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

The sum of these vectors will give us the dipole for the molecule

For a polyatomic molecule we treat the dipoles as 3D vectors

Overlap of Orbitals

The point at which the potential energy is a minimum is called the equilibrium bond distance

The degree of overlap is determined by the system’s potential energy

equilibrium bond distance

2s

These new orbitals are called hybrid orbitalsThe 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 sp hybrid orbitals

The combination of an s orbital and a p orbital produces 2 new orbitals called sp orbitals.

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: HCCH

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.

In general delocalized p bonding is present in all molecules where we can draw resonance structures with

the multiple bonds located in different places.

Benzene is an excellent example.  For benzene the p orbitals all overlap leading to a very delocalized electron

system

Example: C6H6 benzene

The one that is lower in energy is called the bonding orbital,

The one higher in energy is called an antibonding orbital.

These two new orbitals have different energies. 

BONDING

ANTBONDING

Moleculuar Orbital (MO) Theory

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

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.

s

p

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