This document discusses bonding types, including ionic and covalent bonding. It introduces Lewis structures as a model to represent covalent bonding using dots to represent valence electrons and lines for bonding pairs. Formal charge is defined as a way to track the distribution of charge within a molecule based on an atom's group number and bonded/nonbonded electrons. Resonance structures are discussed as an extension of Lewis structures to account for delocalized charge.

1. Bonding Types – ionic and covalent
Covalent Bonding – equal/unequal sharing
Bonding Model – Lewis Structures (with
formal charge and resonance structures)

2. How compounds are held together:
covalent vs ionic bonding

3. Electronegativity:
measure of attraction for electrons in a bond

4. EN and Bond Polarity Scale
Very slightly polar No sharing
(called nonpolar) (Ionic bond)
Nonpolar Polar
0 0.4 1.9
Electronegativity Difference

5. Electronegativity Trends

6. Bond vs Molecular Polarity

7. Lewis Structures: model to describe
bonding in covalent molecules
- Valence electrons represented by dots
- Bonding pair represented by line
- Valence e- so 2 for H and 8 for 2nd row elements

8. Formal Charge: way of keeping track* of
where charge is within a molecule
Electrons that contribute to atom’s charge:
1. All of unshared electrons
2. Half of shared electrons
FC = group number – nonbonding e- – ½ shared e-

9. Resonance (when 1 Lewis structure
doesn’t tell the story well)

10. Resonance “rules of thumb”
• Must be valid Lewis structure
• Move electrons, not nuclei
• Number of unpaired electrons (if any) remain
the same
• Major contributor has lowest energy (see
other rules)
• Resonance – most important when charge is
delocalized

11. Comparing Resonance Structures
• As many octets as possible
• As many bonds as possible
• Any negative charges on electronegative
atoms
• As little charge separation as possible

12. Common Bonding Patterns