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Carbon, 6C

Graphite-and-diamond-with-scale.jpg

Graphite (left) and diamond (right), two allotropes of carbon

Carbon
Allotropes graphite, diamond and more (see Allotropes of carbon)
Appearance
  • graphite: black, metallic-looking
  • diamond: clear
Standard atomic weight Ar°(C)
  • [12.009612.0116]
  • 12.011±0.002 (abridged)[1]
Carbon in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson


C

Si
boron ← carbon → nitrogen
Atomic number (Z) 6
Group group 14 (carbon group)
Period period 2
Block   p-block
Electron configuration [He] 2s2 2p2
Electrons per shell 2, 4
Physical properties
Phase at STP solid
Sublimation point 3915 K ​(3642 °C, ​6588 °F)
Density (near r.t.) amorphous: 1.8–2.1 g/cm3[2]
graphite: 2.267 g/cm3
diamond: 3.515 g/cm3
Triple point 4600 K, ​10,800 kPa[3][4]
Heat of fusion graphite: 117 kJ/mol
Molar heat capacity graphite: 8.517 J/(mol·K)
diamond: 6.155 J/(mol·K)
Atomic properties
Oxidation states −4, −3, −2, −1, 0, +1,[5] +2, +3,[6] +4[7] (a mildly acidic oxide)
Electronegativity Pauling scale: 2.55
Ionization energies
  • 1st: 1086.5 kJ/mol
  • 2nd: 2352.6 kJ/mol
  • 3rd: 4620.5 kJ/mol
  • (more)
Covalent radius sp3: 77 pm
sp2: 73 pm
sp: 69 pm
Van der Waals radius 170 pm

Color lines in a spectral range

Spectral lines of carbon

Other properties
Natural occurrence primordial
Crystal structure graphite: ​simple hexagonal

Simple hexagonal crystal structure for graphite: carbon

(black)

Crystal structure diamond: ​face-centered diamond-cubic

Diamond cubic crystal structure for diamond: carbon

(clear)

Speed of sound thin rod diamond: 18,350 m/s (at 20 °C)
Thermal expansion diamond: 0.8 µm/(m⋅K) (at 25 °C)[8]
Thermal conductivity graphite: 119–165 W/(m⋅K)
diamond: 900–2300 W/(m⋅K)
Electrical resistivity graphite: 7.837 µΩ⋅m[9]
Magnetic ordering diamagnetic[10]
Molar magnetic susceptibility diamond: −5.9×10−6 cm3/mol[11]
Young’s modulus diamond: 1050 GPa[8]
Shear modulus diamond: 478 GPa[8]
Bulk modulus diamond: 442 GPa[8]
Poisson ratio diamond: 0.1[8]
Mohs hardness graphite: 1–2
diamond: 10
CAS Number
  • atomic carbon: 7440-44-0
  • graphite: 7782-42-5
  • diamond: 7782-40-3
History
Discovery Egyptians and Sumerians[12] (3750 BCE)
Recognized as an element by Antoine Lavoisier[13] (1789)
Main isotopes of carbon

  • v
  • e

Iso­tope Decay
abun­dance half-life (t1/2) mode pro­duct
11C syn 20.3402(53) min β+ 11B
12C [98.84%99.04%] stable
13C [0.96%1.16%] stable
14C 1 ppt 5.70(3)×103 y β 14N
 Category: Carbon

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Carbon (from Latin carbo ‘coal’) is a chemical element with the symbol C and atomic number 6. It is nonmetallic and tetravalent—its atom making four electrons available to form covalent chemical bonds. It belongs to group 14 of the periodic table.[14] Carbon makes up about 0.025 percent of Earth’s crust.[15] Three isotopes occur naturally, 12C and 13C being stable, while 14C is a radionuclide, decaying with a half-life of about 5,730 years.[16] Carbon is one of the few elements known since antiquity.[17]

Carbon is the 15th most abundant element in the Earth’s crust, and the fourth most abundant element in the universe by mass after hydrogen, helium, and oxygen. Carbon’s abundance, its unique diversity of organic compounds, and its unusual ability to form polymers at the temperatures commonly encountered on Earth, enables this element to serve as a common element of all known life. It is the second most abundant element in the human body by mass (about 18.5%) after oxygen.[18]

The atoms of carbon can bond together in diverse ways, resulting in various allotropes of carbon. Well-known allotropes include graphite, diamond, amorphous carbon and fullerenes. The physical properties of carbon vary widely with the allotropic form. For example, graphite is opaque and black while diamond is highly transparent. Graphite is soft enough to form a streak on paper (hence its name, from the Greek verb «γράφειν» which means «to write»), while diamond is the hardest naturally occurring material known. Graphite is a good electrical conductor while diamond has a low electrical conductivity. Under normal conditions, diamond, carbon nanotubes, and graphene have the highest thermal conductivities of all known materials. All carbon allotropes are solids under normal conditions, with graphite being the most thermodynamically stable form at standard temperature and pressure. They are chemically resistant and require 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 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 a vast number of compounds, with almost ten million compounds described to date,[19] and yet that number is but a fraction of the number of theoretically possible compounds under standard conditions.

Characteristics

Theoretically predicted phase diagram of carbon, from 1989. Newer work indicates that the melting point of diamond (top-right curve) does not go above about 9000 K.[20]

The allotropes of carbon include graphite, one of the softest known substances, and diamond, the hardest naturally occurring substance. It bonds readily with other small atoms, including other carbon atoms, and is capable of forming multiple stable covalent bonds with suitable multivalent atoms. Carbon is known to form almost ten million compounds, a large majority of all chemical compounds.[19] Carbon also has the highest sublimation point of all elements. At atmospheric pressure it has no melting point, as its triple point is at 10.8 ± 0.2 megapascals (106.6 ± 2.0 atm; 1,566 ± 29 psi) and 4,600 ± 300 K (4,330 ± 300 °C; 7,820 ± 540 °F),[3][4] so it sublimes at about 3,900 K (3,630 °C; 6,560 °F).[21][22] Graphite is much more reactive than diamond at standard conditions, despite being more thermodynamically stable, as its delocalised pi system is much more vulnerable to attack. For example, graphite can be oxidised by hot concentrated nitric acid at standard conditions to mellitic acid, C6(CO2H)6, which preserves the hexagonal units of graphite while breaking up the larger structure.[23]

Carbon sublimes in a carbon arc, which has a temperature of about 5800 K (5,530 °C or 9,980 °F). 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, which are weaker reducing agents at room temperature.

Carbon is the sixth element, with a ground-state electron configuration of 1s22s22p2, of which the four outer electrons are valence electrons. Its first four ionisation energies, 1086.5, 2352.6, 4620.5 and 6222.7 kJ/mol, are much higher than those of the heavier group-14 elements. The electronegativity of carbon is 2.5, significantly higher than the heavier group-14 elements (1.8–1.9), but close to most of the nearby nonmetals, as well as some of the second- and third-row transition metals. Carbon’s covalent radii are normally taken as 77.2 pm (C−C), 66.7 pm (C=C) and 60.3 pm (C≡C), although these may vary depending on coordination number and what the carbon is bonded to. In general, covalent radius decreases with lower coordination number and higher bond order.[24]

Carbon-based compounds form the basis of all known 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 rob oxygen from metal oxides to leave the elemental metal. This exothermic reaction is used in the iron and steel industry to smelt iron and to control the carbon content of steel:

Fe
3
O
4
+ 4 C(s) + 2 O
2
→ 3 Fe(s) + 4 CO
2
(g).

Carbon reacts with sulfur to form carbon disulfide, and it reacts with steam in the coal-gas reaction used in coal gasification:

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:

Graphite is one of the softest materials known. Synthetic nanocrystalline diamond is the hardest material known.[25]
Graphite is a very good lubricant, displaying superlubricity.[26] Diamond is the ultimate abrasive.
Graphite is a conductor of electricity.[27] Diamond is an excellent electrical insulator,[28] and has the highest breakdown electric field of any known material.
Some forms of graphite are used for thermal insulation (i.e. firebreaks and heat shields), but some other forms are good thermal conductors. Diamond is the best known naturally occurring thermal conductor
Graphite is opaque. Diamond is highly transparent.
Graphite crystallizes in the hexagonal system.[29] Diamond crystallizes in the cubic system.
Amorphous carbon is completely isotropic. Carbon nanotubes are among the most anisotropic materials known.

Allotropes

Atomic carbon is a very short-lived species and, therefore, carbon is stabilized in various multi-atomic structures with diverse 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,[30][31] carbon nanotubes,[32] carbon nanobuds[33] and nanofibers.[34][35] Several other exotic allotropes have also been discovered, such as lonsdaleite,[36] glassy carbon,[37] carbon nanofoam[38] and linear acetylenic carbon (carbyne).[39]

Graphene is a two-dimensional sheet of carbon with the atoms arranged in a hexagonal lattice. As of 2009, graphene appears to be the strongest material ever tested.[40] The process of separating it from graphite will require some further technological development before it is economical for industrial processes.[41] If successful, graphene could be used in the construction of a space elevator. It could also be used to safely store hydrogen for use in a hydrogen based engine in cars.[42]

A large sample of glassy carbon

The amorphous form is an assortment of carbon atoms in a non-crystalline, irregular, glassy state, 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.[43] 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.

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, forming a 3-dimensional network of puckered six-membered rings of atoms. Diamond has the same cubic structure as silicon and germanium, and because of the strength of the carbon-carbon bonds, it is the hardest naturally occurring substance measured by resistance to scratching. Contrary to the popular belief that «diamonds are forever», they are thermodynamically unstable (ΔfG°(diamond, 298 K) = 2.9 kJ/mol[44]) under normal conditions (298 K, 105 Pa) and should theoretically transform into graphite.[45] But due to a high activation energy barrier, the transition into graphite is so slow at normal temperature that it is unnoticeable. However, at very high temperatures diamond will turn into graphite, and diamonds can burn up in a house fire. The bottom left corner of the phase diagram for carbon has not been scrutinized experimentally. Although a computational study employing density functional theory methods reached the conclusion that as T → 0 K and p → 0 Pa, diamond becomes more stable than graphite by approximately 1.1 kJ/mol,[46] more recent and definitive experimental and computational studies show that graphite is more stable than diamond for T < 400 K, without applied pressure, by 2.7 kJ/mol at T = 0 K and 3.2 kJ/mol at T = 298.15 K.[47] Under some conditions, carbon crystallizes as lonsdaleite, a hexagonal crystal lattice with all atoms covalently bonded and properties similar to those of diamond.[36]

Fullerenes are a synthetic crystalline formation with a graphite-like structure, but in place of flat hexagonal cells only, some of the cells of which fullerenes are formed may be pentagons, nonplanar hexagons, or even heptagons of carbon atoms. The sheets are thus warped into spheres, ellipses, or cylinders. The properties of fullerenes (split into buckyballs, buckytubes, and nanobuds) have not yet been fully analyzed and represent an intense area of research in nanomaterials. The names fullerene and buckyball are given after Richard Buckminster Fuller, popularizer of geodesic domes, 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 C60 buckminsterfullerene).[30] Carbon nanotubes (buckytubes) are structurally similar to buckyballs, except that each atom is bonded trigonally in a curved sheet that forms a hollow cylinder.[31][32] Nanobuds were first reported in 2007 and are hybrid buckytube/buckyball materials (buckyballs are covalently bonded to the outer wall of a nanotube) that combine the properties of both in a single structure.[33]

Of the other discovered allotropes, 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.[48] Similarly, glassy carbon contains a high proportion of closed porosity,[37] but contrary to normal graphite, the graphitic layers are not stacked like pages in a book, but have a more random arrangement. Linear acetylenic carbon[39] has the chemical structure[39] −(C:::C)n−. Carbon in this modification is linear with sp orbital hybridization, and is a polymer with alternating single and triple bonds. This carbyne is of considerable interest to nanotechnology as its Young’s modulus is 40 times that of the hardest known material – diamond.[49]

In 2015, a team at the North Carolina State University announced the development of another allotrope they have dubbed Q-carbon, created by a high energy low duration laser pulse on amorphous carbon dust. Q-carbon is reported to exhibit ferromagnetism, fluorescence, and a hardness superior to diamonds.[50]

In the vapor phase, some of the carbon is in the form of dicarbon (C
2
). When excited, this gas glows green.

Occurrence

Graphite ore, shown with a penny for scale

Carbon is the fourth most abundant chemical element in the observable universe by mass after hydrogen, helium, and oxygen. Carbon is abundant in the Sun, stars, comets, and in the atmospheres of most planets.[51] Some meteorites contain microscopic diamonds that were formed when the Solar System was still a protoplanetary disk.[52] Microscopic diamonds may also be formed by the intense pressure and high temperature at the sites of meteorite impacts.[53]

In 2014 NASA announced a greatly upgraded database for tracking polycyclic aromatic hydrocarbons (PAHs) in the universe. More than 20% of the carbon in the universe may be associated with PAHs, complex compounds of carbon and hydrogen without oxygen.[54] These compounds figure in the PAH world hypothesis where they are hypothesized to have a role in abiogenesis and formation of life. PAHs seem to have been formed «a couple of billion years» after the Big Bang, are widespread throughout the universe, and are associated with new stars and exoplanets.[51]

It has been estimated that the solid earth as a whole contains 730 ppm of carbon, with 2000 ppm in the core and 120 ppm in the combined mantle and crust.[55] Since the mass of the earth is 5.972×1024 kg, this would imply 4360 million gigatonnes of carbon. This is much more than the amount of carbon in the oceans or atmosphere (below).

In combination with oxygen in carbon dioxide, carbon is found in the Earth’s atmosphere (approximately 900 gigatonnes of carbon — each ppm corresponds to 2.13 Gt) and dissolved in all water bodies (approximately 36,000 gigatonnes of carbon). Carbon in the biosphere has been estimated at 550 gigatonnes but with a large uncertainty, due mostly to a huge uncertainty in the amount of terrestrial deep subsurface bacteria.[56] Hydrocarbons (such as coal, petroleum, and natural gas) contain carbon as well. Coal «reserves» (not «resources») amount to around 900 gigatonnes with perhaps 18,000 Gt of resources.[57] Oil reserves are around 150 gigatonnes. Proven sources of natural gas are about 175×1012 cubic metres (containing about 105 gigatonnes of carbon), but studies estimate another 900×1012 cubic metres of «unconventional» deposits such as shale gas, representing about 540 gigatonnes of carbon.[58]

Carbon is also found in methane hydrates in polar regions and under the seas. Various estimates put this carbon between 500, 2500 Gt,[59] or 3,000 Gt.[60]

In the past, quantities of hydrocarbons were greater. According to one source, in the period from 1751 to 2008 about 347 gigatonnes of carbon were released as carbon dioxide to the atmosphere from burning of fossil fuels.[61] Another source puts the amount added to the atmosphere for the period since 1750 at 879 Gt, and the total going to the atmosphere, sea, and land (such as peat bogs) at almost 2,000 Gt.[62]

Carbon is a constituent (about 12% by mass) of the very large masses of carbonate rock (limestone, dolomite, marble and so on). Coal is very rich in carbon (anthracite contains 92–98%)[63] and is the largest commercial source of mineral carbon, accounting for 4,000 gigatonnes or 80% of fossil fuel.[64]

As for individual carbon allotropes, graphite is found in large quantities in the United States (mostly in New York and Texas), 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. Diamond deposits have also been found 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. Diamonds are found naturally, but about 30% of all industrial diamonds used in the U.S. are now manufactured.

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.[65] Thermal neutrons are produced that collide with the nuclei of nitrogen-14, forming carbon-14 and a proton. As such, 1.5%×10−10 of atmospheric carbon dioxide contains carbon-14.[66]

Carbon-rich asteroids are relatively preponderant in the outer parts of the asteroid belt in the Solar System. These asteroids have not yet been directly sampled by scientists. The asteroids can be used in hypothetical space-based carbon mining, which may be possible in the future, but is currently technologically impossible.[67]

Isotopes

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.[16] The isotope carbon-12 (12C) forms 98.93% of the carbon on Earth, while carbon-13 (13C) forms the remaining 1.07%.[16] The concentration of 12C is further increased in biological materials because biochemical reactions discriminate against 13C.[68] In 1961, the International Union of Pure and Applied Chemistry (IUPAC) adopted the isotope carbon-12 as the basis for atomic weights.[69] Identification of carbon in nuclear magnetic resonance (NMR) experiments is done with the isotope 13C.

Carbon-14 (14C) is a naturally occurring radioisotope, created in the upper atmosphere (lower stratosphere and upper troposphere) by interaction of nitrogen with cosmic rays.[70] It is found in trace amounts on Earth of 1 part per trillion (0.0000000001%) or more, mostly confined to the atmosphere and superficial deposits, particularly of peat and other organic materials.[71] 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. The amount 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.[72][73]

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 1.98739 × 10−21 s.[74] The exotic 19C exhibits a nuclear halo, which means its radius is appreciably larger than would be expected if the nucleus were a sphere of constant density.[75]

Formation in stars

Formation of the carbon atomic nucleus occurs within a giant or supergiant star through the triple-alpha process. This requires a nearly simultaneous collision of three alpha particles (helium nuclei), as the products of further nuclear fusion reactions of helium with hydrogen or another helium nucleus produce lithium-5 and beryllium-8 respectively, both of which are highly unstable and decay almost instantly back into smaller nuclei.[76] The triple-alpha process happens in conditions of temperatures over 100 megakelvins 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.

According to current physical cosmology theory, carbon is formed in the interiors of stars on the horizontal branch.[77] When massive stars die as supernova, the carbon is scattered into space as dust. This dust becomes component material for the formation of the next-generation star systems with accreted planets.[51][78] The Solar System is one such star system with an abundance of carbon, enabling the existence of life as we know it.

The CNO cycle is an additional hydrogen fusion mechanism that powers stars, wherein carbon operates as a catalyst.

Rotational transitions of various isotopic forms of carbon monoxide (for example, 12CO, 13CO, and 18CO) are detectable in the submillimeter wavelength range, and are used in the study of newly forming stars in molecular clouds.[79]

Carbon cycle

Diagram of the carbon cycle. The black numbers indicate how much carbon is stored in various reservoirs, in billions tonnes («GtC» stands for gigatonnes 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 from somewhere and dispose of it somewhere else. The paths of carbon in the environment form the carbon cycle.[80] For example, photosynthetic plants draw carbon dioxide from the atmosphere (or seawater) and build it into biomass, as in the Calvin cycle, a process of carbon fixation.[81] Some of this biomass is eaten by animals, while 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; if bacteria do not consume it, dead plant or animal matter may become petroleum or coal, which releases carbon when burned.[82][83]

Compounds

Organic compounds

Structural formula of methane, the simplest possible organic compound.

Correlation between the carbon cycle and formation of organic compounds. In plants, carbon dioxide formed by carbon fixation can join with water in photosynthesis (green) to form organic compounds, which can be used and further converted by both plants and animals.

Carbon can form very long chains of interconnecting carbon–carbon bonds, a property that is called catenation. Carbon-carbon bonds are strong and stable. Through catenation, carbon forms a countless number of compounds. A tally of unique compounds shows that more contain carbon than do not.[84] A similar claim can be made for hydrogen because most organic compounds contain hydrogen chemically bonded to carbon or another common element like oxygen or nitrogen.

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. A hydrocarbon backbone can be substituted by other atoms, known as heteroatoms. Common heteroatoms that appear in organic compounds include oxygen, nitrogen, sulfur, phosphorus, and the nonradioactive halogens, as well as the metals lithium and magnesium. Organic compounds containing bonds to metal are known as organometallic compounds (see below). Certain groupings of atoms, often including heteroatoms, recur in large numbers of organic compounds. These collections, known as functional groups, confer common reactivity patterns and allow for the systematic study and categorization of organic compounds. Chain length, shape and functional groups all affect the properties of organic molecules.[85]

In most stable compounds of carbon (and nearly all stable organic compounds), carbon obeys the octet rule and is tetravalent, meaning that a carbon atom forms a total of four covalent bonds (which may include double and triple bonds). Exceptions include a small number of stabilized carbocations (three bonds, positive charge), radicals (three bonds, neutral), carbanions (three bonds, negative charge) and carbenes (two bonds, neutral), although these species are much more likely to be encountered as unstable, reactive intermediates.

Carbon occurs in all known organic life and is the basis of organic chemistry. When united with hydrogen, it forms various hydrocarbons that are important to industry as refrigerants, lubricants, solvents, as chemical feedstock for the manufacture of plastics and petrochemicals, and as fossil fuels.

When combined with oxygen and hydrogen, carbon can form many groups of important biological compounds including sugars, 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 rubber products. With the addition of phosphorus to these other elements, it forms DNA and RNA, the chemical-code carriers of life, and adenosine triphosphate (ATP), the most important energy-transfer molecule in all living cells.[86]

Inorganic compounds

Commonly carbon-containing compounds which are associated with minerals or which do not contain bonds to the other carbon atoms, halogens, or hydrogen, are treated separately from classical organic compounds; the definition is not rigid, and the classification of some compounds can vary from author to author (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.[87] Dissolved in water, it forms carbonic acid (H
2
CO
3
), but as most compounds with multiple single-bonded oxygens on a single carbon it is unstable.[88] Through this intermediate, though, resonance-stabilized carbonate ions are produced. Some important minerals are carbonates, notably calcite. Carbon disulfide (CS
2
) is similar.[23] Nevertheless, due to its physical properties and its association with organic synthesis, carbon disulfide is sometimes classified as an organic solvent.

