Chapter 10

1. Chapter 10 • Molecular Geometries and Bonding Theories • Lewis structures do not indicate the molecular architecture – the shape of the molecule. • The shape and structure of a molecule determines much of its physical and chemical characteristics.

2. VSEPR Theory • Valence-shell Electron Pair Repulsion • Electron pairs (domains or regions) repel each other completely. • Balloon model.

3. Electron Regions • The number of electron regions around the central atom are counted as: • Each single bond counts as a region. • Each lone pair counts as a region. • A multiple bond counts as a single region.

4. Electron Regions • How many?

5. Electron Pair Geometry (EPG) • Can be from two to six regions. • Thus, only five EPG’s are possible. • Two regions produces a linear EPG.

6. Electron Pair Geometry • Three regions produces a trigonal planar geometry. • Planar = 2D. • Ex) BF3

7. Electron Pair Geometry • Four regions becomes a three dimensional structure based on the tetrahedron. • Formally called tetrahedral with bond angles of 109.5o

8. Electron Pair Geometry • Tetrahedral is very common and symmetrical. • An example is CF4

9. Electron Pair Geometry • Five regions produces a trigonal bipyramidal geometry with two sets of bond angles.

10. Electron Pair Geometry • An example is PCl5

11. Electron Pair Geometry • Six regions produces an octahedral geometry.

12. Electron Pair Geometry • An example is SF6

13. Molecular Geometry (MG) • This is based on the shape of the electron pairs. • When a molecule has no lone pairs, the EPG = MG. • If the molecules DOES have one or more lone pairs, then the shape of the atoms is determined based off of the EPG.

14. Molecular Geometry

15. Molecular Geometry

16. Molecular Geometry

17. Examples • Bent (120), SO2 • Trigonal pyramidal, NH3 • Bent (109.5), H2O • Seesaw, SF4 • T-shaped, ClF3 • Linear, I3- • Square pyramidal, BrF5 • Square planar, XeF4

18. Sketching the Molecules • Simple = Ball and Stick figures • Representing the 3D shapes: • Put as many of the molecules in the same plane as possible including the central atom. Use straight lines for bonds connected to atoms in plane. • For atoms in front of the plane, use a solid wedge. • For atoms behind the plane, use a hashed wedge.

19. 3D Representations

20. Lone Pairs • A non-bonding pair will always take up more space. • This compresses the normal bond angles.

21. Lone Pairs

22. Lone Pairs • This also explains the MG’s for the trigonal bipyramidal family.

23. Shapes of Larger Molecules • A molecule like acetic acid has three central atoms.

24. Shapes of Larger Molecules

25. Polarity • A molecule can contain very polar bonds, but can be non-polar. • An example is CO2.

26. Polarity • On the other hand, sometimes polar bonds DO make a molecule polar. • An example of a polar molecule is H2O.

27. Polarity

28. Polarity • A molecule with a symmetrical distribution of polar bonds will be non-polar. • A molecule with an un-symmetrical distribution of polar bonds will be polar. • presence of lone pairs • different external atoms

29. Polarity

30. Polarity • Polar molecules are attracted to other polar molecules • Because water is a polar molecule, other polar molecules dissolve well in water • and ionic compounds as well • Non-polar molecules do NOT dissolve in water.

31. Valence Bond Theory • How can we explain the formation of the bonds in a molecular compound? • A bond occurs when a valence orbital on one atom overlaps with a valence orbital of another atom.

32. Valence Bond Theory • The H2 molecule – a closer look. nuclear repulsion no inter-action minimum energy

33. Valence Bond Theory • Three (or more) atom molecules cannot be explained by simple overlap of orbitals. • Fact: a bond generally forms between two half-filled orbitals. • Fact: an s-type orbital is spherical, so it could form a bond in any direction. • Fact: the three p-type orbitals are at 90 degree angles to each other.

34. Valence Bond Theory • CH4 – has an EPG and MG of tetrahedral with bond angles of 109.5o. • Valence diagram for C and H before any bonding is:

35. Valence Bond Theory • Solution: promote the paired electron from the s orbital to the empty p orbital. • Solution: mix the one s and three p orbitals together to get a new set of four orbitals all equal in energy. This is called _____________________.

36. Valence Bond Theory • Each hybrid orbital has some s and some p characteristics. • Thus, they look different!

38. Types of Hybrids • Determined from the EPG.

39. Types of Hybrids • Atoms in the third period and beyond have empty d orbitals that can potentially be used for hybridization. • PCl5 – requires five bonds, so need a set of five orbitals. • Once again, must first promote the s electron to an empty d orbital.

40. Types of Hybrids

41. Molecules with Lone Pairs • Ex) NH3 • Ex) H2O • Ex) BrF3

42. Multiple Bonds • Two types of bonds are possible. • 1. Sigma (s) bonds have a cylindrical shape of electron density along the central axis between the two nuclei. s bond

43. Multiple Bonds • 2. Pi (p) bonds have an electron density above and below the central axis. • Are formed by the overlap of two parallel half-filled p-type orbitals.

44. Multiple Bonds • The majority of bonds are sigma bonds. • When a double bond is present, the first bond is a sigma and the second is a pi.

45. Pi bonds

46. Multiple Bonds • For any pi bonds, you MUST use an un-hybridized half-full p-type orbital. • Ex) C2H4 • Ex) CO2

47. Multiple Bonds

48. Pi Bond Significance • Sigma bonds have free-rotation about the central axis. • Ex) C2H4Cl2 • Pi bonds have NO free-rotation due to the fact that they must overlap above and below the central axis. • Ex) C2H2Cl2

49. Pi Bond Significance

50. Isomers • When two compounds share the exact same formula but are different either structurally or spatially, then they are said to be isomers. • Structural isomers • C5H12 • C2H6O