Carbon
Carbon is a remarkable chemical element for many reasons. It is a vital component of all known forms of life—without it, our physical existence and that of other organisms would not be possible. Its different forms include graphite, one of the softest materials, and diamond, one of the hardest substances. Carbon is known to be part of a huge variety of compounds, including some that occur in the Sun, stars, and planetary atmospheres. We use carbon compounds for fuel and the synthesis of a wide range of materials such as plastics, paints, textiles, and pharmaceuticals. Researchers are now investigating the uses of new forms of carbon, known as fullerenes, leading to the developing field of nanotechnology.
On the other hand, our careless disposal of paints, plastics, and other carbon-containing materials has polluted the environment, and our misuse of drugs has caused needless suffering and pain. Such practices are symptomatic of our inadequate understanding of the value of the environment and our own value, and of the principles of the natural world. Efforts are now being made to restore the environment to its former, unpolluted state, but we need to restore our sense of values as well, so that such problems do not continue into the future. We need to use wisdom in terms of the materials we produce, how we use them, and how we dispose of them.
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| Name, Symbol, Number | carbon, C, 6 | ||||||||||||||||||||||||
| Chemical series | nonmetals | ||||||||||||||||||||||||
| Group, Period, Block | 14, 2, p | ||||||||||||||||||||||||
| Appearance | black (graphite) colorless (diamond) 
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| Atomic mass | 12.0107(8) g/mol | ||||||||||||||||||||||||
| Electron configuration | 1s2 2s2 2p2 | ||||||||||||||||||||||||
| Electrons per shell | 2, 4 | ||||||||||||||||||||||||
| Physical properties | |||||||||||||||||||||||||
| Phase | solid | ||||||||||||||||||||||||
| Density (near r.t.) | (graphite) 2.267 g/cm³ | ||||||||||||||||||||||||
| Density (near r.t.) | (diamond) 3.513 g/cm³ | ||||||||||||||||||||||||
| Melting point | ? triple point, ca. 10 MPa and (4300–4700) K (? °C, ? °F) | ||||||||||||||||||||||||
| Boiling point | subl. ? ca. 4000 K (? °C, ? °F) | ||||||||||||||||||||||||
| Heat of fusion | (graphite) ? 100 kJ/mol | ||||||||||||||||||||||||
| Heat of fusion | (diamond) ? 120 kJ/mol | ||||||||||||||||||||||||
| Heat of vaporization | ? 355.8 kJ/mol | ||||||||||||||||||||||||
| Heat capacity | (25 °C) (graphite) 8.517 J/(mol·K) | ||||||||||||||||||||||||
| Heat capacity | (25 °C) (diamond) 6.115 J/(mol·K) | ||||||||||||||||||||||||
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| Atomic properties | |||||||||||||||||||||||||
| Crystal structure | hexagonal | ||||||||||||||||||||||||
| Oxidation states | 4, 2 (mildly acidic oxide) | ||||||||||||||||||||||||
| Electronegativity | 2.55 (Pauling scale) | ||||||||||||||||||||||||
| Ionization energies (more) |
1st: 1086.5 kJ/mol | ||||||||||||||||||||||||
| 2nd: 2352.6 kJ/mol | |||||||||||||||||||||||||
| 3rd: 4620.5 kJ/mol | |||||||||||||||||||||||||
| Atomic radius | 70 pm | ||||||||||||||||||||||||
| Atomic radius (calc.) | 67 pm | ||||||||||||||||||||||||
| Covalent radius | 77 pm | ||||||||||||||||||||||||
| Van der Waals radius | 170 pm | ||||||||||||||||||||||||
| Miscellaneous | |||||||||||||||||||||||||
| Magnetic ordering | diamagnetic | ||||||||||||||||||||||||
| Thermal conductivity | (300 K) (graphite) (119–165) W/(m·K) | ||||||||||||||||||||||||
| Thermal conductivity | (300 K) (diamond) (900–2320) W/(m·K) | ||||||||||||||||||||||||
| Thermal diffusivity | (300 K) (diamond) (503–1300) mm²/s | ||||||||||||||||||||||||
| Mohs hardness | (graphite) 0.5 | ||||||||||||||||||||||||
| Mohs hardness | (diamond) 10.0 | ||||||||||||||||||||||||
| CAS registry number | 7440-44-0 | ||||||||||||||||||||||||
| Notable isotopes | |||||||||||||||||||||||||
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Occurrence
Carbon is an abundant nonmetal that has been known since prehistory. Early peoples produced it in the form of charcoal by burning organic material (such as wood) in insufficient oxygen.