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.[89][90] 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. Likewise, the heavier analog of cyanide, cyaphide (CP), is also considered inorganic, though most simple derivatives are highly unstable. Other uncommon oxides are carbon suboxide (C
3
O
2
),[91] the unstable dicarbon monoxide (C2O),[92][93] carbon trioxide (CO3),[94][95] cyclopentanepentone (C5O5),[96] cyclohexanehexone (C6O6),[96] and mellitic anhydride (C12O9). However, mellitic anhydride is the triple acyl anhydride of mellitic acid; moreover, it contains a benzene ring. Thus, many chemists consider it to be organic.

With reactive metals, such as tungsten, carbon forms either carbides (C4−) or acetylides (C2−
2
) 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,[97] carbon prefers to form covalent bonds. A few carbides are covalent lattices, like carborundum (SiC), which resembles diamond. Nevertheless, even the most polar and salt-like of carbides are not completely ionic compounds.[98]

Organometallic compounds

Organometallic compounds by definition contain at least one carbon-metal covalent bond. A wide range of such compounds exist; major classes include simple alkyl-metal compounds (for example, tetraethyllead), η2-alkene compounds (for example, Zeise’s salt), and η3-allyl compounds (for example, allylpalladium chloride dimer); metallocenes containing cyclopentadienyl ligands (for example, ferrocene); and transition metal carbene complexes. Many metal carbonyls and metal cyanides exist (for example, tetracarbonylnickel and potassium ferricyanide); some workers consider metal carbonyl and cyanide complexes without other carbon ligands to be purely inorganic, and not organometallic. However, most organometallic chemists consider metal complexes with any carbon ligand, even ‘inorganic carbon’ (e.g., carbonyls, cyanides, and certain types of carbides and acetylides) to be organometallic in nature. Metal complexes containing organic ligands without a carbon-metal covalent bond (e.g., metal carboxylates) are termed metalorganic compounds.

While carbon is understood to strongly prefer formation of four covalent bonds, other exotic bonding schemes are also known. Carboranes are highly stable dodecahedral derivatives of the [B12H12]2- unit, with one BH replaced with a CH+. Thus, the carbon is bonded to five boron atoms and one hydrogen atom. The cation [(Ph3PAu)6C]2+ contains an octahedral carbon bound to six phosphine-gold fragments. This phenomenon has been attributed to the aurophilicity of the gold ligands, which provide additional stabilization of an otherwise labile species.[99] In nature, the iron-molybdenum cofactor (FeMoco) responsible for microbial nitrogen fixation likewise has an octahedral carbon center (formally a carbide, C(-IV)) bonded to six iron atoms. In 2016, it was confirmed that, in line with earlier theoretical predictions, the hexamethylbenzene dication contains a carbon atom with six bonds. More specifically, the dication could be described structurally by the formulation [MeC(η5-C5Me5)]2+, making it an «organic metallocene» in which a MeC3+ fragment is bonded to a η5-C5Me5 fragment through all five of the carbons of the ring.[100]

This anthracene derivative contains a carbon atom with 5 formal electron pairs around it.

It is important to note that in the cases above, each of the bonds to carbon contain less than two formal electron pairs. Thus, the formal electron count of these species does not exceed an octet. This makes them hypercoordinate but not hypervalent. Even in cases of alleged 10-C-5 species (that is, a carbon with five ligands and a formal electron count of ten), as reported by Akiba and co-workers,[101] electronic structure calculations conclude that the electron population around carbon is still less than eight, as is true for other compounds featuring four-electron three-center bonding.

History and etymology

The English name carbon comes from the Latin carbo for coal and charcoal,[102] whence also comes the French charbon, meaning charcoal. In German, Dutch and Danish, the names for carbon are Kohlenstoff, koolstof and kulstof respectively, all literally meaning coal-substance.

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 form 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.[103][104]

In 1722, René Antoine Ferchault de Réaumur demonstrated that iron was transformed into steel through the absorption of some substance, now known to be carbon.[105] In 1772, Antoine Lavoisier showed that diamonds are a form of carbon; when he burned samples of charcoal and diamond and found that neither produced any water and that both released the same amount of carbon dioxide per gram.
In 1779,[106] Carl Wilhelm Scheele showed that graphite, which had been thought of as a form of lead, was instead identical with charcoal but with a small admixture of iron, and that it gave «aerial acid» (his name for carbon dioxide) when oxidized with nitric acid.[107] In 1786, the French scientists Claude Louis Berthollet, Gaspard Monge and C. A. Vandermonde confirmed that graphite was mostly carbon by oxidizing it in oxygen in much the same way Lavoisier had done with diamond.[108] Some iron again was left, which the French scientists thought was necessary to the graphite structure. In their publication they proposed the name carbone (Latin carbonum) for the element in graphite which was given off as a gas upon burning graphite. Antoine Lavoisier then listed carbon as an element in his 1789 textbook.[107]

A new allotrope of carbon, fullerene, that was discovered in 1985[109] includes nanostructured forms such as buckyballs and nanotubes.[30] Their discoverers – Robert Curl, Harold Kroto and Richard Smalley – received the Nobel Prize in Chemistry in 1996.[110] The resulting renewed interest in new forms led to the discovery of further exotic allotropes, including glassy carbon, and the realization that «amorphous carbon» is not strictly amorphous.[37]

Production

Graphite

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.[citation needed] 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 19th century, pencils were made 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.[111]

There are three types of natural graphite—amorphous, flake or crystalline flake, and vein or lump. Amorphous graphite is the lowest quality and most abundant. Contrary to science, in industry «amorphous» refers to very small crystal size rather than complete lack of crystal structure. Amorphous is used for lower value graphite products and is the lowest priced graphite. Large amorphous graphite deposits are found in China, Europe, Mexico and the United States. Flake graphite is less common and of higher quality than amorphous; it occurs as separate plates that crystallized in metamorphic rock. Flake graphite can be four times the price of amorphous. Good quality flakes can be processed into expandable graphite for many uses, such as flame retardants. The foremost deposits are found in Austria, Brazil, Canada, China, Germany and Madagascar. Vein or lump graphite is the rarest, most valuable, and highest quality type of natural graphite. It occurs in veins along intrusive contacts in solid lumps, and it is only commercially mined in Sri Lanka.[111]

According to the USGS, world production of natural graphite was 1.1 million tonnes in 2010, to which China contributed 800,000 t, India 130,000 t, Brazil 76,000 t, North Korea 30,000 t and Canada 25,000 t. No natural graphite was reported mined in the United States, but 118,000 t of synthetic graphite with an estimated value of $998 million was produced in 2009.[111]

Diamond

The diamond supply chain is controlled by a limited number of powerful businesses, and is also highly concentrated in a small number of locations around the world (see figure).

Only a very small fraction of the diamond ore consists of actual diamonds. The ore is crushed, during which care has to be taken in order to prevent larger diamonds from being destroyed in this process and subsequently the particles are sorted by density. Today, diamonds are located in the diamond-rich density fraction with the help of X-ray fluorescence, after which the final sorting steps are done by hand. Before the use of X-rays became commonplace, the separation was done with grease belts; diamonds have a stronger tendency to stick to grease than the other minerals in the ore.[112]

Historically diamonds were known to be found only in alluvial deposits in southern India.[113] India led the world in diamond production from the time of their discovery in approximately the 9th century BC[114] to the mid-18th century AD, but the commercial potential of these sources had been exhausted by the late 18th century and at that time India was eclipsed by Brazil where the first non-Indian diamonds were found in 1725.[115]

Diamond production of primary deposits (kimberlites and lamproites) only started in the 1870s after the discovery of the diamond fields in South Africa. Production has increased over time and an accumulated total of over 4.5 billion carats have been mined since that date.[116] Most commercially viable diamond deposits were in Russia, Botswana, Australia and the Democratic Republic of Congo.[117] By 2005, Russia produced almost one-fifth of the global diamond output (mostly in Yakutia territory; for example, Mir pipe and Udachnaya pipe) but the Argyle mine in Australia became the single largest source, producing 14 million carats in 2018.[118][119] New finds, the Canadian mines at Diavik and Ekati, are expected to become even more valuable owing to their production of gem quality stones.[120]

In the United States, diamonds have been found in Arkansas, Colorado and Montana.[121] In 2004, a startling discovery of a microscopic diamond in the United States[122] led to the January 2008 bulk-sampling of kimberlite pipes in a remote part of Montana.[123]

Applications

Pencil leads for mechanical pencils are made of graphite (often mixed with a clay or synthetic binder).

A cloth of woven carbon fibres

The C60 fullerene in crystalline form

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 distilled in refineries by the petrochemical industry to produce gasoline, kerosene, and other products. Cellulose is a natural, carbon-containing polymer produced by plants in the form of wood, cotton, linen, and hemp. Cellulose is used primarily for maintaining structure in plants. 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 molding 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, barbecue grilling, iron smelting, and in many other applications. Wood, coal and oil are used as fuel for production of energy and 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 fiber, made by pyrolysis of synthetic polyester fibers 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 fibers made from PAN have structure resembling narrow filaments of graphite, but thermal processing may re-order the structure into a continuous rolled sheet. The result is fibers with higher specific tensile strength than steel.[124]

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 (smelting). 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.

Diamonds

The diamond industry falls into two categories: one dealing with gem-grade diamonds and the other, with industrial-grade diamonds. While a large trade in both types of diamonds exists, the two markets function dramatically differently.

Unlike precious metals such as gold or platinum, gem diamonds do not trade as a commodity: there is a substantial mark-up in the sale of diamonds, and there is not a very active market for resale of diamonds.

Industrial diamonds are valued mostly for their hardness and heat conductivity, with the gemological qualities of clarity and color being mostly irrelevant. About 80% of mined diamonds (equal to about 100 million carats or 20 tonnes annually) are unsuitable for use as gemstones are relegated for industrial use (known as bort).[125] synthetic diamonds, invented in the 1950s, found almost immediate industrial applications; 3 billion carats (600 tonnes) of synthetic diamond is produced annually.[126]

The dominant industrial use of diamond is in cutting, drilling, grinding, and polishing. Most of these applications do not require large diamonds; in fact, most diamonds of gem-quality except for their small size can be used industrially. Diamonds are embedded in drill tips or saw blades, or ground into a powder for use in grinding and polishing applications.[127] Specialized applications include use in laboratories as containment for high pressure experiments (see diamond anvil cell), high-performance bearings, and limited use in specialized windows.[128][129] With the continuing advances in the production of synthetic diamonds, new applications are becoming feasible. Garnering much excitement is the possible use of diamond as a semiconductor suitable for microchips, and because of its exceptional heat conductance property, as a heat sink in electronics.[130]

Precautions

Pure carbon has extremely low toxicity to humans and can be handled 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. Consequently, once it enters into the body’s 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.[131] 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. Diamond dust used as an abrasive can be harmful if ingested or inhaled. Microparticles of carbon are produced in diesel engine exhaust fumes, and may accumulate in the lungs.[132] In these examples, the harm may result from contaminants (e.g., organic chemicals, heavy metals) rather than from the carbon itself.

Carbon generally has low toxicity to life on Earth; but carbon nanoparticles are deadly to Drosophila.[133]

Carbon may burn vigorously and brightly in the presence of air at high temperatures. Large accumulations of coal, which have remained inert for hundreds of millions of years in the absence of oxygen, may spontaneously combust when exposed to air in coal mine waste tips, ship cargo holds and coal bunkers,[134][135] and storage dumps.

In nuclear applications where graphite is used as a neutron moderator, accumulation of Wigner energy followed by a sudden, spontaneous release may occur. Annealing to at least 250 °C can release the energy safely, although in the Windscale fire the procedure went wrong, causing other reactor materials to combust.

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.

See also

  • Carbon chauvinism
  • Carbon detonation
  • Carbon footprint
  • Carbon star
  • Carbon planet
  • Gas carbon
  • Low-carbon economy
  • Timeline of carbon nanotubes

References

  1. ^ «Standard Atomic Weights: Carbon». CIAAW. 2009.
  2. ^ Lide, D. R., ed. (2005). CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton (FL): CRC Press. ISBN 0-8493-0486-5.
  3. ^ a b Haaland, D (1976). «Graphite-liquid-vapor triple point pressure and the density of liquid carbon». Carbon. 14 (6): 357–361. doi:10.1016/0008-6223(76)90010-5.
  4. ^ a b Savvatimskiy, A (2005). «Measurements of the melting point of graphite and the properties of liquid carbon (a review for 1963–2003)». Carbon. 43 (6): 1115–1142. doi:10.1016/j.carbon.2004.12.027.
  5. ^ «Fourier Transform Spectroscopy of the Electronic Transition of the Jet-Cooled CCI Free Radical» (PDF). Retrieved 2007-12-06.
  6. ^ «Fourier Transform Spectroscopy of the System of CP» (PDF). Retrieved 2007-12-06.
  7. ^ «Carbon: Binary compounds». Retrieved 2007-12-06.
  8. ^ a b c d e Properties of diamond, Ioffe Institute Database
  9. ^ «Material Properties- Misc Materials». www.nde-ed.org. Retrieved 12 November 2016.
  10. ^ Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics 81st edition, CRC press.
  11. ^ Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN 978-0-8493-0464-4.
  12. ^ «History of Carbon and Carbon Materials — Center for Applied Energy Research — University of Kentucky». Caer.uky.edu. Retrieved 2008-09-12.
  13. ^ Senese, Fred (2000-09-09). «Who discovered carbon?». Frostburg State University. Retrieved 2007-11-24.
  14. ^ «carbon | Facts, Uses, & Properties». Encyclopedia Britannica. Archived from the original on 2017-10-24.
  15. ^ «carbon». Britannica encyclopedia.
  16. ^ a b c «Carbon – Naturally occurring isotopes». WebElements Periodic Table. Archived from the original on 2008-09-08. Retrieved 2008-10-09.
  17. ^ «History of Carbon». Archived from the original on 2012-11-01. Retrieved 2013-01-10.
  18. ^ Reece, Jane B. (31 October 2013). Campbell Biology (10 ed.). Pearson. ISBN 9780321775658.
  19. ^ a b Chemistry Operations (December 15, 2003). «Carbon». Los Alamos National Laboratory. Archived from the original on 2008-09-13. Retrieved 2008-10-09.
  20. ^ J.H. Eggert; et al. (Nov 8, 2009). «Melting temperature of diamond at ultrahigh pressure». Nature Physics. 6: 40–43. doi:10.1038/nphys1438.
  21. ^ Greenville Whittaker, A. (1978). «The controversial carbon solid−liquid−vapour triple point». Nature. 276 (5689): 695–696. Bibcode:1978Natur.276..695W. doi:10.1038/276695a0. S2CID 4362313.
  22. ^ Zazula, J. M. (1997). «On Graphite Transformations at High Temperature and Pressure Induced by Absorption of the LHC Beam» (PDF). CERN. Archived (PDF) from the original on 2009-03-25. Retrieved 2009-06-06.
  23. ^ a b Greenwood and Earnshaw, pp. 289–292.
  24. ^ Greenwood and Earnshaw, pp. 276–8.
  25. ^ Irifune, Tetsuo; Kurio, Ayako; Sakamoto, Shizue; Inoue, Toru; Sumiya, Hitoshi (2003). «Materials: Ultrahard polycrystalline diamond from graphite». Nature. 421 (6923): 599–600. Bibcode:2003Natur.421..599I. doi:10.1038/421599b. PMID 12571587. S2CID 52856300.
  26. ^ Dienwiebel, Martin; Verhoeven, Gertjan; Pradeep, Namboodiri; Frenken, Joost; Heimberg, Jennifer; Zandbergen, Henny (2004). «Superlubricity of Graphite» (PDF). Physical Review Letters. 92 (12): 126101. Bibcode:2004PhRvL..92l6101D. doi:10.1103/PhysRevLett.92.126101. PMID 15089689. S2CID 26811802. Archived (PDF) from the original on 2011-09-17.
  27. ^ Deprez, N.; McLachan, D. S. (1988). «The analysis of the electrical conductivity of graphite conductivity of graphite powders during compaction». Journal of Physics D: Applied Physics. 21 (1): 101–107. Bibcode:1988JPhD…21..101D. doi:10.1088/0022-3727/21/1/015. S2CID 250886376.
  28. ^ Collins, A. T. (1993). «The Optical and Electronic Properties of Semiconducting Diamond». Philosophical Transactions of the Royal Society A. 342 (1664): 233–244. Bibcode:1993RSPTA.342..233C. doi:10.1098/rsta.1993.0017. S2CID 202574625.
  29. ^ Delhaes, P. (2001). Graphite and Precursors. CRC Press. ISBN 978-90-5699-228-6.
  30. ^ a b c Unwin, Peter. «Fullerenes(An Overview)». Archived from the original on 2007-12-01. Retrieved 2007-12-08.
  31. ^ a b Ebbesen, T. W., ed. (1997). Carbon nanotubes—preparation and properties. Boca Raton, Florida: CRC Press. ISBN 978-0-8493-9602-1.
  32. ^ a b Dresselhaus, M. S.; Dresselhaus, G.; Avouris, Ph., eds. (2001). Carbon nanotubes: synthesis, structures, properties and applications. Topics in Applied Physics. Vol. 80. Berlin. ISBN 978-3-540-41086-7.
  33. ^ a b Nasibulin, Albert G.; Pikhitsa, P. V.; Jiang, H.; Brown, D. P.; Krasheninnikov, A. V.; Anisimov, A. S.; Queipo, P.; Moisala, A.; et al. (2007). «A novel hybrid carbon material». Nature Nanotechnology. 2 (3): 156–161. Bibcode:2007NatNa…2..156N. doi:10.1038/nnano.2007.37. PMID 18654245. S2CID 6447122.
  34. ^ Nasibulin, A.; Anisimov, Anton S.; Pikhitsa, Peter V.; Jiang, Hua; Brown, David P.; Choi, Mansoo; Kauppinen, Esko I. (2007). «Investigations of NanoBud formation». Chemical Physics Letters. 446 (1): 109–114. Bibcode:2007CPL…446..109N. doi:10.1016/j.cplett.2007.08.050.
  35. ^ Vieira, R; Ledoux, Marc-Jacques; Pham-Huu, Cuong (2004). «Synthesis and characterisation of carbon nanofibers with macroscopic shaping formed by catalytic decomposition of C2H6/H2 over nickel catalyst». Applied Catalysis A: General. 274 (1–2): 1–8. doi:10.1016/j.apcata.2004.04.008.
  36. ^ a b Frondel, Clifford; Marvin, Ursula B. (1967). «Lonsdaleite, a new hexagonal polymorph of diamond». Nature. 214 (5088): 587–589. Bibcode:1967Natur.214..587F. doi:10.1038/214587a0. S2CID 4184812.
  37. ^ a b c Harris, PJF (2004). «Fullerene-related structure of commercial glassy carbons» (PDF). Philosophical Magazine. 84 (29): 3159–3167. Bibcode:2004PMag…84.3159H. CiteSeerX 10.1.1.359.5715. doi:10.1080/14786430410001720363. S2CID 220342075. Archived from the original (PDF) on 2012-03-19. Retrieved 2011-07-06.
  38. ^ Rode, A. V.; Hyde, S. T.; Gamaly, E. G.; Elliman, R. G.; McKenzie, D. R.; Bulcock, S. (1999). «Structural analysis of a carbon foam formed by high pulse-rate laser ablation». Applied Physics A: Materials Science & Processing. 69 (7): S755–S758. Bibcode:1999ApPhA..69S.755R. doi:10.1007/s003390051522. S2CID 96050247.
  39. ^ a b c Heimann, Robert Bertram; Evsyukov, Sergey E. & Kavan, Ladislav (28 February 1999). Carbyne and carbynoid structures. Springer. pp. 1–. ISBN 978-0-7923-5323-2. Archived from the original on 23 November 2012. Retrieved 2011-06-06.
  40. ^ Lee, C.; Wei, X.; Kysar, J. W.; Hone, J. (2008). «Measurement of the Elastic Properties and Intrinsic Strength of Monolayer Graphene». Science. 321 (5887): 385–8. Bibcode:2008Sci…321..385L. doi:10.1126/science.1157996. PMID 18635798. S2CID 206512830.
    • Phil Schewe (July 28, 2008). «World’s Strongest Material». Inside Science News Service (Press release). Archived from the original on 2009-05-31.

