1. Resonance Structures For some molecules (and polyatomic ions) more than one Lewis structure with “identical” patterns of formal charges appears to be plausible. In such cases we often view the actual molecule as a resonance hybrid of the plausible structures. The examples of ozone and sulfur dioxide may be familiar from high school chemistry.

2. Resonance Structures – Bond Orders: For O 3 and SO 2, the two resonance structures suggest a bond order of 1.5 for each molecule. The bond order calculation takes the number of bonds between equivalent atom pairs and divides by the number of bonded atom pairs. In examples such as CO 3 2- there are three equivalent atom pairs (and even more for SO 4 2- ion).

3. Resonance Structures – Symmetry : Molecules whose structure is best represented as an “average” of several contributing resonance structures often are more symmetrical than an individual Lewis structure might suggest. In ozone, for example, both oxygen-oxygen bond lengths are identical. One “side” of the molecule is a mirror image of the other side.

4. Resonance Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10Slide 4 of 48 O O O O O O + + - - Electrostatic potential map of ozone O O O + -½

5. Class Examples 1. Draw a Lewis structure for the BF 3 molecule. Using all single bonds the molecule appears to be “electron deficient” (less than 8 electrons around B. Can a structure or structures be drawn for which the octet rule is satisfied? 2. Draw one or more Lewis structures for the carbonate ion (CO 3 2- ) for which the octet rule is obeyed. Is the concept of resonance useful?

6. Incomplete Octets Copyright © 2011 Pearson Canada Inc. Slide 6 of 48General Chemistry: Chapter 10 B F FF B F FF - + B F FF - +

7. Lewis Diagrams – New Features: In the Lewis structures we’ve drawn so far the octet rule has been satisfied for atoms other than hydrogen (H). In some molecules such as BeF 2 and BF 3 the central atom (Be and B) have fewer than 8 valence shell electrons for plausible Lewis structures. Such molecules are sometimes called electron deficient.

8. Class Examples 1. Draw Lewis structures for the BeF 2 molecule in which (a) the BeF 2 molecule has two single bonds and (b) the BeF 2 molecule has two double bonds. 2. Which of the two structures satisfies the octet rule for both Be and F? 3. Calculate formal charge values for atoms in both the single bonded and double bonded structure. Do formal charge values favour one structure over the other?

9. Lewis Diagrams – New Features: A significant number of Binary nonmetallic compounds have molecules where the central atom forms more than four covalent bonds. Clearly, in such cases the octet rule is also not obeyed. Molecules of such compounds are sometimes called hypervalent. The terms expanded octet and expanded valence shell are also seen.

10. “Expanded Octets” The next two slides show structures for a molecule containing P and an ion containing S where both P and S have more than the expected octet of “valence shell” electrons. What do both of these structures have in common? An extreme cases of hypervalent molecules is IF 7 (14 valence shell electrons).

11. Expanded octets Copyright © 2011 Pearson Canada Inc. Slide 11 of 48General Chemistry: Chapter 10 P Cl P Cl Cl Cl Cl S F F F F F F

12. Expanded Valence Shells Copyright © 2011 Pearson Canada Inc. Slide 12 of 48General Chemistry: Chapter 10

13. Molecular Geometries: Simple chemical bonding concepts allow us to predict the shapes of most molecules. Finer details, such as lengths of individual bonds “must” be found by experiment. The shape of a molecule usually determines whether a molecule is electrically polar (i.e. has a permanent electric dipole moment). Electrical polarity of individual molecules largely determines the bulk physical properties of compounds (groups of molecules).

14. Molecular Geometries In living systems the geometry of a particular small molecule often determines whether the molecule will interact with an active site in an enzyme. Sizes and shapes of molecules are important throughout chemistry, biochemistry, biology and engineering. The simplest molecules (diatomics) are all linear (neglecting the electrons!). However, such molecules have different sizes (bond lengths).

15. Diatomic Molecules – Bond Lengths Molecular Formula Bond Length (pm) H 2 74.1 HCl 127.4 HBr 141.5 CO 112.8 CS 153.5 RbBr* 294.5 AgCl* 228.1 PbS* 228.7 *These are all gas phase “covalently bonded” molecules and not ionic crystals.

16. Nonlinear Molecules – Geometries: For nonlinear molecules we must know both bond lengths and bond angles to describe a molecule’s geometry. In many cases we will not be concerned with bond lengths as the molecular shape will determine both a molecule’s polarity and physical properties. The structure of the fundamental biological molecule, water, is shown on the next slide. The bond angle value is about 106 degrees.

17. The Shapes of Molecules Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10 Slide 17 of 48 FIGURE 10-9 Geometric shape of a molecule

18. Valence Shell Electron Pair Repulsion Model – VSEPR: Very often the shape of a molecule can be predicted if a satisfactory Lewis structure for the molecule is first drawn. VSEPR theory assumes that the pairs of electrons around a “central” atom can arrange themselves in space to minimize coulombic forces between pairs of electrons. Both bonding and lone pairs of electrons must be considered. Symmetric molecular structures are often seen.

19. Balloon analogy to valence-shell electron-pair repulsion FIGURE 10-10 Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10Slide 19 of 48

20. Methane and Related Molecules In the diagram on the previous page each balloon represents a pair of electrons. In methane, CH 4, and related molecules (SiH 4, GeH 4, SnH 4 each of the electron pairs is a bonding pair. The four bonding pairs (or bonds) point toward the corners of a regular tetrahedron. In molecules such as ammonia, NH 3, four electron pairs point towards the corners of a very slightly distorted tetrahedron.

21. Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10Slide 21 of 48 Methane, Ammonia and Water

22. Several electron-group geometries illustrated FIGURE 10-12 Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10 Slide 22 of 48

23. Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10 Slide 23 of 48 Bond length – distance between nuclei. Bond angle – angle between adjacent bonds. VSEPR Theory – Electron pairs repel each other whether they are in chemical bonds (bond pairs) or unshared (lone pairs). Electron pairs assume orientations about an atom to minimize repulsions. Electron group geometry – distribution of e - pairs. Molecular geometry – distribution of nuclei.

24. Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10Slide 24 of 48 Molecular Geometry as a Function of Electron Group Geometry

25. Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10Slide 25 of 48

26. Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10Slide 26 of 48

27. Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10Slide 27 of 48

28. Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10Slide 28 of 48

29. Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10Slide 29 of 48 Replace this part of table 10.1 this format is not present in the text. Parent table is on page 425 and 426

30. Applying VSEPR Theory Copyright © 2011 Pearson Canada Inc. Slide 30 of 48 General Chemistry: Chapter 10 1.Draw a plausible Lewis structure. 2.Determine the number of e - groups and identify them as bond or lone pairs. 3.Establish the e - group geometry. 4.Determine the molecular geometry.

31. Class Example 1. Draw complete Lewis structures for the PF 3, SeF 6 and IF 5 molecules. Using the Lewis structures as a starting point, employ VSEPR theory in order to determine/predict the three dimensional shapes of these three molecules.

32. Polar Molecules Using the shape as a starting point we can determine whether a molecule is electrically polar. Recall that homonuclear diatomics are nonpolar and heteronuclear diatomics are polar (due to unequal sharing of electrons). Net electrical polarity is reported as an electric dipole moment. Molecules with a permanent dipole moment will orient in the presence of an applied electric field.

33. Molecular Shapes and Dipole Moments Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10 Slide 33 of 48 FIGURE 10-14 Polar molecules in an electric field

34. Polar Molecules – cont’d: Molecules can have a net molecular dipole moment equal to zero even if, in individual bonds, the bonding electrons are not equally shared (bond dipoles cancel – vector addition!). This occurs especially with highly symmetric molecules. Care is required. Thus, carbon dioxide O=C=O is nonpolar (bond dipoles of equal size pointing in opposite directions) while O=C=S is polar. Why?

35. Net Molecular Polarity The polarity of a molecule can be considered using electrostatic potential maps or vector addition of individual bond dipoles. The latter approach is less visually pleasing but is much easier to employ. The next slide contrasts the two approaches.

36. Molecular shapes and dipole Moments FIGURE 10-15 Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 10 Slide 36 of 48

37. Methane and Chloromethanes: Molecule Polar Molecule? CH 4 No CH 3 Cl Yes CH 2 Cl 2 Yes CHCl 3 Yes CCl 4 No

38. Study Slides? The next two slides are meant to help you in your home study. The first slide shows the number of valence electrons possessed by neutral nonmetal atoms. The second slide shows the usual valences of nonmetals. Higher valences are usually observed when a “central atom” is bonded to highly electronegative peripheral atoms (F and O especially!).

39. Nonmetals – Valence Electrons by Group Number Modern Group Number Historical Group Number Element List Valence Electrons per Neutral Atom 14 4AC, Si, Ge, Sn * 4 15 5AN, P, As, Sb 5 16 6AO, S, Se, Te 6 17 7AF, Cl, Br, I 7 * Tin (Sn) is a metal but will form covalent bonds.

40. Nonmetals – Common Valences by Group Number Modern Group Number Historical Group Number Element List Common Valences (Number of Bonds Formed) 14 4AC, Si, Ge, Sn * 4 ** 15 5AN, P, As, Sb 3, 5 16 6AO, S, Se, Te 2, 4, 6 17 7AF, Cl, Br, I 1, 3, 5, 7 * Tin (Sn) is a metal but will form covalent bonds. ** Both single and multiple bonds can be seen together.

41. Study Slides – Footnotes: The usual valences listed on the previous slides will not describe the number of bonds formed in free radicals (unpaired electron species). They also will not “work” in the case of basic molecules that have been protonated. Thus, in the hydronium ion (H 3 O + ), oxygen forms three bonds. In the ammonium ion (NH 4 + ) nitrogen forms four bonds (is tetravalent). Mention coordinate covalent bonds here?

42. Class Examples 1. Draw a Lewis structure for the NH 4 + ion. Is this ion polar? What is its three dimensional shape? Is there a second period hydride with a similar structure? Write a balanced chemical reaction with the reactants being ammonia and H + and the product being NH 4 +. Show the Lewis structure for both reactants and the product. What is “unusual” about the formation of the bond between the proton and ammonia?

43. Class Examples 2. Draw a Lewis structure for the H 3 O + ion. Is this ion polar? What is its three dimensional shape? Is there a second period hydride with a similar structure? Write a balanced chemical reaction with the reactants being water and H + and the product being H 3 O +. Show the Lewis structure for both reactants and the product. What is “unusual” about the formation of the bond between the proton and water?

44. Class Example 1. Draw a Lewis structure or structures for the NO 2 F molecule. Predict the shape of the molecule. Is the molecule polar? What is the nitrogen-oxygen bond order? Is the resonance concept useful here?