1. DEPARTMENT OF CHEMISTRY
GEOMETRY OF MOLECULES
NIKAM N.D.
DEPARTMENT OF CHEMISTRY

11. 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.

13.  Note: The same rules apply for molecules that contain more than one central atom

14. The Dipole
    A dipole arises when two electrical charges of equal magnitude but opposite sign are separated by distance.

15. where Q is the magnitude of the charges and r is the distance
The dipole moment (m)
= Qr 
where Q is the magnitude of the charges and r is the distance

16. For a polyatomic molecule we treat the dipoles as 3D vectors
The sum of these vectors will give us the dipole for the molecule

17. Overlap of Orbitals

18. 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

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

20. Formation of sp2 hybrid orbitals

21. Formation of sp3 hybrid orbitals

22. Hybrid orbitals can be used to explain bonding and molecular geometry

23. 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.

24. Pi bond (p)  A bond where the overlapping regions exist above and below the internuclear axis (with a nodal plane along the internuclear axis).

25. Example: H2C=CH2

26. Example: H2C=CH2

27. Example: HCCH

28. 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.

29. 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.

30. 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.

31. Energy level diagrams / molecular orbital diagrams

32. 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.

33. Example: Li2

34. 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.

36. 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.

37. For B2, C2, and N2 the interaction is so strong that the s2p is pushed higher in energy than p2p orbitals

38. Paramagnetism of O2