Hydrogen
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This article is about the chemistry of hydrogen. For the physics of atomic hydrogen, see hydrogen atom. For the drum machine software, see Hydrogen (software).
1
(none) ← hydrogen → helium
-↑H↓Li
Periodic Table - Extended Periodic Table
General
Name, Symbol, Number
hydrogen, H, 1
Chemical series
nonmetals
Group, Period, Block
1, 1, s
Appearance
colorless
Standard atomic weight
1.00794(7) g·mol−1
Electron configuration
1s1
Electrons per shell
1
Physical properties
Phase
gas
Density
(0 °C, 101.325 kPa)0.08988 g/L
Melting point
14.01 K(−259.14 °C, −434.45 °F)
Boiling point
20.28 K(−252.87 °C, −423.17 °F)
Triple point
13.8033 K, 7.042 kPa
Critical point
32.97 K, 1.293 MPa
Heat of fusion
(H2) 0.117 kJ·mol−1
Heat of vaporization
(H2) 0.904 kJ·mol−1
Heat capacity
(25 °C) (H2)28.836 J·mol−1·K−1
Vapor pressure
P(Pa)
1
10
100
1 k
10 k
100 k
at T(K)
15
20
Atomic properties
Crystal structure
hexagonal
Oxidation states
1, −1(amphoteric oxide)
Electronegativity
2.20 (scale Pauling)
Ionization energies
1st: 1312.0 kJ/mol
Atomic radius
25 pm
Atomic radius (calc.)
53 pm (Bohr radius)
Covalent radius
37 pm
Van der Waals radius
120 pm
Miscellaneous
Thermal conductivity
(300 K) 180.5 m W·m−1·K−1
Speed of sound
(gas, 27 °C) 1310 m/s
CAS registry number
1333-74-0 (H2)
Selected isotopes
Main article: Isotopes of hydrogen
iso
NA
half-life
DM
DE (MeV)
DP
1H
99.985%
H is stable with 0 neutrons
²H
0.0115%
H is stable with 1 neutron
3H
trace
12.32 y
β−
0.019
3He
References
Hydrogen (IPA: /ˈhaɪdrə(ʊ)dʒən/), is a chemical element represented by the symbol H and an atomic number of 1. At standard temperature and pressure it is a colorless, odorless, nonmetallic, tasteless, highly flammable diatomic gas (H2). With an atomic mass of 1.00794 g/mol, hydrogen is the lightest element.
Hydrogen is the most abundant of the chemical elements, constituting roughly 75% of the universe's elemental mass.[1] Stars in the main sequence are mainly composed of hydrogen in its plasma state. Elemental hydrogen is relatively rare on Earth, and is industrially produced from hydrocarbons such as methane, after which most elemental hydrogen is used "captively" (meaning locally at the production site), with the largest markets about equally divided between fossil fuel upgrading (e.g., hydrocracking) and in ammonia production (mostly for the fertilizer market). Hydrogen may be produced from water using the process of electrolysis, but this process is presently significantly more expensive commercially than hydrogen production from natural gas.
The most common naturally occurring isotope of hydrogen, known as protium, has a single proton and no neutrons. In ionic compounds it can take on either a positive charge (becoming a cation composed of a bare proton) or a negative charge (becoming an anion known as a hydride). Hydrogen can form compounds with most elements and is present in water and most organic compounds. It plays a particularly important role in acid-base chemistry, in which many reactions involve the exchange of protons between soluble molecules. As the only neutral atom for which the Schrödinger equation can be solved analytically, study of the energetics and bonding of the hydrogen atom has played a key role in the development of quantum mechanics.
Contents[hide]
1 Nomenclature
2 History
2.1 Discovery of H2
2.2 Role in history of quantum theory
3 Natural occurrence
4 The hydrogen atom
4.1 Electron energy levels
4.2 Isotopes
5 Elemental molecular forms
6 Chemical and physical properties
6.1 Combustion
7 Compounds
7.1 Covalent and organic compounds
7.2 Hydrides
7.3 "Protons" and acids
8 Production
8.1 Laboratory syntheses
8.2 Industrial syntheses
8.3 Biological syntheses
9 Applications
9.1 Hydrogen as an energy carrier
10 See also
11 References
12 Further reading
13 External links//
Nomenclature
Hydrogen, Latin: 'hydrogenium', is from Ancient Greek ὕδωρ (hydor): "water" and (genes): "forming". Ancient Greek γείνομαι (geinomai): "to beget or sire")[2]
The word "hydrogen" has several different meanings;
the name of an element.
an atom, sometimes called "H dot", that is abundant in space but essentially absent on Earth, because it dimerizes.
a diatomic molecule that occurs naturally in trace amounts in the Earth's atmosphere; chemists increasingly refer to H2 as dihydrogen,[3] or hydrogen molecule, to distinguish this molecule from atomic hydrogen and hydrogen found in other compounds.
the atomic constituent within all organic compounds, water, and many other chemical compounds.
The elemental forms of hydrogen should not be confused with hydrogen as it appears in chemical compounds.
History
Discovery of H2
Hydrogen gas, H2, was first artificially produced and formally described by T. Von Hohenheim (also known as Paracelsus, 1493 – 1541) via the mixing of metals with strong acids. He was unaware that the flammable gas produced by this chemical reaction was a new chemical element. In 1671, Robert Boyle rediscovered and described the reaction between iron filings and dilute acids, which results in the production of hydrogen gas.[4] In 1766, Henry Cavendish was the first to recognize hydrogen gas as a discrete substance, by identifying the gas from a metal-acid reaction as "inflammable air" and further finding that the gas produces water when burned. Cavendish had stumbled on hydrogen when experimenting with acids and mercury. Although he wrongly assumed that hydrogen was a liberated component of the mercury rather than the acid, he was still able to accurately describe several key properties of hydrogen. He is usually given credit for its discovery as an element. In 1783, Antoine Lavoisier gave the element the name of hydrogen when he (with Laplace) reproduced Cavendish's finding that water is produced when hydrogen is burned. Lavoisier's name for the gas won out.
One of the first uses of H2 was for balloons, and later airships. The H2 was obtained by reacting sulfuric acid and metallic iron. Infamously, H2 was used in the Hindenburg airship that was destroyed in a midair fire. The highly flammable hydrogen (H2) was later replaced for airships and most balloons by the unreactive helium (He).
Role in history of quantum theory
Because of its relatively simple atomic structure, consisting only of a proton and an electron, the hydrogen atom, together with the spectrum of light produced from it or absorbed by it, has been central to the development of the theory of atomic structure. Furthermore, the corresponding simplicity of the hydrogen molecule and the corresponding cation H2+ allowed fuller understanding of the nature of the chemical bond, which followed shortly after the quantum mechanical treatment of the hydrogen atom had been developed in the mid-1920s.
One of the first quantum effects to be explicitly noticed (but not understood at the time) was a Maxwell observation involving hydrogen, half a century before full quantum mechanical theory arrived. Maxwell observed that the specific heat capacity of H2 unaccountably departs from that of a diatomic gas below room temperature and begins to increasingly resemble that of a monatomic gas at cryogenic temperatures. According to quantum theory, this behavior arises from the spacing of the (quantized) rotational energy levels, which are particularly wide-spaced in H2 because of its low mass. These widely spaced levels inhibit equal partition of heat energy into rotational motion in hydrogen at low temperatures. Diatomic gases composed of heavier atoms do not have such widely spaced levels and do not exhibit the same effect.[5]
Natural occurrence
NGC 604, a giant region of ionized hydrogen in the Triangulum Galaxy
Hydrogen is the most abundant element in the universe, making up 75% of normal matter by mass and over 90% by number of atoms.[6] This element is found in great abundance in stars and gas giant planets. Molecular clouds of H2 are associated with star formation. Hydrogen plays a vital role in powering stars through proton-proton reaction nuclear fusion.
Throughout the universe, hydrogen is mostly found in the atomic and plasma states whose properties are quite different from molecular hydrogen. As a plasma, hydrogen's electron and proton are not bound together, resulting in very high electrical conductivity and high emissivity (producing the light from the sun and other stars). The charged particles are highly influenced by magnetic and electric fields. For example, in the solar wind they interact with the Earth's magnetosphere giving rise to Birkeland currents and the aurora. Hydrogen is found in the neutral atomic state in the Interstellar medium. The large amount of neutral hydrogen found in the damped Lyman-alpha systems is thought to dominate the cosmological baryonic density of the Universe up to redshift z=4.[7]
Under ordinary conditions on Earth, elemental hydrogen exists as the diatomic gas, H2 (for data see table). However, hydrogen gas is very rare in the Earth's atmosphere (1 ppm by volume) because of its light weight, which enables it to escape from Earth's gravity more easily than heavier gases. Although H atoms and H2 molecules are abundant in interstellar space, they are difficult to generate, concentrate, and purify on Earth. Still, hydrogen is the third most abundant element on the Earth's surface.[8] Most of the Earth's hydrogen is in the form of chemical compounds such as hydrocarbons and water.[9] Hydrogen gas is produced by some bacteria and algae and is a natural component of flatus. Methane is a hydrogen source of increasing importance.
The hydrogen atom
Electron energy levels
Main article: Hydrogen atom
Depiction of a hydrogen atom showing the diameter as about twice the Bohr model radius. (Image not to scale)
The ground state energy level of the electron in a hydrogen atom is 13.6 eV, which is equivalent to an ultraviolet photon of roughly 92 nm.
The energy levels of hydrogen can be calculated fairly accurately using the Bohr model of the atom, which conceptualizes the electron as "orbiting" the proton in analogy to the Earth's orbit of the sun. However, the electromagnetic force attracts electrons and protons to one another, while planets and celestial objects are attracted to each other by gravity. Because of the discretization of angular momentum postulated in early quantum mechanics by Bohr, the electron in the Bohr model can only occupy certain allowed distances from the proton, and therefore only certain allowed energies. A more accurate description of the hydrogen atom comes from a purely quantum mechanical treatment that uses the Schrödinger equation or the equivalent Feynman path integral formulation to calculate the probability density of the electron around the proton. Treating the electron as a matter wave reproduces chemical results such as shape of the hydrogen atom more naturally than the particle-based Bohr model, although the energy and spectral results are the same. Modeling the system fully using the reduced mass of nucleus and electron (as one would do in the two-body problem in celestial mechanics) yields an even better formula for the hydrogen spectra, and also the correct spectral shifts for the isotopes deuterium and tritium. Very small adjustments in energy levels in the hydrogen atom, which correspond to actual spectral effects, may be determined by using a full quantum mechanical theory which corrects for the effects of special relativity (see Dirac equation), and by accounting for quantum effects arising from production of virtual particles in the vacuum and as a result of electric fields (see quantum electrodynamics).
In hydrogen gas, the electronic ground state energy level is split into hyperfine structure levels because of magnetic effects of the quantum mechanical spin of the electron and proton. The energy of the atom when the proton and electron spins are aligned is higher than when they are not aligned. The transition between these two states can occur through emission of a photon through a magnetic dipole transition. Radio telescopes can detect the radiation produced in this process, which is used to map the distribution of hydrogen in the galaxy.
Isotopes
Main article: Isotopes of hydrogen
Protium, the most common isotope of hydrogen, has one proton and one electron. Unique among all stable isotopes, it has no neutrons. (see diproton for discussion of why others do not exist)
Hydrogen has three naturally occurring isotopes, denoted 1H, ²H, and ³H. Other, highly unstable nuclei (4H to 7H) have been synthesized in the laboratory but not observed in nature.[10][11]
1H is the most common hydrogen isotope with an abundance of more than 99.98%. Because the nucleus of this isotope consists of only a single proton, it is given the descriptive but rarely used formal name protium.
²H, the other stable hydrogen isotope, is known as deuterium and contains one proton and one neutron in its nucleus. Deuterium comprises 0.0026 – 0.0184% (by mole-fraction or atom-fraction) of hydrogen samples on Earth, with the lower number tending to be found in samples of hydrogen gas and the higher enrichments (0.015% or 150 ppm) typical of ocean water. Deuterium is not radioactive, and does not represent a significant toxicity hazard. Water enriched in molecules that include deuterium instead of normal hydrogen is called heavy water. Deuterium and its compounds are used as a non-radioactive label in chemical experiments and in solvents for 1H-NMR spectroscopy. Heavy water is used as a neutron moderator and coolant for nuclear reactors. Deuterium is also a potential fuel for commercial nuclear fusion.
³H is known as tritium and contains one proton and two neutrons in its nucleus. It is radioactive, decaying into Helium-3 through beta decay with a half-life of 12.32 years.[9] Small amounts of tritium occur naturally because of the interaction of cosmic rays with atmospheric gases; tritium has also been released during nuclear weapons tests. It is used in nuclear fusion reactions, as a tracer in isotope geochemistry, and specialized in self-powered lighting devices. Tritium was once routinely used in chemical and biological labeling experiments as a radiolabel (this has become less common).
Hydrogen is the only element that has different names for its isotopes in common use today. (During the early study of radioactivity, various heavy radioactive isotopes were given names, but such names are no longer used). The symbols D and T (instead of ²H and ³H) are sometimes used for deuterium and tritium, but the corresponding symbol P is already in use for phosphorus and thus is not available for protium. IUPAC states that while this use is common it is not preferred.
Elemental molecular forms
First tracks observed in liquid hydrogen bubble chamber at the Bevatron.
There are two different types of diatomic hydrogen molecules that differ by the relative spin of their nuclei.[12] In the orthohydrogen form, the spins of the two protons are parallel and form a triplet state; in the parahydrogen form the spins are antiparallel and form a singlet. At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form, also known as the "normal form".[13] The equilibrium ratio of orthohydrogen to parahydrogen depends on temperature, but since the ortho form is an excited state and has a higher energy than the para form, it is unstable and cannot be purified. At very low temperatures, the equilibrium state is composed almost exclusively of the para form. The physical properties of pure parahydrogen differ slightly from those of the normal form.[14] The ortho/para distinction also occurs in other hydrogen-containing molecules or functional groups, such as water and methylene.
The uncatalyzed interconversion between para and ortho H2 increases with increasing temperature; thus rapidly condensed H2 contains large quantities of the high-energy ortho form that convert to the para form very slowly.[15] The ortho/para ratio in condensed H2 is an important consideration in the preparation and storage of liquid hydrogen: the conversion from ortho to para is exothermic and produces enough heat to evaporate the hydrogen liquid, leading to loss of the liquefied material. Catalysts for the ortho-para interconversion, such as iron compounds, are used during hydrogen cooling.[16]
A molecular form called protonated molecular hydrogen, or H3+, is found in the interstellar medium (ISM), where it is generated by ionization of molecular Hydrogen from cosmic rays. It has also been observed in the upper atmosphere of the planet Jupiter. This molecule is relatively stable in the environment of outer space due to the low temperature and density. H3+ is one of the most abundant ions in the Universe, and it plays a notable role in the chemistry of the interstellar medium.[17]
Chemical and physical properties
The solubility and adsorption characteristics of hydrogen with various metals are very important in metallurgy (as many metals can suffer hydrogen embrittlement) and in developing safe ways to store it for use as a fuel. Hydrogen is highly soluble in many compounds composed of rare earth metals and transition metals[18] and can be dissolved in both crystalline and amorphous metals.[19] Hydrogen solubility in metals is influenced by local distortions or impurities in the metal crystal lattice.[20]
Combustion
Hydrogen can combust rapidly in air. It burned rapidly in the Hindenburg disaster on May 6, 1937
Hydrogen gas is highly flammable and will burn at concentrations as low as 4% H2 in air. The enthalpy of combustion for hydrogen is – 286 kJ/mol; it combusts according to the following balanced equation.
2 H2(g) + O2(g) → 2 H2O(l) + 572 kJ/mol
When mixed with oxygen across a wide range of proportions, hydrogen explodes upon ignition. Hydrogen burns violently in air. Pure hydrogen-oxygen flames are nearly invisible to the naked eye, as illustrated by the faintness of flame from the main Space Shuttle engines (as opposed to the easily visible flames from the shuttle boosters). Thus it is difficult to visually detect if a hydrogen leak is burning. The Hindenburg zeppelin flames seen in the adjacent picture are hydrogen flames colored with material from the covering skin of the zeppelin which contained carbon and pyrophoric aluminium powder, as well as other combustible materials.[21] (Regardless of the cause of this fire, this was clearly primarily a hydrogen fire since skin of the Zeppelin alone would have taken many hours to burn).[22] Another characteristic of hydrogen fires is that the flames tend to ascend rapidly with the gas in air, as illustrated by the Hindenberg flames, causing less damage than hydrocarbon fires. For example, two-thirds of the Hindenburg passengers survived the hydrogen fire, and many of the deaths which occurred were from falling or from gasoline burns.[23]
H2 reacts directly with other oxidizing elements. A violent and spontaneous reaction can occur at room temperature with chlorine and fluorine, forming the corresponding hydrogen halides: hydrogen chloride and hydrogen fluoride.
Compounds
Further information: Hydrogen compounds
Covalent and organic compounds
While H2 is not very reactive under standard conditions, it does form compounds with most elements. Millions of hydrocarbons are known, but they are not formed by the direct reaction of elementary hydrogen and carbon (although synthesis gas production followed by the Fischer-Tropsch process to make hydrocarbons comes close to being an exception, as this begins with coal and the elemental hydrogen is generated in situ). Hydrogen can form compounds with elements that are more electronegative, such as halogens (e.g., F, Cl, Br, I) and chalcogens (O, S, Se); in these compounds hydrogen takes on a partial positive charge. When bonded to fluorine, oxygen, or nitrogen, hydrogen can participate in a form of strong noncovalent bonding called hydrogen bonding, which is critical to the stability of many biological molecules. Hydrogen also forms compounds with less electronegative elements, such as the metals and metalloids, in which it takes on a partial negative charge. These compounds are often known as hydrides.
Hydrogen forms a vast array of compounds with carbon. Because of their general association with living things, these compounds came to be called organic compounds; the study of their properties is known as organic chemistry and their study in the context of living organisms is known as biochemistry. By some definitions, "organic" compounds are only required to contain carbon (as a classic historical example, urea). However, most of them also contain hydrogen, and since it is the carbon-hydrogen bond which gives this class of compounds most of its particular chemical characteristics, carbon-hydrogen bonds are required in some definitions of the word "organic" in chemistry. (This latter definition is not perfect, however, as in this definition urea would not be included as an organic compound).
