Reflection (physics)
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For other uses, see reflection.
The reflection of Mount Hood in Trillium Lake.
The reflection of Mount Hood in Trillium Lake.
Reflection of a hot air balloon, partially obscured by a jetty.
Reflection of a hot air balloon, partially obscured by a jetty.
Reflection is the change in direction of a wave front at an interface between two different media so that the wave front returns into the medium from which it originated. Common examples include the reflection of light, sound and water waves.
Law of reflection: Angle of incidence = Angle of reflection
Contents
[hide]
* 1 Reflections
* 2 Reflection of light
* 3 Laws of regular reflection
* 4 Other types of reflection
o 4.1 Diffuse reflection
o 4.2 Retroreflection
o 4.3 Complex conjugate reflection
o 4.4 Neutron reflection
o 4.5 Sound reflection
o 4.6 Seismic reflection
* 5 Quantum interpretation
* 6 See also
* 7 External links
[edit] Reflections
Reflections may occur in a number of wave and particle phenomena; these include acoustic (as in common echos and underwater sonar), seismic waves in geologic structures, surface waves in bodies of water, and various electromagnetic waves, most usefully from VHF and higher radar frequencies, progressing upward through centimeter to millimeter-wavelength radar and the various light frequencies and (with special "grazing" mirrors, to X-ray frequencies and beyond to gamma rays.
[edit] Reflection of light
Crepuscular rays and their reflection
Crepuscular rays and their reflection
Diagram of specular reflection
Diagram of specular reflection
Reflection of light may be specular (that is, mirror-like) or diffuse (that is, not retaining the image, only the energy) depending on the nature of the interface. Furthermore, if the interface is between dielectric-conductor or dielectric-dielectric media, the phase of the reflected wave may or may not be inverted, respectively.
Main article: Specular reflection
A mirror provides the most common model for specular light reflection and consists of a glass sheet in front of a metallic coating where the reflection actually occurs. Reflection is enhanced in metals by suppression of wave propagation beyond their skin depths. It is also possible for reflection to occur from the surface of transparent media, such as water or glass. In the diagram, a light ray PO strikes a vertical mirror at point O, and the reflected ray is OQ. By projecting an imaginary line through point O perpendicular to the mirror, known as the normal, we can measure the angle of incidence, θi and the angle of reflection, θr. The law of reflection states that θi = θr, or in other words, the angle of incidence equals the angle of reflection.
An indian triggerfish reflecting in the water surface. In fact, this is caused by total internal reflection.
An indian triggerfish reflecting in the water surface. In fact, this is caused by total internal reflection.
In fact, reflection of light may occur whenever light travels from a medium of a given refractive index into a medium with a different refractive index. In the most general case, a certain fraction of the light is reflected from the interface, and the remainder is refracted. Solving Maxwell's equations for a light ray striking a boundary allows the derivation of the Fresnel equations, which can be used to predict how much of the light reflected, how much is refracted in a given situation. Total internal reflection of light from a denser medium occurs if the angle of incidence is above the critical angle. Total internal reflection is used as a means of focussing waves that cannot effectively be reflected by common means. X-ray telescopes are constructed by creating a converging "tunnel" for the waves. As the waves interact at low angle with the surface of this tunnel they are reflected toward the focus point (or toward another interaction with the tunnel surface, eventually being directed to the a detector at the focus). A conventional reflector would be useless as the X-rays would simply pass through the intended reflector. When light reflects off a material denser (with higher refractive index) than the external medium, it undergoes a 180° phase reversal. In contrast, a less dense, lower refractive index material will reflect light in phase. This is an important principle in the field of thin-film optics. Specular reflection at a curved surface forms an image which may be magnified or demagnified; curved mirrors have optical power. Such mirrors may have surfaces that are spherical or parabolic.
[edit] Laws of regular reflection
Specular reflection at a curved surface of sea foam, which is made out of plankton
Specular reflection at a curved surface of sea foam, which is made out of plankton
If the reflecting surface is very smooth, the reflection of light that occurs is called specular or regular reflection. The laws of reflection are as follows:
1. The incident ray, the reflected ray and the normal to the reflection surface at the point of the incidence lie in the same plane.
2. The angle which the incident ray makes with the normal is equal to the angle which the reflected ray makes to the same normal.
[edit] Other types of reflection
[edit] Diffuse reflection
Diffuse reflection
Diffuse reflection
Main article: Diffuse reflection
When light strikes a rough or granular surface, it bounces off in all directions due to the microscopic irregularities of the interface. Thus, an 'image' is not formed. This is called diffuse reflection. The exact form of the reflection depends on the structure of the surface. One common model for diffuse reflection is Lambertian reflectance, in which the light is reflected with equal luminance (in photometry) or radiance (in radiometry) in all directions, as defined by Lambert's cosine law.
[edit] Retroreflection
Working principle of a corner reflector
Working principle of a corner reflector
Main article: Retroreflector
Some surfaces exhibit retroreflection. The structure of these surfaces is such that light is returned in the direction from which it came. A simple retroreflector can be made by placing three ordinary mirrors mutually perpendicular to one another (a corner reflector). The image produced is the inverse of one produced by a single mirror. A surface can be made partially retroreflective by depositing a layer of tiny refractive spheres on it or by creating small pyramid like structures . In both cases internal reflection causes the light to be reflected back to where it originated. This is used to make traffic signs and automobile license plates reflect light mostly back in the direction from which it came. In this application perfect retroreflection is not desired, since the light would then be directed back into the headlights of an oncoming car rather than to the driver's eyes.
[edit] Complex conjugate reflection
Light bounces exactly back in the direction from which it came due to a nonlinear optical process. In this type of reflection, not only the direction of the light is reversed, but the actual wavefronts are reversed as well. A conjugate reflector can be used to remove aberrations from a beam by reflecting it and then passing the reflection through the aberrating optics a second time.
[edit] Neutron reflection
Materials that reflect neutrons, for example beryllium, are used in nuclear reactors and nuclear weapons. In the physical and biological sciences, the reflection of neutrons off atoms within a material is commonly used to determine its internal structures. [1]
[edit] Sound reflection
When a longitudinal sound wave strikes a flat surface, sound is reflected in a coherent manner provided that the dimension of the reflective surface is large compared to the wavelength of the sound. Note that audible sound has a very wide frequency range (from 20 to about 17000 Hz), and thus a very wide range of wavelengths (from about 20 mm to 17 m). As a result, the overall nature of the reflection varies according to the texture and structure of the surface. For example, porous materials will absorb some energy, and rough materials (where rough is relative to the wavelength) tend to reflect in many directions — to scatter the energy, rather than to reflect it coherently. This leads into the field of architectural acoustics, because the nature of these reflections is critical to the auditory feel of a space. In the theory of exterior noise mitigation, reflective surface size mildly detracts from the concept of a noise barrier by reflecting some of the sound into the opposite direction.
[edit] Seismic reflection
Seismic waves produced by earthquakes or other sources (such as explosions) may be reflected by layers within the Earth. Study of the deep reflections of waves generated by earthquakes has allowed seismologists to determine the layered structure of the Earth. Shallower reflections are used in reflection seismology to study the Earth's crust generally, and in particular to prospect for petroleum and natural gas deposits.
[edit] Quantum interpretation
Light waves incident on a material induce small oscillations of polarisation in the individual atoms, causing each atom to radiate a weak secondary wave (in all directions like a dipole antenna). All of these waves add up to specular reflection (following Hero's equi-angular reflection law) and refraction. Light–matter interaction in terms of photons is a topic of quantum electrodynamics, and is described in detail by Richard Feynman in his popular book QED: The Strange Theory of Light and Matter.
Tuesday, August 5, 2008
CARBON AND ITS COMPOUNDS
Carbon
From Wikipedia, the free encyclopedia
Jump to: navigation, search
For other uses, see Carbon (disambiguation).
6 boron ← carbon → nitrogen
-
↑
C
↓
Si
Periodic table - Extended periodic table
General
Name, symbol, number carbon, C, 6
Chemical series nonmetals
Group, period, block 14, 2, p
Appearance black (graphite)
colorless (diamond)
Standard atomic weight 12.0107(8) g·mol−1
Electron configuration 1s2 2s2 2p2
Electrons per shell 2, 4
Physical properties
Phase solid
Density (near r.t.) (graphite) 1.9-2.3[1] g·cm−3
Density (near r.t.) (diamond) 3.5-3.53[1] g·cm−3
Density (near r.t.) (fullerene) 1.69[1] g·cm−3
Heat of fusion (graphite) ? 100 kJ·mol−1
Heat of fusion (diamond) ? 120 kJ·mol−1
Heat of vaporization 715 kJ·mol−1
Specific heat capacity (25 °C) (graphite)
8.517 J·mol−1·K−1
Specific heat capacity (25 °C) (diamond)
6.115 J·mol−1·K−1
Vapor pressure (graphite) P/Pa 1 10 100 1 k 10 k 100 k
at T/K 2839 3048 3289 3572 3908
Atomic properties
Crystal structure (graphite) hexagonal
Oxidation states 4, 3 [2], 2, 1 [3], 0, -1, -2, -3, -4[4]
(mildly acidic oxide)
Electronegativity 2.55 (Pauling scale)
Ionization energies
(more) 1st: 1086.5 kJ·mol−1
2nd: 2352.6 kJ·mol−1
3rd: 4620.5 kJ·mol−1
Atomic radius 70 pm
Atomic radius (calc.) 67 pm
Covalent radius 77 pm
Van der Waals radius 170 pm
Miscellaneous
Magnetic ordering diamagnetic
Electrical resistivity (graphite) 1.375*10-5 [5]Ω·m
Thermal conductivity (300 K) (graphite)
(80–230) W·m−1·K−1
Thermal conductivity (300 K) (diamond)
(900–2320) W·m−1·K−1
Thermal diffusivity (300 K) (diamond)
(503–1300) mm²/s
Mohs hardness (graphite) 1-2 [6]
Mohs hardness (diamond) 10.0 [6]
CAS registry number 7440-44-0
Selected isotopes
Main article: Isotopes of carbon iso NA half-life DM DE (MeV) DP
15
12C 98.9% 12C is stable with 6 neutrons
13C 1.1% 13C is stable with 7 neutrons
14C trace 5730 y beta- 0.156 14N
References
This box: view • talk • edit
Carbon (pronounced /kɑɹbən/) is a chemical element with the symbol C and atomic number is 6. As a member of group 14 on the periodic table, it is nonmetallic and tetravalent—making four electrons available to form covalent chemical bonds. There are three naturally occurring isotopes, with 12C and 13C being stable, while 14C is radioactive, decaying with a half-life of about 5700 years.[7] Carbon is one of the few elements known to man since antiquity.[8][9] The name "carbon" comes from Latin language carbo, coal, and, in some Romance languages, the word carbon can refer both to the element and to coal.
There are several allotropes of carbon of which the best known are graphite, diamond, and amorphous carbon.[10] The physical properties of carbon vary widely with the allotropic form. For example, diamond is highly transparent, while graphite is opaque and black. Diamond is among the hardest materials known, while graphite is soft enough to form a streak on paper. Diamond has a very low electric conductivity, while graphite is a very good conductor. Also, diamond has the highest thermal conductivity of all known materials under normal conditions. All the allotropic forms are solids under normal conditions but graphite is the most thermodynamically stable.
All forms of carbon are highly stable, requiring high temperature to react even with oxygen. The most common oxidation state of carbon in inorganic compounds is +4, while +2 is found in carbon monoxide and other transition metal carbonyl complexes. The largest sources of inorganic carbon are limestones, dolomites and carbon dioxide, but significant quantities occur in organic deposits of coal, peat, oil and methane clathrates. Carbon forms more compounds than any other element, with almost ten million pure organic compounds described to date, which in turn are a tiny fraction of such compounds that are theoretically possible under standard conditions.[11]
Carbon is the fourth most abundant element in the universe by mass after hydrogen, helium, and oxygen. It is present in all known lifeforms, and in the human body, carbon is the second most abundant element by mass (about 18.5%) after oxygen.[12] This abundance, together with the unique diversity of organic compounds and their unusual polymer-forming ability at the temperatures commonly encountered on Earth, make this element the chemical basis of all known life.
Contents
[hide]
* 1 Characteristics
o 1.1 Allotropes
o 1.2 Occurrence
o 1.3 Isotopes
o 1.4 Formation in stars
o 1.5 Carbon cycle
* 2 Compounds
o 2.1 Inorganic compounds
o 2.2 Organic compounds
* 3 History and etymology
o 3.1 Applications
* 4 Production
o 4.1 Graphite Production
* 5 Precautions
* 6 See also
* 7 References
* 8 External links
[edit] Characteristics
The different forms or allotropes of carbon (see below) include the hardest naturally occurring substance, diamond, and also one of the softest known substances, graphite. Moreover, it has an affinity for bonding with other small atoms, including other carbon atoms, and is capable of forming multiple stable covalent bonds with such atoms. As a result, carbon is known to form nearly ten million different compounds; the large majority of all chemical compounds.[11] Carbon also has the highest melting and sublimation point of all elements.[citation needed] At atmospheric pressure it has no actual melting point as its triple point is at 10 MPa (100 bar) so it sublimates above 4000 K.[citation needed] Carbon sublimes in a carbon arc which has a temperature of about 5800K. Thus, irrespective of its allotropic form, carbon remains solid at higher temperatures than the highest melting point metals such as tungsten or rhenium. Although thermodynamically prone to oxidation, carbon resists oxidation more effectively than elements such as iron and copper that are weaker reducing agents at room temperature.
Diamond and graphite are two allotropes of carbon: pure forms of the same element that differ in structure.
Diamond and graphite are two allotropes of carbon: pure forms of the same element that differ in structure.
Carbon compounds form the basis of all life on Earth and the carbon-nitrogen cycle provides some of the energy produced by the Sun and other stars. Although it forms an extraordinary variety of compounds, most forms of carbon are comparatively unreactive under normal conditions. At standard temperature and pressure, it resists all but the strongest oxidizers. It does not react with sulfuric acid, hydrochloric acid, chlorine or any alkalis. At elevated temperatures carbon reacts with oxygen to form carbon oxides, and will reduce such metal oxides as iron oxide to the metal. This exothermic reaction is used in the iron and steel industry to control the carbon content of steel:
Fe3O4 + 4C(s) → 3Fe(s) + 4CO(g)
with sulfur to form carbon disulfide and with steam in the coal-gas reaction
C(s) + H2O(g) → CO(g) + H2(g).
Carbon combines with some metals at high temperatures to form metallic carbides, such as the iron carbide cementite in steel, and tungsten carbide, widely used as an abrasive and for making hard tips for cutting tools.
The system of carbon allotropes spans a range of extremes:
Synthetic diamond nanorods are the hardest materials known. Graphite is one of the softest materials known.
Diamond is the ultimate abrasive. Graphite is a very good lubricant.
Diamond is an excellent electrical insulator. Graphite is a conductor of electricity.
Diamond is the best known thermal conductor Some forms of graphite are used for thermal insulation (i.e. firebreaks and heatshields)
Diamond is highly transparent. Graphite is opaque.
Diamond crystallizes in the cubic system. Graphite crystallizes in the hexagonal system.
Amorphous carbon is completely isotropic. Carbon nanotubes are among the most anisotropic materials ever produced.
[edit] Allotropes
Main article: Allotropes of carbon
Atomic carbon is a very short-lived species and therefore, carbon is stabilized in various multi-atomic structures with different molecular configurations called allotropes. The three relatively well-known allotropes of carbon are amorphous carbon, graphite, and diamond. Once considered exotic, fullerenes are nowadays commonly synthesized and used in research; they include buckyballs,[13][14] carbon nanotubes,[15] carbon nanobuds[16] and nanofibers[17].[18] Several other exotic allotropes have also been discovered, such as aggregated diamond nanorods,[19] lonsdaleite,[20] glassy carbon,[21] carbon nanofoam[22] and linear acetylenic carbon.[23]
* The amorphous form, is an assortment of carbon atoms in a non-crystalline, irregular, glassy state, which is essentially graphite but not held in a crystalline macrostructure. It is present as a powder, and is the main constituent of substances such as charcoal, lampblack (soot) and activated carbon.
* At normal pressures carbon takes the form of graphite, in which each atom is bonded trigonally to three others in a plane composed of fused hexagonal rings, just like those in aromatic hydrocarbons. The resulting network is 2-dimensional, and the resulting flat sheets are stacked and loosely bonded through weak Van der Waals forces. This gives graphite its softness and its cleaving properties (the sheets slip easily past one another). Because of the delocalization of one of the outer electrons of each atom to form a π-cloud, graphite conducts electricity, but only in the plane of each covalently bonded sheet. This results in a lower bulk electrical conductivity for carbon than for most metals. The delocalization also accounts for the energetic stability of graphite over diamond at room temperature.
Some allotropes of carbon: a) diamond; b) graphite; c) lonsdaleite; d-f) fullerenes (C60, C540, C70); g) amorphous carbon; h) carbon nanotube.
Some allotropes of carbon: a) diamond; b) graphite; c) lonsdaleite; d-f) fullerenes (C60, C540, C70); g) amorphous carbon; h) carbon nanotube.
* At very high pressures carbon forms the more compact allotrope diamond, having nearly twice the density of graphite. Here, each atom is bonded tetrahedrally to four others, thus making a 3-dimensional network of puckered six-membered rings of atoms. Diamond has the same cubic structure as silicon and germanium and, thanks to the strength of the carbon-carbon bonds is the hardest naturally occurring substance in terms of resistance to scratching. Contrary to the popular belief that "diamonds are forever", they are in fact thermodynamically unstable under normal conditions and transform into graphite.[10] But due to a high activation energy barrier, the transition into graphite is so extremely slow at room temperature as to be unnoticeable.
* Under some conditions, carbon crystallizes as lonsdaleite. This form is similar to diamond but has a hexagonal crystal lattice.[20]
* Fullerenes have a graphite-like structure, but instead of purely hexagonal packing, they also contain pentagons (or even heptagons) of carbon atoms, which bend the sheet into spheres, ellipses or cylinders. The properties of fullerenes (split into buckyballs, buckytubes and nanobuds) have not yet been fully analyzed and represents an intense area of research in nanomaterials. The name "fullerene" is given after Richard Buckminster Fuller, developer of some geodesic domes,[citation needed] which resemble the structure of fullerenes. The buckyballs are fairly large molecules formed completely of carbon bonded trigonally, forming spheroids (the best-known and simplest is the soccerball-shaped structure C60 buckminsterfullerene).[13] Carbon nanotubes are structurally similar to buckyballs, except that each atom is bonded trigonally in a curved sheet that forms a hollow cylinder.[14][15] Nanobuds were first published in 2007 and are hybrid bucky tube/buckyball materials (buckyballs are covalently bonded to the outer wall of a nanotube) that combine the properties of both in a single structure.[16]
* Of the other discovered allotropes, aggregated diamond nanorods were synthesised in 2005 and are believed to be the hardest substance known yet.[24] Carbon nanofoam is a ferromagnetic allotrope discovered in 1997. It consists of a low-density cluster-assembly of carbon atoms strung together in a loose three-dimensional web, in which the atoms are bonded trigonally in six- and seven-membered rings. It is among the lightest known solids, with a density of about 2 kg/m³.[25] Similarly, glassy carbon contains a high proportion of closed porosity.[21] But unlike normal graphite, the graphitic layers are not stacked like pages in a book, but have a more random arrangement. Linear acetylenic carbon[23] has the chemical structure[26] -(C:::C)n- .Carbon in this modification is linear with sp orbital hybridisation, and is a polymer with alternating single and triple bonds. This type of carbyne is of considerable interest to nanotechnology as its Young's modulus is forty times that of the hardest known material - diamond.[27]
[edit] Occurrence
Graphite ore
Graphite ore
Raw diamond crystal.
Raw diamond crystal.
Carbon is the fourth most abundant chemical element in the universe by mass after hydrogen, helium, and oxygen. Carbon is abundant in the Sun, stars, comets, and in the atmospheres of most planets. Some meteorites contain microscopic diamonds that were formed when the solar system was still a protoplanetary disk. Microscopic diamonds may also be formed by the intense pressure and high temperature at the sites of meteorite impacts.[28]
"Present day" (1990s) sea surface dissolved inorganic carbon concentration (from the GLODAP climatology)
"Present day" (1990s) sea surface dissolved inorganic carbon concentration (from the GLODAP climatology)
In combination with oxygen in carbon dioxide, carbon is found in the Earth's atmosphere (in quantities of approximately 810 gigatonnes) and dissolved in all water bodies (approximately 36000 gigatonnes). Around 1900 gigatonnes are present in the biosphere. Hydrocarbons (such as coal, petroleum, and natural gas) contain carbon as well — coal "reserves" (not "resources") amount to around 900 gigatonnes, and oil reserves around 150 gigatonnes. With smaller amounts of calcium, magnesium, and iron, carbon is a major component of very large masses carbonate rock (limestone, dolomite, marble etc.).
Coal is a significant commercial source of mineral carbon; anthracite containing 92-98% carbon[citation needed] and the largest source (4000 Gt, or 80% of coal, gas and oil reserves) of carbon in a form suitable for use as fuel.[29]
Graphite is found in large quantities in New York and Texas, the United States, Russia, Mexico, Greenland, and India.
Natural diamonds occur in the rock kimberlite, found in ancient volcanic "necks," or "pipes". Most diamond deposits are in Africa, notably in South Africa, Namibia, Botswana, the Republic of the Congo, and Sierra Leone. There are also deposits in Arkansas, Canada, the Russian Arctic, Brazil and in Northern and Western Australia.
Diamonds are now also being recovered from the ocean floor off the Cape of Good Hope. However, though diamonds are found naturally, about 30% of all industrial diamonds used in the U.S. are now made synthetically.
According to studies from the Massachusetts Institute of Technology, an estimate of the global carbon budget is:[citation needed]
Biosphere, oceans, atmosphere
0.45 x 1018 kilograms (3.7 x 1018 moles)
Crust
Organic carbon 13.2 x 1018 kg
Carbonates 62.4 x 1018 kg
Mantle
1200 x 1018 kg
Carbon-14 is formed in upper layers of the troposphere and the stratosphere, at altitudes of 9–15 km, by a reaction that is precipitated by cosmic rays. Thermal neutrons are produced that collide with the nuclei of nitrogen-14, forming carbon-14 and a proton.
[edit] Isotopes
Main article: Isotopes of carbon
Isotopes of carbon are atomic nuclei that contain six protons plus a number of neutrons (varying from 2 to 16). Carbon has two stable, naturally occurring isotopes.[7] The isotope carbon-12 (12C) forms 98.93% of the carbon on Earth, while carbon-13 (13C) forms the remaining 1.07%.[7] The concentration of 12C is further increased in biological materials because biochemical reactions discriminate against 13C.[30] In 1961 the International Union of Pure and Applied Chemistry (IUPAC) adopted the isotope carbon-12 as the basis for atomic weights.[31] Identification of carbon in NMR experiments is done with the isotope 13C.
Carbon-14 (14C) is a naturally occurring radioisotope which occurs in trace amounts on Earth of up to 1 part per trillion (0.0000000001%), mostly confined to the atmosphere and superficial deposits, particularly of peat and other organic materials.[32] This isotope decays by 0.158 MeV β- emission. Because of its relatively short half-life of 5730 years, 14C is virtually absent in ancient rocks, but is created in the upper atmosphere (lower stratosphere and upper troposphere) by interaction of nitrogen with cosmic rays.[33] The abundance of 14C in the atmosphere and in living organisms is almost constant, but decreases predictably in their bodies after death. This principle is used in radiocarbon dating, invented in 1949, which has been used extensively to determine the age of carbonaceous materials with ages up to about 40,000 years.[34][35]
There are 15 known isotopes of carbon and the shortest-lived of these is 8C which decays through proton emission and alpha decay and has a half-life of 1.98739x10-21 s.[36] The exotic 19C exhibits a nuclear halo, which means its radius is appreciably larger than would be expected if the nucleus was a sphere of constant density.[37]
[edit] Formation in stars
Main articles: Triple-alpha process and CNO cycle
Formation of the carbon atomic nucleus requires a nearly simultaneous triple collision of alpha particles (helium nuclei) within the core of a giant or supergiant star. This happens in conditions of temperature and helium concentration that the rapid expansion and cooling of the early universe prohibited, and therefore no significant carbon was created during the Big Bang. Instead, the interiors of stars in the horizontal branch transform three helium nuclei into carbon by means of this triple-alpha process. In order to be available for formation of life as we know it, this carbon must then later be scattered into space as dust, in supernova explosions, as part of the material which later forms second- and third-generation star systems which have planets accreted from such dust. The Solar System is one such third-generation star system.
