Theories of Covalent Bonding

1. Theories of Covalent Bonding 11.1 Valence Shell Electron Pair Repulsion Theory 11.2 Valence Bond (VB) Theory and Orbital Hybridization 11.3 Molecular Orbital (MO)Theory and Electron Delocalization

2. Chapter 9 Covalent Bonding and Orbital Overlap • Lewis structures and VSEPR do not explain why a bond forms. How do we account for shape in terms of quantum mechanics? • What are the orbitals that are involved in bonding? • We use Valence Bond Theory: • Bonds form when orbitals on atoms overlap. • A covalent bond forms when the orbitals of two atoms overlap and the overlap region, which is between the nuclei, is occupied by a pair of electrons. • There are two electrons of opposite spin in the orbital overlap.

3. Chapter 9 Covalent Bonding and Orbital Overlap Hydrogen, H2

4. Hydrogen fluoride, HF Fluorine, F2

5. Chapter 9 Covalent Bonding and Orbital Overlap • As two nuclei approach each other their atomic orbitals overlap. • As the amount of overlap increases, the energy of the interaction decreases. • At some distance the minimum energy is reached. • The minimum energy corresponds to the bonding distance (or bond length). • As the two atoms get closer, their nuclei begin to repel and the energy increases.

6. Chapter 9 Covalent Bonding and Orbital Overlap • At the bonding distance, the attractive forces between nuclei and electrons just balance the repulsive forces (nucleus-nucleus, electron-electron).

8. Chapter 9 Hybrid Orbitals • Atomic orbitals can mix or hybridize in order to adopt an appropriate geometry for bonding. • Hybridization is determined by the electron domain geometry. sp Hybrid Orbitals • Consider the BeF2 molecule (experimentally known to exist):

9. Chapter 9 Hybrid Orbitals sp Hybrid Orbitals • Be has a 1s22s2 electron configuration. • There is no unpaired electron available for bonding. • We conclude that the atomic orbitals are not adequate to describe orbitals in molecules. • We know that the F-Be-F bond angle is 180 (VSEPR theory). • We also know that one electron from Be is shared with each one of the unpaired electrons from F.

10. Chapter 9 Hybrid Orbitals sp Hybrid Orbitals • We assume that the Be orbitals in the Be-F bond are 180 apart. • We could promote and electron from the 2s orbital on Be to the 2p orbital to get two unpaired electrons for bonding. • BUT the geometry is still not explained. • We can solve the problem by allowing the 2s and one 2p orbital on Be to mix or form a hybrid orbital. • The hybrid orbital comes from an s and a p orbital and is called an sp hybrid orbital.

12. Chapter 9 Hybrid Orbitals sp Hybrid Orbitals • The lobes of sp hybrid orbitals are 180º apart. • Since only one of the Be 2p orbitals has been used in hybridization, there are two unhybridized p orbitals remaining on Be.

13. The sp hybrid orbitals in gaseous BeCl2. Figure 11.2 atomic orbitals hybrid orbitals orbital box diagrams

14. The sp hybrid orbitals in gaseous BeCl2(continued). Figure 11.2 orbital box diagrams with orbital contours

16. The sp2 hybrid orbitals in BF3. Figure 11.3

17. sp2 and sp3 Hybrid Orbitals

18. The sp3 hybrid orbitals in CH4. Figure 11.4

19. Figure 11.5 The sp3 hybrid orbitals in NH3.

20. Figure 11.5 continued The sp3 hybrid orbitals in H2O.

21. Figure 11.6 The sp3d hybrid orbitals in PCl5.

22. The sp3d2hybrid orbitals in SF6. Figure 11.7

23. Key Points Types of Hybrid Orbitals sp sp2 sp3 sp3d sp3d2 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.

25. Step 1 Step 2 Step 3 Figure 11.8 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

26. PROBLEM: Use partial orbital diagrams to describe mixing of the atomic orbitals of the central atom 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.