In inorganic chemistry, hydrides can also serve as bridging ligands that link two metal centers in a coordination complex. This function is particularly common in group 13 elements, especially in boranes (boron hydrides) and aluminum complexes, as well as in clustered carboranes.[9]
Hydrides
Compounds of hydrogen are often called hydrides, a term that is used fairly loosely. To chemists, the term "hydride" usually implies that the H atom has acquired a negative or anionic character, denoted H−. The existence of the hydride anion, suggested by G.N. Lewis in 1916 for group I and II salt-like hydrides, was demonstrated by Moers in 1920 with the electrolysis of molten lithium hydride (LiH), that produced a stoichiometric quantity of hydrogen at the anode.[24] For hydrides other than group I and II metals, the term is quite misleading, considering the low electronegativity of hydrogen. An exception in group II hydrides is BeH2, which is polymeric. In lithium aluminum hydride, the AlH4− anion carries hydridic centers firmly attached to the Al(III). Although hydrides can be formed with almost all main-group elements, the number and combination of possible compounds varies widely; for example, there are over 100 binary borane hydrides known, but only one binary aluminum hydride.[25] Binary indium hydride has not yet been identified, although larger complexes exist.[26]
"Protons" and acids
Oxidation of H2 formally gives the proton, H+. This species is central to discussion of acids, though the term proton is used loosely to refer to positively charged or cationic hydrogen, denoted H+. A bare proton H+ cannot exist in solution because of its strong tendency to attach itself to atoms or molecules with electrons. To avoid the convenient fiction of the naked "solvated proton" in solution, acidic aqueous solutions are sometimes considered to contain the hydronium ion (H3O+) organized into clusters to form H9O4+.[27] Other oxonium ions are found when water is in solution with other solvents.[28]
Although exotic on earth, one of the most common ions in the universe is the H3+ ion, known as protonated molecular hydrogen or the triatomic hydrogen cation.[29]
Production
H2 is produced in chemistry and biology laboratories, often as a by-product of other reactions; in industry for the hydrogenation of unsaturated substrates; and in nature as a means of expelling reducing equivalents in biochemical reactions.
Laboratory syntheses
In the laboratory, H2 is usually prepared by the reaction of acids on metals such as zinc.
Zn + 2 H+ → Zn2+ + H2
Aluminum produces H2 upon treatment with acids but also with base:
2 Al + 6 H2O → 2 Al(OH)3 + 3 H2
The electrolysis of water is a simple method of producing hydrogen, although the resulting hydrogen necessarily has less energy content than was required to produce it. A low voltage current is run through the water, and gaseous oxygen forms at the anode while gaseous hydrogen forms at the cathode. Typically the cathode is made from platinum or another inert metal when producing hydrogen for storage. If, however, the gas is to be burnt on site, oxygen is desirable to assist the combustion, and so both electrodes would be made from inert metals. (Iron, for instance, would oxidize, and thus decrease the amount of oxygen given off.) The theoretical maximum efficiency (electricity used vs. energetic value of hydrogen produced) is between 80 – 94%. Bellona Report on Hydrogen
2H2O(aq) → 2H2(g) + O2(g)
In 2007, it was discovered that an alloy of aluminium and gallium in pellet form added to water could be used to generate hydrogen.[30] The process creates also creates alumina, but the expensive gallium, which prevents to formation of an oxide skin on the pellets, can be re-used. This potentially has important implications for a hydrogen economy, since hydrogen can be produced on-site and does not need to be transported.
Industrial syntheses
Hydrogen can be prepared in several different ways but the economically most important processes involve removal of hydrogen from hydrocarbons. Commercial bulk hydrogen is usually produced by the steam reforming of natural gas.[31] At high temperatures (700 – 1100 °C; 1,300 – 2,000 °F), steam (water vapor) reacts with methane to yield carbon monoxide and H2.
CH4 + H2O → CO + 3 H2
This reaction is favored at low pressures but is nonetheless conducted at high pressures (20 atm; 600 inHg) since high pressure H2 is the most marketable product. The product mixture is known as "synthesis gas" because it is often used directly for the production of methanol and related compounds. Hydrocarbons other than methane can be used to produce synthesis gas with varying product ratios. One of the many complications to this highly optimized technology is the formation of coke or carbon:
CH4 → C + 2 H2
Consequently, steam reforming typically employs an excess of H2O.
Additional hydrogen from steam reforming can be recovered from the carbon monoxide through the water gas shift reaction, especially with an iron oxide catalyst. This reaction is also a common industrial source of carbon dioxide:[31] :CO + H2O → CO2 + H2
Other important methods for H2 production include partial oxidation of hydrocarbons:
CH4 + 0.5 O2 → CO + 2 H2
and the coal reaction, which can serve as a prelude to the shift reaction above:[31] :C + H2O → CO + H2
Hydrogen is sometimes produced and consumed in the same industrial process, without being separated. In the Haber process for the production of ammonia (the world's fifth most produced industrial compound), hydrogen is generated from natural gas.
Hydrogen is also produced in usable quantities as a co-product of the major petrochemical processes of steam cracking and reforming. Electrolysis of brine to yield chlorine also produces hydrogen as a co-product.
Biological syntheses
H2 is a product of some types of anaerobic metabolism and is produced by several microorganisms, usually via reactions catalyzed by iron- or nickel-containing enzymes called hydrogenases. These enzymes catalyze the reversible redox reaction between H2 and its component two protons and two electrons. Evolution of hydrogen gas occurs in the transfer of reducing equivalents produced during pyruvate fermentation to water.[32]
Water splitting, in which water is decomposed into its component protons, electrons, and oxygen, occurs in the light reactions in all photosynthetic organisms. Some such organisms — including the alga Chlamydomonas reinhardtii and cyanobacteria — have evolved a second step in the dark reactions in which protons and electrons are reduced to form H2 gas by specialized hydrogenases in the chloroplast.[33] Efforts have been undertaken to genetically modify cyanobacterial hydrogenases to efficiently synthesize H2 gas even in the presence of oxygen.[34]
Other rarer but mechanistically interesting routes to H2 production also exist in nature. Nitrogenase produces approximately one equivalent of H2 for each equivalent of N2 reduced to ammonia. Some phosphatases reduce phosphite to H2.
Applications
Large quantities of H2 are needed in the petroleum and chemical industries. The largest application of H2 is for the processing ("upgrading") of fossil fuels, and in the production of ammonia. The key consumers of H2 in the petrochemical plant include hydrodealkylation, hydrodesulfurization, and hydrocracking.[35] H2 has several other important uses. H2 is used as a hydrogenating agent, particularly in increasing the level of saturation of unsaturated fats and oils (found in items such as margarine), and in the production of methanol. It is similarly the source of hydrogen in the manufacture of hydrochloric acid. H2 is also used as a reducing agent of metallic ores.
Apart from its use as a reactant, H2 has wide applications in physics and engineering. It is used as a shielding gas in welding methods such as atomic hydrogen welding. H2 is used as the rotor coolant in electrical generators at power stations, because it has the highest thermal conductivity of any gas. Liquid H2 is used in cryogenic research, including superconductivity studies. Since H2 is lighter than air, having a little more than 1/15th of the density of air, it was once widely used as a lifting agent in balloons and airships. However, this use was curtailed after the Hindenburg disaster convinced the public that the gas was too dangerous for this purpose. Hydrogen is still regularly used for the inflation of weather balloons.
Hydrogen's rarer isotopes also each have specific applications. Deuterium (hydrogen-2) is used in nuclear fission applications as a moderator to slow neutrons, and in nuclear fusion reactions. Deuterium compounds have applications in chemistry and biology in studies of reaction isotope effects. Tritium (hydrogen-3), produced in nuclear reactors, is used in the production of hydrogen bombs, as an isotopic label in the biosciences, and as a radiation source in luminous paints.
The triple point temperature of equilibrium hydrogen is a defining fixed point on the ITS-90 temperature scale.
Hydrogen as an energy carrier
Main article: Hydrogen economy
Hydrogen is not an energy source, except in the hypothetical context of commercial nuclear fusion power plants using deuterium or tritium, a technology presently far from development. The sun's energy comes from nuclear fusion of hydrogen but this process is difficult to achieve on earth. Elemental hydrogen from solar, biological, or electrical sources costs more in energy to make than is obtained by burning it. Hydrogen may be obtained from fossil sources (such as methane) for less energy than required to make it, but these sources are unsustainable, and are also themselves direct energy sources (and are rightly regarded as the basic source of the energy in the hydrogen obtained from them).
Molecular hydrogen has been widely discussed in the context of energy, as a possible carrier of energy on an economy-wide scale. A theoretical advantage of using H2 as an energy carrier is the localization and concentration of environmentally unwelcome aspects of hydrogen manufacture from fossil fuel energy sources. For example, CO2 sequestration followed by carbon capture and storage could be conducted at the point of H2 production from methane. Hydrogen used in transportation would burn cleanly, without carbon emissions. However, the infrastructure costs associated with full conversion to a hydrogen economy would be substantial.[36] In addition, the energy density of both liquid hydrogen and hydrogen gas at any practicable pressure is significantly less than that of traditional fuel sources.
See also
[hide] v • d • e Diatomic Elements
HydrogenH2
NitrogenN2
OxygenO2
FluorineF2
ChlorineCl2
BromineBr2
IodineI2
AstatineAt2
Nitrogen
Nitrogen
From Wikipedia, the free encyclopedia
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7
carbon ← nitrogen → oxygen
-↑N↓P
Periodic table - Extended periodic table
General
Name, symbol, number
nitrogen, N, 7
Chemical series
nonmetals
Group, period, block
15, 2, p
Appearance
colorless gas
Standard atomic weight
14.0067(2) g·mol−1
Electron configuration
1s2 2s2 2p3
Electrons per shell
2, 5
Physical properties
Phase
gas
Density
(0 °C, 101.325 kPa)1.251 g/L
Melting point
63.15 K(-210.00 °C, -346.00 °F)
Boiling point
77.36 K(-195.79 °C, -320.42 °F)
Critical point
126.21 K, 3.39 MPa
Heat of fusion
(N2) 0.720 kJ·mol−1
Heat of vaporization
(N2) 5.57 kJ·mol−1
Heat capacity
(25 °C) (N2)29.124 J·mol−1·K−1
Vapor pressure
P/Pa
1
10
100
1 k
10 k
100 k
at T/K
37
41
46
53
62
77
Atomic properties
Crystal structure
hexagonal
Oxidation states
±3, 5, 4, 2(strongly acidic oxide)
Electronegativity
3.04 (Pauling scale)
Ionization energies(more)
1st: 1402.3 kJ·mol−1
2nd: 2856 kJ·mol−1
3rd: 4578.1 kJ·mol−1
Atomic radius
65 pm
Atomic radius (calc.)
56 pm
Covalent radius
75 pm
Van der Waals radius
155 pm
Miscellaneous
Magnetic ordering
diamagnetic
Thermal conductivity
(300 K) 25.83 m W·m−1·K−1
Speed of sound
(gas, 27 °C) 353 m/s
CAS registry number
7727-37-9
Selected isotopes
Main article: Isotopes of nitrogen
iso
NA
half-life
DM
DE (MeV)
DP
13N
syn
9.965 min
ε
2.220
13C
14N
99.634%
N is stable with 7 neutrons
15N
0.366%
N is stable with 8 neutrons
References
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Nitrogen (IPA: /ˈnaɪtrədʒən/) is a chemical element which has the symbol N and atomic number 7. Elemental nitrogen is a colourless, odourless, tasteless and mostly inert diatomic gas at standard conditions, constituting 78.1% by volume of Earth's atmosphere. Nitrogen is a constituent element of all living tissues and amino acids. Many industrially important compounds, such as ammonia, nitric acid, and cyanides, contain nitrogen.
Contents[hide]
1 Notable characteristics of elemental nitrogen
2 Occurrence
3 Isotopes
4 Electromagnetic spectrum
5 History
6 Biological role
7 Reactions toward metals
8 Modern applications
8.1 Molecular nitrogen (gas and liquid)
8.1.1 Liquid nitrogen
9 Nitrogen compounds in industry
9.1 Simple compounds
9.2 Nitrogen compounds of notable economic importance
10 Dangers
11 See also
12 References
13 External links//
[edit] Notable characteristics of elemental nitrogen
Nitrogen is a nonmetal, with an electronegativity of 3.0. It has five electrons in its outer shell and is therefore trivalent in most compounds. The triple bond in molecular nitrogen (N2) is one of the strongest in nature. The resulting difficulty of converting (N2) into other compounds, and the ease (and associated high energy release) of converting nitrogen compounds into elemental N2, have dominated the role of nitrogen in both nature and human economic activities.
At atmospheric pressure molecular nitrogen condenses (liquifies) at 77 K (−195.8 °C) and freezes at 63 K (−210.0 °C) into the beta hexagonal close-packed crystal allotropic form. Below 35.4 K (−237.6 °C) nitrogen assumes the alpha cubic crystal allotropic form. Liquid nitrogen, a fluid resembling water, but with 80.8% of the density, is a common cryogen.
Unstable allotropes of nitrogen consisting of more than two nitrogen atoms have been produced in the laboratory, like N3 and N4.[1] Under extremely high pressures (1.1 million atm) and high temperatures (2000 K), as produced under diamond anvil conditions, nitrogen polymerizes into the single bonded diamond crystal structure, an allotrope nicknamed "nitrogen diamond."[2]
[edit] Occurrence
Nitrogen is the largest single component of the Earth's atmosphere (78.082% by volume of dry air, 75.3% by weight in dry air).
14N is created as part of the fusion processes in stars, and is estimated to be the 7th most abundant chemical element (by mass) in our universe.
Compounds that contain this element have been observed by astronomers, and molecular nitrogen has been detected in interstellar space by David Knauth and coworkers using the Far Ultraviolet Spectroscopic Explorer. Molecular nitrogen is a major constituent of Titan's thick atmosphere, and occurs in trace amounts of other planetary atmospheres.
Nitrogen is present in all living tissues as proteins, nucleic acids and other molecules. It is a large component of animal waste (for example, guano), usually in the form of urea, uric acid, and compounds of these nitrogenous products.
See also: Nitrate minerals and Ammonium minerals
[edit] Isotopes
See also: Isotopes of nitrogen
There are two stable isotopes of nitrogen: 14N and 15N. By far the most common is 14N (99.634%), which is produced in the CNO cycle in stars and the remaining is 15N. Of the ten isotopes produced synthetically, 13N has a half life of ten 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 typically result in 15N enrichment of the substrate and depletion of the product.
0.73% of the molecular nitrogen in Earth's atmosphere is comprised of the isotopologue 14N15N and almost all the rest is 14N2.
[edit] Electromagnetic spectrum
Molecular nitrogen (14N2) is largely transparent to infrared and visible radiation because it is a homonuclear molecule and thus has no dipole moment to couple to electromagnetic radiation at these wavelengths. Significant absorption occurs at extreme ultraviolet wavelengths, beginning around 100 nanometers. This is associated with electronic transitions in the molecule to states in which charge is not distributed evenly between nitrogen atoms. Nitrogen absorption leads to significant absorption of ultraviolet radiation in the Earth's upper atmosphere as well as in the atmospheres of other planetary bodies. For similar reasons, pure molecular nitrogen lasers typically emit light in the ultraviolet range.
Nitrogen also makes a contribution to visible air glow from the Earth's upper atmosphere, through electron impact excitation followed by emission. This visible blue air glow (seen in the polar aurora and in the re-entry glow of returning spacecraft) typically results not from molecular nitrogen, but rather from free nitrogen atoms combining with oxygen to form nitric oxide (NO).
[edit] History
Nitrogen (Latin nitrogenium, where nitrum (from Greek nitron) means "native soda" (see niter), and genes means "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, from the Greek word αζωτος meaning "lifeless". Animals died in it, and it was the principal component of air in which animals had suffocated and flames had burned to extinction. This term has become the French word for "nitrogen" and later spread out to many other languages.
Argon was discovered when it was noticed that nitrogen from air is not identical to nitrogen from chemical reactions.
Compounds of nitrogen were known in the Middle Ages. The alchemists knew nitric acid as aqua fortis (strong water). The mixture of nitric and hydrochloric acids was known as aqua regia (royal water), celebrated for its ability to dissolve gold (the king of metals). The earliest industrial and agricultural applications of nitrogen compounds involved uses in the form of saltpeter (sodium- or potassium nitrate), notably in gunpowder, and much later, as fertilizer, and later still, as a chemical feedstock.
[edit] Biological role
See also: nitrogen cycle
Nitrogen is an essential part of amino acids and nucleic acids, both of which are essential to all life on Earth.
Molecular nitrogen in the atmosphere cannot be used directly by either plants or animals, and needs to be converted to other compounds, or "fixed," in order to be used by life. Precipitation often contains substantial quantities of ammonium and nitrate, both thought to be a result of nitrogen fixation by lightning and other atmospheric electric phenomena. However, because ammonium is preferentially retained by the forest canopy relative to atmospheric nitrate, most of the fixed nitrogen that reaches the soil surface under trees is in the form of nitrate. Soil nitrate is preferentially assimilated by tree roots relative to soil ammonium.
Specific bacteria (e.g. Rhizobium trifolium) possess nitrogenase enzymes which can fix atmospheric nitrogen (see nitrogen fixation) into a form (ammonium ion) which is chemically useful to higher organisms. This process requires a large amount of energy and anoxic conditions. Such bacteria may be free in the soil (e.g. azotobacter) but normally exist in a symbiotic relationship in the root nodules of leguminous plants (e.g. clover or the soya bean plant). Nitrogen fixating bacteria can be symbiotic with a number of unrelated plant species. Common examples are legumes, alders, lichens, casuarina, myrica, liverwort, and gunnera.
As part of the symbiotic relationship, the plant subsequently converts the ammonium ion to nitrogen oxides and amino acids to form proteins and other biologically useful molecules, such as alkaloids. In return for the usable (fixed) nitrogen, the plant secretes sugars to the symbiotic bacteria.
Some plants are able to assimilate nitrogen directly in the form of nitrates which may be present in soil from natural mineral deposits, artificial fertilizers, animal waste, or organic decay (as the product of bacteria, but not bacteria specifically associated with the plant). Nitrates absorbed in this fashion are converted to nitrites by the enzyme nitrate reductase, and then converted to ammonia by another enzyme called nitrite reductase.
Nitrogen compounds are basic building blocks in animal biology. Animals use nitrogen-containing amino acids from plant sources, as starting materials for all nitrogen-compound animal biochemistry, including the manufacture of proteins and nucleic acids. Some plant-feeding insects are so dependent on nitrogen in their diet, that varying the amount of nitrogen fertilizer applied to a plant can affect the birth rate of the insects feeding on it.[3]
Soluble nitrate is an important limiting factor in the growth of certain bacteria in ocean waters. In many places in the world, artificial fertilizers applied to crop-lands to increase yields result in run-off delivery of soluble nitrogen to oceans at river mouths. This process can result in eutrophication of the water, as nitrogen-driven bacterial growth depletes water oxygen to the point that all higher organisms die. Well-known "dead zone" areas in the U.S. Gulf Coast and the Black Sea are due to this important polluting process.
