2. Molecular Orbital Theory (What is it??)  Better bonding model than valence bond theory  Electrons are arranged in “molecular orbitals”  Dealing with valence shell electrons

3. Molecular OrbitalsMolecular Orbitals  Combination of atomic orbitals  2 atomic orbitals------2 molecular orbitals  Molecular region where electrons are likely to be found within a chemical compound  Behave like molecules  Deals with electrons arranged in molecules, NOT atoms

4. Bonding Molecular Orbital ( σ )  Contain electrons involved in chemical bonding  Contribute to bond strength  Increase stability  Lower energy level for electrons  Increased electron density between atoms

5. Anti-bonding Molecular Orbitals ( σ *)  Contain electrons NOT involved in bonding  Electrons hang out away from bond  Decreased stability and bond strength  Higher energy level for electrons  Decreased electron density between atoms

6. **Electrons want to be at a LOW energy level SO---generally pair up and reside in bonding molecular orbitals.

7. How are electrons placed in molecular orbitals? 1)Electrons want to be in the lowest-energy molecular orbitals as possible. 2)Only 2 electrons found in each molecular orbital. 3)Electrons are placed in molecular orbitals by themselves (parallel spins) unless they have to be paired up (opposite spins).

8. Bond OrderBond Order  Determined by molecular orbitals = (# electrons in bonding MO) – (# electrons in anti-bonding MO) 2

9. Example 2: Molecular orbital energy-level diagram H2+H2+

10. What happens when “p” atomic orbitals combine?  Each “p” orbital combines with another “p” orbital—2 molecular orbitals produced  Of the 2 molecular “p” orbitals—  1 lower energy bonding orbital  1 higher energy anti-bonding orbital  One p orbital produces orbital overlap  σ p, σ * p  Other 2 p orbitals overlap in parallel  π p, π* p

11. Example 3: O 2

12. Homework  Read pp. 417-418  Problems #61, 63, 67  Read over Thursday lab procedure