One of the fusion mechanisms powering stars is the carbon-nitrogen cycle.
Rotational transitions of various isotopic forms of carbon monoxide (e.g. 12CO, 13CO, and C18O) are detectable in the submillimeter regime, and are used in the study of newly forming stars in molecular clouds.
[edit] Carbon cycle
Main article: Carbon cycle
Diagram of the carbon cycle. The black numbers indicate how much carbon is stored in various reservoirs, in billions of tons ("GtC" stands for gigatons of carbon; figures are circa 2004). The purple numbers indicate how much carbon moves between reservoirs each year. The sediments, as defined in this diagram, do not include the ~70 million GtC of carbonate rock and kerogen.
Diagram of the carbon cycle. The black numbers indicate how much carbon is stored in various reservoirs, in billions of tons ("GtC" stands for gigatons of carbon; figures are circa 2004). The purple numbers indicate how much carbon moves between reservoirs each year. The sediments, as defined in this diagram, do not include the ~70 million GtC of carbonate rock and kerogen.
Under terrestrial conditions, conversion of one element to another is very rare. Therefore, the amount of carbon on Earth is effectively constant. Thus, processes that use carbon must obtain it somewhere and dispose of it somewhere else. The paths that carbon follows in the environment make up the carbon cycle. For example, plants draw carbon dioxide out of their environment and use it to build biomass, as in carbon respiration or the Calvin cycle, a process of carbon fixation. Some of this biomass is eaten by animals, whereas some carbon is exhaled by animals as carbon dioxide. The carbon cycle is considerably more complicated than this short loop; for example, some carbon dioxide is dissolved in the oceans; dead plant or animal matter may become petroleum or coal, which can burn with the release of carbon, should bacteria not consume it.
[edit] Compounds
[edit] Inorganic compounds
Main article: Compounds of carbon
Commonly carbon-containing compounds which are associated with minerals or which do not contain hydrogen or fluorine, are treated separately from classical organic compounds; however the definition is not rigid (see reference articles above). Among these are the simple oxides of carbon. The most prominent oxide is carbon dioxide (CO2). This was once the principal constituent of the paleoatmosphere, but is a minor component of the Earth's atmosphere today.[38] Dissolved in water, it forms carbonic acid (H2CO3), but as most compounds with multiple single-bonded oxygens on a single carbon it is unstable.[citation needed] Through this intermediate, though, resonance-stabilized carbonate ions are produced. Some important minerals are carbonates, notably calcite. Carbon disulfide (CS2) is similar.
The other common oxide is carbon monoxide (CO). It is formed by incomplete combustion, and is a colorless, odorless gas. The molecules each contain a triple bond and are fairly polar, resulting in a tendency to bind permanently to hemoglobin molecules, displacing oxygen, which has a lower binding affinity.[39][40] Cyanide (CN–), has a similar structure, but behaves much like a halide ion (pseudohalogen). For example it can form the nitride cyanogen molecule ((CN)2), similar to diatomic halides. Other uncommon oxides are carbon suboxide (C3O2),[41] the unstable dicarbon monoxide (C2O),[42][43] and even carbon trioxide (CO3).[44][45]
With reactive metals, such as tungsten, carbon forms either carbides (C4–), or acetylides (C22–) to form alloys with high melting points. These anions are also associated with methane and acetylene, both very weak acids. With an electronegativity of 2.5,[46] carbon prefers to form covalent bonds. A few carbides are covalent lattices, like carborundum (SiC), which resembles diamond.
[edit] Organic compounds
Main article: Organic compound
Structural formula of methane, the simplest possible organic compound
Structural formula of methane, the simplest possible organic compound
Carbon has the ability to form very long chains interconnecting C-C bonds. This property is called catenation. Carbon-carbon bonds are strong, and stable.[citation needed] This property allows carbon to form an almost infinite number of compounds; in fact, there are more known carbon-containing compounds than all the compounds of the other chemical elements combined except those of hydrogen (because almost all organic compounds contain hydrogen too).
The simplest form of an organic molecule is the hydrocarbon—a large family of organic molecules that are composed of hydrogen atoms bonded to a chain of carbon atoms. Chain length, side chains and functional groups all affect the properties of organic molecules. By IUPAC's definition, all the other organic compounds are functionalized compounds of hydrocarbons.[citation needed]
Carbon is the basis for all plastic materials that are used in common household items.
Carbon is the basis for all plastic materials that are used in common household items.
Carbon occurs in all organic life and is the basis of organic chemistry. When united with hydrogen, it forms various flammable compounds called hydrocarbons which are important to industry as chemical feedstock for the manufacture of plastics, petrochemicals and as fossil fuels.
When combined with oxygen and hydrogen, carbon can form many groups of important biological compounds including sugars, celluloses, lignans, chitins, alcohols, fats, and aromatic esters, carotenoids and terpenes. With nitrogen it forms alkaloids, and with the addition of sulfur also it forms antibiotics, amino acids and proteins. With the addition of phosphorus to these other elements, it forms DNA and RNA, the chemical codes of life, and adenosine triphosphate (ATP), the most important energy-transfer molecules in all living cells.
[edit] History and etymology
Please help improve this section by expanding it. Further information might be found on the talk page or at requests for expansion. (January 2008)
The English name carbon comes from the Latin carbo for coal and charcoal,[47] and hence comes French charbon, meaning charcoal. In German, Dutch and Danish, the names for carbon are Kohlenstoff, koolstof and kulstof respectively, all literally meaning coal-substance.
Carl Wilhelm Scheele
Carl Wilhelm Scheele
Antoine Lavoisier in his youth
Antoine Lavoisier in his youth
Carbon was discovered in prehistory and was known in the forms of soot and charcoal to the earliest human civilizations. Diamonds were known probably as early as 2500 BCE in China, while carbon in the forms of charcoal was made around Roman times by the same chemistry as it is today, by heating wood in a pyramid covered with clay to exclude air.[48][49]
In 1722, René A. F. de Réaumur demonstrated that iron was transformed into steel through the absorption of some substance, now known to be carbon.[50] In 1772, Antoine Lavoisier showed that diamonds are a form of carbon, when he burned samples of carbon and diamond then showed that neither produced any water and that both released the same amount of carbon dioxide per gram. Carl Wilhelm Scheele showed that graphite, which had been thought of as a form of lead, was instead a type of carbon.[51] In 1786, the French scientists Claude Louis Berthollet, Gaspard Monge and C. A. Vandermonde then showed that this substance was carbon.[52] In their publication they proposed the name carbone (Latin carbonum) for this element. Antoine Lavoisier listed carbon as an element in his 1789 textbook.[53]
A new allotrope of carbon, fullerene, that was discovered in 1985[54] includes nanostructured forms such as buckyballs and nanotubes.[13] Their discoverers received the Nobel Prize in Chemistry in 1996.[55] The resulting renewed interest in new forms, lead to the discovery of further exotic allotropes, including glassy carbon, and the realization that "amorphous carbon" is not strictly amorphous.[21]
[edit] Applications
Pencil lead for mechanical pencils are made of graphite.
Pencil lead for mechanical pencils are made of graphite.
Sticks of vine and compressed charcoal.
Sticks of vine and compressed charcoal.
A cloth of woven carbon filaments
A cloth of woven carbon filaments
Silicon carbide single crystal
Silicon carbide single crystal
The C60 fullerene in crystalline form
The C60 fullerene in crystalline form
Tungsten carbide milling bits
Tungsten carbide milling bits
Carbon is essential to all known living systems, and without it life as we know it could not exist (see alternative biochemistry). The major economic use of carbon other than food and wood is in the form of hydrocarbons, most notably the fossil fuel methane gas and crude oil (petroleum). Crude oil is used by the petrochemical industry to produce, amongst others, gasoline and kerosene, through a distillation process, in refineries. Cellulose is a natural, carbon-containing polymer produced by plants in the form of cellulose, cotton, linen, hemp. Commercially valuable carbon polymers of animal origin include wool, cashmere and silk. Plastics are made from synthetic carbon polymers, often with oxygen and nitrogen atoms included at regular intervals in the main polymer chain. The raw materials for many of these synthetic substances come from crude oil.
The uses of carbon and its compounds are extremely varied. It can form alloys with iron, of which the most common is carbon steel. Graphite is combined with clays to form the 'lead' used in pencils used for writing and drawing. It is also used as a lubricant and a pigment, as a moulding material in glass manufacture, in electrodes for dry batteries and in electroplating and electroforming, in brushes for electric motors and as a neutron moderator in nuclear reactors.
Charcoal is used as a drawing material in artwork, for grilling, and in many other uses including iron smelting. Wood, coal and oil are used as fuel for production of energy and space heating. Gem quality diamond is used in jewelry, and Industrial diamonds are used in drilling, cutting and polishing tools for machining metals and stone. Plastics are made from fossil hydrocarbons, and carbon fibre, made by pyrolysis of synthetic polyester fibres is used to reinforce plastics to form advanced, lightweight composite materials. Carbon fiber is made by pyrolysis of extruded and stretched filaments of polyacrylonitrile (PAN) and other organic substances. The crystallographic structure and mechanical properties of the fiber depend on the type of starting material, and on the subsequent processing. Carbon fibres made from PAN have structure resembling narrow filaments of graphite, but thermal processing may re-order the structure into a continuous rolled sheet[citation needed]. The result is fibers with higher specific tensile strength than steel.[citation needed]
Carbon black is used as the black pigment in printing ink, artist's oil paint and water colours, carbon paper, automotive finishes, India ink and laser printer toner. Carbon black is also used as a filler in rubber products such as tyres and in plastic compounds. Activated charcoal is used as an absorbent and adsorbent in filter material in applications as diverse as gas masks, water purification and kitchen extractor hoods and in medicine to absorb toxins, poisons, or gases from the digestive system. Carbon is used in chemical reduction at high temperatures. coke is used to reduce iron ore into iron. Case hardening of steel is achieved by heating finished steel components in carbon powder. Carbides of silicon, tungsten, boron and titanium, are among the hardest known materials, and are used as abrasives in cutting and grinding tools. Carbon compounds make up most of the materials used in clothing, such as natural and synthetic textiles and leather, and almost all of the interior surfaces in the built environment other than glass, stone and metal.
[edit] Production
Please help improve this section by expanding it. Further information might be found on the talk page or at requests for expansion. (December 2007)
[edit] Graphite Production
Commercially viable natural deposits of graphite occur in many parts of the world, but the most important sources economically are in China, India, Brazil, and North Korea.[56] Graphite deposits are of metamorphic origin, found in association with quartz, mica and feldspars in schists, gneisses and metamorphosed sandstones and limestone as lenses or veins, sometimes of a metre or more in thickness. Deposits of graphite in Borrowdale, Cumberland, England were at first of sufficient size and purity that, until the 1800s, pencils were made simply by sawing blocks of natural graphite into strips before encasing the strips in wood. Today, smaller deposits of graphite are obtained by crushing the parent rock and floating the lighter graphite out on water.
[edit] Precautions
Pure carbon has extremely low toxicity and can be handled and even ingested safely in the form of graphite or charcoal. It is resistant to dissolution or chemical attack, even in the acidic contents of the digestive tract, for example. Consequently if it gets into body tissues it is likely to remain there indefinitely. Carbon black was probably one of the first pigments to be used for tattooing, and Ötzi the Iceman was found to have carbon tattoos that survived during his life and for 5200 years after his death.[57] However, inhalation of coal dust or soot (carbon black) in large quantities can be dangerous, irritating lung tissues and causing the congestive lung disease coalworker's pneumoconiosis. Similarly, diamond dust used as an abrasive can do harm if ingested or inhaled. Microparticles of carbon are produced in diesel engine exhaust fumes, and may accumulate in the lungs.[58] In these examples, the harmful effects may result from contamination of the carbon particles, with organic chemicals or heavy metals for example, rather than from the carbon itself.
Carbon may also burn vigorously and brightly in the presence of air at high temperatures, as in the Windscale fire, which was caused by sudden release of stored Wigner energy in the graphite core. Large accumulations of coal, which have remained inert for hundred of millions of years in the absence of oxygen, may spontaneously combust when exposed to air, for example in coal mine waste tips.
The great variety of carbon compounds include such lethal poisons as tetrodotoxin, the lectin ricin from seeds of the castor oil plant Ricinus communis, cyanide (CN-) and carbon monoxide; and such essentials to life as glucose and protein.Compounds of carbon
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There is an immense number of distinct compounds that contain carbon atoms. Some sources suggest that this number is close to almost ten million known.[1] However, it is possible that the number is greater.
Contents
[hide]
* 1 Organic compounds
* 2 Inorganic compounds
o 2.1 Compounds with other nonmetals
o 2.2 Compounds with metals
+ 2.2.1 Carbonates and bicarbonates
+ 2.2.2 Carbonyls
+ 2.2.3 Compounds contanining the CN group
+ 2.2.4 Carbides
+ 2.2.5 Other
* 3 Alloys
* 4 Formation of carbon compounds
* 5 References
* 6 See also
[edit] Organic compounds
Main article: Organic compound
Every organic compound contains at least one atom of carbon. The number of these compounds is immense and the described number of defined compounds is close to 10 million. However, an indefinitely larger number of such compounds are theoretically possible.
There are several organic compounds sometimes considered as inorganic: NH2COONH4, COCl2, CSCl2, CS(NH2)2, CO(NH2)2
[edit] Inorganic compounds
See also: Inorganic compounds by element#Carbon
There is a rich variety of carbon chemistry that does not fall within the realm of organic chemistry and is thus called inorganic carbon chemistry.
[edit] Compounds with other nonmetals
Perhaps the best known are the oxides of carbon, carbon dioxide (CO2) and carbon monoxide (CO). Other known oxides are the uncommon carbon suboxide, C3O2, the uncommon dicarbon monoxide, C2O and even the exotic carbon trioxide (CO3).
Other (binary) compounds of carbon with nonmetals include: CS2, β-C3N4, CBr4, CCl4, CF4, COF2, COS, H2C2B10H10,
[edit] Compounds with metals
[edit] Carbonates and bicarbonates
Main articles: Carbonic acid, Carbonate, and Bicarbonate
The only known acid that is derived from the oxides of carbon is the carbonic acid (H2CO3). Upon monodeprotonation of this acid, bicarbonates are formed, which can be further derpotonated to carbonates.
Here is a list of carbonates and bicarbonates: NH4HCO3, (NH4)2CO3, BaCO3, CdCO3, Cs2CO3, Ca(HCO3)3, CaCO3, Ce2(CO3)3, CoCO3, CuCO3, FeCO3, PbCO3, La2(CO3)3, Li2CO3, MgCO3, MnCO3, NiCO3, KHCO3, K2CO3, Ag2CO3, NaHCO3, Na2CO3, SrCO3, ZnCO3
[edit] Carbonyls
Main article: Carbonyl
Carbonyls are coordination complexes between transition metals and carbonyl ligands. Metal carbonyls are complexes that are formed with the neutral ligand CO. These complexes are covalent. Here is a list of some carbonyls: Cr(CO)6, Co2(CO)8, Fe(CO)5, Mn2(CO)10, Mo(CO)6, Ni(CO)4, W(CO)6,
[edit] Compounds contanining the CN group
Main articles: Cyanide, Cyanates, Thiocyanate, and Isocyanate
Other types of inorganic compounds include inorganic salts and complexes of the carbon-containing polyatomic ions cyanide, isocyanide, cyanate, thiocyanate.
NH4SCN, CaNCN, Co(SCN)2, CuCN, (HCNO)x NH2CN HCNO, (CN)2, BrCN, ClCN, HCN, KOCN, KCN, K3Fe(CN)6, K4Fe(CN)6, KSCN, Fe4(Fe(CN)6)3, AgCN, NaOCN, NaCN, Na3Fe(CN)5NO, NaSCN, (SCN)2,
[edit] Carbides
Main article: Carbide
Carbides are binary compounds of carbon with an element that is less electronegative than it. B4C, CaC2 SiC, TaC, TiC, WC,
[edit] Other
The known inorganic chemistry of the allotropes of carbon (diamond, graphite, and the fullerenes) blossomed with the discovery of buckminsterfullerene in 1985, as additional fullerenes and their various derivatives were discovered. One such class of derivatives is inclusion compounds, in which an ion is enclosed by the all-carbon shell of the fullerene. This inclusion is denoted by the "@" symbol. For example, an ion consisting of a lithium ion trapped within buckminsterfullerene would be denoted Li+@C60. As with any other ionic compound, this complex ion could in principle pair with a counterion to form a salt.
[edit] Alloys
There are several alloys that contain carbon of which the best known alloy is carbon steel (see category:steels)). Besides steel, other alloys based on iron and carbon are: anthracite iron, cast iron, pig iron, wrought iron, but also spiegeleisen (which contains also manganese). Stellite is an alloy of carbon with cobalt, chromium and tungsten. To some degree, these alloys could be considered carbides.
[edit] Formation of carbon compounds
In organic chemistry there are 3 important elements: Carbon, Oxygen and Hydrogen. Each of these elements have different kinds of bonds. Carbon atom has tetravalent bonds, Oxygen atoms divalent bonds and Hydrogen monovalent bonds.Organic compound
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Methane is the simplest possible organic compound
Methane is the simplest possible organic compound
An organic compound is any member of a large class of chemical compounds whose molecules contain carbon. For historical reasons discussed below, a few types of compounds such as carbonates, simple oxides of carbon and cyanides, as well as the allotropes of carbon, are considered inorganic. The division between "organic" and "inorganic" carbon compounds while "useful in organizing the vast subject of chemistry...is somewhat arbitrary"[1].
Organic chemistry is the science concerned with all aspects of organic compounds. Organic synthesis is the methodology of their preparation.
Contents
[hide]
* 1 History
* 2 Classification
o 2.1 Natural compounds
o 2.2 Synthetic compounds
* 3 Nomenclature
* 4 Databases
* 5 Structure determination
* 6 See also
* 7 References
[edit] History
The "organic" is a historical name, dating back to 19th century, when it was believed that organic compounds could only be synthesized in living organisms through vis vitalis - the "life-force". The theory that organic compounds were fundamentally different from those that were "inorganic", that is, not synthesized through a life-force, was disproved with the synthesis of urea, an "organic" compound by definition of its known occurrence only in the urine of living organisms, from potassium cyanate and ammonium sulfate by Friedrich Wöhler in the Wöhler synthesis. The kinds of carbon compounds that are still traditionally considered inorganic are those that were considered inorganic before Wöhler's time; that is, those which came from "inorganic" (i.e., lifeless) sources such as minerals.[1]
[edit] Classification
See Organic chemistry#Classification of organic substances
Organic compounds may contain atoms of further elements, so-called heteroatoms. Organometallic compounds constitute a further subsection, characterized by covalent bonds between organic carbon and a metal.
There is also a large number of inorganic carbon compounds to distinguish from organic compounds.
[edit] Natural compounds
An important subset of organic compounds is still extracted from natural sources because they would be far too expensive to be produced artificially. Examples include most sugars, some alkaloids and terpenoids, certain nutrients such as vitamin B12, and in general, those natural products with large or stereoisometrically complicated molecules which are present in reasonable concentrations in living organisms.
Further compounds of prime importance in biochemistry are antigens, carbohydrates, enzymes, hormones, lipids and fatty acids, neurotransmitters, nucleic acids, proteins, peptides and amino acids, vitamins and fats and oils.
[edit] Synthetic compounds
Many polymers, including all plastics are organic compounds.
[edit] Nomenclature
The IUPAC nomenclature of organic compounds slightly differs from the CAS nomenclature.
[edit] Databases
* The CAS database is the most comprehensive repository for data on organic compounds. The search tool SciFinder is offered .
* The Beilstein database contains information on 9.8 million substances, covers the scientific literature from 1771 to the present, and is today accessible via CrossFire. Structures and a large diversity of physical and chemical properties is available for each substance, with reference to original literature.
* PubChem contains 18.4 million entries on compounds and especially covers the field of medicinal chemistry.
There is a great number of more specialized databases for diverse branches of organic chemistry.
[edit] Structure determination
See Structure determination
Today, the main tools are proton and carbon-13 NMR spectroscopy and X-ray crystallography.List of organic compounds
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This page aims to list well-known organic compounds, including organometallic compounds, to stimulate the creation of Wikipedia articles. Note that purely inorganic compounds, minerals, and chemical elements are not included on this list. There are also no generic terms (e.g., carbohydrate) or mixtures of no fixed composition (e.g., naphtha, gasoline). Compounds and enzymes that are overwhelmingly of interest to biochemists, such as Cytochrome c peroxidase, are listed under list of biomolecules.
This list is not necessarily complete or up to date — if you see an article that should be here but isn't (or one that shouldn't be here but is), please update the page accordingly.
For substances with a number prefix such as 2-Butanol or 1,3-Cyclohexadiene, please use the first letter of the name (in this case under B or C) to find the compound. Note that such names usually have the first letter capitalized in a title or at the beginning of a sentence.
Relevant links for chemical compounds are:
* The CAS Substance Databases, which contains information on about 23 million compounds
* ChemIDplus [1] is a useful non-commercial source for chemical lookups
* NIST Chemistry WebBook [2] is a freely available resource compiled by National Institute of Standards and Technology under the Standard Reference Data Program. Apart from chemical structures, it contains a wealth of associated physico-chemical information such as thermochemistry data and spectra
* ChEBI [3], a freely available dictionary of molecular entities focused on ‘small’ chemical compounds
* PubChem [4], maintained by the National Center for Biotechnology Information (NCBI), serves as a repository of chemical compounds from many public and commercial resources
* http://physchem.ox.ac.uk/MSDS/ Material Safety Data Sheets, plus other relevant links
These (commercial) links may also provide useful information:
* Chemfinder [5] is helpful for finding information about a chemical (disable and delete cookies!)
* Sigma Aldrich [6]
* Acros Organics [7]
* Lancaster [8]
* Chemical Suppliers Directory [9]
* ChemSpider [10]has over 20 million structures with chemical names and the ability to download the molfile locally. It includes links to chemical vendors, PubChem, ChEBI and over 100 other sources and is curated by users.
Whilst most compounds are referred to by their IUPAC name, "traditional" names have also been kept where they are in wide use or of significant historical interest.
See also: organic compound, list of compounds, list of inorganic compounds, inorganic compounds by element, list of biomolecules, polyatomic ions, list of elements by name, list of alchemical substances, list of drugs, list of reactions.
Table of contents: A B C D E F G H I J K L M N O P Q R S T U V W X Y Z
[edit] A
For substances with an A- or α- prefix such as α-Terpinene, please see the parent page (in this case Terpinene).
* Abietic acid - C20H30O2
* Acenaphthene
* Acenaphthoquinone
* Acenaphthylene
* Acepromazine
* Acetaldehyde — CH3CHO, also known as ethanal
* Acetamide
* Acetaminophen — C8H9NO2
* Acetaminosalol
* Acetamiprid
* Acetanilide
* Acetic acid — CH3COOH, also known as ethanoic acid, Glacial acetic acid or GAA
* Acetoguanamine
* Acetone — CH3COCH3, or (CH3)2CO
* Acetonitrile
* Acetophenone
* Acetylcholine – (CH3)3N+CH2CH2OCOCH3.