Many saltwater fish manufacture large amounts of trimethylamine oxide to protect them from the high osmotic effects of their environment (conversion of this compound to dimethylamine is responsible for the early odor in unfresh saltwater fish: PMID 15186102). In animals, the free radical molecule nitric oxide (NO), which is derived from an amino acid, serves as an important regulatory molecule for circulation.
Animal metabolism of NO results in production of nitrite. Animal metabolism of nitrogen in proteins generally results in excretion of urea, while animal metabolism of nucleic acids results in excretion of urea and uric acid. The characteristic odor of animal flesh decay is caused by nitrogen-containing long-chain amines, such as putrescine and cadaverine.
Decay of organisms and their waste products may produce small amounts of nitrate, but most decay eventually returns nitrogen content to the atmosphere, as molecular nitrogen.
[edit] Reactions toward metals
N2 reacts spontaneously with few reagents, being resilient to acids and bases as well as oxidants and most reductants. It does however react with lithium metal. Lithium burns in an atmosphere of N2 to give lithium nitride:
6 Li + N2 → 2 Li3N
N2 forms a variety of adducts with transition metals. The first example of a dinitrogen complex is [Ru(NH3)5(N2)]2+ (see figure at right). Such compounds are now numerous, other examples include IrCl(N2)(PPh3)2, W(N2)2(Ph2CH2CH2PPh2)2, and [(η5-C5Me4H)2Zr]2(μ2,η2,η2-N2). These complexes illustrate how N2 might bind to the metal(s) in nitrogenase and the catalyst for the Haber-Bosch Process.[4]
Structure of [Ru(NH3)5(N2)]2+.
[edit] Modern applications
Nitrogen gas is acquired for industrial purposes by the fractional distillation of liquid air, or by mechanical means using gaseous air (i.e. pressurised reverse osmosis membrane or pressure swing adsorption). Commercial nitrogen is often a byproduct of air-processing for industrial concentration of oxygen for steelmaking and other purposes.
[edit] Molecular nitrogen (gas and liquid)
A computer rendering of the nitrogen molecule, N2.
Nitrogen gas has a wide variety of applications, including serving as an 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)
In ordinary incandescent light bulbs as an inexpensive alternative to argon
On top of liquid explosives for safety
The production of electronic parts such as transistors, diodes, and integrated circuits
Dried and pressurized, as a dielectric gas for high voltage equipment
The manufacturing of stainless steel
Use in military aircraft fuel systems to reduce fire hazard, see inerting system
Filling automotive and aircraft tires[5] due to its inertness and lack of moisture or oxidative qualities, as opposed to air, though this is not necessary for consumer automobiles.[6][7]
Nitrogen molecules are less likely to escape from the inside of a tire compared with the traditional air mixture used. Air consists mostly of nitrogen and oxygen. Nitrogen molecules have a larger effective diameter than oxygen molecules and therefore diffuse through porous substances more slowly.[8]
A further example of its versatility 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.
Molecular nitrogen, a diatomic gas, is apt to dimerize into a linear four nitrogen long polymer. This is an important phenomenon for understanding high voltage nitrogen dielectric switches because the process of polymerization can continue to lengthen the molecule to still longer lengths in the presence of an intense electric field. A nitrogen polymer fog is thereby created. The second virial coefficient of nitrogen also shows this effect as the compressibility of nitrogen gas is changed by the dimerization process at moderate and low temperatures.
Nitrogen tanks are also replacing carbon dioxide as the main power source for paintball markers. The downside is that nitrogen must be kept at higher pressure than CO2, making N2 tanks heavier and more expensive.
Solid nitrogen ice in a small plastic beaker with melting liquid flowing off. The nitrogen has been frozen by immersion in liquid helium[9]
[edit] Liquid nitrogen
Liquid nitrogen (liquid density at the triple point is 0.807 g/mL) is produced industrially in large quantities by fractional distillation of liquid air and is often referred to by the abbreviation, LN2. It is a cryogenic fluid which is potentially capable of causing instant frostbite on contact with living tissue (see precautions). When appropriately insulated from ambient heat, liquid nitrogen serves as a compact and readily transported source of nitrogen gas without pressurization. Further, its ability to maintain temperatures far below the freezing point of water (it boils at 77 K, which equals −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 cryopreservation of blood, reproductive cells (sperm and egg), and other biological samples and materials (see image below)
Liquid nitrogen may be used to prepare "home-made" ice cream, as these students are doing.
the cryonic preservation of humans and pets in the hope of future reanimation.
in the study of cryogenics
for demonstrations in science education
as a coolant for highly sensitive sensors and low-noise amplifiers
in dermatology for removing unsightly or potentially malignant skin lesions such as warts and actinic keratosis
as a cooling supplement for overclocking a central processing unit, a graphics processing unit, or another type of computer hardware
as a cooling medium during machining of high strength materials.
as the working fluid in a binary engine
as a means of final disposition of the dead, known as promession.
as a method of freezing water pipes in order to work on them in situations where a tap is not available to block water flow to the work area.
A tank of liquid nitrogen, used to supply a cryogenic freezer (for storing laboratory samples at a temperature of about -150 Celsius).
[edit] Nitrogen compounds in industry
[edit] Simple compounds
See also the category Nitrogen compounds.
The main neutral hydride of nitrogen is ammonia (NH3), although hydrazine (N2H4) is also commonly used. Ammonia is more basic than water by 6 orders of magnitude. In solution ammonia forms the ammonium ion (NH4+). Liquid ammonia (b.p. 240 K) is amphiprotic (displaying either Brønsted-Lowry acidic or basic character) and forms ammonium and the less common amide ions (NH2-); both amides and nitride (N3-) salts are known, but decompose in water. Singly, doubly, triply and quadruply substituted alkyl compounds of ammonia are called amines (four substitutions, to form commercially and biologically important quarternary amines, results in a positively charged nitrogen, and thus a water-soluble, or at least amphiphilic, compound). Larger chains, rings and structures of nitrogen hydrides are also known, but are generally unstable.
Other classes of nitrogen anions (negatively charged ions) are the poisonous azides (N3-), which are linear and isoelectronic to carbon dioxide, but which bind to important iron-containing enzymes in the body in a manner more resembling cyanide. Another molecule of the same structure is the colorless and relatively inert anesthetic gas dinitrogen monoxide (N2O), also known as laughing gas. This is one of a variety of oxides, the most prominent of which are nitrogen monoxide (NO) (known more commonly as nitric oxide in biology), a natural free radical molecule used by the body as a signal for short-term control of smooth muscle in the circulation. Another notable nitrogen oxide compound (a family often abbreviated NOx) is the reddish and poisonous nitrogen dioxide (NO2), which also contains an unpaired electron and is an important component of smog. Nitrogen molecules containing unpaired electrons show an understandable tendency to dimerize (thus pairing the electrons), and are generally highly reactive.
The more standard oxides, dinitrogen trioxide (N2O3) and dinitrogen pentoxide (N2O5), are actually fairly unstable and explosive-- a tendency which is driven by the stability of N2 as a product. 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, and is a fairly strong oxidizing agent.
Nitrogen can also be found in organic compounds. Common nitrogen functional groups include: amines, amides, nitro groups, imines, and enamines. The amount of nitrogen in a chemical substance can be determined by the Kjeldahl method.
[edit] Nitrogen compounds of notable economic importance
Molecular nitrogen (N2) in the atmosphere is relatively non-reactive due to its strong bond, and N2 plays an inert role in the human body, being neither produced or destroyed. In nature, nitrogen is converted into biologically (and industrially) useful compounds by some living organisms, notably certain bacteria (i.e. nitrogen fixing bacteria – see Biological role above). Molecular nitrogen is also released into the atmosphere in the process of decay, in dead plant and animal tissues. The ability to combine or fix molecular nitrogen is a key feature of modern industrial chemistry, where nitrogen and natural gas are converted into ammonia via the Haber process. Ammonia, in turn, can be used directly (primarily as a fertilizer, and in the synthesis of nitrated fertilizers), or as a precursor of many other important materials including explosives, largely via the production of nitric acid by the Ostwald process.
The organic and inorganic salts of nitric acid have been historically important as stores of chemical energy. They include important compounds such as potassium nitrate (or saltpeter, important historically for its use in gunpowder) and ammonium nitrate, an important fertilizer and explosive (see ANFO). Various other nitrated organic compounds, such as nitroglycerin and trinitrotoluene, and nitrocellulose, are used as explosives and propellants for modern firearms. Nitric acid is used as an oxidizing agent in liquid fueled rockets. Hydrazine and hydrazine derivatives find use as rocket fuels. In most of these compounds, the basic instability and tendency to burn or explode is derived from the fact that nitrogen is present as an oxide, and not as the far more stable nitrogen molecule (N2) which is a product of the compounds' thermal decomposition. When nitrates burn or explode, the formation of the powerful triple bond in the N2 which results, produces most of the energy of the reaction.
Nitrogen is a constituent of molecules in every major drug class in pharmacology and medicine. Nitrous oxide (N2O) was discovered early in the 19th century to be a partial anesthetic, though it was not used as a surgical anesthetic until later. Called "laughing gas", it was found capable of inducing a state of social disinhibition resembling drunkenness. Other notable nitrogen-containing drugs are drugs derived from plant alkaloids, such as morphine (there exist many alkaloids known to have pharmacological effects; in some cases they appear natural chemical defences of plants against predation). Nitrogen containing drugs include all of the major classes of antibiotics, and organic nitrate drugs like nitroglycerin and nitroprusside which regulate blood pressure and heart action by mimicking the action of nitric oxide.
[edit] Dangers
Rapid release of nitrogen gas into an enclosed space can displace oxygen, and therefore represents an asphyxiation hazard. This may happen with few warning symptoms, since the human carotid body is a relatively slow and poor low-oxygen (hypoxia) sensing system.[10] An example occurred shortly before the launch of the first Space Shuttle mission in 1981, when two technicians lost consciousness and died after they walked into a space located in the Shuttle's Mobile Launch Platform that was pressurized with pure nitrogen as a precaution against fire. The technicians would have been able to exit the room if they had experienced early symptoms from nitrogen-breathing.
When breathed at high partial pressures (more than about 3 atmospheres, encountered at depths below about 30 m in scuba diving) nitrogen begins to act as an anesthetic agent. As such, it can cause nitrogen narcosis, a temporary semi-anesthetized condition of mental impairment similar to that caused by nitrous oxide.
Nitrogen also dissolves in the bloodstream and body fats, and rapid decompression (particularly in the case of divers ascending too quickly, or astronauts decompressing too quickly from cabin pressure to spacesuit pressure) can lead to a potentially fatal condition called decompression sickness (formerly known as caisson sickness or more commonly, the "bends"), when nitrogen bubbles form in the bloodstream, nerves, joints, and other sensitive or vital areas.
Direct skin contact with liquid nitrogen causes severe frostbite (cryogenic burns) within seconds, though not instantly on contact, depending on form of liquid nitrogen (liquid vs. mist) and surface area of the nitrogen-soaked material (soaked clothing or cotton causing more rapid damage than a spill of direct liquid to skin, which for a few seconds is protected by the Leidenfrost effect).
[edit] See also
[hide] v • d • e Diatomic Elements
HydrogenH2
NitrogenN2
OxygenO2
FluorineF2
ChlorineCl2
BromineBr2
IodineI2
AstatineAt2
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7
carbon ← nitrogen → oxygen
-↑N↓P
Periodic table - Extended periodic table
General
Name, symbol, number
nitrogen, N, 7
Chemical series
nonmetals
Group, period, block
15, 2, p
Appearance
colorless gas
Standard atomic weight
14.0067(2) g·mol−1
Electron configuration
1s2 2s2 2p3
Electrons per shell
2, 5
Physical properties
Phase
gas
Density
(0 °C, 101.325 kPa)1.251 g/L
Melting point
63.15 K(-210.00 °C, -346.00 °F)
Boiling point
77.36 K(-195.79 °C, -320.42 °F)
Critical point
126.21 K, 3.39 MPa
Heat of fusion
(N2) 0.720 kJ·mol−1
Heat of vaporization
(N2) 5.57 kJ·mol−1
Heat capacity
(25 °C) (N2)29.124 J·mol−1·K−1
Vapor pressure
P/Pa
1
10
100
1 k
10 k
100 k
at T/K
37
41
46
53
62
77
Atomic properties
Crystal structure
hexagonal
Oxidation states
±3, 5, 4, 2(strongly acidic oxide)
Electronegativity
3.04 (Pauling scale)
Ionization energies(more)
1st: 1402.3 kJ·mol−1
2nd: 2856 kJ·mol−1
3rd: 4578.1 kJ·mol−1
Atomic radius
65 pm
Atomic radius (calc.)
56 pm
Covalent radius
75 pm
Van der Waals radius
155 pm
Miscellaneous
Magnetic ordering
diamagnetic
Thermal conductivity
(300 K) 25.83 m W·m−1·K−1
Speed of sound
(gas, 27 °C) 353 m/s
CAS registry number
7727-37-9
Selected isotopes
Main article: Isotopes of nitrogen
iso
NA
half-life
DM
DE (MeV)
DP
13N
syn
9.965 min
ε
2.220
13C
14N
99.634%
N is stable with 7 neutrons
15N
0.366%
N is stable with 8 neutrons
References
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Nitrogen (IPA: /ˈnaɪtrədʒən/) is a chemical element which has the symbol N and atomic number 7. Elemental nitrogen is a colourless, odourless, tasteless and mostly inert diatomic gas at standard conditions, constituting 78.1% by volume of Earth's atmosphere. Nitrogen is a constituent element of all living tissues and amino acids. Many industrially important compounds, such as ammonia, nitric acid, and cyanides, contain nitrogen.
Contents[hide]
1 Notable characteristics of elemental nitrogen
2 Occurrence
3 Isotopes
4 Electromagnetic spectrum
5 History
6 Biological role
7 Reactions toward metals
8 Modern applications
8.1 Molecular nitrogen (gas and liquid)
8.1.1 Liquid nitrogen
9 Nitrogen compounds in industry
9.1 Simple compounds
9.2 Nitrogen compounds of notable economic importance
10 Dangers
11 See also
12 References
13 External links//
[edit] Notable characteristics of elemental nitrogen
Nitrogen is a nonmetal, with an electronegativity of 3.0. It has five electrons in its outer shell and is therefore trivalent in most compounds. The triple bond in molecular nitrogen (N2) is one of the strongest in nature. The resulting difficulty of converting (N2) into other compounds, and the ease (and associated high energy release) of converting nitrogen compounds into elemental N2, have dominated the role of nitrogen in both nature and human economic activities.
At atmospheric pressure molecular nitrogen condenses (liquifies) at 77 K (−195.8 °C) and freezes at 63 K (−210.0 °C) into the beta hexagonal close-packed crystal allotropic form. Below 35.4 K (−237.6 °C) nitrogen assumes the alpha cubic crystal allotropic form. Liquid nitrogen, a fluid resembling water, but with 80.8% of the density, is a common cryogen.
Unstable allotropes of nitrogen consisting of more than two nitrogen atoms have been produced in the laboratory, like N3 and N4.[1] Under extremely high pressures (1.1 million atm) and high temperatures (2000 K), as produced under diamond anvil conditions, nitrogen polymerizes into the single bonded diamond crystal structure, an allotrope nicknamed "nitrogen diamond."[2]
[edit] Occurrence
Nitrogen is the largest single component of the Earth's atmosphere (78.082% by volume of dry air, 75.3% by weight in dry air).
14N is created as part of the fusion processes in stars, and is estimated to be the 7th most abundant chemical element (by mass) in our universe.
Compounds that contain this element have been observed by astronomers, and molecular nitrogen has been detected in interstellar space by David Knauth and coworkers using the Far Ultraviolet Spectroscopic Explorer. Molecular nitrogen is a major constituent of Titan's thick atmosphere, and occurs in trace amounts of other planetary atmospheres.
Nitrogen is present in all living tissues as proteins, nucleic acids and other molecules. It is a large component of animal waste (for example, guano), usually in the form of urea, uric acid, and compounds of these nitrogenous products.
See also: Nitrate minerals and Ammonium minerals
[edit] Isotopes
See also: Isotopes of nitrogen
There are two stable isotopes of nitrogen: 14N and 15N. By far the most common is 14N (99.634%), which is produced in the CNO cycle in stars and the remaining is 15N. Of the ten isotopes produced synthetically, 13N has a half life of ten 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 typically result in 15N enrichment of the substrate and depletion of the product.
0.73% of the molecular nitrogen in Earth's atmosphere is comprised of the isotopologue 14N15N and almost all the rest is 14N2.
[edit] Electromagnetic spectrum
Molecular nitrogen (14N2) is largely transparent to infrared and visible radiation because it is a homonuclear molecule and thus has no dipole moment to couple to electromagnetic radiation at these wavelengths. Significant absorption occurs at extreme ultraviolet wavelengths, beginning around 100 nanometers. This is associated with electronic transitions in the molecule to states in which charge is not distributed evenly between nitrogen atoms. Nitrogen absorption leads to significant absorption of ultraviolet radiation in the Earth's upper atmosphere as well as in the atmospheres of other planetary bodies. For similar reasons, pure molecular nitrogen lasers typically emit light in the ultraviolet range.
Nitrogen also makes a contribution to visible air glow from the Earth's upper atmosphere, through electron impact excitation followed by emission. This visible blue air glow (seen in the polar aurora and in the re-entry glow of returning spacecraft) typically results not from molecular nitrogen, but rather from free nitrogen atoms combining with oxygen to form nitric oxide (NO).
[edit] History
Nitrogen (Latin nitrogenium, where nitrum (from Greek nitron) means "native soda" (see niter), and genes means "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, from the Greek word αζωτος meaning "lifeless". Animals died in it, and it was the principal component of air in which animals had suffocated and flames had burned to extinction. This term has become the French word for "nitrogen" and later spread out to many other languages.
Argon was discovered when it was noticed that nitrogen from air is not identical to nitrogen from chemical reactions.
Compounds of nitrogen were known in the Middle Ages. The alchemists knew nitric acid as aqua fortis (strong water). The mixture of nitric and hydrochloric acids was known as aqua regia (royal water), celebrated for its ability to dissolve gold (the king of metals). The earliest industrial and agricultural applications of nitrogen compounds involved uses in the form of saltpeter (sodium- or potassium nitrate), notably in gunpowder, and much later, as fertilizer, and later still, as a chemical feedstock.
