Covalent Bonding: Orbitals

1. Covalent Bonding: Orbitals

2. 9.1 Hybridization and the localized electron model • Remember that the localized electron model says that molecules are a collection of atoms bound by sharing electrons and represented by Lewis structures. • We are going to look at which atomic orbitals are used to share electrons and form bonds.

3. CH4 - Methane • Let’s look at methane • Draw its Lewis structure • Write the electron configuration for the electrons involved in bonding for… • H • C

4. There are two options here • Using carbon’s 2s and 2p orbitals to bond with the 1s orbital from Hydrogen, would mean that there are two different kinds of C—H bonds. • Because the C 2p orbitals are  we would expect that the orbitals would be at 90o angles • But we know that the C—H bonds are ALL equal and we know that a tetrahedral molecule has bond angles of 109.5, so now what?

5. Hybridization • Hybridization is a mixing of two or atomic more orbitals to form a new set of “hybrid” orbitals. (often explains the geometry) • Mix at least 2 nonequivalent atomic orbitals (s and p) and get hybrid orbitals which have a different shape from the original atomic orbital

6. Hybridization cont. • Number of hybrid orbitals is equal to the number of pure atomic orbitals used in the hybridization process • Covalent bonds are formed by: • Overlap of hybrid orbitals with atomic orbitals • Overlap of hybrid orbitals with other hybrid orbitals

7. sp3 So how does this work for methane? Video first. • Let’s look at it in terms of energy first, then orbitals. 2p Hybridization 2s

8. Now the orbitals, formation of sp3 hybrid orbitals

9. Why?? • Why do we do this? • “spend” energy to promote e to • To get more unpaired electrons (4 in sp3) • To make more bonds • Because making bonds is exothermic • Lowers the energy of the molecule (average of the orbitals involved)

10. Formation of covalent bonds in CH4 When the 1s orbital from Hydrogen overlaps with the sp3 orbital from C you get what is called a  (sigma) bond. Sigma bonds (single bond): • form from orbitals whose lobes point toward each other • Increased e- density along the internuclear axis

11. Look at BF3 • Draw its Lewis structure • Write the electron configuration for the electrons involved in bonding for… • B • F

12. 2p sp2 Hybridization Energy for sp2 hybridization 2p 2s

14. sp hybridization with pi bonds O C O

15. What kind of bonds does CO2 have?2 sigma bonds AND 2 pi bonds

16. Pi bonds • Electron pairs shared in the space above and below the sigma bond. • 2p orbital that is perpendicular to the hybrid orbitals on the atoms.

17. Remember that steric number and hybridization go hand in hand