Base Pairing in DNA
Crown Ether, C12H24O6 (18-Crown-6)
Red = O
Grey = C
White = H
Purple = K
Ionic Radii
Li+ = 0.68 Å
Na+ = 0.97 Å
K+ = 1.33 Å
Rb+ = 1.47 Å
Cavity Size (O-O Dist.)
= 1.40 Å
K+ fits best
Rules for Predicting Molecular
Geometry
1. Sketch the Lewis structure of the molecule or ion
2. Count the electron pairs and arrange them in the
way that minimizes electron-pair repulsion.
3. Determine the position of the atoms from the way
the electron pairs are shared.
4. Determine the name of the molecular structure
from the position of the atoms.
5. Double or triple bonds are counted as one bonding
pair when predicting geometry.
Note: The same rules apply for molecules that contain more than one
central atom
The Dipole
A dipole arises when two electrical charges of equal
magnitude but opposite sign are separated by distance.
The dipole moment (m)
m= Qr
where Q is the magnitude of the charges and r is the distance
For a polyatomic molecule we treat the dipoles as 3D vectors
The sum of these vectors will give us the dipole for the
molecule
Overlap of Orbitals
The degree of overlap is determined by the system’s potential energy
equilibrium bond distance
The point at which the potential energy is a minimum is called
the equilibrium bond distance
Formation of sp hybrid orbitals
The combination
2s of an s orbital and a p orbital produces 2
new orbitals called sp orbitals.
These new orbitals are called hybrid orbitals
The process is called hybridization
What this means is that both the s and one p orbital are
involved in bonding to the connecting atoms
Formation of sp2 hybrid orbitals
Formation of sp3 hybrid orbitals
Hybrid orbitals can be used to explain bonding
and molecular geometry
Multiple Bonds
Everything we have talked about so far has only dealt with
what we call sigma bonds
Sigma bond (s)  A bond where the line of electron density
is concentrated symmetrically along the line connecting the
two atoms.
Pi bond (p)  A bond where the overlapping regions exist
above and below the internuclear axis (with a nodal plane
along the internuclear axis).
Example: H2C=CH2
Example: H2C=CH2
Example: HCCH
Delocalized p bonds
When a molecule has two or more resonance structures, the pi
electrons can be delocalized over all the atoms that have pi
bond overlap.
Example: C6H6 benzene
Benzene is an excellent example. For benzene the p orbitals
all overlap leading to a very delocalized electron system
In general delocalized p bonding is present in all molecules
where we can draw resonance structures with the multiple
bonds located in different places.
Moleculuar Orbital (MO) Theory
ANTBONDING
These two new orbitals have
different energies.
BONDING
The one that is lower in energy is called the bonding orbital,
The one higher in energy is called an antibonding orbital.
Energy level diagrams / molecular
orbital diagrams
MO Theory for 2nd row diatomic molecules
Molecular Orbitals (MO’s) from Atomic Orbitals (AO’s)
1. # of Molecular Orbitals = # of Atomic Orbitals
2. The number of electrons occupying the Molecular orbitals is equal to the
sum of the valence electrons on the constituent atoms.
3. When filling MO’s the Pauli Exclusion Principle Applies (2 electrons per
Molecular Orbital)
4. For degenerate MO’s, Hund's rule applies.
5. AO’s of similar energy combine more readily than ones of different
energy
6. The more overlap between AOs the lower the energy of the bonding
orbital they create and the higher the energy of the antibonding orbital.
Example: Li2
MOs from 2p atomic orbitals
s
p
1) 1 sigma bond through overlap of orbitals along the internuclear axis.
2) 2 pi bonds through overlap of orbitals above and below (or to the sides)
of the internuclear axis.
Interactions between the 2s and 2p orbitals
The s2s and s2p
molecular orbitals
interact with each
other so as to
lower the energy
of the s2s MO and
raise the energy of
the s2p MO.
For B2, C2, and N2 the interaction is so strong that the s2p is
pushed higher in energy than p2p orbitals
Paramagnetism of O2