[edit] Biological role
See also: nitrogen cycle
Nitrogen is an essential part of amino acids and nucleic acids, both of which are essential to all life on Earth.
Molecular nitrogen in the atmosphere cannot be used directly by either plants or animals, and needs to be converted to other compounds, or "fixed," in order to be used by life. Precipitation often contains substantial quantities of ammonium and nitrate, both thought to be a result of nitrogen fixation by lightning and other atmospheric electric phenomena. However, because ammonium is preferentially retained by the forest canopy relative to atmospheric nitrate, most of the fixed nitrogen that reaches the soil surface under trees is in the form of nitrate. Soil nitrate is preferentially assimilated by tree roots relative to soil ammonium.
Specific bacteria (e.g. Rhizobium trifolium) possess nitrogenase enzymes which can fix atmospheric nitrogen (see nitrogen fixation) into a form (ammonium ion) which is chemically useful to higher organisms. This process requires a large amount of energy and anoxic conditions. Such bacteria may be free in the soil (e.g. azotobacter) but normally exist in a symbiotic relationship in the root nodules of leguminous plants (e.g. clover or the soya bean plant). Nitrogen fixating bacteria can be symbiotic with a number of unrelated plant species. Common examples are legumes, alders, lichens, casuarina, myrica, liverwort, and gunnera.
As part of the symbiotic relationship, the plant subsequently converts the ammonium ion to nitrogen oxides and amino acids to form proteins and other biologically useful molecules, such as alkaloids. In return for the usable (fixed) nitrogen, the plant secretes sugars to the symbiotic bacteria.
Some plants are able to assimilate nitrogen directly in the form of nitrates which may be present in soil from natural mineral deposits, artificial fertilizers, animal waste, or organic decay (as the product of bacteria, but not bacteria specifically associated with the plant). Nitrates absorbed in this fashion are converted to nitrites by the enzyme nitrate reductase, and then converted to ammonia by another enzyme called nitrite reductase.
Nitrogen compounds are basic building blocks in animal biology. Animals use nitrogen-containing amino acids from plant sources, as starting materials for all nitrogen-compound animal biochemistry, including the manufacture of proteins and nucleic acids. Some plant-feeding insects are so dependent on nitrogen in their diet, that varying the amount of nitrogen fertilizer applied to a plant can affect the birth rate of the insects feeding on it.[3]
Soluble nitrate is an important limiting factor in the growth of certain bacteria in ocean waters. In many places in the world, artificial fertilizers applied to crop-lands to increase yields result in run-off delivery of soluble nitrogen to oceans at river mouths. This process can result in eutrophication of the water, as nitrogen-driven bacterial growth depletes water oxygen to the point that all higher organisms die. Well-known "dead zone" areas in the U.S. Gulf Coast and the Black Sea are due to this important polluting process.
Many saltwater fish manufacture large amounts of trimethylamine oxide to protect them from the high osmotic effects of their environment (conversion of this compound to dimethylamine is responsible for the early odor in unfresh saltwater fish: PMID 15186102). In animals, the free radical molecule nitric oxide (NO), which is derived from an amino acid, serves as an important regulatory molecule for circulation.
Animal metabolism of NO results in production of nitrite. Animal metabolism of nitrogen in proteins generally results in excretion of urea, while animal metabolism of nucleic acids results in excretion of urea and uric acid. The characteristic odor of animal flesh decay is caused by nitrogen-containing long-chain amines, such as putrescine and cadaverine.
Decay of organisms and their waste products may produce small amounts of nitrate, but most decay eventually returns nitrogen content to the atmosphere, as molecular nitrogen.
[edit] Reactions toward metals
N2 reacts spontaneously with few reagents, being resilient to acids and bases as well as oxidants and most reductants. It does however react with lithium metal. Lithium burns in an atmosphere of N2 to give lithium nitride:
6 Li + N2 → 2 Li3N
N2 forms a variety of adducts with transition metals. The first example of a dinitrogen complex is [Ru(NH3)5(N2)]2+ (see figure at right). Such compounds are now numerous, other examples include IrCl(N2)(PPh3)2, W(N2)2(Ph2CH2CH2PPh2)2, and [(η5-C5Me4H)2Zr]2(μ2,η2,η2-N2). These complexes illustrate how N2 might bind to the metal(s) in nitrogenase and the catalyst for the Haber-Bosch Process.[4]
Structure of [Ru(NH3)5(N2)]2+.
[edit] Modern applications
Nitrogen gas is acquired for industrial purposes by the fractional distillation of liquid air, or by mechanical means using gaseous air (i.e. pressurised reverse osmosis membrane or pressure swing adsorption). Commercial nitrogen is often a byproduct of air-processing for industrial concentration of oxygen for steelmaking and other purposes.
[edit] Molecular nitrogen (gas and liquid)
A computer rendering of the nitrogen molecule, N2.
Nitrogen gas has a wide variety of applications, including serving as an 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)
In ordinary incandescent light bulbs as an inexpensive alternative to argon
On top of liquid explosives for safety
The production of electronic parts such as transistors, diodes, and integrated circuits
Dried and pressurized, as a dielectric gas for high voltage equipment
The manufacturing of stainless steel
Use in military aircraft fuel systems to reduce fire hazard, see inerting system
Filling automotive and aircraft tires[5] due to its inertness and lack of moisture or oxidative qualities, as opposed to air, though this is not necessary for consumer automobiles.[6][7]
Nitrogen molecules are less likely to escape from the inside of a tire compared with the traditional air mixture used. Air consists mostly of nitrogen and oxygen. Nitrogen molecules have a larger effective diameter than oxygen molecules and therefore diffuse through porous substances more slowly.[8]
A further example of its versatility 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.
Molecular nitrogen, a diatomic gas, is apt to dimerize into a linear four nitrogen long polymer. This is an important phenomenon for understanding high voltage nitrogen dielectric switches because the process of polymerization can continue to lengthen the molecule to still longer lengths in the presence of an intense electric field. A nitrogen polymer fog is thereby created. The second virial coefficient of nitrogen also shows this effect as the compressibility of nitrogen gas is changed by the dimerization process at moderate and low temperatures.
Nitrogen tanks are also replacing carbon dioxide as the main power source for paintball markers. The downside is that nitrogen must be kept at higher pressure than CO2, making N2 tanks heavier and more expensive.
Solid nitrogen ice in a small plastic beaker with melting liquid flowing off. The nitrogen has been frozen by immersion in liquid helium[9]
[edit] Liquid nitrogen
Liquid nitrogen (liquid density at the triple point is 0.807 g/mL) is produced industrially in large quantities by fractional distillation of liquid air and is often referred to by the abbreviation, LN2. It is a cryogenic fluid which is potentially capable of causing instant frostbite on contact with living tissue (see precautions). When appropriately insulated from ambient heat, liquid nitrogen serves as a compact and readily transported source of nitrogen gas without pressurization. Further, its ability to maintain temperatures far below the freezing point of water (it boils at 77 K, which equals −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 cryopreservation of blood, reproductive cells (sperm and egg), and other biological samples and materials (see image below)
Liquid nitrogen may be used to prepare "home-made" ice cream, as these students are doing.
the cryonic preservation of humans and pets in the hope of future reanimation.
in the study of cryogenics
for demonstrations in science education
as a coolant for highly sensitive sensors and low-noise amplifiers
in dermatology for removing unsightly or potentially malignant skin lesions such as warts and actinic keratosis
as a cooling supplement for overclocking a central processing unit, a graphics processing unit, or another type of computer hardware
as a cooling medium during machining of high strength materials.
as the working fluid in a binary engine
as a means of final disposition of the dead, known as promession.
as a method of freezing water pipes in order to work on them in situations where a tap is not available to block water flow to the work area.
A tank of liquid nitrogen, used to supply a cryogenic freezer (for storing laboratory samples at a temperature of about -150 Celsius).
[edit] Nitrogen compounds in industry
[edit] Simple compounds
See also the category Nitrogen compounds.
The main neutral hydride of nitrogen is ammonia (NH3), although hydrazine (N2H4) is also commonly used. Ammonia is more basic than water by 6 orders of magnitude. In solution ammonia forms the ammonium ion (NH4+). Liquid ammonia (b.p. 240 K) is amphiprotic (displaying either Brønsted-Lowry acidic or basic character) and forms ammonium and the less common amide ions (NH2-); both amides and nitride (N3-) salts are known, but decompose in water. Singly, doubly, triply and quadruply substituted alkyl compounds of ammonia are called amines (four substitutions, to form commercially and biologically important quarternary amines, results in a positively charged nitrogen, and thus a water-soluble, or at least amphiphilic, compound). Larger chains, rings and structures of nitrogen hydrides are also known, but are generally unstable.
Other classes of nitrogen anions (negatively charged ions) are the poisonous azides (N3-), which are linear and isoelectronic to carbon dioxide, but which bind to important iron-containing enzymes in the body in a manner more resembling cyanide. Another molecule of the same structure is the colorless and relatively inert anesthetic gas dinitrogen monoxide (N2O), also known as laughing gas. This is one of a variety of oxides, the most prominent of which are nitrogen monoxide (NO) (known more commonly as nitric oxide in biology), a natural free radical molecule used by the body as a signal for short-term control of smooth muscle in the circulation. Another notable nitrogen oxide compound (a family often abbreviated NOx) is the reddish and poisonous nitrogen dioxide (NO2), which also contains an unpaired electron and is an important component of smog. Nitrogen molecules containing unpaired electrons show an understandable tendency to dimerize (thus pairing the electrons), and are generally highly reactive.
The more standard oxides, dinitrogen trioxide (N2O3) and dinitrogen pentoxide (N2O5), are actually fairly unstable and explosive-- a tendency which is driven by the stability of N2 as a product. 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, and is a fairly strong oxidizing agent.
Nitrogen can also be found in organic compounds. Common nitrogen functional groups include: amines, amides, nitro groups, imines, and enamines. The amount of nitrogen in a chemical substance can be determined by the Kjeldahl method.
[edit] Nitrogen compounds of notable economic importance
Molecular nitrogen (N2) in the atmosphere is relatively non-reactive due to its strong bond, and N2 plays an inert role in the human body, being neither produced or destroyed. In nature, nitrogen is converted into biologically (and industrially) useful compounds by some living organisms, notably certain bacteria (i.e. nitrogen fixing bacteria – see Biological role above). Molecular nitrogen is also released into the atmosphere in the process of decay, in dead plant and animal tissues. The ability to combine or fix molecular nitrogen is a key feature of modern industrial chemistry, where nitrogen and natural gas are converted into ammonia via the Haber process. Ammonia, in turn, can be used directly (primarily as a fertilizer, and in the synthesis of nitrated fertilizers), or as a precursor of many other important materials including explosives, largely via the production of nitric acid by the Ostwald process.
The organic and inorganic salts of nitric acid have been historically important as stores of chemical energy. They include important compounds such as potassium nitrate (or saltpeter, important historically for its use in gunpowder) and ammonium nitrate, an important fertilizer and explosive (see ANFO). Various other nitrated organic compounds, such as nitroglycerin and trinitrotoluene, and nitrocellulose, are used as explosives and propellants for modern firearms. Nitric acid is used as an oxidizing agent in liquid fueled rockets. Hydrazine and hydrazine derivatives find use as rocket fuels. In most of these compounds, the basic instability and tendency to burn or explode is derived from the fact that nitrogen is present as an oxide, and not as the far more stable nitrogen molecule (N2) which is a product of the compounds' thermal decomposition. When nitrates burn or explode, the formation of the powerful triple bond in the N2 which results, produces most of the energy of the reaction.
Nitrogen is a constituent of molecules in every major drug class in pharmacology and medicine. Nitrous oxide (N2O) was discovered early in the 19th century to be a partial anesthetic, though it was not used as a surgical anesthetic until later. Called "laughing gas", it was found capable of inducing a state of social disinhibition resembling drunkenness. Other notable nitrogen-containing drugs are drugs derived from plant alkaloids, such as morphine (there exist many alkaloids known to have pharmacological effects; in some cases they appear natural chemical defences of plants against predation). Nitrogen containing drugs include all of the major classes of antibiotics, and organic nitrate drugs like nitroglycerin and nitroprusside which regulate blood pressure and heart action by mimicking the action of nitric oxide.
[edit] Dangers
Rapid release of nitrogen gas into an enclosed space can displace oxygen, and therefore represents an asphyxiation hazard. This may happen with few warning symptoms, since the human carotid body is a relatively slow and poor low-oxygen (hypoxia) sensing system.[10] An example occurred shortly before the launch of the first Space Shuttle mission in 1981, when two technicians lost consciousness and died after they walked into a space located in the Shuttle's Mobile Launch Platform that was pressurized with pure nitrogen as a precaution against fire. The technicians would have been able to exit the room if they had experienced early symptoms from nitrogen-breathing.
When breathed at high partial pressures (more than about 3 atmospheres, encountered at depths below about 30 m in scuba diving) nitrogen begins to act as an anesthetic agent. As such, it can cause nitrogen narcosis, a temporary semi-anesthetized condition of mental impairment similar to that caused by nitrous oxide.
Nitrogen also dissolves in the bloodstream and body fats, and rapid decompression (particularly in the case of divers ascending too quickly, or astronauts decompressing too quickly from cabin pressure to spacesuit pressure) can lead to a potentially fatal condition called decompression sickness (formerly known as caisson sickness or more commonly, the "bends"), when nitrogen bubbles form in the bloodstream, nerves, joints, and other sensitive or vital areas.
Direct skin contact with liquid nitrogen causes severe frostbite (cryogenic burns) within seconds, though not instantly on contact, depending on form of liquid nitrogen (liquid vs. mist) and surface area of the nitrogen-soaked material (soaked clothing or cotton causing more rapid damage than a spill of direct liquid to skin, which for a few seconds is protected by the Leidenfrost effect).
[edit] See also
[hide] v • d • e Diatomic Elements
HydrogenH2
NitrogenN2
OxygenO2
FluorineF2
ChlorineCl2
BromineBr2
IodineI2
AstatineAt2
Oxygen
Oxygen
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8
nitrogen ← oxygen → fluorine
-↑O↓S
Periodic table - Extended periodic table
General
Name, symbol, number
oxygen, O, 8
Chemical series
nonmetals, chalcogens
Group, period, block
16, 2, p
Appearance
colorless (gas)very pale blue (liquid)
Standard atomic weight
15.9994(3) g·mol−1
Electron configuration
1s2 2s2 2p4
Electrons per shell
2, 6
Physical properties
Phase
gas
Density
(0 °C, 101.325 kPa)1.429 g/L
Melting point
54.36 K(-218.79 °C, -361.82 °F)
Boiling point
90.20 K(-182.95 °C, -297.31 °F)
Critical point
154.59 K, 5.043 MPa
Heat of fusion
(O2) 0.444 kJ·mol−1
Heat of vaporization
(O2) 6.82 kJ·mol−1
Heat capacity
(25 °C) (O2)29.378 J·mol−1·K−1
Vapor pressure
P/Pa
1
10
100
1 k
10 k
100 k
at T/K
61
73
90
Atomic properties
Crystal structure
cubic
Oxidation states
−2, −1(neutral oxide)
Electronegativity
3.44 (Pauling scale)
Ionization energies(more)
1st: 1313.9 kJ·mol−1
2nd: 3388.3 kJ·mol−1
3rd: 5300.5 kJ·mol−1
Atomic radius
60 pm
Atomic radius (calc.)
48 pm
Covalent radius
73 pm
Van der Waals radius
152 pm
Miscellaneous
Magnetic ordering
paramagnetic
Thermal conductivity
(300 K) 26.58 m W·m−1·K−1
Speed of sound
(gas, 27 °C) 330 m/s
CAS registry number
7782-44-7
Selected isotopes
Main article: Isotopes of oxygen
iso
NA
half-life
DM
DE (MeV)
DP
16O
99.76%
O is stable with 8 neutrons
17O
0.038%
O is stable with 9 neutrons
18O
0.21%
O is stable with 10 neutrons
References
This box: view • talk • edit
For other uses, see Oxygen (disambiguation).
In science, oxygen (IPA: /ˈɒkˑsəˑdʒɪn/) is a chemical element with the chemical symbol O and atomic number 8. The word oxygen derives from two roots in Greek, οξύς (oxys) (acid, lit. sharp) and -γενής (-genēs) (producer, lit. begetter). It was recognized in 1777 by Antoine Lavoisier, who coined the name oxygen from the Greek roots mentioned above because he erroneously thought that it was a constituent of all acids. (The definition of acid has since been revised). Oxygen has a valency of 2. On Earth it is usually bonded to other elements covalently or ionically. Examples for common oxygen-containing compounds include water (H2O), sand (silica, SiO2), and rust (iron oxide, Fe2O3).
Diatomic oxygen (O2) is one of the two major components of air (20.95%). It is produced by plants during photosynthesis, and is necessary for aerobic respiration in animals. It is toxic to obligate anaerobic organisms and was a poisonous waste product for early life on Earth.
Triatomic oxygen (ozone, O3) forms through radiation in the upper layers of the atmosphere and acts as a shield against UV radiation.
Contents[hide]
1 Characteristics
1.1 Allotropes
2 Applications
3 History
4 Biological role
5 Occurrence
6 Production
7 Compounds
8 Isotopes
9 Precautions
9.1 Toxicity of O2
9.2 Toxicity and antibacterial use of other chemical oxygen forms
9.3 Combustion hazard
10 See also
11 References
12 External links//
Characteristics
The colour of liquid oxygen is a blue similar to sky blue. The phenomena are not related; the colour of the sky is due to Rayleigh scattering.
Dioxygen, O2, is a gas at standard conditions, consisting of 2-atom molecules. Elemental oxygen is most commonly encountered in this form, as 21% of Earth's atmosphere. Note that the double bond depicted here is an oversimplification; see triplet oxygen.
Ozone, O3, is a gas at standard conditions, consisting of 3-atom molecules. This oxygen allotrope is rare on Earth and is found mostly in the stratosphere.
The common allotrope of elemental oxygen on Earth, O2, is known as dioxygen.
At standard temperature and pressure, oxygen exists as a diatomic molecule with the formula O2, in which the two oxygen atoms are bonded to each other with the electron configuration of triplet oxygen. This bond has a bond order of two, and is thus often very grossly simplified in description as a double bond.[1] Triplet oxygen is the ground state of the oxygen molecule. The electron configuration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals. These orbitals are classified as antibonding, so the diatomic oxygen bond is weaker than the diatomic nitrogen bond, where all bonding molecular orbitals are filled. Though unpaired electrons are commonly associated with high reactivity in chemical compounds, triplet oxygen is relatively (and fortunately) nonreactive by comparison with most radicals.
