH₃ has the lower melting point, boiling point, density, viscosity, dielectric constant and electrical conductivity; this is due at least in part to the weaker H bonding in NH₃ and the fact that such bonding cannot form cross-linked networks since each NH₃ molecule has only 1 lone-pair of electrons compared with 2 for each H₂O molecule. The ionic self-dissociation constant of liquid NH₃ at −50 °C is approx. 10^-33 mol²·l^-2.

Solubility of salts

Liquid ammonia is an ionizing solvent, although less so than water, and dissolves a range of ionic compounds including many nitrates, nitrites, cyanides and thiocyanates. Most ammonium salts are soluble, and these salts act as acids in liquid ammonia solutions. The solubility of halide salts increases from fluoride to iodide. A saturated solution of ammonium nitrate contains 0.83 mol solute per mole of ammonia, and has a vapour pressure of less than 1 bar even at 25 °C.

Solutions of metals

:See also: Solvated electron, metallic solution Liquid ammonia will dissolve the alkali metals and other electropositive metals such as Ca, Sr, Ba Eu and Yb. At low concentrations (< 0.06 mol/L), deep blue solutions are formed: these contain metal cations and solvated electrons, free electrons which are surrounded by a cage of ammonia molecules. These solutions are very useful as strong reducing agents. At higher concentrations, the solutions are metallic in appearance and in electrical conductivity. At low temperatures, the two types of solution can coexist as immiscible phases.

Redox properties of liquid ammonia

The range of thermodynamic stability of liquid ammonia solutions is very narrow, as the potential for oxidation to dinitrogen, E° (N₂ + 6NH₄⁺ + 6e⁻ 8NH₃), is only +0.04 V. In practice, both oxidation to dinitrogen and reduction to dihydrogen are slow. This is particularly true of reducing solutions: the solutions of the alkali metals mentioned above are stable for several days, slowly decomposing to the metal amide and dihydrogen. Most studies involving liquid ammonia solutions are done in reducing conditions: although oxidation of liquid ammonia is usually slow, there is still a risk of explosion, particularly if transition metal ions are present as possible catalysts.

Detection and determination

Ammonia and ammonium salts can be readily detected, in very minute traces, by the addition of Nessler's solution, which gives a distinct yellow coloration in the presence of the least trace of ammonia or ammonium salts. Sulfur sticks are burnt to detect small leaks in industrial ammonia refrigeration systems. Larger quantities can be detected by warming the salts with a caustic alkali or with quicklime, when the characteristic smell of ammonia will be at once apparent. The amount of ammonia in ammonium salts can be estimated quantitatively by distillation of the salts with sodium or potassium hydroxide, the ammonia evolved being absorbed in a known volume of standard sulfuric acid and the excess of acid then determined volumetrically; or the ammonia may be absorbed in hydrochloric acid and the ammonium chloride so formed precipitated as ammonium hexachloroplatinate, (NH₄)₂PtCl₆.

Safety precautions

Toxicity

The toxicity of ammonia solutions does not usually cause problems for humans and other mammals, as a specific mechanism exists to prevent its build-up in the bloodstream. Ammonia is converted to carbamoyl phosphate by the enzyme carbamoyl phosphate synthase, and then enters the urea cycle to be either incorporated into amino acids or excreted in the urine. However fish and amphibians lack this mechanism, as they can usually eliminate ammonia from their bodies by direct excretion. Ammonia even at dilute concentrations is highly toxic to aquatic animals, and for this reason it is classified as dangerous for the environment.

Household use

Solutions of ammonia (5–10% by weight) are used as household cleaners, particularly for glass. These solutions are irritating to the eyes and mucous membranes (respiratory and digestive tracts), and to a lesser extent the skin. They should never be mixed with chlorine-containing products, for example household bleach, as a variety of toxic and carcinogenic compounds are formed (e.g., chloramine, hydrazine).

