1. Aufbau Principle and Hund’s Rule

2. Electron Configurations and the Periodic Table

3. Effective Nuclear Charge
Simplest (Zeff = Z – core electrons)
Sodium: Zeff = 11 – 10 = 1 [Cl: 7]
Use Bohr’s equation, substituting Zeff for Z and using nfinal = ∞ for IE.
Sodium: Zeff = [Cl: 2.93]
Slater’s rule: differing effects of electrons in different orbitals
Sodium: Zeff = 2.2 [Cl: 6.1]

4. Three Trends Atomic radius Ionization Energy Electron Affinity
Energy required to remove an electron from a gaseous atom or ion.
Electron Affinity
Energy change associated with the addition of an electron to a gaseous atom or ion.

6. IE vs. Atomic Number

7. IE vs. Atomic Number

8. IE vs. Atomic Number

9. IE vs. Atomic Number

10. Successive IEs

11. EA vs. Atomic Number

12. Boxes of different dimensions

13. Forming a molecule
Simply combining atomic orbitals.

14. Simplest Model: Valence Bond Theory

15. But What About… Methane (CH4)? Carbon monoxide (CO)?
Why not CH2 (since we have two unpaired 2p electrons)?
Why is this a tetrahedral (since the p orbitals are orthogonal to each other) with four equivalent bonds in the Lewis structure?
Carbon monoxide (CO)?
Why is there a triple bond (we can’t see this from atomic Lewis structures)?
Nitrogen monoxide (NO)?
Odd number of electrons.
Oxygen (O2)? (O2+)?
Why is O2 magnetic?
Why is the bond strength of O2+ greater than that of O2?

16. Hybridization (Part of VBT)
The 2s and 2p orbitals are similar in energy.
Recall that they are identical for hydrogen.
Combining the wavefunctions for 2s and 2p gives a solution to the Schrödinger equation.
The s orbital is combined with p orbitals – the number of p orbitals used depends on the molecule to be formed.
The resulting (new) atomic orbitals are hybridized orbitals.
A molecule is not simply atoms stuck together.