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Carbon General

Name:Carbon Symbol:C
Type:Non-Metal, Carbon group Atomic weight:12.011
Density @ 293 K:2.267 g/cm3 (graphite), 3.513 g/cm3 (diamond) Atomic volume:5.31 cm3/mol (graphite), 3.42 cm3/mol (diamond)

Carbon has been known since ancient times in the form of soot, charcoal, graphite and diamonds. Ancient cultures did not of course realize that these substances were different forms of the same element. 'Carbon' is derived from the Latin carbo, meaning charcoal.

Antoine Lavoisier named carbon and he carried out early experiments to reveal its nature. In 1694 he pooled resources with other chemists to buy a diamond, which they placed in a closed glass jar. They focused the sun's rays on the diamond with a magnifying glass and the diamond burnt and disappeared. Lavoisier noted that the overall weight of the jar was unchanged. He concluded the diamond was composed of the same element as charcoal was - carbon. When it burnt, the diamond had combined with oxygen to form carbon dioxide. (1), (2)

In 1779 Carl Scheele showed that graphite burnt to form carbon dioxide and so must be a form of carbon.(3)

In 1796 Smithson Tennant established that diamond was pure carbon and not a compound of carbon and it burnt to form only carbon dioxide. Tennant also proved that when equal weights of charcoal and diamonds were burnt, they produced the same amount of carbon dioxide. (4)

In 1855 Benjamin Brodie produced pure graphite from carbon in his laboratory, proving it was a form of carbon.(4)

Although it had been previously attempted, in 1955 Francis Bundy and coworkers at General Electric finally demonstrated that graphite could be transformed to diamond at high temperature and high pressure.(5)

In 1985 Robert Curl, Harry Kroto and Richard Smalley discovered fullerenes, a new form of carbon in which the atoms are arranged in soccer-ball shapes.(6), (7)

Carbon States

State (s, l, g):solid
Melting point:3823 K (3550 °C) Boiling point:4300 K (4027 °C)

Carbon Energies

Specific heat capacity: 0.71 J g-1 K-1 (graphite), 0.5091 J g-1 K-1 (diamond) Heat of atomization:717 kJ mol-1
Heat of fusion:117 kJ mol-1 (graphite) Heat of vaporization :710.9 kJ mol-1
1st ionization energy:1086.5 kJ mol-1 2nd ionization energy:2352.6 kJ mol-1
3rd ionization energy:4620.5 kJ mol-1 Electron affinity:121.55 kJ mol-1

Carbon Oxidation & Electrons

Shells:2,4 Electron configuration: [He] 2s2 2p2
Minimum oxidation number: -4 Maximum oxidation number:4
Min. common oxidation no.: -4 Max. common oxidation no.:4
Electronegativity (Pauling Scale):2.55 Polarizability volume:1.8 Å3

Carbon Appearance & Characteristics

Structure:hexagonal layers (graphite), tetrahedral (diamond) Color:black (graphite), transparent (diamond)
Hardness:0.5 mohs (graphite), 10.0 mohs (diamond)
Harmful effects:

Pure carbon has very low toxicity. Inhalation of large quantities of carbon black dust (soot/coal dust) can cause irritation and damage to the lungs.


Carbon can exist in several allotropes, including graphite, diamond, amorphous carbon, fullerines and nanotubes. (The structures of eight allotropes are shown at the bottom of this page.)

Interestingly, graphite is one of the softest substances and diamond was thought, until recently, to be the hardest naturally occurring substance.

An extremely rare allotrope of carbon, Lonsdaleite, has been calculated, in pure form, to be 58% stronger than diamond. Lonsdaleite is made when meteorites containing graphite hit another body, such as Earth. The high temperatures and pressures of the impact transform the graphite into Lonsdaleite, a diamond-like substance that retains graphite's hexagonal structure.

Carbon has the highest melting/sublimation point of all the elements and, in the form of diamond, has the highest thermal conductivity of any element. This is the origin of the slang term "ice" - diamond, at room temperature, carries heat away from your warmer skin faster than any other material and so feels cold to touch.


Carbon (coal) is used as a fuel. 

Graphite is used as a lubricant, for pencil tips, high temperature crucibles, dry cells and electrodes.

Diamonds are used in jewelry and - because they are so hard - in industry for cutting, drilling, grinding, and polishing.

Carbon black is used as the black pigment in printing ink.

Carbon can form alloys with iron, of which the most common is carbon steel. The 14C  radioactive isotope is used in archaeological dating. Carbon compounds are important in many areas of the chemical industry.

Carbon forms a vast number of compounds with hydrogen, oxygen, nitrogen and other elements. Its ability to form long-chained, complex compounds has resulted in carbon acting as the basis of all life on Earth.

The outstanding physical properties - for example thermal conductivity and strength - of new carbon allotropes, such as nanotubes, show enormous potential for future development.

Carbon Reactions

Reaction with air:vigorous, ⇒ CO2 Reaction with 6 M HCl:none
Reaction with 15 M HNO3:mild, w/ht ⇒ C6(CO2H)6 (mellitic/graphitic acid) Reaction with 6 M NaOH:none

Carbon Compounds

Oxide(s):CO , CO2 Chloride(s):CCl4
Hydride(s):CH4 and many CxHy

Carbon Radius

Atomic radius:70 pm Ionic radius (1+ ion):pm
Ionic radius (2+ ion):pm Ionic radius (3+ ion):pm
Ionic radius (2- ion):pm Ionic radius (1- ion):pm

Carbon Conductivity

Thermal conductivity:25-470 W m-1 K-1 (graphite) 470 W m-1 K-1 (diamond) Electrical conductivity:0.07 x 106 S cm-1

Carbon Abundance & Isotopes

Abundance earth's crust: 200 parts per million by weight, 344 parts per million by moles
Abundance solar system: 3,000 parts per million by weight, 300 parts per million by moles
Cost, pure: $2.4 per 100g
Cost, bulk:$ per 100g

Carbon can be obtained by burning organic compounds with insufficient oxygen. The four main allotropes of carbon are graphite, diamond, amorphous carbon and fullerines. Natural diamonds are found in kimberlite from ancient volcanoes. Graphite can also be found in natural deposits. Fullerenes were discovered as byproducts of molecular beam experiments in the 1980's. Amorphous carbon is the main constituent


13 whose half-lives are known, with mass numbers 8 to 20. Of these, two are stable, 12C and 13C. Isotope 14C, with a half-life of 5730 years, is widely used to date carbonaceous materials such as wood, archeological specimens, etc for ages up to about 40,000 years.

Carbon Other


1. Robert E. Krebs, The history and use of our earth's chemical elements: a reference guide., (2006) p192. Greenwood Publishing Group
2. John Emsley, Nature's building blocks: an A-Z guide to the elements., (2002) p95. Oxford University Press.
3. Jessica Elzea Kogel, Industrial minerals & rocks: commodities, markets, and uses., (2006) p507. SME.
4. Amanda S. Barnard, The diamond formula: diamond synthesis--a gemmological perspective., (2000) p3. Butterworth-Heinemann
5. Robert M. Hazen, The diamond makers., (1999) p145. Cambridge University Press.
6. Jonathan W. Steed, Jerry L. Atwood, Supramolecular Chemistry., (2009) p423. Wiley.
7. Nobel Prize for Chemistry, 1996

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