  41. ^ Sanderson, Bill (2008-08-25). «Toughest Stuff Known to Man : Discovery Opens Door to Space Elevator». nypost.com. Archived from the original on 2008-09-06. Retrieved 2008-10-09.
  42. ^ Jin, Zhong; Lu, Wei; O’Neill, Kevin J.; Parilla, Philip A.; Simpson, Lin J.; Kittrell, Carter; Tour, James M. (2011-02-22). «Nano-Engineered Spacing in Graphene Sheets for Hydrogen Storage». Chemistry of Materials. 23 (4): 923–925. doi:10.1021/cm1025188. ISSN 0897-4756.
  43. ^ Jenkins, Edgar (1973). The polymorphism of elements and compounds. Taylor & Francis. p. 30. ISBN 978-0-423-87500-3. Archived from the original on 2012-11-23. Retrieved 2011-05-01.
  44. ^ Rossini, F. D.; Jessup, R. S. (1938). «Heat and Free Energy of Formation of Carbon Dioxide and of the Transition Between Graphite and Diamond». Journal of Research of the National Bureau of Standards. 21 (4): 491. doi:10.6028/jres.021.028.
  45. ^ «World of Carbon – Interactive Nano-visulisation in Science & Engineering Education (IN-VSEE)». Archived from the original on 2001-05-31. Retrieved 2008-10-09.
  46. ^ Grochala, Wojciech (2014-04-01). «Diamond: Electronic Ground State of Carbon at Temperatures Approaching 0 K». Angewandte Chemie International Edition. 53 (14): 3680–3683. doi:10.1002/anie.201400131. ISSN 1521-3773. PMID 24615828. S2CID 13359849.
  47. ^ White, Mary Anne; Kahwaji, Samer; Freitas, Vera L. S.; Siewert, Riko; Weatherby, Joseph A.; Ribeiro da Silva, Maria D. M. C.; Verevkin, Sergey P.; Johnson, Erin R.; Zwanziger, Josef W. (2021). «The Relative Thermal Stability of Diamond and Graphite». Angewandte Chemie International Edition. 60 (3): 1546–1549. doi:10.1002/anie.202009897. ISSN 1433-7851. PMID 32970365. S2CID 221888151.
  48. ^ Schewe, Phil & Stein, Ben (March 26, 2004). «Carbon Nanofoam is the World’s First Pure Carbon Magnet». Physics News Update. 678 (1). Archived from the original on March 7, 2012.
  49. ^ Itzhaki, Lior; Altus, Eli; Basch, Harold; Hoz, Shmaryahu (2005). «Harder than Diamond: Determining the Cross-Sectional Area and Young’s Modulus of Molecular Rods». Angew. Chem. Int. Ed. 44 (45): 7432–5. doi:10.1002/anie.200502448. PMID 16240306.
  50. ^ «Researchers Find New Phase of Carbon, Make Diamond at Room Temperature». news.ncsu.edu. 2015-11-30. Archived from the original on 2016-04-06. Retrieved 2016-04-06.
  51. ^ a b c Hoover, Rachel (21 February 2014). «Need to Track Organic Nano-Particles Across the Universe? NASA’s Got an App for That». NASA. Archived from the original on 6 September 2015. Retrieved 2014-02-22.
  52. ^ Lauretta, D.S.; McSween, H.Y. (2006). Meteorites and the Early Solar System II. Space science series. University of Arizona Press. p. 199. ISBN 978-0-8165-2562-1. Archived from the original on 2017-11-22. Retrieved 2017-05-07.
  53. ^ Mark, Kathleen (1987). Meteorite Craters. University of Arizona Press. ISBN 978-0-8165-0902-7.
  54. ^ «Online Database Tracks Organic Nano-Particles Across the Universe». Sci Tech Daily. February 24, 2014. Archived from the original on March 18, 2015. Retrieved 2015-03-10.
  55. ^ William F McDonough The composition of the Earth Archived 2011-09-28 at the Wayback Machine in Majewski, Eugeniusz (2000). Earthquake Thermodynamics and Phase Transformation in the Earth’s Interior. ISBN 978-0126851854.
  56. ^ Yinon Bar-On; et al. (Jun 19, 2018). «The biomass distribution on Earth». PNAS. 115 (25): 6506–6511. Bibcode:2018PNAS..115.6506B. doi:10.1073/pnas.1711842115. PMC 6016768. PMID 29784790.
  57. ^ Fred Pearce (2014-02-15). «Fire in the hole: After fracking comes coal». New Scientist. 221 (2956): 36–41. Bibcode:2014NewSc.221…36P. doi:10.1016/S0262-4079(14)60331-6. Archived from the original on 2015-03-16.
  58. ^ «Wonderfuel: Welcome to the age of unconventional gas» Archived 2014-12-09 at the Wayback Machine by Helen Knight, New Scientist, 12 June 2010, pp. 44–7.
  59. ^ Ocean methane stocks ‘overstated’ Archived 2013-04-25 at the Wayback Machine, BBC, 17 Feb. 2004.
  60. ^ «Ice on fire: The next fossil fuel» Archived 2015-02-22 at the Wayback Machine by Fred Pearce, New Scientist, 27 June 2009, pp. 30–33.
  61. ^ Calculated from file global.1751_2008.csv in «Index of /ftp/ndp030/CSV-FILES». Archived from the original on 2011-10-22. Retrieved 2011-11-06. from the Carbon Dioxide Information Analysis Center.
  62. ^ Rachel Gross (Sep 21, 2013). «Deep, and dank mysterious». New Scientist: 40–43. Archived from the original on 2013-09-21.
  63. ^ Stefanenko, R. (1983). Coal Mining Technology: Theory and Practice. Society for Mining Metallurgy. ISBN 978-0-89520-404-2.
  64. ^ Kasting, James (1998). «The Carbon Cycle, Climate, and the Long-Term Effects of Fossil Fuel Burning». Consequences: The Nature and Implication of Environmental Change. 4 (1). Archived from the original on 2008-10-24.
  65. ^ «Carbon-14 formation». Archived from the original on 1 August 2015. Retrieved 13 October 2014.
  66. ^ Aitken, M.J. (1990). Science-based Dating in Archaeology. pp. 56–58. ISBN 978-0-582-49309-4.
  67. ^ Nichols, Charles R. «Voltatile Products from Carbonaceous Asteroids» (PDF). UAPress.Arizona.edu. Archived from the original (PDF) on 2 July 2016. Retrieved 12 November 2016.
  68. ^ Gannes, Leonard Z.; Del Rio, Carlos Martı́nez; Koch, Paul (1998). «Natural Abundance Variations in Stable Isotopes and their Potential Uses in Animal Physiological Ecology». Comparative Biochemistry and Physiology – Part A: Molecular & Integrative Physiology. 119 (3): 725–737. doi:10.1016/S1095-6433(98)01016-2. PMID 9683412.
  69. ^ «Official SI Unit definitions». Archived from the original on 2007-10-14. Retrieved 2007-12-21.
  70. ^ Bowman, S. (1990). Interpreting the past: Radiocarbon dating. British Museum Press. ISBN 978-0-7141-2047-8.
  71. ^ Brown, Tom (March 1, 2006). «Carbon Goes Full Circle in the Amazon». Lawrence Livermore National Laboratory. Archived from the original on September 22, 2008. Retrieved 2007-11-25.
  72. ^ Libby, W. F. (1952). Radiocarbon dating. Chicago University Press and references therein.
  73. ^ Westgren, A. (1960). «The Nobel Prize in Chemistry 1960». Nobel Foundation. Archived from the original on 2007-10-25. Retrieved 2007-11-25.
  74. ^ «Use query for carbon-8». barwinski.net. Archived from the original on 2005-02-07. Retrieved 2007-12-21.
  75. ^ Watson, A. (1999). «Beaming Into the Dark Corners of the Nuclear Kitchen». Science. 286 (5437): 28–31. doi:10.1126/science.286.5437.28. S2CID 117737493.
  76. ^ Audi, Georges; Bersillon, Olivier; Blachot, Jean; Wapstra, Aaldert Hendrik (1997). «The NUBASE evaluation of nuclear and decay properties» (PDF). Nuclear Physics A. 624 (1): 1–124. Bibcode:1997NuPhA.624….1A. doi:10.1016/S0375-9474(97)00482-X. Archived from the original (PDF) on 2008-09-23.
  77. ^ Ostlie, Dale A. & Carroll, Bradley W. (2007). An Introduction to Modern Stellar Astrophysics. San Francisco (CA): Addison Wesley. ISBN 978-0-8053-0348-3.
  78. ^ Whittet, Douglas C. B. (2003). Dust in the Galactic Environment. CRC Press. pp. 45–46. ISBN 978-0-7503-0624-9.
  79. ^ Pikelʹner, Solomon Borisovich (1977). Star Formation. Springer. p. 38. ISBN 978-90-277-0796-3. Archived from the original on 2012-11-23. Retrieved 2011-06-06.
  80. ^ Mannion, pp. 51–54.
  81. ^ Mannion, pp. 84–88.
  82. ^ Falkowski, P.; Scholes, R. J.; Boyle, E.; Canadell, J.; Canfield, D.; Elser, J.; Gruber, N.; Hibbard, K.; et al. (2000). «The Global Carbon Cycle: A Test of Our Knowledge of Earth as a System». Science. 290 (5490): 291–296. Bibcode:2000Sci…290..291F. doi:10.1126/science.290.5490.291. PMID 11030643. S2CID 1779934.
  83. ^ Smith, T. M.; Cramer, W. P.; Dixon, R. K.; Leemans, R.; Neilson, R. P.; Solomon, A. M. (1993). «The global terrestrial carbon cycle» (PDF). Water, Air, & Soil Pollution. 70 (1–4): 19–37. Bibcode:1993WASP…70…19S. doi:10.1007/BF01104986. S2CID 97265068. Archived (PDF) from the original on 2022-10-11.
  84. ^ Burrows, A.; Holman, J.; Parsons, A.; Pilling, G.; Price, G. (2017). Chemistry3: Introducing Inorganic, Organic and Physical Chemistry. Oxford University Press. p. 70. ISBN 978-0-19-873380-5. Archived from the original on 2017-11-22. Retrieved 2017-05-07.
  85. ^ Mannion pp. 27–51
  86. ^ Mannion pp. 84–91
  87. ^ Levine, Joel S.; Augustsson, Tommy R.; Natarajan, Murali (1982). «The prebiological paleoatmosphere: stability and composition». Origins of Life and Evolution of Biospheres. 12 (3): 245–259. Bibcode:1982OrLi…12..245L. doi:10.1007/BF00926894. PMID 7162799. S2CID 20097153.
  88. ^ Loerting, T.; et al. (2001). «On the Surprising Kinetic Stability of Carbonic Acid». Angew. Chem. Int. Ed. 39 (5): 891–895. doi:10.1002/(SICI)1521-3773(20000303)39:5<891::AID-ANIE891>3.0.CO;2-E. PMID 10760883.
  89. ^ Haldane J. (1895). «The action of carbonic oxide on man». Journal of Physiology. 18 (5–6): 430–462. doi:10.1113/jphysiol.1895.sp000578. PMC 1514663. PMID 16992272.
  90. ^ Gorman, D.; Drewry, A.; Huang, Y. L.; Sames, C. (2003). «The clinical toxicology of carbon monoxide». Toxicology. 187 (1): 25–38. doi:10.1016/S0300-483X(03)00005-2. PMID 12679050.
  91. ^ «Compounds of carbon: carbon suboxide». Archived from the original on 2007-12-07. Retrieved 2007-12-03.
  92. ^ Bayes, K. (1961). «Photolysis of Carbon Suboxide». Journal of the American Chemical Society. 83 (17): 3712–3713. doi:10.1021/ja01478a033.
  93. ^ Anderson D. J.; Rosenfeld, R. N. (1991). «Photodissociation of Carbon Suboxide». Journal of Chemical Physics. 94 (12): 7852–7867. Bibcode:1991JChPh..94.7857A. doi:10.1063/1.460121.
  94. ^ Sabin, J. R.; Kim, H. (1971). «A theoretical study of the structure and properties of carbon trioxide». Chemical Physics Letters. 11 (5): 593–597. Bibcode:1971CPL….11..593S. doi:10.1016/0009-2614(71)87010-0.
  95. ^ Moll N. G.; Clutter D. R.; Thompson W. E. (1966). «Carbon Trioxide: Its Production, Infrared Spectrum, and Structure Studied in a Matrix of Solid CO2«. Journal of Chemical Physics. 45 (12): 4469–4481. Bibcode:1966JChPh..45.4469M. doi:10.1063/1.1727526.
  96. ^ a b Fatiadi, Alexander J.; Isbell, Horace S.; Sager, William F. (1963). «Cyclic Polyhydroxy Ketones. I. Oxidation Products of Hexahydroxybenzene (Benzenehexol)» (PDF). Journal of Research of the National Bureau of Standards Section A. 67A (2): 153–162. doi:10.6028/jres.067A.015. PMC 6640573. PMID 31580622. Archived from the original (PDF) on 2009-03-25. Retrieved 2009-03-21.
  97. ^ Pauling, L. (1960). The Nature of the Chemical Bond (3rd ed.). Ithaca, NY: Cornell University Press. p. 93. ISBN 978-0-8014-0333-0.
  98. ^ Greenwood and Earnshaw, pp. 297–301
  99. ^ Scherbaum, Franz; et al. (1988). ««Aurophilicity» as a consequence of Relativistic Effects: The Hexakis(triphenylphosphaneaurio)methane Dication [(Ph3PAu)6C]2+«. Angew. Chem. Int. Ed. Engl. 27 (11): 1544–1546. doi:10.1002/anie.198815441.
  100. ^ Ritter, Stephen K. «Six bonds to carbon: Confirmed». Chemical & Engineering News. Archived from the original on 2017-01-09.
  101. ^ Yamashita, Makoto; Yamamoto, Yohsuke; Akiba, Kin-ya; Hashizume, Daisuke; Iwasaki, Fujiko; Takagi, Nozomi; Nagase, Shigeru (2005-03-01). «Syntheses and Structures of Hypervalent Pentacoordinate Carbon and Boron Compounds Bearing an Anthracene Skeleton − Elucidation of Hypervalent Interaction Based on X-ray Analysis and DFT Calculation». Journal of the American Chemical Society. 127 (12): 4354–4371. doi:10.1021/ja0438011. ISSN 0002-7863. PMID 15783218.
  102. ^ Shorter Oxford English Dictionary, Oxford University Press
  103. ^ «Chinese made first use of diamond». BBC News. 17 May 2005. Archived from the original on 20 March 2007. Retrieved 2007-03-21.
  104. ^ van der Krogt, Peter. «Carbonium/Carbon at Elementymology & Elements Multidict». Archived from the original on 2010-01-23. Retrieved 2010-01-06.
  105. ^ Ferchault de Réaumur, R.-A. (1722). L’art de convertir le fer forgé en acier, et l’art d’adoucir le fer fondu, ou de faire des ouvrages de fer fondu aussi finis que le fer forgé (English translation from 1956). Paris, Chicago.
  106. ^ «Carbon». Canada Connects. Archived from the original on 2010-10-27. Retrieved 2010-12-07.
  107. ^ a b Senese, Fred (2000-09-09). «Who discovered carbon?». Frostburg State University. Archived from the original on 2007-12-07. Retrieved 2007-11-24.
  108. ^ Giolitti, Federico (1914). The Cementation of Iron and Steel. McGraw-Hill Book Company, inc.
  109. ^ Kroto, H. W.; Heath, J. R.; O’Brien, S. C.; Curl, R. F.; Smalley, R. E. (1985). «C60: Buckminsterfullerene». Nature. 318 (6042): 162–163. Bibcode:1985Natur.318..162K. doi:10.1038/318162a0. S2CID 4314237.
  110. ^ «The Nobel Prize in Chemistry 1996 «for their discovery of fullerenes»«. Archived from the original on 2007-10-11. Retrieved 2007-12-21.
  111. ^ a b c USGS Minerals Yearbook: Graphite, 2009 Archived 2008-09-16 at the Wayback Machine and Graphite: Mineral Commodity Summaries 2011
  112. ^ Harlow, G. E. (1998). The nature of diamonds. Cambridge University Press. p. 223. ISBN 978-0-521-62935-5.
  113. ^ Catelle, W. R. (1911). The Diamond. John Lane Company. p. 159. discussion on alluvial diamonds in India and elsewhere as well as earliest finds
  114. ^ Ball, V. (1881). Diamonds, Gold and Coal of India. London, Truebner & Co. Ball was a Geologist in British service. Chapter I, Page 1
  115. ^ Hershey, J. W. (1940). The Book Of Diamonds: Their Curious Lore, Properties, Tests And Synthetic Manufacture. Kessinger Pub Co. p. 28. ISBN 978-1-4179-7715-4.
  116. ^ Janse, A. J. A. (2007). «Global Rough Diamond Production Since 1870». Gems and Gemology. XLIII (Summer 2007): 98–119. doi:10.5741/GEMS.43.2.98.
  117. ^ Marshall, Stephen; Shore, Josh (2004-10-22). «The Diamond Life». Guerrilla News Network. Archived from the original on 2008-06-09. Retrieved 2008-10-10.
  118. ^ Zimnisky, Paul (21 May 2018). «Global Diamond Supply Expected to Decrease 3.4% to 147M Carats in 2018». Kitco.com. Retrieved 9 November 2020.
  119. ^ Lorenz, V. (2007). «Argyle in Western Australia: The world’s richest diamantiferous pipe; its past and future». Gemmologie, Zeitschrift der Deutschen Gemmologischen Gesellschaft. 56 (1/2): 35–40.
  120. ^ Mannion pp. 25–26
  121. ^ «Microscopic diamond found in Montana». The Montana Standard. 2004-10-17. Archived from the original on 2005-01-21. Retrieved 2008-10-10.
  122. ^ Cooke, Sarah (2004-10-19). «Microscopic Diamond Found in Montana». Livescience.com. Archived from the original on 2008-07-05. Retrieved 2008-09-12.
  123. ^ «Delta News / Press Releases / Publications». Deltamine.com. Archived from the original on 2008-05-26. Retrieved 2008-09-12.
  124. ^ Cantwell, W. J.; Morton, J. (1991). «The impact resistance of composite materials – a review». Composites. 22 (5): 347–62. doi:10.1016/0010-4361(91)90549-V.
  125. ^ Holtzapffel, Ch. (1856). Turning And Mechanical Manipulation. Charles Holtzapffel.
    Internet Archive Archived 2016-03-26 at the Wayback Machine
  126. ^ «Industrial Diamonds Statistics and Information». United States Geological Survey. Archived from the original on 2009-05-06. Retrieved 2009-05-05.
  127. ^ Coelho, R. T.; Yamada, S.; Aspinwall, D. K.; Wise, M. L. H. (1995). «The application of polycrystalline diamond (PCD) tool materials when drilling and reaming aluminum-based alloys including MMC». International Journal of Machine Tools and Manufacture. 35 (5): 761–774. doi:10.1016/0890-6955(95)93044-7.
  128. ^ Harris, D. C. (1999). Materials for infrared windows and domes: properties and performance. SPIE Press. pp. 303–334. ISBN 978-0-8194-3482-1.
  129. ^ Nusinovich, G. S. (2004). Introduction to the physics of gyrotrons. JHU Press. p. 229. ISBN 978-0-8018-7921-0.
  130. ^ Sakamoto, M.; Endriz, J. G.; Scifres, D. R. (1992). «120 W CW output power from monolithic AlGaAs (800 nm) laser diode array mounted on diamond heatsink». Electronics Letters. 28 (2): 197–199. Bibcode:1992ElL….28..197S. doi:10.1049/el:19920123.
  131. ^ Dorfer, Leopold; Moser, M.; Spindler, K.; Bahr, F.; Egarter-Vigl, E.; Dohr, G. (1998). «5200-year old acupuncture in Central Europe?». Science. 282 (5387): 242–243. Bibcode:1998Sci…282..239D. doi:10.1126/science.282.5387.239f. PMID 9841386. S2CID 42284618.
  132. ^ Donaldson, K.; Stone, V.; Clouter, A.; Renwick, L.; MacNee, W. (2001). «Ultrafine particles». Occupational and Environmental Medicine. 58 (3): 211–216. doi:10.1136/oem.58.3.211. PMC 1740105. PMID 11171936.
  133. ^ Carbon Nanoparticles Toxic To Adult Fruit Flies But Benign To Young Archived 2011-11-02 at the Wayback Machine ScienceDaily (Aug. 17, 2009)
  134. ^ «Press Release – Titanic Disaster: New Theory Fingers Coal Fire». www.geosociety.org. Archived from the original on 2016-04-14. Retrieved 2016-04-06.
  135. ^ McSherry, Patrick. «Coal bunker Fire». www.spanamwar.com. Archived from the original on 2016-03-23. Retrieved 2016-04-06.

Bibliography

  • Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  • Mannion, A. M. (12 January 2006). Carbon and Its Domestication. Springer. pp. 1–319. ISBN 978-1-4020-3956-0.

External links

  • Carbon on In Our Time at the BBC
  • Carbon at The Periodic Table of Videos (University of Nottingham)
  • Carbon on Britannica
  • Extensive Carbon page at asu.edu (archived 18 June 2010)
  • Electrochemical uses of carbon (archived 9 November 2001)
  • Carbon—Super Stuff. Animation with sound and interactive 3D-models. (archived 9 November 2012)

«Element 6» redirects here. For the company, see Element Six.