Laboratory use of ammonia solutions

The hazards of ammonia solutions depend on the concentration: "dilute" ammonia solutions are usually 5–10% by weight (<5.62 mol/L); "concentrated" solutions are usually prepared at >25% by weight. A 25% (by weight) solution has a density of 0.907 g/cm³, and a solution which has a lower density will be more concentrated. The European Union classification of ammonia solutions is given in the table. :S-Phrases: , , , , . The ammonia vapour from concentrated ammonia solutions is severely irritating to the eyes and the respiratory tract, and these solutions should only be handled in a fume hood. Saturated ("0.880") solutions can develop a significant pressure inside a closed bottle in warm weather, and the bottle should be opened with care: this is not usually a problem for 25% ("0.900") solutions. Ammonia solutions should not be mixed with halogens, as toxic and/or explosive products are formed. Prolonged contact of ammonia solutions with silver, mercury or iodide salts can also lead to explosive products: such mixtures are often formed in qualitative analysis, and should be acidified and diluted before disposal once the test is completed.

Laboratory use of anhydrous ammonia (gas or liquid)

Anhydrous ammonia is classified as toxic (T) and dangerous for the environment (N). The gas is flammable (autoignition temperature: 651 °C) and can form explosive mixtures with air (16–25%). The permissible exposure limit (PEL) in the United States is 50 ppm (35 mg/m³), while the IDLH concentration is estimated at 300 ppm. Repeated exposure to ammonia lowers the sensitivity to the smell of the gas: normally the odour is detectable at concentrations of less than 0.5 ppm, but desensitized individuals may not detect it even at concentrations of 100 ppm. Ammonia reacts violently with the halogens, and causes the explosive polymerization of ethylene oxide. It also forms explosive compounds with compounds of gold, silver, mercury, germanium or tellurium, and with stibine. Violent reactions have also been reported with acetaldehyde, hypochlorite solutions, potassium ferricyanide and peroxides. Anhydrous ammonia corrodes copper- and zinc-containing alloys, and so brass fittings should not be used for handling the gas. Liquid ammonia can also attack rubber and certain plastics.

Reference

# Baker, H. B. (1894). J. Chem. Soc. 65: 612.

Bibliography

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External links

- [http://www.ilo.org/public/english/protection/safework/cis/products/icsc/dtasht/_icsc04/icsc0414.htm International Chemical Safety Card 0414] (anhydrous ammonia)
- [http://www.ilo.org/public/english/protection/safework/cis/products/icsc/dtasht/_icsc02/icsc0215.htm International Chemical Safety Card 0215] (aqueous solutions)
- [http://www.cdc.gov/niosh/npg/npgd0028.html NIOSH Pocket Guide to Chemical Hazards]
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- [http://www.inrs.fr Institut national de recherche et de securite] (in French)
- [http://www.ammoniaspills.org Emergency Response to Ammonia Fertilizer Releases (Spills)] for the Minnesota Department of Agriculture
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- [http://www.compchemwiki.org/index.php?title=Ammonia Computational Chemistry Wiki] Category:Nitrogen compounds Category:Hydrides Category:Bases Category:Nitrogen metabolism Category:Household chemicals Category:Refrigerants ms:Ammonia ja:アンモニア simple:Ammonia

Cyanide

A cyanide is any chemical compound that contains the cyano group C≡N, with the carbon atom triple-bonded to the nitrogen atom. Inorganic cyanides contain the highly-toxic cyanide ion CN^-, and are the salts of the acid hydrogen cyanide (HCN). Organic cyanides contain the cyano group single-bonded to another carbon atom, and are also known as nitriles. The cyano group is also known as the nitrile group.

Appearance and odour

Hydrogen cyanide is a colorless gas with a faint, bitter, almond-like odour. Nearly 40% of the population is unable to smell cyanide at all because they lack the necessary gene for smelling the odour. Sodium cyanide (NaCN) and potassium cyanide (KCN) are both white solids or powder with a bitter, almond-like odour in damp air.