* Acetylene — C2H2
* N-Acetylglutamate
* Acetylsalicylic Acid also known as Aspirin
* Acid fuchsin
* Acridine — C13H9N
* Acridine orange
* Acrolein
* Acrylamide — C3H5NO
* Acrylic acid — CH2=CHCOOH
* Acrylonitrile
* Acryloyl chloride
* Acyclovir
* Adamantane
* Adenosine
* Adipamide
* Adipic acid
* Adiponitrile
* Adipoyl dichloride
* Adonitol
* Adrenaline, epinephrine
* Adrenochrome
* Aflatoxin
* Alanine
* Albumins
* Alcian blue
* Aldosterone
* Aldrin
* Aliquat 336
* Alizarin
* Allantoic acid
* Allantoin
* Allethrin
* Allyl propyl disulfide
* Allylamine
* Allyl chloride
* Amido black 10b
* p-Aminobenzoic acid (PABA)
* Aminodiacetic acid
* Aminoethylpiperazine
* 5-Amino-2-hydroxybenzoic acid
* Aminophylline
* 5-Aminosalicylic acid
* Aminothiazole
* Amiodarone
* Amiton
* Amobarbital
* Amoxicillin — C16H19N3O5S.3H2O
* Amphetamine
* Amyl nitrate
* Amyl nitrite — C5H11A.ONO
* Anethole
* Angelic acid
* Anilazine
* Aniline — C6H5-NH2
* Aniline hydrochloride
* Anisole
* Anisoyl chloride
* Anthanthrene
* Anthracene – (C6H4CH)2
* Anthramine
* Anthranilic acid
* Anthraquinone
* Anthrone
* Antipyrine
* Aprotinin
* Arabinose
* Arginine
* Aroclor (polychlorinated biphenyls)
* Ascorbic acid (vitamin C)
* Asparagine
* Asparagusic acid
* Aspartame
* Aspartic acid
* Asphidophytidine
* Astrablue
* Atrazine
* Auramine o
* Aureine
* Avobenzone
* Azadirachtin A — C35H44O16
* Azathioprine
* Azelaic acid
* Azinphos-methyl
* Aziridine
* Azithromycin
* Azo violet
* Azobenzene
* Azulene
* Azure a
[edit] B
For substances with a B- or β- prefix such as β-Pinene, please see the parent page (in this case Pinene).
* Bacillomycin
* Barbital
* Barbituric acid
* Behenic acid
* Benomyl
* Benzaldehyde
* Benzalkonium chloride
* Benzamide
* Benzanthrone
* Benzene — C6H6
* Benzethonium chloride
* Benzidine
* Benzil
* Benzilic acid
* Benzimidazole
* Benzisothiazolinone
* Benzisoxazole
* Benzo(a)anthracene
* Benzo(c)cinnoline
* Benzo(a)pyrene
* Benzo(c)phenanthrene
* Benzo(e)fluoranthene
* Benzo(e)pyrene
* Benzo(ghi)perylene
* Benzo(j)fluoranthene
* Benzo(k)fluoranthene
* Benzo(c)thiophene
* Benzocaine
* Benzofuran
* Benzoic acid
* Benzoin
* Benzothiazole
* Benzothiophene
* Benzotriazole
* Benzoxazole
* Benzoyl chloride
* Benzyl alcohol
* Benzyl chloroformate
* Benzylamine
* Benzyldimethylamine
* Benzylidene acetone
* Betaine
* Betulin
* Butylated hydroxytoluene (BHT) – C6H2(OH)(CH3)(C(CH3)3)2
* Biotin (Vitamin H)
* Biphenyl
* 2,2'-Bipyridyl = 2,2'-Bipyridine
* 1,8-Bis(dimethylamino)naphthalene (Proton-sponge, Aldrich trademark name)
* Bis(chloromethyl) ether
* Bismarck brown y
* Bisphenol A
* Biuret
* Borneol
* Brassinolide
* Brilliant cresyl blue
* Bromacil
* Bromoacetic acid
* Bromobenzene
* 2-Bromo-1-chloropropane
* Bromocresol purple
* Bromocyclohexane
* Bromoform
* Bromomethane — BrCH3
* Bromophenol blue
* 2-Bromopropane
* Bromothymol blue
* Bromotrifluoromethane
* Brucine
* Buckminsterfullerene
* Buspirone
* 1,3-Butadiene
* Butadiene resin
* Butane — C4H10
* Butene
* 2-Butoxyethanol
* Butylamine = n-Butylamine
* Butyllithium
* 2-Butyne-1,4-diol
* Butyraldehyde
* Butyrophenone
* Butyryl chloride
[edit] C
For substances with an c- or cis- prefix such as cis-3-hexenal, you may find these listed under the parent name letter (in this case "H"), as is the norm in chemical catalogues.
* Cacodylic acid
* Cacotheline
* Cadaverine — NH2(CH2)5NH2
* Cadinene
* Cafestol
* Caffeine
* Calcein
* Calciferol (Vitamin D)
* Calcitonin
* Calmodulin
* Calreticulin
* Camphene
* Camphor
* Cannabinol
* Caprolactam
* Caprolactone
* Capsaicin
* Captan
* Captopril
* Carbazole
* Carbazol-9-yl-methanol (N-(Hydroxymethyl)carbazole)
* Carbofuran
* Carbonyl fluoride
* Carboplatin
* Carboxypolymethylene
* Carminic acid
* Carnauba wax
* Carnitine
* Cartap
* Carvacrol
* Carvone
* Castor oil
* Catechol
* Cedar wood oil
* Cefazolin
* Cefotaxime
* Ceftriaxone
* Cellulose
* Cellulose acetate
* Cetrimide
* Cetyl alcohol
* Chloracetyl chloride
* Chloral
* Chloral hydrate
* Chlorambucil
* Chloramine-T
* Chloramphenicol
* Chloranilic acid
* Chlordane
* Chlorhexidine gluconate
* Chloro-m-cresol
* Chloroacetic acid
* 4-Chloroaniline (p-Chloroaniline)
* Chlorobenzene
* 2-Chlorobenzoic acid (o-Chlorobenzoic acid)
* Chlorodifluoromethane
* Chloroethene — C2H3Cl
* Chlorofluoromethane
* Chloroform — CHCl3
* Chloromethane
* 2-Chloro-2-methylpropane (tert butyl chloride)
* Chloronitroaniline
* Chloropentafluoroethane
* Chloropicrin
* Chloroprene
* Chloroquine
* Chlorostyrene
* Chlorothiazide
* Chlorotrifluoromethane
* Chlorotrimethylsilane
* Chloroxuron
* Chlorpyrifos
* Chlorthiamide
* Cholesterol
* Choline
* Chromotropic acid
* Cilostazol
* Cinchonine
* Cinnamaldehyde
* Cinnamic acid
* Cinnamyl alcohol
* Cinnoline — C4H4N2
* cis-2-butene
* cis-3-Hexenal
* cis-3-Hexen-1-ol
* Citral
* Citric acid — C3H4OH(COOH)3
* Citronella oil
* Citronellal
* Citrulline
* Clobetasone
* Clopidol
* Cloxacillin — C19H17ClN3O5S*Na*H2O
* Cobalamin (Vitamin B12)
* Cocamidopropyl
* Colchicine
* Collagen
* Collodion
* Congo red
* Coniine
* Coomassie blue
* Coronene
* Coumarin
* Creatine
* Cresol
* Cresyl violet
* Crotonaldehyde
* 18-Crown-6
* Crystal violet
* Cubane
* Cumene
* Cuneane
* Cupferron
* Cuscohygrine
* Cyanogen
* Cyanogen chloride
* Cyanoguanidine
* Cyanuric acid
* Cyanuric chloride
* Cyclodecane
* α-Cyclodextrin
* Cyclododecane
* Cycloheptatriene
* 1,3-Cyclohexadiene
* 1,4-Cyclohexadiene
* Cyclohexane
* Cyclohexanol
* Cyclohexanone
* Cyclohexene
* Cyclonite - (CH2-N-NO2)3
* Cyclooctatetraene
* Cyclopentadiene — C5H6
* Cyclopentane
* Cyclopentanol
* Cyclopentanone
* Cyclopentene
* Cypermethrin
* Cysteamine
* Cysteine
* Cystine
* Cytosine — C4H5N3O
[edit] D
For substances with a d- or D- prefix such as D-alanine or DL-alanine, please see the parent page (in this case alanine).
* DABCO
* DDT
* Decaborane
* Decabromodiphenyl ether
* Decahydronaphthalene
* Decane — C10H22
* Dehydroacetic acid
* Dehydrocholic acid
* Deltamethrin
* Demeton
* Denatonium
* Dexamethazone
* Dextran
* Dextrin
* 3,3'-Diaminobenzidine
* Di-t-butyl peroxide
* Diacetylene
* Diazinon
* Diazomethane
* 1,2-Dibromoethane
* Dibucaine hydrochloride
* Dichloroacetic acid
* p-Dichlorobenzene
* Dichlorobutane
* Dichlorodifluoromethane
* Dichlorodimethylsilane
* 1,2-Dichloroethane
* Dichlorofluoromethane
* Dichlorophen
* 2,4-Dichlorophenoxyacetic acid
* Dichlorotrifluoroethane
* Dichlorvos
* Diclofenac sodium
* Dicofol
* Dicrotophos
* Dicyclopentadiene
* Dieldrin
* Diethanolamine
* Diethion
* Diethyl aluminium chloride a Lewis acid
* Diethylamine
* Diethylene glycol
* Diethylenetriamine
* Diethyl ether
* Difluoromethane
* Digitonin
* Dihydrocortisone
* Diisoheptyl phthalate
* Diisopropyl ether
* Diketene
* Dimethicone
* Dimethylamine
* N,N-Dimethylacetamide
* N,N-dimethylaniline
* 1,2-Dimethylbenzene (o-Xylene)
* 1,3-Dimethylbenzene (m-Xylene)
* 1,4-Dimethylbenzene (p-Xylene)
* N,N-dimethylformamide
* Dimethyldiethoxysilane
* Dimethylglyoxime
* Dimethylmercury
* Dimethyl sulfoxide
* Dinoseb
* Dioctyl phthalate
* Dioxane
* Dioxathion
* Dioxin
* Diphenylacetylene (Tolane)
* Diphenylmethanol (Benzhydrol)
* Diquat
* Direct Blue 1
* Disulfiram
* Disulfoton
* Dithranol
* 2,6-Di-tert-butylphenol
* 2,6-Di-tert-butyl-4-methylphenol
* 2,6-Di-tert-butylpyridine
* Diuron
* Divinylbenzene
* Docosane
* Dodecane
* Dodecylbenzene
* Domperidone
* Dopamine
* Doxylamine succinate
[edit] E
* EDTA (Ethylenediamine-N,N,N',N'-tetraacetic acid)
* Eicosane
* Endosulfan
* Endrin
* Eosin
* Ephedrine
* Epibromohydrin
* Epinephrine — C9H13NO3
* Erucic acid — CH3(CH2)7CH=CH(CH2)11COOH
* Erythritol
* Estradiol
* Ethacridine lactate
* Ethane — C2H6
* 1,2-Ethanedithiol — C2H4(SH2)2
* Ethanol — CH3CH2OH
* Ethene — C2H4
* Ethidium bromide
* Ethyl acetate
* Ethylamine
* Ethyl 4-aminobenzoate (Ethyl p-aminobenzoate)
* Ethylbenzene
* Ethyl chloride
* Ethylene
* Ethylene glycol — OHCH2CH2OH
* Ethylene oxide
* Ethyl formate
* 2-Ethyl-1-hexanol
* Eugenol
[edit] F
* Farnesol
* Ferrocene
* Fipronil
* Flunixin
* Fluoranthene
* Fluorene
* 9-Fluorenone
* Fluorescein
* Fluorobenzene
* Fluoroethylene
* Fluoxetine
* Folic acid (Vitamin M)
* Fonofos
* Formaldehyde — HCHO
* Formamide
* Formanilide
* Formic acid — HCOOH
* Formoterol
* Fructose
* Fumaric acid
* Furan (furane)
* Furfural
* Furfuryl alcohol
* Furfurylamine
* Furylfuramide
[edit] G
* Gadopentetate - also known as Magnevist
* Galactose
* Gamma-aminobutyric acid
* Gamma-butyrolactone
* Gamma-hydroxybutyrate (GHB)
* Geraniol
* Gibberellic acid
* Gluconic acid
* Glucose — C6H12O6
* Glutamic acid (glutamate)
* Glutamine
* Glutaraldehyde
* Glutaric acid
* Glutathione
* Glyburide
* Glycerin (glycerol)
* Glycerol (glycerin)
* Glycerophosphoric acid
* Glycidol
* Glycine — NH2CH2COOH
* Glycogen
* Glycolic acid
* Glyoxal
* Guaiacol
* Guanidine
* Guanine
* Guanosine
[edit] H
* Halothane
* Hematoxylin
* HEPES
* Heptadecane
* Heptane — C7H16
* Hexabromocyclododecane
* Hexachloropropene
* Hexadecane
* Hexafluoro-2-propanol
* Hexafluoro-2-propanone
* Hexafluoroethane
* Hexafluoropropylene
* Hexamethyldewarbenzene
* Hexamethyldisilazane
* Hexamethylenimine
* Hexamethylolmelamine
* Hexamine – (CH2)6N4
* Hexane — C6H14
* Hexanitrodiphenylamine
* Hexanoic acid
* cis-3-Hexanal
* cis-3-Hexen-1-ol
* Hippuric acid
* Histidine — NH2CH(C4H5N2)COOH
* Histamine
* Homoarginine
* Homocysteine
* Homocystine
* Homotaurine
* Hydrochlorothiazide
* Hydrocinnamic acid
* Hydroquinone
* Hydroxyproline
* 5-Hydroxytryptamine
* Hygrine
[edit] I
* Ibuprofen
* Imazapyr
* Imidazole
* Imiquimod
* Indazole
* Indene
* Indigo
* Indole
* Indoline
* Indole-3-acetic acid
* Inositol
* Iodoxybenzene
* Ionone
* Ipratropium bromide
* Isatin
* Isoamyl isobutyrate
* Isobenzofuran
* Isoborneol
* Isobornyl acetate
* Isoflurane
* Isoindole
* Isoleucine
* Isomelamine
* Isooctanol
* Isophthalic acid
* Isopropanol – (CH3)2CHOH, also known as IPA, IsoPropyl alcohol
* Isoquinoline
* Isoxazole
* Itraconazole
[edit] J
* Jasmone
* Jenner's stain
[edit] K
* Kanamycin
* Kepone alcohol
* Keratin
* Ketene
* Kojic acid
[edit] L
For substances with an l- or L- prefix such as L-alanine or DL-alanine, please see the parent page (in this case alanine).
* Lactic acid — CH3CH(OH)COOH
* Lactose
* Lauric acid
* Lauryl alcohol
* LDA (Lithium diisopropylamide)
* Leucine
* Levulinic acid
* Limonene
* Linalool
* Linoleic acid
* Linolenic acid
* Lipoamide
* Lithium diisopropylamide
* Loratadine
* LSD
* Luminol
* 2,6-Lutidine
* Lycopene
* Lysine
[edit] M
For substances with an m- or meta- prefix such as m-cresol, meta-cresol or metacresol that are not listed below, please look for a more generic page (in this case cresol). For substances with a meso- prefix such as meso-tartaric acid or mesotartaric acid that are not listed below, please see the parent page (in this case tartaric acid).
* Malachite green
* Malathion
* Maleic anhydride
* Malic acid
* Maltose
* Mandelonitrile
* Mannide monooleate
* Mannose
* Mauveine
* MDMA
* Mecoprop
* MEK
* Melatonin
* Meldola's blue
* Meloxicam
* Menthol
* 2-Mercaptoethanol
* 2-Mercaptopyridine
* Merocyanine
* Mesityl oxide
* Mesitylene – (CH3)3-C6H3
* Mesotartaric acid
* Metaldehyde
* Metamizole (dipyrone)
* Methane — CH4
* Methanesulfonic acid
* Methanol — CH3OH
* Methionine
* Methomyl
* 4-Methoxybenzaldehyde (anisaldehyde)
* Methoxychlor
* Methoxyflurane
* Methyl acetate
* Methyl-2-cyanoacrylate
* Methyl ethyl ketone (MEK)
* Methyl isobutyl ketone (MIBK)
* Methyl isocyanate — CH3-N=C=O
* Methyl methacrylate
* Methyl tert-butyl ether (MTBE)
* Methylal
* Methylamine
* 2-Methylbenzoic acid (o-Toluic acid)
* 4-Methylbenzoic acid (p-Toluic acid)
* Methyl chloroformate
* Methylcyclohexane
* Methylene blue — C16H18ClN3S
* Methylhydrazine
* Methylmercury
* Methylmorpholine
* 2-Methylpropene (isobutylene)
* N-Methylpyrrolidone — C5H9NO
* Methyltriethoxysilane
* Methyltrimethoxysilane
* Metoprolol
* Metronidazole
* Michler's ketone
* Milrinone
* Monocrotophos
* Monosodium glutamate
* Mordant red 19
* Morpholine
* MTBE
* Murexide
* Mustard gas — C4H8Cl2S
* Myrcene
[edit] N
For substances with an n- or normal- prefix such as n-pentane that are not listed below, please see the parent page (in this case pentane).
For substances with an N- prefix (meaning on nitrogen) such as N,N-dimethylformamide, if these are not listed below please see the parent page (in this case dimethylformamide).
* n-Nonadecane
* n-Tetradecylbenzene
* Naphthalene — C10H8
* Naphthoquinone (Vitamin K)
* 2-Naphthylamine
* Neomycin
* Niacin or nicotinic acid (Vitamin B3)
* Nicotine
* Niflumic acid
* Nile red
* Nimesulide
* Nitrilotriacetic acid
* Nitrobenzene
* Nitrocellulose
* Nitroethane
* Nitrofen
* Nitrofurantoin
* Nitroglycerine — C3H5(NO2)3
* Nitromethane
* Nitrosobenzene
* N-Nitroso-N-methylurea
* Nitrosomethylurethane
* Nominine
* Nonacosane
* Nonane — C9H20
* Noradrenaline, norepinephrine
* Norephidrine
* Norcarane
* Norleucine
* Nujol
* NMN
[edit] O
For substances with an o- or ortho- prefix such as o-cresol, ortho-cresol or orthocresol that are not listed below, please look for a more generic page (in this case cresol).
* Octabromodiphenyl ether
* Octane — C8H18
* 1-Octanethiol
* Octanoic acid
* 4-Octylphenol
* Oleic acid
* Orcin
* Orcinol
* Ornithine
* Orotic acid
* Oxalic acid
* Oxalyl Chloride — C2O2Cl2
* Oxamide
* Oxazole
* Oxolinic acid
* Oxymetholone
[edit] P
For substances with an p- or para- prefix such as p-cresol, para-cresol or paracresol that are not listed below, please look for a more generic page (in this case cresol).
* p-nitro benzal dehyde
* PABA
* Paclitaxel
* Palmitic acid
* Pantothenic acid (Vitamin B5)
* Para red
* Parachlorometaxylenol (PCMX)
* Paraformaldehyde
* Parathion
* Pelargonic acid
* Pentabromodiphenyl ether
* Pentachlorobiphenyl
* Pentachlorophenol
* Pentadecane
* Pentaerythritol
* Pentaethylene glycol
* Pentafluoroethane
* Pentane — C5H12
* Pentetic acid
* Perfluorotributylamine
* Permethrin
* Peroxyacetic acid
* Perylene
* Petroleum ether
* Phenacetin
* Phenacyl bromide
* Phenanthrene
* Phenanthrenequinone
* Phencyclidine
* Phenethylamine
* Phenobarbital (c-iv)
* Phenol — C6H5OH
* Phenol red, sodium salt
* Phenolphthalein
* Phenothiazine
* Phenylacetic acid
* Phenylacetylene
* Phenylalanine
* p-Phenylenediamine
* Phenylhydrazine
* Phenylhydroxylamine
* Phenyllithium
* 4-Phenyl-4-(1-piperidinyl)cyclohexanol (PPC)
* Phenylthiocarbamide — C7H8N2S
* Phloroglucinol
* Phorate
* Phthalic anhydride
* Phthalic acid
* Phytic acid
* 4-Picoline
* Picric acid — C6H2(OH)(NO2)3
* Pimelic acid
* Pinacol
* Piperazine
* Piperidine
* Piperonal
* Piperylene
* Pivaloyl chloride
* Polyacrylonitrile
* Polyamide 6 = Nylon 6
* Polybenzimidazole - Polybenzimidazole fiber
* Polyethylenimine
* Polygeline
* Polyisobutylene
* Polypropylene
* Polypropylene glycol
* Polystyrene
* Polyurethane
* Polyvinyl acetate
* Polyvinyl alcohol
* Polyvinyl chloride
* Polyvinylidene chloride
* Polyvinylidene fluoride (PVDF)
* Polyvinylpyrrolidone = Poly vinyl pyrrolidone (PVP)
* Porphyrin
* Potassium clavulanate — C8H8KNO5
* Potassium 2-ethyl hexanoate — C8H15KO2
* Prednisone
* Primaquine
* Procaine
* Progesterone
* Prolactin (PRL)
* Proline
* Propane — C3H8
* Propanoic acid
* 2-Propanone
* Propargyl alcohol
* Propiconazole
* Propiolactone
* Propiolic acid
* Propionaldehyde
* Propionitrile
* Propoxur
* Proton-sponge (Aldrich trademark name)
* Purine
* Putrescine — C4H12N2
* Pyrazine
* Pyrazole
* Pyrene
* Pyrethrin
* Pyridazine
* Pyridine — C5H5N
* Pyridinium tribromide
* 2-Pyridone
* Pyridoxal
* Pyridoxine or pyridoxamine (Vitamin B6)
* Pyrilamine
* Pyrimethamine
* Pyrimidine — C4H4N2
* Pyrocatechol violet
* Pyroglutamic acid
* Pyrrole
* Pyrrolidine
* Pyruvic acid
[edit] Q
* Quinaldine
* Quinazoline
* Quinhydrone
* Quinoline
* Quinone
* Quinoxaline
[edit] R
* Raffinose
* Resorcinol
* Retinene
* Retinol (Vitamin A)
* Rhodanine
* Riboflavin (vitamin B2)
* Ribofuranose
* Ribose
* Ricin
* Rosolic acid
* Rotane
* Rotenone
[edit] S
For substances with an s- or secondary- prefix such as s-butyllithium or sec-butyllithium that are not listed below, please see the parent page (in this case under B, butyllithium).
* Saccharin
* Safrole
* Salicin
* Salicylaldehyde
* Salicylic acid
* Salvinorin A
* Sarin
* Sclareol
* Sebacic acid
* Sebacoyl chloride
* Selacholeic acid
* Selenocysteine
* Selenomethionine
* Serine
* Serine kinase
* Serotonin
* Shikimic acid
* Sildenafil (also known as Viagra)
* Skatole
* Snakeroot oil
* Sorbic acid
* Sotolone
* Spermidine
* Squalene
* Stearic acid
* Strychnine
* Styrene
* Succinic anhydride
* Sucrose (sugar)
* Sulfanilamide
* Sulfanilic acid
* Sulforhodamine b
* Suxamethonium chloride
[edit] T
For substances with an t- or tertiary- prefix such as t-butyllithium or tert-butyllithium that are not listed below, please see the parent page (in this case under B, butyllithium). For substances with an t- or trans- prefix such as *trans-2-Butene, you may find these listed under the parent name letter (in this case "B"), as is the norm in chemical catalogues.
