1. Using Bond Energies to Estimate  H° rxn We often use average bond energies to estimate the  H rxn –works best when all reactants and products in gas state Bond breaking is endothermic,  H(breaking) = + Bond making is exothermic,  H(making) = −  H rxn = ∑ (  H(bonds broken)) + ∑ (  H(bonds formed)) 1 Tro: Chemistry: A Molecular Approach, 2/e

2. Example: Estimate the enthalpy of the following reaction Bond breaking 1 mole C─H +414 kJ 1 mole Cl─Cl +243 kJ total +657 kJ Bond making 1 mole C─Cl −339 kJ 1 mole Cl─H −431 kJ total −770 kJ  H rxn = ∑ (  H(bonds broken)) + ∑ (  H(bonds made))  H rxn = (+657 kJ) + (−770 kJ)  H rxn = −113 kJ Tro: Chemistry: A Molecular Approach, 2/e2 BondBond Energy (kJ/mol) C-H414 Cl-Cl243 C-Cl339 H-Cl431

3. Break 1 mol C─H +414 kJ 1 mol Cl─Cl +243 kJ Make 1 mol C─Cl−339 kJ 1 mol H─Cl−431 kJ Tro: Chemistry: A Molecular Approach, 2/e3

4. Practice – Estimate the enthalpy of the following reaction HH + OO HOOH 4

5. Finishing up Covalent Bonding Electronegativity of Atoms; Polarity of Bonds (Section 9.6) –(Note: Polarity of Molecules, Section 10.5, will be covered in Ppt 25 & PS12) Theories of Covalent Bonding –Valence Bond Theory (VBT)_Section 10.6 –Molecular Orbital Theory (Way better [but more complex] model! Not covered in this course [10.8]) VBT and Hybrid Orbitals (Section 10.7)

6. Electronegativity (of Atoms) and Polarity (of Bonds) See separate 1-page handout sheet See Section 9.6, but don’t worry about how electronegativity values were obtained by Pauling or about calculation of “% ionic character”. Be aware that Sec. 9.6 does not as clearly distinguish electronegativity (as a properties of an atom) from polarity (as a property of a bond) as I would like. See Figures 9.8 and 9.9 (in PowerPoint)

7. Electronegativity (EN) A number (no units!) which attempts to reflect how strongly an atom attracts bonding electrons. –Bigger EN = Greater “pulling power” IS A PROPERTY OF ATOMS Trends as IE 1 does (except for noble gases) –Increases to the right (b/c Z eff increases) –Decreases as you go down a column (r increases; farther away means less pull) –F has greatest EN (4.0)

8. Figure 9.8 Pauling Electronegativities

9. Polarity The degree to which a covalent bond (or later, a whole molecule) has two differently charged ends. –Greater polarity = Greater charge difference IS A PROPERTY OF BONDS (later, molecules) Results from a difference in EN values of two bonded atoms. –“unequal sharing” No polarity = “nonpolar” (equal sharing)

10. Table 9.1,Fig. 9.9 Relationship Between Electro- negativity DIFFERENCE and Bond Type (i.e., Polarity) Nonpolar

11. Bond Polarity EN Cl = 3.0 3.0 − 3.0 = 0 Pure Covalent EN Cl = 3.0 EN H = 2.1 3.0 – 2.1 = 0.9 Polar Covalent EN Cl = 3.0 EN Na = 0.9 3.0 – 0.9 = 2.1 Ionic Tro: Chemistry: A Molecular Approach, 2/e

12. Valence Bond Theory Assumes that EACH ATOM making a bond has (uses) one of its ATOMIC orbitals to make the bond Each of these atomic orbitals contains one electron (“singly occupied orbital”) Bonding occurs only if the two orbitals “overlap” each other in the space between the atoms (better “orbital overlap” = stronger covalent bond)

13. Two Basic Kinds of Bonds in VBT (distinguished by the type of overlap)  (sigma) bond –Overlap is along the bond axis (i.e., orbitals generally point “at” one another [except for s] )  Rotation around the bond does not change overlap  (pi) bond –Overlap is above and below the bond axis (i.e., “sideways” overlap)  Rotation around the bond does change overlap

14. See Board for O 2 and N 2

15. (Without hybridization) (With sp 3 hybridization) Why hybrid orbitals were conceived