Carbon, 6C

Graphite-and-diamond-with-scale.jpg

Graphite (left) and diamond (right), two allotropes of carbon

Carbon
Allotropes graphite, diamond and more (see Allotropes of carbon)
Appearance
  • graphite: black, metallic-looking
  • diamond: clear
Standard atomic weight Ar°(C)
  • [12.009612.0116]
  • 12.011±0.002 (abridged)[1]
Carbon in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson


C

Si
boron ← carbon → nitrogen
Atomic number (Z) 6
Group group 14 (carbon group)
Period period 2
Block   p-block
Electron configuration [He] 2s2 2p2
Electrons per shell 2, 4
Physical properties
Phase at STP solid
Sublimation point 3915 K ​(3642 °C, ​6588 °F)
Density (near r.t.) amorphous: 1.8–2.1 g/cm3[2]
graphite: 2.267 g/cm3
diamond: 3.515 g/cm3
Triple point 4600 K, ​10,800 kPa[3][4]
Heat of fusion graphite: 117 kJ/mol
Molar heat capacity graphite: 8.517 J/(mol·K)
diamond: 6.155 J/(mol·K)
Atomic properties
Oxidation states −4, −3, −2, −1, 0, +1,[5] +2, +3,[6] +4[7] (a mildly acidic oxide)
Electronegativity Pauling scale: 2.55
Ionization energies
  • 1st: 1086.5 kJ/mol
  • 2nd: 2352.6 kJ/mol
  • 3rd: 4620.5 kJ/mol
  • (more)
Covalent radius sp3: 77 pm
sp2: 73 pm
sp: 69 pm
Van der Waals radius 170 pm

Color lines in a spectral range

Spectral lines of carbon

Other properties
Natural occurrence primordial
Crystal structure graphite: ​simple hexagonal

Simple hexagonal crystal structure for graphite: carbon

(black)

Crystal structure diamond: ​face-centered diamond-cubic

Diamond cubic crystal structure for diamond: carbon

(clear)

Speed of sound thin rod diamond: 18,350 m/s (at 20 °C)
Thermal expansion diamond: 0.8 µm/(m⋅K) (at 25 °C)[8]
Thermal conductivity graphite: 119–165 W/(m⋅K)
diamond: 900–2300 W/(m⋅K)
Electrical resistivity graphite: 7.837 µΩ⋅m[9]
Magnetic ordering diamagnetic[10]
Molar magnetic susceptibility diamond: −5.9×10−6 cm3/mol[11]
Young’s modulus diamond: 1050 GPa[8]
Shear modulus diamond: 478 GPa[8]
Bulk modulus diamond: 442 GPa[8]
Poisson ratio diamond: 0.1[8]
Mohs hardness graphite: 1–2
diamond: 10
CAS Number
  • atomic carbon: 7440-44-0
  • graphite: 7782-42-5
  • diamond: 7782-40-3
History
Discovery Egyptians and Sumerians[12] (3750 BCE)
Recognized as an element by Antoine Lavoisier[13] (1789)
Main isotopes of carbon

  • v
  • e

Iso­tope Decay
abun­dance half-life (t1/2) mode pro­duct
11C syn 20.3402(53) min β+ 11B
12C [98.84%99.04%] stable
13C [0.96%1.16%] stable
14C 1 ppt 5.70(3)×103 y β 14N
 Category: Carbon

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| references

Carbon (from Latin carbo ‘coal’) is a chemical element with the symbol C and atomic number 6. It is nonmetallic and tetravalent—its atom making four electrons available to form covalent chemical bonds. It belongs to group 14 of the periodic table.[14] Carbon makes up about 0.025 percent of Earth’s crust.[15] Three isotopes occur naturally, 12C and 13C being stable, while 14C is a radionuclide, decaying with a half-life of about 5,730 years.[16] Carbon is one of the few elements known since antiquity.[17]

Carbon is the 15th most abundant element in the Earth’s crust, and the fourth most abundant element in the universe by mass after hydrogen, helium, and oxygen. Carbon’s abundance, its unique diversity of organic compounds, and its unusual ability to form polymers at the temperatures commonly encountered on Earth, enables this element to serve as a common element of all known life. It is the second most abundant element in the human body by mass (about 18.5%) after oxygen.[18]

The atoms of carbon can bond together in diverse ways, resulting in various allotropes of carbon. Well-known allotropes include graphite, diamond, amorphous carbon and fullerenes. The physical properties of carbon vary widely with the allotropic form. For example, graphite is opaque and black while diamond is highly transparent. Graphite is soft enough to form a streak on paper (hence its name, from the Greek verb «γράφειν» which means «to write»), while diamond is the hardest naturally occurring material known. Graphite is a good electrical conductor while diamond has a low electrical conductivity. Under normal conditions, diamond, carbon nanotubes, and graphene have the highest thermal conductivities of all known materials. All carbon allotropes are solids under normal conditions, with graphite being the most thermodynamically stable form at standard temperature and pressure. They are chemically resistant and require 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 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 a vast number of compounds, with almost ten million compounds described to date,[19] and yet that number is but a fraction of the number of theoretically possible compounds under standard conditions.

Characteristics

Theoretically predicted phase diagram of carbon, from 1989. Newer work indicates that the melting point of diamond (top-right curve) does not go above about 9000 K.[20]

The allotropes of carbon include graphite, one of the softest known substances, and diamond, the hardest naturally occurring substance. It bonds readily with other small atoms, including other carbon atoms, and is capable of forming multiple stable covalent bonds with suitable multivalent atoms. Carbon is known to form almost ten million compounds, a large majority of all chemical compounds.[19] Carbon also has the highest sublimation point of all elements. At atmospheric pressure it has no melting point, as its triple point is at 10.8 ± 0.2 megapascals (106.6 ± 2.0 atm; 1,566 ± 29 psi) and 4,600 ± 300 K (4,330 ± 300 °C; 7,820 ± 540 °F),[3][4] so it sublimes at about 3,900 K (3,630 °C; 6,560 °F).[21][22] Graphite is much more reactive than diamond at standard conditions, despite being more thermodynamically stable, as its delocalised pi system is much more vulnerable to attack. For example, graphite can be oxidised by hot concentrated nitric acid at standard conditions to mellitic acid, C6(CO2H)6, which preserves the hexagonal units of graphite while breaking up the larger structure.[23]

Carbon sublimes in a carbon arc, which has a temperature of about 5800 K (5,530 °C or 9,980 °F). 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, which are weaker reducing agents at room temperature.

Carbon is the sixth element, with a ground-state electron configuration of 1s22s22p2, of which the four outer electrons are valence electrons. Its first four ionisation energies, 1086.5, 2352.6, 4620.5 and 6222.7 kJ/mol, are much higher than those of the heavier group-14 elements. The electronegativity of carbon is 2.5, significantly higher than the heavier group-14 elements (1.8–1.9), but close to most of the nearby nonmetals, as well as some of the second- and third-row transition metals. Carbon’s covalent radii are normally taken as 77.2 pm (C−C), 66.7 pm (C=C) and 60.3 pm (C≡C), although these may vary depending on coordination number and what the carbon is bonded to. In general, covalent radius decreases with lower coordination number and higher bond order.[24]

Carbon-based compounds form the basis of all known 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 rob oxygen from metal oxides to leave the elemental metal. This exothermic reaction is used in the iron and steel industry to smelt iron and to control the carbon content of steel:

Fe
3
O
4
+ 4 C(s) + 2 O
2
→ 3 Fe(s) + 4 CO
2
(g).

Carbon reacts with sulfur to form carbon disulfide, and it reacts with steam in the coal-gas reaction used in coal gasification:

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:

Graphite is one of the softest materials known. Synthetic nanocrystalline diamond is the hardest material known.[25]
Graphite is a very good lubricant, displaying superlubricity.[26] Diamond is the ultimate abrasive.
Graphite is a conductor of electricity.[27] Diamond is an excellent electrical insulator,[28] and has the highest breakdown electric field of any known material.
Some forms of graphite are used for thermal insulation (i.e. firebreaks and heat shields), but some other forms are good thermal conductors. Diamond is the best known naturally occurring thermal conductor
Graphite is opaque. Diamond is highly transparent.
Graphite crystallizes in the hexagonal system.[29] Diamond crystallizes in the cubic system.
Amorphous carbon is completely isotropic. Carbon nanotubes are among the most anisotropic materials known.

Allotropes

Atomic carbon is a very short-lived species and, therefore, carbon is stabilized in various multi-atomic structures with diverse 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,[30][31] carbon nanotubes,[32] carbon nanobuds[33] and nanofibers.[34][35] Several other exotic allotropes have also been discovered, such as lonsdaleite,[36] glassy carbon,[37] carbon nanofoam[38] and linear acetylenic carbon (carbyne).[39]

Graphene is a two-dimensional sheet of carbon with the atoms arranged in a hexagonal lattice. As of 2009, graphene appears to be the strongest material ever tested.[40] The process of separating it from graphite will require some further technological development before it is economical for industrial processes.[41] If successful, graphene could be used in the construction of a space elevator. It could also be used to safely store hydrogen for use in a hydrogen based engine in cars.[42]

A large sample of glassy carbon

The amorphous form is an assortment of carbon atoms in a non-crystalline, irregular, glassy state, 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.[43] 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.

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, forming a 3-dimensional network of puckered six-membered rings of atoms. Diamond has the same cubic structure as silicon and germanium, and because of the strength of the carbon-carbon bonds, it is the hardest naturally occurring substance measured by resistance to scratching. Contrary to the popular belief that «diamonds are forever», they are thermodynamically unstable (ΔfG°(diamond, 298 K) = 2.9 kJ/mol[44]) under normal conditions (298 K, 105 Pa) and should theoretically transform into graphite.[45] But due to a high activation energy barrier, the transition into graphite is so slow at normal temperature that it is unnoticeable. However, at very high temperatures diamond will turn into graphite, and diamonds can burn up in a house fire. The bottom left corner of the phase diagram for carbon has not been scrutinized experimentally. Although a computational study employing density functional theory methods reached the conclusion that as T → 0 K and p → 0 Pa, diamond becomes more stable than graphite by approximately 1.1 kJ/mol,[46] more recent and definitive experimental and computational studies show that graphite is more stable than diamond for T < 400 K, without applied pressure, by 2.7 kJ/mol at T = 0 K and 3.2 kJ/mol at T = 298.15 K.[47] Under some conditions, carbon crystallizes as lonsdaleite, a hexagonal crystal lattice with all atoms covalently bonded and properties similar to those of diamond.[36]

Fullerenes are a synthetic crystalline formation with a graphite-like structure, but in place of flat hexagonal cells only, some of the cells of which fullerenes are formed may be pentagons, nonplanar hexagons, or even heptagons of carbon atoms. The sheets are thus warped into spheres, ellipses, or cylinders. The properties of fullerenes (split into buckyballs, buckytubes, and nanobuds) have not yet been fully analyzed and represent an intense area of research in nanomaterials. The names fullerene and buckyball are given after Richard Buckminster Fuller, popularizer of geodesic domes, 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 C60 buckminsterfullerene).[30] Carbon nanotubes (buckytubes) are structurally similar to buckyballs, except that each atom is bonded trigonally in a curved sheet that forms a hollow cylinder.[31][32] Nanobuds were first reported in 2007 and are hybrid buckytube/buckyball materials (buckyballs are covalently bonded to the outer wall of a nanotube) that combine the properties of both in a single structure.[33]

Of the other discovered allotropes, 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.[48] Similarly, glassy carbon contains a high proportion of closed porosity,[37] but contrary to normal graphite, the graphitic layers are not stacked like pages in a book, but have a more random arrangement. Linear acetylenic carbon[39] has the chemical structure[39] −(C:::C)n−. Carbon in this modification is linear with sp orbital hybridization, and is a polymer with alternating single and triple bonds. This carbyne is of considerable interest to nanotechnology as its Young’s modulus is 40 times that of the hardest known material – diamond.[49]

In 2015, a team at the North Carolina State University announced the development of another allotrope they have dubbed Q-carbon, created by a high energy low duration laser pulse on amorphous carbon dust. Q-carbon is reported to exhibit ferromagnetism, fluorescence, and a hardness superior to diamonds.[50]

In the vapor phase, some of the carbon is in the form of dicarbon (C
2
). When excited, this gas glows green.

Occurrence

Graphite ore, shown with a penny for scale

Carbon is the fourth most abundant chemical element in the observable universe by mass after hydrogen, helium, and oxygen. Carbon is abundant in the Sun, stars, comets, and in the atmospheres of most planets.[51] Some meteorites contain microscopic diamonds that were formed when the Solar System was still a protoplanetary disk.[52] Microscopic diamonds may also be formed by the intense pressure and high temperature at the sites of meteorite impacts.[53]

In 2014 NASA announced a greatly upgraded database for tracking polycyclic aromatic hydrocarbons (PAHs) in the universe. More than 20% of the carbon in the universe may be associated with PAHs, complex compounds of carbon and hydrogen without oxygen.[54] These compounds figure in the PAH world hypothesis where they are hypothesized to have a role in abiogenesis and formation of life. PAHs seem to have been formed «a couple of billion years» after the Big Bang, are widespread throughout the universe, and are associated with new stars and exoplanets.[51]

It has been estimated that the solid earth as a whole contains 730 ppm of carbon, with 2000 ppm in the core and 120 ppm in the combined mantle and crust.[55] Since the mass of the earth is 5.972×1024 kg, this would imply 4360 million gigatonnes of carbon. This is much more than the amount of carbon in the oceans or atmosphere (below).

In combination with oxygen in carbon dioxide, carbon is found in the Earth’s atmosphere (approximately 900 gigatonnes of carbon — each ppm corresponds to 2.13 Gt) and dissolved in all water bodies (approximately 36,000 gigatonnes of carbon). Carbon in the biosphere has been estimated at 550 gigatonnes but with a large uncertainty, due mostly to a huge uncertainty in the amount of terrestrial deep subsurface bacteria.[56] Hydrocarbons (such as coal, petroleum, and natural gas) contain carbon as well. Coal «reserves» (not «resources») amount to around 900 gigatonnes with perhaps 18,000 Gt of resources.[57] Oil reserves are around 150 gigatonnes. Proven sources of natural gas are about 175×1012 cubic metres (containing about 105 gigatonnes of carbon), but studies estimate another 900×1012 cubic metres of «unconventional» deposits such as shale gas, representing about 540 gigatonnes of carbon.[58]

Carbon is also found in methane hydrates in polar regions and under the seas. Various estimates put this carbon between 500, 2500 Gt,[59] or 3,000 Gt.[60]

In the past, quantities of hydrocarbons were greater. According to one source, in the period from 1751 to 2008 about 347 gigatonnes of carbon were released as carbon dioxide to the atmosphere from burning of fossil fuels.[61] Another source puts the amount added to the atmosphere for the period since 1750 at 879 Gt, and the total going to the atmosphere, sea, and land (such as peat bogs) at almost 2,000 Gt.[62]

Carbon is a constituent (about 12% by mass) of the very large masses of carbonate rock (limestone, dolomite, marble and so on). Coal is very rich in carbon (anthracite contains 92–98%)[63] and is the largest commercial source of mineral carbon, accounting for 4,000 gigatonnes or 80% of fossil fuel.[64]

As for individual carbon allotropes, graphite is found in large quantities in the United States (mostly in New York and Texas), 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. Diamond deposits have also been found 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. Diamonds are found naturally, but about 30% of all industrial diamonds used in the U.S. are now manufactured.

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.[65] Thermal neutrons are produced that collide with the nuclei of nitrogen-14, forming carbon-14 and a proton. As such, 1.5%×10−10 of atmospheric carbon dioxide contains carbon-14.[66]

Carbon-rich asteroids are relatively preponderant in the outer parts of the asteroid belt in the Solar System. These asteroids have not yet been directly sampled by scientists. The asteroids can be used in hypothetical space-based carbon mining, which may be possible in the future, but is currently technologically impossible.[67]

Isotopes

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.[16] The isotope carbon-12 (12C) forms 98.93% of the carbon on Earth, while carbon-13 (13C) forms the remaining 1.07%.[16] The concentration of 12C is further increased in biological materials because biochemical reactions discriminate against 13C.[68] In 1961, the International Union of Pure and Applied Chemistry (IUPAC) adopted the isotope carbon-12 as the basis for atomic weights.[69] Identification of carbon in nuclear magnetic resonance (NMR) experiments is done with the isotope 13C.

Carbon-14 (14C) is a naturally occurring radioisotope, created in the upper atmosphere (lower stratosphere and upper troposphere) by interaction of nitrogen with cosmic rays.[70] It is found in trace amounts on Earth of 1 part per trillion (0.0000000001%) or more, mostly confined to the atmosphere and superficial deposits, particularly of peat and other organic materials.[71] 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. The amount 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.[72][73]

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 1.98739 × 10−21 s.[74] The exotic 19C exhibits a nuclear halo, which means its radius is appreciably larger than would be expected if the nucleus were a sphere of constant density.[75]

Formation in stars

Formation of the carbon atomic nucleus occurs within a giant or supergiant star through the triple-alpha process. This requires a nearly simultaneous collision of three alpha particles (helium nuclei), as the products of further nuclear fusion reactions of helium with hydrogen or another helium nucleus produce lithium-5 and beryllium-8 respectively, both of which are highly unstable and decay almost instantly back into smaller nuclei.[76] The triple-alpha process happens in conditions of temperatures over 100 megakelvins 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.

According to current physical cosmology theory, carbon is formed in the interiors of stars on the horizontal branch.[77] When massive stars die as supernova, the carbon is scattered into space as dust. This dust becomes component material for the formation of the next-generation star systems with accreted planets.[51][78] The Solar System is one such star system with an abundance of carbon, enabling the existence of life as we know it.

The CNO cycle is an additional hydrogen fusion mechanism that powers stars, wherein carbon operates as a catalyst.

Rotational transitions of various isotopic forms of carbon monoxide (for example, 12CO, 13CO, and 18CO) are detectable in the submillimeter wavelength range, and are used in the study of newly forming stars in molecular clouds.[79]

Carbon cycle

Diagram of the carbon cycle. The black numbers indicate how much carbon is stored in various reservoirs, in billions tonnes («GtC» stands for gigatonnes 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 from somewhere and dispose of it somewhere else. The paths of carbon in the environment form the carbon cycle.[80] For example, photosynthetic plants draw carbon dioxide from the atmosphere (or seawater) and build it into biomass, as in the Calvin cycle, a process of carbon fixation.[81] Some of this biomass is eaten by animals, while 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; if bacteria do not consume it, dead plant or animal matter may become petroleum or coal, which releases carbon when burned.[82][83]

Compounds

Organic compounds

Structural formula of methane, the simplest possible organic compound.

Correlation between the carbon cycle and formation of organic compounds. In plants, carbon dioxide formed by carbon fixation can join with water in photosynthesis (green) to form organic compounds, which can be used and further converted by both plants and animals.

Carbon can form very long chains of interconnecting carbon–carbon bonds, a property that is called catenation. Carbon-carbon bonds are strong and stable. Through catenation, carbon forms a countless number of compounds. A tally of unique compounds shows that more contain carbon than do not.[84] A similar claim can be made for hydrogen because most organic compounds contain hydrogen chemically bonded to carbon or another common element like oxygen or nitrogen.

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. A hydrocarbon backbone can be substituted by other atoms, known as heteroatoms. Common heteroatoms that appear in organic compounds include oxygen, nitrogen, sulfur, phosphorus, and the nonradioactive halogens, as well as the metals lithium and magnesium. Organic compounds containing bonds to metal are known as organometallic compounds (see below). Certain groupings of atoms, often including heteroatoms, recur in large numbers of organic compounds. These collections, known as functional groups, confer common reactivity patterns and allow for the systematic study and categorization of organic compounds. Chain length, shape and functional groups all affect the properties of organic molecules.[85]

In most stable compounds of carbon (and nearly all stable organic compounds), carbon obeys the octet rule and is tetravalent, meaning that a carbon atom forms a total of four covalent bonds (which may include double and triple bonds). Exceptions include a small number of stabilized carbocations (three bonds, positive charge), radicals (three bonds, neutral), carbanions (three bonds, negative charge) and carbenes (two bonds, neutral), although these species are much more likely to be encountered as unstable, reactive intermediates.

Carbon occurs in all known organic life and is the basis of organic chemistry. When united with hydrogen, it forms various hydrocarbons that are important to industry as refrigerants, lubricants, solvents, as chemical feedstock for the manufacture of plastics and petrochemicals, and as fossil fuels.

When combined with oxygen and hydrogen, carbon can form many groups of important biological compounds including sugars, 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 rubber products. With the addition of phosphorus to these other elements, it forms DNA and RNA, the chemical-code carriers of life, and adenosine triphosphate (ATP), the most important energy-transfer molecule in all living cells.[86]

Inorganic compounds

Commonly carbon-containing compounds which are associated with minerals or which do not contain bonds to the other carbon atoms, halogens, or hydrogen, are treated separately from classical organic compounds; the definition is not rigid, and the classification of some compounds can vary from author to author (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.[87] Dissolved in water, it forms carbonic acid (H
2
CO
3
), but as most compounds with multiple single-bonded oxygens on a single carbon it is unstable.[88] Through this intermediate, though, resonance-stabilized carbonate ions are produced. Some important minerals are carbonates, notably calcite. Carbon disulfide (CS
2
) is similar.[23] Nevertheless, due to its physical properties and its association with organic synthesis, carbon disulfide is sometimes classified as an organic solvent.

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.[89][90] 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. Likewise, the heavier analog of cyanide, cyaphide (CP), is also considered inorganic, though most simple derivatives are highly unstable. Other uncommon oxides are carbon suboxide (C
3
O
2
),[91] the unstable dicarbon monoxide (C2O),[92][93] carbon trioxide (CO3),[94][95] cyclopentanepentone (C5O5),[96] cyclohexanehexone (C6O6),[96] and mellitic anhydride (C12O9). However, mellitic anhydride is the triple acyl anhydride of mellitic acid; moreover, it contains a benzene ring. Thus, many chemists consider it to be organic.