Singlet oxygen, a name given to several higher energy species of molecular oxygen in which all the electron spins are paired, is much more reactive towards common organic molecules. In nature, singlet oxygen is commonly formed from water during photosynthesis, using the energy of sunlight. It is also produced by the immune system as a source of active oxygen. Carotenoids in photosynthetic organisms and possibly also in animals, play a major role in absorbing energy from singlet oxygen and converting it to the unexcited ground state, before it can cause harm to tissues.
Liquid O2 and solid O2 are clear substances with a light sea-blue color. In normal triplet form they are paramagnetic due to the spin magnetic moments of the unpaired electrons in the molecule, and the negative exchange energy between neighboring O2 molecules. Liquid oxygen is attracted to a magnet to a sufficient extent that a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet, in laboratory demonstrations. Liquid O2 is usually obtained by the fractional distillation of liquid air.
Oxygen is slightly soluble in water, but naturally occurring dissolved amounts are enough to support animal life (see below).
O2 has a bond length of 121 pm and a bond energy of 498 kJ/mol.[2]
Allotropes
Ozone, the less common triatomic allotrope of oxygen, is a poisonous gas with a distinct, sharp odor. It is thermodynamically unstable toward the more common dioxygen form. It is formed continuously in the upper atmosphere of the Earth by short-wave UV radiation, and also functions as a shield against UV radiation reaching the ground. Ozone has recently been found to be produced by the immune system as an antimicrobial (see below). Liquid and solid O3 (ozone) have a deeper blue color than ordinary oxygen, and they are unstable and explosive.
A newly 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.[3][4]
Applications
Uptake of oxygen from the air is the essential purpose of respiration, so oxygen supplementation has found use in medicine (as oxygen therapy). People who climb mountains or fly in non-pressurized aeroplanes sometimes have supplemental oxygen supplies; the reason is that increasing the proportion of oxygen in the breathing gas at low pressure acts to augment the inspired oxygen partial pressure nearer to that found at sea-level.
A home oxygen concentrator in situ in an emphysema patient's house. The model shown is the DeVILBISS LT 4000.
A notable application of oxygen as a very low-pressure breathing gas, is in modern spacesuits, where use of nearly pure oxygen at a total ambient pressure of about one third normal, results in normal blood partial pressures of oxygen. This trade-off of breathing gas content and needed pressure is important for space applications, because the issue of flexible spacesuits working at Earth sea-level pressures remains a technological challenge of aerospace technology. Oxygen is used in welding (such as the oxyacetylene torch), and in the industrial production of steel and methanol. Also, liquid oxygen finds use as a classic oxidizer in rocket propulsion.
Oxygen presents two spectrophotometric absorption bands peaking at 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.[citation needed] 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 a possibility to monitor the carbon cycle from satellites on a global scale.
Oxygen, as a supposed mild euphoric, has a history of recreational use (see oxygen bar). However, the reality of a pharmacological effect is doubtful, a metabolic boost being the most plausible explanation. Controlled tests of high oxygen mixtures in diving (see nitrox) and other activities, even at higher than normal pressures, demonstrated no particular effects on humans other than promotion of an increased tolerance to aerobic exercise.
In the 19th century, oxygen was often mixed with nitrous oxide to temper its analgesic effect. A stable 50% gaseous mixture (Entonox) is commonly used in medicine today as an analgesic. However, the common basic anaesthetic mixture is 30% oxygen with 70% nitrous oxide; the pain-suppressing effects, obviously, are due to the nitrous oxide and not to oxygen.
History
Oxygen was first described by Michał Sędziwój, a Polish alchemist and philosopher in the late 16th century. Sędziwój thought of the gas given off by warm niter (saltpeter) as "the elixir of life".[5]
Oxygen was more quantitatively discovered by the Swedish pharmacist Carl Wilhelm Scheele some time 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 discoveries in 1775 and Scheele in 1777; consequently Priestley is usually given the credit. Both Scheele and Priestley produced oxygen by heating mercuric oxide.
Scheele called the gas 'fire air' because it was the only known supporter of combustion. It was later called 'vital air' because it was and is vital for the existence of animal life.
The gas was named by Antoine Laurent Lavoisier, after Priestley's publication in 1775, from Greek roots meaning "acid-former". As noted, the name reflects the then-common incorrect belief that all acids contain oxygen. This is also the origin of the Japanese name of oxygen "sanso" (san=acid, so=element).
Oxygen was first time condensed in 1883 by professors of Jagiellonian University - Zygmunt Wróblewski (Polish chemist) Karol Olszewski (Polish physicist and chemist).
Biological role
Delayed oxygen build-up in earth's atmosphere and oceans in reaction to the evolution of oxygenic photosynthesis: A) no oxygen produced by biosphere, B) oxygen produced, but absorbed in oceans and by seabed rock, C) oxygen starts to gas out of the oceans, but is absorbed by land surfaces and formation of ozone layer.
Fluctuations of oxygen levels in the atmosphere over the past 500+ million years, with accompanying events: 1) Radiation of animal phyla (Cambrian explosion) - 2) First land plants - 3) Ordovician-Silurian extinction events - 4) Huge forests form on land, first land animals and seed plants - 5) Coal formation, first conifers, insect and amphibian giantism - 6) Low ocean levels, supercontinent Pangaea forms - 7) Permian-Triassic extinction event - 8) First primitive flowering plants and dinosaurs - 9) Triassic-Jurassic extinction event - 10) Age of dinosaurs - 11) Radiation of flowering plants - 12) Cretaceous-Tertiary extinction event - 13) Radiation of mammals
Molecular oxygen, O2, is essential for cellular respiration in all aerobic organisms. It is used as electron acceptor in the mitochondria to generate chemical energy in the form of adenosine triphosphate (ATP) during oxidative phosphorylation. During this reaction, oxygen is reduced to water. Conversely, free oxygen is produced in the biosphere through photolysis (light-driven oxidation and splitting) of water during photosynthesis in cyanobacteria, green algae and plants, thus closing the biological water-oxygen redox cycle.
Before the evolution of water oxidation in photosynthetic bacteria, oxygen was almost nonexistent in earth's atmosphere. Free oxygen first appeared in significant quantities during the Paleoproterozoic era (between 2.5 billion years ago and 1.6 billion years ago) as a product of the metabolic action of early anaerobes (archaea and bacteria). These organisms developed the mechanism of oxygen evolution between 3.5 and 2.7 billion years ago. At first, the produced oxygen dissolved in the oceans and reacted with iron. It started to "gas out" of the oxygen-saturated waters about 2.7 billion years ago as evident in the rusting of iron-rich terrestrial rocks starting around that time. The amount of oxygen in the atmosphere increased gradually at first and shot up rapidly around 2.2 to 1.7 billion years ago to about 10% of its present level.[6]
The development of an oxygen-rich atmosphere was one of the most important events in the history of life on earth. The presence of large amounts of dissolved and free oxygen in the oceans and atmosphere may have driven most of the anaerobic organisms then living to extinction during the oxygen catastrophe about 2.4 billion years ago. However, the high electronegativity of O2 creates a large potential energy drop for cellular respiration, thus enabling organisms using aerobic respiration to produce much more ATP than anaerobic organisms. This makes them so efficient that they have come to dominate earth's biosphere.[7] Photosynthesis and cellular respiration of oxygen allowed for the evolution of eukaryotic cells and ultimately complex multicellular organisms such as plants and animals.
The atmospheric abundance of free oxygen in later geological epochs and its gradual increase up to the present has been largely due to synthesis by photosynthetic organisms. Over the past 500 million years, oxygen levels fluctuated between 15 and 35% per volume. Towards the end of the Carboniferous era (coal age) about 300 million years ago, atmospheric oxygen levels reached a maximum of 35% by volume, allowing insects and amphibians with limiting respiratory systems to grow much larger than today's species. Today, oxygen is the second most common component of the earth's atmosphere (about 21% by volume) after nitrogen. About three quarters of the free element is being produced by algae and green microorganisms in the oceans, and one quarter from terrestrial plants[citation needed].
Occurrence
Annual mean sea surface dissolved oxygen for the World Ocean. Note more oxygen in cold water near the poles. [8]
Oxygen is the third most abundant chemical element in the universe by mass, after hydrogen and helium (see chemical element). Some of this oxygen was produced during stellar nucleosynthesis as a step in the CNO-II branch of the CNO cycle. However oxygen is primarily produced in massive stars. In stars with at least four times the Sun's mass, 16O nuclei are produced during the Carbon burning process. 16O can also be produced in stars with at least 8 times the Sun's mass as a result of photodisintegration during the Neon burning process.[9]
Oxygen is the most common component of the Earth's crust (49% by mass),[10] the second most common component of the Earth as a whole (28% by mass), the most common component of the world's oceans (86% by mass), and the second most common component of the Earth's atmosphere (20.947% by volume), second to nitrogen.
Elemental oxygen occurs not only in the atmosphere, but also as solution in the world's water bodies. At 25° C under 1 atm of air, a litre of water will dissolve about 6.04 cc (8.63 mg, 0.270 mmol) of oxygen, whereas sea water will dissolve about 4.9 cc (7.0 mg, 0.22 mmol). At 0° C the solubilities increase to 10.29 cc (14.7 mg, 0.460 mmol) for water and 8.0 cc (11.4 mg, 0.36 mmol) for sea water. This difference has important implications for ocean life, as polar oceans support a much higher density of life due to their oxygen content. [11]
See also Silicate minerals, Oxide minerals.
Production
Hoffman electrolysis apparatus used in electrolysis of water
Main article: Oxygen evolution
In nature, free oxygen is produced by the light-driven splitting of water during oxygenic photosynthesis in cyanobacteria, green algae and plants.[12] Algae produce about 73 to 87 percent of the net global production of oxygen, which makes it available to humans and other animals for respiration.[13] Another major source of oxygen is trees. Trees can absorb carbon dioxide at the rate of 26 pounds per year - especially young trees that are still growing - while releasing oxygen into the air. (Global Relief-Georgia).[14]
In the laboratory and industrially, oxygen can be produced through electrolysis of water into oxygen and hydrogen. A similar method is the electrocatalytic oxygen evolution from oxides and oxoacids. Chemical catalysts can be used as well, such as in chemical oxygen generators or oxygen candles that are used as part of the life support equipment on spacecraft and submarines.
Industrially, oxygen is typically produced in bulk quantity as a liquid produced by distillation from atmospheric air. In large quantities, the price of liquid oxygen (2001) is approximately $0.21/kg [15]. Since the primary cost of production is the energy cost of liquifying the air, the production cost will change as energy cost varies.
In the modern era, oxygen is increasingly obtained by non-cryogenic technolgies such as pressure swing adsorption (PSA) and vacuum-pressure swing adsorption (VPSA) technolgies [1].
Compounds
Due to its electronegativity, oxygen forms chemical bonds with almost all other elements hence the original definition of oxidation. The only elements known to escape the possibility of oxidation are a few of the noble gases, and fluorine. However, many noble metals (common examples: gold, platinum) resist direct chemical combination with oxygen, and substances like gold oxide must be formed by an indirect route.
The most familiar oxygen compound is water. Other well-known examples include silica (found in sand, glass, rock, etc.), and the compounds of carbon and oxygen, such as carbon dioxide (CO2), alcohols (R-OH), carbonyls, (R-CO-H or R-CO-R), 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. Phosphorus is biologically important in its oxygenated form as the phosphate (PO43−) ion. Many metals bond with oxygen atoms, such as iron in iron(III) oxide (Fe2O3), commonly called rust.
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.
One unexpected oxygen compound is dioxygen hexafluoroplatinate O2+PtF6−. It was discovered when Neil Bartlett was studying the properties of PtF6. He noticed a change in color when this compound was exposed to atmospheric air. Bartlett reasoned that xenon should be oxidized by PtF6. This led him to the discovery of xenon hexafluoroplatinate Xe+PtF6−.
See also: Category:Oxygen compounds
Isotopes
Main article: isotopes of oxygen
Oxygen has seventeen known isotopes with atomic masses ranging from 12.03 u to 28.06 u. Three are stable, 16O, 17O, and 18O, of which 16O is the most abundant (over 99.7%). The radioisotopes all have half-lives of less than three minutes. Nonetheless, 15O is used in positron emission tomography.
An atomic weight of 16 was assigned to oxygen prior to the definition of the unified atomic mass unit based upon 12C. Since physicists referred to 16O only, while chemists meant the naturally abundant mixture of isotopes, this led to slightly different atomic weight scales.
Precautions
Toxicity of O2
Main article: oxygen toxicity
Oxygen can be toxic at elevated partial pressures. Since oxygen partial pressure is the fraction of oxygen times the total pressure, elevated partial pressures can occur either from high oxygen fraction in breathing gas, or from high breathing gas pressure, or a combination of both. Oxygen toxicity usually begins to occur at partial pressures more than 0.5 atmospheres, or 2.5 times the normal sea-level oxygen partial pressure of about 0.2 atmospheres or bars. This means that at sea-level pressures, mixtures containing less than 50% oxygen are essentially non-toxic. However in medical applications (such as in ventilation gas mixtures in hospital applications) mixtures containing more than 50% oxygen can be expected to show lung toxicity, causing slow damage to the lungs over periods of days, with the rate of damage rising rapidly from mixtures between 50% and 100% oxygen. On the other hand, breathing 100% oxygen in space applications (such as in some modern spacesuits, or in early spacecraft such as the Apollo spacecraft), causes no damage due to the low total pressures (30% to 33% sea-level) used.[16] In the case of spacesuits, oxygen partial pressure in the breathing gas is typically about 0.30 bar (1.4 times normal), and oxygen partial pressure in the astronaut's blood (due to downward adjustments due to water vapor and CO2 in the alveoli) is close to sea-level normal of 0.2 bar.
In deep scuba diving and surface supplied diving and when using equipment which can provide high partial pressures of oxygen, such as rebreathers, oxygen toxicity to the lungs can occur, just as in medical applications. Due to the higher total pressures in these applications, the fraction of oxygen which produces lung damage may be considerably less than 50%. More importantly, under pressures higher than normal sea-level, a far more serious form of oxygen toxicity in the central nervous system may lead to generalized seizures. This form of oxygen toxicity usually occurs after several hours exposure to oxygen partial pressures over about 1.4 atmospheres (bars) (i.e. 7 times normal), with the time decreasing for higher pressures above this, and with great variation from person to person. At over three bars of oxygen partial pressure (15 times normal), seizures typically occur within minutes.
Toxicity and antibacterial use of other chemical oxygen forms
Certain derivatives of oxygen, such as ozone (O3), singlet oxygen, hydrogen peroxide, hydroxyl radicals and superoxide, are also highly toxic. Cells have developed various mechanisms to protect against all of these toxic compounds. For instance, the naturally-occurring glutathione can act as an antioxidant, as can bilirubin which is normally a breakdown product of hemoglobin. To protect against the destructive nature of peroxides, nearly every organism on earth has developed some form of the enzyme catalase, which very quickly disproportionates hydrogen peroxide into water and dioxygen. Another nearly universally present enzyme in living organisms (except for a few species of bacteria which use Mn2+ ions directly for the job) is superoxide dismutase. This family of enzymes disproportionates superoxide to oxygen and peroxide, which is then in turn dealt with, by catalase.
Immune systems of higher organisms have long made use of reactive forms of oxygen which they produce. Not only do antibodies catalyze production of peroxide from oxygen, it is now known that immune cells produce peroxide, superoxide, and singlet oxygen in the course of an immune response. Recently, singlet oxygen has been found to be a source of biologically-produced ozone: this reaction proceeds through an unusual compound dihydrogen trioxide, also known as trioxidane, (HOOOH) which is an antibody-catalyzed product of singlet oxygen and water. This compound in turn disproportionates to ozone and peroxide, providing two powerful antibacterials. The body's range of defense against all of these active oxidizing agents is hardly surprising, then, given their "deliberate" employment as antimicrobial agents in the immune response.[17]
Oxygen derivatives are prone to form free radicals, especially in metabolic processes. Because they can cause severe damage to cells and their DNA before they are dealt with, they form part of many theories of carcinogenesis and aging.
Combustion hazard
Highly concentrated sources of oxygen promote rapid combustion and therefore are fire and explosion hazards in the presence of fuels. Oxygen itself is not the fuel, but as a reactant, concentrated oxygen may allow combustion to proceed dangerously rapidly. The fire that killed the Apollo 1 crew on a test launchpad spread so rapidly because the capsule was pressurized with pure oxygen as would be usual in an actual flight, but to maintain positive pressure in the capsule, this was at slightly more than atmospheric pressure instead of the ⅓ normal pressure that would be used in flight. (See partial pressure.)
Hazards also apply to compounds of oxygen with a high oxidative potential, such as high concentration peroxides, chlorates, perchlorates, and dichromates; they also can often cause chemical burns.
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8
nitrogen ← oxygen → fluorine
-↑O↓S
Periodic table - Extended periodic table
General
Name, symbol, number
oxygen, O, 8
Chemical series
nonmetals, chalcogens
Group, period, block
16, 2, p
Appearance
colorless (gas)very pale blue (liquid)
Standard atomic weight
15.9994(3) g·mol−1
Electron configuration
1s2 2s2 2p4
Electrons per shell
2, 6
Physical properties
Phase
gas
Density
(0 °C, 101.325 kPa)1.429 g/L
Melting point
54.36 K(-218.79 °C, -361.82 °F)
Boiling point
90.20 K(-182.95 °C, -297.31 °F)
Critical point
154.59 K, 5.043 MPa
Heat of fusion
(O2) 0.444 kJ·mol−1
Heat of vaporization
(O2) 6.82 kJ·mol−1
Heat capacity
(25 °C) (O2)29.378 J·mol−1·K−1
Vapor pressure
P/Pa
1
10
100
1 k
10 k
100 k
at T/K
61
73
90
Atomic properties
Crystal structure
cubic
Oxidation states
−2, −1(neutral oxide)
Electronegativity
3.44 (Pauling scale)
Ionization energies(more)
1st: 1313.9 kJ·mol−1
2nd: 3388.3 kJ·mol−1
3rd: 5300.5 kJ·mol−1
Atomic radius
60 pm
Atomic radius (calc.)
48 pm
Covalent radius
73 pm
Van der Waals radius
152 pm
Miscellaneous
Magnetic ordering
paramagnetic
Thermal conductivity
(300 K) 26.58 m W·m−1·K−1
Speed of sound
(gas, 27 °C) 330 m/s
CAS registry number
7782-44-7
Selected isotopes
Main article: Isotopes of oxygen
iso
NA
half-life
DM
DE (MeV)
DP
16O
99.76%
O is stable with 8 neutrons
17O
0.038%
O is stable with 9 neutrons
18O
0.21%
O is stable with 10 neutrons
References
This box: view • talk • edit
For other uses, see Oxygen (disambiguation).
