Chapter 9

1. Chapter 9 Covalent Bonding: orbitals

2. Topics • Hybridization and the localized electron model • The molecular orbital model • Bonding in homonuclear diatomic molecules • Bonding in heteronuclear diatomic molecules • Combining the localized electron and molecular orbital models

3. 9.1 Hybridization and localized electron model How do atoms share electrons between their valence shells? The localized electron bonding model • A covalent bond is formed by the • pairing of two electrons with opposing • spins in the region of overlap of atomic • orbitals between two atoms • Overlap: the two orbitals share a • common region in space • This overlap region has high • electron charge density • The more extensive the overlap • between two orbitals, the stronger • is the bond between two atoms

4. According to the model: • For an atom to form a covalent bond it must have an unpaired electron • Number of bonds formed by an atom should be determined by its number of unpaired electrons

5. Overlap Of 2 1s 2 2p How does Lewis theory explain the bonds in H2 and F2? Sharing of two electrons between the two atoms. H2 (1s1) (1s1) F2 (1s22s22p5) (1s22s22p5) Localized electron model – bonds are formed by sharing of e- from overlapping atomic orbitals.

6. If a 2s electron is promoted to an empty 2p orbital, then four unpaired electrons can give rise to four bonds These four orbitals become mixed, or hybridized to form bonds Hybridization of Atomic Orbitals Based on ground-state electron configuration, carbon should have only two bonds

7. Hybridization of Atomic Orbitals Most of the electrons in a molecule remain in the same orbital locations that they occupied in the separated atoms Bonding electrons are localized in the region of atomic orbital overlap

8. Hybridization ? • Two or more atomic orbitals are mixed to produce a new set of orbitals (blended orbitals) • Number of hybrid orbitals = number of atomic orbitals mixed

9. sp3 Hybridization Occurs most often for central atom only The total number of hybrid orbitals is equal to the number of atomic orbitals combined

10. sp3 Hybridization • s and p orbitals of the valence electrons are blended. • one s orbital is combined with 3 p orbitals. • sp3hybridization has tetrahedral geometry.

11. sp3 Hybridization The carbon atom in methane (CH4) has bonds that are sp3 hybrids Note that in this molecule carbon has all single bonds

12. sp3 In terms of energy 2p Hybridization Energy 2s

13. H C H H H Methane building blocks C H H H C H H VSEPR

14. 3 sp 1s 2s 2px 2py 2pz sp3 sp3 sp3 sp3 y Hybridize Promote x 109.5o z Methane: Carbon

15. 2p 2 sp3 2s 1s sp3 Orbital Hybridization in NH3 3 Equivalent half-filled orbitals are used to form bonds with 3H atoms. The 4thsp3holds the lone pair 7N

16. Bonding in Ammonia Ammonia (NH3) is similar to CH4 except the lone pair of electrons occupies the 4th hybrid orbital 7N: 1s2 2s2 2p3

17. O H H H H 2 sp3 2p 8O 2s How about hybridization in H2O?

18. sp2 Hybridization Consider BF3 5B The empty 2p orbital remains unhybridized sp2 is comprised of one 2s orbital and two 2p orbitals to produce a set of threesp2 hybrid orbitals

19. Formation of sp2 Hybrid Orbitals

20. Hybrid orbitals and geometry The geometric distribution of the three sp2 hybrid orbitals is within a plane, directed at 120o angles This distribution gives a trigonal planar molecular geometry, as predicted by VSEPR

21. sp2 Hybridization scheme is useful in describing double covalent bonds, e.g. Ethylene in CH2=CH2 Unhybridized orbital

22. 10.5

23. Nonhybridized p-orbitals

24. sp2 Hybridization scheme is useful in describing double covalent bonds, e.g. Ethylene

25. Sigma  and pi bonds • Sigma bond is formed when two orbitals each with a single electron overlap (Head-to-head overlap). Electron density is concentrated in the region directly between the two bonded atoms • Pi-bond is formed when two parallel p-orbitals overlap side-to-side • The orbital consists of two lopes one above the bond axis and the other below it. • Electron density is concentrated in the lopes • Electron density is zero along the line joining the two bonded atoms

26. H H C C H H

27. sp2 hybridization • C2H4 • Double bond acts as one pair. • trigonal planar • Have to end up with three blended orbitals. • Use one s and two p orbitals to make sp2 orbitals. • Leaves one p orbital perpendicular.

