Crystal Binding (Bonding) Continued
More on Covalent Bonding
• Consider 2 close Cl atoms. Each has electronic shells
=
2
1s
2
2s
2
6
p
2
3s
5
3p
• If they move close until their outer orbitals
overlap, the atoms can share 2 e- & "fill" the
remaining 3p shell of each Cl.
• The electronic energy there is lowered, which
causes the orbitals to stay overlapped; resulting
in a strong bond in Cl2 .
This is the covalent or
shared electron bond.
Covalent Chemical Bonds
Hybrid Orbitals - Carbon
Atom:  |  |   Diamond: |  |   
1s
2s
Diamond
C-C-C angle
= 109o 28’
2p
1s
2(sp3)
• The 2(sp3) orbital is tetrahedrally shaped.
Larger Overlap  Stronger Bond
Covalent Bonds are Directional
• Each C is tetrahedrally coordinated with 4
others (& each of them with 4 others...)
• The C-C-C bond angle is fixed at 109o 28'
(max. overlap)
• Note the Face-Centered Cubic lattice
The directional character of the bonds,
 lower coordination & symmetry, density
Hybrid Orbitals - Carbon
Alternatively:
Atom:  |  |     Graphite:  | 
1s 2s
2p
1s
|
2(sp2) 2p
• As many know, C is “flexible” in the sense
that it can participate in many different kinds
of bonding.
• In fact, many atoms in the center of the
Periodic Table with partially filled valence
shells are variable in how they bond (this
includes Si)
Covalent Chemical Bonds
Graphite Structure
• The 3 2(sp2) orbitals are coplanar & 120o apart
• The orbital overlap is similar to that in diamond
within the planes (strong too!).
Belongs to the Hexagonal Crystal Class
• Note the p-bonding between the remaining 2p's
• This results in delocalized e- 's in 2p orbitals
which results in electrical conductivity only
within sheets.
• There are other hybrids as well (dsp2 in CuOplanar X) e- may resonate in bonds of nonidentical atoms & give a partial ionic character if
one much more e-neg than other
Covalent-Network and
Molecular Solids
• Diamonds are an example of a covalent-network
solid in which atoms are covalently bonded to
each other.
– They tend to be hard and have high melting points.
Covalent-Network and
Molecular Solids
• Graphite is an example of a molecular solid in
which atoms are held together with van der Waals
forces.
– They tend to be softer and have lower melting points.
• Covalent Bonds occur between atoms that are
“sharing” electrons:
• Form covalent compounds. There is a “tug of war” for the
electrons. There can be single, double & triple covalent bonds:
• Single bond – a bond in which 2 atoms share a pair
of electrons. Double bond – bond that involves 2
shared pairs of e-. Triple bond – bond that involves 3
shared pairs of e• Combinations of atoms of non-metallic atoms are
likely to form covalent bonds
• Groups 4A, 5A, 6A, and 7A
• Summarized by G. Lewis in the octet rule sharing of eoccurs if atoms achieve noble gas configuration,
• H2 is an exception to this rule
Column (Group) Trends
• Halogens form single covalent bonds in
their diatomic molecules (ex: F – F)
• Chalcogens form double covalent bonds in
their diatomic molecules (ex: O = O)
• Phicogens form triple covalent bonds in
their diatomic molecules ( N = N )\
• The Carbon group tends to form 4 bonds
with other atoms
• As we just briefly saw for C, covalent bonding can
be explained using electron configurations and
orbital boxes
• Double and triple covalent bonds
– Oxygen forms a double bond in a diatomic molecule
• It is an exception to the octet rule, 2 unpaired e-.
– Nitrogen forms a triple bond in a diatomic molecule
• Satisfies the octet rule, all e- are paired
• Multiple covalent bonds can form between unlike
atoms (ex: CO2, CH3OH)
Molecular Orbitals
• As we briefly showed, when 2 atoms covalently
bond, their atomic orbitals overlap to produce
molecular orbitals (orbitals that apply to the entire molecule)
• The molecular orbital model of bonding requires that
the number of molecular orbitals equal the number of
overlapping atomic orbitals
• When 2 atomic orbitals overlap, 2 molecular orbitals
are created
• One is called a bonding orbital, the other is called an
anti-bonding orbital
• The anti-bonding orbital has a higher energy that the
atomic orbitals from which it formed
Molecular Orbitals
• When H2 forms, the 1s atomic orbitals overlap
• 2 electrons are available for bonding (see next slide)
• The energy of the e- in the bonding molecular orbital
is lower than the e- in the atomic orbitals of the
separate H atoms
• Electrons seek the lowest energy level, so they fill
the bonding molecular orbital
• This makes a stable covalent bond between the H
atoms
• The anti-bonding orbital is empty
• Sigma and pi bonds are caused by the overlapping of
“s” and “p” orbitals
Covalent Bonding of 2 H Atoms
 H2 Molecule
Interaction
Potential
Hybrid Orbitals
• In orbital hybridization, several atomic orbitals
mix to form the same total number of
equivalent hybrid orbitals
• One 2s and three 2p orbitals of a carbon atom
overlap to form an sp3 hybrid orbital
• These are at the tetrahedral angle of 109.5o
• Four sp3 orbitals of carbon overlap with the 1s
orbitals of the four hydrogen atoms
• This allows for a great deal of overlap, which
results in the formation of 4 C-H sigma bonds
• These are unusually strong covalent bonds
Bond Polarity
• Covalent bonds involve the sharing of electrons
• However, they can differ in how the bonds are shared
• Depends on the kind and number of atoms joined together
• When electrons are shared equally, a nonpolar
covalent bond is formed
• When the atoms share the electron unequally, a polar
covalent bond is formed
• The more electronegative element will have the
stronger electron attraction and will acquire a
slightly negative charge
• The less electronegative element will acquire a
slightly positive charge