In science, oxygen (IPA: /ˈɒkˑsəˑdʒɪn/) is a chemical element with the chemical symbol O and atomic number 8. The word oxygen derives from two roots in Greek, οξύς (oxys) (acid, lit. sharp) and -γενής (-genēs) (producer, lit. begetter). It was recognized in 1777 by Antoine Lavoisier, who coined the name oxygen from the Greek roots mentioned above because he erroneously thought that it was a constituent of all acids. (The definition of acid has since been revised). Oxygen has a valency of 2. On Earth it is usually bonded to other elements covalently or ionically. Examples for common oxygen-containing compounds include water (H2O), sand (silica, SiO2), and rust (iron oxide, Fe2O3).
Diatomic oxygen (O2) is one of the two major components of air (20.95%). It is produced by plants during photosynthesis, and is necessary for aerobic respiration in animals. It is toxic to obligate anaerobic organisms and was a poisonous waste product for early life on Earth.
Triatomic oxygen (ozone, O3) forms through radiation in the upper layers of the atmosphere and acts as a shield against UV radiation.
Contents[hide]
1 Characteristics
1.1 Allotropes
2 Applications
3 History
4 Biological role
5 Occurrence
6 Production
7 Compounds
8 Isotopes
9 Precautions
9.1 Toxicity of O2
9.2 Toxicity and antibacterial use of other chemical oxygen forms
9.3 Combustion hazard
10 See also
11 References
12 External links//
Characteristics
The colour of liquid oxygen is a blue similar to sky blue. The phenomena are not related; the colour of the sky is due to Rayleigh scattering.
Dioxygen, O2, is a gas at standard conditions, consisting of 2-atom molecules. Elemental oxygen is most commonly encountered in this form, as 21% of Earth's atmosphere. Note that the double bond depicted here is an oversimplification; see triplet oxygen.
Ozone, O3, is a gas at standard conditions, consisting of 3-atom molecules. This oxygen allotrope is rare on Earth and is found mostly in the stratosphere.
The common allotrope of elemental oxygen on Earth, O2, is known as dioxygen.
At standard temperature and pressure, oxygen exists as a diatomic molecule with the formula O2, in which the two oxygen atoms are bonded to each other with the electron configuration of triplet oxygen. This bond has a bond order of two, and is thus often very grossly simplified in description as a double bond.[1] Triplet oxygen is the ground state of the oxygen molecule. The electron configuration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals. These orbitals are classified as antibonding, so the diatomic oxygen bond is weaker than the diatomic nitrogen bond, where all bonding molecular orbitals are filled. Though unpaired electrons are commonly associated with high reactivity in chemical compounds, triplet oxygen is relatively (and fortunately) nonreactive by comparison with most radicals.
Singlet oxygen, a name given to several higher energy species of molecular oxygen in which all the electron spins are paired, is much more reactive towards common organic molecules. In nature, singlet oxygen is commonly formed from water during photosynthesis, using the energy of sunlight. It is also produced by the immune system as a source of active oxygen. Carotenoids in photosynthetic organisms and possibly also in animals, play a major role in absorbing energy from singlet oxygen and converting it to the unexcited ground state, before it can cause harm to tissues.
Liquid O2 and solid O2 are clear substances with a light sea-blue color. In normal triplet form they are paramagnetic due to the spin magnetic moments of the unpaired electrons in the molecule, and the negative exchange energy between neighboring O2 molecules. Liquid oxygen is attracted to a magnet to a sufficient extent that a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet, in laboratory demonstrations. Liquid O2 is usually obtained by the fractional distillation of liquid air.
Oxygen is slightly soluble in water, but naturally occurring dissolved amounts are enough to support animal life (see below).
O2 has a bond length of 121 pm and a bond energy of 498 kJ/mol.[2]
Allotropes
Ozone, the less common triatomic allotrope of oxygen, is a poisonous gas with a distinct, sharp odor. It is thermodynamically unstable toward the more common dioxygen form. It is formed continuously in the upper atmosphere of the Earth by short-wave UV radiation, and also functions as a shield against UV radiation reaching the ground. Ozone has recently been found to be produced by the immune system as an antimicrobial (see below). Liquid and solid O3 (ozone) have a deeper blue color than ordinary oxygen, and they are unstable and explosive.
A newly 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.[3][4]
Applications
Uptake of oxygen from the air is the essential purpose of respiration, so oxygen supplementation has found use in medicine (as oxygen therapy). People who climb mountains or fly in non-pressurized aeroplanes sometimes have supplemental oxygen supplies; the reason is that increasing the proportion of oxygen in the breathing gas at low pressure acts to augment the inspired oxygen partial pressure nearer to that found at sea-level.
A home oxygen concentrator in situ in an emphysema patient's house. The model shown is the DeVILBISS LT 4000.
A notable application of oxygen as a very low-pressure breathing gas, is in modern spacesuits, where use of nearly pure oxygen at a total ambient pressure of about one third normal, results in normal blood partial pressures of oxygen. This trade-off of breathing gas content and needed pressure is important for space applications, because the issue of flexible spacesuits working at Earth sea-level pressures remains a technological challenge of aerospace technology. Oxygen is used in welding (such as the oxyacetylene torch), and in the industrial production of steel and methanol. Also, liquid oxygen finds use as a classic oxidizer in rocket propulsion.
Oxygen presents two spectrophotometric absorption bands peaking at 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.[citation needed] 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 a possibility to monitor the carbon cycle from satellites on a global scale.
Oxygen, as a supposed mild euphoric, has a history of recreational use (see oxygen bar). However, the reality of a pharmacological effect is doubtful, a metabolic boost being the most plausible explanation. Controlled tests of high oxygen mixtures in diving (see nitrox) and other activities, even at higher than normal pressures, demonstrated no particular effects on humans other than promotion of an increased tolerance to aerobic exercise.
In the 19th century, oxygen was often mixed with nitrous oxide to temper its analgesic effect. A stable 50% gaseous mixture (Entonox) is commonly used in medicine today as an analgesic. However, the common basic anaesthetic mixture is 30% oxygen with 70% nitrous oxide; the pain-suppressing effects, obviously, are due to the nitrous oxide and not to oxygen.
History
Oxygen was first described by Michał Sędziwój, a Polish alchemist and philosopher in the late 16th century. Sędziwój thought of the gas given off by warm niter (saltpeter) as "the elixir of life".[5]
Oxygen was more quantitatively discovered by the Swedish pharmacist Carl Wilhelm Scheele some time 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 discoveries in 1775 and Scheele in 1777; consequently Priestley is usually given the credit. Both Scheele and Priestley produced oxygen by heating mercuric oxide.
Scheele called the gas 'fire air' because it was the only known supporter of combustion. It was later called 'vital air' because it was and is vital for the existence of animal life.
The gas was named by Antoine Laurent Lavoisier, after Priestley's publication in 1775, from Greek roots meaning "acid-former". As noted, the name reflects the then-common incorrect belief that all acids contain oxygen. This is also the origin of the Japanese name of oxygen "sanso" (san=acid, so=element).
Oxygen was first time condensed in 1883 by professors of Jagiellonian University - Zygmunt Wróblewski (Polish chemist) Karol Olszewski (Polish physicist and chemist).
Biological role
Delayed oxygen build-up in earth's atmosphere and oceans in reaction to the evolution of oxygenic photosynthesis: A) no oxygen produced by biosphere, B) oxygen produced, but absorbed in oceans and by seabed rock, C) oxygen starts to gas out of the oceans, but is absorbed by land surfaces and formation of ozone layer.
Fluctuations of oxygen levels in the atmosphere over the past 500+ million years, with accompanying events: 1) Radiation of animal phyla (Cambrian explosion) - 2) First land plants - 3) Ordovician-Silurian extinction events - 4) Huge forests form on land, first land animals and seed plants - 5) Coal formation, first conifers, insect and amphibian giantism - 6) Low ocean levels, supercontinent Pangaea forms - 7) Permian-Triassic extinction event - 8) First primitive flowering plants and dinosaurs - 9) Triassic-Jurassic extinction event - 10) Age of dinosaurs - 11) Radiation of flowering plants - 12) Cretaceous-Tertiary extinction event - 13) Radiation of mammals
Molecular oxygen, O2, is essential for cellular respiration in all aerobic organisms. It is used as electron acceptor in the mitochondria to generate chemical energy in the form of adenosine triphosphate (ATP) during oxidative phosphorylation. During this reaction, oxygen is reduced to water. Conversely, free oxygen is produced in the biosphere through photolysis (light-driven oxidation and splitting) of water during photosynthesis in cyanobacteria, green algae and plants, thus closing the biological water-oxygen redox cycle.
Before the evolution of water oxidation in photosynthetic bacteria, oxygen was almost nonexistent in earth's atmosphere. Free oxygen first appeared in significant quantities during the Paleoproterozoic era (between 2.5 billion years ago and 1.6 billion years ago) as a product of the metabolic action of early anaerobes (archaea and bacteria). These organisms developed the mechanism of oxygen evolution between 3.5 and 2.7 billion years ago. At first, the produced oxygen dissolved in the oceans and reacted with iron. It started to "gas out" of the oxygen-saturated waters about 2.7 billion years ago as evident in the rusting of iron-rich terrestrial rocks starting around that time. The amount of oxygen in the atmosphere increased gradually at first and shot up rapidly around 2.2 to 1.7 billion years ago to about 10% of its present level.[6]
The development of an oxygen-rich atmosphere was one of the most important events in the history of life on earth. The presence of large amounts of dissolved and free oxygen in the oceans and atmosphere may have driven most of the anaerobic organisms then living to extinction during the oxygen catastrophe about 2.4 billion years ago. However, the high electronegativity of O2 creates a large potential energy drop for cellular respiration, thus enabling organisms using aerobic respiration to produce much more ATP than anaerobic organisms. This makes them so efficient that they have come to dominate earth's biosphere.[7] Photosynthesis and cellular respiration of oxygen allowed for the evolution of eukaryotic cells and ultimately complex multicellular organisms such as plants and animals.
The atmospheric abundance of free oxygen in later geological epochs and its gradual increase up to the present has been largely due to synthesis by photosynthetic organisms. Over the past 500 million years, oxygen levels fluctuated between 15 and 35% per volume. Towards the end of the Carboniferous era (coal age) about 300 million years ago, atmospheric oxygen levels reached a maximum of 35% by volume, allowing insects and amphibians with limiting respiratory systems to grow much larger than today's species. Today, oxygen is the second most common component of the earth's atmosphere (about 21% by volume) after nitrogen. About three quarters of the free element is being produced by algae and green microorganisms in the oceans, and one quarter from terrestrial plants[citation needed].
Occurrence
Annual mean sea surface dissolved oxygen for the World Ocean. Note more oxygen in cold water near the poles. [8]
Oxygen is the third most abundant chemical element in the universe by mass, after hydrogen and helium (see chemical element). Some of this oxygen was produced during stellar nucleosynthesis as a step in the CNO-II branch of the CNO cycle. However oxygen is primarily produced in massive stars. In stars with at least four times the Sun's mass, 16O nuclei are produced during the Carbon burning process. 16O can also be produced in stars with at least 8 times the Sun's mass as a result of photodisintegration during the Neon burning process.[9]
Oxygen is the most common component of the Earth's crust (49% by mass),[10] the second most common component of the Earth as a whole (28% by mass), the most common component of the world's oceans (86% by mass), and the second most common component of the Earth's atmosphere (20.947% by volume), second to nitrogen.
Elemental oxygen occurs not only in the atmosphere, but also as solution in the world's water bodies. At 25° C under 1 atm of air, a litre of water will dissolve about 6.04 cc (8.63 mg, 0.270 mmol) of oxygen, whereas sea water will dissolve about 4.9 cc (7.0 mg, 0.22 mmol). At 0° C the solubilities increase to 10.29 cc (14.7 mg, 0.460 mmol) for water and 8.0 cc (11.4 mg, 0.36 mmol) for sea water. This difference has important implications for ocean life, as polar oceans support a much higher density of life due to their oxygen content. [11]
See also Silicate minerals, Oxide minerals.
Production
Hoffman electrolysis apparatus used in electrolysis of water
Main article: Oxygen evolution
In nature, free oxygen is produced by the light-driven splitting of water during oxygenic photosynthesis in cyanobacteria, green algae and plants.[12] Algae produce about 73 to 87 percent of the net global production of oxygen, which makes it available to humans and other animals for respiration.[13] Another major source of oxygen is trees. Trees can absorb carbon dioxide at the rate of 26 pounds per year - especially young trees that are still growing - while releasing oxygen into the air. (Global Relief-Georgia).[14]
In the laboratory and industrially, oxygen can be produced through electrolysis of water into oxygen and hydrogen. A similar method is the electrocatalytic oxygen evolution from oxides and oxoacids. Chemical catalysts can be used as well, such as in chemical oxygen generators or oxygen candles that are used as part of the life support equipment on spacecraft and submarines.
Industrially, oxygen is typically produced in bulk quantity as a liquid produced by distillation from atmospheric air. In large quantities, the price of liquid oxygen (2001) is approximately $0.21/kg [15]. Since the primary cost of production is the energy cost of liquifying the air, the production cost will change as energy cost varies.
In the modern era, oxygen is increasingly obtained by non-cryogenic technolgies such as pressure swing adsorption (PSA) and vacuum-pressure swing adsorption (VPSA) technolgies [1].
Compounds
Due to its electronegativity, oxygen forms chemical bonds with almost all other elements hence the original definition of oxidation. The only elements known to escape the possibility of oxidation are a few of the noble gases, and fluorine. However, many noble metals (common examples: gold, platinum) resist direct chemical combination with oxygen, and substances like gold oxide must be formed by an indirect route.
The most familiar oxygen compound is water. Other well-known examples include silica (found in sand, glass, rock, etc.), and the compounds of carbon and oxygen, such as carbon dioxide (CO2), alcohols (R-OH), carbonyls, (R-CO-H or R-CO-R), 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. Phosphorus is biologically important in its oxygenated form as the phosphate (PO43−) ion. Many metals bond with oxygen atoms, such as iron in iron(III) oxide (Fe2O3), commonly called rust.
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.
One unexpected oxygen compound is dioxygen hexafluoroplatinate O2+PtF6−. It was discovered when Neil Bartlett was studying the properties of PtF6. He noticed a change in color when this compound was exposed to atmospheric air. Bartlett reasoned that xenon should be oxidized by PtF6. This led him to the discovery of xenon hexafluoroplatinate Xe+PtF6−.
See also: Category:Oxygen compounds
Isotopes
Main article: isotopes of oxygen
Oxygen has seventeen known isotopes with atomic masses ranging from 12.03 u to 28.06 u. Three are stable, 16O, 17O, and 18O, of which 16O is the most abundant (over 99.7%). The radioisotopes all have half-lives of less than three minutes. Nonetheless, 15O is used in positron emission tomography.
An atomic weight of 16 was assigned to oxygen prior to the definition of the unified atomic mass unit based upon 12C. Since physicists referred to 16O only, while chemists meant the naturally abundant mixture of isotopes, this led to slightly different atomic weight scales.
Precautions
Toxicity of O2
Main article: oxygen toxicity
Oxygen can be toxic at elevated partial pressures. Since oxygen partial pressure is the fraction of oxygen times the total pressure, elevated partial pressures can occur either from high oxygen fraction in breathing gas, or from high breathing gas pressure, or a combination of both. Oxygen toxicity usually begins to occur at partial pressures more than 0.5 atmospheres, or 2.5 times the normal sea-level oxygen partial pressure of about 0.2 atmospheres or bars. This means that at sea-level pressures, mixtures containing less than 50% oxygen are essentially non-toxic. However in medical applications (such as in ventilation gas mixtures in hospital applications) mixtures containing more than 50% oxygen can be expected to show lung toxicity, causing slow damage to the lungs over periods of days, with the rate of damage rising rapidly from mixtures between 50% and 100% oxygen. On the other hand, breathing 100% oxygen in space applications (such as in some modern spacesuits, or in early spacecraft such as the Apollo spacecraft), causes no damage due to the low total pressures (30% to 33% sea-level) used.[16] In the case of spacesuits, oxygen partial pressure in the breathing gas is typically about 0.30 bar (1.4 times normal), and oxygen partial pressure in the astronaut's blood (due to downward adjustments due to water vapor and CO2 in the alveoli) is close to sea-level normal of 0.2 bar.
In deep scuba diving and surface supplied diving and when using equipment which can provide high partial pressures of oxygen, such as rebreathers, oxygen toxicity to the lungs can occur, just as in medical applications. Due to the higher total pressures in these applications, the fraction of oxygen which produces lung damage may be considerably less than 50%. More importantly, under pressures higher than normal sea-level, a far more serious form of oxygen toxicity in the central nervous system may lead to generalized seizures. This form of oxygen toxicity usually occurs after several hours exposure to oxygen partial pressures over about 1.4 atmospheres (bars) (i.e. 7 times normal), with the time decreasing for higher pressures above this, and with great variation from person to person. At over three bars of oxygen partial pressure (15 times normal), seizures typically occur within minutes.
Toxicity and antibacterial use of other chemical oxygen forms
Certain derivatives of oxygen, such as ozone (O3), singlet oxygen, hydrogen peroxide, hydroxyl radicals and superoxide, are also highly toxic. Cells have developed various mechanisms to protect against all of these toxic compounds. For instance, the naturally-occurring glutathione can act as an antioxidant, as can bilirubin which is normally a breakdown product of hemoglobin. To protect against the destructive nature of peroxides, nearly every organism on earth has developed some form of the enzyme catalase, which very quickly disproportionates hydrogen peroxide into water and dioxygen. Another nearly universally present enzyme in living organisms (except for a few species of bacteria which use Mn2+ ions directly for the job) is superoxide dismutase. This family of enzymes disproportionates superoxide to oxygen and peroxide, which is then in turn dealt with, by catalase.
Immune systems of higher organisms have long made use of reactive forms of oxygen which they produce. Not only do antibodies catalyze production of peroxide from oxygen, it is now known that immune cells produce peroxide, superoxide, and singlet oxygen in the course of an immune response. Recently, singlet oxygen has been found to be a source of biologically-produced ozone: this reaction proceeds through an unusual compound dihydrogen trioxide, also known as trioxidane, (HOOOH) which is an antibody-catalyzed product of singlet oxygen and water. This compound in turn disproportionates to ozone and peroxide, providing two powerful antibacterials. The body's range of defense against all of these active oxidizing agents is hardly surprising, then, given their "deliberate" employment as antimicrobial agents in the immune response.[17]
Oxygen derivatives are prone to form free radicals, especially in metabolic processes. Because they can cause severe damage to cells and their DNA before they are dealt with, they form part of many theories of carcinogenesis and aging.