* Tabun — C2H5OP(O)(CN)N(CH3)2
* Tannic acid
* Tannin
* Tartaric acid
* Tartrazine
* Taurine
* Terephthalic acid
* Terephthalonitrile
* p-Terphenyl
* α-Terpineol
* Testosterone
* Tetrachlorobiphenyl
* Tetrachloroethylene
* Tetrachloromethane (carbon tetrachloride) – CCl4
* Tetradecane
* Tetraethylene glycol
* Tetrafluoroethene
* Tetrahedrane
* Tetrahydrofuran
* Tetrahydronaphthalene
* Tetramethrin
* Tetramethylsilane (TMS, standard for NMR)
* Tetramethylurea
* Tetranitromethane
* Tetrathiafulvalene (TTF)
* Tetrazine — C2H2N4 (hypothetical)
* Tetrodotoxin
* Tetryl — C7H5N5O8
* Thalidomide
* Thiamine (Vitamin B1) – C12H17ClN4OS·HCl
* Thiazole
* Thioacetamide
* Thiolactic acid
* Thiophene
* Thiophosgene
* Thiourea
* Thiram
* Thorin
* Threonine
* Thrombopoietin
* Thymidine
* Thymine
* Thymol
* Thymolphthalein
* Thyroxine (T4)
* Tiglic acid
* Tinidazole
* Tocopherol (Vitamin E)
* Toluene — C6H5CH3
* Toluene diisocyanate
* p-Toluenesulfonic acid
* o-Toluic acid (2-Methylbenzoic acid)
* p-Toluic acid (4-Methylbenzoic acid)
* Toxaphene
* Triangulane
* Triazole
* Tributyl phosphate
* Tributylamine
* Tributylphosphine
* Trichloroacetic acid
* Trichloroacetonitrile
* 1,1,1-Trichloroethane
* Trichloroethylene
* Trichlorofluoromethane
* 2,4,6-Trichloroanisole
* 2,4,6-Trichlorophenol
* Tris
* Tricine
* Triclabendazole
* Triclosan
* Tricosane
* Tridecane
* Tridecanoic acid
* Triethylaluminium
* Triethylamine
* Triethylamine hydrochloride — C6H15N·HCl
* Triethylene glycol
* Triethylenediamine
* Trifluoroacetic acid (TFA)
* 1,1,1-Trifluoroethane
* 2,2,2-Trifluoroethanol
* Trifluoromethane
* Trimellitic anhydride
* Trimethoxyamphetamine
* Trimethyl phosphite
* Trimethylamine
* Trimethylbenzene
* 2,2,4-Trimethylpentane (isooctane)
* Trinitrotoluene (TNT) – C6H2(NO2)3CH3 or 2,4,6-trinitrotoluene
* Tri-o-cresyl phosphate
* Triphenyl phosphate
* Triphenylamine
* Triphenylantimony
* Triphenylene
* Triphenylmethane
* Triphenylmethanol
* Triphenylphosphine
* Tropane
* Tropinone
* Trypan blue
* Tryptophan
* Tyrosine — C9O3H11N
[edit] U
* Umbelliferone
* Undecanol
* Uracil
* Urea — CO(NH2)2
* Urethane
* Uric acid — C5H4N4O3
* Uridine
* Usnic acid
[edit] V
* Valine
* Valium
* Valproic acid
* Vanillin
* Venlafaxine
* Vigulariol
* Vinyl acetate
* Vinyl fluoride
* Vinylidene chloride
* Violanthrone-79 (16,17-bis(octyloxy)anthra[9,1,2-cde]benzo[rst]pentaphene-5,10-dione)
* Vitamin A (retinol)
* Vitamin B
* Vitamin B1 (thiamine)
* Vitamin B2 (riboflavin)
* Vitamin B3 (niacin or nicotinic acid)
* Vitamin B4 (adenine)
* Vitamin B5 (pantothenic acid)
* Vitamin B6 (pyridoxine or pyridoxamine)
* Vitamin B12 (cobalamin)
* Vitamin C (ascorbic acid)
* Vitamin D (calciferol)
* Vitamin E (tocopherol)
* Vitamin F
* Vitamin H (biotin)
* Vitamin K (naphthoquinone)
* Vitamin M (folic acid)
* Vitamin P (niacin or nicotinic acid)
* Vitamin S
[edit] W
* Warfarin
[edit] X
* Xanthan gum
* Xanthone
* Xylene
* Xylene cyanole ff
* Xylenol orange
* Xylose
* Xylyl bromide
[edit] Y
* Yohimbine hydrochloride - C21H26N2O3
* Yohimbinic acid monohydrate
[edit] Z
* Zingiberene
From Wikipedia, the free encyclopedia
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For other uses, see Carbon (disambiguation).
6 boron ← carbon → nitrogen
-
↑
C
↓
Si
Periodic table - Extended periodic table
General
Name, symbol, number carbon, C, 6
Chemical series nonmetals
Group, period, block 14, 2, p
Appearance black (graphite)
colorless (diamond)
Standard atomic weight 12.0107(8) g·mol−1
Electron configuration 1s2 2s2 2p2
Electrons per shell 2, 4
Physical properties
Phase solid
Density (near r.t.) (graphite) 1.9-2.3[1] g·cm−3
Density (near r.t.) (diamond) 3.5-3.53[1] g·cm−3
Density (near r.t.) (fullerene) 1.69[1] g·cm−3
Heat of fusion (graphite) ? 100 kJ·mol−1
Heat of fusion (diamond) ? 120 kJ·mol−1
Heat of vaporization 715 kJ·mol−1
Specific heat capacity (25 °C) (graphite)
8.517 J·mol−1·K−1
Specific heat capacity (25 °C) (diamond)
6.115 J·mol−1·K−1
Vapor pressure (graphite) P/Pa 1 10 100 1 k 10 k 100 k
at T/K 2839 3048 3289 3572 3908
Atomic properties
Crystal structure (graphite) hexagonal
Oxidation states 4, 3 [2], 2, 1 [3], 0, -1, -2, -3, -4[4]
(mildly acidic oxide)
Electronegativity 2.55 (Pauling scale)
Ionization energies
(more) 1st: 1086.5 kJ·mol−1
2nd: 2352.6 kJ·mol−1
3rd: 4620.5 kJ·mol−1
Atomic radius 70 pm
Atomic radius (calc.) 67 pm
Covalent radius 77 pm
Van der Waals radius 170 pm
Miscellaneous
Magnetic ordering diamagnetic
Electrical resistivity (graphite) 1.375*10-5 [5]Ω·m
Thermal conductivity (300 K) (graphite)
(80–230) W·m−1·K−1
Thermal conductivity (300 K) (diamond)
(900–2320) W·m−1·K−1
Thermal diffusivity (300 K) (diamond)
(503–1300) mm²/s
Mohs hardness (graphite) 1-2 [6]
Mohs hardness (diamond) 10.0 [6]
CAS registry number 7440-44-0
Selected isotopes
Main article: Isotopes of carbon iso NA half-life DM DE (MeV) DP
15
12C 98.9% 12C is stable with 6 neutrons
13C 1.1% 13C is stable with 7 neutrons
14C trace 5730 y beta- 0.156 14N
References
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Carbon (pronounced /kɑɹbən/) is a chemical element with the symbol C and atomic number is 6. As a member of group 14 on the periodic table, it is nonmetallic and tetravalent—making four electrons available to form covalent chemical bonds. There are three naturally occurring isotopes, with 12C and 13C being stable, while 14C is radioactive, decaying with a half-life of about 5700 years.[7] Carbon is one of the few elements known to man since antiquity.[8][9] The name "carbon" comes from Latin language carbo, coal, and, in some Romance languages, the word carbon can refer both to the element and to coal.
There are several allotropes of carbon of which the best known are graphite, diamond, and amorphous carbon.[10] The physical properties of carbon vary widely with the allotropic form. For example, diamond is highly transparent, while graphite is opaque and black. Diamond is among the hardest materials known, while graphite is soft enough to form a streak on paper. Diamond has a very low electric conductivity, while graphite is a very good conductor. Also, diamond has the highest thermal conductivity of all known materials under normal conditions. All the allotropic forms are solids under normal conditions but graphite is the most thermodynamically stable.
All forms of carbon are highly stable, requiring high temperature to react even with oxygen. The most common oxidation state of carbon in inorganic compounds is +4, while +2 is found in carbon monoxide and other transition metal carbonyl complexes. The largest sources of inorganic carbon are limestones, dolomites and carbon dioxide, but significant quantities occur in organic deposits of coal, peat, oil and methane clathrates. Carbon forms more compounds than any other element, with almost ten million pure organic compounds described to date, which in turn are a tiny fraction of such compounds that are theoretically possible under standard conditions.[11]
Carbon is the fourth most abundant element in the universe by mass after hydrogen, helium, and oxygen. It is present in all known lifeforms, and in the human body, carbon is the second most abundant element by mass (about 18.5%) after oxygen.[12] This abundance, together with the unique diversity of organic compounds and their unusual polymer-forming ability at the temperatures commonly encountered on Earth, make this element the chemical basis of all known life.
Contents
[hide]
* 1 Characteristics
o 1.1 Allotropes
o 1.2 Occurrence
o 1.3 Isotopes
o 1.4 Formation in stars
o 1.5 Carbon cycle
* 2 Compounds
o 2.1 Inorganic compounds
o 2.2 Organic compounds
* 3 History and etymology
o 3.1 Applications
* 4 Production
o 4.1 Graphite Production
* 5 Precautions
* 6 See also
* 7 References
* 8 External links
[edit] Characteristics
The different forms or allotropes of carbon (see below) include the hardest naturally occurring substance, diamond, and also one of the softest known substances, graphite. Moreover, it has an affinity for bonding with other small atoms, including other carbon atoms, and is capable of forming multiple stable covalent bonds with such atoms. As a result, carbon is known to form nearly ten million different compounds; the large majority of all chemical compounds.[11] Carbon also has the highest melting and sublimation point of all elements.[citation needed] At atmospheric pressure it has no actual melting point as its triple point is at 10 MPa (100 bar) so it sublimates above 4000 K.[citation needed] Carbon sublimes in a carbon arc which has a temperature of about 5800K. Thus, irrespective of its allotropic form, carbon remains solid at higher temperatures than the highest melting point metals such as tungsten or rhenium. Although thermodynamically prone to oxidation, carbon resists oxidation more effectively than elements such as iron and copper that are weaker reducing agents at room temperature.
Diamond and graphite are two allotropes of carbon: pure forms of the same element that differ in structure.
Diamond and graphite are two allotropes of carbon: pure forms of the same element that differ in structure.
Carbon compounds form the basis of all life on Earth and the carbon-nitrogen cycle provides some of the energy produced by the Sun and other stars. Although it forms an extraordinary variety of compounds, most forms of carbon are comparatively unreactive under normal conditions. At standard temperature and pressure, it resists all but the strongest oxidizers. It does not react with sulfuric acid, hydrochloric acid, chlorine or any alkalis. At elevated temperatures carbon reacts with oxygen to form carbon oxides, and will reduce such metal oxides as iron oxide to the metal. This exothermic reaction is used in the iron and steel industry to control the carbon content of steel:
Fe3O4 + 4C(s) → 3Fe(s) + 4CO(g)
with sulfur to form carbon disulfide and with steam in the coal-gas reaction
C(s) + H2O(g) → CO(g) + H2(g).
Carbon combines with some metals at high temperatures to form metallic carbides, such as the iron carbide cementite in steel, and tungsten carbide, widely used as an abrasive and for making hard tips for cutting tools.
The system of carbon allotropes spans a range of extremes:
Synthetic diamond nanorods are the hardest materials known. Graphite is one of the softest materials known.
Diamond is the ultimate abrasive. Graphite is a very good lubricant.
Diamond is an excellent electrical insulator. Graphite is a conductor of electricity.
Diamond is the best known thermal conductor Some forms of graphite are used for thermal insulation (i.e. firebreaks and heatshields)
Diamond is highly transparent. Graphite is opaque.
Diamond crystallizes in the cubic system. Graphite crystallizes in the hexagonal system.
Amorphous carbon is completely isotropic. Carbon nanotubes are among the most anisotropic materials ever produced.
[edit] Allotropes
Main article: Allotropes of carbon
Atomic carbon is a very short-lived species and therefore, carbon is stabilized in various multi-atomic structures with different molecular configurations called allotropes. The three relatively well-known allotropes of carbon are amorphous carbon, graphite, and diamond. Once considered exotic, fullerenes are nowadays commonly synthesized and used in research; they include buckyballs,[13][14] carbon nanotubes,[15] carbon nanobuds[16] and nanofibers[17].[18] Several other exotic allotropes have also been discovered, such as aggregated diamond nanorods,[19] lonsdaleite,[20] glassy carbon,[21] carbon nanofoam[22] and linear acetylenic carbon.[23]
* The amorphous form, is an assortment of carbon atoms in a non-crystalline, irregular, glassy state, which is essentially graphite but not held in a crystalline macrostructure. It is present as a powder, and is the main constituent of substances such as charcoal, lampblack (soot) and activated carbon.
* At normal pressures carbon takes the form of graphite, in which each atom is bonded trigonally to three others in a plane composed of fused hexagonal rings, just like those in aromatic hydrocarbons. The resulting network is 2-dimensional, and the resulting flat sheets are stacked and loosely bonded through weak Van der Waals forces. This gives graphite its softness and its cleaving properties (the sheets slip easily past one another). Because of the delocalization of one of the outer electrons of each atom to form a π-cloud, graphite conducts electricity, but only in the plane of each covalently bonded sheet. This results in a lower bulk electrical conductivity for carbon than for most metals. The delocalization also accounts for the energetic stability of graphite over diamond at room temperature.
Some allotropes of carbon: a) diamond; b) graphite; c) lonsdaleite; d-f) fullerenes (C60, C540, C70); g) amorphous carbon; h) carbon nanotube.
Some allotropes of carbon: a) diamond; b) graphite; c) lonsdaleite; d-f) fullerenes (C60, C540, C70); g) amorphous carbon; h) carbon nanotube.
* At very high pressures carbon forms the more compact allotrope diamond, having nearly twice the density of graphite. Here, each atom is bonded tetrahedrally to four others, thus making a 3-dimensional network of puckered six-membered rings of atoms. Diamond has the same cubic structure as silicon and germanium and, thanks to the strength of the carbon-carbon bonds is the hardest naturally occurring substance in terms of resistance to scratching. Contrary to the popular belief that "diamonds are forever", they are in fact thermodynamically unstable under normal conditions and transform into graphite.[10] But due to a high activation energy barrier, the transition into graphite is so extremely slow at room temperature as to be unnoticeable.
* Under some conditions, carbon crystallizes as lonsdaleite. This form is similar to diamond but has a hexagonal crystal lattice.[20]
* Fullerenes have a graphite-like structure, but instead of purely hexagonal packing, they also contain pentagons (or even heptagons) of carbon atoms, which bend the sheet into spheres, ellipses or cylinders. The properties of fullerenes (split into buckyballs, buckytubes and nanobuds) have not yet been fully analyzed and represents an intense area of research in nanomaterials. The name "fullerene" is given after Richard Buckminster Fuller, developer of some geodesic domes,[citation needed] which resemble the structure of fullerenes. The buckyballs are fairly large molecules formed completely of carbon bonded trigonally, forming spheroids (the best-known and simplest is the soccerball-shaped structure C60 buckminsterfullerene).[13] Carbon nanotubes are structurally similar to buckyballs, except that each atom is bonded trigonally in a curved sheet that forms a hollow cylinder.[14][15] Nanobuds were first published in 2007 and are hybrid bucky tube/buckyball materials (buckyballs are covalently bonded to the outer wall of a nanotube) that combine the properties of both in a single structure.[16]
* Of the other discovered allotropes, aggregated diamond nanorods were synthesised in 2005 and are believed to be the hardest substance known yet.[24] Carbon nanofoam is a ferromagnetic allotrope discovered in 1997. It consists of a low-density cluster-assembly of carbon atoms strung together in a loose three-dimensional web, in which the atoms are bonded trigonally in six- and seven-membered rings. It is among the lightest known solids, with a density of about 2 kg/m³.[25] Similarly, glassy carbon contains a high proportion of closed porosity.[21] But unlike normal graphite, the graphitic layers are not stacked like pages in a book, but have a more random arrangement. Linear acetylenic carbon[23] has the chemical structure[26] -(C:::C)n- .Carbon in this modification is linear with sp orbital hybridisation, and is a polymer with alternating single and triple bonds. This type of carbyne is of considerable interest to nanotechnology as its Young's modulus is forty times that of the hardest known material - diamond.[27]
[edit] Occurrence
Graphite ore
Graphite ore
Raw diamond crystal.
Raw diamond crystal.
Carbon is the fourth most abundant chemical element in the universe by mass after hydrogen, helium, and oxygen. Carbon is abundant in the Sun, stars, comets, and in the atmospheres of most planets. Some meteorites contain microscopic diamonds that were formed when the solar system was still a protoplanetary disk. Microscopic diamonds may also be formed by the intense pressure and high temperature at the sites of meteorite impacts.[28]
"Present day" (1990s) sea surface dissolved inorganic carbon concentration (from the GLODAP climatology)
"Present day" (1990s) sea surface dissolved inorganic carbon concentration (from the GLODAP climatology)
In combination with oxygen in carbon dioxide, carbon is found in the Earth's atmosphere (in quantities of approximately 810 gigatonnes) and dissolved in all water bodies (approximately 36000 gigatonnes). Around 1900 gigatonnes are present in the biosphere. Hydrocarbons (such as coal, petroleum, and natural gas) contain carbon as well — coal "reserves" (not "resources") amount to around 900 gigatonnes, and oil reserves around 150 gigatonnes. With smaller amounts of calcium, magnesium, and iron, carbon is a major component of very large masses carbonate rock (limestone, dolomite, marble etc.).
Coal is a significant commercial source of mineral carbon; anthracite containing 92-98% carbon[citation needed] and the largest source (4000 Gt, or 80% of coal, gas and oil reserves) of carbon in a form suitable for use as fuel.[29]
Graphite is found in large quantities in New York and Texas, the United States, Russia, Mexico, Greenland, and India.
Natural diamonds occur in the rock kimberlite, found in ancient volcanic "necks," or "pipes". Most diamond deposits are in Africa, notably in South Africa, Namibia, Botswana, the Republic of the Congo, and Sierra Leone. There are also deposits in Arkansas, Canada, the Russian Arctic, Brazil and in Northern and Western Australia.
Diamonds are now also being recovered from the ocean floor off the Cape of Good Hope. However, though diamonds are found naturally, about 30% of all industrial diamonds used in the U.S. are now made synthetically.
According to studies from the Massachusetts Institute of Technology, an estimate of the global carbon budget is:[citation needed]
Biosphere, oceans, atmosphere
0.45 x 1018 kilograms (3.7 x 1018 moles)
Crust
Organic carbon 13.2 x 1018 kg
Carbonates 62.4 x 1018 kg
Mantle
1200 x 1018 kg
Carbon-14 is formed in upper layers of the troposphere and the stratosphere, at altitudes of 9–15 km, by a reaction that is precipitated by cosmic rays. Thermal neutrons are produced that collide with the nuclei of nitrogen-14, forming carbon-14 and a proton.
[edit] Isotopes
Main article: Isotopes of carbon
Isotopes of carbon are atomic nuclei that contain six protons plus a number of neutrons (varying from 2 to 16). Carbon has two stable, naturally occurring isotopes.[7] The isotope carbon-12 (12C) forms 98.93% of the carbon on Earth, while carbon-13 (13C) forms the remaining 1.07%.[7] The concentration of 12C is further increased in biological materials because biochemical reactions discriminate against 13C.[30] In 1961 the International Union of Pure and Applied Chemistry (IUPAC) adopted the isotope carbon-12 as the basis for atomic weights.[31] Identification of carbon in NMR experiments is done with the isotope 13C.
Carbon-14 (14C) is a naturally occurring radioisotope which occurs in trace amounts on Earth of up to 1 part per trillion (0.0000000001%), mostly confined to the atmosphere and superficial deposits, particularly of peat and other organic materials.[32] This isotope decays by 0.158 MeV β- emission. Because of its relatively short half-life of 5730 years, 14C is virtually absent in ancient rocks, but is created in the upper atmosphere (lower stratosphere and upper troposphere) by interaction of nitrogen with cosmic rays.[33] The abundance of 14C in the atmosphere and in living organisms is almost constant, but decreases predictably in their bodies after death. This principle is used in radiocarbon dating, invented in 1949, which has been used extensively to determine the age of carbonaceous materials with ages up to about 40,000 years.[34][35]
There are 15 known isotopes of carbon and the shortest-lived of these is 8C which decays through proton emission and alpha decay and has a half-life of 1.98739x10-21 s.[36] The exotic 19C exhibits a nuclear halo, which means its radius is appreciably larger than would be expected if the nucleus was a sphere of constant density.[37]
[edit] Formation in stars
Main articles: Triple-alpha process and CNO cycle
Formation of the carbon atomic nucleus requires a nearly simultaneous triple collision of alpha particles (helium nuclei) within the core of a giant or supergiant star. This happens in conditions of temperature and helium concentration that the rapid expansion and cooling of the early universe prohibited, and therefore no significant carbon was created during the Big Bang. Instead, the interiors of stars in the horizontal branch transform three helium nuclei into carbon by means of this triple-alpha process. In order to be available for formation of life as we know it, this carbon must then later be scattered into space as dust, in supernova explosions, as part of the material which later forms second- and third-generation star systems which have planets accreted from such dust. The Solar System is one such third-generation star system.
One of the fusion mechanisms powering stars is the carbon-nitrogen cycle.
Rotational transitions of various isotopic forms of carbon monoxide (e.g. 12CO, 13CO, and C18O) are detectable in the submillimeter regime, and are used in the study of newly forming stars in molecular clouds.
[edit] Carbon cycle
Main article: Carbon cycle
Diagram of the carbon cycle. The black numbers indicate how much carbon is stored in various reservoirs, in billions of tons ("GtC" stands for gigatons of carbon; figures are circa 2004). The purple numbers indicate how much carbon moves between reservoirs each year. The sediments, as defined in this diagram, do not include the ~70 million GtC of carbonate rock and kerogen.
Diagram of the carbon cycle. The black numbers indicate how much carbon is stored in various reservoirs, in billions of tons ("GtC" stands for gigatons of carbon; figures are circa 2004). The purple numbers indicate how much carbon moves between reservoirs each year. The sediments, as defined in this diagram, do not include the ~70 million GtC of carbonate rock and kerogen.
Under terrestrial conditions, conversion of one element to another is very rare. Therefore, the amount of carbon on Earth is effectively constant. Thus, processes that use carbon must obtain it somewhere and dispose of it somewhere else. The paths that carbon follows in the environment make up the carbon cycle. For example, plants draw carbon dioxide out of their environment and use it to build biomass, as in carbon respiration or the Calvin cycle, a process of carbon fixation. Some of this biomass is eaten by animals, whereas some carbon is exhaled by animals as carbon dioxide. The carbon cycle is considerably more complicated than this short loop; for example, some carbon dioxide is dissolved in the oceans; dead plant or animal matter may become petroleum or coal, which can burn with the release of carbon, should bacteria not consume it.
[edit] Compounds
[edit] Inorganic compounds
Main article: Compounds of carbon
Commonly carbon-containing compounds which are associated with minerals or which do not contain hydrogen or fluorine, are treated separately from classical organic compounds; however the definition is not rigid (see reference articles above). Among these are the simple oxides of carbon. The most prominent oxide is carbon dioxide (CO2). This was once the principal constituent of the paleoatmosphere, but is a minor component of the Earth's atmosphere today.[38] Dissolved in water, it forms carbonic acid (H2CO3), but as most compounds with multiple single-bonded oxygens on a single carbon it is unstable.[citation needed] Through this intermediate, though, resonance-stabilized carbonate ions are produced. Some important minerals are carbonates, notably calcite. Carbon disulfide (CS2) is similar.
The other common oxide is carbon monoxide (CO). It is formed by incomplete combustion, and is a colorless, odorless gas. The molecules each contain a triple bond and are fairly polar, resulting in a tendency to bind permanently to hemoglobin molecules, displacing oxygen, which has a lower binding affinity.[39][40] Cyanide (CN–), has a similar structure, but behaves much like a halide ion (pseudohalogen). For example it can form the nitride cyanogen molecule ((CN)2), similar to diatomic halides. Other uncommon oxides are carbon suboxide (C3O2),[41] the unstable dicarbon monoxide (C2O),[42][43] and even carbon trioxide (CO3).[44][45]
With reactive metals, such as tungsten, carbon forms either carbides (C4–), or acetylides (C22–) to form alloys with high melting points. These anions are also associated with methane and acetylene, both very weak acids. With an electronegativity of 2.5,[46] carbon prefers to form covalent bonds. A few carbides are covalent lattices, like carborundum (SiC), which resembles diamond.
