1. THE “MAKEUP” LECTURE TOPICS Molecular Polarity (8.7) Introduction to Bonding Theories (9.1) Valence Bond Theory (9.2) December 1, 2009

2. Molecular Polarity Polarity = uneven distribution of charge  Bond polarity = Electrons drawn closer to the more electronegative atom  Molecular polarity = Molecule as a whole has a net separation of charge  A polar molecule must have polar bonds  A molecule is polar if the directions of the polar bonds don’t cancel/offset eachother  Nonpolar Examples: CH 4, CO 2  Polar Examples: H 2 O, NH 3

3. Nonpolar Examples: CH 4, CO 2

4. Polar Examples: H 2 O, NH 3

5. Why is Polarity Important? Polarity dictates many molecular properties  Physical state (solid, liquid, gas)  CO 2 (44 g/mol) vs. H 2 O (18 g/mol)  Solubility

6. Chapter 9- Chemical Bonding Theories Valence Bond Theory: Uses Lewis Structures  Bonds form using shared electrons between overlapping orbitals on adjacent atoms.  Orbitals arrange around central atom to avoid each other.  Two types of bonds: sigma (  ) and pi (  ).  Qualitative, visual- good for many atom systems in ground state Molecular Orbital Theory: Uses MO Diagrams  Orbitals on atoms “mix” to make molecular orbitals, which go over 2 or more atoms.  Two electrons can be in an orbital.  Quantitative- needed to describe excited states

7. Sigma (  ) Bonding Orbitals on bonding atoms overlap directly between bonding atoms

8. Consider VSEPR Shapes and bonding: Sigma (  ) Bonding

9. What’s wrong with this picture? Atoms bond by having their valence orbitals overlap

10. Bonding orbitals are not the same shape as atomic orbitals Electron configurations: H = 1s 1 C = 1s 2 2s 2 2p 2 2s 2p z 2p x 2p y Orbitals in CH 4 Atomic orbitals change shape when they make molecules

11. Hybrid Orbitals