Chapter 3Simple Bonding Theory

• Lewis Dot Structures–Resonance–Formal Charge

• VSEPR: the subtle effects

Lewis Dot Structures

1. Count valence electrons2. Arrange atoms3. Add bonds4. Add lone pairs5. Convert lone pairs to bonding pairs (octet rule and exceptions)

Lewis Dot Structures

• Examples:CO2

SO3

N2OXeF4

ClF3

PCl6–

Why does the octet rule work?

More complex

NO2

NO

Formal Charge

= Group # - #unshared electrons on atom - # bonds to atom

Example: O3

Resonance

Example: SO3

Resonance and Formal Charge

Example: SCN-

Resonance and Formal Charge

Example: SCN-

Octet Rule vs. Pi Bonding Trends

BeF2 and

BF3

Octet Rule vs. Formal Charge

• Always follow octet ruleExceptions? SO4

2-

VSEPR

•Maximize “personal space”• CO2, SO3, SO4

2–, PCl5, SF6

• Lone pairs vs. bonding pairs?• Single bonds vs. multiple bonds?• Electronegativity effects

VSEPR

Lone Pair Effects!

Pi Bonds vs. Lone Pairs: Guess these bond angles.

Pi Bonds vs. Lone Pairs

Pi Bonds vs. Lone Pairs

Which take up more room: lone pairs or a pi bond?

Electronegativity Effects

Molecule X-P-X Angle o

PF3 97.8

PCl3 100.3

PBr3 101

Explain this trend.

Electronegativity Effects

Molecule H-X-H Angle o H2O 104.4H2S 92.1H2Se 90.6

Explain this trend.

Molecule X-As-X Angle o AsF3 AsCl3 AsBr3

Predict this trend.