1. Covalent Bonding: orbitals

2. Hybridization - The Blending of Orbitals = = + +s orbitalp orbital sp orbital

3. We have studied electron configuration notation and the sharing of electrons in the formation of covalent bonds. Methane is a simple natural gas. Its molecule has a carbon atom at the center with four hydrogen atoms covalently bonded around it. What Proof Exists for Hybridization? Lets look at a molecule of methane, CH 4.

4. What is the expected orbital notation of carbon in its ground state? (Hint: How many unpaired electrons does this carbon atom have available for bonding?) Can you see a problem with this? Carbon ground state configuration

5. The first thought that chemists had was that carbon promotes one of its 2s electrons… …to the empty 2p orbital. Carbon’s Empty Orbital

6. The fourth bond is between a 2s electron from the carbon and the lone 1s hydrogen electron. Such a bond would have slightly less energy than the other bonds in a methane molecule. Unequal bond energy

7. In the case of methane, they call the hybridization sp 3, meaning that an s orbital is combined with three p orbitals to create four equal hybrid orbitals. These new orbitals have slightly MORE energy than the 2s orbital… … and slightly LESS energy than the 2p orbitals. sp 3 Hybrid Orbitals

8. Here is another way to look at the sp 3 hybridization and energy profile… sp 3 Hybrid Orbitals

9. While sp 3 is the hybridization observed in methane, there are other types of hybridization that atoms undergo. These include sp hybridization, in which one s orbital combines with a single p orbital. This produces two hybrid orbitals, while leaving two normal p orbitals sp Hybrid Orbitals

10. Another hybrid is the sp 2, which combines two orbitals from a p sublevel with one orbital from an s sublevel. One p orbital remains unchanged. sp 2 Hybrid Orbitals

11. 1s2s3s 2p3p3d SF 6 FFF F F F

12. Hybridization Involving “d” Orbitals Beginning with elements in the third row, “d” orbitals may also hybridize dsp 3 = five hybrid orbitals of equal energy d 2 sp 3 = six hybrid orbitals of equal energy

13. Hybridization and Molecular Geometry Forms Overall Structure Hybridization of “A” AX 2 Linearsp AX 3, AX 2 ETrigonal Planarsp 2 AX 4, AX 3 E, AX 2 E 2 Tetrahedralsp 3 AX 5, AX 4 E, AX 3 E 2, AX 2 E 3 Trigonal bipyramidal dsp 3 AX 6, AX 5 E, AX 4 E 2 Octahedrald 2 sp 3 A = central atom X = atoms bonded to A E = nonbonding electron pairs on A

14. σ BONDS

15. Sigma and Pi Bonds Sigma (  ) bonds exist in the region directly between two bonded atoms. Pi (  ) bonds exist in the region above and below a line drawn between two bonded atoms. Single bond1 sigma bond Double Bond1 sigma, 1 pi bond Triple Bond1 sigma, 2 pi bonds

16. Sigma and Pi Bonds Single Bonds Ethane 1  bond

17. Sigma and Pi Bonds: Double bonds Ethene 1  bond 1  bond

18. Sigma and Pi Bonds Triple Bonds Ethyne 1  bond 1  bond

19. CH 4 - Methane 4 single bonds

20. Double bonds Use unhybridized p orbitals to share electrons

21. Sigma Bond - (σ) is a bond resulting from the head-on overlap of atomic orbitals. The region of electron density is along and cylindrically around an imaginary line connecting the bonded atoms. Pi Bond - (π) a bond resulting from side on overlap of atomic orbitals. The regions of electron sharing are on opposite sides of an imaginary line connecting the bonded atoms and parallel to this line. You can't have a Pi bond without first having a sigma bond.

23. Triple Bond = 1 sigma and 2 Pi bonds

24. The De-Localized Electron Model Pi bonds (  ) contribute to the delocalized model of electrons in bonding, and help explain resonance Electron density from  bonds can be distributed symmetrically all around the ring, above and below the plane.