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Transcript

  • 1. Chapter 2 continue Slidecast #3
  • 2. Last time.
    • Left off addressing orbitals
    • Now we can begin to appreciate bonding in organic.
  • 3. What happens when bonding occurs?
  • 4. Sigma ( σ ) bonding
    • Look at H 2 formation
    • As each hydrogen’s 1s orbitals come together if there wave functions are in phase with each other it will result in a bonding molecular orbital (bonding MO). Fig 2-6
    • This resulting bonding is centered along the nuclei resulting in a cylindrically symmetrical bond or a sigma bond
  • 5. Lets look @ Anti-bonding
  • 6. Bonding cont.
    • If the bonding is out of phase with respect to each other a antibonding orbital is produced. Fig 2-7
    • In other words they will cancel out there overlapping resulting in a node since there is no electron density
  • 7. Bonding cont.
    • In terms of there MO energy diagram
      • A bonding orbital will have LOWER energy than the respective anti-bonding orbital.
      • Stable moleulces usually have vacant antibonding orbital.
      • Lower energy is good.
  • 8. Additional types of sigma bonding
  • 9. Pi bonding
  • 10. Pi ( π )bonding
    • Pi bonding occurs when two p orbitals decide to overlap perpendicular to the sigma bond
    • The electron density at this point is centered above and below the nuclei
    • The overlap is not linear so pi MO’s are not cylindrically symmetrical
  • 11. Double bond
    • Double bonds require 4 e - ’s. fig 2-10
      • First pair go into the σ bond (the purple)
      • Second pair go into the π bonding MO above and below the σ bond.
    • This explanation is simplistic not talking about hybrid atomic orbitals.
  • 12. VSEPR
    • What is VSEPR?
      • V alence S hell E lectron- P air R epulsion
    • Why is it important?
      • Allows us to predict geometry by accounting for bond angles.
  • 13. To hybridize?
    • Shape of the molecules from the previous figure CANNOT result from simple s and p orbitals seen in general chemistry
    • To explain shapes (VSEPR) of organic molecules we assume that the s and p orbitals form hybrid atomic orbtials that separate the electron pairs more effectively in space. And put more electron density in the bonding region between the nuclei.
  • 14. sp hybridized orbital
    • Gives a linear conformation
    • 180° bond angles
  • 15.
    • Fig 2-13 example
  • 16. sp 2 hybridized orbitals
    • Give trigonal planar conformation
    • 120° bond angles
  • 17.
    • Pg 48 solutions figure
  • 18. sp 3 hybridized orbitals
    • Gives tetrahedral conformation
    • 109.5° bond angles
  • 19. Drawing 3-D Molecules If given a simple molecule you should be able to draw it 3-D.
  • 20. General Rules of Hybridization and Geometry
    • General rules we should consider for determining hybridization. These 3 are:
      • Both the sigma bonding electrons and lone pairs can occupy hybrid orbitals. The number of hybrid orbitals on an atom is computed by adding the number of sigma bonds and the number of lone pairs of electrons on that atom.
      • Use the hybridization and geometry that give the widest possible separation of the calculated number of bonds and lone pairs.
      • If two or three pairs of electrons form a multiple bond between two atoms, the first bond is a sigma bond formed by a hybrid orbital. The second bond is a pi bond, consisting of two lobes above and below the sigma bond formed by two p orbitals. The third bond of a triple bond is another pi bond, perpendicular to the first pi bond.
  • 21. Example 1. Acetylene
  • 22. Example 2. Acetaldehyde
  • 23. End.