Carbon: Properties, Structure, Obtaining, Uses

The carbon is a non – metallic chemical element whose chemical symbol is C. named after coal, vegetable or mineral, where its atoms define various structures. Many authors qualify it as the King of the elements, as it forms a wide range of organic and inorganic compounds, and also occurs in a considerable number of allotropes.

And if this is not enough to refer to it as a special element, it is found in all living beings; all its biomolecules owe their existence to the stability and strength of the CC bonds and their high tendency to concatenate. Carbon is the element of life, and with its atoms their bodies are built.

The wood of trees is composed mainly of carbohydrates, one of the many compounds rich in carbon. Source: Pexels.

The organic compounds with which biomaterials are built consist practically of carbon skeletons and heteroatoms. These can be seen with the naked eye in the wood of the trees; and also, when lightning strikes them and roasts them. The remaining inert black solid also has carbon; but it is charcoal.

Thus, there are “dead” manifestations of this element: charcoal, a product of combustion in oxygen-poor environments; and mineral coal, a product of geological processes. Both solids look alike, they are black, and they burn to generate heat and energy; although with different yields.

From this point on, carbon is the 15th most abundant element in the earth’s crust. No wonder when millions of tons of coal are produced annually. These minerals differ in their properties depending on the degree of impurities, placing anthracite as the highest quality mineral coal.

The earth’s crust is not only rich in mineral coal, but also in carbonates, especially limestone and dolomites. 
And regarding the Universe, it is the fourth most abundant element; I mean, there is more carbon out there on other planets.

Article index

  • one

    Carbon history

    • 1.1


    • 1.2


  • two

    Properties (edit)

    • 2.1

      Graphite vs diamond

  • 3

    Electronic structure and configuration

    • 3.1


    • 3.2

      Oxidation numbers

    • 3.3

      Molecular geometries

    • 3.4

      Amorphous or crystalline solids

  • 4


  • 5


  • 6

    Risks and precautions

  • 7


Carbon history


Carbon may be as old as the earth’s crust itself. Since time immemorial, ancient civilizations have encountered this element in its many natural presentations: soot, charcoal, charcoal, charcoal, diamonds, graphite, coal tar, anthracite, etc.

All those solids, although they shared the dark tones (with the exception of diamond), the rest of their physical properties, as well as their composition, differed remarkably. Back then it was impossible to claim that they essentially consisted of carbon atoms.

It was thus that throughout history, coal was classified according to its quality at the time of burning and providing heat. And with the gases formed by its combustion, masses of water were heated, which in turn produced vapors that moved turbines that generated electrical currents.

Carbon was unexpectedly present in charcoal produced by burning trees in closed or hermetic spaces; in the graphite with which the pencils were made; in diamonds used as gems; he was responsible for the hardness of the steel.

Its history goes hand in hand with wood, gunpowder, city lighting gases, trains and ships, beer, lubricants and other essential objects for the advancement of humanity.


At what point were scientists able to associate the allotropes and minerals of carbon to the same element? Coal was seen as a mineral, and it was not thought of as a chemical element worthy of the periodic table. The first step should have been to show that all these solids were transformed into the same gas: carbon dioxide, CO 2 .

Antoine Lavoisier in 1772, using a wooden frame with large lenses, focused the sun’s rays on samples of charcoal and a diamond. He discovered that neither of them formed water vapors but CO 2 . He did the same with the soot and got the same results.

Carl Wilhelm Scheele in 1779, found the chemical relationship between charcoal and graphite; that is, both solids were composed of the same atoms.

Smithson Tennant and William Hyde Wollaston in 1797 methodologically verified (through reactions) that the diamond was actually composed of carbon when producing CO 2 in its combustion.

With these results, light was soon thrown on graphite and diamond, solids formed by carbon, and therefore, of high purity; unlike the impure solids of coal and other carbonaceous minerals.

Properties (edit)

The physical or chemical properties found in carbonaceous solids, minerals, or materials are subject to many variables. Among them are: the composition or degree of impurities, the hybridizations of the carbon atoms, the diversity of the structures, and the morphology or size of the pores.

When describing the properties of carbon, most texts or bibliographic sources are based on graphite and diamond.

Why? Because they are the best known allotropes for this element and represent high purity solids or materials; that is, they are practically made of nothing more than carbon atoms (although with different structures, as will be explained in the next section).

The properties of charcoal and mineral coal differ in their origins or compositions, respectively. For example, lignite (low carbon) as a fuel crawls compared to anthracite (high carbon). And what about the other allotropes: nanotubes, fullerenes, graphenes, grafins, etc.

However, chemically they have one point in common: they oxidize with an excess of oxygen in CO 2 :

C     + O => CO 2

Now, the speed or temperature they require to oxidize are specific to each of these allotropes.

Graphite vs diamond

A brief comment will also be made here regarding the very different properties for these two allotropes:

Table in which some properties of the two crystalline allotropes of carbon are compared. Source: Gabriel Bolívar.

Electronic structure and configuration


Relationship between hybrid orbitals and possible structures for carbon. Source: Gabriel Bolívar.

The electron configuration for the carbon atom is 1s 2 2s 2 2p 2 , also written as [He] 2s 2 2p 2 (top image). This representation corresponds to its ground state: the carbon atom isolated and suspended in such a vacuum that it cannot interact with others.

It can be seen that one of its 2p orbitals lacks electrons, which accepts an electron from the lower energy 2s orbital through electronic promotion; and thus, the atom acquires the ability to form up to four covalent bonds through its four sp 3 hybrid orbitals .

Note that all four sp 3 orbitals are energy degenerate (aligned on the same level). Pure p orbitals are more energetic, which is why they are placed above the other hybrid orbitals (to the right of the image).