The name carbon comes from the French word charbone, which in turn is derived from the Latin carbo, meaning charcoal. In German and Dutch, the names for carbon are Kohlenstoff and koolstof, respectively, both of which literally mean "coal-stuff."
Pure carbon can occur in a variety of forms known as allotropes, such as graphite and diamond, described below. In addition, carbon can bind to itself and various other elements to form compounds.
On planet Earth, carbon and its compounds are widely dispersed in the air, water, and land. They are also present in the Sun and stars, and it is thought that carbon atoms were first formed by nuclear reactions—particularly, the collision of three alpha particles (helium nuclei)—in the interior of stars. Moreover, carbon is found in comets and in the atmospheres of most planets in the solar system. Some meteorites contain microscopic diamonds that may have been formed when the solar system was in its formative stages.
In the United States, graphite is found in large quantities in New York and Texas. It is also abundant in Russia, Mexico, Greenland, and India.
Natural diamonds occur in the mineral 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. Other deposits have been found in Arkansas, Canada, the Russian Arctic, Brazil, and Northern and Western Australia.
Additional allotropic forms of carbon, classified as fullerenes, were discovered as byproducts of molecular beam experiments in the 1980s.
Physical and chemical characteristics
The chemical symbol for carbon is C, and its atomic number (the number of protons in the nucleus of each atom) is 6. Each carbon atom is capable of forming strong chemical bonds, called covalent bonds, with up to four other atoms. Carbon is therefore said to be tetravalent.
Carbon has the highest melting/sublimation point of all elements. If heated at atmospheric pressure, carbon does not melt but sublimates (the solid form vaporizes) above 4000 K. Thus it remains a solid at higher temperatures than metals with the highest melting points, including tungsten and rhenium.
Carbon has two stable, naturally occurring isotopes: carbon-12 (12C) and carbon-13 (13C), at relative abundances of approximately 98.9 percent and 1.1 percent, respectively. In 1961, the International Union of Pure and Applied Chemistry adopted the isotope carbon-12 as the basis for atomic weights.
One naturally occurring radioisotope, carbon-14 (14C), has a half-life of about 5730 years. It has been used extensively for radioactive dating of carbon-containing materials. Twelve other radioisotopes have been produced artificially. Of these, the shortest lived is 8C, which has a half-life of 1.98739x10-21 seconds.
Allotropes
Carbon can exist in different structural states known as allotropes. Two well-known allotropes are graphite, one of the softest substances, and diamond, the hardest naturally occurring mineral.
Graphite has a crystalline structure in which the carbon atoms are arranged in layers of flat sheets. Each sheet, called graphene, consists of a network of hexagonal (six-membered) rings in which each carbon atom is bonded to three other atoms. While each sheet is extremely stable, it interacts loosely with adjacent sheets, so that they can readily slide past one another.
Diamond, which is formed in the Earth at very high pressures, adopts a different crystalline structure. Each carbon atom is bonded to four others, which together form a tetrahedral shape surrounding the central carbon. The overall structure is a three-dimensional network of puckered six-membered rings of carbon atoms. In this manner, each diamond crystal regardless of size can be thought of as a single molecule of carbon atoms. At room temperature, a diamond crystal undergoes an extremely slow transition to graphite.