With reactive metals, such as tungsten, carbon forms either carbides (C4−) or acetylides (C2−
2
) 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,[97] carbon prefers to form covalent bonds. A few carbides are covalent lattices, like carborundum (SiC), which resembles diamond. Nevertheless, even the most polar and salt-like of carbides are not completely ionic compounds.[98]

Organometallic compounds

Organometallic compounds by definition contain at least one carbon-metal covalent bond. A wide range of such compounds exist; major classes include simple alkyl-metal compounds (for example, tetraethyllead), η2-alkene compounds (for example, Zeise’s salt), and η3-allyl compounds (for example, allylpalladium chloride dimer); metallocenes containing cyclopentadienyl ligands (for example, ferrocene); and transition metal carbene complexes. Many metal carbonyls and metal cyanides exist (for example, tetracarbonylnickel and potassium ferricyanide); some workers consider metal carbonyl and cyanide complexes without other carbon ligands to be purely inorganic, and not organometallic. However, most organometallic chemists consider metal complexes with any carbon ligand, even ‘inorganic carbon’ (e.g., carbonyls, cyanides, and certain types of carbides and acetylides) to be organometallic in nature. Metal complexes containing organic ligands without a carbon-metal covalent bond (e.g., metal carboxylates) are termed metalorganic compounds.

While carbon is understood to strongly prefer formation of four covalent bonds, other exotic bonding schemes are also known. Carboranes are highly stable dodecahedral derivatives of the [B12H12]2- unit, with one BH replaced with a CH+. Thus, the carbon is bonded to five boron atoms and one hydrogen atom. The cation [(Ph3PAu)6C]2+ contains an octahedral carbon bound to six phosphine-gold fragments. This phenomenon has been attributed to the aurophilicity of the gold ligands, which provide additional stabilization of an otherwise labile species.[99] In nature, the iron-molybdenum cofactor (FeMoco) responsible for microbial nitrogen fixation likewise has an octahedral carbon center (formally a carbide, C(-IV)) bonded to six iron atoms. In 2016, it was confirmed that, in line with earlier theoretical predictions, the hexamethylbenzene dication contains a carbon atom with six bonds. More specifically, the dication could be described structurally by the formulation [MeC(η5-C5Me5)]2+, making it an «organic metallocene» in which a MeC3+ fragment is bonded to a η5-C5Me5 fragment through all five of the carbons of the ring.[100]

This anthracene derivative contains a carbon atom with 5 formal electron pairs around it.

It is important to note that in the cases above, each of the bonds to carbon contain less than two formal electron pairs. Thus, the formal electron count of these species does not exceed an octet. This makes them hypercoordinate but not hypervalent. Even in cases of alleged 10-C-5 species (that is, a carbon with five ligands and a formal electron count of ten), as reported by Akiba and co-workers,[101] electronic structure calculations conclude that the electron population around carbon is still less than eight, as is true for other compounds featuring four-electron three-center bonding.

History and etymology

The English name carbon comes from the Latin carbo for coal and charcoal,[102] whence also comes the French charbon, meaning charcoal. In German, Dutch and Danish, the names for carbon are Kohlenstoff, koolstof and kulstof respectively, all literally meaning coal-substance.

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 form 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.[103][104]

In 1722, René Antoine Ferchault de Réaumur demonstrated that iron was transformed into steel through the absorption of some substance, now known to be carbon.[105] In 1772, Antoine Lavoisier showed that diamonds are a form of carbon; when he burned samples of charcoal and diamond and found that neither produced any water and that both released the same amount of carbon dioxide per gram.
In 1779,[106] Carl Wilhelm Scheele showed that graphite, which had been thought of as a form of lead, was instead identical with charcoal but with a small admixture of iron, and that it gave «aerial acid» (his name for carbon dioxide) when oxidized with nitric acid.[107] In 1786, the French scientists Claude Louis Berthollet, Gaspard Monge and C. A. Vandermonde confirmed that graphite was mostly carbon by oxidizing it in oxygen in much the same way Lavoisier had done with diamond.[108] Some iron again was left, which the French scientists thought was necessary to the graphite structure. In their publication they proposed the name carbone (Latin carbonum) for the element in graphite which was given off as a gas upon burning graphite. Antoine Lavoisier then listed carbon as an element in his 1789 textbook.[107]

A new allotrope of carbon, fullerene, that was discovered in 1985[109] includes nanostructured forms such as buckyballs and nanotubes.[30] Their discoverers – Robert Curl, Harold Kroto and Richard Smalley – received the Nobel Prize in Chemistry in 1996.[110] The resulting renewed interest in new forms led to the discovery of further exotic allotropes, including glassy carbon, and the realization that «amorphous carbon» is not strictly amorphous.[37]

Production

Graphite

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.[citation needed] 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 19th century, pencils were made 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.[111]

There are three types of natural graphite—amorphous, flake or crystalline flake, and vein or lump. Amorphous graphite is the lowest quality and most abundant. Contrary to science, in industry «amorphous» refers to very small crystal size rather than complete lack of crystal structure. Amorphous is used for lower value graphite products and is the lowest priced graphite. Large amorphous graphite deposits are found in China, Europe, Mexico and the United States. Flake graphite is less common and of higher quality than amorphous; it occurs as separate plates that crystallized in metamorphic rock. Flake graphite can be four times the price of amorphous. Good quality flakes can be processed into expandable graphite for many uses, such as flame retardants. The foremost deposits are found in Austria, Brazil, Canada, China, Germany and Madagascar. Vein or lump graphite is the rarest, most valuable, and highest quality type of natural graphite. It occurs in veins along intrusive contacts in solid lumps, and it is only commercially mined in Sri Lanka.[111]

According to the USGS, world production of natural graphite was 1.1 million tonnes in 2010, to which China contributed 800,000 t, India 130,000 t, Brazil 76,000 t, North Korea 30,000 t and Canada 25,000 t. No natural graphite was reported mined in the United States, but 118,000 t of synthetic graphite with an estimated value of $998 million was produced in 2009.[111]

Diamond

The diamond supply chain is controlled by a limited number of powerful businesses, and is also highly concentrated in a small number of locations around the world (see figure).

Only a very small fraction of the diamond ore consists of actual diamonds. The ore is crushed, during which care has to be taken in order to prevent larger diamonds from being destroyed in this process and subsequently the particles are sorted by density. Today, diamonds are located in the diamond-rich density fraction with the help of X-ray fluorescence, after which the final sorting steps are done by hand. Before the use of X-rays became commonplace, the separation was done with grease belts; diamonds have a stronger tendency to stick to grease than the other minerals in the ore.[112]

Historically diamonds were known to be found only in alluvial deposits in southern India.[113] India led the world in diamond production from the time of their discovery in approximately the 9th century BC[114] to the mid-18th century AD, but the commercial potential of these sources had been exhausted by the late 18th century and at that time India was eclipsed by Brazil where the first non-Indian diamonds were found in 1725.[115]

Diamond production of primary deposits (kimberlites and lamproites) only started in the 1870s after the discovery of the diamond fields in South Africa. Production has increased over time and an accumulated total of over 4.5 billion carats have been mined since that date.[116] Most commercially viable diamond deposits were in Russia, Botswana, Australia and the Democratic Republic of Congo.[117] By 2005, Russia produced almost one-fifth of the global diamond output (mostly in Yakutia territory; for example, Mir pipe and Udachnaya pipe) but the Argyle mine in Australia became the single largest source, producing 14 million carats in 2018.[118][119] New finds, the Canadian mines at Diavik and Ekati, are expected to become even more valuable owing to their production of gem quality stones.[120]

In the United States, diamonds have been found in Arkansas, Colorado and Montana.[121] In 2004, a startling discovery of a microscopic diamond in the United States[122] led to the January 2008 bulk-sampling of kimberlite pipes in a remote part of Montana.[123]

Applications

Pencil leads for mechanical pencils are made of graphite (often mixed with a clay or synthetic binder).

A cloth of woven carbon fibres

The C60 fullerene in crystalline form

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 distilled in refineries by the petrochemical industry to produce gasoline, kerosene, and other products. Cellulose is a natural, carbon-containing polymer produced by plants in the form of wood, cotton, linen, and hemp. Cellulose is used primarily for maintaining structure in plants. 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 molding 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, barbecue grilling, iron smelting, and in many other applications. Wood, coal and oil are used as fuel for production of energy and 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 fiber, made by pyrolysis of synthetic polyester fibers 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 fibers made from PAN have structure resembling narrow filaments of graphite, but thermal processing may re-order the structure into a continuous rolled sheet. The result is fibers with higher specific tensile strength than steel.[124]

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 (smelting). 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.

Diamonds

The diamond industry falls into two categories: one dealing with gem-grade diamonds and the other, with industrial-grade diamonds. While a large trade in both types of diamonds exists, the two markets function dramatically differently.

Unlike precious metals such as gold or platinum, gem diamonds do not trade as a commodity: there is a substantial mark-up in the sale of diamonds, and there is not a very active market for resale of diamonds.

Industrial diamonds are valued mostly for their hardness and heat conductivity, with the gemological qualities of clarity and color being mostly irrelevant. About 80% of mined diamonds (equal to about 100 million carats or 20 tonnes annually) are unsuitable for use as gemstones are relegated for industrial use (known as bort).[125] synthetic diamonds, invented in the 1950s, found almost immediate industrial applications; 3 billion carats (600 tonnes) of synthetic diamond is produced annually.[126]

The dominant industrial use of diamond is in cutting, drilling, grinding, and polishing. Most of these applications do not require large diamonds; in fact, most diamonds of gem-quality except for their small size can be used industrially. Diamonds are embedded in drill tips or saw blades, or ground into a powder for use in grinding and polishing applications.[127] Specialized applications include use in laboratories as containment for high pressure experiments (see diamond anvil cell), high-performance bearings, and limited use in specialized windows.[128][129] With the continuing advances in the production of synthetic diamonds, new applications are becoming feasible. Garnering much excitement is the possible use of diamond as a semiconductor suitable for microchips, and because of its exceptional heat conductance property, as a heat sink in electronics.[130]

Precautions

Pure carbon has extremely low toxicity to humans and can be handled 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. Consequently, once it enters into the body’s 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.[131] 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. Diamond dust used as an abrasive can be harmful if ingested or inhaled. Microparticles of carbon are produced in diesel engine exhaust fumes, and may accumulate in the lungs.[132] In these examples, the harm may result from contaminants (e.g., organic chemicals, heavy metals) rather than from the carbon itself.

Carbon generally has low toxicity to life on Earth; but carbon nanoparticles are deadly to Drosophila.[133]

Carbon may burn vigorously and brightly in the presence of air at high temperatures. Large accumulations of coal, which have remained inert for hundreds of millions of years in the absence of oxygen, may spontaneously combust when exposed to air in coal mine waste tips, ship cargo holds and coal bunkers,[134][135] and storage dumps.

In nuclear applications where graphite is used as a neutron moderator, accumulation of Wigner energy followed by a sudden, spontaneous release may occur. Annealing to at least 250 °C can release the energy safely, although in the Windscale fire the procedure went wrong, causing other reactor materials to combust.

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.

See also

  • Carbon chauvinism
  • Carbon detonation
  • Carbon footprint
  • Carbon star
  • Carbon planet
  • Gas carbon
  • Low-carbon economy
  • Timeline of carbon nanotubes

References

  1. ^ «Standard Atomic Weights: Carbon». CIAAW. 2009.
  2. ^ Lide, D. R., ed. (2005). CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton (FL): CRC Press. ISBN 0-8493-0486-5.
  3. ^ a b Haaland, D (1976). «Graphite-liquid-vapor triple point pressure and the density of liquid carbon». Carbon. 14 (6): 357–361. doi:10.1016/0008-6223(76)90010-5.
  4. ^ a b Savvatimskiy, A (2005). «Measurements of the melting point of graphite and the properties of liquid carbon (a review for 1963–2003)». Carbon. 43 (6): 1115–1142. doi:10.1016/j.carbon.2004.12.027.
  5. ^ «Fourier Transform Spectroscopy of the Electronic Transition of the Jet-Cooled CCI Free Radical» (PDF). Retrieved 2007-12-06.
  6. ^ «Fourier Transform Spectroscopy of the System of CP» (PDF). Retrieved 2007-12-06.
  7. ^ «Carbon: Binary compounds». Retrieved 2007-12-06.
  8. ^ a b c d e Properties of diamond, Ioffe Institute Database
  9. ^ «Material Properties- Misc Materials». www.nde-ed.org. Retrieved 12 November 2016.
  10. ^ Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics 81st edition, CRC press.
  11. ^ Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN 978-0-8493-0464-4.
  12. ^ «History of Carbon and Carbon Materials — Center for Applied Energy Research — University of Kentucky». Caer.uky.edu. Retrieved 2008-09-12.
  13. ^ Senese, Fred (2000-09-09). «Who discovered carbon?». Frostburg State University. Retrieved 2007-11-24.
  14. ^ «carbon | Facts, Uses, & Properties». Encyclopedia Britannica. Archived from the original on 2017-10-24.
  15. ^ «carbon». Britannica encyclopedia.
  16. ^ a b c «Carbon – Naturally occurring isotopes». WebElements Periodic Table. Archived from the original on 2008-09-08. Retrieved 2008-10-09.
  17. ^ «History of Carbon». Archived from the original on 2012-11-01. Retrieved 2013-01-10.
  18. ^ Reece, Jane B. (31 October 2013). Campbell Biology (10 ed.). Pearson. ISBN 9780321775658.
  19. ^ a b Chemistry Operations (December 15, 2003). «Carbon». Los Alamos National Laboratory. Archived from the original on 2008-09-13. Retrieved 2008-10-09.
  20. ^ J.H. Eggert; et al. (Nov 8, 2009). «Melting temperature of diamond at ultrahigh pressure». Nature Physics. 6: 40–43. doi:10.1038/nphys1438.
  21. ^ Greenville Whittaker, A. (1978). «The controversial carbon solid−liquid−vapour triple point». Nature. 276 (5689): 695–696. Bibcode:1978Natur.276..695W. doi:10.1038/276695a0. S2CID 4362313.
  22. ^ Zazula, J. M. (1997). «On Graphite Transformations at High Temperature and Pressure Induced by Absorption of the LHC Beam» (PDF). CERN. Archived (PDF) from the original on 2009-03-25. Retrieved 2009-06-06.
  23. ^ a b Greenwood and Earnshaw, pp. 289–292.
  24. ^ Greenwood and Earnshaw, pp. 276–8.
  25. ^ Irifune, Tetsuo; Kurio, Ayako; Sakamoto, Shizue; Inoue, Toru; Sumiya, Hitoshi (2003). «Materials: Ultrahard polycrystalline diamond from graphite». Nature. 421 (6923): 599–600. Bibcode:2003Natur.421..599I. doi:10.1038/421599b. PMID 12571587. S2CID 52856300.
  26. ^ Dienwiebel, Martin; Verhoeven, Gertjan; Pradeep, Namboodiri; Frenken, Joost; Heimberg, Jennifer; Zandbergen, Henny (2004). «Superlubricity of Graphite» (PDF). Physical Review Letters. 92 (12): 126101. Bibcode:2004PhRvL..92l6101D. doi:10.1103/PhysRevLett.92.126101. PMID 15089689. S2CID 26811802. Archived (PDF) from the original on 2011-09-17.
  27. ^ Deprez, N.; McLachan, D. S. (1988). «The analysis of the electrical conductivity of graphite conductivity of graphite powders during compaction». Journal of Physics D: Applied Physics. 21 (1): 101–107. Bibcode:1988JPhD…21..101D. doi:10.1088/0022-3727/21/1/015. S2CID 250886376.
  28. ^ Collins, A. T. (1993). «The Optical and Electronic Properties of Semiconducting Diamond». Philosophical Transactions of the Royal Society A. 342 (1664): 233–244. Bibcode:1993RSPTA.342..233C. doi:10.1098/rsta.1993.0017. S2CID 202574625.
  29. ^ Delhaes, P. (2001). Graphite and Precursors. CRC Press. ISBN 978-90-5699-228-6.
  30. ^ a b c Unwin, Peter. «Fullerenes(An Overview)». Archived from the original on 2007-12-01. Retrieved 2007-12-08.
  31. ^ a b Ebbesen, T. W., ed. (1997). Carbon nanotubes—preparation and properties. Boca Raton, Florida: CRC Press. ISBN 978-0-8493-9602-1.
  32. ^ a b Dresselhaus, M. S.; Dresselhaus, G.; Avouris, Ph., eds. (2001). Carbon nanotubes: synthesis, structures, properties and applications. Topics in Applied Physics. Vol. 80. Berlin. ISBN 978-3-540-41086-7.
  33. ^ a b Nasibulin, Albert G.; Pikhitsa, P. V.; Jiang, H.; Brown, D. P.; Krasheninnikov, A. V.; Anisimov, A. S.; Queipo, P.; Moisala, A.; et al. (2007). «A novel hybrid carbon material». Nature Nanotechnology. 2 (3): 156–161. Bibcode:2007NatNa…2..156N. doi:10.1038/nnano.2007.37. PMID 18654245. S2CID 6447122.
  34. ^ Nasibulin, A.; Anisimov, Anton S.; Pikhitsa, Peter V.; Jiang, Hua; Brown, David P.; Choi, Mansoo; Kauppinen, Esko I. (2007). «Investigations of NanoBud formation». Chemical Physics Letters. 446 (1): 109–114. Bibcode:2007CPL…446..109N. doi:10.1016/j.cplett.2007.08.050.
  35. ^ Vieira, R; Ledoux, Marc-Jacques; Pham-Huu, Cuong (2004). «Synthesis and characterisation of carbon nanofibers with macroscopic shaping formed by catalytic decomposition of C2H6/H2 over nickel catalyst». Applied Catalysis A: General. 274 (1–2): 1–8. doi:10.1016/j.apcata.2004.04.008.
  36. ^ a b Frondel, Clifford; Marvin, Ursula B. (1967). «Lonsdaleite, a new hexagonal polymorph of diamond». Nature. 214 (5088): 587–589. Bibcode:1967Natur.214..587F. doi:10.1038/214587a0. S2CID 4184812.
  37. ^ a b c Harris, PJF (2004). «Fullerene-related structure of commercial glassy carbons» (PDF). Philosophical Magazine. 84 (29): 3159–3167. Bibcode:2004PMag…84.3159H. CiteSeerX 10.1.1.359.5715. doi:10.1080/14786430410001720363. S2CID 220342075. Archived from the original (PDF) on 2012-03-19. Retrieved 2011-07-06.
  38. ^ Rode, A. V.; Hyde, S. T.; Gamaly, E. G.; Elliman, R. G.; McKenzie, D. R.; Bulcock, S. (1999). «Structural analysis of a carbon foam formed by high pulse-rate laser ablation». Applied Physics A: Materials Science & Processing. 69 (7): S755–S758. Bibcode:1999ApPhA..69S.755R. doi:10.1007/s003390051522. S2CID 96050247.
  39. ^ a b c Heimann, Robert Bertram; Evsyukov, Sergey E. & Kavan, Ladislav (28 February 1999). Carbyne and carbynoid structures. Springer. pp. 1–. ISBN 978-0-7923-5323-2. Archived from the original on 23 November 2012. Retrieved 2011-06-06.
  40. ^ Lee, C.; Wei, X.; Kysar, J. W.; Hone, J. (2008). «Measurement of the Elastic Properties and Intrinsic Strength of Monolayer Graphene». Science. 321 (5887): 385–8. Bibcode:2008Sci…321..385L. doi:10.1126/science.1157996. PMID 18635798. S2CID 206512830.
    • Phil Schewe (July 28, 2008). «World’s Strongest Material». Inside Science News Service (Press release). Archived from the original on 2009-05-31.