28. 2p sp2 Hybridization In terms of energy 2p Energy 2s

29. The geometric distribution of the two sp hybrid orbitals is linear , directed at 180o angles This distribution gives a linear molecular geometry sp Hybridization H-Be-H 1s2 2s2 4Be

30. 2p sp Hybridization In terms of energy 2p Energy 2s

31. This hybridization scheme is useful in describing triple covalent bonds HC CH in acetylene Hybridization in Unhybridized orbitals

33. Carbon–Carbon Triple Bonds

34. Hybridization in molecules containing multiple bonds • The extra electron pairs in double or triple bonds have no effect upon the geometry of molecules • Extra electron pairs in multiple bonds are not located in hybrid orbitals • Geometry of a molecule is fixed by the electron pairs in hybrid orbitals around the central atom • All unshared electron pairs • Electron pairs forming single bonds • One (only one) electron pair in a multiple bond

35. CO2 O C O • C can make two s and two p • O can make one s and one p

36. This hybridization allows for expanded valence shell compounds – typical for group 5A elements, e.g., 15P dsp3/sp3d Hybridization A 3s electron can be promoted to a 3d subshell, which gives rise to a set of five sp3d hybrid orbitals Central atoms without d-orbitals, N, O, F, do not form expanded octet

37. Cl Cl Cl P Cl Cl PCl5 Phosphorus Pentachloride Cl P Cl Cl Cl P Cl Cl VSEPR

38. 3 sp d sp3d sp3d sp3d sp3d sp3d Neon 2 3s 3px 3py 3pz dxz dyz dxy dx2-y2 dz2 Hybridized 90o Promoted 120o TrigonalBipyrimidal 120o Phosphorus Pentachloride: Phosphorus

39. A 3s and a 3p electron can be promoted to the 3d subshell, which gives rise to a set of six sp3d2 hybrid orbitals d2sp3/sp3d2 hybridization This hybridization allows for expanded valence shell compounds – typically group 6A elements, e.g., S

40. Predicting Hybridization Schemes In hybridization schemes, one hybrid orbital is produced for every simple atomic orbital involved Write a plausible Lewis structure for the molecule or ion • Use the VSEPR method to predict the electron-group geometry of the central atom • Count # of e-pairs around the atom • Multiple bond is counted as one pair • Choose the hybrid set having same number of orbitals

41. Success of the localized electron model • Overlap of atomic orbitals explained the stability of covalent bond • Hybridization was used to explain the molecular geometry predicted by the localized electron model • When lewis structure was in adequate, the concept of resonance was introduced to explain the observed properties

42. Weakness of the localized electron model • It incorrectly assumed that electrons are localized and so the concept of resonance was added • Inability to predict the magnetic properties of molecules like O2 (molecules containing unpaired electrons) • No direct information about bond energies

43. 9.2 Molecular Orbital Theory • Molecular orbitals (MOs) aremathematical equationsthat describe the regions in a molecule where there is ahigh probability of finding electrons • Atomic orbitals of atoms are combined to give a new set of molecular orbitals characteristic of the molecule as a whole • The number of atomic orbitals combined equals the number of molecular orbitals formed. • (Two s-orbitals Two molecular orbitals)

44. Molecular orbitals • Two atomic orbitals combine to form a bonding molecular orbital and an anti-bonding MO*. • Electrons in bonding MO’s stabilize a molecule • Electrons in anti-bonding MO’s destabilize a molecule • For the orbitals to combine, they must be of comparable energies. e.g., 1s(H) with 2s(Li) is not allowed • The molecular orbitals are arranged in order of increasing energy. • The electronic structure of a molecule is derived by feeding electrons to the molecular orbitals according to same rule applied for atomic orbitals

45. Molecular orbitals • Each molecular orbital can hold a maximum of two electrons with opposite spins • Electrons go into the lowest energy molecular orbital available • Hund’s rule is obeyed Molecular orbital model will be applied only to the diatomic molecules of the elements of the first two periods of the Periodic Table

46. The hydrogen molecule Destructive interference HB HA Electron density is concentrated away from internuclear region Bonding MO1 = enhanced region of electron density Electron density is concentrated between nuclei Constructive interference Formation of molecular orbitals by combination of 1s orbitals Antibonding MO2 = region of diminished electron density

47. Energy level diagram in hydrogen (H2). Bonding molecular orbitalhas lower energy and greater stability than the atomic orbitals from which it was formed. antibonding molecular orbitalhas higher energy and lower stability than the atomic orbitals from which it was formed.

48. The Molecular Orbital Model • We use labels to indicate shapes, and whether the MO’s are bonding or antibonding. • MO1 = s1s • MO2 = s1s* (* indicates antibonding) • We can write them the same way as atomic orbitals • H2 = s1s2

49. Bond order • Bond Order =1/2 (bonding e – antibonding e) (number of bonds) • Higher bond order = stronger bond

50. Bond order for H2