16. Why Hybrid Orbitals? (“pure” [unhybridized] vs. hybrid orbitals) s, p, d, (and f) orbitals are “pure” atomic orbitals. –If these were the only orbitals used for bonding, only 180  and 90  bond angles could be explained! (see board example) p orbitals are perpendicular to one another Concept of “hybrid orbitals” was conceived –From mathematical combination of 2 or more “pure” orbitals on same atom –Have names that “show” the number and kinds of pure orbitals used to create them sp, sp 2, sp 3, sp 3 d, etc. –Explains the angles predicted by VSEPR (and observed)… 120  and 109 

17. “Conservation of orbitals” Idea, & Shapes of hybrid orbitals When you hybridize atomic orbitals, you always get the SAME NUMBER of orbitals that you started out with. –two pure atomic orbitals => two hybrid orbitals (in a “set”) –three pure orbitals => a set of three hybrid orbitals Every hybrid orbital has two “lobes”, but one is bigger than the other, so the smaller one is typically ignored (not drawn)

18. Shapes (Continued) Orbitals in a “set” of hybrids will all have the same shape, but point in different directions –similar idea to that of a set of three p orbitals: p x, p y, p z

19. Tro, Figure 10.7 s + p + p + p yields four sp 3 orbitals

20. s + p = two sp hybrid orbitals From McMurry text

21. s + p + p = three sp 2 hybrid orbitals (like Fig. 10.8 in Tro)

22. Copyright © Houghton Mifflin Company. All rights reserved. 9–22 Figure 9.10, Zumdahl When An s and Two p Orbitals Are Hybridized to Form a Set of Three sp 2 Orbitals, One p Orbital Remains Unchanged and is Perpendicular to the Plane of the Hybrid Orbitals

23. 10-minute YouTube Video http://www.youtube.com/watch?v=d1E18tBTlBg “molecular shape and orbital hybridization” This will help you visualize what is happening better than I can show on the board or in a static picture (as in PowerPoint). It will not help you determine which hybrid orbitals are used by a given atom—that is most easily done using an LDS [as I will describe]

24. Consistent with VSEPR Predictions It turns out that the “directions” in which the hybrid orbitals in a set “point” match precisely with the VSEPR geometries!!!

26. Final Generalizations (“It turns out that….”) The # hybrid orbitals needed always = the # of e - - clouds (from LDS/VSEPR model) Hybrid orbitals are ONLY used for  -bonds (or lone pairs) The only way to get  bonds for simple molecules (i.e., in this class) is to have/use “leftover” (pure) p orbitals –i.e., it only happens when sp or sp 2 hybrids are used for the  -bonding Any single bond is a  bond; any bond AFTER a single bond is a  bond –A double bond is always 1  and ond  –A triple bond is always 1  and two 

27. Example on Board H 2 CO

28. Strategy / Approach

30. Tro notation: hybridization and bonding scheme O C H H Wedge and dash depiction

31. Copyright © Houghton Mifflin Company. All rights reserved. 9–31 Figure 9.11 The Sigma Bonds in Ethylene

32. Copyright © Houghton Mifflin Company. All rights reserved. 9–32 Figure 9.12 A Carbon-Carbon Double Bond Consists of a  and a  Bond

33. Copyright © Houghton Mifflin Company. All rights reserved. 9–33 Figure 9.13 (Zumdahl) (a)The Orbitals Used to Form the Bonds in Ethylene (b) The Lewis Structure for Ethylene

35. Tro shows dichloroethene (bonding analogous to ethene except for Cl’s)

36. sp hybrids used for either two double bonds or a triple bond

37. Copyright © Houghton Mifflin Company. All rights reserved. 9–37 Figure 9.17 The Orbitals of an sp Hybridized Carbon Atom

52. Copyright © Houghton Mifflin Company. All rights reserved. 9–52 Figure 9.21 A Set of dsp 3 Hybrid Orbitals on Phosphorus Atom

53. Copyright © Houghton Mifflin Company. All rights reserved. 9–53 Figure 9.22b The Orbitals Used to Form the Bonds in PCl 5

54. Copyright © Houghton Mifflin Company. All rights reserved. 9–54 Figure 9.23 An Octahedral Set of d 2 sp 3 Orbitals on Sulfur Atom

55. Copyright © Houghton Mifflin Company. All rights reserved. 9–55 Xenon