carbon - Applications, History and Etymology, Allotropes, Organic compounds, Precautions

C, element 6. It has two main forms: graphite (the stable form, very soft and black with a density of c.2 g/cm³) and diamond (the hardest substance known, density 3·5 g/cm³). Various molecular forms called fullerenes were discovered in the 1980s. Both graphite and diamond melt above 3500°C. Coal is mainly graphite, but also contains amorphous (non-crystalline) carbon, occurring as well as ‘carbon black’ (soot). Graphite can be grown into carbon fibres, used for strengthening plastics. Carbon also occurs naturally in compounds, particularly as carbonates, and in the atmosphere as carbon dioxide. In virtually all its compounds, it is covalently bonded, and shows a valency of 4. Carbon compounds are the basis of all living matter, and form the subject matter of organic chemistry. There are two simple oxides, carbon monoxide (CO), a very poisonous gas, boiling point ?191ºC, formed from the incomplete combustion of carbon and hydrocarbons, and carbon dioxide O=C=O, the product of complete combustion. Carbon dioxide is in turn the raw material for photosynthesis, regenerating combustibles. Solid carbon dioxide, or dry ice, sublimes at ?78°C without passing through a liquid phase. It accounts for less than 0·03% of the gases of the atmosphere.

6 boron ← carbon → nitrogen - ↑ C ↓ Si Periodic Table - Extended Periodic Table
General
Name, Symbol, Number	carbon, C, 6
Chemical series	nonmetals
Group, Period, Block	14, 2, p
Appearance	black (graphite) colorless (diamond)
Atomic mass	12.0107(8) g/mol
Electron configuration	1s 2p2
Electrons per shell	2, 4
Physical properties
Phase	solid
Density (near r.t.)	(graphite) 2.267 g·cm−3
Density (near r.t.)	(diamond) 3.513 g·cm−3
Melting point	? 10 MPa and (4300–4700) K (4027–4427 °C, 7280–8000 °F)
Boiling point	? 4000 K (3727 °C, 6740 °F)
Heat of fusion	(graphite) ? 100 kJ·mol−1
Heat of fusion	(diamond) ? 120 kJ·mol−1
Heat of vaporization	? 355.8 kJ·mol−1
Heat capacity	(25 °C) (graphite) 8.517 J·mol
Heat capacity	(25 °C) (diamond) 6.115 J·mol
Vapor pressure (graphite) P/Pa 1 10 100 1 k 10 k 100 k at T/K   2839 3048 3289 3572 3908
Atomic properties
Crystal structure	hexagonal
Oxidation states	4, 2 (mildly acidic oxide)
Electronegativity	2.55 (Pauling scale)
Ionization energies (more)	1st: 1086.5 kJ·mol−1
	2nd: 2352.6 kJ·mol−1
	3rd: 4620.5 kJ·mol−1
Atomic radius	70 pm
Atomic radius (calc.)	67 pm
Covalent radius	77 pm
Van der Waals radius	170 pm
Miscellaneous
Magnetic ordering	diamagnetic
Thermal conductivity	(300 K) (graphite) (119–165) W·m
Thermal conductivity	(300 K) (diamond) (900–2320) W·m
Thermal diffusivity	(300 K) (diamond) (503–1300) mm²/s
Mohs hardness	(graphite) 1-2
Mohs hardness	(diamond) 10.0
CAS registry number	7440-44-0
Selected isotopes
Main article: Isotopes of carbon iso NA half-life DM DE (MeV) DP 12C 98.9% C is stable with 6 neutrons 13C 1.1% C is stable with 7 neutrons 14C trace 5730 y beta- 0.156 14N
References

Carbon (IPA: /ˈkɑːbən/) is a chemical element in the periodic table that has the symbol C and atomic number 6. Moreover, it has a great affinity for bonding with other small atoms, including other carbon atoms, and its small size makes it capable of forming multiple bonds.

Applications

Carbon is a very important component of all known living systems, and without it life as we know it could not exist (see alternative biochemistry).

History and Etymology

Carbon was discovered in prehistory and was known to the ancients, who manufactured it by burning organic material in insufficient oxygen (making charcoal).

Allotropes

The allotropes of carbon are the different molecular configurations that pure carbon can take.

The three relatively well-known allotropes of carbon are amorphous carbon, graphite, and diamond.

In its amorphous form, carbon is essentially graphite but not held in a crystalline macrostructure.

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.

At very high pressures carbon forms an allotrope called diamond, in which each atom is bonded to four others.

Fullerenes have a graphite-like structure, but instead of purely hexagonal packing, also contain pentagons (or possibly heptagons) of carbon atoms, which bend the sheet into spheres, ellipses or cylinders. Fullerenes: Structure: comparatively large molecules formed completely of carbon bonded trigonally, forming spheroids (of which the best-known and simplest is the buckminsterfullerene or buckyball, because of its soccerball-shaped structure).

Carbon fibers are similar to glassy carbon.

Between amorphous carbon and nanotubes:

Amorphous carbon is among the easiest materials to synthesize, but carbon nanotubes are extremely expensive to make. Amorphous carbon is completely isotropic, but carbon nanotubes are among the most anisotropic materials ever produced.

Organic compounds

The most prominent oxide of carbon is carbon dioxide, CO2.

The other oxides are carbon monoxide, CO, the uncommon carbon suboxide, C3O2, the uncommon dicarbon monoxide, C2O and even carbon trioxide, CO3.

With reactive metals, such as tungsten, carbon forms either carbides, C to form alloys with very high melting points.

Carbon chains

Carbon has the ability to form long chains with interconnecting C-C bonds.

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. The paths that carbon follows in the environment are called the carbon cycle. For example, plants draw carbon dioxide out of the environments and use it to build biomass as in carbon respiration.

The exotic 19C exhibits a Nuclear halo

Precautions

Carbon is relatively safe.