27. 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.

28. Chapter 9 Hybrid Orbitals Hybridization Involving d Orbitals • Since there are only three p-orbitals, trigonal bipyramidal and octahedral electron domain geometries must involve d-orbitals. • Trigonal bipyramidal electron domain geometries require sp3d hybridization. • Octahedral electron domain geometries require sp3d2 hybridization. • Note the electron domain geometry from VSEPR theory determines the hybridization.

29. Chapter 9 Hybrid Orbitals Summary Draw the Lewis structure. Determine the electron domain geometry with VSEPR. Specify the hybrid orbitals required for the electron pairs based on the electron domain geometry.

32. Chapter 9 Multiple Bonds • -Bonds: electron density lies on the axis between the nuclei. • All single bonds are -bonds. • -Bonds: electron density lies above and below the plane of the nuclei. • A double bond consists of one -bond and one -bond. • A triple bond has one -bond and two -bonds. • Often, the p-orbitals involved in -bonding come from unhybridized orbitals.

33. Multiple Bonds

34. both C are sp3 hybridized s-sp3 overlaps to  bonds sp3-sp3 overlap to form a  bond relatively even distribution of electron density over all  bonds The  bonds in ethane(C2H6). Figure 11.9

35. Chapter 9 Multiple Bonds • Ethylene, C2H4, has: • one - and one -bond; • both C atoms sp2 hybridized; • both C atoms with trigonal planar electron pair and molecular geometries.

36. overlap in one position -  p overlap -  electron density The  and  bonds in ethylene (C2H4). Figure 11.10

37. Chapter 9 Multiple Bonds

39. Chapter 9 Multiple Bonds • When triple bonds form (e.g. N2) one -bond is always above and below and the other is in front and behind the plane of the nuclei.

40. Chapter 9 Multiple Bonds • Consider acetylene, C2H2 • the electron pair geometry of each C is linear; • therefore, the C atoms are sp hybridized; • the sp hybrid orbitals form the C-C and C-H -bonds; • there are two unhybridized p-orbitals; • both unhybridized p-orbitals form the two -bonds; • one -bond is above and below the plane of the nuclei; • one -bond is in front and behind the plane of the nuclei.

41. overlap in one position -  p overlap -  The  and  bonds in acetylene (C2H2). Figure 11.11

42. Chapter 9 Multiple Bonds

43. Chapter 9 Multiple Bonds

44. Chapter 9 Multiple Bonds Delocalized  Bonding • So far all the bonds we have encountered are localized between two nuclei. • In the case of benzene • there are 6 C-C  bonds, 6 C-H  bonds, • each C atom is sp2 hybridized, • and there are 6 unhybridized p orbitals on each C atom.

45. Chapter 9 Multiple Bonds Delocalized  Bonding

46. Chapter 9 Multiple Bonds Delocalized  Bonding • In benzene there are two options for the 3  bonds • localized between C atoms or • delocalized over the entire ring (i.e. the  electrons are shared by all 6 C atoms). • Experimentally, all C-C bonds are the same length in benzene. • Therefore, all C-C bonds are of the same type (recall single bonds are longer than double bonds).

47. Chapter 9 Multiple Bonds General Conclusions • Every two atoms share at least 2 electrons. • Two electrons between atoms on the same axis as the nuclei are  bonds. • -Bonds are always localized. • If two atoms share more than one pair of electrons, the second and third pair form -bonds. • When resonance structures are possible, delocalization is also possible.

48. Chapter 9 Molecular Orbitals • Some aspects of bonding are not explained by Lewis structures, VSEPR theory and hybridization. (E.g. why does O2 interact with a magnetic field?; Why are some molecules colored?)‏ • For these molecules, we use Molecular Orbital (MO) Theory. • Just as electrons in atoms are found in atomic orbitals, electrons in molecules are found in molecular orbitals.

49. 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 Bond in Molecules SOLUTION: bond bonds

50. 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).