It is interesting to note that diamond and graphite have certain diametrically opposite features. Diamond is transparent and hard, acts as an excellent electrical insulator, and can be used as an abrasive. By contrast, graphite is opaque and soft, conducts electricity, and is a good lubricant. Diamond crystallizes in the cubic system but graphite crystallizes in the hexagonal system.
Carbon that lacks an overall crystalline structure is known as amorphous carbon. Although it is possible to produce carbon that is entirely amorphous (noncrystalline), material that is described as amorphous usually includes small crystals of graphite or diamond. Amorphous carbon is the main constituent of charcoal, soot (lampblack or carbon black), and activated carbon.
In addition, several exotic allotropes have been synthesized or discovered, including fullerenes, lonsdaleite, and aggregated diamond nanorods. The carbon atoms in these allotropes have different structural arrangements.
Fullerenes can take the form of a hollow sphere, ellipsoid, or tube. Spherical fullerenes are also called buckyballs, and cylindrical fullerenes are known as buckytubes or carbon nanotubes. The most well-known buckyball is Buckminsterfullerene (containing 60 carbon atoms per sphere), named after Richard Buckminster Fuller, the architect who developed the geodesic dome.
A fullerene is composed of a sheet of linked hexagonal rings (similar to a sheet in graphite), but the sheet also contains some pentagonal (and sometimes heptagonal) rings that allow the sheet to curve and form a sphere, ellipse, or tube.
Lonsdaleite is thought to form when meteoric graphite falls to Earth. The impact's heat and stress transform the graphite into a structure similar to diamond, but graphite's hexagonal crystal lattice is retained. Also known as hexagonal diamond, lonsdaleite is transparent and brownish yellow in color.
- Chaoite A mineral supposedly formed in meteorite impacts.
- Lonsdaleite (a corruption of diamond). Structure: similar to diamond, but forming a hexagonal crystal lattice.
{* Amorphous carbon (a glassy substance). Structure: an assortment of carbon molecules in a non-crystalline, irregular, glassy state.}
- Carbon nanofoam (an extremely light magnetic web). Structure: a low-density web of graphite-like clusters, in which the atoms are bonded trigonally in six- and seven-membered rings.
- Aggregated diamond nanorods, the most recently discovered allotrope and the hardest substance known to man.
'Glassy carbon' is isotropic and contains a high proportion of closed porosity. Unlike normal graphite, the graphitic layers are not stacked like pages in a book, but have a more random arrangement.
Carbon fibers are similar to glassy carbon. Under special treatment (stretching of organic fibers and carbonization) it is possible to arrange the carbon planes in direction of the fiber. Perpendicular to the fiber axis there is no orientation of the carbon planes. The result is fibers with a higher specific strength than steel.
A nanofoam allotrope has been discovered which is ferromagnetic.
Carbon allotropes include:
- Amorphous carbon
- Carbon nanofoam (discovered in 1997)
- Carbon nanotube
- Diamond
- Fullerene
- Graphite
- Lonsdaleite
- Aggregated diamond nanorods (synthesised in 2005)
At normal pressures, carbon takes the form of graphite, in which each atom is bonded to three others in a plane composed of fused hexagonal rings, just like those in aromatic hydrocarbons. The two known forms of graphite, alpha (hexagonal) and beta (rhombohedral), both have identical physical properties, except for their crystal structure. Graphites that naturally occur have been found to contain up to 30% of the beta form, when synthetically-produced graphite only contains the alpha form. The alpha form can be converted to the beta form through mechanical treatment and the beta form reverts back to the alpha form when it is heated above 1000 °C.
Because of the delocalization of the pi-cloud, graphite conducts electricity. The material is soft and the sheets, frequently separated by other atoms, are held together only by van der Waals forces, so easily slip past one another.
