1. Molecular Geometry & Bonding Theories
Molecular Shapes
VSEPR Model
Molecular Polarity
Covalent Bonding
Hybrid Orbitals
Multiple Bonds

2. Introduction
Molecules have shapes and sizes that are defined by the angles and distances between nuclei of atoms.
Shape, size, and strength and polarity of bonds determine properties of a substance.
We start with Lewis structures to determine the number and types of bonds between atoms.

3. Electron Domain
Electron domain: a region in which e-’s will most likely be found.
Bonding pair: e- domain between two atoms.
Non-bonding pair (lone pair): e- domain located mainly on one atom.
An e- domain consists of a nonbonding pair, a single bond, or a multiple bond.

4. Determining Number of Electron Domains

5. VSEPR Model = “Valence Shell Electron Pair Repulsion” Model
e-’s are negatively charged, so they repel each other.
VSEPR Model says that the best arrangement of a given number of e- domains is the one that minimizes the repulsions among them.

6. VSEPR Model Using Balloons
Electron domain geometry: arrangement of e- domains around a central atom.

7. Electron-Domain Geometries

8. Molecular Geometry The actual spacial arrangement of atoms.
Molecular geometry is determined from e--domain geometry.
To predict molecular shapes with VSEPR model, draw Lewis structure, count e--domains, then use the arrangement to determine molecular geometry.

9. Molecular Shapes from Electron-Domain Geometries

10. More Molecular Shapes from Electron-Domain Geometries

12. Molecular Geometry Practice
SeCl2
Predict the molecular geometries: O3
SnCl3-
CO32-

13. Lone Pair Electrons on Trigonal Bipyramidal Structures
Nonbonding pairs always occupy equatorial positions on a trigonal bipyramidal structure.
Can you explain why?

14. Lone Pair Electrons on Octahedral Structures
Nonbonding pairs always occupy the axial positions first.
Why?

16. Effect of Non-bonding Electrons & Multiple Bonds
Non-bonding pairs and multiple bonds exert greater repulsive forces on adjacent domains and tend to compress bond angles.
Lone pair > triple bond > double bond > single bond

17. Bond Angle of Water
Water’s H-O-H bond angle is always 104.5º due to two lone electron pairs.

18. Bond Angle of Methane and Ammonia
NH3 has a smaller bond angle than methane due to the lone pair of electrons.

19. Bond Dipoles
Bond dipoles and dipole moments are vectors (magnitude + direction).
Overall dipole moment of a molecule is the sum of its bond dipoles.

20. Polarity of Water
If the water molecule were linear, water would NOT be polar.

22. Polar or Nonpolar?
BrCl
Yes. All diatomics with polar bonds are polar molecules.
SO2
Yes. Bent molecule. O’s more neg.
NF3
Yes. Trigonal bipyramidal geometry.
BCl3
No. Symmetry in trigonal planar geometry
SF6
No. Symmetry in octahedral arrangement.

23. Molecular Shape & Molecular Polarity
Molecular polarity has a significant effect on physical and chemical properties.
For a molecule with more than two atoms, dipole moment depends on polarities of individual bonds and the molecular geometry.

24. Valence-Bond Theory
Covalent bonds form when a valence atomic orbital of one atom merges with that of another atom.
The orbitals overlap (share a region of space).
The overlap allows two e-’s of opposite spin to share common space.

25. Formation of Bonds in H2
The bond in H2 forms from the overlap of two 1s orbitals from two hydrogen atoms.

26. Bonds From Orbital Overlap

27. Hybridization
To explain geometries, we assume atomic orbitals mix to form new hybrid orbitals.
Hybridization: the process of mixing and changing atomic orbitals as atoms approach to form bonds.

28. Hybridization: sp
Linear arrangement of e- domains means sp hybridization.
One s-orbital and one p-orbital hybridize to form two equivalent sp hybrid orbitals.

30. Hybridization: sp2
Trigonal planar arrangement means sp2 hybridization.
One s-orbital and two p-orbitals hybridize to form three sp2 hybrids.

32. Hybridization: sp3 Tetrahedral arrangement means sp3 hybridization.
Four sp3 hybrids form from one s-orbital and three p-orbitals.

35. Sigma () Bond
Sigma () bond: e- density concentrated symmetricaly around the line connecting the nuclei (internuclear axis).
Single bonds are  bonds.
Can be made from s- or p-orbitals.
Allows rotation at bond.

36. Pi (∏) Bond
∏ bond: covalent bond in which there is a side-to-side overlap of p-orbitals.
p-orbitals are perpendicular to internuclear axis.
Overlap regions lie above and below internuclear axis.
Less orbital overlap than in  bonds, so ∏ bonds are weaker.
Does NOT allow rotation around bond.

37. Double & Triple Bonds
Double bonds consists of one  bond and one ∏ bond.
Triple bonds consist of one  bond and two ∏ bonds.
In a double bond, one set of p-orbitals overlap above and below the internuclear axis.
In a triple bond, the second set of p-orbitals overlap in front of and behind the internuclear axis.

38. Bonding in Ethylene
Ethylene is planar.

39. Bonding in Acetylene

40. Delocalized ∏ Bonding
Resonance structures with ∏ bonds have delocalization of electrons.
The electrons in these bonds extend over more than two bonded atoms.