Topic 12 : Advanced Bonding Concepts

1. Topic 12 : Advanced Bonding Concepts • LECTURE SLIDES • VSEPR shape theory • Bond Polarity • Molecular Polarity Kotz & Treichel, 9.6- 9.8

2. Molecular and Polyatomic Ion Shapes Once a Lewis structure is drawn, the three - dimensional geometry of the species can easily be determined by utilizing the “valence shell electron pair repulsion theory” called “VSEPR”: “VSEPR” theory is based on the tendency of negatively charged regions to repel each other and align as far apart as possible, resulting in predictable shapes for any covalently bonded species.

3. To utilize “VSEPR”, the number of regions of electron • density around the central atom in the species is • counted. • Count as “one region”: • Single Bonds • Unshared Pairs • Multiple bonds between same two atoms

4. Examples of “four regions”: “two regions”: “three regions”:

5. Basic Shapes predicted by VSEPR: Two regions: Three Regions: Bond Angles Geometry

6. Five Regions: Four Regions: Six Regions:

8. Before we begin, some guidelines about forming double and triple bonds in Lewis structures: C, N, O, S form double and triple bonds and never show incomplete octets (less than 8 e’s) Metals, metalloids, and halogensdo not as a rule form multiple bonds. Compounds containing these elements will often show an incomplete octet around the central atom.

9. Type One: Two Regions Examples: BeCl2, CO2, NO2+

12. Type Two: Three Regions NO3-, NO2-

15. Black orbital indicates pair of unshared e’s NOTE: “molecular geometry” (bonds only): BENT

16. Group Work 12.1: Geometry, 2,3 Regions 1. Do a Lewis Structure for HCN and CH2O. 2. Draw each molecule “to shape” 3. Describe geometry and bond angles for each

17. Type Three: Four Regions CH2Cl2, NH3, H2O, NH4+

19. Note: molecular geometry, trigonal pyramid

20. As is turns out, unshared pairs of electrons around the central atom are not held in place between two atoms as bonded pairs are. They tend to occupy more space and to be somewhat more “repulsive” than bonded pairs. When grouped with bonded pairs to tiny atoms like H, they tend to distort the bond angles, pushing the bonded pairs closer together. The bond angles in ammonia are closer to 107o.

21. Note: molecular geometry: BENT

23. GROUP WORK 12.2: Geometry, 3, 4 Regions Do Lewis structure and assign shape and bond angles: CO32-, SiCl4

24. Type Four: Five Regions PF5 , ClF3 ,IF2-, SF4

25. Bond angles, each “axial” F, 90o from trigonal plane Bond angles in “equatorial” position 120o

26. Note: “T-shaped”; unshared pairs always“equatorial”

27. Note: “linear” molecular geometry

28. Note: “Seesaw” molecular geometry

29. Type Five: Six Regions SF6, IF5, XeF4

32. Note: molecular geometry “square pyramidal”

33. Note: “Square planar”

34. GROUP WORK 12.3: Geometry, 5, 6 Regions Do Lewis structure and assign shape and bond angles: ICl4+, XeOF2, ICl4- Note: O in XeOF2 is equatorial, experimental evidence

35. To see relevance of “shape work”, let’s turn next to bond and molecular polarity. To help examine this topic we turn back to the property of “electronegativity”:

36. ELECTRONEGATIVITY The trends in ionization energies and electron affinities can be thought of as summarized in a single property called “electronegativity” (en or X). Electronegativityis a unit-less set of assigned values on a scale of 0 --> 4 describing the ability of an atom to attract electrons to itself. The values reaches a maximum at fluorine, with an X =4. Nonmetals have thelargest values,metals the lowest. Noble gases have no assigned Xvalue.

37. Most active non-metals Most active metals

38. ELECTRONEGATIVITY VALUES, MAIN GROUP ELEMENTS

39. We have classified bonds “ionic” and “covalent”, depending on whether electron pairs are shared or electrons are completely transferred from one atom to another. In actuality, there is no sharp dividing line between the two types but rather a continuum: Evenly shared electrons Unevenly shared electrons Transferred electrons To determine where a bond lies in this “continuum”, it is useful to consider the difference in electronegativity ( X) between the two atoms making up the bond:

40. When the difference (X) is close to zero, sharing is fairly even and electrons are not much closer to one atom than the other. The bonds are considered “non-polar.” When the difference is above zero to about 1.7, the electrons are closer to the more electronegative atom and partial charge buildup, polarization, develops. When a metal or polyatomic cation is present and the (X) is 1.7or higher,ionic bonding becomes the more likely bond type and valence electrons aretransferredto the more electronegative atom.

42. So, we need to consider a third more specialized type of bond, “the polar covalent bond:” This type of bond will be the important factor to be considered when we look at molecular polarity, which arises from molecular shape and bond polarity. The polar molecular in turn will exhibit different solubilities and boiling points than non polar molecules.

43. Let us consider the bond between H and Cl in a molecule of hydrogen chloride (only hydrochloric acid when in water!): Orbital between H and Cl E pair closer to Cl, more electronegative

44. The electron cloud from the pair of shared electrons is more dense closer to the chlorine, and much less dense closer to the hydrogen. The bond has become “polarized”: it has developed a region (or “pole”) of partial positive charge buildup and a region (or “pole”) of partial negative buildup.

45. Major portion of electron density

46. “partially negative” “partially positive” Arrow to indicate polar bond, pointing to more (-) atom

47. The molecule has only one bond, and it is polar. This makes the entire molecule a “dipole”, one which has a positive and negative pole and will align in an electrical or magnetic field: All diatomic molecules with polarized bonding between the two atoms are DIPOLES.

48. Other examples of diatomic dipoles:

49. MOLECULAR POLARITY, LARGER MOLECULES • All diatomic molecules with polar bond(s) are dipoles, • but the situation is not so simple for larger molecules. • There are two factors to consider: • Are the bonds polar? • Are they arranged so that the center of positive • charge and the center of negative charge do not • “coincide”?

50. BOND POLARITY MOLECULAR POLARITY