- Graphite is soft and is used in pencils
- Diamond is the hardest mineral known to man (although aggregated diamond nanorods are now believed to be even harder), but graphite is one of the softest.
Compounds of carbon
Each carbon atom can form covalent bonds with other carbon atoms and various other elements, of which the most common are hydrogen, oxygen, nitrogen, and chlorine. In so doing, it forms compounds that consist of long, open-ended chains and ring-shaped molecules.
Scientists are aware of nearly 10 million carbon compounds, constituting the vast majority of all known chemical compounds. Most compounds of carbon are classified as organic compounds, which form the basis for the field of organic chemistry.
When combined with only hydrogen atoms, carbon forms various hydrocarbons, including methane, propane, butane, benzene, and naphthalene. Natural sources of hydrocarbons are coal, petroleum, and natural gas.
When bound to oxygen alone, carbon forms carbon dioxide and carbon monoxide. Carbon dioxide in the atmosphere is vital for plants to perform photosynthesis, and, as a greenhouse gas, it helps retain heat from the Sun and keep our planet warm enough to sustain life.
When attached to both oxygen and hydrogen, carbon forms many groups of compounds, including those classified as alcohols, esters, aldehydes, and ketones.
Thousands of carbon compounds—including carbohydrates, fats, proteins, and nucleic acids—are produced in and used by living systems. Other compounds—particularly the carbonates of metals such as calcium, magnesium, and iron—are major components of rocks, including limestone, dolomite, and marble.
Applications
Carbon is a vital component of all known living systems, and without it life as we know it could not exist (see alternative biochemistry). The major economic use of carbon is in the form of hydrocarbons, most notably the fossil fuels methane gas and crude oil (petroleum). Crude oil is used by the petrochemical industry to produce, amongst others, gasoline and kerosene, through a distillation process in refineries. Crude oil forms the raw material for many synthetic substances, many of which are collectively called plastics.
Other uses
- The isotope Carbon-14 was discovered in February 27 1940 and is used in radiocarbon dating.
- Graphite is combined with clays to form the 'lead' used in pencils.
- Diamond is used for decorative purposes, and also as drill bits and other applications making use of its hardness. Diamonds have long been considered rare and beautiful.
- Carbon is added to iron to make steel.
- Carbon is used as a neutron moderator in nuclear reactors.
- Graphite carbon in a powdered, caked form is used as charcoal for cooking, artwork and other uses.
- Activated charcoal is used in medicine (as powder or compounded in tablets or capsules) to absorb toxins or poisons from the digestive system.
The chemical and structural properties of fullerenes, in the form of carbon nanotubes, has promising potential uses in the nascent field of nanotechnology. Nanoparticles might however be toxic.
Organic compounds
The most prominent oxide of carbon is carbon dioxide, CO2. This is a minor component of the Earth's atmosphere, produced and used by living things, and a common volatile elsewhere. In water it forms trace amounts of methanoic acid, HCO2H, but as most compounds with multiple single-bonded oxygens on a single carbon it is unstable. Through this intermediate, though, resonance-stabilized carbonate ions are produced. Some important minerals are carbonates, notably calcite. Carbon disulfide, CS2, is similar.
The other oxides are carbon monoxide, CO, and the uncommon carbon suboxide, C3O2. Carbon monoxide 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 haemoglobin molecules, so that the gas is highly poisonous. Cyanide, CN-, has a similar structure and behaves a lot like a halide ion; the nitride cyanogen, (CN)2, is related.
With reactive metals, such as tungsten, carbon forms either carbides, C-, or acetylides, C22- to form alloys with very high melting points. These anions are also associated with methane and acetylene, both very weak acids. All in all, with an electronegativity of 2.5, carbon prefers to form covalent bonds. A few carbides are covalent lattices, like carborundum, SiC, which resembles diamond.
