Nature of Chemical Bonds within the allotropes:

 

How does carbon form different structures?

The electronic configuration of the carbon atom allows it to form a number of hybridized atomic orbitals . Carbon atoms in the elemental substances (e.g., diamond, graphite, & Buckminsterfullerenes) bonds to each other covalently, by the sharing of electron pairs. The covalent bonds have directional properties. This in turn gives carbon the ability to adapt into various molecular and crystalline structures. The nature of these bonds underlie the varied chemical properties and physical properties of the carbon allotropes.

 

Carbon, like many of the first-row elements of the Periodic Table, has atomic orbitals that can hybridize. This is because the s-orbital and p-orbitals of carbon's second electronic shell have very similar energies. As a result, carbon can adapt to form chemical bonds with different geometries.

The shape of the carbon atomic orbitals are shown below in a stereo format.

 

Hybridization of Carbon Atomic Orbitals

Within diamond, one s-orbital and three p-orbitals undergo a SP³ hybridization.

The geometry of the hybridized orbital is tetrahedral. This is the reason why each carbon atom within diamond has four nearest neighbors.

(1) The SP3 hybridized orbital has a tetrahedral symmetry

 

 

 

 

Within graphite, one s-orbital and two p-orbitals undergo a SP2 hybridization.

The geometry of the hybridized orbital is trigonal planar. This is the reason why each carbon atom within graphite has three nearest neighbors within the graphite sheets. One of the p-orbitals is left unaffected. This last p-orbital overlap with those from neighboring carbon atoms, in a sideways manner, to form the distributed pi-bonds that reside above and below each graphite sheet.

(2) SP2 having a trigonal planar symmetry

These hybridized atomic orbitals form chemical bonds to form diamond, graphite, and buckminsterfullerenes.

A covalent chemical bond is formed between two atoms when their orbitals overlap and share a pair of electrons. When the orbitals overlap along an axis between the atoms (internuclear axis), they form a sigma bond. In this type of bonding the electron density is highest in the space between the atoms.

Overlap between two S orbitals to form a sigma bond (green).

Overlap between two P orbitals to form a sigma bond (green).

For P orbitals, sideways overlapping is also possible. This results in the formation of pi bonds. The regions of highest electron density for a pi bond occur in pairs, parallel to the internuclear axis, above and below or right and left of the connecting atoms.

A triple bond consists of a sigma bond (green) and two pi bonds (red).

Multiple covalent bonds between atoms contain a sigma bond and one or more pi bonds.

Type of bonding	No. of sigma bonds	No. of pi bonds
single	1	0
double	1	1
triple	1	2

A three-dimensional network of sigma bonds connects the carbon atoms in diamond. The carbon atoms in graphite are connected a hexagonal network of sigma bonds with electrons from pi bonds distributed above and below the plane of atoms.

How do the chemical bonds within graphite affect its properties?

The electrical conductivity and softness of graphite can be related to graphite's crystalline structure. Crystalline graphite consists of parallel sheets of carbon atoms, each sheet containing hexagonal arrays of carbon atoms. Each carbon atom exhibits sp² hybridization to form sigma bonds with three nearest neighbors, separated by 1.4210 Angstroms, in the layer. The bonds between the carbons within the layer are stronger than those in diamond. There is also distributed pi bonding between the carbon atoms in the sheet. This delocalized pi system is responsible for the electrical conductivity of graphite. The interaction of the loosely held electrons, in the pi bonds, with light is responsible for its black color. The softness and lubricating nature of graphite arises from the weak binding of the carbon sheets by weak Van der Waals forces.

 

How do the chemcial bonds within diamond affect its properties?

The hardness of diamond can also be attributed to its crystalline structure and the strength of the chemical bonds between its carbon atoms. Each carbon atom has four nearest neighbors to which it is bonded by sigma bonds, separated by a distance of 1.5445 Angstroms. The bond angles are all 109 degrees, typical of sp3 hybridization. The resultant interlocking network of covalent bonds makes the structure very rigid. Diamond, as it turns out, is the hardest material on earth. Since the valence electrons in diamond are involved in the formation of sigma bonds and no delocalized pi bonds, diamond has very poor electrical conductivity. The electrons within diamond are tightly held within the bonds between the carbon atoms. These electrons absorbs light in the ultraviolet region but not in the visible region, so pure diamond appears clear to the human eye.