Lecture Notes by Ken Marr Chapter 11 (Silberberg 3ed) Covalent Bonding: Valence Bond Theory and

1. Lecture Notes by Ken Marr Chapter 11(Silberberg 3ed) Covalent Bonding: Valence Bond Theory and Molecular Orbital Theory 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

2. Valence Bond Theory • Covalent Bonds • Result from the overlap of valence shell atomic orbitals to share an electron pair • s, p, or hybrid orbitals may be used to form covalent bonds e.g. Predict theOrbitals used for bonding in: H2, HF, H2S, F2

4. Examples of s and p Orbitals involved in Bonding • Overlap of s orbitals • H2 • Overlap of s and p orbitals • HF • H2S • Overlap of p orbitals • F2

5. Hybrid Orbitals • The use of only s and p orbitals does not explain bonding in most molecules!!! e.g. BeCl2, CH4 , H2O hybrid orbitals are used in these cases • Hybrid Orbitals are used to hold bonding and nonbonding electrons! • s, p, and d orbitals may hybridize to form to form hybrid orbitals

6. How to Determine an Atom’s Hybridization Write Lewis structure for the molecule or ion, then... • Determine number of electron pairs around the atom in question • One orbital is needed for each electron pair sphybridization provides 2 orbitals sp2hybridization provides 3 orbitals sp3 hybridization provides.....?...........orbitals sp3dhybridization provides......?..........orbitals sp3d2hybridization provides....?.............orbitals

8. Examples of Hybrid Orbitals • Example of sp hybrid orbitals: • BeCl2

13. Examples of Hybrid Orbitals • Example of sp2 hybrid orbitals BF3

15. Hybrid Orbitals • Examples of sp3 hybrid orbitals CH4, C2H6, H2O, NH3 • Example of sp3d hybrid orbitals PCl5 • Example of sp3d2 hybrid orbitals SF6

21. Bond Angle ~92 o

22. Mode of Orbital OverlapSigma vs. Pi Bonds • Sigma Bonds (s-bond) • Head to head overlap of s, p, or hybrid orbitals • Responsible for the framework of a molecule Single bond = one s bond

23. Mode of Orbital OverlapSigma vs. Pi Bonds • Pi bonds (p-Bonds) • Side to side overlap of p orbitals • Restrict rotation • Double bond = one s bond+ one p bond • Triple Bond = one s bond+ two p bonds

24. Examples: Sigma vs. Pi Bonds • Ethane • Ethylene (ethene) • Effect of p-bonding on rotation about the s-bond? • Acetylene (ethyne) • Nitrogen • Formaldehyde

29. Predict the hybrid orbitals used in the following • Nitrogen gas, N2 • Formaldehyde: H2CO • Carbon dioxide, CO2 • Carbon monoxide, CO • Sulfur dioxide, SO2

30. One option for SO2 o o S = [Ne] 3s2 3px2 3py1 3pz1 This structure is... • Favored by formal charge • Requires ?? hybridization Big Problems with this Structure.. • How many unhybridized p-orbitals are available for p bonding? • How many p-orbitals are needed?

31. Another Option for SO2 S = [Ne] 3s2 3px2 3py1 3pz1 • ??? Hybridization • Bond order? • Resonance? o o

32. Resonance: Delocalization of electrons • Shifting of p-bond electrons without breaking the s- bond • Although not favored by formal charge, B.O. = 1.5 o o o o Resonance

33. Molecular Orbital Theory • p-electron pair found in molecular orbital formed from the overlap of p-orbitals • B.O. = 1.5 • same as measured B.O. • S – O bond length is intermediate between S – O and S = O bond lengths o o

34. Strengths and Weaknesses of Valence Bond Theory • VB Theory Molecules are groups of atoms connected by localized overlap of valence shell orbitals • VB, VSEPR and hybrid orbital theories work well together to explain the shapes of molecules • But……VB theory inadequately explains… • Magnetic property of molecules • Spectral properties of molecules • Electron delocalization • Conductivity of metals

35. Molecular Orbital Theory • The electrons in a molecule are found in Molecular Orbitals of different energies and shapes • Just as an atom’s electrons are located in atomic orbitals of different energies and shapes • MOs spread over the entire molecule • Major drawbacks of MO Theory • Based on Quantum theory • Calculations are based on solving very complex wave equations major approximations are needed! • Difficult to visualize

36. Advantages of MO Theory • VB Theory incorrectly predicts that.... • O2 is diamagnetic with B.O. = 2 or.... • O2 is paramagnetic with B.O. = 1 • MO Theory correctly predicts that.... • O2 is paramagnetic with B.O. = 2 • VB Theory requires resonance structures to explain bonding in certain molecules and ions • MO Theory does not have this limitation

37. Formation of Molecular Orbitals • MO’s form when atomic orbitals overlap • Bonding MOs • Result from constructive interference of overlapping electron waves • Stabilize a molecule by concentrating electron density between nuclei • MO’s  more stable than AO’s  delocalize electron charge over a larger volume

38. Overlap of standing electron Waves Constructive interference Destructive interference

39. (High e- density) (Low e- density) Fig. 11.13

41. H2 is more stable than the separate atoms

42. Antibonding MOs • Antibonding MOs • Result from destructive interferenceof overlapping electron waves • Reduce electron density between nuclei • Destabilize a molecule • Higher in energy than bonding MOs of the same type

43. Using MO Theory to Calculate Bond Order • VB definition of Bond Order.... • Number of electron pairs shared between 2 nuclei • MO Theory B.O. = ½ (No. Bonding e-- No. Antibonding e-) • Meaning of B.O. • B.O. > 0, then molecule more stable than separate atoms • B.O. = 0, then zero probability of bond formation • The greater the B.O., the stronger the bond

44. Why Do Some Molecules Exist and Others Do Not? • Why do H2 and He21+ exist , but He2 does not? Recall…..Bonding results only if there is a net decrease in PE • Molecules with equal numbers of Bonding and antibonding electrons are unstable...Why?...... • Antibonding MOs raise PE more than Bonding MOs lower PE

45. Use MO theory to predict if the following can form • Hydride ions: H2 1- and H2 2- • Li2 , Li2 1+ , Li2 2+ , Li2 1- • Be2 , Be2 1+ , Be2 2+ , Be2 1-

46. In He2, the antibonding electrons in s1s* cancel the PE lowering of s1s

49. Sigma vs Pi Molecular Orbitals • s Molecular Orbitals form when..... • s - atomic orbitals overlap • p - atomic orbitals overlap head to head • pMolecular Orbitals form when..... • p - atomic orbitals overlap side to side • Why are s-bonds more stable than p-bonds?