Occurrence and use

Cyanides can be produced by certain bacteria, fungi, and algae, and are found in a number of foods and plants. In plants, cyanides are usually bound to sugar molecules in the form of cyanogenic glycosides. Cassava roots (aka manioc), an important potato-like food grown in tropical countries, contain cyanogenic glycosides and must be processed prior to consumption (usually by extended boiling). Fruits that have a pit, such as cherries and apricots, often contain either cyanides or cyanogenic glycosides in the pit. Apple seeds do as well. Bitter almonds, from which almond oil and flavouring is made, also contain a cyanogenic glycoside, amygdalin. Hydrogen cyanide is contained in vehicle exhaust and in tobacco smoke. Because the smoke of some burning plastics contains hydrogen cyanide, house fires often result in cyanide poisonings of the inhabitants. A deep blue pigment called Prussian blue, used in the making of blueprints, is iron (III) ferrocyanide (hence the name cyanide, from cyan, a shade of blue). It produces hydrogen cyanide when exposed to acids. Gold and silver cyanides are among the very few soluble forms of the metals, and thus are used in mining, electroplating, metallurgy and jewelry for chemical gilding, buffing, and extraction of gold. (See also below under Mining.) Furthermore, cyanides and hydrogen cyanide are used in the production of chemicals including plastics, photographic development, and fumigating ships. In organic synthesis, cyanides are often used to lengthen the carbon chain: RX + CN^- → RCN (Nucleophilic Substitution) # RCN + H₃O⁺ → RCOOH (Hydrolysis) # RCN + LiAlH₄ → RCH₂NH₂ (under reflux in dry ether, followed by addition of H₃O⁺) # RCN + NaBH₄ → RCOH In the above cases, the number of carbon atoms of the main chain is increased by one. Potassium ferrocyanide is used to achieve a blue colour on cast bronze sculptures during the final finishing stage of the sculpture. On its own, it will produce a very dark shade of blue and is often mixed with other chemicals to achieve the desired tint and hue. It is applied using a torch and paint brush while wearing the standard safety equipment used for any patina application; rubber gloves, safety glasses and a respirator. The actual amount of cyanide in the mixture varies according to the recipes used by each foundry. Two cyanide ions can bond to each other via their carbon atoms, forming the gas cyanogen (NC-CN).

Effects on the human body

To deal with the cyanides contained in many foods, the body has an enzyme (rhodanide synthetase) which can convert small amounts of cyanides to the harmless sulfur-containing thiocyanate (SCN⁻). The cyanide ion is also a component of vitamin B₁₂, where it is one of the ligands for the cobalt ion. In larger amounts, cyanides are harmful. Symptoms of moderate poisoning include vomiting, convulsions, deep breathing, shortness of breath and anxiety; more serious cases result in convulsions, loss of consciousness, and death after apnea and cardiac arrest due to hypoxemia. The lethal dose for adults is 200–300 mg of potassium or sodium cyanide, or 50 mg of hydrogen cyanide. When inhaled, HCN is lethal within a few minutes at concentrations of 300 ppm. Exposure to lower levels of cyanide over a long period (e.g., after use of cassava roots as a primary food source in tropical Africa) results in increased blood cyanide levels. These may result in weakness of the fingers and toes, difficulty walking, dimness of vision, deafness, and decreased thyroid gland function, but chemicals other than cyanide may contribute to these effects. Skin contact with cyanide can produce irritation and sores. It is not known whether cyanides can directly cause birth defects in people. Birth defects were seen in rats that ate diets of cassava roots. Effects on the reproductive system were seen in rats and mice that drank water containing sodium cyanide. There are medical tests to measure blood and urine levels of cyanide; however, small amounts of cyanide are not always detectable in blood and urine. Tissue levels of cyanide can be measured if cyanide poisoning is suspected, but cyanide is rapidly cleared from the body, so the tests must be done soon after the exposure. An almond-like odor in the breath may alert a doctor that a person was exposed to cyanide.

