Chapter 11 Theories of Covalent Bonding

1. Chapter 11 Theories of Covalent Bonding

2. Theories of Covalent Bonding 11.1 Valence Bond (VB) Theory and Orbital Hybridization 11.2 The Mode of Orbital Overlap and the Types of Covalent Bonds 11.3 Molecular Orbital (MO)Theory and Electron Delocalization

3. The Central Themes of VB Theory Basic Principle A covalent bond forms when the orbtials of two atoms overlap and are occupied by a pair of electrons that have the highest probability of being located between the nuclei. Themes A set of overlapping orbitals has a maximum of two electrons that must have opposite spins. The greater the orbital overlap, the stronger (more stable) the bond. The valence atomic orbitals in a molecule are different from those in isolated atoms.

4. Hydrogen, H2 Hydrogen fluoride, HF Fluorine, F2 Atomic Orbital Overlap Orbital overlap and spin pairing in diatomic molecules

5. Key Points Types of Hybrid Orbitals Hybrid Orbitals The number of hybrid orbitals obtained equals the number of atomic orbitals mixed. The type of hybrid orbitals obtained varies with the types of atomic orbitals mixed. sp sp2 sp3 sp3d sp3d2

6. The sp hybrid orbitalsin gaseous BeCl2 atomic orbitals hybrid orbitals orbital box diagrams

7. The sp hybrid orbitalsin gaseous BeCl2 (continued) orbital box diagrams with orbital contours

8. The sp2 hybrid orbitals in BF3

9. The sp3 hybrid orbitals in CH4

10. The sp3 hybrid orbitals in NH3

11. The sp3 hybrid orbitals in H2O

12. The sp3d hybrid orbitals in PCl5

13. The sp3d2hybrid orbitals in SF6

15. Step 1 Step 2 Step 3 Figure 10.1 Figure 10.12 Table 11.1 The conceptual steps from molecular formula to the hybrid orbitals used in bonding. Molecular shape and e- group arrangement Molecular formula Lewis structure Hybrid orbitals

16. PROBLEM: Use partial orbital diagrams to describe mixing of atomic orbitals on the central atoms leads to hybrid orbitals in each of the following: PLAN: Use the Lewis structures to ascertain the arrangement of groups and shape of each molecule. Postulate the hybrid orbitals. Use partial orbital box diagrams to indicate the hybrid for the central atoms. SAMPLE PROBLEM 11.1 Postulating Hybrid Orbitals in a Molecule (a) Methanol, CH3OH (b) Sulfur tetrafluoride, SF4 SOLUTION: (a) CH3OH The groups around C are arranged as a tetrahedron. O also has a tetrahedral arrangement with 2 nonbonding e- pairs.

17. hybridized C atom hybridized O atom single C atom single O atom hybridized S atom S atom SAMPLE PROBLEM 11.1 Postulating Hybrid Orbitals in a Molecule continued (b) SF4 has a seesaw shape with 4 bonding and 1 nonbonding e- pairs.

18. Types of Covalent Bonds Sigma () Bonds - Bonding that results from the end-to-end overlap is called a sigma bond. It has the highest electron density along the axis between the two nuclei. Single bonds are sigma bonds. Pi () Bonds - Bonds that result from the side-to-side overlap of unhybridized p orbitals. The electron density is above and below the axis between the two nuclei.(This is why multiple bonds counted as one group of electrons in VSEPR theory) The multiple part of multiple bonds are  bonds.In a double bond, there is one  and one  bond. In a triple bond, there is one  and two  bonds.

19. both C are sp3 hybridized s-sp3 overlaps to s bonds sp3-sp3 overlap to form a s bond relatively even distribution of electron density over all s bonds The s bonds in ethane.

20. overlap in one position - s p overlap -  electron density The s and p bonds in ethylene (C2H4)

21. overlap in one position - s p overlap -  The s and p bonds in acetylene (C2H2)

22. PROBLEM: Describe the types of bonds and orbitals in acetone, (CH3)2CO. PLAN: Use the Lewis structures to ascertain the arrangement of groups and shape at each central atom. Postulate the hybrid orbitals taking note of the multiple bonds and their orbital overlaps. sp3 hybridized sp3 hybridized sp2 hybridized SAMPLE PROBLEM 11.2 Describing the Bonding in Molecules with Multiple Bonds SOLUTION: bond bonds

23. CIS TRANS Restricted rotation of p-bonded molecules Rotation about the C-C bond can’t take place without breaking the  electron overlap.

