Hybridization

Atoms can form bonds with each other by sharing unpaired electrons such that
each bond contains two electrons. In Topic A1, we identified that a carbon atom
has two unpaired electrons and so we would expect carbon to form two bonds.
However, carbon forms four bonds! How does a carbon atom form four bonds
with only two unpaired electrons?
So far, we have described the electronic configuration of an isolated carbon
atom. However, when a carbon atom forms bonds and is part of a molecular structure,
it can emixf the s and p orbitals of its second shell (the valence shell). This is
known as hybridization and it allows carbon to form the four bonds which we
observe in reality.
There are three ways in which this mixing process can take place.
œ the 2s orbital ismixedwith all three 2p orbitals. This isknownas sp3 hybridization;
œ the 2s orbital is mixed with two of the 2p orbitals. This is known as sp2
hybridization;
œ the 2s orbital is mixed with one of the 2p orbitals. This is known as sp
hybridization.

In sp3 hybridization, the s and the p orbitals of the second shell are ‘mixed’
to form four hybridized sp3 orbitals of equal energy.
Each hybridized orbital contains a single unpaired electron and so four
bonds are possible.
Each sp3 orbital is shaped like a deformed dumbbell with one lobe much
larger than the other. The hybridized orbitals arrange themselves as far
apart from each other as possible such that the major lobes point to the corners
of a tetrahedron. sp3 Hybridization explains the tetrahedral carbon in
saturated hydrocarbon structures.
Sigma (σ) bonds are strong bonds formed between two sp3 hybridized carbons
or between an sp3 hybridized carbon and a hydrogen atom. A σ bond
formed between two sp3 hybridized carbon atoms involves the overlap of
half filled sp3 hybridized orbitals from each carbon atom. A σ bond formed
between an sp3 hybridized carbon and a hydrogen atom involves a halffilled
sp3 orbital from carbon and a half-filled 1s orbital from hydrogen.
Nitrogen, oxygen, and chlorine atoms can also be sp3 hybridized in organic
molecules. This means that nitrogen has three half-filled sp3 orbitals and can
form three bonds which are pyramidal in shape. Oxygen has two half-filled
sp3 orbitals and can form two bonds which are angled with respect to each
other. Chlorine has a single half-filled sp3 orbital and can only form a single
bond. All the bonds which are formed are σ bonds.
adapted from G. L. Patrick
Department of Chemistry and Chemical Engineering,
Paisley University, Paisley, Scotland

ATOMIC STRUCTURE OF CARBON

Atomic orbitals

The atomic orbitals available for the six electrons of carbon are the s orbital
in the first shell, the s orbital in the second shell and the three p orbitals in
the second shell. The 1s and 2s orbitals are spherical in shape. The 2p
orbitals are dumbbell in shape and can be assigned 2px, 2py or 2pz depending
on the axis along which they are aligned.

Energy levels

The 1s orbital has a lower energy than the 2s orbital which has a lower
energy than the 2p orbitals. The 2p orbitals have equal energy (i.e. they
are degenerate).

Electronic configuration

Carbon is in the second row of the periodic table and has six electrons which
will fill up lower energy atomic orbitals before entering higher energy
orbitals (aufbau principle). Each orbital is allowed a maximum of two electrons
of opposite spin (Pauli exclusion principle). When orbitals of equal
energy are available, electrons will occupy separate orbitals before pairing
up (Hund’s rule). Thus, the electronic configuration of a carbon atom is 1s2
2s2 2px
1 2py
1.

Covalent bonding

A covalent bond binds two atoms together in a molecular structure and is formed
when atomic orbitals overlap to produce a molecular orbital – so called because
the orbital belongs to the molecule as a whole rather than to one specific atom. A
simple example is the formation of a hydrogen molecule (H2) from two hydrogen
atoms. Each hydrogen atom has a half-filled 1s atomic orbital and when the atoms
approach each other, the atomic orbitals interact to produce two MOs (the number
of resulting MOs must equal the number of original atomic orbitals, Fig. 1).

The MOs are of different energies. One is more stable than the original atomic
orbitals and is called the bonding MO. The other is less stable and is called the
antibonding MO. The bonding MO is shaped like a rugby ball and results from
the combination of the 1s atomic orbitals. Since this is the more stable MO, the
valence electrons (one from each hydrogen) enter this orbital and pair up. The
antibonding MO is of higher energy and consists of two deformed spheres. This
remains empty. Since the electrons end up in a bonding MO which is more stable
than the original atomic orbitals, energy is released and bond formation is
favored. In the subsequent discussions, we shall concentrate solely on the bonding
MOs to describe bonding and molecular shape, but it is important to realize
that antibonding molecular orbitals also exist.

adapted from G. L. Patrick
Department of Chemistry and Chemical Engineering,
Paisley University, Paisley, Scotland