Mechanism of toxicity and treatment

The cyanide ion kills all aerobic organisms by shutting down cellular respiration. It interrupts the electron transport chain in the inner membrane of the mitochondrion because it binds more strongly than oxygen to the Fe³⁺ (ferric iron ion) in cytochrome a₃, preventing this cytochrome from combining electrons with oxygen. Contrary to popular belief, cyanide does not bind well to ferrous hemoproteins, such as hemoglobin, the mechanism that makes carbon monoxide toxic. One of the therapies for cyanide poisoning is to convert part of the hemoglobin of the blood from ferrous (2+ charge) hemoglobin to ferric (3+ charge); this creates a pool of binding potential that can divert cyanide from the cytochromes it poisons. This is done with the compound 4-dimethylaminophenyl (4-DMAP), or with sodium nitrite. The United States standard cyanide antidote kit uses a small inhaled dose of amyl nitrite followed by intravenous sodium nitrite. This converts a portion of the hemoglobin's iron from ferrous iron to ferric iron, converting the hemoglobin into methemoglobin. Cyanide is more strongly drawn to methemoglobin than to the cytochrome oxidase of the cells, effectively pulling the cyanide off the cells and onto the methemoglobin. Once bound with the cyanide, the methemoglobin becomes cyanmethemoglobin. Therapy with nitrites is not innocuous. The doses given to an adult can potentially cause a fatal methemoglobinemia in children or may cause profound hypotension. Treatment of children affected with cyanide intoxication must be individualized and is based upon their body weight and hemoglobin concentration. The next part of the cyanide antidote kit is sodium thiosulfate, which is administered intravenously. The sodium thiosulfate and cyanmethemoglobin become thiocyanate, releasing the hemoglobin, and the thiocyanate is excreted by the kidneys. Alternative methods of treating cyanide intoxication are used in other countries. For example, the method in France is to use hydroxycobalamin (a form of vitamin B₁₂), which combines with cyanide to form the harmless vitamin B_12a cyanocobalamin. Cyanocobalamin is eliminated through the urine. Hydroxycobalamin works both within the intravascular space and within the cells to combat cyanide intoxication. This contrasts with methemoglobin, which acts only within the vascular space as an antidote. Administration of sodium thiosulfate improves the ability of the hydroxycobalamin to detoxify cyanide poisoning. This treatment is considered so effective and innocuous that it is administered routinely in Paris to victims of smoke inhalation to detoxify any associated cyanide intoxication. However it is relatively expensive and not universally available. 4-Dimethylaminophenol (4-DMAP) has been proposed in Germany as a more rapid antidote than nitrites and with (reportedly) lower toxicity. It is used currently by the German military and by the civilian population. In humans, intravenous injection of 3 mg/kg of 4-DMAP will produce 35% methemoglobin levels within 1 minute. Cobalt salts have also been demonstrated as effective in binding cyanide. One current cobalt-based antidote available in Europe is dicobalt-EDTA, sold as Kelocyanor®. This agent chelates cyanide as the cobalticyanide. This drug provides an antidote effect more quickly than formation of methemoglobin, but a clear superiority to methemoglobin formation has not been demonstrated. Cobalt complexes are quite toxic, and there have been accidents reported in the UK where patients have been given dicobalt-EDTA by mistake based on a false diagnoses of cyanide poisoning.

Use as a poison

The cyanide ion, if used as poison, is generally delivered in the form of gaseous hydrogen cyanide or in the form of potassium cyanide (KCN) or sodium cyanide (NaCN). Zyklon B, the poison gas used in Nazi gas chambers during the Holocaust, works by delivering hydrogen cyanide gas. Cyanide is also the compound used in U.S. gas chambers for execution. Cyanide salts are sometimes used as fast-acting suicide devices. When they reach the stomach acids, cyanide ions are released; therefore they work faster on an empty stomach. Famous cyanide salt suicides include:
- Erwin Rommel
- Adolf Hitler (likely, see article on Hitler's death)
- Joseph Goebbels
- Hermann Göring
- Heinrich Himmler
- Alan Turing
- Odilo Globocnik
- People's Temple mass suicide in Jonestown
- Grigori Rasputin (used in attempting to kill) Cyanides were stockpiled in both the Soviet and the United States chemical weapons arsenals in the 1950s and 1960s. During the Cold War, the Soviet Union was thought to be planning to use hydrogen cyanide as a "blitzkrieg" weapon to clear a path through the opposing front line, knowing that the harmful gas itself would dissipate and allow unprotected access to the captured zone. However, as a military agent, cyanide was not considered very effective, since cyanide is lighter than air and requires a significant dose in order to incapacitate or kill. Some spy agents also carried glasses with cyanide in the frames. If they were caught by the enemy, they could casually chew the frame, releasing the cyanide, and die before being tortured or getting information taken from them.

