Chapter 9 - Molecular Geometry

• Molecular Geometries
• Lewis structures DO NOT indicate shape.
• VSEPR model
• Identify number of high electron density regions (HEDR) from Lewis structure
• Unshared electron pairs
• Bonds
• note double & triple bonds count as 1 HEDR!
• Determine electronic shape from #HEDR
• 2 - linear (180)
• 3 - trigonal planar (120)
• 4 - tetrahedral (109.5)
• 5 - trigonal bipyramidal (120 (equatorial) and 90 (axial))
• 6 - octahedral (90)
• Determine molecular shape from electronic shape
• Cannot 'see' unshared electron pairs - distorts shape of what we observe.
• Tables 9.2 and 9.3
• unshared pairs take up more space than bonds - distorts expected angle slightly
• NH3 - bond angle 107
• H2O - bond angle 104.5
• multiple bond takes up more room than single bond
• CH2O - HCH @ 116; HCO @ 122
• Molecules w/no 'central atom'
• work w/1 atom at a time

• Polarity of molecules
• Must have 2 things:
• Polar bonds and
• asymmetry
• symmetric - polarity 'cancels'.
• Polar molecules have a 'dipole moment'
• measurable polarity of the molecule

• Covalent Bonding and Orbital Overlap
• Valence bond theory
• orbitals 'overlap'
• Non-hybrid bonds (H2, HCl)
• Potential energy curve (Figure 9.13)
• optimum distance between nuclei
• Hybrid Orbitals
• BeF2 - heated molecules form
• F - Be - F Dot structure
• Predict linear
• s hybridized w/p to get sp hybridization
• BF3 - one s and two p - sp2 hybridization
• CF4 - one s and three p - sp3 hybridization
• PCl5 - one s, three p, and one d - sp3d hybridization
• SF6 - one s, three p, and two d - sp3d2 hybridization
• Cannot have (d) or (e) with 2nd period elements - no d orbitals in 2nd shell.
• Determine hybridization
• Draw Lewis structure
• Count HEDR
• add orbitals until # of orbitals = # HEDR

• Multiple Bonds
• first bond - bond
• head-to-head overlap of hybridized or non-hybridized orbitals
• strongest one
• Additional bonds - bonds
• side-to-side overlap of non-hybridized p orbitals
• not as strong individually as bond.
• p orbitals available if not used for hybridization.
• Delocalization
• alternate double-single bonds can distribute electrons across length of conjugation.
• chemical evidence to support
• bond length
• energetic (absorption)

• Molecular Orbitals
• Two AO combine to form two MO
• Bonding orbital - lower energy than AOs
• Antibonding orbital - higher energy than AOs
• labeled with *
• Bond order - (Bonding electrons - antibonding electrons)
• can have fractional bond orders
• bond order = 0 means no bond exists
• H2 and He2
• Li2 and Be2
• B2 - F2
• two types of p bond - bond (1) and bond (2)
• bond pushed up in energy because of interaction w/2s
• Paramagnetism - one or more unpaired electrons
• attracted to magnetic field
• Diamagnetism - no unpaired electrons
• repelled by magnetic field
• O2
• Lewis structure shows no unpaired electrons
• MO theory shows 2 unpaired electrons
• O2 is magnetic!

Learning goals:

1. Relate the number of high electron density regions around a central atom to the electronic shape around that atom.

2. Predict the geometrical structure of a molecule or ion from its Lewis structure.

3. Explain why unshared electron pairs exert a greater repulsive interaction on other pairs than do shared pairs.

4. Predict from the molecular shape and the electronegativities of the atoms involved, whether a molecule can have a dipole moment (be polar).

5. Explain the concept of hybridization and its relationship to geometrical structure.

6. Assign a hybrid orbital set to the valence orbitals of an atom in a molecule, knowing the number and geometrical arrangement of the atoms to which it is bonded.

7. Formulate the bonding in terms of bonds and bonds from its Lewis structure.

8. Explain the concept of delocalization in bonds.

9. Explain the concept of orbital overlap, and the reason why overlap may be zero in some cases because of symmetry.

10. Describe how molecular orbitals are formed by the overlap of atomic orbitals.

11. Explain the relationship between bonding and antibonding orbitals.

12. Construct the molecular-orbital energy-level diagram for a diatomic molecule or ion built from elements of the first or second row, and predict the bond order and number of unpaired electrons.

13. Understand the difference between diamagnetism and paramagnetism, and describe the reason for the occurrence of each from a molecular level.

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