  41. ^ Sanderson, Bill (2008-08-25). «Toughest Stuff Known to Man : Discovery Opens Door to Space Elevator». nypost.com. Archived from the original on 2008-09-06. Retrieved 2008-10-09.
  42. ^ Jin, Zhong; Lu, Wei; O’Neill, Kevin J.; Parilla, Philip A.; Simpson, Lin J.; Kittrell, Carter; Tour, James M. (2011-02-22). «Nano-Engineered Spacing in Graphene Sheets for Hydrogen Storage». Chemistry of Materials. 23 (4): 923–925. doi:10.1021/cm1025188. ISSN 0897-4756.
  43. ^ Jenkins, Edgar (1973). The polymorphism of elements and compounds. Taylor & Francis. p. 30. ISBN 978-0-423-87500-3. Archived from the original on 2012-11-23. Retrieved 2011-05-01.
  44. ^ Rossini, F. D.; Jessup, R. S. (1938). «Heat and Free Energy of Formation of Carbon Dioxide and of the Transition Between Graphite and Diamond». Journal of Research of the National Bureau of Standards. 21 (4): 491. doi:10.6028/jres.021.028.
  45. ^ «World of Carbon – Interactive Nano-visulisation in Science & Engineering Education (IN-VSEE)». Archived from the original on 2001-05-31. Retrieved 2008-10-09.
  46. ^ Grochala, Wojciech (2014-04-01). «Diamond: Electronic Ground State of Carbon at Temperatures Approaching 0 K». Angewandte Chemie International Edition. 53 (14): 3680–3683. doi:10.1002/anie.201400131. ISSN 1521-3773. PMID 24615828. S2CID 13359849.
  47. ^ White, Mary Anne; Kahwaji, Samer; Freitas, Vera L. S.; Siewert, Riko; Weatherby, Joseph A.; Ribeiro da Silva, Maria D. M. C.; Verevkin, Sergey P.; Johnson, Erin R.; Zwanziger, Josef W. (2021). «The Relative Thermal Stability of Diamond and Graphite». Angewandte Chemie International Edition. 60 (3): 1546–1549. doi:10.1002/anie.202009897. ISSN 1433-7851. PMID 32970365. S2CID 221888151.
  48. ^ Schewe, Phil & Stein, Ben (March 26, 2004). «Carbon Nanofoam is the World’s First Pure Carbon Magnet». Physics News Update. 678 (1). Archived from the original on March 7, 2012.
  49. ^ Itzhaki, Lior; Altus, Eli; Basch, Harold; Hoz, Shmaryahu (2005). «Harder than Diamond: Determining the Cross-Sectional Area and Young’s Modulus of Molecular Rods». Angew. Chem. Int. Ed. 44 (45): 7432–5. doi:10.1002/anie.200502448. PMID 16240306.
  50. ^ «Researchers Find New Phase of Carbon, Make Diamond at Room Temperature». news.ncsu.edu. 2015-11-30. Archived from the original on 2016-04-06. Retrieved 2016-04-06.
  51. ^ a b c Hoover, Rachel (21 February 2014). «Need to Track Organic Nano-Particles Across the Universe? NASA’s Got an App for That». NASA. Archived from the original on 6 September 2015. Retrieved 2014-02-22.
  52. ^ Lauretta, D.S.; McSween, H.Y. (2006). Meteorites and the Early Solar System II. Space science series. University of Arizona Press. p. 199. ISBN 978-0-8165-2562-1. Archived from the original on 2017-11-22. Retrieved 2017-05-07.
  53. ^ Mark, Kathleen (1987). Meteorite Craters. University of Arizona Press. ISBN 978-0-8165-0902-7.
  54. ^ «Online Database Tracks Organic Nano-Particles Across the Universe». Sci Tech Daily. February 24, 2014. Archived from the original on March 18, 2015. Retrieved 2015-03-10.
  55. ^ William F McDonough The composition of the Earth Archived 2011-09-28 at the Wayback Machine in Majewski, Eugeniusz (2000). Earthquake Thermodynamics and Phase Transformation in the Earth’s Interior. ISBN 978-0126851854.
  56. ^ Yinon Bar-On; et al. (Jun 19, 2018). «The biomass distribution on Earth». PNAS. 115 (25): 6506–6511. Bibcode:2018PNAS..115.6506B. doi:10.1073/pnas.1711842115. PMC 6016768. PMID 29784790.
  57. ^ Fred Pearce (2014-02-15). «Fire in the hole: After fracking comes coal». New Scientist. 221 (2956): 36–41. Bibcode:2014NewSc.221…36P. doi:10.1016/S0262-4079(14)60331-6. Archived from the original on 2015-03-16.
  58. ^ «Wonderfuel: Welcome to the age of unconventional gas» Archived 2014-12-09 at the Wayback Machine by Helen Knight, New Scientist, 12 June 2010, pp. 44–7.
  59. ^ Ocean methane stocks ‘overstated’ Archived 2013-04-25 at the Wayback Machine, BBC, 17 Feb. 2004.
  60. ^ «Ice on fire: The next fossil fuel» Archived 2015-02-22 at the Wayback Machine by Fred Pearce, New Scientist, 27 June 2009, pp. 30–33.
  61. ^ Calculated from file global.1751_2008.csv in «Index of /ftp/ndp030/CSV-FILES». Archived from the original on 2011-10-22. Retrieved 2011-11-06. from the Carbon Dioxide Information Analysis Center.
  62. ^ Rachel Gross (Sep 21, 2013). «Deep, and dank mysterious». New Scientist: 40–43. Archived from the original on 2013-09-21.
  63. ^ Stefanenko, R. (1983). Coal Mining Technology: Theory and Practice. Society for Mining Metallurgy. ISBN 978-0-89520-404-2.
  64. ^ Kasting, James (1998). «The Carbon Cycle, Climate, and the Long-Term Effects of Fossil Fuel Burning». Consequences: The Nature and Implication of Environmental Change. 4 (1). Archived from the original on 2008-10-24.
  65. ^ «Carbon-14 formation». Archived from the original on 1 August 2015. Retrieved 13 October 2014.
  66. ^ Aitken, M.J. (1990). Science-based Dating in Archaeology. pp. 56–58. ISBN 978-0-582-49309-4.
  67. ^ Nichols, Charles R. «Voltatile Products from Carbonaceous Asteroids» (PDF). UAPress.Arizona.edu. Archived from the original (PDF) on 2 July 2016. Retrieved 12 November 2016.
  68. ^ Gannes, Leonard Z.; Del Rio, Carlos Martı́nez; Koch, Paul (1998). «Natural Abundance Variations in Stable Isotopes and their Potential Uses in Animal Physiological Ecology». Comparative Biochemistry and Physiology – Part A: Molecular & Integrative Physiology. 119 (3): 725–737. doi:10.1016/S1095-6433(98)01016-2. PMID 9683412.
  69. ^ «Official SI Unit definitions». Archived from the original on 2007-10-14. Retrieved 2007-12-21.
  70. ^ Bowman, S. (1990). Interpreting the past: Radiocarbon dating. British Museum Press. ISBN 978-0-7141-2047-8.
  71. ^ Brown, Tom (March 1, 2006). «Carbon Goes Full Circle in the Amazon». Lawrence Livermore National Laboratory. Archived from the original on September 22, 2008. Retrieved 2007-11-25.
  72. ^ Libby, W. F. (1952). Radiocarbon dating. Chicago University Press and references therein.
  73. ^ Westgren, A. (1960). «The Nobel Prize in Chemistry 1960». Nobel Foundation. Archived from the original on 2007-10-25. Retrieved 2007-11-25.
  74. ^ «Use query for carbon-8». barwinski.net. Archived from the original on 2005-02-07. Retrieved 2007-12-21.
  75. ^ Watson, A. (1999). «Beaming Into the Dark Corners of the Nuclear Kitchen». Science. 286 (5437): 28–31. doi:10.1126/science.286.5437.28. S2CID 117737493.
  76. ^ Audi, Georges; Bersillon, Olivier; Blachot, Jean; Wapstra, Aaldert Hendrik (1997). «The NUBASE evaluation of nuclear and decay properties» (PDF). Nuclear Physics A. 624 (1): 1–124. Bibcode:1997NuPhA.624….1A. doi:10.1016/S0375-9474(97)00482-X. Archived from the original (PDF) on 2008-09-23.
  77. ^ Ostlie, Dale A. & Carroll, Bradley W. (2007). An Introduction to Modern Stellar Astrophysics. San Francisco (CA): Addison Wesley. ISBN 978-0-8053-0348-3.
  78. ^ Whittet, Douglas C. B. (2003). Dust in the Galactic Environment. CRC Press. pp. 45–46. ISBN 978-0-7503-0624-9.
  79. ^ Pikelʹner, Solomon Borisovich (1977). Star Formation. Springer. p. 38. ISBN 978-90-277-0796-3. Archived from the original on 2012-11-23. Retrieved 2011-06-06.
  80. ^ Mannion, pp. 51–54.
  81. ^ Mannion, pp. 84–88.
  82. ^ Falkowski, P.; Scholes, R. J.; Boyle, E.; Canadell, J.; Canfield, D.; Elser, J.; Gruber, N.; Hibbard, K.; et al. (2000). «The Global Carbon Cycle: A Test of Our Knowledge of Earth as a System». Science. 290 (5490): 291–296. Bibcode:2000Sci…290..291F. doi:10.1126/science.290.5490.291. PMID 11030643. S2CID 1779934.
  83. ^ Smith, T. M.; Cramer, W. P.; Dixon, R. K.; Leemans, R.; Neilson, R. P.; Solomon, A. M. (1993). «The global terrestrial carbon cycle» (PDF). Water, Air, & Soil Pollution. 70 (1–4): 19–37. Bibcode:1993WASP…70…19S. doi:10.1007/BF01104986. S2CID 97265068. Archived (PDF) from the original on 2022-10-11.
  84. ^ Burrows, A.; Holman, J.; Parsons, A.; Pilling, G.; Price, G. (2017). Chemistry3: Introducing Inorganic, Organic and Physical Chemistry. Oxford University Press. p. 70. ISBN 978-0-19-873380-5. Archived from the original on 2017-11-22. Retrieved 2017-05-07.
  85. ^ Mannion pp. 27–51
  86. ^ Mannion pp. 84–91
  87. ^ Levine, Joel S.; Augustsson, Tommy R.; Natarajan, Murali (1982). «The prebiological paleoatmosphere: stability and composition». Origins of Life and Evolution of Biospheres. 12 (3): 245–259. Bibcode:1982OrLi…12..245L. doi:10.1007/BF00926894. PMID 7162799. S2CID 20097153.
  88. ^ Loerting, T.; et al. (2001). «On the Surprising Kinetic Stability of Carbonic Acid». Angew. Chem. Int. Ed. 39 (5): 891–895. doi:10.1002/(SICI)1521-3773(20000303)39:5<891::AID-ANIE891>3.0.CO;2-E. PMID 10760883.
  89. ^ Haldane J. (1895). «The action of carbonic oxide on man». Journal of Physiology. 18 (5–6): 430–462. doi:10.1113/jphysiol.1895.sp000578. PMC 1514663. PMID 16992272.
  90. ^ Gorman, D.; Drewry, A.; Huang, Y. L.; Sames, C. (2003). «The clinical toxicology of carbon monoxide». Toxicology. 187 (1): 25–38. doi:10.1016/S0300-483X(03)00005-2. PMID 12679050.
  91. ^ «Compounds of carbon: carbon suboxide». Archived from the original on 2007-12-07. Retrieved 2007-12-03.
  92. ^ Bayes, K. (1961). «Photolysis of Carbon Suboxide». Journal of the American Chemical Society. 83 (17): 3712–3713. doi:10.1021/ja01478a033.
  93. ^ Anderson D. J.; Rosenfeld, R. N. (1991). «Photodissociation of Carbon Suboxide». Journal of Chemical Physics. 94 (12): 7852–7867. Bibcode:1991JChPh..94.7857A. doi:10.1063/1.460121.
  94. ^ Sabin, J. R.; Kim, H. (1971). «A theoretical study of the structure and properties of carbon trioxide». Chemical Physics Letters. 11 (5): 593–597. Bibcode:1971CPL….11..593S. doi:10.1016/0009-2614(71)87010-0.
  95. ^ Moll N. G.; Clutter D. R.; Thompson W. E. (1966). «Carbon Trioxide: Its Production, Infrared Spectrum, and Structure Studied in a Matrix of Solid CO2«. Journal of Chemical Physics. 45 (12): 4469–4481. Bibcode:1966JChPh..45.4469M. doi:10.1063/1.1727526.
  96. ^ a b Fatiadi, Alexander J.; Isbell, Horace S.; Sager, William F. (1963). «Cyclic Polyhydroxy Ketones. I. Oxidation Products of Hexahydroxybenzene (Benzenehexol)» (PDF). Journal of Research of the National Bureau of Standards Section A. 67A (2): 153–162. doi:10.6028/jres.067A.015. PMC 6640573. PMID 31580622. Archived from the original (PDF) on 2009-03-25. Retrieved 2009-03-21.
  97. ^ Pauling, L. (1960). The Nature of the Chemical Bond (3rd ed.). Ithaca, NY: Cornell University Press. p. 93. ISBN 978-0-8014-0333-0.
  98. ^ Greenwood and Earnshaw, pp. 297–301
  99. ^ Scherbaum, Franz; et al. (1988). ««Aurophilicity» as a consequence of Relativistic Effects: The Hexakis(triphenylphosphaneaurio)methane Dication [(Ph3PAu)6C]2+«. Angew. Chem. Int. Ed. Engl. 27 (11): 1544–1546. doi:10.1002/anie.198815441.
  100. ^ Ritter, Stephen K. «Six bonds to carbon: Confirmed». Chemical & Engineering News. Archived from the original on 2017-01-09.
  101. ^ Yamashita, Makoto; Yamamoto, Yohsuke; Akiba, Kin-ya; Hashizume, Daisuke; Iwasaki, Fujiko; Takagi, Nozomi; Nagase, Shigeru (2005-03-01). «Syntheses and Structures of Hypervalent Pentacoordinate Carbon and Boron Compounds Bearing an Anthracene Skeleton − Elucidation of Hypervalent Interaction Based on X-ray Analysis and DFT Calculation». Journal of the American Chemical Society. 127 (12): 4354–4371. doi:10.1021/ja0438011. ISSN 0002-7863. PMID 15783218.
  102. ^ Shorter Oxford English Dictionary, Oxford University Press
  103. ^ «Chinese made first use of diamond». BBC News. 17 May 2005. Archived from the original on 20 March 2007. Retrieved 2007-03-21.
  104. ^ van der Krogt, Peter. «Carbonium/Carbon at Elementymology & Elements Multidict». Archived from the original on 2010-01-23. Retrieved 2010-01-06.
  105. ^ Ferchault de Réaumur, R.-A. (1722). L’art de convertir le fer forgé en acier, et l’art d’adoucir le fer fondu, ou de faire des ouvrages de fer fondu aussi finis que le fer forgé (English translation from 1956). Paris, Chicago.
  106. ^ «Carbon». Canada Connects. Archived from the original on 2010-10-27. Retrieved 2010-12-07.
  107. ^ a b Senese, Fred (2000-09-09). «Who discovered carbon?». Frostburg State University. Archived from the original on 2007-12-07. Retrieved 2007-11-24.
  108. ^ Giolitti, Federico (1914). The Cementation of Iron and Steel. McGraw-Hill Book Company, inc.
  109. ^ Kroto, H. W.; Heath, J. R.; O’Brien, S. C.; Curl, R. F.; Smalley, R. E. (1985). «C60: Buckminsterfullerene». Nature. 318 (6042): 162–163. Bibcode:1985Natur.318..162K. doi:10.1038/318162a0. S2CID 4314237.
  110. ^ «The Nobel Prize in Chemistry 1996 «for their discovery of fullerenes»«. Archived from the original on 2007-10-11. Retrieved 2007-12-21.
  111. ^ a b c USGS Minerals Yearbook: Graphite, 2009 Archived 2008-09-16 at the Wayback Machine and Graphite: Mineral Commodity Summaries 2011
  112. ^ Harlow, G. E. (1998). The nature of diamonds. Cambridge University Press. p. 223. ISBN 978-0-521-62935-5.
  113. ^ Catelle, W. R. (1911). The Diamond. John Lane Company. p. 159. discussion on alluvial diamonds in India and elsewhere as well as earliest finds
  114. ^ Ball, V. (1881). Diamonds, Gold and Coal of India. London, Truebner & Co. Ball was a Geologist in British service. Chapter I, Page 1
  115. ^ Hershey, J. W. (1940). The Book Of Diamonds: Their Curious Lore, Properties, Tests And Synthetic Manufacture. Kessinger Pub Co. p. 28. ISBN 978-1-4179-7715-4.
  116. ^ Janse, A. J. A. (2007). «Global Rough Diamond Production Since 1870». Gems and Gemology. XLIII (Summer 2007): 98–119. doi:10.5741/GEMS.43.2.98.
  117. ^ Marshall, Stephen; Shore, Josh (2004-10-22). «The Diamond Life». Guerrilla News Network. Archived from the original on 2008-06-09. Retrieved 2008-10-10.
  118. ^ Zimnisky, Paul (21 May 2018). «Global Diamond Supply Expected to Decrease 3.4% to 147M Carats in 2018». Kitco.com. Retrieved 9 November 2020.
  119. ^ Lorenz, V. (2007). «Argyle in Western Australia: The world’s richest diamantiferous pipe; its past and future». Gemmologie, Zeitschrift der Deutschen Gemmologischen Gesellschaft. 56 (1/2): 35–40.
  120. ^ Mannion pp. 25–26
  121. ^ «Microscopic diamond found in Montana». The Montana Standard. 2004-10-17. Archived from the original on 2005-01-21. Retrieved 2008-10-10.
  122. ^ Cooke, Sarah (2004-10-19). «Microscopic Diamond Found in Montana». Livescience.com. Archived from the original on 2008-07-05. Retrieved 2008-09-12.
  123. ^ «Delta News / Press Releases / Publications». Deltamine.com. Archived from the original on 2008-05-26. Retrieved 2008-09-12.
  124. ^ Cantwell, W. J.; Morton, J. (1991). «The impact resistance of composite materials – a review». Composites. 22 (5): 347–62. doi:10.1016/0010-4361(91)90549-V.
  125. ^ Holtzapffel, Ch. (1856). Turning And Mechanical Manipulation. Charles Holtzapffel.
    Internet Archive Archived 2016-03-26 at the Wayback Machine
  126. ^ «Industrial Diamonds Statistics and Information». United States Geological Survey. Archived from the original on 2009-05-06. Retrieved 2009-05-05.
  127. ^ Coelho, R. T.; Yamada, S.; Aspinwall, D. K.; Wise, M. L. H. (1995). «The application of polycrystalline diamond (PCD) tool materials when drilling and reaming aluminum-based alloys including MMC». International Journal of Machine Tools and Manufacture. 35 (5): 761–774. doi:10.1016/0890-6955(95)93044-7.
  128. ^ Harris, D. C. (1999). Materials for infrared windows and domes: properties and performance. SPIE Press. pp. 303–334. ISBN 978-0-8194-3482-1.
  129. ^ Nusinovich, G. S. (2004). Introduction to the physics of gyrotrons. JHU Press. p. 229. ISBN 978-0-8018-7921-0.
  130. ^ Sakamoto, M.; Endriz, J. G.; Scifres, D. R. (1992). «120 W CW output power from monolithic AlGaAs (800 nm) laser diode array mounted on diamond heatsink». Electronics Letters. 28 (2): 197–199. Bibcode:1992ElL….28..197S. doi:10.1049/el:19920123.
  131. ^ Dorfer, Leopold; Moser, M.; Spindler, K.; Bahr, F.; Egarter-Vigl, E.; Dohr, G. (1998). «5200-year old acupuncture in Central Europe?». Science. 282 (5387): 242–243. Bibcode:1998Sci…282..239D. doi:10.1126/science.282.5387.239f. PMID 9841386. S2CID 42284618.
  132. ^ Donaldson, K.; Stone, V.; Clouter, A.; Renwick, L.; MacNee, W. (2001). «Ultrafine particles». Occupational and Environmental Medicine. 58 (3): 211–216. doi:10.1136/oem.58.3.211. PMC 1740105. PMID 11171936.
  133. ^ Carbon Nanoparticles Toxic To Adult Fruit Flies But Benign To Young Archived 2011-11-02 at the Wayback Machine ScienceDaily (Aug. 17, 2009)
  134. ^ «Press Release – Titanic Disaster: New Theory Fingers Coal Fire». www.geosociety.org. Archived from the original on 2016-04-14. Retrieved 2016-04-06.
  135. ^ McSherry, Patrick. «Coal bunker Fire». www.spanamwar.com. Archived from the original on 2016-03-23. Retrieved 2016-04-06.

Bibliography

  • Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  • Mannion, A. M. (12 January 2006). Carbon and Its Domestication. Springer. pp. 1–319. ISBN 978-1-4020-3956-0.

External links

  • Carbon on In Our Time at the BBC
  • Carbon at The Periodic Table of Videos (University of Nottingham)
  • Carbon on Britannica
  • Extensive Carbon page at asu.edu (archived 18 June 2010)
  • Electrochemical uses of carbon (archived 9 November 2001)
  • Carbon—Super Stuff. Animation with sound and interactive 3D-models. (archived 9 November 2012)

Углерод в таблице менделеева занимает 6 место, в 2 периоде.

Символ C
Номер 6
Атомный вес 12.0096000
Латинское название Carboneum
Русское название Углерод

Как самостоятельно построить электронную конфигурацию? Ответ здесь

Электронная схема углерода

C: 1s2 2s2 2p2

Короткая запись:
C: [He]2s2 2p2

Одинаковую электронную конфигурацию имеют
атом углерода и
N+1, O+2

Порядок заполнения оболочек атома углерода (C) электронами:
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d →
5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p.

На подуровне ‘s’ может находиться до 2 электронов, на ‘s’ — до 6, на
‘d’ — до 10 и на ‘f’ до 14

Углерод имеет 6 электронов,
заполним электронные оболочки в описанном выше порядке:

2 электрона на 1s-подуровне

2 электрона на 2s-подуровне

2 электрона на 2p-подуровне

Степень окисления углерода

Атомы углерода в соединениях имеют степени окисления 4, 3, 2, 1, -1, -2, -4.

Степень окисления — это условный заряд атома в соединении: связь в молекуле
между атомами основана на разделении электронов, таким образом, если у атома виртуально увеличивается
заряд, то степень окисления отрицательная (электроны несут отрицательный заряд), если заряд уменьшается,
то степень окисления положительная.

Ионы углерода

Валентность C

Атомы углерода в соединениях проявляют валентность IV, III, II, I.

Валентность углерода характеризует способность атома C к образованию хмических связей.
Валентность следует из строения электронной оболочки атома, электроны, участвующие в образовании
химических соединений называются валентными электронами. Более обширное определение валентности это:

Число химических связей, которыми данный атом соединён с другими атомами

Валентность не имеет знака.

