1. Lecture 23 © slg CHM 151 TOPICS : 1. Molecular Shapes 2. Polyatomic ion Shapes 3. Introduction to Bond Polarity

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”: “three regions”: “two regions”:

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

6. Four Regions: Five 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 halogens do 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: BeCl 2, CO 2, NO 2 +, HCN

13. Type Two: Three Regions NO 3 -, NO 2 -, CH 2 O

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

18. Type Three: Four Regions CH 2 Cl 2, NH 3, H 2 O, NH 4 +

20. Note: molecular geometry, trigonal pyramid

21. 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 107 o.

22. Note: molecular geometry: BENT

24. GROUP WORK: Do Lewis structure and assign shape and bond angles: CO 3 2-, SiCl 4

27. Type Four: Five Regions PF 5, ClF 3, IF 2 -, SF 4

28. Bond angles in triangle, 120 o Bond angles, each “axial” F, 90 o from trigonal plane

29. Note: “T-shaped”; unshared pairs always trigonal planar

30. Note: “linear” molecular geometry

31. Note: “Seesaw” molecular geometry

32. Type Five: Six Regions SF 6, IF 5, XeF 4

35. Note: molecular geometry “square pyramidal”

36. Note: “Square planar”

37. GROUP WORK: Do Lewis structure and assign shape and bond angles: ICl 4 +, XeOF 2, ICl 4 -

41. 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”: Unit 5, Lecture 24, next week!!

42. ELECTRONEGATIVITY The trends in ionization energies and electron affinities can be thought of as summarized in a single property called “electronegativity” (en or X). Electronegativity is 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 the largest values, metals the lowest. Noble gases have no assigned X value.

43. Most active metals Most active non-metals

44. The electronegativity values are quite useful in evaluating bond type and what we will term “bond polarity,” which arises when electrons are shared unevenly. In summation: Metals: larger size, lower ionization potential, lower electron affinity, and lower electronegativity; tend to form positive ions Non-Metals: smaller size, higher ionization potential, higher electron affinity, higher electronegativity; function as anions in ionic compounds.

45. 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 between the two atoms making up the bond:

46. When the difference is less than 0.5, sharing is fairly even and electrons are not much closer to one atom than the other. When the difference is between 0.5 and about 1.5, the electrons are closer to the more electronegative atom and partial charge buildup, polarization, develops. When electronegativity difference is greater than 1.5 or so, ionic bonding becomes the more likely type and valence electrons are transferred to the more electronegative atom.

47. 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. The relationship between these is what we will next examine.