[edit] Organic compounds
Main article: Organic compound
Structural formula of methane, the simplest possible organic compound
Structural formula of methane, the simplest possible organic compound
Carbon has the ability to form very long chains interconnecting C-C bonds. This property is called catenation. Carbon-carbon bonds are strong, and stable.[citation needed] This property allows carbon to form an almost infinite number of compounds; in fact, there are more known carbon-containing compounds than all the compounds of the other chemical elements combined except those of hydrogen (because almost all organic compounds contain hydrogen too).
The simplest form of an organic molecule is the hydrocarbon—a large family of organic molecules that are composed of hydrogen atoms bonded to a chain of carbon atoms. Chain length, side chains and functional groups all affect the properties of organic molecules. By IUPAC's definition, all the other organic compounds are functionalized compounds of hydrocarbons.[citation needed]
Carbon is the basis for all plastic materials that are used in common household items.
Carbon is the basis for all plastic materials that are used in common household items.
Carbon occurs in all organic life and is the basis of organic chemistry. When united with hydrogen, it forms various flammable compounds called hydrocarbons which are important to industry as chemical feedstock for the manufacture of plastics, petrochemicals and as fossil fuels.
When combined with oxygen and hydrogen, carbon can form many groups of important biological compounds including sugars, celluloses, lignans, chitins, alcohols, fats, and aromatic esters, carotenoids and terpenes. With nitrogen it forms alkaloids, and with the addition of sulfur also it forms antibiotics, amino acids and proteins. With the addition of phosphorus to these other elements, it forms DNA and RNA, the chemical codes of life, and adenosine triphosphate (ATP), the most important energy-transfer molecules in all living cells.
[edit] History and etymology
Please help improve this section by expanding it. Further information might be found on the talk page or at requests for expansion. (January 2008)
The English name carbon comes from the Latin carbo for coal and charcoal,[47] and hence comes French charbon, meaning charcoal. In German, Dutch and Danish, the names for carbon are Kohlenstoff, koolstof and kulstof respectively, all literally meaning coal-substance.
Carl Wilhelm Scheele
Carl Wilhelm Scheele
Antoine Lavoisier in his youth
Antoine Lavoisier in his youth
Carbon was discovered in prehistory and was known in the forms of soot and charcoal to the earliest human civilizations. Diamonds were known probably as early as 2500 BCE in China, while carbon in the forms of charcoal was made around Roman times by the same chemistry as it is today, by heating wood in a pyramid covered with clay to exclude air.[48][49]
In 1722, René A. F. de Réaumur demonstrated that iron was transformed into steel through the absorption of some substance, now known to be carbon.[50] In 1772, Antoine Lavoisier showed that diamonds are a form of carbon, when he burned samples of carbon and diamond then showed that neither produced any water and that both released the same amount of carbon dioxide per gram. Carl Wilhelm Scheele showed that graphite, which had been thought of as a form of lead, was instead a type of carbon.[51] In 1786, the French scientists Claude Louis Berthollet, Gaspard Monge and C. A. Vandermonde then showed that this substance was carbon.[52] In their publication they proposed the name carbone (Latin carbonum) for this element. Antoine Lavoisier listed carbon as an element in his 1789 textbook.[53]
A new allotrope of carbon, fullerene, that was discovered in 1985[54] includes nanostructured forms such as buckyballs and nanotubes.[13] Their discoverers received the Nobel Prize in Chemistry in 1996.[55] The resulting renewed interest in new forms, lead to the discovery of further exotic allotropes, including glassy carbon, and the realization that "amorphous carbon" is not strictly amorphous.[21]
[edit] Applications
Pencil lead for mechanical pencils are made of graphite.
Pencil lead for mechanical pencils are made of graphite.
Sticks of vine and compressed charcoal.
Sticks of vine and compressed charcoal.
A cloth of woven carbon filaments
A cloth of woven carbon filaments
Silicon carbide single crystal
Silicon carbide single crystal
The C60 fullerene in crystalline form
The C60 fullerene in crystalline form
Tungsten carbide milling bits
Tungsten carbide milling bits
Carbon is essential to all known living systems, and without it life as we know it could not exist (see alternative biochemistry). The major economic use of carbon other than food and wood is in the form of hydrocarbons, most notably the fossil fuel methane gas and crude oil (petroleum). Crude oil is used by the petrochemical industry to produce, amongst others, gasoline and kerosene, through a distillation process, in refineries. Cellulose is a natural, carbon-containing polymer produced by plants in the form of cellulose, cotton, linen, hemp. Commercially valuable carbon polymers of animal origin include wool, cashmere and silk. Plastics are made from synthetic carbon polymers, often with oxygen and nitrogen atoms included at regular intervals in the main polymer chain. The raw materials for many of these synthetic substances come from crude oil.
The uses of carbon and its compounds are extremely varied. It can form alloys with iron, of which the most common is carbon steel. Graphite is combined with clays to form the 'lead' used in pencils used for writing and drawing. It is also used as a lubricant and a pigment, as a moulding material in glass manufacture, in electrodes for dry batteries and in electroplating and electroforming, in brushes for electric motors and as a neutron moderator in nuclear reactors.
Charcoal is used as a drawing material in artwork, for grilling, and in many other uses including iron smelting. Wood, coal and oil are used as fuel for production of energy and space heating. Gem quality diamond is used in jewelry, and Industrial diamonds are used in drilling, cutting and polishing tools for machining metals and stone. Plastics are made from fossil hydrocarbons, and carbon fibre, made by pyrolysis of synthetic polyester fibres is used to reinforce plastics to form advanced, lightweight composite materials. Carbon fiber is made by pyrolysis of extruded and stretched filaments of polyacrylonitrile (PAN) and other organic substances. The crystallographic structure and mechanical properties of the fiber depend on the type of starting material, and on the subsequent processing. Carbon fibres made from PAN have structure resembling narrow filaments of graphite, but thermal processing may re-order the structure into a continuous rolled sheet[citation needed]. The result is fibers with higher specific tensile strength than steel.[citation needed]
Carbon black is used as the black pigment in printing ink, artist's oil paint and water colours, carbon paper, automotive finishes, India ink and laser printer toner. Carbon black is also used as a filler in rubber products such as tyres and in plastic compounds. Activated charcoal is used as an absorbent and adsorbent in filter material in applications as diverse as gas masks, water purification and kitchen extractor hoods and in medicine to absorb toxins, poisons, or gases from the digestive system. Carbon is used in chemical reduction at high temperatures. coke is used to reduce iron ore into iron. Case hardening of steel is achieved by heating finished steel components in carbon powder. Carbides of silicon, tungsten, boron and titanium, are among the hardest known materials, and are used as abrasives in cutting and grinding tools. Carbon compounds make up most of the materials used in clothing, such as natural and synthetic textiles and leather, and almost all of the interior surfaces in the built environment other than glass, stone and metal.
[edit] Production
Please help improve this section by expanding it. Further information might be found on the talk page or at requests for expansion. (December 2007)
[edit] Graphite Production
Commercially viable natural deposits of graphite occur in many parts of the world, but the most important sources economically are in China, India, Brazil, and North Korea.[56] Graphite deposits are of metamorphic origin, found in association with quartz, mica and feldspars in schists, gneisses and metamorphosed sandstones and limestone as lenses or veins, sometimes of a metre or more in thickness. Deposits of graphite in Borrowdale, Cumberland, England were at first of sufficient size and purity that, until the 1800s, pencils were made simply by sawing blocks of natural graphite into strips before encasing the strips in wood. Today, smaller deposits of graphite are obtained by crushing the parent rock and floating the lighter graphite out on water.
[edit] Precautions
Pure carbon has extremely low toxicity and can be handled and even ingested safely in the form of graphite or charcoal. It is resistant to dissolution or chemical attack, even in the acidic contents of the digestive tract, for example. Consequently if it gets into body tissues it is likely to remain there indefinitely. Carbon black was probably one of the first pigments to be used for tattooing, and Ötzi the Iceman was found to have carbon tattoos that survived during his life and for 5200 years after his death.[57] However, inhalation of coal dust or soot (carbon black) in large quantities can be dangerous, irritating lung tissues and causing the congestive lung disease coalworker's pneumoconiosis. Similarly, diamond dust used as an abrasive can do harm if ingested or inhaled. Microparticles of carbon are produced in diesel engine exhaust fumes, and may accumulate in the lungs.[58] In these examples, the harmful effects may result from contamination of the carbon particles, with organic chemicals or heavy metals for example, rather than from the carbon itself.
Carbon may also burn vigorously and brightly in the presence of air at high temperatures, as in the Windscale fire, which was caused by sudden release of stored Wigner energy in the graphite core. Large accumulations of coal, which have remained inert for hundred of millions of years in the absence of oxygen, may spontaneously combust when exposed to air, for example in coal mine waste tips.
The great variety of carbon compounds include such lethal poisons as tetrodotoxin, the lectin ricin from seeds of the castor oil plant Ricinus communis, cyanide (CN-) and carbon monoxide; and such essentials to life as glucose and protein.Compounds of carbon
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There is an immense number of distinct compounds that contain carbon atoms. Some sources suggest that this number is close to almost ten million known.[1] However, it is possible that the number is greater.
Contents
[hide]
* 1 Organic compounds
* 2 Inorganic compounds
o 2.1 Compounds with other nonmetals
o 2.2 Compounds with metals
+ 2.2.1 Carbonates and bicarbonates
+ 2.2.2 Carbonyls
+ 2.2.3 Compounds contanining the CN group
+ 2.2.4 Carbides
+ 2.2.5 Other
* 3 Alloys
* 4 Formation of carbon compounds
* 5 References
* 6 See also
[edit] Organic compounds
Main article: Organic compound
Every organic compound contains at least one atom of carbon. The number of these compounds is immense and the described number of defined compounds is close to 10 million. However, an indefinitely larger number of such compounds are theoretically possible.
There are several organic compounds sometimes considered as inorganic: NH2COONH4, COCl2, CSCl2, CS(NH2)2, CO(NH2)2
[edit] Inorganic compounds
See also: Inorganic compounds by element#Carbon
There is a rich variety of carbon chemistry that does not fall within the realm of organic chemistry and is thus called inorganic carbon chemistry.
[edit] Compounds with other nonmetals
Perhaps the best known are the oxides of carbon, carbon dioxide (CO2) and carbon monoxide (CO). Other known oxides are the uncommon carbon suboxide, C3O2, the uncommon dicarbon monoxide, C2O and even the exotic carbon trioxide (CO3).
Other (binary) compounds of carbon with nonmetals include: CS2, β-C3N4, CBr4, CCl4, CF4, COF2, COS, H2C2B10H10,
[edit] Compounds with metals
[edit] Carbonates and bicarbonates
Main articles: Carbonic acid, Carbonate, and Bicarbonate
The only known acid that is derived from the oxides of carbon is the carbonic acid (H2CO3). Upon monodeprotonation of this acid, bicarbonates are formed, which can be further derpotonated to carbonates.
Here is a list of carbonates and bicarbonates: NH4HCO3, (NH4)2CO3, BaCO3, CdCO3, Cs2CO3, Ca(HCO3)3, CaCO3, Ce2(CO3)3, CoCO3, CuCO3, FeCO3, PbCO3, La2(CO3)3, Li2CO3, MgCO3, MnCO3, NiCO3, KHCO3, K2CO3, Ag2CO3, NaHCO3, Na2CO3, SrCO3, ZnCO3
[edit] Carbonyls
Main article: Carbonyl
Carbonyls are coordination complexes between transition metals and carbonyl ligands. Metal carbonyls are complexes that are formed with the neutral ligand CO. These complexes are covalent. Here is a list of some carbonyls: Cr(CO)6, Co2(CO)8, Fe(CO)5, Mn2(CO)10, Mo(CO)6, Ni(CO)4, W(CO)6,
[edit] Compounds contanining the CN group
Main articles: Cyanide, Cyanates, Thiocyanate, and Isocyanate
Other types of inorganic compounds include inorganic salts and complexes of the carbon-containing polyatomic ions cyanide, isocyanide, cyanate, thiocyanate.
NH4SCN, CaNCN, Co(SCN)2, CuCN, (HCNO)x NH2CN HCNO, (CN)2, BrCN, ClCN, HCN, KOCN, KCN, K3Fe(CN)6, K4Fe(CN)6, KSCN, Fe4(Fe(CN)6)3, AgCN, NaOCN, NaCN, Na3Fe(CN)5NO, NaSCN, (SCN)2,
[edit] Carbides
Main article: Carbide
Carbides are binary compounds of carbon with an element that is less electronegative than it. B4C, CaC2 SiC, TaC, TiC, WC,
[edit] Other
The known inorganic chemistry of the allotropes of carbon (diamond, graphite, and the fullerenes) blossomed with the discovery of buckminsterfullerene in 1985, as additional fullerenes and their various derivatives were discovered. One such class of derivatives is inclusion compounds, in which an ion is enclosed by the all-carbon shell of the fullerene. This inclusion is denoted by the "@" symbol. For example, an ion consisting of a lithium ion trapped within buckminsterfullerene would be denoted Li+@C60. As with any other ionic compound, this complex ion could in principle pair with a counterion to form a salt.
[edit] Alloys
There are several alloys that contain carbon of which the best known alloy is carbon steel (see category:steels)). Besides steel, other alloys based on iron and carbon are: anthracite iron, cast iron, pig iron, wrought iron, but also spiegeleisen (which contains also manganese). Stellite is an alloy of carbon with cobalt, chromium and tungsten. To some degree, these alloys could be considered carbides.
[edit] Formation of carbon compounds
In organic chemistry there are 3 important elements: Carbon, Oxygen and Hydrogen. Each of these elements have different kinds of bonds. Carbon atom has tetravalent bonds, Oxygen atoms divalent bonds and Hydrogen monovalent bonds.Organic compound
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Methane is the simplest possible organic compound
Methane is the simplest possible organic compound
An organic compound is any member of a large class of chemical compounds whose molecules contain carbon. For historical reasons discussed below, a few types of compounds such as carbonates, simple oxides of carbon and cyanides, as well as the allotropes of carbon, are considered inorganic. The division between "organic" and "inorganic" carbon compounds while "useful in organizing the vast subject of chemistry...is somewhat arbitrary"[1].
Organic chemistry is the science concerned with all aspects of organic compounds. Organic synthesis is the methodology of their preparation.
Contents
[hide]
* 1 History
* 2 Classification
o 2.1 Natural compounds
o 2.2 Synthetic compounds
* 3 Nomenclature
* 4 Databases
* 5 Structure determination
* 6 See also
* 7 References
[edit] History
The "organic" is a historical name, dating back to 19th century, when it was believed that organic compounds could only be synthesized in living organisms through vis vitalis - the "life-force". The theory that organic compounds were fundamentally different from those that were "inorganic", that is, not synthesized through a life-force, was disproved with the synthesis of urea, an "organic" compound by definition of its known occurrence only in the urine of living organisms, from potassium cyanate and ammonium sulfate by Friedrich Wöhler in the Wöhler synthesis. The kinds of carbon compounds that are still traditionally considered inorganic are those that were considered inorganic before Wöhler's time; that is, those which came from "inorganic" (i.e., lifeless) sources such as minerals.[1]
[edit] Classification
See Organic chemistry#Classification of organic substances
Organic compounds may contain atoms of further elements, so-called heteroatoms. Organometallic compounds constitute a further subsection, characterized by covalent bonds between organic carbon and a metal.
There is also a large number of inorganic carbon compounds to distinguish from organic compounds.
[edit] Natural compounds
An important subset of organic compounds is still extracted from natural sources because they would be far too expensive to be produced artificially. Examples include most sugars, some alkaloids and terpenoids, certain nutrients such as vitamin B12, and in general, those natural products with large or stereoisometrically complicated molecules which are present in reasonable concentrations in living organisms.
Further compounds of prime importance in biochemistry are antigens, carbohydrates, enzymes, hormones, lipids and fatty acids, neurotransmitters, nucleic acids, proteins, peptides and amino acids, vitamins and fats and oils.
[edit] Synthetic compounds
Many polymers, including all plastics are organic compounds.
[edit] Nomenclature
The IUPAC nomenclature of organic compounds slightly differs from the CAS nomenclature.
[edit] Databases
* The CAS database is the most comprehensive repository for data on organic compounds. The search tool SciFinder is offered .
* The Beilstein database contains information on 9.8 million substances, covers the scientific literature from 1771 to the present, and is today accessible via CrossFire. Structures and a large diversity of physical and chemical properties is available for each substance, with reference to original literature.
* PubChem contains 18.4 million entries on compounds and especially covers the field of medicinal chemistry.
There is a great number of more specialized databases for diverse branches of organic chemistry.
[edit] Structure determination
See Structure determination
Today, the main tools are proton and carbon-13 NMR spectroscopy and X-ray crystallography.List of organic compounds
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This page aims to list well-known organic compounds, including organometallic compounds, to stimulate the creation of Wikipedia articles. Note that purely inorganic compounds, minerals, and chemical elements are not included on this list. There are also no generic terms (e.g., carbohydrate) or mixtures of no fixed composition (e.g., naphtha, gasoline). Compounds and enzymes that are overwhelmingly of interest to biochemists, such as Cytochrome c peroxidase, are listed under list of biomolecules.
This list is not necessarily complete or up to date — if you see an article that should be here but isn't (or one that shouldn't be here but is), please update the page accordingly.
For substances with a number prefix such as 2-Butanol or 1,3-Cyclohexadiene, please use the first letter of the name (in this case under B or C) to find the compound. Note that such names usually have the first letter capitalized in a title or at the beginning of a sentence.
Relevant links for chemical compounds are:
* The CAS Substance Databases, which contains information on about 23 million compounds
* ChemIDplus [1] is a useful non-commercial source for chemical lookups
* NIST Chemistry WebBook [2] is a freely available resource compiled by National Institute of Standards and Technology under the Standard Reference Data Program. Apart from chemical structures, it contains a wealth of associated physico-chemical information such as thermochemistry data and spectra
* ChEBI [3], a freely available dictionary of molecular entities focused on ‘small’ chemical compounds
* PubChem [4], maintained by the National Center for Biotechnology Information (NCBI), serves as a repository of chemical compounds from many public and commercial resources
* http://physchem.ox.ac.uk/MSDS/ Material Safety Data Sheets, plus other relevant links
These (commercial) links may also provide useful information:
* Chemfinder [5] is helpful for finding information about a chemical (disable and delete cookies!)
* Sigma Aldrich [6]
* Acros Organics [7]
* Lancaster [8]
* Chemical Suppliers Directory [9]
* ChemSpider [10]has over 20 million structures with chemical names and the ability to download the molfile locally. It includes links to chemical vendors, PubChem, ChEBI and over 100 other sources and is curated by users.
Whilst most compounds are referred to by their IUPAC name, "traditional" names have also been kept where they are in wide use or of significant historical interest.
See also: organic compound, list of compounds, list of inorganic compounds, inorganic compounds by element, list of biomolecules, polyatomic ions, list of elements by name, list of alchemical substances, list of drugs, list of reactions.
Table of contents: A B C D E F G H I J K L M N O P Q R S T U V W X Y Z
[edit] A
For substances with an A- or α- prefix such as α-Terpinene, please see the parent page (in this case Terpinene).
* Abietic acid - C20H30O2
* Acenaphthene
* Acenaphthoquinone
* Acenaphthylene
* Acepromazine
* Acetaldehyde — CH3CHO, also known as ethanal
* Acetamide
* Acetaminophen — C8H9NO2
* Acetaminosalol
* Acetamiprid
* Acetanilide
* Acetic acid — CH3COOH, also known as ethanoic acid, Glacial acetic acid or GAA
* Acetoguanamine
* Acetone — CH3COCH3, or (CH3)2CO
* Acetonitrile
* Acetophenone
* Acetylcholine – (CH3)3N+CH2CH2OCOCH3.
* Acetylene — C2H2
* N-Acetylglutamate
* Acetylsalicylic Acid also known as Aspirin
* Acid fuchsin
* Acridine — C13H9N
* Acridine orange
* Acrolein
* Acrylamide — C3H5NO
* Acrylic acid — CH2=CHCOOH
* Acrylonitrile
* Acryloyl chloride
* Acyclovir
* Adamantane
* Adenosine
* Adipamide
* Adipic acid
* Adiponitrile
* Adipoyl dichloride
* Adonitol
* Adrenaline, epinephrine
* Adrenochrome
* Aflatoxin
* Alanine
* Albumins
* Alcian blue
* Aldosterone
* Aldrin
* Aliquat 336
* Alizarin
* Allantoic acid
* Allantoin
* Allethrin
* Allyl propyl disulfide
* Allylamine
* Allyl chloride
* Amido black 10b
* p-Aminobenzoic acid (PABA)
* Aminodiacetic acid
* Aminoethylpiperazine
* 5-Amino-2-hydroxybenzoic acid
* Aminophylline
* 5-Aminosalicylic acid
* Aminothiazole
* Amiodarone
* Amiton
* Amobarbital
* Amoxicillin — C16H19N3O5S.3H2O
* Amphetamine
* Amyl nitrate
* Amyl nitrite — C5H11A.ONO
* Anethole
* Angelic acid
* Anilazine
* Aniline — C6H5-NH2
* Aniline hydrochloride
* Anisole
* Anisoyl chloride
* Anthanthrene
* Anthracene – (C6H4CH)2
* Anthramine
* Anthranilic acid
* Anthraquinone
* Anthrone
* Antipyrine
* Aprotinin
* Arabinose
* Arginine
* Aroclor (polychlorinated biphenyls)
* Ascorbic acid (vitamin C)
* Asparagine
* Asparagusic acid
* Aspartame
* Aspartic acid
* Asphidophytidine
* Astrablue
* Atrazine
* Auramine o
* Aureine
* Avobenzone
* Azadirachtin A — C35H44O16
* Azathioprine
* Azelaic acid
* Azinphos-methyl
* Aziridine
* Azithromycin
* Azo violet
* Azobenzene
* Azulene
* Azure a
[edit] B
For substances with a B- or β- prefix such as β-Pinene, please see the parent page (in this case Pinene).
* Bacillomycin
* Barbital
* Barbituric acid
* Behenic acid
* Benomyl
* Benzaldehyde
* Benzalkonium chloride
* Benzamide
* Benzanthrone
* Benzene — C6H6
* Benzethonium chloride
* Benzidine
* Benzil
* Benzilic acid
* Benzimidazole
* Benzisothiazolinone
* Benzisoxazole
* Benzo(a)anthracene
* Benzo(c)cinnoline
* Benzo(a)pyrene
* Benzo(c)phenanthrene
* Benzo(e)fluoranthene
* Benzo(e)pyrene
* Benzo(ghi)perylene
* Benzo(j)fluoranthene
* Benzo(k)fluoranthene
* Benzo(c)thiophene
* Benzocaine
* Benzofuran
* Benzoic acid
* Benzoin
* Benzothiazole
* Benzothiophene
* Benzotriazole
* Benzoxazole
* Benzoyl chloride
* Benzyl alcohol
* Benzyl chloroformate
* Benzylamine
* Benzyldimethylamine
* Benzylidene acetone
* Betaine
* Betulin
* Butylated hydroxytoluene (BHT) – C6H2(OH)(CH3)(C(CH3)3)2
* Biotin (Vitamin H)
* Biphenyl
* 2,2'-Bipyridyl = 2,2'-Bipyridine
* 1,8-Bis(dimethylamino)naphthalene (Proton-sponge, Aldrich trademark name)
* Bis(chloromethyl) ether
* Bismarck brown y
* Bisphenol A
* Biuret
* Borneol
* Brassinolide
* Brilliant cresyl blue
* Bromacil
* Bromoacetic acid
* Bromobenzene
* 2-Bromo-1-chloropropane
* Bromocresol purple
* Bromocyclohexane
* Bromoform
* Bromomethane — BrCH3
* Bromophenol blue
* 2-Bromopropane
* Bromothymol blue
* Bromotrifluoromethane
* Brucine
* Buckminsterfullerene
* Buspirone
* 1,3-Butadiene
* Butadiene resin
* Butane — C4H10
* Butene
* 2-Butoxyethanol
* Butylamine = n-Butylamine
* Butyllithium
* 2-Butyne-1,4-diol
* Butyraldehyde
* Butyrophenone
* Butyryl chloride
[edit] C
For substances with an c- or cis- prefix such as cis-3-hexenal, you may find these listed under the parent name letter (in this case "H"), as is the norm in chemical catalogues.