24. The Central Themes of MO Theory A molecule is viewed on a quantum mechanical level as a collection of nuclei surrounded by delocalized molecular orbitals. Atomic wave functions are summed to obtain molecular wave functions. If wave functions reinforce each other, a bonding MO is formed (region of high electron density exists between the nuclei). If wave functions cancel each other, an antibonding MO (*) is formed (a node of zero electron density occurs between the nuclei).

25. Amplitudes of wave functions added Amplitudes of wave functions subtracted. An analogy between light waves and atomic wave functions.

26. Contours and energies of the bonding and antibonding molecular orbitals (MOs) in H2.

27. The MO diagram for H2

28. s*1s Energy 1s 1s 1s 1s s1s AO of He AO of He+ AO of He AO of He MO diagram for He2+ and He2 s*1s Energy s1s MO of He+ MO of He2 He2 bond order = 0 He2+ bond order = 1/2(exists)

29. PROBLEM: Use MO diagrams to predict whether H2+ and H2- exist. Determine their bond orders and electron configurations. PLAN: Use H2 as a model and accommodate the number of electrons in bonding and antibonding orbitals. Find the bond order. s s 1s 1s 1s 1s AO of H- AO of H AO of H AO of H s s SAMPLE PROBLEM 11.3 Predicting Species Stability Using MO Diagrams SOLUTION: bond order = 1/2(1-0) = 1/2 bond order = 1/2(2-1) = 1/2 H2+ does exist H2- does exist configuration is (s1s)2(s2s)1 MO of H2- MO of H2+ configuration is (s1s)1

30. s*2s s*2s 2s 2s 2s 2s s2s s2s s*1s s*1s 1s 1s 1s 1s s1s s1s Bonding in s-block homonuclear diatomic molecules. Be2 Li2 Energy Li2 bond order = 1(is observed) Be2 bond order = 0(not observed)

31. Contours and energies of s and p MOs through combinations of 2p atomic orbitals

32. Relative MO energy levels for Period 2 homonuclear diatomic molecules. without 2s-2p mixing with 2s-2p mixing MO energy levels for O2, F2, and Ne2 MO energy levels for B2, C2, and N2

33. MO occupancy and molecular properties for B2 through Ne2

34. PROBLEM: As the following data show, removing an electron from N2 forms an ion with a weaker, longer bond than in the parent molecules, whereas the ion formed from O2 has a stronger, shorter bond: PLAN: Find the number of valence electrons for each species, draw the MO diagrams, calculate bond orders, and then compare the results. N2 N2+ O2 O2+ Bond energy (kJ/mol) 945 841 498 623 Bond length (pm) 110 112 121 112 SAMPLE PROBLEM 11.4 Using MO Theory to Explain Bond Properties Explain these facts with diagrams that show the sequence and occupancy of MOs. SOLUTION: N2 has 10 valence electrons, so N2+ has 9. O2 has 12 valence electrons, so O2+ has 11.

35. 2p 2p 2p 2p s2s s2s SAMPLE PROBLEM 11.4 Using MO Theory to Explain Bond Properties continued N2 N2+ O2 O2+ 2p antibonding e- lost bonding e- lost 2p 2p 2p s2s s2s bond orders 1/2(8-2)=3 1/2(7-2)=2.5 1/2(8-4)=2 1/2(8-3)=2.5

36. The lowest energy p-bonding MOs in benzene and ozone.

37. s 1s 2px 2py 2p s AO of H AO of F The MO diagram for HF Energy MO of HF

38. s*s *p 2p 2p sp p s*s 2s 2s AO of N AO of O ss The MO diagram for NO Energy possible Lewis structures MO of NO

39. End of Chapter 11

40. Figure 10.1 The steps in converting a molecular formula into a Lewis structure. Place atom with lowest EN in center Molecular formula Step 1 Atom placement Add A-group numbers Step 2 Sum of valence e- Draw single bonds. Subtract 2e- for each bond. Step 3 Give each atom 8e- (2e- for H) Remaining valence e- Step 4 Lewis structure

41. Figure 10.12 The steps in determining a molecular shape. See Figure 10.1 Molecular formula Step 1 Lewis structure Count all e- groups around central atom (A) Step 2 Electron-group arrangement Note lone pairs and double bonds Step 3 Count bonding and nonbonding e- groups separately. Bond angles Step 4 Molecular shape (AXmEn)