Combustion hazard
Highly concentrated sources of oxygen promote rapid combustion and therefore are fire and explosion hazards in the presence of fuels. Oxygen itself is not the fuel, but as a reactant, concentrated oxygen may allow combustion to proceed dangerously rapidly. The fire that killed the Apollo 1 crew on a test launchpad spread so rapidly because the capsule was pressurized with pure oxygen as would be usual in an actual flight, but to maintain positive pressure in the capsule, this was at slightly more than atmospheric pressure instead of the ⅓ normal pressure that would be used in flight. (See partial pressure.)
Hazards also apply to compounds of oxygen with a high oxidative potential, such as high concentration peroxides, chlorates, perchlorates, and dichromates; they also can often cause chemical burns.
Salt
Salt
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From Wikipedia, the free encyclopedia
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For other uses, see Salt (disambiguation).
For the chemical properties of salt, see sodium chloride; for the chemical term, see salt (chemistry).
Salt is mostly sodium chloride (NaCl). This salt shaker also contains grains of rice, which some use to prevent caking
Brine being boiled down to pure salt in Zigong, China
Salt is a mineral essential for animal life, composed primarily of sodium chloride. Salt for human consumption is produced in different forms: unrefined salt (such as sea salt), refined salt (table salt), and iodized salt. It is a crystalline solid, white, pale pink or light grey in color, normally obtained from sea water or rock deposits. Edible rock salts may be slightly greyish in color due to this mineral content.
Sodium and chlorine, the two components of salt, are necessary for the survival of all living creatures, including humans, but they need not be consumed as salt, where they are found together in very concentrated form. Some isolated cultures, such as the Yanomami in South America, have been found to consume little salt.[1] Salt is involved in regulating the water content (fluid balance) of the body. Salt flavor is one of the basic tastes. Salt cravings may be caused by trace mineral deficiencies as well as by a deficiency of sodium chloride itself.
Overconsumption of salt can increase the risk of health problems, including high blood pressure. In food preparation, salt is used as a preservative and as a seasoning.
Contents[hide]
1 History
2 In religion
3 Forms of salt
3.1 Unrefined salt
3.2 Refined salt
3.3 Table salt
3.3.1 Iodized salt
3.3.2 Fluorinated salt
3.4 Salty condiments
4 Health effects
5 Recommended intake
6 Labeling
7 Campaigns
8 Salt substitutes
9 Production trends
10 See also
11 References
12 External links
12.1 Salt and health
13 Further reading//
[edit] History
See main article: History of salt
At the dawn of civilization, salt's preservative ability eliminated dependency on the seasonal availability of food, allowed travel over long distances, and was a vital food additive. However, because salt (NaCl) was difficult to obtain, it became a highly valued trade item throughout history. Until the 1900s, salt was one of the prime movers of national economies and wars. Salt was often taxed; research has discovered this practice to have existed as early as the 20th century BC in China. By the Middle Ages, caravans consisting of as many as forty thousand camels traversed four hundred miles of the Sahara bearing salt, sometimes trading it for slaves.[citation needed]
The first registers of salt use were produced around 4000 B.C. in Egypt, and later in Greece and Rome. Salt was very valuable and used to preserve and flavor foods. In Ancient Rome, salt was used as a currency. The Latin word salarium; meaning a payment made in salt, is the root of the word "salary." Unfortunately for those paid with salt, it was easily ruined by rain and other weather conditions. Payments to Roman workers and soldiers were made in salt.[2] Salt was also given to the parents of the groom in marriage until the 8th century. [attribution needed]
From the Phoenicians dates the evidence of harvesting solid salt from the sea. They also exported it to other civilizations. As a result of the increased salt supply from the sea, the value of salt depreciated. The harvest method used was flooding plains of land with seawater, then leaving the plains to dry. After the water dried, the salt which was left was collected and sold.
In the Mali Empire, merchants in 12th century Timbuktu—the gateway to the Sahara Desert and the seat of scholars—valued salt (NaCl) enough to buy it for its weight in gold; this trade led to the legends of the incredibly wealthy city of Timbuktu, and fueled inflation in Europe, which was importing the salt.[3]
During his protests in India, Mohandas Gandhi performed the famous salt march to challenge the British-imposed monopoly on salt.
[edit] In religion
Among the ancients, as with ourselves, "sol" (sun) and "sal" (salt) were considered essential to the maintenance of life.
There are thirty-five references (verses) to salt in the Bible (King James Version), the most familiar probably being the story of Lot's wife, who was turned into a pillar of salt when she disobeyed the angels and looked back at the wicked city of Sodom (Genesis 19:26). In the Sermon on the Mount, Jesus also referred to his followers as the "salt of the earth". The apostle Paul also encouraged Christians to "let your conversation be always full of grace, seasoned with salt" (Colossians 4:6) so that when others enquire about their beliefs, the Christian's answer generates a 'thirst' to know more about Christ. Salt is mandatory in the rite of the Tridentine Mass. Salt is used in the third item (which includes an Exorcism) of the Celtic Consecration (refer Gallican rite) that is employed in the Consecration of a Church. Salt may be added to the water "where it is customary" in the Roman Catholic rite of Holy water. The earliest Biblical mention of salt appears to be in reference to the destruction of Sodom and Gomorrah (Genesis 19:24-26)When King Abimelech destroyed the city of Shechem, held to have occurred in the thirteenth century BCE., he is said to have "sowed salt on it," this phrase expressing the completeness of its ruin. (Judges 9:45.)
In the native Japanese religion Shinto, salt is used for ritual purification of locations and people, such as in Sumo Wrestling.
In Aztec mythology, Huixtocihuatl was a fertility goddess who presided over salt and salt water.
[edit] Forms of salt
[edit] Unrefined salt
Main articles: Sea salt, Halite, and Fleur de sel
Different natural salts have different mineralities, giving each one a unique flavor. Fleur de sel, natural sea salt harvested by hand, has a unique flavor varying from region to region.
Some assert that unrefined sea salt is more healthy than refined salts.[4] However, completely raw sea salt is bitter due to magnesium and calcium compounds, and thus is rarely eaten. Other people think that raw sea and rock salts do not contain sufficient iodine salts to prevent iodine deficiency diseases like hypothyroidism.[5]
[edit] Refined salt
Salt evaporation pond in Alexandria, Egypt
Refined salt, which is most widely used presently, is mainly sodium chloride. Food grade salt accounts for only a small part of salt production in industrialised countries (3% in Europe[6]) although world-wide, food uses account for 17.5% of salt production[7]. The majority is sold for industrial use, from manufacturing pulp and paper to setting dyes in textiles and fabric, to producing soaps and detergents, and has great commercial value.
Salt mounds in Bolivia. Salt is harvested in the traditional method.
The manufacture and use of salt is one of the oldest chemical industries.[8] Salt is also obtained by evaporation of sea water, usually in shallow basins warmed by sunlight;[9] salt so obtained was formerly called bay salt, and is now often called sea salt or solar salt. Today, most refined salt is prepared from rock salt: mineral deposits high in salt.[citation needed] These rock salt deposits were formed by the evaporation of ancient salt lakes.[10] These deposits may be mined conventionally or through the injection of water. Injected water dissolves the salt, and the brine solution can be pumped to the surface where the salt is collected.
After the raw salt is obtained, it is refined to purify it and improve its storage and handling characteristics. Purification usually involves recrystallization. In recrystallization, a brine solution is treated with chemicals that precipitate most impurities (largely magnesium and calcium salts).[11] Multiple stages of evaporation are then used to collect pure sodium chloride crystals, which are kiln-dried.
Since the 1950's it has been common to add a trace of sodium hexacyanoferrate II to the brine, this acts as an anticaking agent by promoting irregular crystals.[12] Other anticaking agents (and potassium iodide, for iodised salt) are generally added after crystallization.[citation needed] These agents are hygroscopic chemicals which absorb humidity, keeping the salt crystals from sticking together. Some anticaking agents used are tricalcium phosphate, calcium or magnesium carbonates, fatty acid salts (acid salts), magnesium oxide, silicon dioxide, calcium silicate, sodium alumino-silicate, and alumino-calcium silicate. Concerns have been raised regarding the possible toxic effects of aluminium in the latter two compounds,[citation needed] however both the European Union and the United States Food and Drug Administration (FDA) permit their use.[13] The refined salt is then ready for packing and distribution.
[edit] Table salt
Single-serving salt packets
Table salt is refined salt, 99% sodium chloride.[14][15] It usually contains substances that make it free flowing (anticaking agents) such as sodium silicoaluminate or magnesium carbonate. It is common practice to put a few grains of uncooked rice in salt shakers to absorb extra moisture when anticaking agents are not enough.
[edit] Iodized salt
Main article: History of iodised salt
Iodized salt (BrE: iodised salt), table salt mixed with a minute amount of sodium iodide, iodate, or sometimes potassium iodide, is used to help reduce the chance of iodine deficiency in humans. Iodine deficiency commonly leads to thyroid gland problems, specifically endemic goiter. Endemic goiter is a disease characterized by a swelling of the thyroid gland, usually resulting in a bulbous protrusion on the neck. While only tiny quantities of iodine are required in a diet to prevent goiter, the United States Food and Drug Administration recommends (21 CFR 101.9 (c)(8)(iv)) 150 microgrammes of iodine per day for both men and women, and there are many places around the world where natural levels of iodine in the soil are low and the iodine is not taken up by vegetables.
Today, iodized salt is more common in the United States, Australia and New Zealand than in Britain. Table salt is also often iodized—a small amount of potassium iodide (in the US) or potassium iodate (in the EU) is added as an important dietary supplement. Table salt is mainly employed in cooking and as a table condiment. Iodized table salt has significantly reduced disorders of iodine deficiency in countries where it is used.[16] Iodine is important to prevent the insufficient production of thyroid hormones (hypothyroidism), which can cause goitre, cretinism in children, and myxedema in adults.
[edit] Fluorinated salt
In some European countries where fluoridation of drinking water is not practiced, fluorinated table salt is available. In France, 35% of sold table salt contains sodium or potassium fluoride.[17] Another additive, especially important for pregnant women is Folic acid (B vitamin) giving the table salt a yellow color.
[edit] Salty condiments
In many Asian cultures, table salt is not traditionally used as a condiment.[18]However, condiments such as soy sauce, fish sauce and oyster sauce tend to have a high salt content and fill much the same role as a salt-providing table condiment that table salt serves in western cultures.
[edit] Health effects
Sodium is one of the primary electrolytes in the body. All three electrolytes (sodium, potassium, and calcium) are available in unrefined salt, as are other vital minerals needed for optimal bodily function. Too much or too little salt in the diet can lead to muscle cramps, dizziness, or even an electrolyte disturbance, which can cause severe, even fatal, neurological problems.[19] Drinking too much water, with insufficient salt intake, puts a person at risk of water intoxication. Salt is even sometimes used as a health aid, such as in treatment of dysautonomia.[20]
People's risk for disease due to insufficient or excessive salt intake varies due to biochemical individuality. Some have asserted that while the risks of consuming too much salt are real, the risks have been exaggerated for most people, or that the studies done on the consumption of salt can be interpreted in many different ways.[21] [22]
Excess salt consumption has been linked to:
exercise-induced asthma.[23] On the other hand, another source counters, "…we still don't know whether salt contributes to asthma. If there is a link then it's very weak…".[24]
heartburn[25].
osteoporosis: One report shows that a high salt diet does reduce bone density in girls.[26]. Yet "While high salt intakes have been associated with detrimental effects on bone health, there are insufficient data to draw firm conclusions." ([27], p3)
Gastric cancer (Stomach cancer) is associated with high levels of sodium, "but the evidence does not generally relate to foods typically consumed in the UK." ([27], p18) However, in Japan, salt consumption is higher.[28]
hypertension (high blood pressure): "Since 1994, the evidence of an association between dietary salt intakes and blood pressure has increased. The data have been consistent in various study populations and across the age range in adults." ([27] p3). "The CMO [Chief Medical Officer] of England, in his Annual Report (DH, 2001), highlighted that people with high blood pressure are three times more likely to develop heart disease and stroke, and twice as likely to die from these diseases than those with normal levels."([27], p14). Professor Dr. Diederick Grobbee claims that there is no evidence of a causal link between salt intake and mortality or cardiovascular events.[29]. One study found that low urinary sodium is associated with greater risk of myocardial infarction among treated hypertensive men [30].
left ventricular hypertrophy (cardiac enlargement): "Evidence suggests that high salt intake causes left ventricular hypertrophy, a strong risk factor for cardiovascular disease, independently of blood pressure effects." ([27] p3) "…there is accumulating evidence that high salt intake predicts left ventricular hypertrophy." ([31], p12) Excessive salt (sodium) intake, combined with an inadequate intake of water, can cause hypernatremia. It can exacerbate renal disease.[19]
edema (BE: oedema): A decrease in salt intake has been suggested to treat edema (BE: oedema) (fluid retention).[32][19]
duodenal ulcers and gastric ulcers[33]
A large scale study by Nancy Cook et al shows that people with high-normal[1] blood pressure who significantly reduced the amount of salt in their diet decreased their chances of developing cardiovascular disease by 25% over the following 10 to 15 years. Their risk of dying from cardiovascular disease decreased by 20%.[34][35]
[edit] Recommended intake
Sea salt and peppercorns.
A salt mill for sea salt.
This section summarizes the salt intake recommended by the health agencies of various countries. Recommendations tend to be similar. Note that targets for the population as a whole tend to be pragmatic (what is achievable) while advice for an individual is ideal (what is best for health). For example, in the UK target for the population is "eat no more than 6 g a day" but for a person is 4 g.
Intakes can be expressed variously as salt or sodium and in various units.
1 g sodium = 1,000 mg sodium = 42 mmol sodium = 2.5 g salt
United Kingdom: In 2003, the UK's Scientific Advisory Committee on Nutrition (SACN) recommended that, for a typical adult, the Reference Nutrient Intake is 4 g salt per day (1.6 g or 70 mmol sodium). However, average adult intake is two and a half times the Reference Nutrient Intake for sodium. "Although accurate data are not available for children, conservative estimates indicate that, on a body weight basis, the average salt intake of children is higher than that of adults." SACN aimed for an achievable target reduction in average intake of salt to 6 g per day (2.4 g or 100 mmol sodium) — this is roughly equivalent to a teaspoonful of salt. The SACN recommendations for children are:
0–6 months old: less than 1 g/day
7–12 months: 1 g/day
1–3 years: 2 g/day
4–6 years: 3 g/day
7–10 years: 5 g/day
11–14 years: 6 g/day
SACN states, "The target salt intakes set for adults and children do not represent ideal or optimum consumption levels, but achievable population goals."[27]
Republic of Ireland: The Food Safety Authority of Ireland endorses the UK targets "emphasising that the RDA of 1.6 g sodium (4 g salt) per day should form the basis of advice targeted at individuals as distinct from the population health target of a mean salt intake of 6 g per day."([31], p16)
Canada: Health Canada recommends an Adequate Intake (AI) and an Upper Limit (UL) in terms of sodium.
0–6 months old: 0.12 g/day (AI)
7–12 months: 0.37 g/day (AI)
1–3 years: 1 g/day (AI) 1.5 g/day (UL)
4–8 years: 1.2/day (AI) 1.9 g/day (UL)
9–13 years: 1.5 g/day (AI) 2.2 g/day (UL)
14–50 years: 1.5 g/day (AI) 2.3 g/day (UL)
51–70 years: 1.3 g/day (AI) 2.3 g/day (UL)
70 years and older: 1.2 g/day (AI) 2.3 g/day (UL)[36]
New Zealand
Adequate Intake (AI) 0.46 – 0.92 g sodium = 1.2 – 2.3g salt
Upper Limit (UL)) 2.3 g sodium = 5.8 g salt[37]
Australia: The recommended dietary intake (RDI) is 0.92 g–2.3 g sodium per day (= 2.3 g–5.8 g salt)[38]
USA: The Food and Drug Administration itself does not make a recommendation[39] but refers readers to Dietary Guidelines for Americans 2005. These suggest that US citizens should consume less than 2,300 mg of sodium (= 2.3 g sodium = 5.8 g salt) per day. [40]
[edit] Labeling
UK: The Food Standards Agency defines the level of salt in foods as follows: "High is more than 1.5g salt per 100g (or 0.6g sodium). Low is 0.3g salt or less per 100g (or 0.1g sodium). If the amount of salt per 100g is in between these figures, then that is a medium level of salt." In the UK, foods produced by some supermarkets and manufacturers have ‘traffic light’ colors on the front of the pack: Red (High), Amber (Medium), or Green (Low).[41]
USA: The FDA Food Labeling Guide stipulates whether a food can be labelled as "free", "low", or "reduced/less" in respect of sodium. When other health claims are made about a food (e.g. low in fat, calories, etc.), a disclosure statement is required if the food exceeds 480mg of sodium per 'serving.'[42]
[edit] Campaigns
In 2004, Britain's Food Standards Agency started a public health campaign called "Salt - Watch it", which recommends no more than 6g of salt per day; it features a character called Sid the Slug and was criticised by the Salt Manufacturers Association (SMA).[43] The Advertising Standards Authority did not uphold the SMA complaint in its adjudication.[44]. In March 2007, the FSA launched the third phase of their campaign with the slogan "Salt. Is your food full of it?" fronted by comedienne Jenny Eclair.[45]
The Menzies Research Institute in Tasmania, Australia, maintains a website [46] dedicated to educating people about the potential problems of a salt-laden diet.
[edit] Salt substitutes
Main article: Salt substitute
Salt intake can be reduced by simply reducing the quantity of salty foods in a diet, without recourse to salt substitutes. Salt substitutes have a taste similar to table salt and contain mostly potassium chloride, which will increase potassium intake. Excess potassium intake can cause hyperkalemia. Various diseases and medications may decrease the body's excretion of potassium, thereby increasing the risk of hyperkalemia. If you have kidney failure, heart failure or diabetes, seek medical advice before using a salt substitute. A manufacturer, LoSalt, has issued an advisory statement[47] that people taking the following prescription drugs should not use a salt substitute: Amiloride, Triamterene, Dytac, Spironolactone, Aldactone, Eplerenone and Inspra.