Carbon chains
Carbon has the ability to form long chains with interconnecting C-C bonds. This property is called Catenation. Carbon-Carbon bonds are fairly strong, and abnormaly stable. This property is important as it allows carbon to form a huge number of compounds; if fact, there are more known carbon-containing compounds than all the other compounds of the chemical elements combined!
The simplest form of an organic molecule is the hydrocarbon - a large family of organic molecules that, by definition, are composed of hydrogen atoms bonded to a chain of carbon atoms. Chain length, side chains and functional groups all affect the properties of organic molecules.
Carbon cycle
Under terrestrial conditions, conversion of one isotope to another is very rare. Therefore, for practical purposes, the amount of carbon on Earth is constant. Thus, processes that use carbon must obtain it from somewhere, and dispose of it somewhere. The paths that carbon follows in the environment are collectively called the carbon cycle. For example, plants draw carbon dioxide out of the environments and use it to build biomass. Some of this biomass is eaten by animals, where some of it is exhaled as carbon dioxide. The carbon cycle is considerably more complicated than this short loop; for example, some carbon dioxide is dissolved in the oceans; dead plant or animal matter may become sedimentary rock, and so forth.
Isotopes
Carbon has two stable, naturally-occurring isotopes: carbon-12, or 12C, (98.89%) and carbon-13, or 13C, (1.11%), and one unstable, naturally-occurring, radioisotope; carbon-14 or 14C. There are 15 known isotopes of carbon and the shortest-lived of these is 8C which decays through proton emission and alpha decay. It has a half-life of 1.98739x10-21 s.
In 1961 the International Union of Pure and Applied Chemistry adopted the isotope carbon-12 as the basis for atomic weights.
Carbon-14 has a half-life of 5730 y and has been used extensively for radioactive dating of carbonaceous materials.
Precautions
Carbon is relatively safe. Inhalation of fine soot in large quantities can be dangerous. Carbon may catch fire at very high temperatures and burn vigorously (as in the Windscale fire).
There are a tremendous number of carbon compounds; some are lethally poisonous (cyanide, CN-), and some are essential to life (dextrose).
ReferencesISBN links support NWE through referral fees
- On Graphite Transformations at High Temperature and Pressure Induced by Absorption of the LHC Beam, J.M. Zazula, 1997
- WebElements.com and EnvironmentalChemistry.com per the guidelines at Wikipedia's WikiProject Elements
See also
- Organic chemistry
- Inorganic chemistry of carbon
- Allotropes of carbon
- Diamond
- Material properties of diamond
- Fullerene
- Carbon nanotube
External links
- Los Alamos National Laboratory – Carbon
- WebElements.com – Carbon
- It's Elemental – Carbon
- – Carbon Fullerene and other Allotropes models by Vincent Herr
- Extensive Carbon page at asu.edu
- Electrochemical uses of carbon
- Computational Chemistry Wiki
- Origin of Carbon
af:Koolstof ar:كربون bg:Въглерод bn:কার্বন ca:Carboni cs:Uhlík cy:Carbon da:Carbon de:Kohlenstoff et:Süsinik el:Άνθρακας es:Carbono eo:Karbono eu:Karbono fa:کربن fr:Carbone gd:Gualan gl:Carbono (elemento) gu:કાર્બન ko:탄소 hr:Ugljik io:Karbo id:Karbon ia:Carbon (elemento) is:Kolefni it:Carbonio he:פחמן la:Carbonium lv:Ogleklis lt:Anglis li:Koolstof hu:Szén mk:Јаглерод mi:Waro ms:Karbon nl:Koolstof nds:Kohlenstoff ja:炭素 no:Karbon (grunnstoff) nn:Karbon oc:Carbòni pl:Węgiel (pierwiastek) pt:Carbono ru:Углерод simple:Carbon sk:Uhlík sl:Ogljik sr:Угљеник su:Karbon fi:Hiili sv:Kol th:คาร์บอน tr:karbon vi:Cacbon uk:Вуглець wa:Carbone zh:碳
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