1 / 29

Molecular Geometry

Molecular Structure. Molecular Geometry. Overlap of Orbitals. The degree of overlap is determined by the system’s potential energy. equilibrium bond distance. The point at which the potential energy is a minimum is called the equilibrium bond distance. Formation of sp hybrid orbitals

fern
Télécharger la présentation

Molecular Geometry

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. Molecular Structure Molecular Geometry

  2. Overlap of Orbitals

  3. The degree of overlap is determined by the system’s potentialenergy equilibrium bond distance The point at which the potentialenergy is a minimum is called the equilibrium bond distance

  4. Formation of sp hybrid orbitals The combination of an s orbital and a p orbital produces 2 new orbitals called sp orbitals. 2s These new orbitals are called hybrid orbitals The process is called hybridization What this means is that both the s and one p orbital are involved in bonding to the connecting atoms

  5. Formation of sp2hybrid orbitals

  6. Formation of sp3hybrid orbitals

  7. Hybrid orbitals can be used to explain bonding and molecular geometry

  8. Multiple Bonds Everything we have talked about so far has only dealt with what we call sigma bonds Sigma bond (s)  A bond where the line of electron density is concentrated symmetrically along the line connecting the two atoms.

  9. Pi bond (p)  A bond where the overlapping regions exist above and below the internuclear axis (with a nodal plane along the internuclear axis).

  10. Example: H2C=CH2

  11. Example: H2C=CH2

  12. Example: HCCH

  13. A. VSEPR Theory • Valence Shell Electron Pair Repulsion Theory • Electron pairs orient themselves in order to minimize repulsive forces. • The most stable arrangement of groups attached to a central atom is the one that has the maximum separation of electron pairs(bonded or nonbonded).

  14. Lone pairs repel more strongly than bonding pairs!!! A. VSEPR Theory • Types of e- Pairs • Bonding pairs - form bonds • Lone pairs - nonbonding e-

  15. Bond Angle A. VSEPR Theory • Lone pairs reduce the bond angle between atoms.

  16. Know the 8 common shapes & their bond angles! B. Determining Molecular Shape • Draw the Lewis Diagram. • Tally up e- pairs on central atom. • double/triple bonds = ONE pair • Shape is determined by the # of bonding pairs and lone pairs.

  17. BeH2 C. Common Molecular Shapes 2 total 2 bond 0 lone LINEAR 180°

  18. BF3 C. Common Molecular Shapes 3 total 3 bond 0 lone TRIGONAL PLANAR 120°

  19. SO2 C. Common Molecular Shapes 3 total 2 bond 1 lone BENT <120°

  20. CH4 C. Common Molecular Shapes 4 total 4 bond 0 lone TETRAHEDRAL 109.5°

  21. NH3 C. Common Molecular Shapes 4 total 3 bond 1 lone TRIGONAL PYRAMIDAL 107°

  22. H2O C. Common Molecular Shapes 4 total 2 bond 2 lone BENT 104.5°

  23. PCl5 C. Common Molecular Shapes 5 total 5 bond 0 lone TRIGONAL BIPYRAMIDAL 120°/90°

  24. SF6 C. Common Molecular Shapes 6 total 6 bond 0 lone OCTAHEDRAL 90°

  25. F P F F D. Examples • PF3 4 total 3 bond 1 lone TRIGONAL PYRAMIDAL 107°

  26. OCO D. Examples • CO2 2 total 2 bond 0 lone LINEAR 180°

More Related