[edit] Production trends
Salt output in 2005
Salt is produced by evaporation of seawater or brine from other sources, such as brine wells and salt lakes, and by mining rock salt, called halite. In 2002, total world production was estimated at 210 million metric tonnes, the top five producers being the United States (40.3 million tonnes), China (32.9), Germany (17.7), India (14.5), and Canada (12.3).[48] Note that these figures are not just for table salt but for sodium chloride in general.
[edit] See also
A ship loading salt from a terminal.
Sodium chloride
Old Salt Route
Sea salt
Smoked salt
Kosher salt
History of salt
Fleur de sel
Curing (food preservation)
• Ten things you didn't know about Wikipedia •
From Wikipedia, the free encyclopedia
Jump to: navigation, search
For other uses, see Salt (disambiguation).
For the chemical properties of salt, see sodium chloride; for the chemical term, see salt (chemistry).
Salt is mostly sodium chloride (NaCl). This salt shaker also contains grains of rice, which some use to prevent caking
Brine being boiled down to pure salt in Zigong, China
Salt is a mineral essential for animal life, composed primarily of sodium chloride. Salt for human consumption is produced in different forms: unrefined salt (such as sea salt), refined salt (table salt), and iodized salt. It is a crystalline solid, white, pale pink or light grey in color, normally obtained from sea water or rock deposits. Edible rock salts may be slightly greyish in color due to this mineral content.
Sodium and chlorine, the two components of salt, are necessary for the survival of all living creatures, including humans, but they need not be consumed as salt, where they are found together in very concentrated form. Some isolated cultures, such as the Yanomami in South America, have been found to consume little salt.[1] Salt is involved in regulating the water content (fluid balance) of the body. Salt flavor is one of the basic tastes. Salt cravings may be caused by trace mineral deficiencies as well as by a deficiency of sodium chloride itself.
Overconsumption of salt can increase the risk of health problems, including high blood pressure. In food preparation, salt is used as a preservative and as a seasoning.
Contents[hide]
1 History
2 In religion
3 Forms of salt
3.1 Unrefined salt
3.2 Refined salt
3.3 Table salt
3.3.1 Iodized salt
3.3.2 Fluorinated salt
3.4 Salty condiments
4 Health effects
5 Recommended intake
6 Labeling
7 Campaigns
8 Salt substitutes
9 Production trends
10 See also
11 References
12 External links
12.1 Salt and health
13 Further reading//
[edit] History
See main article: History of salt
At the dawn of civilization, salt's preservative ability eliminated dependency on the seasonal availability of food, allowed travel over long distances, and was a vital food additive. However, because salt (NaCl) was difficult to obtain, it became a highly valued trade item throughout history. Until the 1900s, salt was one of the prime movers of national economies and wars. Salt was often taxed; research has discovered this practice to have existed as early as the 20th century BC in China. By the Middle Ages, caravans consisting of as many as forty thousand camels traversed four hundred miles of the Sahara bearing salt, sometimes trading it for slaves.[citation needed]
The first registers of salt use were produced around 4000 B.C. in Egypt, and later in Greece and Rome. Salt was very valuable and used to preserve and flavor foods. In Ancient Rome, salt was used as a currency. The Latin word salarium; meaning a payment made in salt, is the root of the word "salary." Unfortunately for those paid with salt, it was easily ruined by rain and other weather conditions. Payments to Roman workers and soldiers were made in salt.[2] Salt was also given to the parents of the groom in marriage until the 8th century. [attribution needed]
From the Phoenicians dates the evidence of harvesting solid salt from the sea. They also exported it to other civilizations. As a result of the increased salt supply from the sea, the value of salt depreciated. The harvest method used was flooding plains of land with seawater, then leaving the plains to dry. After the water dried, the salt which was left was collected and sold.
In the Mali Empire, merchants in 12th century Timbuktu—the gateway to the Sahara Desert and the seat of scholars—valued salt (NaCl) enough to buy it for its weight in gold; this trade led to the legends of the incredibly wealthy city of Timbuktu, and fueled inflation in Europe, which was importing the salt.[3]
During his protests in India, Mohandas Gandhi performed the famous salt march to challenge the British-imposed monopoly on salt.
[edit] In religion
Among the ancients, as with ourselves, "sol" (sun) and "sal" (salt) were considered essential to the maintenance of life.
There are thirty-five references (verses) to salt in the Bible (King James Version), the most familiar probably being the story of Lot's wife, who was turned into a pillar of salt when she disobeyed the angels and looked back at the wicked city of Sodom (Genesis 19:26). In the Sermon on the Mount, Jesus also referred to his followers as the "salt of the earth". The apostle Paul also encouraged Christians to "let your conversation be always full of grace, seasoned with salt" (Colossians 4:6) so that when others enquire about their beliefs, the Christian's answer generates a 'thirst' to know more about Christ. Salt is mandatory in the rite of the Tridentine Mass. Salt is used in the third item (which includes an Exorcism) of the Celtic Consecration (refer Gallican rite) that is employed in the Consecration of a Church. Salt may be added to the water "where it is customary" in the Roman Catholic rite of Holy water. The earliest Biblical mention of salt appears to be in reference to the destruction of Sodom and Gomorrah (Genesis 19:24-26)When King Abimelech destroyed the city of Shechem, held to have occurred in the thirteenth century BCE., he is said to have "sowed salt on it," this phrase expressing the completeness of its ruin. (Judges 9:45.)
In the native Japanese religion Shinto, salt is used for ritual purification of locations and people, such as in Sumo Wrestling.
In Aztec mythology, Huixtocihuatl was a fertility goddess who presided over salt and salt water.
[edit] Forms of salt
[edit] Unrefined salt
Main articles: Sea salt, Halite, and Fleur de sel
Different natural salts have different mineralities, giving each one a unique flavor. Fleur de sel, natural sea salt harvested by hand, has a unique flavor varying from region to region.
Some assert that unrefined sea salt is more healthy than refined salts.[4] However, completely raw sea salt is bitter due to magnesium and calcium compounds, and thus is rarely eaten. Other people think that raw sea and rock salts do not contain sufficient iodine salts to prevent iodine deficiency diseases like hypothyroidism.[5]
[edit] Refined salt
Salt evaporation pond in Alexandria, Egypt
Refined salt, which is most widely used presently, is mainly sodium chloride. Food grade salt accounts for only a small part of salt production in industrialised countries (3% in Europe[6]) although world-wide, food uses account for 17.5% of salt production[7]. The majority is sold for industrial use, from manufacturing pulp and paper to setting dyes in textiles and fabric, to producing soaps and detergents, and has great commercial value.
Salt mounds in Bolivia. Salt is harvested in the traditional method.
The manufacture and use of salt is one of the oldest chemical industries.[8] Salt is also obtained by evaporation of sea water, usually in shallow basins warmed by sunlight;[9] salt so obtained was formerly called bay salt, and is now often called sea salt or solar salt. Today, most refined salt is prepared from rock salt: mineral deposits high in salt.[citation needed] These rock salt deposits were formed by the evaporation of ancient salt lakes.[10] These deposits may be mined conventionally or through the injection of water. Injected water dissolves the salt, and the brine solution can be pumped to the surface where the salt is collected.
After the raw salt is obtained, it is refined to purify it and improve its storage and handling characteristics. Purification usually involves recrystallization. In recrystallization, a brine solution is treated with chemicals that precipitate most impurities (largely magnesium and calcium salts).[11] Multiple stages of evaporation are then used to collect pure sodium chloride crystals, which are kiln-dried.
Since the 1950's it has been common to add a trace of sodium hexacyanoferrate II to the brine, this acts as an anticaking agent by promoting irregular crystals.[12] Other anticaking agents (and potassium iodide, for iodised salt) are generally added after crystallization.[citation needed] These agents are hygroscopic chemicals which absorb humidity, keeping the salt crystals from sticking together. Some anticaking agents used are tricalcium phosphate, calcium or magnesium carbonates, fatty acid salts (acid salts), magnesium oxide, silicon dioxide, calcium silicate, sodium alumino-silicate, and alumino-calcium silicate. Concerns have been raised regarding the possible toxic effects of aluminium in the latter two compounds,[citation needed] however both the European Union and the United States Food and Drug Administration (FDA) permit their use.[13] The refined salt is then ready for packing and distribution.
[edit] Table salt
Single-serving salt packets
Table salt is refined salt, 99% sodium chloride.[14][15] It usually contains substances that make it free flowing (anticaking agents) such as sodium silicoaluminate or magnesium carbonate. It is common practice to put a few grains of uncooked rice in salt shakers to absorb extra moisture when anticaking agents are not enough.
[edit] Iodized salt
Main article: History of iodised salt
Iodized salt (BrE: iodised salt), table salt mixed with a minute amount of sodium iodide, iodate, or sometimes potassium iodide, is used to help reduce the chance of iodine deficiency in humans. Iodine deficiency commonly leads to thyroid gland problems, specifically endemic goiter. Endemic goiter is a disease characterized by a swelling of the thyroid gland, usually resulting in a bulbous protrusion on the neck. While only tiny quantities of iodine are required in a diet to prevent goiter, the United States Food and Drug Administration recommends (21 CFR 101.9 (c)(8)(iv)) 150 microgrammes of iodine per day for both men and women, and there are many places around the world where natural levels of iodine in the soil are low and the iodine is not taken up by vegetables.
Today, iodized salt is more common in the United States, Australia and New Zealand than in Britain. Table salt is also often iodized—a small amount of potassium iodide (in the US) or potassium iodate (in the EU) is added as an important dietary supplement. Table salt is mainly employed in cooking and as a table condiment. Iodized table salt has significantly reduced disorders of iodine deficiency in countries where it is used.[16] Iodine is important to prevent the insufficient production of thyroid hormones (hypothyroidism), which can cause goitre, cretinism in children, and myxedema in adults.
[edit] Fluorinated salt
In some European countries where fluoridation of drinking water is not practiced, fluorinated table salt is available. In France, 35% of sold table salt contains sodium or potassium fluoride.[17] Another additive, especially important for pregnant women is Folic acid (B vitamin) giving the table salt a yellow color.
[edit] Salty condiments
In many Asian cultures, table salt is not traditionally used as a condiment.[18]However, condiments such as soy sauce, fish sauce and oyster sauce tend to have a high salt content and fill much the same role as a salt-providing table condiment that table salt serves in western cultures.
[edit] Health effects
Sodium is one of the primary electrolytes in the body. All three electrolytes (sodium, potassium, and calcium) are available in unrefined salt, as are other vital minerals needed for optimal bodily function. Too much or too little salt in the diet can lead to muscle cramps, dizziness, or even an electrolyte disturbance, which can cause severe, even fatal, neurological problems.[19] Drinking too much water, with insufficient salt intake, puts a person at risk of water intoxication. Salt is even sometimes used as a health aid, such as in treatment of dysautonomia.[20]
People's risk for disease due to insufficient or excessive salt intake varies due to biochemical individuality. Some have asserted that while the risks of consuming too much salt are real, the risks have been exaggerated for most people, or that the studies done on the consumption of salt can be interpreted in many different ways.[21] [22]
Excess salt consumption has been linked to:
exercise-induced asthma.[23] On the other hand, another source counters, "…we still don't know whether salt contributes to asthma. If there is a link then it's very weak…".[24]
heartburn[25].
osteoporosis: One report shows that a high salt diet does reduce bone density in girls.[26]. Yet "While high salt intakes have been associated with detrimental effects on bone health, there are insufficient data to draw firm conclusions." ([27], p3)
Gastric cancer (Stomach cancer) is associated with high levels of sodium, "but the evidence does not generally relate to foods typically consumed in the UK." ([27], p18) However, in Japan, salt consumption is higher.[28]
hypertension (high blood pressure): "Since 1994, the evidence of an association between dietary salt intakes and blood pressure has increased. The data have been consistent in various study populations and across the age range in adults." ([27] p3). "The CMO [Chief Medical Officer] of England, in his Annual Report (DH, 2001), highlighted that people with high blood pressure are three times more likely to develop heart disease and stroke, and twice as likely to die from these diseases than those with normal levels."([27], p14). Professor Dr. Diederick Grobbee claims that there is no evidence of a causal link between salt intake and mortality or cardiovascular events.[29]. One study found that low urinary sodium is associated with greater risk of myocardial infarction among treated hypertensive men [30].
left ventricular hypertrophy (cardiac enlargement): "Evidence suggests that high salt intake causes left ventricular hypertrophy, a strong risk factor for cardiovascular disease, independently of blood pressure effects." ([27] p3) "…there is accumulating evidence that high salt intake predicts left ventricular hypertrophy." ([31], p12) Excessive salt (sodium) intake, combined with an inadequate intake of water, can cause hypernatremia. It can exacerbate renal disease.[19]
edema (BE: oedema): A decrease in salt intake has been suggested to treat edema (BE: oedema) (fluid retention).[32][19]
duodenal ulcers and gastric ulcers[33]
A large scale study by Nancy Cook et al shows that people with high-normal[1] blood pressure who significantly reduced the amount of salt in their diet decreased their chances of developing cardiovascular disease by 25% over the following 10 to 15 years. Their risk of dying from cardiovascular disease decreased by 20%.[34][35]
[edit] Recommended intake
Sea salt and peppercorns.
A salt mill for sea salt.
This section summarizes the salt intake recommended by the health agencies of various countries. Recommendations tend to be similar. Note that targets for the population as a whole tend to be pragmatic (what is achievable) while advice for an individual is ideal (what is best for health). For example, in the UK target for the population is "eat no more than 6 g a day" but for a person is 4 g.
Intakes can be expressed variously as salt or sodium and in various units.
1 g sodium = 1,000 mg sodium = 42 mmol sodium = 2.5 g salt
United Kingdom: In 2003, the UK's Scientific Advisory Committee on Nutrition (SACN) recommended that, for a typical adult, the Reference Nutrient Intake is 4 g salt per day (1.6 g or 70 mmol sodium). However, average adult intake is two and a half times the Reference Nutrient Intake for sodium. "Although accurate data are not available for children, conservative estimates indicate that, on a body weight basis, the average salt intake of children is higher than that of adults." SACN aimed for an achievable target reduction in average intake of salt to 6 g per day (2.4 g or 100 mmol sodium) — this is roughly equivalent to a teaspoonful of salt. The SACN recommendations for children are:
0–6 months old: less than 1 g/day
7–12 months: 1 g/day
1–3 years: 2 g/day
4–6 years: 3 g/day
7–10 years: 5 g/day
11–14 years: 6 g/day
SACN states, "The target salt intakes set for adults and children do not represent ideal or optimum consumption levels, but achievable population goals."[27]
Republic of Ireland: The Food Safety Authority of Ireland endorses the UK targets "emphasising that the RDA of 1.6 g sodium (4 g salt) per day should form the basis of advice targeted at individuals as distinct from the population health target of a mean salt intake of 6 g per day."([31], p16)
Canada: Health Canada recommends an Adequate Intake (AI) and an Upper Limit (UL) in terms of sodium.
0–6 months old: 0.12 g/day (AI)
7–12 months: 0.37 g/day (AI)
1–3 years: 1 g/day (AI) 1.5 g/day (UL)
4–8 years: 1.2/day (AI) 1.9 g/day (UL)
9–13 years: 1.5 g/day (AI) 2.2 g/day (UL)
14–50 years: 1.5 g/day (AI) 2.3 g/day (UL)
51–70 years: 1.3 g/day (AI) 2.3 g/day (UL)
70 years and older: 1.2 g/day (AI) 2.3 g/day (UL)[36]
New Zealand
Adequate Intake (AI) 0.46 – 0.92 g sodium = 1.2 – 2.3g salt
Upper Limit (UL)) 2.3 g sodium = 5.8 g salt[37]
Australia: The recommended dietary intake (RDI) is 0.92 g–2.3 g sodium per day (= 2.3 g–5.8 g salt)[38]
USA: The Food and Drug Administration itself does not make a recommendation[39] but refers readers to Dietary Guidelines for Americans 2005. These suggest that US citizens should consume less than 2,300 mg of sodium (= 2.3 g sodium = 5.8 g salt) per day. [40]
[edit] Labeling
UK: The Food Standards Agency defines the level of salt in foods as follows: "High is more than 1.5g salt per 100g (or 0.6g sodium). Low is 0.3g salt or less per 100g (or 0.1g sodium). If the amount of salt per 100g is in between these figures, then that is a medium level of salt." In the UK, foods produced by some supermarkets and manufacturers have ‘traffic light’ colors on the front of the pack: Red (High), Amber (Medium), or Green (Low).[41]
USA: The FDA Food Labeling Guide stipulates whether a food can be labelled as "free", "low", or "reduced/less" in respect of sodium. When other health claims are made about a food (e.g. low in fat, calories, etc.), a disclosure statement is required if the food exceeds 480mg of sodium per 'serving.'[42]
[edit] Campaigns
In 2004, Britain's Food Standards Agency started a public health campaign called "Salt - Watch it", which recommends no more than 6g of salt per day; it features a character called Sid the Slug and was criticised by the Salt Manufacturers Association (SMA).[43] The Advertising Standards Authority did not uphold the SMA complaint in its adjudication.[44]. In March 2007, the FSA launched the third phase of their campaign with the slogan "Salt. Is your food full of it?" fronted by comedienne Jenny Eclair.[45]
The Menzies Research Institute in Tasmania, Australia, maintains a website [46] dedicated to educating people about the potential problems of a salt-laden diet.
[edit] Salt substitutes
Main article: Salt substitute
Salt intake can be reduced by simply reducing the quantity of salty foods in a diet, without recourse to salt substitutes. Salt substitutes have a taste similar to table salt and contain mostly potassium chloride, which will increase potassium intake. Excess potassium intake can cause hyperkalemia. Various diseases and medications may decrease the body's excretion of potassium, thereby increasing the risk of hyperkalemia. If you have kidney failure, heart failure or diabetes, seek medical advice before using a salt substitute. A manufacturer, LoSalt, has issued an advisory statement[47] that people taking the following prescription drugs should not use a salt substitute: Amiloride, Triamterene, Dytac, Spironolactone, Aldactone, Eplerenone and Inspra.
[edit] Production trends
Salt output in 2005
Salt is produced by evaporation of seawater or brine from other sources, such as brine wells and salt lakes, and by mining rock salt, called halite. In 2002, total world production was estimated at 210 million metric tonnes, the top five producers being the United States (40.3 million tonnes), China (32.9), Germany (17.7), India (14.5), and Canada (12.3).[48] Note that these figures are not just for table salt but for sodium chloride in general.
[edit] See also
A ship loading salt from a terminal.
Sodium chloride
Old Salt Route
Sea salt
Smoked salt
Kosher salt
History of salt
Fleur de sel
Curing (food preservation)
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