1. An amazing thing about the universe - It works in a way that sometimes when things come together, they stick… Sections 2.4-2.8 Sections 2.4-2.8 H H H H Ch 2 Structure & Bonding “spin-pairing”

2. Valence Bonding (Localized) vs. Molecular Orbital Theory (Delocalized)

3. Orbital Overlap (Localized Bonding) Bonding orbitals are constructed by combining atomic orbitals from adjacent atoms. In VB Theory, usually only worry about the valence electrons In VB Theory, usually only worry about the valence electrons From Quantum Mechanics: orbitals can add or subtract; therefore constructive or destructive interference is possible

4. Orbital Overlap: As two H atoms approach, the overlap of their 1s atomic orbitals increases. The wave amplitudes add, generating a new orbital with high electron density between the nuclei. Localized bonding… H 2

5. Sigma bonding: The  bond is actually symmetry notation means bonding is directed along internuclear axis means bonding is directed along internuclear axis sigma bonds have a C ∞ axis sigma bonds have a C ∞ axis

6. Pi bonding: The  bond is also found in symmetry notation means bonding is above & below the internuclear axis means bonding is above & below the internuclear axis

7. PH 3 Phosphine is a colorless, highly toxic gas with bond angles of 93.6°. Describe the bonding in PH 3.

8. Oh shoot… What about methane? What is the electron configuration of methane?

9. s and p hybridization Promotion: excitation of an electron to a higher energy orbital in the course of bond formation – not real exactly Hybridization: mathematical mixing (linear combinations) of valence atomic orbitals to achieve new equal energy degenerate orbitals

10. Methane hybridization

11. CH 4 is tetrahedral Therefore, 2s and the 2p x, 2p y, & 2p z must hybridize new orbitals are called sp 3 new orbitals are called sp 3 An inner atom with a steric number of 4 has tetrahedral electron group geometry and can be described using sp 3 hybrid orbitals.

12. General Features of Hybridization 1.The # of valence orbitals generated by hybridization equals the # of valence AOs participating in hybridization. 2.The steric number of an inner atom uniquely determines the number and type of hybrid orbitals. 3.Hybrid orbitals form localized bonds by overlap with atomic orbitals or with other hybrid orbitals. 4.There is no need to hybridize orbitals on outer atoms, because atoms do not have limiting geometries. The bonds formed by all other outer atoms can be described using valence p orbitals.

14. Isolobal: When analogous fragments on differing molecules have closely similar bonding patterns

15. Molecular Orbital Theory: When overlapping AOs just won’t cut it. So far, bonding has been described as overlapping AOs or hybrid orbitals However, *all* electrons from each bonding atom feel the presence of the others Bonding in the diatomic molecules second row elements can be explained in two ways. Localized Bonding (LB) Theory, a.k.a. Valence Bond Theory Localized Bonding (LB) Theory, a.k.a. Valence Bond Theory Molecular Orbital (MO) Theory Molecular Orbital (MO) Theory MO Theory assumes pure s and p AOs of the atoms in a molecule combine to produce orbitals that are spread out, or delocalized, over several atoms, leading to MOs assumes pure s and p AOs of the atoms in a molecule combine to produce orbitals that are spread out, or delocalized, over several atoms, leading to MOs One advantage over VB Theory: correctly explains electronic structures of molecules which do not follow Lewis Dot structure. One advantage over VB Theory: correctly explains electronic structures of molecules which do not follow Lewis Dot structure.

16. 4 Principles of MO Theory 1 st Principle: the total # of MOs produced by a set of interacting AOs is equal to the # of interacting orbitals 2 nd Principle: the bonding MO is lower in energy than the parent AOs & the antibonding MO is higher in energy (LCAO) To explain this, let’s look at the hydrogen molecule, H 2

17. Bonding and Antibonding in H 2 Each hydrogen atom contributes a 1s orbital These orbitals can be added or subtracted Addition: Bonding MO (  1s ) Addition: Bonding MO (  1s ) Subtraction: Bonding MO (  1s * ) Subtraction: Bonding MO (  1s * )

19. Principles of MO Theory 3 rd Principle: electrons of the molecule are assigned to orbitals of successively higher energy according to the Aufbau principle and Hund’s Rule Electrons occupy the lowest energy orbitals first. Electrons occupy the lowest energy orbitals first. Atoms are most stable with the highest number of unpaired electrons Atoms are most stable with the highest number of unpaired electrons

20. NOW: Use a molecular orbital diagram to predict if it is possible to form the He 2 + cation. molecular orbital diagram of He 2

21. Bond Order in MO Theory Bond order allows us to represent the net amount of bonding between two atoms. The higher the bond order, the more stable the structure The higher the bond order, the more stable the structure

22. One More Principle… 4 th Principle: atomic orbitals combine to form molecular orbitals most effectively when the atomic orbitals are of similar energy & symmetry i.e. a 1s will not bond with a 2s. i.e. a 1s will not bond with a 2s.

23. Second-Row Diatomic Molecules subtractive additive - +

24. NOTE NUMBERING!

29. Orbital Mixing In B 2, the overlap of 2s and 2p z orbitals stabilizes  s and destabilizes  p The amount of mixing depends on the energy difference between the 2s and 2p atomic orbitals. Mixing is largest when the energies of the orbitals are nearly the same 2 cases: 1) Z ave ≤ 7, 2) Z ave > 7 (qualitative)

30. Heteronuclear Diatomic Molecules Let’s examine the MO diagram of NO …and of CO…

31. Evidence for Antibonding Orbitals