* Cacodylic acid
* Cacotheline
* Cadaverine — NH2(CH2)5NH2
* Cadinene
* Cafestol
* Caffeine
* Calcein
* Calciferol (Vitamin D)
* Calcitonin
* Calmodulin
* Calreticulin
* Camphene
* Camphor
* Cannabinol
* Caprolactam
* Caprolactone
* Capsaicin
* Captan
* Captopril
* Carbazole
* Carbazol-9-yl-methanol (N-(Hydroxymethyl)carbazole)
* Carbofuran
* Carbonyl fluoride
* Carboplatin
* Carboxypolymethylene
* Carminic acid
* Carnauba wax
* Carnitine
* Cartap
* Carvacrol
* Carvone
* Castor oil
* Catechol
* Cedar wood oil
* Cefazolin
* Cefotaxime
* Ceftriaxone
* Cellulose
* Cellulose acetate
* Cetrimide
* Cetyl alcohol
* Chloracetyl chloride
* Chloral
* Chloral hydrate
* Chlorambucil
* Chloramine-T
* Chloramphenicol
* Chloranilic acid
* Chlordane
* Chlorhexidine gluconate
* Chloro-m-cresol
* Chloroacetic acid
* 4-Chloroaniline (p-Chloroaniline)
* Chlorobenzene
* 2-Chlorobenzoic acid (o-Chlorobenzoic acid)
* Chlorodifluoromethane
* Chloroethene — C2H3Cl
* Chlorofluoromethane
* Chloroform — CHCl3
* Chloromethane
* 2-Chloro-2-methylpropane (tert butyl chloride)
* Chloronitroaniline
* Chloropentafluoroethane
* Chloropicrin
* Chloroprene
* Chloroquine
* Chlorostyrene
* Chlorothiazide
* Chlorotrifluoromethane
* Chlorotrimethylsilane
* Chloroxuron
* Chlorpyrifos
* Chlorthiamide
* Cholesterol
* Choline
* Chromotropic acid
* Cilostazol
* Cinchonine
* Cinnamaldehyde
* Cinnamic acid
* Cinnamyl alcohol
* Cinnoline — C4H4N2
* cis-2-butene
* cis-3-Hexenal
* cis-3-Hexen-1-ol
* Citral
* Citric acid — C3H4OH(COOH)3
* Citronella oil
* Citronellal
* Citrulline
* Clobetasone
* Clopidol
* Cloxacillin — C19H17ClN3O5S*Na*H2O
* Cobalamin (Vitamin B12)
* Cocamidopropyl
* Colchicine
* Collagen
* Collodion
* Congo red
* Coniine
* Coomassie blue
* Coronene
* Coumarin
* Creatine
* Cresol
* Cresyl violet
* Crotonaldehyde
* 18-Crown-6
* Crystal violet
* Cubane
* Cumene
* Cuneane
* Cupferron
* Cuscohygrine
* Cyanogen
* Cyanogen chloride
* Cyanoguanidine
* Cyanuric acid
* Cyanuric chloride
* Cyclodecane
* α-Cyclodextrin
* Cyclododecane
* Cycloheptatriene
* 1,3-Cyclohexadiene
* 1,4-Cyclohexadiene
* Cyclohexane
* Cyclohexanol
* Cyclohexanone
* Cyclohexene
* Cyclonite - (CH2-N-NO2)3
* Cyclooctatetraene
* Cyclopentadiene — C5H6
* Cyclopentane
* Cyclopentanol
* Cyclopentanone
* Cyclopentene
* Cypermethrin
* Cysteamine
* Cysteine
* Cystine
* Cytosine — C4H5N3O
[edit] D
For substances with a d- or D- prefix such as D-alanine or DL-alanine, please see the parent page (in this case alanine).
* DABCO
* DDT
* Decaborane
* Decabromodiphenyl ether
* Decahydronaphthalene
* Decane — C10H22
* Dehydroacetic acid
* Dehydrocholic acid
* Deltamethrin
* Demeton
* Denatonium
* Dexamethazone
* Dextran
* Dextrin
* 3,3'-Diaminobenzidine
* Di-t-butyl peroxide
* Diacetylene
* Diazinon
* Diazomethane
* 1,2-Dibromoethane
* Dibucaine hydrochloride
* Dichloroacetic acid
* p-Dichlorobenzene
* Dichlorobutane
* Dichlorodifluoromethane
* Dichlorodimethylsilane
* 1,2-Dichloroethane
* Dichlorofluoromethane
* Dichlorophen
* 2,4-Dichlorophenoxyacetic acid
* Dichlorotrifluoroethane
* Dichlorvos
* Diclofenac sodium
* Dicofol
* Dicrotophos
* Dicyclopentadiene
* Dieldrin
* Diethanolamine
* Diethion
* Diethyl aluminium chloride a Lewis acid
* Diethylamine
* Diethylene glycol
* Diethylenetriamine
* Diethyl ether
* Difluoromethane
* Digitonin
* Dihydrocortisone
* Diisoheptyl phthalate
* Diisopropyl ether
* Diketene
* Dimethicone
* Dimethylamine
* N,N-Dimethylacetamide
* N,N-dimethylaniline
* 1,2-Dimethylbenzene (o-Xylene)
* 1,3-Dimethylbenzene (m-Xylene)
* 1,4-Dimethylbenzene (p-Xylene)
* N,N-dimethylformamide
* Dimethyldiethoxysilane
* Dimethylglyoxime
* Dimethylmercury
* Dimethyl sulfoxide
* Dinoseb
* Dioctyl phthalate
* Dioxane
* Dioxathion
* Dioxin
* Diphenylacetylene (Tolane)
* Diphenylmethanol (Benzhydrol)
* Diquat
* Direct Blue 1
* Disulfiram
* Disulfoton
* Dithranol
* 2,6-Di-tert-butylphenol
* 2,6-Di-tert-butyl-4-methylphenol
* 2,6-Di-tert-butylpyridine
* Diuron
* Divinylbenzene
* Docosane
* Dodecane
* Dodecylbenzene
* Domperidone
* Dopamine
* Doxylamine succinate
[edit] E
* EDTA (Ethylenediamine-N,N,N',N'-tetraacetic acid)
* Eicosane
* Endosulfan
* Endrin
* Eosin
* Ephedrine
* Epibromohydrin
* Epinephrine — C9H13NO3
* Erucic acid — CH3(CH2)7CH=CH(CH2)11COOH
* Erythritol
* Estradiol
* Ethacridine lactate
* Ethane — C2H6
* 1,2-Ethanedithiol — C2H4(SH2)2
* Ethanol — CH3CH2OH
* Ethene — C2H4
* Ethidium bromide
* Ethyl acetate
* Ethylamine
* Ethyl 4-aminobenzoate (Ethyl p-aminobenzoate)
* Ethylbenzene
* Ethyl chloride
* Ethylene
* Ethylene glycol — OHCH2CH2OH
* Ethylene oxide
* Ethyl formate
* 2-Ethyl-1-hexanol
* Eugenol
[edit] F
* Farnesol
* Ferrocene
* Fipronil
* Flunixin
* Fluoranthene
* Fluorene
* 9-Fluorenone
* Fluorescein
* Fluorobenzene
* Fluoroethylene
* Fluoxetine
* Folic acid (Vitamin M)
* Fonofos
* Formaldehyde — HCHO
* Formamide
* Formanilide
* Formic acid — HCOOH
* Formoterol
* Fructose
* Fumaric acid
* Furan (furane)
* Furfural
* Furfuryl alcohol
* Furfurylamine
* Furylfuramide
[edit] G
* Gadopentetate - also known as Magnevist
* Galactose
* Gamma-aminobutyric acid
* Gamma-butyrolactone
* Gamma-hydroxybutyrate (GHB)
* Geraniol
* Gibberellic acid
* Gluconic acid
* Glucose — C6H12O6
* Glutamic acid (glutamate)
* Glutamine
* Glutaraldehyde
* Glutaric acid
* Glutathione
* Glyburide
* Glycerin (glycerol)
* Glycerol (glycerin)
* Glycerophosphoric acid
* Glycidol
* Glycine — NH2CH2COOH
* Glycogen
* Glycolic acid
* Glyoxal
* Guaiacol
* Guanidine
* Guanine
* Guanosine
[edit] H
* Halothane
* Hematoxylin
* HEPES
* Heptadecane
* Heptane — C7H16
* Hexabromocyclododecane
* Hexachloropropene
* Hexadecane
* Hexafluoro-2-propanol
* Hexafluoro-2-propanone
* Hexafluoroethane
* Hexafluoropropylene
* Hexamethyldewarbenzene
* Hexamethyldisilazane
* Hexamethylenimine
* Hexamethylolmelamine
* Hexamine – (CH2)6N4
* Hexane — C6H14
* Hexanitrodiphenylamine
* Hexanoic acid
* cis-3-Hexanal
* cis-3-Hexen-1-ol
* Hippuric acid
* Histidine — NH2CH(C4H5N2)COOH
* Histamine
* Homoarginine
* Homocysteine
* Homocystine
* Homotaurine
* Hydrochlorothiazide
* Hydrocinnamic acid
* Hydroquinone
* Hydroxyproline
* 5-Hydroxytryptamine
* Hygrine
[edit] I
* Ibuprofen
* Imazapyr
* Imidazole
* Imiquimod
* Indazole
* Indene
* Indigo
* Indole
* Indoline
* Indole-3-acetic acid
* Inositol
* Iodoxybenzene
* Ionone
* Ipratropium bromide
* Isatin
* Isoamyl isobutyrate
* Isobenzofuran
* Isoborneol
* Isobornyl acetate
* Isoflurane
* Isoindole
* Isoleucine
* Isomelamine
* Isooctanol
* Isophthalic acid
* Isopropanol – (CH3)2CHOH, also known as IPA, IsoPropyl alcohol
* Isoquinoline
* Isoxazole
* Itraconazole
[edit] J
* Jasmone
* Jenner's stain
[edit] K
* Kanamycin
* Kepone alcohol
* Keratin
* Ketene
* Kojic acid
[edit] L
For substances with an l- or L- prefix such as L-alanine or DL-alanine, please see the parent page (in this case alanine).
* Lactic acid — CH3CH(OH)COOH
* Lactose
* Lauric acid
* Lauryl alcohol
* LDA (Lithium diisopropylamide)
* Leucine
* Levulinic acid
* Limonene
* Linalool
* Linoleic acid
* Linolenic acid
* Lipoamide
* Lithium diisopropylamide
* Loratadine
* LSD
* Luminol
* 2,6-Lutidine
* Lycopene
* Lysine
[edit] M
For substances with an m- or meta- prefix such as m-cresol, meta-cresol or metacresol that are not listed below, please look for a more generic page (in this case cresol). For substances with a meso- prefix such as meso-tartaric acid or mesotartaric acid that are not listed below, please see the parent page (in this case tartaric acid).
* Malachite green
* Malathion
* Maleic anhydride
* Malic acid
* Maltose
* Mandelonitrile
* Mannide monooleate
* Mannose
* Mauveine
* MDMA
* Mecoprop
* MEK
* Melatonin
* Meldola's blue
* Meloxicam
* Menthol
* 2-Mercaptoethanol
* 2-Mercaptopyridine
* Merocyanine
* Mesityl oxide
* Mesitylene – (CH3)3-C6H3
* Mesotartaric acid
* Metaldehyde
* Metamizole (dipyrone)
* Methane — CH4
* Methanesulfonic acid
* Methanol — CH3OH
* Methionine
* Methomyl
* 4-Methoxybenzaldehyde (anisaldehyde)
* Methoxychlor
* Methoxyflurane
* Methyl acetate
* Methyl-2-cyanoacrylate
* Methyl ethyl ketone (MEK)
* Methyl isobutyl ketone (MIBK)
* Methyl isocyanate — CH3-N=C=O
* Methyl methacrylate
* Methyl tert-butyl ether (MTBE)
* Methylal
* Methylamine
* 2-Methylbenzoic acid (o-Toluic acid)
* 4-Methylbenzoic acid (p-Toluic acid)
* Methyl chloroformate
* Methylcyclohexane
* Methylene blue — C16H18ClN3S
* Methylhydrazine
* Methylmercury
* Methylmorpholine
* 2-Methylpropene (isobutylene)
* N-Methylpyrrolidone — C5H9NO
* Methyltriethoxysilane
* Methyltrimethoxysilane
* Metoprolol
* Metronidazole
* Michler's ketone
* Milrinone
* Monocrotophos
* Monosodium glutamate
* Mordant red 19
* Morpholine
* MTBE
* Murexide
* Mustard gas — C4H8Cl2S
* Myrcene
[edit] N
For substances with an n- or normal- prefix such as n-pentane that are not listed below, please see the parent page (in this case pentane).
For substances with an N- prefix (meaning on nitrogen) such as N,N-dimethylformamide, if these are not listed below please see the parent page (in this case dimethylformamide).
* n-Nonadecane
* n-Tetradecylbenzene
* Naphthalene — C10H8
* Naphthoquinone (Vitamin K)
* 2-Naphthylamine
* Neomycin
* Niacin or nicotinic acid (Vitamin B3)
* Nicotine
* Niflumic acid
* Nile red
* Nimesulide
* Nitrilotriacetic acid
* Nitrobenzene
* Nitrocellulose
* Nitroethane
* Nitrofen
* Nitrofurantoin
* Nitroglycerine — C3H5(NO2)3
* Nitromethane
* Nitrosobenzene
* N-Nitroso-N-methylurea
* Nitrosomethylurethane
* Nominine
* Nonacosane
* Nonane — C9H20
* Noradrenaline, norepinephrine
* Norephidrine
* Norcarane
* Norleucine
* Nujol
* NMN
[edit] O
For substances with an o- or ortho- prefix such as o-cresol, ortho-cresol or orthocresol that are not listed below, please look for a more generic page (in this case cresol).
* Octabromodiphenyl ether
* Octane — C8H18
* 1-Octanethiol
* Octanoic acid
* 4-Octylphenol
* Oleic acid
* Orcin
* Orcinol
* Ornithine
* Orotic acid
* Oxalic acid
* Oxalyl Chloride — C2O2Cl2
* Oxamide
* Oxazole
* Oxolinic acid
* Oxymetholone
[edit] P
For substances with an p- or para- prefix such as p-cresol, para-cresol or paracresol that are not listed below, please look for a more generic page (in this case cresol).
* p-nitro benzal dehyde
* PABA
* Paclitaxel
* Palmitic acid
* Pantothenic acid (Vitamin B5)
* Para red
* Parachlorometaxylenol (PCMX)
* Paraformaldehyde
* Parathion
* Pelargonic acid
* Pentabromodiphenyl ether
* Pentachlorobiphenyl
* Pentachlorophenol
* Pentadecane
* Pentaerythritol
* Pentaethylene glycol
* Pentafluoroethane
* Pentane — C5H12
* Pentetic acid
* Perfluorotributylamine
* Permethrin
* Peroxyacetic acid
* Perylene
* Petroleum ether
* Phenacetin
* Phenacyl bromide
* Phenanthrene
* Phenanthrenequinone
* Phencyclidine
* Phenethylamine
* Phenobarbital (c-iv)
* Phenol — C6H5OH
* Phenol red, sodium salt
* Phenolphthalein
* Phenothiazine
* Phenylacetic acid
* Phenylacetylene
* Phenylalanine
* p-Phenylenediamine
* Phenylhydrazine
* Phenylhydroxylamine
* Phenyllithium
* 4-Phenyl-4-(1-piperidinyl)cyclohexanol (PPC)
* Phenylthiocarbamide — C7H8N2S
* Phloroglucinol
* Phorate
* Phthalic anhydride
* Phthalic acid
* Phytic acid
* 4-Picoline
* Picric acid — C6H2(OH)(NO2)3
* Pimelic acid
* Pinacol
* Piperazine
* Piperidine
* Piperonal
* Piperylene
* Pivaloyl chloride
* Polyacrylonitrile
* Polyamide 6 = Nylon 6
* Polybenzimidazole - Polybenzimidazole fiber
* Polyethylenimine
* Polygeline
* Polyisobutylene
* Polypropylene
* Polypropylene glycol
* Polystyrene
* Polyurethane
* Polyvinyl acetate
* Polyvinyl alcohol
* Polyvinyl chloride
* Polyvinylidene chloride
* Polyvinylidene fluoride (PVDF)
* Polyvinylpyrrolidone = Poly vinyl pyrrolidone (PVP)
* Porphyrin
* Potassium clavulanate — C8H8KNO5
* Potassium 2-ethyl hexanoate — C8H15KO2
* Prednisone
* Primaquine
* Procaine
* Progesterone
* Prolactin (PRL)
* Proline
* Propane — C3H8
* Propanoic acid
* 2-Propanone
* Propargyl alcohol
* Propiconazole
* Propiolactone
* Propiolic acid
* Propionaldehyde
* Propionitrile
* Propoxur
* Proton-sponge (Aldrich trademark name)
* Purine
* Putrescine — C4H12N2
* Pyrazine
* Pyrazole
* Pyrene
* Pyrethrin
* Pyridazine
* Pyridine — C5H5N
* Pyridinium tribromide
* 2-Pyridone
* Pyridoxal
* Pyridoxine or pyridoxamine (Vitamin B6)
* Pyrilamine
* Pyrimethamine
* Pyrimidine — C4H4N2
* Pyrocatechol violet
* Pyroglutamic acid
* Pyrrole
* Pyrrolidine
* Pyruvic acid
[edit] Q
* Quinaldine
* Quinazoline
* Quinhydrone
* Quinoline
* Quinone
* Quinoxaline
[edit] R
* Raffinose
* Resorcinol
* Retinene
* Retinol (Vitamin A)
* Rhodanine
* Riboflavin (vitamin B2)
* Ribofuranose
* Ribose
* Ricin
* Rosolic acid
* Rotane
* Rotenone
[edit] S
For substances with an s- or secondary- prefix such as s-butyllithium or sec-butyllithium that are not listed below, please see the parent page (in this case under B, butyllithium).
* Saccharin
* Safrole
* Salicin
* Salicylaldehyde
* Salicylic acid
* Salvinorin A
* Sarin
* Sclareol
* Sebacic acid
* Sebacoyl chloride
* Selacholeic acid
* Selenocysteine
* Selenomethionine
* Serine
* Serine kinase
* Serotonin
* Shikimic acid
* Sildenafil (also known as Viagra)
* Skatole
* Snakeroot oil
* Sorbic acid
* Sotolone
* Spermidine
* Squalene
* Stearic acid
* Strychnine
* Styrene
* Succinic anhydride
* Sucrose (sugar)
* Sulfanilamide
* Sulfanilic acid
* Sulforhodamine b
* Suxamethonium chloride
[edit] T
For substances with an t- or tertiary- prefix such as t-butyllithium or tert-butyllithium that are not listed below, please see the parent page (in this case under B, butyllithium). For substances with an t- or trans- prefix such as *trans-2-Butene, you may find these listed under the parent name letter (in this case "B"), as is the norm in chemical catalogues.
* Tabun — C2H5OP(O)(CN)N(CH3)2
* Tannic acid
* Tannin
* Tartaric acid
* Tartrazine
* Taurine
* Terephthalic acid
* Terephthalonitrile
* p-Terphenyl
* α-Terpineol
* Testosterone
* Tetrachlorobiphenyl
* Tetrachloroethylene
* Tetrachloromethane (carbon tetrachloride) – CCl4
* Tetradecane
* Tetraethylene glycol
* Tetrafluoroethene
* Tetrahedrane
* Tetrahydrofuran
* Tetrahydronaphthalene
* Tetramethrin
* Tetramethylsilane (TMS, standard for NMR)
* Tetramethylurea
* Tetranitromethane
* Tetrathiafulvalene (TTF)
* Tetrazine — C2H2N4 (hypothetical)
* Tetrodotoxin
* Tetryl — C7H5N5O8
* Thalidomide
* Thiamine (Vitamin B1) – C12H17ClN4OS·HCl
* Thiazole
* Thioacetamide
* Thiolactic acid
* Thiophene
* Thiophosgene
* Thiourea
* Thiram
* Thorin
* Threonine
* Thrombopoietin
* Thymidine
* Thymine
* Thymol
* Thymolphthalein
* Thyroxine (T4)
* Tiglic acid
* Tinidazole
* Tocopherol (Vitamin E)
* Toluene — C6H5CH3
* Toluene diisocyanate
* p-Toluenesulfonic acid
* o-Toluic acid (2-Methylbenzoic acid)
* p-Toluic acid (4-Methylbenzoic acid)
* Toxaphene
* Triangulane
* Triazole
* Tributyl phosphate
* Tributylamine
* Tributylphosphine
* Trichloroacetic acid
* Trichloroacetonitrile
* 1,1,1-Trichloroethane
* Trichloroethylene
* Trichlorofluoromethane
* 2,4,6-Trichloroanisole
* 2,4,6-Trichlorophenol
* Tris
* Tricine
* Triclabendazole
* Triclosan
* Tricosane
* Tridecane
* Tridecanoic acid
* Triethylaluminium
* Triethylamine
* Triethylamine hydrochloride — C6H15N·HCl
* Triethylene glycol
* Triethylenediamine
* Trifluoroacetic acid (TFA)
* 1,1,1-Trifluoroethane
* 2,2,2-Trifluoroethanol
* Trifluoromethane
* Trimellitic anhydride
* Trimethoxyamphetamine
* Trimethyl phosphite
* Trimethylamine
* Trimethylbenzene
* 2,2,4-Trimethylpentane (isooctane)
* Trinitrotoluene (TNT) – C6H2(NO2)3CH3 or 2,4,6-trinitrotoluene
* Tri-o-cresyl phosphate
* Triphenyl phosphate
* Triphenylamine
* Triphenylantimony
* Triphenylene
* Triphenylmethane
* Triphenylmethanol
* Triphenylphosphine
* Tropane
* Tropinone
* Trypan blue
* Tryptophan
* Tyrosine — C9O3H11N
[edit] U
* Umbelliferone
* Undecanol
* Uracil
* Urea — CO(NH2)2
* Urethane
* Uric acid — C5H4N4O3
* Uridine
* Usnic acid
[edit] V
* Valine
* Valium
* Valproic acid
* Vanillin
* Venlafaxine
* Vigulariol
* Vinyl acetate
* Vinyl fluoride
* Vinylidene chloride
* Violanthrone-79 (16,17-bis(octyloxy)anthra[9,1,2-cde]benzo[rst]pentaphene-5,10-dione)
* Vitamin A (retinol)
* Vitamin B
* Vitamin B1 (thiamine)
* Vitamin B2 (riboflavin)
* Vitamin B3 (niacin or nicotinic acid)
* Vitamin B4 (adenine)
* Vitamin B5 (pantothenic acid)
* Vitamin B6 (pyridoxine or pyridoxamine)
* Vitamin B12 (cobalamin)
* Vitamin C (ascorbic acid)
* Vitamin D (calciferol)
* Vitamin E (tocopherol)
* Vitamin F
* Vitamin H (biotin)
* Vitamin K (naphthoquinone)
* Vitamin M (folic acid)
* Vitamin P (niacin or nicotinic acid)
* Vitamin S
[edit] W
* Warfarin
[edit] X
* Xanthan gum
* Xanthone
* Xylene
* Xylene cyanole ff
* Xylenol orange
* Xylose
* Xylyl bromide
[edit] Y
* Yohimbine hydrochloride - C21H26N2O3
* Yohimbinic acid monohydrate
[edit] Z
* Zingiberene
CHEMICAL BONDS
Chemical bond
From Wikipedia, the free encyclopedia
Jump to: navigation, search
A chemical bond is the physical process responsible for the attractive interactions between atoms and molecules, and that which confers stability to diatomic and polyatomic chemical compounds. The explanation of the attractive forces is a complex area that is described by the laws of quantum electrodynamics. In practice, however, chemists usually rely on quantum theory or qualitative descriptions that are less rigorous but more easily explained to describe chemical bonding. In general, strong chemical bonding is associated with the sharing or transfer of electrons between the participating atoms. Molecules, crystals, and diatomic gases—indeed most of the physical environment around us—are held together by chemical bonds, which dictate the structure of matter.