If there are three hybrid orbitals, it is because an unhybridized p orbital remains ; therefore, they are three sp 2 orbitals . And when there are two of these hybrid orbitals, two p orbitals are available to form double or triple bonds, being the hybridization of the sp carbon.

Such electronic aspects are essential to understand why carbon can be found in infinities of allotropes.

Oxidation numbers

Before continuing with the structures, it is worth mentioning that, given the valence electron configuration 2s 2 2p 2 , carbon can have the following oxidation numbers: +4, +2, 0, -2 and -4.

Why? These numbers correspond to the assumption that an ionic bond exists such that you form the ions with the respective charges; that is, C 4+ , C 2+ , C 0 (neutral), C 2- and C 4- .

For carbon to have a positive oxidation number, it must lose electrons; And to do so, it necessarily has to be bonded to very electronegative atoms (like oxygen).

Meanwhile, for carbon to have a negative oxidation number, it must gain electrons by bonding to metallic atoms or less electronegative than it (such as hydrogen).

The first oxidation number, +4, means that the carbon has lost all its valence electrons; the 2s and 2p orbitals remain empty. If the 2p orbital loses its two electrons, the carbon will have an oxidation number of +2; if you gain two electrons, you will have -2; and if you gain two more electrons by completing your valence octet, -4.


For example, for CO 2 the oxidation number of carbon is +4 (because oxygen is more electronegative); while for CH 4 , it is -4 (because hydrogen is less electronegative).

For CH 3 OH, the oxidation number of carbon is -2 (+1 for H and -2 for O); while for HCOOH, it is +2 (check that the sum gives 0).

Other oxidation states, such as -3 and +3, are also likely, especially when it comes to organic molecules; for example, in the methyl groups, -CH 3 .

Molecular geometries

The upper image not only showed the hybridization of the orbitals for the carbon atom, but also the resulting molecular geometries when several atoms (black spheres) were linked to a central one. This central atom to have a specific geometric environment in space, must have the respective chemical hybridization that allows it.

For example, for the tetrahedron the central carbon has sp 3 hybridization ; because such is the most stable arrangement for the four sp 3 hybrid orbitals . In the case of sp 2 carbons , they can form double bonds and have a flat trigonal environment; and so these triangles define a perfect hexagon. And for an sp hybridization, the carbons adopt a linear geometry.

Thus, the geometries observed in the structures of all allotropes are simply governed by tetrahedra (sp 3 ), hexagons or pentagons (sp 2 ), and lines (sp).

Tetrahedra define a 3D structure, while hexagons, pentagons and lines, 3D or 2D structures; The latter come to be the planes or sheets similar to the walls of the honeycombs:

Wall with hexagonal designs of a honeycomb in analogy to planes composed of sp2 carbons. Source: Pixabay.

And if we fold this hexagonal wall (pentagonal or mixed), we will obtain a tube (nanotubes) or a ball (fullerenes), or another figure. The interactions between these figures give rise to different morphologies.

Amorphous or crystalline solids

Leaving aside the geometries, hybridizations, or morphologies of the possible structures of carbon, its solids can be globally classified into two types: amorphous or crystalline. And between these two classifications their allotropes are distributed.

Amorphous carbon is simply one that presents an arbitrary mixture of tetrahedra, hexagons or lines, incapable of establishing a structural pattern; such is the case of coal, charcoal or activated charcoal, coke, soot, etc.

While the crystalline carbon consists of structural patterns formed by any of the proposed geometries; for example, diamond (three-dimensional network of tetrahedra) and graphite (stacked hexagonal sheets).


Carbon can be pure as graphite or diamond. These are found in their respective mineralogical deposits, scattered throughout the globe and in different countries. That is why some nations are more exporters of one of these minerals than others. In short, “you have to dig the earth” to get the carbon.

The same applies to mineral coal and its types. But this is not the case with charcoal, since a body rich in carbon must first “perish”, either under fire, or an electric lightning; of course, in the absence of oxygen, otherwise CO 2 would be released .

An entire forest is a carbon source like charcoal; not only for its trees, but also for its fauna.

In general, samples containing carbon must undergo pyrolysis (burning in the absence of oxygen) to release some of the impurities as gases; and thus, a solid rich in carbon (amorphous or crystalline) remains as a residue.


Again, like the properties and structure, the uses or applications are consistent with the allotropes or mineralogical forms of carbon. However, there are certain generalities that can be mentioned, in addition to some well-known points. Such are:

-Carbon has been used for a long time as a mineral reducing agent in obtaining pure metals; for example, iron, silicon and phosphorus, among others.

-It is the cornerstone of life, and organic chemistry and biochemistry are the studies of this reflection.

-It has also been a fossil fuel that allowed the first machines to start their gears. Similarly, carbon gas was obtained from it for the old lighting systems. Coal was synonymous with light, heat and energy.

-Mixed as an additive with iron in different proportions allowed the invention and improvement of steels.

-Its black color took place in art, especially graphite and all the writings made with its strokes.

Risks and precautions

Carbon and its solids do not pose any health risk. Who has cared about a bag of coal? They are sold in droves within the aisles of some markets, and as long as there is no fire nearby, their black blocks will not burn.

Coke, on the other hand, can pose a risk if its sulfur content is high. When it burns, it will release sulfur gases that, in addition to being toxic, contribute to acid rain. And although CO 2 in small amounts cannot suffocate us, it does have a huge impact on the environment as a greenhouse gas.

From this perspective, carbon is a “long-term” danger, since its combustion alters the climate of our planet.

And in a more physical sense, solids or carbonaceous materials if they are pulverized are easily transported by air currents; and consequently, they are introduced directly to the lungs, which can irreparably damage them.

For the rest, it is very common to consume “charcoal” when some food is cooked.


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