1 / 19

Lecture 13

Lecture 13. Valence Bond Theory 3.4 – 3.7 17-September Assigned HW 3.34, 3.36, 3.40, 3.42 Due: Monday 20-Sept. Review 2.13-2.15 D4-D5. We can predict the shape of molecules using 4 rules of VSEPR Separate electron density as much as possible – this includes lone pairs and bonds

winda
Download Presentation

Lecture 13

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Lecture 13 Valence Bond Theory 3.4 – 3.7 17-September Assigned HW 3.34, 3.36, 3.40, 3.42 Due: Monday 20-Sept

  2. Review 2.13-2.15 D4-D5 • We can predict the shape of molecules using 4 rules of VSEPR • Separate electron density as much as possible – this includes lone pairs and bonds • We don’t care about bond order – single bonds are the same as triple bonds when we assign shape • When picking a shape, only bonds matter. Lone pairs are important for determining orbital layout • LP-LP > LP-atom > atom-atom • Stepwise: • Lewis Structure • VSEPR Formula • Predict orbital shape (electron arrangement) Linear, trigonal planar, tetrahedral, trigonalbipyramidal, octahedral, pentagonal bipyramidal • Consider LP influence • VSEPR formula  AXnEm

  3. Shapes of Molecules

  4. Forming a Bond What are covalent bonds? E How do we describe the electron in an atom? 2s 1s + Orbital Overlap

  5. Forming a Bond – sigma bonds What are covalent bonds? E + E E 1s 1s

  6. Forming a Bond – sigma bonds E 1s E E 1s Molecular Orbital

  7. Forming a Bond – sigma bonds What about H-Cl E Why don’t I include 1s electrons for Cl? E E 2p 2s + 1s Why don’t we include the other 2s and 2p electrons?

  8. Forming a Bond – sigma bonds What about H-Cl Note that only ONE of the 2p orbitals from Cl is involved in the sigma bond. What are the others doing?

  9. Forming a Bond – pi bonds N2 E E E 2p 2p 2s 2s Where are the lone pairs? Can we form 3 sigma bonds?

  10. Forming a Bond – double and triple bonds p s

  11. Forming a Bond – double and triple bonds p s

  12. Forming a Bond – Atomic Hybridization CH4 E Carbon 2p E E 2s 2p 2s So we’ll have one sigma bond that is more stable? How do we get the four bonds from this? Form a HYBRID orbital

  13. Forming a Bond – Atomic Hybridization Remember that orbitals are from wavefunctions….and describe electrons – which are waves And waves can interfere constructively or destructively 4 possible combinations of wave interference for 4 different waves – these are our 4 hybrid orbitals. E

  14. Forming a Bond – Atomic Hybridization E

  15. Forming a Bond – Atomic Hybridization 109.5° We needed one s and 3 p orbitals to make this hybrid

  16. Atomic Hybridization Ammonia E sp3 3 individual Hydrogen 1s orbitals

  17. Atomic Hybridization Ethane So we have a sigma bond between the carbon sp3 hybrid and the 1s orbital of hydrogen

  18. Atomic Hybridization

  19. Atomic Hybridization Ethene Shape? sp2

More Related