2. Molecular Orbital Theory

3. Molecular Orbitals Just as atomic orbitals belong to a particular atom, a molecular orbital belongs to molecules as a whole Each orbital is filled with 2 electrons There are 2 types of bonding orbitals: sigma and pi

4. Sigma Bond (  )  A Sigma Bond is when 2 atomic orbitals combine to form a molecular orbital that is symmetrical around the axis S orbitals overlapping P orbitals overlapping end-to-end

5. Pi Bond (  )  Pi bonding electrons are likely to be found in a loop- shape above and below the axis Pi bonds are weaker than sigma bonds because they overlap less Pi bonds are weaker than sigma bonds because they overlap less P orbitals overlapping side-by-side

6. BONDING ANTIBONDING Since orbitals are waves of electrons, they can combine either constructively (forming a bonding molecular orbital), or destructively (forming an antibonding* molecular orbital).

7. MO Theory 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) and Aufbau’s principle Applies (start at the lowest energy) 4. For equal energy MO’s, Hund's rule applies.

8. Bond Order  To find bond order add total electrons in the bonding orbitals and subtract total electrons in antibonding orbitals, then divide by 2

9. Example: Li 2

10. MO Diagram for O 2 http://www.chem.uncc.edu/faculty/murphy/1251/slides/C19b/sld027.htm

11. Molecular Orbital Diagram (HF)