Квантовые числа C

Квантовые числа определяются последним электроном в конфигурации,
для атома C эти числа имеют значение N = 2, L = 1, Ml = 0, Ms = +½

Видео заполнения электронной конфигурации (gif):

Как записать электронную схему углерода

Результат:
электронная схема углерода

Энергия ионизации

Чем ближе электрон к центру атома — тем больше энергии необходимо, что бы его оторвать.
Энергия, затрачиваемая на отрыв электрона от атома называется энергией ионизации и обозначается Eo.
Если не указано иное, то энергия ионизации — это энергия отрыва первого электрона, также существуют энергии
ионизации для каждого последующего электрона.

Энергия ионизации C:
Eo = 1086 кДж/моль

— Что такое ион читайте в статье.


Перейти к другим элементам таблицы менделеева

Где C в таблице менделеева?

Таблица Менделеева

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Строение атома углерода


Строение атома углерода

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Углерод (С) – шестой элемент периодической таблицы Менделеева с атомным весом 12. Элемент относится к неметаллам и имеет изотоп 14С. Строение атома углерода лежит в основе всей органической химии, т. к. все органические вещества включают молекулы углерода.

Атом углерода

Положение углерода в периодической таблице Менделеева:

  • шестой порядковый номер;
  • четвёртая группа;
  • второй период.

Положение углерода в таблице Менделеева

Рис. 1. Положение углерода в таблице Менделеева.

Опираясь на данные из таблицы, можно заключить, что строение атома элемента углерода включает две оболочки, на которых расположено шесть электронов. Валентность углерода, входящего в состав органических веществ, постоянна и равна IV. Это значит, что на внешнем электронном уровне находится четыре электрона, а на внутреннем – два.

Из четырёх электронов два занимают сферическую 2s-орбиталь, а оставшиеся два – 2p-орбиталь в виде гантели. В возбуждённом состоянии один электрон с 2s-орбитали переходит на одну из 2p-орбиталей. При переходе электрона с одной орбитали на другую затрачивается энергия.

Таким образом, возбуждённый атом углерода имеет четыре неспаренных электрона. Его конфигурацию можно выразить формулой 2s12p3. Это даёт возможность образовывать четыре ковалентные связи с другими элементами. Например, в молекуле метана (СН4) углерод образует связи с четырьмя атомами водорода – одна связь между s-орбиталями водорода и углерода и три связи между p-орбиталями углерода и s-орбиталями водорода.

Схему строения атома углерода можно представить в виде записи +6C)2)4 или 1s22s22p2.

Строение атома углерода

Рис. 2. Строение атома углерода.

Физические свойства

Углерод встречается в природе в виде горных пород. Известно несколько аллотропных модификаций углерода:

  • графит;
  • алмаз;
  • карбин;
  • уголь;
  • сажа.

Все эти вещества отличаются строением кристаллической решётки. Наиболее твёрдое вещество – алмаз – имеет кубическую форму углерода. При высоких температурах алмаз превращается в графит с гексагональной структурой.

Кристаллические решётки графита и алмаза

Рис. 3. Кристаллические решётки графита и алмаза.

Химические свойства

Атомное строение углерода и его способность присоединять четыре атома другого вещества определяют химические свойства элемента. Углерод реагирует с металлами, образуя карбиды:

  • Са + 2С → СаС2;
  • Cr + C → CrC;
  • 3Fe + C → Fe3C.

Также реагирует с оксидами металлов:

  • 2ZnO + C → 2Zn + CO2;
  • PbO + C → Pb + CO;
  • SnO2 + 2C → Sn + 2CO.

При высоких температурах углерод реагирует с неметаллами, в частности с водородом, образуя углеводороды:

С + 2Н2 → СН4.

С кислородом углерод образует углекислый газ и угарный газ:

  • С + О2 → СО2;
  • 2С + О2 → 2СО.

Угарный газ также образуется при взаимодействии с водой:

C + H2O → CO + H2.

Концентрированные кислоты окисляют углерод, образуя углекислый газ:

  • 2H2SO4 + C → CO2 + 2SO2 + 2H2O;
  • 4HNO3 + C → CO2 + 4NO2 + 2H2O.

Активность углерода возрастает при нагревании. При низких температурах элемент относительно стабилен.

Заключение

Что мы узнали?

Углерод – типичный неметалл с шестью электронами на s- и р-орбиталях. В активном состоянии приобретает валентность IV и способен присоединять четыре атома вещества. Углерод может быть представлен в виде угля, сажи, графита, алмаза. Элемент реагирует с металлами, неметаллами, кислотами, кислородом, оксидами.

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Углерод, свойства атома, химические и физические свойства.

С 6  Углерод

12,0096-12,0116*     1s2s2p2

Углерод — элемент периодической системы химических элементов Д. И. Менделеева с атомным номером 6. Расположен в 14-й группе (по старой классификации — главной подгруппе четвертой группы), втором периоде периодической системы.

Атом и молекула углерода. Формула углерода. Строение атома углерода

Изотопы и модификации углерода

Свойства углерода (таблица): температура, плотность, давление и пр.

Физические свойства углерода

Химические свойства таллия. Взаимодействие углерода. Химические реакции с углеродом

Получение углерода

Применение углерода

Таблица химических элементов Д.И. Менделеева

Атом и молекула углерода. Формула углерода. Строение атома углерода:

Углерод (лат. Carboneum) – химический элемент периодической системы химических элементов Д. И. Менделеева с обозначением С и атомным номером 6. Расположен в 14-й группе (по старой классификации – главной подгруппе четвертой группы), втором периоде периодической системы.

Углерод – неметалл.

Углерод обладает способностью образовывать полимерные цепочки, что порождает огромный класс соединений на основе углерода, называемых органическими, которых значительно больше, чем неорганических.

Углерод обозначается символом С.

Как простое вещество углерод при нормальных условиях представляет собой матово-чёрное (графит) либо прозрачное (алмаз) вещество. Встречаются и иные аллотропные модификации углерода, имеющие иной внешний вид.

Молекула углерода одноатомна.

Химическая формула углерода С.

Электронная конфигурация атома углерода 1s2 2s2 2p2. Потенциал ионизации (первый электрон) атома углерода равен 1086,45 кДж/моль (11,2602880(11) эВ).

Строение атома углерода. Атом углерода состоит из положительно заряженного ядра (+6), вокруг которого по двум оболочкам движется 6 электронов. При этом 2 электрона находятся на внутреннем уровне, а 4 электрона – на внешнем. Поскольку углерод расположен во втором периоде, оболочек всего две. Первая – внутренняя оболочка представлена s-орбиталью. Вторая – внешняя оболочка представлены s- и р-орбиталями. На внешнем энергетическом уровне атома углерода на 2s-орбитали находятся два спаренных электрона, на 2p-орбитали находятся два неспаренных электрона. В свою очередь ядро атома углерода состоит из 6 протонов и 6 нейтронов. Углерод относится к элементам p-семейства.

Радиус атома углерода (вычисленный) составляет 67 пм.

Атомная масса атома углерода составляет 12,0096-12,0116 а. е. м.

Содержание углерода в земной коре составляет 0,18 %, в морской воде и океане – 0,0028 %.

Изотопы и модификации углерода:

Свойства углерода (таблица): температура, плотность, давление и пр.:

Подробные сведения на сайте ChemicalStudy.ru

100 Общие сведения  
101 Название Углерод
102 Прежнее название
103 Латинское название Carboneum
104 Английское название Carbon
105 Символ С
106 Атомный номер (номер в таблице) 6
107 Тип Неметалл
108 Группа
109 Открыт Известен с древних времен
110 Год открытия 3750 г. до н.э.
111 Внешний вид и пр. Матово-чёрный (графит) либо прозрачный (алмаз) минерал
112 Происхождение Природный материал
113 Модификации
114 Аллотропные модификации более 10 аллотропных модификаций углерода. Из них наиболее известны:

алмаз,

графит,

– аморфный углерод,

– лонсдейлит,

карбин,

графен,

фуллерен,

углеродные нанотрубки,

углеродные нановолокна,

– астралены,

– стеклоуглерод,

колоссальные углеродные трубки,

– углеродные нанопочки,

углеродная нанопена,

– Q-углерод,

графин

115 Температура и иные условия перехода аллотропных модификаций друг в друга
116 Конденсат Бозе-Эйнштейна
117 Двумерные материалы Графен, графин
118 Содержание в атмосфере и воздухе (по массе) в составе углекислого газа и метана, содержание которых в атмосфере и воздухе 0,046 % и 0,000084 % соответственно
119 Содержание в земной коре (по массе) 0,18 %
120 Содержание в морях и океанах (по массе) 0,0028 %
121 Содержание во Вселенной и космосе (по массе) 0,5 %
122 Содержание в Солнце (по массе) 0,3 %
123 Содержание в метеоритах (по массе) 1,5 %
124 Содержание в организме человека (по массе) 23 %
200 Свойства атома  
201 Атомная масса (молярная масса)* 12,0096-12,0116 а. е. м. (г/моль)
202 Электронная конфигурация 1s2s2p2
203 Электронная оболочка K2 L4 M0 N0 O0 P0 Q0 R0

Электронная оболочка углерода

204 Радиус атома (вычисленный) 67 пм
205 Эмпирический радиус атома 70 пм
206 Ковалентный радиус* 76 пм
207 Радиус иона (кристаллический) C4+

29 (4) пм,

30 (6) пм

(в скобках указано координационное число – характеристика, которая определяет число ближайших частиц (ионов или атомов) в молекуле или кристалле)

208 Радиус Ван-дер-Ваальса 170 пм
209 Электроны, Протоны, Нейтроны 6 электронов, 6 протонов, 6 нейтронов
210 Семейство (блок) элемент p-семейства
211 Период в периодической таблице 2
212 Группа в периодической таблице 14-ая группа (по старой классификации – главная подгруппа 4-ой группы)
213 Эмиссионный спектр излучения Спектр_Углерода
300 Химические свойства  
301 Степени окисления -4 , -3 , -2 , -1 , 0 , +1, +2, +3, +4
302 Валентность II, IV
303 Электроотрицательность 2,55 (шкала Полинга)
304 Энергия ионизации (первый электрон) 1086,45 кДж/моль (11,2602880(11) эВ)
305 Электродный потенциал
306 Энергия сродства атома к электрону 121,7763(1) кДж/моль (1,2621226(11) эВ) – углерод 12C,

121,7755(2) кДж/моль (1,2621136(12) эВ) – углерод 13C

400 Физические свойства
401 Плотность* 1,8-2,1 г/см3 (при 20 °C и иных стандартных условиях, состояние вещества – твердое тело) – аморфный углерод,

2,267 г/см3 (при 20 °C и иных стандартных условиях, состояние вещества – твердое тело) – графит,

3,515 г/см3 (при 20 °C и иных стандартных условиях, состояние вещества – твердое тело) – алмаз

402 Температура плавления
403 Температура кипения
404 Температура сублимации 3642 °C (3915 K, 6588 °F) – графит
405 Температура разложения 1000 °C (1273 K, 1832 °F) – алмаз. Продукты разложения алмаза – графит
406 Температура самовоспламенения смеси газа с воздухом
407 Удельная теплота плавления (энтальпия плавления ΔHпл)
408 Удельная теплота испарения (энтальпия кипения ΔHкип) 715  кДж/моль (сублимация)
409 Удельная теплоемкость при постоянном давлении
410 Молярная теплоёмкость* 8,517Дж/(K·моль) – графит,

6,155 Дж/(K·моль) – алмаз,

411 Молярный объём 5,314469 см³/моль – графит,

3,42 см³/моль – алмаз

412 Теплопроводность 119-165 Вт/(м·К) (при стандартных условиях) – графит,

900-2300 Вт/(м·К) (при стандартных условиях) – алмаз

500 Кристаллическая решётка
511 Кристаллическая решётка #1 α-графит
512 Структура решётки Гексагональная

Кристаллическая решетка углерода_графит

513 Параметры решётки a = 2,46 Å, c = 6,71 Å
514 Отношение c/a 2,73
515 Температура Дебая
516 Название пространственной группы симметрии P63/mmc
517 Номер пространственной группы симметрии 194
521 Кристаллическая решётка #2 Алмаз
522 Структура решётки Кубическая алмазная

Кристаллическая решетка углерода_алмаз

523 Параметры решётки a = 3,567 Å
524 Отношение c/a
525 Температура Дебая 1860 K
526 Название пространственной группы симметрии Fd_ 3m
527 Номер пространственной группы симметрии 225
900 Дополнительные сведения
901 Номер CAS 7782-42-5 – графит,

7782-40-3 – алмаз

Примечание:

201* Указан диапазон значений атомной массы в связи с различной распространённостью изотопов данного элемента в природе.

206* Ковалентный радиус углерода согласно [1] составляет для sp3 – 77 пм, для sp2 – 73 пм, sp – 69 пм, согласно [3] составляет 77 пм.

401* Плотность графита согласно [3] составляет 2,25 г/см3 (при 0 °C и иных стандартных условиях, состояние вещества – твердое тело).

410* Молярная теплоемкость графита согласно [3] составляет 8,54 Дж/(K·моль).

Физические свойства углерода:

Химические свойства углерода. Взаимодействие углерода. Химические реакции с углеродом:

Получение углерода:

Применение углерода:

Таблица химических элементов Д.И. Менделеева

  1. 1. Водород
  2. 2. Гелий
  3. 3. Литий
  4. 4. Бериллий
  5. 5. Бор
  6. 6. Углерод
  7. 7. Азот
  8. 8. Кислород
  9. 9. Фтор
  10. 10. Неон
  11. 11. Натрий
  12. 12. Магний
  13. 13. Алюминий
  14. 14. Кремний
  15. 15. Фосфор
  16. 16. Сера
  17. 17. Хлор
  18. 18. Аргон
  19. 19. Калий
  20. 20. Кальций
  21. 21. Скандий
  22. 22. Титан
  23. 23. Ванадий
  24. 24. Хром
  25. 25. Марганец
  26. 26. Железо
  27. 27. Кобальт
  28. 28. Никель
  29. 29. Медь
  30. 30. Цинк
  31. 31. Галлий
  32. 32. Германий
  33. 33. Мышьяк
  34. 34. Селен
  35. 35. Бром
  36. 36. Криптон
  37. 37. Рубидий
  38. 38. Стронций
  39. 39. Иттрий
  40. 40. Цирконий
  41. 41. Ниобий
  42. 42. Молибден
  43. 43. Технеций
  44. 44. Рутений
  45. 45. Родий
  46. 46. Палладий
  47. 47. Серебро
  48. 48. Кадмий
  49. 49. Индий
  50. 50. Олово
  51. 51. Сурьма
  52. 52. Теллур
  53. 53. Йод
  54. 54. Ксенон
  55. 55. Цезий
  56. 56. Барий
  57. 57. Лантан
  58. 58. Церий
  59. 59. Празеодим
  60. 60. Неодим
  61. 61. Прометий
  62. 62. Самарий
  63. 63. Европий
  64. 64. Гадолиний
  65. 65. Тербий
  66. 66. Диспрозий
  67. 67. Гольмий
  68. 68. Эрбий
  69. 69. Тулий
  70. 70. Иттербий
  71. 71. Лютеций
  72. 72. Гафний
  73. 73. Тантал
  74. 74. Вольфрам
  75. 75. Рений
  76. 76. Осмий
  77. 77. Иридий
  78. 78. Платина
  79. 79. Золото
  80. 80. Ртуть
  81. 81. Таллий
  82. 82. Свинец
  83. 83. Висмут
  84. 84. Полоний
  85. 85. Астат
  86. 86. Радон
  87. 87. Франций
  88. 88. Радий
  89. 89. Актиний
  90. 90. Торий
  91. 91. Протактиний
  92. 92. Уран
  93. 93. Нептуний
  94. 94. Плутоний
  95. 95. Америций
  96. 96. Кюрий
  97. 97. Берклий
  98. 98. Калифорний
  99. 99. Эйнштейний
  100. 100. Фермий
  101. 101. Менделеевий
  102. 102. Нобелий
  103. 103. Лоуренсий
  104. 104. Резерфордий
  105. 105. Дубний
  106. 106. Сиборгий
  107. 107. Борий
  108. 108. Хассий
  109. 109. Мейтнерий
  110. 110. Дармштадтий
  111. 111. Рентгений
  112. 112. Коперниций
  113. 113. Нихоний
  114. 114. Флеровий
  115. 115. Московий
  116. 116. Ливерморий
  117. 117. Теннессин
  118. 118. Оганесон

Таблица химических элементов Д.И. Менделеева

Источники:

  1. https://en.wikipedia.org/wiki/Carbon
  2. https://de.wikipedia.org/wiki/Kohlenstoff
  3. https://ru.wikipedia.org/wiki/Углерод
  4. http://chemister.ru/Database/properties.php?dbid=1&id=216, http://chemister.ru/Database/properties.php?dbid=1&id=217
  5. https://chemicalstudy.ru/uglerod-svoystva-atoma-himicheskie-i-fizicheskie-svoystva/

Примечание: © Фото https://www.pexels.com, https://pixabay.com

углерод атомная масса степень окисления валентность плотность температура кипения плавления физические химические свойства структура теплопроводность электропроводность кристаллическая решетка
атом нарисовать строение число протонов в ядре строение электронных оболочек электронная формула конфигурация схема строения электронной оболочки заряд ядра состав масса орбита уровни модель радиус энергия электрона переход скорость спектр длина волны молекулярная масса объем атома
электронные формулы сколько атомов в молекуле углерода
сколько электронов в атоме свойства металлические неметаллические термодинамические 

Коэффициент востребованности
1 936

Содержание статьи

  • Историческая справка.
  • Аллотропия.
  • Строение атома углерода.
  • Стандартная атомная масса.
  • Химические свойства углерода и некоторых его соединений.
  • Субоксид углерода
  • Электронное строение оксидов углерода.
  • Угольная кислота.
  • Карбонаты.
  • Галогениды углерода.
  • Реакции графита.
  • Карбиды.
  • Азотпроизводные углерода.
  • Карбонилы.
  • Углеводороды.

УГЛЕРОД, С (carboneum), неметаллический химический элемент IVA группы (C, Si, Ge, Sn, Pb) периодической системы элементов. Встречается в природе в виде кристаллов алмаза (рис. 1), графита или фуллерена и других форм и входит в состав органических (уголь, нефть, организмы животных и растений и др.) и неорганических веществ (известняк, пищевая сода и др.).

Углерод широко распространен, но содержание его в земной коре всего 0,19%.

Рис. 1. СТРУКТУРА алмаза (а) и графита (б).
 

Рис. 1. СТРУКТУРА алмаза (а) и графита (б).

Углерод широко используется в виде простых веществ. Кроме драгоценных алмазов, являющихся предметом ювелирных украшений, большое значение имеют промышленные алмазы – для изготовления шлифовального и режущего инструмента.

Древесный уголь и другие аморфные формы углерода применяются для обесцвечивания, очистки, адсорбции газов, в областях техники, где требуются адсорбенты с развитой поверхностью. Карбиды, соединения углерода с металлами, а также с бором и кремнием (например, Al4C3, SiC, B4C) отличаются высокой твердостью и используются для изготовления абразивного и режущего инструмента. Углерод входит в состав сталей и сплавов в элементном состоянии и в виде карбидов. Насыщение поверхности стальных отливок углеродом при высокой температуре (цементация) значительно увеличивает поверхностную твердость и износостойкость. См. также СПЛАВЫ.

В природе существует множество различных форм графита; некоторые получены искусственно; имеются аморфные формы (например, кокс и древесный уголь). Сажа, костяной уголь, ламповая сажа, ацетиленовая сажа образуются при сжигании углеводородов при недостатке кислорода. Так называемый белый углерод получается сублимацией пиролитического графита при пониженном давлении – это мельчайшие прозрачные кристаллики графитовых листочков с заостренными кромками.

Историческая справка.

Графит, алмаз и аморфный углерод известны с древности. Издавна известно, что графитом можно маркировать другой материал, и само название «графит», происходящее от греческого слова, означающего «писать», предложено А.Вернером в 1789. Однако история графита запутана, часто за него принимали вещества, обладающие сходными внешними физическими свойствами, например молибденит (сульфид молибдена), одно время считавшийся графитом. Среди других названий графита известны «черный свинец», «карбидное железо», «серебристый свинец». В 1779 К.Шееле установил, что графит можно окислить воздухом с образованием углекислого газа.

Впервые алмазы нашли применение в Индии, а в Бразилии драгоценные камни приобрели коммерческое значение в 1725; месторождения в Южной Африке были открыты в 1867. В 20 в. основными производителями алмазов являются ЮАР, Заир, Ботсвана, Намибия, Ангола, Сьерра-Леоне, Танзания и Россия. Искусственные алмазы, технология которых была создана в 1970, производятся для промышленных целей.

Аллотропия.