Bonds vary widely in their strength. Generally covalent and ionic bonds are often described as "strong", whereas hydrogen bonds and van der Waals bonds are generally considered to be "weak". Care should be taken because the strongest of the "weak" bonds can be stronger than the weakest of the "strong" bonds.
Examples of Lewis dot-style chemical bonds between carbon C, hydrogen H, and oxygen O. Lewis dot depictures represent an early attempt to describe chemical bonding and are still widely used today.
Examples of Lewis dot-style chemical bonds between carbon C, hydrogen H, and oxygen O. Lewis dot depictures represent an early attempt to describe chemical bonding and are still widely used today.
Contents
[hide]
* 1 Overview
* 2 History
* 3 Valence bond theory
* 4 Molecular orbital theory
* 5 Comparison of valence bond and molecular orbital theory
* 6 Bonds in chemical formulas
* 7 Strong chemical bonds
o 7.1 Covalent bond
o 7.2 Polar covalent bond
o 7.3 Ionic bond
o 7.4 Coordinate covalent bond
o 7.5 Bent bonds
o 7.6 Three- and four-center bonds
o 7.7 One- and three-electron bonds
o 7.8 Aromatic bond
o 7.9 Metallic bond
* 8 Intermolecular bonding
o 8.1 Permanent dipole to permanent dipole
o 8.2 Hydrogen bond
o 8.3 Instantaneous dipole to induced dipole (van der Waals)
o 8.4 Cation-pi interaction
* 9 Electrons in chemical bonds
* 10 References
* 11 External links
[edit] Overview
The electrons of atoms are electromagnetically attracted by the nuclei of atoms, due to the opposite electric charge of electrons and nuclei. Chemical bonds are characterised by physical states in which a few electrons move partly from one atom to one or more other atoms, driven by the achievement of a lower state of energy from this motion. This lowering of energy is caused by a rearrangement of charges, usually resulting in net decrease in the average distance between the electrons of all the bonded atoms, and their nuclei. The transfer of charge caused by the movement of the electron from one atom to another, also causes the participating atoms (which may number from two to many) to be attracted to one another electromagnetically. The attractive force between atoms is the bond.
Chemical bonds, for the sake of simplicity, are classically assigned characteristics of two major types: covalent and ionic.
In a simplified view of a pure covalent type bond, the bond forms as a few electrons farthest from their atomic nuclei become more attracted to the region of space between two nuclei. In this region, negatively-charged electrons experience attraction from the positively-charged protons from more than one nucleus. This causes some electrons to spend a high probability of their time in the interatomic space. In turn, the nuclei are stabilized in position by the pull from these shared electrons during the fraction of time that the bonding electrons reside between the atoms. Although such bonding electrons do not spend all of their time between atoms, when they spend more time between a given pair of atoms than otherwise, they constitute chemical bonds. Nuclei fixed by such bonds may vibrate, but they are pulled toward each other by the mutual forces of the bonding electrons pulling them together, yet prevented from approaching too closely by their own charge, or else by the mutual repulsion of other inner electrons, which are held so closely and tightly to individual nuclei that they cannot be shared to any important degree.
In a simplified view of a pure ionic type bond, one or more outer electrons are not shared between atoms, but donated from one atom to another. In such a bond, the structure of the electron cloud of one of the nuclei contains an available space for another electron, which allows an additional electron to experience a greater net attraction from the nucleus than is experienced by outer electrons in a neighboring atom toward their own nucleus. This difference in available states causes effective transfer of one or more electrons from one atom to another atom, where they can be more tightly bound. This transfer causes the donating atom to assume a net positive charge, and the other to assume a net negative charge; the atoms thus become positive or negatively charged ions. The bond then results from electrostatic attraction between these ionized atoms.
Most bonds have a mixture of covalent and ionic character, as bonding electrons are shared between atoms, but shared somewhat unevenly. All bonds can be explained by quantum theory, but in practice, simplification rules allow chemists to predict the strength, directionality, and polarity of bonds. The octet rule and VSEPR theory are two examples. More sophisticated theories are valence bond theory which includes orbital hybridization and resonance, and the linear combination of atomic orbitals molecular orbital method which includes ligand field theory. Electrostatics are used to describe bond polarities and the effects they have on chemical substances.
[edit] History
Main articles: History of chemistry and History of the molecule
Early speculations into the nature of the chemical bond, from as early as the 12th century, supposed that certain types of chemical species were joined by a type of chemical affinity. In 1704, Isaac Newton famously outlined his atomic bonding theory, in "Query 31" of his Opticks, whereby atoms attach to each other by some "force". Specifically, after acknowledging the various popular theories, in vogue at the time, of how atoms were reasoned to attach to each other, i.e. “hooked atoms”, “glued together by rest”, or “stuck together by conspiring motions”, Newton states that he would rather infer from their cohesion, that:
“ Particles attract one another by some force, which in immediate contact is exceedingly strong, at small distances performs the chemical operations, and reaches not far from the particles with any sensible effect. ”
In 1819, on the heels of the invention of the voltaic pile, Jöns Jakob Berzelius developed a theory of chemical combination stressing the electronegative and electropositive character of the combining atoms. By the mid 19th century, Edward Frankland, F.A. Kekule, A.S. Couper, A.M. Butlerov, and Hermann Kolbe, building on the theory of radicals, developed the theory of valency, originally called “combining power”, in which compounds were joined owing to an attraction of positive and negative poles. In 1916, chemist Gilbert N. Lewis developed the concept of the electron-pair bond, in which two atoms may share one to six electrons, thus forming the single electron bond, a single bond, a double bond, or a triple bond:
In Lewis' own words:
“ An electron may form a part of the shell of two different atoms and cannot be said to belong to either one exclusively. ”
That same year, Walther Kossel put forward a theory similar to Lewis' only his model assumed complete transfers of electrons between atoms, and was thus a model of polar bonds. Both Lewis and Kossel structured their bonding models on that of Abegg's rule (1904).
In 1927, the first mathematically complete quantum description of a simple chemical bond, i.e. that produced by one electron in the hydrogen molecular ion, H2+, was derived by the Danish physicist Oyvind Burrau.[1] This work showed that the quantum approach to chemical bonds could be fundamentally and quantitatively correct, but the mathematical methods used could not be extended to molecules containing more than one electron. A more practical, albeit less quantitative, approach was put forward in the same year by Walter Heitler and Fritz London. The Heitler-London method forms the basis of what is now called valence bond theory. In 1929, the linear combination of atomic orbitals molecular orbital method (LCAO) approximation was introduced by Sir John Lennard-Jones, who also suggested methods to derive electronic structures of molecules of F2 (fluorine) and O2 (oxygen) molecules, from basic quantum principles. This molecular orbital theory represented a covalent bond as a orbitals formed by combining the quantum mechanical Schrödinger atomic orbitals which had been hypothesized for electrons in single atoms. The equations for bonding electrons in multi-electron atoms could not be solved to mathematical perfection (i.e., analytically), but approximations for them still gave many good qualitative predictions and results. Most quantitative calculations in modern quantum chemistry use either valence bond or molecular orbital theory as a starting point, although a third approach, Density Functional Theory, has become increasingly popular in recent years.
In 1935, H. H. James and A. S. Coolidge carried out a calculation on the dihydrogen molecule that, unlike all previous calculation which used functions only of the distance of the electron from the atomic nucleus, used functions which also explicitly added the distance between the two electrons.[2] With up to 13 adjustable parameters they obtained a result very close to the experimental result for the dissociation energy. Later extensions have used up to 54 parameters and give excellent agreement with experiment. This calculation convinced the scientific community that quantum theory could give agreement with experiment. However this approach has none of the physical pictures of the valence bond and molecular orbital theories and is difficult to extend to larger molecules.
[edit] Valence bond theory
Main article: Valence bond theory
In the year 1927, valence bond theory was formulated which argued essentially that a chemical bond forms when two valence electrons, in their respective atomic orbitals, work or function to hold two nuclei together, by virtue of system energy lowering effects. In 1931, building on this theory, chemist Linus Pauling published what some consider one of the most important papers in the history of chemistry: “On the Nature of the Chemical Bond”. In this paper, building on the works of Lewis, and the valence bond theory (VB) of Heitler and London, and his own earlier work, he presented six rules for the shared electron bond, the first three of which were already generally known:
1. The electron-pair bond forms through the interaction of an unpaired electron on each of two atoms.
2. The spins of the electrons have to be opposed.
3. Once paired, the two electrons cannot take part in additional bonds.
His last three rules were new:
4. The electron-exchange terms for the bond involves only one wave function from each atom.
5. The available electrons in the lowest energy level form the strongest bonds.
6. Of two orbitals in an atom, the one that can overlap the most with an orbital from another atom will form the strongest bond, and this bond will tend to lie in the direction of the concentrated orbital.
Building on this article, Pauling’s 1939 textbook: On the Nature of the Chemical Bond would become what some have called the “bible” of modern chemistry. This book helped experimental chemists to understand the impact of quantum theory on chemistry. However, the later edition in 1959 failed to address adequately the problems that appeared to be better understood by molecular orbital theory. The impact of valence theory declined during the 1960s and 1970's as molecular orbital theory grew in popularity and was implemented in many large computer programs. Since the 1980s, the more difficult problems of implementing valence bond theory into computer programs have been largely solved and valence bond theory has seen a resurgence.
[edit] Molecular orbital theory
Main article: Molecular orbital theory
Molecular orbital theory (MO) uses a linear combination of atomic orbitals to form molecular orbitals which cover the whole molecule. These are often divided into bonding orbitals, anti-bonding orbitals, and non-bonding orbitals. A molecular orbital is merely a Schrödinger orbital which includes several, but often only two nuclei. If this orbital is of type in which the electron(s) in the orbital have a higher probability of being between nuclei than elsewhere, the orbital will be a bonding orbital, and will tend to hold the nuclei together. If the electrons tend to be present in a molecular orbital in which they spend more time elsewhere than between the nuclei, the orbital will function as an anti-bonding orbital and will actually weaken the bond. Electrons in non-bonding orbitals tend to be in deep orbitals (nearly atomic orbitals) associated almost entirely with one nucleus or the other, and thus they spend equal time between nuclei or not. These electrons neither contribute nor detract from bond strength.
[edit] Comparison of valence bond and molecular orbital theory
In some respects valence bond theory is superior to molecular orbital theory. When applied to the simplest two-electron molecule, H2, valence bond theory, even at the simplest Heitler-London approach, gives a much closer approximation to the bond energy, and it provides a much more accurate representation of the behavior of the electrons as chemical bonds are formed and broken. In contrast simple molecular orbital theory predicts that the hydrogen molecule dissociates into a linear superposition of hydrogen atoms and positive and negative hydrogen ions, a completely unphysical result. This explains in part why the curve of total energy against interatomic distance for the valence bond method lies above the curve for the molecular orbital method at all distances and most particularly so for large distances. This situation arises for all homonuclear diatomic molecules and is particularly a problem for F2, where the minimum energy of the curve with molecular orbital theory is still higher in energy than the energy of two F atoms.
The concepts of hybridization are so versatile, and the variability in bonding in most organic compounds is so modest, that valence bond theory remains an integral part of the vocabulary of organic chemistry. However, the work of Friedrich Hund, Robert Mulliken, and Gerhard Herzberg showed that molecular orbital theory provided a more appropriate description of the spectroscopic, ionization and magnetic properties of molecules. The deficiencies of valence bond theory became apparent when hypervalent molecules (e.g. PF5) were explained without the use of d orbitals that were crucial to the bonding hybridisation scheme proposed for such molecules by Pauling. Metal complexes and electron deficient compounds (e.g. diborane) also appeared to be well described by molecular orbital theory, although valence bond descriptions have been made.
In the 1930s the two methods strongly competed until it was realised that they are both approximations to a better theory. If we take the simple valence bond structure and mix in all possible covalent and ionic structures arising from a particular set of atomic orbitals, we reach what is called the full configuration interaction wave function. If we take the simple molecular orbital description of the ground state and combine that function with the functions describing all possible excited states using unoccupied orbitals arising from the same set of atomic orbitals, we also reach the full configuration interaction wavefunction. It can be then seen that the simple molecular orbital approach gives too much weight to the ionic structures, while the simple valence bond approach gives too little. This can also be described as saying that the molecular orbital approach is too delocalised, while the valence bond approach is too localised.
The two approaches are now regarded as complementary, each providing its own insights into the problem of chemical bonding. Modern calculations in quantum chemistry usually start from (but ultimately go far beyond) a molecular orbital rather than a valence bond approach, not because of any intrinsic superiority in the former but rather because the MO approach is more readily adapted to numerical computations. However better valence bond programs are now available.
[edit] Bonds in chemical formulas
The 3-dimensionality of atoms and molecules makes it difficult to use a single technique for indicating orbitals and bonds. In molecular formulas the chemical bonds (binding orbitals) between atoms are indicated by various different methods according to the type of discussion. Sometimes, they are completely neglected. For example, in organic chemistry chemists are sometimes concerned only with the functional groups of the molecule. Thus, the molecular formula of ethanol (a compound in alcoholic beverages) may be written in a paper in conformational, 3-dimensional, full 2-dimensional (indicating every bond with no 3-dimensional directions), compressed 2-dimensional (CH3–CH2–OH), separating the functional group from another part of the molecule (C2H5OH), or by its atomic constituents (C2H6O), according to what is discussed. Sometimes, even the non-bonding valence shell electrons (with the 2-dimensionalized approximate directions) are marked, i.e. for elemental carbon .'C'. Some chemists may also mark the respective orbitals, i.e. the hypothetical ethene−4 anion (\/C=C/\ −4) indicating the possibility of bond formation.
[edit] Strong chemical bonds
Typical bond lengths in pm
and bond energies in kJ/mol.
Bond lengths can be converted to Å
by division by 100 (1 Å = 100 pm).
Data taken from[3].
Bond Length
(pm) Energy
(kJ/mol)
H — Hydrogen
H–H 74 436
H–C 109 413
H–N 101 391
H–O 96 366
H–F 92 568
H–Cl 127 432
H–Br 141 366
C — Carbon
C–H 109 413
C–C 154 348
C=C 134 614
C≡C 120 839
C–N 147 308
C–O 143 360
C–F 135 488
C–Cl 177 330
C–Br 194 288
C–I 214 216
C–S 182 272
N — Nitrogen
N–H 101 391
N–C 147 308
N–N 145 170
N≡N 110 945
O — Oxygen
O–H 96 366
O–C 143 360
O–O 148 145
O=O 121 498
F, Cl, Br, I — Halogens
F–H 92 568
F–F 142 158
F–C 135 488
Cl–H 127 432
Cl–C 177 330
Cl–Cl 199 243
Br–H 141 366
Br–C 194 288
Br–Br 228 193
I–H 161 298
I–C 214 216
I–I 267 151
S — Sulfur
C–S 182 272
These chemical bonds are intramolecular forces, which hold atoms together in molecules. In the simplistic localized view of bonding, the number of electrons participating in a bond (or located in a bonding orbital) is typically multiples of two, four, or six, respectively. Even numbers are common because electrons enjoy lower energy states, if paired. Substantially more advanced bonding theories have shown that bond strength is not always a whole number, depending on the distribution of electrons to each atom involved in a bond. For example, the carbons in benzene are connected to each other with about 1.5 bonds, and the two atoms in nitric oxide NO, are connected with about 2.5 bonds. Quadruple bonds are also well known. The type of strong bond depends on the difference in electronegativity and the distribution of the electron orbital paths available to the atoms that are bonded. The larger the difference in electronegativity, the more an electron is attracted to a particular atom involved in the bond, and the more "ionic" properties the bond is said to have ("ionic" means the bond electron(s) are unequally shared). The smaller the difference in electronegativity, the more covalent properties (full sharing) the bond has.
[edit] Covalent bond
Main article: Covalent bond
Covalent bonding is a common type of bonding, in which the electronegativity difference between the bonded atoms is small or non-existent. Bonds within most organic compounds are described as covalent. See sigma bonds and pi bonds for LCAO-description of such bonding.
[edit] Polar covalent bond
Main article: Polar covalent bond
Polar covalent bonding is intermediate in character between a covalent and an ionic bond.
[edit] Ionic bond
Main article: Ionic bond
Ionic bonding is a type of electrostatic interaction between atoms which have a large electronegativity difference. There is no precise value that distinguishes ionic from covalent bonding but a difference of electronegativity of over 2.0 is likely to be ionic and a difference of less than 1.5 is likely to be covalent.[4] Ionic bonding leads to separate positive and negative ions. Ionic charges are commonly between −3e to +3e.
[edit] Coordinate covalent bond
Main article: Coordinate covalent bond
Coordinate covalent bonding, sometimes referred to as dative bonding, is a kind of covalent bonding, in which the covalent bonding electrons originate solely from one of the atoms, the electron-pair donor or Lewis base but are approximately equally shared in the formation of a covalent bond. This concept is somewhat fading as chemists increasingly embrace molecular orbital theory. Examples of coordinate covalent bonding occur in nitrones and ammonia borane. The arrangement is different from an ionic bond in that the electronegativity difference is small, resulting in covalency. They are shown by an arrow unlike other bonds. This arrow shows its head towards the electron acceptor or lewis acid and its tail towards the lewis base. This bond is seen in ammonium.
[edit] Bent bonds
Main article: Bent bond
Bent bonds, also known as banana bonds, are bonds in strained or otherwise sterically hindered molecules those binding orbitals are forced into a banana-like form. Bent bonds are often more susceptible to reactions than ordinary bonds.
[edit] Three- and four-center bonds
Main articles: three-center two-electron bond, three-center four-electron bond, and four-center two-electron bond
In three-center two-electron bonds three atoms share two electrons in bonding. This type of bonding occurs in electron deficient compounds like diborane. Each such bond (2 per molecule in diborane) contains a pair of electrons which connect the boron atoms to each other in a banana shape (shown as a more sharply angled section in the stick model at right), with a proton (nucleus of a hydrogen atom) in the middle of the bond, sharing electrons with both boron atoms.
Three-center four-electron bonds are used to explain the bonding in hypervalent molecules. In certain cluster compounds so-called four-center two-electron bonds also have been postulated. These bonds are not to be confused with "two-center four-electron bonds", which is one way of describing conventional double bonds between two atoms. The MO description of such double bonds would distinguish one bond, with electron density concentrated along the line between atoms, as a σ bond, but the other bond, with electron density concentrated in lobes on either side of the axial line, as a π (pi) bond.
[edit] One- and three-electron bonds
Bonds with one or three electrons can be found in radical species, which have an odd number of electrons. The simplest example of a 1-electron bond is found in the hydrogen molecular cation, H2+. One-electron bonds often have about half the bond energy of a 2-electron bond, and are therefore called "half bonds". However, there are exceptions: in the case of dilithium, the bond is actually stronger for the 1-electron Li2+ than for the 2-electron Li2. This exception can be explained in terms of hybridization and inner-shell effects. [5]
The simplest example of three-electron bonding can be found in the helium dimer cation, He2+, and can also be considered a "half bond" because, in molecular orbital terms, the third electron is in an anti-bonding orbital which cancels out half of the bond formed by the other two electrons. Another example of a molecule containing a 3-electron bond, in addition to two 2-electron bonds, is nitric oxide, NO. The oxygen molecule, O2 can also be regarded as having two 3-electron bonds and one 2-electron bond, which accounts for its paramagnetism and its formal bond order of 2.[6]
Molecules with odd-electron bonds are usually highly reactive. These types of bond are only stable between atoms with similar electronegativities.[6]
[edit] Aromatic bond
Main article: Aromaticity
In most cases, the locations of electrons cannot be simplified to simple lines (place for two electrons) or dots (a single electron). In aromatic bonds which occur in planar rings of atoms where the 4n+2 rule determines whether ring molecules would show extra stability.
In benzene, the prototypical aromatic compound, 18 bonding electrons bind 6 carbon atoms together to form a planar ring structure. The bond "order" (average number of bonds) between the different carbon atoms may be said to be (18/6)/2=1.5, but in this case the bonds are all identical from the chemical point of view. They may sometimes be written as single bonds alternating with double bonds, but the view of all ring bonds as being equivalently about 1.5 bonds in strength, is much closer to truth.
In the case of heterocyclic aromatics and substituted benzenes, the electronegativity differences between different parts of the ring may dominate the chemical behaviour of aromatic ring bonds, which otherwise are equivalent.
[edit] Metallic bond
Main article: Metallic bond
In metallic bonding, bonding electrons are delocalized over a lattice of atoms. By contrast, in ionic compounds, the locations of the binding electrons and their charges are static. Because of delocalization or the free moving of electrons, it leads to the metallic properties such as conductivity, ductility and hardness.
[edit] Intermolecular bonding
There are four basic types of bonds that can be formed between two or more (otherwise non-associated) molecules, ions or atoms. Intermolecular forces cause molecules to be attracted or repulsed by each other. Often, these define some of the physical characteristics (such as the melting point) of a substance.
[edit] Permanent dipole to permanent dipole
Main article: Intermolecular force
A large electronegativity difference between two strongly bonded atoms within a molecule causes a dipole to form (a dipole is a pair of permanent partial charges). Dipoles will attract or repel each other.
[edit] Hydrogen bond
Main article: Hydrogen bond
In some ways this is an especially strong example of a permanent dipole, as above. However, in the hydrogen bond, the hydrogen proton comes closer to being shared between target and donor atoms, in a three-center two-electron bond like that in diborane. Hydrogen bonds explain the relatively high boiling points of liquids like water, ammonia, and hydrogen fluoride, compared with their heavier counterparts in the same periodic table column.
[edit] Instantaneous dipole to induced dipole (van der Waals)
Main article: van der Waals forces
Instantaneous dipole to induced dipole, or van der Waals forces, are the weakest, but also the most prolific—occurring between all chemical substances. Imagine a helium atom: At any one point in time, the electron cloud around the (otherwise neutral) atom can be thought to be slightly imbalanced, with momentarily more negative charge on one side. This is referred to as an instantaneous dipole. This dipole, with its slight charge imbalance, may attract or repel the electrons within a neighbouring helium atom, setting up another dipole. The two atoms will be attracted for an instant, before the charge rebalances and the atoms move on.
[edit] Cation-pi interaction
Main article: Cation-pi interaction
Cation-pi interactions occur between the localized negative charge of π orbital electrons, located above and below the plane of an aromatic ring, and a positive charge.
[edit] Electrons in chemical bonds
Many simple compounds involve covalent bonds. These molecules have structures that can be predicted using valence bond theory, and the properties of atoms involved can be understood using concepts such as oxidation number. Other compounds that involve ionic structures can be understood using theories from classical physics.
In the case of ionic bonding, electrons are mainly localized on the individual atoms, and electrons do not travel between the atoms very much. Each atom is assigned an overall electric charge to help conceptualize the molecular orbital's distribution. The forces between atoms (or ions) are largely characterized by isotropic continuum electrostatic potentials.
By contrast, in covalent bonding, the electron density within a bond is not assigned to individual atoms, but is instead delocalized in the MOs between atoms. The widely accepted theory of the linear combination of atomic orbitals (LCAO) helps describe the molecular orbital structures and energies based on the atomic orbitals of the atoms they came from. Unlike pure ionic bonds, covalent bonds may have directed anisotropic properties. These may have their own names, too, such as Sigma and Pi bond.
Atoms can also form bonds that are intermediates between ionic and covalent. This is because these definitions are based on the extent of electron delocalization. Electrons can be partially delocalized between atoms, but spend more time around one atom than another. This type of bond is often called polar covalent. See electronegativity.
Thus, the electrons in a molecular orbital (or 'in a polar covalent, or in a covalent bond') can be said to be either localized on certain atom(s) or delocalized between two or more atoms. The type of bond between two atoms is defined by how much the electron density is localized or delocalized among the atoms of the bonds.