In fiction

Poisoning by cyanide also figures prominently in crime fiction, for example Agatha Christie's Sparkling Cyanide (also entitled Remembered Death); cyanide is the instrument of one murder in The Big Sleep by Raymond Chandler. See also: Victims of poisoning.

Mining

Cyanide salts are used in silver and gold mining, called the cyanide process. The high-grade ore is finely ground and mixed with the cyanide solution (concentration of about two kilogram NaCN per tonne (1000 kilograms)); low-grade ores are stacked into heaps and sprayed with cyanide solution (concentration of about one kilogram NaCN per tonne (1000 kilograms)). The precious-metal cations bind to the cyanide anions and form a soluble cyanide. The pregnant liquor is separated from the leftover dirt, which is discarded to a tailing pond or spent (the recoverable gold having been removed) heap. The metal is recovered from the pregnant solution with zincdust that replaces the gold in solution or by absorption onto activated carbon. This process can result in environmental and health problems. Cyanide is highly reactive; it decomposes rapidly in sunlight. It can mobilize some heavy metals like mercury (if mercury is present). Gold can also be associated with arsenopyrite, which is similar to iron pyrite (fool's gold), with iron atoms replaced by arsenic.

Fishing

Cyanides are used to capture live fish near coral reefs for the aquarium and seafood markets. This illegal fishing occurs mainly in the Philippines, Indonesia and the Caribbean to supply the 2 million marine aquarium owners in the world. In this method, a diver uses a large, needleless syringe to squirt a cyanide solution into areas where the fish are hiding, stunning them so that they can be easily gathered. Many fish caught in this fashion die immediately, or in shipping. Those that survive to find their way into pet stores often die from shock, or from massive digestive damage. The high concentrations of cyanide on reefs so harvested has also resulted in cases of cyanide poisioning among local fishermen and their families. Environmental organizations decry the practice, as do responsible aquarists and aquarium dealers. To prevent the trade of illegally-caught aquarium fish, the Marine Aquarium Council has created a certification in which the tropical fish are caught legally with nets only. To ensure authenticity, "MAC-Certified marine organisms bear the MAC-Certified label on the tanks and boxes in which they are kept and shipped." [http://www.aquariumcouncil.org/subpage.asp?section=13 MAC Certification].

External links

- [http://www.inchem.org/documents/cicads/cicads/cicad61.htm CICAD 61: Hydrogen cyanide and cyanides]
- [http://www.inchem.org/documents/antidote/antidote/ant02.htm#SubSectionNumber:1.13.1 IPCS/CEC Evaluation of antidotes for poisoning by cyanides]
- [http://www.atsdr.cdc.gov/MHMI/mmg8.html ATSDR medical management guidelines for cyanide poisoning (US)]
- [http://www.hse.gov.uk/pubns/misc076.htm HSE recommendations for first aid treatment of cyanide poisoning (UK)]
- [http://agmed.sante.gouv.fr/htm/10/piratox4/piratox4.pdf Fiche piratox de prise en charge thérapeutique (France)] (in French)
- [http://www.aquariumcouncil.org/default.asp Marine Aquarium Council Home Page]
- [http://www.storysmith.net/Terrorism.htm Cyanide intoxication], by Charles Stewart
- Category:Toxicology Category:Chemical weapons ja:シアン

Non-metal

Together with the metals and metalloids, a nonmetal is one of three categories of chemical elements as distinguished by ionisation and bonding properties. These properties stem from the fact that nonmetals are highly electronegative, i.e. they gain valence electrons from other atoms more readily than they give them up. The nonmetals are Halogens, Noble gases and the following elements in order of atomic number (nonmetals as chemical series):
- hydrogen (H)
- carbon (C)
- nitrogen (N)
- oxygen (O)
- phosphorus (P)
- sulfur (S)
- selenium (Se) Most nonmetals are found at the upper right of the periodic table. The exception is