Если структурные единицы вещества (атомы для одноатомных элементов или молекулы для полиатомных элементов и соединений) способны соединяться друг с другом в более чем одной кристаллической форме, это явление называется аллотропией. У углерода три аллотропические модификации – алмаз, графит и фуллерен. В алмазе каждый атом углерода имеет 4 тетраэдрически расположенных соседа, образуя кубическую структуру (рис. 1,а). Такая структура отвечает максимальной ковалентности связи, и все 4 электрона каждого атома углерода образуют высокопрочные связи С–С, т.е. в структуре отсутствуют электроны проводимости. Поэтому алмаз отличается отсутствием проводимости, низкой теплопроводностью, высокой твердостью; он самый твердый из известных веществ (рис. 2). На разрыв связи С–С (длина связи 1,54 Å, отсюда ковалентный радиус 1,54/2 = 0,77 Å) в тетраэдрической структуре требуются большие затраты энергии, поэтому алмаз, наряду с исключительной твердостью, характеризуется высокой температурой плавления (3550° C).

Рис. 2. СРАВНИТЕЛЬНАЯ ТВЕРДОСТЬ МИНЕРАЛОВ по шкале Мооса (качественная оценка) (1) и по шкале истинной относительной твердости (количественная оценка) (2). По кривой 2 видно, насколько алмаз тверже наиболее твердых материалов

Другой аллотропической формой углерода является графит, сильно отличающийся от алмаза по свойствам. Графит – мягкое черное вещество из легко слоящихся кристалликов, отличающееся хорошей электропроводностью (электрическое сопротивление 0,0014 Ом·см). Поэтому графит применяется в дуговых лампах и печах (рис. 3), в которых необходимо создавать высокие температуры. Графит высокой чистоты применяют в ядерных реакторах в качестве замедлителя нейтронов. Температура плавления его при повышенном давлении равна 3527° C. При обычном давлении графит сублимируется (переходит из твердого состояния в газ) при 3780° C.

Рис. 3. СХЕМА КОКСОВОЙ ПЕЧИ для коксования битумных углей в отсутствие воздуха (газовоздушная обогревающая смесь не попадает в камеру, где находится коксуемый уголь).

Структура графита (рис. 1,б) представляет собой систему конденсированных гексагональных колец с длиной связи 1,42 Å (значительно короче, чем в алмазе), но при этом каждый атом углерода имеет три (а не четыре, как в алмазе) ковалентные связи с тремя соседями, а четвертая связь (3,4 Å) слишком длинна для ковалентной связи и слабо связывает параллельно уложенные слои графита между собой. Именно четвертый электрон углерода определяет тепло- и электропроводность графита – эта более длинная и менее прочная связь формирует меньшую компактность графита, что отражается в меньшей твердости его в сравнении с алмазом (плотность графита 2,26 г/см3, алмаза – 3,51 г/см3). По той же причине графит скользкий на ощупь и легко отделяет чешуйки вещества, что и используется для изготовления смазки и грифелей карандашей. Свинцовый блеск грифеля объясняется в основном наличием графита.

Волокна углерода имеют высокую прочность и могут использоваться для изготовления искусственного шелка или другой пряжи с высоким содержанием углерода.

При высоких давлении и температуре в присутствии катализатора, например железа, графит может превращаться в алмаз. Этот процесс реализован для промышленного получения искусственных алмазов. Кристаллы алмаза растут на поверхности катализатора. Равновесие графит алмаз существует при 15 000 атм и 300 K или при 4000 атм и 1500 K. Искусственные алмазы можно получать и из углеводородов.

К аморфным формам углерода, не образующим кристаллов, относят древесный уголь, получаемый нагревом дерева без доступа воздуха, ламповую и газовую сажу, образующуюся при низкотемпературном сжигании углеводородов при недостатке воздуха и конденсируемую на холодной поверхности, костяной уголь – примесь к фосфату кальция в процессе деструкции костной ткани, а также каменный уголь (природное вещество с примесями) и кокс, сухой остаток, получаемый при коксовании топлив методом сухой перегонки каменного угля или нефтяных остатков (битуминозных углей), т.е. нагреванием без доступа воздуха. Кокс применяется для выплавки чугуна, в черной и цветной металлургии. При коксовании образуются также газообразные продукты – коксовый газ (H2, CH4, CO и др.) и химические продукты, являющиеся сырьем для получения бензина, красок, удобрений, лекарственных препаратов, пластмасс и т.д. Схема основного аппарата для производства кокса – коксовой печи – приведена на рис. 3.

Различные виды угля и сажи отличаются развитой поверхностью и поэтому используются как адсорбенты для очистки газа, жидкостей, а также как катализаторы. Для получения различных форм углерода применяют специальные методы химической технологии. Искусственный графит получают прокаливанием антрацита или нефтяного кокса между углеродными электродами при 2260°С (процесс Ачесона) и используют в производстве смазочных материалов и электродов, в частности для электролитического получения металлов.

Строение атома углерода.

Ядро наиболее стабильного изотопа углерода массой 12 (распространенность 98,9%) имеет 6 протонов и 6 нейтронов (12 нуклонов), расположенных тремя квартетами, каждый содержит 2 протона и два нейтрона аналогично ядру гелия. Другой стабильный изотоп углерода – 13C (ок. 1,1%), а в следовых количествах существует в природе нестабильный изотоп 14C с периодом полураспада 5730 лет, обладающий b-излучением. В нормальном углеродном цикле живой материи участвуют все три изотопа в виде СO2. После смерти живого организма расход углерода прекращается и можно датировать С-содержащие объекты, измеряя уровень радиоактивности 14С. Снижение b-излучения 14CO2 пропорционально времени, прошедшему с момента смерти. В 1960 У.Либби за исследования с радиоактивным углеродом был удостоен Нобелевской премии. См. также ДАТИРОВКА ПО РАДИОАКТИВНОСТИ.

В основном состоянии 6 электронов углерода образуют электронную конфигурацию 1s22s22px12py12pz0. Четыре электрона второго уровня являются валентными, что соответствует положению углерода в IVA группе периодической системы (см. ПЕРИОДИЧЕСКАЯ СИСТЕМА ЭЛЕМЕНТОВ). Поскольку для отрыва электрона от атома в газовой фазе требуется большая энергия (ок. 1070 кДж/моль), углерод не образует ионные связи с другими элементами, так как для этого необходим был бы отрыв электрона с образованием положительного иона. Имея электроотрицательность, равную 2,5, углерод не проявляет и сильного сродства к электрону, соответственно не являясь активным акцептором электронов. Поэтому он не склонен к образованию частицы с отрицательным зарядом. Но с частично ионным характером связи некоторые соединения углерода существуют, например, карбиды. В соединениях углерод проявляет степень окисления 4. Чтобы четыре электрона смогли участвовать в образовании связей, необходимо распаривание 2s-электронов и перескок одного из этих электронов на 2pz-орбиталь; при этом образуются 4 тетраэдрические связи с углом между ними 109°. В соединениях валентные электроны углерода лишь частично оттянуты от него, поэтому углерод образует прочные ковалентные связи между соседними атомами типа С–С с помощью общей электронной пары. Энергия разрыва такой связи равна 335 кДж/моль, тогда как для связи Si–Si она составляет всего 210 кДж/моль, поэтому длинные цепочки –Si–Si– неустойчивы. Ковалентный характер связи сохраняется даже в соединениях высокореакционноспособных галогенов с углеродом, CF4 и CCl4. Углеродные атомы способны предоставлять на образование связи более одного электрона от каждого атома углерода; так образуются двойная С=С и тройная СєС связи. Другие элементы также образуют связи между своими атомами, но только углерод способен образовывать длинные цепи. Поэтому для углерода известны тысячи соединений, называемых углеводородами, в которых углерод связан с водородом и другими углеродными атомами, образуя длинные цепи или кольцевые структуры. См. ХИМИЯ ОРГАНИЧЕСКАЯ.

В этих соединениях возможно замещение водорода на другие атомы, наиболее часто на кислород, азот и галогены с образованием множества органических соединений. Важное значение среди них занимают фторуглеводороды – углеводороды, в которых водород замещен на фтор. Такие соединения чрезвычайно инертны, и их используют как пластичные и смазочные материалы (фторуглероды, т.е. углеводороды, в которых все атомы водорода замещены на атомы фтора) и как низкотемпературные хладагенты (хладоны, или фреоны, – фторхлоруглеводороды).

В 1980-х годах физиками США был обнаружены очень интересные соединения углерода, в которых атомы углерода соединены в 5- или 6-угольники, образующие молекулу С60 по форме полого шара, имеющего совершенную симметрию футбольного мяча. Поскольку такая конструкция лежит в основе «геодезического купола», изобретенного американским архитектором и инженером Бакминстером Фуллером, новый класс соединений был назван «бакминстерфуллеренами» или «фуллеренами» (а также более коротко – «фазиболами» или «бакиболами»). Фуллерены – третья модификация чистого углерода (кроме алмаза и графита), состоящая из 60 или 70 (и даже более) атомов, – была получена действием лазерного излучения на мельчайшие частички углерода. Фуллерены более сложной формы состоят из нескольких сотен атомов углерода. Диаметр молекулы С60 ~ 1нм. В центре такой молекулы достаточно пространства для помещения большого атома урана. См. также ФУЛЛЕРЕНЫ.

Стандартная атомная масса.

В 1961 Международные союзы теоретической и прикладной химии (ИЮПАК) и по физике приняли за единицу атомной массы массу изотопа углерода 12C, упразднив существовавшую до того кислородную шкалу атомных масс. Атомная масса углерода в этой системе равна 12,011, так как она является средней для трех природных изотопов углерода с учетом их распространенности в природе. См. АТОМНАЯ МАССА.

Химические свойства углерода и некоторых его соединений.

Некоторые физические и химические свойства углерода приведены в статье ЭЛЕМЕНТЫ ХИМИЧЕСКИЕ. Реакционная способность углерода зависит от его модификации, температуры и дисперсности. При низких температурах все формы углерода достаточно инертны, но при нагревании окисляются кислородом воздуха, образуя оксиды:

Мелкодисперсный углерод в избытке кислорода способен взрываться при нагревании или от искры. Кроме прямого окисления существуют более современные методы получения оксидов.

Субоксид углерода

C3O2 образуется при дегидратации малоновой кислоты над P4O10:

C3O2 имеет неприятный запах, легко гидролизуется, вновь образуя малоновую кислоту.

ФИЗИЧЕСКИЕ СВОЙСТВА ОКСИДОВ УГЛЕРОДА

 

CO

CO2

C3O2

Молекулярная масса

28,010

44,010

68,030

Температура кипения, °С

–192

–56,6

–7

Температура замерзания, °С

–199

–78,5

–111,3

Плотность, г/л (при 0° С)

1,250

1,977

1,114

Растворимость, объем/объем воды (при 0° С)

0,03

1,7

Реагирует с водой

Монооксид углерода(II) СО образуется при окислении любой модификации углерода в условиях недостатка кислорода. Реакция экзотермична, выделяется 111,6 кДж/моль. Кокс при температуре белого каления реагирует с водой: C + H2O = CO + H2; образующаяся газовая смесь называется «водяной газ» и является газообразным топливом. СO образуется также при неполном сгорании нефтепродуктов, в заметных количествах содержится в автомобильных выхлопах, получается при термической диссоциации муравьиной кислоты:

Степень окисления углерода в СО равна +2, а поскольку углерод более устойчив в степени окисления +4, то СО легко окисляется кислородом до CO2: CO + O2 → CO2, эта реакция сильно экзотермична (283 кДж/моль). СО применяют в промышленности в смеси с H2 и другими горючими газами в качестве топлива или газообразного восстановителя. При нагревании до 500° C CO в заметной степени образует С и CO2, но при 1000° C равновесие устанавливается при малых концентрациях СO2. CO реагирует с хлором, образуя фосген – COCl2, аналогично протекают реакции с другими галогенами, в реакции с серой получается сульфид карбонила COS, с металлами (M) СO образует карбонилы различного состава M(CO)x, являющиеся комплексными соединениями. Карбонил железа образуется при взаимодействии гемоглобина крови с CO, препятствуя реакции гемоглобина с кислородом, так как карбонил железа – более прочное соединение. В результате блокируется функция гемоглобина как переносчика кислорода к клеткам, которые при этом погибают (и в первую очередь поражаются клетки мозга). (Отсюда еще одно название СО – «угарный газ»). Уже 1% (об.) СO в воздухе опасен для человека, если он находится в такой атмосфере более 10 мин. Некоторые физические свойства СО приведены в таблице.

Диоксид углерода, или оксид углерода(IV)CO2 образуется при сгорании элементного углерода в избытке кислорода c выделением тепла (395 кДж/моль). CO2 (тривиальное название – «углекислый газ») образуется также при полном окислении СО, нефтепродуктов, бензина, масел и др. органических соединений. При растворении карбонатов в воде в результате гидролиза также выделяется СО2:

Такой реакцией часто пользуются в лабораторной практике для получения CO2. Этот газ можно получить и при прокаливании бикарбонатов металлов:

при газофазном взаимодействии перегретого пара с СО:

при сжигании углеводородов и их кислородпроизводных, например:

Аналогично окисляются пищевые продукты в живом организме с выделением тепловой и других видов энергии. При этом окисление протекает в мягких условиях через промежуточные стадии, но конечные продукты те же – СO2 и H2O, как, например, при разложении сахаров под действием ферментов, в частности при ферментации глюкозы:

Многотоннажное производство углекислого газа и оксидов металлов осуществляется в промышленности термическим разложением карбонатов:

CaO в больших количествах используется в технологии производства цемента. Термическая стабильность карбонатов и затраты теплоты на их разложение по этой схеме возрастают в ряду CaCO3 < SrCO3 < BaCO3. Углекислый газ – химически неактивное соединение, однако в некоторых случаях он может поддерживать процесс горения, например, магний горит в среде СО2. Материалы, горящие при низких температурах – дерево, нефтепродукты, бумага и др., – не горят в среде СО2. Поэтому, а также из-за большей, чем у воздуха, плотности (СO2 в 1,5 раза тяжелее), углекислый газ используют в огнетушителях. В твердом состоянии СО2 известен как «сухой лед». При высоких концентрациях CO2 опасен для человека, так как блокирует доступ кислорода (см. также ПОЖАРНАЯ ПРОФИЛАКТИКА И ПРОТИВОПОЖАРНАЯ ЗАЩИТА).

Электронное строение оксидов углерода.

Электронное строение любого оксида углерода можно описать тремя равновероятными схемами с различным расположением электронных пар – тремя резонансными формами:

Все оксиды углерода имеют линейное строение.

Угольная кислота.

При взаимодействии СO2 с водой образуется угольная кислота H2CO3. В насыщенном растворе CO2 (0,034 моль/л) только часть молекул образует H2CO3, а бóльшая часть CO2 находится в гидратированном состоянии CO2ЧH2O.

Карбонаты.

Карбонаты образуются при взаимодействии оксидов металлов с CO2, например, Na2O + CO2 Na2CO3.

За исключением карбонатов щелочных металлов, остальные практически нерастворимы в воде, а карбонат кальция частично растворим в угольной кислоте или растворе CO2 в воде под давлением:

CaCO3 + H2O + CO2 Ca(HCO3)2

Эти процессы происходят в подземных водах, протекающих через пласт известняка. В условиях низкого давления и испарения из грунтовых вод, содержащих Ca(HCO3)2, осаждается CaCO3. Так происходит рост сталактитов и сталагмитов в пещерах. Окраска этих интересных геологических образований объясняется присутствием в водах примесей ионов железа, меди, марганца и хрома. Углекислый газ реагирует с гидроксидами металлов и их растворами с образованием гидрокарбонатов, например:

NaOH + CO2 → NaHCO3

которые при нагревании разлагаются с выделением СО2:

2NaHCO3 → Na2CO3 + H2O + CO2

Карбонат натрия, или соду, производят в содовой промышленности в больших количествах преимущественно методом Сольве:

Другим методом соду получают из CO2 и NaOH

Карбонат-ион CO32– имеет плоское строение с углом O–C–O, равным 120°, и длиной СО-связи 1,31 Å (см. также ЩЕЛОЧЕЙ ПРОИЗВОДСТВО).

Галогениды углерода.

Углерод непосредственно реагирует с галогенами при нагревании, образуя тетрагалогениды, но скорость реакции и выход продукта невелики. Поэтому галогениды углерода получают другими методами, например, хлорированием дисульфида углерода получают CCl4:

CS2 + 2Cl2 ® CCl4 + 2S

Тетрахлорид CCl4 – негорючее вещество, используется в качестве растворителя в процессах сухой чистки, но не рекомендуется применять его как пламегаситель, так как при высокой температуре происходит образование ядовитого фосгена (газообразное отравляющее вещество). Сам ССl4 также ядовит и при вдыхании в заметных количествах может вызвать отравление печени. СCl4 образуется и по фотохимической реакции между метаном СH4 и Сl2; при этом возможно образование продуктов неполного хлорирования метана – CHCl3, CH2Cl2 и CH3Cl. Аналогично протекают реакции и с другими галогенами.

Реакции графита.

Графит как модификация углерода, отличающаяся большими расстояниями между слоями гексагональных колец, вступает в необычные реакции, например, щелочные металлы, галогены и некоторые соли (FeCl3) проникают между слоями, образуя соединения типа KC8, KC16 (называемые соединениями внедрения, включения или клатратами). Сильные окислители типа KClO3 в кислой среде (серной или азотной кислоты) образуют вещества с большим объемом кристаллической решетки (до 6 Å между слоями), что объясняется внедрением кислородных атомов и образованием соединений, на поверхности которых в результате окисления образуются карбоксильные группы (–СООН) – соединения типа оксидированного графита или меллитовой (бензолгексакарбоновой) кислоты С6(COOH)6. В этих соединениях отношение С:O может изменяться от 6:1 до 6:2,5.

Карбиды.

Углерод образует с металлами, бором и кремнием разнообразные соединения, называемые карбидами. Наиболее активные металлы (IA–IIIA подгрупп) образуют солеподобные карбиды, например Na2C2, CaC2, Mg4C3, Al4C3. В промышленности карбид кальция получают из кокса и известняка по следующим реакциям:

Карбиды неэлектропроводны, почти бесцветны, гидролизуются с образованием углеводородов, например

CaC2 + 2H2O = C2H2 + Ca(OH)2

Образующийся по реакции ацетилен C2H2 служит исходным сырьем в производстве многих органических веществ. Этот процесс интересен, так как он представляет переход от сырья неорганической природы к синтезу органических соединений. Карбиды, образующие при гидролизе ацетилен, называются ацетиленидами. В карбидах кремния и бора (SiC и B4C) связь между атомами ковалентная. Переходные металлы (элементы B-подгрупп) при нагревании с углеродом тоже образуют карбиды переменного состава в трещинах на поверхности металла; связь в них близка к металлической. Некоторые карбиды такого типа, например WC, W2C, TiC и SiC, отличаются высокой твердостью и тугоплавкостью, обладают хорошей электропроводностью. Например, NbC, TaC и HfC – наиболее тугоплавкие вещества (т.пл. = 4000–4200° С), карбид диниобия Nb2C – сверхпроводник при 9,18 К, TiC и W2C по твердости близки алмазу, а твердость B4C (структурного аналога алмаза) составляет 9,5 по шкале Мооса (см. рис. 2). Инертные карбиды образуются, если радиус переходного металла < 1,3 Å, а при большем радиусе образуются реакционноспособные карбиды типа Mn3C, Ni3C3, Fe3C.

Азотпроизводные углерода.

К этой группе относится мочевина NH2CONH2 – азотное удобрение, применяемое в виде раствора. Мочевину получают из NH3 и CO2 при нагревании под давлением:

Дициан (CN)2 по многим свойствам подобен галогенам и его часто называют псевдогалоген. Дициан получают мягким окислением цианид-иона кислородом, пероксидом водорода или ионом Cu2+: 2CN ® (CN)2 + 2e.

Цианид-ион, являясь донором электронов, легко образует комплексные соединения с ионами переходных металлов. Подобно СО, цианид-ион является ядом, связывая жизненно важные соединения железа в живом организме. Цианидные комплексные ионы имеют общую формулу [M(CN)x]–0,5x, где х – координационное число металла (комплексообразователя), эмпирически равно удвоенному значению степени окисления иона металла. Примерами таких комплексных ионов являются (строение некоторых ионов приведено ниже) тетрацианоникелат(II)-ион [Ni(CN)4]2–, гексацианоферрат(III) [Fe(CN)6]3–, дицианоаргентат [Ag(CN)2]:

Карбонилы.

Монооксид углерода способен непосредственно реагировать со многими металлами или ионами металлов, образуя комплексные соединения, называемые карбонилами, например Ni(CO)4, Fe(CO)5, Fe2(CO)9, [Fe(CO)4]3, Mo(CO)6, [Co(CO)4]2. Связь в этих соединениях аналогична связи в описанных выше цианокомплексах. Ni(CO)4 – летучее вещество, используется для отделения никеля от других металлов. Ухудшение структуры чугуна и стали в конструкциях часто связано с образованием карбонилов. Водород может входить в состав карбонилов, образуя карбонилгидриды, такие, как H2Fe(CO)4 и HCo(CO)4, проявляющие кислотные свойства и реагирующие со щелочью:

H2Fe(CO)4 + NaOH → NaHFe(CO)4 + H2O

Известны также карбонилгалогениды, например Fe(CO)X2, Fe(CO)2X2, Co(CO)I2, Pt(CO)Cl2, где Х – любой галоген (см. также МЕТАЛЛООРГАНИЧЕСКИЕ СОЕДИНЕНИЯ).

Углеводороды.

Известно огромное количество соединений углерода с водородом (см. ХИМИЯ ОРГАНИЧЕСКАЯ).

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