From Wikipedia, the free encyclopedia
Jump to: navigation, search
A chemical bond is the physical process responsible for the attractive interactions between atoms and molecules, and that which confers stability to diatomic and polyatomic chemical compounds. The explanation of the attractive forces is a complex area that is described by the laws of quantum electrodynamics. In practice, however, chemists usually rely on quantum theory or qualitative descriptions that are less rigorous but more easily explained to describe chemical bonding. In general, strong chemical bonding is associated with the sharing or transfer of electrons between the participating atoms. Molecules, crystals, and diatomic gases—indeed most of the physical environment around us—are held together by chemical bonds, which dictate the structure of matter.
Bonds vary widely in their strength. Generally covalent and ionic bonds are often described as "strong", whereas hydrogen bonds and van der Waals bonds are generally considered to be "weak". Care should be taken because the strongest of the "weak" bonds can be stronger than the weakest of the "strong" bonds.
Examples of Lewis dot-style chemical bonds between carbon C, hydrogen H, and oxygen O. Lewis dot depictures represent an early attempt to describe chemical bonding and are still widely used today.
Examples of Lewis dot-style chemical bonds between carbon C, hydrogen H, and oxygen O. Lewis dot depictures represent an early attempt to describe chemical bonding and are still widely used today.
Contents
[hide]
* 1 Overview
* 2 History
* 3 Valence bond theory
* 4 Molecular orbital theory
* 5 Comparison of valence bond and molecular orbital theory
* 6 Bonds in chemical formulas
* 7 Strong chemical bonds
o 7.1 Covalent bond
o 7.2 Polar covalent bond
o 7.3 Ionic bond
o 7.4 Coordinate covalent bond
o 7.5 Bent bonds
o 7.6 Three- and four-center bonds
o 7.7 One- and three-electron bonds
o 7.8 Aromatic bond
o 7.9 Metallic bond
* 8 Intermolecular bonding
o 8.1 Permanent dipole to permanent dipole
o 8.2 Hydrogen bond
o 8.3 Instantaneous dipole to induced dipole (van der Waals)
o 8.4 Cation-pi interaction
* 9 Electrons in chemical bonds
* 10 References
* 11 External links
[edit] Overview
The electrons of atoms are electromagnetically attracted by the nuclei of atoms, due to the opposite electric charge of electrons and nuclei. Chemical bonds are characterised by physical states in which a few electrons move partly from one atom to one or more other atoms, driven by the achievement of a lower state of energy from this motion. This lowering of energy is caused by a rearrangement of charges, usually resulting in net decrease in the average distance between the electrons of all the bonded atoms, and their nuclei. The transfer of charge caused by the movement of the electron from one atom to another, also causes the participating atoms (which may number from two to many) to be attracted to one another electromagnetically. The attractive force between atoms is the bond.
Chemical bonds, for the sake of simplicity, are classically assigned characteristics of two major types: covalent and ionic.
In a simplified view of a pure covalent type bond, the bond forms as a few electrons farthest from their atomic nuclei become more attracted to the region of space between two nuclei. In this region, negatively-charged electrons experience attraction from the positively-charged protons from more than one nucleus. This causes some electrons to spend a high probability of their time in the interatomic space. In turn, the nuclei are stabilized in position by the pull from these shared electrons during the fraction of time that the bonding electrons reside between the atoms. Although such bonding electrons do not spend all of their time between atoms, when they spend more time between a given pair of atoms than otherwise, they constitute chemical bonds. Nuclei fixed by such bonds may vibrate, but they are pulled toward each other by the mutual forces of the bonding electrons pulling them together, yet prevented from approaching too closely by their own charge, or else by the mutual repulsion of other inner electrons, which are held so closely and tightly to individual nuclei that they cannot be shared to any important degree.
In a simplified view of a pure ionic type bond, one or more outer electrons are not shared between atoms, but donated from one atom to another. In such a bond, the structure of the electron cloud of one of the nuclei contains an available space for another electron, which allows an additional electron to experience a greater net attraction from the nucleus than is experienced by outer electrons in a neighboring atom toward their own nucleus. This difference in available states causes effective transfer of one or more electrons from one atom to another atom, where they can be more tightly bound. This transfer causes the donating atom to assume a net positive charge, and the other to assume a net negative charge; the atoms thus become positive or negatively charged ions. The bond then results from electrostatic attraction between these ionized atoms.
Most bonds have a mixture of covalent and ionic character, as bonding electrons are shared between atoms, but shared somewhat unevenly. All bonds can be explained by quantum theory, but in practice, simplification rules allow chemists to predict the strength, directionality, and polarity of bonds. The octet rule and VSEPR theory are two examples. More sophisticated theories are valence bond theory which includes orbital hybridization and resonance, and the linear combination of atomic orbitals molecular orbital method which includes ligand field theory. Electrostatics are used to describe bond polarities and the effects they have on chemical substances.
[edit] History
Main articles: History of chemistry and History of the molecule
Early speculations into the nature of the chemical bond, from as early as the 12th century, supposed that certain types of chemical species were joined by a type of chemical affinity. In 1704, Isaac Newton famously outlined his atomic bonding theory, in "Query 31" of his Opticks, whereby atoms attach to each other by some "force". Specifically, after acknowledging the various popular theories, in vogue at the time, of how atoms were reasoned to attach to each other, i.e. “hooked atoms”, “glued together by rest”, or “stuck together by conspiring motions”, Newton states that he would rather infer from their cohesion, that:
“ Particles attract one another by some force, which in immediate contact is exceedingly strong, at small distances performs the chemical operations, and reaches not far from the particles with any sensible effect. ”
In 1819, on the heels of the invention of the voltaic pile, Jöns Jakob Berzelius developed a theory of chemical combination stressing the electronegative and electropositive character of the combining atoms. By the mid 19th century, Edward Frankland, F.A. Kekule, A.S. Couper, A.M. Butlerov, and Hermann Kolbe, building on the theory of radicals, developed the theory of valency, originally called “combining power”, in which compounds were joined owing to an attraction of positive and negative poles. In 1916, chemist Gilbert N. Lewis developed the concept of the electron-pair bond, in which two atoms may share one to six electrons, thus forming the single electron bond, a single bond, a double bond, or a triple bond:
In Lewis' own words:
“ An electron may form a part of the shell of two different atoms and cannot be said to belong to either one exclusively. ”
That same year, Walther Kossel put forward a theory similar to Lewis' only his model assumed complete transfers of electrons between atoms, and was thus a model of polar bonds. Both Lewis and Kossel structured their bonding models on that of Abegg's rule (1904).
In 1927, the first mathematically complete quantum description of a simple chemical bond, i.e. that produced by one electron in the hydrogen molecular ion, H2+, was derived by the Danish physicist Oyvind Burrau.[1] This work showed that the quantum approach to chemical bonds could be fundamentally and quantitatively correct, but the mathematical methods used could not be extended to molecules containing more than one electron. A more practical, albeit less quantitative, approach was put forward in the same year by Walter Heitler and Fritz London. The Heitler-London method forms the basis of what is now called valence bond theory. In 1929, the linear combination of atomic orbitals molecular orbital method (LCAO) approximation was introduced by Sir John Lennard-Jones, who also suggested methods to derive electronic structures of molecules of F2 (fluorine) and O2 (oxygen) molecules, from basic quantum principles. This molecular orbital theory represented a covalent bond as a orbitals formed by combining the quantum mechanical Schrödinger atomic orbitals which had been hypothesized for electrons in single atoms. The equations for bonding electrons in multi-electron atoms could not be solved to mathematical perfection (i.e., analytically), but approximations for them still gave many good qualitative predictions and results. Most quantitative calculations in modern quantum chemistry use either valence bond or molecular orbital theory as a starting point, although a third approach, Density Functional Theory, has become increasingly popular in recent years.
In 1935, H. H. James and A. S. Coolidge carried out a calculation on the dihydrogen molecule that, unlike all previous calculation which used functions only of the distance of the electron from the atomic nucleus, used functions which also explicitly added the distance between the two electrons.[2] With up to 13 adjustable parameters they obtained a result very close to the experimental result for the dissociation energy. Later extensions have used up to 54 parameters and give excellent agreement with experiment. This calculation convinced the scientific community that quantum theory could give agreement with experiment. However this approach has none of the physical pictures of the valence bond and molecular orbital theories and is difficult to extend to larger molecules.
[edit] Valence bond theory
Main article: Valence bond theory
In the year 1927, valence bond theory was formulated which argued essentially that a chemical bond forms when two valence electrons, in their respective atomic orbitals, work or function to hold two nuclei together, by virtue of system energy lowering effects. In 1931, building on this theory, chemist Linus Pauling published what some consider one of the most important papers in the history of chemistry: “On the Nature of the Chemical Bond”. In this paper, building on the works of Lewis, and the valence bond theory (VB) of Heitler and London, and his own earlier work, he presented six rules for the shared electron bond, the first three of which were already generally known:
1. The electron-pair bond forms through the interaction of an unpaired electron on each of two atoms.
2. The spins of the electrons have to be opposed.
3. Once paired, the two electrons cannot take part in additional bonds.
His last three rules were new:
4. The electron-exchange terms for the bond involves only one wave function from each atom.
5. The available electrons in the lowest energy level form the strongest bonds.
6. Of two orbitals in an atom, the one that can overlap the most with an orbital from another atom will form the strongest bond, and this bond will tend to lie in the direction of the concentrated orbital.
Building on this article, Pauling’s 1939 textbook: On the Nature of the Chemical Bond would become what some have called the “bible” of modern chemistry. This book helped experimental chemists to understand the impact of quantum theory on chemistry. However, the later edition in 1959 failed to address adequately the problems that appeared to be better understood by molecular orbital theory. The impact of valence theory declined during the 1960s and 1970's as molecular orbital theory grew in popularity and was implemented in many large computer programs. Since the 1980s, the more difficult problems of implementing valence bond theory into computer programs have been largely solved and valence bond theory has seen a resurgence.
[edit] Molecular orbital theory
Main article: Molecular orbital theory
Molecular orbital theory (MO) uses a linear combination of atomic orbitals to form molecular orbitals which cover the whole molecule. These are often divided into bonding orbitals, anti-bonding orbitals, and non-bonding orbitals. A molecular orbital is merely a Schrödinger orbital which includes several, but often only two nuclei. If this orbital is of type in which the electron(s) in the orbital have a higher probability of being between nuclei than elsewhere, the orbital will be a bonding orbital, and will tend to hold the nuclei together. If the electrons tend to be present in a molecular orbital in which they spend more time elsewhere than between the nuclei, the orbital will function as an anti-bonding orbital and will actually weaken the bond. Electrons in non-bonding orbitals tend to be in deep orbitals (nearly atomic orbitals) associated almost entirely with one nucleus or the other, and thus they spend equal time between nuclei or not. These electrons neither contribute nor detract from bond strength.
[edit] Comparison of valence bond and molecular orbital theory
In some respects valence bond theory is superior to molecular orbital theory. When applied to the simplest two-electron molecule, H2, valence bond theory, even at the simplest Heitler-London approach, gives a much closer approximation to the bond energy, and it provides a much more accurate representation of the behavior of the electrons as chemical bonds are formed and broken. In contrast simple molecular orbital theory predicts that the hydrogen molecule dissociates into a linear superposition of hydrogen atoms and positive and negative hydrogen ions, a completely unphysical result. This explains in part why the curve of total energy against interatomic distance for the valence bond method lies above the curve for the molecular orbital method at all distances and most particularly so for large distances. This situation arises for all homonuclear diatomic molecules and is particularly a problem for F2, where the minimum energy of the curve with molecular orbital theory is still higher in energy than the energy of two F atoms.
The concepts of hybridization are so versatile, and the variability in bonding in most organic compounds is so modest, that valence bond theory remains an integral part of the vocabulary of organic chemistry. However, the work of Friedrich Hund, Robert Mulliken, and Gerhard Herzberg showed that molecular orbital theory provided a more appropriate description of the spectroscopic, ionization and magnetic properties of molecules. The deficiencies of valence bond theory became apparent when hypervalent molecules (e.g. PF5) were explained without the use of d orbitals that were crucial to the bonding hybridisation scheme proposed for such molecules by Pauling. Metal complexes and electron deficient compounds (e.g. diborane) also appeared to be well described by molecular orbital theory, although valence bond descriptions have been made.
In the 1930s the two methods strongly competed until it was realised that they are both approximations to a better theory. If we take the simple valence bond structure and mix in all possible covalent and ionic structures arising from a particular set of atomic orbitals, we reach what is called the full configuration interaction wave function. If we take the simple molecular orbital description of the ground state and combine that function with the functions describing all possible excited states using unoccupied orbitals arising from the same set of atomic orbitals, we also reach the full configuration interaction wavefunction. It can be then seen that the simple molecular orbital approach gives too much weight to the ionic structures, while the simple valence bond approach gives too little. This can also be described as saying that the molecular orbital approach is too delocalised, while the valence bond approach is too localised.
The two approaches are now regarded as complementary, each providing its own insights into the problem of chemical bonding. Modern calculations in quantum chemistry usually start from (but ultimately go far beyond) a molecular orbital rather than a valence bond approach, not because of any intrinsic superiority in the former but rather because the MO approach is more readily adapted to numerical computations. However better valence bond programs are now available.
[edit] Bonds in chemical formulas
The 3-dimensionality of atoms and molecules makes it difficult to use a single technique for indicating orbitals and bonds. In molecular formulas the chemical bonds (binding orbitals) between atoms are indicated by various different methods according to the type of discussion. Sometimes, they are completely neglected. For example, in organic chemistry chemists are sometimes concerned only with the functional groups of the molecule. Thus, the molecular formula of ethanol (a compound in alcoholic beverages) may be written in a paper in conformational, 3-dimensional, full 2-dimensional (indicating every bond with no 3-dimensional directions), compressed 2-dimensional (CH3–CH2–OH), separating the functional group from another part of the molecule (C2H5OH), or by its atomic constituents (C2H6O), according to what is discussed. Sometimes, even the non-bonding valence shell electrons (with the 2-dimensionalized approximate directions) are marked, i.e. for elemental carbon .'C'. Some chemists may also mark the respective orbitals, i.e. the hypothetical ethene−4 anion (\/C=C/\ −4) indicating the possibility of bond formation.
[edit] Strong chemical bonds
Typical bond lengths in pm
and bond energies in kJ/mol.
Bond lengths can be converted to Å
by division by 100 (1 Å = 100 pm).
Data taken from[3].
Bond Length
(pm) Energy
(kJ/mol)
H — Hydrogen
H–H 74 436
H–C 109 413
H–N 101 391
H–O 96 366
H–F 92 568
H–Cl 127 432
H–Br 141 366
C — Carbon
C–H 109 413
C–C 154 348
C=C 134 614
C≡C 120 839
C–N 147 308
C–O 143 360
C–F 135 488
C–Cl 177 330
C–Br 194 288
C–I 214 216
C–S 182 272
N — Nitrogen
N–H 101 391
N–C 147 308
N–N 145 170
N≡N 110 945
O — Oxygen
O–H 96 366
O–C 143 360
O–O 148 145
O=O 121 498
F, Cl, Br, I — Halogens
F–H 92 568
F–F 142 158
F–C 135 488
Cl–H 127 432
Cl–C 177 330
Cl–Cl 199 243
Br–H 141 366
Br–C 194 288
Br–Br 228 193
I–H 161 298
I–C 214 216
I–I 267 151
S — Sulfur
C–S 182 272
These chemical bonds are intramolecular forces, which hold atoms together in molecules. In the simplistic localized view of bonding, the number of electrons participating in a bond (or located in a bonding orbital) is typically multiples of two, four, or six, respectively. Even numbers are common because electrons enjoy lower energy states, if paired. Substantially more advanced bonding theories have shown that bond strength is not always a whole number, depending on the distribution of electrons to each atom involved in a bond. For example, the carbons in benzene are connected to each other with about 1.5 bonds, and the two atoms in nitric oxide NO, are connected with about 2.5 bonds. Quadruple bonds are also well known. The type of strong bond depends on the difference in electronegativity and the distribution of the electron orbital paths available to the atoms that are bonded. The larger the difference in electronegativity, the more an electron is attracted to a particular atom involved in the bond, and the more "ionic" properties the bond is said to have ("ionic" means the bond electron(s) are unequally shared). The smaller the difference in electronegativity, the more covalent properties (full sharing) the bond has.
[edit] Covalent bond
Main article: Covalent bond
Covalent bonding is a common type of bonding, in which the electronegativity difference between the bonded atoms is small or non-existent. Bonds within most organic compounds are described as covalent. See sigma bonds and pi bonds for LCAO-description of such bonding.
[edit] Polar covalent bond
Main article: Polar covalent bond
Polar covalent bonding is intermediate in character between a covalent and an ionic bond.
[edit] Ionic bond
Main article: Ionic bond
Ionic bonding is a type of electrostatic interaction between atoms which have a large electronegativity difference. There is no precise value that distinguishes ionic from covalent bonding but a difference of electronegativity of over 2.0 is likely to be ionic and a difference of less than 1.5 is likely to be covalent.[4] Ionic bonding leads to separate positive and negative ions. Ionic charges are commonly between −3e to +3e.
[edit] Coordinate covalent bond
Main article: Coordinate covalent bond
Coordinate covalent bonding, sometimes referred to as dative bonding, is a kind of covalent bonding, in which the covalent bonding electrons originate solely from one of the atoms, the electron-pair donor or Lewis base but are approximately equally shared in the formation of a covalent bond. This concept is somewhat fading as chemists increasingly embrace molecular orbital theory. Examples of coordinate covalent bonding occur in nitrones and ammonia borane. The arrangement is different from an ionic bond in that the electronegativity difference is small, resulting in covalency. They are shown by an arrow unlike other bonds. This arrow shows its head towards the electron acceptor or lewis acid and its tail towards the lewis base. This bond is seen in ammonium.
[edit] Bent bonds
Main article: Bent bond
Bent bonds, also known as banana bonds, are bonds in strained or otherwise sterically hindered molecules those binding orbitals are forced into a banana-like form. Bent bonds are often more susceptible to reactions than ordinary bonds.
[edit] Three- and four-center bonds
Main articles: three-center two-electron bond, three-center four-electron bond, and four-center two-electron bond
In three-center two-electron bonds three atoms share two electrons in bonding. This type of bonding occurs in electron deficient compounds like diborane. Each such bond (2 per molecule in diborane) contains a pair of electrons which connect the boron atoms to each other in a banana shape (shown as a more sharply angled section in the stick model at right), with a proton (nucleus of a hydrogen atom) in the middle of the bond, sharing electrons with both boron atoms.
Three-center four-electron bonds are used to explain the bonding in hypervalent molecules. In certain cluster compounds so-called four-center two-electron bonds also have been postulated. These bonds are not to be confused with "two-center four-electron bonds", which is one way of describing conventional double bonds between two atoms. The MO description of such double bonds would distinguish one bond, with electron density concentrated along the line between atoms, as a σ bond, but the other bond, with electron density concentrated in lobes on either side of the axial line, as a π (pi) bond.
[edit] One- and three-electron bonds
Bonds with one or three electrons can be found in radical species, which have an odd number of electrons. The simplest example of a 1-electron bond is found in the hydrogen molecular cation, H2+. One-electron bonds often have about half the bond energy of a 2-electron bond, and are therefore called "half bonds". However, there are exceptions: in the case of dilithium, the bond is actually stronger for the 1-electron Li2+ than for the 2-electron Li2. This exception can be explained in terms of hybridization and inner-shell effects. [5]
The simplest example of three-electron bonding can be found in the helium dimer cation, He2+, and can also be considered a "half bond" because, in molecular orbital terms, the third electron is in an anti-bonding orbital which cancels out half of the bond formed by the other two electrons. Another example of a molecule containing a 3-electron bond, in addition to two 2-electron bonds, is nitric oxide, NO. The oxygen molecule, O2 can also be regarded as having two 3-electron bonds and one 2-electron bond, which accounts for its paramagnetism and its formal bond order of 2.[6]
Molecules with odd-electron bonds are usually highly reactive. These types of bond are only stable between atoms with similar electronegativities.[6]
[edit] Aromatic bond
Main article: Aromaticity
In most cases, the locations of electrons cannot be simplified to simple lines (place for two electrons) or dots (a single electron). In aromatic bonds which occur in planar rings of atoms where the 4n+2 rule determines whether ring molecules would show extra stability.
In benzene, the prototypical aromatic compound, 18 bonding electrons bind 6 carbon atoms together to form a planar ring structure. The bond "order" (average number of bonds) between the different carbon atoms may be said to be (18/6)/2=1.5, but in this case the bonds are all identical from the chemical point of view. They may sometimes be written as single bonds alternating with double bonds, but the view of all ring bonds as being equivalently about 1.5 bonds in strength, is much closer to truth.
In the case of heterocyclic aromatics and substituted benzenes, the electronegativity differences between different parts of the ring may dominate the chemical behaviour of aromatic ring bonds, which otherwise are equivalent.
[edit] Metallic bond
Main article: Metallic bond
In metallic bonding, bonding electrons are delocalized over a lattice of atoms. By contrast, in ionic compounds, the locations of the binding electrons and their charges are static. Because of delocalization or the free moving of electrons, it leads to the metallic properties such as conductivity, ductility and hardness.
[edit] Intermolecular bonding
There are four basic types of bonds that can be formed between two or more (otherwise non-associated) molecules, ions or atoms. Intermolecular forces cause molecules to be attracted or repulsed by each other. Often, these define some of the physical characteristics (such as the melting point) of a substance.
[edit] Permanent dipole to permanent dipole
Main article: Intermolecular force
A large electronegativity difference between two strongly bonded atoms within a molecule causes a dipole to form (a dipole is a pair of permanent partial charges). Dipoles will attract or repel each other.
[edit] Hydrogen bond
Main article: Hydrogen bond
In some ways this is an especially strong example of a permanent dipole, as above. However, in the hydrogen bond, the hydrogen proton comes closer to being shared between target and donor atoms, in a three-center two-electron bond like that in diborane. Hydrogen bonds explain the relatively high boiling points of liquids like water, ammonia, and hydrogen fluoride, compared with their heavier counterparts in the same periodic table column.
[edit] Instantaneous dipole to induced dipole (van der Waals)
Main article: van der Waals forces
Instantaneous dipole to induced dipole, or van der Waals forces, are the weakest, but also the most prolific—occurring between all chemical substances. Imagine a helium atom: At any one point in time, the electron cloud around the (otherwise neutral) atom can be thought to be slightly imbalanced, with momentarily more negative charge on one side. This is referred to as an instantaneous dipole. This dipole, with its slight charge imbalance, may attract or repel the electrons within a neighbouring helium atom, setting up another dipole. The two atoms will be attracted for an instant, before the charge rebalances and the atoms move on.
[edit] Cation-pi interaction
Main article: Cation-pi interaction
Cation-pi interactions occur between the localized negative charge of π orbital electrons, located above and below the plane of an aromatic ring, and a positive charge.
[edit] Electrons in chemical bonds
Many simple compounds involve covalent bonds. These molecules have structures that can be predicted using valence bond theory, and the properties of atoms involved can be understood using concepts such as oxidation number. Other compounds that involve ionic structures can be understood using theories from classical physics.
In the case of ionic bonding, electrons are mainly localized on the individual atoms, and electrons do not travel between the atoms very much. Each atom is assigned an overall electric charge to help conceptualize the molecular orbital's distribution. The forces between atoms (or ions) are largely characterized by isotropic continuum electrostatic potentials.
By contrast, in covalent bonding, the electron density within a bond is not assigned to individual atoms, but is instead delocalized in the MOs between atoms. The widely accepted theory of the linear combination of atomic orbitals (LCAO) helps describe the molecular orbital structures and energies based on the atomic orbitals of the atoms they came from. Unlike pure ionic bonds, covalent bonds may have directed anisotropic properties. These may have their own names, too, such as Sigma and Pi bond.
Atoms can also form bonds that are intermediates between ionic and covalent. This is because these definitions are based on the extent of electron delocalization. Electrons can be partially delocalized between atoms, but spend more time around one atom than another. This type of bond is often called polar covalent. See electronegativity.
Thus, the electrons in a molecular orbital (or 'in a polar covalent, or in a covalent bond') can be said to be either localized on certain atom(s) or delocalized between two or more atoms. The type of bond between two atoms is defined by how much the electron density is localized or delocalized among the atoms of the bonds.
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