Ch. 10a: Chemical Bonding II: Molecular Shapes

1. Ch. 10a: Chemical Bonding II: Molecular Shapes Dr. Namphol Sinkaset Chem 200: General Chemistry I

2. I. Chapter Outline • Introduction • Lewis Structures • Resonance • Exceptions • VSEPR Theory • Molecular Polarity

3. I. Importance of Shape • In condensed phases (liquids/solids), molecules are in close proximity, so they interact constantly. • The 3-D shape of a molecule determines many of its physical properties. • We want to be able to predict 3-D shape starting from just a formula of a covalent compound.

4. II. Lewis Structures • The first step to getting the 3-D shape of a molecule is getting the correct 2-D structure. • The 2-D structure will be the basis of our 3-D shape assignment. • A 2-D representation of the bonding in molecule is known as a Lewis structure.

5. II. Steps for Drawing Lewis Structures • Determine total # of valence e-. • Place atom w/ lower Group # (lower electronegativity) as the central atom. • Attach other atoms to central atom with single bonds. • Fill octet of outer atoms. (Why?) • Count # of e- used so far. Place remaining e- on central atom in pairs. • If necessary, form higher order bonds to satisfy octet rule of central atom. • Allow expanded octet for central atoms from Period 3 or lower.

6. II. Lewis Structure Practice • Draw correct Lewis structures for NF3, CO2, SeCl2, PF6-, PI5, IF2-, IF6+, and H2CO.

7. III. Multiple Valid Lewis Structures • Sometimes more than one Lewis structure can be drawn for the same molecule. • For example, ozone (O3).

8. III. Resonance Forms • Resonance forms are also known as resonance structures. • Resonance forms have the same relative placement of atoms, but different locations of bonding and lone e- pairs.

9. III. Resonance Hybrid • Neither resonance form is a true picture of the molecule. • The molecule exists as a resonance hybrid, which is an average of all resonance forms. • In a resonance hybrid, e- are delocalized over the entire molecule.

10. III. Sample Problem • Draw the resonance forms of the carbonate anion.

11. III. Important Resonance Forms • If all resonance forms have the same surrounding atoms, then each contributes equally to the resonance hybrid. • If this is not the case, then one or more resonance forms will dominate the resonance hybrid. • How can we determine which forms will dominate?

12. III. Formal Charge • formal charge: the charge an atom would have if bonding e- were shared equally formal charge = (# valence e-) – (unshared e- + ½ shared e-)

13. III. Formal Charges in O3 • We calculate formal charge for each atom in the molecule. • For oxygen atom A (on the right), there are 6 valence e-, 4 unshared e-, and 4 shared e-. The formal charge for this O atom is 0. • NOTE: sum of all formal charges must equal the overall charge of the molecule!

14. III. Using Formal Charges • Formal charges help us decide the most important resonance forms when we consider to the following guidelines: • Smaller f.c.’s are better than larger f.c.’s. • Same sign f.c.’s on adjacent atoms is undesirable. • Electronegative atoms should carry higher negative f.c.’s.

15. III. Sample Problem • Find the dominant resonance structures for the sulfate anion.

16. IV. Exceptions to the Octet Rule • We’ve already discussed expanded valence cases, but there are other exceptions as well. • e- deficient atoms like Be and B, e.g. BeCl2 and BF3. • Compounds w/ odd # of e-’s: free radicals. Examples include NO and NO2. • Expanded valence – when d orbitals are used to accommodate more than an octet.

17. V. VSEPR Theory • From a correct Lewis structure, we can get to the 3-D shape using this theory. • VSEPR stands for valence shell electron pair repulsion. • The theory is based on the idea that e- pairs want to get as far away from each other as possible!

18. V. VSEPR Categories • There are 5 categories from which all molecular shapes derive.

19. V. Drawing w/ Perspective • We use the conventions below to depict a 3-D object on a 2-D surface.

20. V. Determining 3-D Shape • The 5 categories are a starting point. • To determine the 3-D shape of a molecule, we consider the # of atoms and the # of e- pairs that are associated w/ the central atom. • All the possibilities for molecular geometry can be listed in a classification chart.

21. V. Linear/Trigonal Planar Geometries • First, we have the linear and trigonal planar categories.

22. V. Tetrahedral Geometries

23. V. Trigonal Bipyramidal Geometries

24. V. Octahedral Geometries

25. V. Steps to Determine Molecular Shape • Draw Lewis structure. • Count # of bonds and lone pair e-’s on the central atom. • Select geometric category. • Place e-’s and atoms that lead to most stable arrangement (minimize e- repulsions). • Determine 3-D shape.

26. V. Trig Bipy is Special • In other categories, all positions are equivalent. • In trig bipy, lone pairs always choose to go equatorial first. • Why?

27. V. Distortion of Angles • Lone pair e-’s take up a lot of room, and they distort the optimum angles seen in the geometric categories.

28. V. Some Practice • Draw the molecular shapes for SF4, BeCl2, ClO2-, TeF5-, ClF3, NF3.

29. VI. Molecular Polarity • Individual bonds tend to be polar, but that doesn’t mean that a molecule will be polar overall. • To determine molecular polarity, you need to consider the 3-D shape and see if polarity arrows cancel or not.

30. VI. Sample Problem • Determine the molecular geometry of IF2- and state whether it is polar or nonpolar.