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CH 10: Molecular Geometry

CH 10: Molecular Geometry. Renee Y. Becker Valencia Community College CHM 1045. Molecular Shapes: VSEPR. The approximate shape of molecules is given by Valence-Shell Electron-Pair Repulsion (VSEPR) . Step 01: Count the total electron groups.

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CH 10: Molecular Geometry

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  1. CH 10: Molecular Geometry Renee Y. Becker Valencia Community College CHM 1045

  2. Molecular Shapes: VSEPR • The approximate shape of molecules is given by Valence-Shell Electron-Pair Repulsion (VSEPR). • Step 01: Count the total electron groups. • Step 02: Arrange electron groups to maximize separation. • Groups are collections of bond pairs between two atoms or a lone pair. • Groups do not compete equally for space:Lone Pair > Triple Bond > Double Bond > Single Bond

  3. Molecular Shapes: VSEPR • Two Electron Groups: Electron groups point in opposite directions.

  4. Molecular Shapes: VSEPR • Three Electron Groups:Electron groups lie in the same plane and point to the corners of an equilateral triangle.

  5. Molecular Shapes: VSEPR • Four Electron Groups: • Electron groups point to the corners of a regular tetrahedron.

  6. Molecular Shapes: VSEPR Five Electron Groups: Electron groups point to the corners of a trigonal bipyramid.

  7. Molecular Shapes: VSEPR • Six Electron Groups: Electron groups point to the corners of a regular octahedron.

  8. Molecular Shapes: VSEPR

  9. Example 1: VSEPR • Draw the Lewis electron-dot structure and predict the shapes of the following molecules or ions: O3 H3O+ XeF2 PF6– XeOF4 AlH4– BF4– SiCl4 ICl4– AlCl3

  10. Valence Bond Theory 1. Covalent bonds are formed by overlapping of atomic orbitals, each of which contains one electron of opposite spin. 2. Each of the bonded atoms maintains its own atomic orbitals, but the electron pair in the overlapping orbitals is shared by both atoms. 3. The greater the amount of orbital overlap, the stronger the bond.

  11. Valence Bond Theory • Linus Pauling: Wave functions from s orbitals & p orbitals could be combined to form hybrid atomic orbitals.

  12. sp hybrid:

  13. sp2 hybrid:

  14. sp2 hybrid (π bond):

  15. sp3 hybrid:

  16. sp3d hybrid:

  17. sp3d2 hybrid:

  18. Hybridization Easy Way

  19. Molecular Orbital Theory • The molecular orbital (MO) model provides a better explanation of chemical and physical properties than the valence bond (VB) model. • Atomic Orbital: Probability of finding the electron within a given region of space in an atom. • Molecular Orbital: Probability of finding the electron within a given region of space in a molecule.

  20. Molecular Orbital Theory • Additive combination of orbitals (s) is lower in energy than two isolated 1s orbitals and is called a bonding molecular orbital.

  21. Molecular Orbital Theory • Subtractive combination of orbitals (s*) is higher in energy than two isolated 1s orbitals and is called an antibonding molecular orbital.

  22. Molecular Orbital Theory • Molecular Orbital Diagram for H2:

  23. Molecular Orbital Theory • Molecular Orbital Diagrams for H2– and He2:

  24. Molecular Orbital Theory • Additive and subtractive combination of p orbitals leads to the formation of both sigma and pi orbitals.

  25. Molecular Orbital Theory • Second-Row MO Energy Level Diagrams:

  26. Molecular Orbital Theory • MO Diagrams Can Predict Magnetic Properties:

  27. Molecular Orbital Theory • Bond Order is the number of electron pairs shared between atoms. • Bond Order is obtained by subtracting the number of antibonding electrons from the number of bonding electrons and dividing by 2. BO = Bonding electrons – antibonding electrons 2

  28. Example 2: Molecular Orbital Theory • The B2 and C2 molecules have MO diagrams similar to N2. What MOs are occupied in B2 and C2, and what is the bond order in each? Would any of these be paramagnetic?

  29. B2 B has 5 electrons So B2 has 10 elec Core electrons don’t count toward BO *2p *2p 2p 2p *2s 2s *1s 1s

  30. C2 Carbon has 6 electrons so C2 has 12 electrons *2p *2p 2p 2p *2s 2s *1s 1s

  31. Dipole moment • Based on electronegativity differences between atoms in a molecule • The most electronegative atom is partially negative • The less electronegative atom is partially positive • The dipole moment is the average of all dipoles in the molecule • Exception (C-H bonds do not have a dipole)

  32. If a molecule has a dipole moment it is polar • A compound could contain polar bonds but the molecule could be non-polar because there is no dipole moment • Bond dipoles can cancel each other out

  33. Example 3: Dipole moment • Draw the dipole moment for the following molecules, are they polar? HCl NH3 CHCl3 H2O SF6 CCl4 CO2 CH2Cl2

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