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Chapter 3 Simple Bonding Theory

Chapter 3 Simple Bonding Theory. Lewis Dot Structures Resonance Formal Charge VSEPR: the subtle effects. Lewis Dot Structures. 1. Count valence electrons 2. Arrange atoms 3. Add bonds 4. Add lone pairs 5. Convert lone pairs to bonding pairs (octet rule and exceptions).

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Chapter 3 Simple Bonding Theory

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  1. Chapter 3Simple Bonding Theory • Lewis Dot Structures • Resonance • Formal Charge • VSEPR: the subtle effects

  2. Lewis Dot Structures 1. Count valence electrons 2. Arrange atoms 3. Add bonds 4. Add lone pairs 5. Convert lone pairs to bonding pairs (octet rule and exceptions)

  3. Lewis Dot Structures • Examples:CO2SO3N2OXeF4ClF3PCl6–

  4. Why does the octet rule work?

  5. More complex NO2 NO

  6. Formal Charge = Group # - #unshared electrons on atom - # bonds to atom Example: O3

  7. Resonance Example: SO3

  8. Resonance and Formal Charge Example: SCN-

  9. Resonance and Formal Charge Example: SCN-

  10. Octet Rule vs. Pi Bonding Trends BeF2 and BF3

  11. Octet Rule vs. Formal Charge • Always follow octet ruleExceptions? SO42-

  12. VSEPR • Maximize “personal space” • CO2, SO3, SO42–, PCl5, SF6 • Lone pairs vs. bonding pairs? • Single bonds vs. multiple bonds? • Electronegativity effects

  13. VSEPR

  14. Lone Pair Effects!

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

  16. Pi Bonds vs. Lone Pairs

  17. Pi Bonds vs. Lone Pairs Which take up more room: lone pairs or a pi bond?

  18. Electronegativity Effects Molecule X-P-X Angle o PF3 97.8 PCl3 100.3 PBr3 101 Explain this trend.

  19. Electronegativity Effects Molecule H-X-H Angle o H2O 104.4 H2S 92.1 H2Se 90.6 Explain this trend. Molecule X-As-X Angle o AsF3 AsCl3 AsBr3 Predict this trend.

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