Nature of Chemical Bonds within the allotropes:
How does carbon form different
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
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
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
Overlap between two P orbitals to form a sigma
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
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.