1. Carbon’s valence electrons?

2. Hybrid Orbitals  Mixing of valence shell orbitals to form new similar orbitals for bonding electrons.

3. Formation of sp 3 Hybrid Orbitals

4. Formation of sp 2 Hybrid Orbitals

5. Formation of sp Hybrid Orbitals

6. sp hybrid orbital – from an s and a p orbital sp 2 hybrid orbital – from 1 s and 2 p orbitals sp 3 hybrid orbital – from 1 s and 3 p orbitals...

7. # of Lone Pairs + # of Bonded Atoms HybridizationExamples sp sp 2 sp 3 sp 3 d sp 3 d 2 2345623456 BeCl 2 BF 3 NH 3, H 2 O PCl 5 SF 6

8. Sigma (  ) Bond Electron pair shared in an area between the atoms. Pi (  ) Bond Forms double and triple bonds by sharing electron pair(s) in the space above and below the σ bond. Uses the unhybridized p orbitals.

9. The Orbitals for CO 2

10. Sigma (  ) and Pi Bonds (  ) Single bond 1 sigma bond Double bond 1 sigma bond and 1 pi bond Triple bond 1 sigma bond and 2 pi bonds How many  and  bonds are in the acetic acid molecule CH 3 COOH? C H H CH O OH  bonds = 6 + 1 = 7  bonds = 1

11. Bonding Order (non-MO method) Number of bonding electron pairs between two bonded atoms The larger the bond order, the stronger the bond Bond order = total number of bonding pairs number of bonding locations Single bonds: bond order = 1 Double bonds: = 2 Triple bonds: = 3 Resonance: an average of the resonance structures

12.  For a molecule to be polar, at least one bond must be polar.

13.  If there is no dipole, the molecule is nonpolar  If there is a dipole (usually asymmetrical), the molecule is polar

14. .. :H:.. │.. :Cl−C −Cl: ˙ ˙ │ ˙ ˙ :H: ˙˙ Polar or nonpolar?

15.  Bond polarity: electronegativity difference  Molecule polarity: 1) must have a polar bond 2) asymmetrical shape

16. Which of the following molecules have a dipole moment? H 2 O, CO 2, SO 2, and CH 4 O H H dipole moment polar molecule S O O CO O no dipole moment nonpolar molecule dipole moment polar molecule C H H HH no dipole moment nonpolar molecule

17. Polarity of Molecules Dipole Moments of Polyatomic Molecules Example: CO 2 : There is no C-O dipole because of the molecule’s shape. H 2 O: there is an overall dipole (separation of central positions of partial positive and negative charges) because the molecule is bent.

18. Polarity of Molecules Dipole Moments of Polyatomic Molecules

19. Optional: Energy levels of bonding and antibonding molecular orbitals in hydrogen (H 2 ). A bonding molecular orbital has lower energy and greater stability than the atomic orbitals from which it was formed. An antibonding molecular orbital has higher energy and lower stability than the atomic orbitals from which it was formed.

22. 1.The number of molecular orbitals (MOs) formed is always equal to the number of atomic orbitals combined. 2.The more stable the bonding MO, the less stable the corresponding antibonding MO. 3.The filling of MOs proceeds from low to high energies. 4.Each MO can accommodate up to two electrons. 5.Use Hund’s rule when adding electrons to MOs of the same energy. 6.The number of electrons in the MOs is equal to the sum of all the electrons on the bonding atoms. Optional: Molecular Orbital (MO) Configurations

23. bond order = 1 2 Number of electrons in bonding MOs Number of electrons in antibonding MOs ( - ) bond order ½10½ Optional

24. Molecular orbital theory – bonds are formed from interaction of atomic orbitals to form molecular orbitals. O O No unpaired e - Should be diamagnetic Experiments show O 2 is paramagnetic

25. Delocalized molecular orbitals are not confined between two adjacent bonding atoms, but actually extend over three or more atoms.